Gaining and losing oxygen

Metals react with oxygen to make metal oxides. Because the metal gains oxygen, this is an oxidation reaction:

The opposite is reduction — the loss of oxygen. When copper oxide is heated with carbon, the copper oxide is reduced:

In terms of oxygen: OXidation = gain of OXygen; reduction = loss of oxygen. These two always happen together.

Oxidised or reduced?

Ordering metals by reactivity

The reactivity of a metal depends on its tendency to form positive ions. We arrange metals in a reactivity series, deduced from their reactions with water and dilute acids. The non-metals carbon and hydrogen are included for comparison:

Why it matters: a more reactive metal displaces a less reactive metal from its compound — and a metal's position relative to carbon and hydrogen decides how it is extracted and how it reacts with acid.

Will it displace?

A more reactive metal displaces a less reactive one. Tap the correct outcome each time.

Reduction with carbon vs electrolysis

Most metals are found in the Earth as compounds (ores) and must be extracted. How we do it depends on the metal's position in the reactivity series, compared with carbon:

• Less reactive than carbon (e.g. zinc, iron, copper) → can be extracted by reduction with carbon. The metal oxide loses oxygen to the carbon.
• More reactive than carbon (e.g. potassium, sodium, calcium, magnesium, aluminium) → carbon can't do it, so we use electrolysis.

Gold and other very unreactive metals are found as the metal itself (native), so they need no chemical extraction at all.

How would you extract it?

Why does carbon work for iron but not aluminium?

Redox in terms of electrons

At Higher tier, oxidation and reduction are defined by electrons, not just oxygen. Remember OIL RIG:

In a displacement reaction the more reactive metal loses electrons (oxidised); the less reactive metal ion gains electrons (reduced). The half equations show this:

Watch out (OIL RIG): it feels backwards, but losing electrons is oxidation and gaining electrons is reduction. The electrons lost must equal the electrons gained, so they cancel out of the full ionic equation.

Loss or gain of electrons?

Why do only some acid reactions fizz?

Acid + metal → salt + hydrogen

Acids react with some metals (e.g. magnesium, zinc, iron) to make a salt and hydrogen gas:

The salt's name comes from the acid: hydrochloric acid → chloride, sulfuric acid → sulfate, nitric acid → nitrate.

(HT) This is a redox reaction: the metal atoms lose electrons (oxidised) and the hydrogen ions gain electrons (reduced) to form H₂.

What salt forms?

Acid + base → salt + water

Acids are neutralised by bases. The base can be a metal oxide, a metal hydroxide (an alkali if soluble) or a metal carbonate:

Watch out (naming): the salt's second word comes from the acid (chloride / sulfate / nitrate); the first word comes from the metal in the base.

Balance the neutralisation

Making a pure, dry soluble salt

A soluble salt can be made by reacting an acid with an insoluble solid (a metal, metal oxide, hydroxide or carbonate). The classic example is copper sulfate from sulfuric acid + copper oxide:

• Warm the dilute acid (Bunsen burner), then add the insoluble solid a little at a time.
• Keep adding until no more reacts (excess solid sits unreacted) — this ensures all the acid is used up.
• Filter off the excess solid to leave a salt solution.
• Crystallise: evaporate some water (water bath/electric heater), then leave to cool so crystals form. Pat dry.

Why excess solid? Using more base than the acid can dissolve guarantees the acid is fully neutralised — and the leftover solid is easy to filter out.

Getting a pure salt

Acids, alkalis and pH

Acids produce hydrogen ions, H⁺ in water; alkalis contain hydroxide ions, OH⁻. The pH scale (0–14) measures acidity, read with universal indicator or a pH probe:

In neutralisation between an acid and an alkali, H⁺ and OH⁻ combine to make water:

Watch out (pH): pH 7 is neutral. A lower pH means more acidic (not less!). pH above 7 is alkaline.

Reading the pH scale

Why do only some acid reactions fizz?

Titration: finding reacting volumes

A titration measures exactly what volume of acid reacts with a known volume of alkali. An indicator shows the end-point. Use strong acids (hydrochloric, sulfuric, nitric) and strong alkalis.

• Use a pipette to put 25.0 cm³ of alkali into a conical flask. Add a few drops of indicator.
• Fill a burette with acid and read the start volume.
• Add the acid while swirling, then drop by drop near the end, until the indicator just changes colour. Read the final volume.
• Titre = final reading − start reading. Repeat until titres agree within 0.10 cm³ (concordant), then take their mean.

Higher tier: from the reacting volumes and one known concentration you can work out the other concentration in mol/dm³ and g/dm³ (see Quantitative chemistry).

Strong vs weak acids

This is about how much an acid ionises in water — not the same as concentration:

• A strong acid is completely ionised in solution (e.g. hydrochloric, nitric, sulfuric).
• A weak acid is only partially ionised (e.g. ethanoic, citric, carbonic).

As the pH falls by 1 unit, the H⁺ concentration rises by a factor of 10. So pH 2 has ten times the H⁺ of pH 3.

Don't confuse two pairs: dilute/concentrated = amount of acid in a volume; weak/strong = degree of ionisation. A weak acid can still be concentrated.

Strong, weak, dilute?

Why does water change the product?

Splitting compounds with electricity

When an ionic compound is melted or dissolved, its ions are free to move — it becomes an electrolyte. Passing a current makes the ions move to the electrodes, where they are discharged as elements:

Watch out (cathode): at the cathode, positive ions gain electrons (reduction). "Cat-ions go to the cat-hode." At the anode, negative ions lose electrons (oxidation).

Rules for the products

With aqueous solutions, water also provides H⁺ and OH⁻ ions, so there is competition. The rules (inert electrodes):

• Cathode (–): hydrogen is produced unless the metal is less reactive than hydrogen (e.g. copper) — then the metal is deposited.
• Anode (+): oxygen is produced unless a halide (Cl⁻, Br⁻, I⁻) is present — then the halogen forms.

Reactive metals like aluminium are extracted by electrolysis of the molten oxide:

The carbon anode burns away: oxygen made at the positive electrode reacts with the hot carbon to form CO₂, so the anode must be replaced regularly.

Cathode or anode?

Electrolysis of aqueous copper chloride and sodium sulfate. Tap a product, then the electrode where it forms.

Why does water change the product?

Half equations at the electrodes

At Higher tier you write half equations for each electrode. At the cathode ions gain electrons (reduction); at the anode ions lose electrons (oxidation):

For molten lead bromide the half equations are:

Check the balance: atoms and charge must balance. In 2H⁺ + 2e⁻ → H₂ the 2 positive charges are cancelled by 2 electrons, giving neutral H₂.

Complete the half equation

Electrolysis of copper sulfate

In required practical 3 you electrolyse aqueous solutions with inert (carbon) electrodes, eg aqueous copper(II) sulfate, and observe what forms:

• Cathode (–): a coating of copper appears — copper is less reactive than hydrogen, so the metal is deposited.
• Anode (+): bubbles of oxygen are released (no halide is present).

Investigation idea: changing the current or concentration changes the mass of copper deposited or the volume of gas collected — a good hypothesis to test.

Products at the electrodes

The key ideas to know

You've covered AQA 4.4 — reactivity of metals, reactions of acids (including titrations, Chemistry only), and electrolysis. Press Finish to see your score.