University of Southern California Dissertations and Theses

A study of the possibilities of volumetric microchemical analysis

(USC Thesis Other)

A study of the possibilities of volumetric microchemical analysis

PDF

Download

Open document

Flip pages

Download a page range

Download transcript

Copy asset link

Request this asset

Transcript (if available)

Content A Study of the Possibilities
of Yolumetric Microchemical Analysis
A Thesis
Presented to
The Faculty of the Department of Chemistry
University of Southern California
In Partial Fulfillment
of the Requirements for the Degree
Master of Science
fey
William Charles AtBrinson
June 1939
UMI Number: EP41512
All rights reserved
INFORMATION TO ALL USERS
The quality of this reproduction is. dependent upon the quality of the copy submitted,
In the unlikely event that the author did not send a complete manuscript
and there are.missing pages, these will be noted. Also, if material had to be removed,
a note will indicate the deletion.
UMI EP41512
Published by ProQuest LLC (2014). Copyright in the Dissertation held by the Author.
Microform Edition © ProQuest LLC.
All rights reserved. This work is protected against
unauthorized copying under Title 17, United States Code
ProQuest LLC.
7.89 East Eisenhower Parkway
P.O. Box 1346
Ann Arbor, Ml 48106-1346
< L *fo &Z1&
This thesis, w ritten by
William Charles Atkinson
under the direction of h.%3. F aculty Committee,
and a p p ro v e d by a ll its m em bers, has been
presented to and accepted by the C ouncil on
Graduate Study and Research in p a rtia l f u l f i ll
m ent o f the re q u ire m e n ts f o r the degree o f
Master of Science
D a te June 1939
Faculty Committee
Chairman
Secretary
TABLE OF CONTENTS
CHAPTER PAGE
I. INTRODUCTION
II* THEORETICAL CONSIDERATIONS
III. THE STANDARD THIOSULFATE SOLUTION
Preparation 8
Stability 9
Standardization Against Copper 10
IT. .. END-POINT SHARPENERS
Experimental Comparison 16
Msero-Volumet ric Method 18
Micro-Volumetric Method 19
V. ANALYSIS OF COPPER ORE
Procedure Recommended 22
Experimental Results 26
VI. SUMMARY
BIBLIOGRAPHY
CHAPTER I
INTRODUCTION
This investigation was carried out to ascertain the pos
sibilities of developing a procedure of volumetric micro-
analysis of copper and its ores.. The various phases of the
analysis were investigated , especially those having to do
with the end-point determination. Copper solutions cause
considerable difficulty in obtaining a stoichiometrical anal
ysis due to the fleeting end-point when titrated with sodium
thiosulfate solution.
The general method of analysis was that of allowing cop
per salt in a buffered solution to react with potassium
iodide under certain conditions. The iodine in the iodide
is displaced by the copper, and the former is subsequently
titrated with sodium thiosulfate solution of known copper
titre, This is known as a modification of the de Haen method
proposed early in the nineteenth century. The later proce
dures are often spoken of as the Low method for volumetric
determination of copper.
The procedure was studied in two phases; namely (1) the
macro method using approximately 0,1 normal thiosulfate as
the standard solution, and (2} a developed semi-micro method
using approximately 0,01 normal thiosulfate.
In order to obtain true titre values of the two thio
sulfate solutions electrolytic copper of the highest purity
was used as the primary standard. Small, samples ranging from
4. to 5Q milligrams were carefully weighed on a Christian
Becker Semi-Micro balance sensitive to 0.01 milligrams.
Weighings were direct to 0.0L milligram and a consider
ation of swings was taken. When the average of the swings
was eight or more divisions the hundreth milligram was in
creased or decreased one digit. This technique of weight
relationship was carried out in all the analytical proced
ures.
To facilitate the weighings the foil was cut into very
small pieces. By choosing a combination of the small pieces
of foil, the weights of copper could be reproduced within
one milligram. All weighings were made on a small tared
watch crystal using brass and aluminum weights.
CHAPTER II
THEORETICAL CONSIDERATIONS-
The general Ionic reaction may be given as:
2Cu:+* + 41“ *=* Ig +- Cu2I2
I2 + 2Na2S20-3 ' — » Ha2S406 + 2NaI
This has been proven to tafce place In stoichiometrieal pro
portions under the proper conditions by Low,-*- Hillebrand and
Lundell,^ and Rieman and Hues s.,3 The latter state that equi
librium condition is attained only if the reaction is below
a pH value of 4.5.
As this reaction is essentially one based on oxidation
potentials of the elements copper replacing iodine, Rieman
and Neuss give the following theoretical explanation:
Ex s 619 j-Z9 .6, log { *2 j ?
2 1“ Ig +- 2e
E0u== 150 * 59.1 log
2 Cu -h 2e 2Cu*
Or combined:
2 Cu+ + 51“ 2 CulJ. + Ig
^Low, Tech♦ . Methods of Ore Analysis., 78 (1927)
2
Hillebrand and Lundell, Applied Inorg. Analysis., 199
(1929)
rz
Rieman and Nuess, Quantitative Analysis, 189 (1937)
4.
The oxidation potential is higher for the triiodide ion, so
equilibrium would be expected to be shifted to the left.
However this anomality of equilibrium shift to the right is
found in the large concentration of iodide ion that should
be used. This drives the equilibrium to the right decreas
ing the concentration of' the free iodine triiodide ion, caus
ing a decrease of the potential of the iodine Ej. The shift
of the equilibrium to the right causes a decrease of the cop
per ion, thus causing an increase in the copper potential ECu.
The addition of a thiocyanate causes a further precipita
tion of CuSCN which has a very low solubility product. This
causes a further increase in the copper potential. Eoote and
Vance4 recognized this fact and they state that the addition
of ammonium thiocyanate causes the reaction to take place in
stoichiometrieal proportions of about 1:1600 which is well
within analytical accuracy.
Park5 states that in addition to the above potentials the
presence of the hydrogen ion catalyzes the reaction. This be
lief is corroborated by Whitehead and Miller^ who further state
4Foote and Vance, Jour, of Am. Chem. Soo., 57, 845 (1935)
5Park, B., Ind. and Eng. Chemistry, Anal. Ed., 3,77, (1931)
P \
Whitehead and Miller, Ind. and Eng. Chemistry, Anal. Ed.,
5, 15, (1933)
5
that considerable concentrations of the hydrogen ion have
little effect on the titration results. This is entirely
consistent with the explanation of Rieman and Nuess who state
a minimum hydrogen concentration must be present to have a
complete reaction between the copper and the iodide.
The speed of reaction is of some importance to obtain
practical results. Berthemot7 states that in measurements
of the speed of reaction of copper sulfate with concentrated
potassium iodide, the speed is directly proportional to the
concentration of the former and the square of the potassium
iodide. In support of the theory of excess iodide being
present Traube8 suggests that if not more than four moles of
potassium iodide are present for every two moles of copper
sulfate the precipitate will consist of a mixture of cupric
and cuprous iodides. However, Walker and Dover9 state that
the precipitate contains a complex poly-iodide chiefly as
Cul4.
Excess KI is recommended by Gooch^-8 to complete the re
action and to favor the formation of cuprous iodide, which
is shown to be true with the more concentrated solutions.
7Mellor, J. W., Treatise on Inorg. and Theo. Chemistry, ,
III, 201 (1923);
8Ibid>
9Loc. Cit.
10Gooch, Frank A., Methods in Chemical Analysis, 1st Id.,
120, (1912)
6
Gooch, and Heath.11 in addition state that the volume of
the liquid being titrated has a considerable effect on the
results. This is probably due to the dilution of the potas
sium iodide solution added, thus lowering the concentration.
They further state that the maximum amount of sulfuric or
hydrochloric acid in the solution should not be over two
milliliters in a solution of fifty milliliters. This would
cause an added liberation of iodine if the mineral acid con
tent were raised.
The decomposition of thiocyanogen may cause considerable
theoretical and practical error in analysis, especially if
the work; is of a high precision. This was stated by Rieman
and Nuess,12 and according to them the decomposition will
take place if ammonium thiocyanate is added too early in the
titration.
The reaction is:
2(CNS>~ (CHS) 2 + 2e
5 (CHS) 2 f 4HgQ 5HCNS f 4HCH + H2S04
This error is negligible if the major part of the iodine is
reduced before the addition of the thiocyanate.
Some of the commercial grades of potassium iodide of
rather low purity contain varying quantities of potassium
•^Gooch and Heath, Jm. Jour, of Science, 4, XXIV, 67
12Loc. Cit.
iodate. As the titrations ware carried out in an acid so
lution (pH 3.?} the possibility of the iodate ion reacting
with some of the free iodine liberated by the copper may
cause some error. The ionic reaction is:
IO3 f 51 * 6H* 3Ig f- SHgO
However the better grades of iodide were found to be prac
tically free from iodate.
CHAPTER III
THE STANDARD THIOSULFATE SOLUTION
An approximately O.Ol normal standard solution of sodium
thiosulfate was prepared by dissolving 2.5 grams (theoreti
cal 2.4819} of chemically pure Na2sg03.5Hg0 in a few milli
liters of recently boiled distilled water. A small amount
of sodium carbonate (50 mgm.) was added to the solution which
was made up to a liter volume with freshly boiled water. The
solution was allowed to stand for two weeks in the dark lab
oratory desk before standardization.
For the preparation of the stronger 0.1 Normal solution
the same procedure was used with the weight of the pure salt
being increased to between 24.8 and 24.9 grams.
The sodium carbonate tends to decrease the ionization of
the sodium thiosulfate and hence prevents the decomposition
of the solution and the liberation of free sulfur in the so
lution. Using this method no free sulfur was observed in any
of the standard solutions. Some of them stood as long as
several months.
Other preservatives for the standard thiosulfate solution
used are chiefly 1$ amyl alcohol as recommended by Mayr and
Kersohbaum13, and the use of borax to adjust the pH of the
^Mayr and Kersohbaum, Z. Anal. Chem. 73, 321 (1928)
solution between 9.0 and 9,5. This latter is the suggestion
of Hallet.14
THE STABILITY OF THE SOLUTION
Pure potassium dichromate was recrystallized from dis
tilled water, ground in a mortar and dried for four hours
in the laboratory oven at 110°C. The dried crystals were
then weighed out on the semi-micro balance and dissolved in
distilled water to obtain a solution of as nearly 0.01 nor^
mal as possible. Due to the difficulty of weighing the ex
act quantity, the normality was calculated.
Five milliliter portions of the standard dichromate so
lution were pipetted into Erlenmyer Flasks of 50 milliliter
capacity, one milliliter of 50$ acetic acid added and one
milliliter of 50$ potassium, iodide added. The solution was
allowed to stand for five minutes in a dark cool place and
titrated with sodium thiosulfate solution.
The dichromate solution had a normality of 0.009998. A
five milliliter sample liberated iodine that took 5.15 mil
liliter of the thiosulfate to titrate. The calculated titre
would then be:
0.009998 x 5.00 * N X 5.15
N « 0.00971
^Scott, Standard Methods of Analysis, 2498 (1939)
10
Similar titrations were carried out on the same solution
every day for ten days. No variation was found in the titre
of the sodium thiosulfate. This assumes that the dichromate
solution previously prepared was stable.
The buret used in this series of determinations was of
10 milliliter capacity graduated to 0,05 milliliter. Small
er divisions could be quite accurately estimated.
While the stability of the sodium thiosulfate was not
studied after a ten day period, it is entirely probable that
it would be stable much longer, especially if it contained
a trace of sodium carbonate or one of the other stabilizers
mentioned in an earlier section of this paper,
STANDARDIZATION AGAINST COPPER
The method,of titration of copper to determine the titre
of the sodium thiosulfate is essentially the one given by
Willard and Furman, - 1 * 5 Engelder*6 and Sutton* * 7 with some per
sonal modifications suitable for the volume and quantity of
*-5Willard and Furman, QuantItatlve Analysis. 196 (1933}
^6Engelder, Carl, Quantitative Analysis, 228 (1929)
17
Sutton, Volumetrio Analysis. XII, 220 (1935}
IX
reagents used.
The first series of titrations carried out gave very
fleeting end points and hence no concordance of the various
runs. In these the copper was dissolved in 25# nitric acid*
then evaporated until near dryness. Water was added, fol
lowed by acetic acid and ammonium hydroxide. With the quant
ities first used pH determinations were made. The value was
found to be approximately 4.7, which is entirely inconsistent
with the theory of the reaction.
Using the Beckmann p H Meter, an experimental buffer so
lution was prepared using the above reagents, and in addition
sulfuric acid was added to increase the active hydrogen ion
concentration until the value was 3.6 at room temperature
(23°C)* To attain this condition the weighed amount of cop
per (5-10 mgms of copper for 0.01 Normal thiosulfate) was
dissolved in nitric acid (25#), a few grains of talc added
as suggested by Peters^8, and evaporated to near dryness over
the mierobumer. The talc functions as "boiling chips" and
aids in smooth boiling of the solution thus preventing spat
tering. 0.5 milliliter of sulfuric acid (98#) was then add
ed and the solution allowed to cool to laboratory temperature.
About 0.8 milliliter of ammonium hydroxide (30#) was added.
This quantity of hydroxide was varied to give approximately
■^Woodman, Food Analysis. 258 (1931)
12
the same hue of blue color due to the formation of a complex
copper-ammonia ion. To the above solution* in a 50 milliliter
Erlenmyer Flask, 1 milliliter of 50$ potassium iodide was ad
ded. The iodide was tested for the presence of iodate accord
ing to the method given in the Pharmacopoea.After titra
tion the pH value had changed less than 0,1 in most cases.
The earlier titrations were carried out using a 3$ so
lution of potassium iodide. One milliliter of this was used
for about a 5 mgm* sample of copper. The end-point was ap
proached as determined by the straw color of the solution, and
0.5 milliliter of a 1$ starch solution was added. The titra
tion was then carried out in the usual way. These results were
found to fluctuate due to a fleeting end-point. The values ob
tained, expressed as each milliliter of thiosulfate equivalent
to milligrams of copper varied between 0.584 and 0.680.
A 3 milliliter solution of the 5Q$ iodide was next used.
The larger amount of iodide is in conformity with the theory
of Gooch,and Ferenkes and Koch.^ They suggest that approx-
13
Pharmacopoea of the United States, Eleventh Revision
S W ’OTSfcl------------------------ -----
20
Gooch, Frank A. Methods of Chemical Analysis. 1st Ed.
120 (1912)
21
Ferenkes and Koch, Journal of American Chemlcal Society
XXXIX, 1229
13
imately 1*5 grams of EX should be used for an 0.003 gram
sample of copper. Cantoni and Rosenstein^2 state that a
five-fold or larger amount of potassium iodide would cause
a quantitative reaction to take place. The values obtained
were somewhat better in this case, as the range was limited
between 0.620 and Q.680.
The probability of obtaining the correct titre of the
thiosulfate solution was not likely using this technique,
unless some method was used to determine the end-point more
exactly. Several methods have been suggested for the use of
end-point sharpeners. These were investigated and will be
evaluated*
22
Gooch, Frank A*, Loc. Cit.
CHAPTER IT
END-POINT SHARPENERS
The "fleeting end-points" obtained in titration of the
copper solutions necessarily oaused much concern* This pro
blem has been recognized by Foote and Vance, Park, Low, Rie
man and NTuess and others* Each investigated the problem and
found that various chemicals added near the apparent end
point served to sharpen the color change.
Perhaps the most classical choice of end-point sharpeners
is that of Rieman and Nuess* In this method they suggest the
use of ammonium thiocyanate near the apparent end-point to
cause further precipitation of the copper ton and lessen the
possibility of copper existing in solution to affect the ti
tration.
Theoretically, if the copper were entirely removed by
combined precipitation of cuprous iodide and copper thiocyanate,
the end-point could be determined in the usual way using a
weak starch solution. In this case the titration could be car
ried out to the characteristic sea-green color. However some
difficulty arises in actual laboratory practice as the copper
held in solution causes an off color and a false end-point.
The second method studied was that suggested by Park2®
15
using potassium acid phthalate near the end of the titration*
Using the same technique of preparing the buffer solutions
as stated previously, the titration was carried out until the
end-point was close* A few crystals of potassium acid phtha-
late was then added, and the titration was resumed to the color
change of the starch* A minimum time limit of two minutes was
arbitrarily set, and if no blue color returned within that
time this was called the true point of titration.
Low24 suggests the use of a soluble silver salt. Silver
nitrate in a weak concentration was used. Again the copper
titration was carried out in the same way until the end-point
was approached as shown by the early change of the starch so
lution. A small amount of silver nitrate was added CO.2 ml.
of 0.5 M. Solution} and the titration completed.
The silver nitrate reacts with some of the excess iodine
and causes a precipitate of white silver iodide. The white
color tends to neutralize the slight purple color of the
cuprous iodide, and aids in the determination of the end
point. As a very small amount of silver nitrate is added,
and as its combination Is rather with iodide ion than with
the free iodine, it is supposed to cause no error in the ti
tration.
24
Low, Albert H., Technical Methods of Ore Analysis.
10th Ed., 76 (19271
16
Friedrich Mohr25 suggested the use of sodium carbonate
in excess, and later suggested the use of ammonium carbon
ate to establish a more permanent end-point. This work was
done very early (1859) in the development of the iodide method
of titration and its reliability is to he questioned.
EXPERIMENTAL
Two series of titrations were carried out using the
three end-point sharpeners; namely a macro method using 0.1
normal sodium thiosulfate, and a semi-micro method using a
0.1 normal sodium thiosulfate. The titrations were essent
ially like the standardization of the thiosulfate solution
against copper, with the exception that near the end-point
the appropriate sharpener was used*
Both the ammonium thiocyanate and the potassium acid
phthalate were used as the salt, and the silver nitrate was
used in a weak solution. Using the macro method were the
total liquid volume was approximately sixty milliliters,
enough salt was added to be in excess of the weight of cop
per. This was approximately fifteen milligrams. A 0.2
25Washbura, E. “Theory and Practice of Iodimetric
Determination of Arsenous Acid.** Jour. Am. Chem. Soc.,
30, 31 (19081
17
milliliter portion of silver nitrate was used in each deter
mination. The silver nitrate was 0.5 Molar*
In the semi-micro method the volume at the end of the ti
tration was approximately eight to twelve milliliters* The
quantity of salt added was proportionally smaller being about
two to three milligrams. A weaker solution of silver nitrate
was used for the smaller volume determinations. The 0.5 mo
lar solution was diluted ten times with water and 0.2 milli
liter of this solution was used in the check on Lowrs method.
The buret used for the titration was of 10 milliliter
capacity and graduated to 0.05 milliliter. The graduations
were far enough apart to estimate approximately half of this
value so in most cases the readings were at least 0.03 milli
liters fine.
To further the degree of accuracy of the drops from the
buret the "split drop technique of titration” was used as
suggested by Hallet.26 This is essentially allowing a small
drop to collect on the tip of the buret, and before it has
enough weight to fall a small glass rod is touched to it and
mixed into the solution being titrated. The same results
might be obtained by drawing to tip out to a finer hole over
a very small hot flame. However this would slow the speed
of titration and it was not used for this reason.
S6Scott, Loo* Git.
Table I gives the results of a series of titrations a-
gainst known weights of copper using the various end-point
sharpeners* The copper was weighed on the semi-mi crobalanee,
dissolved in nitric acid* the buffer solution adjusted as
previously discussed* potassium iodide added and the liber
ated iodine was titrated using starch solution as the indi
cator. Near the end-point the sharpener was added. The
values are in terms of milligrams of copper equivalent to
each milliliter of approximately hundreth normal sodium thio
sulfate, The weights of copper taken varied between four and
eight milligrams.
TABLE 1
HERONS Sharpener AgNG3 Sharpener Phthalate Sharp.
4 0.649 8.658 ©.643
Q .642 0.661 0.639
0.644 0.659 0.670
0.645 0.679 0.634
0.651 0.663 0.641
0.643 0.662 0.651
0.640 0.658 0.634
0.647 0.660 0.632
0.645 0.664 0.645
0.647 0.661 0.645
0.643 0.663 0.641
0.647
0.648
Average Value:
0.645 0.662 0.643
Max. Deviation
from. Mean:
40.006 fO.017 fQ.027
It
Table II shows the average value obtained in a similar
determination using approximately 0.1 normal sodium thio
sulfate solution. In this case the weight of copper taken
varied between 50 and 52 milligrams. The same buret was us
ed and a similar method of titration.
TABLE II
NH^GHS Sharpener
6.5t
6 .33
6.39
6.40
6.59
6.41
6.51
6.48
Average Yalue:
6.47
Max. Deviation
from Mean:
AgHQjj Sharpener
6.61
6.65
6.63
6.69
6.67
6.61
6.62
6.66
6.64
Pthalate Sharp.
6.54
6.48
6.48
6.45
6.56
6 .56
6.51
6.50
6.52
6.51
tO. 12 f0.05 fO.05
20
In. table I, it will be seen that the highest titre of the
hundreth normal thiosulfate solution was obtained using an
end-point sharpener of silver nitrate as suggested by Low.
The titres obtained using ammonium thiocyanate and potassium
phthalate are in virtual agreement with each other, having
respective average values of 0.645 and 0.643. The use of po
tassium phthalate did not give as consistent results, as the
single determination averages vary considerably.
Standardizing the tenth-normal thiosulfate using the same
series of end-point sharpeners generally gave the same rela
tionships. It is shown that the highest titre value is again
obtained using the silver nitrate as an end-point sharpener.
The titres of the solution obtained using the ammonium thio
cyanate and potassium phthalate are however reversed to the
values obtained using the weaker thiosulfate solution. These
values are 6.4? and 6.51 respectively. They are in close
agreement with each other.
The essentially 10:1 ratio in the values of the two so
lutions is coincidental to the preparation of the two solu
tions, as each was prepared independently. However as they
were prepared with the same sodium thiosulfate (O.P.) weigh
ed on an analytical balance some similarity would be expect
ed. The value obtained with the tenth normal solution is
approximately ten times the value obtained with the weaker
21
thiosulfate solution.
The normality of the approximately hundreth-normal thio
sulfate solution was calculate to he 0.00971 (Page 9). Each
milliliter of this solution would then he equivalent to 0.611
milligrams CO.000611 Grams) of pure copper. This solution
was standardized against a known normality potassium dichro-
mate solution. However the titre of the sodium thiosulfate
obtained by direct titration against eleotrolytie copper was
upwards of 0.640 (Table I). This is entirely a different
value than that which should be theoretically obtained.
Peters*27 in his long and very complete study of the com
parison of the volumetric standardization of the conditions
of the iodine method of copper analysis, states that Bray and
MaeKay found a discrepancy in the values obtained when sodium
thiosulfate was standardized against pure copper foil and those
obtained in the standardization against a standard dichromate
solution. They further state the results were found to be con
cordant among themselves. Similar results were found in this
series of titrations.
P7
Peters. Amos. Journal of the American Chemical Society.
422 (1912)
zz
CHAPTER V
ANALYSIS OF COPPER ORE
Ore containing copper is at best a mixture of many other
minerals as iron, arsenic, lead, bismuth, etc. As these
metals have potentials considerably higher than that of
iodine they would tend to displace the iodine from the io
dide and cause the resulting titrations to be high.
The method of Low2® was adopted to remove the interfering
elements from a solution of ore allowing subsequent titra
tion with sodium thiosulfate solution. Essentially the ore
is taken into an acid solution to decompose all the elements
to simple chlorides, sulfates and nitrates. The solution is
then treated with pure aluminum metal. The aluminum has a
higher potential than copper and, by exerting an osmotic
ionic pressure, would tend to precipitate the copper from the
solution. An equivalent quantity of aluminum would then go
into solution. The free copper is treated in a buffered so
lution and titrated with sodium thiosulfate standardized
against known weights of copper.
28
Low, 0£. Cit. 79
23
The exact treatment of the copper ore was aa follows:
To 0*1 gram of ore contained in a small Erlenmyer
flask of 50 ml* capacity is added 4 ml* of hydrochlor
ic acid and 2 ml, of nitric acid, A few ml, of water
is added and the mixture is boiled over a direct mi-
orobumer flame to decompose the ore. Replacement of
the acids is made until decomposition is complete.
When all the salts are in solution, 0,8 ml, of con
centrated sulfuric acid is added and the solution is
again boiled until the sulfuric acid fumes strongly.
The solution is allowed to cool, a few milliliters
of water is added and it is again heated and filtered
while warm through #41 Watman filter paper to remove
any insoluble lead sulfate. As this filtration is
difficult, it is best to filter while warm or even
hot. The filter is washed with hot water several
times and the filtrate is collected in a beaker.
To the filtrate is added a small piece of pure
sheet aluminum approximately three quarters of an
inoh square and the solution is brought to boil
ing for ten minutes or longer* Low recommends ten
minutes but it was found better to boil for a long
er time to precipitate all the copper from the solu
tion. To test for complete precipitation a drop of
24
fresh hydrogen sulfide water is added. A clear,
or at most slight brownish color indicates that
the copper is precipitated. When precipitation
is complete the copper is filtered off and the
beaker is washed with hydrogen sulfide water.
The aluminum is also carefully washed. The hy
drogen sulfide acts as a precipitating agent for
the oopper and also to keep the metalic copper from
oxidizing to a soluble salt that would be lost in
such treatment. The last traces of copper are re
moved from the aluminum by the addition of 2 ml. of
nitric acid which are pWed through the filter. When
the washing is complete the original clean flask is
placed under the funnel and the paper is opened to
allow the precipitate to flow through. Warm water
may be used to wash the filter paper.
As the solution would contain arsenic and other
interfering elements, bromine water is added to the
solution until a straw color is produced. The brom
ine oxidizes the arsenic to its highest valence form.
A similar reaction takes place with the other inter
fering elements. The solution is then boiled to re
move the bromine and is evaporated to approximately
10 ml. The addition of talc was found to aid in smooth
boiling.
25
To the concentrated solution 2 ml* of acetic acid
(50$! are added followed by concentrated ammonium hy
droxide until a deep blue color is formed* It was
found that approximately 1 ml* of the hydroxide was
needed in most cases* To remove any excess ammonia
the solution is again brought to a quick boil and
cooled to laboratory temperature.
Potassium iodide (5Q$) was added until a clear io
dine colored solution was produced. 1 ml. of excess
was added* The iodine was immediately titrated with
sodium thiosulfate having a known copper titre*
When the straw color was approached in the titration
a drop of weak starch solution (1j200J was added and
again carefully titrated. Kear the end-point ammonium
thiocyanate crystals (approx. 1 mgm*J were added and
the titration resumed to the end-point.
26
EXPERIMENTAL
A low assay copper ore,, previously analyzed by other
methods, was analyzed according to the above method* The
previous analysis gave a copper value of 5*12$
Weight of Ore
in Grains
0.10074
0.10036
0.1004?
Q.10022
Ml. Thiosulfate Used
Titre 1 ml - 0.642
?.?4
7.42
7.60
7.43
$ Copper
Beterrained.
4.93
4.71
4.90
4.76
Average Value 4.725$
CHAPTER VI
SUMMARY
The possibilities of volumetric microanalysis as
applied to copper ore have been demonstrated. An analy
tical proceedure has been recommended for the analysis
of small samples of ore which may also be used to deter
mine the titre value of standard solutions of sodium
thiosulfate. The following conclusions may be drawn
from the analytical proceedures that have been carried
out in this investigation.
The presence of sodium carbonate in very small
concentrations is a good preservative for sodium thio
sulfate solutions of low normality.
The same titre of sodium thiosulfate was not
obtained when the solution was standardized against pot
assium dichromate solution of known normality as was ob
tained when standardized against pure electrolytic copper
foil.
Excess iodide ion must be present to promote a stol-
chiometrical reaction in the titration of the liberated io
dine.
Ammonium thiocyanate and bipotassium phthalate give
approximately the same titre values when used as end-point
sharpeners in the standardization of thiosulfate solutions
against pure electrolytic coppers Silver nitrate sharpen
er gives a considerably higeer value than either of the two*
28
Hundreth-normal sodium thiosulfate solution may be
successfully sed in the titration of copper solutions us
ing the iodide method.
The analysis of copper ores by a suggested semi-
micro volumetric method gives values that are somewhat
lower than those obtained by other methods.
The results show that when applying the microchem
ical method of analysis the figures may be a little low,
but they are definitely of the same order as those obtained
by the more cumbersome macro methods. The great adavntage
of the micro, or rather semi-micro method for the analysis
of prospect samples has been clearly demonstrated. The
whole apparatus and the needed supplies, exclusive of the
-balance,—could—be—packed—in~a—contain©r_not-Iarger than an
ordinary cigar box. In place of the somewhat large semi
micro balance actually used, a small portable assay balance
would serve equally well.
This method, then, could be readily used in the field
at great distances from the fully equipped chemical labor
atory.
BIBLIOGRAPHY
29
A. 0* A. C., Methods of Analysis. Committe on Editing. Wash
ington, (Fourth Edition), 1935 p. 511.
Bray and McKay. Journal American Chemical Society. 32:1193.
1910. * '
Engelder, Carl, Quantitative Analysis. John Wiley and Sons,
Hew York, p. W l i m j ------ '---
Ferenkes and Koch. Journal American Chemical Society. XXXII.
1229. ’
Foote and Vance, Ind. and Eng. Chemical Analysis Ed., 8:119,
(19361
, Journal American Chemical Society, 57:845, (1935)
Gooch, Frank A., Methods of Chemical Analysis, John Wiley
and Sons, Hew ‘ York, (First Edition), p. 120 (1912)
Gooch and Heath, American Journal Science, 4:24:67 *
Heath, C. L., The Analysis of Copper. Its Alloys and Ores,
(First Edition)
Heath, F. H., American Journal Science, 25:153
Hillebrand and Lundell, Applied Inor. Analysis, John Wiley
and Sons, New York, p. 199 (1929)
Kolthoff and Sandell, Volumetric Analysis. John Wiley and Sons,
New York, pp. 585, 60l (1936)
Low, A. H., Teohnical Methods of Ore Analysis, John Wiley and
Sons Hew York, p. 78 (1927T^
Mayr and Kersehhaum, Z. Anal. Cheat., 73, 321 (1928)
Mellor, J. W., Comp. Treatise on Inorganic and Theoretical
Chemistry. Longmans Green and Co.7 London lit:261 (1923)
Park, B., Ind. and Eng. Chemical Analysis Ed.. 3:77 (1931)
Peters, Amos W., "The Sources of Error and the Electrolytic
Standardization of the Conditions of the Iodine Method
of Copper Analysis.” Journal American Chemical Society,
422 (1912)
30
Pharmaoopoea of United States. (Eleventh Revision, p. 304 (1930)
Rieman and Huesa * Quantitative Analysis, Mac Gram and Hill. Hew
York, p. 18? (1937)
Scott. Standard Methods of Analysis. D. Van Kostrand Co.. New
YorF~p " 246'5' Tl'9^'9) '------
Sutton, Volumetric Analysis. P. Blaokistonrs Sons & Co.. Phtla-
delpKT^Txirr ZT&SSTTZW)
Treadwell-and Hall, Analytical Chemistry, John Wiley and Sons,
Hew York, pp. 59¥7
Washburn, Journal American Chemical Society, 30:31 (1908)
Whitehead and Miller. Ind. and Eng. Chemical Analysis id,,
5:15 (1933) "
Willard and Furman, Quantitative Analysis, D. Van Nostrand Co.,
Hew York, p. 196 (1933}
Woodman, Food Analysis, MacGfaw and Hill Publishers, Hew York,
p. 258 (1931)

Linked assets

University of Southern California Dissertations and Theses

Conceptually similar

PDF

An investigation of the application of 3-aminophthalhydrazide to the forensic detection of blood stains

PDF

A study of the quantitative spectrographic determination of beryllium

PDF

Critical comparison of the quantative methods for determination of alcohol in brain

PDF

Critical study of sources of error in recognized analytical chemical procedures

PDF

A study of the molybdate of cerium as a means of separation from the other rare earths

PDF

Comparative methods of casein extraction

PDF

A microchemical analysis of paraffin

PDF

Basic studies in the analytical chemistry of beryllium

PDF

A study of some sulfates and double sulfates of beryllium

PDF

A critical study of the accuracy of various weighing forms of calcium

PDF

Studies in the production of inulin

PDF

The extraction of titanium dioxide from its ores

PDF

An investigation of hydrogen overvoltage of alloys of chromium and nickel

PDF

The separation of lanthanum from other rare-earths

PDF

A study of methods for the determination of thallium in toxicological analysis

PDF

A rapid method for the determination of beryllium in ores

PDF

A study of the behavior of inulin in solution

PDF

A study of vitamin G in avocados

PDF

A study of some oxalates and double oxalates of beryllium

PDF

A study of the Mohr and Volhard methods of determination of chloride content in brines

Asset Metadata

Creator Atkinson, W. C (author)

Core Title A study of the possibilities of volumetric microchemical analysis

Degree Master of Science

Degree Program Chemistry

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag chemistry, analytical,OAI-PMH Harvest

Language English

Contributor Digitized by ProQuest (provenance)

Advisor Brinton, Paul H.M.P. (committee chair), [illegible] (committee member), Harrison, Bruce M. (committee member)

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c17-788100

Unique identifier UC11348147

Identifier EP41512.pdf (filename),usctheses-c17-788100 (legacy record id)

Legacy Identifier EP41512.pdf

Dmrecord 788100

Document Type Thesis

Rights Atkinson, W. C.

Type texts

Source University of Southern California (contributing entity), University of Southern California Dissertations and Theses (collection)

Access Conditions The author retains rights to his/her dissertation, thesis or other graduate work according to U.S. copyright law. Electronic access is being provided by the USC Libraries in agreement with the au...

Repository Name University of Southern California Digital Library

Repository Location USC Digital Library, University of Southern California, University Park Campus, Los Angeles, California 90089, USA