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Chapter 21: Electrochemistry

Table of Contents. Chapter 21: Electrochemistry. 21.1 – Voltaic Cells. 21.2 – Types of Batteries. 21.3 – Electrolysis. Table of Contents. Chapter 21: Electrochemistry. 21.1 – Voltaic Cells. Basic Assessment Questions. Try it out!.

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Chapter 21: Electrochemistry

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  1. Table of Contents Chapter 21: Electrochemistry 21.1 – Voltaic Cells 21.2 – Types of Batteries 21.3 – Electrolysis

  2. Table of Contents Chapter 21: Electrochemistry 21.1 – Voltaic Cells

  3. Basic Assessment Questions Try it out! Identify what is oxidized, what is reduced, the oxidizing agent & the reducing agent. Zn + NiSO4 → Ni + ZnSO4 Ox: Zn Red: Ni2+ Ox ag: Ni2+ Red ag: Zn • Define • Redox reactions • Oxidation • Reduction • Half-reactions

  4. Electrochemistry: Basic Concepts • Electrochemistryis using chemistry to create electricity • Electrical current: The flow of electrons in a particular direction • Redox reactions can be used to produce an electrical current. • This is what occurs in a battery—one form of an electrochemical cell

  5. Electrochemistry: Basic Concepts Definitions • An electrochemical cell is a device that uses redox reactions to create electricity OR electric energy to cause chemical reactions • A voltaic or galvanic cell converts chemical energy to electrical energy by a spontaneous redox reaction. It is a type of electrochemical cell.

  6. Balanced: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) Total ionic: Zn(s) + Cu2+(aq) + SO42-(aq) → Zn2+(aq) + SO42-(aq) + Cu(s) Net ionic: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) red ox Half Reactions: Zn(s) → Zn2+(aq) + 2e- Cu2+(aq) + 2e-→ Cu(s) Half Reactions: Zn(s) → Zn2+(aq) + 2e-

  7. Electrochemical Cell Zinc strip Copper strip Cannot transfer e- SO42- SO42- 1M Zn2+ 1M Cu2+ Zn(s) → Zn2+(aq) + 2e- Cu2+(aq) + 2e-→ Cu(s)

  8. Electrochemical Cell Allows e- transfer Copper wire (+) charge builds up in one solution, (-) charge builds up in the other

  9. Electrochemical Cell Allows ions to pass from one side to the other Ions pass through plugs in the bridge, but solutions do not mix Salt bridge KCl

  10. Electrochemical Cell e- flow K+ Cl- e- e- Cl- K+ e- e- Zn Cl- K+ Cu Zn2+ Cu2+ Zn → Zn2+ + 2e- Cu2+ + 2e- → Cu

  11. What is this called? What does it do? Why can redox reactions create electricity? e- flow KCl Zn Cu 1M Zn2+ 1M Cu2+ Zn → Zn2+ + 2e- Cu2+ + 2e- → Cu

  12. Voltaic/Galvanic Cell e- flow KCl electrodes electrodes half-cells Zn Cu 1M Zn2+ 1M Cu2+ electrolytes

  13. Electrochemistry: Basic Concepts Definitions • Half-cells: Where oxidation and reduction reactions separately take place (2 parts) • Electrode: an object in the half-cell that conducts electrons to or from another substance • Electrodes are immersed in electrolytes: usually a solution of ions

  14. Voltaic/Galvanic Cell Voltaic Cell e- flow KCl Zn Cu oxidation ANODE reduction CATHODE 1M Zn2+ 1M Cu2+ Zn → Zn2+ + 2e- Cu2+ + 2e- → Cu

  15. Electrochemistry: Basic Concepts Definitions • Anode:The electrode where oxidation takes place • Cathode:The electrode where reduction takes place

  16. Electrochemistry: Basic Concepts Practice • In the next slide, identify… • Salt bridge • What is oxidized • What is reduced • Cathode • Anode • Direction of e- flow through wire

  17. Voltaic/Galvanic Cell oxidized reduced

  18. Energy in a cell • We need to be able to… • find out how much energy we can get out of a cell • write the balanced equations of cells from half-reactions (tomorrow)

  19. Electrochemistry: Basic Concepts Equations of Cells • Standard Reduction Potentials: lists which are more likely to take electrons reduced Cu2+ + 2e- → Cu Cu2+ | Cu • More (+) means it’s more likely to take electrons

  20. Electrochemistry: Basic Concepts Copper-Hydrogen cell under standard conditions H2 (g) + Cu2+(aq) → 2H+(aq) + Cu(s) (net ionic equation) Cell notation H2 | H+ || Cu2+ | Cu reduction half-cell CATHODE oxidation half-cell ANODE salt bridge and wire

  21. Electrochemistry: Basic Concepts Cell Potential • To find a voltaic cell’s standard potential (how many volts it can use): Eºcell = Eºreduction - Eºoxidation • HINT: Write down everything. Show all work. When you do it in your head, you get confused.

  22. Electrochemistry: Basic Concepts Examples:Find the cell potential of a copper-zinc cell. Cu2+(aq) + Zn(s) → Cu(s) +Zn2+(aq) Step 1: Cell notation: Zn | Zn2+ || Cu2+ | Cu Step 2: ID ox and red ox red Step 3: Find reduction potentials Zn2+|Zn -0.762 VCu2+ | Cu +0.342 V Step 4: Equation Eºcell = Eºred - Eºox

  23. Electrochemistry: Basic Concepts Examples: Find the cell potential of a copper-zinc cell. Step 4: Equation Eºcell = Eºred - Eºox = EºCu2+|Cu - EºZn2+|Zn = 0.342 V – (-0.762 V) = +1.104 V The standard reduction potential of the copper-zinc cell is +1.104 V

  24. Additional Assessment Questions Try it out! Calculate the cell potential. 2Ag+(aq) + Co(s) → Co2+(aq) + 2Ag(s) + 1.08V

  25. Electrochemistry: Basic Concepts Potential Difference • A spontaneous reaction: proceeds naturally. Voltaic cells are always spontaneous. • A non- spontaneous reaction: requires an outside influence to proceed. (eg. forcing the electrons in the opposite direction) • In a Electrochemical cell… • when Eº = (+), it is spontaneous • when Eº = (-), it is not spontaneous

  26. Electrochemistry: Basic Concepts Spontaneity in a Cell • spontaneous • spontaneous • not spontaneous • spontaneous • not spontaneous • Eºcell = + 1.322 V • Eºcell = + 0.214 V • Eºcell = - 0.071 V • Eºcell = + 0.421 V • Eºcell = - 1.125 V

  27. Additional Assessment Questions Question 2b Calculate the cell potential to determine if each of these redox reactions is spontaneous. • Mn2+ + 2Br- → Br2 + Mn • Fe2+ + Sn2+→ Fe3+ + Sn • Ni2+ + Mg → Mg2+ + Ni • Pb2+ + Cu+→ Pb + Cu2+ -2.251 V, nonspontaneous -0.908 V, nonspontaneous +2.115 V, spontaneous -0.279 V, nonspontaneous

  28. Practice problems • Voltaic Cells handout • Due tomorrow at the end of class • Finished? Read ch 21 – it really will help! • I will not be covering 21.2 or 21.3 in class, but there are WebAssign questions on it - you’re going to have to read it! Part 1 Oxidation half: Cu → Cu2+ + 2e- Reduction half: Ag+ + e- → Ag

  29. Activity • At your tables, use whatever you have in your backpacks to make model of an electrochemical cell • Be able to explain to another group what each object represents in the “electrochemical cell” and what it does in the “cell” • You have 3 minutes

  30. Electrochemistry: Basic Concepts Sometimes, you must find a balanced reaction of cells from only reduction half-reactions. We will go through steps to follow and examples

  31. Electrochemistry: Basic Concepts Equations of Cells • Standard Reduction Potentials: lists which are more likely to take electrons reduced Cu2+ + 2e- → Cu Cu2+ | Cu • When comparing standard reduction potentials in a cell, the ½ reaction that is… • more (+) is reduction • more (-) is oxidation

  32. Electrochemistry: Basic Concepts 1) Identify which reaction is oxidation and which is reduction, 2) write the balanced equation EºAg+|Ag = +0.800 V EºNi2+|Ni = -0.257 V reduction oxidation • Ag+(aq) + e- → Ag(s) • Ni2+ + 2e- → Ni(s) Ni(s) → Ni2+ + 2e- 2 2 Ag+(aq) + e- → Ag(s) 2 • Ni(s) + 2Ag+(aq) → Ni2+ + 2Ag(s)

  33. Electrochemistry: Basic Concepts 1) Identify which reaction is oxidation and which is reduction, 2) write the balanced equation • Magnesium in a solution of Mg2+ • Lead in a solution of Pb2+ • Mg2+(aq) + 2e- → Mg(s) • Pb2+(aq) + 2e- → Pb(s)

  34. Electrochemistry: Basic Concepts 1) Identify which reaction is oxidation and which is reduction, 2) write the balanced equation EºMg2+|Mg = -2.372 V EºPb2+|Pb = -0.126 V ox red • Mg2+(aq) + 2e- → Mg(s) • Pb2+(aq) + 2e- → Pb(s) Mg(s) → Mg2+(aq) + 2e- Pb2+(aq) + 2e- → Pb(s) Mg(s) + Pb2+ → Mg2+(aq) + Pb(s)

  35. Table of Contents Chapter 21: Electrochemistry 21.2 – Types of Batteries 21.3 – Electrolysis In class we will cover only what you need to know for the test. Please use your book to answer questions for WebAssign.

  36. Electrochemistry: Basic Concepts 21.2: Batteries • Alessandro Volta: • If one cell generates a current, several cells should make a larger current • He piled several cells together to make the first battery • Batteries are one or more electrochemical cells in a single package that generates electrical current

  37. Electrochemistry: Basic Concepts 21.3: Electrolysis • Electrolysisis using electric energy to bring about a chemical reaction. • Electrolytic cell: An electrochemical cell in which electrolysis occurs. • Electrolysis forces a current in the reverse direction (a nonspontaneous reaction) by passing an electric current through it (recharging)

  38. Electrochemistry: Basic Concepts Voltaic Cell Electrolytic Cell Voltage source e- e- e- e- An electrolytic cell is just the opposite of a voltaic cell

  39. Electrochemistry: Basic Concepts Electroplating • Reduction of silver ions onto cheaper metals forms silverplate. Click box to view movie clip.

  40. Practice problems • Finish Ch 21 Practice Problems (due at the end of class) • Begin WebAssign (due Monday, 11 pm) • You will need your book (21.2 & 21.3) for some questions on WebAssign. Now is a great time to look those up. • Review for test is Monday! • Test on Ch 10, 20, 21 is Tuesday!

  41. End of Ch 21

  42. Electrochemistry: Basic Concepts The Electrolysis Process Click box to view movie clip.

  43. Electrochemistry: Basic Concepts Electrolytic Cell

  44. Basic Assessment Questions Distinguish between a voltaic cell and an electrolytic cell. spontaneous redox reaction; nonspontaneous redox reaction by electrolysis

  45. Electrochemistry: Basic Concepts Electroplating • Reduction of silver ions onto cheaper metals forms silverplate. Click box to view movie clip.

  46. Table of Contents Chapter 21: Electrochemistry 21.2 – Types of Batteries In class we will cover batteries only briefly. You will not be tested on this. Don’t take notes. Please use your book to answer questions for WebAssign.

  47. Electrochemistry: Basic Concepts Batteries • Alessandro Volta: • If one cell generates a current, several cells should make a larger current • He piled several cells together to make the first battery • Batteries are one or more electrochemical cells in a single package that generates electrical current

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