4 They use ultrafast two-dimensional infrared spectroscopy to characterise the bonding in HF 2 – in aqueous solution, and see a smooth crossover from a conventional hydrogen bond (with the hydrogen located asymmetrically, covalently bonded to one fluorine) to a ‘short, strong’ hydrogen bond in which the hydrogen occupies a shallow double potential well. While this was recognised decades ago, still this bond is being debated, as a recent paper by Andrei Tokmakoff of the University of Chicago and coworkers attests. The latter bond is striking in that it can become symmetrical, with the hydrogen located at the midpoint of the fluorine atoms. While conventional hydrogen bonds (like those seen in water) feature a localised potential energy minimum, the short strong hydrogen bonds of HF 2 – form a double potential well and approach covalency 3 And while the hydrogen bonds in liquid water are about 25 times weaker than the covalent bonds in H 2O molecules, others – like that in HF 2 – – are stronger and shorter, and generally regarded as delocalised, non-classical three-centre bonds. But it’s been long recognised that there’s some covalency in the hydrogen bond too.
Students typically learn now that a hydrogen bond is essentially electrostatic: a positively polarised hydrogen attached to an electron-withdrawing entity such as oxygen or fluorine is attracted to an electron lone pair on another molecule. Pauling’s uncertainty is understandable although the hydrogen bond celebrated its centenary last year, the debate still hasn’t ceased.
It wasn’t just Pauling’s status that secured the idea, but the fact that he tied it to crystallographic data. Pauling’s 1939 book The Nature of the Chemical Bond cemented the hydrogen bond as a part of the chemist’s lexicon, although he oscillated for a time between viewing it as a purely electrostatic phenomenon between ions or as a resonance between ionic and covalent forms. In a 1925 paper with his graduate student Sterling Hendricks, he showed how such a bond might explain the ion HF 2 –, with the hydrogen sandwiched between fluorines. Instrumental in the acceptance of the hydrogen bond was the intervention of Linus Pauling. 2 But this weird sort of bonding didn’t get a name until Lewis himself, having been won round, wrote in his seminal 1923 book Valence and the Structure of Atoms and Molecules that ‘the most important addition to my theory of valence lies in the suggestion of what has become known as the hydrogen bond’. So it was that in 1920 Wendell Latimer, a young lecturer at Berkeley, and the postdoc Worth Rodebush published a paper on Lewis theory in which they laid out Huggins’ suggestion, crediting him in a footnote. Lewis, however, was sceptical of the whole notion.Īs is often the case, younger researchers were more ready to accept new ideas that older scientists resisted. ‘Huggins,’ he said, ‘there are several interesting ideas in this paper, but there is one you’ll never get chemists to believe: the idea that a hydrogen atom can be bonded to two other atoms at the same time.’ 1 Huggins had advanced this strange idea using the electron-sharing theory of chemical bonding proposed by Gilbert Lewis, depicting, for example, a dimer of hydrogen fluoride in which the four atoms were arranged in a square, with each hydrogen bonded to both fluorines. Might the professor accept these notes as the term paper, with a title added?īray did, but with caveats. In desperation, he showed Bray rough notes he’d made about some of the ‘unsolved problems in chemistry’ that Bray had discussed.
His professor William Bray required his students to write a term paper to pass his course, but Huggins hadn’t yet done that, and the deadline was approaching. In May 1919, an undergraduate chemistry student named Maurice Huggins at the University of California at Berkeley was panicking.
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