Electricity and Chemistry Problem Set

Introduction

Electrochemistry questions can mix cell labels, solution ions and electrical measurements. The reliable first move is to decide what happens at each electrode. Then use Q = It, Q/F and half-reaction coefficients only if a quantitative product amount is requested.

Core explanation

An electrode question begins with reaction classification. Oxidation occurs at the anode and reduction at the cathode in any cell. For a typical electrolytic cell, the cathode is negative and anode positive because of the external power source. A typical galvanic cell delivering power has the reverse usual signs. If an electron appears on the left of a half-reaction, reduction is occurring; on the right, oxidation is occurring. Check atoms and total charge, not only the word attached to an electrode.

Product questions require phase and material. Molten NaCl can give Na at the cathode and Cl₂ at the anode. Concentrated aqueous brine under chlor-alkali conditions gives H₂ at the cathode, Cl₂ at the anode and NaOH in solution. Aqueous CuSO₄ with an inert anode can plate copper and oxidize water; with a copper anode it can transfer copper from one electrode to the other. A problem lacking phase or electrode material may need an assumption stated rather than a single unqualified prediction.

For charge calculations, multiply amperes by seconds for coulombs, divide by approximately 96,485 C mol⁻¹ for electron moles, then divide by the electron coefficient in the desired half-reaction. A 2.00 A current for 965 s passes 1930 C, about 0.0200 mol electrons. Cu²⁺ reduction could form 0.0100 mol Cu ideally; Ag⁺ reduction could form 0.0200 mol Ag ideally. Their atom amounts differ because the electron coefficients differ.

If current efficiency is 80%, multiply the theoretical target amount by 0.80, assuming the efficiency refers to that reaction and no separate recovery loss is specified. Do not apply the factor twice. A mass measurement after product loss is not necessarily a direct faradaic-efficiency measurement.

Check each answer independently: charge units must be C, electron amount mol e⁻, product amount mol and product mass g. A reduction product cannot require more metal atoms than allowed by its electron balance. The complete cell's electrons cancel when half-reactions are combined, and elements and charges must balance.

Step-by-step reasoning

1. State cell type, phase and electrode materials. 2. Write possible electrode half-reactions and choose the stated model. 3. Balance atoms, charge and exchanged electrons. 4. For amounts, calculate I × seconds, then divide by F. 5. Convert electron moles to product and apply any stated efficiency.

Visual explanation

Draw a branching route: cell description → electrodes and half-reactions → overall products. A second branch starts at current and time → charge → electron moles → product moles → mass. Join the branches at the electron coefficient.

Real-world analogy

A recipe problem needs both a count of ingredients available and a recipe specifying parts per dish. Electrical charge supplies the electron inventory; the half-reaction is the recipe connecting electrons to products.

Real-world example

A plating lab may measure current and time, weigh the cathode and compare expected and actual mass. The comparison tests reaction selectivity, handling and the validity of the chosen half-reaction model.

Why?

Why choose the electrode reaction before dividing electron moles by a number? The half-reaction tells whether one, two or more electrons are required per target atom or molecule.

Common misconception

“An electric charge directly tells grams of product.” It gives electron amount only after division by F; a balanced chemical ratio and molar mass are still required.

Worked example

An ideal Cu²⁺ plating cell runs at 1.93 A for 1000 s. Q = 1930 C, so n(e⁻) ≈ 1930/96,485 = 0.0200 mol. Cu²⁺ + 2e⁻ → Cu gives 0.0100 mol Cu, or about 0.635 g at 63.5 g mol⁻¹. At 75% current efficiency, expected electrochemical deposit is about 0.476 g. If a copper anode supplies the matching oxidation, it ideally loses 0.0100 mol Cu only if all anode current also follows that pathway.

Quick check

1. Which quantity comes immediately after Q = It in a charge-to-product calculation? Answer: Divide charge by the Faraday constant to obtain moles of electrons before applying a product half-reaction.

Exam focus

Write phase and electrode assumptions. Keep electrical conversions and chemical coefficients on separate lines so a factor-of-two error is visible.

Advanced insight

Current, potential, concentration and surface state can vary during a run. A simple constant-current, single-reaction problem is a controlled model of a more complex electrochemical system.

Summary

Electrochemistry problems combine reaction labels with charge accounting. Choose products from the actual cell, balance half-reactions, then convert I and t to electrons and products. Check units, conservation and efficiency assumptions.

Practice questions

1. A 1.00 A current runs for 100 s. How much charge passes? Answer: Q = It = 1.00 × 100 = 100 C. 2. Which electrode reaction deposits copper? Answer: Cathode reduction Cu²⁺ + 2e⁻ → Cu deposits copper metal. 3. Why is aqueous NaCl not treated like molten NaCl for cathode product prediction? Answer: Water is present and can be reduced to hydrogen, whereas molten NaCl has no water competitor.