 |
 |
1 of 20 |  |
 |
 |
R-1
EXPERIMENT R: THE NERNST EQUATION AND ITS
APPLICATIONS
OBJECTIVES
Experiment Q dealt with half-cells at standard conditions. In this experiment measurements will be
made using half-cells at nonstandard conditions. Measurements of cell voltages on a series of cells
with decreasing concentrations will be compared to values predicted by the Nernst equation. The
function of the reaction quotient in the Nernst equation will also be examined. This will suggest a
method for obtaining an equilibrium constant for a redox reaction. Finally, cell voltage measurements
and the Nernst equation will be used to determine metal ion concentrations for some test solutions.
These concentrations will then be used to calculate the solubility product of a salt and to calculate the
formation constant of a complex ion.
THEORETICAL CONSIDERATIONS
I. The Nernst Equation
In Experiment Q (Redox Reactions) you measured the difference between half-cell potentials using a
voltmeter. You then used these measurements to produce a displacement series. A number of
standard electrode potentials (half-cell potentials), E°, written as reduction processes have been
tabulated in the Chemistry Data Sheet. These values are for standard cells (concentrations of the
dissolved species are 1.00 M and the pressures of gases are
1 atmosphere). The potential difference (voltage) across any standard cell which is constructed from
two half-cells can be calculated from the tabulated values of E° for the half-cells.
When the concentrations of dissolved species are not 1.00 M and the pressures of gases differ from 1
atmosphere, the cells are nonstandard; the potential difference or voltage across a nonstandard cell
cannot be determined directly from the E° values of standard half-cells. The potentials of the half-cells
vary with the concentrations and pressures of species involved as described by the Nernst equation,
which at ~25 °C has the form:
