# Rate of Reaction and Factors
Some reactions are over in an instant (like explosions), while others take millions of years (like rusting). Understanding what controls the speed of a chemical reaction is crucial in industry and everyday life. In this guide, we'll learn about measuring rates, collision theory, and the factors that speed up or slow down reactions.
1. What Is Rate of Reaction?
The rate of reaction is a measure of how quickly reactants are used up or products are formed.
The "quantity" can be measured as:
- Mass (in grams) — measured using a balance
- Volume of gas (in cm³) — measured using a gas syringe or collection over water
- Colour change — using a colorimeter or observation
- Time for a precipitate to obscure a mark — disappearing cross experiment
2. Collision Theory
For a reaction to occur, particles must:
- Collide with each other
- Collide with enough energy (≥ activation energy)
- Collide in the correct orientation
Collisions that meet all these conditions are called successful collisions.
Rate of reaction depends on the frequency of successful collisions.
Anything that increases the frequency or energy of collisions will increase the rate.
3. Factors Affecting Rate
3.1 Temperature
Effect: Increasing temperature increases the rate.
Why (collision theory):
- Particles have more kinetic energy → move faster
- Particles collide more frequently
- A greater proportion of collisions have energy ≥ activation energy
- More successful collisions per second
Exam key point: State BOTH increased frequency AND increased energy of collisions.
As a rough rule: a 10°C increase approximately doubles the rate.
3.2 Concentration (for solutions)
Effect: Increasing concentration increases the rate.
Why:
- More particles in the same volume
- Particles are closer together
- Collisions are more frequent
- More successful collisions per second
3.3 Pressure (for gases)
Effect: Increasing pressure increases the rate.
Why:
- Particles are pushed closer together into a smaller volume
- Collisions are more frequent
- Same effect as increasing concentration
3.4 Surface Area (for solids)
Effect: Increasing surface area increases the rate.
Why:
- Smaller pieces (e.g. powder) expose more surface to the other reactant
- More particles are available for collision at the surface
- Collisions are more frequent
| Form | Surface Area | Rate |
|---|---|---|
| Large lump | Small | Slow |
| Small pieces | Medium | Medium |
| Powder | Large | Fast |
3.5 Catalysts
Effect: A catalyst increases the rate without being used up.
Why:
- A catalyst provides an alternative reaction pathway with a lower activation energy
- More particles have energy ≥ the (lower) activation energy
- More successful collisions per second
Key facts about catalysts:
- Not used up in the reaction (can be reused)
- Not included in the balanced equation
- Only a small amount is needed
- Specific to particular reactions
- Do not change the products or the amount of product — only speed things up
Common catalysts:
| Catalyst | Reaction |
|---|---|
| Iron | Haber process (N₂ + H₂ → NH₃) |
| Manganese dioxide | Decomposition of H₂O₂ |
| Vanadium pentoxide | Contact process (making H₂SO₄) |
| Enzymes | Biological reactions |
4. Measuring Rate of Reaction
Method 1: Gas Syringe
Collect the gas produced and measure its volume over time.
- Plot volume (y-axis) vs time (x-axis)
- Steeper gradient = faster rate
Method 2: Mass Loss
Place the reaction on a balance and record mass over time as gas escapes.
- Plot mass loss vs time
- Steeper initial gradient = faster rate
Method 3: Disappearing Cross (Sodium Thiosulfate + HCl)
- Sulfur precipitate makes the solution go cloudy
- Time how long until a cross drawn on paper beneath the flask disappears
- Shorter time = faster rate
- Rate =
5. Interpreting Rate Graphs
Volume of Gas vs Time
- Steep curve = fast rate
- Gentle curve = slow rate
- Flat line = reaction finished
- The total volume at the end tells you about the amount of product
Comparing experiments:
- Higher temperature: steeper curve, reaches same final volume faster
- More concentrated: steeper curve, same final volume reached faster
- Catalyst: steeper curve, same final volume reached faster
- More reactant: same initial rate, but higher final volume
Worked Example: Explaining the Effect of Temperature
Question: Explain why increasing the temperature increases the rate of reaction between magnesium and hydrochloric acid.
At higher temperatures, particles have more kinetic energy and move faster. This means collisions between Mg and HCl particles occur more frequently. Also, a greater proportion of collisions have energy equal to or greater than the activation energy, so more collisions are successful. Therefore, the rate increases.
Worked Example: Surface Area
Question: A student reacted calcium carbonate with hydrochloric acid. In experiment A, large marble chips were used. In experiment B, powdered calcium carbonate was used. Same mass was used in both. Compare the results.
Experiment B (powder) would react faster because the powder has a larger surface area exposed to the acid. More acid particles can collide with the calcium carbonate at any time, increasing the frequency of collisions. Both experiments produce the same total volume of CO₂ (same mass of reactant).
Worked Example: Calculating Rate
Question: 50 cm³ of gas was collected in 20 seconds. Calculate the mean rate of reaction.
7. Practice Questions
- Define "rate of reaction."
- State and explain four factors that affect rate of reaction.
- A student added zinc to acid at 25°C and at 50°C. Sketch a graph showing gas volume vs time for both experiments.
- Explain, using collision theory, why a catalyst increases the rate.
- In the disappearing cross experiment, a student measured times of 60 s, 30 s, and 15 s for three temperatures. Calculate the rate () for each and describe the trend.
Want to check your answers and get step-by-step solutions?
8. Common Misconceptions
| Misconception | Reality |
|---|---|
| Catalysts are used up | Catalysts are not consumed — they can be reused |
| Higher temperature means more collisions only | Higher temperature also means collisions have more energy |
| Powdering a solid gives more product | It gives the same total product, just faster |
| Catalysts increase the energy of particles | Catalysts lower the activation energy, providing an alternative pathway |
9. Exam Tips
- Always explain rate changes using collision theory
- For temperature: mention BOTH more frequent collisions AND more energetic collisions
- For surface area, concentration, pressure: focus on frequency of collisions
- For catalysts: mention alternative pathway and lower activation energy
- When describing graphs, use the word gradient (steeper gradient = faster rate)
Summary
- Rate = amount of product formed ÷ time
- Collision theory: reactions need successful collisions (sufficient energy + correct orientation)
- Rate increases with: higher temperature, higher concentration, higher pressure, larger surface area, catalysts
- Catalysts lower activation energy via an alternative pathway
- Rate can be measured by gas volume, mass loss, or time for a visible change