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An overview of enthalpy changes during chemical reactions, explaining the concepts of exothermic and endothermic reactions, and discussing temperature changes, experiments, and calorimetry. It covers the effects of temperature and time on reactions, and includes examples and calculations.
Typology: Lecture notes
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Enthalpy Changes All substances contain chemical energy, called enthalpy. Like any energy it is measured in Joules (previously energy was measured in Calories). When reactions happen, energy is given out or taken in – these are enthalpy changes.
In an EXOTHERMIC reaction: Chemical energy (enthalpy) is being turned into heat energy which is transferred to the surroundings, so the temperature we measure increases.
Experiments to investigate temperature changes CONTROL
For each repeat of the experiment:
Q = m x c x ΔT Energy released = 100 x 4.2 x 25 = 10,500 J (or 10.5 kJ) Comparing enthalpy changes To compare reactions, we can calculate the energy change per mole of reactant used, which we call the molar enthalpy change , Δ H. To do this we work out the energy released as before, then divide by the moles reactant used. Δ H = -Q / moles Why do you think this equation needs a – sign in front of Q?
ΔH is the chemical energy (enthalpy) in the chemical substances. In an exothermic reaction enthalpy decreases when chemical energy is turned into heat energy and given out. So ΔT increases, and has a Q positive sign. This shows that ΔH has an opposite sign to Q and ΔT. Chemical energy Reaction Heat energy Chemical energy Heat energy Δ H (negative) Δ T (positive)
Example: A 0.5g sample of ethanol is burnt, raising the temperature of 250g of water from 10˚C to 28˚C. When 0.8g of butane is burnt, the temperature of 250g of water increases from 10˚C to 40˚C. c = 4.2 J/g/˚C Which fuel produces the most heat energy per mole of fuel burnt? For ethanol: Q = (4.2 x 250 x 18) = 18,900 J moles = mass/RFM = 0.5/46 = 0. ΔH = -Q/moles = -18,900/0. = -1,738,730 J/mol or -1,739 kJ/mol For butane: Q = (4.2 x 250 x 30) = 31500 J moles = mass/RFM = 0.8/58 = 0. ΔH = -Q/moles = -31500/0. = -2,284,264 J/mol or -2,284 kJ/mol
Practical Which fuel produces more energy per mole of fuel burnt, hexanol (C 6 H 13 OH) or ethanol (C 2 H 5 OH)?
Practical Compare the molar enthalpy change from these two reactions:
What happens during a reaction We know that when a reaction takes place, bonds are broken, and new bonds formed. Consider the reaction between hydrogen and chlorine molecules to make hydrogen chloride: H 2
Cl Cl List the bonds that have to be broken List the bonds that have to be made
So bond breaking requires heat energy to be taken from the surroundings and used to break the bonds. The surroundings get cooler. Bond breaking is ENDOTHERMIC When new bonds form, energy is given out. This causes the surroundings to heat up. Bond forming is EXOTHERMIC. (The amount of energy given out is equal to the bond energy for that bond) Look at the list of bonds broken and made in the reaction from before **H 2
One H-H bond is broken: 436 (kJ/mol) One Cl-Cl bond is broken: 242 TOTAL ENERGY TAKEN IN = 678 kJ/mol Two H-Cl bonds are made 2 x 431 TOTAL ENERGY GIVEN OUT = 862 kJ/mol Overall, is energy taken in or given out? The molar enthalpy change for reactions can be calculated: Δ H = energy taken in – energy given out In this example the molar enthalpy change is - kJ/mol. Overall, energy is being given out so this is an exothermic reaction.
Energy level diagram - exothermic These show the relative energy levels (enthalpy) of reactants and products: H 2
Reaction profile diagram This adds information to an energy level diagram, showing activation energy and ΔH: (reactants) (products) Δ H (-ve) Enthalpy activation energy Exothermic H 2
