Group 2 Compounds, Thermal Decomposition and Solubility

AS · 11 min

The compounds of the Group 2 metals show three trends that Cambridge examines every year: how the oxides, hydroxides and carbonates behave with water and acids; how easily the nitrates and carbonates decompose on heating; and how the solubilities of the hydroxides and sulfates change down the group. This note covers all three with equations, observations and data, then uses them for the reacting-mass calculations that often appear alongside, and for predictions about compounds you have not met. The underlying explanations of thermal stability and solubility are A Level content and are flagged where they appear.

Oxides with water and acids

The Group 2 oxides are white, ionic solids containing $\ce{O^2-}$ ions, so they are basic oxides.

With water

\ce{MO(s) + H2O(l) -> M(OH)2}

oxide	reaction with water	pH of resulting mixture
$\ce{MgO}$	reacts only slightly; $\ce{Mg(OH)2}$ is almost insoluble	about 9
$\ce{CaO}$	reacts vigorously and exothermically (hisses, steam may form); white solid of $\ce{Ca(OH)2}$ , slightly soluble	about 11–12
$\ce{SrO}$	vigorous; $\ce{Sr(OH)2}$ more soluble	about 13
$\ce{BaO}$	vigorous; $\ce{Ba(OH)2}$ dissolves	about 13–14

The pH rises down the group because the hydroxides become more soluble, so more $\ce{OH-}$ ions enter solution. Calcium oxide is quicklime; adding water to it makes calcium hydroxide, slaked lime.

With dilute acids

Basic oxides neutralise acids, forming a salt and water:

\ce{MO(s) + 2HCl(aq) -> MCl2(aq) + H2O(l)}

\ce{MO(s) + H2SO4(aq) -> MSO4 + H2O(l)}

All the chlorides are soluble. With sulfuric acid, $\ce{MgSO4}$ is soluble, but for Ca, Sr and Ba the sulfate is sparingly soluble or insoluble, so it may coat the solid and slow or stop the reaction, just as with the metals.

Hydroxides with water and acids

The hydroxides are bases. Those that dissolve give alkaline solutions; their solubility increases down the group (see below).

\ce{M(OH)2(s) + 2HCl(aq) -> MCl2(aq) + 2H2O(l)}

\ce{M(OH)2(s) + H2SO4(aq) -> MSO4 + 2H2O(l)}

The ionic equation for each neutralisation in solution is $\ce{H+(aq) + OH-(aq) -> H2O(l)}$ .

Carbonates with water and acids

The Group 2 carbonates are white solids that are insoluble in water (magnesium carbonate very slightly soluble) and do not react with it.

With dilute hydrochloric acid they fizz, giving carbon dioxide and a colourless solution:

\ce{MCO3(s) + 2HCl(aq) -> MCl2(aq) + H2O(l) + CO2(g)}

With dilute sulfuric acid, magnesium carbonate reacts completely, but calcium, strontium and barium carbonates react briefly and then stop as the insoluble sulfate coats the solid:

\ce{BaCO3(s) + H2SO4(aq) -> BaSO4(s) + H2O(l) + CO2(g)}

Thermal decomposition of the carbonates and nitrates

Carbonates

On heating, the Group 2 carbonates decompose to the oxide and carbon dioxide:

\ce{MCO3(s) -> MO(s) + CO2(g)}

For example, $\ce{CaCO3(s) -> CaO(s) + CO2(g)}$ , the reaction in a lime kiln. There is little to see (white solid to white solid), but the gas turns limewater milky.

Nitrates

On heating, the Group 2 nitrates decompose to the oxide, nitrogen dioxide and oxygen:

\ce{2M(NO3)2(s) -> 2MO(s) + 4NO2(g) + O2(g)}

Observations: the white crystals may first melt (or release water, if hydrated), then brown fumes of nitrogen dioxide are given off, a glowing splint is relit by the oxygen, and a white solid (the oxide) remains.

Key result

Thermal stability increases down Group 2.

$\ce{MgCO3}$ and $\ce{Mg(NO3)2}$ decompose most easily (at the lowest temperature).
$\ce{BaCO3}$ and $\ce{Ba(NO3)2}$ need the highest temperature; barium carbonate barely decomposes in a Bunsen flame.
So the further down the group, the more strongly the compound must be heated to decompose it.

Tip

The explanation of this trend is A Level content: down the group the cation gets larger, so its charge density decreases and it polarises (distorts the electron cloud of) the large carbonate or nitrate ion less. The C–O (or N–O) bonds in the anion are weakened less, so the compound is more stable. At AS you need only describe the trend and write the equations.

Solubility of the hydroxides and sulfates

The two trends run in opposite directions, which is why they are often tested together.

	Mg	Ca	Sr	Ba
solubility of hydroxide / mol per 100 g water (approximate)	$2 \times 10^{-5}$	$1.5 \times 10^{-3}$	$3.4 \times 10^{-3}$	$1.5 \times 10^{-2}$
solubility of sulfate / mol per 100 g water (approximate)	$1.8 \times 10^{-1}$	$4.7 \times 10^{-3}$	$7.1 \times 10^{-5}$	$9.4 \times 10^{-7}$

Key result

Hydroxides: solubility increases down the group. $\ce{Mg(OH)2}$ is almost insoluble; $\ce{Ba(OH)2}$ is soluble.
Sulfates: solubility decreases down the group. $\ce{MgSO4}$ is soluble; $\ce{BaSO4}$ is insoluble.

The explanation, in terms of lattice energy and enthalpy change of hydration, is A Level content. At AS you must state the trends and use them.

Consequences you need to know

Test for sulfate ions. Add dilute acid then aqueous barium chloride or barium nitrate: a white precipitate of barium sulfate forms. $\ce{Ba^2+(aq) + SO4^2-(aq) -> BaSO4(s)}$ . The acid first removes carbonate or sulfite ions, which would also give a white precipitate.
Limewater is a saturated solution of calcium hydroxide (slightly soluble, so dilute). Carbon dioxide turns it milky: $\ce{Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)}$ .
Barium meal. Barium compounds are toxic, but barium sulfate is so insoluble that it can be swallowed safely to show the gut on X-rays.
Magnesium hydroxide in indigestion remedies (milk of magnesia) neutralises excess stomach acid; its very low solubility means the suspension is only mildly alkaline.
Agriculture. Powdered limestone ( $\ce{CaCO3}$ ) and slaked lime ( $\ce{Ca(OH)2}$ ) are spread on acidic soils to raise the pH: $\ce{Ca(OH)2 + 2H+ -> Ca^2+ + 2H2O}$ and $\ce{CaCO3 + 2H+ -> Ca^2+ + H2O + CO2}$ .

Worked examples

Equations for thermal decomposition

Write equations for the thermal decomposition of calcium carbonate and of magnesium nitrate. State two observations for the nitrate.

Solution

$\ce{CaCO3(s) -> CaO(s) + CO2(g)}$

$\ce{2Mg(NO3)2(s) -> 2MgO(s) + 4NO2(g) + O2(g)}$

Observations for the nitrate: brown gas (fumes) given off; a glowing splint relights; a white solid remains. (Any two.)

Check the nitrate equation: Mg 2 = 2; N 4 = 4; O $2 \times 6 = 12$ on the left, $2 + 8 + 2 = 12$ on the right.

Mass of oxide from a carbonate

Calculate the mass of calcium oxide formed when $5.00\ \text{g}$ of calcium carbonate is heated to constant mass. ( $A_r$ : Ca 40.1, C 12.0, O 16.0)

Solution

$n(\ce{CaCO3}) = \dfrac{5.00}{100.1} = 0.04995\ \text{mol}$

$1 : 1$ , so $n(\ce{CaO}) = 0.04995\ \text{mol}$ ; $M(\ce{CaO}) = 56.1\ \text{g mol}^{-1}$ .

Mass $= 0.04995 \times 56.1 = 2.80\ \text{g}$ .

Volume of gas from a nitrate

$2.96\ \text{g}$ of anhydrous magnesium nitrate is heated until it has fully decomposed. Calculate the total volume of gas produced at room conditions, and the volume of oxygen in it. ( $A_r$ : Mg 24.3, N 14.0, O 16.0)

Solution

$M(\ce{Mg(NO3)2}) = 24.3 + 2(14.0 + 48.0) = 148.3\ \text{g mol}^{-1}$

$n = \dfrac{2.96}{148.3} = 0.01996\ \text{mol}$

$\ce{2Mg(NO3)2 -> 2MgO + 4NO2 + O2}$ : 2 mol of nitrate gives 4 mol $\ce{NO2}$ and 1 mol $\ce{O2}$ , so 2.5 mol of gas per mole of nitrate.

Total gas $= 2.5 \times 0.01996 = 0.04990\ \text{mol}$ ; volume $= 0.04990 \times 24.0 = 1.20\ \text{dm}^3$ .

Oxygen $= 0.5 \times 0.01996 = 0.009980\ \text{mol}$ ; volume $= 0.240\ \text{dm}^3$ ( $240\ \text{cm}^3$ ).

Exam-style: identifying a carbonate by heating

$2.00\ \text{g}$ of a Group 2 carbonate, $\ce{MCO3}$ , is heated to constant mass. The mass of the residue is $1.404\ \text{g}$ . Identify M. ( $A_r$ : C 12.0, O 16.0)

Solution

Mass lost $=$ mass of $\ce{CO2}$ $= 2.00 - 1.404 = 0.596\ \text{g}$

$n(\ce{CO2}) = \dfrac{0.596}{44.0} = 0.01355\ \text{mol} = n(\ce{MCO3})$

$M(\ce{MCO3}) = \dfrac{2.00}{0.01355} = 147.7\ \text{g mol}^{-1}$

$A_r(\ce{M}) = 147.7 - 60.0 = 87.7$ , so M is strontium ( $A_r$ 87.6).

"Heated to constant mass" means heated, cooled and weighed repeatedly until two consecutive masses agree, so that decomposition is complete.

Exam-hard: purity of limestone and sources of error

A $10.0\ \text{g}$ sample of limestone (impure calcium carbonate) is heated strongly in a crucible to constant mass. The mass decreases by $3.96\ \text{g}$ . The impurities do not decompose.

(a) Calculate the percentage by mass of calcium carbonate in the limestone.

(b) A second student heated the same limestone with an ordinary Bunsen burner for only five minutes and obtained a lower percentage. Explain why, and suggest how the procedure should be improved.

( $M$ : $\ce{CaCO3}$ 100.1, $\ce{CO2}$ 44.0)

Solution

(a) $n(\ce{CO2}) = \dfrac{3.96}{44.0} = 0.0900\ \text{mol} = n(\ce{CaCO3})$

Mass of $\ce{CaCO3} = 0.0900 \times 100.1 = 9.01\ \text{g}$

Percentage $= \dfrac{9.01}{10.0} \times 100 = 90.1\%$

(b) Calcium carbonate needs strong heating to decompose. After a short time with a Bunsen burner, some $\ce{CaCO3}$ remained undecomposed, so less $\ce{CO2}$ was lost, the mass loss was too small and the calculated percentage too low. Improvement: heat strongly (roaring flame, lid ajar) and heat to constant mass: reweigh after further heating until two consecutive masses agree.

(c) Thermal stability increases down Group 2, so barium carbonate decomposes only at a much higher temperature than a Bunsen burner can provide. It would barely decompose, giving almost no mass change.

Watch out

The nitrate equation needs a 2 in front: $\ce{2M(NO3)2 -> 2MO + 4NO2 + O2}$ . " $\ce{M(NO3)2 -> MO + 2NO2 + 1/2O2}$ " is acceptable, but not unbalanced versions.
Thermal stability increases down the group: barium compounds are the hardest to decompose. Students often reverse this.
Hydroxide solubility increases down the group but sulfate solubility decreases. Do not mix them up.
Group 2 carbonates do not dissolve in water. "Calcium carbonate dissolves in water to give an alkaline solution" is wrong.

Exam tip

"Describe the trend in thermal stability" needs a direction and an indication of what it means: "thermal stability increases down the group; a higher temperature is needed to decompose the compound".
When a question asks for observations on heating a nitrate, the marks are for brown gas and glowing splint relights. " $\ce{NO2}$ produced" is a conclusion, not an observation.
In "heat to constant mass" calculations, state that mass lost equals mass of gas evolved, then use moles of gas.
Predictions for Sr or Ra compounds must use the correct trend: more soluble hydroxide, less soluble sulfate, more stable carbonate.

Practical skills

To compare the thermal stability of Group 2 carbonates, heat equal amounts (in moles) of each in a hard-glass test-tube using the same Bunsen flame, and pass the gas through limewater. Measure the time taken for the limewater to turn milky. The dependent variable is that time; control the moles of carbonate, the volume of limewater, the flame height and the distance from the flame. For nitrates, time how long it takes for brown gas to appear, in a fume cupboard ( $\ce{NO2}$ is toxic). Sources of error: judging "milky" or "brown" is subjective; the tubes may be heated unevenly; suck-back of limewater into the hot tube (remove the delivery tube before stopping heating).

Summary

Oxides $\ce{MO}$ are basic: with water form $\ce{M(OH)2}$ (pH rises down the group); with acids form salt + water.
Carbonates are insoluble; with acids they give salt + water + $\ce{CO2}$ ; with $\ce{H2SO4}$ , Ca, Sr and Ba carbonates stop reacting because of the insoluble sulfate.
$\ce{MCO3 -> MO + CO2}$ ; $\ce{2M(NO3)2 -> 2MO + 4NO2 + O2}$ (brown gas, glowing splint relights).
Thermal stability of carbonates and nitrates increases down the group.
Solubility of hydroxides increases down the group; solubility of sulfates decreases.
Uses: $\ce{BaSO4}$ meal; $\ce{BaCl2}$ test for sulfate; limewater test for $\ce{CO2}$ ; $\ce{Mg(OH)2}$ antacid; $\ce{CaCO3}$ and $\ce{Ca(OH)2}$ to neutralise acidic soil.

Practice

Question

Write an equation for the reaction of calcium oxide with water. Give the common names of the reactant and product.
Write an equation, with state symbols, for the reaction of strontium carbonate with dilute hydrochloric acid.
Write an equation for the thermal decomposition of barium nitrate, and state the colour of the gas produced.
State the trend in solubility of the Group 2 hydroxides and of the Group 2 sulfates.
Explain why a solution of barium chloride is used in the test for sulfate ions, and why dilute hydrochloric acid is added first.
Calculate the volume of carbon dioxide, at room conditions, given off when $5.00\ \text{g}$ of calcium carbonate is completely decomposed. ( $M(\ce{CaCO3}) = 100.1$ )
Equal masses of magnesium carbonate and barium carbonate are heated in identical tubes. Predict which produces carbon dioxide first and explain how you could show this experimentally.
A mixture of magnesium sulfate and barium sulfate is shaken with water and filtered. Suggest which compound is found in the filtrate and which remains on the filter paper, and explain.
$4.00\ \text{g}$ of a mixture of calcium carbonate and calcium oxide is heated to constant mass. The final mass is $3.12\ \text{g}$ . Calculate the mass of calcium carbonate in the original mixture. ( $M$ : $\ce{CaCO3}$ 100.1, $\ce{CO2}$ 44.0)
Strontium hydroxide is a white solid. Predict (a) whether it is more or less soluble than calcium hydroxide, (b) the approximate pH of a saturated solution compared with limewater, and (c) the products and an equation for the thermal decomposition of strontium nitrate. Justify each prediction.

Answers

$\ce{CaO(s) + H2O(l) -> Ca(OH)2(s)}$ . Calcium oxide is quicklime; calcium hydroxide is slaked lime.
$\ce{SrCO3(s) + 2HCl(aq) -> SrCl2(aq) + H2O(l) + CO2(g)}$ .
$\ce{2Ba(NO3)2(s) -> 2BaO(s) + 4NO2(g) + O2(g)}$ . Nitrogen dioxide is brown.
Hydroxides: solubility increases down the group. Sulfates: solubility decreases down the group.
Barium sulfate is very insoluble, so $\ce{Ba^2+}$ ions give a white precipitate with even small amounts of sulfate: $\ce{Ba^2+(aq) + SO4^2-(aq) -> BaSO4(s)}$ . The acid is added first to remove carbonate (and sulfite) ions, which would otherwise also form white precipitates with barium ions. Barium sulfate does not dissolve in the acid, but barium carbonate would.
$n(\ce{CaCO3}) = 5.00 / 100.1 = 0.04995\ \text{mol} = n(\ce{CO2})$ ; $V = 0.04995 \times 24.0 = 1.20\ \text{dm}^3$ .
Magnesium carbonate, because thermal stability increases down the group and magnesium carbonate decomposes at a lower temperature. Pass the gas from each tube into the same volume of limewater and time how long it takes to turn milky, using the same flame and the same heating position (and ideally equal moles rather than equal masses).
Magnesium sulfate dissolves (sulfates are more soluble higher in the group), so it is in the filtrate. Barium sulfate is insoluble and remains on the filter paper.
Mass lost $= 4.00 - 3.12 = 0.88\ \text{g}$ of $\ce{CO2}$ . $n(\ce{CO2}) = 0.88 / 44.0 = 0.0200\ \text{mol} = n(\ce{CaCO3})$ . Mass of $\ce{CaCO3} = 0.0200 \times 100.1 = 2.00\ \text{g}$ .
(a) More soluble: solubility of Group 2 hydroxides increases down the group, and Sr is below Ca. (b) Higher pH than limewater (about 13 compared with about 12), because more hydroxide dissolves, giving a higher $\ce{[OH-]}$ . (c) $\ce{2Sr(NO3)2(s) -> 2SrO(s) + 4NO2(g) + O2(g)}$ : strontium oxide, nitrogen dioxide and oxygen, like all Group 2 nitrates; it needs a higher temperature than calcium nitrate because thermal stability increases down the group.

Oxides with water and acids

With water

With dilute acids

Hydroxides with water and acids

Carbonates with water and acids

Thermal decomposition of the carbonates and nitrates

Carbonates

Nitrates

Solubility of the hydroxides and sulfates

Consequences you need to know

Worked examples

Practice

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