1 Factors affecting rates

Collision theory The Activation Energy is defined as the minimum energy which particles need to collide to start a reaction 3.2.2. Reactions Rates reactants products Activation Energy: EA ∆H Energy Reactions can only occur when collisions take place between particles having sufficient energy. The energy is usually needed to break the relevant bonds in one or either of the reactant molecules. This minimum energy is called the Activation Energy Effect of Increasing Concentration and Increasing Pressure At higher concentrations(and pressures) there are more particles per unit volume and so the particles collide with a greater frequency and there will be a higher frequency of effective collisions. Measuring Reaction Rates The rate of reaction is defined as the change in concentration of a substance in unit time Its usual unit is mol dm-3s -1 When a graph of concentration of reactant is plotted vs time, the gradient of the curve is the rate of reaction. The initial rate is the rate at the start of the reaction where it is fastest Reaction rates can be calculated from graphs of concentration of reactants or products In the experiment between sodium thiosulphate and hydrochloric acid we usually measure reaction rate as 1/time where the time is the time taken for a cross placed underneath the reaction mixture to disappear due to the cloudiness of the Sulphur . Na2S2O3 + 2HCl
2NaCl + SO2 + S + H2O This is an approximation for rate of reaction as it does not include concentration. We can use this because we can assume the amount of Sulphur produced is fixed and constant. Initial rate = gradient of tangent time concentration Note: If a question mentions a doubling of concentration/rate then make sure you mention double the number of particles per unit volume and double the frequency of effective collisions. Comparing rate curves Need to calculate/ compare initial moles of reactants to distinguish between different finishing volumes. e.g. the amount of product is proportional to the moles of reactant Different volumes of the same initial concentrations will have the same initial rate (if other conditions are the same) but will end at different amounts The higher the concentration/ temperature/ surface area the faster the rate (steeper the gradient) Amount of product e.g. Volume of gas Effect of Catalysts Definition: Catalysts increase reaction rates without getting used up. Explanation: They do this by providing an alternative route or mechanism with a lower activation energy so more molecules have energy above activation energy Comparison of the activation energies for an uncatalysed reaction and for the same reaction with a catalyst present. reactants Activation Energy: uncatalysed ∆H Progress of Reaction EA catalysed products Catalysts speed up the rate of reaction. This means that the use of a catalyst may mean lower temperatures and pressures can be used. This can save energy costs as there is reduced energy demand for providing high temperature and less electrical pumping costs for producing pressure. This can mean fewer CO2 emissions from burning of fossil fuels, Benefits of Catalysts Catalysts can enable different reactions to be used, with better atom economy and with reduced waste, or fewer undesired products or less use of hazardous solvents and reactants. Catalysts are often enzymes, generating very specific products, and operating effectively close to room temperatures and pressures Heterogeneous catalysts are usually solids whereas the reactants are gaseous or in solution. The reaction occurs at the surface of the catalyst. Heterogeneous catalysis When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species. Homogeneous catalysis A heterogeneous catalyst is in a different phase from the reactants A homogeneous catalyst is in the same phase as the reactants