1. Energetics

2. Enthalpy Change ∆H Chemical energy is a special form of potential energy that lies within chemical bonds. Chemical bonds are the forces of attraction that bind together atoms in compounds. When chemicals react to form new substances, bonds break in the reactants and new bonds are formed as products are made. This process changes the chemical energy of the substances. Enthalpy, H, is the thermal energy that is stored in a chemical system. It is impossible to measure directly the enthalpy of the reactants or products. However, we can measure the energy absorbed or released to the surroundings during a chemical change. The form of this energy can vary but chemists usually measure energy changes in reactions by monitoring thermal energy. Thermal energy can be monitored using temperature changes. Enthalpy Change ∆H – the energy change measured at constant pressure.

3. You will often hear the system and the surroundings being discussed when enthalpy changes are considered: The system is the actual chemical reaction; that is, the atoms and bonds involved. The surroundings are everything else. If thermal energy is released, the amount of energy that leaves a chemical system is exactly the same as the amount that goes into the surroundings. No energy is lost. It just transfers from one place to another, and some energy might change from one form to another. This is called the law of conservation of energy.

4. This means that: heat loss in a chemical system = heat gain to the surroundings (accompanied by a temperature increase) heat gain in a chemical system = heat loss from the surroundings (accompanied by a temperature decrease). In some reactions, the products of a reaction have more chemical energy than the reactants. In others, the products have less chemical energy than the reactants. An enthalpy change, ΔH, is: the heat exchange with the surroundings during a chemical reaction, at constant pressure; the difference between the enthalpy of the products and the enthalpy of the reactants: ΔH = H products − H reactants

5. In general, all chemical reactions either release heat (exothermic reactions) or absorb heat (endothermic reactions) with the enthalpy changes usually measured in kJ mol −1 (kilojoules per mole). If an enthalpy change has a negative sign in front of it, the enthalpy change is exothermic, as the system is losing heat to the surroundings. If an enthalpy change has a positive sign in front of it, the enthalpy change is endothermic, as the system is gaining heat from the surroundings

6. For chemists to carry out meaningful calculations, all values must be measured under an agreed set of conditions. These are known as standard conditions. Standard conditions are: 100 kPa – this is the same as 100 000 Pa (0.986 atm). 273 K – remember, add 273 to convert from °C to K. All substances are in their standard states, i.e. the most stable form. For example, carbon would be graphite rather than diamond. Standard conditions are given the o symbol.

7. Enthalpy change of reaction is the energy change associated with a given reaction. It has the symbol ΔH o r, where: Δ is the Greek letter delta, which means large change H is enthalpy or energy in a system r stands for reaction o means under standard conditions.

8. A thermodynamic equation includes both the balanced chemical equation and enthalpy data. Consider the endothermic thermal decomposition of copper carbonate, which can be represented by the following thermodynamic equation: CuCO 3 (s) → CuO(s) + CO 2 (g) ΔH o r = +1146 kJ mol −1

9. Standard Enthalpy of Formation ΔH o f This is the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions. It has the symbol ΔH o f, where f stands for formation. This is an important definition to learn

10. When writing a thermodynamic equation you may need to use fractions to balance the equation, as you must ensure there is only 1 mole of product. Consider the exothermic reaction for the formation of water: H 2 (g) + ½ O 2 (g) → H 2 O(l) ΔH o f = −286 k J mol −1 The standard enthalpy change of formation for an element in its standard state is 0 kJ mol −1. This is because according to the definition there is no change (as no compound is formed), so no energy is released or taken in.

11. Standard Enthalpy of Combustion ΔH o c This is the enthalpy change that takes place when 1 mole of a substance is completely combusted in oxygen under standard conditions with all reactants and products in their standard states. It has the symbol ΔH o c, where c stands for combustion. This is an important definition to learn

12. The standard enthalpy of combustion for ethane can be represented by the following thermodynamic equation: C 2 H 6 (g) + 3½ O 2 (g) → 2CO 2 (g) + 3H 2 O(l) ΔH o c = −1560 kJ mol −1 The enthalpy of combustion of hydrogen is the same thermodynamic equation as the enthalpy of formation of water.

13. Calorimetry You cannot measure enthalpy directly but you can measure enthalpy change. The temperature of a chemical system is monitored. If there is a temperature rise, an exothermic reaction has caused heat to be gained by the surroundings. If a temperature drop is recorded, an endothermic reaction has taken in heat from the surroundings. Calorimetry uses a mathematical relationship to calculate enthalpy change from experimental quantitative data. The expression used is: q = mcΔT where: q is the heat exchanged with the surroundings, usually expressed in joules (J) m is the mass of the substance heated or cooled, usually expressed in grams (g) c is the specific heat capacity of the substance that is heated or cooled; it is the energy required to raise the temperature of 1 g of a substance by 1 K, usually expressed as J g −1 K −1 ΔT is the change in temperature, measured in kelvin (K).

14. Calorimetry uses a mathematical relationship to calculate enthalpy change from experimental quantitative data. The expression used is: q = mcΔT where: q is the heat exchanged with the surroundings, usually expressed in joules (J) m is the mass of the substance heated or cooled, usually expressed in grams (g) c is the specific heat capacity of the substance that is heated or cooled; it is the energy required to raise the temperature of 1 g of a substance by 1 K, usually expressed as J g −1 K −1 ΔT is the change in temperature, measured in kelvin (K).

15. Insulated container (‘coffee cup’) calorimetry – where two liquids or a solid and a liquid are mixed in an insulated container and the maximum temperature change is recorded. Add a measured mass of the first liquid reactant. Take the temperature every minute until it is stable. This usually takes around 4 minutes. At 5 minutes, add the second reactant. Do not take or record the temperature for the fifth minute. Monitor the temperature of the reaction mixture every minute for a further 5 minutes. Plot a graph to infer the maximum temperature change generated by the reaction.

16. Errors include heat transfer from surroundings (usually loss), approximation in specific heat capacity of solution. The method assumes all solutions have the heat capacity of water, neglecting the specific heat capacity of the calorimeter- we ignore any heat absorbed by the apparatus, reaction or dissolving may be incomplete or slow and density of solution is taken to be the same as water. These errors are usually small.

17. Conducting container (‘copper calorimeter’) calorimetry – A substance is combusted and the heat energy is transferred into known mass of water. When a fuel is combusted, the heat energy can be used to increase the temperature of a known mass of water. Measure the starting mass of the fuel. Add a known mass of water to a copper calorimeter. Mount the copper calorimeter over the fuel and take the starting temperature of the water. Combust the fuel for a few minutes and take the final temperature of the water. Take the mass of the unused fuel and calculate the mass of the fuel burnt.

18. Errors include heat losses from calorimeter, incomplete combustion of fuel, incomplete transfer of heat, which can be significant. Evaporation of fuel after weighing, heat capacity of calorimeter not included, measurements not carried out under standard conditions, which are usually minor.

19. Bomb calorimetry - A bomb calorimeter is a sophisticated piece of equipment that minimises heat loss as much as possible. It uses pure oxygen, to ensure complete combustion is achieved.

20. Hess’s Law In a chemical reaction, new substances are made and an energy or enthalpy change occurs. Some enthalpy changes are impossible to measure directly using calorimetry. For these reactions, enthalpy cycles can be used to indirectly calculate the enthalpy change. Hess’ law states that the enthalpy change in a chemical reaction is independent of the route it takes.

21. Hess’s Law Consider a chemical reaction where reactant chemicals A and B form product C. This can be written as a balanced chemical equation: A + B → C Imagine that sometimes A and B can react to make D, and then D can react to become C. This can be shown by two balanced chemical equations: A + B → D, then D → C

22. These chemical reactions can be joined in an enthalpy cycle. An enthalpy cycle shows alternative routes between reactants and products. A simple enthalpy cycle should look like a triangle. Each corner of the shape should have chemicals. All the corners should have the same amount of each atom, so you may need to add numbers to balance them. The sides of the triangle are made of arrows. These show the direction of a chemical reaction, from the reactants to the products, and the enthalpy change can be written on the arrow.

24. The enthalpy change to go directly from A and B to C is the same as if you go through the intermediate D. So, Δ 1 H = Δ 2 H + Δ 3 H. Hess’ law states that the enthalpy change in a chemical reaction is independent of the route it takes. An enthalpy cycle is a pictorial representation showing alternative routes between reactants and products.

25. If you are a good mathematician, you may not need to draw the cycle. With enthalpy of combustion data, the unknown enthalpy change can be calculated using the following formula: ΔH = ΣΔH c (reactants) - ΣΔH c (products) With enthalpy of formation data, the unknown enthalpy change can be calculated using the following formula: ΔH = ΣΔH f (products) – ΣΔH f (reactants) where Σ is the Greek letter sigma and means ‘sum of’.

26. Bond enthalpies Energy is needed to break a chemical bond and the same amount of energy is released when the same bond is made. You get information about the strength of a chemical bond from its bond enthalpy. Bond enthalpies tell you how much energy is needed to break each different bond. You can then compare the strengths of different bonds. Note that energy is needed to break bonds – the change is endothermic. When bonds form, the same quantity of energy is released – the change is exothermic.

27. The equations below show the bond enthalpies for H–H and H–Cl bonds: H–H(g) → 2H(g) ΔH = +436 kJ mol −1 H–Cl(g) → H(g) + Cl(g) ΔH = +432 kJ mol −1 The H–H bond enthalpy value is always the same because a H–H bond can only ever exist in a H 2 molecule. Similarly, the H–Cl bond enthalpy applies only to a HCl molecule. Unlike H–H and H–Cl bonds, some bonds can occur in different molecules. For example, almost every organic molecule contains C–H bonds. The C–H bond strength will vary across the different environments in which it is found. These are averaged over a number of typical chemical species containing that type of bond.

28. A chemical reaction is often modelled as a three-step process: 1. Reactant bonds are broken. This process takes in energy and so is endothermic. 2. Atoms rearrange to form products. 3. Product bonds are formed. This releases energy and is an exothermic change. In an endothermic reaction, more energy is needed to break the reactant bonds than is released when the product bonds are made. Overall the reaction takes in energy. In an exothermic reaction, more energy is released when product bonds are formed than is needed to break reactant bonds. Overall the reaction releases energy.

29. You can predict the enthalpy change for a reaction by using the average bond enthalpy data for the reactants and products. The average bond enthalpy is the mean energy needed for 1 mole of a given type of gaseous bonds to undergo homolytic fission. The actual bond enthalpy is individual to each molecule. However, an average value is taken across a variety of molecules. So, the enthalpy change for a reaction can be calculated from average bond enthalpy data using the following expression: ΔH = Σ(bond enthalpies of reactants) − Σ(bond enthalpies of products) where Σ is the Greek letter sigma and means ‘sum of ’.

30. The enthalpy change for an incomplete combustion reaction is difficult or impossible to measure directly. This is because there is often complete and incomplete combustion happening simultaneously. However, using average bond enthalpy data, this enthalpy change can be calculated. Remember, though, that the values are averages for the bond in a variety of environments in different molecules, so it is not as accurate as direct measurement using precise and sensitive equipment like a bomb calorimeter.