The equation in the margin just above tells us that the sign and magnitude of the energy ΔG are the only things that matter in deciding whether an equilibrium goes in one direction or another. If ΔG is negative the equilibrium will favour the products and if ΔG is large and negative the reaction can go to completion. The table on p. 244 tells us that it is enough for ΔG to be only about –10 kJ mol⁻¹ to get complete reaction. But we haven’t yet considered what ΔG actually corresponds to physically. To do this we need to introduce our second equation. The free energy of a reaction, ΔG, is related to two other quantities, the enthalpy of reaction, ΔH, and the entropy of reaction, ΔS, by the equation:

ΔG = ΔH – TΔS

As before, T is the temperature of the reaction in kelvin. Enthalpy, H, is a measure of heat, and the change in enthalpy, ΔH, in a chemical reaction is the heat given out or taken up in that reaction. Reactions which give out heat are called exothermic, and have negative ΔH; reactions which take in heat are called endothermic and have positive ΔH. Since break ing bonds requires energy and making bonds liberates energy, the enthalpy change gives an indication of whether the products have more stable bonds than the starting materials or not. Entropy, S, is a measure of the disorder in the system, so ΔS represents the entropy difference—the change in disorder—between the starting materials and the products. More disorder gives a positive ΔS; less disorder a negative ΔS. So ΔG represents a combination of heat and disorder. But what does this mean for you as a chemist wanting to get a reaction to work the way you want it to? We know that for a favour able change (i.e. an equilibrium favouring products) ΔG must be negative—in fact the more negative the better, as this gives a larger equilibrium constant. Since ΔG = ΔH – TΔS, we get a large, negative ΔG most readily if:

1. ∆H is negative, i.e. the reaction is exothermic. and

2. ∆S is positive (and hence –T∆S is negative), i.e. the reaction becomes more disordered.

Of course, we can still get a negative ΔG from an endothermic reaction (i.e. from a positive ΔH) but only if the reaction products are more disordered than the starting materials; likewise, a reaction which becomes more ordered as it proceeds can still be favourable, but only if it is exothermic to compensate for the loss of entropy. Because of the factor T multiplying the entropy term, both the equilibrium constant K (which depends on ΔG) and the relative importance of the two quantities (ΔH and ΔS) will vary with temperature (entropy changes are more important at higher temperatures). We’ll now look at some examples to see how this works in practice.

مواضيع ذات صلة

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

How the equilibrium constant varies with the difference in energy between reactants and products

Stability and energy levels

Substitution of C=O for C=C: a brief look at the Wittig reaction

Cyanide will attack iminium ions: the Strecker synthesis of amino acids

Lithium aluminium hydride reduces amides to amines

Iminium ions can react as electrophilic intermediates

Secondary amines react with carbonyl compounds to form enamines

Iminium ions and oxonium ions