ChemistryUnit 1 – Atomic Theory and Structure of atom TOPICS Atomic Theory and Structure of atom Discovery of Electrons,Properties of Cathode Rays, Characteristics of an electron, Discovery of Proton, Goldstein’s experiment’ Properties of Anode rays (Canal Rays), Discovery of Neutron, Atomic structure- Postulates of Daltons atomic theory, Thomson’s Model of an Atom, .Rutherford’s Model of an Atom & its drawbacks, Bohr’s Model of an Atom & limitations, Bohr’s model of atom revisited, Atomic spectrum of hydrogen, Merits of Bohr’s atomic model,, Arrangement of electrons in the atoms, Atomic orbitals - Quantum Numbers- Principal quantum number, Azimuthal quantum number magnetic quantum number, spin quantum number, Electronic Configuration of Atoms of Elements - Aufbau principle, Pauli’s Exclusion principle, Hund’s Rule Shape of orbitals - s,p. -Valence electrons, Atomic Number, Mass Number, Isotopes, Isobars, Isotones, Representation of an element. Learning Outcome At the end of the unit the student must be able to • Know the properties of electron, proton, and neutron. • Understand Dalton’s Atomic theory and its draw backs • Understand Thomson’s model of an atom • Explain Rutherford’s model of an atom and its draw backs • Describe Bohr’s model of an atom and its limitations • Explain Atomic spectrum of hydrogen • Give the arrangement of electrons in the atom • Know the various atomic orbitals and their quantum numbers such as principal quantum number, magnetic quantum number, spin quantum number • Explain electronic configuration of atoms of first 20 elements. • Know principals and rules which provide guidance to write electronic configuration - Aufbau principle, Pauli’s Exclusion principle, Hund’s Rule. • Understand various shapes of orbitals- s, p • Understand the terms valence electron, atomic number, atomic mass , Isotopes, Isobars, Isotones, • Represent an element with nucleon and proton 1.0. Atomic Theory 1.1. Introduction: Any object which occupies space and has mass is known as matter. There are three kinds of matter viz., solids, liquids and gases. On the basis of chemical constitution, matter can be further classified as elements, compounds and mixtures. An element is a substance which cannot be split up into two or more simple substances by any physical or chemical means. It is now clear that element is made up of atoms. John Dalton in 1808, proposed his famous atomic theory. He defined atom as the smallest chemically indivisible particle of an element which cannot exist free. They exist in combination with the same element or another element. Atoms are the building blocks of all the matter around us. FPCHE001KB 1 www.smartlearnwebtv.com Chemistry Extensive research has been done to understand the composition of matter in the twentieth century by J .J. Thomson, Rutherford, Neils Bohr, Chadwick and many other scientists. All these led to the conclusion that Dalton’s concept “an atom is indivisible” is not correct. Atoms are found to be made up of electrons, protons and neutrons. These are known as sub-atomic particles. 1.1.1. Discovery of Electrons J.J. Thomson in 1897, through his ‘discharge tube’ experiments confirmed the presence of electrons in an atom. A discharge tube is a long glass tube having two metal plates sealed at two ends. These metal plates are known as electrodes. The electrode which is connected to the positive terminal of the battery is known as anode (positive electrode) and the negative electrode which is connected to the negative terminal of the battery is called cathode (negative electrode). The discharge tube has a side tube through which air (other gases) can be pumped out by using a vacuum pump. Thomson carried out his experiments using this discharge tube. 1. When high voltage is passed through the discharge tube which contains air or any gas, at normal atmospheric pressure, no electricity flows through the air in the discharge tube. 2. If the pressure of air inside the discharge tube is reduced to 1 mm of mercury and when high voltage is applied, electricity flows through air and light is emitted by the air in the discharge tube. The color of the light depends on the nature of the gas inside the tube. 3. When the pressure of the air inside the discharge tube is reduced to about 0.001 mm mercury and a high voltage is applied between the electrodes, the emission of light by air stops. The discharge tube appears dark. The walls of the discharge tube, opposite to the cathode glows with greenish light. − + Cathode Discharge tube Air at very low pressure Green glow Anode To vacuum pump High voltage generator + − − − − − Figure 1. Production of Cathode Rays This experiment reveals that the rays which are emitted by the cathode are called cathode rays. These rays travel from cathode to anode in the form of streaks of light. The cathode rays bombard the glass which results in the glowing of the tube. FPCHE001KB 2 www.smartlearnwebtv.com Chemistry 1.1.2. Properties of Cathode Rays The important properties of cathode rays are: 1.1.2.1. They travel in straight line: This is proved by the following experiment. When an opaque object is placed in the path of cathode rays in a discharge tube, a shadow of the object appears at the end of the cathode. Cathode Cathode rays − Anode (Metal cross) Shadow of the metal cross + Figure 2. Cathode rays cast shadows of the object places in their path This is possible only if the cathode rays travel in straight lines. 1.1.2.2. Cathode rays are a beam of particles having mass and kinetic energy: A paddle wheel placed in the path of the cathode rays, rotates. This is possible only when the cathode rays consist of particles moving with certain velocity. i.e., they have mass and kinetic energy. Cathode Cathode rays Anode Light paddle wheel Figure 3. Cathode rays can rotate a light paddle wheel placed in their path 1.1.2.3. Cathode rays are negatively charged: When an electric field is applied in the path of cathode rays, they get deflected towards the positive plate of the electric field. Cathode Cathode rays − + − − − − − − − − Electric field − Anode Figure 4. Effect of electric field on cathode rays. (They are deflected towards the positive plate of electric field) FPCHE001KB 3 www.smartlearnwebtv.com Chemistry This shows that the cathode ray particles are negatively charged. Cathode rays are also deflected by magnetic field. 1.1.2.4. Cathode ray beam deviated from its path in the presence of a magnetic field: When magnetic field is applied perpendicular to the path of cathode rays the beam gets deflected towards a pole. Figure 5. Effect of magnetic field on cathode rays 1.1.2.5. The nature of the cathode rays are independent of the nature of the gas taken in the discharge tube: This means that the same type of rays with similar mass and charge having similar properties are produced irrespective of the nature of the gas (i.e., H2 (or) O2 or Cl2 or any gas) taken in the tube. 1.1.3. Characteristics of an electron 1. An electron is a negatively charged particles. 2. The mass of an electron is 1 1840 of a hydrogen atom. The absolute mass of an electron is 9 × 10 –28 g. 3. The relative charge of an electron is –1. The absolute charge is 1.6 × 10 –19 coulombs. (It has been found that this is the smallest negative charge carried by the particle). 1.2. Discovery of Proton If an atom consists of an electron, it should have some positively charged particles to maintain electrical neutrality. These positively charged particles are known as protons. The existence of protons in an atom is shown by Goldstein. 1.2.1. Goldstein’s Experiment Goldstein used a perforated cathode (i.e., a cathode which has small holes) in the discharge tube experiment. The mass of the cathode rays (m) and the charge on them (e) have been determined experimentally and the e m ratio is found to be the same for all gases. This shows that atoms of all kinds contain the same negatively charged particles. These are known as electrons and are the fundamental constituent of atoms. FPCHE001KB 4 www.smartlearnwebtv.com In which metal atom did Rutherford discover the nucleus? Cathode Porous anode (with only one hole), Discharge tube High voltage source Switch Semi circular magnet Zinc sulphide screen Chemistry + − Positive rays Discharge tube Air at very low pressure Red glow Anode High voltage generator − + + + + + To vacuum pump + Perforated cathode Figure 6. Production of anode rays or positive rays. When very high voltage (about 10,000 volts) is applied to the discharge tube containing air at about 0.001mm Hg a faint glow is obtained behind the cathode. (See figure 6 above) These rays are believed to emanate from the anode, moving in the direction opposite to the cathode rays. Hence, these rays are known as anode rays or positive rays. 1.2.2. The properties of Anode (or Positive) Rays (Canal rays) 1. The anode rays consist of positively charged particles. These rays are deflected by the electric and magnetic fields. 2. They travel in straight lines. 3. They have mass and kinetic energy. 4. The e m ratio depends on the nature of the gas taken in the discharge tube. 5. The mass of an anode ray particle is almost equal to the mass of the atom from which it is formed. 6. The particle produced from hydrogen is the lightest and is known as proton. A proton has a mass +1 units. 1.3. Discovery of Neutron: Because protons and electrons have equal but opposite charges, a neutral atom must contain equal numbers of protons and electrons. But solving this mystery led to another: the mass of an atom (except hydrogen atoms)is known to be greater than the combined masses of the atom’s protons and electrons Hoping to find a reason for the rest of the mass scientists began to search for a third subatomic particle About 30 years after the discovery of the electron, Irene Juliet-Curie (the daughter of the famous scientists Marie and Pierre Curie) discovered that when alpha particles hit a sample of beryllium, a beam that could go through almost anything was produced. The British scientist James Chadwick found that this beam was not deflected by electric or magnetic fields. He concluded that the particles carried no electric charge. Further investigation showed that these neutral particles, which were named neutrons are part of all atomic nuclei (except the nuclei of most hydrogen atoms). In general, a neutron is represented as ‘n’. The mass of an atom is therefore given by the sum of the masses of protons and neutrons present in the nucleus. FPCHE001KB 5 www.smartlearnwebtv.com In what respects cathode rays and anode rays differ Chemistry Figure 7 Structure of an atom Electron Proton Neutron (i) Discoverer J.J. Thomson Goldstein Chadwick (ii) Mass 1 1840 of the mass of a hydrogen atom mass is equal to that of a hydrogen atom mass is same as that of a proton (iii) Charge has uni negative charge has uni positive charge no charge (iv) symbol denoted by ‘e’ ‘p’ ‘n’ Table 1 Facts about the three sub-atomic particles. 1.4. Atomic Structure 1.4.1. Postulates of Daltons atomic theory. (1) All elements are composed of tiny indivisible particles called atoms (2) Atoms of the same element are identical. Atoms of any other element are different from those of any other element. (3) Atoms of different elements combine in simple whole number ratios and this combination leads to form chemical compounds (4) In chemical reaction atoms are combined and rearrangement can occur but never changed into atoms of another element. (5) Atoms of the same element are identical in their physical and chemical properties while atoms of the different element differ in their physical and chemical properties. 1.4.2. Drawbacks of Dalton’s atomic theory Dalton’s contention that all elements composed of tiny indivisible particles is not true. Atoms are divisible. The second criteria that atoms of the same element are also not identical in all respects. For example – atoms of the same element may possess different relative masses. Another drawback is, by his theory one cannot explain the concept of a molecule which is a simple combination of atoms. For example the O2 molecule formation from two individual atoms of Oxygen FPCHE001KB 6 www.smartlearnwebtv.com Chemistry 1.5. The structure of an Atom To overcome the drawbacks of Dalton’s atomic theory many scientists proposed various atomic models. J.J. Thomson was the first one to propose a model for the structure of an atom. 1.5.1. Thomson’s Model of an Atom Thomson proposed the model of an atom to be similar to that of a Christmas pudding. The electrons, in a sphere of positive charge, were like currants (dry fruits) in a spherical Christmas pudding. We can also think of a watermelon, the positive charge in the atom is spread all over like the red edible part of the watermelon, while the electrons are studded in the positively charged sphere, like the seeds in the watermelon. Positive sphere Electron − − − − − − − − − − Figure 8. Thomson’s model of an atom Figure 9 Atom is electrically neutral Thomson proposed that 1. An atom consists of a positively charged sphere and the electrons are embedded in it. (Figure 8) 2. The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral (Figure 9). Although Thomson’s model explained that atoms are electrically neutral, the results of experiments carried out by other scientists could not be explained by this model, FPCHE001KB 7 www.smartlearnwebtv.com Chemistry 1.5.2. Rutherford’s Model of an Atom Figure 10. Scattering of α -particles by a gold foil Ernest Rutherford was interested in knowing how the electrons are arranged within an atom. Rutherford designed an experiment for this. In this experiment, fast moving alpha (α )-particles were made to fall on a thin gold foil. • He selected a gold foil because he wanted a layer as thin as possible. This gold foil was about 1000 atoms thick. • α -particles are doubly charged helium ions. Since they have a mass of 4 u, the fast moving α -particles have a considerable amount of energy. • It was expected that α -particles would be deflected by the sub-atomic particles in the gold atoms. Since the α -particles were much heavier than the protons, he did not expect to see large deflections. But, the α -particle scattering experiment gave totally unexpected results. The following observations were made: (iii) Most of the fast moving α -particles passed straight through the gold foil. (ii) Some of the α -particles were deflected by the foil by small angles. (iii) Surprisingly one out of every 12000 particles appeared to rebound. In the words of Rutherford, “This result was almost an incredible as if you fire a 15-inch shell at a piece of tissue paper and it comes back and hits you”. Activity Let us think of an activity in an open field to understand the implications of this experiment. Let a child stand in front of a wall with his eyes closed. Let him throw stones at the wall from a distance. He will hear a sound when each stone strikes the wall. If he repeats this ten times, he will hear the sound ten times. But if a blind-folded child were to throw stones at a barbed-wire fence, most of the stones would not hit the fencing and no sound would be heard. This is because there are lot of gaps in the fence which allow the stone to pass through them. Following a similar reasoning, Rutherford concluded from the α -particle scattering experiment that (i) Most of the space inside the atom is empty because most of the α -particles passed through the gold foil without getting deflected. (ii) Very few particles were deflected from their path indicating that the positive charge of the atom occupies very little space. (iii) A very small fraction of α -particles were deflected by 180°, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom. FPCHE001KB 8 www.smartlearnwebtv.com Chemistry From the data he also calculated that the radius of the nucleus is about 10 5 times less than the radius of the atom. On the basis of his experiment, Rutherford put forward the nuclear model of an atom which had the following features. (i) There is a positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus. (ii) The electrons revolve around the nucleus in well-defined orbits. (iii) The size of the nucleus is very small as compared to the size of the atom. 1.5.3. Drawbacks of Rutherford’s model of the atom The orbital revolution of the electron is not expected to be stable. Any particle in a circular orbit would undergo acceleration. During acceleration, charged particles would radiate energy. Thus, the revolving electron would lose energy and finally fall into the nucleus. If this were so, the atom should be highly unstable and hence matter would not exist in the form that we know. But we know that atoms are quite stable. 1.5.4. Bohr’s Model of an Atom In order to overcome the objections raised against Rutherford’s model of the atom, Neils Bohr put forward the following postulates about the model of an atom: (i) Only certain special orbits known as discrete orbits of electrons, are allowed inside the atom. (ii) While revolving in discrete orbits the electrons do not radiate energy. These orbits or shells are called energy levels. Nucleus K shell (n = 1) L shell (n = 2) M shell (n = 3) N shell (n = 4) Figure 11. A few energy levels in an atom These orbits or shells are represented by the letters K, L, M, N, … or the numbers, n = 1, 2, 3, 4, 1.5.5. Limitation of Bohr Model (i) He could not explain the spectra of atoms containing more than one electron (ii) Bohr assumed that the electron revolves around the nucleus in a definite and well-defined path called orbit which was soon proved wrong. Though Bohr model explains the stability of the atom, it could not explain certain experimental observations viz., the atomic spectrum of atoms other than hydrogen. For the present, it is suffice to know that Bohr model cannot explain the present day understanding of the atom. However, it should be appreciated, that the Bohr model provided the basic structure to understand the arrangement of electrons, protons and neutrons in an atom. FPCHE001KB 9 www.smartlearnwebtv.com Chemistry 1.5.6. Bohr’s model of atom Revisited Next in the development of atomic structure is the model of atom as given by Neils Bohr. His postulates regarding the structure of atom are (i) Electrons revolve in certain orbits called stationary states. These are called stationary states because an electron moving along these orbits do not radiate energy. The energy remains the same. (ii) Each stationary state is associated with a definite amount of energy. So these are also called energy levels. These are called K, L, M, N …….shells and are also denoted by numbers 1, 2, 3, 4……….etc., (iii) The energy level nearest to the nucleus is the lowest. With increasing distance from the nucleus in order, the energy of the level increases. An electron in the lowest energy level is said to be in the ground state. (iv) If energy is given to an electron, it jumps to a higher energy level. Now the electron is said to be in an excited state. It is not stable there. It gives up the energy and falls to the ground state. Figure 12 Various energy level of electron in an atom 1.5.7. Atomic spectrum of hydrogen: Bohr explained his idea of stationary orbits based on line spectrum of hydrogen. When an electron absorbs energy, it goes to any one of the higher energy levels depending on the quantum of energy absorbed. When it emits the excess energy absorbed, it falls back to the ground state in one or more jumps. Corresponding to each jump, a line is observed in the spectrum of hydrogen at a definite frequency. Hydrogen produced by different methods give similar lines at the same frequencies in its spectrum. This supports the postulate of Neils Bohr that only certain orbits are present in the hydrogen atom which have definite quantities of energies associated with them. Also it has been found that the calculated frequency of any spectral line of hydrogen spectrum from the Bohr’s formula is in excellent agreement with the experimentally determined value. [Frequency (γ ) of spectral line is related to the energy emitted by the equation, Energy = hν , where h is a constant]. 1.5.8. Merits of Bohr’s model (i) Bohr could calculate the energy of electron in different orbits. (ii) He could also calculate the radius of the orbit of hydrogen’s electron 1.5.9. Arrangement of electrons in the atoms The Bohr model of atom helps us to understand how the electrons are arranged in different energy levels. The energy levels are represented by circles. The energy of the energy level FPCHE001KB 10 www.smartlearnwebtv.com Chemistry increases as it moves away from the nucleus. The number of electrons in an energy level is given by the formula 2n 2 , where ‘n’ is the number of the orbit. Thus for K shell = n = 1 and no. of electrons = 2 This means that the first energy level nearer to the nucleus can hold a maximum of two electrons. The next energy level (L shell)for which n =2, can hold a maximum of 8 electrons. The next higher energy level i.e., the M shell for which n =3, can hold a maximum of 18 electrons and so on. + 1 2 3 4 5 K L M N O Energy levels Nucleus Figure 13. Energy levels around the nucleus While filling the electrons in various energy levels, it must be remembered that the outermost energy level can have a maximum of eight electrons. For example, chlorine has 17 electrons. Of these 17 electrons, 2 electrons are present in the K shell, 8 electrons in L shell and the remaining 7 electrons in M shell. Diagrammatic representation of the arrangement of electron for 20 elements is given below: FPCHE001KB 11 www.smartlearnwebtv.com Chemistry FPCHE001KB 12 www.smartlearnwebtv.com Chemistry Figure 14. Electronic configuration of first twenty elements The outermost electrons shell (energy level) is known as valence shell and the electrons present in it are known as valence electrons. Refer to arrangement of electrons for carbon atom. For this atom, the L shell is valence shell and the four electrons present in it are known as valence electrons. Similarly you will find that in sodium atom, the ‘M’ shell is the valence shell and it has one valence electron. The arrangement of electrons in various energy levels is known as the electronic configuration of atoms. The electronic configuration of twenty elements is given below: Element Symbol Atomic Number Distribution of electrons Electronic configuration Hydrogen H 1 1 1s 1 Helium He 2 2 1s 2 Lithium Li 3 2,1 1s 2 ,2s 1 Beryllium Be 4 2,2 1s 2 ., 2s 2 Boron B 5 2,3 1s 2 , 2s 2 , 2p 1 Carbon C 6 2,4 1s 2 , 2s 2 2p 2 Nitrogen N 7 2,5 1s 2 , 2s 2 , 2p 3 Oxygen O 8 2,6 1s 2 , 2s 2 , 2p 4 Fluorine F 9 2,7 1s 2 , 2s 2 , 2p 5 Neon Ne 10 2,8 1s 2 , 2s 2 , 2p 6 Sodium Na 11 2,8,1 1s 2 , 2s 2 , 2p 6 ,3s 1 Magnesium Mg 12 2,8,2 1s 2 , 2s 2 2p 6 , 3s 2 Aluminium Al 13 2,8,3 1s 2 , 2s 2 2p 6 , 3s 2 3p 1 Silicon Si 14 2,8,4 1s 2 , 2s 2 2p 6 , 3s 2 3p 2 FPCHE001KB 13 www.smartlearnwebtv.com Chemistry Phosphorus P 15 2,8,5 1s 2 , 2s 2 2p 6 , 3s 2 3p 3 Sulphur S 16 2,8,6 1a 2 ,2s 2 2p 6 , 3s 2 3p 4 Chlorine Cl 17 2,8,7 1s 2 , 2s 2 2p 6 3s 2 3p 5 Argon Ar 18 2,8,8 1s 2 , 2s 2 2p 6 ,3s 2 3p 6 Potassium K 19 2,8,8,1 1s 2 , 2s 2 2p 6 , 3s 2 3p 6 ,4s 1 Calcium Ca 20 2,8,8,2 1s 2 ,2s 2 2p 6 , 3s 2 3p 6 ,4s 2 Table 2 Electronic configuration of first twenty elements . 1.6. Atomic orbitals Bohr’s model of atom underwent great changes due to two reasons (i) Heisenberg’s uncertainty principle (ii) Dual nature (particle and wave nature) of electrons. We will proceed to consider the changes brought about to Bohr’s model without going into the details of the above stated reasons. The path of an electron around the nucleus is no longer a definite path or orbit as conceived by Bohr. It is called orbital. An atomic orbital is defined as three dimensional space or region around the nucleus in which the probability of finding the electron is maximum. Bohr’s model of atom underwent a lot of changes and a Quantum mechanical model involving atomic orbitals was evolved. Bohr called his permitted orbits or stationary states as shells and named them K, L, M, N………..etc., These are given numbers 1, 2, 3, 4…………etc., Each shell contains one or more sub-shells or orbitals. The number of sub shells is equal to the number of the shell. For K shell n=1 and it contains s sub shell. For L shell n=2 and it contains s, p sub shells For M shell n=3 and it contains s, p, d sub shells For N shell n=4 and it contains s, p, d, f sub shells. The maximum number of electrons that can be accommodated in a shell is given by 2n 2 where n is the number of the shell (1, 2, 3 …….etc) So the maximum number of electrons that can be accommodated in K shell (n = 1) 2 × 1 2 = 2 L shell (n = 2) 2 × 2 2 = 8 M shell (n = 3) 2 × 3 2 = 18 N shell (n = 4) 2 × 4 2 = 32. Orbit Orbital 1. It is a definite path around the nucleus along which the electron moves. It is a region or space around the nucleus where the probability of finding the electron is maximum 2. It represents movement of electron in one plane. It represents three dimensional space around the nucleus 3. Its shape is circular. Its shape may be spherical (for s orbital) dumb- bell (p-orbitals) or any other shape (d, f) FPCHE001KB 14 www.smartlearnwebtv.com Chemistry 4. The position and velocity of the electron can be found precisely at any instant. It is impossible to determine the position and velocity of electron at any instant with certainty. Table 3 Differences between orbit and orbital 1.6.1. Quantum Numbers The numbers which designate and distinguish various atomic orbitals and electrons present in an atom are called quantum numbers. In an atom, the state of each electron is different with respect to the nucleus. In order to define the state of the electron completely, four quantum numbers are used. They are (a) Principal quantum number (n) (b) Azimuthal quantum numbers () (c) Magnetic quantum number (m) (d) Spin quantum number (s) 1.6.2. Principal Quantum Number (n) 1. Principal quantum number determines the energy shell in which the electron is revolving around the nucleus. It is also known as major energy level. 2. It is denoted by the symbol n and may have any integral value except zero. i.e., it can have the value n = 1, 2, 3, …, etc. 3. The value n = 1 denotes that the electron is in the first shell (K shell) The value n = 2 denotes that the electron is in the second shell (L shell) The value n = 3 denotes that the electron is in the third shell (M shell) The value n = 4 denotes that the electron is in the fourth shell (N shell) 4. As the distance of the electron from the nucleus increases, its energy becomes higher and higher. 5. The maximum number of electrons in a major energy level is given by 2n 2 . Principal quantum number ‘n’ Shell Designation Maximum number of electrons (2n 2 ) 1 2 3 4 5 K L M N O 2 8 18 32 50 Table 4. Maximum number of electrons in an shell 1.6.3. Azimuthal Quantum Number or Orbital Quantum Number () 1. It represents the sub shell to which the electron belongs. 2. It is denoted by the symbol . Its value depends on the principal quantum number n. It may have any value ranging from 0 to (n – 1) Principal quantum number ‘n’ Value = (n – 1) Name of the sub shells or orbital 1 0 1s 2 0 2s 1 2p FPCHE001KB 15 www.smartlearnwebtv.com Chemistry 3 0 3s 1 3p 2 3d 4 0 4s 1 4p 2 4d 3 4f Table 5. Name of sub shells or orbitals 3. The value = 0 denotes that the electron is in the sub shell or s orbital. The value = 1 denotes that the electron is in the p sub shell or p orbital. The value = 2 denotes that the electron is in the d sub shell or d orbital The value = 3 denotes that the electron is in the f sub shell or f orbital 1.6.4. Magnetic Quantum Number (m) 1. It represents the orientation of an atomic orbital in space. 2. It is denoted by the symbol m. The possible value which m can have depends upon the value of . It may have all the integral values between – to + through 0 that is the total number of values of m would be (2 + 1). 3. Its value tells the orientations of orbital in space. The value of m = 0 denotes that the orbital has no orientation. The value of m = 1 denotes that it has three orbital with three types of orientations. The value of m = 2 denotes that it has five orbital with five types of orientations. Principal Quantum number ‘n’ value = (n – 1) m value (– ..0.. + ) Name of the sub shells or orbital with orientation 1 0 0 1s 2 0 0 2s 1 –1, 0, +1 2px, 2py, 2px 3 0 0 3s 1 –1, 0, +1 3px, 3py, 3pz 2 –2, –1, 0, +1, +2 3dxy, 3dxz, 3dyz, 2 2 2 z x y 3d , 3d − Table 6 Magnetic quantum number with their orientations 1.6.5. Spin Quantum Number(s) 1. It represents the direction of the spin of the electrons. 2. It is denoted by the symbol s. The electron may spin in the clockwise ↑ direction or anticlockwise ↓ direction. And hence it can have only two values namely either + 1/2 or – 1/2. 3. Two electrons with the same sign of spin are said to have parallel spins and are represented by ↓↓(or) ↑↑while those having opposite spins are said to have anti parallel spins ↑↓and are known as paired up electrons. The above four quantum numbers give the position of any electron in the major energy level, the orientation of electron in the orbital and the direction of its spin. The various states that an electron can occupy are summarized in the table given below. Principal Quantum Number n Azimuthal Quantum Number Magnetic Quantum number m Total Number of electrons 2n 2 n = 1 K shell = 0; (1s) m = 0 2 × 1 2 = 2 FPCHE001KB 16 www.smartlearnwebtv.com Chemistry n = 2 L shell = 0; (2s) m = 0 2 × 2 2 = 8 = 1; (2p) m = –1, 0, +1 n = 3 M shell = 0; (3s) m = 0 2 × 3 2 = 18 = 2; (3p) m = –1, 0, +1 = 0; (4s) m = 0 n = 4 N shell = 1; (4s) m = 0 2 × 4 2 = 32 = 1; (4p) m = –1, 0, +1 = 2; (4d) m = –2, –1, 0, +1, +2 = 3; (4f) m = –3, –2, –1, 0, +1, +2, +3 Table 7 The various states that an electron can occupy 1.7. Electronic Configuration of Atoms of Elements Distribution of electrons in different orbitals of the atom of an element is called electronic configuration. 1.7.1. Principles and rules which provide guidance to write the electronic configuration. The electronic configuration of an atom is written using the guidelines of the following principles and rules. 1. Aufbau principle 2. Pauli’s exclusion principle. 3. Hund’s rule 1.7.1.1 Aufbau Principle “Electrons are filled in the increasing order of energy level” According to this principle first the electrons occupy the orbitals with lowest energy. This is decided by the sum of the principle quantum number and azimuthal quantum number. This is called (n + ) rule. Rule 1: The electrons first occupy that orbital for which (n + ) value is lowest. Rule 2: When (n + ) values for two orbitals are equal, then the electrons first occupy the orbital with lower value of n. Illustration of (n +) rule of Aufbau Principle • For 1s orbital n + = 1 + 0 = 1, and for 2s orbital n + = 2 + 0 = 2. Therefore according to rule 1, first the electrons occupy 1s orbital, then 2s orbital. • For 2p orbital n + = 2 + 1 = 3, and for 3s orbital n + = 3 + 0 = 3, the values of n + are equal. Now according to rule 2, first the electrons will occupy 2p orbitals then 3s orbital. Following in the (n + ) of Aufbau principle, the orbitals in increasing order of energy are arranged as: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p 1.7.1.2. Pauli’s Exclusion Principle FPCHE001KB 17 www.smartlearnwebtv.com Chemistry “In an atom no two electrons can have the same set of four quantum numbers”. Illustration of Pauli’s exclusion principle 1. In an atom if one electron is assigned a set of four quantum numbers n = 1, = 0, m = 0, s = 1 2 + , then other electrons cannot be assigned the same set of quantum numbers. 2. If three quantum numbers for two electrons are the same, then these electrons must have different fourth quantum number. n m s 1 First electron 1 0 0 2 1 Second electron 1 0 0 2 + − l 1.7.1.3. Hund’s Rule “Among the orbitals of same energy, electrons do not start pairing, until all these orbitals are singly occupied”. Hund’s rule is also called as the principle of minimum pairing and the principle of maximum multiplicity. 1.8. Electronic Configuration of some Individual Elements The electronic configuration of an atom can easily be written with the help of the rules discussed above. The only additional requirement is that, we should have the atomic number (Z) of the atom under consideration. The atomic number (Z) of an element gives the number of electrons present in a neutral atom of that element. After knowing the number of electrons, we can place them one by one in different available orbitals in accordance to the rules discussed above. 1.8.1. Illustration In Hydrogen atom, there is only one electron which occupies 1s orbital and the electronic state is represented by 1 1s Hydrogen 1s ↑ 2 1s Helum1s ↑↓ The third electron in Lithium would occupy 2s orbital which has the minimum energy in this shell. In the atom beryllium, the fourth electron completes the 2s orbital and thus with boron, the fifth electron must enter 2ps orbital. x y z 2 1 2p 2p 2p 1s 2s Lithium1s 2s ↑↓ ↑ 1.9. Shape of Orbitals An orbital is a three-dimensional space around the nucleus in which the probability of finding an electron is maximum. The spatial distribution of electronic cloud or electron density decides the shape of an orbital. The electronic cloud may not be have uniform density every where in an orbital. • In an orbital, there may be regions of higher probability where electron cloud is dense and there may also be regions of low probability where electron cloud is not dense. • An orbital may also contain one or more points or planes where the probability of finding electron is zero. At such points or planes, electron density is found to be equal to zero. Such FPCHE001KB 18 www.smartlearnwebtv.com Different Chemistry points are called nodal points and such planes are termed as nodal planes. The shapes of some orbitals are discussed below. 1.9.1. s orbital 1. These orbitals are spherically symmetric around the nucleus, i.e., probability of finding the electron at a particular distance from the nucleus is the same in all directions. 2. 1s orbital does not contain any node and is the smallest of all subsequent s orbitals. The size of an s orbital increases with increase in the value of n. 3. 2s orbital is larger in size as compared to 1s orbital possesses a node. The 3s orbital is still larger in size and contains two nodes. The shapes of 1s, 2s and 3s orbitals are shown in the given figure. 1s Orbital 2s Orbital 3s Orbital Figure15 Shapes of 1s, 2s and 3s orbitals Nodes are the region in which probability of finding electrons (Ψ 2 ) is zero. 1.9.2. p orbitals 1. p-orbital has three orientations i.e., probability of finding p-electron is along mutually perpendicular X, Y and Z axis. These orbitals are thus named as px, py and pz orbital. 2. In px orbital, the electron density is distributed along X-axis while in py and pz orbitals, the electron density distribution are along Y and Z axes respectively. 3. Each p orbital is dumb bell shaped and consists of two lobes of electron cloud which extend outwards and away from the nucleus along the axial line. 4. A nodal plane exists between the two lobes. Along this plane, the probability of finding electron (Ψ 2 ) is zero and consequently the electron density is also zero. 5. In each p orbital, the point at which the two lobes meet together is a nodal point. It is the point from which the nodal plane passes. The shapes of 2px, 2py and 2pz orbitals are shown below. Figure 16. Shapes of 2px, 2py and 2pz orbitals Example: An atomic orbital has = 3. What are the possible values of m? Solution: When = 3, the possible values of m are –3, –2, –1, 0, +1, +2, and +3 1.10. Atomic models of some atoms FPCHE001KB 19 www.smartlearnwebtv.com Chemistry Figure 17. atomic models of some atoms 1.10.1. Valence electron The outermost shell of an atom is called valence shell. The number of electrons in the valence shell are called valence electrons. Electronic configuration not only gives the arrangement of electrons outside the nucleus but also gives the number of valence electrons. This number decides the chemical reactivity. Atoms combine with one another to attain eight electrons in the valence shell which is the noble gas configuration. This is achieved by lending, borrowing or sharing the valence electrons among the atoms. The number of valence electrons of an atom is called the valency (or maximum valency) of the element. Examples (i) Oxygen has the electronic configuration 1s 2 , 2s 2 , 2p 4 . It has 6 electrons in its valence shell. It can borrow 2 electrons to make it 8 so that it gets the configuration of neon. Hence the valency of oxygen is 2. (ii) Sulfur which has the configuration similar to oxygen as 1s 2 , 2s 2 2p 6 , 3s 2 3p 4 also has 6 electrons in its valence shell. Its valency is 2 and maximum valency is 6. (iii) Potassium has the configuration 1s 2 , 2s 2 2p 6 , 3s 2 3p 6 4s 1 . By losing one electron from its fourth shell it gets an outermost shell of 8 electrons and gets the configuration of Argon. Hence its valency is 1. (iv) Helium has 1s 2 which is the maximum number of electrons for the first shell. It is considered stable electronic configuration. Hence it has no reactivity. (v) Neon and argon have 8 electrons in their valence shells. Hence these are also chemically inert. 1.11.1. Atomic Number We know that protons are present in the nucleus of an atom. It is the number of protons of an atom, which determines its atomic number. It is denoted by ‘Z’. All atoms of an element have the same atomic number, Z. In fact, elements are defined by the number of protons they possess. For hydrogen, Z = 1, because in hydrogen atom, only one proton is present in the nucleus. Similarly, for carbon, Z = 6. Therefore, the atomic number is defined as the total number of protons present in the nucleus of an atom. 1.11.2. Mass Number Mass of an atom is practically due to protons and neutrons alone. These are present in the nucleus of an atom. Hence protons and neutrons are also called nucleons. Therefore, the mass of an atom resides in its nucleus. For example, mass of carbon is 12 u because it has 6 protons and 6 neutrons, 6 u +6 u =12 u. Similarly, the mass of aluminium is 27 u (13 protons +14 neutrons). The mass number is defined as the sum of the total number of protons and neutrons FPCHE001KB 20 www.smartlearnwebtv.com Chemistry present in the nucleus of an atom. In the notation for an atom, the atomic number, mass number and symbol of the element are to be written as MassNumber Atomic Number Symbol of element For example, nitrogen is written as 14 7 N 1.11.3. Atomic number and mass number In 1913, Mosley devised an experiment to determine the positive charge on the nucleus. He stated, based on his experiments, that atoms of different element have different and characteristic positive charges. He called this atomic number. Since protons are the positively charged particles in the nucleus, atomic number is the number of protons in the nucleus of an atom. Since the atoms are electrically neutral atoms have electrons equal to the number of protons. 1.11.4. Atomic number Z = Number of protons = Number of electrons. Since the electrons have negligible mass, the mass of an atom is due to protons and neutrons. Protons and neutrons together are called nucleons. The mass number of an atom is the sum of the number of protons and number of neutrons. 1.11.5. Mass number A = Number of protons + Number of neutrons Mass number is also called nucleon number. So we can write, in an atom Number of protons = Z Number of electrons = Z Number of neutrons = A – Z Problem Find out the number of protons, neutrons and electrons of a) Sodium b) Argon c) Fluorine Solution a) Sodium has atomic number (Z) 11 and mass number 23. Number of protons = Z = 11 Number of electrons = Z = 11 Number of neutrons = A-Z = 23 – 11 = 12 b) Argon has atomic number 18 and mass number 40 Number of protons = Z = 18 Number of electrons = Z = 18 Number of neutrons A-Z = 40-18 = 22. c) Fluorine has atomic number 9 and mass number 19 Number of protons = Z = 9 Number of electrons Z= 9 Number of neutrons =A-Z =19-9 = 10 1.11.6. Representation of an atom An atom is represented with its atomic number and mass number by writing its symbol showing the atomic number below and mass number above. Example 8O 16 , 11Na 23 , 20Ca 40 1.12.1. Isotopes FPCHE001KB 21 www.smartlearnwebtv.com Distinguish between atomic number and mass number Chemistry In 1913, Soddy was the first to notice the presence two atoms of lead with same atomic number but different mass number. He coined the word isotopes for them. Isotopes are atoms of the same element having the same atomic number but different mass number Isotopes are defined as the atoms of the same element, having the same atomic number but different mass numbers.. For example, take the case of hydrogen atom, it has three atomic species namely protium ( ) 1 1 H , deuterium ( ) 2 1 H or D and tritium ( ) 3 1 H or T . x Proton electron x Proton electron Neutron x Proton electron Neutrons Figure 18. Three atoms species of hydrogen The atomic number of each one is 1, but the mass number is 1, 2 and 3 respectively. Other such examples are (i) carbon, 12 6 C and 14 6 C, (ii) chlorine, 35 17 Cl and 37 17 Cl , etc., Many elements consist of a mixture of isotopes. Each isotope of an element is a pure substance. The chemical properties of isotopes are similar but their physical properties are different. Chlorine occurs in nature in two isotopic forms, with masses 35 u and 37 u in the ratio of 3 : 1. Obviously, the question arises: What should we take as the mass of chlorine atom? Let us find out. The mass of an atom of any natural element is taken as the average mass of all the naturally occurring atoms of that element. If an element has no isotopes, then the mass of its atom would be the same as the sum of protons and neutrons in it. But if an element occurs in isotopic forms, then we have to know the percentage of each isotopic form and then the average mass is calculated. The average atomic mass of chlorine atom, on the basis of above data, will be 75 25 35 37 100 100 ] | ` × + × ] . , ] 105 37 142 35.5 u 4 4 4 | ` · + · · . , This does not mean that any one atom of chlorine has a fractional mass of 35.5 u. It means that if you take a certain amount of chlorine, it will contain both isotopes of chlorine and the average mass is 35.5 u. 1.12.2. Examples (i) Hydrogen exists as three isotopes – hydrogen, deuterium and tritium. All the three have the same atomic number (1) but mass numbers as 1,2,3 respectively. These can be represented as 1H 1 , 1H 2 or 1D 2 , 1H 3 or 1T 3 (ii) Chlorine has two isotopes with atomic number 17 and mass numbers 35 and 37. 17Cl 35 , 17Cl 37 (iii) Carbon isotopes are 6C 12 , 6C 13 and 6C 14 (iv) Oxygen isotopes are 8O 16 , 8O 17 , 8O 18 (v) Uranium isotopes are 92U 234 , 92U 235 , 92U 238 , 92U 239 FPCHE001KB 22 www.smartlearnwebtv.com Chemistry 1.12.3. Characteristics of isotopes • Isotopes have the same number of protons and electrons, but different number of neutrons. • Since the isotopes have the same number of electrons they have the same electronic configuration and possess same chemical properties. • Isotopes differ in physical properties. • Since isotopes have same electronic configuration their position in the periodic table is also the same. • All isotopes need not be radioactive . There are non- radioactive isotopes also. (eg) Deuterium. • Isotopes are identified by the instrument – mass spectrograph. Mass number which is the sum of the number of protons and number of neutrons is a whole number but atomic weight can be whole number or fractional. It is due to the presence of isotopes. The relative atomic mass or relative atomic weight is the average of the mass number of the isotopes based on their relative abundance. 1.12.4. Isobars Atoms of same mass number but different atomic numbers are called isobars. Since atomic numbers are different , isobars are atoms of different elements. Isobars have different number of protons and electrons but same number of nucleons Example: 18Ar 40 , 19K 40 , 20Ca 40 52Te 130 , 54Xe 130 , 56Ba 130 Let us consider two elements – calcium, atomic number 20 and argon, atomic number 18. The number of electrons in these atoms is different, but the mass number of both these elements is 40. That is, the total number of nucleons is the same in the atoms of this part of elements. Atoms of different elements with different atomic numbers, which have the same mass number, are known as isobars. 1.12.5. Isotones Atoms having same number of neutrons but different mass numbers are called isotones. Examples: 14Si 30 , 15P 31 , 16P 32 19K 39 , 20Ca 40 Number of neutrons in 19K 39 is 39 – 19 = 20 Number of neutrons in 20Ca 40 is 40 – 20 = 20 Since the atomic number or number of protons are different the number of nucleons (protons + neutrons) are different. 1.12.6. Applications Since the chemical properties of all the isotopes of an element are the same, normally we are not concerned about taking a mixture. But some isotopes have special properties which find them useful in various fields. Some of them are: 1. An isotope of uranium is used as a fuel in nuclear reactors. 2. An isotope of cobalt is used in the treatment of cancer. 3. An isotope of iodine is used in the treatment of goitre FPCHE001KB 23 www.smartlearnwebtv.com