Potentiometry is the field of electroanalytical chemistry in which potential is measured under the conditions of no current flow. The measured potential may then be used to determine the analytical quantity of interest, generally the concentration of some component of the analyte solution. The potential that develops in the electrochemical cell is the result of the free energy change that would occur if the chemical phenomena were to proceed until the equilibrium condition has been satisfied.
This is the method in which the potential between two electrodes is measured while the electric current (usually nearly zero) between the electrodes is controlled. In the most common forms of Potentiometry, two different types of electrodes are used. The potential of the indicator electrode varies, depending on the concentration of the analyte while the potential of the reference electrode is constant
An electrochemical cell is a device used for generating an electromotive force (voltage) and current from chemical reactions, or the reverse, inducing a chemical reaction by a flow of current. The current is caused by the reactions releasing and accepting electrons at the different ends of a conductor.
Electrochemical cell consists of:
Are usually metal strips/wires connected by an electrically conducting wire.
Is a U-shaped tube that contains a gel permeated with a solution of an inert electrolyte. Allows anions and cations to move between electrode compartments
Is the electrode where oxidation takes place.
Is the electrode where reduction takes place.
The potential difference between the cathode electrode potential and the anode electrode potential is the potential of the electrochemical cell.
Ecell =Ecathode - Eanode
The cell potential can be predicted through the use of electrode potentials (the voltages of each half-cell). The difference in voltage between electrode potentials gives a prediction for the potential measured.
Cell potentials have a possible range of about zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts, because the very powerful oxidizing and reducing agents which would be required to produce a higher cell potential tend to react with the water.
An electrochemical cell consists of two half-cells. Each half-cell consists of an electrode, and an electrolyte. The two half-cells may use the same electrolyte, or they may use different electrolytes. The chemical reactions in the cell may involve the electrolyte, the electrodes or an external substance (as in fuel cells which may use hydrogen gas as a reactant). In a full electrochemical cell, ions, atoms, or molecules from one half-cell lose electrons (oxidation) to their electrode while ions, atoms, or molecules from the other half-cell gain electrons (reduction) from their electrode. A salt bridge is often employed to provide electrical contact between two half-cells with very different electrolytes—to prevent the solutions from mixing.
Convention for expressing the cell:
An electrochemical cell showing the two half cell connected by a salt bridge or porous barrier, such as:
Anode Half-Cell || Cathode Half-Cell
Electrode | Anode Soln || Cathode Soln | Electrode
Zn(s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu(s)
Pt(s) | H2 (1 atm) | H+ (1 M) || Fe3+(aq), Fe2+(aq) | Pt(s)
Electrons flow from anode to cathode. Anode is placed on left by convention.
Types of Electrochemical cell:
There are two types of electrochemical cells. Spontaneous reactions occur in galvanic (voltaic) cells; nonspontaneous reactions occur in electrolytic cells. Both types of cells contain electrodes where the oxidation and reduction reactions occur. Oxidation occurs at the electrode termed the anode and reduction occurs at the electrode called the cathode.
i) Galvanic or Voltaic Cells
The redox reaction in a galvanic cell is a spontaneous reaction. For this reason, galvanic cells are commonly used as batteries. Galvanic cell reactions supply energy which is used to perform work. The energy is harnessed by situating the oxidation and reduction reactions in separate containers, joined by an apparatus that allows electrons to flow. A common galvanic cell is the Daniell cell, shown below.
When zinc metal placed in CuSO4
solution, following reaction take place:
Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s)
Reduction: Cu+2 + 2e-1 Cu
Overall: Zn(s) + Cu+2 Zn+2 + Cu(s)
The salt bridge is necessary to complete the circuit and maintain charge neutrality.
Electrons are transferred from Zn° to Cu+2 through a wire producing electrical energy.
ii) Electrolytic Cells
The redox reaction in an electrolytic cell is nonspontaneous. Electrical energy is required to induce the electrolysis reaction. An example of an electrolytic cell is shown below, in which molten NaCl is electrolyzed to form liquid sodium and chlorine gas. The sodium ions migrate toward the cathode, where they are reduced to sodium metal. Similarly, chloride ions migrate to the anode and are oxided to form chlorine gas. This type of cell is used to produce sodium and chlorine. The chlorine gas can be collected surrounding the cell. The sodium metal is less dense than the molten salt and is removed as it floats to the top of the reaction container.
Types of electrodes in Potentiometry:
Potentiometric electrodes measure:
Activity not concentration
Concepts to review:
Activity and affect factors
There are two electrodes are required in Potentiometry
Indicator Electrode :
Potential responds to activity of species of interest
Reference Electrode :
Chosen so that it's potential is independent of solution composition.
Used for half of the cell to determine the potential of the analyte of interest.
Maintains a fixed potential (i.e. reference, stable over time) –in contrast, the indicator electrode responds to the analyte activity and Follows Nernst equation
Reference electrode completes the cell but dose not respond to analyte .its usually separated from the test solution (analyte) by salt bridge
Examples of reference electrodes
1) Saturated Calomel Electrode (SCE):
Hg2Cl2(s)+ 2e−↔ 2Hg(l)+ 2Cl−
or Hg|Hg2Cl2(sat'd, KCl) ||...
Structure of Saturated Calomel Electrode (SCE):
2) Silver/Silver Chloride Electrode:
This generally consists of a cylindrical glass tube containing a 4 Molar solution of KCl saturated with AgCl. The lower end is sealed with a porous ceramic frit which allows the slow passage of the internal filling solution and forms the liquid junction with the external test solution. Dipping into the filling solution is a silver wire coated with a layer of silver chloride (it is chloridised) which is joined to a low-noise cable which connects to the measuring system.
In electrochemical terms, the half-cell can be represented by:
Ag / AgCl (Satd), KCL (Satd)
And the electrode reaction is:
AgCl (s) + e- = Ag (s) + Cl-
The electrode potential for this half-cell is + 0.2046 V relative to the Standard Hydrogen Electrode at 25°C
Double Junction Reference Electrodes.
One problem with reference electrodes is that, in order to ensure a stable voltage, it is necessary to maintain a steady flow of electrolyte through the porous frit. Thus there is a gradual contamination of the test solution with electrolyte ions. This can cause problems when trying to measure low levels of K, Cl, or Ag, or when using other ISEs with which these elements may cause interference problems. In order to overcome this difficulty the double junction reference electrode was developed. In this case the silver / silver chloride cell described above forms the inner element and this is inserted into an outer tube containing a different electrolyte which is then in contact with the outer test solution through a second porous frit. The outer filling solution is said to form a "salt bridge" between the inner reference system and the test solution and is chosen so that it does not contaminate the test solution with any ions which would affect the analysis.
Commonly used outer filling solutions are:
Potassium nitrate - for Br, Cd, Cl, Cu, CN, I, Pb, Hg, Ag, S, SCN.
Sodium chloride - for K,
Ammonium sulphate - for N03,
Magnesium sulphate - for NH4,
Note that double junction reference electrodes are named after their outer filling solutions.
One disadvantage with double junction reference electrodes is that they introduce an extra interface between two electrolytes and thus give the opportunity for an extra liquid junction potential to develop.
Liquid Junction Potentials.
It must be noted that the standard voltage given by a reference electrode is only correct if there is no additional voltage supplied by a liquid junction potential formed at the porous plug between the filling solution and the external test solution. Liquid junction potentials can appear whenever two dissimilar electrolytes come into contact. At this junction, a potential difference will develop as a result of the tendency of the smaller and faster ions to move across the boundary more quickly than those of lower mobility. These potentials are difficult to reproduce, tend to be unstable, and are seldom known with any accuracy; so steps must be taken to minimise them. Using 4 Molar KCL as the inner filling solution has the advantage that the K+ and Cl- ions have nearly equal mobilities and hence form an equi-transferrent solution. Also, in the single junction electrodes, the electrolyte concentration is much higher than that of the sample solution thus ensuring that the major portion of the current is carried by these ions. A third factor in minimising the junction potential is the fact that there is a small but constant flow of electrolyte out from the electrode thus inhibiting any back-diffusion of sample ions - although this is less important with modern gel electrolytes.
As indicated above, all these problems are doubled when double junction reference electrodes are used and an additional problem arises in the case of the last three listed above (Sodium Chloride, Ammonium Sulphate, Magnesium Sulphate) because the filling solutions are not equi-transferrent and hence have a stronger tendency to form liquid junction potentials. It must be noted here that Nico2000 Ltd have recently introduced a novel Lithium Acetate reference electrode which overcomes most of these problems and can be used with all but one of the ELIT range of ISEs. This is because it contains ions which are very nearly equi-tranferrent and which do not interfere with any of the commonly used ISEs - the only known interference being that of acetate on the carbonate electrode.
It must be noted that the E0 factor in the Nernst equation is the sum of all the liquid junction potentials present in the system and any variation in this during analyses can be a major source of potential drift and error in measurements.
The majority of pH electrodes are produced in the form of combination electrodes in which the reference system is housed in the same cylindrical body as the sensor head. This produces a simple, compact unit for immersing in the test solution and has the added advantage that the two cells are in close proximity (with the reference cell normally completely surrounding the sensor element) - thus minimising the effect of any stray electrostatic fields or any inhomogeneity in the test solution. The main disadvantage of this arrangement is the fact that it is the reference element which is the most likely to cause problems or fail, long before the ISE head does, but the whole unit has to be replaced when failure does occur.
In contrast to pH electrodes, some ISEs are produced as mono-electrodes for use with separate reference systems. One reason for this is because ISE membranes have a far lower impedance than pH sensors and are less susceptible to stray electrostatic fields. Thus it is not necessary to screen the sensor head by surrounding it with the reference system. More importantly, the membranes and internal construction of ISEs are generally far more expensive than pH sensors and it is much more cost-effective to have separate units in which the reference system can be replaced independently from the ISE.
Multiple Electrode Heads: Separable Combinations.
A new concept for combination electrodes has recently been introduced. Both the ISEs and the reference electrodes are made in the form of 8mm diameter tubes fitted with a gold plated plug-in connector. These can be inserted separately into special multiple electrode heads which are fitted with the cables and connectors for attaching to the measuring system. The rigid plastic head ensures that the ISE and reference system remain firmly linked together at a regular distance apart during operation, but either one can easily be replaced in the event of failure or need to change the analysis. Moreover, the replacement electrodes are relatively inexpensive compared to conventional electrodes because they do not incorporate the expensive low-noise cables.
Indicator Electrode :
Potential responds to activity of species of interest
Metal Indicator Electrodes
These electrodes develop an electric potential in response to a redox reaction at the metal surface.
These measurements are made against a reference electrode (which has a known, stable potential).
Typically, Pt or Au are used for the metal indicator electrode as they are inert (do not contribute to the reaction).
Metal electrodes of the “first kind”respond directly to changing activity of electrode ion –this type of electrode is not selective for a specific
i) Electrode with the first kind:
A metal in contact with a solution containing its cation
Eg: sliver metal electrode dipping in solution of silver nitrate
The half reaction of electrode Ag+ +e Ag
The potential of electrode E =E0 Ag/Ag+ - 0.0591log 1/[aAg+ ]
Increasing cation activity aAg+ always causes the electrode potential increase.(the case for any electrode measuring a cation)
Ecell = Eind - Eref
This kind of electrode used to monitor the activity of metal ion in solution .
ii)Electrode with the second kind:
The general form of this electrode is M/MX/Xn- .example sliver silver chloride Ag/AgCl/Cl- .
The half reaction of electrode is AgCl(s) +e = Ag(S) +Cl-
The potential of electrode is E =E0 AgCl/Ag -0.0591log aCl-
This electrode can be used to measure the activity of Cl- and Ag+ ions in solution.
Increasing anion activity always causes the electrode potential decrease. The activity of aAg+ is determinded by Ksp and Cl-,
the Ksp = aAg+ x aCl- then aCl- =Ksp /aAg+
E = E0Agcl,Ag+ - 0.0591ksp /aAg+
E= E0Agcl,Ag+ - 0.0591ogksp – 0.0591log 1/aAg+
The Ag electrode used to monitor other anions that form soluble salt with Ag- such as I-,Br- and S-2
iii)Redox electrode :
Is an electrode made from electron conductive material and characterized by high chemical stability in studied solutions. It uses for potentiometric aims (measuring of redox potential of a specific redox system in solution) and for electrochemical studies (investigation of electrochemical kinetics of interfacial processes).
Correlation of an electrode potential and redox system
composition can be realized by Nernst
Ox + e- = Red
An inert metal Pt is in contact with solution containing the soluble oxidized and reduced forms of the half reaction
The half reaction: Ma+ + ne = M(a-n)+
The potential of this electrode is
E = E0Ma+, M(a-n)+ -0.0591/n log aM(a-n)+/a Ma+
MnO-4 + 8 H+ + 5e- àß Mn2+ + 4 H2O
The E0 of an inert electrode is determined by the ratio of reduced and oxidized species .
Avery important example of this type of electrode is hydrogen electrode Pt/H2,H+
H+ + e= 1/2H2
E= E0H+,H2 - 0.059log (pH2)1/2 /aH+
P is presures
In 1 atmosphere the place of activities in above equation since the E0 equal zero then E =E0 -0.059log 1/aH+ =-logPH
Structure of hydrogen electrode:
The hydrogen electrode is easy to build. All that is required is to bubble hydrogen through an acid solution of known pH so that it is saturated with hydrogen. An (ideal) noble metal electrode is placed into the solution. Platinized platinum is generally used to insure a large electrochemical surface area and rapid equilibrium conditions. The pressure of the hydrogen must also be known.
is a transducer (sensor) which converts the activity of a specific ion dissolved in a solution into an electrical potential which can be measured by a voltmeter or pH meter. The voltage is theoretically dependent on the logarithm of the ionic activity, according to the Nernst equation. The sensing part of the electrode is usually made as an ion-specific membrane, along with a reference electrode. Ion-selective electrodes are used in biochemical and biophysical research, where measurements of ionic concentration in an aqueous solution are required, usually on a real time basis
Ion Selective Electrodes (including the most common pH electrode) work on the basic pricipal of the galvanic cell (Meyer Hoff and Opdycke). By measuring the electric potential generated across a membrane by "selected" ions, and comparing it to a reference electrode, a net charge is determined (Figure 2). The strength of this charge is directly proportional to the concentration of the selected ion. The basic formula is given for the galvanic cell:
Ecell = Eise - Eref
Where the potential for the cell is equivelent to the potential of the ISE minus the potential of the reference electrode.
The above figure demonstrates the flow of current through the reference electrode to the ISE, completing the circuit. Modified and reproduced with permission from Orion Research.
Ion-selective electrodes are used in a wide variety of applications for determining the concentrations of various ions in aqueous solutions. The following is a list of some of the main areas in which ISEs have been used.
Pollution Monitoring: CN, F, S, Cl, NO3 etc., in effluents, and natural waters.
Agriculture: NO3, Cl, NH4, K, Ca, I, CN in soils, plant material, fertilisers and feedstuffs.
Food Processing: NO3, NO2 in meat preservatives.
Salt content of meat, fish, dairy products, fruit juices, brewing solutions.
F in drinking water and other drinks.
Ca in dairy products and beer.
K in fruit juices and wine making.
Corrosive effect of NO3 in canned foods.
Detergent Manufacture: Ca, Ba, F for studying effects on water quality.
Paper Manufacture: S and Cl in pulping and recovery-cycle liquors.
Explosives: F, Cl, NO3 in explosive materials and combustion products.
Electroplating: F and Cl in etching baths; S in anodising baths.
Biomedical Laboratories: Ca, K, Cl in body fluids (blood, plasma, serum, sweat).
F in skeletal and dental studies.
Education and Research: Wide range of applications.
1) When compared to many other analytical techniques, Ion-Selective Electrodes are relatively inexpensive and simple to use and have an extremely wide range of applications and wide concentration range.
2) The most recent plastic-bodied all-solid-state or gel-filled models are very robust and durable and ideal for use in either field or laboratory environments.
3) Under the most favorable conditions, when measuring ions in relatively dilute aqueous solutions and where interfering ions are not a problem, they can be used very rapidly and easily (e.g. simply dipping in lakes or rivers, dangling from a bridge or dragging behind a boat).
4) They are particularly useful in applications where only an order of magnitude concentration is required, or it is only necessary to know that a particular ion is below a certain concentration level.
5) They are invaluable for the continuous monitoring of changes in concentration: e.g. in potentiometric titrations or monitoring the uptake of nutrients, or the consumption of reagents.
6) They are particularly useful in biological/medical applications because they measure the activity of the ion directly, rather than the concentration.
7) In applications where interfering ions, pH levels, or high concentrations are a problem, then many manufacturers can supply a library of specialized experimental methods and special reagents to overcome many of these difficulties.
8) With careful use, frequent calibration, and an awareness of the limitations, they can achieve accuracy and precision levels of ± 2 or 3% for some ions and thus compare favourably with analytical techniques which require far more complex and expensive instrumentation.
9) ISEs are one of the few techniques which can measure both positive and negative ions.
10) They are unaffected by sample color or turbidity.
11) ISEs can be used in aqueous solutions over a wide temperature range. Crystal membranes can operate in the range 0°C to 80°C and plastic membranes from 0°C to 50°C.
Types of ion-selective membrane
There are four main types of ion-selective membrane used in ion-selective electrodes: glass, solid state, liquid based and compound electrode.
Glass membranes are made from an ion-exchange type of glass (silicate of chalcogenide). This type of ISE has good selectivity, but only for several single-charged cations; mainly H+, Na+, and Ag+. Chalcogenide glass also has selectivity for double-charged metal ions, such as Pb2+, and Cd2+. The glass membrane has excellent chemical durability and can work in very aggressive media. A very common example of this type of electrode is the pH glass electrode.
Glass membrane electrodes are very good indicator electrodes in Potentiometry
Must exercise care in calibration and in maintaining integrity of glass membrane
Some errors exist & are unavoidable ,Glass electrodes available for Na+, K+,NH4+, Rb+, Cs+, Li+, Ag+ (cations only) by varying glass composition
Combination electrodes combine pH & ref.
E = K’ – 0.0591 pH
Combine with reference electrode and meter
Half cell voltage proportional to pH
Intercept is K’, no Eo
Calibrate with buffers
Crystalline membranes are made from mono- or polycrystalline of a single substance. They have good selectivity, because only ions which can introduce themselves into the crystal structure can interfere with the electrode response. Selectivity of crystalline membranes can be for both cation and anion of the membrane-forming substance. An example is the fluoride selective electrode based on LaF3 crystals.
Two general types -crystalline and non-crystalline membranes
Glass -silicate glasses for H+, Na+
Liquid -liquid ion exchanger for Ca2+
Immobilized liquid -liquid/PVC matrix for Ca2+and NO3-
Single crystal -LaF3for F-Polycrystalline
Mixed crystal –AgS for S2-and Ag
Fluoride selective electrode:
The Fluoride electrode is a typical example of the first type. Here the membrane consists of a single lanthanum fluoride crystal which has been doped with europium fluoride to reduce the bulk resistivity of the crystal. It is 100% selective for F- ions and is only interfered with by OH- which reacts with the lanthanum to form lanthanum hydroxide, with the consequent release of extra F- ions. This interference can be eliminated by adding a pH buffer to the samples to keep the pH in the range 4 to 8 and hence ensure a low OH- concentration in the solutions.
Solid State Membrane Electrodes
Detection limits depend on solubility of the solid state membrane
Ksp for AgCl = approx. 10-10
Therefore solubility is 10-5 M or membrane starts to produce ions of interest in solution
Mixed crystals improve this somewhat but it is still a limitation
Interferences or poisoning by high affinity ions
Can polish electrodes to remove fouling
Selectivity coefficient = electrode response ratio
Ion-exchange resins are based on special organic polymer membranes which contain a specific ion-exchange substance (resin). This is the most widespread type of ion-specific electrode. Usage of specific resins allows preparation of selective electrodes for tens of different ions, both single-atom or multi-atom. They are also the most widespread electrodes with anionic selectivity. However, such electrodes have low chemical and physical durability as well as "survival time". An example is the potassium selective electrode, based on valinomycin as an ion-exchange agent.
Liquid Membrane Electrodes
Principle of Ca2+ electrode is the same as for glass electrode, however, since Ca2+ is
Divalent n = 2 _ Nernstian slope = 29.5 mV per 10 fold change in concentration
Detection limit for Ca2+ is approx. 10-5 M
– Independent of pH from 5.5 to 11
– 50 times better for Ca2+ than for Mg2+
– 1000 times better for Ca2+ than Na+ or K+
Other liquid membrane electrodes available
Calcium Electrode is good example
Liquid ion exchanger
– Water immiscible
organic compound with phosphate
groups selective for Ca2+ in a hydrophobic membrane
1-Analytical chemistry an introduction, 7th edition, Harcourt College, 2000
2- Gray D.christian analytical chemistry, 5th, John Willely& Sons,inc 1994