Electrolysis Cells & Faraday's Laws Cheat Sheet

A printable reference covering electrolysis cells, electrode reactions, charge, current, Faraday’s constant, and mass deposited for grades 10-12.

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Electrolysis uses electrical energy to force a nonspontaneous redox reaction to occur. This cheat sheet helps students track electron flow, identify anode and cathode reactions, and connect electric current to chemical change. It is especially useful for solving plating, gas production, and molten or aqueous electrolysis problems. Students need these tools to move confidently between balanced half-reactions, charge, moles of electrons, and mass.

Key Facts

In an electrolytic cell, oxidation occurs at the anode and reduction occurs at the cathode.
In electrolysis, the anode is positive and the cathode is negative because the power supply pulls electrons from the anode and pushes electrons to the cathode.
Electric charge is calculated by $Q = It$ , where $Q$ is charge in coulombs, $I$ is current in amperes, and $t$ is time in seconds.
Moles of electrons are found using $n_{e^-} = \frac{Q}{F}$ , where $F = 96485\ \text{C mol}^{-1}$ .
Faraday’s law for deposited mass is $m = \frac{ItM}{nF}$ , where $M$ is molar mass and $n$ is electrons transferred per ion.
For a metal ion $M^{n+}$ , the cathode half-reaction is $M^{n+} + ne^- \rightarrow M(s)$ .
For electrolysis problems, always convert time to seconds using $t(\text{s}) = t(\text{min}) \times 60$ or $t(\text{s}) = t(\text{h}) \times 3600$ .
The stoichiometric ratio between product and electrons comes from the balanced half-reaction, not from the coefficient of the compound alone.

Vocabulary

Electrolytic cell: A cell that uses electrical energy to drive a nonspontaneous redox reaction.
Anode: The electrode where oxidation occurs and electrons are produced by the reacting species.
Cathode: The electrode where reduction occurs and electrons are gained by the reacting species.
Faraday’s constant: The charge carried by one mole of electrons, equal to $F = 96485\ \text{C mol}^{-1}$ .
Electroplating: A process in which electrolysis deposits a thin layer of metal onto an object.
Half-reaction: A balanced equation showing either oxidation or reduction, including the electrons transferred.

Common Mistakes to Avoid

Confusing anode and cathode signs, because galvanic and electrolytic cells have different electrode polarities. In electrolysis, the anode is positive and the cathode is negative.
Forgetting to convert time to seconds, which makes $Q = It$ wrong because amperes mean coulombs per second. Convert minutes or hours before calculating charge.
Using the wrong value of $n$ in $m = \frac{ItM}{nF}$ , because $n$ must come from the balanced half-reaction. For $Cu^{2+} + 2e^- \rightarrow Cu(s)$ , $n = 2$ .
Treating $F$ as charge for one electron, which is incorrect because $F$ is charge for one mole of electrons. Use $n_{e^-} = \frac{Q}{F}$ to find moles of electrons.
Ignoring competing reactions in aqueous electrolysis, because water can be oxidized or reduced along with dissolved ions. Always consider which species is easier to discharge at each electrode.

Practice Questions

1 A current of $2.50\ \text{A}$ passes through molten $NaCl$ for $30.0\ \text{min}$ . What charge $Q$ passes through the cell?
2 How many grams of copper are deposited when $3.00\ \text{A}$ flows through $CuSO_4(aq)$ for $20.0\ \text{min}$ , using $Cu^{2+} + 2e^- \rightarrow Cu(s)$ and $M_{Cu} = 63.55\ \text{g mol}^{-1}$ ?
3 How long, in minutes, is needed to deposit $5.00\ \text{g}$ of silver from $Ag^+(aq)$ using a current of $1.50\ \text{A}$ , given $Ag^+ + e^- \rightarrow Ag(s)$ and $M_{Ag} = 107.87\ \text{g mol}^{-1}$ ?
4 Explain why the cathode in an electrolytic cell is negative even though reduction always occurs at the cathode.

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