# Question 9701f

May 5, 2016

${\text{C"_ ((s)) + 4"HNO"_ (3(aq)) -> 2"H"_ 2"O"_ ((l)) + 4"NO"_ (2(aq)) + "CO}}_{2 \left(g\right)} \uparrow$

#### Explanation:

You're dealing with a redox reaction in which carbon reduces hot, concentrated nitric acid to nitrogen dioxide, ${\text{NO}}_{2}$, while being oxidized to carbon dioxide, ${\text{CO}}_{2}$, in the process.

The balanced chemical equation that describes this reaction looks like this

stackrel(color(blue)(0))("C")_ ((s)) + 4"H"stackrel(color(blue)(+5))("N")"O"_ (3(aq)) -> 2"H"_ 2"O"_ ((l)) + 4stackrel(color(blue)(+4))("N")"O"_ (2(aq)) + stackrel(color(blue)(+4))("C")"O"_(2(g))# $\uparrow$

Here carbon is being oxidized from a $\textcolor{b l u e}{0}$ oxidation state in elemental carbon to a $\textcolor{b l u e}{+ 4}$ oxidation state in carbon dioxide.

Nitrogen, on the other hand, is being reduced from a $\textcolor{b l u e}{+ 5}$ oxidation state in nitric acid to a $\textcolor{b l u e}{+ 4}$ oxidation state in nitrogen dioxide.