About this experiment
A galvanic cell turns a chemical reaction into electricity. Put zinc in zinc sulfate and copper in copper sulfate, join the metals with a wire and the solutions with a salt bridge, and electrons flow from the zinc to the copper. Each metal has its own tendency to lose electrons, measured by its electrode potential; the difference between the two gives the voltage. Pair different metals, change the concentrations and the temperature, then connect a load and watch the cell run down as the reaction proceeds.
Learning goals
- Identify the anode, the cathode and the direction of electron flow in a galvanic cell
- Calculate the standard cell potential from standard electrode potentials
- Write the half-equations and the overall equation for a cell
- Use the Nernst equation to explain how concentration and temperature change the voltage
- Explain the role of the salt bridge and why a cell eventually goes flat
How to use it
- 1With zinc and copper at 1 M, read the voltmeter. Check it against E°(Cu²⁺/Cu) - E°(Zn²⁺/Zn).
- 2Swap the copper for silver, then for magnesium on the other side. Which pair gives the biggest voltage, and which electrode is the anode each time?
- 3Go back to zinc and copper. Lower the copper ion concentration to 0.01 M. How much does the voltage change? Compare with the Nernst equation.
- 4Put copper on both sides with 1 M and 0.001 M. Why is there a voltage when both electrodes are the same metal?
- 5Press Connect a load and discharge. Watch the electrodes, the colour of the copper solution and the voltage graph. When does the voltage drop quickly?
- 6Make the load resistance smaller. How does that change the current and the time the cell lasts?
The model
Assumptions
- Electrode potentials use standard values at 25 °C and the Nernst equation with concentrations in place of activities.
- Each beaker holds 100 mL. The salt bridge has a resistance, taken together with the solutions as an internal resistance of 2 Ω.
- The current cannot be larger than diffusion can supply: it is limited to 2 A for each mol/L of ions at the cathode. This is why the voltage falls quickly once the cathode ions run out.
- Side reactions, such as zinc reacting directly with the acid in the solution, are ignored. Time runs at one hour per second while the cell discharges.
Try this in class
- Q1Which metal in the list is the strongest reducing agent? How can you tell from the electrode potentials?
- Q2Write the overall equation for a cell made of magnesium and silver. What is its standard cell potential?
- Q3Why does a cell stop working if the salt bridge is removed?
- Q4Why does the voltage of a zinc and copper cell barely change during most of its life, then drop suddenly?
- Q5A concentration cell has the same metal on both sides. Which side is the cathode, and what happens to the two concentrations as it discharges?
Frequently asked questions
Why is the anode negative in a galvanic cell, when it is positive in electrolysis?
The anode is always where oxidation happens. In a galvanic cell the oxidation releases electrons into the metal, so the anode becomes the negative terminal. In electrolysis an external supply pulls electrons out of the anode, so it is connected to the positive terminal.
What does the salt bridge do?
It lets ions move between the solutions to keep each one electrically neutral, without letting them mix. Negative ions move towards the anode and positive ions towards the cathode.
Why does the voltmeter reading fall when a load is connected?
With current flowing, part of the EMF is used to push the current through the internal resistance of the cell. The voltage across the load is the EMF minus I times r.