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 mass of a substance 1 C. degree to the amount of heat required to change the temperature of the same mass of water 1 C. degree. By definition, it requires 1 calorie to raise the temperature of the gram of water 1°C. 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 gravity | Specific heat | |
| Gold | 19.3 | 0.032 |
| Mercury | 13.6 | 0.033 |
| Copper | 8.9 | 0.093 |
| Brass | 8.4-8.9 | 0.094 |
| Nickel | 8.57 | 0.11 |
| Iron | 7.5+ | 0.1125 |
| Aluminum | 2.67 | 0.218 |
| Glass | 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, 100°C. Suppose that 300 g. of water at 18°C. be placed in a calorimeter, and that the hot iron ball on being placed in the water raises its temperature to 23.5°C. 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 76.5°C. (100°-23.5°) gave out 1650 calories. Then 1 g. of iron in cooling 76.5°C. Would give out 8.25 calories or 1 g. of iron cooling 1°C. 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 0°C. in one dish and a pound of water at 0°C. 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 80°C., or 176°F.
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 0°C. in 1 g. of water at 80°C., the ice would be melted and the water would be cooled to 0°C.
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 mass of crystals, at the same time rising markedly in temperature. The amount of heat now liberated must enter the sodium acetate when the mass of crystals is melted again.