BSC 1010C
General Biology I
Dr. Graeme Lindbeck
glindbeck@valenciacollege.edu

Water and the Fitness of the Environment

Outline

The polarity of water molecules results in hydrogen bonding
Organisms depend on the cohesion of water molecules
Water contributes to Earth's habitability by moderating temperatures
1. Heat and Temperature
2. Water's High Specific Heat
3. Evaporative Cooling
Oceans and lakes don't freeze solid because ice floats
Water is the solvent of life
1. Hydrophilic and Hydrophobic Substances
2. Solute Concentration in Aqueous Solutions
Organisms are sensitive to changes in pH
1. Dissociation of Water Molecules
2. Acids and Bases
3. The pH Scale
4. Buffers
Acid precipitation threatens the fitness of the environment

Water contributes to the fitness of the environment to support life.

Life on earth probably evolved in water.
Living cells are 70%-95% H₂O
Water covers about ³/₄ of the earth.
In nature, water naturally exists in all three physical states of matter - solid, liquid and gas.

Water's extraordinary properties are emergent properties resulting from water's structure and molecular interactions.

I. The polarity of water molecules results in hydrogen bonding

Water is a polar molecule. Its polar bonds and asymmetrical shape give water molecules opposite charges on opposite sides.

Hydrogen bonding orders water into a higher level of structural organization.

The polar molecules of water are held together by hydrogen bonds.
Positively charged H of one molecule is attracted to the negatively charged O of another water molecule.
Each water molecule can form a maximum of four hydrogen bonds with neighboring water molecules.

Water has extraordinary properties that emerge as a consequence of its polarity and hydrogen-bonding. Some of these properties are that water:

has cohesive behavior
resists changes in temperature
has a high heat of vaporization and cools surfaces as it evaporates
expands when it freezes
is a versatile solvent

II. Organisms depend on the cohesion of water molecules

Cohesion = Phenomenon of a substance being held together by hydrogen bonds.

Though hydrogen bonds are transient, enough water molecules are hydrogen bonded at any given time to give water more structure than other liquids.
Contributes to upward water transport in plants by holding the water column together. Adhesion of water to vessel walls counteracts the downward pull of gravity.

Surface tension = Measure of how difficult it is to stretch or break the surface of a liquid.

Water has a greater surface tension than most liquids; function of the fact that at the air/H₂O interface, surface water molecules are hydrogen bonded to each other and to the water molecules below.
Causes H₂O to bead (shape with smallest area to volume ratio and allows maximum hydrogen bonding).

III. Water contributes to Earth's habitability by moderating temperatures

Heat and Temperature

Kinetic energy = The energy of motion.

Heat = Total kinetic energy due to molecular motion in a body of matter.

Calorie (cal) = Amount of heat it takes to raise the temperature of one gram of water by one degree Celsius.

Kilocalorie (kcal or Cal) = Amount of heat required to raise the temperature of one kilogram of water by one degree Celsius (1000 cal).

Temperature = Measure of heat intensity due to the average kinetic energy of molecules in a body of matter.

Celsius Scale at Sea Level
100^oC (212^oF) = water boils
37^oC (98.6^oF) = human body temperature
23^oC (72^oF) = room temperature
0^oC (32^oF) = water freezes

Scale Conversion
^oC = 5(^oF -32)/9
^oF = 9^oC/5 +32
^oK = ^oC + 273

Water's High Specific Heat

Water has a high specific heat, which means that it resists temperature changes when it absorbs or releases heat.

Specific heat = Amount of heat that must be absorbed or lost for one gram of a substance to change its temperature by one degree Celsius.

Specific heat of water = One calorie per gram per degree Celsius (1 cal/g/^oC).

As a result of hydrogen bonding among water molecules, it takes a relatively large heat loss or gain for each 1 ^oC change in temperature.
Hydrogen bonds must absorb heat to break, and they release heat when they form.
Much absorbed heat energy is used to disrupt hydrogen bonds before water molecules can move faster (increase temperature).

A large body of water can act as a heat sink - absorbing heat from sunlight during the day and summer and releasing heat during the night and winter as the water gradually cools. As a result:

Water, which covers three-fourths of the planet, keeps temperature fluctuations within a range suitable for life.
Coastal areas have milder climates than inland.
The marine environment has a relatively stable temperature.

Evaporative Cooling

Vaporization (evaporation) = Transformation from liquid to a gas.

Molecules with enough kinetic energy to overcome the mutual attraction of molecules in a liquid, can escape into the air.

Heat of vaporization = Quantity of heat a liquid must absorb for 1g to be converted to the gaseous state.

For water molecules to evaporate, hydrogen bonds must be broken which requires heat energy.
Water has a relatively high heat of vaporization at the boiling point (540 cal/g or 2260 J/g; Joule = 0.239 cal).

Evaporative cooling = Cooling of a liquid's surface when a liquid evaporates.

The surface molecules with the highest kinetic energy are most likely to escape into gaseous form; the average kinetic energy of the remaining surface molecules is thus lower.

Water's high heat of vaporization:

Moderates the earth's climate.
Þ Solar heat absorbed by tropical seas dissipates when surface water evaporates (evaporative cooling).
Þ As moist tropical air moves poleward, water vapor releases heat as it condenses into rain.
Stabilizes temperature in aquatic ecosystems (evaporative cooling).
Helps organisms from overheating by evaporative cooling.

IV. Oceans and lakes don't freeze solid because ice floats

Because of hydrogen bonding, water is less dense as a solid than it is as a liquid. Consequently, ice floats.

Water is densest at 4^oC.
Water contracts as it cools to 4^oC.
As water cools from 4^oC to freezing (0^oC), it expands and becomes less dense than liquid water (ice floats).
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.

Expansion of water contributes to the fitness of the environment for life:

Prevents deep bodies of water from freezing solid from the bottom up.
Since ice is less dense, it forms on the surface first. As water freezes it releases heat to the water below and insulates it.
Makes the transitions between seasons less abrupt. As water freezes, hydrogen bonds form releasing heat. As ice melts, hydrogen bonds break absorbing heat.

V. Water is the solvent of life

Solution = A liquid that is a homogenous mixture of two or more substances.

Solvent = Dissolving agent of a solution.

Solute = Substance dissolved in a solution.

Aqueous solution = Solution in which water is the solvent.

Water is a versatile solvent owing to the polarity of the water molecule.

Ionic compounds dissolve in water.
- Charged regions of polar water molecules have an electrical attraction to charged ions.
- Water surrounds individual ions, separating and shielding them from one another.
Polar compounds in general, are water-soluble.
- Charged regions of polar water molecules have an affinity for oppositely charged regions of other polar molecules.
Nonpolar compounds (which have symmetric distribution in charge) are NOT water soluble.

Hydrophilic and Hydrophobic Substances

Ionic and polar substances are hydrophilic, but non-polar compounds are hydrophobic.

Hydrophilic = (Hydro=water; philo=loving) Property of having an affinity for water.

Some large hydrophilic molecules can absorb water without dissolving.

Hydrophobic

Solute Concentration in Aqueous Solutions

Most biochemical reactions involve solutes dissolved in water. There are two important quantitative properties of aqueous solutions: solute concentration and pH.

Molecular weight = Sum of the weight of all atoms in a molecule (expressed in daltons).

Mole = Amount of a substance that has a mass in grams numerically equivalent to its molecular weight in daltons.

For example, to determine a mole of sucrose (C₁₂H₂₂O₁₁):

Calculate molecular weight:
C = 12 dal	12 dal x 12 =	144 dal
H = 1 dal	1 dal x 22 =	22 dal
O = 16 dal	16 dal x 11 =	176 dal

Express it in grams (342 g).

Molarity = Number of moles of solute per liter of solution.

For example, to make a 1M sucrose solution, weigh out 342 g of sucrose and add water up to 1L.

Advantage of measuring in moles:

Rescales weighing of single molecules in daltons to grams, which is more practical for laboratory use.
A mole of one substance has the same number of molecules as a mole of any other substance (6.02 x 1023 ; Avogadro's number).
Allows one to combine substances in fixed ratios of molecules.

VI. Organisms are sensitive to changes in pH

Dissociation of Water Molecules

Occasionally, the hydrogen atom that is shared in a hydrogen bond between two water molecules, shifts from the oxygen atom to which it is covalently bonded to the unshared orbitals of the oxygen atom to which it is hydrogen bonded.

Only a hydrogen ion (proton with a +1 charge) is actually transferred.
Transferred proton binds to an unshared orbital of the second water molecule creating a hydronium ion (H₃O⁺).
Water molecule that lost a proton has a net negative charge and is called a hydroxide ion (OH^-).
H₂O + H₂O ® H₃O⁺ + OH^-
By convention, ionization of H₂0 is expressed as the dissociation into H⁺ and OH^-.
H₂O ® H⁺ + OH^-
Reaction is reversible.
At equilibrium, most of the H₂O is not ionized.

Acids and Bases

At equilibrium in pure water at 25°C:

Number of H⁺ ions = number of OH^- ions.
[H⁺] = [OH^-] = 1/10,000,000 M = 10^-7 M
Note that brackets indicate molar concentration.

ACID
Substance that increases the relative [H⁺] of a solution.
Also removes OH^- because it tends to combine with H⁺ to form H₂O
For example: (in water) HCl ® H⁺ + Cl^-

BASE
Substance that reduces the relative [H⁺] of a solution.
May alternately increase [OH^-]
For example:	A base may reduce [H⁺] directly: NH₃ + H⁺ ® NH₄⁺
For example:	A base may reduce [H⁺] indirectly: NaOH ® Na⁺ + OH^- OH^- + H⁺® H₂O

A solution in which:

[H⁺] = [OH^-] is a neutral solution.
[H⁺] > [OH^-] is an acidic solution.
[H⁺] <[oh^-] is a basic solution.

Strong acids and bases dissociate completely in water.

For example, HCl and NaOH.
Single arrows indicate complete dissociation.
NaOH ® Na⁺ + OH^-

Weak acids and bases dissociate only partially and reversibly.

For example, NH₃ (ammonia) and H₂CO₃ (carbonic acid)

Double arrows indicate a reversible reaction; at equilibrium there will be a fixed ratio of reactants and products,

H₂CO₃
Carbonic
Acid ion

HCO₃^-
Bicarbonate
ion

H⁺
Hydrogen

The pH Scale

In any aqueous solution:

[H⁺][OH^-] = 1.0 x 10^-14
For example:
In a neutral solution, [H⁺] = 10^-7 M and [OH^-] = 10^-7 M.
In an acidic solution where the [H⁺]=10^-5 M, the [OH^-] = 10^-9 M.
In a basic solution where the [H⁺] = 10^-9 M, the [OH^-] = 10^-5 M.

pH scale = Scale used to measure degree of acidity. It ranges from 0 to 14.

pH = Negative log of the [H⁺] expressed in moles per liter.

pH of 7 is a neutral solution.
pH <7 is an acidic solution.
pH > 7 is a basic solution.
Most biological fluids are within the pH range of 6 to 8. There are some exceptions such as stomach acid with pH = 1.5.
Each pH unit represents a tenfold difference (scale is logarithmic), so a slight change in pH represents a large change in actual [H+]

Buffers

By minimizing wide fluctuations in pH, buffers help organisms maintain the pH of body fluids within the narrow range necessary for life (usually pH 6-8).

Buffer = Substance that prevents large sudden changes in pH.

Are combinations of H⁺-donor and H⁺-acceptor forms of weak acids or bases.
Work by accepting H⁺ ions from solution when they are in excess, and by donating H⁺ ions to the solution when they have been depleted.

For example: Bicarbonate buffer.

	response to a rise in pH
H₂CO₃	® ¬	HCO₃	+	H⁺
H⁺ donor (weak acid)	response to a drop in pH	H⁺ acceptor (weak base)		Hydrogen ion

HCl
strong
acid + NaHCO₃ ® H₂CO₃
weak
acid + NaCl

NaOH
strong
base + H₂CO₃ ® NaHCO₃
weak
base + H₂O

VII. Acid precipitation threatens the fitness of the environment

Acid precipitation = Rain, snow or fog more strongly acidic than pH 5.6.

Has been recorded as low as pH 1.5 in West Virginia.
Occurs when sulfur oxides and nitrogen oxides in the atmosphere react with water in the air to form acids which fall to Earth in precipitation.
Major oxide source is the combustion of fossil fuels by industry and cars.

Acid rain affects the fitness of the environment to support life:

Lowers soil pH which affects mineral solubility. May leach out necessary mineral nutrients and increase the concentration of minerals that are potentially toxic to vegetation in higher concentration (e.g. aluminum). This is contributing to the dec line of some European and North American forests.
Lowers the pH of takes and ponds, and runoff carries leached out soil minerals into aquatic ecosystems. This adversely affects aquatic life. For example: In the Western Adirondack Mountains, there are lakes with a pH <5 that have no fish.

What can be done to reduce the problem?

Add industrial pollution controls.
Develop and use anti-pollution devices.

Increase involvement of voters, consumers, politicians and business leaders.

Course Pages maintained by
Dr. Graeme Lindbeck.