Energetics

• Enthalpy change


  The amount of heat energy taken in or given out during any change in a system provided the pressure is constant
• Enthalpy change occurs


  1. Energy is transferred between system and surroundings

  2. The system is the chemicals

  3. The surroundings is everything outside the chemicals
• Exothermic change


  • Energy is transferred from the system (chemicals) to the surroundings
  • The products have less energy than the reactants
  • ∆H is negative
• Endothermic change


  • Energy is transferred from the surroundings to the system (chemicals)
  • They require an input of heat energy e.g. thermal decomposition of calcium carbonate
  • The products have more energy than the reactants
  • ∆H is positive
• Enthalpy changes are normally quoted at standard conditions
• Standard conditions


  • 100 kPa pressure
  • 298 K (room temperature or 25oC)
  • Solutions at 1mol dm-3
  • All substances should have their normal state at 298K
• Standard Enthalpy Change of Formation

  The enthalpy change when 1 mole of the compound is formed from its elements under standard conditions (298K and 100kpa), all reactants and products being in their standard states
• The enthalpy of formation of an element = 0 kJ mol-1
• Standard Enthalpy Change of Combustion

  The enthalpy change that occurs when one mole of a substance is combusted completely in oxygen under standard conditions (298K and 100kPa), all reactants and products being in their standard states
• Incomplete combustion will lead to soot (carbon), carbon monoxide and water. It will be less exothermic than complete combustion.
• Common oxidation exothermic processes are the combustion of fuels and the oxidation of carbohydrates such as glucose in respiration.
• Measuring the Enthalpy Change for a Reaction Experimentally

  1. Calorimetric method

  2. General method
• Errors in the calorimetric method include: energy transfer from surroundings, approximation in specific heat capacity of solution, neglecting the specific heat capacity of the calorimeter, reaction or dissolving may be incomplete or slow, density of solution is taken to be the same as water.
• Calculating the enthalpy change of reaction, ∆H from experimental data

  1. Using q = m x cp x ∆T to calculate energy change for quantities used

  2. Work out the moles of the reactants used

  3. Divide q by the number of moles of the reactant not in excess to give ∆H

  4. Add a sign and unit (divide by a thousand to convert Jmol-1 to kJmol-1)
• The heat capacity of water is 4.18 J g-1K-1. In any reaction where the reactants are dissolved in water we assume that the heat capacity is the same as pure water.
• Enthalpies of combustion can be calculated by using calorimetry. Generally the fuel is burnt and the flame is used to heat up water in a metal cup.
• Errors in measuring enthalpies of combustion using calorimetry include: energy losses from calorimeter, incomplete combustion of fuel, incomplete transfer of energy, evaporation of fuel after weighing, heat capacity of calorimeter not included, measurements not carried out under standard conditions as H2O is gas, not liquid, in this experiment.
• Hess's law


  The total enthalpy change for a reaction is independent of the route by which the chemical change takes place
• Hess's law


  Total enthalpy change for a reaction is independent of the route by which the chemical change takes place
• Hess's law is a version of the first law of thermodynamics, which states that energy is always conserved
• Hess's law


  1. Determine the enthalpy change for a reaction that cannot be measured directly by experiments

  2. Carry out alternative reactions that can be measured experimentally

  3. Use Hess's law cycles to calculate the enthalpy change
• H2 (g) + Cl2 (g) → 2HCl (g)


  • One route is arrow 'a'
  • The second route is shown by arrows ΔH plus arrow 'b'
  • So a = ΔH + b
  • And rearranged
  • ΔH = a - b
• H+ (g) + Br- (g) → HBr (g)


  • One route is arrow 'a' plus ΔH
  • The second route is shown by arrows 'c' plus arrow 'd'
  • So a+ ΔH = c + d
  • And rearranged
  • ΔH = c + d - a
• Using Hess's law to determine enthalpy changes from enthalpy changes of formation


  H reaction = Σ fH products - Σ fH reactants
• Using Hess's law to determine enthalpy changes from enthalpy changes of combustion


  H reaction = Σ cH reactants - Σ cH products
• Elements in standard states have fH = 0
• fH(MgO)= -601.7 kJ mol-1, fH(Al2O3) = -1675.7 kJ mol-1
• cH CO(g) = -283 kJ mol-1, cH H2 (g)= –286 kJ mol-1, cH CH3OH(g)= –671 kJ mol-1
• Mean bond energy


  The enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules
• Mean bond energies are positive because energy is required to break a bond
• The definition of mean bond energy only applies when the substances start and end in the gaseous state
• Using mean bond energies to calculate enthalpy changes

  H = Σ bond energies broken - Σ bond energies made
• Calculated values of enthalpy of combustions will be more accurate if calculated from enthalpy of formation data than if calculated from average bond enthalpies
• Experimental results for enthalpy of combustion will be lower than calculated values due to heat loss and incomplete combustion
• As one goes up a homologous series of alcohols

  The enthalpy of combustion increases by a constant amount