H. Lawrence Clever
Heat of Precipitation
Ernon/ University Atlanta, Georgia
A general chemistry experiment
M a n y freshman textbooks discuss the solutiou process of a solid electrolyte in terms of the ious in the crystal lattice interacting and hydrating with the water molecules and going into solution. A few textbooks give t,hemore quantitative relationship
Heats of solution are difficultto measure but for slightly soluble salts their converse, the heat of precipitation, can be easily measured.' Determining the heat of precipitat,ion of two silver halides, and discussing the results, can be a simple but meaningful experiment for freshman chemistry. The heat of precipitation experiment is carried out in a simple calorimeter constructed by each student from an Erlenmeyer flask, a beaker, and a -6 to 100°C thermometer graduated in O.l°C. Considering the simple apparatus used, the results are quite good. The Experiment Construction of Calorimeter. Place a 250-ml Erlenmeyer flask in a. 500-ml beaker. Pack the space between flask and beaker with wadded-up paper towels and wrap the heaker in a towel. Bore a hole for the thermometer in a rubber stopper .. that tightly fits the Bask. Heat Capacity of the Calmimelm. Pipet 100 ml of 0.25 M HCI into the Erlenmever flask. Put the thermometer in d a c e and cheek the solution temperature until a t two minute intervals the readings do not differ by 0.05". From a graduated cylinder add 1 1 ml of 2.5 M NitOH solution. Replace the thermometer, read the temperahre every 30 seconds and record the highest temperature that develops. h a t m e that the heat of neutralization is 13,630calories/mole. Calculate the heat lilmated in the complete reaction of 100 ml of 0.25 M HCI Answne the solution weighs 111 grams and has a heat capacity of 1 calorielgram. Calculate the calories absorbed by the solution and by the calorimeter per degree temperature rise. Heat of Preripilation. Repeat the above procedure, first with 100 ml of 0.25 M AgNOaand 11 ml of 2.5 M NsCI, then with 100 ml. of 0.25 M AgNO. and 11 ml of 2.5 M NaBr solution. The temperature rise will be in the range of 34' C. Nbte: All solut,ions are prepakd well ahead of time and stored in the laboratory so they will be a t room temperature. The author will supply interested readers with a. sheet indiesting to the student a form for reporting results and outlining ralResults and Discussion
Results of the first group to try this experiment are summarized in Table 1. The experiment has been used several times in somewhat larger classes with Presented before the Chemical Education Section, Southesstern Regional ACS Meeting, Birmingham, Alabama, November 1 o'm L""". ' GURNEY,R. W.,''Ionic h e e 8 8 e s in Sohtion," MoGraw-Hill Book Co., New York, 1953,p. 93.
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equally good results. Agreement of the class average value with the accepted value of the heat of precipitation is good. The experimental results are discussed with the class. The heats of precipitation are arranged on the board in descending order. The average and the average deviation are calculated and compared with the accepted value as in Table 1. Students whose values come near the accepted value or a t least fall within the range of the average deviation take considerable pride in their results. The poorer results are used to start an error discussion. Often carelessness in reading a volume or temperature or some simple arithmetic error is discovered in this discussion. The student's attention is directed to a critical evaluation of other errors. The facts that the density of the solutions is not exactly 1, that the solution heat capacities are not exactly 1, and that there is a heat-of-dilution effect are pointed out and di~cussed. Calorimetry in general can be discussed and more elaborate equipment and techniques can be described or even demoustrated if equipment is available. Table 1.
Heat of Precipitation. (cal/mole)
(Results of 12 groups . . of 2 students each, arranged descending order.)
in
Heat of Precipit&ma Calories/Mole Silver Silver chloride bromide
-15.300*700 - 15,65OG
-20.030
=t750
-20,190°
Average Literature
Negative sign in thermodynamic sense of heat liberated. Most freshman texts give heat liberated apositive sign. a Omitted from averages. Calchlated from data. in National Bureau of Standards Circular 500. AHi [AgS (c)] - AHl (Agt, hyp 1 M ) - AH, (X; hyp 1 M)
The results are further discussed as an exercise in t,hermochemistry. The fact that reversing the reaction reverses the sign of the heat effect is emphasized. Thus heat of solution = -heat of precipitation. Ag+
+ CI-
=
AgCl
AgCl = Ag+
heat of preeip. -15,650 caI/mole
+ Cl-
heat of soh.
15,650 cal/mole
The student is given values of silver halide crystal lattice energies3 and asked to combine them with his
experimental results to calculate which ion (Cl- or Br-) has the larger hydration energy. Alternatively the student is sometimes given the hydration energies of the halide ions and asked to estimate which salt has the higher crystal lattice energy. Related topics can be discussed with freshmen on very qualitative terms. For example, the contributions of Coulombic attractions, London attractions, and Born repulsions3, to the crystal lattice energy can be mentioned. Also the solubility mechanism of crystal vaporized to gaseous ions, then ions hydrated to form aqueous ions, leads directly to a discussion of the influence of solvent dielectric constant and the solvolysis reaction on solubility. Discussing all the points outlined above would take much more time than most instructors could allow for one experiment. Each instructor can use his own discretion as to what subjects are important for discussion with his class. A brief discussion of errors and applications to thermochemical calculations would seem a good minimum of discussion for this experiment.
2 LEUSSINR, D. L., in "Treatise on Analytical Chemistry," edited by KOLTHOPP, I. M., AND ELVING, P. J., Interecience Encyclopedia, h e . , New York, 1959. Part I, vol. 1, Chap. 17, p. 70if.
8 KETBLAAB, J. A. A., "Chemical Constitution," 2nd ed., Elsevier Publishing Co., New York, 1958, Chap. 2. G o u ~ u ,E. S., "Inorganic Reaction and Structure," Henry Halt and Co., New York, 1955, p. 168 and Chap. 12.
Addition of reactions and associated heats to give the heat of a third reaction is demonstrated. The relative solubility of AgCl and AgBr is given and the student is asked to calculate the heat of the reaction: AgCl
+ Br-
=
AgBr
+ CI-
which is the difference AHnwnesr- Ah',.tr.cl = -4540 cal/mole. The results are also discussed in terms of the solubility process of vaporizing the ionic crystal to gaseous ions (crystal lattice energy) and hydrating the gaseous ions (hydration energy).2 Mecs, + X-(., MXc crystal lattice energy, U + energy of hydration X-i.1
+ Hz0
--
Q
X-(w
Thus the heat of solution is the difference of two large negative heat effects: AH..I.. = ZAHM
-U
Volume 38, Number
9,
Sepfember 1961
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471
