Question

In each of the following reactions, identify which element is being oxidized and which is being reduced by assigning oxidation numbers. a. \(2 \mathrm{Al}(s)+3 \mathrm{~S}(s) \rightarrow \mathrm{Al}_{2} \mathrm{~S}_{3}(s)\) b. \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)\) c. \(2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{C}(s) \rightarrow 3 \mathrm{CO}_{2}(g)+4 \mathrm{Fe}(s)\) d. \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}(a q)+14 \mathrm{HCl}(a q) \rightarrow 2 \mathrm{KCl}(a q)+2 \mathrm{CrCl}_{3}(s)+7 \mathrm{H}_{2} \mathrm{O}(l)+3 \mathrm{Cl}_{2}(g)\)

Short answer

d. \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}(a q)+14 \mathrm{HCl}(a q) \rightarrow 2 \mathrm{KCl}(a q)+2 \mathrm{CrCl}_{3}(s)+7 \mathrm{H}_{2} \mathrm{O}(l)+3 \mathrm{Cl}_{2}(g)\) Step 1: Assign Oxidation Numbers for Reactants Assign oxidation numbers to each atom in the reactants: K: +1 (due to its 1+ charge in K2Cr2O7) Cr: +6 (due to its 6+ charge in K2Cr2O7) O: -2 (due to its 2- charge in K2Cr2O7) H: +1 (due to its 1+ charge in HCl) Cl: -1 (due to its 1- charge in HCl) Step 2: Assign Oxidation Numbers for Products Assign oxidation numbers to each atom in the products: K: +1 (due to its 1+ charge in KCl) Cr: +3 (due to its 3+ charge in CrCl3) O: -2 (due to its 2- charge in H2O) H: +1 (due to its 1+ charge in H2O) Cl: -1 (due to its 1- charge in KCl and CrCl3) Step 3: Identify Oxidizing and Reducing Agents Compare the change in oxidation numbers from the reactants to the products: K: +1 → +1 (no change) Cr: +6 → +3 (reduction: being reduced) O: -2 → -2 (no change) H: +1 → +1 (no change) Cl: -1 → -1 (no change)

Step-by-step solution

Assign Oxidation Numbers for Reactants

Starting with the reactants, assign oxidation numbers to each atom. Uncombined elements have an oxidation number of 0: Al: 0 S: 0

Assign Oxidation Numbers for Products

Now, assign oxidation numbers to each atom in the products. For the product, the oxidation numbers of Al and S should reflect their respective oxidation states: Al: +3 (due to its 3+ charge in Al2S3) S: -2 (due to its 2- charge in Al2S3)

Identify Oxidizing and Reducing Agents

Compare the change in oxidation numbers from the reactants to the products: Al: 0 → +3 (oxidation: being oxidized) S: 0 → -2 (reduction: being reduced) b. \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)\)

Assign Oxidation Numbers for Reactants

Assign oxidation numbers to each atom in the reactants: C: -4 (due to its 4- charge in CH4) H: +1 (due to its 1+ charge in CH4) O: 0 (uncombined elemental oxygen)

Assign Oxidation Numbers for Products

Assign oxidation numbers to each atom in the products: C: +4 (due to its 4+ charge in CO2) H: +1 (due to its 1+ charge in H2O) O: -2 (due to its 2- charge in CO2 and H2O)

Identify Oxidizing and Reducing Agents

Compare the change in oxidation numbers from the reactants to the products: C: -4 → +4 (oxidation: being oxidized) H: +1 → +1 (no change) O: 0 → -2 (reduction: being reduced) c. \(2 \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+3 \mathrm{C}(s) \rightarrow 3 \mathrm{CO}_{2}(g)+4 \mathrm{Fe}(s)\)

Assign Oxidation Numbers for Reactants

Assign oxidation numbers to each atom in the reactants: Fe: +3 (due to its 3+ charge in Fe2O3) O: -2 (due to its 2- charge in Fe2O3) C: 0 (uncombined elemental carbon)

Assign Oxidation Numbers for Products

Assign oxidation numbers to each atom in the products: Fe: 0 (uncombined elemental iron) O: -2 (due to its 2- charge in CO2) C: +4 (due to its 4+ charge in CO2)

Identify Oxidizing and Reducing Agents

Compare the change in oxidation numbers from the reactants to the products: Fe: +3 → 0 (reduction: being reduced) O: -2 → -2 (no change) C: 0 → +4 (oxidation: being oxidized)

Key concepts

Oxidation Numbers

Oxidation numbers help us understand how electrons are transferred in chemical reactions. They are theoretical charges assigned to atoms, which can vary based on chemical bonding. We determine oxidation numbers by rules, such as:

• Elements in their pure form, like O₂ or S, have an oxidation number of 0.
• For ions, the oxidation number equals the ion's charge. For example, Al in Al₂S₃ has +3.
• Oxygen typically has an oxidation number of -2, except in peroxides.
• Hydrogen is usually +1, except when bonded with metals.

By assigning these numbers, we can track which atoms lose or gain electrons, crucial for identifying oxidization and reduction.

Oxidizing Agent

An oxidizing agent is a chemical species that causes another species to lose electrons. In simpler words, it is the substance that gets reduced itself. As it gains electrons, its oxidation state decreases. For example, in the reaction of CH₄ and O₂ to form CO₂ and H₂O, oxygen is the oxidizing agent as it goes from 0 to -2. This implies it gains electrons, driving the oxidation of carbon from -4 to +4. Understanding the role of the oxidizing agent allows you to decipher the electron flow in a reaction.

Reducing Agent

A reducing agent is the opposite of an oxidizing agent. It donates electrons to another substance, becoming oxidized itself. Its oxidation number increases as it loses electrons. In the reaction between iron(III) oxide (Fe₂O₃) and carbon, carbon acts as the reducing agent. It starts with an oxidation number of 0 and becomes +4 in carbon dioxide (CO₂). Meanwhile, iron is reduced as it moves from an oxidation state of +3 to 0. Recognizing the reducing agent helps reveal the electron donors within a reaction.

Chemical Reactions

Chemical reactions transform reactants into products through rearrangement of atoms. These processes often involve the exchange of electrons, particularly in redox reactions where oxidation and reduction occur simultaneously. In the reaction a. with Al and S, aluminum is oxidized, increasing its oxidation number from 0 to +3, while sulfur is reduced from 0 to -2. By examining these changes, students can better understand the flow of electrons and the overall chemical transformation. Mastering redox reactions unveils the intricate dance of electrons that powers numerous natural and industrial processes.