Beryllium, the Odd One Out, and the Test for Barium Ions

1A chloride that behaves like a non-metal's

Magnesium chloride melts at 714 °C and its melt conducts electricity well, as an ionic compound should. Beryllium chloride, from the element just above magnesium, melts at only 399 °C, its melt hardly conducts, and it dissolves in some organic solvents. In moist air it fumes, giving off hydrogen chloride, much as aluminium chloride does. Beryllium is a Group 2 metal, yet much of its chemistry is covalent. This lesson explains why, shows how beryllium resembles aluminium, and then turns to a practical question: how do you prove that a solution contains barium ions?

2Why beryllium is different

Three properties of beryllium set it apart from the rest of the group:

• A tiny ion with a high charge density. $\mathrm{B}{\mathrm{e}}^{2+}\backslash mathrm\{Be^\{2+\}\}$ would have a radius of only 45 pm, so its charge density is very high: $\frac{2}{45}=0.044\mathrm{p}{\mathrm{m}}^{-1}\backslash frac\{2\}\{45\}\; =\; 0.044\backslash \; \backslash mathrm\{pm^\{-1\}\}$, against 0.028 for $\mathrm{M}{\mathrm{g}}^{2+}\backslash mathrm\{Mg^\{2+\}\}$. Such an ion polarises any anion so strongly that electron density is shared between the two atoms: the bond becomes largely covalent. A free $\mathrm{B}{\mathrm{e}}^{2+}\backslash mathrm\{Be^\{2+\}\}$ ion never exists in a solid.
• High ionisation energies and electronegativity. Removing two electrons costs 2657 kJ mol⁻¹, and beryllium's electronegativity (1.57) is the highest in the group.
• No d orbitals in the second shell. Beryllium can use only its 2s and 2p orbitals, so it never has more than four pairs of electrons around it: its maximum coordination number is 4, as in $[\mathrm{B}\mathrm{e}({\mathrm{H}}_2\mathrm{O}{)}_4{]}^{2+}\backslash mathrm\{[Be(H\_2O)\_4]^\{2+\}\}$ and $[\mathrm{B}\mathrm{e}(\mathrm{O}\mathrm{H}{)}_4{]}^{2-}\backslash mathrm\{[Be(OH)\_4]^\{2-\}\}$. Magnesium and the heavier metals commonly reach 6.

Figure 7.7: Melting points of the Group 2 chlorides. Beryllium chloride (orange) is far below the ionic chlorides of the other metals. Data: the Wikipedia articles on beryllium, magnesium, calcium, strontium and barium chloride.

3The anomalous compounds of beryllium

Beryllium chloride is covalent. In the hot vapour it is a linear molecule with only four electrons around beryllium, so it is electron-deficient and acts as a Lewis acid. In the solid, each beryllium completes its octet by accepting lone pairs from chlorine atoms of its neighbours, building a long chain (Figure 7.8).

Figure 7.8: Beryllium chloride as a vapour molecule and as a solid chain. Drawn by Ganzaa.

Beryllium chloride is hydrolysed by water. The small, highly charged ion pulls electron density from its water ligands so strongly that they release ${\mathrm{H}}^{+}\backslash mathrm\{H^+\}$, and the solution is acidic:$$[\mathrm{B}\mathrm{e}({\mathrm{H}}_2\mathrm{O}{)}_4{]}^{2+}(\mathrm{a}\mathrm{q})+{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\rightleftharpoons [\mathrm{B}\mathrm{e}({\mathrm{H}}_2\mathrm{O}{)}_3(\mathrm{O}\mathrm{H}){]}^{+}(\mathrm{a}\mathrm{q})+{\mathrm{H}}_3{\mathrm{O}}^{+}(\mathrm{a}\mathrm{q})\backslash mathrm\{[Be(H\_2O)\_4]^\{2+\}(aq)\; +\; H\_2O(l)\; \backslash rightleftharpoons\; [Be(H\_2O)\_3(OH)]^\{+\}(aq)\; +\; H\_3O^\{+\}(aq)\}$$Beryllium oxide and hydroxide are amphoteric, while those of the other metals are purely basic:$$\mathrm{B}\mathrm{e}(\mathrm{O}\mathrm{H}{)}_2(\mathrm{s})+2{\mathrm{H}}^{+}(\mathrm{a}\mathrm{q})\to \mathrm{B}{\mathrm{e}}^{2+}(\mathrm{a}\mathrm{q})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\backslash mathrm\{Be(OH)\_2(s)\; +\; 2H^\{+\}(aq)\; \backslash rightarrow\; Be^\{2+\}(aq)\; +\; 2H\_2O(l)\}$$$$\mathrm{B}\mathrm{e}(\mathrm{O}\mathrm{H}{)}_2(\mathrm{s})+2\mathrm{O}{\mathrm{H}}^{-}(\mathrm{a}\mathrm{q})\to [\mathrm{B}\mathrm{e}(\mathrm{O}\mathrm{H}{)}_4{]}^{2-}(\mathrm{a}\mathrm{q})\backslash mathrm\{Be(OH)\_2(s)\; +\; 2OH^\{-\}(aq)\; \backslash rightarrow\; [Be(OH)\_4]^\{2-\}(aq)\}$$The product in alkali is the tetrahydroxoberyllate ion. Beryllium metal itself dissolves in hot concentrated alkali, giving hydrogen, and it does not react with water or steam.

4The diagonal relationship with aluminium

$\mathrm{A}{\mathrm{l}}^{3+}\backslash mathrm\{Al^\{3+\}\}$ is larger than $\mathrm{B}{\mathrm{e}}^{2+}\backslash mathrm\{Be^\{2+\}\}$ but carries a larger charge, so its charge density ($\frac{3}{53.5}=0.056\mathrm{p}{\mathrm{m}}^{-1}\backslash frac\{3\}\{53.5\}\; =\; 0.056\backslash \; \backslash mathrm\{pm^\{-1\}\}$) is close to beryllium's. That is why beryllium resembles aluminium more than it resembles magnesium.

Table 7.6: Beryllium, magnesium and aluminium compared

5Testing for barium ions

Two tests together identify $\mathrm{B}{\mathrm{a}}^{2+}\backslash mathrm\{Ba^\{2+\}\}$ with confidence.

1. Sulphate test. Add a few drops of dilute sulphuric acid (or sodium sulphate solution) to a solution of the sample. A dense white precipitate of barium sulphate forms at once, and it does not dissolve when dilute hydrochloric acid is added:$$\mathrm{B}{\mathrm{a}}^{2+}(\mathrm{a}\mathrm{q})+\mathrm{S}{\mathrm{O}}_4^{2-}(\mathrm{a}\mathrm{q})\to \mathrm{B}\mathrm{a}\mathrm{S}{\mathrm{O}}_4(\mathrm{s})\backslash mathrm\{Ba^\{2+\}(aq)\; +\; SO\_4^\{2-\}(aq)\; \backslash rightarrow\; BaSO\_4(s)\}$$
2. Flame test. A nichrome wire, cleaned in concentrated hydrochloric acid, is dipped in the solid and held in a hot Bunsen flame: barium salts give a pale (apple) green flame.

Figure 7.9: The sulphate test for barium ions. Drawn by Ganzaa.

Why both tests? Strontium and lead(II) ions also give white sulphate precipitates, and calcium ions do so from concentrated solutions. The flame colour tells them apart.

Table 7.7: Telling Group 2 ions apart