Trends in Chemical Properties of Group 2 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 2 Elements and Their Compounds

5 lessons

Lesson 3 of 5

Oxides, Hydroxides, Sulphates, Carbonates and Nitrates of Group 2

25 min

Poison or medicine?

Barium chloride is a poison if swallowed. Yet hospitals give patients a whole cup of barium sulphate to drink before an X-ray. Both contain $\mathrm{Ba^{2+}}$ ions. The difference is solubility: barium sulphate is so insoluble that almost no barium ions reach the blood. Solubility, and stability to heat, change in regular ways down Group 2, and in this lesson you will learn the trends and the reasons behind them.

Oxides and hydroxides: bases that get stronger down the group

The oxides MgO, CaO, SrO and BaO are ionic and basic: they react with acids to give a salt and water, $\mathrm{MgO(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2O(l)}$ and, except for magnesium oxide, they react readily with water to form hydroxides: $\mathrm{CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(s)} \qquad \Delta H \approx -65\ \mathrm{kJ\,mol^{-1}}$ This is slaking lime: quicklime (CaO) becomes slaked lime, $\mathrm{Ca(OH)_2}$ . So much heat is released that the water can boil. Magnesium oxide reacts only slowly with water, and beryllium oxide not at all; BeO is amphoteric (Lesson 4).

The hydroxides are bases too. The more of the hydroxide that dissolves, the higher the concentration of $\mathrm{OH^-}$ and the higher the pH: a suspension of magnesium hydroxide has a pH of about 10, saturated calcium hydroxide solution (limewater) about 12, and barium hydroxide solution higher still. Basic strength increases down the group.

Definition

Solubility: the mass of a substance that dissolves in a stated mass of solvent to form a saturated solution at a stated temperature, here grams per 100 g of water at 20 °C. A substance is called sparingly soluble when very little dissolves, as with $\mathrm{Ca(OH)_2}$ (0.173 g per 100 g of water).

Solubility of hydroxides and sulphates: opposite trends

Table 7.4: Solubility at 20 °C, in grams per 100 g of water

Cation	Hydroxide	Sulphate
$\mathrm{Mg^{2+}}$	0.00096	35.1
$\mathrm{Ca^{2+}}$	0.173	0.255
$\mathrm{Sr^{2+}}$	0.41 at 0 °C; 1.77 at 40 °C	0.0132
$\mathrm{Ba^{2+}}$	3.89	0.00024

Figure 7.5: Solubility of Group 2 hydroxides and sulphates at 20 °C, on a logarithmic scale (each grid line is ten times the one below). The source gives no 20 °C value for strontium hydroxide; the dashed line marks the gap. Barium hydroxide is listed in the source as its octahydrate. Data: Solubility table, Wikipedia, compiled from standard handbooks.

Hydroxides become more soluble down the group: magnesium hydroxide is almost insoluble; barium hydroxide dissolves well enough to give a strongly alkaline solution.
Sulphates become less soluble down the group: magnesium sulphate (Epsom salt) is very soluble; barium sulphate is one of the most insoluble salts known.

Why the trends run in opposite directions

Dissolving an ionic solid means breaking up its lattice (which needs energy, the lattice energy) and surrounding the ions with water molecules (which releases energy, the hydration energy). Both get smaller as the cation gets larger down the group. Which one falls faster depends on the anion.

With a small anion such as $\mathrm{OH^-}$ , the lattice energy depends strongly on the cation's size and falls quickly down the group, faster than the hydration energy. Dissolving becomes easier: solubility increases.
With a large anion such as $\mathrm{SO_4^{2-}}$ , the distance between ions is dominated by the big anion, so the lattice energy hardly changes. The hydration energy of the cation still falls, so dissolving becomes harder: solubility decreases.

Thermal stability of carbonates and nitrates

Every Group 2 carbonate decomposes on strong heating to the oxide and carbon dioxide: $\mathrm{MCO_3(s) \rightarrow MO(s) + CO_2(g)}$ and every nitrate gives the oxide, brown nitrogen dioxide and oxygen: $\mathrm{2M(NO_3)_2(s) \rightarrow 2MO(s) + 4NO_2(g) + O_2(g)}$ The temperature needed rises down the group (Figure 7.6): magnesium carbonate decomposes in a Bunsen flame, but barium carbonate needs far more heat. The nitrates follow the same order. (Compare Group 1, where only lithium's carbonate and nitrate behave like this.)

Figure 7.6: Approximate decomposition temperatures of the Group 2 carbonates, taken as the temperature at which the pressure of $\mathrm{CO_2}$ over the solid reaches one atmosphere. Data: N. N. Greenwood and A. Earnshaw, Chemistry of the Elements, 2nd edition, 1997; measured values depend on how the test is done.

Why: the polarising power of the cation

A small, highly charged cation has a high charge density. It pulls the electron cloud of the large carbonate or nitrate ion towards itself, polarising it and weakening the bonds inside the anion (as you saw for $\mathrm{Li^{+}}$ in Unit 6). The anion then breaks up more easily into $\mathrm{O^{2-}}$ and $\mathrm{CO_2}$ (or $\mathrm{NO_2}$ and $\mathrm{O_2}$ ). Down the group the cation grows from 72 pm ( $\mathrm{Mg^{2+}}$ ) to 135 pm ( $\mathrm{Ba^{2+}}$ ), its polarising power falls, and the compound becomes more stable. A useful measure is charge divided by radius:

Table 7.5: Charge density of the cations

Ion	$\mathrm{Mg^{2+}}$	$\mathrm{Ca^{2+}}$	$\mathrm{Sr^{2+}}$	$\mathrm{Ba^{2+}}$	$\mathrm{Na^{+}}$ (for comparison)
Ionic radius / pm	72	100	118	135	102
$\frac{\text{charge}}{\text{radius}}$ / pm⁻¹	0.028	0.020	0.017	0.015	0.0098

Every Group 2 ion polarises more strongly than $\mathrm{Na^{+}}$ , which is why Group 2 carbonates decompose and sodium carbonate does not. The hydroxides follow the same pattern: $\mathrm{Mg(OH)_2}$ decomposes to MgO and water on heating, while $\mathrm{Ba(OH)_2}$ survives a Bunsen flame.

Worked example 1: mass change on heating

5.00 g of pure calcium carbonate is heated strongly until its mass stops changing. Calculate the final mass and the mass of carbon dioxide lost. $M(\mathrm{CaCO_3}) = 100.09$ , $M(\mathrm{CaO}) = 56.08\ \mathrm{g\,mol^{-1}}$ .

Solution. $\mathrm{CaCO_3 \rightarrow CaO + CO_2}$ : 1 mol gives 1 mol CaO. $n(\mathrm{CaCO_3}) = \frac{5.00}{100.09} = 0.0500\ \mathrm{mol}$ $m(\mathrm{CaO}) = 0.0500 \times 56.08 = 2.80\ \mathrm{g} \qquad m(\mathrm{CO_2}) = 5.00 - 2.80 = 2.20\ \mathrm{g}$ "Heating to constant mass" makes sure the decomposition is complete.

Worked example 2: gases from a nitrate

Calculate the total volume of gas, at room temperature and pressure (24.0 dm³ mol⁻¹), formed when 1.48 g of magnesium nitrate is fully decomposed. $M(\mathrm{Mg(NO_3)_2}) = 148.33\ \mathrm{g\,mol^{-1}}$ .

Solution. $\mathrm{2Mg(NO_3)_2 \rightarrow 2MgO + 4NO_2 + O_2}$ : 2 mol of nitrate give 5 mol of gas, so 1 mol gives 2.5 mol. $n(\mathrm{Mg(NO_3)_2}) = \frac{1.48}{148.33} = 9.98 \times 10^{-3}\ \mathrm{mol}$ $n(\text{gas}) = 2.5 \times 9.98 \times 10^{-3} = 0.0249\ \mathrm{mol} \qquad V = 0.0249 \times 24.0 = 0.599\ \mathrm{dm^3}$

Warning

Common mistakes. (1) Saying all Group 2 compounds get less soluble down the group: hydroxides do the opposite of sulphates. (2) Saying a large cation polarises more: polarising power falls as the cation gets bigger. (3) Writing $\mathrm{MNO_2}$ as a product: Group 2 nitrates give the oxide, not a nitrite.

Safety. Nitrogen dioxide is toxic: heat nitrates only in a fume cupboard. Slaking lime is very exothermic and the product is caustic: add water in small amounts, wear goggles, and keep lime dust off skin and out of eyes.

Tip

Down Group 2: hydroxides up, sulphates down, stability up. One cause behind the stability trend: the cation gets bigger, so it polarises the anion less.

Crunch Chemistry: Why does the thermal stability of Group 2 metal carbonates increase down the group? (11 min). Watch for how the size of the cation changes its polarising effect on the carbonate ion, and how this links to the trend in Figure 7.6.

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Practice

Write an equation for the reaction of strontium oxide with water.
Which is the more soluble: calcium sulphate or barium sulphate? Which is the more soluble: calcium hydroxide or barium hydroxide?
Write a balanced equation for the thermal decomposition of calcium nitrate, with state symbols.
Explain why barium carbonate decomposes at a higher temperature than magnesium carbonate.
A sample of magnesium carbonate of mass 2.10 g is heated to constant mass. Calculate the mass of magnesium oxide left. $M(\mathrm{MgCO_3}) = 84.32$ , $M(\mathrm{MgO}) = 40.31\ \mathrm{g\,mol^{-1}}$ .
A pupil says: "Magnesium sulphate is used as a laxative, so barium sulphate would be a good laxative too." Evaluate this using Table 7.4.
Sodium carbonate does not decompose in a Bunsen flame, but magnesium carbonate does. Explain, using Table 7.5.

Show answersHide answers

$\mathrm{SrO(s) + H_2O(l) \rightarrow Sr(OH)_2(aq)}$
Calcium sulphate (0.255 g against 0.00024 g per 100 g of water); barium hydroxide (3.89 g against 0.173 g).
$\mathrm{2Ca(NO_3)_2(s) \rightarrow 2CaO(s) + 4NO_2(g) + O_2(g)}$
$\mathrm{Ba^{2+}}$ (135 pm) is much larger than $\mathrm{Mg^{2+}}$ (72 pm), with the same charge, so its charge density is lower. It polarises (distorts) the carbonate ion less, the C–O bonds are weakened less, and more heat is needed to break the ion into $\mathrm{O^{2-}}$ and $\mathrm{CO_2}$ .
$n = \frac{2.10}{84.32} = 0.0249$ mol; $m(\mathrm{MgO}) = 0.0249 \times 40.31 = 1.00$ g.
Wrong. Magnesium sulphate works because it dissolves (35.1 g per 100 g of water). Barium sulphate is almost insoluble (0.00024 g), so it would pass through unchanged; that is exactly why it is safe for X-ray meals.
$\mathrm{Mg^{2+}}$ has charge density 0.028 pm⁻¹, nearly three times that of $\mathrm{Na^{+}}$ (0.0098 pm⁻¹), because it is smaller and doubly charged. It polarises the carbonate ion far more, so magnesium carbonate decomposes at a much lower temperature.

Quick check

Match each compound to the description that fits it.

Tap one on the left, then its partner on the right.

Key points

Group 2 oxides and hydroxides are basic; basic strength increases down the group. CaO + water gives slaked lime with much heat.
Hydroxides: more soluble down the group. Sulphates: less soluble down the group ( $\mathrm{BaSO_4}$ is almost insoluble).
Carbonates give MO + $\mathrm{CO_2}$ ; nitrates give MO + $\mathrm{NO_2}$ + $\mathrm{O_2}$ . Both become more stable down the group.
Reason: larger cations have lower charge density and polarise the anion less.

Sources and credits

Image: Example of a chemical change using quicklime and water by James Charles Philip, via Wikimedia Commons (from the Digital Public Library of America), public domain.
Image: Magnesium hydroxide solution.jpg by Alvy16, via Wikimedia Commons, licensed under CC BY 4.0.
Data: solubilities from the Solubility table, Wikipedia; decomposition temperatures from Greenwood and Earnshaw, Chemistry of the Elements (2nd edition, 1997); ionic radii are Shannon's six-coordinate values.
Video: Why does the thermal stability of Group 2 metal carbonates increase down the group?, Crunch Chemistry, embedded from YouTube.
The text, charts, tables, worked examples and Practice items were written and drawn by Ganzaa.

1Poison or medicine?

2Oxides and hydroxides: bases that get stronger down the group

3Solubility of hydroxides and sulphates: opposite trends

4Why the trends run in opposite directions

5Thermal stability of carbonates and nitrates

6Why: the polarising power of the cation