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Factors affecting reaction rate

Concentration influences the rate of reaction in solutions.

At a lower concentration (left), there are fewer collisions while at a higher concentration (right), there are more collisions.
At a lower concentration (left), there are fewer collisions while at a higher concentration (right), there are more collisions.

As the concentration (measured in moles of a chemical per unit volume) of reactant molecules increases, collisions between reactant molecules become more frequent.

This increases the likelihood of having many effective collisions. The rate of reaction thus increases with concentration.

Acid rain corrodes statues more quickly than normal rainwater because acid rain has higher concentration of $$\ce{H^+}$$ ions.

Pressure influences the rate of reaction in gases.

For gases, pressure is directly related to concentration.

As the pressure increases, collisions between molecules become more frequent. The rate of reaction thus increases with pressure.

When ammonia is produced from hydrogen and nitrogen gas, the pressure is set to 200 atmospheres because the reaction is too slow at lower pressures.

The surface area influences the rate of reaction.

The surface area of solids determines how quickly they will react with liquids or gases.

Increasing the surface area increases the rate of reaction.

A 1 kilogram block of iron takes a long time to convert entirely to rust when submerged in water.

If the same 1 kilogram block of iron was ground into iron powder, it would rust quickly when in water.

Increasing the surface area increases the reaction rate because collisions of reactants occur on the surface of solids.

The rate of collisions is higher when there is more surface area on which molecules can collide.

Iron filings (left) reacts faster than a block of iron (right).
Iron filings (left) reacts faster than a block of iron (right).

The rate of a chemical reaction depends on the activation energy of the reaction. Activation energy is the energy that the reactant molecules must obtain before a reaction can occur.

Reactant molecules must possess the activation energy in order for effective collisions to occur.

The lower the activation energy, the faster the reaction will proceed. This is because effective collisions occur more easily, so there will be more of them.

Hydrogen gas and nitrogen gas react to form ammonia ($$\ce{NH3}$$) very slowly on their own because the activation energy is high.

Few molecules have the energy needed to undergo effective collisions.

$$E_A$$ refers to the activation energy. The graph plots the path of the reaction from reactants (on the left) to products (on the right).
$$E_A$$ refers to the activation energy. The graph plots the path of the reaction from reactants (on the left) to products (on the right).

Temperature affects the rate of a chemical reaction.

At higher temperatures, more molecules possess the activation energy needed for effective collisions.

Molecules at higher temperatures have more kinetic energy (they move faster) than molecules at lower temperatures. These molecules therefore collide more frequently and more forcefully.

Effective collisions occur more frequently at higher temperatures. Increasing the temperature therefore increases the rate of reaction. This is true for both endothermic and exothermic reactions.

The reaction of magnesium with cold acid to form hydrogen gas is slow. Only a few gas bubbles are formed each minute.

Reacting magnesium with hot acid causes intense bubbling, indicating rapid formation of hydrogen gas and thus a fast reaction.

Food spoilage is a result of chemical reactions. Keeping food frozen or chilled slows down the rate of these reactions, so that food can be kept for longer periods of time.
Food spoilage is a result of chemical reactions. Keeping food frozen or chilled slows down the rate of these reactions, so that food can be kept for longer periods of time.

A catalyst is a substance that increases the rate of a reaction and that is neither produced nor consumed in the reaction.

A catalyst may change form during the reaction. However, at the end of the reaction it is chemically unchanged when compared to itself at the start of the reaction.

The blue reaction profile curve indicates a reaction without a catalyst. The green curve represents a catalysed reaction. $$E_A$$ refers to the activation energy.
The blue reaction profile curve indicates a reaction without a catalyst. The green curve represents a catalysed reaction. $$E_A$$ refers to the activation energy.

Catalysts offer an alternative pathway for a reaction (as depicted on an energy profile) that has a lower activation energy than the normal pathway (without a catalyst).

Catalysts speed up reactions by lowering the activation energy barrier.

Effective collisions between molecules become easier to obtain and thus more frequent when a catalyst is used.

The complete decomposition of hydrogen peroxide ($$\ce{H2O2}$$) into oxygen and water can take up to a few days. Adding a catalyst (manganese oxide in this case) allows it to complete in seconds.

Catalysts are used in industries to accelerate reactions that would otherwise proceed too slowly.

In the reduction of an alkene to produce an alkane, a nickel catalyst is used to speed up the process. This reduction process is used to convert oils into solid fats.

Catalysts also exist naturally. Enzymes are biological catalysts that accelerate important reactions in living organisms. They are usually proteins.

Enzymes are very specific in action and can only catalyse a particular class of reactions.

Salivary amylase catalyses the breakdown of starch in the mouth. It cannot be used to break down fats or proteins.

Enzymes can only catalyse select reactions due to the shape of the enzyme (shown in blue). Only a particular group of reactants can fit with the enzyme perfectly.
Enzymes can only catalyse select reactions due to the shape of the enzyme (shown in blue). Only a particular group of reactants can fit with the enzyme perfectly.
Factor Measures taken to increase reaction rate Measures taken to decrease reaction rate
Concentration (mainly for solutions) Increase Decrease
Pressure (for gases) Increase Decrease
Surface area Increase Decrease
Temperature Increase Decrease
Activation energy Decrease (usually with catalyst) Increase