Kinetics

Question 1

What is Activation Energy (Ea)?
(2)

Give the basic definition and explain.

Answer 1

• Activation energy refers to the minimum energy required to break bonds and start a reaction.
• Reactions with low Ea occur more easily.

Question 2

Why do gas molecules have different speeds, and how is this shown?

Answer 2

Gas molecules have different kinetic energies, so they move at different speeds.
This variation is shown by the Maxwell–Boltzmann distribution, which displays the spread of molecular energies:
- Some molecules move slowly (low KE)
- Some move very fast (high KE)
- Most have moderate KE
- Only molecules with KE ≥ activation energy can react on collision

Question 3

What are the defining features of the Maxwell–Boltzmann distribution?

Answer 3

Curve starts at origin → no molecules have zero KE
Rises sharply to a peak (most probable energy, Ep)
Then tails off gradually, never touching x-axis
Total area under curve = total number of molecules
Average energy > most probable energy
Shows distribution, not identical energies

Question 4

How does increasing temperature affect the Maxwell–Boltzmann distribution?

Answer 4

Curve shifts right (higher energies)
Peak lowers and broadens
More molecules have energy ≥ Ea
Total area remains constant (same number of molecules)
Effect: Higher frequency of effective collisions → dramatic increase in reaction rate.

Question 5

What is a catalyst?

Answer 5

A catalyst increases rate by providing an alternative pathway with lower activation energy.
It is not consumed in the reaction and is chemically unchanged at the end.
Lower Ea → more molecules have sufficient energy → more effective collisions.

Question 6

How do catalysts increase reaction rate according to kinetic theory?

Answer 6

Catalysts are specific to certain reactions
They interact with reactant particles to form an alternative pathway
Lower activation energy (Ea)
On a Maxwell–Boltzmann curve, Ea line shifts left → more molecules exceed Ea
Increases frequency of effective collisions

Question 7

What is reaction rate and how is it expressed

Answer 7

Rate = change in concentration (or amount) of reactant/product per unit time.
Examples:
- Rate = Δ[reactant]/Δt
- Rate = Δ[product]/Δt
Units include mol dm⁻³ s⁻¹, g s⁻¹, cm³ s⁻¹.

Question 8

How is loss of mass used to measure reaction rate?

Answer 8

• Used when gas escapes from a reaction
• Place reaction on balance
• As gas forms, mass decreases
• Measure mass at time intervals
• Must use fume cupboard if toxic gases form

Question 9

How is volume of gas measured?

Answer 9

Use a gas syringe:
- Collect gas over time
- Measure volume at set intervals
- Very accurate for fast reactions

Question 10

Describe the precipitate method for measuring rate.

Answer 10

Used when a solid forms (e.g., sulfur):
- Place flask on top of a black cross
- Mix reactants
- Time how long until cross is obscured
- Shorter time → faster reaction.
- Keep same cross and observer for fair test.

Question 11

Explain what happens in the thiosulfate–acid reaction and how rate is measured.

Answer 11

Reaction produces sulfur:
Na₂S₂O₃ + 2HCl → 2NaCl + SO₂ + S + H₂O
Sulfur precipitate clouds solution → cross disappears.
Method:
Warm solutions to control temperature
Mix and time disappearance
Higher temperature → faster rate.

Question 12

Describe the process for calculating average rate from a graph.

Answer 12

Draw line of best fit
Pick two points
Draw triangle (rise/run)
Calculate gradient = Δy / Δx
Include correct units (e.g., cm³ min⁻¹)

Question 13

What two conditions must be met for a successful collision?

Answer 13

1. Correct orientation
2. Enough kinetic energy (≥ Ea)
   Most collisions fail because:
   - Insufficient energy
   - Wrong orientation
   - Therefore most collisions are ineffective.

Question 14

Why does increasing surface area increase rate?

Answer 14

More exposed particles → more frequent collisions → more effective collisions → faster rate.
Example: powdered solid > lump.

Question 15

Why does increasing concentration increase rate?

Answer 15

Particles closer together → more frequent collisions → increased number of effective collisions.

Question 16

Why does increasing pressure increase rate?

Answer 16

Affects gases only.
Higher pressure → particles forced closer → more frequent collisions → faster rate.

Question 17

Why does increasing temperature increase reaction rate?

Answer 17

• Particles move faster → more collisions
• More have KE ≥ Ea → more effective collisions
• Maxwell–Boltzmann curve shifts right
• Rate increases dramatically.

Question 18

How does a catalyst appear on a Maxwell–Boltzmann diagram?

Answer 18

• The curve shape stays the same
• Ea (catalysed) line is drawn to the left of the original
• Larger shaded area beyond Ea → more molecules can react