Acids, Bases & Salts

Acids & bases

Acids taste sour, feel corrosive and turn litmus red. Bases feel soapy, neutralise acids and turn litmus blue. Behind those everyday properties is a single trade: an acid releases H⁺; a base absorbs it.

Properties of acids

The salt name comes from two places: the first word from the metal or base, the second from the acid. Hydrochloric acid gives chlorides, sulfuric acid gives sulfates, nitric acid gives nitrates — so sodium hydroxide plus hydrochloric acid gives sodium chloride.

Properties of bases & alkalis

Bases are usually metal oxides or hydroxides. Those that are water-soluble — NaOH, KOH, Ca(OH)₂ and aqueous ammonia — are called alkalis. Bases neutralise acids to give salt and water; alkalis do the same from solution. A distinctive reaction of alkalis: they release ammonia gas from ammonium salts on warming, a useful laboratory route to NH₃.

ExtendedThe proton view

A deeper definition sees acids as proton (H⁺) donors and bases as proton acceptors. In water, HCl donates its H⁺ to a water molecule; NaOH releases OH⁻ which then accepts H⁺ from the acid. Neutralisation reduces to H⁺ + OH⁻ → H₂O. Strong acids (HCl, HNO₃, H₂SO₄) ionise fully in water — every molecule releases its H⁺. Weak acids (ethanoic acid, carbonic acid) ionise only partially: at any moment only a small fraction of molecules have released their proton.

The pH scale & indicators

The pH scale converts “how acidic?” into a single number between 0 and 14. The further from 7, the stronger the acid or alkali.

Reading the scale

pH 0–6 is acidic: the lower the number, the stronger the acid. Pure water is neutral at pH 7. pH 8–14 is alkaline: the higher the number, the stronger the alkali. Stomach acid sits at pH 1–2, vinegar at pH 3, blood at 7.4, soap around pH 10, bleach at pH 13.

Under the proton view, low pH means a high concentration of H⁺ ions and a low concentration of OH⁻; high pH means the reverse. At pH 7 exactly, the two concentrations are equal.

Indicators

Universal indicator is a blend that gives a continuous rainbow across the whole scale — ideal for a rough reading against a colour chart, but too gradual for precise work. Litmus is convenient as a test paper but its change is smeared across pH 5–8. Thymolphthalein and methyl orange each switch sharply at their own end-point, making them the standard choices when the exact moment of neutralisation matters.

ExtendedUniversal indicator colours

The progression is worth remembering: strong acid (pH 0–2) is red; weak acid (pH 3–6) fades through orange to yellow; neutral (pH 7) is green; weak alkali (pH 8–11) shifts through blue; strong alkali (pH 12–14) is violet.

Classifying oxides

Oxides fall into four categories based on what they react with. The pattern: metal oxides tend to be basic, non-metal oxides tend to be acidic, and a small group sits on the fence between the two.

Acidic and basic oxides

Most non-metal oxides are acidic. CO₂ dissolves in water to give a weakly acidic solution (carbonic acid); SO₂ and NO₂ in the atmosphere dissolve in rain water to produce acid rain. All acidic oxides react with alkalis to form a salt and water. Most metal oxides are basic and react with acids to give a salt and water, e.g. MgO + 2HCl → MgCl₂ + H₂O; the soluble ones (Na₂O, K₂O) dissolve in water to give alkaline solutions.

ExtendedAmphoteric oxides

A handful of oxides react with both acids and alkalis. Aluminium oxide is the standard example: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O, and Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄. Zinc oxide and lead(II) oxide behave similarly — forming salts with acids and complex salts with alkalis.

Preparing soluble salts

Making a salt in the lab uses one of two general methods, chosen by the state of the base. If the base is insoluble, add it as a solid in excess and filter; if the base is an alkali, use a titration.

Method A — excess solid

The base (or metal, or insoluble carbonate) is added to the acid in excess so all the acid is used up and the surplus solid can be filtered off; the salt solution left behind is then evaporated and crystallised. The core sequence: (1) warm the dilute acid gently; (2) add the solid a little at a time, stirring, until no more reacts; (3) filter to remove the excess solid and keep the filtrate; (4) evaporate until crystals begin to form on a cold glass rod; (5) leave in a warm place to crystallise, decant excess liquid, dry with filter paper.

Method B — titration

When the base is an alkali it is already in solution, so there is nothing to filter. Instead, acid and alkali are combined in exactly the right proportions: (1) fill a burette with the acid, pipette a known volume of alkali into a flask with a few drops of indicator; (2) add acid until the indicator changes colour and record the volume; (3) repeat without indicator, adding the same volume of acid to the same volume of alkali; (4) evaporate and crystallise the salt solution. The indicator is only used to find the neutralisation volume, so no dye contaminates the salt.

Insoluble salts & solubility rules

If the salt you want is insoluble, none of the soluble-salt methods will work — the product would drop out as sludge before you could filter cleanly. Instead, use precipitation: two soluble salt solutions that swap partners.

Precipitation

Mix two solutions that between them contain the two ions of the required insoluble salt — the insoluble product forms as a fine solid which can then be filtered, washed and dried, while everything else stays dissolved and is washed away. To make barium sulfate: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq). The BaSO₄ precipitates as a white solid; the NaCl stays in solution.

ExtendedSolubility rules

ExtendedIonic equations

The full equation includes spectator ions (Na⁺ and Cl⁻ above) which don’t take part. Stripping them out gives the net ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

Identifying ions & gases

Identifying an unknown cation uses at most two tests: a flame test if the metal is Group 1 or 2, or a precipitate test with NaOH (and if necessary, ammonia). Anions are named by what precipitates or what gas is released, and five simple gas tests round out the toolkit.

Flame tests

Test with aqueous NaOH, then ammonia

The ammonium ion, NH₄⁺, is identified separately: warm the solution with NaOH and ammonia gas is released, turning damp red litmus blue.

Anion tests

Extended adds the sulfite ion, SO₃²⁻: adding dilute acid releases SO₂, which turns acidified KMnO₄ from purple to colourless.

Gas tests