Chemical Change & Rate of Reaction
Revise Chemical Change & Rate of Reaction for Chemistry 0620 (O Level) — revision notes and instant AI marking.
Chemical Change & Rate of Reaction
Cambridge IGCSE Chemistry
Quick Summary
Physical vs Chemical
Physical changes don't create new substances; chemical changes do and are usually hard to reverse.
Five Key Factors
Concentration, temperature, surface area, pressure, and catalysts all affect how fast reactions go.
Collision Theory
Reactions need particle collisions with enough energy (activation energy) to break bonds and form new substances.
Rate Measurement
Track how fast reactants disappear or products form using mass loss, gas volume, or colour change.
Graph Interpretation
Reaction rate is the gradient of a product-vs-time curve. Steeper at the start, then levels off as reactants run out.
Limiting Reactant
The reactant that runs out first determines how much product you can make.
1. Physical & Chemical Changes
What's the difference?
This is one of the first concepts examiners test because students often mix them up. Here's the clearest way to think about it:
No new substances are created. The matter stays the same chemically, just looks or feels different. Easy to reverse.
Examples: Melting ice into water, dissolving salt in water, tearing paper, mixing oil and water.
New substances with different properties are formed. Bonds break and reform. Usually hard or impossible to reverse.
Examples: Burning paper to ash, rusting iron, cooking an egg, acid reacting with magnesium.
Signs of a Chemical Reaction
Examiners love asking you to describe what happens. Watch for these:
🎨 Colour Change
One of the easiest signs. For example, copper is orange-brown. When it reacts with silver nitrate (colourless), you get silver (grey) and copper nitrate (blue). Two clear colour changes = chemical reaction happened.
🌡️ Temperature Change
Exothermic: Releases heat. Combustion, neutralisation, rusting. The container gets hot.
Endothermic: Absorbs heat. Photosynthesis, dissolving ammonium chloride. The container gets cold (sometimes used in cold packs).
💨 Effervescence (Fizzing)
Gas is produced. Usually hydrogen gas when a metal reacts with acid, or carbon dioxide when an acid reacts with a carbonate.
Other signs:
Light being produced (like burning magnesium burning super bright), smell changing, pH changing.
Practice Question 1
2. Rates of Reaction — What Makes Reactions Go Faster?
This is a big exam topic. You need to know four factors that affect rate, and more importantly, why they affect rate using collision theory.
How much product is made (or reactant used up) per unit time. Usually measured in cm³/s or g/s.
Factor 1: Concentration (or Pressure for Gases)
Higher concentration = faster reaction.
Why? If there are more particles in the same volume, they bump into each other more often. More collisions = more successful collisions per second = faster reaction.
On a graph: The curve for high concentration is steeper at the start (faster initial rate) and levels off sooner (finishes quicker). Both lines reach the same height eventually (same amount of product) because you're using the same amount of reactant—just changing how fast it reacts.
Factor 2: Temperature
Higher temperature = much faster reaction.
Why? Warmer particles move faster, so they collide more frequently AND with more energy. More collisions with enough energy to break bonds = more successful collisions.
Rule of thumb: For every 10°C increase, the rate roughly doubles. That's huge. A small temperature change has a big effect.
On a graph: Same pattern as concentration—steeper at the start for the higher temperature line.
Factor 3: Surface Area
Powdered reactant reacts faster than lumps. Larger surface area = faster reaction.
Why? If a solid is powdered, more of its surface is exposed. More surface available for the other reactant to hit = more frequent collisions.
Quick maths: A 2 cm cube has surface area 24 cm². Cut it into 8 smaller cubes, and the total surface area is now 48 cm²—doubled! Same amount of solid, twice the surface exposed.
Factor 4: Catalysts
A catalyst speeds up a reaction without being used up or changed. It's the same at the end as the start.
How it works: A catalyst lowers the activation energy. Think of it like providing an easier route for the reaction. More particles have enough energy to react.
Key point: A catalyst does NOT change how much product is made—just how fast it's made. And you use very little catalyst; a tiny amount does a huge job.
On a graph: The curve with a catalyst is steeper at the start and levels off sooner, but both reach the same final amount of product.
Practice Question 2
3. Collision Theory — Why Things React (Extended Tier)
This is the theory that explains all four factors above. It's the "why" behind the "what."
For a reaction to happen: (1) Particles must collide with each other. (2) The collision must have enough energy to break bonds. (3) The minimum energy needed is called the activation energy.
Successful vs Unsuccessful Collisions
Successful Collision
Particles hit head-on with enough energy (≥ activation energy). Bonds break, atoms rearrange, and new products form.
Unsuccessful Collision
Particles either don't hit hard enough, or they bounce at the wrong angle. No reaction happens; they just bounce apart unchanged.
The fact: Most collisions are unsuccessful. Only a tiny fraction have enough energy. That's why temperature is so powerful—a small increase in heat means WAY more collisions have enough energy.
Why Each Factor Affects Rate (The Deep Explanation)
Concentration → More particles per unit volume → More collisions per second
If you double the concentration, you double the number of particles. Double the collisions. Double the successful collisions. Rate doubles.
Temperature → Particles move faster → More collisions AND higher energy per collision
This is the big one. Higher KE means two things: (1) Faster particles = more frequent collisions. (2) Higher energy per collision = more surpass the activation energy threshold. Both effects combined = huge rate increase.
Surface Area → More solid surface exposed → More contact points between reactants
A larger surface area means more area available for reactant particles to collide with the solid. More collision sites = more collisions per second.
Catalyst → Lowers activation energy → Higher proportion of collisions have enough energy
A catalyst doesn't change the energy of colliding particles. Instead, it provides an alternative reaction pathway that requires less energy. Suddenly, many more collisions are successful, even without heating.
4. Investigating the Rate of a Reaction
Examiners often describe an experiment and ask you to explain what's happening or predict results. Know these three classic setups.
Method 1: Measuring the Volume of Gas Produced
Classic Experiment: Sodium thiosulfate + dilute hydrochloric acid → sulfur precipitate (yellow solid) + other products
Setup:
- Measure 50 cm³ of sodium thiosulfate solution
- Mix with 5 cm³ of dilute HCl in a flask
- The solution becomes cloudy (yellow sulfur forms)
- Start a stopwatch and look down from above
- Stop when you can no longer see a cross drawn on paper under the flask
- Record the time. The faster the cloudiness appears, the faster the rate.
Why this works: The rate of sulfur formation = the rate of reaction. Slower reaction = longer time to become opaque. Faster reaction = shorter time.
Method 2: Measuring Gas Volume (Gas Syringe or Inverted Measuring Cylinder)
Classic Experiment: Magnesium + dilute HCl → magnesium chloride + hydrogen gas
Setup: Use a gas syringe or upside-down measuring cylinder filled with water to collect the hydrogen gas produced.
Why this works: You measure gas volume at regular time intervals and plot a graph. Volume increases as the reaction proceeds. The steeper the curve, the faster the reaction.
To change variables: Use different concentrations of acid (concentration factor), different sizes of magnesium pieces (surface area factor), different temperatures (temperature factor).
Method 3: Measuring Mass Loss
Setup: Place a flask with a reaction on a balance. Gas escapes through cotton wool (which lets gas out but prevents solids from splattering). Measure the mass at regular intervals.
Why this works: As gas escapes, the flask gets lighter. The rate of mass loss tells you the reaction rate.
Real Experiment: Testing the Effect of Concentration
Experiment design: Same as the sodium thiosulfate method above, but use different concentrations of sodium thiosulfate (e.g., dilute different amounts with water).
Results: Higher concentration → faster reaction → cross disappears sooner.
Why? More solute particles per unit volume → more frequent collisions → more successful collisions.
Practice Question 3
5. Interpreting Data and Rate Calculations
What Graphs Tell You
Shape of the Curve (Product vs Time)
At the start: Steep gradient = fast reaction (lots of reactant, lots of successful collisions).
In the middle: Gradient becomes less steep = reaction slows down (some reactant used up, fewer collisions).
At the end: Horizontal line (gradient = 0) = reaction stopped (one reactant ran out, no more collisions possible).
Comparing Two Reactions
Curve 1 is steeper at the start than Curve 2? Curve 1 is faster (higher initial rate).
Curve 1 reaches horizontal sooner? Curve 1 finishes faster.
Curve 1 is higher at the end? Curve 1 made more product (maybe a different amount of reactant, or a limiting reactant difference).
Calculating Rate at a Specific Time
To find the rate at a specific point, you draw a tangent to the curve (a straight line that just touches the curve at that point).
Example: A graph shows volume of CO₂ (y-axis) vs time in seconds (x-axis). At 20 seconds, the tangent line goes from (10 seconds, 5 cm³) to (30 seconds, 29 cm³).
Rate = (29 − 5) / (30 − 10) = 24 / 20 = 1.2 cm³/s
Limiting Reactants and Product Amount
The limiting reactant is the one that runs out first and determines how much product you can make.
On a graph: If you double the concentration of Reactant A but keep Reactant B the same, you might get twice as much product (if B becomes the new limiting reactant). Or you might get the same product (if A was always in excess). The shape of the curve and how high it reaches tells you what's limiting.
Practice Question 4
What to Memorise
Concepts Checklist
Tick off each concept as you master it. Use this to focus your revision.
Exam Tips — Common Mistakes & What Examiners Look For
Rate is how FAST the reaction goes. Amount is how MUCH product is made. A catalyst makes reactions faster but doesn't make more product. Concentration makes reactions faster but doesn't make more product (total is determined by the limiting reactant).
Just saying "higher concentration = faster rate" gets you a fact. Examiners want the mechanism: "Higher concentration means more particles per unit volume, so collisions happen more frequently, so more successful collisions per second, so faster rate." Use collision theory.
Wrong. A catalyst is not used up. Same mass at the end as the start. It provides an alternative pathway but isn't part of the final products.
When two curves are compared: (1) Steeper initial gradient = faster rate. (2) Reaching horizontal sooner = reaction finished faster. (3) Different final heights = different total product (different amounts of limiting reactant or different reactants).
When explaining temperature, don't just say "particles move faster." Say: "Particles have more kinetic energy. This increases both collision frequency AND the proportion of collisions with energy ≥ activation energy."
• Explain how [factor] affects rate using collision theory.
• Compare two graphs and explain the differences.
• Design an experiment to test [variable].
• Draw a tangent and calculate rate.
• Describe what you would observe in this reaction (colour, temperature, gas, etc.).
Always link back to collision theory. More collisions? Explain why. More successful collisions? Explain which part of the mechanism changed (number, frequency, energy, activation energy). Don't just state facts—explain mechanisms.
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