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Unit 13: Electrochemistry (Link to Prentice Hall Text: Chapters 22 & 23) Name:____________________________________________________________________________________ Date Due Assignments Page Number: Problem Numbers Assignment 1: Balancing and Identifying REDOX Reactions 673: 30, 32, 34 Assignment 2: Electrochemical Cells 701: 22, 34, 35, 38, 42 A. Atoms Compete for Electrons (Electrochemistry) Defining Oxidation and Reduction When chemical bonds form, electrons are either lost, gained or shared. In REDOX reactions, electrons are lost or gained. Oxidation Is Loss, Reduction Is Gain, “OIL RIG” Oxidation: - Loss of electrons. - Metals are more easily oxidized. - They are termed “reducing agents.” - If a substance is oxidized, its oxidation number increases. Reduction: - Gain of electrons. - Non-metals are more easily reduced. - They are termed “oxidizing agents.” - If a substance is reduced, its oxidation number decreases. B. Oxidation Numbers Rules for Assigning Oxidation Numbers: (1) Oxidation numbers for atoms that are free elements are always _________. (2) The oxidation number of monatomic ions are the same as the charge on the __________. (3) The sum of the oxidation numbers in a compound is always __________. (4) The sum of the oxidation numbers in a polyatomic ion is equal to ________________. (5) The oxidation number of Group 1 metals is always __________. (6) The oxidation number of Group 2 metals is always __________. (7) Oxygen almost always has a ______ oxidation state.* Unless in a peroxide (8) Halogens usually have a ______ oxidation state.* (9) Hydrogen usually has a ______ oxidation state.* * There are rare exceptions to these rules. Practice Assign oxidation numbers to each element in the following compounds or polyatomic ions. To really keep on your toes, see if you can name them as well. 1. MgBr2 9. CuSO4 17. K2Cr2O7 2. Cu 10. Cr 18. Al2O3 3. Fe2O3 11. H2CO3 19. Fe(NO3)3 4. AlN 12. Ba(NO3)2 20. SrCO3 5. SO3 13. NF3 21. Na2SO3 6. PO43- 14. CO 22. Ca(ClO)2 7. Cr2O72- 15. CO2 23. H2O 8. HClO2 16. CH4 24. K2S C. Identifying REDOX Reactions Identifying REDOX Reactions Using Oxidation Numbers When a REDOX reaction occurs, there must be an element that is reduced (gained electrons) and an element that is oxidized (lost electrons). Are the following reactions REDOX reactions? HINT: Find oxidation numbers of each element. (1) 2H2 + O2 2H2O (2) HCl + NH3 NH4Cl Oxidation Is Loss, Reduction Is Gain, “OIL RIG” Identifying Oxidation and Reduction Half Reactions Oxidation and reduction occur simultaneously. Elements and charges must be balanced for each half-reaction. 1. Zn + CuCl2 Cu + ZnCl2 Oxidation Half: Reduction Half: 2. Na + Cl2 NaCl Oxidation Half: Reduction Half: 3. Ca + O2 CaO Oxidation Half: Reduction Half: 4. MnO2 + HCl MnCl2 + H2O + Cl2 Oxidation Half: Reduction Half: 5. Cu + Ag+ Ag + Cu2+ Oxidation Half: Reduction Half: D. Balancing REDOX Reactions Using Half-Reactions Steps to Balancing REDOX Reactions When a REDOX reaction occurs, there must be an element that is reduced (gained electrons) and an element that is oxidized (lost electrons). The number of electrons lost in the oxidation must equal the number of electrons gained in the reduction. (1) (2) (3) (4) (5) (6) (7) (8) Assign oxidation numbers. Decide what is being oxidized and what is being reduced. Write the half reaction for each oxidation and reduction. Balance all atoms Balance the charge with electrons in each half reaction. Balance the number of electrons gained with the number of electrons lost. Add the two half reactions. Simplify each equation by canceling out things that are the same on both sides. Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons), “OIL RIG” Oxidation Step: Oxidation number of element increases. Reduction Step: Oxidation number of element decreases. Hints: 1. 2. There should be no electrons in your final equation. The number of each element should be the same on both sides of the arrow when you are finished. Balancing REDOX Reactions Oxidation and reduction occur simultaneously. Elements and charges must be balanced for each half-reaction. Electrons must be the same on either sides of the arrow. 1. Zn + CuCl2 Cu + ZnCl2 Oxidizing Agent Reducing Agent Oxidizing Agent Reducing Agent Oxidation Half: Reduction Half: Balanced Reaction: 2. Ni + Sn4+ Ni2+ + Sn Oxidation Half: Reduction Half: Balanced Reaction: 3. Ni + Sn+4 Ni+2 + Sn Oxidizing Agent Reducing Agent Oxidizing Agent Reducing Agent Oxidizing Agent Reducing Agent Oxidizing Agent Reducing Agent Oxidizing Agent Reducing Agent Oxidation Half: Reduction Half: Balanced Reaction: 4. Hg + Ag+1 Ag + Hg+2 Oxidation Half: Reduction Half: Balanced Reaction: 5. KClO3 KCl + O2 Oxidation Half: Reduction Half: Balanced Reaction: 6. H2S + O2 SO2 + H2O Oxidation Half: Reduction Half: Balanced Reaction: 7. H+ + Sn Sn+2 + H2 Oxidation Half: Reduction Half: Balanced Reaction: Electrochemistry Page |6 E. Electrochemical Cells The Basics of an Electrochemical Cell Every electrochemical cell (contains two compartments where REDOX reactions can occur) contains 2 electrodes. An electrode is a strip of metal that facilitates the loss or gain of electrons. There are two types of electrodes, known as the cathode and the anode. Anode Cathode Where ___________________ occurs. Where ___________________ occurs. Electrons always flow from the ___________________ to the ___________________. If you have 2 electrodes made of different metals connected, how can you tell which one is oxidized and which one is reduced? Table J! Oxidation If the metal is ______________ on Table J. (More active = easier to lose electrons) Reduction If the metal is ______________ on Table J. (Less active = harder to lose electrons) Voltaic or Galvanic Electrochemical Cells In a Galvanic/Voltaic cell, reactions occur ________________________; therefore, no additional energy must be added. Spontaneous Reaction - Reactions that proceed without the addition of energy. - Predicting Spontaneous Reactions (Table J) - If the metal by itself if higher up on the chart than the ion in the compound. Electrochemistry Page |7 - In a voltaic cell, the anode is denoted with a NEGATIVE sign. In a voltaic cell, the cathode is denoted with a POSTIVE sign. - Each half cell is connected by a salt bridge. o What is a salt bridge? - Negative ions always flow from the _____________________________ to the _________________________. - Positive ions always flow from the _____________________________ to the _________________________. - The electrodes are connected by a metal wire which allows for the flow/movement of electrons. Remember: Electrons always flow from the anode to the cathode. - A voltmeter usually measures the electron potential in Volts. Electrochemistry Page |8 Determining Spontaneity Which of the following reactions will take place spontaneously? a. Ni (SO4) + Pb b. Sr(CO3) + Sn Ni + Pb(SO4) Sr + Sn(CO3) c. Au (PO4) + Al Au + Al (PO4) d. Fe(OH) + Cu Fe + Cu(OH) Labeling Electrochemical Cell Processes Between each of the following electrodes: Species Oxidized Calcium and Iron Silver and Nickel Magnesium and Lead Copper and Silver Species Reduced Anode Metal With Sign Cathode Metal With Sign Direction of eFlow Direction of Cation Flow in the Salt Bridge Electrochemistry Page |9 Electrolytic Cells Reactions that occur in an electrolytic cell occur ____________________________________________. You must force the reaction to occur by adding electricity. In Electrolytic Cells: - The anode is POSITIVE. - The cathode is NEGATIVE. ____________________ ____________________ Why would you ever want to force a nonspontaneous reaction? (a) To obtain pure metals Many metals are only found as compounds in nature. Electrolysis can lead to a deposit of the pure metal on the cathode. (b) To recharge a battery A car battery powers the car through a spontaneous reaction, but what can you do if the battery dies? (c) To coat one metal on top of another one, as with jewelry, or exhaust pipes. a. To make something look more expensive or shinier b. To improve corrosion resistance Electrochemistry P a g e | 10 Using Electrolysis for Electroplating An example of electrolytic cells in which an electric current is used to plate an object. The substance to be plated is at the _____________________________________. The substance that does the plating is at the ______________________________. If I wanted to gold plate a necklace, where should I put the gold and where should I put the necklace in the electrolytic cell? Why is it necessary to use an electrolytic cell to gold plate almost all other metals instead of a voltaic cell? Would I need to use an electrolytic cell to plate nickel on top of an iron pipe? How do you know? Electrochemistry P a g e | 11 Labeling Electrolytic Cells Label the following for each of the electrolytic cells below. a. Anode b. Cathode c. Reduction Electrode d. Oxidation Electrode e. Show electron flow using arrows. f. Label ions. g. Has the anode increased or decreased in mass? h. Has the cathode increased or decreased in mass? (a) Zn/Cu in CuSO4 (b) Ni/Ag in AgNO3 (c) (d) Ag and a ring in AgNO3 Cu/Pb in CuCl2 Electrochemistry P a g e | 12 Calculating Cell Potentials It is possible to calculate the potential difference between two half-cells using standard reduction potentials. All standard reduction potentials are tabulated relative to the Standard Hydrogen Electrode (SHE). The more positive the reduction potential value, the more likely the substance is to undergo reduction. Electrochemical Potential or Electromotive Force (emf) E°cell = E°red (cathode) - E°red (anode) Determining Cell Potentials Predict the voltages or emf produced by the following cells. 1. Zn/Zn+2//Fe+2/Fe 2. Mn/Mn+2//Br2/Br- Electrochemistry 3. Ni/Ni+2//Hg2+2/Hg 4. Cu/Cu+2//Ag+/Ag 5. Using standard reduction potentials, calculate the standard emf for each of the following reactions: P a g e | 13 (a) H2 (g) + I2 (s) 2H+ + 2I- (b) Ni(s) + 2Ce+4 (aq) Ni+2 (aq) + 2Ce3+(aq) (c) Cr(s) + 2Cr+3 (aq) 3cr+2 (aq) (d) 2Al3(aq)- + 3Cd(s) 2Al(s) + 3Cd2+(aq)] 6. A voltaic cell is constructed: one electrode compartment has an aluminum strip in contact with a solution of Al(NO3)3, and the other is a standard hydrogen electrode. (a) Write the half reactions involved and determine which electrode is the anode and which is the cathode. (b) Will the aluminum strip gain or lose mass as the cell operates? (c) Write a balanced equation for the overall cell reaction. (d) What is the standard emf of the cell?