Get Ahead in Class 9 Chemistry Chapter 7 Electrochemistry | In-Depth Notes and Success-Oriented Tests Preparation Guide

Score High in Chemistry Class 9 Board Exams  & Entry Tests with Detailed Notes and Answer Key

    Are you a Class 9 student struggling to understand the complex concepts of Electrochemistry? Look no further! Our comprehensive answers to Grade IX level Chemistry Chapter 7 will provide you with a clear and concise understanding of the most important topics in Electrochemistry. With a focus on long and detailed answers, you'll be able to grasp key concepts and phenomena in an easy and useful way, ensuring success in board exams like BISE, Federal Boards, as well as competitive tests like ETEA, NMDCAT, NEET, CSS, PCS, NTS, CTS, Army Cadet Colleges, and more. In this article, we'll cover important topics such as Old/Classical and Modern concepts of Oxidation and Reduction reactions, Oxidation State, the Rules of assigning oxidation numbers, Numerical examples of finding oxidation numbers of compounds, Redox reactions, Electrochemical Cells, Electrolytic refining of copper, Batteries, Electrolysis of fused NaCl, Manufacturing of NaOH from brine solution, Corrosion, and its prevention, the Process of rusting of Iron, and Electroplating. Get ready to deepen your understanding of Electrochemistry and achieve success in your exams and tests with our informative and useful guide. 

Class 9th Level Chemistry Notes on "Electrochemistry" 
From the Notes Library of H.E.S (Health, Education, and Skills)

What do you mean by Oxidation and Reduction reactions? Explain your answer.

Oxidation

A Chemical reaction that involves in addition of Oxygen or removal of Hydrogen is called an Oxidation reaction.

Explanation

    The following two concepts explain Oxidation reactions.

A. Old Classical concept

    According to this concept Oxidation reactions occur when Oxygen is added or Hydrogen is removed in a reaction. For example

i. Addition of Oxygen: Consider the following reaction

        C + O₂         ---->       CO₂

    In this reaction Carbon-di-oxide (CO₂) is formed when Oxygen is added to Carbon, and so is an Oxidation reaction.

ii. Removal of Hydrogen: Consider the following reaction

        2 NH₃ + 3 Cl₂        --->        N₂ + 6HCl

    In the above reaction Hydrochloric Acid (HCl) is formed only when Hydrogen is removed from Ammonia (NH₃), so is an Oxidation reaction.

B. Modern Electronic Concept

    According to this concept Oxidation reactions occur when electrons are lost in a reaction. The substance that loses electrons is said to be Oxidized.

For example:            Consider the following reaction

        Fe⁺²   ---->          Fe⁺³ + e-

Here Fe⁺² is Oxidized into Fe⁺³.

Reduction

A Chemical reaction that involves in addition of Hydrogen or the removal of Oxygen is called a Reduction reaction.

The following two concepts explain Reduction reactions.

A. Old Classical concept

    According to this concept Reduction reaction occur when Hydrogen is added or Oxygen is removed in a reaction.

For example: 

i. Addition of Hydrogen: Consider the following reaction

        N₂ + 3H₂     ---->          2NH₃

    In this reaction Ammonia (NH₃) is formed when Hydrogen is added to Nitrogen, so is a reduction reaction.

ii. Removal of Oxygen: Consider the following reaction

        2 H₂O      ---->       2H₂ + O₂

    In the above reaction Water (H₂O) molecules remove Oxygen to liberate Hydrogen, and so is a reduction reaction.

B. Modern Electronic Concept

    According to this concept Reduction reaction occur when electrons are gained in a reaction. The substance that gains electrons is said to be Reduced. For example, consider the following reaction

        2Na + Cl₂    --->     2 NaCl

    In this reaction Sodium when get Oxidized from Sodium ion (Na⁺) and Chlorine when getting Reduced by Chlorine ion (Cl^-). This reaction is called Oxidation Reduction reaction or simply Redox reaction because the loss and gain of electrons occur simultaneously.

Note:  From the above discussion it is clear that Oxidation and Reduction are totally opposite of each other.

Define Oxidation State? What rules should be followed while assigning the Oxidation Number?

Oxidation State

The positive or negative apparent charge on an atom of an element in a compound or an ion is called Oxidation Number.

    It may be +ive, -ive, or even zero. Unlike ionic charges, oxidation numbers do not have an exact meaning but are useful in naming compounds, writing formulae, and balancing chemical equations. The colors of solutions change, with the change of oxidation state i.e. Cr⁺² is blue, Cr⁺³ is green while Cr⁺⁶ is orange in color. The oxidation state can be described by the Oxidation Number

Rules for assigning the Oxidation Number

Following rules must be strictly followed when assigning the Oxidation Number

I. Oxidation number of elements in the free state: The Oxidation number of all the elements in the Free State (not in compound form) is Zero i.e. the Oxidation number of Cl₂, Na, O_2, etc. is Zero.

II. Oxidation number of simple ions: The Oxidation number of simple ions is the same as to charge on it e.g. Sodium has a +1 charge so +1 will be its Oxidation number. Similarly, the oxidation number of Ca⁺², Al^+3, and Br-1 are +2, +3, and -1 respectively.  

III. Oxidation number of Hydrogen in its compounds: The oxidation number of Hydrogen in its compounds is +1, except in Metal Hydrides where it is -1 e.g. in NaH       Na=+1 and H=-1 (because NaH is a metal hydride).

IV. Oxidation number of Oxygen in its compounds: The oxidation number of Oxygen in its compounds is -2, except in Per Oxides where it is -1 and for OF₂ it is +2 while it is -1/2 in the case of superoxide.

V. Oxidation number of elements of Group-IA, IIA, and IIIA: The oxidation numbers of Group - IA, IIA, and IIIA are +1, +2, and +3 respectively. 

VI. Oxidation number of Halogens: The oxidation number of Halogens in their binary compounds is -1.

VII. Oxidation number of neutral molecules: In all neutral molecules, the algebraic sum of the Oxidation numbers of all elements is Zero. For example, in H₂SO₄  2 (+1 ) + 1 (+6) + 4 (-2) = 0

VIII. Oxidation number of neutral molecules: In ions, the algebraic sum of oxidation numbers is equal to the charge on the ion. For example, in MnO₄^-1   1 (+7) + 4 (-2) = -1

IX. Oxidation number and electronegativity: More electronegative element has negative Oxidation number.

X. Oxidation number in different compounds: The same element may show different oxidation numbers in different compounds. For example, in CO oxidation number of Carbon is “+2” and that of Oxygen is “-2” but in CO₂ the oxidation number of Carbon is “+4” and that of Oxygen is “-4”.

Oxidation Number	is of
-2	Oxygen in its compounds
-1	Simple Anionic elements, Metal Hydrides, Per-Oxides, Halogens,
0	Elements in free state, Neutral molecules
+1	Simple Cationic elements, Hydrogen in its compounds, Group-IA elements
+2	Simple Cations with +2 charge, OF2, Group-IIA elements,
+3	Simple Cations with +3 charge, Group-IIIA elements

 

Example: Determine the oxidation number of Phosphorus in Phosphorous pentachloride (PCl₅)     

Given data:          Substance = PCl₅                                                          

Required:             Oxidation number of Phosphorous  (P) in PCl₅   =?                     

Solution:             

Rule  As we know that the oxidation number of Chlorine is “-1” because it has a -1 charge, so     

The oxidation number of Chlorine= -1

Let the Oxidation number of P (Phosphorous) in PCl₅ = X                                                 

As we know that in neutral molecules the algebraic sum of the oxidation numbers is Zero, therefore in PCl₅ 

Oxidation number of P + 5 (Oxidation number of Cl) = 0

= X+5 (-1) = 0                                                                                             

= X – 5 =  0                                                                                                

= X= +5                                                                                            

So the Oxidation number of P (Phosphorous) in PCl₅ = +5

Example:     What is the Oxidation number of Nitrogen in Nitrite ion in NO₂^-1?

Given data:          Substance = NO₂^-1                                                                 

Required:             Oxidation number of Nitrogen (N) = ?                                 

Solution:             

Let the Oxidation number of Nitrogen (N) in NO₂^-1 = X                                                      

As we know that in ions, the algebraic sum of oxidation numbers is equal to the charge on the ion, therefore in NO₂^-1       

Oxidation number of N + 2(Oxidation number of O) = -1

= X+2 (-2) = -1                                                                                            

= X – 4 = -1                                                                                                

= X = -1 + 4                                                                                                

X = +3                                                                                               

So the Oxidation number of Nitrogen in NO₂^-1 = +3

Example:     What are the three possible oxidation states of Sulphur?

    Sulfur belongs to Group-VIA of the periodic table, so due to six valence electrons its maximum oxidation state may be “+6”. While in the free state, its oxidation state is “0”. Similarly, its oxidation states may be +4, +2, and -2.

Example:     What is the oxidation state of Tin (Sn) in SnCl₄?

Given data: Substance = SnCl₄                                                                           

Required:   Oxidation number of Sn (Tin) in SnCl₄ (Tin chloride)  = ?   

Solution:    

Rule  As we know that the oxidation number of Chlorine is “-1” because it has a -1 charge, so the  Oxidation number of Chlorine= -1

Let the Oxidation number of Sn (Tin) in SnCl₄ = X                                           

As we know that in neutral molecules the algebraic sum of the oxidation numbers is Zero, therefore in SnCl₄        

Oxidation number of Sn + 4(Oxidation number of Cl) = 0

= X + 4 (-1) = 0                                                                                           

= X – 4 = 0                                                                                                 

= X = +4                                                                                  

So the Oxidation number of Sn (Tin) in SnCl₄ = +4

Example: Calculate the Oxidation number of S in K₂SO₄.

Given data: Substance = K₂SO₄                                                                          

Required:   Oxidation number of S (Sulphur) in K₂SO₄ (Sulphur Chloride)              

Rule  As we know that the oxidation number of K (Potassium) is “+1” because it has a +1 charge, so

The oxidation number of K= +1, similarly the Oxidation number of Oxygen= -2

Let the Oxidation number of Sulphur (S) in K₂SO₄ = X                                     

As we know that in neutral molecules the algebraic sum of the oxidation numbers is Zero, therefore in K₂SO₄       

2 (Oxidation number of K) + 1 (Oxidation number of S) + 4 (Oxidation number of Oxygen) = 0

= 2 (+1) + X + 4(-2) = 0                                                                              

= +2 + X – 8 = 0                                                                                         

= X – 6 = 0                                                                                                 

= X = +6     

So the Oxidation number of Sulphur (S) in K₂SO₄ = +6

Describe Oxidizing and Reducing agents in detail.

A. Oxidizing Agents

Oxidizing Agent is a substance (acts as an agent) that oxidize other substances and themselves get reduced.

Explanation

    According to the classical concept, the oxidizing agent may be any substance that supplies Oxygen to other substances.

1. Remove Hydrogen from other substances.
2. Remove electrons from other substances.
3. The oxidation state of oxidizing agent decreases.

    For example, H₂, C, H₂S, and CO are some Reducing agents.

Example:    Consider the reaction below.

Zn⁰ + Cu⁺² O^-2   →  Zn⁺²O^-2 + Cu⁰

Solution: According to the above reaction
The Oxidation number of Cu in CuO = +2
The Oxidation number of Cu = +0                  

    It means that the oxidation state of Cu has decreased from “+2” to “0”. So CuO acts as an oxidizing agent in the above reaction. Similarly, Cu of CuO has removed electrons from Zn i.e. 

        Zn⁰   → Zn⁺² + 2 e^- 

So it is an oxidizing agent.

Example:    Consider the reaction below.

        Sn⁺² + Cl₂⁰   →  Sn⁺⁴ +2 Cl^-1   
Solution:    According to the above reaction

The Oxidation number of Cl2 = 0
The Oxidation number of Chloride ion Cl-1 = -1 

    It means that the oxidation state of Cl2 has decreased from “0” to “-1”. So Chlorine acts as an oxidizing agent in the above reaction. Similarly, Cl has removed electrons from Sn+2, so it is an oxidizing agent.

B. Reducing Agents

A substance (act as an agent) that Reduces other substance and itself get Oxidized is called a Reducing Agent.

Explanation

    According to the classical concept, a reducing agent is one which

1. Remove Oxygen from other substances.
2. Provide Hydrogen to other substances.
3. The oxidation number of the reducing agents increases during a redox reaction.

    For example, Hydrogen Sulphide (H₂S), Na, Al, Mg, etc.

Example     Consider the reaction below.

        Br₂⁰  + H₂⁺² S^-2 O^-2   →  2 H⁺¹Br^-1 + S⁰
According to the above reaction
The Oxidation number of S in H2S = -2
The Oxidation number of free Sulphur = 0

    It means that the oxidation state of S has increased from “-2” to “0”. So H₂S acts as a reducing agent in the above reaction. Similarly, S has provided electrons to Br, so it is a reducing agent.

Some other examples of Oxidizing and Reducing agents

Oxidizing agents	Reducing agents
·         Bromine (Br2) ·         Chlorine (Cl2) ·         Concentrated Sulphuric acid (H2SO4) ·         Nitric acid (HNO3) ·         Oxygen (O2) ·         Potassium permanganate (KMnO4) ·         Potassium dichromate (K2Cr2O7)	·         Carbon (C) ·         Carbon monoxide (CO) ·         Hydrogen (H2) ·         Hydrogen Sulphide (H2S) ·         Potassium Iodide (KI) ·         Sulphur dioxide (SO2)

What do you mean by Oxidation-Reduction (Redox) reactions, explain your answer? Give some reactions to illustrate your answer.

Oxidation Reduction Reaction

That type of reaction in which gain and loss of electrons take place simultaneously is called an Oxidation Reduction reaction or Simply Redox reaction. 

    Examples of the Oxidation Reduction reaction: Following are some of the Oxidation Reduction reactions

1.   2 Fe⁺³ + Mg   →    2 Fe⁺² + Mg⁺²

2.   Cu + 2 H₂SO₄    →  CuSO₄ + SO₂ + 2 H₂O

3.   2 HNO₃ + 6 HI   →  2 NO + +3 I₂ + 2H₂O

Explanation

    Every Redox reaction is considered as made up of the following two half-reactions. To understand these, consider the following reactions.

2 Na + I₂  →  2 Na+ + 2 I-

1^st half-reaction: In the 1^st half reaction loss of electrons take place, which is called the Oxidation reaction. The oxidation reaction for the above reaction can be written as

                        2 Na      →     2 Na⁺ + 2 e-

2^nd half-reaction: In the 2^nd half-reaction gains of electrons take place, which is called the Reduction reaction. The reduction reaction for the above reaction can be written as

        I₂ + 2e-   →      2I-

    Collectively the above two half-reactions can be written as 

        2 Na   →     2Na⁺ + 2e^-  (Oxidation reaction)

        I₂ + 2e^-   →      2I-          (Reduction reaction)

Other examples of Redox reaction.

        i.    K⁰ + Ag+    →     K+ + Ag⁰

The two half-reactions are

        K    →   K+ + e-        (Oxidation)

        Ag+ + e-    →    Ag⁰   (Reduction)

        ii.   Sn⁺² + I₂   → Sn⁴+ + 2I-

The two half-reactions are

        I₂ + 2e-  →     2I         (Oxidation)

        Sn²⁺ →  Sn⁴⁺ + 2e-   (Reduction)

Define Electrochemical cells. Describe its two types in detail.

Electrochemical cells

A device in which the interconversion of chemical and electrical energy takes place is called an Electrochemical cell.

Types of Electrochemical cells

Following are the two types of Electrochemical cells.

• Electrolytic cells, and
• Galvanic cell/ Voltaic cell

1. Electrolytic cell

A device in which a non-spontaneous chemical reaction is carried out is called an Electrolytic cell.                              (OR)                      

A cell in which an electric current is used to produce a redox reaction is called an electrolytic cell.

Construction of Electrolytic cell

An electrolytic cell consists of

a.  Vessel:  Electrolytic cell consists of a vessel that contains electrolytes.

b. Metallic plates. An electrolytic cell consists of two metallic cells through which electrons enter or leave the cell. Their metallic plates are called Electrodes.

The electrodes are connected to a direct current source. The electrode connected to the positive terminal of the battery is called the anode while that connected to the negative terminal is called a cathode.

Working of the electrolytic cell

     When an electric current is passed the ions in the electrolyte move toward their respective electrodes. The anion liberates electrons at the anode. These electrons pass through the outer circuit to the cathode. The cations, which surround the cathode, consume these electrons. Hence, the number of electrons lost is always equal to the number of electrons gained. Thus, oxidation takes place at the anode while reduction takes place at the cathode. 

Example:  When an electric current is passed from the fused Sodium Chloride (NaCl), the following reaction takes place

        2 NaCl_(s)    →   Na⁺_(aq) + 2 Cl^-_(aq)       

At anode:            2Cl^-_(aq) →   Cl₂ + 2 e^-      (Oxidation)

At cathode:         2Na⁺_(aq) + 2 e^- →   2 Na  (Reduction)

Overall reaction: 2 Na⁺ + 2Cl^- →   2Na + Cl₂

2. Galvanic or Voltaic cell

The cell which produces electric current by undergoing an Oxidation Reduction reaction is called a Galvanic or Voltaic cell.

Construction of Galvanic or Voltaic cell

a: Two half cells

    A galvanic cell is made up of two half-cells, one for Oxidation and another for Reduction. One half-cell consists of a Zinc rod dipped in 1 Molar ZnSO₄ solution and the other half-cell consists of a copper rod placed in 1 molar copper Sulphate solution.

b: Salt bridge

    The two half cells are internally connected by a salt bridge and externally by a wire to which Galvanometer or Voltmeter is connected to measure the current. The salt bridge is a U-shape tube that is filled with electrolyte gel, such as KCl, K₂SO_4, or Na₂SO₄ solution called the agar.

Functions of salt bridge

• To allow electrical contact between the two solutions.
• To prevent the mixing of the two solutions.
• To keep electrical neutrality in each half cell.

Working of Galvanic cell/ Voltaic cell

    The Zinc metal has the tendency to lose electrons more readily than copper when the circuit is completed. As a result, oxidation takes place at the Zinc electrode. The electrons flow from the zinc electrode through the external circuit to the copper electrode. These electrons (coming from the zinc electrode) are picked by copper ions of the Cathodic solution and are deposited as copper atoms at the cathode. Electrons travel in the external circuit, while ions move through the salt bridge and in this way electric current is produced. 

This cell is just the reverse of the electrolytic cell. Daniel's cell is the best example of a galvanic cell.

At anode:            Zn_(s) →   Zn⁺² + 2 e^-        (Oxidation)

At cathode:         Cu⁺²_(aq) + 2e^- →   Cu_(s)    (Reduction)

Overall reaction: Zn + Cu⁺⁺  →   Zn⁺² + Cu

Describe the construction and purpose of Daniel's cell in detail.

Daniel cell

Purpose of Daniel's cell

    Daniel's cell is an example of a Galvanic cell therefore its purpose is to produce electric current by Oxidation and Reduction reaction.

Construction of Daniel's cell

    The construction of Daniel's cell can be explained under the following headings

I. Rods of Daniel Cell

    Daniel's cell consists of Zinc (Zn) rod dipped in the aqueous solution of Zinc Sulphate (ZnSO₄) and a Copper (Cu) rod dipped in a solution of Copper Sulphate (CuSO₄).

II. Connection between Solutions

    The two solutions are connected through a U-shaped Salt bridge filled with Potassium Chloride (KCl).

III. Voltmeter

    To measure the current a Voltmeter is connected to electrodes.

Process of Current production         

    After the circuit is completed, Zn is oxidized to Zn²⁺ ions and liberated electrons on the electrode. Electrons travel through the outer circuit towards the Cu electrode. Cu²⁺ ions present in the solutions take electrons from the electrode and are deposited as Cu metal at the electrode. Thus two half-reactions can be written as

Anodic half reaction         Zn_(s)  →     Zn²⁺ + 2e-     (Oxidation)

Cathodic half reaction      Cu²⁺_(aq) + 2e-    →   Cu_(s) (Reduction)

Net reaction         Zn_(s) + Cu²⁺_(aq)   →   Zn²_(aq) + Cu_(s)      

    Both the half-reactions take place simultaneously, electrons travel in the external circuit while ions travel through the salt bridge and in this way electricity is produced. Most of the Organic compounds are Non-Electrolytes like Sugar, Glucose, and Urea. 

How can be Copper refined? Describe the Electrolytic refining of Copper in detail.

Electrolytic refining of Copper

Importance of Copper

    Copper is one of the most important metals being used in making electrical cables, utensils, and different alloys like Brass and Bronze, etc.

Why Copper is refined electrolytically?

    Copper when used for electrical purposes must be highly pure, but when it is extracted from its Sulphide ore, it is called blister copper. Blister copper is only 99% pure, because of the impurities like Aurium (Au), Argentum (Ag), and Platinium (Pt), so it is further purified electrolytically.

Process of electrolytic refining of Copper

The electrolytic cell consists of

a.   Anode: It is made from an impure blister of Copper, and

a. Cathode: It consists of a thin sheet of pure Copper.

These two electrodes are then suspended in a solution of Copper Sulphate (CuSO₄) and Sulphuric Acid (H₂SO₄). By passing a current of 0.3 V and at 50 ℃ Copper from the impure Anode dissolves to give Cu²⁺ ions, which are then reduced to metallic Copper and deposited in Cathode, as given below

At Anode   Cu   →     Cu²⁺_(aq) + 2e-        (Oxidation)

At Cathode.  Cu²⁺_(aq) + 2e-     →     Cu    (Reduction)

    During this process, the less active metals, like Au & Ag remain undissolved and settle down at the bottom of the cell as anode sludge. Which is processed to recover these precious metals. By electrolytic refining of Copper, up to 99.99% pure Copper can be obtained.

Define battery. What are the various types of batteries? Explain the Dry cell in detail. 

Battery

A cell in which chemical energy is converted into electrical energy is called a battery.  (OR)

A group or combination of Galvanic cells joined in series is called a battery.   

    For example, car batteries consist of six or more identical Galvanic cells connected in series. A battery is a self-contained, chemical power pack that can produce a limited amount of electrical energy.

Types of batteries

(i)      Primary batteries:         These batteries are not reversible and once discharged are discarded e.g. Dry cell.

(ii)     Secondary batteries:  These are reversible batteries and can be recharged again e.g. lead storage battery.

(iii)    Solar batteries:   These are photochemical cells and generate energy.

(iv)    Fuel batteries:     These are super batteries and have high energy efficiency.

Dry cell/ Leclanchi cell

    This cell was designed by George Leclanchi in 1887. The dry cell is acidic in nature. It consists of an outer “Zn” casing which acts as the anode. It is lined inside with a moist paper which prevents Zinc from coming in contact with the outer reactants but allows the diffusion of ions. A graphite rod is placed in the center of the container which act as a cathode. The container is filled with a paste of Ammonium Chloride (NH₄Cl), Manganese dioxide (MnO₂), and Carbon. The cell is waterproofed with wax. The voltage produced by the dry cell is 1.25V to 1.50V.

At anode:            Zn →  Zn⁺² + 2 e^-                                                  (Oxidation)

At cathode:         2 MnO₂ + 2 NH₄⁺ + 2 e^- →  2 Mn₂ O₃ + 2 NH₃ + H₂O    (Reduction)

Overall reaction: Zn + 2 MnO₂ + 2 NH₄⁺ + 2 e^- →   2 Mn₂O₃ + 2 NH₃ + H₂O

Describe the Electrolysis of fused NaCl in detail.

Or      How does Sodium (Na) is get purified/ produced in Down's cell? Explain the process.

Electrolysis of fused NaCl

    History:      Sodium metal was first discovered by Sir Humphrey Davey in 1807, by the electrolysis of fused NaOH.

Purpose of electrolysis of fused NaCl

    To obtain Sodium (Na) in its purest form from the electrolysis of fused NaCl.

The cell used in the electrolysis of fused NaCl

Down's cell is used for the electrolysis of fused NaCl. Down’s cell was designed by J.C Down.

Construction of Down's cell

 Down's cell consists of    

• Steel tank: The cell consists of a steel container lined inside with firebricks
• Steel electrodes: The cell also consists of two steel electrodes which are introduced through the side wall that acts as Cathode. Cathode is made of copper or iron.
• Graphite rod: The graphite rod emerges from the bottom to act as Anode.
• Cathode apartment: The cathode and anode are separated by an iron screen, to prevent the mixing of the two products of electrolysis i.e. Sodium and Chlorine gas. The molten sodium is collected in a cathode apartment where it rises to the top and is taped off through a pipe.

Process of electrolysis of fused NaCl/ Working of Down cell/ Manufacturing of Na metal from NaCl:

    The process of electrolysis of fused NaCl can be described in the following two steps

a. Oxidation reaction

    First of all, fused NaCl (electrode) is added to the cell which is a strong electrolyte that ionizes to get its electrons free and move to opposite electrodes. When current is passed the Na+ ions move Cathode, pick electrons, and deposit these as Na metal. i.e.

Cathodic half reaction        2 Na⁺ + 2e-   →    2Na

b. Reduction reaction

    The Cl- ions move towards the anode, give their electron to the electrode, and change to a neutral Cl atom. Cl being very reactive forms Cl₂ molecules and liberate Cl₂ at Anode. i.e.

Anodic half reaction      2 Cl-   →   Cl₂ + 2e-

So net overall reaction can be written as 

    2 Na⁺ + 2 Cl-  →  2 Na + Cl₂

Describe the electrolysis of the aqueous solution of NaCl.
Or      How Sodium Hydroxide (NaOH) is manufactured from brine solution?

Electrolysis of aqueous solution of NaCl

30% NaCl solution is called brine solution.

    Commercially electrolysis of NaCl is carried out in Nelson’s cell. Commercially NaOH (Sodium hydroxide) is prepared by the electrolysis of NaCl.

Construction of Nelson’s cell

Nelson’s cell consists of

Graphite anode

U-shaped perforated steel cathode lined inside with asbestos (mixed silicates of Ca & Mg) i.e. CaSiO₃.3MgSiO₃.

Rectangular steel tank, having a catch basin at the bottom.

Working of Nelson’s cell

    The Graphite anode is suspended in a U-shaped steel cathode having a brine solution. The saturated brine solution ionizes as follows

        NaCl   →  Na+_(aq) +Cl-_(aq)

        H₂O_(l)  →  H+_(aq) + OH-_(aq)

    By passing current the positive ions (Na+ and H+) move towards Cathode. H+ has more tendency to pick electrons and form hydrogen atom and then combine to form H₂ gas (molecule).

At Cathode:      2H+_(aq) + 2e-   →    H_2(g)     (Reduction)

    Cl- ions move towards Anode and give electrons to electrodes and are converted into Chlorine atoms. Chlorine atoms then combine covalently to form chlorine gas (molecule).

At Anode    2Cl-_(aq)    →  Cl_2(g) + 2e-       (Oxidation)

    Na+ ions are not reduced instead they combine with OH- to form Caustic Soda (NaOH) or Sodium Hydroxide.

        Na+_(aq) + OH-_(aq)   →   NaOH_(aq)

    The solution gets gradually alkaline with the formation of NaOH, which is collected at the bottom of the cell while Cl₂(g) is released at the Anode and H₂(g) at the Cathode.

What is Corrosion? How corrosion can be prevented?

Corrosion

The slow eating away of the metal by chemicals present in its environment is called Corrosion.

Explanation

    The term corrosion and rust are almost synonymous. Rusting of Iron is the best example of Corrosion that we notice in our daily life. When the Iron rust is corroded, which means that it is slowly eaten up due to some Oxidation-Reduction reaction. In this way, Iron is converted into its Oxide.

Prevention of environmental corrosion

Corrosion can be prevented by

I. Metallic coating

When a thin coating of one metal is applied over another metal by spraying, galvanizing (deposition of Zn on metal by dipping), or through electroplating, it is called Metallic coating.

Iron, utensils can be protected from rusting by Nickel, chromium, or tin plating.

II. Paints coating

Coating the metal surfaces with paints, oils or grease prevents corrosion. Modern paints contain a combination of chemicals called stabilizers. These stabilizers provide prevention against corrosion and other atmospheric effects.

III. Alloying

Stainless steel is corrosion resistant which is an alloy of Ni, Cr, and Si. In this way, metal can be protected. Steel is a solid mixture of iron, chromium, and Nickel. Similarly, Brass is an alloy of Zinc and Copper.

IV. Removal of stains

As the region of stains in an iron act as the site of corrosion, so by removing stains, corrosion can be prevented.

V. Cathodic protection

This method is used to protect iron in buried fuel tanks and pipelines. An active metal like Zn or Mg is connected by a wire to the pipeline or tank to be protected. It is because Zn and Mg are better reducing agents than iron. Electrons are supplied by Zn or Mg instead of iron, thus protecting the iron from oxidation.

As oxidation occurs, Zn or Mg anode dissolves, so it must be replaced periodically. Ship hulls are protected in a similar way by attaching bars (rods) of Titanium metal to the steel hull. In salty water, Titanium act as an anode and is oxidized instead of a steel hull.

Hull:  The hull is the lightweight body of a ship or boat. Above the hull is a superstructure and deckhouse.

VI. Corrosion inhibitors

Certain substances inhibit corrosion to form e.g. Glycerin and Polyethylene.

Explain how rusting of Iron occurs?  Or    Describe the process of rusting of Iron.

Rusting of Iron

What is rust?

When pure white Iron is exposed to moist air, it converts into a reddish-brown mass called Rust.  Chemically rust is hydrated Iron (III) Oxide. So the corrosion of iron is called rust.

Conditions necessary for rusting

Air in the atmosphere.

Weakly acidic atmosphere.

A thin film of water/moisture on the metal’s surface.

Chemistry of rusting (OR) How rust is formed?

For rust formation, moisture air should be around the metal. The impurities are responsible for the formation of small electrolytic cells, with an Anode of pure Iron and a Cathode of impure or strained portion. Iron is Oxidized at Anode producing Iron Fe(II) ions and electrons. It moves along the surface of the metal to Cathode where it reacts with Water and Oxygen to form Hydroxide ions.

Anode                  2Fe   →    2Fe⁺²_(aq) + 4e-

Cathode               2H₂O + O₂ + 4e-   →   4OH-

    Fe(II) hydroxide is further oxidized by atmospheric Oxygen to form hydrated Fe(III) oxide, rust

Fe²⁺_(aq)                  Fe⁺³_(aq) + e-

Fe³⁺_(aq) + 3OH-                  Fe(OH)_3(s)

    The rust mass being porous in nature can't further prevent atmospheric action. In the atmosphere, rust-causing agents are Oxygen, carbon-di-oxide, and moisture.  

Define Electroplating. What is the purpose of electroplating? Describe the process of electroplating with a diagram.

Electroplating

A process in which a thin layer of one metal is deposited on another metal electrolytically is called Electroplating.

Purpose of electroplating

The purpose of electroplating is given as follows.

1. Protection:  To protect the inner metal from atmospheric effect and so from corrosion. For example, Ni and Cr are deposited over Iron to prevent it from corrosion.
2. Repair:  To weld the broken parts of machinery by depositing the metal on it.
3. Decoration: Electroplating is done to enhance its beauty. For example, Gold, Silver, and platinum over the surface of an inferior metal to increase its beauty.

Procedure of electroplating

    The substance on which another metal is deposited is made as a cathode, which is first washed with water and then with Caustic soda. The metal to be deposited is made as an anode. The soluble salt of the metal to be deposited is used as an electrolyte. The electrodes are connected to the battery. When an electric current is passed, the metal atoms from the anode are deposited over the cathode.