Properties of water

1. W a t e r –
C h a p t e r 3

2. Properties of Water
• Polar molecule
• Cohesion and
adhesion
• High specific heat
• Density – greatest
at 4o
C
• Universal solvent
of life

3. Polarity of Water
• In a water molecule two hydrogen atoms
form single polar covalent bonds with an
oxygen atom. Gives water more structure
than other liquids
– Because oxygen is more electronegative, the
region around oxygen has a partial negative
charge.
– The region near the two hydrogen atoms has a
partial positive charge.
• A water molecule is a polar molecule with
opposite ends of the molecule with opposite
charges.

4. • Water has a variety of unusual properties
because of attractions between these polar
molecules.
– The slightly negative regions of one molecule are
attracted to the slightly positive regions of nearby
molecules, forming a hydrogen bond.
– Each water molecule
can form hydrogen
bonds with up to
four neighbors.
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
Fig. 3.1

5. HYDROGEN BONDS
• Hold water molecules
together
• Each water molecule can
form a maximum of 4
hydrogen bonds
• The hydrogen bonds
joining water molecules
are weak, about 1/20th
as
strong as covalent bonds.
• They form, break, and
reform with great
frequency
• Extraordinary Properties
that are a result of hydrogen
bonds.
– Cohesive behavior
– Resists changes in
temperature
– High heat of vaporization
– Expands when it freezes
– Versatile solvent

6. Organisms Depend on Cohesion
• Cohesion is responsible for the
transport of the water column in
plants
• Cohesion among water molecules
plays a key role in the transport of
water against gravity in plants
• Adhesion, clinging
of one substance to
another, contributes
too, as water adheres
to the wall of the
vessels.
Hydrogen bonds hold the substance
together, a phenomenon called cohesion

7. • Surface tension, a measure of the force
necessary to stretch or break the surface of a
liquid, is related to cohesion.
– Water has a greater surface tension than most other
liquids because hydrogen bonds among surface
water molecules resist stretching or breaking the
surface.
– Water behaves as if
covered by an invisible
film.
– Some animals can stand,
walk, or run on water
without breaking the
surface.
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
Fig. 3.3

8. Moderates Temperatures on Earth
• What is kinetic energy?
• Heat?
• Temperature?
• Calorie?
• What is the difference
in cal and Cal?
• What is specific heat?
Celsius Scale at Sea Level
100o
C Water boils
37o
C Human body
temperature
23o
C Room temperature
0o
C Water freezes
Water stabilizes air temperatures by absorbing heat from
warmer air and releasing heat to cooler air.
Water can absorb or release relatively large amounts of heat
with only a slight change in its own temperature.

9. Three-fourths of the earth is covered
by water. The water serves as a
large heat sink responsible for:
2. Prevention of temperature
fluctuations that are outside the
range suitable for life.
3. Coastal areas having a mild
climate
4. A stable marine environment
Specific Heat is the amount of heat that must be
absorbed or lost for one gram of a substance to
change its temperature by 1o
C.

10. Evaporative Cooling
• The cooling of a
surface occurs when
the liquid evaporates
• This is responsible for:
– Moderating earth’s
climate
– Stabilizes
temperature in
aquatic ecosystems
– Preventing organisms
from overheating

11. Density of Water
• Most dense at 4o
C
• Contracts until 4o
C
• Expands from 4o
C to
0o
C
The density of water:
2. Prevents water from freezing from the bottom up.
3. Ice forms on the surface first—the freezing of the
water releases heat to the water below creating
insulation.
4. Makes transition between season less abrupt.

12. – When water reaches 0o
C, water becomes locked into
a crystalline lattice with each molecule bonded to to
the maximum of four partners.
– As ice starts to melt, some of the hydrogen bonds
break and some water molecules can slip closer
together than they can while in the ice state.
– Ice is about 10% less dense than water at 4o
C.
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
Fig. 3.5

13. Solvent for Life
• Solution
– Solute
– solvent
• Aqueous solution
• Hydrophilic
– Ionic compounds
dissolve in water
– Polar molecules
(generally) are water
soluble
• Hydrophobic
– Nonpolar compounds

14. Most biochemical reactions
involve solutes dissolved in water.
• There are two important
quantitative proprieties of aqueous
solutions.
–1. Concentration
–2. pH

15. Concentration of a Solution
• Molecular weight – sum of the weights of all atoms in
a molecule (daltons)
• Mole – amount of a substance that has a mass in
grams numerically equivalent to its molecular weight
in daltons.
• Avogadro’s number – 6.02 X 1023
– A mole of one substance has the same number of molecules
as a mole of any other substance.

16. Molarity
The concentration of a material in solution is called its molarity.
A one molar solution has one mole of a substance dissolved
in one liter of solvent, typically water.
Calculate a one molar solution of sucrose, C12H22O16.
C = 12 daltons
H = 1 dalton
O = 16 daltons
12 x12 = 144
1 x 22 = 22
16 x 11 = 176
342For a 2M solution?
For a .05 M solution?
For a .2 M solution?

17. • Occasionally, a hydrogen atom shared by two
water molecules shifts from one molecule to the
other.
– The hydrogen atom leaves its electron behind and is
transferred as a single proton - a hydrogen ion (H+
).
– The water molecule that lost a proton is now a
hydroxide ion (OH-
).
– The water
molecule with
the extra proton
is a hydronium
ion (H3O+
).
Dissociation of Water Molecules
Copyright © 2002 Pearson Education, Inc., publishing as Benjamin Cummings
Unnumbered Fig. 3.47

18. • A simpler way to view this process is that a
water molecule dissociates into a hydrogen
ion and a hydroxide ion:
– H2O <=> H+
+ OH-
• This reaction is reversible.
• At equilibrium the concentration of water
molecules greatly exceeds that of H+
and
OH-
.
• In pure water only one water molecule in
every 554 million is dissociated.
– At equilibrium, the concentration of H+
or OH-
is 10-7
M (25°C) .

19. Acids and Bases
• An acid is a substance that
increases the hydrogen ion
concentration in a solution.
• Any substance that reduces the
hydrogen ion concentration in a
solution is a base.
– Some bases reduce H+
directly by
accepting hydrogen ions.
• Strong acids and bases complete
dissociate in water.
• Weak acids and bases dissociate
only partially and reversibly.

20. pH Scale
• The pH scale in any aqueous solution :
– [ H+
] [OH-
] = 10-14
• Measures the degree of acidity (0 – 14)
• Most biologic fluids are in the pH range
from 6 – 8
• Each pH unit represents a tenfold
difference (scale is logarithmic)
– A small change in pH actually indicates a
substantial change in H+
and OH-
concentrations.

21. Problem
How much greater is the [ H+
] in a
solution with pH 2 than in a solution with
pH 6?
Answer:
pH of 2 = [ H+
] of 1.0 x 10-2
= 1/100 M
pH of 6 = [ H+
] of 1.0 x 10-6
= 1/1,000,000 M
10,000 times greater

22. Buffers
• A substance that eliminates large sudden
changes in pH.
• Buffers help organisms maintain the pH of
body fluids within the narrow range
necessary for life.
– Are combinations of H+
acceptors and
donors forms in a solution of weak acids
or bases
– Work by accepting H+
from solutions
when they are in excess and by donating
H+
when they have been depleted.

23. Acid Precipitation
• Rain, snow or fog with more strongly acidic than
pH of 5.6
• West Virginia has recorded 1.5
• East Tennessee reported 4.2 in 2000
• Occurs when sulfur oxides and nitrogen oxides
react with water in the atmosphere
– Lowers pH of soil which affects mineral
solubility – decline of forests
– Lower pH of lakes and ponds – In the
Western Adirondack Mountains, there are
lakes with a pH <5 that have no fish.