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1962 Studies of Analytical Separations. Darrell James Donaldson Louisiana State University and Agricultural & Mechanical College
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DONALDSON, Darrell James, 1935- STUDIES OF ANALYTICAL SEPARATIONS.
Louisiana State University, Ph.D., 1962 Chemistry, analytical
University Microfilms, Inc., Ann Arbor, Michigan STUDIES OF ANALYTICAL SEPARATIONS
A Dissertation
Submitted to the Graduate Faculty of the Louisiana State University and Agricultural and Mechanical College in partial fulfillment of the requirements for the degree of Doctor of Philosophy
in
The Department of Chemistry
by D arrell James Donaldson B.S., Loyola University, 1957 August, 1962 ACKNOWLEDGMENT
The writer wishes to express his sincere appreciation to
Dr. P h ilip W. West under whose d irectio n th is work was performed, for h is h elp fu l guidance during the experimental work and in the preparation of this manuscript.
The author would like also to express his thanks to
Dr. Maurice M. Vick for his many valuable suggestions and to the
National Science Foundation under whose sponsorship part of this work was performed.
i i TABLE OF CONTENTS PAGE
ACKNOWLEDGEMENT...... i i
ABSTRACT...... v i i
I. INTRODUCTION ...... 1
I I . EXPERIMENTAL...... 14
A. Solvent Extraction of Thiocyanate ...... 14
1. Reagents...... 14
2. Procedures...... 16
3. Factors Involved in the Extraction of Thiocyanate...... 19
4. Selectivity...... 24
5. Determination of Thiocyauate in the Organic Phase . 29
6. Recovery...... 29
7. Development of a Procedure for the Separation and Determination of Thiocyanate ...... 31
B. Separation and Detection of Antimony ...... 32
1. General Considerations ...... 32
2. Reagents...... 33
3. Procedures...... 33
4. Development of a Test for Antimony...... 39
C. Separation and Detection of Calcium, Strontium and B ariu m...... 40
1. General Considerations ...... 40
2. Reagents...... 41
3. Procedures ...... 41
4. Reaction of Calcium with Ferrocyanide ...... 43
5. Analysis Scheme...... 47
i i i PAGE
I I I . DISCUSSION...... 50
IV. SEUECTED BIBLIOGRAPHY...... 53
V. APPENDIX ...... 57
VI. VITA ...... 76
O
iv LIST OF TABLES
PAGE
I. Solvent Study...... 19
II. Equilibration Study ...... 20
III. Effect of Acidity on Thiocyanate Extraction...... 21
IV. Effect of Initial Concentration of Thiocyanate on its Extraction ...... 22
V. Effect of Different Mole Ratios ofZ%CS on the Extraction . 23
VI. Effect of TBP on Extraction of Thiocyanate...... 23
VII. Temperature Effect ...... 24
VIII. Effect of KOH Concentraion on the Recovery ;df Thiocyanate...... 30
IX. Resin S t u d y...... 34
X. Effect of Ammonium Chloride on Precipitation of Calcium . . 44
XI. Effect of Ammonium Chloride on Precipitation of Calcium . . 44
XII. Effect of Ferrocyanide Concentration on the PteidlpTtatlon of Calcium...... 45
XIII. Acidity...... 46
XIV. Titrimetric Standardization of Thiocyanate ...... 57
XV, Solvent Study...... 58
XVI. Effect of Acidity on Thiocyanate Extraction A. 0.1M N itric A c i d ...... 59
XVII. Effect of Acidity on Thiocyanate Extraction B. 0.5M N itric A c i d ...... 60
XVIII. Effect of Acidity on Thiocyanate Extraction C. 1M Nitric A cid ...... 61
XIX. Effect of Acidity on Thiocyanate Extraction D. 1.5M N itric A c i d ...... 62
XX. Effect of Initial Concentration of Thiocyanate on its Extraction A. 0,05 M/1 63 PAGE
XXI. Effect of Initial Concentration of Thiocyanate on Its Extraction B. 0.01 M /1 ...... 64
XXII. Effect of Initial Concentration of Thiocyanate on Its Extraction C. 0.0002 M /1 ...... 65
XXIII. Effect of Initial Concentration of Thiocyanate on Its Extraction D. 0.0001 M /1 ...... 66
XXIV. Temperature Effect A. 20°C ...... 67
XXV. Temperature Effect B. 40°C ...... 68
XXVI. Effect of KOH Concentration on Recovery of Thiocyanate A. 0 .5 M ...... 6 9
XXVII. Effect of KOH Concentration on Recovery of Thiocyanate B. 1M...... 70
XXVIII. Effect of KOH Concentration on Recovery of Thiocyanate C. 4M ...... 71
XXIX. Standardization of Calcium ...... 72
Graph No. 1: Spectrophotometric Calibration Curve for Thiocyanate...... 73
Graph No. 2: Spectrophotometric Calibration Curve for Z i n c ...... 74
Graph No. 3: Plot of Log TBP Concentration vs Log Distribution Ratio ...... 75
v i ABSTRACT
Solvent extraction has been extensively applied to the separation of cations but its application to anion separation has been for the most part ignored. The present study was undertaken to determine the conditions under which the extraction of thiocyanate was favorable.
Different solutions of thiocyanate were shaken with 1:1 tri-n-butyl phosphate (TBP) - methyl isobutyl ketone (MIBK), and after separation of the layers the aqueous phase was colorimetrically analyzed for thiocyanate.
Zinc was chosen as a coordinating metal ion to form extractable ion association species with thiocyanate. Addition of zinc in various concentrations was found to have no effect on the extraction. The distribution ratio rose with increasing nitric acid concentration until a constant value was reached at 1M. The extractability of thiocyanate was found to increase with decreasing temperature and with increasing thiocyanate concentration. Recovery and subsequent deter mination of thiocyanate was achieved by shaking the organic phase with potassium hydroxide after it had been treated with carbon tetrachloride.
Coextraction and selectivity of the system were studied with regard to eighteen anions. Molecular thiocyanic acid was responsible for the extraction. The extracting species was TBP.HNCS.
A procedure for the separation and determination of thiocyanate by solvent extraction was developed.
Improvements on the q u a lita tiv e an alysis scheme devised by West and Vick have been made. In the above mentioned scheme, bismuth was sometimes carried over with the antimony precipitate and a black
v i i p recip ita te formed when sodium su lfid e was added which hid the orange +3 antimony precipitate. A procedure, based on the oxidation of Sb to
Sb+"* by potassium dichromate, was developed for the detection of
antimony in the presence of bismuth.
Calcium and strontium were also found to be troublesome in the
analysis scheme of West and Vick. Sometimes these ions were not
completely separated by the alcohol-chromate treatment and thus gave
false tests for each other or were completely missed. The separation
of these ions was improved by the p recip ita tio n of calcium with potassium ferrocyanide in the presence of excess ammonium chloride.
Strontium was subsequently detected by the formation of a white
precipitate, strontium sulfate, when ammonium sulfate was added.
v i i i I . INTRODUCTION
It has been known for a long time that a solute will distribute
itself between two different liquids of slight mutual solubility in a manner related to their solubilities in the respective individual phases. Originally this phenomenon was of interest primarily to the
organic chemist. Although examples of extractability of inorganic com
pounds in non-aqueous solvents were reported from time to time in the
lite r a tu r e , only after 1940 did th is technique become commonly employed
in inorganic chemistry. In 1842, Piligot (28) noted the extraction of uranyl nitrate from aqueous solution into ether. Half a century later,
Rothe (30) recognized the possibility of extracting iron with ethyl
ether from a hydrochloric acid medium. Many more investigations could be cited but since no attempt w ill be made to give a comprehensive
coverage of the work in th is f ie ld , the reader is referred to the book by Morrison and Freiser (26).
With the advent of atomic energy programs in a number of countries,
interest in solvent extraction for inorganic substances was renewed on
a large scale. Today, solvent extractions are used in trace analysis,
production of high-purity metals, separation of chemically similar
elements and in radiochemistry for the production of carrier-free
tracers. Because of its simplicity, speed, clean separation and
applicability to analysis of complex mixtures, solvent extraction is
an important tool to the analytical chemist.
Solvent extraction has been extensively applied to the separation
and determination of cations. A scheme for the separation and micro-
1 identification of thirty-five metallic ions in a single unknown based on a combination of solvent extraction and the ring oven has been devised by West and Mukherji (41) . Other applications of cation ex traction systems have been discussed in a recent book by Morrison and
Freiser, "Solvent Extraction In Analytical Chemistry" (26) and in a review article in Analytical Chemistry (27).
The extraction of anions, however, has remained essentially unexplored. A few investigations of anion extractions as complexes or as acids are mentioned in the literature. Although in 1897, Tanret (36) studied the conditions under which nitric acid could be partitioned into ether, it was not until late 1950 that interest in this field was renewed. In 1954, Rynasiewicz, Sleeper and Ryan (32) developed a procedure for the extraction of boron in sodium metal with ethyl alcohol followed by its colorimetric determination with curcumin. Coursier,
Hure and Platzer (6) described a method for separating traces of boron by extraction of tetraphenylarsonium fluoborate with chloroform.
The distribution of nitric acid between water and various organic solvents has been widely studied. Alcock and co-workers (1) examined the partition of nitric acid between water and solvent mixtures of tri-n-butyl phosphate (TBP) with non-polar diluents such as kerosene.
They attributed the solubility of nitric acid in the organic solvents to the formation of the species HNO^.TBP. Rozen and Khurkhorina (31) calculated the thermodynamic functions of the system. Extractions of nitric acid with dibutyl ether (14, ^9), cyclohexanone and methyl isobutyl ketone (13), triheptylamine (43), diethylene glycol dibutyl ether (38), benzene and toluene (17) were studied. Very l i t t l e information on su lfa te extraction was found in the
literature. Ducret and Ratouis (9) described a method for the determination of traces of sulfate by exchange with thiocyanate on a resin and subsequent extraction of the ion association complex of the released thiocyanate and methylene blue. Boirie (3) examined the extraction of sulfuric acid using tri-n-octylamine under various conditions.
The efficiency of the extraction of phosphoric acid as the molybdate complex from aqueous solution into ethyl acetate was studied with regard to concentration of the two constituents, hydrogen ion concentration and solutions with excess iron and vanadium (12). The extraction into acetophenone of the molybdophosphate complex with
safranin was reported later in the literature (8). The difficult separation of phosphate from arsenate was achieved by Ross and
Hahn (29) using a solvent composed of twenty percent butyl alcohol
in chloroform.
Extraction as a means of purifying strong acids such as picric, perrhenic, thiocyanic, perchloric and chloroplatinic acids, was
reported by Tubalat (37) . The extraction of the perrhenate ion itself has been studied widely of late. Its extraction into TBP with
different acids has been reported by three different people (4, 19, 20).
For a long time long chain amines have been used for the extraction
of organic acids from lubricants, but recently they have been applied
to the extraction of inorganic acids such as hydrofluoric, nitric,
sulfuric and phosphoric acids (25) . The two most promising amines
found were tri-n-benzylamine and methyl di-n^octylamine in chloroform. 4
The extraction of the halides as their acids and complexes was reported in the literature. West and Lorica (40) described an extraction method for the separation of iodide using cadminum as the complexing agent. Efficiency up to 96% was reported. Bock (2) using triphenyltin hydroxide was able to extract into benzene such ions as iodide, chloride, bromide, phosphate, arsenate, chromate, selenate and vanadate. He was also able to separate chloride from sulfate in a ratio of 1:500, Moffett, Simmler and Patratz (24) described a procedure for the separation of fluoride from sulfate by extraction of the tetraphenylstibonium fluoride from aqueous solution into carbon tetrachloride. Iron and aluminum were found to inhibit the extraction. The mercuric salts of iodide, bromide and chloride were found to extract into tri-n-butyl phosphate (35), Extraction systems of some of their acids such as hydrochloric, perchloric and hydrobromic into tri-n-butyl phosphate were studied by Kertes and
Kertes (21, 22, 23),
The purpose of this brief literature survey was to give some insight into the present position of anion extraction into non-aqueous solvents. In order to have a better understanding of solvent extraction and some of its terms, a discussion of its principles is necessary,
1. EXTRACTION THEORY
Extraction theory is founded on two principles, Gibb's classical phase rule and Nernst's distribution law.
The phase rule relating the number of independent variables (F) needed to define a system at equilibrium, the number of components (C) and the number of phases (P) is expressed mathematically as:
F » C - P + 2 (1) 5
The system studied:involves essentially the distribution of a third component between two immiscible solvents, therefore, C » 3 and,F = 2. At constant temperature and pressure the system becomes univariant, F • 1. This means that the concentration of the third component is independently variable in but one of the two so lv en ts, the equilibrium concentration in the other phase being fixed by the concentration in the first phase. This relationship of a solute between two solvents is quantitatively expressed by the distribution law with the further restraint that this ratio is independent of the total concentration.
The basic equilibrium expression, assuming constant temperature and pressure, for the distribution of a solute between two immiscible solvents is given by the equation
/il =/ 2 ( 2) where^U 1® chemical potential and the subscripts refer to two different solvents. Applying the equation
JU =JX° + RT In m + RT In y (3) where jm ° represents the chemical potential of the hypothetical standard state m, the molality of the solute, y, the molal activity coefficient and R and T the gas constant and the absolute temperature, respectively, to equation (2) one has
o ° j + RT In m^ + RT In yj = + ^ m2 + ^ ^
Rearrangement of the equation above gives the following expression for the molal distribution coefficient,
V . " - S _ , * 1 exp •< K ' < 5 ) “ ml 02 If the solute does not significantly alter the mutual solubility of the solvent pair, the exponential term at constant temperature is a constant, K' and the equation reduces to (6 )
At infinite dilution, the phases approach ideality and the ratio of the activity coefficients approaches unity. In this case, becomes constant and equivalent to K* . Moderate variations in distribution coefficient with concentration are usually attributed to changes in the activity coefficient as shown by the preceeding equation.
Because of chemical interactions of the solute such as association, polymerization or reaction with the solvent which affect the distribution coefficient, a more easily obtainable quantity called the distribution ratio, D, is used to describe the extraction. It is mathematically defined as follows:
Total analytical concentration of solute in phase (1) ” Total analytical concentration of solute in phase (2)
For systems in which the distributing species has one and the same form in both phases, the distribution ratio, assuming Yj and approach unity becomes equivalent to the distribution coefficient.
The term percentage extracted, %E, because it provides a quick indication of successful separation of substances, is frequently used to characterize solvent extractions. This quantity is mathematically related to the distribution ratio as shown in the following derivation.
Let C = molar concentration
M = molecular weight of the solute
Gq= grams of solute in the organic phase Gw = grams of solute In aqueous phase
V = volume of the organic phase o = volume of the aqueous phase
Since percent is defined as the part divided by the whole times one hundred, one can write that „ (Weight of the solute extracted) ,Q0 0 (Weight of solute extracted + Weight of solute left) (7) or G %E = °+"g -- x 100 (8) o w
Dividing numerator and denominator by the quantity one obtains
GO^MV V
%E ’ (G0O + gJ/w W o V w * 1 0 0 <9 >
c /„ _ ° __ w______x 100 (10) Co/V + °w/v w
Division by C^, gives
(^o^V C ) x 100 (11) %E = \ w w / C°/c V + °w/c V ) \ w w w w /
Substitution of D for the quantity Cq/C, distribution ratio, gives When the volumes of the two phases are equal the equation reduces to
D x 100 (14) %E D + 1
The extractability of a species into organic solvents from aqueous media often can be predicted from it s s o lu b ility r a tio in the two solven ts,
To equate the two unconditionally can lead to serious error on account of two major fa cto rs. One arises from a change in a c tiv ity c o e ffic ie n t of the solute with concentration changes, while the other factor arises from the effect of the presence of the second solvent on the solubility of the solute in the first solvent.
It is generally true that a polar substance tends to dissolve in a polar solvent and a non-polar substance in a non-polar solvent.
Therefore, ionic compounds would not be expected to extract from aqueous solutions into non-polar organic solvents. A closer look at the process of dissolution is beneficial.
Using the Born-Haber cycle for the dissolution of an ionic halide,
MX , in water, one can write n MX (g) (g)
(aq) +nX“ (flq) in which I, is the energy absorbed for the process of dissolution, H is the solvation energy and U is the crystal or lattice energy. Thus, for the cycle one obtains
L = U + H (15) with solubility decreasing as U + H becomes more positive. The solvation energy increases with increasing dielectric constant of
the solvent. This general trend is qualitatively shown by Born's equation.
» - d L (i - f ) (16)
r ■ radius of ion
z = ionic charge
( = dielectric constant of the solvent whereas the lattice energy does not vary. Therefore, the extraction of ionic substances would be favored by solvents having high dielectric constant.
Since the solvents normally employed in solvent extraction are of
low dielectric constant, the charge of the ion must be neutralized prior to extraction. This is accomplished by ion association and many
times also by coordination and chelation.
A simple example is furnished by the anion of a weak acid in a
solution containing excess hydrogen ion such that the equilibrium is
shifted to species in which the overall charge is decreased. Since
the fraction of the extractable molecular species in the aqueous phase depends on the hydrogen ion concentration, the distribution ratio w ill show a dependence on the latter. The nature of this dependence is
easily shown by considering the following equilibria (5)
H+ + A “ > HA
v . (H + >
0 = - .. (HAw + AV (l9 )
Division by HA and substitution of the dissociation constant into J w equation (19) gives <«►> D - - (20)
When (IT,')< « the equation reduces to
K (H+) D = „ P , ,■ • (21) HA or
log D = -pH + pK^ + log Kp (22)
At high (H*), molecular HA is then the dominant species in both phases and a constant value of D is to be expected, D^K^.
CHELATION
If, for example, H Is a metal having a coordination number of four, its chelation may be represented by the following equation 11
An expression describing the extraction behavior of a metal chelate system can be derived by taking into account the formation of the metal chelate and the appropriate acid or base dissociation constant of the chelating agent.
ION ASSOCIATION
Ion a sso cia tio n complexes are formed by the union of oppositely charged ions to form ion pairs. Both cation and anion complexes may be illustrated by the following equations
(1) M+n + bB MB^n (cation complex) I + n MB, + nX%===» (MB, . nx") b b ✓
(2) M+n + (n + a) X «...... y:ViX.^+ g (anion complex)
MX~a . + aY+, : aY+, MX_a ) n + a n + a 2. STABILITY OF UNCHARGED COMPLEXES
The strength of the bond formed between the central ion and the
ligand coordinated to it is a measure of the stability of the complex.
Stability, therefore, depends on certain properties of the central ion and the ligand. With regard to the central ion, stability increases with ionic potential, depth of the available orbitals for coordination, oxidation state and electronegativity. In contrast, since there are so many different types of ligands, no definite rules can be formulated for
them. In general, factors which increase the basicity of the ligand
increase the strength of the complex. Steric effects, however, must be considered in complex formation. These factors are well illustrated for a ligand by consideration of a series of amines. Although in the series
NH^, RNF^* anc* t*ie number of methyl groups in creases, thereby
increasing basicity, the complex formation tendency decreases. This is 12 attributed to the steric hindrance caused by the R groups.
As previously shown, an ion can be extracted from aqueous solution into a non-aqueous medium if it can be transformed into a relatively unionized form as exists in the association of oppositely charged ions in ion pairs or clusters of higher order. The forces holding together the ion pair are coulombic; stability depending on how much stronger these forces are as compared to the solvation energy op the solvent.
The existence and behavior of such complexes was first proposed by Bjerrum and later confirmed by Fuoss and Kraus. The theory develops from the consideration of the factors that determine their formation.
The simplest model is assumed. The ions are assumed to be rigid unpolarizable spheres contained in a fixed dielectric constant medium.
Non-polar quantum bonds between ions as well as ion-solvent inter actions are excluded. ^
From these considerations, Bjerrum (18) arrived at the following expression for the formation constant, K,
= 4 71 N 1000 where
2
Y = r 6 KT
Z2 I *2 b = \h A € KT N = Avogrado's number K - Boltzman's constant Z-^abd Z2 = Ionic charge 13
e = electronic charge
£= dielectric constant
r = variable distance between ions
a = distance of closest approach
Fuoss (15) by applying the Boltzman method to a model in which the solvent is a continuum derived a simpler and more mathematically sounder equation.
K = 4 ^ N fl3 eb 3000 (24)
From this relationship, it can be seen that as the dielectric constant of the solvent rises the formation constant decreases. Therefore, ion pairs are stable to bombardment by solvent molecules in solvents of low dielectric constant such as tri-n-butyl phosphate.
The dependence of the formation constant of the ion pair on temperature is complex because the dielectric constant of the solvent is also temperature dependent. For solvents of high dielectric constant, the value of 6 decreases markedly with increasing temperature so that the value of ( t decreases. In such systems an increase in temperature favors a greater degree of ionic associaton. For solvents of low dielectric constant, increasing the temperature has little effect on £ ,
The quantity £t increases because of the rise in temperature (T), with the results that ion pair formation becomes less. I I , EXPERIMENTAL
A. SOLVENT EXTRACTION OF THICYANATE
REAGENTS
1. Sodium thiocyanate, 0.1 M
2. Nitric acid, 5 M
3. Potassium bromide, 0.1 M
4. Potassium iodide, 0.1 M
5. Sodium flu o rid e, 0.1 M
6. Potassium cyanide, 0.1 M
7. Sodium o x a la te, 0.1 M
8. Sodium pyrophosphate, 0,1 M
9. Sodium metaborate, 0.1 M
10. Sodium n i t r i t e , 0 .1 M
11. Sodium s u l f i t e , 0.1 M
12. Sodium th io su lfa te , 1 M, 0.1 M
13. Sodium monohydrogen phosphate, 0.1 M
14. Potassium ferrocyanide, 0.1 M
15. Potassium ferricyanide, 0.1 M
16, Sodium Hydroxide, 4 M, 1 M, 0.5 M
17. Silver nitrate, 0.1 M
18. Potassium permanganate, 0,01 M
19. Fluorescein (1% in alcohol)
20. Dichlorofluorescein (0,17. in alcohol)
21, Chromotropic acid (0.05% in conc. su lfu ric acid)
22. Iron III chloride, 0.1 M
to Curcumin (boiled alcohol filtrate diluted with water)
24. Iodine - sodium azide (1 gram of sodium azide per 100 mis. of 0.1 M iodine solution) 15
25. Sodium n itrop ru ssid e, 1%
26. Alpha - naphthylamine in glacial acetic acid
27. Benzidine (0.05 gram in 1:10 acetic acid solution)
28. Sulfanilic acid (1 gram per 100 mis. of 30 % acetic acid)
29. Barium ch lo rid e, 0 .1 M
30. Arsenic trioxide, 0.2%
31. Malachite green, 2% (boiled water filtrate mixed with glacial acetic acid)
32. Acetic acid - hydrogen peroxide (2 parts 6% hydrogen peroxide to 1 part acetic acid)
33. Bromine (saturated water solu tion)
34. Benzidine hydrochloride
35. Zinc, 0.1 M
36. Sulfosalicylic acid (0.3% in alcohol)
37. Dithizone (0.0047, in carbon tetrachloride)
38. Starch, 0.257,
39. Sulfuric acid (concentrated)
40. Glacial acetic acid
41. Pyridine
42. Ethyl alcohol, 95%
43. Ethyl acetate
44. Methyl isobutyl ketone (MIBK)
45. Tri-n-butyl phosphate (TBP)
46. Carbon tetrachloride
47. Kerosene 16.
2. PROCEDURES
a. PREPARATION OF SOLUTIONS
1. Sodium Thiocyanate, 0.1 M
Approximately 8 grams of sodium thiocyanate were dissolved
in a liter of distilled water. Suitable aliquots of this solution were used for various amounts of thiocyanate. Standardization was made by the tritrimetric silver nitrate procedure.
2. Silver Nitrate, 0.1 M
Approximately 17 grams of silver nitrate were dissolved in a liter of distilled water. The solution was thoroughly shaken, stored
in a brown bottle and subsequently standardized using sodium chloride as a primary standard.
3. Standard Zinc Solution, 0.1 M
6.5380 grams of reagent grade 20 - mesh zinc were dissolved
in a dilute nitric acid solution and then diluted with distilled water
to the mark of a liter volumetric flask. Suitable aliquot samples were diluted to prepare the various zinc solutions used in the course of
these experiments.
4. Acetate Buffer
Equal volumes of 2M sodium acetate and 2N acetic acid were mixed and shaken with a 0.017o carbon tetrachloride dithizone solution until
the color of the carbon tetrachloride layer remained unchanged. The
so lu tio n then was filte r e d to remove droplets of carbon tetrach lorid e.
5. TBP - MIBK (1:1)
Equal volumes of reagent grade methyl isobutyl ketone and tri- n-butyl phosphate were mixed. 17
h. STANDARDIZATION OF SILVER NITRATE
About 1.5 grams of pure sodium ch loride, dried for an hour a£
110°C, were weighed accurately into a 250-ml. volumetric flask.
Twenty-five mis. samples were pipetted into a 250-ml. flask.
Twenty-five mis. of water, 5 drops of 0.1% alcoholic dichloro-
fluorescein indicator and 0.1 gram of dextrin then were added.
The solution was titrated with silver nitrate, swirling vigorously,; until the mixture had acquired a permanent pink color.
c. COMPARISON OF SODIUM THIOCYANATE AND STANDARD SILVER NITRATE
SOLUTIONS
A 25-mls. sample of sodium thiocyianate was pipetted into a 250- ml. flask. To it, 25 mis. of water, 5 drops of 0.1% alcoholic
dichlorofluorescein indicator and 0.1 gram of dextrin were added.
Standard s ilv e r n itr a te was measured from a buret into the solu tion with vigorous shaking. At the end point, the solution turned a
permanent pink.
d. SPECTROPHOTOMETRIC DETERMINATION OF THIOCYANATE (34)
Suitable aliquots of sodium thiocyanate solution containing from
0.5 to 3 micrograms were used. To a 2 mis. neutral or slightly acid
sample, 0.2 ml. of a saturated bromine solu tion was added to form
CNBr. The excess bromine was destroyed by the addition of 0,2 ml. of
2% arsenic trioxide solution. In a separate test tube, 3 mis. of a
constant boiling mixture of pyridine - water (made by distilling at
93°C a mixture of 57 mis. of pyridine and 43 mis. of water) and 0.6 ml. 18 of a 5% solution of benzidine hydrochloride in 2% hydrochloric acid were mixed. The solutions were mixed, and after allowance of ten minutes for the production of the red color, their optical density was read on a Beckman DU Spectrophotometer at a wavelength of 520 nju.
The color was stable for th irty minutes.
e . SPECTROPHOTOMETRIC DETERMINATION OF ZINC (33)
To 10 mis. of an acidic solution, containing from 6 to 24 micrograms, 5 mis. of acetate Suffer and 1.00 ml. of a 1M sodium thiosulfate solution were added. The resultant solution was shaken for two minutes with two five mis. portions of 0.004% dithizone solution in carbon tetrachloride. The aqueous layer was washed with carbon tetrachloride until the latter remained colorless. The com bined extracts were transferred to a 50 ml. light protected volumetric flask and diluted to the mark with carbon tetrachloride. The absorption of the zinc complex was read against a similar treated blank at 535 m/a with a Beckman DU Spectrophotometer.
f . EXTRACTION OF THIOCYANATE
Five m illiliters of sodium thiocyanate were pipetted into a
60-ml. separatory funnel. Five molar nitric acid was added to bring
the solution to the desired acidity. Sufficient water was added to make the volume ten m illiliters. To this solution 10 mis. of 1:1
TBP-MJBK were added and the mixture was shaken for about a minute.
The separatory funnel then was immersed in a constant temperature bath set at 30°C where it was allowed to come to equilibrium.
After the attainment of equilibrium, the two phases were separated.
The aqueous phase was diluted to 25 mis. in a volumetric flask. Suitable aliquots were analyzed for thiocyanate as described in section (II, 2d). 19
The non-aqueous layer was treated with 20 mis. of carbon tetra chloride and shaken vigorously for about a minute with 20 mis. of 1M potassium hydroxide to recover the thiocyanate. After equilibration, the phases were separated, and the aqueous phase was acidified with 2M nitric acid and diluted to the mark of a 500 ml. volumetric flask.
Appropriate samples were analyzed for thiocyanate as previously described.
3. FACTORS INVOLVED IN THE EXTRACTION OF THIOCYANATE
a. CHOICE OF SOLVENT
Since iron forms a blood red complex with thiocyanate, a very dilute solution of iron thiocyanate in an approximate 1M nitric acid solution was used to survey qualitatively the extractability of thiocyanate in ethers, alcohols, ketones and esters, ^fter a comparison of the merits of these solvents, it was decided to study quantitatively the extractive efficiency in pure tri-n-butyl phosphate (TBP) and in it s mixtures \jfith eth yl acetate and with methyl isobutyl ketone(MIBK) .
TABLE I
SOLVENT STUDY
Symbols sig n ifica n ce
A - 10 m is. of TBP
B - 10 m is. of 1:1 TBP - MIBK
C - 10 mis. of 1:1 TBP - Ethyl acetate
D - 20 mis. of 1:1 TBP - MIBK
E - 20 mis. of 1:1 TBP - Ethyl acetate 20
SAMPLE THIOCYANATE CONCENTRATION %E INITIAL LEFT EXTRACTED
A 535 jig. 14 ^ jg . 5 2 l^ jg . 97
B 535 20 515 96
C 535 28 507 95
D 535 9 526 98
E 535 14 521 97
A 1:1 TBP - MIBK mixture was selected because of its extractive
efficiency, its relative insensitivity to acidity changes and its
rapid separation. A volume of 10 mis. was chosen because, as w ill be seen later, the complete operation of extraction and recovery
could be achieved in the same separatory funnel,
b . EQUILIBRATION TIME
Percentage extracted was studied using different times for
equilibration, ranging from one minute to two hours. The results
are presented in the table below.
TABLE I I
EQUILIBRATION TIME
TIME THIOCYANATE CONCENTRATION 7oE INITIAL LEFT EXTRACTED 1 m in. 535 jig. 19 Jig. 516 jig. 96
1 /4 h r . 535 21 514 96
1/2 h r. 535 20 515 96
1 h r. 535 19 516 96
1 1/4 h r . 535 20 515 96
2 h r. 535 20 515 96 21
The system was found insensitive to time. %E determined after one minute was allowed for separation of layers was somewhat erratic.
Therefore, any time from a 1/4 hr. to 2 hrs. could be used but
1/2 hr. was chosen as a convenient standard.
c . ACIDITY
TABLE I I I
EFFECT OF ACIDITY ON THIOCYANATE EXTRACTION
NITRIC ACID THIOCYANATE CONCENTRATION %E CONCENTRATION INITIAL LEFT EXTRACTED
0.1 M 535 yig. 69 466 yjg. 87
0 .5 535 26 509 95
1.0 535 21 514 96
1.5 535 20 515 96
As the nitric acid concentration was increased from 0.1 M to
1.5 M the efficiency of the extraction was increased until a constant value was obtained. At nitric acid concentration above 2M there was a tendency for thiocyanate to decompose. This behavior depicted in the pfeceeding table was in accordance with the theory presented6 n page 9 .
d. EFFECT OF INITIAL CONCENTRATION OF THIOCYANATE ON THE EXTRACTIVE EFFICIENCY
The concentration of thiocyanate was varied from very dilute solutions containing about 50^ig. to solutions normally encountered in analysis. The data appearing in the succeeding table shows that almost complete separation of thiocyanate could be achieved with a single extraction for solutions as dilute as 0.0001M/1. The extractive efficience of thiocyanate was seen to rise slightly as the thiocyanate concentration was increased. 22
TABLE IV
EFFECT OF INITIAL CONCENTRATION OF THIOCYANATE ON ITS EXTRACTION
MOLAR CONCENTRATION M/1 0 .0 5 0.01 0.001 0.0002 0.0001
Initial (NCS ) 30 m gs. 6 m gs. 535 jxg. 107 jig. 54 jig.
LEFT 1.02 0 .2 3 20 5 3 .5
EXTRACTED 28.98 5 .7 7 515 102 5 1 .5
7oE 97 96 96 95 94
e . EFFECT OF ZINC CONCENTRATION ON THE EXTRACTION OF THIOCYANATE
In the separation of metallic ions an excess of anion is added to form an extractable species. The reverse process seemed to be a logical approach to the extraction of the anion system under discussion. Zinc was chosen becuase it forms a colorless complex with thiocyanate and because its distribution coefficient is somewhat insensitive to changes in thiocyanate concentration. It was first thought that by the addition of excess zinc an extractable species of zinc thiocyanate would be formed, favorable to its partition into the organic solvent. In the analysis, for zinc by the dithizone method, of the aqueous phase after extraction, it was found that essentially all the zinc remained. Various concentrations of zinc were used. The results are summarized in the following table. TABLE V
EFFECT OF DIFFERENT MOLE RATIOS OF Zn/NCS ON THE EXTRACTION
MOLE RATIO Zn/NCS 10/1 4 /1 2/1 1/1
INITIAL (NCS) 535 jig. 535 535 535
LEFT (NCS) 20 jig. 19 20 20
%E (NCS) 96 96 96 96
INITIAL (Zinc) 9 .8 m gs. 3 .9 1.9 1
%E (Z inc) Trace None None Noi
f . EFFECT OF TBP CONCENTRATION ON THE EXTRACTION OF THIOCYANATE
Solutions ranging from 1 to 50% TBP in an inert solvent such as
kerosene were investigated.
TABLE VI
EFFECT OF TBP ON EXTRACTION OF THIOCYANATE
7oTBP THIOCYANATE CONCENTRATION d ist r ib u t : INITIAL LEFT EXTRACTED RATIO
1 1152 jig. 710 jug. 442 jig. 0 .6 2
2 1152 500 652 1.30
3 1152 410 742 1.81
4 1152 330 822 2 .4 9
5 1152 280 872 3 .1 2
10 1152 167 985 5 .8 9
20 1152 88 1064 12.1
30 1152 60 1092 1 8 .2
50 1152 33 1119 3 3 .9
The distribution ratio increased as the concentration of TBP increased.
A plot of log TBP conc. vs log distribution ratio gave a straight line
with unit slope (Graph 3) . g . TEMPERATURE EFFECT
The. temperature was varied from 20 to 40°C. The extractive efficiency, assuming that the small changes noted are significant, was found to decrease as the temperature was increased. This effect
is in accordance with Bjerrum's equation. It predicts that for solvents of low dielectric constant as TBP, the percentage extracted falls off as the temperature rises because of a decrease in ion pair
form ation .
TABLE VII
TEMPERATURE EFFECT
TEMPERATURE THIOCYANATE CONCENTRATION 7„E INITIAL LEFT EXTRACTED
20°C 535 16 j i g. 5 1 9 ^ ig . 97
30 535 20 515 96
40 535 28 507 95
4 . SELECTIVITY
A qualitative study of the extractability of some of the common anions was conducted under the same conditions used in the
study of the thiocyanate system that is from a 1M nitric acid aqueous phase. The study covered the following anions:
1. Sulfide 10. Thiosulfate
2. Sulfite 11. Orthophosphate
3. Sulfate 12. Cyanide
4. Nitrate 13. Chloride
5. Nitrite 14. Bromide 6 . Ferrocyanide 15. Iod id e
7 . Ferricyanide 16. Borate
8. O xalate 17. F lu o rid e
9 . Pyrophosphate 18. A ceta te
a . PROCEDURES FOR DETECTION OF THE ANIONS
1. SULFIDE (1)
A drop of the alkaline extract was mixed on a spot plate with a drop of 1% sodium nitroprusside solution. A more or less intense violet color appeared depending on the amount of sulfide present.
2. Sulfite (11)
A drop of 2% malachite green solution was added to a drop of neutral test solution. The solution was decolorized in the presence of sulfite.
3. S u lfa te
A white precipitate formed upon the addition of a dilute barium chloride solution was the basis of this qualitative test.
4. Nitrate (11)
A drop of the recovered aqueous solution was mixed on a spot plate with one drop of a 0.05% chromotropic acid solution. A few drops of concentrated sulfuric acid then were added. The production of a yellow color indicated the presence of nitrate.
5 , N it r it e
The alkaline extract was made slightly acid with acetic acid.
A drop of this solution was mixed with a drop of sulfanilic acid solution followed by a drop of (X -naphthylamine solution. The appearance of a pink color indicated the presence of nitrite. 6 . Ferrocyanide
The recovered e x tr a c t was made s l i g h t l y a c id w ith 2M n i t r i c acid. The presence of ferrocyanide was indicated by the formation of a blue ring upon the addition of a drop of dilute iron III chloride solution to a drop of the test solution placed on filter p a p er.
7. Ferricyanide
The recovered aqueous phase was neutralized with 2M nitric acid. A drop of this solution was mixed on a spot plate with benzidine acetate solution. A blue precipitate or coloration appeared, depending on the ferricyanide content.
8. O xalate
The decolorizing action of a dilute solution of permanganate on an acidic oxalate solution was the basis of its detection.
9. Thiosulfate (11)
Thiosulfate was detected by the iodine-azide reaction. On mixing a drop of the test solution with a drop of iodine-azide reagent on a watch glass, decolorization with a vigorous development of bubbles ensued.
10, Orthophosphate (11)
The recovered a lk a lin e phase was made s l i g h t l y a c id w ith 2M
nitric acid. A drop of this solution was placed on filter paper followed by a drop of molybdate and a drop of benzidine solution.
Then the paper was held over an open bottle of ammonium hydroxide to neutralize the acid. A blue stain appeared, its intensity de pendent on the phosphate content. 11. Cyanide (11)
The presence of cyanide was detected by the complete clari fication of the brown color of a copper II sulfide suspension.
12. Chloride
The formation of a white precipitate upon the addition of a
0.1 M silver nitrate solution indicated the presence of chloride.
13. Bromide (11)
The fluorescein test was used to detect bromide. A drop of its neutral solution was placed on a filter paper, dried and spotted with 1:2 acetic acid - hydrogen peroxide solution. The spot was treated with a drop of 1% alcoholic fluorescein solution.
A red ring was developed on warming.
14. Iod id e
Iodide was detected by its oxidation to iodine. A drop of the acidified recovered aqueous phase was mixed on a spot plate with a drop of starch solution and a drop of 10% potassium n itrite solution. The appearance of a blue color indicated the presence o f io d id e .
15. Fluoride
The decolorizing action of fluoride or. a dilute solution of ferric III thiocyanate was used for its detection. 16. Borate (11)
A drop of the test solution, acidified with hydrochloric acid, was placed on curcuma paper and dried at 100°C. A red brown fleck, which turns blue to greenish - black on treatment with 1% sodium
hydroxide, indicated the presence of borate.
17. Pyrophosphate
This anion was detected by its decolorizing action on a dilute
solution of iorn III thiocyanate.
18. Acetate
The aqueous solution recovered from the organic phase was
neutralized with 2M nitric acid. To a drop of this neutral solution,
iron III chloride was added. The formation of the red-brown complex
of iron III acetate served as indication of the presence of acetate.
b . RESULTS
Varying degrees of extractability were encountered and may be
classified as follows:
1. Anions extracted with approximately the same efficiency as
thiocyanate: None
2. Anions extracted moderately (40-60%):
bromide, chloride, oxalate, acetate, nitrate and iodide
3. Anions extracted very poorly or not at all (0-10%)
fluoride, thiosulfate, sulfite, sulfate, nitrite,
ferrocyanide, sulfide, cyanide, borate, orthophosphate,
pyrophosphate and ferricyanide
No evidence of coextraction was found. 29
5 . DETERMINATION OF THIOCYANATE IN THE ORGANIC PHASE
The Aldrich's method based on the formation of CNBr which in
turns reacts with amines in pyridine solution to produce a red color was slightly modified. Sulfosalicylic acid, because of its solubility
in organic solvents as ethyl alcohol, replaced arsenic trioxide as a
means of destroying the excess bromine. A 0.2% alcoholic bromine
solution which is stable for short periods of time, was used to form
cyanogen bromide. The following procedure was used. To 2 mis. of the
organic phase containing 1 to 3 micrograms, 0,2 ml. of alcoholic bromine
solution was added, followed by the addition of 0.2 ml. of sulfosalicylic
acid solution. In a separate test tube, 3 mis. of constant boiling
pyridine - water solution and 0.6 ml. of 5% benzidine hydrochloride
solution in 2% hydrochloric acid were mixed. The two solutions were
mixed and allowed to stand twenty minutes for the production of the red
color. The optical density was read at 520 mp. on the Beckman DU
Spectrophotometer. Care must be taken to insure a minimum evaporation
of solvent. The color was stable for about twenty minutes. Taking into
consideration the error introduced by dilution of the sample followed
by use of only a few m is., the accuracy of the method was good, 2-3%.
6 . RECOVERY
Dilute base and acid solutions were found to be ineffective in
the recovery of thiocyanate from the organic phase. The reverse
technique of solvent extraction, wherein a solvent mixture is used to
increase extraction, was employed, A third solvent, carbon tetrachloride,
miscible with the organic phase but a very poor extractant of thiocya
nate, was used. The addition of twice as much carbon tetrachloride
as the volume of the TBP-MIBK layer proved effective. To the resultant 30 solution potassium hydroxide was added, forming a two phase system.
The non-aqueous phase, becsiaBe of the addition of carbon tetrachloride, became the more dense phase. The mixture was shaken vigorously for about a minute and then allowed to stand until the layers separated.
The cloudiness which appeared when carbon tetrachloride was added to the non-aqueous layer disappeared. The non-aqueous layer was drawn off and washed once with a few mis. of distilled water. The aqueous layer was tested for thiocyanate by the Aldrich's method.
This method of recovery proved to be very efficient.
a . EFFECT OF VARIOUS KOH CONCENTRATIONS ON THE RECOVERY OF THIOCYANATE
The effect of different potassium hydroxide concentrations on the recovery of thiocyanate was studied. The results are summarized in the table below.
TABLE V III
EFFECT OF KOH CONCENTRATION ON THE RECOVERY OF THIOCYANATE
KOH Solution Percent Recovered
20 mis. of 0.5M 93
20 m is . o f 1M 96
20 mis. of 4M 96
The data shows that a single extraction, using an alkaline solution o f 1M or greater, achieved essentially complete recovery of the thiocyanate. 7 . DEVELOPMENT OF A PROCEDURE FOR THE SEPARATION AND DETERMINATION
OF THIOCYANATE
From the results found in part II, the following procedure for the separation and determination of thiocyanate is recommended. The thiocyanate solution is pipetted into a 60 ml. separatory funnel.
Nitric acid and water are added to adjust the acid concentration in the range of 1 to 1.5M and the volume to 10 mis. The mixture then is shaken with 10 mis. of a 1:1 TBP-MIBK solution for about a minute.
Sufficient time is allowed for the layers to separate before the aqueous layer is drawn off. The recovery of the thiocyanate is achieved by treating the organic phase with twice its volume of carbon tetrachloride and 20 mis. of 1M potassium hydroxide. The mixture is shaken vigorously for about a minute. After separation of the layers, the aqueous layer is analyzed for thiocyanate by either titration with silver nitrate, if the concentration of thiocyanate is sufficiently high, or colorim etrically.
31 32
B. SEPARATION AND DETECTION OP ANTIMONY
1. GENERAL CONSIDERATIONS
In the basic benzoate scheme of analysis described by West and
Vicfr (42), the ions Al+^, Fe+^, Cr+^, Bi+^, Sn^ and Sb+^ are grouped together. They are precipitated at a pH of 4 by a mixture of sodium benzoate and ammonium benzoate in their respective forms, BiOCl,
Fe(OH)Bzg, A1(PH)Bz2, Cr(0H)Bz2> H^nO^ and SbOCl. The precipitate is then treated with nitric acid; bismuth, iron, aluminum and chromium are dissolved while antimony and tin remain precipitated,
Hydrochloric acid is added to dissolve the antimony and tin precipi tates. Antimony is detected by the formation of an orange precipitate o f Sb2S^ w ith Na2S, while tin is detected by its reduction and sub sequent formation of a white precipitate with mercury II chloride.
There is a tendency for bismuth ion not to dissolve completely in nitric acid but to remain precipitated along with antimony and tin; when sulfide is added to test for antimony a black precipitate forms which hides the orange antimony precipitate.
This study was conducted to find a test for antimony in the presence of bismuth and tin. Two different approaches were tried.
One approach involved the investigation of the feasibility of separation of bismuth, tin and antimony by ion exchange while the other involved the investigation for a reaction specific for antimony or for one that could be made specific by choice of proper masking agents or medium. 2. REAGENTS +3 1 . Bismuth nitrate (10 mgs. of Bi per ml.) +3 2. Antimony III chloride (10 mgs. of Sb per ml.) +4 3 . Tin IV chloride (10 mgs. of Sn per ml.)
4. Potassium dichromate (0.05M, 0.01M, 0.005M, 0.003M, 0.002M)
5 . Sodium sulfide, 5%
6 . Sodium fluoride (saturated solution)
7. Sodium bromide, 2M
8 . Sodium iodide, 1.5M
9 . Malonic acid, 3%
10. Pyrocatechol violet, 0.05%
11. Ferricyanide, 0.1M
12. Thiourea, 2M
PROCEDURES
a . ION EXCHANGE
Dawson and Magee (7) separated bismuth, antimony and tin on an anion exchange resin, Amberlite IRA-400 malonate form. The resin was prepared by placing Amberlite IRA-400 in a 3% malonic acid solution and allowing it to stand with intermittent stirring for about three hours. It was filtered and washed with distilled water until the washings were neutral to litmus. Excess bromide, iodide and fluoride were added to the test solution containing bismuth and antimony to form the anionic complex for exchange with the anion resin. Sodium sulfide solution, 5% was used to test for the cations.
33 34
1, FLUORIDE
In a depression of a spot plate, 4 drops of 5M nitric acid and
6 drops of a saturated sodium fluoride solution were added to a drop of the test solution. Approximately 100 resin beads were added to the resulting solution which after stirring about a minute, was allowed to stand for 15 - 20 minutes. The aqueous phase was drawn off with a small bore pipet and tested for the cation. The beads were washed several times with distilled water and dried before they were treated with sodium sulfide. The aqueous phase of the antimony test solution when treated with sodium sulfide gave an orange pre cipitate Sb2S2 > while that of bismuth remained clear. The beads, if bismuth was present, turned black when sodium sulfide was added, while if antimony was present they remained yellow. Various resins were used and the data presented in the table below.
TABLE IX
RESIN STUDY RESIN BEADS AQUEOUS PHASE Antimony Bismuth Antimony Bismuth
Dowex 2-8x Orange ppt. C lear No change Black 50-100 Mesh
Amberlite IRA-400 Orange ppt. Very slight No change Black Malonate form b la ck p p t.
Amberlite IRA-400 Orange ppt. Slightly No change Black b la ck p p t.
Amberlite IRA-45 Orange ppt. Black ppt. No change B lack
2 . BROMIDE
To a test solution about 5M in nitric acid, 6-7 drops of sodium bromide (2M) were added followed by the addition of about 100 beads. The solution was vigorously stirred and allowed to stand for approximately twenty minutes. Two exchange resins were tried,
Dowex 2 - 8X and Amberlite IRA - 400 malonate form. The aqueous solution when tested with sodium sulfide gave a black precipitate if bismuth was present and an orange precipitate if antimony was present. Therefore, antimony could not be detected in the presence of bismuth. The beads in the case of bismuth turned black when sulfide was added. Amberlite IRA - 400 malonate form was more e f f i c i e n t .
3. IODIDE
Bismuth and antimony in a solution of approximately 2M nitric acid and containing excess potassium iodide (1.5M) were absorbed on the resin, Amberlite IRA - 400 malonate form. Further study showed that tin did not exchange with the resin. If the beads, after washing with distilled water, were tested with pyrocatechol violet, they changed from yellow to blue - green if antimony was present. If bismuth was present the reaction was masked because of the orange color of bismuth tetraiodide imparted to the beads before treatment with pyrocatechol violet. The test was not sen sitive, only 200 jig. could be identified.
b . SPECIFIC REACTION FOR ANTIMONY
After consideration of the data presented on ion exchange, the method was eliminated because of its insensitivity, complexity and slowness of separation. Attention was then focused on finding
either a specific reaction or one which could be made specific by using masking agents or different media. Reactions with different inorganic anions euch as borate, silicate, carbonate, selenate, sul fate, phosphate, pyrophosphate, tungstate, chlorate, bromate, iodate, 36 molybdate, ferricyanide, ferrocyanide, chromate and dichroraate were
screened together with different masking agents such as EDTA, fluoride,
tartrate, oxalate and thiourea. Two reagents, ferricyanide and dich-
romate, showed promise and were examined further.
1. FERRICYANIDE
When ferricyanide was added to a solution of antimony or bismuth
a yellowish precipitate appeared. No precipitate appeared for tin.
Different masking agents were tried to prevent the precipitation of bismuth with ferricyanide. In the presence of fluoride or EDTA, no
precipitate formed for either bismuth or antimony. Thiourea was
found to mask the reaction of only bismuth but tin was then found to
react with ferricyanide to form a yellow precipitate. As the acidity
of the medium was raised no precipitate formed between antimony and
ferricyanide while that between tin and ferricyanide formed but at
a slower rate. The sensitivity of the tin reaction decreased from
60 ^ig. in a slightly acidic solution to 300jig. in a 3M hydrochloric
acid medium. Under the same acid condition, 3M hydrochloric acid,
0.2 mg. of antimony or bismuth gave no precipitate after standing
for over 1 1/2 hours. This reaction was discarded because it lacked
sensitivity and specificity.
2 . DICHROMATE
When chromate was added to either bismuth or tin a yellow
precipitate occurred, but if antimony was present the yellow color
of chromate changed to blue - green. In acid solution, the chromates
of both tin and bismuth were soluble and did not interfere with the
color change reaction of antimony. Since chromate in acid solution 37 is actually dichromate, the reaction of antimony, tin and bismuth with dichromate was studied with regard to acidity, concentration of metal and concentration of dichromate.
a . ACIDITY
In solutions slightly acidic to 6M hydrochloric acid, 175jig. o f antimony were required to produce the color change. Therefore, the sensitivity of the reaction was independent of acidity. The specificity was, as mentioned before, somewhat dependent on acidity.
b . SENSITIVITY
The term " s e n s it iv it y " has had many co m p letely d if f e r e n t m eanings.
It has been used to signify the smallest detectable quantity of material and to express the relationship between the quantity of detectable material and the amount of the solvent present; it has also been applied to the reagent instead of the reaction. In order to judge the sensitivity of a reaction both the identification limit and dilution limit must be expressed. The identification lim it, first proposed by Feigl (10), is defined as the smallest quantity of material expressed in micrograms that can be detected no matter what the volume
is for a particular reaction under specific conditions. The identi
fication lim it alone is not an adequate appraisal of the sensitivity of a reaction, for a reaction is obviously more sensitive if l^ig. of
substance can be detected in 10 mis. than in 1 ml. Therefore, a know
ledge of the volume of the solvent is important in defining the
sensitivity of a reaction. This quantity is incorporated in the di
lution lim it, defined as the reciprocal value of the degree of dilution 38
prevailing at a given identification lim it. Mathematically it is
expressed as
Volume of solution (in ml.) x 10^ Dilution limit = — ------Identification limit (inyig.)
Experimentally, it is determined by testing a given analytical reaction
against increasingly dilute solution until there is no detection of a
reaction. Therefore, in judging the sensitivity of a reaction, the
identification limit and dilution limit must be stated. The smaller
the identification limit and the larger the dilution lim it, the more
sensitive the reaction is.
The sensitivity of the reaction of dichromate and antimony was
found to be dependent on the concentration of dichromate. The data
shown below shows that as the concentration of dichromate is decreased,
the identification limit is decreased. The identification limit was
found to be 30 jig. and the dilution limit 10,000.
DICHROMATE AMOUNT OF ANTIMONY CONCENTRATION NEEDED FOR DECOLORIZATION
0.05M 1000 jag.
0 .0 1 125
0 .005 70
0 .0 0 3 3 50
0.0025 30 39
4 . DEVELOPMENT OF TEST FOR ANTIMONY
From the results in the preceeding section, the following procedure
for the detection of antimony is recommended. Since the precipitates of antimony and tin, SbOCl and H^SnO^ * in the basic benzoate group are dissolved in 6M hydrochloric acid and since the sensitivity of
the reaction is independent of acidity, the test with dichromate is conducted in 6M hydrochloric acid. One drop of 0.005M potassium
dichromate is added. A positive test for antimony, in the presence of bismuth and tin is indicated by the change of the orange color
of dichromate to a blue - green color depending on the quantity of antimony present. C. SEPARATION AND DETECTION OF CALCIUM, STRONTIUM AND BARIUM
1. GENERAL CONSIDERATIONS
Calcium, barium and strontium, together with magnesium and lead, are precipitated as their fluorides by the addition of ammonium
fluoride in the basic benzoate scheme of analysis devised by West and
Vick (42) . Lead fluoride alone is soluble in a mixture of potassium hydroxide - potassium carbonate, thus separating it from the other ions.
The fluoride precipitate is dissolved by a saturated boric acid solution
in a fairly high hydrochloric acid solution. Magnesium is separated
from barium, calcium and strontium by utilization of the solubility of
its carbonate in the presence of high ammonium concentration. The carbonates of barium, calcium and strontium are dissolved in acetic acid
and then treated with potassium chromate to precipitate barium as its chromate. The filtrate is made alkaline with ammonium hydroxide and more potassium chromate is added. Subsequent addition of ethyl alcohol
precipitates strontium as its chromate, thus separating it from calcium.
Calcium and strontium were found to be troublesome, sometimes not being completely separated by the alcohol - chromate treatment and thus
giving false tests for each other or being completely missed. It was,
therefore, decided to improve the separation of calcium from strontium.
The chemical properties of calcium, strontium and barium are similar
except for their reaction with chromate, oxalate and sulfate and also
their reaction with ferrocyanide in the presence of ammonium chloride.
The reaction of calcium with potassium ferrocyanide showed the most
promise and was further studied as a method of separating calcium
40 41 from strontium and subsequent detection of strontium.
2 . REAGENTS
1. Calcium nitrate, 0.1M
2 . Strontium nitrate, 0.1M
3. Barium nitrate, 0.1M
4 . Potassium ferrocyanide, 0.1M
5 . Ammonium s u l f a t e , 1M
6 . Sodium sulfate, 1M
7. Potassium sulfate, 1M
8. Potassium oxalate, 1M
9 . Potassium chromate, 1M
10. Potassium permanganate, 0.1M
3 . PROCEDURES
a . PREPARATION OF SOLUTION
1. Potassium Permanganate, 0.1M
Approximately 3.2 grams of potassium permanganate were dissolved in a liter of distilled water. The solution was heated to boiling and kept slightly below the boiling point for an hour.
After standing at room temperature for a day, the liquid was filtered through a sintered - glass filter crucible and stored in a brown bottle to protect it from the light. Standardization was made by titration with pure sodium oxalate. 42
2. Calcium Nitrate, 0.1M
About 23.6 grams of CaCNQ^^.^I^O were dissolved in a liter of distilled water. Suitable aliquots of this solution were used for various amounts of calcium. Standardization was achieved by precipi tation of the calcium with oxalate, followed by dissolution of the precipitate in sulfuric acid and subsequent titration against a standard potassium permanganate solution.
b . STANDARDIZATION OF POTASSIUM PERMANGANATE
About 0.2 gram of pure sodium oxalate, dried at 110°C was weighed accurately into a 250 - ml. flask. Sixty m illiliters of water and 15 mis. of 1:8 sulfuric acid were added. The solution was heated to 90°C and titrated slowly with potassium permanganate to the first faint permanent pink tinge, while swirling the solution constantly.
c . STANDARDIZATION OF CALICUM
A ten mis. sample of calcium containing from 0,05 to 0.09 gram was pipetted into a 600 - ml. beaker. The solution was diluted to 200 mis. with distilled water, and a few drops of methyl red indicator were added. Five m illiliters of concentrated hydrochloric acid were added in excess, followed by the addition of 50 mis. of a warmamaonium oxalate s o lu tio n co n ta in in g 3 grams o f (NH^) £^2^4 * ^ 0 • The s o lu tio n was h eated to 70 - 80°C and 6 M ammonium hydroxide was added by drops, with stirring until the color changed from red to yellow. The precipitate was digested at room temperature for an hour and then filtered by decantation, keeping as much as possible the precipitate in the beaker. The precipitate was washed several times with cold 0.01 M ammonium oxalate and then was dissolved by pouring 50 mis. of hot 3N hydrochloric acid through the filter paper. The filter paper was washed several times with very small volumes first with a very dilute hydrochloric acid solution
(0.1N) and then with water. The filtrates were combined and diluted to 200 mis. One gram of ammonium oxalate in 200 mis. of water together with a few drops of methyl red were added to the combined filtrates. The solution was heated, and calcium was precipitated as previously described. After washing the preci pitate with very small volumes of distilled water, it was treated w ith 1:8 sulfuric acid until all the precipitate was dissolved.
The solution was diluted with water until the sulfuric acid con centration was 1 to 1.5N, heated to 80°C and titrated slowly with standard permanganate to the first appearance of a faint permanent pink tinge.
4 . REACTION OF CALCIUM WITH FERROCYANIDE
The reaction of calcium with potassium ferrocyanide in the presence of excess ammonium chloride was studied under different conditions with regard to the percentage of calcium precipitated.
Acidity, concentration of ferrocyanide and quantity of ammonium chloride were varied.
Atomic absorption was used to measure the quantity of calcium remaining after precipitation. Very small quantities of calcium could be determined by this method. Essentially the decrease in intensity of the calcium line imparted by the calcium source due to the absorption by the atomized calcium sample is measured. A calibration curve for calcium was constructed. The data from atomic absorption was checked by precipitation of calcium as the oxalate, followed by dissolution of it in dilute sulfuric acid and subsequent titration with standard potassium permanganate. Good agreement was found between the two. 44
a . EFFECT OF AMMONIUM CHLORIDE ON THE PRECIPITATION OF CALCIUM
The percentage precipitation of calcium was studied with three different quantities of ammonium chloride, 0,25g, 0.5g and lg. In the solution in which 0.25g of ammonium chloride was added, no crystals remained on the bottom of the tube. The data for precipitation of calcium with 1 and 2 mis, of 0,1M ferrocyanide are summarized in the succeeding tables.
TABLE X
EFFECT OF AMMONIUM CHLORIDE ON PRECIPITATION OF CALCIUM
1 ml. of 0.1 M ferrocyanide
AMOUNT OF CALCIUM CONCENTRATION 7„Ppt. NH.C1 ADDED INITIAL LEFT 4 0 . 25g 6. 2mgs. 2500 60
0 .5 0 6.2 2000 68
1 6.2 1800 74
TABLE XI
EFFECT OF AMMONIUM CHLORIDE ON PRECIPITATION OF CALCIUM
2 mis. of 0 .1M ferrocyanide
AMOUNT OF CALCIUM CONCENTRATION %Ppt. NH.CL ADDED INITIAL LEFT 4 0 .2 5 g 6 . 2m gs. 2250 73
0 .5 6.2 425 94
1 6.2 500 92 45
From these results one can see that the percentage precipitated
Increased slightly with increasing amounts of ammonium chloride.
For the solution containing 0.25g of ammonium chloride, the increase in the percentage precipitated with double amounts of ferrocyanide was slight as compared to 0.5 and 1 gram samples of ammonium chloride. This effect could be explained by the fact that the solution was not sat urated with ammonium chloride.
b . EFFECT OF FERROCYANIDE CONCENTRATION ON THE PRECIPITATION OF CALCIUM
Concentrations of ferrocyanide, ranging from quantities below the theoretical amount needed to precipitate the calcium (assuming a 1:1 ratio of calcium to ferrocyanide) to solutions containing more than double the necessary amount, were examined.
TABLE XII
EFFECT OF FERROCYANIDE CONCENTRATION ON THE PRECIPITATION OF CALCIUM
MLS. OF 0.1M FERRO- CALCIUM CONCENTRATION % P p t. CYANIDE ADDED INITIAL IEFT
0 . 8m l. 6 . 2m gs. 3 . lm gs. 50
1 6.2 2000 y.%. 68
1 .5 6.2 625 90
2 6.2 425 94
3 6.2 300 95
Ammonium c h lo r id e = 0 .5 g
As seen in the preceeding table, increase of the ferrocyanide concent tration increased the extent of precipitation until at 2 mis. essentially all the calcium was precipitated. 46
c . EFFECT OF ACIDITY ON THE PRECIPITATION OF CALCIUM
Solutions of calcium varying from slightly acidic to 9N in acetic acid were treated with a constant amount of ferrocyanide and excess ammonium chloride to observe the effect, if any, on the completeness of the calcium precipitation.
TABLE X III
ACIDITY
ACIDITY CALCIUM CONCENTRATION 7o. Ppt. INITIAL LEFT
Slightly acid 6.2 m gs. 2000 / g . 68
6N HAc 6.2 1250 80
9N HAc 6.2 500 92
Ferrocyanide = 1 ml. of 0.1M Ammonium c h lo r id e = 0 .5 g Slightly acid 6.2 625 90
3N HAc 6.2 625 90
6N HAc 6.2 500 92
9N HAc 6.2 75 97
Ferrocyanide = 1.5 ml. of 0 .1M Ammonium c h lo r id e = 0 .5 g
As the acidity increased, the percentage precipitated increased, until at a concentration of 9N HAc essentially all the calcium was precipitated,
d. SENSITIVITY
Concentrations of calcium as low as 10jig . in a 0.1 ml. of solution could be detected. At a concentration of lO^ig. and acidity of 9N HAc the reaction was somewhat slow but if the acetic acid concentration was raised to 12 N or greater, the reaction was faster, precipitation occurr^ ing in a matter of seconds. The dilution lim it as defined on page (38) was 10,000, Strontium and barium in milligram concentration did not i n t e r f e r e .
5 . ANALYSIS SCHEME
Barium was separated from calcium and strontium by precipitation of its chromate from an acidic solution. There was a tendency when ferrocyanide was added to a solution containing chromate for it to acquire a brownish color of varying intensity which at times could cause confusion on the part of a student in the detection of the creamy white precipitate of calcium. With this mind, it was decided to separate the chromate from the metals. This could be easily achieved by precipitating their carbonates in an alkaline medium. Therefore, the solution was made alkaline with ammonium hydroxide and heated almost to boiling. Ammonium carbonate reagent was added, and after allowing it to stand for three minutes, the solution was centrifuged.
The carbonate precipitate was then dissolved in 9N HAc, after which ammonium chloride was added until there was a noticeable excess remaining at the bottom of the tube after shaking. To this solution, ferrocyanide was added to precipitate calcium. Hydrochloric, nitric and sulfuric acids of different concentrations were tried to dissolve the calcium precipitate. Hydrochloric acid had hardly any effect on the precipitate but sulfuric acid dissolved the precipitate. However,
in preparing the solution for the calcium test with oxalate, a white precipitate, presumably calcium sulfate, was formed. Nitric acid was found to dissolve the precipitate easily. Therefore, the pre cipitate was dissolved in a minimum volume of 6M nitric acid and
im m ediately made a lk a lin e w ith6 M ammonium hydroxide. Acidified solutions of ferrocyanide have a tendency upon standing or heating to turn blue - green or sometimes to form a blue precipitate resembl ing Prussian blue. The solution was acidified with 4M monochloro- acetic acid and then 10 drops in excess were added. Five drops of
0.2M (NH^^O^ were added and the solution was warmed very gently in a beaker of hot water. A white precipitate confirmed the presence of calcium.
To the solution remaining after the precipitation of calcium with ferrocyanide, 1M cadmium nitrate was added until no white precipitate of Cd^Fe(CN)g formed. To the filtrate, 5 drops of 1M ammonium sulfate was added. The solution was warmed on a hot water bath. A white precipitate, SrSO^, indicated the presence of stro n tiu m .
Solutions containing various combinations of barium* calcium and strontium from concentration of the individual metals ranging from 0.1M to 0.01M were run according to the scheme outlined be low. Perfect agreement was found. OUTLINE BaCOj) SrCOj> CaCO^
D isso lv e in 6M HAc. Add 2 drops in excess. Add 3 drops 3M NH^Ac and 2 drops o f 1H KLCrO, Centrifuge.
m . S o ln . BaCrO^ y e llo w
D isso lv e in 6M HC1. A green Make a lk a lin e w ith 15M NH4OH flame indicates Ba . Add using 1 drop in excess. 1 drop of 0.1M (NH4) 2S04 . A Heat almost to boiling, add 5 white ppt. indicates BaS04< drops of (NH4)2CO^ reagent. If a white ppt. dissolves Allow to stand 3 minutes and in 6M NaOH, no Ba i s centrifuge. Wash ppt. once p resen t with distilled water.
P p t. S o ln . jSrCO^jCaCO^
D isso lv e in 9M HAc. Add s o lid NH^Cl D iscarc until after shaking an excess remains at the bottom of the test tube, Add 10 drops of 0.1M ferrocyanide. Centrifuge and test for completeness of ppt. by adding another drop of ferrocyanide reagent. Centrifuge
S o ln . CaKNH4Fe(CN) 6
Add 6 drops of 1M Cd(NO^)2soln. D isso lv e in minimum volume o f 6M Test for completeness of ppt. nitric acid. Make alkaline with by adding another drop of 6M NH40H. Then make a c id ic w ith Cd(NO^)2 reagen t 4 M C1CH2C00H (monochloracetic acid). Add 10 drops in excess. Add 5 drops of 0.2M and warm very gently in a water bath. Stir vigorously. Centrifuge. The w h ite ppt. is CaCo0 . . 2 4 P p t. S o ln .
D iscard Add 5 drops of 1M (NH4) 2S04 and warm on a h ot w ater b a tn . The w h ite p p t. i s SrSO 4* I I I . DISCUSSION
Golub and Ivanchenko (16) studied the thiocyanate complexes of zinc by conductometrie and potentiometric techniques. The following species, Zn (CNS)2> Zn (CNS)+1, Zn (CNS)^1 and Zn (CNS)^2, were found in aqueous solution. A measure of the potential of the solution with increasing thiocyanate concentration showed the re lative amounts of the species present and also their stability.
The extraction of zinc by using thiocyanate (26) was found to be very efficient and practically constant at low thiocyanate con centrations. It was, therefore, decided to extract thiocyanate as one of its complexes with zinc. At an acidity of 1 - 1.5M nitric acid, zinc is not extracted but 96% of the thiocyanate is extracted.
Therefore, the extraction of thiocyanate is attributed to its acid,
HNCS.
The fundamental equilibrium for the extraction of HNCS is
+ NCS~ + qTBPQ^ = ± HNCS-qTBPo
Neglecting activity coefficients, one can write a mass action, K, as v _ (HNCS'qTBP)o (H+) ( n CS- ) (TBP)q w w o and a distribution ratio, D, as
_ (HNCS-qTBP)o (n c sT " w
= K (H+) (TBP)q w o For a constant aqueous phase condition and low TBP concentration, activity coefficients in the organic phase can be neglected and DJ** (TBP)q o 50 51
Therefore, the slope of the curve log D vs log TBP gives the solvation number (q). From graph 3, the slope of the curve is unity. Therefore, the solvation-number q is 1 and the extracting species is TBP,HNCS.
The slight increase of percentage of thiocyanate extracted with increasing thiocyanate may be accounted for by the mass action and ionic strength effect. Increase in ionic strength favors ion association, while the mass action effect favors the formation of the undissociated acid, HNCS.
Using equation (20) _ (H+) Kp
(lrf> + KHA the dissociation constant is found to be 0.50 for a 0.05M/1 thiocyanate solution. This is in agreement with the constant cal culated if 92% of the acid is considered ionized. Therefore, thiocyanic acid is a strong acid.
The theory of Bjerrum or Fuoss predicts that an increase in temperature w ill cause a decrease in ion association in the TBP-MIBK phase, while in the aqueous phase it w ill cause an increase in ion association. The data in table (VII) shows that the percentage ex tracted decreases with increasing temperature, Therefore, of the two opposing effects, the degree of ion association in TBP-MIBK phase determines the efficiency of extraction. Fuoss' equation (24) can be, written in the following form
log K = log C + where log C denotes the coefficient of the exponential term and A 2 the constant e /ak. For solvents of low dielectric constant as
TBP-MIBK, £ changes only slightly with temperature (T). Thus the increase or decrease of ion pair formation is caused mainly by the change in temperature. For a change in temperature of 10 de grees there w ill be only a small change in log K. This may account for the observed small differences in the percentage extracted of thiocyanate at the three different temperatures. SELECTED BIBLIOGRAPHY
1. Alcock, K. , Grimley, S. S., Healy, T. V., Kennedy, J. and McKay, H. A. C ., "The E x tr a c tio n o f N itr a te s by T r i-n -B u ty l Phosphate," Trans. Far. Soc., 52, 39 (1956).
2. Bock, R. and Burkhardt, P., "Augschlitteln von Anionen Mit NichtwMsserigen Lbsungen von Triphenylzinnhydroxyd," Angew. Chem. , J J , H h ( l9 6 l) .
3. Boirie, C., "Extraction of Sulfates by Long-Chain Ammines," Comm. Energie Atomic (France) Rappt. No. 1262 (i960).
k. Colton, R., "The Extraction of Perrhenic Acid with Tri-n-Butyl Phosphate," At. Energy Research Estab., Gt. B rit. Report. R3823 (1961).
5. Cotton, F ., Progress in Inorganic Chemistry. New York: Interscience Publishers, Inc., i960. II.
6. Coursier, J., Hure, J. and Platzer, R., "Separation of Traces of Boron by Extraction of Tetraphenylarsonium Fluoboratei" Anal. Chim. Acta, 13. 379 (1955)•
7. Dawson, J. and Magee, R, J., "The Determination of Antimony, Tin and Bismuth," Mikrochim. Acta, 330 ( 1958).
8. Ducret, L. and Drouillas, M., "Dosages En Analyse Min^rale Par Extraction a l ’Aide De Cations Colofes. III. Dosages Par Extraction a l ’Aide de Colorants Basiques. C. Dosage De Traces De Phosphate Par La Safranine," Anal. Chim. Acta, 21, 86 (1959)•
9 . Ducret, L. and Ratouis, M., "Dosages En Analyse Minerale Par Extraction a l ’Aide De Cations Colores. III. Dosages Par Extraction a 1’ Aide De Colorants Basiques. D. Dosage De Traces De Sulfate Par Le Blueu De Methylene," Anal. Chim. Acta, 21, 91 (1 9 5 9 ).
10. Feigl, F. Chemistry of Specific. Selective and Sensitive Reaction. New York: Academic Press, 19^9 •
11. . Spot Tests in Inorganic Analysis. New York: Elsevier Publishing Co., 1958.
53 54
12. Fisher, W., Rlidiger, P. and Abendroth, H. , "Uber Die Abtrennung Kleiner PhosphorsHurejnengen von Eisen and Vanadin Durch Verteilung und Durch Ionenaustausch," Anal. Chim. Acta, 13, 38 (1955)-
13. Fomin, V. V., Margunov, A. F. and Korobov, I. V., "Extraction of Nitric Acid with Cyclohexone and Methyl Isobutyl Ketone," Zhur. Neorg. Khim., 5 , 1846 ( 1966).
14. Fomin, V. V. and Maslora, R. N., "Extraction of Nitric Acid with Dibutyl Ether," Zhur. Neorg. Khim. , 6, 481 (1961).
15. Fuoss, R. M., "Ionic Association. III. The Equilibrium Between Ion Pairs and Free Ions," J. Am. Chem. Soc., 80, 5059 (1958).
16. Golub, 0. A. M. and Ivanchenko,, C. D. , "Study of Thiocyanate Complexes of Zinc in Solution," Zhur. Neorg. Khim., 333 (1958).
17. Hardy, C. J. Greenfield, B. F. and Scargill, D., "The Extraction of N itric Acid from Aqueous Solution by Organic Solvent: The Dimerisation of Nitric Acid in the Organic Solvent," J. Chem. Soc., 90 ( 1961) .
18. Harned, H. S. and Owen, B. B. The Physical Chemistry of Electror lytic Solution. Second Edition. New York: Rheinhold Publishing C o ., 1950.
19. Kertes, A. S. and Beck, A. J. "Solvent Extraction of Septivalent Rhenium. Part I, Heterogeneous Equilibria in the System, Aqueous Nitric Acid - Potassium Perrhenate - Tri-n-Butyl Phosphate," J . Chem. S o c ., 1921 ( I 9 6 I ) .
2 0 . ______., "Solvent Extraction of Septivalent Rhenium. Part II. Heterogeneous Equilibria in the System, Aqueous Nitric Acid - Potassium Perrhenate - Tri-iso-octylamine," J. Chem. S o c . , 1926 (1961) .
21. Kertes, A. S. and Kertes, V., "Solute-Solvent Interaction in the System Hydrochloric Acid-Water-Tri-n-Butyl Phosphate," J. Inorg. and N u c l. Chem. , 14, 104 ( i 960) .
2 2 . ______. , "Solvent Extraction of Mineral Acids. IV. Solute-Solvent Interaction in the System Perchloric Acid-Water- Tri-n-Butyl Phosphate," .J. Appl. Chem. (London), 10, 287 (i960) . 55
23. . > "Solvent Extraction of Mineral Acids. Part III. Solute-Solvent Interaction in the System Hydrobromic A cid -W ater-T ri-n -B u tyl Phosphate," Can. .J. Chem. , 3 8 , 612 ( i 960) .
2k. Moffett, K. D., Simmler, J. R. and Potratz, H, A., "Procedure for Solvent Extraction of Fluoride Ion From Aqueous Medium," Anal. Chem. , 2 8 , 1356 (1 9 5 6 ).
25. Moore, F. L ., "Long-Chain Amines, Versatile Acid Extractants," A n a l. Chem. . 2 9 . 1660 (1 9 5 7 ).
26. Morrison, G. H. and Freiser, H. Solvent Extraction in Analytical Chemistry. New York: J. Wiley and Sons, Inc., 1957.
27. Morrison, G. H, and Freiser, H., "Review of Fundamental Development of Analysis, Extraction," Anal. Chem. . 5k, 6kR ( I962) .
28. P iligot, E., "Recherches Sur 1*Uraniiimtf'Ann. Chim. Phys., jj, 7 (I8k2).
29. Ross, H. H. and Hahn, R. B., "A Study of the Separation of Phosphate Ion From Arsenate Ion by Solvent Extraction," Talanta, J, 276 (I96I).
30. Rothe, J. W., "The Action of Hydrochloric Acid on Iron in Presence o f E th er," Chem. News, 6 6 , 182 ( 1892) .
31. Rozen, A. and Khorkhorina, L. P., "On The Thermodynamics of Extraction of Tri-n-Butyl Phosphate," Zhur. Neorg. Khim. , 2, 1956 (1957).
32. Rynasiewicz, J., Sleeper, M. P. and Ryan, J. W., "Boron in Sodium Metal," Anal. Chem. , 26, 935 (195k).
33. Sandell, E. B. Colorimetric Determination of Traces of Metals. Second Edition. New York: Interscience Publishers, Inc., 1950. II, p. 619.
3k. Snell, F. D. and Snell, C. T. Colorimetric Methods of Analysis. Third Edition. New York: D. van Nostrand Co., Inc., 19^9- H> p. 78k.
35. Specker, H. and Kloppenburg, H. G., "Untersuchungen Zur Extraktion Von Q u eck -S ilb erh alogen id en Mit T rib u ty lp h o sp h a t," Z. A n a l. Chem. , 183, 81 (I96I). 56
3 6 . Tanret, M. C., "Action Des Acids Nitrique, Sulfurique, Chlorhydrique Et Phosphorique Etendus sur Les N itrates en Presence de 1’Ether," Bull. Soc. Chim. (France), 1J, k97 (1897).
37. Tubalat, S., "Extraction and Purification of Strong Acids," U. S. 2,855,29^, C. A., a . 106771 (1957).
3 8 . Vdovenko, V. M. and Alekseeva, N. A., "The Distribution of Nitric Acid Between Aqueous Solutions and Diethylene Glycol Dibutyl Ether," Radiochimiyfr. _1, k50 (1959) •
3 9 . Vdovenko, V. M. and Krivokhatskii, A. S., "Extraction of Nitric Acid by Dibutyl Ether," Radiochimiya, JL,5 ^ ^ (1959)* kO. West, P. W. and Lorica, A. S., "A Solvent Extraction Method for the Separation of Iodide," Anal. Chim. Acta, 25, 28 (I9 6 I). kl. West, P. W. and Mukherji, A. K., "Separation and Microidentification of Metallic Ions by Solvent Extraction and Ring Oven Techniques," Anal. Chem. , 3 1 , 9^7 (1959). k2. West, P. W. and Vick, M. M. Qualitative Analysis and Analytical Chemical Separations. Second Edition. NeW York: Macmillan Co., 1959.
^3 . Zakharov-Natsissov, 0. I. and Ochkin, A. V., "Extraction of Nitric Acid by Triheptylamine," Zhur. Neorg. Khim. , 6, 1956 (1961). IV . APPENDIX
TABLE XIV
TITRIMETRIC STANDARDIZATION OF THIOCYANATE
Standardization of silver nitrate with sodium chloride
I II III IV
Weight of sodium chloride 0.2288 0.2288 0.2456 0.2456
Volume of silver nitrate 36.05 36.02 38.58 38.65
Normality of silver nitrate 0.1085 0.1085 0.1086 0.1087
Average normality 0.1086
Average deviation 0.7%°
Comparison of silver nitrate and sodium thiocyanate solutions
I II III IV
Silver nitrate volume 23.92 18.98 19.01 19.04
Sodium thiocyanate volume 25.14 19.95 19.96 20.02
Normality sodium thiocyanate 0.1034 0,1033 0,1033 0.1033
Average normality 0.1033
5.99 grams of NaNCS/liter
Average deviation 0.5 Yoo
57 58
TABLE XV
SOLVENT STUDY
SOLVENT EXTRACTION EFFICIENCE
Ether n-butyl ether n-butyl alcohol
Isobutyl alcohol
Octyl alcohol
Tributyl citrate
Chloroform
Carbon tetrachloride
Ethyl acetate
Amyl a c e ta te
Methyl isobutyl ketone
Amyl a c e ta t e - Methyl isobutyl ketone
Tri-n-butyl phosphate
Meaning of Symbols
Poor +
Moderate ++
Good +++ TABLE XVI
EFFECT OF ACIDITY ON THIOCYANATE EXTRACTION A. O.IM NITRIC ACID
THIOCYANATE CONCENTRATION
INITIAL LEFT EXTRACTED
1 535 jag. 68 jig. 467 jig.
2 535 66 469
3 535 70 465
4 535 69 466
5 535 71 466
6 535 68 467
7. 535 69 466
Average 535 69 466
Temperature 30°C TABLE XV (II
EFFECT OF ACIDITY ON THIOCYANATE EXTRACTION
B. 0.5M NITRIC ACID
THIOCYANATE CONCENTRATION
INITIAL LEFT EXTRACTED
1 535 ^ig. 25^ig. 510 jig.
2 535 25 510
3 535 26 509
4 535 27 508
5 535 26 509
6 535 28 507
7 535 26 509
Average 535 26 509
Temperature 30°C 61
TABLE X V III
EFFECT OF ACIDITY ON THIOCYANATE EXTRACTION
Ci 1M NITRIC ACID
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 535 ^ig. 20yjg. 515 jig. 96
2 535 21 514 96
3 535 21 514 96
4 535 19 516 96
5 535 22 513 96
6 535 21 514 .9 6
7 535 20 515 96
Average 535 21 514 96
Temperature 30°C 62
TABLE XIX
EFFECT OF ACIDITY ON THIOCYANATE EXTRACTION
D. 1.5M NITRIC ACID
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 535 jig. 20 jig . 515 jig . 96
2 535 18 517 97
3 * 535 20 515 96
4 535 19 516 96
5 535 20 515 96
6 535 21 514 96
7 535 20 515 96
Average 535 20 515 96
Temperature 30°C TABLE XX
EFFECT OF INITIAL CONCENTRATION OF THIOCYANATE ON ITS EXTRACTION
A. 0 .0 5 M/1
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 30 m gs. 1 mg. 29 m gs. 97
2 30 1 .1 2 8 .9 96
3 30 1.05 1 28.95 97
4 30 0 .9 8 29.02 97
5 30 1.03 28.97 97
Average 30 1.02 28.98 97
Temperature 30°C
Nitric acid concentration 1.5M 64
TABLE XXI
EFFECT OF INITIAL CONCENTRATION OF THIOCYANATE ON ITS EXTRACTION
B. 0.01M /L
THIOCYANATE CONCENTRATION %E
INITIAL LAST EXTRACTED
1 6 mgs. 2 5 0 jjg. 5 .7 5 mgs. 96
2 6 200 5 .8 0 96
3 6 260 5 .7 4 96
4 6 300 5.70 95
5 6 160 5.84 97
Average 6 230 5.77 96
Temperature 30°C
Nitric acid concentration 1.5M TABLE XXII
EFFECT OF INITIAL CONCENTRATION OF THIOCYANATE ON ITS EXTRACTION
C. 0 .0 0 0 2 M/1
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 107 ^ig. 5 102 ^jg. 95
2 107 6 101 94
3 107 5 102 95
4 107 5 102 95
5 107 6 101 94
Average 107 102 95
Temperature 30°C
Nitric acid concentration 1.5M TABLE XXIII
EFFECT OF INITIAL CONCENTRATION OF THIOCYANATE ON ITS EXTRACTION
D. 0 .0 0 0 1 M/1
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 54 jig. 3 jxg. 51 jig. 95
2 54 4 50 92
3 54 2.5 51.5 95
4 54 3 51 92
5 54 4 50 95
Average 54 3;5 51.5 94
Temperature 30°C
Nitric acid concentration 1.5M 67
TABLE XXIV
TEMPERATURE EFFECT
A. 20°C
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 535 jig. 16 jig. 519 jig. 97
2 535 16 519 97
3 535 16 519 97
4 535 16 519 97
5 535 17 518 97
6 535 16 519 97
7 535 16 519 97
8 535 16 519 97
Average 535 16 519 97
Nitric acid concentration 1.5M 68
tabu : xxv
TEMPERATURE EFFECT
B. 40°C
THIOCYANATE CONCENTRATION %E
INITIAL LEFT EXTRACTED
1 535 jig. 26 jig. 509 jig. 95
2 535 25 510 95
3 535 28 507 95
4 535 26 509 95
5 535 31 504 94
6 535 26 509 95
7 535 -■35 500 93
8 535 30 505 95
Average 535 28 507 95
Nitric acid concentration 1.5M TABLE XXVI
EFFECT OF KOH CONCENTRATION ON RECOVERY OF THIOCYANATE
A. 0.5M
PRESENT FOUND RECOVERED
* 535 ^ig. 490 jag. 92 2 535 475 90
3 535 500 94
4 535 500 94
5 535 485 92
6 535 500 94
A verage 535 492 93 70
TABLE XXVII
EFFECT OF KOH CONCENTRATION ON RECOVERY OF THIOCYANATE
B. 1M
THIOCYANATE CONCENTRATION PERCENT
PRESENT FOUND RECOVERED
1 535 jig. 520 j i g . 97
2 535 510 95
3 535 515 96
4 535 515 96
5 535 505 94
6 535 515 95
A verage 535 513 96 71
tabu : x x v iii
EFFECT OF KOH CONCENTRATION ON RECOVERY OF THIOCYANATE
C. 4M
THIOCYANATE CONCENTRATION PERCENT
PRESENT FOUND RECOVERED
1 535 jig. 505 jig. 94
2 535 520 97
3 535 510 95
4 535 515 96
5 535 500 94
6 535 515 96
Average 535 511 96 72
TABLE XXIX
STANDARDIZATION OF CALCIUM
Standardization of potassium permanganate with sodium oxalate
I I I I I I
Weight of Na2C204 0.2015g 0.2108g 0.1930g
Volume of potassium 23.99 25.18 21.63 permanganate
Sodium oxalate titer 0,00840g 0.00846g 0.00837g of permanganate
Average Titer = 0.00841g Average Normality = 0.1254
Deviation = 4%#
Calcium titer of potassium permanganate 0.292 x 0,00841 0.00246g
Standardization of calcium
I I I I I I
Sample volume 10 mis. 10 mis. 10 mis.
Volume of potassium permanganate 25.55 25.05 25.31
Weight of calcium in sample 0.0628g 0.0623g 0.0616g
Average weight of calcium in 10 mis. sample = 0.0622g
Normality of calcium solution = 0.1554N
Deviation = 7 %o 73
0 .7
0.6
0 .5
0 .4 A
0 .3
0.2
0.1
Micrograms of Thiocyanate
Graph No. 1. Spectrophotometric Calibration Curve 74
0.6
0 .5
0 .4 A
0 .3
0.2
0.1
X 10 15 20 25 Micrograms of Zinc
Graph No. 2, Spectrophotometric Calibration Curve 75
.8
1.5
1.2
0 .9
0.6
0 .3
0
0 .4 0.8 1.2 1.6 2.0
Log TBP Concentration
Graph No. 3. Plot of log TBP vs log distribution ratio VITA
Darrell Donaldson was born at New Orleans, Louisiana on February
11, 1935. He received his elementary education in the parochial school system at New Orleans, Louisiana. He graduated from Jesuit High School in New Orleans in May of 1953. He entered Loyola University in New
Orleans in September of 1953 and received his B.S. degree w ith a major in chemistry in 1957.
In September, 1957, he entered the Graduate School of Louisiana
State University and is now a candidate for the degree of Doctor of
Philosophy.
76 EXAMINATION AND THESIS REPORT
Candidate: Darrell James Donaldson
Major Field: Chemistry
Title of Thesis: "Studies of Analytical Separations"
Approved:
Mfejor Professor and Chairman
Dean of the Graduate School
EXAMINING COMMITTEE
// / / I'( ( , x . .
j - % ' * r ^
Date of Examination: July 27; I962