# OCR AS/A Level Chemistry B (Salters)

# Energetics

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## Introduction

### Overview

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

- Content: A clear outline of the content covered by the delivery guide;
- Thinking Conceptually: Expert guidance on the key concepts involved, common difficulties students may have, approaches to teaching that can help students understand these concepts and how this topic links conceptually to other areas of the subject;
- Thinking Contextually: A range of 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)**

*Learners should be able to demonstrate and apply their knowledge and understanding of:*

**DF(d)** the terms: *exothermic, endothermic, standard conditions, (standard) enthalpy change of reaction**(\(\displaystyle \Delta_rH\)), (standard) enthalpy change of combustion (\(\displaystyle \Delta_cH\)), (standard) enthalpy change of formation (\(\displaystyle \Delta_fH\)), (standard) enthalpy change of neutralisation (\(\displaystyle \Delta_{neut}H\))*

**DF(e)** the term *average bond enthalpy* and the relation of bond enthalpy to the length and strength of a bond; bond-breaking as an endothermic process and bond-making as exothermic; the relation of these processes to the overall enthalpy change for a reaction

**DF(f)** techniques and procedures for measuring the energy transferred when reactions occur in solution (or solids reacting with solution) or when flammable liquids burn; the calculation of enthalpy changes from experimental results

**DF(g)** the determination of enthalpy changes of reaction from enthalpy cycles and enthalpy level diagrams based on Hess’ law

**O(a)** the factors determining the relative solubility of a solute in aqueous and non-aqueous solvents

**O(b)** the terms *hydrated ions, enthalpy change of solution, lattice enthalpy* and *enthalpy change of hydration of ions,* and:

**(i)** the solution of an ionic solid in terms of enthalpy cycles and enthalpy level diagrams involving these terms

**(ii)** use of these enthalpy cycles to perform calculations

**(iii)** techniques and procedures for measuring the energy transferred in experiments involving enthalpy changes in solution

**O(c)** the dependence of the lattice enthalpy of an ionic compound and the enthalpy change of hydration of ions on the charge density of ions

**O(d)** qualitative entropy changes (of the system); entropy as a measure of the number of ways that molecules and their associated energy quanta can be arranged

**O(e)** qualitative predictions of the \(\displaystyle \Delta_{\rm{sys}}S\) for a reaction in terms of:

**(i)** the differences in magnitude of the entropy of a solid, a liquid and a gas

**(ii)** the difference in number of particles of gaseous reactants and products

**O(f)** the expressions: \(\displaystyle \Delta_{\rm{tot}}S=\Delta_{\rm{sys}}S+\Delta_{\rm{surr}}S\) and \(\displaystyle \Delta_{\rm{surr}}S=-\Delta{H}/{T}\); calculations using these expressions; the relation of the feasibility of a reaction to the sign of \(\displaystyle \Delta_{\rm{tot}}S\)

**O(g)** calculation of \(\displaystyle \Delta_{\rm{sys}}S\) for a reaction given the entropies of reactants and products

## Thinking Conceptually

### Overview

**APPROACHES TO TEACHING THE CONTENT**

*Enthalpy*

Ideas about enthalpy are first introduced in *Elements of life (EL)* when students begin to engage with the ionisation enthalpy of atoms. At this point, the focus is on understanding the periodic trends in attraction between the nucleus and the outer electrons, and an in-depth understanding of enthalpy is not required. At this point, it will suffice to equate the term enthalpy with energy.

The formal study of enthalpy is covered in *Developing fuels (DF)*. As well as the formal definition of enthalpy as the energy transferred to or from the surroundings as a reaction takes place, there are several aspects of enthalpy that the learning covers:

- specific enthalpy changes (\(\displaystyle \Delta_rH\), \(\displaystyle \Delta_cH\), \(\displaystyle \Delta_fH\), \(\displaystyle \Delta_{neut}H\))
- average bond enthalpy and its use to estimate enthalpy changes in reactions
- practical determination of enthalpy changes
- using Hess’s Law.

It is important from the outset that learners define and learn the specific enthalpy changes so that they know the amounts and conditions that the definitions are concerned with. It is a teaching challenge to encourage learners to learn these definitions thoroughly and find a ‘fun’ platform for them to do so. Some suggestions for activities are included.

Using molecular model kits to model organic models can be used to help learners engage with the idea that energy is needed to break bonds, and also to support them to count and identify each bond that is broken and made, perhaps in the context of a combustion reaction. Data tables of bond enthalpies can be used as a data source for learners to discuss patterns, or do further research to find patterns and trends in bond enthalpies and bond lengths. Encouraging learners to engage directly with data tables also reinforces their use. Any activity which increases use of data tables supports learners’ independence; the skill of looking up data will also support them in examinations, where a *Data Sheet* is provided. (Note though that the Chemistry B (Salters) *Data Sheet* does not include bond enthalpy data; any such data required in examinations will be provided in relevant questions.)

Engaging with practical determination of enthalpy values is important in making the learning ‘concrete’ for learners, particularly as concepts relating to enthalpy are abstract. Measuring temperature changes and seeing enthalpy in action is useful in making it clear that a study of enthalpy has a direct practical application. The section on ‘Thinking Contextually’ has suggestions for contexts to include to increase the relevance of practical work in this area. Learners may consider why practical measurements of enthalpy changes intrinsically involve errors due to the difficulty of isolating the system from its surroundings.

Hess’s Law calculations are challenging, even for more able learners. One suggested approach is to break the task down so that learners begin by using templates before starting to construct diagrams independently. This part of the learning may build on the practical experience of the learners, with them considering that it is useful to both measure enthalpy changes practically, and also be able to predict their theoretical values. Learners may consider why theoretical calculated values are limited by the use of average enthalpy data and by the use of standard states for measurement.

In the second year of the A Level course, ideas about enthalpy are extended in *Oceans (O)*. The context used to study these concepts is the solubility of ionic solids and the entropy changes that happen during dissolving and other state changes in the hydrosphere.

As with the earlier study of enthalpy, it is important that learners know the definitions of the main enthalpy changes so that they can use values correctly in calculations. For dissolving ionic solids, the important changes are enthalpy change of solution, lattice enthalpy and enthalpy change of hydration of ions. Similar techniques to those used earlier to support learners to learn definitions can be used here. It is also valuable at this point to revisit the enthalpy changes learnt in *DF*, as well as the ionisation enthalpy, which will be required in enthalpy level diagrams connected with lattice enthalpy.

Learners use enthalpy level diagrams to show how the enthalpy changes form a cycle and how these can be used to explain the solubility of ionic solids based on the overall value for enthalpy changes of solution. Again, this area is rich in specific chemical language with specific meanings and learners need to be fluent in using the definitions of the different enthalpy changes involved in constructing and making use of the cycles. The real challenge is for learners to engage with the diagrams at the necessary level of detail. Learners can consider the solubility of ionic solids in aqueous and non-aqueous solvents by considering that non-aqueous solvents do not have the benefit of large, exothermic enthalpy changes of hydration to ‘balance’ the large energy input needed to separate ions due to the lattice enthalpy.

The dependence of lattice enthalpy on the charge density of ions can be presented as a data analysis task with students looking for patterns in data relating to lattice enthalpy for a range of ionic solids.

**Entropy**

The study of entropy may be introduced qualitatively initially, with the idea that the particles in the universe naturally continually change their arrangements. Chemists quantify this by considering the number of ways molecules and energy can be arranged (this links to the Maxwell Boltzmann distributions used to explain rates of reaction). Many ‘everyday’ examples illustrate this; the state of teenage bedrooms or the behaviour of young children in a playground (with, at any one moment, some children being motionless and others running around energetically). Even when simple systems of just a few molecules are considered, the number of possible arrangements increases quickly as more molecules are added or if a mixture of molecules is introduced.

These simple ideas follow on to thinking about how an ‘everyday’ appreciation of entropy applies to chemical systems. Learners consider patterns in entropy values for solids, liquids, gases and mixtures in terms of the number of ways particles can be arranged, and why increasing the size of a sample increases its entropy. Finally, the discussion moves to how these values can be used to calculate the system entropy changes for a reaction, using data about the entropy values of the reactants and products.

Learners can be led to appreciate that a total positive entropy change indicates that a spontaneous reaction will happen. This explains why many endothermic reactions (such as common salt dissolving in water) are spontaneous; although energy is indeed taken in, the entropy increases due to the large number of ways that ions in solution can be arranged compared to that of the solid.

The final extension of learners’ understanding is to consider how the total entropy change also must take into account the entropy change that happens in the surroundings, which inevitably depends on the change in temperature during the reaction. This can be appreciated qualitatively, in that exothermic changes increase the temperature and hence the entropy of the surroundings and that this can balance out any negative entropy change that happens to a system to give a positive total entropy change. Mathematically, this is expressed using the equations

\(\displaystyle \Delta_{\rm{tot}}S=\Delta_{\rm{sys}}S+\Delta_{\rm{surr}}S\)

and

\(\displaystyle \Delta_{\rm{surr}}S=-\Delta{H}/{T}\)

Learners may consider the spontaneous formation of ice or water at temperatures below and above the freezing point of water to illustrate how these equations work.

**Common misconceptions or difficulties students may have**

It is a difficult idea to explain to students that we cannot measure the actual enthalpy that a molecule ‘contains’, but that we can follow enthalpy changes during reactions. An everyday metaphor is how we can talk about somebody’s spending habits, or how much they have spent on a particular shopping trip, without knowing how much they have in their bank account.

Learners have difficulty in remembering the values in the enthalpy definitions. Which reactant or product does the ‘one mole’ refer to in a particular definition? This makes Hess’s Law diagrams difficult. Time spent learning the definitions is vital. Perhaps even old fashioned ‘quick tests’ as starters might help.

Some learners, even at A Level, still confuse positive and negative enthalpy changes. Learners must think in terms of ‘system’ and ‘surroundings’. By convention, in standard enthalpy change experiments involving solutions in e.g. a Styrofoam cup, the contents of the cup are taken to be the system and the cup itself the boundary between system and surroundings. This may lead to exothermic reactions being described as endothermic, if a temperature rise in the contents of the cup is interpreted as the system absorbing heat from the surroundings. It may help to explain that, in such experiments, an attempt is made to create a system that is isolated from the surroundings. Therefore, if a temperature rise is observed, this represents the energy released by the system that would be transferred to the surroundings if the system were not isolated.

This distinction between system and surroundings is also of central importance in studying entropy later. It is difficult for learners to intuitively understand that although the overall *energy* change is the same, the overall *entropy* increases. It is often easier for learners to grasp this concept for exothermic changes first; the energy ‘spreads out’ from the system to the surroundings and increases the number of ways particles can be arranged (increasing the entropy). The overall entropy change of reaction is determined from the sum of the entropy changes of the system and the surroundings. To understand how total entropy change affects the spontaneous nature of some reactions, learners need to separate the system and the surroundings so that they can see whether the entropy changes for each balance out into a positive or negative entropy change overall.

In applying enthalpy changes to the dissolving of ionic solids in *O*, learners encounter similar challenges to those in *DF*. They need to be able to define the different enthalpy changes so that they can apply them properly in calculation, for example ensuring that the change goes in the right direction on the enthalpy level diagram. Many learners find it difficult to ‘envisage’ the states and the signs of the changes. A common difficulty, for example, is remembering the negative sign in front of values for lattice enthalpy. A good understanding is required of what these processes mean; in this case, two distinct collections of charged ions become a neutral solid held together by electrostatic interactions. From discussion of the topic of bonding, learners should appreciate that this is an energetically more favourable arrangement, that therefore energy is released, which means a negative enthalpy change.

An everyday understanding of entropy is the ‘tendency to disorder’. Although this is useful in terms of considering that there are many disordered arrangements and only a limited number of ordered arrangements for particles, it is not scientifically correct. Learners should be encouraged to talk in terms of numbers of ways of arranging particles rather than disorder.

**Conceptual links to other areas of the specification – useful ways to approach this topic to set learners up for topics later in the course.**

An understanding of enthalpy is important to learning about entropy in *O*. In particular, it is important that learners engage with positive and negative enthalpy changes and learn that, as chemists, we think in terms of ‘the system’ and ‘the surroundings’ when we are talking about transfers of energy.

Ideas about energy and its effect on particles is also important in understanding rates of reaction. Enthalpy level diagrams are also met in the study of rates and catalysis (including enzymes). These ideas are studied in *The ozone story (OZ)*, *Chemical industry (CI)* and *Polymers and life (PL)*.

### Enthalpy changes transition guide

### Learning definitions 1 and 2

**Learning definitions 1**

Before tackling the calculations, learners need to know the definitions of the key enthalpy changes. For DF these are \(\displaystyle \Delta_rH\), \(\displaystyle \Delta_cH\), \(\displaystyle \Delta_fH\), \(\displaystyle \Delta_{neut}H\). These are very difficult to remember to the necessary level of detail.

One approach is to make a PowerPoint slide of each definition. Use the Animations menu to make the definition of each term disappear after 1 minute. Ask learners to sit with ‘hands on heads’ (at their age they find this hilarious) and tell them they have one minute to learn the definition. When the definition disappears after 1 minute, learners write it out quickly from memory.

Now check their versions by plenary. Have they used the term ‘enthalpy change’ (rather than energy needed/given out)? One mole? Of what? Standard states mentioned?

This makes a fun, repeated plenary or starter.

**Learning definitions 2**

Give learners cards for them to make their own flashcards. Each card should have the symbol for the enthalpy change on one side (e.g. \(\displaystyle \Delta_cH\)) and its definition on the other ‘The enthalpy change when…..’.

These can be used for several starter or plenary activities. For example

- (easy) learners work in pairs to match symbols (from one learner’s set) with definitions (from the other); next one learner holds the cards and reads a definition to their partner who has to say the name of the enthalpy change.

- (harder) one learner says the name of the enthalpy change and their partner has to give the full definition (this can be played as ‘Taboo’ with a timer).

### Starters for ten (Royal Society of Chemistry)

*DF*.

### Energy changes in neutralisation (Royal Society of Chemistry)

### System and surroundings

To support learners in looking at chemical systems, play the ‘where has it gone’ game.

Look at different examples of enthalpy changes and ask about where heat energy is being transferred *from* and where it is going *to*? Then ask learners to define the system and the surroundings in each case. Is the change exothermic or endothermic from the system’s point of view? Introduce the idea that for ‘open’ systems the surroundings extend to the whole universe.

These examples can be shown as images (sourced by internet image capture) in a PowerPoint or as mini-demos.

Possible examples: a burning jet of methane, a spray of air freshener on a watch glass on the back of a ‘volunteer’s’ hand (it feels very cold as it evaporates), hot water pipes heating water in a water boiler, a cup of tea cooling.

### Bond enthalpy calculations (Doc Brown)

### Group data analysis: Patterns in bond enthalpies

Ask learners to look up the average bond enthalpies and bond lengths for a series of related bonds, e.g.:

- C–C, C=C, C≡C (to show the effect of multiple bonds on bond length and enthalpy).
- H–F, H–Cl, H–Br, H–I (to show the effect of bond length).

Ask them to work in groups to put forward ideas to answer questions such as

- ‘What is the relationship between bond length and bond strength?’

- ‘What effect do multiple bonds have on bond length and strength?’

- ‘Do these patterns hold true for other sets of bonds?’ (This last question asks learners to decide for themselves what bonds to consider.)

To find the information they need, learners can use a data book or table, (see Resource for example).

This approach has the benefit of leading students to be independent at finding information and looking for patterns.

### Charge density and lattice enthalpy A Level Only

Introduce this topic by asking learners to look up data about the lattice enthalpy and ionic radii of a series of related compounds, e.g. the chlorides of group 1 and group 2. (WebElements online has all the data they will need: see Resource for link.)

Ask learners to consider the questions:

- How does the charge on an ion affect the lattice enthalpy?

- How does the ionic radius affect the lattice enthalpy?

- Is the same pattern true for compounds where the negative ion is different? (They should decide themselves which compounds they need to look at to investigate this idea.)

This approach has the advantage of encouraging independence at finding information and stimulating students to think models out for themselves. This leads to introducing charge density.

Clear, visual representations of some of the trends are given in the PowerPoint presentation (see Resource).

### An introduction to entropy (BBC) A Level Only

### Using entropy values for reactants and products to calculate entropy changes of reaction (Royal Society of Chemistry) A Level Only

### Quantum casino (Royal Society of Chemistry) A Level Only

The website in the resource has a number of simulations and videos that can be used to support the teaching of entropy in the classroom (and goes as far as the calculations of total entropy change) as well as tutorials which learners could use to support self-study.

The site has an extensive, indexed list of tutorials (which explain entropy like a text book with embedded videos), simulations (to show entropy in terms of ‘number of ways’ that particles can be arranged and videos of reactions). The full site explores Gibbs free energy, which is not needed for this specification, but this is in a separate section that can be left out.

### Entropy for the gifted and talented (Royal Society of Chemistry) A Level Only

## 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. Studying enthalpy in terms of its use in fuel efficiency is the main theme in *Developing fuels (DF)*, but there are other contexts that are also useful to consider, such as:

- self-heating or self-cooling food and drinks cans

- hand warmers

- sherbet

- cooling packs for sports

- lime-based building materials such as cement

All these contexts can be used for competitive ‘egg race’ group challenge activities.

It appears easy to find contexts for introducing entropy, but 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.

### Determination of the enthalpy change of neutralisation

### Determination of an enthalpy change of reaction by Hess' Law

### Determination of enthalpy changes of combustion

### Periodicity crossword

### Exploding methane tin (Nuffield Foundation)

*DF*as it shows how much energy is given out in a ‘familiar’ reaction between methane and oxygen.

### Making sherbet

An ‘egg race’ type activity for enthalpy. Tell learners that sherbet contains citric acid and sodium hydrogencarbonate (usually with sugar and jelly crystals for flavour). Support them to write and balance the equation for the reaction between the two main compounds.

The ‘best’ sherbet has the ‘biggest’ endothermic reaction between the two compounds. It gives the biggest ‘cool hit’ on the tongue when it gets wet. Ask learners to use the equation to work out the ‘best’ ratio by mass of compounds to use. They now test this practically by adding their mixtures to fixed volumes of water and measuring the temperature change. Which learners get the coolest mixture? Learners can extend this by investigating mixtures just either side of their calculated ‘optimum’ ratio.

The lesson could end with looking at some sherbet recipes on the internet and discussing whether you agree with the proportions given.

### Calculating enthalpy from experimental data in the context of the Thames Tideway Tunnel

### Chemistry and sport: Shooting (Royal Society of Chemistry)

### Hot dinner from a can (Royal Society of Chemistry)

### 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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