HEAT. 687 HEAT. increased its melting-point is lowered. Therefore, if two blocks of ice at 0° are pressed together, the ice under pressure at the point of contact has a melting-point less than 0°, but being at 0" is at a temperature which is higher than its melt- ing-point, and so melts. The water whieh results from this is at a temijerature lower than 0°; and so as it flows out from under the pressure it freezes again, because now its freezing-point is 0°. The motion of glaciers depends on this phenomenon. The presence of nuclei greatly facilitates the processes of solidification, boiling, and condensa- tion. A drop or a bubble cannot be formed with- out some nucleus (see Capillakity) ; and a liquid can be cooled far below its freezing-point if there is no nucleus and if the liquid is not disturbed or jarred. A liquid exposed in an open vessel will evaporate, i.e. will pass slowly into the form of vapor, at all temperatures below its boiling-point. If, however, a large closed vessel is placed over the one containing the liquid, the evaporation will soon cease apparently — there is equilibrium between the liquid and the vapor. The process does not in reality stop: but the evaporation of the liquid is exactly balanced by the condensation of the vapor. Experiments show that this state of equilibrium is reached at a cer- tain temperature when the pressure of the vapor is a definite quantity, viz. at the 'boiling-point' for that pressure : if the temperature is changed, so is the pressure which corresponds to equilib- rium. If the temperature in the above ex- periment is lowered, some vapor will condense; illustrations of this are the formation of dew, for- mation of drops of water on ice-pitchers, etc. Again, if in the above experiment the volume of the space over the liquid is made smaller, some of the vapor will be condensed. Tims there are two methods for the liquefaction of a gas or vapor: to lower the temperature and to de- crease the volume. It was established by An- drews in 1809 that a vapor cannot be liquefied by any decrease in volume, however great, un- less the temperature is below a certain limit, which is different for different substances, and which is called the 'critical temperature.' But by lowering the temperature sufficiently and by making the volume small enough, all known gases have been liquefied with the exception of helium. (See LiQ^TEFACTiON OF GASE.S.) The critical temperatures of a few gases are given below : Critical Temperatures If an unsaturated solution has its temperature lowered below 0°, the freezing-point of the solu- tion will finally be readied, pure ice will separate out, leaving the solution more concentrated: tlio freezing-point of this solution is lower than that of the first; and so as the temiKjrature gets lower and lower, the solution becomes more and more concentrated until finally it is saturated. It now heat-energy is removed, ice will form, but salt will be deposited also in cfiuivalcnt amounts; tliis mixture of ice and salt is called the 'cryo- hydrate,' and the temperature of its formation is — 22° C. (Other salts and liquids have dilVerent erj'ohydrates and ditl'erent temijeratures of for- mation.) In this cryohydratc of common salt there are 23.8 parts by weight of the salt and 76.2 parts of ice. Therefore, if a mixture is made of salt and ice in this proportion, it will form a solid whose melting-point is — 22° C. ; and if it actually is at the temperature' —5° or — 10°, it will of course melt, and in so doing heat-energy will be abstracted from siirroxmding bodies, because energy is required for two reasons, to melt the ice and to make the salt dissolve in the water. For this reason such a mixture is called a 'freezing-mixture.' The fact that heat-energy is required to make a solid melt or a liquid evaporate is familiar from many experiments. Similarly heat-energy is liberated when a liquid freezes or vapor con- denses. The number of calories corresponding to the change in state of one gram of a substance under a definite pressure is called the latent heat for that change at the given pressure. Similarly, when one substance dissolves in another, there are as a rule changes in temperature showing that heat-energy is liberated or absorbed. The 'heat of solution' is defined to be the number of calories produced when one gram of a substance dissolves in a great mass of a given solvent, a quantity so great that any further increase in it would not affect the heat-energ>' liberated or absorbed. Values for latent heats and heats of solution are given in the following tables: Latent Heats Carbon dioxide 300.92 Sulpiiur dioxide 156. Sulpliuric ether 194.4 Water 365 Ammonia 130 If substances are dissolved in a pure liquid, both its freezing and its boiling points are al- tered; the foi-mer is lowered and the latter is raised, the amount of the change varying di- rectly with the quantity of substance dissolved. In most cases, however, both the solid and the vapor formed are those of the pure liquid. If the dissolved substance is volatile, then it will evaporate also. Common salt and water seire as an illustration. If salt is added in small amounts to a vessel of water, a time will come when the water will no longer hold the salt in solution, but will deposit it: the solution is said to be 'satu- rated.' The amount of salt required to produce saturation varies directly with the temperature. 8TTB8TANCE
- Fusion
StlBBTANCE Vaporiza- tion Water 80.02 9.37 30.86 2.82 21.07 Water 636.6 Ammonia 297. 92.9 Ether 911: Silver Cliloroform 68.5 Hydrop:en — 234 Oxygen — 118 Nitrogen — 146 Heats op Solution is Water Ammonia gas -t-495.6 Etlivl alcohol + 65.3 Siilplmric acid -fl82.6 Caustic potash 4-223.3 Sodium chloride — 18.22 Potassium chloride — 69.7 Silver chloride — 110. -f Means rise in temperature. — Means fall in temperature. There are. of course, other changes of state than those mentioned : among the.se are sublima- tion, when a solid passes directly into the state of vapor: dissociation of a gas, when the mole- cules of a gas break up into other parts; etc. It is found, however, in all these cases that there will be equilibrium at a definite temperature only when a certain pressure is reached, and con- versely. Transfer of Heat. There are three processes
