# How would you balance: P4 + NO3- --> H2PO4- + NO?

Nov 25, 2015

By the standard redox reactions: oxidation and reduction half equations.

#### Explanation:

Oxidation of elemental phosphorus (${P}^{0} \rightarrow {P}^{V}$):

$\frac{1}{4} {P}_{4} + 4 {H}_{2} O \rightarrow {H}_{2} P {O}_{4}^{-} + 6 {H}^{+} + 5 {e}^{-}$ $\left(i\right)$

Reduction of nitrate anion to nitrous oxide (${N}^{+ V} \rightarrow {N}^{+ I I}$):

$N {O}_{3}^{-} + 4 {H}^{+} + 3 {e}^{-} \rightarrow N O + 2 {H}_{2} O$ $\left(i i\right)$

We add $3 \times \left(i\right) + 5 \times \left(i i\right)$: (why do I choose these multiples?)

$\frac{3}{4} {P}_{4} + 5 N {O}_{3}^{-} + 2 {H}^{+} + 2 {H}_{2} O \rightarrow 5 N O \uparrow + 3 {H}_{2} P {O}_{4}^{-}$

Is this equation balanced? How do you know? This is not a reaction I would choose to do: these are nasty reactants and nasty products, which would really get up your nostrils.

If you have been having trouble with these redox reactions, it might seem that I have been pulling reactants and products out of my backside. But looking at the starting conditions, you have the starting materials and the products already specified. To balance the reactions, all I needed to do was (i) balance the mass (for every reactant particle there is a corresponding product particle), and (ii) balance the charge. Chemical reactions always preserve mass and charge.