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Aluminum oxide, or A14C3, is formed when aluminum and oxygen ions in gaseous form react. In the reaction, two moles of Al3+ ions in the gaseous state react with three moles of O2- ions in the gaseous state. The resulting oxide is a dense solid crystalline form. Its free energy is 366.3 kcal mole-1.
Heat of formation is a function of the number of bonds formed. Metal oxides, like magnesium oxide, have a lower enthalpy of formation than iron oxide. When the amount of stored energy in the products is more than that in the reactants, the enthalpy change is positive. Usually, chemically "stable" compounds have a negative enthalpy change. This is because reactive substances have a tendency to react until they do. For example, hydrogen gas reacts explosively with gaseous chlorine Cl2. However, it is important to note that the enthalpy of combustion is also a measure of the amount of reaction.
Meichsner and Roth calculated the heat of formation of A14C3. They used a standard calorimeter and omitted water from the bomb. Their results suggest that the enthalpy of combustion of aluminum carbide is roughly -40 kcal mole-1. However, their calculations were not specific enough to determine the form of the A12O3 product.
Recent studies have taken into account the formation of A12OC and A14O4C. However, they have not considered thermal functions for these elements. Therefore, to calculate the heat of formation of A14C3 using equilibrium data, we need to estimate the thermal functions of both A12OC and A14O4C.
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