Dong-Sun Lee / cat - lab / SWU 2012-Fall version

Chapter 22

Electrogravimetry and

coulometry

Example of an electroplated object is the Oscar (brittanium amalgam: Sn, Cu, Sb) which is given to recipients of Academy Awards. Award for Best Actress in 2007, Festival de Cannes.

Secret Sunshine

Ampère, André (1775-1836)

French mathematician and physicist who extended Oersted's results by showing that the deflection of a compass relative to an electrical current obeyed the right hand rule. Ampère argued that magnetism could be explained by electric currents in molecules, and invented the solenoid, which behaved as a bar magnet. Ampère also showed that parallel wires with current in the same direction attract, those with current in opposite directions repel. He dubbed the study of currents electrodynamics, and also developed a wave theory of heat. Ampère maintained that magnetic forces were linear, but this proposition was questioned and disproved by Faraday.

Electricity

The flow of an electric current, usually in a wire or other solid conductor, but possibly in a plasma or other conducting medium.

The energy in electricity can be converted into other forms and thus used to do mechanical work.

The amount of charge passing a given point per unit time from electric flow is called the current, while the energy per unit charge of the flow is called the voltage (or electric potential).

A configuration of components through which electricity is made to flow is called an electric circuit.

Electrochemical Cells : Galvanic and Electrolytic Cells

Oxidation-reduction or redox reactions take place in electrochemical cells. An electrochemical cell can be created by placing metallic electrodes into an electrolyte where a chemical reaction either uses or generates an electric current. There are two types of electrochemical cells. Electrochemical cells which generate an electric current are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells. In other electrochemical cells an externally supplied electric current is used to drive a chemical reaction which would not occur spontaneously. Such cells are called electrolytic 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.

Potentiometry vs electrogravimetric and coulometric analysis

1) Potentiometry , Redox titrations :

spontaneous electrochemical reactions

2) Electrogravimetry, coulometry :

nonspontaneous redox reactions by external source of electricity

Electrodes & Charge

The anode of an electrolytic cell is positive (cathode is negative), since the anode attracts anions from the solution.

However, the anode of a galvanic cell is negatively charged, since the spontaneous oxidation at the anode is the source of the cell's electrons or negative charge.

The cathode of a galvanic cell is its positive terminal.

In both galvanic and electrolytic cells, oxidation takes place at the anode and electrons flow from the anode to the cathode.

Cell convention

- The metal in the half-reaction where oxidation is occurring is called the Anode

- The metal in the half-reaction where reduction is occurring is called the Cathode.

- The cathode is often labeled with a "+"; "this electrode attracts electrons" .

- The anode is often labeled with a "–"; "this electrode repels electrons" .

- These definitions apply to both galvanic and electrolytic cells.

Anode

Is where oxidation occurs

Is where electrons are produced

Is what anions migrate toward

Has a negative sign

Cathode

Is where reduction occurs

Is where electrons are consumed

Is what cations migrate toward

Has a positive sign

Remember the following:

RED CAT & AN OX

· REDuction occurs at the CAThode OXidation occurs at the ANode

Galvanic or Voltaic Cells

Voltaic cell is a simple device with which chemical energy is converted into electrical energy. Voltaic cell is any cell that generates an electric current by an oxidation –reduction reaction.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

e

V+–Anode

ZnCathode Cu

ZnSO4 CuSO4Porous fritted disk (liquid junction)

A voltaic cell in a circuit consists of :

1) a negative electrode -anode

2) a positive electrode - cathode

3) an electrolyte

Cell potential : a measure of difference electron energy between the two electrodes

Open-circuit potential (zero-current potential) : can be calculated by thermodynamic data (Eo of half reactions)

Cathode (Red) : Cu2+ + 2e = Cu (s) + 0.337 V

Anode (Ox) : Zn(s) = Zn2+ + 2e – 0.763 V

Net reaction : Zn (s) + Cu2+ = Zn2+ +Cu (s) +1.100 V

Movement of charge in a galvanic cell :

left-to right flow of positive ions

right-to left flow of negative ions

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.

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.

Principles of electrolysis

Electrolysis is the process in which a reaction is driven in its nonspontaneous direction by the application of an electric current.

Endergonic reaction G>0

NOTE: electrolysis is the process of driving an electrochemical reaction in its non‑spontaneous direction through the application of voltage/current.

To accurately assess voltage-current relationships, we must consider some sources of non-Nernstian potentials :

Ohmic (solution) potential, Concentration polarization, Overpotential.

Electrolysis experiment

Direct current (dc) is current that is always in one direction; it is unidirectional. The direction of alternating current (ac) reverses periodically.

DC voltage sources are often given the battery symbol with + and – polarities.

An arrow through the battery indicates that the source voltage is variable and can be changed to another dc value.

Cathode (working electrode):

2 Cu 2+ + 2e = Cu(s)

Anode (counter electrode):

H2O = ½O2(g) + 2H+ + 2e

Net reaction:

H2O + Cu2+ = Cu(s) + ½O2(g) + 2H+

Ag | AgCl(s), Cl– (0.200M), Cd 2+ (0.00500M) | Cd

Cd 2+ + 2Ag(s) + 2Cl– Cd(s) + 2 AgCl(s)

Reduction

Working electrode operates as a cathode when apply a potential somewhat more negative than a thermodynamic potential of – 0.734V.

An electrolytic cell for determining Cd2+.

(a) Current=0.00mA.

(b) Schematic of cell in (a) with internal resistance of cell represented by a 15.0Ω resistor and Eapplied increased to give a

current of 2.00mA.

Current and potential changes during an electrolysis

Whenever current flows, three factors act to decrease the output voltage of galvanic cell or to increase the applied voltage needed for electrolysis.

1) Ohmic potential ; Ohmic drop The voltage needed to force current (ions) to flow through the cell.

Eohmic = IR

The output voltage of a galvanic cell is decreased by IR.

Egalvanic = Eequilibrium – IR

The magnitude of the applied voltage for an electrolysis cell must be more negative than the thermodynamic cell potential by IR in order for current flow.

Eapplied = Ecathode – Eanode – IR

= Ecell – IR

Ex. Ag | AgCl(s), Cl– (0.200M), Cd2+ (0.00500M) | Cd

Assume that the internal resistance of the cell is 15.0 .

Calculate the potential that must be applied to cause an electrolytic current

of 2.00 mA to develop.

Cd2+ + 2e Cd (s) Eo = – 0.403 V

AgCl (s) + e Ag (s) + Cl– Eo = 0.222 V

Ecathode = – 0.403 – (0.05916 / 2) log (1 / 0.00500) = – 0.471 V

Eanode = 0.222 – (0.05916 / 1) log (0.200 / 1) = 0.263 V

Eapplied = Ecell – IR = Ecathode – Eanode – IR

= – 0.471 V – 0.263 V – (2.00 × 10–3 A) × 15.0

= – 0.764 V

2) Polarization effects

I = (Ecell – Eapplied) / R = – (Eapplied / R) + (Ecell / R)

Overvoltage () is the potential difference between the theoretical cell potential from Eapplied = Ecell – IR and the actual cell potential at a given level of current.

Eapplied = Ecell – IR –

The term polarization refers to the deviation of the electrode potential from the value predicted by the Nernst equation on the passage of current. Cells that exhibit nonlinear behavior at higher currents exhibit polarization, and the degree of polarization is given by an overvoltage or overpotential.

Experimental current/voltage curve for operation of the cell shown in Figure 22-1. Dashed line is the theoretical curve assuming no polarization. Overvoltage ∏ is the potential difference between the theoretical curve and the experimental.

Factors that influenced polarization

1> electrode size, shape, and composition

2> composition of the electrolyte solution

3> temperature and stirring rate

4> current level

5> physical state of species involved in the cell reaction

Two categories of polarization phenomena

1> Concentration polarization

2> Kinetic polarization

Concentration Polarization

Concentration polarization occurs because of the finite rate of mass transfer from the solution and an electrode surface.

The electrode potential depends on the concentration of species in the region immediately surrounding the electrode.

When ions are not transported to or from an electrode as rapidly as they are consumed or created, we say that concentration polarization exists. That is, concentration polarization means that [X]s [X]o, , where [X]o is the concentration of X in the bulk solution and [X]s is concentration of X in the immediate vicinity of the electrode surface.

Reactants are transported to and products away from an electrode by three mechanisms:

(1) diffusion, (2) migration, (3) convection (as a result of stirring, vibration, or temperature gradients)

To decrease concentration polarization :

(1) Raise the temperature.

(2) Increase stirring

(3) Increase electrode surface area.

(4) Change ionic strength

Pictorial diagram (a) and concentration vs. distance plot (b) showing concentration change at the surface of a cadmium electrode. As Cd2+ ions are reduced to Cd atoms at the electrode surface, the concentration of Cd2+ at the surface becomes smaller than the bulk concentration. Ions then diffuse from the bulk of the solution to the surface as a result of the concentration gradient. The higher the current, the larger the concentration gradient until the surface concentration falls to zero, its lowest possible current, called the limiting current, is obtained.

Current-potential curve for electrolysis showing the linear or ohmic region, the onset of polarization, and the limiting current plateau. In the limiting current region, the electrode is said to be completely polarized, since its potential can be changed widely without affecting the current.

Migration involves the movement of ions through a solution as a result of electrostatic attraction between the ions and the electrodes.

Migration of analyte species can be minimized by having a high concentration of an inert electrolyte, called a supporting electrolyte, present in the cell.

The motion of ions through a solution because of the electrostatic attraction between the ions and electrodes is called migration.

(a) Electrogravimetric analysis. Analyte is deposited on the large Pt gauze electrode. If analyte is to be oxidized rather than reduced, the polarity of power supply is reversed so that deposition still occurs on the large electrode. Apparatus for electrodeposition of metals without cathode-potential control. Note that this is a two-electrode cell.

(b) Outer Pt gauze electrode. (c) Opened inner Pt Pt gauze electrode designed to be spun by a mortor in place of magnetic stirring.

Electrogravimetric analysis

Electrodeposition analysis in which the quantities of metals deposited may be determined by weighing a suitable electrode before and after deposition.

Tests for completion of the deposition :

1) disappearance of color

2) deposition on freshly exposed electrode surface

3) qualitative test for analyte in solution

In practice, there may be other electroactive species that interfere by codeposition with the desired analyte.

Two general types of electrolytic procedures :

1) Electrogravimetry without potential control

2) Controlled-potential ; potentiostatic method

1) Residual current :

Initially, a small residual current is present, owing both to oxidation and reduction of impurities and to a small amount of the intended electrolysis.

2) At a sufficient negative voltage, the desired electrolysis is the main reaction. The voltage is shifted from –Eeq by the overpotential for oxygen formation. It requires about 1V of extra potential to overcome the barrier to O2 formation at the anode.

3) At more negative voltages, the current is linear with respect to voltage, according to Ohm’s law. The slope tells us the resistance of the cell.

4) Electrolysis of solvent :

At –4.6V, reduction of water to H2 begins, and the current shoots upward.

Observed current-voltage relationship for electrolysis of 0.2 M CuSO4 in 1M HClO4 under N2.

Current-voltage relationship for the electrolytic techniques.

A = C / s I = C/ t

equivalents deposited = coulombs / 96,487

= It / 96,487

equivalents = grams / equivalent wt.

= (grams)(n) / formula wt.

(grams)(n) / formula wt. = It / 96,487

t = 96,487(grams (n) / I (formula wt.)

From this equation, the time required to deposit a given weight of metal

can be estimated. Determine current and time in electrolytic techniques

Example

If a current of 0.17 A flows for 16 minutes through the cell in the Figure how many grams of Cu will have been deposited ?

n = I · t / F

= (0.17 C/s) (16 min 60 sec/min) / (96,485C/mol)

= 1.69× 10–3 mol

Cu 2 e

63.546 g = 1 mol

x g = (1.69× 10–3 mol) / 2 x = 0.054 g

Example : electrogravimetric deposition of cobalt.

A solution containing 0.40249 g of CoCl2·xH2O was exhaustively electrolyzed to deposit 0.09937 g of metallic cobalt on a platinum cathode.

Co2+ + 2e Co(s)

Calculate the number of moles of water per mole of cobalt in the reagent.

Solution : Co a.w. = 58.9332 deposited wt. = 0.09937 g = x moles

CoCl2 m.w. = 129.8391 sample wt = y g

x = 0.001686 moles y = 0.21893 g

H2O = CoCl2·xH2O – CoCl2 = 0.40249– 0.21893 = 0.18356 g

H2O m.w. = 18.01528 = 1 mole 0.18356 = z moles

z = 0.010189 moles

(moles of H2O) / ( moles of Co) = 0.010189 / 0.001686 6.043 CoCl2·6H2O

E (cathode) becomes more negative with time when electrolysis is conducted in a two-electrode cell with a constant voltage between the electrodes.

Changes in current at constant applied voltage electrolysis

In constant voltage procedure for applied potential is set high enough to lower the metal ion concentration to its desired level and low enough to prevent the evolution of hydrogen or the deposition of another metal.

In general, constant voltage electrolysis is not selective. Any solute more easily reduced than H+ will be electrolyzed.

The current decreases with time owing to depletion of copper ions in the solution as well as to an increase in cathodic concentration polarization. In fact, with the onset of concentration polarization, the current decrease becomes exponential intime, as shown in right Figure.

It = Ioe–kt

where Io is initial current and It is the current t min after the onset of polarization.

The initial current is high because the concentration of electroactive species at each electrode surface is large. Once the electrolysis begins, these species are partially deposited at the electrode surfaces and the electrodes become polarized.

Change in current with time during electrolysis.

(a) Current ; (b) IR drop and cathode potential change during electrolytic deposition of copper at a constant applied cell potential.

The current and IR drop decrease steadily with time.

The cathode potential shifts negatively to offset the decrease in IR drop.

At point B, the cathode becomes depolarized by the reduction of hydrogen ions. Metals that deposit at point A or D interfere with copper because of co-deposition. A metal that deposits a point C does not interfere.

Changes in voltage at constant current

The current and voltage of an electrolysis cell cannot be kept constant simultaneously. Constant current electrolysis is the least selective mode of electrolysis

Controlled potential electrolysis

A three electrode cell can be used to maintain a constant cathode potential and thereby greatly increase the selectivity of the electrolysis.

Working electrode : where the analytical reaction occurs.

Auxiliary electrode : the other electrode needed for current flow

Reference electrode (SCE) : the third electrode, used to measure the potential of the

working electrode.

In controlled potential electrolysis, there is a constant voltage between the working and reference electrodes. The voltage between the working and auxiliary electrodes dose change. In constant voltage electrolysis there is a constant voltage between the working and auxiliary electrodes.

Controlled-potential electrogravimetry

The electrolysis current passes between the working electrode and a counter electrode. The counter electrode has no effect on the reaction at the working electrode.

Reduction reactions occur at working electrode potentials ( measured with respect to the reference electrode ) that are more negative than that required to start the reaction. Oxidations occur when the working electrode is more positive than necessary to start the reaction.

Controlled potential means that a constant potential difference is maintained between the working and reference electrodes.

Constant voltage means that a constant potential difference is maintained between the working and auxiliary electrodes.

Controlled potential affords high selectivity, but the procedure is slower than constant voltage electrolysis.

A potentiostat maintains the working electrode potential at a constant value relative to a reference electrode.

Circuit used for controlled-potential electrolysis with a three-electrode cell

Charges in cell potential (A) and current (B) during a controlled-potential deposition of copper. The cathode is maintained at –0.36V (vs. Lingane, Anal. Chem.. Acta, 1948, 2, 590.)

A mercury cathode for the electrolytic removal of metal ions from solution.

Apparatus for controlled-potential electrolysis. The digital voltmeter monitors the potential between the working and the reference electrode. The voltage applied between the working and the counter electrode is varied by adjusting contact C on the potentiometer to maintain the working electrode (cathode in this example) at a constant potential versus a reference electrode. The current in the reference electrode is essentially zero at all times. Modern potentiostats are fully automatic and often computer controlled. The electrode symbols shown are the currently accepted notation.

Coulometry (charge measurement)

Coulometric procedures are concerned with the quantity of electricity which flows in any given electrochemical cell and the relationship between this quantity and the amount of the reactants and/or products.

Instead of weighing the substance plated on the electrode, coulometry is based on measuring the number of electrons that participate in a chemical reaction.

Coulometry is more versatile than electrodeposition, because they include both electrochemical reactions in which a gas is formed and those in which both the reactant and the product are soluble species.

Coulometry is based on Faraday’s law, which states that one faraday of electricity will react with one equivalent weight of a reactant and will yield one equivalent weight of a product. One coulomb is the quantity of electricity transported in 1 second by a constant current of 1 ampere. If a constant current of I ampres flows for t seconds, the number of coulombs q is given by the expression q = It.

If the current is not constant with time such as controlled potential coulometry, the quantity of electricity is more difficult to determine, requiring integration of the current with respect to time q = o

t I(t)dt.

English bookbinder who became interested in electricity. He obtained an assistantship in Davy's lab, then began to conduct his own experiments. He wrote a review article on current views about electricity and magnetism in 1821, for which he reproduced Oersted's experiment. He was one of the greatest experimenters ever. Because he was self trained, however, he had no grasp of mathematics and could therefore not understand a word of Ampère's papers. In the course of his experiments, Faraday discovered that a suspended magnet would revolve around a current bearing wire, leading him to propose that magnetism was a circular force. He also discovered magnetic optical rotation, invented the dynamo (a device capable of converting electricity to motion) in 1821, discovered electromagnetic induction in 1831, and devised the laws of chemical electrodeposition of metals from solutions in 1857.

He formulated the second law of electrolysis: "the amounts of bodies which are equivalent to each other in their ordinary chemical action have equal quantities of electricity naturally associated with them." He published many of his results in the three-volume Experimental Researches in Electricity (1839-1855). One of his most important contributions to physics was his development of the concept of a field to describe magnetic and electric forces in 1845. He first suggested that current produces a electric "tension" which produced an "electrotonic state," or polarization of matter molecules, and was responsible for transmitting the electric force. He experimented with dielectrics in a capacitor. After further experimentation, he abandoned the concept of electrotonic forces in favor of "lines of force." He maintained that these lines could be made visible in a magnet using iron filings. Faraday was an advocate of the law of conservation of energy, believing that possibility of "the production of any one [power] from another, or the conversion of into another."

Faraday, Michael

(1791-1867)

Faraday’s law

The current or charge passed in a redox reaction is proportional to the moles of the reaction’s reactants and products.

The coulomb (C) is the amount of charge required to produce 0.00111800 g of silver metal from silver ions.

A coulomb is equivalent to an A•s ;

thus for a constant current, I, the charge, q, is given as

q = I · t

coulombs = amperes · seconds

The charge on an electron is defined as 1.6022 × 10–19 coulombs.

Total charge, q, in coulombs, passed during an electrolysis is related to the amount of analyte by Faraday’s law

q = n · F

where F = 96,485.3415 C/mol

n = q / F = I · t / F

The two ways to do coulometry

(a) constant current coulometry

(b) controlled potential coulometry

‑ both a little tricky, but (b) less tricky

Whichever technique is used, the governing equation for calculating the amount of analyte present is:

meq analyte = meq charge generated

One equivalent of chemical change is the change brought about by 1 mol of electrons.

Type of coulometry

1) Constant current coulometry : controlled-currnet coulometry:

coulometric titration

If we know the current, it is only necessary to measure the time needed for complete reaction in order to count the coulombs :

q = It.

Electrons are the reagent in a coulometric titration.

2) Controlled potential coulometry : Potentiostatic coulometry

The initial current is high, but decreases exponentialy as the analyte concentration decreases. Since the current is not constant, the coulombs must be measured by integrating the current over the time of the reaction :

q = ot I(t)dt.

I

Timet= t0t= 0

Current-time curve for contolled potential coulometry

Time

I

q = It

Current-time curve for contolled current coulometry

Electrolysis cells for potentiostatic coulometry. Working electrode: (a)platinum gauze, (b)mercurypool. (Reprinted with permission from J. E. Harrar and C. L Pomernacki, Anal. Cham., 1973,45,57. Copyright 1973 American Chemical Society.)

q(C) = F (C/eq) × n(eq/mol) × Cx(mol/L) × Vx(L)

normally, we do not measure charge directly, but rather current (the rate of charge flow)

I [amps(=C/sec)] = q(C)/t(sec)

thus

q(C) = I(amps) × t(sec)

or, if the current is not constant

q = ot I(t)dt.

For a current that varies with time, the quantity of charge Q in a time t is the shared area under the curve, obtained by integration of the current-time curve.

Standard Calomel Reference Electrode for general purpose applications CRR11 Reference Electrode

Temperature Range 0-60oC

Junction Ceramic Frit

Reference Hg/Hg2Cl2 Reference Electrolyte Saturated

KCl

Length 103mm Diameter 7mm

Features 1m of cable terminating in a BNC

100ml Jacketed Cell Assembly

Available Replacements

Cell Body 124-100ml-body Cell Top 124-100ml-top

Reference Electrodes

Silver/Silver Chloride CRR11/Ag/DRG1726

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Working Electrodes Platinum UMMPTB11/10/19/A=60 Gold UMMAUR11/10/19/A=60

http://www.sycopel.com/electrodes/crr11.html

Selecting a constant potential in controlled- potential coulometry

Cu2+ (aq) + 2e = Cu(s)

More (+)

More (–)

O2

H2O

EoO2/H2O = 1.229 V

Cu 2+

Cu

EoCu2+/Cu = 0.342 V

H3O+

H2

EoH3O+/H2 = 0.000 V

E

Ladder diagram for aqueous solutions of Cu2+

Application of controlled-potential coulometry

To dtermine more than 55 elements in inorganic compounds.

The electrolytic determination (and synthesis) of organic compounds.

Ex. Trichloroacetic acid, and picric acid are quantitatively reduced at mercury cathode whose potential is suitably controlled.

Picric acid(2,4,6-trinitrophenol) is an explosive compound, yellow dye, antiseptics.

Coulometric titration of Fe(II)

At a Pt anode: Fe2+ Fe3+ + e

Auxiliary reagent : Ce3+ Ce4+ + e

cf. Increase in anode potential cause the decomposition of water

H2O O2 + 4H+ + 4e

Ce4+ + Fe2+ Ce3+ + Fe3+

Redox indicator : 1,10-phenathroline

Example

2Br– Br2 + 2e

Br

Br2 + Br

Br2 is generated by the Pt anode.

Cyclohexene 82.146 g = 1 mole

0.6113 mg = x moles x = 0.01488 mmole

For 0.01488 mmol of cyclohexene to react, 0.02976 mmol of electrons must flow.

q = nF = It t = nF / I

t = (0.02976×10-3 mol) ( 96485 C/mol) / (4.825 ×10-3 C/s) = 595.1 s

Apparatus for coulometric titration of cyclihexane with Br2.

Conceptual diagram of a coulometric titration apparatus. Commercial coulometric titrators are totally electronic and usually computer controlled.

Potentiostat

A potentiostat is an electronic device that controls the voltage difference between a working electrode and a reference electrode. Both electrodes are contained in an electrochemical cell. The potentiostat implements this control by injecting current into the cell through an auxiliary (counter), or counter, electrode.

In almost all applications, the potentiostat measures the current flow between the working and auxiliary electrodes. The controlled variable in a potentiostat is the cell potential and the measured variable is the cell current.

Galvanostat

An electronic instrument that controls the current ( from one to several hundred mA) through an electrochemical cell at a preset value, as long as the needed cell voltage and current do not exceed the compliance limits of the galvanostat. Also called "amperostat."

EG&G PAR 314 Potentiostat / Galvanostat Multiplexer

A typical coulometric titration cell

A cell for the external coulometric generation of acid and base.

Selected applications of coulometric titrations

A commercial digital chloridometer. This coulometric titrator is designed to determine chloride ion in such clinical samples as serum, urine, and sweet. It is used in the diagnosis of cystic fibrosis. The chloridometeris also used in food and environmental laboratories. (Courtesy of Labconco Crop., Kansas City, MO.)

If the volumes of the standard solution and the unknown solution are the same, concentrations can be substituted for number of moles in this equation. A commercial coulometric titrator called a chloridometer is shown in Figure 22F-2.

Other popular methods to determine chloride are ion-selective electrodes (see section 21D), photometric titrations (see Section 26A-4), and isotope dilution mass spectrometry.

Q& A

Thanks

Dong-Sun Lee / CAT / SWU