AP Biology Ch. 3

1. AP Biology Ch. 3 Water and the Fitness of the Environment

2. The Molecule that Supports All of Life • Water is the substance that makes possible life as we know it here on Earth. • All organisms familiar to us are made mostly of water and live in an environment dominated by water.

3. Life on Earth began in Water • Life evolved in water for 3 billion years before spreading onto land. Modern life, even terrestrial life, remains tied to water.

4. All Living Organisms Require Water • Humans can survive several weeks without food, but only one week without water. • The abundance of water is the reason Earth is habitable.

5. The Structure of Water Leads to its Properties • Because water is a polar molecule, the two ends have opposite charges. • The polarity of water molecules results in hydrogen bonding—a very important property in biology.

6. Hydrogen Bonds Form Between Water Molecules • The charged regions of a polar water molecule are attracted to oppositely charged parts of neighboring molecules. • Each molecule can hydrogen-bond to multiple partners, and these associations are constantly changing.

7. Four Emergent Properties of Water are Important for Life • Cohesion • Ability to moderate temperature • Expansion upon freezing • Versatility as a solvent

8. Cohesion • Cohesion-water is polar, so water is attracted to water. • Linkages between water molecules make water more structured than other molecules • The collective hydrogen bonds hold the substance together, a phenomenon called cohesion

9. Adhesion • Adhesion= the clinging of one substance to another • Example: Water will cling to any substance that has a charge (either + or -) Why does water do this?

10. Water Movement in Plants is due to both Cohesion & Adhesion • Evaporation of water from leaves pulls water upward from the roots • Cohesion helps hold together the column of water within the cells • Adhesion of water to cell walls by hydrogen bonds helps resist the downward pull of gravity

11. 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 greater surface tension than most other liquids. • Water behaves as though it is coated with an invisible film. Water molecules are held together in a network of hydrogen bonds.

12. Moderation of Temperature • Water moderates air temperature by absorbing heat from air that is warmer and releasing stored heat to air that is cooler. Temperature= a measure of heat intensity that represents the average kinetic energy of the molecules When 2 objects of different temperatures are brought together, heat passes to the cooler object

13. Fig. 3-5 San Bernardino 100° Burbank 90° Santa Barbara 73° Riverside 96° Los Angeles (Airport) 75° Santa Ana 84° Palm Springs 106° 70s (°F) 80s Pacific Ocean 90s 100s San Diego 72° 40 miles By absorbing or releasing heat, oceans moderate coastal climates. In this example from an August day in Southern California, the relatively cool ocean reduces coastal temperatures by absorbing heat.

14. Water’s High Specific Heat • Specific Heat=the amount of heat that must be absorbed or lost for 1 gram of that substance to change its temperature by 1 degree Celsius. • Water has a specific heat of 1, much greater than most other substances The specific heat of water is 1 calorie/gram/degree Celsius

15. High Specific Heat of Water—Benefits for Living Things • Water must absorb a great deal of heat in order to change its temperature even one degree. • Near large bodies of water, the temperature is greatly moderated. • High specific heat of water stabilizes ocean temperatures. The water that covers Earth’s surface moderates its temperatures, allowing life

16. Evaporative Cooling • Water is held together with hydrogen bonds, so it resists evaporation until a lot of heat is added (water has a high heat of vaporization) • As water evaporates, the remaining liquid is cooler because the molecules with the most energy have left.

17. Ice Floats! • In ice, each molecule is hydrogen-bonded to four neighbors in a 3-dimensional crystal • Because the crystal is spacious, ice has fewer molecules than an equal volume of liquid water. • Ice is less dense, so it floats on top of liquid water

18. Fig. 3-6a Hydrogen bond Liquid water Hydrogen bonds break and re-form Ice Hydrogen bonds are stable

19. Water is the Solvent of Life • Water is a very versatile solvent due to its polarity • Water will dissolve any ionic compound easily—the – and + charges will be attracted to the water molecules -/+ sides. • Water will dissolve nonionic (but polar) molecules by surrounding them, forming Hydrogen bonds with them.

20. Hydrophilic & Hydrophobic • Hydrophilic=any substance that has an affinity for water (anything with a + or – charge) • Hydrophobic=any substance that does not have an affinity for water (anything that is nonionic, nonpolar) Oil spills such as the one that killed this bird can be very hazardous to ocean life.

21. Solute Concentrations in Aqueous Solutions • Most of the chemical reactions in organisms involve solutes dissolved in water (aqueous solutions). • To understand such reactions, we must know how many atoms and molecules are involved (the concentration of solutes in the water solution)

22. Step 1: Figure out the molecular mass • Molecular mass=the mass of each atom in a given molecule. • Example: Sucrose (C12H22O11) has a molecular mass of 342 daltons. (Carbon atomic mass is 12 x 12 = 144; Hydrogen atomic mass is 1 x 22 = 22; Oxygen atomic mass is 16 x 11 = 176; 144 + 22 + 176 = 342) • If we measure 342 daltons of sucrose, however, it would be extremely difficult! So, we usually measure it in moles

23. Step 2: Convert molecular mass to moles • Just as a dozen objects always means 12, a mole represents an exact number of objects (6.02 X 1023) which is called Avogadro’s number. • There are 6.02 X 1023daltons in 1 g. • Once we determine the molecular mass of a molecule such as sucrose, we can use the same number (342) with the unit gram to represent the mass of 6.02 X 1023 molecules of sucrose. 342 grams is 1 mole of sucrose.

24. Step 3: Measure the correct amount • A mole of one substance has exactly the same number of molecules as a mole of another substance. • If you are asked to measure one mole of sucrose, what would you do? • If you are asked to measure one mole of ethyl alcohol, C2H6O, what would you do? • To make a liter (L) of solution containing 1 mole of sucrose dissolved in water, what would you do?

25. Molarity • To make a 1-molar (1 M) solution of sucrose, we would measure out 342 g of sucrose and gradually add water, while stirring, until the sugar was completely dissolved. We would then add enough water to bring the total volume of the solution to 1L. • Molarity = the number of moles of solute per liter of solution, is the concentration used most often by biologists for aqueous solutions. • You Try: How would you make a 1 M solution of sodium chloride (NaCl)? How about a .5 M solution?

26. Shifting Hydrogen Atoms • Occasionally, a hydrogen atom in water shifts from one water molecule to another. When this happens, it most often leaves behind its electron. • What is transferred, then, is a hydrogen ion (H+). • The water molecule that lost a proton is now a hydroxide ion (OH-). • The proton binds to another water, making a hydronium ion H3O+.

27. Acids and Bases • The reaction 2H2O H3O+ + OH - is reversible. It reaches a state of dynamic equilibrium when water molecules dissociate at the same rate as they re-form. • The concentration of each ion in pure water is 10-7 . • H+ and OH- are very reactive. Changes in their concentrations can drastically affect a cell’s proteins and other molecules. • Biologists use the pH scale to measure the levels of H+ and OH- in a solution.

28. pH Scale In any aqueous solution, the product of the H+ and OH- ions is 10-14 The concentration of H+ is 10-7 and the concentration of OH- is 10-7. This is known as pH 7.

29. Acids • What would cause an aqueous solution to have an imbalance of H+ and OH- concentrations? • Acid=a substance that increases the H+ concentration • An acidic substance has more H+ ions than pure water. Any pH less than 7 is acidic. Acids that can be eaten have a sour taste. Acids can be corrosive.

30. Bases • Base=A substance that reduces the H+ ion concentration is called a base. • Bases have less H+ ions (and usually have more OH- ions) in solution. • Any solution with pH more than 7 is basic. Bases usually have a bitter taste. In high concentrations, bases can be as corrosive as strong acids.

31. Fig. 3-UN5 0 Acidic [H+] > [OH–] Acids donate H+ in aqueous solutions Neutral [H+] = [OH–] 7 Bases donate OH– or accept H+ in aqueous solutions Basic [H+] < [OH–] 14

32. Buffers • The internal pH of most living cells is close to 7. • Even a slight change in pH can be harmful. • Buffers are substances that minimize changes in the concentrations of H+ and OH- ions. • They do so by accepting H+ ions when they are in excess or donating them when lacking. An important buffer in blood is carbonic acid H2CO3.

33. Water Quality • Considering the dependence of all life on water, contamination of rivers, lakes, seas, and rainwater is a dire environmental problem. • Many threats to water quality have been posed by human activities.

34. Acid Precipitation • Acid precipitation refers to rain, snow, or fog with a pH lower than 5.2. • Uncontaminated rain is about 5.6 • The burning of fossil fuels is a major source of sulfur oxides and nitrous oxides. These combine with water to form strong acids.