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the surface of bubbles which exist within the body of the liquid. The only reason why vaporization takes place so much more rapidly at the boiling temperature than just below it is that the evaporating surface is enormously increased as soon as the bubbles form. The reason why the temperature cannot be raised above the boiling point is that the surface always increases, on account of the bubbles, to just such an extent that the loss of heat because of evaporation is exactly equal to the heat received from the fire.

228. Distillation. Let water holding in solution some aniline dye be boiled in B (Fig. 180). The vapor of the liquid will pass into the tube T, where it will be condensed by the cold water which is kept in continual circulation through the jacket J. The condensed water collected in P will be seen to be free from all traces of the color of the dissolved aniline.



We learn, then, that when solids are dissolved in liquids, the vapor which rises from the solution contains none of the dissolved substance. Sometimes it is the pure liquid in P which is desired, as in the manufacture of alcohol, and sometimes the solid which remains in B, as in the manufacture of sugar. In the white-sugar industry it is necessary that the evaporation take place at a low temperature, so that the sugar may not be scorched. Hence the boiler is kept partially exhausted by means of an air pump, thus enabling the solution to boil at considerably reduced temperatures.

229. Fractional distillation. When both of the constituents of a solution are volatile, as in the case of a mixture of alcohol and water, the vapor of both will rise from the liquid. But the one which has the lower boiling point, that is, the higher

FIG. 180. Distillation

vapor pressure, will predominate. Hence, if we have in B (Fig. 180) a solution consisting of 50% alcohol and 50% water, it is clear that we can obtain in P, by evaporating and condensing, a solution containing a much larger percentage of alcohol. By repeating this operation a number of times we can increase the purity of the alcohol. This process is called fractional distillation. The boiling point of the mixture lies between the boiling points of alcohol and water, being higher the greater the percentage of water in the solution.

Gasoline and kerosene are separated from crude oil, and different grades of gasoline are separated from each other by fractional distillation.


1. A fall of 1° C. in the boiling point is caused by rising 960 ft. How hot is boiling water at Denver, 5000 ft. above sea level?

2. How may we obtain pure drinking water from sea water?

3. After water has been brought to a boil, will eggs become hard any quicker when the flame is high than when it is low?

4. The hot water which leaves a steam radiator may be as hot as the steam which entered it. How, then, has the room been warmed?

5. In a vessel of water which is being heated fine bubbles rise long before the boiling point is reached. Why is this so? How can you distinguish between this phenomenon and boiling?

6. When water is boiled in a deep vessel, it will be noticed that the bubbles rapidly increase in size as they approach the surface. Give two reasons for this.

7. Why are burns caused by steam so much more severe than burns caused by hot water of the same temperature?

8. How many times as much heat is required to convert any body of boiling water into steam as to warm an equal weight of water 1° C.? 9. How many B. T. U. are liberated within a radiator when 10 lb. of steam condense there?

10. In a certain radiator 2 kg. of steam at 100° C. condensed to water in 1 hr. and the water left the radiator at 90° C. How many calories were given to the room during the hour?

11. How many calories are given up by 30 g. of steam at 100° C. in condensing and then cooling to 20° C.? How much water will this steam raise from 10° C. to 20° C.?


230. Cooling by solution. Let a handful of common salt be placed in a small beaker of water at the temperature of the room and stirred with a thermometer. The temperature will fall several degrees. If equal weights of ammonium nitrate and water at 15° C. are mixed, the temperature will fall as low as - 10° C. If the water is nearly at 0° C. when the ammonium nitrate is added, and if the stirring is done with a test tube partly filled with ice-cold water, the water in the tube will be frozen.

These experiments show that the breaking up of the crystals of a solid requires an expenditure of heat energy, as well when this operation is effected by solution as by fusion. The reason for this will appear at once if we consider the analogy between solution and evaporation; for just as the molecules of a liquid tend to escape from its surface into the air, so the molecules at the surface of the salt are tending, because of their velocities, to pass off, and are only held back by the attractions of the other molecules in the crystal to which they belong. If, however, the salt is placed in water, the attraction of the water molecules for the salt molecules aids the natural velocities of the latter to carry them beyond the attraction of their fellows. As the molecules escape from the salt crystals two forces are acting on them, the attraction of the water molecules tending to increase their velocities, and the attraction of the remaining salt molecules tending to diminish these velocities. If the latter force has a greater resultant effect than the former, the mean velocity of the molecules after they have escaped will be diminished and the solution will be cooled. But if the attraction of the water molecules amounts to more than the backward pull of the dissolving molecules, as when caustic potash or sulphuric acid is dissolved, the mean molecular velocity is increased and the solution is heated.

231. Freezing points of solutions. If a solution of one part of common salt to ten of water is placed in a test tube and immersed in a "freezing mixture" of water, ice, and salt, the

temperature indicated by a thermometer in the tube will not be zero when ice begins to form, but several degrees below zero. The ice which does form, however, will be found, like the vapor which rises above a salt solution, to be free from salt, and it is this fact which furnishes a key to the explanation of why the freezing point of the salt solution is lower than that of pure water. For cooling a substance to its freezing point simply means reducing its temperature, and therefore the mean velocity of its molecules, sufficiently to enable the cohesive forces of the liquid to pull the molecules together into the crystalline form. Since in the freezing of a salt solution the cohesive forces of the water are obliged to overcome the attractions of the salt molecules as well as the molecular motions, the motions must be rendered less, that is, the temperature must be made lower, than in the case of pure water in order that crystallization may occur. From this reasoning we should expect that the larger the amount of salt in solution the lower would be the freezing point. This is indeed the case. The lowest freezing point obtainable with common salt in water is - 22° C., or -7.6° F. This is the freezing point of a saturated solution.

232. Freezing mixtures. If snow or ice is placed in a vessel of water, the water melts it, and in so doing its temperature is reduced to the freezing point of pure water. Similarly, if ice is placed in salt water, it melts and reduces the temperature of the salt water to the freezing point of the solution. This may be one, or two, or twenty-two degrees below zero, according to the concentration of the solution. Therefore, whether we put the ice in pure water or in salt water, enough of it always melts to reduce the whole mass to the freezing point of the liquid, and each gram of ice which melts uses up 80 calories of heat. The efficiency of a mixture of salt and ice in producing. cold is therefore due simply to the fact that the freezing point of a salt solution is lower than that of pure water.

The best proportions are three parts of snow or finely shaved ice to one part of common salt. If three parts of calcium chloride are mixed with two parts of snow, a temperature of – 55° C. may be produced. This is low enough to freeze mercury.


1. When salt water freezes, the ice formed is free from salt. What effect, then, does freezing have on the concentration of a salt solution?

2. A partially concentrated salt solution which has a freezing point of

– 5° C. is placed in a room which is kept at — 10° C. Will it all freeze? 3. Explain why salt is thrown on icy sidewalks on cold winter days. 4. Give two reasons why the ocean freezes less easily than the lakes. 5. Why does pouring H2SO, into water produce heat, while pouring the same substance upon ice produces cold?

6. Why will a liquid which is unable to dissolve a solid at a low temperature often do so at a higher temperature? (See § 230.)

7. When the salt in an ice-cream freezer unites with the ice to form brine, about how many calories of heat are used for each gram of ice melted? Where does it come from? If the freezing point of the salt solution were the same as that of the cream, would the cream freeze?


233. The modern steam engine. Thus far in our study of the transformations of energy we have considered only cases in which mechanical energy was transformed into heat energy. In all heat engines we have examples of exactly the reverse operation, namely, the transformation of heat energy back into mechanical energy. How this is done may best be understood from a study of various modern forms of heat engines. The invention of the form of the steam engine which is now in use is due to James Watt, who, at the time of the invention (1768), was an instrument maker in the University of Glasgow.

The operation of such a machine can best be understood from the ideal diagram shown in Fig. 181. Steam generated in the boiler B by the fire F passes through the pipe S into

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