Chemical reactions

Physical and chemical changes

A physical change alters the form of a substance but not its chemical identity. Ice melting, water boiling, salt dissolving and glass shattering are all physical changes: no new substance is made, the particles are simply rearranged or spread further apart, and the change is usually easy to reverse.

A chemical change produces one or more new substances with different properties from the starting materials. Burning magnesium, rusting iron and neutralising an acid with an alkali are chemical changes. The atoms are rearranged into new combinations, which is why the new substance has a different formula from the old one.

You can tell the two apart from the evidence. A chemical change usually gives one or more of: a colour change, a gas given off, a solid precipitate formed, a change in temperature, or a new smell. A physical change such as melting shows none of these, only a change of state.

The clearest test is reversibility. Melting, freezing, dissolving and boiling can all be reversed by changing the conditions back, but a chemical change is normally not reversed by simple physical means. Heating blue hydrated copper(II) sulfate turns it white and drives off the water, and adding water brings the blue colour back, but this is still a chemical change because a new compound, anhydrous copper(II) sulfate, has been formed.

• Physical change: no new substance, usually reversible, chemical identity unchanged.
• Chemical change: a new substance is formed, usually irreversible, atoms rearranged.
• Evidence of chemical change: colour change, gas, precipitate, temperature change, new smell.
• Dissolving is physical; burning, rusting and neutralising are chemical.

Collision theory: why reactions happen at all

For a reaction to take place the reacting particles must collide, and they must collide with at least a certain minimum energy. That minimum is the activation energy, Ea. A collision with enough energy is an effective collision and leads to a reaction; a gentle collision simply bounces the particles apart.

The rate of reaction is the number of effective collisions per unit time. Anything that increases the number of effective collisions per second increases the rate. There are only three ways to do that: raise the number of particles per unit volume, raise the frequency of collisions between particles, or raise the kinetic energy of the particles so that more of them carry at least the activation energy.

A catalyst works by a different route. It provides an alternative pathway with a lower activation energy, so a greater fraction of the collisions that are already happening now carry enough energy to react. The catalyst itself is not changed or used up.

This one idea explains every factor in this topic. If you can say what a change does to the number, the frequency or the energy of collisions, you can explain any rate question.

• Particles must collide with energy greater than or equal to the activation energy.
• Rate = number of effective collisions per unit time.
• More particles per unit volume and a higher temperature both increase the collision frequency.
• A catalyst lowers the activation energy, so more collisions are effective.

The factors that affect the rate of reaction

Concentration and pressure. Increasing the concentration of a solution puts more particles into the same volume, so there are more particles per unit volume, collisions are more frequent and the rate increases. On a graph of rate against concentration the line rises. For gases, increasing the pressure squeezes the same number of particles into a smaller volume, which is exactly the same effect: more particles per unit volume, more frequent collisions and a faster rate.

Surface area. A solid can only react where its surface touches the other reactant. Breaking a lump into smaller pieces, or using a powder instead of chips, exposes far more surface. There are more collisions per unit time, so the rate increases, even though the total mass of solid is unchanged.

Temperature. Raising the temperature gives every particle more kinetic energy. The particles move faster, so they collide more often, and each collision is more energetic, so a larger fraction of collisions exceed the activation energy. Both effects increase the rate, which is why temperature has such a large effect.

Catalyst. Adding a catalyst lowers the activation energy, so more collisions are effective and the rate increases. Removing it reverses the effect. Enzymes are biological catalysts and behave in exactly the same way.

Measuring the rate of reaction

Loss of mass. When a gas escapes, the mass of the reaction mixture falls. Put a flask of marble chips and dilute hydrochloric acid on a balance and record the mass every 30 seconds. Carbon dioxide escapes and the reading drops. The steeper the curve on a mass-time graph, the faster the reaction.

Volume of gas collected. Use a gas syringe or an inverted measuring cylinder. Magnesium ribbon and dilute hydrochloric acid give off hydrogen, which you can collect and read off at intervals. The volume-time graph rises quickly at first and then levels off when the reaction finishes.

Disappearing cross. Mix sodium thiosulfate solution with dilute acid; a pale yellow precipitate of sulfur slowly forms and the mixture turns cloudy. Time how long it takes before a black cross drawn on a tile under the flask can no longer be seen. The rate is proportional to 1 / time.

Interpreting the graph. The gradient of a rate graph is the rate at that moment. A steep gradient means a fast reaction, and a curve that flattens to a plateau means the reaction has stopped because a reactant has been used up. A curve that starts steeper, or reaches the plateau sooner, shows a faster reaction, and the same final level means the same amount of product was formed.

• Measure the rate by the loss in mass when a gas escapes, or by the volume of gas collected.
• The gradient of the curve gives the rate at that instant.
• The steeper the initial gradient, the faster the reaction.
• A plateau means the reaction has finished because a reactant has been used up.

Reversible reactions and dynamic equilibrium

Some reactions are reversible: the products can react together to re-form the reactants. A reversible reaction is written with the equilibrium symbol ⇌. The forward reaction goes from left to right and the reverse reaction goes from right to left.

Heating blue hydrated copper(II) sulfate, CuSO4.5H2O, drives off the water of crystallisation and leaves white anhydrous copper(II) sulfate: CuSO4.5H2O(s) ⇌ CuSO4(s) + 5H2O(l). Adding water reverses the change and the blue colour returns. Cobalt(II) chloride behaves in the same way: pink hydrated CoCl2.6H2O turns blue anhydrous CoCl2 on heating, and the pink colour returns when water is added, which is why blue cobalt(II) chloride paper is used to test for water.

If a reversible reaction happens in a closed system, where nothing can escape, it eventually reaches dynamic equilibrium. At equilibrium the rate of the forward reaction equals the rate of the reverse reaction, so the concentrations of the reactants and products stop changing. Both reactions are still happening - dynamic means they have not stopped - they simply cancel each other out.

Equilibrium does not mean equal amounts. There may be much more product than reactant, or the other way round; what is constant is the concentrations, not their ratio. Changing the conditions moves the equilibrium to a new position. Raising the temperature favours the endothermic direction, because the system absorbs the extra heat. Raising the pressure favours the side with fewer gas molecules, because that relieves the pressure. Adding a reactant pushes the equilibrium towards the products. A catalyst speeds up the forward and reverse reactions equally, so it reaches equilibrium faster but does not change the position of equilibrium or the final yield.

The Haber process and the Contact process

The Haber process makes ammonia, which is used to make fertilisers: N2(g) + 3H2(g) ⇌ 2NH3(g). The nitrogen comes from the air and the hydrogen comes from methane, which is natural gas. The typical conditions are about 450 °C, 200 atm (20 000 kPa) and an iron catalyst.

The forward reaction is exothermic and goes from four gas molecules to two, so a lower temperature and a higher pressure would both give a better yield of ammonia at equilibrium. Neither extreme is used, because the conditions are a compromise. At a low temperature the rate would be far too slow, so 450 °C is chosen to keep the rate high enough. A very high pressure would give a better yield but needs thicker, more expensive equipment and is more dangerous, so 200 atm is used. The iron catalyst speeds up both directions equally, so the reaction reaches equilibrium faster without changing the yield.

The Contact process makes sulfur trioxide, which is used to make sulfuric acid: 2SO2(g) + O2(g) ⇌ 2SO3(g). The sulfur dioxide comes from burning sulfur or roasting sulfide ores, and the oxygen comes from the air. The typical conditions are about 450 °C, 2 atm (200 kPa) and a vanadium(V) oxide catalyst.

The forward reaction is again exothermic and goes from three gas molecules to two, so a lower temperature and a higher pressure would give a better yield of sulfur trioxide. A pressure of only 2 atm is used because the yield is already high enough at this pressure, so the extra cost and hazard of high pressure are not worth it. The temperature of 450 °C is again a compromise between a good yield and a usable rate, and the vanadium(V) oxide catalyst lowers the activation energy so the reaction is fast enough at that temperature.

• Haber: N2(g) + 3H2(g) ⇌ 2NH3(g), about 450 °C, 200 atm, iron catalyst.
• Contact: 2SO2(g) + O2(g) ⇌ 2SO3(g), about 450 °C, 2 atm, vanadium(V) oxide catalyst.
• High pressure and low temperature favour the product in both processes, because both forward reactions are exothermic and lose gas molecules.
• The chosen conditions are a compromise between yield, rate, cost and safety.

Redox: oxidation, reduction and redox agents

A redox reaction is one in which oxidation and reduction happen at the same time. If one substance is oxidised, another must be reduced, because the oxygen or the electrons lost by one are gained by the other.

At Core level, oxidation is gain of oxygen and reduction is loss of oxygen. In 2Mg + O2 -> 2MgO the magnesium is oxidised because it gains oxygen. At Supplement level you must also use electrons and oxidation numbers. Oxidation is loss of electrons and reduction is gain of electrons, remembered as OIL RIG. Oxidation also means an increase in oxidation number, and reduction means a decrease.

An oxidation number shows the charge an atom would have if the bonding were fully ionic. Four rules let you work one out: the oxidation number of an element in its uncombined state is zero; the oxidation number of a monatomic ion is the same as the charge on the ion; the sum of the oxidation numbers in a compound is zero; and the sum of the oxidation numbers in an ion equals the charge on the ion. Roman numerals in names, such as copper(II) oxide and iron(III) oxide, tell you the oxidation number of the metal. For example, in the manganate(VII) ion, MnO4-, let the oxidation number of manganese be x. Then x + 4(-2) = -1, so x = +7. In sulfur dioxide, SO2, x + 2(-2) = 0, so the sulfur is +4.

An oxidising agent is a substance that oxidises another substance and is itself reduced; it takes electrons from the other substance. Common oxidising agents are oxygen, chlorine, hydrogen peroxide and acidified potassium manganate(VII). A reducing agent is a substance that reduces another substance and is itself oxidised; it gives electrons to the other substance. Common reducing agents are hydrogen, carbon, carbon monoxide and reactive metals such as zinc and magnesium. The direction is the trap - the name tells you what the substance does to the other substance, not what happens to itself. Two colour changes let you spot a redox reaction: acidified potassium manganate(VII) is reduced from purple MnO4- to almost colourless Mn2+, so a solution that turns it from purple to colourless is a reducing agent, while aqueous potassium iodide is oxidised from colourless iodide ions to brown iodine, so a solution that turns it brown is an oxidising agent.

• Oxidising agent: oxidises another substance and is itself reduced.
• Reducing agent: reduces another substance and is itself oxidised.
• Acidified potassium manganate(VII) goes from purple to colourless when reduced.
• Aqueous potassium iodide goes from colourless to brown when oxidised to iodine.