STRUCTURE OF ATOM
Sub-atomic particles Discovery of Electron Electron was discovered by J J Thomson by Cathode ray discharge tube experiment. A cathode ray tube is made of glass containing two thin pieces of metal (electrodes) sealed in it. The electrical discharge through the gases could be observed only at very low pressures and at very high voltages. When a very high voltage (about 10,000 volts) is applied between the two electrodes, no electric discharge occurs at normal pressure. When the pressure of the gas inside the tube is less than 1 mm of mercury, a dark space appears near the cathode. When the pressure is reduced to 0.01 mm Hg, it fills the whole tube. When the pressure is further reduced, the electric discharge passes between the electrodes and the tube begins to glow. This is due to the striking of some invisible rays from the cathode. These rays which start from the cathode and move away from it, in straight lines are called cathode rays or cathode ray particles. These rays can be further checked by making a hole in the anode and coating the tube behind anode with phosphorescent material like zinc sulphide. When these rays strike the zinc sulphide coating, a bright spot on the coating is developed. Properties of Cathode Rays (i) The cathode rays start from cathode and move towards the anode. (ii) They are invisible, but their behaviour can be observed with the help of fluorescent or phosphorescent materials. (iii) In the absence of electrical or magnetic field, these rays travel in straight lines. (iv) In the presence of electrical or magnetic field, the cathode rays behave similar to that of negatively charged particles. From this it is clear that the cathode rays consist of negatively charged particles called electrons. (v) The characteristics of cathode rays (electrons) do not depend upon the material of electrodes and the nature of the gas present in the cathode ray tube. Charge to Mass Ratio of Electron J.J. Thomson measured the ratio of electrical charge (e) to the mass of electron (me ) by using cathode ray tube and applying electrical and magnetic field perpendicular to each other as well as to the path of electrons. The amount of deviation of the particles from their path in the presence of electrical or magnetic field depends upon: (i) the magnitude of the negative charge on the particle (greater the magnitude of the charge on the particle, greater is the deflection). (ii) the mass of the particle (lighter the particle, greater the deflection). ATOMIC STRUCTURE Page 2 (iii) the strength of the electrical or magnetic field (the deflection of electrons from its original path increases with the increase in the voltage across the electrodes, or the strength of the magnetic field). In the absence of electric or magnetic field, the cathode rays hit the screen at point B. When only electric field is applied, the electrons deviate from their path and hit the cathode ray tube at point A. Similarly, when only magnetic field is applied, electron strikes the cathode ray tube at point C. By carefully balancing the electrical and magnetic field strength, it is possible to bring back the electron beam to the point B. By carrying out accurate measurements, Thomson was able to determine the value of e/me as: e/me = 1.758 × 1011 C kg–1 Where me is the mass of the electron in kg and e is the magnitude of the charge on the electron in coulomb (C). Charge on the Electron (e) R.A. Millikan determined the charge on the electrons by a method known as ‘oil drop experiment’. He found that the charge on the electron to be – 1.6022 × 10–19 C. Mass of electron (me) The mass of the electron (me) was determined as follows: Mass of electron (me) = e = 1.6022 x 10-19 e/me 1.758 x 1011 = 9.1 ×10–31 kg Discovery of Protons E.Goldstein modified the discharge tube experiment by perforated (with small holes) cathode. After evacuating the tube and on applying high voltage, he found that some rays were emitting behind the cathode and moving in the opposite direction of cathode rays. These rays deflect to the negative plate of electric field. So they carry positive charge and were called anode rays or canal rays. Properties of Canal rays The characteristics of canal rays are: (i) They depend on the nature of gas present in the cathode ray tube. These are simply the positively charged gaseous ions. (ii) The charge to mass ratio of the particles is found to depend on the gas from which these originate. (iii) Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge. (iv) The behaviour of these particles in the magnetic or electrical field is opposite to that observed for cathode rays. The smallest and lightest positive ion was obtained from hydrogen and was called proton. Discovery of Neutrons Neutrons were discovered by Chadwick by bombarding a thin sheet of beryllium by α-particles. 4Be9 + 2He4 → 6C 12 + 0n 1 They are electrically neutral particles having a mass slightly greater than that of the protons. Characteristics of sub-atomic particles Sub atomic particle Symbol Discoverer Absolute Charge (in Coulomb) Relative charge Mass/kg Electron e J J Thomson -1.6022 x 10-19 -1 9.01x10-31 Proton p E.Goldstein +1.6022x10-19 +1 1.6726x10-27 Neutron n James Chadwick 0 0 1.675x10-27 Some important terms relating to Atomic structure Atomic Number: It is the number of protons present in the nucleus or number of electrons present outside the nucleus. It is denoted by the symbol ‘Z’. Atomic number (Z) = nuclear charge or number of protons (p) = number of electrons (e) Mass Number: It is the total number of protons and neutrons in atom. Or, it is the total number of nucleons in an atom. It is denoted by ‘A’. i.e. Mass number (A) = no. of protons (p) + no. of neutrons (n) or, A = p + n By knowing atomic number and mass number, we can calculate the number of neutrons as: n = A – Z ATOMIC STRUCTURE Page 3 If an element X has the atomic number Z and the mass number A, it is denoted as: or ZX A Isotopes, Isobars and Isotones Isotopes are atoms with same atomic number but different mass number. That is they contain same number of protons but different number of neutrons. Hydrogen has three isotopes Protium (1H 1 ), Deuterium (1H 2 or 1D 2 ) and Tritium (1H 3 or 1T 3 ). Among these Protium is the ordinary hydrogen and Tritium is the radioactive isotope of Hydrogen. The number of protons, neutrons and electrons present in the 3 types of hydrogen are: Isotope Number of protons Number of electrons Number of Neutrons Protium 1 1 0 Deuterium 1 1 1 Tritium 1 1 2 Almost all the elements have isotopes. All the isotopes of a given element have same chemical properties, but they differ in their physical properties. Isobars are atoms of different elements having same mass number but different atomic number. i.e. they have different number of protons but have equal sum of the protons and neutrons. e.g. 6C 14 and 7N 14 18Ar40 and 20Ca40 Isotones are atoms having same number of neutrons but have different atomic numbers. Some examples are: Isotones p e n 6C 14 6 6 8 7N 15 7 7 8 8O 16 8 8 8 ATOM MODELS Thomson’s Model of Atom J. J. Thomson proposed the first atom model, which is known as the plum pudding or raisin pudding or watermelon model. According to this model, an atom has a spherical shape in which the positive charge is uniformly distributed. The electrons are distributed in it, just like the seeds are distributed in a water melon or plums are distributed in a pudding. An important feature of this model is that the mass of the atom is assumed to be uniformly distributed over the atom. Also the total positive charge in an atom is equal to the total negative charge and hence the atom is electrically neutral. Rutherford’s Nuclear Model of Atom Rutherford proposed an atom model based on his α–particle scattering experiment. He bombarded a very thin gold foil with α–particles. The Experiment: A stream of high energy α–particles from a radioactive source was directed at a thin gold foil. The thin gold foil had a circular fluorescent zinc sulphide screen around it. Whenever α–particles struck the screen, a tiny flash of light was produced at that point. Observations: The important observations made by Rutherford are: 1. Most of the α– particles passed through the gold foil without any deviation. 2. A small fraction of the α–particles was deflected by small angles. 3. A very few α– particles (∼1 in 20,000) bounced back, that is, were deflected by nearly 180°. Conclusions: From the above observations, Rutherford made the following conclusions: 1. Since most of the α–particles passed through the foil without any deviation, most space in the atom is empty. 2. A few positively charged α– particles were deflected. This is because the positive charge of the atom is concentrated in a very small volume at the centre called nucleus. ATOMIC STRUCTURE Page 4 3. The volume occupied by the nucleus is negligibly small as compared to the total volume of the atom. The radius of the atom is about 10–10 m, while that of nucleus is 10–15 m. On the basis of above observations and conclusions, Rutherford proposed the nuclear model (Planetary model) of atom. According to this model: 1. All the positive charge and most of the mass of the atom were concentrated in an extremely small region called nucleus. 2. Electrons are revolving round the nucleus with a very high speed in circular paths called orbits. 3. Electrons and the nucleus are held together by electrostatic forces of attraction. Drawbacks or Limitations of Rutherford’s atom model 1. According to the electromagnetic theory of Max Planck, charged particles in accelerated motion should emit energy in the form of radiation. So an electron in orbit will emit radiation. So the energy of the electron decreases and finally it falls to the nucleus through a spiral path. But this does not happen. So Rutherford’s model cannot explain the stability of the atom. 2. He cannot explain the electronic structure of atom. That is about the distribution of electrons around the nucleus and the energies of the electrons. Wave nature of Electromagnetic Radiation James Maxwell suggested that when electrically charged particle moves under acceleration, alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the forms of waves called electromagnetic waves or electromagnetic radiation (emr). These are the radiations associated with electric and magnetic fields. E.g. light. The important characteristics of these radiations are: 1. The oscillating electric and magnetic fields are perpendicular to each other and both are perpendicular to the direction of propagation of the wave. 2. The electromagnetic waves do not require a medium for propagation and can move in vacuum. 3. There are many types of electromagnetic radiations, which differ from one another in wavelength (or frequency). These constitute electromagnetic spectrum. The important electromagnetic radiations in the increasing order of wavelength are: Cosmic rays, Gamma rays, X-rays, Ultra-violet rays, Visible light, Infra red rays, Microwaves, Radio waves. 4. All electromagnetic radiation travel through vacuum with a constant speed of 3x108 m/s. Some important terms relating to electromagnetic radiations 1. Frequency (ν ): It is defined as the number of waves that pass through a given point in one second. The SI unit for frequency is hertz (Hz, s–1). 2. Wavelength (λ) : It is the distance between two adjacent crusts or two adjacent troughs. Its unit is m or cm. 3. Wave number (ῡ ): It is defined as the number of wavelengths per unit length. It is the reciprocal of wavelength. Its unit is m–1 or cm-1 . The frequency (ν ), speed of light (c) and the wave length (λ)are related to each other as: c = ν λ Particle Nature of Electromagnetic Radiation: Planck’s Quantum Theory Some of the experimental phenomenon like diffraction and interference can be explained by the wave nature of the electromagnetic radiation. But some phenomena like black body radiation, photoelectric effect, variation of heat capacity of solids with temperature, line spectra of atoms etc. could not be explained by wave nature of emr. Black body radiation An ideal body which emits and absorbs all frequencies of radiations is called a black body and the radiation emitted by such a body is called black body radiation. The frequency distribution of radiation emitted from a black body depends only on its temperature. At a given temperature, intensity of radiation emitted increases with decrease of wavelength, reaches a maximum value at a given wavelength and then starts decreasing. The phenomenon of black body radiation was first explained by Max Planck by his Quantum theory. According to this theory: 1. Atoms and molecules could emit (or absorb) energy not in a continuous manner, but discontinuously in small packets of energy called quanta or photons. 2. The energy (E ) of a quantum of radiation is proportional to its frequency (ν ) It is expressed by the equation, E = hν Where ‘h’ is known as Planck’s constant and its value is 6.626×10–34 J s. ATOMIC STRUCTURE Page 5 Photoelectric effect It is the phenomenon of ejection of electrons by certain metals (like potassium, rubidium, caesium etc.) when light of suitable frequency incident on them. The electrons ejected are called photoelectrons. This phenomenon was first observed by H.Hertz. The important characteristics of photoelectric effect are: 1. The electrons are ejected from the metal surface as soon as the beam of light strikes the surface. i.e., there is no time lag between the striking of light beam and the ejection of electrons from the metal surface. 2. The number of electrons ejected is proportional to the intensity or brightness of light. 3. For each metal, there is a minimum frequency (known as threshold frequency [ν0]) below which photoelectric effect is not observed. 4. The kinetic energy of the ejected electrons is directly proportional to the frequency of the incident light. Explanation of photoelectric effect A satisfactory explanation to photoelectric effect was first given by Albert Einstein using Planck’s Quantum theory. According to him, when a photon of sufficient energy strikes the metal surface, it transfers its energy to the electron of the atom of the metal instantaneously and the electron is ejected without any time lag. A part of the energy is used to eject the electron from the metal surface (i.e. to overcome the attractive force of the nucleus [work function, hν0]) and the other part is given to the ejected electron in the form of kinetic energy. Greater the energy possessed by the photon, greater will be transfer of energy to the electron and greater the kinetic energy of the ejected electron. Since the striking photon has energy equal to hν and the minimum energy required to eject the electron is hν0 (also called work function, W0) then the difference in energy (hν – hν0) is transferred as the kinetic energy of the photoelectron. Following the law of conservation of energy principle, the kinetic energy of the ejected electron is given by K.E = hν - hν0 Or, hν = hν0 + ½ mev 2 Where me is the mass of the electron and v is the velocity of the ejected electron. A more intense beam of light contains larger number of photons, so the number of electrons ejected is also larger. Dual Behaviour of Electromagnetic Radiation Electromagnetic radiations possess both particle and wave nature. This is known as dual nature of Electromagnetic Radiation. Atomic spectrum When a ray of white light is passed through a prism, we get a series of coloured bands called spectrum. This spectrum is called continuous spectrum, because here violet merges into blue, blue into green and so on. Similarly, when electromagnetic radiation interacts with matter, atoms and ATOMIC STRUCTURE Page 6 Johannes Rydberg proposed an equation for finding the wave number of the different lines in Hydrogen spectrum. The expression is: ῡ = 1/ λ = 109677 (1/n1 2 -1/n2 2 ) cm-1 Where n1 = 1, 2, 3,….. and n2 = n1 + 1, n1 + 2, …… The different spectral lines, their n1 and n2 values and their spectral region are: Series Spectral region n1 n2 Lyman Ultra violet 1 2,3,4…. Balmer Visible 2 3,4,5…. Paschen Infra red 3 4,5,6….. Brackett Infra red 4 5,6,7…. Pfund Infra red 5 6,7,8….. BOHR’S MODEL FOR HYDROGEN ATOM The general features of the structure of hydrogen atom and its spectrum was first explained by Niels Bohr. The important postulates of his theory are: 1. The electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy. These paths are called orbits or stationary states or allowed energy states. These energy levels are numbered as 1,2,3 etc or designated as K, L, M, N, etc. These numbers are known as Principal quantum numbers. 2. The energy of an electron in the orbit does not change with time. However, when an electron absorbs energy, it will move away from the nucleus (i.e. to a higher energy level) and when it loses energy, it will move towards the nucleus (i.e. to a lower energy level). 3. The radius of orbits can be given by the equation: rn = a0 n 2 where a0 = 52.9 pm. Thus the radius of the first stationary state is 52.9 pm (called the Bohr radius). As n increases, the value of r will increase. 4. The energy of electron in an orbit is given by the expression: En = -RH (1/n2 ), where n = 1,2,3…… and RH is a constant called Rydberg constant. Its value is 2.18x10-18 J. The energy of the lowest state (the ground state) is given by E1 = – 2.18×10–18J. As the value of n increases, the energy of the electron also increases. 5. The frequency of radiation absorbed or emitted when transition occurs between two stationary states that differ in energy by ΔE, is given by: ν = ΔE = E2 – E1 h h Where E1 and E2 are the energies of lower and higher energy levels respectively. This expression is commonly known as Bohr’s frequency rule. 6. The angular momentum of an electron is an integral multiple of h/2π. i.e. mevr = nh 2π Where me is the mass of electron, v is the velocity of electron and r is the radius of Bohr orbit. n = 1,2,3......... Thus an electron can move only in those orbits whose angular momentum is an integral multiple of h/2π. So only certain fixed orbits are allowed. Significance of negative energy of electron When the electron is free from the influence of nucleus, its energy is taken as zero. In this situation, the electron is at the orbit with n = ∞. When the electron is aZracted by the nucleus and is present in orbit n, the energy is emiZed and its energy is lowered. That is the reason for the presence of negative sign in equation. Explanation of Line Spectrum of Hydrogen According to Bohr atom model, radiation is absorbed if the electron moves from lower energy to higher energy level and radiation is emitted if the electron moves from higher orbit to lower orbit. The energy gap between the two orbits is given by equation: ΔE = E2 – E1 But E1 = -RH and E2 = -RH n1 2 n2 2 Therefore, ΔE = RH [1/n1 2 – 1/n2 2 ] = 2.18 x 10-18 [1/n1 2 – 1/n2 2 ] The frequency associated with the absorption and emission of the photon can be given as ATOMIC STRUCTURE Page 7 ν = ΔE = RH 1 – 1 h h n1 2 n2 2 = 2.18 x 10-18 1 - 1 6.626 x 10-34 n1 2 n2 2 = 3.29 x 1015 [1/n1 2 – 1/n2 2 ] The wave number (ῡ) = 1/ λ = ν = RH 1 - 1 C hc n1 2 n2 2 = 3.29 x 1015 1 - 1 3 x 108 n1 2 n2 2 = 1.09677 x 107 1 - 1 m-1 = 109677 1 - 1 cm-1 n1 2 n2 2 n1 2 n2 2 In case of absorption spectrum, n2 > n1 and the term in the bracket is positive and energy is absorbed. On the other hand, in case of emission spectrum n1 > n2, ΔE is negative and energy is released. Limitations of Bohr Atom Model Bohr atom model could explain the stability and line spectra of hydrogen atom and hydrogen like ions (e.g. He+ , Li2+, Be3+ etc). But it has the following limitations: 1. It could not explain the fine spectrum of hydrogen atom. 2. It could not explain the spectrum of atoms other than hydrogen. 3. It was unable to explain the splitting of spectral lines in the presence of electric field (Stark effect) and in magnetic field (Zeeman effect). 4. It could not explain the ability of atoms to form molecules by chemical bonds. 5. It did not consider the wave character of matter and Heisenberg’s uncertainty principle. Dual Behaviour of Matter – de Broglie’s equation de Broglie proposed that like radiation, matter also exhibit dual behaviour i.e., both particle and wave like properties. This means that electrons should also have momentum as well as wavelength. He gave the following relation between wavelength (λ) and momentum (p) of a material particle. λ = h = h mv p Where m is the mass of the particle, v is the velocity and p is the momentum. The above equation is known as de Broglie’s equation. Just like electromagnetic radiations, an electron beam also undergoes diffraction. This is an evidence for the wave nature of electrons. An electron microscope works on the principle of wave nature of electron. According to de Broglie, every moving object has a wave character. The wavelengths associated with ordinary objects are so short (because of their large masses) that their wave properties cannot be detected. The wavelengths associated with electrons and other subatomic particles (with very small mass) can be detected experimentally. Heisenberg’s Uncertainty Principle Werner Heisenberg proposed the uncertainty principle which is the consequence of dual behaviour of matter and radiation. It states that “it is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of a moving microscopic particle like electron”. Mathematically, it can be given as in equation: Δx. Δp ≥ h 4π Or, Δx.mΔv ≥ h 4π Or, Δx.Δv ≥ h 4πm Where Δx is the uncertainty in position and Δp (or, Δv) is the uncertainty in momentum (or velocity) of the particle. ATOMIC STRUCTURE Page 8 If the position of the electron is known with high degree of accuracy (Δx is small), ATOMIC STRUCTURE Page 9 For example, when n = 1, value of Ɩ is only 0. For n = 2, the possible value of Ɩ can be 0 and 1. For n = 3, the possible Ɩ values are 0,1 and 2. Ɩ= 0 represents s orbital, Ɩ = 1 represents p orbital, Ɩ = 2 represents d orbital and Ɩ = 3 represents f orbital The number of sub shells in a principal shell is equal to the value of n. For example, When n = 1, Ɩ= 0. i.e. K shell contains only one sub shell - s sub shell when n = 2, Ɩ = 0 and1. i.e. L shell contains two sub shells - s and p sub shells when n = 3, Ɩ = 0, 1 and 2. i.e. M shell contains three sub shells – s, p and d sub shells when n = 4, Ɩ = 0, 1, 2 and 3. i.e. N shell contains four sub shells – s, p,d and f sub shells 3. Magnetic Quantum Number (m or mƖ) It gives information about the orientation of orbitals in space. For a given ‘Ɩ’ value, there are 2Ɩ+1 possible values for m and these values are given by : m = – Ɩ to 0 to + Ɩ Thus for Ɩ = 0, mƖ = 0 [2(0)+1 = 1].. i.e. s sub shell contains only one orbital called s orbital. For Ɩ = 1, mƖ = –1, 0 and +1 [2(1)+1 = 3]. i.e. p subshell contains three orbitals called p orbitals (px, py and pz). For Ɩ = 2, mƖ = –2, –1, 0, +1 and +2, [2(2)+1 = 5]. i.e. d subshell contains five orbitals called d orbitals (dxy, dxz, dyz, dx 2 - y 2 and dz 2 ) 4. Spin Quantum Number (s or ms) It is the only experimental Quantum number and it gives the spin orientation of electrons. This spin may be either clockwise or anticlockwise. So the values for s may be +½ or -½. +½ represents clock-wise spin and-½ represents anticlockwise spin. Shapes of orbitals 1. s-orbitals For s-orbitals, Ɩ = 0 and hence mƖ = 0. So there is only one possible orientation for s orbitals. They are spherically symmetrical. The plots of probability density (ψ2 ) against distance from the nucleus (r) for 1s and 2s atomic orbitals are as follows: For 1s orbital the probability density is maximum at the nucleus and it decreases with increase in r. But for 2s orbital the probability density first decreases sharply to zero and again starts increasing. After reaching a small maximum it decreases again and approaches zero as the value of r increases. The region where the probability density (ψ2 ) reduces to zero is called nodal surface or node. For 1s orbital, there is no node, for 2s orbital there is only one node, for 3s orbital there are 2 nodes and so on. In general, for an ns-orbital there are (n – 1) nodes. All the s-orbitals are spherically symmetrical and their size increases with increase in n. The boundary surface diagrams for 1s, 2sand 3s orbitals are as follows: 2. p-orbitals For p-orbitals, Ɩ = 1 and mƖ = -1, 0, +1. i.e., there are three possible orientations for p orbitals. So there are 3 types of porbitals – px, py and pz . Each p orbital consists of two lobes. The probability density function is zero on the plane where the two lobes touch each other. ATOMIC STRUCTURE The size, shape and energy of the three orbitals are identical. They differ only in the orbital, the lobes are along the x-axis, for p have dumb-bell shape. The number of radial nodes for p-orbitals are given by 4p orbital and so on. Besides the radial nodes passing through the nucleus (origin). For example, in angular nodes and number of angular nodes is Number of radial nodes = n - Ɩ – 1 Number of angular nodes = Ɩ Total number of nodes = n-1 The boundary surface diagrams for three 2p orbitals are as follows: 3. d-orbitals For d-orbitals, Ɩ = 2 and mƖ = -2, -1, 0, +1 and +2. i.e., there are five possible orientations for d orbitals. So there are 5 types of d-orbitals. They are dxy, dxz, dyz, d of the fifth one, dz 2 , is dumb-bell having a circular collar in the xy orbitals the number of radial nodes is 2 and the total number of nodes is n as follows: 4. f-orbitals For f-orbitals, Ɩ = 3 and mƖ = -3, -2, there are 7 types of f-orbitals. They Rules for Filling of electrons in various orbitals The filling of electrons into the orbitals of different atoms takes place according to the 3 rules Pauli’s exclusion principle and the Hund’s rule of maximum multiplicity. 1. Aufbau principle: The German word aufbau means ‘build up’. The building up of orbitals m electrons. It states that the orbitals are filled in order of their increasing energies the lowest energy orbital and then to higher energy orbitals. This rule has two sub rules: a) The various orbitals are filled in the increasing order of their (n+ b) If two orbitals have the same (n+ The increasing order of orbitals is as 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5 The size, shape and energy of the three orbitals are identical. They differ only in the orientation of the lobes axis, for py, they are along the y-axis and for pz , they are along the z orbitals are given by (n –2), that is number of radial node is 1 for 3p orbital, ial nodes, the probability density functions for the np orbitals nucleus (origin). For example, in the case of pz orbital, xy-plane is a nodal plane number of angular nodes is given by ‘Ɩ’. The boundary surface diagrams for three 2p orbitals are as follows: 1, 0, +1 and +2. i.e., there are five possible orientations for d orbitals. So there are 5 , dx 2 -y 2 and dz 2 . The shapes of the first four d-orbitals are double dumb bell having a circular collar in the xy-plane. The five d-orbitals have equivalent energies. For d orbitals the number of radial nodes is 2 and the total number of nodes is n-2. Boundary surface diagrams for d 2, -1, 0, +1, +2 and +3. i.e., there are seven possible orientations for f orbitals. So They have diffused shapes. electrons in various orbitals into the orbitals of different atoms takes place according to the 3 rules Pauli’s exclusion principle and the Hund’s rule of maximum multiplicity. The German word aufbau means ‘build up’. The building up of orbitals means the filling up of orbitals with the orbitals are filled in order of their increasing energies. In other words, electrons first occupy the lowest energy orbital and then to higher energy orbitals. This rule has two sub rules: The various orbitals are filled in the increasing order of their (n+Ɩ) value. If two orbitals have the same (n+Ɩ) values, the orbital with the lower n value is filled first. The increasing order of orbitals is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s... Page 10 orientation of the lobes. For px , they are along the z-axis. All the p-orbitals , that is number of radial node is 1 for 3p orbital, two for density functions for the np orbitals are zero at the plane, plane is a nodal plane. These are called 1, 0, +1 and +2. i.e., there are five possible orientations for d orbitals. So there are 5 orbitals are double dumb-bell and that orbitals have equivalent energies. For d2. Boundary surface diagrams for d-orbitals are 1, 0, +1, +2 and +3. i.e., there are seven possible orientations for f orbitals. So into the orbitals of different atoms takes place according to the 3 rules - aufbau principle, eans the filling up of orbitals with . In other words, electrons first occupy Ɩ) values, the orbital with the lower n value is filled first. ATOMIC STRUCTURE Page 11 2. Pauli’s Exclusion Principle It states that no two electrons in an atom can have the same set of four quantum numbers. i.e. an orbital can accommodate a maximum of only 2 electrons with opposite spin. If 2 electrons have same values for n, Ɩ and m, they should have different values for s. i.e. if s = +½ for the first electron, it should be -½ for the second electron. 3. Hund’s rule of maximum multiplicity It states that electron pairing takes place only after partially filling all the degenerate orbitals. Orbitals having same energies are called degenerate orbitals. For example the electronic configuration of N is 1s2 2s2 2px1 py1 pz1 and not 1s2 2s2 2px2 py1 . Electronic Configuration of Atoms The distribution of electrons into various orbitals of an atom is called its electronic configuration. The electronic configuration of different atoms can be represented in two ways. (i) sa p b d c ...... notation (ii) Orbital diagram The electrons in the completely filled shells are known as core electrons and the electrons in the outer most shell are called valence electrons. Stability of Completely Filled and Half Filled Subshells For atoms having half filled or completely filled electronic configurations have extra stability compared to other atoms. This is due to their symmetrical distribution of electrons and greater exchange energy. For example, the electronic configuration of Cr is [Ar] 3d5 4s1 and not 3d4 4s2 . This is because d5 represents a half filled configuration and has extra stability. Similarly for Cu the electronic configuration is [Ar] 3d104s1 and not 3d9 4s2 .
Sub-atomic particles Discovery of Electron Electron was discovered by J J Thomson by Cathode ray discharge tube experiment. A cathode ray tube is made of glass containing two thin pieces of metal (electrodes) sealed in it. The electrical discharge through the gases could be observed only at very low pressures and at very high voltages. When a very high voltage (about 10,000 volts) is applied between the two electrodes, no electric discharge occurs at normal pressure. When the pressure of the gas inside the tube is less than 1 mm of mercury, a dark space appears near the cathode. When the pressure is reduced to 0.01 mm Hg, it fills the whole tube. When the pressure is further reduced, the electric discharge passes between the electrodes and the tube begins to glow. This is due to the striking of some invisible rays from the cathode. These rays which start from the cathode and move away from it, in straight lines are called cathode rays or cathode ray particles. These rays can be further checked by making a hole in the anode and coating the tube behind anode with phosphorescent material like zinc sulphide. When these rays strike the zinc sulphide coating, a bright spot on the coating is developed. Properties of Cathode Rays (i) The cathode rays start from cathode and move towards the anode. (ii) They are invisible, but their behaviour can be observed with the help of fluorescent or phosphorescent materials. (iii) In the absence of electrical or magnetic field, these rays travel in straight lines. (iv) In the presence of electrical or magnetic field, the cathode rays behave similar to that of negatively charged particles. From this it is clear that the cathode rays consist of negatively charged particles called electrons. (v) The characteristics of cathode rays (electrons) do not depend upon the material of electrodes and the nature of the gas present in the cathode ray tube. Charge to Mass Ratio of Electron J.J. Thomson measured the ratio of electrical charge (e) to the mass of electron (me ) by using cathode ray tube and applying electrical and magnetic field perpendicular to each other as well as to the path of electrons. The amount of deviation of the particles from their path in the presence of electrical or magnetic field depends upon: (i) the magnitude of the negative charge on the particle (greater the magnitude of the charge on the particle, greater is the deflection). (ii) the mass of the particle (lighter the particle, greater the deflection). ATOMIC STRUCTURE Page 2 (iii) the strength of the electrical or magnetic field (the deflection of electrons from its original path increases with the increase in the voltage across the electrodes, or the strength of the magnetic field). In the absence of electric or magnetic field, the cathode rays hit the screen at point B. When only electric field is applied, the electrons deviate from their path and hit the cathode ray tube at point A. Similarly, when only magnetic field is applied, electron strikes the cathode ray tube at point C. By carefully balancing the electrical and magnetic field strength, it is possible to bring back the electron beam to the point B. By carrying out accurate measurements, Thomson was able to determine the value of e/me as: e/me = 1.758 × 1011 C kg–1 Where me is the mass of the electron in kg and e is the magnitude of the charge on the electron in coulomb (C). Charge on the Electron (e) R.A. Millikan determined the charge on the electrons by a method known as ‘oil drop experiment’. He found that the charge on the electron to be – 1.6022 × 10–19 C. Mass of electron (me) The mass of the electron (me) was determined as follows: Mass of electron (me) = e = 1.6022 x 10-19 e/me 1.758 x 1011 = 9.1 ×10–31 kg Discovery of Protons E.Goldstein modified the discharge tube experiment by perforated (with small holes) cathode. After evacuating the tube and on applying high voltage, he found that some rays were emitting behind the cathode and moving in the opposite direction of cathode rays. These rays deflect to the negative plate of electric field. So they carry positive charge and were called anode rays or canal rays. Properties of Canal rays The characteristics of canal rays are: (i) They depend on the nature of gas present in the cathode ray tube. These are simply the positively charged gaseous ions. (ii) The charge to mass ratio of the particles is found to depend on the gas from which these originate. (iii) Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge. (iv) The behaviour of these particles in the magnetic or electrical field is opposite to that observed for cathode rays. The smallest and lightest positive ion was obtained from hydrogen and was called proton. Discovery of Neutrons Neutrons were discovered by Chadwick by bombarding a thin sheet of beryllium by α-particles. 4Be9 + 2He4 → 6C 12 + 0n 1 They are electrically neutral particles having a mass slightly greater than that of the protons. Characteristics of sub-atomic particles Sub atomic particle Symbol Discoverer Absolute Charge (in Coulomb) Relative charge Mass/kg Electron e J J Thomson -1.6022 x 10-19 -1 9.01x10-31 Proton p E.Goldstein +1.6022x10-19 +1 1.6726x10-27 Neutron n James Chadwick 0 0 1.675x10-27 Some important terms relating to Atomic structure Atomic Number: It is the number of protons present in the nucleus or number of electrons present outside the nucleus. It is denoted by the symbol ‘Z’. Atomic number (Z) = nuclear charge or number of protons (p) = number of electrons (e) Mass Number: It is the total number of protons and neutrons in atom. Or, it is the total number of nucleons in an atom. It is denoted by ‘A’. i.e. Mass number (A) = no. of protons (p) + no. of neutrons (n) or, A = p + n By knowing atomic number and mass number, we can calculate the number of neutrons as: n = A – Z ATOMIC STRUCTURE Page 3 If an element X has the atomic number Z and the mass number A, it is denoted as: or ZX A Isotopes, Isobars and Isotones Isotopes are atoms with same atomic number but different mass number. That is they contain same number of protons but different number of neutrons. Hydrogen has three isotopes Protium (1H 1 ), Deuterium (1H 2 or 1D 2 ) and Tritium (1H 3 or 1T 3 ). Among these Protium is the ordinary hydrogen and Tritium is the radioactive isotope of Hydrogen. The number of protons, neutrons and electrons present in the 3 types of hydrogen are: Isotope Number of protons Number of electrons Number of Neutrons Protium 1 1 0 Deuterium 1 1 1 Tritium 1 1 2 Almost all the elements have isotopes. All the isotopes of a given element have same chemical properties, but they differ in their physical properties. Isobars are atoms of different elements having same mass number but different atomic number. i.e. they have different number of protons but have equal sum of the protons and neutrons. e.g. 6C 14 and 7N 14 18Ar40 and 20Ca40 Isotones are atoms having same number of neutrons but have different atomic numbers. Some examples are: Isotones p e n 6C 14 6 6 8 7N 15 7 7 8 8O 16 8 8 8 ATOM MODELS Thomson’s Model of Atom J. J. Thomson proposed the first atom model, which is known as the plum pudding or raisin pudding or watermelon model. According to this model, an atom has a spherical shape in which the positive charge is uniformly distributed. The electrons are distributed in it, just like the seeds are distributed in a water melon or plums are distributed in a pudding. An important feature of this model is that the mass of the atom is assumed to be uniformly distributed over the atom. Also the total positive charge in an atom is equal to the total negative charge and hence the atom is electrically neutral. Rutherford’s Nuclear Model of Atom Rutherford proposed an atom model based on his α–particle scattering experiment. He bombarded a very thin gold foil with α–particles. The Experiment: A stream of high energy α–particles from a radioactive source was directed at a thin gold foil. The thin gold foil had a circular fluorescent zinc sulphide screen around it. Whenever α–particles struck the screen, a tiny flash of light was produced at that point. Observations: The important observations made by Rutherford are: 1. Most of the α– particles passed through the gold foil without any deviation. 2. A small fraction of the α–particles was deflected by small angles. 3. A very few α– particles (∼1 in 20,000) bounced back, that is, were deflected by nearly 180°. Conclusions: From the above observations, Rutherford made the following conclusions: 1. Since most of the α–particles passed through the foil without any deviation, most space in the atom is empty. 2. A few positively charged α– particles were deflected. This is because the positive charge of the atom is concentrated in a very small volume at the centre called nucleus. ATOMIC STRUCTURE Page 4 3. The volume occupied by the nucleus is negligibly small as compared to the total volume of the atom. The radius of the atom is about 10–10 m, while that of nucleus is 10–15 m. On the basis of above observations and conclusions, Rutherford proposed the nuclear model (Planetary model) of atom. According to this model: 1. All the positive charge and most of the mass of the atom were concentrated in an extremely small region called nucleus. 2. Electrons are revolving round the nucleus with a very high speed in circular paths called orbits. 3. Electrons and the nucleus are held together by electrostatic forces of attraction. Drawbacks or Limitations of Rutherford’s atom model 1. According to the electromagnetic theory of Max Planck, charged particles in accelerated motion should emit energy in the form of radiation. So an electron in orbit will emit radiation. So the energy of the electron decreases and finally it falls to the nucleus through a spiral path. But this does not happen. So Rutherford’s model cannot explain the stability of the atom. 2. He cannot explain the electronic structure of atom. That is about the distribution of electrons around the nucleus and the energies of the electrons. Wave nature of Electromagnetic Radiation James Maxwell suggested that when electrically charged particle moves under acceleration, alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the forms of waves called electromagnetic waves or electromagnetic radiation (emr). These are the radiations associated with electric and magnetic fields. E.g. light. The important characteristics of these radiations are: 1. The oscillating electric and magnetic fields are perpendicular to each other and both are perpendicular to the direction of propagation of the wave. 2. The electromagnetic waves do not require a medium for propagation and can move in vacuum. 3. There are many types of electromagnetic radiations, which differ from one another in wavelength (or frequency). These constitute electromagnetic spectrum. The important electromagnetic radiations in the increasing order of wavelength are: Cosmic rays, Gamma rays, X-rays, Ultra-violet rays, Visible light, Infra red rays, Microwaves, Radio waves. 4. All electromagnetic radiation travel through vacuum with a constant speed of 3x108 m/s. Some important terms relating to electromagnetic radiations 1. Frequency (ν ): It is defined as the number of waves that pass through a given point in one second. The SI unit for frequency is hertz (Hz, s–1). 2. Wavelength (λ) : It is the distance between two adjacent crusts or two adjacent troughs. Its unit is m or cm. 3. Wave number (ῡ ): It is defined as the number of wavelengths per unit length. It is the reciprocal of wavelength. Its unit is m–1 or cm-1 . The frequency (ν ), speed of light (c) and the wave length (λ)are related to each other as: c = ν λ Particle Nature of Electromagnetic Radiation: Planck’s Quantum Theory Some of the experimental phenomenon like diffraction and interference can be explained by the wave nature of the electromagnetic radiation. But some phenomena like black body radiation, photoelectric effect, variation of heat capacity of solids with temperature, line spectra of atoms etc. could not be explained by wave nature of emr. Black body radiation An ideal body which emits and absorbs all frequencies of radiations is called a black body and the radiation emitted by such a body is called black body radiation. The frequency distribution of radiation emitted from a black body depends only on its temperature. At a given temperature, intensity of radiation emitted increases with decrease of wavelength, reaches a maximum value at a given wavelength and then starts decreasing. The phenomenon of black body radiation was first explained by Max Planck by his Quantum theory. According to this theory: 1. Atoms and molecules could emit (or absorb) energy not in a continuous manner, but discontinuously in small packets of energy called quanta or photons. 2. The energy (E ) of a quantum of radiation is proportional to its frequency (ν ) It is expressed by the equation, E = hν Where ‘h’ is known as Planck’s constant and its value is 6.626×10–34 J s. ATOMIC STRUCTURE Page 5 Photoelectric effect It is the phenomenon of ejection of electrons by certain metals (like potassium, rubidium, caesium etc.) when light of suitable frequency incident on them. The electrons ejected are called photoelectrons. This phenomenon was first observed by H.Hertz. The important characteristics of photoelectric effect are: 1. The electrons are ejected from the metal surface as soon as the beam of light strikes the surface. i.e., there is no time lag between the striking of light beam and the ejection of electrons from the metal surface. 2. The number of electrons ejected is proportional to the intensity or brightness of light. 3. For each metal, there is a minimum frequency (known as threshold frequency [ν0]) below which photoelectric effect is not observed. 4. The kinetic energy of the ejected electrons is directly proportional to the frequency of the incident light. Explanation of photoelectric effect A satisfactory explanation to photoelectric effect was first given by Albert Einstein using Planck’s Quantum theory. According to him, when a photon of sufficient energy strikes the metal surface, it transfers its energy to the electron of the atom of the metal instantaneously and the electron is ejected without any time lag. A part of the energy is used to eject the electron from the metal surface (i.e. to overcome the attractive force of the nucleus [work function, hν0]) and the other part is given to the ejected electron in the form of kinetic energy. Greater the energy possessed by the photon, greater will be transfer of energy to the electron and greater the kinetic energy of the ejected electron. Since the striking photon has energy equal to hν and the minimum energy required to eject the electron is hν0 (also called work function, W0) then the difference in energy (hν – hν0) is transferred as the kinetic energy of the photoelectron. Following the law of conservation of energy principle, the kinetic energy of the ejected electron is given by K.E = hν - hν0 Or, hν = hν0 + ½ mev 2 Where me is the mass of the electron and v is the velocity of the ejected electron. A more intense beam of light contains larger number of photons, so the number of electrons ejected is also larger. Dual Behaviour of Electromagnetic Radiation Electromagnetic radiations possess both particle and wave nature. This is known as dual nature of Electromagnetic Radiation. Atomic spectrum When a ray of white light is passed through a prism, we get a series of coloured bands called spectrum. This spectrum is called continuous spectrum, because here violet merges into blue, blue into green and so on. Similarly, when electromagnetic radiation interacts with matter, atoms and ATOMIC STRUCTURE Page 6 Johannes Rydberg proposed an equation for finding the wave number of the different lines in Hydrogen spectrum. The expression is: ῡ = 1/ λ = 109677 (1/n1 2 -1/n2 2 ) cm-1 Where n1 = 1, 2, 3,….. and n2 = n1 + 1, n1 + 2, …… The different spectral lines, their n1 and n2 values and their spectral region are: Series Spectral region n1 n2 Lyman Ultra violet 1 2,3,4…. Balmer Visible 2 3,4,5…. Paschen Infra red 3 4,5,6….. Brackett Infra red 4 5,6,7…. Pfund Infra red 5 6,7,8….. BOHR’S MODEL FOR HYDROGEN ATOM The general features of the structure of hydrogen atom and its spectrum was first explained by Niels Bohr. The important postulates of his theory are: 1. The electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy. These paths are called orbits or stationary states or allowed energy states. These energy levels are numbered as 1,2,3 etc or designated as K, L, M, N, etc. These numbers are known as Principal quantum numbers. 2. The energy of an electron in the orbit does not change with time. However, when an electron absorbs energy, it will move away from the nucleus (i.e. to a higher energy level) and when it loses energy, it will move towards the nucleus (i.e. to a lower energy level). 3. The radius of orbits can be given by the equation: rn = a0 n 2 where a0 = 52.9 pm. Thus the radius of the first stationary state is 52.9 pm (called the Bohr radius). As n increases, the value of r will increase. 4. The energy of electron in an orbit is given by the expression: En = -RH (1/n2 ), where n = 1,2,3…… and RH is a constant called Rydberg constant. Its value is 2.18x10-18 J. The energy of the lowest state (the ground state) is given by E1 = – 2.18×10–18J. As the value of n increases, the energy of the electron also increases. 5. The frequency of radiation absorbed or emitted when transition occurs between two stationary states that differ in energy by ΔE, is given by: ν = ΔE = E2 – E1 h h Where E1 and E2 are the energies of lower and higher energy levels respectively. This expression is commonly known as Bohr’s frequency rule. 6. The angular momentum of an electron is an integral multiple of h/2π. i.e. mevr = nh 2π Where me is the mass of electron, v is the velocity of electron and r is the radius of Bohr orbit. n = 1,2,3......... Thus an electron can move only in those orbits whose angular momentum is an integral multiple of h/2π. So only certain fixed orbits are allowed. Significance of negative energy of electron When the electron is free from the influence of nucleus, its energy is taken as zero. In this situation, the electron is at the orbit with n = ∞. When the electron is aZracted by the nucleus and is present in orbit n, the energy is emiZed and its energy is lowered. That is the reason for the presence of negative sign in equation. Explanation of Line Spectrum of Hydrogen According to Bohr atom model, radiation is absorbed if the electron moves from lower energy to higher energy level and radiation is emitted if the electron moves from higher orbit to lower orbit. The energy gap between the two orbits is given by equation: ΔE = E2 – E1 But E1 = -RH and E2 = -RH n1 2 n2 2 Therefore, ΔE = RH [1/n1 2 – 1/n2 2 ] = 2.18 x 10-18 [1/n1 2 – 1/n2 2 ] The frequency associated with the absorption and emission of the photon can be given as ATOMIC STRUCTURE Page 7 ν = ΔE = RH 1 – 1 h h n1 2 n2 2 = 2.18 x 10-18 1 - 1 6.626 x 10-34 n1 2 n2 2 = 3.29 x 1015 [1/n1 2 – 1/n2 2 ] The wave number (ῡ) = 1/ λ = ν = RH 1 - 1 C hc n1 2 n2 2 = 3.29 x 1015 1 - 1 3 x 108 n1 2 n2 2 = 1.09677 x 107 1 - 1 m-1 = 109677 1 - 1 cm-1 n1 2 n2 2 n1 2 n2 2 In case of absorption spectrum, n2 > n1 and the term in the bracket is positive and energy is absorbed. On the other hand, in case of emission spectrum n1 > n2, ΔE is negative and energy is released. Limitations of Bohr Atom Model Bohr atom model could explain the stability and line spectra of hydrogen atom and hydrogen like ions (e.g. He+ , Li2+, Be3+ etc). But it has the following limitations: 1. It could not explain the fine spectrum of hydrogen atom. 2. It could not explain the spectrum of atoms other than hydrogen. 3. It was unable to explain the splitting of spectral lines in the presence of electric field (Stark effect) and in magnetic field (Zeeman effect). 4. It could not explain the ability of atoms to form molecules by chemical bonds. 5. It did not consider the wave character of matter and Heisenberg’s uncertainty principle. Dual Behaviour of Matter – de Broglie’s equation de Broglie proposed that like radiation, matter also exhibit dual behaviour i.e., both particle and wave like properties. This means that electrons should also have momentum as well as wavelength. He gave the following relation between wavelength (λ) and momentum (p) of a material particle. λ = h = h mv p Where m is the mass of the particle, v is the velocity and p is the momentum. The above equation is known as de Broglie’s equation. Just like electromagnetic radiations, an electron beam also undergoes diffraction. This is an evidence for the wave nature of electrons. An electron microscope works on the principle of wave nature of electron. According to de Broglie, every moving object has a wave character. The wavelengths associated with ordinary objects are so short (because of their large masses) that their wave properties cannot be detected. The wavelengths associated with electrons and other subatomic particles (with very small mass) can be detected experimentally. Heisenberg’s Uncertainty Principle Werner Heisenberg proposed the uncertainty principle which is the consequence of dual behaviour of matter and radiation. It states that “it is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of a moving microscopic particle like electron”. Mathematically, it can be given as in equation: Δx. Δp ≥ h 4π Or, Δx.mΔv ≥ h 4π Or, Δx.Δv ≥ h 4πm Where Δx is the uncertainty in position and Δp (or, Δv) is the uncertainty in momentum (or velocity) of the particle. ATOMIC STRUCTURE Page 8 If the position of the electron is known with high degree of accuracy (Δx is small), ATOMIC STRUCTURE Page 9 For example, when n = 1, value of Ɩ is only 0. For n = 2, the possible value of Ɩ can be 0 and 1. For n = 3, the possible Ɩ values are 0,1 and 2. Ɩ= 0 represents s orbital, Ɩ = 1 represents p orbital, Ɩ = 2 represents d orbital and Ɩ = 3 represents f orbital The number of sub shells in a principal shell is equal to the value of n. For example, When n = 1, Ɩ= 0. i.e. K shell contains only one sub shell - s sub shell when n = 2, Ɩ = 0 and1. i.e. L shell contains two sub shells - s and p sub shells when n = 3, Ɩ = 0, 1 and 2. i.e. M shell contains three sub shells – s, p and d sub shells when n = 4, Ɩ = 0, 1, 2 and 3. i.e. N shell contains four sub shells – s, p,d and f sub shells 3. Magnetic Quantum Number (m or mƖ) It gives information about the orientation of orbitals in space. For a given ‘Ɩ’ value, there are 2Ɩ+1 possible values for m and these values are given by : m = – Ɩ to 0 to + Ɩ Thus for Ɩ = 0, mƖ = 0 [2(0)+1 = 1].. i.e. s sub shell contains only one orbital called s orbital. For Ɩ = 1, mƖ = –1, 0 and +1 [2(1)+1 = 3]. i.e. p subshell contains three orbitals called p orbitals (px, py and pz). For Ɩ = 2, mƖ = –2, –1, 0, +1 and +2, [2(2)+1 = 5]. i.e. d subshell contains five orbitals called d orbitals (dxy, dxz, dyz, dx 2 - y 2 and dz 2 ) 4. Spin Quantum Number (s or ms) It is the only experimental Quantum number and it gives the spin orientation of electrons. This spin may be either clockwise or anticlockwise. So the values for s may be +½ or -½. +½ represents clock-wise spin and-½ represents anticlockwise spin. Shapes of orbitals 1. s-orbitals For s-orbitals, Ɩ = 0 and hence mƖ = 0. So there is only one possible orientation for s orbitals. They are spherically symmetrical. The plots of probability density (ψ2 ) against distance from the nucleus (r) for 1s and 2s atomic orbitals are as follows: For 1s orbital the probability density is maximum at the nucleus and it decreases with increase in r. But for 2s orbital the probability density first decreases sharply to zero and again starts increasing. After reaching a small maximum it decreases again and approaches zero as the value of r increases. The region where the probability density (ψ2 ) reduces to zero is called nodal surface or node. For 1s orbital, there is no node, for 2s orbital there is only one node, for 3s orbital there are 2 nodes and so on. In general, for an ns-orbital there are (n – 1) nodes. All the s-orbitals are spherically symmetrical and their size increases with increase in n. The boundary surface diagrams for 1s, 2sand 3s orbitals are as follows: 2. p-orbitals For p-orbitals, Ɩ = 1 and mƖ = -1, 0, +1. i.e., there are three possible orientations for p orbitals. So there are 3 types of porbitals – px, py and pz . Each p orbital consists of two lobes. The probability density function is zero on the plane where the two lobes touch each other. ATOMIC STRUCTURE The size, shape and energy of the three orbitals are identical. They differ only in the orbital, the lobes are along the x-axis, for p have dumb-bell shape. The number of radial nodes for p-orbitals are given by 4p orbital and so on. Besides the radial nodes passing through the nucleus (origin). For example, in angular nodes and number of angular nodes is Number of radial nodes = n - Ɩ – 1 Number of angular nodes = Ɩ Total number of nodes = n-1 The boundary surface diagrams for three 2p orbitals are as follows: 3. d-orbitals For d-orbitals, Ɩ = 2 and mƖ = -2, -1, 0, +1 and +2. i.e., there are five possible orientations for d orbitals. So there are 5 types of d-orbitals. They are dxy, dxz, dyz, d of the fifth one, dz 2 , is dumb-bell having a circular collar in the xy orbitals the number of radial nodes is 2 and the total number of nodes is n as follows: 4. f-orbitals For f-orbitals, Ɩ = 3 and mƖ = -3, -2, there are 7 types of f-orbitals. They Rules for Filling of electrons in various orbitals The filling of electrons into the orbitals of different atoms takes place according to the 3 rules Pauli’s exclusion principle and the Hund’s rule of maximum multiplicity. 1. Aufbau principle: The German word aufbau means ‘build up’. The building up of orbitals m electrons. It states that the orbitals are filled in order of their increasing energies the lowest energy orbital and then to higher energy orbitals. This rule has two sub rules: a) The various orbitals are filled in the increasing order of their (n+ b) If two orbitals have the same (n+ The increasing order of orbitals is as 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5 The size, shape and energy of the three orbitals are identical. They differ only in the orientation of the lobes axis, for py, they are along the y-axis and for pz , they are along the z orbitals are given by (n –2), that is number of radial node is 1 for 3p orbital, ial nodes, the probability density functions for the np orbitals nucleus (origin). For example, in the case of pz orbital, xy-plane is a nodal plane number of angular nodes is given by ‘Ɩ’. The boundary surface diagrams for three 2p orbitals are as follows: 1, 0, +1 and +2. i.e., there are five possible orientations for d orbitals. So there are 5 , dx 2 -y 2 and dz 2 . The shapes of the first four d-orbitals are double dumb bell having a circular collar in the xy-plane. The five d-orbitals have equivalent energies. For d orbitals the number of radial nodes is 2 and the total number of nodes is n-2. Boundary surface diagrams for d 2, -1, 0, +1, +2 and +3. i.e., there are seven possible orientations for f orbitals. So They have diffused shapes. electrons in various orbitals into the orbitals of different atoms takes place according to the 3 rules Pauli’s exclusion principle and the Hund’s rule of maximum multiplicity. The German word aufbau means ‘build up’. The building up of orbitals means the filling up of orbitals with the orbitals are filled in order of their increasing energies. In other words, electrons first occupy the lowest energy orbital and then to higher energy orbitals. This rule has two sub rules: The various orbitals are filled in the increasing order of their (n+Ɩ) value. If two orbitals have the same (n+Ɩ) values, the orbital with the lower n value is filled first. The increasing order of orbitals is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s... Page 10 orientation of the lobes. For px , they are along the z-axis. All the p-orbitals , that is number of radial node is 1 for 3p orbital, two for density functions for the np orbitals are zero at the plane, plane is a nodal plane. These are called 1, 0, +1 and +2. i.e., there are five possible orientations for d orbitals. So there are 5 orbitals are double dumb-bell and that orbitals have equivalent energies. For d2. Boundary surface diagrams for d-orbitals are 1, 0, +1, +2 and +3. i.e., there are seven possible orientations for f orbitals. So into the orbitals of different atoms takes place according to the 3 rules - aufbau principle, eans the filling up of orbitals with . In other words, electrons first occupy Ɩ) values, the orbital with the lower n value is filled first. ATOMIC STRUCTURE Page 11 2. Pauli’s Exclusion Principle It states that no two electrons in an atom can have the same set of four quantum numbers. i.e. an orbital can accommodate a maximum of only 2 electrons with opposite spin. If 2 electrons have same values for n, Ɩ and m, they should have different values for s. i.e. if s = +½ for the first electron, it should be -½ for the second electron. 3. Hund’s rule of maximum multiplicity It states that electron pairing takes place only after partially filling all the degenerate orbitals. Orbitals having same energies are called degenerate orbitals. For example the electronic configuration of N is 1s2 2s2 2px1 py1 pz1 and not 1s2 2s2 2px2 py1 . Electronic Configuration of Atoms The distribution of electrons into various orbitals of an atom is called its electronic configuration. The electronic configuration of different atoms can be represented in two ways. (i) sa p b d c ...... notation (ii) Orbital diagram The electrons in the completely filled shells are known as core electrons and the electrons in the outer most shell are called valence electrons. Stability of Completely Filled and Half Filled Subshells For atoms having half filled or completely filled electronic configurations have extra stability compared to other atoms. This is due to their symmetrical distribution of electrons and greater exchange energy. For example, the electronic configuration of Cr is [Ar] 3d5 4s1 and not 3d4 4s2 . This is because d5 represents a half filled configuration and has extra stability. Similarly for Cu the electronic configuration is [Ar] 3d104s1 and not 3d9 4s2 .
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