Combustion Calorimetry
Authors: O'Hare P. A. G. , Ribeiro Maria D. M. C. , Ribeiro Manuel A. V. 
Published: October 2012

In the absence of direct experimental information, one often must rely on knowledge of thermodynamic properties to predict the chemical behavior of a material under operating conditions. For such applications, the standard molar Gibbs energy of formationis particularly powerful; it is frequently derived by combining the standard molar entropy of formationand the standard molar enthalpy of formation, often as functions of temperature and pressure. It follows, therefore, that the standard molar enthalpy of formation,, is among the most valuable and fundamental thermodynamic properties of a material. This quantity is defined as the enthalpy change that occurs upon the formation of the compound in its standard state from the component elements in their standard states, at a designated reference temperature, usually (but not necessarily) 298.15 K, and at a standard pressure, currently taken to be 100 kPa.

Many methods have been devised to obtainexperimentally. Those include the so‐called second‐ and third‐law treatments of Knudsen effusion and mass‐spectrometric results from high‐temperature vaporization observations, as well as EMF results from high‐temperature galvanic cell studies (Kubaschewski et al., 1993). Each of those techniques yieldsfor the studied process at a given temperature. To derive values of, the high‐temperature results formust be combined with auxiliary thermodynamic information (heat capacities and enthalpy increments at elevated temperatures). Frequently, it turns out that the latter properties have not been measured, so they must be estimated (Kubaschewski et al., 1993), with a consequent degradation in the accuracy of the derived. Furthermore, high‐temperature thermodynamic studies of the kind outlined here often suffer from uncertainties concerning the identities of species in equilibrium. Another potential source of error arises from chemical interactions of the substances or their vapors with materials of construction of Knudsen or galvanic cells. All in all, these approaches do not, as a rule, yield precise values of. Sometimes, however, they are the only practical methods available to the investigator.

Alternative procedures involve measurements of the enthalpies (or enthalpy changes) of chemical reactions,of a substance and the elements of which it is composed. Experimentally, the strategy adopted in such determinations is dictated largely by the chemical properties of the substance and by the equipment available in the laboratory. The most obvious approach involves direct combination of the elements, an example of which is the combustion of gaseous H2in O2to form H2O. In the laboratory, however, it is often impossible under experimental conditions to combine the elements to form a particular compound quantitatively. An alternative, albeit less direct route has been used extensively for the past century or more, and involves measurements of the enthalpy change associated with a suitable chemical reaction of the material of interest. Here, the compound is “destroyed” rather than formed. Such chemical reactions have involved,inter alia, dissolution in a mineral acid or molten salt (Marsh and O'Hare, 1994) and the construction of an appropriate thermochemical cycle, thermal decomposition to discrete products (Gunn, 1972), or combustion in a gas such as oxygen or a halogen. This last method is the subject matter of this article.