Water in an open container will eventually dry up. This process is called evaporation. Evaporation is a change of phase from liquid to gas that takes place at the surface of the liquid. Evaporation occurs when molecules of water gain enough kinetic energy to break free of their bonds with other water molecules and fly off. Because the molecules that gain kinetic energy leave the liquid (and molecules that lose kinetic energy remain in the liquid), evaporation cools liquids.

The opposite of evaporation is condensation-the changing of a gas into a liquid. When gas molecules near the surface of a liquid are attracted to the liquid, they strike the liquid with increased kinetic energy. They transmit this energy to the liquid as they join the liquid, warming it up. So condensation is a heating process.

Warm air rises. As it rises, it expands. As it expands, it cools. As it cools, water vapor in the air begins to stick together instead of bouncing off one another. If there are larger molecules or ions present, the water vapor condenses around them, forming clouds.

Under the right conditions, evaporation can take place beneath the surface of a liquid. When this happens, bubbles of vapor are formed, and they float to the surface to escape. Change of phase throughout a liquid rather than only at the surface is called boiling. Bubbles in the liquid can form only when the pressure of the vapor within the bubble is large enough to resist the pressure of the surrounding liquid. Bubbles that do not have enough vapor pressure burst. Boiling is a cooling process. Water is cooled by the boiling process sat the same rate energy is added to the water. This means a boiling liquid's temperature is constant.

Consider two people holding hands. Suppose they start jumping around randomly. It becomes harder for them to keep holding hands. If they jumped violently enough, keeping hold would be impossible. When a solid is heated, its molecules gain more and more kinetic energy, and they move faster. Eventually they move fast enough to break their bonds and the solid melts. Freezing is the opposite of this process. As energy is taken away from a liquid, its molecules slow down and form new bonds. The liquid becomes a solid. Ice molecules are crystals. Pressure on crystals will crush them into a liquid phase. When the pressure is removed, the liquid will refreeze. This phenomenon of melting under pressure and re-freezing when the pressure is reduced is called reduction. Consider an ice block with a wire through it. If you attach weights to the end of the wire, the pressure of the wire will melt the ice, passing through the ice. But the ice will refreeze after the wire passes through it. So the wire will pass through the ice, and the ice will look just as before.

If we continuously heat a solid it will change phase and become a liquid. The liquid will then change phase and become a gas. Energy is required for both liquefication and vaporization. Conversely, energy must be taken away from a substance to change its phase in the direction from gas to liquid to solid.

Theoretically, there is no upper limit to temperature. Matter changes from solid to liquid to gas to just atoms and finally to plasma as heat is added. There is a lower limit to temperature, though. Experiments have shown that a gas decreases its volume by 1/273 from 0 degrees Celsius. So if a substance decreases its temperature from 0 degrees Celsius to -273 degrees Celsius, its volume will be zero. A negative volume is impossible, so -273 degrees Celsius (0 degrees Kelvin) is the lowest possible temperature. It is called absolute zero.

When the law of energy conservation is enlarged to include heat, we call it the first law of thermodynamics. It states that when heat flows to or from a system, the system gains or loses an amount of energy equal to the amount of heat transferred. More specifically, the first law means that heat added to a system=increase in internal energy+external work done.

The process of compression or expansion of a gas sop that no heat enters or leaves the system is said to be adiabatic. Adiabatic processes can be achieved by either thermally insulating a system from its surroundings or by performing the process so rapidly that heat has little time to enter or leave. When work is done on a gas by adiabatically compressing it, the gas gains internal energy and becomes warmer. If a gas adiabatically expands, it loses internal energy and cools. Perform this experiment: with your mouth open, blow on your hand. You will feel warm air. Now pucker your lips and blow, so that the air leaving your mouth has to expand. The air will be noticeably cooler.

Consider a hot object placed next to a colder object in a insulated region. We know the hot object will give energy to the cold object and cool down. The cold object will heat up, and the two will continue this process until they are in thermal equilibrium. This transfer does not break the first law of thermodynamics, because there is the same amount of heat before the transfer of energy as after. What if the heat flowed from the cold object to the hot object? This would not violate the first law either. But it would violate the second law of thermodynamics: Heat will never of itself flow from a cold object to a hot object.

A heat engine is any device that changes internal energy into mechanical work. The basic idea behind a heat engine is that mechanical work can be obtained only when heat flows from a high temperature to a low temperature. In considering heat engines, we talk about reservoirs. Heat flows out of a hot reservoir to a cold reservoir called a sink. Heat engines always gain heat from a high temperature reservoir, convert some of the gained energy into mechanical work, and then expel the remaining energy as heat to the sink. Only some of the energy can be transferred to mechanical energy. The ideal efficiency of a heat engine depends on the difference in temperature between the hot reservoir and the cold sink. Ideal efficiency =T_hot-T_cold/T_hot.

The first law of thermodynamics states that energy can be neither created nor destroyed. The second law qualifies this by adding that in transformations, energy goes from more useful forms to less useful forms. Organized energy turns into disorganized energy. Push a crate across the floor and all your work will have turned into heat energy. This heat energy cannot be used to move the crate again. The quality of energy has been lowered. We see that the second law of thermodynamics can be stated another way: natural systems tend to proceed toward a state of greater disorder. Disorganized energy can only be changed to organized energy at the expense of some organizational input work.

Entropy is a measure of the amount of disorder in a system. The second law of thermodynamics tells us that entropy is always increasing. The laws of thermodynamics can be expressed as game rules: because you can't get more energy out of a system than you put in, you can't win; because you can't get out as much useful energy as you put in, you can't break even; and because entropy in the universe is always increasing, you can't get out of the game.