Carbonates, Nitrates and Solubility, and Why Lithium Is Different

1The gas that makes mandazi rise

Baking powder, used to make cakes and many recipes for mandazi rise, contains sodium hydrogencarbonate. In the hot oil or oven it decomposes and releases carbon dioxide, which forms bubbles in the dough. Sodium carbonate, washing soda, does not decompose at all in a kitchen or even in a Bunsen flame. Lithium carbonate does. This lesson compares what heat does to the carbonates, hydrogencarbonates and nitrates of Group 1, explains the pattern with the idea of polarisation, and shows why lithium so often behaves more like magnesium than like sodium.

2Heating the compounds: what happens

Heat a small amount of solid in a boiling tube and test any gas given off: carbon dioxide turns limewater milky; oxygen relights a glowing splint; nitrogen dioxide is a brown, toxic gas.

Table 3.1: The effect of heat (Bunsen burner temperatures)

So the thermal stability of the carbonates and nitrates increases down the group, and lithium's compounds are the least stable.

3Why: polarisation of the anion

The carbonate ion, $\mathrm{C}{\mathrm{O}}_3^{2-}\backslash mathrm\{CO\_3^\{2-\}\}$, and the nitrate ion, $\mathrm{N}{\mathrm{O}}_3^{-}\backslash mathrm\{NO\_3^\{-\}\}$, are large, with electron clouds that are easily distorted. A cation next to them attracts that cloud towards itself. The smaller the cation, the closer it can get and the stronger the pull. The distortion weakens the C–O (or N–O) bonds inside the anion, so less energy is needed to break off $\mathrm{C}{\mathrm{O}}_2\backslash mathrm\{CO\_2\}$ (or $\mathrm{N}{\mathrm{O}}_2\backslash mathrm\{NO\_2\}$) and the compound decomposes at a lower temperature.

$\mathrm{L}{\mathrm{i}}^{+}\backslash mathrm\{Li^\{+\}\}$ (76 pm) is by far the smallest Group 1 ion, so it polarises anions most and its carbonate, nitrate and hydroxide decompose. Down the group the ions get larger ($\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ 102, ${\mathrm{K}}^{+}\backslash mathrm\{K^\{+\}\}$ 138 pm), their charge density falls, and their compounds become more stable. The same idea will explain the Group 2 carbonates in Unit 7, where the 2+ charge makes the effect much stronger.

4Solubility of Group 1 compounds

Almost all compounds of sodium and potassium are soluble in water: chlorides, sulfates, nitrates, carbonates and hydroxides. This is why sodium and potassium salts are used whenever a soluble compound is needed, and why there is no simple precipitation test for $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ or ${\mathrm{K}}^{+}\backslash mathrm\{K^\{+\}\}$ (Lesson 4 uses flame tests instead). The exceptions are lithium compounds with small, highly charged anions: lithium fluoride, lithium carbonate and lithium phosphate are only sparingly soluble, because their lattice energies are large. The solubility of the hydroxides increases down the group, LiOH being the least soluble.

5Lithium: the odd one out

Table 3.2: Lithium resembles magnesium (its diagonal neighbour) more than sodium

The reason is size: $\mathrm{L}{\mathrm{i}}^{+}\backslash mathrm\{Li^\{+\}\}$ (76 pm) is almost the same size as $\mathrm{M}{\mathrm{g}}^{2+}\backslash mathrm\{Mg^\{2+\}\}$ (72 pm) and much smaller than $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ (102 pm), so it polarises anions strongly and its compounds have more covalent character, more like magnesium's. This resemblance between an element and the one diagonally below and to its right is called a diagonal relationship. Lithium reacts with nitrogen even at room temperature: $6\mathrm{L}\mathrm{i}(\mathrm{s})+{\mathrm{N}}_2(\mathrm{g})\to 2\mathrm{L}{\mathrm{i}}_3\mathrm{N}(\mathrm{s})6\backslash mathrm\{Li(s)\}\; +\; \backslash mathrm\{N\_2(g)\}\; \backslash rightarrow\; 2\backslash mathrm\{Li\_3N(s)\}$.