reaction kinetics

• Rate equation for the reaction A + B → C: rate = k[A]​m​[B]​n​ (m is the order of the reaction with respect to A, n is the order of reaction with respect to B)
• Order of reaction: the power to which the concentration of a reactant is raised in the rate equation (could be 0, 1, or 2)
• Rate constant: the constant (k) linking rate of reaction and the concentration of reactants, changes with temperature and use of a catalyst but remains the same if only concentrations of reactants are changed
• Half-life: the time taken for the concentration of the reactants to halve
• Rate-determining step: the stage with the slowest rate in a multi-step reaction
• Intermediate: a substance formed between the reactants and the product during a chemical reaction
• If a reaction is 0 order with respect to reactant P, the rate is unaffected by changing the concentration of P
• If a reaction is first order with respect to reactant Q, doubling the concentration of Q will double the rate (rate ∝ [Q] or rate = k[Q])
• The overall order of a reaction is equal to the sum of all the orders of the reactants
• Concentration-Time Graphs: to determine the order of a reaction, observe the shape of the concentration-time graph and compare it to known graphs
• Rate-Concentration Graphs: the order can be deduced using a rate-concentration graph by observing its shape
• Initial Rate Experiments: orders of reaction can be determined using experimental data relating to initial rates
• Half-life of First Order Reactions: the half-life of a first order reaction is constant
• Reaction Mechanisms: describe the steps involved in making and breaking bonds during a reaction
• When constructing a reaction mechanism, the powers in the rate equation indicate the number of molecules of each substance involved in the slowest step
• Any intermediates generated in the slowest step must be reactants in another step as they are not present in the balanced overall equation
• Example mechanism: Nitrogen dioxide and carbon monoxide react to form nitrogen monoxide and carbon dioxide
• Rate equation for the reaction: rate = k[NO2]2
• From the rate equation, the reaction is zero order with respect to CO(g) and second order with respect to NO2(g)
• In the rate-determining step of the example mechanism, 2 molecules of NO2 are involved
• Activation energy is the minimum amount of energy required for a reaction to occur
• Boltzmann distribution for gases shows the relative energies of particles, with only particles above the activation energy being able to react when they collide
• Increasing temperature leads to more collisions and a greater proportion of particles having energy above the activation energy, resulting in more frequent successful collisions and an increased reaction rate
• As temperature increases, the rate constant increases, leading to an increase in the rate of reaction
• A catalyst speeds up the rate of a reaction without being chemically changed at the end
• In the Haber process, iron is used as a heterogeneous catalyst to produce ammonia from nitrogen and hydrogen
• Catalytic converters use platinum, palladium, and rhodium as heterogeneous catalysts to convert carbon monoxide and nitrogen oxides from car exhausts into less harmful compounds
• Sulfur dioxide is oxidized in the atmosphere to form sulfur trioxide, catalyzed by the homogeneous catalyst nitrogen dioxide
• Iron ions can be used as a homogeneous catalyst for the reaction between I- and S2O82- ions