Energetics

• If an enthalpy change occurs then energy is transferred between system and surroundings
• System

  The chemicals
• Surroundings

  Everything outside the chemicals
• Exothermic change


  Energy is transferred from the system to the surroundings
• The products have less energy than the reactants in an exothermic change
• In an exothermic reaction the ∆H is negative
• Endothermic change


  Energy is transferred from the surroundings to the system
• Endothermic changes require an input of heat energy
• The products have more energy than the reactants in an endothermic change
• In an endothermic reaction the ∆H is positive
• Enthalpy changes are normally quoted at standard conditions
• Standard conditions

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


  The enthalpy change when 1 mole of the compound is formed from its elements under standard conditions
• Symbol for standard enthalpy change of formation


  ΔfH
• Standard Enthalpy Change of Formation reactions

  • Mg (s) + Cl2 (g) → MgCl2 (s)
  • 2Fe (s) + 1.5 O2 (g) → Fe2O3 (s)
• 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
• Symbol for standard enthalpy change of combustion


  ΔcH
• Incomplete combustion will lead to soot (carbon), carbon monoxide and water
• Incomplete combustion is less exothermic than complete combustion
• Activation Energy (EA)

  The minimum energy required for a reaction to occur
• Common oxidation exothermic processes include the combustion of fuels and the oxidation of carbohydrates such as glucose in respiration
• Enthalpy change

  The amount of heat energy taken in or given out during any change in a system provided the pressure is constant
• Measuring the Enthalpy Change for a Reaction Experimentally

  1. Calorimetric method

  2. Take readings at regular time intervals

  3. Extrapolate the temperature curve

  4. Measure temperature of reactants before addition

  5. Use energy change equation
• If the reaction is slow then the exact temperature rise can be difficult to obtain as cooling occurs simultaneously with the reaction

  
• Energy change equation


  Q (J) = m (g) x cp (J g-1K-1) x ΔT (K)
• This equation will only give the energy for the actual quantities used
• General method for calorimetry

  1. Wash equipment

  2. Dry the cup

  3. Put polystyrene cup in a beaker

  4. Measure out desired volumes

  5. Clamp thermometer in place

  6. Measure initial temperatures

  7. Transfer second reagent to cup

  8. Stir mixture

  9. Record temperature every minute
• One type of experiment is one in which substances are mixed in an insulated container and the temperature rise measured
• Errors in calorimetric method

  • Energy transfer from surroundings
  • Approximation in specific heat capacity
  • Neglecting specific heat capacity of calorimeter
  • Reaction may be incomplete or slow
  • Density of solution assumed to be the same as water
• Calculating the enthalpy change of reaction, ΔH from experimental data

  1. Calculate energy change

  2. Work out moles of reactants

  3. Divide energy change by moles

  4. Add sign and unit
• 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
• Also assume that the solutions have the density of water, which is 1 g cm-3
• Example calculation steps

  1. Calculate energy change

  2. Calculate moles of CuSO4

  3. Calculate ΔH
• 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
• 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
• On an energy level diagram, the directions of the arrows can show the different routes a reaction can proceed by