Assigning Oxidation Numbers RULES RULES Examples Examples 2Na + Cl2  2NaCl Na = 0 or written Na0 Cl2 = 0 or written Cl20 2Na + Cl2  2NaCl 1. Each Uncombined Element has an oxidation number = 0 Na = 0 or written Na0 Cl2 = 0 or written Cl20 Monatomic ions have an oxidation number equal to the ionic charge. Al3+ then Al = +3 In KCl: K=+1 In CaCl2: Ca=+2 In MgO; Mg=+2 In Li2O: Li=+1 2. The metals of Group 1 always have an oxidation number of +1 and the metals of Group 2 always have an oxidation number of +2. Group 1 = +1 Group 13 = +3 Group 2 = +2 In HF: F = -1 In O2F: F = -1 In CaCl2: Cl = -1 In HClO; oxygen is more electronegative than Cl  Cl=+1 3. Fluorine is ALWAYS = -1 The other halogens are also = -1 when they are the most electronegative element in the compound

Assigning Oxidation Numbers RULES RULES Examples Examples In HF; H=+1 In H2SO4; H=+1 In CaH2: H=-1 In LiH: H=-1 In HF; H=+1 In H2SO4; H=+1 In CaH2: H=-1 In LiH: H=-1 4. Hydrogen is +1 in a compound EXCEPT when it is combined with a METAL. When combined with a METAL, Hydrogen = -1 In H2O: O=-2 In OF2: O=+2 In Na2O2 (sodium peroxide): O=-1 In H2O: O=-2 In OF2: O=+2 In Na2O2 (sodium peroxide): O=-1 5. Oxygen is -2 EXCEPT when with Fluorine, then Oxygen = +2 EXCEPT in peroxide ion (O22-) , then oxygen = -1 6. Most elements will be the charge on the periodic table. Zn = +2 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 In NaCl: (+1) + (-1) = 0 In CaCl2: (+2) + 2(-1) = 0 In Al2(SO4)3: 2(+3)+3(-2) =0 In SO42-: S + 4(-2) = -2 S must = +6

OXidation-REDuction Reactions (also called REDOX) WHAT ARE…. OXidation-REDuction Reactions (also called REDOX) OXIDATION REDUCTION What Happens to the electrons? Electrons are LOST (LEO = loss of electrons=oxidation) Electrons are GAINED GER: Gain of Electrons=Reduction LEO the lion goes GER What Happens to the charge? Charge goes up (The oxidation number on the atom increases) Charge goes down (the oxidation number on the Atom decreases) Write the reaction Include: Balanced # of atoms Balanced charge Ex: Mg+Cl2MgCl2 Oxidation Half reaction is: Mg  Mg2+ + 2e- Ex: Mg+Cl2MgCl2 Reduction Half reaction is: Cl2 + 2e-  2Cl- Ex: Hg2+ + 2I-  Hg + I2 Reduction Half reaction is: Hg2+ + 2e-  Hg Ex: Hg2+ + 2I-  Hg + I2 Oxidation Half reaction is: 2I-  I2 + 2e- 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

OXidation-REDuction Reactions (also called REDOX) WHAT ARE…. OXidation-REDuction Reactions (also called REDOX) OXIDATION REDUCTION Also called the… Reducing agent (the thing that gets oxidized is doing the reducing so..) Oxidizing Agent (the thing that gets reduced is doing the oxidizing so..) ELECTRONS LOST ELECTRONS GAINED EVERYTHING IS ELECTRICALLY NEUTRAL!! ALL CHARGES MUST BALANCE (AND CANCEL)! Identifying redox reactions Assign oxidation numbers to both sides of the equation 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

How to Balance a Redox reaction: Assign oxidation #’s to all of the atoms. Identify which are oxidized & which are reduced. Use one bracketing line to to connect the atoms that undergo oxidation & another to connect those that undergo reduction. Make the total increase in oxidation # equal to the total decrease in oxidation # by using the appropriate coefficient. Make sure that the equation is balanced for both atoms & charges. Ex. +3 -2 +2 -2 0 +4 -2 Fe2O3(s) + CO(g) Fe(s) + CO2(g) -3 reduction 2 x (-3) = -6 3 2 +2 oxidation 3 x (+2) = +6

IS an application of REDOX reactions ELECTROCHEMISTY IS an application of REDOX reactions Energy is produced or applied from REDOX reactions Two different ways ELECTROCHEMICAL ELECTOLYTIC Electricity FORCES a chemical change (apply the electricity) Chemicals PRODUCE electricity (I.e., metals & solutions) HALF CELL /Half Reaction: »Show only 1/2 half of the reaction Either oxidation rxn OR reduction rxn CELL POTENTIALS: »Voltages that are produced by each cell If E0 = 0.00, have DEAD BATTERY (equilibrium) » BOTH mass, energy and charge are conserved!! (equal on both sides) Strip of metal Solution that Contains that Metal ions Al Al3+

ELECTROCHEMISTY Spontaneous Reactions: USE TABLE J TO PREDICT The neutral metal must be higher than the metal ion with which it’s reacting Zn + Cu2+  Cu + Zn2+ On table J, Zn is higher than Cu = spontaneous Cu + Mg2+  Cu2+ + Mg: NOT SPONTANEOUS, Cu not higher than Mg NEEDS BATTERY! Electrons naturally flow from higher metal to lower metal

ELECTROCHEMICAL CELL (aka: Voltaic cells or Galvonic cells) Chemical rxn produces electricity SPONTANEOUS: -G Chemistry electricity + volts and +E0 i.e. making a battery How do we tell if it is spontaneous? Salt Bridge (U-tube) Salt Bridge: ions migrate in both directions Metals as electrodes Zn, Cu Electrons flow from higher metal to lower metal on TABLE J Zn higher than Cu CATHODE: reduction occurs (Cu) Positive (+) ANODE: oxidation occurs (Zn) Negative (-) Zn0 gives up electrons Zn mass Ion Concentration 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-

“Plating” or coating an ELECTROLYTIC CELL (NEEDS A BATTERY) Apply electrical energy to make a chemical rxn happen. Need a battery NOT SPONTANEOUS!! +G (-) volts TWO MAJOR TYPES Electrolysis Separating a compound Into its elements Electroplating “Plating” or coating an Element onto an object EXAMPLE OF THE TWO MAJOR TYPES: BRINE SOLUTION (NaCl): 2NaCl  2Na + Cl2(g) Battery - +  Cl- H+ OH- Na+ H2(g) Cl2(g) Anode Ox. 2Cl-  Cl2 + 2e- Cathode Red. 2H+ +2e- H2 e- e- ELECTROLYSIS: +1 -2 2H2O  2H2 + O2 Anode Oxidation Cathode Red. 2H+ +2e- H2 O-2 O2(g) + 2e- e- e- 

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

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

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

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!!!

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