Magnesium nitride

It is said again and again that nitrogen reacts chemically “like a dead dog”. No wonders, why it is used as a protective gas in the food and chemical industry. Most often, it is used to prevent the entry of highly reactive oxygen.

If you still want to persuade it to react, you have to spend a lot of energy, for example in a thunderstorm lightning (so-called air combustion). Another option is to add a suitable catalyst, to reduce the activation energy. This is what the fertilizer industry does when it produces ammonia from nitrogen and hydrogen. This is the well-known Haber-Bosch process.

Nature, however, has found a much more elegant way: plants use the ability of many bacteria to bind the nitrogen from the air with the help of a sophisticated enzymatic reaction. The nitrogen made chemically available, is required by plants to build up their biomass – proteins, nucleotides and nucleic acids as well as many vitamins.

In the laboratory we can only force the nitrogen to react. Magnesium, which we got to know early on as a very good burning metal, helps us here. For us it is important: When magnesium is burned, it not only reacts with oxygen, but also with the nitrogen in the air.

           2 Mg + O2 ———> 2 MgO + Energy

           3 Mg + N2 ———> Mg3N2 + Energy

We show this with the following experiment.

Experiment: Burning magnesium shavings

Please read our disclaimer and inform yourself about the dangers.

We pile not too little magnesium shavings on a fire-proof base (e.g. on a brick) and ignite the top of the pile with the Bunsen burner. Put on protective goggles or sunglasses, do not look into the bright flame! At first, they start to burn surprisingly slowly, but the flame and heat take off after some time. Finally, we let the mass cool down well.

The white product is taken apart after cooling. The pile is pretty tough! Inside the heap we can see a yellowish to gray, crystalline mass: magnesium nitride.

We can prove the formation of nitride by allowing the reaction product to react with water. Ammonia (and not important for us: other nitrogen-hydrogen compounds as well) are released in the process.

Experiment 2: Investigation of magnesium nitride

We put a sample of the yellow mass in a beaker (100 ml). We cover this with a watch glass, on the underside of which we have previously put a damp litmus indicator paper. We’ll add some water to the sample. The mixture foams up. The indicator paper instantly turns deep blue. We do the smell test. (Caution! “Sniff” by fanning the gases!)

There is a pungent smell of ammonia. The indicator paper turns blue because ammonia forms a hydroxide ions with water.

           Mg3N2 + 6 H2O —> 3 Mg(OH)2 + 2 NH3

           NH3 + H2O —> NH4+ + OH

By the way, in the past this was actually used to produce ammonia. That was an expensive process, since magnesium had to be obtained beforehand by electrolysis!

The question remains why the pile is strangely stratified after burning: the white oxide is clearly visible on the outside, the yellow-gray nitride on the inside. And that although there is four times more nitrogen in the air than oxygen! The reason is: The reaction with oxygen is preferred because it releases more energy. Only when the oxygen is used up does the nitrogen get its chance – and that’s right in the middle of the pile.