Rates of Reaction

Subject: Chemistry
Topic: 7
Cambridge Code: 0620 / 0971 / 5070

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Reaction Rate

Reaction rate - Change in concentration per unit time

$\text{Rate} = \frac{\text{Change in concentration}}{\text{Time}} = \frac{Δc}{Δt}$

Units: mol/dm³/s or mol/(L·s)

Measuring Reaction Rate

Methods:

1. Measure concentration over time
2. Measure volume of gas produced
3. Measure mass loss
4. Measure time for color change

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Factors Affecting Reaction Rate

1. Concentration

Increase concentration → Rate increases

• More particles per unit volume
• Greater chance of collisions
• Proportional relationship (for some reactions)

Rate ∝ [Concentration]

2. Temperature

Increase temperature → Rate increases significantly

• Particles move faster (kinetic energy ↑)
• More collisions per unit time
• Higher proportion of collisions have sufficient energy
• Roughly doubles every 10°C (rule of thumb)

Temperature most effective factor

3. Pressure (Gases)

Increase pressure → Rate increases

• Decreases volume
• Molecules closer together
• More frequent collisions

4. Splitting into Smaller Pieces (Surface Area)

Increase surface area → Rate increases

• Larger exposed surface
• More collisions with other substances
• Particularly important for solid reactants

Powder reacts faster than lump

5. Catalysts

Catalyst: Substance that speeds up reaction without being consumed

• Mechanism: Lower activation energy
• Properties: Unchanged at end, reusable
• Examples: MnO₂ (for H₂O₂), Ni (for hydrogenation)

Important: Catalysts increase rate but don't shift equilibrium position

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Collision Theory

Collision Theory - Explains how reactions occur

Requirements for Reaction

1. Particles must collide
2. Collision must have sufficient energy (≥ activation energy)
3. Collision must be in correct orientation

Activation Energy

Activation energy (Ea) - Minimum energy needed for reaction to occur

• Energy barrier to overcome
• Different for different reactions
• Catalyst lowers it

Effective Collisions

Effective collision: Sufficient energy AND correct orientation

$\text{Reaction rate} = \text{(Number of collisions)} \times \text{(Fraction with} \geq E_a) \times \text{(Fraction correctly oriented)}$

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Energy Diagrams

Exothermic Reaction

Energy
^
|     Reactants
|       ___
|      /   \___Products
|     /
|____/______________ Reaction progress
  E_a

• Products lower energy than reactants
• ΔH negative
• Activation energy needed to start

Endothermic Reaction

Energy
^
|            ___Products
|      ___  /
|     /   \/
|    /  Reactants
|___/______________ Reaction progress
    E_a

• Products higher energy than reactants
• ΔH positive
• Activation energy needed to start

Effect of Catalyst

• Lowers activation energy
• Doesn't change reactants/products energies
• Lowers Ea, increases rate
• Shifts equilibrium position (equilibrium favors products/reactants equally)

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Reaction Order

Order - Power to which concentration raised in rate equation

$\text{Rate} = k[A]^m[B]^n$

where k = rate constant, m and n = orders

First Order

Rate ∝ [A]

• Half-life constant
• Examples: Decomposition, radioactive decay

Second Order

Rate ∝ [A]² or Rate ∝ [A][B]

• Half-life decreases
• Example: 2NO + O₂ → 2NO₂

Zero Order

Rate independent of concentration

• Rate = k (constant)
• Unusual but exists

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Initial Rate Method

Finding reaction orders experimentally

Procedure

1. Vary one concentration, keep others constant
2. Measure initial rate (slope at t=0)
3. Compare rates for different concentrations

Example

Experiment	[A]	[B]	Rate
1	0.1	0.1	0.01
2	0.2	0.1	0.02
3	0.1	0.2	0.04

• Changing [A]: doubles [A] → doubles rate = first order in A
• Changing [B]: doubles [B] → quadruples rate = second order in B

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Equilibrium vs Rate

Important distinction:

• Catalyst speeds BOTH forward and reverse reactions equally
• Equilibrium position unchanged
• Rate at equilibrium increases from both directions
• Equilibrium reached faster

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Key Points

1. Rate: Change in concentration per time
2. Concentration, temperature, surface area, pressure affect rate
3. Catalyst lowers activation energy
4. Collision theory: Activation energy + orientation required
5. Reaction order determined experimentally
6. Temperature most effective variable factor
7. Catalyst speeds up reaction without being consumed

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Practice Questions

1. Plot concentration vs time graphs
2. Calculate reaction rates
3. Explain factor effects on rate
4. Identify reaction orders
5. Predict rate changes
6. Draw energy diagrams
7. Determine rate equations

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Revision Tips

• Know all factors affecting rate
• Collision theory explanation
• Activation energy concept
• Catalyst action (doesn't change ΔH)
• Reaction order determination
• Rate calculations
• Graph interpretation
• Temperature most effective