Acids, Bases and Salts

When an acid dissolves, its molecules ionise: HCl(aq) → H⁺(aq) + Cl⁻(aq). Common acids are hydrochloric acid, HCl, nitric acid, HNO₃, sulfuric acid, H₂SO₄, and ethanoic acid, CH₃COOH, the acid in vinegar. The H⁺ ions give all acids the same properties: they taste sour, turn blue litmus red, and their solutions conduct electricity because they contain ions.

5.1.1The three reactions of acids

Learn these three reactions. Each forms a salt, named after the acid: sulfates from sulfuric acid, nitrates from nitric acid and chlorides from hydrochloric acid.

The strength of an acid is how far it ionises in water (Figure 5.1).

Hydrochloric, nitric and sulfuric acids are strong: HCl(aq) → H⁺(aq) + Cl⁻(aq). Ethanoic acid and citric acid are weak. Only some ethanoic acid molecules ionise, and the ions can join again, so the equation has a reversible arrow: CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq).

5.3.1The hydroxide ion

A hydroxide ion, OH⁻, is an oxygen atom covalently bonded to a hydrogen atom, with one extra electron. The extra electron gives it a charge of 1−, and its oxygen keeps three lone pairs. Being negative and having lone pairs to spare, OH⁻ strongly attracts H⁺ ions and joins with them to form water:

This is what makes a hydroxide a base: it removes the H⁺ ions that make a solution acidic. The oxide ion, O²⁻, in a metal oxide does the same, taking two H⁺ ions to form one water molecule.

Most bases, such as copper(II) oxide and magnesium hydroxide, are insoluble. Sodium hydroxide, potassium hydroxide and calcium hydroxide dissolve, so they are alkalis. Aqueous ammonia is also an alkali, although it contains no metal: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq). Like acids, alkalis are strong or weak. Sodium hydroxide is fully ionised, so it is a strong alkali. Ammonia is only partly ionised, so it is a weak alkali. Alkalis feel soapy and turn red litmus blue.

5.3.2The two reactions of bases

The pH scale, usually from 0 to 14, shows how acidic or alkaline a solution is (Figure 5.3).

• pH < 7: more H⁺ than OH⁻, so the solution is acidic. The lower the pH, the more acidic.
• pH = 7: equal H⁺ and OH⁻, so the solution is neutral, like pure water.
• pH > 7: more OH⁻ than H⁺, so the solution is alkaline.

Each step of one pH unit is a tenfold change in the concentration of H⁺. A solution of pH 4 has 1000 times more H⁺ than one of pH 7.

pH can be measured with a pH meter or with an indicator, a dye whose colour depends on pH. Universal Indicator is a mixture of dyes that gives a different colour at each pH. Other indicators have one colour change over a narrow range:

Most plants grow best in neutral or slightly acidic soil. Soil that is too acidic is treated with a base such as calcium hydroxide (slaked lime), which neutralises the excess acid.

How an oxide behaves depends on whether its element is a metal or a non-metal (Figure 5.4).

The definitions so far are the Arrhenius theory: an acid dissociates in water to give H⁺(aq), and a base gives OH⁻(aq). It works only in water, and it cannot explain why ammonia, with no OH group, is a base. Two wider theories fix this.

A proton is far too small to exist on its own. When an acid dissolves, it gives its proton to a lone pair on a water molecule, forming the hydronium ion, H₃O⁺ (Figure 5.5). H⁺(aq) and H₃O⁺(aq) mean the same thing. A Brønsted–Lowry base must have a lone pair to accept the proton.

Two species that differ by one H⁺ are a conjugate acid–base pair. When an acid gives away its proton, what is left is its conjugate base. When a base accepts a proton, it becomes its conjugate acid. In HCl + H₂O → H₃O⁺ + Cl⁻, Cl⁻ is the conjugate base of HCl, and H₃O⁺ is the conjugate acid of H₂O.

Water can act as an acid or a base, depending on what it reacts with. An acid that can donate one proton per molecule, such as HCl or HNO₃, is monobasic. H₂SO₄ is dibasic and H₃PO₄ is tribasic.

Ammonia and boron trifluoride show why this happens (Figure 5.6). Boron is in Group 13, so in BF₃ it shares its three valence electrons and ends up with only six electrons round it, two short of an octet. The nitrogen in NH₃ has a lone pair that is not used in bonding. Nitrogen gives this pair to boron, forming a dative bond (Chapter 3), and boron reaches its octet: NH₃ + BF₃ → H₃N→BF₃.

• NH₃ donates the electron pair, so it is the Lewis base.
• BF₃ accepts the electron pair, so it is the Lewis acid.

No proton moves, so this is an acid–base reaction only in the Lewis sense. Every Brønsted–Lowry base is also a Lewis base, because it accepts a proton with a lone pair.

So far, every salt has come from an acid. There is a second way. When two solutions are mixed and the cation from one and the anion from the other form an insoluble salt, the salt appears as a solid in the liquid, a precipitate (Figure 5.7). This is a precipitation reaction:

Na⁺ and NO₃⁻ are spectator ions and stay in solution. The precipitate is separated by filtration (Chapter 1).

To choose how to make a salt, first decide whether it dissolves in water.

There are three methods (Figure 5.8).

• Acid + excess insoluble solid, for a soluble salt: add a metal, insoluble base or insoluble carbonate to the acid until no more reacts. The excess solid makes sure all the acid is used, and it is then filtered off. Heat the filtrate until it is saturated, cool it to crystallise the salt, filter, wash the crystals with a little cold water and dry them. Very reactive metals are too dangerous, and unreactive metals do not react.
• Titration, for a soluble salt from soluble starting materials: sodium, potassium and ammonium salts. Neither reactant can be filtered off, so the exact volumes are found by titration with an indicator (Section 5.8). The reaction is repeated with the same volumes and no indicator, and the salt is crystallised.
• Precipitation, for an insoluble salt: mix a solution containing its cation with a solution containing its anion. Ba(NO₃)₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaNO₃(aq). Filter, wash the precipitate with distilled water and dry it.

A titration measures the volume of one solution that reacts exactly with a fixed volume of another (Figure 5.9). A pipette delivers a fixed volume of one solution into a conical flask with a few drops of indicator. The other solution is added from a burette until the indicator just changes colour.

• A standard solution is one whose concentration is accurately known.
• The equivalence point is reached when the reactants have reacted exactly in the ratio of the equation. The end-point is when the indicator changes colour. A good indicator changes colour at the equivalence point.

Some salts crystallise with water molecules held in their lattice, called water of crystallisation. Copper(II) sulfate crystals are CuSO₄·5H₂O. A titration can find the number of water molecules.

5.8.1Choosing an indicator Core · A Level

At the equivalence point the pH changes very quickly. The indicator must change colour within that sudden change, which depends on whether the acid and base are strong or weak.

5.8.2Back-titration Core · A Level

An insoluble solid, such as calcium carbonate or magnesium hydroxide, cannot be titrated directly. Instead, a known excess of acid is added so that all the solid reacts. The acid left over is then titrated with a standard alkali. The acid that reacted with the solid is found by difference (Figure 5.10).