Your revision route
What you’ll learn
- Use elapsed time and the change in a reactant or product quantity to obtain a mean reaction rate.
- Draw and interpret quantity–time graphs and use tangents to compare rates.
- Explain concentration, pressure, surface area, temperature and catalyst effects using collision theory.
- Carry out both gas-volume and colour-or-turbidity concentration investigations safely.
- Describe reversible reactions, opposite energy changes and dynamic equilibrium in a closed system.
- Higher Tier: calculate instantaneous rate from a tangent and use mol/s where appropriate.
- Higher Tier: predict equilibrium changes caused by concentration, temperature and gas pressure.
Build the big picture
Key ideas
Rate is change divided by time
A reaction graph is a speedometer built from chemistry: quantity changes vertically while time advances horizontally.
- Mean rate = quantity of reactant used ÷ time, or quantity of product formed ÷ time.
- Mass measurements give g/s; gas-volume measurements give cm³/s.
- On a quantity–time graph, a steeper gradient means a faster rate.
- The final plateau shows total quantity, not the rate at that instant.
The punchline: Write the measured quantity and unit before calculating the rate.
A tangent captures local curve steepness
A curve changes steepness continuously; a tangent borrows its direction at one chosen point.
- All routes: draw a tangent that just touches the curve at the required point and follows its local direction.
- All routes: compare slopes qualitatively; the steeper tangent represents the faster reaction rate.
- Higher Tier: at a specified time, choose two well-separated points on the tangent and calculate vertical change ÷ horizontal change.
- Higher Tier: keep the graph's axis units in the numerical gradient; a quantity measured in moles gives mol/s.
The punchline: All routes use tangent steepness as a rate measure; Higher Tier calculates the numerical gradient at a specified time.
Five levers can change reaction rate
Chemists speed reactions by crowding particles, exposing more surface, heating collisions or lowering the energy barrier.
- In the simple collision model, doubling particle concentration can approximately double collision frequency and rate when other conditions stay fixed; this is not a universal rate law.
- Increasing gas pressure, or increasing a solid's surface area by using smaller pieces, generally increases rate.
- Increasing temperature generally increases rate.
- A suitable catalyst increases rate without being used up overall.
The punchline: Name the changed factor, then explain its collision mechanism.
From changed condition to changed rate
A full explanation connects the condition to collision frequency or energy, then to the number of successful collisions each second.
- Change conditionIncrease concentration, gas pressure, surface area or temperature, or add a suitable catalyst.
- Change collisionsConcentration, pressure and surface area raise relevant collision frequency; temperature raises frequency and energy; a catalyst lowers activation energy.
- Change successesMore collisions per second have energy at least equal to the activation energy.
- Change rateReactant is used or product forms more quickly, so the graph is initially steeper.
Most collisions are merely meetings
Particles react only when a collision has suitable orientation and at least the activation energy required for rearrangement.
- More particles in a volume or more exposed solid surface increases collision frequency.
- Higher temperature makes particles move faster, causing more frequent collisions and a larger energetic fraction.
- Activation energy is the minimum collision energy needed for reaction.
- Rate depends on successful collisions per second, not the total number of particles alone.
The punchline: Complete the chain: factor → collisions → successful collisions per second → rate.
A catalyst opens a lower pass through the energy hill
The starting and finishing energy levels stay put; only the activation barrier is lowered by an alternative pathway.
- A catalyst changes reaction rate but is not used up overall and does not appear in the overall equation.
- Catalyst choice depends on the reaction, while living systems use enzymes to perform this role.
- A reaction profile shows a lower activation-energy peak for the catalysed pathway.
- Higher Tier: at equilibrium a catalyst speeds forward and reverse reactions equally, so it does not move the equilibrium position.
The punchline: Lower activation energy; do not claim the catalyst gives particles more energy.
Make the model move
Interactive checkpoint
Touch the science. Change a state, build a route or test a relationship.
Measure the chemical pace
Calculate a mean rate
Change product quantity and time. The calculator returns quantity formed per second.
rate = quantity ÷ time
Mean rate2 cm³/s
Mean rate = quantity formed ÷ time. Units follow the measured quantity: cm³/s for gas volume, g/s for mass or Higher Tier mol/s.
Make the mechanism explicit
Match each factor to its particle-level effect
Pair the changed condition with the immediate collision-theory reason for a faster reaction.
Different factors act through collision frequency, collision energy or activation energy. Temperature affects both frequency and energy.
Disturb the balance
Higher Tier: how does equilibrium respond?
Switch conditions for N₂(g) + 3H₂(g) ⇌ 2NH₃(g), whose forward reaction is exothermic.
The system uses some added H₂ by favouring the forward reaction until a new equilibrium forms.
1 of 4 states explored
Predict the shift that counteracts the change. These are qualitative changes in relative equilibrium amounts, not claims that a reaction stops.
One practical needs two different clocks
A gas syringe counts product continuously; a disappearing cross marks one subjective turbidity threshold.
- Investigate concentration using a gas-volume method and a separate colour-or-turbidity method.
- Keep temperature, reactant volumes, apparatus and solid surface area controlled as appropriate.
- Gas data can produce a full curve; a disappearing-cross result often uses rate proportional to 1/time.
- Repeat conditions, calculate means and identify anomalous results.
The punchline: State exactly what is measured and how it represents rate.
Some reaction pathways run in reverse
In a reversible reaction, products can react to re-form reactants; changing conditions can favour either direction.
- Use the reversible arrow ⇌ rather than two unrelated one-way equations.
- Reversing the reaction flips the sign of the energy transfer but not its magnitude.
- Changing conditions can change the direction of a reversible reaction.
The punchline: Reverse direction means opposite energy transfer of equal magnitude.
Equilibrium is busy, not frozen
At dynamic equilibrium, molecules keep reacting in both directions while bulk concentrations remain constant.
- Keep every reacting substance inside the apparatus.
- Once equilibrium is established, neither direction is faster than the other.
- Equal rates do not require equal concentrations of reactants and products.
- Constant macroscopic amounts hide continuing particle-level change.
The punchline: Say equal rates in a closed system, not equal amounts or stopped reactions.
How a product–time curve tells its story
The curve's gradient gives rate. Its final plateau gives total product, which is a different quantity.
- Early steep curveReactant concentration is high, so successful collisions occur frequently.
- Curve flattensIn a typical batch reaction, reactants are used up, so their concentration falls; collision frequency and rate usually fall too.
- PlateauFor a reaction going to completion, a plateau means a reactant is used up. At dynamic equilibrium, bulk amounts stay constant while both microscopic reactions continue.
- TangentAll routes: a tangent's slope is a measure of rate; a steeper tangent means a faster rate. Higher Tier: calculate its numerical gradient at a specified time.
Higher Tier: three ways to disturb equilibrium
Le Chatelier's Principle predicts the response that counteracts an imposed change, given the reaction information.
- ConcentrationIncreasing the concentration of a dissolved or gaseous reactant, or decreasing the concentration of a product, favours the forward direction.
- TemperatureIncreasing temperature favours the endothermic direction; decreasing it favours the exothermic direction.
- Gas pressureIncreasing pressure favours the side with fewer gaseous molecules; equal gas counts give no shift.
Higher Tier: equilibrium pushes back
After a disturbance, the equilibrium mixture shifts in the way that partially reduces its effect.
- Use the supplied equation and energy information before predicting a shift.
- Increasing a reactant concentration favours the direction that uses that reactant.
- The equilibrium position changes until a new dynamic equilibrium is established.
- Le Chatelier's Principle predicts direction qualitatively; it does not calculate the final amounts.
The punchline: Name the change, the counteracting direction and the relative product effect.
Higher Tier: heat and pressure favour different equation features
Temperature follows energy direction; pressure counts gaseous molecules. Mixing those rules produces confident nonsense.
- Increasing temperature favours the endothermic direction; decreasing temperature favours the exothermic direction.
- Increasing gas pressure favours the side with fewer gaseous molecules in the balanced equation.
- Decreasing gas pressure favours the side with more gaseous molecules.
- If both sides contain equal numbers of gaseous molecules, pressure does not change equilibrium position.
The punchline: Use energy labels for temperature and gas coefficients for pressure.
Words worth knowing
Key definitions
- rate of reaction
- The change in quantity of a reactant or product per unit time.
- activation energy
- The minimum collision energy required for a reaction to occur.
- catalyst
- A substance that changes reaction rate without being used up overall.
- reversible reaction
- A process capable of running backwards, so substances made in the forward direction can become the starting materials again.
- dynamic equilibrium
- A closed-system state in which opposing reactions proceed equally quickly while observable amounts stay constant.
- Higher Tier: equilibrium position
- Where the equilibrium mixture lies—towards proportionally more reactants or more products.
- Higher Tier: Le Chatelier's Principle
- A rule for predicting the direction an equilibrium mixture shifts after its conditions are disturbed.
Calculate with confidence
Equations
Mean rate from product
mean rate = quantity of product formed ÷ time
Average product formation per unit time over a measured interval.
| Symbol | Meaning | Unit |
|---|---|---|
| rate | mean rate | g/s, cm³/s or HT mol/s |
| quantity | product formed | g, cm³ or HT mol |
| time | elapsed time | s |
Exam tip: Use the quantity and time units from the question.
Higher Tier: instantaneous rate
instantaneous rate = Δquantity ÷ Δtime along a tangent
The tangent gradient estimates the rate at one point on a curved graph.
| Symbol | Meaning | Unit |
|---|---|---|
| Δquantity | vertical change along tangent | axis unit |
| Δtime | horizontal change along tangent | s |
Exam tip: Use a large triangle whose points lie on the tangent, not the curve.
Relative rate for a fixed endpoint
relative rate ∝ 1 ÷ time
When every trial reaches the same visual endpoint, shorter time represents a faster rate.
| Symbol | Meaning | Unit |
|---|---|---|
| time | time to fixed colour or turbidity endpoint | s |
Exam tip: This gives a relative rate unless the measured quantity change is known.
Follow it step by step
Processes to remember
How to explain a rate factor
- State the factor and the direction in which it changes.
- Describe its effect on particle spacing, exposed surface, movement or activation energy.
- State how collision frequency or the energetic fraction changes.
- Conclude with successful collisions per second and the rate effect.
Exam tip: Only use simple doubling proportionality when other conditions are fixed, and state that it is an approximate collision-model prediction, not a universal rate law.
How to draw and interpret a tangent
- All routes: draw a tangent that just touches the curve at the requested point and follows its local direction.
- All routes: compare tangent steepness; the steeper tangent represents the faster rate.
- Higher Tier: for a numerical rate at a specified time, mark two widely separated points on the tangent.
- Higher Tier: calculate vertical change ÷ horizontal change and state the axis-derived unit; a mole quantity gives mol/s.
Exam tip: Higher Tier: do not use the origin unless the tangent actually passes through it; use two well-separated points on the tangent.
Higher Tier: how to predict an equilibrium shift
- Write the imposed concentration, temperature or pressure change.
- Identify the direction that counteracts it using the supplied equation.
- State whether relative product amount increases or decreases.
- For pressure, count gaseous coefficients only; for temperature, identify the endothermic direction.
Exam tip: A catalyst reaches equilibrium faster but does not shift its position.
See the thinking
Worked example
Higher Tier: worked example — rate from a tangent
Higher Tier: a tangent to a gas-volume curve passes through 20 s, 30 cm³ and 70 s, 80 cm³. Calculate the instantaneous rate.
- Calculate the vertical change: 80 − 30 = 50 cm³.
- Calculate the horizontal change: 70 − 20 = 50 s.
- Gradient = 50 cm³ ÷ 50 s = 1.0 cm³/s.
- Keep the unit because it states which quantity changes each second.
Answer: The instantaneous rate is 1.0 cm³/s.
The selected points lie on the tangent, so their gradient estimates the curve's rate at the tangent point. A wide triangle reduces the impact of reading uncertainty.
Protect the marks
Common mistakes
Watch out: Using the final product amount as the reaction rate.
Do this instead: Rate is a gradient or quantity change divided by time; plateau height is total product.
Watch out: Saying temperature only increases collision frequency.
Do this instead: Particles also have more kinetic energy, so a larger fraction of collisions exceeds activation energy.
Watch out: Saying a catalyst gives particles more energy.
Do this instead: It provides an alternative pathway with lower activation energy.
Watch out: Saying equilibrium means the reactions stop.
Do this instead: In sealed apparatus both directions remain active; at equilibrium neither proceeds faster.
Watch out: Saying equilibrium has equal reactant and product amounts.
Do this instead: The rates are equal; the equilibrium amounts need not be.
Watch out: Using all molecules when applying the pressure rule.
Do this instead: Higher Tier: count gaseous molecules only from the balanced equation.
Plan it like the exam
Required practicals
Investigate how concentration affects reaction rate
Combined Science and separate Chemistry
Aim: Test concentration using both a gas-volume method and a separate colour-or-turbidity method.
Method
- Develop the hypothesis that increasing reactant concentration increases rate because particles collide more often, producing more successful collisions each second.
- Before every trial, use fresh reactants and clean, rinsed and dried apparatus; for turbidity, confirm the cross is visible before mixing.
- Gas method: measure 25.0 cm³ of teacher-approved dilute hydrochloric acid at a stated concentration, such as 0.50, 1.00 or 1.50 mol/dm³, into a conical flask.
- Add 0.20 g of same-size marble chips (calcium carbonate), immediately fit the bung to a freely moving 100 cm³ gas syringe, and start timing at first contact using the same start-and-seal routine each time.
- Record the carbon dioxide volume at regular intervals until the reading plateaus; repeat for each acid concentration.
- Turbidity: make a sodium thiosulfate concentration series by replacing some solution with water, keeping total volume and viewing depth fixed. Add the same dilute hydrochloric acid volume each time and start timing.
- Stop timing when the cross is no longer visible through the same depth of mixture, viewed from the same position.
- Repeat each concentration and calculate means. For gas curves at all tiers, draw a tangent at time zero and compare its steepness.
- Higher Tier: calculate the initial tangent gradient in cm³/s.
- For turbidity, calculate relative rate = 1/mean time.
- Plot the Higher Tier gas method's numerical initial rate or the turbidity method's relative rate against concentration, then test the hypothesis.
Variables
- Independent
- concentration of one reactant
- Dependent
- gas volume over time or time to a fixed turbidity endpoint
- Controls
- temperature
- reactant volumes
- solid mass and surface area for gas method
- same apparatus and endpoint rule
Analysis: Gas at all tiers: draw a tangent at time zero on CO₂ curves and compare steepness. Higher Tier: calculate tangent gradient = Δvolume ÷ Δtime in cm³/s. Turbidity: relative rate = 1/mean time in s⁻¹, not an absolute rate. Equal plateaus need equal limiting CaCO₃ and excess acid.
Safety
- Wear eye protection and use only teacher-approved dilute acid and reactant quantities.
- Sodium thiosulfate with acid releases sulfur dioxide; use a well-ventilated laboratory, small quantities and the teacher's local risk controls.
- Keep the gas syringe pointed away from people, check that it moves freely and do not use sealed apparatus without an expansion path.
- Do not lean directly over the reaction vessel; clear spills using the teacher's instructions.
Improvements
- Use a light sensor for turbidity to replace a subjective visual endpoint.
- Use a water bath or temperature probe to keep temperature consistent.
- Check gas connections for leaks and use a data logger where available.
- Repeat conditions in a randomised order and calculate means after checking anomalies.
Try it before you move on
Quick check
Say your answer first, then open the card to check it.
48 cm³ of gas forms in 24 s. What is the mean rate?
Answer: 2.0 cm³/s
Rate = 48/24 = 2.0 cm³/s.
Why does powdered solid usually react faster than equal-mass lumps?
Answer: It has a larger surface area, causing more frequent collisions.
More solid particles are exposed to the other reactant.
How does a catalyst increase rate?
Answer: It provides an alternative pathway with lower activation energy.
A larger fraction of collisions can then be successful.
What is equal at dynamic equilibrium?
Answer: The forward and reverse reaction rates.
Reactant and product amounts stay constant but need not be equal.
Higher Tier: for N₂(g) + 3H₂(g) ⇌ 2NH₃(g), which side does increased pressure favour?
Answer: The ammonia side, because it has fewer gaseous molecules in the balanced equation.
The reactant side has four gaseous molecules and the product side has two. In general, increased pressure favours the side with fewer gaseous molecules; equal gas counts give no shift.
Good questions, clear answers
Frequently asked questions
Does a faster reaction always make more product?
No. Rate describes how quickly product forms. If the same fixed amount of one reactant is fully used up, two conditions can reach the same final product amount at different speeds.
Why does temperature have a larger effect than collision frequency alone suggests?
Particles collide more often and with a wider spread of energies, so a much larger fraction can meet or exceed activation energy.
Why is the disappearing-cross method less objective?
A person judges when the cross vanishes. A light sensor, fixed viewing conditions, repeats and consistent solution depth reduce this uncertainty.
Why must equilibrium use a closed system?
If substances escape, their concentrations keep changing for an external reason and a stable balance of forward and reverse rates may not be maintained.
Higher Tier: does a catalyst increase equilibrium yield?
No. It lowers activation energy for both directions, so equilibrium is reached faster without changing its position.
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