Collision theory

1. For a reaction to happen, the reactant particles must first collide with each other.
2. Most collisions do not lead to a reaction, because the particles do not have enough energy and simply bounce apart.
3. A collision only leads to a reaction if the colliding particles have kinetic energy equal to or greater than the activation energy, $E_a$.
4. The rate of reaction depends on the number of successful collisions per unit time.
5. Four ideas are used to explain changes in rate:
   1. The number of particles per unit volume, which describes how crowded the reacting particles are.
   2. The frequency of collisions, which is the number of collisions between particles per unit time.
   3. The kinetic energy of the particles, which increases as they move faster.
   4. The activation energy, which is the energy barrier that colliding particles must overcome.

Concentration and pressure

1. Increasing the concentration of a solution increases the number of particles per unit volume.
2. The particles are closer together, so they collide more frequently.
3. More frequent collisions mean more successful collisions per unit time, so the rate increases.
4. Increasing the pressure of a gas squeezes the same number of particles into a smaller volume, which also increases the number of particles per unit volume.
5. The gas particles therefore collide more frequently, so the rate increases in the same way as for concentration.
6. Neither change gives the particles more energy, so the proportion of collisions that are successful stays the same.

Surface area

1. Only the particles on the surface of a solid can collide with particles of the other reactant.
2. Cutting a solid into smaller pieces exposes more of its particles at the surface.
3. More particles are available to collide, so there are more frequent collisions and the rate increases.
4. The energy of the particles does not change, so it is the frequency of collisions, not the proportion that succeed, that increases.

Temperature

1. Increasing the temperature gives the particles more kinetic energy, so they move faster.
2. Faster particles collide more frequently.
3. A greater proportion of the particles also have energy equal to or greater than the activation energy.
4. A greater proportion of collisions are therefore successful, so there are many more successful collisions per unit time.
5. The increase in the proportion of successful collisions is the bigger effect, which is why a small rise in temperature can make a large difference to the rate.

Catalysts and activation energy

1. A catalyst provides an alternative reaction pathway with a lower activation energy.
2. The energy of the particles is unchanged, but a greater proportion of collisions now have energy equal to or greater than the lower activation energy.
3. There are more successful collisions per unit time, so the rate increases.
4. On a reaction pathway diagram, the catalysed pathway has a lower peak, while the energies of the reactants and products stay the same.
5. The overall energy change of the reaction, $\Delta H$, is therefore the same with or without the catalyst.
6. Enzymes work in the same way, lowering the activation energy of reactions in living cells so that they are fast at body temperature.
7. At high temperatures an enzyme's structure changes, so it can no longer provide the lower-energy pathway.
8. Removing a catalyst means the reaction must again follow the pathway with the higher activation energy, so the rate falls.

Evaluating methods for measuring rate

1. Gas syringe: this measures the volume of gas directly, usually to the nearest $1{cm}^3$, and works for any gas.
   1. Some gas escapes in the moment between adding the reactant and fitting the bung, so the first readings are slightly low.
   2. A syringe can stick, and it cannot hold more than its maximum volume, often $100{cm}^3$.
2. Upturned measuring cylinder or burette over water: this is cheap and simple to set up.
   1. It is unsuitable for gases that dissolve in water, such as carbon dioxide or ammonia, because some of the gas dissolves instead of being collected.
3. Loss in mass on a balance: this is quick to read and can record many readings, especially with a balance linked to a computer.
   1. It works well for carbon dioxide, which has $M_r=44$, because the mass lost is large enough to measure.
   2. It works poorly for hydrogen, which has $M_r=2$, because the mass lost is too small to measure accurately.
   3. A cotton wool plug is needed so that acid spray does not escape and add to the loss in mass.
4. Disappearing cross: this needs very little apparatus.
   1. Deciding when the cross disappears is a judgement, so the timings vary between observers.
   2. It gives only one time for each experiment, rather than a full curve showing how the rate changes.
5. Ways to improve any of these methods include:
   1. Repeating each experiment and calculating a mean, so that unusual results can be spotted and ignored.
   2. Using a water bath so that the temperature stays constant.
   3. Using a light sensor and data logger in place of the eye for the disappearing cross.
   4. Hanging the solid on a thread or placing it in a small tube inside the flask, so that the bung is already in place when the reaction starts.