Structure of Atom PDF

STRUCTURE OF AN ATOM DEFINTION OF AN ATOM  ATOM IS THE SMALLEST PARTICLE OF AN ELEMENT.  THE TERM ATOM WAS DERIVED FROM THE GREEK WORD “A-TOMIO” MEANS UNCUTTABLE OR INDIVISIBLE Discovery of Electron – Cathode rays  Electron was discovered by J J Thomson by Cathode ray discharge tu...

STRUCTURE OF AN ATOM DEFINTION OF AN ATOM  ATOM IS THE SMALLEST PARTICLE OF AN ELEMENT.  THE TERM ATOM WAS DERIVED FROM THE GREEK WORD “A-TOMIO” MEANS UNCUTTABLE OR INDIVISIBLE Discovery of Electron – Cathode rays  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. CATHODE RAY DISCHARGE TUBE CATHODE RAYS  Under very low pressure and very high voltage, the discharge tube begins to glow. This is due to the striking of some invisible rays from the cathode (CATHODE RAYS).  The rays which start from the cathode and move in straight lines are called cathode rays or cathode ray particles. Properties of Cathode Rays  The cathode rays start from cathode and move towards the anode.  They are invisible, but they can be observed with the help of fluorescent or phosphorescent materials.  In the absence of electrical or magnetic field, these rays travel in straight lines.  In the presence of electrical or magnetic field, the cathode rays behave like negatively charged particles.  The characteristics of cathode rays do not depend upon the material of the electrodes and the nature of the gas present in the cathode ray tube. ELECTRONS – THE FUNDAMENTAL PARTICLES  Cathode rays contains negatively charged particles called electrons.  Since the properties of cathode rays do not depend on the nature of the gas and the electrode, electrons are universal. Charge to Mass (e/m)Ratio of Electron  The charge to mass ratio of electron was determined by J.J Thomson.  He applied electrical and magnetic field perpendicular to each other as well as to the path of electrons. Observations made by J.J Thomson  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’.  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’. The e/m value of electron  J J Thomson calculated the e/m value of electron from the strength of electric and Magnetic field.  The e/m value of electron is 1.758 × 10 11 C/kg (C- Coulomb) 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)  Mass of electron (me) = e = 1.6022 x 10-19 e/me 1.758 x 1011  i.e. me = 9.1 ×10–31 kg Discovery of Protons  Protons were discovered by E.Goldstein.  He used modified discharge tubes with perforated (with small holes) cathode.  At very low pressure and very high voltage, he found that some rays were emitting behind the cathode.  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  They depend on the nature of gas present in the cathode ray tube.  The charge to mass ratio of the particles is found to depend on the gas from which these originate.  Some of the positively charged particles carry a multiple of the fundamental unit of electrical charge.  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.  He bombarded a thin sheet of beryllium by α-particles.  4 Be9 + 2 He4 → 6 C12 + 0 n1  They are electrically neutral particles having a mass slightly greater than that of the protons Characteristics of sub-atomic particles 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. THOMSON’S MODEL OF ATOM  According to this model, an atom is a positively charged sphere in which electrons are uniformly distributed like the seeds of a water melon.  The mass of the atom is uniformly distributed throughout the atom.  The total positive charge in an atom = 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. Rutherford’s Nuclear Model of Atom  The Experiment:  A stream of high energy α–particles from a radioactive source was directed to 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. Rutherford’s Nuclear Model of Atom  OBSERVATIONS: 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 (approximately 1 in 20,000) were deflected by nearly 180°. Rutherford’s Nuclear Model of Atom  CONCLUSIONS:  Since most of the α–particles passed through the foil without any deviation, most space in the atom is empty.  A few 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.  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. Rutherford’s Nuclear Model of Atom  Postulates of Rutherford’s Atom Model: 1. All the positive charge and most of the mass of an atom are 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.  Limitations of Rutherford’s atom model: 1. Rutherford’s model could not explain the stability of the atom. 2. He could not explain the electronic structure of atom.  Atomic Number (Z) It is the number of protons or number of electrons present in an atom. It is denoted by the symbol ‘Z’. Atomic number (Z) = number of protons (p) = number of electrons (e)  MASS NUMBER (A) It is the total number of protons and neutrons in atom. Or it is the total number of nucleons in an atom. ( Protons and neutrons in an atom are together called nucleons.) Mass number (A) = no. of protons (p) + no. of neutrons (n)  ISOTOPES Isotopes are atoms with same atomic number but different mass number. They contain same number of protons but different number of neutrons. Isotopes have similar chemical properties but have different physical properties. Hydrogen has three isotopes Protium (1H1) Deuterium (1H2 or 1D2) Tritium (1H3 or 1T3).  No. of No. of No. of Isotopes Symbol protons electrons neutrons Protium 1H1 1 1 0 Deuterium H2 or 1D2 1 1 1 1 Tritium 1 H3 or 1T3 1 1 2  ISOBARS 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 number of nucleons (sum of the protons and neutrons). e.g. 6C14 and 7N14 18 Ar 40 and 20 Ca 40  ISOTONES Isotones are atoms having same number of neutrons but have different atomic numbers. e.g. 6C14 , 7N15, 8O16 etc.  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. Characteristics of Electromagnetic radiations  The oscillating electric and magnetic fields are perpendicular to each other and both are perpendicular to the direction of propagation of the wave.  The electromagnetic waves do not require a medium for propagation and can move in vacuum.  All the electromagnetic radiations travel through vacuum with a constant speed of 3 x 108 m/s.  Electromagnetic spectrum  It is a series in which different types of electromagnetic radiations are arranged in the increasing order of their wavelength (or decreasing order of frequency).  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.  Some terms related to emr  Frequency (ν )  It is the number of waves passing through a given point in one second.  The SI unit for frequency is hertz (Hz).  Wavelength (λ) It is the distance between two adjacent crusts or two adjacent troughs. Its unit is m or cm. Wave number (ῡ ) It is the number of wavelengths per unit length. It is the reciprocal of wavelength. i.e. ῡ = 1/ λ Its unit is m–1 or cm-1. The frequency (ν ), speed of light (c) and the wave length (λ)are related as c= ν λ  Black body radiation  Anideal 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.  Planck’s Quantum Theory  The phenomenon of black body radiation was first explained by Max Planck by his Quantum theory. According to this theory: Atoms and molecules could emit (or absorb) energy discontinuously in small packets of energy called quantum 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 h= 6.626×10–34 J s (joule-second).  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. Photoelectric effect characteristics of photoelectric effect  The electrons are ejected from the metal surface as soon as the beam of light strikes the surface.  The number of electrons ejected is proportional to the intensity or brightness of light. For each metal, there is a minimum frequency (known as threshold frequency [ν0]) below which photoelectric effect is not observed. The kinetic energy of the ejected electrons is directly proportional to the frequency of the incident light.  Photoelectric effect - Explanation  Photoelectric effect was first explained by Albert Einstein using Planck’s Quantum theory.  According to him, when a photon of sufficient energy strikes the metal surface, it suddenly transfers its energy to the electron 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 [this energy is known as work function, hν0]. The other part of energy 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 mev2 Where me is the mass of the electron v is the velocity of the ejected electron. Qn. Calculate energy of one mole of photons of radiation whose frequency is 5 x 10 ^14 Hz. Qn. The threshold frequency V for a metal is 0 7.0×10^14 s. Calculate the kinetic energy of -1 an electron emitted when radiation of frequency V =1.0 ×10^15 s hits the metal. –1  Particle nature could explain the 1. black body radiation 2. photoelectric effect satisfactorily but could account for the phenomena of interference and diffraction.  Dual Behaviour of Electromagnetic Radiation  Electromagneticradiations possess both particle and wave nature. This is known as dual nature of Electromagnetic radiation.  SPECTRUM a ray of white light is spread out into a series of coloured bands called spectrum. CONTINUOUS SPECTRUM The spectrum of white light, that we can see, ranges from violet to red. Such a spectrum is called continuous spectrum  ATOMIC SPECTRUM  Thespectrum produced by an excited atom or molecule is called atomic spectrum.  Atomic spectra are of two types – 1. Emission spectrum 2. Absorption spectrum.  Emission spectrum The spectrum of radiation emitted by a substance that has absorbed energy is called an emission spectrum. To produce an emission spectrum, energy is supplied to a sample by heating it or irradiating it and the wavelength (or frequency) of the radiation emitted is recorded.  ATOMIC SPECTRUM  Theemission spectra of atoms contains some lines with dark spaces between them. So they are called line spectra or atomic spectra.  Each element has a unique line emission spectrum. So line emission spectra are also called finger print of atoms.  Line spectrum of Hydrogen  Thehydrogen spectrum consists of mainly five series of lines. They are:  1. Lyman series  2.Balmer series  3.Paschen series  4.Brackett series  5. Pfund series.  SERIES SPECTRAL REGION LYMAN U.V. REGION BALMER VISIBLE REGION PASCHEN INFRA RED BRACKETT SERIES INFRA RED PFUND INFRA RED  TheBalmer series is the only series that we can be visible.  RYDBERG EQUATION  Rydberg equation for finding the wave number of different lines in Hydrogen spectrum is: ῡ = 1/ λ =109677 (1/n12 -1/n22) cm-1  Where n1 = 1, 2, 3,….. n2 = n1 + 1, n1 + 2, …… 109677 = Rydberg constant  BOHR’S MODEL FOR HYDROGEN ATOM  The important postulates are:  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.  Theenergy 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).  Thefrequency of radiation absorbed or emitted when transition occurs between two stationary states that differ in energy by ΔE, is given by:  Where E1 and E2 are the energies of lower and higher energy levels respectively. This expression is commonly known as Bohr’s frequency rule.  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  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. Results of Bohr’s theory 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.  The radius of orbits can be given by the equation: rn = a 0 n 2 where a0 = 52.9 pm.  Thusthe radius of the first stationary state is 52.9 pm (called the Bohr radius). As n increases, the value of r will increase.  The energy of electron in an orbit is given by the expression: En = -RH (1/n2).  where n = 1,2,3…… RH =Rydberg constant= 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. Qn. Calculate the energy associated with the first orbit of He. What is the + radius of this orbit? BOHR’S MODEL FOR HYDROGEN ATOM  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 attracted by the nucleus and is present in an orbit n, the energy is emitted. So it becomes less than zero. That is the reason for the presence of negative sign in equation. Limitations of Bohr Atom Model 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 could not explain Stark effect and 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.  Quantum mechanical model Two important developments which contributed significantly in the formulation of such a model were : 1. Dual behaviour of matter, 2. Heisenberg uncertainty principle.  Dual Behaviour of Matter  deBroglie 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.  Thewaves associated with matter is called Matter waves.  de Broglie’s equation  De Broglie proposed a relation between wavelength (λ) and momentum (p) of a material particle. This equation is known as de Broglie’s equation.  The equation is: λ= h = h mv p Where m is the mass of the particle, v is the velocity and p is the momentum (p=mv). This equation is applicable to all material bodies. But for macroscopic bodies the wavelength is too small to detect.  Heisenberg’s Uncertainty  Principle 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 by the 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  Significance of Uncertainty Principle  Heisenberg Uncertainty Principle is significant only for motion of microscopic objects.  Accordingto this Principle, we cannot determine the exact position and momentum of an electron.  Thusit rules out the existence of definite paths or orbits of electrons.  We can only say the probability of finding an electron at a given point. Failure of Bohr Model In Bohr model, an electron is regarded as a charged particle moving in well defined circular orbits The wave character of the electron is not considered in Bohr model. Further, an orbit is a clearly defined path and this path can completely be defined only if both the position and the velocity of the electron are known exactly at the same time Bohr model of the hydrogen atom, therefore, ignores dual behaviour of matter and also contradicts Heisenberg uncertainty principle. Quantum mechanics is a theoretical science that deals with the study of the motions of the microscopic objects that have both observable wave like and particle like properties.  QUANTUM MECHANICAL MODEL OF ATOM  Erwin Schrodinger and Werner Heisenberg proposed a new model of atom called Quantum mechanics. The fundamental equation of quantum mechanics was developed by Schrödinger and is known as Schrödinger equation. The equation is: Ĥ ψ = Eψ where Ĥ is a mathematical operator called Hamiltonian operator, E is the total energy of the system (K.E + P.E) ψ is called the wave function.  Significance of ψ:  The wave function (ψ) is a mathematical function and it has no physical meaning. But ψ2 has some physical significance. It gives the probability of finding an electron at a point within an atom. So ψ2 is known as probability density. From the value of ψ2, it is possible to predict the probability of finding the electron around the nucleus.  Quantum Numbers  These are certain numbers used to explain the size, shape and orientation of orbitals.  Or,Quantum numbers are the address of an electron. There are four quantum numbers. They are: 1. Principal Quantum number (n) 2. Azimuthal Quantum number (Ɩ) 3. Magnetic Quantum number (m or mƖ) 4. Spin Quantum number (s or ms)  Principal Quantum number (n) It gives the following informations:  the size the orbit.  the energy of electron in an orbit.  the shell in which the electron is found.  the average distance between the electron and the nucleus. The possible values of n are 1, 2, 3, 4, etc. n= 1 represents K shell n = 2 represents L shell n = 3 represents M shell etc. Azimuthal Quantum Number [Subsidiary Quantum number] (Ɩ) It gives the following informations: 1. the shape of the orbital. 2. the sub shell or sub level in which the electron is located. 3. the orbital angular momentum of the electron. the possible value of Ɩ are 0, 1, 2, 3,.......... (n-1). For any n Ɩ = 0 to n-1  Whenn = 1, Ɩ= 0. i.e. K shell contains only one sub shell - s sub shell.  whenn = 2, Ɩ = 0 and1. i.e. L shell contains two sub shells - s and p.  whenn = 3, Ɩ = 0, 1 and 2. i.e. M shell contains three sub shells – s, p and d.  whenn = 4, Ɩ = 0, 1, 2 and 3. i.e. N shell contains four sub shells – s, p,d and f. Magnetic Quantum Number (m)  It gives the orientation of orbitals in space.  For a given ‘Ɩ’ value, there are (2Ɩ+1) possible values for m.  The possible values for m are:– Ɩ to 0 to + Ɩ. When Ɩ = 0, mƖ = 0. i.e. s-sub shell contains only one orbital called s orbital. For Ɩ = 1, mƖ = –1, 0 and +1. i.e. p subshell contains three orbitals called p orbitals (px, py and pz). For Ɩ = 2, mƖ = –2, –1, 0, +1 and +2. i.e. d subshell contains five orbitals called d orbitals (dxy, dxz, dyz, dx2- y2 and dz2)  Spin Quantum Number (s or ms )  It is the only experimental Quantum number.  It gives the spin orientation of electrons.  It has 2 values: +½ or -½.  +½ represents clock-wise spin  -½ represents anticlock-wise spin.  Shapes of orbitals  s-orbital  Fors-orbitals, Ɩ = 0 and hence m=0. So there is only one possible orientation for s orbitals.  They are spherically symmetrical.  All s-orbitals have same shape but they have different size.  Shape of s-orbitals  1s, 2sand 3s orbitals are: 1111  Shape of s-orbitals  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 around the nucleus and it decreases with increase in r.  Butfor 2s orbital the ψ2 first decreases sharply to zero and again starts increasing.  After reaching a small maximum it decreases again and approaches zero  Node:  The region where the probability density (ψ2) reduces to zero is called node.  Thereare two types of Nodes: Radial node and angular node.  Number of radial nodes = n - Ɩ – 1  Number of angular nodes = Ɩ  Total number of nodes = n-1  Shape of 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 p-orbitals – px, py and pz.  Each p orbital consists of two lobes.  The boundary surface diagrams for three 2p orbitals are as follows:  The no. of radial nodes for p-orbitals = n–l-1.  i.e. for 2p orbital, n = 2 and l= 1. So no. of radial nodes =0, for 3p orbital, it is 1 and so on.  The probability density functions for the p- orbitals are zero at the plane passing through the origin (angular node).  For example, in the case of pz orbital, xy- plane is a nodal plane.  The number of angular nodes for p-orbitals = 1.  Shape of d-orbitals  For d-orbitals, Ɩ = 2 and m = -2, -1, 0, +1 and +2.  i.e., there are 5 types of d-orbitals. They are dxy, dxz, dyz, dx2-y2 and dz2. The shapes of the first four d- orbitals are double dumb-bell. The shape of dz2 orbital is dumb-bell having a circular collar in the xy- plane.  Shape of f-orbitals  For f-orbitals, Ɩ = 3 and m = -3, - 2, -1, 0, +1, +2 and +3.  i.e., there are seven possible orientations for f orbitals.  So there are 7 types of f-orbitals [fx3, fy3, fz3, fxyz, fx(y2-z2), fy(z2- x2), fz(x2-y2)].  The orbitals having the same energy are called degenerate. px, py and pz dxy, dxz, dyz, dx2-y2 and dz2.  Rules for Filling of electrons in various orbitals  Various orbitals are filled according to the 3 rules: 1. Aufbau principle 2. Pauli’s exclusion principle 3. Hund’s rule of maximum multiplicity.  Aufbau principle  The German word aufbau means ‘build up’.  It states that the orbitals are filled in order of their increasing energies. . This rule has two sub rules: i. The various orbitals are filled in the increasing order of their (n+Ɩ) value. ii. If two orbitals have the same (n+Ɩ) values, the orbital with the lower n value is filled first The increasing order of energy for various orbitals are: 1s,2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, ……………..  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.  Hund’s rule of maximum multiplicity.  It states that electron pairing takes place only after partially filling all the degenerate orbitals. For example the electronic configuration of N is 1s2 2s2 2px1py1pz1 and not 1s2 2s2 2px2py1. Orbitals having same energies are called degenerate orbitals.  Electronic Configuration of Elements  Thedistribution of electrons into various orbitals of an atom is called its electronic configuration.  Theelectrons in the completely filled shells are known as core electrons.  The electrons in the outer most shell are called valence electrons. Electronic Configuration  The maximum no. of electrons filled in various sub shells are: Sub shell Max. no. of electrons s -sub shell 2 p -sub shell 6 d -sub shell 10 f -sub shell 14  Element Atomic No. Electronic configuration  H 1 1s1  He 2 1s2  Li 3 1s2 2s1  Be 4 1s2 2s2  B 5 1s2 2s2 2p1  C 6 1s2 2s2 2p2  N 7 1s2 2s2 2p3  O 8 1s2 2s2 2p4  F 9 1s2 2s2 2p5  Ne 10 1s2 2s2 2p6  Element Atomic No. Electronic configuration  Na 11 1s2 2s2 2p63s1  Mg 12 1s2 2s2 2p6 3s2  Al 13 1s2 2s2 2p6 3s2 3p1  Si 14 1s2 2s2 2p6 3s2 3p2  P 15 1s2 2s2 2p6 3s2 3p3  S 16 1s2 2s2 2p6 3s2 3p4  Cl 17 1s2 2s2 2p6 3s2 3p5  Ar 18 1s2 2s2 2p6 3s2 3p6  K 19 1s2 2s2 2p6 3s2 3p6 4s1  Ca 20 1s2 2s2 2p6 3s2 3p6 4s2  Element Atomic No. Electronic configuration  Sc 21 1s2 2s2 2p6 3s2 3p6 4s2 3d1  Ti 22 1s2 2s2 2p6 3s2 3p6 4s2 3d2  V 23 1s2 2s2 2p6 3s2 3p6 4s2 3d3  Cr 24 1s2 2s2 2p6 3s2 3p6 4s1 3d5  Mn 25 1s2 2s2 2p6 3s2 3p6 4s2 3d5  Fe 26 1s2 2s2 2p6 3s2 3p6 4s2 3d6  Co 27 1s2 2s2 2p6 3s2 3p6 4s2 3d7  Ni 28 1s2 2s2 2p6 3s2 3p6 4s2 3d8  Cu 29 1s2 2s2 2p6 3s2 3p6 4s1 3d10  Zn 30 1s2 2s2 2p6 3s2 3p6 4s2 3d10  Electronic Configuration using noble gas configuration  The electronic configuration can be simplified by using noble gas configuration as follows: Noble Gas Atomic Sub shell Number filled after this confgn. He 2 2s Ne 10 3s Ar 18 4s Kr 36 5s Xe 54 6s Rn 86 7s Na 11 1s2 2s2 2p63s1 Ne 10 1s2 2s2 2p6 Na 11 [Ne] 3s1 Ar 18 1s2 2s2 2p6 3s2 3p6 K 19 1s2 2s2 2p6 3s2 3p6 4s1 K 19 [Ar] 4s1  Element Atomic No. Electronic configuration  Na 11 [Ne] 3s1  P 15 [Ne] 3s2 3p3  Cl 17 [Ne] 3s2 3p5  Ca 20 [Ar] 4s2  Sc 21 [Ar] 4s2 3d1 or, [Ar] 3d1 4s2  V 23 [Ar] 4s2 3d3 or, [Ar] 3d3 4s2  Cr 24 [Ar] 4s1 3d5 or, [Ar] 3d5 4s1  Ni 28 [Ar] 4s2 3d8 or, [Ar] 3d8 4s2  Cu 29 [Ar] 4s1 3d10 or, [Ar] 3d10 4s1  Zn 30 [Ar] 4s2 3d10 or, [Ar] 3d10 4s2  Extra Stability of Half Filled and Completely Filled Subshe  Atoms having half filled or completely filled electronic configurations have extra stability.  Thisis due to their symmetrical distribution of electrons and greater exchange energy. Exceptional configuration of Cr and Cu  The electronic configuration of Cr is [Ar] 3d 54s1 and not 3d44s2.  This is because of the extra stability of half filled d 5 configuration.  Similarly for Cu the electronic configuration is [Ar] 3d 104s1 and not 3d94s2.  This is due to the extra stability of completely filled d 10 configuration.