REVIEW OF AMERICAN CHEMICAL RESEARCH. VOL.
M. T. Bogert, E. M. Chamot, B. S. Cushman, Benton Dales. L. M. Dennis, A. H. Gill,
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WILLIAMA. NOYES,Editor. REVIEWERS : H. M. Goodwin. W. F. Hillebrand, L. P. Kinnicutt, H. W. Lawson, G. N. Lewis, H. N. hlccoy,
A. A. Noyes, J. W. Richards, S. P. Sadtler, J. 0. Schlotterbeck, W. H. Seaman, F. P. Underhill.
GENERAL AND PHYSICAL CHEnlSTRY. On the Electrical Conductivity of Solutions in Amyl Amine. BY LOUISKAHLEXBERG AND OTTOE. RUHOFF. 1 . Phys. Chem., 7, 254-258.-Amyl amine, a substance of small dielectric constant (4. 5 ) , and very small electrical conductivity, dissolves readily a number of salts. The solutes may be divided into two classes. One represented by copper oleate, does not increase t h e conductivity of the solvent. The other produces solutions which are comparatively good conductors. Of thisclass three were studied, silver nitrate, cadmium iodide and ferric chloride. T h e conductivity was found to be small compared with aqueous solutions b u t much larger than that of solutions in other solvents of similarly low dielectric constant, such as chloroform and ether. Each of the first two salts gave a maximum of molecular conductivity with dilution. T h e same result might have been obtained with ferric chloride, if concentrated solutions could have been tried readily, fLr in this case as in t h e others t h e molecular conductivity of dilute solutions decreased very rapidly with dilution. For example, the value for cadmium iodide is 2 0 0 times as great in normal as in fifth-normal solution. This very important phenomenon, which has been observed in other cases, especially in non-aqueous solutions, the authors regard a s a strong argument against t h e Arrhenius dissociation theory and they hint a t a theory of their own which, however, they do not develop. T h e phenomena attending the solution of the salts point to a union of amyl amine with the solids to produce new compounds analogous to hydrated salts. Comparing amyl amine with methyl amine and ammonia, the authors conclude that in this homologous series t h e dissociating power decreases with increasing molecular weight. G. N. LEWIS.
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R e v i e w of A m e r i c a n Chemical Research.
The Reliability of the Dissociation Constant as a Means of Determining the Identity and Purity of Organic Compounds. BY HEYWARD SCUDDER. J . Phys. L-hem., 7, 26g-2gg.-Reviening the literature concerning the dissociatioii constants of organic acids, and citing a large number of data found by different i l l restigators, the author shows tliat in many cases the value of the dissociation constant is a moderately reliable aid to the identification of acids but that tlie agreement lietween the values of the constant at different dilutions is not a satisfactory guide as to the purity of a compound, on account of the possibility of large experimental error and on account of the fact tliat many pure organic acids show a marked increase in the constant on dilution. C;. N. LEWIS. Note on the Identification of Basic Salts. BY IV. LASH MILLERA N D FRAXK B. KEXRICK. J. Phys. Chem., 7, 2 jg-26s. --The authors show how the phase rule map be used to give information concerning the constitution of such precipitates as basic salts. If, for example, a system of three components is studied, such as bismuth nitrate, nitric acid and water, there will not be, in general, more than two solid phases in equilibrium with a solution a t any given temperature and pressure. If, therefore, in different experiments, the solutions are identical in composition, while the composition of the precipitate varies, the precipitate is a mixture of two phases. If the solutions differ in coniposition while the precipitates do not, the precipitate is a siiigle cheniical compound. If both solutions and precipitates vary, the precipitate is a single phase of variable composition, a solid solution. I t is further shown how similar considerations niay be applied to more complicated systems. Since these principles hold only for systems in equilibrium, the practical obstacle to their application lies in the very slow attainment of equilibrium in the majority of cases in which such methods would be needful. G. S . LEVIS. The Compensation Method of Determining the Rate of Oxidation of Hydrogen Iodide. BY JAMES M. BELL. J . Phys. Cheur., 7 , 61-83.-If a reaction produces a substance A , such that it may be titrated against a substance B, the speed of the reaction m a y be determined by adding in advance a known quantity of B, together with an indicator, and then finding the time required to produce enough X to neutralize it. This is called by the author the compensation method. H e discusses a number of researches on the rate of oxidation of liydro,yen iodide in which this method has been used with tliiosulphate a s compeiisator, and points out that the method is applicable only in case I that the thiosulphate reacts only Lvith tlie iodine liberated by the action of the oxidizing agent on tlie hydriodic acid, and ( 2 ) that its presence in no way influences the course of that action.” Experiments are described which show that when H,O, is tlie oxidizer I ‘ (
Geiteral and Physical Chemistry.
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a s in the researches of Harcourt and Esson, these conditions are probably fulfilled, but in other cases the thiosulphate is itself oxidized a t a considerable rate. Equations are developed for the rate of reaction, assuming that both iodide and thiosulphate are oxidized siniultaneously and independently, but these prove inapplicable to the cases studied in which chloric, bromic and chromic acids were used to oxidize. I n all these cases the thiosulphate is oxidized but at a rate which is retarded by the presence of iodide and dependent on the concentration of the latter. This retardation is sometimes sufficient to make the compensation method approximately reliable. “ I n addition, the paper contains a number of measurements on the rate of oxidation of sodium thiosulphate by chloric and chromic acids, with particular reference to the effect produced by adding catalytic agents, and by changing the concentrations of the thiosiilphate, the acid, and the oxidizing agent. T h e results do not suggest any simple ‘mechanism’ for the reaction.” G. N. LEWIS.
The Rate of Reaction in Solutions Containing Potassium Iodide, Potassium Chlorate and Hydrochloric Acid. BY W. C. BRAT. /. Phys, Chem., 7 , gz-r17.-This paper, apparently written in some haste. contains a considerable number of experimental data from which t h e following conclusions are drawn. T h e reaction in question is greatly catalyzed by ferrous sulphate. If, in the iodometric determination of chlorates, ferrous sulphate be added, the use of a digesting bottle may be dispensed with. T h e rate of the original reaction is proportional to the concentration of the chlorate, and to the square of the concentration of the hydrogen ion (within limits). I t is accelerated by the chlorine ion by an amount proportional to the concentration of the latter. Changing the concentration of the iodide produces strange results; as the concentration increases the rate at first diminishes and then increases almost linearly. T h e speed of reaction is doubled by 5.6” rise of temperature. I n explanation of his experimental results the author suggests two reactiogs occurring siniultaneously. H e offers, however, no argiinients for the reactions which he chooses. G. N. LEWIS.
The Rate of Oxidation of Potassium Iodide by Chromic Acid. BY RALPH E. DELURY. 1. Phys. Chem., 7, 239-253.--A study of the rate of the reaction between bichromate, iodide and sulphuric acid at j o ” showed that this rate is proportional to the concentration of the bichromate and approximately to the square of the concentration of the acid. With changing amounts of iodide, the rate changes more than it would if proportional to the concentration of the iodide and less than if proportional to the square of that concentration. T h e author suggests that this may indicate the simultaneous occurrence of two reactions, but he suspends judgment until further experiments are made. A series
502
Review of .47nerica7~ Chemical Research.
of experiments a t oo gave results similar to those obtained at 30'. The temperature has an unusually small influence upon the reaction velocity, the quotient for 10' being only 1.4. Experiments were made on the rate of the reaction in the presence of a nuiiiber of other salts, of which ferric sulphate alone possessed pronounced catalyzing power. G . X'. LEWIS.
Note on the Variation of the Specific Heat of ilercury with Temperature. Experiments by the Continuous-Flow Method of Calorimetry. BY H. T. BARTES A N D H. L.COOSE. P ~ J v . Rezj., 16, 65-~1.--With the precise method of continuous flow ithis Rev., 23, 189) the authors have carefully executed a fen. experiments on the specific heat of mercury and especially its change with the temperature. For Js, the specific heat expressed in joules per degree, or niayers, the following equation is obtained : Js, ==-0.140154 - 4.462 >< IO-'.? j- 0 . O i j : IO-'^'. Subject to a very slight correction, not yet applied, these figures refer to the hydrogen thermometer. T h e decrease i n specific heat with the temperature is in close agreement with the earlier work of Winkelinann and of Milthaler. G . S . LEWIS. ~ v :
Some Optical Properties of Iodine, 111. BY WM. IY.COBphys. Rev., 17, 51-5g.-continuing his investigation of the absorption spectrum of iodine under varying conditions (this Rev. , 2 5 , 268) the author was enabled, by means of cells made of pure, thin quartz plates, to study the ultra-red spectrum as far as wave-length 16p, at which point rock salt begins to absorb heavily. Solutions in ethyl alcohol, i n acetic acid and in both chloroform and carbon disulphide were studied as types, respectively, of the brown, of the red-brown, and the violet solutions. I n the first two cases the solutions absorb more than the pure solvents, and absorption bands due to the iodine appear. In the two last solutions, the transmission curves coincide throughout with those of the solvents, showing that iodine in violet solutions is transparent through the whole ultra-red. G. S . LEWIS. LEKTz.
MINERALOGICAL AND GEOLOGICAL CHEnISTRY. The Origin of Coral Reefs as Shown by the Maldives. BY J. STANLEY GARDIXER.ATJZ. J . Sci., 16. 203-213 ; figure. UT. F. HILLEBKAKD. The Occurrence of the Texas flercury Minerals. B r BENJAF. HILL. Am. J . S i . , 16, 2 j ~ - ~ j ~ . - - T h e s minerals e are found in both the Upper and Lower Cretaceous rocks of Terlingua, Brewster County, in proximity to volcanics. T h e deposits of the Lower Cretaceous, n.hich are at present the most important, occur in decomposed and brecciated zones in the Edwards and MIS
