Chapter 9 - Enthalpy Changes

• Enthalpy: Measure of heat energy in a chemical system
• System: Refers to the atoms, molecules, or ions making up the chemicals
• Enthalpy change: H(products) – H(reactants)
• Exothermic

  • Chemical system releases heat energy to the surroundings
  • ΔH is negative
  • Temperature of the surroundings increases
• Endothermic

  • Chemical system absorbs heat energy from the surroundings
  • ΔH is positive
  • Temperature of the surroundings decreases
• Activation Energy: minimum energy required for a reaction to take place
• Standard Conditions

  Temperature: 298 K

  Pressure: 100 kPa

  Concentration: 1 mol dm-3

  Standard states: State at RTP
• Standard Enthalpy change of a reaction:

  Enthalpy change when a reaction occurs in the molar quantities shown in a chemical equation under standard conditions, with all reactants and products in their standard states
• Enthalpy change of formation:

  Enthalpy change when 1 mole of a compound is formed from its elements under standard conditions, with all reactants and products in their standard states.
• Enthalpy change of combustion:

  Enthalpy change when one mole of a substance reacts completely with oxygen under standard conditions with all reactants and products in their standard states
• Enthalpy change of neutralisation:

  Enthalpy change when one mole of H+ reacts with one mole of OH- to form one mole of water with all reactants and products in their standard states
• Calculating Energy Changes:

  q = mcΔT

  q = Energy change in J

  m = Mass of water (g) (density of water is 1 g cm-3 so 250 cm3 would be 250 g)

  c = Specific Heat Capacity of water = 4.18 J g-1 K-1

  ΔT = Change in temperature (add 273 to convert between °C to K)
• Calculating number of moles of a burning fuel:

  n = m / Mr
• Enthalpy Change of Combustion:

  ΔCH = q / n
• Practical Procedure

  • Measure a known volume of water and transfer to a copper can
  • Measure the temperature of the wate
  • Measure a mass of fuel in a spirit burner
  • Heat water for a 30-40 °C rise
  • Extinguish fuel and measure the final temperature and mass of the spirit burner
• Explain why there is theoretical and experimental variation

  1. Heat loss to the surroundings – q would be lower than the actual value meaning ΔcH would be lower than expected

  2. Incomplete combustion. Less energy released so q is lower than expected and so is ΔcH

  3. Evaporation of fuel from the wick, n is higher than expected so ΔcH is lower than expected

  4. Non standard conditions. The data book value is standard conditions whereas the practical is not (water formed is gas not liquid)
• Improving practical

  • Draught excluder to minimise heat loss
  • Copper can over a glass beaker as it is a better conductor of heat
  • Ensure calorimeter is closer to the flame to minimise heat loss
• Determination of an enthalpy change of reaction

  • Polystyrene cup (insulator to keep heat in)
  • Lid (insulator to keep heat in)
  • Place polystyrene cup in a beaker (create a layer of air – insulator)
  • Transfer a known volume of solution using a volumetric pipette to polystyrene cup
  • Measure temperature of solution
  • Add a known mass of solid or known volume of solution
• Improving practical:

  Plot a cooling curve, extrapolate to determine a more accurate change in temperature
• Average bond enthalpy:

  Energy required to break one mole of a specified type of bond in a gaseous molecule
• • Energy is always required to break bonds
  • Bond enthalpies are always endothermic so have a positive value
• Limitations of average bond enthalpies:

  • Bond enthalpies are calculated from an average of a bond in different environments
  • Needs all reactants and products to be gaseous, whereas, some will be solid/liquid
• Bond Breaking and Bond Making

  Bond breaking: Energy required to break bonds: Endothermic

  Bond Making: Energy released to form bonds: Exothermic
• Calculating enthalpy change from average bond enthalpies:

  ΔrH = Σ(bond enthalpies in reactants) – Σ(bond enthalpies in products)
• Hess’s Law:

  If a reaction can take place by two routes, and the starting and finishing conditions are the same, the total enthalpy change is the same for each route.
• Calculating the enthalpy change of formation

  • Construct an enthalpy cycle between the reactants, products and the elemental intermediates
  • Both arrows point up, because we are converting elements into compounds
  • Work out the subtotal of each arrow
  • ΔCH = -ΔH1 + ΔH2
• Calculating the enthalpy change of combustion

  • Construct an enthalpy cycle between the reactants, products and the products of combustion
  • Both arrows point down, because we are burning substances to make carbon dioxide and water
  • Work out the subtotal of each arrow
  • ΔH = ΔH1 - ΔH2