14.3 Electrolytic Cells - PowerPoint Presentation

Slide1

14.3 Electrolytic Cells

Electrolysis is the process of converting electrical energy into chemical energy.Voltaic cells produce electrical energy and are viewed as exothermic reactions.Electrolytic cells use electrical energy and are viewed as endothermic reactions.Slide2

Electrolytic cells require a flow of electrons from an

outside energy source such as from a power supply or battery.Positive Enet values for electrochemical cells indicate that those reactions occur spontaneously.Electrolytic cells have negative Enet values which indicate that a minimum voltage must be applied to force a non-spontaneous reaction to occur.In electrolytic cells the cathode is negative

and the

anode is

positive

.

(L.E.O.P.A. / G.E.R.N.C.)Slide3

Rechargeable Batteries

Nickel-Cadmium (Ni-Cad) BatteriesBy using specific chemicals one can create a battery that can be easily recharged.At the cathode, the SOA is Ni(OH)(s).At the anode, the SRA is Cd(s).The net voltage of each cell is 1.25 V.Acts as a voltaic cell when it discharges.Acts as an electrolytic cell when it recharges.Slide4

Example 1: Aqueous Solutions

Write half reactions and the net reaction. Compute the minimum voltage for the electrolysis of Pb(NO3)2(aq). Slide5

Example 1: Aqueous Solutions

Write half reactions and the net reaction. Compute the minimum voltage for the electrolysis of Pb(NO3)2(aq). SOA OAPb2+(aq) / NO3–(aq) / H2O(l)

not on tables

SRA

GERNC:

2 (

Pb

2+

(

aq

)

+ 2 e

–



Pb

(s)

)

LEOPA:

2 H

2

O

(l)



O

2(g)

+ 4 H

+

(

aq

)

+ 4 e

–

Net:

2

Pb

2+

(

aq

)

+ 2 H

2

O

(l)



2

Pb

(s)

+ O

2(g)

+ 4 H

+

(

aq

)

E

°

cell

=

E

°

r

-

E

°

r

cathode anode

= (-0.13) - (+1.23)

=

E



net

= – 1.36 V



V

min

= + 1.36 VSlide6

Apparatus:Slide7

Power Supply

C(s)Pb(s)

C

(s)

e

–

e

–

V

min

= 1.36 V

Pb

2+

(aq)

NO

3

–

(aq)

Solution becomes acidic.

H

+

(aq)

O

2(g)

Apparatus:Slide8

Example 2: Aqueous Electrolysis

Draw and label an electrolytic cell for a solution of nickel (II) chloride.Slide9

Example 2: Aqueous Electrolysis

Draw and label an electrolytic cell for a solution of nickel (II) chloride. SOA OANi 2+(aq) / Cl –(aq) / H2O(l) RA SRA***But, Cl –(aq)

oxidizes faster than H

2

O

(l)

***

GERNC:

Ni

2+

(

aq

)

+ 2 e

–

 Ni(s)

LEOPA: 2 Cl

–(aq)

 Cl2(g) + 2 e –

Net: Ni 2+(

aq) + 2

Cl

–(aq

)

 Ni(s)

+ Cl2(g)Slide10

E

°cell = E°r - E°r cathode anode = (-0.26) - (+1.36) Enet = – 1.62 V  Vmin = + 1.62 VSlide11

Power SupplySlide12

C

(s)Cl2(g)bubblesproduced

C

(s)

e

–

e

–

V

min

= 1.62 V

Ni

2+

(aq)

Cl

–

(aq)

Ni

(s)

electroplating

Anode

Cathode

Power SupplySlide13

Example 3: Electrolysis of Water

Slide14

Example 3: Electrolysis of Water

SOA = H2O(l) SRA = H2O(l)GERNC: 2 (2 H2O(l) + 2 e –  H2(g) + 2 OH –

(

aq

)

)

LEOPA:

2 H

2

O

(l)



O

2(g)

+ 4 H

+

(aq) + 4 e

– Redox: 6 H

2O(l)  2 H

2(g) + O2(g) + 4 H +(aq

) + 4 OH –

(aq)

4 H2

O(l)

Simplify the waters:

2 H2O(l)

 2 H

2(g) + O2

(g) Slide15

E

°cell = E°r - E°r cathode anode = (-0.83) - (+1.23) Enet = – 2.06VTherefore the minimum voltage needed to force this non-spontaneous reaction to occur is +

2.06

 V.Slide16

Electrolysis of Molten Compounds

When ionic compounds are melted the "melt" contains liquid ions.The "melt" is a good electrical conductor and can undergo electrolysis.No water is present.Slide17

Example 1: Write half and net equations and draw a labeled diagram for the electrolysis of molten zinc chloride.

Species present: Zn 2+(l) / Cl –(l) cation anionGERNC: Zn 2+(l) + 2 e –  Zn(l) LEOPA: 2 Cl –(l) 

Cl

2(g)

+ 2 e

–

Net:

Zn

2+

(l)

+ 2 Cl

–

(l)



Zn

(l) + Cl2(g) Slide18

Anode

Power Supply

The "Melt"

Zn

2+

(l)

Cl

–

(l)

Zn

(l)

Cathode

e

–

e

–

High Temp Source

Cl

2(g)Slide19

Example 2: Describe the electrolysis of molten gallium oxide.

Species present: Ga 3+(l) / O 2–(l)GERNC: 4 ( Ga 3+(l) + 3 e –  Ga(l) )LEOPA: 3 ( 2 O 2–(l) 

O

2(g)

+ 4 e

–

)

Net:

4 Ga

3+

(l)

+ 6 O

2–

(l)



4 Ga

(l) + 3 O2(g) Slide20

initial volume

Ga

3+

(l)

O

2–

(l)

Ga

(l)

Cathode

e

–

e

–

High Temp Source

O

2(g)

Anode

Power SupplySlide21

Why is the electrolysis of molten compounds so important?

The technique discovered by Sir Humphrey Davy became efficient in the late 1800's.If an aqueous solution of a salt containing an oxidizing agent that is a metallic ion weaker than water is electrolyzed, the water will be reduced at the cathode to form hydrogen gas.If a molten salt is electrolyzed, the metallic ion is reduced.Electrolysis permits production of active metals such as:Cs +(l) + e –  Cs

(l)

Li

+

(l)

+ e

–



Li

(l)

Al

3+

(l)

+ 3 e –  Al

(l)Ca 2+(l) + 2 e

–  Ca(l) Slide22

Industrial Electrolysis

1) Chlor-Alkali Electrolytic CellsNaCl(aq) saturated brine is pumped from an underground salt dome at Dow Chemical in Fort Saskatchewan.NaCl(aq) is pumped between huge electrodes.Cathode: 2 H2O(l) + 2 e –  H2(g)

+ 2 OH

–

(

aq

)

Anode:

2

Cl

–

(

aq

)



Cl2(g) + 2 e – Net:

2 H2O(l) + 2 Cl

–(aq)

 H2(g) + 2 OH –(aq

) + Cl2(g)

Na +

(

aq) remains in solution with the

OH –

(aq)

: NaOH(

aq) Slide23
Slide24

Uses

a) Cl2(g) is used to makedisinfectant for drinking waterbleaches ( NaOCl(aq) , Ca(OCl)2(s) plastics (polyvinyl chloride

)

pesticides

(

2,4 - D

)

solvents

(

C

2

Cl

4

- dry-cleaning

)Slide25

b) H

2(g) is used to makeammoniahydrogen peroxideMargarinec) NaOH(s) - Caustic Soda ( lye ) is used to makecellophanepulp and paperaluminumdetergentsSlide26

2) Downs Process

Molten NaCl(l) is electrolyzed.Cathode: Na +(l) + e –  Na(l)

Anode:

2

Cl

–

(l)



Cl

2(g)

+ 2 e

–

Note:

NaCl

(s)

is dissolved in molten CaCl

2(l) to reduce NaCl's melting point (805 C)

.

Cl

2(g)

NaCl

(l)

Na

(l)

Carbon anode

Cathode

CathodeSlide27

Electrolysis of Molten Sodium ChlorideSlide28

Uses

Sodium is cooled to form a solid and used in sodium vapour lamps and as a coolant in some nuclear power reactors.Chlorine is sold for commercial use already discussed.Slide29

3) Hall-

Heroult ProcessBauxite (Al2O3(s)) is dissolved in molten cryolite (Na3AlF6(l)) at about 970 C.Cathode: 4(Al 3+(l) (cryolite) + 3 e

–



Al

(l)

)

Anode:

3(

2

O

2–

(l)

(

cryolite

)  O2(g) +

4 e – )Net: 4 Al

3+ (cryolite) + 6 O 2-

(cryolite) → 4 Al(l) + 3 O2

(g)Overall the reaction is a decomposition of aluminum oxide:

2 Al2O

3

(s) → 4 Al(s) + 3 O

2(g)Slide30
Slide31
Slide32

Refining of Metals

Electrorefining is the process of using an electrolytic cell to obtain high-grade metals at the cathode from an impure metal at the anode.Electrowinning is the process of using an electrolytic cell to reduce metal cations from a molten or aqueous electrolyte at the cathode.Slide33

Electroplating

It is plating of a metal at the cathode of an electrolytic cell.

Example:

Cathode (spoon): Ag

+

(

aq

) + 1e

-

→

Ag(s)

Anode: Ag(s)

→

1e

-

+ Ag

+

(aq) Slide34

Voltaic Cells

Electrolytic Cells

1)

2)

3)

4)Slide35

Porous

Barrieranions (–)cations (+)

e

–

e

–

Cathode

SOA is Reduced

GER

P

C

Anode

SRA is Oxidized

LEO

N

A

e

–

e

–

Anode

SRA is Oxidized

LEO

P

A

Cathode

SOA is Reduced

GER

N

C

anions (–)

cations (+)

D.C. Power Supply

5)

6)Slide36

Voltaic Cells

Electrolytic Cells

1)

Energy Conversion:

Chemical Electrical

Energy Conversion:

Electrical Chemical

2)

Spontaneous

Non Spontaneous

3)

Exothermic

Produces electricity

Endothermic

Absorbs electrical energy from a "power supply"

4)

E



net

= "+"

- internally driven

E



net

= "–" 

V

min

 = + E

- externally drivenSlide37

Porous

Barrieranions (–)cations (+)

e

–

e

–

Cathode

SOA is Reduced

GER

P

C

Anode

SRA is Oxidized

LEO

N

A

e

–

e

–

Anode

SRA is Oxidized

LEO

P

A

Cathode

SOA is Reduced

GER

N

C

anions (–)

cations (+)

D.C. Power Supply

5)

6)

Examples:

a) alkaline cell

b) lead - acid car battery

Examples:

a) at DOW: NaCl

(aq)

Electrolysis

b) ElectroplatingSlide38

Read pgs. 639 – 650

pgs. 640, 644, 645, 649 Practice #'s 1, 2, 3, 5, 6, 12, 13, 14, 16pg. 651 Section 14.3 Questions #’s 4, 9, 15 Homework: