January 24 of last season. The number of bad days on which little or no sun was recorded lowers the average for the period under examination. It amounts to 7.05 mgms. per hour, whilst in January 1893, taking only the bright days, the quantity was 9.34. I understand that the ordinary meteorological record was kept as usual, and that the data are to be found in the Alpine Post' for the days on which these experiments were carried out. I append, however, a brief note as to the atmospheric conditions in a separate column. It has been remarked in the Engadine that the amount of sunlight during the winter months has been diminishing, but so far as I have been able to ascertain, there is no accurate information upon which a statement of this character can be based. The average value per hour obtained from my experiments last year for the month of January, including the bad days, amounted to 8:41 mgms. of iodine per 100 c.c., per hour, whilst, as already mentioned, the mean from the above experiments, extending over a longer period, is 7·05. This shows that the average amount of light was less during the 1893-4 season than during the previous one, but I do not think that from the records of two seasons any deductions can legitimately be made. I hope to make arrangements for a continuance of this work during the ensuing winter, and should be glad to hear of similar work being carried on at some of the other centres in the High Alps. 7. Interim Report of the Committee on the Formation of Haloids from Pure Materials. 8. Interim Report on the Bibliography of Solution.-See Reports, p. 246. SATURDAY, AUGUST 11. The Section did not meet. MONDAY, AUGUST 13. 1. A joint meeting with Section A was held, at which Lord RAYLEIGH, Sec.R.S., and Professor W. RAMSAY, F.R.S., gave a preliminary account of a New Gaseous Constituent of Air. The following Papers were read :— 2. On Schuller's Yellow Modification of Arsenic.' 3. On the Electrolysis of Glass. By Professor W. C. ROBERTS-AUSTEN, F.R.S. 4. On the Relations between the Viscosity of Liquids and their Chemical Nature. By Dr. T. E. THORPE, F.R.S., and J. W. RODGER. During the past four years the authors have been making observations on the viscosity of liquids with the view of establishing relationships between this property on the one hand and molecular weight and molecular structure on the other. They have employed the tube-method of measuring the viscosity coefficient, and in the case of each liquid the observations extend over the temperature range between 0° and the ordinary boiling-point. More than eighty liquids have now been examined. For liquids which do not appear to contain molecular aggregates, at any temperature, the following conclusions may be drawn :- 1. In homologous series the viscosity coefficient is greater the greater the molecular weight. 2. The coefficient of a normal compound is greater than that of the isomeric iso-compound. 3. The coefficient of an allyl compound is intermediate to those of the corresponding normal propyl and iso-propyl compounds, and, in general, constitution exerts a regular effect on the viscosity coefficient. Liquids which appear to contain complex molecules in certain cases do not obey these rules. Formic and acetic acids are exceptions to Rule 1. The alcohols do not conform to Rules 2 and 3. In general, the effect of temperature upon viscosity is much greater for complex than for simple liquids. In both classes of liquids the behaviour of the initial members of several homologous series does not accord with that of higher homologues. In attempting to quantitatively connect viscosity with chemical nature, the authors have used two magnitudes the molecular viscosity and the molecular viscosity-work-which may be derived from the viscosity coefficient. If n be the viscosity coefficient and v be the molecular volume, the molecular viscosity is nv, or the product of 7 and the molecular area. The molecular viscosity-work is nv, the product of 7 and the molecular volume. The values of these magnitudes have been examined at three different series of temperatures of comparison-viz., the ordinary boiling-points, the corresponding temperatures of 0°-6, and temperatures of equal slope or points on the viscosity curves at which temperature is exercising the same effect upon the viscosity of each liquid. On ascribing definite partial values to the atoms and the different modes of atom-linkage-the iso-grouping, double linkage, the ring-grouping, &c.—it has been found possible to calculate the viscosity magnitudes of the great majority of the simple liquids. The results obtained at equal slope are by far the most precise, but even here the alcohols, water, and, to a less extent, the acids are anomalous, doubtless on account of the influence of molecular complexity. A strong point in favour of the new method of using equal slope as a condition of comparison is that the stoichiometric relationships obtained at any one value of the slope appear to be general, and thus to be independent of the particular value of the slope at which the comparisons are made. This conclusion applies even to complex liquids like water and the acids, but not to the alcohols, which of all the liquids examined exhibit, at all of the systems of comparison, the most exceptional behaviour. ' Published in the Chemical News, lxx., p. 139, Sept. 21, 1894. 2 Published in full in the Phil. Trans., 1894. 5. Some Experiments on the Rate of Progress of Chemical Change. By DR. J. H. GLADSTONE, F.R.S. In the last February number of the Philosophical Magazine' Mr. Veley pointed out four stages of a chemical reaction: First, the commencement; second, the period of inertness or reluctance, followed by acceleration; third, of constancy; fourth, of diminution of velocity.' This reminded the author of various old experiments which had never been published, and he returned to the subject with a view of seeing whether this period of inertness followed by acceleration occurred in such simple cases as that of reciprocal decomposition of salts, and whether, where it did occur, it was capable of any explanation on known principles. The reciprocal decomposition of potassio-platinum-chloride and potassium-iodide slowly produces the iodine salt which makes itself manifest by its deep red colour. In examining this in various ways the action always appeared most rapid at first and gradually slackened till a balance of the salts in solution was obtained. In such cases, on the contrary, as that of the formation of bitartrate of potassium or calcium, where a larger amount of the products is formed than can be kept in solution, it is some time before crystals make their appearance, and then they come with a rush, gradually diminishing to the end of the reaction. This seems to be due partly to the phenomenon of super-saturation and partly to the necessity of rapid redistribution of the acids and bases when the bitartrate produced is thrown out of the field of action and reaction. In cases where almost insoluble salts are formed, such as strontium sulphate, the liquid becomes milky almost at once, a constant redistribution being necessitated by the separation of the insoluble salt, and the curves representing the course of the action closely resembled those of the platinum salt given above. There are unquestionably many cases in which there is very little appearance of action at first, but afterwards it comes on rapidly, and then, of course, diminishes. The formation of zinc-methyl was quoted, but a more interesting instance was the reaction between cuprous oxide and silver nitrate in rather weak solution. At first little or nothing is seen; after a while long filaments of metallic silver shoot forth, the reaction becoming very rapid, until the silver solution is very much weakened, when, of course, it proceeds more slowly. In explanation of this we may conceive of some possible induction,' or charging up of the metallic oxide, or the influence of the local rise of temperature, or the greater scope for voltaic action between the growing silver and the copper compound. The author, however, did not insist on any particular explanation, but gave the facts as a contribution to the general subject. 6. The Determining of the Freezing-point of Water, van't Hoff's Constant, Arrhenius' Law of Dissociation, Ostwald's Law of Dilution. By Dr. MEJER WILDERMANN. I have already given an account, in the Physical Section, of the method devised, in concert with the late P. B. Lewis, for accurately determining the freezing-point of aqueous solutions which freeze at temperatures just below 0° C. The depression of the freezing-point of a solution of any concentration is stated in degrees below the freezing-point of water. The freezing-point of water and of extremely dilute solutions is very difficult to determine with accuracy. Under ordinary circumstances a cap of ice forms round the bulb of the thermometer; if the formation of this cap is prevented, the freezing-point of water determined by my thermometer divided to 0.001 is higher by 0° 0015 to 0° 0017 than when the cap exists and a constant error in the determination of the freezing-point is not removed. The method of determining the freezing-point of very dilute solutions which was devised by my late friend P. B. Lewis, and my investigations of the freezingpoint of water, and of extremely dilute solutions, give us a means of submitting van't Hoff's constant, Arrhenius law of dissociation, and Ostwald's law of dilution to a more accurate verification. It is well known that it was van't Hoff who first drew attention to the fact that the equations representing the generalisations arrived at by Boyle, Gay Lussac, and Avogadro in the case of gases are equally applicable to dissolved substances if the osmotic pressure of the dissolved molecules be substituted for the pressure of the gas. While van't Hoff was able to establish a thermodynamic relation between the osmotic pressure of a dissolved substance and the molecular lowering of vapour pressure, the molecular lowering of the freezing-point of solutions furnishes a rational basis for the empirical generalisations of Raoult, and of Babo and Wüllner. In van't Hoff's thermodynamical argument the solutions are assumed to be very dilute, and Lence experimental verification is specially important for the case of such solutions. 0.02 T2 I have found that in the case of aqueous solutions of sugar and urea the agreement between the value calculated by van't Hoff by means of the equation (which must equal 1·89 if W = 79 kal., and is 1·87 if W = 80 kal.) and the observed value of the constant is excellent. Even in the case of alcohol the t = W values do not vary by more than 1 per cent. from 1-87-a difference which may be accounted for by the difficulty of determining exactly the percentage of alcohol in a solution from its density. That this is the case is shown by the fact that the value 1.84 or 1.85 is observed for all concentrations. It is also possible to calculate van't Hoff's constant without determining the freezing-point of water in the following way: Sugar, urea, and alcohol are not electrolytes, i.e., are in water only very slightly dissociated. We can, therefore, determine the relation between concentration and depression of freezing-point, starting from a solution of any convenient concentration, where an ice cap is not formed, instead of from pure water, and thus eliminate the influence of any error in the determination of the freezing-point of water. From these observations made with my thermometers divided to 0°.01 and 0·001 I have been able to establish van't Hoff's constant by a second independent method. Also, if the results obtained by Loomis with a thermometer reading to 001 are similarly treated, the van't Hoff's constant becomes evident in the case of sugar, less evident in the case of water and alcohol, though the variations are so great that the probable error is greater than he suspected, and the concentration of the solution was probably wrongly determined. We proceed to the generalisation of Arrhenius. Van't Hoff showed by four different methods that a law analogous to that of Avogadro was valid for solution of non-electrolytes like cane-sugar. It then became of importance to account for exceptional cases in which the depression of the freezing-point was abnormal, and in particular the cases of salts, acids, and bases in aqueous solutions. The explanation was given when Arrhenius showed that by two independent, quite different methods, the observation of the lowering of the freezing-point and of the electrical conductivity of a solution, the same value could be obtained for the factor i, which denotes the ratio of the pressure actually exerted by the substance to the pressure which the substance would exert if it consisted entirely of undissociated molecules. This law, which is of special importance owing to the light thrown by the dissociation-theory on various physical and chemical problems, must, like those of van't Hoff already mentioned, be more valid in very dilute solutions, and should at first be verified for them. For sugar, urea, alcohol, which are bad conductors of electricity, we have found normal depression and a constant 1.89 or 1.87. For KCl, SO H2, dichloracetic acid, trichloracetic acid, and nitrobenzoic acid, which are good conductors and show at the same time abnormal depression, I found that the degrees of the dissociation from the lowering of freezing-point and from the electrical conductivity are nearly the same. It is obviously desirable that Ostwald's dilution law, one of the laws of the action of masses, and a most important foundation for the theory of dissociation, should be verified by determinations of freezing-points, just as it has been verified by determinations of electrical conductivity; and for the reasons already stated the experimental verification is most important in the case of the most dilute solutions. The effect of experimental error in the calculation is here very considerable, and the freezing point methods hitherto in use have not been sufficiently delicate to verify the dilution law. The more accurate method already referred to has to a large |