Chemistry 12 Unit 5- Electrochemistry



Chemistry 12 Unit 5- Electrochemistry

Lesson 1: Redox Reactions

Oxidation – Reduction

Definitions: (species means atom, ion or molecule)

Oxidation – a species undergoing oxidation loses electrons

(charge becomes more positive)

Reduction – a species undergoing reduction gains electrons

(charge becomes more negative)

Oxidizing agent – The species being reduced

(gains electrons, causes the other one to be oxidized)

Reducing agent – The species being oxidized

(loses electrons, causes the other one to be reduced)

2 e-

E.g.) Cu2+ (aq) + Zn (s) ( Cu(s) + Zn2+(aq)

Oxidizing agent

|LEO says GER |

|Losing Electrons is |Gaining Electrons |

|Oxidization |is Reduction |

| | |

Redox – Short for Oxidation – Reduction

Charge on neutral atom or molecule = 0

Oxidation – Charge gets more + (loses electrons)

Reduction – Charge gets more – (gains electrons)

Reduction (charge decreases)

E.g.) Pb2+(aq) + Mg0(s) ( Pb0(s) + Mg2+(aq)

Oxidation (Charge increases)

Question

In the reaction:

2Fe2+ + Cl2 ( 2Fe3+ + 2Cl-

Identify:

a) The Oxidizing Agent: _____________________

b) The species being oxidized:________________

c) The reducing agent:______________________

d) The species being reduced:________________

e) The species gaining electrons:______________

f) The species losing electrons:_______________

g) The product of oxidation___________________

h) The product of reduction___________________

Do Ex. 1 (a-e) pp. 192 SW

Half-Reactions

-Redox reactions can be broken up into oxidation & reduction half reactions.

e.g.) Redox rx: Pb2+(aq) + Zn(s) ( Pb(s) + Zn2+(aq)

The Pb2+ (loses/gains) __________ 2 electrons.

Reduction Half-rxn: Pb2+(aq) + 2e- ( Pb(s)

Oxidation Half-rxn:

Note: Half-rx’s always have e-‘s, redox (oxidation-reduction) reactions never show e-‘s!

Given the redox reaction:

F2(g) + Sn2+(aq) ( 2F-(aq) + Sn4+(aq)

Write the oxidation half-rxn:

__________________________

Write the reduction half-rxn:

__________________________

Do ex. 2 a-c on p. 192 SW

Oxidation Numbers

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Practice

1. Given the redox reaction:

2MnO4- + 3C2O42- + 4H2O ( 2MnO2 + 6CO2 + 8OH-

Find:

a) The species being reduced: _____________.

b) The species undergoing oxidation: _____________.

c) The oxidizing agent: ________________.

d) The reducing agent: ________________.

e) The species gaining electrons: ______________.

f) The species losing electrons: ______________.

2. Given the balanced redox reaction:

3S + 4HNO3 ( 3SO2 + 4NO + 2H2O

Find:

a) The oxidizing agent: _____________.

b) The reducing agent: _____________.

c) The species being reduced: ___________.

d) The species being oxidized: ___________.

e) The species losing electrons: ____________.

f) The species gaining electrons: ____________.

g) The product of oxidation: ____________.

h) The product of reduction: ____________.

3. Given the following:

6Br2 + 12KOH ( 10KBr + 2KBrO3 + 6H2O

Find:

a) The oxidizing agent: ______________.

b) The reducing agent: ______________.

c) The species undergoing oxidation: _____________.

d) The species being reduced: ________________.

e) The product of oxidation: _______________.

f) The product of reduction: _______________.

Read p. 193-195 of SW. Do Exercise 3 on p. 194 of SW.

Do Exercises 4, 5 and 6 on p. 194-195 of SW.

Lesson 2: Predicting Spontaneous Reactions

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Assignment: Using data to make your own simple Redox table

Example problem:

1) Four metals A, B, C, & D were tested with separate solutions of A2+, B2+, C2+ & D2+. Some of the results are summarized in the following table:

|Solution |

|Metal |A2+ |B2+ |C2+ |D2+ |

|A | |(1) no reaction |(2) reaction | |

|B | | | |(4) no reaction |

|D |(3) reaction | | | |

List the ions in order from the strongest to weakest oxidizing agent.

Using data

Another example –

Four non-metal oxidizing agents X2, Y2, Z2 and W2 are combined with solutions of ions: X-, Y-, Z- and W-.

The following results were obtained;

1) X2 reacts with W- and Y- only.

2) Y- will reduce W2

List the reducing agents from strongest to weakest

Do Exercises 14,15,16 & 18 on p. 200 of SW.

Lesson 3: Balancing Half-Reactions

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Balancing Redox Reactions Using Half-Reactions

1) Break up Rx into 2 half-rx’s.

2) Balance each one (in acidic or basic as given)

3) Multiply each half rx by whatever is needed to cancel out e-‘s

(Note: balanced half-rx have e-‘s (on left reduction on right oxidation) Balanced redox don’t have e-‘s)

4) Add the 2 half-rx’s canceling e-‘s and anything else (usually H2O’s, H+’s or OH-‘s) in order to simplify.

Example: U4+ + MnO4- ( Mn2+ + UO22+ (acidic)

Try this one:

SO2 + IO3- ( SO42- + I2 (basic solution)

-See examples p.205-207 in SW

Quick notes

-Some redox equations have just one reactant

- Use this as the reactant in both half-rx’s.

- These are called “self-oxidation-reduction” or Disproportionation reactions.

Eg) Br2 ( Br-- + BrO3- (basic) (found in some hot tubs)

Half rx’s are:

|Br2 ( Br-- | Br2 ( BrO3- |

| | |

Answer: ___________________________________________________

Do Ex 24 a-m p. 207

Extension: Balancing using Oxidation Numbers

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Practice

1. U4+ + MnO4- ( Mn2+ + UO22+ (acidic)

2. SO2 + IO3- ( SO42- + I2 (basic solution)

3. Br2 ( Br-- + BrO3- (basic)

-This is optional

- As long as one method (not guessing!) works for you that’s fine. (This method or half-rx method.)

- Read examples p. 271-272 SW

- Do any ex 10 a-d, f, j, m & check with key

Lesson 4: Redox Titrations

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Read p. 210-212 carefully – go over the examples! Do ex 26 & 29 p. 213-214 SW.

Lesson 5: Electrochemical Cells

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3. Demo: Cu/Zn Cell

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Consider the following cell:

The voltage on the voltmeter is 0.45 volts.

a) Write the equation for the half-reaction taking place at the anode. Include the Eo.

______________________________________________ Eo: _________v

b) Write the equation for the half-reaction taking place at the cathode.

______________________________________________ Eo: _________v

c) Write the balanced equation for the redox reaction taking place as this cell operates. Include the Eo.

______________________________________________Eo: __________

d) Determine the reduction potential of the ion X2+.

Eo: __________v

e) Toward which beaker (X(NO3)2) or (Cr(NO3)3) do NO3- ions migrate?

_______________________ _

f) Name the actual metal “X” ________________________________

Consider the following cell:

The initial cell voltage is 1.20 Volts

a) Write the equation for the half-reaction which takes place at the cathode. Include the Eo

____________________________________________________ Eo= ____ ___v

b) Write the equation for the half-reaction taking place at the anode:

____________________________________________________ Eo= ____ ___v

c) Write the balanced equation for the overall redox reaction taking place. Include the Eo.

____________________________________________________ Eo= ____ ___v

d) Find the oxidation potential for Cd: Eo= ____ ___v

e) Find the reduction potential for Cd2+: Eo= ____ ___v

f) Which electrode gains mass as the cell operates? _______

g) Toward which beaker (AgNO3 or Cd(NO3)2) do K+ ions move? _______

h) The silver electrode and AgNO3 solution is replaced by Zn metal and Zn(NO3)2 solution.

What is the cell voltage now? __________Which metal now is the cathode? _________

Consider the following electrochemical cell:

a) Write the equation for the half-reaction taking place at the nickel electrode. Include the Eo

____________________________________________________ Eo= ____ ___v

b) Write the equation for the half-reaction taking place at the Cu electrode. Include the Eo.

____________________________________________________ Eo= ____ ___v

c) Write the balanced equation for the redox reaction taking place.

____________________________________________________ Eo= ____ ___v

d) What is the initial cell voltage? _________________ _

e) Show the direction of electron flow on the diagram above with an arrow with an “e-“ written above it.

f) Show the direction of flow of cations in the salt bridge using an arrow with “Cations” written above it.

Challenge:

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Lesson 6: Standard Reduction Potentials

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Read SW. p. 215-224

Do Ex. 35 p. 217 and Ex. 36 a-d & 37-45 on p. 224-226 of SW

Lesson 7: Electrolysis

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3 Types

Type 1:

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Type 2:

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Overpotential Effect

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Type 3:

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Lesson 8: Applications of Electrochemistry- Part 1: Electroplating and Electrorefining

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Demo: Copper Plating

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Anode Half Reaction:

Cathode Half Reaction:

Overall Reaction:

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Lesson 9: Applications of Electrochemistry- Part 2: CORROSION

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NOTE: For the same element: The more positive species is always the Oxidizing Agent.

Eg.) A2+ A

OA

RA

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1M X(NO3)2

1M Cr(NO3)3

Metal “X”

Balance each ½ rx

U4+ ( UO22+

(Major (U) balanced already)

Oxygen ( U4+ + 2H2O ( UO22+

Hydrogen ( U4+ + 2H2O ( UO22+ + 4H+

Charge ( U4+ + 2H2O ( UO22+ + 4H+ + 2e-

MnO4- ( Mn2+

(Major (Mn) balanced already)

MnO4- ( Mn2+ + 4H2O

MnO4- + 8H+ ( Mn2+ + 4H2O

MnO4- + 8H+ + 5e- ( Mn2+ + 4H2O

Cr

e-

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e-

Cd

Ag

Cd(NO3)2

AgNO3

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Cu

Ni

Cu(NO3)2

Ni(NO3)2

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