This dissertation has been 62—754 microfilmed exactly as received
CRABB, Norman Theodore, 1933- POLARIZATION STUDIES OF A DROPPING GALLIUM ELECTRODE.
The Ohio State University, Ph.D., 1961 Chemistry, analytical
University Microfilms, Inc., Ann Arbor, Michigan Copyright by
Norman Theodore Crabb
1 9 6 2 POLARIZATION STUDIES OF A
DROPPING GALLIUM ELECTRODE
DISSERTATION
Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University
By
Norman Theodore Crabb, B. S., M. Sc.
The Ohio State University 1961
Approved by
Department or unemistry PREFACE
Although the chemical literature contains a multitude of entries for the dropping mercury electrode, the number of entries for dropping gallium electrode was found to be limited to two.
Therefore, the initial purpose of this study was development of a dropping gallium electrode that could be used to contribute to a better understanding of the electrochemistry of gallium.
When a successful dropping gallium electrode had been developed, it was decided to employ this electrode in investigations of the polarization behavior of gallium.
Finally, it is hoped that this investigation may be used as a basis for the extension of the study of the electrochemical properties of gallium, and contribute toward the determination of the composition of gallium rich alloys.
List of Abbreviations
S.C.E, saturated calomel electrode N.C.E. normal calomel electrode N.H.E, normal hydrogen electrode °C. degrees centigrade C bulk concentration in moles per liter M moles per liter D.G.E, dropping gallium electrode D.M.E. dropping mercury electrode p.s.i. pounds per square inch
All potentials in this study with the D.G.E. are given with respect to the S.C.E. at 35°C.
ii Ac knowledgment s
I hereby wish to express my thanks to Dr. William MacNevin for his interest and guidance throughout the course of this research. I also wish to express my appreciation for the financial aid available to me through the Department of Chemistry in the form of teaching and research assistantshipb .
In addition, I express my gratitude to my wife, Nancy, for her patience and understanding and for the typing of this dissertation.
iii TABLE OF CONTENTS
Page
PREFACE ii
List of Abbreviations Ac knovledgments
LIST CF TABLES vi
LIST CF ILLUSTRATIONS vii
INTRODUCTION 1
Chemical Characteristics of Gallium Analytical Chemistry of Gallium Qualitative Tests for Gallium Electrochemistry of Gallium The Standard Electrode Potential of Gallium Polarographic Studies of Gallium Solubility of Metals in Gallium
STATEMENT CF P ROBLEM...... l4
Chapter I . EXPERIMENTAL PROCEDURE 15
Dropping Gallium Electrode The Dropping Gallium Electrode Apparatus Assembly, Operation, and Cleaning of Dropping Gallium Electrode Electrolysis Cells Reagents
II. HYDROLYSIS OF GALLIUM 28
III. OXIDATION CF GALLIUM 41
Electrochemical Oxidation of the Dropping Gallium Electrode Chemical Oxidation of the Dropping Gallium Electrode
iv TABLE CF CONTENTS (continued)
Page
Chapter IV. CATHODIC CURRENT-VOLTAGE CURVES AT THE DROPPING GALLIUM ELECTRODE...... 62
Reduction of Gallium Nitrate The Reduction of Hydrogen Ion The Cathodic Curve Produced in 2 Molar Sodium Perchlorate with Added Sodium Broraate
V. CURRENT-VOLTAGE CURVES AND CONCENTRATION POLARIZATION; CRITERIA FOR REVERSIBILITY...... 89
Reversibility
VI. ANOMALOUS ELECTROCAPILLARY CURVES AT THE DROPPING GALLIUM ELECTRODE...... 93
VII. THIOCYANATES...... 115
Reactions of Gallium Salts In 7*5 M Potassium Thiocyanate
SUMMARY AND CONCLUSIONS...... 123
LITERATURE CITED ...... 126
AUTOBIOGRAPHY...... 131
v LIST CF TABLES
Table Page
1. Amphoteric Behavior of* Gallium S a l t s ...... 3
2. Polarography of Gallium on Dropping Mercury Electrode...... 11
3- Analysis of Gallium Perchlorate 6-hydrate...... 27
b. Reciprocal Slopes of Anodic Current-Voltage C u r v e s ...... U5
5. Current-Concentration Dependence for Gallium Nitrate at the D.G.E. in Acidified 1 M Potassium Chloride...... 67
6. Square Root of Applied Pres sure-Current Dependence for 0.0003 M Hydrochloric Acid in 1 M Potassium Chloride Measured at -1.3 Volts (S.C.E.)...... 73
7- Current-Concentration Dependence for the Reduction of Hydrogen at the D.G.E. in 1 M Potassium Chloride...... 78
8. Current-Concentration Dependence of Broraate at the D.G.E...... 82 1, 9. (p.s.i.)2 Versus d Dependence for 0.00015, 0.00025, and 0.000*+5 M Sodium Broraate ...... 86
10. The Effect of Pressure on the Drop Time of the Dropping Gallium Electrode ...... 97
11. Data for m ^ t ^ a t Open Circuit for the D.G.E...... 100
12. Drop Time Data for Electrocapillary Curves Obtained with the D.G.E. at 15 p.s.i. Nitrogen Pressure...... 101
13. Electrocapillary Data Obtained by Muratzejev and Gorodetzkaya...... 112
vi LIST CF ILLUSTRATIONS
Figure Page
1. Dropping Gallium Electrode Apparatus...... 18
2. Electrolysis Cell ...... 24
3. Current-Voltage Curves of Gallium Nitrate and Hydrochloric Acid in 1 M Potassium Chloride ...... 30
4. Typical Current-Voltage Curves Illustrating the Complexity Produced as the Result of Hydrolysis of Gallium Nitrate in 1 M Potassium Chloride...... 32
5. An Example of the Complicated Current-Voltage Curve Obtained When the Concentration of Added Hydrochloric Acid vas Insufficient to Suppress Hydrolysis of Gallium Nitrate in 1 M Potassium Ch l o r i d e...... 35
6. Current-Voltage Curves Illustrating the Hydrolysis of Gallium Nitrate in 1 M Potassium Chloride...... 37
7. Anodic Dissolution Curves of Gallium in 1 Molar Potassium Chloride ...... 43
8. Residual Current Curves at the Dropping Gallium Electrode in 7.5 M Potassium Thiocyanate ...... 47
'. The Influence of Nitrate Ion at the Dropping Gallium Electrode in 7-5 M Potassium Thiocyanate...... 50
10. The Influence of Nitrate Ion at the Dropping Gallium Electrode in 7-5 M Potassium Thiocyanate...... 53
11. The Effect of Gelatin on the Nitrate Maximum in 7.5 M Potassium Thiocyanate...... 55
12. Current-Voltage Curves Obtained in 2 M Sodium Perchlorate with Added Sodium Bromate, Potassium Nitrate, Sodium Nitrite ...... 58
13. An Example of the Interfering CathodicCurrent Observed in 7*5 M Potassium Thiocyanate ...... 63
vii LIST CF ILLUSTRATIONS (continued)
Figure Page
3.1+. Combined Anodic-Cathodic Curve for 0.000126 M r.alllim Nitrate and 0.0065 M Hydrochloric Acid in 1 M Potassium Chloride...... 65
15. Current-Concentration Dependence for Gallium Nitrate in Acidified 1 M Potassium Chloride with 0.002$ Triton X-100 ...... 68
16. Current-Voltage Curves Resulting from the Reduction of Hydrogen Ion in 1 M Potassium Chloride plus 0.002$ Triton X - 1 0 0 ...... 7^
17. Square Root of Applied Pressure-Current Dependence for Hydrogen Ion in 1 M Potassium Chloride Measured at -1.3 Volts (S.C.E.) at 35°C...... 76
18. Current-Concentration Dependence for Hydrogen Ion in 1 M Potassium Chloride...... 79
19. Current-Concentration Dependence for Bromate Ion in 2 Molar Sodium Perchlorate...... 83
20. Square Root of Applied Pressure-Current Dependence for Sodium Bromate in 2 M Soditm Perchlorate Measured at -1.1+ Volts (S.C.E.) at 35°C...... 87
21. Drop Time-Applied Pressure Relationship for the D.G.E. in 2 M Sodium Perchlorate ...... 9^
22. Electrocapillary Curves by the Drop Time Method in 7*5 M Potassium Thiocyanate...... 102
2 3 . Electrocapillary Curves by the Drop Time Method in 2 M Sodium Perchlorate...... 10l+
2l+. Electrocapillary Curves by the Drop Time Method in 1 M Potassium Chloride...... 107
25. Electrocapillary Curves by the Drop Time Method in 1 M Potassium Chloride...... 109
26. A Plot of the Electrocapillary Data Obtained by Muratzejew and Gorodetzhaya...... 113
27. Ultraviolet Absorption Spectra of the Solutions Above the Yellow Precipitates Formed in 7-5 M Potassium Thiocyanate ...... 121
viii INTRODUCTION
The existence of "ecka-aluminum," the element given the name gallium, was predicted by D. I. Mendeleeff^ in 1871.
XD. I. Mendeleeff, Ann., SuTrolanentband 8, 2, 196 (1 8 7 1).
L. de Boisbaudran2 also predicted the existence of gallium from a spectral line observed while examining the spectra of some zinc ores.
2L. de Boisbaudran, Ann. Chim. et Phys., 10, 100 (l877)*
In 1877, de Boisbaudran3 isolated the metal by electrolysis of caustic
3L. de Boisbaudran, Chem. News, 35. 1^8, 157, 167 (1877). solutions of zinc slag. De Boisbaudran gave the element its name, gallium, after the Latin name of his native land, Gaul.
Gallium is one of three metals, mercury, gallium, and cesium, that are liquids at room temperature. The following list of physical
1 2 properties of gallium is only a portion of an extensive list of properties enumerated by Hampel.**
**C. A. Hampel, "Rare Metals Handbook," Rheinhold Publishing Company, New York, 195^, PP* l1+7-159-
melting point 29.75°C.
boiling point 2000°C.
density 20°C. 5-907 29.85°C. 6.095
viscosity at 97-7°C. - O.Ol6l2 poise
surface tension, H^, 30°C. 735 dynes/cm.
interfacial tension 613 dynes/cm. in 1 M KC1 plus 0.1 M HC1
An extensive bibliography of gallium has been prepared and main tained by P. de la Breteque. 5 > 6 This comprehensive bibliography is
5p. de la Breteque, "Galllum-Bjbliographle," Neuhausen am Rheinfall, Switzerland, 1958.
^P. de la Breteque, "Gallium Bulletin Bibliographic," Marseille, France, i9 6 0. supplemented by several good reviews on gallium metal, gallium salts, and solutions of gallium salts.
P. de la Breteque, "Etudes Sur le Gallium," Lausanne Imprimerie Vandoise, Marseille, France, 1956.
^E. Einecke, "Das Gallium," Verlag von Leopold Voss, Leipzig, 1937.
%.S. Atomic Energy Commission, "Liquid Metals Handbook," Washington, Government Printing Office, 1950.
10W. D. Wilkinson, "Properties of Gallium," ANL-1+109 (1 9 5 8). 3
Chemical Characteristics of Gallium
Gallium is a member of group III, sub-group B of the periodic family of elements. The unusual liquid range of gallium gives It an
anomalous physical behavior in the al tmlnum group. Chemically,
gallium Is similar to aluminum. Moeller and Klng^- have summarized
^T. Moeller and G. L. King, J. of Physical and Colloid Chem.. 5k, 999 (1950).
the amphoteric behavior of some of the common gallium salts.
TABLE 1
AMPHOTERIC BEHAVIOR CP GALLIUM SALTS
pH of precipitation pH of dissolution of incidence precipitate
GaBr^ 5-15 1 1 .5 0
GaCl^ 5.05 1 1 .2k
Ga(N0 3 ) 3 5 . H li.o 6
Ga(SO^) 2 .8 0 1 1 .6 2
Gallium is only known to exhibit the plus three oxidation state
in aqueous solutions. Gallium (III) chloride has been prepared by
J. A. Bridgman in a manner similar to the preparation of aluminum
A. Bridgman, J. Am. Chem. Soc., ^0 , 1531 (1918).
trichloride. Gallium metal was exposed to anhydrous chlorine gas,
whereupon the metal burst into flame and the gallium trichloride sublimed onto the vails of the glass vessel. E. H. S w i f t 1 3 has shown
13E. H. Swift, J. Am. Chem. Soc., h6, 237*+ (192*0* that the covalent character of galllun trichloride permits extraction of toe salt from acidified aqueous solutions into diethylether.
Other oxidation states of gallium have been stabilized by compound formation. Einecke® reported the preparation of gallium (il) sulfide by reacting gallium (III) sulfide with hydrogen at 800° centi grade. Gallium (il) chloride has been prepared by Laubengayer and
Schirmer1^ by the reduction of gallium (ill) chloride on hot gallium
li4A. W. Laubengayer and F. B. Schlrmer, J. Am. Chem. Soc., 62, 1 5 7 8 (19U0 ). metal. An interesting gallium (i) salt has been prepared by Corbett and M c M u l l e n .^5 The reaction of gallium metal and gallium (ill)
^ J. D. Corbett and R. K. McMullen, J . Am. Chem. Soc., 7 8, 2906 (1 9 5 6). chloride in fused aluminum trichloride produces the compound
GatGaCl^) which is a white solid melting at 175° centigrade.
Analytical Chemistry of Gallium
The main source of supply of gallium is from zinc ores. These oreB are found to contain about 0.07 percent gallium. The caustic
"iron muds" obtained during the processing of the zinc ores are dehydrated, dissolved in hydrochloric acid, and the gal H im and iron chlorides are extracted with diethylether. After separation of the 5 ether and aqueous phases, the ether is distilled from the chlorides.
Solution of the chlorides in strong caustic removed the iron con
taminants. Electrolysis of the caustic solution produces a good grade of gallium. The metal is further purified by recrystallization.
Aqueous solutions of gallium may be analyzed quantitatively for
gallium by precipitating gallium as the oxide. Lundell and Hoffman^
1 ^G. E. F. Lundell and J. F. Hoffman, J. Research N.B.S.. 15, *+09 (1935).
recommend precipitation of the hydrous oxide with tannin or ammonium
hydroxide with subsequent ignition to gallium oxide.
The 8-hydroxyquinoline complex of gallium has been used by
Geilman and Wrigge^ to determine gallium gravimetrically. Warm acid
17W. Geilman and W. Wrigge, Z . anorg. u . algem. Chem., 209. 129 (1932).
solutions of gallium and 8-hydroxyquinoline are neutralized with
sodium acetate. In the region of pH 7 the yellow gallium quinolate
precipitates from the solution. 1 8 Moeller and Cohen observed the ultraviolet spectrum of gallium
^T . Moeller and A. J. Cohen, Anal. Chem.. 22, 686 (1 9 5 0).
quinolate at *+00 myi using chloroform as the solvent. The fluorescent
spectrum of gallium quinolate has been examined by Collat and Rogers^
^ J . W. Collat and L. B. Rogers, Anal. Chem., 2 7, 9 6 (1955).
as a method of determining gallium in the presence of aluminum. 6
Milner2® has shown that small concentrations of gallium, 0.2 to
2®G. W. C. Milner, Analyst. 8 0, 77 (1956).
O.U milligrams per liter, can he determined hy a direct titration with ethylenediamine-tetraacetic acid. Gallocyanine, C1,-H^pNpO,., is used as the indicator for the titration. Since the indicator forms colored compounds with other metal cations, the gallium must firBt be separated by extraction of the chloride into ether or by extraction from thiocycnate solutions into a mixture of ether and tetrahydro- furan.
Potassium ferrocyanide has been used as a precipitant for gallium by Browning and Porter2^ and by Fetter and Swinehart22 for the
2^P. E. Browning and L. E. Porter, Am. J. Sci., j+U, 22 (1917).
^ N . R. Fetter and D. F. Swinehart, Anal. Chem., 28, 122 (1956). amperometric titration of gallium. The precipitation is carried out in an acid solution, pH 2, and proceeds at a rate sufficiently slowly that a temperature of 50°C. is required during the amperometric titration.
Gallium has been determined qualitatively and quantitatively by the spectrographic method since its discovery by de Boisbaudran. 2
L. W. Str o c k 2 3 made an interesting variation of the spectrographic
2 3l. w. Strock, Metals Technol., 12, April (I9U5). method by using beryllium oxide as an internal standard and estimating the concentration of gallium by way of the spectral line at 29U3 S. 7
Qualitative Testa for Gallium
A red-brown precipitate Is formed if an acidified solution of gallium is added to an acidified solution of potassium ferrocyanide, potassium bromate, and manga nous chloride. Browning and Porter2^ reported the reaction as a qualitative test for gallium but did not confirm the composition of the precipitate. Oh H. Onishi has shown that the rhodamine-B complex of gallium
2**H. Onishi, Anal. Chem.. 2J, 832 (1955). can be extracted into benzene, where a pink fluorescence can be observed upon radiation by ultraviolet light. Titanous chloride must be added to the test solution to remove interfering oxidizing agents.
Electrochemistry of Gallium
Gallium metal was first prepared by L. de Boisbaudran^ by elec trolysis of caustic solutions of zinc slag. Other researchers2^ *26,27
25L. Schucht, Chem. Ztg.. k, 292 (1 8 8 0).
26L. Ehrlich, ibid.. % 28 (1 8 8 5).
2?G. Kunnert, ibid.. 2, 1826 (1 8 8 5). following the method of de Boisbaudran have reported varying conditions for the separation of the metal from caustic solutions by electrolysis. An important observation made by Uhler and Browning2®
28h. s. Uhler and P. E. Browning, Am. J. Sci.. k2, 389 (1 9 1 6). concerning the electrolysis of caustic galliate solutions was that the presence of nitrates interfered with the electrochemical reduction to gallium metal. Shebba and Pugh2^ employed electrolysis of gallim
2% . Shebba and W. Pugh, J. Chem. Soc., 1371 (1937)- from basic solutions as a means of obtaining pure gallium for their studies of amalgams of gallium.
The electrochemical preparation from acid solution has been reported by Dennis and B r i d g m a n 3° and by Richards and B o y e r , 31 however,
M. Dennis and J. A. Bridgman, J . A m . Chem. Soc., ^0, 1531 (1918).
3*rr. V. Richards and S. Boyer, ibid.. Ul, 133 (1 9 1 9). the electrolysis of basic galliate solutions remains as the industrial method for preparation of the metal.
Ammonium sulfate was investigated by Reichel3^ as the electrolyte
32E. Reichel, Z. Anal. Chem.. 8j, 321 (1933).
for the electrodeposition of gallium. The use of liquid ammonia as the plating bath led to the postulation that the electrochemical
reduction of Ga(lll) to Ga(o) proceeds stepwise through a OafI) ■5 ■a intermediate-^ in liquid ammonia. Fused aluminum chloride was
33a. D. McElroy, J. Kleinberg, and A. W. Davidson, J. Am. Chem. Soc., Jh, 736 (1952). 9 employed by Verdieck and Yntema3^ as a medium for the electrochemical
Verdieck and L. F. Yntema, J. Phys. Chem.. **8, 268 (19^*0• preparation of gallium metal.
An electrochemical separation of gallium Into a mercury pool was performed by Bock and Hackstein3^ in vhich 100 milligrams of
3^R. Bock and K. Hackstein, Z. Anal. Chem., 138, 339 (1953). gallium was completely deposited into the mercury pool when 0 .0 5 M sulfuric acid and 20 grams per liter of potassium sulfate was used as the electrolyte.
The Standard Electrode Potential of Gallium
A review of the electrochemistry of gallium by Fogg3^ gives
36H. Fogg, Trans. Electro chem. Soc., 6 6, 107 (193*0. credit to Regnauld3^ for being one of the first investigators of the
3^J. Regnauld, Conrpt. rend., 8 6, 577 (1 8 7 8). standard electrode potential of gallium. A value of -O.25 volts
(N.H.E.) was reported for the standard electrode potential of gallium by Richards and Boyer.3^ These authors used liquid and solid gallium electrodes and observed, but did not correct for, hydrolysis of gallium ion. S. von Bergkampf3® reported a value of-0.52 volts
3®S. von Bergkampf, Z . Electrochem.. 3 8, 8^7 (1932). 10
(N.H.E.) for the standard potential of the gallium electrode. A comprehensive study of the standard potential of the gallium electrode was made by Stelling.39 He reported in this study that he vaB unable
^O tto Stelling, Z. Electrochem.. Ul, 712 (1935). to obtain null-point potentials in gallium chloride solutions using either liquid or solid gallium electrodes. Stelling reported values of -0.909 to -0.902 (N.H.E.) for the standard potential of gallium as the result of studies of a dropping gallium amalgam electrode. In addition, he concluded that gallium was more noble in the liquid state. Challenger^0 reported a value of -0 .57(8) volts (N.H.E.) for
G. Challenger, Ph.D. dissertation, Harvard University, 19^5* the standard electrode potential of gallium from his studies of a dropping gallium electrode. A more recent study of the standard electrode potential of gallium is that reported by Saltman and
Nachtrieb^ In which the cell Ga/GaX^, HX/Pt(H^) was used. These
Ulv. Saltman and N. Nachtrieb, J. Electrochem. Soc., 100, 126 (1953)- authors observed that liquid gallium became passive more rapidly than did solid gallium. A solid gallium electrode was potentiometrically reversible with a standard potential of -O.5 6 volts (N.H.E.) corrected for hydrolysis. Polanographic Studies of Gallium
The following table is a summary of the available literature concerning the polarographic investigations of gallium.
TABLE 2
POLAROGRAFHY OF GALLIUM ON DROPPING MERCURY.ELECTRODE
-El (S.C.E.) at 25°C. or Electrolyte as indicated Reference
ho 1. HC1 solution no reaction S. Takagl. J. Chem. Soc.. London, 301 (1928).
2. 1 M NHlCI f 1HNH, 1 .6 ^3s. Zelter, Collection 1 M KC1 J 1.23 Czechoslov. Chem. Cammun., 10“3 M HC1 ♦ ii, 319 (1932). 10"2 M KC1 1 .0 8
3 . weakly alkaline 1.23 (N.H.E.) ^J. Heyrovsky, Collection 1 molar alkali 1.00 (N.H.E.) Czechoslov. Chem. Commun., 1, 198 (1935).
h. 0.1 M KC1 1 .1 U0G. Challenger, Ph.D. dissertation, Harvard University, 19^5-
5. NH^GafSCj^ t ^5p. Zunan, Collection 10"3 M HC1 1.08 (N.C.E.) Czechoslov. Chem. Commun.. 1 M KCN I .2 9 12, 1107 (1950).
6 . 2 M HQAc ♦ 1+6M. A. De Sesa, D. N. 2 M NH^QAc no reaction Hume, and D. D. DeFord, 0 .01^: gelatin Anal. Chem.. 25. 983 (1953) Lv 7. 0.5 M KSCN 0.85 'V. A. Tslmmergakl and A. pH 1.7-3.7 stable I. Vovenko, Ukrain. Khim. maximum, Zhurn., 22. 566 (195*0. surfactantb eliminate wave completely TABLE 2 (continued) 12
-El (S.C.E.) at 25°C. or Electrolyte as Indicated Reference
8. 10"3 M NaCl, pH 3-5 1.065 ^®H. Zittle, Diss. Abstr., 10"2 M BaClo 1.05 3J+, 1303 (195*0- 10"2 M ( C H ^ N C l 1 .0 8 10- 2 M KgSOj, 1 .0 6 10-2 M NaClOj, 1 .0 6 Acids at pH 3*50 HC1, HCIO^, HNO,, H 2so4 0.92
9. 0.001 N salicylic N. Vinogradova and acid 0 .8 8 N. N. Chudinova. Zavodskaya pH 2 . 8 - 3 A irreversible Lab.. 2 2 , 1280 (1956).
10. 0.5 M KC1, pH 5-35- 5°J. Szonntagh and J. Roz- 6.35 1.05 manith. Acta Chim. Acad. 0 .005^ gelatin Sci. Hungary, 9, 99 (1956).
11. 0.1 M NaF 1 .1+2 Micka, Collection 5 M NHo ♦ 1 M NH^Cl 1.58 Czechoslov. Chem. Commun.. 0 .1 M KN0 -, 1 .2 1 21, 862 (1956). 0.05 M Na^SO^ 1 .2 1 0.1 M NH^F 1 .3 8 0.2 M NH^F 1 .1+1
12. 0,1 M CyHeOoHa, 52A. I. Zelyankaya and N.V. 0.1 M NaCl 0 .9 8 8 Bausova. Zhur. Fiz. Khim., pH 2.5-1+.5 3 1 ,1+1+0 (195717
13. fused LiCl one step -’3H. A. Laitinen, C. H. Liu, (tungeten Ga(lll)— ► and W. S. Ferguson, Ansi. electrode) Ga(0) Chem.. 30. 1266 (1953).
1*+. 1.1 x 10”2 M Super functional -’Si. A. Cooney and J. H. chrome Garnet Y, group Saylor, Anal. Chim. Acta, pH 5.53 reduction 21, 276 (1959). 0 .3^, 0.53 1.05 55 15. 0.32 M (NH1,)2C20i| 1 .1 8 V. K. Kuznetsova, Zhur. pH 7 A Neorg. Khim., *+, 1+6 (1959).
16. 7.5 M KSCN 0 .9 0 0 M. MacNevin and E. D. reversible Moorhead, J. Am. Chem. Soc., 8 1, 6382 (1959). 13
Solubility of Metals in Gallium
Gallium metal has an unusual liquid range. The metal melts at
29.75°C. and boils at 2000° centigrade. This long liquid range made gallium attractive as a coolant for nuclear reactorshowever,
57R. I. Jaffee, R. M. Evans, E. A. Fromm and B. W. Gonser, "Gallium in Nuclear Reactors," AECD-3317, I9U9. difficulties appeared early in the study of gallium as a reactor coolant. The neutron cross section of gallium is rather high, 2.2 barns per atom. This means that the coolant, gallium, would be absorbing neutrons necessary for the maintenance of a chain reaction.
In addition, at temperatures in the range of 600 to 1000°C., gallium dissolved most prospective container materials. This unusual property has been summarized in a statement In the "Liquid Metals Handbook. "9
Gallium is more aggressive in its attack on most metals at a given temperature than any other molten metal (that has been studied). It can be contained at high tem peratures only in some of the refractory oxides, quartz, graphite, and such metals as tungsten and tantalum.
The interest in gallium as a reactor coolant decreased when the properties of high neutron croBB section and the highly corrosive nature of gallium could not be overcome.
A bibliography of gallium alloys is contained in the work of de la Breteque.5*6,7 STATEMENT CP PROBLEM
It was evident from a survey of the literature that the development of the dropping gallinn electrode had not progressed to a level that would permits its use as a reliable tool for the investigation of reactions of gallium.
From the electrochemical point of view the dropping gallium electrode presents a new metal surface for investigation. The dropping gallium electrode has the advantage of a renewed surface with each drop, thus reducing the passivization difficulties common with gallium pool electrodes.
Since the chemical literature did not contain evidence that a dropping gallium electrode had been used to investigate the polarization behavior of gallium, the problem became resolved into the following principal parts:
1. The development of a dropping gallium electrode
2. A study of the polarization behavior of gallium using
a dropping gallium electrode
11+ CHAPTER I
EXPERIMENTAL raOCEDURE
Dropping Gallium Electrode
The first reported use of a D.G.E. was the work of Challenger**®
in which he used a squeeze bulb to force the gallium through the capillary. The electrode was used in potentiometric studies of the
standard potential, of the gallium electrode.
A D.G.E. has been reported by Giguere and Lamontagne.^® The
A. Giguere and D. Lamorrtagne, Science. 120. 390 (195*0 • gallium was forced through the capillary by air under pressure. These workers report that the only successful curve obtained was for gallium falling into air saturated potassium chloride. A cathodic curve was
obtained in this solution with a half-wave potential of -1 .2 6 volts
(S.C.E. at 35°C»)* This wave did not appear in the air-free solu tions. These authors, in conjunction with consultation with J. J.
Lingane, concluded that this cathodic curve was the result of the
reduction of gallic ion formed at the surface of the drop by the oxygen present in the solution. The anodic curve reported was continuous from -0 .6 to -1.0 volts. These authors were not able to obtain reproducible results with their apparatus.
15 16
A D.G.E. has been used by Grahame to study the differential capacity of the gallium solution Interface as compared with that of m e r c u r y . 59 No Information is given about the operation of the
59D. C. Grahame, Anal. Chem.. ^[0, 1736 (1958). electrode other than the gallium was forced through the capillary by a pressure of nitrogen.
At the outset of the development of a D.G.E. in this laboratory it was believed from statements made by various authors^* ^ that
60E. D. Moorhead, Ri.D. dissertation, The Ohio State University, 1959.
^ G . H. Wagner and W. H. Gitzen, J. Chem. Ed.. 29. 162 (1952). liquid gallium vets glass. This belief in the wetability of glass by gallium led to the attempted construction of plastic reservoirs and plastic capillaries. The plastics investigated were polyethylene,
Teflon, and Penton. Attempts to create the bore of the capillary around a 38 B and S gauge Nichrome wire were not successful.
Subsequently, an all-glass apparatus was coated with Desicote to retard the adherence of gallium to the surfaces of the container and capillary. During the work it was noted, both experimentally and In the literature, that pure oxide-free gallium does not vet
0<=S. Boyer, U. S. Patent 1,576,219, March 9, 1926. glass. Consequently, an all-glass apparatus could be used if oxygen was prevented from coming in contact vith pure oxide-free gallium. 17
The Dropping Gallium Electrode Apparatus
The glass apparatus, Figure 1, for the D.G.E. consisted of three parts: (l) reservoir, (2) capillary, and (3) pressure apparatus. A pair of 1^/35 standard taper glasB Joints was used for the reservoir.
The inner Joint (A) had a one-inch length of one millimeter I.D. by one-fourth inch O.D. capillary tubing sealed to the end opposite to the standard taper. This length of capillary tubing had a 7/25 standard taper outer Joint sealed to it. The outer Joint (B) had a tungsten wire contact sealed into a length of six millimeter glass tubing with epoxy resin cement. A second length of tubing had a two-way stopcock to which were attached the nitrogen pressure line and the vacuum line (H). A -water Jacket (C) was held in place by rubber stoppers.
The capillary used had an outside diameter of about one-fourth inch and an inBide diameter of seventy to eighty microns, 0.070-0.080 millimeters. The lumen of the capillaries was measured with a microscope having a reticule marked in 0 .0 1 millimeter divisions.
All capillaries were four centimeters in length. Epoxy resin cement was used to seal the capillary into a 7 /2 5 standard taper outer
Joint (D).
Oxygen was removed from dry pumped high purity nitrogen (E) by passage over copper gauze (F) heated to ^50° centigrade. The pressure applied to the D.G.E. was controlled by the diaphragm valve on the nitrogen tank and needle valves (g ) on either side of the pressure gauge (H). The pressure gauge operated from zero to thirty pounds FIGURE 1
DROPPING GALLIUM ELECTRODE APPARATUS
A. Gallium reservoir - inner 1^/35 standard taper Joint
B. Gallium reservoir - outer 1^/35 standard taper Joint
C . Water Jacket
D. Capillary in an outer 7/25 standard taper Joint
E. Nitrogen tank
F . Copper gauze heated to *+50°C.
G. Needle valves
H. Pressure gauge
K. Surge tank
M. Connection to mercury trap to atmosphere
N. Connection to vacuum
18 M Mercury ^ Trap Atmosphere
->. N Vacuum
Tf ,," k Surge \ Tungsten Tonk 1 Wire \IOOO m . H
k B I r S - I Needle I Valves
A
CD O h CL O F o
° X Sv ci
Capillary 20 per square inch. Pressures normally used were in the range of fifteen pounds per square inch. A surge tank (K) of about one liter capacity was placed between the nitrogen tank and D.G.E. to help maintain a constant applied pressure. An escape to the atmosphere was provided through a mercury trap at (M). This trap prevented air being drawn into the D.G.E. as the nitrogen pressure was reduced.
The apparatus was checked for leaks by pressurizing the system to 15 p.s.i. and then closing the needle valve on the nitrogen tank and the stopcock on the dropping gallium electrode. In the event that the apparatus did not maintain the applied pressure (15 p.s.i.), a thick soap solution was applied externally to the suspected areas.
Soap bubbles were produced over the site of a leak. When the pressure inside the apparatus remained constant for about one hour, the apparatus was considered to be free of leaks.
Assembly, Operation, and Cleaning of Dropping Gallium Electrode
The D.G.E. was prepared for assembly by applying Kronig Cement to the warmed standard taper joints on the lower, reservoir, section of the assembly. The outer standard taper joint containing the capillary was warmed and the Joint was sealed on to the reservoir. A length of
Tygon tubing from a nitrogen cylinder was slipped over the capillary, and nitrogen was forced through the capillary. While the reservoir was being swept with nitrogen, a quantity of frozen gallium, eighteen grams, was melted under twenty milliliters of acetone containing from two to four drops of concentrated hydrochloric acid. A separatory 2 1 funnel fitted with a Teflon stopcock was used to transfer the galllum- acetone mixture to the reservoir. The upper section of the assembly, with nitrogen and vacuum tubing attached, was warmed and cemented to the reservoir section. Nitrogen was flowing up througi the capillary during these operations.
After the apparatus had been assembled, the gallium was frozen with the apparatus in a horizontal position and the acetone-hydro chloric acid mixture drawn off by vacuum. The stopcock to the nitrogen pressure supply was closed during the evacuation and nitrogen was flowing continuously through the capillary. The evacuation was
continued for fifteen minutes, after which the stopcock to the vacuum line was closed and nitrogen passed down into the reservoir and out through the capillary with the Tygon tubing removed. Finally, the water jacket was fitted over the reservoir. The Jacket was used to
circulate warm water to melt the gallium prior to its use as a dropping electrode.
During the evacuation of the gallium reservoir, the polarographic
cell was mounted in the water bath and the test solution added to the cell. Nitrogen was bubbled through the twenty-five milliliter test
solution for thirty minutes and through the ten milliliter test
solution for ten minutes. This time period also permitted the temperature of the test solution to equilibrate to the temperature of the water bath. The gallium in the reservoir waB melted by circu lating water from the water bath through the water Jacket. The gallium reservoir wa s held in a horizontal position during the melting of the gallium and nitrogen was flowing out through the capillary. 22
When the gallium had heen melted and the test solution was outgassed, the dropping reservoir vas inverted and the capillary inserted under the surface of the test solution. All vater circulating pumps were turned off, electrical connection vas made to the Electrochemograph, and the nitrogen pressure increased to force gallium through the capillary into the test solution.
In order to change test solutions, provision had to be made to either retard or stop the flow of gallium. An attempt vas made to freeze the gallium in the capillary. This vas not successful since gallium expands on freezing and therefore cracked the capillary. The alternate method chosen vas to reduce the nitrogen pressure above the gallium with the capillary tip submerged below the surface of distilled vater in a small flask. The nitrogen pressure vas reduced
slovly to prevent sudden decompression of the D.G.E. Sudden decom pression of the D.G.E. draws the liquid metal up the capillary along with the electrolyte, frequently plugging the capillary. The combined
effect of the expansion of the compressed liquid gallium^1 and the
equilibration of pressures forces the gallium up through the capillary*
All polarographie curves were recorded on a Leeds and Northrup
"Electrochemograph Type-E." Drop times were taken manually by averaging the time required for ten drops to fall from the capillary.
Values for the weight of a drop of gallium were determined by weigh
ing 1 0 drops of gfallium collected in a small glass cup in the test
solution. The gallium was drawn from the cup into a test tube where
it was frozen, washed with water and acetone, dried, and weighed. 23
The dropping electrode assembly vas cleaned after each use. All wax used in sealing ground glass Joints vas removed by melting the wax, soaking the warm glass In a mixture of benzene and toluene, and rinsing with acetone. The reservoir section of the assembly vas soaked In alcoholic potassium hydroxide, rinsed with vater and acetone, and dried in an oven. Warm dilute hydrochloric acid, vater, and acetone were drawn up through the capillary to clean and dry it.
Electrolysis Cells
Two different electrolysis cells were used; one is illustrated in Figure 2. This cell vas used during the work with the D.G.E. in solutions of 7 .5 M potassium ttiocyanate and 2 M sodium perchlorate.
The salt bridge (C) vas a potassium chloride bridge during the work with potassium thiocyanate. The investigation of sodium perchlorate solutions required that the salt bridge (C) be changed to a sodium perchlorate-sodium chloride bridge to prevent the precipitation of potassium perchlorate at the solution-salt bridge interface. The resistance of the cell vas measured with a conductivity bridge and found to be about 300 ohms. The other cell, a polarographic H-cell with a potassium chloride salt bridge, had a resistance of about 200 ohms. These resistances were of such magnitude that their influence on the potential of the cell vas negligible. FIGURE 2
ELECTROLYSIS CELL
A. 30 milliliter sintered glass filter funnel
B. Glass cup to catch gallium drops
C. Salt bridge
D. Receiver for gallium drops for the determination of m
E. Salt bridge to S.C.E.
F, Reservoir for preliminary outgassing of solutions
G. Nitrogen tank
H. Copper gauze heated to ^50°C.
I. Mercury trap to the atmosphere
K. Copper ring and springs - used to hold Lucite cap In place
L. Rubber "0" ring
M. Apertures for capillary and reagents
N. Rubber tubing seal
P. Pinch clamps - used to control the direction of flow of nitrogen
2k 25 26
Reagents
Gallium metal, 99-999? pure, was obtained from the Eagle-PItcher
Company. The gallium nitrate used was a "spectroscopically" pure salt obtained from the Johnson, Mathey and Company, Ltd. The potassium thiocyanate was analytical reagent grade material obtained from both the Malinckrodt and Baker Chemical Companies. A first lot of material from the Malinckrodt Chemical Company did not show the presence of substances that would cause interfering reactions at the
D.G.E. However, subsequent samples of potassium thiocyanate from both companies showed the presence of an interfering substance at the D.G.E. which could not be removed by recrystallization. The
sodium perchlorate was the anhydrous "reagent grade" salt obtained from the G. P. Smith Chemical Company. Solutions of this salt could
not be used until they had been filtered to remove excessive amounts
of foreign matter. The potassium chloride was analytical reagent
grade used without further purification. The standard acid solutions,
1 M perchloric and 1 M hydrochloric acids, were obtained from The
Ohio State University Reagents Laboratory.
Gallium perchlorate 6-hydrate was prepared following the method
of Foster.Gallium nitrate 8-hydrate was boiled In JO?.- perchloric
63l. s. Foster, J. Am. Chem. Soc.. 6l, 3122 (1939)*
acid for one hour. The perchloric acid-gallium solution was yellow
when hot, and white fumes of perchloric acid were visible above the
solution. As the solution cooled, it became colorless and white crystals of gallium perchlorate precipitated from the solution. The crystals vere filtered on a sintered glass filter and washed with
70$ perchloric acid. The crystals were dried in a drying pistol under vacuum for seventy-two hours at 130° centigrade. The temperature was maintained by boiling chlorobenzene. The resulting product was analyzed for gallium by precipitation as the
8-hydro xyquinolate^ and for perchlorate as potassium perchlorate.
TABLE 3
ANALYSIS OF GALLIUM PERCHLORATE 6-HYDRATE
Theoretical Values Experimental Values
Ga cioh~ Ga C10u‘ percent percent percent percent
1 h.6h 6 2 .6 5 IU.3 0 6 2 .2 3
1U.UU 62.39 CHAPTER H
HYDROLYSIS CF GALLIUM
Hydrolysis is the reaction of the dissolved salt with vater to give an excess of either hydrogen or hydroxyl ions. An excess of hydrogen ions is expected with gallium since it is a Levis acid and its chemistry is similar to that of aluminum. Hydrolysis constants^
. Fricke and K. Meyring, Z. anorg u. allgem. Chem.. 3^5 (1928). have been determined for the reactions:
(1) Ga(lll) ♦ H20 • Ga(OH)** ♦ H+, K'h = 4.03 x 10"3
(2) Ga(OH)* 4 t H20 » Ga(0H)2* ♦ H+, K"h ■ 3-16 x 10“ ^
These reactions show that a decrease in pH occurs as hydrolysis progresses.
An examination of the hydrolysis of gallium perchlorate^ by the
^ C . S. Patterson, S. Young Tyree Jr., and Kerro Knox, J. Am. Chem. Soc., J2, 2195 (1955). osmotic pressure method has shown that gallium perchlorate is a very strong highly hydrated electrolyte and that ion pair formation is negligible. Additional evidence vas presented suggesting a polymeric hydroxy-gallium compound. Polymeric hydroxy-gallium aggregates have
23 29 also been observed In gallium solutions as the pH vas Increased by addition of sodium hydroxide. 66
^ P . Moeller and G. L. King, J. Am. Chem. Soc., j k f 1355 (1952).
Experimentally, a solution of 0.21 millimolar gallium perchlorate in vater had a "natural" pH of 2.5 and solutions Of 0.3 to 0.9 milli molar gallium nitrate in 1 M potassium chloride had a pH range of
3.5 to 3-3.
An examination of current-voltage curves at the D.G.E. of gallium nitrate in 1 M potassium chloride at the "natural pH" shoved a hydrogen vave resulting from the hydrolysis of gallium nitrate.
Figure 3 shows the curve produced by the addition of 0.0002 M hydro chloric acid (III), and (il) the curve resulting from the addition of a 0.000065 M gallium nitrate to 1 M potassium chloride. Curve IV is a combined curve of 0.0002 M hydrochloric acid and O.OOOO65 M gallium nitrate in 1 M potassium chloride. These curves show that gallium nitrate is extensively hydrolyzed and that the height of the gallium vave is increased with added hydrochloric acid.
Figure 1+ is an example of the complex curve formed in 0.001 M or greater gallium nitrate in 1 M potassium chloride without added acid
(HCl) to suppress hydrolysis. Curve I vas obtained from a solution of 0.0016 M gallium nitrate. The curve vas obtained within three minutes of the preparation of the solution. The solution vas prepared by dissolving 6.282 milligrams of gallium nitrate 8-hydrate in one hundred milliliters of outgassed 1 M potassium chloride (high purity nitrogen for one hour). Curve II is the same solution twenty-four FIGURE 3
CURRENT-VOLTAGE CURVES OF GALLIUM NITRATE AND HYDROCHLORIC ACID IN 1 M POTASSIUM CHLORIDE
CURVE I . . . . Residual Current Curve 1 M Potassium Chloride
CURVE II . . . O.OOOO65 M Gallium Nitrate
CURVE I D . . . 0.0002 M Hydrochloric Acid
CURVE IV . . . Curve III plus Curve II
30 microamperes ■frrooro-fcaoDOro *o> ■frrooro-fcaoDOro
-E vs. S.C. E. at 35° C FIGURE U
TYPICAL CURRENT-VOLTAGE CURVES ILLUSTRATING THE COMPLEXITY PRODUCED AS THE RESULT CF HYDROLYSIS OF GALLIUM NITRATE IN 1 M POTASSIUM CHLORIDE
CURVE I . . . . 0.0016 M Gallium Nitrate After 3 Minutes
CURVE II . . . 0.0016 M Gallium Nitrate After 2k Hours
CURVE II . . . 1 M Potassium Chloride
32 u> U) + microamperes M © © U» o i
m vs. S.C.E. Qt 35° C 3^ hours later. A comparison of the two curves shows that they are essentially the same shape and height, thus confirming that hydrolysis apd galliin complex formation Is rapid and complete within a few minutes.
An unusual "spike" is seen at -0.95 volts on the rapidly rising gallium curve. Since this "spike" had a life of only two or three drop times, more detailed observations of its nature could not be made with the automatic recording polarograph. The "spike" did not appear when current ranges greater than 30 microamperes were used. This was probably due to the averaging influence placed upon the current measuring system at larger current ranges. However, the "spike" accompanied by a 30 millivolt plateau reappeared in solutions of
0 .0 0 3 5 millimolar gallium nitrate in which the ratio of added hydrogen ion (HCl) to added gallium ion vas 5 (Figure 5). This "spike" could be the result of the reduction of solvated Ga(lll) and the following current the result of the reduction of hydrated and complexed gallium species until the potential (-1 .1 8 volts) is reached where hydrogen ion produced as the result of hydrolysis is reduced.
The addition of a polyether surfactant (0.002^ Triton X-100) to a
0 .0 0 5 6 ml111 molar solution of gallium nitrate in 1 M potassium chloride,
Figure 6, produced Curve II. The solution had attained its "natural pH" by hydrolysis. Gallium is being reduced at the beginning (-0-9 volts) section of the curve and begins to show a limiting current until hydrogen Ion begins to reduce at about -1.18 volts. The condi tions at the drop surface at this point are those of a rapidly rising pH and the subsequent formation of a nonreducible hydroxy-gallium FIGURE 5
AN EXAMPLE CF THE COMPLICATED CURRENT-VQLTAGE CURVE OBTAINED WHEN THE CONCENTRATION CF ADDED HYDROCHLORIC ACID WAS INSUFFICIENT TO SUPPRESS HYDROLYSIS CF GALLIUM NITRATE IN 1 M POTASSIUM CHLCRIDE
CURVE I . . . . 0.0035 M Gallium Nitrate, [ V ] /EGa3*] « 5
35 + + microamperes
vs. S.C.E. ot 35°C FIGURE 6
CURRENT-VOLTAGE CURVES ILLUSTRATING THE HYDROLYSIS CF GALLIUM NITRATE IN 1 M POTASSIUM CHLORIDE
CURVE I . . . . I M Potassium Chloride
CURVE XI . . . O.OOO56 M Gallium Nitrate plus 0.002£ Triton X-lOO added to Curve I
CURVE III . . . Curve II plus 0.020 M Hydrochloric Acid
37 OoSC 3 D S -M microamperes + 39 compound which blocks the electrode and causes the current to fall.
Once the applied potential has became sufficiently negative, a cathodic curve is again observed for the reduction of hydrogen ion.
Additions of hydrochloric acid to this solution increased the overall height of the curve until a sufficiently large concentration of hydro chloric acid had been added so that a sharp maximum appeared at -1 .0 volts. This cathodic maximum was later used as the diffusion current value to determine the *-d/C relationship for the reduction of gallium at the dropping gallium electrode.
Since hydrolysis of Ga(lll) produces a complicating interference
in the interpretation of the gallium curve, hydrochloric acid was added to a 0 .0 0 0 1 3 M gallium solution to suppress hydrolysis.
Hydrochloric acid was added in a stepwise manner such that the ratio
[h+] / increased in the order, 0, 10, 20, 30 . . . 80. The observed cathodic maximum in the vicinity of -1 .0 volts rose until the hydrogen ion-gallium ratio became 50. This ratio is the same as that UA observed by Zittel for the reduction of gallium at the dropping mercury electrode. Acidified gallium solutions provided an additional advantage over solutions of "natural pH" because the multi-curve appearing in solutions of "natural pH" was reduced to a two-step
curve in the acidified solution.
The cathodic currents observed in acidified gallium solutions with concentrations greater than 0 .O0 U M were so large that they could not be recorded on the 100 microampere range of the recorder.
However, the addition of 0.002^ Triton X-100 reduced the magnitude of 4o the current to a value that could be recorded for solutions whose concentrations were as large as 0 .0 0 8 molar.
An approximate hydrolysis constant may be calculated from
Figure 3 . The equilibrium concentrations of Ga(lll) and H* are
0.00000^6 M and 0.00015 M respectively. The initial concentration of gallium nitrate added was O.OOOOO65 molar. These values show that gallium nitrate was 8 7.7f- hydrolyzed in this solution. If the assumption is made that the major product of hydrolysis is Ga(GH)*+, then the following reaction can be used to calculate an approximate hydrolysis constant.
(3) G a ( m ) - H20 ■ Ga(CK)*'' ♦ H*
The value for a hydrolysis constant can be calculated from the following equation.
00 K„ = Men)"] [h *1 ~ jH+P [Ga(lII)] [Ga(III)]
Substitution of the experimentally determined values for H* and
Ga(lll) in equation 3 gives a value for a hydrolysis constant of
Jr.7 x lCf^. This is in good agreement with the value ^.03 x 10“3 determined by Fricke and Meyring^ for K*^. CHAPTER III
OXIDATION OF GALLIUM
Electrochemical Oxidation of the Dropping Gallium Electrode
Anodic polarization studies in 1 M sodium chloride, sodium sulfate, and sodium perchlorate were carried out "by Schwabe^ on solid
6?K. Schvabe, Z. -phys. Chem., 211, 170 (1959). and liquid gallium electrodes. These studies shoved than an anion dependence for anodic polarization on gallium was in the order
SO^" >010^" >C1~. This series showed that chloride exhibits little polarization influence on the anodic dissolution of gallium. Per chlorate ion showed a considerable polarization effect until oxygen evolution occurred at -1 .2 volts on solid gallium, but showed only a minimal polarization effect on liquid gallium. Sulfate showed a large polarization effect on solid and liquid gallium. Schvnbe concluded that anodic dissolution through protective layers is inhibited only when the layer is firmly attached to metal surface.
Anodic polarization studies by Stelling68 in acid solutions
Stelling, Z . Electrochem., ^1, 779 (1935). showed a passive region, d (current density) ^ 0 , in the anodic d E. hi k2 curve at -0.75 to -O.6 5 volts (0.1 M calomel electrode). The pool electrode at this potential had a rough light gray surface. An additional passive region was observed at 0.0 volts. No explanation vas given for these two passive regions except that gallium was going
Into solution at 0.0 volts. Chlorine evolution was not observed at any potential with the liquid gallium pool electrode. 6q Davidson and Jirik ^ studied the anodic dissolution of gallium
69a. W. Davidson and F. Jirik, J. Am. Chem. Soc., 12., 1700 (1950).
In ammonium acetate and sodium acetate dissolved In anhydrous acetic acid. No chemical oxidation vas observed in this electrolyte. The anodic oxidation of gallium dissolved more metal from the anode than corresponded to Faraday's lav if the product vas assumed to be a trlvalent cation. These authors explained this result in terms of the hypothesis that the primary anode product vas a mixture of mono- and trivalent cations.
Electrochemical oxidation of gallium using the D.G.E. vas observed in potassium thiocyanate, sodium perchlorate and potassium chloride. Anodic current-voltage curves obtained in these electro lytes indicate solution of gallium metal at dissolution potentials,
Ed, of -1.2 volts (7.5 M KSCN), -1.15 volts (2 M NaClO^), and -1.05 volts (l M KDl). The Ed value in 0.01 M sodium hydroxide vas -1.5 volts.
The slopes of the anodic current-voltage curves at the D.G.E. in
7.5 M potassium thiocyanate, 2 M sodium perchlorate, and 1 M potassium chloride are the same and constant from -0.90 to -O.8 5 volts, Figure 7. FIGURE 7
ANODIC DISSOLUTION CURVES OF GALLIUM IN 1 MOLAR POTASSIUM CHLORIDE
CURVE I . . . . 1 M Potassium Chloride
CURVE II . . . 1 M Potassium Chloride plus O.Ol M Hydrochloric Acid
CURVE III . . . 1 M Potassium Chloride plus 0.1 M Hydrochloric Acid
CURVE IV . . . Curve III plus 0.002^ Triton X-100
1*3 microamperes 30 20 0 6 0 4 0 5 70 . 06 . 08 . 10 l.l 1.0 0.9 0.8 0.7 0.6 0.5 0.4 0.5 only vs.— E E S.C.E. forNo.T C ot35 0.6 - E vs. S.C.E. E C at 35 - 0.7 TZ 0.8 0.9 nr
1.0 0 4 0 3 0 2 -- - 0 5 ■ - ■ 40 4 -■ 1.2 l.l 20 30 0 4 50 60 70 20 10 30 50 0 6
microamperes M* At potentials more positive than -O.85 volts the slopes are the same but smaller in magnitude than the slopes at potentials more negative
than -O.8 5 volts. Table k lists the slopes obtained in the three
electrolytes over the potential ranges mentioned. The decrease in
TABLE h
RECIPROCAL SLOPES OF ANODIC CURRENT-VOLTAGE CURVES
Electrolyte Potential Range Slope”^ (volts/microampere)
7.5 M KSCN -0 .9 0 to -O.85 k x 103
2 M NaClO^ 4.5 x 1 0 3
1 M KC1 6 x 1 0 3
7.5 M KSCN -O.85 to -0 .6 0 1 .7 x 10^
2 M NaClO^ 1 .7 x 10k
1 M KC1 1.3 x IQ1*
slope suggests an increase in resistance at the drop solution inter
face in the three electrolytes investigated. An increase in
resistance is an indication of an increase in passivity of the gallium
drop, perhaps through the formation of an oxide film.
Current-voltage curves obtained with the D.G.E. in 1 M potassium
chloride show a change of shape upon addition of hydrochloric acid.
An increase in hydrogen ion concentration shifts the Ed values to
more positive potentials. This shift in Figure 7 has been attributed
to the oxidation of the gallium drop by hydrogen ion and the
subsequent reduction at the D.G.E. of the Ga(lll) formed. k6
The anodically polarized D.G.E. exhibits two anodic limiting currents at -0.86 volts and -0.75 volts In 1 M potassium chloride containing 0.01 and 0.1 M hydrochloric acid (Figure 7)* An anodic
"spike" Is observed at -0.86 volts in these acid solutions. Curve IV shows that the "spike" can be removed by the presence of 0 .002$
Triton X-100. These anodic limiting currents possibly result from the rate of removal of the adherent oxide film by hydrogen ion. The occurrence of two anodic limiting currents can be attributed to the presence of different crystalline forms of gallium oxide^® or
7°L. M. Foster and H. C. Stumpf, J. Am. Chera. Soc., 73» 1590 (1951).
/■c 66 insoluble hydrous oxide polymers 5 * with different acid solu bilities.
Two passivization regions were observed with the D.G.E. as com pared with one by Stelling.^® This can be attributed to the constant renewal of the gallium surface at the D.G.E. as compared with a quiet pool electrode. This study and the work of Stelllng^® Indicates passivization at -0.8 volts, whereas the work of Giguere^® does not indicate any anomaly in the anodic curve at this potential.
Chemical Oxidation of the Dropping Gallium Electrode
Preliminary experiments with the D.G.E. were made by using the cell in Figure 2 with a saturated potassium nitrate salt bridge and
7.5 M potassium thiocyanate as the electrolyte. The residual current curve produced under these conditions, Figure 8, was only about FIGURE 8
RESIDUAL CURRENT CURVES AT THE DROPPING GALLIUM ELECTRODE IN 7.5 M POTASSIUM THIOCYANATE
CURVE I . . . . Potassium Chloride Salt Bridge
CURVE II Potassium Nitrate Salt Bridge no + microamperes o O bo no OD 9n vs. S.C.E. at 35 C 0 .2 volts between the anodic dissolution of gallium at -1 .1 volts and a final current rise at -1.35 to -I.1*? volts. This result would
severely limit the investigation of reactions at a cathodlcally polarized dropping gallium electrode. Suspecting an interference from the potassium nitrate salt bridge, the bridge was changed to a
saturated potassium chloride salt bridge. The residual current curve produced under these conditions showed a considerable slope of 18 microamperes per volt, but did not show a final current rise until
about -1.7 volts. This result indicated that nitrate ion was causing
a catholic current in the vicinity of -1 .3 to - l A volts.
In order to examine further the effect of nitrate ion at the
D.G.E., current-voltage curves were obtained for a solution of
0.001 M gallium nitrate in 7*5 M potassium thiocyanate (Figure 9t
Curve III). Following the curve for the reduction of gallium, a small maximum was observed at -1 .3^ volts which in turn was followed by a
third curve in the vicinity of -1.5 volts. This maximum occurred at a
potential assumed to be that of nitrate ion interference. However,
the multi-curve observed could have been due to the reduction of
complexed species of gallium or to impurities in the gallium nitrate.
Therefore, a 0.001 M solution of gallium perchlorate in 7*5 M
potassium thiocyanate was examined (Curve II). This curve did not
show a maximum at -1.35 volts. The addition of potassium nitrate,
0.03 moles in 25 milliliters, to this solution caused the formation
of a high broad maximum at -1.35 to -1.5 volts (Curve IV). FIGURE 9
THE INFLUENCE OF NITRATE ION AT THE DROPPING GALLIUM ELECTRODE IN 7.5 M POTASSIUM THIOCYANATE
CURVE I . . * . 7-5 M Potassium Thiocyanate
CURVE II . . . 0.001 M Gallium Perchlorate
CURVE Ill . . . 0 .001 M Gallium Nitrate
CURVE IV . . . O.OBO M Potassium Nitrate added to Curve II
50 + microamperes
o ■-j
o b
o <0
b
m in< C/> n m
n
ro vn O H 52
Additional proof of a cathodic curve at the D.G.E. in the presence of nitrate ion was obtained by adding nitrate ion directly to 7.5 M potassium thiocyanate and observing the resulting current- voltage curves (Figure 10, Curves II and III). The stepwise addition of gelatin, 0.0012 , 0 .002l+, 0 .0031, and 0 .005^ percent, did not remove the observed maximum (Figure 11, Curves I, II, III, IV).
These results obtained with potassium nitrate indicated either that nitrate ion was being reduced at the D.G.E. or that the nitrate ion was oxidizing the surface of the gallium drop and that gallium oxide or some basic gallium salt formed was being reduced at the
D.G.E. in the vicinity of -1.35 volt6. Free energy calculations indi cate that gallium should be spontaneously oxidized by nitrate ion^-
T V M. Latimer, "Oxidation Potentials," Prentice-Hall Incor porated, New York, 1953* aB shown by equations 5 and 6.
A F ° Kcal
(5) 8Ga 4 3N03- + 9K20 * ^ a 203 4 3NH1++ t 60H" -536
(6) 8Ga 4 3N03" 4 2IH20 » 90a(OH)3 4 NHU+ + 60H- -609
(7) lOGa if- 6Br03“ 4 3H20 = 5Ga203 4 6Br" 4 60H" -9^5
(8) 5Ga 4 3Br03 4 9H20 = 5Ga(OR)3 4 3Br" 4 30H- -682
(9) 2Ga 4 N02“ 4 3R20 = Ga203 + NH1++ 4 20H~ -316
(10) 2Ga +■N02“ 4 6H20 = 2Ga(0H)3 4 NHU+ 4 20H" -11*5
These equations predict that gallium metal should be oxidized spontaneously by the above oxidizing anions. This is in agreement with observations made of the gallium drops in the presence of these FIGURE 10
THE INFLUENCE OF NITRATE ION AT THE DROPPING GALLIUM ELECTRODE IN 7-5 M POTASSIUM THIOCYANATE
CURVE I . . . * 7.5 M
CURVE II . . . 0.008
CURVE Ill . . . o.oio
53 + microamperes 0 9 0 4 0 6 0 8 0 3 20 0 5 0 7 1.2 1.3 s S. E. t C ° 5 3 at . .E .C S vs. E - 1.4 1.5 1.6 1.7 1.8 5^ 1.9 FIGURE U
THE EFFECT OF GELATIN ON THE NITRATE MAXIMUM IN 7.5 M POTASSIUM THIOCYANATE
CURVE I . . . . 0 .0 1 Q M Potassium Nitrate plus 0 .0012*/;Gelatin
CURVE II . . . 0 .0 1 8 M Potassium Nitrate plus 0 .002L?. Gelatin
CURVE Ill . . . 0 .0 1 8 M Potassium Nitrate plus 0.0031# Gelatin
CURVE IV . . . 0 .0 1 8 M Potassium Nitrate plus 0.005!^- Gelatin
55 VO iA 57
anions where the drops generally remained as individual entities, and
the accumulated gallium on the "bottom of the cell was coated vith a thin film that could have been gallium oxide. In the absence of these
anions, the drops retained their brilliant luster and agglomerated
into a single drop on the bottom of the cell.
The cathodic curve produced in the presence of nitrate ion vas
observed in 7-5 M potassium thiocyanate, 2 M sodium perchlorate and
1 M potassium chloride. In each electrolyte, nitrate ion produced a
curve that had the shape of a polarographic maximum. The position of
the peak of the "maximum" was shifted to more negative potentials as
the nitrate ion concentration vas increased. In addition, the area
under the "maximum" increased as the nitrate ion concentration vas
increased. Gelatin added to solutions of the above three electro
lytes, containing added nitrate ion, failed to remove the "maximum,"
(Figure 11).
Figure 12 is a composite of the current-voltage curves resulting
from the addition of 0.00045 M sodium bromate and 0,005$ gelatin (il),
0.016 M potassium nitrate and 0.001$ gelatin (ill), and 0.0018 M
sodium nitrite and 0 .00d$ gelatin (IV). Th:.s composite shows that the
three oxidative anions produce cathodic currents at nearly the same
potential. The estimated values for the respective half-wave poten
tials are: bromate -1.30 volts, nitrate -1.25 volts, nitrite -1.35
volts. The estimated half-wave potentials for the reduction of the
oxidative anions at the D.M.E. are: bromate in 0.1 M potassium FIGURE 12
CURRENT-VOLTAGE CURVES OBTAINED IN 2 M SODIUM PERCHLORATE WITH ADDED SODIUM BROMATE, POTASSIUM NITRATE, SODIUM NITRITE
CURVE I . • * • 2 M Sodium Perchlorate
CURVE II . . 0.000^5 M Sodium Bromate plus 0.005$ Gelatin
CURVE III . . 0.016 M Potassium Nitrate plus 0.001$ Gelatin
CURVE IV . . 0.0018 M Sodium Nitrite plus 0 .0CU$ Gelatin
58 -(-microamperes 0.9 - - O vs. S.C.E. E - 35°C ot 12 59 60 chloride -0.3 volts,^ nitrate in 0.1 M lanthanum chloride -1.5 volts,nitrite in 0.1 M lanthanum chloride, -I.25 for the first
72E. f. Orlemann and I. M. Kolthoff, J. Am. Chem. Soc.t 6h, 10U2 (19^2 ).
T3S. I. Zhdasov, J. Anal. Chem. U.S.S.R.. 12, 103 (1957). vave and -1.55 for the second wave.A comparison of the potentials observed at the D.G.E. and those observed at the D.M.E. leads to the conclusion that a reduction of the anion is not the potential controlling reaction, but the reduction of some compound formed because of the presence of these oxidative anions. The compound formed could be an oxide or basic salt or gallium formed on the surface of the gallium drop. The observed current is the result of the electrochemical reduction of this film. The sudden decrease in current in the presence of nitrate ion results from the formation of a film of such thickness that the drop is insulated from contact with the solution, thus preventing the passage of current.
A comparison of the results of this study of oxidizing anions at the D.G.E. and the results of Giguere^® of a D.G.E. in air saturated potassium chloride shows similar curves at about -1-3 volts. It could be, then, that a reduction of an oxide film at this potential is what is taking place.
An additional example of chemical oxidation of the D.G.E, was observed in 0.001 M solutions of hydrochloric acid in 1 M potassium chloride. Current-voltage curves of 0.001 M solutions of hydrochloric acid were observed in 1 M potassium chloride. The D.G.E. was inserted 6 l
Into the test solution after the solution had been thoroughly out- gassed. A polarographic curve vas begun for each solution within ten seconds of insertion of the D.G.E. into the test solution (Figure
16). An increase in hydrogen ion concentrations shifts the zero current potential toward more positive values and produces a cathodic current at the potential where Ga(lH) is reduced (area A). The height of this current increases with increasing hydrogen ion concentration. This information suggests that hydrogen ion oxidized the gallium drop to Ga(lll) which is subsequently reduced at the electrode. These observations are in agreement with the position of the standard potential of gallium relative to the normal hydrogen electrode, e.g., hydrogen should spontaneously oxidize gallium metal. CHAPTER IV
CATHODIC CURRENT-VOLTAGE CURVES AT THE DROPPING GALLIUM ELECTRODE
Reduction of Gallium Nitrate
The cathodic current-voltage curve obtained for the reduction of
Ga(lll) in 7*5 M potassium thiocyanate, Figure 9, Curve III, shows a curve for the reduction of Ga(lll) and a curve for the presence of nitrate ion. The combined anodic-cathodic curve for Ga(lll) shows a deflection in crossing the zero current axis. Further studies in new lots of 7•5 M potassium thiocyanate gave residual current curves that had maxima at the potential of the nitrate maximum (Figure 13, Curve
I). Curves II and III are the result of adding 0.01 M and 0.02 M potassium nitrate to the thiocyanate solution. Consequently, a change to a 1 M potassium chloride supporting electrolyte was made to study the reduction of Ga(lll) at the dropping gallium electrode.
Figure 1^ is an example of the combined anodic-cathodic current curve of Ga(lll) at the D.G.E. In acidified 1 M potassium chloride.
Curve II was obtained In a solution of 0.0065 M hydrochloric acid in
1 M potassium chloride. Curve III was obtained in a solution of
0.000126 M gallium nitrate plus O.OO65 M hydrochloric acid in 1 M potassium chloride. The combined anodic-cathodic curve under these conditions does not show an Inflection in crossing the zero current
62 FIGURE 13
AM EXAMPLE CF THE UfTEHFERING CATHODIC CURREOT OBSERVED IN 7-5 M POTASSIUM THIOCYANATE
CURVE I . . * . 7.5 M Potassium Thiocyanate
CURVE II . . . Curve I plus 0.01 M Potassium Nitrate
CURVE Ill . . . Curve I plus 0.02 M Potassium Nitrate
63 + microamperes - - ro hO yi o 01 o
OD I - J I I L FIGURE Ik
C O M B I N E D ANODIC-CATHODIC CURVE FOR 0.000126 M GALLIUM NITRATE AND 0.006^ M HYDROCHLORIC ACID I N 1 M POTASSIUM CHLORIDE
CURVE I . . . . 1 M Potassium Chloride
CURVE II . . . 0.006^’ M Hydrochloric Acid
CURVE III . . . 0 . 0 0 0 1 2 6 M Gallium Nitrate plus O.OO65 M Hydrochloric Acid
65 + microamperes -25 -20 O I - -15 -5 20 25 0.8 0.9 1.0 1.1 vs. S.C.E. E - 35°C of 1.2 1.3 1.4 1.5 66 1.7 67 axis. The value of the zero current potential is a function of the concentration of gallium added as well as a function of the concen tration of the hydrochloric acid added.
The observed maximum in Curve III (Figure 1*+) at -1.0 volts increased with an increasing hydrogen ion concentration. The maximum increased in value until the ratio of to [oa3 tl became 50. The data for the current-concentration dependence of gallium measured at this maximum is given in Table 5. Figure 15, Curve II, shows that
TABLE 5
CURRENT-CONCENTRATION DEPENDENCE FOR GALLIUM NITRATE AT THE D.G.E. IN ACIDIFIED 1 M POTASSIUM CHLORIDE
- Curve I Curve II
[ca3+J x 10^ [h +] x id (p) [h*J x 104 id 0.12 6 .0 3.2 0.24 1 .2 4.3 1 2 .0 5-9 o.6o 3.0 10.3 30.0 15.0 0.85 4.2 13.4 42.0 21.0 1 .2 6 6.5 17.3 65.O 31.5 1.50 7.5 19-5 '’5.0 35-5 1.89 9.5 22.0 95-0 44.4 2.15 10.8 24.0 108.0 50.0 2.40 12.0 27.0 120.0 55-5 FIGURE 15 CURRENT-COICENTRATION DEPENDENCE FOR GALLIUM NITRATE IN ACIDIFIED 1 M POTASSIUM CHLORIDE WITH 0 . 0 0 2 ^ TRITON X-100 CURVE I . . . . Hydrogen to Gallium Ratio is 5 CURVE II . . . Hydrogen to Gallium Ratio is 50 68 i4 microamperes O^CDromO-frODroaiO-^tDro^ O O * o CD K> o> M IN) o rv> * 70 id (measured at the maximum at -1.0 volts) versus C is linear over the concentration range 0.00001 to 0.0002 M gallium nitrate in 1 M potassium chloride plus 0.002^ Triton X-100 where the [h ^] to [pa^J ratio is 50. The minimum at -1.1 volts did not give a linear relation ship between id and C. Figure 15, Curve I shows the results of the id versus C relationship when the [h +J to jca^+3 ratio is 5. This curve shows the influence of hydrolysis &b the concentration of the gallium ion is increased and the hydrogen ion concentration is not added in an amount sufficient to effectively repress hydrolysis. The value of the half-wave potential varied from -0.91 to -0.95 volts over the concentration range of 0 .00002*1 to 0.0002*t M gallium nitrate in 1 M potassium chloride in which the ratio of {jI*J to was 50. The shift of Ei and the shift of the zero current potential suggest that the reduction of Ga(lll) at the D.G.E. under these conditions is not reversible. The measurement of m, mass of gallium flowing per second, was not determined because the reaction at slightly more negative potentials than the peak of the maximum was obviously causing an interfering reaction at the drop surface. An evaluation of n, from log versus E, requires that the electrode reaction be reversible^ This does not seem to be the case with the reduction of gallium at the dropping gallium electrode. Microcoulometry was not a practical method for the determination of n. An approximate value of n was obtained by using the Ilkovic equation that contains terms which are the same or (11) id = k n C m ^ t ^ can be assumed to be the same for the reduction of hydrogen and gallium. The constant, k, will be the same for both cathodic pro- V-K Vi cesses. The capillary characteristics, mT3t , can be assumed to be constant for these reactions. A plot of i^ versus C gives a straight line curve for each of these reactions. The theoretical slope of an - 1 - ^3 % *d versus ^ plot is k n D2 m t . The ratio of the slopes of this plot X , X for the two reactions gives a theoretical ratio of D 2H / pGa A value of one was assumea for n for hydrogen and the literature value^ of 9 -3 ^ x 1 0 "^ cm2/sec for was used to calculate the ^ T . M. Kolthoff and J. J. Lingane, "Polarography," Interscience Publishers, Incorporated, New York, 1952 > pp* 50-52. numerator. The denominator was calculated by using three for n and a value of D^a calculated for another trivalent cation, La(lll), from equivalent conductance measurements at infinite dilution. For La(lll), = 69-5 cm^/ohm - equivalents and D° - 6.19 x 10“^ cm2 'sec. Therefore, the calculated theoretical slope ratio is 1.29* The experimentally determined slope ratio obtained under nearly identical experimental conditions of temperature, electrolyte, and capillary characteristics was 1.25. This is in good agreement with the theoretical slope ratio, which indicates that the first cathodic curve observed at the D.G.E. with gallium nitrate in acidified potassium chloride is a three electron reduction of Ga(lll) to gallium metal. 72 The Reduction of Hydrogen Ion Liquid gallivn metal has a relatively low hydrogen overpotential as compared with mercury, 0 .6& volts (0.001 amps/cm^) for gallium^ and 1 .2 volts (0.01 araps/cm^) for mercury.^ A low overpotential for T5S. G. Christov and S. Rajceva, Naturwiss., hj, 127 (lS>6l). *76 J. J. Lingane, "Eleetroanalytical Chemistry," Interscience Publishers, Incorporated, New York, 1958> P* 209- hydrogen on gallium limits the investigations with the dropping gallium electrode to either neutral or "basic solutions. Grahame^^ reported that the reduction of hydrogen ion on gallium is an irreversible process. Cathodic curves for hydrogen ion have been observed at the dropping mercury electrode. These curves were produced as the result of hydrolysis of metal cations, catalytically by platinum and proteins,and by the direct addition of hydrochloric acid to lanthanum chloride.Cathodic curves for hydrogen ion at the D.G.E. were observed in 1 M potassium chloride. In addition to making measurements of the standard electrode potential of gallium,39 Stelling Investigated electrolysis reactions at a gallium pool electrode.^® An examination of acid solutions showed an Increase in current density, milliamperes per square centi meter, when the potential of the pool was -1.1 to -1.2 volts (0.1 M calomel electrode at 35°C.). This voltage is 0.2 volts more negative than the experimentally determined potential of the gallium electrode. The current density showed a limiting value that was dependent on the 73 acidity of the solution. The increase in current density over this potential range was attributed to the reduction of hydrogen. Figure 16 is an example of current-voltage curves at the D.G.E. of 0.0001 to 0.0005 M hydrochloric acid in 1 M potassium chloride with a 0.002$ Triton X-100. Area A in Figure 16 has been attributed to the auto-oxidation of gallium by hydrogen ion and the subsequent reduction of the Ga(lll) produced. The value for the half-wave potential over this concentration range is -1.18 volts. A plot of the data in Table 6, i^ versus the square root of the applied nitrogen pressure, yields a straight line (Figure 17). This indicates that the reduction of hydrogen ion at the D.G.E. is diffusion controlled. TABLE 6 SQUARE ROOT OF APPLIED PRESSURE-CURRENT DEPENDENCE FOR 0.0003 M HYDROCHLORIC ACID IN 1 M POTASSIUM CHLORIDE MEASURED AT -1 .3 VOLTS (S.C.E.) 1 p.s.i. p.s.i. 2 id ()ia) 11.0 3.32 7.5 13.0 3.60 8.2 15.0 3-37 3.9 17.5 U.18 9.^ 18.5 *+ -30 9.7 19.8 k.U-5 9.9 FIGURE 16 CURRENT-VOLTAGE CURVES RESULTING FRCM THE REDUCTION OF HYDROGEN ION IN 1 M POTASSIUM CHLORIDE PLUS 0.002$ TRITON X-100 CURVE I . . « . 1 M Potassium Chloride CURVE II ... 0.0001 M Hydrochloric Acid CURVE Ill . . . 0.0002 M Hydrochloric Acid CURVE IV . . . 0.0004 M Hydrochloric Acid CURVE V . . . . 0 .0 0 0 5 M Hydrochloric Acid 74 PO ro ro + microomperes + H m vs. S.C.E. ot 35*C FIGURE 17 SQUARE ROOT CF APPLIED PRESSURE-CURRENT DEPENDENCE FOR HYDROGEN ION IN 1 M POTASSIUM CHLORIDE MEASURED AT -1.3 VOLTS (S.C.E.) AT 35°C • 7 6 t4 micro omparas 10.0 6.8 8.2 8.4 9.2 8.0 9.0 8.6 9.8 7.0 9.6 9.4 7.2 7.6 7.4 7.8 . 4.0 3.2 t ot f pla Nirgn ast a ( .I.)'/f S (R ra tu s ra P itrogan N Appliad of Root rt o u q S 3.4 3.6 3.8 4.2 4.4 77 The data for the current-concent rat ion dependence for hydrogen at the D.G.E. is given in Table 7* Figure 18, Curve I, was obtained by measuring i^ from the residual current curve. The fact that this plot did not go through the origin, but intersected the current axis at negative values of i^, supports the hypothesis that hydrogen ion was being consumed in a prior chemical reaction. Curve III was obtained by measuring i^ from the height of the gallium curves observed in area A . ^ This curve shows that the diffusion current of 77 Lewis Meites, "Polarographic Techniques," Interscience Publish ers, Incorporated, New York, 1955, p. 2h6. hydrogen corrected for the current resulting from the reduction of Ga(lll) does intersect at the origin as the concentration of added hydrogen ion becomes zero. TABLE 7 CURRENT-CONCENTRATION DEPENDENCE FOR THE REDUCTION CF HYDROGEN AT THE D.G.E. IN 1 M POTASSIUM CHLORIDE I II III id 0 * ) id (Pa) , [h +] X 10k 1 M KC1 1 M KCl 0.002$ Corrected for Ga^** Triton X-100 0.1 2-7 1.7 2.h 0.2 5.8 5.20 5.2 0.3 9*7 O.U 13.0 10.8 10.2 o.5 13.8 12.7 FIGURE 18 CURRENT-C0NCEI7TRATION DEPENDENCE FOR HYDROGEN ION IN 1 M POTASSIUM CHLORIDE CURVE I . . . . Diffusion Current Measured from the Residual Current Curve CURVE II . . . Diffusion Current Measured from the Residual Current Curve, These Solutions Contain 0.002$ Triton X-100 CURVE III . . . Diffusion Currents for Curve II Measured from the Gallium Limiting Current 79 microamperes i— i— i— i— i— i— i— r j i i i i i i i 8 8 1 An evaluation of n from the slope of log f 1 I versus E Ua ’ gave a value of one; however, this method for the determination of n is limited to reversible systems and the reduction of hydrogen has been shown to be irreversible at the dropping gallium electrode. ^ -ii, // Prom the Ilkovic equation and a measured value of m t , the value of n was 0.63, which was lower than anticipated for a one electron hydrogen ion reduction. However, the value of C used in the Ilkovic equation, the real concentration in the solution, will be different (usually less) from the concentration calculated by dilution of a stock solution. In addition, the value of I^, the numerator in the equation for the calculation of n, has already been shown to have been reduced as a result of auto-oxidation of the gallium metal. These considerations make a low calculated value of n from the Ilkovic equation more understandable. The results of this study at the D.G.E. agree very well with /TO the results obtained by Stelling at a quiet gallium pool; namely, that hydrogen is reduced on gallium at about -1.2 volts (pool electrode) and E^. -1.18 volts (D.G.E.). The Cathodic Curve Produced in 2 Molar Sodium Perchlorate With Added Sodium Bromate Orlemann and Kolthoff^ have examined the reduction of bromate ion at the D.M.E. In acidified potassium chloride and reported an estimated half-wave potential of -O.1* volts (S.C.E. at 25°C.). The cathodic current-voltage curve at the D.G.E. in 2 M sodium perchlorate had an estimated half-wave potential of -1.3 volts (S.C.E. at 35°C»). 8 2 Figure 12, Curve II, is an example of a current-voltage curve obtained with 0.000^5 M Bodlum bromate and 0.005^ gelatin in 2 M sodium per chlorate. The solutions were dilutions of a 0.0015 M sodium bromate solution prepared by accurately weighing the dried sample and diluting to volume with 2 M sodium perchlorate. The observed maximum was not removed by increasing the gelatin concentration. When the gelatin concentration became excessive the entire curve was reduced to an unreliable magnitude. Oxidation of the gallium drop by bromate ion and the subsequent reduction of the oxidation product, either gallium oxide or a basic gallium salt, has been proposed as the mechanism by which the cathodic current originates. Table 8 contains the data obtained for the current-concentration dependence of bromate at the D.G.E. measured at -1.^5 volts and 15 p.s.i. applied pressure. These data have been summarized in Figure 19. TABLE 8 CURRENT-CONCENTRATION DEPENDENCE OF BROMATE AT THE D.G.E. [Br03"] x X0k id (microamperes) 0.75 1.6 2 .2 5 U.6 3-00 7-5 3.75 12.8 ^.50 1 6 .1 FIGURE 19 CURRENT-CONCENTRATION DEPENDENCE FOR BRCMATE ION IN 2 MOLAR SODIUM PERCHLORATE 33 id microamperes o - ro O i « 0) <0 o = ro 01 o> O IC 1" l— i— i— i— i— r l— i— i— i— i— i— i— r OD o ro w I O i * I I I I I I L i i i i i i i i 85 The two slopes observed possibly Indicate that the current is dependent on the rate of formation of the reducible gallium compound formed on the drop surface, which in turn is dependent on the concen tration of bromate added. Table 9 includes the data for the (p.s.i.)? versus id curves in Figure 20. These data indicate that the reaction in the presence of added bromate is diffusion controlled. This is difficult to under stand if the reduction reaction is as has been proposed, the reduction of a solid at the drop surface. An attempt was made to study the effect of increasing acidity (HCIO^) with a constant bromate concentration. An acid concentration of 0.001 M perchloric acid in 0.000^5 M sodium bromate caused the drop time to increase sharply and finally the capillary became plugged. A similar result was observed with nitric acid in 2 M sodium perchlorate. Consequently, acid solutions of these oxidative anions should be avoided, and in general oxidizing agents are to be avoided when working with the dropping gallium electrode. Since the mechanism of the reaction of oxidative anions at the D.G.E. is not known with certainty, a complete evaluation of these results is not possible at this time. TABLE 9 (p.s.i.)^ VERSUS id DEPENDENCE FCFt 0.00015, 0.00025, AND 0.00045 M SODIUM BROMATE Br03-JS 0.00015 M ]= 0.00025 M |BrO3“J= 0.00045 M I P*Y t 1d t *d t *d (sec.) (/») p«8*l« (p.s.i. (sec.) (j») p.s.i. (p.s.i. (sec.) (/») p.s.i. (p.s.i. 5-0 2.2 10 3.16 4.8 4.0 10.5 3.1 7 3.1 12.2 15-5 3.28 4.3 2.7 12 3.^6 3-9 4.4 13.5 3 .6 8 3*3 14.0 14.0 3.54 3.9 2.8 13 3 .6 1 3.5 4.7 15.5 3.94 3-5 14.9 13-5 3.67 3.4 3.1 15 3.87 3.2 5.0 17.0 4.12 3.9 15.5 12.5 3.74 3-1 3.4 16 4.oo 3.0 5.2 1 8 .0 4.24 4.5 1 6 .1 10.75 3.94 3.0 3.6 17 4.12 2.8 5.4 19.3 4.39 2.8 3.8 18 4.24 FIGURE 20 SQUARE ROOT CF APPLIED PRESSURE-CURRENT DEPENDENCE FOR SODIUM BROMATE IN 2 M SODIUM PERCHLORATE MEASURED AT -l.k VOLTS (S.C.E.) AT 35°C. CURVE I . . . . 0.00015 M Sodium Bromate CURVE II . . . 0.00025 M Sodium Bromate CURVE III . . . 0.000^5 M Sodium Bromate 87 a 17 16 15 14 13 12 II 10 9 6 7 6 5 4 3 2 I 0 Square Root of Applied Nitrogen Pressure (PS. I.)' CHAPTER V CURRENT-VOLTAGE CURVES AND CONCENTRATION POLARIZATION; CRITERIA FOR REVERSIBILITY Current-voltage curves obtained with the dropping mercury 7I1 electrode have been explained by the well known Ilkovic equation.1 The equation in Its final form gives an expression for the average current, i, observed during the life of the drop. (12) i = 607 n D^ C m ^ t ^ I - average current = microamperes 607 = constant » (cm?/g^) n = number of electrons gained or lost = (Faradays) J, p , 1 V mole ) D 2 = diffusion coefficient ~ (cm. /sec.)2 C ~ concentration in millimoles per liter ra = weight of mercury/sec. * (mg./sec.) % % t = time for one drop to fall = (sec.) This expression for the average current during the life of a mercury drop must be corrected for the change in density of the metal^1 in order for the equation to be applicable to the dropping gallium electrode. From the above equation it can be seen that the current is inversely proportional to the two-thirds power of the density of the flowing metal. Therefore, the corrected equation for gallium is obtained from the equation for mercury by multiplication of the mercury constant, 6 0 7, by the two-thirds power of the density ratio 8 9 90 of mercury "to gallium (1 3 .6/6 .1 ) . The following equation results: (13) id = 9 9 8 n D^C m^t^ . An equation for the polarographic waves obtained at the dropping mercury electrode for reactions that produce products that are soluble rrQ in the mercury drop was first derived by Heyrovsky and Ilkovic. ' yO ' J. Heyrovsky and D. Ilkovic, Collection Czechoslov. Chem. Com m u n s .. J_3 1 9 8 (1935). cf "Polarography'* pp. 190 et seq. (lb) E„ = Ei - 0 .0 6 6 1 log r i J at 3 5 ° C . The terms in this equation have the same significance for the gallium electrode as for the mercury electrode. These two equations are the equations normally employed in making an analysis of current-voltage curves at the dropping mercury electrode. Reversibility Thermodynamic reversibility, a process which takes place infinitely slowly, such that the real state of the system is only infinitesimally different from an equilibrium state, may or may not be attained in a chemically reversible system.^ Then, ideally, a reaction may not proceed a finite amount in any direction In the case of true reversibility. In the case of galvanic cells, the potentials existent are measured at chemical equilibrium. Here, a very small current may flow without changing the equilibrium potential of the cell by an amount large enough to be measured. The limits of measurement, then, distinguish practical experimental reversibility from ideal theoretical reversibility. 91 Polarographic reversibility has been defined in terms of the equation 1^, the equation describing the polarographic wave. It can be seen from equation l1* that a plot of E versus log f i 1 K - 1 J should produce a straight line with a slope equal to 0.0661 at 35° centigrade. The potential at vhich the log term n becomes zero, i ■ £ id , should be the half-wave potential. The following list of requirements must be met before polarographic reversibility can be assumed:^9 ^ G . W. C. Milner, "The Principles and Applications of Polar ography," Longmans, Green and Company, New York, 1957> pp. 2 5-6 6. 1. The plot of Ej _ vs. log I i (must be a straight line with a slope equal to 0 .0 6 6l . n 2 . The half-wave potential must be independent of concentration of reactants. 3 . The half-wave potential for the anodic process must be equal to the half-wave potential of the cathodic process. U . The combined anodic-cathodic wave should not have any inflection in crossing the zero current axis. 5. The temperature coefficient of the half-wave potential should be nearly constant, * one millivolt per degree. For an j T? irreversible process ^ I0 always large and positive. dT The preceding criteria for polarographic reversibility should be equally valid for the D.G.E. as for the dropping mercury electrode. Therefore, an evaluation of the reactions at the D.G.E. is in order. Figure 9 shows that the reduction of gallium nitrate at the D.G.E. in 7.5 M potassium thiocyanate has an inflection in the combined anodic-cathodic curve in crossing the zero current axis. Conse quently, this system cannot be reversible. Figure l1* shows that the combined anodic-cathodic curve for the reduction of gallium nitrate in acidified 1 M potassium chloride does not have an inflection in crossing the zero current axis. The half wave potential shows considerable variation with a change of concen tration of reactants (see page 70). A comparison of the anodic and cathodlc half-wave potentials is not possible, since the relative supply of gallium metal is infinite, no anodic half-wave potential exists. A plot of E, versus log I 1 was not attempted because U d - a solid product is formed at the drop and this equation is only valid when the reduced species is soluble in the liquid metal. Of the above criteria for reversibility studied, only one, the continuous anodic-cathodic curve, gave results that were in favor of reversibil ity. All of the other criteria tested gave results that were quite different than the expected results for a reversible system. Since the reduction of gallium at the D.G.E. in acidified 1 M potassium chloride is not completely reversible, the effect of temperature on the half-wave potential was not studied. The reduction of hydrogen ion at the D.G.E. has been shown to be irreversible by Grahame.59 Therefore, additional studies of its reversibility were not attempted. CHAPTER VI ANOMALOUS ELECTROCAPILLARY CURVES AT THE DROPPING GALLIUM ELECTRODE Investigations of the capillary characteristics^? of the D.M.E. have shown that m, the mass of mercury flowing per second, measured under otherwise identical conditions is inappreciably changed by a change of salt solutions or the addition of surfactants. The product of mt is very nearly equal to the interfacial tension. Therefore, since m is almost constant, t is fairly closely proportional to the interfacial tension. Drop time measurements with the D.M.E. have been used to study the presence of insoluble films on the mercury drop.^80 Mercurous An I. M. Kolthoff and Y. Oklnaka, J. Am. Chem. Soc., 8 3, 1+7 (1961). halides are known to be only slightly soluble in water. The anodic electrocapillary curves with the D.M.E. in halide solutions show plateaus or potential independent regions attributed to presence of an insoluble insulating film on the drop surface. The return to the normal electrocapillaxy curve at potentials more negative than those at the plateau was attributed to an electrochemical reduction of mercury from the insoluble film and finally the establishment of 93 9^ normal Interface chargee on the cathodlc branch of the electro capillary curve. Early work on the determination of the electrocaplllary curve for gallium was done primarily by the capillary height method and then calculating the interfacial tension. Frumkin and Gorodetzkaya 81 A. N. Frumkin and A. Gorodetzkaya, Z. Fhyslk. Chem.. 1 3 6. 219 (1 9 2 8). determined a value for the electrocaplllary maximum in acidified potassium chloride. Later MuratzeJew and Gorodetzkaya®*5 disagreed 82 A. Muratzejew and A. Gorodetzkaya, Acta Physlochera. U.R.S.S., it, 75 (1936). with the previous work on the basis of a suspected oxide coating on the gallium. This study gave a value of -0.60 volts (N.H.E.) as the potential of the electrocaplllary maximum in 1 M potassium chloride with 0.1 M hydrochloric acid. In addition, the electrocaplllary characteristics were studied in 1 M potassium chloride plus added sur factants, i.e., phenol and lsoamyl alcohol. These authors concluded that capillary active substances shift the electrocaplllary maximum in the same direction as that for mercury. The magnitude of the adsorp tion is different for the two metals with gallium, being less noble than mercury, attaining a larger negative charge on its surface. In 1955 FrumkLn83 repeated his work in 1 M potassium chloride and 83A. Frumkin, Usnehki Khim., 2k, 933 (1955). 1 M hydrochloric acid and agreed with the earlier work of Muratzejew 95 and Gorodetzkaya®^ of -0.60 volts (N.H.E.) for the electrocaplllary maximum. Interfacial energy measurements of the metal-gas interface were made by measuring the angle made by a drop of liquid gallium in contact with a planar surface.®1* The values reported at 30°C. were ®\j. L. Mack, J. Pbys. Chen.. 8i+6 (19^1). for hydrogen-gallium 712 to 732 dynes/cm. and 600 to 800 dynes/cm. for the carbon dioxide-gallium interfacial energy. Electrocaplllary measurements have been made In fused lithium chloride-potassium chloride at ^50° centigrade.®^ The value for the 85 S. Karpachev and A. Stromberg, Acta Iftyslochem. U.R.S.S., 16. 331 (1 9^2 ). electrocaplllary maximum was found to be -0.40 volts (lead reference electrode). The most current investigation of the electrocaplllary character- 5 0 1 sties of gallium was that of Grahame. In this study a dropping gallium electrode was used to make a comparison between the differen tial capacity curves of mercury and gallium. This work can be summarized In the statement: The most remarkable feature of the results is the strong resemblance to those of mercpry. Except for a shift of potential of about O A volts (corresponding to the contact potential of mercury and gallium), the minimum capacity has practically the same value and occurs at about the same value as mercury. Drop time measurements were made in 7-5 M potassium thlocyanate, 2 M sodium perchlorate and 1 M potassium chloride as a method of 9 6 determining the electrocaplllary curve in these electrolytes. The measurements were made at 35°C. with a constant applied nitrogen pressure of 15 pounds per square Inch. A value for the drop time was obtained by averaging the result of the time required for ten drops to fall. The potential at the drop was applied by manually adjusting the potential output of the Leeds and Northrup "Electrochemograph." Anomalous electrocaplllary curves for gallium in the three electrolytes studied, 7 . 5 M potassium thiocyanate, 2 M sodium perchlorate, and 1 M potassium chloride, were obtained by measuring the change In drop time with a change In applied potential. The drop time of the D.G.E. was controlled by the applied nitrogen pressure. The drop time-pressure dependence was measured in 2 M sodium perchlorate. All drop times presented are the average of ten drops. Table 10 contains the data for Figure 21. Curve II, open circuit, and Curve I, -1.45 volts, are similar smooth curves. Over the pressure range studied, the drop time at -1.45 volts remained uniformly smaller than the open circuit drop time. This indicates a reduction of interfaclal tension with an applied potential of -1.45 volts as compared to the Interfaclal tension established under conditions of no applied potential. TABLE 10 THE EFFECT CP PRESSURE ON THE DROP TIME CP THE DROPPING GALLIUM ELECTRODE drop time (sec.) p.s.l. open circuit -1.45 volts 10 5-4 4.8 11 4.4 12 4.5 4.0 13 3.7 14 3*9 3.4 15 3-7 3.2 16 3.4 3.0 17 3.2 2.9 19 3.0 2.8 Additional information about the performance of the dropping electrode in the three electrolytes studied can be obtained from an 2/3 'j, inspection of the values of the product of m t in these electro lytes (Table ll). These values show an order of drop size in these electrolytes, sodium perchlorate> potassium thiocyanate^potassium chloride. FIGURE 21 DROP TIME-APPLIED PRESSURE RELATIONSHIP FOR THE D.G.E. IN 2 M SODIUM PERCHLORATE CURVE I . . . . -1.1+5 volts (S.C.E.) at 35°C. CURVE II . . . Open Circuit at 35°C. 98 Orop Tima (Second*) 0 . 6 2.0 4.0 5.0 3.0 8 10 ple Ntoe Posr ( .I.) S (P Prossur* Nitrogen Applied 12 14 16 18 1 0 0 TABLE 11 % yl ___ DATA FOR m t AT OPEN CIRCUIT FOR THE D.G.E. Drop time Electrolyte in seconds m (mg./sec.) 7 .5 M KSCN 2.75 6.1+3 3.9^ 2 M NaClO^ 3.6 1+.55 k.Qk 1 M KC1 3-03 i+.U 3.2U The data for the electrocaplllary curves obtained at the D.G.E. are given in Table 12. Figure 22 is an example of an electrocaplllary curve obtained in 7-5 M potassium thiocyanate (Curve I). Curve II is the result of adding 0.0095 M potassium nitrate to the thiocyanate solution. The added nitrate only changed the shape of the anodic branch of the curve. This is a normal anion dependence expected in electrocaplllary curves. An unusual plateau (potential independent region) can be seen extending from -1.25 to -1.1+0 volts. An examination of the influence of acid vas not attempted because of the acid sensitivity of thiocyanates. Figure 23 is an example of an electrocaplllary curve observed in 2 M sodium perchlorate (i). In this electrolyte a maximumcan be seen in the curve at -1.05 volts. In addition, two potential independent regions are found from -1 .2 5 to -1 .3 5 volts and from -1.1+5 to -I.60 volts. The addition of 0.001 M hydrochloric acid to the perchlorate solution lowered the drop time but did not improve the shape of the curve. TABLE 12 DROP TIME DATA FOR ELECTROCAPILLARY CURVES OBTAINED WITH THE D.G.E. AT 15 p.s.l. NITROGEN PRESSURE d r o p time In s e c 0 nd s Figure 22 Figure 23 Figure 2U Figure 25 7.5 M KSCN 2 M NaClOj, 1 M KC1 1 M KC1 -E(S.C.E.) 7-5 M KSCN + 0.0095 2 M NaClO^ + 0.001 1 M KC1 + 0.0 0 1 1 M KC1 1 0 .0 1 M KNOj M HCIO^ M HC1 M HC1 0 .7 0 2 .3 0 2.30 3.70 3.6 0 1*.00 1 .8 0 0.75 3.50 3.1*0 3 .8 0 3-20 0 .8 0 2.40 2 .1*0 3.1*0 3.30 3 .6 0 3.60 0 .8 5 3.30 3.20 3 .1+0 3.60 0 .9 0 2.70 2 .8 0 3.90 3.75 3.30 3.1*0 3 .1+0 3.60 0.95 3.50 3-50 3.6 0 3.6 0 1.00 3.05 3.20 4.20 3.70 3.50 3.50 3 .6 0 3.6 0 1 .0 5 3.50 3-50 3 .6 0 3.60 1.10 3-25 3.30 1* .20 3 .6 0 3.50 3.50 3.60 3 .60 1.15 3.30 3.30 3.50 3.!+0 3 .60 3 .6 0 1.20 3.30 3.30 3.90 3 .6 0 3.1+0 3*1+0 3.55 3 .6 0 1.25 3.20 3.20 1.30 3.20 3.20 3.80 3.50 3.30 3 .1*0 3 .5 0 3.70 1.35 3.20 3.20 3 .8 0 1.1+0 3.20 3.20 3.75 3-^0 3 .2 0 3.3 0 3.1+0 3.85 1.1*5 3.10 3.10 3.70 1.50 3.00 3-00 3.TO 3.20 3.10 3 .2 0 3.30 1+.00 1-55 3.70 3.20 1.60 2 .9 0 2.90 3.65 3.20 3.00 3 .1 0 3.20 1.6 5 3.6 0 3.15 1.70 2 .8 0 2.80 3.55 3.10 FIGURE 22 ELECTROCAPILLARY CURVES BY THE EPOP TIME METHOD IN 7-5 M POTASSIUM THIOCYANATE CURVE I . . . . 7 .5 M Potassium Thiocyanate CURVE II . . . Curve I plus 0.0095 M Potassium Nitrate 102 Drop Time (Seconds) 2.0 2.2 2.8 2.5 2.9 3.0 2.4 2.6 2.7 3.4 3.3 3.2 0.7 0.8 0.9 1.0 * SCE o 35*C ot S.C.E. v*. E - l.l 1.2 1.3 .4 1.5 1.6 1.7 103 FIGURE 23 ELECTROC APILLARY CURVES BY THE DROP TIME METHOD IN 2 M SODIUM PERCHLORATE CURVE I . . . . 2 M Sodium Perchlorate CURVE II . . . Curve I plus O.OOl M Perchloric Acid Drop Timo (Second*) NWwW UOioibibiWbif » * i o 6 '“ 'n 'w V b « o » >j o d < 0 o Ki irt “ 1 1 ! 1 1 1 1 1 1 1 1 T ~ o (0 ro co M o (H d o ' o* SOT 1 0 6 Figures 2k, 25 represent the results of electrocaplllary measure ments in 1 M potassium chloride obtained with two different capil laries. A comparison of the curves (i) In each case show good reproducibility. The potential region in which a maximum was observed by Muratzejew and Gorodetzkaya (-O.9 0 volts) is on the anodic limit of a potential independent region extending from -0 .9 5 to -1 .1 5 volts. The drop time increased at potentials more positive than -0.88 volts. This is the potential at which a change in slope was observed in the anodic current-voltage curve. The increase in drop time could be the result of the formation of an oxide film on the drop surface. Gallium oxide adheres to glass, and consequently would increase the drop time. A small depression was observed in the vicinity of -1.2 volts. This is the potential region at which hydrogen ion was reduced at the dropping gallium electrode. This same potential was observed to be the beginning of a plateau in the cathodic branches of the curves obtained in thiocyanate and perchlorate. The addition of hydrochloric acid, 0.001 M Figure 2k Curve II and 0.01 M Figure 25 Curve II, influenced both the anodic and cathodic branches of the curves. On the anodic branch the added acid lowered the interfacial tension. This could be through solution of the proposed oxide film. The interfacial tension increased with added acid at potentials more negative than -1.2 volts. This increase could be due to the formation of a gas-metal interface as the hydrogen ion is reduced on the surface, thus increasing the interfaclal tension. FIGURE 2h ELECTROCAPILLARY CURVES BY THE EROP TIME METHOD IN 1 M POTASSIUM CHLORIDE CURVE I . . . . 1 M Potassium Chloride CURVE II . . . Curve I plus O.OOl M Hydrochloric Acid 90T O.fifi 10 *3'0'« '•* 3 S s 2 5 in n *-» rp i* (Seconds) Tim* Drop K» m c CM CM oo CM FIGURE 25 ELECTROCAPILLARY CURVES BY THE DROP TIME METHOD IN 1 M POTASSIUM CHLORIDE CURVE I . . . . 1 M Potassium Chloride CURVE II . . . Curve I plus 0.01 M Hydrochloric Acid 109 Orop Tim# (Seconds) 3.0 2.6 3.3 3.5 2.9 3.4 3.2 3.6 3.6 4.0 3.7 3.9 2.7 0.6 0.7 0.8 0.9 - E vs. S.C.E. ot ot S.C.E. vs. E - 1.0 1.1 35*C 1.2 1.3 1.4 1.5 1.6 Ill The gallium drops formed in potassium thiocyanate, sodium perchlorate, and potassium chloride vere bright and shiny. The accumulated gallium at the bottom of the cell agglomerated into a single drop at potentials more negative than the potential at which gallium oxidized. However, the accumulated gallium was coated with a film when the drop was anodically polarized. The accumulated gallium was not covered with a film in acid solutions. The film was probably gallium oxide which is soluble in acid. GaB bubbles, probably hydrogen, were observed on the accumulated gallium when the concentration of the acid (HCl) was 0.1 molar. The potential independent regions observed in the electrocapil lary curves of gallium obtained with the D.G.E. indicate that a change in potential in thi6 region has no influence on the drop time and therefore no influence on the interfaclal tension. This signifies that the drop surface is insulated from the solution by a dielectric substance. The potential independent region was most pronounced in potassium chloride. Since the chloride salts of gallium are soluble, the nature of the dielectric film on the drop surface remains unidentifled. The electrocaplllary curve obtained with the D.G.E. in 1 M potassium chloride with 0.01 M hydrochloric acid does not agree with Qp the curve obtained by Muratzejew and Gorodetzkaya. These workers reported data from -0.86 to -1.6 volts (N.C.E. at 35°C.). A plot of this data, Table 13, Figure 26, shows a small depression in the electrocaplllary curve in the vicinity of -1.2 volts. A similar depression was observed in 1 M potassium chloride with the D.G.E. 1 1 2 However, In solutions of 0.01 M hydrochloric acid the electrocaplllary curve obtained with the D.G.E. showed an increase at this potential because of the formation of a gallium-hydrogen gas interface arising from the reduction of hydrogen at the D.G.E. at this potential. TABLE 13 ELECTR 0C APILLARY nATA OBTAINED BY MURATZEJEW AND GORODETZKAYA- 2 -E (N.C.E.) dynes/cm. 0 .8 6 611.5 J 0 .9 0 613.3 1 .0 611.2 0.1 M HCl 1 .1 604.7 1 .2 595.7 ^ f 1.3 589.7 1.4 578.4 0.01 M HCl 1.5 566.4 1 .6 552.1 f Current-voltage measurements at the D.G.E. in acidified 1 M potassium chloride showed that the gallium electrode was spontaneously oxidized by acid (HCl). In addition, hydrogen was electrochemically reduced at the D.G.E. Comparison of these results with the electro capillary curves obtained with the D.G.E. leads to the conclusion that electrocaplllary curves for gallium obtained in acid solutions are not valid. FIGURE 26 A PLOT OF THE ELECTROCAPILLARY DATA OBTAINED BY MURATZEJEW AND GORODETZKAYA82 113 Dyn«*/Cm. 540 580 590 550 520 530 570 560 0 2 6 600 510 610 0.7 0.8 0.9 O.IM HCt O.IM 1.0 _L 1.2 0 1.3 1M HCl 0 01 M 1.6 0.01 M NoOH 1.7 CHAPTER V n THIOCYANATES Thiocyanates are prepared by numerous reactions. The following are only intended to be representative: These preparations and the (15) 2NH3 ♦ CSg * NH^NH2CS2 NH^SCN ■* H2S (16) Na^FefCNjg 4 2Na2S^ = 6NaSCN 4 FeS ♦ NagS (IT) NaNOg t CS2 + 2HgS • NaSCN ♦ S3 -► 2H20 (18) NaNHg ♦ CS2 = NaSCN 4 HgS reactions of thiocyanate and its metal salts have been reviewed by 86 H. E. Williams. The thiocyanates of the alkali and alkaline earth 86h . b . Williams, "Cyanogen Compounds," Edward Arnold and Company, London, 19^8. metals are highly soluble in water, e.g., KSCN is soluble to the extent of 239 grams per 100 ml. of water at 25°C., and in addition 20.3 grams will dissolve in 100 ml. of acetone. The water-soluble salts of the heavy metals are generally soluble in alcohol and acetone, and some are soluble in ether. Concentrated solutions of thiocyanates are known to undergo oxidative attack and are attacked by concentrated strong acidB. oxidative attack: (1 9) 2KSCN * 302 + 2HpO = 2KHSO4 4 2HCN 115 1 1 6 (20) KSCN ♦ 4Br2 + i+HgO = KBr + CNBr i ^30^. 4 6HBr The electrolysis of thiocyanates produces only cyanide and sulfate. concentrated nitric acid: The addition of concentrated nitric acid to moderately strong solutions of thiocyanate is capable of producing a reaction of explosive violence which is accompanied by a rapid evolution of hydrogen cyanide. acid attack in aqueous solution: (21) 3KSCN -l 3HC1 - H g C ^ g S ^ 3KC1 4 HCN (22) 3KSCN 4 3H2SO^ * H2C2N2S3 4 3KHS01* 4 HCN The compound H2C2N2S3 is yellow xanthane hydride, s = c s \ \ H — N ---- C — N — H A red intermediate is formed in these acid reactions that is apparently unstable in favor of the insoluble yellow xanthane hydride. 8tE. V. Zappi, Bull, soc. chim.. Vf, '+53 (1930), Uc, 397 (1932). OQ Kolthoff and Lingane investigated the purification, standard ization, and storage of thiocyanate solutions. These workers 88I. M. Kolthoff and J. J. Lingane, J. Am. Chem. Soc., 57, 2126 (1935). recommend recrystallization from water to remove chlorides and standardization with silver nitrate using either the potentiometric or absorption indicator endpoint. In addition, it was shown that potassium thiocyanate could be dried safely at 150°C. and its aqueous 117 solutions were stable if stored In the dark In a closed bottle. Aqueous solutions of potassium thiocyanate have been reported to undergo decomposition when exposed to sunlight. The product formed was thought to be yellow amorphous pseudo-thiocyanogen, (SCN)X .89 8*D. Ganassini, Boll. Chim. Farm.. 5 8, 457 (1919)- 00 Jones has examined the infrared spectra of aqueous solutions 9°L. H. Jones, J. Chem. Phys.. 2£, 1069 (1956). of thiocyanate and reports that the bond distances show considerable double bond character for the carbon to sulfur bond. This would indicate that the thiocyanide ion is SCN“. Jones also states that the smallness of the absorption coefficients makes it practically impossible to study the effects of complexing on any frequency but that of the C-N vibration in aqueous solution. Aqueous solutions of potassium thiocyanate are transparent in light of visible wavelength. Kiss and Csokan^ examined the ^ A . Kiss and P. Csokan, Z. Fhysik. Chem., A 186, 239 (1931). ultraviolet spectrum of potassium thiocyanate and showed that aqueous solutions of potassium thiocyanate are not transparent beyond 300 millimicrons. Jablczynski and Jablczyska^ have shown that solutions of J. Jablczynski and H. Jablczyska, Rocznlki Chem., 10, 579 (1930). C.A.. 2%, 877s (1931). 118 thiocyanates will decompose when subjected to strong ultraviolet radiation. The decomposition is an endothermic separation of sulfur with the accompanying chemical equilibrium: SCN~ S(0) ♦ CN“ . The ultraviolet spectrum of gallium-thiocyanate solutions does not exhibit any characteristic absorption peaks.93 93r . P. Buck, S. Singhadeja and L. B. Rogers, Anal. Chem.t 26, 2ho (195*0. Reactions of Gallium Salts in 7.5 M Potassium Thiocyanate Solutions of gallium salts in 7.5 M potassium thiocyanate have been observed to become yellow after standing for twenty-four hours. This observation was confirmed during the course of this work. How ever, if these yellow solutions were permitted to stand for three or four days, a yellow amorphous precipitate was observed in solutions of 0.001 M gallium nitrate in 7.5 M potassium thiocyanate. Control solutions of thiocyanate and thiocyanate - 0.003 M potassium nitrate showed no change after several months. The yellow solutions were formed if the solutions were exposed to ambient light or protected from light by storage in a closed cupboard. The rate of formation of the yellow precipitates was increased with an increase in temperature. Thiocyanate solutions containing gallium nitrate reacted much more slowly when the solutions had been outgassed with nitrogen. The outgassed solutions formed precipitates after three week6. Solutions in which the thiocyanate concentration was less than five molar did not form precipitates. The volume of the precipitate increased as the 119 concentration of thiocyanate increased. The concentration of the gallium nitrate added was a constant value of 0.001 molar. The precipitates formed by adding gallium nitrate to 7-5 M potassium thiocyanate produced an amorphous yellow precipitate. Adding the 7.5 M potassium thiocyanate to a weighed quantity of gallium perchlorate 6-hydrate produced a "red flush" on the surface of the solid phase. The "red flush" could be due to the rapid hydrolysis of gallium perchlorate and the formation of a local high concentration of acid which oxidized the thiocyanate to a red intermediate. The red intermediate is unstable in aqueous solutions in favor of yellow xanthane hydride, H^C^N^S^. Two drops of concentrated hydrochloric acid added to 25 milli liters of 7-5 M potassium thiocyanate produced large well formed yellow needles arranged in a star-like cluster. The amorphous precipitates from the gallium solutions viewed under a microscope appeared to be composed of thin yellow wafers. The crystalline precipitate was removed from the solution by filtration. Separation of the amorphous precipitates from solution was accomplished by centrifugation. The precipitates were not soluble in water, carbon disulfide, benzene, acetone, or ethanol. They were soluble in lO' sodium hydroxide. Xanthane hydride and gallium hydroxide are soluble in excess base. 1 2 0 A qualitative ultraviolet examination was made of the solutions above the precipitates, using a Cary recording spectrophotometer (Figure 2 7), The solutions were diluted with 7*5 M potassium thio cyanate until the concentration was one-hundreth of the starting concentration. The reference solution was 7.5 M potassium thiocyanate. Curve III, the solution containing hydrochloric acid, had absorption peaks at 3^2 and 310 millimicrons. Curve II, the solution containing gallium perchlorate, had absorption peaks at 36^, 3^2 and 310 milli microns. Curve I, the solution containing gallium nitrate had absorption peaks at 36^ and 30^ millimicrons. These results show that the peak at 310 millimicrons 1b common to all three solutions and that the peak at 36^ millimicrons is common to the gallium- containing solutions. The solution containing gallium perchlorate exhibited an absorption curve that appears to be a composite acid- galliura-thiocyanate curve. This could be due to greater hydrolysis of gallium perchlorate than gallium nitrate or to residual perchloric acid in the gallium perchlorate precipitate. It appears on the basis of these curves that the yellow solutions above the two kinds of precipitates are not the same. Consequently, it must be assumed that the precipitates formed in the three different solutions are not the same. FIGURE 27 ULTRAVIOLET ABSORPTION SPECTRA CF THE SOLUTIONS ABOVE THE YELLOW PRECIPITATES FORMED IN 7.5 M POTASSIUM THIOCYANATE CURVE I . . . . Gallium Nitrate CURVE II . . . Gallium Perchlorate CURVE III . . . Hydrochloric Acid 1 2 1 Per Cent Transmission & SUMMARY AND CONCLUSIONS A glass apparatus was constructed and used as a D.G.E. Oxide free gallium at 35°C. was forced through a glass capillary by nitrogen pressure. The D.G.E. was operated successfully in 7*5 M potassium thiocyanate, 2 M sodium perchlorate, 1 M potassium chloride, and 0.01 M sodlian hydroxide. The lumen of the capillaries and the gallium drops were larger than those normally encountered with the D.M.E. Consequently, the currents observed at the D.G.E. were larger than would be observed at the D.M.E. in the same test solution. A 0.000065 M gallium nitrate solution in 1 M potassium chloride was found to be 87-7$ hydrolyzed. Equilibrium concentrations of Ga(lll) and H~ were determined from the cathodic curves obtained for these two ions at the D.G.E. and a hydrolysis constant, U.J x 10~3, was determined from these values which is in agreement with the literature value of U.03 x 10“3. Additions of acid (HCl) to suppress hydrolysis produced a maximum in the gallium curve at -1.0 volts which increased in magnitude until a limiting value was obtained when the ratio of [,H*j to ^ G a ^ J was 50. The magnitude of the current at this maximum gave a linear current- concentration plot over the concentration range 0.000012 to 0.0002U M gallium nitrate in 1 M potassium chloride with 0.002^ Triton X-100. The gallium curve had half-wave potential values of -0 .90 to - O .85 123 1 2 4 volts over this concentration range, thus the reduction of gallium at the D.G.E. under these conditions cannot he completely reversible. The reduction of hydrogen Ion at the D.G.E. proceeds irreversibly with a half-wave potential of -1 .1 8 volts over the concentration range 0.0001 to 0.0005 M hydrochloric acid in 1 M potassium chloride. The reduction process was found to be diffusion controlled and followed a linear current-concentration relationship over the concentration range mentioned. Gallinn metal, being less noble than mercury, was found to be subject to chemical oxidation by bromate, nitrate, and nitrite ions producing either galliun oxide or a basic gallium salt at the drop surface which was subsequently electrochemically reduced at about -1.3 volts. Hydrogen ion also oxidized gallium metal, but to Ga(lII), which was electrochemically reduced at about -0.90 volts. An anodically polarized D.G.E. exhibited a decrease in slope of the current-voltage curve at -0.86 volts in 7.5 M potassium thio cyanate, 2 M sodium perchlorate, and 1 M potassium chloride indicating an increase in passivity at this potential. The anodic current- voltage curves obtained with the D.G.E. in acidified 1 M potassium chloride showed two regions of anodic passivization, -0.86 volts and -0 .7 6 volts, which are probably due to two different insulating gallium compounds with different acid solubilities. The electrocapillary curves obtained with the D.G.E. in 7.5 M potassium thiocyanate, 2 M sodium perchlorate, and 1 M potassium chloride are anomalous when compared to electrocapillary curves for mercury. The electrocapillary curves obtained In acidified 1 M potassium chloride with the D.G.E. are not in agreement with the literature measurements. As a result of this study, the validity of electrocapillary measurements made on this reactive surface in acid solutions is dubious. A yellow amorphous precipitate found in gallium solutions of 7.5 M potassium thiocyanate was compared to yellow xanthane hydride and was found not to be the same. The nature of the yellow precipitate in gallium solutions of 7*5 M potassium thiocyanate remains to be identified. LITERATURE CITED 1. D. I. Mendeleeff, Ann., Supplementband 8, 2, 196 (1 8 7 1). 2. L. de Boisbaudran, Ann. Chlm- et Fbys.. 10, 100 (1877)* 3* L. de Boisbaudran, Chem. News, 33, Ik6, 157, 167 (1877). *+. C. A. 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Zappi, Bull, soc. chim., kj, ^53 (1930), M?, 397 (1932). 8 8. I. M. Kolthoff and J. J. Lingane, J. Am. Chem. Soc., 57, 2126 (1935). 8 9. D. Ganassini, Boll. Chim. Farm., 5 8, ^57 (1919). 90. L. H. Jones, J. Chem. Phys., 2 5, 1069 (1956). 91. A. Kiss and P. Csokan, Z. Physlk. Chem., A 186, 239 (1931). 92. J. Jablczynski and H. Jablczyska, Rocznlki Chem., 10, 579 (1930). C.A.. 2£, 877s (1931). 93- R. P. Buck, S. Singhadeja and L. B. Rogers, Anal. Chem., 26, 2U0 (195M. AUTOBIOGRAPHY I, Norman Crabb, was b o m in Kent, Ohio, on February 17, 1933- My high-school education was received at the Kent State University School. My undergraduate study was completed at Kent State University, and I received a Bachelor of Science degree In June 1955* My graduate career at The Ohio State University began in the autumn of 1955* A Master of Science degree was conferred in June 1958 upon completion of a thesis entitled "A Thermogravimetric Investigation of Some Metal Versenates." During my graduate years, I was employed in the Department of Chemistry as a teaching assistant and research assistant. Beginning in the autumn of i9 6 0, I was appointed assistant instructor and continued in this capacity until August 1 9 6 1. 131