This question is derived from a question asked in my school test.

What happens when a magnesium ribbon is heated in air?

My first response was the formation of magnesium oxide $(\ce{MgO})$ when oxygen in air reacts with magnesium at a high temperature which can be expressed in the form of a chemical equation like this:

$$\ce{2 Mg(s) + O2(g) ->[\Delta] 2 MgO(s)},$$

but I was wondering if magnesium could react with any other gas in the air to form a compound with that gas and I found out that magnesium does react with nitrogen in the air to form magnesium nitride too:

$$\ce{3 Mg(s) + N2(g) ->[\Delta] Mg3N2(s)}.$$

What determines whether the heated magnesium ribbon will react with the oxygen in the atmosphere or the nitrogen in the atmosphere?

Two possibilities that I can think of are:

  • composition of the air;
  • temperature.

I don't think that composition is the answer because on average the atmosphere of Earth has more nitrogen than oxygen, so I think that the answer may be temperature.

I'd also like to know how the factor affects the chemical reaction on an atomic level.

  • 2
    $\begingroup$ Magnesium Is known to burn in oxygen, nitrogen and carbon dioxide, if heated enough $\endgroup$
    – Poutnik
    Commented Jun 13, 2020 at 17:10
  • 2
    $\begingroup$ I think this is one of the first methods of obtaining argon from air. $\endgroup$
    – fraxinus
    Commented Jun 13, 2020 at 21:41

5 Answers 5


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!).

  • $\begingroup$ Re "the reaction is thermodynamically and kinetically less favourable": Is the reaction with magnesium and nitrogen actually exergonic? It could be endergonic and the energy is taken from the reaction with oxygen going on at the same time. What are the actual numbers for the two reactions? $\endgroup$ Commented Jun 14, 2020 at 22:34
  • $\begingroup$ @PeterMortensen That's a good question. We can be certain that the reaction causes a reduction in entropy (a gas is turned into a solid), so now we can just check the enthalpy of formation. According to Heat Contents and Heat of Formation of Magnesium Nitride, Ind. Eng. Chem. 1949, 41, 9, 2027–2031, the enthalpy of formation of $\ce{Mg3N2}$ is $\mathrm{-110.2\ kcal\ mol^{-1}}$ in ambient conditions. So the reaction is decidedly exergonic even at low temperatures, and if anything becomes less favourable at high temperatures. $\endgroup$ Commented Jun 14, 2020 at 23:08
  • $\begingroup$ For comparison, the heat of formation of $\ce{MgO}$ appears to be approximately $\mathrm{-143\ kcal\ mol^{-1}}$ in ambient conditions. The difference is surprisingly smaller than I would have expected, even taking into account that the $\ce{Mg3N2}$ lattice is more strongly bound than the $\ce{MgO}$ lattice. $\endgroup$ Commented Jun 14, 2020 at 23:13
  • $\begingroup$ Actually, the article linked above also explicitly discusses the free energy of formation of $\ce{Mg3N2}$. The author estimates an entropy of formation using a semi-empirical equation, and then provides an approximate function of free energy of formation (in $\mathrm{cal\ mol^{-1}}$) versus temperature (in $\mathrm{K}$): $\mathrm{\Delta G(T)=-224,240+189.67T+5.00\times10^{-4}T^2+8.13T\ lnT}$. At ambient conditions, this corresponds to $\mathrm{\Delta G=-153.8\ kcal\ mol^{-1}}$. Therefore, the reaction is clearly exergonic. The formula also estimates the decomposition temperature of 1500 °C $\endgroup$ Commented Jun 14, 2020 at 23:31

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

  1. Encyclopedia of the Alkaline Earth Compounds by Richard C. Ropp
  2. 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.
  • $\begingroup$ Ammonia? Do you mean nitrogen? If not, how does ammonia enter the picture? $\endgroup$ Commented Jun 14, 2020 at 22:40
  • 2
    $\begingroup$ @PeterMortensen it was indeed ammonia. Magnesium nitride was first made by burning magnesium in ammonia atmosphere. Later, it was found out that magnesium nitride can also be made by burning magnesium in air at high temperature. $\endgroup$ Commented Jun 15, 2020 at 4:08
  • $\begingroup$ @peter magnesium could displace hydrogen from ammonia. Then you have magnesium nitride plus the displaced hydrogen. $\endgroup$ Commented Jan 1, 2022 at 12:17

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.


On 'heating' in the 'earth atmosphere' elemental magnesium in what form?

This is, in my assessment, actually quite a good chemistry essay type question, because required clarification on the chemistry is likely needed. This arises from noted ambiguities, as for example, with respect to the intensity of the heat applied, the form of Mg present (as a powder or sheet metal), and even the composition of the experiment's atmosphere.

In the case of minor heating in the presence of an 'earth atmosphere' per Wikipedia, to quote:

Magnesium occurs naturally only in combination with other elements, where it invariably has a +2 oxidation state. The free element (metal) can be produced artificially, and is highly reactive (though in the atmosphere, it is soon coated in a thin layer of oxide that partly inhibits reactivity – see passivation).

So one quick answer with mild conditions with bulk magnesium is simply a protective coating of MgO.

However, as we are dealing with an 'earth atmosphere', which can contain water vapor, and per the same source:

Magnesium reacts with water at room temperature, though it reacts much more slowly than calcium, a similar group 2 metal. When submerged in water, hydrogen bubbles form slowly on the surface of the metal – though, if powdered, it reacts much more rapidly.

So, especially powdered Magnesium in the presence of water vapor at room temperature could create some Magnesium hydroxide and hydrogen gas also per the reaction:

$\ce{Mg (s) + 2 H2O (l) -> Mg(OH)2 (s) + H2 (g)}$

Further, assuming sufficient heat and amenability of the form of Mg to ignition, as per the same source, to quote:

Magnesium is highly flammable, especially when powdered or shaved into thin strips, though it is difficult to ignite in mass or bulk. Flame temperatures of magnesium and magnesium alloys can reach 3,100 °C (5,610 °F),[16] although flame height above the burning metal is usually less than 300 mm (12 in).[17] Once ignited, such fires are difficult to extinguish, because combustion continues in nitrogen (forming magnesium nitride), carbon dioxide (forming magnesium oxide and carbon), and water (forming magnesium oxide and hydrogen, which also combusts due to heat in the presence of additional oxygen). This property was used in incendiary weapons...

which apparently adds to the list carbon and magnesium nitride (Mg3N2).

However, again in the presence of water vapor, the nitride is not stable possibly reacting as follows:

$\ce{Mg3N2(s) + 6 H2O(l) -> 3 Mg(OH)2(aq) + 2 NH3(g)}$

So, lastly add ammonia to the product list (for a grade of 'A').


Forming magnesium nitride comes with some subtleties. Consider these two points:

Yet experimental evidence for magnesium nitride is seen in Nicolau's answer.

The solution to this paradox is the quoted adiabatic temperature assumes a stoichiometric mixture. Deep in a large mound it is entirely possible that magnesium is in excess, thus acting as a heat sink, as well as the oxygen being burned out by the outer part of the mound. That is where you get the right temperature range and nitrogen-rich composition to get the nitride identified in the video.

Magnesium has also been reacted directly with nitrogen to form the nitride [1].


1. F. Zong, C.Meng, Z. Guo, F. Ji, H. Xiao, X. Zhang, J. Ma, H. Ma, "Synthesis and characterization of magnesium nitride powder formed by Mg direct reaction with N2", Journal of Alloys and Compounds 508 (1), 15 October 2010, Pages 172-176.


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