Chemistry

March 25, 2018 | Author: ts8166 | Category: Molecular Orbital, Chemical Bond, Gases, Acid, Molecules
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1. Of the following name/symbol combinations of elements, which one is WRONG? (a) uranium/U (b) sulfur/S (c) nitrogen/N (d) potassium/K (e) iron/I 2.Of the following symbol/name combinations of elements, which one is WRONG? (a) C/carbon (b) B/barium (c) F/fluorine (d) N/nitrogen (e) U/uranium 3. The chemical symbol for manganese is (a) Mn (b) Mo (c) Ma (d) Ga (e) Mg 4. The number 0.005436 has how many significant figures? (a) 7 (b) 3 (c) 4 (d) 5 (e) 6 5. The number 10.00 has how many significant figures? (a) 1 (b) 2 (c) 3 (d) 4 (e) 5 6. What is the volume of a 2.50 gram block of metal whose density is 6.72 grams per cubic centimeter? (a) 16.8 cubic centimeters (b) 2.69 cubic centimeters (c) 0.0595 cubic centimeters (d) 0.372 cubic centimeters (e) 1.60 cubic centimeters 7. A cube of 1.2 inches on the side has a mass of 36 grams. What is the density in g/cm3? (a) 21 (b) 2.2 (c) 30. (d) 1.3 (e) 14 8. Nitric acid is a very important industrial chemical: 1.612 x 1010 pounds of it were produced in 1992. If the density of nitric acid is 12.53 pounds/gallon, what volume would be occupied by this quantity? (1 gallon = 3.7854 liters) (a) 7.646 x 1011 liters (b) 8.388 x 109 liters (c) 1.287 x 109 liters (d) 5.336 x 1010 liters (e) 4.870 x 109 liters 9. Identify the INCORRECT statement. (a) Helium in a balloon: an element (b) Paint: a mixture (c) Tap water: a compound (d) Mercury in a barometer; an element 10. 4. 11. An unused flashbulb contains magnesium and oxygen. After use, the contents are changed to magnesium oxide but the total mass does not change. This observation can best be explained by the (a) Law of Constant Composition. (b) Law of Multiple Proportions. (c) Avogadro's Law. (d) Law of Conservation of Mass. Which answer includes all the following that are chemical changes and not physical changes? I. freezing of water II. rusting of iron III. dropping a piece of iron into hydrochloric acid (H2 is produced) IV. burning a piece of wood V. emission of light by a kerosene oil lamp (a) III and IV (b) II and V (c) I, II, III, IV, and V (d) II, III, and V (e) II, III, IV, and V 12. Which response lists all of the following properties of sulfur that are physical properties and not other properties? I. It reacts with hydrogen when heated. II. It is a yellow solid at room temperature. III. It is soluble in carbon disulfide. IV. Its density is 2.97 g/cubic centimeter V. It melts at 112�C. (a) II, III, IV, and V (b) II, IV, and V (c) I (d) II, III, and IV (e) III, IV, and V Answers to Chapter 1 1. (e) 2. (b) 3. (a) 4. (c) 5. (d) 6. (d) 7. (d) 8. (e) 9. (c) 10. (d) 11. (e) 12. (a) The formula weight of the compound, Al2(SO4)3 18H2O is: (a) 394.4 g (b) 666.4 g (c) 110,900 g (d) 466.8 g (e) 561.2 g 2. The weight of a millimole of (NH4)2HPO4 is: (a) 132 g (b) 114 g (c) 1.14 x 10-3 g (d) 0.132 g (e) 6.02 x 1020 g 3. How many moles of alanine, C3H7NO2, are there in 159 g of alanine? (a) 1.42 x 104 (b) 1.78 (c) 0.992 (d) 0.560 (e) 3.31 How many atoms are in one mole of CH3OH? (a) 6 (b) 6.0 x 1023 (c) 12.0 x 1023 (d) 3.6 x 1024 (e) 3 5. The mass in grams of 2.6 x 1022 chlorine atoms is: (a) 4.4 (b) 11 (c) 0.76 (d) 1.5 (e) 3.2 6. How many aluminum atoms are there in 3.50 grams of Al2O3? (a) 4.13 x 1022 (b) 4.90 x 1022 (c) 2.07 x 1022 (d) 1.68 x 1022 (e) 2.45 x 1022 7. Which one of the samples contains the most atoms? (a) 1 mol of CO2(g) (b) 1 mol of UF6(g) (c) 1 mol of CH3COCH3(l) (d) 1 mol of He(g) (e) all contain the same number of atoms 8. Which one of the samples contains the most molecules? (a) 1 mol of CO2(g) (b) 1 mol of UF6(g) (c) 1 mol of CH3COCH3(l) (d) 1 mol of He(g) (e) all contain the same number of molecules 9. Which one of the samples has the largest mass? (a) 1 mol of CO2(g) (b) 1 mol of UF6(g) (c) 1 mol of CH3COCH3(l) (d) 1 mol of He(g) (e) all have the same mass 10. Which of the following statements is(are) FALSE? 1. The percent by mass of each element in a compound depends on the amount of the compound. 2. The mass of each element in a compound depends on the amount of the compound. 3. The percent by mass of each element in a compound depends on the amount of element present in the compound. (a) 2 and 3 (b) 1 only (c) 1 and 2 (d) 1, 2 and 3 (e) another combination 11. Guanidin, HNC(NH2)2, is a fertilizer. To three significant figures, what is the percent by mass of nitrogen in the fertilizer? (a) 45.2% (b) 49.4% (c) 54.8% (d) 65.1% (e) 71.1% 12. Calculate the percent, by weight, of carbon in 154 g of C4H8O3? (a) 46% (b) 31% (c) 72% (d) 27% (e) 55% Analysis of a sample of a covalent compound showed that it contained 14.4% hydrogen and 85.6% carbon 16. 14. by mass. What is the empirical formula for the compound? (a) CH (b) CH2 (c) CH3 (d) C2H3 (e) none of these An oxide of lead contains 90.65% Pb, by weight. The empirical formula is: (a) Pb (b) PbO (c) Pb3O4 (d) Pb2O3 (e) PbO2 15. A 0.500 g sample of a compound containing only antimony and oxygen was found to contain 0.418 g of antimony and 0.082 g of oxygen. What is the simplest formula for the compound? (a) SbO (b) SbO2 (c) Sb3O4 (d) Sb2O5 (e) Sb2O3 A compound contains, by mass, 40.0% carbon, 6.71% hydrogen, and 53.3% oxygen. A 0.320 mole sample of this compound weighs 28.8 g. The molecular formula of this compound is: (a) C2H4O2 (b) C3H6O3 (c) C2H4O (d) CH2O (e) C4H7O2 What mass of cerussite, PbCO3, would contain 35.0 grams of lead? (a) 27.1 g (b) 45.1 g (c) 42.4 g (d) 35.6 g (e) 51.7 g 17. Answers to Chapter 2 1. (b) 2. (d) 3. (b) 4. (d) 5. (d) 6. (a) 7. (c) 8. (e) 9. (b) 10. (b) 11. (e) 12. (a) 13. (b) 14. (c) 15. (e) 16. (b) 17. (b) 1. Balance the following equation with the smallest whole number coefficients. Choose the answer that is the sum of the coefficients in the balanced equation. Do not forget coefficients of "one." PtCl4 + XeF2 PtF6 + ClF + Xe (a) 16 (b) 22 (c) 24 (d) 26 (e) 32 2. Balance the following equation with the smallest whole number coefficients. Choose the answer that is the sum of the coefficients in the balanced equation. Do not forget coefficients of "one." Cr2(SO4)3 + RbOH Cr(OH)3 + Rb2SO4 13. (a) 10 (b) 12 (c) 13 (d) 14 (e) 15 3. Balance the following equation using minimum integral coefficients: NH3 + O2 NO2 + H2O The stoichiometric coefficient for oxygen gas O2 is: (a) 1 (b) 4 (c) 3 (d) 7 (e) 5 4. When iron pyrite (FeS2) is heated in air, the process known as "roasting" forms sulfur dioxide and iron(III) oxide. When the equation for this process is completed and balanced, using the smallest whole number coefficients, what is the coefficient for "O2"? ___ FeS2 + ___ O2 ___ SO2 + ___ Fe2O3 (a) 2 (b) 4 (c) 7 (d) 8 (e) 11 5. How many moles of KBrO3 are required to prepare 0.0700 moles of Br2 according to the reaction: KBrO3 + 5KBr + 6HNO3 6KNO3 + 3Br2 + 3H2O (a) 0.210 (b) 0.0732 (c) 0.0704 (d) 0.220 (e) 0.0233 6. Which of the following statements is FALSE for the chemical equation given below in which nitrogen gas reacts with hydrogen gas to form ammonia gas assuming the reaction goes to completion? N2 + 3H2 2NH3 (a) The reaction of one mole of H2 will produce 2/3 moles of NH3. (b) One mole of N2 will produce two moles of NH3. (c) One molecule of nitrogen requires three molecules of hydrogen for complete reaction. (d) The reaction of 14 g of nitrogen produces 17 g of ammonia. (e) The reaction of three moles of hydrogen gas will produce 17 g of ammonia. 7. Calcium carbide, CaC2, is an important preliminary chemical for industries producing synthetic fabrics and plastics. CaC2 may be produced by heating calcium oxide with coke: CaO + 3C CaC2 + CO What is the amount of CaC2 which can be produced from the reaction of excess calcium oxide and 10.2 g of carbon? (Assume 100% efficiency of reaction for purposes of this problem.) (a) 18.1 g (b) 28.4 g (c) 20.8 g (d) 19.8 g (e) 27.2 g 8. 12. 11. 9. Calculate the mass of hydrogen formed when 25 grams of aluminum reacts with excess hydrochloric acid. 2Al + 6HCl Al2Cl6 + 3H2 (a) 0.41 g (b) 1.2 g (c) 1.8 g (d) 2.8 g (e) 0.92 g When 12 g of methanol (CH3OH) was treated with excess oxidizing agent (MnO4-), 14 g of formic acid (HCOOH) was obtained. Using the following chemical equation, calculate the percent yield. (The reaction is much more complex than this; please ignore the fact that the charges do not balance.) 3CH3OH + 4MnO43HCOOH + 4MnO2 (a) 100% (b) 92% (c) 82% (d) 70% (e) 55% A commercially valuable paint and adhesive stripper, dimethyl sulfoxide (DMSO), (CH3)2SO, can be prepared by the reaction of oxygen with dimethyl sulfide, (CH3)2S, using a ratio of one mole oxygen to two moles of the sulfide: O2 + 2(CH3)2S 2(CH3)2SO If this process is 83% efficient, how many grams of DMSO could be produced from 65 g of dimethyl sulfide and excess O2? (a) 68 g (b) 75 g (c) 83 g (d) 51 g (e) 47 g The formation of ethyl alcohol (C2H5OH) by the fermentation of glucose (C6H12O6) may be represented by: C6H12O6 2C2H5OH + 2CO2 If a particular glucose fermentation process is 87.0% efficient, how many grams of glucose would be required for the production of 51.0 g of ethyl alcohol (C2H5OH)? (a) 68.3 g (b) 75.1 g (c) 115 g (d) 229 g (e) 167 g The limiting reagent in a chemical reaction is one that: (a) has the largest molar mass (formula weight). (b) has the smallest molar mass (formula weight). (c) has the smallest coefficient. (d) is consumed completely. (e) is in excess. 13. If 5.0 g of each reactant were used for the the following process, the limiting reactant would be: 2KMnO4 +5Hg2Cl2 + 16HCl 10HgCl2 + 2MnCl2 + 2KCl + 8H2O (a) KMnO4 (b) HCl 10. (c) H2O (d) Hg2Cl2 (e) HgCl2 14. What mass of ZnCl2 can be prepared from the reaction of 3.27 grams of zinc with 3.30 grams of HCl? Zn +2HCl ZnCl2 + H2 (a) 6.89 g (b) 6.82 g (c) 6.46 g (d) 6.17 g (e) 6.02 g 15. How many grams of NH3 can be prepared from 77.3 grams of N2 and 14.2 grams of H2? (Hint: Write and balance the equation first.) (a) 93.9 g (b) 79.7 g (c) 47.0 g (d) 120.0 g (e) 13.3 g 16. Silicon carbide, an abrasive, is made by the reaction of silicon dioxide with graphite. SiO2 +3C SiC + 2CO If 100 g of SiO2 and 100 g of C are reacted as far as possible, which one of the following statements will be correct? (a) 111 g of SiO2 will be left over. (b) 44 g of SiO2 will be left over. (c) 82 g of C will be left over. (d) 40 g of C will be left over. (e) Both reactants will be consumed completely, with none of either left over. 17. Calculate the mass of 6.00% NiSO4 solution that contains 40.0 g of NiSO4? (a) 667 g (b) 540 g (c) 743 g (d) 329 g (e) none of these 18. How many grams of water are contained in 75.0 grams of a 6.10% aqueous solution of K3PO4? (a) 75.0 g (b) 73.2 g (c) 70.4 g (d) 68.1 g (e) 62.8 g 19. The mass (in grams) of FeSO4 7H2O required for preparation of 125 mL of 0.90 M solution is: (a) 16 g (b) 25 g (c) 13 g (d) 31 g (e) 43 g 20. What is the molarity of phosphoric acid in a solution labeled 20.0% phosphoric acid (H3PO4) by weight with a density = 1.12 g/mL? (a) 0.98 M (b) 2.3 M (c) 2.7 M (d) 3.0 M (e) 3.6 M 21. How many mL of 17 M NH3 must be diluted to 500.0 mL to make a 0.75 M solution? 3. 2. 24. 23. 22. (a) 13 mL (b) 22 mL (c) 39 mL (d) 73 mL (e) none of these How many grams of Ag2CO3 are required to react with 28.5 mL of 1.00 M NaOH solution? Ag2CO3 +2NaOH Ag2O + Na2CO3 + H2O (a) 7.87 g (b) 3.93 g (c) 15.7 g (d) 10.8 g (e) 8.16 g How many milliliters of 0.200 M NH4OH are needed to react with 12.0 mL of 0.550 M FeCl3? FeCl3 + 3NH4OH Fe(OH)3 + 3NH4Cl (a) 99.0 mL (b) 33.0 mL (c) 8.25 mL (d) 68.8 mL (e) 132 mL When 250. mL of a 0.15 M solution of ammonium sulfide (NH4)2S is poured into 120. mL of a 0.053 M solution of cadmium sulfate CdSO4, how many grams of a yellow precipitate of cadmium sulfide CdS are formed? The other product is (NH4)2SO4. (Hint: Write out and balance the equation. Is this a limiting reagent problem? ) (a) 5.4 g (b) 0.92 g (c) 2.6 g (d) 1.9 g (e) 530 g Answers to Chapter 3 1. (a) 2. (b) 3. (d) 4. (e) 5. (e) 6. (e) 7. (a) 8. (d) 9. (c) 10. (a) 11. (c) 12. (d) 13. (d) 14. (d) 15. (b) 16. (d) 17. (a) 18. (c) 19. (d) 20. (b) 21. (b) 22. (b) 23. (a) 24. (b) 1. What alkaline earth metal is located in period 3? (a) Li (b) Na (c) Ca (d) Mg (e) Sr Which of the following is classified as a metal? (a) Ge (b) As (c) F (d) V (e) Ar Which of the following is a weak acid? (a) H2SO4 (b) HClO3 (c) HF (d) HCl (e) HNO3 4. Which one of the following is likely to be the most soluble base? (c) When ammonia is added to water. 15. What is the net ionic equation for the acid-base reaction that occurs when nitric acid is added to copper(II) hydroxide? (a) H+(aq) + OH-(aq) H2O(l) (b) 2H+(aq) + Cu(OH)2(s) Cu2+(aq) + 2H2O(l) (c) 2HNO3(aq) +Cu(OH)2(s) Cu(NO3)2(s) + 2H2O(l) (d) 2H+(aq) + 2NO3-(aq) + Cu2+(aq) + 2OH-(aq) Cu(NO3)2(s) + 2H2O(l) (e) 2H+(aq) + 2NO3-(aq) + Cu2+(aq) + 2OH-(aq) Cu2+(aq) + 2NO3-(aq) + 2H2O(l) Which of the following statements is FALSE given the following net ionic equation? H3PO4(aq) + 3OH-(aq) PO43-(aq) + 3H2O(l) (a) If all the water evaporated away. (c) The base involved must be a strong soluble base. +3 (c) H2PO2-. is an insoluble base. Consider the following reaction: NH3(g) + H2O(l) NH4+(aq) + OH-(aq) Which one of the following statements is false? (a) The double arrows indicate that ammonia. 16. Consider the following reaction: 4NH3 + 5O2 4NO + 6H2O The element being oxidized and the oxidizing agent are: (a) N and NH3 (b) N and O2 14. . (a) 0 (b) +2 (c) +4 (d) -2 (e) some other value 17. (e) This is classified as a neutralization reaction. +1 (d) SeO32-. (d) All of the IA and IIA metal hydroxides are soluble. Which one of the following salts is insoluble? (a) NH4Cl (b) Ca(NO3)2 (c) BaCO3 (d) Na2S (e) Zn(CH3COO)2 8. (b) The reaction is reversible. an acid reacts with base to produce a salt and H2O. (e)This could be the net ionic equation for H3PO4 reacting with Al(OH)3. +4 (e) Cu(NO3)2. NH3. (c) The base. (a) BaS2O4 (b) BaSO3 (c) BaSO2 (d) BaSO4 (e) BaS 10. (b) The acid involved must be a strong electrolyte.(a) Ca(OH)2 (b) Cu(OH)2 (c) Ga(OH)3 (d) Zn(OH)2 (e) Zr(OH)3 5. is only very slightly soluble in water. (b) Solutions of weak acids always have lower concentrations of H+ than solutions of strong acids. (d) When solutions of NH4Cl and NaOH are mixed. HCl. is a weak electrolyte. The precipitate formed when barium chloride is treated with sulfuric acid is _______ . What is the net ionic equation for the acid-base reaction that occurs when acetic acid and potassium hydroxide solutions are mixed? (a) H+(aq) + OH-(aq) H2O(l) (b) H+(aq) + KOH(s) K+(aq) + H2O(l) (c) CH3COOH(aq) + KOH(s) KCH3COO(aq) + H2O(l) (d) CH3COO-(aq) + H+(aq) + K+(aq) + OH-(aq) K+(aq) + CH3COO-(aq) + H2O(l) (e) CH3COOH(aq) + OH-(aq) CH3COO-(aq) + H2O(l) 13. (c) There are several common acids that are insoluble. the salt remaining could possibly be Na3PO4.and NO3(d) Na+ only (e) Na+ and NO311. What salt is formed in the following acid/base reaction? HClO3 + Ba(OH)2 (a) BaCl2 (b) ClOBa (c) H2O (d) BaClO3 (e) Ba(ClO3)2 9. (e) Ammonia is considered to be a weak base. (e) In a neutralization reaction. but not in the net ionic equation. NH4+ and OHions are produced in a 1:1 ratio. Which of the following statements is FALSE given the following net ionic equation? 2H+(aq) + Cu(OH)2(s) Cu2+(aq) + 2H2O(l) (a) If all the water evaporated away. 12. (e) All weak acids are insoluble. the net ionic equation is always H+ + OHH2O (b) "Spectator ions" appear in the total ionic equation for a reaction. H3PO4. and HNO3 are all examples of strong acids. 7. (d) This is classified as a neutralization reaction. +6 (b) NH3. (c) HF. some ammonia is produced. (b) The acid. The spectator ion(s) in the following reaction is/are: Na2CO3(aq) + Ba(NO3)2(aq) BaCO3(s) + 2NaNO3(aq) (a) Na+ and Ba2+ (b) Ba2+ and CO32(c) CO32. 6. (d) This could be the net ionic equation for HNO3 reacting with Cu(OH)2. the salt remaining could possibly be CuS. Which assignment of oxidation number is INCORRECT for the blinking element? (a) K2Cr2O7. Determine the oxidation number of carbon in K2CO3. Cu(OH)2. Which one of the following statements is FALSE? (a) For the reaction of a strong acid with a strong soluble base. Which one of the following statements is TRUE? (a) One mole of any acid will ionize completely in aqueous solution to produce one mole of H+ ions. (d) Titration is a process which can be used to determine the concentration of a solution. +2 18. (d) the diameter of an electron is approximately equal to that of the nucleus. 2. One of the isotopes consists of atoms having a mass of 84. one with mass = 64. Answers to Chapter 4 1. (e) 13.9 amu (d) 65. How much energy is emitted as the excited electron falls to the lower energy level? (a) 7.3 (b) 2 only (c) 3 only (d) 1. The Heisenberg Principle states that _____________.4 amu (c) 64. (a) 1. Ag+ and Au+ 20. (e) the same masses. (c) 17. 4. (c) the same number of electrons.43 x 10-19 J (b) 5. (c) 8.06 x 1022 s-1 (e) 6. the other of 86.912 amu. Which name/formula combination is WRONG? (a) chlorous acid / HClO2 (b) dinitrogen tetroxide / N2O4 (c) ammonium nitrate / NH4NO3 (d) copper(II) periodate / CuIO4 (e) potassium permanganate / KMnO4 (d) the same mass numbers. Naturally occurring rubidium consists of just two isotopes. 3.2 (e) 2.3 10. Cu2+.67 x 1015 s-1 8. (b) the positively charged parts of atoms are moving about with a velocity approaching the speed of light.676 x 10-7 m. (d) electrons of atoms in their ground states enter energetically equivalent sets of orbitals singly before they pair up in any orbital of the set. (b) 6.50 x 10-6 cm? (a) 2. (e) 11.05 x 10-19 J (d) 3. The emission spectrum of gold shows a line of wavelength 2. (e) electrons travel in circular orbits around the nucleus. 5. (c) 12. (c) the positively charged parts of atoms are extremely small and extremely heavy particles. (a) 20. The neutral atoms of all of the isotopes of the same element have (a) different numbers of protons. (d) 2.3 amu (b) 64.16 x 10-20 J 9. 11.2. (b) the same number of protons. Given that the Activity Series is: Na>Mg>Cu>Ag>Au.10 x 104 s-1 (c) 4.32 amu? (a) 65.30 x 10-20 J (c) 6. . (a) 7. From these experiments he concluded that: (a) electrons are massive particles.60 x 10-20 J (e) 5. (d) 10. (b) two atoms of the same element must have the same number of protons.901 amu. What is the frequency of light having a wavelength of 4. (b) 18. Consider the species 72Zn. (d) 3.23 amu (26. Ag+ and Au+ (e) Na+. (d) 1. (b) equal numbers of neutrons.8 amu 6. (a) 21. 75As and 74Ge. (e) the same mass number. (c) it is impossible to determine accurately both the position and momentum of an electron simultaneously.0 amu (e) 64. (c) 4.84 x 10-12 s-1 (b) 2. Which of the following has a positive charge? (a) proton (b) neutron (c) anion (d) electron (e) atom Rutherford carried out experiments in which a beam of alpha particles was directed at a thin piece of metal foil. (a) 16. and one with mass = 65. What is the percent natural abundance of the heavier isotope? (a) 15% (b) 28% (c) 37% (d) 72% (e) 85% 7. These species have: (a) the same number of electrons.0%). What is the atomic weight of a hypothetical element consisting of two isotopes. (d) the same number of protons and neutrons. (e) 9. (a) no two electrons in the same atom can have the same set of four quantum numbers. (c) the same number of neutrons. (3) An electron can jump from the K shell (n = 1 major energy level) to the M shell (n = 3 major energy level) by emitting radiation of a definite frequency.29 x 1014 s-1 (d) 1. (b) 14. (2) The lowest energy orbits are those closest to the nucleus. (e) charged atoms (ions) must generate a magnetic field when they are in motion. (b) 19. (e) 15. (a) 5. Which name/formula combination is WRONG? (a) phosphorous acid / H3PO4 (b) nitrogen oxide / NO (c) acetate ion / CH3COO(d) sodium chromate / Na2CrO4 (e) calcium hypobromite / Ca(BrO)2 21. which one of the following answers represents the ions that would not be displaced from aqueous solution (reduced) by metallic magnesium? (a) Na+ (b) Cu2+ (c) Cu2+ and Au+ (d) Cu2+. Which of the responses contains all the statements that are consistent with the Bohr theory of the atom (and no others)? (1) An electron can remain in a particular orbit as long as it continually absorbs radiation of a definite frequency.(c) O and NH3 (d) O and O2 (e) H and NH3 19. . Answers to Chapter 5 1. (a) 18. ml = 9. . (e) The magnetic quantum number is related to the orientation of atomic orbitals in space. diamagnetic (d) 0. 2. (a) 2. (c) 3. Which atomic orbital is spherical in shape? (Note: you should know and be able to recognize the shapes of the s orbital. paramagnetic 18. dyz... (b) 20. In the ground state of a cobalt atom there are _____ unpaired electrons and the atom is _____.. Select the term best describing the series of elements: Mn.. (e) 8. ml = 0. (b) l = subsidiary (or azimuthal) quantum number. ml = 0. ml = -1. (n+1) (c) ml = magnetic quantum number. The outer electronic configuration ns2np4 corresponds to which one of the following elements in its ground state? (a) As (b) Ca (c) Cr (d) Br (e) S 17. . The maximum number of electrons that can be accommodated in a sublevel for which l = 3 is: (a) 2 (b) 10 (c) 6 (d) 14 (e) 8 The ground state electron configuration for arsenic is: (a) [Ar] 4s2 4p13 (b) [Kr] 4s2 4p1 (c) 1s2 2s2 2p6 3s2 3p6 3d12 4s2 4p1 (d) 1s2 2s2 2p6 3s2 3p6 4s2 3d8 4p5 (e) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 Which of the following electron configurations is correct for nickel? (a) [Ar] 4s1 3d8 (b) [Kr] 4s1 4d8 (c) [Kr] 4s1 3d8 (d) [Kr] 4s2 3d8 (e) [Ar] 4s2 3d8 16. l = 0. 0. 2. (b) 10. Which element has the largest atomic radius? (a) Li (b) Na (c) Rb (d) F (e) I 4. . Co. l = 2. px. ml = (-l). dxz. (c) the total number of electrons in s orbitals. The atom having the valence-shell configuration 4s2 4p5 would be in: (a) Group VIA and Period 5 (b) Group IVB and Period 4 (c) Group VIB and Period 7 (d) Group VIIA and Period 4 (e) Group VIIB and Period 4 2. .. (a) 3. .. Fe. dx2-y2 and dz2 orbitals. 12. Which one of the following sets of quantum numbers could be those of the distinguishing (last) electron of Mo? (a) n = 4. and pz orbitals.. l = 1. ms = -1/2 (c) n = 4. Which of the following terms accurately describes the energy associated with the process: Li(g) Li+(g) + e(a) electron affinity (b) binding energy (c) ionization energy (d) electronegativity (e) none of these 5.. The species that contains 24 protons. (b) 7.. (e) 15. Cu. (e) 16. and dxy. 8 in the second energy level and 8 in the third energy level. l = 1. paramagnetic (b) 5. 15. . . (d) the total number of electrons in p orbitals. ms = +1/2 (d) n = 5. (c) 19. (+l) (d) ms = spin quantum number. This information does not necessarily tell us: (a) the atomic number of the element. 3. (b) anything about the element's chemical properties. (a) 13. (c) 11. (d) 14.. A neutral atom of an element has 2 electrons in the first energy level. ms = +1/2 19. Ni. How many p electrons are there in an atom of rubidium? (a) 12 (b) 18 (c) 24 (d) 9 (e) 6 20..Which statement about the four quantum numbers which describe electrons in atoms is incorrect? (a) n = principal quantum number. l = 2. (c) 4. 26 neutrons and 22 electrons would be represented by the symbol: (a) 50V3+ (b) 26Cr2+ (c) 50Cr2+ (d) 50Mn2+ (e) none of these 6. . ms = +1/2 or -1/2. (b) 12. (e) 1. Which element has the lowest first ionization energy? (a) He (b) Ne (c) Ar 14. (a) d-transition metals (b) representative elements (c) metalloids (d) alkaline earth metals (e) halogens 3. (c) 5. ms = +1/2 (b) n = 5. (e) the number of neutrons in the nucleus of an atom of the element. 3. diamagnetic (e) 2. paramagnetic (c) 2. (e) 17. (a) 9. ms = -1/2 (e) n = 3. n = 1.. l = 2. (d) 6. ml = +2. py.) (a) 2s (b) 3p (c) 3d (d) 4f (e) they are all spherical 13. Which of the following elements has the greatest attraction for electrons in a covalent bond? (a) Ge (b) As (c) Se (d) Br (e) Bi Which statement is wrong? (a) The atomic weight of carbon is about 12. Which of these isoelectronic species has the smallest radius? (a) Br(b) Sr2+ (c) Rb+ (d) Se2(e) They are all the same size because they have the same number of electrons. (b) located in the outermost occupied major energy level. (c) A phosphorus atom is larger than an antimony atom. 19F. (d) The radius of a sodium atom is larger than that of a sodium cation.and 20Ne. (a) in s orbitals only. (e) Oxygen has a less negative electron affinity than fluorine.3 (b) Br . (c) 4. (b)The sum of the neutrons in all three species is 27. (d) 10.2 (e) Tl . (d) Both 19F. 11. (a) 3. 9. (c) 6. Answers to Chapter 6 1. (b) 9. (d) 8. which of the following statements is correct? (a) All three species contain 10 electrons.and 20Ne contain 20 neutrons. (c) The sum of the protons in all three species is 28. (c) 11.you may be responsible for this. (e) located in the innermost occupied shell. (d) 1 triple bond between C and N. The correct dot formulation for nitrogen trichloride has: 10. 1 N-H bond and 2 lone pairs of electrons on the C atom. (e) 1 triple bond between C and N. Which of the following pairs of elements and valence electrons is incorrect? (a) Al . 1 C-N bond. (b) The most stable ion of lithium is Li+. (c) located closest to the nucleus. (c) 1 C-H bond. 2.(d) Kr (e) Xe 7. Which of the following is the most basic oxide? (a) N2O3 (b) N2O5 (c) P4O6 (d) P4O10 (e) Bi2O5 8. (d) located in d orbitals. (b) 1 C-H bond. (c) 5.) (a) Be (b) B (c) C (d) N (e) O 8. Which one of the compounds below is most likely to be ionic? (a) GaAs (b) ScCl3 (c) NO2 (d) CCl4 (e) ClO2 The correct electron-dot formulation for hydrogen cyanide shows: (a) 2 double bonds and two lone pairs of electrons on the N atom. 3. The valence electrons of representative elements are . (e) none of the above 4. Which element has the highest first ionization energy? (Note: this is an exception to the general trend . 6.7 (c) S . (e) 7. Which of the following does not have a noble gas electron configuration? (or Which of the following is not isoelectronic with a noble gas?) (a) S2(b) Ba+ (c) Al3+ (d) Sb3(e) Sc3+ 5. (d) 13. 2 lone pairs of electrons on the C atom and 3 lone pairs of electrons on the N atom.4 (d) Sr . Which of the following is an ionic hydride? (a) PH3 (b) H2S (c) HI (d) KH (e) CH4 13. 1 C-H bond and 1 lone pair of electrons on the N atom. (d) 2.3 With regard to the species 16O2-. 1 C=N bond. 1 lone pair of electrons on the C atom and 1 lone pair of electrons on the N atom. All of the following properties of the alkaline earth metals increase going down the group except (a) atomic radius (b) first ionization energy (c) ionic radius (d) atomic mass (e) atomic volume 12. 9. Which one of the formulas for ionic compounds below is incorrect? (a) SrCl2 (b) Cs2S (c) AlCl3 (d) Al3P2 (e) CaSe Which is classified as nonpolar covalent? (a) the H-I bond in HI (b) the H-S bond in H2S (c) the P-Cl bond in PCl3 (d) the N-Cl bond in NCl3 (e) the N-H bond in NH3 7. (e) 1.see notes. (b) 12. What hybridization is predicted for sulfur in the HSO3ion? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 10. Which polyatomic ion is incorrectly matched with its ionic geometry? (a) SiCl62. The electronic structure of the SO2 molecule is best represented as a resonance hybrid of ____ equivalent structures.octahedral 4.tetrahedral (e) SO42.angular (d) NH4+ . Draw one of the resonance structures of SO3.. 2 N=Cl bonds and 7 lone pairs of electrons.pyramidal (d) ClO3. (c) 11. (a) 2 (b) 3 (c) 4 (d) 5 (e) This molecule does not exhibit resonance.trigonal planar (d) AsF5 . The formal charge of S is (a) +2 (b) +1 (c) 0 (d) -1 (e) -2 15. (d) two equivalent resonance forms. Which one of the following violates the octet rule? (a) PCl3 (b) CBr4 (c) NF3 (d) OF2 (e) AsF5 7. (c) four single bonds around the central carbon atom. (e) three equivalent resonance forms. (a) NO2. you would see: (a) two double bonds around the central carbon atom. 1 N=Cl bond and 8 lone pairs of electrons. Choose the molecule that is incorrectly matched with the electronic geometry about the central atom. (b) 5. Answers to Chapter 7 1. After drawing the correct Lewis dot structure(s).. (d) 14..tetrahedral 6. (a) 10.10.trigonal bipyramidal (e) SeF6 . Which of the following pairs of molecules and their molecular geometries is WRONG? (a) NF3 .tetrahedral (e) BrO4. (b) 2.. 13. Consider the bicarbonate ion (also called the hydrogen carbonate ion).trigonal bipyramidal (b) PH4+ . What is the hybridization of A? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 9.. the number of lone pairs of electrons around the central oxygen atom is (a) 0 (b) 1 (c) 2 (d) 3 (e) 4 12. What is the total number of electrons in the correct Lewis dot formula of the sulfite ion? (a) 8 (b) 24 (c) 26 (d) 30 (e) 32 (d) NH3 . (e) 9.. (a) 15. (d) 2 N-Cl bonds. (c) 1 N-Cl bond. (a) CF4 .tetrahedral (c) SO32. (c) 12.tetrahedral (e) PF3 . (e) 1. Choose the species that is incorrectly matched with the electronic geometry about the central atom. 8..tetrahedral (b) BeBr2 .tetrahedral (c) ClO2.trigonal planar (b) H2O . Which molecule has a linear arrangement of all component atoms? (a) CH4 (b) H2O (c) CO2 (d) NH3 (e) BF3 5. 14..linear (c) H2O . Which of the following four molecules are polar: PH3 OF2 HF SO3? (a) all except SO3 (b) only HF (c) only HF and OF2 11. (b) three single bonds around the central carbon atom. (d) 6. (d) 7. (a) 4. In the Lewis structure for the OF2 molecule. (a) 13.tetrahedral 3. Which of the following species is planar? (a) NH3 (b) H3O+ (c) SO32(d) PF3 (e) NO3What kind of hybrid orbitals are utilized by the carbon atom in CF4 molecules? (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 A neutral molecule having the general formula AB3 has two unshared pair of electrons on A.trigonal planar (b) ClO4. (a) 3 N-Cl bonds and 10 lone pairs of electrons. (c) 3. (e) 3 N-Cl bonds and 9 lone pairs of electrons.bent (c) BF3 . (b) 3 N=Cl bonds and 6 lone pairs of electrons.tetrahedral . (b) 8.pyramidal 2. 0 (c) 2. 4. Which statement is false? A sigma molecular orbital (a) may result from overlap of p atomic orbitals perpendicular to the molecular axis (side-on). 4 (d) 2. 3 bonds. (d) One carbon is described by sp3 hybridization. (a) 4. (e) Both oxygens are described by sp3 hybridization.5 (b) 2. The F-S-F bond angles in SF6 are ______. 4. What is the correct electron configuration for the molecular ion.0 (c) 1.(d) none of these (e) all of these 11. Carbon monoxide has ten bonding electrons and four antibonding electrons. 1 bond. no bonds. B2+? (a) 1s2 *1s2 2s2 *2s2 2p2 (b) 1s2 *1s2 2s2 *2s2 2py2 (c) 1s2 *1s2 2s2 *2s2 2py2 2pz1 (d) 1s2 *1s2 2s2 *2s2 2p1 2py1 (e) none of the above. (d) 9. and 5 bonds. (e) 7. 4 (c) 1. (b) may result from overlap of p atomic orbitals along the molecular axis (head-on). (a) 109o28' (b) 120o only (c) 90o and 120o (d) 45o and 90o (e) 90o and 180o 13. (a) One carbon is described by sp2 hybridization. (e) may be either bonding or antibonding. (e) 2. 5 (b) 2. the bonding orbitals are always lower in energy than the antibonding orbitals. 5. Therefore it has a bond order of (a) 3 (b) 7 (c) 1 (d) 5/2 (e) 2 3. 2 16. Draw the molecular orbital diagram for the molecular ion. 3 (b) 3. 2. (d) may result from overlap of one s and one p atomic orbitals. (c) 8. Draw a complete line-bond or electron-dot formula for acetic acid and then decide which statement is incorrect. (c) may result from overlap of two s atomic orbitals. and 4 bonds. and 6 bonds. (b) When one considers the molecular orbitals resulting from the overlap of any two specific atomic orbitals. The perchloric acid molecule contains: (a) 13 lone pairs. no bonds. A (pi) bond is the result of the (a) overlap of two s orbitals. PCl3.5 (d) 2. The number of unpaired electrons in the B2 molecule is _____. The number of electrons in the 2p molecular orbital is: (a) 0 (b) 1 (c) 2 (d) 3 (e) 4 6. N2+. (c) 8 lone pairs. 3. A triple bond contains ___ sigma bond(s) and ___ pi bond(s). (b) 12. 5 (e) another combination 14. Which response contains all of the characteristics listed that should apply to phosphorus trichloride. (b) The molecule contains only one bond. Which molecule is nonpolar? (a) H2Se (b) BeH2 (c) PF3 (d) CHCl3 (e) SO2 12. (c) 3. Which one of the following statements is false? (a) Valence bond theory and molecular orbital theory can be described as two different views of the same thing. 17. 1. and 7 bonds. (e) . (e) 17. (c) Molecular orbitals are generally described as being more delocalized than hybridized atomic orbitals. (c) 5. Draw the molecular orbital diagram for B2. (a) 6. (e) 11 lone pairs. What is the bond order in O2+? (a) 3. and no other characteristics? (1) trigonal planar (2) one unshared pair of electrons on P (3) sp2 hybridized at P (4) polar molecule (5) polar bonds (a) 1. (a) zero (b) 1 (c) 2 (d) 3 (e) 4 8. (a) 0. (d) 2 lone pairs. 1 (d) 1. Answers to Chapter 8 1. 2. (b) overlap of an s and a p orbital. (c) overlap of two p orbitals along their axes.5 (e) 0 7. (a) 11. (e) sidewise overlap of two s orbitals. Which of the following is the correct electron configuration for C2? (a) 1s2 2s2 2py2 *1s2 *2s2 *2py2 (b) 1s2 *1s2 2s2 *2s2 2py2 *2pz1 2p1 (c) 1s2 *1s2 2s2 *2s2 2py2 2pz2 (d) 1s2 *1s2 2s2 *2s2 2py1 2pz1 (e) 1s2 *1s2 2s2 *2s2 2py1 *2py1 2pz1 *2pz1 4. (b) 9 lone pairs. (c) 10. 15. (e) 13. and 4 bonds. (d) sidewise overlap of two parallel p orbitals. 2 bonds. (c) The molecule contains four lone pairs of valence electrons. (d) 14. 2 (e) 3. (d) 15. (d) 16. 9. (b) The antibonding molecular orbitals have more of the character of the more electropositive element than of the more electronegative element. due to electrons occupying the resulting bonding molecular orbital. (c) an electron-pair acceptor. (c) 5. Antibonding molecular orbitals are produced by (a) constructive interaction of atomic orbitals. 7. (d) 2. (e) accepts a share in a pair of electrons.ions in aqueous solution. (d) HF is a base and H3O+ is its conjugate acid. (d) produces OH. (c) a source of OH. 9. (e) HF is a base and F. 4. (b) a hydroxide donor.and HOBr (e) H2O and HOBr 5. (d) 9. 4.(d) One of the shortcomings of molecular orbital theory is its inability to account for a triple bond in the nitrogen molecule. (c) Their molecular orbital diagrams are more symmetrical than those of homonuclear diatomic molecules. (b) destructive interaction of atomic orbitals. Answers to Chapter 10 1. (c) 10.+ Al3+ Al(NO3)3 Which one of the following is an amphoteric metal hydroxide? (a) KOH (b) Ba(OH)2 (c) Pb(OH)2 (d) LiOH (e) Mg(OH)2 10. In the equation: HF + H2O H3O+ + F(a) H2O is a base and HF is its conjugate acid. (a) 6. (d) a sourse of H+ ions in water. Arrhenius defined an acid as: (a) a species that can donate a proton. (c) 8.17 g . O2. a base _____ . (a) 9. (c) 4.is its conjugate base. (c) makes available a share in a pair of electrons. (c) the overlap of the atomic orbitals of two negative ions (d) all of these (e) none of these Which statement regarding stable heteronuclear diatomic molecules is false? (a) All have bond orders greater than zero. 3 (b) 1. (b) 10. (d) 7. (e) a species that can accept a pair of electrons. 8. (b) 6. Answers to Chapter 9 (a) 2. a base is defined as: (a) a proton donor. (d) 8. (d) 7. 10. the more polar the resulting bond is.+ H2O (d) H+ + OHH2O (e) 3NO3. For the system shown here: HOBr + OHH2O + OBrBronsted would classify the base species as: (a) OH. 3. According to the Lewis theory.and HOBr (b) H2O and OH(c) OBr. (c) 4. Which is the strongest acid? (a) HClO4 (b) HClO3 (c) HClO2 (d) HClO (e) HF Which of these species is probably the weakest acid? (a) HCl (b) H3PO4 (c) H2PO4(d) HPO42(e) HNO3 Consider the neutralization reactions between the following acid-base pairs in dilute aqueous solutions: (1) CH3COOH + NaOH (2) HNO3 + Mg(OH)2 (3) H3PO4 + Ba(OH)2 (4) HCl + KOH (5) H2CO3 + LiOH For which of the reactions is the net ionic equation: H+ + OHH2O ? (a) 1. (e) One of the shortcomings of valence bond theory is its inability to account for the paramagnetism of the oxygen molecule. (d) The bonding molecular orbitals have more of the character of the more electronegative element than of the less electronegative element. (e) 5. How many grams of Ca(OH)2 are contained in 1500 mL of 0.and OH(d) OBr. 5 (c) 2. (d) a water-former. (b) is a proton donor. (e) 3. (c) 6.0250 M Ca(OH)2 solution? (a) 3. 1. (c) 1. N2. (e) a proton acceptor. 3 (d) 4 (e) 1 Which one of the following represents the net ionic equation for the reaction of nitric acid with aluminum hydroxide? (a) 3H+ + Al(OH)3 Al3+ + 3H2O (b) 3HNO3 + Al(OH)3 Al(NO3)3 + 3H2O (c) HNO3 + OHNO3.ions in water. (b) a species that can accept a proton.is its conjugate acid. (c) HF is an acid and F. (a) is a proton acceptor. In the Bronsted-Lowry system. (b) H2O is an acid and HF is the conjugate base. (a) 3. 2. (e) The greater is the difference in energy between two overlapping atomic orbitals. (b) two on the left. 13.85 g (d) 2. Use the smallest whole number coefficients possible.) (a) 0.34 g (e) 4.32 g (d) 0.099 M (c) 0.20 eq (c) 2.840 L (d) 7.100 N HNO3 is required to neutralize 50.40 eq (d) 3.) (a) 0.5 x 102 mL (b) 1.033 N (b) 0.75 mL of KOH solution for complete neutralization.12 N 8.14 L 3. 10. Balance the following redox equation in acidic solution with the smallest whole number coefficients possible.78 g (c) 1. How many equivalents of phosphoric acid are contained in 300 mL of 4.00 M HCl? (a) 1.0325 M (b) 0.100 M HCl with 100 mL of 0. What volume of 0. A 0. KHP has one acidic hydrogen.540 g 12. What is the sum of the coefficients? Don't forget coefficients of one.+ SO32MnO2 + SO42When this equation is balanced using the smallest whole number coefficients possible. mL 9.00 liters of 3. the unbalanced equation is: MnO4.859 g What is the oxidation number for carbon in CaC2O4? (a) 0 (b) +2 (c) +3 (d) +4 (e) +6 11.0 L (c) 0. Consider the following unbalanced equation in acidic solution: NaClO3 + H2O + I2 HIO3 + NaCl 4.2 x 102 mL 5.6745 gram sample of KHP reacts with 41.5 g of (COOH)2 in 3000 mL of solution? (Assume the (COOH)2 is to be completely neutralized in an acidbase reaction. What volume of 12.0 mL (c) 100. (b) 0.500 M KOH? (a) 0. What is the molarity of the salt produced in the reaction of 200 mL of 0. what is the sum of all the coefficients? (Do not forget coefficients of one.0 mL of 0. When the following equation is balanced with the smallest possible set of integers.090 N (e) 0. what is the sum of all the coefficients? (Do not forget coefficients of one. H2SO4(aq) + HI(aq) I2(s) + SO2(g) (a) 7 (b) 9 (c) 11 (d) 13 (e) 5 For the reaction between permanganate ion and sufite ion in basic solution.) (a) 0.0 mL (b) 75. the number of OH.0 mL of a 0.660 g (e) 0.158 M (b) 0.0472 M (c) 0.150 N solution of Ba(OH)2? (a) 50. Balance the molecular equation for the following redox reaction.6 M HCl must be added to enough water to prepare 5. What is the sum of all the coefficients? (Do not forget coefficients of one.139 M (d) 0.0935 M What volume of 0.066 N (d) 0.061 M 6.4 x 103 mL (c) 83 mL (d) 7.56 L (e) 2.+ Se2MnO2 + Se (in basic solution) (a) 20 (b) 22 (c) 24 (d) 26 (e) 28 16.25 g 2.19 L (b) 21.079 M (e) 0.(b) 2. (e) four on the left.045 N (c) 0.5 x 102 mL (e) 5.00 M phosphoric acid? (Assume the acid is to be completely neutralized by a base. (c) three on the right.0864 M (e) 0. When the following equation is balanced with the smallest possible set of integers.25 M H3PO4 solution? (a) 2. (d) four on the right. mL (d) 125 mL (e) 150.50 M KOH would be required to neutralize completely 500 mL of 0.) Cu + SO42Cu2+ + SO2 (in acidic solution) (a) 9 (b) 10 (c) 11 (d) 12 (e) 13 14.0667 M (d) 0.+ H2S Cr3+ + S (in acidic solution) (a) 13 (b) 24 (c) 19 (d) 7 (e) 29 15. .80 eq 7.ions is (a) two on the right.) MnO4. How many grams of NaOH would be required to neutralize all the acid in 75.270 g (c) 1.600 eq (b) 1.) Cr2O72. Calculate the normality of a solution that contains 4. What is the molarity of the KOH solution? (Molecular weight of KHP = 204 g/mol.60 eq (e) 4.0900 N H2SO4? (a) 0. 100 N nitrous acid. Consider the following unbalanced net ionic equation: NO2. NaNO2.(in acidic solution) (a) 0.615 g of Zn according to the following unbalanced equation: Zn + ClOZn(OH)2 + Cl. (b) 18. (b) Gases can be expanded without limit.00 g of K2Cr2O7 in 200 mL of solution.251 M (e) 0.8 L (e) 80. .34 g (b) 23.0 mL sample of 0.510 N (d) 0.126 M (c) 0. (d) 5. (e) Pressure must be exerted on a sample of a gas in order to confine it.441 N (e) 0.ions oxidize Fe2+ into Fe3+ ions and are reduced to Mn2+ ions under acidic conditions? (a) 4. Which statement is false? (a) The density of a gas is constant as long as its temperature remains constant. How many grams of I2 were contained in the I2 solution? (a) 0.8 mL 25.A 25. which will be used in the following unbalanced reaction? Cr2O72.500 N 22.2 liters under a pressure of 1240 torr at 25oC.70 mL (d) 67.317 g 17.4 mL (b) 37. A sample of oxygen occupies 47. (e) 17. (b) 9.226 N (c) 0. (b) 19.791 N A solution of nitrous acid was standardized in a reaction where HNO2 NO3.+ MnO4NO3. (d) The molecular weight of a gaseous compound is a non-variable quantity.+ Mn2+ (in acidic solution) What is the molarity of a sodium nitrite.00 mL of the solution is required to react with 0. What is the normality of this Na2C2O4 solution? (a) 0.0502 M (d) 0.4 mL (c) 7.+ Mn2+ MnO2 (in basic solution) (a) 19 (b) 16 (c) 13 (d) 11 (e) 7 Answers to Chapter 11 1.236 g (d) 0.3 L (c) 32. (a) 24.216 N solution of KMnO4? The products in the reaction include MnO2 and IO3-. MnO4. (a) 12.397 g (c) 0. (a) 21.1 mL 24.and its concentration was determined to be 0.2 g/eq (e) 562. At what temperature will it occupy 10.150 N KI solution is required to react in basic solution with 34. (a) 8.264 g (b) 0. (b) 2. solution if 30.108 N (d) 0.0 mL of it just reacts with 0.125 N (d) 0. What volume of this 0.2 L A sample of nitrogen occupies 5.8 mL (e) 30.250 M solution of Na2C2O4 is to be used in a reaction in which the C2O42.00 N (c) 0. 2. (c) 20.+ SO32Cr3+ + SO42.0316 M 18.250 N (b) 1.25 L of 0.6 g/eq (d) 375. (b) 1.9 mL (b) 15.410 M (b) 0.100 N nitrous acid solution would be required to oxidation of 0. (b) 14.159 g (e) 0. (b) 15. (a) 25.8 g/eq (c) 187.733 N (b) 0. What volume of a 0. 23. (d) 6.0625 N (e) 0. (a) 3. (b) 10. A 0. (c) 4.6 g 21.(in basic solution) (a) 0.9 mL (c) 12.3 L (d) 47.8 L (b) 29.0833 M NaClO3 reacted with 30.will be oxidized to CO2.366 N (e) 0. (b) 16.2 mL (e) 49.1 g (e) 70.0 liters at the same pressure? (a) 32oC (b) -109oC 3. What mass of KMnO4 must be dissolved to prepare 1.1 mL of a 0.810 N 20.8 g (c) 115 g (d) 19. (e) 22. (d) 7. (e) 25. What is the equivalent weight (in grams) of copper(II) nitrate for use in a reaction involving the conversion of copper(II) to copper metal? (a) 46.537 N (b) 0.8 g/eq 19. What volume would it occupy at 25oC if the pressure were decreased to 730 torr? (a) 27.200 g of CoCl2 to CoCl3 according to the following net ionic equation? Co2+ + HNO2 Co3+ + NO (in acidic solution) (a) 33. What is the sum of all coefficients when the following net ionic equation is balanced using the smallest whole number coefficients possible? Do not forget coefficients of one.238 grams of KMnO4? (a) 0. Calculate the normality of a NaClO solution if 35.50 liters under a pressure of 900 torr at 25oC. (c) Gases diffuse into each other and mix almost immediately when put into the same container.6 mL (d) 98.9 g/eq (b) 93.110 N KMnO4 solution? It is used in the reaction in which MnO4. (c) 11.275 N (c) 0. What is the normality of a K2Cr2O7 solution prepared by dissolving 5. (e) 23. (a) 13.0 mL of an aqueous solution of I2. 46 L (c) 6.6 mL 15.9 (d) 4. II. is approximately: (a) 6.7 (b) 34 (c) 47 (d) 27 (e) 0. and III (e) another combination 5. What pressure (in atm) would be exerted by 76 g of fluorine gas in a 1.4 g of nitrogen gas and 4.(c) 154oC (d) 269oC (e) 370oC 4.6 g/L (e) 0. (c) The average kinetic energies of different gases are different at the same temperature.0 mL at 24oC on a day when the barometric pressure was 736 torr.0oC is 22. 14.9 mL contains water vapor at a pressure of 10.695 g/L 9. Boyle's law requires that I.4 mL (b) 21.00 liters at 35oC and 740 torr.6 mL (c) 36.363 g (e) 0.9 mL (e) 27.00 atm pressure and a temperature of 25. A 0. which of the following is true? (a) The partial pressure of H2 exceeds that of N2 in the container. (b) The partial pressure of N2 exceeds that of H2 in the container. What volume will it occupy at STP? (a) 6.8 g of oxygen gas in a 200 mL container at 57oC? (a) 4. What volume would the hydrogen occupy if it were dry and at STP? The vapor pressure of water at 24.2 (c) 3.50 mol H2(g) and 0.5 (e) 1.1 atm (c) 19. A container with volume 71.580 g sample of a compound containing only carbon and hydrogen contains 0. If helium effuses through a porous barrier at a rate of 4. (b) The actual volume of the gas molecules themselves is very small compared to the volume occupied by the gas at ordinary temperatures and pressures. What is the molecular (true) formula for the compound? (a) CH3 17. Which one of the following statements is not consistent with the kinetic-molecular theory of gases? (a) Individual gas molecules are relatively far apart.720 g/L (b) 0. (e) The theory explains most of the observed behavior of gases at ordinary temperatures and pressures.2 (b) 3. At STP. Under conditions of fixed temperature and amount of gas.95 g/L at -35 oC and 1020 torr? (a) 24 (b) 11 (c) 72 (d) 120 (e) 44 11. (b) C2H6 (c) C2H5 (d) C4H10 (e) C4H12 A mixture of 90.980 g/L (c) 1.0 moles per minute.50 16.100 g of hydrogen.50 liter vessel at -37oC? (a) 26 atm (b) 4.0 grams of argon has a pressure of 250 torr under conditions of constant temperature and volume.0 grams of CH4 and 10.7 (d) 239 (e) 26.129 g (d) 0. What pressure (in atm) would be exerted by a mixture of 1. in grams per liter.95 L (d) 5.59 L (b) 5.67 L (e) 5. (c) The partial pressures of the two gases remain equal.600 atm (d) 84 (e) 8.6 13. After a period of time.50 mol N2(g) is introduced into a 15. (d) There is no net gain or loss of the total kinetic (translational) energy in collisions between gas molecules. How many grams of the gas are in the container? (a) 0.0oC? (a) 0.222 g 10.18 L 6. The partial pressure of CH4 in torr is: (a) 143 (b) 100 (c) 10.030 A sample of hydrogen gas collected by displacement of water occupied 30. 33. (d) The partial pressures of both gases increase above their initial values.20 (b) 0. A mixture of 0.183 g (c) 0. What is the molecular weight of a pure gaseous compound having a density of 4. .2 atm 8.6 mL of the gas has a mass of 0.4 atm and a temperature of 465oC.39 g/L (d) 16.4 torr. P1V1 = P2V2 II.480 g of carbon and 0. What is the density of ammonia gas at 2.0 liter container having a pinhole leak at 30oC. PV = constant III. 12. at what rate (in moles per minute) would oxygen gas diffuse? (a) 0.087 g. The volume of a sample of nitrogen is 6. The density of chlorine gas at STP. P1/P2 = V2/V1 (a) I only (b) II only (c) III only (d) I.8 mL (d) 25. (a) 32.421 g (b) 0.3 7. (e) The partial pressure of H2 in the container increases above the initial value. 00 atm? (a) HF (b) HCl (c) HBr (d) HI (e) H2SO4 6.4 liters of methane at STP react with 64. (e) Real gases do not always obey the ideal gas laws. Which of the following statements is true? 520oC 25. (a) Each flask has the same number of gas molecules. For a gas.4 liters of carbon dioxide at STP can be produced. (c) 23. What type of interparticle forces holds liquid N2 together? (a) ionic bonding (b) London forces (c) hydrogen bonding (d) dipole-dipole interaction (e) covalent bonding 3. (e) If 22. AsH3.0 (d) 8.6 g (d) 24.0 (e) 1. (b) 19. 2KClO3(s) 2KCl(s) + 3O2(g) (a) 7.(c) 2. 20. (d) 5. H2Te (b) AsH3. (c) If 11. Which one of the following statements about the following reaction is false? CH4(g) + 2O2(g) CO2(g) + 2H2O(g) (a) Every methane molecule that reacts produces two water molecules. (a) 25. (e) None of the above is true. (a) 2.0 g. T (b) P.4 L at any temperature and pressure 24. P Answers to Chapter 12 1. AsH3. (a) 8. respectively. CH3NH2. Calculate the weight of KClO3 that would be required to produce 29. (d) 4. V (e) n. at STP. (d) There are twice as many O2 and H2 molecules as Ne atoms. (c) 9. the combined masses of the products will be 80. (c) 20. (c) 11. (e) 22. (b) 1. Which response includes only those compounds that can exhibit hydrogen bonding? CH4. . 22. Which of the following statements is false? (a) The properties of N2(g) will deviate more from ideality at -100oC than at 100oC.5 g 22. (d) If 16.0 g of oxygen. HF (a) AsH3. (b) The velocity of the gas molecules is the same in each flask. (e) 6. H2Te (d) CH3NH2. (d) 12.5 L of oxygen measured at 127oC and 760 torr. (b) If 32. which pair of variables are inversely proportional to each other (if all other conditions remain constant)? (a) P.2 g (c) 14. H2Te. The ideal gas law predicts that the molar volume (volume of one mole) of gas equals: (a) gRT/PV (b) (MW)P/RT (c) 1/2ms-2 (d) RT/P (e) 22. CH3NH2 (c) CH4. Which probably has the lowest boiling point at 1. The normal boiling point of a liquid is 4.2) liters.82 g (b) 12. (d) 15. (d) 21. HF (e) HF.0 liter flasks are filled with H2. T (d) n. the maximum amount of carbon dioxide produced will be 22. V (c) V. (e) 18. A real gas most closely approaches the behavior of an ideal gas under conditions of: (a) high P and low T (b) low P and high T (c) low P and T (d) high P and T (e) STP 19.0 g. the volume of carbon dioxide produced at STP is (44/16)(11. What total gas volume (in liters) at and 880 torr would result from the decomposition of 33 g of potassium bicarbonate according to the equation: 2KHCO3(s) K2CO3(s) + CO2(g) + H2O(g) (a) 56 L (b) 37 L (c) 10 L (d) 19 L (e) 12 L 21. (d) Molecules of an ideal gas are assumed to have no significant volume. (b) 7. (c) 16.0 g of oxygen reacts with excess methane. (b) Van der Waal's equation corrects for the nonideality of real gases. 23. (e) 10. (d) 13. (c) The density of each gas is the same. (e) 3. H2Te Which of the following boils at the highest temperature? (a) CH4 (b) C2H6 (c) C3H8 (d) C4H10 (e) C5H12 5. Three 1. (d) 24. (c) Molecules of CH4(g) at high pressures and low temperatures have no attractive forces between each other. (d) 14. O2 and Ne.2 liters of methane react with an excess of oxygen.0 g of oxygen.0 g of methane react with 64. (b) 17.4 g (e) 73.41 18. What type of intermolecular forces are due to the attraction between temporary dipoles and their induced temporary dipoles? (a) metallic bond (b) London dispersion (c) hydrogen bond (d) ionic bond (e) covalent bond 2. molecular solid (b) C4H10(s). (c) the temperature at which the gas molecules have more kinetic energy than the molecules in the liquid. (b) At 0oC and 1200 torr. melting 3. ionic solid (d) SiC(s). condensation 6.138 J/goC Specific Heat (gas) = 0.56 (c) 6. For water (m.104 J/goC Calculate the amount of heat that must be released to convert 20. In any cubic lattic. an increase in the size of the open vessel containing the liquid (a) 1 and 2 only (b) 1 and 3 only (c) 1 only (d) 2 only (e) 3 only 8. (a) 77 (b) 68 (c) 64 (d) 57 (e) 50 9.0 g of mercury vapor at 387 oC to liquid mercury at 307oC (in kJ). b. 5. (b) Metallic solids exhibit a wide range of melting points because metallic bonds cover a wide range of bond strength. vaporization 4. Y(g) will spontaneously convert to Y(l). freezing (a) 1. 357oC) Heat of fusion = 11.03 J/goC Calculate the amount of heat (in kJ) that must be absorbed to convert 108 g of ice at 0oC to water at 70oC. (a) At the temperature and pressure at point 4.(a) the temperature at which the vapor pressure equals 760 torr.p. (c) At the pressure and temperature of point 1. an increase in the intermolecular forces in the liquid 3. According to the phase diagram given for Compound Y.p. b. For mercury (m. (a) 61. Y(s) will spontaneously convert to Y(g) and no Y(l) is possible. Y exists as a solid.10 10. Which of the following changes would increase the vapor pressure of a liquid? 1.18 J/goC Specific Heat (gas) = 2. (b) the temperature above which the substance cannot exist as a liquid regardless of the pressure. (e) At the temperature and pressure at point 2. (c) The metallic solid can be viewed as positive ions closely packed in a sea of valence electrons. (d) Most molecular solids melt at lower temperatures than metallic solids. metallic solid .p. 7.9 (b) 6. Which one of the following classifications is incorrect? (a) H2O(s).69 (e) 5. what description is correct? 16. sublimation 5.09 J/goC Specific Heat (liquid) = 4. 0oC. covalent solid (e) S(s). Y(s) Y(g). 2.04 (d) 5. and 3 (b) 4. Y( l) Y(g) Where on a phase diagram can you locate conditions under which only one phase exists? (a) at an intersection of two lines (b) at the normal boiling point (c) at an intersection of three lines (d) in an area bounded by lines (e) at the triple point 13. Which of the following phase changes is(are) endothermic? 1. -39oC. an increase in temperature 2. (d) At the pressure and temperature at point 3. molecular solid (c) KF(s). (e) The interactions among the molecules in molecular solids are generally stronger than those among the particles that define either covalent or ionic crystal lattices.141 J/goC Specific Heat (liquid) = 0. 100oC) Heat of fusion = 333 J/g @ 0oC Heat of vaporization = 2260 J/g @ 100oC Specific Heat (solid) = 2.p. 15. an atom lying at the corner of a unit cell is shared equally by how many unit cells? (a) one (b) two (c) eight (d) four (e) sixteen 14. Which of the following compounds would be expected to have the highest melting point? (a) BaF2 (b) BaCl2 (c) BaBr2 (d) BaI2 (e) H2O 12. Which statement is false? (a) Molecular solids generally have lower melting points than covalent solids. (d) the only temperature at which there can be equilibrium between liquid and gas.6 J/g @ -39oC Heat of vaporization = 292 J/g @ 357oC Specific Heat (solid) = 0. (e) the temperature at which the liquid will usually boil. deposition 2. and 6 (c) 1 and 2 only (d) 4 and 6 only (e) some other combination 11. (a) 7. 2. (c) 7.512 oC/m) (Note: If the Kf and Kb are not given on the exam. (a) 1/3 (b) 0. 2 (c) 3. assume it is water. benzene. are dissolved in 15.64 oC (b) 100. (b) 6. 1. (c) 10. C6H14O3. . Kb = 0. C12H22O11.) (a) 101. 9. The oxygen end of water molecules is attracted toward Ca2+ ions. 5.00 g/mL) and 150 g of diglyme.5 m BaCl2 solution.67 m Calculate the mole fraction of C2H5OH in a solution that contains 46 grams of ethanol.840 m (e) 1. (a) 0.2 g of naphthalene. (b) osmotic pressure of the solute. 3.626 oC (d) 100. Which statement(s) is(are) true? 1.42 (c) 1/2 (d) 2/3 (e) none of these 6. HNO3 ? (a) 1. Calculate the molality of a solution that contains 51.2 grams of water. water and methanol.8 oC (e) 104. (b) 9. (b) 2.) (a) 0. (a) 1 only (b) 2 only (c) 3 only (d) 1 and 2 only (e) 1. (d) 13. 5 (e) none of the above 11.69 m 4.750 m (d) 0.51 oC/m) (a) 240 g/mol (b) 150 g/mol (c) 79 g/mol (d) 61 g/mol (e) 34 g/mol 1. (II) A 0.500 m (c) 0. Consider the following pairs of liquids. (e) 15. What are the ideal van't Hoff factors for the following compounds: Ba(OH)2.6% glucose (C6H12O6) by weight?" (Note: If the question does not give the solvent. (a) 17.42 oC 10. The density of CCl4 is 1. 5. (a) 12.34 m (c) 0. 1. (d) mole fraction of solvent. (e) 16.250 m (b) 0. and 64 grams of methanol. Which one of the following substances can be melted without breaking chemical bonds? (a) sodium sulfate (b) zinc chloride (c) sulfur dioxide (d) silicon dioxide (e) diamond Answers to Chapter 13 1. 2. 1. 2. (III) Pure water freezes at a higher temperature than pure methanol. water and hexane (a) 1.8 oC 12.27 grams of sucrose.81 g of a nonvolatile nonelectrolyte in 90. (a) 11. (c) 14. 4. C6H12O6. Which observation(s) reflect(s) colligative properties? (I) A 0.3 oC (d) 100. The vapor pressure of a solution containing a nonvolatile solute is directly proportional to the (a) molality of the solvent. Which pairs are miscible? 1.6 oC (c) 102. CH3OH. 2.60 g/mL. C10H8. (e) 5.52 m (e) 0. in 500 mL of carbon tetrachloride.1 oC (b) 101. The hydrogen end of water molecules is attracted toward Cl. C6H6. 3 (e) 2.26 m (b) 0. 3. (a) only I (b) only II (c) only III (d) I and II (e) I and III The vapor pressure of pure water at 85oC is 434 torr. What is the molality of a solution labeled "8.5 m NaBr solution has a higher vapor pressure than a 0. (c) molarity of the solvent. If 4. C2H5OH. (c) 8.44 m (d) 0.37 oC at 760 mm Hg. (Assume complete dissociation of the salt. K3PO4. 2 (d) 6. (e) mole fraction of solute. what will be the boiling point of the resulting solution? (Kb for water = 0. (d) 4. a nonvolatile substance? (a) 361 torr (b) 390 torr (c) 425 torr (d) 388 torr (e) 317 torr 8.0 g of water boiled at 100.17. Calculate the approximate initial boiling point (in oC) of a solution of 285 g of magnesium chloride in 2. What is the approximate molecular weight of the substance? (For water. 2 only (b) 2 only (c) 1 only (d) 1. Consider the three statements below. 13.ions.5 m NaOH solution freezes at a lower temperature than pure water. you can find them on the back of the exam envelope. 3 only 3.0 kg of water. What is the vapor pressure at 85oC of a solution prepared from 100 mL of water (density 1. and 3 2. Hydration is a special case of solvation in which the solvent is water.) (a) 103. 1.42 oC (c) 99. A solution made by dissolving 9. C6H12 2. CH3OH 3. (b) 3.73 oC (e) 101. and hexane. 1 (b) 2. 0 atm. All of the following have a standard heat of formation value of zero at 25oC and 1.31 J/mol K. 8. (b) 10. CH3OH. Ho for the following reaction at 298 K is -36.000 oC. A system suffers an increase in internal energy of 80 J and at the same time has 50 J of work done on it. The specific heat of water is 4.0 liters by a constant external pressure of 5. The temperature of the water rose to 23.6 kJ Calculate the amount of work done for the conversion of 1. What is the heat change of the system? (a) +130 J (b) +30 J (c) -130 J (d) -30 J (e) 0 J A 5.0 g of an enzyme of molecular weight 98. m. A 17. (a) -314 kJ/mol (b) -789 kJ/mol (c) -716 kJ/mol (d) -121 kJ/mol (e) -69.0 oC 14. (d) 7. was combusted in the presence of excess oxygen in a bomb calorimeter conaining 4000 g of water.0 oC. (c) 14. of camphor = 178. Assume the specific heat of the final solution is 4.) (a) 353 g/mol (b) 285 g/mol (c) 231 g/mol (d) 185 g/mol (e) 166 g/mol 15.639 oC.86 oC (c) -3.68 x 103 J (e) -494 J Ni(CO)4(g) 9.86 oC (b) +1.6 kJ/mol 6. A 250 mL solution containing 21. R. 1.2 kJ (b) +35. The universal gas constant. is 8. (e) 7. The value of R is 8.55 kJ/mol (d) -81.4 oC) (Note: This is a freezing point depression problem .80 x 103 J (b) 8.4 Answers to Chapter 14 1.48 torr (e) 3.4 g of a polymer in toluene had an osmotic pressure of 0. (a) 8.71 torr 16. at 75oC. HNO3(aq) + NaOH(s) NaNO3(aq) + H2O(l) (a) -63.000 g/mol dissolved in water to give 2600 mL of solution at 30. The kJ. (b) 13. (e) 2.000 oC to 29.0 liters to 4.4 kJ/mol (e) -98.93 oC (e) 0.314 J/mol K. (a) 6. an ideal gas is compressed from 6.) (a) -1.3 mg sample of an organic compound (a nonelectrolyte) was ground up with 420 mg of camphor to form a homogeneous mixture melting at 170.2 kJ (c) -36. A coffee cup calorimeter having a heat capacity of 451 J/oC was used to measure the heat evolved when 0.765 oC. (a) 0.68 torr (c) 1.5 kJ/mol 5.0 oC. (c) 12. Which one of the following thermodynamic quantities is not a state function? (a) Gibbs free energy (b) enthalpy (c) entropy (d) internal energy (e) work 2.000 g/mol (d) 32.72 oC (d) -0.00 m NaCl solution? (Kf = 1. The heat capacity of the calorimeter was 2657 J/oC.00 mole of Ni to Ni(CO)4 in the reaction below.86 oC/m) (Assume complete dissociation of the salt. (d) 15.184 J/goC.000 g/mol 4. Assume that the gases are ideal. The temperature of the water increased from 24. (d) 5. (b) 4. What is the apparent formula weight of the organic compound? (Kf of camphor = 37. Calculate the osmotic pressure associated with 50.7 oC/m.000 g/mol (b) 18. (a) 3. What is the apparent formula weight of the polymer? (a) 15.000 g/mol (c) 26.0300 mol of NaOH(s) was added to 1000 mL of 0.055 atm at 27 oC.00 g/mL. and the addition of solid does not appreciably affect the volume of the solution.18 J/goC. How much work is done on the gas? (a) w = +10 liter atm (b) w = -10 liter atm (c) w = +30 liter atm (d) w = -30 liter atm (e) The answer cannot be calculated.0300 M HNO3 initially at 23.What is the freezing point of an aqueous 1.96 torr (d) 2. 1/2 H2(g) + 1/2 Br2(l) HBr(g) Calculate Eo at 298 K.7 kJ/mol (b) -151 kJ/mol (c) -2.000 g/mol (e) 38. Ni(s) + 4 CO (g) (a) 1. (a) -35. Calculate H (in kJ/mol NaNO3) for this reaction.80 x 103 J (d) -8.68 x 103 J (c) -1.484 torr (b) 1.p.4 kJ (d) -37. 3. (e) 16.note the Kf of camphor camphor is the solvent. At a constant temperature. the density of each solution is 1. (d) 9.000 g sample of methanol.6 kJ (e) +37. (c) 11.0 atm except: (a) N2(g) (b) Fe(s) (c) Ne(g) (d) H(g) (e) Hg(l) . Calculate E for the reaction in kJ/mol. 0 kJ 17. Br2(l) Br2(g) (a) 85oC (b) 373oC (c) 177oC (d) 59oC (e) 44oC 20. What is the standard entropy change of the reaction below at 298 K with each compound at the standard pressure? 192 kJ 158 kJ 197 kJ Horxn for the following reaction at 25.2 kJ 18. The heat of vaporization of freon. SBr4(g) S(g) + 2Br2(l) (a) +152 kJ (b) -56. a reaction occurs that results in an increase in the number of moles of gas. 130. .10. III.16 J/K (e) 239 J/K 19. H = 30.6 2NH3(g ) 192. Ho = +115 kJ and So = +125 J/K. Estimate the boiling point of Br2(l) ( S = 93. (c) spontaneous at temperatures less than 600 K. Ho = +30 kJ. CCl2F2.4 kJ/mol (a) -70. What is the change of entropy for one mole of liquid freon when it vaporizes at 25oC? (Hint: The vaporization process is at equilibrium and what is true for G at equilibrium?) (a) 57.5 (a) -198.7 J/K 15.688 J/K (c) 5.7 J/K (e) 384.1 kJ (e) -86. IV. III.7 J/K (b) 0.0 kJ (b) +70. Calculate the standard heat of formation.0 kJ (e) +140.8 kJ (d) -140.2 kJ/mol at 25oC.7 J/K (b) 76.5 -272 -393. a liquid changes to a gas. and IV 16. (d) spontaneous at temperatures greater than 600 K. (a) I only (b) II only (c) III only (d) IV only (e) I.13 x 103 kJ/K (d) 3. Br2(g) + 3F2(g) 2BrF3(g) Bond Bond Energy Br-Br F-F Br-F (a) -516 kJ (b) -410 kJ (c) -611 kJ (d) -665 kJ (e) -720 kJ 14. II. So = +50 J/K.2 J/K (d) -129. a molecule is broken into two or more smaller molecules.3 The entropy will usually increase when I.8 kJ (d) +37.7 kJ (c) +77. (b) nonspontaneous at all temperatures. Therefore the reaction is: (a) spontaneous at all temperatures.5 13.0 J/K). Calculate Hfo (kJ/mol) (a) -263 kJ (b) 54 kJ (c) 19 kJ (d) -50 kJ (e) 109 kJ 12.0 oC: Fe3O4(s) + CO(g) 3FeO(s) + CO2(g) -1118 110. Calculate Ho for the reaction: Na2O(s) + SO3(g) Na2SO4(g) given the following information: Ho (1) Na(s) + H2O(l) (2) Na2SO4(s) + H2O(l) SO3(g) (3) 2Na2O(s) + 2H2(g) (a) +255 kJ (b) -435 kJ (c) -581 kJ (d) +531 kJ (e) -452 kJ NaOH(s) + 1/2 H2(g) 2NaOH(s) + -146 kJ +418 kJ +259 kJ N2(g) + 3H2(g) So298 (J/mol K) 191. given the following information: 2FeS2(s) + 5O2(g) 2FeO(s) + 4SO2(g) Horxn = -1370 kJ Hfo for SO2(g) = -297 kJ/mol Hfo for FeO(s) = -268 kJ/mol (a) -177 kJ (b) -1550 kJ (c) -774 kJ (d) -686 kJ (e) +808 kJ Estimate the heat of reaction at 298 K for the reaction shown. For the reaction.4 kJ/mol Gfo for SO3(g) = -370. For which of the following reactions would the Ho for the reaction be labeled Hfo? (a) Al(s) + 3/2 H2(g) + 3/2 O2(g) Al(OH)3(s) (b) PCl3(g) + 1/2 O2(g) POCl3(g) (c) 1/2 N2O(g) + 1/4 O2(g) NO(g) (d) CaO(s) + SO2(g) CaSO3(s) (e) The Ho for all these reactions would be labeled o Hf . Calculate Go for the reaction given the following information: 2SO2(g) + O2(g) 2SO3(g) Gfo for SO2(g) = -300. A + B C. Calculate Go for the reaction at 25o. 4Na(s) + 2H2O(l) 11. for FeS2(s).32 J/K (c) 303.0 kJ (c) -670.9 kJ. a solid changes to a liquid. For the following reaction at 25oC. given the average bond energies below. is 17. Hfo. II. (d) the rate of consumption of oxygen equals the rate of consumption of water.(e) spontaneous only at 25oC. Which statement is false? (a) The thermodynamic quantity most easily measured in a "coffee cup" calorimeter is H.0 g sample? (a) 686 kJ (b) 519 kJ (c) 715 kJ (d) 597 kJ (e) 469 kJ 23. 21. (e) CO2 is formed twice as fast as ethane is consumed. How much heat is absorbed in the complete reaction of 3.010 M 0. (b) No work is done in a reaction occurring in a bomb calorimeter. (c) between gases should in all cases be extremely rapid because the average kinetic energy of the molecules is great. For a reaction 2A + B Rate = k[A]2[B] 2C. If the concentration of A is doubled and the concentration of B is halved.020 M/s 0. 2. 2 (b) decrease. when the concentration unit is mol/L? (a) s-1 (b) s (c) L mol-1 s-1 (d) L2 mol-2 s-1 (e) L2 s2 mol-2 Given: A + 3B 2C + D This reaction is first order with respect to reactant A and second order with respect to reactant B. (c) H is sometimes exactly equal to E.030 M/s 3 0. Which statement is incorrect? (a) At constant pressure.091 M 24.015 M 0. 3. The decomposition of carbon disulfide.7 x 10-6 s (c) 3. The decomposition of dimethylether at 504 oC is first order with a half-life of 1570 seconds.010 M 0.005 M/s The rate law for the reaction is: (a) Rate = k[NH4+][NO2-] (b) Rate = k[NH4+]2[NO2-]2 (c) Rate = k[NH4+]2[NO2-] (d) Rate = k[NH4+][NO2-]2 (e) none of the above 5. (b) the rate of formation of CO2 equals the rate of formation of water. to carbon monosulfide.1 x 104 s (e) 2.8 x 105 s (d) 6. What are the units of k for the rate law: Rate = k[A][B]2. The speed of a chemical reaction (a) is constant no matter what the temperature is. (c) Gibbs free energy is a state function. E is negative.7 kJ. (c) water is formed at a rate equal to two-thirds the rate of formation of CO2. (b) the order with respect to A is 2 and the order overall is 2.069 M (c) 0.8 x 10-7 s-1 at 1000oC. is 1372 kJ/mol ethanol. Given the following data for this reaction: NH4+(aq) + NO2-(aq) N2(g) + 2H2O(l) EXPT [NH4+] [NO2-] RATE 1 2 0. 4. after 2. H is negative. CS.13 x 105 kJ (c) 5. What fraction of an initial amount of dimethylether remains after 4710 seconds? (a) 1/3 (b) 1/6 (c) 1/8 (d) 1/16 (e) 1/32 9. (d) For an endothermic process. (b) is independent of the amount of contact surface of a solid involved.84 M (b) 0. SiO2(s) + 3C(s) SiC(s) + 2CO(g) (a) 366 kJ (b) 1.0 minutes. (e) H is equal to E for the reaction: 2H2(g) + O2(g) 2H2O(g) (a) the order with respect to A is 1 and the order overall is 1. the reactant concentration is 0. the rate of the reaction would _____ by a factor of _____. and sulfur is first order with k = 2. (c) the order with respect to A is 2 and the order overall is 3. 6. 4 (d) decrease. CS2.2 kJ 22. CS2 CS + S What is the half-life of this reaction at 1000oC? (a) 5. The half-life for a first-order reaction is 32 s. H= E+P V (b) The thermodynamic symbol for entropy is S. The combustion of ethane (C2H6) is represented by the equation: 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(l) In this reaction: (a) the rate of consumption of ethane is seven times faster than the rate of consumption of oxygen.0 x 107 s (b) 4.00 grams of SiO2 with excess carbon in the reaction below? Ho for the reaction is +624. (d) between ions in aqueous solution is extremely rapid because there are no bonds that need to be broken. (e) varies inversely with the absolute temperature. (e) If the work done by the system is greater than the heat absorbed by the system. (e) the order with respect to B is 2 and the order overall is 3. 2 (c) increase. What was the original concentration if. 1. C2H5OH. 4 (e) not change 7. (d) H is often nearly equal to E. with the rate equation: . How much heat (in kJ) would be liberated by completely burning a 20.020 M 0.06 kJ (d) 1.33 x 104 kJ (e) 31.5 x 106 s 8.062 M? (a) 0.010 M 0. (d) the order with respect to B is 2 and the order overall is 2. The standard heat of combustion of ethanol.020 M 0. (a) increase. (5) The change in internal energy is greater than zero. (d) The energy of reaction B must be greater than the energy of reaction A. 2 A correct reaction mechanism for a given reaction usually is: (a) the same as its balanced chemical equation. (b) At the same temperature the rate of reaction B is greater than the rate of reaction A. the following reaction is found to obey the rate law: Rate = k[NOCl]2: 2NOCl 2NO + Cl2 Consider the three postulated mechanisms given below. (e) The rate of reaction A at 25 oC equals the rate of reaction B at 100 oC. is: (1) A + A (2) A2 + A A2 A3 fast. (a) only I 12. Suppose the activation energy of a certain reaction is 250 kJ/mol. Given that a reaction absorbs energy and has an activation energy of 50 kJ/mol. (e) obvious if its activation energy is known. (III) an increase in the percentate of "high energy" collisions with increasing temperature.(d) 0. which of the following statements are correct? (Hint: Draw the potential energy diagram.) (1) The reverse reaction has an activation energy equal to 50 kJ/mol.067 (c) 15.0 (d) 525 (e) 3 x 10-28 18. Then choose the response that lists all those that are possibly correct and no others. (a) k[A] (b) k[B] (c) k[A][B] (d) k[B]2 (e) k[A][B]2 14. Most reactions are more rapid at high temperatures than at low temperatures. equilibrium slow Overall: 2NOCl 2NO + Cl2 (a) 2. (c) obvious if its reaction order is known. This is consistent with: (I) an increase in the activation energy with increasing temperature. Mechanism slow NOCl NO + Cl 1 Cl + NOCl NOCl2 + NO + Cl2 Overall: 2NOCl 2NO + Cl2 Mechanism 2 2NOCl NO NOCl2 NOCl2 + NO + Cl2 slow fast NOCl2 2NO fast fast Overall: 2NOCl 2NO + Cl2 Mechanism 3 NOCl NOCl + Cl Cl2 NO + Cl NO + fast. If reaction A has an activation energy of 250 kJ and reaction B has an activation energy of 100 kJ. (II) an increase in the rate constant with increasing temperatures. the rate law will be: (a) Rate = k[A]2 (b) Rate = k[A][B] (c) Rate = k[A]2[B] (d) Rate = k[A] (e) Rate = k[A]3 . 16. 2A + B D. 3 (b) 3 (c) 1 (d) 2 (e) 1.075 M (e) 0. what is the energy of reaction E for this step? (a) 22 kJ (b) -22 kJ (c) 52 kJ (d) -52 kJ (e) 126 kJ Suppose the reaction: A + 2B following mechanism: Step 1 Step 2 A+B AB + B AB AB2 occurs by the slow 15. What is the activation energy (in kJ) of a reaction whose rate constant increases by a factor of 100 upon increasing the temperature from 300 K to 360 K? (a) 27 (b) 35 (c) 42 (d) 53 (e) 69 19. If the rate constant at T1 = 300 K is k1 and the rate constant at T2 = 320 K is k2. (2) The reverse reaction has an activation energy less than 50 kJ/mol. A possible mechanism for the reaction. (c) The energy of reaction A must be greater than the energy of reaction B. (b) obvious if its heat of reaction is known. AB2 fast Overall A + 2B AB2 The rate law expression must be Rate = _________. 17. (3) The reverse reaction has an activation energy greater than 50 kJ/mol. (4) The change in internal energy is less than zero.) (a) 3 x 10-29 (b) 0. (d) sometimes difficult to prove. which of the following statements must be correct? (a) If reaction A is exothermic and reaction B is endothermic then reaction A is favored kinetically.13 M 10. 13. (Hint: Solve for k2/k1. (a) (1) and (4) (b) (2) and (4) (c) (3) and (4) (d) (2) and (5) (e) (3) and (5) 11. If the activation energy in the forward direction of an elementary step is 52 kJ and the activation energy in the reverse direction is 74 kJ. then the reaction is __ times faster at 320 K than at 300 K. At 300 K. equilibrium slow C+ (3) A3 + B A + C + D fast According to the mechanism. 0357 mole NH3. (e) 8. 2. Which statement is false? (a) If a reaction is thermodynamically spontaneous it may occur rapidly. and HI in a vessel at 445oC has the following concentrations: [HI] = 2. (b) Qc is less than Kc. (c) both the forward and the reverse reactions have stopped.99 (c) 16. (c) Qc is less than Kc.94 (e) 0. 4 (e) 1. which one of the following statements is absolutely true? (a) The reaction is first order with respect to H2S and second order with respect to O2. 4 (c) 1. Evaluate Kc. (c) Activation energy is a kinetic quantity rather than a thermodynamic quantity. (a) the sum of the concentrations of A and B must equal the sum of the concentrations of C and D.0 (d) 8. (a) 22. Which of the following statements are true? (1) Reactions with more negative values of Go are spontaneous and proceed at a higher rate than those with less negative values of Go. (d) always decreases the rate for a reaction. (e) 2. the system is at equilibrium.50 M and [I2] = 0. Which one of the following statements concerning the reaction quotient. (b) changes the equilibrium concentration of the products. (d) 5. 2. it will not occur spontaneously. (II) lower the activation energy. (3) The activation energy for a reaction does not change significantly as temperature changes. Consider the following reversible reaction. (c) does not affect a reaction energy path. the rate of reaction increases. 2H2S(g) + O2(g) 2S(s) + 2H2O(l). (e) 21. (b) The reaction is fourth order overall. is usually about the same as E for a reaction. 2. is TRUE for the above system? (a) Qc = Kc. Answers to Chapter 16 1. (e) 20.202 (b) 1. (d) 6. (a) only I & II (b) only II & III (c) only III & IV (d) only I & III (e) only II & IV 21. 22. Ea.0 (e) 16 5. 25. (b) If a reaction is thermodynamically spontaneous it may occur slowly. At 445oC.0420 mole N2. (d) the reverse reaction has stopped. more HI will be produced. In a 3. (e) 15. 4 24. A catalyst: (a) actually participates in the reaction. (IV) provide a new path for the reaction. Kc for the following reaction is 0.020. (c) 14. (c) 9. (e) The rate law cannot be determined from the information given. (b) 13. (b) the forward reaction has stopped. 2. it must have a low activation energy. more H2 and I2 will be produced.(b) only II (c) only III (d) only I and II (e) only II and III 20. 4 (b) 3.503 4. 2A + 4B 2NH3(g) 23. (d) 18. (e) 1. 3.10 M. If the equilibrium constant for the reaction A + 2B C + 5/2 D has a value of 4. 2HI(g) H2(g) + I2(g) A mixture of H2. When the system A + B C + D is at equilibrium. (e) always increases the activation energy for a reaction. what is the value of the equilibrium constant for the reaction 2C + 5D at the same temperature? (a) 0. (b) 12. (b) 25. Which items correctly complete the following statment? A catalyst can act in a chemical reaction to: (I) increase the equilibrium constant. (4) Reactions usually occur at faster rates at higher temperatures. For the reaction. (d) 3. (c) The rate law is: rate = k[H2S]2[O2]. I2. (d) the activation energy increases.00 liter container. 2SO3(g) 2SO2(g) + O2(g) The conventional equilibrium constant expression (Kc) for the system as described by the above equation is: (a) [SO2]2/[SO3] (b) [SO2]2[O2]/[SO3]2 (c) [SO3]2/[SO3]2[O2] (d) [SO2][O2] (e) none of these 3. (III) decrease E for the reaction. When the concentration of reactant molecules is increased. 3 (d) 2.25 (b) 0. (b) 24. (d) 17.063 (c) 2. (d) 16. the following amounts are found in equilibrium at 400 oC: 0.0 (d) 4. [H2] = 0. 0. (2) The activation energy. (b) 7. (c) the rate constant increases. N2(g) + 3H2(g) (a) 0. (d) If a reaction is thermodynamically nonspontaneous. The best explanation is: As the reactant concentration increases. (d) The rate law is: rate = k[H2S][O2]. (a) 10. (e) the order of reaction increases. (e) neither the forward nor the reverse reaction has stopped. (a) the average kinetic energy of molecules increases. (d) 11. Qc. 3.0 M.516 mole H2 and 0. . (e) If a reaction is thermodynamically spontaneous. (e) 23. (c) 4.0. (a) 1. (e) 19. (b) the frequency of molecular collisions increases. 146 (c) 0.0096 mol/L (c) 0. what will be the new equilibrium concentration of A? A(g) + B(g) (a) 0.876 (b) 9.34 x 10-3 M 8.225 mol (b) 0.1 atm.080 mol/L (e) 0. Kc = 0.20 moles of A. 0.0055 mol/L (b) 0.85 x 10-4 M (d) 4.17 x 10-3 M (c) 1.26 (c) 0.40 mole of PCl3 in a 2. 0. B: 10.10 moles of A and 0.00 liter.060 mol/L of SO2 and 0.020 M (c) 0. C: 3.500 M How many moles of A must be added to increase the concentration of C to 0.11 M (d) 0.94 x 10-3 (b) 0. calculate the equilibrium concentration of NO in mol/L (a) 0.040 mol/L of N2 and 0.26 M 10. For the gas-phase reaction. Consider the following system in a 1.10 mol/L 9.07 M (b) 0. Consider the equilibrium system: 2ICl(s) I2(s) + Cl2(g) Which of the following changes will increase the total amount of of Cl2 that can be produced? (a) removing some of the I2(s) (b) adding more ICl(s) (c) removing the Cl2 as it is formed (d) decreasing the volume of the container (e) all of the above 14. After equilibrium is reached.50 gram sample of pure NOCl is heated at 350oC in a volume of 1.10 moles of B are added to this system. The reversible reaction: 2SO2(g) + O2(g) 2SO3(g) has come to equilibrium in a vessel of specific volume at a given temperature.60 atm. Calculate Kc for the reaction as written. Nitrosyl chloride.00 L container: A(g) + B(g) 2C(g) The equilibrium concentrations at 200oC were determined to be: [B] = 3.87 mol/L (d) 0.0532 (d) 54. Calculate the equilibirum concentration of H2 (and I2). Consider the reaction: N2(g) + O2(g) 2NO(g) Kc = 0. 2H2O(g) 2H2(g) + O2(g) Given that the forward reaction (the conversion of "left-hand" species to "right-hand" species) is endothermic. which of the following changes will decrease the equilibrium amount of H2O? (a) adding more oxygen (b) adding a solid phase calalyst (c) decreasing the volume of the container (the total pressure increases) (d) increasing the temperature at constant pressure (e) adding He gas The conventional equilibrium constant expression (Kc) for the system below is: 2ICl(s) (a) [I2][Cl2]/[ICl]2 (b) [I2][Cl2]/2[ICl] (c) [Cl2] (d) ([I2] + [Cl2])/2[ICl] (e) [Cl2]/[ICl]2 13.010 M (b) 0.700 mol 16.700 M at 200oC? (a) 0. Before the reaction began. (a) 7.040 mol/L.23 mol/L (e) 0. NOCl(g) (a) 0.75 x 10-4 (e) 0.00 M [C] = 0. If a reaction is initiated with 0.40 moles of C and 0.37 mol/L (b) 0.78 x 10-3 M (e) 2.417 mol (d) 0.040 for the system below at 450oC: C(g) + D(g) I2(s) + Cl2(g) (a) 9. When a 1.011 mol/L (d) 0. dissociates on heating as shown below.16 M (c) 0.050 mol/L of O2.040 mol/L of O2.0421 7. At equilibrium.55 x M (b) 1. A quantity of HI was sealed in a tube. Kc = 0. Consider the gas-phase equilibrium system represented by the equation: PCl3(g) + Cl2(g) [A] = 0. What is the equilibrium concentration of O2? (a) 0.20 moles of B.200 M . heated to 425oC and held at this temperature until equilibrium was reached. H2 + I2 10-3 2HI Kc = 54. (e) Qc is greater than Kc. The concentration of HI in the tube at equilibrium was found to be 0. If 0.2%. Consider the reversible reaction at equilibrium at 392oC: 2A(g) + B(g) C(g) The partial pressures are found to be: A: 6. what is the equilibrium concentration of Cl2 in the same system? PCl5(g) (a) 0. NOCl.47 mol/L (c) 0.0 liter container. more HI will be produced.040 M (e) none of these 11.(d) Qc is greater than Kc. Evaluate Kp for this reaction. the percent dissociation is found to be 57.15 mol/L 15.107 (d) 1.40 mole of D.04 M (e) 0.6 at 425oC NO(g) + 1/2 Cl2(g) 12.610 mol (e) 0. a 1. the concentrations of the reactants were 0.0706 mol/L. more H2 and I2 will be produced.040 for the system below at 450oC.70 atm.5 (e) 121 17.0 liter container was found to contain 0.305 mol (c) 0. the concentration of SO3 is 0.030 M (d) 0.40 mole of Cl2 and 0.10 at 2000oC Starting with initial concentrations of 0. 6. 0 x 10-10 M (e) 2.17 (e) 10.30. What is the concentration of hydroxide ions in this solution? (a) 4.2 (b) 4. The pOH of a solution of NaOH is 11. (d) It increases if the concentration of one of the reactants is increased.0 x 10-12 (d) 4. neutral (e) 8. The equilibrium constant at 427oC for the reaction: N2(g) + 3H2(g) 2NH3(g) is Kp = 9. (c) It changes with changes in the temperature. what is the pKa of this acid? (a) 5.0 x 10-7 M (c) pH = 7.53 (e) none of these 6.2 (d) 2.27 (d) 7. (c) 18. What is the approximate pH of a solution labeled 6 x 10-5 M HBr? (a) 4.5 (c) 5.87 (d) 1.PCl5(g) PCl3(g) + Cl2(g) Evaluate Kp for the reaction at 450oC. In a sample of pure water.0 x 10-7 M (b) [OH-] = 1.68 (c) 2.2 9.8 (e) 8. (b) 19.050 M solution of Ba(OH)2 is: (a) 1.0 x 10-13 M (d) 5. The [H3O+] in a 0. Which one of the following is a strong electrolyte? (a) H2O (b) KF (c) HF (d) HNO2 .4 x 10-5. (b) 5.9 (c) 3.2 x 10-9 M (b) 1.4 x 10-6 (e) 13. (c) 10.64 (c) 2.8 Answers to Chapter 17 1. A solution in which [H+] = 10-8 M has a pH of ___ and is ___.0 x 10-5 M (b) 5.72 (b) 7.0 x 10-8 M 11.0 (d) pOH = 7. (b) 16. (a) 8.052 (e) 6. (a) 10. (b) 4. (c) 11. What is the equilibrium constant for a reaction that has a value of Go = -41.9 x 10-15 at 10oC. (e) (CH3)3N 4. (b) It increases if the concentration of one of the products is increased. (e) It may be changed by the addition of a catalyst.6 x 10-12 M (d) 6.80.02 M solution of an unknown weak acid is 3. (c) 14.7 x 10-4 18. (b) 3.3 x 10-10 M (e) 2. (c) 9. Which one of the following is a weak acid? (a) HNO3 (b) HI (c) HBr (d) HF (e) HClO3 2. the equilibrium constant? (a) It always remains the same at different reaction conditions. Which salt is not derived from a strong acid and a strong soluble base? (a) MgCl2 (b) Ba(NO3)2 (c) LiClO4 (d) CsBr (e) NaI 3. (a) 0.40 (b) 0.01 (b) 7. (d) 12. (c) 13.83 The pH of a solution is 4. (d) 15. If Kw is 2. what is the pH of pure water at 10oC? (a) 6. basic (c) -6. For a specific reaction.2 x 10-8 7.7 (b) 4.5 x 10-3 (c) 5. What is the pH of 500 mL of solution containing 0.4 (d) 0. (c) 20.0 x 10-2 M (c) 1.8 kJ at 100oC? (a) 1. The pH of a 0.0124 grams of Ca(OH)2? (a) 11.7.5 19.0 x 10-12 (e) 6. (a) -33 kJ (b) -54 kJ (c) 54 kJ (d) 33 kJ (e) 1. (a) 17. acidic (b) 6.3 J 20. What is the [H+] for this solution? (a) 2.0 (e) [H3O+] = [OH-] 5. which of the following statements can be made about K. basic (d) -8.00 (c) 7.04 (b) 9. (e) 2. (c) 7. (a) 8. only one of the following statements is always true at all conditions of temperature and pressure. Calculate the value of Go for the reaction at 427o.96 (d) 3.0 x 10-3 (b) 2. 1.0 x 10-5 M 8.1 x 105 (c) -5.6 x 10-5 M (c) 3. (b) 6. basic 12. Which one is always true? (a) [H3O+] = 1.8 (d) 9. (d) 22.36 (c) 0.19? (a) 0.26 (d) 9.5 x 10-7 (c) 7. Calculate the pH of a 0. (CH3)3NHCl.00? (a) 0. (c) 25. (e) 17. (d) 24.44 (d) 0. Which of the following solutions has the lowest pH at 25oC? (No calculations required.34.059% ionized.0068 % (c) 0.2 M benzoic acid (e) pure water 16.8 x 10-9 (b) 6. What is the concentration of a sodium acetate solution if the pH of the solution is 9.2 M hypochlorous acid (c) 0.10 M in CH3COOH and 0.2 (d) 11.4 (d) 2.06 (b) 5.20 M NH3 and 0.52 (e) 7.10 M solution of a weak acid.00 22. (a) 21.0 x 10-6 (d) 4.0 x 10-7 (c) 4. Evaluate Ka for the acid. (c) 20.80? (a) 0.50 M solution of NaNO2.022 % Answers to Chapter 18 1. (a) 12.0094 % (d) 0.1 (e) 7.2 15. (c) 6. .28 (b) 0.15 (b) 4.1 M (d) 0.10 M solution of a weak acid.5 x 10-5 (b) 1. is 0.56 (e) 0. (a) 3.38 (e) 8.15 M 25.60 M (e) 0. (d) 16.7 (e) 12.24 (e) 5. (a) 9.43 M (c) 2.2 M HF solution? (a) 2.48 (c) 1.9 (c) 4. What is the percent hydrolysis? (a) 0.2 M sodium hydroxide (b) 0. (a) 18.060 M NH4Cl? (a) 5. What is the concentration of ammonium chloride in a solution if its pH is 4.8 (c) 9. (e) 19.011 % (e) 0. (a) 2.6 14.0031 % (b) 0.25 M (b) 0.2 x 10-6 (e) 3.4 % (b) 4. (c) 14. (a) 13.18 (b) 5.84 % (d) 0.050 M HClO? (a) 5.82 (d) 8.1 (b) 3. (b) 23. (d) 15. 4.18 (d) 5. (e) 12.15 M trimethylammonium chloride. What is the pH of a solution labeled 0.(e) 3. The pH of 0.068 M (e) 0. Which of the following statements is true? 20. (b) 4. (e) 5.30 M (CH3)3N? (a) 9.2 M ammonia (d) 0.22 % 18.0 x 10-10 (d) 5.15 M NH4Cl? (a) 2. What is the approximate pH of a solution labeled 0.082 % (e) 0. is 5. (a) 3.5 (b) 10. Which of the following is true about a 0. HX. (d) 11.) (a) 0. Calculate the ratio [CH3COOH]/[NaCH3COO] that gives a solution with pH = 5.62 (c) 8.12 (c) 5.2 % (c) 0.20 M in NaCH3COO. (a) Which of the following combinations cannot produce a buffer solution? (a) HNO2 and NaNO2 (b) HCN and NaCN (c) HClO4 and NaClO4 (d) NH3 and (NH4)2SO4 (e) NH3 and NH4Br 2. What is the percent ionization of an 1.63 Consider a solution which is 0. What is the pH of a solution composed of 0.45 M (d) 0.10 M (e) both b and d Calculate the hydrolysis constant for the cyanide ion.5 x 10-8 17.30 M (c) 0. (e) 10. (a) 7. 23.89 3. CN-. Which of the following weak acids ionizes to give the strongest conjugate base? (a) HClO (b) CH3COOH (c) HF (d) HNO2 (e) HCN 19.30 M (b) 0. 1. HX? (a) [X-] = 0.35 24. (d) 2.6 x 10-10 (e) none of these 21.59 M What is the pH of 0. (c) 8.10 M (b) pH = 1 (c) [HX] > [H+] (d) [H+] = 0.7 13. A 0. a salt. pink blue yellow colorless none of these 1.100 M HNO3.10 M HF and 200 mL of 0. (2) In the middle of the pH range of its color change a solution containing the indicator will probably be orange. 4 (c) 3.00 M NaOH.84 g (d) 6.7.2 .(a) If a small amount of NaOH is added.44 (e) 3. the H+ ions react with CH3COOH ions.58 (d) 3. 7. The following titration curve is the kind of curve expected for the titration of a ____ acid with a ____ base.63 (d) 10.9.05 M methylamine and 0. Which of the following salts give acidic aqueous solutions? (1) KNO3 (2) KCH3COO (5) (NH4)2SO4 (a) 2. Consider an indicator that ionized as shown below for which its Ka = 1. (a) 1.32 (d) 3.0 mL of 1.00.00 L of 0.8 8.6 6.6 (b) 5. (3) NH4N (7) NaCN .0 mL of 0.5 . 3. How many grams of NaF would have to be added to 2.) (4) At pH = 7. 6 16.96 (b) 9.70 (c) 1.57 6.83 (c) 9. 5 (b) 2.00. 3.0 mL of 0.00 liter. the pH decreases very slightly.00? (a) 300 g (b) 36 g (c) 0.20 8. 5 (c) 2.11 (e) 9. (Hint: Write the equilibrium constant expression for the indicator.0 mL of 0. 8 (e) 1. (3) At pH = 7.39 (e) 4.9 g (e) 60. For which pair is the pH at the equivalence point stated incorrectly? Acid-Base Pair pH at Equiv (a) HCl + NH3 (b) HNO3 + Ca(OH)2 (c) HClO4 + NaOH (d) HClO + NaOH (e) CH3COOH + KOH 12. (c) If a small amount of HCl is added.8 3.20 M ammonia with 0. the pH decreases very slightly. What is the pH of this solution? (a) 8. (6) BaCl2 less than 7 equal to 7 equal to 7 less than 7 greater than 7.95 (d) 2.5 (e) 9.52 (e) 2.96 (c) 3. (b) If NaOH is added. 5. 5 (d) 1.100 M HF to yield a solution with a pH = 4.31 (d) 9. the pH increases. (5) The pH at which the indicator changes color is pH = 4. 4.2 (c) 7. (a) 2. Calculate the pH of the solution resulting from the addition of 20.80 10.00 mole of ammonia and 1. most of the indicator is in the unionized form. a solution containing this indicator (and no other colored species) will be red.12 (c) 1.35 (b) 1.10 M KOH. g Calculate the pH that results when the following solutions are mixed.8 15.3 . To 500 mL of this solution was added 30. 6 (d) 1. 4. 8 (b) 3.20 M hydrochloric acid? (a) 2. the OH.20 M formic acid (2) 55 mL of 0. (d) If HCl is added.ions. 4. What is the pH at the equivalence point in the titration of 100. (e) If more CH3COOH is added.57 (b) 1.11 (c) 4. Which indicator (identified by a letter) could be used to titrate aqueous NH3 with HCl solution? Indicator Acid Range Color Color-Change pH (a) (b) (c) (d) (e) 11.100 M NaOH to 30. 2.00 mole of ammonium chloride to form an aqueous solution with a total volume of 1. 5 (e) another combination 9.5 (e) 4.64 (b) 3.82 (b) 2. A buffer was prepared by mixing 1.4 .10 M sodium formate (3) 110 mL of water (a) 3. What is the approximate pH of a solution prepared by mixing equal volumes of 0.10 M hydrochloric acid? (a) 4. 7.9 Consider the titrations of the pairs of aqueous acids and bases listed on the left.0 (d) 5. Calculate the pH of a solution prepared by mixing 300 mL of 0.ions react with the CH3COO. (a) 1.0 x 10-4 HIn + H2O H3O+ + Inyellow red Which of the responses contains all the true statements and no others? (1) The predominant color in its acid range is yellow.9 13. (1) 35 mL of 0.4. 4.2.53 14. 1 M AgNO3 (b) 0. If PbSO4 were used in an unglazed ceramic bowl. When we mix together. equals (d) equals. is greater than 12. What is the molar solubility. (d) 12. (a) 17. .7 x 10-11 M (d) 5. is less than (e) equals.0 M immediately after mixing. (b) 16.7 x 10-10 M 8. is (a) [Sn2+][OH-] (b) [Sn2+]2[OH-] (c) [Sn2+][OH-]2 (d) [Sn2+]3[OH-] (e) [Sn2+][OH-]3 2. and will continue to precipitate until Qsp _____ Ksp. Ksp Ag2CrO4 BaCrO4 PbCrO4 9.6 x 10-3 M (e) 8. (a) 6. The pH at the equivalence point is _____. Ag3PO4 would be least soluble at 25oC in (a) 0. Ksp = 1.20 M nitrous acid by adding 0. What is the molar solubility of Cu(OH)2? (a) 3. from separate sources.1 M HNO3 (c) pure water (d) 0.7 x 10-5 M (d) 4. (c) 6. using x to represent the molar concentration of silver(I) and y to represent the molar concentration of sulfide.0 x 10-12 2. (b) 10. (e) 7.6 x 10-4 6.5 x 10-5 M 11. is formulated as: (a) xy (b) x2y (c) xy2 (d) x2y2 (e) xy3 3. The solubility product expression for silver(I) sulfide. (b). weak (d) weak. the ions of a slightly soluble ionic salt. The solubility of silver sulfate in water at 100oC is approximately 1.4 x 10-7 M (b) 6. s. The molar solubility of PbCl2 in 0.) (a) greater than 7 (b) equal to 7 (c) less than 7 (d) cannot be determined without more data (not including Ka and Kb) (e) is impossible to predict 7. (b) 13. 1.20 M Pb(NO3)2 solution is: (a) 1. strong (b) weak.1 x 10-10 M (e) 1. (d) 4. What is the solubility product of this salt at 100oC? (a) 5. Consider the following solubility data for various chromates at 25oC.17 x 10-3 M at a certain temperature. (d) 3.4 x 10-7 M (c) 2. equals (b) is less than. (a) is greater than. Calculate Ksp for PbBr2.1 M Na3PO4 (e) solubility in (a).8 x 10-14 The chromate that is the most soluble in water at 25oC on a molar basis is: (a) Ag2CrO4 (b) BaCrO4 (c) PbCrO4 (d) impossible to determine (e) none of these The molar solubility of PbBr2 is 2.(b) 6.0 mL of 0.4 g per 100 mL. weak (e) none of these 17. or (d) is not different 10. is greater than (c) is less than. (c) 5. (Note: This is the titration of a weak acid with a weak base.1 x 10-8 (d) 3. (c) Many lead salts are often used as pigments. strong (c) strong.7 x 10-8 (b) 3.4 x 10-6 (e) 1. (c).5 x 10-7 (c) 8.2 x 10-3 M (c) 1. of Ba3(PO4)2 in terms of Ksp? (a) s = Ksp1/2 (b) s = Ksp1/5 (c) s = [Ksp/27]1/5 (d) s = [Ksp/108]1/5 (e) s = [Ksp/4]5 For Cu(OH)2.1 x 10-5 (e) 3. (a) 8. (d) 14. (d) 9.2 x 10-6 4.0500 M aqueous ammonia to it. (a) CuBr2 and K2CO3 (b) HNO3 and NH4I (c) BaCl2 and KClO4 (a) strong. Sn(OH)2. how many milligrams of lead(II) could dissolve per liter of water? (a) 43 (b) 35 (c) 11 (d) 28 (e) 53 9. (b) 15. (c) 2.0 x 10-10 1. The solubility product expression for tin(II) hydroxide. (b) 11.4 x 10-7 (c) 4.6 x 10-19. Consider the titration of 30. the salt will precipitate if Qsp _____ Ksp.3 x 10-6 (d) 4.7 x 10-4 M (b) 9. Answers to Chapter 19 1. Which of the following pairs of compounds gives a precipitate when aqueous solutions of them are mixed? Assume that the concentrations of all compounds are 1.4 x 10-5 5. oxidation (d) cathode.0 x 10-4 M (e) 1. (e) All voltaic (galvanic) cells involve the use of electricity to initiate nonspontaneous chemical reactions.0 (b) 8. (2) the (4) Cl2 (6) 2 H (7) electrons flow from the electrode to the external circuit (9) oxidation (a) 2.0 (d) 9. (e) 1. is produced at the other. (d) 9. the .8 x 10-7 M (d) 3. 6.5 x 10-7 M (c) 1. (c) 16. 6. (d) Oxidation occurs at the anode.0 x 10-13.98 L (e) 1. What is the concentration of Au+ ions at this point? Ksp for AgCl = 1.0010 M in both Ag+ and Au+. A swimming pool was sufficiently alkaline so that CO2 absorbed from the air produced in the pool a solution which was 2 x 10-4 M in CO32. 8.5 g (c) 0.M. (b) All electrochemical reactions involve the transfer of electrons.0 g 6. 2. 5. 5. CaCO3 and FeCO3 14.2 17. Which of the following responses describe or are applicable to the cathode and the reaction that occurs at the cathode? (1) the positive electrode (3) 2 Cl(5) 2 H2O Cl2 + 2 eO2 + 4 H+ + 4 e- 15. 9 (b) 1.0 hours with a current of 0. (a) 12. (a) 8.84 L (d) 0.0 (c) 6.25 amperes flowing for 10 hours? (a) 12 g (b) 5. reduction (e) cannot tell unless we know the species being oxidized and reduced.9 (d) 10.(d) Na2CO3 and H2SO4 (e) KCl and KNO3 13.8 x 10-10 and for AuCl = 2. (d) 11.40 amperes.5 (b) 10. (c) 2.001 M Pb(NO3)2 and 0. (a) 10. What is the pH of a saturated solution of Mg(OH)2? (a) 3. Which solid will precipitate first if an aqueous solution of Na2CrO4 at 25oC is slowly added to an aqueous solution containing 0. (a) 13.1 (c) 10. 7. A solution is 0. 8. 10 (e) 2. In an electrolytic cell the electrode at which the electrons enter the solution is called the ______ .02 L 7. gaseous hydrogen is evolved at one electrode and gaseous chlorine at the other electrode.0 x 10-10 M (b) 4. Cl2. (e) 6.+ H2 During the electrolysis of aqueous KCl solution using inert electrodes. Molten AlCl3 is electrolyzed for 5.0015 M? (a) 9. Which of the following statements is FALSE? (a) Oxidation and reduction half-reactions occur at electrodes in electrochemical cells. (d) 17. 6 x 10-4 M in Ca2+ and 8 x 10-7 M in Fe2+. (d) 14. 10 5.1 x 10-6 M chemical change that occurs at this electrode is called _______. 8.6 16. How many liters of Cl2 measured at STP are produced when the electrode efficiency is only 65%? (a) 0. (a) 2. (d) 7. precipitation should occur only for: (a) 2 x 10-3 M Mg2+ + 2 x 10-3 M OH(b) 2 x 10-1 M Ba2+ + 2 x 10-3 M F(c) 2 x 10-3 M Ca2+ + 2 x 10-2 M OH(d) 2 x 10-3 M Ca2+ + 2 x 10-3 M OH(e) 2 x 10-4 M Pb2+ + 2 x 10-5 M SO42At what pH will Cu(OH)2 start to precipitate from a solution with [Cu2+] = 0. When equal volumes of the solutions indicated are mixed. The half-reaction that occurs at the anode during the electrolysis of molten sodium bromide is: (a) 2 BrBr2 + 2 e(b) Br2 + 2 e2 Br(c) Na+ + eNa (d) Na Na+ + e(e) 2 H2O + 2 e2 OH.5 (e) 9. 9 (c) 2. (b) 3. 9 (d) 1. 7. 4. oxidation (b) anode.7 g (e) 6.4 (e) 4.100 M Ba(NO3)2 at 25oC? (a) BaCrO4(s) (b) NaNO3(s) (c) PbCrO4(s) (d) Pb(NO3)2(s) (e) none of these 18. (a) anode. What mass (in grams) of nickel could be electroplated from a solution of nickel(II) chloride by a current of 0. The solution around the electrode at which hydrogen gas is evolved becomes basic as the electrolysis proceeds.046 g (d) 2. (8) ele circuit (10) re Answers to Chapter 20 1. (a) 15. 3. If the pool water was originally 4 x 10-3 M in Mg2+.63 L (c) 0. (a) 4. (c) 5. Metallic aluminum is produced at one electrode and chlorine gas. then a precipitate should form of: (a) only MgCO3 (b) only CaCO3 (c) only FeCO3 (d) only CaCO3 and FeCO3 (e) MgCO3. (c) Reduction occurs at the cathode. (c) 18. Some solid NaCl is added slowly until the solid AgCl just begins to precipitate. 6. reduction (c) cathode.55 L (b) 0. 66 V (e) -1. Estimate the equilibrium constant for the system indicated at 25oC. using an average current of 10 amperes at an 80% electrode efficiency? (a) 8. What is the reduction potential for the half-reaction at 25o C: Al3+ + 3eAl.10 M and Eo = -1.Au3+.00 (b) 1.How long (in hours) must a current of 5.3 hours (c) 11 hours (d) 16 hours (e) 5.0020 M. As the cell given below operates.0 M) / Ag (a) Pb2+ + 2 ePb (b) Pb Pb2+ + 2 e(c) Ag+ + eAg + (d) Ag Ag + e(e) none of the above For a voltaic (or galvanic) cell using Ag.00 (d) 0. How many faradays are required to reduce 1.06 V (e) Fe.72 V What is the value of E for the half-cell: MnO4. calculate the potential of the cell.86 V (c) 1. Cr2O72-. (e) The concentration of Ag+ will decrease as the cell operates. What is Go per mole of dichromate ions for the reduction of dichromate ions. 20.3 kJ (e) -53. to Cr3+ by bromide ions.0 amperes be maintained to electroplate 60 g of calcium from molten CaCl2? (a) 27 hours (b) 8.020 12. if [Al3+] = 0. Consider the standard voltaic (or galvanic) cell: Fe.44 V (d) 1. Which of the following statements is(are) true for all voltaic (or galvanic) cells? (I) Reduction occurs at the cathode. How long. 15. 1.0 M) || AgNO3 (1.84 V (b) -1.94 V 14. (b) Electrons will flow through the external circuit from the zinc electrode to the silver electrode. 17. ~1069 3Mg + 2Al3+ Mn2+ (0. such as those diagrammed in your text.2 (e) 12.4 9.250 10. the strip of silver gains mass (only silver) and the concentration of silver ions in the solution around the silver strip decreases. (d) The mass of the zinc electrode will decrease as the cell operates. Ga / Ga3+ (10-6 M) || Ag+ (10-4 M) / Ag (a) 1.4 (c) 16. which of the following statements is incorrect? (a) The zinc electrode is the anode.52 V 16.50 V (b) 1.0 M) and Zn.0 M) half-cells.) (a) +26.080 V 18.010 M) + 8H+ (0.8 (d) 11. Calculate the potential (in volts) for the voltaic (or galvanic) cell indicated at 25oC.58 V (e) 1. Which of the following is the strongest oxidizing agent? (a) Pb2+ (b) I2 (c) Ag+ (d) Pb (e) Cu2+ 11. 3 Mg2+ + 2Al (a) (b) ~1023 (c) ~10-24 (d) ~10-36 (e) ~10-72 In voltaic cells. 13.020 V (b) 1.29 V (b) 0. Br-.21 V (e) 1.4 (b) 5.9 hours 8. in acidic solution? (Hint: Use the standard cell potential.030 V (d) 1. in hours. (c) Reduction occurs at the zinc electrode as the cell operates.68 V (d) -1. 1.0 V (e) 0. the salt bridge _______ .97 V (c) 1. (a) 0.Zn2+ (1.50 (c) 3.44 V (b) Au.3 kJ (b) -145 kJ (c) +145 kJ (d) -26. -0.2 V (c) 0.37 V A concentration cell is constructed by placing identical Cu electrodes in two Cu2+ solutions. If the concentrations of the two Cu2+ solutions are 1. Which answer identifies the cathode and gives the Eo for the cell? (a) Fe. .66 V ? (a) -1. while the strip of lead loses mass and the concentration of lead increases in the solution around the lead strip.45 V (d) 1.60 V (c) -1. (a) is not necessary in order for the cell to work (b) acts as a mechanism to allow mechanical mixing of the solutions (c) allows charge balance to be maintained in the cell (d) is tightly plugged with firm agar gel through which ions cannot pass (e) drives free electrons from one half-cell to the other 21.06 V (d) Au. 1.Ag+ (1.(0.6 kJ 19. Which of the following represents the reaction that occurs at the negative electrode in the above cell? Pb / Pb(NO3)2 (1.20 M) + 5eM) + 4H2O ? (a) 1.Fe2+ versus Au.00 g of aluminum(III) to the aluminum metal? (a) 1.111 (e) 0. would be required for the electroplating of 78 g of platinum from a solution of [PtCl6]2-.0 M and 0. 1.94 V (c) Fe. ) (III) The voltage is less than or equal to zero.' a(n) _____ is chemically converted to a(n) _____. Some metals are found in the uncombined free state while other metals are found in the combined state. (c) 2.3 x 108 coulombs Answers to Chapter 21 1. (b) The active metals can occur in the free state while the less active metals occur in the combined state. (e) none of the above 2. Which of the following is NOT true for the Group 1A elements? (a) Most of them are soft.8 pounds of magnetite. (e) They exhibit a +1 oxidation state in compounds. oceans (d) oceans. oxide (c) hydroxide.7% (e) 23. oxide (d) oxide. (c) Metals with positive reduction potentials can occur in the free state while metals with negative reduction potentials can occur in the combined state.0 x 109 coulombs (c) 7. (c) 13. (d) There is no way we can predict which metals will be free or combined. sulfate (e) phosphate. oceans 3. (c) They are named the alkaline earth metals. salt beds (e) rivers. MgCl2 Mg(l) + Cl2 (a) 4.13 weight percent? Assume 100% recovery. (c) 21. (d) 17. (b) 12.9 x 109 coulombs (d) 1.(II) The anode gains mass during discharge (note: this means operation of the cell. (e) 3. Fe3O4.35 grams (d) 0. (e) 18.8% (d) 16. A reaction sequence for the reduction of one of the iron ores is as follows: 2 C(coke) + O2 2 CO Fe2O3 + 3 CO 2 Fe + 3 CO2 Calculate the amount of coke necessary to produce 800 g of Fe. (e) 20.500 coulombs in one faraday. What is the weight percent iron in this ore? (a) 72. (d) They are excellent conductors of heat and electricity. (b) 14.3 grams (c) 0. In the process known as 'roasting. (c) 22. phosphide 4. (c) 6. 2. and III 22. (b) 9. (a) 4. What is a deciding factor? (a) Metals with negative reduction potentials can occur in the free state while metals with positive reduction potentials occur in the combined state.3% (b) 20. (d) 6. (a) 1. (d) 8. (d) 1.1% (c) 27. (a) 3. (d) 7. (a) sulfide. In the standard notation for a voltaic cell. (c) 11. (c) 15. (d) 2. (c) 16. (b) Their atomic radii increases with increasing molecular weight. 10. (b) 8. whereas insoluble metal compounds tend to be found in the _____ . (a) oceans. Soluble metal compounds tend to be found in the _____. Answers to Chapter 22 1.1% 8.0 x 109 coulombs (b) 2. (a) 10. the double vertical line "||" represents: (a) a phase boundary (b) gas electrode (c) a wire (metal) connection (d) a salt bridge (e) a standard hydrogen electrode (e) Cu 6. (c) 5. Cu3(CO3)2(OH)2? (a) 2+ (b) 1+ (c) 0 (d) 1(e) 2How many coulombs of electricity are required to produce one metric ton (1000 kg) of magnesium? There are 96. II. (b) 19.5 grams 7. (a) 4. (c) 9. . (d) 10. What is the charge on the copper ion in the mineral azurite. earth's crust (b) earth's crust. (e) 5. Which metal can be found as the free element? (a) Na (b) Mn (c) Fe (d) Cr (e) Pt The Hall-Heroult process is used in the production of: (a) Mg (b) Fe (c) Al (d) Au 1. oceans (c) salt beds. (e) 5.2 x 1012 coulombs (e) 5. oxide (b) carbonate. (a) 7.7 grams (b) 1.59 grams (e) 2. (a) only III (b) only II (c) only I (d) II and III (e) I. silvery corrosive metals. (a) 114 g (b) 1030 g (c) 258 g (d) 172 g (e) 544 g 9. How much magnesium can be obtained from one pound of seawater if the concentration of Mg2+ is 0. For every 100 pounds of iron ore there are 27. (c) HI is the largest.84% (b) 9. (e) electrolysis of AlCl3(aq). 1. 3. Some element groups of the periodic table are more likely to contain elements that are gases than other groups.3.8% (e) 20. 5.10% (c) 12.0oC and 742 torr. (c) 2. (c) electrolysis of NaCl(aq). (b) 6. (d) HCl has the lowest boiling point. Chlorine gas is prepared commercially by: (a) electrolysis of carbon tetrachloride.3% (d) 15. (b) HF has the lowest of the H-X bond energies. (b) oxidation of chloride ion with F2(g). ___ double bonds and ___ lone pairs of electrons. Which of the following is NOT true for the halogens? (a) They are nonmetals. (d) Their compounds with metals are generally ionic in nature. (b) 5. What is the electron configuration of Mn3+ ion? (a) [Ar] 4s2 3d10 (b) [Ar] 4s2 3d2 (c) [Ar] 3d5 (d) [Ar] 3d4 (e) [Ar] 4d1 3d3 11. (b) They show the -1 oxidation number in most of their compounds. 6. HX? (a) Their aqueous solutions are acidic. (e) Elemental halogens exist as diatomic molecules. Of the oxyacids listed below.6% 7. (e) HF exhibits hydrogen bonding. . 9. (d) 11. (d) 9. (d) BeO. Of the following oxides. (c) The electronic configuration of their outermost electrons is ns2 np6. the most basic is: (a) MgO. (c) 7. Which of the following properties of the alkaline earth metals decreases with increasing atomic weight? (a) ionic radii (b) ionization energy (c) atomic radii (d) activity (e) atomic number 5. Draw the correct Lewis formula for chlorous acid. which one possesses the greatest acid strength in water? (a) HClO4 (b) H2CO3 (c) H3BO3 (d) HClO (e) HBrO 7. Such a relationship is called: (a) amphoterism (b) an allotropic relationship (c) a diagonal relationship (d) the periodic law (e) an isoelectronic series The nitrate of which of the following cations would exhibit paramagnetism to the GREATEST extent? (a) Co3+ (b) Cr3+ (c) Fe3+ (d) Mn3+ (e) V3+ Answers to Chapter 23 1.0 L of CO2 was collected at 50. (e) 8. A 300 g sample of CaCO3 was heated until 10. (b) 10. Which one of the following does not correctly describe one or all of the hydrogen halides. (c) 4. The structure contains ___ single bonds. (c) 4. and boron resembles silicon. Which element group is the most reactive of all the metallic elements? (a) alkali metals (b) alkaline earth metals (c) coinage metals (d) transition metals (e) Group 2B metals In a surprisingly large number of their properties beryllium resembles aluminum. (e) SO2. What percentage of the CaCO3 had decomposed? (a) 6. (c) P2O3. (d) oxidation of chloride ion with Br2(aq). (b) Na2O. Which of the following groups contains the greatest number of gaseous elements? (a) IA (b) IIA (c) IVA (d) VIA (e) VIII (or 0) 2. What mass of lithium nitride could be formed from 104 g of lithium and excess nitrogen gas? (a) 35 g (b) 60 g (c) 105 g (d) 140 g (e) 174 g 8. (a) 3. 6. The most abundant metal in the earth's crust is: (a) Cu (b) Fe (c) Na (d) Al (e) Ca Which element has the electron configuration [Ar] 3d7 4s2? (a) Fe (b) Co (c) Cr (d) Ti (e) Zn 10. Which of the following substances is the strongest reducing agent? (a) Cl2 (b) Cl(c) Br2 (d) Br(e) I2 4. Which acid listed on the right cannot be obtained by adding water to the substance on the left? (a) H2S2O7 . +1 15. Which of the following statements about sulfuric acid is false? (a) It is a strong acid. (e) 12. the correct method is to add sulfuric acid to water. In coordination chemistry. (a) 7. (e) All are covalent compounds. Which of the following has a pyramidal structure (molecular geometry)? (a) CBr4 (b) PF3 (c) BF3 (d) OF2 (e) BrCl 9. (d) the atom in the ligand that shares an electron pair with the metal. The Lewis acid is (a) [Pt(CN)4]2(b) Na+ (c) Pt (d) Pt2+ (e) CN4. (c) Most are found in sulfide deposits. (c) 17. In which one of the following is the oxidation state of nitrogen given incorrectly? (a) N2O3. The ______ sphere is enclosed in brackets in formulas for complex species.tellurous acid 12. Consider the coordination compound. (b) 14. H2Se and H2Te are all gases at room temperature and atmospheric pressure. 5 (b) 3. Which statement about the Group VIA hydrides is false? (a) H2S. 1. (b) 6. (c) The sulfur atom is sp2 hybridized. +3 (b) N2H4. 10. 5 (e) none of these 8. 7 (c) 1. (b) All are colorless. the donor atom of a ligand is (a) a Lewis acid. 13. and it includes the central metal ion plus the coordinated groups. (e) 2. (a) ligand (b) donor (c) oxidation (d) coordination (e) chelating 2. 1. . (c) All except H2O are toxic.sulfurous acid (e) TeO2 . (b) 15.0% FeS2? (a) 84. Which statement about the Group VIA elements is false? (a) All have an outer electronic configuration of ns2 np4. LIGAND NAME (a) OH(b) CNhydroxo cyanide 5. 2. (b) One mole of sulfuric acid reacts completely with two moles of potassium hydroxide. Consider the coordination compound. (d) H2Po has the lowest boiling point.selenous acid (c) SO3 .4 kg (b) 123 kg (c) 136 kg (d) 144 kg (e) 168 kg 14. +5 (d) NaNO2. 4 (d) 2. 0. (b) 8. 11. (d) Oxygen has the highest boiling point and melting point. (d) It is often present in acid rain. A coordinate covalent bond exists between (a) K+ and CN(b) Cu2+ and CN(c) K+ and [Cu(CN)4]2(d) C and N in CN(e) K+ and Cu2+ Given the list of ligands and their corresponding names. (d) 4. (d) 11. (c) 13. What is the major mineral present in phosphate rock? (a) Ca3(PO4)2 (b) Na2HPO4 (c) Ca10(PO4)6F2 (d) NaH2PO4 (e) Ca10(PO4)6(OH)2 Answers to Chapter 24 1. (e) During the dilution of sulfuric acid. choose the pair that disagree. K2[Cu(CN)4]. (c) 5. Which compound gives photochemical smog a brownish color? (a) NO (b) HNO2 (c) NO2 (d) N2O4 (e) N2O3 17.sulfuric acid (d) SO2 . (e) the atom in the ligand that accepts a share in an electron pair from the metal. What maximum mass of sulfuric acid can be produced from the sulfur contained in 100 kilograms of iron pyrite that is 75. (c) 3. (a) 1. (e) 16. +3 (e) H2N2O2. (b) the counter ion (c) the central metal atom. (b) 9. (b) The electronegativity of Group VIA elements decreases as one goes down the group. (e) Polonium has the smallest first ionization energy. (d) 10.sulfuric acid (b) SeO2 . Na2[Pt(CN)4]. 3.(a) 2. Which of the following does not correctly describe ammonia? (a) pyramidal molecule (b) polar molecule (c) extremely soluble in water (d) forms basic aqueous solutions (e) none of these 16. +2 (c) HNO3. 10. (a) 1. the d electrons on a metal ion occupy the eg set of orbitals before they Which name-formula combination is NOT correct? occupy the t2g set of orbitals. II (b) III. (4) The complex is diamagnetic. V A molecule that cannot be superimposed on its mirror (c) I. (c) linkage isomerism (Crystal Field Theory) Consider the violet-colored (d) reactive isomerism compound. (Valance Bond Theory) Magnetic measurements (Crystal Field Theory) How many unpaired electrons indicate that [Co(OH2)6]2+ has 3 unpaired electrons. 16. Select the correct IUPAC name for: [FeF4(OH2)2](a) diaquatetrafluoroiron(III) ion (b) diaquatetrafluoroferrate(III) ion (c) diaquatetrafluoroiron(I) ion (d) diaquatetrafluoroferrate(I) ion (e) none of these 7. [Cu(OH2)6]2+ has one unpaired electron. It is diamagnetic.(c) Cl(d) H2O (e) NH3 6. 13. 12. (b) [Zn(NH3)2Cl2] (tetrahedral) II. The metal ion is a d5 ion. and no false statements? I. It is octahedral. (Valence Bond Theory) The coordination complex. The ligands are weak field ligands. (b) Diamagnetic metal ions cannot have an odd FORMULA NAME number of electrons. V (a) geometrical isomerism (e) III. 2. (3) The complex is d2sp3 hybridized. (b) K[Cr(NH3)2Cl4] potassium diamminetetrachlorochromate(III) (d) In high spin octahedral complexes. chloro aqua ammine 14. (e) coordination isomerism [Cr(NH3)6]Cl3. 5 15. (d) [Co(NH3)5Cl]2+ (octahedral) IV. (a) I. and is relatively very small. Which of the following statements are true? (1) The complex is octahedral. oct is less (c) [Mn(CN)5]2pentacyanomanganate(II) ion than the electron pairing energy. the hybridization of the metal's orbitals in complex? [Co(OH2)6]2+ is: (a) 0 3 (a) sp (b) 1 (b) sp2d (c) 2 (c) dsp2 (d) 4 3 2 (d) sp d (e) 6 (e) d2sp3 18. [Cr(OH2)6]Cl3 and the yellow compound. (2) The complex is an outer orbital complex. 3. IV (b) optical isomerism 19. (Crystal Field Theory) Consider the complex ion Which one of the following complexes can exhibit [Mn(OH2)6]2+ with 5 unpaired electrons. IV. Which geometrical isomerism? response includes all the following statements that (a) [Pt(NH3)2Cl2] (square planar) are true. (Crystal Field Theory) Which one of the following statements is FALSE? (a) In an octahedral crystal field. (5) The coordination number is 6. dxz and dyz (c) +1 (c) dxz and dyz (d) +2 (d) dxz. which orbitals are raised least [Pt(NH3)3Cl]Cl? in energy? (a) -1 (a) dxy and dx2-y2 (b) 0 (b) dxy. 11. Which of the following statements is false? In which one of the following species does the transition metal ion have d3 electronic configuration? (a) [Cr(NH3)6]3+ (b) [Co(OH2)6]2+ (c) [CoF6]3(d) [Fe(CN)6]3(e) [Ni(OH2)6]2+ . are there in a strong field iron(II) octahedral Therefore. (a) [Co(NH3)4(OH2)I]SO4 tetraammineaquaiodocobalt(III) sulfate (c) Low spin complexes can be paramagnetic. 3 (e) 4. IV image is said to exhibit which of the following? (d) II. (d) [Ni(CO)4] tetracarbonylnickel(0) (e) Ca[PtCl4] calcium tetrachloroplatinate(II) (e) Low spin complexes contain strong field ligands. It is a low spin complex. 5 (d) 2. 5 (c) 2. 9. (c) [Cu(NH3)4]2+ (square planar) III. dyz and dz2 (e) +3 (e) dx2-y2 and dz2 17. 4 (b) 1. (e) [Cu(CN)2](linear) V. Select the correct IUPAC name for: [Co(NH3)6]2+ (a) hexammoniacobaltate(II) ion (b) hexaamminecobaltate(II) ion (c) hexammoniacobalt(II) ion (d) hexaamminecobalt(II) ion (e) hexammoniacobalt ion 8. (Crystal Field Theory) When the valence d orbitals of What is the oxidation number of the central metal the central metal ion are split in energy in an atom in the coordination compound octahedral ligand field. (c) 1. (c) n/p ratios that confer nuclear stability. (c) oct for [Cr(OH2)6]3+ is less than oct for [Cr(NH3)6]3+. (a) 0. (a) neutron emission (b) beta emission Answers to Chapter 25 1. One would predict that it decays via _____. (c) Nuclei with highest binding energies are the most stable nuclei. The mass defect for an isotope was found to be 0. (e) The two complexes absorb their complementary colors. Particle masses in amu are: proton = 1. Calculate the binding energy per nucleon (in units of MeV) for 9Be.69 x 1013 kJ/mol (d) 1.341 amu 9. Calculate the binding energy in kJ/mol of atoms. (d) 11. The actual mass of a 37Cl atom is 36. (b) 6. proton (b) 1. (b) 17. 2+. (c) usually produce low spin complexes and high crystal field splittings. (e) cannot form low spin complexes. (d) usually produce high spin complexes and high crystal field splittings. (b) 13.410 amu/atom. (a) 0. 35Ar.388 amu (c) 0. A positron has a mass number of _____.0005486. (d) 10. (a) 4He (b) 16O (c) 32S (d) 55Mn (e) 238U 6. (b) usually produce low spin complexes and small crystal field splittings. a charge of _____.93 MeV (a) Both chromium metal ions are paramagnetic with 3 unpaired electrons.007277. proton 8. (Crystal Field Theory) Strong field ligands such as CN: (a) usually produce high spin complexes and small crystal field splittings. (a) 12. (d) Einstein postulated the Theory of Relativity in which he stated that matter and energy are equivalent.39 MeV (e) 56. Calculate the mass defect (amu/atom) for a 37Cl atom. lies below the "band of stability: (n/p ratio too low). (d) 3. electron = 0. 0.008665. and a mass equal to that of a(an) _____. .23 x 1023 kJ/mol 4. The "magic numbers" for atoms are (a) numbers of electrons that confer atomic stability.966 amu.(e) none of these 3.01219 amu. for which the atomic mass is 9. proton (c) 0. electron (e) 0. (b) 19.23 MeV (d) 11. (b) numbers of protons and/or neutrons that confer nuclear stability. (d) A solution of [Cr(OH2)6]Cl3 transmits light with an approximate wavelength range of 4000 . (b) Nuclear binding energy is the energy released in the formation of an atom from subatomic particles. Which of the following statements is incorrect? (a) Mass defect is the amount of matter that would be converted into energy if a nucleus were formed from initially separated protons and neutrons. (d) 8.69 x 1010 kJ/mol (b) 1. A radioisotope of argon. Conversion factor for E = mc2 is 931 MeV/amu. Which isotope below has the highest nuclear binding energy per gram? No calculation is necessary. (d) 2. electron (d) 1.23 x 1020 kJ/mol (c) 3. (d) 4. (b) 5. (a) 6.263 amu (d) 0. Emission of which one of the following leaves both atomic number and mass number unchanged? (a) positron (b) neutron (c) alpha particle (d) gamma radiation (e) beta particle Which type of radiation is the least penetrating? (a) alpha (b) beta (c) gamma (d) x-ray (e) neutron 10. 1+. (b) 20. 2+. 7. (a) 16. (b) 7. 20. (a) 14. (e) atomic masses that indicate fissile isotopes. (b) oct for [Cr(NH3)6]3+ is calculated directly from the energy of yellow light. 1+. (1 J = 1 kg m2/s2) (a) 3.623 amu (b) 0.23 x 103 kJ/mol (e) 1. (d) atomic masses that confer nuclear stability. (b) 15. 2. neutron = 1. (c) 9.46 MeV (b) 6. (e) Mass number is the sum of all protons and electrons in an atom.33 MeV (c) 6. 5. (a) 18.4200 angstroms. 40 microgram sample remains after 3.5 years. (b) Element Z will weigh exactly the same as element X when decay is complete (weighed to an infinite number of significant figures). (c) Rate of reaction is independent of the presence of a catalyst.0 minutes. (c) Two light nuclei are combined into a heavier one. (d) In fission reactions. . Its halflife is 20 minutes. 14. Which one of the following would be most likely to undergo thermonuclear fusion? (a) 2H (b) 4He (c) 56Fe (d) 141Ba (e) 235U 21. (b) A neutron is split into a neutron and proton. (d) A proton is split into three quarks. Which one of the following statements about nuclear reactions is false? (a) Particles within the nucleus are involved. A Geiger counter registered 1000 counts/second from a sample that contained a radioactive isotope of polonium. the counter registered 281 counts/second. Consider the case of a radioactive element X which decays by electron (beta) emission with a half-life of 4 days to a stable nuclide of element Z. (b) No new elements can be produced. 20. (d) If element X as an atomic number equal to n. (e) They are often accompanied by the release of enormous amounts of energy. (e) A particle and anti-particle appear in an area of high energy density. The missing term is _____ . If present-day plant life shows 15 dpm/g. then element X has an atomic number equal to n-1. 22.0 g of element X is required to produce 1. After 5. What fraction of the initial number of C-11 atoms in a sample will have decayed away after 80 minutes? (a) 1/16 (b) 1/8 (c) 1/4 (d) 7/8 (e) 15/16 How old is a bottle of wine if the tritium (3H) content (called activity) is 25% that of a new wine? The halflife of tritium is 12.04 micrograms 13.1 yr (c) 25 yr (d) 37. the new isotope formed is: (a) 210At (b) 212At (c) 211Po (d) 211Rn (e) 207Bi 25.5 g of element Z after 8 days (to 2 significant figures).5 yr (e) 50 yr 16. and _____ . (b) Nuclear fission is an energetically favorable process for heavy atoms. 239Pu + alpha particle _____ + neutron (a) 2 115Ag (b) 2 106Rh (c) 235U (d) 233Pa (e) 242Cm 23. (c) 2.25 x 104 years. (e) None of the above.0240 micrograms (e) 1. 19. What is the half-life of this isotope in seconds? (a) 87 (b) 110 (c) 164 (d) 264 (e) 2. When 59Cu undergoes positron emission.(c) positron emission (d) alpha emission (e) fission 11. The 14C activity of some ancient Peruvian corn was found to be 10 disintegrations per minute per gram of C. (e) All nuclear fission reactions are spontaneous.18 micrograms (d) 0. (a) 139Ba (b) 141Ba (c) 139Ce 15.0102 micrograms (b) 0. (d) Rate of reaction is independent of temperature. a neutron is split into a proton and an electron. 56Fe readily undergoes fission. As a result of the process of electron capture ("Kcapture") by 211At.25 x 105 years? (a) 0. Which of the following statements is CORRECT? (a) After 8 days the sample will consist of one-fourth element Z and three-fourths element X. When 235U is bombarded with one neutron. (a) 1/4 yr (b) 3. The half life of 231Pa is 3. Complete and balance the following equation. what is the immediate nuclear product? (a) 59Ni (b) 58Ni (c) 58Cu (d) 59Zn (e) 58Zn 24. Which of the following statements about nuclear fission is always correct? (a) Very little energy is released in fission processes. (a) gas ionization detector (b) cloud chamber (c) fluorescence detector (d) spectrophotometer (e) photographic detector 12. 94Kr. How much of an initial 10. A Geiger-Muller tube is a _____ . fission occurs and the products are three neutrons. (c) Due to its instability. Carbon-11 is a radioactive isotope of carbon. (a) 1455 years (b) 1910 years (c) 3350 years (d) 3820 years (e) 9080 years 18. how old is the Peruvian corn? The half-life of 14C is 5730 years.240 micrograms (c) 2.18 17. Which of the following describes what occurs in the fission process? (a) A heavy nucleus is fragmented into lighter ones. (a) 4. . (e) 23. 1. (c) 18. (d) 6. 2.4-dimethylcyclopentane (b) 1. 11. (c) 14. (a) 21. What makes carbon such a unique element? (a) Elemental carbon comes in two forms.(d) 139Xe (e) 142I Select the correct IUPAC name for: Answers to Chapter 26 1. (a) 13. (a) (a) 1.5-dimethylhexane (e) 3. (e) 7. (b) 22.3-dimethylcyclopentane (c) 2.1.5-dimethylcyclopentane (d) 2. Name the following compound: (a) 5-methyl-5-ethyloctane (b) 5-methyl-5-propylheptane (c) 4-ethyl-4-methyloctane (d) 3-methyl-3-propyloctane (e) 3-methyl-3-propylheptane Select the correct IUPAC name for: (a) 6-ethyl-4-methylcyclohexene (b) 6-ethyl-3-methylcyclohexene (c) 3-ethyl-5-methylcyclohexene (d) 6-ethyl-4-methylcyclohex-1-ene (e) 6. (e) 15. (a) 2-methyl-1-butene (b) 2-ethyl-1-propene (c) 2-ethyl-1-pentane (d) 3-methyl-2-butene (e) pentene 9. The hybridization of carbon atoms in alkanes is (a) sp (b) sp2 (c) sp3 (d) sp3d (e) sp3d2 3.5. carbon can bond to itself to form straight chains.3-dimethylcyclopentane (e) 2.4-dimethylhexane (d) 3. A molecule with the formula C3H8 is a(n): (a) hexane (b) propane (c) decane (d) butane (e) ethane 4. (b) Carbon forms four bonds.3-dimethylbutane (c) 2. (a) 19.4-dialkylcyclohexene (a) 1. (c) 25. What is the IUPAC name of the following compound? The general formula for noncyclic alkenes is: (a) CnH2n+2 (b) CnH2n (c) CnH2n-2 (d) CnHn+2 (e) CnHn The correct name for the compound given below is: 8. (d) 3. (b) 20. (a) 12. (c) Carbon forms covalent bonds rather than ionic bonds. carbon-12 and carbon-13.4-dimethylcyclopentane 7. (c) 16. (d) To a greater extent than any other element. although the ground state configuration would predict the formation of fewer bonds.5-trimethylpentane 6. (b) 2. (e) Carbon has two stable isotopes. (a) 5. (a) 10. Select the best name for: 5.3-trimethylpentane (b) 1-ethyl-1. (c) 8. diamond and graphite. (a) 24. branched chains and rings. (d) 9. (c) 11. (c) 17. Select the correct IUPAC name for: (a) 4-ethyl-cis-3-octene (b) 4-ethyl-trans-3-octene (c) 4-butyl-cis-3-hexene (d) 5-ethyl-trans-5-octene (e) 5-ethyl-cis-5-octene 10. 4-dimethylpentanone (d) 2-carboxyisohexane (e) none of these 24. 15.3-diethylbenzene (c) 1.6-diethyl-3-nonyne (b) 2. carbon-carbon bonds in benzene are delocalized around the ring (b) 1 double bond (c) 2 double bonds (d) 3 double bonds (e) 4 double bonds 14.3-trimethylbutanoic acid (c) 1-hydroxy-2.7-dimethyl-5-nonyne (d) 3. (e) Boiling points of normal primary alcohols increase with increasing molecular weight. (a) 1.2-dimethylbenzene (b) 1.1-dimethylisopentanol (d) 2. Select the best name for: . boiling points of alcohols are much higher than those of corresponding alkanes. (c) The hydroxyl group of an alcohol is nonpolar. Select the IUPAC name for: (CH3)2CHCH(OH)CH2C(CH3)3. (a) pentyl amine (b) methyl-n-propyl amine (c) diethyl amine (d) 2-aminopentane (e) isobutylamine 23. Select the IUPAC name for the compound below.1.6-diethyl-3-heptyne 12. (d) Due to hydrogen bonding.4-dimethylbenzene Which of the following formulas represents an alkene? (a) CH3CH2CH3 (b) CH3CH3 (c) CH3CH2CHCH2 (d) CH3CH2Cl (e) CHCH 16.5-dimethyl-4-hexanol (e) none of these (a) C8H10 (b) C6H6 (c) C6H8 (d) C8H12 (e) C8H6 13.5. The following chemical structure represents a molecule of what molecular formula? 19. What is the name of the following compound? (a) primary amine (b) secondary amine (c) tertiary amine (d) primary amide (e) secondary amide 22.4.4-diethylbenzene (e) 1.4-pentamethylbutanol (c) 1.(a) CH3CH2OH (b) CH3OH (c) CH3CH(OH)CH3 (d) (CH3)C3OH (e) none of these 18.4-dimethylpentanoic acid (b) 1. The systematic name for the compound in Problem 21 is _____ . 20.7-dimethyl-4-nonyne (e) 2.5-diethyl-3-nonyne (c) 3.6-dibromophenol (d) m-dibromophenol (e) o-dibromophenol 17. Which is NOT a physical property of alcohols or phenols? (a) Phenols are generally only slightly soluble in water. Give the IUPAC name of this compound: CH3OCH2CH3. (a) dimethyl ether (b) methoxyethane (c) methylethyloxide (d) propyl ether (e) none of the above 21. (b) The solubilities of normal primary alcohols in water decrease with increasing molecular weight. (a) 2. (a) 2. Para-xylene is the same as: (a) 1.3-dimethylbenzene (d) 1.5-trimethyl-3-hexanol (b) 1. Which one of the following is a secondary alcohol? (a) 2.1.5-dibromophenol (c) 2. The compound below is classified as a _____ .3-dibromophenol (b) 2. How many actual double bonds does the benzene ring possess? (a) None. How many aromatic isomers of dibromobenzene exist? (a) 2 (b) 3 (c) 4 (d) 6 (e) 8 5. (c) 5. (a) 1. The functional group given below is characteristic of organic _____ . (a) aldehyde (b) ester (c) ketone (d) carboxylic acid (e) alcohol 26. C3H7OH 7. (a) 9. (c) 17. CH3CHO (b) acetic acid. (a) 10.N-dimethylformamide (e) dimethylamine 28. (d) have a different content of the isotopes of hydrogen. For which of the compounds below are cis-trans isomers possible? CH3CH=CHCH2CH3 CH3CH=CHCH3 (2) (a) only 2 (b) both 1 and 2 (c) both 2 and 3 (d) all three (e) only 3 8. The best classification for the following compound is: _____ . Which of the following compounds is a functional group isomer of C2H5OH. (e) react vigorously with one another. (a) 24. Answers to Chapter 27 1.3-dimethylbutane (e) cyclohexane 4. (CH3)2O (e) propanol. Two isomeric forms of a saturated hydrocarbon (a) have the same structure. Which of the following does NOT exhibit geometric isomerism? (Hint: draw them!) (a) 4-octene (b) 2-pentene (c) 3-hexene (d) 2-hexene (3) (a) m-chlorobenzoic acid (b) o-chlorobenzaldehyde (c) p-chlorobenzoate (d) m-chlorosalicylic acid (e) none of these 25. Which one of the following compounds is an isomer of CH3CH2CH2CH2OH? (a) CH3CH2CH2OH (b) CH3CH(OH)CH3 (c) CH3CH2CH2CHO (Note: This is one way to write an aldehyde. Which of the following hydrocarbons does not have isomers? (a) C7H16 (b) C6H14 (c) C5H10 (d) C4H8 (e) C3H8 3. CH3CH=CH2 (a) ketones (b) acids (c) aldehydes (d) esters (e) alcohols (1) . (d) 23. (c) 16. CH3COOH (c) diethyl ether. (a) 14. (e) 15. (C2H5)2O (d) dimethyl ether. (b) 22. (b) 7. (b) have different compositions of elements. (d) 28. (d) 12. The compound illustrated below is called _____ . The compound given below is called _____ . (c) 18. (c) 11. (c) 3.(c) 20. (a) acetamide (b) formyl acetamide (c) dimethyl acetate (d) N. (a) 13. (b) 8. (b) 4.) (d) CH3CH2CH2CH3 (e) none of the above 6. (a) 19. The name of the alkane isomer of cis-3-hexene is: (a) 2-methylpentane (b) 3-methylpentane (c) n-hexane (d) 2. (b) 21. ethanol (ethyl alcohol)? (a) ethanal. (b) 27. (a) butyl acetate (b) ethyl pentanoate (c) propyl pentanoate (d) ethyl butanoate (e) butyl ethanoate 27. (b) 26. (c) have the same molecular formula. (c) 6. (a) 25. 2. (d) 2. 3-dichlorobutane (d) 2. (a) alkene (b) alkane (c) alkyne (d) alkyl halide (e) aldehyde A reaction in which a carboxylic acid reacts with a base to form a salt and water is called _____ . then HBr (b) HCl. then Br2 (e) H2. 17. (d) The reaction can be initiated with either sunlight or heat. (2) The eclipsed conformation of a molecule is slightly more stable and energetically favored than the staggered conformation. Which of the following compounds displays optical isomerism? (a) CH2(OH)-CH2(OH) (b) CH3-CHCl-COOH (c) CH2=CHCl (d) CHCl=CHCl (e) CH3-O-C2H5 10. What is the first stable intermediate product when ethanol is oxidized with a mild oxidation agent? (a) CH3COOH (b) CO2 (c) CH3CHO (d) CH3CH2OH (e) CH3OCH3 22. Which of the following statements is FALSE regarding the reaction between Cl2 and C2H6? (a) It is a substitution reaction. Which of the following alcohols forms a ketone when oxidized? (a) 1-propanol (b) methanol (c) 2-methyl-2-propanol (d) 2-propanol (e) all of the above 23. How many moles of sodium hydroxide will react with one mole of: (a) structural isomers (b) geometric isomers (c) conformational structures (d) identical structures (e) optical isomers 13. (b) The reaction will give a single product of C2H5Cl. (a) 1 only (b) 2 only (c) 1 and 2 (d) 2 and 3 (e) 1 and 3 14. What is the sum of the coefficients in the balanced equation for the complete combustion of 2- . (c) The reaction mechanism involves free radicals. 18. then HBr (d) Cl2. Which of the following will undergo an addition reaction with chlorine? (a) CH3CH2CH2CH3 (b) CH3CH2CH=CHCH3 (c) C6H6 (a) 5 (b) 4 (c) 3 (d) 2 (e) 1 21. (d) CH3CH2COOH (e) CH3CH2OH What is the expected product formed from the reaction between 2-butene and Cl2? (a) 1-chlorobutane (b) 2-chlorobutane (c) 2.3-dichlorobutane The reaction of ethyne with which of the following gives CH2Br-CHBrCl? (a) HCl. What is the relationship between the structures shown? 19. 16.2-dichlorobutane (e) 3. How many alcohols are structural isomers with the formula: C5H11OH? (a) 5 (b) 6 (c) 7 (d) 8 (e) 9 12. Ethanol can be oxidized stepwise. then Br2 (c) Cl2. (e) The first step in the mechanism is the cleavage of the Cl-Cl bond to give chlorine atoms. (3) A conformation is one specific geometry of a molecule. Which of the following statements concerning conformations is (are) TRUE? (1) Ethane has an infinite number of conformations. (a) ionization (b) esterification (c) hydrolysis (d) saponification (e) neutralization 20. then Br2 Dehydration of an alcohol leads to the formation of an _____ .(e) 1-hexene 9. 15. How many isomeric alkanes of the molecular formula C5H12 are there? (a) 1 (b) 2 (c) 3 (d) 4 (e) 5 11. Do not forget coefficients of 1. methylbutane? Use smallest whole number coefficients. (c) 2. (a) an acid and an alcohol (b) a ketone and an alcohol (c) an alkane and a ketone (d) only an acid (e) an amine and an acid 25. (b) 18. (d) 12.represents the polymer named _______ . (c) 26. (a) polybutylene (b) polyhexene (c) polypropylene (d) polystyrene (e) polyethylene Answers to Chapter 28 1. (e) 21. (e) . The segment -CH2CH2CH2CH2CH2CH2. (a) 10 (b) 13 (c) 17 (d) 20 (e) 23 The organic starting materials for the preparation of an ester could be _______ . (e) 4. (c) 11. (d) 7. (b) 16. (b) 10. (c) 17. (c) 13. (e) 14. (a) 19. (e) 3. (a) 25. (c) 8. (e) 9. Hydrolysis (saponification) of a fat would yield ______ . (e) 6. (a) water and an alkene (b) ethanol and propanoic acid (c) glycerol and soap (d) ethanol and a soap (e) a triester of glycerol with fatty acids 26. (b) 5. (c) 22. (d) 24. (e) 20. (d) 23.24. (b) 15.
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