I
CORRESPONDENCE
1
I J
Use of Tris(hydroxymethy1)-aminomethane as a Comparison Standard for Thermometric Titration Calorimetry SIR: At present there is no generally accepted comparison standard for thermometric titration calorimeters unless it is the neutralization of NaOH with HC1. The only attractive property of this reaction is the high reaction enthalpy (-13.5 kcal/mole). Even this is partially negated by the variation in reaction enthalpy with reactant concentration (1). Furthermore NaOH solutions react with atmospheric carbon dioxide. A substance which meets the requirements for a comparison standard for thermometric titration calorimeters is THAM [tris(hydroxymethyl)-aminomethane]. This substance is already in the process of being evaluated for use as a comparison standard for conventional solution calorimeters (2). There is a fortuitous difference between the use of THAM in a conventional solution calorimeter and thermometric titration calorimeter. Equation 1 gives the reaction which would be used for a
THAM(c)
+ H+(aq) -,THAM:H+(aq)
(1)
conventional solution calorimeter. Because thermometric titration calorimeters are limited to the use of fluids, one cannot use the reaction shown in Equation 1. The primary standard THAM must be made into a standard solution and then reacted according to Equation 2. Table I summarizes information pertinent to the THAM(aq)
+ H+(aq)
+ THAM:H+(aq)
(2)
use of THAM as a comparison standard for calorimeters. The results show that the use of a standard solution of THAM actually enhances its use as a calorimetry standard. The enthalpy accompanying the dissolution of THAM crystal is on the order of + 4 kcal/mole. By dissolving the crystals first, one obtains an additional 4 kcal of heat evolved per mole neutralized. Additional advantages to be gained from using THAM instead of NaOH are: (a) (b) (c) (d) (e)
can be obtained in a pure form (99.94%); reaction with atmosphere is negligible; can be easily dried at 100 "C and weighed in the air; has a high equivalent weight (121.137 grams); is available from National Bureau of Standards as a certified calorimetric standard.
In addition, titration of THAM with HC1 causes no change in ionic strength. It is for these reasons and the excellent precision obtained for the reaction enthalpy of THAM with HC1 that it should be considered as a comparison standard for thermometric titration calorimeters. (1) J. D. Hale, R. M. Izatt, and J. J. Christensen, J. Phys. Chem., 67, 2605 (1963). (2) S . R. G u m , ibid.,69,2902 (1965).
Table I. Summary of Properties of the System THAM-HCl-H,O Reaction: THAM(c, 6 g) HCl (aq, 0.100M) -+ THAM :H+(aq) Enthalpy of Reaction": -7107.0 f 0.9 calimole Reaction: THAM(aq, 0.0140M) HCl(aq, 1.812M) + THAM: H+(as) Enthalpy of Reaction: - 11.38 5 0.06 kcal/mole Enthalpy of Reaction after correction for heat of dilution, etc. : - 1 1 . 3 5 i 0.06 kcal/mole rneg THAM added meg THAM found % error 3.851 f 0.028 3.816 f 0.024 0.91 a See reference (2).
+
+
EXPERIMENTAL
Reagents. The THAM (Eastman Organic Chemicals, primary standard) was first dried for 2 hours at 100 "C in an Abderhalden drying apparatus before it was used to standardize the HCl (3). A carbonate free NaOH solution previously prepared was diluted and standardized with the acid mentioned above. Apparatus. The thermometric titration calorimeter description is given elsewhere (4). The corrected value for the heat of neutralization of NaOH with HCl as determined with this calorimeter was -13.41 f 0.54 kcal/mole and is in excellent agreement with a recent value by Hale et al. (1) of - 13.42 using a conventional solution calorimeter. Treatment of Data. The data were treated according to the method described by Jordan (5).
EDMOND W. WILSON,JR. DONALDF. SMITH Department of Chemistry University of Alabama University, Ala. 35486 1 Present address, Department of Chemistry, University of Virginia, Charlottesville, Va. 22901
RECEIVED for review August 1 , 1969. Accepted September 4, 1969. Taken in part from the Ph.D. dissertation of E. W. Wilson, University of Alabama, University, Ala., 1968. (3) J. H. Fossum, P. C. Markunas and J. A. Riddick, ANAL. CHEM., 23,491 (1951). (4) Edmond W. Wilson, Jr., and Donald F. Smith, unpublished material, 1969. (5) J. Jordan, Rec. Chem. Progress, 19, 193 (1958).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
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