Two Multipurpose Thermochemical Experiments for General Chemistry R. A. D. Wentworth Indiana University, Bloomington, IN 47401 Thermochemistry is an important topic in general chemistry because i t distinguishes heat from temperature, leads to a discussion of bond-energy concepts, and provides an early eroundwork for suhseauent treatment of thermodvnamics. From a student's poini of view, thermochemistry fs an imnortant tooic because it concerns a kind of enerev that the student has experienced. If an exoeriment accom~aniesthe lectures on thermochemistry,;t deserves an emphasis that reflects the importance of this topic. This emphasis can he aenerated without studying a special or arcane-reaction. In fact, any reaction is a fit subject as long as the reagents are cheap and sufficient heat is evolved or absorbed to he measured with uusophisticated equipment. A survey of some simple hut suitable reactions was ouhlished in 1965 (1). The des>gnof the experiment is probably more important than the type of reaction. One way to achieve the desired emphasis is through a multipurpose experiment. This kind of exoeriment shows that thermochemistrv is not an isolated topi; hut one that has relevance to many other areas of chemistrv. Thermochemical measurements can he used to arrive atconclusions about any other topic that you choose. Moreover, this kind of experiment is challenaing because it provideg,&mething extra for a student to after the calculations have been completed. Two multipurpose ther.mochemical experiments that have been used at Indianauniversity are described in this oaner. One of these exneriments deals with neutralization %le the other deals 'with ligation. Each experiment requires twonested, polystyrene coffee cups, a suitable lid, and a thermometer graduated in degress. A water-impelled magnetic stirrer is recommended, but manual stirring- will suffice. Both of these experiments make use of a tactic that is often used a t Indiana University when students are doing an experiment in which only one numerical result is possible. It is almost inevitable that a student will be aware of the data obtained by neighbors in the laboratory. Rather than trying to avoid problems that may arise from this knowledge, it seems better to require students to know their neighbor's data and use it in their calculations. Thus. each student uses the data from the entire class, rejects suspicious data with an aunrooriate test (2) if desired. and calculates the mean value: ~ d d i t i o n a s&tistical l ankysis can also be used if the students are able to cope with it.
Thermochemlcal Data lor the Neutralization Experimenta
Reaction
AH1k.I mnl-$1
--
1022
Journal of Chemical Education
The Neutrallzatlon Experiment Although this experiment begins with a traditional exothermic neutralization reaction, an endothermic process is also a focal point. Problems involving stoichiometry and Hess'slaw are an integral part of the experiment and provide a challenge to most students. However, the most challenging aspect of the experiment may come a t its end when students attempt to understand and explain thermochemical phenomena associated with strong acids and strong hases. Students are divided into two groups. Each student in a eroun. does the same exneriment and shares his or her ex.. perimental data with everyone else in the gmup. The memhersof the other group do a similar but difterent ex~eriment and share the data wkhin that group. Finally, all ofthe data is pooled. Each student in the first group measures the heat accompanying the neutralization of 50 mL of 2.0 M HN03 with an identical volume of 2.0 M NaOH. A typical mean heat of neutralization is -59 kJ mol-1 and the maximum deviation from the mean is usually no more than i3 kJ mol-I. The student must then calculate the mass of NaN03 that was formed durine neutralization. add that amount of this salt to 100 mL of water, and meas&e the heat of solution. A typical mean value for this endothermic orocess is 20 kJ &' with a maximum deviation from the mean of i 4 kJ mol-I.
The mean heat of solution from the exnerimental data is then compared t o a value obtained from Hess's law. This calculation reouires the ex~erimentalmean heat of neutralization and the heats from the first six reactions in the table. The result of this nontrivial calculation is 21 kJ mol-' if the typical mean value is used as the heat of neutralization. The other group does a similar experiment using 2.0 M HC1,Z.O M KOH, and solid KCI. The typical mean heat of neutralization is identical within experimental error to the heat of neutralization measured by the other group. A typical mean heat of solution for KC1 is 16 kJ mol-I with a maximum deviation from the mean of *4 kJ mol-I. Hess's law is important once again. The heat of solution for KC1 is calculated from the mean heat of neutralization and the heats from the appropriate reactions in Table 1. A value of 14 kJ mol-I can be obtained by using the typical mean heat of neutralization as well as the heats from the table's first and last five reactions. After all of the data is pooled, the students can compare the heats of the two neutralization reactions. They are usually ahle to see without difficulty that the heats are prohahly identical if they account for experimental error. However, it is not always an easy task for them to arrive at a reason for this unexpected (to them) result. If thermochemistry is discussed before concents of strone acids and strone bases have received attention A d before &e use of net io& equations is stronelv inerained. this facet of the experiment can be particuiaily clklengkg. The Llgatlon Experlrnent
Although this experiment has been described previously (3), the f&owing comments should explain its st&egy and perhaps set the stage for others t o develop additional asThe purpose is to determine the maximum number of ethylenediamine ligands that will hind to an aqueous Cu(I1) ion in a reasonably dilute solution. The correct answer is two. althoueh three will hind in more concentrated solutions (4). since tKe student will gain an appreciation for inorganic stereochemistry during the experiment, it can he used during lectures on coordination chemistry. Each student prepares for the experiment by reading about the normal octahedral stereochemistry of Ni(I1) complexes and the elongated octahedral stereochemistry of sixcoordinate Cu(I1) complexes. The elongation is attributed to an idiosyncrasy of Cu(I1) without further elaboration. The Jahn-Teller effect is neither named nor explained. The students are told that the distance between the nitroeen atoms of the bidentate lieand is fixed because the ~ X - C H ~backbone canno; expand or contract. They understand that onlv an experiment will tell them if ethvlenediamine is capable bf spanning three of the 12 edges bf the elongated octahedron with equal ease.
Each student understands that four trials mav he necessary. The first trial consists of measuring the heat evolved or absorbed when 50 mL of 0.50 M CuSOd and 5 mL of 5.0 M ethylenediamine are mixed (mole ratio = 1:l). The volume of the ethylenediamine solution is changed to 10 mL in the second trial (mole ratio = 2:1), 15 mL in the third trial (mole ratio = 311, and, if necessary, 20 mL in the fourth trial (mole ratio = 45). The experiment continuesuntil the heat accompanying one trial is found to be very similar to the heat accompanying the preceding trial. The success of the experiment is hased on the large overall formation constants for the two complex ions (6): Cu(H,0),2' t e n
-
C~(en)@I,0)~~' + 2H20
K, = 3.5 X 101°
where en is ethylenediamine. These equilibrium constants dictate the mole fractions of C U ( H ~ O ) ~C~~+(,e n ) ( H ~ 0 ) ~ ~ + , and Cu(en)~(H20)2~+ under any set of experimental conditions. When the mole ratio is 1:1, the mole fractions are 0.131,0.738, and 0.131, respectively. More importantly, they are
