Chapter 16: Oxidation and Reduction – Study Notes for Introductory Chemistry

Oxidation and Reduction: Fundamental Concepts

Introduction to Redox Chemistry

Oxidation and reduction (redox) reactions are central to many chemical processes, including energy production, corrosion, and biological functions. These reactions involve the transfer of electrons between substances, resulting in changes to their oxidation states.

Definitions and Key Terms

• Oxidation: The loss of electrons by a substance, often resulting in an increase in oxidation state.

• Reduction: The gain of electrons by a substance, resulting in a decrease in oxidation state.

• Oxidizing Agent: The substance that is reduced (gains electrons) and causes oxidation of another substance.

• Reducing Agent: The substance that is oxidized (loses electrons) and causes reduction of another substance.

• Oxidation State (Number): A fictitious charge assigned to each atom in a compound, used for electron bookkeeping.

Mnemonics: OIL RIG (Oxidation Is Loss, Reduction Is Gain) and LEO the lion says GER (Lose Electrons Oxidation, Gain Electrons Reduction).

Examples of Oxidation

• Slow Oxidation: Rusting of iron.

• Rapid Oxidation: Combustion of carbon in natural gas.

Assigning Oxidation States

Rules for Assigning Oxidation States

Oxidation states are assigned according to hierarchical rules:

1. The oxidation state of an atom in a free element is 0.

2. The oxidation state of a monoatomic ion equals its charge.

3. The sum of oxidation states in a neutral molecule is 0; in a polyatomic ion, it equals the ion's charge.

4. Group I metals: +1; Group II metals: +2 in compounds.

5. Nonmetals are assigned oxidation states based on a priority table (e.g., F: -1, O: -2, H: +1, etc.).

Note: Oxidation state is not the same as ionic charge.

Identifying Oxidation and Reduction via Oxidation States

Oxidation is an increase in oxidation state; reduction is a decrease. For example:

Example: In the reaction $\mathrm{Ca(s) + 2 H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)}$, calcium is oxidized (state increases from 0 to +2), hydrogen is reduced (state decreases from +1 to 0).

Redox Reactions in Everyday Life

Bleaching of Hair

Hydrogen peroxide ($\mathrm{H_2O_2}$) acts as an oxidizing agent, oxidizing melanin and thiol groups in hair, resulting in lighter color and increased tangling.

Balancing Redox Equations

Half-Reaction Method

Redox equations are often balanced using the half-reaction method, which involves:

1. Assigning oxidation states to all atoms.

2. Dividing the reaction into oxidation and reduction half-reactions.

3. Balancing each half-reaction for mass and charge (by adding electrons).

4. Equalizing electrons lost and gained by multiplying half-reactions.

5. Adding the half-reactions and verifying balance for mass and charge.

Special procedures exist for acidic and basic solutions, involving addition of $\mathrm{H^+}$, $\mathrm{OH^-}$, and $\mathrm{H_2O}$ as needed.

Example: Aluminum and Silver Redox Reaction

Consider the reaction $\mathrm{Al(s) + Ag^+(aq) \rightarrow Al^{3+}(aq) + Ag(s)}$:

Redox Reactions in Biological and Environmental Processes

Photosynthesis and Respiration

Photosynthesis and respiration are key biological processes involving redox reactions:

• Respiration: Oxidation of glucose to produce energy.

• Photosynthesis: Reduction of carbon dioxide to form glucose using sunlight.

Equations:

• Respiration: $\mathrm{C_6H_{12}O_6 + 6 O_2 \rightarrow 6 H_2O + 6 CO_2 + energy}$

• Photosynthesis: $\mathrm{6 CO_2 + 6 H_2O + energy (sunlight) \rightarrow C_6H_{12}O_6 + 6 O_2}$

The Activity Series of Metals

Predicting Spontaneous Redox Reactions

The activity series ranks metals by their tendency to lose electrons (be oxidized). Metals at the top are most reactive and easily oxidized; those at the bottom are least reactive.

Example: Gold is low on the activity series and resists oxidation, making it suitable for jewelry.

Spontaneity of Redox Reactions

Pairing a metal from the top of the activity series with the reduction of a metal ion from the bottom results in a spontaneous reaction.

Redox Reactions and Acids

Predicting Metal Dissolution in Acid

Metals above hydrogen in the activity series dissolve in acids by reducing $\mathrm{H^+}$ to hydrogen gas. Metals below hydrogen generally do not dissolve in acids, except in the case of nitric acid.

Batteries and Electrochemical Cells

Voltaic (Galvanic) Cells

Voltaic cells generate electrical current from spontaneous redox reactions. The anode is where oxidation occurs (negative), and the cathode is where reduction occurs (positive). Electrons flow from anode to cathode.

Dry-Cell and Lead-Acid Batteries

Dry-cell batteries use zinc as the anode and a carbon rod as the cathode, with MnO2 as the oxidizing agent. Lead-acid batteries, used in automobiles, consist of lead and lead dioxide electrodes immersed in sulfuric acid.

Fuel Cells and Electrolysis

Hydrogen–Oxygen Fuel Cells

Fuel cells use the redox reaction between hydrogen and oxygen to generate electricity, emitting only water as a byproduct. They are promising for environmentally friendly vehicles.

Electrolysis

Electrolysis uses electrical current to drive nonspontaneous redox reactions, such as splitting water into hydrogen and oxygen or plating metals.

Corrosion: Undesirable Redox Reactions

Rusting of Iron

Corrosion is the oxidation of metals, most commonly iron rusting. Rust ($\mathrm{Fe_2O_3}$) flakes off, exposing more iron to oxidation.

Prevention of Corrosion

Corrosion can be prevented by keeping metals dry, applying protective coatings, or using a sacrificial anode (a more active metal that oxidizes in place of iron).

Applications: Fuel-Cell Breathalyzer

Principle of Operation

Fuel-cell breathalyzers measure blood alcohol content by oxidizing ethyl alcohol at the anode and reducing oxygen at the cathode, generating a current proportional to alcohol concentration.

Summary Table: Activity Series of Metals

Metal	Half-Reaction	Reactivity
Li	Li(s) → Li+(aq) + e-	Most reactive
Au	Au(s) → Au3+(aq) + 3e-	Least reactive
Zn	Zn(s) → Zn2+(aq) + 2e-	Reactive
Fe	Fe(s) → Fe2+(aq) + 2e-	Moderate
Cu	Cu(s) → Cu2+(aq) + 2e-	Less reactive
Ag	Ag(s) → Ag+(aq) + e-	Low reactivity
Additional info:	See image_16 for full activity series.

Review and Learning Objectives

• Define and identify oxidation and reduction.

• Identify oxidizing and reducing agents.

• Assign oxidation states.

• Balance redox reactions.

• Predict spontaneous redox reactions.