9.7
Applications of EMF Measurements
Electromotive force (EMF) measurements have a broad range of applications in various fields, including chemistry and physics. The electrochemical seri…
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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Q1: How does standard cell potential relate to Gibbs free energy and equilibrium constant?
Standard cell potential (E°) directly determines the standard Gibbs free energy change (ΔG°) and equilibrium constant (K°) through thermodynamic relationships. These values can be derived using the connection between E° and ΔG°. EMF measurements provide equilibrium constants for redox reactions, solubility products, dissociation constants for complex ions, and ionization constants for weak acids.
Q2: What information does the temperature coefficient of cell potential provide?
The temperature coefficient of standard cell potential (∂E°cell/∂T) provides information about the standard entropy change (ΔS°) of the cell reaction. Combined with Gibbs energy relationships, it allows calculation of standard reaction enthalpy (ΔH°) using the equation ΔG° = ΔH° − TΔS°, where T is absolute temperature.
Q3: How does the electrochemical series predict redox reactions between metals?
The electrochemical series ranks metals by standard electrode potentials, showing that a metal can reduce ions of metals placed above it in the series. For example, zinc cannot reduce magnesium ions but can reduce hydrogen ions. This ordering allows prediction of which redox reactions are thermodynamically favorable.
Q4: What role do activity coefficients play in cell potential measurements?
The difference between actual cell potential and standard cell potential reflects the ions' activity coefficient, which accounts for non-ideal interactions and modifies effective concentrations. Once standard electrode potential is known, mean activity coefficients can be determined by measuring cell potential at specific ion concentrations.
Q5: How can transference numbers be determined from EMF measurements?
The anion transference number can be derived from the EMF ratio of two concentration cells—one with transference and one without—provided the end electrodes are cation-reversible. If electrodes are anion-reversible instead, this EMF ratio yields the cation's transference number. Transference numbers depend on ion migration rates.
Q6: How are pH measurements performed using EMF and electrode systems?
EMF measurements determine pH by utilizing electrodes such as hydrogen, quinhydrone, or glass electrodes. According to the Nernst equation, hydrogen electrode potential depends on solution pH. Combining a hydrogen electrode with a reference electrode like the calomel electrode and measuring cell EMF potentiometrically allows accurate pH assessment.
Q7: What equilibrium constants can be derived from cell EMF measurements?
Cell EMF measurements provide multiple equilibrium constants including 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. All these values derive from the relationship between standard cell potential and Gibbs energy.