Carbon and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Carbon and Its Inorganic Compounds

5 lessons

Lesson 4 of 5

Carbonates, Hydrogencarbonates and Hard Water

25 min

Why the mandazi puffs up

Some cooks add a pinch of baking soda to mandazi dough and nothing else: no vinegar, no lemon juice, no acid of any kind. Yet the dough still swells with bubbles in the hot oil. In S2 you made carbon dioxide by adding an acid to a carbonate. Here there is no acid, only heat. Baking soda is sodium hydrogencarbonate, and heat alone is enough to break it down and release carbon dioxide. Many carbonates and hydrogencarbonates do the same, and the pattern tells you a lot about the metals they contain.

Definition

Carbonate: a salt containing the carbonate ion, $\mathrm{CO_3^{2-}}$ , for example sodium carbonate, $\mathrm{Na_2CO_3}$ , and calcium carbonate, $\mathrm{CaCO_3}$ . Hydrogencarbonate (bicarbonate): a salt containing the $\mathrm{HCO_3^-}$ ion, for example sodium hydrogencarbonate, $\mathrm{NaHCO_3}$ . A hydrogencarbonate is an acid salt: it still has a hydrogen that can be replaced by a metal.

Where carbonates are found

Calcium carbonate is one of the commonest compounds in the Earth's crust. It makes up limestone, chalk and marble, and the shells of snails and the eggs of birds. Magnesium carbonate occurs with it in dolomite. Copper forms the green and blue minerals malachite and azurite, which are basic copper carbonates.

Solubility. The only common carbonates that dissolve in water are those of sodium, potassium and ammonium; all the others are insoluble, which is why calcium carbonate precipitates in so many reactions. Sodium and potassium hydrogencarbonates are soluble solids. Calcium and magnesium hydrogencarbonates exist only in solution: if the water is evaporated, the carbonate is left instead.

The action of heat on carbonates

Experiment. Heat a small amount of each carbonate in a test tube and pass any gas given off through limewater (Figure 4.1). Watch the colour of the solid while it is hot and again after it has cooled.

Figure 4.1: Heating a metal carbonate and testing the gas with limewater.

Warning

Safety. Wear eye protection. Before you stop heating, lift the delivery tube out of the limewater. If the tube cools while its end is still in the liquid, the gas inside contracts and limewater is sucked back into the hot test tube, which can crack it (suck-back). Lead(II) carbonate is toxic and is not used in school; its results are given in Table 4.1 for completeness.

Table 4.1: The effect of heat on some carbonates

Carbonate	Colour before	Colour of residue when hot	Colour when cold	Products
potassium, $\mathrm{K_2CO_3}$	white	white	white	no decomposition
sodium, $\mathrm{Na_2CO_3}$	white	white	white	no decomposition
calcium, $\mathrm{CaCO_3}$	white	white (glows)	white	$\mathrm{CaO}$ + $\mathrm{CO_2}$ , only with very strong heat
magnesium, $\mathrm{MgCO_3}$	white	white	white	$\mathrm{MgO}$ + $\mathrm{CO_2}$
zinc, $\mathrm{ZnCO_3}$	white	yellow	white	$\mathrm{ZnO}$ + $\mathrm{CO_2}$
lead(II), $\mathrm{PbCO_3}$	white	reddish-brown	yellow	$\mathrm{PbO}$ + $\mathrm{CO_2}$
copper(II), $\mathrm{CuCO_3}$	green	black	black	$\mathrm{CuO}$ + $\mathrm{CO_2}$
ammonium, $\mathrm{(NH_4)_2CO_3}$	white	no residue: the whole solid turns to gases		$\mathrm{NH_3}$ + $\mathrm{CO_2}$ + $\mathrm{H_2O}$

Most metal carbonates break down into the metal oxide and carbon dioxide, which turns the limewater milky:

$\mathrm{CuCO_3(s) \rightarrow CuO(s) + CO_2(g)} \qquad \mathrm{ZnCO_3(s) \rightarrow ZnO(s) + CO_2(g)}$

$\mathrm{(NH_4)_2CO_3(s) \rightarrow 2NH_3(g) + CO_2(g) + H_2O(g)}$

The pattern. The carbonates of the most reactive metals, potassium and sodium, do not decompose at the temperature of a Bunsen flame. The further down the reactivity series the metal is, the more easily its carbonate breaks down: calcium carbonate needs very strong heat, while copper(II) carbonate decomposes quickly. A reactive metal forms a very stable carbonate.

Definition

Thermal decomposition: the breaking down of a compound into two or more simpler substances by heat alone. Example: $\mathrm{CaCO_3(s) \rightarrow CaO(s) + CO_2(g)}$ in a lime kiln.

Hydrogencarbonates break down more easily

Even sodium hydrogencarbonate, whose metal is very reactive, decomposes at only about 80 °C to 100 °C. This is what happens to baking soda in hot oil or in an oven:

$\mathrm{2NaHCO_3(s) \rightarrow Na_2CO_3(s) + H_2O(g) + CO_2(g)}$

Calcium hydrogencarbonate breaks down as soon as its solution is heated, depositing calcium carbonate:

$\mathrm{Ca(HCO_3)_2(aq) \rightarrow CaCO_3(s) + H_2O(l) + CO_2(g)}$

Table 4.2: Telling a carbonate from a hydrogencarbonate

Test	Sodium carbonate	Sodium hydrogencarbonate
Add dilute hydrochloric acid; pass the gas into limewater	fizzes; limewater turns milky	fizzes; limewater turns milky
Heat the dry solid; pass any gas into limewater	no gas (only water if the crystals are hydrated)	carbon dioxide (limewater milky) and drops of water at the cool mouth of the tube
Add magnesium sulfate solution to a cold solution	white precipitate of magnesium carbonate at once	no precipitate until the mixture is boiled

The acid test shows only that the solid is a carbonate or hydrogencarbonate; the heating test or the magnesium sulfate test tells you which.

Carbonates and hard water

In S1 you learned that water is hard when it contains dissolved calcium or magnesium ions, and that temporary hardness is removed by boiling while permanent hardness is not. Now you can follow the carbonate chemistry behind it.

How temporary hardness forms. Rainwater contains dissolved carbon dioxide. As it seeps through limestone it slowly converts insoluble calcium carbonate into soluble calcium hydrogencarbonate:

$\mathrm{CaCO_3(s) + H_2O(l) + CO_2(aq) \rightarrow Ca(HCO_3)_2(aq)}$

Over thousands of years this dissolves caves out of limestone rock. Where the water drips from a cave roof it loses carbon dioxide to the air and the reaction runs backwards, leaving calcium carbonate behind. Drip by drop, this builds stalactites hanging from the roof and stalagmites rising from the floor.

Softening water: the chemistry

Boiling (temporary hardness only): $\mathrm{Ca(HCO_3)_2(aq) \rightarrow CaCO_3(s) + H_2O(l) + CO_2(g)}$ . The calcium carbonate forms scale in kettles and pipes.
Adding the right amount of limewater (temporary hardness only, used for large volumes where boiling would cost too much fuel): $\mathrm{Ca(HCO_3)_2(aq) + Ca(OH)_2(aq) \rightarrow 2CaCO_3(s) + 2H_2O(l)}$ . Too much limewater would add calcium ions and make the water hard again.
Adding washing soda, sodium carbonate (both kinds): the carbonate ions precipitate the calcium and magnesium ions, whatever their partner ion: $\mathrm{Ca^{2+}(aq) + CO_3^{2-}(aq) \rightarrow CaCO_3(s)}$ .
Ion exchange (both kinds): the water flows through a column of resin beads that hold sodium ions. The resin (written R) takes in calcium ions and releases sodium ions, which do not harden water: $\mathrm{Na_2\text{R}(s) + Ca^{2+}(aq) \rightarrow Ca\text{R}(s) + 2Na^+(aq)}$ . When the resin is full of calcium it is regenerated by washing it with concentrated salt solution.

Table 4.3: Hard water: for and against

Advantages	Disadvantages
its calcium helps build bones and teeth	wastes soap, which forms scum before it lathers
many people prefer its taste	scale in kettles and boilers wastes fuel, because scale conducts heat poorly
it dissolves less lead from old lead pipes	scale narrows and can block hot-water pipes
it is good for brewing and for animals that build shells	it leaves marks on washed clothes and glassware

Worked example

Problem. 5.0 g of copper(II) carbonate is heated until there is no further change. (a) What mass of black solid remains? (b) What volume of carbon dioxide is given off at room temperature and pressure? (Cu = 63.5, C = 12, O = 16; molar volume $24\,\mathrm{dm^3\,mol^{-1}}$ )

Solution. $\mathrm{CuCO_3(s) \rightarrow CuO(s) + CO_2(g)}$ . $M(\mathrm{CuCO_3}) = 63.5 + 12 + 48 = 123.5\,\mathrm{g\,mol^{-1}}$ ; $M(\mathrm{CuO}) = 79.5\,\mathrm{g\,mol^{-1}}$ .

$n(\mathrm{CuCO_3}) = \frac{5.0}{123.5} = 0.0405\,\mathrm{mol}$

(a) 1 mol $\mathrm{CuCO_3}$ gives 1 mol CuO: mass $= 0.0405 \times 79.5 = 3.2\,\mathrm{g}$ .

(b) 1 mol $\mathrm{CuCO_3}$ gives 1 mol $\mathrm{CO_2}$ : volume $= 0.0405 \times 24 = 0.97\,\mathrm{dm^3}$ .

Check: the mass lost, $5.0 - 3.2 = 1.8\,\mathrm{g}$ , should be the carbon dioxide: $0.0405 \times 44 = 1.8\,\mathrm{g}$ . It is.

Worked example

Problem. A student shakes $50\,\mathrm{cm^3}$ of each water sample with soap solution, added $1\,\mathrm{cm^3}$ at a time, until a lasting lather forms. She repeats the test with fresh samples that have been boiled and cooled. Her results (illustrative) are below. Which sample has only temporary hardness, which only permanent hardness, and which both?

Sample	Soap needed before boiling / $\mathrm{cm^3}$	Soap needed after boiling / $\mathrm{cm^3}$
A (rainwater)	1	1
B (well water)	13	1
C (spring water)	9	9
D (borehole water)	15	6

Solution. Soft water needs about $1\,\mathrm{cm^3}$ of soap, as rainwater (A) shows. B needs much more before boiling but only $1\,\mathrm{cm^3}$ after, so boiling removed all its hardness: temporary only. C needs the same amount before and after, so boiling removed none: permanent only. D improves from 15 to $6\,\mathrm{cm^3}$ but is still hard after boiling: it has both kinds.

Tip

Low in the series, easy to break. K and Na carbonates survive the Bunsen flame; carbonates of metals further down the reactivity series break down more and more easily. Colour clues: copper(II) carbonate goes green to black; zinc oxide is yellow when hot and white when cold.

FuseSchool: Making Hard Water Soft (4 min). Watch how an ion-exchange column swaps calcium ions for sodium ions, and compare it with the washing-soda method above.

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Practice

Write the formulae of the carbonate ion and the hydrogencarbonate ion, and name one salt of each.
Which of these carbonates does not decompose when heated with a Bunsen burner: zinc carbonate, sodium carbonate, copper(II) carbonate? Explain the pattern.
A white solid turns yellow when heated strongly and white again on cooling, and gives a gas that turns limewater milky. Name the solid and write the equation.
What is left in the test tube when ammonium carbonate is heated strongly? Write the equation.
Write the equation for the thermal decomposition of sodium hydrogencarbonate, and explain why it makes cakes rise.
Describe two tests that would tell sodium carbonate from sodium hydrogencarbonate.
What mass of calcium oxide is formed when 25 g of calcium carbonate is completely decomposed? (Ca = 40, C = 12, O = 16)
A town takes its water from boreholes in limestone. Evaluate boiling, washing soda and ion exchange as ways of softening its supply.

Show answersHide answers

$\mathrm{CO_3^{2-}}$ , e.g. sodium carbonate $\mathrm{Na_2CO_3}$ ; $\mathrm{HCO_3^-}$ , e.g. sodium hydrogencarbonate $\mathrm{NaHCO_3}$ .
Sodium carbonate. Sodium is very reactive, and the carbonates of the most reactive metals are the most stable; those of metals lower in the reactivity series (zinc, copper) decompose.
Zinc carbonate: $\mathrm{ZnCO_3(s) \rightarrow ZnO(s) + CO_2(g)}$ (the zinc oxide is yellow when hot, white when cold).
Nothing: ammonium carbonate decomposes completely into gases: $\mathrm{(NH_4)_2CO_3(s) \rightarrow 2NH_3(g) + CO_2(g) + H_2O(g)}$ .
$\mathrm{2NaHCO_3(s) \rightarrow Na_2CO_3(s) + H_2O(g) + CO_2(g)}$ . The carbon dioxide forms bubbles in the dough, which expand in the heat, so it rises.
Heat each dry solid: only the hydrogencarbonate gives carbon dioxide (limewater turns milky). Or add magnesium sulfate solution to cold solutions: only the carbonate gives a white precipitate at once.
$\mathrm{CaCO_3(s) \rightarrow CaO(s) + CO_2(g)}$ ; $n = \frac{25}{100} = 0.25\,\mathrm{mol}$ CaCO₃ gives 0.25 mol CaO; mass $= 0.25 \times 56 = 14\,\mathrm{g}$ .
Water from limestone has mainly temporary hardness. Boiling works but needs a great deal of fuel for a whole town and leaves scale, so it suits single households. Washing soda removes both kinds cheaply, but the precipitate must be settled or filtered out and it adds sodium ions. Ion exchange removes both kinds and gives very soft water, but the resin costs money and must be regenerated with salt. Adding a calculated amount of limewater is often the cheapest for a large supply with temporary hardness.

Sources and credits

Chemistry 2e, 18.6: Occurrence, Preparation, and Properties of Carbonates (the solubility of carbonates and hydrogencarbonates, and how caves, stalactites and stalagmites form), Paul Flowers, Klaus Theopold, Richard Langley and William R. Robinson, OpenStax, via LibreTexts, licensed under CC BY 4.0. Adapted by Ganzaa.
Photos: Malachite and azurite, James St. John, CC BY 2.0; Stalactites at Rawhiti Cave entrance.jpg, Pseudopanax, public domain; both via Wikimedia Commons.
Video: Making Hard Water Soft, FuseSchool – Global Education, embedded from YouTube.
The diagram, tables, worked examples and Practice items were written and drawn by Ganzaa.

1Why the mandazi puffs up

2Where carbonates are found

3The action of heat on carbonates

4Hydrogencarbonates break down more easily

5Carbonates and hard water

6Softening water: the chemistry