1. 17.1 Galvanic Cells (Batteries)
Ch. 17: Electrochemistry
17.1 Galvanic Cells
(Batteries)

2. Galvanic Cells
device that changes chemical energy into electrical energy
uses a reduction/oxidation reaction
reducing agent transfers electrons to the oxidizing agent
oxidation:
loss of electrons
(  ox. state)
reduction:
gain of electrons
(  ox state)

3. Galvanic Cells
if the two half-reactions are combined in the same container, the electron exchange occurs directly as work is released as heat= no electricity
to harness the energy, we keep each half-reaction in a separate container so the electron transfer occurs through a wire = Electricity

4. Galvanic Cell
salt bridge or porous disk used to allow for unrelated ions to move to allow for balance of charge

5. Parts of Galvanic Cell anode: oxidation cathode: reduction Names for
Q. In which direction will the electrons flow?
A.from reducing agent to oxidizing agent
Names for
locations
of each
half-rxn
anode:
oxidation
cathode:
reduction

6. Car Battery: Lead storage battery

7. Cell Potential
cell
driving force on electron to move them through the wire
also called Electromotive Force (emf)
units are Volts (V)
1 V = 1 Joule/Coulombs of charge
skip the last paragraph in too much physics!

8. 17.2 Standard Reduction Potentials
Ch. 17: Electrochemistry
17.2 Standard Reduction Potentials

9. Standard Reduction Potentials
we assign  values to each half-reactions
we can find the total cell by summing the individual potentials for the combination of half-reactions
they are always written as a reduction process so must switch one of them
Always
change the sign of ° when you reverse the direction
do not multiply the ° by an integer used to balance the equation- it is intensive

10. Standard Reduction Potentials
° values for reduction half-reactions with all solutions having 1 M conc and all gases 1 atm

11. Standard Hydrogen Electrode
Pt electrode in contact with 1 M H+ and H2(g) at 1 atm
assigned an  of zero
can calculate others by pairing with this and measuring total cell
2H+(aq) + 2e-(aq)  H2(g)

12. Standard Hydrogen Electrode

13. Example Find the °cell for the reaction
Fe3+(aq) + Cu(s)  Cu2+(aq) + Fe2+(aq)
half reactions:
Fe3+(aq) + e-  Fe2+(aq) °=0.77 V
Cu2+(aq) + 2e-  Cu(s) °=0.34 V
2nd one must be reversed
°cell = -0.34V V = 0.43V

14. Line Notation short hand for describing cells
anode is on the L and cathode is on R
separate anode and cathode with ||
separate phases in one half-rxn with |
if no part of a half-rxn is solid, use Pt electrode
electrodes go on far ends of notation
Mg(s) | Mg2+(aq) || Al3+(aq) | Al(s)
Pt(s) | ClO3-(aq), ClO4-(aq) || MnO4-(aq), Mn2+(aq) | Pt(s)

15. Galvanic Cells run spontaneously in the direction where °cell is +
Describe a Galvanic Cell:
balanced chemical eq. (make sure °cell +)
give the direction of electron flow
assign the anode and cathode
give line notation

16. Example Write balanced equation
Describe the Galvanic Cell based on the following half-reactions:
Ag+ + e-  Ag °=0.80 V
Fe3+ + e-  Fe2+ °=0.77 V
Write balanced equation
Fe eq. must be reversed
Ag+ + Fe2+  Ag + Fe3+ °cell = 0.03V

17. Example Assign cathode and anode Give the direction of electron flow
oxidation: Fe2+  Fe3+ +e-
reduction: Ag+ +e-  Ag
electrons flow from Fe2+ compartment to Ag+ compartment
Assign cathode and anode
anode: oxidation: Fe2+  Fe3+ +e-
cathode: reduction: Ag+ +e-  Ag

18. Line Notation
Pt(s) | Fe2+(aq), Fe3+(aq) || Ag+(aq) | Ag(s)

19. Alkaline Battery

20. Sacrificial Metal to prevent rust on large structure

21. Fuel Cell (Hydrogen fuel)