Redox titrations &
potentiometry
(Mark=3)
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Redox reactions
2Fe3+ + Sn2+  2Fe2+ + Sn4+
half-reactions:
Reduction
2Fe3+ + 2e-  2Fe2+
oxidation
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Sn2+  Sn4+ + 2e-
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Reduction Half-Reaction
E(V)
F2(g) + 2e-  2F-(aq)
2.87
Au3+(aq) + 3e-  Au(s)
1.50
Cl2(g) + 2 e-  2Cl-(aq)
1.36
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O
1.33
O2(g) + 4H+ + 4e-  2H2O(l)
1.23
Ag+(aq) + e-  Ag(s)
0.80
Fe3+(aq) + e-  Fe2+(aq)
0.77
Cu2+(aq) + 2e-  Cu(s)
0.34
Sn4+(aq) + 2e-  Sn2+(aq)
0.15
2H+(aq) + 2e-  H2(g)
0.00
Sn2+(aq) + 2e-  Sn(s)
-0.14
Ni2+(aq) + 2e-  Ni(s)
-0.23
Fe2+(aq) + 2e-  Fe(s)
-0.44
Zn2+(aq) + 2e-  Zn(s)
-0.76
Al3+(aq) + 3e-  Al(s)
-1.66
Mg2+(aq) + 2e-  Mg(s)
-2.37
Li+(aq) + e-  Li(s)
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-3.04
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Red. agent strength increases
Ox. agent strength increases
Standard Reduction Potentials (strength)
3
Balancing of redox reactions.
Under Acidic conditions
1. Identify oxidized and reduced species
Write the half reaction for each.
2. Balance the half rxn separately except H & O’s.
Balance: Oxygen by H2O
Balance: Hydrogen by H+
Balance: Charge by e 3. Multiply each half reaction by a coefficient.
There should be the same # of e- in both half-rxn.
4. Add the half-rxn together, the e - should cancel.
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Balancing of redox reactions.
Under Basic conditions
1. Identify oxidized and reduced species
Write the half reaction for each.
2. Balance the half rxn separately except H & O’s.
Balance: Oxygen by H2O
Balance: Hydrogen by OHBalance: Charge by e 3. Multiply each half reaction by a coefficient.
There should be the same # of e- in both half-rxn.
4. Add the half-rxn together, the e - should cancel.
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Balancing of redox reactions
H2O2 (aq) + Cr2O7-2(aq )  Cr 3+ (aq) + O2 (g) Redox reaction
======================================
1)write 2 half reactions
Half Rxn (red):
Cr2O7-2 (aq)  Cr3+
Half Rxn (oxid):
H2O2 (aq)  O2
2)Atom balance
Cr2O7-2 (aq)  2Cr3+
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H
2O 2
(aq)
 O2
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Balancing of redox reactions
3)Oxygen balance
Half Rxn (red): Cr2O7-2 (aq)  2Cr3+ + 7 H2O
Half Rxn (oxi):
H2O2
(aq)
 O2
4)Hydrogen balance
Half Rxn (red): 14H+ + Cr2O7-2 (aq)  2Cr3+ + 7 H2O
Half Rxn (oxi):
H2O2
(aq)
 O2 + 2H+
5)Electron balance
6e- + 14H+ + Cr2O7-2 (aq)  2Cr3+ + 7 H2O
H2O2
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Balancing of redox reactions
6) Equalize of produced and
consumed electrons
6e- + 14H+ + Cr2O7-2 (aq)  2Cr3+ + 7 H2O
( H2O2
(aq)
 O2 + 2H+ + 2e- ) x 3
7)Multiply each half reaction
8 H+ + 3H2O2 + Cr2O72-  2Cr+3 + 3O2 + 7H2O
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Balance the redox reactions
H+
I2 +S2O32- ⇋ I- +S4O62-
OH-
I2 +S2O32- ⇋ I- +SO42-
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Preadjustment of analyte oxidation state
It is necessary to adjust the oxidation state of the analyte to one that can be titrated
with an auxiliary oxidizing or reducing agent.
Ex.
Preadjustment by auxiliary reagent
Fe(II), Fe(III)
–
4
Fe(II)
Titration
Ce4+
Preoxidation :
2–
Peroxydisulfate ( (NH4)2S)2O8 )
Sodium bismuthate ( NaBiO 3)
Hydrogen peroxide (H2O2)
Prereduction : Stannous chloride ( SnCl2)
Chromous chloride
Jones reductor (zinc coated with zinc amalgam)
Walden reductor ( solid Ag and 1M HCl)
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Redox titrations
2Fe3+ + Sn2+  2Fe2+ + Sn4+
2) In electrochemical cell.
(Potentiometry)
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1) In solution
(visual indicators)
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Iodimetry and iodometry
• Iodimetry:
• a reducing analyte is titrated directly with iodine (to
produce I−).
•
• iodometry :
• an oxidizing analyte is added to excess I− to produce
iodine, which is then titrated with standard thiosulfate
solution.
I- + Cu2+→ I2 + Cu+
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I2 + S2O32- → 2I- + S4O62-
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Preparation of aqoues solution I31) Iodine only dissolves slightly in water. Its solubility is
enhanced by interacting with I-
Standardization of Iodin with Arsenious oxide,
As2O3 + 3H2O = 2H3AsO3
As4O6 + 6H2O = 4H3AsO3
H3AsO3 + I3– + H2O = H3AsO4 + 3I– + 2H+
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standard I31) An excellent way to prepare standard I3- is to add a weighed
quantity of potassium iodate to a small excess of KI. Then
add excess strong acid (giving pH ≈ 1) to produce I3- by
quantitative reverse disproportionation:
2)
Cu + HNO3  Cu2+
Cu2++4I- 2CUI + I2
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Stability of I2 Solutions
• In acidic solutions of I3- are unstable because the
excess I− is slowly oxidized by air:
• In neutral solutions, oxidation is insignificant in the
absence of heat, light, and metal ions.
• At pH ≳ 11, triiodide disproportionates to hypoiodous
acid (HOI), iodate, and iodide.
I2 + OH- ⇌ IO- + I- + H+
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3IO- ⇌ IO3- + 2I-
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Iodimetry
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iodometry
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Sodium thiosulfate, Na2S2O3
Thiosulfate ion is a moderate reducing agent that has been widely used
to determine oxidizing agents by an indirect procedure that involves
iodine as an intermediate. With iodine, thiosulfate ion is oxidized
quantitatively to tetrathionate ion according to the half-reaction:
2S2O3 2–  S4O6 2– + 2e
Eo = 0.08
Ex. Determination of hypochlorite in bleaches [CaCl(OCl)H2O]:
OCl– + 2I– + 2H+  Cl– + I2 + H2O (unmeasured excess KI)
I2 + 2 S2O3 2–  2I– + S4O6 2–
Indicator: soluble starch (-amylose)
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Standardization of thiosulfate solution:
Primary standard : potassium iodate (KIO3), K2Cr2O7, KBrO3
Titration reactions:
KIO3 + 5KI + 6HCl  3I2 + 6KCl + 3 H2O
I2 + 2Na2S2O3  2NaI + Na2S4O6
KIO3

3I2

6Na2S2O3·5H2O
 6 Equivalent
S2O32- +H+ ⇋ HSO3- +S(s)
pH, Microorganisms, Concentration, Cu2+, Sunlight
Stabilizer for sodium thiosulfate solution : Na2CO3
Na2S2O3 + H2O + CO2  Na2CO3 + H2S2O3
H2S2O3  H2SO3 + S
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16-2 Finding the end point
A redox indicator is a
compound
that changes color
when it goes from its oxidized
to its reduced state.
or
For ferroin, with E° = 1.147 V
we expect the color change to occur in the approximate range
1.088 V to 1.206 V with respect SHE
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Starch-Iodine Complex
Starch is the indicator of choice for those procedures
involving iodine because it forms an intense blue colour
with iodine.
Starch is not a redox indicator;
it responds specifically to the presence of I2,
not to a change in redox potential.
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Structure of the repeating
unit of the sugar amylose.
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Permanganate titration  
Oxidation with permanganate :
(Reduction of permanaganate)
KMnO4
Powerful oxidant that the most widely used.
1) In strongly acidic solutions (1M H2SO4 or HCl, pH  1)
MnO4– + 8H+ + 5e = Mn2 + + 4H2 O
Eo = 1.51 V
KMnO4 is a self-indicator.
2) In feebly acidic, neutral, or alkaline solutions
MnO4– + 4H+ + 3e = MnO2 (s) + 2H2 O
Eo = 1.695 V
3) In very strongly alkaline solution (2M NaOH)
MnO4– + e = MnO42 –
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Eo = 0.558 V
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Permanganate titration  
Duration of colour in end point (30 seconds)
MnO4– + 3Mn2+ + 2H2O  5MnO2 + 4H+
K=1*1047
Stability of aqoues solution of MnO4MnO4– + 2H2O  4MnO2 (s) + 3O2 (g) +4OH-
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Standardization of KMnO4 solution
Potassium permanganate is not primary standard, because traces of MnO2
are invariably present.
Standardization by titration of sodium oxalate (primary standard) :
2KMnO4 + 5 Na2(COO)2 + 8H2SO4 = 2MnSO4 + K2SO4 + 5Na2SO4 + 10 CO2 + 8H2O
2KMnO4
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
5 Na2(COO)2
 10 Equivalent
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Preparation of 0.1 N potassium permanganate solution
KMnO4 is not pure.
Distilled water contains traces of organic reducing substances
which react slowly with permanganate to form hydrous managnese dioxide.
Manganesse dioxide promotes the autodecomposition of permanganate.
1)
Dissolve about 3.2 g of KMnO4 (mw=158.04) in 1000ml of water,
heat the solution to boiling, and keep slightly below the boiling point for 1 hr.
Alternatively , allow the solution to stand at room temperature for 2 or 3 days.
2)
Filter the liquid through a sintered-glass filter crucible to remove solid MnO2.
3)
Transfer the filtrate to a clean stoppered bottle freed from grease with cleaning
mixture and standardize it.
4)
Protect the solution from evaporation, dust, and reducing vapors, and keep it in
the dark or in diffuse light.
4)
If in time managanese dioxide settles out, refilter the solution and restandardize it.
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Applications of permanganometry
(1) H2O2
2KMnO4 + 5 H2O2 + 3H2SO4 = 2MnSO4 + K2SO4 + 5O2 + 8H2O
(2) NaNO2
2NaNO2 + H2SO4 =
Na2SO4
+ HNO2
2KMnO4 + 5 HNO2 + 3H2SO4 = 2MnSO4 + K2SO4 + 5HNO3 + 3H2O
(3) FeSO4
2KMnO4 + 510 FeSO4 + 8H2SO4 = 2MnSO4 + K2SO4 + 5Fe2(SO4)3 + 8H2O
(4) CaO
CaO + 2HCl = CaCl2 + H2O
CaCl2 + H2C2O4 = CaC2O4 + 2HCl
(excess oxalic acid)
2KMnO4 + 5 H2C2O4 + 3H2SO4 = 2MnSO4 + K2SO4 + 10CO2 + 8H2O (back tit)
(5) Calcium gluconate
[CH2OH(CHOH)4COO]2Ca + 2HCl = CaCl2 + 2CH2OH9CHOH)4COOH
(NH4)2C2O4 + CaCl2 = CaC2O4 + 2 NH4Cl
CaCl2
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+ H2SO4 =
2KMnO4 + 5 H2C2O4
H2C2O4 + CaSO4
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+ 3H2SO4 = 2MnSO
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Bromatimetry
BrO3– + 5Br– + 6H+ 
3Br2 + H2O
2I– + Br2  I2 + 2Br–
I2 + 2 S2O32–  2I– + S4O62–
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Determining water with the Karl Fisher Reagent
The Karl Fisher reaction :
I2 + SO2 + 2H2O  2HI + H2SO4
For the determination of small amount of water, Karl Fischer(1935) proposed
a reagent prepared as an anhydrous methanolic solution containing iodine,
sulfur dioxide and anhydrous pyridine in the mole ratio 1:3:10. The reaction
with water involves the following reactions :
C5H5N•I2 + C5H5N•SO2 + C5H5N + H2O  2 C5H5N•HI + C5H5N•SO3
C5H5N+•SO3– + CH3OH  C5H5N(H)SO4CH3
Pyridinium sulfite can also consume water.
C5H5N+•SO3– + H2O  C5H5NH+SO4H–
It is always advisable to use fresh reagent because of the presence of
various side reactions involving iodine. The reagent is stored in a desiccantprotected container.
The end point can be detected either by visual( at the end point, the color
changes from dark brown to yellow) or electrometric, or photometric
(absorbance at 700nm) titration
methods. The 81
detection of water by the32
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coulometric technique with Karl Fischer
reagent is popular.
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Potentiometric Methods
A.) Introduction:
1.) Potentiometric Methods: based on measurements of the potential of electrochemical
cells in the absence of appreciable currents (I →0)
2.) Basic Components:
a) reference electrode: gives reference for potential measurement
b) indicator electrode: where species of interest is measured
c) potential measuring device
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Electrodes and Potentiometry


Potential change only dependent on one ½ cell concentrations
Reference electrode is fixed or saturated  doesn’t change!
Ecell=Ecathod-Eanod
Anod is conventionally reference electode
0.059
oxi
E cathod  E cathod –
Log
n
red
2


0.059
oxi
E anod  E anod –
Log
n
red
]  




0
.
222

0
.
05916
log[
Cl
]

 [ Fe
0.05916

E cell  0.771 
log 
 [ Fe 3  ]  
1



Fe3+ +e- Fe2+
AgClReference
Ag + Cl(s) + e →
electrode,
Potential of the cell
only depends on [Fe2+]
& [Fe3+]
[Cl-] is constant
Unknown solution of
[Fe2+] & [Fe3+]
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Pt wire is indicator
electrode whose
potential responds
to [Fe2+]/[Fe3+] 34
Reference Electrodes: (Instead of SHE)
Need one electrode of system to act as a reference against which potential
measurements can be made  relative comparison.

Standard hydrogen electrodes are cumbersome
-
Requires H2 gas and freshly prepared Pt surface
Desired Characteristics:
a) known or fixed potential
b) constant response
c) insensitive to composition of solution under study
d) obeys Nernest Equation
e) reversible
E  0.222  0.05916 log[Cl  ]

>
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
>
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Reference Electrodes
1.) Silver-Silver Chloride Reference Electrode
Eo = +0.222 V
Activity of Cl- not 1E(sat,KCl) = +0.197 V
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Reference Electrodes
2.) Saturated Calomel Reference Electrode (S.C.E)

Hg 2Cl2 ( s )  2e 
2Hg (l )  2Cl 

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Eo = +0.268 V
Activity of Cl- not 1E(sat,KCl) = +0.241 V
Saturated KCl maintains constant [Cl-] even with
some evaporation
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Indicator Electrodes
2 Broad Classes of Indicator Electrodes

1) Metal indicator Electrodes
-
Develop an electric potential in response to a redox reaction at the metal surface
2) Membrane Indicator Electrodes

a) Ion-selective Electrodes
-
Selectively bind one type of ion to a membrane to generate an electric potential
b) Molecular Selective Electrode
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1) Metallic Indicator Electrode (3 Main Types)
a) Metallic Electrodes of the First Kind
b) Metallic Electrodes of the Second Kind
c) Metallic Redox Indicators
a) Metallic Electrodes of the First Kind
i. Involves single reaction
ii. Detection of cathione derived from the metal used in the electrode
iii. Example: use of copper electrode to detect Cu2+ in solution
½ reaction: Cu2+ + 2eEind gives direct measure of Cu2+:
since aCu(s) = 1:
or using pCu = -log aCu2+:
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
Cu (s)
Eind = EoCu – (0.0592/2) log aCu(s)/aCu2+
Eind = EoCu – (0.0592/2) log 1/aCu2+
Eind = EoCu – (0.0592/2) pCu
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b) Metallic Electrodes of the Second Kind
i. Detection of
anion
derived from the interaction with metal ion
(Mn+) from the electrode
ii. Anion forms precipitate or stable complex with metal ion (Mn+)
iii. Example: Detection of Cl- with Ag electrode
½ reaction: AgCl(s) + e- 
Eind gives direct measure of Cl-:
Ag(s) + Cl- EO = 0.222 V
Eind = Eo – (0.0592/1) log aAg(s) aCl-/aAgCl(s)
since aAg(s) and aAgCl(s)= 1
& Eo = 0.222 V:
Eind = 0.222 – (0.0592/1) log aCl-
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c) Metallic Redox Indicators
i. Electrodes made from inert metals (Pt, Au, Pd)
ii. Used to detect oxidation/reduction in solution
iii. Electrode acts as e- source/sink
iv. Example: Detection of Ce3+ with Pt electrode
½ reaction: Ce4+ + eEind responds to Ce4+:
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Ce3+
Eind = Eo – (0.0592/1) log aCe3+/aCe4+
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2) Membrane Indicator Electrodes
i. electrodes based on determination of Molecules,cations or anions
by the selective adsorption of these species to a membrane surface.
ii. Ion Selective Electrodes (ISE) or pIon Electrodes are more common.
iii. Desired properties of ISE’s
1) minimal solubility – membrane will not dissolve in solution during measurement.
– silica, polymers, low solubility inorganic compounds , (AgX) can be used
2) Need some electrical conductivity
3) Selectively binds ion of interest
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Ion-Selective Electrodes
Does not involve a redox process
Responds Selectively to one ion
such as (C+)
Contains a thin membrane capable of
only binding the desired ion



electric potential is generated
by a separation of charge
E  constant 
0.05916
log[C  ]
n
ISE + IRE = Combined electrod
Constant C+ - Inner solution
E2
Constant C+ - membrane
E1
C+ - Un known solution
E=E1-E2
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pH meter // ISE = (glass membrane ) that preferentially binds H
+
1) Combined (glass) electrod === ISE+ Ag/AgCl electrode
2 Electrodes
2) SCE outside electrod
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pH Electrode
pH Electrodes
1.) pH Measurement with a Glass Electrode
Ag(s)|AgCl(s)|Cl-(aq) || H+(aq,outside) H+(aq,inside) , Cl-(aq)|AgCl(s)|Ag(s)
Outer reference
electrode
Eref1
[H+] outside
(analyte solution)
Ej
E1
[H+] inside
Inner reference
electrode
E2
Eref2
Boundary potential
Glassdifference
membrane (Eb) = E1 - E2
Selectively binds H+
Eb = c + 0.059 log[H+]
Eb = c – 0.059 pH
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generated by [H+] difference
across
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glass membrane
Alkali Error
H+ not only cation that can bind to glass surface
- H+ generally has the strongest binding
Get weak binding of Na+, K+, etc
Most significant when [H+] or aH+ is low (high pH)
- usually pH > 11-12
At low aH+ (high pH), amount of Na+ or
K+ binding is significant  increases
the “apparent” amount of bound H+
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Acid Error
Errors at low pH (Acid error) can give readings that are too high
Exact cause not known
- usually occurs at pH > 0.5
Glass Electrodes for Other Cations
change composition of glass membrane
putting Al2O3 or B2O3 in glass
enhances binding for ions other than H+
Used to make ISE’s for Na+, Li+, NH4+
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Redox titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M (1M H2SO4)
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
5ml Before Eq
• Fe3+ + e Fe2+
E0=0.771
[Fe 2 ]
• E  0.771 - 0.059  Log
[Fe3 ]
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  5  2.5

 0.357 N  0.357 M
V1 V 2
105
[ Fe3 ] 
N 2V 2 5  2.5

 0.119 N  0.119M
V1  V 2
105
0.357
• E  0.771 - 0.059  Log
 0.743V
0.119
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Attention 1!!!!!!
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
5ml Before Eq
• Fe2+  Fe3+ + e E0=-0.771
[Fe3 ]
• E  0.771 - 0.059  Log
[Fe 2 ]
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  5  2.5

 0.357 N  0.357 M
V1 V 2
105
[ Fe3 ] 
N 2V 2 5  2.5

 0.119 N  0.119M
V1  V 2
105
0.119
• E  0.771 - 0.059  Log
 0.743V !!!!!!!!!!
0.357
920316
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59
Attention 2!!!!!!
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
5ml Before Eq
• 5Fe2+  5Fe3+ + 5e E0=-0.771
0.059
[Fe3 ]5
• E  0.771  Log
5
[Fe 2 ]5
920316
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  5  2.5

 0.357 N  0.357 M
V1 V 2
105
[ Fe3 ] 
N 2V 2 5  2.5

 0.119 N  0.119M
V1  V 2
105
5
0.059
(0.119)
• E  0.771  Log
 0.743V !!!!!!!!
5
5
(0.357)
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60
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
10 ml
• Fe3+ + e Fe2+
Before Eq
E0=0.771
[Fe 2 ]
• E  0.771 - 0.059  Log
[Fe3 ]
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  10  2.5

 0.227 N  0.045M
V1  V 2
110
[ Fe3 ] 
N 2V 2 10  2.5

 0.227 N  0.045M
V1  V 2
105
0.045
• E  0.771 - 0.059  Log
 0.771V
0.045
920316
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61
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• Fe3+ + e Fe2+
15 ml
Before Eq
E0=0.771
[Fe 2 ]
• E  0.771 - 0.059  Log
[Fe3 ]
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  15  2.5

 0.109 N  0.109 M
V1  V 2
115
[ Fe3 ] 
N 2V 2 15  2.5

 0.326 N  0.326M
V1  V 2
115
0.109
• E  0.771 - 0.059  Log
 0.799V
0.326
920316
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62
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
19 ml
• Fe3+ + e Fe2+
Before Eq
E0=0.771
[Fe 2 ]
• E  0.771 - 0.059  Log
[Fe3 ]
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  19  2.5

 0.021N  0.021M
V1 V 2
119
[ Fe3 ] 
N 2V 2 19  2.5

 0.328 N  0.399M
V1  V 2
119
0.021
• E  0.771 - 0.059  Log
 0.846V
0.399
920316
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63
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• Fe3+ + e Fe2+
19.5 ml Before Eq
E0=0.771
[Fe 2 ]
• E  0.771 - 0.059  Log
[Fe3 ]
[ Fe2 ] 
N1V 1  N 2V 2 100  0.5  19.5  2.5

 0.010 N  0.010M
V1  V 2
119.5
[ Fe3 ] 
N 2V 2 19.5  2.5

 0.408 N  0.408M
V1  V 2
119.5
0.010
• E  0.771 - 0.059  Log
 0.866V
0.408
920316
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64
Titration curve
• 1e+Fe3+  Fe2+
5e+MnO4-+8H+  Mn2+
• E Fe  EMnO4  E eq
-
• E eq  E Fe  E  Fe
20 ml At Eq
[Fe 2 ]
- 0.059  Log
[Fe 3 ]
• E eq  E MnO4-  E  MnO 4 
0.059
[Mn 2 ]
 Log
5
[MnO 4 ][H  ]8
×5
2
[Mn
]


• 5E eq  5E MnO 4 - 0.059  Log
[MnO 4 ][H  ]8
2
2
[Fe
][Mn
]



• 6E eq  E Fe  5E MnO 4 - 0.059  Log
[Fe3 ][MnO 4 ][H  ]8
920316
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65
Titration curve
5X
• 6E eq  E  Fe  5E MnO 4 
Y
[Fe 2 ][Mn 2 ]
- 0.059  Log
[Fe3 ][MnO 4 ][H  ]8
5Y
X
MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• 6E eq  E  Fe  5E MnO4  - 0.059  Log
E eq
920316
1
[H  ]8
/6
E  Fe  5E  MnO4  0.059
1

 Log  8
6
6
[H ]
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66
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
20 ml At Eq
• Fe3+  Fe2+
• Mno4-+8H+  Mn2+
nE 0 ox  mE 0 red
0.059
1
•E 

Log
(m  n)
( m  n)
[H  ]x
5 1.51  1 0.771

 1.387V
5 1
920316
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67
Keq
MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• 1e+Fe3+  Fe2+
E Fe  E

Fe
0.059
[Fe 2 ]5
 Log
5
[Fe 3 ]5
E MnO4-  E  MnO 4 
E  Fe
E

MnO 4 
5e+MnO4-+8H+  Mn2+
0.059
[Mn 2 ]
 Log
5
[MnO 4 ][H  ]8
20 ml At Eq
EFe  EMnO4
0.059
[Fe 2 ]5
0.059
[Mn 2 ]

 Log
 E MnO  Log
3 5
5
5
[Fe ]
[MnO 4 ][H  ]8

4
E
920316

Fe
0.059
[Mn 2 ]
0.059
[Fe 2 ]5

 Log

 Log
 8
5
5
[Fe 3 ]5
[MnO 4 ][H ]
http:\asadipour.kmu.ac.ir 81
slides
68
-
Keq
At Eq
Fe
0.059
[Mn 2 ]
0.059
[Fe 2 ]5

 Log

 Log
 8
5
5
[Fe 3 ]5
[MnO 4 ][H ]
E  MnO 4   E  Fe
0.059
[Mn 2 ]
0.059
[Fe 3 ]5

 Log

 Log
 8
5
5
[Fe 2 ]5
[MnO 4 ][H ]
E

MnO 4 
E

E  redox
0.059
[Mn 2 ]
[Fe 3 ]5

 (Log
 Log
)
2 5
 8
5
[Fe ]
[MnO 4 ][H ]
E  redox
920316
0.059
[Mn 2 ][Fe 3 ]5

 Log
5
[MnO 4 ][H  ]8 [Fe 2 ]5
http:\asadipour.kmu.ac.ir 81
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69
Keq
E  redox
0.059
[Mn 2 ][Fe 3 ]5

 Log
5
[MnO 4 ][H  ]8 [Fe 2 ]5
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
K eq 
E

redox
[Mn 2 ][Fe 3 ]5
[MnO 4 ][H  ]8 [Fe 2 ]5
-
0.059

 Log K eq
5
Log K eq
920316
Log K eq
5E  redox

0.059
(5  1)E  redox

0.059
http:\asadipour.kmu.ac.ir 81
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70
Keq
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• MnO4-+5e-+8H+→ Mn2++ 4H20
•
Fe3+ +e → Fe2 +
• LogK eq
920316
E0=1.51
E0= 0.771
n=5
n=1
5 1(1.51 - 0.771)

 62.5
0.059
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71
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
20.5 ml After Eq
• Mno4-+8H+  Mn2+
0.059
[Mn 2 ]
• E  1.51  Log
5
[MnO 4 ][H  ]8
[ Mn 2 ] 

N1V 1 100  0.5

 0.415 N  0.083M
V1  V 2
120.5
[ MnO4 ] 
N 2V 2  N1V 1 20.5  2.5  100  0.5

 0.010 N  0.002 M
V1  V 2
120.5
0.059
0.083
 Log
 1.491V
8
5
0.002 1
http:\asadipour.kmu.ac.ir 81
• E  1.51 920316
slides
72
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• Mno4
-+8H+ 
Mn2+
21 ml After Eq
0.059
[Mn 2 ]
• E  1.51  Log
5
[MnO 4 ][H  ]8
[ Mn 2 ] 

[ MnO4 ] 
N1V 1 100  0.5

 0.413N  0.083M
V1  V 2
121
N 2V 2  N1V 1 21 2.5  100  0.5

 0.021N  0.004 M
V1  V 2
121
0.059
0.083
• E  1.51  Log
 1.494V
8
5
0.004 1
920316
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73
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
22 ml After Eq
• Mno4-+8H+  Mn2+
0.059
[Mn 2 ]
• E  1.51  Log
5
[MnO 4 ][H  ]8
[ Mn 2 ] 

[ MnO4 ] 
N1V 1
100  0.5

 0.410 N  0.082M
V1  V 2
122
N 2V 2  N1V 1 22  2.5  100  0.5

 0.041N  0.008M
V1  V 2
122
• E  1.51 920316
74
0.059
0.082
 Log
 1.498V
8
5
0.008 1
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slides
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
23 ml After Eq
• Mno4-+8H+  Mn2+
0.059
[Mn 2 ]
• E  1.51  Log
5
[MnO 4 ][H  ]8
[ Mn 2 ] 

[ MnO4 ] 
920316
N1V 1 100  0.5

 0.407 N  0.081M
V1  V 2
123
N 2V 2  N1V 1 23  2.5  100  0.5

 0.061N  0.012M
V1 V 2
123
0.059
0.081
• E  1.51  Log
 1.500V
8
5
0.012 1
75
http:\asadipour.kmu.ac.ir 81
slides
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
25 ml After Eq
• Mno4-+8H+  Mn2+
0.059
[Mn 2 ]
• E  1.51  Log
5
[MnO 4 ][H  ]8
[ Mn 2 ] 

[ MnO4 ] 
920316
N1V 1 100  0.5

 0.4 N  0.08M
V1  V 2
125
N 2V 2  N1V 1 25  2.5  100  0.5

 0.1N  0.02M
V1 V 2
125
0.059
0.08
• E  1.51  Log
 1.503V
8
5
0.02 1
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76
Titration curve
• 100 ml Fe2+ 0.5 M WITH Mno4- 0.5 M
• MnO4-+5Fe2++8H+  Mn2++5Fe3++4H20
• Mno4-+8H+  Mn2+
30 ml After Eq
0.059
[Mn 2 ]
• E  1.51  Log
5
[MnO 4 ][H  ]8
[ Mn 2 ] 

[ MnO4 ] 
920316
N1V 1 100  0.5

 0.385 N  0.077 M
V1  V 2
130
N 2V 2  N1V 1 30  2.5  100  0.5

 0.115 N  0.023M
V1 V 2
130
0.059
0.077
• E  1.51  Log
 1.504V
8
5
0.023 1
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77
920316
0.743
0.771
0.799
0.846
0.866
1.387
1.491
1.494
1.498
1.5
1.503
1.504
Δ2E/ΔV2
E(v)
5
10
15
19
19.5
20
20.5
21
22
23
25
30
ΔE/ΔV
ml of MnO4K
Titration curve
0.0056
0.0056
0
0.01175 0.001537
0.04
0.0565
1.042
2.004
0.208
-1.668
0.006
-0.404
0.004
-0.002
0.002
-0.002
0.0015 -0.00025
0.0002 -0.00026
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78
E
Titration curve data
1.6
1.4
1.2
1
0.8
0.6
0.4
0.2
0
Height is related to Keq
Not related to concentration
0
920316
10
20
ml of titrant
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slides
30
40
79
Titration curve
1.2
1
ΔE/ΔV
0.8
0.6
0.4
0.2
0
0
-0.2
920316
10
20
30
40
ml of titrant
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80
Titration curve
2.5
2
Δ2E/ΔV2
1.5
1
0.5
0
-0.5 0
10
20
30
40
-1
-1.5
-2
920316
ml of titrant
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81