1. The Wonder of
Water

3. Learning Objectives
Describe the distribution of water on Earth and identify sources of drinking water supply
Explain the important physical properties of water in terms of its molecular structure and bonding behavior
Calculate chemical concentrations of dissolved substances (e.g. percent, ppm, molarity)
Distinguish between ionic and covalent compounds
Predict the ionic charge on atomic ions
Predict the chemical formulas of ionic compounds
Identify covalent compounds that are likely to dissolve in water

4. Distribution of Water on Earth

5. Sources of Drinking Water
US:
Virginia:
EPA Public Water System Inventory, FY2000

6. What Water Do You Drink?

7. Water and Heat
Specific heat: 1 calorie = energy to heat 1 gram of water by 1o C = J/g oC
Heat of fusion: heat that must be absorbed to cause melting = 331 J/g for H2O
Heat of vaporization: heat that must be absorbed to change liquid into vapor = 2250 J/g for H2O

8. Electronegativity
measure of an atom’s attraction for electrons in bonds
O atom is more electronegative than H atom…
…so the “shared” electrons tend to hang out there!

9. Charge Density Model of Water Molecule
Since water is bent, this means that one “end” is more negative, and the other more positive.
That is, water is polar.
This in turn makes water molecules bond to each other in a very special way.

12. Structure of Ice
Lots of open spaces!! (decrease density)

13. ↑
~4 ºC

14. Natural Phenomenon: Lake Turnover

15. Water as a Solvent
solvent = a substance that dissolves other substances
solute = substance that is dissolved
solution = a homogenous mixture (if water is the solvent, it’s called an aqueous solution)
electrolytes = solutes that conduct electricity

16. Solute Concentration in Aqueous Solutions
percent
ppm and ppb
molarity (moles/L)

17. Example: Calculating ppm
If 1 liter of seawater contains 35 grams of dissolved NaCl, what is the concentration of NaCl in ppm?
In moles/liter?
Density of water is 1g/ml, or 1000g/L
(35 g/L)*(1000 mg/1 g)*(1 L/106 mg) = 35,000 mg/106 mg
35,000 mg/106 mg = 35,000 ppm
(35 g/L)*(1 mole/58.4 g) = M NaCl

18. Ionic vs. Covalent Compounds
ions = charged atoms or molecules (e.g. Na+)
ionic compounds = solid crystalline substances made of anions and cations (e.g. table salt)
polar compounds = contain covalent bonds with electrons that are not shared equally (e.g. water)

19. Ionic or Covalent? Well, a little of both…
The more electronegativity in the compound, the more likely the compound will dissociate into separate ions
I.e. one atom “borrows” one or more electrons from the other
Even water dissociates this way - there is a percentage of H+ and OH- ions in every glass of water you drink.

20. Sodium Chloride Dissolving in Water

21. Predicting Charge on Ions
(Hint: Look at valence electrons…) Na Cl Mg Li O Al
Na+
Cl-
Mg2+
Li+
O2-
Al3+

23. Polyatomic Ions

24. Formulae for Ionic Compounds
Mg + O
Mg + Cl
Al + O
magnesium sulfate
calcium phosphate
Mg2+ & O2- → MgO
Mg2+ & 2(Cl-) → MgCl2
2(Al3+) & 3(O2-) → Al2O3
Mg2+ & SO42- → MgSO4
3(Ca2+) & 2(PO43-) → Ca3(PO4)2