Electrochemical Cells: Galvanic and Electrolytic, Grade 12
The anode is negative in a galvanic cell and positive in an electrolytic cell. Oxidation happens at the anode in both. If you have those two sentences straight, most of this chapter is already yours.
This page covers galvanic and electrolytic cells, oxidation numbers, the salt bridge, cell notation, standard electrode potentials, calculating Eθcell, predicting electrolysis products, and where all of it is used in South Africa.
The sign flip, and why it trips everyone
The definitions of anode and cathode never change. The signs do.
| Galvanic cell | Electrolytic cell | |
|---|---|---|
| Anode | Negative | Positive |
| Cathode | Positive | Negative |
| Oxidation happens at the | Anode | Anode |
| Reduction happens at the | Cathode | Cathode |
Learn the bottom two rows. They are true in every cell ever built. The top two are consequences, and you can work them out rather than memorise them.
Why the signs flip
In a galvanic cell the chemistry pushes the electrons. Metal atoms give up electrons at the anode, so electrons pile up there. A pile of electrons is a negative terminal.
In an electrolytic cell a battery pushes the electrons. The battery's positive terminal is connected to the anode and pulls electrons out of it. Being connected to a positive terminal makes the anode positive.
So the anode is where electrons leave the electrode into the wire, in both cases. What changes is whether the chemistry or the battery is doing the pushing, and that is what sets the sign.
The old mnemonics still work and they are worth having: OIL RIG for oxidation is loss, reduction is gain, and AN OX and RED CAT for oxidation at the anode, reduction at the cathode.
What an electrochemical cell is
An electrochemical cell is any arrangement that connects a redox reaction to an electric circuit. Every one of them has three things:
| Part | What it does |
|---|---|
| Two electrodes | Conductors where the half-reactions happen. Oxidation at the anode, reduction at the cathode |
| An electrolyte | A solution or molten salt that conducts by moving ions. This is the non-metallic part of the circuit |
| An external circuit | Wire, where the current is moving electrons |
Current flows all the way round, but it is carried by two different things. Electrons in the wire, ions in the solution. That is worth saying out loud because exam questions test it directly.
Galvanic against electrolytic, in full
| Galvanic (voltaic) cell | Electrolytic cell | |
|---|---|---|
| Energy conversion | Chemical to electrical | Electrical to chemical |
| The reaction is | Spontaneous | Non-spontaneous |
| Eθcell | Positive | Negative |
| Power supply | None. The cell is the supply | Required. An external battery |
| Anode | Negative | Positive |
| Cathode | Positive | Negative |
| Electrolytes | Usually two, one per half-cell | Usually one |
| Salt bridge | Yes, to keep the two half-cells separate but connected | No. There is only one compartment |
| Everyday example | Any battery. Torch, phone, car | Electroplating, refining aluminium, making chlorine |
A rechargeable battery is both, at different times. Galvanic while it powers your phone, electrolytic while it charges. That is the cleanest way to see that the two are the same physics running in opposite directions.
Oxidation, reduction and oxidation numbers
A redox reaction is an electron transfer reaction. Oxidation is loss of electrons, reduction is gain, and the two always happen together because electrons have to come from somewhere and go somewhere.
| Term | What it does | What happens to it |
|---|---|---|
| Reducing agent, or reductant | Donates electrons | Is oxidised |
| Oxidising agent, or oxidant | Accepts electrons | Is reduced |
The agent is always the opposite of what happens to it, and learners get this backwards constantly. The reducing agent does the reducing to something else, so it must give electrons away, so it is oxidised.
Assigning oxidation numbers
Work down this list and stop when you can.
| Rule | Example |
|---|---|
| An element on its own is 0 | Cu, O2, S8 are all 0 |
| Fluorine is always −1 | |
| Hydrogen is +1, except with a metal, where it is −1 | +1 in H2O, −1 in NaH |
| Oxygen is −2, except −1 in peroxides and positive with fluorine | −2 in H2O, −1 in H2O2, +2 in OF2 |
| A simple ion takes its own charge | Na+ is +1, Cl- is −1 |
| Everything adds to the overall charge | 0 for a neutral compound, the ion charge for a polyatomic ion |
If an oxidation number changes during a reaction, it is a redox reaction. If none change, it is not, and it is probably acid-base.
The salt bridge, and what it is actually for
Three jobs, and exam papers ask for all three.
- It keeps the two electrolytes from mixing. If they mixed, the reaction would happen directly in the beaker and no electrons would go through the wire
- It completes the circuit. Ions moving through the bridge are the current in the non-metallic part of the circuit
- It keeps both half-cells electrically neutral. Anions move towards the anode, cations towards the cathode
The third one is the one that gets left out. Without it, positive charge builds up in the anode compartment as metal ions form, the build-up opposes further oxidation, and the cell stops within seconds.
A salt bridge is usually a U-tube of potassium nitrate solution plugged with cotton wool, or just a strip of filter paper soaked in it. The salt has to be one that will not react with either electrolyte, which is why potassium nitrate is the standard choice.
Cell notation
Shorthand for a galvanic cell, and it follows one order every time:
anode / anode electrolyte // cathode electrolyte / cathode
| Symbol | Means |
|---|---|
| / single slash | A boundary between two phases, such as solid metal and its dissolved ions |
| // double slash | The salt bridge |
So a zinc-copper cell is written Zn / Zn2+ // Cu2+ / Cu
The anode always goes on the left. If no concentrations are shown, they are assumed to be standard, 1 mol·dm-3. Concentrations go in brackets after the ion when they are needed.
Standard electrode potentials
You cannot measure a single half-cell. A voltmeter needs two connections, so every measurement is a comparison. So chemistry picked one half-cell, called it zero, and measured everything against it.
The standard hydrogen electrode is defined as exactly 0,00 V. It is hydrogen gas at 101 kPa bubbled over platinum in 1 mol·dm-3 acid at 25 °C.
| Standard conditions | Value |
|---|---|
| Concentration of all solutions | 1 mol·dm-3 |
| Pressure of any gas | 101 kPa |
| Temperature | 25 °C |
Connect any half-cell to the hydrogen electrode and the voltmeter reading is that half-cell's standard electrode potential.
| Half-cell | Eθ | Meaning |
|---|---|---|
| Ag+ + e- → Ag | +0,80 V | Gains electrons more readily than hydrogen |
| Cu2+ + 2e- → Cu | +0,34 V | Same, less strongly |
| 2H+ + 2e- → H2 | 0,00 V | The reference |
| Zn2+ + 2e- → Zn | −0,76 V | Gives electrons up more readily than hydrogen |
| Al3+ + 3e- → Al | −1,66 V | Same, much more strongly |
More positive means a stronger tendency to be reduced, which means a better oxidising agent. More negative means a better reducing agent.
Use the table on your own data sheet and nothing else. Published tables round the last digit differently, and the marker is working from the data sheet in the exam pack.
Calculating Eθcell
Eθcell = Eθcathode − Eθanode
which is the same thing as Eθreduction − Eθoxidation, because reduction happens at the cathode.
Work out which half-cell is which before you subtract. In a galvanic cell, the half-cell with the more positive electrode potential is the cathode. That is the whole decision, and everything else follows from it.
Never change the sign of a value when you reverse a half-reaction for this formula. The subtraction handles it. Flipping the sign as well is the single most common calculation error in the topic and it doubles the error.
Worked example: a zinc-silver cell
Zn / Zn2+ // Ag+ / Ag
- Eθ(Ag+/Ag) = +0,80 V and Eθ(Zn2+/Zn) = −0,76 V
- Silver is more positive, so silver is the cathode and zinc is the anode
- Cathode, reduction: Ag+ + e- → Ag
- Anode, oxidation: Zn → Zn2+ + 2e-
- Balance the electrons. The silver half-reaction needs doubling: 2Ag+ + 2e- → 2Ag
- Net cell reaction: Zn + 2Ag+ → Zn2+ + 2Ag
- Eθcell = 0,80 − (−0,76)
- Eθcell = 1,56 V
Positive, so the reaction is spontaneous and this is a working battery.
Worked example: three electrons against two
Al / Al3+ // Cu2+ / Cu
- Eθ(Cu2+/Cu) = +0,34 V and Eθ(Al3+/Al) = −1,66 V
- Copper is more positive, so copper is the cathode
- Cathode: Cu2+ + 2e- → Cu, × 3
- Anode: Al → Al3+ + 3e-, × 2
- Six electrons on each side now. Two and three go to six
- Net: 2Al + 3Cu2+ → 2Al3+ + 3Cu
- Eθcell = 0,34 − (−1,66) = 2,00 V
Balancing the electrons does not change Eθcell. Cell potential is not multiplied when you scale a half-reaction, and learners multiply it constantly. It is a potential difference, not an amount.
Worked example: identifying an unknown metal
A galvanic cell is built with a copper cathode and an unknown metal X as the anode. The voltmeter reads 1,10 V under standard conditions. Identify X.
- Eθcell = Eθcathode − Eθanode
- 1,10 = 0,34 − Eθanode
- Eθanode = 0,34 − 1,10 = −0,76 V
- Reading the table, that is zinc
This is the standard hard question and it is only one line of algebra. Rearrange first, then look up.
Electrolysis and predicting the products
Electrolysis is using electricity to force a non-spontaneous redox reaction to happen. Pass a current through a molten salt or a solution and it decomposes.
A molten salt is straightforward, because there is only one thing to reduce and one thing to oxidise.
Molten lead(II) bromide:
- Cathode, reduction: Pb2+ + 2e- → Pb
- Anode, oxidation: 2Br- → Br2 + 2e-
- Overall: PbBr2 → Pb + Br2
A solution is harder, because water joins in
In an aqueous solution the water itself can be reduced or oxidised, and it competes with the dissolved ions.
- Water reduced: 2H2O + 2e- → H2 + 2OH-
- Water oxidised: 2H2O → O2 + 4H+ + 4e-
So there are usually two candidates at each electrode and you have to work out which wins. List every possible half-reaction with its potential, work out the cell potential for each combination, and the one needing the smallest potential difference is the one that happens.
And in practice you always have to supply more than the calculated voltage. The extra is called the overpotential, it cannot be predicted, and it is why a real electrolysis needs more voltage than the arithmetic says.
Three patterns worth knowing, because they cover most school examples:
| Situation | What happens |
|---|---|
| A very reactive metal ion in solution, like Na+ or K+ | Water is reduced instead. You get hydrogen, not the metal |
| Chloride ions present | Chloride is oxidised in preference to water, so you get chlorine |
| Sulfate or nitrate ions present | Neither is oxidised. Water is, so you get oxygen |
Where this fits in the curriculum
| Subject | Physical Sciences |
|---|---|
| Grade | 12 |
| Term | 3 |
| Topic | Chemical change, electrochemical reactions |
| Status | Examinable theory with two supporting experiments. No formal assessment attached |
| Paper | Chemistry, Paper 2 |
It is one of the largest chapters in the Grade 12 book and it reliably carries a full question in Paper 2, usually built around one cell diagram with six or seven parts hanging off it.
The practical you can actually run
Build galvanic cells from four metals and measure what each pair produces. It is the best practical in the chapter because the prediction and the measurement can be compared directly.
Apparatus
| Item | Qty | Note |
|---|---|---|
| Metal electrode strips: copper, zinc, iron, magnesium | 1 each | Clean them with sandpaper first |
| 100 cm³ beakers | 4 | One per half-cell |
| 1 mol·dm-3 nitrate solution of each metal | 75 cm³ each | The salt must match the electrode |
| Glass U-tube and cotton wool | 1 | For the salt bridge. Filter paper strips work too |
| 1 mol·dm-3 potassium nitrate | 50 cm³ | The salt bridge solution |
| Voltmeter or multimeter, 0 to 3 V | 1 | |
| Leads with crocodile clips | 2 |
Method
- Sand each metal strip clean. An oxide layer will give you a low or unstable reading
- Stand each strip in 75 cm³ of its own nitrate solution, one metal per beaker
- Fill the U-tube with potassium nitrate solution and plug both ends with cotton wool. No air bubbles
- Bridge two beakers with the U-tube and connect the two strips to the voltmeter
- Note the reading and which strip is on the positive terminal. That one is the cathode
- Repeat for every pair, then rank the metals
What you should see
Every pair gives a different voltage, and the biggest one comes from the two metals furthest apart on the table. The readings will come in a little under the calculated values, which is normal and worth discussing rather than hiding.
Predict each voltage before you measure it. A practical where the class has already committed to a number is a completely different lesson from one where they write down whatever the meter says.
Where South Africa uses this
| Process | What happens |
|---|---|
| Aluminium at Richards Bay | Alumina is dissolved in molten cryolite at about 950 °C and electrolysed with carbon electrodes. The cell runs at enormous current and the carbon anodes burn away and are replaced |
| The chlor-alkali industry | Electrolysis of brine gives chlorine, hydrogen and sodium hydroxide. About a tenth of the chlorine made goes into water purification |
| Electroplating | The object to be plated is the cathode, the plating metal is the anode, and the electrolyte carries that metal's ions |
| Refining copper | Impure copper is the anode and it dissolves; pure copper deposits on the cathode |
Electrolysis is enormously electricity-hungry, which in South Africa means coal, which is a fair thing for an exam answer to raise when it asks about environmental impact.
The mistakes that cost the marks
| The mistake | What to do instead |
|---|---|
| Getting the anode sign wrong | Negative in galvanic, positive in electrolytic. Oxidation is at the anode in both |
| Multiplying Eθ when you balance the electrons | Never. It is a potential difference, not an amount. Balance the equation, leave the voltage alone |
| Changing the sign of Eθ when reversing a half-reaction, and subtracting as well | Pick one. Using Eθcathode − Eθanode with table values as printed is the safe route |
| Saying the reducing agent is reduced | It donates electrons, so it is oxidised. The agent is always the opposite of what happens to it |
| Only giving two functions of the salt bridge | There are three, and the one usually missed is keeping the half-cells electrically neutral |
| Writing the cathode on the left in cell notation | Anode always on the left |
| Forgetting that current in the solution is carried by ions | Electrons in the wire, ions in the electrolyte. Electrons do not travel through the solution |
If the practical does not work
| What happens | Why |
|---|---|
| The reading is far below the calculated value | Oxidised electrodes. Sand every strip clean immediately before use. This is the commonest cause by a distance |
| The reading drifts downwards while you watch | Normal. The cell is discharging and the concentrations are changing. Take the reading quickly |
| You read zero | The salt bridge is not making contact, or there is an air bubble in the U-tube, or a crocodile clip is on the glass rather than the metal |
| The reading is negative | The leads are the other way round. Swap them, and note that the strip on the positive terminal is the cathode |
| Readings are erratic | The strips are touching the beaker wall or each other, or the solutions have been contaminated between pairs. Rinse the bridge between runs |
| The magnesium half-cell fizzes | Magnesium reacts slowly with water. Expected, and it is why magnesium readings drift fastest |
How the 50 marks are made up
| Section | Marks |
|---|---|
| Definitions and terminology | 8 |
| Galvanic against electrolytic, including the signs | 10 |
| Oxidation numbers and identifying redox | 6 |
| Half-reactions, net equation and cell notation | 12 |
| Eθcell calculations | 10 |
| The salt bridge and the practical | 4 |
No mark allocation is prescribed for this chapter. The split above is ours, weighted towards the equations and the calculations because that is where a Paper 2 question puts its marks.
Free worksheet and marking memo
Both free, no sign up, straight to the PDF.
- Learner worksheet, 50 marks, with multiple choice, the definitions, oxidation numbers, three full cell calculations, cell notation and the salt bridge
- Marking memorandum, with full working, the mark breakdown line by line and a note on the seven places learners drop marks
Related pages
- Redox reactions investigation. The Grade 11 practical this chapter is built on
- Internal resistance of a battery, Grade 12. What happens once your cell actually delivers current
- Titration, Grade 12. Also Chemistry Paper 2
Apparatus
This is one of the few chapters where we can supply almost everything.
Electrodes for student cells come in the metals this practical needs and are the cheapest way to equip a class. Carbon rod electrodes are what you want for electrolysis rather than galvanic cells, because carbon is inert and will not take part in the reaction.
The electricity range has the multimeters and the leads.
Buy more electrode strips than you think you need. They get sanded before every use, and a strip that has been through a few classes is noticeably thinner. They are consumables, not equipment.