2.3 Salts and Their Everyday Uses

Put a pinch of salt on your tongue and you taste the most familiar chemical in the world. Yet “salt” is a whole family. Bakers use baking soda to make idlis and cakes rise; washerwomen and factories use washing soda; swimming pools and drinking-water plants use bleaching powder; a doctor sets your broken arm in a cast of plaster of Paris. Each of these is a salt, and each comes from the same few ideas you already know: an acid and a base that cancel each other.

In this section you will see how salts are grouped into families, how to tell without any experiment whether a salt solution is acidic, neutral or basic, and how sea salt is turned — by passing electricity through its solution — into three valuable substances: sodium hydroxide, chlorine and hydrogen. Then you will meet the four everyday salts above, learn why some salts keep water inside their crystals, and see how plaster of Paris sets hard.

Salt Families and the pH of Salt Solutions

You met salts in Section 2.1 as the product of a neutralisation reaction. The name of every salt tells where it came from: the first word comes from the base, the second from the acid. Sodium chloride comes from sodium hydroxide and hydrochloric acid; potassium sulphate from potassium hydroxide and sulphuric acid.

Writing the formula of a salt is a matter of making the charges balance (Section 3.2 shows why). The metal ion and the acid's negative ion are used in whatever numbers give a total charge of zero: $\mathrm{Na^{+}}$ with $\mathrm{SO_{4}^{2-}}$ needs two sodium ions, giving $\mathrm{Na_{2}SO_{4}}$; $\mathrm{Ca^{2+}}$ with $\mathrm{Cl^{-}}$ needs two chlorides, giving $\mathrm{CaCl_{2}}$. A salt such as $\mathrm{CuSO_{4}}$ is named after its parts: copper(II) sulphate.

The table shows a pattern. Salts are not all neutral: the pH of a salt solution depends on how strong the parent acid and base were (Section 2.2). The reason is simple. The ions of a strong parent hardly react with water and leave the pH alone. The ion of a weak parent holds on to the $\mathrm{H^{+}}$ or $\mathrm{OH^{-}}$ it came from, and takes some from water, so the solution turns basic or acidic. (The values in the table are typical for $0.1\ \mathrm{mol/L}$ solutions at room temperature.)

Common Salt and the Chlor-alkali Process

Common salt, sodium chloride $\mathrm{NaCl}$, is the salt in your kitchen. It is obtained from sea water, by letting the water evaporate in shallow pans in the sun, and from underground deposits of rock salt, which are mined like coal. Rock salt is often brown because of impurities; it is purified before it is eaten. Salt is essential for life, but it is also one of the most important raw materials of industry.

When electricity is passed through a concentrated solution of sodium chloride in water (called brine), a remarkable reaction takes place: the salt and water change into three new substances, and the solution becomes basic. Two of the products are gases, one at each electrode.

Notice how the equation balances: the left side has two $\mathrm{Na}$, two $\mathrm{Cl}$, four $\mathrm{H}$ and two $\mathrm{O}$, and so does the right side. The mass balances too: $117\ \mathrm{g}$ of $\mathrm{NaCl}$ and $36\ \mathrm{g}$ of water ($153\ \mathrm{g}$) give $80\ \mathrm{g}$ of $\mathrm{NaOH}$, $71\ \mathrm{g}$ of $\mathrm{Cl_{2}}$ and $2\ \mathrm{g}$ of $\mathrm{H_{2}}$ ($153\ \mathrm{g}$).

Baking Soda and Washing Soda

Baking soda, sodium hydrogencarbonate $\mathrm{NaHCO_{3}}$, is made from brine (common salt solution), ammonia and carbon dioxide: $$\mathrm{NaCl(aq) + H_{2}O(l) + CO_{2}(g) + NH_{3}(g) \rightarrow NH_{4}Cl(aq) + NaHCO_{3}(s)}$$ It is a mild, non-corrosive base: its solution has pH about 8.3 (Principle 2.8: weak acid, strong base). It is used in cooking, in antacids (Section 2.2), in soda-acid fire extinguishers and in cleaning.

Washing soda comes from baking soda. Heating baking soda gives sodium carbonate $\mathrm{Na_{2}CO_{3}}$ (soda ash), a white powder. Dissolve it in water and let it crystallise, and it takes ten molecules of water into each crystal: $$\mathrm{Na_{2}CO_{3}(s) + 10H_{2}O(l) \rightarrow Na_{2}CO_{3}\cdot 10H_{2}O(s)}$$ This second, hydrated salt is washing soda. Recovering a salt as purer crystals from its solution in this way is called recrystallisation.

Bleaching Powder, Water of Crystallisation and Plaster of Paris

Bleaching powder is made by passing the chlorine from the chlor-alkali cell over dry slaked lime: $$\mathrm{Cl_{2}(g) + Ca(OH)_{2}(s) \rightarrow CaOCl_{2}(s) + H_{2}O(l)}$$ Its chemical name is calcium oxychloride, with the formula $\mathrm{CaOCl_{2}}$ (often written $\mathrm{Ca(OCl)Cl}$).

Now look at a different kind of salt. Heat a few blue crystals of copper sulphate in a dry boiling tube. The crystals lose their colour and turn white, and drops of water collect on the cool upper part of the tube. Add two or three drops of water to the white powder, and the blue colour returns. The blue crystals therefore contain water, which is part of the crystal, even though the crystals feel dry.

Do not confuse crystal water with dampness. Copper sulphate crystals contain exactly five water molecules for each $\mathrm{CuSO_{4}}$, and they stay blue and dry-looking until you heat them. Pure common salt crystals contain no water of crystallisation at all.

Plaster of Paris is a hydrated salt with only half a molecule of water. It is made by heating gypsum, $\mathrm{CaSO_{4}\cdot 2H_{2}O}$, at $373\ \mathrm{K}$ ($100\ ^{\circ}\mathrm{C}$): $$\mathrm{CaSO_{4}\cdot 2H_{2}O(s) \xrightarrow{373\ \mathrm{K}} CaSO_{4}\cdot\tfrac{1}{2}H_{2}O(s) + 1\tfrac{1}{2}H_{2}O(g)}$$ (Doubling to avoid fractions: $\mathrm{2CaSO_{4}\cdot 2H_{2}O \rightarrow (CaSO_{4})_{2}\cdot H_{2}O + 3H_{2}O}$.) It is a white powder. Mix it with water and it sets, within minutes, into a hard solid mass of gypsum: $$\mathrm{CaSO_{4}\cdot\tfrac{1}{2}H_{2}O(s) + 1\tfrac{1}{2}H_{2}O(l) \rightarrow CaSO_{4}\cdot 2H_{2}O(s)}$$

Exercises