2_4 Group 1 and 2 Compounds: Solubility and Thermal Stability Trends

Once you know how the s-block metals react, the next question is how their compounds behave. Two ideas explain almost every trend. The first is the balance between lattice energy (the energy released when gaseous ions come together as a solid) and hydration energy (the energy released when water molecules surround the ions), which sets solubility. The second is the polarizing power of the cation, which sets how easily a compound breaks down on heating.

Oxides

Alkali metal oxides are strongly basic. The oxide ion takes a proton from water,

\mathrm{O^{2-} + H_2O \rightarrow 2OH^-}

, which is an acid–base reaction, not a redox one, because no oxidation number changes. Group 2 oxides are less basic. BeO is amphoteric and MgO barely dissolves, while CaO, SrO and BaO react with water to give hydroxides. In both groups, basic character increases down the group as the cation becomes larger and less polarizing.

Potassium superoxide has a practical use in breathing apparatus, for example in spacecraft and mine rescue. It absorbs exhaled carbon dioxide and releases oxygen at the same time:

\mathrm{4KO_2(s) + 2CO_2(g) \rightarrow 2K_2CO_3(s) + 3O_2(g)}

Hydroxides

Group 1 hydroxides are very soluble, strongly basic, hygroscopic (they absorb moisture from the air) and stable to heat. LiOH is the exception: it is less soluble and decomposes to

\mathrm{Li_2O}

when strongly heated. Group 2 hydroxides become more soluble down the group. Be(OH)₂ is insoluble and amphoteric, Mg(OH)₂ is very sparingly soluble, Ca(OH)₂ is slightly soluble and Ba(OH)₂ dissolves reasonably well. The hydroxide ion is small, so as the cation grows the lattice energy falls quickly, faster than the hydration energy, and dissolving becomes easier. Mg(OH)₂ and Ca(OH)₂ also lose water on heating to give the oxide.

Two everyday mixtures come from this. A suspension of Mg(OH)₂ is milk of magnesia, a mild antacid. A saturated solution of Ca(OH)₂ is lime water, which turns milky with carbon dioxide because insoluble calcium carbonate forms:

[Ca+2].[OH-].[OH-].O=C=O>>[Ca+2].[O-]C([O-])=O.O

Carbonates

All group 1 carbonates are soluble and, apart from

\mathrm{Li_2CO_3}

, stable at bunsen temperatures. Their solutions are alkaline because carbonate ions take protons from water. Sodium carbonate crystallizes as the decahydrate washing soda,

\mathrm{Na_2CO_3 \cdot 10H_2O}

, which slowly loses water in air to form a white powdery monohydrate. Group 2 carbonates are only very slightly soluble, and all decompose to the oxide and carbon dioxide when heated.

[Ca+2].[O-]C([O-])=O>>[Ca+2].[O-2].O=C=O

Approximate decomposition temperatures of group 2 carbonates

Carbonate	MgCO₃	CaCO₃	SrCO₃	BaCO₃
Decomposes at (°C)	≈ 540	≈ 900	≈ 1280	≈ 1360

Thermal stability rises down the group. A small, doubly charged cation such as

\mathrm{Mg^{2+}}

pulls electron density out of the large carbonate ion, weakening one C–O bond so that

\mathrm{CO_2}

leaves more easily. The small cation also forms an oxide with a very high lattice energy, which makes the product more favourable. A large

\mathrm{Ba^{2+}}

ion distorts the carbonate ion far less, so much more heat is needed.

Nitrates

Every group 1 and group 2 nitrate dissolves in water. Their behaviour on heating splits along the same polarizing-power line. Sodium and potassium nitrates only lose oxygen and form nitrites. Lithium nitrate and the group 2 nitrates break down completely to the oxide, giving off brown nitrogen dioxide and oxygen:

\begin{aligned} \mathrm{2NaNO_3} &\rightarrow \mathrm{2NaNO_2 + O_2} \\ \mathrm{2Mg(NO_3)_2} &\rightarrow \mathrm{2MgO + 4NO_2 + O_2} \end{aligned}

Sulphates

All alkali metal sulphates are soluble. In group 2 the trend runs opposite to the hydroxides, and solubility falls down the group:

\mathrm{BeSO_4}

and

\mathrm{MgSO_4}

dissolve readily,

\mathrm{CaSO_4}

is slightly soluble, and

\mathrm{SrSO_4}

and

\mathrm{BaSO_4}

are almost insoluble. The sulphate ion is large, so changing the cation hardly alters the lattice energy, but hydration energy falls steadily as the cation grows. Barium sulphate's insolubility is why it is safe to swallow for X-ray imaging of the gut and why adding barium chloride is the standard test for sulphate ions.

Summary of group 2 trends going down the group (Mg → Ba)

Compound	Trend down the group	Main reason
Hydroxides	Solubility increases	Lattice energy falls faster than hydration energy
Sulphates	Solubility decreases	Hydration energy falls; lattice energy nearly constant
Carbonates	Thermal stability increases	Larger cation polarizes the anion less
Nitrates	Thermal stability increases	Larger cation polarizes the anion less

Tips for remembering

• Hydroxides up, sulphates down: going down group 2, hydroxides get more soluble while sulphates get less soluble, which is why BaSO₄ is the classic insoluble precipitate.
• Small cation, easy breakdown: small, highly charged cations polarize large carbonate and nitrate ions, so MgCO₃ breaks down far more easily than BaCO₃.
• Na and K nitrates only lose O₂: they stop at the nitrite. Lithium and group 2 nitrates go all the way to the oxide and give brown NO₂.