Chapters 2-4 Study Questions
CHAPTER 3
Smith, T. M., & Smith, R. L. (2015). Elements of Ecology (9th ed.). Boston, MA: Pearson.
3.1 Water Cycles between Earth and the Atmosphere
All marine and freshwater aquatic environments are linked, either directly or indirectly, as components of the water cycle (also referred to as the hydrologic cycle; Figure 3.1)—the process by which water travels in a sequence from the air to Earth and returns to the atmosphere.
Solar radiation, which heats Earth’s atmosphere and provides energy for the evaporation of water, is the driving force behind the water cycle (see Chapter 2). Precipitation sets the water cycle in motion. Water vapor, circulating in the atmosphere, eventually falls in some form of precipitation. Some of the water falls directly on the soil and bodies of water. Some is intercepted by vegetation, dead organic matter on the ground, and urban structures and streets in a process known as interception.
Because of interception, which can be considerable, various amounts of water never infiltrate the ground but evaporate directly back to the atmosphere. Precipitation that reaches the soil moves into the ground by infiltration. The rate of infiltration depends on the type of soil, slope, vegetation, and intensity of the precipitation (see Section 4.8). During heavy rains when the soil is saturated, excess water flows across the surface of the ground as surface runoff or overland flow. At places, it concentrates into depressions and gullies, and the flow changes from sheet to channelized flow—a process that can be observed on city streets as water moves across the pavement into gutters. Because of low infiltration, runoff from urban areas might be as much as 85 percent of the precipitation.
Some water entering the soil seeps down to an impervious layer of clay or rock to collect as groundwater (see Figure 3.1). From there, water finds its way into springs and streams. Streams coalesce into rivers as they follow the topography of the landscape. In basins and floodplains, lakes and wetlands form. Rivers eventually flow to the coast, forming the transition from freshwater to marine environments.
Water remaining on the surface of the ground, in the upper layers of the soil, and collected on the surface of vegetation—as well as water in the surface layers of streams, lakes, and oceans—returns to the atmosphere by evaporation. The rate of evaporation is governed by how much water vapor is in the air relative to the saturation vapor pressure (relative humidity; see Section 2.5). Plants cause additional water loss from the soil. Through their roots, they take in water from the soil and lose it through their leaves and other organs in a process called transpiration. Transpiration is the evaporation of water from internal surfaces of leaves, stems, and other living parts (see Chapter 6). The total amount of evaporating water from the surfaces of the ground and vegetation (surface evaporation plus transpiration) is called evapotranspiration.
Figure 3.2 is a diagram of the global water cycle showing the various reservoirs (bodies of water) and fluxes (exchanges between reservoirs). The total volume of water on Earth is approximately 1.4 billion cubic kilometers (km3) of which more than 97 percent resides in the oceans. Another 2 percent of the total is found in the polar ice caps and glaciers, and the third-largest active reservoir is groundwater (0.3 percent). Over the oceans, evaporation exceeds precipitation by some 40,000 km3. A significant proportion of the water evaporated from the oceans is transported by winds over the land surface in the form of water vapor, where it is deposited as precipitation. Of the 111,000 km3 of water that falls as precipitation on the land surface, only some 71,000 km3 is returned to the atmosphere as evapotranspiration. The remaining 40,000 km3 is carried as runoff by rivers and eventually returns to the oceans. This amount balances the net loss of water from the oceans to the atmosphere through evaporation that is eventually deposited on the continents (land surface) as precipitation (see Figure 3.2).
The relatively small size of the atmospheric reservoir (only 13 km3) does not reflect its importance in the global water cycle. In Figure 3.2, note the large fluxes between the atmosphere, the oceans, and the land surface relative to the amount of water residing in the atmosphere at any given time (e.g., the size of atmospheric reservoir). The importance of the atmosphere in the global water cycle is better reflected by the turnover time of this reservoir. The turnover time is calculated by dividing the size of the reservoir by the rate of output (flux out). For example, the turnover time for the ocean is the size of the reservoir (1.37 × 106 km3) divided by the rate of evaporation (425 km3 per year) or more than 3000 years. In contrast, the turnover time of the atmospheric reservoir is approximately 0.024 year. That is to say, the entire water content of the atmosphere is replaced on average every nine days.
3.2 Water Has Important Physical Properties
The physical arrangement of its component molecules makes water a unique substance. A molecule of water consists of two atoms of hydrogen (H) joined to one atom of oxygen (O), represented by the chemical symbol H2O. The H atoms are bonded to the O atom asymmetrically, such that the two H atoms are at one end of the molecule and the O atom is at the other (Figure 3.3a). The bonding between the two hydrogen atoms and the oxygen atom is via shared electrons (called a covalent bond), so that each H atom shares a single electron with the oxygen. The shared hydrogen atoms are closer to the oxygen atom than they are to each other. As a result, the side of the water molecule where the H atoms are located has a positive charge, and the opposite side where the oxygen atom is located has a negative charge, thus polarizing the water molecule (termed a polar covalent bond; Figure 3.3b).
Because of its polarity, each water molecule becomes weakly bonded with its neighboring molecules (Figure 3.3c). The positive (hydrogen) end of one molecule attracts the negative (oxygen) end of the other. The angle between the hydrogen atoms encourages an open, tetrahedral arrangement of water molecules. This situation, wherein hydrogen atoms act as connecting links between water molecules, is called hydrogen bonding. The simultaneous bonding of a hydrogen atom to the oxygen atoms of two different water molecules gives rise to a lattice arrangement of molecules (Figure 3.3d). These bonds, however, are weak in comparison to the bond between the hydrogen and oxygen atoms. As a result, they are easily broken and reformed.
Water has some unique properties related to its hydrogen bonds. One property is high specific heat—the number of calories necessary to raise the temperature of 1 gram of water 1 degree Celsius. The specific heat of water is defined as a value of 1, and other substances are given a value relative to that of water. Water can store tremendous quantities of heat energy with a small rise in temperature. As a result, great quantities of heat must be absorbed before the temperature of natural waters, such as ponds, lakes, and seas, rises just 1°C. These waters warm up slowly in spring and cool off just as slowly in the fall. This process prevents the wide seasonal fluctuations in the temperature of aquatic habitats so characteristic of air temperatures and moderates the temperatures of local and worldwide environments (see Section 2.7). The high specific heat of water is also important in the thermal regulation of organisms. Because 75–95 percent of the weight of all living cells is water, temperature variation is also moderated relative to changes in ambient temperature.
As a result of the high specific heat of water, large quantities of heat energy are required for it to change its state between solid (ice), liquid, and gaseous (water vapor) phases. Collectively, the energy released or absorbed in transforming water from one state to another is called latent heat (see Section 2.5). Removing only 1 calorie (4.184 joules [J]) of heat energy will lower the temperature of a gram of water from 2°C to 1°C, but approximately 80 times as much heat energy (80 calories per gram) must be removed to convert that same quantity of water at 1°C to ice (water’s freezing point of 0°C). Likewise, it takes 536 calories to overcome the attraction between molecules and convert 1 gram (g) of water at 100°C into vapor, the same amount of heat needed to raise 536 g of water 1°C.
The lattice arrangement of molecules gives water a peculiar density–temperature relationship. Most liquids become denser as they are cooled. If cooled to their freezing temperature, they become solid, and the solid phase is denser than the liquid. This description is not true for water. Pure water becomes denser as it is cooled until it reaches 4°C ( Figure 3.4 ). Cooling below this temperature results in a decrease in density. When 0°C is reached, freezing occurs and the lattice structure is complete—each oxygen atom is connected to four other oxygen atoms by means of hydrogen atoms. The result is a lattice with large, open spaces and therefore decreased density (see Figure 3.3e). When frozen, water molecules occupy more space than they do in liquid form. Because of its reduced density, ice is lighter than water and floats on it. This property is crucial to life in aquatic environments. The ice on the surface of water bodies insulates the waters below, helping to keep larger bodies of water from freezing solid during the winter months.
Because of hydrogen bonding, water molecules tend to stick firmly to one another, resisting external forces that would break their bonds. This property is called cohesion. In a body of water, these forces of attraction are the same on all sides. At the water’s surface, however, conditions are different. Below the surface, molecules of water are strongly attracted to one another. Above the surface is the much weaker attraction between water molecules and air. Therefore, molecules on the surface are drawn downward, resulting in a surface that is taut like an inflated balloon. This condition, called surface tension, is important in the lives of aquatic organisms.
For example, the surface of water is able to support small objects and animals, such as the water striders (Gerridae spp.) and water spiders (Dolomedes spp.) that run across a pond’s surface ( Figure 3.5 ). To other small organisms, surface tension is a barrier, whether they wish to penetrate the water below or escape into the air above. For some, the surface tension is too great to break; for others, it is a trap to avoid while skimming the surface to feed or to lay eggs. If caught in the surface tension, a small insect may flounder on the surface. The nymphs of mayflies (Ephemeroptera spp.) and caddis flies (Trichoptera spp.) that live in the water and transform into winged adults are hampered by surface tension when trying to emerge from the water. While slowed down at the surface, these insects become easy prey for fish.
Cohesion is also responsible for the viscosity of water. Viscosity is the property of a material that measures the force necessary to separate the molecules and allow an object to pass through the liquid. Viscosity is the source of frictional resistance to objects moving through water. This frictional resistance of water is 100 times greater than that of air. The streamlined body shape of many aquatic organisms, for example most fish and marine mammals, helps to reduce frictional resistance. Replacement of water in the space left behind by the moving animal increases drag on the body. An animal streamlined in reverse, with a short, rounded front and a rapidly tapering body, meets the least water resistance. The perfect example of such streamlining is the sperm whale (Physeter catodon; Figure 3.6 ).
Water’s high viscosity relative to that of air is largely the result of its greater density. The density of water is about 860 times greater than that of air (pure water has a density of 1000 kilograms per cubic meter [kg/m3]). Although the resulting viscosity of water limits the mobility of aquatic organisms, it also benefits them. If a body submerged in water weighs less than the water it displaces, it is subjected to an upward force called buoyancy. Because most aquatic organisms (plants and animals) are close to neutral buoyancy (their density is similar to that of water), they do not require structural material such as skeletons or cellulose to hold their bodies erect against the force of gravity. Similarly, when moving on land, terrestrial animals must raise their mass against the force of gravity with each step they take. Such movement requires significantly more energy than swimming movements do for aquatic organisms.
But water’s greater density can profoundly affect the metabolism of marine organisms inhabiting the deeper waters of the ocean. Because of its greater density, water also undergoes greater changes in pressure with depth than does air. At sea level, the weight of the vertical column of air from the top of the atmosphere to the sea surface is 1 kilogram per square centimeter (kg/cm2) or 1 atmosphere (atm). In contrast, pressure increases 1 atm for each 10 m in depth. Because the deep ocean varies in depth from a few hundred meters down to the deep trenches at more than 10,000 m, the range of pressure at the ocean bottom is from 20 atm to more than 1000 atm. Recent research has shown that both proteins and biological membranes are strongly affected by pressure, and animals living in the deep ocean have evolved adaptations that allow these biochemical systems to function under conditions of extreme pressure.
3.3 Light Varies with Depth in Aquatic Environments
When light strikes the surface of water, a certain amount is reflected back to the atmosphere. The amount of light reflected from the surface depends on the angle at which the light strikes the surface. The lower the angle, the larger the amount of light reflected. As a result, the amount of light reflected from the water surface will vary both diurnally and seasonally between the equator and the poles (see Section 2.1 and Figure 2.5 for a complete discussion).
The amount of light entering the water surface is further reduced by two additional processes. First, suspended particles, both alive and dead, intercept the light and either absorb or scatter it. The scattering of light increases the length of its path through the water and results in further attenuation. Second, water itself absorbs light (Figure 3.7). Moreover, water absorbs some wavelengths more than others. First to be absorbed are visible red light and infrared radiation in wavelengths greater than 750 nanometers (nm). This absorption reduces solar energy by half. Next, in clear water, yellow disappears, followed by green and violet, leaving only blue wavelengths to penetrate deeper water. A fraction of blue light is lost with increasing depth. In the clearest seawater, only about 10 percent of blue light reaches to more than 100 m in depth.
These changes in the quantity and quality of light have important implications for life in aquatic environments, both by directly influencing the quantity and distribution of productivity and by indirectly influencing the vertical profile of temperature with water depth (see Section 20.4 and Chapter 24). The lack of light in deeper waters of the oceans has resulted in various adaptations. Organisms of the deeper ocean (200–1000 m deep) are typically silvery gray or deep black, and organisms living in even deeper waters (below 1000 m) often lack pigment. Another adaptation is large eyes, which give these organisms maximum light-gathering ability. Many organisms have adapted organs that produce light through chemical reactions referred to as bioluminescence (see Section 24.10).
Interpreting Ecological Data
Q1. As you dive down in depth from the surface, which wavelength of light is the first to disappear? At approximately what depth would this occur?
Q2. Is it the shorter or longer wavelengths of visible light that penetrate the deepest into the water column? (Refer to Figure 2.1.)
3.4 Temperature Varies with Water Depth
Surface temperatures reflect the balance of incoming and outgoing radiation (see Section 2.1). As solar radiation is absorbed in the vertical water column, the temperature profile with depth might be expected to resemble the vertical profile of light shown in Figure 3.7—that is, decreasing exponentially with depth. However, the physical characteristic of water density plays an important role in modifying this pattern (see Section 3.2, Figure 3.4).
As sunlight is absorbed in the surface waters, it heats up (Figure 3.8). Winds and surface waves mix the surface waters, distributing the heat vertically. Warm surface waters move downward, whereas the cooler waters below move up to the surface. As a result of this vertical mixing, heat is transported from the surface downward and the decline in water temperature with depth lags the decline in solar radiation. Below this mixed layer, however, temperatures drop rapidly. The region of the vertical depth profile where the temperature declines most rapidly is called the thermocline. The depth of the thermocline will depend on the input of solar radiation to the surface waters and on the degree of vertical mixing (wind speed and wave action). Below the thermocline, water temperatures continue to fall with depth but at a much slower rate. The result is a distinct pattern of temperature zonation with depth.
The difference in temperature between the warm, well-mixed surface layer and the cooler waters below the thermocline causes a distinctive difference in water density in these two vertical zones. The thermocline is located between an upper layer of warm, lighter (less dense) water called the epilimnion and a deeper layer of cold, denser water called the hypolimnion (see Figure 3.8; also see Section 21.10 and Figure 21.23). The density change at the thermocline acts as a physical barrier that prevents mixing of the upper (epilimnion) and lower (hypolimnion) layers.
Just as seasonal variation in the input of solar radiation to Earth’s surface results in seasonal changes in surface temperatures (see Section 2.2), seasonal changes in the input of solar radiation to the water surface give rise to seasonal changes in the vertical profile of temperature in aquatic environments (Figure 3.9). Because of the relatively constant input of solar radiation to the water surface throughout the year, the thermocline is a permanent feature of tropical waters. In the waters of the temperate zone, a distinct thermocline exists during the summer months. By fall, conditions begin to change, and a turnabout takes place. Air temperatures and sunlight decrease, and the surface water of the epilimnion starts to cool. As it does, the water becomes denser and sinks, displacing the warmer water below to the surface, where it cools in turn. As the difference in water density between the epilimnion and hypolimnion continues to decrease, winds are able to mix the vertical profile to greater depths. This process continues until the temperature is uniform throughout the basin (see Figure 3.9). Now, pond and lake water circulate throughout the basin. This process of vertical circulation, called the turnover, is an important component of nutrient dynamics in open-water ecosystems (see Chapter 21). Stirred by wind, the process of vertical mixing may last until ice forms at the surface.
Then comes winter, and as the surface water cools to below 4°C, it becomes lighter again and remains on the surface. (Remember, water becomes lighter above and below 4°C; see Figure 3.4.) If the winter is cold enough, surface water freezes; otherwise, it remains close to 0°C. Now the warmest place in the pond or lake is on the bottom. In spring, the breakup of ice and heating of surface water with increasing inputs of solar radiation to the surface again causes the water to stratify.
Because not all bodies of water experience such seasonal changes in stratification, this phenomenon is not necessarily characteristic of all deep bodies of water. In some deep lakes and the oceans, the thermocline simply descends during periods of turnover and does not disappear at all. In such bodies of water, the bottom water never becomes mixed with the top layer. In shallow lakes and ponds, temporary stratification of short duration may occur; in other bodies of water, stratification may exist, but the depth is not sufficient to develop a distinct thermocline. However, some form of thermal stratification occurs in all open bodies of water.
Temperature and density profiles with water depth for an open body of water such as a lake or pond. (a) The vertical profile of temperature might be expected to resemble the profile of light presented in Figure 3.7, but vertical mixing of the surface waters transports heat to the waters below. Below this mixed layer, temperatures decline rapidly in a region called the thermocline. Below the thermocline, temperatures continue declining at a slower rate. The vertical profile can therefore be divided into three distinct zones: epilimnion, thermocline, and hypolimnion. (b) The rapid decline in temperature in the thermocline results in a distinct difference in water density (see Figure 3.4) in the warmer epilimnion as compared to the cooler waters of the hypolimnion, leading to a two-layer density profile—warm, low-density surface water and cold, high-density deep water.
The temperature of a flowing body of water (stream or river), on the other hand, is variable (Figure 3.10). Small, shallow streams tend to follow, but lag behind, air temperatures. They warm and cool with the seasons but rarely fall below freezing in winter. Streams with large areas exposed to sunlight are warmer than those shaded by trees, shrubs, and high banks. That fact is ecologically important because temperature affects the stream community, influencing the presence or absence of cool- and warm-water organisms. For example, the dominant predatory fish shift from species such as trout and smallmouth bass, which require cooler water and more oxygen, to species such as suckers and catfish, which require warmer water and less oxygen (see Figure 24.13).
3.5 Water Functions as a Solvent
As you stir a spoonful of sugar into a glass of water, it dissolves, forming a homogeneous, or uniform, mixture. A liquid that is a homogeneous mixture of two or more substances is called a solution. The dissolving agent of a solution is the solvent, and the substance that is dissolved is referred to as the solute. A solution in which water is the solvent is called an aqueous solution.
Water is an excellent solvent that can dissolve more substances than can any other liquid. This extraordinary ability makes water a biologically crucial substance. Water provides a fluid that dissolves and transports molecules of nutrients and waste products, helps to regulate temperature, and preserves chemical equilibrium within living cells.
The solvent ability of water is largely a result of the bonding discussed in Section 3.2. Because the H atom is bonded to the O atoms asymmetrically (see Figure 3.3), one side of every water molecule has a permanent positive charge and the other side has a permanent negative charge; such a situation is called a permanent dipole (where dipole refers to oppositely charged poles). Because opposite charges attract, water molecules are strongly attracted to one another; they also attract other molecules carrying a charge.
Compounds that consist of electrically charged atoms or groups of atoms are called ions. Sodium chloride (table salt), for example, is composed of positively charged sodium ions (Na+) and negatively charged chloride ions (Cl–) arranged in a crystal lattice. When placed in water, the attractions between negative (oxygen atom) and positive (hydrogen atoms) charges on the water molecule (see Figure 3.3) and those of the sodium and chloride atoms are greater than the forces (ionic bonds) holding the salt crystals together. Consequently, the salt crystals readily dissociate into their component ions when placed in contact with water; that is, they dissolve.
The solvent properties of water are responsible for the presence of most of the minerals (elements and inorganic compounds) found in aquatic environments. When water condenses to form clouds, it is nearly pure except for some dissolved atmospheric gases. In falling to the surface as precipitation, water acquires additional substances from particulates and dust particles suspended in the atmosphere. Water that falls on land flows over the surface and percolates into the soil, obtaining more solutes. Surface waters, such as streams and rivers, pick up more solvents from the substances through and over which they flow. The waters of most rivers and lakes contain 0.01–0.02 percent dissolved minerals. The relative concentrations of minerals in these waters reflect the substrates over which the waters flow. For example, waters that flow through areas where the underlying rocks consist largely of limestone, composed primarily of calcium carbonate (CaCO3), will have high concentrations of calcium (Ca2+) and bicarbonate (HCO3–).
In contrast to freshwaters, the oceans have a much higher concentration of solutes. In effect, the oceans function as a large still. The flow of freshwaters into the oceans continuously adds to the solute content of the waters, as pure water evaporates from the surface to the atmosphere. The concentration of solutes, however, cannot continue to increase indefinitely. When the concentration of specific elements reaches the limit set by the maximum solubility of the compounds they form (grams per liter), the excess amounts will precipitate and be deposited as sediments. Calcium, for example, readily forms calcium carbonate (CaCO3) in the waters of the oceans. The maximum solubility of calcium carbonate, however, is only 0.014 gram per liter of water, a concentration that was reached early in the history of the oceans. As a result, calcium ions continuously precipitate out of solution and are deposited as limestone sediments on the ocean bottom.
In contrast, the solubility of sodium chloride is high (360 grams per liter). In fact, these two elements, sodium and chlorine, make up some 86 percent of sea salt. Sodium and chlorine—along with other major elements such as sulfur, magnesium, potassium, and calcium, whose relative proportions vary little—compose 99 percent of sea salts (Figure 3.11). Determination of the most abundant element, chlorine, is used as an index of salinity. Salinity is expressed in practical salinity units (psu), represented as ‰ and measured as grams of chlorine per kilogram of water. The salinity of the open sea is fairly constant, averaging about 35‰. In contrast, the salinity of freshwater ranges from 0.065 to 0.30‰. However, over geologic timescales (hundreds of millions of years), the salinity of the oceans has increased and continues to do so.
3.6 Oxygen Diffuses from the Atmosphere to the Surface Waters
Water’s role as a solvent is not limited to dissolving solids. The surface of a body of water defines a boundary with the atmosphere. Across this boundary, gases are exchanged through the process of diffusion. Diffusion is the general tendency of molecules to move from a region of high concentration to one of lower concentration. The process of diffusion results in a net transfer of two metabolically important gases, oxygen and carbon dioxide, from the atmosphere (higher concentration) into the surface waters (lower concentration) of aquatic environments
Oxygen diffuses from the atmosphere into the surface water. The rate of diffusion is controlled by the solubility of oxygen in water and the steepness of the diffusion gradient (the difference in concentration between the air and the surface waters where diffusion occurs). The solubility of gases in water is a function of temperature, pressure, and salinity. The saturation value of oxygen is greater for cold water than warm water because the solubility (ability to stay in solution) of a gas in water decreases as the temperature rises. However, solubility increases as atmospheric pressure increases and decreases as salinity increases, which is not significant in freshwater.
Once oxygen enters the surface water, the process of diffusion continues, and oxygen diffuses from the surface to the waters below (because of their lower concentration). Water, with its greater density and viscosity relative to air, limits how quickly gases diffuse through it. Gases diffuse some 10000 times slower in water than in air. In addition to the process of diffusion, oxygen absorbed by surface water is mixed with deeper water by turbulence and internal currents. In shallow, rapidly flowing water and in wind-driven sprays, oxygen may reach and maintain saturation and even supersaturated levels because of the increase of absorptive surfaces at the air–water interface. Oxygen is lost from the water as temperatures rise, decreasing solubility, and through the uptake of oxygen by aquatic life.
Oxygen stratification in Mirror Lake, New Hampshire, in winter, summer, and late fall. The late fall turnover results in uniform temperature as well as uniform distribution of oxygen throughout the lake basin. In summer, a pronounced stratification of both temperature and oxygen exists. Oxygen declines sharply in the thermocline and is nonexistent on the bottom because of its uptake by decomposer organisms in the sediments. In winter, oxygen levels are high in surface water reflecting higher solubility. Formation of ice during winter, however, can greatly reduce diffusion into surface waters. (Adapted from Likens 1985.)
During the summer, oxygen, like temperature (see Section 3.4), may become stratified in lakes and ponds. The amount of oxygen is usually greatest near the surface, where an interchange between water and atmosphere, further stimulated by the stirring action of the wind, takes place (Figure 3.12). Besides entering the water by diffusion from the atmosphere, oxygen is also a product of photosynthesis, which is largely restricted to the surface waters because of the limitations of available light (see Figure 3.7 and Chapter 6). The quantity of oxygen decreases with depth because of the oxygen demand of decomposer organisms living in the bottom sediments (Chapter 21). During spring and fall turnover, when water recirculates through the lake, oxygen becomes replenished in deep water. In winter, the reduction of oxygen in unfrozen water is slight because the demand for oxygen by organisms is reduced by the cold, and oxygen is more soluble at low temperatures. Under ice, however, oxygen depletion may be serious as a result of the lack of diffusion from the atmosphere to the surface waters.
As with ponds and lakes, oxygen is not distributed uniformly within the depths of the oceans (Figure 3.13). A typical oceanic oxygen profile shows a maximum amount in the upper 10–20 m, where photosynthetic activity and diffusion from the atmosphere often lead to saturation. With increasing depth, oxygen content declines. In the open waters of the ocean, concentrations reach a minimum value of 500–1000 m, a region referred to as the oxygen minimum zone. Unlike lakes and ponds, where the seasonal breakdown of the thermocline and resultant mixing of surface and deep waters result in a dynamic gradient of temperature and oxygen content, the limited depth of surface mixing in the deep oceans maintains the vertical gradient of oxygen availability year-round.
Vertical profile of oxygen with depth in the Atlantic Ocean. The oxygen content of the waters declines to a depth known as the oxygen minimum zone. Oxygen increase below this may be the result of the influx of cold, oxygen-rich waters that sank in the polar waters.
The availability of oxygen in aquatic environments characterized by flowing water is different. The constant churning of stream water over riffles and falls gives greater contact with the atmosphere; the oxygen content of the water is high, often near saturation for the prevailing temperature. Only in deep holes or in polluted waters does dissolved oxygen show any significant decline (see Chapter 24, Ecological Issues & Applications).
Even under ideal conditions, gases are not very soluble in water. Rarely is oxygen limited in terrestrial environments. In aquatic environments, the supply of oxygen, even at saturation levels, is meager and problematic. Compared with its concentration of 0.21 liter per liter in the atmosphere (21 percent by volume), oxygen in water reaches a maximum solubility of 0.01 liter per liter (1 percent) in freshwater at a temperature of 0°C. As a result, the concentration of oxygen in aquatic environments often limits respiration and metabolic activity.
3.7 Acidity Has a Widespread Influence on Aquatic Environments
The solubility of carbon dioxide is somewhat different from that of oxygen in its chemical reaction with water. Water has a considerable capacity to absorb carbon dioxide, which is abundant in both freshwater and saltwater. Upon diffusing into the surface, carbon dioxide reacts with water to produce carbonic acid (H2CO3):
CO2+H2O⇋H2CO3CO2 + H2O⇋ H2CO3
Carbonic acid further dissociates into a hydrogen ion and a bicarbonate ion:
H2CO3⇋HCO3−+H+H2CO3 ⇋ HCO3− +H+
Bicarbonate may further dissociate into another hydrogen ion and a carbonate ion:
HCO−3⇋H++CO2−3HCO3− ⇋ H++CO32−
The carbon dioxide–carbonic acid–bicarbonate system is a complex chemical system that tends to stay in equilibrium. (Note that the arrows in the preceding equations go in both directions.) Therefore, if carbon dioxide (CO2) is removed from the water, the equilibrium is disturbed and the equations will shift to the left, with carbonic acid and bicarbonate producing more CO2 until a new equilibrium is reached.
The chemical reactions just described result in the production and absorption of free hydrogen ions (H+). The abundance of hydrogen ions in solution is a measure of acidity. The greater the number of H+ ions, the more acidic is the solution. Alkaline solutions are those that have a large number of OH– (hydroxyl ions) and few H+ ions. The measurement of acidity and alkalinity is called pH, calculated as the negative logarithm (base 10) of the concentration of hydrogen ions in solution. In pure water, a small fraction of molecules dissociates into ions: H2O → H+ + OH–, and the ratio of H+ ions to OH– ions is 1:1. Because both occur in a concentration of 10–7 moles per liter, a neutral solution has a pH of 7 [–log(10–7) = 7]. A solution departs from neutral when the concentration of one ion increases and the other decreases. Customarily, we use the negative logarithm of the hydrogen ion to describe a solution as an acid or a base. Thus, a gain of hydrogen ions to 10–6 moles per liter means a decrease of OH– ions to 10–8 moles per liter, and the pH of the solution is 6. The negative logarithmic pH scale goes from 1 to 14. A pH greater than 7 denotes an alkaline solution (greater OH– concentration) and a pH of less than 7 an acidic solution (greater H+ concentration).
Although pure water is neutral in pH, because the dissociation of the water molecule produces equal numbers of H+ and OH– ions, the presence of CO2 in the water alters this relationship. The preceding chemical reactions result in the production and absorption of H+ ions. Because the abundance of hydrogen ions in solution is the measure of acidity, the dynamics of the carbon dioxide–carbonic acid–bicarbonate system directly affect the pH of aquatic ecosystems. In general, the carbon dioxide–carbonic acid–bicarbonate system functions as a buffer to keep the pH of water within a narrow range. It does this by absorbing hydrogen ions in the water when they are in excess (producing carbonic acid and bicarbonates) and producing them when they are in short supply (producing carbonate and bicarbonate ions). At neutrality (pH 7), most of the CO2 is present as HCO3– ( Figure 3.14). At a high pH, more CO2 is present as CO32– than at a low pH, where more CO2 occurs in the free condition. Addition or removal of CO2 affects pH, and a change in pH affects CO2.
The pH of natural waters ranges between 2 and 12. Waters draining from watersheds dominated geologically by limestone (CaCO3) have a much higher pH and are well buffered as compared to waters from watersheds dominated by acid sandstone and granite. The presence of the strongly alkaline ions sodium, potassium, and calcium in ocean waters results in seawater being slightly alkaline, usually ranging from 7.5 to 8.4.
The pH of aquatic environments can exert a powerful influence on the distribution and abundance of organisms. Increased acidity can affect organisms directly, by influencing physiological processes, and indirectly, by influencing the concentrations of toxic heavy metals. Tolerance limits for pH vary among plant and animal species, but most organisms cannot survive and reproduce at a pH below about 4.5. Aquatic organisms are unable to tolerate low pH conditions largely because acidic waters contain high concentrations of aluminum. Aluminum is highly toxic to many species of aquatic life and thus leads to a general decline in aquatic populations.
Theoretical percentages of carbon dioxide (CO2) in each of the three forms present in water in relation to pH. At low values of pH (acidic conditions), most of the CO2 is in its free form. At intermediate values (neutral conditions) bicarbonate dominates, whereas under alkaline conditions most of the CO2 is in the form of carbonate ions.
Interpreting Ecological Data
Q1. Under conditions of neutral pH, what is the relative abundance of the different forms of CO2?
Q2. The current pH of the ocean is approximately 8.1. what is the dominant form of CO2?
Aluminum is insoluble when the pH is neutral or basic. Insoluble aluminum is present in high concentrations in rocks, soils, and river and lake sediments. Under normal pH conditions, the aluminum concentrations of lake water are low; however, as the pH drops and becomes more acidic, aluminum begins to dissolve, raising the concentration in solution.
3.8 Water Movements Shape Freshwater and Marine Environments
The movement of water—currents in streams and waves in an open body of water or breaking on a shore—determines the nature of many aquatic environments. The velocity of a current molds the character and structure of a stream. The shape and steepness of the stream channel, its width, depth, and roughness of the bottom, and the intensity of rainfall and rapidity of snowmelt all affect velocity. In fast streams, velocity of flow is 50 cm per second or higher (see Chapter 24, Quantifying Ecology 24.1). At this velocity, the current removes all particles less than 5 millimeters (mm) in diameter and leaves behind a stony bottom. High water volume increases the velocity; it moves bottom stones and rubble, scours the streambed, and cuts new banks and channels. As the gradient decreases and the width, depth, and volume of water increase, silt and decaying organic matter accumulate on the bottom. Thus, the stream’s character changes from fast water to slow (Figure 3.15).
Wind generates waves on large lakes and on the open sea. The frictional drag of the wind on the surface of smooth water causes ripples. As the wind continues to blow, it applies more pressure to the steep side of the ripple, and wave size begins to grow. As the wind becomes stronger, short, choppy waves of all sizes appear; as they absorb more energy, they continue to grow. When the waves reach a point where the energy supplied by the wind equals the energy lost by the breaking waves, they become whitecaps. Up to a certain point, the stronger the wind, the higher the waves.
The waves breaking on a beach do not contain water driven in from distant seas. Each particle of water remains largely in the same place and follows an elliptical orbit with the passage of the wave. As a wave moves forward, it loses energy to the waves behind and disappears, its place taken by another. The swells breaking on a beach are distant descendants of waves generated far out at sea.
As the waves approach land, they advance into increasingly shallow water. When the bottom of the wave intercepts the ocean floor, the wavelength shortens and the wave steepens until it finally collapses forward, or breaks. As the waves break onshore, they dissipate their energy, pounding rocky shores or tearing away sandy beaches in one location and building up new beaches elsewhere.
We have discussed the patterns of ocean currents, influenced by the direction of the prevailing winds and the Coriolis effect in Chapter 2 (see Section 2.4). As the warm surface currents of the tropical waters move northward and southward (see map of surface currents in Figure 2.13), they bring up deep, cold, oxygenated waters from below, a process known as upwelling (Figure 3.16a). A similar pattern occurs in coastal regions. Winds blowing parallel to the coast move the surface waters offshore. Water moving upward from the deep replaces this surface water, creating a pattern of coastal upwelling (Figure 3.16b).
Along the equator, the Coriolis effect acts to pull the westward-flowing currents to the north and south (purple solid arrows), resulting in an upwelling of deeper cold waters to the surface. (b) Along the western margins of the continents, the Coriolis effect causes the surface waters to move offshore (purple solid arrows). Movement of the surface waters offshore results in an upwelling of deeper, colder waters to the surface. Example shown is for the Northern Hemisphere.
3.9 Tides Dominate the Marine Coastal Environment
Tides profoundly influence the rhythm of life on ocean shores. Tides result from the gravitational pulls of the Sun and the Moon, each of which causes two bulges (tides) in the waters of the oceans. The two bulges caused by the Moon occur at the same time on opposite sides of Earth on an imaginary line extending from the Moon through the center of Earth (Figure 3.17). The tidal bulge on the Moon side is a result of gravitational attraction; the bulge on the opposite side occurs because the gravitational force there is less than at the Earth’s center. As Earth rotates eastward on its axis, the tides advance westward. Thus, in the course of one daily rotation, Earth passes through two of the lunar tidal bulges, or high tides, and two of the lows, or low tides, at right angles (90° longitude difference) to the high tides.
Tides result from the gravitational pull of the Moon. Centrifugal force applied to a kilogram of mass is 3.38 milligrams (mg). This force on a rotating Earth is balanced by gravitational force, except at moving points on Earth’s surface that are directly aligned with the Moon. Thus, the centrifugal force at point N, the center of the rotating Earth, is 3.38 mg. Point T is directly aligned with the Moon. At this point, the Moon’s gravitational force is 3.49 mg, a difference of 0.11 mg. Because the Moon’s gravitational force is greater than the centrifugal force at T, the force is directed away from the Earth and causes a tidal bulge. At point A, the Moon’s gravitational force is 3.27 mg, 0.11 mg less than the centrifugal force at N. This causes a tidal bulge on the opposite side of Earth.
The Sun also causes two tides on opposite sides of Earth, and these tides have a relation to the Sun like that of the lunar tides to the Moon. Because the Sun has a weaker gravitational pull than the Moon does, solar tides are partially masked by lunar tides—except for two times during the month: when the Moon is full and when it is new. At these times, Earth, Moon, and Sun are nearly in line, and the gravitational pulls of the Sun and the Moon are additive. This combination makes the high tides of those periods exceptionally large, with maximum rise and fall. These are the fortnightly spring tides, a name derived from the Saxon word sprungen, which refers to the brimming fullness and active movement of the water. When the Moon is at either quarter, its pull is at right angles to the pull of the Sun, and the two forces interfere with each other. At those times, the differences between high and low tides are exceptionally small. These are called the neap tides, from an old Scandinavian word meaning “barely enough.”
Tides are not entirely regular, nor are they the same all over Earth. They vary from day to day in the same place, following the waxing and waning of the Moon. They may act differently in several localities within the same general area. In the Atlantic, semidaily tides are the rule. In the Gulf of Mexico and the Aleutian Islands of Alaska, the alternate highs and lows more or less cancel each other out, and flood and ebb follow one another at about 24-hour intervals to produce one daily tide. Mixed tides in which successive or low tides are of significantly different heights through the cycle are common in the Pacific and Indian oceans. These tides are combinations of daily and semidaily tides in which one partially cancels out the other.
Local tides around the world are inconsistent for many reasons. These reasons include variations in the gravitational pull of the Moon and the Sun as a result of the elliptical orbit of Earth, the angle of the Moon in relation to the axis of Earth, onshore and offshore winds, the depth of water, the contour of the shore, and wave action.
The area lying between the water lines of high and low tide, referred to as the intertidal zone, is an environment of extremes. The intertidal zone undergoes dramatic shifts in environmental conditions with the daily patterns of inundation and exposure. As the tide recedes, the uppermost layers of life are exposed to air, wide temperature fluctuations, intense solar radiation, and desiccation for a considerable period, whereas the lowest fringes of the tidal zone may be exposed only briefly before the high tide submerges them again. Temperatures on tidal flats may rise to 38°C when exposed to direct sunlight and drop to 10°C within a few hours when the flats are covered by water.
Organisms living in the sand and mud do not experience the same violent temperature fluctuations as those living on rocky shores do. Although the surface temperature of the sand at midday may be 10°C (or more) higher than that of the returning seawater, the temperature a few centimeters below the sand’s surface remains almost constant throughout the year (see Section 25.3).
3.10 The Transition Zone between Freshwater and Saltwater Environments Presents Unique Constraints
Water from streams and rivers eventually drains into the sea. The place where freshwater mixes with saltwater is called an estuary. Temperatures in estuaries fluctuate considerably, both daily and seasonally. Sun and inflowing and tidal currents heat the water. High tide on the mudflats may heat or cool the water, depending on the season. The upper layer of estuarine water may be cooler in winter and warmer in summer than the bottom—a condition that, as in a lake, will cause spring and autumn turnovers (see Figures 3.9 and 3.12).
In the estuary, where freshwater meets the sea, the interaction of inflowing freshwater and tidal saltwater influences the salinity of the estuarine environment. Salinity varies vertically and horizontally, often within one tidal cycle (Figure 3.18).
Vertical and horizontal stratification of salinity from the river mouth to the estuary at high tide (brown lines) and low tide (blue lines). At high tide, the incoming seawater increases the salinity toward the river mouth. At low tide, salinity is reduced. Salinity increases with depth because lighter freshwater flows over denser saltwater.
Salinity may be the same from top to bottom or it may be completely stratified, with a layer of freshwater on top and a layer of dense, salty water on the bottom. Salinity is homogeneous when currents are strong enough to mix the water from top to bottom. The salinity in some estuaries is homogeneous at low tide, but at high tide a surface wedge of seawater moves upstream more rapidly than the bottom water. Salinity is then unstable, and density is inverted. The seawater on the surface tends to sink as lighter freshwater rises, and mixing takes place from the surface to the bottom. This phenomenon is known as tidal overmixing. Strong winds, too, tend to mix saltwater with freshwater in some estuaries, but when the winds are still, the river water flows seaward on a shallow surface over an upstream movement of seawater, more gradually mixing with the salt.
Horizontally, the least saline waters are at the river mouth and the most saline at the sea (see Figure 3.18). Incoming and outgoing currents deflect this configuration. In all estuaries of the Northern Hemisphere, outward-flowing freshwater and inward-flowing seawater are deflected to the right (relative to the axis of water flow from the river to ocean) because of Earth’s rotation (Coriolis effect; see Section 2.3). As a result, salinity is higher on the left side; the concentration of metallic ions carried by rivers varies from drainage to drainage; and salinity and chemistry differ among estuaries. The portion of dissolved salts in the estuarine waters remains about the same as that of seawater, but the concentration varies in a gradient from freshwater to sea.
To survive in estuaries, aquatic organisms must have evolved physiological or behavioral adaptations to changes in salinity. Many oceanic species of fish are able to move inward during periods when the flow of freshwater from rivers is low and the salinity of estuaries increases. Conversely, freshwater fish move into the estuarine environment during periods of flood when salinity levels drop. Because of the stressful conditions that organisms face in the mixed zones of estuaries, there is often a relatively low diversity of organisms despite the high productivity found in these environments (see Chapter 24).
Ecological Issues & Applications Rising Atmospheric Concentrations of CO2 Are Impacting Ocean Acidity
The exchange of carbon dioxide (CO2) between the atmosphere and the surface waters of the oceans is governed by the process of diffusion, with the net exchange moving CO2 from higher concentrations (atmosphere) to lower concentrations (surface waters) (Section 3.7). Upon diffusing into the surface, the CO2 reacts with the water to produce carbonic acid (H2CO3), which further dissociates into a hydrogen ion (H+) and a bicarbonate ion (HCO3−). The bicarbonate may further dissociate into another hydrogen ion and a carbonate ion (CO32−). In both of these chemical reactions, free hydrogen ions (H+) are produced, the abundance of which is a measure of acidity. The greater the number of H+ ions, the lower the value of pH and the more acidic the solution.
Under current ocean conditions, about 89 percent of the carbon dioxide dissolved in seawater takes the form of a bicarbonate ion, about 10 percent as a carbonate ion, and 1 percent as dissolved gas, and the pH of seawater on the surface of the oceans has remained relatively steady for millions of years at a value of about 8.2 (slightly basic—7.0 is neutral; see Figure 3.14). Since the height of the Industrial Revolution in the 19th century, however, atmospheric concentrations of CO2 have been steadily rising as a result of the burning of fossil fuels (see Chapter 2, Ecological Issues & Applications and Chapter 27). As a consequence, the diffusion gradient of CO2 between the atmosphere and oceans has increased, resulting in an increasing uptake of CO2 into the surface waters. As a consequence, the pH of the surface waters of the oceans has fallen by about 0.1 pH unit from preindustrial times to today (Figure 3.19). Recall from Section 3.7 that the pH scale is logarithmic (log10; thus for every drop of 1 pH unit, hydrogen ion levels increase by a factor of 10), so this 0.1-unit drop in pH is equivalent to about a 25 percent increase in the ocean hydrogen ion concentration. According to estimates from the Intergovernmental Panel on Climate Change (IPCC; see Chapter 2, Ecological Issues & Applications), under the expected trajectory of fossil fuel use and rising atmospheric CO2 concentrations, pH is likely to drop by 0.3–0.4 units by the end of the 21st century and increase ocean hydrogen ion concentration (or acidity) by 100 to 150 percent.
Increased absorption of CO2 by the surface waters of the oceans can potentially impact life in the oceans in a variety of ways, both positive and negative. Photosynthetic algae and plants may benefit from higher CO2 concentrations in the surface waters because elevated CO2 may enhance rates of photosynthesis (see Chapter 6, Ecological Issues & Applications). On the other hand, one of the most important negative impacts of increasing ocean acidity relates to the process of calcification—the production of shells and plates out of calcium carbonate (CaCO3)—which is important to the biology and survival of a wide range of marine organisms.
CaCO3 is formed in marine environments through the reaction of calcium and carbonate ions:
CO32−+C2+a⇋CaCO3CO32− + C2+a⇋ CaCO3
As with the chemical equations describing the formation of bicarbonate and carbonate ions from dissolved CO2, the reaction involved in the formation of CaCO3 proceeds in both directions. As sea water pH declines (acidity increases), carbonate ions (CO32−) function like an antacid to neutralize the H+, forming more bicarbonate (CO32− + H+ ↔ HCO3−). Therefore, declining pH results in an associated decline in carbonate ion concentrations (Figure 3.19a). This decline in carbonate ion concentration shifts the preceding equation in favor of the disassociation of CaCO3 minerals into calcium and carbonate ions. The resulting decline in dissolved CaCO3 minerals can have a significant impact on calcifying species, including oysters, clams, sea urchins, shallow water corals, deep sea corals, and calcareous plankton.
The process of calcification by marine organisms involves the precipitation of dissolved CaCO3 into solid CaCO3 structures, such as coccoliths (individual plates of CaCO3 formed by single-celled algae; Figure 3.20). After they are formed, these structures are vulnerable to dissolution unless the surrounding seawater contains saturating concentrations of CaCO3. As carbonate ions become depleted because of declining pH, seawater becomes undersaturated with respect to two CaCO3 minerals vital for calcification, aragonite and calcite (Figure 3.19b and 3.19c). Current estimates suggest that the oceans are becoming undersaturated with respect to aragonite at the poles, where the cold and dense waters most readily absorb atmospheric CO2, and that under projected rates of CO2 emissions (IPCC), undersaturation would extend throughout the entire Southern Ocean (<60° S) and into the subarctic Pacific by the end of the century (2100). These changes will threaten high-latitude aragonite secreting organisms including cold-water corals, which provide essential fish habitat, and shelled pteropods (free-swimming pelagic sea snails and sea slugs; Figure 3.21), an abundant food source for marine predators
In a review of experimental studies that have examined the response of marine calcifying species to elevated CO2 conditions, Scott Doney of Woods Hole Oceanographic Institute and colleagues found that the degree of sensitivity varies among species, and some species may even show enhanced calcification at elevated CO2 levels (Table 3.1). The researchers found, however, that in the vast majority of species, including every study published that has examined the calcification rates of coral species, elevated CO2 concentrations and the associated decreasing aragonite saturation state had a negative effect on calcification rates. The researchers concluded that “ocean acidification impacts processes so fundamental to the overall structure and function of marine ecosystems that any significant changes could have far-reaching consequences for the oceans of the future.”
Summary
The Water Cycle 3.1
Water follows a cycle, traveling from the air to Earth and returning to the atmosphere. It moves through cloud formation in the atmosphere, precipitation, interception, and infiltration into the ground. It eventually reaches groundwater, springs, streams, and lakes from which evaporation takes place, bringing water back to the atmosphere in the form of clouds. The various aquatic environments are linked, either directly or indirectly, by the water cycle.
The largest reservoir in the global water cycle is the oceans, which contain more than 97 percent of the total volume of water on Earth. In contrast, the atmosphere is one of the smallest reservoirs but has a fast turnover time.
Properties of Water 3.2
Water has a unique molecular structure. The hydrogen atoms are located on the side of the water molecule that has a positive charge. The opposite side, where the oxygen atom is located, has a negative charge, thus polarizing the water molecule. Because of their polarity, water molecules become coupled with neighboring water molecules to produce a lattice-like structure with unique properties.
Depending on its temperature, water may occur in the form of a liquid, solid, or gas. It absorbs or releases considerable quantities of heat with a small rise or fall in temperature. Water has a high viscosity that affects its flow. It exhibits high surface tension, caused by a stronger attraction of water molecules for each other than for the air above the surface. If a body is submerged in water and weighs less than the water it displaces, it is subjected to the upward force of buoyancy. These properties are important ecologically and biologically.
Light 3.3
Both the quantity and quality of light change with water depth. In pure water, red and infrared light are absorbed first, followed by yellow, green, and violet; blue penetrates the deepest.
Temperature in Aquatic Environments 3.4
Lakes and ponds experience seasonal shifts in temperature. In summer there is a distinct vertical gradient of temperature, resulting in a physical separation of warm surface waters and the colder waters below the thermocline. When the surface waters cool in the fall, the temperature becomes uniform throughout the basin and water circulates throughout the lake. A similar mixing takes place in the spring when the water warms. In some deep lakes and the oceans, the thermocline simply descends during turnover periods and does not disappear at all.
Temperature of flowing water is variable, warming and cooling with the season. Within the stream or river, temperatures vary with depth, amount of shading, and exposure to sun.
Water as a Solvent 3.5
Water is an excellent solvent with the ability to dissolve more substances than any other liquid can. The solvent properties of water are responsible for most of the minerals found in aquatic environments. The waters of most rivers and lakes contain a relatively low concentration of dissolved minerals, determined largely by the underlying bedrock over which the water flows. In contrast, the oceans have a much higher concentration of solutes. As pure water evaporates from the surface to the atmosphere, the flow of freshwaters into the oceans continuously adds to the solute content of the waters.
The solubility of sodium chloride is high; together with chlorine, it makes up some 86 percent of sea salt. The concentration of chlorine is used as an index of salinity. Salinity is expressed in practical salinity units (psu; represented as ‰, measured as grams of chlorine per kilogram of water).
Oxygen 3.6
Oxygen enters the surface waters from the atmosphere through the process of diffusion. The amount of oxygen water can hold depends on its temperature, pressure, and salinity. In lakes, oxygen absorbed by surface water mixes with deeper water by turbulence. During the summer, oxygen may become stratified, decreasing with depth because of decomposition in bottom sediments. During spring and fall turnover, oxygen becomes replenished in deep water. Constant swirling of stream water gives it greater contact with the atmosphere and thus allows it to maintain a high oxygen content.
Acidity 3.7
The measurement of acidity is pH, the negative logarithm of the concentration of hydrogen ions in solution. In aquatic environments, a close relationship exists between the diffusion of carbon dioxide into the surface waters and the degree of acidity and alkalinity. Acidity influences the availability of nutrients and restricts the environment of organisms sensitive to acid situations.
Water Movement 3.8
Currents in streams and rivers as well as waves in open sea and breaking on ocean shores determine the nature of many aquatic and marine environments. The velocity of currents shapes the environment of flowing water. Waves pound rocky shores and tear away and build up sandy beaches. Movement of water in surface currents of the ocean affects the patterns of deep-water circulation. As the equatorial currents move northward and southward, deep waters move up to the surface, forming regions of upwelling. In coastal regions, winds blowing parallel to the coast create a pattern of coastal upwelling.
Tides 3.9
Rising and falling tides shape the environment and influence the rhythm of life in coastal intertidal zones.
Estuaries 3.10
Water from all streams and rivers eventually drains into the sea. The place where this freshwater joins and mixes with the salt is called an estuary. Temperatures in estuaries fluctuate considerably, both daily and seasonally. The interaction of inflowing freshwater and tidal saltwater influences the salinity of the estuarine environment. Salinity varies vertically and horizontally, often within one tidal cycle.
Ocean Acidification Ecological Issues & Applications
Rising atmospheric concentrations of carbon dioxide have resulted in increased concentrations in the surface waters of the oceans. The increased carbon dioxide concentrations of the surface waters have resulted in a decline in pH and reduced carbonate concentrations. The reduction in carbonate concentrations has reduced calcium carbonate mineral concentrations that are essential for calcifying marine species.
