Question

The Ostwald process for the commercial production of nitric acid involves the following three steps: $$4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g)$$ $$2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)$$ $$3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(a q)+\mathrm{NO}(g)$$ a. Which reactions in the Ostwald process are oxidation–reduction reactions? b. Identify each oxidizing agent and reducing agent.

Short answer

a. The first two reactions in the Ostwald process are oxidation-reduction reactions. b. In the first reaction, NH₃ is the reducing agent, and O₂ is the oxidizing agent. In the second reaction, NO is the reducing agent, and O₂ is the oxidizing agent.

Step-by-step solution

Assign oxidation states to each element

We will assign oxidation states to all elements in the reactions. To do this, remember the following rules: 1. The oxidation state of an element in its elemental state is 0 (e.g., O₂ or N₂). 2. The more electronegative element is assigned its typical oxidation state (e.g., -2 for oxygen in most compounds, except peroxides). 3. The remaining elements will have the oxidation state that preserves the overall charge neutrality. Here are the oxidation states of elements in each reaction: \( 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g) \) (N: -3, H: +1, O: 0) -----> (N: +2, O: -2, H: +1, O: -2) \( 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g) \) (N: +2, O: -2, O: 0) -----> (N: +4, O: -2, O: -2) \( 3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(a q)+\mathrm{NO}(g) \) (N: +4, O: -2, O: -2, +1, H: +1, O: -2) -----> (H: +1, N: +5, O: -2, O: -2, O: -2, N: +2, O: -2)

Identify the oxidation-reduction reactions

The first two reactions in the Ostwald process involve changes in oxidation states for elements and are therefore oxidation-reduction reactions. The third reaction does not have any changes in oxidation states, so it is not an oxidation-reduction reaction. \( 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g) \) - Oxidation-reduction reaction \( 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g) \) - Oxidation-reduction reaction \( 3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(a q)+\mathrm{NO}(g) \) - Not an oxidation-reduction reaction The answer to part a is that the first two reactions are oxidation-reduction reactions.

Identify the oxidizing and reducing agents

In the first reaction, the oxidation state of nitrogen increases from -3 to +2, which means ammonia (NH₃) is being oxidized. Thus, it is the reducing agent. Oxygen's oxidation state decreases from 0 to -2, indicating that molecular oxygen (O₂) is being reduced. Therefore, oxygen is the oxidizing agent. In the second reaction, the oxidation state of nitrogen increases from +2 to +4, which means nitric oxide (NO) is being oxidized. Thus, it is the reducing agent. Oxygen's oxidation state decreases from 0 to -2, indicating that molecular oxygen (O₂) is being reduced, and it is the oxidizing agent. The answer to part b is: - In the first reaction, NH₃ is the reducing agent, and O₂ is the oxidizing agent. - In the second reaction, NO is the reducing agent, and O₂ is the oxidizing agent.

Key concepts

Oxidation-Reduction Reaction

Understanding oxidation-reduction reactions, or redox reactions, is essential in the context of the Ostwald process for nitric acid production. Redox reactions involve the transfer of electrons between chemical species. During the reaction, one substance gains electrons (is reduced), while the other loses electrons (is oxidized). To identify a redox reaction, look for changes in the oxidation states of elements. If the oxidation state of some elements changes from reactants to products, it implies a redox process. In the Ostwald process:

• The first reaction, \( 4 \mathrm{NH}_{3}(g) + 5 \mathrm{O}_{2}(g) \rightarrow 4 \mathrm{NO}(g) + 6 \mathrm{H}_{2} \mathrm{O}(g) \), involves electron transfer. Nitrogen goes from an oxidation state of -3 to +2, and oxygen goes from 0 to -2, confirming a redox reaction.
• The second reaction, \( 2 \mathrm{NO}(g) + \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{NO}_{2}(g) \), also shows a change where nitrogen's oxidation state changes from +2 to +4, indicating it is again part of an oxidation-reduction reaction.
• The third step in the Ostwald process doesn't show any change in oxidation states, indicating it is not a redox reaction.

Oxidation States

Oxidation states help us track the gain and loss of electrons in chemical reactions. They are essentially the charge an atom would have if all bonds were ionic. Let's simplify how to determine these states.Steps for determining oxidation states:

• Elements in their pure form (like \( \mathrm{O}_2 \) or \( \mathrm{N}_2 \)) have an oxidation state of 0.
• For monoatomic ions, the oxidation state is the ion's charge. For instance, \( \mathrm{Na}^+ \) has an oxidation state of +1.
• In compounds, oxygen typically has an oxidation state of -2 (except in peroxides), and hydrogen is usually +1.
• The sum of oxidation states in a neutral molecule must equal zero, and in an ion, it equals the ion's charge.

In the Ostwald process:

• In \( \mathrm{NH}_3 \), nitrogen is at -3, and hydrogen is +1. This means nitrogen has gained electrons compared to its elemental state.
• Oxygen in \( \mathrm{O}_2 \) is 0, while in \( \mathrm{H}_2 \mathrm{O} \) or \( \mathrm{NO} \), it is found at -2, having gained electrons each time.

Note how changes in these values between reactants and products define redox reactions.

Oxidizing and Reducing Agents

Oxidizing and reducing agents are central to redox chemistry.

• An oxidizing agent accepts electrons and is reduced in the process itself, causing the other reactant to be oxidized.
• A reducing agent donates electrons and is oxidized itself, leading the other substance to be reduced.

In the context of the Ostwald process:

• In the first reaction, \( \mathrm{NH}_3 \) is the reducing agent as it loses electrons (is oxidized) while turning nitrogen from -3 to +2. \( \mathrm{O}_2 \) acts as the oxidizing agent as it accepts electrons (is reduced) changing oxygen from 0 to -2.
• In the second reaction, \( \mathrm{NO} \) is the reducing agent because it is oxidized (from +2 to +4), and \( \mathrm{O}_2 \) is the oxidizing agent as it is reduced to form \( \mathrm{NO}_2 \).

These roles are pivotal in redox reactions, as they drive the transfer of electrons essential for chemical changes.