17.1 Review of Redox Chemistry

Fundamentals of Redox Chemistry

Redox reactions are all about electron transfer between chemical species. They're the basis for batteries, rust, and even how our bodies break down food. Understanding redox helps explain countless everyday processes and sets the foundation for the rest of electrochemistry.

Core Concepts

Oxidation is the loss of electrons by a species. Reduction is the gain of electrons. A helpful mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

These two processes always occur together. You can't have one without the other. If one species loses electrons, another species must be there to accept them.

• The oxidizing agent is the species that accepts electrons (and is itself reduced). It causes oxidation in the other species by taking its electrons.
• The reducing agent is the species that donates electrons (and is itself oxidized). It causes reduction in the other species by giving away its electrons.

This naming convention trips people up. The oxidizing agent doesn't get oxidized; it gets reduced. Think of it from the other species' perspective: the oxidizing agent is the reason oxidation happens to someone else.

Identifying Oxidizing and Reducing Agents

Track oxidation numbers (oxidation states) to figure out what's being oxidized and what's being reduced:

• If a species' oxidation number increases, it lost electrons and was oxidized. That species is the reducing agent.
• If a species' oxidation number decreases, it gained electrons and was reduced. That species is the oxidizing agent.

Common oxidizing agents: hydrogen peroxide ($\ce{H2O2}$), permanganate ion ($\ce{MnO4-}$), chlorine gas ($\ce{Cl2}$)

Common reducing agents: sodium metal ($\ce{Na}$), hydrogen gas ($\ce{H2}$), iron metal ($\ce{Fe}$)

Electrochemical Cells

Electrochemistry studies the interconversion of electrical and chemical energy through redox reactions. There are two main types of electrochemical cells:

• Galvanic cells (also called voltaic cells) generate electrical energy from spontaneous redox reactions. Batteries are a familiar example.
• Electrolytic cells use electrical energy to drive non-spontaneous redox reactions. Electroplating is a common application.

Both types share the same basic setup:

• Two half-cells, each containing an electrode immersed in an electrolyte solution
• The anode is where oxidation occurs (electrons are lost)
• The cathode is where reduction occurs (electrons are gained)
• A salt bridge connects the two half-cells, allowing ions to flow between them so electrical neutrality is maintained
• Standard reduction potentials measure how strongly a species tends to gain electrons (be reduced). You'll use these values extensively later in this unit.

Another mnemonic: An Ox, Red Cat (Anode = Oxidation, Reduction = Cathode).

Balancing Redox Equations

Balancing redox equations can seem tricky, but it's a repeatable step-by-step process. The key idea is to split the overall reaction into two half-reactions, balance each one separately, then combine them.

The Half-Reaction Method (Acidic Solution)

1. Separate the overall reaction into an oxidation half-reaction and a reduction half-reaction.
2. Balance all atoms except H and O in each half-reaction.
3. Balance O atoms by adding $\ce{H2O}$ molecules to whichever side needs oxygen.
4. Balance H atoms by adding $\ce{H+}$ ions to whichever side needs hydrogen.
5. Balance charge by adding electrons ($\ce{e-}$) to the more positive side of each half-reaction.
6. Equalize electrons by multiplying each half-reaction by the appropriate factor so both half-reactions transfer the same number of electrons.
7. Add the two balanced half-reactions together and cancel species that appear on both sides.
8. Check that the final equation is balanced for both atoms and charge.

The most common mistake is forgetting step 6. If the oxidation half-reaction produces 2 electrons but the reduction half-reaction consumes 3, you need to multiply them (by 3 and 2, respectively) before adding. Electrons should completely cancel in the final balanced equation.