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
• Exothermic reaction


  • The products have less energy than the reactants
  • The ∆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
• Endothermic reaction


  • The products have more energy than the reactants
  • The ∆H is positive
• 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
• Standard enthalpy change of formation

  • Mg (s) + Cl2 (g) MgCl2 (s)
  • 2Fe (s) + 1.5 O2 (g) Fe2O3 (s)
• 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
• Standard enthalpy change of combustion

  • CH4 (g) + 2O2 (g) CO2 (g) + 2 H2O (l)
• Incomplete combustion will lead to soot (carbon), carbon monoxide and water. It will be less exothermic than complete combustion.
• Activation energy


  • The energy barrier that must be overcome for a reaction to occur
• As temperature increases


  The rate of reaction increases
• Calorimetric method


  Measuring the enthalpy change for a reaction experimentally by mixing substances in an insulated container and measuring the temperature rise
• 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, and density of solution is taken to be the same as water.
• Calculating the enthalpy change of reaction, ΔH from experimental data

  1. Using q = m x cp x ΔT 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 ΔH

  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.
• We also assume that the solutions have the density of water, which is 1g cm-3. Eg 25 cm3 will weigh 25 g
• Hess's law states that total enthalpy change for a reaction is independent of the route by which the chemical change takes place
• On an energy level diagram the directions of the arrows can show the different routes a reaction can proceed by


  a = ΔH + b
• Interconnecting reactions can also be shown diagrammatically


  a + ΔH = c + d
• H2O is gas, not liquid, in this experiment
• Errors in this method
• Hess's law is a version of the first law of thermodynamics, which states that energy is always conserved
• Hess's law

  1. 2H (g) + 2Cl(g) → H2 + Cl2 → 2HCl (g)

  2. H+ (g) + Br - (g) → H+ (aq) + Br - (aq) → H (g) + Br (g) → HBr (g)
• CuSO4 (aq) → CuSO4 (s) + 5H2O (l) → CuSO4.5H2O (s)
• H reaction

  • 66.1 - 11 = -77.1 kJmol-1
• Hess's law is used to work out the enthalpy change to form a hydrated salt from an anhydrous salt
• Hess's law cycles are used to measure the enthalpy change for a reaction that cannot be measured directly by experiments
• 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
• Reactants and Products
• Combustion Products
• Mean Bond energies

  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
• Mean bond energies are used because every single bond in a compound has a slightly different bond energy
• Using mean bond energies to calculate enthalpy changes

  H = Σ bond energies broken - Σ bond energies made