C4 Chemical changes

Metals react with oxygen to form metal oxides, such as when magnesium burns to form white magnesium oxide. These are oxidation reactions because the metal gains oxygen. Reduction is the loss of oxygen.

When metals react, they form positive ions. The more easily a metal forms its positive ion, the more reactive it is.

The reactivity series of metals, with carbon and hydrogen included for comparison. A metal higher in the series forms positive ions more easily, so it reacts more vigorously and can displace any metal below it.

You can deduce the order of reactivity by observing reactions.

With water: Potassium, sodium and lithium react violently with cold water. Calcium reacts steadily; magnesium reacts very slowly. Zinc, iron and copper do not react.

With dilute acid: Magnesium fizzes fast, zinc and iron bubble more slowly, and copper has no reaction.

Displacement: A more reactive metal displaces a less reactive metal from its compound.

A student leaves an iron nail in silver nitrate solution. Will a reaction happen, and what forms?

Iron is above silver in the reactivity series, so iron is more reactive.

Iron displaces silver from its compound: iron + silver nitrate → iron(II) nitrate + silver.

Crystals of silver form on the nail. If a silver wire were put in iron(II) nitrate solution instead, nothing would happen.

More reactive metals are higher in the series. Only a more reactive metal can displace a less reactive one.

Carbon sits between magnesium and zinc; hydrogen sits between iron and copper.

Reactivity depends on the metal's ability to form positive ions, not on the size of the piece used.

Most metals are found as compounds (such as metal oxides) in rocks called ores. A chemical reaction is needed to extract the metal.

Reduction is the loss of oxygen. Oxidation is the gain of oxygen. During extraction, the metal oxide is reduced (loses oxygen) to form the metal.

The method used to extract a metal depends on its position in the reactivity series.

Very unreactive metals (like gold) are found as the uncombined element itself, so they do not need extracting from an ore.

Metals below carbon in the reactivity series (such as zinc, iron, lead and copper) are extracted by heating their oxides with carbon.

Metals above carbon (such as magnesium and calcium) cannot be extracted by heating with carbon because carbon is less reactive than them. Electrolysis is used instead.

How a metal is extracted depends on its position in the reactivity series. Metals above carbon are extracted by electrolysis, metals below carbon by heating their oxides with carbon, and very unreactive metals such as gold are found as the element itself.

Tin is extracted by heating tin oxide, SnO₂, with carbon. Write the balanced equation and say which substance is oxidised and which is reduced.

The balanced equation is: SnO₂ + C → Sn + CO₂ (1 tin, 2 oxygen and 1 carbon atom on each side).

Carbon gains oxygen to form carbon dioxide, so carbon is oxidised. Tin oxide loses oxygen, so it is reduced.

A common mistake is thinking the metal is oxidised during extraction. The metal oxide loses oxygen, so it is reduced.

Remember that reduction means losing oxygen, not adding something.

Metals above carbon in the reactivity series hold onto their oxygen too strongly for carbon to remove it.

When a reactive metal reacts with a dilute acid, the products are a salt and hydrogen gas.

Metal + acid → salt + hydrogen

Magnesium, zinc and iron react with dilute acids. Copper is less reactive than hydrogen, so it does not react with dilute acids.

The hydrogen gas produced in the reaction can be collected in a test tube. A lit splint is held at the mouth of the tube, and a squeaky pop is heard as the hydrogen burns.

The name of the salt depends on the metal and the acid used. Hydrochloric acid produces chlorides, and sulfuric acid produces sulfates.

Magnesium + hydrochloric acid → magnesium chloride + hydrogen

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

Zinc + sulfuric acid → zinc sulfate + hydrogen

Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

Iron reacts with dilute hydrochloric acid to form iron(II) chloride and hydrogen gas. Write the balanced symbol equation for this reaction.

Fe + 2HCl → FeCl₂ + H₂

Remember that hydrogen exists as diatomic molecules (H₂) in equations, not single atoms.

Iron forms iron(II) salts when it reacts with dilute acids at GCSE.

Look for fizzing and the metal getting smaller as signs that the reaction is happening.

Do not confuse the test for hydrogen with the test for oxygen, which relights a glowing splint.

An acid is neutralised by a base to produce a salt and water. Alkalis are bases that are soluble in water (like sodium hydroxide). Insoluble bases include metal oxides and metal hydroxides.

If a metal carbonate is used, carbon dioxide gas is also produced, which causes fizzing.

To name a salt, combine the metal part from the base with the acid part: hydrochloric acid makes chlorides, nitric acid makes nitrates, and sulfuric acid makes sulfates.

To find the chemical formula of a salt, balance the charges of the ions so the total positive charge equals the total negative charge. If you need more than one of a group ion (like SO₄²⁻), put brackets around it.

To make a soluble salt from an insoluble solid, add the solid to the acid until no more reacts (an excess). This ensures all the acid is used up. Filter off the excess solid, gently heat the solution to evaporate some water, and leave it to crystallise.

What is the balanced equation for the reaction between potassium carbonate and nitric acid?

Potassium carbonate + nitric acid → potassium nitrate + water + carbon dioxide.

Potassium ions are K⁺ and nitrate ions are NO₃⁻, so potassium nitrate is KNO₃.

Two potassium ions come from K₂CO₃, so 2KNO₃ and 2HNO₃ are needed:

K₂CO₃ + 2HNO₃ → 2KNO₃ + H₂O + CO₂

Alkalis are just soluble bases. Both neutralise acids to make a salt and water.

Never write 'hydrogen' as a product of neutralisation; it always produces water.

When making a soluble salt from an insoluble solid, you must add an excess of the solid to ensure all the acid is used up, then filter it.

Acids release hydrogen ions, H⁺(aq), when dissolved in water. Alkalis are bases that dissolve in water, and they contain hydroxide ions, OH⁻(aq).

The pH scale goes from 0 to 14 and shows how acidic or alkaline a solution is. A pH below 7 is acidic, exactly 7 is neutral, and above 7 is alkaline. The further the value is from 7, the more strongly acidic or alkaline the solution is.

The pH scale with typical universal indicator colours. Acids are below pH 7, neutral solutions are pH 7 and alkalis are above pH 7.

A student neutralises a strong acid by adding a strong alkali. How can they use universal indicator to see when the solution is neutralised, and what ionic equation represents this reaction?

They should add a few drops of universal indicator to the acid. It will start red.

As alkali is added, the colour will change through orange and yellow. When it turns green, the solution is neutral (pH 7).

The ionic equation for neutralisation is: H⁺(aq) + OH⁻(aq) → H₂O(l).

A pH of 7 means neutral, but it does not mean the liquid has to be pure water (it could be a neutral salt solution).

Universal indicator gives an approximate whole-number pH based on a colour chart. A pH probe gives a more precise reading as it does not rely on judging colours.

When alkali is added to acid, the pH rises slowly at first, then jumps sharply near neutralisation, before levelling off at a high value.

Electrolysis is the use of a direct electrical current to break down an ionic compound into its elements.

An electrolyte is a liquid or solution that contains ions that are free to move. This includes molten ionic compounds and dissolved ionic compounds in water.

Solid ionic compounds do not conduct electricity because their ions are locked in a giant lattice and cannot move.

In electrolysis, positive ions move to the negative electrode (cathode) and negative ions move to the positive electrode (anode), where they are discharged.

What happens during the electrolysis of molten potassium chloride?

Potassium chloride contains positive potassium ions (K⁺) and negative chloride ions (Cl⁻). When melted, these ions are free to move.

The positive potassium ions move to the negative electrode (the cathode), where they gain electrons to become potassium metal.

The negative chloride ions move to the positive electrode (the anode), where they lose electrons to become chlorine gas.

Word equation: potassium chloride → potassium + chlorine

Symbol equation: 2KCl → 2K + Cl₂

Remember that electrons do not flow through the electrolyte. Electrons flow through the external wires, while ions move through the liquid.

Positive ions (metal ions and hydrogen ions) are attracted to the cathode.

Negative ions (non-metal ions) are attracted to the anode.

Inert electrodes (such as graphite or platinum) are used so that they do not react with the electrolyte or the products.

Electrolysis passes an electric current through a molten ionic compound or a solution to break it down. An ionic compound must be molten (melted) so its ions are free to move and carry charge.

For a molten binary ionic compound, the positive metal ions move to the negative electrode (cathode) and form the metal. The negative non-metal ions move to the positive electrode (anode) and form the non-metal.

Metals that are more reactive than carbon cannot be extracted from their ores by reduction with carbon. They must be extracted using electrolysis.

This is an expensive process because a lot of energy is needed to melt the compounds and to supply the electric current.

Aluminium oxide is dissolved in molten cryolite to lower the melting point. Molten aluminium forms at the carbon cathode lining and collects at the bottom of the cell. Oxygen forms at the carbon anodes and reacts with them to make carbon dioxide, so the anodes wear away.

What forms at each electrode when molten calcium bromide is electrolysed?

Calcium bromide contains positive calcium ions (Ca²⁺) and negative bromide ions (Br⁻).

At the negative cathode, calcium ions gain electrons and form calcium metal.

At the positive anode, bromide ions lose electrons and form bromine.

Cryolite is not a catalyst and it is not the source of aluminium. It is only mixed with the aluminium oxide to lower the melting point of the mixture, saving energy.

In aluminium extraction, oxygen forms at the positive carbon anodes. The oxygen reacts with the hot carbon to make carbon dioxide gas. This means the anodes gradually burn away and must be replaced regularly.

Remember that 'molten' means melted into a liquid by heat, not dissolved in water.

In an aqueous solution, as well as the ions from the dissolved ionic compound, there are also hydrogen ions (H⁺) and hydroxide ions (OH⁻) from the water splitting slightly.

Which ions are discharged at the electrodes depends on the reactivity of the elements.

At the cathode (−): Hydrogen gas is produced if the metal in the compound is more reactive than hydrogen (e.g. sodium, potassium, magnesium). The metal itself is produced only if it is less reactive than hydrogen (e.g. copper, silver).

At the anode (+): Oxygen gas is produced from the hydroxide ions unless the solution contains halide ions (chloride, bromide, iodide). If halide ions are present, the halogen (chlorine, bromine, iodine) is produced instead. Sulfate and nitrate ions are not discharged.

Deciding the products of electrolysing a solution with inert electrodes. At the cathode, hydrogen forms unless the metal is less reactive than hydrogen. At the anode, oxygen forms unless halide ions are present.

Predict the products at the cathode and anode when potassium iodide solution is electrolysed with inert electrodes.

At the cathode: potassium is more reactive than hydrogen, so hydrogen gas is produced.

At the anode: iodide ions are halide ions, so iodine is produced.

Remember that the products from an aqueous solution can be different from a molten compound. Molten sodium chloride produces sodium at the cathode, but aqueous sodium chloride produces hydrogen.

Sulfate and nitrate ions do not produce sulfur or nitrogen; they lead to oxygen being produced.

Hydrogen forms at the cathode, never the anode.

In chemical reactions, oxidation and reduction can be defined in terms of electrons. Oxidation is the loss of electrons and reduction is the gain of electrons. You can remember this using the acronym OIL RIG (Oxidation Is Loss, Reduction Is Gain).

A redox reaction is one where both oxidation and reduction happen at the same time.

You can split a redox reaction into two half equations to see where electrons are lost and gained. For example, when a metal reacts, it loses electrons to form positive ions (it is oxidised).

When you write an ionic equation for a reaction like displacement or a metal reacting with an acid, you leave out the spectator ions. These are the ions that do not change during the reaction, such as sulfate, chloride or nitrate ions.

An iron nail is placed in copper(II) sulfate solution, and copper forms on the nail. Write the balanced ionic equation and identify the species oxidised and reduced.

Word equation: iron + copper(II) sulfate → iron(II) sulfate + copper.

The sulfate ions (SO₄²⁻) do not change, so they are spectator ions and are left out.

Ionic equation: Fe(s) + Cu²⁺(aq) → Fe²⁺(aq) + Cu(s). Atoms balance, and the charge is 2+ on both sides.

Iron atoms lose electrons (Fe → Fe²⁺ + 2e⁻), so iron is oxidised. Copper(II) ions gain electrons (Cu²⁺ + 2e⁻ → Cu), so they are reduced.

When balancing ionic equations, check that the total charge on the left is the same as the total charge on the right. For example, Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s) balances because there is a total 2+ charge on both sides.

Make sure electrons are on the correct side in half equations: on the right for oxidation (loss) and on the left for reduction (gain).

A strong acid (like hydrochloric, nitric or sulfuric acid) ionises completely in aqueous solution. Every dissolved molecule releases a hydrogen ion. For example: HCl(aq) → H⁺(aq) + Cl⁻(aq), which is hydrochloric acid forming hydrogen and chloride ions.

A weak acid (like ethanoic, citric or carbonic acid) only partially ionises. Most molecules stay un-ionised. This is a reversible reaction: CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq), which is ethanoic acid forming ethanoate and hydrogen ions.

Because they release more H⁺ ions, strong acids have a lower pH and react faster than weak acids at the exact same concentration.

Two acid solutions of the same concentration. In the strong acid every molecule has ionised into H⁺ and negative ions; in the weak acid most molecules stay un-ionised, so there are fewer H⁺ ions.

The pH scale is logarithmic. Each decrease of 1 in pH means the concentration of hydrogen ions becomes 10 times greater.

A solution of hydrochloric acid has a pH of 1. It is diluted until its volume is 1000 times larger. What is the new pH?

Diluting by a factor of 1000 makes the hydrogen ion concentration 10³ times smaller.

Each factor of 10 increases the pH by 1, so the pH increases by 3.

New pH = 1 + 3 = 4.

Do not confuse strong/weak with concentrated/dilute. Strong/weak is the fraction of molecules that ionise. Concentrated/dilute is the mass of acid dissolved in a given volume.

You can have a dilute strong acid (a small amount of acid in a large volume of water) or a concentrated weak acid.

When calculating pH changes, remember that diluting an acid makes it less acidic, so its pH increases.

A half equation shows what happens to one type of ion at an electrode during electrolysis. It includes electrons (e⁻) to show that the ion has gained or lost them.

At the cathode (negative electrode), positive ions gain electrons to become neutral atoms. This is reduction. The electrons are written on the left.

At the anode (positive electrode), negative ions lose electrons to become neutral atoms or molecules. This is oxidation. The electrons are written on the right.

A half equation must balance in atoms and in charge. Remember that halogens, oxygen and hydrogen form diatomic molecules (e.g. Cl₂, O₂, H₂).

Write the half equations for the electrolysis of molten zinc chloride.

At the cathode (reduction): Zn²⁺ + 2e⁻ → Zn

At the anode (oxidation): 2Cl⁻ → Cl₂ + 2e⁻

Write the half equations for the electrolysis of aqueous potassium bromide.

Potassium is more reactive than hydrogen, so at the cathode: 2H⁺ + 2e⁻ → H₂

Bromide ions are halide ions, so at the anode: 2Br⁻ → Br₂ + 2e⁻

OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons.

Electrons lost can also be shown as subtracted on the left, e.g. 4OH⁻ − 4e⁻ → O₂ + 2H₂O. Both forms are accepted.

If aqueous solutions contain no halide ions, hydroxide ions (OH⁻) react at the anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻.

Make pure, dry crystals of a soluble salt from an insoluble metal oxide or carbonate and a dilute acid. Example: copper(II) oxide + sulfuric acid → copper(II) sulfate + water.

CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)

1. Warm about 40 cm³ of dilute sulfuric acid gently in a beaker on a tripod and gauze.

2. Add the black copper(II) oxide a spatula at a time, stirring, until some stays unreacted.

3. Filter to remove the excess solid.

4. Heat the filtrate in an evaporating basin over a water bath until crystals start to form.

5. Leave to cool and crystallise, then pat the crystals dry with filter paper.

Warming speeds up the reaction. Excess solid makes sure all the acid is used up; with a carbonate, the fizzing stops.

The unreacted solid is the residue; the salt solution is the filtrate.

A water bath heats gently so the crystals do not spit or decompose, and stopping before dryness gives larger, purer crystals.

Wear eye protection. Dilute sulfuric acid is an irritant; copper(II) oxide and copper(II) sulfate are harmful. Do not boil the acid, as it can spit.

1. Black copper(II) oxide is added to warm dilute sulfuric acid until some is left over. 2. The excess solid is filtered off; the blue filtrate is copper(II) sulfate solution. 3. The filtrate is heated gently over a water bath until crystals start to form.

The excess solid is added so that no acid is left; it is then filtered off because it is unreacted solid.

The salt is dissolved in the filtrate, not trapped in the filter paper.

Do not evaporate to dryness with a strong flame: the crystals spit and may break down.

Investigate what forms at each electrode when different aqueous solutions are electrolysed with inert electrodes.

A typical hypothesis: 'A metal forms at the cathode only if it is less reactive than hydrogen'.

1. Pour the solution into a beaker.

2. Place two inert graphite electrodes into the solution, ensuring they do not touch.

3. Fill two small test tubes with solution and invert one over each electrode to collect any gas.

4. Connect the electrodes to a low-voltage d.c. power pack using leads and crocodile clips.

5. Turn on the power and observe the electrodes. Test any gases collected.

Variables: Independent = the solution. Dependent = product at each electrode. Control = solution volume and concentration, voltage, electrode material, distance between electrodes.

Gas tests: Hydrogen burns with a squeaky pop. Oxygen relights a glowing splint. Chlorine turns damp blue litmus paper red then bleaches it white.

Hazards: Chlorine is toxic; run on a small scale in a well-ventilated room. Copper(II) sulfate is harmful; wear eye protection. Switch off before handling electrodes.

Each graphite electrode sits inside an inverted test tube filled with the solution, so any gas formed collects at the top of the tube. The electrode joined to the negative terminal (−) is the cathode; the one joined to the positive terminal (+) is the anode.

In aqueous solutions, hydrogen forms at the cathode unless the metal is less reactive than hydrogen (e.g. copper, silver). Oxygen forms at the anode unless a halide ion is present.

Electrolysing copper(II) sulfate gives a pink-brown copper coating at the cathode and oxygen at the anode.

Hydrogen is tested with a lit splint, not a glowing one. Chlorine bleaches litmus paper; it does not turn it blue.