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An outline of potentiometry, which is the use of electrodes to measure voltages that provide chemical information. The document covers reference electrodes, indicator electrodes, junction potential, ion-selective electrodes, and pH measurement with a glass electrode. how potentiometry works, the advantages and disadvantages of different types of electrodes, and how to calculate voltage changes. The document also includes examples of titration curves and calibration curves. useful for students studying analytical chemistry or electrochemistry.
Typology: Lecture notes
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Outline:
Potentiometry is the use of electrodes to measure voltages that provide chemical information. The analyte must be an electroactive species (i.e., must be able to donate or accept electrons at one of the electrodes. The unknown solution is converted into a half-cell by adding an electrode (e.g., Pt wire) that can accept or donate electrons from or to the analyte. This is the indicator electrode. This half-cell is connected to another half-cell with a constant composition via a salt bridge. This is the reference electrode. The cell voltage is the difference between the variable potential of the analyte half-cell and the constant potential of the reference electrode.
The cell voltage is the difference E + − E − : The cell voltage is dependent only upon the quotient [Fe2+]/[Fe3+], since the composition of the left half-cell is fixed (saturated solution of KCl); the left half-cell is the reference electrode. The cell and salt bridge enclosed by the dashed line (in the figure on the previous page) can be thought of as a single unit dipped into the analyte solution where:
The silver - silver chloride electrode can be constructed as a thin tube that is dipped into solution (left) or as a double - junction electrode that minimizes contact between analyte solution and KCl from the electrode (right). Double-junction reference electrode. The inner electrode is the same as the one as pictured to the left. The solution in the outer compartment is compatible with analyte solution. For example, if you do not want Cl−^ to contact the analyte, the outer electrode can be filled with KNO 3 solution. The inner and outer solutions slowly mix, so the outer compartment must be refilled periodically with fresh KNO 3 solution.
The calomel electrode is based on the reaction The standard potential for this reaction is +0.268 V. If the cell is saturated with KCl at 25°C, the potential is +0.241 V. A calomel electrode saturated with KCl is called a saturated calomel electrode , abbreviated S.C.E. (and pictured to the right). The advantage in using saturated KCl is that [Cl−] does not change if some liquid evaporates.
If an electrode has a potential of −0.461 V with respect to a calomel electrode, what is the potential with respect to a silver-silver chloride electrode? What would be the potential with respect to the standard hydrogen electrode? Point A: −0.461 V from S.C.E. −0.417 V from Ag | AgCl −0.220 V from S.H.E. Point B: −0.011 V from S.C.E. +0.033 V from Ag | AgCl +0.230 V from S.H.E.
A silver electrode can be used with a reference electrode to measure Ag+^ concentration. The reaction at the Ag indicator electrode is The calomel reference half-cell reaction is The reference potential (E−, not E−o) is fixed at 0.241 V because the reference cell is saturated with KCl. The Nernst equation for the entire cell is The voltage of this cell provides a measure of [Ag+]. Ideally, the voltage changes by 59.16 mV (at 25°C) for each factor-of-10 change in [Ag+].
e.g., A 100.0-mL solution containing 0.1000 M NaCl was titrated with 0.1000 M AgNO 3 , and the voltage of the cell monitored. The equivalence volume is V e = 100.0 mL. Calculate the voltage after the addition of (a) 65.0 and (b) 135.0 mL of AgNO 3.
Junction potential : A small voltage difference (a few mV) that develops at the interface of two dissimilar electrolyte solutions are in contact with one another. The junction potential is found at each end of a salt bridge connecting two half-cells, and places a fundamental limitation on the accuracy of direct potentiometric measurements, because we usually do not know the contribution of the junction to the measured voltage. Why does the junction potential arise? Consider NaCl in distilled H 2 O:
Salt bridges are usually constructed from compounds with ions possessing similar mobilities (e.g., K+^ and Cl-) and therefore small resulting junction potentials Nonetheless, the junction potential of 0.1 M HCl | 3.5 M KCl is 3.1 mV. A pH electrode has a response of 59 mV per pH unit. A pH electrode dipped into 0.1 M HCl will have a junction potential of ca. 3 mV, or an error of 0.05 pH units (12% error in [H+]).
When a few C+^ ions diffuse from the membrane into the outer aqueous phase, there is excess positive charge in this phase. This imbalance creates an electric potential difference that opposes diffusion of more C+^ into the aqueous phase. L: ionphore - ligand chosen to have a high affinity for analyte cation, C+, and low affinity for other ions. R-: hydrophobic anion that is soluble in the membrane and poorly soluble in water (cannot leave membrane for aqueous solution). LC+: Complex in eqb. with free C+. Membrane made of poly(vinyl chloride) impregnated with the plasticizer dioctyl sebacate, plus excess free L, R-, and LC+ Ovals are a visual aid to help in counting charge. A-: insoluble in organic solvents, does not enter membrane
When C+^ ions diffuse from the membrane (region of activity A m) to the outer solution (region of activity A o), there is a change in Gibbs energy: Δ G associated with change in solvation energy on going from organic liquid to the aqueous phase Δ G associated with change when a species diffuses between regions of different activities (concentrations) The driving force for diffusion of C+^ from the membrane to the aqueous solution is the favourable solvation of the ion by water. As C+^ diffuses, there is a buildup of positive charge in the water immediately adjacent to the membrane. The charge separation creates an electric potential difference ( E outer) across the membrane. The difference in Gibbs energy for C+^ in the two phases is Δ G = − nFE outer, where n is the charge of the ion. At equilibrium, for diffusion of C+, the net change in Δ G = 0.
The potential difference between the outer and the inner solutions is If the constant terms are combined, we find that the potential difference across the membrane depends only on the activity of analyte in the outer solution: Converting ln into log and inserting values of R , T and F gives which applies to any ion-selective electrode, including a glass pH electrode. If the analyte is an anion, the sign of n is negative.
The glass electrode used to measure pH is the most common ISE. A typical pH combination electrode has both glass and reference electrodes in one body, and has the line diagram: The reference electrode at the left of the line diagram is the coiled Ag | AgCl electrode in the combination electrode (see next slide). The reference electrode at the right side of the line diagram is the straight Ag | AgCl electrode at the centre of the electrode (see next slide). The salt bridge in the line diagram is the porous plug at the bottom right side of the combination electrode (see next slide). The two reference electrodes measure the electric potential difference across the glass membrane (see next slide).
