Reaction Feasibility (Edexcel International A Level Chemistry)

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Philippa

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Reaction Feasibility

Summary 

  • For ΔStotal to be positive and therefore the reaction feasible: 
    • Both ΔSsystem and ΔSsurroundings are positive 
    • ΔSsurroundings is positive and ΔSsystem is negative, but ΔSsurroundings > ΔSsystem
    • ΔSsurroundings is negative and ΔSsystem is positive, but ΔSsurroundings ΔSsystem 

Feasibility

  • Generally, entropy will increase in the order:
    • solid < liquid < gas
  • Therefore we can determine the change in entropy and the feasibility by considering the change in state of reactants to products 

  • Example 1: the reaction between magnesium and oxygen at 293 K is feasible

Mg (s) + O2 (g) → MgO (s) 

  • This reaction produces a solid from a solid and a gas. Therefore entropy of the system is negative (ΔSsystem < ΔSsurroundings)
  • But since the entropy of the surroundings is very large as this reaction is very exothermic 
  • This outweighs the entropy of the system, so the total entropy is positive therefore the reaction is spontaneous

  • Example 2: the reaction between ethanoic acid and ammonium carbonate 

2CH3COOH (aq) + (NH4)2CO(s) → 2CH3COONH4 (aq) + H2O (l) + CO(g) 

  • This reaction is endothermic, therefore ΔSsurroundings is negative 
  • However, a gas is produced from a solid and liquid, so ΔSsystem is positive
  • ΔSsystem > ΔSsurroundings and the reaction is spontaneous 

Gibbs free energy

  • We can also use Gibbs free energy to work out if a reaction is feasible or not (this is not required as a part of the course) 
  • The feasibility of a reaction is determined by two factors
    • The enthalpy and entropy change
  • The two factors come together in a fundamental thermodynamic concept called the Gibbs free energy (G)
  • The Gibbs equation is:

ΔG = ΔHreaction – TΔSsystem

    • The units of ΔGare in kJ mol1
    • The units of ΔHreactionare in kJ mol1
    • The units of T are in K
    • The units of ΔSsystem are in J K-1 mol1(and must therefore be converted to kJ Kmol1by dividing by 1000)

  • For a reaction to be feasible, ΔG must be equal or less than zero 

Temperature & feasibility

  • We can look at the the values for ΔH and ΔS to determine whether the reaction is spontaneous / feasible at a given temperature (T)
  • The Gibbs equation can explain what will affect the spontaneity / feasibility of a reaction for exothermic and endothermic reactions

Exothermic reactions

  • In exothermic reactions, ΔHreaction is negative
  • If the ΔSsystem is positive:
    • Both the first and second term will be negative
    • Resulting in a negative ΔG so the reaction is feasible
    • Therefore, regardless of the temperature, an exothermic reaction with a positive ΔSsystem will always be feasible

  • If the ΔSsystem is negative:
    • The first term is negative and the second term is positive
    • At very high temperatures, the –TΔSsystemꝋ will be very large and positive and will overcome ΔHreaction
    • Therefore, at high temperatures ΔGꝋ is positive and the reaction is not feasible

  • Since the relative size of an entropy change is much smaller than an enthalpy change, it is unlikely that TΔS > Δas temperature increases
  • These reactions are therefore usually spontaneous at normal conditions

The diagram shows under which conditions exothermic reactions are feasible

Endothermic reactions

  • In endothermic reactions, ΔHreaction is positive
  • If the ΔSsystem is negative:
    • Both the first and second term will be positive
    • Resulting in a positive ΔG so the reaction is not feasible
    • Therefore, regardless of the temperature, endothermic with a negative ΔSsystem will never be feasible

  • If the ΔSsystem is positive:
    • The first term is positive and the second term is negative
    • At low temperatures, the –TΔSsystemꝋ will be small and negative and will not overcome the larger ΔHreaction
    • Therefore, at low temperatures ΔGꝋ is positive and the reaction is not feasible
    • The reaction is more feasible at high temperatures as the second term will become negative enough to overcome the ΔHreaction resulting in a negative ΔG

  • This tells us that for certain reactions which are not feasible at room temperature, they can become feasible at higher temperatures
    • An example of this is found in metal extractions, such as the extraction if iron in the blast furnace, which will be unsuccessful at low temperatures but can occur at higher temperatures (~1500 oC in the case of iron)

The diagram shows under which conditions endothermic reactions are feasible

 Summary of factors affecting Gibbs free energy 

Worked example

The reaction between aluminium oxide and carbon is not feasible at room temperature. 

Al2O3 + 3C(s)  → 2Al(s) + 3CO2 (g)

Given that  ΔH = +1336 kJmol-1 and ΔS= +581 JK-1mol-1 , calculate the temperature at which the reaction becomes feasible.

Answer

  • As both ΔH and ΔS are positive, ΔG will become negative if TΔS > ΔH.
  • The temperature at which this reaction becomes spontaneous can be calculated. This will be when ΔG = 0.
  • If ΔG = 0, then T = ΔH / ΔS*
  • T = begin mathsize 14px style fraction numerator 1336 over denominator 0.581 end fraction end style
  • T = 2299 K

*Don't forget to convert this into kJ

Thermodynamic & Kinetic Stability

Key Terms

  • Thermodynamically stable
    • Depends upon whether or not a reaction is spontaneous
    • The reaction is feasible if the total entropy is positive so there is a reduction in overall entropy
  • Kinetic stability
    • Depends in the rate of reaction and activation energy (Ea)
    • Reactions which have a high Ewill not take place at room temperature 
  • For example, methane will not form carbon dioxide and water unless it is ignited

  • Another example is the decomposition of hydrogen peroxide at 298 K

H2O2 (l) → H2O (l) + ½O2 (g)

  • This reaction has a very large Ea so must be catalysed using manganese dioxide, MnO2
  • If the reaction was left for long enough, the hydrogen peroxide would eventually decompose, however the addition of the MnO2 allows the reaction to take place via an alternative route with a lower Ea

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Philippa

Author: Philippa

Expertise: Chemistry

Philippa has worked as a GCSE and A level chemistry teacher and tutor for over thirteen years. She studied chemistry and sport science at Loughborough University graduating in 2007 having also completed her PGCE in science. Throughout her time as a teacher she was incharge of a boarding house for five years and coached many teams in a variety of sports. When not producing resources with the chemistry team, Philippa enjoys being active outside with her young family and is a very keen gardener.