268 THE ANALYST. EXPERIMENTS ON TEE DETERMINATION O F THE FREE ACIDS OF VINEGAR. BY ALFRED H. ALLEN AND R. BODMER. Read before the Society of Public Analysts, on the 20th Harch, 1878. WITH the view of ascertaining the extent to which the known methods of determining the acids of vinegar could be trusted, we have instituted a series of experiments on repre- sentative mixtures of known amounts of the constituents of vinegar.As a starting point, we prepared a pure acetic acid by distilling the commercial acid after addition of a little soda. The distillate had a density of 1.0396 at 15” C. According to Oudemann, this number corresponds t o 28.67 per cent. of real acetic acid (C, H, 0,), or, according to Mohr, to 29.5 per c a t . The figures of the latter chemist’s table of densities are only carried to three placerr of decimals.Weighed (not measured) quantities of the above sample of acid were next titrated with decinormal caustic soda, using litmus as an indicator. The results shewed 28-54, 28.44, 28.49, and 28.59 per cent. of real acid, the average being 28.515. Another titration, in which a few drops of cupric sulphate were employed instead of litmus (a permanent turbidity being taken as the end of the reaction), gave 23-52 per cent.of acetic acid.” It will be seen, therefore, that the two methods of titration gave extremely close results, and that the amount of acid calculated from the density (by Oudemann’s table) was slightly higher than that found by titration. This result is in accordance with the general opinion that titration of acetic acid gives results slightly below the truth.A dilute acid was next made by mixing a quantity of the above sample with nine times its weight of water. A portion of it was then titrated, (using litmus), when it gave 2.853 per cent. of acid,-almost exactly one-tenth of the original amount. On the other hand the density was 1.0040, which corrcsponds t o 3.20 per cent.of acid according to Oudemann, or 3.0 according to Mohr. Hence, the two parts of Oudemann’s density tables are inconsistent. The fact that the part of the table referring to 28 per cent. acid gives results agreeing fairly with the titration method, while that referring to 3 per cent. gives discordant results, shews pretty clearly the direction of the error. It is, of course, impossible to ascertain the cause of the discrepancy with certainty, but it is worth * Another sample of acid gave 27.27 aa the mean of three titrations in which the end of the reaction was indicated by litmus, and 27-25 as the mean of three in which Cu SO4 was employed.Hence, if there be an error of deficiency introduced by the w e of litmus, the 8ame objection applies to sulphate of copper, and probably other indicators.THE ANALYST.269 notice that if one measure of acid at 29 per cent. were diluted with nine measures of water, the dilute acid would really contain 3.004 per cent., instead of 2-9 per cent. as might often be assumed. These considerations have no relation to, and are in no way affected by, the well- known abnormal density of strong acetic acid.On the whole, we considered that the real amount of acetic acid in the sample was represented most accurately by the result of the titrations, and therefore in the following experiments the acid used is regarded as containing 28-52 per cent. of real acetic acid (Q H, 0,). I n all cases in which acetic acid was to be determined, a weighed (not a measured) quantity of the sample was employed.A. The first process tried was the determination of free acetic acid in presence of free sulphuric acid, by adding excess of' carbonate of barium, boiling well, filtering, and precipitating the barium from the filtrate by dilute sulphuric acid. The amount of Ba SO4 found represents an equivalent amount of acetate of barium formed, and the weight multiplied by *515 gives the acetic acid.HX Taken. Ea' Found, Expt. 1. ,447 grms. . . . . . . . . . . . . . . . . . . *456 grm. Expt. 2. 9397 ,, . . . . . . . . . . . . . . . . . . '406 ,, These experiments shewed, as was to be expected, that free acetic could be readily determined in presence of sulphuric acid. Unfortunately the method is useless in presence of sulphates and many other salts.B. In this case, the above method was modified so a8 to render it applicable to the analysis of acetates, and to free acetic acid in presence of sulphates. A known weight of the sample of acid was neutralized with standard soda, and standard sulphuric acid added in twice the quantity necessary for the conversion of the soda into NaHSO,. The liquid was then distilled nearly to dryness, water added, and the distillation repeated.The distillate was treated with Ba CO,, as in process A. Expt. 3. *548 grms. . . . . . . . . . . . . . . . . . . $61 grms. Expt. 4. 1,126 ,, . . . . . . . . . . . . . . . . . . 1.118 ,, H r Taken. HX Found. If sulphate of silver were added before distillation, the method would be equally accurate in presence of hydrochloric acid and chlorides.Phospharic acid has some- times been employed instead of sulphuric acid, and would, doubtless, be preferable in presence of sugar, &c. C. The next process investigated was that for the determination of free sul- phuric acid in vinegar, by precipitating the sulphates with alcohol. An artificial vinegar was made by adding to dilute acetic acid some caramel, calcium sulphate, potassium sulphate, and a known quantity of standard sulphuric acid, in such quantity that the liquid contained about 6 per ceut.of Ha, and 1 per cent. of H,SO,. 50 C.C. of the (' vinegar" were evaporated to 10 c.c., and treated with 50 C.C. of rectified washed with alcohol, and H,SO, precipitated with results : - &SO4 Added. Expt, 5. -636 grms. Expt. 6.-536 ,, spirit. After standing, the precipitate was filtered off, the filtrate diluted, the alcohol boiled off, and the free barium chloride. By this process we obtained these HzS04 Found. . . . . . . . . . . . . . . . . . . ,655 prms. . . . . . . . . . . . . . . . . . . e53a ,,270 THE ANALYST. Hence, it appears that, provided sufficient alcohol be added, a very exact separation of free from combined sulphuric acid, can be effected.A sample of commercial vinegar treated by the above method, shewed no free sulphuric acid, whilst it contained sulphates corresponding to no less than -159 per cent. (= 111 grp. per gallon) of H2S04. The same vinegar contained 63 grs. per gallon of chlorine. These results, given by vinegar of unknown origin, present a curious resemblance to those obtained by Lethcby from the article manufactured by Messrs.Hill and Evans, of Worcester. No KHSO, is formed. These were :- H2SO4 (as sulphates.) GI. Two experiments were next tried by adding to 50 C.C. of the above sample of commercial vinegar a known amount of standard sulphuric acid, and proceeding as before. I n the Vinegar. ... 11 1 grs. per gallon ....... 50 grs. per gallon. In the Water. ... 119 ,, ,, ... ... 48 ,, 9 7 Expt. 7. For 0.268 grm., H2S04 added, *188 was obtained. Expt. 8. ,, 0.268 ,, 9 , -190 ,, Hence, a considerable and nearly equal loss occurred in both cases, the mean being -079 grms. Of this, -062 is accounted for by the reaction of the sulphuric acid added, upon the chlorides present in the vinegar. The remaining *017 grms.pro- bably reacted on acetates or phosphates. The conclusion to be drawn from the experiments is that the alcohol method will shew the true amount of sulphuric acid existing free in the vinegar, but that will probably be less than the amount added. By adding sulphate of silver before concentrating, the result would indicate the total free mineral acid in terms of sulphuric acid.The process thus modified would be applicable to the determination of free hydrochloric acid. This ie based on the fact that, while acetates are conv'erted into carbonates on ignition, and hence yield an alkaline ash, sulphates and chlorides suffer no similar change. Mr. Hehner further proceeds on the assumption that the presence of acetates in the vinegar is incompatible with that of free hydrochloric or sulphuric acid in the original vinegar, and hence any alkaline reaction of the ash of the vinegar, by proving the presence of an acetate, negatires the possiblity of the presence of a free mineral acid.It was to be expected that the evaporation to dryness and subsequent ignition of a solution containing acetic acid, free sulphuric acid, and a sulphate, would D.We next examined the very convenient process of Mr. 0. Hehner." produce a non-alkaline ash, but the same result seemed by no means certain if a chloride were evaporated with a comparatively large quantity of acetic acid. In this case it was thought probable that the effect of mass would be observed, and that the large proportion of acetic acid would effect more or less decomposition of the chloride, with volatilization of hydrochloric acid and formation of an acetate.To obtain information on this point we made the following experiments :-A solution of common salt, in which the chlorine had been determined by nitrate of silver, was evaporated with a large excess of acetic acid. In some cases the operation was concluded at dryness, in others the solid residue was ignited.In some cases burnt sugar was added. The chlorine in the residue was determined by nitrate of silver. * see ANALYST vol. 1, p. 135.THE ANALYST. 271 Expt. 9. ‘0504 grms. of Na GI, evaporated to dryness with a large excess of acetic acid, gave a Expt. 10. The same experiment repeated, gave absolutely the same amount of Na C1 before and Expt. 11.The same operation, with subsequent careful ignition of the residue, showed a loss of Expt. 12. The same operation as in the last experiment, but with caramel added, shewed a lose Expt. 13. Conditions the same. Loss = 00014 grms. Na C1. It appears from these experiments that the decomposition of the salt is practically nil, by mere evaporation, but that on ignition there is a loss of 2 to 3 per cent.of the total chlorine present. As there can be no free acetic acid present to account for this loss, it is probably due to unavoidable volatilization of the chloride, rather than to its decomposition. On the other hand, when chloride of sodium solution was evaporated with tartaric acid, and the residue ignited, 36 out of 54 milligrsmmes of salt were decomposed, or Q of the total quantity taken.We also evaporated common salt solution with excess of cream- of-tartar. This result was to be expected, as even acid szt&hate of potassium does not react on common salt at moderate temperatures. On igniting the evaporated mixture of sodium chloride and cream of tartar, slight decomposition occurred, in one experimeiit 3 milligrammes, and in another 4 milligrammea of common salt being decomposed, out of the 50 milligrammes added.When a solution containing 50 milligrammes of sodium chloride was evaporated to dryness (but not ignited) with excess of citric acid, the residue gave a weight of AgCl corresponding to only -036’7 grms. of Na C1, shewing a loss of 24.6 per cent. of the common sslt taken. These experiments have a bearing on the method of Mr.W. C. Young for the determination of mineral acids in vinegar.* His process consists in adding excess of Ba CI, t o a known measure of the vinegar. I n a portion of this liquid the chlorine is determined. The rest is evaporated, ignited, and the chlorine determined in the ash. The difference represents the free mineral acid expressed in terms of C1.Acid tartrate of potassium is, of course, a constituent of wine-vinegar, and its presence would cause the determination of mineral acid by the above method to be somewhat too low. residue containing -0496 grms. of Na GI. Loss *0006 grms. after evaporation with acetic acid. *0012 grms. Na C1. of -0014 grms. Na GI. By mere evaporation to dryness no decomposition of the chloride ensued.The presence of free tartaric acid would quite invalidate the results. As citric acid decomposes a chloride on evaporation of a solution containing it, it is clear that Mr. Young is in error in stating that his process is of course applicable to lime- juice or lemon-juice.” Mr. Eehner’s method of determining free mineral acids in vinegar, is dependent on the alkalinity of the ash of the vinegar as compared with the amount of alkali added to the original liquid.‘( I f we add to a measured quantity of the vinegar a known and exactly measured volume of decinormal soda solution, somewhat more than would be necessary to neutralise the total amount of free mineral acid present, evaporate and incinerate, the alkalinity of the ash gives the measure of the quantity of the free aulphuric or hydrochloric acid.” The author’s test experiments are very satisfactory.See ANALYST vol. 3, p. 163,272 THE ANALYST. It is evident that the amount of alkali added must be sufficient to combine with the fixed organic acids present, in addition t o the mineral acids. I f the amount of alkali employed be insufficient, it is necessary to recommence the experiment.For this reason, and from the desire to determine the free acetic and the mineral acid in the same portion of vinegar, we have made some experiments in which enough normal soda was employed to neutralize the whole of the acid, the subsequent manipulation being unchanged. As the amount of alkali used was about 20 times as great as that employed by Mr. Hehner, the tendency t o a slight error in the titration was greatly increased, but the following results show that this modification of the process is capable of all desirable accuracy.An artificial vinegar was made by mixing acetic acid, potassium sulphate, caramel, and a known amount of standard sulphuric acid. A slight excess of standard soda was added, the liquid evaporated, the residue ignited, the ash dissolved in excess of standard acid, and titrated back with alkali.Expt. 18. For -2735 H2S04 taken, . . . . . . . . . -2695 was found. Expt. 19. ,, -2735 ,, ,, . . . . . . . . . -2755 ,, The commercial vinegar already mentioned as containing sulphates equivalent to 11 1 grains per gallon of sulphuric acid, in addition to 63 grains of chlorine, when examined by this process, Ehowed a small minzcs quantity of free sulphuric acids-a result fully oonfirming the alcohol determination.Another experiment was made by adding a definite amount of standard sulphuric acid to 50 c. c. of the above commercial vinegar, and then proceeding a5 before, when we obtained :- H2S04 Taken. H2SOa Found. Expt. 21. . . . . . . . . . -1220 grms. . .. . . . . . -1157 grms. As, however, a small error in the amount of alkali and acid used causes a sensible difference in the result, it is preferable to add (as recommended by Mr. Hehner) only a fraction of the total alkali which would be required for complete neutralization. Under these circum- stances, decinormal solutions can be conveniently employed, and hence greater accuracy in the results obtained. It will be observed that in Expt.21 nearly the full amount of sulphuric acid added is accounted for. Of course the result is really a determination of the free mineral acids (actually existing in the Tinegar) expressed in terms of sulphuric acid, for theoretical considerations and the results of experiments 7 and 8 (made on the same vinegar by the alcohol process) show that a considerable proportion of the free acid was hydrochloric acid. In short, Hehner’s process determines the total amount of free mineral acid, while the alcohol process,-with the use of sulphate of eilver if necessary-enables the relative proportions of the free mineral acids to be ascertained. Hehner’s procesa is in our experience, decidedly the most convenient and accurate in general use, and furnishes a valuable solution of a somewhat difEcdlt problem. R. Warington has employed the Bame plan for the determination of free sulphuric acid in citric acid liquors, and the same principle has been frequently made use of. Hence the process gives fairly accurate results in actual practice. * A minus result is very common with vinegars containing no free mineral acid, and, when beyond the limita of experimental error, is clearly due to the presence of acetate8 or other organic adts (e.9, malatee, lactates, tartrates.)