This is a revision of the previous version. See *Eur. J. Biochem.*, 1996, **240**, 1-14; **242**, 433; *Pure Appl. Chem.* 1994, **66**, 1641-1666 [Copyright IUBMB and IUPAC]. A PDF (1655 kB) of the printed 1994 version is available.

**Abstract**

Chemical equations are normally written in terms of specific ionic and elemental species and balance atoms of elements and electric charge. However, in a biochemical context it is usually better to write them with ionic reactants expressed as totals of species in equilibrium with each other. This implies that atoms of elements assumed to be at fixed concentrations, such as hydrogen at a specified pH, should not be balanced in a biochemical equation used for thermodynamic analysis. However, both kinds of equations are needed in biochemistry. The apparent equilibrium constant *K′* for a biochemical reaction is written in terms of such sums of species and can be used to calculate standard transformed Gibbs energies of reaction Δ_{r}*G′ ^{0}.* This property for a biochemical reaction can be calculated from the standard transformed Gibbs energies of formation Δ

- 1. Preamble
- 2. Introduction
- 3. Basic thermodynamics
- 4. Thermodynamics of chemical reactions
- 5. Legendre transform to introduce the pH as an independent variable in biochemical thermodynamics
- 6. Equations for the standard transformed formation properties of a reactant
- 7. Thermodynamics of biochemical reactions
- 8. Stoichiometry
- 8.1 Stoichiometry of chemical reactions
- 8.2 Stoichiometry of biochemical reactions

- 9. Standard apparent reduction potentials for half reactions of enzyme-catalyzed reactions
- 10. Thermodynamic tables
- 11. Relations between biochemical thermodynamics and enzyme kinetics
- 12. Calorimetric measurements involving biochemical reactions
- 13. Recommendations for reporting experimental results
- 14. List of symbols
- 15. References

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