GCSE Chemistry₈₄₆₂

Metals react with oxygen to produce metal oxides. The reactions are oxidation reactions because the metals gain oxygen.

Students should be able to explain reduction and oxidation in terms of loss or gain of oxygen.

When metals react with other substances the metal atoms form positive ions. The reactivity of a metal is related to its tendency to form positive ions. Metals can be arranged in order of their reactivity in a reactivity series. The metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids.

The non-metals hydrogen and carbon are often included in the reactivity series.

A more reactive metal can displace a less reactive metal from a compound.

Students should be able to:

• recall and describe the reactions, if any, of potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper with water or dilute acids and where appropriate, to place these metals in order of reactivity
• explain how the reactivity of metals with water or dilute acids is related to the tendency of the metal to form its positive ion
• deduce an order of reactivity of metals based on experimental results.

The reactions of metals with water and acids are limited to room temperature and do not include reactions with steam.

AT 6

Mixing of reagents to explore chemical changes and/or products.

Unreactive metals such as gold are found in the Earth as the metal itself but most metals are found as compounds that require chemical reactions to extract the metal.

Metals less reactive than carbon can be extracted from their oxides by reduction with carbon.

Reduction involves the loss of oxygen.

Knowledge and understanding are limited to the reduction of oxides using carbon.

Knowledge of the details of processes used in the extraction of metals is not required.

Students should be able to:

• interpret or evaluate specific metal extraction processes when given appropriate information
• identify the substances which are oxidised or reduced in terms of gain or loss of oxygen.

Oxidation is the loss of electrons and reduction is the gain of electrons.

• write ionic equations for displacement reactions

• identify in a given reaction, symbol equation or half equation which species are oxidised and which are reduced.

Acids react with some metals to produce salts and hydrogen.

(HT only) Students should be able to:

• explain in terms of gain or loss of electrons, that these are redox reactions
• identify which species are oxidised and which are reduced in given chemical equations.
  Knowledge of reactions limited to those of magnesium, zinc and iron with hydrochloric and sulfuric acids.

Acids are neutralised by alkalis (eg soluble metal hydroxides) and bases (eg insoluble metal hydroxides and metal oxides) to produce salts and water, and by metal carbonates to produce salts, water and carbon dioxide.

The particular salt produced in any reaction between an acid and a base or alkali depends on:

• the acid used (hydrochloric acid produces chlorides, nitric acid produces nitrates, sulfuric acid produces sulfates)
• the positive ions in the base, alkali or carbonate.

Students should be able to:

• predict products from given reactants
• use the formulae of common ions to deduce the formulae of salts.

Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates. The solid is added to the acid until no more reacts and the excess solid is filtered off to produce a solution of the salt.

Salt solutions can be crystallised to produce solid salts.

Students should be able to describe how to make pure, dry samples of named soluble salts from information provided.

Acids produce hydrogen ions (H⁺ ) in aqueous solutions.

Aqueous solutions of alkalis contain hydroxide ions (OH^– ).

The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.

A solution with pH 7 is neutral. Aqueous solutions of acids have pH values of less than 7 and aqueous solutions of alkalis have pH values greater than 7.

In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water.

This reaction can be represented by the equation:

Students should be able to:

• describe the use of universal indicator or a wide range indicator to measure the approximate pH of a solution
• use the pH scale to identify acidic or alkaline solutions.

AT 3

This is an opportunity to investigate pH changes when a strong acid neutralises a strong alkali.

The volumes of acid and alkali solutions that react with each other can be measured by titration using a suitable indicator.

Students should be able to:

• describe how to carry out titrations using strong acids and strong alkalis only (sulfuric, hydrochloric and nitric acids only) to find the reacting volumes accurately
• (HT Only) calculate the chemical quantities in titrations involving concentrations in mol/dm³ and in g/dm³ .

A strong acid is completely ionised in aqueous solution. Examples of strong acids are hydrochloric, nitric and sulfuric acids.

A weak acid is only partially ionised in aqueous solution. Examples of weak acids are ethanoic, citric and carbonic acids.

For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.

As the pH decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of 10.

Students should be able to:

• use and explain the terms dilute and concentrated (in terms of amount of substance), and weak and strong (in terms of the degree of ionisation) in relation to acids

AT 8

An opportunity to measure the pH of different acids at different concentrations.

• describe neutrality and relative acidity in terms of the effect on hydrogen ion concentration and the numerical value of pH (whole numbers only).

MS 2h

Make order of magnitude calculations.

When an ionic compound is melted or dissolved in water, the ions are free to move about within the liquid or solution. These liquids and solutions are able to conduct electricity and are called electrolytes.

Passing an electric current through electrolytes causes the ions to move to the electrodes. Positively charged ions move to the negative electrode (the cathode), and negatively charged ions move to the positive electrode (the anode). Ions are discharged at the electrodes producing elements. This process is called electrolysis.

(HT only) Throughout Section 4.4.3 Higher Tier students should be able to write half equations for the reactions occurring at the electrodes during electrolysis, and may be required to complete and balance supplied half equations.

When a simple ionic compound (eg lead bromide) is electrolysed in the molten state using inert electrodes, the metal (lead) is produced at the cathode and the non-metal (bromine) is produced at the anode.

Students should be able to predict the products of the electrolysis of binary ionic compounds in the molten state.

A safer alternative for practical work is anhydrous zinc chloride.

Metals can be extracted from molten compounds using electrolysis. Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon. Large amounts of energy are used in the extraction process to melt the compounds and to produce the electrical current.

Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide and cryolite using carbon as the positive electrode (anode).

• explain why a mixture is used as the electrolyte
• explain why the positive electrode must be continually replaced.

The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved.

At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen.

At the positive electrode (anode), oxygen is produced unless the solution contains halide ions when the halogen is produced.

This happens because in the aqueous solution water molecules break down producing hydrogen ions and hydroxide ions that are discharged.

During electrolysis, at the cathode (negative electrode), positively charged ions gain electrons and so the reactions are reductions.

At the anode (positive electrode), negatively charged ions lose electrons and so the reactions are oxidations.

Reactions at electrodes can be represented by half equations, for example:

2H⁺ + 2e^- → H₂

and

4OH^- → O₂ + 2H₂ O + 4e^-

or

4OH^- – 4e^- → O₂ + 2H₂ O