17.7 Electrolysis

Electrolysis is the process of using electrical energy to force a non-spontaneous chemical reaction to occur. It's how industries extract reactive metals, purify copper, plate jewelry with gold, and produce chemicals like chlorine gas. As a core topic in electrochemistry, electrolysis shows the flip side of galvanic cells: instead of getting electricity from a reaction, you're putting electricity into one.

Electrolysis

Electrochemistry Fundamentals

Electrochemistry studies chemical processes that involve electron transfer. The foundation here is redox reactions, where one species loses electrons (oxidation) and another gains them (reduction). These two processes are described separately as half-reactions, one for each electrode.

For any of this to work, you need an electrolyte: a substance that conducts electricity because it contains freely moving ions. This happens either when an ionic compound is dissolved in water or when it's melted (molten). The conductivity of the solution depends on how many ions are present and how easily they move.

Steps and Components of Electrolysis

An electrolytic cell uses an external power source (like a battery) to drive a reaction that wouldn't happen on its own. Here's what's inside and how it works:

Components:

• An electrolyte (an ionic compound, either dissolved in water or molten)
• Two electrodes submerged in the electrolyte
  • Cathode (negative electrode): where reduction occurs. Cations migrate here and gain electrons.
  • Anode (positive electrode): where oxidation occurs. Anions migrate here and lose electrons.
• An external power source supplying electrical energy

The process, step by step:

1. The electrolyte dissociates into cations (+) and anions (−).
2. The external power source pushes electrons into the cathode and pulls them from the anode.
3. Cations in solution migrate toward the cathode, where they pick up electrons (reduction).
4. Anions migrate toward the anode, where they give up electrons (oxidation).
5. The net result is a non-spontaneous chemical change driven by electrical energy.

Common applications:

• Electroplating: depositing a thin layer of metal (like chromium or silver) onto an object's surface
• Electrolytic refining: purifying metals like copper by dissolving impure metal at the anode and depositing pure metal at the cathode
• Production of elements: extracting highly reactive metals like sodium or aluminum from their molten compounds, since these metals can't be obtained by ordinary chemical reduction

Electrolytic vs. Galvanic Cells

These two cell types are essentially mirrors of each other. Both involve redox reactions and electrodes, but the direction of energy flow is reversed.

One thing that trips students up: reduction always happens at the cathode and oxidation always happens at the anode in both cell types. What changes is the sign (+ or −) of each electrode. In a galvanic cell the cathode is positive; in an electrolytic cell it's negative.

Faraday's Laws in Electrolysis Calculations

Faraday's laws connect the amount of electric charge you pass through a cell to the amount of substance produced or consumed. This is how you solve quantitative electrolysis problems.

First Law: The mass of substance deposited or dissolved at an electrode is directly proportional to the total charge passed through the cell.

$m = kQ$

where $m$ is mass, $k$ is a proportionality constant, and $Q$ is total charge in coulombs.

Second Law: When the same amount of charge passes through different electrolytic cells, the masses of substances produced are proportional to their equivalent weights (molar mass divided by the number of electrons transferred per ion).

The practical formula that combines both laws:

$m = \frac{QM}{nF}$

• $m$ = mass of substance produced (grams)
• $Q$ = total charge passed (coulombs). Calculate this as $Q = I \times t$, where $I$ is current in amps and $t$ is time in seconds.
• $M$ = molar mass of the substance (g/mol)
• $n$ = number of electrons transferred per formula unit (from the balanced half-reaction)
• $F$ = Faraday's constant = 96,485 C/mol of electrons

How to solve an electrolysis problem:

1. Write the balanced half-reaction at the electrode you care about.
2. Identify $n$ (electrons transferred) and $M$ (molar mass of the product).
3. Calculate total charge: $Q = I \times t$.
4. Plug into $m = \frac{QM}{nF}$ to find mass.
5. If the product is a gas, convert moles to volume using the ideal gas law: $PV = nRT$.