I read in a textbook that all alkali metal and alkaline earth metal can react with hydrogen directly to give hydride except for beryllium. I want to know the reason why because I know that $\ce{BeH2}$ exists in a polymeric state.
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2$\begingroup$ See this paper: pubs.acs.org/doi/pdf/10.1021/ja00078a057 $\endgroup$– Nilay GhoshCommented Apr 9, 2021 at 7:51
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$\begingroup$ @NilayGhosh this q never got an answer that could stand up to criticism. Maybe summarize key findings from the reference in an answer that will work? $\endgroup$– Oscar LanziCommented Sep 6, 2022 at 19:54
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$\begingroup$ @OscarLanzi Please go ahead and make into an answer. I can see you had deleted your answer 4 years ago due to lack of reference(?) $\endgroup$– Nilay GhoshCommented Sep 10, 2022 at 7:27
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$\begingroup$ That paper shows evidence for BeH, BeH2, and several gaseous oligomers under heroic conditions. BeH is endothermic deltaG = 67 kcal. that combined with the strong bond energies of H2 and crystal lattice of Ge direct reaction would be kinetically and thermally difficult. The answer seems to be they do react if all the activation energies are removed at a low temperature to prevent decomposition. $\endgroup$– jimchmstCommented Sep 21, 2022 at 3:48
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1 Answer
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As described in Ref. 1, the proposed reaction is endothermic and thus unlikely to be favored. In contrast the other alkaline earth metals give exothermic reactions that ought to proceed at least at low enough temperature. So then the question becomes why does beryllium give an endothermic reaction?
One clue lies in the unusually high melting and boiling points of this metal versus other alkaline-earth metals, which do react directly with hydrogen. From the respective Wikipedia articles we render the following melting and boiling points for the first three alkaline earth metals:
Beryllium: 1287°C, 2469°C
Magnesium: 650°C, 1091°C
Calcium: 842°C, 1484°C
This implies that metallic Beryllium has much stronger bonds than other metallic alkaline earth elements, and a reaction of beryllium metal with hydrogen is thereby hindered by those stronger elemental bonding.
In this respect beryllium resembles not the other alkaline earth metals or even post-transition metals in Group 12, but higher-melting transition metals such as titanium or iron. The lack of a reaction between beryllium and hydrogen thereby fits a more general pattern: high-melting metals, which require breaking strong bonds to form salt-like compounds, do not do so with hydrogen. They either fail to react directly with hydrogen at all (unless they are somehow atomized first, as in Ref. 1), like beryllium or iron; or else they instead form interstitial hydrides that allow most of the metal-metal bonding to be preserved like titanium.
Beryllium is also not alone in being atypically hard to melt or boil. The same thing happens with boron versus the rest of Group 13 and carbon versus the rest of Group 14. In all these cases the small size of the electronic core, which is only a $1s$ subshell, allows strong and deep overlap in the delocalized bonds between the atoms, whereas the larger cores of the later-period congeners limits overlap and makes the heavier elements potentially more reactive.
References
- Thomas J. Tague Jr. and Lester Andrews (1993). "Reactions of beryllium atoms with hydrogen: Matrix infrared spectra of novel product molecules". J. Am. Chem. Soc. 115, 25, 12111–12116. https://doi.org/10.1021/ja00078a057
answered May 7, 2018 at 20:34
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$\begingroup$ It really is fairly simple the product is not particularly stable and reacting from the elements has a large activation energy to supply Ge and H atoms. Ge and H atoms react nicely at temperatures near 0 K, the activation energy is supplied elsewhere. $\endgroup$– jimchmstCommented Sep 23, 2022 at 1:56
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$\begingroup$ More is involved than activation energy. Breaking the strong bonds implied by the high melting point is also a thermodynamic barrier. $\endgroup$ Commented Sep 23, 2022 at 10:53
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$\begingroup$ That is what the activation energy is composed of. Add in a strong H-H bond and finally an endothermic compound why does anyone suspect this reaction to happen at ambient temperatures $\endgroup$– jimchmstCommented Sep 24, 2022 at 3:36