Chapter 02: Atomic Structure

Moseley's Law Statement: This law states that "the square root of the frequency of the X-rays was directly proportional to the atomic number of an element" Mathematical form: V frequency (v) a Z Explanation In 1913, Moseley. observed that when different elements were bombarded with cathode rays, the X-rays of some characteristics were produced Proton number: He concluded that, the number of protons in the nucleus of an atom is called atomic number or proton number. It is represented by symbol 'Z' and it is fundamental DIC 5H property of an element. Nucleon number: The number of protons and neutrons in the nucleus of an atom is collectively called its nucleon number, also called mass number. It is represented by symbol 'A'. Atomic number is related to the mass number by the following equation. A =Z+N Example, an atom of an element X having atomic number Z and mass number A is described as ¿ X like i3 Al. Number of neutrons in an atom can be calculated as N=A-Z Consider 13 Al, Atomic number/proton number (Z) = 13 Mass number/nucleon number (A) = 27 N = A -Z=27 - 13 = 14 Similarly, the number of electrons, protons and neutrons can be justified for an ion as in the following example: i3 Al atom loses three electrons to form Al3+ , then; Number of protons = 13 Number of neutrons = 14 Number of electrons = 13 - 3 = 10 Similarly, 1, Cl gain an electron to form Cl- ion; Number of protons = 17 Number of neutrons = 18 Number of electrons = 17 + 1 = 18 In case electron gain happens by the neutral atoms, say ! Oto s02- 1652- , the number of neutrons, protons and electrons are as follows. Table 2.1 Number of protons, electrons and neutrons in different ions Species Neutrons 8 16 S2- 16 p3- Thus, the atomic number and proton number represent the same concept. EFFECT OF ELECTRIC FIELD ON FUNDAMENTAL PARTICLES

Fundamental Particles: Electron, proton and neutron are said to be fundamental particles. The behaviour of particles in an electric field depends upon their mass and charge. If the beams of electrons, protons and neutrons are allowed to pass one by one at the same speed through an electric field, they show their behaviour as follows: • Neutrons being neutral are not deflected but travel in a straight path perpendicular to the direction of electric field. • Protons being positively charged are deflected towards the negative plate. Electrons being negatively charged are deflected towards the positive plate, to greater 1 extend since they are times lighter than protons. 1836 Beam of Particles ++++++ Fig: Behaviour of proton, electron and neutron in the electric field The amount of deviation from its original direction of movement is measured in two ways. (i) Angle of deflection oc Charge mass The angle of deflection of fundamental particles is directly proportional to the charge to mass ratio. , iP to, P and i6S to Electrons Protons 10 8 18 16 15 18 -p+ De mass (ii) Radius of deflection o charge The radius of deflection is directly proportional to mass to change ratio. This is possible if we imagine that after deflection, the particle moves in a circular path. Hence, the factors affecting the radius of deflection are reciprocal to that for the angle of deflection. The Table (2.2) shows the properties of three fundamental particles electron, proton and neutron present in an atom. Table 2.2 Properties of three fundamental particles Particle Charge (coulomb) +1.6022 × 10-19 Proton 0 Neutron -1.6022 × 10-19 Electron Quick Check 2.1 (a) Calculate the number of neutrons in the following elements. 1K, 15CC, 18 Ar 15 Cl Ans. 11K N= A-Z N= A-Z N= 35-17 N= 39-19 N = 20 N= 18 (b) Which of the electron, proton and neutron deflected the most in the magnetic field? Ans. Electron: It is deflected the most in magnetic field because it has highest charge to mass ratio. Proton: It is less deflected in magnetic field because its charge to mass ratio is less than electron. Neutron: It is not deflected because it is neutral. EXPERIMENTAL EVIDENCES FOR ELECTRONIC CONFIGURATION

Experimental Evidence: The modern theory of electronic structure or atomic spectra originates from the Bohr Model of atom. Evidences for atomic models of the atoms derives principally from following two sources; (i) Atomic spectra and ii) Ionization energies. Atomic Spectra: There are two types of atomic spectra. (a) Atomic emission spectrum: When an element in its gaseous state is heated to high• temperatures or subjected to electrical discharge, radiation of certain wavelengths is emitted. The spectrum of this radiation contains coloured lines and is called atomic emission spectrum. The atomic emission spectrum of hydrogen is shown in figure. (b) Atomic Absorption Spectrum: When a beam of white light is passed through a gaseous sample of an element in cold state, certain wavelengths are absorbed. • The wavelengths of the white light that has been absorbed by the atoms show up as dark lines on the spectrum. Mass (amu) Relative Mass (kg) charge +1 1.0073 1.6726 × 10~ 1.0087 0 1.6750 × 10-27 9.1095 × 10-31 5.4858 × 104 18 Ar N= A-Z N= 40-18 N = 22 • The spectrum of this radiation is called an atomic absorption spectrum. • The atomic absorption spectrum of hydrogen is shown in figure. • The wavelengths of the dark lines in the absorption spectrum are exactly the same as those of coloured lines in the emission spectrum. (a) Emission spectrum of Hydrogen (b) Absorption spectrum of Hydrogen 400 nm Fig: (a) atomic emission spectrum of hydrogen (b) atomic absorption spectrum of hydrogen Relation with electronic configuration • Fixed energy levels: Each element has a unique arrangement of electrons and thus a unique range of fixed energy levels. • Radiation: It follows that the 700 nm wavelengths and frequencies of the radiation absorbed or emitted when electrons jump from one energy level to another must also This uniqueness unique. convinces us to conclude that has its own every element characteristic spectrum. • Identification of element Every element is identified by its characteristic spectrum. • Finger prints: We can say that atomic spectra are the finger prints of the elements. figure shows the emission spectra of some elements. ionization ENERGY AND ENRGY LEVELS

Ionization and Energy levels • A major difference between electrons in different types of shells is their energy. • We can investigate the electronic configuration of the atoms by measuring experimentally the energies of the electrons within them. • This can be done by measuring ionization energies of atoms. Ionization energies are used to investigate the electronic configurations of elements in two ways. (i) Successive ionization energies of the same element (ii) First ionization energies for different elements 700 nm 400 nm Helium Neon Sodium Mercury Fig: Spectral series of various elements as plotted by a spectrometer (i) Successive Ionization Energies of the Same Element An atom of a particular element is studied and measure the energy required to remove each of its electrons, one by one. x →X+ + x→x2++ x2+→X3++ Removal of electrons continuously from an atom until only the nucleus is left. • We call this sequence of ionization energies, is successive ionizațion energies. • The successive ionization energies show clearly the arrangement of electrons in shells around the nucleus. Example: If we take the magnesium atom as an example and measure the energy required to remove successively the first electron, the second, the third and so on. • Graph: We obtain a graph when the ionization energies values are plotted against number of electrons as in figure. • The plot shows that successive ionization energies increase when we move from the valence shell to the inner shell. • First two electrons are removed from the outermost shell and require lower energy for their removal. increase • A large occurs when the third electron is removed • This is because when two electrons of the outer shell have been removed, the next has to be removed from the shell that is very 3rd shell much closer to the electrons nucleus. • The next • seven electrons are removed Logio (ionization energy) successively from the second shell and a gradual increase in 1 2 ionization energy is observed • A similar but much Fig: A plot of the successive ionization energies of Mg more enormous jump occurs when the eleventh and twelfth electrons are removed from the first, innermost shell, right next to the nucleus. • Conclusion: We observe two large jumps in the successive ionization energies. These two large jumps in the series of successive ionization energies. are very good evidence that the electron in the magnesium atoms exist in three different shells. (ii) First ionization energies of different atoms è ДНі, ΔH, ДНіз 1st shell electrons 2nd shell electrons 2nd Jump 1st Jump 12 3 4 5 6 7 8 9 10 11 Numbers of electrons removed Ionization energy and atomic No: The second way in which ionization energies show us the details of electronic configuration is to look at how the first ionization energies of elements vary with atomic numbers. The following figure shows a plot for the first 88 elements. This graph tells us the following (i) All ionization energies are strongly endothermic; it takes energy to separate an electron from an atom. (ii) Decreases in Groups: As we go down a particular group, for example, from helium to neon to argon, or from lithium to sodium to potassium, ionization energies decrease. • The larger the atom, the easier is to separate an electron from it:- • Reason: Actually, down the group, number of shells increases, hold of nucleus on the valence electrons decreases, hence removal of electrons becomes easier. (wii) Increases in Period: The ionization energies generally increase on going across a period • The group 1 elements, the alkali metals, have the lowest ionization energy within each period and the noble gases have the highest. Reason: It is due to the reason that across the period shell number remains same. As the proton number increases, electrons are added in the same shell. Therefore, nucleus attracts the valence electrons more strongly. As a result, ionization energy increases across the period Period 35J-0 2,500 #Period 2 Period 4 Ne 2000 Ar 1500 Kr N H Cl Br C 1000 Be Se N ge Si Ga 500 Na Rb Ionization energy. (kJ mol-1) 36 10 18 Atomic number Fig: A plot of the ionization energies against atomic number Quick Check 2.2 (a) Write equations that describe. (i) 1% ionization energy of calcium Ans. →Ca(g) + le (ii) 3rd ionization energy of potassium →K3+ Ans. (g) (g) te (ill) 2nd ionization energy of lithium • SP In Ra Cs 54 86 - +é Lit! →Lit2 Ans. (8) (iv) 5th ionization energy of sulfur 54+ - → S*+ +é Ans. (8) (8) (b) For the element aluminium (Z=13), draw a sketch graph between the of the successive ionization energies (y-axis) against the number of electrons removed (x- axis). Ans. A graph of successive ionization of Al13 is given: 2nd shell electron 3rd shell electron Log10 (ionization energy) 2nd Jump irenic. 12345618.910111213 Number of electron removed A plot of successive ionization energies of Al. (c) The first ΔH and the second Hiz ionization energies (kJ/mol) of a few eléments are given in table. Element II III Which of the above element is likely to be? (i) A reactive metal (ii) A reactive non-metal (wii) A noble gas (iv) A metal that forms a stable binary halide of the formula AX2 (X = halogen) Ans. (i) Element I: It is a noble gas & has two electrons that is He, because of high 1st ionization energy. (i) Element II: It is lithium that has low 1s ionization energy and high 2nd ionization energy shows that its metal and has one electron in its outermost shell. It is reactive. (i) Element III: Two successive ionization energies 900 and 1760 shows that it is a metal 1st shell electron 1st Jump AHil AHi2 2382 5251 520 7300 900 1760 1680 3380 and has two valence electrons. It is alkaline earth metal and is Be. It can form AX2 i.e. BeX2. (iv) Element IV: The high 1s ionization energy shows that element is non-metal and difficult to lose 1s electron and 2nd electron as well, that should be of smaller sized non-metal. F has 1680 1% ionization energy. QUANTUM NUMBERS

Definition: Quantum numbers are set of numbers that describe the properties of atomic orbitals and the electron within them. Quantum Numbers Demerits of Bohr's Principal and Azimuthal model quantum number: The Bohr's model was a two-dimensional model that used one quantum number to describe the distribution of electrons in the atom. • Merits of Bohr's model: The only important information was the size and energy of the orbit, which was described by the n-quantum number. • Since Schrodinger's model allowed the electron to occupy three-dimensional space, therefore, it required three coordinates, or three quantum numbers, to describe the orbitals in which electrons can be found • The three quantum numbers that come from Schrodinger's wave equation are the principal (n), azimuthal (e), and magnetic (m) quantum numbers. • These quantum numbers describe the size, shape and orientation in space of the orbitals in an atom. Principal Quantum number (n) The principal quantum number, n, can have positive integral values 1,2,3,4... designated by K, L, M, N..., and it corresponds the quantum number (number of orbit) in Bohr's model of the hydrogen atom. • This quantum number, n, describes the size and energy of the orbital • The collection of orbitals with the same values of n is called an electron shell. • The larger 'n' is, the greater the average distance of an electron in the orbital from the nucleus and therefore the larger the orbital. • An increase in 'n' also means that the electron has a higher energy and is therefore less tightly bound to the nucleus. • The principal Quantum number, n, can also be used to calculate the maximum number of electrons in a shell by the formula 2n? •Therefore, shells K, L, M and N can accommodate maximum electrons 2,8,18,32 respectively. Azimuthal Quantum Number (l): Definition: • Azimuthal Quantum number. also called Angular Momentum Quantum Number, t, can have integral values 0 to (n -1) for each value of n. Explanation: This quantum number describes the shape of the orbital. • The number of subshells in a shell. • The number of electrons in a subshell. • The values of " (' are integers that depend on the value of the principal quantum number. If n = 1, there is only one possible value of 'l' i.e, l = 0 (n-1, where n = 1). If n = 2, there are two values of '(' i.e. 0 and 1. If n = 3 there are three values of "l' i.e., 0,1 and 2. If n = 4, there are four values of "l', i.e. 0,1,2 and 3. • The values of (' are designated by the letters s,p,d and f, with which stand for sharp, principle, diffused and fundamental, respectively. • These are the spectral terms used to describe certain features of spectral lines. • The set of orbitals that have the same n and l values is called a subshell. The number of electrons in a subshell can be calculated by the formula 2(2l+1) as given in the table 2.3. Table 2.3 Shapes of orbital 0 Value of l S Orbital designation spherical shape of orbital (dumb bell): (double dumbbell) 2 No. of electrons in a subshell, 2(2l +1) The number of subshells in a shell is equal to its shell number. For example 1st, 2nd, 3rd, and 4th shells have one, two, three and four subshells respectively Table 2.4 Relationship between ni, land subshells Azimuthal~ Principal Quantum Shell number 'n' Quantum Number (l) 1 K 2 L 3 M 4 N. Did you Know! S.Q. How orbitals in subshell are confirmed? Ans. Splitting of small fine lines in the presence of magnetic field (Zeeman's effect) and there three dimensional orientation in space indicate the presence of orbitals in subshells. Magnetic Quantum Number (m) Definition: The magnetic quantum number 'm' describes the orientation of an orbital in space. Example: Within a subshell, the value of m depends on the value of l. For a certain value of l, there are (2l + 1) integral values of m as follows: -l.....o.....+C • The values of m indicate the number of orbitals in a subshell. 3 1 2 f d cloverleaf complicated polar 14 10 No. of subshells in a Subshells shell 1 1s 0 2s 0 2 .. 2p 1 3s 3 3p 1 3d 2 4s 4p 4 2 4d 4f • If l = 0, (s-subshell) there is only one possible value of m, i.e. O. It means s-subshell has only 1 orbital. • If l=1, (p-subshells) there are three values of m; -1,0 and +1. It means p-subshells has three orbitals. • If l=2, (d-subshell), there are five values of m, namely, - 2, - 1, 0; + 1 and + 2. It means d-subshells have 5 orbitals. • If l =3, (f-subshell) there are seven values of m: i.e: -3, -2, -1, 0, +1, +2, +3. It means f- subshells have 7 orbitals. Orbitals of the same subshell have same energy and are called degenerate orbitals. • These degenerate orbitals are differentiated from each other in the presence of magnetic field, hence the name of this quantum number, ie. magnetic quantum number. The relationship between the magnetic quantum numbers is provided in Table 2.5. Table 2.5 Relationship between l and m Azimuthal Quantum number Subshell (P) S 1 d 2 f 3 Spin Quantum Number (s): Definition: The spin quantum number describes the intrinsic angular moment of electron or its spin, whether it is clockwise or anticlockwise. Electrons are thought of spinning around their own axes, as the Earth does. According to electromagnetic theory, a spinning charge generates a magnetic field. It is the motion that causes an electron to behave like a magnet. Figure shows the two possible spinning motions of an electron. • One is clockwise and the other is anticlockwise. To take the electron spin into account, it is necessary to introduce Magnetic Number of Quantum degenerate number (m) orbitals (2/'+1) 0 One orbital -1 Three degenerate p-orbitals +1 +2 +1 Five degenerate d- orbitals +2 +1 Seven degenerate 0 f-orbitals -3 S N N S (a) (b) Fig (a) clockwise (b) counter clockwise spins of an electron a fourth quantum number called the electron spin quantum number (s). 1 - and - Its values are + ž as in figure. • The clockwise spin is represented by an arrow (1) pointing downwards, while the anti-clockwise spin is represented by an arrow (1) pointing upwards. Each orbital can accommodate at the most two electrons provided the two electrons have opposite spins. Thus it takes three quantum numbers to describe an orbital but forth quantum number to differentiate between the two electrons that can occupy an orbital. Did you Know! S.Q. How no. of electrons and no. of sub-shell are calculated? Ans. In the nth principal quantum number, there are n subshells consisting of n' orbitals with a maximum number of 2n? electrons. Quick check 2.3 (a) What information about an electron in an atom can be obtained from: (i) Principal quantum number Ans. The information obtained from principal quantum number is of shells. It is denoted as "n". And n = 1, 2, 3.... (ii) Azimuthal quantum number Ans. This quantum number gives us the information about sub shells. It is denoted as "f". {= n-1. It's values are l= 0, 1, 2, 3, ---- It can be zero. 1 = 0 (s-subshell), {=1(p-subshell) l= 2 (d-subshell), l =3 (f-subshell) (wii) Magnetic quantum number Ans. This quantum number tells us the information about orbitals. Shapes and orientations in space can be determined by this Q.N. It is denoted as "m". It's value are: m =0,+1, +2, +3,-- (iv) Spin quantum number Ans. This quantum number tells us the information about the spinning of electron in an orbital. It is denoted as "s". it's values are +½ and - ½: (b) For on electron(s): (i) If n = 2 and l= 1, how many orientations in space are possible? (ii) If n =3 and l= 2, which shell and subshell does the electron belong to? (wii) If l = 2, find all possible values of m and maximum number of electrons for 'm'? Ans. (i) If n = 2, it is 2nd shell (L), There are three space orientations for p sub-shell, Px, Py and pz (i) If n = 3; it is 3rd shell (M-shell), Then the electrons are present in d-subshell of 3rd shell. (wii) If e = 2, then it is d sub-shell. The possible values for d-subshell of "m" are - 2, -1, 0, + 1 and + 2. Because it has five orbitals. Total electrons that can be placed in all "m" values are 10 in 5 orbitals. It will never be zero (n # 0) l=1, it is p= subshell 1=2, it is d = subshell SHAPES OF ORBITALS

s-orbital: The shape of an 's' orbital is spherical. Explanation • The electronic density around the nucleus in an 's' orbital is uniformly distributed in all directions. • With the increase in the principal quantum number, the size of's orbital also becomes larger. p-orbitals The distribution of electron density for a 2p orbital is shown in figure. Explanation • The electron density is not distributed in a spherically symmetric fashion as in an s orbital. • p orbital has two lobes on any of the axes. • The p orbitals are named as Px Py, and pz accordingly to their axis. 2py orbital 2Px orbital l=1,m = -1 l= 1,m = +1 Fig: Shapes of three 2p orbitals d-orbitals In a given shell, 'd' orbitals have different shapes and orientations in space. • Explanation: The day, dxz, and dye lie in the xy, xz, and yz planes, respectively: y lie along the x and y axes. The lobes of the d* - y The dz orbital has two lobes along the z-axis and a "doughnut" in the xy plane. The shape. of d orbital is complicated or message shape. • The shapes for the five d orbitals (l = 2,) are shown in Figure. Z Z d xy Oyz. Fig: Shapes of five d orbitals → X → X Is 25 (b) n=2,1=0, m= 0 (a) n=1,1 = 0, m =0 Fig: Boundary surface for, (a) Is (b) 2s. 2p, orbital l = 1, m = 0 Z Z X X axe d'-y f-orbitals An f subshell has seven orientations in space, i.e. there are seven f orbitals. However, the shapes off orbitals are very complicated Quick check 2.4 (a) What does an orbital represent according to the wave mechanical model of atom? Ans. According to wave mechanical model of an atom, an orbital shows the space orientation and shape. Like s-orbital is spherical while p-orbitals are dumb-bell (polar), d- orbitals are double dumb-bell while f-orbitals are triple dumb-bell or complicated. (b) There are three orientations of p-orbital due to three values of magnetic quantum number. Justify it. Ans. The existence of three orientations of p orbitals is justified by the magnetic quantum number (m), which is one of the four quantum numbers used to describe the state of an electron in an atom. Here's the explanation: Magnetic Quantum Number (m): • Determines the orientation of the orbital in space. • The possible values of m for p-orbital ranges from -1 to +1 in integer steps. For a p-orbital, l =1, so: m = - 1,0,+1 These three values correspond to three different spatial orientations of p orbitals, which are: Px, Py, pz. Each of these orbitals is oriented along a different axis in three-dimensional space, justifying three distinct orientations for p orbitals. ELECTRONIC CONFIGURATION

Definition: Electronic configuration is the distribution of electrons among available shells, subshells, or orbitals of an atom or ion. Example: The configuration of , Be atom with two electrons in the '1s" subshell and two electrons in the '2s' subshell is written 15'2s? Explanation • In case of subshells, the electronic configuration is described by a notation that lists the subshell symbols, one after the other. • Each symbol has a superscript on the right. This gives the number of electrons in the subshell. • Each group of orbitals in a subshell is labeled by its subshell notation. An electron in an orbital is shown by an arrow. The arrow points upward, when s = +½ and downward 1 when s = -- • The orbital diagram of boron ("'B) is as follows. "'B=1s" 2s" 2p} 2p, 2p; Distribution of electron in shells The electronic configuration of an atom describes the distribution of electrons in its atomic shells. The shells, denoted as K, L, M, N, and so on, correspond to the principal quantum number (n) of the orbitals. Shell capacities: Each shell has a specific capacity for electrons: K shell (n=1): 2 electrons maximum (1s orbital) L shell (n=2): 8 electrons maximum (2s and 2p orbitals) M shell (n=3): 18 electrons maximum (3s, 3p, and 3d orbitals) N shell (n=4): 32 electrons maximum (4s, 4p, 4d, and 4f orbitals) For example Hydrogen (H): 1s' (K shell) Sodium (Na): 1s? 2s 2p6 3s' (K L M) 2 8 1 Distribution of Electrons in sub shells orbitals In case of subshells, the electronic configuration is described by a notation that lists the subshell symbols, one after the other. Each symbol has a superscript on the right. This gives the number of electrons in the subshell. Following rules are applied to fill the orbitals of mufti-electron atoms. Aufbau principle: Aufbau principle is also known as the building up principle. This principle says that the subshells in an atom are filled with electrons in an increasing order of their energy values. Since, the energy of a subshell in the absence of any magnetic field, depends upon the principal quantum number (n) and the azimuthal quantum number (l); n+l Role: The order of filling subshells with electrons may be obtained from the summation (n + l). values are given in Table. Table: (n +l) values of various sub-shells. Azimuthal Principal quantum no. (n) quantum no. (() 0 0 2 1 0 1 3 2 0 4 2 3 Helium (He): 1s? (K shell) Chlorine (Cl) = 1s' 2s 2p6 3s 3p (K LM) 28 7 (n + l) Subshell Value 1s (1 + 0) = 1 2s (2 + 0) = 2 2p (2 + 1) = 3 3s (3 + 0) = 3 3p (3 + 1) = 4 3d (3 + 2) = 5 4s (4 + 0) = 4 (4 + 1) = 5 4d (4 + 2) = 6 4f (4+3) = 7 (a) The subshell having lower (n • + l) value has lower energy and is filled first. Example: 4s orbital has (n +* l) = 4 + 0 = 4 and 3d orbital has (n + l) = 3+2 = 5. Since (n + l) value of 4s orbital is i 4s lower than that of 3d, hence 4s lower subshell has energy than 3d and 4s will be filled first. Orbital filling sequence (b) In case there are two subshells are having equal (n + l) lower Values, then the subs bellied first. Example: Both 4p and 3d Fig: Arrangement of subshells in increasing energy subshells have n + l value = 3+2=5: 3d subshell wıll be preferred to be tilled because or its down valu According to this rule the energy wise arrangement of orbitals should be. So, the order of filling of various subshells with electrons obtained by this rule is given below figure. Example of Be: The electronic configuration of ¿Be atom with two electrons on the 1s subshell and two electrons in the 2s subshell is written 1s' 2s Example of Na: The electronic configuration of 2}Na is 1s' 2s 2p° 3s' electronic configuration of different elements is provided in table. Table 2.6 Electronic configuration of ground states of elements Z= 1 - 36. Z Z Element Configuration H : 1 19 Is' 1s He 20 Li 1s'2s' 21 4 Be 22 1s'2s? 5 B 23 1s'2s2p' 6 C 24 1s 252p2 7 N 25 1s'2s 2p3 8 26 1s'2s 2p4 F 9 1s'2s 2p5 27 10 Ne 28 11 Na 29 •5f 6s 5p 5s •3p 12p 2s Iss order . The subshell Element Configuration 1s'2s 2p°3s'3p'4s' Ca 1s2s2pº3s?3pº4s? Sc 1s'2s 2p°3s'3p°3d'4s? Ti V 1s2s2p°Зs?3pº3d4s? Cr 1s'2s2p63s3p63d4s' Mn 1s2s2p°3s'3p°3d4s? Fe Co Ni Cu 1s:252p63s? 12 Mg 31 Al 1s 25-2p'3533p' 13 32 Si 14 1s'2s 2p°353p2 33 1s 252p°3523p3 15 34 is 252p635 3p* S 16 35 Cl 1s'2s 2p635'3p5 17 36 Ar 1s 2s 2p6353p° 18

Distribution of electrons in orbitals A useful way of representing electronic configurations is a box diagram. Each box represents an orbital as shown in figure. • Each box represents an atomic orbital. • Each orbital con occupy maximum of two electrons. An electron is represented by an arrow. • The boxes (orbitals) can be arranged in order of increasing energy from bottom to top The following rules are obeyed while filling orbitals with electrons. Thus it takes three quantum numbers to describe an orbital but forth quantum number to differentiate between the two electrons that can occupy an orbital. Pauli's Exclusion principle: Statement: No two electrons in an atom can have the same values for all the four quantum numbers, or "Two electrons in an orbital will always have opposite spins" Example: In the first shell of helium (He) atom, there are two electrons. They are present in 1s orbital. According to the concept of quantum numbers and Pauli's exclusion principal, the values of their quantum numbers are: Table 2.7 Values of quantum numbers of two electrons in the same orbital n Electron 1 0 Electron 1 1 Electron 2 The two electrons having the same values of 'n', 'l' and 'm' can have different values of 's. It means that their spins are in the opposite directions. Hund's rule Statement: When degenerate orbitals are available and more than two electrons are to be placed in them, they should be placed in separate orbitals with the same spin rather than in the same orbital with opposite spins. Example: The three p-orbitals, i.e., Px› Py and p, have equal energy. To understand it, let 1s'2s 2p'353p°3d'°45 Zn Ga 1s'2s2p°3s3p°3d'°4s'4p' Ge 1s'2s'2p°3s 3p'3d'°4s'4p2 1s'252p°3s3p"3d'°454p} As Se 1s'2s2p°3s'3p°3d'°4s4p4 1s°2s2p"35'3p63d'°4s'4p) Br 1s°2s2p"3s'3p"3d°454p Kr 2p 2pz 2py 2px 2s 1L 1s Fig: The electronic configuration of boron in terms of orbitals. m +1 (anti clockwise) 2 0 (clockwise) 2 us take an example in which three electrons are to be filled into three p-orbitals. There are two different ways to do this as shown below: Px Py Pr 1L1L I (wrong) According to Hund's rule 2nd one is right, configuration of p-orbital. Explanation: This rule gives an idea for filling electrons into the orbitals having equal energies. According to the Hund's rule, the correct way of filling three electrons in three p orbitals is that in which each orbital is singly occupied Electronic configuration in terms of orbitals in carbon, nitrogen, oxygen is provided in figure. 2P 1 2px 2px 2pz 2py 2s 1L 1L Is 1L 1L carbon nitrogen 18 2s 2p2 1s 2s 2p3 Fig: Electronic configuration in terms of orbitals in carbon, nitrogen and oxygen. The electron configurations of some elements of the periodic table in the light of the various principles of electronic configuration are given in table 2.6. ELECTRONIC CONFIGURATION AND THE PERIODIC TABLE

The arrangement in group The electronic configurations of elements are related to their position in the periodic table. The periodic table is structured so that elements with the same pattern of outer-shell (valence) electronic configuration are arranged in same groups. Location of Elements You can easily write the electronic configuration of an element based on its location in the periodic table. Notice that in the periodic tablé the elements can be grouped in terms of the type of orbital into which the electrons are placed. s-Block elements • On the left are two columns of elements. • These elements, known as the alkali metals and alkaline earth metals (groups 1 and 2), • Those in which the outer-shell s orbitals are being filled • The group 1 and 2 elements all have ns' and ns outer configurations respectively. p-Block elements • When we jump to group 13 elements, we find that they have ns'np' configuration. • On the right is a block of six columns. • These are the elements in which the outermost "p' Px Py Pr 111 II (right) 1 1 1 2py 2px 2pz 2py 2pz 2py 1L 1L Oxygen 15 2s 2p4 orbitals are being filled. Transition elements (d & f-block elements): • In the middle of the table is a block of ten columns that contain the transition metals. • In these, the d orbitals are being filled. • Below the main portion of the table are two rows that contain fourteen columns. • These elements are often referred to as the f-block elements. These are the ones in which the 'f orbitals are being filled. • Recall that the numbers 2, 6, 10, and 14 are precisely the number of electrons that can fill the s, p, d, and f subshells, respectively. Quick check 2.5 (a) With the help of periodic table, write the electronic configuration for the following elements by giving the appropriate noble-gas core plus the electrons beyond it: (i) 48 Cd, (i) syLa Ans. Cd= [Kr] 5s?, 4d° 48 Cd: 57. La: La = [Xe] 6s, 4f' (b) Write the complete electron configuration for antimony (Sb) with atomic number 51. 1s 2s 2p" 3s 3p° 45 3d'° 4p' 5s 4dl°5p3 Ans. Sb: (c) How many unpaired electrons are there in each atom of siSb. Ans. As per given electronic configuration of sySb., three unpaired electrons present in valence orbitals which are: 5p, 5p5p2: (d) Write down the electronic configuration of the following in terms of orbitals: (ii) 14Si (i) 13Al (wii) 28Ni 13Al: 1s 2s 2p' 3s 3p' 14Si: 1s 2s 2p° 3s 3p? 28 Ni: 1s 2s 2p6 3s 3p° 45 308 VALENCE ELECTRONS

Valence electrons Definition: The electrons in an atom in the outermost shell are called valence electrons. Explanation: These are such electrons that are primarily involved in chemical reactions. The similarities among the configurations of valence electrons account for similarities of the chemical properties among groups of elements. Figure shows a periodic table which include the valence shell configurations. Note the similarity in electron configuration within any group (column) of elements. 25 11 12 Na Mg 352 35' 21 26 24 25 20 27 22 19 23 Cr V Ca Sc Fe Mn Co Ti 4523d3 4523d 4s23d2 4523d° 4523d? 4s'3d5 4s? 4s' 43 41 42 40 39 37 44 45 38 Ru Sr. Tc Rh Y Rb Nb Zr Mo 5s^4d2 5s'4ds 5s'4d 5s24d' 5s'4d 5s'4d4 5s'4d 5s' 552 74 76 55 75 57 72 73 56 77 W Re Re Cs Hf Ва Та La Ir 6s'4f*5d' 65'4f45d4 6s'4f*5d5 65 6s'5d' 6s'4F45d? 6s' 89 104 105 106 107 108 109 88 87 Db Hs Ac Ra Rf Bh Mt Sg Fr 751 7s2 61 62 66 64 58 60 59 67 63 65 68 Pr ть Ho Er Ce Sm Pm Dy 63:40 65°40 694P 6544P 63°470 65-41 6s'4F 5d* 96 91 92 93 90 94 100 98 95 Pa Am Th Cm NP Pu Ts'5r 78°501 75°62 7s:55 797500 Fig: The electronic distribution of elements of modern periodic table Classification of elements in periodic table • The main-group or representative elements all have valence-shell configurations ns np'. • They have some choice of 'a' and "b' • In other words, the outer 's' or 'p' subshell is being filled • In the d-block transition elements, a 'd' subshell is being filled. • In the f-block transition elements or inner-transition elements, an 'f subshell is being filled. • There is a definite pattern to the order of filling of the subshells as we go through the elements in the periodic table. • From this we can write down the building-up order. The overall sketch of periodic table is given in figure, shows the blocks of elements. 1s 2s 3s - 3d 4s 5s 4d * 6s 5d ** 6d * ** Fig: Modern periodic table showing s, p, d and f-block elements. He 5 9 10 7 Ne F • 00 N 12s22p 252p? 2s2p' 252p* 2s'2p" 16 18 17 13 14 15 Ar P AI SI 3523p4 3s23p5 3s?3p 3s73p5 3s23p' 3573p? 36 31 28 29 30 32 33 34 35 As -Br Kr Se Ge NI Zn Cu Ga 4s"3d*4p* 14s 3d*4p: 4s 3d 94p* 4523d8 45'3d1° 4s23d'° 4s*3d*4p*| 4s*3d*°4p? 4s*3d"4p? 48 45 53 49 54 47 52 50 51 Xe Te Pd Cd Sb Sn In 4d10 5s-4d105p) 55445058340195( 5634d195p* 5s24d105p 5540l5540'0554095p 86 80 78 85 81 84 82 83 79 At Rn PЬ TI Po Au Pt Hg 115 111 1:12 114 116 118 110 113 117 Ts LV Mc Og Ds FI Cn 71 69 70 Yb Lu 65*4(45d' 684Г 101 102 103 7s°5г* Is 2p 3 p 4 p 5p 6p 4f 5f Figure shows a periodic table indicating this pattern. • In figure the extreme left area, an 'ns' subshell is being filled. • The extreme right area, an np subshell is being filled • In the middle part, an (n - 1) d subshell is being filled. • At the bottom of figure an (n-2)f subshell is being filled. Building up order • Building-up order is starting with the first period, in which the "Is' subshell is being filled. • In the second period, we have '2s'; then staying in the same period but jumping across, we have "2p' • In the third period, wẹ have '3s' and "3p'; in the fourth period, '4s', '3d' and then ' 4p'. • This pattern should become clear enough to visualize with a periodic table. Categories of elements: The classification of elements of periodic table has been done into metals, non-metals, representative, transition elements, periods, groups and blocks (s, p, d, f). This division is done on the basis of electronic configuration and their valence electrons. The chemical properties of various categories of elements discussed above can be easily assessed. Quick Check 2.6 (a) An element has the electronic configuration 1s' 2s' 2p° 3s' 3p° 3d1° 4s 4p6 4dl° 5s' 5p.. (i) Which block in the periodic table does this element belong to? Ans. According to the given electronic configuration, the valence electrons are present in p-subshell, so it belongs to p-block. (i) Which group does it belong to? Ans. It belongs to the p-block and 5p' electrons shows that it belongs to 5th group that is halogen group. Halogen is 5t group of p-block. (wii) Which period does it belong to? Ans. It belongs to the 5'h period, as it has maximum numbers 5s & 5p that shows period number, (iv) Identity this element. Ans. The element, that belongs to halogen and has atomic number 53 is Iodine. (b) Which block in the periodic table does the element with the electronic configuration 1s' 2s' 2p6 3s' 3pó 3df 4s' belong to? Name it. Ans. The given electronic configuration is: 1s 2s 2p° 352 3p' 3d 4sl As it has half-filled d-subshell, so it means it belongs to d-block. One electron in 4s is due to the stability of d-subshell with half-filled rule. The element is Cr-24. ELECTRONIC CONFIGURATION OF IONS AND FREE RADICALS

Electronic configuration of ions Positive ions: Positive ions are formed when electrons are removed from atoms. The sodium ion, Nat (proton number = 11), has 10 electrons. So, its electronic configuration is 1s? 2s 2p° Note that this is the same as the electronic configuration of neon, the element with 10 electrons in each atom. Negative ions: Negative ions are formed when atoms gain electrons. The sulfide ion, Sở- (proton number = 16), has 18 electrons. Its electronic configuration is Is 2s' 2p! 3s? 3p', which is the same as argon, the element with 18 electrons in each atom. Ions of transition elements: In general, electrons in the outer subshell are removed when metal atoms form the positive ions. However, the d-block elements behave slightly differently. Reading across the periodic table from potassium to zinc, the 4s subshell fills before the 3d subshell. But when atoms of a d-block element lose electrons to form ions, the 4s electrons are lost first. Examples Ti: 1s' 2s 2p° 3s 3p' 3d' 4s? Tit ion: 1s 2s 2p° 3s 3p° 3d2 Cratom: 1s 2s 2p° 3s 3p° 3d5 4s' Crit ion: 1s 2s 2pб 3s 3p° 303 Free radicals Definition: "A free radical is a specie that has one or more unpaired electrons". An example of a simple free radical is free chlorine atom :Cl: The electronic configuration of this radical is 1s' 2s 2p° 3s' 3p5. In the 3p-subshell, two orbitals have paired electrons whereas, the third one contains a single unpaired electron. The unpaired electron is shown by a single dot as in Cl. Apart from single atoms, group of atoms can also be free radicals. Example: Some of free radicals are OH®, CH3° etc. Quick Check 2.7 Write electronic configuration for the following ions and free radicals: Ans. (i) Al® (Z=13) Al" : 1s 2s 2p6 3s 3p' 02: 15 2s 2p6 (ii) O^ (Z= 8) (üi) Fe*+ (Z=26) Fe3+: 1s 2s 2p° 3s 3p° 3d5 (iv) Cu?+ (Z= 29) Cu2+: 1s 2s 2p° 3s 3p° 3d° (V) Cu° (Z=29) Cự : 1s 2s 2p6 3s 3p' 3d1° 4s! ELECTRONIC CONFIGURATION AND THE FORMATION OF SEMI- CONDUCTORS

Semiconductors Semiconductors are materials that can conduct electricity under some conditions. They are used in many electronic devices, including smartphones, laptops, and cars. Example: The elements that can act as are semiconductors silicon, arsenic etc. The germanium and formation of semiconductors is possible because of a unique electronic configuration of these elements. Consider the example of Si and explore how it can be converted into a p-type and n-type semiconductors. Silicon as non-semiconductor: The electron configuration of 14Si = 2, 8, 4; meaning that it has 4 electrons in its valence shell. In the pure crystalline form, each Si atom bonded to four other Si atoms. In this form, there is no possibility of electronic conduction through the Si crystal. P-type and N-type semiconductors are formed by "doping" a pure semiconductor material with impurity atoms, where adding trivalent impurities (like boron or aluminum) creates a P-type semiconductor, while adding pentavalent impurities (like phosphorus or arsenic) creates an N-type semiconductor. The difference lies in whether the added impurity creates "holes" (positive charge carriers) in the lattice, leading to P-type, or extra electrons (negative charge carriers) leading to N-type. P-type semiconductor formation Impurity atoms with three valence electrons (like Al) are added to the pure semiconductor. Some of the trivalent atoms take place of the Si atoms in the crystals. The silicon atoms cannot make four bonds due to the lack of electrons. For this reason, there are created holes in the crystal lattice, which act as positive charge carriers as shown in figure. This process creates a positive-type semiconductor or P-type semiconductor. Electrons from an external current source can move through the semiconductor and it can act as a conductor. N-type semiconductor formation: When impurity atoms with five valence electrons (like phosphorus) are added to the pure semiconductor, some of Si atoms are replaced with the pentavalent phosphorus atoms. The Si atoms in the vicinity of these atoms can make four bonds and the fifth electron is an extra electron. These impurity atoms contribute extra electrons to the crystal lattice, which become free to move and act as negative charge carries as in figure. The result is an N-type semiconductor that can conduct electricity when connected to an external source. Covalent bond Silicon valence electrons Silicon atoms Fig: A crystal of pure silicon Complete covalent bond AL C Incomplete Presence of covalent.bond hole • Si = Interinsic semiconductor atom Al = Trivalent impurity atom Formation of P type extrinsic semiconductor Fig: Doping and the formation of P-type and N-type semiconductor free i'm. Complete covalent bond Free electron Si = Interinsic semiconductor atom • P= Pentavalent impurity atom Formation of N type extrinsic semiconductor (b)