SIMULTANEOUS COULCMETRIC ANALTSIS

DISSERTATION Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University

By

RICHARD DONALD McIVER, B«A.

The Ohio State University

195k

Approved by

Adviser ACKNGWLED GMENTS

I wish to express ny sincere appreciation to Dr® William MacNevin, a most competent adviser and friend, to the Proctor and Gamble Comparer and the Al­ lied Chemical Company both of which awarded me fellow­ ships, and to u$r wife who was both an inspiration and help during ray graduate studies® I also wish to thank Dr. Bertsil Baker for his assistance and suggestions and to the Department of Chemistry of The Ohio State University in which I served two years as a Graduate

Assistant.

i

A. 4 8 2 7 2 TABLE OF CONTENTS page

THE PROBLEM...... 1 INI R QDUCTI O N ...... 2 The Faraday and Electrolysis...... 2 Instruments for Measuring the Number of Coulombs * . • • 2

Electrolysis at Constant Current ...... 3 Electrolysis at Controlled Potential ...... 5 i Devices for Controlling Electrode Potential ...... 6

Coulometric A n a l y s i s ...... 7

Constant Current Coulometry ...... * . 8 Controlled Electrode Potential Coulometry...... 10 Coulometry Applied to Indirect Analysis ...... 11 Indirect Analysis and Simultaneous Equations...... 12 Propagation of Errors in Simultaneous Equations ...... 1U

Applications of Coulometry to Indirect Analysis ..... 18

EXPERIMENTAL...... 20 General Apparatus and Materials ...... 20 Apparatus for the Automatic Control of Electrode Potential •• ...... 20 Balance and Weights ...... 2h

Reagents ...... 2U Saturated Calomel Electrode ...... 2$ Barometer...... 25

Hydrogen-Ctxygen Coulometer...... 25

ii Table of Contents (continued) Page Investigations with the Rydrogen-Qxygen Coulometer . 0 . 29 Construction . • • ...... 29 Investigation of Potassium Dichromate Solution for the Coulometer Electrolyte...... 29 Investigation of Current Density and its Effect on Accuracy of the Coulometer...... 30

Results and Conclusions...... o . . 32 Simultaneous Determination of Chloride and Bromide Ions . 3U Preliminary Considerations...... • « 3k

Special Apparatus...... 36 Preliminary Investigations...... • 38 Experimental Procedure • ...... 39 Results and Conclusions...... • ...... 1*0 Simultaneous Determination of Zinc and Cadmium . . . . • 1*1* Preliminary Considerations ...... 14*

Special Apparatus...... 1*5 Preliminary Investigations...... 1*8 Experimental Procedure ...... £1*

Results and Conclusions...... 55 Simultaneous Determination of Lead andTi n ...... 58 Preliminary Considerations...... * 58 Special Apparatus...... 60 Preliminary Investigations...... 62

Experimental Procedure ...... 65 Results and Conclusions ...... • 67

iii Table of Contents (continued) Page

Simultaneous Determination of Silver and. Copper .<>.<>. 69 Preliminary Considerations ...... 69 Special Apparatus ...... 70

Preliminary Investigations ...... 72 Experimental Procedure • ...... 75 Results and Conclusions , ...... 76 The Investigation of Residual Currents ...... 78 Preliminary Considerations ...... 78 Special Apparatus ...... 78 Experimental Procedure ...... 80

Results and Conclusions ...... 82

SUMM&HY...... » ...... 86

BIBLIO GR A P H Y ...... 88 AUTOBIOGRAPHY...... 92

iv THE HtOBLEH

The objective of this research was to extend controlled potential

electrolysis and coulometric analysis to indirect methods of analysis of two or mare substances which occur together and which are difficult to separate* This method of determination, although limited by the relative concentrations of -the two substances, sometimes is a more rapid method than conventional electrolytic or coulcmetric methods and in some cases gives good results where electrolytic separation of the substances is not possible.

1 INTRODUCTION

The Faraday and Electrolysis

In the study of phenomena associated with electrolysis. Michael Faraday (22,23) recognized the equivalence between the quantity of electricity passed through an electrolyte and the amount of chemical reaction brought about at the electrodes. These observations led to the postulation of the now famous laws of electrolysis which bear his name. To produce one equivalent of chemical change at an electrode there is required a definite number of ampere-seconds which are called coulombs. The value of this number, the far ad ay, has been precisely determined to be 9 6,1*9k *2 .3 coulombs in studies with the silver cou­ lometer (55,56,37,58)o Studies with other coulometers give values ranging from 9 6,1*90 + 1*0 for the copper coulometer to 9 6 ,3 9 0for the hydrogen-octygen coulometer. After surveying these results, Birge (2,

3,U) estimated the best value to be 9 6 ,5 0 1 + 1 0 coulombs.

Instruments for Measuring the Number of Coulombs The number of coulombs passing through a circuit is most con­ veniently measured by maintaining a constant current and accurately measuring the time elapsed. Since this is often neither possible nor advisable, it is best to use a coulometer, or voltameter, to measure the total current. A coulometer is simply an instrument of one of two types, the most common of which is an electrolysis cell in which the reaction at one or both electrodes proceeds with 100 percent current efficiency and can be quantitatively measured by weighing, titration, observation of gas volume, stripping a deposit at constant current, etc . Examples of this type are the silver, the copper, the cadmium, the iodine, the hydrogen-axygen, and the coulometric (2 1) coulcmaters . The other type of coulometer is the electromechanical type© This is in the farm of an automatic recorder which describes a curve under which the area is proportional to the current. This area may be meas­ ured by several methods including the paper weight method (1 8), con­ verting to the slope-intercept form on sendlogarithmic graph paper followed by suitable calculation (1,U7), or by mechanical or electronic integration (£,3 6,UUjUE>) •

Electrolysis at Constant Current In the most common method of electrolytic deposition, the total applied potential across the electrolysis cell is adjusted to maintain the current at some suitable value which is sometimes arbitrarily chosen. Under such conditions, only the most noble metal is deposited first but, as the concentration of this metal is reduced, the oxidation potential gradually decreases and the next less noble metal is depos­ ited. In this method there is no good way of determining just when to stop in order to obtain a deposit of only the most noble metal. This subsequent reduction of less noble metals continues until hydrogen is liberated, at which point the deposition of the metals below hydrogen in Table I may be considered essentially complete. The deposition may be extended to seme metals above hydrogen in the table by reducing the acidity of the solution and thus lowering the oxidation potential of the hydrogen or by taking advantage of the higher overvoltage exhibited by hydrogen toward some electrode materials such as copper, silver,

3 and mercury®

Table I (37) STANDARD REDUCTION POTENTIALS

1C5* + e" cs K -2*92£ volts Na+ + e“ S3 Na -2 .7 lit + 3e“ B A1 -1*66 Zn+* + 2e“ a Zn —0 *7 63 Fe** + 2e~ as Fe -O.LtitO Cd** + 2e“ a Cd -0.2t03 Ni 2e” C3 Ni ~0.2£0 Sn** + 2e“ Q Sn -0 .1 3 6 Pb++ + 2e“ a Fb -0 .1 2 6 Normal Hydrogen Electrode 0 .0 0 0 Saturated Calomel Electrode 0 .2U£8 Cu *{* 2e~ C Cu 0.337 Ag* ■{* e“ B Ag 0.7991 Au + 3e“ O Au 1 .5 0

This constant current method is generally limited to the separa­ tion of single metals frcm solution or to the separation of metals be­

low hydrogen in Table I from those above® In seme cases selective deposition of one metal in the presence of another may be effected without controlled electrode potential even

though both are below hydrogen in Table I. This is brought about by altering seme chemical property of one or all of the cations by cdu­ plexing them or by other means* If only one cation forms a complex with the addition of a ccmplexing agent, or if the stability of the

complexes formed by different cations is sufficiently different, it is sometimes possible to deposit selectively one cation in the pres­

ence of the others* Examples are the separation of cadmium from copper

in cyanide solution (2 6) and nickel frcm zinc as the ammina complexes

discussed in Willard, Merritt, and Dean (67) * k Furman (28) has studied the use of 'potential buffers' to limit cathodic reduction to a certain potential range. These buffers are of such nature that they are reduced at a certain potential and thus act in the same manner as lydrogen to limit the reduction at the cathode to ions with cocidation potentials above the oxidation potential of the buffer.

Electrolysis at Controlled Potential The second method of electrolytic deposition is that in which the total applied potential* the potential across the entire electrolysis cell* is kept constant. This method* originated by Brunck (7) lacks the advantages of the speed of the constant current method and the se­ lectivity of the controlled electrode potential method discussed be­ low. A complete discussion serves little purpose in this dissertation* but it is of historical interest because it preceded* and perhaps led to* the development of the controlled electrode potential method. The third method of electrolytic deposition is that in which the potential of the working electrode* visually the cathode* with respect to a reference electrode* is controlled within certain limits so that it is maintained essentially constant. This method* originated by

Sand (60) is advantageous because the potential of the working elec­ trode is the primary factor in determining which of two or more possi­ ble ions will be deposited at the electrode. The potential is main­ tained as nearly constant as possible with respect to a reference electrode by the regulation of the total applied potential of the cell. This compensates for natural changes in the electrode poten­ tials* the changes in the overvoltages on both of the electrodes* and

5 the decrease of the iR drop of the cell* all of which occur during electrolysis * By inserting a reference half cell into the solution and meas­ uring the potential between the working electrode and the reference electrode* it is possible to isolate the effect of the working elec­ trode. This potential between an electrode and a solution containing ions of the electrode material consists of the equilibrium potential of the electrode metal toward the solution of its ions and the over­ voltage caused by the current flow. The expression for equilibrium potential is shown in the equation

E » E0 * (0.059Vn) log (Hn*) at 25°C. (1)

In this equation E is the potential of an electrode in a solution of its ions of activity Mn+* E0 is the standard reduction potential of the electrode in a solution in which all the ions involved in the electrode reaction are at unit activity* and n is the number of elec­ trons involved in the electrode reaction* As the metal ion concentration decreases during an electrolysis the second term of the equation gets smaller and the reduction poten­ tial is decreased (moved frcm bottom toward the-top of Table I)* In this manner the change in concentration of the ion may be followed during electrolysis* or the separation of various ions can be brought about by the control of the potential.

Devices for Controlling Electrode Potential Originally the applied potential of the electrolysis cell was controlled manually to maintain the cathode potential which* by means of a potentiometer set-up* was checked against a calomel electrode* 6 Very close attention, was required and the operator was kept busy until the electrolysis was c ample ted. This tedious manipulation was the fac­ tor which perhaps hindered the growth of constant electrode potential electrolysis mare than any other® It was the advent of automatic con­

trol devices which revived the interest in the method and its many possible applications. Hickling (31) was the first to describe an automatic potential controlling device which he called a ‘potentiootat' . This electronic device was limited to fairly loti currents which prolonged the time of the electrolysis somewhat and it lacked the close control that later devices have. Caldwell, Parker, and Diehl (10) and then Diehl (17) followed with improved models which were also limited because they would keep the cathode frcm becoming more negative but would not com­ pensate far potentials drifting in the opposite direction. These po­ tential drifts sometimes carried the potential beyond the lower limit

set for an electrolysis. Since the development of the Lingane and Jones (39,UU) electro­ mechanical Instrument, many other electronic potentiostats have been devised, each with its own advantages and improvements, but the Lingane potentiostat, controlling equally well in either direction, is satis­ factory for all except the most precise and demanding work. It was this instrument, constructed by Tuthill (65), that was available in this laboratory and was used throughout this work.

Coulometric Analysis The term •coulometric analysis' was introduced by Szebelledy and

Somogyi (6 3) who were pioneers in the application of Faraday's Laws

7 to electa*©analysis. These authors recognized that fry measuring the amount of electricity, expressed in coulombs, required to bring about complete or essentially complete reaction of a substance in solution, the amount of that substance present could be determined. The sub­ stance being determined need not react at the electrode (primary cou­ lometry) but may react with an intermediate which is generated at the electrode (secondary coulometry). The limitations of this method are that only the desired reaction take place during the electrolysis and that this reaction proceed with an accurately known current efficiency, preferably 100 percent. The latter limitation has been met in the analyses thus far performed in that only reactions that proceed with 100 percent current efficiency have been employed.

Since the inception of the coulometric method, the experimenta­ tion has diverged into two branches which are altogether different. These are the controlled potential method, originated by Hickling (31)> and the constant current method of Szebelledy and Somogyi (6 3). In the former case, the electrode potential of one of the electrodes is maintained constant until the current falls to a negligible value in­ dicating that the reaction is complete. In the latter case, the sub­ stance to be determined is electrolysed at constant current until an indicator system indicates the completion of the reaction.

Constant Current Coulometry The constant current methods of coulometry are essentially titri- metric methods in which the volume of reagent required to bring about the desired reaction is replaced by the number of coulombs required to generate the reagent or bring about the desired reaction at the electrode surface. In fact, it has been suggested that chemical pri­ mary standards be replaced by the coulomb for all titrimetric analy­ ses (6 6 ) 0 One of the advantages of this method is the simple method of determining the number of coulombs — accurately measuring the time while the current is held constant. This makes possible the measure­ ment of very small amounts of reaction and thus makes very accurate micro determinations possible. In general the advantages and limita­ tions of this method are the same as those of constant current elec- trodeposition. In addition, an indicator system is needed to detect the end point. The generated reagent not only has to be generated at 100 percent current efficiency, but also must react stoichiometrically with the desired constituent in solution. The generation of the reagent does have the advantage of limiting or stabilising the anode potential so that no interfering substances react and, in addition, generated re­ agents which otherwise are difficult to employ as standards may be used. Examples of such standards are the cuprous ion and free bro­ mine • Investigation of this technique have been carried out to a large extent by Swift and cowarlcers (6,8 ,9,21;,25,^9,50,51,53*59 ,61,69), Furman and cowarlcers (1U,29,30,5U), and numerous other authors. The first two authors worked with applications in which the reagent was generated in the electrolysis cell (internally generated reagent method). Deford and coworlcers (16) and Pitts and coworkers (52) found that there were advantages to generating the reagent in an external system and introducing this reagent into the reaction vessel through 9 a capillary (externally generated reagent method). The general requirements for automatic constant current titration equipment for coulometric titrations have been discussed try Carson (11,

1 2 *1 3) who devised an electronic anticipator circuit which makes pos­ sible the very precise determination of the end point of the titration. Other authors also have explored automatic titrations with success (1*2*

62*68)•

Controlled Electrode Potential Coulometry Hicl&ing (31) reported the determination of copper by measuring the number of coulombs required for its reduction at the platinum cathode* the potential of which was controlled within definite limits. In a similar manner* the iodide ion was determined by its oxidation to free iodine at the anode* the potential of which was also controlled»

Lingane (39 $U0 *Hl5i42 *1*3*14**1*6) contributed much to the develop­ ment of this method both through his skill in devising practical con­ trol and measuring instruments and also his ingenuity in developing rapid and convenient schemes of analysis to be used with the instru­ ments* He also was responsible for the introduction of the use of the mercury cathode into coulometry to take advantage of the ease with which polarographic data may be applied to determine optimum deposi­ tion potentials and also to take advantage of the extended cathodic potential range exhibited by that electrode material* Employing the mercury electrode, he devised coulometric methods for the determina­ tion of lead* copper* and bismuth, Lingane and Small (1*6) developed a coulometric method for the de­ termination of the lialides in an acetate buffer solution by deposition

10 on a silver anode* A mixture of iodide and bromide or a mixture of iodide and chloride ions could be separated by control of the anode potential but bromide could not be separated or determined in the presence of the chloride because of eodeposition of the two halides* MacNevin and Baker (l,li7), using the platinum anode, devised a coulometric method for the determination of iron and of arsenic by the oxidation of ferrous ion to ferric ion and the oxidation of arsenic in the plus three cocidation state to arsenic in the plus five oxidation state, respectively, in one molar sulfuric acid* It was also in this work that the slope-intercept method of current summation was devised*

It is this controlled electrode potential method that fulfills the limitation of having only the desired constituent reacting at the electrode* By controlling the electrode potential closely, the de­ sired reaction may be chosen in preference to others. Any reaction which will proceed below the chosen potential and thus will interfere with the desired reaction may be eliminated by allowing this reaction

to proceed to completion at a potential slightly less than the one at which the desired reaction takes place* Afterwards, the desired reac­ tion can be carried out as a coulometric analysis free from interfer­

ence*

Coulometry Applied to Indirect Analysis It was the article by Lingane and Small (Ij-6) reporting the deter­ mination of halides which inspired the principle of indirect analysis using constant potential and coulometric methods. In this article it is stated that the electrogravlmetric determination of chloride and bromide in the presence of one another is impossible because of some

11 codeposition* Ths question immediately raised in this laboratory was, why couldn’t the two halides be deposited together at 1 0 0 percent cur­ rent efficiency5 their combined weights and the amount of current needed for their combined deposition then serving to give an indirect determination which could be solved by simultaneous equations? This particular problem was the first in a series to be considered in this research problem*

Indirect Analysis and Simultaneous Equations In order to visualise the applicability of indirect methods of analysis based upon double measurement, it was necessary to explore the possibilities of extension of this method as well as its limita­ tions* This Tjas best done by the consideration of simultaneous equa­ tions upon which this method is based* A brief summary of this con­ sideration is given in the following paragraphs. In order to solve for any number of unknowns, thnre must be the same number of simultaneous equations as there are unknowns* In addi­ tion, each unknown must appear at least once in one of the equations* In the simplest case of simultaneous equations there are two unknowns so there must be two equations* These may be represented by the gen­ eral equations aX+bY = K (2)

mX + hi = J (3) in which a, b, m, and n are constants which are fixed by the particular experiment performed to arrive at K and J which in turn represent a total reaction of X grams of one unknoim, x, and Y grams of a second unknown, y, together in a mixture* For example, in a coulogravimetric

12 determination of chloride and bromide ions, X would be the weight of chlorine and Y the weight of bromine. K would simply be the sum of X plus Y, the gain in weight of the silver anode, and a and b would be equal to one. The symbols m and n would represent the volume of gas equivalent to a unit weight of chlorine and bromine respectively. J would be the total volume of gas liberated in the lydrogen-csygen cou- loraeter during the deposition of the chloride and bromide ions.

A double coulometric analysis would be a little more complicated. The symbols a and b would then represent the volume of gas equivalent to a vmit weight of x and y respectively in the first electrolytic step, K the volume of gas liberated in the hydrogen-ooygen coulometer during this first step, m and n the volume of gas equivalent to a unit weight of x and y respectively in the second coulometric step, and J the volume of gas liberated in the hydrogen-o:ygen coulomater during the second step. In order to solve for X and Y, several limitations must be im­ posed. These are listed below. 1. Ifa equals b, then m cannot equal n. 2. Ifm equals n, then a cannot equal b.

3. If a equals m, then b cannot equal n. U. Ifb equals n, then a cannot equal m. If ary of these four constants equals zero, then the others must be real numbers in order to be called a simultaneous determination or solution.

13 Propagation of Errors in Simultaneous Equations

When there are err or a in K and J3 which in this case are the ex­ perimental errors in the indirect determination^ than the best answers for X and T are obtained under certain conditions* The conditions for good results in the case of a coulograviraetric analysis are discussed in the following paragraphs. Considerations of other cases, such as a double coulometric analysis, would be similar. The best results are obtained for both X and Y when they occur in approximately the same amounts in a mixture. This relationship is shown in Figure I and Figure IX. Figure I shows the percent relative errors found in X and Y caused by a -O.lj. milligram error in K and no error in J when they are determined experimentally. The relative er­ rors are shown at various percentages of X in the total mixture X i- Y, the total number of equivalent weights of which is held constant at milliequivalents. The curve labeled X represents the errors for values of X and the curve labeled Y represents the errors in values of Y. Figure II is the same except that K is assumed to be accurately determined and J, the coulometric result, is assumed always to b© -0.2 percent in error. These experimental errors used in calculating the values presented graphically were not arbitrarily chosen but were the average errors en­ countered in the determination of chloride and bromide ions discussed in the experimental portion of this dissertation. The value milli­ equivalents was chosen because experimentally this number of milli­ equivalents causes the displacement of nearly 1 0 0 ml. of electrolyte

1U SI PER CENT RELATIVE ERROR i i i i i i i — — — o p o o o p o p — 4^ PO o CD (T> ro oi\34^ODOt\D4^

PO m

o x X zg- X _ + - < 0 0 _

PER CENT RELATIVE ERROR

ro o 00 ro 00

(V) ■ D O m x o

C 3D m

X 00 in the 100 ml* gas coulometer buret* The larger the volume of gas liberated, the smaller the relative error brought about by a small error in reading the buret.

Both Figure X and Figure II show that, with these experimental errors, the major constituent is always fairly accurately determined but the minor constituent, particularly in the extreme cases, is not as accurately determined. An interesting point which is very important in the consideration of errors in these determinations is shown in Figure III. If both a

-0 *1}. milligram error and a -0 *2 percent coulometric error are encoun­ tered the errors in X and Y caused by each of these errors partially compensates for the other and thus a larger range of percent X in to­ tal X + Y gives satisfactory results for both X and Y than if only one of the errors were encountered. This compensation also occurs with any small errors of the same sign. Figure IV and Figure V show another fact which must be considered. The values of m and n must not be too nearly the same. In other words, the equivalent weights of x and y must be separated by a reasonable amount — in the case of the experimental errors assumed here, about $0 percent or more — in order to obtain satisfactory results for ra­ tios of chloride to bromide from $s 1 to ls£. Figure IV shows the er­ rors in X only at the various equivalent weight ratios shown next to each curve. The curve labeled Cl/Br is the X curve of Figure III* Graph V shows the errors in Y only and the curve labeled Cl/Br is the

Y curve of Figure III* By checking these considerations before proceeding with experi-

16 PER CENT RELATIVE ERROR PER CENT RELATIVE ERROR - - 0.4 -0 - - - -1.5 .5 2 2.0 0.8 0.6 0.2 0.5 1.0 1.2 1.0 1.0 - 33/50 E CENT X IN X T N E C PER + ft 2 + X rC!/8r 0 4 / 3 3 20 E CENT N E C PER IUE HTFIGURE FIGURE FIGURE 0 4 0 4 17

12

0 6 33/30 0 8 5 1 100 0 .5 - oc o 20 4 0 6 0 8 0 1 0 0 tr o - CE LU

LU “0 .5 — Cf/Br > i— 3 3 / 1 0 0 _i LU tr 3 3 / 5 0

i u - 2 0 o 3 3 / 4 0 cr 3 3 / 3 0 lu - 2.5 CL FIGURE V

mental work, it is possible to eliminate certain combinations of sub­ stances unlikely to give satisfactory* results.

Applications of Coulometry to Indirect Analysis In Table II various possibilities are listed for the application of coulometry to indirect analyses* Each individual possibility in the table must be weighed in the light of the discussion of the propagation of errors in simultaneous equations* For example, a coulogravimetric determination of Ni++ and

Co+4' would give poor results for, in terms of equations (2) and (3)> a is equal to b and m is very nearly equal to n because their equiva­ lent weights are almost identical* Likewise, a stepwise double cou­ lometric analysis of Sn'i"*++ and Cu'1"®* in the scheme

18 Table II

Method of Obtaining Method of Obtaining Equation (1) Equation (2)

Weighing X and Y together after Measuring the number of coulombs their simultaneous deposition* required to deposit both X and Y. Weighing X alone after its de­ Measuring tlxe number of coulombs position. required to both deposit X and cocidize or reduce Y fran one oxi­ dation state to another.

Measuring the number of coulombs Measuring the number of coulombs required to simultaneously de­ required to deposit X alone after posit both X and Y* stripping the electrode and re­ moving Y by precipitation, com- plexation, etc• Measuring the number of coulombs Measuring the number of coulombs required to oxidize or reduce required to oocidize or reduce X both X and Y from their original alone to still another oxidation oxidation states to others. state or back to Hie original oneo

Measuring the number of coulombs Measuring the number of coulombs required to oxidize or reduce required to oxidize or reduce both X and Y frcm their original both X and Y to still other oxidation states to others* oxidation states.

Measuring the number of coulombs Measuring the number of coulombB required to oxidize or reduce X required to oxidize or reduce from its original oxidation both X and Y after the original state to another. oxidation or reduction of X.

Sn+++*— - Sn++ v Sn° Cu++ ---^ Cu+ Cu° could not be solved even if such a stepwise reduction were possible because a equals m and b equals n.

19 EKEERIMEMTAL

General Apparatus and Materials

Apparatus for the Automatic Control of Electrode Potential Th© apparatus used for controlling the working electrode in this work was essentially the potentiostat devised by Lingane (39)® The wiring diagram and mechanical details are shown in Figure TI and a drawing of the instrument appears in Figure VII just as it appears in the laboratory, set up for use® This is the instrument constructed by Tuthill (65) in connection with the problem of the electrolytic separation of rhodium from iridium and the electrodeposition of rho­ dium® This is also the instrument which was used by Baker (lflU7) in

the determination of iron and arsenic. The symbols in Figure VI and Figure VII designate the major parts of the instrument. These symbols are listed below with an explanation

of each part. A Multiple stage ammeter with 0 to 0.1 and 0.1 to 1.0 ampere scales 3,C Electronic mercury contact relays set for normal open

CL Gouloraeter EC Electrolysis Cell G Galvanometer relay with a sensitivity of +0.02 volt M Reversible synchronous motor, shaft speed 60 rpm

P-j 110-115 volt, 60 cycle A.C. source

Pg 12 volt D.C® source (too 6—volt Willard Radio Batteries)

20 \z

FIGURE W R G U R E TOT 2 volt D.Co source (Eveready air cell)

R^_ 1 0 0 watt, 2 f? ohm radio potentiometer

P.2 1 0 0 watt, 1 0 0 ohm radio potentiometer

2 watt, 1 0 0 ohm radio potentiometer

sl»s 2 *s3>slj. >s£ * s 6 Switches

SB Salt Bridge SCE Saturated calomel electrode V^ Voltmeter, 0 to 25 volt scale

V2 Voltmeter, 0 to 3 volt scale in 0.02 volt divisions W Worm gear unit connecting motor shaft to the radio potentiometer X,Y,Z Points of contact from potentiostat in Figure XI to cell etc. as indicated in Figure I.

The operation of the instrument can be explained simply by refer­ ence to Figure VI and Figure VII. The electrode potential desired is set on the voltmeter V2 by adjusting the variable resistance R^. When this potential exists between the saturated calomel electrode and the working electrode, there is no galvanometer deflection. But when the potential between the electrode and the calomel half cell is different, the circuit is thrown out of balance and current flows. This causes a deflection in the galvanometer G whose pointer is actually a switch which activates the control circuit. The appropriate relay, depending on the direction of the deflection, then operates the reversible motor M which changes the variable resistance until balance is once again restored. A potentiometer may be placed across points Y and Z in Figure VI

23 to calibrate the voltmeter V2 or to observe fluctuations in the poten­ tial of the working electrode during the electrolyses.

In the determination of chloride and bromide ions, the need was for a controlled anode device so the following changes were made in the potantiostats 1. The polarity of the two B.C. sources was reversed. 2. The leads to the ammeter and the voltmeters and Vg were reversed.

3® The galvanometer relay leads were reversed. To follow the changes in current with time, a Model S IjOQOOO Micramax recorder was connected across a calibrated slide wire in series with the coulometer and the electrolysis cell.

Balance and Weights An Ainsworth Type BB magnetically damped, chainomatic balance was used for all of the weighings in this work. The weights, including the rider and chain on the balance, were calibrated by the Richards* substitution method and corrections were calculated on comparison of the 1 0 gram weight with the standard 10 gram weight. The mass of this standard weight was .established by calibration at the National Bureau of Standards. Corrections have been applied to all weights reported in this work.

Reagents All the chemicals used in this investigation were reagent grade chemicals. Double distilled water with less than two and one-half parts per million of impurities was used in the preparation of the various solutions which served as supporting electrolytes and stand— 2h ards in the analyses.

Saturated Calomel Electrode The saturated calomel electrode was made in a 2£0 ml. wide mouth bottle using reagent grade chemicals. It was checked against other reliable calomel electrodes by means of a potentiometer and was found

to have a potential of 0 .2Uf>8 volt vs. the normal hydrogen electrode at 2$ degrees Centigrade. The temperature coefficient of O.QOl* volt per degree Centigrade between 2 £> and 30 degrees was small enough that no thermostatic con­ trol was necessary in this work. Throughout this dissertation, the saturated calomel electrode will be designated by the abbreviation SCE.

Barometer The barometer used throughout this work was a Cenco Model 76890 brass scale barometer. Appropriate corrections were applied before

calculations were made.

ffydrogen-Qxygen Coulometer In this work, the hydrogen-axygen coulometer described by Lin-

gane (I4.O) was used. This coulometer is-shown in Figure VIII. The

symbols designate the major parts of the coulometer and are listed be­ low with descriptions of each part. A Inner compartment filled with electrolyte B Outer comportment filled with tap water to act as a thermostat C Stopcock F I G U R E E m D Thermometer divided in 0.2 degree intervals E Platinum Electrodes^ one square centimeter inarea F Tygon connecting tube G 100 ml. buret into which electrolyte isforced by gas collecting in the inner compartment A H Buret cover

In the first part of this work, the electrolyte used was 0.5 molar potassium sulfate for which the vapor pressure-temperature data in Table III was used to correct the barometric pressure on the col­ lected gas.

Table III (UO)

Temperature in Vapor Pressure degrees Centigrade

2h 22.2 mm Hg 25 23*3 26 2iw7 27 2 6 .2 28 27 *8 29 29 .U 30 31*2

The accuracy of this coulometer was satisfactory far these anal­ yses if it was cleaned and refilled with freshly prepared electrolyte every day or two* However, in forty-eight hours, and in some cases less, the volume of combined gases evolved was les3 than the 0.1739 ml* per coulomb indicated by both Lingane (1+0 ) ami Lehfeldt (3 8)* In seventy-two hours, the results were more than 1 percent and sometimes as much as H percent lew. 27 The author also had difficulty preventing the growth of mold in the electrolyte in spite of the fact that the water was carefully dis­ tilled, the potassium sulfate was sterilised, and various mold inhibi­ tors were added to the freshly prepared electrolyte. At the same time, the inside of the coulometer became quite dirty, indicated by the maiy droplets of electrolyte adhering to the glass walls of the coulometer after drainage. It was this difficulty which led to the investigations with the bydrogen-osygen coulometer found in the next section.

28 Investigations with the Hydrogen-Qnygen Coulometer

Construction The coulometer used for these investigations was the same as the coulometer s h a m in Figure III except far one minor change. This one had three sets of electrodes in place of the one set sham. The sises o . O of these electrodes were 1 cm , 0.5 cmr, both of sheet platinum, and a wire electrode one half millimeter in diameter and one centimeter in length.

Investigation of Potassium Bichromate Solution for the Coulometer Electrolyte After trying several mold inhibitors, methyl salicylate, formalde­ hyde, and chloroform, which did not prevent mold formation in the po­ tassium sulfate electrolyte, it %-jas decided to try the other electro­ lyte, potassium die hr ornate, which Lehfeldt (3 8) found to give the theoretical volume of gas far the passage of current. It was also felt that this electrolyte would be advantageous in over caning mold growth because of its oxidizing power and its poisonous character*

After 0.25 molar potassium dichrornate had been in the coulometer over a week, mold was not visible and the walls of the coulometer were still apparently clean. Nevertheless the coulometer was cleaned and refilled each week. The calculated vapor pressure-temperature data for this electro­ lyte is given in Table IV. These corrections were applied to the baro­ metric pressure before calculations were made*

29 Table XV

Temperature in Vapor pressure in degrees Centigrade mm. of mercury

2 k 22 .2 25 23-5 26 2 5 .0 27 26.5 28 28.1 29 29.8 30 31

In order* -bo check -the precision of the hydrogen-osqygen coulometer with this electrolyte and with the vapor pressure data given, it was placed in series with a silver coulometer which is shown in Figure IX and several runs were made for comparative purposes. The symbols marIdng the important parts of the coulometer are defined below,,

A Silver Anode B Semipermeable membrane C Platinum gauze electrode D Glass ring to support membrane E Glass cup to support membrane and catch anode 'slime1

F Magnetic stirring bar G Magnet attached to shaft of variable speed motor H 180 ml. electrolytic beaker J Electrolyte - 20$ silver nitrate solution.

Investigation of Current Density and Its Effect on Accuracy of the

Coulometer This coulometer was constructed with three sets of electrodes which could be used interchangeably. Several runs were made with each 30 FIGURE IS 31 set of electrodes using the potassium dichromate electrolyte so that the effect of the current density on the accuracy of the coulometer could be determined« This data is presented in Table V.

Results and Conclusions

Table V

Current Density Electrode in ml. gas Meq. Silver Meq. Gases Relative /area deposited Evolved Error in cm per cmr per sec.

2 . 0 0 . 6 2.926 2.922 — 0 2 . 0 0 . 5 2.710 2 . 7 0 7 - 0 . 1 1 2 . 0 0 . 5 3 - 0 8 1 3 . 0 8 0 - 0 . 1 2 0 . 5 0 1 . 0 3 - 0 7 5 3 - 0 6 0 - 0 . 1 5 o . 5 o 1 . 8 2 . 7 1 6 2.711 -0.17 0 . 5 0 1 . 2 2 . 6 1 8 2 . 6 1 5 -0 .1 1 0 . 1 5 5 - 1 2 . 1 5 5 2 . H 2 - 0 . 5 3 0.15 3 - 3 2 . 5 2 8 2 . 5 1 8 -o«!o o . i 5 2 . 8 2 . 6 0 8 2 . 6 0 0 - 0 . 3 2

These data are also presented graphically in Figure X which is the plot of the percent relative error versus the current density in milliliters of gas evolved per square centimeter per second. It shows that, in general, the lower the current density the better the results and the nearer the theoretical amount of gas is evolved per coulomb. This means that the larger the electrodes, for any current, the better the coulometric results. This investigation also led to the conclusion that the use of 0 .2 5 molar potassium dichromate is more satisfactory in the hydrogen-axygen coulometer than the potassium sul­ fate used previously. It also explains why Lehfeldt (38) observed

32 the liberation of 0 .171*1 ml. of combined gases per coulomb in the hydrogen-oxygen coulometer while Lingane (UO) observed only 0.1739 ml. per coulomb9 for the area of the electrodes used by Lehfeldt were 3 cm^ while those of Lingane were 1 cm? in area.

oc -0 7 O cr £ -0.6 W - 0 .5

-0 .4

CURRENT DENSITY IN ML. GAS/SO. CM./SEC.

FIGURE

33 Simultaneous Determination of Chloride and Bromide Ions

Preliminary Considerations

Lingane and Small (U6) found the oxidation potentials of the halide ions depositing as the silver halides on a silver electrode were such that the controlled potential separation of the halides should be possible* The optimum potentials calculated from the stand­ ard potentials fear chloride, bromide, and iodide are 0 *2 5 v 0, 0*12 v., and “0*10 v» vs* SCE, respectively* Since these reactions take place at one hundred percent current efficiency* it enabled them to devise a successful coulometric analysis for each of the three halides when alone in solution* In addition they were able to determine iodide in the presence of bromide and iodide in th© presence of chloride* How­ ever* codeposition of bromide and chloride ions at a potential below which chloride should deposit made it impossible to determine bromide coulcsnetrically when chloride was present* Because these halide deposits are adherent layers which can be washed and dried without difficulty, the halides can also be deter­ mined electrogravimetrically (32). The same limitations apply to these electrogravimetric determinations as apply to the coulometrlc methods of analysis. With the knowledge that both coulcmetric and gravimetric data could be obtained when these two halides are simultaneously deposited, It was felt that an indirect method of analysis for the two ions could be affected with little change in the procedure for the individual ions. This method calls for the simultaneous deposition of the two

3U halides in question* measuring the number of combined equivalents, M, by using the hydrogen-oxygen coulometer and at the same time, employ­ ing the els ctrogravime trie technique to find the total weight, W, of the two halides. These data lend themselves to the formulation of two simultaneous equations*

Increase in weight of anode, W «= Weight of chlorine, X + Weight of bromine, Y (U)

Number of coulombs required for combined deposition, C,

divided by 96*500 coulombs per equivalent => Total equiva­ lents of chloride plus bromide <=

Weight of chlorine, X Weight of bromine, Y 3$M6 79*92

From these two equations an expression for X and for Y may be found by algebraic manipulation* From equation (U),

Y » W - X* (6)

Substituting equation (6) into equation (5),

M = 3^55 * % r.sf* • (7)

Rearranging the terms,

X - (79*92 M - W) (8)

The total number of equivalents, M, is found experimentally from the volume, V, of gas liberated in the hydrogen-oxygen coulometer dis­ cussed previously.

M a vSTp s 0.1739 X 96,5»ob ^ where V’jj, Fg, and Tjjj represent volume of gas, corrected barometric pressurej and temperature3 respectively, at the time of the experiment and Vgipp is the volume of coulometer gas corrected to standard tempera­ ture and pressure.

Because a small but significant amount (A grams) of silver is transferred to the cathode during the electrolysis, both M and ¥ were corrected in the final expression. This expression is

X = 0.797$ [79.92(0.0231}! Z§3s - 9*27 a ) - (W + A) . (11) L •*■£ -1

This equation can be quickly solved when the five variables are determined and are substituted in the appropriate places. 1 is then found by substituting the values of W and X into equation (6).

Special Apparatus The cell for the simultaneous determination of chloride and bro­ mide ions was the simple cell shown in Figure XI. The symbols on the major parts are defined below. A The anode — a platinum gauze (Slcsnin) electrode plated with a heavy coating of approximately five grams of silver deposited frcm a bath containing 11}.7 grams

KCN, 13.8 grams AgWO^, and 11.1 grams K2CQ3 per one

hundred ml. of distil Ted water. This plating bath recommended by Creighton and Koehler (1$) gives a smooth adherent white deposit of silver. B Small platinum gauze (Slcsnin) cathode C Chloride and bromide mixture in supporting electrolyte

which is 0 .1 molar and 0 .1 molar F I G U R E X L 31 D Electrolysis cell - a l£0 ml. jyrex beaker covered with black taps to make it opaque. E Magnetic stirring bar P Magnet .fastened to shaft of variable speed motor G Opaque cover SB Salt bridge to saturated calomelelectrode* This salt

bridge is filled with saturated KNO3 supported in h%

agar.

H Saturated calomel electrode

Only too determinations were made before the halide layer on the silver anode was stripped by back electrolysis in cyanide solution. Otherwise, the resistance was quite high, the current was low, and the time needed for the run was excessive *

The potential of the anode was controlled by the potentiostat using the saturated calomel electrode as the reference electrode, preliminary Investigations Far complete deposition of the chloride ion, an anode potential of 0.25 v* vs. SCE is required. (The sign convention of the electrode potentials used throughout this paper is in accord with the European system in which the calomel electrode is positive with respect to the normal hydrogen electrode.) At potentials only slightly higher, silver dissolves frera the anode. Therefore the first part of the electrolysis was carried out at 0.22 v. and was increased to 0.25 v. vs. SCE for the last 10 to 20 minutes. In spite of the fact that the control was to 40.01 v., sane silver did dissolve fi*csn the anode and was plated on

38 the cathode» No detectable amount of silver could be found in the sup- parting electrolyte nor could ary trace of precipitated silver halide be found in the electrolysis cell in spite of the fact that from 0 -5 to 1*5 mgo were transferredo The increase in weight of the cathode, representing transferred silver, then was added to the final anode weight and the number of equivalents of silver that this represents is subtracted from the coulometric result. These two corrections were made in the data reported belcw.

Experimental Procedure The silver plated anode end platinum cathode were thoroughly rinsed in double distilled water, dried one hour at 250 degrees Centi­ grade, and weighed<, The coulometer electrolyte was saturated with oxygen and hydrogen by shorting out the electrolysis cell and allow­ ing several hundred milliamps to flow through the coulometer. The electrodes were then placed in position in the electrolysis cell* The sample which contained a total of Lw5 milliequivalents of the combined halides was then pipetted into the opaque cell. Enough supporting electrolyte, 0 .2 molar in both sodium acetate and acetic acid, and distilled water were added to result in 100 ml. of solution

0 .1 molar in both sodium acetate and acetic acid. The electrolysis was then begun with the potential of the anode controlled at 0.22 v. vs. SCE* When the current dropped to 10 to 15 milliamperes, the potential was increased to 0 .2 5 v. for the remainder of the electrolysis which was assumed ccmpleto when the current dropped below 0*5 milliampere • The total time required for such an electroly­ sis was about one hour but the operator time was only a small fraction

39 of that. The electrodes were then removed from the cell;, rinsed with dis­ tilled water,, dried for one hour at 2$0 degrees Centigrade, and weighed* The coulometer volume and barometer were then read*

Results and Conclusions Results of simultaneous determinations of various synthetic mix­ tures of chloride and bromide ions are shown in Table VI. These re­ sults are also shown in Figure XII. The circles in the figure repre­

sent the percent relative errors for chloride in the mixture and the squares represent the same for bromide. The curves labeled Cl and Bl­ are the curves which represent the errors in chlorine and bromine cal­

culated assuming a -0 oi| milligram gravimetric error and a -0 .2^ coulo- metric error. The curves roughly foiler.; the experimentally observed results. However, no results outside of the 0.8^ limit were found when

chlorine was present to the extent of 20 to 80^ in the total of chlo­ rine plus bromine. Th±3 means that one of these halogens may be pres­ ent in five times as great an amount as the other and still be deter­ mined with good accuracy. There is a lower limit to the total number of milligrams of com­

bined halogens even if the above limitation is observed. This was established in an attempt to extend this determination of these two halides to semi-micro samples. Dunton (1?) found that the error of about -0.U milligram was encountered even in the smaller samples* Of course, the percent relative error was increased considerably and the

results were not at all good.

1*0 Table VI

ANAI2SIS OF CHLORIDE AMD BROMIDE MIXTURES

Errors Approx. % Cl Run in total Present, mg. Found Absolute Relative No. Cl + Br Cl Br Cl Br Cl Br Cl Br ll*b 88 11*7.3 12.0 11*7.0 11.5 -O.^ng -0.5mg -0.255 -!*.<$

13 83 11*7*3 30.1 11*6.8 29.9 -0.5 —0.2 -0.3 —0.7

h 71 11*7.3 60.2 11*7.1 60.0 -0.2 , —0.2 —0.1 -0.1*

9 1*6 103.0 120.1* 102.9 120.0 —0.1 -0.1* —0.1 -0,3

17 33 73.6 150.5 73.3 150,3 -0.3 -0.2 —0,1* -0.1

25 33 76.7 150.3 77.0 11*9.9 0.3 -0.1* 0,1* -0.3

2U 33 76.7 150.3 76.2 150.6 -0,5 0.3 —0,6 0,2

8 22 58.9 210.7 59.0 209.9 0.1 —0.8 0.2 -0.1*

12 17 1*1*.2 210.8 10**2 210.0 0.0 -0,8 0.0 -0.1*

15b 12 29.5 211.3 29.2 210.8 -0.3 -0.5 -1.0 -0.2

6 12 29.5 301.2 28.1* 301.7 -1,1 0.5 -3.7 0.2

10 9 29.1* 301.0 28.8 301.6 -0,6 0.6 -2,0 0.2 -2,8 0.0 5 9 29.1* 301.0 28.6 301,1 -0.8 0.1 1.0

PER CENT Cl IN Cl + Br

° 3 0 4 0 6 0 8 0 too oc o CL CL LlJ LU B r > - 1.0

i— 2 LU O CL LU CL - 3.0

- 4.0 FIGURE XU

Because of the successful coulometric determination of iodide in the preoence of either chloride or bromide ion and the subsequent cou­ lometric determination of the other halide following removal of the

iodide (U6), it appears that a mixture of all three of these halides could be analyzed* The suggested procedure would involve the coulo­ metric determination of the iodide by controlling the silver anode po­ tential at 0*06 v. vs. SCE followed by the simultaneous determination

U2 of t-he chloride and "bromide given in the above discussion. The work discussed in this section of the dissertation was pub­ lished in part prior to the completion of the entire research prob­ lem (li8)«

U3 Simultaneous Determination of Zina and Cadmium

Preliminary Considerations

After the successful simultaneous determination of the two halide anions, chloride ion and bromide ion, an extension of the coulogravi­ metric method to cover two cations was investigated* These two ions, sine and cadmium, were chosen arbitrarily so that the method of at­ tacking such a problem could be demonstrated*

The equation for the solution of this problem is obtained simi­ larly to the equation for the mixed halides* Increase in weight of cathode, W a weight of sine, X ❖

+ weight of cadmium, Y (12) Number of coulombs required for combined deposition divided

by 9 6 ,^ 0 0 coulombs per equivalent « total equivalents, M, of sine plus cadmium a weight of zinc, X weight of cadmium, Y x 32.69 56.21

From equation (12),

Y » (W - X). (lU) Substituting equation (lit) into equation (13), M . - X . &z£l . (15) 3 2.6? 5 6 .2 1 K '

Rearranging,

X «* (0.056205 M - W). (16)

As in the calculations for chloride and bromide, the value of M is determined by the volume of gas collected in the gas coulometer, V, at a corrected pressure and temperature of Pg and T-g, respectively.

Uit Substituting equation (15) into equation (16)*

X » 1*3899 f 0*056205(0*0211*1— ^) - W I . (17) L rE J This equation is quickly solved when the four variables are de­ termined and are substituted in the appropriate places® It was decided that a mercury cathode would be best for this de­ termination so that preliminary studies could be made with the polaro- graph in the same manner as the work done by Lingane (1*3) «* The use of the mercury cathode enables one to determine the cathode potentials which need to be maintained for the deposition of the metals from the various electrolytes which might be used and at the same time gives the potential above which hydrogen is liberated or impurities are re­ duced so that this value may not be approached during electrolysis *

Special Apparatus First it was necessary to devise a new type mercury cathode to avoid some of the difficulties encountered with the types mentioned in the literature* These difficulties were in weighing the large glass vessel necessary to hold the mercury and a hundred or more mil­ liliters of solution and finding an analytical balance on which to weigh the two hundred or so grams of merculy that was required to cover the bottom of such a cell* At first* a 15 ml. beaker containing 1*0 to 50 grams of mercury was lowered into the electrolysis cell and contact was made by means of a platinum wire sealed in a Pyreoc tube which dipped into the mer­ cury from above5 the Pyrex tube insulating the wire from the electro­ lyte* A graphite contact was substituted for the platinum when it

US was found that the mercury wetted the platinum, and thus the weight of the pool after electrolysis was light by the same amount as the mer­ cury adhering to the platinum* The graphite did not seem to pick up any of the fresh mercury, but careful examination with a magnifying glass showed tiny mercury or amalgam droplets adhering to the graphite after sine had been deposited in the mercury* The final solution was the construction of a mercury cathode as­ sembly shown in Figure XIII* In this cell the 15> ml* beaker, the platinum or tungsten contact, and a Pyrex insulating tube were all made an integral part of the cathode and were all weighed together* No further trouble was encountered with the gravimetric results after this cell was put into use although the lyrex insulating tube proved to be a little fragile and a glass support was provided* This sup­ port, on which the assembly rested, was suspended from the edge of the electrolysis beaker* In addition this cell made it quite easy to out-gas the solution containing the reducible ions with nitrogen, then to lower the cathode and other electrodes into place* It also allowed room for a magnetic stirring bar under the cathode to circulate the electrolyte vigorously and assure complete deposition. The complete cell is shown in Figure XIII. The symbols marking the various parts of the cell are listed below with descriptions of each part* A Anode - a coil of heavy copper wire plated with a heavy

coat of silver according to the directions given in the chloride-bromide experiment. Best results were

U6 FIGURE s i r 47 obtained when the anode was 8 to 10 ram* from the mer­

cury surface. B Mercury cathode - also shown at the left of the figure just as it is washed, dried, and weighed.

C Stirring propeller. The blades are partially immersed in the mercury surface rotating in such a direction as to force the mercury down into the pool. D Platinum or tungsten contact sealed in the pyrex insu­ lating tube. E Copper lead connected to the platinum or tungsten

F Salt bridge - saturate potassium chloride and h % agar contact with the saturated calomel electrode. This salt bridge was placed as close to the mercury sur­ face as possible to eliminate as much tR drop as possible. G Nitrogen tube from tank of nitrogen with which the solu­

tion was outgassed prior to introducing the mercury cathode and from which nitrogen was passed into the electrolyte throughout the entire electrolysis.

H Magnetic stirring bar J Magnet attached to the chuck of a variable speed motor K Saturated calomel electrode.

Preliminary Investigations Various supporting electrolytes were investigated polarographi- cally using the Leeds and Northrup Electro-Chemograph. This was done to determine which of these electrolytes might be suitable for a coulo­ ir metric determinations Those tried were potassium chloride citric acid (pH U to 5), because of a successful electrogravimetric determination of zinc from this medium, sodium tartrate, because of the successful

U3e of tartrate in other coulometric procedures, ammonia, because of the successful deposition of zinc from ammoniacal solution, and alka­ line cyanide, because of the successful deposition of cadmium from this medium* All these were tried in various concentrations ranging from Oof? molar to 2 £ molar* Polarograms were run Tilth OcOOOii molar zinc and O.OOOU molar cad­ mium first separately then together in each of the electrolytes with Q®01^ gelatin added as a maximum suppressor* The results were of two types. The first of these, which was the type of curve found in the

case of citrate, tartrate, and alkaline cyanide is shown in Figure XIV.

f— z UJ 02 02 2D O

0 O CATHODE POTENTIAL vs. SCE

FIGURE X I2 The increase in current at & is the reduction wave for cadmium,, The half wave potential in both the citrate and tartrate supporting elec­ trolytes was about -0.7 v. vs. SCE. In the alkaline cyanide medium^ the wave height was not as great as that for the preceding two and the half wave potential was shifted about 0„$ v. more negative. The in­ crease in current at C is the wave for the reduction of the hydrogen ion occurring at —1.6 v. vs. SCE for the first two electrolytes and -2.0 v. for the last. In each case this reduction of hydrogen appeared before the reduction of zinc and therefore these electrolytes were im­ mediately eliminated as satisfactory electrolytes for the coulometric

determination of zinc and cadmium. The second type wave is that found for both potassium chloride and ammonia supporting electrolytes. This is shown in Figure XV.

\- 2 LU OC OC z> o

CATHODE POTENTIAL v s . SCE

FIGURE W. $0 The sudden increase in current at A is the reduction wave for cadmium with half wave potentials at -0.67 and -0.80 v. vs. SCE in 1.0 molar potassium chloride and 1.0 molar ammonia respectively. The current rise at B is the reduction wave for the zinc with half wave potentials at -1.07 and -1.35 v. vs. SCE in the two electrolytes. In both cases, the zinc reduced well before the hydrogen wave C at about -1.90 v. The 1.0 molar potassium chloride was chosen as the supporting electrolyte for the attempted coulometric analysis of zinc and cadmium. Preliminary electrolyses were made with zinc and cadmium solutions prepared by dissolving the sulfates of the metals in distilled water to give approximately O.Ol; molar solutions of each. The current-time relationship was followed during these prelimi­ nary electrolysis by means of an automatic Micromax point recorder connected across a calibrated slide wire which was in series with the electrolysis cell* The shape of these curves for both zinc and cad­ mium is shown in Figure XVI. The plateau at the beginning of the electrolysis is only indica­ tive of the fact that the initial current was held at 100 milliamperes, or just slightly above , in order to avoid too high a current density on the rather small electrodes. This high current density causes sil­ ver to be stripped freon the anode and deposited at the cathode. What is more important is that for both cadmium and zinc , separately or to­ gether, the curve levelled off at a significant current ranging from

1.0 to 1$ milliamperes. It then remained essentially constant indefi­ nitely as long as the cathode potential was held constant and the rate of stirring of mercury cathode was held constant. 51 OL O TIME

FIGURE X2 L This residual current made it necessary to correct the volume of gas collected during the electrolysis for this residual current which flows throughout the electrolysis.

This correction was found by one of three methods. The current was read on a sensitive ammeter and, with the time required for the electrolysis, the number of coulombs was calculated. Frcm this value, the amount of gas liberated by this current was calculated and was subtracted from the corrected volume of gas before calculation. This method was scrapped because no ammeter was available to read the cur­ rent to more than two significant figures and because the mercury cathode gained weight when the residual current flowed. The second method was more accurate. It involved replacing the mercury cathode after it had been weighed and resuming the electroly­ sis. The amount of gas and the gain in weight of the cathode meas- 52 ured for the blank, run for the same time and under the same conditions as the original electrolysis,, was simply subtracted frcci the total vol­ ume of gas and gain in weight of the cathode before calculations were made. The third method was essentially the same as the second. How­ ever, the blank was run for only a fraction of the time as the origi­ nal electrolysis and the corrections were calculated. This method proved to be the most satisfactory if it were one third to one half the time of the original electrolysis. No correlation could be drawn between the gain in weight of the cathode — from a few tenths of a milligram to two milligrams — and the amount of gas evolved in the coulcmeter although it was suspected that this current was due to silver transfer. This suspicion was confirmed in part by dissolving a small por­ tion of the mercury pool, both for a zinc electrolysis and for the subsequent blank, in nitric acid, evaporating to dryness, and carrying out a qualitative spectr ographic analysis on the IRL one meter grating spectrograph. Results showed a definite trace of silver in both cases and only a slight trace of zinc for the sample fron the blank. No other impurities could be detected® Calculations showed that if the gain in weight were due entirely to silver it still could not account for allthe gas liberated during the electrolysis of the blank. Later studies will show more of the nature of this residual current. Standard solutions of zinc and cadmium were prepared by dissolv­ ing the pure G.P. metals in just enough lei hydrochloric acid to dis­ solve them, boiling gently to expel excess acid, and diluting with 53 distilled water and dilute hydrochloric to a final volume of one liter with a pH of 6 to 7 o The silver anode was prepared by plating a coil of heavy copper wire with silver from the basic cyanide bath recommended by Creighton and Koehler (15>)« The anodes were used twice before stripping in Isl ammonia and replating.

Experimental Procedure About 5>0 grams of cathodic mercury were placed in the mercury cathode cell after the cell was thoroughly cleaned with nitric acid* rinsed, and dried. The mercury was quickly washed with distilled wa­ ter. This was done by rotating the cathode at an angle so the water reached under the edges of the pool, and then as much water as possi­ ble was removed with a suction bulb with a medicine dropper tip* Next it was washed twice with anhydrous acetone, rotating in the same man­ ner as above each time, and as much of the acetone as possible was re­ moved with the suction bulb. The acetone was allowed to evaporate for

20 minutes and then the cathode assembly was weighed. While this evaporation took place, the solution of zinc and cad­ mium was pipetted into the cell and enough supporting electrolyte and distilled water were added to make the resulting solution one molar in potassium chloride. This solution was outgassed with nitrogen for at

least ten minutes before the introduction of the mercury cathode. The nitrogen, taken directly from the tank was allowed to bubble through­ out the electrolysis, both for the actual run and the subsequent blank. At the same time the electrolyte was outgassed the coulometer was

being saturated with hydrogen and oxygen.

5U The mercury cathode was lowered into the cell and the silver anode, salt bridge, and mercury pool stirrer were all lowered into place. The magnetic stirrer was started and the electrolysis was be­ gun with the cathode potential controlled at -l.f?0 v. vs. SCE. The electrolysis was continued until the current dropped to a constant value. With the potential still applied, the silver anode, salt bridge, and stirrer were moved to one side and the mercury cathode assembly was lifted out. The electrolyte above the mercury in the fifteen milliliter beaker was quickly removed with the suction bulb. The pool was immediately washed just as was done prior to the elec­ trolysis and once again it was weighed twenty minutes after the ace­ tone rinse. During this twenty minutes the coulometer and barometer were read. The blank was run as quickly as possible, using the same proce­ dure, applied cathode potential, electrodes, electrolyte, stirring rate, etc., as for the original electrolysis. This run was continued for some fraction of the time required for the original run, usually half or one-third the time. The gain in weight of the mercury cathode and the volume of gas liberated during the blank run were multiplied by the appropriate factor and were subtracted from the original re­ sults. Calculations were then made.

Results and Conclusions Table VII shows the experimental results obtained from the simul­ taneous determination of synthetic mixtures of zinc and cadmium® The table shows that these elements can be determined with fair accuracy —- particularly the one present in the one present in the greater amount. Table V I I

ANAIZSIS OF ZINC AND CAJKCUM MIXTURES

Errors Approx. % Zn Run in t o t a l Presentj mg, Found, mg. No. Absolute R e la tiv e Zn + Cd Zn Cd Zn Cd Zn Cd Zn Cd

55 100 130.8 none 131.2 none O.lt mg - 0 . # -

56 100 130.8 none 131.3 none 0.5 - O.lt -

58 100 130.8 none 131.2 none O.lt 0.3 -

66 8 k 117.7 22.5 117.5 23.0 —0*2 0.5 mg -0.2 2*2%

6U 67 91.6 U5.o 91.3 ii5.B -0.3 0.8 -0.3 1.8

65 51 91.6 89.9 91.1 90.9 -0.5 1.0 -0.5 1.1

59 37 65.lt 112.lt 66.3 111.5 0.9 -0.9 l.lt —0.8

60 37 ■ 65 .it 112.lt 65.6 112.lt 0.2 0.0 0.3 0.0

61 37 65 112.it 65.5 112.8 0.1 O.lt 0.1 o.lt

63 37 65*U 112.lt 61;.8 113.5 —0.6 0.9 -0.9 1.0

67 25 52.3 157.3 53.8 156.2 1.5 -1.1 2.9 -0.7 U6 0 none 221;.8 none 22lt.lt - —O.lt - -0.2

U7 0 none 22lw8 none 22lt.6 — -0.2 - -0.1 The results are not as consistent as the results for the chloride and bromide analysis partly because of the errors introduced in the han­ dling and weighing of the mercury cathode and the running of a blank which is relatively large because of the rather high residual current.

It will be shown later in this work that this residual current cannot be decreased to an insignificant value because of the rather high ap­ plied cathode potential necessary for the reduction of the cations* This simultaneous analysis of two cations could be extended to cover many other pairs of ions* It is most valuable, of course, when reduction potentials of the two ions are such that the two ions cannot be separated by control of the potential but can also be applied, as it was in this case, to any two elements which can be deposited to­ gether at 100 percent current efficiency and can be weighed together accurately. It may be that this technique can be used to speed up present electroanalyses of mixtures which are presently determined separately.

$7 Simultaneous Determination of Tin and Lead

Preliminary Considerations Lead and tin were chosen for the next pair of ions to be simul­ taneously determined although the coulogravimetric technique had al­ ready been applied to the analysis of deposits of the two (20). They were chosen because their reduction potentials indicated a double coulometric analysis might be performed on them by first oxidizing tin at a platinum anode from the stannous to the stannic ion and then reducing both the stannic ion and the unchanged lead (II) ion to the free metals at the mercury cathode. The oxidation step would give an equation in terms of tin alone and the reduction step would give a second equation in terms of both tin and lead. These two equations could then be solved simultaneously. However, as the work progressed, it was found that air oxidation

of the stannous tin and the inability to find simple chemical means to reduce the stannic ion to the stannous ion without either reducing

some of the stannous ion to free tin or introducing another metallic ion which would interfere with the determination made such a determi­ nation impractical if not impossible. The reduction potentials of all three ions, Sn (IV), Sn (II), and Fb (II), are such that one cannot be reduced electrolytically without all of them being reduced. This necessitated a change of procedure in order to effect a practical analysis of the two elements. In introductory quantitative analysis, the tin is often separated from the rest of the metals in an alloy such as brass because stannic oxide precipitates in nitric acid whereas the other metals form the

58 soluble nitrates and go into solution (27). Methods have been devised to dissolve all the metals, including tin, which might occur in such a mixture (6U). With a knowledge of these facts, it was decided that this chemical means could be used to give two samples of one alloy, one sample containing both the lead and tin ions and the other con­ taining only the lead ions. First the tin (IV) and lead ions could be reduced to the metals coulometrically to determine the total num­ ber of milliequivalents of the two. A coulometric reduction of the lead ions in the second solution would give the number of milliequiva­ lents of lead only. Suitable calculations would give an answer for the tin only. The derivation of equations for these calculations is shown below. Let X be the ireight of lead in a sample of a mixture of lead and tin weighing grams and let Y be the weight of tin in the same sam­ ple. Let 2 be the weight of lead in another sample of the same mix­ ture weighing W2 grams.

x/w-l - z/w2 (1 0 ) z » x(w2Af1) (^)

If the first sample of the mixture is electrplyzed so that the stannic ion is reduced to free tin and the lead (II) ion to free lead, the corrected volume of gas collected, in a coulometer would be due to both reductions. The volume due to thelead reduction, Vp^, is given by the expression

Vph p (200QX/Fb)(96.5)(0.1739). (20)

Similarly, the volume due to the reduction of the tin, Vgn, is given by the expression 59 VSn *■ (UOOOt/Sn(UOOOt/Sn) (96.$ ) (0 .17 39 ) • ( 2 1 )

2000(96 <>£)( 0*1739) (22)

Rearranging equation (22) in terms of Y,

Y = VjCSn) (Sn)(X) (23) 1|000(96.5) (0.17 39) " 2Fb

IT the second sample of the mixture is electroiyzed so that the lead (II) ion is reduced to free lead, the corrected volume of gas collected, V2s would be given by the expression

V2 = (2000Z/Pb)(96.5)(0.1739). (2U)

Substituting equation (19) into equation (2h) and rearranging gives an expression for X.

V2 (Pb)(W1) X <= (2*) 2000( 96 .5 ) ( 0 .1739 ) (W2 )

To solve for X, equation (2f>) is substituted into equation (23)- Rearranging this expression gives the equation

UOOO(96.5)(0.1739)

This equation and equation (8 ) are solved when the experimental values of the unknowns are substituted into the proper places.

Special Apparatus A new mercury cathode cell very similar to that used by Lingane (Ul) was constructed. The advantages of such a cell were discussed in detail in the determination of zinc and cadmium. The complete cell is shown in Figure XVII. The symbols marking FIGURE M L the various parts of the cell are listed below -with a description of each part. A Anode - a large Slomin platinum gauze electrodeplated with a heavy coat of silver according to the direc­ tions given in the chloride-brcanide experiment. B Mercury cathode C Cell cover to maintainan atmosphere of nitrogen over the electrolyte D Stirrer with two sets of blades,, one to stir the mercury .

surface and the other the electrolyte. E Salt bridge to the saturated calomel electrode F Nitrogen bubbler tube G Cell - 180 ml. tall form electrolytic beaker H Stopcock J Tygon tubing

K Levelling bulb L Platinum contact to the cathode.

Preliminary Investigations Solutions for the preliminary investigations were made by dis­ solving SnCl2 *2H2 0, SnCljj/^O, and Fb(NC>3)2 in 0 *£ molar hydrochloric acid to give separate solutions 0 .0b molar each* While the possibility of an oxidation of stannous ion to stannic ion was being considered as a means of obtaining one of the two simul­ taneous equations, polarograms of stannous ions in various supporting electrolytes were run. These were done using the Fisher Elecdropode and a stationary micro platinum electrode, varying the anode potential

62 from 0.0 to 1.0 v. vs. SGE. The supporting electrolytes tried were 0.5 M HCl, 1.0 M HCl, 2.0 M HCl, U.O M HCl, 1.0 M HC1 and 1.0 M NaCl,

and 1.0 M HC1 and 3.0 M NaCl* The solutions were made ujj so that they were 0 .0 0 3 M in stannous ion. In the solutions in which the chloride ion was 2.0 molar or more, a smooth oxidation wave was observed, similar to the usual polaro-

graphic reduction wave in appearance. In the solutions below 2.0 M, no oxidation takes place before chlorine evolution. The half wave potentials for these oxidations were 0.1*1 v. in the 1*.0 M HCl, 0.51 v. in 2.0 M HCl, 0.52 v. in 1.0 M HCl and 3«0 M NaCl, and 0.62 v. in 1.0 M HCl and 1.0 M NaCl, all versus the saturated calomel electrode. Similar polarographic studies with lead (II) show no oxidation of lead and no interference of the lead on the oxidation of the stan­ nous ion* The heights of the oxidation current waves seem to bear a quanti­ tative relationship to the concentration of the stannous ions. This was shown by running a series of polarograms using various concentra­ tions of stannous ion in U.O M HCl and measuring the residual current at 0.70 v. vs. SCE. With no attempt to do further worlc along these lines and without calibration of the Elecdropode, the following re­ sults were observed. For 0.003 M stannous ion the wave height was 75 units, far 0.002 M, k9 units, and far 0.001 M, 25 units. This in­ dicates that the stannous ion could be determined polar o graphic ally. However, due to atmospheric oxidation of this ion, precautions would

have to be taken to assure that the unknown and the standards for such a determination tie re not subject to this oxidation, or the unknown and standards should be made up at the same time and in the same manner, 63 running the polarograma as quickly as possible. After it was found that the tin could be oxidised at a platinum anode, coulometric oxidations of the stannous ion were tried. The samples were roade in the same manner as the samples for the polaro- graphic studies except that the supporting electrolyte was li.O M HCl in this case* The samples were electrolysed in a two compartment cell to avoid a cyclic cxidation-reduction. This cell is the one used by Baker (1,U7) in the oxidation of iron and arsenic. The coulometric determination proceeded very much like the deter­ mination of iron and arsenic with the current dropping to practically aero at the end of th9 electrolysis. However, the results were always low. In fact, the longer the samples stood in solution, the lower the results were* This shows further the effect of the air oxidation on the stannous ion. In spite of such precautions as preliminary outgassing of the acid prior to dissolving the stannous chloride, bubbling nitrogen through the solution during the electrolysis, covering the cell to assure an atmosphere of nitrogen over the solution, and electrolysing immediately after preparation of the solution, the best results which were obtained were about 2 percent low. At this point the original plan was abandoned in favor of the two-solution determination discussed in the preliminary considerations, both because of the difficulty in the oxidation and the closeness of the reduction potentials of all three of the involved ions. As in the previous determination the potentials at which the runs are best carried out were determined by preliminary polarographic studies with the Electro-Chemograph. The choice of the supporting 6h electrolytes was influenced by the fact that lead had already been de­ termined in 0 .5 molar potassium chloride (1*3) and also the possibility of the precipitation of basic stannic chlorides unless the solution were acid. Therefore, various concentrations of hydrochloric acid were tried. These ranged from 0.5 to l*.0 molar. The half wave potentials for all three ions, stannous, stannic, and lead (II), were all within 0.02 v. of one another in each electro­ lyte tried. In 0.5 molar HCl the half wave potentials were about

-Q.U8 v. and they were shifted to more negative values with increasing hydrogen ion concentration to -0.58 v. vs. SCE in the 1*.0 molar acid.

In any case, a cathode potential of 0 .8 v. should be satisfactory for the reductions. The standards for the coulometric determinations were prepared from the C.P. metals. For the lead solution, O.Ol* gram atcens of lead were dissolved in just enough Is10 nitric acid to dissolve it and this solution was diluted to one liter. For the tin solution, O.Ol* gram atoms of tin were placed under

0 .5 molar hydrochloric acid and bromine was added drop ty drop until the tin was all dissolved. The solution was heated to drive off the excess bromine and the solution was diluted with 0 .5 molar hydro­ chloric acid to one liter.

Experimental Procedure Enough 2.0 molar hydrochloric acid was placed in the electrolysis cell to make a final volume of 100 ml. irtien the samples are added.

This solution was outgassed with tank nitrogen for 10 minutes while the by dr o ge n- o:cy ge n coulometer was being saturated. Mercury was then

65 introduced into the cell by opening the stopcock connecting the elec­ trolysis cell to the mercury reservoir1 . k cathode potential of -0.8 v. vs. SCE was applied to the mercury pool and the current was allowed to flow until it reached a constant value* The solution of lead and tin to be investigated was introduced into the cell and the electrolysis was started. The current was allowed to pass until a constant residual current was once again observed. This current was determined by measuring the time required to liberate one milliliter of gas in the coulometer. This made it possible to subtract the volume of gas due to the residual current frcm the volume liberated during the electrolysis and thus served as a blank. Some of the supporting electrolyte was removed — an amount equal to the volume of solution to be introduced in the second step of the electrolysis. The coulometer was again read and the lead solution was introduced into the cell. The procedure for this solution was identical to that for the solution of both tin and lead. The residual current was determined in the same manner as before and the correction was again applied. In scsne cases, when the lead concentration was high, the lead precipitated as the chloride in the supporting electrolyte. However, this did not affect the final result since the precipitate dissolved as the electrolysis proceeded. In the current-time curve this was characterized by a horizontal line of constant current which could be attributed to the concentration remaining constant wliile lead chloride dissolved.

66 The two corrected gas volumes were substituted into the equation and the results were calculated.

Results and Conclusions Table VIII shows the results of the simultaneous determinations of various synthetic mixtures of tin and lead. It shows that* like the previous determinations, the major constituent can be determined accurately. The range of good results far both metals is somewhat narrower than that for the chloride and bromide. This is probably due to the errors introduced by the large blank correction. The time required for such electrolyses when tin was present was of the order of six to seven hours which is too long to make the method practical for the determination of tin. However, it does demonstrate how a two-solution procedure may be used to avoid weighing a mercury cathode when chemical means can be used to eliminate one of the two ions and electrolytic separation is impossible. Since copper has been coulometrically determined (31*i*3) and a coulometric method for the analysis of zinc has been developed in this work in connection with the determination of zinc and cadmium, it appears that a complete coulometric analysis of brass could be de­ veloped using these two determinations and the simultaneous analysis of lead and tin which was developed here.

67 Table V I I I

ANAUSIS OF LEM) AND TIN MIXTURES

Errors W A * p X w Run in t o t a l Present^ mg. Found Absolute R e la tiv e i l O o f Fb + Sn Fb Sn Fb Sn Fb Sn Fb Sn

8ii 100 illii.il - kl7.1 - 2.7 mg - 0.7% -

7h 100 207.2 - 207.7 - 0,5 - 0.2 -

85 100 82.9 - 83.7 - 0.8 - 1.0 -

78 90 207.2 23.5 207.7 23.il o.5 -0,1 mg 0.2 - 06%

82 81i 2ii8.7 U7.0 2ii9.2 I16.8 o.5 -0,2 0.2 -O.ii

87 81 207.2 ii7.0 207.7 1*6.9 0,5 -0.1 0.2 —0.2

80 78 165. B ii7.0 166.3 ii6,9 0.5 -0.1 0.3 -0.2

88 7ii 165.8 58.7 166.3 58,7 0.5 0.0 0.3 0.0

77 59 82.9 58.7 83.7 58.5 0.8 -0.2 1.0 -O.ii

90 ill ill .5 58.7 ii2.3 58.7 0.8 0.0 2.0 0.0

92 0 - 23.5 - 23.il - -0.1 - -0.5 76 0 - 58.7 - 58,8 - 0.1 - 0.2 -0.6 91 0 - 117.ii - 116.8 - —0.6 - Simultaneous Determination of Silver and Copper

Preliminary Considerations Silver and copper were chosen for the next simultaneous determina­ tion in order to demonstrate a procedure which is a two step, one solu­ tion reduction technique. Essentially the same technique was originally intended for the determination of lead and tin but experiments showed this to be impossible for tin and lead and so the two solution tech­ nique was used. In ammoniacal solution polarograms of copper solutions show a two step reduction (35). These two steps are the reduction of the

CuCNH^)^* to the Cu(NH^)J ion at about -0 .2 v. vs. SCE and the reduc­ tion of the Cu(NH^ )2 ion to free copper (amalgam) at about -0.5 v. vs.

SCE. The author could find no information concerning the reduction of silver in ammoniacal solution, but this was found experimentally to be about -0 .2 v. vs. SCE also. This made practical the two step schemes first step second step Ag+ ----- v Ag (amalgam) Cu** — --- ► Cu* ------> Cu (amalgam).

If such a stepwise reduction is possible the following derivation serves to provide the two simultaneous equations for the analysis.

Let X => weight of copper in a sample containing both

copper and silver. (27)

Let Y <= weight of silver in the same sample. (28)

In the first step the volume of gas evolved in the coulometer, V-^, due to the reduction of both metals, is given by ihe equation,

6 9 V-l - (1000)(96.5)(0.1739) , (29)

and in the second step, the volume of gas evolved in the coulometer, V2 , due to the reduction of copper only, is given by the equation,

V2 * (I000)(96.5)(0.1739)(x/Cu). (30)

Rearranging equation (30),

v . (Cu)(Vg> (3l) 1000(96,5) (0.173?)

This equation is the one used to solve for the amount of copper

present. Rearranging equation (29),

y - (Vl)(As)______MM (32) 1000(96.5)(0.1739) Cu

Substituting (31) into (32) and rearranging gives an expression for the weight of silver present.

y = ■■ (Vy - V2) (33) 96,500(0.1739)

The experimentally observed gas volumes are corrected to standard temperature and pressure and are substituted into equations (31) and

(33) for results.

Special Apparatus The cell used far these determinations was the cell shown in Fig­ ure XVIII. The divided compartment type was needed to avoid the oxida­ tion of the complex cuprous ions. The important parts marked by sym­

bols in the figure are defined below.

70 F i g u r e x v m i\ A Cathode - large platinum gauze (Slcrain) electrode B Anode — same as cathode C Sintered glass disc dividing the two compartments D Agar plug - b% agar saturated with potassium nitrate E Salt bridge to saturated calomel electrode F Supporting electrolyte G Supporting electrolyte containing hydrazine dihydrochloride H Layer of benzene or toluene to keep oxygen from solution J Nitrogen bubbler tube K Cover L Separate compartments from 180 ml tall form electrolytic beakers.

Preliminary Investigations As in previous work, preliminary polarographic runs were made with the Leeds Northrup Electro-Chemograph. These runs were carried out with 0.002 molar silver nitrate and 0.002 molar copper sulfate in both 1.0 molar arrmoniuni bydrcod.de and 1.0 molar ammonium hydroxide with

1.0 molar ammonium chloride. The results in both supporting electro­ lytes were almost exactly the same. These results are shown in Fig­ ure XIX. Curve I is the polarogram of the supporting electrolyte. Curve II is the polarogram for the 0.002 molar copper ion. It shows the two reduction steps mentioned in the introductory discussion. Curve III shows the reduction of silver which occurs at the same potential as the first copper reduction. Curve IV shows the polarogram for copper and silver together. It shows the 'smearing' together of the first re-

72 m

k i

M IT

CATHODE POTENTIAL vs. SCE

FIGURE IK

4 duction wave for copper and the reduction wave of silver. These pre­ liminary runs show that an electrolytic separation of the two is pos­ sible by controlled cathode potential methods* but a controlled coulo- metric determination of the separate metals is impossible unless an indirect or simultaneous determination is employed. Polarograms of the silver and. copper solutions were also recorded with a micro platinum electrode using the Fisher Elecdropode• The re­ sults for these solutions appeared almost exactly the same as those with the mercury electrode except that the hydrogen wave occurred at a less negative potential and the second copper wave occurred at about -0.65 v. vs. SCE with the platinum electrode. Since the platinum elec­ trodes are easier to handle and to work with than the mercury cathode it was decided to use the platinum electrodes for these determinations.

In preliminary electrolyses* it was found that both stirring and bubbling of nitrogen during electrolysis gave high residual currents 73 which necessitated a rather large blank correction. Therefore no stir­ ring was used during the electrolyses and after the solution was out-

gassed before the electrolysis was begun, the nitrogen bubbler tube was removed from the electrolyte, washing with outgassed 1,0 molar ammonium hydroxide and ammonium chloride. The tube was placed ju3t above the solution and nitrogen was allowed to flair throughout the electrolysis to keep the space above the electrolyte as oxygen free as possible. This lack of stirring did prolong the electrolysis considerably, but by keeping the volume of the electrolyte as low as possible it was possible to get good results in about it or 5 hours. In addition, an agar plug was found necessary in the tube con­ necting the two campartments to avoid oxidation of the cuprous ions at the junction of the two compartments at the sintered glass plug. Other precautions observed to avoid this oxidation were the removal of traces of coy gen from the tank nitrogen by bubbling the nitrogen through alkaline pyrogallol and the covering of the catholyfce with a thin layer of benzene or toluene. In order to prevent the removal of ammonia from the catholyte during the outgassing, the nitrogen was also passed through 1;1 am­

monia solution. The standard solutions were made by dissolving approximately 0*0b. moles of silver nitrate and O.OU moles of copper sulfate in dou­ ble distilled water and diluting to approximately a liter. The sil­ ver solution was standardized gravimetrically by the method outlined

in Hillebrand, Lund ell, Bright, and Hoffman (33) and the copper solu­ tion was standardized electrogravimetrie ally by the method outlined

in the same reference (3h) • 7k The silver deposit was spongy and non-adherent but if there was no stirring the deposit remained on the electrode and the copper plated over it with no ill effects on the results a

Experimental Procedure About 12$ ml* of 1.0 molar ammonium hydroxide and 1.0 molar am­ monium chloride were placed in the anode compartment and enough 2.0 molar ammonium hydroxide and ammonium chloride and distilled water Ttfere placed in the cathode compartment to make 100 ml. of solution 1.0 molar in each when the unknown m s added. About 2 to 3 grams hy­ drazine dihydrochloride was placed in the anolyte to serve as the anodic depolarizer. The catholyte was outgassed five minutes with the oxygen free nitrogen. A cathode potential of -0*8$ v. vs. SCE was applied to re­ move any reducible impurities. The current was stopped and the un­ known sample was pipetted into the cell and the solution was again outgassed. The nitrogen bubbler tube was removed, washing it care­ fully under a stream of outgassed ammonia-aimnonium chloride solution.

The tube was placed just above the solution and the nitrogen was al­ lowed to flow into the air space above the electrolyte throughout the electrolysis. A layer of outgassed benzene or toluene was put on the catholyte and the electrolysis was started at an applied cathode po­ tential of -0.U5 v. vs. SCE. The electrolysis was allowed to proceed until a constant current was obtained. This residual current was de­ termined by measuring the time required for the evolution of a definite volume of gas in the coulomoter* After the residual current was determined and the coulometer vol-

75 ume was read, the cathode potential was raised to -0.85 v. vs. SCE and the above procedure was repeated. The volume of gas liberated and the residual current in the second step were recorded. These values were used to calculate the results.

Results and Conclusions Table IK shows the results of the simultaneous analyses of vari­ ous synthetic mixtures of copper and silver. It shows that the major constituent can be determined accurately. In fact,, all the mixtures but one gave good results for both copper and silver. This may be at­ tributed to the fact that the relative errors for both the first and second steps are about the same as in subsequent calculations these errors almost compensate for each other. The time required far a complete analysis was of the order of U to $ hours, most of which is not operator time because of the use of the automatic potentiostat. The analysis could be speeded up by stirring the solution but this would increase the residual current and the size of the blank correction. In addition, stirring would tear away some of the non-adherent silver deposit.

76 Table DC

ANALYSIS OF SILVER AT® COPPER MIXTURES

Errors Approx. % kg Run No. in total Present* mg. Found>.*& • Absolute Relative Ag + Cu Ag Cu Ag Cu Ag Cu Ag Cu

119 100 202.8 - 203.3 - 0.5 mg n> 0.2$ -

93 100 81.1 - 80.9 - -0.2 —0.2 -

116 100 81.1 - 81.3 - 0.2 0.2

109-10 87 32U.5 50.6 325.9 50.7 l.lt 0.1 mg 0.!t 0.2$

107-8 80 202.8 5o.6 203.5 50.7 0,7 0.1 0.3 0.2

105-6 62 202.8 126.6 203.3 127.0 0.5 O.lt 0.2 0.3

111-12 39 81.1 126.6 82.1 127.0 1.0 o.U 1.2 0.3

117-18 35 81.1 151.9 8l.lt l52.lt 0.3 0.5 O.it 0.3 g 0 HU-15 25 81.1 202,5 8l.6 201.8 0.5 -0 0.6 —0.3

96 0 - 50.6 - 50.5 - -0.1 - —0.2

103 0 - 126.6 - 126,6 - 0.0 0.0

120 0 - 253.1 - 253.9 - 0.8 - 0.3 The Investigation of Residual Currents

Preliminary Considerations After showing that the relatively high residual current necessi­ tated a blank correction in both of the previous determinations, it was decided that a study of the cause of this residual current would be of value, particularly since Lingane (1*3) apparently had little trouble getting the current to drop below one milliampere at the end of the deposition of metals» In these determinations the only way our procedure varied from Lingane*s procedure was the higher cathode potential and higher hydrogen ion and chloride ion concentration. It also had been observed that the rate of stirring influenced the mag­ nitude of the residual current. Therefore the effect of varying the applied cathode potential at a constant stirring rate and the effect of varying the stirring rate at a constant applied cathode potential tjere studied. In addition the effect of changing the pH and chloride ion concentration, at constant stirring rate and constant cathode potential, were studied.

Special Apparatus A divided compartment cell shown in Figure XX was used in these investigations. The symbols marking the various parts of this cell are listed below with a description of each part. A Anode - large platinum (Slomin) gauze electrode plated with a heavy coat of silver B Two compartments made of 180 ml. tall form electrolytic beakers

78 FIGURE 79 C Sintered glass disc separating the two compartments D Mercury cathode E Nitrogen bubbler tube F Stirrer - one set of blades stirring the pool and one stirring the solution G Salt bridge to the saturated calomel electrode H Stopcock J Tygon tubing IC Levelling bulb L Platinum contact to mercury cathode

The stirrer was a propeller type motor driven stirrer with blades constructed as shown in Figure XXI# It was rotated so that the pitch

SIDE FRONT VIEW VIEW 5mm. tubing

FIGURE XXX

of the blades forced the mercury downward into the pool* The upper set of blades forced the solution downward toward the mercury pool*

Experimental Procedure With a freshly plated silver anode, nitrogen tube, and stirrer in place, the 2 molar supporting electrolyte used in the previous de­ termination was introduced into both compartments of the cell - about 80 20 ml* less in the cathode compartment to compensate for the volume to be taken up by the mercury. The solution in both compartments was outgassed with nitrogen for ten minutes with the stirrer going. The stopcock, H, was then opened and by proper adjustment of the levelling bulb, about 20 ml* of mercury was allowed to flow into the cell. The salt bridge to the saturated calomel electrode was lowered to within one millimeter of the mercury pool. The stirrer was adjusted so that the propeller blades were almost completely immersed in the pool and the stirring speed was adjusted to the desired speed by means of a powerstat con­ trolling the stirring motor. The stirring rate was determined by counting the revolutions of the stirrer during several twenty-second periods and then averaging and converting to revolutions per minute.

The nitrogen was bubbled continuously throughout this investigation. A cathode potential of -1.00 v. vs. SCE was applied for ten min­ utes to remove any reducible impurities. With the stirring rate essentially constant at approximately 500 rpm, the residual current at cathode potentials of -0.25 v., -0.50 v., -0.75 v., and -1.00 v., each versus the saturated calomel electrode, was observed. Three readings of the residual current were taken, five minutes apart, and these were averaged. This procedure was repeated at stirring rates of UOO, 300, 250, 200, and 150 revolu­ tions per minute and finally with no stirring. Next, the depth of the stirrer blades in the mercury pool was varied at a constant cathode potential of -0*50 v. and constant stir­ ring rate of 250 rpm. The depth of the stirrer simply changed the de­

81 gree of agitation of the mercury pool. The stirrer was first adjusted so that it barely touched the mercury surface and reading of the milli- ammeter were taken as in the previous determinations. This was re­ peated after lowering the stirrer in one millimeter steps until it was 6 mm. lower than the original position.

The effect of the chloride concentration on the residual current was studied next by holding the rate of stirring and applied cathode potential constant and introducing potassium chloride solutions which ranged from 0.2$ 1,1 to h*0 M in concentration. The solutions were out­ gassed $ minutes before electrolysis and the current was allowed to flow several minutes before the residual current was read. This pro­ cedure was followed at two potentials, -0.$ v. vs. SCE and -1.00 v. vs. SCE. Finally, the effect of the pH was studied in the same manner as the effect of the chloride ion concentration. In this case the elec­ trolytes were made from ammonium hydroxide and hydrochloric acid and the pH was measured with a Beckman pH meter. Otherwise, the procedure was just the same.

Results and Conclusions Figure XXII shows the effect of varying either the applied cathode potential or the rate of stirring while the other is held constant. The voltage by each curve indicates the applied cathode potential ver­ sus the saturated calomel electrode at which the data was taken. This figure shows that the residual current is a function of both the rate of stirring and the applied cathode potential^ the higher the rate of stirring at any applied cathode potential or the higher the

82 0 G 00 ■ Q STIRRING RATE - IN RPM

RESIDUAL CURRENT IN MlLLIAMPERES FIGURE XX.

POOL IN IN MILLIMETERS DEPTH DEPTH OF STIRRER IN

RESIDUAL CURRENT IN M ILLIAMPERES FIGURE XXTTI applied cathode potential, the higher the residual current* However the relationship is not a direct one — doubling one of these variables with the other held constant does not mean that the residual current will be doubled* Figure XXIII shows the effect of the depth of the stirrer in the mercxiry cathode, which is to 3ay the degree of agitation, on the re­ sidual current* This figure shows that the greater the amount of agitation at a constant applied cathode potential and rate of stirring the higher the residual current* However, for any stirrer, there is a certain depth below which the amount of agitation cannot be increased and therefore there is a maximum residual current obtainable. The chloride ion concentration was found to have only a slight effect on the residual current at either of the two potentials. The experiments showed a slow Increase of residual current frcrn 1.0 milli- amperes in 2*0 M potassium chloride to 1.6 milliamperes in 0.2J? M po­ tassium chloride at -0.£0 v. vs. SCE at a constant starring rate of 200 rpn. At -1.00 v. vs. SCE and the same stirring rate these values were increased to 1.7 and 2.1; milliamperes respectively. These values are small in comparison to the changes in residual current brought about by changes in the rate of stirring and the change in applied cathode potential. The pH was also found to have little effect on the residual cur­ rent. In a range of pH values frcm 0.6 to 12.0, the residual current remained essentially constant at about 0.7 milliamperes at -0.£ v. vs. SCE and 1.2 milliamperes at -1.00 v. vs. SCE, both at a stirring rate of 200 rpm. 8U The use of a divided cell seems to eliminate any possibility of cyclic oxidation and reduction of some impurity in the electrolyte although this was considered at first as a possibility. It can be concluded that it would be impossible to duplicate Lin­ gane 1s work, in terms of residual currents, unless an identical cell, stirrer, potential, etc. were employed. It does explain, however, why the residual current obtained in the simultaneous determination of lead and tin was somewhat higher than that gotten by Lingane in the work on the coulometric determination of lead (ii3) because he used a lower applied cathode potential than was used here, although nothing was mentioned as to the rate of stirring. This also makes it possible in future work to choose conditions for which the smallest residual currents are obtained. However, it is not always best to use the conditions which give the minimum residual current, for these conditions may be ones which will prolong the length of electrolysis excessively. It is therefore best to choose scene rea­ sonable compromise which gives a low residual current but does not ex­ cessively prolong the time required for the electrolysis.

85 SUMMARY

The accomplishments of this research may be summarized as followsj 1. Potassium dichromate solution, 0.23> molar ^ was found to be a better electrolyte than 0.5>0 M potassium sulfate in the hydrogen cou- lometer because it prevents mold formation in the coulometer and thus gives good results for a longer time. For ary current, large elec­ trodes were found to give better results than small ones because the current efficiency is more nearly 100^ the lower the current density at the electrodes. 2. A simultaneous determination of mixtures of chloride and bro­ mide ions was devised by combining the coulometric and gravimetric techniques. The halides were deposited together fran a sodium ace­ tate-acetic acid buffer solution at +0.22 v. vs. SCE. The coulombs required for their combined deposition on a silver anode and the gain in weight of that anode were used to give simultaneous equations which, when solved, gave the amounts of each of the two halides pres­ ent. 3. A similar coulogravimetric technique was developed for the analysis of zinc and cadmium mixtures. These two ions were deposited coulometrically on a mercury cathode at -l.$0 v. vs. SCE and the mer­ cury cathode was weighed to give the combined weights of the metals. This analysis is solved in the same manner as that for the halides. U. A two solution, double coulometric analysis was developed for mixtures of lead and tin. Chemical means were used to obtain two solutions, one containing lead only and the other containing both lead

86 and tin. The lead in the first solution was determined coulometrically at -0.85 v« vs* SCE using the mercury cathode and the total number of milliequivalents of lead and tin in the second solution was determined coulometrically at the same potential. Subtraction of the coulometric results gave the results for the tin alone. 5. A one solution double coulometric technique was developed for mixtures of silver and copper by stepwise reduction at a platinum elec­ trode. Ammoniacal solutions of the mixture were coulometrically elec­ trolysed in the cathode compartment of a divided cell at -0.50 v. vs* SCE so that the silver was deposited and the cupric ammine complex was reduced to the cuprous complex. The potential was raised to -0.85 v. vs. SCE and the cuprous ammine complex was coulometrically reduced to copper at the electrode. Subtrcxction of the coulometric results gave a result for the silver alone.

6 . It was found that large residual currents encountered in the electrolysis with the mercury cathode necessitated blank corrections which amounted to as much as five percent of the total result. The nature of these residual currents was investigated. The residual cur­ rent was found to increase with increasing cathode potential and in­ creasing rate of stirring. It was also found that the changes in re­ sidual currents due to changing chloride ion concentration and pH are negligible in comparison to the changes due to changing potential and

stirring rate.

87 BIBLIOGRAPHY

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91 AUTOBIOGRAPHY

I, Richard Donald Mclver, was born in South Haven, Michigan, December 1, 1929® I received my secondary school education in a number of public schools in and near Ft. Wayne, Indiana. Ifiy under­ graduate training was obtained at John Brown University where I held a presidential scholarship for four years. In 1951* I received the degree Bachelor of Arts, summa cum laude. In the Summer quarter of 1951 I entered The Ohio State University where I held the position Graduate Assistant from the Fall quarter of 1951 until the Spring quarter of 1953 with the exception of the Summer quarter of 1952 when I was a Research Assistant. For the academic year 1953-195H I was awarded the Proctor and Gamble Fellowship and in the Summer quarter

of 195U I was awarded the Allied Chemical Company Fellowship which I held until requirements for the degree Doctor of Philosophy were

completed.

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