STUART R. GUNN
2902
Comparison Standards for Solution Calorimetry’
by Stuart R. Gunn Lawrence Radiation Laboratory, University of California, Livermore, California
(Received February 16, 1966)
Criteria for processes to be used in checking and comparing solution calorimeters are discussed. Measurements have been made of the heats of four processes: dissolution of potassium chloride in water, dissolution of succinic acid in hydrochloric acid, reaction of sulfuric acid with excess sodium hydroxide, and reaction of tris(hydroxymethy1)aminomethane with hydrochloric acid. It is concluded that the last two are most generally satisfactory. Corrections to be applied in different types of solution calorimeters and some sources of instrumental error are discussed.
Introduction Solution calorimetry may be broadly defined as calorimetry of all processes-dissolution, mixing, dilution, and reaction-occurring primarily in a liquid solution. This field has suffered from the lack of any generally accepted standard process for the intercomparison of calorimeters ; many investigations exhibit good internal precision but are not free from suspicion of significant systematic error. Standard substances for calorimetry may be considered to be of two kinds: first, calibrating standards, which are used to transfer the unit of energy from a standardizing laboratory for the purpose of calibrating other calorimeters, and second, comparison standards, which are used to compare the operation of various calorimeters with one another, the actual calibration of each calorimeter being, however, performed electrically or by other means. This second class should perhaps be termed “standard substances” in accordance with the nomenclature for standards proposed by McNish,2 although this seems less descriptive of their function. The outstanding example of the first class is benzoic acid for combustion calorimetry; in the second, the most widely used are probably n-heptane, aluminum oxide, benzoic acid, diphenyl ether, and water for heat capacity ~alorimetry.~ Three important reasons for the utility of a calibrating standard in combustion calorimetry may be cited. (1) The calorimetric instrumentation is rather highly standardized. (2) The heat capacity of the reacting system-benzoic acid and oxygen-is a very small fraction of the energy equivalent of the calorimThe Journal of Physical Chemistry
eter, which is hence altered only slightly in changing from the calibrating system to the experimental system. (3) The combustion reaction in the bomb is very rapid and hence the heat distribution in the calorimeter is primarily a function of its mechanical design and stirring, which also are the same for both calibrations and combustions of unknowns. None of these conditions obtain in solution calorimetry, where the instrumentation is varied widely according to the nature of the problem under investigation, the reacting system is usually the major part of the heat capacity of the calorimetric system, and the reaction kinetics are widely variable. Electrical calibration of the calorimeter will probably always be necessary for the most accurate work. Nevertheless, it would seem to be possible and desirable to establish one or a few comparison standards of fairly broad utility for the intercomparison of solution calorimeters, the function being primarily the detection of systematic errors between calibrations and reaction heats. The following desirable characteristics of the comparison process may be suggested. (1) The process should consist of mixing a weighed amount of a standard substance-solid or liquid-in a frangible bulb or other sample chamber with a large volume of a liquid. (2) The liquid should be preferably pure water; (1) Work performed under the auspices of the U. S. Atomic Energy Commission. (2) A. G . McNish, IRE Trans. Instr., 1-7, 371 (1958). (3) D. G. Ginning8 and G . T. Furukawa, J . Am. Chem. Soc., 75, 522 (1953).
COMPARISON STANDARDS FOR SOLUTION CALORIMETRY
if not, then an aqueous solution, preferably of neutral PH. (3) The process should be rapid and complete. (4) An exothermic process would be more desirable than an endothermic one. (5) The temperature change should be about l o , but the heat of the process should be known withlittle loss of accuracy at conditions giving considerably smaller changes. (6) The heat should be defined at 25”, but the temperature coefficient should be small. (7) The heat should be fairly insensitive to the ratio of standard substance to liquid, and to the concentration of solutes, if any, in the initial liquid. (8) Exposure of the liquid to the atmosphere should have a negligible effect on the heat. (9) The heat should be the same in closed and open calorimeters, or the difference should be readily calculable. (10) No gases should be evolved. (11) The change of vapor pressure of the liquid should be small. (12) The change in volume of the liquid should be small. (13) The volume required of the sample chamber should be no more than 5% of the liquid, and preferably less. (14) The mass of the standard substance should be large enough for convenient and accurate weighing. (15) The heat capacity of the standard substance should be small (this is important in isothermally jacketed calorimeters where the temperature of the standard substance isolated in the sample chamber may lag significantly behind that of the liquid during the foredrift; it is unimportant in adiabatic or isothermal calorimeters). Point 15 generally conflicts with 14. (16) The standard substance should be available in high purity, or be readily purifiable. (17) The standard substance should be nonhygroscopic and nonvolatile and nonreactive with the atmosphere. (18) The standard substance should not be subject to significant energy perturbations from crystal defects or strain or surface energy effects. The heat of solution of KC1 has most often been standad. It has been proposed as a measured many times, but the results are wildly disevaluated the cordant. Mischenko and Kaganovich4 literature, recoinmending a value of 4194 f 3 tal. for the heat Of to give KC1’200H20‘ Somsen, Coops, and Tolk5 discussed the virtues of the system as a standard; from an extensive series of
2903
measurements at lower concentrations they give 4185 j= 2, adjusted to KC1.200H20. Parkere listed 70 investigations from 1872 to 1962 and recommended a “best” value of 4115 f 10 at infinite dilution, or 4195 at KC1.200H20, but even the more modern values are scattered over a range of more than 1%. Sunner and Wadso7 concluded that the system is unsatisfactory. I n the present work, the effect of some variations in the treatment of the KC1 and in the method of temperature measurement in the calorimeter have been investigated. The heat of neutralization of a strong acid by a strong base has sometimes been used or suggested as a standard. It appears, however, that this heat has been more uncertain than sometimes thought. Calorimetric values for the heat of ionization of water (heat of acid-base neutralization corrected to infinite dilution) scattered over a range of ca. 0.5% and furthermore differed systematically by ca. 1% from electrometric values. Recently, however, Vanderzee and Swansons have published precise measurements of the heat of neutralization of sodium hydroxide and perchloric acid and Hale, Izatt, and Christenseng have published slightly less precise but excellently agreeing results for sodium hydroxide-perchloric acid and sodium hydroxide-hydrochloric acid. The discrepancy with respect to electrometric determinations is attributed to small systematic effects in the treatment of e.m.f. data to derive the ionization constant.8 I n the present work, the objective was to measure a neutralization heat very accurately at somewhat higher concentrations, without concern for correction to infinite dilution. Standardization of one of the reagents and introduction of an accurately known amount of it into the calorimeter present some problems compared with the more customary introduction of a weighed sample of a pure compound, but these are not insurmountable. KunzlerlO performed an exhaustive study of the preparation of accurately known sulfuric acid solutions although the work seems to be generally ignored in acidimetry. He concluded that acid prepared by the (4) K. P. Mischenko and Y. Kaganovich, Zh. Prikl. Khim., 22, 1078 (1949). (5) G . Somsen, J. Coops, and M. W. Tolk, Rec. trav. chim., 82, 231 (1963). (6) V. B. Parker, “Thermal Properties of Aqueous Uni-univalent Electrolytes,” National Bureau of Standards Report NSRDS-NBS-2 (1965). (7) S. Sunner and I. Wadso, Acta C h m . Scand., 13, 97 (1959). (8) C. E. Vanderzee and J. A. Swanson, J. Phys. Chem., 67, 2608 (1963). (9) J. D. Hale, R. M. Izatt, and J. J. Christensen, ibid., 67, 2605 (1963). (io) J. E. Kunzler, A ~ ZChem., . 25,93 (1953).
Volume 69, Number 9 September 1966
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constant-boiling procedure was accurate in composition to O.Ol%, and better accuracy was obtainable by some other methods. The reaction of this with excess sodium hydroxide promised to give a rapid and reproducible heat and has been investigated in the present work. There appear to be no other accurate data under comparable conditions for comparison. Succinic acid has been extensively investigated as a secondary standard for combustion calorimetry and appears to be satisfactory, although it has not been widely used. Previous work on the heat of solution is limited and apparently quite inaccurate, but the compound possesses most of the properties desired of a reference material and, at the suggestion of Wilhoit," the heat of solution was further investigated. HC1 (0.1 M ) was chosen as the solvent instead of pure water to repress ionization at low concentrations and permit the linear extrapolation of heats of solution expected for a nonelectrolyte. The dissolution is endothermic, like that of potassium chloride; an endothermic as well as an exothermic comparison reaction may occasionally be desirable. Wadso and Irving12have recently proposed the heat of reaction of tris(hydroxymethy1)aminomethane (or 2amino-2-hydroxymethyl-l,3-propanediol, hereinafter designated THAM) with aqueous hydrochloric acid. THAM has properties qualifying it as a primary acidimetric standard.I3 The reaction seems quite suitable as a calorimetric comparison; several other laboratories are also investigating it. In the present work, the heat has been measured over a range of THAM and acid concentrations. Shomate and Huffman14 determined the heat of solution of M g in 1 M HC1. The reaction has since been used occasionallyto check the operation of calorimeters, especially reaction microcalorimeters; agreement has generally been good, but none of the subsequent determinations are sufficiently precise or exhaustive to add anything to the reliability of the original determination. The particular advantages of the system are that because of the high molar heat much smaller samples can be used and concentration and dAH/dT effects are much smaller. The evolution of hydrogen, carrying a perhaps indefinite amount of water out of constant-pressure calorimeters, is a major disadvantage. ;\lore insidious is the uncertain purity of available metal; while multiply distilled magnesium is quite low in metallic impurities and these can be adequately checked spectrographically, variable inclusions of nonmetals, especialIy oxygen, seem to be more of a problem and these are more difficult to determine. It was originally intended to reinvestigate this system in the present work, but in view of the above difficulties it The Journal of Physical Chemistry
STUARTR. GUNN
now appears that there would be few situations where it would offer any advantage over the H2S04-NaOH or THAM-HC1 reactions.
Experimental Materials. Four batches of potassium chloride were used. The first three were all taken from a single jar of Baker and Adamson reagent grade material. Batch A was oven-dried 18 hr. at loso,batch B was heated 16 hr. at 720" in air in a muffle furnace, and batch C was fused in air in the muffle furnace and cooled slowly. Batch D was Harshaw optical quality, random-sized pieces cleaved into plates about 1 mm. thick. The salt was stored in a desiccator, then weighed into annealed Pyrex bulbs which were evacuated, filled with 5 torr of helium (to promote thermal equilibration during the foredrift of the calorimeter run), and sealed off. Weighings were corrected to in vacuo using a density of 1.98 g. ml.-l. Flame photometric analysis indicated about 0.004 mole % NaCl in both the Baker and Adamson and Harshaw material. The same result was found in a sample from the jar used in previous measurement~~ evidently ~; the 0.028 mole yo reported at that time was erroneous. Constant-boiling sulfuric acid was prepared as described by Kunzler,'O using a still similar to his design. A batch of reagent grade sulfuric acid was first distilled into the receiver, the still was washed out and dried, and the distilled acid was put in the pot. About half of it was then distilled over at a rate of about 5 ml. min.-l while dry nitrogen was flowed slowly through a tee connected to the receiver. The still was allowed to cool; then the cap of the pot was removed while the tee was temporarily blocked to give a slow flow of dry nitrogen through the still, and a joint closed by a stopper with a tube passing through it was quickly placed on the tapered joint. The tube dipped into the acid in the still and, outside, was bent to a vertical termination with constricted tip. A 1-1. Pyrex bottle was fitted with an adapter, replacing the stopper, which had a small-diameter standard taper and cap at the top. The bottle was previously flushed with dry air and weighed against a tare. The cap was then removed, the bottle quickly raised around the tube from the still, and about 100 ml. of acid forced over with nitrogen pressure. The bottle was reweighed and water was (11) R. C.Wilhoit, private communication. (12) I. Wadso and R. 3. Irving, Acta Chem. Scand., 18, 195 (1964). (13) J. H.Fossum, P. C. Markunas, and J. A. Riddick, Anal. Chem., 23,491 (1951). (14) C.H.Shomate and E. H. Huffman,J . Bm. Chem. $oc., 6 5 , 1625 (1943). (15) S. R. Gunn, Rev. Sci. Instr., 29, 377 (1958).
COMPARISON STANDARDS FOR SOLUTION CALORIMETRY
added very slowly, while the bottle was cooled in an ice bath, to the desired composition, and the bottle was again weighed. A Teflon-coated magnetic stirring bar was placed in the bottle, and the solution was thoroughly mixed by stirring and shaking. All glassware, including the calorimeter bulbs, had been kept in contact with concentrated H2S04 or 6 M HzS04 for 2 months before use. Three calibrated sets of weights were intercompared and used. Weighings were corrected to in vacuo using 1.83 g./ml. for the density of constant-boiling sulfuric acid and 1.15 for H~SOC 20H20. Four batches were used; they are described in Table I. The compositions are calculated using the percentage of HzS04given by Kunzlerlo for the pressure of the distillation. Table I: The Sulfuric Acid Solutions
Betoh
A B C
D
Pressure of distillation, mm. (cor.)
Constantboiling acid, g.
Solution, g.
Moles of H a 0 / moles of HzSO4
749 744 754 744
192.705 213.424 189.422 176.765
891.621 982.403 871.851 813.717
20.0013 20.0013 20.0000 20.0031
The calorimeter bulbs were blown on 12/30 standard taper male Pyrex joints and were annealed, cleaned, and constricted about 2.5 cm. above the shoulder. The solution was run in from a thin-tipped pipet coated with a silicone water repellent on the outside of the tip. Some of the bulbs were coated on the inside of the neck at the shoulder to prevent solution from creeping up the neck. The solution was weighed to 0.1 mg. The bulbs were capped during the weighing and showed no evaporation. After being weighed, the cap was removed, the joint greased, a compressed rubber bulb attached, the sample bulb chilled in an ice bath, and the constriction sealed off. This left about 0.9 atm. of air in the gas space, which was 20 to 30% of the total bulb volume. The sodium hydroxide solutions for reaction with sulfuric acid were made up by weight from Hellige carbonate-free standardized concentrates. Succinic acid batch A was J. T. Baker reagent grade, oven-dried at 120" for 17 hr., crushed and sieved to 4580 mesh, and stored in a desiccator over Mg(C104)Z. Batch B was the same reagent recrystallized four times from water, dried over Mg(C104)2at room temperature, crushed, sieved, and further dried to constant weight over Mg(C1O4)2a t room temperature. Batches C and D were supplied by Dr. R. C. Wilhoit: batch C
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was reagent grade material twice recrystallized from water and then sublimed at 0.1 mm. and 130"; batch D was this material further recrystallized from water and dried 4 hr. at 80" and 5 mm. pressure. Batches E, F, G, H, and I were all material of batch B heated in various manners: E was heated in the oven at 110" for 38 hr. ; F and G were portions of unsublimed material heated under high vacuum at 115" for 4 hr. and 130 for 3 hr., respectively; H and I were sublimed under high vacuum at 130" and 140", respectively. Weighings were corrected to in vacuo using a density of 1.572 g./ml.18 THAM batch A was Eastman 99.94% reagent, twice recrystallized from water-methanol solution as described by Irving and Wads012 and Fossum, Markunas, and Riddick,la screened 45-80 mesh, oven-dried 22 hr. at 80", further dried under high vacuum at room temperature for 66 hr., and stored 5 months over calcium chloride. Batch B was kindly supplied by Dr. Wadso (his designation: sample D), and had been purified by a similar procedure and screened 50-100 mesh. It was shipped by air from Sweden in plasticcapped bottles; when stored in shallow dishes in a desiccator over calcium chloride it lost ca. 0.003% in weight, attaining constancy in a few days. Weighings were corrected to in vacuo using a density of 1.35 g./ml." Hydrochloric acid solutions for the succinic acid and THAM work were made up from Hellige standardized concentrates. The 1961 atomic weights were used throughout: KC1, 74.555; H2S04, 98.078; CJ3604, 118.0900; C4H1103N, 121.1372; HzO, 18.0153. The Calorimeter. Rocking bomb calorimeter ID, similar to others previously described,15was used. It is fabricated of coinage gold; the bulb-breaking mechanism is of tantalum. The bomb weighs 3.6 kg. and has an internal volume of 650 ml. Temperature measurements were performed with a 2000-ohm thermistor, d.c. bridge, Liston-Becker amplifier, and recorder. A Leeds and Northrup G-2 Mueller bridge with two 100-ohm standard resistors added externally was used for the earlier work (potassium chloride, succinic acid except runs 22-25, and sulfuric acid runs except batch D) ; the remaining work was done with a Tinsley Type 5415 Wheatstone bridge, which proved to be more satisfactory. In each run the calorimeter was twice calibrated by electrical heating over nearly the same temperature interval as that of the reaction; the thermistor and bridge thus function as a comparison system and need not be accurately calibrated. The (16) "Handbook of Chemistry and Physics," 45th Ed., Chemical Rubber Publishing Co., Cleveland, Ohio, 1964, p. C-548. (17) H. A. Rose and A. V. Camp, Anal. Chem., 27, 1356 (1955).
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combination was, however, calibrated against a platinum resistance thermometer and a group of mercuryin-glass thermometers to establish the reference (initial) temperature of the reaction with an accuracy of 0.01”. Standard cells and standard resistors were calibrated at this laboratory against standards certified by the National Bureau of Standards. A calibrated Rubicon Type B potentiometer was used for the earlier work, done with the Mueller bridge; a Honeywell Model 2802 potentiometer was used for the remainder. The heating interval, usually about 3 min., was controlled by a pendulum-type standard clock; the relay lag and chatter were found oscilloscopically to introduce an error of no more than a few milliseconds. The electrical energy measurements are believed to have an over-all accuracy of 0.01%. The difference of duplicate calibrations of the calorimeter is limited, especially at smaller temperature rises, primarily by the noise level and reproducibility of the thermistor-thermometer system. For the 12 potassium chloride runs, with a temperature rise of ca. 0.9’, the average difference was 0.005% with a maximum of 0.015%. For the earlier sulfuric acid work with similar rises, the differences were similar; for the three runs with rises ca. 0.3’, differences were 0.016, 0.007, and 0.010%. During the succinic acid work except runs 22-25 performance was somewhat poorer. The seven sulfuric acid runs with batch D, having rises ca. 0.9’, showed an average difference of 0.005% and a maximum of 0.008%. The THAM runs, with rises ca. 0.3’, showed average differences of 0.006%, with a maximum of 0.015%; performance was similar at the higher temperature rises. Calculation of the Corrected Temperature Rise. The corrected temperature rise was obtained by the method of D i c k i n ~ o n lwherein * ~ ~ ~ a time t, is found by graphical integration such that
where t i and tf are the initial and final times of the main period and 0 is the calorimeter temperature, and the corrected temperature rise is the difference of the foredrift and afterdrift extrapolated to t,. The thermal leakage modulus of the calorimeter varies with the volume of solution used, but in the present runs ranged largely from 0.0008 to 0.0010 min.-’; for the higher value, an error of 6 sec. in t, will result in an error of O.Olyoin the corrected temperature rise. The temperature 0 which should be used in evaluating the integrals of eq. 1 is the average surface temperature of the calorimeter, not a temperature measured at a single point internally. This is because the heat transThe Journal of Physical Chemistry
STUART R. GUNN
fer from an element of the surface of the calorimeter by radiation and residual gas conduction is proportional to the temperature difference between that surface element and the jacket, and the total heat transfer is proportional to this temperature difference integrated over all surface elements. This conclusion will be somewhat perturbed by conduction through the heater and thermometer leads and the supporting cords, and in other types of calorimeters by such components as hangers and stirrer shafts; if these components contribute a substantial part of the heat transfer, the calorimeter temperature in their vicinity must be weighted more heavily. It is convenient to speak of a “reaction lag,’’ the difference in time between initiating a reaction by breaking the sample bulb and t,, and a “heating lag,” the difference between the midtime of an electrical heating period at uniform power and t,. These may be further specified as “internal” or “surface” to denote whether measured by the regular internal thermometer of the calorimeter or by a special external thermometer designed to measure the average surface thermometer. A series of experiments was performed with a no. 40 copper coil thermometer wrapped on the outside of the bomb and connected to a recording bridge. This has been shown to have negligible lag with respect to the bomb surface. The winding covered the cylindrical portion of the bomb but not the hemispherical ends, but the winding density was increased at the ends of the cylindrical portion and it is believed that the response of the winding should approximate rather closely the average surface temperature. The heating lag ranges from 25 to 45 sec., being longer when the volume of solution is greater and the stirring is evidently less efficient, but is the same within 1 or 2 sec. measured internally and on the surface. Three KC1 dissolutions were performed, using material from batch 1; the internal reaction lags were 113, 115, and 118 sec.; the corresponding surface reaction lags were 53, 57, and 55 sec., respectively. For two sulfuric acid neutralizations, internal reaction lags (neglecting a small bump which occurs immediately upon breaking and is undoubtedly due to the fact that the sample bulb is mounted immediately adjacent to the thermometer well) were 47 and 58 sec.; the corresponding surface reaction lags were the same within 1 sec. A plausible explanation of the difference in lags of KC1 dissolutions can be constructed. Evidently the salt falls to the bottom of the bomb and lies there, mov(18) H.C.Dickinson, Bull. Natl. BUT.Std., 11, 189 (1914). (19) J. Coops, R. S. Jessup, and K. Van Ness, “Experimental Thermochemistry,” F. D. Rossini, Ed., Interscience Publishers, Inc., New York, N. Y.,1956,pp. 28-35.
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COMPARISON STANDARDS FOR SOLUTION CALORIMETRY
ing along the end somewhat as the bomb is rocked, dissolving for about 4 min. This results in a local “cold spot” much cooler than the body of the solution which is “seen” by the internal thermometer; and while most of the bomb wall would be nearly in thermal equilibrium with the body of the solution, this spot is so much colder that the average surface temperature is lower than the internally measured temperature. The surface thermometer is not used in normal operation; the internal reaction lags observed for the 12 runs were for batch A, 114, 129, and 114; bath B, 118, 115, 106; batch C, 149,157, 156; batch D, 141,151, and 134 sec., in order. The results were calculated using one-half of these values to determine t,. An obvious advantage of the sulfuric acid reaction over the potassium chlor?de appears here: the reaction occurs in the bulk of the solution and not at a solid-liquid interface which may become located near the calorimeter wall and introduce a perturbing effect on the temperature rise correction. A similar but smaller effect was observed for dissolution of succinic acid; the results are calculated for t, at three-fourt’hs of the observed internal reaction lags, which ranged from 55 to 150 sec. For THAM, the effect was still smaller; heats are calculated for a t, 6 sec. before the observed internal reaction lags, which were mostly 40 to 50 sec. Since this surface temperature effect might contribute an error to other investigations where in the surface temperature of the calorimeter was not examined, an apparatus was built to approximate a common type of solution calorimeter. A 1-1. spherical flask with about a 2.5-cm. neck was surrounded by an air gap and water jacket and fitted with a propellor at the center producing a downward flow. The flask was filled nearly to the neck. Internal temperature was measured with a mercury-in-glass thermometer, and the surface temperature was measured by a uniformly distributed copper resistance thermometer. Reactants were introduced in some cases by breaking a bulb and in others by pouring through a funnel. The propellor speed was varied. For four KC1 dissolutions, internal and surface reaction lags were: 243, 183; 66, 46; 52, 31 ; and 18, 10 sec. For a sulfuric acid neutralization, the internal lag was 6 sec. and the external 13sec. This last discrepancy is consistent with the metal bomb performance since the glass wall of the flask is quite thick, much thicker than the thermometer bulb, and it indicates that the real difference between the lags for the KC1 dissolutions should be somewhat greater than here indicated. The KC1 in all cases lay at the bottom of the flask; the time required for complete dissolution was highly sensitive to stirrer speed.
These observations indicate the importance of studying the surface temperature behavior of many types of calorimeters in which the thermometer is not distributed over the surface. I n the present instrument, for KCl dissolutions, use of the observed internal reaction lag would have resulted in errors of 0.07 to 0.12%, but in a calorimeter with a higher thermal leakage modulus, the error would be proportionally greater. Another serious surface-temperature effect was discovered during the course of this work: a superheating of the bomb surface in the vicinity of the heater. A cross section of the heater arrangement is shown in Figure 1. The heater well is of the same material as the bomb, coinage gold, and is welded in an off-axis hole through the hemispherical lower end of the bomb. The heater spool was copper. Use of copper was a design error, since it promotes undesirable conduction of heat to the bore and the outer end of the spool. After discovery of this local superheating, a new heater wound on a stainless steel spool was installed. The heater windings were no. 35 Formvar-insulated Manganin, and the leads were no. 26 copper. The space between the winding and the well from the inner end to the holes at point A (Figure 1) was filled with Apiezon W wax by inverting the bomb section, partially filling the well with wax, and screwing in the spool with the assembly heated to melt the wax. The copper leads had three turns around the spool immersed in the wax and were cemented to the inner surface of the spool from points B to C, and also along the outer surface of the bomb from points E to G or H. After installation of the stainless steel heater spool, the end of the heater well was covered by two layers of 0.002-in. gold foil, cemented around the rim at E except for a small notch to pass the leads. Table I1 gives results of some observations of the degree of superheating at the indicated points as observed with a thermocouple taped or otherwise pressed in contact with the surface. The heater power in all cases was 11 w. Point D for the stainless steel spool refers to the center of the gold foil. Point I refers to the thermocouple attached to the heater leads, 2 cm. from the point of contact with the bomb, points G and H, respectively. To determine the effect of this superheating upon the measurement of heats of reaction,batc h D of HzS04. 20H20 was prepared and eight bulbs of this were weighed on the same day. Bulbs 1,4, and 8 were used for the first series with the copper-spool heater and bulbs 2, 3, 5, and 6 for a second series after installation of the stainless steel spool heater. The first series gave results of 32.592, 32.591, and Volume 69,Number 9
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STUART R. GUNN
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Figure 1. Cross section of heater.
Table 11: Superheating of Surface, "C. Point
B D F G H
I
Copper spool
5.6 2.8 2.2
0.5 0.2 -0.7
Stainleas steel spool
... 0.15 0.2
