2.3 - Module 3: Energy


This module provides candidates with a knowledge and understanding of chemical reasoning that underpins the study of physical chemistry.

2.3.1 Enthalpy Changes

* enthalpy changes of reaction, combustion and formation;

* bond enthalpies;

* Hess' law and enthalpy cycles.

2.3.2 Rates and Equilibrium

* collision theory, the Boltzmann distribution and catalysis;

* a qualitative study of reaction rates;

* dynamic equilibrium and le Chatelier's principle.

Links AS Unit F321: Atoms, Bonds and Groups

* 1.1.2 Moles and Equations

* 1.3.2 Group 2 (acid reactions with metals, carbonates and bases)

AS Unit F322: Chains, Energy and Resources

* 2.1.2 Alkanes (combustion of fuels)

* 2.2.1 Alcohols (combustion of alcohols)

2.3.1 Enthalpy Changes

Context and exemplification Assessable learning outcomes

Enthalpy changes: ΔH of reaction, formation and combustion

Candidates should be able to:

(a) explain that some chemical reactions are accompanied by enthalpy changes that can be exothermic (ΔH, negative) or endothermic (ΔH, positive);

(b) describe the importance of oxidation as an exothermic process in the combustion of fuels and the oxidation of carbohydrates such as glucose in respiration;

(c) describe that endothermic processes require an input of heat energy, eg the thermal decomposition of calcium carbonate;

(d) construct a simple enthalpy profile diagram for a reaction to show the difference in the enthalpy of the reactants compared with that of the products;

(e) explain qualitatively, using enthalpy profile diagrams, the term activation energy;

* Standard conditions can be considered as 100 kPa and a stated temperature, 298 K.

(f) define and use the terms:

(i) standard conditions,

(ii) enthalpy change of reaction,

(iii) enthalpy change of formation,

(iv) enthalpy change of combustion;

(g) calculate enthalpy changes from appropriate experimental results directly, including use of the relationship: energy change = mcΔT; Bond enthalpies

(h) explain exothermic and endothermic reactions in terms of enthalpy changes associated with the breaking and making of chemical bonds;

(i) define and use the term average bond enthalpy (ΔH positive; bond breaking of one mole of bonds);

(j) calculate an enthalpy change of reaction from average bond enthalpies;

Hess' law and enthalpy cycles

* Unfamiliar enthalpy cycles will be provided.

(k) use Hess' law to construct enthalpy cycles and carry out calculations to determine:

(i) an enthalpy change of reaction from enthalpy changes of combustion,

(ii) an enthalpy change of reaction from enthalpy changes of formation,

(iii) an enthalpy change of reaction from an unfamiliar enthalpy cycle.

2.3.2 Rates and Equilibrium

Context and exemplification Assessable learning outcomes

Simple collision theory

Candidates should be able to:

(a) describe qualitatively, in terms of collision theory, the effect of concentration changes on the rate of a reaction;

(b) explain why an increase in the pressure of a gas, increasing its concentration, may increase the rate of a reaction involving gases; Catalysts How Science Works 6a: * Benefits of catalysis in terms of possible lower production costs but also implications for their disposal (toxicity). * Details of processes are not required.

(c) state that a catalyst speeds up a reaction without being consumed by the overall reaction;

(d) explain that catalysts:

  • (i) affect the conditions that are needed, often requiring lower temperatures and reducing energy demand and CO2 emissions from burning of fossil fuels,
  • (ii) enable different reactions to be used, with better atom economy and with reduced waste,
  • (iii) are often enzymes, generating very specific products, and operating effectively close to room temperatures and pressures,
  • (iv) have great economic importance, eg iron in ammonia production, Ziegler- Natta catalyst in poly(ethene) production, platinum/palladium/rhodium in catalytic converters (see also 2.4.1.i); (e) explain, using enthalpy profile diagrams, how the presence of a catalyst allows a reaction to proceed via a different route with a lower activation energy, giving rise to an increased reaction rate; The Boltzmann distribution

How Science Works 1:

* The Boltzmann distribution as a theoretical model arising from kinetic theory.

(f) explain qualitatively the Boltzmann distribution and its relationship with activation energy;

(g) describe qualitatively, using the Boltzmann distribution, the effect of temperature changes on the proportion of molecules exceeding the activation energy and hence the reaction rate;

(h) interpret the catalytic behaviour in (e), in terms of the Boltzmann distribution; Dynamic equilibrium and le Chatelier's principle

(i) explain that a dynamic equilibrium exists when the rate of the forward reaction is equal to the rate of the reverse reaction;

(j) state le Chatelier's principle;

(k) apply le Chatelier's principle to deduce qualitatively (from appropriate information) the effect of a change in temperature, concentration or pressure, on a homogeneous system in equilibrium;

(l) explain, from given data, the importance in the chemical industry of a compromise between chemical equilibrium and reaction rate.

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.

  • * Direct enthalpy changes of reaction for simple reactions: Zn + CuSO4 (exo); NaHCO3 + citric acid (endo); NaOH + HCl (exo).
  • * Enthalpy change of combustion of alcohols.
  • * Indirect enthalpy change of reaction: 2KHCO3 ¡÷ K2CO3 + H2O + CO2 indirectly using HCl.
  • * Rate graphs for gas products, eg CaCO3 + HCl; Mg + HCl
  • * Changing equilibrium position with heat: [Cu(H2O)6]2+ . CuCl4 2-
  • * Changing equilibrium position with concentration: Fe3+ and SCN-

2.3.1 - Enthalpy Changes

2.3.2 - Rates and Equilibrium