Chapter 3

Water and Life

Lecture Presentations by
Nicole Tunbridge and
Kathleen Fitzpatrick
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The Molecule That Supports All of Life

▪ Water makes life possible on Earth

▪ Water is the only common substance to exist in the
natural environment in all three physical states of
matter
▪ Water’s unique emergent properties help make Earth
suitable for life
▪ The structure of the water molecule allows it to
interact with other molecules

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Figure 3.1

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Figure 3.1a

Black guillemots, threatened

by climate change

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Concept 3.1: Polar covalent bonds in water
molecules result in hydrogen bonding
▪ In the water molecule, the electrons of the polar
covalent bonds spend more time near the oxygen
than the hydrogen
▪ The water molecule is thus a polar molecule:
The overall charge is unevenly distributed
▪ Polarity allows water molecules to form hydrogen
bonds with each other

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Figure 3.2

δ+
δ+

Polar covalent bond

δ–
δ–

Region of partial Hydrogen bond

δ+
negative charge

δ–
δ+
δ– δ+
δ–
δ+

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Concept 3.2: Four emergent properties of water
contribute to Earth’s suitability for life
▪ Four of water’s properties that facilitate an
environment for life are
▪ Cohesive behavior
▪ Ability to moderate temperature
▪ Expansion upon freezing
▪ Versatility as a solvent

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Cohesion of Water Molecules

▪ Collectively, hydrogen bonds hold water molecules

together, a phenomenon called cohesion
▪ Cohesion helps the transport of water against gravity
in plants
▪ Adhesion is an attraction between different
substances, for example, between water and plant
cell walls

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Figure 3.3

Evaporation pulls water upward.

H 2O

Adhesion

Two types of
water-conducting
cells

Direction Cohesion
of water 300 µm
movement

H 2O
H 2O
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Figure 3.3a

Two types of
water-conducting
cells

300 µm
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Animation: Water Transport in Plants

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▪ Surface tension is a measure of how difficult it is to
break the surface of a liquid
▪ Water has an unusually high surface tension due to
hydrogen bonding between the molecules at the
air-water interface and to the water below

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Figure 3.4

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Moderation of Temperature by Water

▪ Water absorbs heat from warmer air and releases

stored heat to cooler air
▪ Water can absorb or release a large amount of heat
with only a slight change in its own temperature

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Temperature and Heat

▪ Kinetic energy is the energy of motion

▪ The kinetic energy associated with random motion of
atoms or molecules is called thermal energy
▪ Temperature represents the average kinetic energy
of the molecules in a body of matter
▪ Thermal energy in transfer from one body of
matter to another is defined as heat

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▪ A calorie (cal) is the amount of heat required to
raise the temperature of 1 g of water by 1ºC
▪ It is also the amount of heat released when 1 g of
water cools by 1ºC
▪ The “Calories” on food packages are actually
kilocalories (kcal); 1 kcal = 1,000 cal
▪ The joule (J) is another unit of energy;
1 J = 0.239 cal, or 1 cal = 4.184 J

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Water’s High Specific Heat

▪ The specific heat of a substance is the amount of

heat that must be absorbed or lost for 1 g of that
substance to change its temperature by 1ºC
▪ The specific heat of water is 1 cal/(g ∙ ºC)
▪ Water resists changing its temperature because of
its high specific heat

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▪ Water’s high specific heat can be traced to hydrogen
bonding
▪ Heat is absorbed when hydrogen bonds break
▪ Heat is released when hydrogen bonds form
▪ The high specific heat of water minimizes
temperature fluctuations to within limits that
permit life

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Figure 3.5

Burbank San Bernardino

Santa Barbara 73º
90º 100º
Los Angeles Riverside 96º
(Airport) 75º Santa Ana
Palm Springs
70s (ºF) 84º
106º
80s Pacific Ocean 68º
90s
100s San Diego 72º 40 miles

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Evaporative Cooling

▪ Evaporation (or vaporization) is transformation of a

substance from liquid to gas
▪ Heat of vaporization is the heat a liquid must
absorb for 1 g to be converted to gas
▪ As a liquid evaporates, its remaining surface cools, a
process called evaporative cooling
▪ Evaporative cooling of water helps stabilize
temperatures in organisms and bodies of water

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Floating of Ice on Liquid Water

▪ Ice floats in liquid water because hydrogen bonds in

ice are more “ordered,” making ice less dense than
water
▪ Water reaches its greatest density at 4ºC
▪ If ice sank, all bodies of water would eventually
freeze solid, making life impossible on Earth

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Figure 3.6

Hydrogen bond Liquid water:

Hydrogen bonds
break and re-form

Ice:
Hydrogen bonds
are stable

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Figure 3.6a

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▪ Many scientists are worried that global warming is
having a profound effect on icy environments around
the globe
▪ The rate at which glaciers and Arctic sea ice are
disappearing poses an extreme challenge to animals
that depend on ice for their survival

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Figure 3.7

Benefiting from loss of ice:

Phyto- Bowhead
plankton whales Capelin

Harmed by
loss of ice:
Russia
Arctic
ocean
Extent of sea ice in Sept. 2014
Extent of sea ice in Sept. 1979
Polar bears Bering
Strait
North Pole

Greenland

Pacific walrus
Alaska

Black Canada
guillemots
Sea ice in Sept. 2014
Ice lost from Sept. 1979 to Sept. 2014
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Water: The Solvent of Life

▪ A solution is a liquid that is a completely

homogeneous mixture of substances
▪ The 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

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Figure 3.8

Na+
–
+
+ –
– +
– –
Na+
–
+ +
CI– CI–
+
–
– +
–
+
–
–

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▪ Water is a versatile solvent due to its polarity
▪ When an ionic compound is dissolved in water, each
ion is surrounded by a sphere of water molecules
called a hydration shell

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▪ Water can also dissolve compounds made of
nonionic polar molecules
▪ Even large polar molecules such as proteins
can dissolve in water if they have ionic and
polar regions

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Figure 3.9

δ+

δ– δ–

δ+

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Hydrophilic and Hydrophobic Substances

▪ A hydrophilic substance is one that has an

affinity for water
▪ A hydrophobic substance is one that does not have
an affinity for water
▪ Oil molecules are hydrophobic because they
have relatively nonpolar bonds
▪ Hydrophobic molecules related to oils are the major
ingredients of cell membranes

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Solute Concentration in Aqueous Solutions

▪ Most chemical reactions in organisms involve

solutes dissolved in water
▪ When carrying out experiments, we use mass to
calculate the number of solute molecules in an
aqueous solution

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▪ Molecular mass is the sum of all masses of all
atoms in a molecule
▪ Numbers of molecules are usually measured in
moles, where 1 mole (mol) = 6.02 × 1023 molecules
▪ Avogadro’s number and the unit dalton were defined
such that 6.02 × 1023 daltons = 1 g
▪ Molarity (M) is the number of moles of solute
per liter of solution

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Possible Evolution of Life on Other Planets

▪ Biologists seeking life on other planets have

concentrated their search on planets that might have
water
▪ More than 800 planets have been found outside our
solar system; there is evidence that a few of them
have water vapor
▪ In our solar system, Mars has been found to
have water

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Figure 3.10

Dark streaks

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Concept 3.3: Acidic and basic conditions affect
living organisms
▪ A hydrogen atom in a hydrogen bond between two
water molecules can shift from one to the other
▪ The hydrogen atom leaves its electron behind and is
transferred as a proton, or hydrogen ion (H+)
▪ The molecule that lost the proton is now a hydroxide
ion (OH–)
▪ The molecule with the extra proton is now a
hydronium ion (H3O+), though it is often represented
as H+

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▪ Water is in a state of dynamic equilibrium in which
water molecules dissociate at the same rate at which
they are being reformed

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Figure 3.UN01

+ –
H H
O H O O H O
H H H H
2 H 2O Hydronium Hydroxide
ion (H3O+) ion (OH–)

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▪ Though statistically rare, the dissociation of water
molecules has a great effect on organisms
▪ Changes in concentrations of H+ and OH– can
drastically affect the chemistry of a cell

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▪ Concentrations of H+ and OH– are equal in
pure water
▪ Adding certain solutes, called acids and bases,
modifies the concentrations of H+ and OH–
▪ Biologists use the pH scale to describe whether a
solution is acidic or basic (the opposite of acidic)

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Acids and Bases

▪ An acid is a substance that increases the H+

concentration of a solution
▪ A base is a substance that reduces the H+
concentration of a solution
▪ Strong acids and bases dissociate completely
in water
▪ Weak acids and bases reversibly release and accept
back hydrogen ions, but can still shift the balance of
H+ and OH– away from neutrality

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The pH Scale

▪ In any aqueous solution at 25ºC, the product of H+

and OH– is constant and can be written as
[H+][OH–] = 10–14
▪ The pH of a solution is defined by the negative
logarithm of H+ concentration, written as
pH = –log [H+]
▪ For a neutral aqueous solution, [H+] is 10–7, so
pH = –(–7) = 7

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▪ Acidic solutions have pH values less than 7
▪ Basic solutions have pH values greater than 7
▪ Most biological fluids have pH values in the range of
6 to 8

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Figure 3.11
pH Scale
0

1
Battery acid

Increasingly Acidic
2 Gastric juice (in stomach),
lemon juice

[H+] > [OH–]

3 Vinegar, wine,
cola
Acidic 4 Tomato juice
solution Beer
5 Black coffee
Rainwater
6 Urine
Saliva
Neutral 7 Pure water
[H+] = [OH–]
Human blood, tears
8 Seawater
Neutral
Inside small intestine
solution
Increasingly Basic

9
[H+] < [OH–]

10
Milk of magnesia
1
1 Household ammonia
12
Basic
Household
solution 13 bleach
Oven cleaner
14
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Buffers

▪ The internal pH of most living cells is close to 7

▪ Buffers are substances that minimize changes in
concentrations of H+ and OH– in a solution
▪ Most buffer solutions contain a weak acid and its
corresponding base, which combine reversibly with
H+ ions

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Acidification: A Threat to Our Oceans

▪ Human activities such as burning fossil fuels

threaten water quality
▪ CO2 is the main product of fossil fuel combustion
▪ About 25% of human-generated CO2 is absorbed by
the oceans
▪ CO2 dissolved in seawater forms carbonic acid; this
process is called ocean acidification

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Figure 3.12

CO2

CO2 + H2O H2CO3

H2CO3 H+ + HCO3–

H+ + CO32– HCO3–

CO32– + Ca2+ CaCO3

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▪ As seawater acidifies, H+ ions combine with
carbonate ions to produce bicarbonate
▪ Carbonate is required for calcification (production of
calcium carbonate) by many marine organisms,
including reef-building corals
▪ We have made progress in learning about the
delicate chemical balances in oceans, lakes,
and rivers

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Figure 3.UN02a

[mmol CaCO3/(m2 ∙ day)]

Calcification rate

20

10

0
220 240 260 280
[CO32–] (µmol/kg of seawater)
Data from C. Langdon et al., Effect of calcium carbonate saturation state on the calcification
rate of an experimental coral reef, Global Biogeochemical Cycles 14:639–654 (2000).

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Figure 3.UN02b

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Figure 3.UN03

δ–
δ+

δ+ δ–
δ– δ+
δ+ δ–

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Figure 3.UN04

0
Acidic
[H+] > [OH–]
Acids donate H+ in
aqueous solutions.

Neutral
[H+] = [OH–] 7

Bases donate OH–

or accept H+ in
Basic aqueous solutions.
[H+] < [OH–]

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14
Figure 3.UN05

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