F325: Module 1 Specification

5.1.1 How Fast? Rate graphs and orders Concentration–time can be plotted from continuous measurements taken during the course of a reaction (continuous monitoring). Initial rates require separate experiments using different concentrations of one of the reactants. Clock reactions are an approximation of this method. Candidates should be able to: (a) explain and use the terms: rate of reaction, order, rate constant, half-life, ratedetermining step; (b) deduce, from a concentration–time graph, the rate of a reaction and the half-life of a firstorder reaction; (c) state that the half-life of a first-order reaction is independent of the concentration; (d) deduce, from a rate–concentration graph, the order (0, 1 or 2) with respect to a reactant; (e) determine, using the initial rates method, the order (0, 1 or 2) with respect to a reactant; Rate equations; rate constants Integrated forms of rate equations are not required. (f) deduce, from orders, a rate equation of the form: rate = k[A]m[B]n, for which m and n are 0, 1 or 2; (g) calculate the rate constant, k, from a rate equation; (h) explain qualitatively the effect of temperature change on a rate constant and hence the rate of a reaction; Use of rate equations to predict and propose a reaction mechanism. (i) for a multi-step reaction: (i) propose a rate equation that is consistent with the rate-determining step, (ii) propose steps in a reaction mechanism from the rate equation and the balanced equation for the overall reaction. 5.1.2 How Far? Equilibrium Candidates should be able to: (a) calculate, given appropriate data, the concentration or quantities present at equilibrium; (b) deduce, for homogeneous reactions, expressions for the equilibrium constant Kc; (c) calculate the values of the equilibrium constant Kc including determination of units; (d) explain the effect of changing temperature on the value of Kc for exothermic and endothermic reactions; (e) state that the value of Kc is unaffected by changes in concentration or pressure or by the presence of a catalyst. 5.1.3 Acids, Bases and Buffers Brønsted–Lowry acids and bases Candidates should be able to: (a) describe an acid as a species that can donate a proton and a base as a species that can accept a proton (b) illustrate, using ionic equations, the role of H+ in the reactions of acids with metals, carbonates, bases and alkalis (c) describe and use the term conjugate acid– base pairs; Strong and weak acids (d) explain qualitatively, in terms of dissociation, the differences between strong and weak acids; (e) explain that the acid dissociation constant, Ka, shows the extent of acid dissociation; (f) deduce, for weak acids, expressions for Ka and pKa; pH and [H+(aq)] For a weak acid HA, assume: [H+(aq)] = [A(aq)]; equilibrium [HA] = undissociated [HA]. (g) define pH as pH = –log[H+]; [H+] = 10–pH; (h) state and use the expression for the ionic product of water, Kw; (i) calculate pH from [H+(aq)] and [H+(aq)] from pH for: (i) strong monobasic acids, (ii) weak monobasic acids, (iii) strong bases, using Kw; (j) calculate Ka for a weak acid, given appropriate data; Buffers: action, uses and calculations The details of a basic buffer system are not required. The H2CO3/HCO3– buffer is present in blood plasma, maintaining a pH between 7.35 and 7.45. (k) describe a buffer solution as a system that minimises pH changes on addition of small amounts of an acid or a base; (l) state that a buffer solution can be made from a weak acid and a salt of the weak acid, eg CH3COOH/CH3COONa; (m) explain the role of the conjugate acid–base pair in an acid buffer solution, eg CH3COOH/CH3COO–, in the control of pH; (n) calculate the pH of a buffer solution, from the Ka value of a weak acid and the equilibrium concentrations of the conjugate acid–base pair; (o) explain the role of carbonic acid– hydrogencarbonate as a buffer in the control of blood pH; Neutralisation (p) for acid–base titration pH curves for strong and weak acids and bases: (i) interpret, or sketch, their shapes, (ii) explain the choice of suitable indicators for acid–base titrations, given the pH range of the indicator; (q) define and use the term enthalpy change of neutralisation and calculate enthalpy changes from appropriate experimental results

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