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Redox Titrations. -the oxidation/reduction reaction between analyte and titrant. -titrants are commonly oxidizing agents, although reducing.
Typology: Summaries
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ox
red
!
red
ox
4+
2+
!
3+
3+
2
2
3
2+
0
4+
3+
0
4
3+
!
2+
2
2
4+
!
3+
2
2
3+
cell
cathode
anode
Ag+
Cu2+
cathode
anode
Ag+
Cu2+
Ox
Ox
Ox
Ox
Ox
2+
0
Ag+
2
log =
1
0
Cu2+
2
log
[Cu2]
1
0
Ag+
0
Cu2+
[Cu ]
1
log
2
[ ]
1
log
2
2 2 +
!
Ag +
2 eq
2 Cu
0 0
log K
[Ag ]
[Cu ]
log
2( )
=
=
!
E (^) Ag+ E +
2(0.799- 0.337)
[Ag ]
[Cu ]
logK log 2
2
eq
=
=
= 15.
K eq = antilog 15.6 = 4.1 x 10
15 = 4 x 10
15
Fe
2+
4+! Fe
3+
3+
Ce
4+ = E Fe3+
system
In
Ce4+
Fe3+
system
3
2+
4+
3+
[Ce
[Ce
log
Ce eq
[Fe
[Fe
log
Fe eq
[Ce ][Fe ]
[Ce ][Fe ] 0
Fe
Ce eq 4 3
3 2
4 3 + +
3+
3+
2+
4+
f
Fe
f
Ce
eq
4+
[Fe
2+ ] = amt of Ce
4+ left unreacted, therefore added to C Ce4+
calculated from the volumes of the two solutions and subtracted from
Ce3+
Conc of two cerium ion species:
[Ce
3+ ]=
10
500
]
2
25.00x 0.
[Fe!
"
[Ce ]
[Ce ]
log
4
3
010 / 75. 10
500 / 75. 00
log
1
Effect of system variables on redox titration curves
Concentration – independent of analyte and reagent concentrations.
Exception: Electrode potentials dependent upon dilution
!
3
+2e
log
3
3
0 E= E!
num-mol/L
3 , denom-mol/L
Completeness of reaction – the change in E system
in the e.p. region
becomes larger as the reaction becomes more concentrated.
Redox indicators
a. specific indicators – react with one of the participants in the
titration to produce a color, e.g. thiocyanate
b. Oxidation-reduction indicators- respond to the potential of the
system rather than to the appearance or disappearance of some
species during the course of the titration, e.g. methylene blue
Color changes will occur over the range:
Volts
n
0 = ±
where n = # of electrons in the indicator half-reaction
break in the titration curve at the e.p.
≥0.2 V, best detected potentiometrically
2+ preoxidized to MnO 4
interfere in subsequent titration
Preoxidation
peroxydisulfate (S 2
8
2 - ) – requires Ag+ as a catalyst.
+! + +
4
2 -
4
2 -
2 8
S O Ag SO SO Ag
Excess reagent destroyed:
2 -
4
2 -
2 8
boiling
Prereduction
) will reduce Fe
3+ to Fe
2+ in hot HCl
Excess reductant is then destroyed:
Sn
2+
! Sn
4+
HgCl 2
2 Cl
Oxidation with Potassium Permanganate
In strongly acidic solutions, reduced to colorless Mn
2+ :
MnO 4
2+
In neutral or alkaline solution, the product is the brown solid, MnO 2
MnO 4
(s) + 2H 2
In strongly alkaline solution (2 M NaOH), green manganate is
produced:
MnO 4
2 -
Tales 16.3……..see below
Note: permanganate solutions are unstable, therefore not a primary
standard.
4MnO 4
O > 4MnO 2
( MnO 2
catalyses this
reaction )
Permanganate must be standardized for example with oxalate;
2
2
4
CO 2
2e
Overall:
. 2MnO 4 - + 5H 2
2
4
> 2Mn
2+
2
KMnO4 can serve as own indicator, since product Mn
2+ is colorless.
Cerium(IV)
Strong oxidant > Ce(III)
Ce
4+ [Yellow ]+ e
3+ [Colorless]
Note however that the color change not good enough for it to act as
self indicator.
Ce(IV) not found in acid solution as simple aqua ion .. forms
complexes.
Dichromate reactions
Dichromate ion is an oxidizing agent
Cr 2
7
2 -
14H
6e
3+
Dichromate has replace permanganate in many analyses ... notably
iron (II)... it can be prepared as a standard solution and so avoids the
need to standardize as is the case with permanganate.
Iodine Methods
2
Value for E is intermediate can therefore be reduced or
oxidized...iodine can be reduced to iodide by for example As(III),
Sn(II) whilst iodide can be oxidized to iodine by for example
permanganate. Use of iodide as titrant..practical problems ..so add
excess potassium iodide and titrate the liberated iodine with for
example standard thiosulphate solution
Miscellaneous Oxidizing agents
Sodium bismuthate and lead (IV) oxide are strong oxidizing agents.
