1. Chapter 3: Water and the Fitness of the Environment

2. Water

3. Water A common but exceptional substance
Makes life as we know it possible on Earth
Organisms composed mainly of water (70-95%) and live in an environment dominated by water (75% of surface)
Life on Earth began in water and evolved there for 3 billion years before spreading onto land
Water is the only common substance to exist in all 3 physical states (solid, liquid, gas)

4. Figure 3.5x1 Ice, water, and steam

5. Hydrogen Bonding/Polarity
Polar covalent bonds between oxygen and hydrogen due to oxygen’s high electronegativity
Note the bent shape of a water molecule
Anomalous properties of water arise from attractions between polar molecules
CD-ROM Activity 3A

6. Figure 3.1 Hydrogen bonds between water molecules

7. Important Properties of Water
Cohesive behavior
Ability to stabilize temperature
Expansion upon freezing
Versatility as a solvent

8. Cohesion Cohesion Adhesion Surface Tension
Due to hydrogen bonding holding individual molecules together
Hydrogen bonds are made and broken with great frequency
Contributes to the transport of water against gravity in a plant stem
Adhesion
Clinging of one substance to another
Surface Tension
A measure of how difficult it is to stretch or break the surface of a liquid

9. Figure 3.3 Walking on water

10. Stabilization of Temperature
Kinetic Energy - the energy of motion
Heat - the total kinetic energy of a sample
Temperature - the average kinetic energy of a sample
Heat passes from warmer to cooler bodies until temperatures are equal
Note an ice cube works by absorbing heat, not by adding coldness!

11. Stabilization of Temperature
calorie – the amount of energy needed to increase the temperature of 1 gram of water 1 degree centigrade
1000 calories = 1 kilocalorie = 1 Calorie
Specific Heat - the amount of energy needed to increase the temperature of 1 gram of a substance 1 degree centigrade

12. Stabilization of Temperature
Water has a very high specific heat compared to most other substances
Reason = hydrogen bonding
Relevance to life on Earth – a large body of water can absorb and store a huge amount of heat from the sun during the daytime while only warming a few degrees; at night, the gradually cooling water can warm the air.

13. Evaporative cooling

14. Evaporative Cooling
Heat of vaporization – the quantity of heat 1 gram of a liquid must absorb to be converted from a liquid to a gas
Evaporative cooling – you are decreasing the average temperature because the hottest molecules are leaving

15. Ice floats and frozen benzene sinks

16. Expansion Upon Freezing
Ice is less dense than water due to packing caused by hydrogen bonding
This keeps ice from sinking and bodies of water from freezing completely

17. Figure 3.5 The structure of ice (Layer 2)

18. Figure 3.6 Floating ice and the fitness of the environment

19. Versatility as a Solvent
Solution
A liquid that is a completely homogeneous mixture of two or more substances
Solute
The substance that is dissolved in a solution
Solvent
The dissolving agent of the solution
Aqueous solution
A solution in which water is the solvent

20. Versatility as a Solvent
Sphere of hydration formed around dissolved particles
Water is a versatile solvent, but not a universal solvent!
Hydrophilic = water-loving
Hydrophobic = lipid-loving
Note: A compound does not have to be ionic to dissolve in water, it must be polar
Molarity = moles of solute/liter of solution

21. A crystal of table salt dissolving in water

22. Dissociation of Water Molecules
2 H2O ↔ H3O OH-
hydronium hydroxide
ion ion
Although the dissociation of water is reversible and statistically rare (1 in 554 million), it is exceedingly important in the chemistry of life.

23. Ionization of water

24. Dissociation of Water Molecules
H2O ↔ H OH-
hydrogen hydroxide
ion ion
Water can be considered a very weak acid
[H+] = 1 x 10-7M [OH-] = 1 x 10-7M

25. Acids HCl → H+ + Cl- H2CO3 ↔ H+ + HCO3-
Acid = a substance that increases the hydrogen ion concentration of a solution
HCl → H Cl-
HCl is an example of a strong acid. It dissociates completely
H2CO3 ↔ H HCO3-
Carbonic acid is a weak acid that reversibly releases and reaccepts hydrogen ions

26. Bases
Base = a substance that reduces the hydrogen ion concentration of a solution
Some bases do this by accepting a hydrogen atom NH3 + H+ ↔ NH4+
Some bases do this by dissociating to form hydroxide ions NaOH → Na+ + OH-

27. pH For any solution: [H+] x [OH-] = 10-14
pH = -log [H+] Each pH unit represents a 10-fold
difference in [H+] and [OH-] !!!!!!!!
[H+] > [OH-] solution is acidic pH < 7
[H+] = [OH-] solution is neutral pH = 7
[H+] < [OH-] solution is basic pH > 7

28. Figure 3.9 The pH of some aqueous solutions

29. pH of Biological Systems
Most biological fluids are within the range of pH 6 to 8. An exception is the digestive juices of the stomach at pH 2.
In biological systems, slight changes in pH can be harmful. Chemical processes within the body are very sensitive to the concentrations of H+ and OH–
Ex: Blood pH is normally pH A person can survive only minutes if pH drops to 7 or is raised to 7.8
Therefore, biological systems rely on buffers to resist changes in pH.

30. Buffers
Buffers are substances that minimize the changes in concentrations of H+ and OH–
Buffers work by accepting H+ from the solution when they are in excess and donating H+ to the solution when they have been depleted.
Most buffers are composed of a weak acid and its corresponding base.
H2CO3 ↔ H HCO3-
(H+ donor) (H+ acceptor)
acid base
Le Chatelier’s Principle

31. Acid Precipitation Rain, snow or fog with a pH below 5.6
Results from a reaction between water vapor and sulfur oxides or nitrogen oxides produced by the combustion of fossil fuels
Destructive to aquatic systems such as ponds/May also cause serious harm to plants due to changes in soil pH

32. Figure 3.10 The effects of acid precipitation on a forest

33. Figure 3.10x2 Acid rain damage to statuary, 1908 & 1968