A STUDY OF THE DETERMINATION OF BROMIDES 
 IN THE PRESENCE OF EXCESS CHLORIDES 
 
 BY 
 
 MAYOR FARTHING FOGLER 
 B. S. University of Illinois, 1920 
 
 THESIS 
 
 Submitted in Partial Fulfillment of the Requirements for the 
 
 Degree of 
 
 MASTER OF SCIENCE 
 IN CHEMISTRY 
 
 IN 
 
 THE GRADUATE SCHOOL 
 OF THE 
 
 UNIVERSITY OF ILLINOIS 
 
 1921 
 

 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 ■ 
 
 
 
 
 

 UNIVERSITY OF ILLINOIS 
 
 THE GRADUATE SCHOOL 
 
 Augu st 1. 192-1 • 
 
 I HEREBY RECOMMEND THAT THE THESIS PREPARED UNDER MY 
 
 SUPERVISION BY -Mayor Farthing Fogler 
 
 ENTITLED A St udy o f t h e De te rr ri nation of Bromide s in 
 
 th e Pr esence of Excess Chlorides 
 
 BE ACCEPTED AS FULFILLING THIS PART OF THE REQUIREMENTS FOR 
 THE DEGREE OF has tan of Science^ 
 
 Recommendation concurred in* 
 
 Committee 
 
 on 
 
 Final Examination* 
 
 ^Required for doctor’s degree hut not for master’s 
 
 476970 
 
A STUDY OF THE DETERMINATION OF BROMIDES 
 IN THE PRESENCE OF EXCESS CHLORIDES 
 
 MAYOR FARTHING FOGLER 
 B. S. University of Illinois, 1920 
 
 THESIS 
 
 Submitted in Partial Fulfillment of the Requirements for the 
 
 Degree of 
 
 MASTER OF SCIENCE 
 IN CHEMISTRY 
 
 IN 
 
 THE GRADUATE SCHOOL 
 OF THE 
 
 UNIVERSITY OF ILLINOIS 
 
 1921 
 

 
 
 
 
 
 ' Vi* to rr ffc rm r 4 ;r • >iti rcfutt 
 
 
 
 If <>TAIJC1AM{> till 
 
 "tl .O 
 
 ic hi >f '-i * n (Vi nu 
 
Table of Contents. 
 
 Page 
 
 I . Introduction . ... 1 
 
 II. Historical .......... 
 
 Ill . Experimental 
 
 A. Gravimetric Methods ...... 
 
 Precipitation of Silver Bromide in the 
 Presence of Excess of Chlorides 
 Precipitation from Ammonium Hydroxide 
 Solutions ....... 
 
 Alcohol as a Solvent * 
 
 Cause of Low Results ..... 
 
 Empirical Corrections for Low Results 
 
 2 
 
 4 
 
 4 
 
 5 
 
 7 
 
 S 
 
 9 
 
 B. Selective Oxidation Methods . 
 
 Sodium Sulfite as a Reducing Agent 
 Oxidizing Agents ...... 
 
 Fixation of Bromine by Metallic Silver 
 Theory of Oxidation of Bromides and 
 
 Chlorides 
 
 Method Recommended ..... 
 
 14 
 
 15 
 
 16 
 19 
 
 19 
 
 21 
 
 IV. Summary .... 23 
 
 V. Bibliography 25 
 
 Plates . 
 
 1 . 
 
 2 . 
 
 Opposite 
 
 Page 
 
 Solubility of AgBr in NaCl Solution ... 13 
 
 Change of Oxidation Potential with Concentration 20 
 
Digitized by the Internet Archive 
 in 2016 
 
 https://archive.org/details/studyofdeterminaOOfogl 
 
Acknowledgment 
 
 The writer wishes to express 
 his appreciation and sincere 
 thanks to Doctor J. H. Reedy 
 for his assistance in the 
 experimental work, and also 
 for his aid in writing this 
 thesi s . 
 

 
 
 
A Study of the Determination of Bromides 
 In the Presence of Excess 
 of Chlorides . 
 
 I. Introduction* 
 
 A fundamental problem, in the bromine industry is the 
 
 accurate estimation of the bromide content of the brines or 
 * 
 
 bitterns. Since the bromide content is low usually less 
 than 5^ — this problem involves the determination of bro- 
 mides in the presence of a large excess of chlorides. It 
 is important because it is desirable to know just how much 
 bromine is present, so as to add exactly the equivalent 
 amount of oxidizing agent needed to effect its liberation. 
 Various methods in current use which give satisfactory re- 
 sults for high concentrations give very divergent results 
 for low concentrations. The purpose of this investigation 
 is to study the methods for estimating bromine and to ascer- 
 tain, if possible, the cause of their inaccuracy. 
 
 II . Historical ♦ 
 
 Chemical literature contains numerous methods that have 
 been proposed for this analysis, most of them depending on 
 
 ( 1 ) 
 

 
 
 
 
 
 
 
 
 
 . 
 
 © 
 
 . 
 
 ■ 
 
 
 
 
 
 
p 
 
 a selective oxidation of the bromide and its subsequent titra- 
 tion or gravimetric determination. Vortman 3 - who attempted 
 to determine chlorine in the presence of bromine used lead 
 dioxide in the presence of acetic acid to oxidise the bro- 
 mides and then determined gravimetrically the chlorine which 
 remained in the solution. The results were always high be- 
 cause the bromine was not all set free, while if a higher 
 concentration of acetic acid was used traces of chlorine 
 were liberated. Cavazzi s used barium peroxide and sulfuric 
 acid as the oxidizing agent. The bromine was distilled, but 
 a little chlorine was also liberated. Eng.le^ tried to use 
 
 ammonium persulfate and sodium acetate at 70° -80°. Berg- 
 4 
 
 lund used potassium hydrogen sulfate and potassium perman- 
 ganate in the cold, and aspirated air through the solution 
 to remove the bromine, collecting it in sodium hydroxide 
 solution. He reports that the bromine was then determined 
 gravimetrically, though no statement is made as to how the 
 sodium hypobromite formed was reduced. It was found, more- 
 over, that while the oxidizing agent would not liberate chlo- 
 rine if it alone were present, it did liberate chlorine when 
 bromine was present. He avoided this by a double aspiration. 
 Baubigny and Rivals 0 used copper sulfate and potassium per- 
 manganate at the constant temperature of £0°, but with little 
 success, as chlorine was also liberated. Wyss^ attempted 
 to separate all of the halogens by selective oxidation. He 
 first removed the iodine by ferric sulfate, and then added 
 
. 
 
 ' 
 
 . 
 
 - 
 
 ■ 
 
 . 
 
 
 
3 
 
 potassium permanganate and Seated to 60° to liberate the bro- 
 mine which was determined gravimetrically . Ee stated that 
 his method gave good results, but offered no figures to sup- 
 port his claim. White^ used aluminium sulfate and potassium 
 permanganate and stated that bromine was liberated while chlo 
 rine and iodine were not. It is quite inconceivable how an 
 oxidizing agent could have sufficient oxidizing power to lib- 
 erate bromine and at the same time not liberate iodine which 
 is much more readily oxidized. Jannasch and Aschoff- used 
 acetic acid and permanganate tc liberate the bromine but 
 their results were always io w. Bugarszky^ very carefully 
 investigated other methods proposed, and himself used iodic 
 acid to free bromine, but obtained very poor results. An- 
 drews 1 ^ used iodic acid to oxidize the bromide and found it 
 very suitable if the bromine was present in large quantities. 
 But this method is hardly applicable where chlorine is pres- 
 ent in large quantities. Later investigators have found 
 out that the oxidation potential of the oxidizing agent must 
 lie between that of bromine and chlorine. As an oxidizing 
 agent of the correct potential iodic acid has been used, as 
 mentioned above. Gooch 11 used selenic and telluric acids, 
 but neither works well unless chlorine is present in very 
 minute amounts. Skinner and Baughman 155 used hydrogen per- 
 oxide and chromic acid to liberate the bromine, which was 
 removed by aspiration in the cold. They found that chlorine 
 also came over from a saturated solution to the extent of 
 
. 
 
 ' 
 
 . 
 
 - 
 
4 
 
 about ifo. This was remedied by double aspiration. They ob- 
 tained very good results, but the method is rather long and 
 involved. 
 
 None of the above methods are adapted to rapid, accu- 
 rate work, and seme of them admit of errors as large as 2*fc 
 Those that are fairly accurate require several hours or even 
 days for their completion. It would appear, then, that the 
 optimum method has not yet been developed. 
 
 III. Experimental . 
 
 There are in general two methods for determining bro- 
 mides in the presence of an excess of chlorides: (1) The 
 
 gravimetric method, which involves the precipitation of the 
 bromine and all or part of the chlorine as silver halides 
 and the subsequent estimation of the composition of the pre- 
 cipitate, either by indirect analysis or calculation; ( 2 ) 
 
 The selective oxidation of the bromine, using an oxidizing 
 agent that will not liberate chlorine, or at least only a 
 small amount. 
 
 A. Gravimetric Methods . 
 
 Precipitation of Silver Bromide in the Presence of Ex - 
 cess of Chlorldes --A gravimetric method that appears — ■ on 
 its fac6, at least -- to be both rapid and accurate, involves 
 the addition of a known amount of silver nitrate to the ha- 
 lide solution in a quantity somewhat in excess of that needed 
 to completely precipitate the bromine. The amount of bro- 
 
- 
 
 
 
 . 
 
 . 
 
 . 
 
 . 
 
 — 
 
5 
 
 mine present may be calculated from th© following expression: 
 
 Wt. of Br = 1.7976 x Wt . of AgBr + AgCl - 
 
 2.3885 x wt. of Ag. 
 
 The excess of silver nitrate over the amount necessary to 
 preceipitate the bromine should be small. This follows from 
 the fact that, for mixtures of silver bromide and silver 
 Chloride, 
 
 $ AgBr = 100-= - 1 3 .. 00 x - ..-?!■£ 
 
 Wt. of AgBr + AgCl 
 
 Now, in case of an error in the weight of the silver halide 
 precipitate, the effect cn the result will be small if the 
 amount of silver chloride is small; but if the latter is 
 large, the value of the fraction in the above expression 
 will be increased considerably; that is, the magnitude of 
 the inaccuracy will be multiplied as many fold, as the weight 
 of the silver chloride is increased. 
 
 The accuracy of this method depends on the assumption 
 that the bromine is quantitatively precipitated by the sil- 
 ver nitrate before appreciable amounts of silver chloride 
 are formed. This inference in justified by a comparison of 
 the solubility products of the two halides. 
 
 Precipitation from Ammoniu m Hydroxi de Solutions --! t 
 was felt that the above method was open to criticism in that, 
 owing to the speed of the reaction, the silver nitrate might 
 be used up locally in the formation of silver chloride, leav- 
 ing unprecipitated bromides in the solution. It is known 
 

 
 
 
 - 
 
 
 
 
 
 
 
 
 
 
 
 . 
 
 
 
 
 
 : 
 
 
 - 
 
 
 
 - 
 
 - 
 
 
 
6 
 
 that the conversion of silver chloride into silver bromide 
 by means of soluble bromides is low. Hence if the precipi- 
 tation could be made to proceed slowly, the silver bromide 
 precipitation would occur first, and the silver chloride pre- 
 cipitation would not begin until the former is complete. 
 
 In order to do this it is necessary to use a solvent for the 
 silver halides that may be gradually removed, or else to 
 add the silver nitrate very slowly. Ammonium hydroxide was 
 the first solvent to be used. The experiments were carried 
 out as follows: Solutions of potassium bromide and potaseium 
 
 chloride were mixed, a volume of ammonium hydroxide added, 
 and then enough silver nitrate to precipitate all the bro- 
 mine and part of the chlorine. The beaker containing the 
 mixture was heated to boiling and stirred vigorously until 
 all the ammonia was driven out. Any traces were neutralized 
 by the addition of dilute nitric acid, and the precipitate 
 filtered in a Gooch crucible, dried and weighed. Table I 
 shows typical results. 
 
 Table I_. Result s with Ammonium Hydroxide as Solvent . 
 
 Exp. 
 
 Per cent 
 KOI 
 
 Amount of AgBr 
 Theory 
 
 Amount of AgBr 
 Found 
 
 Difference 
 
 1 
 
 Sfo 
 
 1.5480 gm. 
 
 1.5254 gm. 
 
 -.0126 gm 
 
 2 ! 
 
 2$ 
 
 1 . 5480 
 
 1.4850 
 
 -.0630 
 
 5 
 
 2 f 0 
 
 1.5480 
 
 1.5405 
 
 -.0075 
 
 4 
 
 2 fo 
 
 1.5480 
 
 1.4945 
 
 -.0535 
 
 5 
 
 P% 
 
 1 . 5480 
 
 1.4861 
 
 -.0619 
 
 6 
 
 P$ 
 
 1 . 5480 
 
 1.5533 
 
 -.0147 
 
 7 
 
 
 1 . 5480 
 
 1.4666 
 
 -.0814 
 
 8 
 
 9 $ 
 
 1 . 5480 
 
 1.4260 
 
 -.1220 
 
 9 
 
 
 1.5480 
 
 1.4789 
 
 -.0691 
 
 10 
 
 9 $ 
 
 1.5480 
 
 1.4996 
 
 -.0484 
 
 11 
 
 2 % 
 
 1.5480 
 
 1.4771 
 
 -.0709 
 

 * * I 1C1 Bkff- ;’|H^ 
 
 ' ' 
 
 
 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 . 
 
 
 . 
 
 
 — 
 
 
 
 
 
 — 
 
 « . 
 
 
 
 
 > ■— 
 
 
 
 o . 
 
 
 
 . 
 
 
 
 
 
 
 
 • X 
 
 
 • - 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 
 
 ■ 
 
 
 . 
 
 
 
 
 
 . 
 
 
 
 
7 
 
 It will be noted that the results were always low, indi- 
 cating incomplete precipitation of either the bromine or the 
 silver. This fact will be interpreted later. 
 
 Alcohol as a Solvent — It was thought that the silver 
 halides were both tcc soluble in ammonium hydroxide to get 
 consistent results, so some other solvent was sought for. 
 
 It was observed that silver chloride is appreciably soluble 
 
 in ethyl alcohol, while silver bromide is not. So the above 
 determinations were repeated, except 50$ alcohol was used 
 as the solvent. Table II sho?/s the result?.. 
 
 Table II. Results with 50$ Alcohol as Solvent. 
 
 Exp . 
 
 Per cent 
 KC1 
 
 Amount of AgBr 
 Theory 
 
 Amount of AgBr 
 Found 
 
 Difference 
 
 1 
 
 2$ 
 
 1.5480 gm. 
 
 1.5518 gm. 
 
 +.0038 
 
 2 
 
 2$ 
 
 1.5480 
 
 1.5158 
 
 -.0312 
 
 3 
 
 2$ 
 
 1.5480 
 
 1.4956 
 
 -.0524 
 
 4 
 
 2$ 
 
 1.5480 
 
 1.5231 
 
 -.0249 
 
 5 
 
 2$ 
 
 1.5480 
 
 1.5080 
 
 -.0400 
 
 6 
 
 2$ 
 
 1.5480 
 
 1.5100 
 
 -.0380 
 
 7 
 
 2 % 
 
 1.5480 
 
 1.5268 
 
 -.0212 
 
 8 
 
 2%. 
 
 1.5480 
 
 1.5328 
 
 -.0152 
 
 9 
 
 2$ 
 
 1 . 5480 
 
 1.5121 
 
 -.0359 
 
 10 
 
 2$ 
 
 1.5480 
 
 1.5334 
 
 -.0146 
 
 
 Water as a 
 
 Solvent- -Mainly 
 
 for the purpose 
 
 i 
 
 •H 
 
 u 
 
 cfi 
 
 P* 
 
 a 
 
 o 
 
 o 
 
 o 
 
 son 
 
 with the above results, a series of determinations was 
 
 made 
 
 with water 
 
 as the solvent. 
 
 Table III shows 
 
 the results. 
 
 
 Table III. Results with Water as Solvent. 
 
 Exp. 
 
 Per cent 
 KC1 
 
 Amount of AgEr 
 Theory 
 
 Amount of AgBr 
 Found 
 
 Difference 
 
 1 
 
 2$ 
 
 1.5480 
 
 1.5228 
 
 -.0252 
 
 s 
 
 2% 
 
 1.5480 
 
 1.5250 
 
 -.0230 
 
 3 
 
 2$ 
 
 1.5480 
 
 1.5344 
 
 -.0126 
 
 4 
 
 2$ 
 
 1.5480 
 
 1.5098 
 
 -.0382 
 

 . 
 
 
 
 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 • 
 
 
 
 
 
 . 
 
 
 
 
 _ 
 
 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 _ 
 
 
 
 
 
 
 
 • — 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 . 
 
 
 . 
 
 
 
 
 
 ♦ _ .-j 
 
 
 • 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 - 
 
 
 . 
 
 
 
 
 , 
 
 
 
 
The results in the experiments in water solution are 
 distinctly better than with ammonium hydroxide and alcohol 
 as solvents, but still they fall far below the usual stan- 
 dard for accuracy. 
 
 To meet the suggestion that the low results were due 
 to the bromine not being completely precipitated, owing to 
 the slow conversion of silver chloride into silver bromide, 
 the following experiment was made: The silver in 4 oc. of 
 
 normal silver nitrate was completely precipitated as sil- 
 ver chloride and then stirred with an excess of sodium bro- 
 mide solution for six hours, after which it was found that 
 about 75 °jo of the chlorine had been replaced by bromine. In 
 the precipitation of silver from silver nitrate solution 
 by mixed halides, on the other hand, the silver halide is 
 formed in a very highly dispersed condition, which has a 
 far higher solubility than the coagulated form, and in this 
 state the bromine seems to replace the chlorine in a com- 
 paratively short time. It would appear that the magnitude 
 of these dispersed particles is so small that no particle 
 is large enough to permit of the formation of a protective 
 coating of silver bromide and thus prevent the inner part 
 from reacting with the sodium bromide. 
 
 Cause of Low Re suit a - -The cause of the low results in 
 the precipiation methods is to be found in the solvent action 
 of alxaline halides on silver bromide and silver chloride. 
 Schierhclz- 1 - 0 has reported that 100 grams of sodium chloride 
 in concentrated solution dissolve 0.474 grams of silver bro- 
 

 
 
 - 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 . ) 
 
 
 
 
 
 
 
 - 
 
 
 
 
9 
 
 mide, and IOC grams of potassium bromide in concentrated solu- 
 tion dissolve 3.019 grams of this salt, the temperature being 
 15° in both cases. In a similar way, silver chloride is 
 also dissolved by solutions of concentrated alkaline halides. 
 
 Empirical Corrections for Low Result s --The expedient 
 of deriving a set of solubility values for silver bromide 
 in alkaline halides was then taken up. Given these data, 
 empirical corrections could be made on all analytical re- 
 sults. So it was planned to construct a solubility curve 
 for silver bromide in concentrations of sodium chloride vary- 
 ing from if* to saturation. The determinations were made as 
 fellows: 50 cc . of approximately 0.1 N silver nitrate were 
 
 taken and just enough sodium bromide was added to precipi- 
 tate all of the silver as silver bromide. Results are given 
 in Table IV. 
 
 Table IV . Solubili ty of Silver Bromide in Sodium 
 
 Chloride Solution. 
 
 Grams per 
 100 cc » of 
 
 Weight of AgBr 
 
 Weight of AgBr 
 
 Weight of AgBr 
 
 Solution 
 
 Theory 
 
 Found 
 
 Dissolved 
 
 1 
 
 .9354 
 
 .9178 
 
 .0176 
 
 1 
 
 .9354 
 
 .9100 
 
 .0254 
 
 1 
 
 .9354 
 
 .9169 
 
 .0185 
 
 1 
 
 .9324 
 
 .9204 
 
 .0120 
 
 1 
 
 .9324 
 
 .9193 
 
 .0131 
 
 1 
 
 .9324 
 
 .9200 
 
 .0124 
 
 1 
 
 .9798 
 
 .9706 
 
 .0092 
 
 1 
 
 .9798 
 
 .9705 
 
 .0093 
 
 1 
 
 .9798 
 
 .9712 
 
 .0086 
 
 2 
 
 .9354 
 
 .9087 
 
 .0267 
 
 2 
 
 .9354 
 
 .9044 
 
 .0310 
 
 2 
 
 .9354 
 
 .9087 
 
 .0267 
 

 -• 
 
 r ' 
 
 * 
 
 
 
 
 
 
 
 
 . 
 
 . 
 
 
 
 
 . 
 
 
 
 
 . 
 
 
 
 
 
 
Table IV--Cont. 
 
 Grams per 
 100 cc « cf 
 
 Weight of AgBr 
 
 Weight of AgBr 
 
 Weight of AgB 
 
 -L- V V w v 1 • V* i* 
 
 Solution 
 
 Theory 
 
 Found 
 
 Dissolved 
 
 2 
 
 .9324 
 
 .9133 
 
 .0191 
 
 2 
 
 .9324 
 
 .9195 
 
 .0129 
 
 2 
 
 .9324 
 
 .9143 
 
 .0181 
 
 2 
 
 .9798 
 
 .9668 
 
 .0130 
 
 2 
 
 .9798 
 
 .9683 
 
 .0115 
 
 2 
 
 .9798 
 
 .9656 
 
 .0142 
 
 3 
 
 .9354 
 
 .9012 
 
 .0342 
 
 3 
 
 .9354 
 
 .90 §5 
 
 .0329 
 
 3 
 
 .9354 
 
 .9027 
 
 .0327 
 
 3 
 
 .9324 
 
 .9082 
 
 .0242 
 
 3 
 
 .9324 
 
 .9100 
 
 .0226 
 
 3 
 
 .9324 
 
 .9077 
 
 .0247 
 
 4 
 
 .9354 
 
 .8986 
 
 .0368 
 
 4 
 
 .9354 
 
 .8959 
 
 .0395 
 
 4 
 
 .9354 
 
 .8976 
 
 .0378 
 
 4 
 
 .9324 
 
 .9055 
 
 .0269 
 
 4 
 
 .9324 
 
 .9060 
 
 .0264 
 
 4 
 
 .9798 
 
 .9605 
 
 .0293 
 
 4 
 
 .9798 
 
 .9597 
 
 .0301 
 
 4 
 
 .9798 
 
 .9588 
 
 .0316 
 
 5 
 
 .9354 
 
 .8900 
 
 .0454 
 
 5 
 
 .9354 
 
 .8908 
 
 .0446 
 
 5 
 
 .9354 
 
 .8905 
 
 .0449 
 
 5 
 
 .9324 
 
 .9014 
 
 .0310 
 
 5 
 
 .9324 
 
 .9022 
 
 .0302 
 
 6 
 
 .9354 
 
 .8905 
 
 .0449 
 
 6 
 
 .9354 
 
 .8902 
 
 .0446 
 
 6 
 
 .9354 
 
 .8900 
 
 .0448 
 
 6 
 
 .9324 
 
 .9000 
 
 .0324 
 
 6 
 
 .9324 
 
 .8994 
 
 .0330 
 
 7 
 
 .9354 
 
 .8871 
 
 .0483 
 
 7 
 
 .9354 
 
 .8885 
 
 .0469 
 
 7 
 
 .9324 
 
 .8958 
 
 .0368 
 
 7 
 
 .9324 
 
 .8958 
 
 .0368 
 
 7 
 
 .9324 
 
 .8952 
 
 .0374 
 

 
 
 
 
 
 
 
 . 
 
 . 
 
 
 . 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
11 
 
 Table IV-=»Ccnt. 
 
 Grains per 
 1GG cc. of 
 Solution 
 
 Weight of AgBr 
 
 Weight of AgBr 
 
 Weight of AgBr 
 
 Theory 
 
 Found 
 
 Dissolved 
 
 7 
 
 .9798 
 
 .9537 
 
 .0867 
 
 7 
 
 .9798 
 
 .9580 
 
 .0878 
 
 7 
 
 .9798 
 
 .9586 
 
 .0878 
 
 8 
 
 .9354 
 
 .8830 
 
 .0511 
 
 8 
 
 .9354 
 
 .8836 
 
 .0505 
 
 8 
 
 .9384 
 
 .8945 
 
 .0379 
 
 8 
 
 .9384 
 
 .8945 
 
 .0379 
 
 8 
 
 .9384 
 
 .8939 
 
 .0385 
 
 9 
 
 .9354 
 
 .8888 
 
 .0538 
 
 9 
 
 .9354 
 
 .8815 
 
 .0539 
 
 9 
 
 .9354 
 
 .8817 
 
 .0537 
 
 9 
 
 .9798 
 
 .9473 
 
 .0385 
 
 9 
 
 .9798 
 
 .9507 
 
 .0891 
 
 9 
 
 .9798 
 
 .9458 
 
 .0340 
 
 10 
 
 .9354 
 
 .8771 
 
 .0581 
 
 10 
 
 .9354 
 
 .8790 
 
 .0568 
 
 10 
 
 .9354 
 
 .8784 
 
 .0568 
 
 10 
 
 .9758 
 
 .9444 
 
 .Q354 
 
 10 
 
 .9758 
 
 .9444 
 
 .0354 
 
 10 
 
 .9758 
 
 .9440 
 
 .0358 
 
 11 
 
 .9798 
 
 .9400 
 
 .0398 
 
 11 
 
 .9798 
 
 .9416 
 
 .0388 
 
 11 
 
 .9798 
 
 .9389 
 
 .0409 
 
 18 
 
 .9798 
 
 .9396 
 
 .0408 
 
 IS 
 
 .9798 
 
 .9389 
 
 .0409 
 
 17 
 
 .9798 
 
 .9888 
 
 .0570 
 
 17 
 
 .9798 
 
 .9888 
 
 .0516 
 
 17 
 
 .9798 
 
 .9890 
 
 .0508 
 
 SO 
 
 .9798 
 
 .9185 
 
 .0613 
 
 SO 
 
 .9798 
 
 .9184 
 
 .0614 
 
 SO 
 
 .9798 
 
 ,9141 
 
 .0657 
 
 ss 
 
 .9798 
 
 .9140 
 
 .0656 
 
 ss 
 
 .9798 
 
 .9187 
 
 .0671 
 
 ss 
 
 .9798 
 
 .9189 
 
 .0669 
 
12 
 
 Table IV— Cent. 
 
 Grams per 
 ICC 1 cc . of 
 Solution 
 
 Weight of AgBr Weight of AgEr 
 Theory Found 
 
 Weight of AgBr 
 Dissolved 
 
 24 
 
 .9798 
 
 
 . .9079 
 
 
 .0719 
 
 24 
 
 .9798 
 
 
 .9127 
 
 
 .0671 
 
 24 
 
 .9798 
 
 
 .9120 
 
 
 .0678 
 
 26 
 
 .9798 
 
 
 .9042 
 
 
 .0756 
 
 26 
 
 .9798 
 
 
 .9040 
 
 
 .0758 
 
 26 
 
 .9798 
 
 
 .908 2 
 
 
 .0716 
 
 28 
 
 .9798 
 
 
 .8924 
 
 
 .0874 
 
 28 
 
 .9798 
 
 
 .8909 
 
 
 .0887 
 
 28 
 
 .9798 
 
 
 .8937 
 
 
 .0861 
 
 30 
 
 .9798 
 
 
 .8830 
 
 
 .0968 
 
 30 
 
 .9798 
 
 
 .8805 
 
 
 .0993 
 
 30 
 
 .9798 
 
 
 .8797 
 
 
 .1001 
 
 32 
 
 .9798 
 
 
 .8681 
 
 
 .1117 
 
 32 
 
 .9798 
 
 
 .8699 
 
 
 .1099 
 
 34 
 
 .9798 
 
 
 .8540 
 
 
 .1258 
 
 34 
 
 .9798 
 
 
 .8591 
 
 
 .1207 
 
 34 
 
 .9798 
 
 
 .8538 
 
 
 ,1260 
 
 36 
 
 .9798 
 
 
 .8408 
 
 
 .1390 
 
 36 
 
 .9798 
 
 
 .8409 
 
 
 .1389 
 
 36 
 
 .9798 
 
 
 .8450 
 
 
 .1348 
 
 In the 
 
 above table 
 
 the results from 
 
 one and the same 
 
 solution are 
 
 set off in 
 
 blocks 
 
 of three, 
 
 or 
 
 sometimes two. 
 
 It will be noticed that, 
 
 for the most part, 
 
 there is good 
 
 agreement within these blocks. 
 
 For different blocks, how- 
 
 ever, w r here 
 
 solutions of 
 
 sodium 
 
 bromide i 
 
 of 
 
 different con- 
 
 centrations 
 
 were used in 
 
 preparing the silver bromide, and 
 
 where uniform conditions 
 
 were not maintained, the results 
 
 are markedly 
 
 different . 
 
 Take , 
 
 for example. 
 
 the solubilities 
 
 in 1 % sodium 
 
 chloride solution. 
 
 The first 
 
 three detemiina- 
 
 tions are from the same 
 
 concentration of 
 
 sodium bromide, were 
 
 run at the same time, and under conditions that were iden 
 
30 
 
 36 
 
 34 
 
 32 
 
 30 
 
 26 
 
 26 
 
 24 
 
 22 
 
 20 
 
 /8 
 
 / 6 
 
 J4 
 
 /2 
 
 / 0 
 
 6 
 
 6 
 
 4 
 
 2 
 
 0 
 
 G/7)<s. /VaC/ per /OO cc /V a O 
 
— ] 
 
 13 
 
 tical, including concentration, heating, stirring, and so 
 forth . 
 
 This indefiniteness in the solubility of silver bromide 
 is explained by the fact that substance is a colloid, and 
 the size of the particles depends on such conditions as those 
 just mentioned. And in turn, the solubility of a solid de- 
 pends on the size of the particles, the solubility increas- 
 ing enormously with the degree of subdivision. 
 
 Attempts were made to regulate the conditions by carry- 
 ing out all determinations alike. For example, each precip- 
 itate was stirred for a definite time to coagulate the par- 
 ticles, and washed with a definite amount of water. In this 
 way better checks were obtained. If a solution of sodium 
 bromide of the same concentration we re used throughout, good 
 results could probably be obtained and a fairly smooth solu- 
 bility curve drawn for the solubility of silver bromide in 
 various strengths of sodium, chloride solution. Figure I 
 shows three solubility curves plotted from the data in Table 
 IV. Curve I was plotted from solubility measurements on 
 silver bromide made by mixing .0994 N sodium bromide and 
 .1043 N silver nitrate in equivalent amounts; curve II from 
 .1007 N sodium bromide and .0993 N silver nitrate; curve 
 III from .1006 N sodium bromide and .0996 N silver nitrate. 
 
 Other conditions than concentration were practically the 
 same in all cases. 
 
 It will be seen therefore that the amount of silver 
 bromide which dissolves in a certain concentration of sodium 
 

 
 
 
 
 
 
 ' 
 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 
 - 
 
 - 
 
 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 , • I 1 1 
 
 
14 
 
 chloride solution is not constant unless all details of tech- 
 nique are regulated so as to be exactly identical in every 
 case. These, of course, can not be regulated on an unknown 
 sample. Therefore it is not possible to determine bromine 
 in a solution of this kind by precipitating and weighing it 
 and then adding an empirical correction which corresponds 
 to its solubility in that concentration of sodium chloride 
 solution . 
 
 £ • Selective Oxidat ion Meth ods . 
 
 Methods involving the selective oxidation of bromides 
 to free bromine and the subsequent distillation of the lib- 
 erated halogen offer certain advantages, the most important 
 of which is the concentration of the bromine. Later work 
 has shown that it is not always possible to make a "clean” 
 separation from chlorine in this way, since varying amounts 
 of the latter are usually present in the distillate. But 
 it does have the advantage of reducing the chloride concen- 
 tration very considerably. The importance of this point 
 is evident when it is remembered that the halogen content 
 of the distillate is almost always determined gravimetrically , 
 and, as has been shown above, the gravimetric method has the 
 maximum accuracy when the amount of chlorine is a minimum. 
 
 The whole range of oxidizing agents has been explored 
 by investigators in their efforts to find one that will pref- 
 erentially oxidize the bromine without affecting the chlo- 
 rine. Nitrosyl sulfuric acid ("ni trose" ) , chromic acid, 
 acidified potassium permanganate, and so forth, have been 
 
. 
 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
15 
 
 used with more cr less uncertain success. 
 
 An important point in the technique of these distilla- 
 tion methods is that the halogens of the distillate must 
 he absorbed in some suitable reducing agent which will re- 
 convert them into bromides and chlorides, respectively. If 
 silver nitrate were also present in this absorbing medium, 
 the halogens would be immediately precipitated in their 
 final form for weighing. The use of silver nitrate without 
 a reducing agent would lead to low results, owing to the 
 formation of oxidized bromine compounds like hypobromous 
 acid and silver bromate. 
 
 Sodium Sulfite as a Reducing Agent - -The first agent 
 tried was sodium sulfite. The determination was carried 
 out as follows: 5G cc . of 0.1 N sodium bromide was placed 
 
 in a distilling flask and an amount of sodium chloride was 
 added and the solution diluted to about 100 cc. To this 
 was added IS grams of chromic acid and about 2 cc. of 30$ 
 hydrogen peroxide, and the mixture distilled. The solution 
 in which the halogens were to be absorbed consisted of 5 
 cc. of 0.1 N silver nitrate and the calculated amount of 
 sodium sulfite. The distillation was run for about 30 min- 
 utes. The precipitate was coagulated by stirring, collected 
 in a Gooch crucible, dried and weighed. Results were as 
 follows: 
 
 Br Present 
 
 Br Found 
 
 1.0562 gm 
 1.0562 
 
 9999 gm 
 9853 
 
- 
 
 f 
 
 
 
 ■ 
 
 
 
 
 
 . 
 
 
 ■ 
 
16 
 
 The results in these determinations were low because the 
 sodium sulfite reduced some of the silver nitrate to free 
 silver on heating. In order to avoid this complication a 
 series of experiments were conducted in which the absorbing 
 solution was not heated. Results were: 
 
 Br Present Br Found 
 
 .9865 gm 
 .9865 
 .9865 
 .9865 
 
 .9930 gm, 
 .9870 
 .9943 
 .9898 
 
 These values are high -- presumably because varying amounts 
 of chlorine were carried over during the distillation. 
 
 The next determinations were made by absorbing the 
 bromine in sodium sulfite solution, then acidifying with 
 sulfuric acid and boiling, so as to expell the sulfur diox- 
 ide. Silver nitrate solution was then added, and the bro- 
 mine determined in the usual way. Only a small excess of 
 oxidizing agent was used in the reaction flask beyond what 
 was theoretically required to liberate the bromine, so as 
 not to oxidize the chlorine if possible. Results: 
 
 Br Present Br Found 
 
 .9865 gto. .9822 gm. 
 
 .9865 .9843 
 
 These determinations show that the oxidation of the bromides 
 was not complete, and that an excess of oxidizing agent is 
 necessary to effect its quantitative expulsion. 
 
 Cxi dizing Agents — The chromic acid-hydrogen peroxide 
 mixture is not a very satisfactory oxidizing medium, be- 
 cause if the mixture is heated long enough to drive off all 
 

 
 
 
 . 
 
 . 
 
 . 
 
 . 
 
 
 . 
 
 . 
 
 . 
 
 
 - 
 
 . • 
 
 
 
 
17 
 
 of the bromine large amounts of chromyl chloride distill 
 over, and thus too large an excess of silver nitrate must 
 be used to precipitate all of the halogen present. 
 
 Another oxidizing agent experimented with was hydrogen 
 peroxide in the presence of sulfuric acid. It was observed 
 that hydrogen peroxide in the presence of concentrated sul- 
 furic acid will liberate bromine from sodium bromide solu- 
 tion at 50° - 50°, or even a lower temperature, while the 
 chlorine in sodium chloride solution appears unaffected until 
 the solution is brought to a boil. This would suggest that 
 it might be possible to liberate bromine almost quantitatively 
 from a solution of the halides by this reagent. In order to 
 ascertain if all the bromide could be oxidized quickly and 
 also to find out if it would be completed absorbed in sil- 
 ver nitrate solution to which hydrogen peroxide had been 
 added to act as a reducing agent, a series of experiments 
 was made in which no chlorine at all was present. The reac- 
 tions involved are as follows: 
 
 S JSaBr + H g 0 2 + H 2 S0 4 -* Na 2 SG 4 + 2 H 2 0 + Br 2 . 
 
 This is an oxidation reaction, the hydrogen peroxide acting 
 as the oxidizing agent. In the following reaction it con- 
 ducts itself as a reducing agent: 
 
 P. AgNOg + Er g + H g G g -» P AgBr + 2 BNOg + 0 g . 
 
 These experiments were carried out in the following manner: 
 
 20 cc. of the sodium bromide solution was placed in an as- 
 pirator bottle and 2 cc. of 30 ^ hydrogen peroxide and 5 cc. 
 of concentrated sulfuric acid added. The liberated bromine 
 

 - 
 
18 
 
 was then removed by aspiration and collected in nitrogen bulbs 
 containing 25 cc. of dilute silver nitrate solution and about 
 2 cc. of 30$ hydrogen peroxide. The solution in the aspira- 
 tor bottle became colorless after aspiration for about one- 
 half hour. The results were as follows: 
 
 Br Present 
 
 Br Pound 
 
 .4012 gm. 
 .4012 
 .4012 
 .4012 
 
 .4027 gm. 
 .4013 
 .4018 
 .4023 
 
 Why these results should be 
 
 somewhat high has not been ac- 
 
 counted for. However, it should be noted that this method 
 gives decidedly the most satisfactory results of all those 
 tried. 
 
 The following determinations were made on solutions of 
 sodium bromide in 7$ sodium chloride, using only a slight ex 
 ce3s of silver nitrate to collect the bromine: 
 
 Br Present 
 
 Br Found 
 
 .4012 gm. 
 
 .4012 
 
 .4012 
 
 .3656 gm. 
 
 .3683 
 
 .3678 
 
 The following were made using a larger excess of silver ni 
 trate: 
 
 Br Present 
 
 Br Pound 
 
 ,4012 gm. 
 .4012 
 .4012 
 .4012 
 
 .4119 gm. 
 .4147 
 .4110 
 .4163 
 
 This demonstrates the fact that even in fairly dilute solu- 
 tions of sodium chloride the sulfuric acid will liberate 
 hydrogen chloride, which may be more or less oxidized to 
 

 
 
 
 
 
 
 
 
 
 
 
 
 
 • 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 
 ' 
 
 . '1 - 
 
 . 
 
 - .. . 
 
 
 .f 
 
 . 
 
 . 
 
 . 
 
 . 
 
 
19 
 
 free chlorine. At any rate the silver precipitate contains 
 silver chloride. 
 
 Fixation of Bromine b£ Metallic Sjiver --An interesting 
 possibility suggested itself in the expedient of absorbing the 
 bromine by passing the aspirated gases through a tube con- 
 taining finely divided silver. The principle upon which this 
 conjecture was based is the very general belief that bromine 
 will react with metallic silver while hydrochloric acid will 
 not. The procedure was as follows: Very fine crystals of 
 
 pure silver were prepared by electrolyzing silver nitrate 
 between a silver anode and a platinum cathode. These were 
 carefully washed and dried, and placed in a tube through 
 which the gases were to be drawn. Upon trial, however, tak- 
 ing no precautions to dry the gases, the increase in weight 
 was too high, indicating that moist hydrochloric acid will 
 react with fine divided silver. Upon introducing between 
 the tube and the flask wash bottles containing concentrated 
 sulfuric acid to dry the gases, the increase in weight was 
 practically nil, showing that silver will not combine with 
 dry bromine. 
 
 Theory of Oxidation of Bromides and Chlorides - -The fact 
 that no one has found an oxidizing agent that will selectively 
 oxidize the bromine finds a complete explanation in thermo- 
 dynamic considerations. According to thermodynamics, the 
 chemical work done in these oxidations depends wholly upon 
 the "initial and final states" and is "independent of the 
 
Loy<y/-/ f/?/77s of O/ ft/ f/o/?s 
 
 /oo /ooo / 
 
 A60 
 
20 
 
 path" -- that is, the oxidizing agent used. By "initial and 
 final states" is meant particularly concentration conditions* 
 This doctrine teaches, therefore, that the expenditure of 
 energy in effecting the oxidation of a certain amount of 
 bromide ions at definite concentration to bromine is a con- 
 stant, and that a certain "thermodynamic potential" is nec- 
 essary for the process. As a measure of this thermodynamic 
 potential may be used the electrical potential necessary to 
 bring about electrically the oxidation, which of course is 
 equal to the change in "free energy" involved. It is pos- 
 sible, of course, that a particular oxidizing agent might be 
 somewhat "faster" than another agent, and in this way show 
 selective action. But for similar ions like Br“ and 01”, 
 this is not likely. 
 
 Figure II shows the relation between the concentrations 
 of the halogen ions, Br" and Cl”, and the oxidizing poten- 
 tial necessary to convert them into free halogens. These 
 graphs have been calculated by means of the well-known 
 Nernst formula from the values of the normal bromine and 
 chloine electrodes. The temperature is assumed to be °5° C. 
 
 The horizontal axis shows the dilutions (that is, the recip- 
 rocals of the concentrations) expressed in terms of their 
 logarithms. Using 0.1 milligram of bromine as the limit of 
 analytical error, the potential necessary to reduce the Br”- 
 ion concentration (using 100 cc . of solution) to that point 
 has been calculated by the figure to be 1.36 volts, referred 
 to the normal hydrogen electrode as 0.0 volts. This potential, 
 

 
 
 
 
 
 
 
 
 
 
 
 
 
 .• 
 
 
 - 
 
 
 '■ 
 
 
 
 
 ■ 
 
 * 
 
 
 
 
21 
 
 however, as also shown by the figure, is sufficient to oxi- 
 dize Cl“=ions of concentrations greater than 0.6 N sodium 
 chloride -- that is, a concentration of 3.5 <f> by weight. 
 
 This shows very decidedly why selective oxidation methods 
 will not bring about a complete separation of bromine from 
 moderate concentrations of sodium chloride like 5 $ -- which 
 is a weak strength as brines and bitterns go. For higher con' 
 centrations, approaching saturated sodium chloride solution, 
 liberation of the chlorine may begin when 100 cc. of the 
 solution still contain as much as 4.5 milligrams of unoxi- 
 dized bromide. 
 
 Method Recommended --The results obtained in thi3 work 
 as discussed above indicate very conclusively that bromine 
 can not be separated quantitatively from an excess of chlo- 
 rides by selective oxidation and distillation, since oxida- 
 tion of the chloride begins before the bromine is completely 
 expelled. The best method is the following: The halide mix- 
 
 ture is placed in an aspirator bottle and 2 cc. of 30^ hy- 
 drogen peroxide and about 5 cc. of concentrated sulfuric 
 acid added. The halogens are collected in a known amount 
 of dilute silver nitrate containing a volume of pure hydro- 
 gen peroxide -- for example, 2, cc. of the 3 Oft product. After 
 the aspiration has continued for about three quarters of an 
 hour, the vessel containing the absorbing medium is discon- 
 nected and placed on a water bath until the solution "bright- 
 ens." Now dilute potassium chloride solution is carefully 
 added drop by drop as long as a precipitate forms. The pre- 
 
~ 
 
 
 
 
 
 
 - 
 
 ■ 
 
22 , 
 
 cipitat© is then treated in the usual way, and its composi- 
 tion calculated by the expression on page 5. 
 

 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
23 
 
 IV . Summary , 
 
 I, Indirect methods for calculating bromides in the 
 presence of chlorides are accurate only when the chloride 
 concentration is low. 
 
 P. Slow precipitation of silver halides from ammonium 
 hydroxide and from alcoholic solutions give unsatisfactory 
 results in the presence of excess chlorides. 
 
 3. Low results for bromine are due to the solubility 
 of silver bromide in alkaline halides. 
 
 4. Empirical corrections for low results are not feas- 
 ible, since the solubility of silver bromide varies with 
 its state of subdivision. 
 
 5. In selective oxidation processes the halogens are 
 distilled into an absorption medium; sodium sulfite was stu- 
 died as a reducing agent for this medium and proved trouble- 
 some by reducing the silver nitrate at high temperatures. 
 
 6. Hydrogen peroxide was found very satisfactory as 
 an oxidizing agent for liberating the bromine, and as a 
 reducing agent for reconverting it to the bromide form. 
 
 7. An excess of oxidizing agent is desirable in the 
 liberation of the bromine from the halide solution. 
 
 8. There is always a simultaneous liberation of chlo- 
 rine along with the bromine in the oxidation of solutions of 
 bromides and chlorides, particularly if the latter are pres- 
 ent in quantity. 
 
 9. The simultaneous liberation of bromine and chlorine 
 

 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 
 . 
 
 
 
 
 
 
 
 
 1 
 
 
 
 
 
 
24 
 
 is explained from thermodynamic considerations, which indi- 
 cate that there is no oxidizing agent that will effect a 
 satisfactory selective oxidation of bromine in the presence 
 of chlorides in excess. 
 
 10. A new procedure of estimating bromine in the pre3 
 ence of chlorides is suggested. 
 

 
 
 
 
 
 
 
 
 
25 
 
 V. Bibliography . 
 
 1. Vortman: Z. anal. Chem., 25, 172 (1886). 
 
 2. Cavazzi: Gazz. Chim. Ital., 13, 174. 
 
 3. Rngle: Compt. Rend., 118, 1263 (1894). 
 
 4. Berglund: Z. anal. Cham,, 24, 184 (1885). 
 
 5. Baubigny and Rivals: Compt. Rend., 125, 527 (1827). 
 
 5. Wyes: Rept. anal. Chem., 5, 238 (1885). 
 
 7. White: Chem. News, 1, 14 4 (1892). 
 
 8. Jannasch and Aschoff: Z. anorg. Chem., 1, 144 (1892). 
 
 9. 'Bugarszky: Z. anorg. Chem., 10, 387 (1897). 
 
 10. Andrews: J. Am. C. S., 25, 809 (1903). 
 
 11. Gooch: J. Am. C. S., 29, 275 (1907); Am. J. Sci., 35, 
 
 54 (1913). 
 
 12. Skinner and Baughman: J, Ind. Sng. Chem., 11, 954 (1919). 
 
 13. Schierholz: Sitzber. K. Akad. Wi3s. (Vienna), 101, 2b, 
 
 4 (1890). 
 
 14. Abegg, Auerbach and Luther: Messungen elektromotori scher 
 
 Kraefte galvani scher Ketten (1911), p. 200.