9/27/2014
Why study water?
Water and Life
• All life occurs in water – inside and
outside the cell.
• Water is the molecule that supports all
life.
Chapter 3
Pgs. 46 - 57
Water and Polar Covalent Bonds
• Water molecules are shaped like a wide V – one O
and two H molecules bonded covalently
• The electrons of the covalent bonds spend more
time closer to O than H, making them polar
covalent bonds
– O has a partial negative charge while each H has a
partial positive charge
– This makes the water molecule polar
• Because the entire water molecule is polar, it is
attracted to other polar molecules.
– This leads to hydrogen bonding, which is responsible
for the various properties that water has
Water and Hydrogen Bonding
• Water molecules are polar, and are thus
attracted to each other
– The slightly positive H of one molecule is
attracted to the slightly negative O of a nearby
molecule
• In liquid form, H-bonds are very fragile; they
form, break, and re-form frequently
– Each H-bond only lasts a fraction of a second,
but the molecules are constantly forming new
H-bonds with a succession of partners
Emergent Properties of Water
• Cohesion and adhesion
– Surface tension, capillary action
• Moderation of temperature
– High specific heat – water stores heat
– High heat of vaporization – water heats and cools slowly
– Evaporative cooling
• Floating of ice on liquid water
– Lower density as a solid
• Solvent of life
– Many molecules dissolve in water
– Hydrophilic vs. hydrophobic
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Cohesion and Adhesion
• Cohesion is the bonding of water molecules sticking to neighboring
water molecules through H-bonds
• Because of H-bonds, water molecules always stay close to each other;
multiple H-bonds linking water molecules make the substance more
structured
– Cohesion is directly related to surface tension – a measure of how difficult
it is to stretch or break the surface of a liquid
• How is it helpful?
Moderation of Temperature
• Heat and temperature
• High specific heat
• Evaporative cooling
– Contributes to the transport of water and dissolved
nutrients against gravity in plants (capillary action)
• Adhesion also helps – the
clinging of water molecules to
other substances (water clings to
cell walls, fights gravity); also
causes meniscus in glassware
– How does H2O get to the tops of
trees? Transpiration is built on
cohesion and adhesion.
Heat and Temperature
• The amount of heat is a measure of matter’s total
kinetic energy due to motion of its molecules,
influenced by volume.
• Temperature is a measure of heat intensity that
represents the average kinetic energy of molecules,
regardless of volume.
• Heat passes from warm  cool until the same
temperature has been reached; cold does not travel!
• We always use the Celsius scale
– Water freezes at 0° and boils at 100°
– Body temperature is about 37°
– Room temperature is usually 20-25°
More Heat + Temp. Vocabulary
• Calorie (cal): amount of heat it takes to
raise the temperature of 1 g of water by
1°C
• Kilocalorie (kcal, Cal): 1000 calories;
amount of heat it takes to raise the
temperature of 1 kg of water by 1°C
• Joule (J): one joule = 0.239 cal; one cal =
4.184 J
High Specific Heat
How is High Specific Heat Good?
• Specific heat: amount of heat that must be
absorbed/lost for 1 g of that substance to change
temperature by 1°C
• A large body of water can absorb and store heat
from the sun in the daytime and summer while
only warming a few degrees.
• At night and during the winter, the cooling water
can warm the air and stabilize the local climate.
• Ocean temperatures are stabilized this way as
well, which creates a favorable environment for
marine life.
• Because organisms are
primarily made of water, they
are better able to resist changes
in their own temperature.
– Specific heat of water is 1 calorie per gram and per
degree Celsius – 1 cal/g•°C
• Water can absorb or lose a substantial amount of
heat before its temperature changes
• This property is due to H-bonding – heat breaks
the H-bonds before the molecules can move
faster and H-bonds must re-form before the
molecules start moving more slowly
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Evaporative Cooling
• Heat of vaporization is the quantity of heat a liquid
must absorb for 1 g of it to be converted from liquid 
gas (vaporization, or evaporation)
– H-bonds must be broken before liquid water can be
converted into water vapor gas
– Helps moderate Earth’s climate
• As liquid evaporates, the surface cools down
(evaporative cooling) – the “hottest” molecules with
the greatest kinetic energy are most likely to leave as
gas
– Stabilizes temperature in lakes and ponds
and prevents terrestrial organisms
from overheating (sweating)
– These H-bonds keep the molecules farther apart at
0°C than at 4°C, so water expands as it freezes – it
becomes 10% less dense through this change
• What would happen if ice sank?
Solute Concentrations
• Molecular mass = sum of the mass of all atoms in a molecule (daltons)
• Mole (mol) = exact number of molecules: 6.02 x 1023, called Avogadro’s
number
• Water is not a universal solvent, but it is a very
versatile one – it can dissolve many substances,
including ionic and nonionic polar molecules
– Polarity makes water a good solvent – polar H2O molecules
surround + and - ions
– Solvents dissolve solutes, creating solutions
– 6.02 x 1023 daltons = 1 g
– Once you determine the molecular mass of a molecule, use the same number with the unit
gram to represent 1 mol
• Solution: a liquid that is a completely
homogeneous mixture of two or more substances
• Molarity: number of moles of solute per liter of solution
– Importance: in biology, we measure and create solutions with specific molarity (the number
of moles of solute per liter of solution), such as a 1M solution of HCl
– Solvent: dissolving agent / Solute: substance being
dissolved
– Aqueous solution: water is the solvent
• M 1 V1 = M 2 V2
– Ex. How much concentrated 18M sulfuric acid is needed to prepare 100mL of a 2M
solution?
(18M) (V1)= (2M)(100mL)
V1= (2M)(100mL) / (18M) = 11 mL
– Ex. To how much water should 100 mL of 18 M HCl be added to produce 1.5M solution?
(18M) (100mL)= (1.5M)(V2)
V2= (18M)(100mL) / (1.5M) = 1200 mL
1200mL-100mL= 1100mL
• Hydrophilic vs. hydrophobic
– Hydrophilic: any substance that has an affinity for water;
polar
• Substances do not have to dissolve to be hydrophilic, such as a
colloid (stable suspension of fine particles in a liquid)
– Hydrophobic: substances that are nonionic and nonpolar
that seem to repel water
Chemical
Explanation
• Water is one of the few
substances that are less
dense as a solid than as a liquid
• At 0°C, water molecules
become locked into a
crystal-like lattice where each
molecule is H-bonded to four partners
– All bodies of water would freeze solid, making life
impossible – floating ice insulates the water below,
allowing life to exist under a frozen surface
Solvent of Life
H2O Property
Floating of Ice on Liquid Water
Examples of
Benefits to Life
Cohesion
•polar
•H-bond
•like-like
↑gravity plants, trees
Adhesion
•H-bond
•unlike-unlike
plants xylem
bloodveins
Surface Tension
•diff. in stretch
•break surface
•H-bond
bugswater
Specific Heat
•Absorbs & retains E
•H-bond
oceanmod temp
protect marine life
Evaporation
•liquidgas
•KE
Cooling
Homeostasis
Solvent of life
•Polarityionic
•H-bond
Good dissolver
solvent
pH
• Occasionally a hydrogen atom may shift from one water
molecule to another
• This H atom leaves its electron behind – this creates a
hydrogen ion (H+) – a single proton with a charge of 1+
• The other water molecule that lost a proton is now a
hydroxide ion (OH-), with a charge of 1• The proton binds to a separate water molecule, making
that molecule a hydronium ion (H3O+)
– H+ won’t exist independently in solution – it will always be
associated with another water molecule
• Concentrations of H+ and OH- are equal in pure water,
but adding acids and bases disrupts this balance
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H+ Ion
Concentration
pH
• pH is ONLY a measure of [H+]
• Equilibrium of pure water:
– Hydrogen ion [H+]= 1 x 10-7
– Hydroxide ion [OH-]= 1 x 10-7
– [H+] + [OH-] = 1 x 10-14
• What is the pH of a solution with [H+] of 10-10?
• What is the pH of a solution with [OH-] of 10-8?
• Please remember that pH declines as [H+]
increases and each pH unit represents a 10x
difference in [H+] and [OH-]
Acids
• Acid: substance that increases the H+ ion
concentration of a solution
– It donates H+ (hydrogen ions) to water to form
hydronium ions [H3O+] (this is called the
ionization of water)
– pH of 0 – 6.99 (low pH [high H+])
– Tastes sour
– Turns litmus paper red
– Strong acids completely dissociate
to form ions
100
pH
pH Scale
0
Examples of Solutions
Hydrochloric acid
10–1
1
tenfold change
in H+ ions
10–2
2
Stomach acid, Lemon juice
10–3
3
Vinegar, cola, beer
pH1  pH2
10-1  10-2
10–4
4
Tomatoes
10–5
5
Black coffee, Rainwater
10–6
6
Urine, Saliva
10–7
7
Pure water, Blood
10 times more H+
10–8
8
Seawater
pH10  pH8
10-10  10-8
10–9
9
Baking soda
10–10
10
Great Salt Lake
10–11
11
Household ammonia
10–12
12
Household bleach
10–13
13
Oven cleaner
10–14
14
Sodium hydroxide
10 times less H+
pH8  pH7
10-8  10-7
100 times more
H+
The Threat of Acid Precipitation
• Refers to rain, snow, or fog with a pH
lower than pH 5.6
• Is caused primarily by
the mixing of different
pollutants with water
in the air
• Can damage life in
Earth’s ecosystems
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Bases
• Base: substance that reduces [H+] of a
solution
– It accepts/removes H+ from solution
– Donates hydroxide (OH-) ions
– pH of 7.01 – 14 (high pH [low H+])
– Tastes bitter, feels slimy
– Turns litmus paper blue
– Strong bases completely
dissociate to form ions
pH and [H+] conversions
Ex. pH= 8 or [H+]= 1 x 10-8
Ex. pH = 10, [H+]= ?
Ex. [H+]= 1 x 10-8 , pH?
pH= -log 10 [H+]
[H+] = 10-pH
pOH=-log10 [OH-]
pH+ pOH= 14
pH+ pOH= 1 x 10-14
Ex. [H+]= 1 x 10-4 [OH-]= 1 x 10-10
Ex. [H+]= 1 x 10-12 , [OH-]= ?
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pH=2 is 10 times
stronger than pH=3
pH =2 is 100 times
stronger than pH=4
pH=2 is 1000 times
stronger than 5
Ex. If pH is increased
from pH=4 to pH=2,
what does that mean?
Ex. If your pH=4, what is
your pH if you have 100
times less [H+]?
Buffers
• Internal pH of most cells is 7, and a slight change
can be detrimental
– pH affects molecule shape  shape affects function
 molecular function affects cellular function
• Buffer: substance that minimizes changes in [H+]
and [OH-] – it resists changes in pH
– These substances allow biological fluids to
maintain a relatively constant pH despite the
addition of acids or bases
– Donates H+ when [H+] falls, and absorbs H+
when [H+] rises
• Carbonic acid-bicarbonate buffer system and
your blood
Bicarbonate Buffer System in Blood
• Maintains blood pH between 7.38 and 7.42
• HCO3- = Bicarbonate (weak base)
• H2CO3 = Carbonic acid (weak acid)
• These two are in equilibrium
ACTION
EFFECT
• Strenuous exercise
• Fatty acid metabolism
• Acidic drug overdose (aspirin)
INCREASE IN [‘H+]
DECREASE pH
• Hyperventilation
REMOVE CO2
BUFFER’S RESPONSE
Equilibrium shifts to left
H2O + CO2 ← H2CO3 ← HCO3- + H+
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