Lab, Electrolysis & Faraday's Law (Copper Coulometer) · AP Chemistry

In an electrolytic cell you push electric current through a solution to drive a reaction that would not happen on its own. Here you plate copper metal onto a clean electrode from a copper(II) sulfate solution, measure the current and the time, and use Faraday's Law to predict the mass of copper deposited: then weigh the electrode and compare your prediction to reality.

CED 9.7 (Electrolysis & Faraday's Law) + 9.10 (Cell Potential & Electrolysis) · Unit 9. ~1 block · POST-EXAM lab: run after the Unit 9 test. FORMATIVE category: scored on the shared AP Chem lab rubric.

Background CED 9.7 · 9.10

A galvanic (voltaic) cell uses a spontaneous redox reaction to produce electricity. An electrolytic cell does the opposite: an external power supply forces a non-spontaneous reaction by pushing electrons. In both, oxidation happens at the anode and reduction happens at the cathode: but in an electrolytic cell the power supply makes the cathode the negative terminal, where cations are reduced.

Your cell: two copper strips dipped in acidified 0.5 M CuSO₄(aq), wired to a DC supply. At the cathode (the strip wired to the −/negative terminal) copper ions are reduced and plate out as solid copper; at the anode the copper metal is oxidized and dissolves, replenishing the Cu²⁺ in solution:

Cathode (reduction): Cu²⁺(aq) + 2 e⁻ → Cu(s)

Anode (oxidation): Cu(s) → Cu²⁺(aq) + 2 e⁻

Because the same current flows everywhere in a series circuit, every electron that leaves the anode arrives at the cathode, so the copper that plates out is set entirely by how much charge you pass.

Faraday's Law. The amount of substance produced at an electrode is proportional to the electric charge passed. Charge is current × time:

Q = I · t (coulombs = amps × seconds)

n(e⁻) = Q / F, F = 96485 C/mol e⁻ (1 mol of electrons carries 96485 C)

n(Cu) = n(e⁻) / 2 (2 electrons per Cu²⁺) → mass Cu = n(Cu) × 63.55 g/mol

The 2 comes straight from the cathode half-reaction: it takes two electrons to deposit one copper atom. That stoichiometric link is the whole point of the lab.

Why acidify with a little H₂SO₄? The acid keeps the solution conducting and suppresses hydroxide/oxide films on the copper, so the plating stays clean and even. It does not change the copper half-reaction.

Safety read before you start

Copper(II) sulfate solution, irritant/toxic. CuSO₄ irritates skin and eyes and is harmful if swallowed. Splash goggles + nitrile gloves whenever you handle the solution or a wet electrode.

Dilute sulfuric acid, corrosive. The solution is acidified with a little H₂SO₄. Wipe spills immediately and wash hands after every transfer, even if you think none spilled.

Low-voltage DC, but treat it with respect. Keep the alligator leads and your hands dry. Never let the two leads or electrodes touch each other: shorting the supply makes the wires hot. Switch the supply off before you connect or disconnect anything.

PPE, every student, the whole period: splash goggles, lab coat or apron, nitrile gloves when handling solution or electrodes. Tie back hair. No food or drink.

Disposal: all blue (Cu²⁺) solution and copper rinse goes in the labelled copper waste container, never down the sink. Spent ethanol rinse to the labelled solvent waste. Spills: alert Dr. Mulinta, neutralize acid with the sodium bicarbonate kit at the bench.

Materials & Equipment low-gear setup

Everything is everyday school chem stock plus a DC supply and an ammeter. The class shares power supplies and meters at the bench stations.

Per group: DC power supply (or 6 V battery) · ammeter (or multimeter set to DC amps, wired in series) · 2 copper electrode strips · 3 alligator-clip leads · ~0.5 M CuSO₄(aq) acidified with H₂SO₄ · 250 mL beaker · stopwatch / phone timer · fine sandpaper or steel wool (to clean electrodes) · wash bottle of distilled water · small bottle of ethanol (or acetone) for the final rinse · paper towel.

Shared at stations: analytical balance (± 0.001 g) · copper waste container · solvent waste container · sodium-bicarbonate spill kit.

Mrs. Kathy (lab tech) will pre-set: the acidified ~0.5 M CuSO₄ in a 250 mL beaker at each station; the two copper electrodes; the DC power supply wired in series with the ammeter (leads off / supply off) so you only attach the electrodes; the balance zeroed; the copper/solvent waste containers and bicarbonate spill kit. Check with Dr. Mulinta before switching the supply on.

Procedure clean → mass → plate → re-mass

Part 1: Prepare and mass the cathode

Choose one copper strip to be your cathode (the one you will plate onto). Lightly clean both faces with fine sandpaper or steel wool until bright, then rinse with distilled water and dry fully with a paper towel. Handle it by the top edge only from now on, fingerprints add mass and oils block plating.
Mass the clean, dry cathode on the analytical balance to ± 0.001 g. Record as mass of cathode (before).
Lightly clean the second strip (the anode) the same way so it dissolves evenly. It does not need to be massed.

Part 2: Assemble the cell

With the power supply OFF, clip the cathode to the negative (−) terminal side of the circuit and the anode to the positive (+) side. The ammeter is already wired in series, confirm with Dr. Mulinta.
Suspend both electrodes in the acidified 0.5 M CuSO₄ so they are well submerged but not touching each other and not touching the bottom. Keep them parallel, a few cm apart.
Have Dr. Mulinta check the circuit before you switch on.

Part 3: Run the electrolysis

Switch the supply ON and immediately start the stopwatch. Adjust the supply so the ammeter reads a steady ~0.50 A (any steady value between about 0.3 and 0.6 A is fine, record the actual reading).
Watch the ammeter through the run and nudge the supply to hold the current steady. Record the current you are holding as I. If it drifts, note the average value you actually maintained.
Run for a measured time of about 20–30 minutes (e.g. 1800 s). When time is up, switch the supply OFF and stop the watch at the same instant. Record the elapsed time t in seconds.
Disconnect the cathode. Lift it out without scraping it on the beaker rim.

Part 4: Rinse, dry, and re-mass the cathode

Gently dip-rinse the cathode in distilled water to remove clinging CuSO₄ solution, do not rub the plated surface, the fresh copper flakes off easily.
Follow with a quick rinse of ethanol (or acetone); this displaces water so the strip dries fast without spotting.
Dry gently in air (or a very light pat at the top edge only). Make sure it is completely dry.
Mass the dry plated cathode to ± 0.001 g. Record as mass of cathode (after).
Return all solution and rinses to the labelled waste containers. Wipe down your station.

Data Table fill in lab

Current held, I (A)	__________ A
Elapsed time, t (s)	__________ s
Mass of cathode, before (± 0.001 g)	__________ g
Mass of cathode, after (± 0.001 g)	__________ g
Mass deposited (measured) = after − before	__________ g

The mass deposited is what you compare against Faraday's prediction. Keep the before/after to the full ± 0.001 g, the deposited mass is a small difference of two larger numbers, so every milligram matters.

Analysis Questions show all work · units · sig figs

Answer on your own paper / in your typed report. Show every formula and substitution, full credit needs the math, not just the answer.

Charge passed. Using your measured current I and time t, calculate the total charge Q in coulombs (Q = It).
Moles of electrons. Convert the charge to moles of electrons using n(e⁻) = Q/F, F = 96485 C/mol e⁻.
Moles of copper. From the cathode half-reaction Cu²⁺ + 2 e⁻ → Cu, find the moles of copper deposited.
Predicted mass. Convert moles of copper to a predicted mass in grams (M = 63.55 g/mol). This is your Faraday prediction.
Measured mass. State the mass actually deposited (after − before) from your data table.
Percent error. Compare prediction and measurement: % error = |measured − predicted| ÷ predicted × 100%.
Electrode identity. Which electrode gained mass and which lost mass? Tie each to its half-reaction (reduction at the cathode, oxidation at the anode) and state which terminal of the supply each was wired to.
Conservation of charge. Explain why the moles of electrons that reduced Cu²⁺ at the cathode must equal the moles of electrons released by oxidation at the anode. What does that say about the change in the solution's overall [Cu²⁺] during the run?
Reasoning / scaling. If you had run the same current for twice as long, what mass of copper would you predict? If instead you doubled the current and kept the time the same, how does the deposited mass change? Justify with Q = It.
Error analysis. Identify one source of error that would make your measured deposited mass read too HIGH relative to the Faraday prediction, and one that would make it read too LOW. For each, explain the direction of the effect through Q = It and the relation mass = (It/F)/2 × 63.55.

Lab · Electrolysis & Faraday's Law