4U-CH-notes-ch-9

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CHAPTER 9: E;ECTRIC CELLS
4 U CHEMISTRY
KEY IDEAS P 650
- introduction to electric cells and their uses.
9.1: Oxidation and Reduction
- reduction – metals from ores see ex p 652
- reducing agents – cause ores to be reduced to metals ex: CO, C and H2
- oxidation – metals to metal compounds , first with oxygen
- oxidizing agents – cause metals to be made into metal compounds, ex: O2 Cl2 Br2
p 653 practice # 1, 2, 3, 4.
Electron Transfer Theory
- chemical reactions involve transfer of e- , usually to make the molecules more stable
- single displacement reactions easy to split into half reactions
- Reduction: e- gained, oxidation # goes down
Ag+1 (aq) + e-  Ag (s)
- Oxidation: e- lost, oxidation # goes up
Cu (s)
 Cu+2 (aq) + 2 e- loss of e- and gain of e- happen simultaneously
- reaction must be balanced for charge as well as for # atoms
- Summary p 656
p 656-7 practice # 7, 8, 9, 10, 11
Oxidation States:
- net electric charge an atom WOULD have if the e- were attached to the more electronegative atom ie:
if all the bonds were ionic
- = oxidation #
- rules: see table 1 p 658 and Summary p 659
- p 659 practice # 12, 13, 14, 15, 16, 17.
Oxidation Numbers and Redox Reactions
- oxidation – Ox # increases
- reduction – Ox # decreases
- Redox reaction – both reduction and oxidation occur simultaneously
- p 662 practice # 18, 19, 20, 21.
- section 9.1 p 663 practice # 1, 2, 3, 4, 5, 6, 7, 8, 10.
9.2: Balancing Redox Equations
- 3 reasons for having a balanced chemical equation: ________________________________________
- _________________________________________________________________________________
Oxidation # method: the entire equation is given
1. write ox #
2. write half reactions for oxidation reaction and reduction reaction
3. balance the atoms with changed ox #
4. find total # e- being lost / gained
5. multiple each half reaction by the necessary multiple so that the e- lost = e- gained
6. put coefficients from half reactions into original equation and balance the rest by inspection
Ion-Electron method: may take place in acidic or basic solutions and may not have the entire equation
given:
1. separate the equation into half reactions
2. for each half reaction separately: balance the atoms being oxidized / reduced
3. balance O by adding water molecules
4. balance H by adding H ions
5. balance charge by adding e6. multiply each half reaction by the necessary factor so e- on left in one half reaction = e- on right in
other half reaction
7. add the half reactions and cancel any molecules and e- that are on both sides
8. FOR BASIC reactions: add OH-1 to both sides to neutralize any H+ ions
9. cancel out any extra water molecules formed
p 673 practice # 6, 7, 8
section 9.2 practice p 673 # 3, 4, 6a, 7
9.3: Predicting Redox Reactions:
Oxidizing and Reducing Agents – definitions
The Spontaneity Rule: a reaction will be spontaneous if the oxidizing agent on the left of the table is
HIGHER than the reducing agent on the right. see table appendix C11 p
Predicting Redox Reactions in Solution:
- copy blue box rules p 680
- Summary p 682
p 682 practice # 25, 26
section 9.3 practice # 1, 2, 3, 4, 5, 68, 9, 10, 11.
9.4: Technology of Cells and Batteries: read for interest
- electric cell:________________________________________________________________________
- battery: _________________________________________________________________________
Basic Cell Design and Properties:
- electrodes: ________________________________________________________________________
- electrolyte: ________________________________________________________________________
- Electric potential difference (voltage) ___________________________________________________
- Volts: ___________________________________________________________________________
- Electric current: ___________________________________________________________________
- Amperes: ________________________________________________________________________
- Coulombs: _______________________________________________________________________
- see Summary p 687
- p 687 practice # 1, 2, 3, 4
Consumer, Commercial and Industrial Cells
see table p 689 for common cell types
- primary cells: ______________________________________________________________________
- secondary cells: ___________________________________________________________________
ex: lead storage battery for a vehicle see p 689
Fuel Cells:
Ballard Fuel Cell
Aluminum Air Cell
9.5: Galvanic ( Electrochemical) Cells:
- the substances that can spontaneously react in a reaction like a single displacement reaction are
separated so that the e- that are transferred in the redox rxn are forced to travel through a wire, and so
are available to do work.
- half cell: __________________________________________________________________________
- salt bridge: ________________________________________________________________________
- porous cup: _______________________________________________________________________
A Theoretical Description of a Galvanic Cell:
- copy table 1 p 696
- cathode: reduction rxn of the strongest oxidizing agent (HIGHEST on the LEFT of the table)
SOAC / GERC – SEE P 697
- anode: oxidation rxn of the strongest reducing agent ( LOWEDT on the RIGHT of the table)
- copy diagram of the copper with silver cell p 697
Galvanic Cells with Inert Electrodes
- used when the solution, not the metal reacts, and is just a surface to deliver electrons
- ex: Carbon (graphite) rod or Platinum metal rod
Summary p 700
p 700 practice # 1, 2, 3, 4, 5, 6, 7b, 8.
Standard Cells and Cell Potentials
- standard cell: ____________________________________________________________________
- standard cell potential Eo: ___________________________________________________________
- standard reduction potential: Ero
- reference half cell: the hydrogen cell against which all other potentials are measured see diagram p
702
Measuring Standard Reduction Potentials
- difference between the reduction potentials of the two half cells
- see ex p 702 – 704
- Eo = Eor cathode - Eor anode
- if E is +, then reaction is spontaneous with the reactions as written
Summary p 705
Summary p 707
p 708 practice # 10, 11, 12.
9.5 section practice p 709 # 1, 2, 3, 4, 5, 6, 7, 8.
9.6: Corrosion:
- corrosion: ________________________________________________________________________
- any metal on the reduction table below the oxygen half reaction will corrode
- ex: Al . It corrodes to form Al2O3 (s) but the oxide sticks well to the metal “sealing” the metal from
further oxidation
Rusting of Iron:
- iron, water and oxygen necessary
- Fe(s)  Fe+2 + 2 eand ½ O2 + H2O + 2 e- 2 OH –1
- then iron II hydroxide oxidized to iron III hydroxide and iron III oxide – rust
- see diagram p 711
- rust forms faster in acid rain – standard redox reaction with hydrogen ions
- electrolytes speed up rusting ex: _____________________________________________________
Corrosion Prevention:
ex: _______________________________________________________________________________
- coating with tin: if tin is broken, cell set up with tin and iron and the iron oxidezes, faster than without
the tin
- Galvanizing: coat iron with Zinc – if zinc breaks, zinc oxidizes and iron is protected
Cathodic Protection:
- oxidation – loss of e-, happens at the anode
- to protect iron – force it to accept e- by attaching the neg end of a generator or battery to it
- need constant electrical supply for it to work
- used for pipelines and culverts
- sacrificial anode: a metal that oxidizes more readily than iron
- ex: Zn or Mg
- anode metal gradually oxidizes but iron is preserved
- Zn used to preserve ship hulls, Mg used to preserve hot water tanks in houses
9.6 practice p 714 # 1, 3, 4, 7, 8.
Problem Set: p 726 you must do # 2a, 3, 4ace, 6, 7ac, 8ace, 10ace, 13, 14, 15ac, 18, 20b, 22, 25a.
you should be able to do # 1, 2bc, 4bd, 7b, 8bd, 10bdf, 11a, 12, 15b, 16, 17, 20a, 21,
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