February 2013

# Electron affinity, the oft-forgotten — yet essential – periodic property

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Geoff Rayner-Canham, Grenfell Campus, Memorial University, Corner Brook NL

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“Ionic compounds form because metals want to give up valence electrons and nonmetals want to gain valence electrons” – a convenient fiction for students starting out in chemistry. But the part of the statement referring to metals is as fictitious as the Tooth Fairy. Let us start by looking at the ionization energy and electron affinity of a gaseous sodium atom:

Na(g) → Na+(g) + e− ΔE = +496 kJ∙mol−1
Na(g) + e− → Na−(g) ΔE = −53 kJ∙mol−1

Thus it takes nearly a half of a megajoule of energy per mole to remove an electron from gaseous sodium but addition of an electron releases energy. In other words, sodium “wants” to acquire an electron!

There was a nice analogy, for which I have lost the original source, that: “sodium no more ‘wants’ to donate an electron than you would happily donate your purse or wallet to a mugger.” That is, it is a competition for the electron, for example, with fluorine for which the electron affinity is much higher. So ionic bonding is not benign, but atomic “nature red in tooth and claw,” in other words, a fight for those valence electrons.

F(g) + e− → F−(g) ΔE = −328 kJ∙mol−1

In this article, we will focus on the importance of electron affinity,but as I have discussed previously, lattice energy plays a very significant role in the formation of solid ionic compounds.1

The table below shows the trend in electron affinities. A plot of electron affinities of gaseous atoms (credit: Wikipedia, DePiep)

It can be seen that the alkali metals have higher electron affinities than the alkaline earth metals. There are the three noticeable “spikes” which can be accounted for by completion, or half-completion of the valence levels as follows:

Alkali metals: ns1 → ns2
Group 14 elements: ns2np2 → ns2np3
Halogens: ns2np5 → ns2np6

### The sodide ion

If the formation of the sodide ion, Na−, is energetically favoured, then compounds containing that ion should be feasible. It was in 1974 that James L. Dye and his researchers at Michigan State University synthesized the first known compound containing the sodide ion.2 Dye realized that, in the solid phase, there was little energy needed for the formation of the sodium cation-anion pair:

2 Na(s) → Na+(s) + Na−(s)

The key, then, was to find a way of keeping the two ions separated. To do this, he caged the sodium ion in a bicyclic diaminoether, commonly known as 2,2,2-crypt. The synthesis was successful and gold-coloured crystals of [Na+(C18H36N2O6)]Na− were produced.

From the crystal structure, the radius of the sodide ion was calculated to be 217 pm, close to that of the iodide ion, and the sodide compound has a structure similar to that of the analogous iodide: [Na+(C18H36N2O6)]I− .

In 1987, Concepción and Dye synthesized a very simple compound of the sodide ion: [Li(diaminoethane)2]+Na−.3 It is a concern of mine that the teaching of chemistry is so “fossilized” that a discovery of decades ago has still not altered the mindsets of chemistry instructors. And it is not because the information is obscure, for even Wikipedia has a page on the
alkalides.4

### The auride ion

Looking at the plot of electron affinities, gold stands out as an obvious candidate for anion formation. In fact, the first evidence for the formation of an auride came in 1937 by the equimolar mixing of cesium and gold.5 This transparent yellow compound was shown in 1959 not to be an alloy, but to be Cs+Au−, with a sodium chloride crystal structure. Since then, several other auride compounds have been synthesized, including Cs3AuO,
which has the perovskite crystal structure and contains (Cs+)3(Au−)(O2−).

As can be seen from the plot, the element preceding gold, platinum, has a high electron affinity, too. Thus it should come as no surprise that there is an increasing chemistry of the platinide ion, Pt2− including cesium latinide, Cs2Pt.

### Commentary

For beginners in chemistry, the idea of give and take of electrons resulting in ionic compound formation makes sense. The difficulty arises in dissuading them of the erroneous concept when they advance to a more in-depth study of bonding.

### References

1. G. Rayner-Canham, “Chemical compounds: Why some are nonexistent,” Chem 13 News, pages 4-5, February 2012.
2. F.J. Tehan, B.L. Barnett and J.L. Dye, “Alkali Anions,” Journal of American Chemical Society, 1974, 96, pages 7203-8
3. R. Concepción and J.L. Dye, “Li+(en)2∙Na−: A Simple CrystallinSodide,” Journal of American Chemical Society, 1987, 109, pages 7203-4.
4. “Alkalide,” http://en.wikipedia.org/wiki/Alkalide, accesses December 12, 2012.
5. See: M. Jansen, “Effects of relativistic Motion of Electrons on the Chemistry of Gold and Platinum,”Solid State Sciences, 2005, 7, pages 1464-74.

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