1. Oxidation-Reduction (Redox) Reactions
In oxidation-reduction (abbreviated as “redox”) reactions, electrons are transferred from one reactant to another.
Oxidation I
Lose electrons
Reduction
Gain electrons

2. Redox Reactions In the reaction between Na and Cl2:
Na lost an electron, it has been oxidized
electron (e-)
Cl- Cl
Cl gained an electron, it has been reduced
2 Na (s) + Cl2 (g)  2 NaCl (s)

3. Redox Reactions What about the reaction between Al and O2?
O gained two electrons, it has been reduced
O2-
electrons (e-) Al
Al lost 3 electrons, it has been oxidized
Al3+
Al (s) + O2 (g)  Al2O3 (s)
4 Al (s) + 3 O2 (g)  2 Al2O3 (s)

4. Oxidation Numbers
Oxidation Number (or Oxidation State): actual or hypothetical charge of an atom in a compound if it existed as a monatomic ion
Common Oxidation Numbers:
H+ = +1 Cl- = -1
O2- = -2 Al = 0
Na = 0 Na+ = +1
Oxidation numbers can also be assigned to atoms with in a more complex molecule.

5. Assigning Oxidation Numbers
1. The oxidation number of an element in its natural form is 0. Examples: the oxidation number is zero for each element in H2, O2, Cl2, P4, Na, etc.
2. The oxidation number of a monatomic ion is the charge on the ion. Examples: Na3N, the ions are Na+ and N3–, so oxidation #’s: Na = +1 and N = -3. In Al2O3, the ions are Al+3 and O2–, so oxidation #’s: Al = +3 and O = -2
3. In a compound or polyatomic ion,
Group I elements are always +1.
Group II elements are always +2.
Fluorine is always -1.
Oxygen is usually -2 (except in the peroxide ion, O22–, when O is -1)
Hydrogen is usually +1 (except when it is with a metal, like NaH or CaH2, then it is -1)
4. In a neutral compound, the sum of all oxidation numbers must equal 0. In a polyatomic ion, the sum of all oxidation numbers must equal the charge.

6. Assigning Oxidation Numbers
Examples: Determine the oxidation number for each element in the following:
CrO42–: Cr: ____, O: ____
H2SO4: H: ____, S: ____, O: ____
NO3-: N: ____, O: ____
CaCr2O7: Ca: ____, Cr: ____, O: ____
C2O42–: C: ____, O: ____
C3H8: C: ______________, H: ____

7. Redox Reactions In a redox reaction:
One reactant Loses Electrons/is Oxidized (LEO)
Another reactant Gains Electrons/is Reduced (GER)
An easy way to remember is “LEO the lion goes GER!” (Though I prefer OIL RIG, it’s your choice).
The element or reactant that is oxidized is the reducing agent.
The element or reactant that is reduced is the oxidizing agent.

8. Examples Zn(s) + AgNO3(aq)  Zn(NO3)2(aq) + Ag(s)
Al(s) HCl(aq)  AlCl3(aq) H2(g)
C2H2(g) O2(g)  CO2(g) H2O(g)

9. Examples Ca(s) + H2O(l)  Ca(OH)2(aq) + H2(g)
H2O2(aq) Mn(OH)2(aq)  Mn(OH)3(aq)

10. Solution Concentration
solution: homogeneous mixture of substances present as atoms, ions, and/or molecules
  solute: component present in smaller amount
solvent: component present in greater amount
Note: Unless otherwise stated, the solvent for most solutions considered in this class will almost always be water!
Aqueous solutions are solutions in which water is the solvent.

11. How do we measure concentration?
A concentrated solution has a large quantity of solute present for a given amount of solution.
A dilute solution has a small quantity of solute present for a given amount of solution.
SOLUTION CONCENTRATION = amount of solute amount of solvent
The more solute in a given amount of solution  the more concentrated the solution
Example: Explain the difference between the density of pure ethanol and the concentration of an ethanol solution.

12. How do we measure concentration?
Concentration can be measured a number of ways:
ppm (parts per million) – one part in a million parts
ppb (parts per billion) – one part in a billion parts
g/kg (grams per kilogram) – one gram solute per one kilogram of solvent
The chemical standard most used is Molarity
Molarity = moles of solute liters of solution
units: M (molar = mol/L)

13. Solution Concentration
Find the molarity of a solution prepared by dissolving 1.25 g of KOH in mL of solution.
Find the molarity of a solution prepared by dissolving 5.00 g of copper(II) sulfate in mL of solution

14. Ion Concentrations
When an ionic compound is dissolved in water, the concentration on the individual ions is based on their molecular formula…
For example:
1 M NaCl solution contains 1 M Na+ and 1 M Cl-
2 M NaCl solution contains 2 M Na+ and 2 M Cl-
1 M CaCl2 solutions contains 1 M Ca2+ and 2 M Cl-
2 M CaCl2 solutions contains 2 M Ca2+ and 4 M Cl-

15. Solution Concentration
Indicate the concentration of barium and chloride ions in a 1.00M barium chloride solution.
Indicate the molarity of each ion in the solutions indicated below:
In a 0.125M Na2SO4(aq) solution
[Na+]=____________ and [SO42-]=____________.
b. In a 0.500M Fe(NO3)3(aq) solution
[Fe3+]=____________ and [NO3–]=___________.
c. In a 1.250M Al2(SO4)3(aq) solution
[Al+3]=____________ and [SO42-]=___________.

16. Solving Concentration Problems
Keep in mind that if molarity and volume are both given, you can calculate # of moles since:
volume  molarity = volume (in L)  moles of solute
liters of solution
so volume units will cancel  # of moles!
If you are given volume and molarity for a solution, multiply them together to get # of moles!

17. Solving Concentration Problems
Calculate the mass of NaCl needed to make 1.00 L of a 1.00 M solution.

18. Preparing Solutions

19. Examples
Calculate the mass of barium hydroxide required to make mL of a 0.500M barium hydroxide solution. What volume (in mL) of a 0.125M silver nitrate solution contains 5.00 g of silver nitrate?

20. Examples
Calculate the molarity of hydroxide ion in a solution prepared by diluting 50.0 mL of 1.50M potassium hydroxide with mL of 0.500M calcium hydroxide.