A2 unit F325. Module 1: Rates, Equilibrium and pH

Specification

This module provides candidates with a quantitative study of physical chemistry. The material covered in this module links many areas of chemistry and explains many chemical phenomena. For example, the qualitative treatment of reaction rates and equilibria encountered at AS is developed within a quantitative and graphical context.

  • Practical Skills are assessed using OCR set tasks.

The practical work suggested below may be carried out as part of skill development. Centres are not required to carry out all of these experiments:

  • * Continuous monitoring of a reaction for 'quantity' against time graph: CaCO3/HCl: monitoring gas or mass loss; H+/I2/(CH3)2CO: monitoring [I2] using a colorimeter.
  • * Clock reactions for determination of orders and rate constants by initial rates and rate concentration graphs: Iodine clocks: I2/S2O8 2- or H+/H2O2/I2 with S2O3 2- (or vitamin C); Thiosulfate clock: HCl/S2O3 2-.
  • * The effect of temperature on reaction rates (clock reactions above are suitable).
  • * Determination of Kc for ethanoic acid/ethyl ethanoate equilibrium.
  • * Use of pH meters to: measure pHs of strong and weak acids; investigate buffer solutions.
  • * Generating an acid-base titration curve with a data logger, www.chemit.co.uk/uploads/java/ .

5.1.1 How Fast?

* orders, rate equations, rate constants;

* continuous and initial rate methods;

* rate-determining step.

5.1.2 How Far?

* equilibrium concentrations;

* the equilibrium constant, Kc.

5.1.3 Acids, Bases and Buffers

* acid-base equilibria;

* strength of acids including Ka;

* pH determination, titration curves and buffer solutions.

AS Unit F321: Atoms, Bonds and Groups 1.1.3 Acids; 1. 3.2 Group 2 (acid reactions with metals, carbonates and bases).

AS Unit F322: Chains, Energy and Resources 2.3.2 Rates and Equilibrium.

5.1.1 How Fast?

Context and exemplification Assessable learning outcomes

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;

Rate-determining step

How Science Works 1, 7a:

* 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?

Context and exemplification Assessable learning outcomes

Equilibrium

* Candidates will not be required to solve quadratic equations.

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

Context and exemplification Assessable learning outcomes

Bronsted-Lowry acids and base

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 (see also unit F321: 1.1.3.a,h);

(b) illustrate, using ionic equations, the role of H+ in the reactions of acids with metals, carbonates, bases and alkalis (see also unit F321: 1.1.3.g; 1.1.4.f);

(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 (see also unit F322: 2.3.1.f,g).

Practical Skills are assessed using OCR set tasks.

The practical work suggested below may be carried out as part of skill development. Centres are not required to carry out all of these experiments:

* Continuous monitoring of a reaction for 'quantity' against time graph: CaCO3/HCl: monitoring gas or mass loss; H+/I2/(CH3)2CO: monitoring [I2] using a colorimeter.

* Clock reactions for determination of orders and rate constants by initial rates and rate concentration graphs: Iodine clocks: I2/S2O8 2- or H+/H2O2/I2 with S2O3 2- (or vitamin C); Thiosulfate clock: HCl/S2O3 2-.

* The effect of temperature on reaction rates (clock reactions above are suitable).

* Determination of Kc for ethanoic acid/ethyl ethanoate equilibrium.

* Use of pH meters to: measure pHs of strong and weak acids; investigate buffer solutions.

* Generating an acid-base titration curve with a data logger, www.chemit.co.uk/uploads/java/ .


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