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

Oxidation and Reduction

Introduction to Redox Chemistry

Oxidation and reduction (redox) reactions are fundamental 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.

Fuel Cell Technology and Redox Reactions

Fuel cells generate electricity by harnessing redox reactions, particularly the transfer of electrons from hydrogen to oxygen. In hydrogen–oxygen fuel cells, hydrogen is oxidized (loses electrons) and oxygen is reduced (gains electrons), producing water as the only emission.

• Oxidation: Loss of electrons (hydrogen in fuel cells).

• Reduction: Gain of electrons (oxygen in fuel cells).

• Application: Fuel cells power electric vehicles with water as the emission.

Definitions of Oxidation and Reduction

Oxidation and reduction can be defined in several ways:

• Oxidation: Loss of electrons or gain of oxygen.

• Reduction: Gain of electrons or loss of oxygen.

• Oxidizing agent: Substance that is reduced (gains electrons).

• Reducing agent: Substance that is oxidized (loses electrons).

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

Oxidation States: Electron Bookkeeping

Oxidation states (or numbers) are assigned to atoms in compounds to track electron transfer during reactions. The rules for assigning oxidation states are hierarchical:

• Free element: Oxidation state is 0.

• Monatomic ion: Oxidation state equals its charge.

• Sum of oxidation states in a neutral molecule is 0; in a polyatomic ion, equals the ion's charge.

• Group I metals: +1; Group II metals: +2.

• Nonmetals: Assigned according to a priority table (e.g., fluorine, oxygen, hydrogen).

Identifying Oxidation and Reduction Using Oxidation States

Oxidation is an increase in oxidation state; reduction is a decrease. Assigning oxidation states to each atom in a reaction helps identify which species are oxidized and reduced.

• Example: Calcium reacts with water:

Everyday Chemistry: Hair Bleaching

Hydrogen peroxide (H2O2) is an oxidizing agent used in hair bleaching. It oxidizes melanin (hair pigment) and thiol groups in hair proteins, resulting in lighter and more tangled hair.

Balancing Redox Equations

The Half-Reaction Method

Redox equations are often balanced using the half-reaction method, which separates the overall reaction into oxidation and reduction half-reactions. Each half-reaction is balanced for mass and charge, then combined.

• Assign oxidation numbers to all atoms.

• Divide into two half-reactions.

• Balance each for mass and charge (add electrons as needed).

• Equalize electrons lost and gained by multiplying half-reactions.

• Add half-reactions together, canceling electrons.

• Verify balance for mass and charge.

Balancing Redox Reactions in Acidic and Basic Solutions

Special procedures are used for reactions in acidic or basic solutions:

• Acidic: Balance O with H2O, H with H+.

• Basic: Balance O with H2O, H with H+, then neutralize H+ with OH−.

Chemistry in the Environment: Photosynthesis and Respiration

Energy for Life

Photosynthesis and respiration are key biological processes involving redox reactions. Plants capture solar energy to reduce carbon dioxide into glucose, while animals oxidize glucose to release energy.

• Respiration: Oxidation of carbon in glucose.

• Photosynthesis: Reduction of carbon dioxide.

The Activity Series: Predicting Spontaneous Redox Reactions

Activity Series of Metals

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.

• Pairing a top half-reaction with the reverse of a bottom half-reaction yields a spontaneous reaction.

• Metals below hydrogen do not dissolve in acids (except with 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 carbon as the cathode, with MnO2 as the oxidizing agent. Lead-acid batteries (used in cars) consist of lead and lead dioxide electrodes immersed in sulfuric acid.

Fuel Cells and Electrolysis

Fuel cells use spontaneous redox reactions to generate electricity. Electrolysis uses electrical current to drive nonspontaneous redox reactions, such as splitting water into hydrogen and oxygen.

Corrosion: Undesirable Redox Reactions

Rusting and Metal Protection

Corrosion is the oxidation of metals, most commonly iron rusting. Iron oxide flakes off, exposing more iron to oxidation. Other metals, like aluminum, form protective oxide coatings.

• Paint and coatings prevent corrosion by keeping metals dry.

• Sacrificial anodes (e.g., zinc) protect iron by oxidizing in its place.

Everyday Chemistry: Fuel-Cell Breathalyzer

Application of Redox Reactions

Fuel-cell breathalyzers use redox reactions to measure blood alcohol levels. Ethyl alcohol is oxidized at the anode, and oxygen is reduced at the cathode. The electrical current produced is proportional to the alcohol concentration.

Summary Table: Activity Series of Metals

Metal	Half-Reaction	Reactivity
Li	Li(s) → Li+(aq) + e−	Most reactive
Au	Au(s) → Au+(aq) + e−	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−	Least reactive

Additional info: Table entries are representative; see image_16 for full activity series.

Key Equations (LaTeX Format)

• $\mathrm{Ca(s) + 2 H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)}$

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

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

• $\mathrm{Zn(s) + 2 H^+(aq) \rightarrow Zn^{2+}(aq) + H_2(g)}$

• $\mathrm{2 Fe(s) + O_2(g) + 2 H_2O(l) \rightarrow 2 Fe(OH)_2(s)}$

• $\mathrm{2 H_2O(l) \rightarrow 2 H_2(g) + O_2(g)}$

Learning Objectives

• Define and identify oxidation and reduction.

• Identify oxidizing agents and reducing agents.

• Assign oxidation states.

• Balance redox reactions.

• Predict spontaneous redox reactions.