GCSE Chemistry Acids & salts revision

Solutions run from strongly acidic (pH 0 — battery acid) through neutral (pH 7 — pure water) to strongly alkaline (pH 14 — drain cleaner). What the number actually reports is the concentration of hydrogen ions, H⁺(aq): acids dump H⁺ into water, and the more they release, the lower the pH. Alkaline solutions are rich in hydroxide ions, OH⁻(aq), and high on the scale. Universal indicator is a dye mixture that steps through red → orange → yellow → green → blue → purple as pH climbs; a pH probe reads the number directly and is more precise.

A strong acid ionises completely in water: every HCl molecule becomes H⁺ + Cl⁻. A weak acid — ethanoic acid in vinegar, citric acid in lemon juice — ionises only partially: most molecules stay whole, with just a fraction releasing H⁺. Strength is about the degree of ionisation, a fixed property of the acid. Concentration is about how much acid is dissolved per volume — you can have a dilute strong acid or a concentrated weak acid, and the two labels are independent.

The consequences show up in reactions. At equal concentration, a weak acid has a lower H⁺ concentration, so it reacts more slowly — magnesium fizzes gently in vinegar but furiously in hydrochloric acid of the same molarity — and shows a higher pH. But given enough time, a weak acid can still dissolve the metal completely; there is plenty of un-ionised acid waiting to release more H⁺ as it is used up.

The naming logic: the acid fixes the negative part of the salt — hydrochloric acid makes chlorides, sulfuric acid makes sulfates, nitric acid makes nitrates — while the other reagent supplies the metal ion. Copper oxide + sulfuric acid gives copper sulfate; sodium hydroxide + nitric acid gives sodium nitrate; zinc carbonate + hydrochloric acid gives zinc chloride. With a reactive metal, the hydrogen pops out as a gas that gives a squeaky pop with a lit splint.

At the ionic level (Higher), every neutralisation by a hydroxide is the same half-reaction: H⁺(aq) + OH⁻(aq) → H₂O(l). The spectator ions — Na⁺ and Cl⁻, say — simply watch. Carbonates bring their own drama: CO₃²⁻ + 2H⁺ → H₂O + CO₂, the fizzing that identifies limestone in an instant.

1. 11. Warm the dilute acid

   Gently heat about 25 cm³ of dilute sulfuric acid in a beaker — warming speeds the reaction, but never boil the acid.

2. 22. Add the insoluble base in excess

   Stir in black copper(II) oxide a spatula at a time. It dissolves as it reacts, and the blue colour deepens. Keep adding until some black solid stays undissolved — that proves the acid has all reacted.

3. 33. Filter

   Filter the warm mixture: the filtrate is blue copper sulfate solution; the residue is the excess copper oxide. Filtration removes the impurity you deliberately created.

4. 44. Evaporate to crystallisation point

   Heat the filtrate gently until about a third of the water has gone and crystals just begin to form at the edges.

5. 55. Cool and dry

   Leave the concentrated solution to cool slowly so blue copper sulfate crystals grow. Pat them dry between filter papers.

Acidic soil locks up nutrients, so farmers spread calcium hydroxide (slaked lime) — an alkali — to raise the pH toward neutral. Too much makes the soil too alkaline, so the dose matters. Industrial waste acids are neutralised before release; indigestion tablets use mild bases such as calcium carbonate or magnesium hydroxide to mop up excess stomach HCl (that gentle carbonate fizz). Each application is the single reaction H⁺ + OH⁻ → H₂O dressed in different clothes.