Influencing Factors Revision Notes for HSC SSCE Chemistry

Influencing Factors

Understanding what affects the speed of chemical reactions is crucial in chemistry. Various factors can make reactions go faster or slower, and knowing these helps us control reactions in the lab and in industry.

What affects reaction rates?

Several key factors influence how quickly a chemical reaction proceeds. These factors work differently depending on whether the reaction occurs in one phase (homogeneous) or between different phases (heterogeneous).

Concentration

When we increase the concentration of a reactant, we generally increase the rate of the reaction. This happens because there are more reactant particles available to collide and react with each other.

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This explains why reactions typically slow down as they proceed - the reactants are being used up, so their concentration decreases over time.

Temperature

Temperature has a powerful effect on reaction rates. For most reactions, the rate increases as temperature rises. In some cases, a temperature increase of just $10°\text{C}$ can double the rate of reaction. This is because particles move faster at higher temperatures, leading to more frequent and more energetic collisions.

Catalysts

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A catalyst is a substance that increases the rate of a reaction without undergoing permanent chemical change itself. This is a remarkable property - the catalyst helps the reaction go faster but emerges unchanged at the end.

For example, consider the reaction between hydrogen peroxide and acidified iodide solution:

$\text{H}_2\text{O}_2(\text{aq}) + 2\text{I}^-(\text{aq}) + 2\text{H}^+(\text{aq}) \rightarrow \text{I}_2(\text{aq}) + 2\text{H}_2\text{O}(\text{l})$

This reaction proceeds much faster when sodium molybdate solution ( $\text{Na}_2\text{MoO}_4$ ) is added as a catalyst. Similarly, cobalt(II) chloride can catalyse the oxidation of hydrogen tartrate by hydrogen peroxide. The catalyst speeds up the reaction but doesn't appear in the stoichiometric equation because it's not consumed.

Homogeneous reactions

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A homogeneous reaction is one where all reactants exist in the same phase. The reactants form a homogeneous mixture, either as:

A mixture of gases
A solution (all dissolved substances)

For example, the reaction between nitrogen oxide and oxygen gases ( $\text{NO} + \text{O}_2$ ) occurs entirely in the gas phase. Similarly, the hydrogen peroxide and iodide reaction shown above occurs entirely in solution.

Factors affecting homogeneous reactions

Three main factors influence the rate of homogeneous reactions:

Concentration of reactants (whether in solution or gas phase)
Nature and concentration of any catalyst present
Temperature

These three factors work together to determine how quickly a homogeneous reaction proceeds.

Heterogeneous reactions

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A heterogeneous reaction occurs at the interface between two different phases. This might be between:

A gas and a solid
A solution and a solid
Other combinations of different phases

The key point is that reaction happens at the boundary where the two phases meet.

Common examples of heterogeneous reactions

Several familiar laboratory reactions are heterogeneous:

Zinc metal with hydrochloric acid: The solid zinc reacts with the acid solution to produce hydrogen gas and zinc chloride in solution
Decomposition of hydrogen peroxide: When hydrogen peroxide solution decomposes on the surface of solid manganese dioxide, it forms oxygen gas and water
Marble chips with hydrochloric acid: Solid calcium carbonate (limestone) reacts with acid solution to produce carbon dioxide gas and calcium chloride solution

Additional factors for heterogeneous reactions

Heterogeneous reactions are affected by all the same factors as homogeneous reactions (concentration, temperature, and catalysts), but two additional factors become important:

Surface area (state of division)

The surface area of the solid has a major impact on reaction rate. Breaking a solid into smaller pieces dramatically increases its total surface area. This means more solid is in contact with the other phase (gas or solution), so more reaction can occur at once.

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Worked Example: Effect of Surface Area

Consider calcium carbonate reacting with hydrochloric acid:

Crushed limestone reacts much faster than large lumps. Even though the total mass is the same, the crushed material has a much larger surface area exposed to the acid. More calcium carbonate particles are in contact with the acid solution, so more collisions can occur, and carbon dioxide is produced more rapidly.

Stirring

Stirring (or agitation) of the mixture increases the rate of heterogeneous reactions. This is because stirring:

Brings fresh reactant from the solution to the solid surface
Prevents solid particles from settling
Replaces reactant that has been used up at the surface

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Without stirring, used-up reactant near the solid surface must be replaced by slow diffusion. Stirring speeds up this replacement, keeping the reaction going faster.

For instance, if you stir crushed limestone in acid, you keep the limestone dispersed throughout the solution rather than letting it settle at the bottom, which maintains a higher reaction rate.

Summary of heterogeneous reaction factors

For heterogeneous reactions, five factors affect the rate:

Concentration of reactants
Temperature
Nature and concentration of catalyst
State of division of the solid (surface area)
Rate of stirring (or whether stirring occurs at all)

Industrial applications

Heterogeneous reactions are extremely important in industrial chemistry. Several major industrial processes rely on heterogeneous reactions, demonstrating the practical significance of understanding these reaction factors.

Ammonia synthesis

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Industrial Example: The Haber Process

The Haber process combines nitrogen and hydrogen gases using an iron catalyst to produce ammonia:

$\text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3$

Ammonia is essential for making fertilisers, nitric acid, and explosives. The iron catalyst provides a solid surface where the gas-phase reaction can occur more rapidly.

Hydrogenation of vegetable oils

Liquid vegetable oils can be converted to semi-solid margarines by adding hydrogen in the presence of finely divided nickel metal as a catalyst. The nickel surface allows the hydrogen to add to double bonds in the oil molecules, making them more solid.

Cracking in oil refining

Large hydrocarbon molecules from crude oil are broken down (cracked) into smaller, more useful molecules for petrol. This reaction occurs in the gas phase but requires solid catalysts to proceed at a practical rate and temperature.

Catalytic converters

Modern car exhausts contain catalytic converters that remove harmful pollutants. These devices use solid catalysts (often platinum or palladium) to speed up reactions that convert:

Carbon monoxide to carbon dioxide
Nitrogen oxides to nitrogen
Unburnt fuel to carbon dioxide and water

The exhaust gases pass over the solid catalyst surface, where these important environmental reactions occur rapidly.

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Key Points to Remember:

Concentration: Higher concentration of reactants generally means faster reactions
Temperature: Increasing temperature speeds up most reactions; sometimes just $10°\text{C}$ can double the rate
Catalysts: These special substances speed up reactions without being permanently changed themselves
Homogeneous reactions: Occur in one phase and are affected by concentration, temperature, and catalysts
Heterogeneous reactions: Occur at interfaces between phases and are also affected by surface area and stirring - breaking solids into smaller pieces and stirring mixtures both increase reaction rates

Influencing Factors (HSC SSCE Chemistry): Revision Notes