Acids, Bases & Salts

Acids are substances that taste sour and turn blue litmus red. Bases taste bitter, feel soapy to touch, and turn red litmus blue. (You should never taste a laboratory chemical to test it — this is only how these substances are known to behave.)

Acids and bases are all around us. Many acids come from fruits and food, while many bases are used in cleaning agents. In water, acids give H⁺ ions and bases give OH⁻ ions, but for this section just focus on their properties and how indicators respond to them.

• Common acids: hydrochloric acid (HCl), sulphuric acid (H₂SO₄), nitric acid (HNO₃), acetic or ethanoic acid (CH₃COOH, the acid in vinegar), citric acid (in citrus fruits such as lemon and orange) and carbonic acid (H₂CO₃, present in soft drinks).
• Common bases: sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)₂) and ammonium hydroxide (NH₄OH).

An indicator is a substance that tells us whether a solution is acidic or basic by changing its colour (or smell). The common indicators you must know are:

• Litmus — a natural dye extracted from lichens. It is purple in a neutral solution, turns red in acids and blue in bases.
• Methyl orange — red in acids and yellow in bases (orange when neutral).
• Phenolphthalein — colourless in acids and pink in bases.

Some substances change their smell in acidic or basic solutions. These are called olfactory indicators. Examples are onion, vanilla and clove. For instance, the smell of onion or vanilla cannot be detected in a basic solution but stays in an acidic one. Such indicators are helpful for testing solutions with visually impaired students.

Acids do not just taste sour and turn blue litmus red. They take part in three important chemical reactions you must know for the exam. In each one, an acid gives a salt plus one other product. Learn the pattern first, then the examples.

1. Acid + Metal → Salt + Hydrogen gas. An acid reacts with an active metal (like zinc, magnesium, aluminium or iron) to form a salt and to release hydrogen gas.

• Example: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
• Test for hydrogen: bring a burning candle (lighted splint) near the gas. It burns with a 'pop' sound. This confirms the gas is hydrogen.

2. Acid + Metal carbonate or Metal hydrogencarbonate → Salt + Water + Carbon dioxide. Here the acid gives three products, and you see brisk bubbling of CO₂.

• With a carbonate: Na₂CO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g)
• With a hydrogencarbonate: NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g)

3. Acid + Metal oxide → Salt + Water. Metal oxides are basic oxides, so an acid neutralises them just like a base. There is no gas here — only a salt and water.

• Example: CuO(s) + 2HCl(aq) → CuCl₂(aq) + H₂O(l)
• The black copper(II) oxide dissolves and the solution turns blue-green because of the CuCl₂ formed.

Limewater test for CO₂. To prove that the gas from reaction 2 is carbon dioxide, pass it through limewater (a solution of Ca(OH)₂). The limewater turns milky because insoluble calcium carbonate forms:

• Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l) — turns milky
• If you keep passing excess CO₂, the milkiness clears because soluble calcium hydrogencarbonate forms: CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq)

Just like acids, bases take part in a few important reactions. Learn these three patterns and you can answer almost any question on the chemical behaviour of bases.

1. Reactive metal + alkali → salt + hydrogen gas. Some metals, such as zinc, react with strong alkalis like sodium hydroxide and set free hydrogen gas. Zinc gives sodium zincate:

Zn(s) + 2NaOH(aq) → Na₂ZnO₂(aq) + H₂(g)

Notice that hydrogen gas is released here — just as it is when a metal reacts with a dilute acid.

2. Base (alkali) + non-metal oxide → salt + water. Non-metal oxides, such as carbon dioxide, react with bases to give a salt and water:

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

2NaOH(aq) + CO₂(g) → Na₂CO₃(aq) + H₂O(l)

An acid also reacts with a base to give a salt and water. Since a non-metal oxide reacts with a base in exactly the same way, we conclude that non-metal oxides are acidic in nature (they are called acidic oxides).

3. Neutralisation. Neutralisation is the reaction between an acid and a base to give a salt and water. The base cancels the effect of the acid, and the acid cancels the effect of the base:

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

2HCl(aq) + Ca(OH)₂(aq) → CaCl₂(aq) + 2H₂O(l)

In short, always remember the master line: Acid + Base → Salt + Water.

• Reactive metal + alkali → salt + H₂ gas (e.g. zinc → sodium zincate).
• Base + non-metal oxide → salt + water (this proves non-metal oxides are acidic).
• Acid + base → salt + water (this is neutralisation).

Every acid, when dissolved in water, gives out hydrogen ions, H⁺. These H⁺ ions are what make a solution acidic. For example, hydrochloric acid ionises in water as HCl(aq) → H⁺(aq) + Cl^-(aq).

But H⁺ ions never stay on their own in water. Each H⁺ ion joins with a water molecule to form a hydronium ion, H₃O⁺: HCl + H₂O → H₃O⁺ + Cl^-. So it is really the hydronium ion, H₃O⁺, that is responsible for acidic behaviour.

This is why an acid shows acidic behaviour only in the presence of water — the H⁺ ions can form only when water is there. Dry hydrogen chloride gas has no free H⁺ ions, so it does not turn dry blue litmus paper red. Add water, and the acidity appears.

Bases behave the opposite way. A base that dissolves in water is called an alkali, and it gives out hydroxide ions, OH^-. For example, NaOH(s) → Na⁺(aq) + OH^-(aq). These OH^- ions make the solution basic.

How strong an acid or base is depends on how many ions it produces:

• A strong acid or strong base ionises completely in water, so it gives a large number of H⁺ or OH^- ions.
• A weak acid or weak base ionises only partly, so it gives far fewer H⁺ or OH^- ions.

Dilution: Mixing an acid or a base with water is called dilution. It lowers the concentration of ions (H₃O⁺ or OH^-) per unit volume. This mixing is highly exothermic — a lot of heat is released. So you must always add acid to water slowly, with stirring. If you add water to a concentrated acid, the heat produced can make the mixture splash out violently and cause severe burns.

The pH scale is a simple way to measure how acidic or basic (alkaline) a solution is. It runs from 0 to 14. A pH less than 7 means the solution is acidic, a pH equal to 7 means it is neutral (like pure water), and a pH greater than 7 means it is basic or alkaline.

Acidity depends on hydrogen ions. The more hydrogen ions (H⁺) a solution contains, the lower its pH and the more acidic it is. As the H⁺ concentration falls (and OH⁻ rises), the pH climbs towards the basic end. The 'p' in pH comes from the German word 'potenz', which means power — so pH is really a measure of the power of hydrogen ions in a solution.

To read pH quickly we use a universal indicator — a mixture of dyes that shows a different colour at each pH value. As a rough guide, red or orange means strongly to weakly acidic, green means neutral, and blue to violet means basic. We just match the colour to a standard chart to find the pH.

pH is not only a lab idea — it quietly controls a great deal of everyday life:

• Body and digestion: our stomach produces hydrochloric acid (HCl) to help digest food. Too much of it causes indigestion and pain. An antacid — a mild base such as milk of magnesia, Mg(OH)₂ — neutralises the excess acid: Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l).
• Tooth decay: it begins when the pH in the mouth falls below 5.5. Bacteria act on left-over sugar and produce acids that eat into the tooth enamel. Cleaning the mouth with a basic toothpaste neutralises this acid and helps prevent decay.
• Soil and crops: plants grow well only within a certain pH range, so farmers test the soil's pH and treat it when it is too acidic or too basic.
• Self-defence in nature: a bee or an ant sting injects methanoic (formic) acid, HCOOH, which causes burning pain. Rubbing on a mild base like baking soda, NaHCO₃, gives relief: HCOOH(aq) + NaHCO₃(aq) → HCOONa(aq) + H₂O(l) + CO₂(g). Stinging nettle leaves also inject methanoic acid.
• Acid rain: when the pH of rain water falls below 5.6, it is called acid rain. Flowing into rivers, it lowers their pH and makes survival difficult for aquatic life, and over time it also damages buildings and monuments.

A salt is the compound formed when an acid reacts with a base. This reaction is called neutralisation. The H⁺ ions of the acid join the OH⁻ ions of the base to form water, and the remaining ions make the salt. In general: acid + base → salt + water. For example, HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l).

Salts can be grouped into families. Members of a family share a common radical. For example, the chloride family — NaCl, KCl, CaCl₂, NH₄Cl — all contain the chloride radical (Cl⁻). The sodium family — NaCl, Na₂SO₄, NaNO₃, Na₂CO₃ — all contain the sodium radical (Na⁺).

The pH of a salt solution depends on the strength of the acid and the base that formed it:

• Strong acid + strong base → a salt whose solution is neutral (pH = 7). Example: NaCl, made from HCl (strong acid) and NaOH (strong base).
• Strong acid + weak base → a salt whose solution is acidic (pH < 7). Example: NH₄Cl, made from HCl (strong acid) and NH₄OH (weak base).
• Weak acid + strong base → a salt whose solution is basic (pH > 7). Examples: Na₂CO₃ (from carbonic acid H₂CO₃ and NaOH) and CH₃COONa (from acetic acid CH₃COOH and NaOH).

Common salt (sodium chloride, NaCl) is obtained from sea water and from underground deposits of rock salt, formed when ancient seas dried up. It is far more than table salt — it is the raw material for making many important chemicals such as sodium hydroxide, baking soda, washing soda and bleaching powder.

The chlor-alkali process passes electricity through concentrated sodium chloride solution (brine). Its name comes from the products — chlorine and the alkali sodium hydroxide:

2NaCl(aq) + 2H₂O(l) → 2NaOH(aq) + Cl₂(g) + H₂(g)

• Chlorine (Cl₂) is released at the anode — used in water treatment (disinfection) and to make bleaching powder, PVC and hydrochloric acid.
• Hydrogen (H₂) is released at the cathode — used as a fuel and in the hydrogenation of oils to make vanaspati (margarine).
• Sodium hydroxide (NaOH) forms in the solution near the cathode — used to make soaps and detergents, and in paper making.

Common salt (NaCl) is the raw material for three very useful salts. For each one, learn its formula, how it is made and its uses.

1. Bleaching powder — calcium oxychloride, CaOCl₂. It is made by passing chlorine over dry slaked lime, Ca(OH)₂:

Ca(OH)₂(s) + Cl₂(g) → CaOCl₂(s) + H₂O(l)

Uses of bleaching powder:

• for bleaching cotton and linen in the textile industry, and wood pulp in paper factories;
• as an oxidising agent in many chemical industries;
• as a disinfectant to make drinking water free from germs (safe to drink).

2. Baking soda — sodium hydrogencarbonate, NaHCO₃. It is a mild, non-corrosive base made from common salt. On heating, it decomposes:

2NaHCO₃(s) → Na₂CO₃(s) + H₂O(g) + CO₂(g)

Uses of baking soda:

• In baking: baking powder is baking soda mixed with a mild edible acid such as tartaric acid. When heated or moistened, the baking soda releases CO₂, whose bubbles make cakes and bread rise and turn soft and spongy. The tartaric acid neutralises the sodium carbonate formed, so the food does not taste bitter.
• As an antacid: being alkaline, it neutralises excess acid in the stomach and relieves indigestion.
• In soda-acid fire extinguishers.

3. Washing soda — sodium carbonate decahydrate, Na₂CO₃·10H₂O. Heating baking soda gives sodium carbonate; recrystallising this sodium carbonate gives washing soda:

Na₂CO₃(s) + 10H₂O(l) → Na₂CO₃·10H₂O(s)

Uses of washing soda:

• as a cleaning agent at home, and in the glass, soap and paper industries;
• in making other sodium compounds, such as borax;
• to remove the permanent hardness of water.

Remember: both baking soda and washing soda are salts of a strong base (NaOH) and a weak acid (H₂CO₃), so their solutions are basic (alkaline).

Water of crystallisation is the fixed number of water molecules present in one formula unit of a salt. Such salts still look and feel dry, but this water is chemically fixed inside the crystal. It often gives the salt its shape and colour.

The classic example is copper sulphate crystals, CuSO₄·5H₂O, which are blue. Each formula unit holds 5 water molecules. On heating, the crystals lose this water and turn white (anhydrous copper sulphate, CuSO₄). If you add a few drops of water to the white powder, it turns blue again. This shows that the water of crystallisation is responsible for the blue colour.

• Copper sulphate (blue vitriol): CuSO₄·5H₂O
• Washing soda: Na₂CO₃·10H₂O
• Gypsum: CaSO₄·2H₂O

Plaster of Paris (POP) is calcium sulphate hemihydrate, CaSO₄·½H₂O, a white powder. It is made by heating gypsum at about 373 K (100 °C), where gypsum loses most of its water of crystallisation: CaSO₄·2H₂O → CaSO₄·½H₂O + 1½H₂O. The formula ½H₂O means that two formula units of CaSO₄ share one water molecule.

When POP is mixed with water, it takes up water and sets into a hard solid in a few minutes, changing back into gypsum: CaSO₄·½H₂O + 1½H₂O → CaSO₄·2H₂O. Because of this, POP is used for making casts to support fractured bones, making moulds and toys, smoothing walls and decoration. Since it reacts with moisture in the air and sets, POP must always be stored in a moisture-proof container.