Question

In each of the following reactions, identify which element is oxidized and which is reduced by assigning oxidation numbers. a. \(2 \mathrm{HNO}_{3}(a q)+3 \mathrm{H}_{2} \mathrm{S}(g) \rightarrow\) \(2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(g)+3 \mathrm{S}(s)\) b. \(2 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)\) c. \(2 \mathrm{ZnS}(s)+3 \mathrm{O}_{2}(g) \rightarrow 2 \mathrm{ZnO}(s)+2 \mathrm{SO}_{2}(g)\) d. \(\mathrm{CH}_{4}(g)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{CH}_{3} \mathrm{Cl}(g)+\mathrm{HCl}(g)\)

Short answer

In the given reactions, the following elements are oxidized and reduced: a. Oxidation: S in H₂S (from -2 to 0); Reduction: N in HNO₃ (from +5 to +2) b. Oxidation: O in H₂O₂ (from -1 to 0); Reduction: O in H₂O₂ (from -1 to -2) c. Oxidation: S in ZnS (from -2 to +4); Reduction: O in O₂ (from 0 to -2) d. Oxidation: C in CH₄ (from -4 to -2); Reduction: Cl in Cl₂ (from 0 to -1)

Step-by-step solution

Understand the rules for assigning oxidation numbers

Here are some general rules for assigning oxidation numbers to atoms in a compound: 1. The oxidation number of an element in its elemental form (like \(O_2\) or \(Zn\)) is 0. 2. The oxidation number of a monatomic ion is equal to its charge (e.g., \(Na^+\) has an oxidation number of +1). 3. The sum of the oxidation numbers in a compound must be equal to the overall charge of the compound. 4. In general, Oxygen has an oxidation number of -2 (except for peroxides when it is -1), and Hydrogen has an oxidation number of +1 (except in some metal hydrides when it is -1). 5. More electronegative elements in a compound will have a negative oxidation number, while less electronegative elements will have a positive oxidation number. Now let's apply these rules to each reaction to find out which atom is being oxidized and which is being reduced.

Assign oxidation numbers to reaction (a)

First, assign oxidation numbers to each atom of the chemical equation for reaction (a) using the rules mentioned above: \(2 \mathrm{HNO}_{3} \rightarrow 2 \mathrm{NO} + 4 \mathrm{H}_{2} \mathrm{O} + 3 \mathrm{S}\) Oxidation numbers: \(\mathrm{H}^{+1}\mathrm{N}^{+5}\mathrm{O}^{-2}\) in \(\mathrm{HNO}_3\), \(\mathrm{H}^{+1}\mathrm{S}^{-2}\) in \(\mathrm{H_2 S}\), \(\mathrm{N}^{+2}\mathrm{O}^{-2}\) in \(\mathrm{NO}\), \(\mathrm{H}^{+1}\mathrm{O}^{-2}\) in \(\mathrm{H_2 O}\), and \(\mathrm{S}^{0}\) in \(\mathrm{S}\). Oxidation: \(\mathrm{S}\) in \(\mathrm{H_2 S}\) (from -2 to 0) Reduction: \(\mathrm{N}\) in \(\mathrm{HNO_3}\) (from +5 to +2)

Assign oxidation numbers to reaction (b)

Next, assign oxidation numbers to each atom of the chemical equation for reaction (b) using the rules mentioned above: \(2 \mathrm{H}_{2} \mathrm{O}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O} + \mathrm{O}_{2}\) Oxidation numbers: \(\mathrm{H}^{+1}\mathrm{O}^{-1}\) in \(\mathrm{H_2 O_2}\), \(\mathrm{H}^{+1}\mathrm{O}^{-2}\) in \(\mathrm{H_2 O}\), and \(\mathrm{O}^{0}\) in \(\mathrm{O_2}\). Oxidation: \(\mathrm{O}\) in \(\mathrm{H_2 O_2}\) (from -1 to 0) Reduction: \(\mathrm{O}\) in \(\mathrm{H_2 O_2}\) (from -1 to -2)

Assign oxidation numbers to reaction (c)

Next, assign oxidation numbers to each atom of the chemical equation for reaction (c) using the rules mentioned above: \(2 \mathrm{ZnS} + 3 \mathrm{O}_{2} \rightarrow 2 \mathrm{ZnO} + 2 \mathrm{SO}_{2}\) Oxidation numbers: \(\mathrm{Zn}^{+2}\mathrm{S}^{-2}\) in \(\mathrm{ZnS}\), \(\mathrm{O}^{0}\) in \(\mathrm{O_2}\), \(\mathrm{Zn}^{+2}\mathrm{O}^{-2}\) in \(\mathrm{ZnO}\), and \(\mathrm{S}^{+4}\mathrm{O}^{-2}\) in \(\mathrm{SO_2}\). Oxidation: \(\mathrm{S}\) in \(\mathrm{ZnS}\) (from -2 to +4) Reduction: \(\mathrm{O}\) in \(\mathrm{O_2}\) (from 0 to -2)

Assign oxidation numbers to reaction (d)

Finally, assign oxidation numbers to each atom of the chemical equation for reaction (d) using the rules mentioned above: \(\mathrm{CH}_{4} + \mathrm{Cl}_{2} \rightarrow \mathrm{CH}_{3} \mathrm{Cl} + \mathrm{HCl}\) Oxidation numbers: \(\mathrm{C}^{-4}\mathrm{H}^{+1}\) in \(\mathrm{CH_4}\), \(\mathrm{Cl}^{0}\) in \(\mathrm{Cl_2}\), \(\mathrm{C}^{-2}\mathrm{H}^{+1}\mathrm{Cl}^{-1}\) in \(\mathrm{CH_3 Cl}\), and \(\mathrm{H}^{+1}\mathrm{Cl}^{-1}\) in \(\mathrm{HCl}\). Oxidation: \(\mathrm{C}\) in \(\mathrm{CH_4}\) (from -4 to -2) Reduction: \(\mathrm{Cl}\) in \(\mathrm{Cl_2}\) (from 0 to -1)

Key concepts

Oxidation Numbers

Understanding oxidation numbers is essential in redox chemistry. They help us identify how elements undergo oxidation and reduction.
The oxidation number, sometimes called an oxidation state, reflects the number of electrons an atom gains or loses during a chemical reaction.
This concept is crucial because it helps determine which atoms are oxidized and which are reduced in a reaction.

• The oxidation number of any pure element is 0. For example, \( \text{O}_2 \) or \( \text{H}_2 \) is zero.
• For monoatomic ions, the oxidation number is equal to the ion's charge.
  For instance, in \( \text{Na}^+ \), the oxidation number is +1, while in \( \text{Cl}^- \), it's -1.
• Within compounds, more electronegative atoms receive negative oxidation numbers.
  For example, oxygen typically has an oxidation number of -2 (unless in peroxides like hydrogen peroxide where it's -1).
• The sum of oxidation numbers in a neutral compound must be 0, whereas in polyatomic ions, the total must match the ion's overall charge.

Learning and applying these rules enables us to systematically assign oxidation numbers in reactions and determine shifts in electron density.

Redox Chemistry

Redox chemistry is a fascinating area that looks into reactions where electron transfer happens. It's the umbrella under which oxidation-reduction reactions belong.
These reactions are pivotal in both nature and industrial processes, offering insights into energy conversion, metallurgy, and even biological systems such as respiration.

• Redox stands for reduction and oxidation.
• Reduction refers to the gain of electrons, resulting in a decrease in oxidation number.
• Oxidation is the loss of electrons, leading to an increase in oxidation number.

Take reaction \( a \), where \( \text{S} \) is oxidized from \( -2 \) to \( 0 \) and \( \text{N} \) is reduced from \( +5 \) to \( +2 \).
These changes highlight the core principle of redox: a balance between electron loss and gain.
Recognizing these electron shifts is fundamental to mastering chemical reactions at the molecular level.

Chemical Reactions

Chemical reactions showcase the transformation of substances via the breaking and forming of bonds. Redox reactions are a key class of these transformations, with widespread applications.

• They follow the principle of mass conservation, meaning the amount of matter remains constant, even as substances change.
• Chemical equations, such as \( c \) \( 2 \text{ZnS} + 3 \text{O}_2 \rightarrow 2 \text{ZnO} + 2 \text{SO}_2 \), represent these transformations by detailing reactants and products.
• In redox reactions, besides the change in substance identity, there's an electron transfer between species.
  This electron exchange is detected via changes in oxidation numbers.

This understanding helps break down complex reactions into simpler steps, making it easier to deduce which substances are oxidized and reduced, as seen in each assigned equation.

Oxidation and Reduction Processes

These processes are central in redox chemistry, both involving the movement of electrons.
Each implies specific changes for species within a chemical reaction as demonstrated through reaction \( d \).

• Oxidation: When an entity loses electrons, it undergoes oxidation. For example, in \( \text{CH}_4 \) in reaction \( d \), carbon changes from \( -4 \) to \( -2 \), indicating it lost some electrons.
• Reduction: This refers to the gain of electrons, as seen with \( \text{Cl}_2 \) in the same reaction, which goes from an oxidation number of \( 0 \) to \( -1 \).
  This demonstrates the incorporation of electrons.
• Though seemingly opposite, these processes are interdependent. In any reaction, oxidation and reduction occur simultaneously.

This relationship ensures that whenever electrons are lost by one substance, they are gained by another, maintaining the balance of charge and matter in chemical processes.