Silver-Bromide Galvanic Cell. Loys Joseph Nunez Jr Louisiana State University and Agricultural & Mechanical College
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1960 Thermodynamic Studies in Anhydrous Ethanol. The yH drogen - Silver - Silver-Bromide Galvanic Cell. Loys Joseph Nunez Jr Louisiana State University and Agricultural & Mechanical College
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Recommended Citation Nunez, Loys Joseph Jr, "Thermodynamic Studies in Anhydrous Ethanol. The yH drogen - Silver - Silver-Bromide Galvanic Cell." (1960). LSU Historical Dissertations and Theses. 585. https://digitalcommons.lsu.edu/gradschool_disstheses/585
This Dissertation is brought to you for free and open access by the Graduate School at LSU Digital Commons. It has been accepted for inclusion in LSU Historical Dissertations and Theses by an authorized administrator of LSU Digital Commons. For more information, please contact [email protected]. Thermodynamic Studies in Anhydrous Ethanol. 'The Hydrogen - Silver - Silver Bromide Galvanic Cell
A Dissertation
Submitted to the Graduate Faculty of the Louisiana State University and Agricultural and Mechanical College in partial fulfillment of the requirements for the degree o: Doctor of Philosophy in The Department of Chemistry
by Loys Joseph Nunez j M.S., Louisiana State University, 1955 January I960 Acknowledgment
The author wishes to express his appreciation for the advice of Dr. Marion C. Day. He is also grateful for the financial aid received as the recipient of a Lewis Gottlieb Fellowship. TABLE OF CONTENTS Page LIST OF TABLES ...... iv LIST OF FIGURES...... v ABSTRACT...... vi I INTRODUCTION ...... 1 II THE IONIZATION CONSTANT OF HYDROGEN BROMIDE IN ETHA N O L ...... 7 III PURIFICATION OF ETHANOL A, Methods...... 13 B. Experimental Procedure 14 IV EXPERIMENTAL APPARATUS AND PROCEDURE A. The Galvanic Cell...... 17 B. Preparation of the Silver, Silver- Bromide Electrodes ...... 20 C. Reagents Used* . • ...... 21 D. Procedure...... * ...... 22 V THE DETERMINATION OF THE STANDARD POTENTIAL. . . . 25 VI IONIZATION CONSTANTS OF WEAK ACIDS IN ETHANOL. . . 32 VII DISCUSSION OF RESULTS...... 35 BIBLIOGRAPHY ...... 78 VITA ...... 81 LIST OF TABLES TABLE PAGE I Conductivity Data of H. Goldschmidt and P. Dahll ...... 38 II Functions calculated from the data of H. Goldschmidt and P. Dahll .... 39 III Physical Properties of Anhydrous Ethanol...... 40 Observed electromotive force of the cell IV-XV Runs 1-12...... 41-52
(Eq - E*) and (E* - E*)* as a function of the concentration for the value of the parameter a : XVI ft = 4 . 0 ...... 53 XVII X = 5.0 55 XVIII X Z 6 . 5 ...... 57 XIX X - 8 . 0 ...... 59 XX-XXVII Observed electromotive force of the cell. Runs A-l - A-G...... 61-68 XXVIII-XXXV Observed electromotive force of the cell. Runs B-l-B-8...... 69-76 XXXVI The function -0.0591 log as depen dent on the Ionic Strength,UJ, for Acetic Acid...... 77 LIST OF FIGURES FIGURE PAGE 1. The Data of Goldschmidt. Equivalent Conductivity vs. the Square Root of the Concentration ...... & 2. The Data of Goldschmidt F{Z)//i vs. Cf2A/F(Z) ...... 11 3. Schematic Diagram of Potentiometer-Cell Circuitry •••••••...... 18 4. The Galvanic C e l l ...... 19 5. Corrected Eobs> vs. 4- log Molality...... 24 6. Corrected Eobs^ vs. — log Normality...... 26
7. Concentration vs. — {Eq — E^) ...... 29
S. Concentration vs. — (Eq — E^)*...... 31
9. The Ionic Strength,U) , vs. 0.0591 log k ! for Acetic Acid ...... * 3 3
10. The Ionic Strength, IV, vs. 0.0591 log for Benzoic A c i d ...... 34
v ABSTRACT
The standard electrode potential of the silver- silver bromide electrode was determined in anhydrous ethanol for the first time. The value was found to be — 0.1940 +1 0.0005 volts on the molar concentration scale and — 0.1S60 — 0.0005 volts on the molal concentration scale, the standard potential being expressed as an oxidation potential. The silver-silver bromide standard electrode potential in ethanol is more negative than the silver-silver chloride standard electrode potential in ethanol, when both are determined by the method used by H. Taniguchi and G. Janz (23). The difference in these two standard potentials in ethanol is 0.1002 volts as contrasted to 0.1493 volts in water. Potentiometric measurements were made at a temperature of 25°C on a galvanic cell without liquid junction consisting of a silver-silver bromide electrode and a platinum black hydrogen electrode. The silver-silver bromide electrodes were pre pared by a new modification of existing procedures, which eliminated the need for platinum wire in the electrodes. The conductivity data of H. Goldschmidt and P. Dahll (10) were independently evaluated in order to obtain an approximation for the molar thermodynamic ionization constant of hydrogen bromide in ethanol. This was found to be 0.01S7. vi The thermodynamic molal ionization constants of acetic and benzoic acids in anhydrous ethanol were evaluated. The ionization constants of benzoic and acetic acids in ethanol were found to be 1.01x10“^ and 2*05x10"’'^ respectively. The value of the ionization' constant of benzoic acid in ethanol is in substantial agreement with that found by other workers using conductivity methods. No value for the ionization con stant of acetic acid in ethanol is available in the literature for comparison.
vii I
INTRODUCTION
The measurement of the electromotive force of galvanic cells without liquid junction has been frequently- reported in the litej^ature of the recent past for aqueous systems, particularly in connection with the determination of the ionization constant of dissolved weak acids (20). Until recently non-aqueous measurements have been less frequently reported and are more difficult to interpret due to the lack of an adequate theoretical approach and a clear understanding of the effects of solvent impurities. A recent study of a non-aqueous galvanic cell without liquid junction was made by H. Taniguchi and G. Janz (23) using anhydrous ethanol as the solvent with reversible hydrogen and silver-silver chloride electrodes. These authors used the best known procedures to ensure the purity of ethanol, although small quantities of benzene may well have been present in their solvent, and they were the first ones to consider the degree of dissociation of hydrogen chloride in ethanol in evaluating the standard electrode potential of their system. They accepted the value of the ionization constant of hydrogen chloride in ethanol as calculated by the method of R. Fuoss and G. Kraus (8), (9) from the conductivity data of I. Bezman and F. Verhoek (2). By using the extended terms of the Debye- 1 Huckel Theory, as tabulated by T. Gronwell, V. La Mer and K„ Sandved (11), Taniguchi and Janz calculated the activity coefficients of the ions, evaluated the degree of disso ciation of hydrogen chloride in ethanol, and computed the standard electrode potential, E0 , for the cell by a method which is illustrated by H. Earned and B. Owen (14). Taniguchi and Janz obtained an EQ value of -0.09383 in terms of the molar concentration scale. This value differs some what from previously reported values which varied from -0.0864 volts (20) to 0.00977 volts (21). C. Le Bas (18) has shown that some of the dis crepancies in E0 values may be attributed to the prescence of small quantities of water in the ethanol. Other dis crepancies may be due to methods of graphical extrapolation which place emphasis on points in regions,of high dilution where experimental error tends to be high. Attention was drawn to the bromide system by the fact that the salts, sodium bromide and potassium bromide, are considerably more soluble in anhydrous ethanol than the corresponding chlorides as shown in Table I. It was felt that practical applications of anhydrous ethanol as a solvent may favor the bromide system rather than the chloride system due to solubility considerations, but it was found that t-hese advantages are off-set by the instability found in the regions of relatively high concentration of the bromide system in ethanol. The aqueous cell: Ag; AgBr (s), HBr (m), H^; Pt has been investigated by A. Keston (17) and by H. Harned, A. Keston, and J. Donelson (12), (13)* The present inves— tigation is concerned with a cell analogous to the one shown above, using anhydrous ethanol as the solvent. The cell reaction is: AgBr 4- l/2H£=* Ag + HBr Considerable attention has been paid to the purification of ethanol; the small quantities of benzene which may have been present in the solvent of Taniguchi and Janz have been carefully removed, even though their effect on the potential of the cell is not accurately known.
It was felt that the theoretical approach of Taniguchi and Janz was the most desirable even though the value of E0 will be dependent on the value of K, the ionization constant of hydrogen bromide in ethanol, which must be obtained from conductivity data. The conductivity data of H. Goldschmidt and P. Dahll (10), presented in the literature of 1925* allows an approximation of the ionization constant of hydrogen bro mide in ethanol to be obtained by a method unknown to those workers at the time of publication. The method of determining the standard electrode potential of the cell, the same as that used by Taniguchi
and Janz, essentially reduces to plotting the function 4 , t {E0 — E^J versus the concentration of hydrogen bromide for a given value of the parameter a where (E^ — Ew) is defined as:
(Eq - Ej = Eobs<+ 0.1163 logoCf^C (1-1) where Eobs z Corrected observed E.M.F. of cell Ot, Z degree of dissociation HBr in ethanol fKz activity coefficient, dependent on Ka - m^ m£_ . Yh * T a” mHA YHA where m = molality V ~ molal activity coefficient The expression for the electromotive force of the cell may be written as follows: E = Ec - ET In mH+V H+mBr_ YBr- Combining the two equations yields: E— E0 + R T In mHAmBr- Z- -RT In Yh+ ^Br" ^HA F------F------mA- Yh+ Ya~ — RT In k ! (1-4) F A The left hand side of the equation may be readily evaluated. As the dilution approaches infinite dilution, the activity coefficients approach unity and the first term on the right becomes zero. According to D. A, Mac Innes (20) the left hand side of the equation has been found to be a linear function of the ionic strength. There fore, when the left hand side of the equation is plotted versus the ionic strength, the intercept is proportional to the thermodynamic dissociation constant, KA. In dilute aqueous solution it is frequently possible to neglect the extended terms of the Debye - Hiickel Theory and use only the Debye approximation for calculating the activity coefficient. This equation is of the general form: — l o g f = a ^ n r - 1 + Bai TfT where C = Concentration f = activity coefficient A = constant B = constant a^ = mean distance of closest approach For very dilute aqueous solutions the equation reduces to: — log f = A VTT Due to the low dielectric constant of alcohol, it is not possible to neglect the extended terms of the Debye- Huckel Theory when working with this solvent in concentrations comparable to dilute aqueous systems. Therefore, all cal culations of the activity coefficient presented in this investigation have used two extended terms, that is, the terms of the 5th power have been included in the calcula tions. Once the E0 value has been obtained the activity coefficient may be calculated from the experimental data at any concentration by means of Eq. (I—1)• In lieu of a formal section on review of the literature, the pertinent references appear in this section and the body of the dissertation. II THE IONIZATION CONSTANT OF HYDROGEN BROMIDE IN ETHANOL The conductivity data of H. Goldschmidt and P. Dahll (10) have been used to determine the ionization constant, K, of HBr in ethanol by the method of R. Fuoss and C. Kraus (8), (9)* These data are presented in Table I. The value of conductivity at infinite dilution reported by Goldschmidt was not used in the calculations. The method allows an approximate value to be used, based on a simple extrapolation to zero concentration. A plot of the con ductivity data is shown in Fig.l. Essentially, the method is based on a solution by successive approximation of the Onsager Equation. This equation is believed to hold exactly up to concentrations where ionic interactions of higher order than pairwise become appreciable. According to R. Fuoss and C. Kraus the equation may be written: A =Y(A> — «3cr) (ii-D where: A z equivalent conductance Aor equivalent conductance at infinite dilution C z concentration, m/1. Y = degree of dissociation. and oCamay be defined as: X - S. 18x105A0 4 . 82 (DT) 3/2 tl (DT) 1/2 where: D I dielectric constant T z absolute temperature = viscosity of the solvent 7 80 THE DATA OF GOLDSCHMIDT. EQUIVALENT CONDUCTIVITY vs. THE SQUARE ROOT OF THE CONCENTRATION. 70 60 50 40 0 4 8 12 16 20 24 28 32 YcT x io2 FIGURE 1 ca We may define the ionization constant, K: K : C V 2 f 2 -L± (II-2) where & - and activity coefficient of the positive ion. In order to facilitate the solution of the various equations R. Fuoss has defined a function Z such that: Z 5 O C A Z 3^2 NfCA (II—3) and F(Z) = 1 - Z(l-Z(l-Z(l-...)-£)-4)-l (II-4) Then the solution for ^ , that is, the best approximation of Y , is given immediately by: Y = —A (n-5) A 0 (F - log f : b J jjT where: B = 0.4343 e2 N e2 \l/2 2DkT \1000 Dkt / /8-IT N e2 \ 1/2 \1COO Dkt ) and [H*] = [fir-] where e Z electronic charge N Z Avogadrofs Number k Z- Boltzmann Constant D z Dielectric constant of solvent T = absolute temperature aj_ = ionic radius 10 For our purposes it was possible to neglect the term in the denominator, Eq.(II-6) reducing to: - iog10 = 2B'ire_ as an approximation. For ethanol: 2 B : $.90 Substituting Eq.(II-5) into Eq.(II-2) the following expression is obtained: k ( l — A ^ = cf2 / A 2 V A 0F(Z) ) \ A § F(z)2 Upon rearrangement this reduces to: F (Z) = 1______f Cf2A S|. 1 A lKA02) { F (Z) A 0 Thus if F(Z )/A is plotted versus Gf^A/F{Z) a straight line of slope ( V kA o ^ and intercept V A 0 should be obtained. Such a plot is shown for the data of Goldschmidt in Fig.2. A 0 was found to be S3.47 and the value of the molar ionization constant, was calculated to be 0.0187. Numerical values of the functions plotted in Fig.2 are tabulated in Table II. F (Z) values were taken as those conveniently tabulated by Fuoss (8) for any given value of Z. The value of the molar ionization constant K^gr = 0.0187 in ethanol may be compared to the molar ionization constant « 0*0113 in ethanol, both calculated by the same method. Since both HBr and HC1 have long been assumed to be strong acids in ethanol, it is not surprising that the ionization constants are of the same order of magnitude. Qualitatively, at least, it would be expected that HBr would THE DATA OF GOLDSCHMIDT 1300 LO o 1200 9 10 11 1412 c A f 2 xioo F FIGURE 2 F X l O 5 1300 1200 H AAO GOLDSCHMIDT OF DATA THE F(z) IUE 2 FIGURE F 11 12 12 be a stronger acid than HC1 due to the greater size of the bromide ion, which would allow greater dissociation of the proton. Ill PURIFICATION OF ETHANOL A. Methods Traces of benzene present in commercially available U.S.I. absolute ethyl alcohol, U.S.P. grade, are of sufficient quantity to prevent the use of this grade of alcohol as a glass in phosphorescence studies and as a solvent where the highest degree of purity is required. Spectroscopic data of the author indicate that benzene may be present in quantities higher than 0.001%. Anhydrous alcohol may be prepared from 95% ethanol by a variety of methods some of which depend on depressing the vapor pressure of water (6). Usually the starting material is sufficiently free from benzene so that this impurity does not show up in the final product, but the removal of other impurities often necessitates con siderable effort. Calcium oxide treatment (22) and various modifications of this method have been used for purifying 95% alcohol. J. Woolcock and H. Hartley {2k) as well as P. Danner and J. Hildebrand (5) used ethanol treated in this manner for potentiometric studies. Other workers have followed calcium oxide treatment by the Bjerrum method (19), which involves the addition of magnesium. By combining known procedures a simplified method has been obtained, using U.S.I. absolute ethyl alcohol as a 13 Ik starting material, which produces ethanol of a high degree of purity that may be used as a solvent in potentiometric studies. The physical properties of this alcohol compare favorably with data reported in the literature. Some of the results obtained are presented in Table III. Taniguchi and Janz (23) report a density of 0.73056 at 25°C for pure ethanol, although it is probable that this value is a misprint since a density of 0.73506 at 25°C is reported for pure ethanol in the International Critical Tables (16). A. Clow and G. fearson (3), L. Mukherjee (21), and L. Harris (15) determined the purity of alcohol by means of spectroscopic data, observing transmission of ultraviolet light in the region 210 — 250 m)i. Transmission in both the infrared and ultraviolet regions, as measured by the Beckman DK Spectrophotometer, have been regarded as the most important criteria of purity of alcohol prepared by the method described below: B. Experimental Procedure Modifying the method of Bjerrum, magnesium alcoholate was prepared by adding 25g of magnesium metal and 5g< of iodine to 150 ml of dry ethanol (pre-purified and containing less than 1% water). The mixture was allowed to react for 3-10 hours ar room temperature until no liquid remained in the flask. 15 Two per cent water by volume was added to U.S.I. ethyl alcohol, U.S.P. grade, and the mixture was fractionally distilled. When a fractionating column of approximately ten theoretical plates was used, the first 25%> of distillate contained most of the benzene as the azeotropic mixture of the following composition: water, 7.5%; ethanol, 18,5%; benzene, 74.0$. The distillation was stopped at this point and the solid magnesium alcoholate mixture introduced into 900 ml of the distillation residue. This was refluxed for approximately one-half hour until free of water, followed immediately by fractional distillation. The first and last 10%o cuts of the distillate were discarded and the inter mediate fraction of pure ethanol was retained. Absorption of water from the atmosphere was prevented during refluxing and distillation by the use of silica gel drying towers. The distillation which removes benzene as well as the final distillation of pure ethanol was monitored by means of the Beckman Model DK Spectrophotometer. Using water as a blank less than 0.000B$ of benzene in ethanol may be easily detected in the ultraviolet region with a cell path length of 1 cm. An impurity in the first and last cuts of the final distillation also appears in the region 335— 200 mj*. The same instrument will also detect small quantities of water in ethanol in the wavelength region of 1950 mjju A cell path length of 0.5 cm. may be used with 16 anhydrous ethanol as a blank. It was found that less than 0.02 wtof water in ethanol could be easily detected in this manner. It is important that once the magnesium alcoholate is introduced into the alcohol-water mixture, the resulting distillation be conducted without interruption and as rapidly as possible; otherwise undesirable side reactions will con siderably reduce the yield of pure ethanol. Best results were obtained using magnesium turnings of 99.6$ purity, marketed for use in the Grignard Reaction by the J. T. Baker Chemical Co., Phillipsburg, N. J. IV EXPERIMENTAL APPARATUS AND PROCEDURE A. The Galvanic Cell A schematic diagram of the cell circuitry is presented in Fig.3* A. Leeds and Northrup Potentiometer was used. The cell itself, illustrated in Fig.4* consisted of a four neck flask in which a hydrogen electrode of the Hildebrand type occupied the center position. Silver- silver bromide electrodes were in adjacent positions. Thus the same hydrogen electrode served as "reference" when potentiometric measurements were made alternately on both silver-silver bromide electrodes. The platinum black surface of the hydrogen electrode was immersed approximately one inch below the surface of the liquid, the bubbles emerging from the electrode at about the same height. The rate of flow of hydrogen was 1-2 bubbles/second. The platinum black was deposited on the surface of the electrode from a chloroplatinic acid solution con taining a little lead acetate. A platinum electrode was used as the anode and a current of 200-400 ma. was passed for 3 minutes. The electrode was replatinized every 2 or 3 runs or whenever the system showed signs of instability. Prepurified hydrogen, sold by the Matheson Co., Inc., passed first through a hydrogen catalytic purifier marketed by Engelhard Industries, Inc., Newark, N. J. 17 POTENTIOMETER CIRCUIT FIGURE 3 19 TO HYDROGEN HYDROGEN ELECTRODE Ag,AgBr ELECTRODE THE CELL FIGURE 4 20 It then went through a silica gel column and into a bubbling tower filled with a liquid of identical composition as that present in the celle From the bubbling tower the hydrogen entered the cell and passed over the platinum surface of the electrode. Exit of the gas into the atmosphere was through a capillary tube. The cell was immersed in a Sargent constant temperature water bath, maintained at a temperature of 25,25°C i 0,05°. As room temperature exceeded 25°C, a Sargent water bath cooler was employed which circulated cool water through coils immersed in the constant tempera ture bath, B, Preparation of the Silver-Silver Bromide Electrodes Silver-silver bromide electrodes of the thermal electrolytic type were prepared initially by use of a platinum wire as a holder, A mass of porous silver was deposited on a helix of platinum wire by the thermal decomposition of a paste of silver oxide. The silver was then converted to silver bromide by electrolysis. Elec— • trodes of this type showed marked instability, presumably due to cracks in the silver surface exposing areas of platinum which may function as small hydrogen electrodes. The following adaptation was therefore chosen. Silver wire of 1 ram. diameter was sealed in one end of a 21 glass tube of slightly larger diameter than the silver wire* Approximately one centimeter of silver wire was allowed to protrude from the glass tube on both ends, one end serving as a means of electrical contact to the external circuit and the other end serving as the surface of the electrode. One end was heated in a flame until molten, a spherical droplet being formed. After cooling, a paste of very pure silver-oxide (1) was applied to the silver surface and this was heated again until a spongy surface of very pure metallic silver had covered the surface. The electrode was then anodized in dilute KBr solution as usual, until a spongy layer of AgBr had covered the surface of the elec trode. All silver-silver bromide electrodes used in this investigation were prepared in this manner. C. Le Bas (IB) has shown that these electrodes gave results identical to those observed with electrodes prepared by the thermal electrolytic method using the platinum wire spiral. C. Reagents Used Anhydrous hydrogen bromide as obtained from the Matheson Co., Inc. was used directly from the cylinder. Anhydrous acetic acid was prepared by distilling a mixture containing 1% acetic anhydride and glacial acetic acid of 99,1% purity. A middle fraction of 20-40% was collected and used as anhydrous acetic acid. 22 Benzoic acid, sodium bromide and sodium benzoate were all of U.S.P. or Analytical Reagent grade and were recrystallized from anhydrous ethanol. Sodium acetate •3H2O was heated until molten in order to obtain the anhydrous salt. All operations which would expose anhydrous materials to the atmosphere were conducted in a nitrogen atmosphere dry box desiccated with phosphorous pentoxide. The fractional distillation apparatus used was of approximately 10 theoretical plates. All connections were glass to glass with no stopcock grease. Drying towers of silica gel were placed on all openings to the atmosphere. D. Procedure After completion of a run, aliquots of the ethanol-hydrogen bromide solution were withdrawn by pipette from the cell and titrated with standardized sodium hydroxide solution using phenolphthalein as an indicator. The normality was then converted to molality by the appropriate conversion factor. The concentrations of salts and weak acids were determined directly on a weight basis. Both salts and alcohol were weighed on an analytical balance in a desiccated atmosphere. The salts and weak acids were added to the anhydrous ethanol in a dry box and were allowed to remain there, stoppered, for several hours until all solid particles c 23 went into solution. No electromotive force measurements were recorded until the system had reached an apparent equilibrium, which usually required several hours after initiation of hydrogen bubbling. The criterion of equilibrium was that no appre ciable drift in readings in one direction could be observed during a period of one hour, although most measurements recorded in the tables extended over a period of time considerably greater than one hour. It was found impossible to obtain reliable readings in regions of concentration greater than that shown in Fig.5. V THE DETERMINATION OF THE STANDARD POTENTIAL The observed potential of the cell, corrected to one atmosphere pressure of hydrogen, is presented in Fig.6 as a function of the logarithm of the concentration of hydrogen bromide expressed as m/l. The raw data from which this curve was obtained are presented in Tables IV-XV. The empirical equations of the line within the concentration ranges studied may be expressed as: EE — 0.1021 log Molarity — 0.1327 (V-l) E = — 0.1021 log Molality — 0.1221 , The above equations were obtained by the method of least squares, the standard deviation being: 0.00165. Considerable deviation from a straight line is to be expected in the more concentrated regions but the insta bility of the system did not permit an accurate evaluation of points in this area. The expression for the density of HBr- ethanol mixtures as a function of the normality is as follows: D = 0.07993 (N)-f0.7855 (V-2) where D z density, g/cc N = normality The above expression, which was obtained from a least squares treatment of experimental density points, holds in the range of concentration from 0.0 to 0.5 normal. 25 ta obs . 40 120 -80 -40 80 2.8 2.4 CORRECTED E0bs. vs. -log MOLALITY -log vs. E0bs. CORRECTED FIGURE 5 FIGURE -log m m -log 2.0 1.6 1.2 -e- JO CO u o 80 “ X 10 -40 120 160 40 2.8 2.6 2.4 CORRECTED Eobs^ vs. -log NORMALITY -log vs. Eobs^ CORRECTED 2.2 -log NORMALITY -log 2.0 IUE 6 FIGURE 1.8 1.6 1.4 1.2 o ro 1.0 Molalities were calculated by means of this expression. The method of Taniguchi and Janz for determining the standard potential of the cell involves plotting the function — (Eq - E*) vs. the concentration of HBr and extrapolating the line to zero concentration. For any given value of a, the distance of closest approach of the Debye-Huckel Theory, — (E^ - E*) is the summation of the terms E, the corrected observed potential of the cell, and 0.1183 logocfK C, where these terms are defined in Eq. (1-1). The quantities oc and fee must be obtained by successive approximations, the activity coefficient then becoming a function of the degree of dissociation and ionization constant, K of hydrogen bromide in ethanol. The activity coefficient is first found by means of the following equation "which includes the extended terms of T. Gronwall, V. La Mer and K. Sandved (11): Inf =: / -e2Z2 \ / 1 U x \ I (V-3) J [ 2 j [ 1+x/ ' e2Z2 'N 2m+l (l Y(x) — 2m V (x) kTDa I [2 2m+l l2m+l m=l ' where: x = 1^ a and: \[2 z 8TTNe2Z2C 103kTD where: C z concaitration m/l. D = dielectric constant of ethanol, 24-3 T r absolute temperature N Z Avogadrofs Number k = Boltzmann constant e = electronic charge unit, e.s.u. a = distance of closest approach Z = valence of ions £ = activity coefficient 28 and x< x) and Y W are complicated functions of x tabulated by Gronwall, LaMer, and Sandved. Once the activity coefficient has been evaluated for a given value of the parameter a, the degree of disso ciation, DC - 1 •_K \l/2\ (V-4) 2 Cf± * )1/2] where: K = thermodynamic ionization constant. Once DC has been calculated, the activity coefficient, f*, may be recalculated on the basis ofotC to give a better approximation of the function f*. These data are presented in Tables XVI —-XIX for values of a 5, 6.5, and 8.0 angstrom units. Plots of these data are shown in Fig.7. The theoretical line of zero slope is exhibited by the function only in the region of concentration less than 3 m/1. Therefore this plot is not as unambiguous as that presented by Taniguchi and Janz, as their function exhibited a straight line of zero slope over a larger concentration range. This is perhaps because the hydrogen bromide system is less ideal since the bromide ion is considerable larger than the chloride ion. Since it was observed that OC, the degree of dissociation approaches unity as the concentration approaches zero, a new function was defined by the author such that: (E* - E «)' = E-f 0.1183 log foe C. This makes DC--1, but f* is still calculated as a function of 20.2 cs o 3—I X w 19.8 ► o w 5.0A 4.0A 19.4 . o 6.5A 8.0A 19.0 1 8 12 16 20 CONCENTRATION, MOLES/l XlO^ FIGURE 7 30 the approximate values of ot given by the ionization constant expression, Eq. (V-4)« This function is plotted at values of a equal to 4, 5* and 6.5 angstroms in Fig.8. It can be seen that the curves are smooth and extrapolate to the same value of E0 as shown in Fig.7. The value of E0 has been taken as — 0.1940 on the molar concentration scale and — 0.1816 on the molal concentration scale. The reliability may be estimated to be + 0.0005. Once the E0 value has been determined, the mean molal activity coefficients may readily be calculated from the following equation: log Xt " Eo — E molality (V-5) 0.1183 where I mean molal activity coefficient Eq = standard electrode potential E z corrected observed E.M.F. 20 - E*) q -(E 10 FIGURE 8 CONCENTRATION vs. CONCENTRATION, MOLES/l.X103 5.0A 5A 6. 4.0A __ 19.00 18.80 20.00 zoi X 4(*a - o3)- VI IONIZATION CONSTANTS OF WEAK ACIDS IN ETHANOL Once the standard potential of the cell has been obtained, the ionization constant of a weak acid dissolved in ethanol may be determined by potentiometric measurements. These data are presented in Tables XI - XXXVI for benzoic and acetic acids. The method utilizes Eq. (1-4). When the left hand side of the equation, the function 0.0591 log K^, is plotted versus the ionic strength,Ui, at regions of high dilution the activity coefficients reduce to unity and the intercept equals — ^ / F lnK^. The thermodynamic ionization constant, K^, may then be readily calculated from the intercept. The ionic strength,60 , is defined as: 10 = l/2Zc±Z? (VI-1) where C I- concentration Z = valence = 1 for HBr These data are presented in Figs.9 and 10 and Tables XXXVII and XXXVIII. Considerable scattering is observed in the points of Figs.9 and 10. The function was assumed to be a straight line and a least squares treatment was used. Extrapolation to zero ionic strength gave an intercept equal to — > RT InK^. Values of the molal HF ionization constant calculated from the intercept were 1.01 x 10*"^ for benzoic acid and 2.05 x 10”^ for acetic acid. 32 - < - 0 0 o I .0591 log 59.5 60.0 60.5 61.0 0 THE IONIC STRENGTH,CO STRENGTH,CO vs. IONIC THE 8 0.0591 log log 0.0591 IONIC STRENGTH , X10 STRENGTH U3 IONIC IUE 9 FIGURE K a ACETIC ACID ACETIC 12 164 20 VjJ < - (M O o i .05915 log 60.0 59.5 58.0 58.5 i i i i i i i i i i i i i i i i i i i i i i I i i THE IONid STRENGTH, U), vs. 0.0591 log log 0.0591 vs.STRENGTH, U), IONid THE 7 BENZOIC ACID. BENZOIC IONIC STRENGTH, ii> STRENGTH, IONIC |0? X , 8 FIGURE 10 FIGURE 9 10 11 12 O 13 KA for 4 5 6 17 16 15 14 -P- U) VII DISCUSSION OF RESULTS Evaluation of the data of Goldschmidt has yielded 0.01&7 for the molar ionization constant of hydrogen bromide in ethanol. This constant is reasonable compared to 0,0113, the molar ionization constant of hydrogen chloride in ethanol used by Taniguchi and Janz. D. A. Mac Innes (20) has done an independent evaluation of the data of Goldschmidt and arrived at an ionization constant of 0.022 for hydrogen bromide in ethanol. He also reported however, an ionization constant of 0.015 for hydrogen chloride in ethanol, a value somewhat higher than that used by Taniguchi and Janz. Therefore,the values used in this investigation, 0.01S7, appears to be a very reasonable approximation. The standard electrode potentials in aqueous solutions as reported by Fc Daniels (4 ) give values of — 0.2223 volts for the Ag,AgCl;Cl- electrode reaction and — 0.073 volts for the Ag,AgBr;Br- electrode reaction. This is to be compared in ethanol with — 0.1940 volts for the Ag,AgBr;Br- electrode reaction and — 0.093&3 volts for the Ag,AgCl;Cl- electrode reaction as determined by Taniguchi and Janz. Thus there is a difference of 0.1493 volts in water and 0.1002 volts in ethanol. The silver-silver bromide standard electrode potential is the more negative in both aqueous and ethanol systems. 35 The determination of the ionization constant of benzoic acid affords an indirect check on the value of the standard potential of the silver-silver bromide electrode. A pK value of 10.0 was obtained by the author which may be compared to a pk value of 10.13 used by J. Elliott and M. Kilpatrick (7) in their studies, although these authors mention there is still some disagreement on the correct value of the pK of benzoic acid in ethanol* Another value of 10.43 has been reported. The experimentally determined pK of benzoic acid in ethanol (as calculated with the value of Eq found in this investigation) appears to be in good agreement with the most acceptable value of the pK of benzoic acid as determined by other workers using different systems. The ionization constants of acetic and benzoic acids in water are 1.753x10“^ and 6 . 3 0 x 1 0 “ ^ respectively as tabulated by D. A. Mac Innes (20). The values found in this investigation in ethanol were 2 .05x10“^ for acetic and l.OlxlO”10 for benzoic acids. Thus in ethanol, acetic acid has a slightly smaller ionization constant, just as it has the smaller ionization constant in water. A major drawback of the silver-silver bromide, hydrogen electrode galvanic cell is the inability to obtain reliable results in HBr concentration regions greater than 0.07 molar. The exact cause of instability has not been determined but it is not improbable that silver bromide itself becomes more soluble due to complex formation. The greatest source of error in applying the method for calculating the standard potential of the cell was the departure from ideality of the system at relatively low concentrations as shown in Fig.7* Thus, emphasis had to be placed on points which had been determined at low con centrations, necessitating the extrapolation of the straight line presented in Fig.6. The extrapolation to infinite dilution was not as unambiguous as in the hydrogen chloride system, however the results appear to be entirely reasonable. 38 TABLE I Conductivity Data of H. Goldschmidt and P. Dahll The Conductivity of hydrogen bromide in absolute ethyl alcohol. V X X 10 42.2 42.2 20 47.6 47.6 40 52.7 52.9 80 53.1 58.3 160 63.0 63.2 320 67.3 68.0 640 72.6 72.9 1230 75.6 75.8 2560 78.4 73.3 160-640 38.9 89.4 320-1230 88.9 89.1 640-2560 88.8 88.8 Average 88.9 1 - x where: x = concentration, gram v equivalent weights/1. X Z Equivalent Conductance 39 TABLE II Functions Calculated from the data of H. Goldschmidt and P. Dahll Y» degree of No. C j m/1 • F(Z) dissociation 1 0.00625 0.1624 0.82073 0.911 2 0.00312 0.1194 0.37214 0.923 3 0.00156 0.0870 0.90873 0.947 4 0,00781 0.0629 0.93495 0.959 0.00391 0.0453 0.95362 0.976 No. f2 F/A C A ^ / F 1 0.359 0.01301 0.1723 2 0.483 0.01284 0.1174 3 0.593 0.01250 0.0737 4 0.690 0.01235 0.0436 5 0.767 0.01213 0.0247 Intercept z 1198 A o = 3 3 . 4 7 9 Limiting Slope = 0.767x10“^ K 0.0187 m/1 TABLE III Physical Properties of Anhydrous Ethanol Experimentally Literature Determined, 25°C Values Density O.7S5O 0.7S506* Viscosity 1.077 Gentipoises 1.07^** Refractive Index 1.3635 at 20.0°C. 1.3624** 15^9.3ju ) at 13°C * J. Woolcock and H. Hartley (24) ** International Critical Tables (16) TABLE IV Observed Electromotive Force of the Cell Run 1 Normality HBr 1 0.00255 Molality HBr r 0.00325 Corrected Barometric Pressure - 767.1 mm Hg Equilibrium Time, Observed E.M.F., volt minutes Ag-AgBr Electrode #5 #2 0 0.1325 0.1325 15 0.1315 0.1315 30 0.1320 0.1320 45 0.1319 0.1319 60 0.1325 0.1325 Average: 0.1321 0.1321 0.1321 Corrected E, volts 0.1330 TABLE V Observed Electromotive Force of the Cell Run 2 Normality HBr “ 0.0119 Molality HBr = 0.0151 Corrected Barometric pressure — 766.£ mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag, AgBr Electrode #5 0 0.0620¥ 0.0620 10 0.0623 0.0623 20 0.0621 0.0621 30 0.0621 0.0621 40 0.0621 0.0621 50 0.0620 0.0620 60 0.0622 0.0622 Average: 0.0621 0.0621 0.0621 Corrected E, Volts 0.0630 43 TABLE VI Observed Electromotive Force of the Cell Run 3 Normality HBr z 0.01B04 Molality HBr = 0.0230 Corrected barometric pressure 762.0 mm Hg brium Time, Observed E.M.F., volts nutes Ag-AgBr Electrode ULr\ #5 0 0.0439 0.0439 10 0.0441 0.0441 o o O O LfM> . 20 * . o o o o • • • • 30 -P--P- 0.0441 0.0441 O O O 0.0444 0.0444 o o 1—1 0.0439 0.0441 0.0441 0.0446 20 0.0444 0.0446 30 0.0443 0.0446 40 0.0444 0.0446 Average 0.0443 0.0444 0.04435 Corrected E, volts 0.0454 TABLE VII Observed Electromotive Force of the Cell Run 4 Normality HBr = 0.0536 Molality HBr z 0.0682 Corrected Barometric Pressure 762.0 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode #2 #5 0 -0.0022 -0.0022 15 -0.0022 -0.0021 25 -0.0021 -0.0021 40 -0.0022 -0.0022 60 -0.0020 -0.0020 85 -0.0020 -0.0020 Average: -0.0021 i I i o oo o MH o o i«o ro . . . Corrected E, volts —0.0011 TABLE VIII Observed Electromotive Force of the Cell Run ■5 Normality HBr = 0.0286 Molality HBr = 0.0364 Corrected barometric pressure 760.5 mm Hg Llibrium Time, Observed E.M.!F., volts minutes Ag-AgBr Electrode #2 #5 0 0.0200 0.0204 15 0.0198 0.0206 40 0.0196 0.0204 65 0.0198 0.0202 75 0.0195 0.0200 0 0.0200 0.0200 20 0,0204 0.0204 30 0.0204 0.0204 Average: 0.0199 C.0203 0.0201 Corrected E, volts 0.0211 46 TABLE IX Observed Electromotive Force of the Cell Run 6 Normality HBr I 0,0416 Molality HBr = 0.0530 Corrected barometric pressure 764*5 mm Hg ilibrium Time, Observed E.M.F., volts .minutes Ag-AgBr Electrode $2 #5 0 0.0090 0.0089 25 0.0088 0.0069 35 0.0069 0.0069 125 0.0064 0.0060 135 0.0066 0.0089 150 0.0093 O.OO65 0 0.0065 0.0085 20 O.OO65 0.0085 85 0.0063 0.0063 Average 0.006? 0.0086 0.00665 Corrected E, volts 0.0097 TABLE X Observed Electromotive Force of the Cell Run 7 Normality HBr r 0.0346 Molality HBr = 0.0440 Corrected barometric pressure 764.5 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode $2 #5 0 0.0126 0.0126 15 0.0127 0.0127 35 0.0127 0.0127 55 0.0128 0.0126 70 0.0123 0.0127 Average 0.0127 0.0127 0.0127 Corrected E, volts 0.0137 TABLE XI Observed Electromotive Force of the Cell Run 8 Normality HBr = 0.0692 Molality HBr Z 0.0881 Corrected barometric pressure 755*4 m Hg Equilibrium Time Observed E.M.F., volts minutes Ag-AgBr Electrode #2 #5 0 -0.0147 -0.0147 15 -0.0145 -0.0146 25 -0.0143 i i i o o o 40 —0.0144 • • o o H H 55 -0.0142 -0.0142 65 -0.0145 -0.0145 75 -0.0143 -0.0143 Average -0.0144 -0.0145 -0.01445 Corrected E, volts -0.0134 TABLE XII Observed Electromotive Force of the Cell Run 9 Normality HBr z 0.0496 Molality HBr = 0.0631 Corrected barometric pressure 763 mm Hg Equilibrium Time, Observed" E.M.F., volts minutes Ag-AgBr Electrode #2 #5 0 0.0002 0.0002 25 0.0004 0.0004 40 -0.0004 -0.0004 55 -0.0009 —0.0004 75 -0.0003 -0.0004 95 -0.0004 -0.0004 110 -0.0004 0.0005 125 0.0012 0.0012 Average 0.0000 0.0001 0.0000 Corrected E, volts 0.0010 TABLE XIII Observed Electromotive Force of the cell Run 10 Normality HBr = 0.022$ Molality HBr =: 0.0290 Corrected barometric pressure 757-$ mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode :2 #5 0 0.033$ 0.0333 40 0.0342 0.0342 55 0.033$ 0.033$ 80 0.0339 0.0341 95 0.0341 0.0344 110 0.0341 0.0344 120 0.0340 0.0344 130 0.0342 0.0344 Average 0.0340 0.0341 0.03405 Corrected E, volts 0.0351 TABLE XIV Observed Electromotive Force of the Cell Run H Normality HBr = 0.00807 Molality HBr = 0.0103 Corrected barometric pressure 757.6 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode h #6 0 0 .0 S06 0.0810 40 0.0310 0.0818 50 0.0307 0.0807 65 0.0807 0.0807 65 0.0310 0.0808 120 0.0336 0.0836 130 0.0810 0.0810 150 0.0810 0.0810 160 0.0814 0.0814 170 0.0314 0.0814 Average 0.0812 0.0813 0.08125 Corrected E, volts 0.0824 TABLE XV Observed Electromotive Force of the Cell Run 12 Normality HBr r 0.00229 Molality HBr = 0.00291 Corrected barometric pressure 757.6 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr electrode h #6 0 0.1355 0.1375 15 0.1348 0.1348 55 0.1350 0.1357 90 0.1345 0.1350 105 0.1350 0.1350 130 0.1344 0.1344 150 0.1350 0.1350 175 0.1352 0,1364 200 0.1351 0.1360 Average 0.1349 0.1355 0.1352 Corrected E, volts 0.1363 TABLE XVI (Eo - Eot) and (Eo - EpJ^as a function of the concentration a — 4*0 No. C,moles/l x leP [gX^ (x)-21 ^ (x)] 10^ Jj^X^ (x)-4X^ (x:)J 1 0.0009 0.07092 -0.09594 -0.1030 2 0.0016 0.09456 -0.14105 -0.1371 3 0.0025 0.1182 -0.18476 -0.1599 4 0.0036 0.1416 -0.22530 -0.1705 5 0.0049 0.1655 -0.26194 -0.1701 6 0.0064 0.1961 -0.2941 -0.1608 7 0.0100 0.2364 -0.3453 -0.1232 6 0.0144 0.2837 -0.3805 -0.0728 9 0.0196 0.3310 -0.4023 -0.0202 -■> No. J. oc x * Eobs. (Calculated) 1 0.806 0.970 0.06984 0.1783 2 0.752 0.955 0.09240 0.1528 3 0.704 0.940 0.1146 0.1330 4 0.662’ 0.925 0.1364 0.1168 5 0.624 0.915 0.1583 0.1031 6 0.591 0.900 0.1794 0.0913 7 0.535 0.880 0.2217 0.0715 8 0.469 0.865 0.2636 0.0553 9 0.451 0.845 0.3043 0.0417 54 TABLE XVI Continued ur» No. lo3[iX3 tx)-2Y3 (x)]0C 105[&X5 (x )-4Y5 U-)|* 1 -0.09373 -0.10110 0.808 2 -0.13687 -0.13443 0.757 3 -0.17825 -0.15710 0.710 4 -0.21633 -0.16680 0.672 5 -0.25130 -0.17123 0.635 6 -0.28148 -0.16543 0.604 7 -0.33121 -0.13847 0.551 6 -0.36747 -0.09895 0.507 9 -0.39144 -0,05453 0.472 No. -(Eo-E*) -(Eo-E*)1 1 0.1945 0.1929 2 0.1946 0.1923 3 0.1955 0.1923 4 0.1968 0.1928 5 0.1980 0.1934 6 0.1995 0.1941 7 0.2023 0.1957 8 0.2049 0.1974 9 0.2076 0.1989 TABLE XVII (Eo - E*) and (E6 - Eot)1 as a function of the concentration & = 5.0 lo. C,moles/l X 10^ [iX^ (x)-2Xcj (x)J 10^ [j^X5(x)-4Yt 1 0.0009 0.08865 -0.12982 -0.12969 2 0,0016 0.1182 -0.16191 -0.15987 3 0.0025 0.1478 -0.23498 -0.17135 4 0.0036 0.1773 -0.27858 -0.16642 5 0.0049 0.2069 -0.31538 -0.14902 6 0.0064 0.2364 -0.34529 -0.12322 7 0.0100 0.2955 -0.38705 -0.05959 8 0.0144 0.3546 -0.40884 0.00472 No. f oc Xot Eobs. (Calculated) 1 0.816 0.970 0.08732 0.1783 2 0.768 0.950 0.1149 0.1528 3 0.723 0.940 0.1432 0.1330 4 0.685 0.920 0.1700 0.1168 5 0.651 0.910 0.1974' 0.1031 6 0.620 0.895 0.2236 0.0913 7 0.569 , ,0.870 0.2757 0.0715 8 0.525 0.&50 0.3269 0.0553 56 TABLE XVII Continued No. 103[jX3 (x)-2Y3 (x)J(< 105[lX5{x)-4Y5(x))« 1 -0.12714 -0.12771 0.313 2 -0.17379 -0.15733 0.772 3 -0.22756 -0.17069 0.730 k -0.26351 -0.16910 0.694 5 -0.30433 -0.15572 0.661 6 -0.33309 -0.13521 O'. 633 7 -0.37555 -0.03172 0.534 3 -0.40030 -0.02469 0.544 No. -(Eo - E*) -(e £ t- E*)' 1 0.1939 0.1923 2 0.1939 0.1913 3 0.1942 0.1910 4 0.1953 0.1911 5 0.1963 0.1914 6 0.1974 0.1917 7 0.1999 0.192? 3 0.2022 0.1933 TABLE XVIII (Eo - E*) and (Eo - E*) * as a function of the concentration a r 6.5 No. G,moles/1 X 10? (x)-2X3 (xj) 10^ [|x 5 (x)-4Y5(xj) 1 0.0009 0.1152 -0.17933 -0.15756 2 0.0016 0.1537 -0•24422 -0.17150 3 0.0025 0.1921 -0.29786 -0.15905 4 0.0036 0.2305 -0.33989 -0.12694 5 0.0049 0.2689 -0.37109 -0.08924 6 0.0064 0.3073 -0.39284 -0.04637 7 0.0100 0.3842 -0.41385 0.03377 8 0.0144 0.4610 -0.41335 0.0947 9 0.0196 0.5378 -0.39937 0.1334 No. f oc x* Eobs. (Calculated) 1 0.825 0.970 0.1135 0.1783 2 0.780 0.955 0.1502 0.1528 3 0.741 0.935 0.1857 0.1330 4 0.706 0.920 0.2211 0.1168 5 0.675 0.900 0.2551 0.1031 6 0.647 O .885 0.2891 0.0913 7 0.600 0.865 0.3573 0.0715 8 0.561 0.835 0.4213 0.0553 9 0.529 0.810 0•4840 0.0417 58 TABLE XVIII Continued No. 103[|X3 (x )-2Y3 (x )) 1 - 0.17626 -0.15626 0.827 2 -0.23664 -0.17166 0.764 3 -0.26971 -0.16037 0.747 4 -0.33062 -0.13647 0.714 5 -0.36096 -0.09695 0.686 6 -0.36363 -0.06577 0.660 7 -0.40946 0.01027 0.615 6 -0.41574 0.0650 0.580 9 -0.41025 0.1063 0.551 No. -(Eo - Epc) -(Eo — Ek )» 1 0.1933 0.1918 2 0.1928 0.1906 3 0.1932 0.1897 4 0.1939 0.1896 5 0.1950 0,1895 6 0.1959 0.1896 7 0.1975 0.1901 8 0.1999 0.1906 9 0.2017 0.1909 59 TABLE XIX (Eo - E*) as a function of the concentration g = 3.0 No. C,moles/l x 10^ [|X^ (x)-2Y^ (x)J l O ^ X ^ (x)-4X(xjj 1 0.0009 0.1418 -0.22530 -0.17049 2 0.0016 0.1891 -0.29413 -0.16083 3 0.002$ 0.2364 -0.34529 -0.12322 4 0.0036 0.2837 -0.38048 -0.07282 5 0.0049 0.3310 -0.40225 -0.02023 6 0.0064 0.3782 -0.41313 -0.02812 7 0.0100 0.i?28 -0.41184 -0.01019 8 0.0144 0.5674 -0.39164 0.01430 9 0.0196 0.6619 -0.36205 0.01581 No. f OC Eobs. (Calculated) 1 0.831 0.970 0.1397 0.1783 2 0.789 0.950 0.1844 0.1528 3 0.753 0.930 0.2279 0.1330 4 0.720 0.915 0.2712 0.1168 5 0.692 0.895 0.3131 0.1031 6 0.666 0.885 0.3559 0.0913 7 0.623 0.850 0.4359 0.0715 8 0.$88 0.820 0.5141 0.0553 9 0.558 0.790 0.5884 0.0417 60 TABLE XIX Continued No. 103 [£x3(x)-2Y3(x)] 105[iX5(x)-4Y5(x)]^ 1 -0.22189 -0.17023 0.833 2 -0.28803 -0.16543 0.793 3 -0.33735 -0.12942 0.759 4 -0.37264 -0.08804 0.729 5 -0.39540 -0.04335 0.702 6 -0.40914 -0.01027 0.678 7 -0.41559 0.0803 0.639 s -0.40469 0.1264 0.607 9 -0.3&563 0.1464 0.581 No. -(E; - Ej 1 0.1929 2 0.1925 3 0.1926 4 0.1932 5 0.1940 6 0.1945 7 0.1964 8 0.1984 9 0 .2003 TABLE XX Observed Electromotive Force of the Cell Run A-l Molality NaBr = 0.00610 Molality HAc “ 0.00539 Molality NaAc = 0.00626 Corrected Barometric Pressure 760.2 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode #5 #2 0 0.5520 0.5525 20 0.5490 0.5500 60 0.5518 0.5516 90 0.5490 0.5488 120 0.5486 0.5466 140 0.5486 0.5481 165 0.5473 0.5477 180 0.5492 0.5481 215 0.5480 0.5483 245 0.5498 0.5486 300 0.5482 0.5490 Average. 0.5492 0.5492 0.5492 Corrected E, volts 0.5502 TABLE XXI Observed Electromotive Force of the Cell Run A-2 Molality NaBr - 0.00721 Molality HAc = 0.00662 Molality NaAc = 0.00748 Corrected Barometric Pressure 759*7 mm Hg Equilibrium Time Observed E.M.F., volts minutes Ag-AgBr Electrode #6 0 0.5361 0.5384 15 0.5378 0.5405 30 0.5381 0.5405 45 0.5378 0.5390 60 0.5378 0.5400 75 0.5375 0.5386 90 0.5385 0.5404 110 0.5375 0.5405 125 0.5401 0.5415 150 0.5381 0.5386 Average 0.5379 0.5398 0.5389 Corrected E, volts 0.5399 TABLE XXII Observed Electromotive Force of the Cell Run A—3 Molality NaBr = 0.01019 Molality HAc = 0.00891 Molality NaAc = 0.00953 Corrected barometric pressure 756.8 mm Hg Equilibrium Time, Observed E.M.F., vo3 minutes Ag-AgBr Electrode . h #6 0 0.5283 0.5273 10 0.5277 0.5284 30 0.5290 0,5281 40 0.5283 0.5286 65 0.5300 0.5290 BO 0.5290 0.5293 100 0.5290 0.5281 125 0.5290 0.5305 145 0.5294 0.5291 160 0.5275 0.5300 170 0.5301 0.5296 180 0.5280 0.5295 190 0.5286 0.5280 Average 0.5288 0.5289 0.5288 Corrected E, volts 0.5300 TABLE XXIII Observed Electromotive Force of the Cell Rim A-4 Corre Molality NaBr s: 0.00190 Molality HAc s 0.00220 Molality NaAc = 0.00212 Corrected barometric Pressure 757.7 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode |?2 $5 0 0.5781 0.5804 15 0.5796 0.5805 30 0.5794 0.5811 45 0.5804 0.5800 75 0.5804 O.58OO 90 0.5810 0.5796 105 0.5602 0.5804 125 0.5813 0.5798 140 O . 58O 5 0.5605 155 0.5810 0.5802 170 0.5804 0.5811 190 0.5806 O.58OI 210 0.5811 0.5810 Average O .5803 O .5804 0.5803 Corrected E, volts O. 58I5 TABLE XXIV Observed Electromotive Force of the Cell Run A-5 Molality NaBr = 0,00373 Molality HAc = 0.00375 Molality NaAc = 0.00411 Corrected barometric pressure 758*3 nun Hg Equilibrium Time, Observed E.M.F., vo] minutes Ag-AgBr Electrode #5 n 0 0.5676 0.5650 20 0.5666 0.5650 55 0.5674 0.5651 75 0.5673 0.5660 130 0.5668 0.5654 145 0.5671 0.5650 165 0.5662 0.5652 ISO 0,5654 0.5644 195 0.5650 0.5645 210 0.5670 0.5660 230 0.5645 0.5645 245 0.5645 0.5648 260 O.566O 0.5650 Average 0.5663 0.5651 0.5657 Corrected E, volts 0.5668 TABLE XXV Observed Electromotive Force of the Cell Run A-6 Molality NaBr z 0.00546 Molality HaC = 0.00939 Molality NaAc = 0.00576 Corrected barometric pressure 753.1 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode #5 #2 0 0.5379 0.5369 20 0.5331 0.5374 40 0.5373 0.5366 65 0.5331 0.5369 30 0.5370 0.5371 95 0.5370 0.5370 110 0.5370 0.5363 135 0.5365 0.5365 153 0.5357 0.5360 200 , 0.5369 0.5363 225 0.5376 0.5363 240 0.5373 0.5367 Average O .5372 0.5367 0.5370 Corrected E, volts 0.5331 TABLE XXVI Observed Electromotive Force of the Cell Run A-7 Molality NaBr z 0.00458 Molality HAc Z 0,00457 Molality NaAc = 0.00517 Corrected barometric pressure 757#5 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode #5 #2 0 0.5539 0.5595 20 0.5592 0.5530 55 0.5535 0.5534 75 0.5564 0.5534 85 0.5533 0.5531 100 0.5571 0.5564 120 0.5533 0.5594 135 0.5535 0.5532 1555 0.5530 0.5532 170 0.5534 0.5535 205 0.5576 0.5590 220 0.5599 0.5594 Average 0.5536 0*5536 0.5536 Corrected E, volts 0.5597 TABLE XXVII Observed Electromotive Force of the Cell Run A-8 Molality NaBr = 0.00297 Molality HAc = 0.00238 Molality NaAc = 0.00302 Corrected barometric pressure 757.0 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode r r2 r r ? 0 0.5758 0.5764 15 0.5765 0.5755 60 0.5760 0.5760 75 0.5778 0.5790 90 0.5775 0.5788 105 0.5784 0.5778 120 0.5776 0.5784 135 0.5776 0.5775 150 0.5770 0.5786 165 0.5775 0.5770 180 0.5766 0.5783 195 0.5784 0.5786 Average 0.5772 0.5777 0.5775 Corrected E, volts 0.5786 TABLE XXVIII Observed Electromotive Force of the Cell Run B-l Molality NaBr z 0.00420 Molality Benzoic Acid ~ 0.00290 Molality Socium benzoate = 0.00315 Corrected barometric pressure 760.9 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode if 5 $2 0 0.5492 0.5492 20 0.5430 0.5494 35 0.5475 0.5495 50 0.5437 0.5494 70 0.5495 0.5516 35 0.5473 . 0.5473 100 0.5431 0.5493 115 0.5482 0.5490 125 0.5430 0.5491 135 0.5479 0.5437 150 0.5475 0.5490 Average 0.5432 0.5493 0.5437 Corrected E, volts 0.5497 TABLE XXIX Observed Electromotive Force of the Cell Run B-2 Molality NaBr = 0.00709 Molality benzoic acid - 0.00554 Molality sodium benzoate = O.OO5I6 Corrected barometric pressure 757*6 mm Hg .librium Time, Observed E.M.F., voj minutes Ag-AgBr Electrode #5 #2 0 0.5263 0.5274 20 0.5270 0.5270 30 0.5268 0.5285 40 0.5271 0.5271 50 0.5260 0.5260 65 0.5255 0.5250 75 O .5263 0.5270 90 0.5261 0.5261 105 0.5260 0.5270 125 0.5250 0.5250 195 0.5250 0.5259 235 0.5250 0.5250 Average 0.5260 0.5264 0.5262 Corrected E, volts 0.5273 TABLE XXX Observed Electromotive Force of the Cell Run 13-3 Molality NaBr - 0.00999 Molality benzoic acid I 0.00828 Molality sodium benzoate - 0.00701 Corrected barometric pressure 761.5 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode ir5 $2 0 0.5203 0.5198 30 0.5200 0.5190 50 0.5193 0.5190 65 0.5195 O .5186 85 0.5204 0.5194 120 0.5198 0.5198 240 0.5193 0.5190 Average 0.5198 0.5192 0.5195 Corrected E, volts 0.5205 TABLE XXXI Observed Electromotive Force of the Cell Run B-4 Molality NaBr = 0.00221 Molality benzoic acid = 0.00140 Molality sodium benzoate = 0.00132 Corrected barometric pressure 762.0 mm Hg Equilibrium Time, Observed E.M.F., vo] minutes Ag-AgBr Electrode #5 :0 0.5627 0.5634 15 0.5644 0.5625 30 0.5635 0.5634 45 0.5636 0.5636 60 O. 56I6 0.5639 85 0.5644 0.5644 110 0.5635 0.5649 130 0.5629 0.5625 155 0.5664 0.5650 175 O .5625 O .5627 205 O .5630 0.5640 Average 0.5635 0.5637 O .5636 Corrected E, volts 0.5646 TABLE XXXII Observed Electromotive Force of the Cell Run B-5 Molality NaBr = 0.00558 Molality benzoic acid I 0.00379 Molality sodium benzoate = 0.00409 Corrected barometric pressure 762.0 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode #5 #2 0 0.5425 0.5405 20 0.5430 0.5412 40 0.5420 0.5428 70 0.5425 0.5415 85. 0.5415 0.5415 100 0.5418 0.5415 125 0.5425 0.5417 145 0.5439 0.5410 160 0.5408 0.5420 180 0.5417 0.5415 190 0.5410 0.5410 Average 0.5421 0.5415 0.5418 Corrected E, volts 0.5428 TABLE XXXIII Observed Electromotive Force of the Cell Run B-6 Molality NaBr " 0.00715 Molality benzoic acid z 0.00640 Molality sodium benzoate = 0.00541 Corrected barometric pressure 762.0 mm Hg Equilibrium Time, Observed E.M.F., vo3 minutes Ag-AgBr Electrode #5 #2 0 0.5250 0.5273 20 0.5242 0.5269 35 0.5240 0.5267 50 0.5240 0.5264 65 0.5243 0.5260 95 0.5236 0.5263 335 0.5256 0.5265 345 0.5247 0.5264 360 0.5258 0*5260 375 0.5251 0.5263 410 0.5264 0.5264 425 0.5255 0.5264 445 0.5260 0.5260 470 0.5256 0.5264 485 0.5284 0.5279 Average 0.5252 0.5265 0.5259 Corrected E, volts 0.5269 TABLE XXXIV Observed Electromotive Force of the Cell Run B-7 Molality NaBr Z.0.002S4 Molality benzoic acid = 0.00242 Molality sodium benzoate = 0.00203 Corrected barometric pressure 757.7 mm Hg Equilibrium Time, Observed E.M.F., volts minutes Ag-AgBr Electrode #5 #2 0 0.5525 0.5505 20 0.5510 0.5505 35 0.5525 0.5515 45 0.551S 0.5514 60 0.5520 0.5520 75 0.5515 0.551^ S5 0.5520 0.5515 95 0.5515 0.5525 105 0.5505 0.5521 115 0.5515 0.5520 125 0.5519 0.5515 135 0.551S 0.5512 Average 0.5517 0.5515 0.5515 Corrected E, volts 0.5526 TABLE XXXV Observed Electromotive Force of the Cell Run B-3 Molality NaBr r 0,00377 Molality benzoic acid - 0,00391 Molality sodium benzoate “ 0,00339 Corrected barometric pressure 760,$ mm Hg Equilibrium Time, Observed E.M.F., vo] minutes Ag-AgBr Electrode #5 #2 0 0.5440 0.5450 15 0.5452 0.5445 30 0.5451 0.5446 45 0.5456 0•5444 60 0.5456 0.5446 90 0.5456 0.5446 10$ 0.5452 0•5444 H 5 0.5445 0.5445 135 0,5443 0.5443 165 0.5455 0.5440 Average 0.5451 0.5445 0.5453 . Corrected E, volts 0.5463 77 TABLE XXXVI The function -0.0591 log as dependent on the ionic strength, ACETIC ACID Run No. mHA mBr 60 -0.0591 log mAc A-l 0.00525 0.01236 0.5969 A-2 0.00638 0,014699 0.5917 A-3 0.00948 0.01977 0.5919 A-4 0.00197 0.00402 0.6031 A-5 0.00340 0.00784 0.6024 A-6 0.00937 0.01122 0.5997 A— 7 0.00405 0.00975 0.5998 A-8 0.00234 0.00599 0.6046 ^Acetic Acid, Ethanol z 2.05x10“^^- BENZOIC ACID B-l 0.00387 0.00735 0.5886 B— 2 0.00761 0.01225 0.5836 B-3 0.0117 0.01702 0.5878 B-4 0.00234 0.00353 0.5906 B-5 0.00517 0.00967 0.5892 B-6 0.00846 0.01256 0.5859 B-7 0.00339 0.00487 0.5881 B-8 0.00435 0*00716 0.5882 ^Benzoic Acid, Ethanol = l.OlxlO”^-0 BIBLIOGRAPHY Bates, Roger G. Electrometric pH Determinations New York: Johnwiiey and Sons, Inc., 1954* p. 202 Bez-man, I. I. and Verhock, F. H. The Conductance of Hydrogen Chloride and Ammonium Chloride in Ethanol-Water Mixtures. J. Am. Chem. Soc.. 67. 1330 (1945) Clow, A. and Pearson, G. Pure Ethyl Alcohol for Absorption Spe ctrophotometry. Nature. 144. 208 (1939) Daniels, Farrington Outlines of Physical Chemistry New Yorki”"3ohn Wiley and Sons, Inc., 1952 p. 447 Danner, P. S. and Hildebrand, J. H. The Degree of Ionization of Ethyl Alcohol. 1. From Measurements of Conductivity J. Am. Chem. Soc., 44* 2824 -{1922) Encvlopedia of Chemical Technology TJew York: Tnterscience Publishers, Inc., 1947 Elliott, J. H. and Kilpatrick, M. The Effect of Substituents on the Acid Strength of Benzoic Acid. II. in Ethyl Alcohol J. Phys. Chem. 45* 466 (1941) Fuoss, R. M. Solution of the Conductance Equation. Soc.. ££, 488 (1935) Fuoss, R. M. and Kraus, C. A. Properties of Electrolytic Solutions. II. The Evaluation of « and of K for Incompletely Dissociated Electrolytes. J. Am. Chem. Soc.. 55_. 476 (1933) Goldschmidt, H. and Dahll, P. Die Leitfahigkeit und die katalytische Wirkung der drei starken Halogen- Wasserstoffsauren in methyl - und athylalkolischer Losung. 2. physik. Chem. Leipsig. 114. 1 (192$) 79 11. Gronwall, T. H., LaMer, V. K. and Sandved, K. Uber den Einfluss der sogennanten hoheren Glieder in der Debye-Huckelschen Theorie der losungen starker Elektrolyte Phvsik. Zi, 22, 358 (1923) 12. Harned, H. S. and Donelson, J. G. The Thermodynamics of Ionized Water in Lithium Bromide Solutions. Chem. Soc.. 1230 (1937) 13* Harned, H. S., Keston, A. S. and Donelson, J. G. The Thermodynamics of Hydrobromic Acid in Aqueous Solutions from Electromotive Force Measurements. Anu Chem. Soc. j>3, 989 (1936) 14. Harned, H. S. and Owen, B. B. The Physical Chemistry of Electrolytic Solutions. New York: Reinhold Publishing Corporation, 1950, ( Second Edition) 15* Harris, L. Purification arid Ultraviolet Transmission of Ethyl Alcohol. ii Anu Chem. Soc.. ££, 1940 (1933) 16. International Critical Tables New York: McGraW-Hill Book Company, 1923 17- Keston, A. S. A Silver-Silver Bromide Electrode Suitable for Measurements in Very Dilute Solutions. Chem. Soc. j>2, 1671 (1935) 18s LeBas, Carlyle. Master*s Thesis Louisiana State University, 1959 19* Lund, H. and Bjerrum, J, Eine einfache Methods Zur Darstellung wasserfreier Alkohole Ber.. 64B, 210 (1931) 20. Maclnnes, D. A. The Principles of Electrochemistry New York: Reinhold Publishing Corporation 1939 go 21. Mukherjee, L. M. The Standard Potential of the Silver- Silver Chloride Electrode in Ethanol, it Phys. Chem.. £8, 1042 (1954) 22. Osborne, N» S., McKelvy, E. C. and Bearce, H. W. Density and Thermal Expansion of Ethyl Alcohol and its mixtures with water. Bull. Bur. Standards. £ (Mo. 3), 327 (1913) 23. Taniguchi, H. and Janz, G. T. The Thermodynamics of Hydrogen Chloride in Ethyl Alcohol from Electromotive Force Measurements. J. Phys. Chem.. 61, 688 (1957) 24. Woolcock, J. W. and Hartley, H. The Activity coefficients of Hydrogen Chloride in Ethyl Alcohol. Phil. Mag.. 5 (7th Series), 1133 (1928) VITA Loys Joseph Nunez, Jr. was born in New Orleans, Louisiana on March 18, 1926. He received his elementary education in the parochial schools of New Orleans and Madisonville, Louisiana. He was graduated with honors from St. Paul*s College (high school), Covington, Louisiana, in May 1944 and in August, 1947 he received the Bachelor of Science Degree from Tulane University, majoring in Chemistry. He was then in the employment of Cities Service Refining Corporation and Cities Service Research and Development Corporation for six years. In September of 1953 he entered the Graduate School of Louisiana State University and received the Master of Science degree in Chemistry in May, 1955* He spent the academic year of 1957-5# as a Fulbright Student at the University of Tubingen, Ttibingen, West Germany. He is at present a candidate for the degree of Doctor of Philosophy. 81 EXAMINATION AND THESIS REPORT Candidate: Loys J. Nunez Major Field: Chemistry Title of Thesis: Thermodynamic Studies in Anhydrous Ethanol - Silver^ Silver Bromide Galvanic Cell Approved: i f , I Major Professor and C h a i r m// a n the Graduate School EXAMINING COMMITTEE: Date of Examination: August 3, 19^9