R. J. Kokes, M. K. Dorfman,
and 1. Mathia The Johns Hopkins University Baltimore, Maryland
Experiments for general chemistry 111
Calorimetry
O n e of the experiments performed in the general chemistry laboratory a t this university involves the use of an ice calorimeter with which the students measure the heat capacities of nickel and copper and the heats of neutralization of both a strong and a weak acid by sodium hydroxide. Many calorimeters designed for student use have been reported in the All of these are more complicated in design than the device described below, and most of them are capable of a greater accuracy than is necessary for a general chemistry experiment. The apparatus used in the freshman course is shown in Figures l(a) and ( b ) . The reaction vessel, which has a total volume of about 50 ml, consists of a tube with a brass lower portion cemented to a glass top. This tube extends through one of the holes in a rubber stopper into a small Dewar flask which serves as the calorimeter. A piece of capillary-bore glass tubing 50mm long projects through the other hole in the stopper into the Dewar, which is filled with ice and water in such a way that the water level is a t the top of this tubing. When the calorimeter is assembled, the stopper is pushed down to make a tight seal a t the mouth of the flask. The flask is then set upright in a plastic wastebasket filled with crushed ice to further insulate it from the surroundings. Prior to the run, a shell of air-free ice is formed around the calorimeter by the addition of a small amount of dry ice. After about 15 minutes water is substituted for the remaining dry ice to melt any ice formed inside the reaction vessel. When this water is removed, the calorimeter is ready for a calibration and reaction run, and these can be carried out without renewing the shell of ice around the reaction vessel. The principle of operation of the calorimeter is simple. The density of ice is lower than that of water, and as the heat liberated by an exothermic process in the reaction vessel melts the ice in the vacuum flask, the level of water in the capillary tubing drops until equilibrium is reached in the calorimeter. The change in The new freshman laboratory course at Johns Hopkina has been described in THIS JOURNAL, 38. 16 (1962). Many of the experiments are innovations in an introductory course. This series of articles describes the experimental procedures in some detail. ' MOWERY, D. F.,JR.,J. CHEM.EDUC.,34,244 (1957). 'SLARAUGH, W. H., J. CHEM.EDUC.,33,519-20 (1956). J P a ~ ~ s D. o ~B., , MILLER,J. G., AND LUCASSE,W. W., J. CHEM.EDUC.,20,319 (1943). I., AND WRIGAT,R. H., J. CAEM.EDUC., 16, 'CAMERON, 510-12 (1941). . . ' L I ~ ~ N GR., ~ ~AND N , HORWITZ, W., J. CHEM.EDUC., 16, 287-90 (1939).
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water level is read from a scale fixed to the capillary tubing, and the readings are translated into calories from a calibration of the apparatus. With this type of calorimeter, a steady slow drop in water level due to heat leak from the surroundings is observed. Therefore, scale readings are taken a t regular intervals during the experiment and the data are represented graphically. In this plot, the gentle, parallel initial and final slopes represent heat leak; the distance between them is a measure of the heat given off in the reaction vessel.
Figwe 1.
(a) Student calorimeter;
(b) calorimeter components,
The calibration of the calorimeter, which is done immediately before each experiment, involves adding 10 ml of water a t room temperature to the reaction vessel and noting the change in height in the capillary which accompanies the known temperature change of the water added. Room temperature is measured by a thermometer with a 0.1 degree scale; the heat capacity of water is given in the laboratory instructions. I n a typical run, a calibration factor of 2.77 cal/mm of scale was obtained. The Heat Capacities of Nickel and Copper In the first calorimetry experiment, the students measure the heat capacity of a metal, either nickel or copper shot. About 100 g of shot a t room temperature is used. After taking scale readings for a few minutes, students rapidly add the shot to the reaction vessel and continue to record the change in water level for about 30 min. From the known weight of metal, the known temperature change, the calibration factor of the calorimeter, and the observed drop in water level, they calculate the specific heat of the metal. Figure 2 shows a typical cooling curve, obtained using copper shot. Results reported by a randomly selected group of 25 students are given in Table 1, together with the heat capacity values from the literature.
Table 1.
-
The Average Results of 25 Students in the Calarimetrv , Exaeriments ~~
Quantity
C* of copper, cal deg-I g-I C. of nickel, cal deg-I g-I
AHw of HISOl by NaOH, kcal
mole-'
AHN of KHCOI by NaOH, kcal mole-1
Students 0.090*0.004 0.108 i 0 . 0 0 2 14.3
Literature 0.09218 0.105'
&0.7
14.55
9.18 i 1 . 2
9.20
The Heats o f Neutralization of Acids
I n the second experiment with the calorimeter, the students find the heat of neutralization of 1 M sulfuric acid by 1 M sodium hydroxide and of 1 M potassium hydrogen carbonate by 1 M sodium hydroxide. The procedure is the same in both cases. They place a 20milliliter sample of 1 M NaOH in the reaction vessel, allow it to cool to O°C, and then readjust the water
level in the capillary tubmg. They next add 10 ml of 1 M HzSOI (or KHC08), which has been chilled to 0°, and record the depression in water level, which in this case is a measure of both the heat gained during the transfer of the acid to the calorimeter, and the heat of neutralization. A correction is made for the heat of transfer by subtracting the change observed when 10 ml of water initially a t 0°C are transferred to 20 ml of water in the reaction vessel; this correction is accurate only if the students are consistent in the way they perform both runs. Results obtained by the students in this experiment are also included in Table 1. The values reported in the literature for the heats of neutralization were obtained a t 25'C. In order to compare them with the students' results, obtained a t O°C, a calculated correction was made with the help of heat capacity data found in the literat~re.'.~For AHHN of HzSOa the correction was $0.75 kcal mole-'; for A?& of KHC03 it was +0.50 kcal mole-'. These corrections have been added to the values in the table. The sample of student results in the table shows a precision of 2 4 % in the heat capacity measurements, and a deviation of about the same magnitude from values taken from the literature. The values reported for the heats of neutralization of sulfuric acid and bicarbonate ion have deviations of 5 and 13y0 respectively; the agreement of the students' average values with values taken from the literature is surprisingly good. "Handbook of Chemistry and Physics," 25th ed., Chemical Rubber Publishing Company, Cleveland, Ohio, 1941-42, p. 1668. PAUL,M. A., "Principles of Chemical Thermodynamics," 1st ed., McGrew-Hill Book Co., New York, 1951, p. 189. Footnote 7 and ChemicalA bsiracts, 49,4383f (1955).
Volume 39, Number 2, February 1962
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