Grade 12 Unit 4
Grade 12 Unit 4
Grade 12 Unit 4
2H2O(l) →H2(g)+O2(g)
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+ -
e– e–
e– e–
Anode Cathode
Cl- H+
H+
Cl-
Na+
OH- 2Cl-
Na+ NaOH OH- Na+
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Learning objective
Mention the industrial application of electrochemistry
Explain electroplating and Electrorefining
Explain how electrolysis is used in the production of some metals, nonmetals and compounds
Focused Questions
1. What are the purposes of electroplating and electrorefining?
2. Where do we put the plating metal and the metal to be plated, during electroplating?
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How it works?
What is electrorefining?
Electrorefining is defined as the process of refining or purification of impure metals by electrolysis.
It is another important application of elctrochemistry.
In an electrorefining process, the anode is the impure metal and the impurities must be lost during the
passage of the metal from the anode to the cathode during electrolysis.
For example, purification of impure copper is one common application of electrorefining method
Cathode reaction:
Cu2+(aq) + 2e- → Cu (s)
Cu2+ SO4-2
Cu2+
During the process, the size of the impure
SO4-2 Cu2+
Cu2+
Cu2+
copper anode decreases and that of the
Impurities SO4-2 pure copper cathode increases.
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Alkali metals, alkaline earth metals and Aluminum are extracted from their compound by
electrolysis.
The method of extraction depends on the reactivity of the metal being extracted.
This can be discovered using the reactivity series.
increasing reactivity
Extracted using
electrolysis
Therefore, electrolytic reduction rather than chemical
reduction is often used to obtain active metals from their
compounds.
Extracted by
reduction with
carbon
Why is it not possible to extract aluminium by heating its ore with carbon?
Aluminium ore (bauxite) has a very high melting point (2050 ºC).
For electrolysis, the ore is dissolved in a compound called cryolite (Na3AlF6), which lowers the
melting point to 700 ºC.
What redox processes occur at the electrodes during the electrolysis of aluminium oxide (Al2O3)?
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1. Fluorine (F2)
Fluorine is prepared by the electrolysis of a molten mixture of KF and HF.
Anode: oxidation
2F- (aq) F2 (g) + 2e- Overall cell reaction:
Cathode: reduction 2F- (aq) + H+ (aq) F2 (g) + H- (aq)
H+ (aq) + 2e- H- (aq)
At the Anode:
anions OH- and Cl- move toward the anode
The position of OH- ions in the electrochemical series is higher
than that of Cl- ion.
Hydroxide ions are preferentially discharged (oxidized) to form
oxygen gas.
4OH- (aq) 2H2O (l)+ O2 (g) + 4e-
At the Cathode:
Figure:- electrolysis of dilute sodium chloride cations Na+ and H+ moves toward the anode
The position of H+ ions in the electrochemical series is lower
than that of Na+ ions.
Hydroxide ions are preferentially discharged (reduced) form
hydrogen gas.
2H+ (aq) + 2e- 2H2(g)
Overall reaction
4OH- (aq) +4H+ (aq) 2H2O (l)+ O2 (g) + 4H2(g)
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2. Chlorine (Cl2)
Chlorine is manufactured by the electrolysis of a molten sodium chloride or concentrated sodium chloride
(brine) solution.
Draw and label a Zn-Cu galvanic cell, using ZnSO4 and CuSO4 solutions
Identify the cathode, anode and electrolyte of a given Galvanic cell
Calculate the electrode potentials and cell potentials using Nernst equation
Compare and contrast electrolytic and Galvanic cells
This means the reaction will proceed on its own without any external energy or
influence.
In galvanic cell, the redox reaction involves the transfer of electrons from reducing
agent to oxidizing agents.
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Focused Questions
1. What are the anode and the cathode in a Zn-Cu Galvanic cell?
2. Write the half-reactions at each half-cells of a Zn-Cu Galvanic cell .
3. How does the salt bridge maintain the electroneutrality of a solution?
4. What makes a Galvanic cell different from an electrolytic cell?
Such a cell was invented by the British Chemist John Daniell in 1836.
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3. Electrolytes
are substances that conduct electricity, either in an
aqueous solution or in a molten state
–V+ Cathode
Anode e– e–
Salt bridge
KCl
Zn electrode Cu electrode
Zn2+ Zn
Left hand Right hand
Cl– K+ compartment
compartment
Cu Cu2+
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Electrons flow from negative electrode to Electrons flow from positive electrode to
Direction of
positive electrode negative electrode
electricity
(- ive to + ive) (+ ive to - ive )
Electrolyte are kept separate from one another, and are The cathode and anode are in the same
solution(s) connected by a salt bridge electrolyte
Cell Potential Ecell > 0 and DG < 0 Ecell < 0 and DG > 0
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Exercise 4.8
Q1. Write the cell notation for a Galvanic cell consisting of an Al electrode placed in 1 M
Al(NO3)3 solution and a Ag electrode placed in a 1 M AgNO3 solution.
Answer
Q2. Write the anode and cathode half-cell reaction of the following of galvanic cells
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What is the difference between standard reduction potential and cell potential?
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• Eºred is the electrode potential measured under standard condition relative to the Standard
hydrogen electrode.
Anode
• The electrode chosen as a reference is called the standard hydrogen electrode (SHE).
•
• The reduction half-reaction chosen as the reference is
2H+ (aq) + 2e- H2 (g), Eº= 0.00 V
0.76 V
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More positive Eº
More positive Eº
More easily electron is gained
More negative Eº
More easily electron is lose
More easily oxidized
Better reducing agents
Example:-
More negative Eº
More positive Eº
More easily electron is gained
Increasing Strength as Oxidizing agents
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Justification
1. Under standard state conditions, any species on the left of a given cell reaction reacts spontaneously
with a species that appears on the right of any half-cell reaction
2. The more positive the reduction potential, the greater the tendency to accept electrons. As a result, when two half-cells
are coupled, the reaction with higher (more positive) reduction potential proceeds as reduction, while the other proceeds
as oxidation.
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We use batteries throughout our day-to-day lives. Cell phones use lithium-ion batteries, cars use lead-acid batteries
(Figure 13.1 (left)), while silver-oxide batteries (Figure 13.1 (top right)) are used in watches. Some batteries are
rechargeable (Figure 13.1 (bottom right)), while others cannot be recharged and have to be thrown away.
A battery consists of multiple electrochemical cells. And within each cell there are electrochemical reactions taking
place. This can be seen in the lemon battery experiment shown in Figure 13.2. Each lemon is a cell in the battery, which
consists of three lemon cells. The reactions the copper and zinc undergo in the lemons are electrochemical reactions,
and a current is produced.
Figure 13.2: A current is produced by connecting lemons with zinc and copper metal. The reactions taking place here are elect rochemical reactions.
Electrochemical reactions, and electrochemical cells are covered in this chapter. Before going into any more detail
however, it is important to revise oxidation and reduction, as well as redox reactions and how to balance them, as these
concepts are very important in electrochemistry.
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