Thermothemital I N V ~ S ~ ~ ~ Ufor ~ ~aO N S First-year College Chemistry Course
H. A. Neidig, H. Schneider' and T. G. Teates2
Lebanon Volley College Annville, Pennsylvania
1
A survey of existing literature
The development of a new first-year course in chemistry requires that attention be given not only to the lecture content but also to the laboratory experiments to be used. One of the recent trends in laboratory programs has been to introduce the students to the basic concepts of physical chemistry through a problem-solving experiment involving quantitative techniques. Another practice is to use a series of experiments within a program which involve the sequential development of ideas, concepts, and techniques. The inclusion of thermochemical investigations in the laboratory program of a first-year college chemistry course constitutes a useful way of introducing the students to some of the basic concepts of thermodynamics. Numerous articles appear in the literature that describe experiments involving thermochemistry. Each of these experiments serves its intended purpose well. The effectiveness of these experiments can he increased if several of them are used in a sequential ordei at various places in the over-all laboratory program. A series of vertically developed experiments could be constructed using the major classifications of thermometry; calorimetry; phase changes; enthalpies of dilution, solution, and mixing; enthalpies of reactions; and the relationship between enthalpy, free energy, and entropy. The purpose of this article is to review some of the experiments that appear in the literature involving thermochemistry. Many of these experiments can he modified for use in a first-year course. Thermometry
Neidig and Yingling (1) report an investigation in which the students are given the problem of comparing the thermal energy capacity of several solutions and water. Time-temperature data for the cooling of water and of aqueous solutions of potassium chloride and potassium iodide are used to find the rate of cooling. A comparison of the results can be used to develop the idea of heat capacity. A number of questions are raised that suggest further work such as comparing the thermal energy capacity of solutions of potassium bromide, calcium chloride, and iron(II1) chloride and investigating such liquids as benzene, glycerol, and dichloroacetic acid. 1 Present address: Eau Gallic High School, Eau Gallie, Florida. Present address: Longwood College, FarmviIIe, Virginia.
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An experiment (S) is described in which the students investigate the temperature-changing capacity of a system. The effect of varying the volumes of distilled water and of hydrochloric acid on the change in temperature and on the change in thermal energy is found. This experiment introduces the student to the extrapolation of plotted time-temperature data to obtain the temperature change (3). A discussion of the graphical representation of the data establishes the relationship of heat and temperature. Another approach to presenting thermometry would be to collect data to be used to find the heat capacity of a metal. Such an experiment is described in a number of laboratory manuals of which that given by Plane and Sienko (4) is typical. Blackwell and Fosdick (5) present an experiment in which a temperature change is used to measure the effect which blood catalase has in accelerating the decomposition of hydrogen peroxide. Within certain limits, the degree of acceleration of the reaction is proportional to the catalase activity. Calorimetry Construction o f a Calorimeter
There are numerous articles appearing in the literature describing procedures and equipment that can be used to obtain precise thermochemical data. Skinner,
Sturtevant, and Sunner (6) describe the design and constmction of reaction calorimeters. Sturtevant (7) presents some general remarks on calorimetry including a description of various types of calorimeters. Armstrong (8) describes a number of different types of calorimeters in conjunction with a historical discussion of calorimetry. A number of articles have appeared in Tms JOURNAL describing various types of calorimeters (9). While the previous references describe elaborate calorimeters that would not be available in sufficient quantity for the students in a first-year course, the ice calorimeters described by Mahan (lo), Dunics (If), Kokes (I@, and Vallee (15) could be nsed for such a course. Rogen (14) describes a calorimeter which is a polyethylene bottle placed in a polyethylene block. Although it is customary in a first-year chemistry course to use two beakers with some type of insulation between them as a calorimeter, commercially available Thermokups (styrofoam coffee cups) serve quite well as calorimeters without the use of any insulation.
energy. The differencebetween the calculated electric energy transferred to the system and the calculated thermal energy is the calorimeter constant. An inexpensive heater that can be used for the electrical calibration of a calorimeter is reported by Rogers (14). An experiment suggested by Davidson ($0) is designed for the study of heat flow and would be very useful in presenting the basic ideas and techniques of both thermometry and calorimetry. The reaction chamber of the calorimeter is a 250-ml beaker with a vented top for thermometer and stirrer. Blocks of heat-insulating material of Foamglasl are prepared to fit the beaker. Cooling curves and warming curves are prepared from the data and used to compare the heat flow of insulated and uninsulated beakers. Davidson suggestsadditional experiments that could be done including finding the effect of removing the top of the calorimeter, of stirring, of using various size jackets and beakers, and of using different volumes of water or different liquids in the calorimeter.
Colibrotion of the Colorimeter
A discussion of the theoretical aspects of thermochemical studies, includmg the basis for calculations and of general thermochemical laboratory operations, is given by Pattison, Miller, and Lucasse ($1).
One type of experiment that can be nsed to introduce students to the fundamental techniques of calorimetry is to find the calorimeter constant (water equivalent or heat capacity) of a calorimeter. In one such experiment (Ba), the students discover that the thermal energy available or required for agivenprocess isnot completely transferred to or from a system. Two samples of water of equal volume but each a t a different temperature are mixed. From time-temperature data, the change in temperature is obtained by extrapolation. The heat transferred is compared to the heat available and the difference is the calorimeter constant. The calorimeter constant is defined as that quantity of heat per degree change in temperature (calories per degree) that is lost or gained by the calorimeter assembly and the environment during the calorimetric determination. Dunicz (11) describes an experiment using a vacuum flask calorimeter and the ice fusion method. The data obtained show the relationship between the water equivalent of the calorimeter and the quantity of ice added to 50 ml of water. Livingston (15) gives the difference between the memured and the literature enthalpies for the reaction of sodium hydroxide solution and hydrochloric acid as a means of finding the heat capacity of the calorimeter. Mishcheuko and Kaganovich (16) report the use of the solutionof potassium chlorideinwater to calibrateasolution calorimeter, but Gunn (17) further explored this approach and points out that the process is endothermic. Palmer (18) suggests the dilution of sulfuric acid for finding the calorimeter constant for a calorimeter. Sturtevant (7a) discusses the measurement of electric energy associated with a calorimeter heater circuit suggesting the various apparatus required for precision measurements. Skinner, Sturtevant, and Sunner (6a) consider the calibration of a calorimeter using electric energy and give two schematic diagrams for useful apparatus. While the two previous methods require elaborate equipment, Thompson (19) describes an experiment using simple apparatus. The students are asked to find if the law of conservation of energy is upheld when electric energy is transformed to thermal
Thermochemical Studies
Entholpy of Voporizotion and Condensation
An experiment is presented in which a test tube or a small distilling flask serves as a generator and the vapors are condensed in a simple calorimeter ( t b ) . From the data collected, the enthalpy of condensation is calculated and is assumed to he equal to the enthalpy of vaporization. Britton ( B ) describes an experiment using a copper calorimeter containing a short piece of glass tubing through which the wat,er vapors pass into the room during the determination. From the mass of water lost and the rate of heating during the loss of vapors, it is reported that the enthalpy of vaporization can be calculated. Another experiment is reported by Owen (923) in which steam is passed into a weighed amount of water until the water boils. From the data, the enthalpy of condensation is calculated. Stafford (24) describes an experiment in which the enthalpy of vaporization of iodine is calculated from equilibrium constants. These constants are obtained from vapor pressure measurements. Literature free energy functions and enthalpy changes calculated from heat capacity or spectroscopic data are used. Enfholpy of Melting
A number of laboratory manuals include an experiment on the enthalpy of fusion of ice of which that given by McLellan and Tucker (925) is representative. The calorimeter is a 250-ml beaker placed on corks in a 600-ml beaker using an asbestos sheet as a cover. The enthalpy of fusion of ice is found in an experiment described by Storer ($6) using a polished calorimeter with an immersion heater. When the heater is turned on, a quantity of dried ice is added to a mass of water. The time-temperature data are plotted and from the graph the rate of heating and time of melting are Av&,ble from Armstrong Cork Co. Volume 42, Number I , Jonuory 1965
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obtained. The mass of ice used is found using the mass of the initial and final states. These data are used to calculate the latent heat of fusion of ice. Heat Capacity
A very precise method for obtaining the heat capacity of a liquid is reported by Greene (9). The equipment recommended would not be practical in a first year course, but the procedure given could be followed. Heat is supplitvl to tl;f, liquid h\;pasing elrctric enrrgy into it. Wlwn R k~:nnnquantityof Ilwt isil~ld~d to the liquid, tlie observed temperature change is used to calculate the heat capacity of the liquid. A simple flow calorimeter is described by Pierce (9) that can be used for the direct determination of the heat capacity of gases. Storer (26) suggests the use of two balanced calorimeters, equally insulated and heated to find the heat capacity of a solution. A theoretical consideration of the procedure is given. A calorimeter and procedure are described by Charles (27) in which the reaction of sulfuric acid and sodium hydroxide is used rather than an electrical heater. An interesting investigation could be developed from the experiment presented by Fell ($8) involving the heat capacities and densities of the i-propyl alcohol-water and the t-butyl alcohol-water systems. The procedure involves obtaining the density using the density bottle method and the heat capacity by electrically heating 150-ml of the mixture in a vacuum flask. The data are plotted to show the change in heat capacity and in density as a function of composition. Representative data are given for both systems. Fell suggests that the n-propyl alcohol-water and the acetone-water system can also be used to advantage for this experiment. Kokes, Dorfman, and Mathia (12) give a procedure for obtaining the heat capacities of nickel and copper using an ice calorimeter. With this procedure, a group of 25 randomly selected students obtained for the heat capacity of copper 0.090 =t0.004 cal deg-'g-I as compared to a literature value of 0.0921 cal deg-'g-' and of nickel 0.108 r 0.002 cal deg-'g-I as compared to a literature value of 0.105 cal deg-'g-'. Enthalpy of Dilution
Wolthuis, Leegwater, and Vanderploeg ($9) report an experiment in which the enthalpy of dilution is found for adding sulfuric acid to water a t a number of d i e r e n t concentrations. The calorimeter used was a pint-size Thermos bottle. The students prepare a standardization curve from these data and use the curve to find the composition of an unknown mixture of sulfuric acid and water using enthalpy data. Using two unknown acid mixtures of 63.6 and 82.6% acid composition, the class averages were 64.2 and 82.9%, with maximum deviations of 3.5 and 0.9%, respectively. A discussion of the addition of sulfuric acid to water in terms of enthalpy changes is given by Horton (30). The integral enthalpy of solution is presented in such a way that these enthalpies can beused similar to enthalpy of formations. The difference in the enthalpy changes for mixing different concentrations of sulfuric acid and water is discussed with appropriate equations and enthalpies. Another enthalpy of dilution experiment is suggested by Pattison, Miller, and Lucasse (21). The enthalpy 28
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Journal of Chemical Education
change of the dilution of 95.6% (by mass) ethyl alcohol by six different proportions of water is found. A vacuum Thermos bottle is used as the calorimeter and the alcohol is poured into the water in the calorimeter. Using the data obtained, the enthalpy change is calculated a t the final temperature. A sample calculation, representative data, and a graph of the enthalpy change as a function of temperature are given. The use of volume change and enthalpy changes to consider the structure of aqueous alcohol solutions is presented as an investigation (zc). In addition, volume and thermochemical data are obtained for adding an alcohol to a solution of sodium hydroxide solution and these data are compared with those obtained for the alcohol-water systems. The alcohols suggested for use include methyl, ethyl, n-propyl, i-propyl, and n-butyl alcohol. Enthalpy of Mixing
Zaslow (31) describes an experiment in which the enthalpies of mixing are used to consider the ideality of pairs of liquids and to discuss hydrogen bonding. A Dewar flask was used as a calorimeter. For the systems studied, the observed temperature change was reported to be chloroform-acetone, 13'C, carbon tetrachlorideacetone, -l°C, and chloroform-tetrachloride, -2"C, and the average enthalpy of mixing chloroform and acetone found by the students was -0.95 kcal/mole. Enthalpy of Solution
An excellent discussion of integral heat of solution and of differential heat of solution is given by Daniels, et al. (52). An experiment is described for studying the enthalpy of solution of potassium nitrate. The apparatus to be used is a vacuum bottle containing an electric immersion heater. The procedure is described in detail, and the calculations are discussed for obtaining the euthalpy of solution, the integral heat of solution, and the dierential heat of solution. Furt,herwork that is suggested includes &ding the enthalpy of solution of urea, phenol, and the compound (NH2)C0.2CsHsOH, and calculating the enthalpy of formation of the compound. Another suggestion is to calculate the enthalpy of hydration of calcium chloride from the experimentauy found enthalpy of solution of calcium chloride and of calcium chloride six hydrate. Shoemaker and Garland (Sa) also give an experiment involving the study of the dissolution of potassium nitrate. Palmer (18a) describes an experiment in which the enthalpies of solution of sodium carbonate and of sodium carbonate ten hydrate are found and used to calculate the enthalpy of hydration of sodium carbonate. Representative data is given in Table 1. Table 1 .
Enthalpier of Solution
-AH(saln) Expt.
Na2COa Na.CO.. 10HzO MgS04 MgS04.7Hn0 CuS04 CuS06.5Hg0
-6.9 +15.15 -20.0 +3.46 f15.4 -2.61
(kcd mole-')Lit. -6.6 +15.21 -20.3 +3.8 +15.8 -2.75
The use of sodium acetate and sodium acetate three hydrate is suggested by Pattison, Miller, and Lucasse
(21) for an enthalpy of solution experiment. The quantities of reagents to be used, heat capacities and densities of the solutions, and literature enthalpies are given. Another experiment (2d) involving the enthalpy of solution of a hydrated salt and the corresponding anhydrous salt uses the resulting enthalpies to discuss the degree of interaction in the system. The enthalpy of solution is calculated from data obtained from the dissolution of magnesium sulfate and of magnesium sulfate seven hydrate. In addition, the enthalpy of hydration of magnesium sulfate is calculated. Suggested additional work includes the investigation of copper(I1) sulfate and copper(I1) sulfate five hydrate. Representative data are given in Table 1. Enthalpy of Neutralization
An elementary investigation of a neutralization reaction involves the use of thennochemical data to find the stoichiometry of a reaction (Be). The calorimeter used is a Thermokup. A continuous variation study of the sodium hydroxide-hydrochloric acid-water system provides data that suggest the stoichiometry of the reaction. An experiment is presented in which enthalpies of neutralization are considered in view of the experimentally determined conductivity of a series of acids and bases (55). Thermochemical data are obtained for hydrochloric, nitric, and acetic acids individually with potassium hydroxide and for sulfuric, nitric, acetic, and hydrochloric acids with sodium hydroxide. Livingston (15a) suggests using sodium hydroxide solutions with such substances as sodium hydrogen sulfate and hydrochloric, acetic, monochloroacetic, propionic, or meta boric acids. A special solution calorimeter is recommended. A detailed discussion of the neutralization process is given by Miller, Lowell, and Lucasse (34) including the consideration of enthalpies of dissociation and dilution. The representative data reported for the suggested experiment include those which are given in Table 2. Table 2 .
Arias
Enthalpies of Neutralization
Sodium hydroxide Ammonium hydroxide Avg AH Avg AH (kcal Temp, (kcal Temp, mole-1) T mole-') C
-13.07 -13.01 -12.65 Acetic -12.67 Monochloroacetio -13.73 Oxdic -13.87c -14.03 Tartaric -12.699 -12.e7' -11.21 -13.90
Sulfamic
Range of 27to30 26 29.1 24.6 28.7 25.2 24.3 25.9 21.7 24.7
-11.44 -11.43 -10.94 -11.43 -12.2 -12.01 -12.90 -11.32= -11.29
Flange of 26to29 26.2 29 24.3 27.8 25.1 25.7 24.7
-Use of 0.25 M base and 0.50 M acid.
I n comparing the enthalpies, the differences were attributed to such factors as the temperature coefficient of ionization and the magnitude of heat capacity of the resultimg reaction mixture. The data from this experiment provide a convenient way to discuss many of the factors associated with the neutralization process. Kokes, Dorfrnan, and Mathia (18) present an experiment in which an ice calorimeter is used to obtain data
to calculate enthalpies of neutralization. The enthalpy of neutralization using the stated procedure gave 14.3 =t 0.7 kcal mole-' for sulfuric acid and sodium hydroxide as compared to 14.55 kcal mole-' appearing in the literature and for potassium hydrogen carbonate and sodium hydroxide 9.18 1.2 kcal mole-' was obtained as compared to 9.20 kcal mole-' given in the literature. Palmer (18b) presents an experiment in which a simple apparatus is used to obtain data for calculating the enthalpy of neutralization for sodium hydroxide with phosphoric acid, sodium hydrogen phosphate and potassium dihydrogen phosphate. I n addition, another experiment is given in which the enthalpy change for the reaction of aniline and acetic acid is used to calculate the dissociation constants for acetic acid and for anilime. Enthalpy of Precipitation
The reaction of aluminum sulfate or of magnesium sulfate with sodium hydroxide is reported by Pattison, Miller, and Lucasse (21) as an experiment from which data can be obtained to calculate enthalpy of precipitation. The enthalpy of precipitation of magnesium hydroxide is reported as 0.2 kcal mole-' and of aluminum hydroxide as -31.9 kcal mole-'. Clever (55) presents an experiment in which the enthalpies of precipitation of silver chloride and of silver bromide are related to the enthalpies of solution, which in turn are related to the hydration and crystal lattice energies. Other topics that are suggested for discussion include the contributions of Coulombic attractions, London attraction, and Born repulsion to crystal lattice energy, and the mechanism of solution. The article describes the constmction of the calorimeter using a 250-ml Erlenmeyer flask and gives a procedure for the calibration of the calorimeter. The represent* tive data obtained using the described procedure are given in Table 3. Daniels et al. (Sfa) suggest an experiment in which the student is given the enthalpy of formation of lead(I1) iodide and asked to design a suitable procedure to be used to find the enthalpy of formation of Pb2+(aq)and 21-(aq) at infinite dilution in water. Table 3.
Enthalpy of Precipitation
Average AH (kcal mole-') Calculated AH (kcd mole-')
Silver chloride
Silver bromide
-15.30 =t 0 . 7 0 15.65
-20.03 *0.75 -20.19
Enthalpy of lonizotion
Shoemaker (Sb) presents an experiment in which the enthalpy of ionization of water and of the second ionization of malonic acid are found. Using a Dewar flask equipped with a heating unit as a calorimeter, data are collected for the reaction of sodium hydroxide and hydrochloric acid and of sodium hydrogen malonate and sodium hydroxide. Directions are given for the calculations. Entholpy of Reaction
Three reactions are suggested by Pattison, Miller, and Volume 42, Number
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Lucasse ($1). Time-temperature data are obtained and used to calculate enthalpy changes for the decomposition of hydrogen peroxide in the presence of manganese(1V) oxide, for the reaction of calcium carbide with hydrochloric acid, and for the reaction of potassium bromate(V) and hydrobromic acid in excess hydrochloric acid. These data are given in Table 4. The reactions of sodium chloride with calcium nitrate, with barium nitrate, and with mercury(I1) nitrate are presented by Livingston (15b) as an experiment for studying the enthalpy of ionic reactions. An experiment is described by Daniels et a1 (SBb) in which the enthalpy of reaction a t infmite dilution is calculated for the reaction of mercury(I1) nitrate and sodium chloride. The theoretical aspects of the experiment are discussed in detail. Two organic reactions are given by Pattison, Miller, and Lucasse ($1). Following the procedure given, the enthalpy change for the reaction of acetic anhydride with sodium hydroxide is given in Table 4. A proce dure is also given for the reaction of acetone with hydroxylamine hydrochloride. Modifications of the procedure are reported to alter the enthalpy change appreciably. Table 4.
Enthalpies of Reaction
System
---AH Expt
&O,(MnOd HCI CaC. KBrO. HBr(HC1) (CHaCO)20 NaOH
-23.2 -58.7 -49.7 -38.5
++ +
(kcill mole-*)----. Lit -22.8 -60.8 -51.2to - 5 3 1 -43.7
The enthalpy changes are found for the reactions of the alkaline earth oxides with water and with hydrochloric acid (Sf). The enthalpy data obtained are discussed in terms of the periodicity of the alkalme earth metals. Entholpy of Displacement
Charlesworth and Patch (56) describe an experiment involving the reaction of zinc and copper(I1) sulfate using a Thermos bottle as a calorimeter. The experimental enthalpy change for the reaction is reported as 49.05 kcal mole-' compared to a literature enthalpy of 50.11 kcal mole-'. An experiment is described by Hilbourne (57) in which the activity of magnesium, zinc, and iron is related to the enthalpy changes for the reactionof these metals with hydrochloric acid. Mahan (10) gives a procedure using an ice calorimeter and calculations for finding the enthalpy change for the reaction of magnesium and hydrochloric acid. Entholpy of Complex Ion Formation
Two reactions are suggested by Pattison, Miller, and Lucasse ($1) for studying the enthalpy of complex ions. The enthalpy change for the reaction of mercury(I1) nitrate and sodium chloride is reported as -12.7 kcal mole-' compared to the literature enthalpies which vary from -11.2 to -12.7 kcal mole-'. For the reaction of copper(I1) sulfate with excess ammonia, the reported euthalpy change is -19.3 kcal mole-' compared to the
literature enthaluy of - 19.5 kcal mole-' for the formation of ~ u ~ 0 4 . 5 ~ ~ ~ . Daniels et a1 (S2c) suggest the possibility of investigating the reaction of njckel(I1) ion with a large excess of ammonia to obtain data for calculatine the enthal~vof formationof acomplex ion fromitsconsiituents. ?Lgling, Neidig, and Teates (58) suggest a thermochemical study of the mercury(I1) nitrate-potassium iodide-water system using a continuous variation experimental design to obtain data that can be used to consider the stoichiometry of the various species present at different compositions. Enthalpy of Formation
An investigation is described (2g) in which the enthalpy of formation of solid ammonium chloride is calculated using Hess' Law. Literature enthalpies are used for the formation of aqueous ammonia and of hydrochloric acid from the elements. Enthalpy changes are calculated from experimental data for the dissolution of solid ammonium chloride and for the reaction of ammonia water with hydrochloric acid. Mahan (10) describes an experiment in which the enthalpy of formation of magnesium oxide is calculated using Hess' Law. Using an ice calorimeter, data are obtained and used to calculate the enthalpy change for the reaction of magnesium and hydrochloric acid and of magnesium oxide and hydrochloric acid. From these experimental enthalpy changes and those for the formation of magnesium and of water from the elements (as found in the literature), the enthalpy of formation of magnesium oxide is calculated. A test of Hess' Law is the indicated object of an experiment reported by Matthews ($9). Data are obtained to calculate the enthalpy change for the reaction of lithium and water and of lithium hydroxide and hydrochloric acid. These enthalpy changes are used to calculate the change for the reaction of lithium and hydrochloric acid, which is reported to be 59.5 kcal mole-' (using the suggested procedure) compared to a literature enthalpy change of 59.0 lccal mole-'. The experiment could be modified so that the enthalpy of formation of lithium chloride could be calculated. Anexperiment (SSa) is presented in which the enthalpy change is calculated from experimental data for the dissolution of solid sodium hydroxide in water, the reaction of solid sodium hydroxide with hydrochloric acid, and the reaction of a solution of sodium hydroxide with hydrochloric acid. The enthalpy changes for the three reactions are compared and examined in t e r m of Hess' Law. This experiment could be altered to involve the calculations of the enthalpy of formation of sodium chloride. Relationship of Enthalpy, Free Energy, and Entropy
An experiment (2h) is givenin whichdataare obtaiued using a The~mokupas a calorimeter from which the enthalpy change for the reaction of zinc and silver nitrate is found. From the measured electric potential diiereuce of the zinc-zinc nitrate/silver-silver nitrate electrochemical cell, the free energy change is calculated. The entropy change for the system is found using the calculated enthalpy and free energy changes. A study
is also suggested for the zinc-copper(I1) sulfate-wa,ter system. The enthalpy change is calculated from experimental data for the reaction of zinc and copper(I1) sulfate and the free energy change from data for the zinczinc sulfate/copper-copper(I1) sulfate electrochemical cell. From the enthalpy and free energy change, the entropy change for the system is calculated. The experimental and literature data for these two systems are given in Table 5. Table 5.
Enthalpy, Free Energy, and Entropy Changes AH
Svstem
(kc61molecL) Exnt Lit
AG (kcalmole-I Exnt Lit
AS (cal mole-'deg-') E x ~ t Lit
A second experiment (29 is given in which the enthalpy changes are calculated, from data obtained by using Thermokups as calorimeters, for the reaction of sodium hydroxide with each of a series of acids (acetic, chloroacetic, dichloroacetic, and trichloroacetic acids). These enthalpy changes are corrected for enthalpies of neutralization, giving the euthalpies of ionization for the acids. From the half-titration of each acid, the enthalpies of ionization are calculated for the acids. By using the enthalpy of ionization and the free energy of ionization, the enthalpy of ionization is calculated. The thermodynamic functions of the acids are compared and used to explain the observed difference in the degree of association of the acids. The representative data reported include that presented in Table 6. Table 6. Experimental Thermodynamic Constants for the Ionization of Acetic and Chloroacetic Acids AHo
Acids
(kcal mole-')
AGO (kcd mole-')
AS'
( 4 mole-' dee-1)
Literature Cited
( 1 ) NEIDIG, H. A,, AND YINGLING, R. T.,Chemi817y, 37 ( I ) , 29 (1964). ( 2 ) "Investigating Chemical Systems," Teachers Guide, Wehster Division, McGraw-Hill Book Company, Inc., New York, 1963, p. 8-1; ($a) p. 9-1; ( b )p. 21-1; (&) p. 45-1; (Zd) p. 27-1; (Se) p. 36-1; (Bj) p. 2 4 1 ; ( B g ) p. 22-1; (M) p. 29-1; and (Zi) p. 32.1. D. P., AND GARLAND, C. W., "Experiments in ( 3 ) SHOEMAKER, Physical Chemistry," McGraw-Hill Book Company, h e . , New York, 1962, p. 110-111; (Sa) 122; and (Sb) p. 117. ( 4 ) PLANE,R. A,, AND SIENKO,M. J., "Experimentd Chemistry," 2nd ed., McGraw-Hill Book Company, Ine., 1961, p. 55.
I BLACKWELL. L.. THIS JOURNAL. . R. &., . . AND FOSDICK, . 32,588 .
(1955). SKINNER,H. A,, "Experimental Thermochemistry," Vol. 2, Interscience Publishers, John Wiley and Sons, New York, 1962, chap. 9; and (6a)p. 177. 8 WEISSBEROER, A., "Physicd Methods of Organic Chemistry," Vol. 1, Part 1, 3rd ed., Interscience Publishers, John Wiley and Sans, New York, 1959, p. 53552; and (7a) p. 530. ARMSTRONG, G. T., THIS JOURNAL, 41,297 (1964). PIERCE,P. E., THIS JOURNAL, 39, 338 (1962); HORNYAK, F. M., ibid., 38, 97 (1961); O'HARA,W. F., Wu, C. H., AND HEPLER,L. G., ibid., 38, 512 (1961); ZUTTY,N. L., AND HERBRANDSON, H. F., ibid., 35,260 (1958); MOWERY, JR.,D. F., ibid., 34, 244 (1957); SLABAUGH, W. H., ibid., 33, 519 (1956); GREENE,S. A,, ibd., 32, 577 (1955); C H ~ L ER.S G., ibid., 31, 577 (1954); CAMERON, I., AND WRIGHT,R. H., ibid., 18, 510 (1941); LIVINGSTON, R., AND HORwITz. W.. ibid.. 16. 287 (1939): . . and V O N KLOOSTER, H. s., ibid., 12; 285(1936).' MA~AN B., H., THIS JOURNAL, 37, 634 (1960). DUNICZ, B. L., THIS JOURNAL, 37, 635 (1960). KOKES,R. J., DORFMAN, M. K., AND MATHIA,T., THIS JOURNAL, 39,90 (1962). VALLEE,R. E., Rev. Sei. Instr., 33,856-8 (1962). ROGERS, M. J. W., Seh. Sc. Rev., 45, 142 (1963). LIVINGSTON. R.. "Phvsica-Chemical Emeriments." Mac, p. i27; ( 4 p. 123; millan om piny, N ~ W~ o r k 1939, and (15b) p. 118. MISCHENKO, K. P., AND KAGANOVICH, Y. Y., Zhur. Priklad. Khim., 22, 1078 (1949). GUNN,S. R., Rev. Sei. Ins*., 29, 377 (1958). PALMER,W. G., "Experimentd Physical Chemistry," Cambridge University Press, London, England, 1954, p. 159; (18a)p. 170; (186) p. 164. THOMPSON, D., Cmnecticut Science Teachers Association Newsletter, June, 1962, p. 10. DAVIDsoN, D. L., THIS JOURNAL, 22, 38 (1945). PAWSON,D. B., MILLER,J. G., AND LUCASSE, W. W., THIS JOURNAL, 20,319 (1943). BRITTON, K. G., Seh. Se. Rev., 28, (105), 241 (1947). OWEN,B., The Science Teacher, 21, 153 (1954). STAPFORD, F . E., THIS JOURNAL, 40,249 (1963). MCLELLAN. C. R.. A N D- TUCKER.W. C.. "Exuerimental en& 'chemistry," Scott, Fbresman' and 'Campany, Faidawn, N. J., 1961, p. 23. STORER, J. A,, Sch. Se. Rev., 32, (116), 104 (1950). CHARLES, R. G., THIS JOURNAL, 31, 577 (1954). FELL,J. I., Seh. Se. Reu., 37, (131), 141 (1955). WOLTHUIS, E.. LEEGWATER, A., AND VANDERPLOEG, J.9 THIS JOURNAL, 38, 472 (1961). HORTON, W. S., THIS JOURNAL, 23, 393 (1946). ZASLOw, B., THIS ~ O U R N A L ,37, 578 (1960). DANIELS,F., ET AL., "Experimentsl Phy~icalChemistry," McGraw-Hill Book Company, 6th ed., 1962, p. 22; (%a) p. 38; (SZb) p. 29; and (SZc)p. 38. "Chemistry," W. H. Freeman and Co., 1963, p. 48; and (SSa)p. 39. MILLER,J. G., LOWELL, A. I., AND LUCASSE, W. W., THIS JOURNAL, 24, 121 (1947). C L E ~ RH., L.,THIS JOURNAL, 38, 470 (1961). CHARLESWORTH, bf. E., AND PATCH,E. M., Sch. SC.Rw., 13, 256 (1932). HILBOORNE, E. W., THIS JOURNAL, 5, 733 (1928). YINGLIN~R. T., NEIDIQ, H. A,, AND TEATES,T. G.3 Chemistry, 37 ( 4 ) , 24 (1964). MA-EWS, P. G., Seh. Sc. Reu., 45, 194 (1963). 1
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