Topic 6: The rate and extent of chemical change

Some reactions are over in a flash while others take years. In this topic you will learn how chemists measure reaction rate, why factors such as concentration, temperature and catalysts change it, and how collision theory explains all of these effects. You will also meet reversible reactions and the idea of dynamic equilibrium.

Measuring the rate of reaction

The rate of a reaction tells you how quickly reactants are used up or products are made. You can follow a reaction by measuring either the amount of reactant that disappears or the amount of product that forms over time. Common methods include measuring the volume of gas given off with a gas syringe, measuring the loss in mass on a balance as gas escapes, and timing how long it takes for a mixture to turn cloudy enough to hide a cross drawn under the flask. Whatever you measure, you record it at regular time intervals and plot the results on a graph, usually amount of product on the y-axis against time on the x-axis. The line is steepest at the start when reactant concentration is highest, then gradually levels off, and becomes flat (horizontal) when the reaction has finished. The mean rate can be worked out by dividing the quantity of reactant used or product made by the time taken.

Calculating mean rate and using tangents (Higher)

The mean rate of reaction is found using the formula: mean rate = quantity of reactant used (or product formed) divided by time taken. Units are usually given as grams per second (g/s), centimetres cubed per second (cm3/s), or moles per second (mol/s). This gives an average over a chosen period, but the actual rate changes from moment to moment. To find the rate at a single instant you draw a tangent to the curve at that point: a straight line that just touches the curve. The gradient of the tangent equals the rate at that time. You calculate the gradient by reading the change in the y-value divided by the change in the x-value (rise over run) for two well-spaced points on the tangent. The tangent is steepest at the start of the reaction, showing the rate is fastest then, and its gradient falls towards zero as the reaction finishes.

Collision theory and activation energy

Collision theory explains why reactions happen and why their rates change. For particles to react they must collide with each other, and they must collide with enough energy. The minimum amount of energy that colliding particles need in order to react is called the activation energy. Collisions that do not have this much energy simply bounce apart unchanged. Anything that increases either the frequency of collisions or the proportion of collisions that have enough energy will increase the rate of reaction. This single idea lies behind every rate factor: concentration, pressure, surface area, temperature and catalysts all work by changing how often particles collide successfully.

Effect of concentration and pressure

Increasing the concentration of a dissolved reactant means there are more reactant particles in the same volume of solution. The particles are closer together, so they collide more often, and a higher collision frequency means a faster rate. The same idea applies to gases: increasing the pressure of a gas squeezes the same number of particles into a smaller volume, so they too are closer together and collide more frequently, increasing the rate. Note that raising concentration or pressure increases how often collisions happen but does not change the energy of each collision.

Effect of surface area

When a solid reacts with a solution or a gas, only particles on the surface of the solid can take part. Breaking a solid into smaller pieces, or grinding it into a powder, increases its surface area to volume ratio. With more surface exposed, more reactant particles are available to be hit, so the frequency of successful collisions rises and the reaction speeds up. This is why a powdered metal reacts much faster than the same mass of metal in one large lump, and it is also why fine dusts such as flour can be dangerously explosive.

Effect of temperature

Raising the temperature has two effects, both of which speed up a reaction. First, the particles gain kinetic energy and move faster, so they collide more frequently. Second, and more importantly, a greater proportion of the particles now have energy equal to or greater than the activation energy, so a larger fraction of collisions are successful. A rough rule of thumb is that the rate of many reactions roughly doubles for every 10 degrees Celsius rise in temperature. Because both collision frequency and collision energy increase, temperature is one of the most powerful ways to change a reaction rate.

Catalysts and enzymes

A catalyst is a substance that speeds up a reaction without being used up itself, so it is not part of the overall reaction equation and can be recovered unchanged at the end. Catalysts work by providing an alternative reaction pathway that has a lower activation energy. With a lower energy barrier, a greater proportion of collisions have enough energy to react, so the rate increases. Different reactions need different catalysts. Enzymes are biological catalysts: large protein molecules that speed up the reactions taking place in living cells, such as the fermentation of sugar into ethanol by yeast. Because catalysts are not consumed, only a small amount is needed and it can be reused, which makes industrial processes cheaper and more sustainable.

Reversible reactions and energy changes

In a reversible reaction the products can react together to remake the original reactants. We show this with the symbol <=> in the equation, for example A + B <=> C + D. The direction of a reversible reaction can be changed by altering the conditions. A good example is heating hydrated copper sulfate, which is blue, to drive off water and form white anhydrous copper sulfate; adding water reverses the change. If a reversible reaction is exothermic in one direction (releasing energy to the surroundings), it must be endothermic by exactly the same amount in the opposite direction (taking in energy). The energy released going one way equals the energy absorbed going the other way.

Dynamic equilibrium and changing conditions (Higher)

When a reversible reaction takes place in a closed container (one that lets nothing in or out), it can reach equilibrium. At equilibrium the forward and reverse reactions are still happening, but they occur at exactly the same rate, so the amounts of reactants and products stay constant. This is called dynamic equilibrium because the reactions never stop, even though nothing appears to change. The relative amounts of reactants and products at equilibrium depend on the conditions. Le Chatelier's principle states that if a condition is changed, the equilibrium shifts to oppose that change. Increasing the concentration of a reactant shifts the equilibrium towards the products. Increasing temperature shifts equilibrium in the endothermic direction; decreasing it favours the exothermic direction. For reactions involving gases, increasing pressure shifts the equilibrium towards the side with fewer molecules of gas.