Inorganic Chemistry: Periodic Properties PDF

## Chapter 3: PERIODIC PROPERTIES It is necessary to have an adequate knowledge of the atomic properties before proceeding to understand the topics like chemical bonding, geometry and shapes of molecules and ions. These atomic properties show a regular variation in the periods and groups of the per...

## Chapter 3: PERIODIC PROPERTIES It is necessary to have an adequate knowledge of the atomic properties before proceeding to understand the topics like chemical bonding, geometry and shapes of molecules and ions. These atomic properties show a regular variation in the periods and groups of the periodic table. Therefore they are also referred to as periodic properties. Some of these properties discussed in this topic are namely: - Atomic and Ionic Radius - Ionization energy - Electron affinity - Electronegativity. ### 3.1 Atomic and Ionic Radius The size of an atom or ion is one of the most important properties of an element. The quantum theory of the atom does not result in precise atomic or ionic radii because the radial distribution curve falls off gradually with increasing distance from the nucleus. It is also not possible to have an isolated atom and the electron cloud is influenced by the presence of other atoms coming under its influence. But still chemists have developed many ways of defining size. The three operational concepts of these radii are: 1. Covalent radius 2. van der Waal's Radius 3. Ionic radius #### (1) Covalent radius: Covalent radius of a non-metallic element is defined as half of the internuclear distance of two neighbouring atoms of the same element in a molecule. Similarly the _metallic radius_ of a metallic element is defined as half the experimentally determined distance between the nucleus of the nearest neighbouring atoms. The radii of a metal atom determined from the covalent metal hydride and from the atomic volume of a metallic phase are found to be the same. Also the size of an atom in free state i.e. atomic radii and its size determined from a single non-polar covalent bond are almost the same. Therefore metallic and covalent radii are synonymous to atomic radii. Atomic radius is thus defined as the size of an atom in its free state. It is determined from the internuclear distance of a covalent molecule. The atoms are considered as hard incompressible spheres which are touching each other in a normal covalent bond. ##### A. Method of Determination of Atomic Radius: The internuclear distance can be very accurately determined by X-ray diffraction or spectroscopic studies. E.g the internuclear distance between H atom in H₂ molecule is 0.74 A°. Half of this distance, viz. 0.37 A° is the covalent radii of H. Thus if the electronegativity difference between atom A and B in a molecule A-B is very small then the internuclear distance between them (d) is the sum of the covalent radii of atom A and B i.e. dA-B = rA + rB Once the atomic radii of some elements are known, others can be calculated. However the bond length between two unlike atoms differing appreciably in their electronegativity is shorter than the sum of their covalent radii. This is due to the polarity developed in the bond, which results in additional electrostatic attraction between the atoms. Greater the difference in electronegativity of the combining atoms shorter will be the bond distance, and shorter will be the covalent radii obtained. Schomake and Stevenson (1941) proposed that the bond distance or internuclear distance d(A-B) between A and B when their electronegativity difference (χA-χB) is large is given by the relation: d(A-B) = rA + rB - 0.09(χA-χB) where rA is the normal covalent radius of A (equal to half of bond distance of A-A) and rB is covalent radius of B (equal to half of bond distance of B-B). One of the modifications of this equation suggested by Porterfield is as follows : dA-B = rA + rB - 0.07(χA-χB) ##### B. Trend in Periodic Table: ###### Variation in Period: In general the atomic radii (covalent radii) decrease on moving from left to right in a given period of main group elements. For example, consider the covalent radii of 2nd period (Table 3.1). The effective nuclear charge (Z) felt by the electron in the outermost shell (calculated by Slater's rule) goes on increasing from Li to F. Since the last electron in all these elements is added to the L shell, one would expect the size to remain the same. But the table reveals that the covalent radii goes on decreasing as the atomic number increases within the period. This is because the greater the Zeff the greater would be the force with which the electron is pulled towards the nucleus. Thus there is more and more contraction of both K and L shells towards the nucleus and hence the size goes on decreasing on moving from left to right across the period. As a result in a given period alkali metal is the largest atom and halogen is the smallest. The same trend is shown by 3rd and 4th period. In case of a noble gas though the Zeff increases, but the size does not decrease and is in fact found to be greater than halogen element. This is because the force holding atoms together in a noble gas is only weak van der Waal force. The radii of atoms thus calculated are van der Waal radii (discussed further) which are larger than the covalent radii of halogen atoms. Table 3.1: Covalent radii of Elements of Second Period of the Periodic Table | Element | Electronic Cont | Za for L electron | Covalent Radius | | :------- | :-------------------------------- | :--------------- | :-------------- | | Li | 1s²2s¹ | 1.30 | 1.34 | | Be | 1s²2s² | 1.95 | 0.90 | | B | 1s²2s²2p¹ | 2.60 | 0.82 | | C | 1s²2s²2p² | 3.25 | 0.77 | | N | 1s²2s²2p³ | 3.90 | 0.75 | | O | 1s²2s²2p⁴ | 4.55 | 0.73 | | F | 1s²2s²2p⁵ | 5.20 | 0.72 | ###### Variation in a group: The covalent radii increase on moving from top to bottom in a given group of the main group elements. E.g the trend shown by alkali metals is given in table 3.2. As we move down a group the Zeff and the number of principal shells goes on increasing. The size of the atom goes on increasing with increase in the number of electron shells. This effect is partially annulled by the contracting effect of increasing Zeff. But the effect of adding new shells is so large that it overcomes the contracting effect of increasing Zeff. Hence the atomic radius increases on moving from top to bottom in a group. Table 3.2: Atomic radii of Alkali Metals. | Alkali Metal | Electronic Conf. | No. of Shells | Atomic radii | | :----------- | :---------------- | :------------ | :------------ | | Li | 1s²2s¹ | 2 | 1.34 | | Na | [Ne]3s¹ | 3 | 1.54 | | K | [Ar]4s¹ | 4 | 1.96 | | Rb | [Kr]5s¹ | 5 | 2.11 | | Cs | [Xe]6s¹ | 6 | 2.55 | #### (ii) van der Waal's Radius: The van der Waal's radius is defined as one half of the distance between the nucleus of two adjacent identical atoms belonging to two neighbouring molecules of an element in the solid state. (fig 3.1). The name van der Waal's radius is used because the forces existing between the molecules are weak van der Waal's forces of attraction. These radii are obtained by the X-ray studies of various elements in the solid state. E.g As in fig 3.1, the internuclear distance between two chlorine atoms of the nearest neighbouring molecule is 3.60 A°, therefore its van der Waal's radius is 1.80 A°. The van der Waal's radius and covalent radius of some common elements are given in table 3.3. Table 3.3: The van der Waal's and Covalent radii of some Elements (in A°) | Element | van der Waal's Radius | Covalent Radius | | :------- | :------------------- | :-------------- | | H | 1.2 | 0.37 | | N | 1.5 | 0.75 | | O | 1.4 | 0.73 | | F | 1.35 | 0.72 | | Cl | 1.80 | 0.99 | | Br | 1.95 | 1.14 | | I | 2.15 | 1.33 | During formation of a covalent bond, the atoms have to come close to each other while van der Waal's forces are weak attractive forces. The molecules are thus not held closely. As a result the covalent radius is smaller than the van der Waal's radius. This also explains why covalent bonds are much stronger than van der Waal's forces. In noble gases, the van der Waal's forces are the only attractive forces. Hence the atomic radii of a noble gas atom are only the van der Waal's radii, which is the distance of closest approach of two adjacent atoms of a particular noble gas in solid state. #### (iii) Ionic Radius: The ionic radius corresponds to the radius of an ion in an ionic crystal. It may be defined as the distance from the nucleus of an ion upto which it has influence on the electron charge clouds. ##### A. Method of determination of Ionic Radius: The absolute size of an ion cannot be determined because of the following: - The internuclear distance in an ionic crystal can be very accurately measured by X-ray studies, but there is no universally accepted formula for approtioning this to the two ions. Several methods have been suggested (The most recent and accurate method is given by Shannon). - The ionic radii need to be corrected if the charge on the ions change. - The electron density around the nucleus change in the presence of the surrounding atomi.e coordination number and geometry. - The assumption that the ions are spherical is probably true for ions from s and p block elements, which attain noble gas configuration. For d-block metal ion it may not be true. - Extensive delocalization of d electron also results in change in radii. The internuclear distance in any ionic compound is accurately determined by the X-ray measurement. This distance is taken as the sum of the radii of the two ions involved i.e For an ionic compound AB. d(A-B) = rA + rB knowing the radius of one ion, that of the other can be calculated. Several methods have been suggested. Pauling's method is most widely followed. He selected some ionic solids namely NaF, KCI, RbBr and CsI. In each of these salts the cation and the anion are isoelectronic ie They have the same number of electrons which corresponds to the noble gas configuration. E.g in NaF, both Na+ and F-have 10 electrons and the configuration is 1s²2s²2p⁶, which is the same as that of Ne. Pauling proposed that in ionic compounds formed by iso-electronic ions, the ratio of the two ionic radii should be inversely proportional to the ratio of the effective nuclear charge (Z) of the two ions. The Zeff is calculated by the Slater's rule (discussed in the earlier chapter). According to this rule Zeff of Na+ and Frion at the peripheries is 6.5 and 4.5 respectively. Hence according to Pauling rNa+ / rF- = Zeff (F⁻) / Zeff (Na⁺) Thus: rNa+ / rF- = 4.5 / 6.5 Also from the X-ray studies the internuclear distance. d(Na*-F-) = rNa+ + rF- = 2.31A° Solving equation (i) and (ii) we have rNa+ = 0.95 A°. rF- = 1.36 A The radii of Na+ thus obtained can be used to calculate the radii of other anions like Br, and I ion, by measuring internuclear distance in NaBr and Nal crystal. Following this procedure the radii of most cations and anions have been determined. Some of these are listed in Table no. 3.4. Table 3.4: Calculated Radii of some Univalent Cations and Anions. | Ions | Radius (A°) | Ions | Radius (A°) | | :-------- | :---------- | :-------- | :---------- | | Li+ | 0.60 | Au+ | 1.37 | | Na+ | 0.95 | Ag+ | 1.26 | | K+ | 1.33 | F- | 1.36 | | Rb+ | 1.48 | Cr- | 1.81 | | Cs+ | 1.69 | Br- | 1.95 | | Cu+ | 0.96 | I- | 2.16 | A positive ion or cation of an atom is smaller than the atom itself as one or more electrons are removed and the outer shell is lost. The remaining inner shell obviously occupies smaller volume. Also with elimination of one or more outer electrons, the total number of electrons decreases but the nuclear charge remains the same, thus increasing the effective nuclear charge and causing contraction of size. E.g. Atomic Radius of Na = 1.57A° Ionic Radius of Na+ = 0.98A" Atomic Radius of Fe = 1.17A" Ionic Radius of Fe+2 = 0.76A" Ionic Radius of Fe+3 = 0.64A" Anions or negative ions are bigger than the corresponding atom. When the anion is formed one or more extra electrons are added to the valence shell, while the nuclear charge is the same. The effective nuclear charge is thus reduced and the electron cloud is expanded and causes an increase in size. E.g. Atomic Radius of Cl = 0.99A° Ionic Radius of Cl = 1.81A The effect of nuclear charge on ionic size is best illustrated by considering the radii of isoelectronic species (Table 3.5). All of them contain the same number of electrons. They differ only in the charge of the nucleus. It is evident from this table that as the nuclear charge increases, the electrons, are held more and more tightly by the nucleus and the ionic size thus decreases. Table 3.5: Radii of Isoelectronic Ions. | Ions | N³- | O²⁻ | F⁻ | Na⁺ | Mg²⁺ | | :---- | :-- | :-- | :-- | :-- | :--- | | No. of electrons | 10 | 10 | 10 | 10 | 10 | | Charge on nucleus | +7 | +8 | +9 | +11 | +12 | | Radius (A") | 1.71 | 1.40 | 1.36 | 0.95| 0.60 | #### B. Trend in Periodic Table: ###### Variation in a period: Within a period the radii of cations of elements of Group 1 to 3 decrease with increase in atomic number across the period, so do the radii of anions of Group 16 to 17 (Table 3.6). This is because the atomic size decreases from left to right in a period. ###### Variation in a group: In general the ionic radii increase with increase in atomic number down the group (Table 3.6), since the atomic radii also increase correspondingly. Table 3.6: Ionic Radii (A°) of some ions. | Gr 1 | Gr 2 | Gr 3 | Gr 16 | Gr 17 | | :----- | :----- | :----- | :----- | :------ | | Li⁺ | Be²⁺ | B³⁺ | O²⁻ | F⁻ | | 0.68 | 0.30 | 0.20 | 1.45 | 1.33 | | Na⁺ | Mg²⁺ | A¹³⁺ | S²⁻ | Cl⁻ | | 0.98 | 0.65 | 0.45 | 1.90 | 1.81 | | K⁺ | Ca²⁺ | Ga³⁺ | Se²⁻ | Br⁻ | | 1.33 | 0.94 | 0.60 | 2.02 | 1.96 | | Rb⁺ | Sr²⁺ | In³⁺ | Te²⁻ | I⁻ | | 1.48 | 1.10 | 0.81 | 2.22 | 2.19 | | Cs⁺ | Ba²⁺ | Tl³⁺ | | | | 1.67 | 1.29 | 0.91 | | | There is a rapid increase in ionic radius as move down the first few elements. But the increase is not so rapid among the last few elements. E.g the increase in ionic radii between Li+ to Na+, and Na+ to K+ is large, but the increase in size between K+ to Rb+ is not much. This is possibly because between K+ and Rb+ comes the 1st transition series, which involves filling up of 10 electrons in (n-1) d orbital. These additional 10 electrons (compared to only 8 electrons of sp orbital) causes an increase in nuclear charge resulting in contraction of atoms and ions. Thus the ion which follows any of the transition series has smaller size. ### 3.2 Ionization Energy The ionization energy (IE) is defined as the minimum amount of energy required to remove the most loosely bound electron from an isolated, gaseous atom. This results in the formation of a cation, thus IE is a direct measure of the ease with which an atom can change into a cation. The process is represented as follows. M(g) + IE → M²⁺(g) + e⁻ #### A. Method of determination of IE: The ionization energy (IE) can be measured experimentally by spectroscopic techniques. Another method is to have the vapours of the element in a discharge tube and connect it to a source of current. The applied voltage is gradually increased. At a certain voltage there is a sudden rise in the current due to the liberation of an electron from a neutral atom producing an ion M⁺. The energy corresponding to this voltage is known as the first ionization energy (IE₁,). On further increase in the applied voltage, the current may again show a sudden rise. This due to the elimination of another electron from a positively charged ion (M⁺) to give a bivalent M²⁺ ion. The energy corresponding to this stage is known as the second ionization energy (IE₂). If the voltage is further increased there may again be a sudden, sharp rise in current corresponding to loss of three or more electrons. The ionization energy is then referred to as third (IE₃) or higher ionization energy. These steps are represented as follows. M → (IE₁)→ M⁺ + e⁻ M⁺ → (IE₂) M²⁺ + e⁻ M²⁺ →(IE₃) M³⁺ + e⁻ The unit of IE is KJ/mole or electron volt, ev (1 ev = 96.49 KJ/mole). The successive ionization energy of some elements are listed in Table 3.7. Table 3.7: Successive IE of some elements (KJ/mole) | Element | Electronic Conf. | (IE₁) | (IE₂) | (IE₃) | | :------- | :---------------- | :----- | :----- | :----- | | Li | 1s² 2s¹ | 520.1 | 7297 | 11810 | | Be | 1s² 2s² | 899.3 | 1758 | 14810 | | B | 1s² 2s² 2p¹ | 800.6 | 2428 | 3660 | | C | 1s² 2s² 2p² | 1086.2| 2353 | 4618 | | N | 1s² 2s² 2p³ | 1402.1| 2855 | 4577 | | O | 1s² 2s² 2p⁴ | 1313.7| 3388 | 5297 | | F | 1s² 2s² 2p⁵ | 1680.8| 3375 | 6045 | | Ne | 1s² 2s² 2p⁶ | 2080.4| 3963 | 6130 | It can be seen that the 2nd ionization energy (IE₂) is very much higher than the 1st ionization energy (IE₁) while the 3rd ionization energy (IE₃) is still higher. This is due to the fact that after the removal of an electron, the number of electrons decrease while the nuclear charge remains the same. As a result the remaining electrons are held more tightly by the nucleus due to very high Zeff developed, and it becomes increasingly difficult to remove the 2nd and the 3rd electron. Thus (IE₃) > (IE₂) > (IE₁). #### B. Factors affection IE: IE depends upon the following factors: 1. **Atomic size**: In a small atom the electrons are tightly held, as there are few electron shells close to the nucleus. Larger the atom more are the number of electron and electron shells lying further away from the nucleus. As a result they are less strongly held by the nucleus. Due to the weak attractive pull, it becomes easier to knock out an electron from the outer shell. Hence larger the atomic size, the smaller is the ionization energy. 2. **Nuclear charge**: With the increase in the nuclear charge the size decreases and also the force of attraction between the nucleus and the outermost electron increases. As a result energy required to pull the outer electron is greater. Hence ionization energy increases with increase in nuclear charge. 3. **Shielding (screening) effect**: The valence shell electrons are shielded from the attractive force of the nucleus by the electrons in the inner shell. This effect is called the shielding or screening effect. Larger the number of electrons in the inner shells, greater is the screening effect. The outer electron therefore feels less nuclear attraction leading to lower IE. This factor explains the large decrease in IE on moving from a noble gas to an alkali metal. In alkali metals the only electron from the ns orbital is easily removed as it is effectively screened from the nuclear attraction by the electrons present in (n-1)th shell. While in noble gas the lost electron is not effectively shielded by the remaining electrons from the n orbital. Thus due to greater shielding effect of the inner (n-1) electrons, alkali metals show lower ionization energy while it is difficult to ionize a noble gas as it also involves breaking a complete shell of an electron. 4. **Type of electron/orbital involved (s, p, d or f)**: The IE also depends on the type of orbital from which the electron is removed. s, p, d and f orbitals have different shape. An s electron penetrates nearer to the nucleus and is therefore more tightly held than a p orbital, while p orbital electrons are more tightly held compared to d electrons, which in turn is more tightly held as compared to f. Thus in a particular principle shell, the s electron experiences greater nuclear attraction as compared to p, d and f electron. Keeping all the other factors constant the IE corresponding to removal of an electron from an orbital decreases in the order ns > np > nd > nf. The increase in the IE is therefore not quite smooth on moving from left to right in a period. Thus Group 2 elements show higher IE when ns electron is removed, while the adjacent Group 3 elements show lower IE as it involves removal of np electrons. 5. **Electronic configuration**: Certain electronic configuration are found to be more stable as a result the corresponding IE is higher. E.g If an atom has a fully filled or exactly half-filled orbital, its ionization energy is higher than expected normally from its position in the periodic table. Thus Be and N in the second period and Mg and P in third period show slightly higher IE due to extra stability of the fully occupied s orbital in Be and Mg and half filled p orbital in N and P. Also the noble gas elements show highest IE in their respective period, due to the highly stable octate configuration. Similarly for Gr I elements the 2nd ionization energy is about 7-14 times the 1 ionization energy, because removal of 2nd electron involves breaking into a filled shell of electrons and this requires very high energy. #### C. Trend in Periodic Table: ###### Variation in period: In general, the ionization energy increases as we move along a period from left to right (fig 3.2). On moving across a period, the atomic size decreases and the nuclear charge increases, consequently the nuclear attractive force felt by the electrons in the outermost shell goes on increasing. It thus becomes increasingly difficult to remove an electron from the outer shell and hence the ionization energy increases. It is maximum for noble gas which has stable octate electronic configuration (ns²np⁶). It requires a great deal of energy to remove an electron from a stable filled shell of electrons. The plot (fig 3.2) shows many irregularities. This is because atomic size is not the only factor which determines the ionization energy. ###### Variation in a group: In general the 1st ionization energy decreases on going from top to bottom in a group. As we move down the group there is a gradual increase in the atomic size with increase of the principal energy shell though the nuclear charge also increases. Also there is increase in the number of inner electrons, which increase the shielding effect. The overall effect of increase in atomic size and shielding effect is more than the effect of increase in nuclear charge. Therefore the electron is less and less tightly held to the nucleus. Hence the 1ª IE decreases down the group. However a departure from this trend occurs in Group 13 (Table 3.8). The expected decrease in IE occurs from B to Al, but the further element Ga, In and Tl show an irregular pattern for the 1st as well as 2nd and 3rd ionization energy. Table 3.8: Ionization Energies (KJ/mole) for Group 13 | | 1st | 2nd | 3rd | | :------- | :-- | :--- | :---- | | B | 801 | 2427 | 3659 | | Al | 577 | 1816 | 2744 | | Ga | 579 | 1979 | 2962 | | In | 558 | 1820 | 2704 | | TI | 589 | 1971 | 2877 | ### 3.3 Electron Affinity (EA) : Electron affinity (EA) of an element is defined as the amount of energy released when an extra electron is added to the gaseous isolated atom. X(g) + e⁻ → X⁻(g) + Energy (EA) Usually only one electron is added forming a univalent anion. The reaction is exothermic (-ΔH) therefore EA has a negative sign. But normally they are shown without the sign. The greater the energy released, greater is the electron affinity. The electron affinity is expressed in eV per atom or KJ/mol-¹. #### L. Method of Determination of Electron Affinity: Since electron affinity cannot be determined experimentally, it is determined by using a thermo-cycle. This cycle was devised by Born and Haber in 1919. It relates the lattice energy of crystals to other thermodynamic data. It is based on Hers Law according to which the overall energy change in a process depends only on the initial and final stage and not on the route. Thus it can be seen that (fig 3.3) the enthalpy of formation ΔHƒ of a NaCl crystal lattice is the algebraic sum of the terms involved in the cycle. [Insert image of the Born-Haber Cycle](null) ΔH₁ = ΔHsub + IE + 1/2ΔHdiss + EA + U All these energies except Lattice energy (U) and Electron affinity (EA) can be measured. By using a known crystal structure the lattice energy can be determined and hence electron affinity can be calculated. Thus for NaCl crystal. ΔH₁ = ΔHsub + IE + 1/2ΔHdiss + EA + U 381.2 + 108.4 + 495.4 + 120.9 + EA + (-757.3) EA = 348.6 J/mole After knowing some electron affinity the cycle can be used to calculate lattice energy of other crystals, which is a useful guide to determine solubility of a crystal. When the lattice energy is high a large amount of solvation energy is required to break the lattice. It is unlikely to have a large enthalpy of solvation, so the substances with high lattice energy will probably be insoluble. The 'noble behaviour' of many transition metals is also due to high heat of sublimation, high IE and less enthalpy of solvation of ion. Lattice energy also provides information about the ionic or covalent nature of bonding. If the lattice energy is calculated theoretically from the Born - Haber cycle, assuming ionic bond, matches well with the experimental value then the bonding is truly ionic and if it does not match then the bonding is not ionic. #### II. Factors affecting Electron Affinity: Electron affinity depends on the atomic size, effective nuclear charge and electronic configuration. 1. **Atomic size**: In case of a smaller atom the attraction of the nucleus felt by the extra electron added is stronger than in the case of a larger atom. Therefore smaller the atom greater is the electron affinity. 2. **Effective nuclear charge**: Greater the effective nuclear charge felt at the periphery of an atom, stronger is the affinity for the extra electron added. Thus with increase in Zeff, the electron affinity increases. 3. **Electronic configuration**: An element having completely filled principal shell like noble gas (ns² np⁶) with octate configuration cannot accept an extra electron, and therefore shows zero electron affinity. Also elements of Group 2 with (ns²) configuration and Group 12 with (n-1)d¹⁰ ns² configuration have completely filled s and d orbitals. Due to stability of these filled orbitals they have very little tendency to accept an extra electron and show very little EA. Similarly elements like N (1s² 2s² 2p³) and P ([Ne] 3s² 3p³) with exactly half filled p orbital, being very stable show zero or very slight electron affinity (Table 3.9). #### III. Trends in Periodic Table: ###### Variation in a Period: In general electron affinity (EA) increases from left to right across a Period. The atomic size and the nuclear charge goes on increasing across a period and hence the colombic attraction for the extra electron added goes on increasing. Thus in a period the alkali metals have lowest electron affinity while halogen at the extreme end shows highest electron affinity. (Table 3.9). The trends in a period are sometimes irregular indicating that atomic size is not the only criterion for determining electron affinity. Thus EA of Be is less than that of Li though the atomic size of Be is smaller than that of Li. Similarly N shows lesser electron affinity than O though the atomic size of O is smaller than that of N. ###### Variation in a Group: On moving down the group the size as well as nuclear charge increases. But the effect of increasing size out weighs the effect of increase in nuclear charge. As a result the Zeff decreases and consequently the EA goes on decreasing down the group. (Table 3.9). However there are some deviations in this trend. The 1st element of Group 16 and Group 17 namely Oxygen and Fluorine show less electron affinity than the next member in the group. This is due to very small size of O and Fas compared to the next element Sand Cl respectively. The addition of an extra electron produces high electron charge density in a relatively compact 2p shell, resulting in strong electron-electron repulsion. This repulsion thus results in low electron affinity. Elements of Group 13 and Group 14, which follow the d block, do not show the expected decreasing trend. It is probably due to the shielding effect of the preceeding d subshell. Table 3.9: Electron Affinity (in KJ/mole) of various Elements | Group 1 | Group 2 | Group 13 | Group 14 | Group 15 | Group 16 | Group 17 | Group 18 | | :------- | :------- | :------- | :------- | :------- | :------- | :------- | :------- | | Li | Be | B | C | N | O | F | Ne | | 59.8 | 26.7 | 122.3 | (0) | ≈0 | 140.9 | 328.6 | 0 | | Na | Mg | Al | Si | P | S | Cl | Ar | | 53.1 | 45.2 | 133.6 | 72.3 | 200.7 | 348.5 | 0 | | K | Ca | Ga | Ge | As | Se | Br | Kr | | 48.4 | 28.7 | 120.0 | 77.2 | 195.0 | 324.7 | (0 | | Rb | Sr | In | Sn | Sb | Te | I | Xe | | 46.9 | 29.0 | 107.9 | 103.0 | 190.1 | 295.5 | 0 | | Cs | Ba | TI | Pb | Br | Po | At | Rn | | 45.5 | 19.2 | 35.1 | 91.0 | ≈0 | 180.0 | 270.0 | 0 | Halogens show a very high electron affinity due to their strong tendency to gain an additional electron and attain the stable octate configuration. #### III. Trends in Periodic Table: ###### Variation in a Period: In general electron affinity (EA) increases from left to right across a Period. The atomic size and the nuclear charge goes on increasing across a period and hence the colombic attraction for the extra electron added goes on increasing. Thus in a period the alkali metals have lowest electron affinity while halogen at the extreme end shows highest electron affinity. (Table 3.9). The trends in a period are sometimes irregular indicating that atomic size is not the only criterion for determining electron affinity. Thus EA of Be is less than that of Li though the atomic size of Be is smaller than that of Li. Similarly N shows lesser electron affinity than O though the atomic size of O is smaller than that of N. ###### Variation in a Group: On moving down the group the size as well as nuclear charge increases. But the effect of increasing size out weighs the effect of increase in nuclear charge. As a result the Zeff decreases and consequently the EA goes on decreasing down the group. (Table