Transcript Types of cells Overview • “Cells” are containers of liquid with electrodes: Source or use of electricity Electrode Cell – + Molten or – + – + aqueous chemicals • In “electrolytic cells”, electricity is used.

Types of cells
Overview
• “Cells” are containers of liquid with electrodes:
Source or use
of electricity
Electrode
Cell
–
+
Molten
or
–
+
–
+
aqueous
chemicals
• 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
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)
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
The electrolytic cell
Cl–
•
•
•
•
•
•
–
–
–
Na+
Cl–
+
+ Na+
+
Electric current forces charges on electrodes
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
Activity
1)
2)
3)
4)
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
5) Submerse the aluminum in the CuCl2(aq)
6) What is produced? (think about the
Cu
chemicals that you started with)
7) Write the redox reactions for what you saw:
___
Cu2++ _e
2 –  ___
Cu GER ___
Al  ___
Al3++ _e
3 – LEO
8) Give the cell reaction 3Cu2+ +2Al3Cu+2Al3+
9) Dump solution down sink. Rinse & dry beaker
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?
5) Try PE 5 (similar to example 17.5)
6) 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)
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
5) Answers: PE 5
Electron flow
Fe (+) Salt bridge Mg (–)
Fe2+
Mg2+
Mg  Mg2+ + 2e– (oxidation - LEO)
Fe2+ + 2e–  Fe (reduction - GER)
Mg + Fe2+  Mg2+ + Fe
6) Answers: Cu and Al
Electron flow
Cu (+) Salt bridge
Cu2+
Al (–)
Al3+
Al  Al3+ + 3e– (oxidation - LEO)
Cu2+ + 2e–  Cu (reduction - GER)
3Cu2+ +2Al3Cu+2Al3+
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