Chapter 3:
Water and the Fitness of
the Environment
AP Biology
Overview: The Molecule That Supports All of Life
•
Water is the biological medium on Earth
–
3/4 of Earth is submerged in water
–
It is the only common substance to exist in the natural environment in all 3
physical states (solid, liquid, gas)
•
•
Most is in liquid form but some is also present as ice and vapor
Life on Earth began in water and is the main reason Earth is habitable:
–
Evolution occurred in water for ~3 billion years before spreading to land
–
Even terrestrial life is tied to water
•
Humans can only survive ~week without water but for quite a few weeks
without food
–
Water is required in many chemical reactions necessary to sustain life
–
Most cells are surrounded by water and themselves are made up of ~70-95%
water
Concept 3.1:
The polarity of water molecules
results in hydrogen bonding
•
Certain properties of water allow it to function as a support to all living organisms
–
Water is a polar molecule
•
Oxygen is more electronegative than
hydrogen, leading to an unequal
distribution of electrons
– This means that the 2 ends of a water
molecule have opposite charges
•
•
Oxygen has partial negative charge
•
Hydrogen has partial positive charge
This polarity allows water molecules to form hydrogen bonds with each
other
Animation: Water Structure
Concept Check 3.1
•
1) What is electronegativity, and how does it affect interactions between
water molecules?
•
2) Why is it unlikely that 2 neighboring water molecules would be arranged
like this:
H H
O
O
H H
•
3) What would be the effect on the properties of the water molecule if
oxygen and hydrogen had equal electronegativity?
Concept 3.2:
Four emergent properties of water
contribute to Earth’s fitness for life
• 4 emergent properties of water contribute to Earth’s suitability
as an environment for life
– These emergent properties result from the hydrogen
bonding that orders molecules into a higher level of
structural organization
•
Cohesive behavior
•
Ability to moderate temperature
•
Expansion upon freezing
•
Versatility as a solvent
Cohesion
•
Though the arrangement of molecules in a sample of liquid water is constantly
changing (remember H-bonds last a few trillionths of a second), many of these
molecules are linked by multiple H-bonds at any given moment
–
These linkages make water more structured than most other liquids
•
Collectively, these H-bonds hold water molecules together, a property
called COHESION
– Cohesion due to H-bonds
contributes to transport of water
and dissolved nutrients against
gravity in plants
–
ADHESION is clinging of one substance
to another, different substance
•
Adhesion of water to cell walls by
H-bonds helps counter the downward
pull of gravity
Animation: Water Transport
•
SURFACE TENSION is 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
Surface tension is related to COHESION:
–
At interface between water and air is an ordered arrangement of water
molecules, each H-bonded to one another and to water below
–
This makes water behave like it is coated
with an invisible film
•
Ex) You can observe surface tension of
water by slightly overfilling a drinking
glass - some water will stand above rim
•
Ex) Some animals can stand, walk, or run on water without breaking the
surface
Moderation of Temperature
•
•
Water moderates air temperature by:
–
Absorbing heat from air that is warmer and
–
Releasing the stored heat to air that is cooler
Water is effective as a “heat bank” because it can absorb or release a relatively large
amount of heat with only a slight change in its own temperature
–
To understand this capability of water, we must look at concepts of heat and
temperature
•
Both heat and temperature are related to kinetic energy
Heat and Temperature
•
Anything that moves has KINETIC ENERGY – the energy of motion
–
Atoms and molecules have kinetic energy because they are always moving
•
•
HEAT is a form of energy
–
For a given body of matter, the amount of heat is a measure of the matter’s
TOTAL kinetic energy due to the motion of its molecules
•
•
The faster a molecule moves, the greater its kinetic energy
Heat therefore depends in part on matter’s volume
Although heat is related to TEMPERATURE, they are not the same thing
–
Temperature is a measure of heat intensity that represents the AVERAGE
kinetic energy of the molecules (regardless of volume)
•
Ex) When water is heated in coffeemaker, the average speed of its
molecules increases, and the thermometer records this as a rise in
temperature of the liquid
•
***NOTE: Although pot of coffee has much higher temperature than water
in a swimming pool, the swimming pool contains more heat because of its
much greater volume
•
Whenever 2 objects of different temperatures are brought together, heat
passes from warmer to cooler object until the 2 objects are the same
temperature
–
Molecules in the cooler object speed up at the expense of the kinetic
energy of the warmer object
• Ex) Ice cube cools drink not by adding coolness to liquid, but by
absorbing heat from liquid as the ice itself melts
–
In science, the CELSIUS SCALE is used to indicate temperature
• At sea level, water freezes at 0 and boils at 100 degrees Celsius
• Temperature of the human body averages 37 o C and comfortable
room temperature is 20-25 o C
•
HEAT is measured in units called CALORIES – the amount of heat it takes
to raise temperature of 1 gram of water by 1 degree Celsius
–
Conversely, it is also the amount of heat 1 gram of water releases
when it cools by 1 degree Celsius
–
A KILOCALORIE (1 kcal – 1000 cal) is quantity of heat required to
raise temperature of 1 kg of water by 1 degree Celsius
• Calories on food packages are actually kcal
–
Another energy unit is the JOULE (J)
• 1 J = 0.239 cal
• 1 cal = 4.184 J
Water’s High Specific Heat
•
Water can stabilize temperature due to its relatively high SPECIFIC HEAT – amount
of heat that must be absorbed or lost for 1 g of that substance to change its
temperature by 1 degree Celsius
•
–
Water: 1 cal/g/ºC
–
Ethyl alcohol: 0.6 cal/g/ºC
Because of high specific heat of water relative to other materials, water will change
its temperature less when it absorbs or loses a given amount of heat
–
Ex) The reason you can burn your fingers by touching side of metal pot on the
stove even though the water is still lukewarm is that specific heat of water is
10X greater than that of iron
•
This means that the same amount of heat will raise temperature of 1 g of
iron much faster than 1 g water
•
***Specific heat can therefore be thought of as a measure of how well a substance
resists changing its temperature when it absorbs or releases heat
• Water’s high specific heat is due to H-bonding
– Heat must be absorbed in order to break hydrogen bonds
– Heat is released when H-bonds form
• A calorie of heat causes only a small change in temperature of
water because much of the heat must first be used to disrupt Hbonds before water molecules themselves will begin to move
faster
•
Water’s high specific heat keeps temperature fluctuations on land and water within
limits that permit life
–
The water that covers
most of Earth can absorb
and store a huge amount
of heat from sun during
the day and during summer by only warming up a few degrees
•
At night and during winter, the gradually cooling water can then warm the
air (lake effect)
–
High specific heat also stabilizes ocean temperatures, creating a favorable
environment for marine life
–
Also, because organisms are made primarily of water, they are more able to
resist changes in their own temperature than if they were made of a liquid with
a lower specific heat – HOMEOSTASIS)
Evaporation
•
If molecules can move fast enough to overcome their attraction to one another, they
can depart the liquid and enter air as gas
–
This transformation from liquid to gas is called vaporization or EVAPORATION
•
Some evaporation occurs at any temperature, even low temperatures
–
•
Ex ) A glass of water at room temperature will eventually evaporate
If a liquid is heated, the average kinetic energy of molecules (temperature)
of molecules increases and evaporation occurs more rapidly
•
HEAT OF VAPORIZATION is quantity of heat a liquid must absorb got 1 g of it to be
converted from liquid to gaseous state
–
For the same reason that water has a high specific heat (H-bonds must first be
broken), it also has a high heat of vaporization relative to most other liquids
•
580 cal of heat needed at 25 o C
•
Double that is needed for alcohol or ammonia
• EFFECTS of Water Evaporation:
– GLOBALLY: helps moderate Earth’s climate
• Solar heat absorbed by tropical seas is consumed
during evaporation of surface water
• As this moist tropical air circulates poleward, it releases
heat as it condenses (forming rain)
– ORGANISMALLY: accounts for severity of steam burns
• These burns are caused by heat energy released when
steam condenses into liquid on skin
Evaporative Cooling
•
EVAPORATIVE COOLING – the cooling of the surface of a liquid that occurs as
liquid evaporates
–
Occurs because “hottest” (greatest kinetic energy) molecules are most likely to
leave as gas, thereby reducing the AVERAGE kinetic energy (temperature) of
the remaining liquid
•
–
–
Think: Removing the fastest 100 runners from a college causes the
average speed of remaining students to decline
Contributes to stability of temperature in lakes and ponds and also provides
mechanism that prevents terrestrial organisms from overheating
•
Ex) Evaporation of water from plant leaves helps keep leaf tissue from
becoming too warm in sunlight
•
Ex) Evaporation of sweat from human skin dissipates body heat and helps
prevent overheating on hot day or during strenuous activity
High humidity on hot days is uncomfortable because the high concentration of
water vapor in air inhibits evaporation of sweat from body
Water Density and Temperature
•
Water is one of the few substances that is less dense as solid than as a liquid (ice
floats in water)
–
While other materials contract when they solidify, water expands (up to 4 o C)
•
–
Water behaves the same way as other liquids above this temperature
Water starts to freeze when its molecules are no longer moving fast enough to
break their hydrogen bonds
•
At 0 o C, water becomes locked into a crystalline lattice, with each water
molecule H-bonded to 4 partners
–
•
–
These H-bonds keep water molecules at an “arm’s length” – far
enough apart to make ice about 10% less dense (10% fewer
molecules for same amount of volume) than liquid water at 4 o C
When ice absorbs enough heat for its temperature to rise above 0 o C, Hbonds become disrupted, allowing crystal structure to collapse and ice to
melt (molecules are then free to slip closer together)
Water’s greatest density is at 4 o C and then it begins to expand as molecules
move faster (again reducing density)
Insulation of Bodies of Water by Floating Ice
•
ENVIRONMENTAL EFFECTS of water density’s temperature-dependence :
–
If ice sank, then eventually all bodies of water would freeze solid, making life as
we know it on Earth impossible
•
–
During summer, only a few inches of the ocean would thaw
Instead, because ice floats on water, it insulates the liquid water below,
preventing it from freezing and allowing life to exist under its frozen surface
•
The Solvent of Life
• A solution is a liquid that is a homogeneous mixture
of substances
– A solvent is the dissolving agent of a solution
– The solute is the substance that is dissolved
– An aqueous solution is one in which water is the
solvent
•
Though water is not a universal solvent (it would then dissolve any container in which
it was stored, including our cells), it is a very VERSATILE solvent
–
–
This is due to the POLARITY of water molecules
•
Hydrogen regions are positively charged and attracted to anions
•
Oxygen regions are negatively charged and attracted to cations
As a result, water molecules will surround individual ions, separating them and
shielding them from one another
•
This sphere of water molecules around
each dissolved ion is called a
HYDRATION SHELL
–
Working inward from the surface of a
salt crystal, for example, water
eventually dissolves all the ions,
resulting in an aqueous solution
•
A compound does not need to be ionic to dissolve in water
–
Many compounds made up of nonpolar ionic molecules (sugars) are also
water-soluble
•
Such compounds dissolve when water molecules surround each of the
solute molecules, forming H-bonds with them
–
Even molecules as large as proteins can dissolve in water if they have ionic
and polar regions on their surface
•
Within organisms, many different kinds of polar compounds are dissolved in water –
blood, liquid within all cells,
plant sap
–
***Water is the
solvent of life
Hydrophilic Substances
•
Any substance that has an affinity for water is said to be HYDROPHILIC
•
•
Hydro – water; philios – loving
In some cases, substances can be hydrophilic without actually dissolving
–
Some molecules in cells are so large that they can’t dissolve; instead they
remain suspended in aqueous liquid of cells
–
•
This type of mixture is an example of a COLLOID – a stable
suspension of fine particles in a liquid
Ex) Cotton, a plant product made of giant molecules called cellulose (has
numerous regions of partial positive and negative charges that can form Hbonds with water)
–
Water adheres to cellulose fibers, which is why cotton towels work
well for drying but do not dissolve in washing machines
–
Cellulose is also present in walls of water-conducting cells in plant,
allowing the adhesion that facilitates water transport in plants
Hydrophobic Substances
•
Substances that repel water are said to be HYDROPHOBIC (phobos –
fearing)
–
Includes nonionic and nonpolar compounds
• Ex) Hydrophobic behavior of oil molecules results from prevalence
of relatively nonpolar bonds (like C-H bonds)
– Hydrophobic molecules related to oil are a major ingredient of
cell membranes
Solute Concentration in Aqueous Solutions
• Most of the chemical reactions in organisms involve solutes
dissolved in water
–
Chemical reactions depend on collisions of molecules and
therefore on the concentration of solutes in an aqueous
solution
• Therefore, to understand these reactions, we have to:
– Know how many atoms and molecules are involved
– Be able to calculate the concentration of solutes in an
aqueous solution (the number of solute molecules in a
volume of solution)
•
When doing experiments, we use atomic mass and numbers of molecules to
calculate molecular mass (sum of the masses of all atoms in a molecule)
• Ex) sucrose (C12H22O11) = (12 x 12) + (22 x 1) +( 11 x 16) = 342
daltons
–
Weighing (in grams) out small numbers of molecules is not practical
• We usually measure substances in units called MOLES (mol)
– 1 mol = 6.02 X 1023 molecules (Avogadro’s Number)
– 6.02 X 1023 daltons = 1 gram
• To obtain 1 mol of sucrose in lab, we therefore weigh out 342
grams
•
Measuring molecules in moles makes it convenient for scientists working in the lab to
combine substances in fixed ratios of molecules:
•
Ex) If molecular mass of substance A is 342 daltons and that of substance
B is 10 daltons, then 342 g of A will have same number of molecules as 10
gB
•
How would we make 1 L of solution made of 1 mol of sucrose dissolved in water?
•
Measure out 342 grams of sucrose and enough water to bring total volume
of solution up to 1 L, creating a 1-molar (M) solution of sucrose
–
MOLARITY – number of moles of solute per liter of solution
•
This is the unit of concentration used most often by biologists for aqueous
solutions
Concept Check 3.2
•
1) Describe how properties of water contribute to the upward movement of
water in a tree.
•
2) Explain the saying “It’s not the heat; it’s the humidity.”
•
3) How can the freezing of water crack boulders?
•
4) If you were a pharmacist, how would you make a 0.5-molar (0.5 M)
solution of sodium chloride (NaCl)? (The atomic mass of Na is 23 daltons
and that of Cl is 35.5 daltons)
•
5) A water strider’s legs (see Figure 3.4) are coated with a hydrophobic
substance. What might be the benefit? What would happen if the substance
were hydrophilic?
Concept 3.3:
Acidic and basic conditions affect
living organisms
• Occasionally, a hydrogen atom participating in an H-bond between 2
water molecules shifts from one water molecule to the other
–
When this happens, the hydrogen atom leaves an electron
behind, and what is actually transferred is a hydrogen ion (H+)
– a single proton with a +1 charge
• The water molecule that lost the proton is now a hydroxide
ion (OH-) with a charge of –1
–
The proton binds to other water molecule, making a hydronium
ion (H3O+)
• Hydronium ions ,however, are often represented simply as
H+
•
The dissociation of water is a reversible reaction that reaches a state of dynamic
equilibrium
–
At equilibrium, water molecules dissociate at the same rate that they are being
reformed from H+ and OH•
At this equilibrium point, however, the concentration of water molecules
greatly exceeds the concentrations of H+ and OH-
–
In pure water, only 1 water molecule in every 554 million is dissociated)
•
The concentration of each ion in pure water is 10-7 M
–
This means that there is only 1 ten-millionth of a mole of H+ and OHions per liter of pure water
•
Though the dissociation of water is statistically rare, it is very important in
the chemistry of life:
–
H+ and OH- ions are very reactive
• Changes in their concentrations can drastically effect a cell’s
proteins and other
–
In pure water, the concentration of H+ and OH- ions are equal
• Adding certain kinds of solutes, called acids and bases, modifies
the concentrations of these ions
– Biologists use the pH scale to describe how acidic or basic a
solution is
Acids
• When acids dissolve in water, they donate additional H+ ions to
the solution
– ACID – substance that increases hydrogen ion
concentration of a solution
• This results in an acidic solution, having one more H+
than OH- ion
– Ex) Hydrochloric acid (HCl) dissociates to form
chloride ions in water
Bases
•
A substance that reduces the hydrogen ion concentration of a solution is a
BASE
–
Some bases reduce H+ concentration directly by accepting H+ ions
• Ex) NH3 (ammonia) acts as a base when unshared electron pair in
nitrogen’s valence shell attracts a hydrogen ion from the solution,
resulting in an ammonium ion (NH4+)
–
Other bases reduce H+ concentration indirectly by dissociating to form
hydroxide ions, which then combine with hydrogen ions and form water
• Ex) NaOH dissociates into its ions Na+ and OH-
–
In either case, the base reduces that H+ concentration, resulting in a
basic solution
•
A solution that has equal concentrations of H+ and OH- is said to be neutral
•
Acids or bases that dissociate completely when mixed with water (like HCl or
NaOH) are considered to be strong acids or bases
• Thus, the chemical reaction showing their dissociation is
represented with a single arrow
–
In contrast, weak acids or bases (like NH3) have fixed ratios of products
and reactants at equilibrium (NH4+ and NH3)
• These chemical reactions are thus represented as a reversible
reaction with double arrows
•
In any aqueous solution at 25o C, the product of the H+ and OH- concentrations is
constant at 10-14
–
This can be written as [H+][OH-] = 10-14
•
In this equation, brackets indicate molar concentration
–
In a neutral solution:[10-7][10-7] = 10-14
–
In an acidic solution, [H+] increases while [OH-] decreases an
equivalent amount, since multiplying these concentrations will always
equal 10-14
•
•
Ex) [H+][OH-] = [10-5][10-9]=10-14
***Whenever we know the concentration of OH- or H+ in an aqueous solution, we
can deduce the concentration of the other because of this constant relationship***
–
This constant relationship expresses the behavior of acids and bases in an
aqueous solution:
•
Acids not only add hydrogen ions to solution, but also removes hydroxide
ions because of the tendency of H+ to combine with OH- to form water
•
A base has the opposite effect: increases OH- concentration but also
reduces H+ concentration by formation of water
The pH Scale
•
Because H+ and OH- concentration of solutions can vary by a factor of 100 trillion or
more, scientists have developed a way to express this variation more conveniently
than mol/L
–
The pH scale compresses the range of H+ and OH- concentrations using
logarithms
•
pH is defined as the negative logarithm (base 10) of hydrogen ion
concentration
–
pH = -log[H+]
–
For neutral solutions: -log 10-7 = -(-7) = 7
–
Notice that pH decreases as hydrogen ion concentration increases:
•
Ex) –log 10 -5 = -(-5) = 5
•
The pH of neutral solutions at 25 o C is 7 (midpoint of pH scale)
–
pH values less than 7 denote an acidic solution
•
–
•
pH for basic solutions is above 7
Most biological fluids are within the range of 6-8
–
•
The lower the number, the more acidic
the solution
Some exceptions include strongly acidic
digestive juice of human stomach (pH = 2)
Remember that each pH unit represents a 10-fold
difference in H+ and OH- concentration
•
–
Ex) pH = 3 is not twice as acidic as
pH = 6 but 1000X (10 x 10 x 10) more
acidic
This means when the pH of a solution
changes slightly, the actual concentrations
of H+ and OH- in solution actually change
substantially
•
The internal pH of most living cells is close to 7
–
Even slight changes in pH can be harmful because chemical processes of the
cell are very sensitive to concentrations of hydrogen and hydroxide ions
•
The pH of human blood is close to 7.4 (slightly basic)
–
A person cannot survive for more than a few minutes if blood pH drop
to 7 or rises to 7.8
–
A chemical system therefore exists in blood that maintains a stable pH
(homeostasis)
–
If you add 0.01 mol of a strong acid to a liter of pure water, the pH drops from 7
to 2
•
If same amount is added to liter of blood, pH only decreases from 7.4 to
7.3 (Why?)
Buffers
•
The presence of substances called buffers allows for relatively constant pH
in biological fluids despite addition of acids or bases
–
BUFFERS: substances that minimize changes in H+ and OHconcentrations by accepting hydrogen ions from solution when they are
in excess and donating hydrogen atoms when they are depleted
• Most buffer solutions contain a weak acid and its corresponding
base which combine reversibly with hydrogen ions
• There are several buffers that contribute to pH stability in human
blood and other biological solutions (like carbonic acid – H2CO3)
Threats to Water Quality on Earth
•
Considering the dependence of all life on water, contamination of rivers, lakes, seas,
and rain by human activities is a huge environmental problem
–
The burning of fossil fuels is a major source of sulfur oxides and nitrous oxides,
which react with water in the air to form strong acids that fall back to Earth with
rain or snow
•
ACID PRECIPITATION – rain, snow, or fog with a pH lower (more acidic)
than 5.2
–
–
Uncontaminated rain has pH of ~5.6; it is slightly acidic owing to the
formation of carbonic acid from carbon dioxide and water
Electric power plants that burn coal
produce more of these oxides than
any single source
•
Winds carry the pollutants away
and acid rain may fall 100s of km
away from industrial centers
•
Acid precipitation can damage
life in lakes and streams, and
has an adverse effect on soil chemistry
•
CO2 (the main product of fossil fuel combustion) causes other problems as well:
–
Its release into the atmosphere has been increasing steadily and is expected to
double by 2065 relative to 1880 levels
–
About ½ of CO2 stays in atmosphere, acting like a reflective blanket over the
planet that prevents heat from radiating to outer space – GREENHOUSE
EFFECT
•
–
–
A portion of this CO2 is taken up by trees and other photosynthetic
organisms but about 30% is absorbed b oceans
When CO2 dissolves in seawater, it reacts with water to form carbonic acid
(H2CO3)
•
Almost all the carbonic acid in turn dissociates, producing protons and a
balance between 2 ions, bicarbonate (HCO3- and carbonate (CO3 2-)
•
As sea water acidifies due to these extra protons, the balance shifts toward
bicarbonate, lowering the concentration of carbonate
Many studies have shown that calcification, the production of calcium
carbonate by corals and other organisms is directly affected by carbonate
concentration
•
This can affect coral reef formations, which house a great diversity of
organisms
•
What is the effect of carbonate ion concentration on coral reef calcification?
–
–
One of the best known and longest studies on coral reef calcification was
carried out at ecosystem center in Arizona known as Biosphere-2
•
This center includes an artificial coral reef system in which temperature
and chemistry of sea water can be controlled and manipulated
•
This system was used to test the effects of varying concentration of
carbonate on rate of calcification in coral reef
It led to the prediction that expected doubling
of CO2 emissions by 2065 could lead to a
40% decrease in coral reef calcification
•
EXPERIMENT: For almost 4 years, researchers
varied carbonate concentration in seawater under
controlled conditions and measured rate of
calcification by reef organisms
•
RESULTS: calcification rate was observed to be
lower at lower concentrations of carbonate ions
•
CONCLUSION: 40% decrease in rate of coral
reef calcification by 2065 relative to pre-industrial
levels, endangering coral reefs
Concept Check 3.3
•
1) Compared with a basic solution at pH 9, the same volume of an acidic
solution at pH 4 has _____ times as many hydrogen atoms (H+).
•
2) HCl is a strong acid that dissociates in water: HCl
H+ + Cl-
What is the pH of 0.01 M HCl?
•
3) Acetic acid (CH3COOH) can be a buffer, similar to carbonic acid. Write
the dissociation reaction, identifying the acid, base, H+ acceptor, and H+
donor.
•
4) Given a liter of pure water and a liter solution of acetic acid, what would
happen to the pH if you added 0.01 mol of a strong acid to each? Use the
reaction equation to explain the results.
You should now be able to:
1. List and explain the four properties of water
that emerge as a result of its ability to form
hydrogen bonds
2. Distinguish between the following sets of
terms: hydrophobic and hydrophilic
substances; a solute, a solvent, and a
solution
3. Define acid, base, and pH
4. Explain how buffers work
Copyright © 2008 Pearson Education, Inc., publishing as Pearson Benjamin Cummings