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:

1. Particles must collide with sufficient energy (activation energy)

2. 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

1. 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

2. 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

3. 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

4. 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