Periodic Table & Periodicity

Periodic Table & Periodicity S.No. 1. 2. 3. Long form of the Periodic Table or Moseley’s Periodic Table or Modern Periodic Table : Introduction DISCRIPTION Proposed by Moseley (i) In the long form of periodic table there is contribution of Ramsey, Werner, Bohr and Bury. (ii)This table is also referred to as Bohr's table since it follows Bohr's scheme Contribution of the arrangements of elements into four types based on electronic configuration of elements The modern periodic table consits of horizontal rows (periods) and vertical column (groups). Based on Atomic number (i) Moseley did an experiment in which he bombarded high speed electrons on different Metal surfaces and obtained X-rays(electromagnetic rays). He observed regularities in the characteristic X-ray spectra of the elements and found that plot  vs. Z (atomic number) is straight line while  vs. A (atomic weight) is not, and  = a(Z – b), where a and b are constants that are same for all elements and  is frequency of X-rays. Thus he concluded that atomic number is more fundamental property than atomic weight. 4. 5. 6. 7. ⚫ ⚫ ⚫ ⚫ Experiment (ii) Moseley concluded that the physical and chemical properties of the elements are periodic function of their atomic number. The physical and chemical properties of elements are periodic function of their Modern Periodic atomic number. So the elements are arranged in order of increasing atomic Law number, the elements with similar properties comes after regular intervals. The repetition of the properties of elements after regular intervals when the elements are arranged in the order of increasing atomic number is called periodicity. Periodicity (a) In a period, the ultimate orbit remain same, but the number of electrons gradually increases. (b) In a group, the number of electrons in the ultimate orbit remains same, but the values of n increases. The periodic repetition of the properties of the elements is due to the recurrence of similar valence shell electronic configuration after certain regular Cause of intervals. For example, alkail metals have same electronic configuration ns1, Periodicty therefore, have similar properties. In the periodic table, elements with similar properties occur at intervals of 2, 8, 8, 18, 18 and 32. These numbers are called as magic numbers. The modern periodic table consists of horizontal rows (periods) and vertical column (groups). Periods: There are seven periods numbered as 1, 2, 3, 4, 5, 6 and 7. Each period consists of a series of elements having same valence shell. Each period corresponds to a particular principal quantum number of the valence shell present in it. Each period starts with an alkali metal having outermost electronic configuration as ns 1. Each period ends with a noble gas with outermost electronic configuration ns 2np6 except helium having outermost electronic configuration as 1s2. Periodic Table & Periodicity ⚫ ⚫ Each period starts with the filling of new shell. The number of elements in each period is twice the number of atomic orbitals available in shell that is being filled. For illustration–  st period shortest period having only two elements. Filling of electrons takes place in the first shell, for which, n = 1, = 0 (s-subshell) and m = 0. Only one orbital (1s) is available and thus it contains only two elements.  3rd period (short period) having only eight elements. Filling of electrons takes place in the third shell. For which, n = 3, = 0, 1, 2 and number of orbitals 1 3 5 (3s) (3p) (3d) –––––––––––––––– Total number of orbitals 9 –––––––––––––––– But the energy of 3d orbitals are higher than 4s orbitals. Therefore, four orbitals (one 3s and three 3p orbitals) corresponding to n = 3 are filled before filling in 4s orbital (next energy level). Hence 3rd period contains eight elements not eighteen elements. Groups : There are eighteen groups numbered as 1, 2, 3, 4, 5, ........... 13, 14, 15, 16, 17, 18. Group consists of a series of elements having similar valence shell electronic configuration. Table-1 Periods (1)st n = 1 (2)nd n = 2 (3)rd n = 3 (4)th n = 4 (5)th n = 5 (6)th n = 6 (7)th n = 7 S–Block Number of elements 2 8 8 18 18 32 32 Called as Very short period Short period Short period Long period Long period Very long period Very long period p–Block Elements Elements 1 IA 18 VIII A 1 H 1.007 2 II A 3 Li 6.941 11 Na 22.98 19 K 39.08 37 Rb 85.46 55 Cs 132.90 4 Be 9.012 12 Mg 24.30 20 Ca 40.078 38 Sr 87.62 56 Ba 137.27 87 Fr 223 88 Ra 226 d –Block Elements 3 III B 4 IV B 21 22 Sc Ti 44.959 47.88 39 40 Y Zr 88.905 91.224 57 72 La* Hf 138.905 178.49 89 Ac** 227 104 Rf 261.11 5 VB 6 VI B 7 VII B 23 V 50.9415 41 Nb 92.906 73 Ta 180.947 24 Cr 51.996 42 Mo 95.94 74 W 183.85 25 Mn 54.938 43 Tc 98 75 Re 186.207 105 Ha 262.114 106 Sg 263.118 107 Bh 262.12 8 VIII 9 VIII 10 VIII 26 27 28 Fe Co Ni 55.84 55.933 58.693 44 45 46 Ru Rh Pd 101.07 102.905 106.42 76 77 78 Os Ir Pt 190.2 192.22 195.08 108 Hs 265 109 Mt 266 110 Ds 269 11 IB 12 II B 29 30 Cu Zn 63.546 65.39 47 48 Ag Cd 107.868 112.411 79 80 Au Hg 196.666 200.59 111 Rg 272 112 Cn 285 13 III A 14 IV A 15 VA 16 VI A 17 VII A 2 He 4.002 5 B 10.811 13 Al 26.981 31 Ga 69.723 49 In 114.82 81 Tl 204.383 6 C 12.011 14 Si 28.085 32 Ge 72.61 50 Sn 118.710 82 Pb 207.2 7 N 14.006 15 P 30.973 33 As 74.921 51 Sb 121.757 83 Bi 207.980 8 O 15.999 16 S 32.006 34 Se 78.96 52 Te 127.60 84 Po 209 9 F 18.998 17 Cl 35.452 35 Br 79.904 53 I 126.904 85 At 210 10 Ne 20.179 18 Ar 39.948 36 Kr 83.80 54 Xe 132.29 86 Rn 222 113 Uut 284 114 Fl 289 115 Uup 288 116 Lv 292 117 Uus 118 Uuo 294 Inner - Transition Metals (f-Block elements) *Lanthanides **Actinides 58 Ce 140.115 59 Pr 140.907 60 Nd 144.24 61 Pm 145 90 Th 232.038 91 Pa 231 92 U 238.028 93 Np 237 62 63 Sm Eu 150.36 151.965 94 Pu 244 95 Am 243 64 Gd 157.25 65 Tb 158.925 66 Dy 162.50 67 Ho 164.930 68 Er 167.26 69 Tm 168.934 70 Yb 173.04 71 Lu 174.967 96 Cm 247 97 Bk 247 98 Cf 251 99 Es 252 100 Fm 257 101 Md 258 102 No 259 103 Lr 260 Periodic Table & Periodicity Classification of the Elements :                    It is based on the type of orbitals which receives the differentiating electron (i.e., last electron). (a) s-block elements When shells upto (n – 1) are completely filled and the last electron enters the s-orbital of the outermost (nth) shell, the elements of this class are called s-block elements. Group 1 & 2 elements constitute the s-block. General electronic configuration is [inert gas] ns1-2 s-block elements lie on the extreme left of the periodic table. This block includes metals. (b) p-block elements When shells upto (n – 1) are completely filled and differentiating electron enters the p-orbital of the nth orbit, elements of this class are called p-block elements. Group 13 to 18 elements constitute the p-block. General electronic configuration is [inert gas] ns2 np1-6 p-block elements lie on the extreme right of the periodic table. This block includes some metals, all nonmetals and metalloids. (c) d-Block elements When outermost (nth) and penultimate shells (n – 1)th shells are incompletely filled and differentiating electron enters the (n – 1) d orbitals (i.e., d-orbital of penultimate shell) then elements of this class are called d-block elements. Group 3 to 12 elements constitute the d-block. General electronic configuration is [inert gas] (n – 1)d1–10ns1–2 (except, palladium which has valence shell electron configuration 4d105s0). All the transition elements are metals and most of them form coloured complexes or ions. d-block elements are classified into four series as given below. (1) Ist transition series i.e. 3d series contains 10 elements and starts from 21Sc–30Zn. Filling of electrons takes place in 3d sub-shell. (2) IInd transition series i.e. 4d series contains 10 elements and starts from 39Y–48Cd. Filling of electrons takes place in 4d sub-shell. (3) IIIrd transition series i.e. 5d series contains 10 elements and starts from 57La, 72Hf–80Hg. Filling of electrons takes place in 5d sub-shell. (4) IVth transition series i.e. 6d series contains 10 elements and starts from 89Ac, 104Rf–112Uub. Filling of electrons takes place in 6d sub-shell (incomplete series). Those elements which have partially filled d-orbitals in neutral state or in any stable oxidation state are called transition elements. (Zn, Cd, Hg, Uub are not transition elements) (d) f-Block elements When n, (n – 1) and (n – 2) shells are incompletely filled and last electron enters into f-orbital of antepenultimate i.e., (n – 2)th shell, elements of this class are called f-block elements. General electronic configuration is (n – 2)f1–14(n – 1)d0-1ns2 All f-block elements belong to 3rd group. They are metals Within each series, the properties of the elements are quite similar. The elements coming after uranium are called transuranium elements. They are also called as inner-transition elements as they contain three outer most shell incomplete and were also referred to as rare earth elements since their oxides were rare in earlier days. The elements of f-blocks have been classified into two series. 1. st inner transition or 4 f-series, contains 14 elements 58Ce to 71Lu. Filling of electrons takes place in 4f subshell. 2. IInd inner transition or 5 f-series, contains 14 elements 90Th to 103Lr. Filling of electrons takes place in 5f subshell. The actinides and lanthanides have been placed at the bottom of the periodic table to avoid the undue expansion of the periodic table. Periodic Table & Periodicity Ex-1 Sol. Elements A, B, C, D and E have the following electronic configurations : A : 1s2 2s2 2p1 B : 1s2 2s2 2p6 3s2 3p1 C : 1s2 2s2 2p6 3s2 3p3 2 2 6 2 5 2 2 6 2 6 2 D : 1s 2s 2p 3s 3p E : 1s 2s 2p 3s 3p 4s Which among these will belong to the same group in the periodic table ? Out of these, elements A and B will belong to the same group of the periodic table because they have same outer electronic configuration, ns2 np1. Prediction of period, group and block:    (a) (b) (c) (d) Period of an element corresponds to the principal quantum number of the valence shell. The block of an element corresponds to the type of subshell which receives the last electron. The group is predicted from the number of electrons in the valence shell or/and penultimate shell as follows. For s-block elements, Group number = the number of valence electrons For p-block elements, Group number = 10 + number of valence electrons For d-block elements, Group number = number of electrons in (n – 1) d sub shell + number of electrons in valence shell. For f block element – (58-71) 4f   - All f block elements are belongs to group number 3. (92-103) 5f  Two exceptions to this categorization: Strictly, helium belongs to the s-block but its positioning in the p-block along with other group 18 elements in justified because it has a completely filled valence shell (1s2) and as a result, exhibits properties & characteristics of other noble gases. Hydrogen has only one s-electron and hence can be placed in group 1 (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 (halogen family) elements. Because it is a special case, we shell place hydrogen separately at the top of the periodic table. Isotopes in modern periodic table: Periodic Table & Periodicity As isotopes of a elements have same atomic number and in modern periodic table the elements have been allotted places based on their atomic number. All the isotopes an elements have been assigned the same position in the modern periodic table. Ex-2 Sol. An element X with Z = 112 has been recently discovered. What is the electronic configuration of the element ? To which group and period will it belong ? (a) The electronic configuration of element X is [Rn]86 5f14 6d107s2 (b) It belongs to d-block as last electron enters in d subshell. (c) As number of electrons in (n – 1)d subshell and valence shell is equal to twelve i.e. 10 + 2. So it belongs to group 12. (d) It belongs to period 7 of the periodic table as principal quantum number of valence shell is 7 (i.e., 7s2). Metals and nonmetals :   The metals are characterised by their nature of readily giving up the electron(s) and from shining lustre. Metals comprises more than 78% of all known elements and appear on the left hand side of the periodic table. Metals are usually solids at room temperature (except mercury, gallium). They have high melting and boiling points and are good conductors of heat and electricity. Oxides of metals are generally basic in nature (some metals in their higher oxidation state form acid oxides e.g. CrO 3). Nonmetals do not lose electrons but take up electrons to form corresponding anions. Nonmetals are located at the top right hand side of the periodic table. Nonmetals are usually solids, liquids or gases at room temperature with low melting and boiling points. They are poor conductors of heat and electricity. Oxides of nonmetals are generally acidic in nature. Metalloids (Semi metals) :    It can be understood from the periodic table that nonmetallic character increases as we move from left to right across a row. It has been found that some elements which lie at the border of metallic and nonmetallic behavior, possess the properties that are characteristic of both metals and nonmetals. These elements are called semi metals or metalloids. The metalloids comprise of the elements Si, Ge, As, Sb, Se and Te. Oxides of metalloids are generally amphoteric in nature. Diagonal relationship : Some elements of certain groups of 2nd period resemble much in properties with the elements of third period of next group i.e. elements of second and third period are diagonally related in properties. This phenomenon is known as diagonal relationship. For example, the similarity between lithium (the first member of group 1) and magnesium (the second element in group 2) is called a diagonal relationship. Diagonal relationship also exist between other pairs of elements Be and Al, B and Si as shown in figure; 2nd period Li Be B C 3rd period Na Mg Al Si Diagonal relationship arises because of ; (i) on descending a group, the atoms and ions increase in size. On moving from left to right in the periodic table, the size decreases. Thus on moving diagonally, the size remains nearly the same. (Li = 1.23 Å & Mg = 1.36 Å; Li+ = 0.76 Å & Mg2+ = 0.72 Å) (ii) it is sometimes suggested that the diagonal relationship arises because of diagonal similarity in electronegativity values. (Li = 1.0 & Mg = 1.2; Be = 1.5 & Al = 1.5; B = 2.0 & Si = 1.8) (iii) Be and Al also show a diagonal relationship. In this case sizes are not so close (Be 2+ = 0.45 Å and Al3+ = 0.535 Å) but the charge / ionic radius ratio is nearly similar because the charges are 2+ and 3+ respectively. (Ionic charge) Charge /Ionic radius ratio = (Ionic radius) Periodic Table & Periodicity Shielding Effect and Zeff The periodicity of atomic properties : (1) Effective nuclear charge (Zeff): The outer electron experiences two opposing force : (i) Force of attraction from nucleus. (ii) Force of repulsion from inner electrons. Suppose inner electrons are not present there then force of attraction experience by test electron must be greater than it actually experienced because electronic repulsion weakens the force of attraction. This reduced nuclear charge or nuclear actually experienced by an electron is termed as effective nuclear charge. Zeff = Z −  Zeff = Effective nuclear charge Z = Atomic number  = Screening constant or shielding constant or slater constant Effective nuclear charge is not same for all the electrons present in an atom and it varies with distance between electron and nucleus. This effect in which inner electrons shield valence electron from the attraction due to nucleus is called shielding effect. It is also known as screening effect as inner electron create a protective screen between nucleus and valence electrons. Due to shielding effect valence electron experiences less attraction from the side of nucleus as if inner electrons are absent. s-orbitals have the largest screening effect for a given n value since s electrons are closer to the nucleus. Screening effect decreases as s-orbital > p-orbital > d-orbital > f-orbital. th n shell (valence shell ) th (n–1) shell (n–2) th shell e− Z Electronic Repulsion Force of Attraction e− e− Electronic Repulsion Test electron Periodic Table & Periodicity Table-2 List of common oxidation sate of an element in Periodic Table 1 1 H +1 –1 3 Li +1 18 2 He 2 3 Be +2 11 Na +1 12 Mg +2 19 K +1 20 Ca +2 37 Rb +1 55 Cs +1 13 5 B +3 14 6 C +4 15 7 N +5 –3 +2 +4 –4 etc. +3 +1 –3 0 etc. 15 P +5 +3 +1 –3 16 8 O +2 – 1/2 –1 –2 17 9 F –1 16 S +6 +4 +2 –2 17 Cl +7 +5 +3 +1 0 –1 35 Br +7 +5 +3 +1 –1 13 Al +3 14 Si +4 –4 31 Ga +3 32 Ge +4 –4 33 As +5 +3 –3 34 Se +6 +4 –2 38 Sr +2 49 In +3 +1 50 Sn +4 +2 51 Sb +5 +3 –3 52 Te +6 +4 –2 56 Ba +2 81 Tl +3 +1 82 Pb +4 +2 83 Bi +5 +3 84 Po 3 21 Sc +2 +3 4 22 Ti +2 +3 +4 5 23 V +2 +3 +4 +5 6 24 Cr +2 +3 +4 +5 +6 7 25 Mn +2 +3 +4 +5 +6 +7 8 26 Fe +2 +3 +4 +5 +6 9 27 Co +2 +3 +4 +5 10 28 Ni +2 +3 +4 11 29 Cu +1 +2 12 30 Zn +2 53 I +7 +5 +3 +1 0 –1 85 At 10 Ne 18 Ar 0 36 Kr +4 +2 0 54 Xe +8 +6 +4 +2 0 86 Rn Periodic Table & Periodicity * Bold mark oxidation number are general stable oxidation number of an element in compound state. Section (A) : Atomic and Ionic radius (1) Atomic radius : It is distance between outermost electron and nucleus. X-ray diffraction, electron diffraction method and nuclear magnetic resonance (NMR) spectrum methods are used to determine internuclear distance or bond length. Atomic radius depends on the type of chemical bond between atoms in a molecule. These are : 1. Covalent radius 2. Ionic radius 3. Metallic radius 4. Vander waal’s radius Type of S.No. DISCRIPTION radius 1. Covalent radius SBCR(Single Bonded Covalent Radius) Halogen Covalent radius is calculated in covalent compounds. Covalent compound’s form covalent bonds. Covalent bonds are formed by overlapping of atomic orbitals. Type of Covalent radius. d(X−X) / pm Molecule Structure Model (gas phase) Fluorine F2 Chlorine Cl2 Bromine Br2 F—F 143pm Cl — Cl 199pm Br — Br 228pm 143 199 228 d(X−X) / pm (solid phase) 149 198 227 Periodic Table & Periodicity Iodine I2 I—I 266 266pm 272 Covalent radius is the one-half of the internuclear distance between two singly bonded two nuclei (of like atoms). (i) In Homo atomic Molecules Internuclear distance of A–A(A2) molecule is (dA–A) and covalent radius is rA then dA–A = rA + rA dA–A = 2rA d A− A rA = 2 e.g. – In Cl2 molecule, internuclear distance is 1.98 Å so rCl = 1.98 = 0.99 Å 2 SBCR of O, N and C etc. elements can be determined by taking H 2O2, N2H4, C2H6 respectively. (a) For Hetero atomic molecule with not E.N. difference. dA–B = rA + rB (ii) Hetero atomic molecule 2. Example : For A–B molecule Electronegativities of A and B are approximately equal e.g. C–I E.N. of C and I are approx equal (2.5) internuclear distance of C–I is 2.13Å and rC is 0.77Å. Solution : dC–I = rC + rI  rI = 2.13 – 0.77 = 1.36Å (b) Heteroatomic molecule with  E.N. difference more If in a diatomic molecule electronegativities of A–B have more difference. Then actual bond length will be reduced. Ionic Radius (i) Cationic radius When an neutral gaseous atom loses electron it converts into cation. Radius = Atomic radius > Cationic radius Reason: After loosing electron number of electron reduces, but number of protons remains same, due to this Zeff increases, hence electrons pulls towards nucleus and atomic radius decreases, moreover after losing all the electrons from the outer most shell, penultimate shell becomes ultimate shell which is nearer to nucleus so size decreases. Size of cation  1 Magnitude of +ve charge or z eff Example (ii) Anionic Size of cation = Fe > Fe+2 > Fe+3 Size of cation = Pb+2 > Pb+4 Size of cation = Mn > Mn+2 > Mn+3 > Mn+4 > Mn+5 > Mn+6 > Mn+7 When neutral gaseous atom gains electron it converts into anion. Radius = Anionic radius > Atomic radius Reason: In an anion electron are more than protons so effective nuclear charge reduces, and inter electronic repulsion increases, which also increases screening effect. So distance between electron and nucleus increases and size of anion also increases. Sizeof anion  Magnitudeof − ve charge Example : Size of F– > F why ? Sol. radius Proton electron Z e  F 9 9 F– 9 10 9 =1 9 9 = 0.9 10 As Zeff of F– is less than F so size of F– > F Periodic Table & Periodicity (iii) Size of iso electronic species Li+ Na+ K+ Rb+ Be+2 Mg+2 Ca+2 Sr+2 N3– Al+3 Sc+3 Y+3 Cs+ Ba+2 La+3 O2– S2– Se2– Te2– H– F– Cl– Br– I– He Ne Ar Kr Xe Those species having same number of electron but different nuclear charge forms isoelectronic series. For isoelectronic species the atomic radius increases with decrease in nuclear charge. Ions K+ Ca2+ S2– Cl – Atomic number 19 20 16 17 electron 18 18 18 18 Example Order of radius = S2– > Cl– > K+ > Ca2+ Order of radius = P3– > S2– > Cl– Order of radius = N3– > O2– > F– > Na+ > Mg2+ > Al3+ Order of radius = Au+  Hg2+  Tl3+  Pb4+ Metallic radius 3. It is one-half of the nuclear distance between two adjacent metallic atoms in crystalline lattice structure. so metallic radius > Covalent radius. Metallic radius  1 Metallic bond strength More metallic radius → loose crystal packing → less bond strength (body centered packing) Less metallic radius → tight crystal packing FCC → High bond strength (Hexagonal close packing) Periodic Table & Periodicity Those atoms (like noble gases) which are not bonded with each other, experiences a weak attractive force to come nearer. The half of the distance between the nuclei of adjacently placed atoms in solid state of a noble gas is Vander Waal’s radius. Vander Waal radius > Covalent radius. Inert gas have only Vander Waal radius In molecules of nonmetals both covalent and Vander Waal radius exists. 4. Van der Waal’s radius Radius = Vander Waal radius > Metallic Radius > Covalent Radius Van der waal force of attraction  Molecular weight OR atomic weight (in inert gases) In a period from left to right Vander Waal radius decreses. In a group from top to bottom its values increases. 5. Factors affecting atomic size (i) Effective nuclear charge (ii) Number of shells (iii) screening effect (iv) Magnitude of +ve charge (v) Magnitude of -ve charge (vi) Bond order Atomic radius  1 Effective nuclear charge Atomic radius  No. of shells Example Li > Be > B > C > N > O > F Example Li < Na < K < Rb < Cs Atomic radius  Screening effect Atomic radius  1 Magnitudeof + vecharge Example Mn > Mn+2 > Mn+3 > Mn+4 Atomicradius  Magnitudeof − ve charge Atomic radius  1 Bond order Example O < O– < O–2 Example N–N > N=N > NN Periodic Table & Periodicity Table-3 Variation in a Period (left to right) : Variation in a Group (top to bottom) : Nuclear charge (Z) increases by one unit Nuclear charge (Z) increases by more than one unit Effective nuclear charge (Zeff) also increases Effective nuclear charge (Zeff) almost remains constant because of increased screening effect of inner shells electrons. But number of orbitals (n) remains constant But number of orbitals (n) increases. As a result, the electrons are pulled closer to the nucleus by the increased Zeff. 1 rn  Z* Hence atomic radii decrease with increase in atomic number in a period from left to right. The effect of increased number of atomic shells overweighs the effect of increased nuclear charge. As a result of this the size of atom increases from top to bottom in a given group.  Irregularties In the transition series (e.g. in first transition series), the covalent radii of the elements decrease from left to right across a row until near the end when the size increases slightly. On passing from left to right, extra protons are placed in the nucleus and the extra electrons are added. The orbital electrons shield the nuclear charge incompletely. Thus the nuclear charge attracts all the electrons more strongly, hence a contraction in size occurs. The radii of the elements from Cr to Cu, are very close to one another because the successive addition of d-electrons screen the outer electrons (4s) from the inward pull of the nucleus. As a result of this, the size of the atom does not change much in moving from Cr to Cu. Table-4  Ex-3 Sol. Ex-4 Sol. Element Sc Ti V Cr Mn Fe Co Ni Cu Zn Atomic radius (pm) 144 132 122 117 117 117 116 115 117 125 The elements of group 3 belonging to d-block show the expected increase in size as that found in case of s-block and p-block elements. However, in the subsequent groups that is 4 th onwards (upto 12th group), there is increase in size between the first and second member, but hardly any increase between the second and third elements of the given group. There are 14 lanthanide elements between lanthanum and hafnium, in which the antepenultimate 4f shell of electrons (exert very poor shielding effect) is filled. There is a gradual decrease in size of the 14 lanthanide elements from cerium to lutetium. This is called lanthanide contraction. This lanthanide contraction cancels out the normal size increase on descending a group in case of transition elements. Atomic radius of Li is 1.23 Å and ionic radius of Li + is 0.76 Å. Calculate the percentage of volume occupied by single valence electron in Li. 4 Volume of Li = × 3.14 × (1.23)3 = 7.79 Å3 (–Li = 1s22s1) 3 4 Volume of Li+ = × 3.14 × (0.76)3 = 1.84 Å3 (–Li+ = 1s2) 3  Volume occupied by 2s subshell = 7.79 – 1.84 = 5.95 Å3. 5.95  % Volume occupied by single valence electron i.e., 2s electron = ×100 = 76.4% 7.79 Select from each group the species which has the smallest radius stating appropriate reason. (a) O, O–, O2– (b) P3+, P4+, P5+ (a) O is having smallest radius. Anion is larger than its parent atom. Also the anion of the same atom with higher negative charge is bigger in size as compared to anion with smaller negative charge as proton to electron ratio decreases thus attraction between valence shell electrons and nucleus decreases. Hence electron cloud expands. (b) The ionic radius decreases as more electrons are ionized off that is as the valency increases. So the correct order is P5+ < P4+ < P3+. Periodic Table & Periodicity Ex-5 Sol. Mg2+ is smaller than O2– in size, though both have same electronic configuration. Explain ? Mg2+ and O2– both are isoelectronic i.e., have same number of electrons. But Mg 2+ having 12 protons in its nucleus exerts higher effective nuclear charge than O2– having 8 protons and thus valence shell as well as inner shells electrons are more strongly attracted by the nucleus in Mg 2+ resulting smaller size than O2–. Section (B) : Ionisation energy (2) Ionisation Energy : Ionisation energy (IE) is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a cation. (IE ) 1 M(g) ⎯⎯⎯ → M+(g) + e– ;   M+ (g) + IE2 ⎯⎯ → M2+ (g) + e– M2+ (g) + IE3 ⎯⎯ → M+3 (g) + e– st IE1, IE2 & IE3 are the I , IInd & IIIrd ionization energies to remove electron from a neutral atom, monovalent and divalent cations respectively. In general, (IE)1 < (IE)2 < (IE)3 < .............. because, as the number of electrons decreases, the attraction between the nucleus and the remaining electrons increases considerably and hence subsequent ionization energies increase. Ionization energies are determined from spectra and are measured in kJ mol–1, k Cal mol–1, eV (electron volt). Factors Influencing Ionisation energy Variation in ionization energies in a period and group may or not be regular and can be influenced by the following factors. (A) Size of the Atom : Ionisation energy decreases with increase in atomic size. As the distance between the outer most electrons and the nucleus increases, the force of attraction between the valence shell electrons and the nucleus decreases. As a result, outer most electrons are held less firmly and lesser amount of energy is required to knock them out. For example, ionisation energy decreases continuously in a group from top to bottom with increase in atomic size. But in 13th & 14th group the ionisation energy does not decrease continuously and order is as follows: For 13th group: B > Tl > Ga > Al > In For 14th group: C > Si > Ge > Pb > Sn (B) Nuclear Charge : The ionisation energy increases with increase in the nuclear charge. This is due to the fact that with increase in the nuclear charge, the electrons of the outer most shell are more firmly held by the nucleus and thus greater amount of energy is required to pull out an electron from the atom. For example, ionisation energy increases as we move from left to right along a period due to increase in nuclear charge. (C) Shielding or screening effect : The electrons in the inner shells act as a screen or shield between the nucleus and the electrons in the outer most shell. This is called shielding effect. The larger the number of electrons in the inner shells, greater is the screening effect and smaller the force of attraction and thus ionization energy (IE) decreases. (D) Penetration effect of the electron : The ionization energy also depends on the type of electron which is removed. s, p, d and f electrons have orbitals with different shapes. An s electron penetrates closer to the nucleus, and is therefore more tightly held than a p electron. Similarly p-orbital electron is more tightly held than a d-orbital electron and a d-orbital electron is more tightly held than an f-orbital electron. If other factors being equal, ionisation energies are in the order s > p > d > f. For example, ionisation energy of aluminium is comparatively less than magnesium because outer most electron is to be removed from 3p-orbital (having lesser penetration effect) in aluminium where as in magnesium it will be removed from 3s-orbital (having larger penetration effect) of same energy level. (E) Electronic Configuration: If an atom has exactly half-filled or completely filled orbitals, then such an arrangement has extra stability. Periodic Table & Periodicity    Ex-6 Sol. Ex-7 Sol. Ex-8 Sol. Ex-9 Sol. The removal of an electron from such an atom requires more energy then expected. For example, first ionisation energy of beryllium is greater than boron because beryllium has extra stable outer most completely filled outer most 2s orbital while boron has partially filled less stable outer most 2p-orbital. Be (Z = 4) 1s2, 2s2 B (Z = 5) 1s2, 2s2, 2p1 Similarly noble gases have completely filled electronic configurations and hence they have highest ionisation energies in their respective periods. Metallic or electropositive character of elements increases as the value of ionisation energy decreases. The relative reactivity of the metals in gaseous phase increases with the decrease in ionisation energy. The reducing power of elements in gaseous phase increases as the value of ionisation energy decreases. Amongst alkali metals, the lithium is strongest reducing agent in aqueous solution. From each set, choose the atom which has the largest ionization enthalpy and explain your answer with suitable reasons. (a) F, O, N (b) Mg, P, Ar (a) Fluorine (F) has the largest ionization enthalpy because in moving from left to right in a period, atomic size decreases and electrons are held more tightly. Since F has the smallest size and maximum nuclear charge. It has the largest ionization enthalpy among these elements. (b) Argon (Ar) has the largest ionization enthalpy as argon has extra stable fully filled configuration. First and second ionisation energies of magnesium are 7.646 eV and 15.035 eV respectively. The amount of energy in kJ needed to convert all the atoms of magnesium into Mg 2+ ions present in 12 mg of magnesium vapour will be ? [Given 1 eV = 96.5 kJ ml–1]. Total energy needed to convert one Mg atom into Mg2+ gas ion, = IEI + IEII = 22.681 eV = 2188.6 kJ mol–1.  12 mg of Mg = 0.5 × 10–3 mole.  Total energy = 0.5 × 10–3 × 2188.6 = 1.0943 kJ Ans. Following graph shows variation of ionization energies with atomic number in second period (Li–Ne). Value of ionization energies of Na(11) will be : (A) above Ne (B) below Ne but above O. (C) below Li (D) between N and O. Na is 3rd period element and is bigger than Li. The distance between the nucleus and outer most electron is more as compared to Li. Thus the outer most electron is loosely bound with nucleus and removal of electron is easier. So option (C) is correct. M(g) ⎯→ M+ (g) + e– ; M(g) ⎯→ M2+ (g) + 2e– ; Which is/are correct statement(s) ? (A) IE1 of M(g) is 100 eV (C) IE2 of M(g) is 250 eV. M(g) ⎯→ M+(g) + e– IE1 of M M+ ⎯→ M2+ + e– M ⎯→ M2+ + 2e– H = 100 eV. H = 250 eV. (B) IE1 of M+ (g) is 150 eV. (D) IE2 of M(g) is 150 eV. IE2 of M but IE1 of M+ (IE1 + IE2). Ans. (A,B,D) Periodic Table & Periodicity Section (C) : Electron gain enthalpy (3) Electron Affinity: Electron affinity is conventionally defined as the energy released when an electron is added to the valence shell of an isolated gaseous atom. F + e– ⎯→ F– Ea = 328 kJ mol–1 A positive electron affinity indicates that the ion X– has a lower, more negative energy than the neutral atom X. The second electron gain enthalpy, the enthalpy change for the addition of a second electron to an initially neutral atom, invariably positive because the electron repulsion out weighs the nuclear attraction. Exothermic O(g) + e–(g) ⎯⎯⎯⎯⎯ → O– (g) ; Endothermic O–(g) + e–(g) ⎯⎯⎯⎯⎯⎯ → O2– (g) ; Ea = 140.9 kJ mol–1 ; egH = – 140.9 kJ mol–1 Ea = – 744 kJ mol–1 ; egH = 744 kJ mol–1 An element has a high electron affinity if the additional electron can enter a shell where it experiences a strong effective nuclear charge. The electron gain enthalpy egH, is the change in standard molar enthalpy when a neutral gaseous atom gains an electron to form an anion. X (g) + e– (g) ⎯→ X– (g) Now a days both electron affinity and electron gain enthalpy terms are used. Both are same in magnitude but opposite in sign. E.A. = +ve (Exothermic):– egH < 0. E.A. = –ve (Endothermic):– egH > 0. The units of EA and Electron gain enthalpy are ev/atom or Kcal / mole or KJ / Mole. Electron gain enthalpy provides a measure of the ease with which an atom adds an electron to form anion. Electron gain may be either exothermic or endothermic depending on the elements. When an electron is added to the atom and the energy is released, the electron gain enthalpy is negative and when energy is needed to add an electron to the atom, the electron gain enthalpy is positive. Although the electron gain enthalpy is the thermodynamically appropriate term, much of inorganic chemistry is discussed in terms of a closely related property, the electron affinity, E a of an element which is the difference in energy between the gaseous atoms and the gaseous ions. Ea = E (X, g) – E (X–, g)  Group 17 elements (halogens) have very high negative electron gain enthalpies (i.e. high electron affinity) because they can attain stable noble gas electronic configuration by picking up an electron.  Across a period, with increase in atomic number, electron gain enthalpy becomes more negative because left to right across a period effective nuclear charge increases and consequently it will be easier to add an electron to a small atom.  As we move in a group from top to bottom, electron gain enthalpy becomes less negative because the size of the atom increases and the added electron would be at larger distance from the nucleus. In general the electron gain enthalpies for some third period element (e.g. P, S, Cl) are more negative than the corresponding second period members (e.g. N, O, F). This is due to the smaller size of the atoms of the second period elements which would produced larger electron-electron repulsions for the additional electron.   Group 1 e.g. H Group 0 e.g. H H –73 He +48 Li –60 O –141 F –328 Ne +116 Na –53 S –200 Cl –349 Ar +96 K –48 Se –195 Br –325 Kr +96 Rb –47 Te –190 I –295 Xe +77 Cs –46 Po –174 At –270 Rn +68 Group 16 e.g. H Group 17 e.g. H Noble gases have large positive electron gain enthalpies because the electron has to enter the next higher energy level leading to a very unstable electronic configuration. Negative electron gain enthalpy of O or F is less than S or Cl. This is due to the fact that when an electron is added to O or F, the added electron goes to the smaller n = 2 energy level and experiences significant repulsion from the other electrons present in this level. In S or Cl, the electron goes to the Periodic Table & Periodicity    larger n = 3 energy level and consequently occupies a larger region of space leading to much less electron-electron repulsion. Electron gain enthalpies of alkaline earth metals are very less or positive because the extra electron is to be added to completely filled s-orbitals in their valence shells. Nitrogen has very low electron affinity because there is high electron repulsion when the incoming electron enters an orbital that is already half filled. 1 (i) Electron affinity  Atomic size (ii) Electron affinity  Effective nuclear charge (zeff) 1 (iii) Electron affinity  Screening effect (iv) Stability of half filled and completely filled orbitals of a subshell is comparatively more and the addition of an extra electron to such an system is difficult and hence the electron affinity value decreases. Ex-10 Consider the elements N, P, O and S and arrange them in order of increasing negative electron gain enthalpy. Sol. Order of increasing negative electron gain enthalpy is N < P < O < S. For detail refer text. Ex-11 Why do halogens have high electron gain enthalpies (i.e. – egH) ? Sol. The valence shell electronic configuration of halogens is ns 2np5 and thus they require one electron to acquire the stable noble gas configuration ns 2np6. Because of this they have strong tendency to accept an additional electron and hence have high electron gain enthalpies. Ex-12 Which will have the maximum value of electron affinity O x ,Oy, Oz [x, y and z respectively are 0, –1 and –2] ? (A) Ox (B) Oy (C) Oz (D) All have equal. Sol. Being neutral atom oxygen will have higher electron affinity as there is electrostatic repulsion between additional electron and negative ion in case of O– and O2–. So option (A) is correct. Ex-13 The amount of energy when million atoms of iodine are completely converted into I – ions in the vapour state according to the equation, I (g) + e– (g) → I– (g) is 5.0 × 10–13 J. Calculate the electron gain enthalpy of iodine in terms of kJ mol –1 and eV per atom. Sol. The electron gain enthalpy of iodine is equal to the energy released when 1 mole of iodine atoms in vapour state are converted into I– ions. 5.0  10 –13  6.023  1023 =– = – 30.1 × 104 J = – 301 kJ. 106 –301 Electron gain enthalpy of iodine in eV per atom = = –3.12. 96.5 Ex-14 Account for the large decrease in electron affinity between Li and Be despite the increase in nuclear charge. Sol. The electron configurations of Li and Be are [He]2s 1 and [He]2s2, respectively. The additional electron enters the 2s orbital of Li but the 2p orbital of Be and hence is much less tightly bound. In fact, the nuclear charge is so well shielded in Be that electron gain is endothermic. Periodic Table & Periodicity Section (D) : Electronegativity (4)     Electronegativity: Electronegativity is a measure of the tendency of an element to attract shared electrons towards itself in a covalently bonded molecules. The magnitude of electronegativity of an element depends upon its ionisation potential & electron affinity. Higher ionisation potential & electron affinity values indicate higher electronegativity value. With increase in atomic size the distance between nucleus and valence shell electrons increases, therefore, the force of attraction between the nucleus and the valence shell electrons decreases and hence the electronegativity values also decrease. With increase in nuclear charge force of attraction between nucleus and the valence shell electrons increases and, therefore, electronegativity value increases In higher oxidation state, the element has higher magnitude of positive charge. Thus, due to more positive charge on element, it has higher polarising power. Hence with increase in the oxidation state of element, its electronegativity also increases. Charge on cation  electronegativity of the atom. The electronegativity also increases as the s-character in the hybrid orbitals increases. Hybrid orbital sp3 sp2 sp s-character 25% 33% 50% Electronegativity increases Electronegativity of some elements according to pauling scale. Elements Electronegativity H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Ne 0.9 P 2.1 S 2.5 Cl 3.0 Br 2.8 I 2.5 Table-5 Variation of electronegativity in a group Variation of electronegativity in a period On moving down the groups, Z increases but Zeff almost remains constant, number of shells (n) increases, rn (atomic radius) increases. Therefore, electronegativity decreases moving down the groups. While moving across a period left to right, Z, Zeff increases & rn decreases. Therefore, electronegativity increases along a period. There is no direct method to measure the value of electronegativity, however, there are some scales to measure its value. (a) Pauling’s scale : Linus Pauling developed a method for calculating relative electronegativities of most elements. According to Pauling * − E E  = XA – XB = 0.208 E. A −B A−A B −B EA-B = Bond enthalpy/ Bond energy of A – B bond. EA - A = Bond energy of A – A bond EB –B = Bond energy of B – B bond (All bond energies are in kcal / mol) * All bond energies are in kJ / mol. − E E  = X – X = 0.1017 E. A B A −B A−A B −B Mulliken’s scale : Electronegativity  (chi) can be regarded as the average of the ionisation energy (IE) and the electron affinity (EA) of an atom (both expressed in electron volts). IE + EA M = 2 [Paulings’s electronegativity P is related to Mulliken’s electronegativity M as given below. P = 1.35 (M)1/2 – 1.37 Mulliken’s values were about 2.8 times larger than the Pauling’s values.]* * Only for reference, student need not memorize it. (b) Periodic Table & Periodicity APPLICATION OF ELECTRONEGATIVITY : (a) (b) Nomenclature : Compounds formed from two nonmetals are called binary compounds. Name of more electronegative element is written at the end and ‘ide’ is suffixed to it. The name of less electronegative element is written before the name of more electronegative element of the formula. For example Correct formula Name (a) I+ Cl– Iodine chloride (b) Cl+ F– Chlorine fluoride (c) Br+ Cl– Bromine chloride (d) IBr Iodine bromide (e) OF2 Oxygen difluoride (f) Cl2O Dichlorine oxide Effect on bond length: When the  EN increases between two atoms in a molecule then bond length between that atoms decreases. Order of  EN:– HF > HCl > HBr > HI Order of bond length:– HF < HCl < HBr < HI