A water molecule consists of an oxygen atom and two hydrogen
atoms, which are attached at an angle of 105°. Not shown are two pairs of
electrons on the bottom that form a similar angle in a plane perpendicular to
this view. This asymmetrical arrangement accounts for the many unusual
properties of water, such as the fact that it expands when it freezes.
Water, common name applied
to the liquid state of the hydrogen-oxygen compound H2O. The ancient
philosophers regarded water as a basic element typifying all liquid substances.
Scientists did not discard that view until the latter half of the 18th century.
In 1781 the British chemist Henry Cavendish synthesized water by detonating a
mixture of hydrogen and air. However, the results of his experiments were not
clearly interpreted until two years later, when the French chemist Antoine
Laurent Lavoisier proved that water was not an element but a compound of oxygen
and hydrogen. In a scientific paper presented in 1804, the French chemist
Joseph Louis Gay-Lussac and the German naturalist Alexander von Humboldt
demonstrated jointly that water consisted of two volumes of hydrogen to one of
oxygen, as expressed by the present-day formula H2O.
Almost all the hydrogen
in water has an atomic weight of 1. The American chemist Harold Clayton Urey
discovered in 1932 the presence in water of a small amount (1 part in 6000) of
so-called heavy water, or deuterium oxide (D2O); deuterium is the
hydrogen isotope with an atomic weight of 2. In 1951 the American chemist
Aristid Grosse discovered that naturally occurring water contains also minute
traces of tritium oxide (T2O); tritium is the hydrogen isotope with
an atomic weight of 3.
|
PROPERTIES OF WATER
 |
Temperature
Change Versus Heat Added: Water
The graph represents the temperature change that occurs when
heat is added to water. At 0° C and at 100° C, you can add heat to water
without changing its temperature. This “latent heat” breaks bonds that hold the
molecules together but does not increase their kinetic energy. Note that
approximately seven times more heat must be added to evaporate one gram of
water than to melt it. This is represented by the relative lengths of the horizontal
portions of the graph. The slopes of the inclined lines represent the number of
degrees that the temperature changes for each calorie of heat that is added to
one gram. The reciprocal of this number is the amount of heat that must be
added to make the temperature of one gram change by one degree. This is called
the specific heat.
Pure water is an odorless,
tasteless liquid. It has a bluish tint, which may be detected, however, only in
layers of considerable depth. Under standard atmospheric pressure (760 mm of
mercury, or 760 torr); the freezing point of water is 0° C (32° F) and its
boiling point is 100° C (212° F). Water attains its maximum density at a
temperature of 4° C (39° F) and expands upon freezing. Like most other liquids,
water can exist in a supercooled state; that is, it may remain a liquid
although its temperature is below its freezing point. Water can easily be
cooled to about -25° C (-13° F) without freezing, either under laboratory conditions
or in the atmosphere itself. Supercooled water will freeze if it is disturbed,
if the temperature is lowered further, or if an ice crystal or other particle
is added to it. Its physical properties are used as standards to define the
calorie and specific and latent heat and in the metric system for the original
definition of the unit of mass, the gram.
Hydrogen
Bonding in Water
Hydrogen bonds are chemical bonds that form between
molecules containing a hydrogen atom bonded to a strongly electronegative atom
(an atom that attracts electrons). Because the electronegative atom pulls the
electron from the hydrogen atom, the atoms form a very polar molecule, meaning
one end is negatively charged and the other end is positively charged. Hydrogen
bonds form between these molecules because the negative ends of the molecules
are attracted to the positive ends of other molecules, and vice versa. Hydrogen
bonding makes water form a liquid at room temperature.
Water is one of the best-known
ionizing agents. Because most substances are somewhat soluble in water, it is
frequently called the universal solvent. Water combines with certain salts to
form hydrates. It reacts with metal oxides to form acids. It acts as a catalyst
in many important chemical reactions.
|
OCCURRENCE
|
Water is the only substance
that occurs at ordinary temperatures in all three states of matter, that is, as
a solid, a liquid, and a gas. As a solid, or ice, it is found as glaciers and
ice caps, on water surfaces in winter, as snow, hail, and frost, and as clouds
formed of ice crystals. It occurs in the liquid state as rain clouds formed of
water droplets, and on vegetation as dew; in addition, it covers three-quarters
of the surface of the earth in the form of swamps, lakes, rivers, and oceans.
As gas, or water vapor, it occurs as fog, steam, and clouds. Atmospheric vapor
is measured in terms of relative humidity, which is the ratio of the quantity
of vapor actually present to the greatest amount possible at a given
temperature. Water occurs as moisture in the upper portion of
the soil profile, in which it is held by capillary action to the particles of
soil. In this state, it is called bound water and has different characteristics
from free water. Under the influence of gravity, water accumulates in rock
interstices beneath the surface of the earth as a vast groundwater reservoir
supplying wells and springs and sustaining the flow of some streams during
periods of drought.
|
WATER IN LIFE
|
Water is the major constitutent
of living matter. From 50 to 90 percent of the weight of living organisms is
water. Protoplasm, the basic material of living cells, consists of a solution
in water of fats, carbohydrates, proteins, salts, and similar chemicals. Water
acts as a solvent, transporting, combining, and chemically breaking down these
substances. Blood in animals and sap in plants consist largely of water and
serve to transport food and remove waste material. Water also plays a key role
in the metabolic breakdown of such essential molecules as proteins and
carbohydrates. This process, called hydrolysis, goes on continually in living
cells.
|
NATURAL WATER CYCLE
 |
Water
Cycle
Hydrology is the science
concerned with the distribution of water on the earth, its physical and
chemical reactions with other naturally occurring substances, and its relation
to life on earth; the continuous movement of water between the earth and the
atmosphere is known as the hydrological cycle. Under several influences, of
which heat is predominant, water is evaporated from both water and land
surfaces and is transpired from living cells. This vapor circulates through the
atmosphere and is precipitated in the form of rain or snow.
On striking the surface
of the earth, the water follows two paths. In amounts determined by the
intensity of the rain and the porosity, permeability, thickness, and previous
moisture content of the soil, one part of the water, termed surface runoff,
flows directly into rills and streams and thence into oceans or landlocked
bodies of water; the remainder infiltrates into the soil. A part of the
infiltrated water becomes soil moisture, which may be evaporated directly or
may move upward through the roots of vegetation to be transpired from leaves.
The portion of the water that overcomes the forces of cohesion and adhesion in
the soil profile percolates downward, accumulating in the so-called zone of
saturation to form the groundwater reservoir, the surface of which is known as
the water table. Under natural conditions, the water table rises intermittently
in response to replenishment, or recharge, and then declines as a result of
continuous drainage into natural outlets such as springs.
|
COMPOSITION
|
Because of its capacity
to dissolve numerous substances in large amounts, pure water rarely occurs in
nature.
During condensation and
precipitation, rain or snow absorbs from the atmosphere varying amounts of
carbon dioxide and other gases, as well as traces of organic and inorganic
material. In addition, precipitation carries radioactive fallout to the earth's
surface.
In its movement on and
through the earth's crust, water reacts with minerals in the soil and rocks.
The principal dissolved constituents of surface and groundwater are sulfates,
chlorides, and bicarbonates of sodium and potassium and the oxides of calcium
and magnesium. Surface waters may also contain domestic sewage and industrial
wastes. Groundwaters from shallow wells may contain large quantities of
nitrogen compounds and chlorides derived from human and animal wastes. Waters
from deep wells generally contain only minerals in solution. Almost all
supplies of natural drinking water contain flourides in varying amounts. The
proper proportion of flourides in drinking water has been found to reduce tooth
decay. Seawater contains, in addition to concentrated amounts of
sodium chloride, or salt, many other soluble compounds, as the impure waters of
rivers and streams are constantly feeding the oceans. At the same time, pure
water is continually lost by the process of evaporation, and as a result the
proportion of the impurities that give the oceans their saline character is
increased.
|
WATER PURIFICATION
|
Suspended and dissolved
impurities present in naturally occurring water make it unsuitable for many
purposes. Objectionable organic and inorganic materials are removed by such
methods as screening and sedimentation to eliminate suspended materials;
treatment with such compounds as activated carbon to remove tastes and odors;
filtration; and chlorination or irradiation to kill infective microorganisms.
In aeration, or the saturation
of water with air, water is brought into contact with air in such a manner as
to produce maximum diffusion, usually by spraying water into the air in
fountains. Aeration removes odors and taste caused by decomposing organic
matter, and also industrial wastes such as phenols and volatile gases such as
chlorine. It also converts dissolved iron and manganese compounds into
insoluble hydrated oxides of the metals which may then be readily settled out.
Hardness of natural waters
is caused largely by calcium and magnesium salts and to a small extent by iron,
aluminum, and other metals. Hardness resulting from the bicarbonates and
carbonates of calcium and magnesium is called temporary hardness and can be
removed by boiling, which also sterilizes the water. The residual hardness is
known as noncarbonate, or permanent, hardness. The methods of softening
noncarbonate hardness include the addition of sodium carbonate and lime and
filtration through natural or artificial zeolites which absorb the
hardness-producing metallic ions and release sodium ions to the water. Sequestering agents in detergents serve to
inactivate the substances that make water hard.
Iron, which causes an
unpleasant taste in drinking water, may be removed by aeration and
sedimentation or by passing the water through iron-removing zeolite filters, or
the iron may be stabilized by addition of such salts as polyphosphates. For use
in laboratory applications, water is either distilled or demineralized by
passing it through ion-absorbing compounds.
|
WATER
DESALINIZATION
 |
Water
Desalinization Technique
Flash evaporation is the most widely used method of water desalinization.
The seawater is heated and then pumped into a low-pressure tank, where the
water is partially vaporized. The water vapor is then condensed and removed as
pure water. This process is repeated many times (three stages are shown). The
remaining liquid, called brine, contains a large amount of salt and is removed
and often processed for minerals. Note that the incoming seawater is used to
cool the condensers in each evaporator. This design conserves energy since the
heat released when the vapor condenses is used to heat the next batch of
seawater.
To meet the ever-increasing
demands for fresh water, especially in arid and semiarid areas, much research
has gone into finding efficient methods of removing salt from seawater and
brackish waters. In the U.S., desalinization research is directed by the Bureau
of Reclamation, Department of the Interior. Several processes are being
developed to produce fresh water cheaply.
Three of the processes
involve evaporation followed by condensation of the resultant steam and are
known as multiple-effect evaporation, vapor-compression distillation, and flash
evaporation. The last-named method, the most widely used, involves heating
seawater and pumping it into lower-pressure tanks, where the water abruptly
vaporizes (flashes) into steam. The steam then condenses and is drawn off as
pure water. In 1967, Key West, Florida, opened a flash-evaporation plant and
thus became the first city in the U.S. to draw its fresh water from the sea.
Freezing is an alternate
method, based on the different freezing points of fresh and salt water. The ice
crystals are separated from the brine, washed free of salt, and melted into
fresh water. In another process, called reverse osmosis, pressure is used to
force fresh water through a thin membrane that does not allow the minerals to
pass. Reverse osmosis is still undergoing intensive development.
Electrodialysis is being used to desalt brackish waters. When salt dissolves in
water, it splits into positive and negative ions, which are then removed by
electric current through anion and cation membranes, thus depleting the salt in
the product water. Although developmental work on electrodialysis is
continuing, a number of commercial plants are in operation. In 1962 Buckeye,
Arizona, became the first town to have all its water supplied by its own
electrodialysis-desalting plant, which provides about 2,460,000 liters (about
650,000 gallons) of water daily at a cost of about $1 per 6300 liters (1670 gallons).
One major problem in desalinization
projects is the cost of producing fresh water. Using conventional fuels, plants
with a capacity of 3.8 million liters (1 million gallons) per day or less
produce water at a cost of $1 or more per 3800 liters (1000 gallons). More than
500 such plants are in operation, with a total capacity of nearly 473 million
liters (nearly 125 million gallons) a day; however, their high costs limit
their use to areas of great water scarcity. Water from conventional sources,
such as wells and reservoirs, is sold for less than 30 cents per 3800 liters
delivered to the home, and water for irrigation is usually priced at less than
5 cents per 3800 liters. The dual-purpose atomic power and water-desalting
plants now being planned are designed to produce fresh water for between 20 and
30 cents per 3800 liters.
Most experts expect more
immediate results from efforts to purify mildly brackish water that contains
between 1000 and 4500 parts per million of minerals, compared to 35,000 parts per
million for ocean water. Because water is potable if it contains fewer than 500
parts per million of salt, the cost of desalting brackish water is
correspondingly less than it is for desalting seawater.
