Electromotive force (EMF) measurements have a broad range of applications in various fields, including chemistry and physics. The electrochemical series, an arrangement of elements in order of their standard electrode potentials, can be determined through EMF measurements. Elements with lower standard potentials can reduce ions of elements with higher standard potentials.
The standard cell potential, E°, allows for the calculation of the standard reaction Gibbs energy, ΔG°, and the equilibrium constant, K°. Once E° has been calculated, these values can be derived using the relation that connects them to the standard potential of a cell. Equilibrium constants derived from cell EMF measurements encompass redox K° values, solubility products, dissociation constants for complex ions, the ionization constant of water, ionization constants for weak acids, and constants for ion-pair-formation equilibria.
The temperature coefficient of the standard cell potential, ∂E°cell/∂T, provides information about the standard entropy change, ΔS°, of the cell reaction. The standard enthalpy change, ΔH°, can be derived from the relationship ΔG° = ΔH° − TΔS°, where T is the absolute temperature.
Once the standard potential of an electrode in a cell is known, it can be used to determine mean activity coefficients by measuring the cell potential with the ions at the concentration of interest.
Finally, the transport or transference number t can be determined from EMF measurements, and it depends on ion migration. The ratio of the EMFs of two concentration cells, one with transference and the other without transference, gives the transference number of the anion or cation, depending on whether the end electrodes are reversible with respect to the cation or anion, respectively.
In addition to these applications, EMF measurements play a key role in determining pH, utilizing electrodes such as hydrogen, quinhydrone, or glass electrodes. According to the Nernst equation, the potential of a hydrogen electrode relies on the pH of the solution. This pH can be accurately assessed by combining the hydrogen electrode with a reference electrode, like the calomel electrode, and measuring the EMF of the cell potentiometrically.
The standard cell potential helps find out the Gibbs free energy change and then the equilibrium constant of the reaction using thermodynamic relationships.
The temperature coefficient of the standard cell potential, its temperature derivative, allows calculation of the standard entropy change of the cell reaction; combining this with Gibbs energy and its relation to EMF yields the standard reaction enthalpy.
The difference between the cell potential and its standard value reflects the ions’ activity coefficient, which accounts for non-ideal interactions and modifies effective concentrations.
The electrochemical series ranks metals by standard electrode potentials, showing that a metal can reduce ions of metals placed above it; for example, zinc cannot reduce magnesium ions but can reduce hydrogen ions.
The anion transference number, describing anion current and its effect on concentration and potential gradients, can be derived from the EMF ratio of two concentration cells—one with transference and one without—provided the end electrodes are cation-reversible.
But, if the end electrodes are anion-reversible, this EMF ratio yields the cation’s transference number.
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