5.7: Heat Properties of Water

Of the properties of water, perhaps none are more important than the unique responses of water, ice, and water vapor to the application and removal of heat.

Heat Energy and Phase Changes

All substances can exist in three different phases: solid, liquid, or gas. The 3 phases of water and their properties are shown in Figure 5-8. In a solid, attractive forces between molecules (van der Waals forces, and hydrogen bonds if present) are strong enough to ensure that the molecules stay firmly fixed in place relative to each other, even though the molecules (or atoms in a pure element) vibrate. Heating the solid increases the strength of these vibrations. When the heat energy of each molecule is sufficient to overcome most of the attractive forces, the solid melts and becomes liquid. In a liquid, the molecules have enough energy to vibrate, rotate, and translate (temporarily move about in relation to each other). However, they do not have enough energy to escape completely from the attractive forces of their neighbors. If more heat is added to a liquid, each molecule eventually has enough energy to break free of the attractive forces. The compound then becomes a gas, in which each molecule is free to move about by itself. If heat is removed, the process is reversed; gas becomes liquid, and then solid, as heat is progressively lost.

Figure 5-8. Characteristics of solid, liquid, and gaseous water.

Table 5-8. Characteristics of Water

	Solid (ice)	Liquid (water)	Gas (steam)
Motion	Limited to vibrations in almost fixed positions	Molecules moving around, can flow	Complete freedom of motion, can flow
Shape	Fixed, definite, rigid	Same as that of the lower part of the container, variable	Same as that of the closed container, highly variable
Volume	Almost constant, variest only slightly with temperature	Almost constant, varies slightly with temperature	Same as that of the closed container, variable

Molecules of different chemical compounds have van der Waals forces of different strengths. As a result, each compound requires a certain characteristic quantity of heat energy to convert from solid to liquid or from liquid to gas. The stronger the attractive force between the molecules, the more heat energy each molecule must have to break free of this force. Therefore, the temperature at which a solid melts (the melting, or freezing, point) and the temperature at which a liquid vaporizes (the boiling, or condensation, point) increase as the van der Waals attractive force increases.

Freezing and Boiling Points

To convert solid water (ice) to liquid water and liquid water to gaseous water (water vapor), the heat energy supplied must overcome both the van der Waals attractive force and the much stronger attractive force of the hydrogen bond. Therefore, the amount of heat energy that each molecule of water must have to become free to rotate and move in relation to adjacent molecules (that is, to change from ice to water) is much greater than it would be if the hydrogen bonds were not present. Similarly, the amount of heat energy that each water molecule must have to free itself completely and enter the gaseous state is much greater than it would be without the hydrogen bond.

The extra energy needed to break the hydrogen bond causes water’s freezing point and boiling point temperatures to be anomalously high. If there were no hydrogen bond, water would freeze at about –90°C and boil at about –70°C, and all water on the Earth would be gaseous. There would be no oceans or life as we know it. Comparison of the boiling and freezing points of water with those of other hydrogen compounds that are formed with elements that, like oxygen, have two electrons short of a full outer electron shell illustrates the anomalous nature of water (Fig. 5-9). Molecules in which hydrogen is combined with elements whose atoms are similar in size to oxygen but have a different number of outer-shell electrons (hydrogen fluoride, HF; ammonia, NH₃) are also polar. They, too, have hydrogen bonds and anomalously high melting and boiling points. As discussed later in this chapter, adding salt to water raises the boiling point and lowers the freezing point.

Figure 5-9. Freezing and boiling points of hydrogen compounds. The lines on the chart connect compounds of hydrogen with elements that have the same number of electrons in their outer shell but a different number of filled inner shells. If there were no hydrogen bonds, both the boiling point and the melting point of those compounds would be expected to increase progressively along these lines as the number of filled shells increased. However, water (H₂O), hydrogen fluoride (HF), and ammonia (NH₃) all have strong hydrogen bonds between their polar molecules, whereas there is no appreciable hydrogen bond in the compounds of hydrogen that are formed with sulfur (S), selenium (Se), tellurium (Te), chlorine (Cl), bromine (Br), iodine (I), phosphorus (P), arsenic (As), and antimony (Sb). In this latter group of compounds, the hydrogen bond is virtually absent because the atom combined with the hydrogen is much larger, and, therefore, the polarity of the molecules is much less. Since the hydrogen bond must be broken to convert solid to liquid or liquid to gas, the boiling points (a) and melting (freezing) points (b) of water, hydrogen fluoride, and ammonia are much higher than would be expected if they were not held by hydrogen bonds.

Heat Capacity and Latent Heat

In addition to the freezing and boiling points, the numerical values of three other related heat properties of water are anomalously high because of the hydrogen bond: the heat capacity, the latent heat of fusion, and the latent heat of vaporization. These properties express the quantity of heat per specified quantity of a substance needed to raise the temperature, to convert solid to liquid, and to convert liquid to gas, respectively. Heat is a form of energy, and as such, it is quantified by a unit of energy. The SI unit of energy is the joule (abbreviated J). This SI unit is not yet in common usage in the U.S., although it is widely used elsewhere. Many students will be more familiar with the calorie (abbreviated cal) widely used as the unit of energy in the U.S. In this text, we will use the calorie, and we will also identify the equivalent value in joules.

To understand the concepts of heat capacity and latent heat, we can consider the sequence of events that occurs when we add heat to ice (Fig. 5-10). Within the solid ice, the individual molecules vibrate, but the vibrations are suppressed by the hydrogen bonds. Adding heat energy to the ice increases the intensity of vibration of each water molecule and increases the temperature of the ice. The temperature of 1 g of ice is increased by 1°C when approximately 0.5 cal (2.1 J) is added. As heat continues to be added, the temperature continues to rise until the ice reaches its melting point. Heat that raises (or lowers) the temperature of a substance to which it is added (or removed) is called sensible heat. The quantity of heat needed to increase the temperature of a specified quantity of a substance is called the “heat capacity.” For ice, the heat capacity is approximately 0.5 cal•g^–1•°C^–1 (2.1 J•g^–1•°C^–1). That is, one-half of a calorie (2.1 J) is needed to raise the temperature of 1 g of ice by 1°C.

Figure 5-10. Water’s latent heat of fusion (melting) is the amount of heat that must be added to convert ice to water at the freezing (melting) point temperature. Water’s latent heat of vaporization is the amount of heat that must be added to water to convert water to water vapor at boiling point temperature. Both of these are high because of the need to provide energy to overcome the hydrogen bond attraction. As latent heat is added to change the phase, there is no change in temperature. In the reverse processes, the same amounts of latent heat that were added are released when water vapor condenses and when water freezes. The heat capacity of water is the amount of heat required to raise the temperature of 1 g of liquid water by 1°C and is measured in cal•g^–1•°C^–1 (J•g^–1•°C^–1). The same amount of heat that was absorbed during warming is released as water cools.

After the ice has reached the melting point temperature, the temperature of the solid-liquid (ice-water) mixture does not increase again until an additional 80 calories (334 J) of heat energy has been added for each gram of ice (Fig. 5-10). That is, as heat is added, the ice is progressively converted to water. Heat added to a substance that does not raise its temperature but instead changes its state (solid to liquid or liquid to gas) is called “latent” heat. The heat added to melt a specified quantity of ice is the latent heat of fusion (melting). The term latent is used because the heat is stored in the molecules of the liquid water and is released when the water is refrozen (fused). Hence, the latent heat of fusion of ice is 80 cal•g^–1 (334 J•g^–1).

If we continue to add heat after all the ice is converted to water, the heat is again taken up as sensible heat and the temperature of the water rises. The heat capacity of water is about twice that of ice, or 1 cal•g^–1•°C^–1. That is, 1 cal (4.2 J) of heat energy must be added to raise the temperature of 1 g of water by 1°C.

After 100 cal•g–1 (418 J•g–1) of heat energy has been added to the water (starting as melted ice water at 0°C), the water temperature reaches the boiling point. With the continued addition of heat, the temperature does not rise again until an additional 540 cal•g^–1 (2260 J•g^–1) of heat energy has been added and all the water has been converted to a gas. Hence, the latent heat of vaporization of water is 540 cal•g^–1 (2260 J•g^–1). Finally, if we add more heat to the gaseous water (water vapor), its temperature rises by about 1°C for each additional 0.5 cal•g^–1 (2.1 J•g^–1), so the heat capacity of water vapor is about 0.5 cal•g^–1•°C^–1 (2.1 J•g^–1•°C^–1).

The latent heats and heat capacity of water are very high primarily because hydrogen bonds are stronger than van der Waals forces. More heat energy is needed to overcome the attraction between molecules in solid, liquid, and gaseous forms of water than in substances whose molecules are bound only by van der Waals forces. Consequently, the heat capacity of liquid water is the highest of all liquids other than liquid ammonia and is higher than that of all solids. The latent heat of fusion of water is the highest of all substances other than ammonia; and the latent heat of vaporization of water is the highest of all known substances.

Implications of the High Heat Capacity and Latent Heats of Water

The anomalously high heat capacity and latent heats of water have many implications. In our everyday experience, we rely on the high latent heat of fusion to keep our iced drinks cold. When we add ice cubes to a cold drink, the drink remains cold for a long time, until all the ice has melted. The drink, a mixture of ice and water, remains cold as it gains heat from its surroundings, until it has absorbed 80 calories (334 J) for each gram of ice, converting the ice to water. Subsequently, when all the ice has melted, the drink warms much faster because each calorie (each 4.2 J) gained can raise the temperature of 1 g of water by 1°C.

The high heat capacity of water is illustrated by the behavior of hot drinks. A hot drink cools much more slowly than, for example, an empty cup or glass taken out of a hot dishwasher. The reason is that more heat must be lost per unit weight of water than of any other substance (other than liquid ammonia) to lower its temperature by a given number of degrees.

The high heat capacity of water allows large amounts of heat energy from the sun to be stored in ocean waters without causing much of a temperature change. The heat is released to the atmosphere when atmospheric temperatures fall. Because water has this heat-buffering capability, coastal locations have milder climates than inland locations (Chap. 7). The heat-buffering capacity of the oceans and the transfer of latent heat between ice, oceans, and atmosphere are important to global climatic control and to studies of the effects of releases of gases that contribute to the greenhouse effect (Chap. 7).

In the polar ocean and coastal regions, water’s high latent heat of fusion acts to control the air and water temperatures in much the same way that ice cubes cool an iced drink. During winter, latent heat is released to the atmosphere as water freezes to form more sea ice. Conversely, in summer, heat added to the polar regions is used to melt the ice. Because the heat lost or gained is latent heat, not sensible heat, the ocean surface water and sea ice remain at or close to the freezing point throughout the year. Because air temperatures are partially controlled by ocean temperatures, the annual climatic temperature ranges in polar ocean regions are small (Chap. 7, CC5).

Evaporation

Liquids, such as water, can be converted to a gas at temperatures below their boiling point by a process known as “evaporation.” For example, puddles of cold water can evaporate from the ground after a rainstorm. Although the average energy level of water molecules remains the same unless there is a change of temperature, the energy level of individual molecules varies around this average. Some molecules that temporarily possess higher amounts of energy can overcome their hydrogen bonds and escape completely from the liquid phase (evaporate).

Water evaporating from the oceans carries its latent heat of vaporization into the atmosphere and releases it when the water subsequently condenses as rain. Because the latent heat of vaporization is high, water can transport large quantities of heat energy from ocean to atmosphere, where that energy can be redistributed geographically. For example, areas of the Earth where evaporation is slow and precipitation is high tend to be warmed, while areas where evaporation is fast tend to be cooled (Chap. 7, CC5).

Water’s high latent heat of vaporization facilitates cooking and can be observed in a kitchen. If we heat a pot of water on a stove, the water can be brought to a boil fairly quickly, but then, even if we do not change the heat setting, it takes much longer for the pot to boil dry. It takes only 100 cal•g^–1 (418 J•g^–1) to heat water from its freezing point to its boiling point, but more than five times as much, 540 cal•g^–1 (2260 J•g^–1), to vaporize or convert the water to steam. Water’s high latent heat of vaporization is also apparent when we bake food for substantial lengths of time at temperatures well above 100°C and find that the food is still moist.

Because evaporating molecules gain the energy to escape by colliding with molecules that remain behind in the liquid, the average energy level (temperature) of the water molecules that remain behind is reduced. This is why we feel colder when we first climb out of a swimming pool or shower. While we are wet, water molecules evaporate using heat energy gained from our skin, and the air and leaving behind colder skin. After toweling off, we have less water on our skin and the amount of evaporation is reduced. Another example of this phenomenon is the wind chill factor used in weather forecasts. Increasing the speed of wind blowing across a wet or damp surface, such as the skin, increases the evaporation rate and thus increases the rate of surface cooling.

At temperatures below the boiling point, the average water molecule has less energy than it does at the boiling point. Therefore, more heat must be supplied to evaporate each molecule of water than is needed to vaporize it at the boiling point, and the latent heat of vaporization increases with decreasing temperature. For example, the latent heat of vaporization at 20°C is 585 cal•g^–1 (2449 J•g^–1) which is 45 cal•g^–1 (189 J•g^–1) higher than at 100°C.