Title: Water - A Maverick Compound
Full Citation
Permanent Link: http://ufdc.ufl.edu/WL00004713/00001
 Material Information
Title: Water - A Maverick Compound
Physical Description: Book
Language: English
Publisher: Life Science Library
Spatial Coverage: North America -- United States of America -- Florida
Abstract: Jake Varn Collection - Water - A Maverick Compound
General Note: Box 28, Folder 13 ( Water - 1966 ), Item 2
Funding: Digitized by the Legal Technology Institute in the Levin College of Law at the University of Florida.
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Bibliographic ID: WL00004713
Volume ID: VID00001
Source Institution: Levin College of Law, University of Florida
Holding Location: Levin College of Law, University of Florida
Rights Management: All rights reserved by the source institution and holding location.

Full Text




IN ITS VERY ORDINARINESS, water is extraordinary. It is everywhere. In
the form of oceans, ice fields, lakes and rivers it covers nearly three quar-
ters of the earth's surface; these bodies together contain more than 324
million cubic miles of water. Beneath the earth, permeating soil and rock,
lies some two million cubic miles more in the form of groundwater. In the
earth's atmosphere is another 3,100 cubic miles of water, mostly vapor.
This abundance of water was present when the earth was born, and
most scientists believe that life was conceived in the planet's primeval
oceans. Water continues to support all life-some very simple organisms
can exist without air, but none can grow without water. It has given rise
to great civilizations, and sometimes it has been responsible for their de-
struction. Over hundreds of millions of years, it has been one of the most
powerful agents in shaping and reshaping the face of the earth. Frozen
into creeping glaciers, it chisels the landscape, gouging out vast depres-
sions and lake beds, switching river courses and carrying soil and boul-
ders over tremendous distances. Falling as rain or in flowing rivers, it
levels mighty mountains, creates broad valleys and steep canyons, and
weathers the hardest rocks. As pounding waves or lapping surf, it gnaws
constantly at coastlines, transforming the profiles of islands and conti-
nents. It determines the climate, forms the soil in which crops and for-
ests take root and, as steam or hydroelectric power, it drives the ma-
chines of modern technology. It is an indispensable ingredient in nearly
all manufacturing processes, from the baking of bread to the manufac-
ture of semiconductors for transistor radios.
As a substance, water is odorless, colorless and tasteless. Yet it plays
an unusual role in the affairs of the world because the seeming insipidity
of its properties is deceiving. As a chemical it is unique. It is a compound
of great stability, a remarkable solvent and a powerful source of chemical
energy. It draws away from most organic substances but is strongly at-
tracted by most inorganic materials, including itself; in fact, its own mol-
ecules cling together more tenaciously than those of certain metals. When
frozen into a solid it expands, instead of contracting as nearly all other
substances do, and the lighter solid floats on the heavier liquid-with as-
tonishing consequences. It can absorb and release more heat than most
common substances. In many physical and chemical properties-such as.
its freezing and boiling temperatures-water is an oddity, an exception
to the rules. And nearly every one of these exceptions is woven into the
fabric of human life, either naturally, as in the process of digestion, or
artificially, as in the steam engine.
All of water's oddities can be traced to its molecular structure. The
combination of two atoms of hydrogen and one of oxygen that comprises
water (H20) forms a surprisingly sturdy molecule. Tremendous energy is
needed to break apart water. In fact, until some 180 years ago water was
believed to be an indivisible element rather than a chemical compound.
The converse of water's stability is even more intriguing. For the same
reason that hydrogen and oxygen atoms resist being pulled apart, they
willingly join together. Any little nudge-striking a match, for instance

)* ""

atoms of hydrogen and an oxygen atom fi
their electron orbits by sharing elections.
hydrogen atom, with one electron spinning
around its nucleus, needs one more electr
to become stable. The larger oxygen atom
with six electrons in its outer shell, needs
two more to fill its orbit. When the three
unstable atoms pool their electrons (below
the result is a stable molecule of water.


-will mate them. The water that "steams" kitchen windows is synthe-
sized in the flames of the stove as hydrogen atoms from the cooking gas
unite with oxygen from the air. Even the human body synthesizes water
-about two quarts per week-in the process of metabolizing its food.
Although an exorbitant sum of energy must be absorbed for the disso-
ciation of water, the same sum is released during its formation. About
14 ounces of pure oxygen and 114 ounces of pure hydrogen, for example,
when combined to produce one gallon of water, provide enough energy
to keep a 60-watt bulb burning for 270 hours. The hydrogen-oxygen reac-
tion is such a good source of energy that it was put to practical use in the
fuel cell that first served as a long-term power generator aboard the Gem-
ini V spacecraft.

The tie that binds the atoms
This prodigious energy comes entirely from the powerful force which
binds two hydrogen atoms to one oxygen atom in the water molecule.
The connection is established between the electrons that make up the
outer parts of the atoms, and is a strong link called a covalent bond.
The hydrogen atom has a single "shell" around its nucleus and al-
though this shell contains a single electron, it has room for two. The out-
er shell of the oxygen atom, with room for eight electrons, contains but
six. These unfilled shells are not stable-their energetic electrons are
/ precariously held and quick to join with others to fill all the room in a
shell. The filled shell is the stable form; once it is created it firmly resists
being torn apart.
The oxygen atom can fill its shell by adding the electrons from two hy-
drogen atoms. At the same time, two electrons from the oxygen atom
join the shells of the two hydrogen atoms, filling them. That is, the three
atoms share their electrons, endowing the water molecule with its re-
markable stability.
The covalent bond is the basis for other characteristics of water-its
prowess as a solvent, for instance. This quality arises from the shape of
11 the molecule. When two hydrogen atoms link with an oxygen atom, the
Each connection produces a lopsided molecule, with the hydrogen atoms held
on on one side of the oxygen atom and at an angle of 105" to each other-
somewhat like the ears on a rabbit's head. One effect of this misshapen
structure is an unequal distribution of electric charges. The hydrogen
'). "side" of the water molecule becomes positively charged, while the oxy-
gen side becomes negatively charged. Thus the molecule becomes a di-
pole-the electrical equivalent of a bar magnet. One side is charged dif-
ferently from the other, just as a magnet has different poles at each end.
A dipole reacts to electrical charges much as a bar magnet does to mag-
netism. Its positive side will be attracted to negative charges, its nega-
tive side to positive charges. And the electrical force resulting from these
charges will counterbalance the influence of other charges. This effect be-
comes noticeable when water touches certain kinds of compounds.
In many substances, the atoms are held together not by covalent

._( j

bonds but by a simple electrical attraction. Table salt-sodium chloride
-is an example. In this compound, each atom carries opposite electrical
charges and therefore attracts the other, to hold the salt molecule to-
gether. The molecule will break apart if this attraction between charged
atoms, or ions, is blocked.
If a water molecule begins to wedge its way between the two ions that
form salt, its dipole effect will cancel some of the electrical attraction be-
tween the ions. The weaker attraction permits the ions to move apart,
making more room for water and its disruptive dipole effect. In this way
water works its way between such ions, cancels their mutual attraction
and separates them. The separated ions are then totally surrounded by
water-dissolved. Many compounds that are held together by this simple
electrical bond, called an ionic bond, readily dissolve in water.
Of all the substances naturally occurring on the face of the earth, water
comes closest to being the universal solvent. It is, in fact, such a good
solvent that perfectly pure water is very rare, if indeed it occurs at all
in nature. The water sipped from a glass may contain, among other
things, an infinitesimal number of glass molecules. The very rain, as it
condenses and descends, dissolves materials such as atmospheric gases.
Wherever it lands, it dissolves still other substances. About half of all
the chemical elements are dissolved in natural waters-some of them
only as traces, some of them in abundance; every trickle, puddle, lake or
sea on earth is an aqueous solution. Seawater is a quite concentrated
one; hundreds of organic and inorganic substances, of metals and non-
metals, make up the sum total of its "salt."

A nonconforming substance
The lopsidedness which endows the water molecule with such potency
as a solvent is also indirectly the source of other exceptional properties.
The most significant of these is often overlooked because it is so simple:
ice floats. By all the rules of physical behavior it should not. Almost
every substance, whether solid, liquid or gas, will shrink in volume as its
temperature goes down. As it contracts, it grows more dense. Thus in its
liquid form it is heavier than as a gas, and its solid form is heavier than
its liquid.
Water follows this rule precisely as a gas and, as a liquid, for 96 per
cent of the way down the temperature range to its freezing point, shrink-
ing steadily all the way. But at 39" F. something happens. As cooling
continues the water expands and gets lighter, and as it freezes into a
solid at 32* F. it becomes still lighter, until it has finally gained about
9 per cent in volume.
However inconvenient this expansion of ice may be for the house-
holder faced with burst water pipes after a sudden winter freeze, it is
fortunate for the rest of the world.
If water behaved like other freezing liquids, there would soon be no
life on the earth, for the water would be irrevocably locked in eternal
ice on the beds of seas, lakes and streams. As it is, when winter comes,

ice forms and floats on the surface of bodies of water, forming an insulat-
ing skin which protects the water beneath from further freezing. If ice
were heavier than water, it would sink to the bottom and gradually
build up from there. Before long the lakes and Arctic seas which now are
only superficially covered with ice would be frozen solid, with perhaps
thin layers of liquid water over the ice where it melted during the warm-
est seasons. Most of the world's water supply would become unusable
to plants, animals or man.
One of the most spectacular effects to follow such a circumstance
would be enormous changes in the world's climate. In this icebound
world the daily fluctuations of temperature would amount to hundreds
of degrees, seasonal variations would be even more radical, and the
winds that blow around the world would be parched and searing. For
the climate of the world is tempered by the ability of water to soak up
and store the sun's heat and to release it slowly.

Heat in freezing water
In the last half of the 18th Century a Scottish chemist, Joseph Black,
observed with precision part of this mechanism. As recounted by the
editor of Black's posthumous Lectures on the Elements of Chemistry:
"Since a fine winter day of sunshine did not at once clear the hills of
snow, nor a frosty night suddenly cover the ponds with a thick cake of
ice, Dr. Black was already convinced that much heat was absorbed and
fixed in the water which slowly trickled from the wreaths of snow; and
on the other hand, that much heat emerged from it while it was as slowly
changing into ice. For, during a thaw, a thermometer will always sink
when removed from the air into melting snow; and during severe frost
it will rise when plunged into freezing water. Therefore, in the first case,
the snow is receiving heat, and in the last, the water is allowing it to
emerge again."
From observations such as this, Black discovered two important prop-
erties of water. First, he recognized its very large heat capacity, its ability
to absorb heat. Heat capacity is expressed in terms of the amount of
heat required to raise a given quantity of a substance by a given number
of degrees. An enormous amount of heat is needed to warm water-a fact
obvious to every housewife who has burned her hand on the pot handle
while the water inside the pot was still cool. The iron of a pot will heat
up nearly 10 times faster than water; it requires that much less heat to
raise its temperature a given number of degrees.
The second of Black's discoveries was the strange fact of so-called
latent heat. It is called latent because it produces no change in tempera-
ture; the heat goes entirely into changing the form of a substance. When
a solid melts, for example, it absorbs a certain amount of heat-the exact
amount depending on the substance-without any increase in tempera-
ture until it is entirely melted. If the process is then reversed, the op-
posite reaction takes place: as the liquid substance freezes it gives up
heat without any lowering of temperature, the amount given up being

_ L

exactly equal to the heat it previously had absorbed without a raising
of temperature.
Water's latent heat is unusually high. To convert ice completely into
water-with no change in temperature-requires as much heat as would
be needed to bring to the boiling point the same amount of tepid water.
This is far more heat than most other common substances require for
melting. Iron, for example, melts at 2,763" F.; once the solid is at
this temperature, the addition of enough heat to melt a pound of water
will melt eight times as much iron. When water again is frozen into ice,
it will in the process give up the same heat it absorbed in melting, and
this released heat can keep the surroundings warm.
Because of water's latent heat capacity, a tubful of water placed in a
greenhouse on a night of freezing cold will act as a reserve supply of heat.
Some of its water will be frozen by morning, but the heat it releases dur-
ing freezing keeps the air inside warmer than the air outside. People
living on the seashore experience milder winter temperatures than those
who live inland, though both may be swept by the same storms. On the
other hand, ice melting in an old-fashioned refrigerator draws heat out
of the surrounding food. Similarly, ice cubes chill a drink not so much
because the ice is cold but because it melts, absorbing heat from the
surrounding fluid in the process.
When a substance evaporates or condenses it gains or loses energy just
as it does when melting or freezing. Raise the liquid's temperature to the
boiling point and there will be a pause during which heat is absorbed,
without a rise in temperature, solely to transform the liquid into gas. If
the gas is later steadily cooled, there will also be a pause in temperature
decline at the boiling point, though heat will continue to be given off.
This pause will endure, and the temperature will remain the same, un-
til all the gas has been liquefied.
The latent heat of evaporation, or its reverse, the latent heat of con-
densation, is greater for water than it is for most other substances. It
takes more than five times as much heat to change boiling-hot water into
steam as it does to bring freezing water to a boil. And condensing steam
releases exactly the same amount of heat in returning to the liquid state.

Energy from the skies
As a result of water's latent heat, each molecule of water vapor in the
atmosphere and each droplet of moisture in a cloud is an airborne
bundle of heat energy. It is the dynamic flow of water vapor in our at-
mosphere that creates global climates and local weather.
The energy locked in water vapor can be observed in the building-up
of puffy cumulus clouds into towering thunderheads on a summer day.
As the water vapor in the clouds cools and condenses, huge amounts of
heat energy are released and the atmosphere inside the clouds boils over
with stormy convection currents. The updrafts and downdrafts inside
a thunderhead reach hurricane force; the total amount of energy released
by a summer thunderstorm is equal to that of a large atomic bomb and

200 -



-170 -

H20 (18)


H2O (18)

span between its freezing and boiling points-is
curiously out of step with chemical theory,
as illustrated by this chart. Substances similar
to water (H20) in structure-H2Te, H2Se,
H2S-descend the scale in a regular pattern:
the lower the molecular weight, the lower and
narrower the temperature range. Water, with
the lowest of the four molecular weights,
belongs at the bottom (pale blue bar). Instead it
is highest of all. Chemists believe that the
strong bonds between the water molecules
themselves make it freeze sooner but boil at a
higher point than would otherwise be normal.

could cause the same destruction if its energy could be concentrated as
it is in a bomb.
Even the temperatures at which water freezes and boils are out of
the ordinary. Water freezes at 32" F. and boils at 212* F. This does not
fit the pattern set by similar compounds. Most related substances boil
and freeze at predictable temperatures: an orderly progression of boil-
ing and freezing points which increase as molecular weight increases. In
a group of four compounds chemically related to water, water is the
lightest and should have the lowest boiling and freezing points; if water
followed the pattern of its chemical sisters it would boil at about -132 F.
and freeze at about -148' F.-there could be no liquid water, but only
steam at temperatures found on earth. Happily for life on this planet,
water is the maverick in the group; its boiling and freezing points are not
the lowest but the highest.
The strange relationship between water and heat can be traced to
details of water's molecular structure. The lopsided molecule, with elec-
trical charges concentrated on opposite sides, is attracted to other mole-
cules with similar distributions of charges. Negative to positive, these
polar molecules link together like so many submicroscopic magnets. This
electrical tie, called a hydrogen bond, most readily connects one water
molecule to another-the positively charged hydrogen side of one mole-
cule hitches to the negatively charged oxygen side of a near neighbor.

The ice crystal's airy pattern

Since each molecule has two hydrogen positive "terminals" but only
one oxygen negative "terminal," the connections build characteristic pat-
terns, such as the six-pointed structures that are familiar in snowflakes.
A snowflake's six-pointed star is one modification of a more common
form taken by ice crystals. This form is created as each water molecule
bonds with four other water molecules, which in turn bond with still
others. Thus a system of connections is set up that results in a solid
crystal shaped like an elongated pyramid. The inside of the pyramid
contains no atoms-it is empty space. As a result it is an airy, light-
weight structure. That is why solid water-ice-is lighter than liquid
water. In solidifying it shapes itself onto an open skeleton established
by hydrogen bonds and held rigid by those bonds.
Only in its solid form are all molecules of water linked together by
hydrogen bonds. And the pattern they establish in the solid crystal
makes ice float. It accounts for the expansion of water frozen into ice
and the odd expansion of liquid water between 32 F. and 39" F.
When ice is heated to the melting point, some of the hydrogen bonds
break and the patterned arrangement of molecules begins to collapse.
When this happens the molecules can move more closely together: the
liquid water is denser than the ice from which it was formed. This in-
crease in density, caused by the packing of molecules, continues from
the melting point of ice, 32* F. to about 39" F. Within this small temper-
ature range some molecules are still bound together in the ice pattern

while others, their hydrogen bonds broken, are free to move around with
increasing speed as the temperature is raised. At 39" F. more hydrogen
bonds are broken, and enough molecules can speed up to cause water to
expand (become less dense) with increasing temperature, the way other
substances do.
The tenacious hydrogen bonds-the strongest type of bonds between
molecules-are also responsible for the other useful oddities of water's
behavior. Their strength is indicated by the large amounts of energy
-heat-that are required to break them. That is why so much heat must
be supplied to raise the temperature of water, and why its freezing and
boiling points are so abnormally high.
But perhaps the strangest result of the hydrogen bond is water's unu-
sual ability to climb inside tubes, in seeming defiance of gravity. Every-
one has observed how the edge of the water in a drinking glass curves
slightly upward, forming a distinct liplike rim as it clings to and climbs
up the sides of the glass. This tendency of liquids to rise along the surface
of a solid material is called capillarity. In a very fine tube capillary ac-
tion is sufficient to lift a column of water against the force of gravity,
sometimes to considerable heights, and it is important to the movement
of water through soil, the feeding of plants from their roots, and the
circulation of the blood.
The tube literally pulls the water up by forming hydrogen bonds be-
tween oxygen or nitrogen atoms in the tube material and molecules of
water. This attraction raises the water edge. But simultaneously bonds
within the water itself are pulling the water surface taut, trying to keep
it flat. These opposing actions work in sequence to lift the water inside
the tube by a sort of hand-over-hand process. First the water edge rises,
then the tension in the water tries to level the surface, and that brings
more water near the edge to be lifted higher. The sequence ends only
when so much water has been raised that its weight pulling downward
balances the capillary force pulling upward on the water edge.

Tension at the faucet rim
The powerful tension that hydrogen bonds create on a water surface
can be seen most clearly at a dripping faucet. The horizontal film of wa-
ter that first appears at the faucet's opening acts as if it were a circular
piece of very thin transparent rubber. Like an elastic membrane, it slow-
ly bulges as the weight of water it encloses grows greater. But it does
not break. Instead it seems at last to tear itself away from the faucet rim
and to snap around a freely falling drop which, if it were not distorted by
air pressure, would be a perfect sphere. Of all possible shapes the sphere
is the one having the smallest surface per unit volume. It is the shape in
which the falling drop can most tightly, closely pull itself together.
There, in the homely shape of a falling drop, are demonstrated the
molecular forces that give water its peculiar properties-those rare
qualities that make it the one substance most important to the affairs of
this planet.

I C1 41~ 4 -'

The Unpredictable

Water Molecule

If water, the most common substance on earth, suddenly
began to behave as its molecular makeup suggests, life
would be overwhelmed by a series of unparalleled disasters.
Blood would boil in the body, plants and trees would wither
and die, and the world would be transformed into an arid
waste. But water molecules are bound together in ways un-
like those of any other compound; for this reason they pos-
sess properties that are unique and paradoxical.
For example, water is one of the very few substances that
are heavier as liquids than as solids. As a liquid, it can
creep uphill despite the force of gravity. Water is so benign
that immensely diversified forms of life can thrive within
it-and so corrosive that, given sufficient time, it will dis-
integrate the toughest metal. Although it seems to change
its form with miraculous ease-sometimes existing simul-
taneously as a solid, a liquid and a gas around the same
river or lake-water actually must yield or absorb prodi-
gious amounts of energy to produce these transformations.
In fact, the energy it would take to melt even a small ice-
berg could drive a large ship across the Atlantic 100 times.

Water appears in all three of its physical states as quickly cool and condense into tiny water drop-
a hot stream of liquid sculpts a jagged hole in lets which make up the cloud of mist rising above
a block of ice. Some of the water molecules im- the ice block. Whenever water takes the form
mediately disperse to form an invisible gas, then of ice, some liquid and gas are always present.

it 4



An Ironclad

Molecular Bond

Hydrogen and oxygen have so great
an affinity for each other that, given
even the slightest nudge, they come
together violently, forming water and
releasing great quantities of energy.
In 1937 the huge dirigible Hinden-
burg exploded over Lakewood, New
Jersey, when its hydrogen, ignited by
a spark, fused with the oxygen in the
air; amid the explosive release of en-
ergy, water was produced.
Conversely, it takes a great deal of
energy to split water into its compo-
nents. In fact, in ancient times wa-
ter was considered a basic, inde-
structible element of the universe.
Not until Henry Cavendish startled
the scientific community in 1783 by
synthesizing the water molecule did

it become clear that the substance is
actually a compound made up of one
part oxygen and two parts hydrogen.
The reason water was long thought
to be a single element was that the
sturdy water molecule remains intact
even when frozen solid or heated to
temperatures at which many other
compounds disintegrate. For the at-
oms of the water molecule are laced
together by powerful bonds (below),
which can be severed only by the
most aggressive agents-such as elec-
tric energy or certain chemicals. One
such chemical is potassium; when
even a small lump of potassium is
dropped into water, it pulls the mole-
cules apart so violently that the con-
tainer of water may actually explode.

Water's unique character is a result of the
bonds that tie its two elements together. The
way in which an oxygen atom is linked to two
hydrogen atoms to form water is shown at top.
The couplings of their atomic particles are akin
to keys slipping into locks; the fit is so perfect

that water is one of nature's most stable com-
pounds. When these bonds between elements
are broken by electric energy or chemicals
(bottom), the oxygen and hydrogen regain their
separate identities and are free to seek alliances
with other elements, creating new compounds.

Splitting the water molecule with electricity
is a traditional laboratory demonstration. The
two columns of liquid water in the picture
above almost fill their respective tubes. The bal-
loons attached to the top of the tubes are de-
flated. When the current is turned on (right),
oxygen atoms break away at the positive ter-
minal, or blue wire, and bubble upward into
the blue balloon. Hydrogen atoms are released
at the negative terminal, the red wire, and rise
into a red balloon. Because two hydrogen at-
oms are released for every oxygen atom. the
red balloon inflates twice as fast. Using stand-
ard household current, it would take about 10
years to disintegrate a bathtub full of water.


* I-,

-' .1* '':

ii. -.4

*; -- :-,

A Network

of Nimble Molecules

Once formed, water molecules join
with each other in a unique way, cre-
ating the liquid latticework shown
in the diagram at far right. The con-
nection between water molecules is
called a hydrogen bond. When water
is in the form of ice, these bonds
hold the molecules in a more or less
rigid pattern. But in a liquid state
this structure gives way to a chaotic
molecular square dance in which

groups of molecules take turns whirl-
ing about with one another, breaking
their bonds, and finding new groups
to form partners with.
When water is heated the pace in-
creases until the bonds, no longer
able to keep their partners at arm's
length, snap, and the molecules fly off
as gas. It is these bonds that pull wa-
ter's surface into a taut sheet-a phe-
nomenon known as surface tension.

A waterdrop, shown above forming on the lip
of a faucet, is given its shape by the hydrogen
bonds pulling its molecules inward toward the
center of the drop-one manifestation of sur-
face tension. Owing to the force of gravity, the
drop initially takes a tear shape; it becomes
spherical as it falls, and ultimately flattens out
before hitting the sink because of air resistance.

In a liquid state a water molecule can establish
hydrogen bonds with four of its immediate
neighbors (right) except at the surface, where
there are no water molecules above it. As is
shown in the diagram at far right, surface mole-
cules form bonds only below and to the sides.

The surface of the water in a tumbler can sup-
port amazing weight-in this case a heavy
metal grid-if the object is flat enough to take
full advantage of surface tension. If the grid had
been placed in the water edge first, however.
its weight would have been concentrated on too
few hydrogen bonds, and it would have sunk.



,. < -*i. =,r,= ...

r"*--*B ,i ,, ^ *'

Mi~il~faiaiiiliiiii'liir* i t ^ V ~ lias

Water's bag of tricks bulges with sur-
prises. One of these is water's ability
to creep uphill under certain condi-
tions. Without this characteristic,
known as capillary action, the flow of
nutriments to plants and trees would
stall in the soil, and blood, which is
largely water, would never complete
its circuit of the body.
The explanation for this phenome-
non lies in the nature of water mole-
cules. Bound to each other in almost
every direction, they also bind to a
variety of other substances, such as
glass, clay or soil. In fact, almost any

solid that has oxygen in it will lure
the hydrogen in water. Thus, the sur-
face of the dyed water in the tubes
above is like a chain made up of hy-
drogen links. When the molecules at
the edge reach for and adhere to the
molecules of glass just above them,
they haul the rest of the chain along
with them. The surface, in turn, pulls
the entire body of water to a new
level. The molecules at the edge now
repeat the process, and the water
smoothly continues its ascent. It
ends only when the downward pull
of gravity is too great to overcome.

The Climbing


Rising by capillary action, water slowly ascends
in a series of tubes. It reaches higher levels in
the narrower tubes, which contain less water
for each molecule at the edge to lift. Doubling
the diameter of the tube adds twice the edge
for the molecules to adhere to-but it also
adds four times as much water to be hauled.

Attracted by the oxygen atoms in glass, water
molecules adhere to the sides of a glass tube,
stretching the surface into a crescent shape.
As indicated by the network of hydrogen bonds,
the edge molecules pull the others with them.

The Universal


Water is close to being the all-purpose
chemical solvent; given enough time,
it will dissolve almost any inorganic
substance. In fact, about half of the
known elements are found dissolved
in the earth's waters.
Without water's property of solu-
bility, nutrition could not go on: all
living organisms depend on water to
dissolve the substances they feed on.
The roots of plants cannot absorb
food in the soil unless it is in solution,
and humans' food must be dissolved
before it can enter the bloodstream.
Water molecules in contact with
foreign substances act like cowboys
cutting cattle from a herd-they force
their way between clusters of parti-
cles, break them apart and hold them
at bay. Water's capacity for such ac-
tion is staggering: a gallon of water
(eight pounds) dissolves 70 pounds
of the fertilizer ammonium nitrate.


A solid compound (purple) dropped into liq-
uid water is quickly broken up by water mole-
cules (blue), which squeeze between the solid
particles, separate them from one another and
surround the liberated particles with a protec-
tive shield that prevents them from regrouping.

Each tray of this scale holds a test tube of water
plus six chemicals, but in one case the chemi-
cals have been added to the water. Since the
amounts of chemicals are equal, the trays bal-
ance-but, surprisingly, the volume of water
at right is not increased by the added matter.


A granulated solid disintegrates as its surface
molecules, broken away from the main body,
are surrounded by molecules of water. As wa-
ter attacks the solid in this fashion, streamers of
dissolving particles appear and the solid then
changes its state as surely as if it were melted.

Equipping Water

to Carry Current

There is a good chance that a man
who plugs his electric razor into a
socket while standing in the bath will
receive a walloping jolt of current.
Yet, oddly enough, water itself is a
very poor conductor of electricity; in
fact, no current at all can pass through
it when it is pure and distilled. But
when there are impurities in water
-as is most often the case-the liquid
is endowed with the properties it
needs to conduct electricity.
The reason is that electric current
requires free charged particles to car-
ry it through a medium such as wa-
ter. These particles (usually traces of
dissolved salts) are abundant in im-
pure water. In distilled water the
atoms of oxygen and hydrogen are so

perfectly fitted together that such
particles do not exist.
The experiment shown below dem-
onstrates how water is transformed
into a conductor. In the beaker at left,
distilled water has been substituted
for a length of wire that would com-
plete an electric circuit. Since there
are no free charged particles in the
water, the electric energy is halted
and cannot complete its journey to
the light bulb. When impurities are
added (center and right), the electric
energy has the material it needs to
bridge the gap between the two ter-
minals. The impurities break down as
they dissolve, providing free charged
particles. Set in motion by the elec-
tric energy, they complete the circuit.

To complete an electric circuit, current must pass through the water between the two wires in the beaker (left). When salt (shown dyed)




The three panels at right illustrate the way im-
purities increase water's capacity to conduct
electricity, as shown in the photographs below.
At left are pure water molecules. At center salt
crystals are introduced. They immediately dis-
solve, yielding ions, or particles, of sodium (tan)
and chlorine (yellow), which are encircled by
water molecules. As indicated by the arrows.
the negatively charged chlorine ions head for
the positive terminal; the positively charged so-
dium heads for the negative terminal. As more
salt is added (right), more ions are set in motion.

is added (center) to provide carriers, the current can flow and light the bulb. More salt (right) produces more current and a brighter glow.


A Lightweight

among Solids

Water, which often appears to fol-
low a set of natural laws all its own,
behaves most outlandishly when it
forms ice. For one thing, unlike most
other compounds, it is lighter in this
solid form than it is as a liquid. As
a result, it floats when it freezes. If
this did not happen in nature and ice
were heavier than water, it would
continuously sink to the bottom,
where the sun's rays could not melt
it. Slowly an ice pack would build
upward until the world's oceans, riv-
ers and lakes became frozen solid.
Even as it changes from a liquid

As w .in ary llud molecule; of hIeaiec w ~ler
i elltt area aboate are *.ori .ser3bl\ less, aense
Ihar. cooler moleiiIle; YVei he mosi extreme
di.leren,:s in warer denstl oi:iur vilh.n a de
gree ofl acn oiher Ine lead, dense waler mole
:cuIs are Iho e in Ce ith e bEIOLI/I [he mosI
denp5e 3re Ihie in I[e I[rea3rer, ol exlremelv
cold wal'r ibrotinl Ireshly smelled oilj of ice

Heated from below by a Bunsen burner, water
at the bottom of a tank grows lighter and quick-
ly rises to the top. But as it rises, moving far-
ther away from the source of heat, it immedi-

ately begins to cool. When it reaches the sur-
face, the warmer water behind it nudges it out
of the way. Still cooling and becoming more
dense, the water falls, creating a circular flow.

to a solid, water acts contrary to ex-
pectations. At first it follows the uni-
versal pattern of cooling: it contracts,
and grows heavier and more dense.
But when cooled below 39" F., it
suddenly begins to expand and grow
lighter and less dense. The reason
for this odd turnabout lies, again,
in the hydrogen bonds that exist
between water molecules. As they
cool, the molecules slow up and be-
gin crowding together. At 32" F., the
bonds bring them to a halt and fix
them at arm's length from one an-
other in lightweight crystals of ice.


A piece of ice, lighter and less dense than liq- forms the streamers shown radiating from the
uid water, floats on the surface of a water tank. ice block. This brief period of great density oc-
But as the ice melts, relinquishing its airy crys- curs because as ice melts, the hydrogen bonds
talline structure, it becomes more dense and between molecules begin to collapse and the

molecules crowd closely together. At 390 F.
they reach maximum density. Above 39 F. the
heat sets the molecules moving more rapidly
and the water expands and becomes lighter.

A Compound

Slow to Boil

Homeowners whose water pipes have
burst when the temperature sudden-
ly fell need no other proof that wa-
ter, unlike most liquids, releases tre-
mendous energy when it freezes.
Conversely, it must absorb a great
deal of energy-in the form of heat-
before its temperature is raised even
slightly. An iron kettle used to boil
water will be blistering hot long be-
fore the water in it is lukewarm. This
property accounts for water's use as a
cooling agent in automobile engines:
it soaks up an enormous amount of
heat without boiling.
In this manner, large bodies of
water and the moisture in the at-
mosphere can regulate extremes of


A violent explosion (right) can occur when wa-
ter is frozen in a confined space. A quarter-inch-
thick cast-iron container-shown in cross sec-
tion at the top of the page-is filled with water
(above, left) and placed in a beaker of dry ice
and alcohol (above). As the water freezes and
expands, a tremendous amount of energy is ex-
erted against the walls of the container. Finally
the container explodes, hurling some frag-
ments deep into a steel door 20 feet away.

The water molecules in ice (left) are held in a
relatively rigid geometric pattern by their hy-
drogen bonds, producing an open, porous struc-
ture. In liquid water (far left), which has fewer
bonds, more molecules can occupy the same
space, making liquid water more dense than ice.

temperature, absorbing heat on hot
days, and giving off heat on cold days.
Where there is little natural water,
as on the desert, temperatures can
range from a searing 140" F. to well
below freezing at night.
Before the temperature of a sub-
stance can be raised, its molecules
must be prodded into vigorous mo-
tion. But in water molecules, the firm
grip of the hydrogen bonds must first
be loosened-a task which requires
considerable amounts of heat. If the
hydrogen bonds did not put up such
firm resistance, water would boil at
temperatures lower than those at the
North Pole, and all of the world's
water would immediately evaporate.


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