The combustion of the hydrocarbon shown below, the major component of petrol (gasoline) trivially known as ‘isooctane’, proceeds with ΔG = –1000 kJ mol⁻¹ at 298 K.

The table on p. 244 shows that even a ΔG of only –50 kJ mol⁻¹ gives rise to a huge equilibrium constant: –1000 kJ mol⁻¹ gives an equilibrium constant of 10175 (at 298 K), a number too vast to contemplate (there are ‘only’ about 1086 atoms in the observable universe). This value of ΔG (or the corresponding value for the equilibrium constant) suggests that isooctane simply could not exist in the presence of oxygen. Yet we put it into the fuel tanks of our cars every day—clearly something is wrong. Since isooctane can exist in an atmosphere of oxygen despite the fact that the equilibrium position really would be completely on the side of the combustion products, the only conclusion we can draw must be that a mixture of isooctane and oxygen cannot be at equilibrium. A small burst of energy is needed to reach equilibrium: in a car engine, the spark plug provides this energy and combustion occurs. Without this burst of energy, the petrol is stable and no combustion occurs (as you will ruefully be aware if you have ever tried to start a car with a fl at battery). The mixture of petrol and oxygen is said to be thermodynamically unstable with respect to the products of the reaction, CO₂ and H₂O, but kinetically stable. We can be certain that they are thermodynamically unstable because even if the same small energy burst were applied to the products CO₂ and H₂O, they would never convert back to petrol and oxygen. Kinetically stable means that although the mixture could convert to a more stable set of products, it doesn’t do so because an energy barrier separates it from those products. An energy level diagram for a reaction such as the combustion of isooctane is shown below. The products are more stable (lower in energy) than the reactants, but to become the products, the reactants have to overcome a barrier to reaction. This barrier is called the activation energy and is usually given the symbol Ea or ΔG+.

If a reaction cannot proceed until the reactants have sufficient energy to overcome the activation energy barrier, it is clear that, the smaller the barrier, the easier it will be for the reaction to proceed. Likewise, the more energy we give the starting materials in the form of temperature, the more likely it is that they will collide with sufficient combined energy to cross the activation energy barrier. Unlike equilibria, which can change in either direction, reaction rates always increase at higher temperatures. A word of warning, however: heating is not all good for the chemist—not only does it speed up the reaction we want, it will also probably speed up lots of other reactions that we don’t want, including perhaps decomposition of the product! We shall see how we can get round this, but fi rst we shall take a closer look at what determines how fast a reaction takes place.

مواضيع ذات صلة

Catalysis

Reaction intermediates

Why some reactions are done at low temperature

The route from reactants to products: the transition state

Some reactions are reversible on heating : cracking

Equilibrium constants vary with temperature

Enthalpy versus entropy—some examples

Energy, enthalpy, and entropy: ΔG, ΔH, and ΔS

Typical method for making an ester

How to make the equilibrium favour the product you want The direct formation of esters

A small change in ΔG makes a big difference in K

The sign of ΔG tells us whether products or reactants are favoured at equilibrium