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$\ce{HNO3}$ always reduces its anion to give nitrogen containing compounds ( $\ce{NO2}$, $\ce{NO}$, and even $\ce{N2O}$) when reacting with metals, but with $\ce{Mg}$ and $\ce{Mn}$ it releases $\ce{H2}$. Is there a specific reason for this variation?

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It's because of the activity of hydrogen ions in comparison to the metals.

Certain metals like $\ce{Mg}$, say, will be oxidized by the hydrogen ions. Note that $\ce{Mg^2+}$ has a reduction potential of -2.37V, which is less than $\ce{H+}$'s reduction potential of 0V (not accounting for overvoltage and what not). Hence, $\ce{H2}$ gas is produced spontaneously.

However, metals like $\ce{Cu}$ (which is the example given by your Wikipedia link) and $\ce{Ag}$ have a reduction potential greater than 0V. This means that H+ will no longer be able to oxidize these metals. Instead, the $\ce{NO3-}$ ion will serve as the oxidizing agent.

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Actually, it usually doesn't. It produces all types of products ranging from $\ce{NO2}$ to $\ce{NH4NO3}$ with exact products depending on concentration and conditions.

The reason for it is high oxidation power of nitric acid, exceeding that of hydrogen cations, so nitric acid reduces first. As result, to produce some nitrogen, the oxidation power must be brought below that of hydrogen cations. I'm not aware if it is actually possible, but possibly high dilution may help.

As for Mg, it reacts with hot water, so it's unclear if the reaction of $\ce{Mg}$ with diluted nitric acid counts. Not aware about $\ce{Mn}$ though.

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Because reaction of Metal+Acid= Salt + hydrogen $$\ce{Mg +2HNO3 -> Mg(NO3)2 +H2}$$

$\ce{Mg}$ has a charge of +2 while nitric acid ($\ce{HNO3}$) has a charge of -1

Metals always react with acids to form salt and hydrogen gas

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  • $\begingroup$ thats just the point! Nitric reacts with ALL metals to release NO, NO2,N2O BUT releases H2 with Mn and Mg. Why the VARIATION? $\endgroup$ – user14982 Mar 15 '15 at 10:30
  • $\begingroup$ It also releases hydrogen with other metals not just with Mn or Mg $\endgroup$ – hamza ahsan Mar 15 '15 at 10:35
  • $\begingroup$ actually, the wikipedia reference suggested by binary geek agrees w/ my facts, but doesn't supply a reason. $\endgroup$ – user14982 Mar 15 '15 at 18:39

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