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Nel 4U: UNIT 7 LESSONS “ELECTROCHEMISTRY”
(Chap 9-pg 650-726 & Chap 10-pg 728-760)
1. Rationale:
Oxidation & Reduction reactions occur in many chemical systems. Examples
include the rusting of iron, the displacement of metals and photosynthesis in
plants. All of these reactions involve the transfer of electrons from one chemical
species to another. Because “electrons” are involved in all these reactions, these
changes are described as electrochemical changes and the study of these
changes is called electrochemistry.
UNIT 7 LESSONS – LESSON 1: “ELECTROCHEMISTRY”
2. REVIEW
a) Oxidation – Reduction Reactions (from SCH3U) – (9.1-pg 652)
DEMO lab activity: 1Write rxn eq’ns - Lets go over “oxidation” of Mg & Fe
in each case and the “oxidizing agent” or the substance “reduced”.
i)
Mg + O2 
( burn a pc of Mg) pg 194
ii)
Mg + HCl 
( drop pc of Mg into HCl sol’n – test gas) pg 206
iii)
Fe + CuSO4 
( drop pc Fe wool into sol’n - observe) pg 209-10
Read pg 652-656 * definitions on pg 655
Practice: # 2 on pg 653
*b) Activity Series (from SCH3U) – what it is? / what it means? - check 3U handout
** Note to self:
-23. Oxidation Numbers ( pg 657-659)
1. What are they ? What are they used for ?
2. Rules for assigning Oxidation Numbers
i)
The Oxidation # of all elements is “0”
ii)
The Oxidation # of any simple ion = charge
iii)
The Oxidation # of any compound = “0”
iv)
The Oxidation # of any compound ion = overall charge
v)
The Oxidation # of Hydrogen = +I, oxygen = -II
3. Do several examples (Handout Note)
ASSIGN : 7.1 (Handout)
UNIT 7 LESSONS – LESSON 2: “ELECTROCHEMISTRY”
4. Using Oxidation Numbers to balance REDOX rx equations (see pg 660-661 & 9.2 pg 664-673)
a) Lets do a few examples – Others assigned to be done over next 2 weeks
(Handout Note – Balancing Redox Reactions)
** Note to Self:
ASSIGN 7.2 – 7.3 + bonus
Quiz: Balancing REDOX equations (10 days later)
-3UNIT 7 LESSONS – LESSON 3: “ELECTROCHEMISTRY”
5. Electrochemical Cells ( Galvanic Cells ) (9.5 – pg 695-709)
*Technology of Cells & Batteries (9.4 – pg 685-694)
DEMO: Zinc/Copper EC cell
- Principles of the 2 half rxns being separate
- What if we hook up the volt meter – (Any current/voltage??)
- Principles of “salt bridge”
-
 see “Zn/Cu Electrochemical Cell”
see pg 696
**Note to self:
6. Cell Potentials of Electrochemical Cells ( E°)
pg 701-709
Define: Standard Oxidation Potential / Std Reduction Potential (
E°)
-4*
Standard 1/2 cell  H2 //H+ cell
 see Std Hydrogen cell” - pg 702
 Zinc 1/2 cell connected to std 1/2 cell
 see “Zn & Std H+ cell”
 Copper 1/2 cell connected to std 1/2 cell - pg 703
-5 Formation of the Std Oxidation Potential/reduction Potential List
 see “Standard Reduction Potentials”
Reference the “Senior Chemistry Data Sheet” – Std Oxidation Potentials
 Zinc 1/2 cell connected to the copper 1/2 cell
 see “Zn / Cu galvanic cell”
 Copper/Silver Cell & Nickel/Copper Cell
 see “galvanic cells”
HW: ASSIGNMENT QUESTION 7.4
-6UNIT 7 LESSONS – LESSON 4: “ELECTROCHEMISTRY”
7. Practical applications of EC Cells – Batteries (9.4
 Lead storage battery  Demo using car battery
 Zinc-Chloride dry cell  demo with cut up batteries
– pg 685-694)
(see pg 689)
(see pg 688)
 Nickel-Cadmium batteries  rechargeables
 Specialized batteries –Fuel Cells -
ASSIGNMENT QUESTION 7.5
(see pg 691-693)
-7-
UNIT 7 LESSONS – LESSON 5: “ELECTROCHEMISTRY”
8. Using Cell Potentials of EC Cells to predict “spontaneity”
 Looking for a positive E° – electrons will spontaneously flow from a substance that
will oxidize in the presence of the appropriate oxidizing agent ( ie the oxidizing agent
must be below the substance being oxidized on the Std Oxidation Potential list).
 Do examples:
a) Can you stir chlorine water with a nickel plated spoon?
b) Can you stir a KMnO4 sol’n with a chromium spoon?
c) Will copper react with hydrochloric acid? DEMO
d) Will copper react with nitric acid? – DEMO
e) Will copper react with sodium nitrate?
ASSIGN 7.6-7.11
-8UNIT 7 LESSONS – LESSON 6: “ELECTROCHEMISTRY”
9. Effect of Concentration on Cell Potentials – Nernst Eq’n
(pg 706-707)
 Lets illustrate how LeChatelier predicts that as time proceeds & the [reactants]
drops and [products] rises that E° should also drop.
Remember: ** it changes – it is not a constant.
 To calculate Ecell at any concentration we use the NERNST equation:
p
Ecell = E° – 0.0592/ne * log10 [products] / [reactants]
r
** explain “ne” as # electrons transferred
EX: Use the Nernst & Zn/Cu cell with [Cu2+] = 1x104 M & [Zn2+]=1x10-5 M
** Calculating concentrations using the Nernst Eq’n
see examples on overhead  see overhead – “Using the Nernst”
ASSIGN 7.12-7.14
-9-
UNIT 7 LESSONS – LESSON 7: “ELECTROCHEMISTRY”
10. Using Electrochemical Principles – REDOX Titrations
During any REDOX rxn:
# moles electrons lost must = # moles electrons gained
[subst. oxidized] * vol. * ne- lost = [subst. reduced] * vol. * ne- gained
*10b. LAB ACTIVITY Redox Titrations (2)
Complete Lab 7-1:
i)
ii)
find molar mass and then # moles water on ferrous diammonium disulfate
find molar mass and then # moles water on ferrous sulfate
Complete Lab 7-2:
i)
ii)
Find the % sodium hypochlorite in laundry bleach
Find the % sodium hypochlorite in “pool shock” & do cost analysis
Requires:
*Fe(NH4)2(SO4)2 .6 H2O
41.2 g/L
*FeSO4.6H2O
……. g/L
*KMnO4
*conc H2SO4/H3PO4 80%/20% mixture
Laundry bleach
Pool shock
Na2S2O3.5H2O
0.6 M KI
9.95 g /100 mL
40% acetic acid
40 mL glacial/60 mL water
ASSIGN: lab write-up + lab questions
-10-
UNIT 7 LESSONS – LESSON 8: “ELECTROCHEMISTRY”
11. CONNECTIONS: Cell Potentials & Thermodynamics / equilibrium
A) E° and Ke
As a reaction proceeds the concentration of all reactants drops and concentration
of all products increase until the system reaches a state of Equilibrium where the
system can not deliver any more useful energy ( work) so E = 0
p
0 = E° – 0.0592/n * log10 [products] / [reactants]
So
log10Ke = nE°/0.0592
or
r
(Nearnst)
(since [prod] & [react] are now constant)
Ke = 10nE°/0.0592
Do example – find Ke value for Zn/Cu cell
( 1.4 x 1037 )
B) E° and ∆G°:
Recall from Equilibrium/Thermodynamics that
∆G° = –RT ln Ke
or ∆G° = –RT (2.3)log10Ke
so
log10Ke = –∆G°/2.3 RT
so:
but
log10Ke = nE°/0.0592
as well
∆G°/nE° = –2.3RT/0.0592
(grouping all constants)
∆G°/nE° = –96 500 coul / mol e–
( volt=J/coul)
or
∆G°/nE° = – 1 Faraday
or
∆G° = –nFE°
-11-
Do example – find ∆G° value for Zn/Cu cell
( –212.3 kJ)
Do example – find Ke & ∆G° value for:
NiO2 + 2Cl– + 4 H+  Cl2 + Ni2+ + 2 H2O E° = 0.32 v
( 6.5 x 1010 )
( –62 kJ)
ASSIGN 7.15-7.18
-12-
UNIT 7 LESSONS – LESSON 9: “ELECTROCHEMISTRY”
12. Electrolysis (Chapter 10
– pg 728-753)
Passing electricity through a molten ionic compound or a solution of any electrolyte causes
a REDOX reaction to occur – called electrolysis.
** Comparison of Galvanic vs Electrolytic cells see chart on pg 731
Ex 1: Electrolysis of molten sodium chloride – see diagram pg 738
* go over ‘anode’ & “cathode” and “cell reaction”
Ex 2: Electrolysis of potassium bromide(aq)
* do “cell reaction” – DEMO
Ex 3: Electrolysis of copper(II) bromide (aq)
* do “cell reaction” – DEMO
** READ Summary -pg 733-34 Do example 1/2 pg 734-35 
ASSIGN 7.19
-13-
UNIT 7 LESSONS – LESSON 10: “ELECTROCHEMISTRY”
13. Stoichiometric relationships in Electrolysis (10.3
– pg 747)
*Quantitative Electrolysis
1 mol Ag+ + 1 mol e– (1 Faraday)  1 mol Ag
1 mol Cu2+ + 2 mol e– (2 Faraday)  1 mol Cu
1 mol Al3+ + 3 mol e– (3 Faraday)  1 mol Al
* use
Q (coulombs) = I (amp) * t (s)
then
# Faradays = Q/96500
1 mole of electrons will
reduce 1 mol of Ag but
only 0.5 mole of Cu
and 0.33 mole of Al .
to find the # coulombs
# moles red = #F / Δ charge
moles reduced  # grams
see EX 1 pg 749 (calculating moles of electrons)
see EX 2 pg 749 (calculating time)
see Sample problem pg 750 (calculating mass from current)
see Example pg 751 (calculating current from mass)
ASSIGN 7:20-7.24
-14-
14. Industrial Applications of Electrolysis (10.2 –
discuss
a) electroplating ( pg 743 )
pg 737-745)
DEMO
b) production of Sodium/Aluminum ( pg 738-739 )
 see overhead – “Aluminum production”
c) Refining of Metals ( pg 742)
ASSIGN : summary question: 7:25
MAJOR TEST #7: Electrochemistry