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Write an unbalanced equation for the following: Magnesium metal is added to dilute nitric acid.

My attempt: Nitric acid commonly oxidizes metals and gets itself reduced in the process - to $\ce{NO}$ if the solution is dilute, and to $\ce{NO2}$ if it is concentrated. Therefore, I thought that the reaction would proceed as follows: $\ce{Mg + HNO3 -> Mg^2+ + NO + H2O}.$

However, my book says that instead of being reduced to $\ce{NO},$ the nitric acid is actually reduced to ammonium in the following fashion: $\ce{Mg + HNO3 -> Mg^2+ + NH4^+ + H2O}$

Why is the reaction with dilute nitric acid different specifically for Magnesium (as compared to nitric acid reactions with other metals, such as copper) ?

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    $\begingroup$ There is nothing specific about Mg. Other active metals (Al, for instance) would behave the same way. Then again, in fact you'll get a bunch of products containing nitrogen in all oxidation states from -3 to +4; which of them would be the major one depends on many things. $\endgroup$ Oct 16, 2017 at 7:32
  • $\begingroup$ Reading ‘dilute’, I would have assumed that the proton reduces magnesium liberating hydrogen (and leaving the nitrate untouched). $\ce{NO}$ is produced by half-concentrated nitric acid as far as I can remember. $\endgroup$
    – Jan
    Oct 16, 2017 at 13:24
  • $\begingroup$ chemistry.stackexchange.com/questions/6531/… $\endgroup$
    – Mithoron
    Oct 16, 2017 at 14:00

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