The third chapter in the science textbook for the SSLC curriculum is titled Reactivity Series and Electrochemistry. Metals can be arranged using the reactivity series according to how readily they give up electrons. The metal families and transitional elements are discussed in this section along with their relative positions in the reactivity series. Although they are precious metals, gold and platinum are more significant to chemistry students and researchers as electrochemical catalysts for extracting water from carboxylic acid derivatives like esters.
The Kerala SSLC exam syllabus' third chapter was just made public. Reactivity Series and Electrochemistry are the topics of this chapter. This article can assist you better understand the concept if you are preparing for the Kerala SSLC exam. All of the chapter's questions are answered here, along with thorough explanations.
| Board | SCERT, Kerala |
| Text Book | SCERT Based |
| Class | SSLC |
| Subject | Chemistry Solution |
| Chapter | Chapter 3 |
| Chapter Name | Reactivity Series and Electrochemistry |
| Category | Kerala SSLC |
Kerala Syllabus SSLC Class 10 Chemistry Textbook Solution Chapter 3 Reactivity Series and Electrochemistry
- Chapter 1: Periodic Table and Electronic Configuration
- Chapter 2: Gas Laws Mole Concept
- Chapter 3: Reactivity Series and Electrochemistry
- Chapter 4: Production of Metals
- Chapter 5: Compounds of Non-Metals
- Chapter 6: Nomenclature of Organic Compounds and Isomerism
- Chapter 7: Chemical Reactions of Organic Compounds
SSLC Chemistry Part I
SSLC Chemistry Part II
Chapter 3 Reactivity Series and Electrochemistry Textbook Solution
Let us Assess
Let Us Assess
Question 1.The solutions of ZnSO4, FeSO4, CuSO4, and AgNO3 are taken in four different test tubes. Suppose, an iron nail is kept immersed in each one.
• In which test tube the iron nail undergoes a colour change?
• What is the reaction taking place here?
• Justify your answer. (Refer reactivity series of metals).
Answer:
(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.
Question 2.
Compare the electrolysis of molten potassium chloride and solution of potassium chloride. What are the processes taking place at the cathode and the anode?
Answer:
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-
Question 3.
You are given a solution of AgNO3, a solution of MgSO4, an Ag rod and an Mg ribbon. How can you arrange a Galvanic cell using these? Write down the reactions taking place at the cathode and the anode.
Answer:
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.
Extended Activities
Question 1.Keep two carbon rods immersed in a copper sulphate solution. Then pass electricity through the soon.
(i) At which electrode does a colour change occur-anode or cathode?
(ii) Is there any change in the blue colour of the copper sulphate solution?
(iii) Write down the chemical equations for the changes occurring here.
Answer:
(i) The colour change occurs at the cathode. The electrolyte CuSO4 dissociates into Cu2+ and SO42-. 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 (SO42-) 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-
Question 2.
When acidified copper sulphate solution is electrolyzed oxygen is obtained at the anode. What arrangements are to be made for this? Find the element deposited at the cathode.
Answer:
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)
Question 3.
How many Galvanic cells can be made by using the metals Ag, Cu, Zn, and Mg? When Galvanic cells are made using the metals given, what will be the nature of reactions in each cell? (Reactivity: Mg>Zn>Cu>Ag).
Answer:
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.
Question 4.
You are familiar with some materials which are used in various secondary cells.
Make a list of different types of chemical cells. Analyse how they influence the environment.
Answer:
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 ecosystem. 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.
SSLC Chemistry Textbook Solution
- Chapter 1: Periodic Table and Electronic Configuration
- Chapter 2: Gas Laws Mole Concept
- Chapter 3: Reactivity Series and Electrochemistry
- Chapter 4: Production of Metals
- Chapter 5: Compounds of Non-Metals
- Chapter 6: Nomenclature of Organic Compounds and Isomerism
- Chapter 7: Chemical Reactions of Organic Compounds
SSLC Chemistry Part I
SSLC Chemistry Part II
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