CHAPTER 2hopkins/WES/wesch02.pdf · Lake Nyos could absorb a tremendous amount of dissolved CO 2 (perhaps 5 times the maximum per unit volume in a typical carbonated beverage). Culminating

CHAPTER 2

WATER’S UNIQUE PROPERTIES• Case-in-Point• Driving Question• Phases of Water

Temperature and HeatFreezing and MeltingVaporization and CondensationSublimation and Deposition

• Why is Water so Unique?• Water’s Specific Heat• Water as a Solvent

Acid DepositionDissolved Oxygen

• Water Density and TemperatureFresh WaterSeawater

• Water’s Surface Tension• Conclusions• Basic Understandings• ESSAY: Lake Turnover

Case-in-Point

On the evening of 1 August 1986, Lake Nyos inCameroon, Central Africa, literally exploded, killing1700 residents of nearby villages. Although Lake Nyosoccupies a volcanic crater, the explosion was no volcaniceruption. The cause of the disaster at Lake Nyos hasbeen likened to what happens when the cap is removedfrom a bottle of carbonated beverage after it has beenshaken vigorously.

Lake Nyos is about 2 km (1.3 mi) long, 1.2 km(0.8 mi) wide, and 185 m (610 ft) deep. A chamber ofhot molten material (magma) some 80 km (50 mi)beneath the lake releases carbon dioxide (CO2) some ofwhich seeps upward into the lake bottom waters. Carbondioxide dissolves in lake water much as it does whenbeverages are carbonated. The maximum amount of CO2

that can dissolve in water increases with rising pressure(and falling temperature). Water at the bottom of LakeNyos is under tremendous pressure produced by the

weight of the overlying water. Every 10 m (33 ft) ofwater exerts a pressure equivalent to the average airpressure at sea level. At the bottom of Lake Nyos, thepressure was more than 18 times the usual air pressure atsea level. Under these conditions, the bottom waters ofLake Nyos could absorb a tremendous amount ofdissolved CO2 (perhaps 5 times the maximum per unitvolume in a typical carbonated beverage).

Culminating a steady long-term input of CO2,sometime prior to August 1986 the bottom waters ofLake Nyos became saturated with carbon dioxide—theycould hold no more. A short-term jolt, perhaps alandslide or minor earthquake, caused a drop in pressureand triggered a catastrophic release of CO2 from thesaturated lake waters, producing a fountain of foamingwater that rose nearly 90 m (300 ft) into the air above thelake surface. A cloud of freshly released CO2, denserthan air, spread along the ground and into low-lying

20 Chapter 2 WATER’S UNIQUE PROPERTIES

villages surrounding the lake. Carbon dioxideconcentration was well above the threshold for humantoxicity and most victims died from asphyxiation.

Driving Question:

How do water’s unique properties influence thefunctioning of the Earth system?

Water, an abundant and important component of theEarth system, is the central focus of this book. Althoughwater is a very common substance, it has some veryuncommon properties compared to substances of similarmolecular size (or chemical composition). Water’sunusually high freezing and boiling temperatures,coupled with the usual variations in temperature nearEarth’s surface, mean that water coexists on Earth in allthree phases, as solid (ice), liquid, and gas (water vapor).Furthermore, water frequently changes phase and duringthese transitions, unusually large amounts of heat energyare either absorbed from the environment or released tothe environment. The uniqueness of these and otherproperties of water (i.e., high specific heat) is derivedfrom its molecular structure and the influence ofhydrogen bonding.

In this chapter we examine water’s uniqueproperties, the fundamental reasons for those properties,and some of the implications for the functioning of theEarth system. We return to water’s unique properties insubsequent chapters as we examine the flow of water andenergy in the global water cycle.

Phases of Water

The water molecule (H2O) consists of two hydrogen (H)atoms bonded to an oxygen (O) atom (Figure 2.1). Thischemical bonding is relatively strong so that the watermolecule resists decomposition. In the atmosphere, forexample, only very intense solar radiation at highaltitudes can break a water molecule into its constituentatoms. Water is the only component of the Earth systemthat occurs naturally in all three phases, as solid (ice orsnow), liquid, and gas (water vapor). The type andamount of molecular activity distinguish the phases ofwater. When water changes phase, its level of molecular

activity either increases or decreases depending on thetype of phase change.

In the solid phase, ice is crystalline and watermolecules vibrate about fixed locations within a crystallattice (Figure 2.2). (A crystal lattice is a 3-dimensionalframework consisting of a repeated pattern of atoms ormolecules.) For this reason an ice cube or any otherpiece of ice retains its shape. The ordered arrangementof water molecules in the crystal lattice is responsible forthe hexagonal (six-sided) structure of ice crystals. Also,the open structure of the ice lattice explains why thedensity of ice is about 90% of the density of liquid waterso that ice floats in water.

FIGURE 2.2

Crystal lattice of ice.

FIGURE 2.1The water molecule is comprised of two hydrogen atoms and oneoxygen atom.

Chapter 2 WATER’S UNIQUE PROPERTIES 21

Research published in the mid-1990s confirmedearlier speculation that an ice crystal is not uniformly

rigid, and this finding may explain why ice is slippery.Apparently the outermost layer of water molecules in anice crystal is much more mobile (molecules vibratefaster) than the inner molecules. While still bonded tothe crystal lattice as a solid, the outermost molecules aremore disordered and behave more like a liquid. Thisoutermost layer of molecules can be thought of as asurface film that reduces the frictional resistance of icemaking it slippery. It is widely but erroneously believedthat objects such as ice skates readily glide across an icesurface because of pressure-induced melting; that is,pressure exerted on the ice suppresses the melting pointof ice and a thin film of water forms on which an objectslides. The surface becomes slippery because liquidwater offers less frictional resistance than ice. Actualmeasurements demonstrate that the magnitude of theapplied pressure is insufficient to lower the melting pointenough to account for the slipperiness of an ice surface(unless the ice is already at a temperature very close tomelting).

Water molecules exhibit much greater activityin the liquid phase than in the solid phase. In the liquidphase, water molecules exhibit vibrational, rotational,and translational motions as illustrated in Figure 2.3.This greater freedom of movement explains why liquidwater takes the shape of its container. In the vaporphase, water molecules have maximum freedom ofmovement and disperse randomly into an empty space orgaseous environment by translation. Water vaporintroduced into a room quickly spreads throughout theroom’s entire volume.

Water readily changes phase, contributing to thedynamic nature of the Earth system. When waterchanges phase heat is either absorbed from theenvironment or released to the environment (Figure 2.4).Melting, evaporation, and sublimation are phase changesthat absorb heat. Phase changes that release heat to thesurroundings are freezing, condensation, and deposition.Before examining these phase changes in detail, we needto review the distinction between heat and temperature.

TEMPERATURE AND HEAT

From everyday experience, we know thattemperature and heat are closely related concepts.Heating a pan of soup on the stove raises the temperatureof the soup and dropping an ice cube into a warmbeverage lowers the temperature of the beverage.

FIGURE 2.3Vibrational, rotational, and translational motion of the watermolecule.

FIGURE 2.4Heat is either absorbed from or released to the environment whenwater changes phase.


Although sometimes used interchangeably, temperatureand heat are distinctly different concepts.

All matter is composed of atoms or moleculesthat are in continual vibrational, rotational, and/ortranslational motion. The energy represented by thismotion is referred to as kinetic molecular energy or justkinetic energy, the energy of motion. In any substance,atoms or molecules actually move about with a range ofkinetic energy. Temperature is directly proportional tothe average kinetic energy of atoms or moleculescomposing a substance. At the same temperature, oneliter of water has the same average kinetic molecularenergy as five liters of water.

Internal energy encompasses all the energy in asubstance, that is, the kinetic energy of atoms andmolecules plus the potential energy arising from forcesbetween atoms or molecules. If two objects are atdifferent temperatures (different average kineticmolecular energies) and are brought into contact, energywill be transferred between the two objects; we call thisenergy in transfer, heat. Heat transferred from an objectreduces the internal energy of that object whereas heatabsorbed by an object increases its internal energy. Heattransferred to or from water brings about a change intemperature or phase.

Differences in temperature rather thandifferences in internal energy govern the direction ofheat transfer. Heat energy is always transferred from awarmer object to a colder object. Heat is not necessarilytransferred from an object having greater internal energyto an object with less internal energy. Consider a hotmarble (at 40 oC) that is dropped into 5 liters of coldwater (at 5 oC). The water has much more internalenergy than the marble, but heat is transferred from thewarmer marble to the cooler water.

The following illustration makes clearer thedistinction between heat and temperature. A cup ofwater at 60 oC (140 oF) is much hotter than a bathtub ofwater at 30 oC (86 oF). That is, the average kineticmolecular energy of water molecules is greater at 60 oCthan at 30 oC. Although lower in temperature, the greatervolume of water in the bathtub means that it containsmore total kinetic molecular energy than does the cup ofwater. If in both cases the water is warmer than itsenvironment, energy (i.e., heat) is transferred from waterto its surroundings. Much more heat energy must beremoved from the bathtub water than from the cup ofwater for both to cool to the same temperature. Hence,

the cup of hot water will cool down to room temperaturemuch faster than the water in the bathtub.

Through the years, scientists have devisedvarious scales that express the temperature of an objectby a number representing the degree of warmth.Temperature scales were originally derived using thefreezing and boiling points of water as fixed points ofreference. Eighteenth century scientists who developedtemperature scales used water for this purpose because itwas readily available and inexpensive. In this book weuse the Celsius temperature scale primarily with theFahrenheit equivalent in parentheses. The Celsius scale,devised by the Swedish astronomer Anders Celsius in1742, has the numerical convenience of a 100-degreeincrement between the melting point of ice (0 oC) and theboiling point of fresh water at average sea level airpressure (100 oC). The United States is one of only afew nations still using the Fahrenheit temperature scale,introduced by the German physicist Gabriel Fahrenheitin 1714. The Fahrenheit temperature scale features a180-degree increment between the melting point of ice(32 oF) and the boiling point of fresh water at average sealevel air pressure (212 oF).

On the Kelvin scale, temperature is the numberof kelvins above absolute zero (a theoretical point whereall molecular motion ceases and no radiation is emitted).Absolute zero corresponds to –273.15 oC and –459.67oF). Nothing can be colder than absolute zero, so theKelvin scale has no negative temperatures. Formulas forconverting among the Celsius, Fahrenheit, and Kelvintemperature scales are given in Table 2.1.

A convenient unit of heat energy is the calorie,defined as the amount of heat needed to raise thetemperature of 1 gram of water 1 Celsius degree(technically, from 14.5 oC to 15.5 oC). (The calorie usedto measure the energy content of food is actually 1000heat calories or 1 kilocalorie.) Although the preferredunit of energy in any form, including heat, is the joule(J), we generally use the calorie (cal) in this book

TABLE 2.1Temperature conversion formulas

oF = 9/5 oC + 32 o

oC = 5/9 (oF – 32 o)K = 5/9 (oF + 459.67)K = oC + 273.15


because of its numerical convenience. Furthermore,thermal characteristics of water were used to define thecalorie. For conversion purposes, one calorie equals4.1868 J and one joule equals 0.239 cal.

FREEZING AND MELTING

Fresh water cooled to 0 oC (32 oF) and subjectedto additional loss of heat energy normally will freeze. Amixture of fresh water and ice has an equilibriumtemperature of 0 oC (32 oF). Adding heat to the mixture(warming) causes ice to melt, whereas removing heat(cooling) causes water to freeze. For that reason, 0 oC iscalled the freezing point of fresh water. As discussed inChapter 5, cloud droplets can cool well below the usualfreezing point while remaining liquid. Such clouddroplets are composed of supercooled water.Supercooled water droplets are common components ofcertain types of clouds and fog. Additionally, substancesdissolved in water suppress the equilibrium temperatureof a mixture of ice and water to temperatures below thefreezing point of fresh water.

The freezing point of seawater varies withsalinity, a measure of the mass (grams) of dissolved saltsin a kilogram of seawater. The original unit of salinitywas parts per thousand (‰). At first, salinity wasdetermined by evaporating seawater and weighing theresidual, and in later years by chemical analysis (i.e.,titration). Today, salinity is based on measurements ofelectrical conductivity of seawater and the standard unitof salinity is the practical salinity unit (psu).Numerically, one psu is essentially equivalent to 1.0 partper thousand. Seawater’s freezing point at a salinity of30 psu is –1.627 oC and –2.196 oC at a salinity of 40 psu.For seawater of average salinity (35 psu), the freezingpoint is –1.9 oC.

When water freezes, latent heat is released tothe environment and for ice to melt an equivalent amountof latent heat is absorbed from the environment (Figure2.4). The term latent means hidden—a reference to thefact that this heat energy is used only to change the phaseof water and not the temperature of the water. If heat isadded to a mixture of ice and water at 0 oC (32 oF), theice-water mixture remains at constant temperature untilall the ice melts. All available heat is used to bring aboutphase change. Whether freezing or melting is takingplace, the latent heat involved is commonly called latentheat of fusion.

Growers sometimes take advantage of the latentheat released when water freezes to protect their cropsfrom potential frost-damage during cold weather.Growers spray their plants with a fine water mist whenthe environmental temperature drops to 0 oC and lower.Mist freezes on the surface of leaves, fruits, andvegetables. Someone might question how a coating ofice could possibly help plants survive subfreezing airtemperatures. As long as ice is actually forming, plantsurfaces will remain at 0 oC because the latent heatreleased during water’s phase change from liquid to solidhelps keep the temperature of the plants from droppingbelow 0 oC. Stabilization of the temperature at or near 0oC often prevents crop damage because the activelygrowing tissues of most crop species are not injuredunless their internal temperature drops to between –5 oCand -1 oC (23 oF and 30 oF). Nonetheless, this freeze-prevention strategy requires careful monitoring. As longas sprinkling and freezing continue, the temperature ofthe ice remains at about 0 oC. If sprinkling werediscontinued before the air temperature climbs to meltingtemperatures, then heat is conducted from the plant to theice and leaf temperatures fall to potentially lethal levels.Also, care must be taken that the ice burden does notbecome so great that plants are damaged by excessiveweight of the ice. For this reason, the sprinkling methodis best suited to protecting low-growing crops such ascucumbers, strawberries, and cranberries, althoughsprinkling is also used in citrus groves.

VAPORIZATION AND CONDENSATION

At the interface between liquid water and air(e.g., lake or sea surface), water molecules continuallychange phase: some crossing the interface from water toair and others from air to water. If more water moleculesenter the atmosphere as vapor than return as liquid, a netloss occurs in liquid water mass. This process is knownas evaporation. Evaporation explains the drying of soiland disappearance of puddles following a rain shower.On the other hand, if more water molecules return to thewater surface as a liquid than escape as vapor, net gain ofliquid water mass results. This process is calledcondensation. Water vapor condenses on the coldsurface of an aluminum beverage can on a humidsummer day. Heat is absorbed from the environmentduring evaporation and released to the environmentduring condensation (Figure 2.4). Heat absorbed or


released during these phase changes is known as thelatent heat of vaporization (or condensation).

All of us have experienced evaporative cooling.We are chilled upon stepping out of a shower orswimming pool. Water droplets evaporate from the skinsurface and absorb heat, lowering the skin’s temperature.The human body uses evaporative cooling to guardagainst an unhealthy rise in the temperature of vitalorgans (e.g., heart, lungs) during hot and humid weather(Chapter 4). We perspire and the evaporation of tinybeads of sweat cools the skin. On a global scale,evaporative cooling is the most important processwhereby heat at Earth’s surface is transferred to theatmosphere (Chapter 10). When water evaporates atEarth’s surface, water vapor moves into the atmospherewhere it subsequently may condense into clouds. Heatthat was absorbed in the evaporation of water at thesurface is later released to the atmosphere duringcondensation. This heat transfer mechanism is veryimportant in powering storms, especially tropical stormsand hurricanes.

Whereas melting or freezing in an ice/watermixture occurs at a specific temperature (0 oC for freshwater), evaporation of water can occur at anytemperature. The latent heat of fusion is 80 calories (335J) per gram at 0 oC (lower for supercooled water). Thelatent heat of vaporization varies with temperature from597 calories (2500 J) per gram at 0 oC (32 oF) to 540calories (2260 J) per gram at 100 oC (212 oF).

SUBLIMATION AND DEPOSITION

At the interface between ice and air (e.g., thesurface of a snow cover), water molecules are alsocontinually changing phase: directly from ice to vaporand from vapor to ice. If more water molecules enter theatmosphere as vapor than make the transition to ice, a netloss of ice mass occurs. Sublimation is the processwhereby ice or snow becomes vapor without firstbecoming a liquid. Sublimation explains the gradualdisappearance of a snow cover even though the airtemperature remains well below freezing. On the otherhand, if more atmospheric water molecules transition toice than move from ice to vapor, a net gain of ice massresults. Deposition is the process whereby water vaporbecomes ice without first becoming a liquid. During acold winter night, the formation of frost on automobilewindows is an example of deposition.

Heat is absorbed from the environment duringsublimation and heat is released to the environmentduring deposition (Figure 2.4). Latent heat involved insublimation or deposition must equal the total amount ofheat absorbed or released during the combined solid-liquid plus liquid-vapor phase changes. Sublimationrequires the latent heats of fusion plus vaporization,known as the latent heat of sublimation. Depositionreleases to the environment an equivalent amount oflatent heat, that is, the latent heat of deposition. Themagnitude of the latent heats of sublimation anddeposition are remarkably uniform, varying only from677 calories (2835 J) per gram at 0 oC to 678 calories(2839 J) per gram at –30 oC.

Why is Water so Unique?

Compared to other naturally occurring substances,water’s thermal properties are unique. For one, water’sfreezing and boiling points are exceptionally high.Based only on water’s molecular weight and the freezingand boiling temperatures of chemically relatedsubstances, fresh water should freeze at about –90 oC(-130 oF) and boil at about –68 oC (-90 oF). Actually,fresh water’s freezing point is 0 oC (32 oF) and its boilingpoint is 100 oC (212 oF) at average sea level pressure.The unusually high values of water’s freezing andboiling temperatures permit water to occur naturally inall three phases in the Earth system (Figure 2.5).Furthermore, an unusually great quantity of heat energyis involved when water changes phase, and a relativelylarge amount of heat is needed to change the temperatureof water. All of these unique thermal properties of waterhave important implications for the functioning of theEarth system.

Why does water exhibit these unusual thermalproperties? Ultimately, the answer lies in the structure ofthe water molecule. Within the water molecule, bondingbetween hydrogen and oxygen atoms involves sharing oftwo electrons, one from each hydrogen atom and anotherfrom the oxygen atom. (An electron is a negativelycharged subatomic particle.) The oxygen atom has astronger attraction for the shared electrons than do thehydrogen atoms so that the oxygen acquires a smallnegative charge and the hydrogen is left with a smallpositive charge. The 105-degree angle formed by thearrangement of the hydrogen-oxygen-hydrogen atomsproduces a charge separation in the water molecule


(Figure 2.1). Molecules having a separation of positiveand negative charges are described as dipolar.

Opposite electrical charges attract so that, liketiny magnets, neighboring dipolar water molecules linktogether. The positively charged (hydrogen) pole of onewater molecule attracts the negatively charged (oxygen)pole of another water molecule; this special attractiveforce is known as hydrogen bonding. Hydrogenbonding is roughly 10 to 50 times weaker than the bondsbetween the hydrogen and oxygen atoms in individualwater molecules. Each water molecule can formhydrogen bonds in three directions because the moleculehas three potential sites for hydrogen bonding.

Hydrogen bonding inhibits changes in theinternal energy of individual water molecules. As heat istransferred to or from water, the accompanying change ininternal energy is unusually small. Recall from earlier inthis chapter that temperature is proportional to averagekinetic molecular energy, part of the internal energy of asubstance. Addition or removal of heat from a volume ofwater is accompanied by relatively little change in watertemperature; that is, the specific heat of water isrelatively high (discussed below). In addition, greateramounts of heat, and unusually high temperatures arerequired for water to reach its melting and boiling points.Hydrogen bonding also means that exceptionally largeamounts of heat energy are needed to change the phaseof water.

In a general sense, for any substance to changephase from solid to liquid or liquid to gas, heat energymust be supplied to break intermolecular forces ofattraction. Van der Waals force is the relatively weakelectrostatic attraction between the atomic nuclei of onemolecule and the electrons of another molecule. Van derWaals forces are stronger in the solid and liquid phases(when molecules are closest together). In the case ofwater, a change in phase from solid (ice or snow) toliquid or from liquid to vapor, requires an addition ofheat energy in amounts sufficient to overcome not onlyvan der Waals forces but also the much strongerhydrogen bonds. For this reason, water has unusuallygreat latent heat requirements.

For water to change from solid (ice or snow) toliquid, heat energy must be supplied to break some of thehydrogen bonds that maintain water in the crystalline(solid) phase (Figure 2.2). This heat is the latent heat offusion, discussed above. Not all hydrogen bonds arebroken as water changes from solid to liquid; that is,numerous small clusters of bonded water moleculespersist into the liquid phase. Considerably greateramounts of heat are required for water to vaporize (byeither evaporating or boiling) because essentially allhydrogen bonds must be broken. For this reason, themagnitude of water’s latent heat of vaporization is morethan seven times the magnitude of water’s latent heat offusion. When phase changes are reversed, that is whenwater vapor changes to liquid (or solid) or liquid changesto ice, unusually large quantities of latent heat arereleased to the environment.

In summary, in the solid and liquid phases,dipolar water molecules are linked by hydrogen bonds.Hydrogen bonding inhibits changes in the internal energyof water so that water requires or releases unusually greatquantities of heat when changing phase. Hydrogenbonding also means that greater additions or losses ofheat accompany a change in water temperature ascompared to other substances.

Water’s Specific Heat

The temperature change associated with an input (oroutput) of a specified quantity of heat varies from onesubstance to another. The amount of heat that will raisethe temperature of 1 gram of a substance by 1 Celsiusdegree is defined as the specific heat of that substance.Joseph Black, a Scottish chemist, first proposed the

Ice

Liquid

Vapor

Triple Point

273.16

6.11

Temperature (°K)

Vap

or P

ress

ure

(mb)

FIGURE 2.5Phase diagram of water. At the triple point, water existssimultaneously as a solid, liquid, and vapor.


concept of specific heat in 1760. The specific heat of allsubstances is measured relative to that of liquid water,which is defined as 1 calorie per gram per Celsius degree(at 15 oC). The specific heat of ice is 0.5 calorie pergram per Celsius degree (near 0 oC). The specific heatsof some other familiar substances are listed in Table 2.2.The variation in specific heat from one substance toanother implies that different materials have differentcapacities for storing internal energy.Upon absorbing the same amount of heat energy, asubstance with a high specific heat experiences a smallerincrease in temperature (warms less) than a substancehaving a low specific heat. Because of hydrogenbonding, water has an unusually high specific heat, infact the highest specific heat of any naturally occurringcommon substance. From Table 2.2, water’s specificheat is about 5 times that of dry sand. One calorie ofheat will raise the temperature of one gram of water by 1Celsius degree whereas one calorie of heat will raise thetemperature of one gram of dry sand by 5 Celsiusdegrees. This contrast in specific heat helps explain whyat the beach in summer the sand feels considerably hotterto bare feet than the water. This also largely explainswhy wet sand feels cooler than dry sand.

Water’s unusually high specific and latent heatsare illustrated by a simple experiment conducted at sealevel (Figure 2.6). A one-gram ice cube initially at –20oC (-4 oF) is heated to 0 oC (32 oF). Every 1-Celsius

degree (2-Fahrenheit degree) rise in temperature of theice cube requires an addition of 0.5 calorie of heat energy(specific heat of ice). A total of 10 calories of heat (20Co × 0.5 cal/Co) are required to warm the one-gram icecube to 0 oC (32 oC). Once it reaches 0 oC (32 oF), thetemperature of the water substance remains constantwhile 80 calories of heat (latent heat of fusion) are addedto completely melt the 1-gram ice cube. Liquid water isthen heated to 100 oC (212 oF). Every 1-Celsius degree(2-Fahrenheit degree) rise in temperature of the 1 gramof liquid water requires an addition of 1.0 calorie of heatenergy (specific heat of water). Assuming that none ofthe water evaporates, a total of 100 calories of heat (100Co × 1.0 cal/Co) are required to warm the one-gram ofliquid water to 100 oC (212 oF). At 100 oC, the watervaporizes requiring an input of 540 calories of heat(latent heat of vaporization).

Water’s relatively great capacity to storeinternal energy is the reason water is a highly effectivecooling agent. Coal-fired and nuclear electric powerplants routinely pipe water from lakes and rivers throughthe facility to absorb and carry off waste heat. Heatedwater is discharged into the source waterway orevaporated from cooling towers or storage lagoons.

Water’s exceptional capacity to store heat alsohas important implications for weather and climate. Alarge body of water (such as the ocean or Great Lakes)can significantly influence the climate of downwindlocalities. The most persistent influence is on airtemperature. Compared to an adjacent landmass, a bodyof water does not warm as much during the day (or insummer) and does not cool as much at night (or inwinter). In other words, a body of water exhibits agreater resistance to temperature change, called thermalinertia, than does a landmass. Whereas the difference inspecific heat between land and water is the major reasonfor this contrast in thermal inertia, differences in heattransport also contribute. Sunlight penetrates water tosome depth and is absorbed (converted to heat) through asignificant volume of water. But sunlight cannotpenetrate the opaque land surface and is absorbed only atthe surface. Furthermore, circulation of ocean and lake-waters transports heat through great volumes of water,whereas heat is conducted only very slowly into soil.The input (or output) of equal amounts of heat energycauses a land surface to warm (or cool) more than theequivalent surface area of a body of water.

Air temperature is regulated to a considerableextent by the temperature of the surface over which air

TABLE 2.2Specific Heats of Some Familiar Substances

Substance Specific Heat (cal per gper Celsius degree)

Water 1.000Ice (at 0 oC) 0.478Wood 0.420Aluminum 0.214Brick 0.200Granite 0.192Sand 0.188Dry aira 0.171Copper 0.093Silver 0.056Gold 0.031

aAt constant volume


resides or travels. Air over a large body of water tends totake on similar temperature characteristics as the surfacewater. Winds blow from sea to land and from land tosea. Places immediately downwind of the oceanexperience much less contrast between average winterand summer temperatures and the climate is described asmaritime. Places at the same latitude but well inlandexperience a much greater temperature contrast betweenwinter and summer and the climate is described ascontinental.

Consider an example of the contrast betweencontinental and maritime climates. The latitude ofStevens Point, WI (44.5 degrees N) is about the same asthat of Newport, OR (44.6 degrees N) so that theseasonal variation in the amount of solar radiationstriking Earth’s atmosphere (due to astronomical factors)is about the same at both places. Stevens Point issituated far from the moderating influence of the oceanand its climate is continental. Stevens Point’s averagesummer (June, July, August) temperature is 19.9 oC(67.9 oF) and its average winter (December, January,

February) temperature is -8.2 oC (17.2 oF), giving anaverage summer-to-winter seasonal temperature contrastof 28.1 Celsius degrees (50.7 Fahrenheit degrees).Newport, on the other hand, is located on the Oregoncoast, immediately downwind of the Pacific Ocean; itsclimate is maritime. The average summer temperature atNewport is 13.6 oC (56.5 oF) and the average wintertemperature is 7.0 oC (44.6 oF), giving an averageseasonal temperature contrast of only 6.6 Celsius degrees(11.9 Fahrenheit degrees).

The moderating influence of the ocean is alsoevident in the contrast in climate between WesternEurope and Eastern North America. At mid-latitudes,prevailing winds blow from west to east so that themaritime influence of the North Atlantic Ocean is muchmore apparent in Western Europe than Eastern NorthAmerica. In the same latitude belt, winters areconsiderably milder in Western Europe than in EasternNorth America. Consider, for example, the contrast inaverage January temperatures for Montreal, Quebec(45.5 degrees N) versus London, England (51.5 degrees

FIGURE 2.6Heating a 1-gram ice cube causes the ice to change phase.


N). In January, the average daily high temperature is-6.1 oC (21 oF) at Montreal and 6.7 oC (44 oF) in London.The January average daily low temperature is -14.4 oC (6oF) at Montreal and 1.7 oC (35 oF) in London. AlthoughLondon is considerably further north than Montreal, itsJanuary temperatures are significantly milder.

The greater thermal inertia of water bodiescompared to landmasses is responsible for twodistinctive seasonal regimes on the Great Lakes (Figure2.7). These regimes are based on the averagetemperature difference between water and land. LakeMichigan’s average surface temperatures are lower thanthe average air temperature over adjacent land surfacesfrom about mid-March into early or mid-August, withthe greatest temperature contrast usually between mid-May and early June. The relatively cool lake surfacechills the overlying warmer air inhibiting thedevelopment of convective weather systems such asthunderstorms. Fog forms over the lake when warmhumid air streams over the cold water surface and coolonshore lake-breezes often develop during the daytimealong the coast.

Lake Michigan’s average surface temperaturesare higher than the average surface temperatures ofadjacent land areas from about late August into middle orlate March. The greatest temperature contrast usuallyoccurs from late November into early December. Therelatively warm lake surface heats and supplies watervapor to the overlying cooler air spurring cloud

development. At times very cold air streams across therelatively warm lake water surface triggering substantiallake-effect snowfall on the downwind shores.

Water as a Solvent

Pure water is unknown in the Earth system, that is, waterfree of all dissolved and suspended materials does notoccur naturally. This is because water is an excellentsolvent (sometimes referred to as the universal solvent).The dipolar nature of the water molecule favors thesolution of both ionic and non-ionic substances. Manyinorganic materials (primarily salts) are bonded ionically,whereas many organic chemicals have non-ionic(covalent) bonds. River water, groundwater, and oceanwater dissolve some of the bedrock or sediment (bothorganic and inorganic) that water contacts. Some Earthmaterials dissolve in water more readily than other Earthmaterials. Formation of huge subsurface caverns isdramatic evidence of the ability of groundwater todissolve limestone (CaCO3).

Consider what happens when a pinch ofcommon household table salt (sodium chloride, NaCl) isadded to water. In salt’s crystalline form, ionic bondshold the positively charged sodium ions (Na+) and thenegatively charged chloride (Cl-) ions together (Figure2.8A). (An ion is an electrically charged atom or

Jan Feb Mar Apr May

Surface watertemperature

Mea

n te

mpe

ratu

re

Jun Jul Aug Sep Oct Nov Dec

Airtemperature

FIGURE 2.7A schematic drawing of the mean monthly temperature variations of Lake Michigan surface waters and adjacent land areas.


molecule.) Once in the water, however, the hydrogen-bonded complexes of water molecules greatly reduce theforce of attraction between oppositely charged sodiumand chloride ions. That is, the strength of ionic bondingbetween sodium and chloride diminishes so that thecompound readily dissociates into sodium and chlorideions (Figure 2.8B). Sodium ions are attracted to thenegatively charged pole of the water molecule andchloride ions are attracted to the positively charged poleof the water molecule. That is, salt dissolves in water.

We are well aware from personal experiencethat solids (such as salt and sugar) dissolve in water.What may be less obvious is that gases also dissolve inwater. An example is carbon dioxide (CO2) dissolvedunder pressure in a carbonated beverage. When we opena can of cola, for example, tiny bubbles of carbondioxide escape, giving the drink its fizz. Some of thedissolved CO2 becomes carbonic acid (H2CO3), a weakacid that is responsible for the beverage’s tart taste. Themost abundant gases dissolved in surface seawater are

nitrogen (N2), oxygen (O2), and carbon dioxide (CO2).Gases (and particles) that dissolve in rain can impact thequality of precipitation. Carbonic acid formed by carbondioxide dissolved in water is important in the chemicaldisintegration of bedrock (chemical weathering). Also,oxygen dissolved in water is essential for the survival ofaerobic aquatic life.

ACID DEPOSITION

Rain and snow are normally slightly acidicbecause they dissolve some atmospheric carbon dioxide,producing weak carbonic acid. Where air is pollutedwith oxides of sulfur or oxides of nitrogen, however,these gases interact with moisture in the atmosphere toproduce tiny droplets of sulfuric acid (H2SO4) and nitricacid (HNO3). These acid droplets dissolve inprecipitation, increasing its acidity. Precipitation thatfalls through such contaminated air may become 200times more acid than normal. Even in the absence ofprecipitation, sulfuric acid droplets convert to tiny acid(sulfate) particles that reduce visibility and may causehealth problems in humans when inhaled. Acid particleseventually settle to Earth’s surface as dry deposition.The combination of acid precipitation and dry depositionis known as acid deposition.

The pH scale encompasses the range of acidityand alkalinity, where pH is a measure of the hydrogenion (H+) concentration (Figure 2.9). By definition, purewater has a pH of 7, which is considered neutral; a pHabove 7 is increasingly alkaline while a pH below 7 isincreasingly acidic. Rain falling through an unpollutedatmosphere dissolves carbon dioxide and has an averagepH of 5.6. On this basis, a pH of 5.6 is taken as thethreshold for acid precipitation; that is, rain or snowhaving a pH less than 5.6 is described as acid rain (orsnow). Small amounts of naturally occurring acids(other than that produced by atmospheric CO2) lower thenormal pH of precipitation closer to 5.0, arguing for arevision of the current criterion for acid rain and snow.The pH scale is logarithmic; that is, each unit incrementon the scale corresponds to a tenfold change in acidity oralkalinity. A drop in pH from 5.6 to 3.6, for example,represents a one hundred fold (10 × 10) increase inacidity.

Gene E. Likens was one of the first scientists tosound the alarm regarding the potential impact of acidrain on the functioning of aquatic ecosystems. Likensand his colleagues reported an increase in the acidity of

FIGURE 2.8Dissociation of sodium chloride (table salt) in water.


rainfall over portions of the United States between 1955and 1973. Measurements made by the NationalAtmospheric Deposition Program in the United Statesand the Canadian Network for Sampling Precipitationlater confirmed and updated Liken’s findings. Ingeneral, the mean annual pH of precipitation is greaterthan 5 west of the Mississippi River and less than 5 eastof the Mississippi. Rain and snow tend to be most acidicin northeastern United States and adjacent portions ofCanada, where average annual pH since 1963 rangesbetween 4.05 and 4.3, with individual storms havingproduced rainfall with pH values as low as 2 to 3. In thisregion, sulfuric acid is suspected to be responsible for55% to 75% of excess acidity.

Most acid deposition is derived from gasesemitted as byproducts of fuel combustion for electricpower, industry, and motor vehicles. Coal burning forelectric power generation is the principal source of sulfuroxides, whereas high-temperature industrial processesand internal combustion engines (in motor vehicles)produce nitrogen oxides.

Where soils are thin and the bedrock is non-carbonate (does not contain carbonate ions), acid rainand snowmelt are not neutralized and may adverselyimpact waterways. (Neutralization is a chemical reactionbetween acid and alkaline substances to produce a saltand water, with the pH of the solution approaching 7.)For this reason, acid deposition is potentially a muchgreater problem where the bedrock consists of granitethan where the bedrock is limestone or dolostone. Aciddeposition lowers the pH of lakes and streams, disruptingthe reproductive cycles of fish. Furthermore, acid rainsleach metals (such as aluminum) from the soil, washingthem into lakes and streams where they may harm fish,microorganisms, and aquatic plants. Lowered pH isresponsible for the decline and even elimination of fishpopulations in some lakes and streams in Norway,Sweden, eastern Canada, and the northeastern UnitedStates.

DISSOLVED OXYGEN

Aquatic organisms, like terrestrial organisms,require oxygen for cellular respiration. Recall fromChapter 1 that cellular respiration is the processwhereby organisms break down food and release energyin a usable form. Essential for survival of aquaticanimals (e.g., fish, shellfish) is dissolved oxygen (DO),oxygen that is dissolved in water. (Note that dissolvedoxygen is not the O in the H2O.) Dissolved oxygenlevels in surface waters (i.e., rivers, lakes, ocean) varyconsiderably depending on water’s capacity to dissolveoxygen, factors governing the rate of transfer of oxygenbetween the atmosphere and the water body, and thesupply of organic matter in the water body.

Oxygen (O2), the second most abundantatmospheric gas, enters an aquatic system through theair-water interface and via photosynthesis by aquaticplants. The maximum amount of oxygen that candissolve in water (saturation concentration) depends ontemperature. Oxygen and most other gases are moresoluble in cold water than warm water. As shown inFigure 2.10, at saturation the amount of dissolved oxygen

FIGURE 2.9Scale of acidity and alkalinity (pH scale).


declines from about 15 milligrams per liter at 0 oC (32 oF)to less than 7 milligrams per liter at 35 oC (95 oF). Allother factors being equal, we would expect less dissolvedoxygen in water bodies in summer than in the coolerseasons. On the other hand, in cold-climate regions,winter ice cover on lakes is a barrier to oxygen transferand results in lower dissolved oxygen levels.

Turbulence in waterfalls, river rapids, and damspillways as well as waves on the surface of open bodiesof water enhance oxygen transfer from air to water(unless the water is already saturated with oxygen). Thecross-sectional profile of a waterway also affects oxygentransfer. A wide, shallow stretch of a river presents alarger surface area for oxygen uptake than does a narrow,deep segment. Furthermore, the amount of oxygenproduced per unit area by photosynthesis depends on thedensity of aquatic plants. (Recall from Chapter 1 thatphotosynthesis is the process whereby green plants usesunlight, water and carbon dioxide to manufacture theirfood and generate oxygen as a byproduct.)

Dissolved oxygen is removed from waterprimarily through cellular respiration by decomposerorganisms. Decomposers are microorganisms, primarilybacteria and fungi, that consume the remains of deadplants and animals as a food source. Fish, shellfish, andzooplankton remove smaller amounts of dissolvedoxygen through cellular respiration. Warm water

enhances decomposer activity, increasing the rate atwhich they remove dissolved oxygen through cellularrespiration.

The amount of organic waste available todecomposers also affects the level of dissolved oxygen inaquatic systems. In most natural waterways, the supplyof organic waste is relatively small and the concentrationof dissolved oxygen usually remains relatively constantat more than 5 milligrams per liter, the minimum levelconsidered essential for survival of the most sensitivefish, such as trout. When saturated with oxygen, naturalwaters typically contain between 10 and 14 milligrams ofdissolved oxygen per liter. Industrial and municipalwastewater, however, usually contain relatively highconcentrations of organic waste and when dumped into awaterway may cause serious depletion of dissolvedoxygen. Abundant organic waste spurs the growth ofdecomposer populations, which consume large quantitiesof dissolved oxygen. Many aquatic organisms, includingthe more desirable species of fish, cannot tolerate theresulting low levels of dissolved oxygen.

The amount of dissolved oxygen thatdecomposers require to break down organic waste in agiven volume of water is known as biochemical oxygendemand (BOD). A major source of BOD is humanwaste. Sewage entering a treatment plant typically has aBOD level of about 100 to 200 milligrams per liter. Thatwater is likely to contain only a few milligrams ofdissolved oxygen per liter so that microbialdecomposition of sewage quickly depletes the water’ssupply of dissolved oxygen. Decomposition of the dailywaste produced by one person requires all the dissolvedoxygen in about 9000 liters (2200 gal) of water. Waterused for processing organic materials such as paper,vegetables, and meat also acquires substantial levels ofBOD. Some concentrated industrial wastewater hasBOD levels that exceed by one thousand times the BODlevels in sewage water. Another source of high-levelBOD waste is runoff from livestock feedlots. Runofffrom farmland where animal waste is spread on fields isalso BOD-rich.

At mid latitudes, discharge of organic waste hasits greatest negative impact on aquatic life during warmsummer months. At that time of year stream flow isoften relatively low so that organic waste is less diluted.Also, as pointed out earlier, oxygen is less soluble inwarm water and decomposers have higher metabolicrates (and greater oxygen demand) at higher watertemperatures.

15

10

5

100 20Temperature (°C)

Dis

solv

ed O

xyge

n at

Sat

urat

ion

(mg/

L)

30 400

FIGURE 2.10Variation in the saturation concentration of dissolved oxygen withwater temperature.


What is the impact of changing dissolvedoxygen levels on aquatic life? Consider an example.Suppose a discharge pipe delivers substantial amounts oforganic waste to a river. Upstream from the pipe, theriver water is oxygen-rich and supports a variety of fish,shellfish, and other organisms. Fishing is fine.Immediately downstream from the pipe, bacteria(decomposers) begin to consume the organic waste.Bacterial populations grow rapidly in direct proportion tothe concentration of organic matter, and soon they areremoving dissolved oxygen at a faster pace than it isnaturally replenished (from the atmosphere and viaphotosynthesis). Dissolved oxygen levels may decline tolevels that can support only species of rough fish (e.g.,carp, gar, and catfish), which tolerate relatively lowdissolved oxygen levels. Fishing is less than desirable.As the organic waste flows downstream, however,decomposers consume more of the waste and as theamount of organic waste declines, BOD levels alsodecrease. Oxygen transfer from the atmosphere plusphotosynthesis slowly reduces the dissolved oxygendeficit. Eventually the rate of oxygen replacementexceeds the rate of removal and dissolved oxygengradually returns to what it was upstream of thedischarge pipe. With restoration of oxygen-rich aquaticconditions, a diverse and desirable community of fishthrives.

Aquatic systems are resilient and graduallyrecover from inputs of organic waste. On occasion,however, a single BOD discharge is so great thatdecomposer populations increase to the point that theyremove all the dissolved oxygen from the water. Such adrastic shift from an aerobic environment (with oxygen)to an anaerobic environment (without oxygen) triggersmajor changes in species composition. Aerobicdecomposers (microorganisms that require oxygen) arereplaced by anaerobic decomposers (those that cannottolerate oxygen). Furthermore, the byproducts of aerobicand anaerobic decay are distinctly different. Decayproducts of aerobic decomposers are mainly carbondioxide (CO2), water, nitrate (NO3

-), and sulfate (SO42-),

which are not usually harmful. On the other hand, someof the decay products produced by anaerobicdecomposers are potentially dangerous to humans andother species. Products of anaerobic decay includemethane (CH4) which is potentially explosive whenmixed with oxygen, hydrogen sulfide (H2S), which isrecognized by its rotten-egg odor, and ammonia (NH3).Surface waters that develop anaerobic conditions become

a putrid, turbid, decaying mess, in stark contrast to theinviting oxygen-rich waters teaming with aquatic life.

Water Density and Temperature

Density is defined as mass per unit volume, usuallyexpressed as grams per cubic centimeter. Objects placedin water that are less dense than water will float to thesurface and materials that are denser than water will sink.Fresh water density varies primarily with temperaturewhereas seawater density varies chiefly with temperatureand salinity. Water is essentially incompressible so thatpressure arising from the weight of water does notsignificantly impact its density.

FRESH WATER

Most substances contract when cooled andexpand when heated; that is, their density increases withfalling temperature and decreases with risingtemperature. As the average kinetic molecular energy ofa substance decreases (i.e., as the temperature falls), thesame number of molecules occupy a progressivelysmaller volume. However, for fresh water, the situationis not quite that simple (Figure 2.11). As the temperatureof fresh water falls steadily from say 25 oC (77 oF), thewater contracts and its density increases. Fresh waterdensity reaches a maximum at about 4 oC (39.2 oF) and,with additional cooling (below 4 oC), the water expands(its density decreases).

.99995

1.00000

.99990

50 10 15Temperature (°C)

Den

sity

(g/

cm3 )

.99985

FIGURE 2.11Variation in the density of fresh water with temperature.


An increasing number of ice-like molecularclusters that continually form and break-up in the liquidis responsible for the anomalous behavior of fresh waterdensity at temperatures below 4 oC. These clustersoccupy more volume than the unorganized watermolecules in the liquid. Recall that ice crystals are openhexagonal (six-sided) structures with widely spacedwater molecules (Figure 2.2). As the water temperaturefalls from 4 oC, the decrease in water density caused bythe presence of ice-like molecular clusters more thanoffsets the increase in water density that accompanies areduction in kinetic molecular activity.

Whereas most liquids contract when theysolidify, ice expands so that the density of ice is about90% of liquid water. (This expansion has someinteresting consequences as anyone who has had frozenwater pipes can attest.) As water freezes, its moleculeschemically bond into an open hexagonal crystallinestructure and its density drops dramatically. The uniquetemperature-density behavior of fresh water explainswhy less dense ice floats in more dense water and lakesfreeze from the top down rather than the bottom up. Ifice were denser than liquid water, ice that forms at theair-water interface would sink and in winter lakes in coldclimates would freeze solid from the bottom up,destroying all aquatic life.

In autumn, lakes begin cooling at the air-waterinterface when the temperature of the overlying air falls

below that of lake surface. Lake-surface water cools andcontracts, reaching maximum density at 4 oC, and sinksto the bottom. This process of surface cooling andsinking is repeated until the entire lake has a uniformtemperature of 4 oC (and uniform density). With furthercooling of surface waters, water density decreases and at0 oC an ice cover begins to form. Once ice forms, itcontracts (becomes denser) as its temperature falls. Butno matter how cold it is, ice remains less dense thanliquid water.

As described in this chapter’s Essay, thetemperature-density relationship of fresh water also playsan important role in the spring and fall turnover of mid-latitude lakes. Lake turnover is necessary forreplenishment of dissolved oxygen in lake waters.

SEAWATER

At constant temperature, the density of seawaterincreases with increasing salinity because the atomicmass of dissolved salts is greater than that of watermolecules. Hence, less dense fresh water floats on moredense seawater. The salinity of seawater also affects thetemperature of maximum density and the freezingtemperature for the same reason: adding dissolvedmaterials such as salt apparently interferes with theformation of ice-like clusters. As shown in Figure 2.12,the temperature of maximum density decreases linearly

FIGURE 2.12Variation in water’s temperature of maximum density versus salinity (in psu).


with increasing salinity. At any salinity less than 24.7psu, the maximum density of water occurs at atemperature above the freezing point (which alsodecreases linearly with increasing salinity). At a salinityof 24.7 psu, the temperature of maximum density is thesame as the freezing temperature (–1.33 oC). At anysalinity greater than 24.7 psu, the maximum-densitytemperature is lower than water’s freezing point. Thedensity of seawater of average salinity (35 psu) variesinversely with temperature; that is seawater densityalways increases with falling temperature and decreaseswith rising temperature.

Freezing of seawater affects the salinity ofsurrounding water. During the freezing process,dissolved substances are largely excluded from thecrystalline structure of growing ice, and are concentratedin the surrounding surface water, lowering its freezingpoint and increasing its density. Even small amounts ofseawater that is trapped between crystals in thedeveloping ice mass eventually migrate downward(toward higher temperature) to the sea. Sea ice thereforefreshens with time.

As we will see in Chapter 8, differences in thedensity of seawater, caused mainly by variations intemperature and salinity, are very important controls ofthe vertical circulation of ocean water. Denser seawatersinks and less dense seawater rises.

Water’s Surface Tension

Water can fill a container to a level that is slightly higherthan the rim, forming a convex surface (Figure 2.13).Individual raindrops hold together as they strike the soilsurface making them more effective weathering agents.An insect known as a water strider literally walks onwater. And a soap bubble forms a nearly perfect sphere.All these phenomena are the consequence of water’sstrong surface tension.

Surface tension, the attraction betweenmolecules at or near the surface of a liquid that acts toreduce the surface area, is another property of waterarising from hydrogen bonding. The density of watermolecules in the air above the water surface is relativelylow so that water molecules in the surface layer are morestrongly attracted (by hydrogen bonding) to each otherand to the layer of molecules immediately below. Infact, among liquids water is second only to mercury inthe strength of its surface tension.

Water adheres to the surface of many differentmaterials including glass, soil, and rock. Consider, forexample, water partially filling a narrow glass tube.Water climbs up the sides of the glass tube until theforces of adhesion to the tube balance the weight of thewater in the tube. The water forms a concave surfacebecause the adhesive forces at the water-glass interfaceare stronger than the attractive forces among moleculesat the center of the tube. This phenomenon is calledcapillarity. Capillary behavior is common in soil,sediment or rock where groundwater rises above thewater table through tiny openings and fissures (Chapter6).

FIGURE 2.13Surface tension allows water to fill a container slightly higher thanits rim.


ConclusionsWater is a common ingredient of the Earth system withvery unusual properties. Most of these properties arisefrom the structure of the water molecule and hydrogenbonding. Water’s capacity to absorb large quantities ofheat has important implications for climate, contributingto the moderating influence of the ocean and large lakeson the average temperature of downwind localities.Water is referred to as the universal solvent because itreadily dissolves inorganic and organic solids so thatchemically pure water does not occur in nature. Wateralso dissolves gases; in fact, dissolved oxygen isessential for the survival of most aquatic organisms. Thedensity of water varies with temperature and the amountof dissolved materials (salinity). The maximum densityof fresh water occurs at about 4 Celsius degrees above itsfreezing point so that freshwater lakes freeze from thetop down. Another unique property of water is itsunusually strong surface tension—an additionalconsequence of hydrogen bonding.

In subsequent chapters, we will revisit theunusual properties of water as we examine the role ofwater in the functioning of the Earth system. We beginour study of water in the Earth system by investigatingthe global water cycle in the next chapter.

Basic Understandings

• Water is the only common constituent of the Earthsystem that occurs naturally in all three phases, assolid (ice or snow), liquid, and vapor at or nearEarth’s surface. These three phases aredistinguished by level of kinetic-molecular activity.

• Water continually changes phase. Melting,evaporation, and sublimation absorb heat from theenvironment whereas freezing, condensation, anddeposition release heat to the environment.

• Temperature is directly proportional to the averagekinetic energy of the atoms or molecules composinga substance. Heat is the name given to the energytransferred from a warmer object to a colder object.

• The freezing point of fresh water is 0 oC. Dissolvedsalts suppress the freezing point of seawater byabout 1.5 to 2 Celsius degrees depending on salinity.Latent heat in the amount of 80 calories (335 J) pergram is released to the environment when waterfreezes at 0 oC and the same amount is absorbed

from the environment when ice melts. This is calledthe latent heat of fusion.

• Heat absorbed during evaporation or released duringcondensation is known as the latent heat ofvaporization (or condensation). Evaporation canoccur at any temperature that water exists as a liquidand requires an input of 597 calories (2500 J) pergram at 0 oC to 540 calories (2260 J) per gram at100 oC. The latent heat of vaporization is greaterthan the latent heat of fusion because all hydrogenbonds must be broken in the phase change fromliquid to vapor.

• Water’s unique thermal properties stem from itsmolecular structure. The water molecule is dipolar;that is, its oxygen atom acquires a small negativecharge and its hydrogen atoms have a small positivecharge. Opposite charges attract so water moleculeslink together with the positively charged hydrogenof one water molecule attracting the negativelycharged oxygen of a neighboring water molecule.This attraction between neighboring watermolecules is known as hydrogen bonding.

• Hydrogen bonding inhibits changes in the internalenergy of water molecules and thereby explainswater’s unusually high freezing and boiling points,latent heats of fusion and vaporization, and specificheat. Hydrogen bonding is also responsible forwater’s excellent dissolving properties.

• Water’s unusually high specific heat has importantimplications for weather and climate. Compared toan adjacent landmass, a water body does not warmas much in summer or during the day and does notcool as much in winter or at night. The moderatinginfluence of the ocean on air temperature is felt inthe maritime climates of downwind localities. Inplaces having a maritime climate winters are milderand summers are cooler than at places in the samelatitude belt having a continental climate.

• Water is an excellent solvent, dissolving both solidsand gases. Salts readily dissociate into ions whenthey enter water. Some gases and solids thatdissolve in rain or snow may turn precipitationexcessively acidic.

• Oxygen dissolved in surface waters is essential foraquatic life. Oxygen enters water directly from theatmosphere and as a byproduct of photosynthesis.Turbulence of the water, the surface area of awaterway, temperature, and density ofphotosynthetic organisms govern the input of


dissolved oxygen. Dissolved oxygen is removedfrom water via cellular respiration. An excessiveload of organic waste can seriously deplete thesupply of dissolved oxygen and stress or eveneliminate aquatic life.

• The maximum density of fresh water occurs at atemperature of about 4 oC and water densitydecreases with further cooling. Ice, being less densethan water, floats and lakes freeze from the topdown. The freezing point of seawater lowers withincreasing salinity, and the temperature of water’smaximum density also drops with increasingsalinity.

• Hydrogen bonding also accounts for water’sexceptionally strong surface tension. Surfacetension explains why water can fill a vessel toslightly above its rim and the capillary movement ofwater.


ESSAY: Lake Turnover

____________________________________________________________________

The unusual way water density varies with temperature plays an important role in the autumn and spring turnover offreshwater lakes in temperate latitudes. Lake turnover is important for organisms living in lakes because this processreplenishes the dissolved oxygen supply of those water bodies. Bright summer sun penetrates a lake to shallow depths,warming the surface layer of water. Meanwhile, the dark deep water does not benefit from solar heating and remainscold, with an average temperature perhaps as low as 4 oC. This stable stratification (layering) of warmer, less densesurface water and colder, denser water at depth persists through most of the summer, with little mixing between the twolayers. The upper layer of the stratified lake is known as the epilimnion and the lower layer is known as thehypolimnion, with the transition zone between the two layers called the thermocline. The thermocline features a rapiddrop in temperature with depth.

In a stratified lake, oxygen that is supplied by the atmosphere and photosynthesis replenishes the dissolvedoxygen supply of only the epilimnion. In the hypolimnion, cellular respiration by decomposers and other organismsremove dissolved oxygen. Without transfer of oxygen from the epilimnion, dissolved oxygen levels in the hypolimnionsteadily diminish. If lake stratification were to persist, the dissolved oxygen concentration could decline to a level thatwould severely stress cold-water fish species (e.g., trout, whitefish) living in the hypolimnion. Fortunately, these fishand other inhabitants of the hypolimnion usually do not die because the lake stratification eventually breaks down andthe dissolved oxygen content of the hypolimnion is replenished.

As summer gives way to autumn, the sun is lower in the sky, daylight becomes shorter, and air temperaturesdrop. Heat is lost from the warm surface waters to the overlying cool air and space. Eventually the temperature of theepilimnion cools to that of the hypolimnion, the thermocline disappears, and the lake has a uniform temperature anddensity from top to bottom. Winds blowing across the lake transport surface water from the upwind shore toward thedownwind shore. Oxygen-depleted water wells up from below on the upwind shore and oxygen-rich surface watersinks at the downwind shore (see Figure). Fall turnover of lake-waters brings oxygen-depleted water from the lakebottom to the surface where it is exposed to the atmosphere and replenished in dissolved oxygen.

Cooling during late autumn and early winter eventually drops the surface water temperature to 4 oC and lower.The lake begins to stratify with the coldest and least dense water at the surface. With continued cooling, the surfacewater temperature drops to 0 oC and a skim of ice forms on the lake surface. In the Northern U.S. and Canada, lakes areice-covered for at least part of the winter and the ice may thicken to tens of centimeters depending on winter airtemperatures. Lake-water temperature typically varies from 0 oC just under the ice to 4 oC at the lake bottom. Ice formsa barrier that prevents exchange of oxygen between the lake-water and atmosphere, and biological processes within thelake (e.g., cellular respiration) cause a gradual decline in dissolved oxygen levels. By the time the ice finally melts inspring, the lake waters may be seriously depleted of dissolved oxygen.

With the arrival of spring, days get warmer and surface waters warm. Eventually, the temperature of the surfacewater reaches 4 oC and once again the temperature and density of the lake become uniform from top to bottom. Just asin autumn, strong winds trigger turnover of lake waters and the mixing that accompanies spring turnover bringsoxygen-depleted water to the surface where oxygen is replenished directly from the atmosphere and via photosynthesis.As spring gives way to summer, lake stratification is reestablished.

In mid latitudes, fall and spring turnovers are essential for re-supplying the dissolved oxygen content of lake-bottom waters and permit the survival of cold-water species. Lake turnovers also play an important role in recyclingnutrients, especially nitrogen and phosphorus compounds, from bottom sediments to the overlying water, where theybecome available to aquatic plants, especially algae. However, in tropical climates surface water temperatures may neverdrop as low as 4 oC and the seasonal variation in air temperature is minimal so that lake-turnover is a rare event. This isone of the reasons why such a tremendous amount of carbon dioxide was able to build up in the bottom waters of LakeNyos, as described in this chapter’s Case-in-Point.


ESSAY FIGURETurnover of mid latitude lakes in spring and fall replenishes the supply of dissolved oxygen in bottom waters.


Documents

CHAPTER 2hopkins/WES/wesch02.pdf · Lake Nyos could absorb a tremendous amount of dissolved CO 2 (perhaps 5 times the maximum per unit volume in a typical carbonated beverage). Culminating