17.1 Review of Redox Chemistry

Redox reactions are all about electron swapping between chemicals. They're the basis for batteries, rust, and even how our bodies break down food. Understanding redox helps explain countless everyday processes and technologies.

Balancing redox equations can seem tricky, but it's just a step-by-step process. By breaking reactions into half-reactions and following a specific order, you can tackle even complex redox problems. This skill is crucial for predicting and controlling chemical reactions.

Fundamentals of Redox Chemistry

Fundamentals of redox chemistry

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• Redox reactions involve transfer of electrons between chemical species (electron transfer)
  • Oxidation is loss of electrons by a species
  • Reduction is gain of electrons by a species
• Oxidation and reduction occur simultaneously in a redox reaction
  • Oxidizing agent is reduced as it accepts electrons
  • Reducing agent is oxidized as it donates electrons
• Electrochemistry studies interconversion of electrical and chemical energy through redox reactions
  • Galvanic cells (voltaic cells) generate electrical energy from spontaneous redox reactions (batteries)
  • Electrolytic cells use electrical energy to drive non-spontaneous redox reactions (electroplating)

Oxidizing vs reducing agents

• Oxidizing agent accepts electrons and is reduced in the reaction
  • Causes oxidation of another species by taking its electrons
  • Examples: hydrogen peroxide ($\ce{[H2O2](https://www.fiveableKeyTerm:H2O2)}$), permanganate ion ($\ce{[MnO4-](https://www.fiveableKeyTerm:MnO4-)}$), chlorine gas ($\ce{[Cl2](https://www.fiveableKeyTerm:Cl2)}$)
• Reducing agent donates electrons and is oxidized in the reaction
  • Causes reduction of another species by giving it electrons
  • Examples: sodium metal ($\ce{[Na](https://www.fiveableKeyTerm:Na)}$), hydrogen gas ($\ce{[H2](https://www.fiveableKeyTerm:H2)}$), iron metal ($\ce{[Fe](https://www.fiveableKeyTerm:Fe)}$)
• Changes in oxidation numbers help identify oxidizing and reducing agents
  • Increase in oxidation number indicates oxidation occurred (reducing agent)
  • Decrease in oxidation number indicates reduction occurred (oxidizing agent)

Electrochemical cells

• Consist of two half-cells connected by a salt bridge
• Each half-cell contains an electrode (anode or cathode) immersed in an electrolyte solution
• Anode is the site of oxidation, where electrons are lost
• Cathode is the site of reduction, where electrons are gained
• Salt bridge allows ion flow to maintain electrical neutrality
• Standard reduction potential measures the tendency of a species to be reduced

Balancing Redox Equations

Balancing redox equations

• Separate overall reaction into two half-reactions
  • Oxidation half-reaction shows species losing electrons
  • Reduction half-reaction shows species gaining electrons
• Balance each half-reaction in this order:
  1. Balance all atoms except H and O
  2. Balance O atoms by adding $\ce{H2O}$ molecules
  3. Balance H atoms by adding $\ce{H+}$ ions
  4. Balance charge by adding electrons ($\ce{[e-](https://www.fiveableKeyTerm:e-)}$)
• Multiply half-reactions by factors to equalize electrons transferred
• Add balanced half-reactions and cancel out common terms
• Check that final equation is balanced for both atoms and charge