# Why is the standard enthalpy of formation of aluminium oxide higher than magnesium oxide?

\begin{align} &\text{Aluminium oxide:} &\ce{4Al + 3O_2 &-> 2Al2O3} &(\Delta H_\mathrm f &= \pu{-1675 kJ/mol})\\[1.5em] &\text{Magnesium oxide:} &\ce{2Mg + O_2 &-> 2MgO} &(\Delta H_\mathrm f &= \pu{-602 kJ/mol})\\[1.5em] &\text{Iron (III) oxide:} &\ce{4Fe + 3O_2 &-> 2Fe2O3} &(\Delta H_\mathrm f &= \pu{-824 kJ/mol})\end{align}

Question:

1. Why is the enthalpy of formation for aluminium oxide higher than that of magnesium oxide, even though magnesium is higher in the reactivity series?
2. Why is enthalpy of formation for magnesium oxide lower than iron oxide even though its more reactive? Is this due to the number of bonds formed? Thus, magnesium is less exothermic as less bonds are formed?
3. What determines reactivity? Why is magnesium more reactive than aluminium and why is aluminium more reactive than iron? Does it have anything to do with valency of the electrons?

So for the aluminum oxidation you should divide all the coefficients by three to get a reaction with one mole of oxygen, and then the enthalpy of formation is similarly divided by three. You would then get $$-558\text{ kJ/mol }\ce{O2}$$ versus $$-602\text{ kJ/mol }\ce{O2}$$ for magnesium oxidation.