C5.2 - Controlling reactions

What is rate of reaction?

How quickly reactants are converted into products

Formula for rate of reaction

rate = amount of reactant used or product formed ÷ time

Rate triangle

[DRAW: rate on top, amount + time at bottom]

Units for rate

g/s, cm³/s, mol/s depending on measurement

How can rate be measured?

By change in mass, volume of gas, or formation of precipitate

Measuring rate by mass change

Reaction releasing gas → mass decreases over time

Mass change diagram

[DRAW: conical flask on balance losing mass as gas escapes]

Measuring rate by gas volume

Collect gas in syringe or measuring cylinder over water

Gas volume diagram

[DRAW: flask → tube → gas syringe]

Measuring rate by precipitate formation

Reaction forms solid → obscures mark under flask

Precipitate example

Na₂S₂O₃ + 2HCl → 2NaCl + SO₂ + S + H₂O

Cross method diagram

[DRAW: flask with cross visible → cross disappears as precipitate forms]

What do rate graphs show?

How amount of reactant or product changes over time

gradient of tangent = rate at that point

Reactant graph shape

Downward curve (reactant decreases)

Product graph shape

Upward curve (product increases)

Gradient meaning

Steeper gradient = faster rate

When graph levels off

Reaction finished

Mean rate formula

mean rate = change in quantity ÷ change in time

Example mean rate

48 cm³ gas in 40 s → 48 ÷ 40 = 1.2 cm³/s

Rate at a particular point

Draw tangent → gradient = rate

Tangent diagram

[DRAW: curve with tangent line at chosen point]

Comparing rates

Steeper curve = faster reaction

Rate from reaction time

rate = 1 ÷ reaction time

Factors affecting rate

Concentration/pressure, temperature, surface area, catalysts

Increasing concentration or pressure

More particles per volume → more frequent collisions

Concentration/pressure graph

[DRAW: high vs low concentration curves]

Increasing temperature

Particles have more kinetic energy → more frequent + energetic collisions

At higher temperature → more particles have energy ≥ activation energy

Temperature graph

[DRAW: high vs low temperature curves]

Increasing surface area

greater surface area → more collisions per second

Surface area diagram

[DRAW: large cube vs small cubes]

Collision theory

Reactions occur when particles collide with enough energy to overcome activation energy

Activation energy definition

Minimum energy needed for reaction to occur

Successful vs unsuccessful collisions

Successful = energy ≥ activation energy

Collision diagrams

[DRAW: particles colliding with/without enough energy]

Effect of concentration/pressure

More particles → higher collision frequency → faster rate

Effect of temperature

Higher kinetic energy → more collisions with sufficient energy

Effect of surface area

More exposed particles → more collisions

Effect of catalyst

Provides alternative pathway with lower activation energy

Catalyst definition

substance that speeds up rate of reaction but does not get used up and can be reused

Catalyst energy profile diagram

[DRAW: energy vs reaction coordinate; catalysed curve lower peak]

Catalyst industrial importance

Allows lower temperature + pressure → saves energy + cost

Catalyst equilibrium effect

Does not change position of equilibrium

What is an enzyme?

Biological catalyst speeding up reactions in living organisms

How enzymes work

Substrate fits into enzyme’s active site forming enzyme–substrate complex

Enzyme specificity

Each enzyme catalyses one type of reaction

High temperature → enzyme denatures (active site changes shape)

Factors affecting enzyme activity

Temperature, pH, substrate concentration