QD Ft-Ws DflD EXCHANGE Oxidation and Reduction without the Addition of Acid I. The Reaction between Ferrous Sulfate and Potassium Dichromate II. The Reaction between Stannous Chloride and Potassium Dichromate A DISSERTATION SUBMITTED TO THE FACULTY OF THE GRADUATE SCHOOL OF TH IXIYERSITY OF PITTSBURGH IN CONFORMITY WITH THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY BY JOSHUA CHITWOOD WITT PITTSBURGH 1915 EASTON, PA.: ESCHENBACH PRINTING Co. 1916 Oxidation and Reduction without the Addition of Acid I. The Reaction between Ferrous Sulfate and Potassium Bichromate II. The Reaction between Stannous Chloride and Potassium Dichromate A DISSERTATION SUBMITTED TO THE FACULTY OF THE GRADUATE SCHOOL OF THE UNIVERSITY OF PITTSBURGH IN CONFORMITY WITH THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY BY JOSHUA CHITWOOD WITT PITTSBURGH 1915 EASTON, PA.: ESCHENBACH PRINTING Co. 1916 ACKNOWLEDGMENT. The writer is greatly indebted to Dr. Marks Neidle, under whose super- vision this work has been carried out, for his kind consideration and help- fulness. OXIDATION AND REDUCTION WITHOUT THE ADDITION OF ACID. I. THE REACTION BETWEEN FERROUS SULFATE AND POTASSIUM BICHROMATE. BY JOSHUA C. WITT. The first use of the reaction between ferrous salts and dichromate for the determination of iron was made by Penny. 1 In the method, as described in his paper, a sample of "iron stone" was dissolved in hydro- chloric acid, and the iron reduced by adding sodium sulfite in excess. After boiling off the excess sulfurous acid, he titrated with dichromate solution, using potassium ferricyanide as an outside indicator. It is interesting to note that this method is essentially the same as that in use today for the determination of iron in iron ore. Whenever the reaction between ferrous salts and dichromate has been studied a mineral acid has been added. Penny employed excess of free acid in dissolving the iron ore, and the equation for the reaction demands free acid for the formation of the normal salts of potassium, chromium, and iron. Since no mention of any investigation of the reaction in the absence of free acid could be found in the literature, it was decided to perform a few preliminary experiments in which a quantity of ferrous sulfate was titrated by o.i N dichromate, with and without acid. It was considered preferable to weigh out a separate portion of ferrous sulfate for each titration, rather than to keep a standard solution of the salt. As soon as a portion was weighed out it was rapidly transferred to a beaker containing water, and titrated at once with the dichromate solution, using potassium ferricyanide as an outside indicator. The following results were obtained showing the effect of acid: FeSO4.7HO (g.). Cc. KjCrjOr. HiSO 4 . 0.8 29.52 Excess present 0.8 29.54 Excess present 0.8 36.16 None present 0.8 36.25 None present The end point is obtained when the amount of ferrous salt remaining at the time the drop test is made is insufficient to affect the indicator. When no acid is added, an excess of dichromate is required to give an end point, which means that with the theoretical amount of dichromate necessary to completely oxidize the ferrous sulfate, enough of the latter remains to affect the indicator, i. e., the reaction is incomplete. A brown precipitate appears after a few drops of the dichromate have been added. The following results show the effect of the volume of ferrous sulfate solution on the titration, 0.8 g. of salt being used in each experiment, 1 Brit. Assoc. Rep., [2] 1850, 58, 59. Cc. water Cc. dichromate Cc. dichromate added to PeSO. sol. with HjSO*. sol. without HjS o 29.67 29.89 5 :.... 29.88 15 30.03 30 30.62 IOO 32 . 12 looo 53-00 This increase in the dichromate was to be expected, since the reaction is slower the greater the volume, and larger amounts of dichromate are required to drive the reaction to the end point. When no water is added the result of the titration is nearer theoretical, and in several experiments, in which more than 0.8 g. was taken and the solid titrated, the result was exactly the theoretical. We may therefore conclude that the precipitate formed does not adsorb the ferrous ion appreciably. Adsorption of ferrous ion would vitiate the results on the velocity of the reaction. Considerable difficulty was encountered in finding the end point at the higher concentrations when the titration was made in the absence of sulfuric acid. The brown precipitate had a tendency to mask the end point. To overcome this, when the end point was nearly reached, it was found necessary to filter a few drops of the mixture each time be- fore it was applied to the indicator. Measurement of the Velocity of the Reaction. The problem which presented itself at this point was to find a method of determining the unoxidized ferrous salt or unreduced dichromate in a solution containing ferric salts, ferrous salts, chromic salts, and dichromate. Three methods suggested themselves: (1) To stop the reaction by the addition of ammonium hydroxide, filter the precipitated hydroxides of iron and chromium, and determine the chromium in the precipitate. (2) To add ammonium hydroxide as in (i) and titrate the unchanged dichromate in the filtrate. (3) To precipitate the unchanged dichromate with lead acetate, dis- solve the precipitate of the reaction in acetic acid, and determine chro- mate in the residue. The second method, being more direct and therefore more accurate, was adopted. It is well known that ferrous salts, in common with salts of other di- valent metals, cannot be completely precipitated by ammonium hydroxide in the presence of ammonium salts in consequence of the repression of hydroxyl ion by the latter. In order to be certain of completely precipi- tating ferrous iron, the necessary conditions were investigated. It was found that ferrous salts may be completely precipitated with ammonium hydroxide provided, (1) The solution is neutral. (2) No ammonium salts are present to begin with. (3) The concentration is sufficiently low. (4) The solution is boiled and the precipitate allowed to settle before filtration is attempted. If 0.5 g. portions of ferrous sulfate were dissolved in various volumes of water, and an excess of ammonium hydroxide added, the precipitation was complete only when the volume was at least 100 cc. Solutions. Potassium Dichromate. A o.i N solution, standardized against iron wire, was kept in a ten-liter bottle fitted with a siphon. All air entering the bottle came through a cotton plug to avoid contamination. Sodium Thiosulfate. A o.oi N solution was standardized each time before using against the dichromate solution. Ferrous Sulfate. At first it was thought advisable to make up a stand- ard solution of ferrous sulfate and attempt to protect it from oxidation, but it was finally decided to use the dry salt and weigh out a portion for each determination. To avoid difficulty from any variation in quality, a fresh pound bottle of the c. P. salt was taken, and used for all the work. The surface layer was discarded, a weighing bottle filled and kept in the balance until used, then refilled when necessary. In weighing out a sample, a slight excess was placed on the balance and the stopper of the weighing bottle replaced at once. The excess salt was removed as quickly as possible and discarded to avoid any possible contamination. The salt was analyzed from time to time and found to remain constant in compo- sition, as shown by the following results obtained with 0.8 g. samples: Date April 5. May 3. June 9. Titration with 0.1014 N K 2 Cr 2 O 7 29.53 cc. 29.67 cc. 29.63 cc. Manipulation. A large, electrically controlled bath was maintained at 30 == 0.05. A ten-liter bottle of distilled water was kept in this bath that no delay might be caused by waiting for water to assume the correct temperature. Nearly as much water as was needed for the experiment was placed in a liter flask corrected for temperature, and a given amount of standard dichromate solution was run into an Erlenmeyer flask. Both flasks were immersed in the bath and allowed to assume constant temperature. In the meantime a portion of the ferrous sulfate was weighed and rapidly transferred to the liter flask. When solution was complete, the dichromate was added and the volume adjusted. The flask was kept in the bath and, at intervals, 100 cc. portions were removed and run into beakers containing excess of ammonium hydroxide. The precipitate formed by the ammonium hydroxide was very finely divided and would pass very readily through the filter. It was 8 found necessary to let it stand some time preferably over night or to boil it a few minutes before a complete filtration could be made. It was preferable to filter at once, without heating, but no method could be found which gave the desired result. The precipitate passed through an alundum Gooch, and would not settle when kept in a centrifuge for 15-20 minutes. In order to determine whether the potassium dichromate still in solution was in any way affected by the precipitated ferrous iron, 25 cc. of o.i N dichromate was added to 0.8 g. of ferrous sulfate dissolved in water in a liter flask, and, after introducing an excess of ammonium hydroxide, the mixture was made up to volume. A number of 100 cc. portions were withdrawn and placed in beakers. They were filtered at various intervals and titrated with o.oi TV thiosulfate by the method already described. Some were boiled before filtering, and others were filtered in the cold. The following are the results obtained : Time. Titration, 0.01 N thiosulfate. Remarks. 2 hours o . 90 cc. Not boiled 2 hours i . oo Boiled 24 hours o . 94 Not boiled 24 hours i . oo Boiled It is seen from the above that the final result is not altered by allowing the mixture to stand for many hours, or by boiling, before filtration. Results. All measurements were made at 30. The ferrous sulfate and dichromate solution were in the ratio of 0.8 g. of the former to 25 cc. of the latter, or 2.878 mols to 0.4225 mol. It was not thought advisable to attempt any measurements with more dichromate than would be re- quired for the normal end point, since in this case a very large volume of o'.oi N thiosulfate would be required. The ratio of the reacting substances was maintained constant. The volume of dichromate reduced is repre- sented by x and that unreduced by a x. TABLE I. Total Volume Containing 25 cc. of 0.1014 N K 2 Cr 2 O 7 and 0.8 g. FeSO 4 . 100 cc. 250 cc. 500 cc. 1,000 cc. 2,000 cc. 4,000 cc. 5,000 cc. utes. a x. x. a *. x. a x. x. a *. x. a x. x. a x. x. a x. x. I 0.06 24.94 0.22 24.78 0.35 24.65 1.50 23.50 1.93 23.07 5 0.03 24.97 0.18 24.82 0.34 24.66 1.46 23.50 1.99 23.01 2.69 23.31 15 0.04 24.96 0.23 24.77 0.30 24.70 0-93 24.07 1.74 23.26 1.65 23.35 2.20 23.80 30 0.04 24.96 0.22 24.78 0.29 24.71 0.80 24.20 1.36 23.64 1.57 23.43 2.20 23.80 60 0.04 24.96 0.19 24.81 0.26 24.74 0.61 24.37 i .06 23.94 2.31 23.69 One series of experiments was made with method number three as a check. The work was carried on in the same way up to the time when 100 cc. portions were removed from the liter flask. In this case they were run into beakers containing lead acetate solution, which precipitated the sulfate ion and the chromate ion. The mixture was then acidified with a few cubic centimeters of acetic acid and boiled to dissolve all the iron salts. The lead salts were then filtered out and the lead chromate dis- solved in dilute hydrochloric acid. The resulting dichromate was titrated, after cooling, with o.oi N thiosulfate. The results given in Table II compare satisfactorily with those previously obtained and given in Table I. TABLE II. Minutes. Method III. Method II. 5 1-49 1.46 15 0-98 0.93 60 0.59 0.6i Comments on Velocity Measurements and the Order of the Reaction. From Table I it is seen that in the titration of 0.8 g. of ferrous sulfate with dichromate, the reaction is 99.8% complete at the end of one minute, provided the final volume is 100 cc. This statement may be made even though in our experiments the dichromate taken was a little less than equivalent to the ferrous sulfate. At all other concentrations except the most dilute, the reaction is more than 90% complete at the end of the first minute. The data as obtained are not of a nature to permit ready calculation of the order of the reaction, although those in the last column of Table I seemed sufficiently regular to justify an attempt at such a calculation. No constant could be obtained by assuming the reaction to be of the first order with respect to each of the reacting substances, of the first order with respect to one and of the second with respect to the other, and, finally, of the second order with respect to both. Our conclusion, therefore, is that this reaction is probably of an order higher than the fourth. The rate of oxidation of ferrous sulfate by dichromate with the addition of more than the sulfuric acid required by the normal equation has been investigated by Benson. 1 It is stated in his conclusions that the rate is proportional to the second power of the concentration of ferrous salt, and to the second power of that of the acid, and that the order is variable with respect to the dichromate. Benson also found that the order is much retarded by the presence of ferric salts. If the velocity of this re- action is strictly proportional to the square, or any other power, of the concentration of acid added it should be zero when no acid is employed. There can be no question, however, that the velocity is proportional to some power of the hydrogen-ion concentration, in which case the ve- locity of the reaction without the addition of acid is due to the hydrogen- ion concentration arising from the hydrolytic dissociation of both dichro- mate and ferrous salt. The concentration of hydrogen ion must play an 1 /. Phys. Chem., i, i (1903). 10 important part in the reaction, even in very low concentration. Our reaction is most probably accompanied by a change in the concentration of hydrogen ion, which was disregarded in our velocity calculations. For this reason, we can not conclude with certainty that the reaction is of an order higher than the fourth. The great velocity of the reaction without the addition of acid is partly due to the fact that less than one-third of the iron remains in solution as ferric salt, which has a retarding influence, while the remainder pre- cipitates in the form of hydrous ferric oxide and adsorbed ferric sulfate. The Products of the Reaction. Preliminary experiments showed that all the brown precipitate, ultimately formed when solutions of dichromate and ferrous sulfate are mixed, does not come down instantly,but gradually, reminding one of the precipitation of suspensoids by small quantities of electrolytes. Upon filtering the mixture after it had stood for several days, the filtrate still yielded apparently the same precipitate on standing. The precipitation, it was found, could be rendered complete by boiling, when a reddish brown, gelatinous precipitate, resembling ferric hydroxide, appeared. One-tenth of the equivalent weights of potassium dichromate and ferrous sulfate were dissolved in water and the solutions mixed, diluted nearly to a liter and heated to boiling for several minutes to bring about com- plete precipitation. After cooling, the mixture was made up to a liter exactly. The precipitate was brown and very abundant. The superna- tant liquid had the purplish green color characteristic of chromium salts. It was thought that heating the mixture might have some effect on the reaction. To settle this point, another solution was made up exactly like the one already described, except that it was not heated. After standing over night, it was filtered and both precipitate and filtrate were analyzed along with those of the first mixture, giving practically the same results. Although the precipitation was not complete, the differ- ence was practically negligible. The work on this second solution was dropped, therefore, and only the first carried on. The brown precipitate from the first mixture was dried to constant weight at 100-105, giving a very hygroscopic amorphous powder. This solid and also the filtrate were analyzed for SO 3 , Cr 2 O 3 , and Fe 2 O 3 . The SOs was determined as BaSO4. To determine iron and chromium, the hydrochloric acid solution was neutralized with sodium hydroxide and the chromium oxidized by sodium peroxide. After boiling, the ferric hydroxide was filtered out and washed with hot water. The precipitate was then dissolved in hot hydrochloric acid, reprecipitated with sodium hydroxide, again treated with sodium peroxide, filtered and washed. The two filtrates were combined, boiled, acidified with hydrochloric acid (5 cc. in excess) and again boiled for some time. After cooling, 10 cc. of Grams. Gram Grams equivalents. originally present. SOs.... Fe . 6.241 , 1.176 1. 122 0.1560 0.0631 o . 0647 PRECIPITATE. 8.006 5-590 1-733 Cr Fe 2 O 3 Grams. 6.313 Gram equivalents. 0.2369 0.0353 o . 0440 o.oon Cr 2 O 3 o 894. SOs I 7SS Loss on ignitic Undetermined n (except SO 3 ) . . 2 . 239 (K 2 0).. . 0.067 II a 10% potassium iodide solution was added, and the solution titrated with o. i N thiosulfate, using starch as the indicator. The iron was again dissolved, brought nearly to dryness on the hot plate, reduced by stannous chloride and titrated with o.i N potassium permanganate. FILTRATE. Percentage of total in precipitate. 22.05 78.96 35-26 Percentage. 56.03 7-93 15-58 19.87 0.59 Further Investigation of the Precipitate. It is seen that the pre- cipitate contains quantities of all the salts produced in the reaction. In order to ascertain the nature of these adsorbed salts, a weighed portion of the precipitate was boiled in water for some minutes, and filtered. The filtrate was made up to 250 cc., and 25 cc. portions removed for anal- ysis. It was found that the nitrate contained 2. 1 1% SO 3 , calculated on the basis of the amount of precipitate taken; or, 13.54% of the SO 3 present in the original precipitate had been removed by the first boiling. A 50 cc. portion of the filtrate, analyzed for iron, gave 0.69%, approxi- mately the amount required to correspond to the formula Fe 2 (SO 4 ) 3 . We may therefore conclude that the adsorbed salt is mainly Fe 2 (SO 4 )3. Discussion of Results on the Products of the Reaction. The value 35.26% for the amount of chromium in the precipitate suggests that one-third of the total is precipitated as Cr 2 O3 and the remaining 1.93% adsorbed as Cr 2 (SO 4 ) 3 . If we add the number of equivalents corresponding to the adsorbed potassium sulfate and chromium sulfate, and subtract the sum from the total number of equivalents of SO 3 in the precipitate, the result gives the number of equivalents of Fe 2 (SO 4 )3 adsorbed. This value is 0.0407, which, added to the number of equivalents of Fe 2 (SO 4 ) 3 in the filtrate (0.0631), gives the number of equivalents of this salt formed in the reaction (0.1038). The mixture contained sufficient iron for 0.3 equiva- lents of Fe 2 (SO 4 ) 3 . Thus two-thirds of the iron forms hydrous ferric oxide, and one- third forms ferric sulfate. The following equation completely harmonizes with the above results: 3K 2 Cr 2 7 + i8FeS0 4 + (x + 6y)H 2 O = Cr 2 O 3 .*H 2 + 2Cr 2 (SO 4 ) 3 -f 3Fe 2 (SO 4 ) 3 12 where Cr 2 O 3 .#H 2 O and Fe 2 O 3 .;yH 2 O stand for the colloidal oxides of chro- mium and iron, each carrying adsorbed water. The products of the reaction between potassium dichromate and ferrous sulfate without the addition of acid are: potassium sulfate, chromium sulf ate and colloidal chromic oxide in the molar ratio of 2:1; and ferric sulfate and colloidal ferric oxide in the molar ratio of 1:2. The colloids are precipitated by the sulfate ion in the solution. The normal ionic reaction is written Cr 2 O 7 " + 6Fe++ + ^H ^1 2Cr+++ + 6Fe+++ + 7H 2 O. We believe that the reaction without acid proceeds in the same way, the hydrogen ion being derived from the water. H 2 ^1 H+ + OH". As hydrogen ion is consumed by the reaction, more is formed, and at the same time hydroxyl ion accumulates. Soon the concentration of hydroxyl ion is sufficient to exceed the solubility products of the hydroxides of iron and chromium, and the colloidal hydrous oxides are formed. Fe+++ + 3 QH- 5 Fe(OH) 3 ; 2 Fe(OH), + (y 3 )H 2 O ^ Fe 2 O 3 .jH 2 O. Cr+++ + 30H- ^ Cr(OH) 3 ; 2 Cr(OH) 3 + (* 3 )H 2 O Z Cr 2 O 3 .*H 2 O. Summary. 1. The stoichiometric relations in the reaction between potassium dichromate and ferrous sulfate are the same with or without acid. 2. The experimental conditions for the complete precipitation of ferrous iron by ammonium hydroxide have been found and employed to determine dichromate in a mixture also containing ferrous, ferric, and chromium salts. 3. Without acid the reaction is instantaneous, except in very dilute solutions. 4. Disregarding the change of hydrogen-ion concentration accompanying the reaction, the order is higher than the fourth. 5. The rate of the reaction, with acid, can not be proportional to the second power of the concentration of acid added, for then it should be zero without acid. 6. The products of the reaction are the sulfates of potassium, chromium, and iron, and the colloidal hydrous oxides of iron and chromium. The latter are precipitated by the sulfate ion, and adsorb a large quantity of ferric sulfate and smaller quantities of the other two sulfates. 7. The equations for the reaction have been formulated. - OXIDATION AND REDUCTION WITHOUT THE ADDITION OF ACID. H. THE REACTION BETWEEN STANNOUS CHLORIDE AND POTASSIUM DICHROMATE. A CONTRIBUTION TO COLLOID-CHEMISTRY. It may be stated from the results of the first paper of this thesis that colloidal hydrous oxides or hydroxides are obtained in an oxidation- reduction reaction, in which acid must be added for the formation of normal salts, if the stoichiometric relation is the same without acid as with acid. If the reaction involves ions which are good precipitants of the colloids formed, precipitation takes place; otherwise, hydrosols are obtained. The equation for the reaction between stannous chloride and potassium dichromate, with acid, is 3SnCl 2 + K 2 Cr 2 7 + i 4 HCl == aSnCU + 2CrCl 3 + yH 2 O + 2KC1, where, it is seen, fourteen mols of hydrochloric acid per mol of dichromate are necessary to form the normal salts of tetravalent tin, trivalent chromium, and of potassium. The object of this investigation was to determine whether the stoichiometric relation between dichromate and stannous chloride is the same, i. e., one mol of the former oxidizing three mols of the latter, and what substances are formed when no acid is added. The Stoichiometric Relation. Samples of commercial c. P. stannous chloride of about 0.4 g. each were rapidly transferred to beakers from a weighing bottle, dissolved in 50 cc. of water, and the solutions titrated with standard dichromate, some after adding 10 cc. of concentrated hydrochloric acid and others without the addition of any acid. Ferrous ammonium sulfate solution containing potassium thiocyanate was employed as an outside indicator. At first the results of the titrations without acid seemed to be slightly higher than those with acid, which, as in the titrations of ferrous sulfate with dichr ornate without the addition of acid, would indicate that the reaction was not instantaneous. Further investigation, however, showed that the hydrochloric acid alone gave a faint pink color with the indicator, which was caused by the ferric iron in the slightly oxidized ferrous sulfate of the indicator being brought into solution by the acid. To overcome this difficulty, the titrations in which acid was used were run until a drop of the solution gave a darker tint with the indicator than acid alone. The results with and without acid were then almost identical. Therefore the oxidizing power of dichromate towards stannous chloride is not affected if the reaction takes place without the addition of acid. Furthermore, the reaction is practically instantaneous in dilute solutions, or, in the titrations referred to above, the results should be higher without acid than with acid. It is not surprising, however, that this should be the case, for the reaction may be regarded as compounded of two, each of which has a very great velocity, namely, that between ferrous salt and dichromate and that between ferric salt and stannous salt. When no acid is used, stannous chloride is considerably hydrolyzed in the concentrations employed in our titrations, giving milky, opalescent, solutions. The dichromate rapidly reduced the turbidity, which dis- appeared after a few cubic centimeters had been added. It is also in- teresting to note that at the end of these titrations, the solutions possessed a peculiar, faint and yet distinct, fruity odor, which was not observed in the acid titrations. We have been unable to discover the cause of this odor. In titrating the solid salt with o.i N dichromate, an olive-green solution having no turbidity is obtained immediately. The Products of the Reaction. The percentage of stannous tin contained in the solid chloride was estimated by titration with o. i N dichromate in the presence of an excess of hydrochloric acid, and the amount containing an equivalent weight in grams calculated. This quantity, 119.6 g., was dissolved in about 300 cc. of water contained in a liter flask. An equivalent weight of potassium dichromate (49.03 g.), enough to completely oxidize the stannous chloride, was dissolved in 200-300 cc. of water contained in a beaker. The di- chromate solution was gradually added to the stannous chloride solution, the mixture being shaken vigorously to secure homogeneity. During the process of mixing, brownish and greenish blue gelatinous masses were formed and at one point the entire mixture became a jelly; but when all the dichromate had been added a perfectly clear, deep, olive- 15 green solution resulted, which in sufficient depth appeared red by trans- mitted light, natural or artificial. The mixture was diluted to a liter and aliquot portions removed for investigation. Treatment with Ethyl Alcohol. One hundred cubic centimeters were evaporated to dryness on a steam bath and dried to constant weight in an air oven. The residue was treated with 95% ethyl alcohol, which dis- solved all but a white crystalline substance slightly tinged with green. This alcohol-insoluble matter was filtered off by suction through a Biichner funnel and washed with 95% alcohol, but it could not be en- tirely freed from the slight coloration due to an adsorbed chromium compound. Alcohol-Insoluble Matter. The residue from 100 cc. of the original mixture was dissolved in water and made up to 250 cc. Twenty-five cubic centimeter portions were removed for the estimation of tin, chromium and chlorine. The tin was determined by precipitation with hydrogen sulfide and ignition to stannic oxide; the chromium by addition of ammonium hydroxide to the hydrogen sulfide filtrate; and the chlorine by precipitation with silver nitrate. The potassium was obtained by subtracting the amount in the alcohol-soluble matter from the total employed in the reaction. Alcohol-Soluble Matter. The alcohol solution was evaporated to dryness on a steam bath, and the residue, dried to constant weight in an air oven at 120, ground and analyzed for potassium, chlorine, chromium and tin. The methods for the tin and chlorine were the same as those above, while the potassium was determined as potassium chloride, and the chromium by fusion with sodium peroxide and titrating the resulting chromate with thiosulfate. The results calculated to totals for the entire mixture are as follows: Alcohol-insoluble matter. Alcohol-soluble matter. K Cr m . Sn iv . Grams. Gram equivalents. Grams. Gram equivalents. JLUIH.1 gram equiv. 12. II 0.3097 0.92 0.0235 0.3332 II. 12 0.3133 19.01 0.5361 0.8494 0.15 0.0087 17.15 0.9894 0.9981 0-45 O.OI5I 60.39 2.0300 2.0451 The quantities of the elements contained in the entire mixture are: potassium, 0.3333 equivalent; chromium, i equivalent; chlorine and tin, i and 2 equivalents, respectively, provided the stannous chloride did not contain stannic tin. It will be remembered that the weight of stannous salt containing one-half the molecular weight of unoxidized chloride was employed. The results show that 0.0451 equivalent of stannic tin was present, with which 0.0226 equivalent of chlorine was associated. Thus, the total chlorine should be 1.0226 equivalents. i6 It is evident that the substance separated in the alcohol treatment is potassium chloride, which therefore is one of the products of the reaction. The constituents of the alcohol-soluble matter cannot be determined from the analysis alone. It may be observed, however, that there is a deficiency of chlorine of nearly one-sixth the total. This loss was in- curred when the alcohol-soluble matter was dried in the air oven, and was due to the decomposition of stannic chloride or chromic chloride, or both. Since no tin was lost in the process of drying, hydrated stannic chloride could not have been present to any appreciable degree, for hydrated stannic chloride volatilizes considerably when heated. 1 Hydrated chromic chloride on the other hand does yield hydrochloric acid when heated to 120 C. 2 We are thus led to the conclusion that the alcohol-soluble matter consists of a mixture of the hydrous oxides of tin and chromium, and hydrated chromic chloride. If stannic chloride were a product of the reaction, it could be extracted by means of carbon bisulfide. A portion of the original mixture was evaporated to dryness on a water bath, the residue powdered, introduced into a thimble, and extracted with carbon bisulfide for about eighteen hours, but no trace of tin could be found in the solvent. Dialysis of the reaction mixture gave further evidence that it does not contain stannic chloride. Dialysis. Fifty cubic centimeters of the original solution were dialyzed in a parchment paper bag suspended in a beaker filled with distilled water to the level of the solution in the bag. In a short time the external liquid was colored bluish green and considerable osmosis had taken place. Fresh water was placed in the beaker every day until it was not per- ceptibly colored after standing twenty-four hours. The accumulated diffusate for this period was concentrated and tested for tin, but none was found, thus proving the absence of stannic chloride in the mixture. The dialysis was continued in order to free the hydrosol as far as possible from electrolytes. Excessive dilution by osmosis was avoided by keeping the level inside the membrane several centimeters higher than outside, which resulted in the concentration of the colloid as the osmosis dimin- ished. In about five weeks when the diffusate was giving only a faint test for chloride ion, the contents of the bag set to a perfectly clear gel of a beautiful, emerald green in reflected light, and a deep red in trans- mitted light. In a second dialysis where no effort was made to keep the solution in the membrane from increasing in volume, 250 cc. in- creased to 1 1 oo cc. in five weeks, giving a clear hydrosol resembling the 1 Gmelin-Kraut, 4, I, 313. 1 Itnd. t 3, 1, 439- 17 hydrogel. The hydrosol may be boiled down to a very viscous con- sistency and on being dehydrated over sulfuric acid becomes a firm gel. Analysis of the gel showed that it contained all the tin, practically one- half of the chromium and a negligible quantity of chlorine. These re- sults, together with those previously obtained, enable us to formulate the reaction as follows : 2K 2 Cr 2 7 + 6SnCl 2 + (6* + ?)H 2 O ^ 4KC1 + 6SnO 2 .#H 2 O + Cr 2 O 3 .;yH 2 O + 2CrCU + 2HC1 Thus written the equation expresses the fact that dialysis yields a mixed hydrosol of hydrous stannic and chromic oxides in the molecular ration of 6 SnO 2 to i Cr 2 O 3 . The reaction, with acid, is written ionically Cr 2 7 - + 3811++ + HH+ Z^ 2 Cr+++ + 3 Sn ++ +++ 7 H 2 O. (i) Though no acid is added in our reaction, hydrogen ion is present, due to the hydrolytic dissociation of the dichromate and stannous chloride. Equation i, therefore, represents the reaction when no acid is added. Aqueous solutions contain hydroxyl ion in equilibrium with hydrogen ion according to the equation: H+ + OH- ^ H 2 (2) As the hydrogen ion is removed by reaction (i) this equilibrium is dis- turbed, the more so as the available hydrogen is very limited and the concentration of hydroxyl ion correspondingly increases. When the latter accumulates sufficiently, the solubility products of stannic and chromic hydroxides are exceeded and the colloidal hydrous oxides are formed 3 OH- ^ Cr(OH) 3 ; 2Cr(OH) 3 + (y 3 )H 2 O ^ Cr 2 O 3 .;yH 2 O. 4 OH- ^ Sn(OH) 4 ; Sn(OH) 4 + (x 2)H 2 O ^SnO 2 .^H 2 O . The formation of the hydrosols maintains the concentration of hydroxyl ion within perfectly definite limits, which insures a definite minimal concentration of hydrogen ion sufficient for reaction (i), which there- fore continues to completion. Conclusions. 1. The stoichiometric relations in the reaction between potassium dichromate and stannous chloride are the same with or without acid. 2. The products of the reaction are potassium and chromium chlorides, and stannic and chromic hydrous oxides in colloidal solution. 3. A clear mixed hydrosol of stannic and chromic hydrous oxides, approximately in the molar ratio of 6 SnO 2 to i Cr 2 O 3 , may be ob- tained by adding an equivalent amount of dichromate solution to stan- nous chloride and dialyzing the mixture. The hydrosol will contain all of the tin and practically one-half of the chromium used in the re- action. 4. The equations for the reaction have been formulated. VITA. Joshua Chitwood Witt was born at Connersville, Indiana, 1884. He attended school at Connersville and Liberty, Indiana, and was graduated from high school in 1903. He entered Butler College, Indianapolis, in 1904, receiving the degree of Bachelor of Arts in 1908, with chemistry as his major subject. The same year he took up the study of chemistry and bacteriology at the University of Chicago, which culminated in the degree of Bachelor of Science. From 1909 to 1911, he did special work in mechanical engineering at Armour Institute of Technology. In 1911 he entered the University of Pittsburgh to pursue graduate work in chemistry and physics, receiving the degree of Master of Science the following year. BATB AN INITIAL FINE OF 25 CENTS OVERDUE. W6~8 1993 LD 21-100m-8,'34 Photomount Pamphlet Binder Gaylord Bros. Makers Syracuse, N. Y. PAT. JAN 21, 1908 U.C. BERKELEY LIBRAR m UNIVERSITY OF CALIFORNIA LIBRARY