NOTES
Feb., 193J
M and (CeH&PO= 1.60 X A I ; therefore, a plot of 1/OD-OD where OD is the optical density, vemus time is the proper derived function for a second-order reaction with the reactants a t equal equivalent con~eiitrations.~ The upper abscissa has been added to show the t.emperature in the reaction cell as a function of time. The slope of Fig. 1 a t any given temperature is then corrected for the change in density of the solvent with temperature and for a small change in the molar extinction coefficient of sulfur with temperature. The resultant Arrhenius plot (Fig. 2) of log corrected slope versim 1/ T gives an activation energy of 16.5 f 0.4 kcal./mole, in agreement with 16.0 f 0.2 kcal./mole.
309
THE CHELATING TENDENCY OF RIBOFLAVIN' BY T ~ 0 a r . kR. ~ HARKINS AND HENRY FREISER~ Received August 4 , 1958
The structural similarity of riboflavin (I) and 8hydroxyquinoline (11) has been noted previously by Albert3 in accounting for the ability of the riboflavin to complex various metal ions. His results indicated an unusual metal stability order in that iron(I1) (log KI = 7.1) formed a more stable complex thaii did copper(I1) (log KI = 6.5). Inasmuch as 8-hydroxyquinoline chelates follow the usual stability order, this result was unexpected. Of course, it is possible that the presc~HiiO4
TABLE I RATECONSTANTS I N BENZENE T e m p . , 'C.O
k2, 1. mole-'
sec.-'b
Ref. methode
7.50 x 10-4 0 Ultraviolet 4.40 x 10-3 8,9 Ultraviolet titr. 11.3 x 10-3 8,9 Ultraviolet titr. a =t0.02. Rate = k?(Ss)((C6H&P). Ultraviolet analysis of sulfur, titr. iodometric analysis of the phosphine. 7.35 25.00 35.00
The main requirements for the use of this method are (i) the rate of reaction is moderately slow a t the lowest temperature; (ii) the temperature in the cell must be uniform; (iii) the activation energy does not vary with temperature; (iv) the boiling point of the solvent cannot be exceeded; and (v) Beer's law must be obeyed even though the molar extinction coefficient may change with temperature. Requirement (ii) restricts the size of the cell and the rate of heating. If the activation energy varies greatly with temperature one would be unable to obtain the correct order of the reaction and the rate constants. The main disadvantage of this method is that only a small percentage of reaction is used to determine the rate at any one temperature. The advantage is obtaining the rate, the order of the reaction, the frequency factor and the activation energy in a single rapid experiment.
6H
I
I
OH I1
ence of a sterically hindering group in riboflavin (the fused benzene ring) would have a sufficiently greater effect on copper(I1) which is smaller thaii iron(I1). The sterically hindering group present in 2-methyl-8-hydroxyquinoline was shown to have a greater effect on the chelates of the smaller metal ions.4 Another interesting observation concerning metal-riboflavin complexes was made by Foye and Lange6 who prepared a series of such complexes whose composition corresponded to a two to one mole ratio of metal to riboflavin. The stoichiometry and thermodynamics of the formation of metal-riboflavin reported in this paper was undertaken to evaluate these observations. Experimental
Stock solutions of approximately 0.01 M metal ions were prepared by dissolving their reagent grade perchlorates (G. Frederick Smith Co.) in water. The copper(I1) and cobalt( 11)solutions were standardized by electrodeposition. The nickel( 11) solution was standardized by precipitation with dimethylglyoxime. Zinc(I1) was standardized gravimetrically as ZnNH4P04. The iron(I1) solution was preExperimental pared by dissolving high purity iron wire in perchloric acid The purification of sulfur, triphenylphosphine and benzene under an inert atmosphere. Riboflavin, obtained ofrom the Nutritional Biochemicals has been reported previously.* The use of the ultraviolet Anal. Calcd. C, 54.25; H, 5.36. absorption spectrum of sulfur to study this reaction a t a Corp., was dried a t 110 constant temperature and with conventional techniques Found: C, 55.50; H, 4.99. The titration apparatus and rocedure have been prewill be subject to a forthcoming publication.@ The thermostated cell compartment for a Beckman D U spectrophotom- viously described .e A slight mosification in preparing the solution for titration was undertaken to facilitate the diseter has been discussed.1O The brass jacket was carefully made to ensure good thermal contact with a square Corex solution of riboflavin. Fifty-five milliliters of water was cell. To the top of the cell was sealed a 7 mm. Pyrex t8ube added to a weighed quantity of the reagent. A small meas(15 cm. in length) through which a multi-junction thermo- ured volume of standard base was added to bring about solution after which the perchloric acid and metal perchlorcouple was placed into the cell just above the light path The output of the thermocouple was applied to a Speedomax ate were added and the titration proceeded in the customary recorder. The reactants were mixed and placed in the cell ( I ) Abstracted from the thesis submitted by T. R. Harkins in (total volume 1 to 2.6 4). Water, circulated a t the rate of partial fulfillment of the requirements for the Ph.D. degree at the one gallon per minute through the compartment, was slowly University of Pittsburgh, June, 1956. heated. Optical density measurements were manually re(2) Department of Chemistry, University of Arizona, Tucson. corded every 30 seconds at 345 mp. The heating rate does (3) A. Albert, Biochem. J . , 64, 646 (1953). not enter into the graphic analysis of the data but averaged (4) W. D. Johnston and H. Freiser, Anal. Chim. Acta, 11, 201 0.5 degree per minute to cover the range of 23 to 77". Other (1954). heating rates can be used to cover only a 25' increase. (5) W. 0. Foye and W. E. Lange, J . A m . Chem. SOC.,7 6 , 2199
.
(10) P. D. Bnrtlett and R. E. Davis,
(1958).
J . A m . Chem. SOC.,80, 2513
(1954). (0) €1. Freiser, R. 11952).
G. Charles and W. D. J o h n s t o n , i t i d . , 7 4 , 1383
310 2.0
NOTES
Vol. 63
assuming a value for log Kr (in this case 12.31, arid using the expression
1
pR = 1/2 log 1.5
n
- log Kr\
to evaluate pR a t various values of a. Curve B represents the results of a titration in which the riboflavin-metal rat,io is five to one. This curve has flattened to approximately the slope of curve B and is displaced in the direction of lower log K g 1.0 values. The remainder 'of the titrations were carried out a t the higher riboflavin-metal ratios to avoid interference by hydrolytic phenomena. Were it not for the relatively low solubility of riboflavin, even higher ratios would have been em0.5 ployed. NO result is reported for iroa(I1) because the titration curves for this metal both in the presence and absence of riboflavin were identical. This indicated that hydrolysis is a more favorable reac1 I 0 tion than complex formation. The stability con4 5 6 7 stant previously reported3 is of doubtful validity pR-. Fig. 1.-Formation curves copper(I1)-riboflavin complex. and probably can be attributed to hydrolysis. Thus, the unusual metal stability sequence in fashion. The results of the measurements at several tem- which iron(I1) appears before copper(I1) may not peratures as well a s the thermodynamic quantities calcu- be claimed for riboflavin. lated therefrom are summarized in Table I. The results reported here do show a displaced metal stability sequence reminiscent of what was TABLE I found for 2-methyl-8-hydro~yquinoline.~ The exCHELATE FORMATION DATAFOR RIBOFLAVIN I N WATER traordinary sensitivity of nickel to the influence of (5: 1 Ratio of Riboflavin to Metal) steric hindrance to chelate formation exhibited Temp., 2 log AH3 AS. with 2-methyl-8-hydroxyquinoline would seem to Metal ' C . Log K I Kav AF, ked. e.& account for its displacement below zinc with riboCu(I1) 10 6.0 11.7 flavin whose fused benzene ring in juxtaposition 25 5.9 11.6 -15.8 -3 44 to the nitrogen involved in chelation offers steric 40 5.9 11.5 hindrance. A comparison of the AH values with 25a 6 . 2 12.4 those of 4-methyl-8-hydroxyquinoli11e as was made Ni(I1) 10 4.2 7.8 with 2-methyl-8-hydroxyquinoline4 reveals a simi25 4.2 7 . 8 -10.6 $3 46 lar trend of decreasing difference with increasing 40 4.1 8.0 ionic radius of the metal, Also the A S values found Co(11) 10 4.1 7.6 here are very close to those observed with the 225 4.3 7 . 8 -10.6 $5 52 methyl-8-hydroxyquinoline. 40 4.3 8.1 The isolation of metal-riboflavin complexes in Zn(I1) 10 4.9 9.5 which there is a 2: 1 molar ratio of metal to ribo25 4.9 9 . 6 -13.1 + I 47 flavin could not be predicted from the data obtained in this investigation. The titration data 40 4.9 9.6 indicate that two moles of hydrogen ion are liberReagent + Hf 10 10.02 ated per mole of metal ion in the reaction with 25 9.69 -13.2 -8.3 16 riboflavin. Since the metals are divalent, this 40 9.40 would satisfy their neutralization of charge require2: 1ratio of riboflavin to metal. ments. Structures correspondillg to the formulas assigned to complexes isolated in basic solution by Discussion Initially the titrations were carried out using a Foye and Lang5 (e.g., MsR.H20) would be diftwo to one mole ratio of riboflavin to metal as ficult to visualize on this basis. The stoichiometry Albert had done.3 Under these conditions, the re- of two metal ions per riboflavin molecule might he sults for copper(I1) agreed with those he reported. accounted for with one metal ion chelating with the However, the titration curve indicated an overlap oxine type grouping and the other metal ion bindof the complex formation with hydrolysis. Fur- ing to the ribitol group, To account for a neutral ther, the formation curve obtained indicated that species being formed, however, a negative group something was amiss since its slope was too steep. such as hydroxyl would have to be bound to the I n Fig. 1, curve A which represents the results of metal attached to the flavin portion of the molethe titration with a 2 to 1 riboflavin-metal ratio is cule. compared with curve C which has been calculated Acknowledgment.-The authors gratefully acon the basis of simultaneous addition of two moles knowledge the financial assistance of the U. S. of reagent to the metal. Curve C was derived by Public Health Service. I
NOTES
Fcb., 1959 THERMAL CONDUCTIVITY OF POLYCRYSTALLINE BORON1 BY CLAUDE P. TALLEY Experiment Incorporated, Richmond 2 , Virginia Receii,ed Augzrst 8 , 1958
To the author’s knowledge there have been no previous reports of the thermal conductivity of boron. In the course of preparing samples of boron to study its oxidation properties, cylindrical polycrystalline boron rods about 1 mm. in diameter and several centimeters long were made which contained a 0.025 mni. diameter tungsten core. Such specimens are well suited to the measurement of thermal conductivity, and it was decided to take this opportunity to obtain a t least an order-of-magnitude estimate of this property for boron. The boron rods were prepared by the reduction of boron tribromide by hydrogen near a 0.025 mni. diameter tungsten filament a t about 1250”. Wet chemical analysis for total boron showed that the boron content exceeded 99% by weight. The main impurity in the rods was that due to the tungsten core, which amounted to about 0.7% by weight. Emission spectrographic analysis indicated small amounts of Cs, Fe, Cu, Mg and Si amounting to a total of 0.02% by weight. The thermal conductivity of a polycrystalline boron rod was measured in the following manner. The tiny 0.025 mm. diameter tungsten core was heated electrically and served as the heat source. The circuit consisted of battery, variable resistor, ammeter, voltmeter and the specimen immersed in a constant temperature water-bath. At the temperatures used in these experiments the electrical conductivity of the boron was negligible, and therefore practically all the electrical current that passed through the rod flowed through the tungsten filament located in the core. Electrical contact wns made to the tungsten by fusing on platinum leads a t each end of the rod. Heat generated in the tungsten was conducted through the cylindrical casing of boron and into a surrounding water-bath. The temperature of the inside surface of the boron was assumed to be equal to the tempernture of the tungsten. The temperature of the tungsten was obtained by computing its resistance from the measured voltage and current supplied to the rod and comparing this resistance with a previously experimentally determined curve on the same rod, a t negligible power input, of resistance versus temperature. The temperature of the outside surface of the boron vas nasnmed equal to that of the stirred nrater-bath. In a t>,pical esperinient Kith a 1.19 mm. diameter by 27.4 min. length rod, the power supplied to the 0.025 mm. diameter tungsten core wns 0.iO cal./sec., the resistance of the tungsten corresponded to a temperature of 80°, and the ternperature of the water-bath was 22”. In these preliminary experiOC./cm.) was ments an average value of 0.003 cal./(sec. ohtained for the thermal conductivity between 20 and 80’. The equ:it8ionfor the calculation of t h i thermal conductivity IVaS2
I‘ =
Q In r h 27rL(T1 - Tz)
where
K = thermal conductivity
Q = rate of heat conduction
rz = outside radius of boron rod rl = radius of the tungsten core L = length of boron rod TI = temperature a t T I Tz = temperature a t rz (1) This work was supported by the Office of Naval Research. (2) W. H. RlcAdams, ”Heat Transmission,” 3rd Ed., McGraw-Hill Rook Co., Inc., New York, N. Y., 1954, Chapter 2.
31 1
Heat loss a t the ends of the rod was neglected because of the high length-to-diameter ratio of the rod. By electrically probing a polished cross section of a polycrystalline boron rod for electrical resistance, it was found that the tungsten filament in the center had remained essentially unaltered during deposition, and therefore rI was taken as equal to the starting radius of the tungsten filament. For example, a t a distance of about 0.010 mm. from the center of the rod the resistance was only a few ohms, whereas a t about 0.025 mm. from the center of the rod the resistance increased to about 150,000 ohms. Also the tungsten core appeared under the microscope to be a maximum of 0.035 mm. in diameter. Using this same technique a value was obtained for the thermal conductivity of Pyrex glass rods containing a central 0.025 mm. diameter tungsten filament which agreed with published values within a factor of three. Considering that the thermal conductivity of Pyrex IS about the same as obtained on polycrystalline boron and considering the accuracy of measurement of the individual quantities, the thermal conductivity value obtained for boron is thought to be accurate to within a factor of three also.
T H E DESICCATION AND DENSITY OF ACETONE B Y K.s. HOWARD AND F.P. PIKE Department of Chemical Engineering North Carolina Slate College Raleigh, N. C . Received August 1 1 , 1068
The dehydration of otherwise pure acetone has been a troublesome problem for many years. The ordinary inert desiccants, such as CaC12and CaSO4, are extremely slowacting1 and are ineffective2 in the complete removal of water, since snlall amounts of water are retained very tenaciously by the acetone phase. Formation of an addition compound between acetone and sodium iodide3 with subsequent regeneration of ((anhydrous” acetone was used by Young4 t o obtain a product with a density of 0.79053 g./nil. at 20”. Timmermans2 has used PzO5 for dehydration, a process which involves great loss of material through condensation reactions and requires an efficient distillation column for separation of the product. Iiiterpolation of the Timmermans data gives a density of 0.7904 g./ml. a t 20°, and this value was accepted for many years as the density of anhydrous acetone. However, Thirion and Craven5 used desiccation by acetic anhydride, folloved by distillation, to obtain a product with an average density at 20“ of 0.78990 f 0.00006 g./ml., significantly lower than these previous values. The results of Thomas and McAllister’ confirm this lower density. The current work was begun in an attempt to find a simple procedure, employing ordinary laboratory equipment, for complete dehydration of acetone. Use of the sodium iodide adduct was no more successful here than it had been iii Young’s hands.4 Dehydration by CaH2 or acetic anhydride led to great material loss and separation problems. An attempt to titrate the water by modifications of the Karl Fischer technique,6 followed by distil(1) K. T. Thomas and R. A. MoAllister. A . I . Ch. E. J . , 3 , 161 (1957). (2) J. Timmermans, “Physico-Chemical Constants of Pure Organic Compounds,” Elsevier Publ. Co., Inc., New York, N. Y . , 1950, p. 354. (3) K. Shipsey and E. A. Werner, J. Chem. S o c . , 103, 1255 (1913). (4) W. Young, J . Soc. Chem. I n d . , 5 2 , 449 (1933). ( 5 ) P. Thirion and E. C. Craven, J . A p p l . Chem., 2 , 210 (1952). (6) J. Mitchell, Jr., and D. 11. Smith, “Application of the Karl F~scllerReagent to Quantitative Analyses Involving Water,” InterBcience Publishers, Inc., New York, N. Y., 1948.
