The Oxidation Potential of Postassium Ferrocyanide-potassium Ferricyanide

8/10/2019 The Oxidation Potential of Postassium Ferrocyanide-potassium Ferricyanide

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THE

OXIDATION POTENTIAL OF

THE

SYSTEM POTASSIUM

VARIOUS IONIC STRENGTHS

I.

M. K O L T H O F F AND

WILLIAM

J. TOMSICEK

School of Chemis t ry , U nivers i ty of Minnesota , Minneapol is , Minnesota

Received January

85 936

FERROCYANIDE-POTASSIUM F E R ~ I C Y A N I D E

AT

The oxidation potential of the ferrocyanide-ferricyanide system has

been determined by a number of investigators

(1,

3 ,

8, 9, 10, 11,

12,

13).

After the introduction of the Debye-Huckel theory of strong electrolytes,

this system becomes of special interest, since we are dealing here with

highly unsymmetrical salts of high valence type. If potassium ferro-

cyanide and potassium ferricyanide behave like strong electrolytes, the

oxidation potential should be greatly affected by a change of the ionic

strength of the solution. In the first place, the purpose of this study was

to determine the potential of the potassium ferrocyanide-potassium

ferricyanide system a t varying ionic strengths and to extrapolate the value

to an ionic strength of aero; in other words, to determine the normal po-

tential of the system. In addition, the potential of

a

very dilute ferro-

ferricyanide solution was determined in the presence of different neutral

salts at varying ionic strengths, in order to test the applicability of the

Debye-Huckel equations.

At extremely small ionic strengths, the relation between the activity

coefficient of an ion and the ionic strength of the solution is given by the

expression

:

-log

f

= 0 . 5 ~ ~

i

(1)

at 25C. in water, in which z is the valence of the ion, and 1 1 the ionic

strength. The oxidation potential E of the system ferrocyanide-ferri-

cyanide a t 25C. then is given by:

CFeT;- f a

=

0 + 0.0591 log

CFeOC- f 4

(2)

From the exper imen tal pa r t of a thesis submitted by Will iam J. Tomsicek to

the Gradua te School of t he Unive rsity

of

Minnesota in part ial fulf i l lment

of

t h e

requirements for the degree of Doctor of Philosophy, 1934.

945


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946

I. XI

KOLTHOFF

AND WILLIAM

J. TOMSICEK

The normal potential

eo

denotes the potential referred to the normal

hydrogen electrode in

a

system in which the activity of the ferricy-

anide

aF ;-

is

equal to that of ferrocyanide

aFi;--.

CF;;-

and

CF~;;--

represent the corresponding concentrations, whereas fs and f 4 represent

the activity coefficients of the ferricyanide and the ferrocyanide ions.

If the limiting Debye-Hudkel expression (equation

1)

holds a t extremely

small ionic strengths and the system contains equimolecular amounts

of potassium ferricyanide and potassium ferrocyanide, it

is

found from

equations

1

and 2 that:

54

Therefore if the limiting Debye-Huckel expression holds, the measured

potential E should change by 0.2068 volt for one unit change in the square

root of the ionic strength.

The practical work in this study involves the use of

a

cell with liquid

junction, the ferro-ferricyanide half-cell being measured against the

quinhydrone electrode in a mixture containing 0.01 of an equivalent of

hydrochloric acid and 0.09 of an equivalent of potassium chloride per liter,

the saturated potassium chloride-agar salt bridge being used for making

electrolytic contact between the two half-cells. No correction has been

applied for the liquid junction potential, which is very small in dilute

solutions containing potassium ferrocyanide and potassium ferricyanide,

but may be greater in the presence of larger amounts of neutral salts.

The introduction of the liquid junction potential, however, does not invali-

date the conclusions arrived a t in this paper.

EXPERIMENTAL PART

Materials used

IGFe(CN)a.3Hz0. A C. P . product of potassium ferrocyanide was

recrystallized twice from conductivity water and kept over deliquescent

sodium bromide hydrate. An analysis of the salt showed that it had the

theoretical composition.

K3Fe(CN)s

A

C . P . product of potassium ferricyanide was recrystallized

twice from conductivity water and dried over anhydrous calcium chloride.

The various salts used in this work had been analyzed by

W.

Bosch

and had been used in a previous study 4). Conductivity water was used

throughout this work.

Apparatus and method for the measurement

of

the potential

The potential of the ferro-ferricyanide system was measured in a Pyrex

cell as shown in figure

1,

a piece of bright platinum gauze serving as elec-


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OXIDATION POTENTIAL OF FERROCYANIDE-FERRICYANIDE 947

trode. One terminal of the potassium chloride-agar salt bridge was

placed in the side well b, thus preventing diffusion of potassium chloride

from the bridge into the main body of the solution. Nitrogen gas from a

tank was introduced through e. Oxygen gas was removed from the nitro-

gen by passing the gas through electrically heated copper gauze a t 500C.

The solution in the standard reference half-cell 0.01 N hydrochloric acid,

0.09 N potassium chloride saturated with quinhydrone) was prepared fresh

every day. The normal potential of the quinhydrone electrode is 0.6990

volt a t 25C. Assuming that the paH of the acid mixture in the quin-

FIG. FIG.2

F I G .

1

THE

ELL

FIG.2.

R a t i o

of

K s F e C N ) e t o K I F e C N ) s : o , r a ti o

1 : l ;

A,

ra t i o

1O:l;

0

ra t i o

1:lO

D. H., calculated from simple Debye-Huckel expression.

hydrone half-cell

is

equal t o 2.0755, we find that t he potential of the lat ter

against the normal hydrogen electrode is equal to 0.5764 volt a t 25C.2

All the measurements were made in a thermostat a t 25 C. f 0.05'.

Various salt bridges were used, all yielding the same values.

The measure-

ments were made with a Leeds and Northrup student potentiometer.

For

the dilution experiments a stock solution containing 0.1 M potassium

ferrocyanide and

0.1

M

potassium ferricyanide was carefully prepared by

weight from the pure salts. This stock solution was kept in the dark and

Recent ly Guggenhe im and Schindle r

(J.

Phys . Chem.

38,533

1934)) gave evidence

tha t t he paH of t he s t anda rd ac id mix tu re used in t he qu inhydrone e l ec t rode is

equa l t o 2.10.


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948 I . M . KOLTHOFF A N D WILLIAM J . TOMSICEK

prepared fresh each day. The solutions from 0.1 to

0.004

molar were

found to give the same potential in air as in a nitrogen atmosphere.

The

potential of the 0.004 molar solution in air referred to the normal hydrogen

electrode was

0.4009

volt after

5

minutes and

0.4011

volt after

60

minutes.

The same solution in a nitrogen atmosphere gave readings of

0.4011

and

0.4012 volts after 5 and 60 minutes respectively.

More dilute solutions

gave higher readings in air than in nitrogen.

The potential of the

0.0004

molar solution was measured a t least ten times during the course of the

investigation. In

a

nitrogen atmosphere, the values found after 5 minutes

TABLE

Ox idat ion potential of equimolecular mixtures of pota ssiu m ferrocyanide and po tassi um

ferricyanide

M

0 . 1 *

0.04

0.02

0 . 0 1

0.007

0.004

0.002

0.001

0.0008

0.0004

0.0002

0.0001

0.00008

0.00006

0.00004

P

1 . 6

0.64

0.32

0.16

0.112

0.064

0.032

0.016

0.0128

0.0064

0.0032

0.0016

0.00128

0.00096

0.00064

1.265

0 . 8

0.5657

0 .4

0.334

0.253

0.173

0.1265

0.1131

0.08

0.0566

0 .04

0.0358

0.031

0.0253

E

AGAINST S T A N D A R D

QUINHYDRONE)

0.1178

0.1362

0.1490

0.1610

0.1670

0.1753

0.1856

0.1930

0.1950

0.2010

0.2O50

0.2100

0,2112

0.2122

0.2145

E

AGAINST

N O R M A L

ELECTRODI)

H Y D R O Q E N

0.4586

0.4402

0.4276

0.4154

0.4094

0.4011

0.3908

0.3834

0.3814

0,3754

0,3714

0.3664

0.3652

0.3642

0.3619

* M =

0.1

designates that the concentrations of both potassium ferrocyanide and

potassium ferricyanide are equal to 0.1 mole per liter.

remained unchanged for periods of twelve hours and more. The various

readings agreed within ~ t 0 . 0 0 0 3olt, the average being 0.3754 volt. The

reproducibility of measurements with solutions from 0.0004 to 0.00006

molar was within 0.0005 volt. Each of the solutions was prepared fresh

and measured a t least four times.

Light was found to have

a

distinct effect on solutions whose concentra-

tions were

0.0004

M or less, the E.M.F. tending to increase in light. All

measurements were therefore made in

a

darkened room. Under these

conditions, the potentials of even the most dilute mixtures remained

constant for a t least one hour.


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0 4450

0.4275

0.4153

0,4043

0.3915

0,3836

0.3775

0.3711

Experimental results

Table

1

gives the average of the results of measurements with equi-

molecular mixtures

of

potassium ferrocyanide and potassium ferricyanide,

p representing the ionic strength. The value of O was found by plotting

the measured values of E against on large cross section paper and

TABLE 2

Osidation potentials measured

in

a mixture containing

K8Fe CN)6

nd

K,Fe CN)s

in

the ratio

O : i

KaFe CN)o

M

0 . 1

0.04

0.02

0 01

0.004

0.002

0 001

0.004

KdFe CN)o

M

0 01

0.004

0.002

0.001

0.0004

0.0002

0 001

0.00004

T O T AL

p

0 .7

0 .28

0 . 1 4

0.07

0.028

0.014

0.007

0.0028

0.8366

0.529

0.3742

0.2646

0.1673

0.1183

0.0837

0.0530

E

A G A I N S T

DRONE)

S T A N D A R D

Q U I N A Y -

0.0723

0.0898

0 1020

0.1130

0.1258

0.1337

0.1398

0.1462

E

A G A I N S T

E L E C T R O D E )

N O R M A L

H Y D R O G E N

0.5041

0.4866

0.4744

0,4634

0,4506

0.4427

0.4366

0.4302

CAL CUL AT E D

TABLE

3

Oxidation potentials measured

in

a mixture containing

K3Fe CN)8

nd

KIFe CN)8

KaFe CN)e

M

0..01

0.004

0.002

0.001

0.0004

0.0002

0.0001

0.00004

KpFe CN)s

M

0 . 1

0.04

0.02

0 . 0 1

0.004

0.002

0.001

0.0004

in

the ratio : l O

T O T AL

p

1 06

0.424

0.212

0.106

0.0424

0.0212

0,0106

0.00424

1.0295

0.6511

0.4604

0.3256

0.2083

0.1456

0.1029

0.0651

E

AO AI NST

DRO NI )

S T A N D A R D

Q U I N H Y -

0.1895

0,2068

0.2188

0.2295

0.2420

0.2495

0.2555

0.2615

E

AGAINST

ELECTRODE)

N O R M A L

H Y D R O Q E N

0.3869

0.3696

0,3576

0.3469

0.3344

0.3269

0.3209

0.3149

60

CAL CUL AT E D

0.4460

0.4287

0,4167

0.4060

0 .3935

0.3860

0.3800

0.3740

extrapolating to an ionic strength of zero. It was found to be equal to

0.3560 volt.

The straight line repre-

sents the change of

E

assuming that the limiting Debye-Huckel expression

holds (equation 3).

In addition, the @oxidation potentials were measured in mixtures con-

taining ratios of potassium ferrocyanide and potassium ferricyanide of

10:

1

and

1

:10. They were recalcu-

The data are plotted in figure

2.

The data are given in tables

2

and 3.


6/10

950

I .

M.

KOLTHOFF

A N D

WILLIAM

J.

TOMSICEK

, . . . . .

z


7/10


AGAINST

ELECTBODE)

N O R M A L

H Y D R O G E N

0.4187

0.4097

0.3987

0.3918

0.3864

0.3809

lated

on

the basis of a ratio of the concentrations of 1

:

and correspond to

the e; values plotted in figure

2.

These E values are identical with the

values of the potential

E

measured in the equimolecular mixtures

of

ferro-

cyanide and ferricyanide. The reproducibility of the measurements in

the very dilute solutions containing unequal molecular ratios of ferro-

cyanide and ferricyanide i s not as good as of those reported in table 1;

therefore, the extrapolated value of

e

a t an ionic strength of zero is less

reliable in the former cases.

T h e e f ec t of neutral salts u p o n the potential

In all of the following determinations a solution containing 0.0004 molar

potassium ferrocyanide and 0.0004 molar potassium ferricyanide, freshly

prepared by dilution of a 0.01 molar mixture, was used.

Ten ml. of the

CONCENTRA-

T I O N O F

Nan C I T R A T E

hi

0.0833

0.0416

0.0166

0.0083

0.00416

0.00166

T OT A L p

0.5064

0.2564

0.1064

0.0564

0.0314

0.0164

In

C ON C E N T R A -

T I O N O F

NarPzOi

M

0.05

0.025

0.01

0.005

0.025

0.001

TABLE 6

hence of

salts on

the

oxida t ion potent ial

~~

E

A G A I N S T

N O R M A L

H Y D R O G E N

E L E C T R O D E )

0.4262

0.4154

0.4034

0.3953

0.3889

0.3826

CONCENTRA-

T I O N O F

MgSOi

ti

0.125

0,0625

0.025

0.0125

0.00625

0.0025

E

AGAINST

E L E C T R OD E )

N O R M A L

H Y D R O O E N

0.4584

0.4474

0.4344

0.4246

0.4152

0.4028

latter was diluted with conductivity water in a 2 5 0 4 . volumetric flask,

a

weighed amount of pure salt added, and the flask filled up to the mark.

The results are given in tables 4, 5, and

6

Total refers to the sum

of the ionic strengths of the added salt and of the 0.0004 molar ferro-

cyanide-ferricyanide mixture p = 0.0064). From equation

3

i t is found

that

f 3 E 0

log - =

4

0.0591

The values of log 3/f4 thus derived in various salt solutions are plotted in

figure 3 against di . The straight line again gives the values calculated

with the assumption t hat the limiting Debye and Huckel expression holds

a t extreme dilutions.

DI S CUS S I ON O F RE S UL T S

1 The generally accepted value of the normal potential of the ferro-

A t an ionic strength

erricyanide electrode of 0.44 volt is much too high.


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952

I M.

KOLTHOFF

A N D WILLIAM J TOMSICEK

of zero, a value of

0.356

volt was derived in this paper. From a practical

viewpoint i t is of interest to mention th at the oxidation potential increases

very rapidly with the increasing ionic strength and that it even can exceed

the value of

0.44

in equimolecular mixtures of ferricyanide and ferrocyanide.

2.

Even at infinite dilutions, the behavior of the system is not in har-

mony with the postulates of the simple Debye-Huckel expression. The

slope of the curve giving the change of the oxidation potential or of log

f3/f4 plotted against the square root of the ionic strength isgreater than

o a i

FIG.

3

FIG.3 .

a,

CsCl; b, RbC l; c , KC l and NHcC1; d , LiCl . D . H., calculated from simple

Debye-Htickel expression.

FIG.4.

a,

M g NO s)z; b , BaCl2; c , Ca N Oa )n; d, SrC12; e, NazSOa;

f , Nas

c i t r a t e ;

g ,

NarPzO,.

D.

H ., ca lcu la ted

f r o m

simple Debye-Huckel expression.

that calculated on the basis of the Debye-Huckel limiting equation.

It

is

impossible

to

account for this anomaly on the basis of ionic size, using the

present form of the Debye-Huckd theory, for, as V. K. La Mer 6) states,

absurd negative values of u would be demanded at very high dilutions

followed by positive values in more concentrated solutions. Deviations

of experimental data from the theoretically predicted curves have been

described by various authors, a discussion of which is given in a paper by

La Mer, Gronwall, and Greiff

7).

Gronwall, La Mer, and Sandved 2)

have shown that these discrepancies disappear if the influence of higher


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O X I D A T I O N P O T E N T I A L O F FERROCYANIDE-FERRICYANIDE 953

terms of the Debye-Huckel theory in the case of unsymmetrical valence

type electrolytes is taken into account. On the basis of the extended

Debye-Huckel equation, values are found which fit the experimental data

without assuming ion association or incomplete dissociation of the strong

electrolytes. Undoubtedly, in a quantitative interpretation of the data

found in this study, the extended equation of Gronwall, La Mer, and

Sandved should be applied, since we are dealing with highly unsymmetric

valence type electrolytes. Still, we have evidence to believe tha t even the

extended equation does not account quantitatively for the results obtained,

and that potassium ferrocyanide has to be considered as an incompletely

dissociated electrolyte. In a subsequent paper, it will be shown that the

curve obtained in a study of the potential of the potassium molybdo-

molybdicyanide electrode, a system very similar to that of ferro-ferri-

cyanide, does not intersect with the straight line calculated from the

simple Debye-Huckel expression, but is found below this line even a t

extreme dilutions. In addition it was found tha t the fourth dissociation

of molybdocyanic acid

HMo(CN)s--- H+ + Mo(CN)s----

is complete whereas that of ferrocyanic acid

HFe(CN)a---

H+

+ Fe(CN)a----

is incomplete. This means tha t the proton combines with the ferrocyanide

ion tof~rmHFe(CN)~-- -, ndi tis quite plausible tha t other cations behave

similarly.

I n

the study of the influence of salts upon the potential of

a

very dilute potassium ferrocyanide-potassium ferricyanide mixture

described in this paper i t was found that the effect is virtually independent

of

the type of the anions. Potassium bromide, chloride, and nitrate have

an identical effect a t the same ionic strength; the same is true for sodium

chloride, nitrate, and perchlorate on the one hand and sodium sulfate,

oxalate, carbonate, and phosphate on the other.

With the

alkali cations i t decreased in the order

Cs,

Rb,

K

= NH,, Na = Li, and we

conclude that the degree of dissociation

of

the corresponding ferrocyanides

decreased in the same order. The dissociation becomes more incomplete

with the increasing valence of the cations, the effect of the various alkaline

earths being of about the same order. This larger effect of the divalent

ions is especially pronounced at the smaller ionic strengths.

Since the concentration of the cation is of primary importance, it is

easily understood why the oxidation potentials found in potassium ferro-

cyanide-potassium ferricyanide mixtures of various ratios and recalculated

on the basis of a ratio

of

1:

1

are not the same at the same ionic strength

The type of cation, however, has a very pronounced effect.


10/10

954 I. M. KOLTHOFF AND WILLIAM J. TOMSICEK

(figure

2).

The potentials found increase from the mixture with a ratio of

10

ferrocyanide to 1 ferricyanide to t ha t with a ratio of

1

to 10. In the

former, the potassium-ion concentration is much smaller than in the latter

a t the same ionic strength. In a

similar way, it is explained why the 1-2

valence types of electrolytes (sodium sulfate, carbonate, etc.) have

a

smaller effect than the

1-1

valence type of salts (sodium chloride, etc.).

SUMMARY

1. The normal potential of the ferrocyanide-ferricyanide electrode is

equal to 0.3560 volt at

25C.

2. The change of the potential of a very dilute ferrocyanide-ferricyanide

solution with increasing ionic strength is greater than calculated on the

basis of the simple Debye-Huckel expression. This is partly explained

by incomplete dissociation of alkali and alkaline earth ferrocyanides.

3 .

For the same valence type of salts the anion effect upon the potential

is the same for different anions a t the same ionic strength.

A

pronounced

cation effect was observed, the effect decreasing in the order

Cs,

Rb, K =

NHI Na =Li for the alkali ions and being of about the same order for the

alkaline earth ions. The latter, especially a t the smaller ionic strengths,

have

a

much greater effect than the univalent cations.

R E F E R E N C E S

1)

FREDENHAGEN,

. : Z. anorg. a llgem. Chem.

29,

396 1902).

2) GRONWALL,.

H.,

L A M E R ,V. K . , AND SANDVED,. : Phys ik .

Z.

29,558 1928).

3) KOLTHOFF,. M . : Z . anorg. a llgem. Chem. 110, 143 1920).

4) KOLTHOFF,. M., AND BOSCH,W . :

J.

Phys . Chem. 36, 1685 1932).

5) KOLTHOFF,

.

M . : T h e D e te r m in a ti o n

of

pH, Elec trometr ic Ti t ra t ions .

6) L A MER,

V.

K.: Trans . Am. Elec trochem SOC. 1, 543 1927).

7) L A M E R ,

V.

K., GRONWALL,. H., AND GREIFF,L. J.: J. Phy s . Chem. 36,2245

(8) LEWIS,G. N.,

AND

SAROENT,. W.:

J.

Am. Chem. SOC. 1,355 1909).

9) LINHA RT, .

A , :

J. Am. Chem. SOC.39, 615 1917).

J o h n

Wiley and Sons, New

York

1931).

1931).

10)

MULLER,

E.:

Z. physik. Chem.

88,

46 1914).

(11)

SCHAUM , . , AND LINDE,R. v . D . : Z. Elektrochem. 9,407 1903).

12) SCHOCH,

.

P.: J. Am. Chem. SOC.

6,

1422 1904).

13) SCHOCH,. P.,

AND

FELSING,. A , : J. Am. Chem. SOC.38, 1928 1916).

Documents

The Oxidation Potential of Postassium Ferrocyanide-potassium Ferricyanide