Chemistry - On heating in the Earth's atmosphere, can magnesium react with nitrogen to form magnesium nitride?
Solution 1:
A large pile of grey magnesium powder, when lit in air, produces a smouldering pile which cools down to reveal a crusty white solid of magnesium oxide. However, if you break apart the mound, you can find something quite strange in the middle - a clearly brownish powder that wasn't there before.
Seeing is believing! The author of the video also has a clever idea to identify the brown solid. By adding water and placing some moist pH paper above the puddle, it clearly shows the transfer of some alkaline substance across the gap. This is ammonia gas, $\ce{NH3}$, whose presence is explained by the hydrolysis of magnesium nitride:
$$\ce{Mg3N2(s) + 6H2O(l) -> 3 Mg(OH)2(aq) + 2 NH3(g)}$$
It is important that the pH paper not come in direct contact with the water used to hydrolyze the magnesium oxide, as $\ce{Mg(OH)2}$ is itself also basic, and could also be formed by reaction with either $\ce{MgO}$ or $\ce{Mg}$ directly. Only $\ce{Mg3N2}$ produces a basic gas which forms an alkaline solution in water.
As you can see, magnesium metal does react directly with molecular nitrogen ($\ce{N2}$) when burned in air. However, the reaction is thermodynamically and kinetically less favourable than the reaction with molecular oxygen ($\ce{O2}$). This is almost certainly due to the extreme strength of the bond between nitrogen atoms in molecular $\ce{N2}$, whose bond dissociation energy of $\mathrm{945\ kJ\ mol^{-1}}$ is one of the strongest in all of chemistry, second only to the bond in carbon monoxide. For comparison, the bond dissociation energy of molecular $\ce{O2}$ is drastically lower, at $\mathrm{498\ kJ\ mol^{-1}}$.
So why did the Chem13 magazine article referenced in Aniruddha Deb's answer not find any magnesium nitride? It is likely that 1 g of magnesium metal is far too little for the experiment run under their conditions. It takes a significant amount of "sacrificial" magnesium to completely consume the oxygen in its surroundings. Only once practically all the oxygen is consumed (and while the pile of magnesium is still hot enough from the reaction between magnesium and oxygen) will the remaining magnesium metal react with the nitrogen in air. Alternatively, the reaction would have to be performed in an oxygen-free environment. Magnesium metal is such a strong reductant that many substances can act as an oxidant for it, including pure $\ce{CO2}$ (also shown in the video above) and water (never put out a magnesium fire with water!).
Solution 2:
I don't think that composition is the answer because on an average, the atmosphere of Earth has more Nitrogen than Oxygen, so I think that the answer may be temperature.
Indeed, temperature is an important factor for this reaction and the reaction carries at a specific temperature. It was intensively studied in the late 19th century$\ce{^{[2]}}$. It was predicted that the reaction begins at $\pu{450 ^\circ C}$ and proceeds most intensely at $\ce{600-700 ^\circ C}$ at atmospheric pressure of ammonia($\pu{1003 kPa}$ at $\pu{25 ^\circ C}$). The temperature dependence of the reaction was found to be parabolic in nature. Later, it was established that magnesium nitride can be formed by heating magnesium in air by means of a gas burner. Researchers at that time suggested that a relatively high temperature was needed for the reaction to get going specifically in the range of $\pu{700-900 ^\circ C}$. Interaction of magnesium with nitrogen started at $\pu{780-800 ^\circ C}$ and within 4-5 hrs at a temperature of $\pu{800-850^\circ C}$, the nitride is formed with nitrogen content of $\ce{27.3-27.6 {%}}$ which corresponded to theoretical nitrogen content in $\ce{Mg3N2}$ i.e $\ce{27.4{%}}$.
Notes and references
- Encyclopedia of the Alkaline Earth Compounds by Richard C. Ropp
- Discovered in 1854 by Saint-Claire Deville during a study of sublimation of magnesium in air. In 1885, it was synthesized by heating magnesium in the atmosphere of ammonia.
Solution 3:
Nicolau Saker Neto's answer does a more accurate job of answering the question. Do give that answer a read as well.
Exactly the same question was pubilshed in University of Waterloo's Chem13 magazine. More details can be found in the link but the conclusion was:
Given that no evidence of $\ce{Mg3N2}$ formation could be found, it appears that the hydration step is not necessary and only makes the experiment more difficult. Not only could no ammonia be detected by smell; within the precision of the electronic scale ($\pu{0.01 g}$) the results were consistent with pure $\ce{MgO}$ being the product.
NOTE: Wikipedia mentions a contrary result
In fact, when magnesium is burned in air, some magnesium nitride is formed in addition to the principal product, magnesium oxide.
However, Wikipedia does not seem to provide a citation for the same. In this case, I would believe in the first reference more than the Wiki article.