Electrochemical Analysis
Introduction:
In 1791 Luigi Galvanic discovered electrical activity in the nerves of the frogs that he was dissecting. He thought that electricity was of animal origin and could be found only in living tissues. A few years later, in 1800 Alessandro Volta discovered that electricity could be produced through inorganic means. In fact, by using small sheets of copper and zinc and cloth spacers soaked in an acid solution, he built a battery - the first apparatus capable of producing electricity. Naysayers were quick to predict that electricity would never serve a useful purpose. Obviously they were very wrong. Electricity has a central role in our lives and to this day Electrochemistry is a standard course of study.
While listening to lessons on Electrochemistry, many students may wonder why it was ever invented, if it was really ever necessary to invent it and if the world would be better off without it. With the small experiments that follow, we hope to make peace between these students and the study of Electrochemistry. These fun and simple experiments can teach the fundamental concepts of Electrochemistry without asking much of the student. As you will see, many of these demonstrations are easily adapted to various configurations and each can be done independently or as part of the full curriculum.
POROUS VASE - An actual porous vase made for the purpose may be difficult to acquire. It is used to prevent the quick mixing of various solutions, while permitting the exchange of ions. For our purposes you can adapt a terracotta pot of the type used in gardening simply by plugging the hole in the bottom with molten wax and allowing it to cool. Another even more economical answer lies in constructing a barrier of paper. As shown in figure 4, roll the paper to form a cylinder and glue it in place on the bottom of the main container using a silicone adhesive such that liquids cannot pass between the two areas defined by the paper. A barrier of just one sheet would be too permeable, therefore use at least three layers of paper when building this device.
Controlling and Measuring Current and Potential:
Electrochemical measurements are made in an electrochemical cell, consisting of two or more electrodes and associated electronics for controlling and measuring the current and potential. In this section the basic components of electrochemical instrumentation are introduced. Specific experimental designs are considered in greater detail in the sections that follow.
The simplest electrochemical cell uses two electrodes. The potential of one of the electrodes is sensitive to the analyte’s concentration and is called the working, or indicator electrode. The second electrode, which is called the counter electrode, serves to complete the electric circuit and provides a reference potential against which the working electrode’s potential is measured. Ideally the counter electrode’s potential remains constant so that any change in the overall cell potential is attributed to the working electrode. In a dynamic method, where the passage of current changes the concentration of species in the electrochemical cell, the potential of the counter electrode may change over time. This problem is eliminated by replacing the counter electrode with two electrodes: a reference electrode, through which no current flows and whose potential remains constant; and an auxiliary electrode that completes the electric circuit and through which current is allowed to flow.
Although many different electrochemical methods of analysis are possible there are only three basic experimental designs: (1) measuring the potential under static conditions of no current flow; (2) measuring the potential while controlling the current; and (3) measuring the current while controlling the potential. Each of these experimental designs, however, is based on Ohm’s law that a current, i, passing through an electric circuit of resistance, R, generates a potential, E; thus
E = iR
Each of these experimental designs also uses a different type of instrument. To aid in understanding how they control and measure current and potential, these instruments are described as if they were operated manually. To do so the analyst observes a change in current or potential and manually adjusts the instrument’s settings to maintain the desired experimental conditions. It is important to understand that modern electrochemical instruments provide an automated, electronic means of controlling and measuring current and potential. They do so by using very different electronic circuitry than that shown here.
Auxiliary electrode: The third electrode in a three-electrode cell that completes the circuit.
Ohm’s law: The statement that the current moving through a circuit is proportional to the applied potential and inversely proportional to the circuit’s resistance (E = iR).
Potentiometer: A device for measuring the potential of an electrochemical cell without drawing a current or altering the cell’s composition.
Potentiometric Methods of Analysis:
In potentiometry the potential of an electrochemical cell is measured under static conditions. Because no current, or only a negligible current, flows while measuring a solution’s potential, its composition remains unchanged. For this reason, potentiometry is a useful quantitative method. The first quantitative potentiometric applications appeared soon after the formulation, in 1889, of the Nernst equation relating an electrochemical cell’s potential to the concentration of electroactive species in the cell.
When first developed, potentiometry was restricted to redox equilibria at metallic electrodes, limiting its application to a few ions. In 1906, Cremer discovered that a potential difference exists between the two sides of a thin glass membrane when opposite sides of the membrane are in contact with solutions containing different concentrations of H3O+. This discovery led to the development of the glass pH electrode in 1909. Other types of membranes also yield useful potentials. Kolthoff and Sanders, for example, showed in 1937 that pellets made from AgCl could be used to determine the concentration of Ag+. Electrodes based on membrane potentials are called ion-selective electrodes, and their continued development has extended potentiometry to a diverse array of analytes.
Potentiometric Measurements:
Potentiometric measurements are made using a potentiometer to determine the difference in potential between a working or, indicator, electrode and a counter electrode. Since no significant current flows in potentiometry, the role of the counter electrode is reduced to that of supplying a reference potential; thus, the counter electrode is usually called the reference electrode. In this section we introduce the conventions used in describing potentiometric electrochemical cells and the relationship between the measured potential and concentration.
Potentiometric Electrochemical Cells:
Note that the electrochemical cell is divided into two half-cells, each containing an electrode immersed in a solution containing ions whose concentrations determine the electrode’s potential. This separation of electrodes is necessary to prevent the redox reaction from occurring spontaneously on the surface of one of the electrodes, short-circuiting the electrochemical cell and making the measurement of cell potential impossible. A salt bridge containing an inert electrolyte, such as KCl, connects the two half-cells. The ends of the salt bridge are fixed with porous frits, allowing ions to move freely between the half-cells and the salt bridge, while preventing the contents of the salt bridge from draining into the half-cells. This movement of ions in the salt bridge completes the electric circuit.
By convention, the electrode on the left is considered to be the anode, where oxidation occurs
Zn(s) → Zn2+(aq) + 2e– and the electrode on the right is the cathode, where reduction occurs
Ag+(aq) + e– → Ag(s)
The electrochemical cell’s potential, therefore, is for the reaction
Zn(s) + 2Ag+(aq) → 2Ag(s) + Zn2+(aq)
Also, by convention, potentiometric electrochemical cells are defined such that the indicator electrode is the cathode (right half-cell) and the reference electrode is the anode (left half-cell).
Shorthand Notation for Electrochemical Cells:
Although Figure 3 provides a useful picture of an electrochemical cell, it does not provide a convenient representation. A more useful representation is a shorthand, or schematic, notation that uses symbols to indicate the different phases present in the electrochemical cell, as well as the composition of each phase. A vertical slash (|) indicates a phase boundary where a potential develops, and a comma (,) separates species in the same phase, or two phases where no potential develops. Shorthand cell notations begin with the anode and continue to the cathode. The electrochemical cell, for example, is described in shorthand notation as
Zn(s) | ZnCl2 (aq, 0.0167 M) || AgNO3 (aq, 0.100 M) | Ag(s)
The double vertical slash (||) indicates the salt bridge, the contents of which are normally not indicated. Note that the double vertical slash implies that there is a potential difference between the salt bridge and each half-cell.MPLE 11.1
Example. What are the anodic, cathodic, and overall reactions responsible for the potential in the electrochemical cell shown here? Write the shorthand notation for the electrochemical cell.

The oxidation of Ag to Ag+occurs at the anode (the left-hand cell). Since the solution contains a source of Cl–, the anodic reaction is:
Ag(s) + Cl–(aq) → AgCl(s) + e–
The cathodic reaction (the right-hand cell) is the reduction of Fe3+ to Fe2+Fe3+(aq) + e–→ Fe2+(aq)
The overall cell reaction, therefore, is:
Ag(s) + Fe3(aq) + Cl–(aq) → AgCl(s) + Fe2+(aq)
The electrochemical cell’s shorthand notation is
Ag(s) | HCl (aq, 0.100 M), AgCl (sat’d) ||
FeCl2 (aq, 0.0100 M), FeCl3 (aq, 0.0500 M) | Pt
Note that the Pt cathode is an inert electrode that carries electrons to the reduction half-reaction. The electrode itself does not undergo oxidation or reduction.
Potential and Concentration in the Nernst Equation:
The potential of a potentiometric electrochemical cell is given as:
Ecell = Ec – Ea where Ec and Ea are reduction potentials for the reactions occurring at the cathode and anode. These reduction potentials are a function of the concentrations of those species responsible for the electrode potentials, as given by the Nernst equation