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Chapter 4: Problem 96

The commercial production of nitric acid involves the following chemical reactions: $$ \begin{aligned} 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) & \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g) \\ 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) & \longrightarrow 2 \mathrm{NO}_{2}(g) \\ 3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) & \longrightarrow 2 \mathrm{HNO}_{3}(a q)+\mathrm{NO}(g) \end{aligned} $$ (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

Expert verified
(a) All three reactions are redox reactions. (b) In Reaction 1, \(N\) is oxidized and \(O\) is reduced. In Reaction 2, \(N\) is oxidized and \(O\) is reduced, and in Reaction 3, \(N\) undergoes both oxidation and reduction.

Step by step solution

01

Assign Oxidation Numbers

First, let's assign oxidation numbers to each element in each reaction. Reaction 1: \(4 \,\text{NH}_3(g) + 5 \,\text{O}_2(g) \longrightarrow 4\, \text{NO}(g) + 6 \,\text{H}_2\text{O}(g)\) \(N: -3, +2; H: +1; O: 0, -2\) Reaction 2: \(2\, \text{NO}(g) + \text{O}_2(g) \longrightarrow 2\, \text{NO}_2(g)\) \(N: +2, +4; O: 0, -2\) Reaction 3: \(3 \,\text{NO}_2(g) + \text{H}_2\text{O}(l) \longrightarrow 2 \,\text{HNO}_3(aq) + \text{NO}(g)\) \(N: +4, +5, +2; H: +1; O: -2\)
02

Identify Redox Reactions

Now that we have assigned oxidation numbers, let's find out which of these reactions are redox reactions by looking for changes in oxidation numbers. Reaction 1: \(N\) changes from -3 to +2 (oxidation), and \(O\) changes from 0 to -2 (reduction). So, Reaction 1 is a redox reaction. Reaction 2: \(N\) changes from +2 to +4 (oxidation), and \(O\) changes from 0 to -2 (reduction). So, Reaction 2 is also a redox reaction. Reaction 3: \(N\) changes from +4 to +5 and +2 (both oxidation and reduction). So, Reaction 3 is a redox reaction as well. (a) All three reactions are redox reactions.
03

Identify Elements Undergoing Oxidation and Reduction

Let's now identify the elements undergoing oxidation and reduction in each redox reaction. Reaction 1: Oxidation: \(N\) in NH3, as oxidation number increases from -3 to +2. Reduction: \(O\) in O2, as oxidation number decreases from 0 to -2. Reaction 2: Oxidation: \(N\) in NO, as oxidation number increases from +2 to +4. Reduction: \(O\) in O2, as oxidation number decreases from 0 to -2. Reaction 3: Oxidation: \(N\) in NO2, as oxidation number increases from +4 to +5 (in HNO3) Reduction: \(N\) in NO2, as oxidation number decreases from +4 to +2 (in NO) (b) In Reaction 1, \(N\) is oxidized and \(O\) is reduced. In Reaction 2, \(N\) is oxidized and \(O\) is reduced, and in Reaction 3, \(N\) undergoes both oxidation and reduction.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Understanding Oxidation Numbers
Oxidation numbers are like scorekeepers in chemical reactions. They help us keep track of how many electrons are lost or gained by elements. For any element in its pure form, the oxidation number is zero. However, in compounds, this number can vary based on the rules of chemistry.
  • For hydrogen, it's usually +1.
  • Oxygen typically has an oxidation number of -2.
  • For nitrogen, it can change depending on the compound, as seen in reactions.
Assigning oxidation numbers to each atom in the given reactions helps us identify which atoms are gaining or losing electrons. This process is key to understanding redox reactions and is the first step in figuring out who gets oxidized and who gets reduced.
Exploring Chemical Reactions
Chemical reactions involve the transformation of substances through breaking and forming chemical bonds. In the context of the commercial production of nitric acid, several chemical reactions occur, each with a specific role in transforming reactants into products.
In these reactions:
  • Ammonia (NH鈧�) reacts with oxygen (O鈧�) to form nitrogen monoxide (NO) and water (H鈧侽).
  • Nitrogen monoxide further reacts with more oxygen to produce nitrogen dioxide (NO鈧�).
  • Finally, nitrogen dioxide reacts with water to produce nitric acid (HNO鈧�) and more nitrogen monoxide.
Each reaction is a step in the industrial synthesis of nitric acid and involves shifts in electron configurations, which are crucial to understanding their reactivity and behavior.
The Process of Oxidation and Reduction
Oxidation and reduction are two sides of the same chemical process, often referred to as "redox" reactions. In a redox reaction, one element's oxidation number increases, indicating oxidation, while another decreases, indicating reduction.
Let's break this down for each reaction:
  • In the first reaction, nitrogen in ammonia undergoes oxidation (from -3 to +2), and oxygen undergoes reduction (from 0 to -2).
  • In the second, nitrogen in NO is oxidized (from +2 to +4), with oxygen reduced (from 0 to -2).
  • In the third, nitrogen in NO鈧� is involved in both its oxidation (from +4 to +5) and reduction (from +4 to +2) as it forms HNO鈧� and NO, respectively.
This interplay ensures that the total electron count remains constant in the reaction, adhering to the law of conservation of matter. Understanding these electron exchanges is foundational to mastering chemistry.

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Most popular questions from this chapter

What does it mean to say that ions are solvated when an ionic substance dissolves in water?

Which element is oxidized and which is reduced in the follow- ing reactions? $$ \begin{array}{l} \text { (a) } \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g) \\ \text { (b) } 3 \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Al}(s) \longrightarrow \\ \text { (c) } \mathrm{Cl}_{2}(a q)+2 \mathrm{NaI}(a q) \longrightarrow \mathrm{I}_{2}(a q)+2 \mathrm{NaCl}(a q) \\ \text { (d) } \mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(l) \end{array} $$

We have seen that ions in aqueous solution are stabilized by the attractions between the ions and the water molecules. Why then do some pairs of ions in solution form precipitates? \([\) Section 4.2\(]\)

You choose to investigate some of the solubility guidelines for two ions not listed in Table \(4.1,\) the chromate ion \(\left(\mathrm{CrO}_{4}^{2-}\right)\) and the oxalate ion \(\left(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\right) .\) You are given \(0.01 \mathrm{M}\) solutions (A, B, C, D) of four water-soluble salts: $$ \begin{array}{lll} \hline \text { Solution } & \text { Solute } & \text { Color of Solution } \\ \hline \text { A } & \mathrm{Na}_{2} \mathrm{CrO}_{4} & \text { Yellow } \\ \mathrm{B} & \left(\mathrm{NH}_{4}\right)_{2} \mathrm{C}_{2} \mathrm{O}_{4} & \text { Colorless } \\ \mathrm{C} & \mathrm{AgNO}_{3} & \text { Colorless } \\ \mathrm{D} & \mathrm{CaCl}_{2} & \text { Colorless } \\ \hline \end{array} $$ When these solutions are mixed, the following observations are made: $$ \begin{array}{lll} \hline \text { Expt } & \text { Solutions } & \\ \text { Number } & \text { Mixed } & \text { Result } \\ \hline 1 & \mathrm{~A}+\mathrm{B} & \text { No precipitate, yellow solution } \\\ 2 & \mathrm{~A}+\mathrm{C} & \text { Red precipitate forms } \\ 3 & \mathrm{~A}+\mathrm{D} & \text { Yellow precipitate forms } \\ 4 & \mathrm{~B}+\mathrm{C} & \text { White precipitate forms } \\ 5 & \mathrm{~B}+\mathrm{D} & \text { White precipitate forms } \\ 6 & \mathrm{C}+\mathrm{D} & \text { White precipitate forms } \end{array} $$ (a) Write a net ionic equation for the reaction that occurs in each of the experiments. (b) Identify the precipitate formed, if any, in each of the experiments.

Using solubility guidelines, predict whether each of the following compounds is soluble or insoluble in water: (a) \(\mathrm{MgBr}_{2}\), (b) \(\mathrm{PbI}_{2}\), (c) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{CO}_{3}\) (e) \(\mathrm{ZnSO}_{4}\). (d) \(\mathrm{Sr}(\mathrm{OH})_{2}\),
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