In voltaic cell diagrams, voltmeters are shown attached to the external circuit. The voltmeter measures voltage or the standard cell potential.

The standard cell potential is the maximum electric potential difference of a cell; it represents the energy difference between the cathode half-cell and the anode half-cell. It can be measured empirically using a voltmeter as the cell operates or it can be calculated theoretically using reduction potentials.

\( \mathrm { E^\circ_{net} = E^\circ_{cathode} - E^\circ_{anode} } \)

For example, consider the theoretical standard cell potential of the silver-copper voltaic cell (the cell featured in the Lesson 1.2 animation).

First, for each half-reaction, find the half-reactions and reduction potentials on the "Table of Selected Standard Electrode Potentials" in your data booklet (page 7).

Next, find the difference between the reduction potentials of these two half-cells.

\( \mathrm { E^\circ_{net} = E^\circ_{cathode} - E^\circ_{anode} } \)

\( \mathrm { E^\circ_{net} = +0.80V - (+0.34 V) } \)

The positive cell potential indicates that the net reaction is spontaneous - a requirement for all voltaic cells.

It is impossible to determine the reduction potential for a half-cell in isolation because electron transfer occurs only when two half-cells are connected. How, then, did scientists determine reduction potentials for the half-cells indicated on the table?

Actually, scientists determined reduction potentials by selecting a reference half-cell arbitrarily and assigning it a value of 0.00 V.

All other half-cells were connected to the standard reference hydrogen half-cell (one at a time) and the potential difference was recorded for each.

Note: Any standard half-cell could have been chosen as the "reference half-cell", which is point zero on the reduction potential scale. Of course, if a different half-cell had been selected as the reference half-cell, individual reduction potentials would all be different. However, the standard cell potentials would remain unaffected.

For example, what if the standard lithium cell were chosen as the reference half-cell with its reduction potential defined as 0.00 V? We would have to add 3.04 V to each reduction potential on the chart.

\( \mathrm { E^{\circ}_{cell} = +3.38V - (+2.28V) = +1.10V } \)

The original reduction potential would give the following results:

\( \mathrm { E^{\circ}_{cell} = +0.34V - (-0.76V) = +1.10V } \)

Sometimes the measured cell potential is not the same as the predicted standard cell potential. Non-standard conditions for concentration, temperature, or pressure along with the purity of the substance and presence of oxide coatings or other circumstances influence the empirical cell potential.

Check Your Understanding

Complete the Self-Check Questions below. Click the link to check your work.

Self-Check Question 1

Comment on the following observation: "A higher cell potential is observed with reactants that are greatly separated on the table of half-reactions."

Self-Check Question 2

Calculate the reduction potential for the following half-cells if the iodine half-cell had been chosen as the reference half-cell instead of the hydrogen half-cell.

1. \( \mathrm { Cu^{2+}(aq) + 2e^- \rightarrow Cu(s) } \)

2. \( \mathrm { Zn^{2+}(aq) + 2e^- \rightarrow Zn(s) } \)

Self-Check Question 3

Calculate the standard cell potential for the image shown below using hydrogen as the reference cell. Recalculate the standard cell potential using iodine as the reference cell.