Chapter 12.5 - The Electrolytic cell

In the previous section, we have seen that electricity can be produced from chemical reactions. Is the reverse possible? That is., Can we make a reaction to occur by passing electricity? In this section, we will see the answer.

1. Take some solid cupric chloride (CuCl₂) in a beaker. 
2. Take two graphite rods. Bring them into contact with the cupric chloride. But the rods should not touch each other. 
3. Connect one rod to the positive terminal of the battery. Connect the other rod to the negative terminal through a volt meter.
4. We can see that there is no passage of current. That means, solid cupric chloride does not conduct electricity.

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Now let us modify the experiment:
• Make the cupric chloride into a solution by adding water. The rest of the arrangements are the same. This is shown in fig.12.10 below:

• In this arrangement, the voltmeter will show that current is passing through the circuit. 
So how does aqueous solution of cupric chloride conduct electricity?
We will write the answer in steps:
1. When solid CuCl₂ is made into an aqueous solution, it dissociates into Cu⁺² and Cl^-1 ions. 
• So an aqueous solution of CuCl₂ is actually a mixture of Cu⁺² and Cl^-1 ions. 
2. When the switch is turned on, the electrons begin to flow from the negative terminal of the battery to the positive terminal. 
• For that, the electrons first reach the graphite rod which is connected to the negative terminal of the battery. That rod becomes negatively charged. 
• So we can say that the rod connected to the negative terminal of the battery is the negative electrode in this experiment. 
3. The positive Cu⁺² ions, present in the solution will move towards this negative electrode. See fig.12.11 below:

• Each Cu⁺² ion will receive two electrons, and they will become Cu atom. Let us write the equation:
Cu⁺² (aq) + 2e^-1 ⟶ Cu⁰ (s)
• This is a reduction reaction. Because, the oxidation number of Cu is reduced from +2 to zero. 
• In other words:
Gaining electrons is reduction. So this is a reduction reaction  
• The newly formed copper atoms stick to the surface of the electrode.
• Also, the number of Cu⁺² ions present in the solution will go on decreasing. So the blue colour of the solution (which is due to the presence of Cu⁺² ions) will fade 
• We have seen earlier (when we discussed galvanic cells) that, the electrode at which reduction takes place is the cathode. 
• So in the present experiment, the negative electrode is the cathode. 
[Recall that, in the galvanic cell, the negative electrode was the anode. Because there, oxidation takes place at the negative electrode]. 
■ It is interesting to note the following:
• In the present experiment: Negative electrode ⟶ reduction takes place ⟶ cathode
• In galvanic cell: Negative electrode ⟶ oxidation takes place ⟶ anode
This is shown in the fig.12.12 below:

■ Note:
• An electrode may be positive or negative
    ♦ If oxidation takes place at that electrode, it is the anode
    ♦ If reduction takes place at that electrode, it is the cathode
4. Now let us continue our discussion:
• The negatively charged Cl^-1 ions in the solution move towards the positive electrode. See fig.12.11 above. 
• They donate their electrons to that electrode. Thus the circuit is completed. Let us write the equation:
2Cl^-1 (aq) ⟶ 2Cl ⁰ (s) + 2e^-1
• This is a oxidation reaction. Because the oxidation number of Cl increases from -1 to zero. 
In other words:
Losing electrons is oxidation. So this is an oxidation reaction.
• So chlorine is formed at this electrode. They come out as bubbles from near this electrode. 
• We have seen that, the electrode at which oxidation takes place is the anode. 
• So in the present experiment, the positive electrode is the anode. [Recall that, in the galvanic cell, the positive electrode was the cathode. Because there, reduction takes place at the positive electrode]. 
■ It is interesting to note the following:
• In the present experiment: Positive electrode ⟶ oxidation takes place ⟶ anode
• In galvanic cell: Positive electrode ⟶ reduction takes place ⟶ cathode
This is shown in fig.12.13 below:

5. From (3) and (4) we can say that:
■ Both in the present reaction and in the previous galvanic cell,
• The electrode at which oxidation takes place is the anode    
• The electrode at which reduction takes place is the cathode
■ But positive and negative electrodes are interchanged:
• In the present reaction, anode is the positive electrode and cathode is the negative electrode
• In the previous galvanic cell, anode is the negative electrode and cathode is the positive electrode

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• So we saw how the circuit is completed and current is allowed to pass, when the solid is made into an aqueous solution. 

• The passage of current was possible because ions acquired mobility in the aqueous state. 
    ♦ In solid state, ions are held firmly in position. They cannot move. 
• This information tells us about another possibility:
If the cupric chloride is heated to a molten state, then also, the ions will acquire mobility. So the circuit will be completed. 
• That means, instead of aqueous solution, we can use the molten state of the salt also, for passing electricity.
■ Electrolytes are substances which conduct electricity in molten states or in aqueous solutions and undergo a chemical change. Acids, alkalies and salts are electrolytes in their molten state or in aqueous solution.
■ The process of chemical change taking place in an electrolyte by passing electricity is called electrolysis
■ So now we can define a new type of cell: The electrolytic cell:
In an electrolytic cell, external electricity is allowed to pass through an electrolyte through two electrodes, as a result of which, a chemical reaction takes place in the electrolyte.
• In other words, an electrolytic cell is an apparatus used to carry out electrolysis

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The following table 12.2 shows the difference between Galvanic cell and Electrolytic cell

	GALVANIC CELL	ELECTROLYTIC CELL	Remarks
1	Redox reaction produces electricity	Electricity produces redox reaction
2	Oxidation takes place at anode	Oxidation takes place at anode	Whatever be the type of cell, anode is where oxidation takes place
3	Reduction takes place at cathode	Reduction takes place at cathode	Whatever be the type of cell, cathode is where reduction takes place
4	Anode is the negative electrode	Anode is the positive electrode
5	Cathode is the positive electrode	Cathode is the negative electrode

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In the next section, we will see electrolysis of water. 

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Chapter 12.3 - The Zinc - Copper Galvanic cell

In the previous section, we saw reason for the displacement of silver from a solution of silver nitrate. In this section, we will see how metals can be used to make electricity.

1. Take two beakers. Add 100 mL of 1 M zinc sulphate solution into one beaker. 
2. Add 100 mL of copper sulphate solution into the other. See fig.12.5 below:
[We have seen the method to prepare 1 M solution of any given salt. See details here]

Fig.12.5

3. Immerse a zinc rod in the zinc sulphate solution. 
4. Immerse a copper rod in the copper sulphate solution. 
5. Connect the negative terminal of a voltmeter to the zinc rod. Connect the positive terminal to the copper rod. 
6. Connect the solutions in the two beakers by a salt bridge. A long strip of filter paper soaked in KCl solution can be used instead of the salt bridge.

■ Method of making a salt bridge: • Make a paste by mixing two items: (i) gelatin (ii) Potassium chloride (KCl)     ♦ agar agar gel can be used instead of gelatin     ♦ Potassium nitrate (KNO3) or Ammonium chloride (NH4Cl) can be used instead of KCl • Fill this paste into an U-tube • Close the ends of the tube with cotton balls • The salt bridge is ready to use. We will see it's function later in this section

7. Now observe the change in the voltmeter reading.
From the voltmeter, it is clear that electricity is produced in the experiment. Let us see the reason:
8. In the previous experiments we have seen that zinc is more reactive than copper. 
So Zn in the zinc rod, loses two electrons and become. Zn²⁺. The equation can be written as:
Zn⁰ (s) ⟶ Zn⁺² (aq) + 2e^-1
• This is a oxidation reaction. Because the oxidation number of Zn increases from zero to +2.
• In other words:
Losing electrons is oxidation. So an oxidation reaction is taking place here.
■ The electrode at which oxidation takes place is called the anode.
9. The newly formed Zn⁺² ions gets detached from the zinc rod and goes into the solution. But the released electrons stick to the Zn rod
• Note that many Zn⁺² ions are already present in the zinc sulphate solution because, the aqueous zinc sulphate solution exists as a mixture of Zn⁺² ions and (SO₄)^-2 ions.
10. Because of the released electrons, the Zn rod becomes negatively charged. These free electrons reach the copper rod through the external circuit
11. These electrons which reach the copper rod flows through the copper rod and reaches the copper sulphate solution. 
• In the copper solution, Cu⁺² ions are present. These ions receive the electrons and become Cu atoms. The equation can be written as:
Cu⁺² (aq) + 2e^-1 ⟶ Cu⁰ (s)
• This is a reduction reaction. Because the oxidation number of Cu⁺² decreases from +2 to zero.
• In other words:
Gaining electrons is reduction. So a reduction reaction is taking place here.
■ The electrode at which reduction takes place is called the cathode.
12. So both oxidation and reduction takes place in this reaction. Thus it is a redox reaction.
• The transfer of electrons produced by the redox reaction causes the flow of electric current.
• If a flow of electric current is obtained from a device, we can call it a cell.
• The cell which converts chemical energy to electrical energy through redox reaction is called Galvanic cell or Voltaic cell.
• They are named after the scientists Luigi Galvani and Alessandro Volta who achieved early developments in such cells.

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In such cells, we can see two interesting facts:

Fact 1:
• Out of the two metals zinc and copper, zinc is more ready to donate electrons. 
• So oxidation takes place at the zinc rod. 
• We have seen that, the electrode at which oxidation takes place is called the anode.
So we can say this:
■ In the galvanic cell, the more reactive metal will become the anode
The opposite can also be written:
■ In a galvanic cell, the less reactive metal will become the cathode
Fact 2:
• We have seen that the free electrons are first formed when the Zn atoms become Zn⁺² ions. 
• This happens at the anode. 
■ Thus in a galvanic cell, the electrons flow is from the anode to cathode
The above facts are shown in the fig.12.6 below: 

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Corrosion of metals

At this stage, we can have a short discussion about the basics of 'corrosion of metals'. 

• We have seen that the zinc atoms become Zn⁺² ions and leave the zinc rod. If the process continues, the zinc rod will soon become useless.

• The zinc atoms left the rod because favorable conditions were available on the surface of the zinc rod.

    ♦ We provided those favorable conditions by connecting it with a copper rod, providing zinc sulphate and copper sulphate solutions etc.,

• If such favorable conditions occur naturally around the surface of any metal, it's atoms would surely leave. 

• For example, the oxygen and water vapour present in the atmospheric air can provide favorable conditions for certain metals. 

• The atoms of those metals will then become ions and will leave the original metal surface. 
• The released electrons will be received by oxygen. The oxygen become negative ions. 
• The 'positive metal ions' and 'negative oxygen ions' together will form a new compound. The rust which we see on the surface of iron is such a compound.           

■ The process of conversion of a metal into it's compounds by continuous interaction with atmospheric air and water vapour is termed corrosion of metals. This is an electrochemical reaction. 
• Metals like potassium, sodium, zinc, copper etc., also undergo corrosion.

Cathodic protection

• In the above discussion, we have seen that the zinc which is more reactive than copper has become the anode.
    ♦ Atoms of the anode loses electrons and become ions. 
• If in the place of copper, we use a metal which is more reactive than zinc, then that metal will become anode. 
    ♦ Zinc will become the cathode and will be saved. 
• So we must make the metal to be protected, the cathode
• This is the basis of cathodic protection.
■ Cathodic protection is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. In the simplest form, the metal to be protected is connected to a more easily corroded 'sacrificial metal' to act as the anode
• We will see more technical details of this method in higher classes

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Now we will discuss about the salt bridge
• Consider the cathode compartment in the above galvanic cell. 
    ♦ We have a copper rod in a copper sulphate solution. 
• The copper sulphate solution always exists as a mixture of Cu⁺² and (SO₄)^-2 ions. 
• We have seen that, the electrons reaching the copper rod will be received by the Cu⁺² ions, and they will become Cu atoms. 
• So the positive charges in the solution will continuously decrease. In other words, there will be an excess negative charge (due to the (SO₄)^-2 ions) in the solution. 
• So these ions flow from the copper compartment to the zinc compartment through the salt bridge.
• This will help to maintain the electrical neutrality of the solution

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So we have completed this discussion on Galvanic cell using Zinc and copper. In the next section, we will see another Galvanic cell. 

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