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M a n y freshman textbooks discuss thesolutiou process of a
solid electrolyte in terms of theious in the crystal lattice
interacting and hydrating withthe water molecules and going into
solution. A fewtextbooks give t,he more quantitative
relationship
H. Lawrence CleverErnon/ University
Atlanta, Georgia
Heats of solution are difficult o measure but for
slightlysoluble salts their converse, the heat of precipitation,can
be easily measured.' Determining the heat ofprecipitat,ion of two
silver halides, and discussing theresults, can be a simple but
meaningful experiment forfreshman chemistry.The heat of
precipitation experiment is carried outin a simple calorimeter
constructed by each studentfrom an Erlenmeyer flask, a beaker, and
a -6 to 100Cthermometer graduated in O.lC. Considering thesimple
apparatus used, the results are quite good.
Heat of PrecipitationA ge ne ra l c he mi s t ry e x pe r i me
nt
The ExperimentConstruction of Calorimeter. Place a 250-ml
Erlenmeyer flaskin a. 500-ml beaker. Pack the space between flask
and beakerwith wadded-up paper towels and wrap the heaker in a
towel.Bore a hole for the thermometer in a rubber stopper that
tightlyfits th e Bask.Heat Capacity of the Calmimelm. Pipet 100 ml
of 0.25 MHCI into the Erlenmever flask. Pu t the thermometer in
daceand cheek the solution temperature until a t two minute
intervalsthe readings do not di ffer by 0.05". From a graduated
cylinderadd 1 1 ml of 2.5 M NitOH solution. Replace the
thermometer,read the temperahre every 30 seconds and record the
highesttemperature that develops.h a t m e tha t the heat of
neutralization is 13,630 alories/mole.Calculate the heat lilmated
in the complete reaction of 100 mlof 0.25M HCI Answne the solution
weighs 111 grams and hasa heat capacity of 1 calorielgram.
Calculate the caloriesabsorbed by th e solution and by t he
calorimeter per degree tem-perature rise.Heat of Preripi lat ion.
Repeat the above procedure, first with100 ml of 0.25M AgNOa and 11
ml of 2.5M NsCI, then with 100ml. of 0.25 M AgNO. and 11 ml of 2.5
M NaBr solution. Thetemperature rise will be in the range of 34'
C.
equally good results. Agreement of the class averagevalue with
the accepted value of the heat of precipita-tion is good.The
experimental results are discussed with the class.The heats of
precipitation are arranged on th e boardin descending order. The
average and the averagedeviation are calculated and compared with
the ac-cepted value as in Table 1. Students whose valuescome near
the accepted value or a t least fall within therange of the average
deviation take considerable pridein their results. The poorer
results are used to startan error discussion. Often carelessness in
reading avolume or temperature or some simple arithmetic erroris
discovered in this discussion. The student's atten-tion is directed
to a critical evaluation of other errors.The facts tha t the
density of the solutions is not ex-actly 1, that the solution heat
capacities are not exactly1, and that there is a heat-of-dilution
effect are pointedout and di~cussed. Calorimetry in general can be
dis-cussed and more elaborate equipment and techniquescan be
described or even demoustrated if equipment isavailable.
Table 1 . Heat of Precipitation. (cal/mole)(Results of 12 groups
of 2 students each, arranged in. -
descending order.)Heat of Precipit&ma Calories/MoleSilver
Silverchloride bromide
-15.300*700 -20.030 =t 50 AverageNbte: All solut,ions are pre
pak d well ahead of time and -stored in the laboratory so they will
be at room temperature. - 5, 5OG -20,190 LiteratureThe author will
supply interested readers with a.sheet indiest- Negative sign in
thermodynamic sense of hea t liberated.ing to the student a form
for reporting results and outlining ral- Most freshman texts give
heat liberated apositive sign.Results and Discussion
Results of the first group to tr y this experiment aresummarized
in Table 1. The experiment has beenused several times in somewhat
larger classes withPresented before the Chemical Education Section,
Southesst-ern Regional ACS Meeting, Birmingham, Alabama,
November
1o'mL""".' GURNEY, . W.,'Ionic h ee 88e s in Sohtion,"
MoGraw-HillBook Co., New York, 1953,p. 93.
a Omitted from averages.Calchlated from data. in National Bureau
of StandardsCircular 500.AH i [AgS c) ] - AH l (Agt, hyp 1M ) - A H
, (X; hyp 1 M)
The results are further discussed as an exercise
int,hermochemistry. The fact th at reversing the reactionreverses
the sign of the heat effect is emphasized. Thusheat of solution =
-heat of precipitation.Ag+ + CI- = AgCl heat of preeip. -15,650
caI/mole
AgCl = Ag+ + Cl- heat of soh . 15,650 cal/mole470 / Journd of
Chemical Education
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Addition of reactions and associated heats to give theheat of a
third reaction is demonstrated. The relativesolubility of AgCl and
AgBr is given and the student isasked to calculate the heat of the
reaction:AgCl + Br- = AgBr + CI-
which is the difference AHnwnesr- Ah',.tr.cl =- 540 cal/mole.The
results are also discussed in terms of the solubilityprocess of
vaporizing the ionic crystal to gaseous ions(crystal lattice
energy) and hydrating the gaseous ions(hydration energy).2Mecs,+
X-(., -MXc crystal lattice energy, U
+ Q energy of hydrationX-i.1 + Hz0 - - ( wThus the heat of
solution is the difference of two largenegative heat effects:
AH..I.. = ZAHM - UThe student is given values of silver halide
crystallattice energies3 and asked to combine them with his
experimental results to calculate which ion (Cl- orBr-) has the
larger hydration energy. Alternativelythe student is sometimes
given the hydration energiesof the halide ions and asked to
estimate which salthas the higher crystal lattice energy.Related
topics can be discussed with freshmen onvery qualitative terms. For
example, the contribu-tions of Coulombic attractions, London
attractions,and Born repulsions3, to the crystal lattice energycan
be mentioned. Also th e solubility mechanism ofcrystal vaporized to
gaseous ions, then ions hydratedto form aqueous ions, leads
directly to a discussion ofthe influence of solvent dielectric
constant and thesolvolysis reaction on solubility.Discussing all
the points outlined above would takemuch more time than most
instructors could allowfor one experiment. Each instructor can use
his owndiscretion as to what subjects are important for dis-cussion
with his class. A brief discussion of errors andapplications to
thermochemical calculations wouldseem a good minimum of discussion
for this experiment.
2 LEUSSINR,D. L., in "Treatise on Analytical Chemistry,"edited
by KOLTHOPP,. M., AND ELVING,. J., InterecienceEncyclopedia, h e .
, New York, 1959. Part I, vol. 1, Chap. 17,p. 70if.
8 KETBLAAB,. A. A., "Chemical Constitution," 2nd ed.,Elsevier
Publishing Co., New York, 1958, Chap. 2.Gou~u ,E. S. , "Inorganic
Reaction and Structure," HenryHalt and Co., New York, 1955, p. 168
and Chap. 12.
Volume 38, Number 9, Sepfember 1961 / 471
