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1961 Studies on the Extractability of Anions. Apolinar S. Lorica Louisiana State University and Agricultural & Mechanical College
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Recommended Citation Lorica, Apolinar S., "Studies on the Extractability of Anions." (1961). LSU Historical Dissertations and Theses. 674. https://digitalcommons.lsu.edu/gradschool_disstheses/674
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LORICA, Apolinar S., 1927- STUDIES ON THE EXTRACTABILJTY OF ANIONS.
Louisiana State University, Ph.D., 1961 Chem istry, analytic al
University Microfilms, Inc., Ann Arbor, Michigan STUDIES ON THE EXTRACTABILITY OF ANIONS
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 Apolinar S • Lorica B.S. Chem., University of the Philippines, 1954 June, 1961 ACKNOWLEDGMENT
The author wishes to express his gratitude to Dr. Philip W. West under whose guidance this work was conducted.
Discussions with Drs. R. V. Nauman, B. R. Sant, A. K.
Mukherji, Alberto Llacer, M. C. Day, and Joel Selbln proved very helpful; they are gratefully acknowledged.
This acknowledgment will not be complete without mention of the financial support from the Continental Oil Company and the Standard
Oil Company (ESSO), for which the author is very grateful.
11 TABLE OF CONTENTS
ACKNOWLEDGMENT...... il
ABSTRACT...... I/ll
GENERAL CONSIDERATIONS...... 1
EXPERIMENTAL...... 13
A. Qualitative Studies ...... 13
B. Study of a Specific System Extraction of Iodide ...... 28
1. Experimental Procedures ...... 29
2. Various Factors Affecting the Extraction of Iodide . 34
3. Sources of Error...... 39
4. The Hydrogen Ion Concentration ...... 39
5 . The Cadmium to Iodide Molar R atio ...... 40
6 . The Initial Concentration of Cadmium Iodide . . . . 41
7. The Effect of Temperature...... 42
8 . Selectivity ...... 42
DISCUSSION OF RESULTS...... 44
DEVELOPMENT OF A PROCEDURE FOR THE EXTRACTION
METHOD OF IODIDE DETERMINATION...... 49
APPENDIX...... 52
SELECTED BIBLIOGRAPHY...... 81
VITA......
Ill LIST OF TABLES
Page
I. Extraction of Anions - A Qualitative Survey ...... 25
II. Effect of Sodium Hydroxide Concentration on the Recovery of Iodide ...... 36
III. Effect of Sodium Hydroxide Concentration on the Titration of Iodide ...... 37
IV. Equilibration Time and Percentage Extraction ...... 38
V. Effect of Hydrogen Ion Concentration on the Extraction of Iodide ...... 39
VI. Effect of Initial Cd/1 Molar Ratio on the Extraction of Io d id e ...... 41
VII. Effect of Initial Concentration on the Extraction of Io d id e ...... 41
VIII. Effect of Temperature on the Extraction of Iodide .... 42
DC. Gravimetric Standardization of Cadmium Nitrate ...... 52
X. Potassium Iodate Solutions ...... 53
XI. Effect of Sodium Hydroxide Concentration on the Titration of Iodide:
A. B lan k ...... 54
B. 0.1 N Sodium Hydroxide ...... 55
C. IN Sodium Hydroxide ...... 56
D. 10% and 2% Sodium Hydroxide ...... 57
XII. Effect of Sodium Hydroxide Concentration on the Recovery of Iodide:
A. 0.1 N Sodium Hydroxide ...... 58
B. IN Sodium Hydroxide ...... 59
lv Page
C . 10% and 2% Sodium H y d ro x id e ...... 60
XIII. Recovery of Iodide with 1 N NaOH:
A. 20 mis. of 1 N NaOH ...... 61
B. 40 mis. of 1 N NaOH ...... 61
C. 60 mis. of 1 N NaOH ...... 62
XIV. Effect of Hydrogen Ion Concentration on the Extraction of Iodide:
A. pH - 0 .3 ...... 63
B. pH - 0 .6 ...... 64
C . pH - 1 ...... 65
D . pH - 2 ...... 6 6
E. pH - 3 ...... 67
F. pH - 6 ...... 6 8
XV. Effect of CdA Ratio on the Extraction of Iodide:
A. Cd/l - 1/1 ...... 69
B. Cd/l = 2 / 1 ...... 70
C . Cd/l = 4 / 1 ...... 71
XVI. Effect of Initial Concentration of Cdl 2 on th e Extraction of Iodide ;
A. 0.001 M Cadmium Iodide ...... 72
B. 0.003 M Cadmium Iodide ...... 73
C. 0.006 M Cadmium Iodide ...... 74
D. 0.01 M Cadmium Iodide ...... 75
E. 0.03 M Cadmium Iodide ...... 76
F. 0.1 M Cadmium Iodide ...... 77
v Page
XVII. Effect of Temperature on the Extraction of Iodide:
A. 2 5 ° C ...... 78
B. 4 0 ° C ...... 79
Graph No* I: Spectrophotometric Calibration Curve
for Cadmium ...... 80
vl ABSTRACT
The extractabillty of fourteen anions has been Investigated. The
procedure used In the preliminary qualitative survey was to add a slight
excess of certain cations to the solution of the anion In order to form weak electrolytes or lon-associatlon complexes that were extractable.
The three types of solvents used for the subsequent extractions were an
ether (diethyl ether), a ketone (methyl lsobutyl ketone), and an ester
(ethyl acetate). The following anions were extracted well by one or
more of the solvents: chloride, bromide, iodide, thlocyanate, cyanide,
and nitrate. Fluoride was only slightly extracted. Spot tests also
showed extraction of the following anions by diethyl ether: borate,
oxalate, orthophosphate, thlosulfate, sulfite, nitrite, and ferrocyanide.
The extraction of iodide with cadmium as coordinating cation was
chosen for detailed study. Extraction efficiency decreased with dilu
tion of the cadmium iodide solution, but down to 0.006 M, it was
still appreciable. Such factors as temperature, pH of the aqueous
phase, and the initial cadmium to iodide mole ratio had only a slight
effect. A 1:1 mixture of tributyl phosphate and methyl lsobutyl ketone
was used for the immiscible solvent.
The existence of the following extractable species has been
postulated: (2H+, C dlJ, (H+, Cdlp, Cdl2, (Cdl+, OH"), and
(Cdl+ , NO 3 ).
An experimental procedure for the determination of iodide by the
extraction technique was developed. vii I. GENERAL CONSIDERATIONS
Until recent times only the organic chemist* engaged in Isolating a reaction product in high purity* used solvent extraction as a basic tool. The insolubility of Inorganic compounds in organic solvents* the limited number of available solvents* and what Conant called "our slavish devotion to water" probably blocked serious consideration of the use of solvents in Inorganic analysis. There were early isolated
Instances* however* of the use of an immiscible solvent in inorganic analytical separations. In 1842 Morrison and Freiser (15) noted the use of diethyl ether in the purification of uranium for pitchblende by
Peligot and the discovery by Rothe half a century later* of the ex- tractability of iron (III) chloride from strongly acidic solutions.
Sherrill (2)* while studying the halides of mercury* observed the ex traction of mercury (11) bromide into benzene. Many more investiga tions over the years could be cited, but since no comprehensive historical survey is being attempted here, the interested reader is re ferred to the book by Morrison and Freiser (15). Suffice to say that re newed interest in solvent extraction came only recently* with the increasing demands of modern technology mainly responsible for the revival. Today, among other uses* organic solvents are used in the recovery of expensive nuclear fuels* in trace analysis* and in the separation of chemically similar entities. Solvent extraction* because of its simplicity* its applicability to trace analysis and the analysis of complex mixtures, and its speed, is currently one of several techniques frequently employed in analytical chemistry.
Non-aqueous solvents have found extensive application in the separation of cations. A scheme for the separation and micro identifi cation of metallic ions combining the merits of solvent extraction and the ring oven has already been devised by West and Mukherjl (22).
Cation extraction systems were systematically treated by Morrison and
Freiser in their book, "Solvent Extraction In Analytical Chemistry."
The extraction of anions, however, has been largely ignored.
Only a few investigations that used organic solvents to separate anions are mentioned in the literature. Rynaslewicz, Sleeper, and Ryan (16) developed a method for extracting with ethyl alcohol small amounts of boron in sodium metal. In 1955, Coursier, Hure, and Platzer ( 6 ) studied the distribution of tetraphenylarsonium fluoborate between water and chloroform, in the course of developing a method for the isolation and determination of mlcrogram quantities of boron.
Alcock, et aK (2) studied the distribution of nitric acid between water and solvent mixtures of tributyl phOBphate (TBP) and non-polar diluents like kerosene. They attributed the solubility of nitric acid in the solvent mixture to the formation of the species, HNOg.TBP. In ac cordance with the principle of "like dissolves like," solubility data showed that as the amount of trlbutyl phosphate in the solvent mixtures was Increased, the solubility of polar substances was enhanced.
Ducret and co-workers (7) took advantage of the solubility In organic solvents of complexes of some anions with organic compounds to
separate traces of sulfate and phosphate. Sulfate was exchanged with thiocyanate adsorbed on a resin and the association complex subse
quently formed between the released thiocyanate and added methylene
blue was extracted with 1 , 2 -dichloroethane. Solvent mixtures of
aromatic ketones, specifically, acetophenone, and volatile solvents
like o-dlchlorobenzene, were used to extract the association complex
of phosphate with safranlne.
A few terms introduced in the preceding paragraphs now necessi
tate a discussion of the principles of liquid-liquid extraction.
1. EXTRACTION THEORY
Extraction theory is founded on two outstanding principles of
chemical thermodynamics, namely, the celebrated phase rule of Gibbs
and the distribution law of Nemst.
The mathematical formulation of the phase rule is as follows:
F-C-P+2
where F. is the number of variables needed to describe the system uni
quely, Q_is the number of components and, P.is the number of phases.
At constant temperature and pressure, the equation reduces to
F - C - P
In the present study, the systems were composed of two phases,
and the temperature and pressure were held constant, so that the
equation further simplifies to F - C - 2
For systems of five components such es the ones used in this work (TBP, methyl isobutyl ketone, water, cadmium iodide, and sulfuric acid) three variables are needed for unique description. These may be the concen tration of cadmium iodide in one phase, the pH of the aqueous solution, and the volume ratio of the mixed solvents, TBP and methyl isobutyl ketone (MiBK) • Since the volume ratio of the mixed solvents was fixed at 1 : 1 , only a knowledge of the concentration of cadmium iodide in and the pH of the aqueous solution after equilibration is necessary.
At constant temperature and pressure, the equilibrium condition for a solute distributed between two immiscible solvents is given by the equation,
/ l * > * 2 or expanding these expressions,
ju^ + RT In mx + RT In Y j + RT In m 2 + RT In Y \ (2) where th e ^ is the chemical potential in phases 1. and Z j ^ is the chemical potential at the hypothetical standard state of one molal; m^ is the concentration in molality; Yj is the molal activity coeffi cient; R and X are the constant, and the absolute temperature, respectively. Rearrangement of the preceding equation gives
m Y Kd " ^ " T z exp‘{f 2 '/ ' i )/RT (3 ) The exponential term is constant provided the presence of the solute does not affect the mutual solubilities of the immiscible solvent pair. Equation (3) may then be written thus, m« Yy KD - JLIk (4 )
At Infinite dilution, the activity coefficients approach unity and the distribution coefficient, Kj j , becomes constant and equivalent to
Variation in the distribution coefficient is usually due to changes in the activity coefficients as shown by equation (4).
In practice, solutions of such concentration that equation (4) be comes Inapplicable are the rule rather than the exception. A quantity related to the distribution coefficient thus supersedes the latter in utility. It is called the distribution ratio, £., and its mathematical ex pression is as follows:
13 * Total concentration in organic phase Total concentration in aqueous phase
£ approaches Kj) at infinite dilution.
Some prefer to report the results of extraction studies in percentage of solute extracted, %E. because It gives a quick indication of the degree of success attainable in the Isolation of the compound with a given sol vent. The mathematical relation between the distribution ratio, £ , and the percentage extraction, %E. is shown by the following derivation:
Let GQ ■ grams of solute in the organic phase
Gw ■ grams of solute in the aqueous phase
M * molecular weight of the solute
VQ = volume of the organic phase Vw = volume of the aqueous phase C - molar concentration Now,
* E - ;w t. JS^&SStSfmfka a s m * 1 0 0 G° X 100 Go + Gw Division of the numerator and denominator by the quantity, MVQ Vw gives:
G q/M V0 VW “ (G0 + G w)/ m V0Vw * 100
“ (c ^ rrfe ^ r x 100 If we now divide the numerator and denominator by Cw , we obtain,
%E C(/ Cw^w ------x ioo (Co/CwVw) + (CW/C WV0) Substituting p for the quantity, Co/Cy,,
D/Vw %E - -— ————— -— x 100 (D/Vw) + l/V 0
« ------2 — x 1Q0 D + VW/V Q When the volume of the aqueous phase, Vw, Is equal to the volume of the organic phase, VQ, a much simpler equation results:
D %E ------x 100 D + 1 The distribution of a solute between two immiscible solvents, one of which as far as the present study is concerned will always be water, finds a simple explanation in the difference in its solubility In the two liquids. It Is not by any means the complete explanation be
cause there are other factors Involved: the effect of the presence of the second solvent on the solubility of the solute in the first, and the change in the activity coefficients, to name two principal ones.
Solution theory has not grown to the point where it can say with certainty what substance will dissolve in which liquid. However, there are a few generalisations which have been very useful. Worthy of first mention is the familiar principle of "like dissolves like" which adequately explains the solubility of metal chelates in organic solvents.
An Ionic salt, like sodium chloride, is known to be more soluble in a solvent of high dielectric constant like water than in another solvent of much lower dielectric constant; the explanation is the relatively smaller work required in the former to separate the oppositely charged ions of the salt. When the molecules of the polar solvent possess a strong enough attraction for the ions of solute to separate them, the solvent molecules also remain bonded to the individual ions. The energy made available by this process is the solvation energy, H_, which according to the Born equation (14),
r * radius of ion € “ ionic charge D * dielectric constant of solvent should increase with increase in the dielectric constant of the solvent.
Because the solvents used in solvent extraction are as a rule of low dielectric constant, it may be inferred from the aforementioned generalizations that extraction is dependent on the prior formation of an uncharged species soluble in the solvent. Such an entity may be formed by coordination, chelation, or ion association. These processes 8 are now illustrated by the following equations:
M+n + nR~ ------» MRn where ^ is an n-valent metal ion and is an anion.
CHELATION
i M+n + •(X), n B<
g .i8 an atom having basic character in the Lewis sense; 00p is usually a carbon skeleton.
If, for example, M*n has a coordination number of 4. and n is 2., the metal chelate will have the formula:
,B' ,B 1- (x): M. (50, B< 'B;
ION ASSOCIATION
The formation of charged coordination complexes precedes the actual association of ions. +n (a) M+n + bB MBb (cation complex)
MBb+n + n -> (MBb+n , n5C”) (ion pair)
(b) M+n + (n + a) X' (anion complex)
-a + aY+ (ion pair) MX n + a ■* (»Y+ ■ MXn * «>
2. STABILITY OF UNCHARGED COMPLEX
The stability of the uncharged complex is a very important factor
in its extractability. For a coordination complex, stability is synonymous with the strength of the bond between the central metal ion and the ligands coordinated to it. Several factors determine the strength of this bond.
Regarding the contribution of the metal ion, the strength of the coordinate bond Increases with the ionic potential, the depth of the orbitals avail able for coordinate bonding, the oxidation state, and the electronegativity.
On the other hand, for the ligand, since there are different kinds, no definite rules can be laid down. A ligand may be one of the following types (3): a) it may ha/e one or more free pair of electrons or b) it may have gi. bonding electrons instead. To the latter type belong ethylene and cyclopentadlenyl ion. The ligands of group (a) may or may not be able to form gi. bonds with the central metal ion depending on whether they have vacant orbitals or orbitals that can be vacated, or whether they have additional gi electrons that can be shared with vacant metal orbitals •
To group (a) belong fluoride (no vacant orbitals), iodide (has vacant orbitals), and hydroxyl (has additional pi electrons) ions.
Stability of halide complexes of small electropositive ions like iron decreases with Increasing atomic weight, i_. e,., with decreasing electronegativity of the halide. The stability of the halide complexes of cadmium and mercury (II), on the other hand, follows the reverse order, 1“ > Br” > Cl” > F“. Significantly, the first three halides have vacant orbitals in their valence shells available for gi. bonding with the £ electrons of the metal ions mentioned. Other factors, probably more significant than dd-pl bonding (3) may be involved. In the case of iodide, its greater polarlzability may contribute greatly to 10 the lon-induced dipole binding with the cadmium and mercury (II) ions.
More probably, the heat of hydration, that decreases from fluoride down to iodide may be large enough, as in fact it is for the fluoride, to cause the reversal in the order of stability of the complexes.
In solution, coordination usually results in charged complexes which must subsequently associate with oppositely charged ions to give species of zero net charge to be extractable. Intuitively, it is easy to see that since the bond holding the ion association complex is purely electrostatic, stability is a question of how much greater is the cqulombic attraction between the ions of the pair, compared with the average thermal energy of a solvent molecule.
The theory of ion pair formation was originally proposed by
Bjerrum in 1926. Bjerrum compared the ratio of the measured activity, a^, to the stoichiometric concentration, c, with the activity coefficient given for the same system by the Debye-Huckel theory and noted that the former was smaller. He attributed the decrease in activity to associa tion of oppositely charged ions into pairs and proceeded to calculate the fraction of electrolyte associated into pairs (10). The final equation is usually seen in this form:
.2
N* Avogadro's number € * ionic charge a ■ distance between centers of ion pair k ■ Boltzmann constant T - absolute temperature D * dielectric constant 11
A mathematically sounder derivation by Fuoss (10) resulted In the following simpler equation:
, _ 4lTN«3eb 3000 The symbols have the same meaning; £ is the base of natural logarithms.
The value of h>ls an index of the extent of Ion pair formation. From the mathematical definition, It is seen to represent the ratio of the electro-
c 2 static potential energy, of the two Ions In contact to the average thermal energy of a solvent molecule, ]&. When £ is 5.60 & and remains constant at 25°C ., |> is equal to -1^- . In water, b is approximately 1., which means that the ion pair is not very Btable. On the other hand, for a solvent like trlbutyl phosphate of dielectric constant 8 ,, b is approxi mately JL2. and the Ion pair Is stable to bombardment by solvent molecules.
Because an ion pair possesses a dipole field, it can be attracted by either anions or cations. In fact, triple ion clusters are expected to form in the dielectric constant range, 1 0 - 2 0 .
The effect of temperature is intimately connected with its effect on the dielectric constant. For solvents of high dielectric constant, an in crease in temperature occasions a relatively greater decrease In the di electric constant so that the net effect is a lowering in the value of the quantity, (DT). In terms of either equation, this means a greater degree of ion association. For solvents of low dielectric constant, on the other hand, increasing the temperature has very little effect on the dielectric constant. The quantity (DT) increases because of the temperature in crease and the extent of ion pair formation becomes less. 12
3 . PURPOSE OF THE PRESENT STUDY
The present study was undertaken as a means of establishing the possible usefulness of solvent extraction in the separation of anions.
Although the separation of anions (19, 20, 24) has always been a diffi cult problem, and although solvent extraction methods have become in creasingly popular in the separation of cations, the possibility of com- blnihg the two fields has remained Essentially unexplored.
In the separation of metal ions as ion association complexes or as simple coordination compounds, a large excess of anions is used to insure the formation of the higher complexes and create favorable con ditions for ion association. Thus, iron (III) is extracted with ethyl ether
(15) from aqueous solutions 6 .^ in hydrochloric acid. The reverse process, i_. e,., addition of a slight excess of catlonB to a solution of a given anion to produce extractable species, seemed to be the logical approach to the investigation of the extractability of anions •
The halides and pseudohalides had been extensively applied in the extraction of cations. In fact, except for the nitrate system,
Morrison and Freiser (15) gave extended tabulated summaries only for halide and pseudohalide ion association systems. We decided, there fore, to start the present work with them. If the results were encouraging, the investigation would be extended to the other anions • A detailed study of one system would then be undertaken after a successful exploratory qualitative survey. II. EXPERIMENTAL
A. QUALITATIVE STUDIES
Preliminary qualitative studies were deemed necessary to establish the practicability of using solvent extraction to separate anions. The cation extraction systems listed by Morrison and Freiser (15) furnished the starting points In the choice of the cation, the solvent, and the acidity of the solution.
All extractions, except in the case of thlosulfate, nitrite, and sulfite, were In acidic medium. The exceptions were ions that are un stable at low pH so that extractions had to be made from solutions that were either very slightly acidic, neutral, or very slightly alkaline.
The cations chosen were good coordinating centers and were readily available. They were silver (I), cadmium, cerium (IV), cobalt (II), chromium (III), mercury (II), zinc, Iron (III), gold (III), antimony (III), tin (IV), and uranium (VI).
Three types of solvents were employed, namely, an ether (diethyl ether), a ketone (methyl lsobutyl ketone), and an ester (ethyl acetate).
Occasionally, 1:1 mixtures of tributyl phosphate (TBP) with one of the aforementioned solvents were tried. Although TBP is a good solvent for many ion association complexes, it is not very convenient to use be cause of its viscous nature. In one instance, lsoamyl alcohol (the vapors of which are harmful to the eyes) was tried and found to give a very favorable distribution ratio for the anion (iron (III) thiocyanate
13 14 com plex).
To detect the anion, a few drops of the extract were subjected to the ring oven treatment (21) preliminary to the Identification with suitable color reagents. * For the chloride ion, because there was no color reac tion applicable, the residue obtained by evaporation of the extract was dissolved In water, and the resulting aqueous solution was tested for the ion with dilute sliver nitrate. Evaporation of the methyl lsobutyl ketone (b. p ., 115.8°C) extract was, however, time-consuming.
As long as sensitive color reactions were available for detecting the anion, and provided the anion complex was heat stable, the ring oven was used and stripping was not a problem.
The detection of the anion In the extract required repeated extrac tion of the aqueous layer with fresh solvent to obtain a rough estimate of the extraction efficiency.
REAGENTS USED
1. Hydrochloric acid, 6 N
2. Nitric acid, 6 N
3. Sulfuric acid, 6 N --
4. Potassium bromide, 0.1 N
5. Potassium iodide, 0.1 N
*In the ring oven technique, a drop of test solution is placed on the center of a filter paper set on top of an electrically heated metal ring. Addition of a suitable reagent then either fixes the Ion In question in the center or drives It to the periphery. The heated ring causes rapid evaporation of the spreading liquid and subsequent deposition of the Ion In the form of a ring. 15
6 . Sodium Fluoride, 0.1 N
7. Potassium thiocyanate, 0.1 N
8 . Potassium cyanide, 0.1 N
9. Sodium oxalate 0.1 N
10. Sodium metaborate, 0.1 N
11. Sodium nitrite, 0.1 N
12. Sodium sulfite, 0.1 N
13. Sodium thlosulfate 0.1 N
14. Sodium monohydrogen phosphate, 0.1 N
15. Potassium ferrocyanlde, 0.1 M
16. Cadmium chloride, 0.1 M
17. Cobalt (II) nitrate, 0.1 M
18. Cerium (IV) sulfate (0.02 N in 3.6 N sulfuric acid)
19. Chromium (III) chloride, 0.1 M
20. Iron (III) nitrate, 0.1 M
21. Mercury (II) nitrate, 0.1 M
22. Lead nitrate.., 0.1 M
23. Antimony (III) chloride, 0.1 M
24. Gold (III) chloride (10 mgs. per milliliter of solution)
25. Silver nitrate, 0.1 M
26. Tin (IV) chloride, 0.1 M
27. Zinc chloride, 0.1 M
28. Uranyl nitrate, 0.005 M
29. Diethyl ether 16
30. Ethyl acetate
31* Methyl lsobutyl ketone (MiBK)
32. Tri-&-butyl phosphate (TBP)
33. Xsoamyl alcohol
34. Potassium permanganate, 0.01 M
35. Fluorescein (saturated solution In 1:1 alcohol)
36* Chromotropic acid (0.05% In conc. sulfuric acid)
37. Curcumln (boiled alcohol extract diluted with water)
38. Iodlne-sodlum aside solution (1 gram of sodium aside per 100 mis. of 0.1 N iodine solution)
39. Alpha-naphthylamine In glacial acetic acid (boiled water extract mixed with glacial acetic acid)
40. Malachite green, 2%
41. Molybdate~qulnollne solution (See p. 24)
42. Palladium (II) chloride, 1%
43. Sulfanlllc acid in acetic acid (1 gram of sulfanllic acid per 100 m is. of 30% a c e tic acid)
PROCEDURE
To 2 or 3 dropB of 0.1 M anion solution In a test tube, 2 or 3 drops of 6 N HC1 or HNO3 and 1 or 2 drops of 0.1 M cation solution were added and the volume was made up to about 1 ml. with distilled water. An equal volume of immiscible solvent was poured In, and the mixture agitated and the phases separated In a Carlton plpet (22). After a few minutes, when the two liquid phases had separated and both had become clear, each solution was pumped out into separate test tubes. 17
A few drops of the extract were subjected to the ring oven tech nique, The ring obtained was then spotted with the specific reagent for the anion.
The survey covered the following anions:
1 . C hloride 8 . Borate
2 . Bromide 9 . N itrate
3 . Iodide 1 0 . N itrite
4. Fluoride 1 1 . Sulfite
S. Thiocyanate 1 2 . Thio sulfate
6 . C yanide 13. Orthophosphate
7. O xalate 14. Ferrocyanlde
WSSVLIS 1. CHLORIDE
a. With mercury (II) .
West and Tron (23) studied this system using a number of solvents.
Here is a summary of their findings:
Solvent Percentage Extraction (One Pass)
TBP ...... 98
Ethyl acetate ...... 82
Amyl a c eta te ...... 76
MiBK ...... 76
M1PK...... 80
Methyl ethyl ketone ...... 73 18
Acetylacetone ...... 76
jv-Butyl alcohol ...... 74
jv-Amyl a lc o h o l ...... 65
Ethyl ether ...... 62
H-Butyl ether ...... 19
Carbon tetrachloride 2
A mixture (1:1) of ethyl acetate and TBP and pH's below 6 were recom mended for the extraction of chloride.
b. With iron (in)
In all three solvents (ethyl ether, MIBK, and ethyl acetate), extraction yields were fair.
To test for chloride, the ether and ethyl acetate extracts were treated with a few drops of water, and the organic solvents were evaporated off. The aqueous solution left was acidified with dilute nitric acid before the addition of 0.1 M sliver nitrate.
The time-consuming evaporation of the M1BK was later circumvented by the technique of adding another solvent, poor at extracting the anion but mlsclble with MiBK, followed by treatment with sodium hydroxide solution. Carbon tetrachloride was used as the second solvent. Test for the chloride was made on the aqueous phase after acidification with dilute nitric acid.
c. With antimony (V)
Fair yields were obtained with ether and MIBK. With ethyl acetate, poorer extraction was observed* 19
d. With gold (in)
Good extraction was noted with each solvent •
e. With cadmium
The three solvents extracted fair amounts of chloride.
A 1:1 TBP-M1BK mixture did not appear to be more efficient.
2 . BROMIDE
The fluorescein test (9) was used for qualitative detection.
The ring on the filter paper was spotted with acetic acld-hydrogen peroxide (1:2) solution. The spot was dried and then treated with a drop of 1% alcoholic fluorescein solution. A red ring appeared on drying If bromide was present.
a. With mercury (II)
The organic solvents yielded bromide in plentiful amounts.
West and Tron (23) reported 97% extraction for a Hg/fer ratio of
1:2 using a 1:1 TBP-ethyl acetate mixture.
b. With cadmium
No bromide could be detected in the ether extract.
Positive results were obtained with the MIBK and ethyl acetate extracts but further test of the strippings with dilute silver nitrate gave only a faint turbidity.
c. With gold (III)
Extraction with each of the three solvents was very efficient. The yellow compound formed upon mixing of Au(in) and bromide solutions was
completely transferred into the upper non-aqueous phase. 20
d. With Iron (III)
Positive fluorescein tests were given by the extracts • A check with dilute silver nitrate, however, failed to show even a slight turbidity.
3 . IODIDE
The qualitative detection was made with 1% palladous chloride solu tion added directly to the ring obtained in the ring oven. A black-brown ring appeared in the presence of iodide,
a. With cadmium
High extraction efficiency was found with each of the three solvents.
b. With antimony (III), mercury (II) or lead (II)
The yellow metal iodide precipitate was almost completely trans ferred into the non-aqueous phase in one pass.
1:1 mixtures of each of the solvents with TBP were likewise very e ffe c tiv e .
c. With gold (III)
The light brown metal iodide showed a strong preference for the organic solvent. Visual evidence was the complete transfer of the color to the upper organic solvent layer.
4 . THIOCYANATE
Its presence was demonstrated by its formation of the familiar
blood-red ferric thiocyanate complex with iron (III) •
To the ether and ethyl acetate extracts, a small volume of water
was added, and the mixture was left in a water bath to boil off the 21 organic solvent. After evaporation, the sides of the test tube were washed with water. A few drops of ferric nitrate solution were then added.
A blood-red color was formed In the presence of thiocyanate.
The MIBK extract was diluted with twice Its volume of carbon tetrachloride and shaken with 1 ml. of 10% sodium hydroxide. The alkaline aqueous layer was made acidic for the ferric thiocyanate test.
a. With cadmium
Good extraction was noted with each solvent.
b. With tin (IV)
Extraction efficiency was as good as that with Cd.
c. With Iron (III)
A fair amount of thiocyanate was found In each solvent.
Isoamyl alcohol extracted the ferric thiocyanate very well. Com
plete transfer of the color into the alcohol phase was noted.
d. With zinc
Ether gave the best result. Extraction with the other two solvents
was only fair.
e. Mercury (II)
The thiocyanate was extracted equally well by the solvents.
5 . NITRATE
To the ether and ethyl acetate extracts a small amount of water was
added, and the mixtures were set to evaporate in a water bath. The
aqueous solution left was tested for the nitrate. 22
The MiBK extract was subjected to the same treatment that was out lined under thiocyanate (above). The aqueous solution recovered was tested for nitrate with a 0.05% solution of chromotropic acid in concen trated sulfuric acid. A yellow color attributed to a nltro derivative of chromotropic acid (9) indicated the presence of nitrate,
a • With cerium (IV), gold (III) or uranium (VI)
Extraction efficiency was better than fair for all three solvents.
6 . FLUORIDE
A dilute solution of the ferric thiocyanate complex was used for
qualitative detection. A drop of the blood-red solution added to the ring
of fluoride on the filter paper gave a colorless ring clearly visible in the
surrounding redness.
a. With tin (IV)
Extraction with the three solvents was very poor.
b. With iron (III) or borate
Extent of extraction was only slightly better than with the preceding
sy ste m .
7 . CYANIDE
White precipitate with silver nitrate was the basis of qualitative
test. Strippings from the three solvents were made acidic with dilute
nitric acid and then treated with dilute silver nitrate,
a. With cadmium
MiBK gave a better result than either diethyl ether or ethyl acetate. 23
1:1 mixtures of TBP and each of the solvents used In the present survey were remarkably efficient.
The following anions were also shown by spot tests to be extracted using the cations and solvents indicated opposite them:
8 . BORATE - with fluoride In acid medium and diethyl ether.
Procedure For Detection (9):
A few drops of the aqueous solution containing boron recovered from the organic solvent were treated in a porcelain crucible with a very small quantity of silica and 2 drops of concentrated sulfuric acid. Heating, which was done carefully, was stopped when fumes of sulfur trioxlde were evolved. The residue was washed with a few drops of water, and the washings were used In the test for borate.
A drop of the borate solution, acidified with hydrochloric acid, was placed on curcuma paper and dried at 100°C. A red-brown fleck which turned blue to greenish-black on treatment with 1% sodium hydroxide
showed the presence of boric acid.
Curcuma paper was prepared by soaking quantitative filter paper
in tincture of curcuma and drying it. Tincture of curcuma was prepared
by boiling 20 grams of curcuma with 50 mis. of alcohol, filtering the
resulting mixture and diluting the filtrate with 50 mis. of water.
9. OXALATE - with iron (III) In acidic medium and diethyl ether.
The decolorizing action of oxalate on dilute potassium perman
ganate in acid medium was used to detect the anion. 24
10. . ORTHOPHOSPHATE - with iron (m) in addle medium and diethyl
tthar or isopropyl ether.
Proooduro For Detection (25) t
Tho reagent wot prepared by dissolving 10 grams of N*2Mo 0 4 «2H20
In 26 mis. of wator and than adding 11 mis* of oonoantmtad hydroohlorio sold and 10 mis. of oonosntratad nitric sold* Anothar solution was pro- parad oomposad of 1.4 mis* of qulnollna In a mlxtura of 15 mis. water,
7 mis* oonosntratad hydroohlorio sold# and € mis* oonosntratad nltrlo sold* Tho two solutions wars than mixed, and 60 mis* of watsr wars addad. Tha rasulting mlxtura was flltarad. Tha filtrate was usad to tast for orthophosphata in tha ring* A bright yallow ring davalopad In tha prssanoa of orthophosphata*
11* THIOSULFATE - with sllvar (I) In vary slightly aoldlo madlum, and
dlathyl athar or lxl M1BK-TBP.
Thlosulfata was datactad by tha lodlna-aslda raaotlon (9) • On
mixing a drop of tha tast solution recovered from tha extracts with a drop
of lodlne-sodlum aside solution, decolorlsatlon with concomittant forma tion of bubbles ensued in tha presence of thlosulfata*
Tha lodlne-sodlum aside solution was prepared by dissolving
1 gram of sodium aside in 100 mis* of 0.1 N Iodine solution.
12* SULFITE - with chromium (HI) In neutral solution and dlathyl ether.
Test By Decolorlsatlon of Malachite Green (9)t
Spotting of the ring deposit on the filter paper with 2% malachite 25 green solution gave a colorless ring Indicative of the presence of sulfite.
13. NITRITE - w ith cobalt (n) in neutral solution and diethyl ether.
Spot T est:
To the deposit on the filter paper, a drop of a 1:1 mixture of sul-
fanilic acid and aloha-naphthvlamlne was added. A pink ring showed the
presence of nitrite.
14 • FERROCYANIDE - w ith th e H+ and diethyl e th e r.
Spot T est:
A blue ring was produced upon addition of a drop of dilute iron (III)
nitrate solution to the deposit formed in the ring oven.
A table summarising the results of the qualitative survey on the
extractability of anions Is now presented.
TABLE I
Extraction of Anions - A Qualitative Survey
Meaning of symbols:
anion extracted as shown by spot test o, anion not extracted + , anion extracted slightly ++, anion extracted moderately +++, anion extracted very efficiently
Ether MIBK Ethvl Acetate
1 . CHLORIDE
a . Hg(ll) ++ +++ +++
b* Fe(III) ++ ++ ++ 26
TABLE I (continued)
Ether M1BK-. J$hriA£2$a$S
c . Sb(V) ++ ++ +
d . Au(III) +++ *hR- +++
e . Cd ++ ++ ++
2 . BROMIDE
a . H b. Fe(ni) O O O c . Au(IIl) +++ +++ +++ d . Cd O + + 3* IODIDE a . Hg(II) +++ +++ +++ b. Sb(in) +++ +++ -H-+ c . PbdD +++ +++ +++ d . Au(m ) +++ +++ +-H- e . Cd +++ +++ +++ 4 . THIOCYANATE a . Hg(II) +++ +++ +++ b . Fe(IIl) ++ ++ ++ c . SndV) +++ +++ +++ d . Zn +++ ++ ++ e . Cd +++ +++ +++ 27 TABLE I (continued) Ether MIBK EthvlAcetate 5 . NITRATE a . Ce(IV) ++ ++ ++ b . Au(m) + + + + ++ c . U(VD ++ ++ ++ 6 . FLUORIDE a. Fa (III) + + + b . Sn(IV) + + + c . B(m) + + + — 7 . CYANIDE a . Cd + +++ + 6 . BORATE a . F 9 . OXALATE a . Fe(III) 10. ORTHOPHOSPHATE a . Fe(IIl) 11. THIOSULFATE a . Ag(I) 12. SULFITE a . Cr(III) 13. NITRITE a . Co(II) 14. FERROCYANIDE II. EXPERIMENTAL (continued) B. STUDY OF A SPECIFIC SYSTEM - EXTRACTION OF IODIDE REAGENTS USED 1. Cadmium metal, mosey 2. Cadmium Iodide (wide range of concentrations) 3. Carbon tetrachloride 4 • Chloroform 5. Dlthlsone (5 mgs. per 100 mis. of chloroform) 6 . Ethyl alcohol, 95% 7 . Hydrochloric acid , 12 N 8 . Hydroxy lamina hydrochloride, 20% 9 . 8 -hydroxyqulnollne (2 grams per 100 mis. of alcohol) 10. MIBK 11. Nitric acid, 16 N 12. Potassium lodate, 0.1 N 13. Potassium permanganate, saturated solution 14. Salicylic acid 15. Silver nitrate, 0.1 N 16. Sodium hydroxide (0.1 N, 2%, 1 N, 10%) 17. Sodium nitrite, 1.5 M 18. Sodium potassium tartrate, 20% 19. Sodium thlosulfate, 0.05 N 20. Starch, 0.25% 28 29 21. Sulfuric acid (wide range of concentrations) 22. Trlbutyl phosphate 23. Urea, 5 M 1. EXPERIMENTAL PROCEDURES a . PREPARATION OF SOLUTIONS L • Cadmium Io d id e, 0. OS M Approximately 18 grams of cadmium iodide were dissolved in water, and the solution was diluted to one liter. By successive dilutions, 0.03 M, 0.01 M, 0.006 M, 0.003 M, and 0.001 M solutions were prepared. 2. Cadmium Iodide, 0.1 M To the solution containing 36*38 grams of cadmium iodide, 3 drops of 6 N sulfuric acid were added to opunteract hydrolysis. For buffer ing action, a gram of sodium acetate was dissolved in the solution. The pH of the solution was 5.1. There was no evidence of iodide oxidation even after a month of standing. 3. Cadmium nitrate solution Approximately 35 grams of mossy cadmium metal were treated with concentrated nitric acid (approximately 1 0 0 mis.) until solution was complete. Then from 10-15 mis. of 5 M urea were added to destroy oxides of nitrogen. The resulting clear liquid was diluted to 500 mis. to give a final solution containing about 70 mgs. of cadmium per milliliter. Standardization was made by the gravimetric oxlnate procedure. 4, Standard Potassium lodate Solution Tan grama of potaaslum lodata previously dried for an hour at 150°-160°C were weighed to the nearest milligram and dissolved in water. The solution was then diluted to 2 liters. 5 • Sodium Thlosulfata Stock Solution Twenty-five grams of sodium thlosulfata pentahydrate were dissolved in enough previously boiled water for a liter of solution. A centigram of sodium carbonate was added as preservative. For the prepara tion of 0.05 N solution, equal volumes of stock solution and previously boiled distilled water were mixed. 6 • Starch Indicator Solution 8 tarch paste containing 1 . 2 grams of starch was added, with vigorous stirring, Into 500 mis. of boiling water. One-half gram of salicylic add was added as preservative. 7, TBP-MIBK (1:1) Equal volumes of methyl lsobutyl ketone and trlbutyl phosphate were mixed. b . STANDARDIZATION OF CADMIUM NITRATE (12) Five milliliters of the cadmium nitrate solution were pipetted into a 250-ml. beaker. Sodium carbonate solution was added drop by drop until a slight turbidity was obtained. The precipitate was dissolved by the addition of 1:1 acetic acid drop by drop. A solution of 3-5 grams of sodium acetate in a small volume of water was poured carefully into the clear solution in the beaker. Cadmium was precipitated as the bright 31 yellow oxinate with a slight excess of 2% oxlne in 95% ethyl alcohol* The mixture was digested on an electric hot plate and subsequently filtered under suction through a sintered glass crucible* The precipitate was washed first with warm water, then with cold water. The cadmium oxinate was dried to constant weight at 150°-160°C. The weight of the precipitate multiplied by the chemical factor, 0.281, gave the weight of cadmium in 5 mis* of cadmium nitrate solution* c* COMPARISON OF CADMIUM IODIDE AND STANDARD POTASSIUM IODATE SOLUTIONS Ten milliliters of cadmium Iodide solution were pipetted Into a 1-liter Erlenmeyer flask. Sixty milliliters of 1 N sodium hydroxide were added, and the solution was diluted to about 120 mis. The solution, after being cooled in lce-water, was made 6 N In hydrochloric add by the addition of enough of the concentrated acid* Twenty milliliters of carbon tetrachloride were added as Indicator. Standard potassium lodate was measured from a buret Into the solution with vigorous shaking* At the end point, the faint pink carbon tetrachloride layer turned colorless. The Iodide titer of the standard potassium lodate was calculated, d . COMPARISON OF SODIUM THIOSULFATE AND STANDARD POTASSIUM IODATE SOLUTIONS Approximately 10 mis. of potassium lodate were transferred from a buret into a 250-ml. iodine flask. Ten milliliters of 1 N sulfuric acid, 50 mis. of water, and 10 mis. of 10% potassium Iodide were added in that order. The solution was titrated with sodium thlosulfate. One ml. 32 of 0.1 N silver nitrate was added midway In the titration to sharpen the endpoint with starch solution (4). The iodine titer of the sodium thlosulfate solution was calculated. e . SPECTROPHOTOMETRIC DETERMINATION OF CADMIUM Suitable aliquots of cadmium iodide solution containing 7-25 micrograms of cadmium were used. Two milliliters of 20% hydroxylamine hydrochloride and 3 mis. of 20% sodium potassium tartrate were added and the resulting solution diluted with an equal volume of 1 0 % sodium hy droxide. Small volumes of dlthizone in chloroform were used to extract cadmium as the dithlsonate. A colorless chloroform layer signalled the end of the extraction. The combined extracts in another separatory funnel were washed twice with 10 mis. of 2% sodium hydroxide. Finally/ the extract was shaken with 10 mis. of water acidified with 2 drops of 6 N hydrochloric acid. The green chloroform layer was transferred into a 50-ml. vol umetric flask and diluted to the mark with chloroform. The absorbancy was read at 600 nyi and slit width of 0.0375 mm. with a Beckman DU Spectrophotometer. f . DETERMINATION OF SEMI-MICRO QUANTITIES OF IODIDE A 10-milllliter solution of iodide in an iodine flask was treated with a few milliliters of 1 0 % sodium hydroxide before dilution to about 50 mis. An excess of saturated potassium permanganate solution was added/ and the mixture was kept simmering for about 5 minutes. When the mixture had cooled a little, 2-3 mis. of 1:1 sulfuric acid were poured 33 in to acidify tho mixture. Addition of 1.5 M sodium nitrite drop by drop directly into the colored solution resulted in decolorlsatlon. The clear solution was set aside to cool before the titration with standard sodium thlosulfata was done. g . EXTRACTION OF IODIDE Ten milliliters of cadmium iodide solution were pipetted into a 60-ml. separatory funnel. Dilute sulfuric acid was then added to bring— the solution to the desired pH. To this acidified solution, an equal volume of 1:1 TBP-MXBK was added, and the mixture shaken for about thirty seconds to insure contact between the two immiscible liquids. The separatory funnel containing the mixture was immersed in a thermostat kept at 30°C, where the system was allowed to come to equilibrium. After equilibration, the non-aqueous layer was treated with 2 0 mis. of carbon tetrachloride, and the mixture then shaken with 2 0 mis. of 1 N sodium hydroxide to recover the cadmium and iodide. Copious precipitation of white cadmium hydroxide resulted, and the Iodine was re leased in the form of I*, possibly including small amounts of OX~. The CC I 4 -MIBK-TBP layer was washed twice with 20-ml. por tions of 1 N sodium hydroxide, and the washings were combined with the first one containing the cadmium hydroxide precipitate. This mixture w as made 6 N in regard to hydrochloric acid, and the iodide determined by a conventional iodine monochloride procedure. ► Cadmium was determined in the aqueous layer by use of a dlthlaone extraction procedure (II-B, le , p. 32) on suitable aliquots 34 of the solution. The cadmium In the non-aqueous layer and the unex tracted Iodide were obtained by difference. Cadmium Iodide solutions used for the Initial concentration studies were standardised by hot permanganate oxidation of the iodide followed by thiosulfate titration (n-B, 1 f; II- B, Id). 2 . VARIOUS FACTORS INVOLVED IN THE DETERMINATION OF IODIDE BY EXTRACTION Careful consideration of the results of the qualitative survey and other factors led to the decision to investigate the extraction of iodide In detail. The qualitative survey showed that several cations could be used in the formation of the uncharged extractable species. However, the precipitate formed with lead (II), mercury (II), and antimony (III) made clean separation of the layers difficult and argued against these metal Ions. Mercury (II) had another limitation: percentage extraction drops considerably (23) as the Hg:I ratio deviates from the 1:2 optimum. For obvious economic reasons, gold (III) eliminated itself from serious consideration. Only cadmium remained: it was chosen. Cadmium has several advantages over the other metal iciis mentioned above. Its Iodide not only dissolves in water readily to give reasonably stable solutions, but it is also non-hygroscoplc. Cadmium iodide can be dried at 100-110°C, and therefore, standard solutions can be pre pared from accurately weighed samples. 35 a. CHOICE OF 80LVENTS Waat and Tron (23) found that a mixture of TBP and ethyl acetate gave eaaentlally complete extraction of the halides of mercury in a single pass. After a comparison of the relative merits of the solvents used in the qualitative survey individually, and their separate mixtures with TBP, a 1:1 TBP-M1BK mixture was selected because of its extractive efficiency and its relative insensitivity to changes in acidity. b . JtBCOVERY OF IODIDE For the recovery of extracted iodide, the usual techniques used in extraction, such as shaking the non-aqueous layer with dilute acids or dilute bases, were tried and found ineffective. Finally, a new ap proach was introduced, that of adding a third solvent, mlsclble with the non-aqueous solvents but very poor at extracting iodide. The addi tion of about twice as much carbon tetrachloride as the volume of the TBP-M1BK layer produced dramatic results. A yellow color developed after a minute or two of standing. At this point, addition of sodium hydroxide gave a two-phase system composed of an upper aqueous layer and a lower non-aqueous layer. A large proportion of carbon tetra chloride had made the non-aqueous phase the more dense one. After a minute of vigorous shaking, a white precipitate of cadmium hydroxide formed in the aqueous phase, and the yellow color disappeared from the TBP-M1BX-CC14 layer. The recovery of the Iodide from the non-aqueous phase by the solvent displacement technique Just outlined proved to be remarkably efficient. It Is of interest to note that this Is a reverse of 36 the established practice of using solvent mixtures to effect Increase In solvent extraction efficiencies. c . EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE RECOVERY OF IODIDE Recovery of the extracted Iodide was tried with various sodium hydroxide solutions. The results are indicated In the table below. TABLE II Effect of NaOH Concentration on the Recovery of Iodide NaOH Solution fffretOttflS .RtgflYtty 60 mis. of 0.1 N NaOH 93 60 mis. of 1 N NaOH 95 20 mis. of 10% NaOH 40 mis. of 2% NaOH 9 3 The foregoing data suggested the use of 1 N NaOH. Stripping was done with three 20-ml. volumes of base. The recovery per 20-ml. of 1 N NaOH Is summarized below: Volume of 1 N NaOH Percentage Recovery 20 m is. 93 40 m is. 94 60 m is. 95 Data will be presented later showing that 95% represents essentially complete recovery of the extracted iodide (see Tables XIV-D/ E, F of the Appendix) . 37 d . EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE TITRATION OF IODIDE The effect of sodium hydroxide on the titration of Iodide with potassium lodate (n -B , 1c, p . 31) was studied next. The following table diows that while moderate amounts of base can be tolerated, the presence of larger quantities causes higher values for the Iodide. TABLE III Effect of NaOH Concentration on the Titration of Iodide Solution ja wiFpunfl (Mflii) ftihttvf Brcac 1 0 m is. C dl 2 6 6 mis. water 123.5 76 mis. conc. HC1 1 0 m is. C d l 2 5.5 mis. water 123.8 0.24% 60 mis. 0.1 N NaOH 76.5 mis. conc. HC1 10 m is. C dl2 1 ml. water 123.6 0.24% 60 mis. IN NaOH 81 mis. conc. HC1 10 m is. C d l2 20 mis. 10% NaOH 126.2 2.1% 40 mis * 2% NaOH 82 mis. conc. HC1 38 Apparently, the presence of too much sodium chloride slows down the reaction between the lodate and the Iodide or Iodine • e . EQUILIBRATION TIME AND PERCENTAGE EXTRACTION The relation between the equilibration time and percentage extraction was studied using time Intervals differing by half an hour. On the basis of the results shown below, an equilibration time of one hour was chosen. TABLE IV Equilibration Time and Percentage Extraction Time Iodide Present teflldff % Extraction 1 / 2 hour 132 mgs. 125 m gs. 95 1 hour 132 mgs. 126 mgs. 96 1 1 / 2 hours 139 mgs. 133 m gs. 96 2 hours 139 m gs. 132 m gs. 95 2 1 / 2 hours 139 m gs. 132 m gs. 95 3 hours 139 mgs. 131 m gs. 94 The values obtained for a if 1 hour were relatively less erratic, although even that for 1/2 hcui was not bad. As a matter of fact, any time Interval between 1 / 2 hour and 2 1 / 2 hours could be used. Atmospheric oxidation of the Iodide In the TBP-M1BK layer, as evidenced by Its yellow color which deepened with time, and the sub sequent loss of the iodine were probably responsible for the decrease in percentage extraction with time of equilibration. 39 3. SOURCES OF ERROR The principal source of error was the separation of the two phases. This operation might be responsible for an error of about 2 per cent. The spectrophotometrlc determination of cadmium was another source of error (—2 per cent). Solutions had to be diluted several times to get them into the microgram range, and then, only a few milliliters of the final solution were used in the spectrophotometrlc determination. The atmospheric oxidation of the iodide was already mentioned. As the pH of the solution was Increased, this error became smaller, be cause oxidation was less. Solving for the Iodide in the aqueous layer and the cadmium in the non-aqueous layer was also subject to error. The decision to do it by difference came after material balance calculations for some samples did not show great error. There seemed to be no point In going through a whole series of time-consuming operations that would give a value not much different from that easily obtainable by difference. 4 . THE HYDROGEN ION CONCENTRATION TABLE V Effect of Hydrogen Ion Concentration on the Extraction of Iodide ______Eli______. 0 .3 0 . 6 _L _2_ _2 _ _6_ Percentage I Extracted 96 95 95 94 94 94 Percentage Cd Extracted 6 8 70 82 93 95 96 Cd/I Mole Ratio In Extract 1/ 2 . 8 1 /2 .7 1 /2 .4 1 / 2 1 / 2 1 / 2 pH of Aqueous Phase +0 . 2 + 0 . 1 + 0 . 1 +0.5 +0.3 -3.2 40 As shown in Table V, the system was not very sensitive to changes in pH. A twenty-fold change in pH (0.3 to 6 ) occasioned only about 2 per cent decrease in the iodide extracted. At pH values less than 0.3, there was extensive atmospheric oxidation of the iodide. On the other hand, starting at pH 2, a whitish emulsion formed In the non-aqueous layer with shaking, which took longer to break with further Increase In pH. A yellow tinge was observed in the aqueous layer of the system a t pH 6 , which could only be Interpreted as evldenoe of oxidation of the Iodide. It is well known that addle iodide solutions turn yellow because of the presence of free iodine released by the action of either atmospheric or dissolved oxygen. The final pH («^3) of the aqueous phase, however, was quite high for oxidation to occur appreciably. The only likely ex planation was the presence of localized regions of low pH where oxida tion of the iodide by dissolved oxygen was appreciable. Unless otherwise stated, the Initial concentration of the cadmium iodide solution was approximately 0.05 M^he temperature, 30°C, and the pH of the solution, 0.6. 5 . THE CADMIUM TO IODIDE MOLAR RATIO If any metal ion is to be of practical use in the separation of anions by solvent extraction, its presence in excess should not appre ciably decrease the extraction efficiency. Fortunately, the cadmium ion proved to be very satisfactory for iodide extractions when considered on this basis. 41 TABLE VI Effect of Initial Cd/I Mole Ratio on the Extraction of Iodide C d A Mole Ratio In Aqueous Phase l / l 1 / 1 l / l A /1 Percentage I Extracted 95 95 92 89 Percentage Cd Extracted 70 50 24 2 0 Cd/1 Mole Ratio in Extract 1 /2 .7 1/ 1 . 8 1 /1 .9 1 /1 .3 6 . THE INITIAL CONCENTRATION OF THE CADMIUM IODIDE The data below show that starting with solution strengths approxi mating those usually encountered In analysis (0.05-0.1 M), almost com plete separation of the Iodide Ion using cadmium as the coordinating cation can be attained In three extractions. TABLE VII Effect of Initial Concentration of Cdl 2 on the Extraction of Iodide Initial Concentration of Cadmium Iodide ______0.001 0.003 0.006 0.01 0.03 0.05 0.1 M P ercen tag e I Extracted 61 77 87 91 95 95 96 P ercentage Cd Extracted 60 53 6 6 69 70 70 75 Cd/I Mole Ratio in Extract 1/2 1/3.7 1/2.6 1/2.7 1/2.7 1/2.7 1/2.6 It can be seen that extraction is still appreciable even for solu tions as dilute as 0.006 M. The data above were obtained at 30°C and pH 0.6. Greater efficiencies should result at higher acidities. 7 . EFFECT OF TEMPERATURE Because the solvent mixture Is inflammable, the study of temperature effects was not extended beyond 40°C. Moreover, the system (assuming that the small differences, which are smaller than the experimental error, are significant) seems to behave in accordance with BJerrum's theory. For solvents of low dielectric constant, theory pre dicts a decrease in ion association with Increase In temperature. Since the extractable species are presumably Ion association complexes, the decrease in percentage iodide extracted as the temperature increases to o 40 C is in agreement with theory. TABLE VIII Effect of Temperature on the Extraction of Iodide Temperature 30° 40° C 1(0 Icn o 1 Percentage I Extracted 97 95 94 Percentage Cd Extracted 68 70 68 Cd/I Mole Ratio in Extract 1/2.9 1 /2 .7 1/2 .8 8 . SELECTIVITY A qualitative survey of the common anions was conducted to find out which of them would be extracted under the same conditions. Varying degrees of extractability were encountered and may be classified according to the following groups: 1. Anions that were extracted with approximately the same efficiency as the iodide: CNS" and CN~ 2. Anions extracted to a moderate degree: C l~, Br“ , C 2 0 4 , and NO 3 3 • Anions that were extracted very poorly or not at all: F“ , S 2 Oj“, 804", POj3, Ac- / S“ , F®(CN)“4, and BOj 3 No evidence of coextraotlon was found. III. DISCUSSION OF RESULTS Cadmium Iodide solutions are known to contain the following ionic and molecular species: Cd**, Cdl+, Cdl 2 / Cdl^, Cdl4, and I~. The simple cadmium and iodide ions are probably found only in negligible amounts, since cadmium iodide is a weak electrolyte, except in very dilute solutions (18). Working with neutral solutions of ionic strength of 3, Leden (13) found the iodide complexes of cadmium present in the following decreasing order of abundance: Cdl^~, Cdg, Cdl2, and Cdl+. Alberty and King (1), on the other hand, obtained from mobility data, using solutions of ionic strength of 0.5 and total iodide concentration of 0.1 M, the following values for the cadmium iodide complexes: C dl4“ 62 per cent CdlJ 16 per cent C d ^ . . 1 2 per cent C dl+ ...... 6 per cent Cd++ ...... negligible Their graphs of fraction of complex vs. total iodide concentration from where the above values were obtained showed that the relative amounts of the different complexes varied with the total iodide concentration. Harris (11) reported from Ion exchange experiments, using 0.0666 M Cdl2, the presence of Cdl4”, Cdl2, Cdl^, and Cdl+, in that increasing order of abundance. 44 In the present studies the solutions were acidic. Undoubtedly, all the different complexes of cadmium with iodide would still be present. The question is which forms predominate over the others ? For simplicity, the formulas of the cadmium iodide complex ions were written unhydrated, although it is known that molecules of water are coordinated to the central ion. Cadmium Iodide, for example, is really C dff^O )^* With this In mind, it Is possible that the addition of hydrogen ions, which have a great affinity for water molecules, makes it relatively easier for the coordinated water to leave the sphere of Influence of the central metal ion and for the simple iodide ions to take over the vacated posi tions. In solutions of low pH, say 0.3 or 0.6, the principal cadmi'im iodide complexes are probably, in decreasing order of abundance, CdlJ, C d l 2 and C d^”. At higher pH values, the following are probably present to a greater degree: Cdl2 , and Cdl+. The data in Table V point to the following as the uncharged species extracted from solutions of low pH: (H+, C dlp, Cdl2, and (2H+, CdlJ") The Increased participation of the last named ion association complex should be expected with further Increase in acidity. From pH 2 on up to higher pH values, the extracted species are probably Cdl 2 , and (Cdl+, OH”) with the latter making its greatest participation at pH 6 as evidenced by the Increase in the acidity of the aqueous phase, an in crease which is too great to attribute to an enhanced solubility of water in the TBP-MiBK phase. The relative insensitivity to pH change may be due to the existing 46 equilibria among the different complexes of cadmium and Iodide which allow lnterconverslon from one major Ionic species to another one that can form an uncharged extractable complex at the new pH. As long as the higher Ion association complexes are the major species being extracted, the extraction efficiency of iodide is high. Addition of cadmium Ions to vary the Cd:I mole ratio increases the ionis strength of the solution, making extensive ion association possible. However, the concurrent chemical effect limits the Ion association to the lower complexes as shown by the following equations: Cd++ + C d lJ" ■ ■■ » 2 C dl 2 Cd** + Cdl* ------» C dl2 + Cdl+ Cd** + C dl 2 ------» 2 Cdl* The uncharged species being extracted are probably Cdl2, and {Cdl+, NOj), with the latter increasing in Importance with Increase in Cd:I mole ratio. The solution added to Increase the Cd:I mole ratio was cadmium nitrate which explains the appearance of the nitrate ion here. No attempt was made to estimate the amount of nitrate extracted, but its presence was detected. The Increase in extraction with increase in initial concentration of the cadmium iodide solution may be due to the ionlc-strength and mass-action effects. The former favors ion-association complex forma tion which, because of the mass-action effect and the presence of hydrogen ions, would in turn favor the higher complexes, Cdlj and C d lJ". 47 2 C d l 2 ------► Cdlj + Cdl+ 3 C d l 2 ------> Cdl“ + 2 Cdl+ At lower concentrations, the amount of simple cadmium and Iodide ions increases, but apparently, enough complexes are still present to sustain the formation of extractable species • The data In Table VII point to (2H+, CdIJ“), (H+, Cdip and Cdl 2 as the predominant forms being ex- ticted with the participation of (2H+, CdlJ") reaching its peak at 0.003 M. Apparently, the Cdl+ ion, which is the predominant complex at very low cadmium iodide concentrations, is unable to form an extractable uncharged ion association complex, the pH of the solutions (0.6) being so low as to make the formation of (Cdl+, OH”) highly improbable. BJerrum's theory or Fuoss* Improved theory predicts that an in crease In temperature will cause a decrease In ion association In the TBP-M1BK phase. In the aqueous layer, the number of Ion association complexes should increase. Of these two opposing effects, the data seem to Indicate predominance of the former. The small change (smaller than the experimental error) in percentage extraction shown in Table VIII with change of temperature, can only mean that it is the extent of ion association in the TBP-M1BK phase that is significant in determining the efficiency of extraction. The Fuoss equation may be written in the form: log K * log A + -XI- DT where log A represents the coefficient of the exponential term and k.', the constant €.Vk. changes only slightly with change in temperature, for solvents of low dielectric constant. Hence, the Increase or decrease in ion association in the TBP-MiBK phase must be due mainly to the change in the temperature, X* However, for a change of 10 degrees in £, there will only be a small change in log K. This may be the reason for the small differences in the percentage of iodide extracted at the three different temperatures. IV. DEVELOPMENT OF A PROCEDURE FOR THE EXTRACTION METHOD OF IODIDE DETERMINATION Aqueous solutions of potassium Iodide ( t* 0.1 M), acidified with 6 N sulfuric acid to a pH of approximately 0.6, were used. Before a procedure could be recommended, the volume of mixed solvents that should be added for each pass had to be determined. The following volume distributions were tried; 1. Three passes of 10 mis. each. 2. Three passes of 5 mis. each. 3. Ten milliliters for the first pass and 5 mis. each for the second and third passes. The first one was finally abandoned because the large volume that re sulted with the addition of carbon tetrachloride made it impractical and unw ieldy. The results for the volume distributions (2) and (3) are shown below; (2) THREE PASSES OF 5 MLS. EACH OF 1:1 TBP-MiBK Vol. of KIO3 Iodide Present Iodide Found % Rel. Error 1. 2 0.94 m is. 129 m gs. 127 m gs. 1.55 2 . 2 1 . 0 0 m is. 129 m gs. 127 m gs. 1.55 3. 21.05 mis. 129 mgs. 127 m gs. 1.55 4. 20.98 mis. 129 mgs. 127 m gs. 1.55 5. 20.89 mis. 129 mgs. 126 mgs. 2.32 6 . 20.60 m is. 129 m gs. 124 mgs. 3.88 Av. 20.91 mis. 126 mgs. 2.07 49 50 (3) ONE PASS OF 10 MLS* OF 1:1 TBP-M1BK FOLLOWED BY TWO PASSES OF 5 MLS. EACH Iodide Present ifididS^ound % Rel. Error 1 . 20.98 m is. 129 m gs. 127 mgs. 1.55 2 * 2 0 .9 6 m is. 129 m gs. 127 mgs. 1.55 3* 2 1 . 0 0 m is. 129 mgs. 127 mgs. 1.55 4. 20.96 mis. 129 mgs. 127 mgs. 1.55 5. 20.75 mis. 129 mgs. 125 mgs. 3 .10 6 . 21.03 m is. 129 m gs. 127 m gs. 1.55 k u 20.95 m is. 127 mgs. 1.81 The greater error in the first case may be traced to the smaller volume (5 mis.) of mixed solvents used in the first pass. The use of a volume of mixed solvents approximately equal to that of the aqueous phase for the first pass meant relatively smaller error because of easier separation of the immiscible phases. The following procedure, based on the above results and on the findings detailed in Part II of this dissertation, is recommended. PROCEDURE Ten milliliters of the iodide ( r 0.1 M) are pipetted into a 125-ml. separatory funnel. One milliliter of 0.5 M cadmium nitrate and 1 ml. of 6 N sulfuric acid are added, and the solution is mixed well. Ten milliliters of 1:1 TBP-MiBK are next poured in. The two-phase system is shaken for thirty seconds before being set aside to allow separation 51 and clarification of the Immiscible layers. The lower aqueous phase is run Into a 60-ml. separatory funnel and a few drops of 6 N sulfuric acid added. The residual Iodide is ex tracted with two 5-ml. volumes of the mixed solvents. The 5-ml. ex tracts ve combined with the first one in the 125-ml. separatory funnel, and the resulting mixture is diluted with twice its volume of carbon tetrachloride. The Iodide Is recovered by shaking the TBP-MiBK-CCl 4 p h ase three times with 20 mis. of 1 N sodium hydroxide. The combined alkaline strippings In a 1-liter Erlenmeyer flask are made 6 N in hydro chloric acid, and the iodide titrated by the conventional Iodine mono chloride titration with standard potassium lodate. Twenty mis. of carbon tetrachloride were used as indicator. The percentage of iodide in the sample is computed from the formula: % j _ Vol. of KIQ 3 x Iodide Titer of KIO 3 x 100 Weight of Sample V. APPENDIX TABLE IX GRAVIMETRIC STANDARDIZATION OF CADMIUM NITRATE SOLUTION BY PRECIPITATION OP CADMIUM OXINATE JOSls_ W eight of Crucible + 25.1315 25.5913 25.4038 25.1238 25.4154 precipitate W eight of empty 23.9305 24.3911 24.2118 23.9308 24.2118 cru cib le W eight of precipitate 1.2010 1.2002 1.1920 1.1930 1.2036 Av. weight of precipitate » 1.198 grams Weight of Cd Weight of precipitate x — — C dO x2 * 1.198 x 112.4 4 0 0 .6 336.2 mgs. cadmium in 5 mis. of cadmium nitrate so lu tio n . - 67.24 mgs. cadmium per ml. of solution. 52 53 TABLE X POTASSIUM IODATE SOLUTIONS Iodine Titer 1 7.693 mgs 2 7.023 3 6.328 4 6.022 5 6.097 6 6.045 7 6.384 8 7.579 Note: Because of space limitations, in some of the following tables the potassium iodate solution used will just be Indicated by its number (above) In parenthesis. For example, KIO 3 (1) is solution no. 1, above, with an iodine titer of 7.693 mgs. per milliliter. 54 TABLE XI-A EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE TITRATION OF IODIDE A: 10 mis. of Cdl 2 * 06 mis. water, 76 mis. conc. HC1 Volume of KlOq Iodine Titer of KIQg Iodide Found 1 17.62 mis. 7.023 mgs./ml. 123.7 mgs 2 17.64 7.023 123.8 3 17.68 7.023 124.1 4 17 . 65 7.023 123.9 5 16.27 7.579 123.3 6 16.25 7.579 123.1 7 16.22 7.579 122.9 8 16.21 7.579 122.9 Average 123.5 mgs 55 TABLE XI - B EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE TITRATION OF IODIDE B: 10 mis. Odl^t 5.5 mis. H2 ), 60 mis. 0.1 NaOH, 76.5 mis. conc. HC1 Volume of KIO'% Iodide Present Iodide Found Deviation 1 17.69 mis. (2) 123.5 mgs. 124.2 mgs. 0.7 2 17.66 123.5 124.0 0.5 3 17.68 123.5 124.1 0 . 6 4 17.67 123.5 124.0 0.5 5 16.34 (8 ) 123.5 123.8 0.3 6 16.26 123.5 123.2 0.3 7 16.32 123.5 123.7 0 . 2 8 16.29 123.5 123.4 0 . 1 Average: 123.8 mg s . 0.4 Relative Error: 0.24% 56 TABLE XI-C EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE TITRATION OF IODIDE C: 10 mis. Cdl2, 1 ml. water, 60 mis. 1 N NaOH, 81 mis. cone. HC1 Volume of KIO-i Iodide Present Iodide Found Deviation 1 17.62 mis. (2) 123.5 mgs. 123.7 mgs. 0 . 2 2 17.64 123.5 123.8 0.3 3 17.68 123.5 124.1 0 . 6 4 17.66 123.5 124.0 0.5 5 16.15 (8 ) 123.5 122.4 1 . 1 6 16.29 123.5 123.4 0 . 1 7 16.28 123.5 123.4 0 . 1 8 16.32 123.5 123.7 0 . 2 Average: 123 .6 mgs. ■0.3 Relative Error 0.24% TABLE XI-D EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE TITRATION OF IODIDE D: 10 mU. Cdlg, 20 mis. 10% NaOH, 40 mis. 2% NaOH, 82 mis. conc. HC1 Voluma of KIO^ feStttiS-EZSifiSL Iodide Found Deviation 1 17.99 mis. (2) 123.5 mgs. 126.3 mgs. 2 . 8 2 17.98 123.5 126.2 2.7 3 18.04 123.5 126.6 3.1 4 18.03 123.5 126.6 3.1 5 16.63 (8 ) 123.5 126.1 2 . 6 6 16.61 123.5 125.9 2.4 7 16.55 123.5 125.4 1.9 8 16*66 123.5 126.3 2 . 8 Average: 126.2 mgs. 2.7 Relative Error 2 . 1% 58 TABLE XII-A EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE RECOVERY OF IODIDE A: 60 rnla. of 0.1 N NaOH Percentage Volume of XIO q Iodide Preeent Iodide Found Recovery 1 17.69 mis. ( 1) 145.2 mge. 136.1 mgs. 94 2 17.48 145.2 134.5 93 3 17.75 145.2 136.6 94 4 17.72 145.2 136.3 94 5 17.55 145.2 135.0 93 6 17.64 145.2 135.7 93 7 17.61 145.2 135.5 93 8 17.51 145.2 134.7 93 Av.: 17.62 mis. 135.6 mgs. 93 Concentration of Cdl 2 Solution = 0.05 M pH = 0.6 Temperature = 30°C 59 TABLE XII-B EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE RECOVERY OF IODIDE B: 60 mis. of 1 N NaOH Percentage Volume of KIO3 Iodide Present Iodide Found Recovery 1 21.83 mis. (5) 139.3 mgs. 133.1 mgs. 96 2 21.79 139.3 132.9 95 3 21.78 139.3 132.8 95 4 21.59 139.3 131.6 94 5 21.52 139.3 131.2 94 6 20.94 (4) 131.7 126.1 96 7 20.99 131.7 126.4 96 8 20.61 131.7 124.1 94 Average 95 Concentration of Cdl^ Solution = 0.05 M pH = 0.6 Temperature = 30°C 60 TABLE XII-C EFFECT OF SODIUM HYDROXIDE CONCENTRATION ON THE RECOVERY OF IODIDE C: 20 mis. of 10% NaOH, 40 mis. of 2% NaOH Percentage Volume of KIO^ Iodide Present Iodide Found Recovery 1 20.96 mis. (4) 133.3 mgs. 126.2 mgs. 95 2 21.05 133.3 126.7 95 3 2 1 . 0 1 133.3 126.5 95 4 21.03 133.3 126.6 95 5 21.05 133.3 126.7 95 6 21.16 133.3 127.4 96 7 2 1 . 2 0 133.3 127.7 96 8 2 1 . 0 0 133.3 126.4 95 Av.: 2 1 . 1 0 mis. 126.8 mgs. 95 Concentration of Cdlg Solution = 0. OS M pH = 0.6 Temperature * 30°C 61 TABLE XIII RECOVERY OF IODIDE WITH 1 N NaOH A; 20 mis. of 1 N NaOH Percentage Volume of KIO^ Iodide Present Iodide Found Recovery 1 20.41 mis. (4) 131.7 mgs. 122.9 mgs. 93 2 20.50 131.7 123.5 94 3 2 0 . 1 0 131.7 1 2 1 . 0 92 4 19.99 131.7 120.4 91 5 20.63 131.7 124.2 94 6 19.95 131.7 1 2 0 . 1 91 7 20.35 131.7 122.5 93 7 20.43 131.7 123.0 93 A; 20.30 mis. 1 2 2 . 2 mgs. 93 B: 40 mis. of 1 N NaOH 1 20.56 mis. (4) 131.7 mgs. 123.8 m gs. 94 2 20.60 131.7 124.1 94 3 20.42 131.7 123.0 93 4 20.63 131.7 124.2 94 5 20.65 131.7 124.3 94 6 20.58 131.7 123.9 94 7 20.53 131.7 123.6 94 8 2 0 . 2 1 131.7 121.7 92 20.52 mis. 123.6 mgs. 94 62 TABLE XIII (C o n 't.) C: 60 mis. of 1 N NaOH Percentage Volume of KIQg Iodide Present Iodide Found Recovery 1 20.78 mis. (4) 131.7 mgs. 125.1 mgs. 95 2 20.77 131.7 125.1 95 3 20.99 131.7 126.4 96 4 20.61 131.7 124.1 94 5 20.55 131.7 123.7 94 6 20.94 131.7 126.1 96 7 2 1 . 0 0 131.7 126.4 96 8 20.65 131.7 124.3 94 Av.: 20.79 mis. 125.2 mgs. 95 Concentration of Cdlg Solution = 0.05 M pH = 0.6 Temperature * 30°C 63 TABLE XIV - A EFFECT OF HYDROGEN ION CONCENTRATION ON THE EXTRACTION OF IODIDE A: pH - 0.3 ______Non-aqueous Laver ______AVt „ KIO3 (5) 21.77 21.83 21.85 21.91 21.92 21.96 21.87 I(mgs.) 133 133 133 134 134 134 134 %E (I) 96 96 96 96 96 96 96 Cd(mgs?) 43 42 42 43 42 42 42 %E (Cd) 6 8 67 67 69 6 8 6 8 6 8 fid -L- -A. I 2 . 8 2.9 2.9 2 . 8 2 . 8 2 . 8 2 . 8 Aaueous Laver I Left* 6 6 6 5 5 5 5 I (Orlg.) 139 139 139 139 139 139 139 %I (Left) 4 4 4 4 4 4 4 Cd Left 19 2 0 2 0 19 2 0 2 0 2 0 Cd (Orlg.) 62 62 62 62 62 62 62 %Cd (Left) 32 33 33 31 32 32 32 ApH +0.18 +0.17 + 0 .20 +0.19 +0.20 +0 . 2 +0.19 * obtained by difference 64 TABLE XIV - B EFFECT OF HYDROGEN ION CONCENTRATION ON THE EXTRACTION OF IODIDE B: pH » 0.6 ______Non-agueoua Laver ______ VKI03 (5) 21.62 21.79 21.78 21.59 21.79 21.83 2 1 .' I (mg*.) 132 133 133 132 133 133 132 %E (I) 95 95 95 94 95 96 95 Cd(mgst) 44 44 44 44 44 44 44 %E (Cd) 70 70 70 70 70 70 70 0 4 1 -L- 1 1 1 1 I 2.7 2.7 2.7 2.7 2.7 2.7 2.7 Aaueous Laver I Left* 7 6 6 7 6 6 7 I (Orlg.) 139 139 139 139 139 139 139 %I (Left) 5 5 5 6 5 4 5 Cd (Left) 18 18 18 18 18 18 18 Cd (Orlg.) 62 62 62 62 62 62 62 %Cd (Left) 30 30 30 30 30 30 30 A pH + 0 . 1 0 +0.08 +0.05 +0.08 + 0.06 + 0 . 1 0 + 0 . * obtained by difference 65 TABLE XIV - C EFFECT OF HYDROGEN ION CONCENTRATION ON THE EXTRACTION OF IODIDE C: pH - I ______Non-agueous Laver ______1 2 3 _ 4 5 6 Av. VKI03 (3) 19.65 19.64 20.06 19.61 20.10 19.82 19.88 1 (mgs.) 124 126 127 125 127 125 126 %E (I) 94 95 96 95 96 95 95 Cd (mgs?) 48 46 48 46 48 47 47 %E (Cd) 83 80 83 80 84 81 82 £d 1 1 1 X— _ 1 I 2.3 2.4 2.3 2.4 2.3 2^4 2 . Aaueous Laver I Left* 8 6 5 7 5 7 6 I (Orig.) 132 132 132 132 132 132 132 %I (Left) 6 5 4 5 4 5 5 Cd (Left) 1 0 12 10 12 10 11 11 Cd (Orig.) 58 58 58 58 58 58 58 %Cd (Left) 17 2 0 17 2 0 16 19 18 ApH + 0.08 +0 . 1 0 + 0 . +0.09 +0 . 1 0 +0.08 +0.08 * obtained by difference 66 TABLE XIV - D EFFECT OF HYDROGEN ION CONCENTRATION ON THE EXTRACTION OF IODIDE D: p H - 2 ______Non-aautoui Laver ______vn o s (3) 19.49 19.60 19.59 19.75 19.47 19.65 19 I (mgs.) 123 124 124 125 123 124 124 %E (I) 94 94 94 95 93 94 94 Cd (mgs?) 54 54 54 55 54 54 54 %E (Cd) 93 93 93 95 93 93 93 £d 1 1 1 1 1 1 1 I 2 2 2 2 2 2 2 Aqueous Layer I Left 7.1 7.1 7.1 7.1 7.1 7.1 7.1 I (Orig.) 132 132 132 132 132 132 132 %I (Left) 5.4 5.4 5.4 5.4 5.4 5.4 5.4 Cd (Left) 4.4 4.2 4.0 3.8 4.1 4.1 4.1 Cd (Orlg.) 58 58 58 58 58 58 58 %Cd (Left) 7 7 7 5 7 7 7 ApH +0.4 +0 . 6 +0.5 +0.5 + 0.5 +0.5 + 0.5 * obtained by difference 67 TABLE XIV - E EFFECT OF HYDROGEN ION CONCENTRATION ON THE EXTRACTION OF IODIDE E: pH - 3 ______Non-aqueous Laver ______1 2 3 4 5 6 Av vKK> 3 (3) 19.42 19.62 19.63 19.47 19.65 19.72 19, I (mgs.) 123 124 124 123 124 125 124 %E (I) 93 94 94 93 94 95 94 Cd (mgs?) 55 56 55 55 55 55 55 %E (Cd) 95 97 95 95 95 95 95 1 I 1 II I i I 2 2 2 2 2 2 2 Aaueous Layer 00 1 ". I (Left) 7.8 7.9 7.9 7.9 • 7.1 7 I (Orlg.) 132 132 132 132 132 132 132 %I (Left) 5.9 6 6 6 5.9 5.4 5 Cd (Left) 3.3 2.3 2 . 8 3.5 3 3 3 Cd (Orlg.) 58 58 58 58 58 58 58 %Cd (Left) 5 3 5 5 5 5 5 ApH + 0.4 +0 . 2 + 0.3 +0 . 2 +0.3 + 0.3 + 0 * obtained by difference 68 TABLE XIV - F EFFECT OF HYDROGEN ION CONCENTRATION ON THE EXTRACTION OF IODIDE F: pH « 6 Non-aaueous Layer I 2 3 4 5 6 Av. vKIOj(4) 20.45 20.47 20.41 20.59 20.47 20.52 20.41 I (mgs.) 123 123 123 124 123 124 123 %E (I) 94 94 93 94 94 94 94 Cd (mast) 56 56 56 56 56 56 56 %E (Cd) 97 97 97 97 97 97 97 1 1 1 1 1 1 1 1 2 2 2 2 2 2 2 Aaueous Laver I (Left) 7.9 7.9 7.9 ^ 7.8 7.9 7.8 7.9 I (Orig.) 132 132 132 132 132 132 132 %I (Left) 6 6 6 5.9 6 5.9 6 Cd (Left) 2 . 6 2.5 2.5 2 . 6 2 . 6 2 . 6 2 . 6 Cd (Orig.) 58 58 58 58 58 58 58 %Cd (Left) 3 3 3 3 3 3 3 ApH -3.2 -3 .0 -3.2 -3 .4 -3 .2 -3.3 -3.2 * obtained by difference 69 TABLE XV - A EFFECT OF CADMIUM TO IODIDE RATIO ON THE EXTRACTION OF IODIDE A: Cd/I - 1/1 Non-aquaoui Laver vKI0 3 (6) 21.87 21.83 21.83 21.87 21.83 21.78 2 1 . 8 I (mgs,) 132 132 132 132 132 132 132 %E (I) 96 95 95 96 95 95 95 Cd (mgs?) 65 64 63 64 64 64 64 %E (Cd) 51 50 49 50 50 50 50 Cd 1 1 1 1 1 1 1 I 1 . 8 1 . 8 1 . 8 1 . 8 1 . 8 1 . 8 1 . 8 Aaueous Laver I (Left)* 6 6 6 6 6 6 6 I (Orig.) 138 138 138 138 138 138 138 %I (Left) 4 5 5 4 5 5 5 Cd (Left) 64 64 65 64 64 64 64 Cd (Orig.) 128 128 128 128 128 128 128 %Cd (Left) 49 50 51 50 50 50 50 * obtained by difference 70 TABLE XV - B EFFECT OF CADMIUM TO IODIDE RATIO ON THE EXTRACTION OF IODIDE B: Cd/I - 2/1 . 1 2 3 4 5 6 Av, v*ao 3 (6 ) 22.87 22.74 22.63 21.37 21.31 21.27 I (mgs.) 138 138 137 129 129 128 %E (I) 92 92 91 93 93 93 92 Cd (mgs?) 62 6 6 64 54 54 54 %E (Cd) 23 25 24 2 1 21 2 1 22 Cd 1 1 1 1 1 1 1 1 2 . 0 1 . 8 1.9 2 . 1 2 . 1 2 . 1 2 . 0 Aaueous Laver I (Left)* 12 12 13 9 9 10 I (Orig.) 150 150 150 138 138 138 %I (Left) 8 8 9 7 7 7 8 Cd (Left) 206 2 0 2 204 208 208 208 Cd (Orig.) 268 268 268 262 262 262 %Cd (Left) 77 75 76 79 79 79 78 * obtained by difference 71 TABLE XV - C EFFECT OF CADMIUM TO IODIDE RATIO ON THE EXTRACTION OF IODIDE C: Cd/I - 4/1 Non-aaueous Laver 1 2 3 4 5 Av. VKI03 (6) 18.09 18.51 18.48 18.42 18.36 18.: I (mgs.) 109 1 1 2 1 1 2 1 1 1 111 i l l %E (I) 8 8 90 89 89 89 89 Cd (mgs?) 75 77 77 77 79 77 %E (Cd) 19 2 0 2 0 2 0 2 0 2 0 Qd 1 1 1 1 1 1 I 1.3 1.3 1.3 1.3 1 . 2 1.3 Aqueous Layer I (Left)* 16 13 13 14 14 14 I (Orig.) 125 125 125 125 125 125 %I (Left) 12 10 11 11 11 11 Cd (Left) 316 314 314 314 312 314 Cd (Orig.) 391 391 391 391 391 391 %Cd (Left) 81 80 80 80 80 80 * obtained by difference 72 TABLE XVI - A EFFECT OF INITIAL CONCENTRATION OF CADMIUM IODIDE ON THE EXTRACTION OF IODIDE A: 0.001 MCdI2 ______Non-aaueous Laver ______Ay* 1.97 1.97 VN»282°3 1.98 1.97 1.97 I (mgs.) 2.12 2.10 2.10 2.10 2.10 %E (I) 61 61 61 61 61 Cd (mgs?) 0.92 0.92 0.92 0.91 0.92 %E (Cd) 60 60 60 60 60 Qd 1 1 1 1 1 I 2 2 2 2 2 Aqueous Laver I (Left)* 1.33 1.35 1.35 1.35 1.34 I (Orig.) 3.45 3.45 3.45 3.45 3.45 %I (Left) 39 39 39 39 39 Cd (Left) 0.61 0.61 0.61 0.62 0.61 Cd (Orig.) 1.53 1.53 1.53 1.53 1.53 %Cd (Left) 40 40 40 40 40 N a 2 S2 0 3 titer* 1.07 mgs. I * obtained by difference 73 TABLE XVI - B EFFECT OF INITIAL CONCENTRATION OF CADMIUM IODIDE ON THE EXTRACTION OF IODIDE B: 0.003 M Cdl 2 ______Non-aaueous _Laver ______ VNa2 S2 0 3 6.32 6.23 6.24 6.21 6.25 I(mgs.) 6.76 6.67 6 . 6 8 6.64 6.69 %E (I) 78 77 77 77 77 Cd (mgst) 1.6 1.6 1.6 1.6 1.6 %E (Cd) 53 53 53 53 53 1111 I 3.7 3.7 3.7 3.7 3.7 Aqueous Laver I (Left)* 1.88 1.97 1.96 2.00 1.95 I (Orig.) 8.64 8.64 8.64 8.64 8.64 %l (Left) 22 23 23 23 23 Cd (Left) 1.4 1.4 1.4 1.4 1.4 Cd (Orig.) 3 3 3 3 3 %Cd (Left) 47 47 47 47 47 Na2 S2 0 3 titer =1.07 mgs. I * obtained by difference 74 TABLE XVI - C EFFECT OF INITIAL CONCENTRATION OF CADMIUM IODIDE ON THE EXTRACTION OF IODIDE C: 0.006 M Cdl 2 Non-aaueous Laver 1 2 3 4 Av. Vn «2s 2° 3 14.02 14.00 14.11 14.07 14.05 I (mgs.) 15.0 15.0 15.1 15.1 15.0 %E (I) 87 87 87 87 87 Cd (mgs?) 5.0 5.1 5 .0 5.2 5.1 %E (Cd) 65 6 6 65 6 8 6 6 JL. I 2.7 2 . 6 2^7 2 . 6 2 . 6 Aoueous Laver I (Left)* 2.3 2.3 2 . 2 2 . 2 2 . 2 I (Orig.) 17.3 17,3 17.3 17.3 17.3 %I (Left) 13 13 13 13 13 Cd (Left) 2.7 2.6 2.7 2.5 2.6 Cd (Orig.) 7.7 7.7 7.7 7.7 7.7 %Cd (Left) 35 34 35 32 34 Na_S«0_2 2 3 titer = 1.07 m gs. I * obtained by difference 75 TABLE XVI - D EFFECT OF INITIAL CONCENTRATION OF CADMIUM IODIDE ON THE EXTRACTION OF IODIDE D; 0.01 M Cdl 2 ______N w - t w w ______3 4 Av. v N«2820 3 28.51 28.81 28.71 28.69 28.68 I (mgs.) 31.3 31.6 31.5 31.5 31.5 %E (I) 90 91 91 91 91 Cd (mgit) 10.3 10.3 10.3 10.4 10.3 %E (Cd) 67 67 67 68 67 fid 1 JL 1 _L_ • I 2.7 2.7 2.7 2.7 2.7 Aaueous Layer I (Left)* 3.4 3.1 3.2 3.2 3.2 I (Orig.) 34.7 34.7 34.7 34.7 34.7 %I (Left) 10 9.0 9.0 9.0 9.0 Cd (Left) 5.1 5.1 5.1 5.0 5.1 Cd (Orig.) 15.4 15.4 15.4 15.4 15.4 %Cd (Left) 33 33 33 32 33 Na2S2 0 3 titer = 1 . 1 0 mgs. I * obtained by difference 76 TABLE XVI - E EFFECT OF INITIAL CONCENTRATION OF CADMIUM IODIDE ON THE EXTRACTION OF IODIDE E: 0.03 M Cdl 2 Non-aaueous Laver 2 3 4 ? 6 Av. VKI03 (S) 10.74 10.64 10.80 10.80 10.83 (10.40) 1 0 . 8 C I (mgs.) 65.5 6 6 . 1 65.8 65.8 6 6 (63.4) 65.8 %B (I) 94 95 95 95 95 (91) 95 Cd (mgs?) 22.9 21.5 2 0 . 8 21.5 20.7 (21.5) 21.5 %E (Cd) 74 70 6 6 70 67 (70) 70 fid JL o . 1 ( J L ) JL I 2.5 2.7 2 . 8 2.7 2 . 8 1 2 . 6 1 2.7 AflWOTi. L m c . I (Left)* 4.0 3.4 3.7 3.7 3.5 (6 . 1) 3.7 I (Orig.) 69.5 69.5 69.5 69.5 69.5 69.5 69.5 %I (Left) 6 5 5 5 5 (9) 5 Cd (Left) 7.9 9.3 10 9.3 1 0 . 1 (9.3) 9.3 Cd (Orig.) 30.8 30.8 30.8 30.8 30.8 30.8 30.8 %Cd (Left) 26 30 32 30 33 (30) 30 * obtained by difference 77 TABLE XVI - F EFFECT OF INITIAL CONCENTRATION OF CADMIUM IODIDE ON THE EXTRACTION OF IODIDE F: 0.1 M Cdl 2 Noi^aqueous Laver 1 2 3 4 5 6 Av VKI03 (7) 40.24 40.08 40.01 39.97 39.95 39.97 40 I (mgs.) 257 256 255 255 255 255 256 %E (I) 97 96 96 96 96 96 96 Cd (mgs?) 89 8 8 8 8 8 8 8 8 8 8 8 8 %E (Cd) 75 75 75 75 75 75 75 _L. 1 1 1 1 I 2 . 6 2 . 6 2 . 6 2 . 6 2 . 6 2 . 6 7 Aaueous Layer I (Left)* 9.0 1 0 11 11 11 11 10 I (Orig.) 266 266 266 266 266 266 266 %I (Left) 3 4 4 4 4 4 4 Cd (Left) 29 30 30 30 30 30 30 Cd (Orig.) 118 118 118 118 118 118 118 %Cd (Left) 25 25 25 25 25 25 25 * obtained by difference 78 TABLE XVII-A EFFECT OF TEMPERATURE ON THE EXTRACTION OF IODIDE A: 25°C ______Non-agu+bus Layer______1 2 3 4 5 6 Av. vKI0 3 (6) 22.18 22.19 2 2 22.17 22.16 2 2 . 2 0 2 2 1 (mgs.) 134 134 134 134 134 134 134 %E (1) 97 97 97 97 97 97 97 Cd (mgs?) 41 42 41 41 41 42 41 %L (Cd) 67 69 67 67 67 69 6 8 04 1 1 1 I 2.9 2 . 8 2 2.9 2.9 2 . 8 2 Aaueous Laver I (Left)* 4 4 4 4 4 4 4 I (Orig.) 138 138 138 138 138 138 138 %I (Left) 3 3 3 3 3 3 3 Cd (Left) 2 0 19 2 0 2 0 2 0 19 2 0 Cd (Orig.) 61 61 61 61 61 61 61 %Cd (Left) 33 31 33 33 33 31 32 * obtained by difference 79 TABLE XVII - B EFFECT OF TEMPERATURE ON THE EXTRACTION OF IODIDE B: 40°C 1 2 3 4 5_ 6 Av. vK1O 3 0 ) 19.53 19.34 19.30 19.51 19.39 19.59 19.44 I (mgs.) 124 1 2 2 1 2 2 124 123 124 123 %E (I) 94 93 93 94 93 94 94 Cd (mgs?) 40 39 39 40 39 40 40 %E (Cd) 69 67 67 69 67 69 6 8 £ 4 I 2.6 2.8 2.8 2.8 2.6 2.8 2.8 Aaueous Lever I (Left)* 8 1 0 1 0 8 9 8 9 I (Orig.) 132 132 132 132 132 132 132 %I (Left) 6 7 7 6 7 6 6 Cd (Left) 16 19 19 18 19 18 16 Cd (Orig.) 58 58 58 58 58 58 58 %Cd (Left) 31 33 33 31 33 31 32 * obtained by difference 80 20 1 5 Micrograms of Cadmium Graph No. 1. Spectrophotometrlc Calibration Curve SELECTED BIBLIOGRAPHY 1 • Alberty, R. A* and Kln^, E• L. , Am» Chem. Soc», 7^* 517 (1951) • 2. Alcock, K., Grimley, S. S., Healy, T. V., Kennedy, J. and McKay, H. A. C ., Trans. Far. Soc. , 52, 39 (1956). 3. Basolo, F. and Pearson, R. G ., Mechanisms of Inorganic Reactions,, New York: J. W iley and Sons, I n c ., 1957. 4. Briscoe, H. V. A. and Holt, P. F .. Inorganic Mlcroanalysls, London: Edward Arnold and C o ., 1950. 5. Carlton, J. K.. Anal. Chem. . 22, 1072 (195 0). 6 . Coursier, J., Hure, J., and Platzer, R .. Anal. Chlm. Acta, .J3., 379 (1955). 7. Ducret, L. and Drouillas, M., Anal. Chlm. Acta, 21, 8 6 (1959). 8 . Ducret, L. and Ratouis, M., Anal. Chlm. Acta, 21, 91 (1959). 9. Feigl, F., Snot Tests In Inorganic Analysis, New York: Elsevier Publishing C o., 1958. 10. Fuoss R. M. and Accascina, F.. Electrolytic Conductance, New York: Interscience Publishers, Inc., 1959. 11. Harris, L. E. , Ph. _D_. Dissertation. L. S. U., 1961. 12. Hollingshead, R. G. W ., Oxine And Its Derivatives, Vol. I, London: Butterworths Scientific Publications, 1954. 13. Leden. I., Z. phvslk. Chem., A 188, 160 (1941). 14. Moeller, T., Inorganic Chemistry, An Advanced Textbook, New York: J. Wiley and Sons, In c ., 1952. 15. Morrison, G. H. and Freiser, H., Solvent Extraction in Analytical Chemistry, New York: J. Wiley and Sons, Inc., 1957. 16. Rynasiewicz, J., Sleeper, M. P., and Ryan, J. W ., Anal. Chem. , 26., 935 (1954). 82 17. Sherrill, M. S., Z. phvslk. Chem., 43, 705 (1903). 18. Sneed, M. C. and Brasted, R. C. Comprehensive Inorganic Chemistry, Vol. IV, Princeton, N. D. Van Nostrand C o., Inc., 1955. 19. Talmni, I. K. and Lai, M .. Anal, Chlm. Acta. 17, 367(1957). 20. Talmni, I. K. and Lai, M.. Anal. Chim. Acta,21,105 (1959). 21. Weisz, H .. Mlkrochlm. Acta, 140, 376, 785 (1954). 22. West, P. W. and Mukherji, A. K., Anal. Chem., 31, 947 (1959). 23. West, P. W. and Tron, F., Unpublished Studies. 24. West, P. W ., Vick, M ., and Le Rosen, A. L ., Qualitative Analysis and Analytical Separations, New York: The MacMillan C o., 1953. 25. Zechner, S., Mlkrochlm. Acta, 2, 159 (1957). VITA Apolinar S . Lorica was born on July 22 , 1927 In Albay, one of the provinces of the Philippines. He completed the first three years of ele mentary school in Albay and the last four years of it in Manila, where he graduated from Jose Rlzal Elementary School in March. 1940. He attended Florentlno Torres High School, also in Manila, and was graduated in November. 1946 after a two-year Interruption brought about by the Second World War . Then he decided to take up chemistry at the University of the Philippines (U. P.). In his senior year at the U.P.. he served as a student assistant teaching laboratory courses in chemistry. He became a full-fledged member of the faculty shortly after his graduation with a B. S. Chemistry degree in April, 1954. He came to the Louisiana State University (L. S. U.) in September, 1957 for graduate studies in chemistry under a fellowship from the University of the Philippines and a graduate assistantship from Louisiana State University. He is a member of the honor societies of Phi Lambda Upsllon and Sigma XI. Now he is a candidate for the degree of Doctor of Philosophy. 83 EXAMINATION AND TMESIS IffO R Candidate: Apolinar S. Lorlea Major Field: Chenlatry Title of Theaia: STUDIES ON THE EXTRACTABILITY OF ANIONS Approved: TaJ. Major Profeasor and Chairman Dein>«flhe (lraduateSchool EXAMINING COMMITTEE: Date of Examination: February 1, I96I