BIOL 1020 - CHAPTER 3 LECTURE NOTES
Chapter 3: Water and the fitness of the environment : What’s so great about water?
Overview: Life as we know it requires water. All organisms that we know of are made mostly of liquid water, and most of their
metabolism requires an aqueous medium. In addition, many organisms live in liquid water or in an environment dominated by water in
its various states (solid, liquid, or gas). Some numbers:
 cells are typically 70% or more water by mass
 about 75% of the Earth’s surface is covered by liquid water
But then, just being common on the Earth doesn’t make something essential for life. A large percentage of the Earth’s crust is sand,
but we don’t consider sand a requirement for life. What is it about water that makes it so special?
I.
The chemistry of water is dominated by the polar nature of water molecules.
A. oxygen atoms are electron seeking (electronegative), especially compared to hydrogen; thus for an oxygen-hydrogen bond:
1. the oxygen atom has a partial (-) charge
2. the hydrogen atoms have a partial (+) charge
B. the polar character of water allows water molecules to form many (up to 4) hydrogen bonds.
II. What properties of water are important for life?
A. four properties of water are critical for life as we know it, and all of them come in some way from water’s polar nature and the
resulting tendency of water to form hydrogen bonds and similar interactions
1. water is the principal solvent in living things
2. water exhibits both cohesive and adhesive forces
3. water helps maintain a stable temperature
4. ice (solid water) floats in liquid water
B. water is the principal solvent in living things
1. the highly polar character of water makes it an excellent solvent for other polar substances, and for ionic compounds
2. hydrophilic substances – interact readily with water
3. water does not readily dissolve nonpolar (hydrophobic) substances – thus, hydrophobic substances are good components
for membranes
C. water exhibits both cohesive and adhesive forces (due to hydrogen bonding)
1. cohesive forces are caused by the attraction of water molecules to other water molecules, and give water a high surface
tension (the ability of a water surface to withstand some stress)
2. adhesive forces cause water molecules to be attracted to other kinds of molecules; it is how things are made wet
3. capillary action, the tendency of water to move up narrow tubes even against gravity, results from cohesion and
adhesion; living organisms take advantage of this
D. water helps maintain a stable temperature
1. the unusual specific heat of water leads to temperature stability
 specific heat - the amount of energy required to raise the temperature of a specific amount of a substance one degree
Celsius (for water, 1 calorie = heat needed for 1 g of water to raise by one degree Celsius)
 the specific heat of water is much higher than most other substances, due to hydrogen bonding (as a comparison, the
specific heat of sand is about 0.2 calories for 1 g)
 thus, it requires the gain or loss of more energy (heat) to change the temperature of water than it does other
substances
 since much of the ecosphere is water, and most biological organisms are more than 70% water, this property of water
leads to temperature stability (which is critical for most living organisms)
2. the high heat of vaporization of water helps cool the ecosphere and biological organisms
 heat of vaporization is the amount of heat energy required to convert one gram of liquid into the gaseous state
 because of hydrogen bonds in liquid water, water has an extremely high heat of vaporization
 a commonly used unit for measuring energy is the calorie, the amount of heat (energy) required to cause the
temperature of one gram of pure water to rise one degree Celsius
 540 calories are required to convert one gram of liquid water into water vapor
 biological organisms take advantage of this property of water to cool themselves, examples sweating and cooling a
leaf
E. ice floats (ice is less dense than liquid water)
1. liquid water, like most substances, becomes denser as it cools – but only up to a point
2. at 4ºC (under standard atmospheric pressure), water begins to expand as it cools further – that is, it gets less dense from
then on – due to hydrogen bonds becoming locked in place
3. at 0ºC ice freezes into a crystal based on the placement of hydrogen bonds
4. floating ice keeps lakes, etc., from freezing solid and is important for temperature cycling on the planet
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BIOL 1020 - CHAPTER 3 LECTURE NOTES
III. Acids and Bases
A. acids are proton donors
1. an acid is a substance that dissociates in solution to yield hydrogen ions (H +)
 HA (an acid)
H+ + A- (an anion)
2. note that hydrogen has an atomic number of 1, which means that the nucleus has only one proton
 when the atom loses its electron to become a hydrogen ion, all that remains is the proton
 thus, hydrogen ions are sometimes referred to as protons
 therefore, any substance that yields a proton is an acid, or an acid is a proton donor
B. bases are proton acceptors
1. a base is a substance that can accept a proton
2. bases either dissociate in water to produce hydroxide ions and a cation, or split water to form a cation and hydroxide ion:
NaOH
Na+ + OHor
B (a base) + HOH
BH+ + OHC. water tends to slightly dissociate into hydrogen and hydroxide ions (H+ and OH-)
HOH
H+ + OHin pure water, the concentrations of these ions are equal: [H+] = [OH-] = 10-7 M
(note that the designation M stands for molar, the moles of a substance per liter of solution)
2. the product of these remains constant: [H+] x [OH-] = 10-14
 acidic solutions have an elevated [H+], and thus reduced [OH-]
 basic solutions have an elevated [OH-], and thus reduced [H+]
D. the pH scale is a convenient short hand notation to express the proton concentration of a solution
1. the pH of a solution is defined as the reciprocal of the logarithm of the proton concentration in the solution, or -log[H+]
2. pure water (having a proton concentration of 10 -7 M) has pH = 7
3. a pH below 7 is acidic, and a pH above 7 is basic
4. the pH of most living cells is usually around 7.2 to 7.4
E. buffers minimize pH changes
1. weak acids and weak bases serve as buffers
2. living things use buffers to prevent dramatic changes in pH, which can kill them
3. carbonate/bicarbonate is an example of a biologically important buffer system
 CO2 + H2O
H2CO3
H+ + HCO3 because these reactions are equilibrium reactions, as H+ is added to this system bicarbonate (HCO3-) and protons are
consumed to keep the proton concentration constant (pH constant) producing H 2CO3
 if OH- ions are added to the system, H+ ions are consumed forming water. More H+ ions are produced by the
dissociation of H2CO3 to form protons and bicarbonate. This demonstrates how the pH is stabilized as acid or
base is added
IV. Some useful definitions
A. solvent – a liquid into which a substance dissolves
B. solute – the dissolved substance
C. solution = solvent + solute
D. salts – form from acids and bases
1. water is formed
2. the cation of the base and the anion of the acid form the salt
 HCl + NaOH
NaCl + HOH
E. electrolytes are salts, acids, or bases that form ions in water and thus can conduct an electrical current when dissolved in
water (pure water is a poor conductor of electricity, but put in a salt and it becomes an excellent conductor)
F. nonelectrolytes are substances like sugar that dissolve in water but do not become ionic
G. mixtures - a mixture of 2 or more elements and/or compounds; they can be broken down into elements and compounds by
simple physical means. There are two types:
1. heterogeneous mixtures - mixtures that are not of uniform composition throughout - a living organism is a good example
2. homogeneous mixtures - mixtures that are completely uniform throughout - a salt water solution is a good example
1.
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