Chapter 3: WATER AND THE FITNESS OF THE ENVIRONMENT

Chapter 3:

WATER AND THE FITNESS

OF THE ENVIRONMENT

Evelyn I. Milian

Instructor

2012

BIOLOGY I

BIOLOGY I. Chapter 3 – Water and the Fitness of the Environment

Class Collaborative Activity

• The class will be split into small groups of

2-4 students.

• Each small group should work together to

answer the following question and turn in

your answer to the instructor:

What are the functions and the importance

of water for life? – Write a paragraph or a list

of several sentences.

• You can use your previous knowledge,

your textbook, and any other resources

available in the classroom.

Evelyn I. Mi l ian - Instructor 2



Functions and Properties of Water

• Water is one of the most important and abundant compounds in all living systems. All livings things are 70-95% water.

• Water exists in the natural environment in all three physical states of matter: liquid, solid (ice), and gas (water vapor).

• Main Functions of Water:

Solvent: Water dissolves other substances (called the solutes).

Stabilizing Temperature: Water has a high heat capacity, it

can absorb or release a relatively large amount of heat.

Chemical reactions: Water serves as the medium, as reactant

or product in chemical reactions in the body.

Lubricant: Water is a major part of mucus and other lubricating

fluids in the body of an organism.


Properties of Water: A Polar Molecule

• Water is a polar molecule. The red electrons are shared unequally.

Because the oxygen nucleus attracts the shared electrons more strongly,

the oxygen end of a water molecule has a partial negative charge; written

--, and the hydrogen ends have partial positive charges, written +.

Evelyn I . Mi l ian - Ins t ructor 4


Properties of Water: Polarity

• Water is a polar molecule; opposite ends of the molecule have opposite charges.

Oxygen attracts the shared electrons more strongly and is partially negative.

The two hydrogen atoms are partially positive.

The polarity of water molecules results in hydrogen bonding.

• The charged regions of a polar water molecule are attracted to oppositely charged parts of neighboring water molecules. Each molecule can hydrogen-bond to multiple partners, and these associations are constantly changing. The hydrogen bonds form, break, and re-form with great frequency.





Emergent Properties of Water

• The emergent properties of water contribute to

Earth’s fitness for life and water’s key functions.

1. Cohesion, adhesion, surface tension

2. Moderation of temperature by water:

• High heat capacity or high specific heat

• Evaporative cooling

3. Floating of ice on liquid water:

• Expansion upon freezing

4. Versatility as a solvent

• Water is the solvent of life.



Emergent Properties of Water: Cohesion and Adhesion


• Cohesion is the binding together of like molecules, often by hydrogen bonds.

• Cohesion due to hydrogen bonds between water molecules contributes to the transport of water and dissolved nutrients against gravity in plants; it helps hold together the column of water within the cells.

• Adhesion, the attraction between different kinds of molecules, also plays a role. Adhesion of water to cell walls by hydrogen bonds helps resist the downward pull of gravity.

• In plants, evaporation from leaves pulls water upward from the roots through water-conducting cells. Because of the properties of cohesion and adhesion, the tallest trees can transport water more than 100 meters upward.


Emergent Properties of Water: Cohesion and Surface Tension


• Related to cohesion is

surface tension, a

measure of how difficult

it is to stretch or break

the surface of a liquid.

• Water has a high surface

tension because of the

hydrogen bonding of

surface molecules. This

makes water behave as

though coated with an

invisible film.

• Due to water’s high surface tension

resulting from the collective strength of

its hydrogen bonds, some animals can

stand, walk, or run on water without

breaking the surface.


Emergent Properties of Water: Moderation of Temperature

• The ability of water to stabilize temperature stems from its relatively high heat capacity (or high specific heat) due to its hydrogen bonding.

Heat, a form of energy, is a measure of the total amount of kinetic energy (energy of motion) in matter due to motion of its molecules.

• Specific heat is the amount of heat that must be absorbed or lost for 1 g (gram) of a substance to change its temperature by 1C (Celsius).

Compared to most other fluids, water absorbs more heat energy before it gets measurably hotter.

What is the relevance of water’s high heat capacity to life on Earth?

• Large bodies of water can absorb and store a large amount of heat from the sun and warm up only a few degrees. It stabilizes ocean temperatures, creating a favorable environment for marine life. By absorbing or releasing heat, oceans moderate coastal climates.

• The water that covers most of Earth keeps temperature fluctuations on land and in water within limits that permit life!




• Water has a high heat of vaporization, the quantity of heat a

liquid must absorb for 1 g of it to be converted from the liquid to

the gaseous state (evaporation).

• This property of water allows evaporative cooling; the surface

becomes cooler during evaporation, owing to a loss of highly

kinetic molecules to the gaseous state (the “hottest” molecules

are the most likely to leave as gas).

• Evaporative cooling of water contributes to the stability of

temperature in lakes and ponds and also provides a mechanism

that prevents terrestrial organisms from overheating.

For example, evaporation of sweat from human skin dissipates body

heat and helps prevent overheating on a hot day or when excess

heat is generated by strenuous activity.




• (a) Water can be a solid, a liquid, or a gas at naturally occurring environmental temperatures. At room temperature and pressure, water is a liquid. When water freezes and becomes a solid (ice), it gives off heat, and this heat can help keep the environmental temperature higher than expected. On the other hand when water evaporates, it takes up a large amount of heat as it changes from a liquid to a gas. (b) This means that splashing water on the body will help keep body temperature within a normal range.



Emergent Properties of Water: Floating of Ice on Liquid

Water: Expansion Upon Freezing

• Insulation of Bodies of Water by Floating Ice

Water becomes locked into a crystalline lattice as the

temperature falls to 0C (it freezes).

Ice is less dense than liquid water because its more

organized hydrogen bonding causes expansion into a

crystal formation. (Density is the mass per unit volume of a

substance under specified conditions of pressure and temperature).

The lower density causes ice to float, which insulates

water below, preventing it from freezing, and thus

protecting life under the frozen surfaces of lakes and

polar seas.



Emergent Properties of Water: Expansion Upon Freezing


Figure 3.6 – Ice: crystalline structure and floating barrier. In ice, each

molecule is hydrogen-bonded to four neighbors in a three-dimensional crystal.

Because the crystal is spacious, ice has fewer molecules than an equal volume

of liquid water. In other words, ice is less dense than liquid water. Floating ice

becomes a barrier that protects the liquid water below from the colder air. The

marine organism shown here is a type of shrimp called krill.


Emergent Properties of Water: The Solvent of Life

IMPORTANT TERMS

Solution A liquid that is a completely homogeneous mixture of two

or more substances; it contains dissolved substances.

Solvent A substance that dissolves another substance; the

dissolving agent of a solution.

An aqueous solution is one in which water is the solvent.

Solute The substance that is dissolved (by the solvent).

Hydrophilic

(hydro = water;

philic = loving)

Having an affinity for water (“water-loving”). Hydrophilic

compounds or substances interact readily with water.

Polar molecules, with polar covalent bonds (electrons

shared unequally, resulting in slight negative and positive

charges), attract water molecules and dissolve in water.

Hydrophobic

(hydro = water;

phobic = fearing)

Having an aversion to water (tending to coalesce and

form droplets in water, “water-fearing”).

Nonpolar molecules lack polar covalent bonds or have

very few; they do not have positive and negative ends.

These molecules do not dissolve in water or mix poorly. Evelyn I . Mi l ian - Ins t ructor 15


Emergent Properties of Water: Versatility as a Solvent

• Water is an unusually versatile

solvent because its polar

molecules are attracted to

charged (ionic) and polar

substances.

• These substances (solutes)

dissolve in water, resulting in a

solution.

Examples: sugar, NaCl

(sodium chloride, or table salt).

• Most chemical reactions in

organisms involve solutes

dissolved in water.


A salt dissociates into cations and anions, neither of which is H+ or OH--.




Emergent Properties of Water: Versatility as a Solvent

• Figure 3.8 (9th Edition, Campbell) – A water-soluble protein.

This figure shows human lysozyme (purple), a protein found in

tears and saliva that has antibacterial action.



Dissociation of Water: Acids and Bases

• Occasionally, water molecules dissociate and the result is a hydronium ion (H3O

+) and a hydroxide ion (OH--).

• The hydrogen atom leaves an electron behind, and what is actually transferred is a hydrogen ion (H+); so this reaction can be simplified as:

H2O H+ + OH–.

• The dissociation of water is reversible and has important consequences for cells. Hydrogen and hydroxide ions are very reactive and changes in their concentrations can drastically affect a cell’s proteins and other complex molecules. Acidity in the human body affects the rate of most chemical reactions and the concentration of many chemicals.

• In pure water, the concentrations of H+ + OH– are equal, but adding certain kinds of solutes, called acids and bases, disrupts this balance.




Acids and Bases

• Acid

Substance that increases the hydrogen ion concentration of a solution.

An acid dissociates into one or more hydrogen ions (H+) and one or

more anions (negatively charged ions). An acid is a hydrogen ion

donor, or proton donor (because hydrogen ions have a proton, with

positive charge).

• Base

A substance that reduces the hydrogen ion concentration of a solution.

A base dissociates into one or more hydroxide ions (OH--) and one or

more cations (positively charged ions). A base is a hydrogen ion

acceptor , or proton acceptor (hydroxide ions have a strong attraction

for protons).


Acids and Bases

• An acid increases the hydrogen ion (H+) concentration of a solution.

• For example, when hydrochloric acid (HCl) is added to water, hydrogen ions dissociate from chloride ions:

HCl H+ + Cl-

• Strong acids (such as HCl) dissociate completely in water.

• Weak acids partially dissociate in water (reversibly release and accept back hydrogen ions; notice the arrow in both directions). For example, carbonic acid (H2CO3):

H2CO3 HCO3 -- + H+

(carbonic acid) (bicarbonate ion) (hydrogen ion)



Acids and Bases

• A base reduces the hydrogen ion (H+) concentration of a solution.

• Some bases do it directly by accepting H+. For example, ammonia (NH3):

NH3 + H+ NH4+ (ammonium ion)

• Other bases do it indirectly by dissociating to form hydroxide ions (OH--) which then combine with H+ in the solution to form water. For example, sodium hydroxide:

NaOH Na+ + OH--

• Strong bases (such as NaOH) dissociate completely in water.

• Weak bases (such as NH3) partially dissociate in water.



Salts and Water

• A salt is any compound that dissolves easily in water and

releases ions other than H+ and OH--.

• A salt commonly forms when an acid interacts with a base.

HCl (acid) + NaOH (base) NaCl (salt) + H2O

Hydrochloric acid Sodium hydroxide Sodium chloride Water

Na+ Cl-- (ionization)

• Many of the ions release when salts dissolve in fluid are

important components of cellular processes.

Ions of calcium, sodium, and potassium are involved in the functions of

nerve and muscle cells, and help plant cells take up water from the soil.






Effects of Changes in pH: The pH Scale

• pH A measure of hydrogen ion concentration, [H+].

The pH of a solution is defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH = -- log [H+ ]

• Each pH unit represents a change of 10 times in the H+ concentration.

• pH Scale A pH scale measures the relative acidity or alkalinity of a solution.

Extends from 0 to 14 (acidity → alkalinity) and is based on the concentration of H+ in moles per liter (m/L) (1 mole = 6.02 x 1023 molecules; Avogadro’s number).

neutral pH = 7 (midpoint): the concentrations of H+ and OH- are equal. [H+ ] = 10--7 M. Nearly all of life’s chemistry occurs near pH 7.

• log 10--7 = ( 7) = 7

Acidic solution: Has more H+ than OH--; pH below 7. Notice that pH declines as [H+] increases. The lower the pH, the more acidic the solution.

Basic (alkaline) solution: Has more OH- than H+; pH above 7. The higher the pH, the more basic the solution.


The pH Scale and pH

Values of Some Aqueous

Solutions

• pH is a measure of

hydrogen ion

concentration, [H+]. The

pH scale measures the

relative acidity or

alkalinity of a solution.

• Nearly all of life’s

chemistry occurs near pH

7. Most of your body’s

internal environment

(tissue fluids and blood) is

between pH 7.3 and 7.5.

26 Evelyn I . Mi l ian - Ins t ructor


The pH Scale and pH Values of Some Aqueous Solutions




Effects of Changes in pH: Buffers

• Cells must respond fast to even slight changes in pH, because excess H+ or OH– can alter the functions of biological molecules.

• A buffer is a substance that helps maintain a constant pH by minimizing changes in the concentrations of H+ and OH-- in a solution. Buffers can convert strong acids or bases into weak ones.

• A buffer works by accepting hydrogen ions from the solution when they are in excess and donating hydrogen ions to the solution when they have been depleted (they absorb excess H+ or OH– ions).

• * Buffer systems in the body stabilize or maintain the pH of fluids inside and outside cells almost constant; they resist changes in pH.

Example: carbonic acid-bicarbonate buffer system:

Response to a rise in pH

H2CO3 HCO3-- + H+

carbonic acid Response to (bicarbonate; Hydrogen ion (H+ donor) a drop in pH base: H+ acceptor)






References

• Audesirk, Teresa; Audesirk, Gerald & Byers, Bruce E. (2005). Biology: Life on Earth. Seventh Edition. Pearson Education, Inc.-Prentice Hall. NJ, USA.

• Campbell, Neil A.; Reece, Jane B., et al. (2011). Campbell Biology. Ninth Edition. Pearson Education, Inc.-Pearson Benjamin Cummings. CA, USA.

• Enger, Eldon D.; Ross, Frederick C.; Bailey, David B. (2007). Concepts in Biology. Twelfth Edition. The McGraw-Hill Companies, Inc. NY, USA.

• Mader, Sylvia S. (2010). Biology. Tenth Edition. The McGraw-Hill Companies, Inc. NY, USA.

• Presson, Joelle & Jenner, Jan. (2008). Biology, Dimensions of Life. The McGraw-Hill Companies, Inc. NY, USA.

• Solomon, Eldra; Berg, Linda; Martin, Diana W. (2008). Biology. Eighth Edition. Cengage Learning. OH, USA.

• Starr, Cecie. (2008). Biology: Concepts and Applications Volume I. Houston Community College. Thompson Brooks/Cole. OH, USA.

• Tortora, Gerard J.; Derrickson, Bryan. (2006). Principles of Anatomy and Physiology. Eleventh Edition. John Wiley & Sons, Inc. NJ, USA. www.wiley.com/college/apcentral.

• Tortora, Gerard J.; Funke, Berdell R.; Case, Christine L. (2010). Microbiology An Introduction. Tenth Edition. Pearson Education, Inc.-Pearson Benjamin Cummings; CA, USA. www.microbiologyplace.com.


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Chapter 3: WATER AND THE FITNESS OF THE ENVIRONMENT