Question

The commercial production of nitric acid involves the following chemical reactions: $$\begin{array}{rl}4{\mathrm{NH}}_3(g)+5{\mathrm{O}}_2(g)& ⟶4\mathrm{NO}(g)+6{\mathrm{H}}_2\mathrm{O}(g)\\ 2\mathrm{NO}(g)+{\mathrm{O}}_2(g)& ⟶2{\mathrm{NO}}_2(g)\\ 3{\mathrm{NO}}_2(g)+{\mathrm{H}}_2\mathrm{O}(l)& ⟶2{\mathrm{HNO}}_3(aq)+\mathrm{NO}(g)\end{array}$$ (a) Which of these reactions are redox reactions? (b) In each redox reaction identify the element undergoing oxidation and the element undergoing reduction.

Short answer

(a) The first and second reactions are redox reactions. (b) In the first reaction, N undergoes oxidation, and O undergoes reduction. In the second reaction, N undergoes oxidation, and O undergoes reduction.

Step-by-step solution

(a) Identifying redox reactions

: Examine each reaction to determine if it is a redox reaction. 1st Reaction: $4N{H}_3(g)+5O_2(g)\to 4NO(g)+6H_{2}O(g)$ Oxidation states: N in $N{H}_3$: $-3$ H in $N{H}_3$: $+1$ O in $O_2$: $0$ N in $NO$: $+2$ O in $NO$: $-2$ H in $H_{2}O$: $+1$ O in $H_{2}O$: $-2$ Here, we can see that the oxidation state of N changes from -3 to +2 and the oxidation state of O changes from 0 to -2. Thus, this is a redox reaction. 2nd Reaction: $2NO(g)+O_2(g)\to 2N{O}_2(g)$ Oxidation states: N in $NO$: $+2$ O in $NO$: $-2$ O in $O_2$: $0$ N in $N{O}_2$: $+4$ O in $N{O}_2$: $-2$ Here, the oxidation state of N changes from +2 to +4 while the oxidation state of O changes from 0 to -2. Thus, this is also a redox reaction. 3rd Reaction: $3N{O}_2(g)+H_{2}O(l)\to 2HN{O}_3(aq)+NO(g)$ Oxidation states: N in $N{O}_2$: $+4$ O in $N{O}_2$: $-2$ H in $H_{2}O$: $+1$ O in $H_{2}O$: $-2$ H in $HN{O}_3$: $+1$ N in $HN{O}_3$: $+5$ O in $HN{O}_3$: $-2$ N in $NO$: $+2$ O in $NO$: $-2$ Here, we see that the oxidation states of all elements do not change, and thus, this is not a redox reaction.

(b) Identifying elements undergoing oxidation and reduction

: For each redox reaction identified in (a), find the elements undergoing oxidation (increase in oxidation state) and reduction (decrease in oxidation state): 1st Reaction: N changes from -3 (in $N{H}_3$) to +2 (in $NO$), so N undergoes oxidation. O changes from 0 (in $O_2$) to -2 (in $NO$), so O undergoes reduction. 2nd Reaction: N changes from +2 (in $NO$) to +4 (in $N{O}_2$), so N undergoes oxidation. O changes from 0 (in $O_2$) to -2 (in $N{O}_2$), so O undergoes reduction. The 3rd reaction is not a redox reaction, so there is no oxidation or reduction.

Key concepts

Oxidation States

Understanding oxidation states is crucial for determining redox reactions. The oxidation state, often called an oxidation number, represents the degree of oxidation of an atom in a chemical compound. These numbers help in tracking the transfer of electrons during chemical reactions. In simple ions, the oxidation state is equal to the charge of the ion.

For neutral molecules, the sum of oxidation states for all atoms must equal zero. Consider ammonia ($N{H}_3$): nitrogen is in the $-3$ oxidation state, while each hydrogen atom has an oxidation state of $+1$. These oxidation states are essential for identifying redox reactions, where there is a change in oxidation states of certain elements.

• N in $N{H}_3$: $-3$
• H in $N{H}_3$: $+1$
• O in $O_2$: $0$
• N in $NO$: $+2$
• O in $NO$/$H_{2}O$: $-2$

When analyzing a chemical reaction, changes in these numbers indicate which elements are oxidized or reduced.

Chemical Reactions

Chemical reactions involve the transformation of reactants into products. Some reactions involve both oxidation and reduction processes, known as redox reactions. During a redox reaction, one element is oxidized (loses electrons) while another is reduced (gains electrons).

In the nitric acid production process, the first two reactions are redox reactions. For example, in the first reaction involving $N{H}_3$ and $O_2$, nitrogen goes from $-3$ to $+2$, showing oxidation as it loses electrons, while oxygen is reduced from $0$ to $-2$ as it gains electrons. Each reaction requires careful balancing to ensure conservation of mass and charge.

• Oxidation: Increase in oxidation state due to loss of electrons.
• Reduction: Decrease in oxidation state due to gain of electrons.

These changes are tracked using oxidation states, helping to identify which elements undergo changes during the reaction.

Nitric Acid Production

The production of nitric acid is a multi-step chemical process involving redox reactions. Nitric acid ($HN{O}_3$) is essential in various industries, including fertilizers and explosives. The process typically starts with the oxidation of ammonia ($N{H}_3$) to nitrogen monoxide ($NO$) in the presence of oxygen ($O_2$).

This reaction results in nitrogen monoxide, which is further oxidized to nitrogen dioxide ($N{O}_2$). Finally, nitrogen dioxide reacts with water to produce nitric acid and nitrogen monoxide. The step where nitrogen dioxide forms nitric acid does not involve a redox reaction since oxidation states remain unchanged.

Key steps include:

• Ammonia Oxidation: $4N{H}_3+5O_2\to 4NO+6H_{2}O$
• Nitrogen Monoxide Reoxidation: $2NO+O_2\to 2N{O}_2$
• Nitric Acid Formation: $3N{O}_2+H_{2}O\to 2HN{O}_3+NO$

These reactions illustrate the sequential transformations that lead to nitric acid, emphasizing the importance of understanding oxidation-reduction processes in industrial applications.