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Molecular Weight by Boiling Point Elevation

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Molecular Weight by Boiling Point Elevation 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 1 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, T* . Based on the definition of molality it is possible to rearrange b Equation 1 into Equation 2 in order to isolate the solute molecular weight, M. M = [g • Kb] / [G • ∆Tb] 2 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. K = [R • T*2 ] / [∆H ] 3 b b vap The value of T* in Equation 3 is, as mentioned, that of the boiling point of the pure solvent b 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)] 4 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 5 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)] 6 The value of mapp may be evaluated as shown in Equation 7 with g, G, and Mapp as previously defined. mapp = g / [G • Mapp] 7 Equations 6 and 7 may be combined to show that @ = [Mo - Mapp ] / [Mapp • { n - 1}] 8 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. 2 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 super- heating 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 Boiling Point, T* . K K / P (torr) Solvent b b b (deg C at 760 torr) (molal / deg @ 1 atm) (molal/deg/torr) Acetone 56.0 1.71 0.0004 Benzene 80.2 2.53 0.0007 Bromobenzene 155.8 6.20 0.0016 Chloroform 60.2 3.63 0.0009 Ethanol 78.3 1.22 0.0003 Ethyl ether 34.4 2.02 0.0005 Methanol 64.7 0.83 0.0002 Water 100.0 0.51 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). 3 2 Digital thermometer capable of measuring to 0.1 ˚C equipped with a vapor phase temperature probe (a Beckmann thermometer is better if available). 3 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.
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