A cryoscopic study in ethylenediamine

1478

NOTES

The decreasing solubility, in ethanol, with increasing cation size, from 251 g per 100 g for lithium iodide 14~ and 43.3 g per 100 g for sodium iodide I5~ to 0.331 g per 100 g for caesium iodide, all at 25°C, is in contrast to minimal solubilities often observed for alkali metal halides when water is the solvent. This type of pattern has been explained for aqueous solutions I7, s~ and the similar explanation in this ease is that there is possibly a decrease in solvation energy with increasing cation size in contrast to little change in lattice energy. The high solubility of lithium iodide is probably partly explained by polarization. Other factors such as the size of the solvent ethanol molecule preventing its entry between the ions in the crystal lattice, also probably contribute towards the observed trend. tt~ N. A. LANOE(Ed.), Handbook of Chemistry, 6th Edition. Handbook Publishers Inc. (19~.0. ~2~j. H. PERxY(Ed.), Chemical Engineering Handbook, 2nd Edition. McGraw Hill, New York (1941). t3~ C. D. HODO~N (Ed), Handbook of Chemistry and Physics, 41st Edition. Chemical Rubber Publishing Co. (1959). ~4~W. E. S. TURNERand C. C. BtSS~TT,J. Chem. Soc. 1904 (1913). 15) R. G. LARSONand H. HUNT, J. Phys. Chem. 43, 417 (1939). (6) p. T. WALDEN,Z. Phys. Chem. 55, 712 (1906). ~7~L. BREWER,J. Chem. Educ, 36, 446 (1959). Is) M. F. C. L~d~Dand W. H. LEE, Trans. Faraday Soc. 54, 34 (1958).

Department of Chemistry and Biology Welsh College of Advanced Technology Cathays Park, Cardiff

J. D. R. THOMAS

A cryoscopic study in ethylenediamlne (Received 22 February 1962; in revisedform 2 April 1962) IN CONNT.XION with equilibrium studies in ethylenediamine (EDA),O. 27 it became necessary to determine the state of aggregation of several acids and salts. EDA has a very convenient meltingpoint,,- 11 °CC3},and it was decided to use cryoscopy to obtain the necessary data. A literature search failed to reveal any previous cryoscopic study in this solvent. Preliminaryattemptsto obtain useable temperature-time cooling curves were unsuccessful. EDA supercools readily, and on freezing large crystals form which adhere to the cryoscopic cell wall. EDA continues to separate on this adherent coating, forming a hollow shell of frozen solvent rather than separating as individual crystals which can be suspended in the solution by agitation. The net result of this type of crystallization is a cooling curve which shows rapid temperature fluctuations, probably as a result of local temperature gradients in the solution. In addition, even brief exposure of the EDA solution to the atmosphere resulted in significant decrease in freezing point. The cell design shown in Fig. 1 proved to be a satisfactory solution to these problems. The temperature sensitive device was a Veco Type 32A83 thermistor in a conventional equal-arm Wheatstone bridge. Using a 1.000V bridge supply and a 1-00 mV, variable zero, potentiometric recorder, a sensitivity of 3 x 10-4°Cwas attainable. The gas-tight construction of the cryoscopie cell allowed indefinite storage of solutions without a detectable change in freezing point. All solutions were saturated with dry nitrogen by introducing a gas bubbler through joint C of Fig. 1. The gas bubbler was then removed and replaced by the assembly shown in joint C of Fig. 1. This assembly consisted of an outer TS 12/18 joint sealed to a short stem of glass tubing of an inside diameter to just pass another glass tube which had a short piece of platinum wire sealed into the bottom. A sleeve of polyethylene tubing allowed the platinum wire to be adjusted in height without exposing the cell contents to the atmosphere. A large paddle type stirrer entered the cell through a water-jacketed glass bearing and was rotated at ,,~900 rev/min. A cooling curve was obtained by placing the cell ~1~S. BRUCg~NSTE~and L. M. MUgZmPJEE,J. Phys. Chem. 64, 1601 (1960). (2~ L. M. M ~ . Ph.D. Thesis, University of Minnesota (1961). ~3~L. A o v ~ r r i and J. KLBn~rnmto,Non-aqueous Solvents, p.119. J.Wiley, New York (1953).

NOTES

1479

;5/so

FIG. 1. Freezing Point Cell A. T 20/40jointwith Veco Type 32A83 thermistorshowninposition; B. T 20/40joint with Labglass watercooled, precision ground sleeve bearing stirrer shown in position; C, t[ 12/18 joint. Inlet port for nitrogen bubbler (not shown) and platinum (0"010 wire) coldfinger (shown in position): D. liquid level. Cell contents 50 ml. in a water bath held at 9 + 0.02°C. Powdered solid carbon dioxide was dropped inside the tube with the platinum wire, and the wire submerged in the EDA solution as the anticipated freezing point was approached. Crystals of solvent formed on the platinum wire and were distributed through the solution by the vigorous stirring. These crystals greatly reduced supercooling and minimized the tendency of the solvent to solidify as an adherent layer on the cell walls. More than half the cooling curves obtained were satisfactory. On random occasions poor cooling curves were still obtained. In all cases each solution was reexamined until three satisfactory cooling curves were obtained. The average of the three determinations usually agreed to 4-0.0015°C. The thermistor was calibrated by evaluating the constant, B, in the equation l n R r l / R r 2 = B(1/T1 -- 1/7"2) from the measured resistances at the melting points of ice and of very pure acetic acid.t4) This calibration yielded a calculated sensitivity of 0.0954°C/mV of unbalance bridge potential at 10.6 ° compared to the experimental value of 0.0952°C/mV over the range 9.3-12°C~ The solvent handling techniques and the reagents used have been described elscwhere.(~, 2) The freezing point of pure ethylenediamine was determined to be 11.3 °C, as compared to the previously reported value of 11.0°CJ 3~ The molal freezing point depression constant, K f °, was determined using benzene and naphthalene as standards. Using benzene, six solutions in the concentration range 0.05-0.33 molal were studied, yielding K f ° = 2.43 +0.04°C/mole/kg. Three naphthalene solutions, 0-04-0-11 molal, yielded K f ° = 2.44+0.01°C/mole/kg. The mean value obtained giving equal weight to each individual measurement is 2.43 +0.02°C/mole/kg. The latent heat of fusion of EDA is calculated to be 64.5 +0.6 cal/g at 11.3°C, as compared to a calorimetrically determined value of 77 cal/g at 0°C. c4) K. HEss and H . H t m ~ , Ber. Dtsch. Ges. 70, 2205 (1937).

1480

NoTEs

Earlier potentiometric studies of silver cyanidetl) were interpreted in terms of the reaction: 2AgCN ~ " Ag + +AgCN2assuming the silver cyanide was nearly completely present in the form of the undissociated monomer. Cryoscopic studies of silver cyanide in the concentration range 0.018-0.16 molal verified that this compound exists quantitatively as undissociated AgCN. (By quantitatively, we mean the molality found cryoscopically was indistinguishable from the analytical concentration within the experimental error.) Solutions of silver, potassium and sodium iodide were also studied in the concentration range of ,-, 0-02-0.1 M and found to be present quantitatively as undissociated monomers. Silver bromide and chloride, and phenylacetic and acetic acids are not measurably dissociated in E D A in the concentration range of N0-01-0.09 Air. Silver nitrate is measurably dissociated in EDA. The results obtained are given in Table I, m, is the analytical concentration and Zm the molality calculated from the freezing

TABLE I.--CRYOSCOPIC DATA FOR AgNO3 IN E D A mj 0.0158 0.0359 0"0561 0"0809 0.1129

Xm

KAsNO~

(KAgNO~) M, G

0-0230 (6.0 × 10 -3) (1.2 x 10-4) 0.0441 2-4 x 10 -3 4.2 x 10-4 0"0670 2.6 × 10 -3 3 "9 )< 10 -4 0"0938 2"5 × 10-3 3.4 x 10 -4 0"1305 3"3 x 10 -3 3"8 x 10 -4 Mean value KAzNO~ = 2"7 +--0"5 X 10-3 .... (KAgNO~) M, O = 3-8 +0.2 x 10 -4

point depression. The third column lists the value of the dissociation constant of silver nitrate calculated from the freezing point data assuming ideal behavior; i.e., from the expression KAgNO3 = [(Zm--ms)]z/[Y.m--2(Y,m--ms))]. A n estimate of the thermodynamic dissociation constant requires that some relation for the activity coefficient of ions and the ionic strength (/z = [Ag+]) be assumed. SCHAAP et al.¢6) have suggested that the MarshallGrunwald equation would be applicable to E D A solutions. The fourth column in Table 1 gives calculated values of the dissociation constant of silver nitrate (KAsNO3 M,G) using the latter relation to estimate the single ion activity coefficients. The mean value of (KAsNO,) equals 3"8 --+0"2X 10 -4. The dissociation constant of silver nitrate was estimated to be ,~ 6 x 10-4 in our previous potentiometric studytl) and 5.75 x 10-4 by HmnARD and SCHMn)TtT) USing the Fuoss-Kraus method. SIEFKER(8) obtained similar values using a combination of potentiometric and conductometric data. Finally, an attempt was made to demonstrate the existence of K I . A g I and N a I . A g I by studyingmixtures of the alkali and silver iodides. The latter two species have been postulated to interpret data obtained using a silver indicator electrode, tl) As is shown in Table 2. the total molality as found cryoscopically is substantially lower than the sum of the analytical concentrations of the silver and alkali halide. (s~ H. HII~aER and A. WOERNER,Z. Elektrochem. 40, 252 (1934). (6) W. ]3. SCHAAP,R. E. BAYER,J. R. SiEFKER,J. K. KIM, P. W. BREWSTERand F. C. SC~tDT, Record of Chemical Progress, 22, 197 (1961). (7) B. B. HIBBARDand F. C. SCHMmT, J. Amer. Chem. Soe. 77, 225 (1955). ts) j . R. SmFKBR. Ph.D. Thesis, Indiana University (1960).

NOTI~

1481

TABLE2.---CMYOSCOI~ICDATAFOR AgI'KI ANt) AgI.NaI soLUTIONS MI Potassium iodide Sodium iodide

ms, MI(a) 0.1127 0.0803

ms, Agi(a) 0.1138 0.1271

~m (b) 0-2050 0.1880

KMI.AgI 2.6 3-0

(a) ms is molal analytical concentration of species indicated by subscript. (b) 5'.m is total molal concentration calculated by freezing point depression. Assuming ideal behavior, the formation constant, K MLAgI= [MI'AgI]/[MI][AgI] =

[(ms)Mt+(ms)lmd[~m--(ms)Agd--(mAsd[~m--(ms)Mfl, was calculated from the data in Table 2. The results obtained are in fair agreement with the previously obtained values of K KI'Agl 1 "1 and K N ~ ' ~ = 1.3.(1) Aeknowledgment.--This work was sponsored by the Office of Ordinance Research. =

L. D. P~rrn*

STANLEYBRUCKENSTEIN

School of Chemistry University of Minnesota Minneapolis 14, Minn. * Present address: Department of Inorganic and Structural Chemistry, The University, Leeds.

BOOK REVIEWS

H. H. StaLER: Chemistry in Non-Aqueem Solvents vii + 119 pp. Reinhold, New York, 1916. $1.95 THIS monograph appears to be intended for the undergraduate interested in chemistry rather than either the specialist or educated layman. To this end, the author steers a course between dull scholarship and a Sunday supplement popular science approach and, on the whole, is successful. Current ideas in the chemistry of non-aqueous solutions are given with adequate examples. The classical approach to non-aqueous chemistry in terms of water analogs is illustrated by liquid ammonia. The section on sulphuric acid emphasizes the more recent work of R. 5. GmLeSrm and his colleagues, but fits neatly into acid-base concepts involving proton transfer. The other solvents selected for individual attention are sulphur dioxide and nitrogen dioxide a choice which permits exploration of the range of usefulness and the limitations of the solvent-system ionization concept. The book, however, has some flaws which may reduce its value for the teacher and the serious student. It is not completely authoritative; for instance, the data given in the various tables are not provided with references, and many of the constants may be out-of-date. The observant reader, for example, will have difficulty in reconciling the ion product constant for liquid ammonia described as being in the neighborhood of 10-33, with a specific conductance greater than 10-7 ohm-1. Experimenters (V. F. HNIZDA, C. A. KRAUS,J. Amer. Chem. Soc. 71, 1565-75 1949). have obtained liquid ammonia with a specific electrical conductivity below 10 -t0 ohm -lcm-1. It is likely that the pure liquid is even less conducting and that the ion product estimate of PLeSKOVwhich S~J~t quotes is a realistic estimate.