The Heat of Formation of Na2SiF6 A Physical Chemistry Laboratory Experiment T. C. DeVore and T. N. Gallaher James Madison University, Harrisonburg, VA 22807 The concepts of heat (enthalpy) of formation, heat (enthalpy) of reaction, and chemical equilibrium are familiar ones in most general and physical chemistry courses. Classic laboratory experiments to measure these quantities and to explore the relationships between them are available in most physical chemistry laboratory manuals. ( 1 , 2 )While classics, these experiments are viewed as being "antiques of little use" by many students, particularly those who plan t o enter the job market after receiving their BS degrees. In order to provide the illusion of "being relevant", this experiment which uses infrared spectroscopy to measure the heat of formation of NazSiFs has been added to the applied physical chemistry laboratory at JMU. This experiment illustrates several points about science. I t has the student use afamiliar instrument for an unfamiliar application. This shows that there are often many ways to attack problems in science. Since AHr(NazSiF6) was last measured in 1981 (3), it clearly shows that research in thermodynamics is still an active experimental branch of chemistry (i.e., a modern topic). Finally, since early work overlooked the NazSiFs (m)-NanSiF&) phase change a t 833 K and the early data was collected a t both higher and lower temperatures ( 4 , 5 ) ,the need to evaluate literature values critically is apparent.
pressure of the gas is directly proportional to the absorbance if Beer's Law is obeyed. Thus,
where f = (Rlrb) = constant for this experiment. Substituting eq 6 into eq 3 gives
Integrating and rearranging gives
Hence, a plot of In(A7') versus 1/T will also give a straight line with slope equal to -AHIR.
Theory When heated, NalSiFs decomposes via the reaction (3-5)
Apparatus High-temperatureIRcell or IR cell connectedto a furnaceassembly Furnace and power supply IR spectrometer Vacuum system
The equilibrium constant for this reaction is readily seen to he
Several possible cell designs can be envisioned that could be used to do this experiment. The cell used at JMU is shown schematically in Figure 1. The cell hody is an MDC four-way eross-vacuum flange.
K , = psi.po
(1)
where P s ~ F is ~the pressure of SiF4 and Po is the standard pressure (101.32 kPa). The Gibbs-Helmholtz equation, which is derived in most physical chemistry texts (6-8)can be rearranged to give
Substitution of eq 1 into eq 2 gives
Assuming that AH is not a strong function of temperature, the familiar integrated form of this equation shows that a plot of In P s i ~ ~ v e r s11T u s is a straight line with a slope equal to -AHIR. The familiar ideal gas law equation can be rearranged to give P = CRT
low
(4)
quartz furnace
where the concentration C = n1Vand P, R, and T have their usual meaning. Beer's Law states that the absorbance of the sample is related to the concentration (C) by A = cbC
(5)
where cb is a constant for samples of the same molecule taken in the same cell. Combining eqs 4 and 5 shows that the
171 Figure 1. Schematic diagram of the apparahrsused in this experiment. Volume 63 Number 8 August 1986
729
KBr optical windows are attached to the cell using Viton O-rings and connection flanges available from MDC. A Cajun O-ring tube connector that had been silver-soldered into a hale drilled into a blank flange ,. was used to connect the furnace to the IR cell. A sealed 9-nlm quartz t u l w zerved as the sample companmrnf of the iurni(w.The hratingclrmmt wan formed by wrapping a 12-mm quartz tube with Nichrome heating wire. Better performance was obtained by covering the wire with thermal insulation. A chromelalumel thermocouple placed between the quartz tubes was used to measure the temperature. If desired, glass blower's cloth can be substituted for the 12-mm quartz tube. This has the advantage of heatine faster than does the ouartz. which somewhat shortens the time needed t o do the rrprrilnmr. While a Perkin regulated dc pwer supply waa urrd to hrnr the furn~ce,any Variac-typepower aupply raiuble of deliwring 1%20amps of current wuuld nork. ~
Na, S i Fg 1984 DATA 1985 DATA
~
SLOPE
75.Or.4 w
Procedure The experiment is done by placing 1-2 g of NszSiF6 in the sample section of the furnace and evacuating the eell prior to class. The furnace is heated to the lowest desired temperature (-550 K) and the valve is closed to isolate the cell from the vacuum line. After allowing 10-20 min to assure that the sample has reached equilibrium, the IR spectrum of the 1031cm-I band of SiFa is scanned. The temperature is raised without opening the valve and the process is repeated.The maximum temperature that can beused is-675 K. At least five data points can be obtained in a 3-h lab period. A?l (reaction) is determined from a plot of the lnAT versus 1/T. This value is used with the values for the heat of formation of S ~ F I and NaF obtained from the JANAF tables (9)to determine the heat of formation of Na2SiF6(S).If desired, the heat capacity given by Chiotti can be used to correct AHrto 298.15 K. Some observations made about this experiment are 1) Equilibrium is established quickly on heating, but slowly an
Safety ~recautlons Na2SiF6 is somewhat toxic (10) and SiF4 will react with moist air to form HF slowlv. so some care should be taken in handling the chemicals. 1 f t h e instructor loads the cell and evacuates the cell, the students should not he exposed to the chemicals. Since only a small amount of SiF4 is formed durina the experiment, it can he collected in a liquid Np trap in thevacuum line and vented through the fume hood ~ h used Na2SiFBcan he handled using regular solid waste disposal procedures. The usual dangers encountered in handling vacuum lines are also present. Standard vacuum line safety procedures (I, 2) should also he employed. Results The data from lab groups that did this experiment during the spring semester of 1984 and 1985 are shown in Figure 2. The IR svectra to make the plots were obtained in the % transmittance mode 11984)a n i the ahsorhance mode ( 1 9 8 5 ) usina a Nicolct MX-l FT-IR. No corrections to rhe rhermocouple readings were made. A standard linear least-squares
Journal of Chemical Education
Flgure 2. Plot of INAT) versus T-' obtained far the thermal decomposition of Na2SiFe.This data was wllected over a two-year period by students in lhe
Applied Physical Chemistry Laboramry at JMU.
Cornoarlson oi Literature and Exoerimental Valuesa Source
TRange (K)
This work Ref 3 Refs H , 12
550-650 724-833
...
&N
(kK-')
16.0(4) 18.0(9)
. ..
W)(kK-') 349.2(4) 350.9
. ..
4 2 9 8 ) (kK-') 349.9(4) (351.7)b 350.0
All v a i m are in k K V Calmlated as ( ~ n in e this experiment. When mlid solublllty arnectionn are applied as was done In ref 3, avalue of 350 kK-' is obtained.
cooling. Therefore only increasing temperatures should be used unless the cell is evacuated between runs. 2) SiFl adsorbs to stainless steel. The eell should he treated by filling it with a small amount of SiF4 for %4 h prior to use in this experiment. 3) Theoretical calculations indicate that -0.4 g of NazSiFe is adequate to do the experiment in a cell with a 2-L volume once if the cell is not evacuated during the experiment.
730
..
b
treatment was used to compute the slope and standard deviation of the data points. The results from all data points are summarized in the table. Not surnrisinelv. the individual groups varied slightly from the class average data, and the individual standard deviations were sliehtlv lareer. However, all groups were within experimental &o; ofthe literature value for AHi(Na2SiF6) reported by Stull e t al. (11) and Wagmer et al. (12). Literature Clted
k (31 Chiofti, P. J. Loas Comm~nMeLola1981,80.97. (4)
Hantke,G.Z.AngauCham.1921.39.1065.
151 Ceilut,R.Ann.Chim.1945,20,367.
(61 Mwre. W.J. "BasicPhysica1Chemiatry";Prenfi~-Hall: EnglewmdCliffs. NJ, 1983; "
