Group 2

The Group 2 cations: radius, charge density and polarising power

Every Group 2 metal loses the two electrons in its outer s subshell to form a 2+ ion with the electron configuration of the noble gas before it: Mg2+ is [Ne] and Ba2+ is [Xe]. The charge is the same for all four ions, but the size is not. The ionic radius increases steadily down the group from Mg2+ (0.065 nm) to Ca2+ (0.099 nm), Sr2+ (0.113 nm) and Ba2+ (0.135 nm), because each element down the group has one more occupied shell of electrons. The charge stays at 2+ while the volume grows, so the charge density of the cation - the charge packed into each unit of volume - falls from magnesium to barium.

Charge density decides how hard the cation pulls on the electron clouds of the ions around it, so almost every trend in this topic follows from that one change. A cation attracts the electron cloud of a neighbouring anion and distorts it, and this distortion is called polarisation. The ability of a cation to cause it is its polarising power: Mg2+ is small and highly charged, so it polarises strongly, while Ba2+ is large and polarises only weakly.

The anion matters just as much. Large anions such as the carbonate ion, CO3^2-, the nitrate ion, NO3-, and the sulfate ion, SO4^2- spread their outer electrons over a big volume, so those electrons are held loosely and the ion is easily distorted - it is polarisable. Polarisation therefore matters most in those compounds and hardly at all in an oxide, where O^2- is small and holds its electrons tightly. Remember the pairing: a strongly polarising cation with a polarisable anion is the combination that weakens the bonds inside the anion.

Thermal stability of the carbonates

A Group 2 carbonate decomposes on strong heating to the oxide and carbon dioxide: MCO3(s) -> MO(s) + CO2(g), for example CaCO3(s) -> CaO(s) + CO2(g). You meet the reaction in the laboratory, where the carbon dioxide evolved turns limewater milky, and on a huge scale in a lime kiln, where limestone is decomposed at around 1000 °C to make quicklime.

Thermal stability increases down the group. Magnesium carbonate starts to decompose on gentle heating, calcium carbonate needs about 900 °C, and strontium and barium carbonates need temperatures of well over 1000 °C. Said the other way round, the smaller the cation, the more easily its carbonate breaks down.

The reason is polarisation of the large carbonate ion. The cation drags electron density away from one of the C-O bonds in the anion, which is weakened and lengthened, so less energy is needed to break the ion apart and release the carbon dioxide. Mg2+ has the highest charge density of the four cations, so it distorts the carbonate ion most and MgCO3 is the least stable; Ba2+ barely distorts it, so BaCO3 survives to the highest temperature.

Two points decide whether you score here. First, the bond that breaks is inside the carbonate ion, not the ionic bond between the metal ion and the anion, so an answer about the strength of the metal-oxygen bond misses the point. Second, the same idea explains why lithium carbonate decomposes far more readily than sodium or potassium carbonate: Li+ is unusually small for a Group 1 ion and has the high charge density and the polarising power of a Group 2 cation.

Thermal stability of the nitrates

A Group 2 nitrate decomposes on heating to the oxide, brown nitrogen dioxide and oxygen: 2M(NO3)2(s) -> 2MO(s) + 4NO2(g) + O2(g), for example 2Mg(NO3)2(s) -> 2MgO(s) + 4NO2(g) + O2(g). The equation is written with whole numbers because one molecule of oxygen is released when two nitrate ions break down. The brown gas is unmistakable and slightly acidic, so it turns damp blue litmus paper red, and the residue of metal oxide left in the tube is basic and gives an alkaline solution in water.

The trend is the same as for the carbonates: thermal stability increases down the group, so magnesium nitrate decomposes at a lower temperature than barium nitrate.

Nitrates are also less thermally stable than the corresponding carbonates. The nitrate ion is larger than the carbonate ion and carries only one negative charge instead of two, so its charge density is lower and it is more easily polarised. The cation distorts its electron cloud more strongly, an N-O bond is weakened, and the compound breaks down at a lower temperature. Comparing the two decompositions is a favourite question: a carbonate gives one gas, carbon dioxide, while a nitrate gives two, nitrogen dioxide and oxygen.

For contrast, the nitrates of sodium and potassium behave differently: 2NaNO3(s) -> 2NaNO2(s) + O2(g), giving the nitrite rather than the oxide. The Na+ and K+ ions are large and only weakly polarising, so they never distort the nitrate ion enough to break it right down to the oxide. Lithium nitrate, with its small, highly polarising Li+ ion, is the exception among the Group 1 nitrates and decomposes like a Group 2 nitrate to give Li2O, NO2 and O2.

Enthalpy change of solution: lattice enthalpy against hydration enthalpy

Two enthalpy changes decide whether an ionic solid dissolves. The standard lattice enthalpy is the enthalpy change when one mole of the ionic compound is formed from its gaseous ions. It is strongly exothermic because oppositely charged ions attract one another, so it carries a negative sign, for example -2995 kJ mol-1 for Mg(OH)2. The standard enthalpy change of hydration is the enthalpy change when one mole of gaseous ions becomes aqueous ions surrounded by water molecules; it is also exothermic, because ion-dipole attractions are formed with the polar water molecules, for example -1920 kJ mol-1 for Mg2+ and -460 kJ mol-1 for OH-. The standard enthalpy change of solution is the enthalpy change when one mole of the solid dissolves in enough water to give a very dilute solution, and it may be exothermic or endothermic.

Dissolving is a two-step process, and a simple enthalpy cycle gives the answer. Step one separates the lattice into gaseous ions, which costs energy equal to the lattice dissociation enthalpy - the lattice enthalpy written with the opposite sign. Step two hydrates those gaseous ions, which releases energy equal to the sum of the hydration enthalpies. So: enthalpy change of solution = -lattice enthalpy + sum of the enthalpies of hydration, when the lattice enthalpy is given as the negative formation value.

Work through magnesium hydroxide and barium hydroxide with the data above. For Mg(OH)2 the sum of the hydration enthalpies is -1920 + 2(-460) = -2840 kJ mol-1, and the lattice enthalpy is -2995 kJ mol-1, so the enthalpy change of solution is -2840 - (-2995) = +155 kJ mol-1 and dissolving is endothermic. For Ba(OH)2 the sum is -1360 + 2(-460) = -2280 kJ mol-1 and the lattice enthalpy is -2235 kJ mol-1, so the enthalpy change of solution is -2280 - (-2235) = -45 kJ mol-1 and dissolving is exothermic.

Why the hydroxides become more soluble down the group

The hydroxides become more soluble from magnesium to barium. Magnesium hydroxide dissolves only to a very small extent, calcium hydroxide is slightly soluble - a saturated solution in water at room temperature holds about 1.7 g dm-3, and that solution is limewater - strontium hydroxide is appreciably more soluble again, and barium hydroxide dissolves well enough that a saturated solution is a strong alkali of reliable concentration.

Both of the enthalpy changes that decide dissolving fall down the group. The cation gets bigger, so the ions sit further apart in the lattice and the lattice enthalpy becomes less exothermic; and the larger cation attracts water molecules less strongly, so the enthalpy of hydration of the cation also becomes less exothermic. Which way the solubility moves depends on which of the two falls faster, and the size of the anion decides that.

The hydroxide ion is small, so the distance between the cation and the anion matters a great deal to the lattice. Replacing Mg2+ by the much bigger Ba2+ increases that separation sharply, and the lattice enthalpy falls a long way, from -2995 to -2235 kJ mol-1. The enthalpy of hydration of the cation falls too, but not as far, from -1920 to -1360 kJ mol-1. The lattice step therefore loses more energy than the hydration step, and the enthalpy change of solution moves from +155 kJ mol-1 for Mg(OH)2 through +80 for Ca(OH)2 to -45 for Ba(OH)2 - that is, from endothermic towards exothermic - so dissolving becomes progressively more favourable and the solubility rises.

Because more solid dissolves, the concentration of OH- ions in a saturated solution also rises from magnesium to barium, so the alkalinity of the saturated solution increases down the group, even though all four hydroxides are strong bases once they are in solution. Limewater is a weak alkali only because so little calcium hydroxide dissolves in it.

Why the sulfates become less soluble down the group

The sulfates run the other way: they become less soluble from magnesium to barium. Magnesium sulfate is very soluble and its heptahydrate, MgSO4.7H2O, is Epsom salts; calcium sulfate is only slightly soluble, and its dihydrate is gypsum; strontium sulfate is sparingly soluble; and barium sulfate is so nearly insoluble that a white precipitate of it is the standard test for sulfate ions.

The explanation uses the same principle with a differently sized anion. The sulfate ion is large, far larger than the hydroxide ion, so the lattice is built around a big anion and the size of the cation makes comparatively little difference to the sum of the attractions. Going down the group the centres of the ions move a little further apart, but they were already far apart relative to the cation, so the lattice enthalpy falls only slowly. The enthalpy of hydration of the cation, however, falls steeply, because it is the small, highly charged 2+ cation that binds water molecules most strongly. The hydration step therefore loses more than the lattice step, the enthalpy change of solution becomes less exothermic and then endothermic, and the solubility falls.

The two cases are worth comparing side by side, because examiners set them together.

Using the trends: predictions and practical consequences

An exam question rarely asks for one fact on its own. It hands you an unfamiliar element - strontium or radium, for instance - and asks you to predict. Sort the question into one of the two trends and apply it in the right direction. The thermal stability of the carbonates and nitrates increases down the group, so radium carbonate and radium nitrate need a higher temperature to decompose than the barium salts. The solubility of the hydroxides increases down the group, so radium hydroxide is even more soluble than barium hydroxide; the solubility of the sulfates decreases down the group, so radium sulfate is even less soluble than barium sulfate, and radium ions coprecipitate with barium sulfate for exactly that reason.

Then give the reason at the level of the ions, as a chain rather than a single word. For thermal stability: the cation radius increases down the group, so charge density and polarising power fall, the large anion is distorted less, the bond inside the anion is left stronger, and the compound decomposes only at a higher temperature. For solubility: the cation gets bigger, so the lattice enthalpy and the enthalpy of hydration of the cation both become less exothermic, the size of the anion decides which of the two falls faster, the enthalpy change of solution changes in one direction, and the solubility follows it.

The practical consequences are worth quoting. Barium sulfate is swallowed as a barium meal before an X-ray of the digestive system precisely because it is essentially insoluble: soluble barium salts are poisonous, and an insoluble one releases almost no Ba2+ ions into the body while its density makes the gut show up clearly. Limewater, the saturated solution of calcium hydroxide, is the standard test for carbon dioxide. Barium hydroxide solution can be made up as a standard alkali and used to standardise acids, whereas magnesium hydroxide is used as the antacid in milk of magnesia, where its very low solubility keeps the concentration of hydroxide ions low.