4. Since three-quarters of the earth"s surface is covered with water, why is not the air constantly saturated?
5. If the air has the temperature of the body, will fanning the perfectly dry face cool one? Explain. Will the effect be the same if the face is moist? Explain.
6. What is the cause of "Cloud Capped" mountains?
7. Why does the exhaust steam from an engine appear to have so much greater volume on a cold day in winter than on a warm one in summer?
8. What causes an unfrozen pond or lake to "steam" on a very cold day in winter, or on a very cool morning in summer?
9. As the air on a mountain top settles down the sides to places of greater pressure, how will its temperature be affected? its relative humidity? Explain.
10. On our Pacific coast, moist winds blow from the west over the mountains. Where will it rain? Where be dry? Explain.
CHAPTER VIII
HEAT AND WORK
(1) HEAT MEASUREMENT AND SPECIFIC HEAT
=177. Specific Heat.=--In the study of density and specific gravity it is made clear that different substances differ widely in the amount of matter contained in equal volumes, _e.g._, lead is much denser than water. The study of the relative densities of substance is usually considered under the subject of _specific gravity_.
_Specific heat_ as distinguished from specific gravity is concerned with the _capacity_ for heat possessed by different substances. The definition for specific heat is: _The ratio of the amount of heat required to change the temperature of a given ma.s.s of a substance 1 C.
degree to the amount of heat required to change the temperature of the same ma.s.s of water 1 C. degree._ By definition, it requires 1 calorie to raise the temperature of the gram of water 1C. The _specific heat_ therefore of water is taken as one. The specific heat of most substances except hydrogen, is _less_ than that of water, and as a rule, the denser the body the less its specific heat, as may be observed in the following table:
---------+----------+---------- | Specific | Specific | gravity | heat ---------+----------+---------- Gold | 19.3 | 0.032 Mercury | 13.6 | 0.033 Copper | 8.9 | 0.093 Bra.s.s | 8.4-8.9 | 0.094 Nickel | 8.57 | 0.11 Iron | 7.5+ | 0.1125 Aluminum | 2.67 | 0.218 Gla.s.s | 2.5-3.6 | 0.19 Ice | 0.918 | 0.504 Air | 0.00129 | 0.237 Steam | 0.00061 | 0.480 Hydrogen | 0.00009 | 3.409 ---------+----------+----------
=178. Method of Determining Specific Heat.=--The specific heat of a body is usually determined by what is called the _method of mixtures_.
For example, a definite weight of a substance, say a 200-g. iron ball, is placed in boiling water until it has the temperature of the hot water, 100C. Suppose that 300 g. of water at 18C. be placed in a calorimeter, and that the hot iron ball on being placed in the water raises its temperature to 23.5C. The heat received by the water equals 5.5 300 = 1650 calories. This must have come from the heated iron ball. 200 g. of iron then in cooling 76.5C.
(100-23.5) gave out 1650 calories. Then 1 g. of iron in cooling 76.5C. Would give out 8.25 calories or 1 g. of iron cooling 1C.
would yield about 0.11 calorie. The specific heat of the iron is then 0.11. For accurate determination the heat received by the calorimeter must be considered.
=179. Heat Capacity of Water.=--The large capacity for heat shown by water is useful in regulating the temperature of the air near lakes and the ocean. In hot weather the water rises slowly in temperature absorbing heat from the warm winds blowing over it. In winter the large amount of heat stored in the water is slowly given out to the air above.
Thus the climate near the ocean is made more moderate both in winter and summer by the large capacity of water for heat. This large heat capacity of water may seem to be a disadvantage when one is warming it for domestic purposes since it requires so much heat to warm water to boiling. However, it is this capacity that makes hot-water bottles and hot-water heating effective.
If one takes a pound of ice at 0C. in one dish and a pound of water at 0C. in another, and warms the dish of ice by a Bunsen flame until the ice is just melted, and then warms the water in the other dish for the same time, the water will be found to be _hot_ and at a temperature 80C., or 176F.
=180. The Heat of Fusion of Ice.=--This experiment indicates the large amount of heat required to change the ice to water without changing its temperature. As indicated by the experiment, it requires 80 calories to melt 1 g. of ice without changing its temperature or, in other words, if one placed 1 g. of ice at 0C. in 1 g. of water at 80C., the ice would be melted and the water would be cooled to 0C.
=181. Heat Given out by Freezing water.=--Just as 80 calories of heat are required to melt 1 g. of ice, so in freezing 1 g. of water, 80 calories of heat are given out.
The fact that heat is set free or given out when a liquid solidifies may be strikingly shown by making a strong solution of sodium acetate. On allowing this to cool quietly it will come to the room temperature and remain liquid. If now a small crystal of sodium acetate is dropped into the liquid the latter quickly becomes a solid ma.s.s of crystals, at the same time rising markedly in temperature. The amount of heat now liberated must enter the sodium acetate when the ma.s.s of crystals is melted again.
The large amount of heat that must be liberated before water freezes accounts for the slowness of the formation of ice. It is also the reason why the temperature never falls so low in the vicinity of large lakes as it does far inland, the heat given out by the freezing water warming the surrounding air.
The heat that disappears on melting and reappears on solidifying is called the _heat of fusion_. It is sometimes called _latent heat_ since the heat seems to become hidden or latent. It is now believed that the heat energy that disappears when a body melts has been transformed into the _potential energy_ of partially separated molecules. The heat of fusion therefore represents the work done in changing a solid to a liquid without a change of temperature.
=182. Melting of Crystalline and Amorphous Substances.=--If a piece of ice is placed in boiling hot water and then removed, the temperature of the unmelted ice is still 0C. There is no known means of warming ice under atmospheric pressure above its melting point and maintaining its solid state. Ice being composed of ice crystals is called a crystalline body. All crystalline substances have fixed melting points. For example, ice always melts at 0C. The melting points of some common crystalline substances are given below:
_Melting Points of Some Crystalline Substances_
1. Aluminum 658 C.
2. Cast iron 1200 C.
3. Copper 1083 C.
4. Ice 0 C.
5. Lead 327 C.
6. Mercury -39 C.
7. Phenol (carbolic acid) 43 C.
8. Platinum 1755 C.
9. Salt (sodium chloride) 795 C.
10. Saltpeter (pota.s.sium nitrate) 340 C.
11. Silver 961 C.
12. Sodium hyposulphite (hypo) 47 C.
13. Zinc 419 C.
_Non-crystalline or amorphous substances_ such as gla.s.s, tar, glue, etc., do not have well defined melting points as do crystalline bodies.
When heated they gradually soften and become fluid. For this reason gla.s.s can be pressed and molded.
=183. Change of Volume During Solidification.=--The fact that ice floats and that it breaks bottles and pipes in which it freezes shows that water expands on freezing. How a substance may occupy more s.p.a.ce when solid than when liquid may be understood when we learn that ice consists of ma.s.ses of star-shaped crystals. (See Fig. 151.) The formation of these crystals must leave unoccupied s.p.a.ces between them in the ice.
When liquefied, however, no s.p.a.ces are left and the substance occupies less volume. Most substances contract upon solidifying. Antimony and bis.m.u.th, however, expand on solidifying while iron changes little in volume. Only those bodies that expand, or else show little change of volume on solidifying, can make sharp castings, for if they contract they will not completely fill the mold. For this reason gold and silver coins must be stamped and not cast. Type metal, an alloy of antimony and lead, expands on solidifying to form the sharp outlines of good type.
Several important effects of the expansion of water when freezing should be noted. (a) Ice floats, (b) if it sank as soon as formed, lakes and rivers would freeze solid, (c) freezing water is one of the active agents in the disintegration of rocks.
[Ill.u.s.tration: FIG. 151.--Ice crystals.]
[Ill.u.s.tration: FIG. 152.--Melting ice by pressure.]
Since water expands on freezing, pressure would on compressing ice at 0C., tend to turn it into water. Pressure does lower the melting point of ice, so that a little ice may melt when it is subjected to pressure.
On removing the pressure the water freezes. This may be shown by placing a loop of fine piano wire (see Fig. 152) over a piece of ice supported so that a weight may be hung upon the wire. The wire will be found to gradually cut through the ice, the melted ice refreezing above the wire.
Important Topics
1. Specific heat.
2. Heat of fusion of ice.
3. Crystalline substances have fixed melting points.
4. Expansion on freezing, importance.
Exercises
1. What are two advantages in the high heat of fusion of ice?