3.2

• The total chemical energy inside a substance is called the enthalpy (or heat content)
• Enthalpy change


  Represented by the symbol ΔH where Δ= change and H = enthalpy
• Exothermic reactions


  • Products have less energy than the reactants
  • Heat energy is given off by the reaction to the surroundings
  • Temperature of the environment increases
  • Energy of the system decreases
  • Enthalpy decrease during the reaction so ΔH is negative
  • Thermodynamically possible as the enthalpy of the reactants is higher than that of the products
• An energy level diagram shows the energy level of the reactants, transition state(s), energy level of the products, activation energy, and the enthalpy change for the reaction
• Energy level diagram

  • Shows the energy level of the reactants
  • Transition state(s)
  • Energy level of the products
  • Activation energy (E)
  • Enthalpy change for the reaction (ΔH)
  • Overall energy taken in from/given out to the surroundings or the energy difference from reactants to products
• Endothermic reactions


  • Products have more energy than the reactants
  • Heat energy is absorbed by the reaction from the surroundings
  • Temperature of the environment decreases
  • Energy of the system increases
  • Enthalpy increase during the reaction so ΔH is positive
• The energy difference from reactants to products is released out to the surroundings
• Combustion reactions are always exothermic (ΔH is negative)
• The energy level diagram for the reaction of hydrogen with chlorine to form hydrogen chloride gas
• Determining the activation energy

  ΔH for a reaction is +70 kJ mol and Ea for the reverse reaction is +20 kJ mol. E for the forward reaction = (+70 kJ mol) + (+20 kJ mol) = +90 kJ mol
• The activation energy, Ea, and enthalpy change, ΔH, for the complete combustion of methane are +2653 kJ mol and -890 kJ mol respectively
• All thermodynamic measurements are carried out under standard conditions
• Standard conditions for thermodynamic measurements

  • A pressure of 100 kPa
  • A temperature of 298 K (25 C)
  • Each substance involved in the reaction is in its standard physical state (solid, liquid, or gas)
• Calculating the enthalpy change

  1. One mole of water is formed from hydrogen and oxygen, releasing 286 kJ of energy

  2. Calculate ΔH for the reaction: 2H (g) + O (g) → 2H O (I)

  3. Calculate ΔH for the reaction: 4Fe (s) + 3 O (g) → 2 Fe O (s)
• The ΔH of an element in its standard state is zero
• Calorimetry


  1. Calorimetry is the measurement of enthalpy changes in chemical reactions

  2. A simple calorimeter can be made from a polystyrene drinking cup, a vacuum flask, or metal can

  3. A polystyrene cup can act as a calorimeter to find enthalpy changes in a chemical reaction

  4. The energy needed to increase the temperature of 1 g of a substance by 1°C is called the specific heat capacity (c) of the liquid

  5. The specific heat capacity of water is 4.18 J g K

  6. The energy transferred as heat can be calculated by the equation for calculating energy transferred in a calorimeter
• Specific heat capacity calculations

  1. In a calorimetry experiment 2.50 g of methane is burnt in excess oxygen

  2. 30% of the energy released during the combustion is absorbed by 500 g of water, the temperature of which rises from 25°C to 68°C

  3. What is the total energy released per gram of methane burnt?
• Step 2

  q = 500 x 4.18 x 43 = 89 870 J
• Step 1

  1. q = m x c x ΔT

  2. m (of water) = 500 g

  3. c (of water) = 4.18 J g °C

  4. ΔT (of water) = 68 C - 25 C = 43 C
• What is the total energy released per gram of methane burnt?
• The specific heat capacity of water is 4.18 J g °C
• Step 3

  1. This is only 30% of the total energy released by methane

  2. Total energy x 0.3 = 89 870 J

  3. Total energy = 299 567 J
• Step 4

  Energy released by 1.00 g of methane = 299 567 ÷ 2.50 = 120 000 J g (to 3 s.f.) or 120 kJ g
• Aqueous solutions of acid, alkalis and salts are assumed to be largely water so you can just use the m and c values of water when calculating the energy transferred
• The amount of energy required to break one mole of a specific covalent bond in the gas phase is called the bond dissociation energy
• When the temperature falls, the value for ΔH becomes positive suggesting that the reaction is endothermic
• When there is a rise in temperature, the value for ΔH becomes negative suggesting that the reaction is exothermic
• Bond dissociation energy, E, is usually just simplified to bond energy or bond enthalpy
• Bond energies are affected by other atoms in the molecule, so average bond enthalpies are listed in data tables
• As energy is required to break bonds, bond breaking is endothermic. ΔH is positive
• Bond energies are used to find the ΔH of a reaction when this cannot be done experimentally
• As energy is released making new bonds, bond forming is exothermic. ΔH is negative
• The difference between the energy required for bond breaking and the energy released by bond making determines whether an overall reaction is exothermic or endothermic
• Calculating the enthalpy change in the Haber process
• Bond forming is exothermic
• It is important to be aware that the actual bond enthalpy value may differ from the average value
• Haber process


  Producing ammonia from hydrogen and nitrogen
• Bond breaking is endothermic
• Bond energies


  • Given in the table
• The enthalpy of combustion is the enthalpy change when one mole of a substance reacts in excess oxygen to produce water and carbon dioxide