Energetics

1. Energetics

2. Contents Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of reaction (∆Hѳ) State that combustion and neutralization are exothermic processes. Apply the relationship between temperature change and the classification of a reaction as endothermic and exothermic. Deduce, from an enthalpy level diagram, the relative stabilities of reactants and products, and the sign of the enthalpy change for the reaction.

3. Exothermic and Endothermic reactions • Chemical reactions involve the ________ and __________of bonds. • It requires energy to break bonds and energy is given out when new bonds are formed. • If more energy is given out than needs to be put in, then this will be released as heat to the surroundings, and the reaction is said to be___________. • If more energy is required to break the reactant bonds than is given out when the products are formed, then extra heat will be need to be added, and the reaction is said to be____________. making endothermic breaking exothermic

5. Enthalpy • The energy contained in chemical bonds that can be converted into heat is known as enthalpy and is given the symbol H. • It is impossible to measure the actual heat content or enthalpy of a particular substance, but what can be readily measured is the enthalpy change for a reaction, ΔH.

6. Exothermic reactions • For exothermic reactions the heat content (enthalpy) of the products is less than that of the reactants. • By convention the enthalpy change for the reaction, ΔH, is said to be negative (-). Reactants - ΔH Enthalpy Products Extent of the reaction

7. Examples of exothermic reactions • Examples of exothermic reactions include: • the combustion of fuels, such as hydrogen or gasoline (petrol) C8H18(l) + 12½O2(g) → 8CO2(g) + 9H2O(l)ΔH = -5512kJ mol-1 • the neutralization of an acid by a base: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) ΔH -57.3kJ mol-1

8. Endothermic reactions • For endothermic reactions, where the enthalpy of the products is more than the enthalpy of the reactants, the enthalpy change, ΔH, will have a positive value (+). • An example of an endothermic reaction is the combination of nitrogen with oxygen to form nitrogen dioxide. N2(g) + O2(g) → NO2(g) ΔH = +33.9 kJ mol-1 Products Enthalpy + ΔH Reactants Extent of the reaction

9. Another example of an endothermic reaction is the reaction between solid hydrated barium hydroxide and solid ammonium chloride: Ba(OH)2.8H2O(s) + 2NH4Cl(s) → BaCl2.2H2O(s) + 2NH3(g/aq) + 8H2O(l) ΔH = +96.0 kJ mol-1 • This is interesting not only because it is a reaction between two solids, but also because the temperature decrease is such that the flask will often stick to the bench after the reaction. • When there is a change of state from solid to a liquid, or a liquid to a gas at a constant temperature, the process is also endothermic. • For example: Na (s) → Na (g) ΔH = +108 kJmol-1

10. Standard Enthalpy Change of Reaction • The actual amount of heat given out or taken in during a chemical reaction will depend on several factors: • The nature of the reactants and products; different reactions have different values of enthalpy changes, ΔH, as different bonds with different strengths are being broken and formed. • The amount (or concentration) of reactants; the greater the amount that reacts , the greater the heat change. • The states of the reactants and products – changing state involves an enthalpy change, and so will affect the total amount of heat given out or taken in. • The temperature of the reaction • The pressure at which the reaction is carried out

11. All these factors are taken into account by referring to standard enthalpy changes. • The standard enthalpy change of the reaction, ΔH θ, is defined as: • Standard pressure is 101.3 k Pa (1 atmosphere) and standard temperature is any specified temperature, but usually this is taken as 298 K (25oC). • The symbol θ signifies that the enthalpy change has been measured under standard conditions. the enthalpy change when molar quantities of reactants in their normal states react to form products in their normal states under standard conditions of temperature and pressure

12. It is worth noting that enthalpy changes are normally carried out under constant pressure. • If the volume is kept constant then the value for the heat change will be slightly different from that at constant pressure, because some work is involved in expanding or contracting a gas. • Strictly speaking the enthalpy change, ΔH, is the heat change if the pressure is kept constant; if the volume is kept constant then the total change is designated by ΔU.

13. Activity • Carry out the experiment: Measuring some energy changes.

14. Contents • Calculate the heat energy change when the temperature of a pure substance is changed. • Design the suitable experimental procedures for measuring the heat energy changes of reactions. • Calculate the enthalpy change for a reaction using experimental data on temperature changes, quantities of reactants and mass of water • Evaluate the results of experiments to determine enthalpy changes.

15. Measuring enthalpy changes • The general name for any reaction is the standard molar enthalpy change of reaction, ΔH θ. • It is measured in kilojoules per mole, kJ mol-1. • We write a balanced symbol equation for the reaction and then find the heat change for the quantities in moles given by the reaction. • For example, ΔH for: 2NaOH + H2SO4 → Na2SO4 + 2H2O is the enthalpy change when 2 moles of sodium hydroxide react with one mole of H2SO4

16. Heat and Temperature • The words ‘heat’ and ‘temperature’ are often used to mean the same thing in daily conversation but in science they are quite distinct and you must be clear about the difference. • Temperature is a measure of the average kinetic energy of the particles in a system. • As the particles move faster, their average kinetic energy increases and the temperature goes up. • But it does not matter how many particles there are – temperature is independent of the number present. We measure temperature with a thermometer.

17. Heat is a measure of the total energy of all the particles present in a given amount of substance. • It does depend on how much of the substance is present. • The energy of every particle is included. There is no instrument that can measure heat directly. • So a bath of lukewarm water has much more heat than a sparkler even though the temperature is less. • Heat always flows from high to low temperature, so heat will flow from the sparkler into the bath water, even though the water has much more heat than the sparkler.

19. Calculating the enthalpy change of reaction • The enthalpy change of a reaction is the heat given out or taken in as the reaction heats up or cools down. • There is no instrument that measures heat directly. • To measure the enthalpy change we need to know three things: • The mass of the substance that is being heated up or cooled down. • The temperature change • The specific heat capacity of the substance

20. Specific Heat Capacity • The specific heat capacity, c, is the amount of heat needed to raise the temperature of 1g of substance by 1K. • Its units are joules per gram per kelvin, or J g-1 K-1 • Water has a high specific heat capacity and is able to absorb considerable quantities of heat without its temperature increasing significantly. • One gram of water requires 4.18 J of heat energy for the temperature of the water to increase by 1K. • This is known as the specific heat capacity of water, and is often quoted in SI units as; 4.180 J g-1 K-1 • It is important to make sure that the units agree with each other when carrying out enthalpy calculations.

21. Enthalpy change • To calculate enthalpy change we use the equation: • Mass and temperature change are found out from experiment. • The specific heat capacity of water (or dilute aqueous solution) is 4.18J g-1 K-1 Enthalpy change Mass of substance Specific heat capacity Temperature change x x = ΔH = m x c x ΔT

22. Reactions carried out in solution • It is relatively easy to measure heat changes for reactions that take place in solution. • The heat is generated in the solutions themselves and only has to be kept in the calorimeter. • We often use expanded polystyrene beakers for the calorimeters. • These are good insulators (this reduces heat loss through their sides) and they have a low heat capacity so the beaker itself absorbs very little heat. • We usually take the specific heat capacity of dilute solutions to be the same as water, 4.18J g-1 K-1

23. Exothermic or endothermic? A B Watch the video clips. Which reaction is exothermic and which is endothermic? Calculate the enthalpy change for each reaction, show your working.

24. The Simple Calorimeter • We can use this apparatus to find the enthalpy change when a fuel burns. • The enthalpy change when one mole of substance in its standard state is completely burned in oxygen is called the standard molar enthalpy of combustion, ΔHθc • We burn the fuel to heat a known mass of water and then measure the temperature rise of the water. • We assumeall the heat from the fuel goes into the water. • The apparatus used is called a calorimeter.

25. An example • A calorimeter was used to measure the enthalpy change of combustion of methanol: • 0.32g (0.010mol) of methanol was burned and heated 200g of water, the water temperature rose by 4K. • 0.01 mol gives 3344 J so 1 mol would give 334400J or 334 kJ • ΔHc = -334 kJ mol-1(negative because heat is given out) CH3OH(l) 2H2O(l) + 1½ O2(g) → CO2(g) + Heat change = ΔH = m x c x ΔT = x x 200 4.18 4 = 3344J

26. This value is less than half the accepted value since not all the heat reaches the water, can you suggest reasons why?: • There may be draughts that take heat away from the calorimeter • Heat is lost from the top and the sides of the calorimeter. • The beaker and the rest of the apparatus itself gets heated rather than just the water • Not all the methanol burns completely to carbon dioxide. This is incomplete combustion. If there is not enough oxygen available, then we get some carbon monoxide and carbon formed. Soot (carbon) on the bottom of the beaker will show if this has happened.

27. Flame calorimeter • Watch the video and record the results in a table similar to the one below; • Calculate the enthalpy of combustion for each of these reactions

28. Contents • Determine the enthalpy change of a reaction that is the sum of two or three reactions with known enthalpy changes.

29. Hess’s Law • Consider a chemical reaction in which reactant A is converted into product B. A → B • Let us suppose it can either be converted directly into B with an enthalpy change of ΔHx, or can be converted directly into B via an intermediate product C with a total enthalpy change of ΔHy + ΔHz A B C • In other words, the reaction can be proceed by two different pathways.

30. Hess’s Law states that the total enthalpy change of a reaction depends only on the difference between the enthalpy of the products and the enthalpy of the reactants. It is independent of the reaction pathway. • This is a useful law in chemistry, because it enables us to work out some enthalpy change values indirectly. • Consider the standard enthalpy of formation of methane. The equation for this reaction, in which methane is formed from its elements in their normal states under standard conditions, is C(s) + 2H2(g) → CH4(g) ΔHf θ

31. Superficially this looks a simple enough reaction. • We could react carbon with hydrogen and determine the heat change. • However, it is difficult to get carbon to react with hydrogen, and when it does there are many possible compounds of carbon and hydrogen (called hydrocarbons) that could be formed, so the reaction is actually impossible to achieve experimentally. • However, what we can do easily is burn carbon, hydrogen and methane separately and determine the standard enthalpy change of combustion, ΔHcθ, directly for the three substances

32. ΔHf θ This reaction is impossible to achieve experimentally C (s) + 2H2(g) CH4(g) 3. 1. 2. But all three components in the reaction can be reacted with oxygen easily – this is called the enthalpy of combustion H2O (l) CO2(g) According to Hess Law the total enthalpy of the reaction is independent of the reaction pathway 1. C (s) + O2(g) → CO2 (g)ΔHcθ = -393 kJ mol-1 2. H2 (g) + ½O2(g) → H2O (l)ΔHcθ = -286 kJ mol-1 3. CH4(s) + O2(g) → CO2 (g) + H2O(l)ΔHcθ = -890 kJ mol-1

33. Contents • Define the term average bond enthalpy. • Explain in terms of average bond enthalpies, why some reactions are exothermic and others are endothermic.