Electrolytic Cells When a nonspontaneous redox reaction is made to occur by putting electrical energy into the system The battery energy source acts as a pump pushing electrons into the cathode and removing electrons from the anode ID: 538042
Download Presentation The PPT/PDF document "Electrolysis & Understanding" is the property of its rightful owner. Permission is granted to download and print the materials on this web site for personal, non-commercial use only, and to display it on your personal computer provided you do not modify the materials and that you retain all copyright notices contained in the materials. By downloading content from our website, you accept the terms of this agreement.
Slide1
Electrolysis & Understanding Electrolytic Cells:
When a non-spontaneous
redox
reaction is made to occur by putting electrical energy into the system.
The battery (energy source) acts as a “pump” pushing electrons into the cathode and removing electrons from the anode.
To maintain electrical neutrality, a redox reaction must occur within the cell
consume electrons at the cathode
-
Reduction
liberate electrons at the anode
-
OxidationSlide2
Galvanic –vs-Electrolytic Cells:Slide3
Electrolysis Cell”
Anode
Cathode
Ions present for current to flow
DC voltage-
with high enough voltage, chemical reactions will occur at the two electrodes.Slide4
Electrolysis of molten state
Application: purification of metals
Example:
NaCl
(l) -achieved only at 800°C.
Na
+
attracted to cathode (-) and undergoes reductions.
Cl
-
is attracted to the anode (+) and undergoes oxidation.
2Na+ + 2e-
2Na(l) 2Cl-
2e
-
+ Cl
2
(g)_____
2Na
+
+ 2Cl
-
2Na(l) + Cl
2
(g)Slide5
Electrolysis of Aqueous Solutions
Electrolysis in aqueous solutions also includes the presence of H
2
O which may undergo either oxidation or reduction (depending on energy requirements)
Species present: [Na
+
, Cl
-
, H
2
O]
Possible
Reduction
:
Na
+
(
aq
) + e
-
Na(s) 2H2O(l) + 2e- H2(g) + 2OH-(aq)Possible Oxidation: 2Cl-(aq) Cl2(g) + 2e- 2H2O(l) 2H2(g) + O2(g) + 4e- 2H2O(l) + 2Cl-(aq) H2(g) + 2OH-(aq) +Cl2(g)
Since the process is NOT spontaneous, E must have a net (-) value.Compare E(V) for each half reaction to determine what is occurring at each electrode. This cell is unique when we compare the oxidation of Cl- & H2O Slide6
ElectroplatingSlide7
Electrolysis and ElectroplatingElectric current is passed through a solution containing a salt of the metal to be plated.The object to be plated is the cathode and the metal ion is reduced on its surface.Slide8
Calculations & electroplatingBy knowing the # of moles e- that are required and the current flow/time one is able to calculate the mass of metal plated.
Using a solution containing Ag
+
(
aq
) ions, metallic silver is deposited on the cathode. A current of 1.2A is applied for 2.4 hours. What is the mass of silver formed?
(Useful conversions provided)
Charge: 2.4hrs 3600s 1.12A = 9675.8C
1 hr
Mass of Ag: 9676.8 C 1 mole e- 1 mole Ag(s) 107.9g 96,485C 1 mole e- 1mole AgAnswer: 10.8gSlide9
Useful Relationships:Used to relate electricity through an electrolytic cell and the amount of substances produced by the redox process.
Quantity
Unit
Relationship
Conversion Factor
Charge
Coulomb (C)
1C = 1A·S
= 1J/V
1 mole
e- = 96,480C
CurrentAmpere (A)
1A = 1C/s
Potential
Volt (V)
1V = 1J/C
Power
Watt (W)
1W
= 1 J/s
Energy
Joule (J)1J = 1V·CSlide10
Sample Problem:A current of 2.20A is passed through a solution containing Pb2+
for 2.00 hours, with lead metal being deposited at the cathode. What mass of lead is deposited?
2.00 hr. 60 min. 60 sec. 2.20C 1mole e- 1mole
Pb
(s) 207.2 g
Pb
1 hr. 1 min. S 96,500C 2 mole e
- 1mole Pb
= 17.0g Pb Slide11
Sample Problem:Chromium metal can be electroplated from a water solution of potassium dichromate; the reduction half reaction is:
Cr
2
O
7
2-
(
aq
) + 14H+(
aq) + 12 e-
2 Cr(s) + 7 H2
O(l)How many grams of chromium will be plated by 1.00x104C? ( Strategy: Coulombs
mole e- mole Cr mass Cr)
Ans. = 0.898g Cr