Class 10 Chemistry Chapter 5 Notes Kerala Syllabus Electrochemistry Questions and Answers

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SSLC Chemistry Chapter 5 Notes Questions and Answers Pdf Electrochemistry

SCERT Class 10 Chemistry Chapter 5 Electrochemistry Notes Pdf

SSLC Chemistry Chapter 5 Questions and Answers – Let Us Assess

Question 1.
Examine the diagram of a galvanic cell and find out whether the given statements are true or false.

a) As the cell operates, the mass of the zinc rod, Zn(s) decreases.
b) The copper electrode is the anode.
c) Electrons flow from the zinc electrode to the copper electrode through the external circuit.
d) During cell reaction, reduction takes place at the copper electrode.
e) During cell reaction, the concentration of Cu2+ decreases.
Answer:
a) True
b) False: Copper is the cathode (site of reduction). Zinc is the anode.
c) True
d) True
e) True

Question 2.
The chemical equation of the reaction taking place in a galvanic cell is given.
Zn(s) + 2Ag⁺(aq) → Zn²⁺(aq) + 2Ag(s)
a) Represent diagrammatically the above galvanic cell.
b) Which electrode has a negative charge?
c) What is the reaction taking place at each electrode?
Answer:
a)
b) The zinc electrode is negatively charged.
c) Anode: Zn(s) → Zn²⁺ + 2e^– [Oxidation]
Cathode: 2Ag⁺ + 2e^– → 2Ag(s) [Reduction]

Question 3.
A galvanic cell is made up of the following half cells:
• A magnesium electrode immersed in magnesium sulphate solution.
• A nickel electrode immersed in nickel sulphate solution.
The reactions taking place in the half-cells are given.
Mg (s) → Mg²⁺ (aq) + 2e^–
Ni²⁺(aq) + 2e^– → Ni(s)
Draw the cell diagram and label the anode, cathode and direction of flow of electrons.
Answer:

Question 4.
When an aqueous sodium chloride solution is electrolysed, the reaction that takes place at the anode is:
(i) Na⁺(aq) + e^– → Na(s)
(ii) 2H₂O → 4H⁺(aq) + O₂(g) + 4e^–
(iii) H⁺ (aq) + e^– → 1/2H₂(g)
(iv) Cl^– (aq) → 1/2 Cl₂ +e^–
Answer:
The reaction at the anode is:
Cl^– (aq) → Cl₂(g) + e^–

Question 5.
Write the products obtained when the following solutions undergo electrolysis.
(i) Aqueous solution of NaCl
(ii) Aqueous solution of CuSO₄ (using copper electrodes)
Answer:
(i) At Cathode (-): Water is reduced → Hydrogen gas (H₂) is released.
2H₂O + 2e^– → H₂ + 2OH^–
At Anode (+): Chloride ions are oxidised → Chlorine gas (Cl₂) is released.
2Cl → Cl₂ + 2e^–
Products: Hydrogen gas, Chlorine gas, and Sodium hydroxide (NaOH in solution).

(ii) At Cathode (-): Copper ions are reduced → Copper metal is deposited.
Cu²⁺ + 2e^– → Cu
At Anode (+): Copper electrode dissolves (oxidation).
Cu → Cu²⁺ + 2e^–
Products: Copper is deposited on the cathode and at anode copper dissolves as ions.

Question 6.
Brass is an alloy of zinc and copper. When brass comes into contact with saline water, corrosion of metal takes place and zinc gradually dissolves in the solution, leaving copper behind. Explain why zinc dissolves in comparison to copper.
Answer:
Zinc is more reactive than copper as it lies above copper in the reactivity series. When brass is exposed to saline water (which contains electrolytes), zinc atoms lose electrons more readily and go into solution as Zn²⁺ ions. Copper, being less reactive, does not dissolve as easily and remains behind.

Question 7.
List the cells we use in our daily life and classify them into primary and secondary cells.
Answer:
Primary cells (non-rechargeable): Dry cell (used in torches, clocks, toys), Mercury cell, Lithium cell (in calculators, watches).

Secondary cells (rechargeable): Lead-acid battery (car battery), Nickel-Cadmium cell, Lithium-ion battery (in mobiles, laptops).

Question 8.
Define fuel cells and write their advantages.
Answer:
A fuel cell is a device that converts the chemical energy produced by the combustion of fuel directly into electrical energy.

Advantages:

1. High efficiency.
2. Environment-friendly (produces water as a by-product in hydrogen fuel cells).
3. Continuous supply of electricity as long as fuel is provided.

Chemistry Class 10 Chapter 5 Notes Kerala Syllabus Electrochemistry

Question 1.
a) What did you observe?
Answer:
The bulb glows.

b) What is the reason for the glowing of the bulb?
Answer:
The bulb glows because chemical energy is converted into electrical energy in the electrochemical cell.
We can understand that there is a flow of electric current through the bulb.

Separate the copper wires from the bulb and connect them to a galvanometer as shown in the figure given below. When the copper rod is connected to the positive terminal of the galvanometer and the zinc rod to the negative terminal, it can be seen that the pointer of the galvanometer is deflecting towards the right (towards the positive terminal)

This proves that:
• Electrons flow from the zinc rod to → copper rod.
• Electric current flows in the opposite direction (copper → zinc) through the galvanometer.
Connect a voltmeter instead of the galvanometer as shown in the figure below.

SALT BRIDGE

• A salt bridge is a U-shaped glass tube filled with an inert electrolyte.
• Common electrolytes used:
  • Potassium chloride (KCl)
  • Ammonium chloride (NH₄Cl)
  • Potassium nitrate (KNO₃)
• The electrolyte is made into a gel using agar-agar.
• It connects the oxidation and reduction half-cells in a galvanic cell.
• It allows electrical contact between the two solutions without mixing them.
• Maintains electrical neutrality by allowing ion flow.
• Enables smooth operation of the galvanic cell.

Examine the metal rods and the solutions in which they are immersed in the figure below.

Question 2.
Find the voltage between the two electrodes?
Answer:
The voltmeter shows a reading of about 1.1 volts. This means the voltage (potential difference) between zinc and copper electrodes is approximately 1.1 V.
The reaction that occurs at each electrode is given below.
Zinc electrode: Zn(s) → Zn²⁺ (aq) + 2e^–

Question 3.
What type of chemical reaction takes place at the zinc electrode? (Oxidation / Reduction)
Answer:
Oxidation
Copper electrode: Cu²⁺(aq) + 2e^– → Cu(s)

Question 4.
What type of chemical reaction takes place at the copper electrode?
Answer:
Reduction
Anode: The electrode where oxidation occurs is called the anode
Cathode: The electrode where reduction occurs is called the cathode.
The processes that take place in the Daniel cell can be summarised as:
Anode: Zn(s) → Zn²⁺ (aq) + 2e^–
Copper electrode: Cu²⁺(aq) + 2e^– → Cu(s)
Cell reaction: Zn(s) + Cu²⁺ (aq) → Zn²⁺(aq) + Cu(s)
The electric current is produced as a result of a redox reaction taking place in the galvanic cell.

In the Daniel cell, replace the zinc rod in zinc sulphate with aluminium, silver and magnesium rods (plates) and immerse these metals in their own salt solution. Repeat the experiment by connecting each to the copper rod immersed in copper sulphate.

Question 5.
a) List out the direction of flow of electrons, the equation of reaction at each electrode, the cell voltage, and analyse them.

Answer:

b) Which has the higher value of voltage, Zn – Cu galvanic cell or Al – Cu galvanic cell?
Answer:
Al – Cu galvanic cell has the higher voltage.
The tendency of aluminium to get oxidised in redox reactions is greater than that of zinc. That is, aluminium is more reactive than zinc.
Zn – Cu cell voltage = 0.34 – (-0.76) = 1.10 V
Al – Cu cell voltage = 0.34 – (-1.66) = 2.00 V

c) Which has the higher value of voltage, Mg – Cu galvanic cell or Al-Cu galvanic cell?
Answer:
The voltage value for the Mg-Cu galvanic cell is higher.

d) What is the direction of flow of electrons in a galvanic cell made up of a pair of copper and silver electrodes?
Answer:
Cu acts as anode and Ag acts as cathode. So, Electrons flow from copper to silver.

Question 6.
Draw a diagram of the galvanic cell having silver and copper as electrodes and mark the direction of flow of electrons.
Answer:

CELL VOLTAGE
The cell voltage is the potential difference between the two electrodes of a cell when no current is flowing.

Cell Voltage (EMF) = Potential of Cathode – Potential of Anode

Question 7.
Prepare a list of the metals, familiarised through the experiment, in the order of their tendency to get oxidised in redox reactions.

Answer:

In this way, we can construct galvanic cells using different elements as electrodes and compare their reactivities.

REACTIVITY SERIES.
The series in which elements are arranged in decreasing order of their reactivity is known as the reactivity series.

APPLICATION OF THE REACTIVITY SERIES

1. To identify which metal is being displaced
Place a zinc rod in a beaker of copper sulphate solution. Observe for some time.
a) Did the colour of the CuS04 solution change?
Answer:
Yes, the blue colour of the CuS04 solution faded and eventually disappeared.

b) What change was observed in the zinc rod?
Answer:
The zinc rod became coated with a reddish- brown deposit.

c) What may be the reason for the fading of the colour of the solution?
Answer:
A decrease in the concentration of Cu2+ ions is the reason for the fading of the blue colour.

d) With the help of the reactivity series, find out which metal is more reactive, Zn or Cu?
Answer:
Zinc (Zn) is more reactive than copper (Cu).

e) What are the ions fonned when CuSO₄ dissolves in water?
Answer:
Cu²⁺ (copper ions) and SO₄^2- (sulphate ions).

f) What changes have occurred to the Cu²⁺ ions?
Answer:
The Cu²⁺ ions have gained two electrons and have been converted to solid copper atoms (Cu). This process is called reduction.

g) Write the chemical equation of the reaction.
Answer:
Zn + Cu²⁺SO₄^2- → Zn²⁺SO₄^2- + Cu

h) What changes have occurred to Zn?
Answer:
The zinc (Zn) atoms have lost two electrons and have been converted into zinc ions (Zn²⁺). This process is called oxidation.

i) Which metal is displaced here?
Answer:
In the reaction between zinc (Zn) and copper sulfate (CuSO₄), copper (Cu) is the metal being displaced. This is a single displacement reaction, where the more reactive metal, zinc, displaces the less reactive metal, copper, from its compound.

Take a little silver nitrate solution in a beaker and dip a copper rod in it. Observe for some time.
a) Is there any change in the colour of the solution?
Answer:
Yes, the solution gradually turns blue because of the formation of Cu²⁺ ions.

b) What change was observed in the copper rod?
Answer:
A greyish-white layer of silver metal gets deposited on the copper rod.

c) Write the equation of the chemical reaction taking place here.
Answer:
Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag

d) Explain the reaction by comparing the positions of Ag and Cu in the reactivity series.
Answer:
In the reactivity series, copper (Cu) is above silver (Ag), which means copper is more reactive than silver.

e) Complete this chemical equation by assigning oxidation numbers.
Answer:

f) Which ion is responsible for the change in the colour of the solution?
Answer:
The Cu²⁺ ion is responsible (makes the solution blue).

g) Which metal is oxidised here?
Answer:
Copper (Cu) is oxidised.

h) Complete the equation for oxidation.
Answer:
Cu → Cu²⁺ + 2e^–

i) Which one is reduced here?
Answer:
Silver ion (Ag⁺) is reduced to metallic silver (Ag).

j) Write the equation for reduction.
Answer:
Ag⁺ + e^– → Ag

The more reactive metal displaces the less reactive metal from its salt solution.

Question 8.
Silver nitrate solution cannot be stored in a copper vessel. Why?
Answer:
Copper (Cu) is more reactive than silver (Ag) in the reactivity series. So, copper will displace silver from the silver nitrate solution:
Cu + 2AgNO₃ → Cu (NO₃)₂ + 2Ag
As a result, the copper vessel will gradually dissolve (forming blue copper nitrate solution), and silver will get deposited. This reaction would spoil both the vessel and the solution.

2. To identify the oxidising agent and the reducing agent
Question 9.
Consider the displacement reaction given below.

a) Which metal is oxidised here?
Answer:
Magnesium (Mg)

b) Write the equation for oxidation.
Answer:
Mg → Mg²⁺ + 2e^–

c) Which metal is reduced here?
Answer:
Copper (Cu)

d) Write the equation for reduction.
Answer:
Cu²⁺ + 2e^– → Cu

e) Which reaction occurs to the more reactive metal? (Oxidation/Reduction)
Answer:
Oxidation

f) Which reaction occurs to the less reactive metal?(Oxidation/Reduction)
Answer:
Reduction

The more reactive metal gets oxidised while the ion of the less reactive metal gets reduced.

g) What is the oxidising agent and the reducing agent in this reaction?
Answer:
Cu²⁺ is the oxidising agent and Mg is the reducing agent.

The substance that causes oxidation is called an oxidising agent and the substance that causes reduction is called a reducing agent.

h) What happens to the oxidising agent in this redox reaction?
Answer:
The oxidising agent gets reduced

i) What happens to the reducing agent?
Answer:
The reducing agent gets oxidised.

Question 10.
Identity the oxidising agent and reducing agent in the reaction Given below.
2AgNO₃ + Cu → Cu(NO₃)₂ + 2Ag
Answer:
Oxidising agent: AgNO₃ (because Ag+ is reduced) Reducing agent: Cu (because it is oxidised)

In redox reactions, the more reactive metal acts as reducing agent and the less reactive metal acts as an oxidising agent.

3. To identify the displacement of hydrogen from the acid
Take equal amounts of dilute HCl in different test tubes as shown in Figure. Treat equal masses of polished Fe, Mg, Cu, Pb, Zn with dilute hydrochloric acid.

a) In which of the test tubes is hydrogen gas produced?
Answer:
Hydrogen gas is produced in the test tubes containing:

• Fe (Iron)
• Mg (Magnesium)
• Pb (Lead)
• Zn(Zinc)

Not produced in the Cu (Copper) test tube, because copper is below hydrogen in the reactivity series and cannot displace hydrogen from dilute HCl.

b) Is the rate of formation of hydrogen gas the same in all the test tubes?
Answer:
No, the rate of formation of hydrogen gas is not the same in all test tubes.
It depends on the reactivity of the metal with dilute HCl.
Mg > Zn > Fe > Pb > Cu (no reaction)

c) Examine the position of hydrogen and the metals used here in the reactivity series and record your observations.
Answer:
Metals above hydrogen in the reactivity series
(Mg, Zn, Fe, Pb) can displace hydrogen from dilute acids and produce hydrogen gas.
Copper (Cu) is below hydrogen in the reactivity series, so it cannot displace hydrogen from dilute HCl. Hence, no reaction is observed.

The metals placed above hydrogen in the reactivity series can displace hydrogen from dilute acids.

Question 11.
Which of the following metals can displace hydrogen from hydrochloric acid?
[Sodium, gold, silver, aluminium]
Answer:
Sodium (Na) and Aluminium (Al) can displace hydrogen from hydrochloric acid because they are above hydrogen in the reactivity series.

Gold (Au) and Silver (Ag) are below hydrogen in the reactivity series, so they cannot displace hydrogen from hydrochloric acid.

Question 12.
a) Bring a burning candle close to the mouth of the test tubes and record the observations.
Answer:
When you bring a burning candle to the test tube near the cathode, a “pop” sound will be heard, and the gas will ignite. This indicates the presence of hydrogen gas.

When you bring the burning candle to the test tube near the anode, the flame will bum more brightly. This indicates the presence of oxygen, as it supports combustion.

b) Which gas is produced in each test tube?
Answer:
At the cathode (negative electrode): Hydrogen gas (H₂) is produced. This is the test tube with the larger volume of gas.

At the anode (positive electrode): Oxygen gas (O₂) is produced. This is the test tube with the smaller volume of gas.

c) Write the chemical equation of the reaction.
Answer:
The overall chemical equation for the electrolysis of water is:

Here, acidified water decomposes into hydrogen and oxygen. This is because of the flow of electricity from the external electrical source into the solution through the electrodes.

Observe the given figure

Carry out this electrolysis with the help of the teacher.
Question 13.
Record your observations.
a) Which metal is connected to the positive terminal of the battery?
Answer:
Copper

b) Which metal is connected to the negative terminal of the battery?
Answer:
Iron

c) Which solution is used as the electrolyte?
Answer:
Copper sulphate solution

d) What are the ions present in the electrolyte?
Answer:
The ions present are copper cations (Cu²⁺) and sulphate anions (SO₄^2-).

e) Which electrode is connected to the positive terminal of the battery? Anode/Cathode
Answer:
Anode

f) What is the reaction taking place at the anode?
Answer:
Oxidation

g) Complete the equation of the oxidation reaction taking place at the copper plate.
Answer:
Cu → Cu²⁺ + 2e^–

h) Which electrode is connected to the negative terminal of the battery?
Anode/Cathode
Answer:
Cathode

i) What is the reaction taking place at the cathode?
Answer:
The reaction taking place at the cathode is reduction.

j) Which ions will be attracted to the cathode, the iron ring, from the solution?
Answer:
Copper ions (Cu²⁺) will be attracted to the cathode (the iron ring) from the solution.

k) Complete the equation of the reduction reaction taking place here.
Answer:
Cu²⁺ + 2e^– → Cu

The reactions taking place at the anode and cathode:
Anode: Cu → Cu²⁺ + 2e^–
Cathode: Cu²⁺ + 2e^– → Cu
That is, at the anode, copper undergoes oxidation and enters into the solution as Cu²⁺ ions. At the same time, Cu²⁺ ions from the solution are reduced on the surface of the iron ring, which is the cathode, forming a thin coating of copper on it.

Electroplating is the process of coating a layer of one metal onto another metal through electrolysis.

Question 14.
What are the benefits of electroplating?
Answer:
It enhances the beauty of the metal and prevents its corrosion.

Question 15.
Explain the process of electroplating.
Answer:

• The metal that is to be plated should be connected to the positive terminal of the battery.
• The object on which the plating is to be done should be connected to the negative terminal of the battery.
• A salt solution of the metal to be plated is used as the electrolyte

The metals to be coated	Electrolyte
Silver	Silver nitrate solution/Sodium cyanide + Silver Cyanide solution
Gold	Sodium cyanide + Gold cyanide solution

Std 10 Chemistry Chapter 5 Notes – Extended Activities

Question 1.
The following are certain observations about metals A, B, C and D.
(a) When a plate of metal A is placed in a solution containing B²⁺ ions, no reaction is observed.
(b) When the plate of A is placed in a solution containing C⁺ ions, no change occurs.
(c) When a plate of metal D is placed in a solution containing C⁺ ions, a black precipitate of C is formed on the surface of D, and the presence of D²⁺ ions can be detected in the solution.
(d) When a plate of B is placed in a solution of D²⁺ ions, D appears on the surface of B and B²⁺ ions appear in the solution.
List A⁺, B²⁺, C⁺ and D²⁺ in the ascending order of their ability to attract electrons.
Answer:
The tendency to lose an electron varies in the following order.
C < D < B
Therefore, the ascending order of their ability to attract electrons is.
B²⁺ < D²⁺ < C⁺ < A⁺

Question 2.
As shown in the galvanic cell of the given figure, place a silver electrode in a solution of silver nitrate and a lead electrode in a solution of lead nitrate. Connect the two electrodes using a copper wire. Also connect the two breakers using a salt bridge. Next, find the answers to the questions given below.

(a) What is the anode
(b) What is the cathode?
(c) Where does oxidation occur?
(d) Where does reduction occur?
(e) In which direction do electrons flow through the copper wire?
(f) What will be the cell voltage?
(g) Will the cell voltage vary if the two solutions are diluted alternately?
Answer:
a) Anode is where oxidation occurs.
Lead (Pb) electrode

b) Cathode is where reductions occur.
Silver (Ag) electrode

c) At the lead (Pb) electrode

d) At the silver (Ag) electrode

e) From the lead electrode to the silver electrode

f) Standard electrode potentials:

• Pb²⁺/Pb = -0.13 V
• Ag⁺ /Ag = +0.80 V
  Cell Voltage =Potential of Cathode – Potential of Anode = 0.80 V – (-0.13 V) = 0.93 V

g) Yes, the cell voltage will vary if the solutions are diluted. The cell voltage depends on the concentration of the ions. Diluting the solutions changes the ion concentrations, which affects the electrode potential and thus alters the overall cell voltage.

Electrochemistry Class 10 Notes

Electrochemistry Notes Pdf

• Electrochemistry: Study of electricity-related chemical reactions.
• Electrochemical cells: Devices that convert chemical energy to electrical energy or vice versa.
• Types of Electrochemical Cells
  • Galvanic cells: Convert chemical energy to electrical energy.
  • Electrolytic cells: Convert electrical energy to chemical energy.
• Galvanic Cell (Daniell Cell)
  • Oxidation takes place at the anode (Negative electrode).
  • Reduction takes place at the cathode (Positive electrode).
• Electrolytic cell
  • Oxidation takes place at the anode (Positive electrode).
  • Reduction takes place at the cathode (Negative electrode).
• Reactivity Series: Arrangement of elements in decreasing order of reactivity is called the reactivity series.
• Displacement Reactions: More reactive metals displace less reactive metals from salt solutions.
  Zn + CuSO₄ → ZnSO₄ + Cu
• Oxidising & Reducing Agents
  • Oxidising agent: Gains electrons (gets reduced).
  • Reducing agent: Loses electrons (gets oxidised).
• Hydrogen Displacement
  • Metals above hydrogen in the reactivity series displace H₂ from acids.
  • Reactive metals (Mg, Zn, Fe, Pb) react with HCl to release H₂ gas.
  • Unreactive metals (Cu, Ag, Au) do not.
• Electrolysis of molten sodium chloride
  • Ions: Na⁺, Cl^–
  • At Cathode: Na⁺ + e^– → Na (molten)
  • At Anode: 2Cl^– → Cl₂ + 2e^–
  • Products: Na (metal), Cl₂ (gas)
• Electrolysis of aqueous sodium chloride
  • At Cathode: 2H₂O + 2e^– → H₂ + 2OH^–
  • At Anode: 2Cl^– → Cl₂ + 2e^–
  • Products: H₂ gas, Cl₂ gas, NaOH solution
• Electroplating
  • Anode (positive): Metal to be plated (e.g., Cu)
  • Cathode (negative): Object to be plated (e.g., Iron ring)
  • Electrolyte: Salt solution of plating metal (e.g., CuSO₄)
• Benefits
  • Enhances appearance
  • Prevents corrosion
• Applications of electrolysis:
  • Production of metals: Na, K, Ca, Al
  • Production of non-metals: H₂, O₂, Cl₂
  • Manufacture of compounds: NaOH, KOH
  • Electroplating: Jewellery, utensils
  • Purification of metals: Cu, Au
• Different types of cells
  • Primary Cells: Non-rechargeable [Examples: Dry cell, Button cell]
  • Secondary Cells: Rechargeable [Examples: Lead-acid battery, Li-ion battery, Ni-Cd cell]
    Fuel Cells: Fuel cells are galvanic cells that use fuels like H,, CH₄, or CH₃OH to continuously and efficiently convert chemical energy directly into electrical energy in an environmentally friendly manner

INDRODUCTION

In our daily lives, we use many devices, such as vehicles, mobile phones, laptops, and medical equipment. All these works with the help of stored electrical energy from cells or batteries. Chemical reactions inside them produce this energy. Electrochemistry is the branch of chemistry that studies how chemical reactions can produce electricity and how electricity can be used to carry out chemical reactions. The devices that make this possible are called electrochemical cells.

In this unit, we will learn about electrochemical cells, the reactivity series and its applications, electrolytic cells, electrolysis of molten sodium chloride, electrolysis of aqueous sodium chloride solution, electroplating and its uses, and different types of cells

Electrochemical Cells

• Devices that convert chemical energy into electrical energy (or vice versa).
• Made of two electrodes (anode and cathode) and an electrolyte.
• Work on redox reactions where oxidation occurs at the anode and reduction at the cathode.
• Examples: Galvanic (Voltaic) cells and Electrolytic cells.

Electrolysis of Molten Sodium Chloride

• Process of breaking down molten NaCl using electricity.
• At the cathode: Na⁺ ions are reduced to sodium metal.
• At the anode: Cl^– ions are oxidised to chlorine gas.
• Overall: Production of sodium metal and chlorine gas.

Electrolysis of Aqueous Sodium Chloride Solution

• Involves electrolysis of aqueous sodium chloride solution.
• At the cathode: Water is reduced to hydrogen gas.
• At the anode: Cl^– is oxidised to chlorine gas.
• Solution left behind is sodium hydroxide (NaOH).

Electroplating

• Coating one metal with another using electrolysis.
• Cathode: Object to be plated.
• Anode: Metal to be deposited.
• Electrolyte: Salt solution of the coating metal.
• Used for decoration, corrosion protection, and improving durability.

Applications of Electrolysis

• Extraction of metals (e.g., aluminium from alumina, sodium from NaCl).
• Purification of metals (e.g., copper).
• Electroplating (jewellery, utensils, machine parts).
• Production of metals and non-metals.

Different Types of Cells

• Primary Cells: Non-rechargeable [Examples: Dry cell, Button cell]
• Secondary Cells: Rechargeable [Examples: Lead-acid battery, Li-ion battery, Ni-Cd cell]
• Fuel Cells: Generate electricity from a continuous supply of fuel and oxygen (e.g., hydrogen fuel cell).

ELECTROCHEMICAL CELLS
Electrochemistry is the branch of chemistry that deals with the study of the processes that produce electricity through chemical reactions and use electricity to bring about chemical reactions. Electrochemical cells are the devices that make such changes possible. Electrochemical cells can be divided into two types.

1. Galvanic cells
2. Electrolytic cells

GALVANIC CELL
Galvanic cells are devices that convert chemical energy into electrical energy. Daniel cell is an example of a galvanic cell.

METHOD OF CONSTRUCTING A DANIEL CELL
Materials Required:

• Beaker
• Zinc rod
• Copper rod
• Zinc sulphate solution
• Copper sulphate solution
• Salt bridge
• NH₄Cl/KCl/KNO₃
• Copper wires
• Galvanometer
• Voltmeter
• LED bulb

Experimental procedure

• Take two beakers.
  • In one beaker, pour zinc sulphate solution and dip a zinc rod (Half cell 1).
  • In the other beaker, pour copper sulphate solution and dip a copper rod (Half cell 2).
• Make a salt bridge:
  • Roll a piece of filter paper.
  • Add some KCl (Potassium chloride) / NH₄Cl (Ammonium chloride) / KNO₃ (Potassium nitrate) crystals.
  • Sprinkle water to moisten it.
  • Bend it into a ‘U’ shape.
• Place the salt bridge so that it connects the two beakers.
• Connect the zinc rod and copper rod to a bulb using copper wires.

ELECTROLYTIC CELLS
The cells that utilise electrical energy to bring about chemical changes are called electrolytic cells.

EXPERIMENT: ELECTROLYSIS OF produced?
ACIDIFIED WATER
Procedure of the experiment:
Take a plastic cup. Make two holes at the bottom and attach rubber stoppers as shown in the Figure. Insert carbon electrodes (graphite rods from old torch cells can be used for this purpose) into these rubber stoppers. Connect the electrodes to a 6V battery.

Fill the cup with water so that the electrodes are immersed. Add 2 – 3 mL of dilute sulphuric acid. Dilute sulphuric acid is added to increase the ionisation of water and thereby increase its electrical conductivity. Take two test tubes filled with water and invert them over the carbon electrodes. Turn on the switch and leave the experimental setup undisturbed for some time. You will observe bubbles forming at both the electrodes. After some time, remove the test tubes carefully.

ELECTROLYTES
Electrolytes are substances that undergo chemical changes when electricity passes through them.

Electrolytes are substances that conduct electricity in molten state or in an aqueous solution and undergo chemical changes.

ELECTROLYSIS
Electrolysis is the process by which an electrolyte undergoes chemical changes when electricity is passed through it.
Examine the figure that represents the electrolytic process.

Ions are free to move in the liquid state or in an aqueous solution of an electrolyte. These ions are responsible for the electrical conductivity in the electrolyte.

• Acids (like HCl, H₂SO₄, HNO₃, CH₃COOH) contain free ions in aqueous solutions (when mixed with water).
• Alkalies (like KOH, NaOH) also contain free ions in aqueous solutions.
• Salts (NaCl, KCl,.) contain free ions in their aqueous solution and molten state.
  For example, sodium chloride dissociates into Na⁺ and Cf ions in the molten state or in the aqueous solution.

ELECTRODES

• Electrodes are conductors through which electricity enters or leaves the electrolyte.
• During electrolysis:
  • Anode is the electrode connected to the positive terminal of the battery.
  • Cathode is the electrode connected to the negative terminal of the battery.
• Reduction happens at the cathode.
• Oxidation happens at the anode.

ELECTROLYTIC CELL	GALVANIC CELL
Oxidation take place at anode (positive electrode). Reduction take place at cathode (negative electrode).	Oxidation take place at anode (negative electrode). Reduction take place at cathode (positive electrode).

ELECTROLYSIS OF MOLTEN SODIUM CHLORIDE
In the solid state, sodium chloride does not conduct electricity. The reason is that, though the ions can vibrate about their fixed positions, they are not free to move. However, molten NaCl is a good conductor because its ions move freely.

Examine the figure given below.

Two graphite electrodes are connected to a direct current (DC) source through wires. They are dipped in a vessel containing molten sodium chloride. When electricity is passed through them, the following observations can be made.
1. Chlorine gas (Cl₂), having a pale green colour, is liberated at the anode.
2. Silvery molten sodium metal (Na) is formed at the cathode.
a) What are the ions in molten sodium chloride?
Answer:
The ions present are sodium cations (Na⁺) and chloride anions (Cl^–). When sodium chloride (NaCl) is in a molten (liquid) state, its ionic bonds are broken, allowing these ions to move freely.

b) What is the chemical reaction taking place at the anode?[Oxidation/Reduction|
Answer:
Oxidation occurs at the anode (the positive electrode). The chloride anions (Cl^–) lose an electron to form chlorine gas (Cl₂).
Write the equation of the reaction. 2Cl^– → Cl₂ + 2e^–

c) What is the chemical reaction taking place at the cathode? [Oxidation/Reduction]
Answer:
Reduction occurs at the cathode (the negative electrode). The sodium cations (Na⁺) gain an electron to form molten sodium metal (Na).
Write the equation of the reaction. Na⁺ + e^– → Na

d) What are the products obtained as a result of the electrolysis of molten sodium chloride?
Answer:
The products are chlorine gas (Cl₂) at the anode and molten sodium metal (Na) at the cathode.

ELECTROLYSIS OF AQUEOUS SODIUM CHLORIDE
The figure below representing the electrolysis of aqueous sodium chloride solution.

When a suitable voltage is applied across the electrodes of the cell, the following observations can be made.

1. H₂ gas is liberated at the cathode. If phenolphthalein indicator is present in the solution, the solution around the cathode turns pink. That is, the solution becomes basic in nature.

2. Cl₂ gas is liberated at the anode.
At the anode, chloride ions are getting oxidised to Cl₂ gas.
a) Write the equation of the reaction.?
Answer:
2Cl^– → Cl₂ + 2e^–

b) Based on the reactivity series, which has greater oxidising tendency-sodium or hydrogen?
Answer:
Hydrogen has a greater oxidising tendency. In the reactivity series of metals, hydrogen is below sodium, which means it is less reactive.

c) If so, which of the following is getting reduced at the cathode, Na⁺ or H₂O?
Answer:
At the cathode, H₂O molecules are reduced to produce H₂ gas and OH^– ions.
2H₂O + 2e^– → H₂ + 2OH^–
As a result of the processes taking place in the cell gases such as H₂, Cl₂ and an aqueous solution of NaOH are formed. This aqueous solution can
basic in nature. be evaporated to produce solid NaOH.

Electrodes	Chemical change	Product
Anode	2Cl– → Cl2 + 2e–	Chlorine gas
Cathode	2H2O + 2– → H2 + 2OH–	Hydrogen gas

APPLICATION OF ELECTROLYSIS

1. Production of metals
   Metals like sodium, potassium, calcium and aluminium are produced by the electrolysis of some of their compounds.
2. Production of nonmetals
   Electrolysis is used in the industrial production of nonmetals like hydrogen, oxygen and chlorine.
3. Manufacture of compounds
   Electrolysis is used in the production of sodium hydroxide and potassium hydroxide.
4. Electroplating
   Gold plated jewellery, silver plated utensils, chromium plated iron objects etc are produced through electroplating.
5. Purification of metals
   Electrolysis is useful in the purification of metals such as copper, gold etc.

ELECTROCHEMICAL CELLS AS ENERGY SOURCES
A cell is a device that helps to convert the energy released in a chemical reaction directly into electricity. Cells perform two main functions.

1. As portable sources of electrical energy:
   Examples range from the button cells, used in electronic watches, to the lead-acid cells used in vehicles.
2. As storage devices of electrical energy provided by an external source:
   Such cells can be used for powering electric vehicles, emergency power distribution and storing solar energy.

DIFFERENT TYPES OF CELLS
There are two main types of cells – Primary cells and secondary cells.
1. PRIMARY CELLS

• In primary cells, the cell becomes dead, when the electrical energy produced by chemical reaction is used up.
• These cells cannot be recharged and reused.
• Examples of primary cells are dry cells and button cells. [Dry cells are commonly used in clocks and button cells are used in watches]

2. SECONDARY CELLS

• A secondary cell can be recharged and used again.
• The most important type of secondary cell is the lead acid cell used in vehicles.
• The nickel-cadmium cell used in flashlights is also a secondary cell.

Uses of Lithium-ion batteries
Lithium-ion batteries are secondary cells that have become an integral part of our daily life.

• They are used as power sources for devices ranging from smartphones and laptops to electric vehicles.
• Lithium-ion batteries are used in portable electronic devices due to their high energy density, long lifespan and low self-discharge rate.
• Lithium-ion batteries are also used as satellite batteries, a crucial component in the energy systems of spacecrafts.

FUEL CELLS

• Fuel cells are a type of galvanic cell.
• They convert chemical energy produced by the combustion of fuel into electrical energy
• Fuels used include: – Hydrogen, Methane, Methanol
• They operate continuously as long as fuel is supplied.
• Main advantage: They produce electricity with high efficiency.