When something like magnesium nitride forms, you have to supply all the energy needed to form the magnesium ions as well as breaking the nitrogen-nitrogen bonds and then forming N 3 - ions. Nitrogen is fairly unreactive because of the very large amount of energy is required to break the triple bond joining the two atoms in the nitrogen molecule, N 2. Nitrogen is often thought of as being fairly unreactive, and yet all these metals combine with it to produce nitrides, X 3N 2, containing X 2 + and N 3 - ions. Why do these metals form nitrides on heating in air? In this case, though, the effect of the fall in the activation energy is masked by other factors - for example, the presence of existing oxide layers on the metals, and the impossibility of controlling precisely how much heat you are supplying to the metal in order to get it to start burning. The activation energy will fall because the ionization energies of the metals fall. You could argue that the activation energy will fall as you go down the Group and that will make the reaction go faster. The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy. But how reactive a metal seems to be depends on how fast the reaction happens (i.e., Kinetics) - not the overall amount of heat evolved (i.e., Thermodynamics). If anything, there is a slight tendency for the amount of heat evolved to decrease as you go down the Group. The overall amount of heat evolved when one mole of oxide is produced from the metal and oxygen also shows no simple pattern: While it would be tempting to say that the reactions get more vigorous as you go down the Group, but it is not true. There are no simple patterns in the way the metals burn.
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