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The rate of reaction between bromoacetate and thiosulphate ions in dodium chloride and sea water media Dodimead, Allan John 1954

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THE RATE OF REACTION BETWEEN BROMOAGBTATE AND THIOSULPHATE IONS IN SODIUM CHLORIDE AND SEA WATER MEDIA by ALLAN JOHN DODJJEAD A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE in the Department of Chemistry We accept this thesis as conforming to tbe standard required from candidates for the degree of MASTER OF SCIENCE Members of the Department of Chemistry THE UNIVERSITY. OF BRITISH COLUMBIA October, 1954 ABSTRACT Rate measurements of the thiosulphate-bromoacetate reaction have been made in sodium chloride and sea water media at 15.3°, 25.2° and 32.0*'C and in magnesium chloride at 15.3°C, Rates are more simply de-pendent on the concentration and kind of cations in the media than upon the ionic strength. The rate constants are approximately 6% higher in neutral solutions of sea water than in sodium chloride media and the higher rate constants are attributed to the presence in sea water of the cations, K*, Ca"r and Mg** , The energy of activation is 16,000*100 cal./mole and the entropy of activation is-6,6 e,u, for the reaction done in both media. The convergence of the rate curves in magnesium and sodium chloride media at high concentrations may be explained by the presence of undissociated species. The magnitude of the correction to the observed rate depends upon the concentration of the cations and the formation con-stant for the undissociated species. This explanation has been used to account for the decrease in rate in sea water that has been made alkaline. / ACKNOWLEDGEMENT The author wishes to express his gratitude to Dr. M. Kirsch for his generous assistance and encourage-ment throughout the preparation of this thesis and also to the Defence Research Board of Canada for their f i n -ancial assistance. TABLE OF CONTENTS Page INTRODUCTION 1 Historical 1 Theoretical 2 EXPERIMENTAL 9 Preparation of Solutions 10 Rate Measurements 12 Absorbanee Measurements 15 RESUETS 15 Rate Measurements 15 Effect of Varying pH in Sea Water Media 28 Effect of Varying pH in Sodium Chloride Media 35 Summary of the Effect of pH on the Rate Constants 35 Absorbanee Measurements 36 Sodium Bromoacetate and Sodium Chloride Solutions 36 Sodium Bromoacetate and Magnesium Chloride Solutions 36 Sodium Bromoacetate and Sea Water Solutions 38 Sodium Thiosulphate and Sodium Chloride Solutions 38 Sodium Thiosulphate and Sea Water Solutions AO Sodium Thiosulphate and Magnesium Chloride Solutions 4-0 Sodium Thiosulphate and Barium Chloride Solutions 44 Activation Energy 44 Entropy of Activation 47 Page DISCUSSION 4,9 Calculation of Formation Constant 52 SUMMARY 5 5 SUGGESTIONS FOR FURTHER WORK 57 BIBLIOGRAPHY 58 LIST OF FIGURES Page Figure 1 . The influence of ionic strength on the velocity 7 of ionic reactions. Figure 2 . Rate constants in various media at 1 5 , 3 ° C . 2*4 Figure 3 . Reaction rates in various media at 1 5 , 3 ° C , 2 5 Figure A . Rate constants in various media at 25,2°G« 2 6 Figure 5 o Rate constants in sodium chloride and sea water 2 7 media at 3 2 . 0 ? ' 0 , Figure 6 , Effect of pH and reaction concentration on rate 3 0 constants in sea water at 2 5 , 2 ° C , Figure 7 , Graphical determination of rate constants in 3 3 sea water media at 15,3° C . Figure 8 , Graphical determination of rate constants in 3 A sea water media at 25,2°G, Figure 9 . Absorbance of solutions of sodium bromoacetate 3 7 sodium chloride, and a mixture of these components* Figure 1 0 , Absorbance of solutions of sodium thiosulphate, 3 9 sodium chloride, and a mixture of these components. Figure 1 1 , Absorbance of solutions of sodium thiosulphate, Al sea water, and a mixture of these components. Figure 1 2 , Absorbance of solutions of sodium thiosulphate, A 2 sea water, and a mixture of these components. Figure 1 3 • Absorbance of solutions of sodium thiosulphate, 4 3 magnesium chloride, and a mixture of these components. Figure 1 A . Absorbance of solutions of sodium thiosulphate, A 5 magnesium chloride, and a mixture of these components. Figure 1 5 , Absorbance of solutions of sodium thiosulphate, 4 6 barium chloride, and a mixture of these components. Figure 1 6 . Rate constants corrected for the possible presence 5 4 of undissociated species. LIST OF TABLES 1. Elements present in solution in sea water. U. Specific rate constants for the bromoacetate thiosulphate reaction in sodium chloride media at 15.3*0. T T T , Specific rate constants for the bromoacetate thiosulphate reaction in sea water media at 15.3°C, HT. Specific rate constants for bromoacetate thiosulphate reaction in magnesium chloride media at 15.3°C. Y. Specific rate constants for the bromoacetate thiosulphate reaction in sodium chloride media at 25.2°C8 U . Specific rate constants for the bromoacetate thiosulphate reaction in sea water media at 25.2°C< ITU. Specific rate constants for the bromoacetate thiosulphate reaction in sea water media at 25.2°C. imi . Specific rate constants for the bromoacetate thiosulphate reaction in sodium chloride media at 32.0^0. TX. Specific rate constants for the bromoacetate thiosulphate reaction in sea water media at 32.01° C, X. Effect of pH on the specific rate constant in sea water media at 15.3°C» • "XI. Activation Energy. INTRODUCTION HISTORICAL The fi r s t consideration of the ..importance of sea water dates from the time of Thales, who reached the conclusion that a l l things were of water, a l l things come from water, and to water a l l things return* Aristotle many year8 later, made many interesting observations of marine l i f e * During the sixteenth century sea voyages did much to stimulate man's curiosity concern-ing marine phenomena* With the beginnings of the f i r s t scientific societies in the seventeenth century and the periodicals sponsored by them, the interest in oceanograpnic things became very pronounced* Robert Boyle, for one, per-formed a number of experiments on sea water and was the f i r s t to estimate the total solids by evaporation of known quantities of sea water and ignition of the residues* The f i r s t extensive study of the composition of sea water was initiated by Forehammer, Natterer and Dittmar ( l ) * Although Fore hammer an-alyzed a large number of samples, his investigations were not complete, since he did not determine certain of the abundant elements* Natterer made more detailed analyses, but i t was the work of Dittmar that laid the foundation for the present knowledge of the composition of sea water* Dittmar made careful determinations on seventy-3even water samples, representative of a l l oceans, which had been collected on the voyage around the world on H«M«S, "Challenger11* Dittmar's work showed that there were no significant regional differences in the relative concentrations of the major dissolved constituents in sea water* More recent work, using the recent advances in analytical chemistry give results which are not much different from those of Dittmar's* Although these small differences are significant, the importance of Dittmar's work is that i t did show the constancy of the ratios between the concentrations of the major constituents* Now along with more accurate results many different elements have been detected and determined in sea water. The challenge in this field is great since the determination of the chemical nature and concentration of the dis-solved substances may be difficult owing to a high concentration of some of the substances and the presence in very minute quantities of others, or the complex nature of the dissolved materials may require a specially developed technique to determine the concentrations of any constituent• However forty-nine ele-ments are known to occur in sea water. Table I shows the relative amounts of the more abundant elements in order of decreasing concentration in sea water© They are listed as the amounts of the individual elements which occur in water of chlorinity 19 % 0 • The complexity of sea water and the fact that i t serves as a medium for a host of reactions, many of biological importance, suggests that our understanding of such processes may be improved by permitting relatively simple reactions to take place in sea water media. Since the rate of ionic reactions is known to be affected by the nature of the medium, the study presented here involved the investigation of such a reaction permitted to occur in sea water media, THEORETICAL The study of the effect of ions on the activity of other ions has proved to be of great complexity even in simple mixtures of electrolytes and studies have shown that each mixture exhibits individual behaviour even in solutions of low concentrations. In the consideration of sea water the prob-lem is even more complex, since so many ionic species are present, some in very large concentration. Instead of attempting evaluation of activities of individual com-ponents in sea water, i t may be possible to gain some insight into the mutual effect of each of the solutes on one another by using sea water as a medium -3-TABLB 1 ELEMENTS PRESENT IN SOLUTION IN SEA WATER Element mg/kg CI - 19.00 /oo millimoles/litre CI - 19.00 700 /Chlorine 18980 548.30 Sodium 10561 470.15 Magnesium 1272 53.57 Sulphur 884 28.24 Calcium 400 10.24 Potassium 380 9.96 Bromine 65 0.83 Carbon 28 2.34 Strontium 13 0.15 Boron 4.6 0.43 Silicon 0.02-4.0 0.0007-0.14 Fluorine 1.4 0.07 Nitrogen (comp.) 0.01-0.7 0.001-0.05 Aluminum 0.5 0.02 Rubidium 0.2 0.002 Lithium 0.1 0.014 Phosphorus 0.001-0.10 Barium 0.05 Iodine 0.05 Arsenic 0,01-0.02 Iron 0.002-0.02 Manganese .001-0.01 for ionic reactions. A theoretical treatment of ionic reactions by Bronsted (2,p.71-7A) accounted for variations in the velocity of ionic reactions with varying ionic strength. The theory can be presented thus. If we consider the reaction AZ a + B2* ^ X ^ A + Z f t ) * products where A and B are reacting ions with charges Z^, ^ respectively, and X is an intermediate complex with charge (ZA+28). We may write for the fi r s t equilibrium constant Where the a's are activities f's are activity coefficients C's are concentrations. The reaction rate is proportional to the concentration of the intermediate complex X, so that or " d t -fx where k, is independent of concentration. The factor ^ft , the ratio of the activity coefficients of the reactants and the complex, is the correct-ion to the classical rate expression valid for ideal reactants to account for the changes in velocity resulting from a change in the nature of the medium* Such environmental changes are produced by a variation in the concentration of the reactants or by the addition of a chemically inert substance* The observed rate constant, k, is given by d.t * for a second order reaction free of any side or interfering reactions. On substitution into Equation (l) we have k B k.' (2) Since the unstable nature of X precludes its isolation, and the instability of A and B in the presence of one another prevents the determination of £j\ and-ft for the exact conditions under which the reaction proceeds, i t is necessary to attempt a theoretical calculation of these quantities* Such a treatment was provided by Debye and Huckel (3) to account for the deviations of the activity coefficients of ions from unity* The expression developed gives the activity coefficient of an ion as a function of the charge on the ion and the ionic strength of the solution and is given by where j*. is the ionic strength of the solution and is given by Lewis and Ran-dall as equal to 2. 51 C J L Z . ^ where c^ is the concentration of the ions expressed as molality 1^ is the charge on the ions A and B are two constants dependent upon the dielectric constant and temperature of the solution* (Li, is the distance of closest approach of an ion to a central ion* The validity of this expression has been tested by a number of investigations and i t has found to hold for very dilute solutions of electrolytes* More recent extensions of the Debye-Huckel expression have been suggested to ac-count for variations at higher ionic strengths, however the mathematical development is complicated and beyond the scope of this work* From this expression we are able to obtain In . 2 . Z A Z « — F — ~ - (3) -6-Substitution of Equation (3) into Equation (2) yields l a k = In W0 + * 7 A Z F T * £ i f the solution is dilute, in which case the denominator will approximate unity«, or log \ = logk'„+- Z L T - A Z ^ A ' ^ J I where A ' = A - - 0.5o$) for water at 25°C or I09 V. = I03X0 + l . o \ Z A " L ^ ^ j I . ( A ) Equation (4) shows that the rate of an ionic reaction is dependent upon the charges of the reacting species and the ionic strength of the solution* Numerous experiments have been made to confirm the validity of this expression. A summary of the results is given in the well known Livingston diagram O O (Figure l ) . The data show that a straight line of predicted slope is obtained when the logarithm of the specific rate constant is plotted against the square root of the ionic strength for solutions of low ionic strength. The effect of higher-valent ions on the reaction rates has been studied (5, 6)• One reaction considered was the bromoacetate-tblosulphate reaction. S 20 3 =+ BrCHgCOO" Sz03CH2C00= 4 Br" It has been shown (7) that this reaction obeys the Bronsted theory up to an ionic strength of 0.03 in solution of the sodium and potassium salts of the reactants. However at the same ionic strength in the presence of salts of barium, calcium and magnesium, the reaction is faster than in solutions con-taining sodium and potassium salts, with barium having the largest effect on the rate constant. S t i l l larger rate constants axe obtained when the reaction occurs in the presence of small concentrations of trivalent salts such as lan-thanum chloride. For this reaction the increased velocity has been explained by the increase in number of collisions of the reacting ions resulting from a 7-0.60 -0.A0 1 1 1 : 1 0 0.10 ; 0.20 0.30 FIGURE 1. THE INFLUENCE OF IONIC STRENGTH 09 TBE VELOCITY OP IONIC REACTIONS, 1. 2(Co(NH3)cBrr%Hg^* 2H*0 2(Co(HH3Jt HiOr" * HgBr* (Bimolecttlar). No foreign salt added. 2. CHiBrCOO'v SiO^CHiSxOj COO" + Br-as the sodium salt. No foreign salt added. 3. (H^B-COOCiHs j"*- 0B^HtO*CO3* CjEsOH. 4.. HxOa*2H** 2Br~ -* 2Hz0*Br,. 5. (Co(NH3)5 B r O H " -* (Co(HH3)OHf * Br". -8-local increase in their concentration in the neighborhood of the inert Ions of opposite sign. Therefore in the presence of the divalent and trivalent ions there will be an increased accumulation of negative reactant ions around . the cations which wil l provide an enhanced opportunity for collision between the reacting species. More recently the effect on the rate constants of ionic reactions has been studied by the addition of inert salts (8). Such work led to the conclusion that for reaction between ions of the same sign, the effect is caused almost exclusively by the concentration and character of the salt ion of the charge opposite to that of the reactants, and therefore the rate is not simply dependent upon the ionic strength of the solution. The effects can be interpreted in terms of an ion-association-ccnstant and specific rate con-stents for.the associated and non-associated reactants. A quantitative treatment of the ion-pair association has recently been attempted by Wyatt and Davies (9). They investigated the reaction SZC>3 * BrCH^ COO" — > SiOjCH^COO^ 4 Br~ in the presence of barium and calcium salts of the reactants. They took into account a l l the association products that may be present in significant amounts and assumed that the separate reactions proceed independently by molecular mechanisms and contribute additively to the measured velocity* Using dis-sociation constants obtained from other work (10) for BaSj.03 they showed that the increase of observed rate constants in solution of divalent ions compared with monovalent ions at similar ionic strengths may be explained i f the reaction BaSt03 -* BrCHjCCXT — » S^ OjCH,,C00= + Br * Ba** contributes importantly to the measured rate. In the present study, the behaviour of sea water as a reaction medium has been compared with that of sodium chloride solutions by determining the rate at which an ionic reaction occurs. Any difference in the rate constants -9-in these two media is then due to the presence in sea water of ions other than sodium and chloride in sufficient concentration to affect the activity of the reacting species. The ionic reaction chosen for these studies is the thiosulphate-bromoacetate reaction. This reaction has been studied extensively by La Mer (6, 7, 11), Von Kiss and Vass (12) and others (13). It has the following advantages over some other reactions that might have been used. It is second order and free from side reactions. Its rate is susceptible to convenient and accurate measurements. No change in the ionic strength of the solution takes place as the reaction proceeds to completion. The results obtained for this reaction in dilute solutions are in excellent agreement with the Bronsted equation The rate of hydrolysis of the sodium bromoacetate has been determined (11) and is of the order of 1x10 b litre moles' ndn^ ' at 2$°C and can be neglected up to a temperature of 50°C relative to the substitution of thiosulphate for bromide. EXFERU'ENTAL Briefly, the rate measurements were made in the usual way. Measured volumes of standard solutions of sodium bromoacetate and sodium thiosulphate were mixed in the appropriate medium. After a measured time interval at con-stant temperature, an aliquot was withdrawn and quenched with excess iodine* The volume of iodine was than a direct measure of the quantity of thiosulphate that replaced the bromine in the bromoacetate. Various solutions were required for these experiments. Details of their preparation are given below. -10-PEBPARATION OF SOLUTIONS An solutions were prepared using carbon-dioxide-free demineralized water* Sodium Thiosulphate* Reagent grade sodium thiosulphate was used to prepare solutions of a concentration of four times that required in the reaction and were standardized against a standard potassium iodate solution (14, p.334-335)* The solutions kept their titre for a week. Recrystallization of the sodium thiosulphate had no effect on the stability of the solutions nor on the velocity of the reaction. The sodium thiosulphate solutions, used in the back-titration of the excess iodine added to stop the reaction, were made from a 0.1M solution when required. The concentration of these solutions was four-tenths the in-i t i a l concentration of the reactants. Potassium Iodate. Standard potassium iodate solutions were prepared (l4»p«334-335) using analytical grade potassium iodate. Suitable aliquots were used to standardize the sodium thiosulphate solutions. Iodine. A solution containing 4$ potassium iodide was made approximately 0.1N in iodine and from this,dilute solutions, four-tenths the concentra-tion of the reactantSjwere prepared when needed. Starch. Freshly prepared 1$ starch solutions were used. Sodium Bromoacetate. B.D.H. bromoacetic acid was purified by recrystal-lization from benzene and kept dry over phosphorus pentoxide in a vacuum desiccator. The bromine content of the acid was determined by hydrolyzing the acid in an alkaline solution, acidifying, precipitating the bromide with silver nitrate, and weighing the silver bromide. The bromine content was found to be better than 99»5# of theoretical. Purification of the acid by distillation under reduced pressure did not improve the product* The bromoacetic-= acid was weighed and transferred to a lOO-ml. flask. It was then neutralized with standard sodium hydroxide solution using phenolphthalein as an indicator and diluted to a specific volume* The normality calculated from the weighed quantity of acid and that calculated from the titration of the acid always agreed to within better than 1# of one another* From this solution 100 ml* of sodium bromoacetate of the strength required in the reaction was prepared* Sodium Hydroxide* An approximately 0*2-0*3 N solution of carbonate-free sodium hydroxide was prepared and standardized against a weighed quantity of analytical grade potassium acid phthalate using phenolphthalein as the in-dicator* Sodium Chloride. Analytical grade sodium chloride was used in the prepara-tion of a l l solutions. Solutions were made up in volumetric flasks to twice the concentration required in the reaction* Magnesium Chloride. Reagent grade magnesium .carbonate was dissolved in an equivalent amount of hydrochloric acid and made up to volume in a 500 ml. volumetric flask* Sea Water. Sea water solutions were obtained by filtering sea water obtained from Burrard Inlet and concentrating under reduced pressure to twice the con-centration required in the reaction. To lower the pH of the sea water after this process, dry ice or nitric add was added. To standardize the sea water a 25 ml. portion was diluted with 25 ml. of distilled water and 10 ml. a l i -quots titrated with standard silver nitrate solution using potassium chromate as indicator. Duplicate titrations agreed to within 0.02 ml. using a 25 ml. burette* Silver Nitrate. Reagent grade silver nitrate (A9 g.) was dissolved in 1 lit r e of distilled water and standardized against standard sea water using potassium chromate as indicator. - 1 2 -RATB MEASUREMENTS A"n rate measurements of the reaction SaO, • BrCH zCOO-—>CHjCOCP * Br" were carried out in 250-ml. ground-glass stoppered Erlenmeyer flasks fitted with a side arm with a capacity of about 50 ml. as shown in the accompany-ing diagram. With this type of vessel i t was possible to keep the reactants separated while they came to thermal equilibrium and also minimized any time error in mixing at the start of the reaction. A reaction mixture consisted of 25 ml. of each of the reactants, sodium thiosulphate and sodium bromoacetate, and 50 ml. of the reaction medium, which in these studies was either sodium chloride, magnesium chloride or sea water solutions. Twenty-five ml. of sodium thiosulphate solution was pipetted into the side arm of the flask, while to the flask was added 25 ml. of sodium bromoacetate and 50 ml. of the reaction medium. Neglecting the volume change on dilution, the resulting concentration of the reactants was one-quarter of the original concentrations, while the concentration of the reaction medium was one-half the original concentration of the medium. The flasks were then placed in a constant temperature bath and the solutions allowed to attain the temperature of the bath. The reaction was started by simply mixing the reactants in the flask and recording the time. Ten ml. aliquots were taken at suitable intervals and the reaction was quenched by transferring the -13-aliquot to a flask containing an excess of iodine. The final time of the reaction was taken when half the aliquot was delivered by the pipette. The time for half delivery was considered to be five seconds^ Excess sodium thiosulphate was added and the mixture then titrated to the starch-iodine end point. The end point was corrected by appropriate blanks. At least five aliquots were taken at various times from each reaction mixture for the determination of the rate constants. Unless the reaction was particularly fast, i t was possible to follow three separate reaction mixtures at the same time. Instead of standardizing the iodine solutions and sodium thio-sulphate solutions used in quenching the reaction, blank mixtures were prepared which differed from the reaction mixtures in the substitution of distilled water for the sodium bromoacetate used in the latter. A blank mixture was prepared for each reaction mixture. A 10 ml. aliquot of each was titrated with iodine and agreed within 0.03 ml. of iodine. The mean was taken as the i n i t i a l concentration of the reactants. The i n i t i a l con-centration of the reactants and the amount used up during the course of the reaction were expressed in terms of ml. of iodine. A comparison was made between the sodium thiosulphate and iodine solutions used in these titrations• Unless otherwise stated the pH during the reaction was between 6.5-7.5 since results were reproducible in this range. The pH of the reaction mixtures was determined with a Beckman Model 0 pH meter. If the pH change of the reaction was to be followed, two identical reaction mixtures were prepared. One was followed kinetlcally, while the other was followed for pH variation. One sample was taken from the latter and the rate constant determined to ensure that there was agree-ment in the rate constants for the two reactions. - H -Rate measurements were made at 15*3°, 25*2°and 32.0° C in solutions of MgCl^p NaCl and sea water at ionic strengths varying from 0.04 to 1.0* The temperature of the bath was controlled to±0.02°C. Most of the reactions were done with reactant concentrations of 0.01M. The ionic strengths of the sodium chloride and magnesium chloride reactions were calculated from the formula. The ionic strength of the sea water solutions was determined from the expres-sion given by Lyman and Fleming (15). JJ. a 0.00U7 * 0.03592Cl7oo • 0.000068(01 Too)* where CI yoo is the chlorinity expressed in grams of chloride per 1000 gm. of sea water. The value obtained from this expression was added to the con-tribution of the reactants to obtain the total ionic strength of the solution. The total concentration of the cations in the sea water was calcu-lated using only Na*, K**, Mg^and Ca** , since the concentrations of the other ions are small enough to be neglected. The specific rate constant was determined from the second order rate expression to. v o.-* J Where a = b = i n i t i a l concentration of the reactants t « time for the reaction k » specific rate constant x = amount reacted at time t a~x - concentration of reactants at time t The rate constants were calculated by plotting the ratio as abscissa against the time, t, as ordinate. The slope of the line is equal to ka and therefore k - specific rate constant - slope 6L -15-ABSORBAHGS MEASUREMENTS The results of the rate measurements, i t was thought, might be clarified by an examination of the absorption spectra of the reactants and reaction mixtures under various conditions. In particular, i f any complexes were formed amongst the various ionic species present, measurements of the absorbancies of the individual components and the reaction mixture should permit their detection. Accordingly absorption measurements were made on solutions of suitable concentrations in the ultra violet region of the spectrum using a Eeckman model DU Spectrophotometer with a photomultiplier attachment and the hydrogen lamp accessory unit,.. The absorbancy of a solution of sodium bromoacetate or sodium thiosulphate of suitable concentration and the absorbancy of a solution of the medium (sodium chloride, magnesium chloride, or sea water) were measured. The absorbancy of a mixture of these components at the same concentrations as in the simple solutions was then determined. The effect of pH, when pos-sible was also studied in these solutions, A comparison was made between the absorbancy for the mixture and the sum of the absorbancies of the separate solutions. Any difference in these absorbancies was assumed to be due to the formation of some complex. Although this assumption was not directly verified i t could lead to a qualitative explanation of some of the results of the rate measurements• RESULTS RATE MEASUREMENTS The results of the rate measurements for the reaction SiOa + BrCBzCOO" — > SzO3CH2_C0O= + Br~ are given in Tables XT - H. for the three temperatures studied. The logarithm -16 SPECIFIC RATE CONSTANTS FOR THE BROMOACETATE THIOSULPHATE REACTION IN SODIUM CHLORIDE MEDIA AT 15.3°C Initial Concentration of Reactants - 0,01 moles/litre Exp't. No. Total Normality of Cations Ionic Strength IA> k (litres moles-' min.1 ) 1 + log k Initial pH 1 0.08 0.090 0.243 0.385 6.5-7.5 2 0.23 0.240 0.330 0.518 6.5-7.5 3 0.38 0.393 0.375 0.574 6.5-7.5 , 4 0.655 0.674 0.443 0.646 6.5-7.5 5 0.655 0.674 0.444 0.646 6.5-7.5 6 1.03 1.06 0.504 0.702 6.5-7.5 7 1.03 1.06 0.505 0.703 9.0 8 1*03 1,06 0.512 0.709 6.5 9 1.03 1.06 0.510 0.707 8.4 10 1.03 1.06 0.5H 0.710 6.2 11 1.03 1.06 0.504 0.702 8.8 12 1*06 1.10 0.508 0.706 6.5-7.0 13 o u i 0.432 0.388 0.589 6.5-7.0 12, 13 - Initial Concentration of Reactants - 0.02 moles/litre TABLE IH SPECIFIC RATE CONSTANTS FOR THE BROMOACETATE THIOSULPHATE REACTION IN SEA WATER MEDIA AT 15.3°C Initial Concentration of Reactants - 0«01 moles/Litre Total Ionic k Exp't. Noe Normality Strength (litres moles"1 min;*) 1 + log k Initial of Cations pH u 0.1A6 0.172 0.310 0.492 6.5-7.5 15 0.275 0.317 0.367 0.564 6.6 16 0.275 0.317 0.364 0.561 7.8 , 17 0.356 0.412 0.392 0.593 6.5-7.5 18 0.570 0.665 0.399 0.600 8.7 19 0.570 0.665 0.4a 0.644 7.1 20 0.685 0.798 0.437 0.640 8.7 23. 0.685 0.798 0.462 0.666 6.5 22 0.687 0.808 0.469 0.670 6.5-7.5 23 0.904 1.072 1.096 0.504 0.702 7.1 6.5-7.5 24 0.925 0.500 0.699 25 0.305 0.360 0.377 0.576 7.0 26 0.305 0.360 0.377 0.576 7.7 27 0.113 0.140 0.284 0.453 7.4 2 4 , 27 - Initial Concentration of Reactants = 0 . 0 2 moles/litre -18-TABLE XJL. SPECIFIC RATE CONSTANTS FOR BROMOACETATE THIOSULPHATE REACTION IN MAGNESIUM CHLORIDE MEDIA AT 15,3°C Initial Concentration of Reactants » O.Ol moles/litre Exp't. No. Total Normality of Cations Ionic Strength k (litres moles'1 min.') 1 +• log k Initial pH 28 0.110 0.160 0.335 0.525 7.0 29 0.110 0.160 0.334 0.524 6.5-7.0 30 0.230 0.340 0.401 0.603 6.5-7.0 31 0.330 0.490 0.420 0.623 6.5-7.0 32 0.450 0.670 0.437 0.640 6.6 33 0.630 0.940 0.459 0.661 6.5-7.0 34 0.280 0.340 0.384 0.584 6.5-7.0 35 0.880 0.940 0.489 0.688 6.5-7.0 34 - Mixture - NaCl (0.15M) + MgClz. (0.05M) 35 - Mixture - Nad (0.75M) + MgC^ (0.05M) -19-TABLE jj SPECIFIC BATE CONSTANTS FOR TBE BROMOACETATE THIOSULPHATE REACTION IN SODIUM CHLORIDE MEDIA AT 25,2°G Initial Concentration of Reactants - 0*01 moles/litre Initial pH - 6.5 - 7.5 Total Ionic k Exp't No. Normality Strength (litres moles'1 min.-1) 1 * log k of Cations 36 0.030 0.040 0.491 0.692 37 0.030 0.040 0.486 0.687 38 0.030 0.040 0.493 0.692 39 0.130 0.140 0.703 0.846 AO 0.230 0.240 0.839 0.923 41 0.188 0.198 0.789 0.893 42 0.330 0.342 0.923 0.964 A3 0.430 0.443 1.002 0.985 1.000 0.993 44 0.435 0.449 45 0.486 0.500 1.015 1.006 46 0.435 0.449 0.996 0.998 47 0.537 0.553 1.048 1.020 48 0.537 0.553 1.030 1.012 49 0.537 0.553 1.052 1.022 50 0.537 0.553 1.057 1.023 51 0.610 0.628 1.090 1.037 52 0.661 0.680 1.122 1.050 53 0.530 0.546 1.050 1.021 54 0.655 0.673 1.090 1.036 55 0.984 1.009 1.247 1.095 56 1.030 1.060 1.269 1.102 57 1.030 1.060 1.168 1.067 58 1.030 1.060 1.243 1.094 59 1.030 1.060 1.276 1.105 60 1.030 1.060 1.285 1.108 61 1.030 1.060 1.281 1.107 62 1.030 1.060 1.258 1.098 63 1.030 1.060 1.276 1.105 64 1.030 1.060 1.277 1.105 65 1.430 1.479 1.372 1.137 66 1.430 1.479 1.381 1.140 67 0.966 0.986 1.243 1.094 68 0.918 0.952 1.239 1.092 69 0.691 0.719 1.142 1.058 70 0.660 0.680 1.121 1.050 71 0.720 0.786 1.113 1.048 72 1.120 1.180 1.314 1.118 57, 58, 62 - Initial pH - 9.0 67 - Initial Concentration of Reactants '68, '69, 70 - n « « « 71, 7.2 - « n n o 0.005 moles/litre 0.02 moles/litre 0.04 moles/litre -20-TABLE TT SPECIFIC RATE CONSTANTS FOR THE BROMOACETATE THIOSULPHATE REACTION IN SEA WATER MEDIA AT 25*2*0 Initial Concentration of Reactants » 0,01 moles/Litre _ Initial pH - 6.5 - 7.5 Total Ionic k Exp't, No. Normality Strength (litres 1 +- log k of Cations moles'1 min.1 ) 73 0.056 0.071 0.612 0.786 74 0.099 0.119 0.710 0.850 75 0.140 0.165 0.778 0.890 76 0.197 0.230 0.844 0.925 77 0.232 0.268 0.887 0.947 78 0.235 0.272 0.895 0.951 79 0.235 0.272 0.356 0.897 0.952 80 0.306 0.965 0.984 81 0.306 0.356 0.948 0.976 82 0.263 0.306 0.933 0.970 83 0.360 0.420 1.012 1.016 1.005 84 0.380 0.441 1.007 85 0.545 0.637 1.119 1.048 86 0.689 0.810 1.172 1.069 87 0.805 0.950 1.221 1.087 88 0.805 0.950 1.225 1.088 89 1.000 1.115 1.288 1.110 TABLE mi SPECIFIC RATE CONSTANTS FOR THE BROMOACETATE THIOSULPHATE REACTION IN SEA WATER MEDIA AT 25.2°C Initial Concentration of Reactants « O.Ol moles/litre Total Ionio k Exp«t. No. Normality Strength (litres 1 + log k Initial of Cations moles-* min^ ) pH 90 0.545 0.637 1.119 1.048 7.0 91 0.545 0.637 1.020 1.008 8.7 92 0.463 0.540 1.012 1.005 8.7 93 0.648 0.762 1.058 1.023 8.7 94 0.515 0.602 1.048 1.020 8.7 95 0.515 0.602 1.030 1.012 8.7 96 0.570 0.666 1.020 1.008 8.7 97 0,570 0.666 1.033 1.014 8.7 98 1.000 1.115 1.288 1.110 6.9 99 1.000 1.115 1.266 1.102 7.5 100 1.000 1.115 1.219 1.086 8.1 8.0 •101 0.955 1.137 1.240 1.093 102 0.770 0.908 1.161 1.061 8.0 103 0.572 0.667 1.108 1.043 7.7 104 0.545 0.637 1.119 1.048 7.7 105 0.808 0.955 1.185 1.073 7.9 106 0.600 0.706 1.081 1.034 8,7 107 0.517 0.671 1.090 1.037 8.7 108 0.601 0.711 1.110 1.044 8.7 109 0.713 0.847 1.U5 1.047 8.7 110 0.673 0.793 1.137 1.055 8.7 111 0 . 4 0 2 0.477 1.002 1.001 8.7 112 0.631 0.751 1.050 1.020 1.048 8.7 113 0.631 0.751 1.113 8.7 114 0.888 1,076 1.178 1.070 7.9 106, 107 - Initial Concentration of Reactants - 0.02 moles/litre 108 - 1111- " • a a 5 0^03 m o i e s / l i t r e 112 - 114- M 11 • • s 0 . 0 4 moles/litre TABLE v m SPECIFIC RATE CONSTANTS FOR THE BROMOACETATE THIOSULPHATE REACTION IN SODIUM CHLORIDE MEDIA AT 32.05'G Initial Concentration of Reactants « 0*01 moles/litre Initial pH - 6.5 - 7.0 ' Exp't. No. Total Normality of Cations Ionic Strength UO k (litres moles"' min.1) 1 + log k 115 0.080 0.090 1.134 1.054 116 0.130 0.140 1.303 1.114 117 0.230 0.240 1.512 1.180 118 0.380 0.392 1.729 1.237 119 0.530 0.546 1.909 1.280 120 0.750 0.770 2.118 1.325 121 0.930 0.955 2.218 1.345 122 1.030 1.060 2.270 1.356 TABLB g: J. SPECIFIC RATE CONSTANTS FOR THE BROMOACETATE THIOSULPHATE REACTION IN SEA WATER MEDIA AT 32.0? 0 Initial Concentration of Reactants » 0.01 moles/litre Initial pH - 6.5 - 7.5 Exp't. No. Total Normality of Cations Ionic Strength CU) k (litres moles-1 min;1) 1 «*• log k 123 0.082 0.100 1.218 1.085 124 0.084 0.101 1.234 1.091 125 0.163 0.192 1.488 1.172 126 0.297 0.345 1.739 1.240 127 0.426 0.497 1.879 1.273 128 0.805 0.950 2.175 1.337 129 0.805 0.950 2.200 1.342 130 0.899 1.066 2.254 1.352 -23-of the specific rate constants at 15.3°C has been plotted against the square root of the ionic strength (Figure 2) for the different media in which the reaction was conducted. The results at higher temperatures f a l l on very similar curves. For each of the temperatures studied, i t is seen that at low ionic strengths the rate constants for the reaction in sea water are higher than those obtained in sodium chloride. However as the ionic strength of the solution is increased by the addition of solutions of these neutral salts, the rate in sea water gradually approaches that in sodium chloride and falls below the latter above an (ionic strength)^ of about 0.90. The effect of magnesium chloride on the specific rate constant at 15.3°C is shown in Table TJT and included in Figure 2. The results show that, at low ionic strengths, the reaction is faster than in sodium chloride media. This effect of divalent ions on the reaction has been reported previously by other workers (5»6). However, as the ionic strength is increased by the ad-dition of more magnesium chloride the curve approaches that obtained in sodium chloride media and intersects i t at an (ionic strength) of 0.8Q. and continues to l i e below the latter at greater ionic strengths. Following the conclusions of Olsen and Simonson (8) that the re-action rate was dependent upon the kind and concentration of the ions opposite in charge to those of the reacting species, the logarithms of the specific rate constants have been plotted against the normality of the cations in the solutions. Figures 3, A » 5 show that the curves obtained for the three temp-atures studied have the same general shape. At the highest temperature, however, the curves converge slightly at high concentration. The reaction is faster in sea water than in sodium chloride solutions containing the same normality of cations over the entire concentration range studied. A few of La iter's results have been included in Figure A . His data for the sodium salts of the reactants agree well with the author's. However (IONIC STRENGTH)* FIGURE 2. RATE CONSTANTS IN VARIOUS MEDIA AT 15.3°9. NaCl, © ~ — -OMgClj,, Sea Water. » Reaotant Conc.O.OlM • Reactant Conc.0.02M i 0.3 0.A 0.5 0.6 NORMALITY OP CATIONS FIGURE 4. RATE CONSTANTS IN VARIOUS MEDIA AT 25.2° C. o o NaCl, o » Sea Water, La Mer's Results (5) Na Salts - A K Salts - * T -28-there is no agreement when these data and the author's are plotted against the square root of the ionic strength. This disagreement is encountered be-cause of the methods used to raise the ionic strength of the solutions. La Her did not increase the ionic strength by the addition of neutral salts such as sodium chloride and therefore when the sodium concentrations are the same in his and the present experiments, the ionic strengths differ since a larger contribution to the ionic strength is made by the divalent thiosulphate ion then from the monovalent chloride ion. La Mer showed that when the reactants were in the form of their pot-assium salts the reaction was faster than when sodium salts were used (Figure 4). On the strength of a l l these results i t may be concluded that for the reaction studied, the particles responsible for the increase in rate constant in sea water are the potassium, calcium and magnesium ions. Two experiments were made in mixtures of sodium chloride and magne-sium chloride, the concentration (0,05M) of the magnesium being slightly less than that found in sea water of chlorinity 19,00 °/oo. At the lower concen-tration, the rate of reaction is faster than in sea water, as would be expected, since the magnesium concentration is higher in this solution than in sea water of this concentration (Figure 3), At the higher concentration, the rate is slightly lower than in sea water, presumably because the concentration of the magnesium is lowe than that of calcium plus magnesium in sea water. These experiments agree with the suggestion that potassium, magnesium and calcium ions are responsible for the specific rate constants being faster in sea water than sodium chloride solutions. Other ions may have an effect on the reaction rate, but because of their small concentration, the effect is probably slight. EFFECT OF VARYING pH IN SEA WATER MEDIA The effect of pH on the specific rate constant in sea water solutions -29-has been found to be very pronounced. The pH anomaly was brought to our at-tention because the sea water was concentrated by distillation for many of the experiments. Faurholt (16) has shown that in an alkaline solution carbon dioxide may react directly with hydroxyl ion to give the bicarbonate ion. This in turn will react with more hydroxyl ion to yield carbonate ion as shown in the f o l -lowing equations CG^ (dissolved) +• OH" HC03~ HCO3-»- OH" CQ^HjP The carbon dioxide which varies from 1.5 to 2.5 millimoles per l i t r e for normal sea water, a concentration 15 to 30 times that in air, is supplied by solution of CO2. from the atmosphere, by decomposition of dead organic material or by leaching from land by solution of carbonate rocks. The concentration of the sea water removed carbon dioxide, resulting in an increase in the hydroxyl-ion concentration since the equlibruim was shifted to the left according to the above equations. The pH of the sea water was not adjusted in the early experiments in sea water media, with the result that low values were obtained for the rate constants, since i t was assumed that the best straight line second-order plot through the experimental points should also pass through the origin. If the reaction is not closely followed in i t i a l l y , i t is quite possible to draw a straight line through the origin that passes through most of the points. As more experiments were completed at high ionic strengths i t became apparent that there was a considerable drift with time in the calculated specific rate constants. The calculated values increased towards a constant value. The best straight lines drawn through the plots of £ vs. time did not pass through the origin. The graph of log k vs .-{JZ is shown in Figure 6 and i t is seen that the rate at high pH lies far below those obtained in nearly neutral media and are not reproducible. i : — r - 1 1—; 1 r 0.3 O.A 0.5 0.6 0.7 0.8 0.9 1.0 1.1 (IONIC STRENGTH)^ FIGURE 6. EFFECT OF pH AND REACTANT CONCENTRATION ON RATE CONSTANTS IN SEA WATER AT.25.2° C. -31-Several experiments were done at 25.2°C at the same concentration of sea water and the same high pH but using higher concentrations of reactants. A l l the experiments showed an Induction period. However, the rate constants are somewhat higher than those obtained for reaction concentrations of 0.01M (Figure 6) . This increase in rate constant for an increase in the reactant concentration does not appear when the reaction is done in neutral media at 15.3°C as shown in Figure 2. More experiments were done at high and low pH in sea water solutions, the reaction being followed more closely i n i t i a l l y . By lowering the temper-ature to 15.3° C the i n i t i a l part could be followed more easily. The experimental results of one such experiment are listed in Table Z and plotted in Figure- 7 0 Table "X. shows the steady increase in the calculated rate constants, finally reaching a constant value, for sea water of high pH as compared with the con*» stant value obtained in a nearly neutral solution. A plot of the experimental values indicates an induction period of about 10 minutes. Also the value of k obtained from the graph s t i l l differs from 6-956 from that obtained in nearly neutral sea water media. The pH did not change during the reaction regardless of the i n i t i a l pH of the mixture. A few experiments were done in media having a pH between 7.9-8.1. The results show that the rate in this pH range is nearly the same as in neutral solutions (Figure 6) . The results of three reactions done in media of the same concen-tration but varying pH are plotted in Figure 8. The specific rate constants agree to within 1.5/6 up to a pH of 7.5 but beyond this the difference between the rate constants becomes much larger. However in dilute solutions of sea water (CI - 7.66 c/oo) the rate constants show a variation of less than 1$, which is within the limit of the experimental error, between a pH of 6.6 and -32-TABLE X EFFECT OF pH ON THE SPECIFIC RATE CONSTANT IN SEA WATER MEDIA AT 15.3°C pH - 8.7^ pH - 7 .2 Time (min.) k (litres moles"1 minr1 ) Time (min.) X k (litres moles-1 min;' ) a - x a - x 12 0.0141 0.118 16 0.0727 0.455 26 0.0696 0.226 29 0.133 0.459 43 0.141 0.328 46 0.213 0.463 64 0.225 0.352 67 0.307 0.458 84 0.328 0.390 93 0.431 0.464 105 0.407 0.388 109 0.513 0.470 \ -33' TIME (min.) FIGURE 7. o-GRAPHICAL DETERMINATION OF RATE CONSTANTS IN SEA WATER MEDIA AT 15.3° C, O pH - 7.2,« O pH - 8.7 0 20 40 60 80 100 TIME (min.) FIGURE 8. GRAPHICAL DETERMINATION OF RATE CONSTANTS IN SEA WATER MEDIA AT 25.2° C. O OpH - 6o9, •• • pH - 7.5, « • pH - 8.1 7.8 as shown by experiments (15) and (16) in Table H I . EFFECT OF VARYING pH IN SODIUM CHLORIDE MEDIA. Similar studies were carried out in sodium chloride solutions of dif-ferent pH. The pH of these solutions was adjusted by the addition of a drop of sodium hydroxide solution. Results showed no pronounced induction period in solutions of high pH, although the rates were about 2% smaller than those ob-tained in nearly neutral sodium chloride solution (Exp1ts.58,62,Table IT). A reaction carried out at 25.2*C showed a decrease in pH from 9.0 to 7.3 for the time the reaction was studied. A similar decrease in pH was noted in a l l re-action mixtures in sodium chloride media initially of high pH. This change is equivalent to a decrease in the hydroxyl ion concentration of lxlO" 5 moles per litre which is presumably accompanied by an equivalent decrease in bromoacetate concentration as given by the following equation BrCHaC00" •> OH" > CHjQHCOO" * Br" Correcting for the bromoacetate concentration does not account for the observed 2% decrease in rate constant in sodium chloride media of high pH. SUMMARY OF THE EFFECT OF pH ON THE RATE CONSTANTS The investigations of the effect of pH on the rate constants show a number of results. 1. In concentrated sea water of high pH (8.7) there is an induction period; the rate constants are 6-9$ less than those in nearly neutral media. 2. In dilute sea water media the induction period is absent and the rate remains the same up to a pH of 7.80, whereas results are only reproducible in concentrated sea water up to a pH of 7.5". 3. Regardless of its i n i t i a l value there is no change of pH as the reaction progresses in sea water media; however in sodium chloride media the ike pH drops to nearlyA neutral point during the course of the reaction. -36-4. A 2j6 difference is observed in rate constants when the reaction is done in sodium chloride solutions of high (8.7) and low pH (6.5)* 5. A calculation shows that the hydrolysis of the sodium bromoacetate is not sufficient to account for the 2% difference in rate constants observed in sodium chloride media. ABSORBANCE MEASUREMENTS Sodium Bromoacetate and Sodium Chloride Solutions Absorbance measurements were made with solutions of sodium bromo-acetate (0.001M), sodium chloride (l.OM) and a mixture of the two salts. The data, plotted in Figure 9, show that there is no difference between the sum of the absorbancies of the separate solutions and that of the mixture of the salts of similar concentrations between wave lengths of 220 and 360 rmj. when the pH of the solution was approximately 6.0. One drop of NaOH was added to the above solutions, raising the pH to 9.5, and the absorbancies of the three solutions were again measured. There was no significant differences in the absorbancies at pH 6.0 and 9.5. When the sodium bromoacetate concentration was increased to 0.05M, the absorbancies of the mixture remained equal to the sum of those of the components. In solutions of sodium bromoacetate and sodium chloride there does not appear to be any interaction between the sodium and the bromoacetate ions. Sodium Bromoacetate and Magnesium Chloride Solutions Similar measurements were made with solutions of magnesium chloride (0.3M) and sodium bromoacetate (0.001M) between wave lengths of 220 and 330 mji. It was found that there was only a very slight and probably in-significant difference between the absorbancies of the mixture and the sum of the absorbancies of the two salts. Apparently no interaction occurs when sodium bromoacetate is mixed with magnesium chloride. •o -37-0.4 0.3 0.2 0.1 1 1 1 \ 1 1 \\ - \ \ — \ 1 * - - ° — 1 1 I 9 220 230 260 270 240 250 WAVE LENGTH (m/») FIGUBS 9- ABSORBANGE OF SOLUTIONS OF SODIUM BROMOACETATE SODIUM CHLORIDE, AND A MIXTURE OF THESE COMPONENTS. ° o SaBrAo(0.00l|),o o NaCl(l.OM), • » Mixture. -38-It was not possible to raise the pH of these solutions containing magnesium chloride since precipitation occurred above a pH of 7.9. Sodium Bromoacetate and Sea Water Solutions Similar measurements made with sea water, sodium bromoacetate and mixtures of them at pH 5*7 and 8*7 indicate that l i t t l e or no interaction occurs between them* In general, i t appears therefore that sodium bromoacetate does not interact with the solutes present in any of the three reaction media studied under conditions at which its reaction with sodium thiosulphate was examined* Sodium Thiosulphate and Sodium Chloride Solutions Because of the high extinction coefficient of solutions of sodium thiosulphate, solutions of lower concentration than those of sodium bromo-acetate were used* Figure 10 shows the absorbancies of solutions of sodium thiosulphate (0.000AM), sodium chloride (1*0M) and a mixture of the two salts at a pH of 7.0 and 9,3* The absorbancy of each of these solutions between wave lengths of 220 and 280 m>«.. did not change with pH. A "^difference" curve, obtained by subtracting the sum of the absorbancies of the components from that of the mixture at a particular wave length, is also plotted in Figure 10* There is a maximum difference (0*139) at a wave length of 230 mju If i t is assumed that the "difference" curve results because of the removal of thio-sulphate due to incomplete dissociation and that the undissociated species does not absorb at 230 mu* then the percentage thiosulphate removed is .139 x 100 a 15.856* .882 When the experiments were done with solutions of sodium thiosulphate (0.001M) and sodium chloride of the same concentration as before, similar re-sults were obtained. At 230 mjuu, the point of maximum difference, the percent-age thiosulphate removed is .3AO x 100 = 15.3/6. 2.22 -39. 210 220 230 240 250 260 WAVE LENGTH (mp) FIGURE 10. ABSORBANCE OP SOLUTIONS OF SODIUM THIOSULPHATE, SODIUM CHLORIDE, AND A MIXTURE OF THESE COMPONENTS. ° ONatSi03(0.0004M), •>----• NaCl(l.OM), o » Mixture. The "difference• curve represents the absorbancy of the mixture minus the sum of the absorbancies of the individual components. - A O -Other experiments done at high concentrations of sodium thiosulphate (0.01) showed no absorption peak for the undissociated species when the solu-tions were studied up to a wave length of 400 m>*. The absorbancy of the mixture was the same as the sum of the absorbancies of the individual solutions. Sodium Thiosulphate and Sea Water Solutions Similar measurements were made with sodium thiosulphate (0.0008M) and sea water solutions at a pH of 6,1 and 8,7 (Figure 11), A plot of the difference between the absorbancies of the mixture and the sum of the ab-sorbancies of the two components again shows a maximum difference at 230 mu, and the larger difference occurs in solutions of high pH. Experiments made when the concentration of the sodium thiosulphate was 0,001M showed similar results. Similar measurements with 0,000AM sodium thiosulphate at pH 6.4 and 8.4 gave "difference" curves that lay closer together^than those of the more concentrated solutions because the solution of high pH was less strongly alkaline than those mentioned above. It therefore appears that complex formation takes place between the thiosulphate ion and the cations in sea water and to a greater extent at higher pH, Sodium Thiosulphate and Magnesium Chloride Solutions Experiments were done with solutions of magnesium chloride (0,3M) and of sodium thiosulphate of varying concentrations (0,0004-0,0O1M). The results a l l showed that there was a maximum difference at a wave length of 230 mu.. Again on the assumption that undissociated species are present, the percentage thiosulphate removed in this way is 16% (average of 4 experiments). The re-sults of one of the experiments are shown in Figure 13. Although this value is approximately the same as obtained for sodium chloride, nevertheless i t indicates that more thiosulphate is removed in solutions of magnesium chloride, than in sodium chloride of the same concentration. - u -1.4 1.2 1.0 o 0.8 0.6 0.4 0.2 210 PH-6.I + 0.1 fH o to 1 -0.1 -0.2 -0.3 220 230 240 WAVE LENGTH (my) 250 FIGURE 11. ABSORBANCE OF SOLUTIONS OF SODIUM THIOSULPHATE, SEA WATER, AND A MIXTURE OF THESE COMPONENTS. o Na»Sl03(0.0008M), o- o Sea Water(Cl-2025 %s) , «— - Mixture The "difference" curve represents the absorbancy of the mixture minus the sum of the absorbancies of the individual components. -42-1.2 -210 pH-8-4-pH-6-4-I _ l _ + 0.1 x p 5 r-o.i « --0.2 220 230 240 WAVE LENGTH (mu) 250 260 FIGURE 12. ABSORBANCE OF SOLUTIONS OF SODIUM THIOSULPHATE, SEA WATER, AND A MIXTURE OF THESE COMPONENTS* o o NajSjOj J0.000AM), o ° Sea Water(Cl-18.39* ° Mixture The "difference" curve represents the absorbancy of the mixture minus the sum of the absorbancies of the individual components. T - 4 3 -\ I I I I J 220 230 240 250 260 WAVE LENGTH (my) FIGURE 13. ABSORBANCE OF SOLUTIONS OF SODIUM THIOSULPHATE, MAGNESIUM CHLORIDE, AND A MIXTURE OF THESE COMPONENTS. o o Na^Sa03(0.0004M), o -o MgCle(0.3M), <» © Mixture The "difference" curve represents the absorbancy of the mixture minus the sura of the absorbancies of the individual components„ -44-Further experiments were done with solutions of magnesium chloride of varying concentrations (0.15-0.6M) and of sodium thiosulphate (0.01M) up to a wave length of 400 mju The results showed that the "difference11 curve had a slight broad maximum between 275 and 280 mjj., (Figure IA)• With dif-ferent concentrations of magnesium chloride, there were very slight differences (0.04) in the height of the peak. It was not possible to repeat the experiments with solutions of high pH since precipitation occurred. Sodium Thiosulphate and Barium Chloride Solutions Similar experiments, done with barium chloride (0.3M) and sodium thiosulphate (0,0004M) and mixtures of the salts, gave a "difference" curve with a maximum difference at a wave length of 230 mji. (Figure 15). A similar calculation shows that .232 x 100 « 26% of the thiosulphate is undissociated. This value is larger than that for magnesium chloride, indicating the form-ation of a more stable complex. It is also noted in Figure 15 that sodium thiosulphate has an absorbanee maximum at 215 mu, in agreement with the work of others (17)» In general the results show that an interaction occurs between sodium thiosulphate and the cations of the media used in the reaction .mix-tures, with a larger interaction occurring in sea water in alkaline than in neutral solution. ACTIVATION ENERGY An important quantity which determines the specific rate constants is the energy of activation (E A) which appears in the well known Arrhenlus equation where A is a constant independent or only slightly dependent upon temperature and E^ represents the experimental activation energy. -45-270 280 WAVE LENGTH ( FIGURE 14. ABSORBANCE OF SOLUTIONS OF SODIUM THIOSULPHATE, MAGNESIUM CHLORIDE, AND A MIXTURE OF THESE COMPONENTS. o o NajS^O.Ol), • • MgCl2.(0.6M), 9 » Mixture The "difference" curve represents the absorbancy of the mixture minus the sum of the absorbancies of the individual components, 220 230 240 250 260 WAVE LENGTH (m/J FIGURE 15. ABSORBANCE OF SOLUTIONS OF SODIUM THIOSULPHATE, BARIUM CHLORIDE, AND A MIXTURE OF THESE COMPONENTS. ° ° N a ^ O a (O.OOOAM), o o BaCl^ (0.3M), o » Mixture The "difference 0 curve represents the absorbancy of the mixture minus the sum of the absorbancies of the individual components. -47-Equation (5) in logarithmic form becomes logk « log A - Vz3o3RT ( 6 ) A plot of log k vs J- should give a straight line of slope equal to -and therefore E A can be determined graphically. E A can also be calculated from rate data at two different temperatures using the expression log kj_ s J±—( IizZM (7) In Table 21 are indicated the logarithms of the specific rate constants read from the graphs for the reaction carried out in sodium chloride and sea water media. From these values, the energy of activation has been calculated using Equation (7) for the three temperature intervals, 15.3°-25.2°, 15.3°-32.0*1 and 25.2"-32.0:°' • A study of the values of the energy of activation shows they apparently decrease slightly as the ionic strength of the solution is increased in both sodium chloride and sea water media. The average values, however, over the range of the ionic strength studies are not significantly different for the sea water and sodium chloride media. The value of the activation energy for this reaction in the presence of sodium chloride or sea water is 16,000 - 100 calories/mole. ENTROPY OF ACTIVATION .Calculation of the value of the term log A appearing in Equation (6) by using the experimental value of 16,000 cal./mole for the energy of activation gives an average value of 9.95 - 0,05 for both sea water and sodium chloride media. Therefore A has the value 8.92 x 10 . The term A is considered to be equal to the product of two terms, namely P and Z, the probability factor and the collision frequency respective-ly. The average value of the theoretical collision frequency varies but u slightly from reaction to reaction, the average value being about 2.77x10 . Therefore P is approximately 3.24x10 8 Since P «e where AS i s the entropy of activation, the latter has a value of -6.6e.u • -AS-TABLE 2 1 ACTIVATION ENERGY 25.2°- 15.3°C NaCl Sea Water log k log k E log k log k E (25.2') (15.3°) (cal./moles) (25.2°) (15.3°) (cal./moles) Oo3 0.782 0.385 15,800 0.811 0.413 15,900 O.A 0.863 0.455 16,300 0.881 0.480 16,000 0.5 0.928 0.516 16,400 0.940 0.532 16,300 0.6 0.972 0.564 16,300 0.983 0.579 16,100 0.7 1.010 0.603 16,200 1.016 0.612 16,100 0.8 1.041 0.640 16,000 1.042 0.642 15,900 0.9 1.071 0.672 15,900 1.069 0.671 15,900 1.0 1.098 0.702 15,800 1.092 0.697 15,800 Average 16,100 Average 16,000 15.3°- 32.0.*3C 0.3 1.054 0.385 16,100 1.078 0.413 16,000 0.4 1.129 0.455 16,200 1.149 0.480 16,100 0.5 1.185 0.516 16,100 1.202 0.532 16,100 0.6 1.230 0.564 16,000 1.245 0.579 16,000 0.7 1.269 0.603 16,000 1.277 0.612 16,000 0.8 1.301 0.640 15,900 1.301 0.642 15,800 0.9 1.330 0.672 15,800 1.324 0.671 15,700 1.0 1.350 0.702 15,600 1.343 0.697 15,500 Average 16,000 Average 15,900 25.2°- 32.0.^C -0.3 1.054 0.782 16,600 1.078 0.811 16,300 0.4 1.129 0.863 16,300 1.149 0.881 16,300 0.5 1.185 0.928 15,700 1.202 0.940 16,000 0.6 1.230 0.972 15,700 1.245 0.983 16,000 0.7 1.269 1.010 15,800 1.277 1.016 15,900 0.8 1.301 1.041 15,900 1.301 1.042 15,800 0.9 1.330 1.071 15,800 1.324 1.069 15,600 1.0 1.350 1.098 15,400 1.343 1.092 15,300 Average 15,900 Average 15,900 -49-DISCUSSION Rate measurements in a l l media show a wide deviation from those predicted by Bronsted equation. This behavior has often been reported (5,7), Such deviations presumably arise as a result of the inapplicability of the Debye-Huckel expression for activity coefficients in other than the most dilute solutions of electrolytes. More recent work (8) has led to the conclusion that, for reactions between ions of the same sign, the rates are dependent almost exclusively on the concentration and character of the salt ions of the charge opposite to that of the reactants, and are therefore not simply depend-ent upon the ionic strength of the solution. In view of this work the log-arithm of the rate constant was plotted against the normality of the cations, Agreement was then obtained between tbe author's data and those of other workers (Figure 4)« It is also noted that the experimental curves have a slightly dif-ferent shape when the logarithm of the rate constant is plotted against the square root of the ionic strength (Figure 2), The intersection of the curves in this figure is possibly due to the unsuitability of the square root of the ionic strength for the abscissa instead of normality of the cations. The (ionic strength) of 0.5N MgCla and 0.5N NaCl are respectively 1,2 and 0,7, The use of (ionic strength)^ therefore expands the abscissa for MgClg media and leads to intersection of the rate curve in such solutions with that obtain-ed in sodium chloride. The rate curve for sea water solutions will intersects that for sodium chloride solutions for the same reason. Rate measurements show that reaction rates are much faster in magnesium chloride than in sodium chloride solutions of the same normality over the concentration range studied. Similar results have been obtained by other authors (5), for such ions as calcium, magnesium,.barium and lanthanum for -50-this and other reactions. The increase in reaction rate in solutions of polyvalent ions has been qualitatively explained (5) as resulting from the increase in the number of collisions and hence in the probability of reaction because of the abnormal accumulation of negative reactant ions around the cat-ions. It has also shown that the reaction is slightly faster in solutions of potassium ions than in those of sodium ions of the same concentration (5,13). The higher reaction rates observed in sea water media may therefore be due to magnesium, calcium and potassium ions but predominantly to the first two, since their combined concentration in ordinary sea water is approximately 0.07 moles per l i t r e . The concentration of the other ions in sea water is small, and so their effect on the reaction rate is probably negllble* The reaction rates in sodium chloride and sea water media appear to level off with increasing concentration at the same rate while those in mag-nesium chloride media, although considerably faster at low concentration, approach fairly rapidly the sodium chloride and sea water curves at high con-centrations. The results of the absorbanee measurements indicate that there is an interaction between the cations of the media and the thiosulphate ions. It is therefore possible that undissociated species of the type NaS^ O^  , CaS^ Os , MgSz03 are present in the reaction mixtures. Dissociation constants for CaSaOs and BaS^Oj have been reported (10)» If undissociated species are present in the reaction mixtures, then two mutually exclusive explanations of a few anomalies that arise from the present studies may be made. The f i r s t is that of Wyatt and Davies (9) who have shown that the increase in observed rate constants in solutions of calcium and barium ions compared with monovalent ions at similar ionic strengths, may be explained by the reaction BaSi03 + BrCHzC00" — > S^CHjCOO" + Br" * Ba** on the assumption that It is a faster reaction than the purely ionic reaction -51-and contributes measurably to the observed rate. However l i t t l e support is found for this argument. Slator (IB) has reported a considerable re-duction in rate due to a decrease in the concentration of thiosulphate ion when silver nitrate was added to the reaction of ethyl bromoacetate and thiosulphate ions, A calculation showed that i f the concentration of the thiosulphate was taken as the difference between the iodine titre and that corresponding to the complea Ag^S^O^)^ , then i t was found that the rate agreed fairly closely with that obtained in pure sodium thiosulphate. These results Indicate that the removal of some of the thiosulphate by complex formation results in a constant that is lower than that observed in the ab-sence of silver nitrate. It therefore appears that the presence of undis-sociated species, such as MgStO^  CaS^ O^ , Na S 2 0 £ would produce calculated rates lower than their real values since the concentration of thiosulphate used in the calculation is larger than that actually available for reaction. If i t is assumed that the undissociated species are of the follow-ing form Mg"* *• SzOf ^=± MgSz03 then the formation constant can be written It can be shown that — — - fraction of thiosulphate that is free It is possible to correct the observed rate constant for the amount of thio-sulphate removed due to the formation of the undissociated species. The following reactions are considered MgSaOa — * Mg+* + StCy1 (fast reaction) S^Cy +BrAc" > S z 0 3Ac = + Br" (slow reaction) -52-We can therefore write But - d C B c ^ l = V o b s {j»x(ylC Br *c 1 ^ « W^M>S""1*0 (8) where.kt is the corrected rate constant. The magnitude of the correction will be dependent upon the concentration of the cations and the stebility of the undissociated species. CALCULATION OF FORMATION CONSTANT The absorbanee measurements show that in a 1.0M sodium chloride solution at least 15 .5$ of the thiosulphate is undissociated. fraction of thiosulphate undissociated - CNG.Sa.O3] -o.\55 The formation constant for the reaction Na++ St03= ^ NaSjO^ * is given by K - fo^o*! ^ 1 „ o.-va (Kla**3 C ^ f c O ^ l L l.o-K5.45l « Similarly for the reaction MgVSiOf MgSt03 a formation constant can be written 1 / _ [M^S^Qa*] The data showed that in a 0.3M solution of magnesium chloride at least 16% of the thiosulphate was undissociated _ _ 1 = 5 - 2 5 -53-These values are at best minimum values for the formation constants since i t has been assumed that the undissociated species does not absorb; at the wave length of the maximum in the "difference" curves. However these values have been inserted into Equation (8) to correct the observed rate constants in sodium chloride and magnesium chloride media at 15.3sG. The corrected rate values have been plotted in Figure 16. It is seen that the magnesium chloride and sodium chloride curves do not converge as rapidly as when the rate constants had no correction. The sea water curve i f corrected should s t i l l l i e above the sodium chloride curve. Corrections to the rate constants in sea water have not been applied since formation constants for KS^O^" and CaS203 were not determined,. The decrease in rate constants observed in alkaline sea water media can also be attributed to undissociated species. Absorbance measurements indicate that there is a greater difference between the absorbancy of the mixture and that of the sum of absorbancies of sea water and sodium thio-sulphate in alkaline solutions than in nearly neutral solutions. Apparently some of the undissociated species present in sea water are more stable in alkaline solution. Less thiosulphate is directly available for reaction and observed rate constants are lower than in neutral media. Further, absorbance measurements did not indicate an increase in the stability of the undissociated species, NaS^Oj , in alkaline media. Rate constants are only slightly dif-ferent in neutral sodium chloride media than in alkaline media. It therefore appears that the stability of the undissociated species Mg^Oj, and CaStO-j increases with increasing pH. Absorbance measurements were done on solutions of BaCl2(0«3M) and NazS203(0.0004M). The results showed that the undissociated species, BaSj>03, is more stable than MgSi03. The stability of the undissociated species there-'s-. I I I I I 1 1 £.2 0.3 0.A 0.5 0.6 0.7 0.6 NORMALITY OF CATIONS FIGURE 1 6 . RATE CONSTANTS CORRECTED FOR THE POSSIBLE PRESENCE OF UNDISSOCIATED SPECIES NaCl Observed o o MgCl£ Observed ® © Corrected Corrected -55-fore, decreases in the following order, BaS;. O5 > GaSz03 ? MgS^ O,,. The experimental results did not suggest any explanation for the induction period apparent in sea "water of high pH. This is discussed further in a following section* SUMMARY 1* Rate measurements for the bromoacetate-thiosulphate reaction in a l l media (sodium chloride, sea water and magnesium chloride) showed a wide deviation from those predicted by the Bronsted equation. This behaviour is due to the inapplicability of the Debye-Huckel expression for activity coef-ficients in other than the most dilute solutions of electrolytes. 2. In agreement with other work (8), the rates appear to be de-pendent almost exclusively on the concentration and character of salt ions of .charge opposite to that of the reactants. 3. The reaction rates are about 6/6 faster in sea water than in sodium chloride media of the same normality of cations over the concentration range studied. The faster rates are attributed to the potassium, magnesium and calcium ions in sea water. 4. Reaction rates in sodium chloride and sea water appear to level off at the same rate, with increasing concentration while those in magnesium chloride media, although considerably faster at low concentrations approach fairly rapidly the sodium chloride and sea water curves at high concentrations. Results of absorbance measurements indicate an interaction between the cations of the various media and the thiosulphate ions. It has therefore been post-ulated that the .following species exist, NaStOa" , MeS^ , GaSi03 in the reaction mixtures. Since some thiosulphate is removed in their formation, -56-reaction rates are lower than expected. An equation has been derived to cor-rect the observed rate constants for the presence of these species and has the following form where K is the formation constant of the metal-thiosulphate species and M is the concentration of the cation. The magnitude of the correction therefore depends on the stability of the undissociated species and on the concentration of the cation, 5 , Formation constants for the species, NaS^O^' and MgSi03 have been calculated. They are at best minimum values. However when the constants are used to correct the reaction rates in sodium chloride and magnesium chloride the corrected rate curves no longer converge to the same extent as the observed rate curves for these media. 6 , Lower rate constants in sea water solutions of high pH compared with nearly neutral media are also attributed to the presence of undissociated species, particularly MgSfcO^  and CaSj03 since these appear to be more stable with increasing pH, The stability of the undissociated species decrease in the following way B a S ^ > CaS^ Os > EgS-jO*,, 7. The energy of activation is 16,000±100 calories/mole for the reaction in both sodium chloride and sea water media, 8 . The entropy of activation has been calculated to be - 6 . 6 e.u. "57-SUGGSSTIONS FOR FURTHER WORK Studies should be done to attempt to determine the cause for the induction period noted only in sea water media of high pH and chlorinity. Thus far the experimental evidence seems to indicate interaction between the thiosulphate and the media. Possibly experiments should f i r s t be at-tempted with the sodium bromoacetate in the side arm of the flask instead of the sodium thiosulphate as was followed throughout this work. Other ionic reactions might be studied in sea water media to determine whether similar anomalies appear as did in the present studies. -58-BIBLIOGRAPHY 1. Dittmar, W. Report on researches into the composition of ocean water, collected by H.M.S. "Challenger". Challenger Rep»ts, Physics and Chem., .1, 1-251, (1884). 2. Amis, E.S. Kinetics of Chemical Change in Solution. The MacMillan Company, New York, (19A9). 3. Debye, P. and Huckel, E. Physik.Z. 2£, 185-206, (1923). A. Livingston, R. J. Chem. Education. 7, 2887-2903, (1930). 5. La Mer, V.K. and F: essenden, R.W. J. Am. Chem. Soc. 54. 2351-2366, (1932). 6. Bronsted, J.N. and Livingston, R. J. Am. Chem. Soc. 49. 435-446, (1927). 7. La Mer, V.K. J. Am. Chem. Soc. £1, 3341-3347, (1929). 8. Olson, A.R. and Simonson, T.R. J. Chem. Phys. 17, 1167-1173, (1949). 9. Wyatt, P. and Davies, C.W. Trans. Faraday Soc, Aj>, 774-780, (1949). 10. Davies, C.W. and Wyatt, P.A. Trans. Faraday Soc. AJ5, 770-773, (1949). 11. La Mer, V.K. and Kamner, M.E. J. Am. Chem. Soc. £7, 2662-2668, (1935). 12. Von Kiss, A . and Vass, P. Z. anorg. allgem. Chemie. 217. 305-320, (1934). -59-: 13. Kappana, A.N. J. Ind, Chem. Soc. 6, 45-52, (1929). 14. Vbgel, A.I. Quantitative Inorganic Analyses. Longmans, Green and Go. Toronto, (1951). 15. Lyman, J. and Fleming, R.H. J. Marine Res., Sears Foundation Marine Res. 134-146, (1940). 16. Faurholt, C. Dissertation, Copenhagen (1924) as cited by Thompson, T.G. Bull. 85, V7.5, Nat. Research Council, (1932). 17. Buck, R.P., and Singhadeia, A. and Rogers, L.B. Analytical Chem. 26, 1240-1242, (1954). 18. Slator, A. J. Chem. Soc. 8%, 481-494, (1905). 

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