Topic 16 · page 3 of 19

From a reaction to a cell

Split a redox reaction into two half-cells and the electrons have to travel through a wire, so you can measure a voltage. Standard electrode potential is that voltage for one half-cell, under fixed conditions.

13 min read16.216.3(i)16.3(ii)16.3(iii)

Read first: Building redox equations

One beaker or two 16.2

Put a zinc strip into blue copper(II) sulfate solution and the reaction happens at once. Copper coats the zinc, the blue fades, and the solution warms up slightly.

Electrons jump straight from zinc atoms to copper(II) ions. The energy is released mainly as heat, which is no use for electrical work.

Now split the same chemistry into two beakers. Each half-cell holds one half-equation:

The electrons now have nowhere to go except along the wire joining the two metals, so they do electrical work on the way. The two solutions must also be connected, or charge would build up and stop the current. A salt bridge does this: ions move through it, but the solutions do not mix.

The voltmeter reads about 1.10 V. Zinc becomes thinner, so zinc is being oxidised: it is the negative electrode. Copper is being reduced: it is the positive electrode. Electrons leave the cell from the negative electrode.

Standard electrode potential 16.2

There are hundreds of useful redox couples. Building every possible pair would be endless, so chemists measure each one once, against one fixed reference, and list the results. For any new pair you then look up two numbers.

The next page describes the standard hydrogen electrode. Here you only need to know that it is the fixed reference.

The sign of $E^\ominus$ is the polarity of the half-cell compared with the hydrogen electrode:

• Copper is the positive electrode when joined to hydrogen, so $E^\ominus$(Cu²⁺/Cu) = +0.34 V.
• Zinc is the negative electrode when joined to hydrogen, so $E^\ominus$(Zn²⁺/Zn) = −0.76 V.

A couple is named with the oxidised form first: Cu²⁺/Cu, Zn²⁺/Zn, Fe³⁺/Fe²⁺.

Standard conditions 16.3(i)16.3(ii)16.3(iii)

An electrode potential depends on how much of each species is present, so it only means something if the conditions are fixed. They are:

Pure solids such as copper or zinc strips need no stated concentration, and a solution with no gas needs no pressure. A half-cell that has a gas, such as the hydrogen electrode, needs all three.

Check yourself

1. Why does zinc in copper(II) sulfate solution not give an electric current when the two are in one beaker?
2. What is the job of the salt bridge?
3. Define standard electrode potential.
4. State the three standard conditions for measuring an electrode potential.
5. A half-cell is joined to a standard hydrogen electrode and is the negative terminal. Is its $E^\ominus$ positive or negative? Explain.