Advanced level chemistry kinetics notes: Comparing thermodynamic and kinetic stability

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Kinetics-Rates Part 6


6.2 Kinetic stability v thermodynamic stability

Do all reactions happen spontaneously just because they are thermodynamically feasible?

Advanced A Level Kinetics Index

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6.2 Kinetics versus thermodynamic stability and reaction feasibility


6.2 Introductory points

  • Even before rates factors are considered, the feasibility of a reaction is governed thermodynamically, that is in order for a reaction to be able to occur the energy changes must be favourable.

  • In order for a reaction to feasible the overall entropy change ...

    • ΔSθtot must be >=0 and the overall free energy change ΔGθ must be <=0.

    • More on the thermodynamics of feasibility: entropy change  and  free energy change

  • BUT, however feasible a reaction might be thermodynamically, it does not necessarily mean it will happen spontaneously because of kinetic limitations.

  • The speed at which a reaction takes place depends on the many factors described here or on the GCSE rates notes page.

  • However, a feasible reaction that you might expect to take place, may not happen because the activation energy is so high that virtually no molecules have enough kinetic energy to change on collision.

  • This would be described as a kinetically stable mixture but thermodynamically unstable.

The thermochemistry & thermodynamics dealt with in section 6. have been re–written and extended via

Part 1 Thermochemistry – Enthalpy changes * Part 2 Born–Haber Cycle * Part 3 Entropy & Free Energy

6b. Kinetic stability/instability examples

  • A mixture of hydrogen and oxygen (e.g. in air at room temperature) is perfectly stable until a means of ignition, e.g. a lit splint, match or spark etc., is applied. Thermodynamically the mixture is highly unstable with a very negative  free energy ΔGθ and shouldn't exist!

  • However, the activation energy to break the strong H–H or O=O bonds is so high, that they 'happily' co–exist without reacting, because the particle collisions are not energetic enough to cause a reaction. Therefore the mixture is kinetically stable.

  • A high temperature from a match or spark etc., gives enough of the reactant molecules sufficient kinetic energy to overcome the activation energy on collision*.

    • 2H2(g) + 2O2(g) ==> 2H2O(l), ΔGθ = –237 kJmol–1   and   ΔHθ = –286 kJmol–1 

    • Note: A very negative ΔHθ is usually, but not necessarily, indicative of a negative free energy change.

    • *The transition metal palladium can reduce the activation energy so much that it catalyses the spontaneous combustion/combination of hydrogen and oxygen at room temperature!

    • Like the methane–chlorine reaction below, it is free radical chain reaction. The initiating energy produces the first free radicals.

  • A mixture of hydrogen or methane and chlorine is stable in the dark, but exposed to light (particularly ultra–violet) or a high temperature, the mixture explodes to form hydrogen chloride/chloromethane gas.

    • The strong H–H or C–H and Cl–Cl bonds ensure a high activation energy, but the absorption by chlorine (with the weakest bond) of a photon of light energy or high kinetic energy molecules at a high temperature, will start the Cl–Cl bond breaking and so initiating the fast and exothermic free radical chain reaction (explosive!).

    • H2(g) + Cl2(g) ==> 2HCl(g), ΔGθ = –191 kJmol–1   and   ΔHθ = –185 kJmol–1 

      • mechanism: initiation: Cl2 ==> 2Cl

        • propagation: Cl + H2 ==> HCl + H and H + Cl2 ==> HCl + Cl

          • termination: 2Cl ==> Cl2 or 2H ==> H2 or H + Cl ==> HCl

    • CH4(g) + Cl2(g) ==> CH3Cl(g) + HCl(g), ΔGθ = –103 kJmol–1   and   ΔHθ = –99 kJmol–1

  • Hydrogen peroxide is thermally unstable, its aqueous solution is kept in a brown bottle to avoid decomposition initiated by light. It has a decent shelf–life of a few weeks, i.e., reasonably kinetically stable in the short term, until manganese(IV) oxide powder is added and then the rapid exothermic reaction takes place!

    • 2H2O2(aq)  ==> 2H2O(l) + O2(g) 






The thermochemistry & thermodynamics dealt with in section 6. have been re–written and extended via

Part 1 Thermochemistry – Enthalpy changes

  Part 2 Born–Haber Cycle

  Part 3 Entropy & Free Energy

Advanced A Level Kinetics Index

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