Water and the Fitness of the Environment
The Molecule that Supports All of
Life
 Water is the substance that makes possible life as we
know it here on Earth.
 All organisms familiar to us are made mostly of water
and live in an environment dominated by water.
Life on Earth began in Water
 Life evolved in water for 3 billion years before
spreading onto land. Modern life, even terrestrial life,
remains tied to water.
All Living Organisms Require Water
 Humans can survive several weeks without food, but
only one week without water.
 The abundance of water is the reason Earth is
habitable.
The Structure of Water Leads to its
Properties
 Because water is a polar molecule, the two ends have
opposite charges.
 The polarity of water molecules results in hydrogen
bonding—a very important property in biology.
Hydrogen Bonds Form Between
Water Molecules
 The charged regions of a polar water molecule are
attracted to oppositely charged parts of neighboring
molecules.
 Each molecule can hydrogen-bond to multiple
partners, and these associations are constantly
changing.
Four Emergent Properties of
Water are Important for Life
 Cohesion
 Ability to moderate temperature
 Expansion upon freezing
 Versatility as a solvent
Cohesion
 Cohesion-water is polar,
so water is attracted to
water.
 Linkages between water
molecules make water
more structured than
other molecules
 The collective hydrogen
bonds hold the substance
together, a phenomenon
called cohesion
Adhesion
 Adhesion= the clinging
of one substance to
another
 Example: Water will
cling to any substance
that has a charge (either
+ or -) Why does water
do this?
Water Movement in Plants is due
to both Cohesion & Adhesion
 Evaporation of water
from leaves pulls water
upward from the roots
 Cohesion helps hold
together the column of
water within the cells
 Adhesion of water to
cell walls by hydrogen
bonds helps resist the
downward pull of gravity
Surface Tension
 Related to cohesion is
surface tension, a
measure of how difficult
it is to stretch or break
the surface of a liquid.
 Water has a greater
surface tension than
most other liquids.
 Water behaves as though
it is coated with an
invisible film.
Water molecules are held together in
a network of hydrogen bonds.
Moderation of Temperature
 Water moderates air
temperature by
absorbing heat from air
that is warmer and
releasing stored heat to
air that is cooler.
Temperature= a measure of
heat intensity that represents
the average kinetic energy of
the molecules
When 2 objects of different
temperatures are brought
together, heat passes to the
cooler object
Fig. 3-5
Los Angeles
(Airport) 75°
70s (°F)
80s
San Bernardino
100°
Riverside 96°
Santa Ana
Palm Springs
84°
106°
Burbank
90°
Santa Barbara 73°
Pacific Ocean
90s
100s
San Diego 72°
40 miles
By absorbing or releasing heat, oceans moderate coastal climates. In this
example from an August day in Southern California, the relatively cool ocean
reduces coastal temperatures by absorbing heat.
Water’s High Specific Heat
 Specific Heat=the
amount of heat that
must be absorbed or lost
for 1 gram of that
substance to change its
temperature by 1 degree
Celsius.
 Water has a specific heat
of 1, much greater than
most other substances
The specific heat of water is 1
calorie/gram/degree Celsius
High Specific Heat of Water—
Benefits for Living Things
 Water must absorb a
great deal of heat in
order to change its
temperature even one
degree.
 Near large bodies of
water, the temperature is
greatly moderated.
 High specific heat of
water stabilizes ocean
temperatures.
The water that covers Earth’s
surface moderates its
temperatures, allowing life
Evaporative Cooling
 Water is held together
with hydrogen bonds, so
it resists evaporation
until a lot of heat is
added (water has a high
heat of vaporization)
 As water evaporates, the
remaining liquid is
cooler because the
molecules with the most
energy have left.
Ice Floats!
 In ice, each molecule is
hydrogen-bonded to four
neighbors in a 3dimensional crystal
 Because the crystal is
spacious, ice has fewer
molecules than an equal
volume of liquid water.
 Ice is less dense, so it
floats on top of liquid
water
Fig. 3-6a
Hydrogen
bond
Ice
Hydrogen bonds are stable
Liquid water
Hydrogen bonds break and re-form
Water is the Solvent of Life
 Water is a very versatile
solvent due to its polarity
 Water will dissolve any ionic
compound easily—the – and +
charges will be attracted to the
water molecules -/+ sides.
 Water will dissolve nonionic
(but polar) molecules by
surrounding them, forming
Hydrogen bonds with them.
Hydrophilic & Hydrophobic
 Hydrophilic=any
Oil spills such as the one that killed
this bird can be very hazardous to
ocean life.
substance that has an
affinity for water
(anything with a + or –
charge)
 Hydrophobic=any
substance that does not
have an affinity for water
(anything that is
nonionic, nonpolar)
Solute Concentrations in
Aqueous Solutions
 Most of the chemical reactions in organisms involve
solutes dissolved in water (aqueous solutions).
 To understand such reactions, we must know how
many atoms and molecules are involved (the
concentration of solutes in the water solution)
Step 1: Figure out the
molecular mass
 Molecular mass=the mass of each atom in a given
molecule.
 Example: Sucrose (C12H22O11) has a molecular mass
of 342 daltons. (Carbon atomic mass is 12 x 12 = 144;
Hydrogen atomic mass is 1 x 22 = 22; Oxygen atomic
mass is 16 x 11 = 176; 144 + 22 + 176 = 342)
 If we measure 342 daltons of sucrose, however, it
would be extremely difficult! So, we usually measure it
in moles
Step 2: Convert molecular mass
to moles
 Just as a dozen objects always means 12, a mole
represents an exact number of objects (6.02 X 1023)
which is called Avogadro’s number.
 There are 6.02 X 1023 daltons in 1 g.
 Once we determine the molecular mass of a molecule
such as sucrose, we can use the same number (342)
with the unit gram to represent the mass of 6.02 X 1023
molecules of sucrose. 342 grams is 1 mole of sucrose.
Step 3: Measure the correct
amount
 A mole of one substance has exactly the same number
of molecules as a mole of another substance.
 If you are asked to measure one mole of sucrose, what
would you do?
 If you are asked to measure one mole of ethyl alcohol,
C2H6O, what would you do?
 To make a liter (L) of solution containing 1 mole of
sucrose dissolved in water, what would you do?
Molarity
 To make a 1-molar (1 M) solution of sucrose, we would
measure out 342 g of sucrose and gradually add water,
while stirring, until the sugar was completely
dissolved. We would then add enough water to bring
the total volume of the solution to 1L.
 Molarity = the number of moles of solute per liter of
solution, is the concentration used most often by
biologists for aqueous solutions.
 You Try: How would you make a 1 M solution of
sodium chloride (NaCl)? How about a .5 M solution?
Shifting Hydrogen Atoms
 Occasionally, a hydrogen atom in water shifts from one
water molecule to another. When this happens, it most
often leaves behind its electron.
 What is transferred, then, is a hydrogen ion (H+).
 The water molecule that lost a proton is now a hydroxide
ion (OH-).
 The proton binds to another water, making a hydronium
ion H3O+.
Acids and Bases
 The reaction 2H2O
H3O+ + OH - is reversible. It
reaches a state of dynamic equilibrium when water
molecules dissociate at the same rate as they re-form.
 The concentration of each ion in pure water is 10-7 .
 H+ and OH- are very reactive. Changes in their
concentrations can drastically affect a cell’s proteins
and other molecules.
 Biologists use the pH scale to measure the levels of
H+ and OH- in a solution.
pH Scale
In any aqueous
solution, the
product of the H+
and OH- ions is
10-14
The
concentration of
H+ is 10-7 and the
concentration of
OH- is 10-7.
This is known as
pH 7.
Acids
 What would cause an
aqueous solution to have
an imbalance of H+ and
OH- concentrations?
 Acid=a substance that
increases the H+
concentration
 An acidic substance has
more H+ ions than pure
water. Any pH less than 7
is acidic.
Acids that can be eaten have a
sour taste. Acids can be corrosive.
Bases
 Base=A substance that
reduces the H+ ion
concentration is called a
base.
 Bases have less H+ ions
(and usually have more
OH- ions) in solution.
 Any solution with pH
more than 7 is basic.
Bases usually have a bitter taste.
In high concentrations, bases
can be as corrosive as strong
acids.
Fig. 3-UN5
0
Acidic
[H+] > [OH–]
Acids donate H+ in
aqueous solutions
Neutral
[H+] = [OH–]
Basic
[H+] < [OH–]
7
Bases donate OH–
or accept H+ in
aqueous solutions
14
Buffers
 The internal pH of most living
cells is close to 7.
 Even a slight change in pH can
be harmful.
 Buffers are substances that
minimize changes in the
concentrations of H+ and OHions.
 They do so by accepting H+ ions
when they are in excess or
donating them when lacking.
An important buffer in
blood is carbonic acid
H2CO3.
Water Quality
 Considering the
dependence of all life on
water, contamination of
rivers, lakes, seas, and
rainwater is a dire
environmental problem.
 Many threats to water
quality have been posed
by human activities.
Acid Precipitation
 Acid precipitation refers
to rain, snow, or fog with
a pH lower than 5.2.
 Uncontaminated rain is
about 5.6
 The burning of fossil
fuels is a major source of
sulfur oxides and nitrous
oxides. These combine
with water to form
strong acids.