Transcript 분석화학 강의노트 electrogravimetry 전해무게법 분석, 전기량법 분석

Dong-Sun Lee / cat - lab / SWU
2012-Fall version
Chapter 22
Electrogravimetry and
coulometry
Secret Sunshine
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.
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
Cathode
Is where oxidation occurs
Is where reduction occurs
Is where electrons are produced
Is where electrons are consumed
Is what anions migrate toward
Is what cations migrate toward
Has a negative sign
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
Anode
Zn
e
–
V
+
Cathode
Cu
A voltaic cell in a circuit
consists of :
1) a negative electrode -anode
2) a positive electrode - cathode
3) an electrolyte
ZnSO4
Porous fritted disk
(liquid junction)
CuSO4
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
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.
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+
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.
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.
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 Cd 2+ 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.
Electrogravimetric analysis
Electrodeposition analysis in which the quantities of metals deposited may be determined
by weighing a suitable electrode before and after deposition.
(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.
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.)
Determine current and time in
electrolytic techniques
From this equation, the time required
to deposit a given weight of metal
can be estimated.
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
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.
E (cathode) becomes more
negative with time when
electrolysis is conducted in a
two-electrode cell with a
constant voltage between the
electrodes.
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 codeposition. 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 = ot 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.
Faraday, Michael
(1791-1867)
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’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
I
q = It
Time
Current-time curve for contolled current
coulometry
t= 0
t= t0
Time
Current-time curve for contolled potential
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
Calomel CRR11/10/19/A=60
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 (+)
O2
EoO2/H2O = 1.229 V
H2O
Cu 2+
E
EoCu2+/Cu =
0.342 V
Cu
H3O+
EoH3O+/H2 = 0.000 V
H2
More (–)
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
Apparatus for coulometric titration of cyclihexane with Br2.
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
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 22F2.
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