Water has many properties favorable to life

Water is abundant over most of Earth’s surface. Within the temperature range organisms usually encounter, it is liquid. Because water also has an immense capacity to dissolve inorganic compounds, it is an excellent medium for the chemical processes of living systems. In fact, it is hard to imagine a form of life that could exist without water. In this section we will look at how water makes life possible, including water’s thermal properties, its density and viscosity, and its function as a solvent for inorganic nutrients.

Thermal Properties of Water

On Earth, water can be found as a solid (ice), as a liquid, and as a gas (water vapor). No other common substance is liquid under most conditions at Earth’s surface. Pure water—water not containing any dissolved minerals or other compounds—becomes a solid below 0°C and becomes a gas above 100°C at sea level. At higher elevations, the freezing point of water changes very little, but the boiling point can be several degrees lower.

When water contains dissolved compounds, such as salts, its freezing temperature drops below 0°C. This is why road salts are applied to ice- and snow-covered roads; the salts allow the ice and snow to melt at a lower temperature than they otherwise would. Dissolved compounds also raise the boiling point of water above 100°C.

Water stays liquid over a broad range of temperatures, from 0°C to 100°C. The temperature of water remains relatively steady even when heat is removed or added rapidly, as can happen at the air–water interface or at an organism’s surface, such as on the surface of a whale in the ocean. This is because water has a high specific heat, which is the amount of heat required to increase its temperature by 1°C.

Water also resists changing from one state to another. For example, raising the temperature of 1 kg of liquid water by 1°C requires the addition of 1 Calorie of heat. However, converting 1 kg of liquid water into water vapor requires the addition of 540 Calories of heat. Similarly, lowering the temperature of 1 kg of liquid water by 1°C requires the removal of 1 Calorie of heat, but converting that amount of liquid water to ice requires the removal of 80 Calories of heat. In short, liquid water is very resistant to changing states. This resistance helps prevent large bodies of water from freezing solid during winter. In addition, because water transfers heat rapidly, heat tends to spread evenly throughout a body of water, which also slows local changes in temperature.

Another curious, yet fortunate, thermal property of water is the way it changes density with changes in temperature. Water achieves its highest density—that is, its molecules are the most densely packed together—at 4°C. Above and below 4°C, the water molecules are less tightly packed and the water becomes less dense. Below 0°C, pure water is converted into ice, which is less dense than liquid water as you can see in Figure 2.1. As a result of its lower density, ice floats on the surface of liquid water. This means that lakes experiencing cold winters will generally have a layer of 4°C water at their bottom. Above this layer will be water that is less than 4°C and on top of that will be a layer of ice.

Figure 2.1 The density of water. As water cools, the molecules contract and become more dense. Below 4°C, they begin to expand and becomes less dense. Below 0°C, pure water is converted into ice that is even less dense. As a result of its lower density, ice floats on the surface of liquid water.

Photo by Zoonar/Christa Kurtz/age fotostock.

Water’s unusual thermal properties are especially important to aquatic plants and animals. In large water bodies containing fresh water, such as lakes, the bottom of the lake does not freeze, in part because of the insulation that ice provides from very cold air temperatures. The salt in ocean water lowers the freezing point of the water well below 0°C, which prevents oceans from freezing. In both cases, the available water provides a refuge for organisms during periods of cold temperatures.

Density and Viscosity of Water

The adaptations of aquatic organisms often exploit the density and viscosity of water. For example, animals and plants are comprised of bone, proteins, and other materials that are somewhat denser than salt water, and much denser than fresh water. However, organisms can also contain fats and oils that are less dense than water. In some cases, they also possess pockets of air, such as the lungs of the whales that were described at the beginning of this chapter. The combination of the materials that compose an animal’s body and the presence of air pockets determine whether an organism will float or sink in water.

For those organisms that are denser than their surrounding water, a variety of adaptations can either reduce an organism’s density or retard its rate of sinking. For example, many fish species have a gas-filled swim bladder that can adjust in size to make the density of the fish’s body equal to that of the surrounding water. Human divers use this same concept when they wear air-filled inflatable vests to help match the density of the water. Divers can add air to the vest, which allows them to float on the surface of the water. Alternatively, divers can release some of the air to match the density of the water so that they neither float nor sink, or they can let even more air out so that they sink in the water. Some large kelps, like those we saw in Figure 1.8, have gas-filled bulbs that cause their leaf-like blades to float up into the sunlit surface waters. The whales discussed earlier become buoyant when they take a breath of air but a slow release of air bubbles will help them sink to a particular depth. At the other end of the size spectrum, many of the microscopic unicellular algae that float in great numbers in the surface waters of lakes and oceans use droplets of oil as flotation devices (Figure 2.2). Because oils are less dense than water, the algae can use the oil droplets to help offset their natural tendency to sink.

Figure 2.2 Adaptation to water density. These algae (Cyclotella cryptica) are able to float near the water’s surface by using oil droplets that have a lower density than water.

Photo by Bigelow Laboratory National Center for Marine Algae and Microbiota.

Aquatic organisms also possess adaptations to deal with the high viscosity of water. Viscosity is technically defined as the resistance of a fluid to being deformed by a stress, but we can think of viscosity as the thickness of a fluid that causes objects to encounter resistance as they move through it. In response to living in water, fast-moving aquatic animals such as fish, penguins, and whales have evolved highly streamlined shapes that reduce the drag caused by the high viscosity of water and other factors (Figure 2.3). The viscosity of water is higher in cold water than in warm water, which can make swimming in cold water more difficult. Movement in water is even more difficult for smaller animals. However, the same high viscosity that impedes the progress of tiny organisms as they swim in the water also impedes them from sinking. Because these organisms are slightly denser than water, they are prone to sinking due to the force of gravity. To take advantage of water’s viscosity, many tiny marine animals have evolved long, filamentous appendages that cause greater drag in the water. The appendages function like a parachute slowing the fall of a body through air (Figure 2.4).

Figure 2.3 Streamlined shape. Large, fast-moving aquatic organisms like the barracuda (Sphyraena barracuda) have evolved highly streamlined shapes to help them move through the highly viscous water.

Photo by George Grall/National Geographic Stock.

Figure 2.4 Adaptation to water viscosity. Some small aquatic organisms exploit the high viscosity of water by evolving large appendages, such as the antenna and feathery projections of this tiny marine crustacean. These appendages help slow down movement through the viscous water and thereby retard sinking.

Photo by Solvin Zankl/naturepl.com.

Dissolved Inorganic Nutrients

Both aquatic and terrestrial organisms require a variety of nutrients to build necessary biological structures and maintain life processes. Large amounts of hydrogen, carbon, and oxygen are necessary for building most compounds found in organisms. Nitrogen, phosphorus, and sulfur are used in building proteins, nucleic acids, phospholipids, and bones. Other major nutrients—including potassium, calcium, magnesium, and iron—play important roles as solutes and as structural components of bones, woody plant cells, enzymes, and chlorophyll. Certain organisms need additional minor nutrients. For example, diatoms are a group of algae that need silica to construct their glassy shells (Figure 2.5). Similarly, some species of bacteria require the element molybdenum, which makes up part of the enzyme they use to convert nitrogen from the atmosphere (N₂) into ammonia (NH₃).

Figure 2.5 Use of inorganic nutrients. Diatoms, such as this species of Arachnoidiscus, are a type of algae that require minor nutrients such as silica to build a hard, glassy shell. Image is magnified 175 times.

Photo by Steve Gschmeissner/Science Source.

The Solubility of Minerals

Water is a powerful solvent with an impressive capacity to dissolve substances, which makes them accessible to living systems. Because of this property, water also provides a medium in which substances can react chemically to form new compounds.

Water acts as a solvent because of its molecular structure. As you can see in Figure 2.6, water molecules consist of an oxygen atom in the middle and two hydrogen atoms connected in a V-shaped arrangement. This arrangement results in an unequal sharing of electrons: the oxygen end of the water molecule has a slightly negative charge and the hydrogen end has a slightly positive charge. When the two ends of the molecule possess opposite charges, we say that the molecule is polar. Water is a polar molecule: the negative oxygen end of one water molecule is strongly attracted to the positive hydrogen end of another. These forces of attraction are known as hydrogen bonds.

Figure 2.6 Water molecules. Because of the configuration of water molecules, they are negatively charged on the oxygen end and positively charged on the hydrogen end. The attractive forces of these opposite charges, known as hydrogen bonds, allow water molecules to be attracted to each other and to the charged ions of other compounds such as salts and sugars.

The polar nature of water molecules also allows them to be attracted to other polar compounds. Some solid compounds consist of electrically charged atoms, or groups of atoms, called ions. For example, common table salt—sodium chloride (NaCl)—contains positively charged sodium ions (Na⁺) and negatively charged chloride ions (Cl⁻). In their solid form, these ions are arranged in a crystal lattice. In water, however, the charged sodium and chloride ions are attracted by the charges of the water molecules. As shown in Figure 2.7, the attraction of these ions to water molecules is stronger than the attraction that holds the crystal together. As a result, when sodium chloride is added to water, its crystal lattice breaks apart, and water molecules surround the salt ions. In other words, when you put salt in water, it dissolves. This solvent ability of water is not restricted to ionic compounds such as salts; it occurs with any polar compound, including the various types of sugars that organisms commonly use. In contrast, water is not a good solvent for oils and fats because they are nonpolar compounds.

Figure 2.7 Dissolving ions in water. Because water molecules have negative and positive ends, they attract the negatively and positively charged ions, such as the sodium and chlorine ions found in sodium chloride. The forces of attraction to water molecules are stronger than the forces of attraction within the crystal, so the ions separate and become surrounded by water molecules.

The solvent properties of water explain the presence of minerals in streams, rivers, lakes, and oceans. When water vapor in the atmosphere condenses to form clouds, the condensed water in the atmosphere is nearly pure. However, as it falls back to Earth as rain or snow, water acquires some minerals from dust particles in the atmosphere. Precipitation that hits land comes into contact with rocks and soils, and it dissolves some of the minerals in them. These additional minerals are carried toward the ocean.

Water in most lakes and rivers contains a dissolved mineral concentration of 0.01 to 0.02 percent whereas ocean water contains a dissolved mineral concentration of 3.4 percent. Oceans have much higher concentrations of dissolved minerals because mineral-laden water continuously flows in from streams and rivers, and the constant evaporation from the ocean’s surface removes pure water, leaving the minerals behind. Over billions of years, this process has caused increased concentrations of minerals in the oceans.

Every mineral has an upper limit of solubility in water, known as saturation. This limit generally increases with higher temperatures. After a mineral achieves saturation, water cannot hold any more and the mineral precipitates out of solution. For some minerals, such as sodium, ocean concentrations are far below saturation. Most of the sodium that is washed into ocean basins remains dissolved and its concentration in seawater continues to increase. In contrast, the concentrations of other minerals in the oceans commonly exceed their saturation concentrations. For example, calcium ions (Ca⁺²) readily combine with dissolved CO₂ to form calcium carbonate (CaCO₃), and calcium carbonate has a low solubility in ocean water. Over millions of years, the excess calcium carbonate that has washed into the oceans from streams and rivers has subsequently precipitated out of the water. This precipitated calcium carbonate, combined with the calcium carbonate from the bodies of countless tiny marine organisms, has resulted in massive limestone sediments (Figure 2.8). Today, these sediments are important sources of limestone for construction applications such as stone blocks and concrete, for agricultural uses such as fertilizer, and for numerous industrial processes.

Figure 2.8 The formation of limestone sediments. The continuous addition of calcium minerals into oceans from streams and rivers causes calcium to combine with CO₂ to become calcium carbonate. Because calcium carbonate is not very soluble in water, it precipitates out of the water to form massive sediments over millions of years. This site of limestone sediments located in Victoria, Australia, was once under water but is now above the water due to changes in ocean depth.

Photo by Phillip Hayson/Photo Researchers, Inc.

Hydrogen Ions

Among dissolved substances in water, hydrogen ions (H⁺) deserve special mention because they are extremely reactive with other compounds. In pure water, a small fraction of the water molecules (H₂O) break apart into their hydrogen (H⁺) and hydroxide (OH ⁻) ions. The concentration of hydrogen ions in a solution is referred to as its acidity. Acidity is commonly measured as pH, which is defined as the negative logarithm of hydrogen ion concentration (as measured in moles/L, where 1 mole = 6.02 × 10²³ molecules):

As you can see in the pH scale shown in Figure 2.9, water containing a high concentration of hydrogen ions has a low pH value, whereas water containing a low concentration of hydrogen ions has a high pH value. Therefore, we categorize water with low pH values as acidic, water with a mid-range value of 7 as neutral, and water with a high pH value as basic or alkaline. As we will see, natural rain or snow can vary a great deal in pH, depending on the presence of different chemical compounds in the atmosphere that affect the pH value.

Figure 2.9 The relationship between pH and hydrogen ion concentration in water. The pH scale of hydrogen ion concentration extends from 0 (highly acidic) to 14 (highly alkaline). The pH of rainfall can vary a great deal around the world.

Hydrogen ions, because of their high reactivity, dissolve minerals from rocks and soils, enhancing the natural solvent properties of water. For example, in the presence of hydrogen ions, the calcium carbonate found in limestone dissolves readily, according to the following chemical equation:

H⁺ + CaCO₃ → Ca²⁺ HCO₃⁻

Calcium ions are important to life processes, and their presence at high concentrations is vital to organisms such as snails that form shells made of calcium carbonate. Indeed, mollusks are less abundant in streams and lakes that are low in calcium. Therefore, hydrogen ions are essential for making certain nutrients available for life processes. At high concentrations, however, hydrogen ions negatively affect the activities of most enzymes. In addition, high concentrations of hydrogen ions cause many heavy metals to begin dissolving in water. These heavy metals, including arsenic, cadmium, and mercury, are highly toxic to most aquatic organisms.

The normal range of pH of lakes, streams, and wetlands is between 5 and 9. However, some bodies of water can have even lower pH values. Sometimes lower pH conditions have a natural cause. Bogs, for example, are aquatic habitats with vegetation such as sphagnum mosses that release H⁺ ions into the water and thereby make the water more acidic and unsuitable for many other species of plants.

Other bodies of water have low pH as a result of human influences. For example the release of sulfur dioxide (SO₂) and nitrogen dioxide (NO₂) from coal-powered industrial plants became a major environmental issue in the 1960s. At that time, ecologists in Russia, China, northern Europe, the United States, and Canada began to notice that many bodies of water were becoming more acidic and less hospitable to numerous species of fish and other aquatic organisms. They also noticed that trees were dying, particularly in the spruce-fir forests that existed at high elevations in these regions of the world.

It turned out that the areas with more acidic bodies of water and dying trees were all far downwind of industrial areas with coal-powered factories that had tall smokestacks. Years of data collection revealed that the sulfur dioxide and nitrogen dioxide emitted into the atmosphere by these smokestacks were converted to sulfuric acid and nitric acid in the atmosphere. These acids were transported downwind and then came back down to Earth as acid deposition, also known as acid rain. The acid deposition occurred in two forms: as gases and particles that stuck to plants and soil, a form called dry acid deposition, and as rain and snow, a form called wet acid deposition. Acid deposition lowered the pH of precipitation and, as a result, water of unusually low pH was entering streams, lakes, and forests. Most aquatic species cannot tolerate water with a pH lower than 4, so these bodies of water became toxic to many aquatic organisms, which included insects and fish.

In forests, acid deposition has several effects. First, it leaches the calcium out of the needles of conifer trees such as spruce. It also causes increased leaching of soil nutrients that trees require, including calcium, magnesium, and potassium. Finally, acid deposition causes aluminum to dissolve in the water. Although aluminum naturally occurs in the soil, it is typically not in a form that is available to plants. Dissolved aluminum can negatively affect a plant’s ability to take up nutrients. Collectively, these effects of acid deposition make trees more susceptible to the harmful effects of natural stressors that include drought, diseases, and extreme temperatures. In short, while the trees do not die directly from acid deposition they become more susceptible to other causes of death. Because acid deposition interacts with so many other natural stressors, scientists recognize that acid deposition has contributed to the death of trees in North America, Europe, and Asia. However, the complexity of the interactive effects has made it difficult to accurately estimate how much of the observed tree death is directly attributable to acid deposition.

Once scientists understood the causes and consequences of acid deposition, they began to explore solutions. In the United States, legislation required the installation of smokestack scrubbers that force smokestack gases through a slurry of limestone and water, which remove the gases. The use of these scrubbers has caused a major reduction in the amount of acidic compounds going into the atmosphere. The U.S. Environmental Protection Agency (EPA) reports that from 1980 to 2010, emissions of sulfur dioxide declined by 79 percent. At the end of this chapter, in Connecting the Concepts, “The Decline of Coral Reefs,” we discuss another example of how understanding environmental problems related to pH can help us develop effective solutions.