Chapter 3
WATER AND THE FITNESS OF THE ENVIRONMENT
Summary of Chapter 3, BIOLOGY, 10TH ED Campbell, by J.B. Reece et al. 2014.
All organisms are made mostly of water and live in an environment dominated by water.
Water is the biological medium on Earth.
Life on Earth began in water and evolved there for 3 billion years before spreading onto land.
Modern life remains tied to water.
Living cells are 70% - 95 % water.
Water covers about ¾ of the Earth.
In nature, water naturally exists in all three physical states of matter: gas, liquid and solid.
I. POLAR COVALENT BONDS IN WATER RESULTS IN HYDROGEN BONDING.
Water unique behavior depends on its structure and the interaction of its molecules.
Polarity of water molecules results in hydrogen bonding.
The water molecule is shaped like a V.
Oxygen is more electronegative than hydrogen and pulls the electron pair closer to itself.
The hydrogens acquire a partial positive charge.
Water is a polar molecule.
Water molecules attract each other; the slightly positive hydrogen (δ+, delta) of one molecule is
attracted to the slightly negative oxygen (δ-, delta) of a nearby molecule.
The strength of the hydrogen bonds in liquid water is very weak, as
See animation: http://www.stolaf.edu/people/giannini/flashanimat/water/water.swf
II. FOUR EMERGENT PROPERTIES OF WATER
1. Organisms depend on the cohesion of water molecules.
Hydrogen bonds hold the water molecules together. This is called cohesion.
Hydrogen bonds are about 1/20 the strength of covalent bonds when water is its liquid state.
This bonds form, break and reform frequently.
Because at any given time most water molecules are bonded to its neighboring molecules,
water is a more structure liquid than most other liquids.
Water can also cling to other substances that are polar or partially polar. This is called
adhesion.
Cohesion and adhesion are responsible for the movement of water up the plant.
Hydrogen bonds make water to be...
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Cohesive: water molecules are attracted to each other.
Adhesive: water molecules are attracted to molecules on the surface of objects.
The attraction is due to the formation of hydrogen bonds.
Capillary action and surface tension is the result of these attractive forces.
Surface tension is a measure of the difficulty in breaking or stretching the surface of a liquid.
The energy of motion is kinetic energy.
Atoms and molecules are always moving; they have kinetic energy.
Heat is a measure of the total quantity of kinetic energy due to molecular motion in a body of
matter.
A form of energy representing aggregated internal energy of motions of atoms and molecules in
a body. Mass is an important consideration when talking about heat: the greater the mass, the
greater the heat.
2. Moderation of temperature.
a) Temperature and Heat
Water moderates air temperature by absorbing heat from air that is warmer and releasing the
stored heat to air that is cooler.
Heat is the energy that can be transferred between objects of different temperature. It is the
total amount of kinetic energy due to molecular motion in a substance that its bulk is not moving.
Temperature is a measurement of the average kinetic energy of a substance.
Science uses the Celsius scale to measure temperature.
b) Specific heat
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Water has a high specific heat.
1 calorie raises 1 g of water 1oC.
1 cal = 4.184 joules (1 J = 0.239 cal).
Specific heat is refers to the amount of energy required to raise the temperature of 1 gram of a
substance 1oC.
The specific heat of water is 1 cal/gram/1ºC
Another unit used for specific heat is the joule: J (joule).
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1 J = 0.239 cal; 1cal = 4.184 J
Compare to most other substances, water has an unusually high specific heat.
Because of its high specific heat, water will change its temperature less when absorbing or
losing heat.
Specific heat is a measure of how well a substance resists changing its temperature when it
absorbs or releases heat.
c) Evaporative cooling
When water changes its temperature, it absorbs or releases large amounts of heat.
Heat of vaporization is amount of energy that must be absorbed by 1 g of liquid water at 100oC
to change to 1 g of water vapor (gas) at 1oC.
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539.70 cal/1g or 2256J/1g
40700 J/mol (18.0 g/mol) or 2260 J/g
For water to evaporate, hydrogen bonds must be broken, which requires energy.
Evaporative cooling occurs because the molecules leaving the liquid are those with higher
kinetic energy. The molecules that remain have a lower average kinetic energy. They are
cooler.
The diagram on the left shows the uptake of heat by 1 kg of water, as it passes from ice at -50 ºC to steam at
temperatures above 100 ºC, affects the temperature of the sample.
A: Rise in temperature as ice absorbs heat.
B: Absorption of latent heat of fusion: 334 J/g
C: Rise in temperature as liquid water absorbs heat.
D: Water boils and absorbs latent heat of vaporization: 2256 J/g
E: Steam absorbs heat and thus increases its temperature.
The above is an example of a heating curve. One could reverse the process, and obtain a cooling curve. The flat
portions of such curves indicate the phase changes.
The diagram was taken from http://www.kentchemistry.com/links/Energy/HeatFusion.htm
Check: http://en.wikipedia.org/wiki/Cooling_curve
http://www.engineeringtoolbox.com/fluids-evaporation-latent-heat-d_147.html
3) Floating Ice on Liquid Water
Ice insulates bodies of water.
Water expands when it freezes.
Water contracts as it cools down to 4 oC. At this temperature water is densest.
As water continues to cool from 4 oC to 0 oC, it expands and becomes less dense than liquid
water.
Heat of fusion: 79.7 cal/g (or 333.5 J/g).
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1 thermochemical calorie = 4.184 J
When water begins to freeze, the molecules do not have enough kinetic energy to break
hydrogen bonds.
As the crystalline lattice forms, each water molecule forms a maximum of 4 hydrogen bonds,
which keeps water molecules further apart than they would be in the liquid state.
Because of the large space this arrangement takes, ice has fewer water molecules than in equal
volume of liquid water.
Expansion of water contributes to the fitness of the environment for life.
1. Prevents deep bodies of water from freezing solid from the bottom up.
2. Since ice is less dense, it forms on the surface firs. As water freezes, it releases heat to
the water below and insulates it.
3. Makes transitions between seasons less abrupt. As water freezes, hydrogen bonds form
releasing heat. As ice melts, hydrogen bonds break absorbing heat.
Check the structure of ice:
http://www.physicsofmatter.com/NotTheBook/Talks/Ice/Ice.html
http://www.indigo.com/models/gphmodel/molymod-ice-model-26.html
4. Water is a solvent
A solution is a homogeneous mixture of two substances.
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The dissolving agent is the solvent.
The substance that is dissolved is the solute.
An aqueous solution is one in which water is the solvent.
In a solution, the solute is broken down to the ionic or molecular level.
If an ionic crystal is dissolved, positive and negative ions go into solution.
The negative oxygen region of the polar water molecules is attracted to the positive ion.
The positive hydrogen region of the polar water molecules is attracted to the negative ion.
The surrounding water molecules shield the positive and negative ions from one another. The
water molecules form a hydration shell.
Water molecules can form hydrogen bonds with other particles.
a) Hydrophilic and Hydrophobic Substances
Hydrophilic substances interact readily with water.
A hydrophilic substance may or may not dissolve, e.g. cotton is made of cellulose, which is
hydrophilic, but cotton does not dissolve.
Water adheres to the cellulose molecules.
Hydrophobic substances repel water.
Hydrophobic substances are non-ionic and non-polar.
Some cell components are so large that they remain in suspension although they are
hydrophilic. These substances form colloids, a stable suspension of fine solid particles in a
liquid.
b) Solute Concentrations in Aqueous Solutions
MOLES
One mole of a substance contains 6.02 x 1023 atoms or molecules.
6.02 x 1023 is called Avogadro's number.
The mass of one mole of a substance is equivalent to the number of amu of the atom or
molecule of that substance.
Example: What is the mass of one mole of water?
H2O has 18 daltons (16 daltons of oxygen and 2 daltons of hydrogen).
1 mole of H2O contains 6.02 x 1023 molecules and has a mass of 18 grams.
One mole of a substance has the same number of molecules than one mole of another
substance. Both have 6.02 x 1023 molecules but their mass will be different.
Molarity is the number of moles of solute per liter of solution, 1 M = 1 mole/1 liter of solution.
Summary: Four emergent properties of water
1. Cohesion
2. Moderation of Temperature
3. Insulation of bodies of water by floating ice.
4. Solvent of life
III. ACIDIC AND BASIC CONDITIONS AFFECT LIVING ORGANISMS
Water can dissociate into H+ and OH-. This leads to acidic or basic conditions that affect living
organisms.
2H2O ↔ H3O+ + OHThe dissociation of water is reversible.
The hydrogen ion does not exist by itself, as a single proton, but it becomes attached to a water
molecule to form a hydronium ion, H3O+.
The anion OH- is called a hydroxide ion.
At 25ºC, the concentration of hydronium and hydroxide ions is 10-7 M (molar). That means that
only one water molecule in 554 million dissociates.
a) Acids and Bases
Acids are proton (H+) donors; bases are proton acceptors.
Acids tend to increase the number of protons (H+) in a solution.
Bases reduce the hydrogen ion concentration by accepting protons, H+.
NH3 + H+ ↔ NH4+
Many bases dissociate to produce hydroxide ions (HO-), which tend to accept H+.
NaOH → Na+ + OH-
The hydroxide ions then combine with hydrogen ions to form water and the hydrogen ion
concentration becomes reduced.
Strong acids and bases dissociate completely in water.
Weak acids and bases dissociate partially. These substances reversibly release and
reaccept hydrogen and hydroxide ions.
H2CO3 is a weak acid that releases H+ and HCO3-.
In any solution the product of H+ and OH- is constant at 10-14.
[H+] [OH-] = 10-14
In a neutral solution, [H+] = 10-7 and [OH-] = 10-7. Therefore [10-7] [10-7] = 10-14
If enough acid is then added to the solution to increase the H+ concentration up to 10-5 M, then
the concentration of OH- will decline to 10-9. The final product of both concentrations is still 10-14.
b) pH Scale
The pH scale measures the concentration of hydrogen ions (protons) in a solution.
7 is neutral. Below 7 is acid and above 7 is basic or alkaline.
It is based on the negative logarithm base 10 of the hydrogen ion concentration in a neutral
solution, -log 10-7 = -(-7) = 7.
Notice that as pH declines the concentration of H+ increases.
A pH of 10 has an acid concentration of 10-10 and a hydroxide concentration of 10-4. This is a
basic or alkaline solution.
The pH for acid solutions is below 7, and for basic solutions above 7. See fig. 3.10.
c) Buffers
Buffers are combinations of H+ donor and H+ acceptor that form weak acids or bases.
When the H+ are in excess, they accept them; and release H+ when they have been depleted.
Buffer solutions resist changes in pH when acids or bases are added. They minimize changes in
the concentrations of H+ and OH-.
Buffers help maintain the pH of body fluids within the narrow range necessary for life, usually pH
6-8.
Organisms have many weak acids and weak bases that maintain an essential buffering capacity
and avoid pH extremes.
Salts are compounds in which the hydrogen atom of an acid is replaced by some other cation.
Reactions of acids with bases form salts.
Electrolytes are salts, acids and bases that conduct electricity when in solution.
Nonelectrolytes do not conduct water when in solution, e.g. sugar, alcohols.
d) Acidification: A Thread to Water Quality
Acid precipitation is rain, snow or fog with a pH below 5.6.
It has been recorded to reach a pH of 1.5 in West Virginia.
It often results from a reaction in the air between water vapor and sulfur oxides and nitrogen
oxides produced by the combustion of fossil fuels.
Acid precipitation can change the solubility of minerals in the soil, e. g. Ca and Mg are washed
away.
More information about acid rain:
http://www.epa.gov/acidrain/
http://www.enviroliteracy.org/article.php/2.html
Ocean acidification is caused by the abundance of CO2 in the atmosphere. See fig. 3.11.
Summary of chapter 3:
http://www.mansfield.ohio-state.edu/~sabedon/campbl03.htm#hydrophobic_exclusion
1. Polarity of water molecules and hydrogen bonding
2. Four emergent properties of water:
Cohesion; adhesion
Moderation of ambient temperature
 temperature and heat
 specific heat
 heat of vaporization
 evaporative cooling
Insulation of bodies of water by floating ice
Water as a solvent:
 importance to life
 hydrophilic and hydrophobic substances
3. Acidic and Basic Conditions Affect Living Organisms:
 Hydrogen ion
 Hydroxide ion
 Acids and bases
 pH scale
 Buffers
 Acid rain: what is acid rain? Ocean acidification.