(4.4).1 IONIZATION ENERGY

(4.4).1.1 Definition

Ionization energy (IE) is the energy required to remove the highest energy or ‘the most loosely held’ electron in the ground state of isolated atom, molecule or ion. The reactions for ionization of electron can be shown as follows:

H(g) → H⁺(g) + e^–               ΔH = 1310 kJ mol^–1
O₂(g) → O₂⁺(g) + e^–              ΔH = 1370 kJ mol^–1

NO(g)  NO⁺(g) + e^–                ΔH = 900 kJ mol^–1
Cu⁺(g)   Cu²⁺(g) + e^–              ΔH = 1950 kJ mol^–1
F^–(g) + e^–    F(g) + e^–               ΔH = 320 kJ mol^–1

As ionization energy is the energy required to remove an electron from the positive influence of the nucleus, it is always endothermic. If ΔH is the enthalpy change for the ionization reaction, then IE = + ΔH.

(4.4).1.2 Successive Ionizations

If the species to be ionized contains more than one electron, these can be removed successively one after the other. Thus, we have the first, second, third, … ionization energies of the atoms for the ionization of first, second, third, …  electron respectively:

Mg(g) → Mg⁺(g) + e^–           ΔH = 738 kJ mol^–1
Mg⁺(g) → Mg²⁺(g) + e^–       ΔH = 1440 kJ mol^–1
Mg²⁺(g) → Mg³⁺(g) + e^–     ΔH = 7740 kJ mol^–1

Zn(g) → Zn⁺(g) + e^–               ΔH = 906 kJ mol^–1
Zn⁺(g) → Zn²⁺(g) + e^–           ΔH = 1733 kJ mol^–1
Zn²⁺(g) → Zn³⁺(g) + e^–         ΔH = 3831 kJ mol^–1

Since successive ionization of electrons from an atom leads to successive increase of positive charge on the cation formed, the successive ionization energies  increase progressively:

IE₁ < IE₂ < IE₃

The ionization energies of the elements for their common ionizations are given in Table (4.4).1.

Table (4.4).1 Ionization energies of Elements in kJ mol^–1
————————————————————————-
Element               IE₁          IE₂          IE₃          IE₄      IE5        IE6
————————————————————————-

H                           1312
He                         2372      5251
Li                           520.2     7298      11820
Be                         899.2     1757      21007
B                            800.5     2426      3659      25025
C                            1086      2352      4619      6220      37831
N                           1402      2856      4578      7475      9445      53266
O                           1314      3398      5301      7469      10990
F                            1681      3374      6050      11023
Ne                         2081      3952      6122      9371      12177

Na                         495.8     4562      6910
Mg                        737.7     1451      7733      10543
Al                           577.5     1817`     2745      11577
Si                           786.5     1577      3232      4356      16091
P                            1012      1907      2914      4964      6274      21267
S                            999.6     2252      3357      4556      7004      8495
Cl                           1251      2298      3822      5159      6542      9362
Ar                          1520      2666      3931      5771      7238      8781

K                            418.8     3052
Ca                         589.8     1145      4912      6941
Sc                          633.1     1235      2389      7091
Ti                           658.9     1310      2653      4175      9581      11533
V                            650.9     1414      2830      4507      6299      12363
Cr                          652.9     1591      2987      4743      6702      8745
Mn                        717.3     1501      3248      4940      6990      9221
Fe                          762.5     1562      2957      5290      7240      9560
Co                         760.4     1648      3232      4950      7670      9840
Ni                          737.1     1753      3395      5300      7339      10400
Cu                         745.5     1958      3555      5536      7700      9900
Zn                          906.4     1733      3833      5731

Ga                         578.8     1979      2963      6180
Ge                         762        1538      3302      4411      9020
As                          947.0     1798      2735      4837      .6043     12310
Se                          941.0     2045      2974      4144      6590      7880
Br                          1139.5   2103      3470      4560      5760      8550
Kr                          1350.8   2350      3565      5070      6240      7570

Rb                         403.0     2633
Sr                           549.5     1062      4138
Y                            600        1180      1980      5847
Zr                           640.1     1270      2210    3313      7752
Nb                         652.1     1380      2416      3700      4877      9847
Mo                        648.3     1560      2618      4480      5257      6641
Tc                          702        1470      2850
Ru                         710.2     1620      2747
Rh                         719.7     1740      2997
Pd                          804.4     1870      3177
Ag                         731.0     2070      3361
Cd                         867.8     1631      3616
In                           558.2     1821      2704      5210
Sn                          708.6     1421      2943      3930      7456
Sb                          834        1595      2440      4260      5400      10400
Te                          869.3     1790      2698      3610      5668      6820
I                             1008.4   1846      3180
Xe                          1170.4   2046      3099

Cs                          375.7     2234      3400
Ba                         502.7     965.2     3600
La                          538.1     1067      1850      4819
Ce                         534.4     1050      1949      3547      6425
Pr                          527        1020      2086      3761      5551
Nd                         533.1     1040      2130      3900
Pm                        540        1050      2150      3970
Sm                         544.5     1070      2260      3990
Eu                          547.1     1085      2404      4120
Gd                         593.4     1170      1990      4250
Tb                          565.8     1110      2114      3839
Dy                         573        1130      2200      3990
Ho                         581        1140      2204      4100
Er                          589.3     1150      2194      4120
Tm                        596.7     1160      2285      4120
Yb                          603.4     1175      2417      4203
Lu                          523.5     1340      2022      4370      6440
Hf                          658.5     1440      2250      3216
Ta                          761        1500
W                          770        1700
Re                         760        1260      2510      3640
Os                         840        1600
Ir                           880        1600
Pt                          870        1791
Au                         890        1980
Hg                         1007      1810      3300
Tl                           589.4     1971      2878
Pb                          715.6     1450      3081      4083      6640
Bi                           703        1610      2466      4370      5400      8520
Po                         812.1
At                          899
Rn                         1037

Fr                           380
Ra                         509.3     979
Ac                          499        1170
Th                          587        1110      1930      2780
Pa                          568
U                           597.6     1420
Np                         604.5
Pu                          584.7
Am                        578
Cm                        581
Bk                          601
Cf                          608
Es                          619
Fm                         627
Md                        635
No                         642
Lw                         470
Rf                          580
————————————————————————————–
Note: In atomic physics, ionization energies are used in units of eV.
To convert the IE values in kJ mol^–1 into eV, divide by 96.48536.

Figure (4.4).1 Variation of ionization energy with atomic number Z

(4.4).1.3 Factors controlling Ionization Energies

The ionization energy depends on the following factors.

3. Effective Nuclear Charge: Higher the effective nuclear charge, higher is the ionization energy. This includes the effects of inner shell electrons in shielding the outer electrons. It is observed that the first ionization energy of the 2p electron from lithium (520.1 kJ mol^–1) to neon (2081 kJ mol^–1) is similar to that of their Z*
4. Size of Atoms: Ionization energy increases with a decrease in size of the orbital from which electron is to be ionized. This is due to the increase in electrostatic attraction between the nucleus and the electron to be removed.
5. Penetration of Orbital from which Electron is Ionized: The ionization energy (IE) of electron in a shell increases if the orbital becomes more penetrating. (A more penetrating orbital of a shell is closer to the nucleus and has a lower energy.) Therefore, ionization energy for   electron from different subshells is in the order

s > p > d > f.

The first IE for group 13 or group III A atoms is less than that for the group 2 or II A  elements in spite of a higher nuclear charge on the former atoms, the values (in kJ mol^–1) being  899 for beryllium (1s²2s²) but only 800 for  boron (1s²2s²2p¹), and 737 for magnesium (1s²2s²2p¹3s²) but only 577 for aluminium (1s²2s²2p¹3s²3p¹).

In case of group 2 elements, the electron is ionized from the electron pair in the more penetrating and closer orbital closer to nucleus (2s) orbital, whereas for group 13 elements, it is the unpaired electron in the less penetrating p orbital that is ionized. Both the factors increase the stability of the electron in s² pair and thus, explain the observations.

4. Stability of Half-filled and Completely filled Shells: The ionization energy of electron increases if the shell or subshell from which the electron is ionized is half-filled shell (d⁵ or p⁶) or a completely-filled closed shell (d¹⁰ or p⁶). The relative stabilities of these configurations, and hence, ionization energies, increase in the order:

d⁵ < p³ < d¹⁰ < p⁶ < s²p⁶.

This explains the following observations.

• The noble gases of group 18 with a closed shell configuration of ns²np⁶ have very high ionization energies as compared to the corresponding halogens (group 17, ns²np⁵ elements), as seen by the values of ionization energy: F= 1681, Ne = 2080; Cl = 1256, Ar = 1520; Br = 1143, Kr = 1351; and I = 1008, Xe = 1170 (all values in kJ mol^–1).
• The ionization energy of oxygen (1s²2s²2p⁴) (1314 kJ mol^–1) is lower than that of nitrogen (1s²2s²2p³) (1402 kJ mol^–1) because for nitrogen, the electron should be ionized from a more stable half-filled shell (p³), whereas for oxygen, loss of one electron from 2p orbital leads to the formation of a stable half-filled 2p³ Both of these factors reduce the first ionization energy of oxygen.
• Ionization energies of other group 16 elements (S = 1000, Se = 941) are very close to those of the corresponding group 15 elements (P = 1012, As = 947) (all values in kJ mol^–1). This is also due to the same factors as above because the group 16 elements have electronic configuration of ns²np⁴ whereas group 15 elements have the half-filled p³ configuration of ns²np³. Here the effect of increased nuclear charge is just compensated by the stability of the half-filled $p^3$ shells.
• However, for heavy elements, no such stability of p³ configuration is seen. Ionization energy for antimony (Sb) (5s²5p³) = 834 vs tellurium (Te) (5s²5p⁴) = 869, or bismuth (Bi) (6s²6p³) = 703 vs polonium (Po) (6s²6p⁴) = 813  (all values in kJ mol^–1).
• Change in Quantum Shell: Successive ionization energies for a species increase regularly but if the electron to be ionized belongs to a lower quantum shell, ionization energy increases many-fold because of the combined effects of:
• Increase in effective nuclear charge (Z),
• Closed shell stability of lower filled shell, and
• Sudden decrease in the  radius of the lower shell, which now becomes the outermost shell.

Consider the case of magnesium or zinc. The first and the second ionization energies increase regularly as expected. But the third ionization energy of these atoms (Mg = 7733, Zn = 3833 kJ mol^–1) becomes very high, because the third electron comes from the lower quantum shell of electrons.   More examples can be found for other elements, especially the transition elements where the relative changes in ionization energies determine the stabilities of different oxidation states.

The difference in the stability of the closed shells is reflected in the values of ionization energies also. Ionization of electron from the less stable d¹⁰ configuration requires less energy than the loss of electron from the more stable 2s²2p⁶ configuration. This is reflected by the jumps in the ionization energies values:

——————————————————————————–
Atom                    IE­₁             IE₂          IE₃          IE₄ Last Shell broken
——————————————————————————–
Na                         495.8     4562      6910                     3s²3p⁶
Mg                        737.7     1451      7733                      3s²3p⁶
Zn                          906.4     1733      3833                    3d¹⁰
In                           558.2     1821      2704      5210      4d^{10
—————————————————————————————————-}

(4.4).1.4 Variations in Ionization Energies in Periodic Table

The variation of ionization energy or enthalpy of the elements with atomic numbers is shown in Figure (4.4).1,

(4.4).1.4.1 Variation in a Group

In a periodic group, ionization energy decreases with increase in atomic size from top to bottom as expected from changes of the radii of the atoms.

1. For s block elements of group 1 (alkali metals) and group 2 (alkaline earth metals), ionization energy decreases from lithium to cesium and from beryllium to barium as the radius of the atom increases.
2. For the p block elements, the decrease in ionization energy is not regular, especially between third and fourth period post-transition elements. This may be due to the filling up of the $3d$ orbitals in the penultimate inner shell in the fourth period. The 3d electrons cannot  shield the valence shell 4s and 4p electrons completely from the effect of increased nuclear charge. Calculations show that the effective nuclear charge (Z*) of the fourth period elements is more than that of the corresponding elements of third period (Z for Al = 3.5, Ga = 5.0; Si = 4.15, Ge = 5.65; etc.
3. For transition elements, ionization energy decreases slightly from 3d to 4d elements due to size. However, due to the lanthanide contraction, the radius of 5d elements is less than or almost equal to the radius  of corresponding 4d Due to combined effects of (1) smaller radius  and  (2) high nuclear charge, the ionization  energy of 5d elements become equal to or even more than those of the 4d elements.

For the later elements of the two series, the ionization energy becomes almost same.

(4.4).1.4.2 Variation in a Period

 For the representative elements, the ionization energy increases with increasing nuclear charge, except for group 13 (np¹) elements (due to ionization of p electron), and group 16 np⁴) elements due to stability of p³

1. For transition elements, the ionization energy remains almost constant throughout the period. This is because of the combined effects of electron–nucleus attractions and the electron–electron repulsions in the d However, for (n – 1)d⁵ns¹ (group 6) and (n – 1d¹⁰ ns¹) (group 11)  elements, ionization energy decreases due to the larger size of these atoms due to the formation of d⁵ and d¹⁰ configurations.
2. For inner transition elements, the ionization energy increases with atomic number throughout the series as expected, except at samarium (Sm) (Z = 62) and ytterbium (Yb) (Z = 70) due to the stability of half-filled f⁷ and completely filled f¹⁴

(4.4).2 ELECTRON AFFINITY

(4.4).2.1 Definition

Electron affinity (EA) is defined as the energy released} when an electron is added to the gaseous atom or ion.

F(g} + e^– → F^– (g)                EA = 320 kJ mol^–1

 If ΔH is the enthalpy change for the addition of an electron i.e., electron affinity reaction, then EA = –ΔH. Therefore, electron affinity of an atom can be taken as the ionization energy of the corresponding anion:

F^– (g) → F(g) + e^–
IE of F^– =  EA of F = 320 kJ mol^–1

Whereas ionization energies are always concerned with the formation of positive ions, electron affinities are their negative ion equivalent.

Negative electron affinities can be used where electron capture requires energy. In these cases, electron capture can occur only if the impinging electron has a kinetic energy large enough to excite a resonance in the atom + electron system. Conversely electron removal from the anion formed in this way releases energy, which is carried out by the freed electron as kinetic energy. The resulting anions formed this way are always unstable with lifetimes of microseconds to milliseconds, and simply lose electrons spontaneously in a short time.

Successive Electron Affinities: The atoms, molecules or ions can accept more than one electron stepwise, and form multi-charged anions having more than one unit of negative charge, e.g., O^{2– ,} S^2– , SO₃^2– , PO₄^3–.

First electron affinity (EA) for the nonmetals of group 15, 16 and 17 is exothermic and positive (ΔH for the addition of electron is negative), and forms negatively charged anion X^–. Addition of another negatively charged electron to a negatively charged species X^– cannot be spontaneous due to electrostatic repulsions, so that second and higher electron affinities of all the atoms become negative and highly endothermic. They are so highly endothermic that the overall formation of the anion becomes non-spontaneous even for the highly electronegative oxygen and sulphur, which form a stable anion with an octet configuration

O(g) + e^– →  O^–(g) +141 kJ mol^–1                    Exothermic
O^–(g)  + e^– →  O^2–(g) — 780 kJ mol^–1             Endothermic
O(g) + 2e^– →  O^2–(g) — 640 kJ mol^–1               Endothermic

S(g) + e^– →  S^–(g) + 200 kJ mol^–1                            Exothermic
S^–(g)  + e^– →  S^2–(g) – 456 kJ mol^–1                 Endothermic
S(g) + 2e^– →  S^2–(g) – 256 kJ mol^–1                  Endothermic

(4.4).5.2 Trends in Electron Affinities

The electron affinity of an atom depends on the effective nuclear charge and its size, and follows trends similar to those for the ionization energies (Figure (4.4).1).

1. In a group, magnitude of electron affinity increases with increasing size of the atoms, except for the smallest member in the group (see below).
2. Alkali metals} (ns¹ elements) have a positive electron affinity. This is because they release pairing energy when the s² electron pair is formed in the valence shell by accepting an incoming electron.
3. In a period, electron affinity increases with increase in nuclear charge due to the smaller size and increasing nucleus–electrons attractions. Atoms with higher ionization energy have higher electron affinity, which increases in the order:

B (26.99) < C (121.78) < O (141) < F (328.16)

The electron affinity of nitrogen (– 6.8 kJ mol^–1) is less than that of carbon, because the addition of electron to nitrogen atom gives a p⁴ configuration and results in the loss of exchange energy or stability of the half-filled p³ configuration.

4. Electron affinities for the elements of second period are lower than those for the third period elements. This is due to the higher electron density in the valence shell of second period elements, which results in stronger electrostatic repulsions for the incoming electron for smaller atoms (all values in kJ mol^–1).

B (26.99)    C (121.78)      N (– 6.8)      O (141)      F (328.16)
Al (44)          Si (120)           P (75)            S (200)      Cl (345)

6. Most of the transition metals} have a positive electron affinity as they can accommodate the incoming electron in their inner (n – 1)d orbitals, e.g., Cr = 64, Ni = 111, Cu = 123, Ga = 36, Pb = 100 and Au = 223 (all values in kJ mol^–1). However, for atoms with s² or $d^10sd² configuration, the incoming electron has to go to a higher energy level, and electron affinities become non-spontaneous and endothermic. For example, electron affinity is negative for ns² (group 2) elements (Be = –240, Mg = –230 and Ca = –156 kJ mol^–1) and for $(n-1)d^{10}ns^2$ (group 12) elements (Zn =–87  kJ mol^–1, Cd = –56 kJ mol^–1).
7. Gold has a very high electron affinity of 222 kJ mol^–1, may be due to high effective nuclear charge, poor shielding of nucleus by d electrons and small size. Gold dissolves in alkali metals on heating and forms the ionic compounds (Cs⁺Au^–), which contain negatively charged auride
8. The values of EA for inner transition elements are:

     f⁰               f¹                 f²                   f³              f⁴              f⁵
La (53)   Ce (53)  Pr (93), Nd (184), Pm (12.4),  Sm (15)
Ac (34)  Th (114)  Pa (53)  U (51)        Np (46)   Pu (–48)
         f⁶            f⁷             f⁸                   f⁹                    f¹⁰             f¹¹
Eu (11)   Gd (13)    Tb (112)   Dy (34)    Ho (33)     Er (30)
Am (–10) Cm (27) Bk (–165)  Cf (–97)   Es (–29)    Fm (34)
        f¹²                  f¹³                 f¹⁴
Tm (98)      Yb (–1)   Lu (23)
Md (94)   No (–224)  Lw (–30)

Except for Pr (93), Nd (184), Tb (112) and Tm (98) in the 4f series, and Th (114) and Md (89) in the 5f series, the EA are within thermal energy range (42 at 300 K and 55 at 400 K) (all values in kJ mol^–1), or are negative, so that these anions are unstable. No trend is apparent in the series as a whole.

(4.4).2.3 Electron Affinities of Molecules and Radicals

For sake of completeness, the electron affinities (EA) of some common molecules and radicals (which form anions in compounds) are given below. Many more have been listed by Rienstra–Kiracofe et al. (2002). The electron affinities of the radicals OH and SH are the most precisely known of all molecular electron affinities.

OH = 176.3473, SH = 223.3373, NO = 2.5, NO₂ = 219.3
C₂ = 43,42, C₆₀ = 258.92, O₂ = 43.42, O₃ = 202.89,
F₂ = 297, Cl₂ = 227, Br₂ = 244, I₂ = 243.5
SF₆ = 99.4, WF₆ = 338, UF₆ = 488, C₂(CN)₄ = 306