Reactivity Series And Electrochemistry

(i) The iron nail undergoes a colour change in that test tube which has CuSO₄ solution.

Since Fe is more reactive than Cu and Ag but less reactive than Zn because in reactivity series Fe is above to Cu and Ag. So, Fe will displace Cu and Ag from their salts. But colour will change in a test tube of CuSO₄ solution from blue to green. Because AgNO₃ is colourless and Fe(NO₃)₂ is also colourless so there is no colour change in test tube of AgNO₃.

(ii) Fe(s) + CuSO₄(aq) FeSO₄(aq)+ Cu(s)

(iii) Since Fe is more reactive than Cu and Ag but less reactive than Zn because in reactivity series Fe is above to Cu and Ag. So Fe will displace Cu and Ag from their salts. But colour will change in test tube of CuSO4 solution from blue to green. Because AgNO3 is colourless and Fe(NO3)2 is also colourless so there is no colour change in a test tube of AgNO3.

On electrolysis of molten potassium chloride, potassium chloride dissociates into K⁺ and Cl^-. The anode is connected to positive terminal and cathode to the negative terminal of the battery. So, the oxidation occurs at anode and reduction occurs at the cathode. At the cathode, there are only K⁺ ions so they reduced and potassium will deposit on the cathode. And at the anode, there are only Cl^- ions so they will oxidized and chlorine gas will evolve at the anode.

At anode- 2Cl^-(aq) Cl₂(g) + 2e^-

At cathode- K⁺(aq) + e^- K(s)

On electrolysis of a solution of potassium chloride, potassium chloride dissociates into K⁺ and Cl^-. The anode is connected to positive terminal and anode to the negative terminal of the battery. So, the oxidation occurs at anode and reduction occurs at the cathode. At the cathode, there are K⁺ ions and water. So, while comparing K⁺ ions and water, water has greater tendency to get reduced. Hence hydrogen gas is liberated at the cathode. Similarly, at anode, there are Cl^- ions and water. So, on comparing Cl^- ions, and water, oxidation occurs to Cl^-. Hence chlorine gas is liberated at the anode.

At anode- 2Cl^- Cl₂ + 2e^-

At cathode- 2H₂O + 2e^- H₂ + 2OH^-

First we take 100 ml solution of MgSO₄ in one beaker and 100 ml solution of AgNO₃ in the second beaker. Immerse Mg ribbon in MgSO₄ solution and Ag rod in AgNO₃ solution. Connect the negative terminal of a voltmeter to the Mg ribbon and positive terminal to the Ag rod. Connect the solutions in two beakers by a salt bridge.

Oxidation will occur at anode and reduction will occur at the cathode.

Reaction at anode - Mg(s) Mg²⁺(aq) + 2e^-

Reaction at cathode - Ag⁺(aq) + e^-Ag(s)

Mg(s) + 2Ag⁺(aq) Mg²⁺(aq) + Ag(s)

Note-Salt bridge is a U-tube filled with a paste made by mixing gelatin or agar gel and a salt like KCl or KNO₃. This completes the circuit by transfer of ions and maintains the electrical neutrality of the cell.

(i) The colour change occurs at the cathode. The electrolyte CuSO₄ dissociates into Cu²⁺ and SO₄^2-. So at cathode Cu²⁺converts into Cu which deposited on the cathode and the colour will change at the cathode.

Cu²⁺(aq) + 2e^- Cu(s)

(ii) Yes, the intensity of the blue colour of CuSO₄ will decrease as the copper ions are converted to the copper deposit on the cathode.

(iii) Reaction at cathode: Cu²⁺(aq) + 2e^- Cu(s)

Reaction at anode: The negative sulphate ions (SO₄^2-) and the hydroxide ions (OH^-) are attracted to the positive electrode. But the sulphate ion is too stable and nothing happens. So either hydroxide ions or water molecules are oxidised to form oxygen.

4OH^-(aq) 2H₂O(l) + O₂(g) + 4e^-

Or 2H₂O(l) 4H⁺(aq) + O₂(g) + 4e^-

Take copper sulphate (CuSO₄) in a beaker. Add some acid (H₂SO₄) to it and make a copper solution. Immerse two carbon rods in a beaker. Connect a battery to the rods. Oxidation occurs at anode and reduction occurs at the cathode. And we can observe that oxygen is produced at the anode.

4OH^-(aq) 2H₂O(l) + O₂(g) + 4e^-

Or H₂O(l) 4H⁺(aq) + O₂(g) + 4e^-

And at cathode, copper will deposit due to conversion of Cu²⁺ ions to Cu.

Cu²⁺(aq) + 2e^- Cu(s)

In Galvanic cell, the anode is always more reactive than the cathode.

Total 6 Galvanic cells can be made.

(i) Mg-Zn – Anode = Mg and Cathode = Zn

(ii) Mg-Cu – Anode = Mg and Cathode = Cu

(iii) Mg-Ag – Anode = Mg and Cathode = Ag

(iv) Zn-Cu – Anode = Zn and Cathode = Cu

(v) Zn-Ag – Anode = Zn and Cathode = Ag

(vi) Cu-Ag – Anode = Cu and Cathode = Ag

All the reactions in each Galvanic cell are redox in nature because oxidation occurs at anode and reduction occurs at the cathode.

Types of chemical cells-

1- Simple chemical cell

2- Dry cell

(i) Primary cells

a- Leclanche cells

b- Alkaline cells

c- Lithium cells

d- Mercury cells

e- Silver oxide cell

(ii) Secondary cells

a- Nickel-cadmium cell

b- Lithium-ion cell

c- Nickel metal-hydride cell

3- Wet cell

4- Fuel cell

(i) Phosphoric Acid fuel cell (PAFC)

(ii) Proton Exchange Membrane fuel cell

5- Solar cell

6- Electric cell

The effects of chemical cells on the environment are negative. As chemical cells are burned they pollute the air. And when they are thrown into dump areas then their toxic chemicals are absorbed by soil and damaging our natural eco system. These toxic chemicals damages to plant and animal life. These chemical cells also have a large impact on air acidification. The three worst chemical toxins found in chemical cells are lead, cadmium, and mercury. And these have the worst effect on the environment.