1. Oxidation-Reduction I The Basics A. Definitions 1. Oxidation Loss of electrons Electrons are products (right side) Oxidation number increases ex. Na 0 (s) → Na + (aq) + e -

2. Oxidation-Reduction I The Basics A. Definitions 2. Reduction Gain of electrons Electrons are reactants (left side) Oxidation number decreases ex. Cl 2 0 (g) + 2e - → 2Cl - (aq)

3. Oxidation-Reduction B. Assigning Oxidation Numbers 1. Free state elements are 0 Free state means unbonded Na 0 Diatomic molecules are “unbonded” Cl 2 0

4. Oxidation-Reduction 2. Bonded elements require the reference table  Make the naming chart Ca Cl 2 Fe 2 O 3 +2 +2 +3 -2 -6+6

5. Oxidation-Reduction 3. Hydrogen is always +1 H 2 +1 O -2 except  When bonded to a metal (known as a hydride) Na +1 H -1

6. Oxidation-Reduction 4. Oxygen is always -2 H 2 +1 O -2 except  When bonded to fluorine O +2 F 2 -1  In a peroxide H 2 +1 O 2 -1

7. Practice finding oxidation numbers: H 3 PO 4 KNO 3 Na 2 Cr 2 O 7 NH 3 MgCl 2 CuCO 3 PbSO 4 Sn(NO 3 ) 2

8. II Half-Reactions  When writing ½-reactions 1. Assign oxidation numbers for every element 2. Find elements that change oxidation state and write without using any coefficients or subscripts  If zero charge occurs, see if element is diatomic If diatomic, add subscript to “0” charged substance and add coefficient to the “charged” substance 3. Label as oxidized or reduced 4. Balance charge using electrons

9. H 2 S + Ca  H 2 + CaS H  H Ca  Ca +1-20+20-2 +10 0+2 2 2 oxidation reduction + 2e

10. III Electrochemical Cells  An electrochemical cell combines oxidation and reduction half-reactions Two Types  Voltaic cell (battery)  Electrolytic cell

11. A. Voltaic Cell Reaction that produces electricity  Battery  Exothermic reaction  Spontaneous reaction Oxidation site  Anode (-) Reduction site  Cathode (+)

12.  Parts of Cell Electrode  Site where reaction occurs Wire  Pathway for electrons Salt Bridge  Pathway for ions + ions move to the cathode - ions move to the anode

13. B. Electrolytic Cell Reaction that requires electricity  Endothermic reaction  Nonspontaneous reaction Oxidation site  Anode (+) Reduction site  Cathode (-)

14. C. Table J Shows relative oxidation (metals) and reduction (nonmetals)  Higher metal oxidizes the other metal reduces  Higher nonmetal reduces the other nonmetal oxidizes Example  React Barium and Magnesium Ba and Ba 2+ to Mg and Mg 2+ Ba 0 → Ba 2+ + 2e Mg 2+ + 2e → Mg 0