Experimental determination of the boiling point

In the diagrams shown in Section 14.5, the temperature was constant. The equilibrium pressure of the system was then a function of either or y, according to Eqs. (14.9) or (14.12). In those equations, the values of pi and are functions of temperature. If, in Eqs. (14.9) and (14.12), we consider the total pressure p to be constant, then the equations are relations between the equilibrium temperature, the boiling point, and either x or y. The relations T = /(x ) and T = g(yi) are not such simple ones as between pressure and composition, but they may be determined theoretically through the Clapeyron equation or, ordinarily, experimentally through determination of the boiling points and vapor compositions corresponding to liquid mixtures of various compositions. 

Materials. The UFe used in this work was a portion of a larger batch originally obtained from Oak Ridge National Laboratory. Almost two-thirds of the original batch had been distilled away in previous experimental work, presumably contributing to the purification of the UFe from low boiling impurities e.g., HF, CF4, F2). Emission-spectrographic analysis of the material indicated that the predominant impurities were P, at a concentration of <400 p.p.m., and As, B, Cs, Pd, Re, Sb, Sn, and Th, each present at concentrations of <100 p.p.m. Two determinations of the triple point of a sample of the UFe yielded values of 64.1 °C. and 64.2°C. The best literature value 19) for this is 64.05°C. 

Three determinations of the melting point of technetium metal are in reasonably good agreement 2140 20 °C [13, 2200 50 °C [14], and 2162 40°C (15). The average melting point of 2167 °C is near those of neighboring elements in the same period —molybdenum (2610 °C) and ruthenium (2310 °C) —but almost 1000 °C lower than that of rhenium [14]. Ilic boiling point of technetium metal was estimated as 4900 K [16] no experimental value seems to be available. 

Almost all solids are more soluble in a hot than in a cold solvent, and solution crystallization takes advantage of this fact. Thus, if you first dissolve a solid in an amount of hot solvent insufficient to dissolve it when cold, crystals should form when the hot solution is allowed to cool. The extent to which the solid precipitates depends on the difference in its solubility in the particular solvent at temperatures between the extremes used. The upper extreme is determined by the boiling point of the solvent, whereas the lower limit is usually dictated by experimental convenience. For example, an ice-water bath is often used to cool the solution to 0 °C, whereas ice-salt and dry ice-acetone baths are commonly used to cool solutions to -20 °C and -78 °C, respectively (Sec. 2.10). The solid should be recovered with greater efficiency at these temperatures, provided the solvent itself does not freeze. 

Experimental Determination of Boiling-point. Unless only minute quantities of the liquid are available cj. p. 60), the boiling-point is usually determined by simple distillation. For this purpose, the apparatus shown in Fig. 2 is assembled. A distillation flask A of suitable size is fitted to a water-condenser B, the water supply of which is arranged as show-n. An adaptor C is sometimes fitted in turn to the condenser, so that the distillate... 

Iodide. A 0.01 M solution of potassium iodide, prepared from the dry salt with boiled-out water, is suitable for practice in this determination. The experimental details are similar to those given for bromide, except that the indicator electrode consists of a silver rod immersed in the solution. The titration cell may be charged with 10.00 mL of the iodide solution, 30 mL of water, and 10 mL of the stock solution of perchloric acid + potassium nitrate. In the neighbourhood of the equivalence point it is necessary to allow at least 30-60 seconds to elapse before steady potentials are established. 

The same reaction can be applied, not only to the aromatic parent substances, the hydrocarbons, but also to all their derivatives, such as phenols, amines, aldehydes, acids, and so on. The nitration does not, however, always proceed with the same ease, and therefore the most favourable experimental conditions must be determined for each substance. If a substance is very easily nitrated it may be done with nitric acid sufficiently diluted with water, or else the substance to be nitrated is dissolved in a resistant solvent and is then treated with nitric acid. Glacial acetic acid is frequently used as the solvent. Substances which are less easily nitrated are dissolved in concentrated or fuming nitric acid. If the nitration proceeds with difficulty the elimination of water is facilitated by the addition of concentrated sulphuric acid to ordinary or fuming nitric acid. When nitration is carried out in sulphuric acid solution, potassium or sodium nitrate is sometimes used instead of nitric acid. The methods of nitration described may be still further modified in two ways 1, the temperature or, 2, the amount of nitric acid used, may be varied. Thus nitration can be carried out at the temperature of a freezing mixture, at that of ice, at that of cold water, at a gentle heat, or, finally, at the boiling point. Moreover, we can either employ an excess of nitric acid or the theoretical amount. Small scale preliminary experiments will indicate which of these numerous modifications may be expected to yield the best results. Since nitro-compounds are usually insoluble or sparingly soluble in water they can be precipitated from the nitration mixture by dilution with water. 

Boiling point elevations are directly proportional to the molality of a solution, but chemists have found that some solvents are more susceptible to this change than others. The formula for the change in the boiling point of a solution, therefore, contains a proportionality constant, abbreviated K, which is a property determined experimentally and must be read from a table such as Table 13-2. The formula for the boiling point elevation is... 

The IBM machines were used to set up the Antoine constants from determined data. A preliminary C value was obtained from the equation C = 239. — 0.19 /,. A and B were then obtained and new C values either side of the first C used and new A and B values found. In each case above, the boiling points at the experimental pressures were calculated and compared with the determined boiling points. 

Modern inorganic chemistry is a quantitative science. Consequently, when performing experimental work, students must determine the yield of the substances obtained and certain constants such as the boiling points, solubility, and cryohydrate points, and also perform the required calculations with the use of the fundamentals of thermodynamics. 

The slope, — d(2ri) /d0, of the straight lines which characterize combustion of various compounds as well as the lighter petroleum fractions during the evaporation stage has been termed the evaporation constant —i.e., n2 = rr -o — 0/4 and e = 2M/(tp°ti) (7-11, 4 )- Experimentally determined values of this constant are presented in Table IV. The burning lifetimes of fuel drops initially at the boiling point—i.e., with no preheat period—can be calculated by dividing the square of the initial drop diameter by the evaporation constant. 

The vapor pressures of all tour compounds hove been measured up to the boiling point Christensen and Smith have determined the vapor pressure of acetaldehyde up to I50 C The vapor pressures from (he boiling point up to the critical point have been calculated by the method of Miller. When compared with ihc experimental data on dcetuldchyde. the error averaged. ... 

Only the heats ot vtipon7ation at the boiling points huve been experimentally determined 1 u- Kharbanda s nomograph has been used to calculate the heats of vaporization to the critical temperuture. ... 

When reasonable amounts of liquid components are available (> 5 ml) the boiling point is readily determined by slowly distilling the material from a pear-shaped flask in an apparatus assembly shown in Fig. 2.98, and recording the temperature at which the bulk of the compound distils. Due attention should be paid to the experimental procedure which was discussed in detail in Section 2.24. 

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