1. Assigning Oxidation Numbers RULESExamples 2Na + Cl 2  2NaCl Na = 0 or written Na 0 Cl 2 = 0 or written Cl 2 0 RULESExamples 1. Each Uncombined Element has an oxidation number = 0 2Na + Cl 2  2NaCl Na = 0 or written Na 0 Cl 2 = 0 or written Cl 2 0 Monatomic ions have an oxidation number equal to the ionic charge. Al 3+ then Al = +3 2. The metals of Group 1 always have and oxidation number of +1 and the metals of Group 2 always of an oxidation number of +2. Group 1 = +1 Group 2 = +2 In KCl: K=+1 In CaCl 2 : Ca=+2 In MgO; Mg=+2 In Li 2 O: Li=+1 3. Fluorine is ALWAYS = -1 The other halogens are also = -1 when they are the most electronegative element in the compound In HF: F = -1 In O 2 F: F = -1 In CaCl 2 : Cl = -1 In HClO; oxygen is more electronegative than Cl  Cl=+1 Group 13 = +3

2. Assigning Oxidation Numbers RULESExamples In HF; H=+1 In H 2 SO 4 ; H=+1 In CaH 2 : H=-1 In LiH: H=-1 In H 2 O: O=-2 In OF 2 : O=+2 In Na 2 O 2 (sodium peroxide): O=-1 RULESExamples 7.The sum of the oxidation numbers in all compounds must be ZERO 7a. The sum of the oxidation numbers in polyatiomic ions must equal the charge on the ion 4. Hydrogen is +1 in a compound EXCEPT when it is combined with a METAL. When combined with a METAL, Hydrogen = -1 In HF; H=+1 In H 2 SO 4 ; H=+1 In CaH 2 : H=-1 In LiH: H=-1 5. Oxygen is -2 EXCEPT when with Fluorine, then Oxygen = +2 EXCEPT in peroxide ion (O 2 2- ), then oxygen = -1 In H 2 O: O=-2 In OF 2 : O=+2 In Na 2 O 2 (sodium peroxide): O=-1 In NaCl: (+1) + (-1) = 0 In CaCl 2 : (+2) + 2(-1) = 0 In Al 2 (SO 4 ) 3 : 2(+3)+3(-2) =0 In SO 4 2- : S + 4(-2) = -2 S must = +6 6. Most elements will be the charge on the periodic table. Zn = +2

3. OXidation-REDuction Reactions (also called REDOX) WHAT ARE…. OXIDATION REDUCTION Electrons are LOST (LEO = loss of electrons=oxidation) Electrons are GAINED GER: Gain of Electrons=Reduction Charge goes up (The oxidation number on the atom increases) Charge goes down (the oxidation number on the Atom decreases) Ex: Mg+Cl 2  MgCl 2 Oxidation Half reaction is: Mg  Mg 2+ + 2e - Ex: Mg+Cl 2  MgCl 2 Reduction Half reaction is: Cl 2 + 2e -  2Cl - What Happens to the electrons? What Happens to the charge? Write the reaction Include: Balanced # of atoms Balanced charge Ex: Hg 2+ + 2I -  Hg + I 2 Oxidation Half reaction is: 2I -  I 2 + 2e - Ex: Hg 2+ + 2I -  Hg + I 2 Reduction Half reaction is: Hg 2+ + 2e -  Hg What happens to the electrons if you add the two half reactions? NOTE: e - are on right side of arrow NOTE: e - are on left side of arrow LEO the lion goes GER

4. OXidation-REDuction Reactions (also called REDOX) WHAT ARE…. OXIDATION REDUCTION Reducing agent (the thing that gets oxidized is doing the reducing so..) Oxidizing Agent (the thing that gets reduced is doing the oxidizing so..) Also called the… ELECTRONS LOST ELECTRONS GAINED EVERYTHING IS ELECTRICALLY NEUTRAL!! ALL CHARGES MUST BALANCE (AND CANCEL)! Identifying redox reactions 1.Assign oxidation numbers to both sides of the equation 2.Look to see if oxidation numbers change. If they do = redox (if not then not) HINTS: Single replacement reactions are REDOX Double replacement are NOT redox

5. How to Balance a Redox reaction: 1.Assign oxidation #’s to all of the atoms. 2.Identify which are oxidized & which are reduced. 3.Use one bracketing line to to connect the atoms that undergo oxidation & another to connect those that undergo reduction. 4.Make the total increase in oxidation # equal to the total decrease in oxidation # by using the appropriate coefficient. 5.Make sure that the equation is balanced for both atoms & charges. Ex. +3 -2 +2 -2 0 +4 -2 Fe 2 O 3 (s) + CO(g) Fe(s) + CO 2 (g) 3 x (+2) = +6 -3 reduction +2 oxidation 2 x (-3) = -6 3 32

6. IS an application of REDOX reactions Energy is produced or applied from REDOX reactions ELECTROCHEMISTY Two different ways ELECTROCHEMICAL Chemicals PRODUCE electricity (I.e., metals & solutions) ELECTOLYTIC Electricity FORCES a chemical change (apply the electricity) HALF CELL /Half Reaction: »Show only 1/2 half of the reaction Either oxidation rxn OR reduction rxn » BOTH mass and charge are conserved!! (equal on both sides) Strip of metal Solution that Contains that Metal ions Al Al 3+ CELL POTENTIALS: »Voltages that are produced by each cell If E 0 = 0.00, have DEAD BATTERY (equilibrium)

7. ELECTROCHEMISTY Spontaneous Reactions: USE TABLE J TO PREDICT The neutral metal must be higher than the metal ion with which it’s reacting Zn + Cu 2+  Cu + Zn 2+ On table J, Zn is higher than Cu = spontaneous Cu + Mg 2+  Cu 2+ + Mg: NOT SPONTANEOUS, Cu not higher than Mg NEEDS BATTERY! Electrons naturally flow from higher metal to lower metal

8. Salt Bridge (U-tube) ELECTROCHEMICAL CELL (aka: Voltaic cells or Galvonic cells) Chemical rxn produces electricity Chemistryelectricity i.e. making a battery SPONTANEOUS: -  G + volts and +E 0 How do we tell if it is spontaneous? Metals as electrodes Zn, Cu Electrons flow from higher metal to lower metal on TABLE J ANODE: oxidation occurs (Zn) Negative (-) Zn 0 gives up electrons Zn mass Ion Concentration Salt Bridge: ions migrate in both directions CATHODE: reduction occurs (Cu) Positive (+) Cu +2 gains electrons Mass of Cu Ion Concentration AN OX Anode = oxidation Decrease mass /Decrease e - RED CAT Reduction = cathode Increase mass/increase e - Zn higher than Cu

9. ELECTROLYTIC CELL (NEEDS A BATTERY) Apply electrical energy to make a chemical rxn happen. Need a battery NOT SPONTANEOUS!! +  G(-) volts Electrolysis Separating a compound Into its elements Electroplating “Plating” or coating an Element onto an object TWO MAJOR TYPES EXAMPLE OF THE TWO MAJOR TYPES: ELECTROLYSIS: 2H 2 O  2H 2 + O 2 +1 -2 0 0 Anode Oxidation O -2  O 2(g) + 2e - Cathode Red. 2H + +2e -  H 2 e-e- e-e-   BRINE SOLUTION (NaCl): 2NaCl  2Na + Cl 2(g) Battery - + - +  Cl - H +  OH - Na +  H 2(g)  Cl 2(g) Anode Ox. 2Cl -  Cl 2 + 2e - Cathode Red. 2H + +2e -  H 2 e-e- e-e- e-e-

10. FOR ALL ELECTROLYTIC CELLS AN OX Oxidation @ Anode Pos. term. (+) ELECTROLYTIC CELL (NEEDS A BATTERY) ELECTROPLATING (2 nd type) Object to be plated- Always cathode (reduction) Always attached to negative term. Metal used to plate- Always anode (oxidation) Always attached to positive term. Reduction @ Cathode RED CAT Neg. term (-)

11. Similarities & Differences: Cu + Pb +2 ==> not spontaneous. Zn+ Ag+ ==> Spontaneous! Electrochemical/ Voltaic Cell Electrolytic Cell Chemicals producing electricity- EXO (IS Spontaneous) + volts Table J Metal is Higher than the ion (compound) Summary: What’s going on? Electricity forcing a chemical reaction to occur Endo (NOT spontaneous)

12. Cathode + lower on Table J Anode - Higher on Table J Cathode -Attached to - Anode + Attached To + Charge: Positive or Negative? Reduction Oxidation ReductionOxidation An Ox & Red Cat IS a Battery! NEED a Battery! IS a battery or NEEDS a battery?

13. e- flow anode cathode Salt bridge- allows IONS to migrate in both directions e- flow out- in + opposites attract object to be plated is ALWAYS attached to the negative!!!

14. Also known as Voltaic and Galvanic Spontaneous- produces electricity Electron flow predicted by Table J. --From higher to lower metal. Anode is the (-) terminal Cathode is the (+) terminal Makes a battery NOT Spontaneous Needs a battery to force e- to flow in wanted direction Anode is connected to (+) terminal of battery and is NOW (+) electrode Cathode is connected to (-) terminal of battery and is now (-) electrode Used for : Electrolysis-separating compounds into elements Electroplating – coating items Both involve the transfer and movement of e- to make something happen Oxidation occurs at the anode “An Ox” Reduction occurs at the cathode “Red Cat”