Rate of reaction is the change in the amount of a reactant or product per unit time. Average rate is the total change divided by the total time; instantaneous rate is the gradient of the tangent to the curve at a given moment.
Typical units are cm³ s⁻¹, g s⁻¹ or mol dm⁻³ s⁻¹.
This page covers one Form 4 content standard: the concept of rate of reaction. It is the foundation of the whole chapter, because every later idea, the factors that change a reaction’s speed, collision theory and catalysts, is really an explanation of why the rate goes up or down. Get the definition, the way rate is measured, and the units secure here, and the rest of the chapter follows easily.
What “rate of reaction” means
A chemical reaction turns reactants into products over a period of time. The rate of reaction is a measure of how fast that change happens. Formally, it is the change in the quantity of a reactant or a product per unit time.
- As a reaction proceeds, the amount of reactant decreases and the amount of product increases.
- The rate tells you how quickly one of those quantities is changing.
- The rate is normally fastest at the start, slows down as reactants are used up, and becomes zero when the reaction stops (when a reactant is completely used up).
You can measure rate using whichever quantity is easiest to follow in a particular reaction, a volume of gas, a loss of mass, a change in colour, or the time for a precipitate to hide a mark.
Average rate and instantaneous rate
Two versions of “rate” appear in the exam, and you must keep them apart.
Average rate = (total change in quantity) ÷ (total time taken). This gives a single value describing the whole interval or the whole reaction.
Instantaneous rate = the rate at one particular moment, found from the gradient of the tangent to the curve at that point.
Because the reaction slows as it goes, the instantaneous rate keeps falling. The initial rate (at the very start) is the largest, and it equals the gradient of the tangent at time = 0. The average rate over the whole reaction sits somewhere between the fast start and the zero finish.
How rate is measured in the laboratory
The method depends on what visibly changes:
- Volume of gas released, collect the gas in a syringe or over water and record its volume at regular times (e.g. the reaction of a metal or carbonate with acid). Rate has units cm³ s⁻¹.
- Loss of mass, stand the flask on a balance; as gas escapes, the mass falls. Rate has units g s⁻¹. This suits reactions that give off a heavy gas such as carbon dioxide.
- Formation of a precipitate (turbidity), in the sodium thiosulfate and acid reaction, sulfur clouds the mixture; you time how long a cross under the beaker takes to disappear. A shorter time means a faster rate.
- Change in concentration, measured as mol dm⁻³ s⁻¹, the most fundamental unit.
Whichever quantity you choose, the rate is always “that quantity per second”, so the unit always ends in s⁻¹.
Worked example
Question. In the reaction between excess zinc and dilute hydrochloric acid, hydrogen gas is collected. A total of 48 cm³ of gas is produced in the first 40 s. (a) Calculate the average rate of reaction over this period. (b) State and explain how the instantaneous rate at 40 s compares with the rate at the start.
Step 1, Write the formula for average rate. Average rate = (change in volume of gas) ÷ (time taken).
Step 2, Substitute the values. Average rate = 48 cm³ ÷ 40 s.
Step 3, Calculate and add the unit. Average rate = 1.2 cm³ s⁻¹.
Step 4, Compare the moments. The rate at the start is higher than the rate at 40 s. At the start the acid is most concentrated, so the reaction is fastest; by 40 s some acid has been used up, its concentration has fallen, and the reaction has slowed.
Answer. (a) The average rate is 1.2 cm³ s⁻¹. (b) The instantaneous rate at 40 s is lower than the initial rate, because the reactant concentration has decreased as the reaction proceeds.
Practice question
Question. A flask containing marble chips (calcium carbonate) and dilute hydrochloric acid is placed on a balance. Carbon dioxide escapes, and the mass falls by 3.2 g in 80 s. (a) Calculate the average rate of reaction in g s⁻¹. (b) Explain why measuring the loss of mass is a suitable method for this reaction.
Answer. (a) Average rate = change in mass ÷ time = 3.2 g ÷ 80 s = 0.04 g s⁻¹. (b) The reaction releases carbon dioxide, a gas that escapes from the flask, so the mass of the flask and its contents falls measurably over time. Because the loss of mass can be read directly from the balance at regular intervals, it gives a convenient and continuous measure of how fast the reaction is going.
Exam tip
When you define rate of reaction, always name the quantity that changes and add “per unit time”, “the change in volume of gas per unit time”, not just “how fast the reaction is”. In a calculation, always attach the unit (cm³ s⁻¹, g s⁻¹ or mol dm⁻³ s⁻¹); a bare number loses the mark. And remember the shape of the story: the rate is largest at the beginning and falls to zero at the end, so any answer claiming a constant rate throughout is wrong unless a reactant is being continuously replaced.
Reading rate questions well
Rate questions in Paper 2 often give you data, a table of gas volume against time, or a mass reading at two times, and ask for the average rate over a stated interval. The move is always the same: find the change in the quantity between the two times, divide by the time interval, and quote the unit. If the question asks for the rate “at” a particular time rather than “over” an interval, it wants the instantaneous rate, which on a graph is the gradient of the tangent at that point.
A second common task is to compare the speed at two moments and explain the difference. The explanation almost always comes back to concentration: as reactants are consumed, their concentration falls, collisions become less frequent, and the rate drops. That link, falling concentration, slower reaction, is the bridge from this standard into collision theory later in the chapter, so it is worth stating clearly every time.
Where this fits
This is content standard 7.1 of the Rate of Reaction chapter, the first of five. It sets up the factors that change the rate, collision theory and catalysts, and it is examined heavily in the practical paper as part of SPM Chemistry (4541). Reinforce it with the chapter revision notes and practise the calculations in the worked examples. In our online 1-to-1 SPM Chemistry lessons, taught in English, from RM50/hr, our teachers drill the definition-plus-unit habit until a rate answer is automatic, because this is where students most often drop an easy mark.
Quick recap
- Rate of reaction = change in the amount of a reactant or product per unit time.
- Average rate = total change ÷ total time; instantaneous rate = gradient of the tangent.
- Measure by gas volume (cm³ s⁻¹), loss of mass (g s⁻¹), precipitate/turbidity, or concentration (mol dm⁻³ s⁻¹).
- Rate is fastest at the start and falls to zero when a reactant runs out.
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