Transcript Potentiometry

Potentiometry
Dr Hisham E Abdellatef
2010
• It is a method of
analysis in which we
determine the
concentration of an ion
or substance by
measuring the
potential developed
when a sensitive
electrode is immersed
in the solution of the
species to be
determined.
Mo = Mn++ ne
Applying Nernest equation.
Determination of the substances by potentiometric
technique can be carried out by two ways:
1. Direct potentiometry and
2. potentiometric titrations
• The potential of the indicator
electrode cannot be measured alone;
•
1.
2.
3.
4.
For any potentiometric measurement
we must have:
Reference electrode
Indicator electrode.
Potentiometer
Salt bridge to connect the two
electrode solutions and complete the
circuit.
A- Reference electrode
Reference electrode must:
1. Have a constant potential
2. Its potential must be definite
To express any electrode we have to mention:
1. Redox reaction at the electrode surface.
2. Half cell and Nernst equation.
3. Sketch of its design.
4. Any necessary conditions for its preparation.
5. Any necessary precautions for its use.
Standard Hydrogen Electrode
It’s a primary reference electrode.
Its potential is considered to
be zero.
Electrode reaction:
half cell:
pt/ H2 , H+ (1N) 
Eo = zero
d-Limitation
1.
It is difficult to be used and
to keep H2 gas at one
atmosphere during all
determinations.
2.
It needs periodical replating
of Pt. Sheet with Pt. Black
Saturated calomel electrode (S.C.E.)
Hg | Hg2Cl2 (sat’d), KCl (sat’d) | |
electrode reaction in calomel hal-cell
Hg2Cl2 + 2e = 2Hg + 2Cl–
Eo = + 0.268V
E = Eo – (0.05916/2) log[Cl–]2 = 0.244 V
Temperature dependent
Potential of the electrode depends on the chloride ion
Hg2Cl2
⇌ 2 Hg22+ + 2Cl-
Sp Hg2Cl2 = [Hg22+]2 [Cl-]2
Ksp = 1.8 ×10–18
E = Eo – (0.0591/2) log[Cl–]2 = 0.244 V
The crystal structure of calomel(Hg2Cl2),
which has limited solubility in water
(Ksp = 1.8 ×10–18).
Hg2Cl2  Hg22+ + 2Cl–
Saturated KCl = 4.6 M KCl
Ksp = 1.8 ×10–18
KCl
E volt
Saturated
0.241
1M
0.280
0.1 M
0.334
Silver-silver chloride electrode
Ag(s) | AgCl (sat’d), KCl (xM) | |
AgCl(s) + e = Ag(s) + Cl–
Eo = +0.244V
E = Eo – (0.05916/1) log [Cl–]
E (saturated KCl) = + 0.199V (25oC)
• SpAgCl =
[Ag+] [Cl-]
E = Eo – (0.05916/1) log [Cl–]
Disadvantage of silver-silver chloride electrode
• It is more difficult to prepare than SCE.
• AgCI in the electrode has large solubility in
saturated KCl
Advantage of Ag/AgCI electrodes over SCE.
• It has better thermal stability.
• Less toxicity and environmental problems with
consequent cleanup and disposal difficulties.
B- Indicator electrode
its potential is sensitive to the concentration of analyte
Ecell=Eindicator-Ereference
It must be:
(a) give a rapid response and
(b) its response must be reproducible.
Metallic electrodes: where the redox reaction takes
place at the electrode surface.
Membrane (specific or ion selective) electrodes:
where charge exchange takes place at a specific
surfaces and as a result a potential is developed.
1. Electrodes for precipitemetry and
complexometry
a- First-order electrodes for cations:
e.g. in determination of Ag+ a rode or wire of
silver metal is the indicator electrode, it is
potential is:
It is used for determination of Ag+ with Cl-,
Br- and CN-. Copper, lead, cadmium, and
mercury
Example of First-order electrode
Ag+ + e = Ag(s)
Eo = + 0.800V
E = 0.800 – (0.05916/1) log {1/[Ag+]}
b) Second order electrodes for anions
A metal electrode can sometimes be indirectly responsive to
the concentration of an anion that forms a precipitate or
complex ion with cations of the metal.
Ex. 1. Silver electrode
The potential of a silver electrode will accurately reflect the
concentration of
iodide ion in a solution that is saturated with silver iodide.
AgI(s) + e = Ag(s) + I–
Eo = – 0.151V
E = – 0.151 – (0.05916/1) log [I–]
= – 0.151 + (0.05916/1)pI
2. Mercury electrode for measuring the
concentration of the EDTA anion Y4–.
Mercury electrode responds in the presence
of a small concentration of the
stable EDTA complex of mercury(II).
HgY2– + 2e = Hg(l) + Y4–
Eo = 0.21V
E = 0.21 – (0.05916/2) log ([Y4–] /[HgY2–])
K = 0.21 – (0.05916/2) log (1 /[HgY2–])
E = K – (0.05916/2) log [Y4–] = K +(0.05916 / 2) pY
2. Inert electrodes (Indicators electrodes for redox reaction)
Chemically inert conductors such as gold, platinum, or carbon that do
not participate, directly, in the redox process are called inert electrodes.
The potential developed at an inert electrode depends on the nature and
concentration of the various redox reagents in the solution.
Examples:
Ag(s) | AgCl[sat’d], KCl[xM] | | Fe2+,Fe3+) | Pt
Fe3++e = Fe2+
Eo = +0.770V
Ecell = Eindicator – Ereference
= {0.770 – (0.05916/1) log [Fe2+]/[Fe3+]} – {0.222 – (0.05916/1) log [Cl–]}
2) Membrane indicator electrodes
The potential developed at this type of electrode results from an unequal
charge buildup at opposing surface of a special membrane. The charge at each
surface is governed by the position of an equilibrium involving analyte ions,
which, in turn, depends on the concentration of those ions in the solution.
The electrodes are categorized according to the type of membrane they employ :
glass,
polymer,
crystalline,
gas sensor.
The first practical glass
electrode. (Haber and
Klemensiewcz, Z. Phys.
Chem, 1909, 65, 385.
3. indicator electrodes for neutralization reaction
Glass Membrane
Electrode
Composition of glass membranes
70% SiO2
30% CaO, BaO, Li2O, Na2O,
and/or Al2O3
Ion exchange process at glass
membrane-solution interface:
Gl– + H+ = H+Gl–
(a) Cross-sectional view of a silicate glass struture. In addition to the three
Si│O bonds shown, each silicon is bonded to an additional oxygen atom,
either above or below the plane of the paper. (b) Model showing threedimensional structure of amorphous silica with Na+ ion (large dark blue)
and several H+ ions small dark blue incorporated.
Glass Membrane
Electrode
E = K + 0.059 (pH1 - pH2)
K= constant known by the
asymmetry potential.
PH1 = pH of the internal solution 1.
PH2 = pH of the external solution 2.
The final equation is:
E = K - 0.059 pH
Standardization
at pH=7.00 , E = 0 V.
pH 4.00, E= 59.16 mV/pH unit
•Asymmetry potential
•E of the 2 reference electrodes
•pH of the internal solution
•Liquid junction potential
Measurement of pH (cont.)
Ecell = E°cell - (0.0591)log[H+] + constant
• Ecell is directly proportional to log [H+]
electrode
pH Meters
Glass Membrane Electrode
• Advantages of glass electrode:
It can be used in presence of oxidizing, reducing, complexing
• Disadvantage:
1. Delicate, it can’t be used in presence of dehydrating agent e.g. conc.
H2SO4, ethyl alcohol….
2. Interference from Na+ occurs above pH 12 i.e Na+ excghange
together with H+ above pH 12 and higher results are obtained.
3. It takes certain time to come to equilibrium due to resistance of
glass to electricity.
Junction potential :
a small potential that exists at the interface between two electrolyte solutions
that differ in composition.
Development of the junction potential caused by unequal mobilities of ions.
Mobilties of ions in water at 25oC:
Na+ : 5.19 × 10 –8 m2/sV K+ : 7.62 × 10 –8
Cl– : 7.91× 10 –8
To reduce the liquid junction potential to only few
millivolts one has to:
1. Use a sat for preparation of the junction which its
cation and anion have very near mobilities, so that they
move by the same rate e.g. KCl and KNO3. (K+ =74, Cl- =
73 and NO3- = 76)
2. Use high concentration of the salt for preparation of
the bridge, to reduce the effect of difference in rates
of migration of other ions in the electrode solutions.
2. Standard Hydrogen Electrode
electrode reaction:
Pt. balck
H  2e 

 12 H2
Nernst equation
1
E  zero - 0.059log 
[H ]
E = -0.059 pH
When it is connected with NHE as reference
electrode the e.m.f. of the cell :
Ecell = zero –(–0.059 pH)
= 0.059 pH
pH = E / 0.059
Disadvantages:1.
It cannot be used in solution containing oxidising agent which
will oxidiose [ ½ H2 = H+ + e ] or reducing substances which
will reduce [ H+ + e = ½ H2 ] especially in presence of platinum
black
2.
It cannot be used in reactions involving volatile constituent’s e.g.
CO2, as it will be bubbled out by the H2 gas.
3.
It cannot be used in presence of catalytic poisons which will
affect Pt black which catalyses the electrode reaction.
4.
It needs repletion with Pt black.
5.
It is not easy to keep H2 gas at one atmospheric pressure during
all measurements.
3. Antimony electrode Sb/Sb2O3
Electrode reaction:
Sb2O3 + 6 H+ + 6 e = 2 Sb + 3 H2O
Nernst equation
ESb/Sb 2O3
2
0.059
[Sb]
 Eo log
6
[Sb 2O3 ][H  ]6
2

  0.059
0.059
[
Sb
]
1 
25
0


or E  E - 
log
 
log  6 

[ Sb2O3 ]   6
[H ] 
 6
E25 = E0 – 0.059 pH
Advantages
Easy to use, cheep and
durable.
Disadvantages
1. can only be used within
pH range 2 – 8 at lower
pH Sb2O3 dissolves and
at higher pH Sbo
dissolves.
2. It cannot be used in
presence of oxidizing
agents, reducing agents,
complexing agents and
noble metals
Quinhydrone electrode
•
1.
2.
3.
Advantages
It is not affected by catalytic poisons.
Easy to prepare and use.
It comes to equilibrium rapidly.
• Disadvantages:
1. It cannot be used in presence of oxidising agents and
reducing agents
2. The upper limit of the electrode use is pH 8
3. It needs to be used freshly.
Application of potentiometry
1. Direct potentiometric measurements
Eobs = Eref + Eja - Eind
sulfate calibration curve
y = 14427x - 12024
R2 = 0.999
1400000
1200000
peak area
1000000
800000
600000
400000
200000
0
0
10
20
30
40
50
60
concentration (ppm )
by area
Linear (by area)
70
80
90
2. Potentiometric titration
It is used for all types of volumetric analysis: acid base,
precipitimetry, complexometry and redox
It is used when it is not easy or impossible to detect
the end point by ordinary visual methods i.e:
1. For highly coloured or turbid solutions.
2. For very dilute solutions 10-3, 10-6 M.
3. When there is no available indicator
Potentiometric titration
Titration of 2.433mmol of chloride ion with
0.1000M silver nitrate.
(a) Titration curve.
(b) First-derivative curve.
(c) Second-derivative curve.
Application of potentiometric titration in
a)
Neutralization reactions: glass / calomel
electrode for determination of Ph
b) Precipitation reactions: Membrane electrodes for
the determination of the halogens using silver
nitrate reagent
c) Complex formation titration: metal and membrane
electrodes for determination of many cations
(mixture of Bi3+, Cd2+ and Ca2+ using EDTA)
d) Redox titration: platinum electrode For example
for reaction of Fe3+/ Fe2+ with Ce4+/Ce3+
a) Neutralization
reactions:
glass / calomel electrode
for determination of
pH
• Precipitation
reactions:
Membrane electrodes
for the
determination of
the halogens using
silver nitrate
reagent
• Complex formation titration: metal and
E, V vs Ag/AgCl, 1M KCl
membrane electrodes for determination of many
cations (mixture of Bi3+, Cd2+ and Ca2+ using EDTA)
0.12
E/V
E
0.10
0.08
0.06
0.0
0.1
0.2
0.3
0.4
Volume, mL
0.5
0.6
• Redox titration: platinum electrode For example for
reaction of Fe3+/ Fe2+ with Ce4+/Ce3+