(4.4).1 IONIZATION ENERGY
Ionization energy values are typically very high and follow trends throughout the periodic table. The IE increase from bottom to top and left to right in the periodic table. Below is a diagram showing the directions atomic size increase over the periodic table. The ionization energy of an atom is equal to the amount of energy given offwhen an electron is added to an atom. When an electron is added to an atom, wecall the energy given off the electronaffinity(EA). For most atoms, the initial electron affinity isexothermic meaning energy is given off. The ionization energies associated with some elements are described in the Table 1. For any given atom, the outermost valence electrons will have lower ionization energies than the inner-shell kernel electrons. As more electrons are added to a nucleus, the outer electrons become shielded from the nucleus by the inner shell electrons. The first of these quantities is used in atomic physics, the second in chemistry, but both refer to the same basic property of the element. To convert from 'value of ionization energy' to the corresponding 'value of molar ionization energy', the conversion is: 1 eV = 96.48534 kJ/mol. Jul 03, 2019 Ionization energy is the minimum energy required to remove an electron from an atom or ion in the gas phase. The most common units of ionization energy are kilojoules per mole (kJ/M) or electron volts (eV). Ionization energy exhibits periodicity on the periodic table.
(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
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
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
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
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
————————————————————————-
————————————————————————-
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
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
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
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
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
Ionization Energy Periodic Table
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
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
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
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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.
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.
- 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*
- 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.
- 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.
- 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:
Second Ionization Energy Table
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 4d10
—————————————————————————————————-
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 4d10
—————————————————————————————————-
(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.
- 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.
- 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.
- 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
- 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.
- 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:
![Successive ionization energy table Successive ionization energy table](/uploads/1/3/7/1/137181548/932921933.png)
F– (g) → F(g) + e–
IE of F– = EA of F = 320 kJ mol–1
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., O2– , S2– , SO32– , PO43–.
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– → O2–(g) — 780 kJ mol–1 Endothermic
O(g) + 2e– → O2–(g) — 640 kJ mol–1 Endothermic
O–(g) + e– → O2–(g) — 780 kJ mol–1 Endothermic
O(g) + 2e– → O2–(g) — 640 kJ mol–1 Endothermic
S(g) + e– → S–(g) + 200 kJ mol–1 Exothermic
S–(g) + e– → S2–(g) – 456 kJ mol–1 Endothermic
S(g) + 2e– → S2–(g) – 256 kJ mol–1 Endothermic
S–(g) + e– → S2–(g) – 456 kJ mol–1 Endothermic
S(g) + 2e– → S2–(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).
- In a group, magnitude of electron affinity increases with increasing size of the atoms, except for the smallest member in the group (see below).
- 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.
- 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.
- 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)
Al (44) Si (120) P (75) S (200) Cl (345)
- 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 $d10sd2 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).
- 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
- 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)
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
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