Brigham Young University BYU ScholarsArchive
Theses and Dissertations
1974-08-01
Thermodynamic stability constants of silver bromide complexes in water and water-acetone mixtures determined by ion-selective electrode measurements
Arthur Lee Cummings Brigham Young University - Provo
Follow this and additional works at: https://scholarsarchive.byu.edu/etd
BYU ScholarsArchive Citation Cummings, Arthur Lee, "Thermodynamic stability constants of silver bromide complexes in water and water-acetone mixtures determined by ion-selective electrode measurements" (1974). Theses and Dissertations. 8194. https://scholarsarchive.byu.edu/etd/8194
This Dissertation is brought to you for free and open access by BYU ScholarsArchive. It has been accepted for inclusion in Theses and Dissertations by an authorized administrator of BYU ScholarsArchive. For more information, please contact [email protected], [email protected]. &-V f, O 22. tCU,,· !CJ?~ mERMODYNAMICSTABILITY CONSTANTS OF SILVERBROMIDE COMPLEXESIN WATERAND WATER-ACETONE MIXTURES DETERMINEDBY ION-SELECTIVEELECTRODE MEASUREMENTS
A Dissertation Presented to the Department of Chemistry Brigham Young University
In Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy
by Arthur Lee Cummings August 1974 This dissertation, by Arthur Lee Cummi_ngs, is accepted in its present form by the Department of Chemistry of Brigham Young University as satisfying the thesis requirement for the degree of Doctor of Philosophy.
ii ACKNOWLEDGMENTS
Tile author wishes to express his gratitude to his advisory com- mittee, Dr. Keith P. Anderson, Dr. Earl M. Woolley, and Dr. James L. Bills, for their direction, encouragement and sound advice during the period of his graduate study and preparation of this dissertation. While·the author was yet an undergraduate student, Dr. Anderson awakened his interest in graduate work. Since then, Dr. Anderson has directed the author into interesting and challenging research areas and has provided him with many valuable opportunities to develop research and teaching abilities. Tilree other members of the faculty of the Brigham Young Univer- sity Chemistry Department have been of specific assistance to the author. Dr. J. Rex Goates and Dr. Eliot A. Butler have shown concern, given encouragement and provided examples of teaching excellence. Dr. Coran L. Cluff, Assistant Department Chainnan, has been helpful in many, many ways. Tile efforts in the author's behalf of other members of the faculty of the Brigham Young University Chemistry Department, who have been his instructors in the classroom and his supervisors in his employ- ment as a teaching assistant, are gratefully acknowledged. Fellow students have also been helpful, especially those who have occupied rooms 208 and 209 of the Eyring Science Center during the past f~ur years, who have provided an environment conducive to concentration, study and research, and who have put up with the disturbances which have accompanied some of the author's experimental work. A special expression
iii of gratitude is extended to Karl J. Walker, Jr., who during his own graduate research, provided the author with experience which has proven invaluable in the completion of th.is work. Also, Mr. Walker•s frequently expressed, sincere overestimations of the author•s abilities have been very confidence-building. Thanks are due the Brigham Young University Graduate School for financial assistance through the Supplementary Award and the Internship programs. Financial assistance was also received from the John Einar Anderson EndowmentFund, for which the author is very grateful. Gratitude is also expressed to Elizabeth Sykora Prisbrey for her diligent and consciencious typing of the manuscript. Very important to the author have been tb.e patience, faith, and love of bis wife, Joy, and his children, David, Matthew, and Kirsten. Without their support this dissertation may never have been WTitten.
iv TABLEOF CONTENTS
Page ACKNOWLEDGMENTS iii LIST OF TABLES. vii LIST OF FIGURES viii I. INTRODUCTION 1 A. Background 1 B. Scope and Purpose. 2
II. EXPERIMENTAL 4 A. Materials. 4 B. Procedure and Equipment .• 4 Preparation of Solutions Potentiometric Titrations C. Calculations 6 Electrode Calibration and Ks Stability Constants of the Complexes Activity Coefficients Precision of Least Squares Constants The Average Ligand Number D. Error Analysis 14
III. RESULTS 16 A. Behavior of the Silver Ion Selective Electrode 16 B. Equilibrium Constants. . • . . . • . . • • • • 27
V Page IV. DISCUSSION...... ' . . . . 40 A. Comparison of the Results of this Work with the Results of Other Investigations • • • • • • • 40 Electrode Behavior Equilibrimn Constants
B. Suggestions for Further Study 45
C. Summaryand Conclusions 47
APPENDIX•. 48
REFERENCES 64
.,
vi LIST OF TABLES
Table Page 1. Densities and Dielectric Constants of Water-Acetone Mixtures. . . • ...... 12
2. E' and c Values Used in Experiments I-IX .• 16
. 3. Stability Constants Assuming the Presence of AgBr, AgBr2-, and AgBr32- ..•..••.••.•••• 27 4. Stability Constants Assuming the Presence of AgBr, AgBr2-, and AgBr43- . • • • • • • . • • • •• 32 S. Thermodynamic Equilibrium Constants of Silver Bromide Complexes in Acetone-Water Mixtures. • • • . 35 6. Comparison of the Results of this Work with the Results of Other Investigations • • • • • . • • • • • . • • 43
vii LIST OF FIGURES
Figure Page = 0.06108. 18 1. Plot of F0 versus aAg+ at aBr- 2. Plot of ii• versus log CBr- for Experiments I, II, and III. 19 3. Plot of n' versus log CBr- for Experiments IV and V.• 20 4. Plot of n' versus log CBr- for Experiments VI and VII •• 21 s. Plot of ii.- versus log CBr- for Experiments VIII and IX 22 6. Plot of F versus aA +' and F versus aA + at 0 g 0 corr g corr a = 0.06108...... 24 8 r- 7. Plot of D versus log yAg+CAg(total)" 26 8. Plot of log F versus log a for Experiments I, II, o 8r- and I I I...... 28 9. Plot of log F versus log aBr- for Experiments IV and V•• 29 0 10. Plot of log F versus log aBr- for Experiments VI and VII. 30 0 11. Plot of log F versus log aBr- for Experiments VIII and IX • • • • 0 • . • • • • • • • • • • • • • • • • • • • 31
12. Plot of -log aAg+aBr- versus log aBr- for Experiments VI and IX • • • • • • • • • • • • • • • • • • • • • • 34
13. Plot of -log Ks versus Weight Percentage Acetone in Solvent...... 36
14. Plot of log 61 1 versus Weight Percentage Acetone in Solvent. . . ' ...... 37
15. Plot of log B1 2 versus Weight Percentage Acetone in Sol vent. . . ' ...... 38
16. Plot of log 81 3 versus Weight Percentage Acetone in Solvent ...... ' . ·39
17. Plot of log a1 1 versus Weight Percentar,e Acetone in Solvent for This Work.and.for Reference 6, before and after Correction of Data •••••.•.•.. 46
viii I. INTRODUCTION
A. Background Perhaps the simplest evidence of the existence of complexes of silver and bromide ions in solution is the increased solubility of the solid silver bromide in solutions containing either excess silver ion or excess bromide ion. Since Hellwig 1 performed his detailed study of the cationic and anionic complexes of the silver halides in water, several investigators have sought to determine the identity and the stability constants of the silver(!) bromide complexes in water. The method most commonly employed has been the determination of solubility of silver bromide as a function of bromide concentration. Erber; 2 Vouk, Kratohvil, and Tezak; 3 Berne and Leden;q Lieser; 5 and Anderson, Butler, and Woolley6 were among those using this method. Berne and Leden4 also used silver/silver ion electrode potential measurements, as did Chateau and Pouradier. 7 Anderson and co-workers used a silver ion selective mem- brane electrode to determine the activity of silver ion in their saturated solutions. Kratohvil, Tezak, and Vouk8 calculated stability constants of silver bromide complexes in water by graphical analysis of the data of several investigators. In studies of silver bromide complexes in nonaqueous solvents potentiometric titrations have been used more often than solubility measurements. Some of the solvents which have been used are: dimethyl sulfoxide, 9110 sulfolane, 11 dimethyl formamide, 12 N-methylpyrrolidinone, 12 propylene carbonate, 13 and acetone.1 4 -15
1 2
Silver bromide equilibria have also been studied in mixed solvents. Kazaryan.and ~go~6-l8 used ion-selective membrane electrodes to deter- mine the solubility product constant for silver bromide in various mix- tures of water with methanol, ethanol, acetone, n-propanol, and iso-propanol. Kratohvil and Tezak19 studied anionic complexes in mixtures of water with ethanol, methanol, and acetone. Anderson, Butler and Woolley6 report solubility product constants and stability constants for complexes of one silver ion with one, two, and three bromide ions in ethanol-water, methanol-water, dioxane-water, and acetone-water mixtures of 10, 20, 30, 40, and 50 percent by weight organic component. They used solubility measurements in conjunction with silver ion-selective electrode potential measurements. In the 10 percent acetone system no meaningful value could be obtained for the stability constant of the undissociated silver bromide in solution. It has been suggested that the value of the activity coefficient of that species deviated from the assumed value of unity when small amounts of acetone were present in the solvent mixture, and that further study was needed to determine the nature and/or reality of that deviation.
B. Scope and Purpose This dissertation reports an investigation into the complexation of silver and bromide ions in water and acetone-water mixtures of approxi- mately 5, 10, and 15 percent acetone by weight. The stability constants of the complexes were determined by titrations in which total silver and bromide concentrations were determined by addition and dilution, and silver ion activity was measured through the use of a silver ion-selective membrane electrode and a digital voltmeter. The formal concentration of 3
silver was within the r~nge of 10-a to 10- 6 molal. The formal concentra- tion of bromide was varied between 5 x 10-4 and 0.5 molal. All solutions
were unsaturated with respect to all solutes. The solubility product constant for silver(!) bromide was determined in each solvent system by calibration of a bromide ion-selective electrode with respect to bromide ion from approximately 10- 3 to 10-l molal, and with respect to silver ion over the same range of concentration. An extended Debye-Hueckel expression was used to calculate values of the thermodynamic activity
coefficients. A relative deviation least squares method of data treat- ment was used to calculate the values of the thermodynamic stability
constants for each of the bromoargentate species in water and in each solvent mixture. The purposes of the investigation reported here were (a) to deter- mine if ion-selective electrodes could be used to obtain reliable values of the complex stability constants under the above-outlined conditions, (b) to determine the values of the thermodynamic stability constants for
each silver(I)-bromide species present in water and in approximately 5,
10, and 15 percent by weight acetone-water mixtures at 25°C, (c) to determine the thermodynamic solubility product constant of silver(!) bromide in each solvent system at 25°C, and (d) to explain the observation of Anderson, Butler, and Woolley6 concerning the 10 percent acetone system. II. EXPERIMENTAL
A. Materials All chemicals used were reagent grade. Sodium bromide was from Mallinckrodt Chemical Works, lot XVH. Silver nitrate from Wasatch Chemical Company was certified 99.999 percent pure. Gas chromatographic analysis of the acetone used indicated less than 0.2 percent water and showed no other volatile impurities. Solutions were prepared from either doubly distilled or deionized water which had a specific conductance of less than 1.5 x 10- 6 mho/cm at 25°C.
B. Procedure and Equipment Preparation of Solutions Stock solvent mixtures, 1 molal sodium bromide, and 0.5 molal silver nitrate solutions were prepared gravimetrically using a Mettler P1210 top-loading, constant-load balance which had a capacity of 1200 g and was readable to 0.01 gram. Sodium bromide and silver nitrate reagents were dried at 110°C for one hour and allowed to cool in a desiccator prior to preparation of the solutions. A buoyancy correction was made when calculating the concentration of these sodium bromide and silver nitrate stock solutions. All other solutions used in the experi- ments were dilutions of these, and the buoyancy corrections divided out. All other solutions were prepared and measured using either the Pl210 balance or, whenever possible, a Mettler Pl60 top-loading, constant-load balance which had a capacity of 160 grams and was readable to 1 milligram.
All weighings were by difference.
4 s
Potentiometric Titrations Approximately 130 grams of solvent were measured·. into a 250 milliliter three-necked flask, using theMettlerP160 balance and a polyethylene bottle with a screw~top cap. An Orion Model 94-16A Silver- Sulfide Specific Ion Electrode or an Orion Model 94~35 Bromide Specific Ion Electrode, fitted with a rubber stopper, was placed.in one neck of the flask. In another neck was placed an Orion Model 90-02-00 Double- Junction Reference Electrode, fitted with a rubber stopper~ ·Another rubber stopper was placed in the middle neck. The flask was set in a 25.00 ±0.01°C temperature-regulated water bath and the contents were stirred at a constant rate with a magnetic stirrer~ After.the electrodes had been in the stirred solvent for twenty to thirty minutes, titrations · were begun. Additions of titrant solutions were made through the middle neck of the flask using polyethylene gravimetric burets like those de~ scribed by Butler and Swift. 20 Potential differences were measured using an Orion Ionalyzer Model 801 Digital pH Meter, which was readable to 0.1 millivolt. Readings were taken after the potential reading had remained the same for at least one minute.21,22 After a titration was completed the electrodes were removed from the flask~ washed three or four times with distilled-water.-soaked absorbent tissues, and placed in · stirred distilled water for several minutes.· When not in use, the electrodes were stored in tubes of distilled water in the constant temperature bath. Determinations of the solubility product constant made use of the bromide ion-selective electrode. The electrode was calibrated with respect to silver ion by adding increments of standard silver nitrate solutions to approximately 130 milliliters of solvent and recording the 6 potential difference as described above. The electrode was then cal- ibrated with respect to bromide ion by the same procedure using standard sodium bromide solutions. From the slope and intercept of both cal- ibration curves the solubility product constant was calculated. To investigate the complexation equilibria, the silver-sulfide ion-selective electrode was first calibrated with respect to silver ion using silver nitrate solutions. Another titration followed in which
enough 10-s molal silver nitrate solution was added to pure solvent to
produce the desired concentration of silver ion (between 10- 8 and 10- 6
molal). The resulting solution was then titrated with sodium bromide and the potential differences recorded for each point as described above.
In several experiments further increments of 10-s molal silver nitrate
solution were added at various points during the titration.
C. Calculations
Electrode Calibrations and K5 Both the silver-sulfide and the bromide ion-selective electrodes
employ membranes which are selectively permeable to silver ion. 21 - 24 The potential difference, E, between the ion-selective electrode and a refer- ence electrode in a solution containing silver ion is described by Equation 1,
0 E=(E +E +E.) + (c)log aA = E' + (c)log aA (1) r J . g+ . . g+ where E 0 is the standard half-cell pote~tial of the ion-selective elec-
trode, Er is the potential of the reference, Ej is the junction potential,
and aAg+ is the thermodynamic activity of the silver ion. Since the bro- mide ion-selective electrode is a silver bromide membrane, the cell poten- tial in a bromide-containing solution is related to bromide ion activity, 7 aBr- , and to the solubility product constant, K,s by equation 2.
K ·E = E' + c log ~ = E' + c l_og Ks - c l_og aBr- aBr-
= E" + c'log a Br- (2)
The constants E' and c, or E" and c', were calculated by fitting the electrode calibration data to equation 1 or equation 2 by means of a linear least squares computer program. By combining equations 1 and 2, equation 3 is obtained.
E" - E' = ·• log Ks C (3)
Since it was observed that the· absolute values of c and c' were not identical, Equation 4
E" - E' log Ks = [c(-c')]~ (4)
K was used to calculate 8 from the bromide ion-selective electrode cali- bration data.
Stability Constants of the Complexes
The overall formation constant, or stability constant, for e:L,J.. , the formation of the complex Ag.Br.i-j from silver ion and bromide ion is 1 J defined mathematically by Equation 5.
e.. 1,J (5)
The mass balance equation for total silver present in solution is given 8 by Equation 6,
1•a .... . mibj 1,J = i = I i (6) i=l i=l j=O where c indicates molal concentration, mand b stand for the thermodynamic activity of silver and bromide ions, respectively, and y .. is the activ- 1,J ity coefficient of the species Ag.Br.i-j. Subtraction of CA+ from both ·lJ g_ \ sides of Equation 6 and division by cAg+ yields Equation 7.
C -C Ag(total) Ag+ Fo = ------= I y 1· • CAg+ j=l ,J
+ i I (7) i=2 j=l
If Equation 7 is divided by the activity of bromide ion Equation 8 is obtained.
C - C = Ag(total) Ag+ (8) Fl (CAg+)(b)
A general expression for the .i.' ratios is given by Equation 9. i-1 j-k C . i f3. • Ag(total) - CAg+ 1,J YAg+m b Fk = = I }: (CAg+)(b)k i=l j=l Yi ,J . (9)
Inspection of Equation 7 reveals that if no polynuclear species are present (i never more than 1), F is only a function of the bromide 0 activity and a plot of F0 against log b would be the same curve regardless of the total concentration of silver. If i is greater than unity, one can 9 expect a different curve for each different total concentration of silver.
For each data point (increment in a titration) P0 was calculated from the silver-sulfide ion-selective electrode measurements by
Equation 10,
ET - E C (10) F = 10 -1 0 where c and E are defined by Equation 1, and ET is the potential that would have been observed had all the silver been in the form of unbound silver ion. Since CAg(total) was not the same for each point, ET was calculated for each point by Equation 11.
YAg+ CAg(total) n n = + clog C YAg+ Ag(tota1) (11) 1 1
The behavior of the electrode was found to deviate from that predicted by Equation 1 when the formal concentration of silver was less than S x 10- 7 molal. The extent of the deviation as a function of formal silver con- centration was experimentally determined, and a correction was added to
• ET prior to calculation of P0 The corrections are described in detail in section III-A.
The values of the constants were calculated from the F ei,J.. functions using a general relative deviation least squares curve-fitting procedure described by Anderson and Snow25 and modified according to Kohma~•s26 suggestion. This method allows the small absolute errors in small concentrations to have as much influence in determining the shape of the calculated curve as do the large absolute errors in the large 10
concentrations. It also prevents data points with abnormally large error factors from making extremely large contributions. Equations 12
through 16
M Yi =k!lGk xki (12)
M f. = l G'k xk1· (13) 1 k=l
(f.-y.)2 l(RD)2= l 1 1. (14) . . y. 2 1 1 1
2 Cf. -y.) \ ' ·2 t 1 1 ',(GRD).= ~ ( .. )(f.) (15) l: l l. yl 1
(16)
_ describe this method for the general case, where y is a function of M
constants, Gk, and M variables, xki; Gk represents the constants obtained from a fit of the data and the subscript i represents the ith data point.
Equation 16 represents the set of M simultaneous equations obtained by
minimizing the deviation in Equation 15 with respect to each constant Gn.
This set of equations is solved by matrix methods to obtain values of the
constants Gic- An iteration procedure was used in which fi was first set
equal to Yi, the set of Equations 16 solved, fi calculated by Equation 13, 11 and the set of Equations_ 16 solved again. The last two steps were repeated unt.il the va,lue of F.quation 14 ceased to change.
Activity Coefficients
Equation 17 is an extended fom of the Debye-Hueckel equation 27, which was used to calculate the ionic activity coefficient for the ith ionic species.
(z.) 2Sd(I)-% 1 log y. = ------e---- :(17) 1 1 + Rrd(I)%
Iri Equation 17
S = 2.5057 x 102 (l/€ 3)% (18)
d = (2p)~ (19)
R = 2.060 (1/E)½ (20)
I= ~}:icizf (21)
In Equations 17-21, I is the ionic strength on the molality scale, pis the density of the solvent, £ is the dielectric constant of the solvent, r is the "effective ionic diameter" parameter of the ions in solution, expressed in Angstroms, c. and z. are the molal concentration and the J. 1 charge of the ith ionic species in solution. The value of r used in this wor~ is 3. 047.A., so that (Rrd) = 1. 000 for water as solvent. 28 , 29 Values for the densities and dielectric constants of water and the acetone-water mixtures were obtained by interpolation of literature 30, 31 data, and are given in Table 1. The activity coefficients for the uncharged species, AgkBrk, were assumed to be unity.32-34 12
TABLE1 DENSITIESAND DIELECTRIC CONSTANTS OF WATER-ACETONEMIXTURES
Percent Acetone by We_ight p (glee) e:
0 0.997 78.54 5.18 0.992 76.0 9.87 0.986 73.0 14.90 0.978 70.3
Guggenheim and Turgeon29 proposed a further extension of the Debye-Hueckel equation to correct for differences between measured values of-the activity coefficients of many electrolytes and the values calcu-
lated by Equation 17. Values of activity coefficients for the electrode calibration data in water were calculated using Guggenheim and Turgeon's equation and were used to determine values for the constants E' and c
(Equation 1) and E" and c' (Equation 2). The values obtained for the constants were the same, within experimental error, as those obtained when Equation 17 was used to calculate values of the activity coefficients.
Thus Equation 17 was shown to be no less accurate than Guggenheim and Turgeon's equation in the range of concentrations of sodium bromide and silver nitrate used in this work.
Precision of Least Squares Constants To solve for the constants in a function like Equation 12 by the method of least squares, it is assumed that all deviation of the data points from the "best curve" is due to errors in y only. The standard deviation of the constants obtained is thus dependent on the standard 13 deviation of y. 35 Equation 22 is the general expression used to
(22)
(23)
o 2 = (N-M)-1 t (RD)2 (24) rel l y i calculate the relative standard deviation, orel' of a constant, Gk, ob- tained by the relative deviation least squares fit of N data points. In
Equations 22-24, M, RD, and Gk represent the same quantities as they did in Equations 12-16, Dis the determinant of the coefficients of the constants, Gk, in the matrix of equations represented by Equation 16, and
Pis the cofactor determinant of the element Dkk in determinant D. The standard deviation, o, of the l_ogari thm of Gk was calculated by Equation 25. a relG k (25) a= ---(ln 10)
The Average Ligand Number The average ligand number, n, defined by Equation 26, i (3. • m bj j ]., J 2 y •. CBr(total) - C I - Br- i j 1,J (26) n = = i j CAg(total) i IL . m b 1,J X y .. 14 j 1,J can be obtained from F by the Bodlaender equation, 36 - 38 Equation 27, if 0 14
the absence of polynuclear complexes is assumed.
(27)
If the absence of polynuclear complexes cannot be assumed,. n' calculated
by Equation 27 is not the average ligand number, but is a complex function which bears enough resemblance ton to be useful in qualitative compari- son of systems containing similar complexes. n' was actually calculated for each point by Equation 28,
dfE-1 C ] C og YAg+ Ag(total) ii I = -----:---,:,-----',!!-,---..__ __ ....._ (28) d fag CBr-
which was derived by combin~ng Equations 27, 7 and 1.
D. Error Analysis Solutions were prepared·and dispensed gravimetrically with a pre-
cision of between 0.01 and 0.2%. Potential measurements were precise to ±0.1 .millivolt (mV). In the electrode calibrations, an uncertainty of
±0.1% in the activity of silver ion (or bromide ion) and an uncertainty of ±0.1 mV in the measured potential propagate uncertainties of ±0.2 mV
in E' (or E"), ±0.1 iri c (or c') and•±0.03 in log Ks. Experimental standard deviations of E' (or E") ranged from ±0.1 to ±0.3 mV. Experimental standard deviations of c (or c') ranged from ±0.04 to ±0.09 mV. Values of log K from similar experiments showed a reproducibility s of ±0.05.
Equations 29 and 30, der_ived from Equations 10 and 11, illustrate
the effect of the uncertainties in ET, E, c, and CAg(total) on the 15
F • An relative uncertainty of 0 asterisk(*) preceding a quantity means "uncertainty in" that quantity.
*F 0 --w- = (ln 10) *log P (29) I' 0 . 0
*ET *E (ET -E) A *C 1 1 _!. Ag(total) (30) = -- + -- + --- *c + 2 2 C C c An C Ag(total)
The quantity (An/A ) is the quantity which is the argument of the 1 l_ogari thm in Equation 11. Uncertainties of O. 2 mV. in E, 2 mV in ET,
0.1 mV inc, and 0.1% in cAg(total) propagate errors of 0.8%, 8%, 2%, and
1.6%, respectively, in F0 when evaluated for the worst case, i.e. when
(ET-E) (A /An) the values of the quantities and 1 are maximumvalues encountered in the experiments. II I. RESULTS
A. Behavior of the Silver Ion Selective Electrode
The values of the constants E' and c (Equation 1) obtained by calibration of the silver sulfide membrane electrode in silver nitrate solutions are listed in Table 2. The concentrations of the silver nitrate in each determination ranged from approximately 5 x 10-s to 10- 2 molal. The standard deviations (o) in each constant and in the measured potential measurements are also listed in Table 2. Prior to the calibra- tion for Experiment V, potential measurements were made in solutions containing 1 x 10- 7 to 7 x 10- 7 molal silver nitrate. These measurements were found to deviate as much as 2 millivolts from the values calculated by Equation 1 usi_ng the values of E' and c from Experiment V listed in
Table 2.
TABLE2
E' ANDc VALUESUSED IN EXPERIMENTSI-IX
Wt. Percent Experiment E' C 0 Acetone aE, C OE
I 0 .559.5 .1 58.28 .06 . 1 II 0 576.6 .2 58.22 .OS .1 III 0 566.6 .2 58.26 .07 .2 IV 5.18 562.2 .1 58.08 .03 .1 V 5.18 569.2 .2 57.99 .07 .2 VI 9.87 577.6 .03 57.69 .01 .02 VII 9.87 579.0 .2 58.11 .08 .2 VIII 14.90 572. 7 .4 58.05 .13 .3 IX 14.90 582.1 .2 57.84 .11 .2
16 17 Further evidence of the departure of the electrode response from that predicted by Equation 1 is_ given in F_igure 1, a plot of F0 _against silver ion activity at constant bromide ion activity, where the formal concentration of silver nitrate ranged from 2 x 10- 8 to 8 x 10-7 molal. Equation 1 was used to calculate ET (Equation 10). Due to dilution upon addition of silver nitrate solution to the reaction solution, the concen- tration (and therefore the activity) of bromide was not identical at each data point. For reasons discussed below it was assumed that the last point (upper ~ight corner of F_igure 1) was the most accurate point. The activity of bromide ion at each point was corrected to equal the activity of bromide ion at the last point. A corrected value of F was then calcu- 0 lated for each point from the value calculated by Equation 10 (F0 in calc Equation 31).
F = F - (2) (F ) O 0 (t.:) O calc calc (31)
In Equation 31, bis the experimental activity of bromide, and Ab is the difference between band the activity of bromide at the last point. These values of F are plotted in F_igure 1. The relative cha_nge in b is multi- 0 plied by the number 2 in Equation 31 because, as is shown in F_igures 2-5,
F0 is a function of approximately the square of bromide ion activity.
ii' , the change in log F with respect to log CB , F_igures 2-5 are plots of o r-
_against log CBr- . • Equation '7 indicates that Figure 1 should be a straight line of zero slope if only mononuclear complexes are formed, or a concave-upward curve if mono- and polynuclear complexes are formed. The downward concavity of the plot in Figure 1 indicates the electrode was not 18 ....• • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • ....
.... ·~I -w • N • 4'• • • co • • 0 • o-i • '°0 . • 0 • • II •·- * ....N f -w .. • 0 + ('ll• aj~
II) > 0 f.l.. I I • - o-i * • . bO • ,,-j • ~ • N • .... • .... wI • C1' • ....• * • • * - * * * * ~.... • • • • • • • • • • • • • • • • • • • • • • • • • • • • • * •• - w' I I I I' .... • ' ~' ' "'" Cl) u u g"'" 0 "'+ + + + w w w w N ,0 4' • • 0 • • ... CD "' ... 2.5E+OO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • + • • .. .. • • ♦ +itlf • • + ** • • • • + * ,. ♦ * • • * X xx • • X * X X X X X • • )( • l.9E+OO-- • )( + • • • • X • • + • • • • • • • • • • • • * • l.3E+OO-- • • • • • • .x ·x + • • • • • + * X X • • * X + X • • X X + * • • • +)()( * • • * X + • 7.8E-Cl--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• I I I I I I I I -3.0E+00 -2.4E+OO -1.ae+oo -1.2e+oo
Fig. 2.--fi' vs log c r_: *, Experiment I; +, Experiment II; X, Experiment III. 8 .... 10 20
• • • • • • • • • • • • • • • • • • • • • ..• • • • • • • • .·:.~.- •• 4:? * *... * * - * ·•• • 0 + 0 * • + • - w + * -4" • . + ...• > • I • • .µ • • s:: • * ... • Cl) • • a • • •rt.... • * • Cl) • •. - ~ • • u.:i * • +.. • ... • 0 > * • 0 1-1 • + .µ * •-w s:: * • Q\ Cl) + * • • a + • ... •rt.... * • Cl) + • • ' ~ + * • u.:i * • ... + - -IC I.... * r:Q • u • + * bO • 0 0 • 0 .... * + + • -w Cl) -4' > • * • • + • N IS:: • • I I • * • I + tt) . • • * • bO • •rt • a:i. •·- * + • • • • 0 • Q * + • + • ..• • • • • • • • • • ·• • • • • • • • • • - w I • • • • I• • • I • • - I 0 I I I I • 0 0 0 ... C'l'l 0 Q C, 0 + + + I • w w w w N co ('fl ('('I • • • N ...... CD• 2.3E+CO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • * • • • • • • • • + + +. • + • • • • • • • • • • ♦ • • * • l.7E+OO-:- • ** • • + • • * • • • • * • • * • • • * • • * * • + • • ** • • ♦ * * • 1.oe-co--.• * ♦ + + • • + + + + + • • + + + • • • • • .• • • • • • • • • • • • 3.4e-01--~···························•·····•·····••··••··••···•·········•··••••••·•••·•·I I I I I I I I -3.0E+C0 -2.6E+OC -2.lE+OO -l.6E+00
Fig. 4.--fi' vs log c8r_: *, Experiment VI;+, Experiment VII. ...."' 22
• • • • • • • • • • • • • • • • • • • • • • • • • • • • • • * • * * • ... • ... • * ·-• * • • . • * • >< • 0 1-4 • + • 0 .µ • * • + r:: •-w (I) • + C'C"I 13 • * • •r-i • • • 1-1 * .... (I) • • * + ~ • ' '1,:1 • ..+ .. • * + • ...... 1-4 • * - 1-4 • * + 1-4 • > * + .µ • + r:: • 0 (I) • 13 • + u •r-i • + + 1-1 • •-w (I) + • CD ~ * • ....• '1,:1 • I + * • .. • "' + • I * • 1-1 + al * u - bO * 0 + * l""'4 Cl) + > * u u 1r:: + + I I * - C'rtw ... • LI) * N I bO •r-i *+ ~ • • + * • ... • • • * • u • 0 • + • • • • • • • • • • • • • • • • • • • • • • • • • • *• • • . •-w I I I I Cl' t I I I • 0 0 0 ... N Q 0 0 0 I + + I w w+ w w 0 '° N ~ • • • (0• N ...... 23 respondi_ng as expected. However, the magnitude of the deviation appears to bear a smooth relationship to the total concentration of silver nitrate present. Most of the data taken in Experiments I-IX were in solutions contain~ng less than 8 x 10- 7 molal.total silver. Thus results calculated from those data would be falacious unless a correction could be made for the nonlinear response of the electrode revealed by Figure 1. Whether the correction took the form of a correction in ET or in the measured
F potential made no difference as far as calculation of 0 by Equation 10 was concerned. It was considered simplest and of most general application to determine a correction in ET as a function of formal concentration of silver. Equations 32-35 describe how the correction was determined for each point in Figure 1.
= F (32) 0 1ast
= E +Clog Fo (33) corr
" (.34)
(35)
As before, mis used to denote activity of silver ion. The subscript "corr" means "corrected". The subscript "last" refers to the last data point in F_igure 1 (see p . 18). The effect of the corrections on the
Figure 1 data is demonstrated in Figure 6. l.2E+C5--•••••••••••••••••••••••••••••••••~••••••••••••••!•••••••••••••••••••••••••••••• .XX )XX) X X X X X. • • • • • • • • • • • • * • • • • • 8.6E+C4--. • • • • • • • • • • • • • • • • • • • • • 5.0E+C4--. • • • . • • *• • • • • * • • • • • • * • • • • * • l.4E+C4--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• I I I I I I I I 7.8E-l~ 1 .4E-12 2.1e-12 4.0E-12
Fig. 6.--*, Fo vs aA g+ ; X, F o vs aA g+ ; at a r- = 0.0618. corr corr 8 "'.i:,. 25
The difference, D, between the corrected value of ET and the value of ET calculated by Equation 1 is plotted against log YAg+CAg(total} in Figure 7. It was found that the curve could be approximated by a line with a slope of - 43 mV per log unit where log yAg+cAg(total} was less than -7.05, and by a line with a slope of -23 mV per log unit where logy C was between -7.05 and -6.47. For values greater than Ag+ Ag(total) -6.47, D was zero. These linear approximations were used to calculate corrected values of ET for all experimental data. The corrected ET values were then used in Equation 10 to calculate values for F • 0 It is reasonable to assume that at such low formal concentrations of silver (less than 10- 6 molal) no polynuclear complexes are present; thus Figure 1 should be a straight line of zero slope. Because the curve in F_igure 1 appears to be approaching a straight line of zero slope where the formal concentration of silver nitrate is the greatest (upper right corner), the last point can be considered the most accurate point.
To determine the corrections in ET the last point was assumed to lie on the expected straight line of zero slope. Obviously, more data at higher formal concentrations of silver nitrate are needed to determine the validity of this assumption. Experiments are planned for the near future which will supply data for a more accurate correction. However, the correction determined in this work is probably within 2 mV of the accurate correction--an error which would introduce an error of less than
0.04 in the logarithm of the equilibrium constants calculated from the data. 4.7E+Ol--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • • • • • • • • • • • • • • • • 3.lE+Ol--. • • • • • • • • • • • • • • • • • • • 1. 5E+Cl--. • • • • • • • • • • • • • • • • • • *• -l.6E+QO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • I I l I I I I J -7.8E+CC -7.4E+Q() -7.0E+CJO -6.6E+CJC
Fig. 7.--D vs log yAg+CAg(total): *, data; , 1 inear approximation·s. ---- N °' 27 B. Equilibrium Constants The experimental titration data for this work are tabulated in the Appendix.
Figures 8-11 are plots of the logarithm of F versus the logarithm 0 of the bromide ion activity from the data of Experiments I-IX. In order
to fit Equation 7 to these data all but two to four of the terms in the sums were assumed to be negligible or zero. The best fit of the data
from most of the experiments was obtained when undissociated silver(!) bromide, dibromoargentate(I) ion, and'tribromoargentate(I) ion were assumed to be the only complexes present. The logarithms of the values of the constants obtained and their standard deviations are presented in
Table 3 along with the relative standard deviation of F 0 , which is an indication of how well the data were fit by Equation 7, using the values of the constants listed.
TABLE3 STABILITYCONSTANTS ASSUMING THE PRESENCEOF AgBr, AgBr2-, ANDAgBr3 2-
Exper- Wt% tog tog tog No. of iments Acetone 81 , 1 0 81, 2 0 81 , 3 0 0 Data relFO Points
I 0 s.ss .03 6.99 .12 8.06 .06 0.10 15 II 0 5.58 .04 6.83 .18 8.26 .04 0.14 22 III 0 5.55 .02 7.04 .08 8.24 .04 0.08 24 I, II, III 0 5.56 .02 7.02 .07 8.17 .03 0.14 61 IV 5.81 5.28 .02 6.92 .06 8.08 .03 0.10 31 V 5.81 5.41 .01 6.95 .05 8.36 .03 0~05 26 IV,V 5.81 5.32 .02 7.07 .06 8.04 . 05 0.17 57 VI 9.87 5.70 .03 7.18 .12 8.54 .08 0.12 27 VII 9.87 21 VI, VII 9.87 5.82 .OS 7.15 • 26 8.54 .14 0.27 48 VIII 14.90 5.69 .04 7.86 • 04 8.73 .04 0.13 36 IX 14.90 5.79 .03 7.77 .OS 8.95 .06 0.09 24 VIII,IX 14.90 s. 71 .03 7.87 .03 8.73 .03 0.13 60 6.3E+CO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • +. • ++ • • + • • •* * • • .. • X * . • • • XX* • • +X* • • X)( • 4.ae+co-- • +X • • xx• • • X + * • • X • • X * • • + • • X * * • • X + • • + * • • XXX • 3.4E+OO--• +X *X • • xx • • X + • • +* • • )( • • + * • • X • • • • +* • .x • 2.OE+0O--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• I I I I I I I I -3.4E+CO -2.7E+CCJ -2.lE+CO -l.4E+OO
Fig. 8.--Log F vs log a r_: *, Experiment I; +, Experiment II; X, Experiment III. 0 8 N 00 s.ae+co--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• ~•••••••·••• • *• • * • • • • • •• • • • • • • • • .. +• • • ♦ * • • • 4.5e+oo.:.- • + +*• • • • • +* • • + • • • • • + • • + ++ * • • ++ ** • • + + •• • • + +*. • 3.2E+OO-- • + • • + +* * • • • ♦ * * • • + * • • + * .< • * • • +* • • • • + • • * 1.9e+co--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • l I I I I I I 1 -3 .3E+CC -2.7E+OO -2.lE+QO - l .5E+QO Fig. 9.--Log F vs log a _: *, Experiment IV;+, Experiment V. o 8 r N &O 5.4E+OO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • +. • + • • • • + • • • .....* • • • • ♦ • • • •• • • •• • 4.4E+OO-- • • • + • ••• • • + • • • .. • • + + * * • • + * • • + + * • • . ♦ • • *·* • • + * *** • 3.4E+OO-- • + • • • • • + * * * • • • • • * • • + • • * • • • • • • 2.4E+OO--•••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••·* I I I I I I I I -3.3E+CC -2.ae+co -2.2E+OC -1.7E+ Fig. 10.--Log F vs log a r_: *, Experiment VI; +, Experiment VII. 0 8 c,, 0 6.4E+CO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • ••• • *** • • * • • • • * • • • • • • + ** • • • • *+** • s.1e+co--• +* • • * • • *+ • • * • • ••• • • * • • +•• • • ++ • • • • ++ * • 3.9E+OO-- • -+ * * • • +* * • • + *+ * • • + • • • ••• • • + * • • • • * • • • • • 2.6E+CO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• I I I I I I I I -3.lE+OO -2.6E+OO -2.0E+CC -1.SE+OC Fig. 11.--Log Fo vs log a8 r -= *, Experiment VIII;+, Experiment IX. t,.I.... 32 Interpretation of the data assuming the only complexes present were undissociated silver bromide, dibromoargentate ion, and tetrabromo- argentate ion also yielded constants with reasonably good precision. The values of the constants so obtained are listed in Table 4. Attempts to determine all four constants simultaneously yielded at least one negative value. TABLE4 STABILITYCONSTANTS ASSUMING IBE PRESENCE OF AgBr, AgBr2-, ANDAgBr4 3- Exper- Wt% .tog .tog R.og iment Acetone C1 C1 (1 (1 a11, a1,2 a1,4 relF o. 1· 0 5.52 .07 7.04 .10 8.35 .08 0.26 II 0 5.48 .07 7.35 .07 8.29 .08 0.25 III 0 5.48 .03 7.32 .04 8.59 .07 0.12 ·:1, II,III 0 5.47 .03 7.35 .03 8.27 .06 0.19 IV 5.18 5.23 .03 7.15 .03 8.37 .04 0.13 V 5.18 5.37 ~02 7.18 .04 9.03· .06 0.09 IV,V 5.18 5.27 .03 7.23 .03 8.32 .07 0.18 VI 9.87 5.65 .03 7.44 .OS 9.09 .12 0.15 VII 9.87 6.08 .02 7.13 .10 9.22 .06 0.08 VI, VII 9.87 5.79 .OS 7.43 .09 9.10 .17 0.28 VIII 14.90 5.60 .06 8.01 .03 8.91 .07 0.19 IX 14.90 5.73 .04 7.92 .03 9.57 .10 0.12 VIII,IX 14.90 5.63 .04 8.00 .02 8.92 .06 0.17 Whether Table 3 or Table 4 expresses the correct interpretation of the experimental data is difficult to ascertain. For most of the experiments the relative standard deviation in F is slightly lower in 0 Table 3 than in Table 4. It is reasonable to suppose that the formation of complexes proceeds in a stepwise fashion, i.e., that as the concentra- tion of bromide is increased from zero, associated silver bromide first forms, then dibromoargentate ion, followed by tribromoargentate ion, and 33 finally tetrabromoargentate ion. Vouk, Kratohvil, and Tezak, 3 for exam~ pie, have interpreted solubility data at 20°C in this fashion. They con- cluded that the predominant complex was undissociated silver bromide up • • _4 to a bromide concentration of 4 x 10 , dibromoargentate ion between 0.01 and 0.1 molar bromide, tribromoargentate ion from 0.1 to 0.8 molar bromide, and tetrabromoargentate ion from 0.8 to 2.5 molar bromide. The data of Experiment VII, however, yield a negative value for 81 , 2 when attempting to obtain 81,1, B1,2, B1,3; the uncertainty in 81,2 from Experiment VI and from the combination of Experiments VI and VII is higher in Table 3 than in Table 4. The erratic nature of the data of Experiment VI is demonstrated in Figure 4 (a plot of versus log CB ) and in n' r- Figure 12 (a plot of log aA aB versus log aB ). Figure 12 also shows g+ r- r- that the product of the activities of silver ion and bromide ion from Experiment VII is nearly constant over several values of bromide ion activity, implying that the solution was saturated at those points. If the solution was saturated, the value of the formal concentration of sil- ver, and thus the value of F , was not accurately known at those points. 0 The results of Experiments VI and VII, therefore, do not conclusively speak against the assumption of the presence of tribromoargentate ion rather than tetrabromoargentate ion. It was thus decided to interpret the data in terms of undissociated silver bromide, dibromoargentate ion, and tribromoargentate ion. Data at higher concentrations of bromide are needed to determine the value of 81 , 4. The values of the negative logarithm of the solubility product con$tant determined in each solvent system by calibration of the bromide ion selective electrode with respect to silver ion and with respect to bro ide ion are given in Table 5, along with the values of the logarithm 34 • • • • • • • • • • ...• • • • • • • • • • • • • • • • • • • • • • • •·* + • • * • • * + • • • + • • • • * ·- 1--4 • ... 1--4 • + > • * +J • * i:: "' ... u Q) Q * + + -w -~k= * ,... Q) • * * + .... I z + .. * + * • ... * • 1--4 .,.. + • > +J • i:: * + •·- Q) • -~k= * ... • Q) • p. * ... • 0 >< * u '1J * • + + •-w * N ... • • • * + N I • I k * • CQ • al * • bl) * • .....0 -•- II) • • > + • ,.... k • CQ • al • c;.;, + • g bl) * + • + al< '-' • •-wa> bl) • • 0 • N• ..... • * I • + ' .....N . * bl) -~~ + 0 Q... • • • • • • • • • • • • • • • • *• • • • • • I • • • • I • • • I • I·- r TABLE5 1HERM:>DYNAMICEQUILIBRIUM CONSTANTS OF SILVERBROMIDE C COMPLEXESIN ACETONE-WATER MIXTIJRES Wt. % -log log log log Acetone Ks a 81,1 a 81, 2 a 81,3 a 0 12.23 .04 5.56 .02 7.02 .07 8.17 .03 5.18 12.40 .OS 5.32 .02 7.07 .06 8.04 .05 9.87 12.29 .05 5.68 .03 7.31 .09 8.50 .09 14.90 12.34 .OS 5. 71 .03 7.87 .03 8.73 .03 The change in the numerical value of each of the constants in Table 5 as a function of weight percentage acetone is graphically represented in Figures 13-16. 36 • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • * • • • • • • • • • • • ·- • • .... • 0 • + • w • N • - (I) • • r::: • .... 0 • +' • (I) • 0a, (I) bO .... a, * +' r::: • (I) • 0 • ~ (I) • p.. • 0 • +' • g .c: + bO • .... w •n ,0 (I) • ): • • r-• • • Ill • • > • • Ill • • ::,,=: • • bO • • 0 .... r-4 • I • * • • t') • r-4 • g • 0 • bO • + •n w ~ • r- • - • • m • • • • • • .... • • • • • • • • • • • • 0 • • • • • • • • • • • • • • • • • • • • • • • • • • • • • *• • 5.7E+CQ--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• • •• • • • * • • • • • • • • • • • • • 5.6E+OO--. • • • •* • • • • • • • • • • • • • • • s.1te+co--. • • • • • • • • • • • • • • • • • • * • 5.3E+QO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• I I I I I I I 0 3.7E+O(l 7.6E+CO 1 .2E+O 1 Fig. 14.--log B1, 1 vs weight percentage acetone. ~ -..J 38 • • • • • • • • • • • • • • • • • • • • • • • • • • • • • * - ... Q) 0... 6 -w ~ N Q) u ....• Fig.. 16.--log a1, 3 vs weight percentage acetone. ~ IV. 'DISCUSSION A. Comparison of the Results ·of This Work · with Results of Other Investigations Electrode Behavior Hseu and Rechnitz 39 studied the response of an Orion Model 94-16 silver sulfide membrane electrode at 25°±0.02°C in constant ionic strength aqueous solutions containing from 10- 4 to 10- 1 molar silver nitrate and found the slope (c, Equation 1) of the electrode to be 59.0 mVper log CAg+unit. Mueller, West and Mueller 21 reported that the slope of an Orion Model 94-16 silver sulfide membrane electrode at 25.0°C was 59.11 ±0.03 mV in aqueous solutions of silver nitrate ranging from 1.2 x 10- 7 to 1.2 x 10- 6 molar at an ionic strength of 0.1 molar. They placed the accuracy of the detection of silver ion in that range of concentrations at ±0.2%, which corresponds to an uncertainty in measured potential of slightly more than 0.1 mV. Immediately following the acquisition of the silver sulfide mem- brane electrode used in this work it was calibrated in silver nitrate solutions. The observed slope was 59.1 ±0.1 mV per log aAg+ unit. It was then used in solutions containing both silver nitrate and sodium bromide and left in distilled water for two weeks. Again the electrode was calibrated using freshly prepared silver nitrate solutions. The slope was found to be 58.58 ±0.04 mV per log aAg+ unit. Perhaps.the decrease in the value of the slope with age, apparent from Table 2, is due in part to contamination of the silver sulfide membrane with silver bromide. 40 41 Durst 40 demonstrated that an Orion Model 94-16 silver sulfide membrane electrode responded according to Equation 1 throughout the range of silver ion concentrations from 10- 1 down to 10- 26 molar in aqueous solutions either "in the presence of complexed silver ion" or in solutions in which the concentration of silver nitrate was greater than approx- imately 5 x 10- 7 molar. Measurements in 10- 7 and 10- 8 molar silver nitrate solutions were shown to deviate from Equation 1 behavior. Durst offers no explanation for the deviation, but Ross41 states that the lower limit of detection of the silver sulfide electrode is 10- 8 molar due to "the experimental difficulty of preparing extremely dilute solutions of ions without extensive ionic adsorption on and desorption from the sur- faces of the containi_ng vessels and the electrodes." He also notes that "silver ion levels can be followed down to the 10- 20 level during the course of titrations of silver with a number of complexing and precipi- tating reagents." The conclusion appears to be that the electrode responds in a Nernstian fashion (according to Equation 1) in solutions where the concentration of silver ion exceeds approximately 10- 7 molar £!. where the silver(!) in solution is primarily complexed. The deviations from Nernstian behavior observed in this work were therefore expected in solutions containing silver nitrate at levels near 10- 7 molal and no complexing agent, but were not expected in similar solutions to which sodium bromide had been added. The conclusion must then be that to ex- pect Nernstian behavior from the silver sulfide membrane electrode the total formal concentration of silver(!) must exceed approximately 5 x 10- 7 molar. Figures 2-5 indicate that the electrode responded to changes in silver ion activity in the same manner regardless of the formal 42 concentration of silver. The formal concentration of silver ranges from 1 ~ 10- 8 in F_igure 3 to 6 ~ 10- 7 molal in Figure 2, but the plots of ii' (calculated by Equation 29) versus the logarithlll of the bromide concen- tration are nearly identical. It appears, then, that the corrections made on the data in this study are not oniy justifiable, but perhaps constitute a reproducible method of determining stability constants at formal silver ion concentrations of less than 10- 6 molal. It might also be noted that Berne and Leden4 presented three plots of F against silver ion concentration derived from silver/silver 0 ion electrode measurements in aqueous silver nitrate solutions with con- stant bromide concentrations of 3.00, 4.00, and 4.96 molal, respectively. In the region where the formal concentration of silver was less than approximately 5 x 10- 3 molal, each plot resembled Figure 1 •. They recog- nized this as an indication of electrode malfunction and discarded the affected data. Equilibrium Constants In Table 6, values of the logarithm of the solubility product constant and of.the stability constants obtained in this work are com- pared with literature values. The results of studies in which the ionic strength was corrected to or extrapolated to zero (where concentration equals activity) should represent thermodynamic equilibrium constants: constants in terms of thermodynamic activities. The thermodynamic solubility product constant in water reported in this work compares favorably with the values reported in the literature, which range from 12.23 to 12.30, except the value reported by Berne and Leden4 (12.34) and the value reported by 43 TABLE6 COMPARISONOF TIIE RESULTSOF TIIIS WORKw6m TIIE RESULTSOF OTIIERINVESTIGATIONSa, Weight Percent -log log log log Ionic Acetone IC Strength Method Reference s a1 .,1 a1.,2 f31., 3 ' 0 12.23 5.56 7.02 8.17 0 corrc ISEe this work 12.36 varied Agf 43 12.18 varied Ag 44 12.24 0 corr Solg 38.,cf 49 12.27 0 corr Sol 46,cf 49 12.20 ·...,od Ag 47 12.30 -+ 0 Ag 48 12.4 varied Bo 54 8.72 O corr Ag 7 8.85 0 corr Ag 7,cf 8 12.29 O corr cond so 12.62 8.88 5(NaC104) Ag 4 7.23 9.08 5(NaC104) Sol 4 12.11 .1 (NaCl04) Ag 4 12.10 4.15 7.11 7.95 .l(NaCl04) Sol 4 12.34 4.38 7.34 8.00 0 corr Sol 4 9~18 5 (NaCl04) Sol 4,cf 8 4.30 6.64 8.08 .l(NaCl04) Sol 4,cf 8 8.53 0 corr Sol 51 12.28 O corr 52 12.36 varied ISE 16 12.23 0 corr ISE 53 12.46 6.09 6.54 8.64 0 corr Sol,ISE 6 6.10 7.0 8.68 6~ 6.10 7.0 8.68 61 5.18 12.40 5.32 7.07 8.04 0 corr ISE this work 12.29 0 corr ISE 53 8.20 12.4 varied ISE 16 9.87 12.29 5.68 7.31 8.50 0 corr ISE this work 12.35 O corr ISE 53 9.64 12~52 <4 <5.8 <6.4 O corr Sol.,ISE 6h <4 <7. <8 6 5.26 7.95 8.64 6i 14.90 12.34 5. 71 7.87 8.73 0 corr ISE this work 12.44 ISE 53 16.70 12.6 varied ISE 16 44 TABLE6--Continued FOOTNOTES aAll values were reported at 25°C, except Ref. 7 (24°C). bMuchof this compilation came from Ref. 72. cCorrected to zero ionic strength by some formula, as in this work the Debye-Hueckel equation for activity coefficients was used. dExtrapolation to zero ionic strength. eEMFwith an ion-selective electrode. fEMFwith a silver metal electrode. gSolubility. hRecalculated using Eq. 7. iData corrected according to section III-A of this work and results calculated using Eq. 7. Anderson, Butler, and Woolley6 (12.46). The aver~ge difference between the value reported in this work and the nine literature yalues is 0,07 log units. The value of each of the thennodynainic stability constants in water reported in this study is within the range of values reported in the literature. The numerical values of the thermodynamic solubility product constant determined in the acetone-water mixtures, with the exception of the 5.18% acetone ~ystem, appear to increase as function of acetone content. (See Figure 13.) Other studies 6,1G,S3 indicate a similar increase, but of greater magnitude. The only study reported in the literature presenting values for stability constants of silver bromide complexes in acetone-water mixtures within the range of compositions used in this work is the study of Anderson, Butler, and Woolley. 6 Their results from solubility and ion 45 selective electrode data in saturated silver bromide solutions are shown in Table 6, along with the results obtained by fitting their data with Equation 7 before and after correction of the electrode measurements with the correction determined in this work. Approximately half of their data in 9.64 percent acetone required correction in ET values (Equation 10) of greater than 10 millivolts. The formal concentration of silver in these data was in the same range as in the data of Experiments IV and V of this work. Calculation of 61,1, 61,2, and 61,3 from the uncorrected data of Experiments IV and V gave values of the logarithms of the constants of 4.6, 6.0 and 7.3 respectively. The corrections have had the effect of bringing.both the results of Anderson, Butler, and Woolley6 in 9.64 per- cent acetone and the results of Experiments IV and Vin 5.81 percent acetone more in agreement with the values of the constants determined in other acetone-water mixtures. Figure 17 illustrates the effect of the corrections on the logarithm of 61 , l• B. Suggestions for Further Study The electrode correction needs to be more accurately determined by extending the experiment illustrated in Figure 1 to higher formal concentrations of silver. Through a set of similar experiments at diff- erent constant bromide ion activities the presence of polynuclear com- plexes can be determined. Berne and Leden4 found evidence of polynuclear complexes at formal concentrations of silver greater than 5 x 10- 3 molal, but, due to electrode malfunction, were unable to conclude anything about the presence of polynuclear complexes in solutions containing less than 5 x 10- 3 molal total silver. 6.2E+QO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• ~ • • • • • • • • • • • • +•• • •* • • • 5.4E+OO--.+ • • • • • ® • • • • • • • • • • • • • • • 4.6E+OO--. + • • • • • • • • • • • • • • • • • • @ • 3.8E+OO--••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••••• I I I I I I I 0 3.7E+0C 1.6E+OC 1 .2E+O 1 Fig. 17.--log 81 1 vs weight percentage acetone: *, corrected data, this work;©, corrected data, Ref. 6; .r:,. +, uncorrected, this work; 0, uncorrected, Ref. 6. °' 47 Because the nonlinear behavior of the electrode appears to be re- producible, the method used in this study could be employed in the study of the complexes of otheT "insoluble" silver salts in unsaturated solu- tions. C. Summaryand Conclusions The results of an investigation into the feasibility of determin- ation by ion-selective electrode of the thermodynamic solubility product constant of silver(!) bromide and the thermodynamic stability constants of silver(!) bromide complexes in unsaturated water and acetone-water mixtures have been presented. The values of the solubility product con- stant of silver bromide and the stability constants of undissociated silver bromide, dibromoargentate ion, and tribromoargentate ion in water and in approximately 5, 10, and 15 percent acetone-water mixtures at 25°C have been tabulated and found to compare favorably with literature values. The response of a silver sulfide membrane electrode was found to deviate from the Nernstian behavior in solutions where the formal concentration of silver(!) was between approximately 1 x 10-a and 8 x 10- 7 molal. A correction for the nonlinearity was determined as a function of formal silver concentration, and was applied to the data of this study and to the data of another study in which a similar electrode was used. The results were markedly improved in both cases. APPENDIX On the following p_ages the data, from which the stability constants reported in this dissertation were calculated, are listed, followed by data from which the electrode calibration constants (Equation 1 or Equa- tion 2) for the silver sulfide and the bromide ion-selective electrodes were determined. The definitions of the column headings are as follows: -LOGAGN03 -- the negative logarithm of silver nitrate con- centration (molality). -LOG BR- ACT the negative logarithm of bromide ion activity. ACT, COEF-- the activity coefficient for singly charged ions. EMF-- the observed potential difference (millivolts) between the ion-selective electrode and the reference electrode. CORECTN-- the electrode correction used. (See section III-A.) ET -- ET in Equation 10~ LOGFO -- the logarithm of F (Equations 7 and 10). 0 -LOG AG+ ACT the negative logarithm of the silver ion activity. -LOG BR CONC the negative logarithm of the bromide ion molality. -LOG AG CONC the negative logarithm of the silver ion molality. Page numbers given in the headings refer to pages in the research note- book. It was presumed that listings of the computer programs used for data analysis would be of little value to the reader; thus, although the programs were written by the author, they are not listed here. 48 49 EXP. 1. AGBR I~ hATER (PG 228 PTS 3-23) -LOG -LCG ACT. CCR- LCG -lCG AGN03 BR- ACT CCEF EMF ECTN ET FO AG+ ACT 6.7Vt 3.228 .9722 36.6 7.4 170.9 2.304 c;.1oc t:. 1 c;o 2.816 .95€:C 9.5 1.1 17C.4 2.760 9.57C (:.7c;9 2.557 .9415 -1.1 s.1 169.e 3.C46 9.871 ti• E 13 2.338 .9258 -21.0 8.6 169.1 3.262 10.1oe t.829 2.178 .9121 -3CJ.6 9.1 168.3 3.413 u:.2e1 6.841 2.057 .c;co2 -38.Q c;. (: 161.4 3.525 10 .411 6.849 1.837 .e11te -51.9 1c.o 166.9 3.754 10.662 6.854 1.642 .8478 -64.4 10.4 166.3 3.958 10 .e e~ 6.858 1.soa .8269 · -17.6 10. 1 165.1 4.115 11.llt t.861 1.454 .e11e -e2.1 10 .c; 165.5 4.248 11.1c;l: 6.8t6 1.358 .ecce -c;1.a 11.2 165.C 4.509 11.471 l:.1<;6 1.368 .8026 -82.4 .o 192.e 4.722 11 .c 14 6.199 1.327 .1«;52 -87.9 .c 192.4 4.ElG 11.1ce t:.,03 1.266 .783(: -c;t: .a .o 191.€ 4.c;s2 11.261 t:.210 1.189 .7t:83 -108.3 .c 19(l.c; 5.134 11.458 E-.219 1.102 .75C4 -121.5 .o 189.e- 5.341 11.t:8~ 6.224 1.063 .7424 -127.6 .o 1ac;.2 5.436 11.;c;c 6.230 1.026 .7343 -133.e .c 188.(: 5.532 11.€9(: 6., 33 1.008 .7303 -136.8 .o 188.3 5.578 11.947 t:.237 .983 .725C -14CJ.1 .o 101.c; 5.t:38 12 .c 14 t.24t .933 .7136 -14<;.3 .c 186.S 5.769 12.162 ELECTRCCECALIBRATICN SLCPE= 5.828CE+Ol INTERCEPT= 5.5S50E+02 so EXP. 11. AGBRIN WATER (PG 230 PTS 3-25) -LOG -LOG ACT. COR- LCG -LCG AGNC3BR- ACT CCEF E~F ECT~ ET FO AG+ ACT E:.814 3.256 .9731 53.7 e.1 187.3 2.295 9.121 1:.e20 2.886 .9593 29.5 8.4 186.9 2.703 9.541 6.830 2.563 .9419 9.0 a.a 186.3 3.045 c;.c;oc 6.836 2.448 .9341 1.4 c;.c 185.c; 3.169 10 .03~ 6 .f 50 2.211 .9203 -10.2 9. 5 185.2 3.356 10 .242 t:.8t:l 2.ll:6 .c;11c -16.8 c;.e 184.l: 3.460 10.361 t:.8e1 2.040 .ac;a3 -24.9 lC.4 183.7 3.583 lC • 5 l 1 6.8S8 1.959 .eec;4 -30.3 10. c; 183.(l 3.tl:3 10.l:12 t:.900 1.775 .att:1 -42.5 11.2 182.5 3.864 10.e21 t:.S(3 1.651 .8492 -54.7 11.5 182.l 4.(67 ll.C41 6.911 1.441 .8157 -e1.o 12.1 1e1.2 4.5C3 11.502 6.Sl5 1.365 .ec21 -c;1.5 12.3 1eo.e 4.l:76 11.te1 t.211 1.374 .8038 -11.1 .o 2os.1 4.823 11.135 6.221 1.304 .1c;os -e1.c; .o 208.5 4.<;87 11.311 6.226 1.238 .7782 -91.8 .o 201.e 5.145 11.481 6.232 1.180 .7l:6l: -1<11.e .o 201.1 5.3C5 11.l:52 6.238 1.11a .7540 -110 .1 .c 206.3 5.444 11.eo5 6.244 1.073 .7444 -117.9 .o 205.6 5.557 11.929 t:.251 1.024 .7339 -126.2 .o 204.e 5.l:86 12 .en 1 6.258 .983 .7250 -133. 1 .o 204.1 5.792 12.1c;c ·t.zte .932 .713t -1'4 l • 7 .o 203.2 s.c;23 12.33e 6.t.76 .894 .7050 -148.6 .c 202.4 6.028 12.45(: t:.2€6 .855 .6c;t:l -155.5 .o 201.5 6.131 12.~75 ELECTRCCECALlBRATlCN SLCPE= 5.8220E+Ol INTERCEPT= 5.7E60E+02 51 EXP. II 1. AGBR It\ '-ATER (PG 233 PTS 3-2el -LOG -LOG ACT. CCR- LCG -LCG AGNC3B~- ACT CCEF E~F ECTJ\ ET FO AG+ ACT 6.7«;8 3.318 .9749 53.5 1.1 177.7 2.131 6.94C 6.8Cl 3.013 .9646 29.8 7.9 177.4 2.533 9.34«; E.ece 2.713 .9507 1c.o e.2 176.9 2.e65 9.694 6.817 2.495 .9314 -4.2 e.s 176.4 3.C99 9.~44 6.826 2.355 .9272 -13.8 e.c; 175.c; 3.256 10.114 6.830 2.307 .9233 -16.7 c;.o 175.7 3.302 10.166 6.S35 2.250 .9185 -20. 7 c;.2 175.4 3.366 10 .2 36 6.846 2.155 .9C99 -26.6 c;.s 174.c; 3.458 10 .345 6.6 51 2.118 .9063 -2c;.1 c;. 7 174.6 3.4c;7 lQ.3<;1 6.856 2.091 .9CJ36 -30. 8 9.8 174.4 3.523 10 .422 6.8E3 2.045 .ec;ac; -33.7 l EXP. IV. s.1e PCT ACETCNE fPG 238 3-33) -LOG -LOG ACT• CtR- LCG -LCG AGt.C3 BP- ACT CCEF EMF ECTN ET FO AG+ ACT 7.740 3.256 ·.9718 39.5 43.4 155.3 1.993 9.746 1.144 2.955 .9605 19.3 43.7 155.1 2.339 10.101 7.747 2.719 .9530 9.1 44.0 155.C 2.513 10.281 7.751 2.671 .9459 .1 44.3 154.c; 2.656 10.431 7.755 2.586 .94C6 -4.5 44.6 154.e 2.744 10.525 1. 760 2.486 .:;.933e -.11.0 44.9 154.7 2.es3 10 .t43 7.7t5 2.413 .9284 -15.7 45.3 154.6 2.932 10. 7 2~ 1.110 2.342 .9227 -20.4 45.6 154.5 3 .o 11 10 .e If 1.111 2.261 .9162 -24.5 46.Q 154.3 3.079 10 .894 1. 1.e9 2.165 .9G67 -31.9 46.7 154.1 3.202 11.c 34 7.796 2.113 .9Cl5 -34.8 47.2 153.9 3.250 11.c91 7.8C9 2.041 -.8937 -40.4 47.9 153.7 3.342 11.199 7.818 1.994 .8885 -44.2 48.4 153.5 3.404 11.274 7.S26 1.959 .8S43 -47.6 48.8 153.4 3.4l:O 11.34( 1.837 1.918 .8794 -s2.c 49.4 153.2 3.532 11.425 7.846 1.887 .8755 -55.5 49.e 153.C 3.590 11.494 1.848 1.767 .esc;s -65.7 5(l.3 1s2.e 3-.763 11.677 1aso 1.665 .844E -ec.s 50.7 152.7 4.Cl5 11.c;3c; 7.854 1.563 .8283 -92.9 51.2 152.5 4.226 12.161 7.858 1.447 .e084 -un.s 51.9 152.3 4.473 12.424 7.8El 1.399 .1sc;e -113.5 52.2 152.2 4.574 12.532 7.Et5 1.326 .7859 -123.0 52.7 152.0 4.735 12.1c~ 1.810 1.259 .1121 -131.9 53.2 1s1.e 4.€85 12.et1 1.875 1.201 .7t09 -139.7 53.7 151.6 5.Cl6 13 .o lC 1.se1 1.148 .7497 -147.0 54.2 151.5 5.139 13.145 1.ees 1.109 .7414 -152.3 5't. 7 151.3 5.228 13.243 1.ec;1 1.010 .7329 -157.8 55.l 151.2 5.320 13.34!: 6.ac;2 1.012 .7334 -151.5 12.1 lt6.e 5.480 12.501 6.ec;4 1.057 .12c;c; -153.6 12.e 166.6 5.514 12.544 t.ac;1 1.035 .7251 -156.7 12.c; 166.4 5.St:3 12.toc 6.905 .989 .7147 -164.4 13.2 165.~ 5.686 12.737 ELECTRCCECALIBRATICN SLCPE= s.aceoe•o1 INTERCEPT= 5.6220E+02 53 EXP. v. 5.18 PCT ACETCt\E. (PG 241 3-28) -LOG -LOG ACT. COR- LCG -LCG AGNC3eR- ACT CCEF E~F ECTf\ ET f(j AG+ ,c1 7.602 3.274 .912~ 41.7 31.4 165.0 2.121 9.742 1.6Ct 2.980 .9ElE 23.4 37.8 164.c; 2.440 10 .C63 7.ElQ 2.811 .953E 12.9 38.1 164.e 2.620 10 .2 5C l.Et5 2.635 .'9431 2.1 38.5 164.1 2.ec3 10 .444 7.619 2.543 .9378 -3.6 3e.e 164.6 2.c;oo 10 .547 1.624 2.458 .c;:31e -a.a 3c;. 1 16't.4 2.c;ea 10.642 7.630 2.374 .9253 -14.l 39.5 164.3 3.C77 10. 7 4C 1.639 2.272 .9167 -20.1 40. l 164.l 3.187 10.€64 7.645 2.216 .9116 -24.3 40.4 164.C 3.247 10.<;32 7.654 2.146 .9048 -2c;.o 41.0 163.8 3.325 11.022 7.EEl 2.0CJ6 .8997 -32.6 41.4 163.7 3.384 11.092 1.t11 2.040 .ac;36 -37.3 42.0 163.~ 3.462 11.182 7.679 2.004 .aec;s -40.8 42.4 163.3 3.520 11.24c; 1.6e1 1.968 .8854 -44.6 42.8 163.2 3.583 11.322 7.6S5 l.9'35 .ee11t -41.9 43.2 163.0 3.t:37 11.381 7.7(2 1.909 .8783 -5Cl .3 43.l: 162.S 3.676 ll.43~ 1.112 1.875 .e13c; -53.1 44.1 162.7 3.121 11.4c;2 7.724 1.840 .8694 -56.l 44.7 162.5 3.770 11.554 1.7;1 1.820 .8(:67 -58.2 45.l 162.4 3.€(4 11.597 1.734 1.709 .8511 -73.9 45.5 162.2 4.072 11.€75 7.737 1.602 .834€ -et.I 4(:.0 ll:2.C 4.279 12 .c c;4 7.739 1.523 .8218 -c;s.9 4E.4 161.c; 4.446 12.21c 7.742 1.470 .8126 -101. 7 46.e 161.e 4.544 12.375 7.746 l.3'il .1c;e2 -112. 8 47.3 161.f lt.132 12.516 7.749 1.337 .1ee1 -119.6 47.t 161.5 4.€47 12.fc;c; 1.152 1.2c;o .11ec; -126.l 48.0 161.4 4.c;57 12 .e 1e ELECTRCCECALIBRATICN SlCPE= 5.7990E+Ol INTERCEPT= 5.6S20E+02 54 EXF. VI. 9.81 PCT ,\CETCNE (PG 254 3-32) -LOG -LCG ACT. CCR- LCG -LCG AGNC3e~- ACT CCEF Ef'F ECTN ET FO jG+ ,\Cl 6.E57 3.278 .9109 4e.s 9. 1 19C.4 2.459 9.33(1 t.ato 3.023 .9l:12 32.2 c;.3 190.1 2.138 c;.l:15 6.8l:4 2.842 .9526 21.3 C,. 5 1ec;.c; 2.c;22 c;.eo1 6.868 2.719 .c;457 13.8 9 .6 189.6 3.048 9. 940 6.ae1 2.453 .9274 e.1 10.1 188.9 3.134 10 .c 4 7 6.EE4 2.4C8 .9238 3.6 10 .2 188.1 3.209 l(J.121 t:.aee 2.357 .91C,5 -1.2 10 .4 1ee.s 3.288 10.213 l:.8<;3 2.304 .914€ -5.6 10 .5 188.2 3.360 10 .292 t:.ec;a 2.252 .9100 -lC .2 10.1 188.C 3.435 l0.374 6.9Cl: 2.1c;2 .9C41 -1s.1 10. c; 187.l: 3.514 10.463 6.910 2.163 .9011 -17.6 11.1 un.4 3.554 10.50<; 6.915 2.132 .ac;1a -20. 1 11.2 187.2 3.593 10.555 t:.<;22 2.086 .ec;29 -23.6 11.4 186.8 3.l:48 10.t:1«; 6.9~5 2.025 .8E58 -28.4 11.e 186.3 3.122 10.1cc; 6.<;38 1.964 .8783 -34.2 12.c 186.1 3.818 10.812 (:.4i41 1.906 .81cc; -39.6 12.1 1as.e 3.c;o1 10.c;oc; 6.945 1.856 .8642 -44.2 12.3 185.l: 3.983 10.c;91 6 .S48 1.8C9 .8577 -51.9 12.4 185.3 4.112 11.121 6.952 l. 768 .8518 -57.1 12.6 185.1 4.198 11.22c 6.957 1.121 .8446 -63.3 12.e 1a1t.e 4.301 11.331 t:.963 l.669 .8365 -69.3 13.0 184.5 4.399 11.43c; 6. t; 66 1.642 .8322 -12.2 13.1 184.3 4.445 ll.4 EXP. \J11. 9.87 PCl ~CETOt\E CPG 257 4-24) -LOG -LOG ACT. CCR- LCG -LCG AGNC3 eR- ACT COEF EMF ECTN ET FO AG+ ACT 6.343 3.199 .9t81 43.8 .Q 209.6 2.es3 c;. 2 lC E.347 2.919 .95l:5 25.8 .o 209.1 3.154 9.520 6.352 2.743 .9471 14.2 .o 208.5 3.344 9.720 6.3~8 2.581 .9369 4.0 .o 201.9 3.5C9 c;.e9~ 6.3t6 2.444 .926S -4.6 .c 201.2 3.644 10 .C43 l:.!76 2.321 .9164 -12.2 .o 206.3 3.760 10 .174 6.3€5 2.235 • 9(J8 3 -11.1 .o 205.5 3.€31 10.2se 6.3S4 2.111 .c;c1c; -21.e .o 204.9 3. so 1 10.33c; 6.4C3 2.112 .8951 -25.5 .c 2C4.2 3.c;52 10.403 6.4C6 2.021 .ee61 -30.8 .o 203.7 4.035 10 ·" 94 6.410 1.957 .e11s -35.3 .o 2C3.2 4.105 10.572 t:.417 1.852 .8636 -41.9 .2 202.1 4.209 10.tes 6.422 1.791 .8550 -46.0 .5 202 .3 4.214 10.764 l:.430 1. 711 .S431 -52.0 .e 2c1.c; 4.369 1c.e12 l:.439 1.634 .8308 -59.5 1.1 201.3 4.488 11.coe 6.451 1.558 .818C -t:c;.2 1.5 200.6 4.t44 11.1a2 6.452 1.526 .8124 -73.l 1 • 6 200.5 4.708 11.2sc 6.455 1.452 .1see -e3.o 1. c; 200.1 4.€72 11.425 6 -~ 58 1.393 .7E75 -Sl.3 2.1 199.8 5.009 11.571 6.463 1.317 .7725 -Hll.5 2.4 199.3 5.177 11.752 6.468 1.253 .7592 -110.6 2.1 1c;e.c; 5.326 11.c;1~ ELECTRCCECALIBRATICN SLCPE= 5.8110E+Ol INTERCEPT= 5.7<;01E+02 56 EXP. VIII. 11e.c;o PCT ACETCf\'E (PG 249 3-'t 1) . -LOG ·LOG ACT. CCR- LCG -LCG AGNC3eR- ACT CCEF EMF ECTN ET FQ AG ♦ .ict 6.879 3.104 .9f27 24.1 c;. 7 182.l 2.112 9.t07 6.ee3 2.858 .95G9 9.3 9.9 1e1.e 2.'372 9.87E 6.8E6 2.723 .9431 1.3 10 • 1 181.6 3. 105 10 .c 17 t:.ec;o 2.599 .9348 -6.7 10 • 2 181.3 3.238 Hl-.15e 6.8<;4 2.513 .9284 -11.9 10.4 181.0 3.323 10.250 t.ec;a 2.437 .9223 -16.9 10.l: 180.8 3.4()5 10.339 t:.9C4 2.358 .9154 -21.6 lC .e 18C.5 3.481 lC .423 l:.9C9 2.294 .9G94 -26 .2 lQ .c; 1.eo.2 3.~55 10.5Cf 6.917 2.213 .9C13 -32.5 11.2 11c;.e 3.l:57 10.61 c; 6.923 2.163 .896C -36.6 11.4 179.5 3.722 10.693 l: .c; 29 2.114 .ec;os -4() • 5 11.6 179.1 3.184 10.164 6.937 2.064 .8846 -44.0 1 l. c; 178.f 3.e3e 10 .e ze 6.951 1.990 .8755 -49.0 12.3 178.l 3.913 10.c;21 6.965 1.928 .8673 -53.5 12.1 177.5 3.979 11.0C6 6.«;69 1.912 .8651 -55.1 12.e 177.3 4.C03 11 .c 36 6.976 1.833 .8539 -74.8 13.1 176.c; 4.336 ll.38C 6.919 1.aoo .8490 -78.9 13.2 176.7 4.403 11.452 ELECTRGCECALIBRATION SLCPE= 5.8C5CE+Ol lt(TEPCEPT= 5.7270E+02 6.S19 1.8co .a4c;c -73.0 13.2 185.7 4.459 11.soc; t:.9€4 1.747 .8408 -79.9 13.5 185.4 4.514 ll.63~ 6.9S9 1.709 .8347 -es.1 13.c; 185.3 4.f62 11.12c; 6.c;c;3 1.676 .8293 -ec;.6 14.2 185.2 4.737 11.e11 1.oc1 1.618 .8194 -<;7.4 14.e 185.C 4.868 11.956 1.c 10 1.569 .8108 -103.9 15.3 184.8 4.977 12.c1e 1.,22 1.s11 .ecc1 -112.0 16.l 184.5 s.112 12.23C 7.024 1.460 .7<;04 -118.7 16.4 184.4 5.225 12.352 7.026 1.415 .1e1s -124.8 1~.1 184.3 5.329 12.462 7.029 1.370 .7723 -131.0 11.1 184.2 5.434 12.575 7 .c 34 1.300 .1516 -140.6 11.6 184.0 5.596 12. 7 5C 7.036 1.268 .7508 -145.0 17.9 183.9 5 .6 70 12.€3( 7.039 1.230 .1424 -150 .2 18.2 1e3.e 5.151 12.<;2(: 7.042 1.202 .7362 -154.2 18.5 183.7 5.825 13.COC 7.C46 1.168 .7283 -159.0 18.9 183.6 s.c;o5 13.cac; 1.0 lt9 1.138 .7214 -163.2 1 c;• 2 183.5 5.976 13.167 1.054 1.096 .7116 - 16C,. 2 1c; .1 183.3 6.016 13.21€ 6.1t2 1.098 .1121 -162.6 10.0 190.6 6.C88 12.c;c;e t.1t5 1.081 • 7C8C -165.0 10. 1 190.lt 6.127 13.042 t:.767 1.062 .7036 -lt1.7 10 .3 190.2 6.170 12.cc;c 1;.110 1.048 .70C3 -lt9.5 10 .4 1c;o.1 6.199 12.122 1;.112 l .031 .6c;t:1 -112.1 10 .s 189.c; 6.240 13.17C i:. 115 1.011 .6928 -173.8 10.6 189.7 6.267 13.201 ELECTRCCECALIBRATICN SLCPE= 5.8010E+Ol If\'TERCEPT= 5.8140E+02 57 EXP. IX. 14.c;O PCT ACETONE (PG 252 3-26) -LOG -LCG ACT. CCR- LCG -LCG AGNC35R- ACT CCEF EMF ECTJ\ ET FO AG+ ,er 6.816 2.851 .9506 20. l 8.4 195.0 3.C23 9.S62 6.821 2.690 .9410 9.6 e.t 194.7 3.1c;c; 10.047 6.823 2.621 .9364 s.2 e.1 194.5 3.273 10.125 6.828 2.521 .9290 -1.3 8.9 194.2 3.380 1O.24C 6.832 2.432 .9219 -6.7 9.1 193.9 3.469 10.331 6.838 2.355 .9152 -11.6 c;. 3 193.6 3.5'i8 10.42~ 6.844 2.278 .901e -11.1 c;. 5 193.3 3.f37 10.524 6.8~1 2.2oe .9coe -22.4 9.7 192.9 3.723 10 .(: 19 6.et2 2.122 .8914 -3C • 7 10. 1 192.4 3.e57 10.1e:c; 6.870 2.069 .8852 -35.8 Ul .3 l CALIBRATIONOF AGS ELECTRCOEIN hjTER, PG 228 -LCG ACT. -LOG .aCT. -~ CONC CCEF E"'f AG CCNC CCE:F Ef' F 4.314 .c;c;1c; 3C1.3 2.575 .9442 4C1.S 3.964 .c;eac 328.0 2.121 .910 5 433.~ 3.693 .c;e31 343.9 • 1.65<, .esc;e 45c; .c 3.437 .9783 358.7 1.363 .8172 475.C 3.283 .9142 3(:7.6 1.1ac; .7886 4€4.C CALIBIUT ICf\ CF AGS ELECTRCCEIN hATER, PG 230 -LCG ACT. -LCG .ACT. .a(; CONC CCEF Ef'F AG CCf\C CCEF Ef' f 4.476 • c;c;33 315.4 2.542 .9422 - 427.2 4.027 .c;eee 341.7 2.22e .c; 197 444.c; 3.69S .c;e31 361.1 l.95~ .€945 4tC .C 3.403 .9774 378.0 1.541 .e437 4€2.(: 3.185 .9712 3C,(l.6 1.305 .eoao 495.(l 2-:a~ CALIBRATICNOF BR ELECTRCOEIN h.4TER, PG ..~ -LOG ACT. -lCG .ac1. AG CONC CCEF ff,l,f AG CCNC CCEF Ef' F 4.477 .9c;33 3C4.3 2.483 .9384 41c;.3 4.170 .c;c;os 322.3 2.211 .9184 434.t 3.829 .9Sl:0 342.3 1.e21 .e1c;1 45(:.4 3.492 .979(: 3(:}.8 1.533 .8427 472.1 3.252 .9733 375.6 1.191 .7889 4«;0.4 3.071 .<;674 3et.o 2'2 ~ CALIBRATICNOF BR ELECTRCOEIN ti..aTER, FG ..~ -LCG ACT. -LOG ACT. ER CONC CCEF Ef'F BR CCNC CCEF ff,lf 3.307 .974c; 3c; • 5 l.95t .ec;4, -3S.S 2.997 .9f:ltf: 21.1 1.750 .8713 -~1.1 2.719 .9521 4.7 l.43f .e2es -t:9.8 2.429 .c;34c; -12.3 1.1ae • 1885 -83.5 2.229 .c;1,;c; -24.2 .994 .7537 -94.i 2.112 .c;c91 -31.0 59 CALIBRATICI\CF AGS ELECTRCCEIr-. hllTER, PG 2:3 -LCG ACT. -LOG ACT. AG CONC CCEF Eflf AG CCNC CCEF EffF 4.437 .c;c;3c 3C7.5 2.48E .93€7 42C.C 4.112 .c;ec;e 326.6 2.131 .c; 114 440.C 3.836 .c;e61 342.6 1.857 .8838 455.2 3.5C6 • cnc;c; 361.8 1.533 .8426 412.S 3.264 .S131 316.l 1.351 .El53 'if2.f 3.075 .9t75 386.8 CALIBRATICf\CF AGS ELECTRCOEIN WATER,FG 234 -LCG ACT. -LCG ACT. AG CONC CCEF Et'F AG CCI\C CCEF El'F 4.428 .c;c;2c; 3CS.3 3.056 .c;668 3€7.7 4.100 .c;e CALIE!RATIONOF BR ELECTRCOEIN 5.18 PCT ACETCI\E, PG 235 -LCG ACT. -LCG ACT. AG CONC CCEF Ef'-F AG CCNC CCEF Efl f 4.418 • c;c;25 2c;e.a 2.305 .c;22s 42C .5 4.071 .c;eec; 31S.3 2.003 .ec;4c; 437.2 3.767 .9€43 337.1 1 • 7c;c; .e11c; 448.4 3.474 .c;1e2 354.2 1.562 .84C4 461.2 3.373 .9756 35c; .9 1.357 .eoee 472.1 3.275 .9721 3(:5.l: 1.295 .1c;e s 475.3 3.149 .96€7 372.0 60 CALiePATICNCF BR ELECTRCDEIN S.18 PCT ACETCf\E, PG 235 -LOG ACT. -LCG ACT. eR CONC CCEF El'F BR CCNC CCEF E.,F 3.171 .c;E94 lt.O 1.999 .eCJ44 -52.7 2.860 .CJ510 -2.3 1.16c; .e6e1 -t.s.e 2.512 .9376 -22.9 1.525 .8349 -1c;.c; 2.364 .9271 -31.6 1.312 .€015 -c;1.s 2.176 .c;11t -42.5 1.14c; .7731 -ICC.It 2.100 • c;c4t -4t.c; CALIEJcATICNCF AGS ELECTRCOEIN s.10 PCT ACETCI\E, PG 238 -LCG ACT. -LCG ACl. AG CONC CCEF E~F AG CCf\C CCEF E,i F 4.317 .c;c;21 3C6.e 2.484 .c;:357 4 H .2 3.946 .<;€72 33}.c; 2.164 .9105 434.1 3.629 .9€17 3So.e 1.19c; .e11c; 454.~ 3.311 .9757 365.4 1.49c; .8310 47C.4 3.252 .912C 372.6 1.2ec; .1cn5 4el.f 3.136 .«;tB2 379.3 CALIBl CAllfPATION OF BR ELECTRCOEIN 5.18 PCT iCETC"E, PG 240 -LCG ACT. -LCG ~Cl• fP CONC CCEF Et'F BR CCf\C CCEF Ef'F 3.140 .9l:S4 15.8 1.eo2 .e122 -tl.~ 2.758 .9520 -t:.5 1.sec .e42e -13.1 2.4l:7 .934t -23.s la36l: .e 103 -es.1 2.290 .c;214 -33.6 1.1e2 .7790 -c;t:.C 2.148 .c;ocn -41.8 1 .oc;o .1~23 -1c1.1 2.015 .ec;t:1 -4Cj.4 61 CjllBRATJCN CF AGS ELECTRCOEIN 5.18 PCT /lCETCf\E, PG 241 -LCG ACT• -LOG ACT. AG CONC CCEF E~F AG CCNC CCEF Ef'f 4.383 .c;c;22 313.9 2.210 .9146 43e.c; 3.934 .c;e10 340.5 2.011 .8957 44c;.c; 3.644 .9820 357.4 ' 1.695 .8587 4l:7.(l 3.428 .c;11c 31C.O 1.1tt:e .e2t:3 1i1c;.2 3.260 .9723 379.6 1.21c; .795<; 4Ec;.c 2.so1 .9373 422.4 CALiePATICN CF /lGS ELECTRCOEIN c;.e1 PCT /lCETCf\E, FG 255 -LOG ACT. -LCG ACT. /lG COt-.C CCEF Ef'F AG co"c CCEF Ef,lf 4.458 .9c;24 31c;.o 3.36~ .c;73c; 2e2.e 4.136 .9e9C 33e.c 2.473 .9313 433.1 3.931 .9€l:2 350.4 2.154 .9047 45C • E 3.7t:l .c;e33 3EC.2 1.905 .e1e1 4 l:4 • It 3.528 .97€3 373.5 1.623 .8411 479.l: CALifPATICN CF AGS ELECTRCOEIN c;.e7 PCT ACETC"E, FG 257 -LOG ACT. -LOG ACT. AG CONC CCEF E~F AG cc"c CCEF Ef'-F l:.812 .c;c;c;~ H!4. c; 3.691 .9819 364.1 6.292 .c;c;c;1 211.1 3.423 .9755 319.1 t:.121 .9989 221.4 2.502 .9334 432.1 4.475 .9Ci2E 31€.4 2.213 .9102 1i1ie.~ 4.121 .9sec; 33c;.1 1.904 .e11Cj 465.1 3.906 .c;ese 3~1-5 1.543 .8293 484.3 CALIEPATICN CF JIGS ELEClRCOE IN c;.e1 PCT /lCETC~E, FG 259 -LCG ACT. -LCG JICT. AG cor.c CCEF Efl F AG CCf\C CCEF Eflf 4.450 .c;c;23 32C.1 2.470 .9311 433.c; 4.137 .c;ec;1 33€.5 2.14~ .<;038 1t~1.c; 3.848 .c;e4e 355.4 1.951 .884C ~t2.J 3.601 .9EOC 3ECj.7 1.747 .ese3 47~.; 3.412 .9152 ?e.c.s 62 CALie~ATICN Cf BR ELECTRCDEIN c;.e1 PCT /JCETCJ\E, FG 261 -LOG ACT. -LCG ACT. •G CONC CCEF Ef',F AG CCNC CCEF er,,F 4.408 .9c;2c 315.5 2.oe1 .ec;eo 448.C 4.110 .9ee1 333.1 1.874 .e744 4 5c:;• 1 3.702 .c;e21 356.8 '1.684 .e4c;7 470.C 3.482 .'3171 3E<;.6 1.441 .8133 482.Cj 2.341 .<;212 434.0 C/JLlflUTICN 'cf BR ELECTRCDEIN c;.e1 PCT ACETCJ\E, FG 261 -LCG ACT. -l..OG ACT. BR CONC CCEF Efl!F ER CCNC CCEF er,,f 3.201 .9f87 38.4 1.63€ .8433 -,2.2 2.776 .95G3 13.3 l.4Clt .ec11 -E5 .• 4 2.442 .c;291 -6.2 1.213 .7848 -72.E 2.168 .9C6C -22.1 1.062 .7456 -.64 .1 2.051 .ec;43 -28.8 C/JLieRATIONCF eR ELECTRCDEIN 14.<;0 PCT ACETCJ\E, FG 245 -LCG ACT. . -LCG ACT. j(; CONC CCEF Eftf AG CCf\C CCEF Efllf 4.287 .9<;0 3 3Ci5.7 2.420 .9 ..237 412.7 4.057 .9€74 319.l 2.141 .8986 1i2e.c 3.877 .9845 32Ci.6 1.830 .et2t 445.C 3.604 .97C,C 345.5 1.594 .e2c;3 457.E 3.3c;e .c;735 357.4 1.372 .7929 469.4 3.238 .c;t e4 3Ef.6 CAllfRATICN CF 0R flECTRCOE IN 14.<;0 PCT ACETCf\6, FG 246 -LCG ACT• -LCG ~CT. e~ CONC CCEF E~F BR CCf\C - CCEF Ef'f 3.te9 .9ff(: ·1~.9 2 .03c; .es1e -~~.E 2.920 .9!52 -2.3 1 .6ee· .8432 -1<:.C 2.658 .94()(: -1e.2 l.4l:E .eoc;2 -ee.i 2.481 .9264 -2e •.9 1.26(: • 113 c; -c;c;.c; 2.306 .9142 -39.4 1.074 .1312 -1 lC. ~ 2.145 • ec;c;c -49.l .932 .1oe2 -11e.~ 63 CJLIBIUTICN CF BR ELECTRCDEIN 14.«;0 PCT ACETCf\'E, PC: 246 -LCG ACT. -LCG ACT. AE CONC CCEF Efi#,f AG CONC CCEF Ef'f 4.382 .c;c;13 302.4 2.41€ .c;235 414.J 4.102 .c;eec 31e.1 2.13€ .ec;e3 429.5 3.717 .9Sl5 341.C '1.877 .8687 443.t 3 .5Cl5 ~9765 353.2 t.631 .€347 45E.e 3.322 .9112 363.4 1.337 .78E7 472.C 3.187 .9tE5 371.0 CALIBPATICN CF AGS ELECTRCOEIN 14.«;0 PCT ACETC"E, PG 249 -LCG ACT. -lCG ACT. j(: CONC CCEF Ef'F AG CCt\C CCEF Ef'F 4.393 .9c;14 316.7 2.403 .9223 431.4 3.978 .9€62 341.0 2.12c; .€913 44E.5 3.771 .9S26 353.2 1.873 .€681 460 .4 3.507 -.9766 368.1 1 .• 635 .€354 413.C 3.303 .9706 3eC.6 1.423 .eo1s 4E3.1 3.136 .9646 3~0.2 CALl8PATICN OF AGS ELECHlCOE IN 14.«;0 PCT ACETCti.E, FG 252 -LOG ACT. -LCG ACT. AG CONC CCEF EMF AG CCJ\C CCEF Eflf 4.180 .c;ec;c 331.(: 3.181 .c;(:(:3 3c;f.3 3.8CJ7 .9€49 354.5 2.453 .92(:3 431.3 3.733 .c;e1e 364.2 2.163 .c;ooe 453.~ 3.558 .91.1E 374.6 1.esc; .e102 4EE.C 3.370 .c;121 385.5 CALlfPATICN CF AGS ELEC.TRCOEIN 14.c;o PCT ,ceTCttE, FG 253 -LCG ACT. -LCG jCT. JG CONC CCEF Efl.F ~G cc"c CCEF . EflF 4.427 .9«;17 325.0 2.410 .c;230 440.e 4.121 .9ee3 343.1 2.111 .c;o22 453.E 3.ac;q .c;e4c; 35E.l 1.q24 .e744 4E7.~ 3.727 .9Sl7 3tt.l 1.11c; .€559 475.2 3.534 .q773 377.3 1.s1e .e261 4es.c; 3.4CO .C,73t 3es.o REFERENCES 1. Karl Hellwig, Z. Anorg. Chem., 25, 157 (1900). 2. W. Erber, Z. Anorg. Chem., 248, 32 (1941). 3. V. B. Vouk, J. Kratohvil, and B. Tezak, Arhiv. Kemi., 25, 219 (1953). 4. E. Berne and I. Leden, Z. Naturforsch, 8a, 719 (1953). 5. K. H. Lieser, Z. Anorg. Allg. Chem., 292, 97 (1957). 6. K. P. Anderson, E. A. Butler, and E. M. Woolley, J. Phys. Chem., 77, 2564 (1973). 7. H. Chateau and J. Pouradier, Science Ind. Phot., 23, 225 (1952). 8. J. Kratohvil, B. Tezak, and V. B. Vouk, A:thiv. Kemi., 26, 191 (1954). 9. C. Jambon and J. C. Merlin, C. R. Acad. Sci., Ser. C, 272(2), 195-7, (1971). 10. F. Gaizer and I. Szalay, Magy. Kem. Poly., 78(4), 488-91, (1972). 11. M. DellaMonica, U. Lamanna, and L. Senatore, Inorg. Chim. Acta, 2(4), 12. M. Breant, C. Buisson, M. Porteix, J. L. Sue, and J.P. Terrat, J. Electroanal. Chem. Interfacial Electrochern., 24(2-3), 409-15 (1970). 13. J. Courtot-Coupez, and M. L'Her, Bull. Soc. Chim. Fr., 1969(2), 675-80 (1969). 14. D. C. Luers, R. T. Iwamoto, and J. Kleinberg, Inorg. Chem., 5(2), 201-4 (1966). 15. E. L. Macker, Rec. Trav. Chim, 70, 457 (1951). 16. N. A. Kazaryan and E. Pungor, Anal. Chem., 60, 193 (1972). 17. N. A. Kazaryan and E. Pungor, Anal. Chim Acta, 51, 213 (1970). 18. N. A. Kazaryan and E. Pungor, Acta Chim (Budapest), 76(4) 339-43 (1973). 19. J. Kratohvil and B. Tezak, Arhiv. Kemi, 26, 243 (1954). 64 65 20. E. A. Butler and E. H. Swift, J. Chem. Ed., 49, 425 (1972). 21. D. C. Mueller, P. W. West and R.H. Mueller, Anal~ Chem., 41, 2038 (1969). 22. R. G. Bates and M. Alfenaar, "Activity Standards for Ion-Selective Electrodes," in Ion-Selective Electrodes, Ed. by R. A. Durst, Nat. Bur. Stand. (U.S.), Spec. Puhl. 314, pg. 191-214. 23. "Instruction Manual for Silver and Sulfide Ion Acitivity Electrode," Orion Research Inc., Cambridge, 1968, pg. 4. 24. "Instruction Manual for Halide Electrodes," Orion Research Inc., Cambridge, 1967, pg. 16. 25. K. P. Anderson and R. L. Snow, J. Chem. Ed., 44, 756 (1967). 26. T. P. Kohman, J. Chem. Ed., 47, 657 (1970). 27. H. S. Harned and B. B. Owen, "The Physical Chemistry of Electrolyte Solutions," 3rd Ed., Reinholt Pub. Corp., New York, N.Y., 1958, pp. 63, 508. 28. G. N. Lewis and M. Randall, "Thermodynamics," 2nd Ed., Rev. by K. S. Pitzer and L. Brewer, McGraw-Hill, New York, 1961, pp. 345-6. 29. E. A. Guggenheim and J.C. Turgeon, Trans. Faraday Soc., 51, 747 (1955). 30. D. F. Othmer, D. Friedland and Schiebel, Ind. Eng. Chem., 44, 1872 (1952). 31. P. S. Albright, J. Am. Chem. Soc., 59, 2098 (1937). 32. J. H. Jonte and D.S. Martin, J. Am. Chem. Soc., 74, 2052 (1952). 33. K. P. Anderson, E. A. Butler, D. R. Anderson, and E. M. Woolley, J. Phys. Chem., 71, 3566 (1967). 34. K. P. Anderson, E. A. Butler and E. M. Woolley, J. Phys. Chem., 75, 93 (1971). 35. C. Daniel and F. S. Wood, "Fitting Equations to Data," John Wiley and Sons, Inc., New York, 1971, pp. 13-18 .. 36. L. Johansson, Coord. Chem. Rev., 3, 293 (1968). 37. G. Bodlaender, "Festschrift fuer R. Dedekind," Braunschweig, 1901. 38. G. Bodlaender and R. Fittig, Z. Physik. Chern., 39, 597 (1902). 39, T.-M. Hseu and G. A. Rechnitz, Anal. Chem., 40, 1054 (1968). 66 40. R. A. Durst, "Analytical Techniques and Applications of Ion-Selective· Electrodes," in Ion·Selective Electrodes, Ed. by R. A. Durst, Nat. Bur. Stand. (U.S.), Spec. Puhl. 314, pp. 402-3. 41. J. W. Ross, Jr., "Solid State and Liquid Membrane Ion-Selective Electrodes," in Ion Selective·E1ectrodes, Ed. by R. A. Durst, Nat. Bur. Stand. (U.S.), Spec. Publ. 314, .P&• 77. 42. J. Bjernnn, G. Schwarzenbach and L. G. Sillen, "Stability Constants, Part II: Inorganic Ligands," The Chemical Society, London, 1958, pp. 112-114. 43. H. M. Goodwin, Z. Phys. Chem., 13, 577 (1894). 44. A. Thiel, Z. Anorg. Chem., 24, 1 (1900). 45. G. Bodlaender and W. Eberlein, Z. Anorg. Chem., 39, 197 (1904). 46. A. E. Hill, J. Amer. Chem. Soc., 39, 68 (1908). 47. K. Hass and K. Jellinek, Z. Phys. Chem., A, 162, 153 (1933). 48. B. B. Owen and S. R. Brinkley, Jr., J. Amer. Chem. Soc., 60, 2233 (1938). 49. L. G. Sillen, "Solubilities of Silver Halogenides," Report July 1953 to the Analytical Section, I.U.P.A.C. so. J. A. Gledhill and G. M. Malan, Trans. Faraday Soc., 49, 166 (1953). 51. K. S. Lyalikov and V. N. Piskunova, Zhur. Fiz. Khim., 28, 127, 595 (1954). 52. J. A. Gledhill and G. M. Malan, Trans. Faraday Soc., SO, 126 (1954). 53. P. s. Smith and E. M. Woolley, (unpublished results from a new method). 54. R. Ruka and J.E. Willard, J. Phys. Chem., 53, 351 (1949). THERMODYNAMIC STABILITY CONSTANTS OF SILVER BROMIDE COMPLEXES IN WATER AND WATER-ACETONE MIXTURES DETERMINED BY ION-SELECTIVE ELECTRODE MEASUREMENTS Arthur Lee Cummi_ngs Department of Chemistry Ph.D. Degree, August 1974 ABSTRACT Thermodynamic equilibrium constants for silver(!) bromide com plexes were determined by ion-selective membrane electrode potential measurements in water and in mixtures of water and acetone of approxi mately 5, 10 and 15 percent acetone by weight. Values of the solubil ity product constant were determined in each solvent system by calibra tion of a silver bromide membrane electrode with silver nitrate solu tions and with sodium bromide solutions. Stability constants of silver(!) bromide complexes were determined using a silver ion-selective, silver sulfide membrane electrode in unsaturated solutions in which the formal concentrations of silver(!) and of bromide ranged from approximately 8 Ix 10- to 7 x 10-7 and 10- 3 to ro- 1 molal, respectively. Non-Nernstian electrode response was observed in these solutions. The magnitude of the corrections made for the non-linearity was found to be a function of for mal silver concentration. Comparison of the values of the stability con stants obtained with those reported in the literature substantiated the validity of the corrections. The data were interpreted by assuming the · presence of AgBr, AgBr2, and �gl3r3.2�. . ·._ · -