Electrochemistry
That aluminum wire is not spliced to copper is no electrician’s whim: it is two volts of difference in the table of potentials, and with moisture that is a cell that eats the aluminum.
01Electrons that change owners
All of electrochemistry is an oxidation-reduction reaction in which the electrons are forced to pass through a wire instead of being handed over directly. That is the only difference between a nail that rusts and a battery: the path.
| Occurs in | What happens | Electrode | Sign in a galvanic cell |
|---|---|---|---|
| Oxidation | The species loses electrons | Anode | Negative |
| Reduction | The species gains electrons | Cathode | Positive |
The rule for not getting lost: oxidation always occurs at the anode, in a galvanic cell and in an electrolytic cell. What does get reversed between the two is the sign of the terminal, and that is where the most common confusion in the subject comes from.
In a galvanic cell the reaction is spontaneous and generates current: the anode is the negative terminal. In an electrolytic cell an external source forces a reaction that would not occur on its own —electroplating, anodizing, electrolysis— and there the anode is the positive terminal. Oxidation is still at the anode in both cases.
02Electrode potentials and the Nernst equation
Each half-reaction is assigned a standard potential measured against the hydrogen electrode, which by convention was assigned zero. The EMF of the cell is the difference, and it is always built by subtraction: cathode minus anode.
Choose the two metals and the cell solves itself: which one is oxidized, which one is reduced and how much voltage it delivers. The concentration control applies Nernst. Try the zinc–copper pair, which is the Daniell cell.
03Galvanic corrosion: the cell nobody meant to build
Two different metals in electrical contact with moisture between them form a cell, whether you want it or not. The less noble one becomes the anode and corrodes, and the greater the potential difference, the faster it happens.
| Pair in contact | \( \Delta E^{\circ} \) | What happens |
|---|---|---|
| Aluminum with copper | 2.00 V | Very bad: the aluminum gets eaten away. That is why aluminum wire is not spliced to copper without a bimetallic connector |
| Steel with copper | 0.78 V | Bad: the steel pits at the contact |
| Tin with copper | 0.48 V | Tolerable, and it is the pair of every solder joint on a trace |
| Gold over nickel | Protected | That is why a gold-plated contact has a nickel barrier underneath |
The same mechanism is used to advantage: a sacrificial anode of zinc or magnesium connected to a steel structure corrodes in its place. It is galvanic corrosion directed on purpose against a metal that does not matter.
The copper traces of a multilayer board and the columns that connect its layers are made by electroplating copper from a sulfate solution. The amount deposited is set by Faraday's law: \( m = \frac{QM}{nF} \), with \( F = 96485 \) coulombs per mole of electrons. The thickness is controlled by the current and the time, and by nothing else.
04Batteries
A battery is a cell with the reaction chosen to give the voltage and the energy that are needed. The voltage comes from the difference of potentials; the capacity, from how much mass of reactant there is.
| System | Voltage | Why |
|---|---|---|
| Alkaline Zn/MnO₂ | 1.5 V | Not rechargeable; it is the reaction from the stoichiometry unit |
| Lead-acid | 2.0 V per cell | Rechargeable and cheap; six cells give the 12 V of a car |
| Nickel–metal hydride | 1.2 V | Rechargeable, robust, with less energy per kilogram |
| Lithium-ion | 3.6 V | Lithium is the most electropositive metal and the lightest: both at once |
The last row explains why lithium prevailed. With \( E^{\circ} = -3.04 \) V it is the end of the table of potentials, so each cell delivers more than twice the voltage of an alkaline one; and since it is also the lightest solid metal, the energy per kilogram has no competition. It also explains its danger: that same reactivity is what makes a damaged cell catch fire.
05In the lab
Build the zinc–copper cell with a salt bridge and measure the open-circuit EMF. Compare with the theoretical 1.10 V and explain the difference. Then measure with a load and discuss the drop due to internal resistance.
Repeat the measurement, diluting the copper sulfate to one-tenth and one-hundredth. Plot the EMF against the logarithm of the concentration and verify the slope of 0.0592/n volts per decade.
Electroplate copper onto a clean steel part with controlled current. Weigh before and after, calculate the expected mass with Faraday's law and obtain the current efficiency.
06Common mistakes
- Believing that the anode is always the positive one. It is in the electrolytic cell; in the galvanic cell it is the negative one. Oxidation, however, is always at the anode.
- Adding the potentials of the two half-reactions. They are subtracted: cathode minus anode.
- Multiplying \( E^{\circ} \) when balancing. The potential is intensive; the coefficients do not change it.
- Confusing voltage with capacity. One comes from the potentials, the other from the mass of reactant.
- Putting metals that are far apart in the table in contact. That is a cell, and the less noble one corrodes.
- Forgetting the salt bridge. Without it the solutions become charged and the current stops right away.
07Self-assessment
Where does oxidation occur?
At the anode, always: in a galvanic cell and in an electrolytic cell. What changes between the two is the sign of the terminal.
What is the EMF of the Daniell cell?
1.10 V: \( +0.34 \) of copper minus \( -0.76 \) of zinc.
What happens to the EMF when the cell runs down?
\( Q \) grows until it equals \( K \) and the EMF falls to zero. That is a discharged battery.
Why is aluminum not spliced to copper?
Because 2.00 V separates them in the table: they form a cell and the aluminum, which is the anode, corrodes.
What does Faraday's law say?
That the mass deposited is \( m = QM/nF \): it depends only on the charge that passed, with \( F = 96485 \) C/mol.
Why does lithium give so much energy per kilogram?
Because it is the most electropositive metal in the table and also the lightest solid: a lot of voltage and little mass.
08Further reading
- Raymond Chang and Kenneth A. Goldsby. Chemistry. 12th ed. (Spanish edition, “Química”), McGraw-Hill, 2017. Chapter 19: balancing redox reactions, galvanic cells, Nernst, batteries, corrosion and electrolysis.
- Allen J. Bard and Larry R. Faulkner. Electrochemical Methods: Fundamentals and Applications. 2nd ed., Wiley, 2001. The reference of the field, for those who go on to electrochemical sensors or electroplating.
- Clyde F. Coombs and Happy T. Holden. Printed Circuits Handbook. 7th ed., McGraw-Hill, 2016. Copper electroplating of a multilayer board and surface finishes, with their corrosion problems.