At AS you learned the trends; now explain them. Small, highly charged cations polarise big anions and make carbonates and nitrates decompose easily. A tug of war between lattice energy and hydration enthalpy decides whether a sulfate or a hydroxide dissolves.
Group 2 carbonates and nitrates become more thermally stable down the group. A small 2+ cation such as Mg2+ has a high charge density, so it has high polarising power: it distorts the electron cloud of the large CO32− or NO3− ion and weakens its C–O or N–O bonds, so less heat is needed to release CO2 or NO2. Larger cations polarise less. Solubility is linked to the enthalpy change of solution: ΔHsol = ΣΔHhyd − lattice energy. Down the group both the lattice energy and the hydration enthalpies become less exothermic, but at different rates. For sulfates the large anion means the lattice energy changes little while the cation hydration enthalpy falls a lot, so ΔHsol becomes less exothermic and solubility decreases. For hydroxides the small OH− means the lattice energy falls faster, so solubility increases.
Compare the cations next to a carbonate ion in the simulator and watch how much each one distorts the anion; then heat each carbonate and nitrate on the next tab and compare how hot it must get before gas is given off.
| Carbonate | Cation radius / nm | Charge density | ΔH for MCO3 → MO + CO2 / kJ mol−1 |
|---|---|---|---|
| MgCO3 | 0.065 | highest | +101 |
| CaCO3 | 0.099 | ↓ | +178 |
| SrCO3 | 0.113 | ↓ | +235 |
| BaCO3 | 0.135 | lowest | +275 |
Switch to the lattice versus hydration view and watch the two energy terms change as you move down the group for the sulfates. The numbers in the tables below are rounded, illustrative values chosen to show the trend; data-book values differ slightly, and exam questions always give you the data to use.
| Sulfate | ΣΔHhyd / kJ mol−1 | Lattice energy / kJ mol−1 | ΔHsol / kJ mol−1 |
|---|---|---|---|
| MgSO4 | −1920 + (−1160) = −3080 | −2833 | −3080 − (−2833) = −247 |
| BaSO4 | −1360 + (−1160) = −2520 | −2538 | −2520 − (−2538) = +18 |
| change Mg → Ba | 560 less exothermic | 295 less exothermic | 265 less exothermic |
Four quick checks. Each correct answer earns XP and records this skill on your profile.
Which Group 2 carbonate decomposes at the lowest temperature?
What does the cation do to the carbonate ion in thermal decomposition?
How does the solubility of Group 2 sulfates change down the group?
Which expression gives the enthalpy change of solution of MgSO4?
Add sodium hydroxide to the four solutions on the bench: the precipitates run the opposite way to the sulfates, because the numbers in the cycle move in the opposite direction.
| Hydroxide | ΣΔHhyd / kJ mol−1 | Lattice energy / kJ mol−1 | ΔHsol / kJ mol−1 |
|---|---|---|---|
| Mg(OH)2 | −1920 + 2(−460) = −2840 | −2993 | +153 |
| Ba(OH)2 | −1360 + 2(−460) = −2280 | −2228 | −52 |
| change Mg → Ba | 560 less exothermic | 765 less exothermic | 205 more exothermic |
Both terms get less exothermic down the group; the trend depends on which one falls faster, so never explain solubility with just one of them. Polarisation is about the cation distorting the anion, not the anion being attracted to the cation more strongly. And watch the sign in the cycle: with lattice energy defined as formation (exothermic), ΔHsol = ΣΔHhyd minus lattice energy, so −3080 − (−2833) = −247, not −5913.
Unlocks once the four checks above are done. Worth more XP, written to A Level Paper 4 standard.
Hydration enthalpies: Ba2+ −1360, SO42− −1160; lattice energy of BaSO4 −2538 (all kJ mol−1). What is ΔHsol of BaSO4?
Why does the lattice energy of the Group 2 sulfates change only slightly from MgSO4 to BaSO4?
Why are Group 2 hydroxides more soluble down the group?
3.28 g of Ca(NO3)2 (Mr 164.1) is fully decomposed. What total volume of gas forms at room conditions (24.0 dm3 mol−1)?
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