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2.1Unit 2. Energetics, Group Chemistry, Halogenoalkanes and Alcohols

Energetics — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

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2.1 Energetics: enthalpy diagrams, definitions, calorimetry, Hess cycles and bond enthalpies — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Draw enthalpy level and reaction profile diagrams, define standard enthalpy changes, process calorimetry data, apply Hess's law and use bond enthal…

  • Exothermic reaction
  • Endothermic reaction
  • Activation energy, Ea
  • Transition state
  • Standard enthalpy of formation, ΔfH
  • Standard enthalpy of combustion, ΔcH

Part 1 of 10Enthalpy level diagrams

Oxidation reduction oil rig
Oxidation reduction oil rig
Chlorine hot alkali disproportionation
Chlorine hot alkali disproportionation
Enthalpy
  • Enthalpy the total chemical energy (heat content) inside a substance; ΔH is the enthalpy change.
Exothermic reaction
  • Exothermic reaction products have less energy than reactants, heat is given out and ΔH is negative.
  • The temperature of the surroundings rises.
  • Exothermic reactions are thermodynamically possible, but a very slow rate means they may not occur (kinetic control).
Endothermic reaction
  • Endothermic reaction products have more energy than reactants, heat is absorbed and ΔH is positive.
  • The temperature of the surroundings falls.
State symbols must be given because a change of state changes the enthalpy: NaCl(s) → Na+(aq) + Cl−(aq) is +4 kJ mol−1,
  • but NaCl(g) → Na+(g) + Cl−(g) is +500 kJ mol−1.
The system is the reacting substances; the surroundings are everything else, e.g. the flask.

Part 2 of 10Reaction profile diagrams

A reaction profile is like an enthalpy level diagram but also shows the activation energy.
Activation energy, Ea
  • Activation energy, Ea the minimum energy reactant molecules need for a successful collision and to start the reaction.
  • It is the difference between the reactant level and the peak of the curve.
Transition state
  • Transition state the peak of the curve, where bonds are partly broken and partly formed.
  • It is very unstable, cannot be isolated and is higher in energy than reactants and products.
In exothermic reactions the reactants are closer in energy to the transition state, so Ea is lower than for endothermic reactions.
Activation energy is not shown on enthalpy level diagrams.

Part 3 of 10Standard conditions and enthalpy definitions

Standard conditions and enthalpy definitions
Standard conditions
  • A pressure of 100 kPa (101 kPa in older papers).
  • A temperature of 298 K (25 °C).
  • Each substance in its standard physical state.
  • The symbol ⦵ shows standard conditions.
Standard enthalpy of formation, ΔfH
  • Standard enthalpy of formation, ΔfH the enthalpy change when one mole of a compound forms from its elements in their standard states.
Standard enthalpy of combustion, ΔcH
  • Standard enthalpy of combustion, ΔcH the enthalpy change when one mole of a substance burns completely in oxygen.
Standard enthalpy of neutralisation, ΔneutH
  • Standard enthalpy of neutralisation, ΔneutH the enthalpy change when an acid and an alkali react to form one mole of water.
Standard enthalpy of reaction, ΔrH
  • Standard enthalpy of reaction, ΔrH the enthalpy change for the molar quantities shown in the equation.
Standard enthalpy of atomisation, ΔatH
  • Standard enthalpy of atomisation, ΔatH the enthalpy change when one mole of gaseous atoms forms from an element in its standard state.
ΔfH of an element in its standard state is zero, e.g. O2(g).
An enthalpy symbol needs Δ (change), H (enthalpy), ⦵ (standard) and a subscript for the type (c, f, neut, r).
Enthalpy values scale with the equation: doubling the moles doubles ΔH.

Part 4 of 10Calorimetry for reactions in solution

React the solution with an excess of one reagent in a polystyrene cup (or vacuum flask or metal can) and measure the temperature change.
Energy transferred
  • Energy transferred q = mcΔT.
    What the symbols mean: E, Q or q = energy transferred (J) · m = mass · c = specific heat capacity (J kg−1 °C−1; water 4.18 J g−1 °C−1) · ΔT or Δθ = temperature change (°C or K)
    • (kg; g in chemistry calorimetry)
  • The specific heat capacity of water is 4.18 J g−1 K−1.
  • Convert q from J to kJ.
Enthalpy change per mole
  • ΔH = −q ÷ n, where n is the moles of the limiting reagent;
    • a temperature rise gives a negative ΔH and a fall gives a positive ΔH.
Sample method for zinc and copper(II) sulfate
  • Measure 25 cm3 of 1.0 mol dm−3 copper(II) sulfate into the cup.
  • Weigh about 6 g of zinc powder, which is in excess.
  • Record the temperature every half minute for 2.5 minutes while stirring.
  • At 3 minutes add the zinc without recording that reading.
  • Keep stirring and record for a further 6 minutes.
Assumptions
  • The solution has the specific heat capacity and density of water (4.18 J g−1 K−1, 1 g cm−3).
  • The heat capacity of the container is ignored.
  • The reaction is complete and heat losses are negligible.
Sources of error include incorrect masses or volumes and heat loss; the temperature uncertainty is doubled because ΔT uses two readings.

Part 5 of 10Temperature correction graphs

  • A slow reaction loses heat before the maximum temperature is reached, so the true maximum is never measured.
1
Record the temperature of the first reagent for a few minutes to get a steady value.
2
Add the second reagent and keep recording temperature and time.
3
Plot the graph and extrapolate the cooling section back to the time the second reagent was added.
4
Use the extrapolated lines to read the temperature change ΔT.
Temperature correction graphs
Use q = mcΔT with the mass of the total solution, then divide by the moles of the limiting reagent.
If the reagents start at different temperatures, work out the energy for each separately and add.

Part 6 of 10Calorimetry for enthalpy of combustion

1
The heat from burning a fuel raises the temperature of a known mass of water in a simple calorimeter.
2
Record
  • The starting and final temperature of the water.
  • The mass of the spirit burner before and after, to find the mass of fuel burned and its moles.
3
Not all the heat reaches the water.
  • Some is lost to the surroundings and some is absorbed by the calorimeter.
  • Minimise loss by keeping the calorimeter close to the flame, using a lid and draught shields.
4
The main sources of error are heat loss and incomplete combustion.
5
Combustion is exothermic, so ΔcH is negative.

Part 7 of 10Hess's law and Hess cycles

1Hess's law
  • Hess's law the total enthalpy change is independent of the route taken, as long as the initial and final conditions are the same.
2
It follows from conservation of energy and is used for enthalpy changes that cannot be measured directly, e.g. ΔfH of propane.
3From enthalpies of formation
  • ΔrH = ΣΔfH(products) − ΣΔfH(reactants).
  • Write the equation, put the elements with state symbols underneath, draw arrows up to each compound and multiply by the moles.
  • Go from reactants to products, reversing the sign when the route goes against an arrow.
4From enthalpies of combustion
  • ΔrH = ΣΔcH(reactants) − ΣΔcH(products).
  • Write the formation equation, put the combustion products underneath and draw arrows down to them.
  • Balance the atoms of each element in the cycle.
5
Keep each value in brackets so that no minus sign is lost.

Part 8 of 10Bond enthalpies

Bond enthalpies
Bond enthalpy (bond dissociation enthalpy)
  • Bond enthalpy (bond dissociation enthalpy) the energy needed to break a particular bond.
Breaking bonds is endothermic and making bonds is exothermic.
  • The energy released on forming a bond equals the energy needed to break it, with the opposite sign.
Overall enthalpy change
  • More energy released in bond making than needed for bond breaking gives an exothermic reaction with more stable products.
  • More energy needed for bond breaking gives an endothermic reaction.
Average bond enthalpy
  • Average bond enthalpy the energy needed to break one mole of bonds in gaseous molecules, averaged over similar compounds.
  • Bond energies vary with the other atoms in the molecule, so an average is used.
Bonds with low enthalpies break first and react at lower temperatures; bonds with high enthalpies need heat or a catalyst.
ΔrH = Σ(bonds broken) − Σ(bonds formed), using positive values for both sums.
Draw full displayed formulae so no bond is missed, and multiply by the coefficients in the balanced equation.

Part 9 of 10More on this topic

More on this topic
Bond breaking absorbs energy and bond making releases energy.
Using the formation-data rule for combustion data — the subtraction is reversed.

Part 10 of 10More from the specification

1Calculating Enthalpy Changes via Calorimetry
  • In experimental calorimetry where a combustion or neutralisation takes place, heat energy transferred is determined using q = mc Delta T,
  • allowing calculating enthalpy change per mole of limiting reactant from Delta H = -q / n.
2Enthalpy of Combustion of an Organic Liquid
  • Measuring the enthalpy change of combustion of an organic liquid involves burning a known mass in a spirit burner beneath a copper calorimeter containing water and accounting for heat capacity and heat losses.
    • (such as an alcohol)
3Hess's Law Calculations
  • Hess's law states that the total enthalpy change for a chemical reaction is independent of the route taken, enabling enthalpy calculations using standard enthalpy of formation (Delta Hf) or combustion (Delta Hc) cycles.
4Mean Bond Enthalpies and Reaction Energetics
  • Bond-breaking is an endothermic process absorbing energy while bond-making is exothermic,
  • so overall enthalpy change is calculated via Delta H = sum(bonds broken) - sum(bonds formed).

Key terms

TermMeaning
Enthalpy change (ΔH)The heat energy change at constant pressure, in kJ mol−1.
Activation energyThe minimum energy needed for a successful collision.
Transition stateThe unstable peak of a reaction profile where bonds are partly broken and formed.
Standard enthalpy of formationThe enthalpy change when one mole of a compound forms from its elements in their standard states.
Standard enthalpy of combustionThe enthalpy change when one mole of a substance burns completely in oxygen.
Hess's lawThe total enthalpy change is independent of the route taken.
Average bond enthalpyThe energy to break one mole of a bond averaged over similar gaseous compounds.

Avoid these mistakes

1
Confusing endothermic with bond breaking and exothermic with bond making — breaking bonds is endothermic.
2
Forgetting to convert q from J to kJ before dividing by moles.
3
Giving the wrong sign — a temperature rise means a negative ΔH.
4
Omitting state symbols — a change of state changes the enthalpy.
5
Forgetting the coefficients in bond enthalpy sums — multiply each bond by the number in the equation.

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