1. Properties of Water
Section 2.2

2. What do you know about water?

3. Unique Properties of Water:
Most abundant compound in nearly all living things

4. Unique Properties of Water:
Density of solid water vs. liquid water
Expansion on freezing – water forms a crystalline structure that expands and is less dense than its liquid state

5. Unique Properties of Water:
Covers more than 75% of the Earth’s surface
Water 100°C or 212°F
Water 0°C or 32°F
Universal solvent

6. Unique Properties of Water:
Water has a high heat capacity
Heat capacity = the amount of heat energy required to increase its temperature, is relatively high.
Large bodies of water (oceans and lakes) can absorb large amounts of heat with only small changes in temperature.
This protects organisms living within from drastic changes in temperature.

7. What is a polar molecule?
An uneven distribution of electrons
The molecule becomes charged
on each end
It means that the electrons are
NOT shared equally

8. Hydrogen Bonding
Because of their partial positive and negative charges, polar molecules such as water can attract each other.
The attraction between a hydrogen atom on one water molecule and the oxygen atom on another is known as a hydrogen bond.

9. Interaction Between Water Molecules
Negative Oxygen atomof one water molecule is attracted to the Positive Hydrogen atom of another water molecule to form a HYDROGEN BOND

11. Cohesion
Cohesion is when one water molecule sticks to another water molecule  due to hydrogen bonding
It creates surface tension (a measure of strength) at the surface of the liquid
Produces a surface “film” that allows small insects to “walk on water”

12. Surface Tension
A measure of how difficult it is to break the surface of a liquid because of hydrogen bonds.
The surface tension of a liquid results from an imbalance of intermolecular attractive forces, (cohesive forces) between molecules
A molecule in the bulk liquid experiences cohesive forces with other molecules in ALL directions.
A molecule at the surface of a liquid experiences only NET INWARD cohesive forces.

13. Adhesion Adhesion is when a water molecule sticks to another substance
(NOT another water molecule)
A piece of tape adheres to the wall
Capillary action is an example of adhesion (water moving from the roots, through the stem, & into the leaves of a flower)

14. Capillary Action
Water climbing up a small tube because of the adhesion to the sides and the cohesion between water molecules.

15. Adhesion Also Causes Water to …
Attach to a silken spider web
hold onto plant leaves

16. Cohesion and Adhesion
Mercury does not wet glass - the cohesive forces within the drops are stronger than the adhesive forces between the drops and glass. When liquid mercury is confined in a tube, its surface (meniscus) has a convex shape because the cohesive forces in liquid mercury tend to draw it into a drop.
When liquid water is confined in a tube, its surface (meniscus) has a concave shape because water adheres to the surface and creeps up the side.

17. High Heat Capacity
Water’s heat capacity, the amount of heat energy required to increase its temperature, is relatively high.
Because of the multiple hydrogen bonds between water molecules, it takes a large amount of heat energy to cause those molecules to move faster and raise the temperature of the water.
Large bodies of water, such as oceans and lakes, can absorb large amounts of heat with only small changes in temperature. This protects organisms living within from drastic changes in temperature.

18. Universal Solvent Hydrophilic Hydrophobic
A substance that is attracted to water (a polar molecule)
Examples: sugar, alcohol
Hydrophobic
A substance that repels (is afraid of) water (a nonpolar molecule)
Examples: oil
The Chemistry of Water

19. Solutions and Suspensions
Mixtures
Solutions and Suspensions

20. What is a Mixture?
Substance composed of two or more elements or compounds that are mixed together, but NOT chemically combined.
Mixture Examples: salt & pepper; sugar & sand; salt & sand; Earth’s atmosphere; soil

21. Solutions and Suspensions
Mixture – a material that consists of 2 or more elements or compounds physically mixed together but are not chemically combined.

22. Homogeneous Mixture (Solution)
Homogeneous mixture in which one substance is dissolved in another
Examples: sugar in water; salt in water; hot chocolate; tea, Kool-aid
Salt Dissolving in Water

23. Parts of a Solution All solutions contain a solute and a solvent.
Solvent is the substance doing the dissolving.
Example: water
Solute is the substance that is being dissolved.
Example: salt

24. Heterogeneous Mixture (Suspensions)
Mixture of water and non-dissolved materials (heterogeneous)
Examples: blood; oil and water

25. Application: 1 – Choose one of your examples of a
solution. Explain how the parts
(components of the solution) can be
separated.
2 – For the same example, identify the
solute and solvent in the solution.
Explain your reasoning.

26. Example of a Solution: Water is the universal solvent.
The solute (sugar) dissolves in the solvent (water).
The solute (Kool-aid) dissolves in the solvent (water).
The solute (oxygen) dissolves in the solvent (nitrogen).
Water is the universal solvent.

27. Homeostasis a. Makes a good insulator b. Resists temperature change
Ability to maintain a steady state despite changing conditions
Water is important to this process because…
a. Makes a good insulator
b. Resists temperature change
c. Universal solvent
d. Coolant
e. Ice protects against temperature extremes (insulates frozen lakes) 4

28. Acids, Bases, pH, & Buffers
All about the pH Scale

29. The Water Molecule:
Water dissociation animation

30. Acids Compounds that release H+ into solution H+ = hydrogen ion
If [H+] > [OH-] then the
solution is acidic.
HCl  H+ + Cl-

31. Bases If [OH-] > [H+] then the solution is basic
Compounds that release
OH- into solution
OH- = hydroxide ion
If [OH-] > [H+] then the
solution is basic
NaOH  Na+ + OH-

32. The pH scale (power of Hydrogen) as in Hydrogen ions
Measures the relative concentration of H+ and OH-

33. Interpreting the pH Scale
pH is based on a logarithmic scale
A pH of 5 is 10× more concentrated than a pH of 6 (in terms of H+ concentration)
A pH of 9 is 100× more concentrated than a pH of 11 (in terms of H+ concentration)

34. Classifying Acids & Bases
Strong acids have a pH value between 1-3
Weak acids have a pH value between 4-6
Bases:
Strong bases have a pH value between 12-14
Weak bases have a pH value between 8-11

35. Buffers
Weak acids or bases that can react with strong acids or bases to prevent sharp sudden changes in pH
Produced naturally by the body to maintain homeostasis
The pH value in most cells is
The pH of stomach acid is 2
The pH of the blood is 7.4

36. Blood Buffer – pH needs to be 7.4
Double arrows mean
reactions are in
equilibrium
bicarbonate ion (buffer)
(Kidney – carried
away in urine)
Free hydrogen ions
Carbonic Acid
(Lungs – carried
away as CO2)

37. Blood Buffer – pH needs to be 7.4
You can see that if the reactions go to the right, hydrogen ions are released, making the solution more acidic, and if the reactions go to the left, hydrogen ions are sucked up, making the solution more basic.  The question that is probably eating away at you right now is:  "how do I know which way the reactions are going to go-- to the right or to the left?“
Good question!  The bicarbonate ion isn't a very strong acid or base.  It doesn't have to go one way or the other all the time.   Instead, the direction it goes depends on the solution it is in.  You see, if it is in an acidic solution, there are lots of hydrogen ions floating around; in this situation, the presence of tons of hydrogen ions will force the reaction to go to the left.  Do you see why?  Look at the reaction again.  All you need for the reaction to be able to go to the left is available hydrogen ions (see the H+ in the middle that is needed for the reaction to go to the left?).  So if bicarbonate ions find themselves in acidic solutions, they tend to act like the weak base they can be, and suck up the excess hydrogen ions.
Consider what will happen if bicarbonate is put into a basic solution. A basic solution has very few hydrogen ions floating around.  In this condition, the bicarbonate reaction cannot proceed to the left, since no hydrogen ions are available.  Instead, it will go to the right, producing hydrogen ions.  In this manner, bicarbonate ions will tend to act like the weak acid they can be, and release hydrogen ions.

38. Use the structure of a water molecule to explain why it is polar.
Why is water such a good solvent?
What is the difference between a solution and a suspension?
How do hydrogen bonds between water molecules occur?
How does the body counteract a drop in blood pH during strenuous exercise?
If a strong acid is dissolved in pure water will the pH of the solution be greater or less than 7?

39. Water Chemistry
Please draw the structure of a water molecule (H2O).
Label the charges on the molecule that make it polar (negative and positive ends).
See if you can connect the water molecule that you drew to two other molecules of water using hydrogen bonds.
Can you think of any unique properties that water displays?

40. Water Chemistry
Explain why water is “polar” in your own words.
Explain capillary action in your own words.
Explain surface tension in your own words.

41. According to the scale, what substance is neutral?
What solution has an [H+] of 0.01?
Which substance produces the most OH- ions in solution?
How many more times basic is pH 14 compared to pH 11.