# OCR AS/A Level Chemistry A

# Enthalpy changes A level

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## Introduction

### Overview

Delivery guides are designed to represent a body of knowledge about teaching a particular topic and contain:

- Curriculum Content: A clear outline of the content covered by the delivery guide
- Thinking Conceptually: Expert guidance and activities on the key concepts involved, common difficulties learners may have, approaches to teaching that can help learners understand these concepts and how this topic links conceptually to other areas of the subject
- Thinking Contextually: A range of guidance and suggested teaching activities using a variety of themes so that different activities can be selected which best suit particular classes, learning styles or teaching approaches.

## Curriculum Content

### Overview

**Content (from A Level)**

**5.2.1 Lattice enthalpy**

(a) explanation of the term *lattice enthalpy* (formation of 1 mol of ionic lattice from gaseous ions, \(\displaystyle \Delta_{\rm{LE}}H\)) and use as a measure of the strength of ionic bonding in a giant ionic lattice (see also 2.2.2 b–c)

(b) use of the lattice enthalpy of a simple ionic solid (i.e. NaC*l*, MgC*l*_{2}) and relevant energy terms for:

(c) explanation and use of the terms:

*enthalpy change of solution*(dissolving of 1 mol of solute, \(\displaystyle \Delta_{\rm{sol}}H\))

*enthalpy change of hydration*(dissolving of 1 mol of gaseous ions in water, \(\displaystyle \Delta_{\rm{hyd}}H\))

(d) use of the enthalpy change of solution of a simple ionic solid (i.e. NaC*l*, MgC*l*_{2}) and relevant energy terms (*enthalpy change of hydration* and *lattice enthalpy*) for:

(e) qualitative explanation of the effect of ionic charge and ionic radius on the exothermic value of a lattice enthalpy and enthalpy change of hydration.

**5.2.2 Enthalpy and entropy**

(a) explanation that entropy is a measure of the dispersal of energy in a system which is greater, the more disordered a system

(b) explanation of the difference in magnitude of the entropy of a system:

(c) calculation of the entropy change of a system, \(\displaystyle \Delta S\), and related quantities for a reaction given the entropies of the reactants and products

(d) explanation that the feasibility of a process depends upon the entropy change and temperature in the system, \(\displaystyle T\Delta S\), and the enthalpy change of the system, \(\displaystyle \Delta H\)

(e) explanation, and related calculations, of the free energy change, \(\displaystyle \Delta G\), as: \(\displaystyle \Delta G = \Delta H - T\Delta S\) (the Gibbs’ equation) and that a process is feasible when \(\displaystyle \Delta G\) has a negative value

(f) the limitations of predictions made by \(\displaystyle \Delta G\) about feasibility, in terms of kinetics.

## Thinking Conceptually

### Overview

This guide focuses on two key areas of A level Chemistry: Lattice Enthalpy and Entropy.

Lattice Enthalpy builds on the concepts learned in Hess cycles and relies heavily on mathematical understanding. The individual reactions involved – such as atomisation, ionisation and electron affinity – can be hard to visualise and their details easy to mix up and forget. This can lead to a large variety of errors within questions and makes it a very polarising topic.

Entropy and free energy, when boiled down, are fairly straightforward calculations and questions can be attempted with little prior knowledge. However, the ideas behind entropy can be difficult to visualise and explain and there are many pitfalls which catch out even the brightest of learners.

**Approaches to teaching the content**

To explain and build Born–Haber cycles you must first have a clear understanding of each of the individual standard reactions and how to define them. These are: formation, atomisation, 1^{st} and 2^{nd} ionisation, electron affinities and lattice enthalpy. This can be a very dry lesson if just written down and discussed. A good method is to include a competitive edge: have pairs write down your choice of equation and element and score points by spotting each other’s mistakes.

Born–Haber cycles contain endothermic reactions like atomisation and exothermic ones like electron affinity. The negative values can cause confusion; drawing a cycle to scale (see Learner Resource 1) can help learners visualise the cycle before they start calculating it.

Entropy can be easily introduced with some high entropy reactions, for example the reaction of a carbonate and acid. Learners grasp that solids are more ordered than gases and the associated entropy change can then be calculated from standard values. Gibbs energy is a lot harder to visualise and is better explained using the equation, practically the enthalpy and entropy of a system like the one mentioned above can be measured, compared and shown to be feasible.

**Common misconceptions or difficulties learners may have**

A common mistake in Born–Haber cycles is that learners forget to multiply or divide standard values. For example, when completing a Born–Haber cycle for MgC*l*_{2}, the enthalpy change of atomisation (\(\displaystyle \Delta_aH\)) for chlorine is

½C*l*_{2}(g) → C*l*(g) –437 kJ mol^{–1}

But many learners forget that this applies to half a mole of chlorine and would not double it for a whole molecule.

Another common mistake in a cycle would be that learners forget that \(\displaystyle \Delta_{\rm{EA1}} H^\ominus\) for oxygen is exothermic (arrow points down) but \(\displaystyle \Delta_{\rm{EA2}} H^\ominus\) is endothermic (arrow points up). This causes them to add rather than subtract the value and calculate the lattice enthalpy incorrectly.

State symbols are a regular tripping point for learners. Forgetting these in equations relating to enthalpy changes can cost marks, so make a big deal of them.

A common error when calculating free energy using \(\displaystyle \Delta G = \Delta H - T\Delta S\) is that learners forget that the units for entropy (\(\displaystyle \Delta S\)) are J K^{–1} mol^{–1} whereas enthalpy (\(\displaystyle \Delta H\)) is given in kJ mol^{–1}. These means learners must divide their entropy by 1000 before using it in the Gibbs equation.

Another common mistake is to use the temperature in °C rather than converting to K, so using 25°C rather than 298 K.

### Bonding and enthalpy quiz (Suberberg chemistry)

### Born–Haber cycle calculations (ScienceGeek.net)

### Born-Haber cycles (IBChem.com)

### Born–Haber cycle to scale

### Building cycles

### Entropy: Embrace the chaos!

### Endothermic solid–solid reaction (Nuffield Foundation)

## Thinking Contextually

### Overview

The challenge in teaching enthalpy and entropy through context is that learners find these ideas very conceptual. ‘Everyday’ understanding about both terms is confused, particularly in the case of entropy.

With entropy, care needs to be taken not to reinforce the idea that it is the ‘tendency to disorder’, which leads to only a partial understanding of the second law of thermodynamics. Entropy laws are embedded in random number generators and in statistical methods behind the gambling industry; the house always wins because even seemingly random systems follow the rules of entropy so that over time their behaviour can be predicted.

### Entropy demonstration

As learners arrive to the lesson, be silently building a ‘house of cards’ on the front bench. As they watch and start to wonder what you are doing, ‘accidentally’ knock it down and start again (or ask one of them to rebuild it).

Use this as a context for talking about the number of ways the cards can be arranged, and that to keep them in one, chosen order, an input of energy is needed. Ask for other everyday examples of this (photos of ‘everyday’ entropy could be used as a backdrop, e.g. tidy and untidy rooms/desks, ordered groups of people and crowds).

### Entropy dice

This is a ‘game’ type activity: Ask learners to roll two dice 100 times and make a tally chart of the scores generated (from 2 to 12). This should give a normal distribution. Ask students to put forward ideas of why this is (it is because there is only one way that 2 or 12 can be scored. The ‘highest’ tally should fall with the score with the greatest number of ways (there are most ways of scoring 7). This gives a context for why a ‘double 6’ is considered a lucky dice roll.

The analogy to entropy is that the dice are more likely to fall into a state that can be achieved in many ways (a ‘random’ state). This is the same for matter, which has a tendency towards states with many different arrangements.

The experiment is explained in the video (good for teacher background).

## Acknowledgements

### Overview

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