Every chemical reaction proceeds to a certain extent, but not necessarily to completion (all reactants converted into products). Each reaction has a specific equilibrium point, which is related to the relative free energy content of the reactants and products. To understand the principle of equilibrium, let’s consider the following example, which takes place inside most cells but which we can do in the lab. This is the interconversion of glucose 1-phosphate and glucose 6-phosphate, a rearrangement of a phosphate group from one position on the ring of carbon atoms to another:

Glucose 1-phosphate ⇌ glucose 6-phosphate

Imagine that we start out with an aqueous solution of glucose 1-phosphate that has a concentration of 0.02 M. (M stands for molar concentration; see Key Concept 2.4.) The solution is maintained under constant environmental conditions (25°C and pH 7). As the reaction proceeds to equilibrium, the concentration of the product, glucose 6-phosphate, rises from 0 M to 0.019 M, while the concentration of the reactant, glucose 1-phosphate, falls from 0.02 M to 0.001 M, at which point, equilibrium is reached (Figure 8.4). At equilibrium, the reverse reaction, from glucose 6-phosphate to glucose 1-phosphate, progresses at the same rate as the forward reaction.

At equilibrium, then, this reaction has a product-to-reactant ratio of 19:1 (0.019/0.001), so the forward reaction has gone 95 percent of the way to completion (“to the right,” as written above). This result is obtained every time the experiment is run under the same conditions.

The change in free energy (ΔG) for any reaction is related directly to its point of equilibrium. The further toward completion the point of equilibrium lies, the more free energy is released. In an exergonic reaction, ΔG is a negative number. The total value of ΔG also depends on the beginning concentrations of the reactants and products and other conditions such as temperature, pressure, and pH of the solution. Biochemists often calculate ΔG using standard laboratory conditions: 25°C, one atmosphere pressure, one molar (1 M) concentrations of the solutes, and pH 7. The standard free energy change calculated using these conditions is designated ΔG⁰ In our example of the conversion of glucose 1-phosphate to glucose 6-phosphate, ΔG⁰ = –1.7 kcal/mol, or –7.1 kJ/mol.

A large, positive ΔG for a reaction means that it does not proceed to the right (A → B). If the concentration of B is initially high relative to that of A, such a reaction runs “to the left” (A ← B), and at equilibrium nearly all of B is converted into A. A ΔG value near zero is characteristic of a readily reversible reaction: reactants and products have almost the same free energies.

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In Chapters 9 and 10 you will learn the biochemical conversions that extract energy from food and light. In turn, this energy is used to synthesize carbohydrates, lipids, and proteins. All of the chemical reactions carried out by living organisms are governed by the principles of thermodynamics and equilibrium.