Idea Transcript
MOLECULAR WEIGHT BY BOILING POINT ELEVATION BACKGROUND This experiment demonstrates the use of colligative properties. The goal is to measure the molecular weight of a non-volatile solute by determining the concentration dependence of the boiling point elevation of a solution. The solvent used must be one of the compounds commonly referred to as volatile; that is, it must have an appreciable vapor pressure. One of the several useful aspects of colligative properties is the fact that the vapor pressure of volatile solvents is lowered when a non-volatile solute is used to make a solution. The result is that such a solution will necessarily have a higher boiling point than that of the pure solvent. The higher boiling point is due to the fact that a higher temperature is needed in the presence of the non-volatile solute, which is not making any contribution to the solution’s vapor pressure, in order to cause the volatile component of the solution, the solvent, to exert one atmosphere of pressure. It must be remembered that the boiling point elevation being investigated in this experiment is a property of the solution as a whole and, for ideal dilute solutions, is directly proportional to the solute concentration as shown in Equation 1. ∆Tb = m • Kb
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In Equation 1, m is the solution molality and Kb is the boiling point elevation constant which is a function of the solvent not the solute. The value of ∆Tb is the boiling point of the solution minus that of the pure solvent, Tb* . Based on the definition of molality it is possible to rearrange Equation 1 into Equation 2 in order to isolate the solute molecular weight, M. M = [g • Kb] / [G • ∆Tb]
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In Equation 2, ∆Tb is the boiling point elevation of a solution containing g grams of solute in G kilograms of solvent. It is interesting to note that Kb for a solvent may be estimated by Equation 3. Kb = [R • Tb*2 ] / [∆Hvap ]
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The value of Tb* in Equation 3 is, as mentioned, that of the boiling point of the pure solvent while ∆Hvap is the heat of vaporization of the pure solvent. Frequently, for dilute ideal solutions, the boiling points of the pure solvent and that of the solution may be used interchangeably. Equations 1, 2, and 3 assume non-volatile and non-electrolyte solutes in ideal dilute solutions. Since colligative properties are independent of the identity of the solute, depending only on total particle concentration, it is possible to obtain important information about electrolytes by measuring boiling point elevations. This is especially true in the case of weak electrolytes that are only partially dissociated in solution. To the extent that a solute dissociates in solution the net number of actual particles present will increase. A larger number of particles in solution will result in a larger measured boiling point elevation, ∆Tb, because the molality, m, in Equation 1
will be larger. In the case of a weak electrolyte it is more appropriate to re-write Equation 1 as shown in Equation 4. ∆Tb = Kb • mapp = Kb • [g / (G • Mapp)]
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In equation 4, mapp is the “apparent” total particle molality that results from the partial dissociation of the solute. Mapp is a weighted average molecular weight representative of the actual ions and molecules as they exist in solution. The value of mapp may be evaluated in terms of the analytical molality, which is based on formula weight of solute, and the percent of the solute particles which undergo dissociation, @, as shown in Equation 5. mapp = n • @ • mo + (1 - @) • mo
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In equation 5, n represents the number of ions produced by the dissociation of one molecule of solute while mo represents the molality based on the formula weight of the solute ignoring any dissociation. Equation 5 can be rearranged into equation 6 where g and G have the same meaning they did in equation 2 while Mo is the actual formula weight of the solute. mapp = [{n - 1} • @ + 1] • [g / (G • Mo)]
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The value of mapp may be evaluated as shown in Equation 7 with g, G, and Mapp as previously defined. mapp = g / [G • Mapp]
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Equations 6 and 7 may be combined to show that @ = [Mo - Mapp ] / [Mapp • { n - 1}]
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Since Mapp may be evaluated from the boiling point elevation as shown in Equation 4 it is easy to determine the degree of dissociation of a weak electrolyte using Equation 8. Inter-particle Forces In addition to the influence of electrolyte behavior the boiling point of solutions may be affected by the way in which solute particles interact with one another. Typically attractive forces exist between solute particles. These forces can have an influence on the value of the molecular weight as calculated using Equation 2. Equation 2 suggests a possible approach to attempt to correct for such inter-particle forces. A plot can be constructed of calculated molecular weight on the y-axis versus solution molality on the x-axis. Such a plot is extrapolated to zero molality, the y-intercept, to obtain what is referred to as the limiting molecular weight. This infinite dilution extrapolated molecular weight value may be thought of as the value of the molecular weight that would be obtained if only a single molecule were present in solution. Such a technique has the result of minimizing the influence of inter-particle forces.
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Since Equation 2 requires values for the Kb constants, the values for common solvents are shown in Table 1. Notice that since Kb is sensitive to atmospheric pressure, the information is provided in Table 1 so that Kb values may be corrected to the barometric pressure in the laboratory at the time the experiment is conducted. A final note of caution is appropriate. The heart of this experiment is the measurement of how the boiling point of a pure solvent changes as a solution is prepared. The only reliable manner to accomplish this measure of temperature change is to make an actual measurement of the boiling point of the pure solvent at the actual laboratory conditions rather than depending on literature values. Since boiling points show significant pressure variation the boiling point of the pure solvent may be significantly different from the literature value if the barometric pressure in the laboratory is significantly different from standard atmospheric conditions. Failure to actually measure the boiling point of the pure solvent under the conditions that the solution boiling points are measured can give rise to serious errors! Further, it is good to be mindful of the fact that when measuring the boiling point of a solution it is best to measure the actual boiling liquid itself rather than the vapor. When inserting thermometers into boiling liquid it is easy to experience difficulties associated with the superheating of the liquid. For this reason it is best to use an apparatus such as the Cottrell boiling point apparatus or similar equipment. It is also good to use one thermometer throughout in order to avoid any potential problems associated with faulty thermometer calibrations. TABLE 1: boiling point elevation constants
Solvent Acetone Benzene Bromobenzene Chloroform Ethanol Ethyl ether Methanol Water
Boiling Point, Tb* . (deg C at 760 torr) 56.0 80.2 155.8 60.2 78.3 34.4 64.7 100.0
Kb (molal / deg @ 1 atm)
Kb / P (torr) (molal/deg/torr)
1.71 2.53 6.20 3.63 1.22 2.02 0.83 0.51
0.0004 0.0007 0.0016 0.0009 0.0003 0.0005 0.0002 0.0001
PROCEDURE Apparatus and Materials Needed: 1
Cottrell boiling-point apparatus or equivalent (standard reflux/distillation set-up may be used if Cottrell apparatus not available; ask instructor for assistance).
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Digital thermometer capable of measuring to 0.1 ˚C equipped with a vapor phase temperature probe (a Beckmann thermometer is better if available). Suitable solvent and solute as suggested by instructor (benzoic acid, urea, oxalic acid, or naphthalene as the solute and acetone, methanol, or ethanol as the solvent).
Boiling Point Measurement If a Cottrell-type boiling point apparatus is available it will include suitable means for withdrawal of solution aliquots. When a Cottrell apparatus is unavailable it may be necessary to use a reflux/distillation set-up. When using a reflux set-up it is advisable to secure a piece of clean cotton cloth to the bulb of the thermometer. The cotton cloth, as it becomes wetted with condensing vapor, will insure the thermometer bulb is actually in contact with liquid rather than vapor. If a reflex/distillation set-up is used be sure that a two-neck round bottom flask is used in order to facilitate the addition of solute and removal of solution aliquots. A convenient size to use is a 250 ml round bottom flask. The round bottom flask is fitted with a reflex column mounted vertically in one of the necks. The second neck has a stopper inserted At the top of the reflux column an adapter is fitted to permit the positioning of the thermometer as well as to permit the attachment of a guard condenser column to avoid escape into the laboratory of vapor as the solutions are boiled. Water cooling should be used as appropriate, particularly in the guard column. Begin by placing about 250 ml (graduated cylinder) of the solvent into the round bottom flask and assembling the apparatus. Using a safe heat source as directed by the instructor heat the pure solvent to establish a stable reflux situation. Allow sufficient time for a stable temperature indication and record this value as the boiling point of the pure solvent under the current laboratory barometric pressure. Remove the heat source long enough for the boiling to subside then carefully introduce, through the second neck, enough solid solute to produce a solution whose weight percent concentration is 2% to 5%. For some solids, for example benzoic acid, it may be advisable to prepare solid pellets using a pellet press. Your instructor will provide assistance with this if it is necessary. After introduction of solute the boiling point of the solution is measured in the same way in which the boiling point of the pure solvent was measured. Be sure to allow sufficient time for a stable boiling point to be reached. After the boiling point of the solution has been recorded remove the heat source and allow the boiling to subside. At this point a pipet is used to withdraw a sample of solution from the flask through the second neck. It is convenient to remove approximately 10 ml of sample although it is not necessary to accurately measure the actual volume nor that all aliquots are of the same volume. Each aliquot should be placed into a securely sealed weighing bottle and placed in a cold bath to promote rapid cooling. The empty weighing bottle should be previously weighed on an analytical balance. For efficient use of time the weighing bottle and its contents can be set aside as more solution samples are collected. Once several samples have been collected all of the solutions may be analyzed to determine the solution molality. Each subsequent solution is prepared from the previous one by adding, through the second neck of the round bottom flask, enough solute to produce a boiling point increase of about 0.1 to 0.2 degrees.
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Each solution’s boiling point is measured and recorded. An aliquot of each solution is withdrawn and placed in a separate pre-weighed weighing bottle that is then cooled and set aside. Collect aliquots from four or five different solutions as directed by the instructor. Each of these aliquots are analyzed as described below. Once its contents have been cooled to room temperature each bottle and its contents are weighed on an analytical balance. These data permit the weight of solution to be determined. The top is now removed from the weighing bottle and the solvent is evaporated off. Typically a hot water bath may be used to accomplish the solvent evaporation. If directed by your instructor a drying oven may be used to remove the final residual of solvent to leave only solid solute in the weighing bottle. Cool the weighing bottle and the solid solute to room temperature and weigh it. Determine the weight of the solid solute (g in Equation 2). The weight of the solvent (G in Equation 2) may be determined by the difference. CALCULATIONS Equation 2 may be used to calculate the molecular weight of the solute for each of the different solution concentrations. Use g and G for each of the solutions to calculate the molality of the solution. Prepare a graph of molecular weight, M, on the y-axis versus the solution molality, m, on the x-axis. The resultant line may be extrapolated to zero molality to obtain what is referred to as the “limiting molecular weight, Minfinite”. Calculate the degree of dissociation of the solute for each of the solutions whose boiling points have been measured. Prepare a plot of the degree of dissociation, @, on the y-axis versus the molality, mo, on the x-axis and extrapolate to zero concentration to obtain a limit at infinite dilution of the degree of dissociation, @infinite.
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