8.3: The Three States of Water

In our everyday experience, heating or cooling something results in a straightforward change in its temperature. But for water, heating and cooling may bring about changes in its physical state; that is, whether it’s a solid, liquid, or gas.

By definition, solids maintain their own form. Liquids take the shape of the containers in which they are held. Gases also take the shape of their container, but they will expand to fill the container, while liquids will not. A fourth state of matter, plasma, consists of free-moving electrons and ions. Plasmas occur primarily in the upper atmosphere and outer space, although lightning and St. Elmo’s fire—an electrical discharge from the masts of ships—also qualify as plasmas. We won’t deal with plasmas here.

As far as we know, water is the only substance on Earth to occur in all three physical states at the same time:

• ice, the solid form of water
• liquid water, simply referred to as water
• water vapor, the gaseous form of water

Each of these states plays an important role in Earth and ocean processes. The transformations between these states occur at specific temperatures. However, as we shall see, it’s a little more complicated than heating water to its boiling point or cooling it to its freezing point.

Specific Heat: Heat We Can Sense

Water’s enormous importance in global heat exchange comes from its ability to absorb a high amount of heat. In fact, water ranks first in heat-absorbing capability among all liquids at environmental temperatures.

Scientists define the heat-absorbing abilities of a substance as specific heat, the amount of heat required to change the temperature of a given mass of a substance by 1 degree Celsius (°C; scientists use the symbol ° as a shorthand for degrees). Because specific heat changes the temperature of a liquid, we sometimes refer to it as sensible heat, heat that can be detected by instruments (or touch). We use the Celsius scale here to make calculations simple. By definition, the specific heat of water is one calorie per gram per degree Celsius. (Note that “calorie” is spelled with a little c to differentiate it from “Calorie,” used for food, which is 1,000 times greater but often also styled as lowercase). In an honest-to-goodness physics course, you would use official scientific units—the International System of Units, or SI units, also known as the metric system—in which case you would learn that the specific heat of water is 4.186 joules per gram per degree Celsius. But given that it’s hard enough for nonmajors (and even beginning science majors) to grasp the math, we’ll stick to calories. My apologies to the physicists.

To help you visualize specific heat, imagine a spoon holding one cubic centimeter of water—roughly a fifth of a teaspoon—which has a weight of one gram. Warming it (carefully) over a flame and recording its temperature with a thermometer, a temperature increase of 1°C would indicate that you have added one calorie of heat. Another degree increase, another calorie. For each degree of temperature rise, you’ve added one calorie.

If you cool your gram of water, you can watch the temperature fall. One calorie of heat will be lost for every 1°C drop in temperature. If you start with a gram of water at room temperature—22°C (72 degrees Fahrenheit, °F)—and remove one calorie, its temperature will be 21°C. Remove another calorie, and its temperature will be 20°C. The math follows easily if you understand what is happening: remove heat, lower the temperature; add heat, raise the temperature.

Latent Heat: “Hidden” Heat

Specific heat applies to changes in temperature that occur when liquid water gains or loses heat. But what about ice and water vapor? How do the solid and gaseous forms of water interact with heat?

Here’s where water gets really interesting. While the freezing and boiling points of water are set at specific temperatures—0°C (32°F) and 100°C (212°F), respectively—water doesn’t just instantaneously transform into a different state. Additional heat must be removed or added to “rearrange” water molecules into their solid or gaseous form. Because the heat involved in transforming water into its solid or gaseous state does not register as a change in temperature, these forms of heat are referred to as latent heat. The term “latent” means “hidden,” so we can think of latent heat as a kind of “invisible” heat, undetectable with a thermometer.

Two kinds of latent heat must be considered: the latent heat of fusion—the heat required to transform liquid water to a solid and vice versa—and the latent heat of vaporization—the heat required to transform liquid water to water vapor and back again. Before ice or water vapor can form, heat must be removed or added, respectively, to rearrange the water molecules into their new physical states, a solid or a gas.

What’s important to remember is that to make ice—to turn liquid water into solid water—not only must you remove enough calories to lower the temperature to the freezing point, but once you reach freezing, you must remove even more calories of heat to change the liquid to a solid. Similarly, to boil water—to turn liquid water into water vapor—you must add enough heat to raise its temperature to the boiling point and then add even more heat to turn the liquid water into water vapor.

How much heat? A lot! The latent heat of fusion for water is 80 calories per gram. The latent heat of vaporization is even more extraordinary: 540 calories per gram. That’s 5.4 times the amount of heat required to take the liquid temperature of water from 0°C to 100°C. That means you have to remove 80 calories of additional heat after reaching 0°C (32°F) to turn one gram of liquid water into ice. At the other end of the temperature scale, you would need to add 540 calories of additional heat once you reach the boiling point to turn one gram of liquid water into water vapor.

Let’s return to our fifth of a teaspoon of room temperature water. With a temperature of 22°C, you would need to remove 22 calories of heat to lower its temperature to 0°C (32°F). But to turn it into ice, you would need to remove another 80 calories. So, in total, you would need to remove 22+80 calories, or 102 calories, to turn one gram of room temperature water into ice.

To go the other direction, you would need to add 78 calories to reach the boiling point (100 degrees minus 22 degrees is a difference of 78 degrees). Then, you’ll need to add 540 more calories to transform that gram of water into water vapor. That’s a total of 78+540 calories, or 618 calories, to turn one gram of room temperature water into ice.

Why is latent heat important? When tropical water vapor carried by winds condenses at higher latitudes, it releases heat, keeping temperate and polar regions warm. Latent heat proves important in the formation and intensification of storms and hurricanes. It’s integral to the integrity of polar ice caps, which can absorb a tremendous amount of heat. Once gone, the heat that once went into melting ice goes directly into the atmosphere. Finally, latent heat cools your body. When your body heats up during exercise, you sweat. When those sweat molecules gain enough heat, they vaporize, carrying latent heat with them. Heat is transferred more effectively from your body to the surrounding air because of latent heat transfer during sweating. Pretty cool, huh?