RATES OF REACTION
LEARNING OBJECTIVES
Understand what reaction rate is and how to measure it
Identify factors that affect the rate of a reaction
Explain how each factor changes reaction rate using collision theory
Understand the role of catalysts in reactions
Interpret rate vs time graphs
Predict how changing conditions will affect reaction rate
KEY DEFINITIONS
Reaction rate :
The speed at which a chemical reaction occurs (how quickly reactants are used up or products are formed)
Collision theory :
Theory stating that reactions occur when particles collide with sufficient energy (activation energy)
Activation energy :
The minimum energy required for a collision between particles to result in a chemical reaction
Catalyst :
A substance that speeds up a reaction without being used up itself
Effective collision :
A collision between particles with enough energy and correct orientation to produce a reaction
WHAT IS REACTION RATE
Measures how fast a reaction is happening
Can be measured by observing how quickly reactants decrease or products form
Expressed as: amount of substance / time
Examples of fast reactions: combustion, explosions, neutralization
Examples of slow reactions: rusting, decomposition of waste
COLLISION THEORY
For a reaction to occur, particles must collide
Not all collisions result in reaction - only effective collisions do
For effective collision:
Particles must collide with sufficient energy (activation energy)
Particles must be oriented correctly
Higher temperature = more energetic particles = more effective collisions
Higher concentration = more particles in same volume = more frequent collisions
Larger surface area = more particles exposed = more frequent collisions
FACTORS AFFECTING REACTION RATE
TEMPERATURE
Higher temperature increases reaction rate
Why: Particles move faster, collide more frequently and with more energy
More collisions have sufficient energy to react (effective collisions increase)
Even a small increase in temperature causes significant increase in rate
Example: Magnesium reacts slowly with cold dilute acid, much faster with hot acid
CONCENTRATION (or pressure for gases)
Higher concentration increases reaction rate
Why: More particles in the same volume = more frequent collisions
More collision opportunities lead to more effective collisions
Example: Zinc reacts slowly with dilute HCl, much faster with concentrated HCl
Reaction: Zn + 2HCl → ZnCl₂ + H₂
SURFACE AREA
Larger surface area increases reaction rate
Why: More particles exposed means more chance for collisions
Applies when solid reacts with liquid or gas
Example: Powdered calcium carbonate reacts much faster than large lumps with acid
Example: Steel wool burns brightly, iron block does not
CATALYST
A catalyst increases reaction rate without being used up
How: Provides an alternative reaction pathway with lower activation energy
More particles have sufficient energy to react
Same catalyst can be used multiple times
Not consumed in the reaction
Different catalysts for different reactions
Examples: MnO₂ for H₂O₂ decomposition, enzymes in living cells
COMPARISON OF FACTORS AFFECTING RATE
HOW TO MEASURE REACTION RATE
Method 1 : Monitor product formation
Measure volume of gas produced over time
Record mass of solid produced
Observe time taken for colour change
Method 2 : Monitor reactant consumption
Measure volume/mass of reactant remaining
Observe temperature change
Track concentration decrease
Rate calculation :
Rate = amount of product formed / time
Rate = amount of reactant used / time
Units: g/s, cm³/s, mol/s, mol/dm³/s
INTERPRETING RATE vs TIME GRAPHS
Typical graph features :
Time on x-axis, product formed (or reactant left) on y-axis
Curve starts steep (fast rate at start)
Curve gradually flattens (rate slows down)
Curve levels off (reaction stops)
Why rate decreases :
Reactants are used up (lower concentration)
Fewer particles available for collision
Rate of collision decreases over time
Steeper curve = faster reaction
Comparing reactions:
Higher temperature : curve is steeper initially
Higher concentration : curve is steeper initially
With catalyst : curve rises more steeply, reaches end faster
Both reach same final point (same amount of product, same time)
PRACTICAL EXAMPLES
Example 1 : Zinc + Hydrochloric acid
Reaction: Zn + 2HCl → ZnCl₂ + H₂
Product: H₂ gas (measured by gas collection)
Rate increases with: higher temperature, higher HCl concentration, smaller Zn pieces
Catalyst: Not typically used for this reaction
Example 2 : Calcium carbonate + Hydrochloric acid
Reaction: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
Product: CO₂ gas
Rate affected by: particle size (powder vs lump), temperature, acid concentration
Powdered CaCO₃ reacts much faster than chips
Example 3 : Hydrogen peroxide decomposition
Reaction: 2H₂O → 2H₂O + O
Slow reaction without catalyst
With MnO₂ catalyst: reaction becomes very fast (bubbles vigorously)
MnO₂ is not used up, can be recovered and reused
