Question 1

How do I determine the order of reaction from experimental data?

Accepted Answer

To determine the order, you can use the initial rates method or concentration-time graphs. For initial rates, compare how the rate changes when you change the concentration of one reactant while keeping others constant. If doubling concentration doubles the rate, it's first order; if it quadruples, it's second order; if no change, it's zero order. For concentration-time graphs, plot [A] vs time (zero order gives straight line), ln[A] vs time (first order gives straight line), or 1/[A] vs time (second order gives straight line). The one that is linear indicates the order.

Question 2

What is the rate-determining step and how do I find it?

Accepted Answer

The rate-determining step is the slowest step in a multi-step reaction mechanism. It determines the overall rate and the rate equation. To find it, look at the rate equation: the reactants in the rate-determining step must appear in the rate equation. For example, if rate = k[A][B], then the slow step likely involves one molecule of A and one of B. Intermediates (species produced in one step and consumed in another) may appear in the slow step but not in the overall equation. You can deduce the mechanism by ensuring the sum of steps gives the overall equation and that the slow step matches the rate law.

Question 3

How does a catalyst increase reaction rate?

Accepted Answer

A catalyst provides an alternative reaction pathway with a lower activation energy. This means a greater proportion of collisions have energy equal to or greater than the activation energy, so more successful collisions occur per unit time. The catalyst is not consumed; it may form an intermediate with reactants but is regenerated. For example, in the decomposition of hydrogen peroxide, manganese dioxide catalyst lowers the activation energy, speeding up the reaction without being used up.

Question 4

What is the Arrhenius equation and how do I use it?

Accepted Answer

The Arrhenius equation is k = Ae^(-Ea/RT), where k is the rate constant, A is the pre-exponential factor (frequency factor), Ea is activation energy, R is the gas constant (8.314 J mol⁻¹ K⁻¹), and T is temperature in Kelvin. It shows how k varies with temperature and activation energy. To use it, you can take natural logs: ln k = ln A - Ea/(RT). A plot of ln k against 1/T gives a straight line with slope -Ea/R and intercept ln A. This allows you to calculate Ea from experimental data at different temperatures.

Question 5

Why does increasing temperature increase reaction rate?

Accepted Answer

Increasing temperature increases the kinetic energy of particles, so they move faster and collide more frequently. More importantly, it increases the fraction of particles with energy equal to or greater than the activation energy (Ea). According to the Boltzmann distribution, at higher temperatures, the peak shifts to the right and the curve flattens, meaning more molecules have high energy. This exponential increase in the number of successful collisions leads to a significant rate increase, often doubling for every 10°C rise.

Question 6

What is the difference between homogeneous and heterogeneous catalysts?

Accepted Answer

Homogeneous catalysts are in the same phase as the reactants (e.g., all in solution). They often work by forming an intermediate complex. For example, acid catalysts in ester hydrolysis. Heterogeneous catalysts are in a different phase (e.g., solid catalyst with gaseous reactants). They work by adsorbing reactants onto their surface, weakening bonds and lowering activation energy. Examples include iron in the Haber process and platinum in catalytic converters. Heterogeneous catalysts are easier to separate but can be poisoned by impurities.

Kinetics

Topic Overview

Key Concepts

What You Need to Demonstrate

Marking Points

Examiner Tips

Common Mistakes

Frequently Asked Questions

Before You Start

Key Terminology

Likely Command Words

Ready to test yourself?

Related Topics in AQA A-Level Chemistry

Acids and bases (A-level only)

Alcohols

Aldehydes and ketones (A-level only)

Alkanes