UNIT 10 Reduction/Oxidation Reactions & Electrochemistry NOTES
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1 Name Period CRHS Academic Chemistry UNIT 10 Reduction/Oxidation Reactions & Electrochemistry NOTES Quiz Date Lab Dates Exam Date Notes, Homework, Exam Reviews and Their KEYS located on CRHS Academic Chemistry Website:
2 Page 2 of 12 Unit 10 Notes 10.1 OXIDATION NUMBERS Oxidation-Reduction ( Reactions) reactions in which. Oxidation number In the past, we looked at group # on periodic table to find # of valence electrons, then used the octet rule to determine # of electrons lost or gained. Element # of valence e - e - lost or gained oxidation # S Li Al Mg Br However, the majority of elements can have oxidation numbers (remember the transition metals!) The following oxidation numbers are considered FIXED: All elements in free state have an oxidation number of All diatomic molecules (H2, O2, F2, Br2, I2, N2, Cl2) in free state have an oxidation number of Group 1 and 2 metals and aluminum (Al) are always as indicated on the periodic table Oxygen is in a compound (except peroxides) Hydrogen is almost always in a compound, except when bonded to metal, then -1 Rules: examples with Cl Cl2 KCl KClO3 More Practice Assign oxidation numbers to ALL of the elements in each of the compounds or ions below. FeO CuCl2 SO3 2- AlF3 H2SO4 Cr2O7 2- H3PO3 NO2 - NH3
3 10.2 THE MEANING OF OXIDATION AND REDUCTION Unit 10 Notes Page 3 of 12 Oxidation LEO goes GER Reduction All reactions are paired with a reaction. Electrons that are by the reactant that is is by the reactant that is. Example of redox reaction: REDOX: Mg + S MgS Oxidation Half Reaction: Reduction Half Reaction: The two reactions above are called Mg e to become the Mg 2+ cation. S e to become the S 2- anion. Reducing agent the reactant that gives up electrons Mg is the reducing agent and is oxidized (reducing agent is ALWAYS the reactant in oxidation halfreaction) Oxidizing agent the reactant that gains electrons S is the oxidizing agent and is reduced (oxidizing agent is ALWAYS the reactant in reduction halfreaction) 3
4 Page 4 of 12 Unit 10 Notes Practice 1: 4 Al + 3 O2 2 Al2O3 Oxidation half-reaction: Reduction Half Reaction: Reducing Agent: Oxidizing Agent: Practice 2: 2 AgNO3 + Cu Cu(NO3)2 + 2 Ag Oxidation: Reduction: Reducing Agent: Oxidizing Agent: How do I recognize a reaction? The. NOTE! reactions and reactions are NEVER redox! Practice 3: Which of the following reactions represent Redox reactions? 1. 2KNO3 (s) -> 2 KNO2(s) + O2 2. H2(g) + CuO(s) -> Cu(s) + H20 3. NaOH (aq) + HCl -> NaCl + H2O 4. H2(g) + Cl2 - > 2 HCl(g) 5. SO3(g) + H2O(l) -> H2SO4
5 Unit 10 Notes Page 5 of APPLICATIONS OF REDOX REACTIONS The branch of chemistry that deals with electricity related applications of oxidation-reduction reactions is called. Electrochemical Cell any device that converts energy into energy or vice versa. Electrochemical cells oxidation half reactions from reduction half reactions so that the energy produced or required is electrical energy and not heat. Two general types of electrochemical cells: Cells. Cells and Voltaic or Galvanic Cell Spontaneously converts energy into energy (using ). Applications - Basic parts of a voltaic cell 1. 2 each has a immersed in a. This is where separate oxidation / reduction reactions take place. a. the electrode where takes place (lose e - ) b. the electrode where takes place (gain e - ) 2. a solution of ions that conducts e 5
6 Page 6 of 12 Unit 10 Notes 3. a pathway constructed to allow ions to pass from one half-cell to another, filled with an electrolyte 4. metal wire carries electrons ( ) from to Example: zinc-copper voltaic cell Two half-reactions for redox: Oxidation: Reduction Zn + Cu 2+ Zn 2+ + Cu Anode ( ) where takes place is anode piece of is immersed in atoms of shorthand for this half-cell reaction is single represents Cathode ( ) where takes place is cathode piece of is immersed in atoms of shorthand for this half-cell reaction is Electric Cell Potential How do we measure how strong an electrochemical cell is? Electric potential, Ecell a measure of the cell s ability to. Units are Use table of relative Standard Reductions Potentials on Page 9 to calculate E 0 cell, where E 0 cell = E 0 red - E 0 oxid
7 Unit 10 Notes Page 7 of 12 Example: Calculate standard cell potential, E 0 cell, for the Zn-Cu electrochemical cell: E 0 red (Cu e -1 Cu) minus E 0 oxid (Zn Zn e -1 ) equals E 0 cell If E 0 cell is, then reaction is. If E 0 cell is, then reaction is. NOTE: Do not multiply Standard Reduction Potentials by coefficients of balanced redox reactions. 7
8 Page 8 of 12 Unit 10 Notes Example: Ni + Mg 2+ Mg + Ni 2+ a) Write the 2 half reactions and label as oxidation or reduction b) Calculate E 0 cell c) Is this redox reaction spontaneous? Practice 1: 2 Ag + Fe 2+ 2 Ag + + Fe a) Write the 2 half reactions and label as oxidation or reduction b) Calculate E 0 cell c) Is this redox reaction spontaneous? Practice 2: Given a voltaic cell constructed using the following half-reactions Cu e - Cu E 0 Cu = V Al e - Al E 0 Al = V a) Calculate E 0 cell b) Is this redox reaction spontaneous? c) What metal will anode? What metal will be cathode?
9 Unit 10 Notes Page 9 of 12 Electrolytic Cells If electrical energy is required to produce a redox reaction and bring about a chemical change in an electrochemical cell, it is an cell. The redox reaction in an electrolytic cell is and is essentially the of the voltaic cell. Voltaic/Galvanic Cell vs Electrolytic Cell Applications for each include : 9
10 Page 10 of 12 Unit 10 Notes
11 Unit 10 Notes Page 11 of 12 Common Polyatomic Ions List 20 Name Common Polyatomic Ions Ion Name Ion acetate C 2H 3O 2 or CH 3COO hypochlorite ClO ammonium + NH 4 nitrate NO 3 carbonate 2 CO 3 nitrite NO 2 chlorate ClO 3 perchlorate ClO 4 chlorite ClO 2 permanganate MnO 4 chromate 2 CrO 4 phosphate 3 PO 4 cyanide CN phosphite 3 PO 3 dichromate Cr 2O 2 7 silicate 2 SiO 3 hydrogen carbonate HCO 3 sulfate 2 SO 4 hydroxide OH sulfite 2 SO 3 1 = I 2 = II 3 = III 4 = IV 5 = V 6 = VI 11
12 Page 12 of 12 Unit 10 Notes
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