Ch 2 · Rate of Reaction · O Level Chemistry 5070
Topic 2 of 4 in Chemical Reactions — O Level Chemistry 5070 — 38 lessons.
Introduction
Why does a powder burn faster than a lump, and why does food keep longer in a fridge? Rate of reaction is a heavily examined practical topic: you must set up an experiment, read a graph and explain every result using collision theory. The vocabulary is small but must be exact.
Key ideas — Measuring rate and reading rate graphs
Rate is the change in amount of a reactant or product per unit time. Two standard methods follow a reaction that produces a gas, such as calcium carbonate with hydrochloric acid.
🧑🏫 Teacher's explanation
Think of a chemical reaction like a crowd of people trying to shake hands in a busy hallway. The 'rate' is simply how many handshakes happen every minute. If the hallway is packed (high concentration), handshakes happen constantly. If people are walking very fast (high temperature), they bump into each other more often, increasing the rate.
📖 Definitions
🔍 Worked example — Easy
A reaction produces gas. The volume of gas collected is 0 ext cm^3 at t=0 ext s and 40 ext cm^3 at t=100 ext s. Calculate the average rate of reaction in cm^3/s.
🔍 Worked example — Standard
In an experiment measuring mass loss, the mass of a flask decreases from 50.0 g to 48.2 g over 5 minutes. Calculate the rate of mass loss in g/min.
🔍 Worked example — Exam-style
A graph of volume vs time shows a curve that starts steep and levels off. At t=20 s, the volume is 15 cm^3. At t=40 s, the volume is 25 cm^3. Calculate the rate of reaction between 20\ ext s and 40\ ext s.
⚠️ Common errors
🧠 Quick check 1
🧠 Quick check 2
🎯 Exam question
(a) Calculate the average rate of gas production between 0 and 20 seconds. [2] (b) Explain why the volume of gas stays constant after 60 seconds. [1] [3 marks]
Key ideas — Factors affecting rate
how a catalyst works. A catalyst decreases the activation energy by providing an alternative reaction pathway with a lower energy barrier. More of the colliding particles therefore have energy at or above the activation energy, so a greater proportion of collisions are successful and the rate increases. The catalyst does not change the enthalpy change of the reaction.
🧑🏫 Teacher's explanation
Imagine you are at a crowded music festival. If you want to move from one side to the other, your speed depends on how many people are in your way. If the crowd is sparse (low concentration), you move quickly. If the crowd is packed tight (high concentration), you bump into people constantly, slowing you down—but wait, in chemistry, bumping into people is actually what makes the 'reaction' happen! So, more bumps means more reactions.
📖 Definitions
🔍 Worked example — Easy
A reaction produces 20 cm^3 of gas in 40 seconds. Calculate the rate of reaction in cm^3/s.
🔍 Worked example — Standard
A student investigates the reaction between Mg and HCl. The mass of the flask decreases from 50.0 g to 48.2 g in 100 seconds. Calculate the rate of mass loss in g/s.
🔍 Worked example — Exam-style
Explain, using collision theory, why increasing the temperature increases the rate of reaction.
⚠️ Common errors
🧠 Quick check 1
🧠 Quick check 2
🎯 Exam question
Key ideas — Collision theory
explaining each factor. Particles can react only if they collide with energy equal to or greater than the activation energy. The rate depends on the frequency of collisions and the proportion of collisions that are successful.
🧑🏫 Teacher's explanation
Imagine a busy dance floor. For a 'reaction' (a dance move) to happen, two people must bump into each other (a collision) with enough energy to actually start the move. If the dance floor is crowded (high concentration), people bump into each other much more often. If everyone is dancing wildly and fast (high temperature), they bump into each other more often, and when they do bump, they are much more likely to have enough energy to start the move.
📖 Definitions
🔍 Worked example — Easy
Explain why increasing the concentration of a reactant increases the rate of reaction using collision theory.
🔍 Worked example — Standard
A reaction between Mg(s) and HCl(aq) is faster when using Mg powder than when using a Mg ribbon. Explain why.
🔍 Worked example — Exam-style
Calculate the percentage increase in the number of successful collisions if the temperature is increased, assuming the frequency of collisions stays constant but the proportion of particles with energy > E_a doubles.
⚠️ Common errors
🧠 Quick check 1
🧠 Quick check 2
🎯 Exam question
(a) State the effect of using powder instead of ribbon on the rate of reaction. [1] (b) Explain this effect in terms of collision theory. [2] [3 marks]
Worked example
Question. A student adds 2.0 g of marble chips (calcium carbonate) to excess dilute hydrochloric acid and records the volume of carbon dioxide. After 20 s the volume is 30 cm³; after 60 s it is 60 cm³; the volume stops rising at 72 cm³ after 100 s. The experiment is repeated with 2.0 g of powdered marble. (a) Calculate the average rate of reaction during the first 20 s. [1] (b) Describe two differences between the graph for the powder and the graph for the chips. [2] (c) Explain, using collision theory, why the powder reacts faster. [2]
🔍 Worked example — Easy
Calculate the average rate of reaction for the reaction Mg + 2HCl arrow MgCl_2 + H_2 if the volume of H_2 gas increases from 0 cm^3 to 45 cm^3 in 15 s.
🔍 Worked example — Standard
A reaction produces 80 cm^3 of gas in 40 s. If the concentration of the reactant is doubled, the rate increases to 3.0 cm^3/s. Calculate the original average rate.
🔍 Worked example — Exam-style
A student investigates the reaction between CaCO_3(s) and HCl(aq). The mass of the flask decreases from 50.0 g to 48.5 g in 100 s. Calculate the rate of mass loss in g/s.
⚠️ Common errors
Common mistakes
Practice
Answers Add acid to magnesium in a conical flask and immediately fit the bung connected to a gas syringe [M1]; start the stopwatch and record the volume of hydrogen at regular time intervals [M1]; plot volume against time and use the gradient as the rate [A1]. Curve steep at first, then flattening to a horizontal line [B1]; reactants are used up, so their concentration falls [B1]; collisions become less frequent and the reaction slows and then stops [B1]. More acid particles per unit volume [B1]; more frequent collisions with the zinc surface, so more successful collisions per second [B1]. Particles gain kinetic energy and collide more frequently [B1]; a much greater proportion of the particles have energy equal to or greater than the activation energy [B1]; so the proportion of successful collisions rises sharply [B1].