1. Wonderfully Wacky Water

2. Water as a Polar Molecule
Oxygen is more electronegative than Hydrogen
Oxygen pulls electrons more, so they spend more time with Oxygen giving it a slightly negative charge.
The Hydrogens then become slightly positive because the electrons are kickin’ it with Oxygen more.

3. Happy Hydrogen Bonding
Water molecules are lightly bonded to other water molecules by Hydrogen Bonds!
Because water has charges, the opposite charges attract. One molecules + attracts to another molecules -.
Water can make FOUR Hydrogen Bonds
Hydrogen Bonding is behind all of water’s special properties.

5. Properties of Water
Hydrogen bonds are weak, but they re-form quickly in liquid water giving it more structure than other liquids

6. Cohesion
Water molecules sticking together

7. Adhesion Water molecules cling to other molecules
ex. water clings to vessels of plant to help move it up to the leaves against gravity

8. Figure 3.2x Trees

9. Surface tension How difficult it is to break the surface of a liquid.
At the surface between water and air, hydrogen bonds hold water molecules to each other, and to the molecules below. This makes the surface of water act like it has a special invisible coating.

10. Figure 3.3 Walking on water

11. High Specific Heat
Heat: Total Quantity of Kinetic Energy due to molecular motion of matter.
Temperature: Intensity of heat due to the average kinetic energy of the molecules.
Calorie: Amount of heat energy it takes to raise the temperature of 1 gram of water by 1 C.

12. Specific Heat
Amount of heat energy it takes to raise the temp of 1 g of water by 1 C.
How well a substance resists changing its temp when it absorbs or releases heat.

13. High Specific Heat
Water’s specific heat is 1 cal/g/ C. (high compared to other things)
This means it is harder to change the temperature of water than most other substances.
ex. burn your hand on metal pot, but water is still not hot)

14. Heat of Vaporization Amount of heat a liquid must absorb for
1 g to be converted from liquid to gas (at 100 C).
It takes a lot of heat to convert water from a liquid to a gas.

15. Why?
Hydrogen Bonds!!!

16. Hydrogen Bonds
Bonds are weak, BUT must be broken to raise the temperature of water. Heat breaks the bonds!
Bonds must be broken to change water to vapor!
Heat is re-released when the bonds are formed!

17. High Specific Heat
ex.-When you boil water, you are adding heat to the water to break the bonds and change the water to steam. When you put your hand over the steam, you feel heat and get wet as the bonds are re-formed and the state changes

18. High Specific Heat
ex. -water stabilizes air temp by absorbing heat from warmer air, and releasing heat to cooler air; ocean temps are stable for marine life; organisms resist temp changes in their bodies due to high water content

19. Evaporative Cooling
If a molecule of liquid moves fast enough, it can exit as a gas (evaporate)
When liquid evaporates, it cools the surface because the hottest molecules have escaped as a gas
ex.- sweating allows the surface of your body to cool, prevent overheating on a dry day.
Why wouldn’t it work as well on a humid day?

20. Figure 3.4 Evaporative cooling

21. Ice Floats
Water is less dense as solid than a liquid because of HYDROGEN BONDING!!
As water cools, it locks into a crystal bonded to four other molecules. They are far enough apart to make ice 10% less dense than water
ex.- if ice sank, bodies of water would freeze solid...no life....ugh!

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

23. Figure 3.6x1 Floating ice and the fitness of the environment: ice fishing

24. Water is the solvent of life

25. Terms
Solution - liquid that is a completely homogeneous mixture of two or more substances
Solvent - dissolving agent
Solute - substance being dissolved
Aqueous solution - water is the solvent

26. Water is a very versatile solvent
Ionic compounds like salt dissolve in water by attractions between the anion with hydrogen, and the cation with oxygen. Water surrounds the ions and shields them from each other, breaking the bonds and dissolving the substance

27. Figure 3.7 A crystal of table salt dissolving in water

28. Water is a very versatile solvent
Polar molecules like sugar also dissolve in water because water molecules coat them

29. Hydrophilic substances
Have an affinity for water (water lovers). All things that dissolve in water are hydrophilic (ionic/polar molecules). Other molecules are hydrophilic but don’t dissolve (like cotton)
ex-Towel absorbs water but doesn’t dissolve!!

30. Hydrophobic substances
Don’t have an affinity for water (water fearing). Non-ionic and non-polar substances repel water (like oil).
Ex. Hydrophobic substances are used in cell membranes so that they don’t dissolve!!

31. Solute Concentration in Aqueous Solutions
Molecular weight - sum of weights of all atoms in a molecule
1 mole = molecular weight expressed as grams

32. Figure 3.x2 Moles

33. Practice Carbon = 12 daltons Oxygen = 16 daltons Hydrogen = 1 dalton
C12H22O11 (sucrose)
molecular weight = (12x12)+(16x11)+(1x22) = dalton
1 mole of sucrose = 342 grams

34. You try Ethyl alcohol C2H6O molecular weight =
1 mole of ethyl alcohol =

35. Molarity Molarity (M)=number of moles of solute per liter of solution
If something is 1M = 1 mol/liter
For sucrose - add 342 g and fill to 1 liter with water

36. Dissociation of Water Water is able to dissociate into H+ and OH-
H2O ----> H+ + OH-
Reaction is reversible and at equilibrium H2O is heavily favored (much more water in than ions). In pure water the concentration of H+ and OH- is 10-7M.
Even though this is a very small #, it is very important because H+ and OH- are very reactive and adding or taking away these ions can affect a cell dramatically

37. Acids Substances that add H+ to a solution ex. *HCl ---> H+ + Cl-
Since there are now more H+ ions than OH- ions, it is called an acidic solution

38. Bases Substances that reduce the H+ concentration in a solution.
ex. NaOH ----> Na+ + OH-
The OH- ions will combine with H+ to form water, therefore reducing the H+ concentration and making a basic solution.

39. Neutral Solutions
Have equal concentrations of H+ and OH-

40. pH Scale
A convenient way to express the H+ and OH- concentrations in a solution.
*pH = - log [H+]

41. Figure 3.9 The pH of some aqueous solutions

42. neutral solution pH = -log 10-7 = 7
acidic solution = pH less than 7
basic solution = pH greater than 7

43. Moving 1 number on the pH scale makes a 10x change
ex. going from pH 4 to pH 2 = 100x more acidic
You try - How much more acidic is a solution of pH 4 than one of
pH 8??

44. Buffer
Minimize changes in pH when acids or bases are added. These are especially important in the body because slight changes in pH can make enzymes and proteins inactive and cause many problems.
Buffers either accept H+ ions when there are too many, or donate H+ ions when there are not enough.

45. Example of a Buffer in the Body
ex. carbonic acid in blood regulates
blood pH
If blood pH rises
H2CO3 HCO H+
If blood pH drops

46. Acid Precipitation
Rain, snow or fog more acidic than pH 5.6 caused by sulfur oxides and nitrogen oxides in the air from burning of fossil fuels.
These pollutants combine with water in air and return to earth as rain, snow or fog

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