9 - Enthalpy

• universe = system + surroundings
• when a reaction involving an enthalpy change takes place, heat energy is transferred between the system and the surroundings
• enthalpy (H) - the measure of heat energy present in a chemical system
• enthalpy change (ΔH) - the amount of energy released/taken in during a chemical reaction
• exothermic (- ΔH) - a reaction which releases energy to the surroundings
• endothermic (+ ΔH) - a reaction which takes in energy from the surroundings
• label the diagram

  A) exothermic

  B) reactants

  C) products

  D) energy

  E) progress of reaction

  F) activation energy

  G) enthalpy change

  H) negative

  
• label the diagram

  A) endothermic

  B) reactants

  C) products

  D) energy

  E) progress of reaction

  F) activation energy

  G) enthalpy change

  H) positive

  
• activation energy (E a) - the minimum amount of energy required for a reaction to occur (the energy required to break the bonds in the reactants)
• in an exothermic reaction, the temperature of the surroundings increases
• in an endothermic reaction, the temperature of the surroundings decreases
• standard conditions (ΔHΦ):

  • 100 kPa
  • 298 K (25 °C)
  • 1 mol dm-3 - solutions only
  • standard state - the physical state of a substance under standard conditions
• standard enthalpy change of reaction (∆rHΦ)

  • enthalpy change when
  • a reaction occurs in the molar quantities shown in the chemical equation
  • under standard conditions with all products and reactants in standard state
• standard enthalpy change of formation (∆fHΦ)

  • enthalpy change when
  • one mole of a substance is formed from its reacting elements
  • under standard conditions with all products and reactants in standard state
• standard enthalpy change of combustion (∆cHΦ)

  • enthalpy change when
  • one mole of a substance is burned completely in oxygen
  • under standard conditions with all products and reactants in standard state
• standard enthalpy change of neutralisation (∆neutHΦ)

  • enthalpy change when
  • one mole of water is formed from the reaction of an acid and a base
  • under standard conditions with all products and reactants in standard state
• energy change is calculated by the equation:

  q = m c Δ t
• q = m c Δ t

  • q = heat energy transferred to the surrounding in joules
  • m = mass of material changing temperature in grams
  • c = specific heat capacity in J g-1 K-1
  • specific heat capacity of water = 4.18 Jg-1K-1
  • Δt = temperature change in kelvin
  • 0 °C = 273 K
  • a change of 50 °C = a change of 50 K
• a change of 50 °C = a change of 50 K
• 0 °C = 273 K
• average bond enthalpy - the energy required to break one mole of a specified type of bond in a gaseous molecule
• units of average bond enthalpy - kJmol-1
• the larger the value of the average bond enthalpy, the stronger the bond
• limitations of average bond enthalpies:

  • the actual bond enthalpy can vary depending on the chemical environment of the bond
  • C-H in CH4 = +439
  • C-H in C2H6 = +420
  • C-H in C3H8 = +422 and +411
  • average = 413 kJ mol-1
• the energy required to break bonds is endothermic (ΔH is positive)
• the energy released when bonds are made is exothermic (ΔH is negative)
• ΔH is exothermic (negative) if less energy is required to break the bonds in the reactants than is released when making the bonds in the products
• ΔH is endothermic (positive) if more energy is required to break the bonds in the reactants than is released when making the bonds in the products
• explain this graph in terms of making and breaking bonds

  A) breaking

  B) making

  C) exothermic

  
• explain this graph in terms of making and breaking bonds

  A) breaking

  B) making

  C) endothermic

  
• ΔH = Σ(bond enthalpies in reactants) – Σ(bond enthalpies in products)
• why might the enthalpy change calculated from average bond enthalpies be slightly different?

  average bond enthalpies are used, not actual

  the species might not be in standard conditions
• ΔH = break – make
• energy cannot be created or destroyed