www.studyguide.pk 1.1 Chemistry & Measurements Scientists throughout the world use the International System of Units, abbreviated SI, for their measurements. There are seven base units in the SI system. Measurements can involve very large or very small units that do not correspond with the base SI units. SI units are conveniently modified by a series of prefixes that represent multiples of the base units. Thus 1/1000th of a meter (or 0.001m) becomes 1 millimeter or simply 1 mm. Because the numbers chemists use are often very small or very large, it is convenient to express these numbers in scientific notation. Notice also that all measurements contain both a number and the unit of measure. 1.2 Measuring Mass Mass is defined as the amount of matter in an object. The standard SI unit of mass is the kilogram (1 kg weighs 2.205 lb). Smaller mass units are frequently used in chemistry. 1 gram = 0.001 kg = 1.0 x 10–3 kg 1 milligram (1 mg) = 1.0 x 10–3 g 1 microgram (1µg) = 1.0 x 10–6 g Note that the terms mass and weight are used interchangeably, but they do have different www.studyguide.pk meanings. Mass is the amount of matter in an object, while weight is a measure of the pull of gravity on an object. 1.3 Measuring Length The standard unit of length in the SI system is the meter, abbreviated m (1 m = 39.37 in). A meter, like a kilogram, is too large a unit of measure for most chemistry work. Chemists frequently use the following smaller units of measure: 1 centimeter (1 cm) = 0.01 m = 1.0 x 10–2 m 1 millimeter (1mm) = 1.0 x 10–3 m 1 micrometer (1µm) = 1.0 x 10–6 m 1 nanometer (1nm) = 1.0 x 10–9 m 1 picometer (1 pm) = 1.0 x 10–12 m 1.4 Measuring Temperature Of the three common temperature scales, the Kelvin scale is generally used for scientific work. Much of the world uses the Celsius scale, except the United States which uses the Fahrenheit scale. The kelvin (abbreviated K), as the unit of measure is called, is the same physical increment as a Celsius degree (abbreviated °C). The difference between the two units of measure is that the corresponding temperature scales are offset by a fixed amount. The Fahrenheit degree (abbreviated °F) is smaller than the Kelvin and Celsius degrees, and the Fahrenheit scale is offset by a different amount. 2 www.studyguide.pk An understanding of the three temperature scales will help you make temperature conversions without blindly applying a formula. Look at the number of degrees that separate the freezing point of water and the boiling point of water on the Fahrenheit scale: There are 180 degrees. This same temperature interval is separated by 100°C and 100 K. Thus, one °F is 100/180 or 5/9 the size of a kelvin or °C. That is, there are fewer Celsius degrees than Fahrenheit degrees in the same range because Celsius degrees are "fatter". The Fahrenheit scale is offset from the Celsius scale by 32°F. These facts lead to the following conversion equations: The Kelvin scale is offset from the Celsius scale by 273.15. Thus, Let's convert a temperature commonly used in baking—350°F—to °C and to kelvins. Our conversion path will be to convert from °F to °C, then perform a second conversion from °C to K. To convert from °F to °C, we first subtract the offset of 32. 350°F – 32 = 318°F Next we convert the size of the degrees. 318°F (5°C/9°C) = 177°C To convert °C to kelvins, we simply add 273.15, the Kelvin scale offset. 177 + 273.15 = 450 K 1.5 Derived Units: Measuring Volume The seven fundamental SI units are not sufficient to describe units of measurement for such things as area, volume, density, etc. These units are called derived units because they can be expressed using one or more of the seven base units. Volume, the amount of space occupied by an object, is measured in SI units by the cubic meter, abbreviated m3. Smaller, more convenient measures of volume are frequently used. 3 www.studyguide.pk 1 dm3 = 1 liter (1 L) 1 cm3 = 1 milliliter (1 mL) Measuring liquid volume is a common laboratory task. Some of the specialized glassware used in chemistry labs is shown below. Which one is the buret? 4 www.studyguide.pk 1.6 Derived Units: Measuring Density Density is an intensive physical property that relates the mass of an object to its volume. Notice the wide range of densities of common substances listed in the table below. The volume of many substances changes with temperature, so densities, too, are temperature dependent. Density provides a useful link in the laboratory between the mass of a substance and its volume. It is sometimes simpler to use volumetric glassware to measure a particular volume of a substance, and then to convert that volume to mass. 5 www.studyguide.pk ELEMENTS and COMPOUNDS MIXTURES and their separation CHEMICAL REACTIONS and EQUATIONS KEYWORDS ... atom ... chemical change ... chromatography ... compound ... covalency ... distillation (simple/fractional) ... element ... equations - (word, picture, symbol, quizzes) ... formula ... impure/pure ... ionic equations ... ionic valency ... magnet ... mixture ... molecule ... physical change ... products ... reactants ... separating mixtures ... chemical symbols - (elements, formula, in equations) ... state symbols ... valency ... working out formulae ... Introduction and Some keywords (pictures) ATOM An ATOM is the smallest particle of a substance which can have its characteristic properties. BUT remember atoms are built up of even more fundamental sub-atomic particles - the electron, proton and neutron. A MOLECULE is a larger particle formed by the chemical combination of two or more atoms. The molecule may be an element or a compound eg hydrogen H2 or carbon dioxide CO2 and the atoms are held together by covalent bonds. ELEMENT and symbols H I Na Al Fe C Ag U? COMPOUND and FORMULA An ELEMENT is a pure substance made up of only one type of atom*, 92 in the Periodic Table naturally occur from hydrogen H to uranium U. Note that each element has symbol which is a single capital letter like H or U or a capital letter + small letter eg cobalt Co, calcium Ca or sodium Na. Each element has its own unique set of properties but the Periodic Table is a means of grouping similar elements together. They may exist as atoms like the Noble Gases eg helium He or as molecules eg hydrogen H2 or sulphur S8. (more examples applied to equations and see note about 'formula of elements') * At a higher level of thinking, all the atoms of the same element, have the same atomic or proton number. This number determines how many electrons the atom has, and so ultimately its chemistry. Any atom with 27 protons and electrons is cobalt! A COMPOUND is a pure substance formed by chemically combining at least two different elements by ionic or covalent bonding www.studyguide.pk CH4 MIXTURE PURE Compounds can be represented by a FORMULA, eg sodium chloride NaCl (ionic, 2 elements,1of sodium and 1 of chlorine), methane CH4 (covalent, shown on the left has 2 elements in it, 4 of carbon and 1 of hydrogen*) and glucose C6H12O6 (covalent, 3 elements, 6 atoms of carbon, 12 of hydrogen and 6 of oxygen). There must be at least two different types of atom (elements) in a compound.(* the 1 is never written in the formula, no number means 1) Compounds have a fixed composition and therefore a fixed ratio of atoms represented by a fixed formula, however the compound is made or formed. In a compound the elements are not easily separated by physical means, and quite often not easily by chemical means either. The compound has properties quite different from the elements it is formed from. o For example soft silvery reactive sodium + reactive green gas chlorine ==> colourless, not very reactive crystals of sodium chloride. The formula of a compound summarises the 'whole number' atomic ratio of what it is made up of eg methane CH4 is composed of 1 carbon atom combined with 4 hydrogen atoms. Glucose has 6 carbon : 12 hydrogen : 6 oxygen atoms, sodium chloride is 1 sodium : 1 chlorine atom. When there is only one atom of the element, there is no subscript number, the 1 is assumed eg Na in NaCl or C in CH4. When there is more than 1 atom of the same element, a subscript number is used eg the 4 in CH4 meaning 4 hydrogen atoms. Sometimes, a compound (usually ionic), is partly made up of two or more identical groups of atoms. To show this more accurately ( ) are used eg o calcium hydroxide is Ca(OH)2 which makes more sense than CaO2H2 because the OH group is called hydroxide and exists in its own right in the compound. o Similarly, aluminium sulphate has the formula Al2(SO4)3 rather than Al2S3O12, because it consists of two aluminium ions Al3+ and three sulphate ions SO42-. The word formula can also apply to elements. eg hydrogen H2, oxygen O2, ozone O3 (2nd unstable form of oxygen), phosphorus P4, sulphur S8, have 2, 2, 3, 4 and 8 atoms in their molecules. Elements like helium He are referred to as 'monatomic' because they exist as single uncombined atoms. A MIXTURE is a material made up of at least two substances which may be elements or compounds. They are usually easily separated by physical means eg filtration, distillation, chromatography etc. Examples: air, soil, solutions. PURE means that only one substance present in the material and can be an element or compound. A simple physical test for purity and helping identify a compound is to measure the boiling point of a liquid. Every pure substance melts and boils at a fixed temperature. o If a liquid is pure it may boil at a constant temperature (boiling point). o An impure liquid could boil higher or lower than the expected boiling point and over a range of temperature. 2 www.studyguide.pk o o IMPURE PURIFICATION If a solid is pure, it will quite sharply at the melting point. An impure solid melts below its expected melting point and more slowly over a wider temperature range. IMPURE usually means a mixture of mainly one substance plus one or more other substances physically mixed in. The % purity of a compound is important, particularly in drug manufacture. Any impurities present are less cost-effective to the consumer and they may be harmful substances. Materials are purified by various separation techniques. The idea is to separate the desired material from unwanted material. they include: o Filtration to separate a solid from a liquid. You may want the solid or the liquid or both! o Simple distillation to separate a pure liquid from dissolved solid impurities which have a very high boiling point. o Fractional distillation to separate liquids with a range of different boiling points, especially if relatively close together. o Crystallisation to get a pure solid out of a solvent solution of it. o Chromatography can be used on a larger scale than spots' to separate out pure samples from a mixture. Picture examples of Elements, Compounds and Mixtures 3 www.studyguide.pk 4 www.studyguide.pk METHODS of SEPARATING MIXTURES Simple Distillation Distillation involves 2 stages and both are physical state changes. (1) The liquid or solution mixture is boiled to vaporise the most volatile component in the mixture (liquid ==> gas). The ant-bumping granules give a smoother boiling action. (2) The vapour is cooled by cold water in the condenser to condense (gas ==> liquid) it back to a liquid (the distillate) which is collected. This can be used to purify water because the dissolved solids have a much higher boiling point and will not evaporate with the steam. BUT it is too simple a method to separate a mixture of liquids especially if the boiling points are relatively close. Fractional Distillation Fractional distillation involves 2 main stages and both are physical state changes. It can only work with liquids with different boiling points. (1) The liquid or solution mixture is boiled to vaporise the most volatile component in the mixture (liquid ==> gas). The antbumping granules give a smoother boiling action. (2) The vapour passes up through a fractionating column, where the separation takes place (theory at the end). This column is not used in the simple distillation described above. (3) The vapour is cooled by cold water in the condensor to condense (gas ==> liquid) it back to a liquid (the distillate) which is collected. This can be used to separate alcohol from a fermented sugar solution. It is used on a large scale to separate the components of crude oil, because the different hydrocarbons have different boiling and condensation points FRACTIONAL DISTILLATION THEORY: o Imagine green liquid is a mixture of a blue liquid (but. 80oC) and a yellow liquid (bpt. 100oC), As the vapour from the boiling mixture enters the fractionating column it begins to cool and condense. The highest boiling or least volatile liquid tends to condense more ie the yellow liquid (water). The lower boiling more volatile blue liquid gets further up the column. Gradually up the column the blue and yellow separate from each other so that yellow condenses back into the flask and pure blue distills over to be collected. The 1st liquid, the lowest boiling point, is called the 1st fraction and each liquid distills over when the top of the column reaches its particular boiling point to give the 2nd, 3rd fraction etc. o To increase the separation efficiency of the tall fractionating column, it is usually packed with glass beads, short glass tubes or glass rings etc. which greatly increase the surface area for evaporation and condensation. o In the distillation of crude oil the different fractions are condensed out at different 5 www.studyguide.pk points in a huge fractionating column. At the top are the very low boiling fuel gases like butane and at the bottom are the high boiling big molecules of waxes and tar. Paper Chromatography This method of separation is used to see what coloured materials make up eg a food dye analysis. The material to be separated eg a food dye (6) is dissolved in a solvent and carefully spotted onto chromatography paper alongside known colours on a 'start line' (1-5). The paper is carefully dipped into a solvent, which is absorbed into the paper and rises up it. Due to different solubilities and different molecular 'adhesion' some colours move more than others up the paper, so effecting the separation of the different coloured molecules. Any colour which horizontally matches another is likely to be the same molecule ie red (1 and 6), brown (3 and 6) and blue (4 and 6) match, showing these three are all in the food dye (6). It is possible to analyse colourless mixture if the components can be made coloured eg protein can be broken down into amino acids and coloured purple by a chemical reagent called ninhydrin and many colourless organic molecules fluoresce when ultra-violet light is shone on them. FILTRATION EVAPORATION CRYSTALLISATION Filtration use a filter paper or fine porous ceramic to separate a solid from a liquid. It works because the tiny dissolved particles are too small to be filtered BUT any non-dissolved solid particles are too big to go through! Evaporation means a liquid changing to a gas or vapour. In separation, its removing the liquid from a solution, usually to leave a solid. It can be done quickly with gentle heating or left out to 'dry up' slowly. The solid will almost certainly be less volatile than the solvent and will remain as a crystalline residue. Crystallisation can mean a liquid substance changing to its solid form. However, the term usually means what happens when the liquid from a solution has evaporated to a point beyond the solubility limit. Then solid crystals will 'grow' out of the solution. All three of these separation methods are involved in (1) separation of sand and salt mixtures or (2) salt preparations eg from dissolving an insoluble base in an acid. Miscellaneous Separation Methods MAGNET This can be used to separate iron from a mixture with sulphur (see below). It is used in recycling to recover iron and steel from domestic waster ie the 'rubbish' is on a conveyer belt that passes a powerful magnet which pluck's out magnetic materials. 6 www.studyguide.pk PHYSICAL CHANGES These are changes which do not lead to new substances being formed. Only the physical state of the material changes. The substance retains exactly the same chemical composition. Examples ... melting, solid to liquid, easily reversed by cooling eg ice and liquid water are still the same H2O molecules. dissolving, eg solid mixes completely with a liquid to form a solution, easily reversed by evaporating the liquid eg dissolving salt in water, on evaporation the original salt is regained. So freezing, evaporating, boiling, condensing are all physical changes. separating a physical mixture eg chromatography, eg a coloured dye solution is easily separated on paper using a solvent, they can all be re-dissolved and mixed to form the original dye. So distillation, filtering are also physical changes. CHEMICAL CHANGES - REACTIONS - reactants and products Heating iron and sulphur is classic chemistry experiment. A mixture of silvery grey iron filings and yellow sulphur powder is made. The iron can be plucked out with a magnet ie an easily achieved physical separation because the iron and sulphur are not chemically combined yet! They are still the same iron and sulphur. On heating the mixture, it eventually glows red on its own and a dark grey solid called iron sulphide is formed. Both observations indicate a chemical change is happening ie a new substance is being formed. We no longer have iron or sulphur BUT a new compound with different physical properties (eg colour) and chemical properties (unlike iron which forms hydrogen with acids, iron sulphide forms toxic nasty smelling hydrogen sulphide!). iron + sulphur ==> iron sulphide or in symbols: Fe + S ==> FeS AND it is no longer possible to separate the iron from the sulphur using a magnet! So signs that a chemical reaction has happened include: o colour changes, o temperature changes, o change in mass eg some solids when burned in air gain mass in forming the oxide eg magnesium forms magnesium oxide some solids lose mass when heated, eg carbonates lose carbon dioxide in thermal decomposition Therefore a chemical change is one in which a new substance is formed, by a process which is not easily reversed and usually accompanied by an energy (temperature) change. This is summarised as reactants ==> products as expressed in chemical equations in words or symbols. 7 www.studyguide.pk THE CONSTRUCTION OF CHEMICAL EQUATIONS "How to write and understand chemical equations" Seven equations are presented, but approached in the following way o (1a-7a) the individual symbols and formulae are explained o (1b-7b) the word equation is presented to summarise the change of reactants to products o (1c-7c) a balanced 'picture' equation which helps you understand reading formulae and atom counting to balance the equation o (1d-7d) the fully written out symbol equation with state symbols (often optional for starter students) Chemical Symbols and Formula For any reaction, what you start with are called the reactants, and what you form are called the products. o So any chemical equation shows in some way the overall chemical change of ... o REACTANTS ==> PRODUCTS, which can be written in words or symbols/formulae. It is most important you read about formula in an earlier section of this page. empirical formula and molecular formula are dealt with on another page. In the equations outlined below several things have been deliberately simplified. This is to allow the 'starter' chemistry student to concentrate on understanding formulae and balancing chemical equations. Some teachers may disagree with this approach BUT my simplifications are: o the word 'molecule' is sometimes loosely used to mean a 'formula', o the real 3D shape of the 'molecule' and the 'relative size' of the different element atoms is ignored o if the compound is ionic, the ion structure and charge is ignored, its just treated as a formula Molecular and Structural Formulas A molecular formula gives the types and the count of atoms for each element in a compound. An example of a molecular formula is ethane, C2H6. Here the formula indicates carbon and hydrogen are combined in ethane. The subscripts tell us that there are 2 carbon atoms and 6 hydrogen atoms in a formula unit. The structural formula shows the atoms in a formula unit and the bonds between atoms as lines. Single bonds are one line, Double bonds are two lines. Triple bonds are three lines. The Lewis dot structure shows 8 www.studyguide.pk the number of valence electrons and types of bonds in the molecule. Lewis dot structure Ball and stick model Electron pairs that are shared are physically between the symbols for the atoms. Electron pairs that are unshared are called lone pairs. Lone pairs are not between atom symbols. 1a A single symbol means an uncombined single atom of the element, or Fe 1 atom of iron, atom of sulphur (2Fe would mean two atoms, 5S would mean five atoms etc.) or the formula FeS means one atom of iron is chemically combined with 1 atom of sulphur to form the compound called iron sulphide or the formula NaOH means 1 atom of sodium is combined with 1 atom of oxygen and 1 atom of hydrogen to form the compound called sodium hydroxide or the formula HCl means 1 atom of hydrogen is combined with 1 atom of chlorine to form 1 molecule of the compound called hydrochloric acid or the formula NaCl means 1 atom of sodium are combined with 1 atom chlorine to form the compound called sodium chloride or the formula H2O means 2 atoms of hydrogen are chemically combined with 1 atom of oxygen to form the compound called water. or S 1 2a 3a or the symbol Mg means 1 atom of the element called magnesium (see example 2) or 2HCl means two separate molecules of the compound called hydrochloric acid or the formula MgCl2 means 1 formula of the compound called magnesium chloride, made of one atom of magnesium and two atoms of chlorine. atoms or the formula H2 means 1 molecule of the element called hydrogen made up of two joined hydrogen 9 www.studyguide.pk 4a or the formula CuCO3 means one formula of the compound called copper carbonate, made up of one atom of copper is combined with one atom of carbon and three atoms of oxygen to form the compound copper carbonate or the formula H2SO4 means one formula of the compound called sulphuric acid, which is made up of two atoms of hydrogen, one atom of sulphur and four atoms of oxygen or the formula CuSO4 means one formula of the compound called copper sulphate which is made up of one atom of copper, one atom of sulphur and four atoms of oxygen H2O (example 2) or the formula CO2 means one molecule of the compound called carbon dioxide which is a chemical combination of one atom of carbon and two atoms of oxygen. or the formula CH4 means one molecule of the compound called methane which is made of one atom of carbon combined with four atoms of hydrogen or 2O2 means two separate molecules of the element called oxygen, and each oxygen molecule consists of two atoms of oxygen CO2 (see also example 4) 5a also example 2) or 2H2O means two separate molecules of the compound called water (see 6a or the formula Mg(OH)2 is the compound magnesium hydroxide made up of one magnesium, two oxygen and two hydrogen atoms BUT the OH is a particular combination called hydroxide within a compound, so it is best to think of this compound as a combination of an Mg and two OH's, hence the use of the ( ). or 2HNO3 means two separate molecules of the compound nitric acid, each molecule is made up of one hydrogen atom, one nitrogen atom and three oxygen atoms. or the formula Mg(NO3)2 is the compound magnesium nitrate, it consists of a magnesium (ion) and two 'nitrates' (ions), each nitrate consists of one nitrogen and three oxygen atoms, again the 10 www.studyguide.pk nitrate is a particular combination of atoms within a compound and hence the use of ( ) again. 7a 2 and 5) or 2H2O meaning two molecules of the compound water (see also examples or the formula Al2O3 means one formula of the compound called aluminium oxide, made up of two atoms of aluminium Al and three atoms of oxygen O molecules of the compound called sulphuric acid (see also example 4) or 3H2SO4 meaning three or the formula Al2(SO4)3 means one formula of the compound called aluminium sulphate, it consists of two aluminium, three sulphur and twelve oxygen atoms BUT the SO 4 is a particular grouping called sulphate, so it is best to think of the compound as a combination of two Al's and three SO4's called water (see also examples 2 and 5) or 3H2O means three separate molecules of the compound Chemical word equations ==> means the direction of change from reactants =to=> products no symbols or numbers are used in word equations always try to fit all the words neatly lined up from left to right, especially if its a long word equation eg for clarity in example 4, some names are split in two parts using two lines, one under the other, this 'style' helps understanding when it comes to revision! 1b iron + sulphur ==> iron sulphide 2b sodium hydroxide + hydrochloric acid ==> sodium chloride + water 3b magnesium + hydrochloric acid ==> magnesium chloride + hydrogen 4b o copper + sulphuric ==> copper + water + carbon carbonate acid sulphate dioxide 5b methane + oxygen ==> carbon dioxide + water 6b magnesium hydroxide + nitric acid ==> magnesium nitrate + water 7b aluminium oxide + sulphuric acid ==> aluminium sulphate + water 11 www.studyguide.pk Chemical picture equations There are three main points to writing and balancing equations Writing the correct symbol or formula for each equation component. Using numbers if necessary to balance the equation. if all is correct, then the sum of atoms for each element should be the same on both side of the equation arrow ..... o in other words: atoms of products = atoms of reactants This is a chemical conservation law of atoms and later it may be described as the 'law of conservation of mass. o the 7 equations are first presented in 'picture' style and then written out fully with state symbols o The individual formulas involved and the word equations have already been presented above. PRACTICE QUESTIONS - words and symbols o Multiple choice quiz on balancing numbers o Word-fill exercises o Reactions of acids with metals, oxides, hydroxides and carbonates. 1c on average one atom of iron chemically combines with one atom of iron forming one molecule of iron sulphide atom balancing, sum left = sum right: 1 Fe + 1 S = (1 Fe + 1S) two elements chemically combining to form a new compound the reactants are one molecule of sodium hydroxide and one molecule of hydrochloric acid the products are one molecule of sodium chloride and one molecule of water all chemicals involved are compounds atom balancing, sum left = right: ( 1 Na + 1 O + 1 H) + (1 H +1 Cl) = (1 Na + 1 Cl) + (2 H's + 1 O) one atom of magnesium reacts with two molecules of hydrochloric acid the products are one molecule of magnesium chloride and one molecule of hydrogen Mg and H-H are elements, H-Cl and Cl-Mg-Cl are compounds atom balancing, sum left = right: (1 Mg) + (1 H + 1 Cl) + (1 H + 1 Cl) = (1 Mg + 2 Cl's) + (2H's) the reactants are one formula of copper carbonate and one molecule of sulphuric acid the products are one formula of copper sulphate, one molecule of water and one molecule of carbon dioxide all molecules are compounds in this reaction 2c 3c 4c 12 www.studyguide.pk atom balancing, sum left = sum right: (1 Cu + 1 C + 3 O's) + (2 H's + 1 S + 4 O's) = (1 Cu + 1 S + 4 O's) + (2 H's + 1 O) + (1 C + 2 O's) one molecule of methane is completely burned by two molecules of oxygen to form one molecule of carbon dioxide and two molecules of water atom balancing, sum left = sum right: (1 C + 4 H's) + (2 O's) + (2 O's) = (1 C + 2 O's) + (2 H's + 1 O) + (2 H's + 1 O) one formula of magnesium hydroxide reacts with two molecules of nitric acid to form one formula of magnesium nitrate and two molecules of water (all compounds) atom balancing, sum left = sum right: (1 Mg + 2O's + 2 H's) + (1 H + 1 N + 3O's) + (1 H + 1 N + 3O's) = (1 Mg + 2 N's + 6 O's) + (2 H's + 1 O) + (2 H's + 1 O) 5c 6c 7c one formula of aluminium oxide reacts with three molecules of sulphuric acid to form one formula of aluminium sulphate and three molecules of water note the first use of numbers (3) for the sulphuric acid and water! so picture three of them in your head, otherwise the picture gets a bit big! atom balancing, sum left = sum right: (2 Al's + 3 O's) + 3 x (2 H's + 1 S + 4 O's) = (2 Al's + 3 S's + 12 O's) + 3 x (2 H's + 1 O) Chemical symbol equations 1d 2d 3d (rules already stated above) Fe(s) + S(s) ==> FeS(s) atom balancing, sum left = sum right: 1 Fe + 1 S = (1 Fe + 1S) all the reactants (what you start with) and all the products (what is formed) are all solids in this case. When first learning symbol equations you probably won't use state symbols at first (see end note). NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l) atom balancing, sum left = right: (1 Na + 1 O + 1 H) + (1 H +1 Cl) = (1 Na + 1 Cl) + (2 H's + 1 O) Mg(s) + 2HCl(aq) ==> MgCl2(aq) + H2(g) atom balancing, sum left = right: (1 Mg) + 2 x (1 H + 1 Cl) = (1 Mg + 2 Cl's) + (2H's) 13 www.studyguide.pk 4d CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2(g) 5d CH4(g) + 2O2(g) ==> CO2(g) + 2H2O(l) 6d Mg(OH)2(aq) + 2HNO3(aq) ==> Mg(NO3)2(aq) + 2H2O(l) 7d Al2O3(s) + 3H2SO4(aq) ==> Al2(SO4)3(aq) + 3H2O(l) balancing sum left = sum right: (1 Cu + 1 C + 3 O's) + (2 H's + 1 S + 4 O's) = (1 Cu + 1 S + 4 O's) + (2 H's + 1 O) + (1 C + 2 O's) atom balancing, sum left = sum right: (1 C + 4 H's) + 2 x (2 O's) = (1 C + 2 O's) + 2 x (2 H's + 1 O) atom balancing, sum left = sum right: (1 Mg + 2O's + 2 H's) + 2 x (1 H + 1 N + 3 O's) = (1 Mg + 2 N's + 6 O's) + 2 x (2 H's + 1 O) atom balancing, sum left = sum right: (2 Al's + 3 O's) + 3 x (2 H's + 1 S + 4 O's) = (2 Al's + 3 S's + 12 O's) + 3 x (2 H's + 1 O) NOTE 1: means a reversible reaction, it can be made to go the 'other way' if the conditions are changed. Example: o nitrogen + hydrogen ammonia o N2(g) + 3H2(g) 2NH3(g) o balancing: 2 nitrogen's and 6 hydrogen's on both sides of equation Note 2 on the state symbols X(?) of reactants or products in equations (g) means gas, (l) means liquid, (s) means solid and (aq) means aqueous solution or dissolved in water eg carbon dioxide gas CO2(g), liquid water H2O(l), solid sodium chloride 'salt' NaCl(s) and copper sulphate solution CuSO4(aq) VALENCY - COMBINING POWER - FORMULA DEDUCTION (2nd draft) The valency of an atom or group of atoms is its numerical combining power with other atoms or groups of atoms. The theory behind this, is all about stable electron structures! o The combining power or valency is related to the number of outer electrons. o You need to consult the page on "Bonding" to get the electronic background. A group of atoms, which is part of a formula, with a definite composition, is sometimes referred to as a radical. In the case of ions, the charge on the ion is its valency or combining power (list below). To work out a formula by combining 'A' with 'B' the rule is: o number of 'A' x valency of 'A' = number of 'B' x valency of 'B', However it is easier perhaps? to grasp with ionic compound formulae. o In the electrically balanced stable formula, the total positive ionic charge must equal the total negative ionic charge. Example: o Aluminium oxide consists of aluminium ions Al3+ and oxide ions O2o number of Al3+ x charge on Al3+ = number of O2- x charge on O2o the simplest numbers are 2 of Al3+ x 3 = 3 of O2- x 2 (total 6+ balances total 6-) o so the simplest whole number formula for aluminum oxide is Al 2O3 14 www.studyguide.pk Examples of ionic combining power of ions (left, valency = numerical charge value) Examples of covalent combining power of atoms (valencies below) Hydrogen H (1) Chlorine Cl and other halogens (1) Oxygen O and sulphur S (2) Boron B and aluminium Al (3) Nitrogen (3, 4, 5) Carbon C and silicon Si (4) Phosphorus (P 3,5) Examples of working out covalent formulae 'A' (valency) 'B' (valency) deduced formula 1 of carbon C (4) balances 4 of hydrogen H (1) 1 x 4 = 4 x 1 = CH4 1 of nitrogen (3) balances 3 of chlorine Cl (1) 1 x 3 = 3 x 1 = NCl3 1 of carbon C (4) balances 2 of oxygen O (2) 1 x 4 = 2 x 2 = CO2 The diagram on the left illustrates the three covalent examples above for methane CH4 nitrogen trichloride NCl3 carbon dioxide CO2 Examples of working out ionic formulae 'A' (charge=valency) + 2 of Na (1) 1 of Mg 2+ 1 of Fe 3+ 2 of Fe 3+ (2) (3) (3) 'B' (charge=valency) 2 x 1 = 1 x 2 = Na2O - 1 x 2 = 2 x 1 = MgCl2 - 1 x 3 = 3 x 1 = FeF3 balances 1 of O (2) balances 2 of Cl (1) balances 3 of F (1) balances 3 of deduced formula 2- SO42- (2) 2 x 3 = 3 x 2 = Fe2(SO4)3 15 www.studyguide.pk The Structure of Atoms Atoms are the smallest particles of matter whose properties we study in Chemistry. However from experiments done in the late 19th and early 20th century it was deduced that atoms were made up of three fundamental sub-atomic particles (listed below) A Portrait of an Atom The diagram below gives some idea on the structure of an atom, it also includes some important definitions and notation used to describe atomic structure. The atomic number (Z) is also known as the proton number of the nucleus of a particular element. It is the proton number that determines the specific identity of a particular element and its electron structure. The mass number (A) is also known as the nucleon number (N), that is the number of particles in the nucleus of a particular isotope. Protons and neutrons are the nucleons present in the positive nucleus and the negative electrons are held by the positive nucleus in 'orbits' called energy levels or shells. In a neutral atom the number of protons equals the number of electrons. ISOTOPES Isotopes are atoms of the same element with different numbers of neutrons. This gives each isotope of the element a different mass or nucleon number but being the same element they have the same atomic or proton number. There are small physical differences between the isotopes eg the heavier isotope has a greater density or boiling point. However, because they have the same number of protons they have the same electronic structure and identical chemically. Examples are illustrated below. Do NOT assume the word isotope means it is radioactive, this depends on the stability of the nucleus www.studyguide.pk i.e. unstable atoms (radioactive) might be referred to as radioisotopes. Many isotopes are stable and NOT radioactive i.e. most of the atoms that make up you and the world around you! , and are the three isotopes of hydrogen with mass numbers of 1, 2 and 3, with 0, 1 and 2 neutrons respectively, but all have 1 proton. Hydrogen-1 is the most common, there is a trace of hydrogen-2 naturally but hydrogen-3 is very unstable and is used in atomic fusion weapons. and are the two isotopes of helium with mass numbers of 3 and 4, with 1 and 2 neutrons respectively but both have 2 protons. Helium-3 is formed in the Sun by the initial nuclear fusion process. Helium-4 is also formed in the Sun and as a product of radioactive alpha decay of an unstable nucleus. An alpha particle is a helium nucleus, it picks up two electrons and becomes the atoms of the gas helium. and are the two isotopes of sodium with mass numbers of 23 and 24, with 12 and 13 neutrons respectively but both have 11 protons. Sodium-23 is quite stable e.g. in common salt (NaCl, sodium chloride) but sodium-24 is a radio-isotope and is a gamma emitter used in medicine as a radioactive tracer e.g. to examine organs and the blood system. The relative atomic mass of an element is the average mass of all the isotopes present compared to 1/12th of the mass of carbon-12 atom (12C = 12.00000 ie the standard). The Electronic Structure of Atoms (electron configuration, arrangement in shells or energy levels) The electrons are arranged in energy levels or shells around the nucleus and with increasing distance from the nucleus. Each electron in an atom is in a particular energy level (or shell) and the electrons must occupy the lowest available energy level (or shell) available nearest the nucleus. When the level is full, the next electron(s) go into the next highest level (shell) available. There are rules about the maximum number of electrons allowed in each shell and you have to be able to work out the arrangements for the first 20 elements (see the Periodic Table diagrams further down). st o The 1 shell has a maximum of 2 electrons o The 2nd shell has a maximum of 8 electrons o The 3rd shell has a maximum of 8 electrons th th th o The 19 and 20 electrons go into the 4 shell (limit of GCSE knowledge) If you know the atomic (proton) number, you know it equals the number of electrons in a neutral atom, you then apply the rules to work out the electron arrangement (configuration). Examples: diagram, symbol or name of element (Atomic Number = number of electrons in a neutral atom), shorthand electron arrangement 2 www.studyguide.pk On Period 1 On Period 2 On Period 3 On Period 4 (2 elements only) .... .... .... (3 of the 8 elements) (3 of the 8 elements) ===> Kr [2.8.18.8], The Periodic Table and Electronic Structure Below are the electron arrangements for elements 1 to 20 set out in Periodic Table format (Hydrogen and The Transition metals etc. have been omitted). When you move down to the next period you start to fill in the next shell according to the maximum electrons in a shell rule (see previous section) 3 www.studyguide.pk o o o The first element in a period has one outer electron (eg sodium Na 2.8.1), and the last element has a full outer shell (eg argon Ar 2.8.8) Apart from hydrogen (H, 1) and helium (He, 2) the last electron number is the group number. and the number of shells used is equal to the Period Number. Atomic structure diagrams 4 www.studyguide.pk 5 www.studyguide.pk Structure, Bonding and properties 1. Why do atoms bond together? Some atoms are very reluctant to combine with other atoms and exist in the air around us as single atoms. These are the Noble Gases and have very stable electron arrangements eg 2, 2.8 and 2.8.8 because their outer shells are full. The first three are shown in the diagrams below and explains why Noble Gases are so reluctant to form compounds with other elements. (atomic number) and electron arrangement. All other atoms therefore, bond together to become electronically more stable, that is to become like Noble Gases in electron arrangement. Bonding produces new substances and usually involves only the 'outer shell' or 'valency' electrons The phrase CHEMICAL BOND refers to the strong electrical force of attraction between the atoms or ions in the structure. The combining power of an atom is sometimes referred to as its valency and its value is linked to the number of outer electrons of the original uncombined atom COVALENT BONDING - sharing electrons to form molecules with covalent bonds, the bond is usually formed between two non-metallic elements in a molecule. IONIC BONDING - By one atom transferring electrons to another atom. COORDINATE COVALENT BONDING A coordinate covalent bond is special because it involves a shared pair of electrons that came from a single atom. METALLIC BONDING The crystal lattice of metals consists of ions NOT atoms surrounded by a 'sea of electrons' forming giant lattice. These free or 'delocalised' electrons are the 'electronic glue' holding the particles together. There is a strong electrical force of attraction between these mobile electrons (-) and the 'immobile' positive metal ions (+) and this is the metallic bond. An ion is an atom or group of atoms carrying an overall positive or negative charge o eg Na+, Cl-, [Cu(H2O)]2+, SO42- etc. If a particle, as in a neutral atom, has equal numbers of protons (+) and electrons (-) the particle charge is zero ie no overall electric charge. The proton number in an atom does not change BUT the number of associated electrons can! If negative electrons are removed the excess charge from the protons produces an overall positive ion. If negative electrons are gained, there is an excess of negative charge, so a negative ion is formed. The charge on the ion is numerically related to the number of electrons transferred. For any atom or group of atoms, for every electron gained you get a one unit increase in negative charge, for every electron lost you get a one unit increase in the positive The atom losing electrons forms a positive ion (cation) and is usually a metal. The atom gaining electrons forms a negative ion (anion) and is usually a non-metallic element. www.studyguide.pk NOBLE GASES are very reluctant to share, gain or lose electrons to form a chemical bond. For most other elements the types of bonding and the resulting properties of the elements or compounds are described in detail below. In all the electronic diagrams ONLY the outer electrons are shown. 2. Covalent Bonding Covalent bonds are formed by atoms sharing electrons to form molecules. This type of bond usually formed between two non-metallic elements. The molecules might be that of an element ie one type of atom only OR from different elements chemically combined to form a compound. The covalent bonding is caused by the mutual electrical attraction between the two positive nuclei of the two atoms of the bond, and the electrons between them. One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared electrons between the same two atoms gives a double bond and it is possible for two atoms to share 3 pairs of electrons and give a triple bond. The bonding in Small Covalent Molecules The simplest molecules are formed from two atoms and examples of their formation are shown below. The electrons are shown as dots and crosses to indicate which atom the electrons come from, though all electrons are the same. The diagrams may only show the outer electron arrangements for atoms that use two or more electron shells. Examples of simple covalent molecules are … Example 1: 2 hydrogen atoms (1) form the molecule of the element hydrogen H2 and combine to form where both atoms have a pseudo helium structure of 2 outer electrons around each atom's nucleus. Any covalent bond is formed from the mutual attraction of two positive nuclei and negative electrons between them. Example 2: 2 chlorine atoms (2.8.7) form the molecule of the element chlorine Cl2 and combine to form where both atoms have a pseudo argon structure of 8 outer electrons around each atom. All the other halogens would be similar eg F2, Br2 and I2 etc. Valency of halogens is 1 here. Example 3: 1 atom of hydrogen (1) combines with 1 atom of chlorine (2.8.7) to form the molecule of the compound hydrogen chloride HCl 2 www.studyguide.pk and combine to form where hydrogen is electronically like helium and chlorine like argon. All the other hydrogen halides will be similar eg HF, HBr and HI etc. Example 4: 2 atoms of hydrogen (1) combine with 1 atom of oxygen (2.6) to form the molecule of the compound we call water H2O and and combine to form so that the hydrogen atoms are electronically like helium and the oxygen atom becomes like neon. The molecule can be shown as with two hydrogen - oxygen single covalent bonds. Hydrogen sulphide will be similar, since sulphur (2.8.6) is in the same Group 6 as oxygen. Valency of oxygen and sulphur is 2 here. Example 5: 3 atoms of hydrogen (1) combine with 1 atom of nitrogen (2.5) to form the molecule of the compound we call ammonia NH3 three of and one combine to form so that the hydrogen atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule can be shown as hydrogen single covalent bonds ( with three nitrogen - Example 6: 4 atoms of hydrogen (1) combine with 1 atom of carbon (2.4) to form the molecule of the compound we call methane CH4 four of and one of combine to form so that the hydrogen atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule can be shown as with four carbon hydrogen single covalent bonds. SiH4 will be similar because silicon (2.8.4) is in the same group as carbon. 3 www.studyguide.pk All the bonds in the above examples are single covalent bonds. Below are three examples 7-9, where there is a double bond in the molecule, in order that the atoms have stable Noble Gas outer electron arrangements around each atom. Carbon has a valency of 4. Example 7: Two atoms of oxygen (2.6) combine to form the molecules of the element oxygen O2. The molecule has one O=O double covalent bond Example 8: . Oxygen valency 2. One atom of carbon (2.4) combines with two atoms of oxygen (2.6) to form carbon dioxide CO2. The molecule can be shown as covalent bonds Valencies of C and O are 4 and 2 respectively. Example 9: with two carbon = oxygen double Two atoms of carbon (2.4) combine with four atoms of hydrogen (1) to form ethene C2H4. The molecule can be shown as with one carbon = carbon double bond and four carbon - hydrogen single covalent bonds The valency of carbon is still 4. Examples 10-13: The scribbles below illustrate some more complex examples. Ex. 10 nitrogen; Ex. 11 ethane; Ex. 12 chloromethane and Ex. 13 methanol. The valencies or combining power in theses examples are N 3, H 1, C 4, Cl 1 and O 2. 4 www.studyguide.pk MULTIPLE BONDS Why multiple bonds form: Nonmetal atoms bond to reach a stable or low energy condition. This happens when a main-group atom shares enough electrons to achieve a rare gas valence shell. Sometimes the number of electrons needed cannot be provided by sharing electrons simply in single bonds(pairs). Nitrogen molecules are an example of this situation. Nitrogen atoms have five valence electrons. We know that molecules of N2 exist based on a great deal of measurements. If two nitrogen atoms simply formed a single bond the dot structure would look like the illustraton below. Each nitrogen atom would have only six electrons not an octet. The single bond doesn't provide the two nitrogen atoms with an octet. If the octet rule is to be followed, the nitrogen atoms must share more than two electrons. The trial and error method is used to decide how many shared electrons are needed to create a structure where the octet rule is met. Since one bond didn't work the next thing to try is a double bond where the atoms share four electrons. Unfortunately the double bond structure only provides 7 valence shell electrons not eight. Because the single and double bonds didn't do the trick the next thing to try is a triple bond where the nitrogen atoms share six electrons. The count for both atoms when a triple bond is used in the structure shows that the octet rule is met. Multiple bonds are very common in carbon compounds. The carbon atom can form four bonds. These four bonds can be all single,CH3CH3, two single and one double, CH2CH2, two double, one single and one triple, CHCH. Ethane has a single bond carboncarbon bond Ethene has a double carbon-carbon bond Acetylene has a triple carbon-carbon bond 5 www.studyguide.pk The carbon-carbon single bond has a bond length of 154 picometers. The carbon-carbon double bond has a The carbon-carbon triple bond has a bond length of 133 picometers. bond length of 120 picometers. The carbon-carbon bond lengths decrease as the number of shared electrons and bonds increase. This is reasonable because there are more electrons attracted to each nucleus. The repulsions between positive nuclei decrease as more electrons are shared. The relative lengths for the three types of bonds are summarized here. REMEMBER these compare bonds between pairs of atoms like ; C:::C, C::C, C:C or C:::N, C::N, C:N or N:::N, N::N N:N or N::O , N:O The Properties of simple covalent molecular substances small molecules! The electrical forces of attraction, that is the chemical bond*, between atoms in any molecule are strong and most molecules do not change chemically on moderate heating.(* sometimes referred to as the intramolecular bond) However, the electrical forces** between molecules are weak and easily weakened further on heating. These weak attractions are known as **intermolecular forces and consequently the bulk material is not usually very strong. Consequently small covalent molecules tend to be volatile liquids, easily vapourised, or low melting point solids. On heating the inter-molecular forces are easily overcome with the increased kinetic energy gain of the particles and so have low melting and boiling points. They are also poor conductors of electricity because there are no free electrons or ions in any state to carry electric charge. Most small molecules will dissolve in a solvent to form a solution. Large Covalent Molecules and their Properties 6 www.studyguide.pk (macromolecules - giant covalent networks and polymers) It is possible for many atoms to link up to form a giant covalent structure or lattice. The atoms are usually non-metals. This produces a very strong 3-dimensional covalent bond network or lattice. This gives them significantly different properties from the small simple covalent molecules mentioned above. This is illustrated by carbon in the form of diamond. Carbon can form four single bonds so each carbon bonds to four others etc. This type of structure is thermally very stable and they have high melting and boiling points. They are usually poor conductors of electricity because the electrons are not usually free to move as they can in metallic structures. Also because of the strength of the bonding in all directions in the structure, they are often very hard, strong and will not dissolve in solvents like water. Silicon dioxide (silica, SiO2) has a similar 3D structure and properties, shown below diamond. The hardness of diamond enables it to be used as the 'leading edge' on cutting tools. Diamond is an allotrope of carbon. Two other allotropes of diamond are described below. Allotropes are different forms of the same element in the same physical state (3 solid forms of carbon are described here). Carbon also occurs in the form of graphite. The carbon atoms form joined hexagonal rings forming layers 1 atom thick. There are three strong covalent bonds per carbon, BUT, the fourth bond carbon can form from its four outer electrons, is shared between the three bonds shown (this requires advanced level concepts to fully explain, and this bonding situation also occurs in fullerenes described below). The layers are only held together by weak intermolecular forces shown by the dotted lines NOT by strong covalent bonds. Like diamond and silica (above) the large molecules of the layer ensure graphite has typically very high melting point because of the strong 2D bonding network (note: NOT 3D network).. Graphite will not dissolve in solvents because of the strong bonding BUT there are two crucial differences compared to diamond ... o Electrons, from the 'shared bond', can move freely through each layer, so graphite is a conductor like a metal (diamond is an electrical insulator and a poor heat conductor). Graphite is used in electrical contacts eg electrodes in electrolysis. o The weak forces enable the layers to slip over each other so where as diamond is hard material graphite is a 'soft' crystal, it feels slippery. Graphite is used as a lubricant. 7 www.studyguide.pk Bonding in polymers and 1-3 'dimension' concepts in macromolecules The bonding in polymers or plastics is no different in principle to the examples described above, but there is quite a range of properties and the difference between simple covalent and giant covalent molecules can get a bit 'blurred'. o Bonds between atoms in molecules, eg C-C, are called intra-molecular bonds. o The much weaker electrical attractions between individual molecules are called inter-molecular forces. In thermosoftening plastics like poly(ethene) the bonding is like ethane except there are lots of carbon atoms linked together to form long chains. They are moderately strong materials but tend to soften on heating and are not usually very soluble in solvents. The structure is basically a linear 1 dimensional strong bonding networks. Graphite structure is a layered 2 dimensional strong bond network made of layers of joined hexagonal rings of carbon atoms with weak inter-molecular forces between the layers. Thermosetting plastic structures like melamine have a 3 dimensional cross-linked giant covalent structure network similar to diamond or silica in principle, but rather more complex and chaotic! Because of the strong 3D covalent bond network they do not dissolve in any solvents and do not soften on heating and are much stronger than thermoplastics. More on polymers in Oil Notes and Extra Organic Chemistry Notes. 3. Ionic Bonding Ionic bonds are formed by one atom transferring electrons to another atom to form ions. Ions are atoms, or groups of atoms, which have lost or gained electrons. The atom losing electrons forms a positive ion (a cation) and is usually a metal. The overall charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in chemical reactions). The atom gaining electrons forms a negative ion (an anion) and is usually a non-metallic element. The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons. The examples below combining a metal from Groups 1 (Alkali Metals), 2 or 3, with a non-metal from Group 6 or Group 7 (The Halogens) Example 1: A Group 1 metal + a Group 7 non-metal eg sodium + chlorine sodium chloride NaCl or ionic formula Na+Cl- In terms of electron arrangement, the sodium donates its outer electron to a chlorine atom forming a single positive sodium ion and a single negative chloride ion. The atoms have become stable ions, because electronically, sodium becomes like neon and chlorine like argon. Na (2.8.1) + Cl (2.8.7) Na+ (2.8) Cl- (2.8.8) can be summarised electronically as [2,8,1] + [2,8,7] [2,8]+ [2,8,8]- 8 www.studyguide.pk ONE combines with ONE to form The valencies of Na and Cl are both 1, that is, the numerical charge on the ions. NaF, KBr, LiI etc. will all be electronically similar. Example 2: A Group 2 metal + a Group 7 non-metal eg magnesium + chlorine magnesium chloride MgCl2 or ionic formula Mg2+(Cl-)2 In terms of electron arrangement, the magnesium donates its two outer electrons to two chlorine atoms forming a double positive magnesium ion and two single negative chloride ions. The atoms have become stable ions, because electronically, magnesium becomes like neon and chlorine like argon. Mg (2.8.2) + 2Cl (2.8.7) Mg2+ (2.8) 2Cl- (2.8.8) can be summarised electronically as [2,8,2] + 2[2,8,7] [2,8]2+ [2,8,8]-2 ONE combines with TWO to form see * * NOTE you can draw two separate chloride ions, but in these examples a number subscript has been used, as in ordinary chemical formula. Example 3: A Group 3 metal + a Group 7 non-metal eg aluminium + fluorine aluminium fluoride AlF3 or ionic formula Al3+(F-)3 In terms of electron arrangement, the aluminium donates its three outer electrons to three fluorine atoms forming a triple positive aluminium ion and three single negative fluoride ions. The atoms have become stable ions, because electronically, aluminium and fluorine becomes electronically like neon. Valency of Al is, F is 1. Al (2.8.3) + 3F (2.7) Al3+ (2.8) 3F- (2.8) can be summarised electronically as [2,8,3] + 3[2,7] [2,8]3+ [2,8]-3 ONE combines with THREE to form 9 www.studyguide.pk Example 4: A Group 1 metal + a Group 6 non-metal eg potassium + oxygen potassium oxide K2O or ionic formula (K+)2O2- In terms of electron arrangement, the two potassium atoms donates their outer electrons to one oxygen atom. This results in two single positive potassium ions to one double negative oxide ion. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). Valencies, K 1, oxygen 2. Na 2O, Na2S, K2S etc. will be similar. 2K (2.8.8.1) + O (2.6) 2K+ (2.8.8) O2- (2.8) can be summarised electronically as 2[2,8,8,1] + [2,6] [2,8,8]+2 [2,8]2- TWO combine with ONE to form Example 5: A Group 2 metal + a Group 6 non-metal eg calcium + oxygen calcium oxide CaO or ionic formula Ca2+O2- In terms of electron arrangement, one calcium atom donates its two outer electrons to one oxygen atom. This results in a double positive calcium ion to one double negative oxide ion. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). the valency of both calcium and oxygen is 2. MgO, MgS, or CaS will be similar electronically (S and O both in Group 6) Ca (2.8.8.2) + O (2.6) Ca2+ (2.8.8) O2- (2.8) can be summarised electronically as [2,8,8,2] + [2,6] [2,8,8]2+ [2,8]2- ONE combines with ONE to form Example 6: A Group 3 metal + a Group 6 non-metal eg aluminium + oxygen aluminium oxide Al2O3 or ionic formula (Al3+)2(O2-)3 In terms of electron arrangement, two aluminium atoms donate their three outer electrons to three oxygen atoms. This results in two triple positive aluminium ions to three double negative oxide ions. All the ions have the stable electronic structure of neon 2.8. Valencies, Al 3 and O 2. 2Al (2.8.3) + 3O (2.6) 2Al3+ (2.8) 3O2- (2.8) can be summarised electronically as 2[2,8,3] + 3[2,6] [2,8]3+2 [2,8]2-3 TWO combines with THREE to form The Properties of Ionic Compounds 10 www.studyguide.pk The alternate positive and negative ions in an ionic solid are arranged in an orderly way in a giant ionic lattice structure shown on the left. The ionic bond is the strong electrical attraction between the positive and negative ions next to each other in the lattice. The bonding extends throughout the crystal in all directions. Salts and metal oxides are typical ionic compounds. This strong bonding force makes the structure hard (if brittle) and have high melting and boiling points, so they are not very volatile! The bigger the charges on the ions the stronger the bonding attraction eg magnesium oxide Mg2+O2- has a higher melting point than sodium chloride Na+Cl-. Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature. They are hard but brittle, when stressed the bonds are broken along planes of ions which shear away. They are NOT malleable like metals. Many ionic compounds are soluble in water, but not all, so don't make assumptions. The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current. However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion particles are now free. 4. BONDING IN METALS Electron-Sea Model of Metals In the electron-sea model, a metal crystal is considered to be a three-dimensional array of metal cations immersed in a sea of valence electrons. The delocalized valence electrons are free to move throughout the crystal and are not associated with any one particular metal cation. The mobility of the electrons accounts for the high electrical conductivity of metals. Thermal conductivity can also be ascribed to the mobile electrons that conduct heat by carrying kinetic energy from one part of the crystal to another. Three-dimensional delocalized bonding allows the metal to be both malleable and ductile. The crystal lattice of metals consists of ions NOT atoms surrounded by a 'sea of electrons' forming another type of giant lattice. The outer electrons (-) from the original metal atoms are free to move around between the positive metal ions formed (+). These free or 'delocalised' electrons are the 'electronic glue' holding the particles together. There is a strong electrical force of attraction 11 www.studyguide.pk between these mobile electrons (-) and the 'immobile' positive metal ions (+) and this is the metallic bond. This strong bonding generally results in dense, strong materials with high melting and boiling points. Metals are good conductors of electricity because these 'free' electrons carry the charge of an electric current when a potential difference (voltage!) is applied across a piece of metal. Metals are also good conductors of heat. This is also due to the free moving electrons. Non-metallic solids conduct heat energy by hotter more strongly vibrating atoms, knocking against cooler less strongly vibrating atoms to pass the particle kinetic energy on. In metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to transfer the particle kinetic energy more efficiently to 'cooler' atoms. Typical metals also have a silvery surface but remember this may be easily tarnished by corrosive oxidation in air and water. Unlike ionic solids, metals are very malleable, they can be readily bent, pressed or hammered into shape. The layers of atoms can slide over each other without fracturing the structure. The reason for this is the mobility of the electrons. When planes of metal atoms are 'bent' or slide the electrons can run in between the atoms and maintain a strong bonding situation. This can't happen in ionic solids. Note on Alloy Structure 1. Shows the regular arrangement of the atoms in a metal crystal and the white spaces show where the free electrons are (yellow circles actually positive metal ions). 2. Shows what happens when the metal is stressed by a strong force. The layers of atoms can slide over each other 3. and the bonding is maintained as the mobile electrons keep in contact with atoms, so the metal remains intact BUT a different shape. Shows an alloy mixture. It is NOT a compound but a physical mixing of a metal plus at least one other material (shown by red circle, it can be another metal eg Ni, a non-metal eg C or a compound of carbon or manganese, and it can be bigger or smaller than iron atoms). Many alloys are produced to give a stronger metal. The presence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The result is a stronger harder less malleable metal. 5.Coordinate Covalent Bonds(Dative Bonding) 12 www.studyguide.pk A coordinate covalent bond is special because it involves a shared pair of electrons that came from a single atom. Ammonia had a nitrogen atom with an unshared pair of electrons. These can be shared with an electron defficient atom like H +. ammonia ammonia proton ammonium ion Water molecules have two unshared pairs of electrons. These form coordinate covalent bonds with cations that are dissolved in water. This is one reason why water dissolves many ionic solids. The energy released when the water molecules bond to the cations is often enough to break up the ionic solid. water cation water and cation A dissolved cation will form as many as six coordinate covalent bonds. 6. INTERMOLECULAR FORCES Intermolecular forces are those forces that occur between particles (molecules, atoms, or ions). The strength of these forces at a given temperature dictates whether a substance will have the properties of a solid, a liquid, or a gas. The term van der Waals forces encompasses all types of intermolecular forces. All intermolecular forces arise from electrostatic interactions governed by the basic rule that like charges repel and unlike charges attract. Hydrogen Bonds 13 www.studyguide.pk A hydrogen bond is an attractive interaction between a hydrogen that is bonded to a very electronegative atom (O, N, or F) and an unshared electron pair on another electronegative atom. Hydrogen bonds can be quite strong. Substances that form hydrogen bonds have unusually high boiling points due to the extra energy that must be used to separate the molecules. DONE 14 www.studyguide.pk CHEMICAL TESTS INORGANIC TESTS TEST FOR hydrogen gas H2 carbon dioxide gas CO2 oxygen gas O2 Hydrogen chloride gas HCl, in water hydrochloric acid HI TEST METHOD OBSERVATIONS TEST CHEMISTRY lit splint or spill squeaky pop! (might 2H2(g) + O2(g) ==> 2H2O(l) + energy! see condensation on test tube) bubble into turns cloudy – fine Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l) limewater milky white precipitate (aqueous calcium of calcium carbonate hydroxide solution) glowing splint or spill re-ignites it flame (i) blue litmus and (i) litmus turns red, (ii) drop of silver (ii) white precipitate nitrate on the end with silver nitrate of a glass rod Hydrogen As above. In water as above but cream bromide HBr and they are HBr or yellow HI Hydrogen iodide hydrobromic acid precipitate and hydriodic acid. C(in wood) + O2(g) ==> CO2(g) (i) Strongly acid gas, (ii) in water forms chloride ions - hence precipitate with silver nitrate. as above - combination of acid and halide ion tests 1 www.studyguide.pk Sulphur dioxide gas SO2 freshly made potassium dichromate(VI) paper Ammonia gas strong pungent odour*, (i) red NH3 litmus, (ii) fumes conc. hydrochloric acid paper changes from orange to green the dichromate(VI) ion, Cr2O72-(aq) is reduced to the green Cr3+(aq) ion (i) litmus turns blue, (ii) white clouds with HCl fumes. (i) only common alkaline gas and (ii) forms fine ammonium chloride crystals with HCl (*volatile organic aliphatic amines give the same result, and smell more fishy) Chlorine (i) blue litmus, (ii) pungent green gas, (i) (i) non-metal, is acid in aqueous solution and a gas Cl2 drop silver nitrate litmus turns red and powerful oxidising agent, (ii) forms chloride ion in [test (ii) on on the end of a its own is no good, could be glass rod HCl] Iodine solid then is bleached white, (ii) white precipitate water (i) heating, (ii) test (i) purple vapour, (ii) aqueous solution blue black colour with or solid with starch starch solution solution Nitrogen( no simple relatively nasty brown gas IV) oxide unambiguous test (or nitrogen dioxide) strong oxidising agent NO2 Water liquid H2O (i) white anhydrous copper(II) sulphate, (ii) dry blue cobalt chloride paper Carbonate ion CO32- add any dilute (i) turns from white to (i) blue hydrated copper(II) crystals or solution blue, (ii) turns from formed, (ii) hydrated cobalt ion formed blue to pink [Co(H2O)6]2+ fizzing - colourless gas - turns limewater milky cloudy (see above CO2) carbonate/hydrogencarbonate + acid ==> salt + water + carbon dioxide, then white precipitate with limewater. (or hydrogencarbonate HCO3-) strong acid to the suspected carbonate - if colourless gas given off, test with limewater Sulphate ion [sulphate(VI)] SO42- to a solution of the white precipitate of Ba2+(aq) + SO42-(aq) ==> BaSO4(s) suspected sulphate barium sulphate add dilute any soluble barium salt + any soluble sulphate hydrochloric acid ==> barium sulphate and a few drops of barium chloride or nitrate solution Sulphite ion 2[sulphate(IV)] SO3 (i) add dilute hydrochloric acid to the suspected sulphite, (ii) test any gas evolved with fresh potassium (i) acrid choking sulphur dioxide gas formed, (ii) the dichromate paper turns from orange to green (i) sulphite salt + hydrochloric acid ==> chloride salt + sulphur dioxide, (ii) the sulphur dioxide reduces the dichromate(VI) to chromium(III). Note: sulphites do not give ppt. with acidified barium chloride/nitrate because sulphites dissolve in acids. 2 www.studyguide.pk dichromate(VI) paper Sulphid (i) If soluble, add a (i) Black ppt. of lead few drops lead(II) sulphide. e ion ethanoate solution. (ii) Rotten egg smell S2- for (ii) (ii) If solid, add dil. of hydrogen sulphide (i) Pb2+(aq) + S2-(aq) => PbS(s) (ii) MS(s) + 2H+(aq) => M2+(aq) + H2S(g) (e.g. M = Pb, Fe, Cu, Ni etc.) Then reaction (i) above occurs. dangerous hydrogen sulphide formed HCl(aq) acid, test gas with lead(II) ethanoate paper. and the H2S gas turns lead(II) ethanoate paper black. Chloride ion (i) if soluble, add dilute nitric acid and silver nitrate solution, (ii) if insoluble salt, add conc. sulphuric acid, warm if necessary then test gas as for HCl above. (i) white precipitate of silver chloride soluble in dilute ammonia, (ii) get fumes of hydrogen chloride which turn blue litmus red and give a white precipitate with silver nitrate solution (i) Ag+(aq) + Cl-(aq) ==> AgCl(s) , any soluble silver salt + any soluble chloride ==> silver chloride precipitate, (ii) Cl-(s) + H2SO4(l) ==> HSO4-(s) + HCl(g) , then Ag+(aq) + Cl-(aq) ==> AgCl(s) (i) if soluble, add dilute nitric acid and silver nitrate solution, (ii) if insoluble salt, add conc. sulphuric acid, warm if necessary (i) cream precipitate of silver bromide, only soluble in concentrated ammonia, (ii) orange vapour, test for sulphur dioxide. (i) Ag+(aq) + Br-(aq) ==> AgBr(s) any soluble silver salt + any soluble bromide ==> silver bromide precipitate, (ii) bromide ion is oxidised to bromine and the sulphuric acid is reduced to sulphur dioxide (i) if soluble, add dilute nitric acid and silver nitrate solution, (ii) if insoluble salt can heat with conc. sulphuric acid, (ii) get purple fumes of iodine and very smelly hydrogen sulphide, (iii) if soluble, add lead(II) nitrate solution (i) yellow precipitate of silver iodide insoluble in concentrated ammonia, (ii) purple vapour and rotten egg smell!, (iii) a yellow precipitate forms (i) Ag+(aq) + I-(aq) ==> AgI(s) , any soluble silver salt + any soluble iodide ==> silver iodide precipitate, (ii) iodide ion is oxidised to iodine and the sulphuric acid is reduced to hydrogen sulphide, (iii) insoluble lead(II) iodide formed, Pb2+(aq) + 2I-(aq) ==> PbI2(s) Cl- Bromide ion Br- Iodide ion I- Nitrate ion [or (i) boil the nitrate(V)] NO3 suspected nitrate (i) the fumes contain with sodium ammonia, hydroxide solution which turns red litmus and fine aluminium blue, see ammonia powder (Devarda's test details Alloy) (ii) Where the liquids (i) the aluminium powder is a powerful reducing agent and converts the nitrate ion, NO3-, into ammonia gas, NH3 (ii) NO complex of iron(II) formed 3 www.studyguide.pk meet a brown ring forms (ii) Add iron(ii) sulphate solution and then conc. sulphuric acid (the 'brown ring' test) Nitrite ion [or No simple test, (i) in acid solution it decomposes to give nasty brown fumes of NO2, (ii) it nitrate(III)] NO2 - decolourises (purple ==> colourless) acidified potassium manganate(VII), (iii) it liberates iodine from acidified potassium iodide solution, (iv) forms ammonia with hot Al powder/NaOH(aq) and gives 'brown ring' test - see nitrate tests above. Ammonium ion NH4+ no smell at first, add COLD sodium hydroxide solution to the suspected ammonium salt and test any gas turns blue with red litmus smelly ammonia evolved! and red litmus (i) litmus turns red, Hydrogen ion (i) litmus or + universal indicator variety of colours with ie acids! H or univ. ind. strong - red, or pH meter, (ii) + H3O (note: to add a little sodium weak - yellow/orange, completely identify acids you need to test for the anion eg chloride for HCl etc.) Hydroxide ion ie an alkali OH- (note: to completely identify alkalis you need to test for the cation eg sodium for NaOH etc.) Positive metal cations via flame tests (see below for NaOH and NH3 for metal ion tests too) Positive metal cations via sodium hydroxide (NaOH) ammonia gas is evolved: NH4+(aq) + OH-(aq) ==> NH3(g) + H2O(l) (i) pH meter gives a value of less than 7, the lower the pH number the stronger the acid, the higher the H+ concentration, (ii) HCO3-(aq) + H+(aq) ==> H2O(l) + CO2(g) hydrogencarbonate (ii) fizzing with any powder carbonate - test for CO2 as above (i) litmus or universal indicator or pH meter, (ii) add ammonium salt (i) turns litmus blue, variety of colours univ. ind. dark green violet for weak strong, (ii) if strongly alkaline ammonia should be released, see ammonia test for rest of details (i) pH meter gives a value of more than 7, the higher the pH number the stronger the alkali, the higher the OH- concentration, (ii) ammonia gas is evolved: NH4+(aq) + OH-(aq) ==> NH3(g) + H2O(l) The metal salt or other compound is mixed with concentrated hydrochloric acid and a sample of the mixture is heated strongly in a bunsen flame on the end of a cleaned nichrome wire (platinum if you can afford it!) lithium Li+ crimson All colours are due to electronic excitation to a sodium Na+ yellow higher level. You see the light emitted as the electron returns to its lower more stable level. potassium K+ lilac This is the basis of atomic emission and calcium Ca2+ brick absorption spectroscopy. Aluminium, magnesium, red iron and zinc do not produce a useful identifying barium Ba2+ apple flame colour. green Dilute sodium hydroxide solution is added to a solution aluminium ion: Al3+(aq) + 3OH-(aq) ==> Al(OH)3(s) white precipitate * The ppt. is not soluble in excess of the weak alkali ammonia, but dissolves in the strong alkali sodium hydroxide: Al(OH) 3(s) + 3OH-(aq) ==> [Al(OH)6]3-(aq) (amphoteric behaviour) copper(II) Cu2+ blue/green 4 www.studyguide.pk containing the or ammonia (NH3) solutions (both alkalis, suspected ion. giving hydroxide ions, OH-, in their solutions) Both the precipitate formed and the effect of excess alkali are important observations. All precipitates white, unless otherwise stated and all tend to be gelatinous in nature. calcium ion: Ca2+(aq) + 2OH-(aq) ==> Ca(OH)2(s) white ppt. * The ppt. is not soluble in excess of NH3 or NaOH. magnesium ion: Mg2+(aq) + 2OH-(aq) ==> Mg(OH)2(s) white ppt. * The ppt. is not soluble in excess of NH3 or NaOH. You could distinguish Mg from Ca with a flame test. copper(II) ion: Cu2+(aq) + 2OH-(aq) ==> Cu(OH)2(s) ***blue/turquoise ppt. - this does dissolve in excess ammonia to give a deep blue solution. iron(II) ion: Fe2+(aq) + 2OH-(aq) ==> Fe(OH)2(s) dark green ppt.* The ppt. is not soluble in excess of NH3 or NaOH. iron(III) ion: Fe3+(aq) + 3OH-(aq) ==> Fe(OH)3(s) brown ppt.* The ppt. is not soluble in excess of NH3 or NaOH. zinc ion: Zn2+(aq) + 2OH-(aq) ==> Zn(OH)2(s) white ppt. The ppt. dissolves in both excess sodium hydroxide or ammonia to give a clear colourless solution. The test can be repeated with aqueous ammonia solution (sometimes wrongly called 'ammonium hydroxide'). The observations are usually, but not always, similar. ppt. = precipitate. MISCELLANEOUS CATION TESTS: (i) add potassium (i) Pb2+(aq) +2I-(aq) ==>PbI2(s) lead(II) iodide ppt. iodide solution => yellow precipitate (i) Lead(II) ion Metal Carbonates Sometimes heating a metal carbonate strongly to decompose it provides some clues to its identity. Adding acid => CO2 and the colour of the resulting solution (eg blue Cu2+(aq), may also provide clues. The metal ion solution might also give a flame colour or a copper(II) carbonate==> copper(II) oxide + carbon dioxide: CuCO3(s) ==> CuO(s) + CO2(g) [green] ==> [black] + [colourless gas, test with limewater, white precipitate] zinc carbonate==> zinc oxide + carbon dioxide ZnCO3(s) ==> ZnO(s) + CO2(g) [white] ==> [yellow hot, white cold] +[colourless gas, test with limewater, white precipitate] 5 www.studyguide.pk hydroxide precipitate with sodium hydroxide eg copper. ORGANIC TESTS TEST FOR TEST METHOD ALKENE or bubble gas through, or add liquid to, a solution of bromine in hexane or water alkyne any other non-aromatic unsaturated hydrocarbons OBSERVATIONS TEST CHEMISTRY the orange/brown R2C=CR2 + Br2 ==> BrR2C-CR2Br bromine, decolourises, as a saturated RC CR + 2Br2 ==> Br2RC-CRBr2 colourless organic bromo-compound is R = H, alkyl or aryl formed (saturated alkanes give no fast reaction with bromine) 6 www.studyguide.pk The Three States of Matter KEYWORDS: Boiling Boiling point Brownian motion Condensing Cooling curve Diffusion Evaporation Freezing Freezing point Gas particle picture Heating curve Liquid particle picture Melting Melting point Properties of gases Properties of liquids Properties of solids sublimation Solid particle picture The particle model of a Gas Almost no forces of attraction between the particles so they are completely free of each other. Particles widely spaced and scattered at random throughout the container so there is no order in the system. Particles move rapidly in all directions, frequently colliding with each other and the side of the container. With increase in temperature, the particles move faster as they gain kinetic energy. Using the particle model to explain the properties of a Gas Gases have a very low density (‘light’) because the particles are so spaced out in the container. o Density order: solid > liquid >>> gases Gases flow freely because there are no effective forces of attraction between the particles. o Ease of flow order: gases > liquids >>> solids (no real flow in solid unless you powder it!) Gases have no surface, and no fixed shape or volume, and because of lack of particle attraction, they always spread out and fill any container (so gas volume = container volume). Gases are readily compressed because of the ‘empty’ space between the particles. o Ease of compression order: gases >>> liquids > solids (almost impossible to compress a solid) If the ‘container’ volume can change, gases readily expand* on heating because of the lack of particle attraction, and readily contract on cooling. o On heating, gas particles gain kinetic energy, move faster and hit the sides of the container more frequently, and significantly, they hit with a greater force. o Depending on the container situation, either or both of the pressure or volume will increase (reverse on cooling). o Note: * It is the gas volume that expands NOT the molecules, they stay the same size! The natural rapid and random movement of the particles means that gases readily ‘spread’ or diffuse. Diffusion is fastest in gases where there is more space for them to move. www.studyguide.pk o o The rate of diffusion increases with increase in temperature as the particles gain kinetic energy and move faster. Other evidence for random particle movement: When smoke particles are viewed under a microscope they appear to 'dance around' when illuminated with a light beam at 90o to the viewing direction. This is because the smoke particles show up by reflected light and 'dance' due to the millions of random hits from the fast moving air molecules. This is called 'Brownian motion'. At any given instant of time, the hits will not be even, so the smoke particle get a greater bashing in a random direction. If a long glass tube is filled at one with a plug of cotton wool soaked in conc. hydrochloric acid, and a similar plug of conc. ammonia solution at the other end. If left undisturbed and horizontal, despite the lack of tube movement (eg shake to mix), a white cloud forms about 1/3rd along from the conc. acid tube end. What happens is the colourless gases ammonia and hydrogen chloride diffuse down the tube and react to form fine white crystals of the salt ammonium chloride. NH3(g) + HCl(g) ==> NH4Cl(s) Note the rule: The smaller the molecular mass, the faster the molecules move. Therefore the smaller the molecular mass, the faster the gas diffuses. eg Mr(NH3) = 14 + 1x3 = 17, moves faster than Mr(HCl) = 1 + 35.5 = 36.5 AND that's why they meet nearer the HCl end of the tube! So the experiment is not only evidence for molecule movement, its also evidence that different molecular masses move on at different speeds. A coloured gas, that is heavier than air, is put into a gas jar and a second gas jar is placed over it separated with a cover. If the cover is removed than coloured gas diffuses into the colourless air above. It can't be due to convection because the more dense gas starts at the bottom! No 'shaking' or other means of mixing is required. The random movement of both lots of particles is enough to ensure that both gases are completely mixed eventually. This is clear evidence for the process of diffusion due to particle movement. 2 www.studyguide.pk The particle model of a Liquid Much greater forces of attraction between the particles in a liquid compared to gases, but not quite as much as in solids. Particles quite close together but still arranged at random throughout the container, there is a little close range order as you can get clumps of particles clinging together temporarily. Particles moving rapidly in all directions but more frequently collisions with each other than in gases due to shorter distances between particles. With increase in temperature, the particles move faster as they gain kinetic energy, so increased collision rates, increased collision energy and increased rate of diffusion. Using the particle model to explain the properties of a Liquid Liquids have a much greater density than gases (‘heavier’) because the particles are much closer together because of the attractive forces. Liquids usually flow freely despite the forces of attraction between the particles but liquids are not as ‘fluid’ as gases. o Note 'sticky' or viscous liquids have much stronger attractive forces between the molecules BUT not strong enough to form a solid. Liquids have a surface, and a fixed volume (at a particular temperature) because of the increased particle attraction, but the shape is not fixed and is merely that of the container itself. Liquids are not readily compressed because of the lack of ‘empty’ space between the particles. Liquids will expand on heating but nothing like as much as gases because of the greater particle attraction restricting the expansion (will contract on cooling). o Note: When heated, the liquid particles gain kinetic energy and hit the sides of the container more frequently, and more significantly, they hit with a greater force, so in a sealed container the pressure produced can be considerable! The natural rapid and random movement of the particles means that liquids ‘spread’ or diffuse. Diffusion is much slower in liquids compared to gases because there is less space for the particles to move in and more ‘blocking’ collisions happen. Evidence for random particle movement in liquids: o If coloured crystals of eg the highly coloured salt crystals of potassium manganate(VII) are dropped into a beaker of water and covered at room temperature. Despite the lack of mixing, convection etc. the bright purple colour of the dissolving salt slowly spreads throughout all of the liquid but it is much slower than the gas experiment described above. o When pollen grains are viewed under a microscope they appear to 'dance around' when illuminated with a light beam at 90o to the viewing direction. This is because the pollen grains show up by reflected light and 'dance' due to the millions of random hits from the fast moving water molecules. This is called 'Brownian motion' after a botanist called Brown first described the effect. At any given instant of time, the hits will not be even all round the pollen grain, so they get a greater number of hits in a random direction. 3 www.studyguide.pk The particle model of a Solid The greatest forces of attraction are between the particles in a solid and they pack together as tightly as possible in a neat and ordered arrangement. The particles are too strongly held together to allow movement from place to place but the particles vibrate about there position in the structure. With increase in temperature, the particles vibrate faster and more strongly as they gain kinetic energy. Using the particle model to explain the properties of a Solid Solids have the greatest density (‘heaviest’) because the particles are closest together. Solids cannot flow freely like gases or liquids because the particles are strongly held in fixed positions. Solids have a fixed surface and volume (at a particular temperature) because of the strong particle attraction. Solids are extremely difficult to compress because there is no real ‘empty’ space between the particles. Solids will expand a little on heating but nothing like as much as liquids because of the greater particle attraction restricting the expansion (contract on cooling). o The expansion is caused by the increased energy of particle vibration, forcing them further apart. Diffusion is almost impossible in solids because the particles are too strongly held and there are no ‘empty spaces’ for particles to move into. Changes of State 4 www.studyguide.pk Evaporation and Boiling (liquid to gas) On heating particles gain kinetic energy and move faster. In evaporation* and boiling the highest kinetic energy molecules can ‘escape’ from the attractive forces of the other liquid particles. The particles lose any order and become completely free to form a gas or vapour. Energy is needed to overcome the attractive forces in the liquid and is taken in from the surroundings. This means heat is taken in, so evaporation or boiling are endothermic processes. If the temperature is high enough boiling takes place. Boiling is rapid evaporation anywhere in the bulk liquid and at a fixed temperature called the boiling point and requires continuous addition of heat. The rate of boiling is limited by the rate of heat transfer into the liquid. * Evaporation takes place more slowly at any temperature between the melting point and boiling point, and only from the surface, and results in the liquid becoming cooler due to loss of higher kinetic energy particles. 5 www.studyguide.pk Condensing (gas to liquid) On cooling, gas particles lose kinetic energy and eventually become attracted together to form a liquid. There is an increase in order as the particles are much closer together and can form clumps of molecules. The process requires heat to be lost to the surroundings ie heat given out, so condensation is exothermic. o This is why steam has such a scalding effect, its not just hot, but you get extra heat transfer to your skin due to the exothermic condensation on your surface! Melting (solid to liquid) When a solid is heated the particles vibrate more strongly as they gain kinetic energy and the particle attractive forces are weakened. Eventually, at the melting point, the attractive forces are too weak to hold the particles in the structure together in an ordered way and so the solid melts. The particles become free to move around and lose their ordered arrangement. Energy is needed to overcome the attractive forces and give the particles increased kinetic energy of vibration. So heat is taken in from the surroundings and melting is an endothermic process. Freezing (liquid to solid) On cooling, liquid particles lose kinetic energy and so can become more strongly attracted to each other. Eventually at the freezing point the forces of attraction are sufficient to remove any remaining freedom and the particles come together to form the ordered solid arrangement. Since heat must be removed to the surroundings freezing is an exothermic process!!! 6 www.studyguide.pk Cooling and Heating Curves Note the temperature stays constant during the state changes of condensing Tc and freezing Tf. This is because all the energy removed on cooling at these temperatures weakens the inter-particle forces without temperature fall. A cooling curve summarises the changes: gas liquid solid Note the temperature stays constant during the state changes of melting at Tm and boiling at Tb. This is because all the energy absorbed in heating at these temperatures goes into weakening the inter-particle forces without temperature rise. A heating curve summarises the changes: Solid liquid gas 7 www.studyguide.pk Sublimation Sublimation: o This is when a solid, on heating, directly changes into a gas, AND the gas on cooling re-forms a solid directly. They usually involve just a physical change BUT its not always that simple! Theory in terms of particles: o When the solid is heated the particles vibrate with increasing force from the added thermal energy. If the particles have enough kinetic energy of vibration to partially overcome the particleparticle attractive forces you would expect the solid to melt. HOWEVER, if the particles at this point have enough energy at this point that would have led to boiling, the liquid will NOT form and the solid turns directly into a gas. Overall endothermic change, energy absorbed and 'taken in' to the system. o On cooling, the particles move slower and have less kinetic energy. Eventually, when the particle kinetic energy is low enough, it will allow the particleparticle attractive forces to produce a liquid. BUT the energy may be low enough to permit direct formation of the solid, ie the particles do NOT have enough kinetic energy to maintain a liquid state! Overall exothermic change, energy released and 'given out' to the surroundings. Examples: 1. Even at room temperature bottles of solid iodine show crystals forming at the top of the bottle above the solid. The warmer the laboratory, the more crystals form when it cools down at night! 2. I2 (s) I2 (g) (physical change only) The formation of a particular form of frost involves the direct freezing of water vapour (gas). Frost can also evaporate directly back to water vapour (gas) and this happens in the 'dry' extremely cold winters of the Gobi Desert. 3. H2O (s) H2O (g) (physical change only) Solid carbon dioxide (dry ice) is formed on cooling the gas down to less than -78oC. On warming it changes directly to a very cold gas!, condensing any water vapour in the air to a 'mist', hence its use in stage effects. 4. CO2 (s) CO2 (g) (physical change only) On heating strongly in a test tube, the white solid ammonium chloride, decomposes into a mixture of two colourless gases ammonia and hydrogen chloride. On cooling the reaction is reversed and solid ammonium chloride reforms at the cooler top of the test tube. Ammonium chloride + heat NH4Cl(s) NH3(g) + HCl(g) ammonia + hydrogen chloride (this involves both chemical and physical changes) 8 www.studyguide.pk SOLUTIONS & SOLUBILITIES Solutions A solution is a homogeneous mixture consisting of particles 0.1–2.0 nm in diameter. Homogeneous mixtures having larger particles (2–500 nm) are classified as colloids. Suspensions are mixtures with even larger particles, but they are not considered true solutions because they separate upon standing. We usually think of solutions as solids dissolved in liquids, or perhaps a mixture of two liquids, but there are many other kinds of solutions as well. Some examples are provided in Table 11.1. For solutions in which a gas or solid is dissolved in a liquid, the dissolved substance is called the solute and the liquid is called the solvent. When one liquid is dissolved in another, the minor component is usually considered the solute. Solution: a homogeneous mixture containing particles the size of a typical ion or covalent molecule Colloid: a homogeneous mixture containing particles with diameters in the range 2–500 nm Solute: the dissolved substance in a solution Solvent: the major component in a solution www.studyguide.pk Factors Affecting Solubility A solution is saturated when no additional solute can be dissolved. There is a dynamic equilibrium established. A supersaturated solution can form when more than the equilibrium amount of solute is dissolved at an elevated temperature, and then the supersaturated solution is slowly cooled. The amount of solute per unit solvent required to form a saturated solution is called the solute's solubility. A substance's solubility is a characteristic of that substance. Figure 11.5 illustrates the solubilities of some solids, and their temperature dependence. When two liquids are completely soluble in each other they are said to be miscible. The effect of temperature on gas solubility is more predictable than solid solubility. Most gases become less soluble in water as the temperature increases. 2 www.studyguide.pk Pressure has little effect on the solubility of liquids and solids. The solubility of gases is strongly influenced by pressure Saturated solution: a solution containing the maximum possible amount of dissolved solute at equilibrium Supersaturated solution: a solution containing a greater-than-equilibrium amount of solute Solubility: the amount of a substance that dissolves in a given volume of solvent Miscible: mutually soluble in all proportions Henry's law: The solubility of a gas in a liquid at a given temperature is directly proportional to the partial pressure of the gas over the solution 3 www.studyguide.pk 4 www.studyguide.pk METALS IMPORTANCE OF METALS More then three quarters of elements in the periodic table are metals. We cook food, heat the food on metals stoves. We travel in metallic vehicals like car, airplanes, ships. The concreate wall and roofs of houses are reinforced with metal rods. Coins, taps, cutlary, pins, needles paper pins etc. are all made up of metals. PHYSICAL PROPERTIES OF METALS PROPERTY REASON Metals are malleable and ductile i.e. they can be drawn into sheets and wires layers of metals can slide over each other. Metals conduct electricity. they have free moving electrons in their outer most shell. Metals usually have high melting points, high boiling points and high densities Strong metallic bond CHEMICAL PROPERTIES OF METALS Reaction with water Potassium reacts vigrously with cold water to form salt and hydrogen gas. The reaction is so exothermic that the hydrogen gas produced, burn in air. Potassium + water ---------- Potassium hydroxide + Hydrogen 2K(s) + 2H2O (l) ------------- 2KOH(aq) + H2(g) Sodium reacts with cold water in the same way. Sodium + water -------------- Sodium hydroxide + hydrogen gas 2Na(s) + 2H2O (l) ------------- 2NaOH(aq) + H2(g) Calcium reacts readily with cold water and vigrously with hot water to produce salt and hydrogen gas. Calcium + water ------------ Calcium hydroxide + hydrogen gas Ca(s) + 2H2O (l) ------------- Ca(OH)2(aq) + H2(g) www.studyguide.pk Magnesium reacts very slowly with cold water but vigrously with steam to form salt and hydrogen gas. Magnesium + steam --------- Magnesium oxide + hydrogen gas Mg(s) + H2O (g) ------------- MgO(s) + H2(g) Zinc do not react with cold water but reacts slowly with steam to form zinc oxide and hydrogen gas. Zink + steam --------------- Zink oxide + hydrogen gas Zn(s) + H2O (g) ------------- ZnO(s) + H2(g) Iron do not react with cold water but rusting occur very slowly in the presence of oxygen. Red hot iron reacts very slowly with steam to produce salt and hydrogen gas. Iron + steam --------------- Iron oxide + hydrogen 3Fe(s) + 4H2O (g) ------------- Fe3O4(s) + 4H2(g) Copper do not react with water under any condition Silver do not react with water in any condition. Reaction with Hydrochloric acid Potassium and sodium reacts explosively to form salt and hydrogen gas. The reaction is so exothermic that the hydrogen gas produced, burn in air. Potassium + Hydrochloric acid ------ Potassium chloride + hydrogen 2K(s) + 2HCl (aq) ------------- 2KCl(aq) + H2(g) Sodium + hydrochloric acid ------ Sodium chloride + hydrogen 2Na(s) + 2HCl (aq) ------------- 2NaCl(aq) + H2(g) Calcium reacts vigorously` to produce calcium chloride and hydrogen gas. Calcium + hydrochloric acid ------- Calcium chloride + hydrogen gas Ca(s) + 2HCl (aq) ------------- CaCl2(aq) + H2(g) 2 www.studyguide.pk Magnesium reacts very fastly to form magnesium chloride and hydrogen gas. Magnesium + hydrochloric acid -------- Magnesium chloride + hydrogen gas Mg(s) + 2HCl (aq) ------------- MgCl2(aq) + H2(g) Zinc reacts moderately to form zinc chloride and hydrogen gas. Zink + hydrochloric acid ------ Zink chloride + hydrogen gas Zn(s) + 2HCl (aq) ------------- ZnCl2(aq) + H2(g) Iron reacts slowly to produce iron chloride and hydrogen gas. Iron + hydrochloric acid -------- Ironchloride + hydrogen gas Fe(s) + 2HCl (aq) ------------- FeCl2(aq) + H2(g) Copper do not react with dilute HCl Silver do not react with dilute HCl Reaction with oxygen Potassium tarnishes in the presence of oxygen to form potassium oxide K2O Potassium + oxygen ------- Potassium oxide 4K(s) + O2(g) ------------------- 2K2O(s) Sodium burns with a yellow flame to produce odium oxide Na2O Sodium + Oxygen --------------- Sodium Oxide 4 Na(s) + O2(g) ------------- 2 Na2O(s) Copper powder burns with dull red glow to form copper oxide. CuO Copper + Oxygen ----------------- Copper oxide 2Cu(s) + O2(g) ------------------- 2CuO(s) Iron powder or wire burns with a bright yellow flame to form iron oxide Fe3O4 Iron Fe(s) + Oxygen ---------------- + O2(g) ------------------- Iron oxide 2Fe3O4(s) 3 www.studyguide.pk Magnesium burns with a bright white flame to produce white solid magnesium oxide. MgO Magnesium + oxygen ----------- Magnesium oxide 2Mg(s) + O2(g) ------------------- 2MgO(s) REACTIVITY SERIES Metals can be arranged in order of their chemical reactivity. The reactivity series is based on the reaction of metals with water or dilute hydrochloric acid. When metal recats with acid or water, the metal atom lose electron to become ion. Metal(s) + H2O (l) ------------------ Metal+ ion H2 (g) Metal (s) + HCl (aq) ---------------- Metal + ion + + OH- ion + Cl- ion + H2(g) The more readily a metal gives up electrons to form ions, the more reactive it is. A metal high up in the reactivity series Reacts vigorously with chemicals Readily gives up electrons in reactions to form positive ions Corrode easily A metal low down in the reactivity series Does not Reacts vigorously with chemicals Does not Readily gives up electrons in reactions to form positive ions Does not Corrode easily 4 www.studyguide.pk Hydrogen is sometimes placed in the reactivity series. Metals below hydrogen in the series do not react with acids to produce hydrogen gas. Please -------- Potassium Send -------- Sodium Cats ------- Calcium Monkeys -------Magnesium And -------Aluminium Zebras -------Zinc In ------Iron Large ----- Lead Hired ------ Hydrogen Cages ------Copper Make ---- Mercury Sure -----Silver Good -----Gold Padlock ----- Platinium DISPLACEMENT OF METALS Displacement of metals from solutions A more reactive metal will displace the ions of any less reactive metal in the reactivity series, from solution. Zinc + copper (II) sulphate solution ----------- Copper + zinc sulphate solution. Zn (s) + CuSO4 (aq) -------------------------- Cu(s) + ZnSO4 (aq) Zinc displace copper from the copper sulphate solution because it is more reactive than copper and readily give up electrons to form positive ions. The electrons are transferred from zinc atom to copper (II) ions. Cu2+ (aq) + Zn(s) blue solution ---------------- Cu (s) + redish-brown solid Zn2+ (aq) colourless 5 www.studyguide.pk Other examples: Ag+ (aq) + Cu (s) Zn2+ (aq) Mg (s) + ------------------ ---------------- Displacement of metals from metallic oxides by more reactive metals A metal will take oxygen from the oxide of any metal below it in the reactivity series. For example, when magnesium powder and copper (II) oxide powder is heated there is a vigrous exothermic reaction. The magnesium takes oxygen from copper (II) oxide to from magnesium oxide and copper metal. Magnesium Mg(s) + Copper oxide ----------- + CuO(s) ------heat------------- Magnesium oxide + Copper MgO(s) + Cu(l) Thermite reaction reaction. Aluminium + Iron oxide ---------------- Iron + Aluminium oxide 2Al (s) + Fe2O3 (s) -------heat--------- 2Fe(l) + Al2O3 (s) Reaction of metallic oxides with hydrogen Hydrogen can take oxygen from metallic oxides, producing the metal and water. For example when hyrogen is passed over hot lead (II) oxide, lead metal and water are produced. Lead (II) oxide PbO (s) + hydrogen -------------- lead + H2(g) ---------------- Pb(s) Copper (II) oxide + hydrogen CuO (s) -------heat------------ + H2 + + water. H2O (l) ----heat--------- copper + water. Cu (s) + H2O (l) The less reactive the metal, the easier it is for hydrogen to take oxygen from its oxide. The oxides of vary recative metals such as aluminium oxide and sodium oxide cannot be reduced to the metal by hydrogen. Reaction of metallic oxides with carbon. Carbon can take up oxygen from the oxide of metals which are not too high in the reactivity series. For example a mixture of charcoal and copper (II) oxide reacts when heated together 6 www.studyguide.pk Copper (II) oxide 2CuO (s) + Carbon -----------heat---------- copper + C(s) -------heat--------- 2Cu (s) + carbon dioxide. + CO2 (g) The more reactive the metal the more harder it for carbon to take oxygen from its oxide. Iron is more reactive than copper, iron oxide and carbon must be heated very strongly for the reaction to take place. Iron (II) oxide + Carbon -----strong heating------------- 2FeO (s) C (s) --------strong heating---------- + 2Fe Iron + + carbon dioxide. CO2 Carbon is unable to take oxygen from the oxides of very reactive such as calcium and sodium. THE EXTRACTION OF METALS Most of the metals are found as compounds called minerals. Minerals are usually found mixed with large amounts of impurities. These impure minerals are called ores. A ROCK is a mixture of minerals from which useful substances can be made. A MINERAL is a solid element or compound found naturally in the Earth’s crust. A METAL ORE is a mineral or mixture of minerals from which economically viable amounts of metal can be obtained. Two important ores to know: Haematite for Iron [contains iron(III) oxide, Fe2O3] Bauxite for Aluminium [contains aluminium oxide, Al2O3] Some important minerals Name of Mineral Chemical Name Formula Metal extracted Usual method of Extraction Bauxite Aluminium oxide Al2O3 Aluminium Galena Lead sulphide PbS Lead Haematite Sphalerite Iron (III) oxide Zinc Sulphide Fe2O3 ZnS Iron Zinc Copper pyrite Copper iron sulphide CuFeS2 Copper Electrolysis of oxide dissolved in molten cryolite. Sulphide is roasted in air and the oxide produced is Heat oxide with carbon Sulphide is roasted in air and the oxide produced is heated with carbon. Sulphide ore is roasted in air 7 www.studyguide.pk Introduction The Earth's crust contains many different rocks. Rocks are a mixture of minerals and from some we can make useful substances. A mineral can be a solid metallic or non-metallic element or a compound found naturally in the Earth's crust. A metal ore is a mineral or mixture of minerals from which economically viable amounts of metal can be extracted, i.e. its got to have enough of the metal, or one of its compounds, in it to be worth digging out! Ores are often oxides, carbonates or sulphides. They are all finite resources so we should use them wisely! In order to extract a metal, the ore or compound of the metal must undergo a process called reduction to free the metal (i.e. the positive metal ion gains negative electrons to form the neutral metal atom, or the oxide loses oxygen, to form the free metallic atoms). Generally speaking the method of extraction depends on the metals position in the reactivity series. The reactivity series of metals can be presented to include two non-metals, carbon and hydrogen, to help predict which method could be used to extract the metal. o lower Pt Au Ag Cu (H) Pb Sn Fe Zn (C) Al Mg Ca Na K higher in series o RULE: Any element higher in the series can displace any other lower element Metals above zinc and carbon in the reactivity series cannot usually be extracted with carbon or carbon monoxide. They are usually extracted by electrolysis of the purified molten ore or other suitable compound o eg aluminium from molten aluminium oxide or sodium from molten sodium chloride. 8 www.studyguide.pk The ore or compound must be molten or dissolved in a solution in an electrolysis cell to allow free movement of ions (electrical current). Metals below carbon can be extracted by heating the oxide with carbon or carbon monoxide. The non-metallic elements carbon will displace the less reactive metals in a smelter or blast furnace e.g. iron or zinc and metals lower in the series. o Metals below hydrogen will not displace hydrogen from acids. Their oxides are easily reduced to the metal by heating in a stream of hydrogen, though this is an extraction method rarely used in industry. In fact most metal oxides below carbon can be reduced when heated in hydrogen, even if the metal reacts with acid. Some metals are so unreactive that they do not readily combine with oxygen in the air or any other element present in the Earth's crust, and so can be found as the metal itself. For example gold (and sometimes copper and silver) and no chemical separation or extraction is needed. In fact all the metals below hydrogen can be found as the 'free' or 'native' element. Other methods are used in special cases using the displacement rule. A more reactive metal can be used to displace and extract a less reactive metal but these are costly processes since the more reactive metal also has to be produced in the first place! See Titanium or see at the end of the section on copper extraction Sometimes electrolysis is used to purify less reactive metals which have previously been extracted using carbon or hydrogen (eg copper and zinc). Electrolysis is also used to plate one metal with another. The demand for raw materials does have social, economic and environmental implications eg conservation of mineral resources by recycling metals, minimising pollution etc. Historically as technology and science have developed the methods of extraction have improved to the point were all metals can be produced. The reactivity is a measure of the ease of compound formation and stability (ie more reactive, more readily formed stable compound, more difficult to reduce to the metal). o The least reactive metals such as gold, silver and copper have been used for the past 10000 years because the pure metal was found naturally. o Moderately reactive metals like iron and tin have been extracted using carbon based smelting for the past 2000-3000 years. o BUT it is only in the last 200 years that very reactive metals like sodium or aluminium have been extracted by electrolysis. o 21.2 Metallurgy Metallurgy is the combination of science and technology used to extract metals from their ores. Ores are complex mixtures of metal-containing material and useless impurities called gangue. The steps involved in extracting a metal include the following: concentrating the ore, and chemically treating it if necessary reducing the mineral to free metal refining and purifying the metal. The metal may be mixed with other elements to modify its properties or to form an alloy, a metallic solution of two or more elements 9 www.studyguide.pk Concentration and Chemical Treatment of Ores The different physical properties of the mineral and the gangue, such as density and magnetic charge, can be used to concentrate the mineral and remove the gangue. Metal sulfide ores are concentrated by flotation, a process that exploits differences in the ability of water and oil to wet the surfaces of the mineral and the gangue. Mineral particles float to the top of the tank along with soapy air bubbles, while the gangue sinks to the bottom. Ores can also be concentrated by chemical means. In the Bayer process, the Al 2O3 in bauxite is separated from Fe2O3 impurities by treating the ore with NaOH. Roasting, or heating in air, is another chemical treatment used to convert minerals to compounds that are more easily reduced to the metal. 10 www.studyguide.pk Reduction The more active metals are obtained by reducing their ores with a chemical reducing agent. Zinc is obtained by reducing ZnO with coke, a form of carbon. Carbon cannot be used to reduce metals that form stable carbides, such as tungsten. Tungsten(VI) oxide is reduced with hydrogen gas. The most active metals cannot be reduced with chemical reducing agents, so these metals are produced by electrolytic reduction,. Refining The metals obtained from reducing ores generally require purification. Some metals, including zinc, can be purified by distillation. Nickel is purified using the Mond process, a chemical method in which Ni(CO)4 is formed and then decomposed at a higher temperature. The equilibrium shift at the higher temperature favors pure nickel. Extraction of Metal 11 www.studyguide.pk The Extraction of Iron Raw Materials: Iron Ore eg haematite ore [iron(III) oxide, Fe2O3] coke (carbon, C) hot air (for the O2 in it) limestone (calcium carbonate, CaCO3) The solid mixture of haematite ore, coke and limestone is continuously fed into the top of the blast furnace. The coke is ignited at the base and hot air blown in to burn the coke (carbon) to form carbon dioxide in an oxidation reaction (C gains O). The heat energy is needed from this very exothermic reaction to raise the temperature of the blast furnace to over 1000oC to effect the ore reduction. The furnace contents must be he ated. o carbon + oxygen ==> carbon dioxide o C(s) + O2(g) ==> CO2(g) at high temperature the carbon dioxide formed, reacts with more coke (carbon) to form carbon monoxide o carbon dioxide + carbon ==> carbon monoxide o CO2(g) + C(s) ==> 2CO(g) o (note: CO2 reduced by O loss, C is oxidised by O gain) The carbon monoxide is the molecule that actually removes the oxygen from the iron oxide ore. This a reduction reaction (Fe2O3 loses its O, or Fe3+ gains three electrons to form Fe) and the CO is known as the reducing agent (the O remover and gets oxidised in the process). This frees the iron, which is molten at the high blast furnace temperature, and trickles down to the base of the blast furnace. The main reduction reaction is ... o iron(III) oxide + carbon monoxide ==> iron + carbon dioxide o Fe2O3(s) + 3CO(g) ==> 2Fe(l) + 3CO2(g) o note, as in the two reactions above, oxidation and reduction always go together! Other possible ore reduction reactions are ... Fe2O3(s) + 3C(g) ==> 2Fe(l) + 3CO(g) 2Fe2O3(s) + 3C(g) ==> 4Fe(l) + 3CO2(g) The original ore contains acidic mineral impurities such as silica (SiO2, silicon dioxide). These react with the calcium carbonate (limestone) to form a molten slag of e.g. calcium silicate. o calcium carbonate + silica ==> calcium silicate + carbon dioxide o CaCO3 + SiO2 ==> CaSiO3 + CO2 o this is sometimes shown in two stages: CaCO3 ==> CaO + CO2 CaO + SiO2 ==> CaSiO3 The molten slag forms a layer above the more dense molten iron and they can be both separately, and regularly, drained away. The iron is cooled and cast into pig iron ingots OR transferred directly to a steel producing furnace. Iron from a blast furnace is ok for very hard cast iron objects BUT is too brittle for many applications due to too high a carbon content from the coke. So it is converted into steel alloys for a wide range of uses. The waste slag is used for road construction or filling in quarries which can then be landscaped or making cement. 12 www.studyguide.pk 21.3 Iron and Steel The cast iron or pig iron produced in a blast furnace must be purified. In the basic oxygen process, molten iron from the blast furnace is mixed with pure oxygen gas in a furnace lined with basic oxides. The impurities in the iron are oxidized and the acidic oxides react with CaO to yield a molten slag. Phosphorus impurities react in this process to form a calcium phosphate slag. The Extraction of Aluminium Raw materials for the electrolysis process: Bauxite ore of impure aluminium oxide The purified bauxite ore of aluminium oxide is continuously fed in. Cryolite is added to lower the melting point and dissolve the ore. Ions must be free to move to the electrode connections called the cathode (), attracting positive ions eg Al3+, and the anode (+) attracting negative ions eg O2-. When the d.c. current is passed through aluminium forms at the positive cathode (metal*) and sinks to the bottom of the tank. At the negative anode, oxygen gas is formed (non-metal). This is quite a problem. At the high temperature of the electrolysis cell it burns and oxidises away the carbon electrodes to form toxic carbon monoxide or carbon dioxide. So the electrode is regularly replaced and the waste gases dealt with! It is a costly process (6x more than Fe!) due to the large quantities of expensive electrical energy needed for the process. The redox details of the electrode processes: At the negative (-) cathode, reduction 13 www.studyguide.pk [Al2O3 made up of Al3+ and O2- ions] occurs (electron gain) when the positive Carbon (graphite) for the electrodes. aluminium ions are attracted to it. They gain three electrons to change to neutral Al atoms. Cryolite reduces the melting point of the ore and saves energy, because the ions must be free to move to carry the current Al3+ + 3e- ==> Al Electrolysis means using d.c. electrical energy to bring about chemical changes At the positive (+) anode, oxidation takes eg decomposition of a compound to place (electron loss) when the negative form metal deposits or release gases. The oxide ions are attracted to it. They lose two electrical energy splits the compound! electrons forming neutral oxygen molecules. At the electrolyte connections called the anode electrode (+, attracts - ions) and 2O2- ==> O2 + 4ethe cathode electrode (-, attracts + ions). An electrolyte is a conducting melt or Note: Reduction and Oxidation always go solution of freely moving ions which together! carry the charge of the electric current. The overall electrolytic decomposition is ... o aluminium oxide => aluminium + oxygen o 2Al2O3 ==> 4Al + 3O2 o and is a very endothermic process, lots of electrical energy input! The original extraction of copper from copper ores from copper carbonate ores* ... o The ore can be roasted to concentrate the copper as its oxide. o Water is driven off and the carbonate thermally decomposed. o copper(II) carbonate ==> copper oxide + carbon dioxide o CuCO3(s) ==> CuO(s) + CO2(g) o The oxide can be smelted by heating with carbon (coke, charcoal) to reduce the oxide to impure copper, though this method isn't really used much these days (the 'bronze age' method archaeologically!). o copper(II) oxide + carbon ==> copper + carbon dioxide o 2CuO(s) + C(s) ==> 2Cu(s) + CO2(g) from copper sulphide ores ... o copper sulphide ores can roasted in air to form impure copper o nasty sulphur dioxide gas is formed, this must be collected to avoid pollution and can be used to make sulphuric acid to help the economy of the process o copper(I) sulphide + oxygen ==> copper + sulphur dioxide o Cu2S(s) + O2(g) ==> 2Cu(s) + SO2(g) sulphur dioxide is a nasty toxic acidic gas, it is collected and used to make sulphuric acid, helps pay for the extraction process. o or *CuS(s) + O2(g) ==> Cu(s) + SO2(g) * the CuS might be part of an ore like chalcopyrite CuFeS2 which is the principle ore copper is extracted from. * It is also possible to dissolve the carbonate ore or the oxide from roasted ore in dilute sulphuric acid and extracting copper by .... o (1) using electrolysis see purification by electrolysis above. or o (2) by adding a more reactive metal to displace it eg scrap iron or steel is used by 14 www.studyguide.pk adding it to the resulting copper(II) sulphate solution. iron + copper(II) sulphate ==> iron(II) sulphate + copper Fe(s) + CuSO4(aq) => FeSO4(aq) + Cu(s) The Extraction of Titanium by Displacement Titanium ore is mainly the oxide TiO 2, which is converted into titanium tetrachloride TiCl4 The chloride is then reacted with sodium or magnesium to form titanium metal and sodium chloride or magnesium Chloride. This reaction is carried out in an atmosphere of inert argon gas so non of the metals involved becomes oxidised by atmospheric oxygen. TiCl4 + 2Mg ==> Ti + 2MgCl2 or TiCl4 + 4Na ==> Ti + 4NaCl These are examples of metal displacement reactions eg the less reactive titanium is displaced by the more reactive sodium or magnesium. Overall the titanium oxide ore is reduced to titanium metal (overall O loss, oxide => metal) Environmental Impact and Economics of Metal and other Mineral Extraction One of the problems of metal or mineral extraction is balancing ecological, environmental, economic, social advantages. It doesn't matter whether you are mining and processing iron ore or limestone, many of the advantages and disadvantages are common to these operations. Examples of advantages of a country exploiting it's own mineral resources: o Valuable revenue if the mineral or its products are exported. o Jobs for people, especially in poor countries or areas of high unemployment in developed countries. o Wages earned go into the local economy. o Increase in local facilities promoted eg transport systems, roads and recreational and health social facilities. o ? Examples of disadvantages of a country exploiting it's own mineral resources and reduction of its social and environmental impact: o Dust from mining or processing can be reduced by air filter and precipitation systems. o Noise from process operation or transport of raw materials and products. Difficult to deal with, sound-proofing often not practical, but operations can be reduced for unsociable hours eg evening and night! o Pollution can be reduced by cleaning the 'waste' or 'used' air or water of toxic or acidic materials eg carbon monoxide from the blast furnace extraction of iron 15 www.studyguide.pk o o sulphur dioxide gas from copper extraction of its sulphide ore Mining operations will disfigure the landscape BUT it can be re-claimed and 'landscaped' in an attempt to restore the original flora and fauna. ? 4. How can metals be made more useful? Extraction details Aluminium can be made more resistant to corrosion by a process called anodising. Iron can be made more useful by mixing it with other substances to make various types of steel. Many metals can be given a coating of a different metal to protect them or to improve their appearance. Aluminium is a reactive metal but it is resistant to corrosion. This is because aluminium reacts in air to form a layer of aluminium oxide which then protects the aluminium from further attack. o This is why it appears to be less reactive than its position in the reactivity series of metals would predict. For some uses of aluminium it is desirable to increase artificially the thickness of the protective oxide layer in a process is called anodising. o This involves removing the oxide layer by treating the aluminium sheet with sodium hydroxide solution. o The aluminium is then placed in dilute sulphuric acid and is made the positive electrode (anode) used in the electrolysis of the acid. o Oxygen forms on the surface of the aluminium and reacts with the aluminium metal to form a thicker protective oxide layer. Aluminium can be alloyed to make 'Duralumin' by adding copper (and smaller amounts of magnesium, silicon and iron), to make a stronger alloy used in aircraft components (low density = 'lighter'!), greenhouse and window frames (good anti-corrosion properties), overhead power lines (quite a good conductor and 'light'), but steel strands are included to make the 'line' stronger and poorly electrical conducting ceramic materials are used to insulate the wires from the pylons and the ground. The properties of iron can be altered by adding small quantities of other metals or carbon to make steel. Steels are alloys since they are mixtures of iron with other metals or with nonmetals like carbon or silicon. Making Steel: o (1) Molten iron from the blast furnace is mixed with recycled scrap iron o (2) Then pure oxygen is passed into the mixture and the non-metal impurities such as silicon or phosphorus are then converted into acidic oxides (oxidation process) .. eg Si + O2 ==> SiO2, or 4P + 5O2 ==> P4O10 o (3) Calcium carbonate (a base) is then added to remove the acidic oxide impurities (in an acid-base reaction). The salts produced by this reaction form a slag which can be tapped off separately. eg CaCO3 + SiO2 ==> CaSiO3 + CO2 (calcium silicate slag) o Reactions (1)-(3) produce pure iron. o Calculated quantities of carbon and/or other metallic elements such as titanium, manganese or chromium are then added to make a wide range of steels with particular properties. o Because of the high temperatures the mixture is stirred by bubbling in unreactive argon gas! o Economics of recycling scrap steel or ion: Most steel consists of >25% recycled 16 www.studyguide.pk iron/steel and you do have the 'scrap' collection costs and problems with varying steel composition* BUT you save enormously because there is no mining cost or overseas transport costs AND less junk lying around! (NOTE: * some companies send their own scrap to be mixed with the next batch of 'specialised' steel they order, this saves both companies money!) Different steels for different uses: o High % carbon steel is strong but brittle. o Low carbon steel or mild steel is softer and is easily shaped and pressed eg into a motor car body. o Stainless steel alloys contain chromium and nickel and are tougher and more resistant to corrosion. o Very strong steels can be made by alloying the iron with titanium or manganese metal. Steel can be galvanised by coating in zinc, this is physically done by dipping the object into a bath of molten zinc. On removal and cooling a thin layer of zinc is left on. The zinc chemically bonds to the iron via the free electrons of both metals - its all the same atoms to them! It can also be done by electroplating. Steel (and most metals) can be electroplated. o The steel object to be plated is made the negative electrode (cathode) and placed in a solution containing ions of the plating metal. o The positive electrode (anode) is made of the pure plating metal (which dissolves and forms the fresh deposit on the negative electrode). o Nickel, zinc, copper, silver and gold are examples of plating metals. o If M = Ni, Cu, Zn .... At the positive (+) anode, the process is an oxidation, electron loss, as the metal atoms dissolve to form metal(II) ions. M(s) ==> M2+(aq) + 2e at the negative (-) cathode, the process is a reduction, electron gain by the attracted metal(II) ions to form neutral metal atoms. M2+(aq) + 2e- ==> M(s) For silver plating it is Ag+, Ag and a single electron change Any conducting (usually metal) object can be electroplated with copper or silver for aesthetic reasons or steel with zinc or chromium as anticorrosion protective layer. Many other metals have countless uses eg zinc o zinc is used to make the outer casing of zinc-carbon-weak acid batteries. o It is alloyed with copper to make the useful metal brass (electrical plug pins). Brass alloy is stronger and more hardwearing than copper AND not as brittle as zinc. 17 www.studyguide.pk USES OF METALS METAL Aluminium USES a) Structural material for ships, planes, cars, saucepans. b) Overhead electricity cables a) Coating iron to give galvanized iron Zinc b) To make alloys e.g brass (Zn/Cu) and bronze ( Zn/Sn/Cu). Iron Structural amterial for all industries ( in the form of steel) a) Car batteries. Lead b) Solder (Pb/Sn) alloys Copper a) Electric cables b) Pipes c) Alloys d) Coins (Cu/Ni) Tin Coating steel cans or tins. Nickle Electroplating steel PROPERTIES THAT MAKE IT SUITABLE a) strong but light; oxide layer prevents corrosion. b) light but good conductor a) Reactive- gives acrificial protection to iron; does not corrode easily. b) Modifies the properties of other elements. Strong and cheap; properties can be made suitable by alloying. a) Design of battery makes recharging possible. b) low melting point. a) very good conductors b)Very ductile, does not corrode easily c) d) A traditional metal for coins Un reactive and non- toxic. Protevts the steel from rusting Resist corrosion, shiny and attractive to look at. 18 www.studyguide.pk Note on Alloy Structure 1. Shows the regular arrangement of 3. the atoms in a metal crystal and the white spaces show where the free electrons are (yellow circles actually positive metal ions). 2. Shows what happens when the metal is stressed by a strong force. The layers of atoms can slide over each other and the bonding is maintained as the mobile electrons keep in contact with atoms, so the metal remains intact BUT a different shape. Shows an alloy mixture. It is NOT a compound but a physical mixing of a metal plus at least one other material (shown by red circle, it can be another metal eg Ni, a non-metal eg C or a compound of carbon or manganese, and it can be bigger or smaller than iron atoms). Many alloys are produced to give a stronger metal. The presence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The result is a stronger harder less malleable metal. ALLOY Stain less steel Cupronickle Manganese steel Brass Bronze Magnalium Solder COMPOSITION % Fe = 74% Cr = 18 % Ni = 8 % Cu = 75% Ni = 25% Fe = 85 % Mn = 13.8 % C = 1.2 Cu = 70% Zn = 30 % Cu = 90 % Sn = 10 % Al = 90 % Mg = 10 % Pb = 50 % Sn = 50 % SPECIAL PROPERTIES USES Resist corrosion Car parts, kitchen sinks, cutlery Hard wearing, attractive silver color Silver coins Very hard Springs Harder then Copper, does not corrode Harder then brass, does not corrode. Musical instruments, taps Light but strong. Aeroplanes bodies Low mwlting point but form a strong solid Joining wires and pipes. Statues, ornaments. 19 www.studyguide.pk METAL CORROSION and the RUSTING of IRON Iron (or steel) corrodes more quickly than most other transition metals and readily does so in the presence of both oxygen (in air) and water to form an iron oxide. You can do simple experiments to show that BOTH oxygen and water are needed. Put an iron nail into (1) boiled water in a sealed tube; (2) a tube of air and a drying agent; (3) an open test tube with water. Rusting appears overnight with (3) only. Rusting is speeded up in the presence of salt or acid solutions because of an increased concentration of ions. Corrosion is a redox process involving redox electron transfer and ion movement. The rusting metal behaves like a simple cell and more ions enable the current, and hence the electron transfer, to occur more readily. Rusting is overall ... Fe(s) + O2(g) + H2O(l) ==> Fe2O3.xH2O(s) ie rust is a hydrated iron(III) oxide (the equation is not meant to be balanced and the amount of water x is variable, from dry to soggy!). o The reaction proceeds via iron(II) hydroxide Fe(OH)2 which is the oxidised further to the FeO3 o Rusting is an oxidation because it involves iron gaining oxygen (Fe ==> Fe2O3) or iron atoms losing electrons (Fe - 3e- ==> Fe3+. o See more examples of oxidation and reduction below. The rusting of iron is a major problem in its use as a structural material. Iron and steel (alloy of iron) are most easily protected by paint which provides a barrier between the metal and air/water. Moving parts on machines can be protected by a water repellent oil or grease layer. This 'rusting' corrosion can be prevented by connecting iron to a more reactive metal (e.g. zinc or magnesium). This is referred to as sacrificial protection or sacrificial corrosion, because the more reactive protecting metal is preferentially oxidised away, leaving the protected metal intact. The picture illustrates what might be seen after a few days.* Iron or steel can also be protected by mixing in other metals (e.g. chromium) to make non-rusting alloys called stainless steel. The chromium, like aluminium, forms a protective oxide layer. * Theoretically, any iron ions formed by oxidation would be reduced by electrons from the oxidation of the more reactive 'sacrificed' metal. Coating iron or steel with a thin zinc layer is called 'galvanising'. The layer is produced by electrolytic deposition by making the iron/steel the negative cathode or by dipping the iron/steel object in molten zinc (more details). The zinc preferentially corrodes or oxidises to form a zinc oxide layer that doesn't flake off like iron oxide rust does. Also, if the surface is scratched, the exposed zinc again corrodes before the iron and continues to protect it. Steel tin cans are protected by relatively unreacted tin and works well as long as the thin tin layer is complete. HOWEVER, if a less reactive metal is connected to the iron, it then the iron rusts preferentially (try scratching a 'tin' can and leave out in the rain and note the corrosion by the scratch!) 20 www.studyguide.pk Methods of Prevention of Rusting of Iron Covering with Paint Covering with Grease or Oil Covering with Chromium ( Chrome Plating) Covering with Tin ( Tin plating) Covering with Zinc Metal ( Galvanising) Using Blocks of Zinc Metal Making Stainless Steel Using Bocks of Magnesium Metal \Aluminium does not oxidise (corrode) as quickly as its reactivity would suggest. Once a thin oxide layer of Al2O3 has formed on the surface, it forms a barrier to oxygen and water and so prevents further corrosion of the aluminium. Aluminium is a useful structural metal. It can be made harder, stronger and stiffer by mixing it with small amounts of other metals (e.g magnesium) to make alloys. Copper and Lead are both used in roofing situations because neither is very reactive and the compounds formed do not flake away as easily as rust does from iron. Lead corrodes to a white lead oxide or carbonate and copper corrodes to form a basic green carbonate (combination of the hydroxide Cu(OH)2 and carbonate CuCO3 eg seen as green roof on buildings). Both metals have been used for piping but these days lead is considered too toxic and copper is usually used as the stronger, but equally unreactive alloy with zinc, brass. Now of course, most piping is flowing in the plastic direction which doesn't corrode at all! The Group 1 Alkali Metals rapidly corrode in air and need to be stored under oil. 21 www.studyguide.pk Apart from their structural weakness they would hardly used for any outside purpose! DONE 22 www.studyguide.pk The Reactivity Series and Corrosion of Metals The higher the metal in the series, the more reactive it is ie the faster, more vigorous and more exothermic the reaction. At a more theoretical level, the more reactive a metal, the greater tendency it has to form its positive ion. This also implies that the reverse reaction becomes more difficult ie the more reactive a metal, the more difficult it is to extract from its ore and the metal is also more susceptible to corrosion with oxygen and water. The reactivity series can be established by observation of the reaction of metals with water, oxygen or acids (and also from simple cell experiments). A metal in the series, can displace any metal below it in the series, from the less reactive metal's oxide, chloride or sulphate compound. o eg on heating the mixture of magnesium powder and black copper(II) oxide, white magnesium oxide is formed with brown bits of copper: Mg(s) + CuO(s) => MgO(s) + Cu(s) o or adding a metal to a salt solution eg adding magnesium to blue copper(II) sulphate solution, the blue colour fades as colourless magnesium sulphate is formed and brown bits of copper metal form a precipitate: Mg(s) + CuSO4(aq) => MgSO4(aq) + Cu(s) The electron transfer redox theory behind displacemet reactions is explained later. Some general word equations where the metal does react: o (a) metal + cold water => metal hydroxide + hydrogen (metals above aluminium) o (b) heated metal + steam => metal oxide + hydrogen (for metals above tin?) o (c) metal + acid => metal salt + hydrogen if the metal is at least as reactive as lead (see reactivity series list above left) and hydrochloric acid makes a metal chloride and sulphuric acid makes a metal sulphate reactions with nitric acid are complex, the nitrate is formed BUT the gas is rarely hydrogen, and more often an oxide of nitrogen (not usually studied at GCSE level these days). Within the general Reactivity or Activity Series of Metals there are some Periodic Table Trends … o Down Group 1 (I) the "Alkali Metals" the activity increases Cs > Rb > K > Na > Li o Down Group 2 (II) the activity increases eg Ca > Mg o On the same period, the Group 1 metal is more reactive than the group 2 metal, and the group 2 metal is more reactive than the Group 3 metal, and all three are more reactive than the "Transition Metals". eg Na > Mg > Al (on Period 3) and K > Ca > Ga > Fe/Cu/Zn etc. (on Period 4) The reactivity of a metal has an important bearing on the method by which a metal is extracted from its ore. Since prehistoric times, as technology has improved more and more, all metals can now be extracted and comments on when the metals were first isolated and used are added in the table below. If the metal is above carbon, it cannot be o www.studyguide.pk extracted by carbon reduction and must be usually extracted by electrolysis. Two non-metals, carbon and hydrogen, are included in the table for comparison, and are important chemical reference points concerning the method of metal extraction and reactivity towards acids o Metals above carbon cannot usually be extracted by carbon or carbon monoxide reduction and are usually extracted by electrolysis o Metals below carbon in the series can be extracted by heating the oxide with carbon or carbon monoxide. o Metals below hydrogen will not displace hydrogen from acids and can be extracted by heating the oxide in hydrogen. METAL in decreasing reactivity order Caesium Cs Rubidium Rb Potassium K Reactivity and Reactions Burns vigorously with a blue flame when heated in air to form the white powder caesium oxide. 4Cs(s) + O2(g) ==> 2Cs2O(s) Because it is extremely reactive, it explodes with cold water forming the alkali caesium hydroxide and hydrogen gas. 2Cs(s) + 2H2O(l) => 2CsOH(aq) + H2(g) Caesium was first extracted in 1860 by electrolysis of the molten chloride CsCl. Burns vigorously with a red flame when heated in air to form the white powder rubidium oxide. 4Rb(s) + O2(g) ==> 2Rb2O(s) Extremely reactive, can ignite in air, it explodes with cold water forming the alkali rubidium hydroxide and hydrogen gas. 2Rb(s) + 2H2O(l) => 2RbOH(aq) + H2(g) Rubidium was first extracted in 1861 by electrolysis of the molten chloride RbCl. Burns vigorously with a lilac flame when heated in air to form the white powder potassium oxide. 4K(s) + O2(g) ==> 2K2O(s) The reaction of potassium with water - the reaction is the same as for sodium (full description below) BUT it is faster and more exothermic AND so the hydrogen is ignited to give a purple or lilac flame. The hydrogen flame is coloured by the excitation of potassium atoms in the very hot flame. The very rapid reaction with cold water forms the alkali potassium hydroxide and hydrogen gas. 2K(s) + 2H2O(l) => 2KOH(aq) + H2(g) Potassium was first extracted in 1807 by electrolysis of the molten chloride KCl. [top] 2 www.studyguide.pk Sodium Na Lithium Li Calcium Ca Burns vigorously with a yellow flame when heated in air to form the white powder sodium oxide. 4Na(s) + O2(g) ==> 2Na2O(s) (also forms Na2O2) The reaction of sodium with water: the sodium melts to a silvery ball and fizzes as it spins over the water. The rapid exothermic reaction produces a colourless gas which gives a squeaky pop! with a lit splint (hydrogen). Universal indicator will turn from green to purple/violet as the strong alkali sodium hydroxide is formed. The initially sodium floats because it is less dense than water. 2Na(s) + 2H2O(l) => 2NaOH(aq) + H2(g) Sodium was first extracted in 1807 by electrolysis of the molten chloride NaCl Burns vigorously with a red flame when heated in air to form the white powder lithium oxide. 4Li(s) + O2(g) ==> 2Li2O(s) Quite a fast reaction with cold water forming the alkali lithium hydroxide and hydrogen gas. For full description see sodium above, but the reaction is not as fast. 2Li(s) + 2H2O(l) => 2LiOH(aq) + H2(g) Lithium was first extracted in 1821 by electrolysis of the molten chloride LiCl. Burns with a brick red flame when strongly heated in air to form the white powder calcium oxide. 2Ca(s) + O2(g) ==> 2CaO(s) Quite reactive with cold water forming the moderately soluble alkali calcium hydroxide and hydrogen gas. Ca(s) + 2H2O(l) => Ca(OH)2(aq/s) + H2(g) Very reactive with dilute hydrochloric acid forming the colourless soluble salt calcium chloride and hydrogen gas. Ca(s) + 2HCl(aq) => CaCl2(aq) + H2(g) Not very reactive with dilute sulphuric acid because the colourless calcium sulphate formed is not very soluble and coats the metal inhibiting the reaction. Ca(s) + H2SO4(aq) => CaSO4(s) + H2(g) Calcium was first extracted in 1808 by electrolysis of the molten chloride CaCl2. 3 www.studyguide.pk Magnesium Mg Aluminium Al (Carbon C, a nonmetal) Burns with a bright white flame when strongly heated in air to form a white powder of magnesium oxide. 2Mg(s) + O2(g) ==> 2MgO(s) Slow reaction with water forming the slightly soluble alkali magnesium hydroxide and hydrogen gas. Mg(s) + 2H2O(l) => Mg(OH)2(aq/s) + H2(g) With steam, the reaction is faster with heated magnesium and the white powder magnesium oxide is formed with the hydrogen. Magnesium will burn with a bright white flame in steam, if previously ignited in air! Mg(s) + H2O(g) => MgO(s) + H2(g) In fact it will even burn in carbon dioxide forming black specks of carbon! 2Mg(s) + CO2(g) ==> 2MgO(s) + C(s) Very reactive with dilute hydrochloric acid forming the colourless soluble salt magnesium chloride and hydrogen gas. Mg(s) + 2HCl(aq) => MgCl2(aq) + H2(g) Very reactive with dilute sulphuric acid forming colourless soluble magnesium sulphate and hydrogen. Mg(s) + H2SO4(aq) => CaSO4(s) + H2(g) Magnesium was first extracted in 1808 by electrolysis of the molten chloride MgCl2. Surface goes white when strongly heated in air to form aluminium oxide. Theoretically its quite a reactive metal but the oxide layer as an inhibiting effect. This is why it appears to be less reactive than its position in the reactivity series of metals would predict. o 4Al(s) + 3O2(g) ==> 2Al2O3(s) Aluminium has no reaction with water or steam due to a protective aluminium oxide layer of Al 2O3. [note: Chromium behaves chemically in the same way, forming a protective layer of chromium(III) oxide, Cr2O3, and so its anti-corrosion properties are used in stainless steels and chromium plating] The thermit reaction: However the true reactivity of aluminium can be spectacularly seen when its grey powder is mixed with brown iron(III) oxide powder. When the mixture is ignited with a magnesium fuse (high activation energy!), it burns very exothermically in a shower of sparks to leave a red hot blob of iron and white aluminium oxide powder. o aluminium + iron(III) oxide ==> iron + aluminium oxide o 2Al(s) + Fe2O3(s) ==> Al2O3(s) + 2Fe(s) Slow reaction with dilute hydrochloric acid to form the colourless soluble salt aluminium chloride and hydrogen gas. o 2Al(s) + 6HCl(aq) => 2AlCl3(aq) + 3H2(g) The reaction with dilute sulphuric acid is extremely slow to form colourless aluminium sulphate and hydrogen. o 2Al(s) + 3H2SO4(aq) => Al2(SO4)3(aq) + 3H2(g) Aluminium was first extracted in 1825 by electrolysis of its molten oxide Al2O3. Elements higher than carbon ie aluminium or more reactive, must be extracted by electrolysis (or displacing it with an even more reactive metal). Metals below it, ie zinc or a less reactive, can be extracted by reducing the hot metal oxide with carbon. [top] 4 www.studyguide.pk Zinc Zn Iron Fe Surface goes white when strongly heated in air to form zinc oxide (yellow when hot). 2Zn(s) + O2(g) ==> 2ZnO(s) No reaction with cold water. When the metal is heated in steam zinc oxide and hydrogen are formed. Zn(s) + H2O(g) => ZnO(s) + H2(g) Quite reactive with dilute hydrochloric acid forming the colourless soluble salt zinc chloride and hydrogen gas. Zn(s) + 2HCl(aq) => ZnCl2(aq) + H2(g) Quite reactive with dilute sulphuric acid forming the colourless soluble salt zinc sulphate and hydrogen gas. Zn(s) + H2SO4(g) => ZnSO4(s) + H2(g) (this reaction is catalysed by adding a trace of copper sulphate solution) Zinc can be extracted by reducing the hot metal oxide on heating with carbon 2ZnO(s) + C(s) => 2Zn(s) + CO2(g) A zinc coating (galvanising) is used to protect iron from rusting. The more reactive zinc oxidises 1st. Blocks of zinc attached to steel are used as 'sacrificial corrosion'. Zinc was known and used in India and China before 1500 so it must have been extracted like copper or iron by carbon reduction of the oxide, sulphide or carbonate. [top] Surface goes dark grey-black when strongly heated in air to form a tri-iron tetroxide. When steel wool is heated in a bunsen flame it burns with a shower of sparks - large surface area - increased rate of reaction - so even moderately reactive iron has its moments! 3Fe(s) + 2O2(g) ==> Fe3O4(s) No reaction with cold water (rusting is a joint reaction with oxygen). When the metal is heated in steam an iron oxide (unusual formula) and hydrogen are formed. This oxide is 'technically' diiron(III)iron(II) oxide!!!! 3Fe(s) + 4H2O(g) => Fe3O4(s) + 4H2(g) Slow reaction with dilute hydrochloric acid forming the soluble pale green salt iron(II) chloride and hydrogen gas. Fe(s) + 2HCl(aq) => FeCl2(aq) + H2(g) Slow reaction with dilute sulphuric acid forming the soluble pale green salt iron(II) sulphate and hydrogen gas. Fe(s) + H2SO4(g) => FeSO4(s) + H2(g) Iron can be extracted by reducing the hot metal oxide on heating with carbon monoxide formed from carbon in the blast furnace eg Fe2O3(s) + 3CO(g) => 2Fe(s) + 3CO2(g) Fe3O4(s) + 4CO(g) => 3Fe(s) + 4CO2(g) For the past 2500 years. iron has been extracted from pre-historic times using charcoal (C). Known in Anglo-Saxon as 'iron' and in Roman times in Latin as 'ferrum' hence the Fe symbol! [top] 5 www.studyguide.pk Tin Sn Lead Pb Hydrogen H non-metal Copper Cu Slow reaction when heated in air to form white tin(IV) oxide or tin dioxide Sn(s) + O2(g) ==> SnO2(s) No reaction with cold water or when heated in steam. Very slow reaction with dilute hydrochloric acid forming the slightly soluble tin(II) chloride and hydrogen gas. Sn(s) + 2HCl(aq) => SnCl2(aq) + H2(g) Very slow reaction with dilute sulphuric acid forming the colourless slightly soluble tin(II) sulphate and hydrogen gas. Sn(s) + H2SO4(g) => SnSO4(s) + H2(g) Tin can be extracted from its oxide by heating with carbon. Tin has been known from pre-historic times. Known in Anglo-Saxon as 'tin' and in Latin - 'stannum' hence the symbol Sn! [top] Slow reaction when heated in air to form red/yellow lead(II) oxide and tri-lead tetroxide 2Pb(s) + O2(g) ==> 2PbO(s) and 3Pb(s) + 2O2(g) ==> Pb3O4(s) No reaction with cold water or when heated in steam. Very slow and effectively no reaction with dilute hydrochloric acid or dilute sulphuric acid. Lead can be extracted from its oxide by heating with carbon. Probably used from pre-historic times and known in Anglo-Saxon as 'lead' and in Latin 'plumbum' hence the symbol Pb! [top] Non of the metals below hydrogen can react with acids to form hydrogen gas. They are least easily corroded metals and partly accounts for their value and uses in jewellery, electrical contacts etc.[top] Surface blackens when strongly heated in air to form copper(II) oxide. 2Cu(s) + O2(g) ==> 2CuO(s) No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Copper can be extracted by reducing the hot black metal oxide on heating with carbon 2CuO(s) + C(s) => 2Cu(s) + CO2(g) The elemental metal can be found 'native' and was probably first used over 6000 years ago in Turkey by literally beating it out of rocks and into shape (malleable at room temperature!) - no high temperature technology used or available. It has been extracted by carbon reduction of a copper mineral for at least 3000 years. Latin 'cuprum' meaning Cyprus?, anyway that's why its symbol is Cu! [top] 6 www.studyguide.pk No reaction when heated in air. Silver No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Silver can be extracted by BUT can be found 'native' as the element because it is so unreactive. It has been used from prehistoric times in jewellery for 4000 years at least. In Anglo-Saxon it was 'siolfur' meaning 'silver in nature' and in Latin 'argentum' hence its symbol Ag. [top] No reaction when heated in air No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. Gold can be readily extracted from its ores easily by reduction BUT it is usually found 'native' as the element because it is so unreactive and has been used from pre-historic times in jewellery for at least 6000 years. Known in Anglo-Saxon as 'gold'. Gold is rather a soft metal and is 'hardened' by alloying with other metals - pure gold is 24 carat - 22, 18, 15, 12 and 9 carat gold are legalised, meaning it has that carat number/24 as parts of gold as a measure of its purity and value! [top] No reaction when heated in air. No reaction with cold water or when heated in steam. No reaction with dilute hydrochloric acid or dilute sulphuric acid. It seems ironic that despite its apparent lack of 'reactivity' it is a very potent catalyst eg catalytic converter in cars. Spanish 'platina' meant 'silvery in nature'. Like gold, it is a very rare metal but was known by pre-Columbian South American Indians and brought to Europe in about 1750. It is used in expensive jewellery, laboratory ware (eg inert crucible container) and catalytic converters in car exhausts. [top] Ag Gold Au Platinum Pt METAL CORROSION and the RUSTING of IRON Iron (or steel) corrodes more quickly than most other transition metals and readily does so in the presence of both oxygen (in air) and water to form an iron oxide. You can do simple experiments to show that BOTH oxygen and water are needed. Put an iron nail into (1) boiled water in a sealed tube; (2) a tube of air and a drying agent; (3) an open test tube with water. Rusting appears overnight with (3) only. Rusting is speeded up in the presence of salt or acid solutions because of an increased concentration of ions. Corrosion is a redox process involving redox electron transfer and ion movement. The rusting metal behaves like a simple cell and more ions enable the current, and hence the electron transfer, to occur more readily. Rusting is overall ... Fe(s) + O2(g) + H2O(l) ==> Fe2O3.xH2O(s) ie rust is a hydrated iron(III) oxide (the equation is not meant to be balanced and the amount of water x is variable, from dry to soggy!). o The reaction proceeds via iron(II) hydroxide Fe(OH)2 which is the oxidised further to the 7 www.studyguide.pk Fe(OH)3 Rusting is an oxidation because it involves iron gaining oxygen (Fe ==> Fe2O3) or iron atoms losing electrons (Fe - 3e- ==> Fe3+. o See more examples of oxidation and reduction below. The rusting of iron is a major problem in its use as a structural material. Iron and steel (alloy of iron) are most easily protected by paint which provides a barrier between the metal and air/water. Moving parts on machines can be protected by a water repellent oil or grease layer. This 'rusting' corrosion can be prevented by connecting iron to a more reactive metal (e.g. zinc or magnesium). This is referred to as sacrificial protection or sacrificial corrosion, because the more reactive protecting metal is preferentially oxidised away, leaving the protected metal intact. The picture illustrates what might be seen after a few days.* Iron or steel can also be protected by mixing in other metals (e.g. chromium) to make non-rusting alloys called stainless steel. The chromium, like aluminium, forms a protective oxide layer. * Theoretically, any iron ions formed by oxidation would be reduced by electrons from the oxidation of the more reactive 'sacrificed' metal. Coating iron or steel with a thin zinc layer is called 'galvanising'. The layer is produced by electrolytic deposition by making the iron/steel the negative cathode or by dipping the iron/steel object in molten zinc (more details). The zinc preferentially corrodes or oxidises to form a zinc oxide layer that doesn't flake off like iron oxide rust does. Also, if the surface is scratched, the exposed zinc again corrodes before the iron and continues to protect it. Steel tin cans are protected by relatively unreacted tin and works well as long as the thin tin layer is complete. HOWEVER, if a less reactive metal is connected to the iron, it then the iron rusts preferentially (try scratching a 'tin' can and leave out in the rain and note the corrosion by the scratch!) o Aluminium does not oxidise (corrode) as quickly as its reactivity would suggest. Once a thin oxide layer of Al2O3 has formed on the surface, it forms a barrier to oxygen and water and so prevents further corrosion of the aluminium. Aluminium is a useful structural metal. It can be made harder, stronger and stiffer by mixing it with small amounts of other metals (e.g magnesium) to make alloys. Copper and Lead are both used in roofing situations because neither is very reactive and the compounds formed do not flake away as easily as rust does from iron. Lead corrodes to a white lead oxide or carbonate and copper corrodes to form a basic green carbonate (combination of the hydroxide Cu(OH)2 and carbonate CuCO3 eg seen as green roof on buildings). Both metals have been used for piping but these days lead is considered too toxic and copper is usually used as the stronger, but equally unreactive alloy with zinc, brass. Now of course, most piping is flowing in the plastic direction which doesn't corrode at all! The Group 1 Alkali Metals rapidly corrode in air and need to be stored under oil. Apart from their structural weakness they would hardly used for any outside purpose! DONE 8 www.studyguide.pk The Mole Concept Counting by weighing The size of molecule is so small that it is physically difficult if not impossible to directly count out molecules. this problem is solved using a common trick. Atoms and molecules are counted indirectly by weighing. Here is a practical example. You need to estimate the number of nails in a box. You weigh an empty box, 213. g. The weight of the box plus nails is 1340. g. The weight of one nail is 0.450 g. I hope you aren't going to tear open the package and count the nails. We agree that mass of nails = 1340 g - 213 g = 1227. g Number of nails = (1227. grams nails)(1 nail/ 0.450 grams ) = 2726.6 nails = 2730 nails You can count the nails by weighing them. Avogadro's Number and the Mole To calculate real chemical reactions in the laboratory, chemists use a special unit called a mole (abbreviated mol). One mole of a substance is the amount of the substance that is equal in molar mass of the molecular or formula mass of the substance in grams. Thus for ethylene, C 2H4, its molecular weight is 28 and its molar mass is 28 g. In other words, 28 g represents 1 mol of ethylene. One mole contains 6.022 x1023 molecules or formula units. The number 6.022 x 1023 is called Avogadro's number. For ethylene, C2H4, Molecular mass C2H4 = 28. Molar mass C2H4 = 28.0 g 1 mol C2H4 contains 6.022 x 1023 molecules The coefficients of a balanced chemical equation indicate the number of moles of each substance in the reaction. Thus, at the level of moles: C3H8 + 5O2 -------- 3CO2 + 4H2O One mole of propane reacts with five moles of oxygen to form three moles of carbon dioxide and four moles of water. Avogadro's number is an accident of nature. It is the number of particles that delivers a mole of a substance. Avogadro's number = 6.022 x 1023. The reason why the value is an accident of nature is that the mole is tied to the gram mass unit. The gram is a convenient mass unit because it matches human sizes. If we were a thousand times greater in size ( like Paul Bunyan) we would find it handy to use kilogram amounts. This means the kilogram mole would be convenient. The number of particles handled in a kilogram mole is 1000 www.studyguide.pk times greater. The kilo Avogadro number for the count of particles in a kilomole is 6.022 x 1026. If humans were tiny creatures (like Lilliputians) only 1/1000 our present size, milligrams would be more convenient. This means the milligram mole would be more useful. The number of particles handled in a milligram mole (millimole) would be 1/1000 times smaller. The milli Avogadro number for the count of particles in a millimole is 6.022 x 1020. Molecular Mass and Mass Percent Activity It is very helpful to think about chemical reactions in molecular terms. However, it is not practical to carry out reactions at the molecular level. 1 C2H4 molecule + 1 HCl molecule 1 C2H5Cl molecule Because it is not practical to count individual molecules, chemists use ratios of the masses of molecules to carry out reactions. Mass ratios are determined by using the molecular masses of the substances involved in a reaction. Molecular mass provides the mass ratio we need for carrying out reactions. The mass ratio of one HCl molecule to one ethylene molecule is 36.5 to 28 in the following equation. C2H4 + HCl ------------- C2H5Cl More useful, however, is the fact that the mass ratio in grams is also 36.5 to 28.0. If we were to combine 36.5 g HCl with 28 g ethylene in the laboratory, they would react in a 1:1 molecular ratio. Molar mass for elements You are able to read the periodic table and determine the average atomic mass for an element like carbon. The average mass is 12. This is a ridiculously tiny number of grams. It is too small to handle normally. The molar mass of carbon is defined as the mass in grams that is numerically equal to the average atomic weight. This means 1 mole carbon = 12.0 grams carbon. This is the mass of carbon that contains 6.022 x 1023 carbon atoms. Avogadro's number is 6.022 x 1023 particles. This same process gives us the molar mass of any element. 1 mole neon = 20 grams neon, Ne 1 mole sodium = 23 grams sodium, Na 2 www.studyguide.pk Molar Mass for Compound The formulas for compounds are familiar to you. You know the formula for water is H2O. It should be reasonable that the weight of a formula unit can be calculated by adding up the weights for the atoms in the formula. The formula weight for water = weight from hydrogen + weight from oxygen The formula weight for water = 2 H atoms x 1. + 1 O atom x 16. = 18. The molar mass for water = 18. grams water. Example 1. What is the molar mass for sulfur dioxide, SO2 (g), a gas used in bleaching and disinfection processes. 1. Look up the atomic weight for each of the elements in the formula. 1 sulfur atom = 32 1 oxygen atom = 16 2. Count the atoms of each element in the formula unit. . one sulfur atom ; two oxygen atoms 3. The formula weight = weight from sulfur + weight from oxygen 4. The formula weight = 1 sulfur atom x (32 ) + 2 oxygen atoms x (16 ) 5. The formula weight SO2 = 32 + 32. = 64 6. The molar mass SO2 = 64 grams SO2 Example 2: The formula for methane the major component in natural gas is CH4. The formula weight for methane = weight from hydrogen + weight from carbon The formula weight for methane = 4 H atoms x 1. + 1 C atom x 12. = 16. The molar mass for methane = 16.0 grams methane Example 3: The formula for ethyl chloride is CH3 CH2Cl. The formula weight = weight from hydrogen + weight from carbon + weight from chlorine The formula weight = 5 H atoms x 1.0 + 2 C atom x 12.0 + 1 Cl atom x 35.5 = 64.5 The molar mass for ethyl chloride = 64.5 grams Mole, Molar Mass and Mass Conversions Example. How many grams of hydrogen are needed to give 3. moles of hydrogen? 1. Calculate the molar mass for hydrogen. Look up the atomic weight/mass in the periodic table The molar mass for hydrogen is 1 mole H2 = 2 grams H2 3 www.studyguide.pk 2. determine the mass needed to provide 3. moles of hydrogen. 1 mole H2 = 2.0 grams H2; 2 mole H2 = 4.0 grams H2 3 mole H2 = 6.0 grams H2 The practical way is to multiply the molar mass by the number of moles. This converts mole to grams (3 mole H2 )(2 grams H2/ 1 mole H2 ) = 6 grams H2 Example. How many moles of water are in a liter of water? Assume 1 liter = 1 kilogram water 1. Calculate the formula weight(mass) for water, H2O. Look up the atomic weights in the periodic table for H and O. The atomic weight for hydrogen is 1 The atomic weight for oxygen is 16 2. Add up the masses from all the atoms in the formula The formula weight for water is 1 + 1 + 16 = 18 3. Determine the molar mass for water. Molar mass is a mass in grams that is numerically the same as the formula weight. 1 mole H2O = 18.0 grams H2O 4. Convert 1000 grams of water to moles. The "conversion factor" is the molar mass. (1000 grams H2O )(1 mole H2O/ 18. grams H2O) = 55.55 moles H2O Example. How many moles of sulfur dioxide, SO2 (g), are in 2000 grams of the gas? 1. Look up the atomic weights in the periodic table for S and O. The atomic weight for sulphur is 32 The atomic weight for oxygen is 16 2. Calculate the formula weight for SO2 . Add up the masses from all the atoms in the formula The formula weight for sulfur dioxide is 32 S + 2 x (16 O) = 64 SO2 3. Determine the molar mass. Molar mass is a mass in grams that is numerically the same as the formula weight. 1 mole SO2 = 64. grams SO2 4 www.studyguide.pk 4. Convert 2000 grams of SO2 to moles. The "conversion factor" is the molar mass. (2000 grams SO2 )(1 mole SO2 /64. grams SO2 ) = 31.25 moles SO2, Chemical Equations and Mole Relationships The balanced equations are a chemists recipe for producing a product from reactants. The equation tell us the amounts of reactants needed and the amount of product formed. Balanced equations can be viewed at three levels. The first is the molecular level. The second is the mole level. The third is in terms of masses. We will look at mole relationships here. These interpretations of chemical equations are of value because they enable us to make predictions about the outcome of reactions. Example: Burning carbon and carbon containing compounds in air can produce carbon monoxide. Carbon monoxide is poisonous. It is cumulative and even if it doesn't kill it can cause chronic illness and brain damage. 2 C(s) + O2(g) -----> 2 CO(g) This equation can be viewed in terms of the atomic and molecular level. Two atoms of carbon must react with one molecule of oxygen. Two molecules of carbon monoxide are produced. The coefficients in the balanced equation tell the moles of each substance involved in the equation The mole ratios for this equation are C : O : CO 2moles : 1 mole : 2moles The reaction between nitrogen and oxygen to produce nitrogen dioxide is analyzed here. The equation is N2(g) + 2 O2(g)---> 2 NO2(g) 5 www.studyguide.pk Exercise: What is the mol ratio for nitrogen to oxygen? Answer: 1 mole N2 : 2 moles O2 Stoichiometry: Chemical Arithmetic In the laboratory, it is often necessary to convert between moles and mass of a substance. This relationships is called stoichiometry. Example: How many grams of carbon are required to react completely with 100 g Fe2O3? 6 www.studyguide.pk Fe2O3(s) + 3C(s) 2Fe(l) + 3CO(g) Step 1: Write a balanced chemical equation (or check to see that a given equation is balanced). In this case, a balanced chemical equation was given. Organize the information in the problem. It's often helpful to write the amounts given underneath the balanced chemical equation. Fe2O3(s) + 3C(s) ------------2Fe(l) + 3CO(g) 100 g ?g Step 2: Convert grams of a given substance to moles. Remember that substances react in terms of their mole ratios, not their mass ratios. To convert grams of Fe2O3 to moles, we need to know the molar mass of this compound. 2 x 56. g for each mol Fe + 3 x 16.0 g for each mol O = 160 g/mol Step 3: Use coefficients in the balanced chemical equation to find the mole ratio. Relate moles of what you were given to moles of what you are determining using the mole ratio. Step 4: Convert moles to grams using molar mass as a conversion factor. It's always a good idea to check to make sure you have answered the question you were asked. Here you were asked to calculate grams of carbon. Another step or two would be necessary if you had been asked to report your answer in some other unit, such as kg. Reactions with Limiting Amount of Reactants In actual chemical reactions, one or more reactants may be in excess. The limiting reactant will be consumed completely and limit the amount of product formed. The following exercise provides a simplified view of how limiting reactants affect a chemical reaction. Example: 30 g NO2 and 10 g H2O react as shown below. 3 NO2(g) + H2O(l) ---------------- 2 HNO3(l) + NO(g) What is the limiting reactant? When two or more reactants are present, one reactant must be limiting. To determine which reactant is limiting, we need to look at the mole ratio of the 7 www.studyguide.pk reactants involved. The mole ratio of the two reactants is 0.652 mol NO2/0.555 mol H2O = 1.17 According to the stoichiometry of the balanced equation, the mole ratio should be 3 mol NO2/1 mol H2O = 3. We see that there is not enough NO2, and thus NO2 is the limiting reactant. H2O is the excess reactant. b. What amount of HNO3 forms under these conditions? Once the limiting reactant is consumed, no additional product can be formed. We therefore use the limiting reactant to calculate the amount of product. c. What amount of NO2 and H2O remain? All of the limiting reactant is consumed, so no NO2 remains. Stoichiometry will allow us to calculate the amount of H2O remaining by first determining how much H2O reacts. 3.5 Yields of Chemical Reactions In the previous section, it was assumed that the reactions proceed "to completion,"—in other words, that all reactants are converted to products. However, this is not always the case. Side reactions can result in the formation of secondary products. Chemists calculate the percent yield of a reaction by comparing the amount of product formed to the theoretical yield predicted from stoichiometry. 8 www.studyguide.pk Example: What is the theoretical yield of Al2S3 when 10.0 g of aluminum is reacted with excess sulfur according to the equation below? 2Al(s) + 3 S(s) Al2S3(s) First, we need to convert grams of aluminum to moles: Next, we relate moles of aluminum to moles of product using the stoichiometric coefficients as a mole ratio: Finally, we calculate our theoretical yield of Al2S3 in grams. Example: A student performing the reaction above collected 18.7 g Al2S3. What is her percent yield? Percent Yield The percent yield is defined as The predicted yield is determined by the masses used in a reaction and the mole ratios in the balanced equation. This predicted yield is the "ideal". It is not always possible to get this amount 9 www.studyguide.pk of product. Reactions are not always simple. There often are competing reactions. For example, if you burn carbon in air you can get carbon dioxide and carbon monoxide formed. The two reactions occur simultaneously. Some carbon atoms end up in CO and others end up in CO2 Example: What is the percent yield for a reaction if you predicted the formation of 21. grams of C6H12 and actually recovered only 3.8 grams? 1. Recall definition of percent yield. 2. Substitute the actual and predicted yields. 3. Answer: The percent yield is 18 %. Example: A reaction between solid sulfur and oxygen produces sulfur dioxide. The reaction started with 384 grams of S6 (s). Assume an unlimited supply of oxygen. What is the predicted yield and the percent yield if only 680 grams of sulfur dioxide are produced? 1 S6 (s) 6 O2 (g) 6 SO2 (g) Step 1 : Calculate the molar masses for S6 (s) and SO2(g). The oxygen has no effect on the answer because there is more than you need. 1 mole S6 (s)= 193 grams S6 (s); 1 mole SO2(g) = 64 grams SO2(g) Step 2 : Mole ratio method Determine the mole ratio for 1 mole S6 (s) to mole SO2(g) The balanced equation indicates 1 mole S6 (s) to 6 mole SO2(g) Step 3: Calculate the number of moles of S6 (s) 10 www.studyguide.pk moles S6 (s) = [384 g S6 (s)][ 1 mole S6 (s)/ 192 g S6 (s)] = 2 moles S6 (s) Step 4: Calculate the moles of SO2(g) expected using the mole ratio 6 SO2(g) / 1 S6 (s) moles SO2(g) = 2 moles S6 (s)[6 SO2(g) / 1 S6 (s)] = 12 moles SO2(g) Step 5: Calculate the grams of SO2(g) predicted using 1 mole SO2(g) = 64 grams SO2(g) grams SO2(g) = 12 moles SO2(g)[64 grams SO2(g)/1 mole SO2(g)] = 768 g SO2(g) Step 6: Calculate the percent yield using the definition Percent yield = 100[actual yield/ predicted yield] = 100[680 grams SO2(g)/ 768 g SO2(g)]= 89% Concentrations of Reactants in Solution: Moles/dm3 (Molarity) Many chemical reactions occur in solution. In order to make stoichiometric calculations for these reactions, we need to be able to express the concentration of reactants in solution. One of the most useful concentration units is moles/dm3 (molarity abbreviated M). Using moles/dm3 as a conversion factor is quite useful. In the laboratory, solutions are prepared according to several steps. Let's prepare 250 mL of a 0.100 M solution of NaCl. (Unless otherwise noted, solutions are aqueous and water is the solvent.) First, we have to do a calculation. We need to know how many grams of NaCl to weigh. 0.250 L (0.100 mol NaCl/dm3 solution) = 0.0250 mol NaCl (58.5 g/mol) = 1.46 g NaCl. Next, we weigh this amount on a balance and transfer the solid to a 250 mL volumetric flask—a very precise piece of glassware designed to contain only a specific volume of liquid. Finally, we add our solvent—in this case, water—to the flask. First, we add a small amount to dissolve the solute. Then we add water up to the calibration mark on the flask and mix well. 11 www.studyguide.pk Solution Stoichiometry Moles/dm3 serves as a useful link between the volume of a solution and the number of moles of a solute. The flow diagram below summarizes the steps in stoichiometry calcuations involving solutions. Example: How many mL of a 0.90 M solution of HCl is required to react with 4.16 g CaCO3, according to the equation below? CaCO3(s) + 2 HCl(aq)-------CaCl2(aq) + CO2(g) + H2O(l) In this problem, we are given the concentration of the HCl solution. We are given a mass in grams of one of the reactants. So our first step is to convert mass to moles. 4.16 g CaCO3 (1 mol/100 g) = 4.16 x10-2 mol CaCO3 Next, we relate moles of CaCO3 to moles of HCl required using the coefficients in the balanced equation. The reaction ratio is 2:1 respectively 4.16 x 10-2 mol CaCO3 (2 mol HCl / 1 mol CaCO3) = 8.31 x 10-2 mol HCl Now we can convert moles of HCl to volume of HCl using the moles/dm3. 8.31 x 10-2 mol HCl (1 dm3 solution / 0.90 mol HCl) = 9.23 x 10-2 L HCl solution 9.23 x 10-2 L (1000 mL / L) = 92.3 mL HCl solution 12 www.studyguide.pk The Rates of Chemical Reactions KEYWORDS: Activation energy ... Catalysts … Concentration effect … Graphs-gas collection ... Graphs in general ... How reactions happen … Interpreting results ... Light (catalyst) effect ... Measuring rate … … Pressure effect … Rate of reaction … Reaction profiles ... Stirring effect … Surface area effect … Temperature effect What do we mean by Rate and how is it measured? The phrase ‘rate of reaction’ means ‘how fast is the reaction’. It can be measured as the 'rate of formation of product' or the 'rate of disappearance of reactant'. o Rusting is a ‘slow’ reaction, you hardly see any change looking at it!, o weathering of rocks is a very slow reaction, o fermentation of sugar to alcohol is quite slow but you can see the carbon dioxide bubbles forming in the 'froth'! o a ‘fast’ reaction would be magnesium dissolving in hydrochloric acid, o and an explosion and burning/combustion reactions would be described as ‘very fast’! The importance of "Rates of Reaction knowledge": o Time is money in industry, the faster the reaction can be done, the more economic it is. Hence the great importance of catalysts eg transition metals or enzymes. o Health and Safety Issues: Mixtures of flammable gases in air present an explosion hazard (gas reactions like this are amongst the fastest reactions known). eg Methane gas in mines, petrol vapour etc. so knowledge of 'explosion/ignition threshold concentrations', ignition temperatures and activation energies are all important knowledge to help design systems of operation to minimise risks. Flammable fine dust powders can be easily ignited eg coal dust in mines, flour in mills. Fine powders have a large surface area which greatly increases the reaction rate causing an explosion. Any spark from friction is enough to initiate the reaction! A reaction will continue until one of the reactants is used up. To measure the ‘speed’ or ‘rate’ of a reaction depends on what the reaction is, and can what is formed be measured as the reaction proceeds? Two examples are outlined below. When a gas is formed from a solid reacting with a solution, it can be collected in a gas syringe (see diagram below and the graph). o The initial gradient of the graph eg in cm3/min gives an accurate measure of how fast the gaseous product is being formed. o If the reaction is allowed to go on, you can measure the final maximum volume of gas and the time at which the reaction stops. Reactions involving: o (i) metals dissolving in acid eg magnesium + sulphuric acid ==> magnesium sulphate + hydrogen www.studyguide.pk o o o o Mg(s) + H2SO4(aq) ==> MgSO4(aq) + H2(aq) (ii) carbonates dissolving in acids calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(g) and (iii) the manganese(IV) oxide catalysed decomposition of hydrogen peroxide hydrogen peroxide ==> water + oxygen, 2H2O2(aq) ==> 2H2O(l) + O2(g) can all be followed with this method. You can investigate the effects of (a) the solution concentration, (b) the temperature of the reactants, (c) the size of the solid particles (surface area effect), (d) the effectiveness of different catalysts on hydrogen peroxide decomposition. The shape of the graph is quite characteristic (see below). o The reaction is fastest at the start when the reactants are at a maximum (steepest gradient in cm3/min). o The gradient becomes progressively less as reactants are used up and the reaction slows down. o Finally the graph levels out when one of the reactants is used up and the reaction stops. o The amount of product depends on the amount of reactants used. o The initial rate of reaction is obtained by measuring the gradient at the start of the reaction. A tangent line is drawn through the first part of the graph, which is usually reasonably linear from the x,y origin 0,0. and so gives you an initial rate of reaction in cm 3 gas/minute. 2 www.studyguide.pk The rate of a reaction that produces a gas can also be measured by following the mass loss as the gas is formed and escapes from the reaction flask. o The method is ok for reactions producing carbon dioxide or oxygen, o but not very accurate for reactions giving hydrogen (low mass loss). When sodium thiosulphate reacts with an acid, a yellow precipitate of sulphur is formed. o To follow this reaction you can measure how long it takes for a certain amount of sulphur to form. o You do this by observing the reaction down through a conical flask, viewing a black cross on white paper (see diagram below). o The X is eventually obscured by the sulphur precipitate and the time noted. o sodium thiosulphate + hydrochloric acid ==> sodium chloride + sulfur dioxide + water + sulphur o Na2S2O3(aq) + 2HCl(aq) ==> 2NaCl(aq) + SO2(aq) + H2O(l) + S(s) 3 www.studyguide.pk Mix Ongoing Watch stopped By using the same flask and paper X you can obtain a relative measure of the speed of the reaction in forming the same amount of sulphur. The speed or rate of reaction can expressed as 'x amount of sulphur'/time, so the rate is proportional to 1/time for a given set o You can investigate the effects of (a) the hydrochloric acid or sodium thiosulphate concentration (b) the temperature of the reactants. The theory of how reactions happen Reactions can only happen when the reactant particles collide, but most collisions are not successful in forming product molecules. The minority high kinetic energy collisions between particles which do produce a chemical change are called 'fruitful collisions' The reactant molecules must collide with enough energy to break the original bonds so those new bonds in the product molecules can be formed. All the rate-controlling factors are to do with the frequency and energy of reactant particle collision. In the case of temperature, the energy of the collision is even more important than the frequency effect. The particle theory of gases and liquids and the diagrams below will help you understand what is going on. The Factors affecting the Rate of Chemical Reactions The effect of Concentration (see also graphs) If the concentration of any reactant in a solution is increased, the rate of reaction is increased o Increasing the concentration, increases the probability of a collision between reactant particles because there are more of them in the same volume. o Examples ….. 4 www.studyguide.pk Increasing the concentration of acid molecules increases the frequency or chance at which they hit the surface of marble chips to dissolve them (slower faster). Increasing the concentration of reactant A or B will increase the chance or frequency of collision between them and increase the speed of product formation (slower faster). The effect of Pressure If one or more of the reactants is a gas then increasing pressure will effectively increase the concentration of the reactant molecules and speed up the reaction. o because the increased chance of a 'fruitful' collision produces the increase in reaction. The effect of Stirring In doing rate experiments with a solid reactant (eg marble chips-acid solution) or a solid catalyst (manganese(IV) oxide-hydrogen peroxide solution) it is sometimes forgotten that stirring the mixture is an important rate factor. If the reacting mixture is not stirred ‘evenly’ then the reactant concentration in solution becomes much less near the solid, which tends to settle out. At the bottom of the flask the reaction prematurely slows down distorting the overall rate measurement and making the results uneven and therefore inaccurate. 5 www.studyguide.pk The effect of Surface Area - solid particle size If a solid reactant or a solid catalyst is broken down into smaller pieces the rate of reaction increases. The speed increase happens because smaller pieces of the same mass of solid have a greater surface area compared to larger pieces of the solid. Therefore, there is more chance that a reactant particle will hit the solid surface and react. The diagrams below illustrate the acid–marble chip reaction, but they could also represent a solid catalyst in a solution of reactants. The effect of Temperature (see also graphs) When gases or liquids are heated the particles gain kinetic energy and move faster (see diagrams below). The increased speed increases the chance of collision between reactant molecules and the rate increases. However this is not the main reason for the increased reaction speed. 6 www.studyguide.pk Most molecular collisions do not result in chemical change. Before any change takes place on collision, the colliding molecules must have a minimum kinetic energy called the Activation Energy shown on the energy level diagrams below. o Going up and to top 'hump' represents bond breaking on reacting particle collision. The arrow up represents this minimum energy needed to break bonds to initiate the reaction. o Going down the other side represents the new bonds formed in the reaction products. It does not matter whether the reaction is an exothermic or an endothermic energy change. Now when heated molecules have a greater kinetic energy, a greater proportion of them have the required activation energy to react. The increased chance of 'fruitful' higher energy collisions greatly increases the speed of the reaction. 7 www.studyguide.pk The effect of a Catalyst (see also light effect and graphs) I was once asked "what is the opposite of a catalyst?" There is no real opposite to a catalyst, other than the uncatalysed reaction! The word catalyst means changing the rate of a reaction with some other material 'added to' or in 'contact with' the reaction mixture. There are the two phrases you may come across: o a 'positive catalyst' meaning speeding up the reaction o OR a 'negative catalyst' slowing down a reaction Catalysts increase the rate of a reaction by helping break chemical bonds in reactant molecules. This effectively means the Activation Energy is reduced (see diagram below). Therefore at the same temperature, more reactant molecules have enough kinetic energy to react compared to the uncatalysed situation. Although a true catalyst does take part in the reaction, BUT does not get used up and can be reused with more reactants. o It is chemically the same at the end of the reaction but it may change a little physically if its a solid. o In the hydrogen peroxide solution decomposition by the solid black catalyst manganese dioxide, the solid can be filtered off when reaction stops 'fizzing'. o After washing with water, it can be collected and added to fresh colourless hydrogen peroxide solution and the oxygen production 'fizzing' is instantaneous! Note: At the end of the experiment the solution is sometimes stained brown from minute manganese dioxide particles, the reaction is exothermic and the heat has probably caused some disintegration of the catalyst. A solid catalyst might change physically by becoming more finely divided, especially if the reaction is exothermic. Different reactions need different catalysts and they are extremely important in industry: examples .. o nickel catalyses the hydrogenation of unsaturated fats to margarine o iron catalyses the combination of unreactive nitrogen and hydrogen to form ammonia o enzymes in yeast convert sugar into alcohol o zeolites catalyse the cracking of big hydrocarbon molecules into smaller ones o most polymer making reactions require a catalyst surface or additive with the monomer molecules. Enzymes are biochemical catalysts o They have the advantage of bringing about reactions at normal temperatures and pressures which would otherwise need more expensive and energy-demanding equipment. 8 www.studyguide.pk The Effect of Light Light energy (uv or visible radiation) can initiate or catalyse particular chemical reactions. o As well as acting as an electromagnetic wave, light can be considered as an energy 'bullets' called photons and they have sufficient 'impact' to break chemical bonds, that is, enough energy to overcome the activation energy. Examples: o Silver salts are converted to silver in the chemistry of photographic exposure of the film. Silver chloride (AgCl), silver bromide (AgBr) and silver iodide (AgI) are all sensitive to light ('photosensitive'), and all three are used in the production of various types of photographic film to detect visible light and beta and gamma radiation from radioactive materials. Each silver halide salt has a different sensitivity to light. When radiation hits the film the silver ions in the salt are reduced by electron gain to silver Ag+ + e- ==> Ag, the halide ion is oxidised to the halogen molecule 2X- ==> X2 + 2e AgI is the most sensitive and used in X-ray radiography, AgCl is the most sensitive and used in 'fast' film for cameras. o Photosynthesis in green plants: The conversion of water + carbon dioxide ==> glucose + oxygen 6H2O(l) + 6CO2(g) ==> C6H12O6(aq) + O2(g) requires the input of sunlight 9 www.studyguide.pk energy The green chlorophyll molecules absorb the photon energy packets and initiate the chemical changes shown above. Photochemical Smog: This is very complex chemistry involving hydrocarbons, carbon monoxide, ozone, nitrogen oxides etc. Many of the reactions to produce harmful chemicals are catalysed by light energy. o More examples of interpreting graphical results (see two other graphs and notes) the decrease in the amount of a solid reactant with time the increase in the amount of a solid product with time the increase in speed of a reaction with increase in temperature the decrease in reaction time with increase in temperature the increase in the amount of a gas formed in a reaction with time typical graph showing the effect of concentration on the speed of a reaction 10 www.studyguide.pk slower and only half as much of the original gas is formed. The graphs on the left is typical of where a gaseous product is being collected. The middle graph might represent the original experiment 'recipe' and temperature. Then the experiment repeated with variations eg X could be the same recipe but a catalyst added, forming the same amount of product, but faster. Initially, Y and Z might represent progressively lower concentrations. Z could represent taking half the amount of reactants or half a concentration. Its W might represent taking double the quantity of reactants to forming twice as much gas eg same volume of reactant solution but doubling the concentration, so producing twice as much gas with initially double the speed (gradient). DONE 11 www.studyguide.pk Introduction to electrolysis Electrolysis is the process of electrically inducing chemical changes in a conducting melt or solution eg splitting an ionic compound into the metal and non-metal. SUMMARY OF COMMON ELECTRICAL CONDUCTORS and what makes up the circuit? o These materials carry an electric current via freely moving electrically charged particles, when a potential difference (voltage!) is applied across them, and they include: o All metals (molten or solid) and the nonmetal carbon (graphite). This conduction involves the movement of free or delocalised electrons (e- charged particles) and does not involve any chemical change. Any molten or dissolved material in which the liquid contains free moving ions is called the electrolye. Compounds that dissociate to a large extent (70 to 99%) into ions when dissolved in water are classified as strong electrolytes. Compounds that dissociate to a small extent are termed weak electrolytes. Ions are charged particles eg Na+ sodium ion, or Cl- chloride ion, and their movement or flow constitutes an electric current, because a current is moving charged particles. What does the complete electrical circuit consist of? There are two ion currents in the electrolyte flowing in opposite directions + positive cations eg Na attracted to the negative cathode electrode, and negative anions eg Cl attracted to the positive anode electrode, remember no electrons, they flow in metal wires or carbon (graphite)! The circuit of 'charge flow' is completed by the electrons moving around the external circuit eg copper wire, metal or graphite electrode, from the positive to the negative electrode The molten or dissolved materials are usually acids, alkalis or salts and their electrical conduction is usually accompanied by chemical changes eg decomposition. www.studyguide.pk The chemical changes occur at the electrodes which connect the electrolyte liquid containing ions with the external d.c. electrical supply. If the current is switched off, the electrolysis process stops. Liquids that conduct must contain freely moving ions to carry the current and complete the circuit. You can't do electrolysis with an ionic solid!, the ions are too tightly held by chemical bonds and can't flow from their ordered situation! o When an ionically bonded substances are melted or dissolved in water the ions are free to move about. However some covalent substances dissolve in water and form ions. eg hydrogen chloride HCl, dissolves in water to form 'ionic' hydrochloric acid H+Cl-(aq) The solution or melt of ions is called the electrolyte which forms part of the circuit. The circuit is completed by eg the external copper wiring and the (usually) inert electrodes like graphite (form of carbon) or platinum AND electrolysis can only happen when the current is switched on and the circuit complete. ELECTROLYSIS SPLITS a COMPOUND: o When substances which are made of ions are dissolved in water, or melted material, they can be broken down (decomposed) into simpler substances by passing an electric current through them. o This process is called electrolysis. o Since it requires an 'input' of energy, it is an endothermic process. During electrolysis in the electrolyte (solution or melt of free moving ions) ... o positive metal or hydrogen ions move to the negative electrode (cations attracted to cathode), eg in the diagram, sodium ions Na+ , move to the -ve electrode, o and negatively charged ions move to the positive electrode (anions attracted to anode), eg in the diagram, chloride ions Cl-, move to the +ve electrode. During electrolysis, gases may be given off, or metals dissolve or are deposited at the electrodes. o Metals and hydrogen are formed at the negative electrode from positive ions by electron gain (reduction), eg in molten sodium chloride + sodium ions change to silvery grey liquid sodium, Na + e ==> Na o 2 www.studyguide.pk and non-metals eg oxygen, chlorine, bromine etc. are formed from negative ions changing on the positive electrode by electron loss (oxidation), eg in molten sodium chloride chloride ions change to green chlorine gas, 2Cl -2e ==> Cl2. In a chemical reaction, if an oxidation occurs, a reduction must also occur too (and vice versa) so these reactions 'overall' are called redox changes. o You need to be able to complete and balance electrode equations or recognise them and derive an overall equation for the electrolysis. o ELECTROLYSIS OF MOLTEN COMPOUNDS Electrolysis of Molten Lead (II) bromide. Electrolyte: Molten lead (II) bromide. PbBr2 Electrode: carbon ( inert ) Ions present: Pb 2+ + 2Br1- Reaction at cathode ( negative electrode) Lead ions Pb2+ + 2 electrons ---------------- Lead atoms. + 2 e ----------------- Pb Molten lead will collect at the bottom near cathode. Reaction at anode (negative electrode) Bromide ions -------------------- bromine gas 2Br1- ----------------- Br2 + 2 electrons + 2e Bromine gas will be discharged at anode as a reddish brown gas. The over all reaction: Lead Bromide------------ Molten lead PbBr2 (l) ----------------- + Bromine gas Pb (l) + Br2 (g) 3 www.studyguide.pk Electrolysis of Molten Bauxite Electrolyte: Electrodes: Ions present: Molten Aluminum oxide ( Al2O3) Carbon Al3+ & O2- The Extraction of Aluminium The purified bauxite ore of aluminium oxide is continuously fed in. Cryolite is added to lower the melting point and dissolve the ore. Ions must be free to move to the electrode connections called the cathode (-), attracting positive ions eg Al3+, and the anode (+) attracting negative ions eg O2-. When the d.c. current is passed through aluminium forms at the positive cathode (metal*) and sinks to the bottom of the tank. At the negative anode, oxygen gas is formed (non-metal). This is quite a problem. At the high temperature of the electrolysis cell it burns and oxidises away the carbon electrodes to form toxic carbon monoxide or carbon dioxide. So the electrode is regularly replaced and the waste gases dealt with! It is a costly process (6x more than Fe!) due to the large quantities of expensive electrical energy needed for the process. Raw materials for the electrolysis process: The redox details of the electrode processes: Bauxite ore of impure aluminium oxide [Al2O3 made up of Al3+ and O2- ions] Carbon (graphite) for the electrodes. Cryolite reduces the melting point of the ore and saves energy, because the ions must be free to move to carry the current Electrolysis means using d.c. electrical energy to bring about chemical changes eg decomposition of a compound to form metal deposits or release gases. The electrical energy splits the compound! At the electrolyte connections called the anode electrode (+, attracts ions) and the cathode electrode (-, attracts + ions). An electrolyte is a conducting melt or solution of At the negative (-) cathode, reduction occurs (electron gain) when the positive aluminium ions are attracted to it. They gain three electrons to change to neutral Al atoms. Al3+ + 3e- ==> Al At the positive (+) anode, oxidation takes place (electron loss) when the negative oxide ions are attracted to it. They lose two electrons forming neutral oxygen molecules. 2O2- ==> O2 + 4e 4 Note: Reduction and Oxidation always go together! The overall electrolytic decomposition is www.studyguide.pk freely moving ions which carry the charge of the electric current. ... o o o aluminium oxide => aluminium + oxygen 2Al2O3 ==> 4Al + 3O2 and is a very endothermic process, lots of electrical energy input! Reaction at cathode ( negative electrode) Aluminum ions 4Al3+ + 3 electrons ---------------- Molten Aluminum + 12 e ----------------- 4Al (l) Molten aluminum will collect at the bottom near cathode. Reaction at anode (negative electrode) Oxide ions -------------------- oxygen gas 6O2- ----------------- 3O2 + 12 electrons + 12e Oxygen gas will be discharged at anode. The over all reaction: Aluminum oxide ------------- Molten Aluminum + oxygen gas 4Al2O3 (l) ----------------- 4Al (l) + 6O2 (g) ELECTROLYSIS OF SOLUTIONS Electrolysis of Brine Electrolyte: Electrodes: Ions present: Aqueous Sodium chloride NaCl Carbon / Copper H+ , OH-, Na+, Cl-. Sodium chloride solution gives equal volumes of hydrogen gas (-) and green chlorine gas (+) with sodium hydroxide left in solution. However in dilute solution, oxygen gas as well as chlorine gas is produced. 5 www.studyguide.pk Reaction at cathode ( negative electrode) Hydrogen ions 2H+ + 2e ----------------- + 2e --------------------- Hydrogen gas. H2 (g) Hydrogen gas liberated at cathode. Reaction at anode ( positive electrode) Chloride ion ------------- Chlorine gas + 2 electrons 2Cl- (aq) -------------------- Cl2 + 2e Chlorine gas liberated at anode. Over all reaction Sodium chloride + Water -- Hydrogen gas + Chlorine gas + Sodium hydroxide 2NaCl(aq) + H2O (l) ---------------------- H2(g) + Cl2(g) + NaOH (aq). Electrolysis of Concentrated HCl Electrolyte: Electrodes: Ions present: Aqueous Hydrochloric acid HCl Carbon / Copper H+ , OH-, Cl-. Reaction at cathode ( negative electrode) Hydrogen ions 2H+ + 2e ----------------- + 2e --------------------- Hydrogen gas. H2 (g) Hydrogen gas liberated at cathode. Reaction at anode ( positive electrode) Chloride ion ----------------- Chlorine gas + 2 electrons. 2Cl- (aq) -------------------- Cl2 + 2e 6 www.studyguide.pk Chlorine gas liberated at anode. Over all reaction Hydrochloric acid ------------------ Hydrogen gas + Chlorine gas 2HCl(aq) ---------------------- H2(g) + Cl2(g) Electrolysis of dilute sulphuric acid Electrolyte: Electrodes: Ions present: Aqueous Sulphuric acid H2SO4 Carbon / Copper H+ , OH-, SO42-. Reaction at cathode ( negative electrode) Hydrogen ions 2H+ + 2e ----------------- + 2e --------------------- Hydrogen gas. H2 (g) Hydrogen gas liberated at cathode. Reaction at anode ( positive electrode) Hydroxide ions--------------- Oxygen gas + water + 4 electrons 4OH- (aq) -------------------- O2 (g) + 2H2O (l) + 4e- Oxygen gas liberated at anode Over all reaction 2H2O (l) ---------------------- 2H2(g) + O2(g) Electrolysis of copper (II) sulphate solution using copper electrodes. Electrolyte: Electrodes: Ions present: Aqueous Copper sulphate solution (CuSO4) Copper metal H+ , OH-, Cu2+, SO42-. Copper 7 www.studyguide.pk Copper(II) sulphate with copper electrodes, the copper deposits at cathode and the copper dissolves from anode. The blue colour of the Cu2+ ions stays constant because Cu deposited = Cu dissolved. Both involve a 2 electron transfer so it means mass of Cu deposited = mass of Cu dissolving. Reaction at cathode ( negative electrode) Copper ions + 2e ----------------- Copper atoms. Cu2+ + 2e --------------------- Cu (s) Copper metal deposit at cathode. Reaction at anode ( positive electrode) Copper atoms from anode --------------------- Cu(s) -------------------- Copper ions in the solution + 2e Cu2+(aq) + 2e Copper metal of anode will dissolve and added in electrolyte. This method is used to refine copper. Electrolysis of copper (II) sulphate solution using graphite electrodes. Electrolyte: Electrodes: Ions present: Aqueous Copper sulphate solution CuSO4 Graphite H+ , OH-, Cu2+, SO42-. Reaction at cathode ( negative electrode) Copper ions Cu2+ + 2e ----------------- + 2e --------------------- Copper atoms. Cu (s) Copper atom deposit on cathode. Reaction at anode ( positive electrode) Hydroxyl ions --------------- Oxygen gas 4OH-(aq) ----------------------- O2(g) Oxygen gas is liberated at anode. 8 + Water + 4 electrons + 2H2O (l) + 4e www.studyguide.pk ELECTROLYSIS OF SOLUTIONS USING INERT ELECTRODES ELECTROLYTE IONS IN SOLUTIONS Concentrated HCl H+ (aq) PRODUCT AT (-) CATHODE Hydrogen gas Cl- (aq) Na+ (aq) Cl-(aq) Concentrated NaCl solution Dilute Sulphuric H+(aq) SO42- (aq) Acid Copper (II) Cu2+(aq) SO42-(aq) Sulphate Solution PRODUCT (+) ANODE AT Chlorine gas Hydrogen gas Chlorine gas from water Hydrogen gas Oxygen gas from water Copper metal Oxygen gas from water ELECTRODE PRODUCTS FROM DIFFERENT IONS IN AQUEOUS SOLUTION (USING INERT ELECTRODE) CATION K+ Na+ Ca2+ Mg2+ Al3+ Ni2+ Pb2+ H+ Cu2+ Ag+ PRODUCT AT CATHODE Hydrogen gas from water Nickel Lead Hydrogen Copper Silver 9 ANION PRODUCT AT ANODE ClBrI- Chlorine Bromine Iodine SO42- Oxygen from water NO3- Oxygen from water www.studyguide.pk ELECTROPLATING The coating of a metal object with another metal object is called Electroplating. It is carried out in a cell called plating bath. It contains an electrolyte. For silver plating electrolyte is a solution of silver salt. The article to be plated is made the cathode in the cell so that metal ions move to it when current is switched on. GENERAL NOTE ON ELECTROLYSIS: . What does the complete electrical circuit consist of? o There are two ion currents in the electrolyte flowing in opposite directions: positive cations eg Al3+ attracted to the negative cathode electrode, and negative anions eg O2- attracted to the positive anode electrode, BUT remember no electrons flow in the electrolyte, only in the graphite or metal wiring! o The circuit of 'charge flow' is completed by the electrons moving around the external circuit eg copper wire or graphite electrode, from the positive to the negative electrode o This e- flow from +ve to -ve electrode perhaps doesn't make sense until you look at the electrode reactions, electrons released at the +ve anode move round the external circuit to produce the electron rich negative cathode electrode. 10 www.studyguide.pk 4. Simple Cells or batteries In electrolysis, electrical energy is taken in (endothermic) to enforce the oxidation and reduction to produce the products. The chemistry of simple cells or batteries is in principle the opposite of electrolysis. A redox reaction occurs to produce products and energy is given out. It is exothermic, BUT the energy is released as electrical energy and the system shouldn't heat up. A simple cell can be made by dipping two different pieces of metal (of different reactivity) into a solution of ions eg a salt or dilute acid. The greater the difference in reactivity, the bigger the voltage produced. However this is not a satisfactory 'battery' for producing even a small continuous current. One of the first practical batteries is called the 'Daniel cell'. o It uses a half-cell of copper dipped in copper(II) sulphate, o and in electrical contact with a 2nd half-cell of zinc dipped in zinc sulphate solution. o The zinc is the more reactive, and is the negative electrode, releasing electrons because 2+ on it zinc atoms lose electrons to form zinc ions, Zn(s) ==> Zn (aq) + 2eo The less reactive metal copper, is the positive electrode, and gains electrons from the negative electrode through the external wire connection and here .. 2+ the copper(II) ions are reduced to copper atoms, Cu (aq) + 2e ==> Cu(s) o Overall the reactions is: Zn(s) + CuSO4(aq) + ZnSO4(aq) + Cu(s) 2+ 2+ or ionically: Zn(s) + Cu (aq) + Zn (aq) + Cu(s) o The overall reaction is therefore the same as displacement reaction, and it is a redox reaction involving electron transfer and the movement of the electrons through the external wire to the bulb or voltmeter etc. forms the working electric current. The cell voltage can be predicted by subtracting the less positive voltage from the more positive voltage: o eg a magnesium and copper cell will produce a voltage of (+0.34) - (-2.35) = 2.69 Volts o or an iron and tin cell will only produce a voltage of (-0.15) - (-0.45) = 0.30 Volts. o Note (i) the bigger the difference in reactivity, the bigger the cell voltage produced and 11 www.studyguide.pk (ii) the 'half-cell' voltages quoted in the diagram are measured against the H+(aq)/H2(g) system which is given the standard potential of zero volts. Cells or batteries are useful and convenient portable sources of energy but they are expensive compared to what you pay for 'mains' electricity. o 5. Fuel Cells Hydrogen gas can be used as fuel. o It burns with a pale blue flame in air reacting with oxygen to be oxidised to form water. hydrogen + oxygen ==> water or 2H2(g) + O2(g) ==> 2H2O(l) o It is a non-polluting clean fuel since the only combustion product is water. o It would be ideal if it could be manufactured by electrolysis of water eg using solar cells. o Hydrogen can be used to power fuel cells. o It all sounds wonderful BUT, still technological problems to solve for large scale manufacture and distribution of 'clean' hydrogen gas or use in generating electricity AND its rather an inflammable explosive gas! Fuel cells are 'battery systems' in which two reactants can be continuously fed in. The consequent redox chemistry produces a working 12 www.studyguide.pk current. Hydrogen's potential use in fuel and energy applications includes powering vehicles, running turbines or fuel cells to produce electricity, and generating heat and electricity for buildings and very convenient for remote and compact situations like the space shuttle. When hydrogen is the fuel, the product of its oxidation is water, so this is potentially a clean non-polluting and non-greenhouse gas? fuel. Most fuel cells use hydrogen, but alcohols and hydrocarbons can be used. A fuel cell works like a battery but does not run down or need recharging as long as the 'fuel' supply is there. It will produce electricity and heat as long as fuel (hydrogen) is supplied. DIAGRAM and CHEMISTRY A fuel cell consists of two electrodes consisting of a negative electrode (or anode) and a positive electrode (or cathode) which are sandwiched around an electrolyte (conducting salt/acid/alkali solution of free ions). Hydrogen is fed to the (+) anode, and oxygen is fed to the (-) cathode. The platinum catalyst activates the hydrogen atoms/molecules to separate into protons (H+) and electrons (e-), which take different paths to the (+) cathode. o The electrons go through an external circuit, creating a flow of electricity eg to light a bulb. o The protons migrate through the electrolyte and pass through the semi-permeable membrane to the cathode, where they reunite with oxygen and the electrons to produce water. Each cell only produces a small voltage (0.4V?) so many cells can be put 13 www.studyguide.pk together in series to give a bigger working voltage. Note on reverse action: o o o o If there is spare electricity from another source available, you can run the fuel cell in reverse and electrolyse the water to make hydrogen and oxygen (acting as an electrolyser). The two gases are stored, and when electricity or heat needed, the fuel cell can then be re-run using the stored gaseous fuel. this is called a regenerative fuel cell system. You can use solar energy from external panels on the space shuttle to do this, and use the fuel when in the 'darkness of night'. equation summary of hydrogenoxygen fuel cell 1 oxidation 2H2(g) ==> 4H+(aq) + 4e- (at -ve anode electrode*) 2 reduction O2(g) + 4H+(aq) + 4e- ==> 2H2O(l) (at +ve cathode electrode*) 3=1+2 2H2(g) + O2(g) ==> 2H2O(l) DONE 14 www.studyguide.pk Acids, Bases and Salts THE THEORY of ACIDS and ALKALIS and a few technical terms: o Acids are substances that form hydrogen ions (H+(aq)) when dissolved in water eg hydrochloric acid HCl gives H+(aq) and Cl-(aq) ions, sulphuric acid H2SO4 gives 2H+(aq) and SO42- ions and nitric acid HNO3 gives H+(aq) and NO3-(aq) ions. o Alkalis are substances that form hydroxide ions (OH-(aq)) in water eg sodium hydroxide NaOH gives Na+(aq) and OH-(aq) ions, calcium hydroxide Ca(OH)2 gives Ca2+(aq) and 2OH-(aq) ions. Note: an alkali is a base soluble in water. o In water, there are trace quantities of H+ and OH- ions BUT they are of equal concentration and so water is neutral. o In acid solutions there are more H+ ions than OH- ions. o In alkaline solution there are more OH- ions than H+ ions. o Acids dissociate to different extents in aqueous solution. Acids that dissociate to a large extent are strong electrolytes and strong acids. In contrast, acids that dissociate only to a small extent are weak acids and weak electrolytes In a similar manner, bases can be strong or weak depending on the extent to which they dissociate and produce OH– ions in solution. Most metal hydroxides are strong electrolytes and strong bases. Ammonia, NH3, is a weak electrolyte and weak base. o BASES eg oxides and hydroxides are substances that react and neutralise acids to form salts and water. Bases which are soluble in water are called alkalis. Acids Some common acids are listed below: Name Hydrochloric acid Sulphuric acid Nitric acid Ethanoic (acetic) acid Methanoic (formic) acid Citric Acid Formula HCl H2SO4 HNO3 CH3COOH HCOOH C6H8O7 Strong/Weak Strong Strong Strong Weak Weak Weak Where is it found? The stomach, in the lab. Acid rain, car batteries, the lab. Acid rain, in the lab. Vinegar Ant & nettle stings, descalers Citrus fruits Acids taste sour (e.g. vinegar, lemon juice). Acids are harmful to living cells. Aqueous solutions of all acids contain hydrogen ions, H+. www.studyguide.pk Bases Most bases are oxides or hydroxides of metals. Not all bases fit into these categories, however (e.g. ammonia). Some examples of bases are shown below: Name Sodium hydroxide (caustic soda) Calcium hydroxide Magnesium oxide (magnesia) Calcium carbonate Sodium hydrogencarbonate (bicarbonate) Ammonia Formula NaOH Ca(OH)2 MgO CaCO3 NaHCO3 NH3 Where is it found? Oven cleaners, in the lab. Soil lime, limewater Indigestion tablets Limestone, soil lime Baking powder Cleaning fluids, in the lab. Soluble bases are known as alkalis. Aqueous solutions of alkalis contain hydroxide ions, OH-. Reactions of Acids With metals Metals above copper in the reactivity series will react with acids, giving off hydrogen gas. The metal dissolves, forming a salt. e.g. METAL + ACID SALT + HYDROGEN Mg(s) + H2SO4(aq) MgSO4(aq) + H2(g) This is why acids corrode metals, and must be stored in glass containers. With bases (metal oxides and hydroxides) The base dissolves in the acid and neutralises it. A salt is formed. e.g. ACID + BASE SALT + WATER H2SO4(aq) + CuO(s) CuSO4(aq) + H2O(l) With metal carbonates With metal carbonates, much effervescence occurs when they react with acids, as carbon dioxide gas is released too. e.g. ACID + CARBONATE SALT + WATER + CARBON DIOXIDE 2HCl(aq) + CaCO3(s) CaCl2(s) + H2O(l) + CO2(g) 2 www.studyguide.pk Neutralisation Acids are neutralised by bases. If you have indigestion (too much acid in the stomach), you may take a tablet containing a base (e.g. magnesia, MgO). A farmer may spread lime (calcium hydroxide, Ca(OH)2) on fields to make the soil less acidic. What happens during a neutralisation reaction? ACID + ALKALI SALT + WATER e.g. hydrochloric acid + sodium hydroxide sodium chloride + water HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) To understand why water is formed, we must consider what happens to the ions that are present in the reacting solutions: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l) (from HCl) (from NaOH) The H+ ions from the acid combine with the OH- ions from the alkali – we can show this by writing a simple ionic equation: H+(aq) + OH-(aq) H2O(l) Water is formed. So if equal amounts of acid and alkali are mixed, the resulting solution will be neutral. The Na+ and Cl- ions do not take part in the neutralisation reaction, they just remain in solution - they are spectator ions. It is unnecessary to include them in the ionic equation. When an acid and a base are mixed in stoichiometric proportions, their acidic and basic properties disappear as the result of a neutralization reaction. Because the salts that form in neutralization reaction are generally strong electrolytes, we can write the neutralization reaction as an ionic equation. When the spectator ions are removed, the net ionic equation is revealed. 3 www.studyguide.pk This net ionic equation is the same for the neutralization reaction of any strong acid and strong base. For the reaction of a weak acid with a strong base, a similar neutralization occurs. Consider the neutralization of HF with KOH. HF(aq) + KOH(aq) KF(aq) + H2O(l) HF(aq) + K+(aq) + OH-(aq) H2O(l) HF(aq) + OH-(aq) K+(aq) + F-(aq) + F-(aq) + H2O(l) molecular equation ionic equation net ionic equation The weak acid HF is written as a molecular formula because its dissociation is incomplete. Formation of H30+ion The hydrogen ion H+(aq) does not exist as such in aqueous solutions. Hydrogen ions combine with water molecules to give a more stable species, the hydronium ion H3O+, as demonstrated in the following equation: HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Acids can contain different numbers of acidic hydrogens, and can yield different numbers of H3O+ ions in solution. 4 www.studyguide.pk Some important reactions of Bases (alkali = soluble base) Neutralisation with acids is dealt with above. Ammonium salts are decomposed when mixed with a base eg the alkali sodium hydroxide. o eg sodium hydroxide + ammonium chloride ==> sodium chloride + water + ammonia o NaOH + NH4Cl ==> NaCl + H2O + NH3 o The ammonia is readily detected by its pungent odour (strong smell) and by turning damp red litmus blue. o The ionic equation is: NH4+ + OH- ==> H2O + NH3 o This reaction can be used to prepare ammonia gas and as a simple chemical test for an ammonium salt. Alkali's (soluble bases) are used to produce the insoluble hydroxide precipitates of many metal ions from their soluble salt solutions. o eg sodium hydroxide + copper(II) sulphate ==> sodium sulphate + copper(II) hydroxide o 2NaOH(aq) + CuSO4(aq) ==> Na2SO4(aq) + Cu(OH)2(s) a blue precipitate o ionically: Cu2+(aq) + 2OH-(aq) ==> Cu(OH)2(s) o This reaction can be used as a simple test to help identify certain metal ions. ACIDIC, BASIC & AMPHOTERIC OXIDES Oxygen combines with most other elements to form oxides of varying physical chemical character. o On the left and middle of the Periodic Table are the basic metal oxides which react with acids to form salts eg Na2O, MgO, CuO etc. These metal oxides tend to be ionic in bonding character with high melting points. The Group 1 Alkali Metals, and to a less extent, Group 2 oxides, dissolve in water to form alkali solutions. All of them react with , and neutralise acids to form salts. o As you move left to right the oxides become less basic and more acidic. o So on the right you have the acidic oxides of the non-metals CO2, P2O5, SO2, SO3 etc. These tend to be covalent in bonding character with low melting/boiling points. Those of sulphur and phosphorus are very soluble in water to give acidic solutions which can be neutralised by alkalis to form salts. o These oxides are another example of the change from metallic element to non-metallic element chemical behaviour from left to right across the Periodic Table. o BUT life is never that simple in chemistry!: Some oxides react with both acids and alkalis and are called amphoteric oxides. They are usually relatively insoluble and have little effect on indicators. An example is aluminium oxide dissolves in acids to form 'normal' aluminium salts like the chloride, sulphate and nitrate. However, it also dissolves in strong alkali's like sodium hydroxide solution to form 'aluminate' salts. This could be considered as 'intermediate' basic-acidic character in the Periodic Table. Some oxides are neutral, tend to be of low solubility in water and have no effect on litmus, and do not react with acids or alkalis. eg CO 5 www.studyguide.pk carbon monoxide (note that CO2 carbon dioxide is weakly acidic) and NO nitrogen monoxide (note that NO2 nitrogen dioxide is strongly acidic in water). There is no way of simply predicting this kind of behaviour from periodic table patterns! Salts We have seen that when an acid reacts with a base, a salt is formed: e.g. sulphuric acid + sodium hydroxide sodium sulphate + water Here sodium sulphate (Na2SO4) is the salt formed. Salts are ionic compounds. The metal ion is provided by the base (in this case sodium ions, Na+). The non-metal ion is provided by the acid (in this case sulphate ions, SO42-). Note: Ammonia (NH3) is an unusual base - it does not contain a metal. It forms ammonium salts, containing the ammonium ion, NH4+. e.g. NH3(aq) + HNO3(aq) NH4NO3(aq) (ammonium nitrate) Soluble and Insoluble Salts Many ionic salts are soluble in (cold) water. It is useful to know whether or not a salt is soluble in water, as this will influence your choice of method for making it. The table below gives a guide to the solubility of salts: Soluble Salts All common potassium, sodium and ammonium salts All nitrates All common ethanoates 6 www.studyguide.pk All common chlorides, except lead and silver chlorides All common sulphates, except lead, barium and calcium sulphates Methods of making Salts which are water soluble Soluble salts can be made in four different ways: 1) 2) 3) 4) ACID + METAL SALT + HYDROGEN ACID + BASE SALT + WATER ACID + CARBONATE SALT + WATER + CARBON DIOXIDE ACID + ALKALI SALT + WATER Method 1 (Acid + Metal) Not suitable for making salts of metals above magnesium, or below iron/tin in reactivity. e.g. zinc + hydrochloric acid zinc chloride + hydrogen Apparatus used: (1) balance, measuring cylinder, beaker and glass stirring rod; (2) beaker/rod, bunsen burner, tripod and gauze; (3) filter funnel and filter paper, evaporating (crystallising) dish; (4) evaporating (crystallising) dish. (ii) A measuring cylinder is adequate for measuring the acid volume, you do not need the accuracy of a pipette or burette required in method (a). Add excess metal to (warm) acid. Wait until no more H2 is evolved. Filter to remove excess metal. Heat the filtrate to evaporate off water until crystallisation starts. Set aside to cool slowly and crystallise fully. Method 2 (Acid + Base) Useful for making salts of less reactive metals, e.g. lead, copper. e.g. copper(II) oxide + sulphuric acid copper(II) sulphate + water Add excess base to acid. Warm gently. Filter to remove excess base, then continue as in method 1… 7 www.studyguide.pk Method 3 (Acid + Carbonate) Useful particularly for making salts of more reactive metals, e.g. calcium, sodium. e.g. calcium carbonate + nitric acid calcium nitrate + water + carbon dioxide Add excess metal carbonate to acid. Wait until no more CO2 is evolved. Filter to remove excess carbonate, then continue as in method 1… Method 4 (Acid + Alkali) This is useful for making salts of reactive metals, and ammonium salts. It is different from methods 1-3, as both reactants are in solution. This means neutralisation must be achieved, by adding exactly the right amount of acid to neutralise the alkali. This can be worked out by titration e.g. sodium hydroxide + hydrochloric acid sodium chloride + water ammonia + sulphuric acid ammonium sulphate (1) A known volume of acid is pipetted into a conical flask and universal indicator added. The acid is titrated with the alkali in the burette (2) until the indicator turns green. (3). The volume of alkali needed for neutralisation is then noted, this is called the endpoint. (1-3) are repeated with both known volumes mixed together BUT without the contaminating indicator. (4) The solution is transferred to an evaporating dish and heated to partially evaporate the water. (5) The solution is left to cool to complete the crystallisation. (6) The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration (below) and dried (as above). 8 Cl- Mg2+ Cl- www.studyguide.pk Making Insoluble Salts Insoluble salts cannot be prepared by acid-base reactions in the same way as soluble salts. Insoluble salts are prepared by precipitation. This involves mixing solutions of two soluble salts that between them contain the ions that make up the insoluble salt. Here is an example: barium chloride + magnesium sulphate barium sulphate + magnesium chloride (soluble salt) (soluble salt) (insoluble salt) (soluble salt) BaCl2(aq) + MgSO4(aq) BaSO4(s) + MgCl2(aq) When the two solutions are mixed, a white solid precipitate of barium sulphate is formed. Let us consider what happens to the ions involved in the reaction: Ba2+ ClCl- Ba2+ ClClMg2+ SO42- Mg2+ SO42- 1. At the start, the ions of each soluble salt move about freely in solution, in their separate containers. 2. After mixing, the ions of the two solutions are free to collide with each other, so new combinations of ions are possible. 3. If one of the new combinations of ions is an insoluble salt, it will be precipitated from the solution. The other ions simply remain in solution - they are spectator ions, and play no part in the reaction. precipitate of solid BaSO4 9 www.studyguide.pk The ionic equation for the reaction is: Ba2+(aq) + SO42-(aq) BaSO4(s) Notice this equation does not include the spectator ions, only the ions that combine to form the precipitate. Once the reaction is complete, the precipitate can be filtered off, washed with distilled water and dried. This method can also be used to prepare many insoluble metal hydroxides and carbonates. All common hydroxides/carbonates are insoluble except sodium, potassium and ammonium hydroxides/carbonates. e.g. iron(III) chloride + sodium hydroxide iron(III) hydroxide + sodium chloride (rust brown ppt.) FeCl3(aq) + 3NaOH(aq) Fe(OH)3(s) + 3NaCl(aq) The ionic equation for this reaction, which shows the formation of the precipitate whilst missing out the spectator ions is as follows: Fe3+(aq) + 3OH-(aq) Fe(OH)3(s) Types of Salts Normal Salts: Normal salts are formed when all the replaceable hydrogen ions in the acid have been completely replaced by metallic ions. HCl(aq) H2SO4(aq) + NaOH(aq) + ZnO(aq) NaCl(aq) + H2O(l) ZnSO4(aq) + H2O(l) Normal salts are neutral to litmus paper. Acid salts: Acid salts are formed when replaceable hydrogen ions in acids are only partially replaced by a metal. Acid salts are produced only by acids containing more then one replaceable hydrogen ion. Therefore an acid with two replaceable ions e.g. H2SO4 will form only one 10 www.studyguide.pk acid salt, while acid with three replaceable hydrogen ions e.g. H3PO4 will form two different acid salts. H2SO4(aq) + KOH(aq) H3PO4(aq) + NaOH H3PO4(aq) + 2NaOH(aq) KHSO4(aq) + H2O(l) NaH2PO4(aq) + H2O(l) Na2HPO4(aq) + 2H2O(l) An acid salt will turn blue litmus red. In the presence of excess metallic ions an acid salt will be converted into a normal salt as its replaceable hydrogen ions become replaced. Basic Salts: Basic salts contain the hydroxide ion, OH-. They are formed when there is insufficient supply of acid for the complete neutralization of the base. A basic salt will turn red litmus blue and will react with excess acid to form normal salt. Zn(OH)2(s) + HCl(aq) Zn(OH)Cl(aq) Zn(OH)Cl(aq) + HCl(aq) Mg(OH)2(s) + HNO3(aq) Mg(OH)NO3(aq) + HNO3(aq) + H2O(l) ZnCl2(aq) + H2O(l) Mg(OH)NO3(aq) + H2O(l) Mg(NO3)2(aq) + H2O(l) The pH Scale - Acids and Alkalis The colours of solutions with universal indicator The pH scale is a measure of the relative acidity or alkalinity of a solution. To find the pH of a solution an indicator is used like Universal Indicator. An indicator is a substance or mixture of substances that when added to the solution gives different colours depending on the pH of the solution. Universal indicator is a very handy indicator for showing whether the solution is acid, neutral or alkaline and gives the pH to the nearest pH unit. Water is a neutral liquid with a pH of 7 (green). When a substance dissolves in water it forms an aqueous (aq) solution that may be acidic, neutral or alkaline. Acidic solutions have a pH of less than 7, and the lower the number, the stronger the acid is. The colour can range from orange-yellow (pH 3-6) for partially ionised weak acids like ethanoic acid (vinegar) to carbonated water. Strong acids like hydrochloric, sulphuric and nitric are fully ionised and give a pH 1 or less! and a red colour with universal 11 www.studyguide.pk indicator and litmus paper. Neutral solutions have a pH of 7. These are quite often solutions of salts, which are themselves formed from neutralising acids and bases. The 'opposite' of an acid is called a base. Some bases are soluble in water to give alkaline solutions - these are known as alkalis. Alkaline solutions have a pH of over 7 and the higher the pH the stronger is the alkali. Weak alkalis (soluble bases) like ammonia give a pH of 10-11 but strong alkalis (soluble bases) like sodium hydroxide give a pH of 13-14. They give blue/purple colour with universal indicator or litmus paper. NEUTRALISATION usually involves mixing an acid (pH <7) with a base or alkali (pH > 7) which react to form a neutral salt solution of pH 7. INDICATORS. Indicators are the substances that have different colors in acidic and in alkaline solution. Some important indicators are given below S.No. Indicator Color in strongly acidic solution pH at which color changes Color in strongly alkaline solution 1. methyl orange red 4 yellow 2. bromothymol blue yellow 7 blue 3. phenolphthalein colorless 9 red 4. screened methyl orange red 4 green 12 www.studyguide.pk What pH changes go on in a neutralisation reaction? The graphs show how the pH changes when an alkali (soluble base) and an acid neutralise each other. If you add acid to an alkali (ind. = blue), the pH starts at 13 and falls little at first. As you get near to the neutralisation point at pH 7 (ind. = green), the change becomes more dramatic. With excess acid the pH falls and then levels out to about pH1 (ind. = red). If you add alkali to an acid (ind. = red), the pH starts at about 1 and rises little at first. As you get near to the neutralisation point at pH 7 (ind. = green), the change becomes more dramatic. With excess alkali the pH rises and then levels out to about 13 (ind. = blue/violet). Formulae of bases: oxides, hydroxides and carbonates M2O (oxide O2-, soluble, alkali) MOH (hydroxide OH-, soluble, alkali) M2CO3 (carbonate CO32-, soluble mild alkali) Formulae of salts formed: The metal soluble chlorides, (or other sulphates and nitrates ion) involved MCl (chloride, Cl-) M2SO4 (sulphate, SO42-) M = Li, Na, K [usually Group 1], the ion is M+ MNO3 (nitrate, NO3-) 13 www.studyguide.pk MHCO3 (hydrogencarbonate HCO3-, soluble, mild alkali) MO (oxide, often insoluble base) M(OH)2 (hydroxide, often insoluble, alkali if soluble) MCO3 (carbonate, often insoluble) MCl2 (chloride) MSO4 (sulphate) M(NO3)2 (nitrate) Al2O3, Al(OH)3 (insoluble bases, amphoteric) AlCl3, Al2(SO4)3, Al(NO3)3 the alkaline soluble base ammonia, NH3, no stable oxide or hydroxide NH4Cl, (NH4)2SO4, NH4NO3 M = Mg, Ca, Cu, Zn, Fe [often Group 2 or Transition], the ion is M2+ Al, aluminium from Group 3 the ammonium ion in the salts from ammonia, NH4+ FURTHER EXAMPLES of WORD & SYMBOL EQUATIONS for Salt Preparations 1a. copper(II) carbonate + sulphuric acid ==> copper(II) sulphate + water +CO2. 1b. CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2(g) 2a. magnesium hydroxide + hydrochloric acid ==> magnesium chloride + water 2b. Mg(OH)2(s) + 2HCl(aq) ==> MgCl2(aq) + 2H2O(l) 3a. zinc + sulphuric acid ==> zinc sulphate + hydrogen 3b. Zn(s) + H2SO4(aq) ==> ZnSO4(aq) + H2(g) 4a. ammonia + nitric acid ==> ammonium nitrate 4b. NH3(aq) + HNO3(aq) ==> NH4NO3(aq) 5a. zinc oxide + hydrochloric acid ==> zinc chloride + water 5b. ZnO(s) + 2HCl(aq) ==> ZnCl2(aq) + H2O(l) 6a. calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide 6b. CaCO3(s) + 2HCl(aq) ==> CaCl2(aq)+ H2O(l) + CO2(g) 7a. sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide 14 www.studyguide.pk 7b. Na2CO3(s) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g) 8a. sodium hydroxide + hydrochloric acid ==> sodium chloride + water 8b. NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l) 9a. sodium hydrogencarbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide 9b. NaHCO3(s) + HCl(aq) ==> NaCl(aq) + H2O(l) + CO2(g) IONIC EQUATIONS Aqueous Reactions and Net Ionic Equations The equations written up to this point have been molecular equations. All substances have been written using their full chemical formulas as if they were molecules. Because we now know that strong electrolytes dissociate in water to their component ions, it is more accurate to write an ionic equation in which all of the ionic species are shown. In many reactions only certain ions change their 'chemical state' but other ions remain in exactly the same original physical and chemical state. The ions that do not change are called 'spectator ions'. The ionic equation represents the 'actual' chemical change and omits the spectator ions. To write a net ionic equation: 1.Write a balanced molecular equation. 2.Rewrite the equation showing the ions that form in solution when each soluble electrolyte dissociates into its component ions. Only dissolved strong electrolytes are written in ionic form. 3.Identify and cancel the spectator ions that occur unchanged on both sides of the equation. Five types of examples are presented below. When reactions between ions occur, at least one kind of ion is removed from the "field of action". Simply put, its concentration decreases as the reaction proceeds. There are three ways to remove ions: 1. Formation of an insoluble precipitate 2. Formation of a weakly ionized substance, and 3. Oxidation or reduction of an ion 15 www.studyguide.pk Let's examine each way individually: 1) Formation of An Insoluble Precipitate Excess chloride ion "drives" this reaction to the right. 2) Formation of A Weakly Ionized Substance 3) Oxidation or Reduction of An Ion Rules for Writing Ionic Equations 1) Ionic formulas are written for a strong electrolyte in solution e.g 2) Molecular formulas are written for: a) Elements, gases, solids and non-electrolytes, e.g.: 16 www.studyguide.pk b) Weak electrolytes in solution, e.g.: c) Solid strong electrolytes or precipitates, e.g.: Writing Ionic Equations When writing these equations, do so to answer the following three (3) questions: 1. What kind of reaction is it? Double decomposition? Redox? 2. What are the possible products of the reaction? 3. Are any of the possible products or reactants insoluble or weakly ionized? Double Decomposition Reactions (Precipitate and Weak Electrolyte Reactions) Examples a) KCl + NaNO3 NR Even the products would be soluble and, hence, no reaction occurs. Since the hydrogen ion and the nitrate ion are spectators, the net ionic reaction is the result. (1) Acid-base reactions: Acids can be defined as proton donors. A base can be defined as a proton acceptor. o eg any acid-alkali neutralisation involves the hydroxide ion is (base) and this accepts a proton from an acid. HCl(aq) + NaOH(aq) ==> NaCl(aq) + H2O(l) which can be re-written as H+Cl-(aq) + Na+OH-(aq) ==> Na+Cl-(aq) + H2O(l) H+(aq) + OH-(aq) ==> H2O(l) the spectator ions are Cl- and Na+ (2) Insoluble salt formation: An insoluble salt is made by mixing two solutions of soluble compounds to form the insoluble compound in a process called 'precipitation'. o (a) Silver chloride is made by mixing solutions of solutions of silver nitrate and sodium chloride. silver nitrate + sodium chloride ==> silver chloride + sodium nitrate Ag+NO3-(aq) + Na+Cl-(aq) ==> AgCl(s) + Na+NO3-(aq) the ionic equation is: Ag+(aq) + Cl-(aq) ==> AgCl(s) the spectator ions are NO3- and Na+ 17 www.studyguide.pk (b) Silver nitrate and hydrochloric acid -- precipitate formation Lead(II) iodide can be made by mixing lead(II) nitrate solution with potassium iodide solution. lead(II) nitrate + potassium iodide ==> lead(II) iodide + potassium nitrate Pb(NO3)2(aq) + 2KI(aq) ==> PbI2(s) + 2KNO3(aq) the ionic equation is: Pbg2+(aq) + 2I-(aq) ==> PbI2(s) the spectator ions are NO3- and K+ o (c) Calcium carbonate forms on eg mixing calcium chloride and sodium carbonate solutions calcium chloride + sodium carbonate ==> calcium carbonate + sodium chloride CaCl2(aq) + Na2CO3(aq) ==> CaCO3(s) + 2NaCl(aq) ionically: Ca2+(aq) + CO32-(aq) ==> CaCO3(s) the spectator ions are Cl- and Na+ o (d) Barium sulphate forms on mixing eg barium chloride and dilute sulphuric acid barium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid BaCl2(aq) + H2SO4(aq) ==> BaSO4(s) + 2HCl(aq) ionically: Ba2+(aq) + SO42-(aq) ==> BaSO4(s) the spectator ions are CO32- and H+ (3) Redox reaction analysis: o (a) magnesium + iron(II) sulphate ==> magnesium sulphate + iron Mg(s) + FeSO4(aq) => MgSO4(aq) + Fe(s) this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below. The sulphate ion SO42-(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!! Mg(s) + Fe2+(aq) ==> Mg2+(aq) + Fe(s) Mg oxidised by electron loss, Fe2+ reduced by electron gain o (b) zinc + hydrochloric acid ==> zinc chloride + hydrogen Zn(s) + 2HCl(aq) => ZnCl2(aq) + H2(g) the chloride ion Cl- is the spectator ion Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(g) Zn oxidised by electron loss, H+ reduced by electron gain o (c) copper + silver nitrate ==> silver + copper(II) nitrate Cu(s) + 2AgNO3(aq) ==> 2Ag + Cu(NO3)2(aq) the nitrate ion NO3- is the spectator ion Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq) Cu oxidised by electron loss, Ag+ reduced by electron gain o (d) halogen (more reactive) + halide salt (of less reactive halogen) ==> halide salt (of more reactive halogen) + halogen (less reactive) X2(aq) + 2K+Y(aq) ==> 2K+X(aq) + Y2(aq) X2(aq) + 2Y-(aq) ==> 2X-(aq) + Y2(aq) the potassium ion K+ is the spectator ion halogen X is more reactive than halogen Y, F > Cl > Br > I) o 18 www.studyguide.pk X is the oxidising agent (electron acceptor, so is reduced) KY or Y- is the reducing agent (electron donor, so is oxidised) (4) Ion Exchange Resins: Ion exchange polymer resin columns hold hydrogen ions or sodium ions. These can be replaced by calcium and magnesium ions when hard water passes down the column. The calcium or magnesium ions are held on the negatively charged resin. The freed hydrogen or sodium ions do not form a scum with soap. o eg 2[resin]-H+(s) + Ca2+(aq) ==> [resin]-Ca2+[resin]-(s) + 2H+(aq) o or 2[resin]-Na+(s) + Mg2+(aq) ==> [resin]-Mg2+[resin]-(s) + 2Na+(aq) etc. (5) Scum formation with hard water: On mixing hard water with soaps made from the sodium salts of fatty acids, insoluble calcium or magnesium salts of the soap are formed ... o CaSO4(aq) + 2C17H35COONa(aq) ==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq) o or more simply ionically: Ca2+(aq) + 2C17H35COO-(aq) ==> (C17H35COO-)2Ca2+(s) o the spectator ions are SO42- and Na+ DONE 19 www.studyguide.pk Industrial Chemistry Limestone - a very useful material Limestone, is a sedimentary rock formed by the mineral and 'shelly' remains of marine organisms, including coral, in warm shallow fertile seas. It is chemically mainly calcium carbonate and is a useful material that is quarried and used directly as a building material. It reacts with acids - 'fizzing' due to carbon dioxide formation - test with 'limewater' - milky white precipitate. It is a valuable natural mineral resource and is quarried in large quantities in many countries (see environmental impact). Other uses of limestone rock are outlined below and it is an important raw material in the manufacture of cement and glass and iron. Powdered limestone can be used to neutralise acidity in lakes and soils When limestone is heated in a kiln at over 900oC, it breaks down into quicklime (calcium oxide) and carbon dioxide. Both are useful products. This type of reaction is endothermic and an example of thermal decomposition (and other carbonates behave in a similar way). o calcium carbonate (limestone) ==> calcium oxide (quicklime) + carbon dioxide CaCO3(s) CaO(s) + CO2(g) This is a reversible endothermic reaction. To ensure the change is to favour the right hand side, a high temperature of over 900oC is needed as well as the continual removal of the carbon dioxide. Note on other carbonates These also show a similar thermal decomposition to CaCO3 above ... o copper(II) carbonate(green) ==> copper(II) oxide(black) + carbon dioxide o CuCO3(s) ==> CuO(s) + CO2(g) o zinc carbonate(white) ==> zinc oxide(yellow hot, white cold) + carbon dioxide o ZnCO3(s) ==> ZnO(s) + CO2(g) o FeCO3 and MnCO3 behave in a similar way Quicklime reacts very exothermically with water to produce slaked lime (calcium hydroxide). o calcium oxide (quicklime) + water ==> calcium hydroxide (slaked lime) this is a very exothermic reaction, the quicklime 'puffs' up and steam is produced! CaO(s) + H2O(l) ==> Ca(OH)2(s) o o Lime (calcium oxide) and slaked lime (calcium hydroxide) are both used to reduce the acidity of soil on land, they are both faster and stronger acting than limestone powder. They are also used to reduce acidity in lakes and rivers due to acid rain. They are also used to neutralise potentially harmful industrial acid waste including sulphur dioxide in the flue gases of power stations. In the test for carbon dioxide, calcium hydroxide solution (limewater) forms a white milky precipitate of calcium carbonate (back to where you started!). o calcium hydroxide + carbon dioxide ==> calcium carbonate + water o Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l) www.studyguide.pk The oxides and hydroxides readily react with acids. o general word equation: oxide or hydroxide + acid ==> salt + water examples ... calcium oxide + hydrochloric acid ==> calcium chloride + water calcium hydroxide + nitric acid ==> magnesium nitrate + water calcium hydroxide + sulphuric acid ==> calcium sulphate + water CaO(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) Ca(OH)2(s) + 2HNO3(aq) ==> Ca(NO3)2(aq) + H2O(l) Ca(OH)2(s) + H2SO4(aq) ==> CaSO4(aq,s*) +2 H2O(l) * the sulphates of eg calcium and barium are not very soluble, this slows the reaction down! Solubility of calcium compounds (and the chemically similar magnesium): o Magnesium and calcium oxides or hydroxides are slightly soluble in water forming alkaline solutions. They readily react and dissolve in most acids o Magnesium and calcium carbonate are insoluble in water but readily dissolve in most dilute acids like hydrochloric, nitric and sulphuric. Calcium carbonate reacts slowly in dilute sulphuric acid because calcium sulphate is not very soluble and coats the limestone o Equation examples for calcium carbonate (similar for magnesium carbonate) ... o calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide CaCO3(s) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(g) o calcium carbonate + nitric acid ==> calcium nitrate + water + carbon dioxide CaCO3(s) + 2HNO3(aq) ==> Ca(NO3)2(aq) + H2O(l) + CO2(g) o calcium carbonate + sulphuric acid ==> calcium sulphate + water + carbon dioxide CaCO3(s) + H2SO4(aq) ==> CaSO4(aq,s) + H2O(l) + CO2(g) Magnesium and calcium hydrogencarbonate are soluble in water and cause 'hardness' ie scum with 'traditional' non-detergent soaps. Formulae are Mg(HCO3)2 and Ca(HCO3)2 Cement is produced by roasting a mixture of powdered limestone with powdered clay* in a rotary kiln. When cement is mixed with water, sand and crushed rock, a slow chemical reaction produces a hard, stone-like building material called concrete. o * Clay is used directly to make pottery and other ceramics Glass is made by heating a mixture of limestone (CaCO3), sand (mainly silica = silicon dioxide = SiO2) and 'soda' (sodium carbonate, Na2CO3). Limestone is used to remove acidic oxide impurities in the extraction of iron and in making steel. Calcium oxide and calcium hydroxide also react with acids to form salt. 2 www.studyguide.pk Why is sulphuric acid a useful material? How is it made? – Contact Process Because sulphuric acid has so many uses the industrial development of a country is sometimes measured by the amount of sulphuric acid that is used each year. Sulphuric acid is made starting from the element sulphur which is found in the Earth's crust. Sulphuric acid is used as car battery acid and is used to make fertilisers, dyes and detergents. o eg ammonia + sulphuric acid ==> ammonium sulphate (a fertiliser salt) o 2NH3(aq) + H2SO4(aq) ==> (NH4)2SO4(aq) => evaporation to get crystals o Its acid action make it good for cleaning metal surfaces in industry. Sulphuric acid is manufactured from the raw materials sulphur, air and water. (1) Sulphur is burned in air to form sulphur dioxide (exothermic). o In the reaction the sulphur is oxidised (O gain) S(s) + O2(g) ==> SO2(g) Note: Sulphur dioxide itself is a useful chemical in its own right: o It is used as a bleach in the manufacture of wood pulp for paper manufacture o and its toxic nature makes it useful as a food preservative by killing bacteria. (2) In the reactor the sulphur dioxide is mixed with air and the mixture passed over a catalyst of vanadium oxide V205 at a high temperature (about 450°C) and at a pressure of between one and two atmospheres. It is a 2nd exothermic oxidation and is known as the Contact Process. In the reactor the sulphur dioxide is oxidised in the reversible exothermic reaction ... o 2SO2(g) + O2(g) 2SO3(g) The reaction forms sulphur trioxide and the equilibrium is very much to the right hand side ... o despite the reaction being exothermic and a high temperature used (favours reverse reaction R to L, energy change equilibrium rule) o the reaction is favoured by high pressure (pressure equilibrium rule, 3 => 2 gas molecules), but only a small increase in pressure is used to give high yields of sulphur trioxide, because the right hand side is energetically very favourable (quite exothermic) o the use of a catalyst ensures a fast reaction without having to use too a higher temperature which would favour the left hand side (energy change equilibrium 3 www.studyguide.pk rule) (3) The sulphur trioxide is dissolved in concentrated sulphuric acid to form fuming sulphuric acid (oleum). o SO3(g) + H2SO4(l) ==> H2S2O7(l) (4) Water is then carefully added to the oleum to produce concentrated sulphuric acid (98%). o H2S2O7(l) + H2O(l) ==> 2H2SO4(l) o If the sulphur trioxide is added directly to water an acid mist forms which is difficult to contain because the reaction to form sulphuric acid solution is very exothermic! Good anti-pollution measures need to be in place since the sulphur oxides are harmful and would cause local acid rain! To help this situation AND help the economics of the process, any unreacted sulphur dioxide is recycled through the reactor. Concentrated sulphuric acid can be used in the laboratory as a dehydrating agent. Dehydration is the removal of water or the elements of water from a compound. o When added to some organic compounds containing hydrogen and oxygen, e.g. sugar, concentrated sulphuric acid removes the elements of water from the compound leaving carbon. o When added to copper sulphate crystals concentrated sulphuric acid removes the water of crystallisation leaving anhydrous copper sulphate What is titanium and how is it produced? Titanium is a very important metal for various specialised uses. It is more difficult to extract from its ore than other, more common metals. Titanium is a transition metal and is strong and resistant to corrosion. o Titanium alloys are amongst the strongest of metal alloys. It is used in aeroplanes, in nuclear reactor alloys and for replacement hip joints. Titanium is extracted from the raw material is the ore rutile which contains titanium dioxide. The rutile titanium oxide ore is heated with carbon and chlorine to make titanium chloride o TiO2 + 2Cl2 + C ==> TiCl4 + CO2 After the oxide is converted into titanium chloride TiCl4 it is then reacted with sodium or magnesium to form titanium metal and sodium chloride or magnesium Chloride. o This reaction is carried out in an atmosphere of inert argon gas so non of the metals involved becomes oxidised by atmospheric oxygen. o TiCl4 + 2Mg ==> Ti + 2MgCl2 or TiCl4 + 4Na ==> Ti + 4NaCl o These are examples of metal displacement reactions eg the less reactive titanium is displaced by the more reactive sodium or magnesium. o Overall the titanium oxide ore is reduced to titanium metal (overall O loss, oxide => metal) 4 www.studyguide.pk The Synthesis of ammonia - The Haber Process Ammonia gas is synthesised in the chemical industry by reacting nitrogen gas with hydrogen gas. The nitrogen is obtained from air (80% N2). The hydrogen is made by reacting methane (natural gas) and water or from cracking hydrocarbons (both reactions are done at high temperature with a catalyst). o CH4 + H2O ==> 3H2 + CO o eg C8H18 ==> C8H16 + H2 The synthesis equation for this reversible reaction is ... N2(g) + 3H2(g) 2NH3(g) .. which means a dynamic equilibrium will form, so no chance of 100% yield! In forming ammonia 92kJ of heat energy is given out (ie exothermic, 46kJ of heat released per mole of ammonia formed). Four moles of 'reactant' gas form two moles of 'product' gas, so there is net decrease in gas molecules on forming ammonia. So applying the equilibrium rules from section 2 the formation of ammonia is favoured by o using high pressure because you are going from 4 to 2 gas molecules (the high pressure also speeds up the reaction because it effectively increases the concentration of the gas molecules). o and low temperature because is an exothermic reaction, o to try to get the optimum conditions to get the biggest yield of ammonia, o these arguments make the point that the yield* of an equilibrium reaction depends on the conditions used. * The word 'yield' means how much product you get compared to the theoretical maximum possible if the reaction goes 100%. In industry pressures of 200 - 300 times normal atmospheric pressure are used in line with the theory. Theoretically a low temperature would give a high yield of ammonia BUT ... o Nitrogen is very stable molecule and not very reactive ie chemically inert. o To speed up the reaction an iron catalyst is used as well as a higher temperature (eg 400-450oC). o The higher temperature is an economic compromise, ie it is more economic to get a low yield fast, than a high yield slowly! o Note: a catalyst does NOT affect the yield of a reaction, ie the equilibrium position BUT you do get there faster! On cooling the reacted mixture the ammonia liquefies and is removed and stored in cylinders. Any unreacted nitrogen or hydrogen is recycled back through the reactor chamber, nothing is wasted! The Uses of Ammonia (a) Ammonia is used to manufacture nitric acid. Ammonia is oxidised with oxygen from air using a hot platinum catalyst to form nitrogen monoxide and water. 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) 5 www.studyguide.pk The gas is cooled and reacted with more oxygen to form nitrogen dioxide. 2NO(g) + O2(g) 2NO2(g) This is reacted with more oxygen and water to form nitric acid. 4NO2(g)+ O2(g) + 2H2O(l) 4HNO3(aq) Nitric acid is used in dye making processes and artificial nitrogenous fertilisers (see below). (b) Ammonia is used to manufacture 'artificial' nitrogenous fertilisers. Ammonia is a pungent smelling alkaline gas that is very soluble in water. The gas or solution turns litmus or universal indicator blue because it is a soluble weak base or weak alkali (more on theory on the Acids, Bases Salts page or on the Extra Aqueous Chemistry page). The fertiliser salts are made by neutralising ammonia solution with the appropriate acid (more method details on Acids, Bases and Salts, but the equations are given below). The resulting solution is heated, evaporating the water to crystallise the salt eg ammonia + sulphuric acid 2NH3(aq) + H2SO4(aq) ammonium sulphate (NH4)2SO4(aq) AND ammonia + nitric acid NH3(aq) + HNO3(aq) ammonium nitrate NH4NO3(aq) These equations are sometimes written in terms of the ficticious 'ammonium hydroxide' which is, as shown above, quite simply an aqueous solution of ammonia, but this is how it looks in some textbooks! About 1% of the dissolved ammonia forms ammonium and hydroxide o ammonium hydroxide + sulphuric acid ammonium sulphate + water o 2NH4OH(aq) + H2SO4(aq) (NH4)2SO4(aq) + 2H2O(l) o ammonium hydroxide + nitric acid ammonium nitrate + water o NH4OH(aq) + HNO3(aq) NH4NO3(aq) + H2O(l) If ammonium salts are mixed with sodium hydroxide solution, free ammonia is formed (detected by smell and damp red litmus turning blue). o eg ammonium chloride + sodium hydroxide ==> sodium chloride + water + ammonia o NH4Cl + NaOH ==> NaCl + H2O + NH3 Artificial fertilisers are important to agriculture and used on fields to increase crop yields but they should be applied in a balanced manner (see below). o Fertilisers usually contain compounds of three essential elements for healthy and productive plant growth to increase crop yield. They replace nutrient minerals used by a previous crop or enriches poor soil and more nitrogen gets converted into plant protein. nitrogen eg from ammonium or nitrate salts like ammonium sulphate, ammonium sulphate or ammonium phosphate or urea phosphorus eg from potassium phosphate or ammonium phosphate potassium eg from potassium phosphate, potassium sulphate. o Fertilisers must be soluble in water to be taken in by plant roots. 6 www.studyguide.pk Problems with using 'artificial' fertilisers Overuse of ammonia fertilisers on fields can cause major environmental problems as well as being uneconomic. Ammonium salts are water soluble and get washed into the groundwater, rivers and streams by rain contaminating them with ammonium ions and nitrate ions. This contamination causes several problems. Excess fertilisers in streams and rivers cause eutrophication. o Overuse of fertilisers results in appreciable amounts of them dissolving in rain water. o This increases levels of nitrate or phosphate in rivers and lakes. o This causes 'algal bloom' ie too much rapid growth of water plants on the surface where the sunlight is the strongest. o This prevents light from reaching plants lower in the water. o These lower plants decay and the active aerobic bacteria use up any dissolved oxygen. o This means any microorganisms or higher life forms relying on oxygen cannot respire. o All the eco-cycles are affected and fish and other respiring aquatic animals die. o The river or stream becomes 'dead' below the surface as all the food webs are disrupted. Nitrates are potentially carcinogenic (cancer or tumor forming). o The presence in drinking water is a health hazard. o Rivers and lakes can be used as initial sources for domestic water supply. o You cannot easily remove the nitrate from the water, it costs too much! o So levels of nitrate are carefully monitored in our water supply. . The Industrial Electrolysis of Sodium Chloride Solution or brine made from concentrated 'rock salt' solution Summary of the ions involved and what happens to them at the two electrodes negative ions from electrode water product H2(g) <= <= H+(aq) ion left => Summary equation: 2NaCl(aq) + 2H2O(l) ==> 2NaOH(aq) + H2(g) + Cl2(g) OH- ions from salt positive electrode product Cl-(aq) => => Cl2(g) Na+ <= ion left When electricity is passed through the sodium chloride solution (brine) there are three products, (1) hydrogen is formed at the negative electrode (-ve cathode), (2) chlorine at the positive electrode (+ve anode) and (3) sodium hydroxide is left in solution. summary equation: sodium chloride + water ==>sodium hydroxide + hydrogen + chlorine 7 www.studyguide.pk The industrial electrodes must be made of an inert material like titanium which is not attacked by chlorine or alkali. However a simple cell using carbon electrodes can be used to demonstrate the industrial process in the laboratory. The cathode gas gives a squeaky pop with a lit splint - hydrogen. The anode gas turns damp blue litmus red and then bleaches it white - chlorine. SODIUM HYDROXIDE NaOH The electrode equation theory and details The (-) cathode attracts the Na+ and H+ ions. The hydrogen ions are reduced by electron (e-) gain to form hydrogen molecules: 2H+(aq) + 2e- ==> H2(g) The (+) anode attracts the OH- and Clions. The chloride ions are oxidised by electron loss to give chlorine molecules: 2Cl-(aq) ==> Cl2(g) + 2e- Sodium hydroxide is used in the manufacture of soaps, detergents, paper, ceramics and to make soluble salts of organic acids with low solubility in water (eg soluble Aspirin). It isn't a halogen compound, but it is made from the electrolysis of salt solution. Chemical Economics The greater the amount of starting materials (reactants) the greater amount of new substances (products) formed. However in the real world chemical processes are not 100% perfectly efficient! o The amount that you actually make is called the yield. o The percentage % yield = actual yield x 100 / predicted yield o The predicted yield assumes there is no loss of product, ie no waste, and the reaction goes 100% in the desired direction. o If no product is obtained then the yield is 0%! o In reality, yields can typically range from 5% to 95% for a variety of chemical processes. Why aren't processes 100% efficient? Typical reasons are: o Loss in filtration of a solid product, ie some may get through as very fine particles or more likely dissolved in the liquid residue. o Loss in evaporation if the product is a volatile liquid. o Loss in transferring liquids, ie traces left on the sides of containers. o The reaction may be an equilibrium, so its impossible to get 100% yield anyway and this means that the yield of an equilibrium reaction depends on the conditions used. The costs of making new substances depends on: o Price of energy (eg gas, electricity). o Starting materials (reactants). o Labour (wages). o Equipment (chemical plant eg machines, reactors, heat transfer systems). o Speed of manufacture (time efficiency). These cost factors can be analysed in more detail eg o The higher the operating pressure of the reactor, the higher the cost. The engineering is more costly due to eg thicker steel reaction vessel, higher health and safety standards require. o The higher the temperature the higher the energy cost. Fortunately this cost is reduced if the reaction is exothermic and the reaction does go faster at higher temperature. 8 www.studyguide.pk o o o o o Time is money! so catalysts save time and money by speeding up the reaction. The rate of reaction must be high enough to give a reasonable yield in reasonable time eg at least within 24 hours for a continuously working plant. Often with equilibrium reactions, it is possible to recycle unreacted starting materials back through the reactor. The % yield must be high enough at least per day, but an initial low yield is quite acceptable if the unreacted starting materials can be recycled many times on a continuous basis through the reactor. Optimum reaction conditions are geared to the lowest cost situation. This often means 'balancing' the rate of reaction versus the highest % yield. It is often best to get a low yield fast and recycle! Automating the chemical plants with sensors, controls, computer software etc. significantly reduces the wages bill. DONE 9 www.studyguide.pk Introduction to Oxidation and Reduction OXIDATION and REDUCTION - REDOX REACTIONS OXIDATION - definition and examples REDUCTION - definition and examples (a) The gain or addition of oxygen by an atom, molecule or ion eg ... (b) The loss or removal of oxygen from a compound etc. eg ... (1) S ==> SO2 [burning sulphur - oxidised] (1) CuO ==> Cu [loss of oxygen from copper(II) oxide to form copper atoms] (2) CH4 ==> CO2 + H2O [burning methane to water and carbon dioxide, C and H gain O] (2) Fe2O3 ==> Fe [iron(III) oxide reduced to iron] (3) NO ==> NO2 [nitrogen monoxide oxidised to nitrogen dioxide] (3) NO ==> N2 [nitrogen monoxide reduced to nitrogen] (4) SO32- ==> SO42- [oxidising the sulphite ion to the sulphate ion] (4) SO3 ==> SO2 [sulphur trioxide reduced to sulphur dioxide] (c) The loss or removal of electrons from an atom, ion or molecule eg (d) The gain or addition of electrons by an atom, ion or molecule eg ... (1) Fe ==> Fe2+ + 2e- [iron atom loses 2 electrons to form the iron(II) ion] (1) Cu2+ + 2e- ==> Cu [the copper(II) ion gains 2 electrons to form neutral copper atoms) (2) Fe2+ ==> Fe3+ + e- [the iron(II) ion loses (2) Fe3+ + e- ==> Fe2+ [the iron(III) ion gains an 1 electron to form the iron(III) ion] electron and is reduced to the iron(II) ion] (3) 2Cl- ==> Cl2 + 2e- [the loss of electrons by chloride ions to form chlorine molecules] (3) 2H+ + 2e- ==> H2 [hydrogen ions gain electrons to form neutral hydrogen molecules] (e) An oxidising agent is the species that gives the oxygen or removes the electrons (f) A reducing agent is the species that removes the oxygen or acts as the electron donor REDOX REACTIONS - in a reaction overall, oxidation and reduction must go together (g) Redox reaction analysis based on the oxygen definitions (1) copper(II) oxide + hydrogen ==> copper + water o CuO(s) + H2(g) => Cu(s) + H2O(g) o copper oxide reduced to copper, hydrogen is oxidised to water o hydrogen is the reducing agent (removes O from CuO) o copper oxide is the oxidising agent (donates O to hydrogen) (2) iron(III) oxide + carbon monoxide ==> iron + carbon dioxide o Fe2O3(s) + 3CO(g) => 2Fe(l) + 3CO2(g) o the iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide o CO is the reducing agent (O remover from Fe2O3) o the Fe2O3 is the oxidising agent (O donator to CO)] (3) nitrogen monoxide + carbon monoxide ==> nitrogen + carbon dioxide o 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g) o nitrogen monoxide is reduced to nitrogen o carbon monoxide is oxidised to carbon dioxide o CO is the reducing agent and NO is the oxidising agent www.studyguide.pk (4) iron(III) oxide + aluminium ==> aluminium oxide + iron (the thermit reaction) o Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s) o iron(III) oxide is reduced and is the oxidising agent o aluminium is oxidised and is the reducing agent (h) Redox reaction analysis based on the electron definitions (1) magnesium + iron(II) sulphate ==> magnesium sulphate + iron o Mg(s) + FeSO4(aq) => MgSO4(aq) + Fe(s) o this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below. o The sulphate ion SO42-(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!! o Mg(s) + Fe2+(aq) ==> Mg2+(aq) + Fe(s) o the magnesium atom loses 2 electrons (oxidation) to form the magnesium ion, the iron(II) ion gains 2 electrons (reduced) to form iron atoms. o Mg is the reducing agent (electron donor) and the Fe2+ is the oxidising agent (electron remover or acceptor) o Displacement reactions involving metals and metal ions are electron transfer reactions. (2) zinc + hydrochloric acid ==> zinc chloride + hydrogen o Zn(s) + 2HCl(aq) => ZnCl2(aq) + H2(g) o the chloride ion Cl- is the spectator ion o Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(g) o Zinc atoms are oxidised to zinc ions by electron loss, so zinc is the reducing agent (electron donor) o hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules (3) copper + silver nitrate ==> silver + copper(II) nitrate o Cu(s) + 2AgNO3(aq) ==> 2Ag + Cu(NO3)2(aq) o the nitrate ion NO3- is the spectator ion o Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq) o copper atoms are oxidised by the silver ion by electron loss o electrons are transferred from the copper atoms to the silver ions, which are reduced o the silver ions are the oxidising agent and the copper atoms are the reducing agent (4) iron(II) chloride + chlorine ==> iron(III) chloride (5) halogen (more reactive) + halide salt (of less reactive halogen) ==> halide salt (of more reactive halogen) + halogen (less reactive) o X2(aq) + 2KY(aq) ==> 2KX(aq) + Y2(aq) o X2(aq) + 2Y-(aq) ==> 2X-(aq) + Y2(aq) o where halogen X is more reactive than halogen Y, F > Cl > Br > I) o X is the oxidising agent (electron acceptor) o KY is the reducing agent (electron donor) (6) Electrode reactions in electrolysis are electron transfer redox changes o at the negative cathode positive ions are attracted: metal ions are reduced to the metal by electron gain: Mn+ + n e- ==> M n = the numerical charge of the ion and the number of electrons transferred or 2H+(aq) + 2e- ==> H2(g) (for the discharge of hydrogen) 2 www.studyguide.pk o at the positive anode negative ions are attracted: negative non-metal ions are oxidised by electron loss eg for oxide ions: 2O2- - 4e- ==> O2 or 2O2- ==> O2 + 4e for hydroxide ion: 4OH- - 4e- ==> O2 + 2H2O or 4OH- ==> O2 + 2H2O + 4e for halide ions (X = F, Cl, Br, I): 2X- - 2e- ==> X2 or 2X- ==> X2 + 2e- Miscellaneous Redox Notes Redox changes can often be observed as significant colour changes eg o iron + copper(II) sulphate ==> iron(II) sulphate + iron Mg(s) + FeSO4(aq) => MgSO4(aq) + Fe(s) Mg(s) + Fe2+(aq) ==> Mg2+(aq) + Fe(s) Sulphate, SO42-(aq), is colourless BUT a blue to pale green colour change is observed in the solution as the blue copper(II) ion is replaced by the pale green iron(II) ion. o Potassium manganate(VII) is a powerful oxidising agent and an intense purple colour in water due to the MnO4- ion. In acidified solution it changes to an almost colourless* manganese(II) ion, Mn2+ when it oxidises something (* which actually is a very pale pink transition metal ion). o Potassium dichromate(VI) is another strong oxidising agent and is orange due to the dichromate(VI) ion, Cr2O72- ion. When it oxidises something it changes to the green chromium(III) ion, Cr3+. o Potassium iodide is a colourless salt dissolving in water to form a colourless solution. If it is oxidised eg with chlorine a yellow=>orange==>brown colour develops as iodine is formed from the colourless iodide ion. The use of Roman Numerals in names: o This indicates what is called the oxidation state of an atom in a molecule or ion. o It is easy to follow for simple metal ions because it equals the charge on the ion eg the oxidation state of copper in the copper(II) ion is referred to as +2 the more electrons removed from the atom or ion by oxidation, the higher its oxidation state eg Fe2+ - e- ==> Fe3+, gives iron the oxidation state of +3 in the iron(III) ion (via a suitable oxidising agent). but for more complex ions things are not so simple. in manganate(VII) ion, the Mn is in the +7 oxidation state in dichromate(VI) ion, the Cr is in the +6 oxidation state o This topic is dealt with at AS-A2 advanced level chemistry (there is an introduction on another web page) but not for KS4-GCSE! Oxidation Number Oxidation numbers are a useful tool for determining whether a substance has been oxidized or reduced. An element that undergoes a change in oxidation number in the course of a reaction has been oxidized or reduced. Let's learn how to assign oxidation numbers. 3 www.studyguide.pk Assigning Oxidation Numbers 1. At atom in its elemental state has an oxidation number of 0. Na H2 Cl2 S Xe Each atom in these elements has an oxidation number of 0. 2. An atom in a monoatomic ion has an oxidation number identical to its charge. Na+ Ba2+ Al3+ Br– S2– +1 +2 +3 –1 –2 The oxidation number is equal to the charge on the monoatomic ion. 3. An atom in a polyatomic ion or a molecular compound usually has the same oxidation number it would have if it were in a monoatomic ion. a. Elements to the left on the periodic table are "cationlike" and have positive oxidation numbers. b. Elements to the right on the periodic table are "anionlike" and have negative oxidation numbers. Consider NH3. N has an oxidation number of –3; each H has an oxidation number of +1. c. Hydrogen has a +1 oxidation number when bonded to nonmetals, and has a –1 oxidation number when bonded to a metal. NaH (H –1 oxidation number) H2O (H +1 oxidation number) d. Oxygen often has a –2 oxidation number, but can have a –1 oxidation number in the peroxide ion, O22–. H2O (O –2 oxidation number) 4 www.studyguide.pk HOOH (O –1 oxidation number) e. Halogens usually have an oxidation number of –1, Unless bonded to oxygen, when they have a positive oxidation number. HCl (Cl –1 oxidation number) HOCl (Cl +1 oxidation number) 4. The sum of the oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. Oxidizing and Reducing agents Oxidation and reduction always occur together. Whenever one atom loses electrons (is oxidized), another atom must gain those electrons (be reduced). The reactants can be classified as either a reducing agent or an oxidizing agent. Reducing agent causes reduction loses one or more electrons is oxidized oxidation number of atom increases Oxidizing agent causes oxidation gains one or more electrons is reduced oxidation number of atom decreases 5 www.studyguide.pk The Activity Series of the Elements The elements at the top of the table readily give up electrons and are stronger reducing agents. The elements at the bottom give up electrons less readily and are weaker reducing agents. Any element higher in the activity series will react with the ion of any element lower in the activity series. Some Applications of Redox Reactions A vast number of redox reactions occur in industrial and biological processes. A few are summarized here. 1. Combustion is the burning of fuel by oxidation with oxygen in air. Fuels include natural gas, wood, paper, and other organic substances composed of carbon and hydrogen. Some metals also burn in air. CH4(q) + 2 O2(g) CO2(g) + 2 H2O(l) 2Mg(s) + 2 O2(g) 2 MgO(s) 2. Bleaching is the use of redox reactions to decolorize or lighten colored materials. Oxidizing agents used in bleaching include hydrogen peroxide (H2O2) and sodium hypochlorite (NaClO). 3. Batteries are all based on redox reactions. 4. Metallurgy is the science of extracting and purifying metals from their ores. 6 www.studyguide.pk 5. Undesirable oxidation reactions are termed corrosion. The rusting of iron in moist air is a familiar process with enormous economic impact. 4 Fe(s) + 3 O2(g) Fe2O3H2O(s) 6. Respiration is the process of breathing and using oxygen for the many biological redox reactions that occur in living organisms. C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) + energy DONE 7 www.studyguide.pk ORGANIC CHEMISTRY OIL and its many useful PRODUCTS The origin of oil Crude oil is formed from organic material of the remains of plant and animal organisms that lived millions of years ago. These remains form sediments eg at the bottom of seas, and become buried under layers of sedimentary rock. They decay, without air (oxygen), under the action of heat and pressure to form crude oil over millions of years. It is a fossil fuel because it is formed from once living organisms and the Sun is the original source of energy. It is a non-renewable and finite (limited reserves) energy resource because it takes millions of years to form and we burn it faster than its is formed! It is also known as a finite energy resource because it will eventually run out! We do not have unlimited oil reserves! Coal and natural gas (mainly methane CH4 and often found with oil) are also non-renewable fossil fuels formed from the remains of plants or animals. When the fossil fuels are burned the 'carbon', as carbon dioxide, is returned to the living environment, gets used up in photosynthesis, the plant material decays or is eaten by animals, so completing the carbon cycle. The SEPARATION of the crude oil mixture into fractions and the USES of these fractions A fraction is a mixture of a restricted boiling point range of molecules, they have a similar number of carbon atoms and physical properties. The uses of the fractions depend on their physical and chemical properties. www.studyguide.pk Hydrocarbon molecules are only made of a chemical combination of carbon and hydrogen atoms. They are compounds because they consist of atoms of atleast two different elements. THE FRACTIONAL DISTILLATION OF CRUDE OIL Crude oil is a complex mixture of mainly hydrocarbon compound molecules. A mixture consists of two or more elements or compounds which are not chemically combined. The chemical properties of each substance in the mixture is unchanged. This means crude oil can be separated by physical methods, in this case fractional distillation, because they have different boiling and condensation points. The most volatile fraction, ie with the lowest boiling point, boils or evaporates off first and goes to the top of the column. The rest separate out according to their boiling point so that the highest boiling fraction, ie the less volatile with higher boiling points, tend to condense more easily lower down the column. The bigger the molecule, the greater the intermolecular forces, so the higher the boiling point. Chemical bonds are not broken in the process, only the intermolecular force of attraction. names C atoms boiling of in the range fractions molecule in oC USES of the fraction methane gas fuel, C3-4 easily liquefied, portable energy source bottled gas for cooking (butane), higher pressure cylinders (propane) Fuel Gas, LPG, Refinery Gas 1 to 4 -160 to 20oC Gasoline, Petrol 5 to 11 20 to 60oC easily vaporised, highly flammable, easily ignited, car fuel Naphtha 7 to 13 60 to 180oC no good as a fuel, but valuable source of organic molecules to make other things, cracked to make more petrol and alkenes Paraffin, Kerosene 10 to 16 120 to 240oC less flammable than petrol, domestic heater fuel, jet fuel Diesel oil, Gas oil 15 to 25 220 to 250oC Car and larger vehicle fuel Fuel and lubricating 20 to 70 Oils and Waxes Bitumen over 70 not so easily evaporated, not as flammable, safe to store for 250 to central heating oil, quite viscous 350oC (sticky) and can also be used for lubricating oils, clear waxes and polishes forms a thick, black, tough and resistant adhesive on cooling, over used as waterproofing material o 350 C and to sticks rock chips on roofs or road surfaces 2 www.studyguide.pk More on examples of relating the physical properties of the fractions to their uses and dangers down the list the molecule gets bigger, more viscous, higher boiling and less flammable The refinery gas fractions, under pressure, are conveniently pumped to burner systems, but are easily ignited and explosive. Vehicle fuels must be liquid for compact and convenient storage but they must be easily vapourised to mix with air in the engine prior to ignition. The ease of vaporisation does however make them flammable! Paraffin and kerosine are less flammable and safer, but not as easily ignited. Fuel oil is not too viscous to pump to a central heating burner, and it is not very volatile and so not as flammable and dangerous as petrol or diesel etc. for domestic use. Lubricating oil must be quite viscous to stick onto surfaces. Smaller molecules might be more runny but they would evaporate away! It is also water repellent and helps reduce corrosion on moving machine parts. Candle wax is very convenient as a solid for humble lamp (especially in power cuts!), but via a wick, the heat from the flame is sufficient to vaporise the hydrocarbons to burn them. Bitumen is a water repellent solid at room temperature but is readily melted (sometimes too easily in hot weather). Used as base for a road chipping top surface or sometimes directly. It is also used to waterproof roofing felt. 3 www.studyguide.pk 4 www.studyguide.pk The ALKANE series of hydrocarbons Alkanes are a group of hydrocarbon molecules in which all the carbon and hydrogen atoms are only joined by single covalent bonds (eg C-H or C-C). Alkanes are known as saturated molecules because other atoms cannot add to them (compare alkenes further on). The first four in the series are shown. They are not very reactive unless burned! (1) is the molecular formula: a summary of the totals of each atoms of each element in one molecule; (2) is a 'shorthand' version of the structural formula (3); (3) is called the structural formula: it shows how all the atoms are linked with the covalent bonds representation of the structural formula (3) -; (4) is a '3D' methane (1) , (3) (main molecule in natural gas) , (4) ethane (1) (2) (3) (4) propane (1) (2) (3) butane (1) (2) (3) 5 www.studyguide.pk 6 www.studyguide.pk The complete combustion of hydrocarbons When hydrocarbons are burned in air a fast exothermic reaction occurs releasing heat and forming carbon dioxide and water. It is an oxidation reaction due to O gain by C and H. The carbon dioxide is chemically detected with limewater - with which it forms a white precipitate (milky appearance) of calcium carbonate. The water is chemically detected either by (i) anhydrous white copper sulphate turning blue OR (ii) anhydrous blue cobalt chloride paper turning pink. A physical test for water is to measure its boiling point (should be 100oC). Equations for the complete combustion of a hydrocarbon hydrocarbon + oxygen => carbon dioxide + water eg word equation: methane + oxygen => carbon dioxide + water and the symbol equation: CH4(g) + 2O2(g) => CO2(g) + 2H2O(l) (one CO2 for every C and one H2O for every two H's in the hydrocarbon molecule) for ethane the symbol equation (more awkward) is ... 2C2H6(g) + 7O2(g) => 4CO2(g) + 6H2O(l) and for pentane the symbol equations is ... C5H12(l) + 7O2(g) => 5CO2(g) + 6H2O(l) 7 www.studyguide.pk The Incomplete Combustion of hydrocarbons CO If there is not enough oxygen present to completely burn the fuel to carbon dioxide and water other products may form causing pollution and fuel inefficiency. The most common partially burned products are likely to be carbon C (soot) and deadly carbon monoxide CO. It would appear that the hydrogen in the fuel molecules is more easily burned and usually forms water. There is also less heat released compared to complete combustion. o eg CH4(g) + O2(g) => C(s) + 2H2O(l) o or 2CH4(g) + 3O2(g) => 2CO(g) + 4H2O(l) Therefore it is extremely important that any combustion system is as efficient as possible eg gas heaters, furnaces etc. must all have excellent ventilation for complete combustion to harmless water and carbon dioxide. If there is any smell of gas, make sure (i) all appliances are turned off, (ii) all sources of ignition are absent, and (iii) ring the gas board! Faulty gas appliances have led to tragic deaths. Carbon monoxide is colourless and odourless and even low concentrations in the air can be fatal. Carbon monoxide is unfortunately emitted by all car exhausts, though catalytic converters help reduce this by converting nitrogen monoxide (another pollutant) and carbon monoxide into harmless nitrogen and carbon dioxide. o 2NO(g) + 2CO(g) => N2(s) + 2CO2(l) What makes a good fossil fuel? Factors that should be taken into consideration Energy value: eg kJ of heat energy released per kg; Availability: Geographical convenience, oil production levels; Storage: Health and safety issues eg coal very safe, natural gas more dangerous Cost: Extraction, transport, market price Toxicity and Pollution: Greenhouse effect (which produces the least or most CO2/energy released?); sulphur content of fuel (most removed before fuel used to minimise sulphur dioxide and acid rain formation); efficiency of combustion eg minimum carbon monoxide and soot levels Ease of use: 8 www.studyguide.pk The ALKENE hydrocarbons series Alkenes are hydrocarbons containing a double bond as well as single bonds. These are called unsaturated molecules because two atoms can join onto the bond when it opens up. The first two in the series are shown below. They are extremely reactive and important compounds in the chemical industry and are converted into very useful compounds eg plastics. They are made from cracking processes (see below) 1) is the molecular formula: a summary of the totals of each atoms of each element in one molecule; (2) is a 'shorthand' version of the structural or displayed formula (3); (3) is called the structural or displayed formula: it shows how all the atoms are linked with the covalent bonds - ethene (1) , (2) , (3) propene (1) (2a) (2b) (1) (3) butene , (2) A test to distinguish between ALKANE and ALKENE hydrocarbons Hydrocarbons are colourless. Bromine dissolved in water or trichloroethane solvent forms an orange (yellow/brown) solution. When bromine solution is added to both an alkane or an alkene the result is quite different. The alkane solution remains orange - no reaction. However, the alkene decolourises the bromine as it forms a colourless dibromo-alkane compound - see equations .... or Alkenes are unsaturated molecules, atoms can add to them via the C=C double bond, so a reaction occurs. Alkanes are saturated - no double bond - and atoms cannot add - so no reaction. ethene + bromine ==> 1,2-dibromethane 9 www.studyguide.pk .... or propene + bromine ==> 1,2dibromopropane CRACKING a problem!! there isn't enough petrol in crude oil! AND crude oil doesn't have any alkenes in it for plastics! When crude oil has been distilled into useful fractions it is found that the quantities produced do not match the ratio required for commercial needs eg we have an insatiable appetite for petrol and diesel in our cars and there are two many left-overs of the larger molecules which do not make good fuels or have other uses. Fuel oil, naphtha and bitumen in crude oil exceed demand. Also, alkenes are not found in crude oil and they one of the most valuable types of organic molecule in the chemical industry eg to make polymers (plastics) or ethanol (an alcohol). The two deficiencies are remedied by the process of cracking which converts useless big molecules into useful smaller ones. CRACKING is done by heating some of the less used fractions to a high temperature vapour and passing over a suitable hot catalyst. The cracking reaction is an example of thermal decomposition - a reaction that breaks down molecules into smaller ones using heat. The main products from cracking alkanes from oil are smaller alkanes (eg for petrol or diesel) and alkenes (eg for plastics). The equations below illustrate the process, small molecules are used to show the overall molecular change clearly BUT in practice the 'starter' molecules are likely to be more those shown in equations (3) and (4). The cracking involves breaking single carbon-carbon bonds to form the alkanes (saturated hydrocarbons) and alkenes (unsaturated hydrocarbons) products (1) butane ethane ethene.....or 10 www.studyguide.pk (2) butane methane propene lots and lots of other possibilities! eg (3) (4) C8H18 ===> C6H14 + C2H4 (making ethene) C12H26 ===> C9H20 + C3H6 (making propene) The formation of POLYMERS and the USES of PLASTICS - Macromolecules The formation of big polymer molecules called polyalkenes from small molecules called alkenes When catalysed and heated under pressure, alkenes link together when the double bond opens. The spare bonds are used to join up the molecules. The original small molecule is called the monomer and the long molecule is called the polymer, which is the sort of molecule most plastics consist of. The polymer is now a saturated molecule but has the same C:H ratio as the original alkene. So lots of small molecules join up to form a big long molecule in a process called addition polymerisation and the polymers are named as poly(name of original alkene). Examples of Polymer Molecules - formation and uses Poly(ethene) from ethene is a cheap but very useful plastic used for plastic bags and bottles. 11 www.studyguide.pk Poly(propene) from propene is stronger and more hard wearing than polythene and is used for making crates, fibres and ropes. Poly(chloroethene), old name PVC, from chloroethene (vinyl chloride) is tougher than poly(ethene) and very hard wearing and has good heat stability, so is used for covering electrical wiring and plugs. It also replacing metals as gas and water drain pipes and has found a use as artificial leather and readily dyed to bright colours! Polystyrene is made from styrene (another alkene monomer) and is used in a gas expanded form for packaging and insulation. 12 www.studyguide.pk Three problems associated with using Polymers or Plastics Polymers or plastics cannot be easily broken down by microorganisms ie they are not biodegradable. This leads to waste disposal problems and 'non-rotting' litter around the environment and land-fill sites are getting full! When plastic materials burn they can produce highly toxic gases such as carbon monoxide, hydrogen cyanide and hydrogen chloride (particularly from PVC and other plastics containing chlorine and nitrogen). This has caused deaths in house fires and controversial problems with alleged inefficient waste incinerators and they will generally cause environmental problems if burning on waste tips etc. It is difficult to recycle plastics because of separation into the various sorts. It would be beneficial to prolong the life of the finite crude oil reserves AND reduce pollution and space in land-fill sites. New plastics are being developed which are more biodegradable or can be recycled, so will the paper bag and cardboard package make a comeback? (in Eire you have to bring your own bag or buy one, and not necessarily a plastic one!). 13 www.studyguide.pk Oil Products and Environment Problems Our economy, like many other countries has become very dependent on the extraction, sale and use of oil based products. BUT, there is high price to be paid at times whether it be pollution effects or warring countries with oil economics factors. Oil rig accidents, broken pipelines, oil tanker wrecks etc. all have terrible effects on the plant and animal life of the locality as we see from the horrible TV pictures of seabirds coated in oil, and toxic oil slicks covering the beaches and sands. The burning of oil and other fossil fuels is contributing to the 'Greenhouse Effect' of global warming. The extra carbon dioxide forming absorbs and traps sunlight (or more precisely the re-radiated sunlight energy from the Earth's surface) rather like a greenhouse. The effects are predicted to be dramatic eg rising sea levels as polar ice melts causing flooding in low lying land regions, more energy in the global weather system leads to more frequent violent weather patterns, .. Fossil fuels contain the element sulphur or compounds of sulphur. When the fuel is burned the sulphur also burns to form sulphur dioxide. This is an acidic gas and dissolves in rainwater, it then reacts with water and oxygen to form a very dilute solution of sulphuric acid. o Sulphur dioxide is a harmful gas and lung irritant and contributed to 5000 extra deaths in the great 'London Smog' in the 1950's. o The formation of acid rain has several bad effects on the environment eg the low pH causes plant damage, particularly trees, kills certain life forms and so damages eco cycles and food chains in rivers or lakes harming wildlife like trout, increases the 'weathering' corrosion rates of building stone (particularly limestone). High temperature combustion also produces other pollutants including ... o Nitrogen oxides, NO and NO2, which are acidic and contribute further to acid rain (above), and are also involved in the chemistry of 'photochemical smog' - which produces chemicals harmful to respiration and eyes. Many of the reactions are initiated by sunlight. o Carbon monoxide CO, which is toxic, and also involved in the chemistry of 'photochemical smog'. This is formed by inefficient combustion o Unburned hydrocarbons, CxHy, which can be carcinogenic and are also involved in photochemical smog chemistry. But catalytic converters* can significantly reduced these three unwanted emissions (see above for CO and NO removal, and CxHy gets oxidised to CO2 and H2O). * eg using platinum-rhodium transition metal catalysts, these are dispersed on ceramic bed to give a big surface area for the best reaction rate. There are other indirect pollution problems to do with burning fossil fuels: o Lead compounds are added to petrol to improve engine performance. This produces lead compound emissions into the environment. Lead compounds are nerve toxins so it is fortunate they are being phased out in many countries. o Photochemical smog is mentioned in the previous paragraph. And finally, should we using a very valuable source of organic chemicals by merely burning most of it? AND how long will oil reserves last? AND what happens if the oil runs out? Hydrogen gas can be used as fuel and a long-term possible alternative to fossil fuels. o It burns with a pale blue flame in air reacting with oxygen to be oxidised to form water. o hydrogen + oxygen ==> water or 2H2(g) + O2(g) ==> 2H2O(l) o It is a non-polluting clean fuel since the only combustion product is water and so its use would not lead to all environmental problems associated with burning fossil fuels. o It would be ideal if it could be manufactured by electrolysis of water eg using solar cells. o Hydrogen can be used to power fuel cells on the "Extra Electrochemistry" page. 14 www.studyguide.pk n-butane Isobutane Neopentane n-butane Isobutane Neopentane Isopentane n-pentane n-pentane Isopentane 15 www.studyguide.pk Extra ORGANIC CHEMISTRY including introductory aspects of food and drug chemistry on other pages .. [basic Oil Products notes - if not on this page] [enzymes] and on this page .. [1. combustion] [2. alkane/alkene 'families'] [3. alcohols like ethanol] [4. other organic molecule families eg carboxylic acids, esters, polymers] [5. Natural molecules - carbohydrates, sugars, amino acids, proteins, fats, oils etc.] [6. Vitamin C, Drugs and Food Additives] [7. CFC's, Ozone and free radicals] [multiple choice quiz - still under development] and [email query or comment] KEYWORDS: .. addition(alkene reactions) .. alcohols .. alkanes .. alkenes .. amino acids(from hydrolysis) .. analgesics .. aspirin (soluble/insoluble/action) .. carbohydrates .. carboxylic acids .. cellulose .. CFC's .. colourings .. combustion .. cracking .. disaccharide .. drugs .. E numbers .. esters .. ethanoic acid .. ethanol .. fats (saturated/unsaturated) .. fibre .. flavourings .. food additives .. fractional distillation .. free radicals .. fuel .. glucose .. glycerol .. glycine .. homologous series .. hydrolysis(starch, protein, fats) .. ibuprofen .. isomerism/isomers .. macromolecule .. margarine .. monosaccharide .. Nylon .. oils(and fats) .. organic .. ozone .. paracetamol .. plastics/polymers(burning, structure) .. polyamide .. polypeptides .. polysaccharides .. proteins .. preservatives .. saturated .. soaps .. starch .. sucrose .. sugars(from hydrolysis, cyclic structure) .. sweeteners .. Terylene .. thermosoftening .. thermosetting .. triglycerides .. unsaturated .. vinegar .. vitamin C .. [email comment or query] 1. What is produced when organic compounds are burned? Some organic compounds are used as fuels. Other organic compounds, including plastics, are burned as waste. Burning these organic compounds releases gases into the atmosphere. All organic compounds consist partly of carbon atoms. Coal, crude oil, natural gas (methane) and wood contain organic compounds o all are used as fuels, either directly like coal or natural gas, o or indirectly as coke from coal or petrol from crude oil etc., o and apart from wood, they are finite (limited reserve) fossil (from decayed organic material) fuels. Many hydrocarbons are fuels ie a substance burned to release heat energy. When organic compounds are burned in a plentiful supply of air the carbon is oxidised to carbon dioxide and the hydrogen is oxidised to water. In a limited supply of air incomplete combustion occurs forming carbon monoxide and/ or carbon. Carbon monoxide is poisonous because it reduces the capacity of blood to carry oxygen. Combustion equations and tests for combustion products are all on the Oil Notes web page. Each fossil fuel has a different cost, efficiency and cleanliness on burning. Generally speaking in 'cleanliness' the order is methane (natural gas) > alkanes in petrol > heavy oil and from left to right there is also an increase in C/H atom ratio in the molecule so more CO2 produced too. Some notes on other fuels (but they are designed for more advanced level courses) and a fossil fuel survey on Oil Products Notes page The combustion of plastics (and other organic compounds) which contain chlorine and nitrogen produce poisonous fumes when burnt eg choking hydrogen chloride HCl and toxic hydrogen cyanide HCN respectively. Especially where there is a limited supply of air. The combustion products of carbon (toxic CO and CO 2) and hydrogen (H2O) are also formed. Hydrogen gas can be used as fuel. o It burns with a pale blue flame in air reacting with oxygen to be oxidised to form water. o hydrogen + oxygen ==> water or 2H2(g) + O2(g) ==> 2H2O(l) o It is a non-polluting clean fuel since the only combustion product is water and so its use would not lead to all environmental problems associated with burning fossil fuels. o It would be ideal if it could be manufactured by electrolysis of water eg using solar cells. o Hydrogen can be used to power fuel cells on the "Extra Electrochemistry" page. 16 www.studyguide.pk 2. Why are there families of organic compounds? [alkanes and alkenes are introduced on the Oil Notes page and give details of (i) alkane combustion (ii) the reaction of bromine with alkenes and (iii) the basics of alkene polymerisation] 2a INTRODUCTION Organic compounds belong to different families. The compounds in each family have a similar chemical structure and a similar chemical formula. Each family of organic compounds forms what is called a homologous series. Different families arise because carbon atoms readily join together in chains (catenation) and strongly bond with other atoms such as hydrogen, oxygen and nitrogen. The result is a huge variety of 'organic compounds'. The name comes from the fact that most of the original organic compounds studied by chemists came from plants or animals. A homologous series is a family of compounds which have a general formula* and have similar chemical properties because they have the same functional group of atoms (eg C=C alkene, C-OH alcohol or -COOH carboxylic acid). o * Match the general formula pattern with the alkane and alkene examples shown below. members of a homologous series have similar physical properties such as appearance, melting/boiling points, solubility etc. but show trends in them eg steady increase in melting/boiling point with increase in carbon number or molecular mass. The molecular formula represents a summary of all the atoms in the molecule (see examples below). The structural or displayed formula shows the full structure of the molecule with all the individual bonds and atoms shown (though there are different 'sub-styles' of varying detail, see below). 2b ALKANES These are obtained directly from crude oil by fractional distillation (see oil notes). The saturated hydrocarbons form an homologous series called alkanes with a general formula CnH2n+2 Saturated means the molecule has no C=C double bonds, only carbon-carbon single bonds, and so has combined with the maximum number of atoms ie no atoms can add to it. The alkanes don't really have a functional group and have quite a limited chemistry BUT they are still a clearly defined homologous series. Alkane examples: The gases: methane CH4, ethane C2H6, propane C3H8, butane C4H10, liquids: pentane C5H12, hexane C6H14 etc. The first four alkane structures are shown on the oil notes page. Names end in ...ane Carbon always forms 4 bonds with other atoms and hydrogen 1 bond with other atoms eg Propane: molecular formula C3H8, structural and displayed formula styles include ... o or or Isomerism occurs when two or more compounds have the same chemical formula but have different structures. eg for the molecular formula C4H10 there are two possibilities - one 'linear' and one with carbon chain 'branching' ... o butane: or or 17 www.studyguide.pk o and its isomer methylpropane: or or Can you work out the structures of the 3 isomers of C5H12 ? (you will find enough to work out the answers on the advanced level page ALKANES) Isomers show variation in physical properties which depend upon the strength of the intermolecular forces. Intermolecular forces are due to weak electrical attractive forces that exist between all molecules. (a) For a homologous series the strength of intermolecular forces increases as the carbon chain length increases (b) For isomers (same C number), the forces decrease as the amount of chain branching increases. This is because the attractive forces are a function of the potential surface-surface contact ie the compactness of the molecules. o (a) as the chain length increases the surface-surface contact must increase per molecule, o (b) for isomers, with more branching, the chain length decreases and the molecule is more 'compact' reducing the surface-surface contact per molecule. For example ... o (a) from methane ==> ethane ==> propane ==> petrol ===> oils ==> grease ==> waxes etc. the boiling point rises and so does the viscosity (stickiness!) as the carbon chain length increases (trend also indicated by gases ==> liquids ==> solids). o (b) 'linear' butane has a higher boiling point than 'branched' methylpropane (diagrams above). Alkanes and alkenes undergo combustion reactions (see Oil). They are not very reactive unless burned! BUT they will react with strong oxidising chemicals like chlorine when heated or subjected to uv light (you need something to initiate the reaction). o A substitution reaction occurs and a chloro-alkane is formed eg o a hydrogen is swapped for a chlorine and the hydrogen combines with a chlorine atom o ethane + chlorine ==> chloroethane + hydrogen chloride o C2H6 + Cl2 ==> C2H5Cl + HCl + Cl2 ==> o + HCl 2c ALKENES These cannot be obtained directly from crude oil and must be made by cracking (see oil notes). The unsaturated hydrocarbons form an homologous series called alkenes with a general formula CnH2n Unsaturated means the molecule has a C=C double bond to which atoms or groups can add. Alkene examples: Names end in ...ene o ethene C2H4, o propene C3H6, or or or o butene or The alkenes are more reactive than alkanes because of the presence of the carbon=carbon double bond. The alkenes undergo addition reactions in which one of the carbon=carbon double bonds breaks allowing each carbon atom to form a covalent bond with another atom such as hydrogen or bromine. Examples of addition reactions are: with hydrogen under pressure and in the presence of a nickel catalyst to form an 18 www.studyguide.pk alkane o + H2 ==> (ethene + hydrogen ==> ethane) o + H2 ==> (propene + hydrogen ==> propane) Alkenes react by 'addition' with bromine and decolourises the orange bromine water because the organic product is colourless, and this is a simple test to distinguish an alkene from an alkane. o see Oil Notes for equations for ethene and propene Vegetable oils contain unsaturated fats and can be hardened to form margarine by adding hydrogen on to some of the carbon=carbon double bonds using a nickel catalyst. Alkenes can add to themselves by addition polymerisation to form 'plastic' or polymeric materials (see below or oil notes) Alkenes are isomeric with cycloalkanes eg C6H12 can be hexene or cyclohexane o o hexene CH3-CH2-CH2-CH2-CH=CH2 or cyclohexane and note that .... hexene is an unsaturated hydrocarbon with a double bond, the isomeric cyclohexane does not have a double bond and is a saturated hydrocarbon, so a simple bromine test could distinguish the two similar colourless liquids, because only the hexene would decolorise the bromine water test reagent. 3. What is ethanol and how can we make it? What we call alcohol in everyday life is a substance whose chemical name is ethanol. Ethanol is just one member of a family of substances called alcohols which have a C-OH functional group in their structure. Ethanol structure: or or or Ethanol is used as a solvent, as a fuel (can be mixed with petrol), and used to make 'ethyl esters' (see below) as well as the 'potent' chemical present in alcoholic drinks! o The % alcohol in wines, spirits and beer varies from 1-40% o There are health and social issues about the medical and behavioral aspects of alcohol consumption. o Methylated spirit is mainly ethanol but poisonous and nasty tasting chemicals like methanol are added so it is not used as a beverage! Ethanol can be produced by fermentation of sugars. The raw materials are mixed with water and yeast at just above room temperature. The yeast contains enzymes which are biological catalysts. The sugars react to form ethanol and carbon dioxide. The carbon dioxide is allowed to escape and air is prevented from entering the reaction vessel to stop oxidation of ethanol to ethanoic acid ('acetic acid' or vinegar!). When the reaction is over the ethanol is separated from the reaction mixture by fractional distillation. o eg glucose ==enzyme==> ethanol + carbon dioxide 19 www.studyguide.pk C6H12O6(aq) ==> 2C2H5OH(aq) + 2CO2(g) Ethanol, from a solution made from fermented sugar cane, can be concentrated by fractional distillation. In Brazil it is blended with petrol to give an alternative motor vehicle fuel. C2H5OH(l) + 3O2(g) ==> 2CO2(g) + 3H2O(l) The alcohols form an homologous series with the functional group C-OH. It is the presence of this functional group that gives alcohols their characteristic properties. The simplest homologous series of alcohols have the general formula CnH2n+1OH eg Ethanol is shown above and the simplest, lowest carbon number one, is methanol (shown below) o o o o o or or All the alcohols are flammable colourless liquids with a not un-pleasant odour? They all behave chemically in the same way but the boiling point steadily rises with increase in molecule size. The next two are called propanol and butanol, the names end in ...ol o CH3-CH2-CH2-OH or and Alcohols react, reversibly, with carboxylic acids to form esters and water. Ethyl ethanoate is formed by the reaction of ethanoic acid with ethanol eg o ethanoic acid + ethanol ethyl ethanoate + water o + + H2O o its an equilibrium, 2/3rds conversion, and catalysed by a few drops of concentrated sulphuric acid Alcohols react with sodium to form hydrogen. o normal fizzing is observed and the salt product is soluble in the alcohol itself. o eg ethanol + sodium ==> sodium ethoxide + hydrogen o 2C2H5OH + 2Na ==> 2C2H5O-Na+ + H2 Ethanol can be produced by the reaction of steam and ethene in the presence of a strong acid catalyst (Phosphoric acid). The reversible reaction is carried out at a moderately high temperature (eg 300 oC) and a high pressure (eg 60 x atmospheric pressure). The higher temperature and catalyst speed up the reaction and increasing pressure moves the equilibrium to the right (side least gaseous molecules at 300 oC) o + H2O ==> Advantages and disadvantages of the two methods of making ethanol: o advantages of fermentation: cheap and renewable resource like sugar cane (Brazil), sugar beet o disadvantages of fermentation: slow reaction and made by an inefficient batch process, poor quality product eg low aqueous concentration, other organic chemicals formed to and yeast cell residues to remove . o advantages of ethene route: fast and efficient continuous process, relatively pure product, country may have local oil supply (eg North Sea for UK, Middle East countries) o disadvantages of ethene route: using a non-renewable finite resource (crude oil/cracking) Ethanol can be oxidised to form ethanoic acid which is a useful organic chemical. BUT it is this oxidation of ethanol that results in alcoholic drinks turning sour (eg cider, wine) when exposed to air. Ethanoic acid (old name 'acetic acid') is the basis of vinegar and is also used in making esters (eg pear drop essence, or . 20 www.studyguide.pk + O2 ==> + H2O This oxidation can also be done by heating the ethanol with a mixture of sulphuric acid and potassium dichromate(VI) solution. The mixture turns from orange to green. o When burned, ethanol, like any alcohol, forms carbon dioxide and water CH3CH2OH + 2O2 ==> 2CO2 + 3H2O Ethanol can be dehydrated to ethene by passing the alcohol vapour over heated aluminium oxide catalyst. o o ==> o o + H2O This reaction is potentially an important source of organic chemicals eg plastics from a renewable resource since the ethanol can be made by fermentation of carbohydrates etc. Alcohols from propanol upwards, ie from carbon number 3 or greater, will form isomers. o You will find plenty of examples on the advanced organic chemistry page for alcohols. The steroid, cholesterol, contains the alcohol group -OH. Cholesterol is an essential steroid to humans but if too much is produced it can cause heart disease. http://webbook.nist.gov/cgi/cbook.cgi?Name=cholesterol&Units=SI gives the skeletal formula structure of cholesterol (this structure representation is only dealt with at advanced level). All the lines in the structure represent bonds between carbon atoms except the 'dash' for the -OH alcohol group in the bottom right of the molecule. Also note the 'alkene' double bond functional group to the right of the -OH group. 4. What other families of organic compounds are there? 4a. The acids that we find in fruits and in vinegar belong to a homologous series called carboxylic acids and many fragrances and food additives are esters. 4b. Polymers do not form a homologous series but they are all organic compounds having very long molecules. 4a. CARBOXYLIC ACIDS and ESTERS Carboxylic acids form another homologous series and have the functional group -COOH. The structures of the first three members are given below: Names end in ...oic acid. o methanoic acid: o ethanoic acid: or or or or o propanoic acid: or or Vinegar contains ethanoic acid (old name 'acetic acid'), see above in section 3 Alcohols above. It is used as a preservative and in food flavourings. Ethanoic acid is used in the manufacture of the fibre, acetate rayon. Oranges, lemons and many soft drinks contain a carboxylic acid eg citric acid. Aspirin is a carboxylic acid. Aspirin is a drug used for pain relief and is taken regularly by those at risk from heart attacks. Ascorbic acid (vitamin C) is another carboxylic acid and is present in fresh fruit and vegetables. Carboxylic acids are weak acids, typically solutions are around pH3 (yellow-orange-pink with universal indicator). 21 www.studyguide.pk o [theory on Extra Aqueous Chemistry page look for keywords] They react and are neutralised by ... with examples ... o metals to form salts and hydrogen ethanoic acid + magnesium ==> magnesium ethanoate + hydrogen 2CH3COOH + Mg ==> (CH3COO)2Mg + H2 o alkali bases to form a carboxylic acid salt and water eg ... eg ethanoic acid + potassium hydroxide ==> potassium ethanoate + water CH3COOH + KOH ==> CH3COOK + H2O o carbonate and hydrogencarbonate bases to produce a carboxylic acid salt, water and carbon dioxide eg ... eg ethanoic acid + sodium hydrogen carbonate ==> sodium ethanoate + water + carbon dioxide CH3COOH + NaHCO3 ==> CH3COONa + H2O + CO2 OR propanoic acid + sodium carbonate ==> sodium propanoate + water + carbon dioxide 2CH3CH2COOH + Na2CO3 ==> 2CH3CH2COONa + H2O + CO2 Carboxylic acids react with alcohols to form members of another homologous series called esters. Concentrated sulphuric acid acts as a catalyst in this reaction. See the formation of ethyl ethanoate above in section 3. above. and show the structures of other esters made from ethanoic acid: namely methyl ethanoate using methanol, and propyl ethanoate from propyl alcohol o and what would the structure of their original alcohols be and what would the structure of butyl ethanoate be? Esters are usually sweet smelling and widely used as fragrances (eg perfumes) and food flavourings. o o 4b. POLYMERS - synthetic macromolecules Some basic notes on polymers and plastics in the Oil Products Notes. Some important structure, strength and 1D to 3D dimension concepts) in the "Chemical Bonding" notes. Most polymers (plastics) are made from alkene compounds containing the -C=C- bond by addition polymerisation. Poly(chloroethene) is made from chloroethene (old name 'vinyl chloride), CH 2=CHCl but the polymer is generally called polyvinylchloride, PVC. The general equation and the formation of poly(ethene) and poly(propene) are shown on the Oil notes page. The formation of PVC is shown below. Polymers (plastics) consist of a tangled mass of very long molecules in which the atoms are joined by strong covalent bonds to form long chains, but there are much weaker intermolecular forces holding the material together. In thermosoftening plastics like poly(ethene), poly(propene) or poly(chloroethene) PVC, because the inter-molecular attractive forces between the chains are weak, the plastic softens when heated and hardens again when cooled. It also means the polymer molecules can slide over each other. This means they can be easily stretched or moulded into any desired shape. 22 www.studyguide.pk However it is possible to manufacture and process plastics in which the polymer chains are made to line up. This greatly increases the intermolecular forces between the 'aligned' polymer molecules and strong fibre strands of the plastic can be made. o Examples: The addition polymer poly(propene) and the condensation polymers nylon and Terylene When a thermosetting plastic is first heated covalent bonds are formed between adjacent chains of the polymers. These strong covalent cross-linkages give the material a high melting point and greatly increased strength and rigidity. They also prevent thermosetting plastics from being softened with heat and therefore from being stretched or re-shaped. However it does make a much stronger material and not as flammable. On heating strongly they do NOT melt, but tend to char, gradually giving off gases. o Melamine (used in furniture) and many glues are examples of thermosetting polymers. Problems with using plastics are on the Oil Products Notes page. Some important structure, strength and 1D to 3D dimension concepts) in the "Chemical Bonding" notes. 4c. Other Synthetic Polymers - macromolecules Condensation polymerisation involves linking lots of small monomer molecules together by eliminating a small molecule. This is often water from two different monomers, a H from one monomer, and an OH from the other, the 'spare bonds' then link up to form the polymer chain. Nylon (a polyamide) is formed by condensation polymerisation, the structure of nylon represented as ... (the rectangles represent the rest of the carbon chains in each unit) (3 units etc.) This is the same linkage (-CO-NH-) that is found in linked amino acids in naturally occurring macromolecules called proteins, where it is called the 'peptide' linkage. Terylene (a polyester) is formed by condensation polymerisation and the structure of Terylene represented as (3 units etc.) This is the same kind of 'ester linkage' (-COOC-) found in fats which are combination of long chain fatty carboxylic acids and glycerol (alcohol with 3 -OH groups, a 'triol'). Terylene and nylon are good for making 'artificial' or 'man-made' fibres used in the clothing and rope industries. In the manufacturing process the polymer chains are made to line up. This greatly increases the intermolecular forces between the 'aligned' polymer molecules and strong fibre strands of the plastic can be made. Some important structure, strength and 1D to 3D dimension concepts in "Chemical Bonding" notes. 23 www.studyguide.pk 5. Naturally Occurring Molecules Small Molecules <=> Natural Polymers = Macromolecules Carbohydrates, Proteins and Fats are the main nutrient constituents of food. 5a. Carbohydrates Carbohydrates are a whole series naturally occurring molecules based on the elements carbon, hydrogen and oxygen. Historically the name 'carbohydrate' comes from the fact that all their formulae seemed to be based on Cx(H2O)y (see key above) BUT this is not the way to think of their formula. They range from relatively small molecules called monosaccharide (means one basic unit), or disaccharide (two basic units combined) to very large natural polymers or macromolecules called polysaccharides (many units combined). A summary of them is shown in the key diagram above along with some familiar names from biology. Glucose is one of the simpler sugar molecules. The structural formula is shown on the left and you should be able to see that there are 4 bonds to each carbon, 2 to each oxygen and just 1 bond to each hydrogen atom. The right-hand 'shorthand' skeletal formula version uses short straight lines to represent bonds. Most H's and their bonds are not shown, and at AS-A2 level it is assumed you can interpret these structures 'back to' a full structure!, but they are handy for describing large 'biochemical' molecules (see polysaccharide below) 24 www.studyguide.pk Sucrose is a disaccharide made from combining two monosaccharide molecules, glucose and fructose by the elimination of a water. o On hydrolysis sucrose reforms the glucose and fructose. o 2C6H12O6 <=> C12H22O12 + H2O The formation of complex carbohydrates: o These are made of smaller carbon, hydrogen and oxygen based molecules combining together eg starch and cellulose are formed from glucose, molecular formula C6H12O6. Their formation can be described in terms of a large number of sugar units joined together by condensation polymerisation o eg the 'box' diagram below shows 4 units of a natural carbohydrate polymer being formed o Note: Condensation polymerisation means the joining together of many small 'monomer' molecules by eliminating an even smaller molecule between them to form the linkage. eg HO-XXXXX-OH + HO-XXXXX-O-XXXXX-OH + H2O etc. n C6H12O6 ==> (C5H10O5)n + nH2O (where n is a very large number to form the natural polymer) The XXXXX or the [rectangles] below, represent the rest of the carbon chains in each unit (more detail in the 3rd diagram below). plus many H2O etc. This diagram of starch or cellulose is in 'skeletal formula' style and both are polymers of glucose - can you see the connection between each 'unit' and the structure of glucose itself? The resulting natural polymer is called a polysaccharide. Acid hydrolysis of complex carbohydrates (eg. starch) gives simple sugars. o This can be brought about by eg warming starch with hydrochloric acid solution to form glucose. o (C5H10O5)n + nH2O ==> n C6H12O6 (where n is a very large number) 25 www.studyguide.pk The hydrolysis products from polysaccharides can be analysed with paper chromatography. 5b. Proteins and Amino Acids Amino acids are carboxylic acids (like ethanoic acid) but one of the hydrogen atoms of the 2nd carbon atom is substituted with an amino group (a nitrogen + two hydrogens gives -NH2). Another hydrogen on the same 2nd carbon can be substituted with other groups of atoms (R) to give a variety of amino acids. or The simplest is aminoethanoic acid or 'Glycine' is another amino acid called 2-aminopropanoic acid or 'Alanine'. All amino acids have the general structure H2N-CH(R)-COOH. o R can vary, think of it as the 'Rest of the molecule! o R = H for Glycine, R = CH3 for Alanine. They polymerise together, by condensation polymerisation, forming proteins or polypeptides. o The peptide linkage is formed by elimination of water between two amino acids. o HNH-CH(R)-COOH + HNH-CH(R)-COOH ==> H2N-CH(R)-CO-HN-CH(R)-COOH + H2O etc. so ... o n H2N-CH(R)-COOH ==> -NH-CO-CH(R)-NH-CO-CH(R)-NH-CO-CH(R)-NH-CO-CH(R)- etc. n units long o So proteins are condensation polymers of amino acids. Proteins have the same (amide) linkages as nylon but with different units. Proteins are an important component of tissue structure and enzymes are protein molecules. When proteins are heated with aqueous hydrochloric acid or sodium hydroxide solution they are hydrolysed to amino acids. o see chromatography below, about how amino acids are identified in proteins. 5c. Fats, Oils and Margarine Oils and Fats are an important way of storing chemical energy in living systems and are also a source of essential longchain fatty acids. Most of them are esters of the tri-alcohol ('triol') glycerol (systematic name propane-1,2,3-triol, but that can wait until AS-A2 level!). The carboxylic acids which combine with the glycerol are described as 'long-chain fatty acids'. The resulting ester is called a 'triester' or 'triglyceride'. The 'long-chain fatty acids' can be saturated, with no C=C double bonds, and so forming saturated oils or fats (1st diagram below of the triglyceride formed from palmitic acid). 26 www.studyguide.pk The 'long-chain fatty acids' can be unsaturated, with one or more C=C double bonds, and so forming unsaturated oils or fats (2nd diagram below of the triglyceride formed from oleic acid). Some sub-notes on Oil and Fat Structure: (health issues dealt with further down) o They have the same linkages as Terylene but with different units. o They are not as big as polymer molecules, but a lot bigger than a single petrol or sugar molecule. o There can be 1 to 3 different saturated or unsaturated fatty acid components, so lots of variation possible in structure of the oil or fat. The diagrams just assume three molecules of the same 'fatty' acid. o Monounsaturated fats have one C=C double bond in them, polyunsaturated fats usually have at least three C=C bonds in their molecular structure. o For the same molecular size in terms of carbon number, unsaturated fats have slightly lower intermolecular forces because the C=C double bond produces a kink in the carbon chain and they can't pack as closely together as the saturated molecules. This gives unsaturated fats a lower melting point and so they tend to occur as eg vegetable oils rather than saturated low melting solids from meat and dairy products. o However, this means these unsaturated oils are not as conveniently 'spreadable' as 'butter'. To overcome this problem, 'margarine' was invented. The vegetable oils are reacted with hydrogen using the gas and a nickel catalyst. This reaction adds hydrogen atoms to the double bonds making a more saturated and more 'spreadable' solid fat we call 'margarine'. The reaction for any double bond is: >CH=CH< + H2 == Ni ==> -CH2-CH2 BUT it does mean that it is more like animal fat now but various blendes have been developed to suit your dietary needs or desires! 27 www.studyguide.pk 'Traditional' soap is a product of the hydrolysis of fats. o 'Soapy' soaps (not modern detergents) are the sodium salts of long chain fatty acids formed by heating fatty oils with sodium or potassium hydroxide to hydrolyse them. o This reaction breaks the fat molecule down into one glycerol molecule (triol alcohol) and three sodium salts of the long chain carboxylic fatty acids. Since fats and oils are important to our diet, there is the ever present danger of over-consumption (speaking as someone who loves chips and spicy crisps!). So there are health and social, as well as 'molecular' issues to address! o We need oils and fats as sources of important essential fatty acids. o We need both saturated and unsaturated fats or oils. The main sources of saturated fats are from meat and dairy products eg 'dripping' and butter. The main sources of unsaturated fats are plant oils eg olive oil. o It is recommended that we do not overdo the fat intake but we do need both saturated and unsaturated fats. However, too much saturated fat raises cholesterol levels and is not too good for the heart. If you wish to know more about fats and oils there is plenty on the web! eg .... o http://www.healthnet.org.uk/diet/fatoil.htm o http://biology.clc.uc.edu/courses/bio104/lipids.htm o http://www.nutristrategy.com/fatsoils.htm 5d. Chromatography Hydrolysis means breaking down a molecule with water to form two or more products. o Hydrolysis is accelerated if the substance is heated with acid or alkali solutions. When proteins are heated with aqueous acid they are hydrolysed to amino acids. Acid hydrolysis of complex carbohydrates (eg. starch) gives simple sugars. Chromatography is useful in separating and identifying the products of hydrolysis of carbohydrates and proteins. o The hydrolysis can be done by boiling the carbohydrate or protein with hydrochloric acid. o The hydrolysed mixture is then 'spotted' onto the pencil base line of the chromatography paper. Known sugars or amino acids are also spotted onto the base line too. The prepared paper is then placed vertically in a suitable solvent, which rises up the paper. o Since the products are colourless, the dried chromatogram is treated with another chemical to produce a coloured compound. Ninhydrin produces purple spots with amino acids and resorcinol makes coloured spots with sugars. o You can then tell which amino acids made up the protein or the sugars from which the carbohydrate was formed. The number of different spots tells you how many different amino acids or sugars made up the natural macromolecule. Spots which horizontally match the standard known molecule spots confirm identity. Starch gives one spot because only glucose is formed on hydrolysis. (C5H10O5)n + nH2O ==> n C6H12O6 (where n is a very large number) o More details on chromatography. 28 www.studyguide.pk 6. Vitamins, Drugs and Food Additives (NOT FINISHED) Vitamins are particular essential molecules with particular roles in living systems which are NOT proteins, carbohydrates, fats or mineral salts. One of the most important ones in any diet is Vitamin C or Ascorbic Acid. Its structure is related to 'simple' sugars but humans are one of the few mammals that are unable to synthesise vitamin C. o It is essential for healthy tissue and one of its functions is the removal of dangerously reactive chemical species called free radicals (see further on). o Vitamin C is present in fruit and vegetables but the amount is reduced by prolonged storage and cooking.. o 250 years ago, as many as 2/3 of a ship's crew died from vitamin C deficiency causing scurvy. In 1747 it was decided to give sailors citrus fruits to recover from scurvy but wasn't until 200 years later that vitamin C was recognised. o In contrast to the other water-soluble vitamins, vitamin C has no clear cut role as a catalyst or part of an enzyme. It does, however, have a range of other important functions: Collagen formation. Vitamin C in collagen formation which is found wherever tissues require strengthening, especially in tissues with a protective, connective, or structural function. Collagen is critical to the maintenance of bone and blood vessels and is essential in wound healing. Antioxidant activity. Ascorbic acid can act as an antioxidant by donating electrons and hydrogen ions, and reacting with reactive oxygen species or free radicals. Iron absorption. Vitamin C is important for the effective absorption of iron and reduces iron(III) Fe 3+ to iron(I) Fe2+. It helps in the synthesis of vital cell compounds. During times of physical and emotional stress, as well as during infection, there is increased production of oxygen radicals. Therefore there is increased reliance on vitamin C's activity as an antioxidant. Vitamin C is vital for the function of the immune system, but the effectiveness of large doses of vitamin C in preventing and alleviating the symptoms of the common cold is still a matter of debate. o Two of the earliest signs of deficiency (prevention of collagen synthesis) relate to its roles in maintaining the integrity of blood vessels. The gums around the teeth bleed more easily, and the capillaries under the skin break spontaneously producing tiny haemorrhages. If you are short of vitamin C for say 20 days, scurvy can develop and is characterised by further haemorrhaging, muscles depletion, rough-brown-dry-scaly skin, deep bruising. Wounds fail to heal properly and bone fails to rebuild properly too and you are further likely to suffer from anaemia and infections. o SO EAT yer fruit and veg 'guys' (as well as a few crisps!) AND keep yer health and still pass those dreaded exams!!!! Drugs can be defined as an externally administered substances which modifies or affects chemical reactions in the body, usually for the bodies greater well-being. Poisons can be defined in the same way, but hopefully not intentionally and have undesired effects! o An analgesic is a drug used to reduce pain and is type of anti-inflammatory agent. o The molecular structure of three well known analgesics are shown in the diagram below. o All are used for 'headache' treatment, and hopefully using this website and others will help minimise their use! 29 www.studyguide.pk The central hexagonal ring of 6 carbon atoms is called a 'benzene' or 'aromatic' ring. The 4th outer electron of carbon (group 4) is delocalised, so the expected 4th bond per C atom forms part of a 'communal' system (more on this at advanced level, but the covalence rule of 4 for carbon is not broken!, you have seen this situation before, check out graphite. You can show a benzene ring as a simple hexagon with a circle in it from + NaOH ==> + H2 O The modern pharmaceutical industry has its origins in herbal and other traditional medicine. eg An extract of willow herb extract can be made from the leaves, bark and seeds of the willow tree. Amongst other ailments it was given to help curing feverish headaches and relief of pain in childbirth. When ingested the body hydrolyses and oxidises the naturally occurring 'precursor' molecule to form salicylic acid* which is the 'active' molecule in the body. in the 1890's the German chemist Hoffmann experimented with various chemical modifications of salicylic acid and found the best and chemically stable form was 'aspirin' (shown below). He tried the variations on his own father! who survived to provide valuable 'clinical trials' - hardly acceptable these days! * 'Oil of winter green' from certain plants is the methyl ester of salicylic acid and has similar 'medicinal effects'. Aspirin (and the others shown) are not very soluble in water. Soluble aspirin is made by neutralising the carboxylic acid with the alkali sodium hydroxide to make the much more soluble sodium salt of the acid. The reaction, using skeletal formula, is shown in the diagram below the three analgesic drug structures. New drugs and testing them: It costs a lot of money to develop a new medicine so the price charged by the pharmaceutical company must cover the cost of research, production and marketing. Patents are taken out to protect the company's commercial interests in the new medicine. There can be a range of formulations of a particular medicine when you buy it over the counter eg tablet of 100% aspirin, soluble aspirin (via Na + salt of the acid from neutralisation) and aspirin might form part of a mixture including substances that have other beneficial effects. The main point here is that aspirin, like many drugs, can have multi-functional effects, hopefully all beneficial. BUT this, sadly, is not always the case, because with any new drug there is always the danger of unknown side-effects. Therefore there is a tremendous responsibility on pharmaceutical companies to ensure the development of safe and effective drugs. Lots of time and money spent on discovering and developing new drugs and there are lots of factors to consider: o From the discovery of a potentially useful molecule, sometimes called the 'lead molecule', which can be from natural source or produced in some other project etc. o Is there room in the commercial market place for it? o Do research to see if its safe, otherwise further development is a waste of time and money or if not safe, can its molecular structure be modified? o Can the modification be safe? and more effective? o In what form, can it be/needs to be, administered in? for clinical trials. o Carefully clinical trials in various phases, noting particularly if any side-effects which may be harmful. 30 www.studyguide.pk o o o o Do you test new drugs on animals? - an emotive issue, can non-animal testing always allow the safe development of new products? Do you test new drugs on patients in a life threatening situation, give them a last chance at some risk? Patient health and safety issues versus very big drug company commercial interests are a matter of public concern. Any new drug must finally pass all the tests before legally licensed for patient consumption ... sadly, the 'drug companies' and the 'powers to be' do not always get it right (eg phthalidomide), but do not the benefits outweigh the occasional tragedy which we should do our best to avoid? Food additives are chemicals added to food to give particular effects eg colourings, flavourings, preservation and sweetening. o Colourings to make food more attractive, to fit in with the consumers perception of what it should look like. o Flavourings to make food more 'tasty', less 'bland', and to fit in with the consumers perception of what it should taste like. o Preservatives are to increase the 'shelf-life' of packaged food, decrease risk of food poisoning. o Sweeteners counter bitterness or pander to our taste! E-numbers are reference numbers used by the European Union to help identification of food additives. o All food additives used in the European Union are identified by an E-number. o The "E" stands for "Europe" or "European Union". o Normally each food additive is assigned a unique number, though occasionally, related additives are given an extension (eg a,b,i or ii etc.) to another E-number. o The Commission of the European Union assigns E-numbers after the additive is cleared by the Scientific Committee on Food (SCF), the body responsible for the safety evaluation of food additives in the European Union. A summary is given below. E100-199 food colours E200-299 preservatives E300-399 anti-oxidants, phosphates, and complexing agents E400-499 thickeners, gelling agents, phosphates, emulsifiers E500-599 salts and related compounds E600-699 flavourings E700-899 not used for food additives (used for animal feed additives!) E900-999 surface coating agents, gases, sweeteners E1000-1399 E14001499 miscellaneous additives starch derivatives E-numbers are only used for substances added directly to food products, so contaminants, enzymes and processing aids, which may be classified as additives in the USA, are not included in the E-number system. There is an EU directive on food labeling which requires food additives to be listed in the product ingredients whenever they are added for technological purposes. o This includes colouring, sweetening and favor enhancement as well as for preservation, thickening, emulsifying and the like. o Ingredients must be listed in descending order of weight, which means that are generally found close to the end of the list of ingredients. o However, substances used in the protection of plants and plant products, flavorings and substances added as nutrients (e.g., minerals, trace elements or vitamins) do not need to be included in the ingredient list. o Because of this, some substances that are regulated as food additives in other countries may be exempt from 31 www.studyguide.pk the food additive definition in the EU. 7. CFC's and Free Radicals (NOT FINISHED) If enough energy is supplied by heat or by visible/uv electromagnetic radiation, or the is weak enough, a covalent bond can break in two ways. This illustrated with the molecule chloromethane CH3Cl. o Unevenly where the electron bond pair can stick with one fragment and a positive and negative ion form. eg CH3Cl ==> CH3+ + Cl- (at AS-A2 level this is called heterolytic bond fission) o shows what happens to the molecule or evenly, where the bonding pair of electrons are equally divided between two highly reactive fragments called free radicals. Free radicals are characterised by having an unpaired electron not involved in a chemical bond. The . means the 'lone' electron on the free radical, which is not part of a bond anymore, and wants to pair up with another electron to form a stable bond - that's why free radicals are so reactive! eg CH3Cl ==> CH3. + .Cl (at AS-A2 level this is called homolytic bond fission) shows what happens to the molecule In the stratosphere small amounts of unstable ozone O3 (trioxygen) are formed by free radical reactions. The chemistry of free radicals is important in the current environmental issue of ozone layer depletion. o Chlorofluorocarbons (CFC's for shorthand) are organic molecules containing carbon, fluorine and chlorine o eg dichlorodifluoromethane has the formula CCl2F2 o They are very useful low boiling organic liquids or gases, until recently, extensively used in refrigerators and aerosol sprays eg repellents. o They are relatively unreactive, non-toxic and have low flammability, so in many ways they are 'ideal' for the job they do. o However it is their chemical stability in the environment that eventually causes the ozone problem but first we need to look at how ozone is formed and destroyed in a 'natural cycle'. This presumably has been in balance for millions of years and explains the uv ozone protection in the upper atmosphere. o Ozone is formed in the stratosphere by free radical reactions. 'ordinary' stable oxygen O2 (dioxygen) is split (dissociates) into two by high energy ultraviolet radiation (uv photon energy 'wave packets) into two oxygen atoms (which are themselves radicals) and then a 'free' oxygen atom combines with an oxygen molecule to form ozone. O2 + uv ==> 2O. then O. + O2 ==> O3 The ozone is a highly reactive and unstable molecule and decomposes into dioxygen when hit by other uv light photons. The oxygen atom radical can do several things including help forming O 3, and O2. O3 + uv ==> O2 + O. This last reaction is the main uv screening effect of the upper atmosphere and the ozone absorbs a lot of the harmful incoming uv radiation from the Sun. If the ozone levels are reduced more harmful uv radiation reaches the Earth's surface and can lead to medical problems such as increased risk of sunburn and skin cancer and it also accelerates skin aging processes. There is strong evidence to show there are 'holes' in the ozone layer with potentially harmful 32 www.studyguide.pk o o o effects, so back to the CFC problem! The chemically very stable CFCs diffuse up into the stratosphere and decompose when hit by ultraviolet light (uv) to produce free radicals, including free chlorine atoms, which themselves are highly reactive free radicals. eg CCl2CF2 ==> CClF2. + Cl. (note the C-Cl bond is weaker than the C-F bond) The formation of chlorine atom radicals is the root of the problem because they readily react with ozone and change it back to much more stable ordinary oxygen. O3 + Cl. ==> O2 + ClO. bye bye ozone! and no uv removed in the process! Therefore many countries are banning the use of CFCs, but not all despite the fact that scientists predict it will take many years for the depleted ozone layer to return to its 'original' O 3 concentration. 33