Heating the nitrates (Group 1)
Most nitrates tend to decompose on heating to the metal oxide, brown fumes of nitrogen dioxide, and oxygen. For example, a typical Group 2 nitrate like magnesium nitrate decomposes this way:
2Mg(NO3)2(s)→2MgO(s)+4NO2(g)+O2(g)
In Group 1, lithium nitrate behaves in the same way, producing lithium oxide, nitrogen dioxide, and oxygen as shown:
4LiNO3(s)→2Li2O(s)+4NO2(g)+O2(g)
The other Group 1 nitrates, however, do not decompose completely at regular laboratory temperatures. They produce the metal nitrite and oxygen, but no nitrogen dioxide:
2XNO3(s)→2XNO2(s)+O2(g)
Each of the nitrates from sodium to cesium decomposes in this way; the only difference is in the temperature required for the reaction to proceed. For larger metals, the decomposition is more difficult and requires higher temperatures.
