ELECTROCHEMICAL METHODS OF ANALYSIS
POTENTIOMETRY
Electrochemical cell is vessel with investigated solution in which dipped electrode.
Potentiometers
Measuring the potential of an electrochemical cell under conditions of zero current is accomplished using a potentiometer.
Schematic diagram of a manual potentiostat: C = counter electrode; W = working electrode; SW = slide-wire resistor; T = tap key; i = galvanometer.
A schematic diagram of a manual potentiometer is shown in Figure. The current in the upper half of the circuit is
where EPS is the power supply’s potential, and Rab is the resistance between points a
and b of the slide-wire resistor. In a similar manner, the current in the lower half of the circuit is
where Ecell is the potential difference between the working electrode and the counter
electrode, and Rcb is the resistance between the points c and b of the slide-wire resistor.
When
no current flows through the galvanometer and the cell potential is given by
To make a measurement the tap key is pressed momentarily, and the current is noted at the galvanometer. If a non zero current is registered, then the slide wire is adjusted and the current remeasured. This process is continued until the galvanometer registers a current of zero. Using the tap key minimizes the total amount of current allowed to flow through the cell. Provided that the total current is negligible, the change in the analyte’s concentration is insignificant. For example, a current of 10–9 A drawn for 1 s consumes only about 10–14 mol of analyte. Modern potentiometers use operational amplifiers to create a highimpedance voltmeter capable of measuring potentials while drawing currents of less than 10–9 A.
Galvanostats
A galvanostat is used for dynamic methods, such as constant-current coulometry, in which it is necessary to control the current flowing through an electrochemical cell.
A schematic diagram of a manual constant-current galvanostat is shown in Figure.
Schematic diagram of a galvanostat: R = resistor; i = galvanometer; A = auxiliary electrode; W = working electrode; R = reference electrode; V = voltmeter or potentiometer (optional).
If the resistance, R, of the galvanostat is significantly larger than the resistance of the electrochemical cell, and the applied voltage from the power supply is much greater than the cell potential, then the current between the auxiliary and working electrodes is equal to
The potential of the working electrode, which changes as the composition of the electrochemical cell changes, is monitored by including a reference electrode and a high-impedance potentiometer.
Potentiostats
A potentiostat is used for dynamic methods when it is necessary to control the potential of the working electrode.
Figure shows a schematic diagram for a manual potentiostat that can be used to maintain a constant cell potential.
Schematic diagram of a manual potentiostat: SW = slide-wire resistor; A = auxiliary electrode; R = reference electrode; W = working electrode;V = voltmeter or potentiometer; i = galvanometer.
The potential of the working electrode is monitored by a reference electrode connected to the working electrode through a high-impedance potentiometer. The desired potential is achieved by adjusting the slide-wire resistor connected to the auxiliary electrode. If the working electrode’s potential begins to drift from the desired value, then the slide-wire resistor is manually readjusted, returning the potential to its initial value. The current flowing between the auxiliary and working electrodes is measured with a galvanostat. Modern potentiostats include waveform generators allowing a time-dependent potential profile, such as a series of potential pulses, to be applied to the working electrode.
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Electrochemical cell |
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Methods without potential imposing |
Methods with potential imposing |
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– galvanic cell |
– electrolytic cell – conductometric cell |
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with two different metals connected by a salt bridge or a porous disk between the individual half-cells |
decomposes chemical compounds by means of electrical energy, in a process called electrolysis (except for conductometric cell) |
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
A schematic diagram of a typical potentiometric electrochemical cell is shown in Figure.
Electrochemical cell for potentiometry
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
and the electrode on the right is the cathode, where reduction occurs
The electrochemical cell’s potential, therefore, is for the reaction
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 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
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.
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
The cathodic reaction (the right-hand cell) is the reduction of Fe3+ to Fe2+
The overall cell reaction, therefore, is
The electrochemical cell’s shorthand notation is
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. The Nernst Equation.
The potential of a potentiometric electrochemical cell is given as
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
where E° is the standard-state reduction potential, R is the gas constant, T is the temperature in Kelvins, n is the number of electrons involved in the reduction reaction, F is Faraday’s constant, and Q is the reaction quotient. Under typical laboratory conditions (temperature of
where E is given in volts.
Using equation the potential of the anode and cathode in are
Note, again, that the Nernst equations for both Ec and Ea are written for reduction reactions. The cell potential, therefore, is
Substituting known values for the standard-state reduction potentials and the concentrations of Ag+ and Zn2+, gives a potential for the electrochemical cell of
In potentiometry, the concentration of analyte in the cathodic half-cell is generally unknown, and the measured cell potential is used to determine its concentration. Thus, if the potential for the cell in Figure 11.5 is measured at +1.50 V, and the concentration of Zn2+ remains at
Solving for [Ag+] gives its concentration as
Despite the apparent ease of determining an analyte’s concentration using the Nernst equation, several problems make this approach impractical. One problem is that standard-state potentials are temperature-dependent, and most values listed in reference tables are for a temperature of
Another problem is that the Nernst equation is a function of activities, not concentrations. As a result, cell potentials may show significant matrix effects. This problem is compounded when the analyte participates in additional equilibria. For example, the standard-state potential for the Fe3+/Fe2+ redox couple is +0.767 V in
A more serious problem is the presence of additional potentials in the electrochemical cell, not accounted for by equation. In writing the shorthand notation for the electrochemical cell, for example, we use a double slash ( ) for the salt bridge, indicating that a potential difference exists at the interface between each end of the salt bridge and the solution in which it is immersed. The origin of this potential, which is called a liquid junction potential, and its significance are discussed in the following section.
Liquid Junction Potentials
A liquid junction potential develops at the interface between any two ionic solutions that differ in composition and for which the mobility of the ions differs. Consider, for example, solutions of
Since the concentration of HCl on the left side of the membrane is greater than that on the right side of the membrane, there is a net diffusion of H+ and Cl– in the direction of the arrows. The mobility of H+, however, is greater than that for Cl–, as shown by the difference in the lengths of their respective arrows. As a result, the solution on the right side of the membrane develops an excess of H+ and has a positive charge (Figure b). Simultaneously, the solution on the left side of the membrane develops a negative charge due to the greater concentration of Cl–. The difference in potential across the membrane is called a liquid junction potential, Elj.
The magnitude of the liquid junction potential is determined by the ionic composition of the solutions on the two sides of the interface and may be as large as 30–40 mV. For example, a liquid junction potential of 33.09 mV has been measured at the interface between solutions of
Ecell = Eind – Eref + Elj
Since the junction potential is usually of unknown value, it is normally impossible to directly calculate the analyte’s concentration using the Nernst equation.
POTENTIOMETRIC CELL, REFERENCE AND INDICATING ELECTRODES, MEASUREMENT OF EMF
!!! Direct measurement of potential value is impossible, therefore apply system of electrodes and compare potential of one electrode with potential of the second electrode
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.
Indicator electrode is the electrode whose potential is a function of the analyte’s concentration (also known as the working electrode).
Counter electrode is the second electrode in a two-electrode cell that completes the circuit.
Reference electrode is an electrode whose potential remains constant and against which other potentials can be measured.
Classification of electrodes
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Electrodes |
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I sorts |
ІІ sorts |
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Reversible on cation |
Reversible on anion |
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The electrode potential is defined by cation activity which is like to the metal dipped into electrolyte |
The electrode potential is defined by anion activity |
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– Metal electrodes – membrane or ion-selective electrodes – A glass electrode – A hydrogen electrode – quinhydrone electrode |
– An covered electrode slightly soluble salt of those cation, from which made electrode (silver-chloride and calomel (SCE) electrode) – Gas electrodes (oxygen, chlorine-chloride, etc.) |
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Oxidation-reduction electrodes (precious metals) The potential depends from ratio of oxidized and reduced forms of redox-pair |
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The characteristic of electrodes
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Electrode |
Nernst Equations |
Possibilities |
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Quinhydrone electrode |
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рН=0-8.0 |
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The glass |
Е0 is depends from asymmetry potential and equilibrium constant H+(solution) « H+(glass) |
1.рН=0-12.0-13.0 2. possible presence Ox and Red 3. possibility measurement in colloidal solutions |
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Hydrogen |
Е = – 0,058рН. |
inconveniences of application |
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Silver |
E=E0(Ag+/Ag)+0,058 lgaAg+ |
I sorts |
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Mercury |
E=E0(Hg2+/Hg)+0,029 lgaHg2+. |
I sorts |
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Silver-chloride Ag,AgCl KClsat. |
E(Ag+/Ag) = E0(Ag+/AgCl) –
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ІІ sorts (reference electrodes) |
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Calomel Hg,Hg2Cl2 KClsat. |
E(Hg22+/Hg2Cl2) =E0(Hg22+/2Hg) +
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ІІ sorts (reference electrodes) |
INDICATING ELECTRODES: METAL, MEMBRANE
Metallic Indicator Electrodes
The potential of the indicator electrode in a potentiometric electrochemical cell is proportional to the concentration of analyte. Two classes of indicator electrodes are used in potentiometry: metallic electrodes, which are the subject of this section, and ion-selective electrodes, which are covered in the next section.
The potential of a metallic electrode is determined by the position of a redox reaction at the electrode–solution interface. Three types of metallic electrodes are commonly used in potentiometry, each of which is considered in the following discussion.
Electrodes of the First Kind
When a copper electrode is immersed in a solution containing Cu2+, the potential of the electrode due to the reaction
Cu2+(aq) + 2e↔ Cu(s)
is determined by the concentration of copper ion.
If the copper electrode is the indicator electrode in a potentiometric electrochemical cell that also includes a saturated calomel reference electrode
then the cell potential can be used to determine an unknown concentration of Cu2+ in the indicator half-cell
Metallic indicator electrodes in which a metal is in contact with a solution containing its ion are called electrodes of the first kind. In general, for a metal M, in a solution of Mn+, the cell potential is given as
where K is a constant that includes the standard-state potential for the Mn+/M redox couple, the potential of the reference electrode, and the junction potential. For a variety of reasons, including slow kinetics for electron transfer, the existence of surface oxides and interfering reactions, electrodes of the first kind are limited to Ag, Bi, Cd, Cu, Hg, Pb, Sn, Tl, and Zn. Many of these electrodes, such as Zn, cannot be used in acidic solutions where they are easily oxidized by H+.
Electrodes of the Second Kind
An electrode of the first kind involving an Mn+/M redox couple will respond to the concentration of another species if that species is in equilibrium with Mn+.
For example, the potential of a silver electrode in a solution of Ag+ is given by
If the solution is saturated with AgI, then the solubility reaction
determines the concentration of Ag+; thus
where Ksp,AgI is the solubility product for AgI.
shows that the potential of the silver electrode is a function of the concentration of I–. When this electrode is incorporated into a potentiometric electrochemical cell
the cell potential is
where K is a constant that includes the standard-state potential for the Ag+/Ag redox couple, the solubility product for AgI, the potential of the reference electrode, and the junction potential.
When the potential of an electrode of the first kind responds to the potential of another ion that is in equilibrium with Mn+, it is called an electrode of the second kind. Two common electrodes of the second kind are the calomel and silver/silver chloride reference electrodes. Electrodes of the second kind also can be based on complexation reactions. For example, an electrode for EDTA is constructed by coupling a Hg2+/Hg electrode of the first kind to EDTA by taking advantage of its formation of a stable complex with Hg2+.
Redox Electrodes
Electrodes of the first and second kind develop a potential as the result of a redox reaction in which the metallic electrode undergoes a change in its oxidation state. Metallic electrodes also can serve simply as a source of, or a sink for, electrons in other redox reactions. Such electrodes are called redox electrodes. The Pt cathode is an example of a redox electrode because its potential is determined by the concentrations of Fe2+ and Fe3+ in the indicator half-cell. Note that the potential of a redox electrode generally responds to the concentration of more than one ion, limiting their usefulness for direct potentiometry.
Membrane Electrodes
If metallic electrodes were the only useful class of indicator electrodes, potentiometry would be of limited applicability. The discovery, in 1906, that a thin glass membrane develops a potential, called a membrane potential, when opposite sides of the membrane are in contact with solutions of different pH led to the eventual development of a whole new class of indicator electrodes called ionselective electrodes (ISEs). Following the discovery of the glass pH electrode, ionselective electrodes have been developed for a wide range of ions. Membrane electrodes also have been developed that respond to the concentration of molecular analytes by using a chemical reaction to generate an ion that can be monitored with an ion-selective electrode. The development of new membrane electrodes continues to be an active area of research.
Membrane Potentials Ion-selective electrodes, such as the glass pH electrode, function by using a membrane that reacts selectively with a single ion. Figure shows a generic diagram for a potentiometric electrochemical cell equipped with an ion-selective electrode.
The shorthand notation for this cell is
where the membrane is represented by the vertical slash ( ) separating the two solutions containing analyte. Two reference electrodes are used; one positioned within the internal solution, and one in the sample solution. The cell potential, therefore, is
where Emem is the potential across the membrane. Since the liquid junction potential and reference electrode potentials are constant, any change in the cell’s potential is attributed to the membrane potential.
Interaction of the analyte with the membrane results in a membrane potential if there is a difference in the analyte’s concentration on opposite sides of the membrane.
One side of the membrane is in contact with an internal solution containing a fixed concentration of analyte, while the other side of the membrane is in contact with the sample. Current is carried through the membrane by the movement of either the analyte or an ion already present in the membrane’s matrix. The membrane potential is given by a Nernst-like equation
where [A]samp and [A]int are the concentrations of analyte in the sample and the internal solution, respectively, and z is the analyte’s charge. Ideally, Emem should be zero when the concentrations of analyte on both sides of the membrane are equal. The term Easym, which is called an asymmetry potential, accounts for the fact that the membrane potential is usually not zero under these conditions.
Assuming a temperature of
where K is a constant accounting for the potentials of the reference electrodes, any liquid junction potentials, the asymmetry potential, and the concentration of analyte in the internal solution. Equation is a general equation, and applies to all types of ion-selective electrodes.
Selectivity of Membranes
Membrane potentials result from a chemical interaction between the analyte and active sites on the membrane’s surface. Because the signal depends on a chemical process, most membranes are not selective toward a single analyte. Instead, the membrane potential is proportional to the concentration of all ions in the sample solution capable of interacting at the membrane’s active sites. Equation can be generalized to include the contribution of an interferent, I,
where zA and zI are the charges of the analyte and interferent, and KA,I is a selectivity coefficient accounting for the relative response of the interferent. The selectivity coefficient is defined as
where [A]E and [I]E are the concentrations of analyte and interferent yielding identical cell potentials. When the selectivity coefficient is 1.00, the membrane responds equally to the analyte and interferent. A membrane shows good selectivity for the analyte when KA,I is significantly less than 1.00.
Selectivity coefficients for most commercially available ion-selective electrodes are provided by the manufacturer. If the selectivity coefficient is unknown, it can be determined experimentally. The easiest method for determining KA,I is to prepare a series of solutions, each of which contains the same concentration of interferent, [I] add, but a different concentration of analyte. A plot of cell potential versus the log of the analyte’s concentration has two distinct linear regions
When the analyte’s concentration is significantly larger than KA,I[I]add, the potential is a linear function of log [A], as given by equation. If KA,I[I] add is significantly larger than the analyte’s concentration, however, the cell potential remains constant. The concentration of analyte and interferent at the intersection of these two linear regions is used to calculate KA,I.
REFERENCE ELECTRODES: SILVER-CHLORIDE AND CALOMEL (SCE)
Potentiometric electrochemical cells are constructed such that one of the half-cells provides a known reference potential, and the potential of the other half-cell indicates the analyte’s concentration. By convention, the reference electrode is taken to be the anode; thus, the shorthand notation for a potentiometric electrochemical cell is
Reference Indicator
and the cell potential is
Ecell = Eind – Eref + Elj
The ideal reference electrode must provide a stable potential so that any change in Ecell is attributed to the indicator electrode, and, therefore, to a change in the analyte’s concentration. In addition, the ideal reference electrode should be easy to make and to use. Three common reference electrodes are discussed in this section.
Calomel Electrodes
Calomel reference electrodes are based on the redox couple between Hg2Cl2 and Hg (calomel is a commoame for Hg2Cl2).
Hg2Cl2(s) +2e↔ 2Hg(l) + 2Cl–(aq)
The saturated calomel electrode (SCE), which is constructed using an aqueous solution saturated with KCl, has a potential at
Hg(l) Hg2Cl2 (sat’d), KCl (aq, sat’d)
The SCE has the advantage that the concentration of Cl–, and, therefore, the potential of the electrode, remains constant even if the KCl solution partially evaporates. On the other hand, a significant disadvantage of the SCE is that the solubility of KCl is sensitive to a change in temperature. At higher temperatures the concentration of Cl– increases, and the electrode’s potential decreases. For example, the potential of the SCE at
Silver/Silver Chloride Electrodes
Another common reference electrode is the silver/silver chloride electrode, which is based on the redox couple between AgCl and Ag.
AgCl(s) + e↔ Ag(s) + Cl–(aq)
As with the saturated calomel electrode, the potential of the Ag/AgCl electrode is determined by the concentration of Cl– used in its preparation.
E = E°AgCl/Ag – 0.05916 log [Cl–] = +0.2223 – 0.05916 log [C1–]
When prepared using a saturated solution of KCl, the Ag/AgCl electrode has a potential of +0.197 V at
A typical Ag/AgCl electrode consists of a silver wire, the end of which is coated with a thin film of AgCl. The wire is immersed in a solution that contains the desired concentration of KCl and that is saturated with AgCl. A porous plug serves as the salt bridge. The shorthand notation for the cell is
Ag(s) AgCl (sat’d), KCl (x M)
where x is the concentration of KCl.
In comparison to the SCE the Ag/AgCl electrode has the advantage of being useful at higher temperatures. On the other hand, the Ag/AgCl electrode is more prone to reacting with solutions to form insoluble silver complexes that may plug the salt bridge between the electrode and the solution.
ELECTRODES FOR DEFINITION pН: HYDROGEN, GLASS, QUINHYDRONE, ANTIMONIC.
Standard Hydrogen Electrode
The standard hydrogen electrode (SHE) is rarely used for routine analytical work, but is important because it is the reference electrode used to establish standard-state potentials for other half-reactions. The SHE consists of a Pt electrode immersed in a solution in which the hydrogen ion activity is 1.00 and in which H2 gas is bubbled at a pressure of 1 atm. A conventional salt bridge connects the SHE to the indicator half-cell. The shorthand notation for the standard hydrogen electrode is
Pt(s), H2 (g, 1 atm) H+ (aq, a = 1.00)
and the standard-state potential for the reaction
2H+(aq) + e↔ H2(g)
is, by definition, 0.00 V for all temperatures. Despite its importance as the fundamental reference electrode against which all other potentials are measured, the SHE is rarely used because it is difficult to prepare and inconvenient to use.
Standard hydrogen electrode scheme: 1. platinized platinum electrode; 2. hydrogen gas; 3. acid solution with an activity of H+=1 mol/L; 4. hydroseal for prevention of oxygen inteference reservoir via which the second half-element of the galvanic cell should be attached
Glass Ion-Selective Electrodes
The first commercial glass electrodes were manufactured using Corning
giving rise to the membrane’s selectivity for H+. The transport of charge across the membrane is carried by the Na+ ions. The potential of glass electrodes using
over a pH range of approximately 0.5–9. Above a pH of 9–10, the glass membrane may become more responsive to other cations, such as Na+ and K+.
Replacing Na2O and CaO with Li2O and BaO extends the useful pH range of glass membrane electrodes to pH levels greater than 12.
Glass membrane pH electrodes are often available in a combination form that includes both the indicator and the reference electrode. The use of a single electrode greatly simplifies the measurement of pH. An example of a typical combination electrode is shown in Figure
The response of the
Since the typical thickness of the glass membrane in an ion-selective electrode is about
The composition of a glass membrane changes over time, affecting the electrode’s performance. The average lifetime for a glass electrode is several years.
Crystalline Solid-State Ion-Selective
Electrodes Solid-state ion-selective electrodes use membranes fashioned from polycrystalline or single-crystal inorganic salts. Polycrystalline ion-selective electrodes are made by forming a thin pellet of Ag2S, or a mixture of Ag2S and either a second silver salt or another metal sulfide.
The pellet, which is 1–2 mm in thickness, is sealed into the end of a nonconducting plastic cylinder, and an internal solution containing the analyte and a reference electrode are placed in the cylinder. Charge is carried across the membrane by Ag+ ions.
The membrane potential for a Ag2S pellet develops as the result of a difference in the equilibrium position of the solubility reaction
on the two sides of the membrane. When used to monitor the concentration of Ag+ ions, the cell potential is
The membrane also responds to the concentration of S2–, with the cell potential given as
If a mixture of an insoluble silver salt and Ag2S is used to make the membrane, then the membrane potential also responds to the concentration of the anion of the added silver salt. Thus, pellets made from a mixture of Ag2S and AgCl can serve as a Cl– ion-selective electrode, with a cell potential of
Membranes fashioned from a mixture of Ag2S with CdS, CuS, or PbS are used to make ion-selective electrodes that respond to the concentration of Cd2+, Cu2+, or Pb2+. In this case the cell potential is
where [M2+] is the concentration of the appropriate metal ion.
Several examples of polycrystalline, Ag2S – based ion-selective electrodes. The selectivity of these ion-selective electrodes is determined by solubility. Thus, a Cl– ion-selective electrode constructed using a Ag2S/AgCl membrane is more selective for Br– (KCl–/Br– = 102) and I– (KCl–/l– = 106) since AgBr and AgI are less soluble than AgCl. If the concentration of Br– is sufficiently high, the AgCl at the membrane–solution interface is replaced by AgBr, and the electrode’s response to Cl– decreases substantially. Most of the ion-selective electrodes can be used over an extended range of pH levels. The equilibrium between S2– and HS– limits the analysis for S2– to a pH range of 13–14. Solutions of CN–, on the other hand, must be kept basic to avoid the release of HCN.
The membrane of a F– ion-selective electrode is fashioned from a single crystal of LaF3 that is usually doped with a small amount of EuF2 to enhance the membrane’s conductivity. Since EuF2 provides only two F– ions, compared with three for LaF3, each EuF2 produces a vacancy in the crystal lattice. Fluoride ions move through the membrane by moving into adjacent vacancies. The LaF3 membrane is sealed into the end of a nonconducting plastic tube, with a standard solution of F–, typically
The membrane potential for a F– ion-selective electrode results from a difference in the solubility of LaF3 on opposite sides of the membrane, with the potential given by
One advantage of the F– ion-selective electrode is its freedom from interference. The only significant exception is OH– (KF–/OH– = 0.1), which imposes a maximum pH limit for a successful analysis.
Below a pH of 4 the predominate form of fluoride in solution is HF, which, unlike F–, does not contribute to the membrane potential. For this reason, an analysis for total fluoride must be carried out at a pH greater than 4.
Unlike ion-selective electrodes using glass membranes, crystalline solid-state ion-selective electrodes do not need to be conditioned before use and may be stored dry. The surface of the electrode is subject to poisoning, as described earlier for a Cl– ISE in contact with an excessive concentration of Br–. When this happens, the electrode can be returned to its original condition by sanding and polishing the crystalline membrane.
Liquid-Based Ion-Selective Electrodes
Another approach to constructing an ion-selective electrode is to use a hydrophobic membrane containing a selective, liquid organic complexing agent. Three types of organic liquids have been used: cation exchangers, anion exchangers, and neutral ionophores.
When the analyte’s concentration on the two sides of the membrane is different, a membrane potential is the result. Current is carried through the membrane by the analyte. One example of a liquid-based ion-selective electrode is that for Ca2+, which uses a porous plastic membrane saturated with di-(n-decyl) phosphate.
As shown in next Figure, the membrane is placed at the end of a nonconducting cylindrical tube and is in contact with two reservoirs. The outer reservoir contains di-(n-decyl) phosphate in di-n-octylphenylphosphonate, which soaks into the porous membrane. The inner reservoir contains a standard aqueous solution of Ca2+ and a Ag/AgCl reference electrode. Calcium ion-selective electrodes are also available in which the di-(n-decyl) phosphate is immobilized in a polyvinyl chloride (PVC) membrane, eliminating the need for a reservoir containing di-(n-decyl) phosphate.
A membrane potential develops as the result of a difference in the equilibrium position of the complexation reaction
on the two sides of the membrane, where (m) indicates that the species is present in the membrane. The cell potential for the Ca2+ ion-selective electrode is
The selectivity of the electrode for Ca2+ is very good, with only Zn2+ showing greater selectivity.
An electrode using a liquid reservoir can be stored in a dilute solution of analyte and needs no additional conditioning before use. The lifetime of an electrode with a PVC membrane, however, is proportional to its exposure to aqueous solutions. For this reason these electrodes are best stored by covering the membrane with a cap containing a small amount of wetted gauze to maintain a humid environment. The electrode must then be conditioned before use by soaking in a solution of analyte for 30–60 min.
Gas-Sensing Electrodes
A number of membrane electrodes have been developed that respond to the concentration of dissolved gases. The basic design of these electrodes is shown in Figure
and consists of a thin membrane separating the sample from an inner solution containing an ion-selective electrode.
The membrane is permeable to the gaseous analyte, but is not permeable to nonvolatile components in the sample matrix. Once the gaseous analyte passes through the membrane, it reacts in the inner solution, producing a species whose concentration can be monitored by an appropriate ion-selective electrode.
For example, in the CO2 electrode, CO2 reacts in the inner solution to produce H3O+.
The change in the concentration of H3O+ is monitored with a pH ion-selective electrode, for which the cell potential is given by equation. The relationship between the concentration of H3O+ and CO2 is given by rearranging the equilibrium constant expression for reaction; thus
where K is the equilibrium constant. If the amount of HCO3– in the internal solution is sufficiently large, then its concentration is unaffected by the presence of CO2 and remains constant.
where K¢ is a constant that includes the constant for the pH ion-selective electrode, the equilibrium constant for reaction 11.10, and the concentration of HCO3–.
The composition of the inner solution changes with use, and both it and the membrane must be replaced periodically.
Gas-sensing electrodes are stored in a solution similar to the internal solution to minimize their exposure to atmospheric gases.
Potentiometric Biosensors
Potentiometric electrodes for the analysis of molecules of biochemical importance can be constructed in a fashion similar to that used for gas-sensing electrodes. The most common class of potentiometric biosensors are the so-called enzyme electrodes, in which an enzyme is trapped or immobilized at the surface of an ion-selective electrode.
Reaction of the analyte with the enzyme produces a product whose concentration is monitored by the ion-selective electrode.
Potentiometric biosensors have also been designed around other biologically active species, including antibodies, bacterial particles, tissue, and hormone receptors.
One example of an enzyme electrode is the urea electrode, which is based on the catalytic hydrolysis of urea by urease
In one version of the urea electrode, shown in Figure 11.16, an NH3 electrode is modified by adding a dialysis membrane that physically traps a pH 7.0 buffered solution of urease between the dialysis membrane and the gas-permeable membrane. When immersed in the sample, urea diffuses through the dialysis membrane, where it reacts with the enzyme urease. The NH4+ that is produced is in equilibrium with NH3
which, in turn, diffuses through the gas-permeable membrane, where it is detected by a pH electrode. The response of the electrode to the concentration of urea is given by
Another version of the urea electrode (Figure)
immobilizes the enzyme in a polymer membrane formed directly on the tip of a glass pH electrode. In this case, the electrode’s response is
Few potentiometric biosensors are commercially available. As shown in Figures, however, available ion-selective and gas-sensing electrodes may be easily converted into biosensors.
DIRECT POTENTIOMETRY: THE THEORY OF THE QUANTITATIVE ANALYSIS, METHODS OF THE QUANTITATIVE ANALYSIS, APPLICATION.
Direct potentiometry is based on direct application of Nernst equation for calculation of activity or concentration of substance which take part in electrode reaction on experimentally measured of EMF or electrode potential
Theoretical basis of the quantitative analysis:
Potentiometric electrochemical cells are on structed such that one of the half-cells provides a known reference potential, and the potential of the other half-cell indicates the analyte’s concentration.
Ecell = Eind – Eref + Elj
where: Eind – potential of indicator electrode
Eref – potential of reference electrode
Elj – liquid junction potential (a potential that develops at the interface between two ionic solutions that differ in composition, because of a difference in the mobilities of the ions)
Methods of quantitative analysis:
1. A method of calibration chart (at constant ionic strength of solution and constant activity coefficients will be identical, and liquid junction potential will aspire to zero)
The basic equation will be:
рСМ = аЕ + b
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lg CMe
If concentration and structure of investigated and standard solutions similar to each other so it is possible to use concentration:
At Е = 0, lnCst/Cx = 0, then Сst = Сх.
To carry out this condition (Е = 0) is possible to reach by dilution of standard solution by background electrolyte solution or addition of more concentrated standard solution.
3. Method of additives
Before addition of standard solution (EMF)
And after addition:
where ΔС – incremental value of concentration
The difference of electrode potentials is equal:
POTENTIOMETRIC TITRATION: AN ESSENCE, REACTIONS, TITRATION CURVE, APPLICATION.
Potentiometric titration is a technique similar to direct titration of a redox reaction. No indicator is used, instead the potential across the analyte, typically an electrolyte solution is measured. To do this, two electrodes are used, an indicator electrode and a reference electrode. The indicator electrode forms an electrochemical half cell with the interested ions in the test solution. The reference electrode forms the other half cell, holding a consistent electrical potential. The overall electric potential is calculated as Ecell = Eind – Eref + Esol. Esol is the potential drop over the test solution between the two electrodes. Ecell is recorded at intervals as the titrant is added. A graph of potential against volume added can be drawn and the end point of the reaction is half way between the jump in voltage.
Ecell depends on the concentration of the interested ions with which the indicator electrode is in contact. For example, the electrode reaction may be
Mn++ne–—–>M
As the concentration of Mn+ changes,the Ecell changes correspondingly. Thus the potentiometric titration involve measurement of Ecell with the addition of titrant. types of potentiometric titration: acid-base titration (total alkalinity and total acidity), redox titration (HI/HY and cerate), precipitation titration (halides), and complexometric titration.
The scheme of equipment for potentiometric titration
§ Potentiometric titration is based on definition of a equivalence point by results of potentiometric measurements.
http://www.youtube.com/watch?v=g5z6EaT46iA
Requirements to reactions:
– stoichiometry
– quantitatively
– rapidly
– Selectively
– There should be fitted corresponding electrode
Types of used reactions:
– Acid-base
– Precipitation
– Complexing
– Oxidation-reduction
TITRATION CURVES:
Integral curve (the acid-base titration)
Differential curve (the acid-base titration)
curves on Gran method
curves of double differentiation
Features of potentiometric titration:
§ Separate titration of acid mixture is probably if acid constants differ more than on 4 order (accuracy to 1 %)
§ Separate titration of mix of strong and weak acids is probably if acid constants differ more than on 6 order (accuracy to 0,1 %)
§ Separate titration of a mix of oxidizers (reducers) is probably if potentials differ not less than on 0,4 V
§ Separate titration of a mix on precipitation method is possible, if solubility products of precipitates differ not less than on 3 order
Advantages of potentiometric titration
1. High accuracy (accuracy of titration to 0,1 % at ΔЕ=0,01mV)
2. High sensitivity
3. Possibility of definition of several substances
4. Titration possibility in the painted and muddy mediums
5. Possibility of usage of nonaqua mediums
6. Possibility of automation of titration processes
7. rapid analysis method
Lacks:
2. Not always fast stabilisation of potential
Usage of potentiometry in the analysis of chemical compounds and drugs
– The analysis of substances (streptocide, ofloxacin, trimethoprimum)
– The analysis of ready medicinal forms (a tablet of Biseptolum, streptocide, clotrimazolum etc.)
– Definition of рН substances, drugs
VOLTAMMETRY: A METHOD ESSENCE, ELECTRODES, CONDITIONS ANALYSIS.
In voltammetry a time-dependent potential is applied to an electrochemical cell, and the current flowing through the cell is measured as a function of that potential. A plot of current as a function of applied potential is called a voltammogram and is the electrochemical equivalent of a spectrum in spectroscopy, providing quantitative and qualitative information about the species involved in the oxidation or reduction reaction.
The earliest voltammetric technique to be introduced was polarography, which was developed by Jaroslav Heyrovsky (1890–1967) in the early 1920s, for which he was awarded the Nobel Prize in chemistry in 1959. Since then, many different forms of voltammetry have been developed. Before examining these techniques and their applications in more detail, however, we must first consider the basic experimental design for making voltammetric measurements and the factors influencing the shape of the resulting voltammogram.
Voltammetric Measurements
Although early voltammetric methods relied on the use of only two electrodes, modern voltammetry makes use of a three-electrode potentiostat. A time-dependent potential excitation signal is applied to the working electrode, changing its potential relative to the fixed potential of the reference electrode. The resulting current between the working and auxiliary electrodes is measured. The auxiliary electrode is generally a platinum wire, and the SCE and Ag/AgCl electrode are common reference electrodes.
Several different materials have been used as working electrodes, including mercury, platinum, gold, silver, and carbon. The earliest voltammetric techniques, including polarography, used mercury for the working electrode. Since mercury is a liquid, the working electrode often consists of a drop suspended from the end of a capillary tube (Figure).
Mercury electrodes: (a) hanging mercury drop electrode; (b) dropping mercury electrode; (c) static mercury drop electrode.
hanging mercury drop electrode is an electrode in which a drop of Hg is suspended from a capillary tube.
dropping mercury electrode is an electrode in which successive drops of Hg form at the end of a capillary tube as a result of gravity, with each drop providing a fresh electrode surface.
static mercury drop electrode is an electrode in which successive drops of Hg form at the end of a capillary tube as the result of a mechanical plunger, with each drop providing a fresh electrode surface.
In the hanging mercury drop electrode, or HMDE, a drop of the desired size is formed by the action of a micrometer screw that pushes the mercury through a narrow capillary tube. In the dropping mercury electrode, or DME, mercury drops form at the end of the capillary tube as a result of gravity. Unlike the HMDE, the mercury drop of a DME grows continuously and has a finite lifetime of several seconds. At the end of its lifetime the mercury drop is dislodged, either manually or by gravity, and replaced by a new drop.
The static mercury drop electrode, or SMDE, uses a solenoid-driven plunger to control the flow of mercury. The SMDE can be used as either a hanging mercury drop electrode or as a dropping mercury electrode. A single activation of the solenoid momentarily lifts the plunger, allowing enough mercury to flow through the capillary to form a single drop. To obtain a dropping mercury electrode the solenoid is activated repeatedly. A mercury film electrode consists of a thin layer of mercury deposited on the surface of a solid carbon, platinum, or gold electrode.
The solid electrode is placed in a solution of Hg2+ and held at a potential at which the reduction of Hg2+ to Hg is favorable, forming a thin mercury film. Mercury has several advantages as a working electrode. Perhaps its most important advantage is its high overpotential for the reduction of H3O+ to H2, which allows for the application of potentials as negative as -1 V versus the SCE in acidic solutions, and –2 V versus the SCE in basic solutions. A species such as Zn2+, which is difficult to reduce at other electrodes without simultaneously reducing H3O+, is easily reduced at a mercury working electrode. Other advantages include the ability of metals to dissolve in the mercury, resulting in the formation of an amalgam, and the ability to easily renew the surface of the electrode by extruding a new drop. One limitation to its use as a working electrode is the ease with which Hg is oxidized. For this reason, mercury electrodes cannot be used as at potentials more positive than –0.3 V to +0.4 V versus the SCE, depending on the composition of the solution.
Advantages of dropping mercury electrode:
§ The surface is constantly updated.
§ Range of used potentials from -2.5 V to +0,2 V
§ The electrode area is small and currents big, so that potential of a mercury dropping electrode differs from the equilibrium potential necessary for carrying out of electrochemical reaction. This phenomenon is called as electrode polarisation, and from it there was a name polarography.
Solid electrodes constructed using platinum, gold, silver, or carbon may be used over a range of potentials, including potentials that are negative and positive with respect to the SCE. For example, the potential range for a Pt electrode extends from approximately +1.2 V to –0.2 V versus the SCE in acidic solutions and from +0.7 V to -1 V versus the SCE in basic solutions. Solid electrodes, therefore, can be used in place of mercury for many voltammetric analyses requiring negative potentials and for voltammetric analyses at positive potentials at which mercury electrodes cannot be used. Except for the carbon paste electrode, solid electrodes are fashioned into disks that are sealed into the end of an inert support and are in contact with an electrical lead.
The carbon paste electrode is made by filling the cavity at the end of the inert support with a paste consisting of carbon particles and a viscous oil. Solid electrodes are not without problems, the most important of which is the ease with which the electrode’s surface may be altered by the adsorption of solution species or the formation of oxide layers. For this reason solid electrodes need frequent reconditioning, either by applying an appropriate potential or by polishing.
Current in Voltammetry
When an analyte is oxidized at the working electrode, a current passes electrons through the external electric circuitry to the auxiliary electrode, where reduction of the solvent or other components of the solution matrix occurs. Reducing an analyte at the working electrode requires a source of electrons, generating a current that flows from the auxiliary electrode to the cathode. In either case, a current resulting from redox reactions at the working and auxiliary electrodes is called a faradaic current. In this section we consider the factors affecting the magnitude of this faradaic current, as well as the source of any nonfaradaic currents.
Sign Conventions
Since the reaction of interest occurs at the working electrode, the classification of current is based on this reaction. A current due to the analyte’s reduction is called a cathodic current and, by convention, is considered positive.
Anodic currents are due to oxidation reactions and carry a negative value.
Influence of Mass Transport on the Faradaic Current
There are three modes of mass transport that influence the rate at which reactants and products are transported to and from the electrode surface: diffusion, migration, and convection.
Diffusion from a region of high concentration to a region of low concentration occurs whenever the concentration of an ion or molecule at the surface of the electrode is different from that in bulk solution. When the potential applied to the working electrode is sufficient to reduce or oxidize the analyte at the electrode surface. The volume of solution in which the concentration gradient exists is called the diffusion layer.
Without other modes of mass transport, the width of the diffusion layer, d, increases with time as the concentration of reactants near the electrode surface decreases.
The contribution of diffusion to the rate of mass transport, therefore, is time-dependent.
Convection occurs when a mechanical means is used to carry reactants toward
the electrode and to remove products from the electrode. The most common means of convection is to stir the solution using a stir bar. Other methods include rotating the electrode and incorporating the electrode into a flow cell.
The final mode of mass transport is migration, which occurs when charged particles in solution are attracted or repelled from an electrode that has a positive or negative surface charge. Thus, when the electrode is positively charged, negatively charged particles move toward the electrode, while positively charged particles move toward the bulk solution. Unlike diffusion and convection, migration only affects the mass transport of charged particles.
Influence of the Kinetics of Electron Transfer on the Faradaic Current The rate of mass transport is one factor influencing the current in a voltammetric experiment.
The ease with which electrons are transferred between the electrode and the reactants and products in solution also affects the current. When electron transfer kinetics are fast, the redox reaction is at equilibrium, and the concentrations of reactants and products at the electrode are those specified by the Nernst equation. Such systems are considered electrochemically reversible. In other systems, when electron transfer kinetics are sufficiently slow, the concentration of reactants and products at the electrode surface, and thus the current, differ from that predicted by the Nernst equation. In this case the system is electrochemically irreversible.
Nonfaradaic Currents Faradaic currents result from a redox reaction at the electrode surface. Other currents may also exist in an electrochemical cell that are unrelated to any redox reaction. These currents are called nonfaradaic currents and must be accounted for if the faradaic component of the measured current is to be determined.
The most important example of a nonfaradaic current occurs whenever the electrode’s potential is changed. In discussing migration as a means of mass transport, we noted that negatively charged particles in solution migrate toward a positively charged electrode, and positively charged particles move away from the same electrode. When an inert electrolyte is responsible for migration, the result is a structured electrode–surface interface called the electrical double layer, or EDL, the exact structure of which is of no concern in the context of this text. The movement of charged particles in solution, however, gives rise to a short-lived, nonfaradaic charging current. Changing the potential of an electrode causes a change in the structure of the EDL, producing a small charging current.
Residual Current Even in the absence of analyte, a small current inevitably flows through an electrochemical cell. This current, which is called the residual current, consists of two components: a faradaic current due to the oxidation or reduction of trace impurities, and the charging current.
QUALITATIVE POLAROGRAPHIC ANALYSIS: POLAROGRAPHIC WAVE, HALFWAVE POTENTIAL AND FACTORS OF INFLUENCE ON IT
Shape of Voltammograms
The shape of a voltammogram is determined by several experimental factors, the most important of which are how the current is measured and whether convection is included as a means of mass transport. Despite an abundance of different voltammetric techniques, several of which are discussed in this chapter, only three shapes are common for voltammograms (Figure).
The voltammogram in Figure a is characterized by a current that increases from the background residual current to a limiting current at potentials at which the analyte is oxidized or reduced. Since the magnitude of a faradaic current is inversely proportional to d, a limiting current implies that the thickness of the diffusion layer remains constant. The simplest method for obtaining a limiting current is to stir the solution (Figure 11.32), which can be accomplished with a magnetic stir bar, or by rotating the electrode. Voltammetric techniques that include convection by stirring are called hydrodynamic voltammetry. When convection is absent, the thickness of the diffusion layer increases with time, resulting in a peak current in place of a limiting current (Figure b).
In the voltammograms in Figures a and b, the current is monitored as a function of the applied potential. Alternatively, the change in current following a change in potential may be measured. The resulting voltammogram, which is shown in Figure c, also is characterized by a peak current.
QUANTITATIVE POLAROGRAPHIC ANALYSIS: ILKOVICH EQUATION AND CALCULATION OF CONTENTS OF SUBSTANCE BY RESULTS POLAROGRAPHY, APPLICATION.
Quantitative and Qualitative Aspects of Voltammetry
Earlier we described a voltammogram as the electrochemical equivalent of a spectrum in spectroscopy. In this section we consider how quantitative and qualitative information may be extracted from a voltammogram. Quantitative information is obtained by relating current to the concentration of analyte in the bulk solution.
Qualitative information is obtained from the voltammogram by extracting the standard-state potential for the redox reaction.
Determining Concentration Let’s assume that the redox reaction at the working electrode is
O + ne↔ R
and that initially only O is present in the bulk solution. The current is determined by the rate at which O diffuses through the fixed diffusion layer.
On halfwave potential influence:
§ The nature of defined substance – depolarizer;
§ The nature of indifferent electrolyte -depolarizer;
§ рН of medium;
§ Presence of ions or molecules which can form with depolarizer complex.
Polarography The earliest voltammetric experiment was normal polarography at a dropping mercury electrode. Iormal polarography the potential is linearly scanned, producing voltammograms such as that shown in Figure
Although polarography takes place in an unstirred solution, a limiting current is obtained because the falling Hg drops mix the solution. Each new Hg drop, therefore, grows in a solution whose composition is identical to that of the initial bulk solution.
Oscillations in the current are due to the growth of the Hg drop, which leads to a time-dependent change in the area of the working electrode. The limiting current, which is also called the diffusion current, may be measured from the maximum current, imax, or from the average current, iavg. The relationship between the concentration of analyte, CA, and the limiting current is given by the Ilikovic equation
where n is the number of electrons transferred in the redox reaction, D is the
analyte’s diffusion coefficient, m is the flow rate of the Hg, and t is the drop time.
The half-wave potential, E1/2, provides qualitative information about the redox reaction.
Normal polarography has been replaced by various forms of pulse polarography, 16 several examples of which are shown in Figure 11.36. Differential pulse polarography (Figure b), for example, uses a series of potential pulses characterized by a cycle of time t, a pulse time of tp, a potential pulse of DEp, and a potential step per cycle of DEs. Typical experimental conditions for differential pulse polarography are t » 1 s, tp » 50 ms, DEp » 50 mV, DEs » 2 mV. The current is measured twice, for approximately 17 ms before the forward pulse and for approximately 17 ms before the reverse pulse. The difference in the two currents gives rise to a peak-shaped voltammogram. Other forms of pulse polarography include normal pulse polarography (Figure a), staircase polarography (Figure c), and square-wave polarography (Figure 11.36d). Limiting and peak currents are directly proportional to the concentration of analyte, and half-wave and peak potentials can be used for qualitative purposes. The popularity of pulse polarography is due to a substantial improvement in sensitivity and detection limits from those in normal polarography.
Polarography is used extensively for the analysis of metal ions and inorganic anions, such as IO3– and NO3–. Organic compounds containing easily reducible or oxidizable functional groups also can be studied polarographically. Functional groups that have been used include carbonyls, carboxylic acids, and carbon–carbon double bonds.
Hydrodynamic Voltammetry In polarography a limiting current is obtained because each falling drop of mercury returns the solutioear the electrode to its initial composition. As noted earlier, a limiting current is also obtained whenever the solution is stirred during the analysis. The simplest means of stirring the solution is with a magnetic stir bar. More commonly, however, stirring is achieved by rotating the electrode.
In hydrodynamic voltammetry current is measured as a function of the potential applied to a solid working electrode. The same potential profiles used for polarography, such as a linear scan or a differential pulse, are used in hydrodynamic voltammetry. The resulting voltammograms are identical to those for polarography, except for the lack of current oscillations resulting from the growth of the mercury drops. Because hydrodynamic voltammetry is not limited to Hg electrodes, it is useful for the analysis of analytes that are reduced or oxidized at more positive potentials.
Stripping Voltammetry
One of the most important quantitative voltammetric techniques is stripping voltammetry, which is composed of three related techniques: anodic, cathodic, and adsorptive stripping voltammetry. Since anodic strip- ping voltammetry has found the widest application, we consider it in the greatest detail.
Quantitative Applications
Quantitative voltammetry has been applied to a wide variety of sample types, including environmental samples, clinical samples, pharmaceutical formulations, steels, gasoline, and oil.
Multicomponent Analysis
One advantage of voltammetry as a quantitative method of analysis is its capability for analyzing two or more analytes in a single sample. As long as the components behave independently, the resulting voltammogram for a multicomponent mixture is a summation of their respective individual voltammograms. If the separation between the half-wave potentials or peak potentials is sufficient, each component can be determined independently as if it were the only component in the sample (Figure).
The minimum separation between the half-wave potentials or peak potentials for the independent analysis of two components depends on several factors, including the type of electrode and the potential- excitation signal. For normal polarography the separation must be at least ±0.2–0.3 V, and for differential pulse voltammetry a minimum separation of ±0.04–0.05 V is needed.
Methods of quantitative analysis
§ Method of calibration chart
§ Method of additives
§ Comparison method
§ Theoretically on Ilkovich equation.
I = 607 nD1/2m2/3t1/6 × C
СХ = I / (607 nD1/2m2/3t1/6)
AMPEROMETRIC TITRATION: AN ESSENCE, TYPES OF CURVES, APPLICATION.
Amperometric titration
Amperometric titration refers to a class of titrations in which the equivalence point is determined through measurement of the electric current produced by the titration reaction. It is a form of quantitative analysis.
Types of reactions:
§ Precipitation;
§ Comlexing;
§ The oxidation-reduction
Background
Consider a solution containing the analyte, A, in the presence of some conductive buffer. If an electrolytic potential is applied to the solution through a working electrode, then the measured current depends (in part) on the concentration of the analyte. Measurement of this current can be used to determine the concentration of the analyte directly; this is a form of amperometry. However, the difficulty is that the measured current depends on several other variables, and it is not always possible to control all of them adequately. This limits the precision of direct amperometry.
If the potential applied to the working electrode is sufficient to reduce the analyte, then the concentration of analyte close to the working electrode will decrease. More of the analyte will slowly diffuse into the volume of solution close to the working electrode, restoring the concentration. If the potential applied to the working electrode is great enough (an overpotential), then the concentration of analyte next to the working electrode will depend entirely on the rate of diffusion. In such a case, the current is said to be diffusion limited. As the analyte is reduced at the working electrode, the concentration of the analyte in the whole solution will very slowly decrease; this depends on the size of the working electrode compared to the volume of the solution.
What happens if some other species which reacts with the analyte (the titrant) is added? (For instance, chromate ions can be added to oxidize lead ions.) After a small quantity of the titrant (chromate) is added, the concentration of the analyte (lead) has decreased due to the reaction with chromate. The current from the reduction of lead ion at the working electrode will decrease. The addition is repeated, and the current decreases again. A plot of the current against volume of added titrant will be a straight line.
After enough titrant has been added to react completely with the analyte, the excess titrant may itself be reduced at the working electrode. Since this is a different species with different diffusion characteristics (and different half-reaction), the slope of current versus added titrant will have a different slope after the equivalence point. This change in slope marks the equivalence point, in the same way that, for instance, the sudden change in pH marks the equivalence point in an acid-base titration.
The electrode potential may also be chosen such that the titrant is reduced, but the analyte is not. In this case, the presence of excess titrant is easily detected by the increase in current above background (charging) current.
Advantages
The chief advantage over direct amperometry is that the magnitude of the measured current is of interest only as an indicator. Thus, factors that are of critical importance to quantitative amperometry, such as the surface area of the working electrode, completely disappear from amperometric titrations.
The chief advantage over other types of titration is the selectivity offered by the electrode potential, as well as by the choice of titrant. For instance, lead ion is reduced at a potential of -0.60 V (relative to the saturated calomel electrode), while zinc ions are not; this allows the determination of lead in the presence of zinc. Clearly this advantage depends entirely on the other species present in the sample.
Indicating electrodes
§ The platinum;
§ The graphite;
§ Other firm electrodes;
§ Mercury dropping electrode.
Application:
§ amperometry is used for definition cations and anions, organic substances.
§ It is applied to substance mix titration at a combination of titration conditions.
§ Titration of the diluted solutions of 10-5 mol/L and less.
§ Titration of very painted and muddy solutions.
