electroanalytical – introduction ch 21, 7 th e, wmds) four basic electroanalytical methods of...

Electroanalytical – IntroductionCh 21, 7th e, WMDS)

Four Basic Electroanalytical Methods of Analysis

1) Potentiometric

2) Voltammetric, Polarographic, Amperometric

3) Electrolysis: electrogravimetric and coulometric

4) Conductiometric

We will briefly describe each one, then focus individually

on each.


1) Potentiometric - The electrical potential of a galvanic (spontaneous cell) is related to the concentration of the analyte by the Nernst Equation. Examples of potentiometric include pH measurements and the use of ion-selective electrodes. Potentiometric methods may also be used to monitor the course of redox, acid/base/complexometric, and precipitation reactions.


2) Voltammetric, Polarographic, Amperometric – in these methods, the electrical current at an electrode in an electrolysis (nonspontaneous) cell is related to the concentration of the analyte. They vary according to the parameter varied and the quantity measured. Generally, calibration curves are constructed to show the relationship between the measured quantity and the concentration of the analyte.


3) Electrolysis: electrogravimetry and coulometry- electrolysis is done on the nonspontaneous electrochemical cell by an external power supply.

a - In electrogravimetric methods the sought analyte is electrochemically deposited onto a previously weight electrode (generally the cathode) and the amount of analyte determined by the increase in the mass on the electrode.


3) Electrolysis: electrogravimetry and coulometry b - In coulometric methods, the amount of an electrochemical species oxidized or reduced is related to the amount of current, time and number of electrons in that half cell reaction. The number of coulombs (C) of electricity Q = it where i is current in amperes and t time in sec. One Faraday (F) (96,485 C) of charge is equivalent to 1 mole of electrons, so in combination with the half cell one can quantify the amount of electrochemical work done.


4) Conductiometric – The electrical conduction of a solution is related to the number of ions present in solution. Method is especially useful to monitor the purity of deionized water. It is also useful to monitor titrations if the electrical conductance of the titrant and analyte are significant different, or if the product at the equivalence is different than the reactants.


Potentiometry – In order to make potentiometric measurements one must have both a reference and an indicator electrode.

Reference ElectrodesThe point of beginning for the construction and definition of reference electrodes is with the standard hydrogen electrode (SHE), the primary reference standard electrode. It is based on the half-cell reaction

H+ + e = ½H2. Whenever the [H+ ] = 1.00 m and pH2 = 1.00 atm (or in SI, 1 bar), this electrode has a defined potential of 0.0000 volts.


Potentiometry – Reference Electrodes SHEUnfortunately the SHE is not a very convenient reference electrode. 1) It is quite fragile2) hazardous because of the hydrogen gas3) difficult to maintain the standard acid solution at 1.00 m.

Although it is not used in the laboratory for measurements, it is the primary reference electrode against which all other electrodes are ultimately standardized.


Potentiometry – Reference Electrodes Secondary Reference Electrodes.

There are 2 commonly used secondary reference electrodes. The first is the saturated calomel electrode (SCE) whose half-cell reaction is

Hg2Cl2(s) + 2e = 2Hg(l) + 2Cl-(aq). Measured against the SHE with saturated KCl as the electrolyte, its potential is 0.2445 volts @ 25oC. The potential is also dependent of the concentration of chloride, so the saturated KCl is the most commonly used. Literature gives the potential at other temperatures, as well as with other concentrations of KCl.


Potentiometry – Reference Electrodes Secondary Reference Electrodes.

The second widely used laboratory reference electrode is the silver-silver chloride electrode whose half-cell reaction is

AgCl(s) + e = Ag(s) + Cl-(aq). The potential of the Ag-AgCl system is 0.2223 volts @ 25oC and a saturated solution of KCl is used as its electrolyte. Literature gives the potential at other temperatures.


Indicator Electrodes

Indicator electrodes are classified according to the process that produces the electrode potential.

Metal indicator electrodes develop a potential dependent on the position of the equilibrium of the redox half-reaction at the surface of the electrode. Membrane indicator electrodes develop a potential determined by the difference in the concentration of particular ions across a special thin layer known as a membrane.


Metal indicator electrodes are classified as either first-order or second-order electrode.

A first order electrode involves the metal in contact with its own ions, such as Ag, Ag+ or Zn, Zn. Only a few metals give reproducible potentials as first-order electrodes due to crystalline irregularities on their surfaces and the ease of forming oxides.

A second-order electrode is one that responds to the presence of precipitating or complexing ions. For example, silver wire could serve as the indicator electrode for chloride. The use of second-order indicator electrodes are limited.


The most commonly used indicator electrodes are known as inert. These electrodes are not involved in the half-cell reactions of the electrochemical species. Typical inert electrodes are platinum, gold, and carbon. Inert electrodes are responsive to any reversible redox system; these are widely used in potentiometric work.


There are several different types of membrane indicator electrodes according to the type of membrane and the specific use of that electrode. Membranes may be made of glass, polymers, or crystals. The first and most studied membrane indicator electrode is the glass electrode, most commonly used as the pH electrode. The membrane is a very thin layer of specially composition glass, generally containing CaO, BaO, Li2O and Na2O as well as SiO2. The internal solution is of known constant hydrogen ion activity, such as pH 7.


When this system is placed in an external solution other hydrogen ion activity, a Nernstian potential develops across the glass membrane. Although the response follows the Nernst equation

E = 0.05916/1 log ([H+]external /[H+]internal)there are three important facts that must be considered.


Factors in the Use of Glass electrodes

1 - The electrical resistance of the glass membrane is very high (~ 100 M ), so special high input resistance voltmeters

(electrometers with Rinput ~ 106 M ), are required to obtain accurate potential readings.



2 - In use, a small error, known as the asymmetry potential develops across the membrane so that frequent re-standardization of the electrode is required. Standardization consists of placing the electrode in a buffer of known pH and then electrically adjusting the meter to give the reading of that buffer.



The composition and structure of the glass membrane affects the selectivity of the glass membrane electrode. Membrane electrodes are not exclusive in their response to one single ion, but response to other ions of similar size and charge, so the potential may be a combination of the presence of several ions. The response for a desired ion A to the response to B is given by the selectivity coefficient k B/A .


Selectivity Coefficient k B/A . Small values of k B/A, such as 0.001 are desirable. This meansthat only at equal concentration of A and B, only 1/1000 of thepotential comes from B. A typical glass electrode for pHmeasurements may have appreciable Na ion error, especially athigh pH (when [H+] is low). Low sodium error electrodes areavailable when needed.


Ion-selective electrodes (ISE)

In addition to the glass electrode for pH measurements, thereis a wide range of membrane electrodes for the measurementof specific ions. Although sometimes called ion-specificelectrodes one must realize that they are really selective fordefinite ions, but have the same limitations of selectivity as theGlass electrodes discussed earlier. They should really betermed as ion-selective electrodes (ISE). All of the types of membranes described earlier are used in their construction.


Polymer membrane ISE, similar to the Calcium ISE


The Flouride ISE, a solid crystal membrane ISE;


A calibration curve for the F ISE; other electrodes would have different y-intercept values. The theoretical slope is 59.16/n where n is the charge of the ion.


Ion Conc Range, M

Membrane

pH Range

Interferences

F– 10 –6 – 1 LaF3 5 – 8 OH–

Cl– 10 –4 – 1 AgCl 2 – 11 CN–, S–2, I–, S2O3

–2, Br –

Br– 10 –5 – 1 AgBr 2 – 12 CN–, S–2, I–

I– 10 –6 – 1 AgI 3 – 12 S–2

CN– 10 –6 – 10 –2 AgI 11 – 13 S–2 , I–

S–2 10 –5 – 1 Ag2S 13 – 14

Typical Solid-state (Crystal-membrane) ISE (from Harris, 2e, p 316)


At the present time the following Ion-selective electrodes have been developed:

NH4+, Ca+2, Mg+2, K+, Cl–, S–2, F–, ClO4

–, NO3–, Br –, I– ,

CN–, Ag+, Cu+2, Pb+2, Cd+2 ;

In addition to the above, by using permeable membranes over glass electrodes that measure the amounts of various gases or ofspecific metabolites have been developed, such as the CO2, urea, or glucose electrodes.


A CO2 secondary membraneelectrode


The measurement of the potential of ion-selective electrodes isaccomplished by making an electrochemical cell consisting ofthe ion-selective electrode and a suitable reference electrode.

Emeasured = Eindicator – Ereference. Since the potential of the reference electrode is known,

Eindicator = Ecell + Ereference .There are three different ways that potentiometric measurements may be done in the laboratory:1 - Direct2 - Electrode Calibration 3 - Standard Addition


1 - Direct – the measured potential is related to the concentration of the analyte by the Nernst equation,

0.05916/n pAwhere pA is the –log [A]. Although this method is straight forward and simple, it does not account for matrix effects, i.e., the effect of other ions, etc in the sample. This method is useful for continuous monitoring and is thus used in the analysis of industrial and natural processes such as effluent streams.


2 - Electrode Calibration – electrode response is calibrated against solutions of known [analyte]. Often a plot is constructed of Emeasured vs pA and used as both a check of range of linearity of the response of the electrode, and to use as a calibration curve.


Standard Addition – in this method the electrode is used to first measure the unknown sample and then a small known amount of the analyte is added and the measurement taken. The incremental change in potential was caused by the known addition. This then allows the analyst to calculate the concentration of the analyte in the original sample. This technique is best to assure that the ionic strength (I or ) of the solution (and therefore the activity) is not causing a significant error.


Standard Addition –.The following relationship is useful in the calculation of the concentration of the analyte using the standard addition method;

E = E2 – E1 = S log ((C1 + C/C1)) Eqn 22.18

where S is the slope (emf/pC) .


Standard Addition –.Alternately, I find the following relationship from Hargis to be friendlier in working with real data:

[A] = ___________[S] x VS________

_ [(VA + Vs) 10 –n(E/0.0592 ] – VA

where A is the analyte, S the standard solution, V volume, n number of electrons, and E = E2 – E1.

Hargis, “Analytical Chemistry “, p 344, ©1988, Prentice-Hall


Sample Standard - Addition Problem The wastewater from an industrial processing plant was routinely analyzed for lead as required by EPA. A Pb ISE and a SCE reference were placed in a 50.0 mL sample. The potential was measured to be –0.118 volts; after the addition of 5.00 mL of a 0.00600 M solution of Pb+2 standard solution was added, the potential was –0.109 volts. What is the approximate concentration of Pb+2 in the wastewater? Also express the [Pb+2] in ppm Solution: [Pb+2] = (6.00 x 10–3)(5.00) /[(50.0 + 5.00) 10–(2)(-0.118-(-0.109))/0.0592] – 50)

= (3.00 x 10–2) /(55.0 x 100.304 – 50.0)= (0.0300/(110.7 – 50.0) = 0.000494 M

= (0.000494 mmol/mL) x 207.2mg/mmol = 0.1024 mg/mL ≈ 0.1024 mg/g = 102.4 g/g = 102 ppm

electroanalytical – introduction ch 21, 7 th e, wmds) four basic electroanalytical methods of...

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