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THS

ELECTROMETRIC TETPATITON
or URWUM erH CERIC SULFATE ’

THESIS 3703 THE DEGREE OF M. s‘

Mabel Florey \Vilson
1930 .

 

    

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JELECTROMETRIC TITRATION
OF URANIUM

WITH
CERIC SULFATE

THESIS

summltted to the Faculty
of
MICHIGAN STATE COLLEGE
In Partial Fulfillment of the
Requirements for the Degree
of

Mister of Science

By

Mabel Florey wileon
June, 1930

In a previous publication (This Journal, 33,
1484 (1922))it was shown by Ewing and Eldridge that the
electrometric titration curves for uranium and
potassium permanganate, and potassium dichromate had
two inflection points. The first denoted the complete
oxidation of trivalent uranium to the tetravalent form
and the second point, the complete oxidation of the
tetravalent form. Gustavson and Knudson (Gustavson
and Knudson, This Journal, 55, 2756 (19222)de Muller
and Plath (Muller and Plath, z. Elektrochem., 33
(1925)) titrated uranium electrometrically, but did
not note the first inflection point. In commenting on
this point of difference it was correctly pointed out
by Kdlthoff and Furnam (I.M. Koltheff and N. H. Furnmm
"Potentiometric Titrations", iggg. P 245-4) that the
presence of carbon dioxide and the correct acidity must
be obtained in order to form and titrate trivalent
uranium. In the previous investigation a special
titration cell was used which was designed to eliminate
oxygen and the titration was only successful when made

in an atmosphere of carbon dioxide or nitrogen.

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PREPARATIOE OF SOLUTIONS

The ceric sulfate solution was prepared by
the same method as that used by Willard and Young
(H. H. Willard and Philena Young, This Journal, _5_o_,
1522 (1928) ). Approximately 500 g. of U.S.P. cerium
oxalate were ignited in a platinum crucible, in the
Open air until all gas was given off, and then for 10
hours in an oven at BOO-625°C. This yielded about
' 150 g. of reddish brown ceric oxide, which was heated
approximately 5 hours at 125°-150°C, with 210 ml.of
sulfuric acid (1.84 sp. gr.) until a yellow paste was
formed. It was then dissolved in about 5.5 liters of
distilled water, by heating at 75-80°C and stirring
constantly for an hour. This was filtered hot through
the undiesolved substance.

The ceric sulfate solution was standardized
against 0.05 N sodium oxalate solution made from a
Bureau of Standards product in a carefully calibrated
flask, at 25.0°C. This ceric sulfate solution was
found to be 0.05326 N.

The uranium solution was prepared by
dissolving 28.285 g. of uranyl acetate in a little
distilled water containing 9.8 g. sulfuric acid and

carefully diluting to 2 liters.

The uranium solution was standardized by
evaporating two 20 ml.portions of stock solution
to dryness and igniting to the oxide. From this
data the normality of the uranium solution was
found to be 0.07115.

An approximately 0.1 N solution of ferrous
ammonium sulfate solution was prepared with an excess
of free sulfuric acid. It was standardised against
the eerie sulfate solution and was found to be

0.1041 No
APPARATUS

The e.m.f. of the titration cell was
determined with a potentiometer which read to one-
half of a millivclt.

The titration vessel consisted of a
400 ml beaker with a large cork stapper with six
holes for the following pieces:- a burette, a zinc
reductor, a mechanical stirrer, an inlet for an inert
gas and two electrodes. The standard electrode was a

tenth normal calomel half cell and a small bright

platinum wire was used for the other electrode.

EXPLR IMEN TAL PRO CEDURE

1.06 ml of concentrated sulfuric acid was
added to 50 ml of the uranium solution and evaporated
to dryness. This was diluted to 50 m1., heated
nearly to boiling and forced through a 50 cm zinc
reductor with nitrogen or carbon dioxide. The zinc
reductor was washed with 2% sulfuric acid until the
total volume of solution in the titrating vessel was
100 ml. This solution while warm was titrated with
the standard cerium sulfate reagent. During this
procedure nitrogen or carbon dioxide was run through
continuously.

The first end point was approached rather
slowly and usually required about 5 ml of the cerium
solution. After each addition it was necessary to
wait a few moments for an equilibrium reading. After
the first end point was established the cerium
solution could be added quite rapidly until near the
second end point. The behavior of the reaction as the
second end point is approached is quite similar to

that of any other cerium electrometric titration where

equilibrium conditions must be carefully noted.

The volume of ceric sulfate represented
by the difference between the two end points was
carefully computed, corrected for calibration
errors, etc., and from this value the concentration
of the uranium solution was found.

Graph 1 (Fig. l) is typical of the results
of many titrations which were made. I

In Table 1 doagiven the results of seven

consecutive titratione.

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TABLE I

The total volume of the solution titrated

was 100 ml and it contained 2 percent sulfuric acid.

 

 

 

 

 

ml Ceric
Ceric Sulfate Sulfate Solution

Titration Uranium

Number Solution byyexp't calculated ml uranium
1 29.95 ml 59.88 ml 59.86 ml 1.556
2 29.85 " 59.85 ” 59.86 " 1.554
5 29.85 " 59.98 " 59.86 " 1.559
4 29.85 ” 59.84 " 59.86 " 1.555
5 29.85 " ‘59.64 " 59.86 " ‘ (1.527)
6 29.85 ” 59.95 " ‘ 59.86 ” 1.558
7 29.85 " 59.86 " 59.86 " 1.556

Mean 59.87 ” 1.556

In experiment number 5 trouble was encountered

with the mechanical stirrer.

DISCUSSION OF THE RESULTS.

The stock solution of uranium used in
this investigation was prepared from the acetate.
Irregular results were obtained in the first
titations. The second end point was uncertain and
the results were low. This difficulty was not
overcome until the acetate was removed. This was
accomplished by evaporating the solution with
sulfuric acid until the acetate had disapeared.

The concentrated solution was diluted with water Just
previous to reduction. .

The color of the freshly reduced acid
uranium solution was a dark olive green, and was
slightly opaque. Upon addition of ceric sulfate
solution it gradually cleared and became a very light
green color at the first end point. is oxidation
proceeded the color became less intense and gradually
changed to a greenish-yellow at the second end point.

Due care must be taken in this titration
to make certain that the e.m.f. represents an
equilibrium condition. Furnam (N. F. Furman, This
Journal, £9; 755, (1928) ). in titrating ferrous

sulfate allowed one to four minutes after each

addition before making the e.m.f. measurements.
In titrating uranium with cerium even more time was
required for the e.m.f. reading especially at the

first end point.
The initial e.m.f. of the reduced solution

was usually -600 to -450 mv. ’This changed
gradually about 100 mv.to the first end point when
an abrupt change of approximately 400 mv. was noted.
A small fraction of a drop of ceric sulfate solution
is sufficient to cause this abrupt change. The
second end point was also very definite and usually
the addition of a small fraction of a drop of the
eerie sulfate solution was sufficient to cause the
e.m.f. to increase 500 or 600 mv.

Higher concentrations of sulfuric acid,
6%, 155-5 and 50% by volume, were later titrated. The
amount of uranium solution, the total volume and
conditions of titration remained the same. In all
cases both end points were definite, and the amount
of cerio sulphate used between and points was
theoretical. Formerly reported titrations (gwing
andegldridge, ibid) stated that high acid concentrations
caused the voltage to rise rapidly and the second end

point consequently disappeared. In all cases of the

higher acid concentrations Just mentioned, the shape .
of the curves remained the same, and all three
corresponded closely to the curves obtained from
the 2% titration data. The only noticeable effect
was the time required for equilibrium to be
established at the end points, is the percent
acid was increased, a larger time was required for
equilibrium. This was particularly noticeable at
the second end point, which was fairly rapid in the
2% solution. In this reapect the work of Gustavson
99'

w
and Knudson (Gustavson.and9Knudson, ibid) has been

confirmed.

TITRATION 0E URANIUM WITH CERIUM SULFATE IN
HYDROCHLORIC ACID SOLUTION.

Hydrochloric acid.may also be used in the
titration of uranium with cerium sulfate. Titrations
were made using 2M., 4M., 6M., and a combination of
2H. for first end point, with an increase to 4M.
immediately after reaching the first end point.

In each case, thirty ml. of the uranium
solution was heated to boiling with half the acid,
and put through the reduotor. This was washed with

10

the remainder of the acid in solution, making a
total volume of 100 ml. for titrating. In each
case, more reduction to trivalent uranium was
noted than in the case of the sulphuric acid
titrations. The color of the reduced solution
was a very dark red brown and was Opaque. Upon
the addition of cerio sulfate, the color rapidly
changed to the same olive green, noted in the
case of sulphuric acid. Colors at the end points
were identical.

In no case was the titration with
hydrochloric acid as satisfactory as 2% sulfuric.
The great difference in effect of these two acids
on the titrations, was upon the change of voltage
at the end points, and the time required for
equilibrium to be established.

Graph number 2 gives a comparison between
the deflections in voltage at the two end points
in 2% sulfuric acid and 4M. hydrochloric acid titrations.

It will be noticed that the scope of the
first end point is 100 mv. with hydrochloric acid
and compared with approximately 400 in the case of
sulfuric acid; 450”for hydrochloric compared with

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11

675hfor sulfuric at the second end point.

In all cases where hydrochloric acid
was used, the first end point was too slow to be
practical. In the 2M., the first end point covered
about 400 mv., in the 4M., 100 ~ 200 mv., and in
6M., 60 - 75 mv. With higher acid concentrations,
the time required for equilibrium was greater.

This was also true of the sulfuric acid solutions,
although it did not seem to affect the voltage greatly
in case of the sulfuric acid solutions. In exact
contrast to the first end point, the stronger the
hydrochloric acid, the sharper and quicker the
second end point. This is also in exact contract
to the second end point when sulfuric acid was used.
In 2M. hydrochloric acid, the equilibrium was very
slow, and the change in voltage was approximately
500 - 400 mv; in am. it was 400 - 500 mv. and more
rapid; in em. which compared with 2% sulfuric acid,
the change was 600 mv. and more rapid.

By trying varying combinations of
S3Hi‘1t1es of hydrochloric acid on the two end points,
it was found that a concentration of 2M. hydrochloric
acid increased to 4M. after the first end point,

gave the best conditions for titration and the most

constant results.

12

A few titrations were made with a
combination of the uranium with ferrous ammonium
sulphate. Thirty ml. of the uranium solution were
prepared and reduced as in the 2% sulfuric acid
titrations. 'Before washing the reductor, 5 ml. of
the iron solution were added. The titration was
performed in 100 m1. volume, 2% sulfuric acid.

Three end points were obtained: one when
trivalent uranium was oxidized to tetravelent, another
when tetravelent was oxidized to hexavalent, and
a third when the ferrous iron was oxidized to
ferric iron. Because of the closely correSponding
voltage of the second and third oxidations, it is
necessary to plot also,.%%:: the change in.millivolts
per change in ml. This gave three definite and
points. The amount between the first and second
corresponded closely to the theoretical amount of
ceric sulphate.necessary to oxidize the uranium
present; and the amount between the second and
third, very nearly to amount required to oxidize
the ferrous iron present. Graph number 3 shows
this titration.

Nitrogen was used to exclude the air from

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13

the cell during the titrations. This is essential

in eucoeeafully.titrating the small amount of

trivalent uranium present in the reduced solution,

due to the rapid oxidation of trivalent uranium in

air, as shown by McCoy and Bunzel dfiocoy and
ww.Bunzel, This Journal, gl, 367 (1909) ).

14

SUMMARY

1. When a hot acid solution of uranium sulfate is
reduced in a Jones reductor and is
titrated in an atmosphere of nitrogen
with ceric sulfate as an oxidizing agent,
two end points are obtained. The amount
of uranium oxidized between the two end
points correSponds exactly to the amount
of uranium present. 2% sulfuric acid by
volume in a total volume of 100 ml. was
found to be the best conditions for
titration.

2. Uranium acetate in hydrochloric acid solution,
when treated in the same way, also gives
two end points, difference between which
is the amount of uranium present. A
concentration of am. for the first and
4M.for the second gave the best results.

3. Sharper end points were obtained by the titrations
in sulfuric acid, and less time was
required for the e.m.f. to come to

equilibrium, than in hydrochloric acid.

15

More consistent results were also obtained
with 2% sulfuric acid.

4. when 5 m1. of ferrous ammonium sulfate were added
to the reduced 2% uranium solution, three
and points were obtained; (1) when
trivalent uranium was oxidized to
tetravalent, (2) when tetravalent was
oxidized to hexavelent, and (3) when

ferrous iron was oxidized to ferric.

16

BIBLIOGRAPHY OF THE ELECTROMETRIC TITRATIONS 0F
URANIUM.ARD 0F CEBIC SULFATE.

1. (”Determination of Uranium and Vanadium"),
A. N. Finn.
J. Am. Chem. 800., _2_§, 1443, (1906).

2. ("The Speed of Oxidation, by Air, of Uranaue
Solutions, with a Note on the velumetric
Determination of Uranium").

H. M. MCCOY and H. H. Bunzel.
J. m. Chem. 300., Q, 567, (1909).

3. ("The Volumetric Determination of Cerium in
Cerite and Monozitefi)
F. J. Metzger and M. Heidelberger.
J. Am. Chem. 300., 93, 642, (1910).

4. ("The Estimation of Uranium and Phosphorus"),
H. D. Newton and J. L. H hes.
J. Am. Chem. 300., El, 17 1, (1915).

5. ("The Volumetric Determination of Cerium by
Means of Potassium Permanganate."),
Victor Lenher and C. C. Meloche.
J. Am. Chem. Soc., gs, 66, (1916).

6. ("The Eleotrometric Determination of Uranium
with Potassium Permanganat. and Potassium
Dichromate")

D. T. Ewing and E. F. Eldridge.
J. M. Chem. $00., .22, 1484, (1922).

7. ("The Successive Electrometric Titration
of Iron, Vanadium and Uranium").
R. G. Gustavson and C. M. Knudson.
J. Am. Chem. $00. 3.2., 2757, (1922).

8. ("The Determination of Uranium")
F. E. T. Ismael]. and H. B. KflOWles.
J. Am. Chem. 800., g, 2657, (1925).

l7

9. ("Applications of Cerio Sulfate in Volumetric
Analysis. I. The Preparation and Stability of
Cerio Sulfate (Containing Free Sulfuric Acid).
II. Potentiometric Study of the Reactions
between Cerio and Ferrous or Oxalate Ion.
Application to the Standardization of Cerio
Solutions. III. The Potentiometrio Determination
of Cerium")

N. Howell Furman.
J. Am. Chem. 800., fig, 755, (1928).

lO. ("Cerio Sulfate as a Volumetric Oxidising Agent.
The Volumetric Determination of Cerium").
H. H. Hillard and Philena Young.
J. Am. Chem. 300., £2, 1579, (1928).

11. ("Applications of Ceric Sulfate in Volumetric
Analysis").
N. Howell Furman.
J. Am. Chem. 300., 29-, 1675, (1928).

12. ("Cerio Sulfate as a Volumetric Oxidizing Agent.
VIII. The Determination of Chromium in the
Presence of Manganese, Iron and Vanadium").

H. H. Hillard and Philena Young.
J. m. Chem. 300., El, 139, (1929).

15. ("Cerio Sulfate as a Volumetric Oxidizing A ent.
IX. Preparation and Stability of Solutions" .
H. H. Millard and :Philena Young.
J. Am. Chem. 3°C., 21, 14°, (1929).

14. ("Applications of Cerio Sulfate in Volumetric
Analysis. V. Potentiometric Study of the
Reaction between Ferrocyanide and Cerio Ions").
N. Howell Furman and Oliver Evans.

J. Am. Chem. 300., 2;, 1128, (1929).

15. ("Applications of Cerio Sulfate in Volumetric
Analysis. VI. Oxidation of Hydrogen Peroxide
by Cerio Sulfate. Indirect Determination of
Lead").

N. Howell Furman and John H. wallace, Jr.
J. M. Chem. 8°C., 2, 1449, (1929).

16.

17.

18.

19.

20.

21.

18

("Cerio Sulfate as a Volumetric Oxidizing Agent.
X. The Determination of Thallium").

H. A. Hillard and Philena Young.

J. Am. Chem. Soo.,'§§, 56, (1950).

("Cerio Sulfate as a Volumetric Oxidizing Agent.
XI. The Oxidation of Organic Acids").

H. H. Hillard and Philena Young.

J. Am. Chem. 300., _5__2_, 1:52, (1950).

("Cerio Sulfate as a Volumetric Oxidising Agent.
XII. The Determination of Tellurous Acit").

H. H. Hillard and Philena Young.

J. m. Chem. 300., EEO _555, (1930).

(”Cerio Sulfate as a Volumetric Oxidising Agent.
XIII. The Determination of Mercurous Mercury)

H. H. Hillard and Philena Young.

J. Am. Chem. 300., _5__2_’ 557, (1930).

("The Use of Iodine Monochloride End Point in
Volumetric Analysis. 11. The Titration of
Arsenious Acid with Permanganate and with Cerio
Sulfate").

Ernest H. Swift and Carter H. Gregory.

J. Am. Chem. $00. 2%; 401, (1950).

("Applications of Cerio Sulfate in Volumetric
Analysis. VII. The Oxidation of Hydroquinore
by Cerio Sulfate”).

N. Howell Eurman and John H. Wallace, Jr.

J. Am. Chem. 800., 2;, 1443, (1930).

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