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Types of cells

Types of cells. Source or use of electricity. Electrode. Cell. –. Molten or aqueous chemicals. +. –. +. –. +. Overview. “Cells” are containers of liquid with electrodes:. In “electrolytic cells”, electricity is used to force chemicals to undergo a redox reaction

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Types of cells

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  1. Types of cells

  2. Source or use of electricity Electrode Cell – Molten or aqueous chemicals + – + – + Overview • “Cells” are containers of liquid with electrodes: • In “electrolytic cells”, electricity is used to force chemicals to undergo a redox reaction • In “galavanic cells”, electricity is produced spontaneously from a redox reaction

  3. Assignment • Read pg. 695 - 697. Answer these questions: 1) What in this room is a product of electrolysis? 2) Are ions at the anode gaining or losing electrons? What about at the cathode? 3) How is the conduction of electricity in a wire different from in an electrolytic cell. 4) Will electricity be conducted indefinitely through an electrolytic cell? Explain. 5) 697 gives the cell reaction for the electrolysis of NaCl. Write half reactions and the cell reaction for the electrolysis of HF(aq)

  4. Answers 1) Aluminum cans, copper wires 2) Anode = oxidation = loss of electrons (LEO) Cathode= reduction= gain of electrons (GER) 3) In wire, electricity means flow of electrons past metal atoms. In electrolysis, electricity means the movement of ions. 4) No. Only until all of the ions are used up. 5) 2H+(aq) + 2e– H2 (cathode / reduction) 2F–(aq)  F2 + 2e– (anode / oxidation) 2H+(aq) + 2F–(aq)  H2(g) + F2(g) Read over study note

  5. – + – + – + The electrolytic cell • Electric current forces charges on electrodes Na+ Na+ Cl– Cl– • Na+ is attracted to negative, Cl– to positive • Na+ takes up an electron: Na+(aq) + e– Na • Cl– gives up an electron: 2Cl–(aq)  Cl2 + 2e– • Thus electricity continues to flow • Pure Na is deposited, Cl2 gas is produced

  6. Activity • Add a scoop of CuCl2 to a 50 mL beaker • Add about 30 mL of distilled water • Stir until the CuCl2 is completely dissolved • Remove a piece of aluminum (about 5 cm square) from the role of aluminum foil • Submerse the aluminum in the CuCl2(aq) • What is produced? (think about the chemicals that you started with) • Write the redox reactions for what you saw: ___ + _e– ___ ___  ___ + _e– • Give the cell reaction • Dump solution down sink. Rinse & dry beaker Cu Cu2+ 2 Cu GER Al Al3+ 3 LEO 3Cu2++2Al3Cu+2Al3+

  7. Assignment (read 17.5) 1) Where in the room is there a galvanic cell? 2) In fig.17.12, is a solution with Cu2+ needed for the Cu half-cell to conduct? Is a solution containing Ag+ needed for the Ag half-cell? 3) Looking at 17.12, which electrode is losing electrons, which is gaining electrons, which is reduction, which is oxidation? 4) How do anodes and cathodes differ between electrolytic and galvanic cells? • Try PE 5 (similar to example 17.5) • You saw that Cu and Al react. How can these be used in a galvanic cell to produce energy (i.e. draw a diagram as in PE 5)

  8. Assignment 1) Batteries (computer, walkman) are galvanic 2) Cu2+ in unnecessary since it comes off of the Cu electrode, Ag+ is needed because this is deposited onto the Ag electrode 3) Cu is losing electrons to become Cu2+(oxidation), Ag+ is gaining electrons to become Ag (reduction). 4) In both, oxidation occurs at the anode and reduction at the cathode. Electrolytic: anode = + ve, cathode = – ve. Galvanic: anode = – ve, cathode = + ve

  9. Electron flow 5) Answers: PE 5 Mg  Mg2+ + 2e– (oxidation - LEO) Fe2+ + 2e–  Fe (reduction - GER) Mg + Fe2+  Mg2+ + Fe Fe (+) Salt bridge Mg (–) Fe2+ Mg2+

  10. Electron flow 6) Answers: Cu and Al Al  Al3+ + 3e– (oxidation - LEO) Cu2+ + 2e–  Cu (reduction - GER) 3Cu2++2Al3Cu+2Al3+ Cu (+) Salt bridge Al (–) Cu2+ Al3+ For more lessons, visit www.chalkbored.com

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