# 14.3 Phase change and latent heat  (Page 2/9)

 Page 2 / 9
Heats of fusion and vaporization Values quoted at the normal melting and boiling temperatures at standard atmospheric pressure (1 atm).
L f L v
Substance Melting point (ºC) kJ/kg kcal/kg Boiling point (°C) kJ/kg kcal/kg
Helium −269.7 5.23 1.25 −268.9 20.9 4.99
Hydrogen −259.3 58.6 14.0 −252.9 452 108
Nitrogen −210.0 25.5 6.09 −195.8 201 48.0
Oxygen −218.8 13.8 3.30 −183.0 213 50.9
Ethanol −114 104 24.9 78.3 854 204
Ammonia −75 108 −33.4 1370 327
Mercury −38.9 11.8 2.82 357 272 65.0
Water 0.00 334 79.8 100.0 2256 At $\text{37}\text{.}0º\text{C}$ (body temperature), the heat of vaporization ${L}_{v}$ for water is 2430 kJ/kg or 580 kcal/kg 539 At $\text{37.}0º\text{C}$ (body temperature), the heat of vaporization ${L}_{v}$ for water is 2430 kJ/kg or 580 kcal/kg
Sulfur 119 38.1 9.10 444.6 326 77.9
Lead 327 24.5 5.85 1750 871 208
Antimony 631 165 39.4 1440 561 134
Aluminum 660 380 90 2450 11400 2720
Silver 961 88.3 21.1 2193 2336 558
Gold 1063 64.5 15.4 2660 1578 377
Copper 1083 134 32.0 2595 5069 1211
Uranium 1133 84 20 3900 1900 454
Tungsten 3410 184 44 5900 4810 1150

Phase changes can have a tremendous stabilizing effect even on temperatures that are not near the melting and boiling points, because evaporation and condensation (conversion of a gas into a liquid state) occur even at temperatures below the boiling point. Take, for example, the fact that air temperatures in humid climates rarely go above $\text{35}\text{.}0º\text{C}$ , which is because most heat transfer goes into evaporating water into the air. Similarly, temperatures in humid weather rarely fall below the dew point because enormous heat is released when water vapor condenses.

We examine the effects of phase change more precisely by considering adding heat into a sample of ice at $-\text{20º}\text{C}$ ( [link] ). The temperature of the ice rises linearly, absorbing heat at a constant rate of until it reaches $0º\text{C}$ . Once at this temperature, the ice begins to melt until all the ice has melted, absorbing 79.8 cal/g of heat. The temperature remains constant at $0º\text{C}$ during this phase change. Once all the ice has melted, the temperature of the liquid water rises, absorbing heat at a new constant rate of . At $\text{100º}\text{C}$ , the water begins to boil and the temperature again remains constant while the water absorbs 539 cal/g of heat during this phase change. When all the liquid has become steam vapor, the temperature rises again, absorbing heat at a rate of .

Water can evaporate at temperatures below the boiling point. More energy is required than at the boiling point, because the kinetic energy of water molecules at temperatures below $\text{100º}\text{C}$ is less than that at $\text{100º}\text{C}$ , hence less energy is available from random thermal motions. Take, for example, the fact that, at body temperature, perspiration from the skin requires a heat input of 2428 kJ/kg, which is about 10 percent higher than the latent heat of vaporization at $\text{100º}\text{C}$ . This heat comes from the skin, and thus provides an effective cooling mechanism in hot weather. High humidity inhibits evaporation, so that body temperature might rise, leaving unevaporated sweat on your brow.

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using a micro-screw gauge,the thickness of a piece of a A4 white paper is measured to be 0.5+or-0.05 mm. If the length of the A4 paper is 26+or-0.2 cm, determine the volume of the A4 paper in: a). Cubic centimeters b). Cubic meters
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E=MC^2
study of matter and energy and an inter-relation between them.
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that's how the mass and energy are related in stationery frame
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Kinetic energy is the energy due to montion of waves,electrons,atoms, molecule,substances an object s.
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Energy, it always remains there in a physical system. it can only take the form either in motion (kinetic energy) or in rest (potential energy)
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