Constitution of Ketolactonic Acid.

It has been seen that this acid is formed from the compound

COOC,H.

One of the (O.C2H,) groups and an additional H-atom must then be removed from the first compound to form the alcohol. As regards the hydrogen-atom, the only difference between the two ethers is that an H-atom in the one is replaced by CH, in the other; but in the latter case no crystalline acid is formed. It is therefore at least probable that this hydrogen-atom is the one removed. It is also almost certain that it is the (OC2H¿) group farthest from this H-atom which is split off, because if so an unsaturated lactone group would be formed thus

and the acid appears in its behaviour with the carbonate and hydrate of barium to contain such a group.

Moreover the splitting off of alcohol in this way would resemble the formation of an unsaturated anhydride noticed by Thorne on distillation of a-ethyl-3-acetopropionic acid, which anhydride (or at least its next homologue) resembles the lactones markedly in its solubility in water.

Ketolactonic acid would then be obtained from its ether by simple saponification.

The barium salts, or for the sake of clearness the acids derived from ketolactonic ether, would then be―

(1.) Ketolactonic acid yielding a barium salt when treated with barium carbonate

(2.) Probably B-ethylacetosuccinic acid, obtained by the action of barium hydrate in the cold, that is by simple addition of H2O—

(3.) Probably a-ethyl-B-acetopropionic acid by the action of barium hydrate at 100°, that is by removal of CO, from the last acid

Since ketolactonic acid is decomposed by boiling with barium hydrate, it is easy to understand why Thorne did not obtain it, since he saponified the ether by boiling with caustic potash, so that the acid would be decomposed as soon as formed.

Before closing this paper I wish to return my sincere thanks to Professor Fittig for his ever ready advice and assistance in carrying out this investigation.

XXV.-On the Constitution of Molecular Compounds. The Molecular
Weight of Basic Ferric Sulphate.

By SPENCER UMFREVILLE PICKERING, B.A. Oxon, Lecturer in
Chemistry at Bedford College.

In representing basic compounds, two methods of notation are commonly adopted: the one consists in representing the compound as a combination of molecules of the metallic oxide with molecules of the non-metallic oxide, as, for instance, 4CuO,SO; the other, in representing it as a combination of molecules of the metallic oxide with the normal sulphate; thus, CuSO4,3CuO. In using this latter form we hint at an analogy between a basic salt and a hydrated salt, the metallic oxide in the former occupying a position similar to that which the water does in the latter.

Taking, now, the only known basic ferric sulphate, we may represent it either as 2Fe2O3, SO3, or as Fe2(SO4)3,5 Fe2O3; the former indicating its molecular weight to be 400, the latter indicating it to be 1200; if, therefore, a determination of the weight of its molecule showed it to be 400, it would render the latter method of representation impossible.

The only means of determining the molecular weight of a stable

solid, such as basic ferric sulphate, which promised any success, was in ascertaining, if possible, the unit of water removable from a hydrated specimen of it.

In 1879 the author performed a few experiments on basic ferric sulphate, which indicated that definite hydrates were formed, not only by exposing this substance to dry air at various temperatures, but also by exposing it to air saturated with moisture; the results of these preliminary experiments were communicated to the Ashmolean Society of Oxford, February, 1880. Somewhat similar experiments on certain metallic oxides were performed by C. F. Cross (Chem. Soc. J., Trans., 1879, 796), who showed that definite hydrates of oxides were obtainable by the action of moist air on them, and also that the hydrates thus obtained in any particular case depended on the temperature to which the oxide had previously been exposed. It will be seen below that similar results are obtained with basic salts.

Two samples of basic ferric sulphate were prepared by the action of a defecit of sodic carbonate on the normal sulphate (see this Journal, Trans., 1880, 807), and dried by exposure to air, without the application of heat. The first sample was used in experiments A and B, the second in experiments C and D.

A glass boat, containing a portion of the basic salt, was placed in a tube passing through a water-bath, and through this tube a current of thoroughly dried air was drawn by means of a water-pump.

In no case was the sulphate found to contain the least trace of carbonate; nor was it found to give off any traces of sulphuric acid (tested in experiment C, which lasted for 44 days).

The exposure to moist air was effected in the earlier experiments by passing a very slow current of moist air over the sulphate, and, in the latter experiments, by leaving the sulphate under a bell-jar in presence of a dish containing water, the whole being placed in a cellar where the temperature was constant. The temperatures mentioned below are the highest reached during the heating in any particular

case.

The letters given in the first column of the accompanying table refer to the experiments, and the numbers, to the order of the various operations in each. In experiment A the weight of anhydrous substance taken (as determined by subsequent analysis) was 0.4264 gram, and H2O would therefore correspond to 0.0064 gram, taking the triple formula: in B 2.0806 grams taken, H2O = 0·312 gram; in C 0·8163 gram taken, H2O = 0·01224 gram; and in D 0.81604 gram taken, H2O = 0·012237 gram. In the last two experiments the current of air was much slower than in the first two.

That the constant weights attained by the basic salt correspond to definite bydrates there can be little doubt. Altogether 28 constant

A(7) Fe2(SO4)3,5Fe03,8H2O In dry air at 100° C.

gained
weight.

12 hours

1341 34 1341 34

C (2)

12 days

8 days 1341 44

A(9)

Fe(SO4)3,5е03,11H2O

at 17° C. after being heated at 100° C. 26 hours

1394.66

1395 22 -0.56

A (10)

B (12)

93° C.

7 days

1 day

1396.66

+ 0.84 +1.44

A(3) Fe2(SO4)3,5 Fe2O2,12HO

at

72° C.

1412.90

1413 18-0.28

A (5)

at 17° C. after being heated at 100° C. 22 hours

B (10) Fe2(SO4)3,5Fe03,13H,O

B(9) Fe2(SO4)3.5 Fe2O3,14H2O

A1 Fe2(SO4)3,5Fe0.16H2O

18° C. after being heated at 85° C. 70 hours

C (1)

at

35° C.

4 days 1485.34

+0.32

weighings were obtained, and all of these correspond to the theoretical numbers within the limits of probable error, the greatest difference between the observed and calculated weights being only + 0.16 per cent., and the average differences being + 0·04 and -0.03 per cent. of the total weights.

Now of the 14 different hydrates thus obtained, four only can be represented by the simpler formula 2Fe2O3, SO3, H2O, while the remaining 10 necessitate the adoption of the triple formula

Fe2(SO4)3,5 Fe2O3,уH2O.

Even had there been obtained but one hydrate of this latter class, it would, in my opinion, have shown that the empirical molecular weight must be trebled; but the existence of so large a number of them must put the question beyond all doubt.

These experiments prove with almost equal certainty that no more complex formula than the triple one is required: for between the extreme hydrates actually obtained, 33 different hydrates are possible if the molecular weight be 1200 (corresponding to the triple forraula), and if the molecular weight were double, namely, 2400, 66 hydrates would be possible. Now of these possible hydrates 14 have been obtained, all of which correspond to the molecular weight 1200, and not to 2400. The odds, therefore, in favour of the true weight being 1200 will be found to be about 95,000: 1; and, making allowance for the fact that many of these hydrates were obtained several times, these odds will be increased to as much as 14,400,000,000,000: 1; in other words, it is practically a certainty that the molecular weight of basic ferric sulphate is 1200, and that it may therefore be represented by the formula Fe2(SO4)3,5 Fe2O3, H2O, but not by any simpler or more complex formula.

XXVI.-The Phenates of Amido-bases.

By R. S. DALE, B.A., and C. SCHORLEMMER, F.R.S.

OUR hope to get hold of the intermediate compounds supposed to be capable of existing between aurin and pararosaniline, has not yet been fulfilled (this Journal, 1879 [i], 148). A long series of experiments has convinced us that, on heating aurin with aqueous or alcoholic ammonia, the action proceeds at once to the formation of pararosaniline, which can easily be detected, even if only a trace has been formed, by dissolving the product in a little ammonia, adding hot water, and dipping some silk into the liquid. It will soon be dyed magenta-red, as aurin does not dye in an alkaline solution. "Red