Electronic affinity

The electron affinity or electroaffinity is defined as the energy released when a neutral gaseous atom in its ground state captures an electron and forms a mononegative ion.

X(g)+e− − Δ Δ X− − (g)+Eea{displaystyle mathrm {X(g)+e^{-}} longrightarrow mathrm {X^{-}(g)}} +E_{rm {ea},}.

Since it is about released energy, since normally when inserting an electron in an atom the attractive force of the nucleus predominates, it has a negative sign. In the cases in which the energy is absorbed, when the repulsive forces win, they will have a positive sign; E_ea is commonly expressed in the International System of Units, in kJ·mol^-1.

We can also resort to the opposite process to determine the first electronic affinity, since it would be the energy consumed in extracting an electron from the mononegative anionic species in the gaseous state of a certain element; evidently the corresponding enthalpy E_ea has a negative sign, except for noble gases and alkaline earth metals. This process is equivalent to that of the ionization energy of an atom, so the E_ea would be, by this formalism, the ionization energy of zero order.

This property helps us to predict which elements will easily generate stable anionic species, although other factors must not be relegated: type of counterion, solid state, ligand-dissolution, etc.

Methods for determining electron affinity

In many cases it can be measured directly by using electron beams that strike atoms in the gas phase. In a less precise way, it can be estimated by extrapolation of the values of the different ionization energies available to the atom considering: E_i,1, E_i,2, etc.

General trends

The electron affinity increases when the size of the atom decreases, the screen effect is not strong or when the atomic number increases. Seen another way: electron affinity increases from left to right, and from bottom to top, just like electronegativity does. In the traditional periodic table it is not possible to find this information.

The elements of the p block, and specifically those of group 17, are those with the highest electronic affinities, while the atoms with external configurations s² (Be, Mg, Zn), s²p⁶ (Ne, Ar, Kr) along with those that the set of p orbitals (N, P, As) are half-filled and are those with the lowest E_ea. The latter demonstrates the quantum stability of these electronic structures that do not admit to being easily disturbed. The elements with the highest E_ea are fluorine and its closest neighbors O, S, Se, Cl and Br -a notable increase in the effective nuclear charge defined in this area of the periodic table-, except for the noble gases that have a closed electronic structure of high stability and each electron that is inserted into them must be placed in an empty upper shell.

We are going to highlight some aspects related to the E_ea that are inferred by the position and zone of the element in the periodic table:

• The elements located on the right side of the periodic table, block p, are those of favorable electronic affinities, manifesting their clearly non-metallic character.
• The highest electronic affinities are for group 17, followed by group 16 elements.
• It is surprising that the fluoride has less affinity than the chlorine, but by placing an electron in the F, a smaller atom than the Cl, repulsive forces must be overcome between the electrons of the valence layer. From the chlorine the trend is expected according to the greater distance from the external electrons to the nucleus.
• Nitrogen has an electronic affinity far below its neighboring elements, both from the period and from its group, which is due to its seedling layer that is very stable.
• The remaining elements of group 15 do have more favorable electronic affinities, despite the stability of the seed layer, because the increase in size makes that outer layer separate from the core by other means.
• We must also highlight the role of hydrogen, since its affinity is not very high, but enough to generate ion H^- which is very stable in ionic hydra and complex species. Here we can also apply the argument analogous to that of the fluoride, because we have an even smaller atom and we want to add an electron by overcoming the repulsive forces of the electron. 1¹.
• Regarding the block d must be fixed in the special case of gold as its electronic affinity -223 kJ·mol⁻¹is comparable to that of iodine with -295 kJ·mol⁻¹, so it is feasible to think of the anion Au^-. It has been possible to synthesize ionic gold compounds of the RbAu and CsAu type, with the participation of the most electropositive alkaline metals. In them is reached the noble pseudogas configuration of the Hg (of 6s¹ to 6s²) for ion Au^- (lanthanide contract + maximum relativistic contraction in the Au).

Group

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

Period

1

H -73

He 21

2

Li - 60.

Be 19

B -27

C -122

N 7

O -141

F -328

Ne 29

3

Na - 53

Mg 19

Al -43.

Yeah. -134

P - 72

S - 200

Cl -349

Ar 35

4

K - 48

Ca 10

Sc -18

Ti -8

V - 51

Cr - 64.

Mn

Fe -16

Co - 64.

Ni -112

Cu -118

Zn 47

Ga -29

Ge -116

As -78

Separate -195

Br -325

Kr 39

5

Rb - 47

Mr.

And - 30

Zr -41

Nb -86.

Mo - 72

Tc - 53

Ru -101

Rh - 110.

Pd - 54

Ag -126

Cd 32

In -29

Sn -116

Sb -103

You -190

I -295

Xe 41

6

Cs -45.

Ba

Lu

Hf

Ta -31

W - 79

Re -14

You -106

Go -151

Pt -205

Au - 223

Hg 61

Tl - 20

Pb - 35

Bi -91

Po -183

At -270

Rn 41

7

Fr - 44

Ra

Lr

Rf

Db

Sg

Bh

Hs

Mt

Ds

Rg

Cn

Nh

Fl

Mc

Lv

Ts

Og

Periodic table of electronic affinities (with the reverse sign, for the previous definition), in kJ/mol