INTRODUCTION TO ELECTROANALYTICAL CHEMISTRY


INTRODUCTION TO ELECTROANALYTICAL CHEMISTRY
Wiliam R.Heineman
INTRODUCTION
Electrochemistry involves the measurement of electrical signals associated  with chemical system that are incorporated into an electrochemical cell. The cell consists of two or more electrodes that function as transducers between the chemical system and an electrical system in which the electrical parameters of voltage and current can be measured or controlled. Electrochemistry plays key roles in many areas of chemistry : analysis, thermodynamics, synthetic, kinetics, energy conversion, biological electron transport, and nerve impulse conduction to name a few.
Electro analytical chemistry makes use of electrochemistry for the purposes of analysis. In this application the magnitude of a voltage or current signal originating from an electrochemical cell is related o the activity or concentration of  a particular species in the cell. A host of electroanalytical techniques has been developed for this purpose. These techniques have certain features  that in some situations make electroanalysis advantages  comparison with other analytical methods. Excellent detection limits coupled with a  wide dynamic range are exhibited by many techniques. An operating of 10-8-10-3 M is typical of some techniques. Methods are available for the rapid determination of the relatively high concentrations found for species such as blood electrolytes (Na+, Cl-, HCO3-) as well as trace levels of species such as heavy metals in crop samples and drug metabolite in blood and urine samples. Measurements can generally be made on very small volumes of sample, for example, in the micro liter range. The combination of low detection limits with micro liter volume samples enable amounts of analyte at the picomole level to be measured routinely in some instances. Measurements on nanoliter volume have been demonstrated in research laboratories! Electroanalysis is an inexpensive technique in comparison with many instrumental methods. Electrochemistry benefits from the feature that the signals originating from an electrochemical cell are electrical in nature. By comparison, many analytical techniques require conversion of the signal for measurement. Electroanalysis lends itself to measurement made in vivo. For example, miniature electrochemical censors are used to measure pH and pOH in the bloodstream of patients with in-dwelling catheters.
The utility of many electrochemical techniques extends beyond qualitative and quantitative analysis into the fundamental characterization of chemical process. For example, electroanalytical methods are commonly used to study the redox properties of inorganic, organic, biological compounds.
The purposes of this chapter is to provide a fundamental beckground for understanding the modern electroanalytical techniques that are discussed in Chapter 28-32.
27.2 ELECTROCHEMICAL CELL : CONCEPTS, TERMS, AND, SYMBOLS.
An electrochemical cell typically consists of two electrodes immersed in a solution of ionsin solvent, that is, an electrolyte solution. A representative cell that serves as a suitable illustration is shown in Fig.27.1. The cell consist of two half-cells. An electrode of Zn immersedin an aqueous solution of  ZnSO4 and an electrode of Cu in an aqueous solution of CuSO4. The two sulfate solution are  connected by a salt bridge, which consist of a tube containing an aqueous  solution of concentrated KCl. The  ends of the tube are plugged with porousfrits that allow the passage of ions, but prevent the contents of the tube from rapidly draining out. The purposes of the salt bridge is two fold : to allow electrical contact based on ionic  conductance to be maintained between the electrolytes in the two containers and two prevent the two solutions from mixing. The wires attached to the electrodes in Fig.27.1a can be connected to a digital voltmeter (position A) or to each other (B), whereas the electrodes in Fig.27.1b are connected to an external powersupply. The behavior of the of the electrochemical cell can be examined by performing the three experiments outlined below.

Cell Potential. In the first experiment the potential of the electrochemical cell (Ecell) is measured by connecting the wire leads to the digital voltmeter (position A). Since this voltmeter requires neglible current, the potential measurement can be made without perturbing the electrochemical cell (i.c. witout allowing any electrochemical reactions to occur).
The cell potential is a measure of the different in electron energy between the two electrodes. The electron energy of each electrode is related to the driving force for specific redox reactions that occur at the electrode-solution interface. For the example in Fig.27.1a, one electrode reaction (as shown at the bottom of the Zn electrode) is the oxidant of Zn to form Zn2+. The electrons released. By this process contribute to an excess negative charge on the Zn electrodes as indicated by the (-) sign at the top of the electrode. The other electrode reaction (as shown at the bottom of the Cu electrode) is the reduction of Cu2+ to form Cu. The electrons consumed by this process contribute to a positive charge on the Cu electrode, as indicated at the top of electrode. Thus, the potential of an electrode depends on an excess or deficit of charge on the electrode. The energy required to add or substract an e electron from the electrodes can be expressed in terms of this potential. With an excess of negative. An excess in positive charge corresponds to low electron energy an a positive potential. The digital voltmeter responds to the magnitude of the electron energy difference between the two electrodes, which is Ecell.
The difference of potential between two points is measured by the work necessary to carry a unit positive charge from one point to the  other. The unit in which  potential is generally expressed is the volt (V) or milivolt (mV). Typical values for electrochemical cells range from a few tenths of a milivolt to a few volts. A volt is that potential difference against which one joule of work is done in the transfer of one coulomb of charge. Ohm’s Law states that the volt is the electrical potential which when applied to a resistor of one ohm causes a current of one ampere. These and other terms, units, constants, symbols, and conversions in electrochemistry are listed in Table 27.1
Electrolysis in a Galvanic Cell : Current , Charge, and Faraday’s Law. In the second experiment the two wire leads are connected (position B). This completion of the electrical circuit for the electrochemical cell is accompanied by the onset of several phenomena : electrolysis at the two electrodes, electrons flowing through the wire leads, and ions flowing through the salt bridge. The cell is now functioning as a galvanic cell, which means that it  is producing energy spontaneously. The electrolysis reactions at the two electrodes consists of reduction-oxidation (redox) reactions that can be described as shown at the bottom of Fig.27.1a. The electrode at which oxidation occurs  is termed the anode, which in this cell is the Zn electrode. The electrode at which reduction occurs is termed the cathode, which is the Cu electrode. The electrochemical cell can be thought of in terms of the two half-cells, namely the anode compartment and the cathode compartment. The overall chemical reaction is simply the sum of the two half-cell reactions. In this cell the overall chemical reaction is the reduction of copper ion to copper metal by zinc metal. Since Zn reduces Cu2+  to Cu, Zn is termed the reduction in the redox reaction. Conversely , Cu2+ is termed the oxidant, since it oxidizes Zn to Zn2+.
Table 27.1
Examination of the wire leads shows electrons to be flowing from the Zn anode to the Cu cathode as a result of the two half-cell reaction and in response to the difference in electron energies of the two electrodes. The flow of electrons  is termed current, which is the rate of transfer of electricity. Current is an expression of the rate at which the two electrode processes are occurring. Since the conversion of each Zn atom to a Zn2+ ion releases 2e, the current is a direct measure of the rate at which the oxidation process is occurring. The same argument applies to the rate of reduction of Cu2+ at the Cu electrode.
The practical unit for current is the ampere (A), which is the transfer of one coulomb per second. This corresponds to the passage of 1.05 x 105 moles of electrons per second. Since the current involved in most electroanalytical techniques is very small, miliamperes (mA), microamperes (ยตA), and nanoamperes (nA) are commonly used units.
Charge is a quantity of electricity. The charge passed during a period of time can be to obtained by integration of the current during that period :
RUMUS
The unit for charge is the coulomb ©, which corresponds to 1.05 x 10-5 moles of electrons.
In electroanalytical chemistry, it is important to relate the amount of electrical charge passed through an electrochemical cell to the quantity of material that has undergone electrolysis. Faraday’s law gives the relationship.
Q= Nfn
Where F is the Faraday constant (96,486 C/mol), N is the number of moles electrolyzed,  and n is the number of electrons involved in the electrode or redox reaction (n can be taken as a dimensionless parameter that defines the electron stoichiometry of the reaction). Differentiation of  Eqs. (27.1) and (27.1) gives the expression:
RUMUS
Which clearly delineats  the direct relationship between the rate at which electricity is moved across the electrode-solution interface (dQ/dt) and how fast the chemistry is accomplished at the interface (dN/dt).
In an electrochemical cell, the current at the anode must equal the current at the cathode or huge potential differences would result from the loss of electroneutrality . Thus, the electrons removed from the anode compartment replaced in the cathode compartment. In order of maintain electroneutrality in each half-cell. The flow of charge by the movement of electrons in the electrode and wire leads of the cell must be accompanied by the  flow of ions in the solution components of the cell, that is, the electron current must equal the ionic current. The salt bridge provides the path for ion flow to maintain electrical neutrality in each half-cell compartment. The movement of ions in this particular cell is shown diagrammatically at the two ends of the salt – bridge tube (Fig.21.1a). In the Zn2+ compartment, the development of a positive electrostatic charge from the electrolytic production of Zn2+ is prevented by the movement of Zn2+ into the salt bridge and of Cl- out of the bridge. At the other end of the bridge, K+ moves out an SO2-4 moves into the bridge to prevent the buildup of a negative electrostatic charge in the Cu2+ half-cell compartment due to the loss of Cu2+ by reduction to Cu. An overall movement of positive ionic charge from the anode half-cell to the cathode half-cell and of negative charge in the reverse direction occurs to maintain electrical neutrality in both half-cell compartments and the salt bridge. The proportion of ionic charge carried by each type of ion (K+, Zn2 ,Cl , and SO24) depends on the charge, mobility, and concentration of each ion. From a practical, point of view, removing the salt bridge halts electrolysis in the electrochemical cell just as effectively as disconnecting the two wire leads. In one case the path for ions is disrupted, in the other the path for electrons is disrupted. Both are essential for completion of the circuit for the cell.
As the electrolysis in the electrochemical cell proceeds, the overall chemical reaction is approaching a position of equilibrium. The attainment of equilibrium is signated by the measurement of zero voltage and current in what is now a “dead” battery.

Electrolysis in an electrolyte cell.
The last experiment involves the application of a potential to the electrochemical cell from an external power supply, as shown in Fig.27ib. If an external voltage of the appropriate polarity is impressed on the electrochemical cell the overall chemical reaction can be forced in the opposite direction of that observed for the galvanic cell, as sown at the bottom of Fig.27ib. In so doing, we find that the Cu electrode is now the anode and the Zn electrode is cathode. The electrochemical cell is now functioning as an electrolytic cell in that it is consuming electrical power, which is driving the chemical reaction. The external power supply is forcing the electron energy of the Zn electrode to a sufficiently high level to cause reduction of Zn2+ at the electrode, correspondingly, the electron energy of the Cu electrode is diminished to the point that electrons can be removed to form Cu2+.
Schematic Representation of Cell.
The portrayal of an electrochemical cell can be simplified by a schematic representation based on symbols. The commonly used symbols are as follows :
(,) two species in the same phase or of the same phase type but where no potential is developed : (/) phase boundary at which a potential may developed; and (//) salt bridge, which has two phase boundary at which potentials may developed. The chemical components of the cell are identified by the usual symbols with appropriate activities or concentrations indicated in parentheses. By convention, the anode half-cell is written left of the salt-bridge symbol. Also by convention the anode is considered the negative (-) electrode in a galvanic cell, but it is the positive (+) electrode in an electrolyte cell.
CELL POTENTIALS AND THE NERNST EQUATION
Free Energy and cell potential.
The free-energy change (G) for a chemical reaction is a measure of the thermodynamic driving force for the reaction to occur. The free-energy change for the chemical reaction of an electrochemical cell is related to the equilibrium cell potential by the equation :
RUMUS
By the thermodynamic convention, the cell reaction is spontaneous when Ecell is positive and Gcell is negative. If Ecell is negative an Gcell is positive, the reaction is not spontaneous as written and will proceed in the reserved direction. The cell is at equilibrium when Ecell and Gcell are zero.
The free-energy change of a reaction and the potential of an electrochemical cell are said to be the standard free-energy change Gcell and the standard cell potential Ecell respectively, when all of the reactants and products are in the standard state, which is defined as unit activity . The standard free-energy change is related to the standard cell potential by the equation.
RUMUS
Since analytical measurements are rarely, if ever, made under the conditions of standard state, an expression that relates cell potential to activity is of utmost  importance.
Nernst Equation
Consider a general chemical reaction that is comprised of two-half-cell redox reactions as shown below.
RUMUS
The free energy change for the overall reaction, Eq (27.8), can be related to the standard free energy change by:
RUMUS
Where a is the activity of the constituent of the cell reaction, R is the gas constant, and T is temperature. Substitution for G From Eqs (27.4) and (27.5) and division by the term-nF gives the Nersnt equation.
RUMUS
Which is one of the most important equation in electroanalytical chemistry. This general expression can be converted into a commonly used from of the Nersnt equation at room temperature (25oC) by conversion to a base  ten logarithm and substation  for the constant terms (R=8,314 J mol-1K-1, T=298 k, F= 96,485 C mol-1) to give a constant in front of the log of 0,0591/n with units of JC-1 or V.
RUMUS
The Nersnt equation enables to standard potential of cell to be used to calculate the cell potential for reactant and product activities  that are different from those of standard state. The activities of soluble ionic or molecular species should be expressed in terms of moles per liter. The activity for solvent and a pure solid are taken as unity; the activity of a gas is taken a the partial pressure in atmosphere.
Figure 27.2 (left ordinate) shows how the potential of an electrochemical cell at 25oC is influenced by the ratio of activities in the logarithmic term. When the activities are such that the logarithmic term is zero, the cell potential is the standard potential Ecell. As the logarithmic  term increase, the cell potential increases. The slope of the plot, -0,0591/n, is determined by the electron stoichiometry of the redox reaction.
FORMAT CELL POTENTIAL
Since analytical chemists are generally more interested in measuring concentration than activity, it is useful to express activity  in terms of concentration C and activity coefficient where :
RUMUS
Which the substituted into Eq (27.10) gives:
RUMUS
This general expression can be converted into a commonly used form of the Nerst equation by combining the activity coefficient terms and Ecell to give the formal cell potential Ecell :
RUMUS
The appropriate substitutions for C are concentration (ml L-1) for soluble ionic and molecular species, parcial pressure (atm) for gases, and unity for solvent and pure solids.
EQUILIBRIUM CONSTANT
As mentioned above, electrolysis in a cell with the two electrodes connected will proceed until the position of equilibrium for the overall chemical reaction has been reached. At equilibrium, the free energy of the cell has been expended so that Gcell =0 (and Ecell=0), and the logarithmic term of activities becomes the equilibrium constant K for the reaction. Equation (27.9) can then the written ;
RUMUS
Substitution of Eq (27.5) gives
RUMUS
The equilibrium constant for a chemical reaction can be calculated from Ecell of the appropriate electrochemical cell.

CALCULATION OF CELL POTENTIAL FROM HALF-CELL POTENTIAL
An electrochemical cell based on the redox reactioning Eq (27.8) consists of two half-cells with the half cell reactions described by Eq (27.8) and (27.7). The free energy change for the overall reaction in Eqs (27.8) in equal to the sum of the free energy changes for the two half-cell reactions :
RUMUS
Since tables of standard redox potential tabulate Eo in terms of a reduction reaction, it is convenient to express the free-energy for the oxidation reaction G in terms of a reduction reaction (-G), which has the opposite sign. Using this convention, one would express the half-cell reactions for Eq (27.8) as follows :
RUMUS
So that,
RUMUS
Substitutions of Eq (27.4) and division by - gives ;
RUMUS
Because the electrode to be considered as the anode is written as the left half-cell in the schematic diagram of an electrochemical cell. Eq (27.3) can be expressed in more general term as :
RUMUS
Equation (27.24) is the convention established by IUPAC for the calculation of cell potentials. In this convention the half-cells are not specifically identified as being the anode or cathode. This is advantageous in that the true identity of the anode is generally not known until the result of the calculation is known or the cell is actually constructed and its potential measured. Thus, couching the calculation of Ecell in the framework of Eright a Eleft for a schematically represented call is generally less confusing than dealing with Ecathode  and Eanode, which may have to be reassigned after the result of the calculation are known.
The potential of a cell is easily calculated by the first calculating the potential of each half-cell, both written in terms of reduction reactions. For example, the appropriate expressions for Eright and Eleft for the general case considered in Eqs (27.18-27.20) would be ;
RUMUS
Where Eo is the standard electrode potential for each half-cell. Once Eright and Eleft are calculated Eq (27.24) is applied to calculate the cell potential.

TABLES OF STANDARD ELECTRODE POTENTIALS
As described in the previous section, the potential of an electrochemical cell can be calculated by taking the difference in the potentials of two half-cells. For the purpose of calculating half-cell potentials, it is useful to have a table of standard electrode potentials, Eo, for various redox systems. One difficulty of a fundamental nature in obtaining data for such a table is the inability to measure the potential difference between two half-cells! The question “what is the potential of the half –cell Cu2+ (0,1M)/Cu?” is somewhat analogous to the question :How far away is Cincinnati?” In answering the latter question one must ask “How far away is Cincinnati from where : Lubbock, Durham, the planet Jupiter?” In other words one must establish a point of reference from which the distance is to be measured. The same is true of the first question, which should also be rephrased : “What is the potential of the half cell Cu2+ (0,1M)/Cu with respect to some other half-cell?” In the case of electrochemical cells, it is useful to select a half-cell against which other half-cells can be measured or referenced. Such a half-cell is called a reference electrode.
A short tabulation of Eo values for representative half-cell reactions is shown in Table 27.2. The half –cell reaction for the reduction of H+ is assigned a value of 0V. As such, this couple is the “reference electrode” for the table and is called the standard hydrogen electrode (SHE) . The measurement of half-cell potentials for other redox couples consist of preparing the half cell for each redox couple under the conditions of standard state. The potential of the cell is then measured with respect to the SHE. By convention, the potential is assigned a positive sign if the half-cell functions as a cathode versus the SHE. A negative sign is given if the half-cell functions as an anode. Also by convention, all of the half-cell reactions in the table are written as reductions. The sign tell how the redox  couple reacts with the respect to the SHE. A positive sign indicates that the reaction for the redox couple is the reduction when connected to the SHE, where as a negative sign denotes that an oxidation would take place.
Some formal electrodes potentials, Eo are also given in the table. Note that the electrolyte concentration must be stipulated for each value of Eo.

0 komentar:

Posting Komentar

Diberdayakan oleh Blogger.