UNIT 10: REDOX How can we assign oxidation numbers? How can we recognize a RedOx reaction? How can we identify which species is oxidized/reduced? How can we write half reactions? How can we balance equations using the half reaction method? How can we balance RedOx reactions in acidic and basic solutions? What are free radicals? What are electrode potentials? How can we determine if a redox reaction will occur using Table J? What are the differences and similarities between voltaic and electrolytic cell? -Honors Chemistry Ms. Argenzio

AIM: How can we assign oxidation numbers? Oxidation is defined as the loss of electrons (the charge is increased). Reduction is defined as the gain of electrons (the charge is reduced). Helpful hint LEO says GER LEO – Loss of Electrons is Oxidation GER – Gain of Electrons is Reduction

AIM: How can we assign oxidation numbers? Oxidation Numbers: - all atoms are assigned oxidation numbers (states) - can be positive (metal in a compound), negative (nonmetal in a compound), or neutral (single atom) - identify how many electrons are either gained (nonmetals in a compound) or lost (metals in acompound) by an atom or ion in a reaction

Rules for assigning oxidation numbers: 1. Any uncombined element has an oxidation state of zero 2. Monatomic = atomic charge Na + Cl - 3. Group 1 metals = + 1 in compounds Group 2 metals = + 2 in compounds 4. Fluorine is - 1 in compounds 5. Halogens (Group 17) are - 1 in compounds if they’re the most electronegative element 6. Hydrogen is + 1 in compounds but is - 1 when combined with a metal

Rules for assigning oxidation numbers: 7. Oxygen is - 2 in compounds Oxygen is + 2 with fluorine (OF 2 ) since fluorine is more electronegative Oxygen is - 1 in H 2 O 2 (hydrogen peroxide) 8. The sum of all oxidation states in compounds is zero 9. The sum of all oxidation states in polyatomic ions is the charge of the polyatomic ion on the reference table (Table E)

CONCLUSION QUESTIONS! DETERMINE THE OXIDATION NUMBERS FOR EACH OF THE FOLLOWING: 1.N 2 2.CO 3.FeO 4.NO 5.MgO 6.K 7.Fe 2 O 3 8.CO 2 9.NO 2 CHALLENGE 1.Na 2 CO 3 2.CaSO 4 3.OF 2

Redox reactions are used for electrochemistry and are driven by a change in charge Oxidation – more (+) charge Reduction – more (-) charge Redox Reaction types: 1.A + B  AB (Ex. 2H 2 + O 2  2H 2 O) 2.AB  A + B (Ex. 2NaCl  2Na + Cl 2 3.A + BC  AC + B (Ex. Zn + Cu(NO 3 ) 2  Zn(NO 3 ) 2 + Cu) AIM: How can we recognize RedOx reactions?

Rules for recognizing RedOx reactions: 1. If a single element is on either side of the equation it must be a RedOx reaction 2. Assign oxidation numbers to all elements and see if any have changed from reactant to product 3. If the oxidation number increased, it lost electrons, therefore it was oxidized 4. If the oxidation number decreased, it gained electrons, therefore it was reduced

Rules for Identifying Which Species is Oxidized and Which Species is Reduced: 1.If charge becomes more (+) going from left to right  oxidized & lost electrons O L 2. If charge becomes more (-) going from left to right  reduced & gained electrons R G

Rules for Identifying Which Species is Oxidized and Which Species is Reduced: EX) Cu + 2AgNO 3  Cu(NO 3 ) 2 + 2Ag - Determine oxidation numbers for each - If there is a polyatomic ion that remains constant on both sides just look up the charge of the polyatomic ion Agents: Species that is oxidized  reducing agent Species that is reduced  oxidizing agent Spectator ion: Ion that does not change its charge

Directions: 1.Assign oxidation states to each species in the reaction 2.Identify the species being oxidized and species being reduced 3.Identify the reducing agent, oxidizing agent, and spectator ion (if there is one) Ex1) Cu + O 2  CuO EX2) Al + Cl 2  AlCl 3 EX3) N 2 + H 2  NH 3 EX4) H 2 + O 2  H 2 O EX5) Na + CaCl 2  NaCl + Ca

AIM: How can we write half reactions? Redox reactions may be split into 2 half-reactions, one for oxidation and one for reductions Examples: Oxidation Fe(s) -> Fe 3+ (aq) + 3e - Reduction Fe 3+ (aq) + 3e - -> Fe(s)

Do NOW: For each of the following: - Make sure each balanced - Assign oxidation numbers - Write the oxidation half rxn/reduction half rxn -Identify the oxidizing agent and reducing agent -Identify the spectator ion (if there is any) EX1) 2Li +Zn(NO 3 ) 2  2LiNO 3 +Zn EX2) 2K + Cl 2  KCl EX3) SnCl 2 + 2FeCl 3  SnCl 4 + 2FeCl 2

ELECTROCHEMISTRY: The branch of chemistry that deals with the relations between electrical and chemical phenomena. Electrochemical Cell – Voltaic Cell *Spontaneous Electrolytic Cell Electroplating *Non Spontaneous

Redox Potentials and Metal/Non metal Activity TABLE J  The higher up a substance is the more reactive it is  For metals the more reactive means higher potential to be oxidized  For non metals the more reactive means a higher potential to be reduced Ex) Which metal has the greatest potential to be oxidized? Ex) Which non metal has the greatest potential to be reduced?

Redox Potentials and Metal/Non metal Activity Ex) Which reaction is more likely to occur? 1.Cu +2 + Al 0  Cu 0 + Al +3 2.Cu 0 +Al +3  Cu +2 + Al 0

Redox Potentials and Metal Activity Ex) Which reaction is more likely to occur? 1.Cl F -1  2Cl -1 + F F Cl -1  2F -1 +Cl 2 0

Wednesday 5/7/2014 A-Day AIM: Electrochemistry DO NOW: For each of the following: 1. Check to see if it is balanced 2. Assign Oxidation numbers 3. Determined the oxidized/reduced species/spectator ion 4. Write the oxidation and reduction half reactions ***WORK QUIETLY AND INDEPENDENTLY (USE YOUR NOTES )**** EX 1) Al + O 2  Al 2 O 3 EX 2) Mg + ZnCl 2  MgCl 2 + Zn

VOLTAIC CELL: electrochemical cells in which spontaneous chemical reactions produce a flow of electrons

PARTS OF A VOLTAIC CELL: 1.2 half cells (2 beakers) 2.2 electrodes (anode and cathode) 3.Salt Bridge – MIGRATION OF IONS!!! (upside down U-TUBE) 4.Wire – flow of ELECTRONS! 5.Optional (voltmeter, switch)

AIM: How can we describe an electrochemical cell DO NOW: Using your notes from yesterday and your review books Label the following voltaic cell

VOLTAIC CELL RED CAT sat on AN OX REDUCTION occurs at the CATHODE OXIDATION occurs at the ANODE Remember from Table J - the higher metal will be oxidized; anode - the lower metal is the site of reduction; cathode **FLOW OF ELECTRONS IS ALWAYS FROM ANODE TO CATHODE!!!!!***** (RED CAT GETS FAT)

ELECTROLYTIC CELL Electrolysis: Electricity used to force a chemical reaction to occur

ELECTROLYTIC CELL Electroplating: Electrolysis can be used to electroplate metals onto a surface