Question

Balance the following equation:
 $ N{H_3} + {O_2} \to NO + {H_2}O $

Answer 1

Hint: To answer this question you must recall the steps for balancing a chemical equation. A balanced equation is that equation in which the number of atoms of each element are equal on the either side of the arrow.

Complete step by step solution
In the given reaction, the oxidation state of nitrogen changes from $ - 3 $ to $ + 2 $ . The oxidation states of the rest of the elements, i.e oxygen and hydrogen remain the same in the reaction. Hence the given reaction is not a reduction- oxidation (redox) reaction. It is an oxidation reaction. Ammonia is oxidized or burned to produce nitrogen monoxide.
First, we write the given unbalanced equation, known as the skeleton equation of the given chemical reaction. You must be familiar to all the reactants and products.
 $ N{H_3} + {O_2} \to NO + {H_2}O $
We balance the equation by hit and trial method. The number of hydrogen atoms is not balanced in the equation. Balancing it, we get,
 $ 2N{H_3} + {O_2} \to NO + 3{H_2}O $
Now, the number of nitrogen atoms must be balanced. The equation becomes,
 $ 2N{H_3} + {O_2} \to 2NO + 3{H_2}O $
The number of oxygen atoms in the reaction is unbalanced. Balancing, we get,
 $ 2N{H_3} + \frac{5}{2}{O_2} \to 2NO + 3{H_2}O $
The above equation is completely balanced, but fractional coefficients are not preferably written. So, we write,
 $ \Rightarrow $ $ 4N{H_3} + 5{O_2} \to 4NO + 6{H_2}O $ .

Note
This reaction of ammonia with oxygen to give nitrogen monoxide is an exothermic reaction and a large amount of heat is evolved which further encourages the reaction to continue producing nitrogen monoxide readily. This reaction is the first step of the Ostwald’s process for the production of nitric acid from ammonia. The reaction takes place in presence of catalysts like platinum and rhodium because of high efficiency.