⚡ Oxidation, Reduction & Oxidation Numbers Explained! 🔬 | Master Chemistry Basics 🧪💥
📘Types of Chemical Reactions According to Loss or Gain of Electrons
Oxidation and reduction reactions are electrochemical reactions. In electrochemistry oxidation and reduction reactions involve transfer of electrons. The chemical reactions in which chemical energy changes into electrical energy or Vice Versa are called electrochemical reactions.
1. Oxidation (De-electronation)
2. Reduction (Electronation)
📘Difference between Oxidation and Reduction
📘Difference between Oxidation and Reduction
📘Oxidation or De-electronation
📚1. Oxidation as addition of oxygen
Oxidation is defined as a reaction of addition of oxygen to a substance either other elements or compounds (to produce oxide) e.g. the rusting of iron, burning of magnesium, carbon etc. in air are typical examples of oxidation.
📚2. Oxidation as a removal of Hydrogen (dehydrogenation)
Oxidation is the process of removal of hydrogen from a compound. In organic chemistry, removal of hydrogen from a compound is termed as dehydrogenation.
📚3. Oxidation as addition of electronegative element
Oxidation is a chemical reaction in which an electronegative element is added into any chemical species (atom, molecule or ion).
📚4. Oxidation as Removal of electropositive element
Oxidation is a chemical reaction in which an electropositive element is removed from any chemical species (atom, molecule or ion).
📚5. Oxidation as loss or removal of electrons (Electronic definition)
The most concise definition and broader view of oxidation reaction is in terms of the electron transfers.
According to modern electronic concept, the process or a reaction in which a substance (i.e. atom, molecule or ion) loses one or more electrons (which is manifested by an increase in its oxidation number) is called Oxidation. Oxidation involves in producing or increasing the positive charge on the species or decreasing its negative charge
e.g.
The substance that donates electrons is oxidized but it acts as a reducing agent.
📚6. Oxidation as an increase in oxidation number
In terms of oxidation number concept, the process in which the oxidation state of an element is increased is called Oxidation.
e.g.
(i) In the following redox reaction between zinc powder and dilute hydrochloric acid, zinc has been oxidized to zinc chloride because its oxidation number has been increased from 0 to +2.
(ii) In the following redox reaction between carbon and oxygen gas, carbon has been oxidized to carbon dioxide because its oxidation number has been increased from 0 to +4.
📘Reduction or Electronation
📚1. Reduction as addition of hydrogen (Hydrogenation)
Reduction is the process of addition of hydrogen to a substance. In organic chemistry, addition of hydrogen in a substance is termed as hydrogenation.
📚2. Reduction as removal of oxygen
Reduction is the process or a reaction (just opposite to oxidation) which involves the removal of oxygen (or electronegative atom) from substances (e.g. oxides).
The most of such reactions are simple displacement reactions. In such reactions, oxides are reduced to free elements (usually metal) while other substances (usually an element) are oxidized to their respective oxides.
📚3. Reduction as addition of an electropositive element
Reduction is a chemical reaction in which an electropositive element is added to any chemical species (atom, molecule or ion).
📚4. Reduction as removal of an electronegative element
Reduction is a chemical reaction in which an electronegative element is removed from chemical species (atom, molecule or ion).
📚5. Reduction as gain or addition of electrons/De-electronation (Electronic definition)
The most concise definition and broader view of reduction reactions is in terms of the electron transfers.
According to modern electronic concept, the process or a reaction in which a substance (i.e. atom, molecule or ion) gains one or more electrons (which is manifested by a decrease in its oxidation number) is called Reduction. Reduction involves in producing or increasing the negative charge on the species or decreasing its positive charge
The substance that gains electrons is reduced but it acts as an oxidizing agent.
e.g.
📚6. Reduction as a decrease in oxidation number
In terms of oxidation number concept, the process in which the oxidation state of an element is decreased is called Reduction
e.g.
In this reaction, reduction of Br₂ occurs due to decrease in its oxidation number from 0 to –1.
📘Oxidizing Agent or Oxidant OR Oxidizer
📚Complete definition
2. accepts hydrogen
3. gains one or more electrons
4. undergoes decrease in oxidation number
5. causes oxidation
6. is reduced
📚1. Oxidizing Agent as donor of oxygen
Some other examples of oxidant are illustrated by their decomposition or their reactions with other reagents.
📚2. Oxidizing Agent as Acceptor of Hydrogen
Oxidizing agent is a substance that removes or accepts hydrogen from a substance.
📚3. Oxidizing Agent as Electron Acceptor or Electron recipient
(ii) Cl₂ molecules decreases its oxidation number from 0 to −1, so it acts as an oxidizing agent.
📚4. Test for Oxidizing Agent by KI solution
📚5) Examples of Oxidizing Agent
Following are the examples of Oxidant:
📘Reducing Agent or Reductant or Reducer
📚Complete Definition
2. loses hydrogen
3. loses one or more electrons
4. oxidation number of atom increases
5. causes reduction
6. is oxidized
📚1. Reducing Agent as Acceptor of Oxygen
Some other examples of reducing agent are:
📚2. Reducing Agent as Donor of Hydrogen
📚3) Reducing Agent as Electron Donor
A substance that donates or loses one or more electrons is called a reducing agent or reductant. Thus it is a substance that reduces the other substance (by supplying electrons to it) while itself gets oxidized (by losing electrons). Stated differently, reducing agent undergoes increase in oxidation number during a reaction.
e.g.
(i) Zn atom loses electrons during reaction & thus it is a reducing agent.
📚4) Identification Test for Reducing Agent by acidified potassium dichromate
The resulting colour change of orange to green by the addition of acidified potassium dichromate (an oxidant) to a reducing agent can be used as a test for reducing agent.
📘Difference between Oxidizing and Reducing Agent
📘Some Oxidizing agent and Reducing Agent
📘Types of Chemical Reactions According to Electron Transfer
There are three types of chemical reactions based on oxidation and reduction:
1. Non- Redox Reaction
2. Redox Reaction or Oxidation-Reduction Reaction (ORR)
3. Auto Redox Reaction or Self Oxidation-Reduction Reaction (ARR)
📘Non- Redox Reaction
Definition of Non-Redox Reactions
The chemical reaction in which there is no electron transfer i.e. no substance is oxidized or reduced not undergoing change in oxidation number is called non-redox reaction.
General Examples of Non-Redox Reactions
1. Neutralization
2. Hydrolysis
3. Precipitation reactions
4. Acid displacement reactions
5. Base Displacement Reactions
6. Some decomposition reactions
7. All double decomposition reactions
8. Some molecular addition reactions
📘Oxidation-Reduction Reactions (ORR) or Redox Reactions
Can you tell?
In any redox reaction, the specie that oxidizes the other substance and itself gets reduced is known as oxidizing agent. It is a substance which gains electrons during a reaction from other substance undergoing decrease in oxidation number thereby oxidizing it. Stronger oxidizing agents are found in the lower region of ECS. Thus strongest oxidizing agent is fluorine.
Oxidizing agent may be defined as a substance supplying oxygen or electronegative element, removing hydrogen or electropositive element and accepting electrons thereby decreasing oxidation number. Oxidizing agent
1. Gives nascent oxygen2. accepts hydrogen
3. gains one or more electrons
4. undergoes decrease in oxidation number
5. causes oxidation
6. is reduced
📚1. Oxidizing Agent as donor of oxygen
Oxidizing agent is a substance (element or compound) that releases or supplies oxygen or nascent (or atomic) oxygen either on decomposition or on treatment with other substance.
e.g.
in following reaction, CuO being donor of oxygen is acting as oxidizing agent supplying oxygen to H₂ and thus itself reduces to Cu while H₂ being acceptor of oxygen is acting as reducing agent adding oxygen and thus itself oxidizes to H₂O.
Some other examples of oxidant are illustrated by their decomposition or their reactions with other reagents.
📚2. Oxidizing Agent as Acceptor of Hydrogen
Oxidizing agent is a substance that removes or accepts hydrogen from a substance.
📚3. Oxidizing Agent as Electron Acceptor or Electron recipient
A substance (in a redox chemical reaction) that accepts or gains or receives one or more electrons from other substance (called the reductant or reducer) is known an oxidizing agent or oxidant. Stated differently, oxidizing agent undergoes decrease in oxidation number. Thus it is a substance that oxidizes the other substance (by removing electrons from it) while itself gets reduced (by accepting electrons from the other substance) to a lower oxidation state.
e.g.
(i) Br₂ molecule accepts electron during reaction and thus it acts as an oxidant.
(ii) Cl₂ molecules decreases its oxidation number from 0 to −1, so it acts as an oxidizing agent.
(iii) Zinc reacts with dilute sulphuric acid to form zinc sulphate and hydrogen gas. In this redox reaction, H₂SO₄ acts as an oxidizing agent accepting electrons (undergoing increase in oxidation number) form zinc and thus reduces to H₂ gas while zinc acts as reducing agent by donating electrons (undergoing decreases in oxidation number) and thus oxidizes to Zn²⁺ ions. [SO₄²⁻ ions being spectator ions, do not appear in net equation].
📚4. Test for Oxidizing Agent by KI solution
The resulting colour change of colourless to brown by the addition of KI (a reductant) to an oxidizing agent is a test for oxidizing agent.
Aqueous potassium iodide (KI) is used to test for the presence of an oxidizing agent. KI is colourless. If a drop of KI is added to a solution containing an oxidizing agent, a brown solution will be formed. The solution turns brown because the iodide ions (I⁻) is oxidized to iodine (I₂), by the oxidizing agent. The iodide ion is colourless but aqueous iodine is brown.
Starch-iodide paper can also be used to test for the presence of oxidizing agents. Oxidizing agents change the colour of moist Starch-iodide paper from white to blue. This is because the iodine produced reacts with the starch to give a blue colour.
📚5) Examples of Oxidizing Agent
Following are the examples of Oxidant:
📘Reducing Agent or Reductant or Reducer
📚Complete Definition
In any redox reaction, the specie that reduces the other substance and itself gets oxidized is known as Reducing Agent or Reductant. It is a substance which loses electrons and give to the other for reduction undergoing oxidation. Strongest reducing agent are located on upper region of ECS. Thus strongest reducing agent is Li.
A reducing agent may be defined as a substance supplying hydrogen or electropositive element, removing oxygen or electronegative element and donating electrons thereby increasing oxidation number. Reducing agent
1. accepts nascent oxygen2. loses hydrogen
3. loses one or more electrons
4. oxidation number of atom increases
5. causes reduction
6. is oxidized
📚1. Reducing Agent as Acceptor of Oxygen
Reducing agent is a substance (element or component) that accepts oxygen or nascent (or atomic) oxygen (released by oxidizing agent).
e.g.
In following reaction, C being acceptor of oxygen is acting as reducing agent receiving oxygen from ZnO and thus itself oxidizes to CO while ZnO being donor of oxygen is acting as oxidizing agent supplying oxygen and thus itself reduces to Zn.
Some other examples of reducing agent are:
📚2. Reducing Agent as Donor of Hydrogen
Reducing agent is a substance (element or compound) that releases nascent (or atomic) hydrogen either on decomposition or on treatment with other substance.
📚3) Reducing Agent as Electron Donor
A substance that donates or loses one or more electrons is called a reducing agent or reductant. Thus it is a substance that reduces the other substance (by supplying electrons to it) while itself gets oxidized (by losing electrons). Stated differently, reducing agent undergoes increase in oxidation number during a reaction.
e.g.
(i) Zn atom loses electrons during reaction & thus it is a reducing agent.
(ii) Na atom loses electrons during reaction & thus it is a reducing agent.
(iii) Magnesium reacts with dilute sulphuric acid to form magnesium sulphate and hydrogen gas. In this redox reaction, H₂SO₄ acts as an oxidizing agent accepting electrons (undergoing increase in oxidation number) from magnesium and thus reduces to H₂ gas while magnesium acts as reducing agent by donating electrons (undergoing decreases in oxidation number) and thus oxidizes to Zn²⁺ ions. [SO₄²⁻ ions being spectator ions, do not appear in net equation].
📚4) Identification Test for Reducing Agent by acidified potassium dichromate
The resulting colour change of orange to green by the addition of acidified potassium dichromate (an oxidant) to a reducing agent can be used as a test for reducing agent.
Acidified potassium dichromate (VI) can be used to test for the presence of a reducing agent. Acidified potassium dichromate (VI) is made by adding dilute sulphuric acid to aqueous potassium dichromate (VI). The colour of Acidified potassium dichromate (VI) solution changes from orange to green in the presence of a reducing agent. The half reaction is shown here:
In this reaction, the dichromate (VI) ion (Cr₂O₇²⁻) is reduced to the chromium (III) ion (Cr³⁺). Cr₂O₇²⁻ loses oxygen and the oxidation state of chromium decreases from +6 to +3.
📘Difference between Oxidizing and Reducing Agent
📚Quick Recap of Key Concepts:📘
Oxidation is the loss of electrons, resulting in an increase in the oxidation number.
Reduction is the gain of electrons, resulting in a decrease in the oxidation number.
Oxidizing Agent: The substance that accepts electrons (gets reduced).
Reducing Agent: The substance that donates electrons (gets oxidized).
📚Criteria of deciding oxidizing and reducing nature of compounds📘
A substance acts only as an oxidizing if the oxidation number of one of its element (central atom) is in its highest oxidation state and as a reducing agent if the oxidation number of one of its element is in its lowest oxidation state. However, if the oxidation number of one of the elements of a substance is in its intermediate oxidation state, it can act both as an oxidizing as well as a reducing agent.
For example;
1. The oxidation number of N in HNO₃ is maximum i.e.+5, therefore, it can act only as an oxidizing agent by accepting one or more electrons e.g.
Here, the oxidation number of N decreases from +5 in HNO₃ to +4 in NO₂ and hence it acts as an oxidizing agent.
2. The oxidation number of S in H₂S is least i.e. -2 and hence it can act only as a reducing agent by losing one or more electrons. Hence the oxidation number of S increases from -2 in H₂S to 0 in elemental sulphur and hence it acts as a reducing agent
3. The oxidation number of N in HNO₂ is intermediate (+3), it is neither maximum (+5) nor minimum (-3), therefore, it can act both as an oxidizing as well as a reducing agent. E.g.
Here, the oxidation number of N increases from +3 in HNO₂ to +5 in HNO₃, therefore, it acts a reducing agent 
Here, the oxidation number of N decreases from +3 in HNO₂ to +2 in NO, therefore, it acts as a an oxidizing agent.
📘Factors Affecting Strength of Oxidizing and Reducing Agent
1. A substance can act as oxidizing agent if the oxidation number of one of its element is maximum.
e.g.
HNO₃ in which O.N of N is +5 which is its maximum oxidation state is a strong oxidizing agent
2. The more the electronegativity of central element and the more is its oxidation number, the more is the oxidizing power.
e.g.
3. Oxyanions are stronger oxidizing agents in acidic solution than in basic or neutral solution.
4. A substance can act as reducing agent if the oxidation number of one of its element is minimum.
e.g. SnCl₂ (O.N of Sn =+2 which is least), FeSO₄ (O.N of Fe =+2), Na₂S₂O₃ (O.N of S =+2), H₂S (O.N of S = −2), H₂C₂O₄ (O.N of C =+3) etc.
6. A substance that can act as both reducing as well as oxidizing agent if oxidation number of one its element is in between its maximum and the minimum oxidation number value.
e.g.
HNO₂ (O.N of N =+3 which is intermediate of +5 and 0).
Oxidation number and Acid Strength
The greater the oxidation number of the central element, the greater is the acid strength.
📘Types of Chemical Reactions According to Electron Transfer
There are three types of chemical reactions based on oxidation and reduction:
1. Non- Redox Reaction
2. Redox Reaction or Oxidation-Reduction Reaction (ORR)
3. Auto Redox Reaction or Self Oxidation-Reduction Reaction (ARR)
📘Non- Redox Reaction
Definition of Non-Redox Reactions
The chemical reaction in which there is no electron transfer i.e. no substance is oxidized or reduced not undergoing change in oxidation number is called non-redox reaction.
General Examples of Non-Redox Reactions
1. Neutralization
2. Hydrolysis
3. Precipitation reactions
4. Acid displacement reactions
5. Base Displacement Reactions
6. Some decomposition reactions
7. All double decomposition reactions
8. Some molecular addition reactions
📘Oxidation-Reduction Reactions (ORR) or Redox Reactions
Can you tell?
🔥Why does cut apple turn brown when exposed to air?
🔥Why does old car bumper change colour?
🔥Why do new batteries becomes useless after some days?
🔥Why does silver tarnish over time?
🔥Why does the copper on a statue turn green after a while?
🔥Why does hydrogen peroxide bubble when it touches a wound?
All these changes involve Redox which is an abbreviation used for the terms “oxidation and reduction”. A large number of reactions such as respiration, rusting, combustion of fuel, fermentation of sugar etc involve redox reactions.
Definition of Oxidation-Reduction (Redox) Reaction (ORR)
Redox reactions are also called electron-transfer reactions since electrons are transferred from the reductant to oxidant.
Oxidation and reduction always occur simultaneously during a chemical reaction. The chemical reactions in which oxidation and reduction occur simultaneously are called oxidation-reduction reactions (ORR) or redox reactions. In other words, these are the reactions in which increase and decrease in oxidation number of same or different atoms occurs.
These reactions comprising of simultaneous oxidation and reduction.
In terms of electron transfer, a redox reaction is defined as the process in which electrons are transferred from one substance (reducing agent) to another (oxidizing agent).
Explanation
All oxidation and reduction reactions are complimentary of one another and occur simultaneously, one cannot take place without the other. No single oxidation and no single reduction process are known. Oxidation-reduction reactions involve two opposing but complementary processes. These processes can never occur singly i.e. every oxidation must necessarily be accompanied by its opposing process reduction and vice versa. The simultaneously oxidation and reduction reactions are generally termed as redox reactions. The substance which brings reduction is known as reducing agent while a substance which brings oxidation is known as oxidizing agent.
Daily life examples of Redox reactions
The reactions taking place in batteries are redox reactions. Redox reactions take place in the batteries such that electrons transferred can pass through some external circuit so that they produce electric current
Digestion and metabolism of food which takes place in our body in order to supply us the energy required to perform work is also takes place through a series of redox reactions.
Ordinary bleach oxidize the substances that stain fabric, this makes them colourless and easier to remove from fabric.
Redox couple
A redox couple is defined as having together the oxidized and reduced forms of a substance taking part in an oxidation or reduction half reaction. Represented as Zn²⁺/Zn and Cu²⁺/Cu.
🔑Example No. 1 of Redox Reaction
The addition reaction between H₂ gas and Br₂ to form hydrogen bromide is an example of redox reaction. In this reaction H₂ has been oxidized because its oxidation number has been increased so H₂ is a reducing agent while Br₂ has been reduced because its oxidation number has been decreased so Br₂ is an oxidizing agent. Thus it is a Redox Reaction during which oxidation and reduction takes place simultaneously.
🔑Example No. 2 of Redox Reaction
🔑Example No. 3 of Redox Reaction
🔑Example No. 4 of Redox Reaction
🔑Example No. 5 of Redox Reaction
The displacement reaction between H₂S and Cl₂ to form S and HCl is an example of redox reaction. In this reaction, H₂S oxidizes to S as it loses hydrogen (loss of H₂ = oxidation), so H₂S is reducing agent, while Cl₂ reduces to HCl as it ads hydrogen (gain of H₂ = reduction), so Cl₂ is oxidizing agent.
In terms of electron transfers, Cl₂ oxidizes H₂S to S by accepting two electrons from it while itself reduces to HCl by gaining these two electrons. Thus Cl₂ acts as oxidizing agent as it gains electrons while H2S is reducing agent as it donates electrons.
🔑Example No. 6 of Redox Reaction
The reaction between NH₃ and NO to form N₂ and H₂O is a redox reaction. In this reaction, NH₃ oxidizes to N₂ as it loses hydrogen (loss of H₂ = oxidation), so NH₃ is reducing agent while NO also reduces to N₂ as it loses oxygen (loss of oxygen), so NO is oxidizing agent.
In terms of electron transfer, NO oxidizes NH₃ to N₂ by accepting two electron (or four electrons per N₂) from it while itself reduces to N₂ by gaining three electrons from NH₃ (or six electrons per N₂). Thus NO acts as oxidizing agent as it gains electrons while NH₃ is reducing agent as it donates electrons.
🔑Example No. 7 of Redox Reaction
In above example iron undergoes reduction from +3 to +2 and tin undergoes oxidation from +2 to +4.
📘🌟General Examples of Redox Reactions🌟
Following are the examples of Redox reactions:
📘🌟Types of Redox Reactions
Redox reactions are divided into two main types.
(i) Inter molecular Redox Reactions
(ii) Intra molecular Redox Reactions
(iii)Disproportionation
(iv) Comproportionation reaction (reverse of Disproportionation; HClO + Cl‒ → Cl₂ + OH‒)
📘⚡Inter molecular Redox Reactions
In such redox reactions, one molecule of reactant is oxidized whereas molecule of other reactant is reduced. In this case, one substance is oxidized and another is reduced. In following reaction, HCl is oxidized while MnO₂ is reduced.
📘⚡Intra molecular Redox Reactions
In such redox reactions, one atom of a molecule is oxidized and other atom of same molecule is reduced. In this case, one element of the compound is reduced while another element of the same compound is oxidized.
Examples
In the decomposition of KClO₃, its Cl is reduced to KCl and O is oxidized to O₂.
In the decomposition of (NH4)₂Cr₂O7, its Cr is reduced to Cr₂O₃ and N is oxidized is oxidized to N₂.
📘⚡Auto-Redox reaction/Self-Redox reactions/ Disproportionation Reaction/Disputation Reaction
Definition
It is an important and special type of redox reaction in which a single substance (specie) undergoes simultaneous oxidation and reduction i.e. it occurs when a same element is both oxidized and reduced simultaneously (i.e. in the meantime). A specie undergoing auto-redox reaction is said to be disproportionate.
Disproportionation, also called disputation reaction, is basically a redox reaction involving simultaneous reduction and oxidation of atoms of the same element of a substance of intermediate oxidation state from one oxidation state to two different oxidation states forming two compounds, one with higher and one with lower oxidation states. So a species is simultaneously reduced and oxidized to form two different products.
Example
Reason
The requirement for disproportionation reaction to occur is, the element undergoing disproportionation should exhibit minimum three different oxidation states and the element must be less stable in a particular oxidation state from which it can be both oxidized as well as reduced to relatively more stable oxidation states.
Examples of Auto-Redox Reactions
1. Decomposition or Disproportionation of potassium chlorate to potassium perchlorate and potassium chloride
2. Decomposition of nitrogen (III) oxide into nitric oxide and nitrogen dioxide
3. Decomposition of hydrogen peroxide into water and oxygen
Decomposition reaction of hydrogen peroxide into water and oxygen involves disproportionation of oxygen. In this auto-redox reaction, the relatively less stable oxygen of peroxide in the -1 oxidation state disproportionates into relatively more stable compounds i.e. water and dioxygen changing its oxidation state to the -2 oxidation state in water and zero oxidation state in oxygen gas at the same time.
4. Dissolution of chlorine gas in water (Reaction of chlorine gas with water)
5. Photolysis of Mercurous chloride into mercuric chloride and mercury
Upon UV-irradiation, Mercurous chloride or mercury(I) chloride undergoes disproportionation. under UV light to give mercury and mercuric chloride. The Hg₂²⁺ ion is oxidized to Hg²⁺ and reduced to Hg.
6. Dissolution of nitrogen dioxide in water
When nitrogen dioxide in which oxidation state of nitrogen is +5 reacts with water (Ostwald process), it undergoes disproportionation reaction resulting in the formation of both nitric acid and nitrous acid (or nitric oxide; O.S of N = +2) wherein nitrogen has oxidation states +5 and +3 respectively. In this reaction, nitrogen of NO₂ with +4 oxidation state is simultaneously oxidized to nitric acid (+5 oxidation state) and reduced to nitrous acid or NO (with oxidation state +3 or +2). Thus, it is a disproportionation reaction.
7. Decomposition of Cuprous chloride into cupric chloride and copper
Decomposition of Cuprous chloride into cupric chloride and copper involves disproportionation of copper. When cuprous chloride in which oxidation state of copper is +1 is heated it is decomposed and simultaneously oxidized to copper chloride changing the oxidation state of copper from +1 to +2 and reduced to elemental copper changing the oxidation state of copper from +1 to 0. Thus, this is a disproportionation reaction
8. Dissolution of metal superoxides with water
9. Disproportionation of Phosphorus to phosphine and hypophosphite in alkaline medium.
Phosphorus disproportionates to phosphine and hypophosphite in alkaline medium. In this case, one P atom is reduced to -3 oxidation number (in PH₃) and three P atoms get oxidized to +1 (in NaH₂PO₂).
10. Auto-redox Reactions of chlorine gas with dilute or conc Alkalis (sodium hydroxide & lime water)
11. Cannizaro’s reaction/Auto-redox Reactions of formaldehyde with conc. Alkalis
The self-addition oxidation reduction and disproportionation Reaction in which two molecules of aldehyde lacking a-hydrogen are disproportinated into carboxylic acid (which form salt with alkali) and alcohol is known as Cannizaro’s Reaction.
Aldehydes lacking a-hydrogen like formaldehyde and benzaldehyde (do not show aldol condensation) undergo self-redox reaction in presence of aqueous alkali, two molecules of such aldehydes disproportionate and simultaneously oxidize and reduce one another into acid and alcohol respectively.
For example
Formaldehyde on heating with conc. Solution of strong alkali like NaOH undergoes self-oxidation reduction reaction in one molecule of formaldehyde is reduced to methanol and the other is oxidized to formic acid that forms salt with alkali.
⚡Comproportionation reaction/synproportionation (opposite of the disproportionation)
Comproportionation reaction is the opposite of disproportionation reaction. In this reaction, two reactants with the same element in different oxidation states combine to form the same element in the intermediate oxidation state.
The reverse of disproportionation, such as when a compound in an intermediate oxidation state is formed from precursors of lower and higher oxidation states, is called comproportionation,
Example
Ag²⁺(aq) + Ag(s) → 2Ag⁺(aq)
Another variant on disproportionation is radical disproportionation, in which two radicals form an alkene and an alkane.
2CH₃−CH₂ ⟶ H₂C=CH₂ + H₃C−CH₃
🎨 ⚡ Redox Reaction & Oxidation Number Quiz: Test Your Knowledge on Oxidation and Reduction! 🧪 🔴
1. Which of the following is true about oxidation?
🟥 A) It involves the gain of electrons
🟦 B) It involves the loss of electrons
🟩 C) It involves a decrease in oxidation number
🟨 D) It involves the gain of protons
2. In a redox reaction, if an element’s oxidation number decreases, what is happening to it?
🟥 A) It is being oxidized
🟦 B) It is being reduced
🟩 C) It is undergoing a chemical equilibrium
🟨 D) It is remaining unchanged
3. Which of the following substances is oxidized in the reaction:
2Mg + O₂ → 2MgO₂
🟥 A) Magnesium (Mg)
🟦 B) Oxygen (O₂)
🟩 C) Magnesium oxide (MgO)
🟨 D) None of the above
4. In the reaction
2Na + Cl₂ → 2NaCl₂
What happens to sodium (Na)?
🟥 A) It gains electrons and is reduced
🟦 B) It loses electrons and is oxidized
🟩 C) It remains unchanged
🟨 D) It gains protons
5. Which of the following is the oxidizing agent in the reaction:
Cu²⁺ + 2ē → Cu
🟥 A) Cu²⁺
🟦 B) Cu
🟩 C) 2ē
🟨 D) None of the above
6. Which of the following elements undergoes reduction in the reaction:
2Fe³⁺ + 6ē → 2Fe
🟥 A) Iron (Fe³⁺)
🟦 B) Iron (Fe)
🟩 C) Oxygen (O₂)
🟨 D) Electrons (ē)
7. In a redox reaction, which element is most likely to be the reducing agent?
🟥 A) An element with a high electronegativity
🟦 B) An element with a low electronegativity
🟩 C) An element with a high atomic number
🟨 D) An element that gains electrons
8. Which of the following reactions demonstrates both oxidation and reduction?
🟥 A) 2H₂ + O₂ → 2H₂O₂
🟦 B) NaCl → Na+Cl₂
🟩 C) N₂+3H₂ → 2NH₃
🟨 D) 2KCl → 2K+Cl₂
9. What happens to the oxidation number of oxygen in the reaction:
2H₂O → O₂+ 4H⁺ + 4ē
🟥 A) The oxidation number of oxygen decreases
🟦 B) The oxidation number of oxygen increases
🟩 C) The oxidation number of oxygen remains the same
🟨 D) Oxygen does not participate in this reaction
10. In a redox reaction, if a substance is oxidized, what happens to its oxidation number?
🟥 A) It increases
🟦 B) It decreases
🟩 C) It remains the same
🟨 D) It becomes zero
11. In the reaction:
2Fe²⁺ + Cl₂ → 2Fe³⁺ + 2Cl⁻
Which species is reduced?
🟥 A) Fe²⁺
🟦 B) Fe³⁺
🟩 C) Cl₂
🟨 D) Cl⁻
12. Which of the following has the highest tendency to act as an oxidizing agent?
🟥 A) Na⁺
🟦 B) F₂
🟩 C) H₂
🟨 D) Cl⁻
13. In acidic medium, the reaction:
MnO₄⁻ + 8H⁺ + 5ē → Mn²⁺ + 4H₂O
The change in oxidation number of Mn is:
🟥 A) +7 to +2
🟦 B) +2 to +7
🟩 C) +4 to +2
🟨 D) +6 to +3
14. Which of the following represents disproportionation reaction?
🟥 A) 2H₂+O₂→2H₂O₂
🟦 B) Cl₂ + 2NaOH → NaCl + NaOCl + H₂O
🟩 C) Zn + Cu²⁺ → Zn²⁺ + Cu
🟨 D) Fe²⁺ → Fe³⁺ + ē
15. Which of the following is not a redox reaction?
🟥 A) CaCO₃ → CaO + CO₂
🟦 B) 2Mg + O₂ → 2MgO
🟩 C) Zn + 2HCl → ZnCl₂+H₂
🟨 D) 2Na + Cl₂ → 2NaCl
16. In the reaction:
2Al + 3Cu²⁺ → 2Al³⁺ + 3Cu
Which statement is correct?
🟥 A) Al is reduced, Cu is oxidized
🟦 B) Al is oxidized, Cu is reduced
🟩 C) Both Al and Cu are oxidized
🟨 D) Both Al and Cu are reduced
17. Which of the following elements is most likely to act as a reducing agent in aqueous solution?
🟥 A) Cl₂
🟦 B) Na
🟩 C) O₂
🟨 D) Fe³⁺
18. In the reaction:
2H₂O₂ → 2H₂O+O₂
Hydrogen peroxide acts as:
🟥 A) Only oxidizing agent
🟦 B) Only reducing agent
🟩 C) Both oxidizing and reducing agent
🟨 D) Neither oxidizing nor reducing agent
19. Which of the following correctly describes the change in oxidation number of nitrogen in:
NH₃ → NO
🟥 A) –3 to +2
🟦 B) –3 to +1
🟩 C) –3 to 0
🟨 D) –3 to +3
20. Which of the following is the oxidizing agent in the reaction:
Zn + 2Ag⁺ → Zn²⁺ + 2Ag
🟥 A) Zn
🟦 B) Ag⁺
🟩 C) Zn²⁺
🟨 D) Ag
21. Which of the following is the reducing agent in the reaction:
Zn+H₂SO₄ → ZnSO₄+H₂
🟥 A) Zn
🟦 B) H₂SO₄
🟩 C) H₂
🟨 D) SO₄²⁻
22. In the reaction:
2K + 2H₂O → 2KOH + H₂
Potassium undergoes:
🟥 A) Reduction
🟦 B) Oxidation
🟩 C) Neutralization
🟨 D) No change
23. Which of the following is not an oxidizing agent?
🟥 A) KMnO₄
🟦 B) O₃
🟩 C) H₂O₂
🟨 D) Na
24. In the reaction:
Cl₂ + 2Fe²⁺ → 2Fe³⁺ + 2Cl⁻
Which statement is correct?
🟥 A) Cl₂ is oxidized, Fe²⁺ is reduced
🟦 B) Cl₂ is reduced, Fe²⁺ is oxidized
🟩 C) Both are oxidized
🟨 D) Both are reduced
25. Which of the following represents a redox reaction in basic medium?
🟥 A) MnO₄⁻ → Mn²⁺
🟦 B) MnO₄⁻ → MnO₂
🟩 C) Cr₂O₇²⁻ → Cr³⁺
🟨 D) SO₂ → SO₄²⁻
26. Which of the following is a disproportionation reaction?
🟥 A) 2H₂ + O₂ → 2H₂O
🟦 B) 2H₂O₂ → 2H₂O + O₂
🟩 C) Zn + Cu²⁺ → Zn²⁺ + Cu
🟨 D) Fe²⁺ → Fe³⁺ + ē
27. Which of the following has the highest oxidation state?
🟥 A) Mn in MnO₂
🟦 B) Mn in KMnO₄
🟩 C) Mn in MnCl₂
🟨 D) Mn in MnSO₄
28. In the reaction:
2H₂ + CO → CH₃OH
Which element is reduced?
🟥 A) H
🟦 B) C
🟩 C) O
🟨 D) None
29. Which of the following is the oxidizing agent in the reaction:
2Al + 3Br₂ → 2AlBr₃
🟥 A) Al
🟦 B) Br₂
🟩 C) Al³⁺
🟨 D) Br⁻
30. Which of the following correctly represents the oxidation number change in:
SO₂ → SO₄²⁻
🟥 A) +4 to +6
🟦 B) +6 to +4
🟩 C) –2 to –4
🟨 D) +2 to +6
31. In the reaction:
Cr₂O₇²⁻ + 14H⁺ + 6ē → 2Cr³⁺ + 7H2O
The oxidation number of Cr changes from:
🟥 A) +6 to +3
🟦 B) +3 to +6
🟩 C) +7 to +2
🟨 D) +4 to +2
32. Which of the following is a redox reaction but not an acid–base reaction?
🟥 A) CaO + 2HCl → CaCl₂ + H₂O
🟦 B) Zn + 2HCl → ZnCl₂+H₂
🟩 C) NaOH + HCl → NaCl + H₂O
🟨 D) CH₃COOH + NaOH → CH₃COONa + H₂O
33. Which of the following represents a disproportionation of oxygen?
🟥 A) 2H₂ + O₂ → 2H₂O
🟦 B) 2H₂O₂ → 2H₂O + O₂
🟩 C) O₃ → O₂ + O
🟨 D) 2CO + O₂ → 2CO₂
34. Which of the following is the oxidizing agent in the reaction:
2Fe³⁺ + Sn²⁺ → 2Fe²⁺ + Sn⁴⁺
🟥 A) Fe³⁺
🟦 B) Fe²⁺
🟩 C) Sn²⁺
🟨 D) Sn⁴⁺
35. Which of the following correctly represents the oxidation number of chlorine in ClO₄⁻?
🟥 A) +3
🟦 B) +5
🟩 C) +7
🟨 D) +4
36. Which of the following is not a redox reaction?
🟥 A) AgNO₃ + NaCl→ AgCl + NaNO₃
🟦 B) 2Mg + O₂ → 2MgO₂
🟩 C) Zn + CuSO₄ → ZnSO₄ + Cu
🟨 D) 2Na + Cl₂ → 2NaCl
37. In the reaction:
2H₂S + SO₂ → 3S + 2H₂O
Which statement is correct?
🟥 A) H₂S is oxidized, SO₂ is reduced
🟦 B) H₂S is reduced, SO₂ is oxidized
🟩 C) Both are oxidized
🟨 D) Both are reduced
38. Which of the following elements shows variable oxidation states due to d‑orbital involvement?
🟥 A) Na
🟦 B) Mg
🟩 C) Fe
🟨 D) Al
39. In the reaction:
2KClO₃ → 2KCl+3O₂
The oxidation number of chlorine changes from:
🟥 A) +5 to –1
🟦 B) +5 to +1
🟩 C) +5 to +3
🟨 D) +5 to +7
40. Which of the following is the reducing agent in the reaction:
2Na + 2H₂O → 2NaOH + H₂
🟥 A) Na
🟦 B) H₂O
🟩 C) NaOH
🟨 D) H₂
✅ Answers & Reasons
1. B) Loss of electrons → Oxidation = electron loss, oxidation number increases.
2. B) Reduced → Decrease in oxidation number = gain of electrons.
3. A) Magnesium → Mg → Mg²⁺ (oxidized), O₂ → O²⁻ (reduced).
4. B) Sodium oxidized → Na loses electron → Na⁺.
5. A) Cu²⁺ → Gains electrons, acts as oxidizing agent.
6. A) Fe³⁺ → Gains electrons → reduced to Fe.
7. B) Low electronegativity → Metals lose electrons easily → reducing agents.
8. A) 2H₂ + O₂ → 2H₂O → H oxidized, O reduced → redox.
9. B) Oxygen increases → From –2 in H₂O to 0 in O₂ → oxidation.
10. A) It increases → Oxidation = loss of electrons → higher oxidation number.
11. C) Cl₂ → Cl₂ gains electrons → reduced to Cl⁻.
12. B) F₂ → Highest electronegativity, strongest oxidizing agent.
13. A) +7 to +2 → Mn reduces from +7 in MnO₄⁻ to +2 in Mn²⁺.
14. B) Cl₂ + NaOH → Same element (Cl) is oxidized and reduced → disproportionation.
15. A) CaCO₃ → CaO + CO₂ → No change in oxidation numbers → not redox.
16. B) Al oxidized, Cu reduced → Al → Al³⁺ (oxidized), Cu²⁺ → Cu (reduced).
17. B) Na → Low electronegativity, easily loses electrons → reducing agent.
18. C) Both oxidizing and reducing agent → H₂O₂ can oxidize or reduce depending on reaction.
19. A) –3 to +2 → N in NH₃ (–3) → N in NO (+2).
20. B) Ag⁺ → Gains electrons → reduced, hence oxidizing agent.
21. A) Zn → Zn loses electrons → reducing agent.
22.B) Oxidation → K → K⁺, loses electrons.
23.D) Na → Na is a reducing agent, not oxidizing.
24.B) Cl₂ reduced, Fe²⁺ oxidized → Cl₂ → Cl⁻, Fe²⁺ → Fe³⁺.
25.B) MnO₄⁻ → MnO₂ → In basic medium, permanganate reduces to MnO₂.
26.B) H₂O₂ → H₂O + O₂ → Same element (O) oxidized and reduced.
27.B) Mn in KMnO₄ → Mn = +7 (highest).
28.B) C → C in CO (+2) → C in CH₃OH (–2).
29.B) Br₂ → Gains electrons → reduced → oxidizing agent.
30.A) +4 to +6 → S in SO₂ (+4) → S in SO₄²⁻ (+6).
31.A) +6 to +3 → Cr in dichromate (+6) reduces to Cr³⁺.
32.B) Zn + HCl → Redox (Zn oxidized, H⁺ reduced), not acid–base.
33.B) H₂O₂ → H₂O + O₂ → O is both oxidized (–1 → 0) and reduced (–1 → –2).
34.A) Fe³⁺ → Gains electrons → reduced → oxidizing agent.
35.C) +7 → Cl in ClO₄⁻ is +7.
36.A) AgNO₃ + NaCl → Double displacement, no oxidation number change.
37.A) H₂S oxidized, SO₂ reduced → S in H₂S (–2 → 0), S in SO₂ (+4 → 0).
38.C) Fe → Transition metal, variable oxidation states.
39.A) +5 to –1 → Cl in KClO₃ (+5) → Cl in KCl (–1).
40.A) Na → Loses electrons → reducing agent.
🔥🌟Oxidation Number (O.N) OR Oxidation State (Oxi. No.)🧪💡
🔍Definition # 1 of Oxidation Number
In covalent bond formation the electrons are not transferred as in ionic bond formation, but partial transfer of electronic charge takes place, known as electron shift. The oxidation number method always assumes that there is a complete transfer of electrons from a less electronegative atom to a more electronegative atom.
Oxidation Number is a fictitious charges assigned to the atom of an element in a covalently bonded molecule by arbitrary conventions that results when the electrons in a covalent bond are assigned to the more electronegative atom (assuming the bonding were ionic) making certain the law of charge conservation is strictly obeyed. Oxidation Number is a purely a hypothetical number without any theoretical justification and it does not correspond (coincide with) to the real (actual) charge on the atoms, except in the special case of simple ionic compounds. Oxidation number of an atom in a molecule or ion is the hypothetical or real charge present on an atom due to electronegativity difference.
🔍Definition # 2 of Oxidation Number
The number of charges an atom would have in a molecule of a compound or polyatomic ion if bonding electrons were transferred completely in the direction indicated by the difference in electronegativity. Thus oxidation number reflects the number of electrons transferred in a covalent molecule or polyatomic ion. It is the number of electrons lost or gained by an atom of an element during its change from free state into a particular compound.
🔍Definition # 3 of Oxidation Number
The apparent charge (i.e. not real), either positive or negative or zero, on an atom of element in a molecule of a compound or in a polyatomic ion (radical) that results when the electrons in a covalent bond are assigned to the more electronegative atom is called oxidation number or oxidation state. Its value may be positive, negative or zero even fractional value ranges -4 to +7 (+8) depending upon the charge of combined atoms in the molecule or ion. (In Ni (CO)₄ oxidation number of Ni is zero).
🔍Definition # 4 of Oxidation Number
It is the fictitious charge that an atom appears to have in a given species when the bonding electrons are counted towards more electronegative atom i.e. it is the hypothetical charge an atom would possess in a compound if the bonding were completely ionic.
🔍Oxidation number as the degree of Oxidation
Oxidation number is the number with positive or negative sign which indicates the extent to which an element has been oxidized or reduced i.e. it shows the number of electrons which an atom has lost or gained as a result of bonding. The oxidation state is a “measure (or indicator) of the degree of oxidation” of an atom in a chemical compound. (In writing oxidation numbers, we will write the sign before the number to distinguish them from actual electronic charges, which we write with number first). It is the fundamental key to understanding redox reactions, reaction mechanisms, catalysis, etc.
🔍Examples of Oxidation Number
➡️1. Oxidation number of Mn in KMnO₄ is +7.
➡️2. Oxidation number of Cr in K₂Cr₂O₇ is +6.
➡️3. Oxidation number of Ni in Ni(CO)₄ is 0.
➡️4. Oxidation number of O in OF₂ is +2.
➡️5. Oxidation number of O in KO₂ is –½.
➡️6. Oxidation number of O in H₂O₂ is –1.
🔍Basis of Assigning Oxidation Number
The oxidation number for an element in a covalent compound is by taking the oxidation number to be equal to the charge that the element would carry, if all the bonds in the compound were regarded as ionic instead of covalent. In doing this, a shared pair or electrons between two atoms is assigned to the atom with the greater electronegativity. Or, if the two atoms are alike, the shared pair is split between the two, one electron being assigned to each atom. The resulting charges on the various atoms when the bonding electrons are so assigned are the oxidation numbers of the atoms.
🔍Covalency
It is the number of hydrogen atoms which can combine with a given atom. It is equal to the number of single bonds which an atom can form. It is also equal to the number of electrons an atom can share.
🔍Oxidation State
It is the oxidation number per atom.
🔍Difference between Valency and Oxidation number
Valency is a different term than oxidation number though sometimes the valency and the oxidation number of an element are same in a compound.
⚡1. Oxidation number is just the apparent charge (not necessarily actual) over the atom when the electrons are counted according to the arbitrary rules i.e. oxidation number is the number with positive or negative sign which indicates the extent to which an element has been oxidized or reduced. While valency is mere a number without positive or negative sign which expresses the combining or displacing tendency of an atom of an element and valency of an element is given by the number of electrons it actually loses or gains or shares during the formation of a compound.
⚡2. The oxidation no. of an atom may be in fraction, whereas the valency is always in whole number. The oxidation number of an atom in a compound may be zero but valency of an element cannot be zero (except noble gases).
⚡3. The oxidation state of an element may vary in its different compounds whereas in most of the cases, the valency of an element is constant.
⚡4. Valency and oxidation states of carbon in its different compounds give a good example to differentiate the two concepts. In CH₄, CH₃Cl, CH₂Cl₂, CHCl₃ and CCl₄, the valency of carbon is always four (due to sharing of four electrons) but its oxidation numbers is -4, -2, 0, +2 and +4 respectively.
🔍Oxidation State Vs Valency
🔍Some Important Points on Oxidation Number
➡️1. Oxidation number may be fractional.
➡️2. Oxidation number is positive in metallic elements
➡️3. Oxidation number is positive or negative in non-metallic elements
➡️4. Oxidation number is represented in Roman numbers in parenthesis (brackets) after the symbol of the metal in compounds
🔍Stock Notation
Representation of oxidation state of element by Roman numerals within parenthesis is known as stock notation i.e. Expressing the oxidation state of a metal by Roman numerals like I, II, III etc. within parenthesis is called stock notation.
e.g.
⚡Fe(II) SO₄ or FeSO₄ → Iron(II) sulphates (or ferrous sulphate)
⚡Fe(III) or FeCl₃ → iron(III) chloride (or ferric chloride)
⚡Au(III) Cl₃ or AuCl₃ → Gold(III) chloride (or auric chloride)
⚡Sn(II) Cl₂ or Sn(II) Cl₂ → Tin(II) chloride (or stannous chloride)
⚡Hg(II)Cl₂ or Hg(II)Cl₂ → Mercury(II) chloride (or mercuric chloride)
⚡Na₂CrO₄ → sodium chromate(VI)
➡️5. The oxidation number of metals in amalgams and metal carbonyls i.e. Ni(CO)₄, Fe(CO)₃, Cr(CO)₆ etc. is zero.
➡️6. In allotropic forms like diamond, graphite etc. oxidation number is 0.
➡️7. In case of coordinate bond, it gives +2 value of oxidation number to less electronegative atom and -2 values to more electronegative atom when coordinate bond is directed form less electronegative atom to more electronegative atom.
➡️8. If coordinate bond is directed from more electronegative to less electronegative atom then its contribution be zero for both the atoms.
➡️9. Oxidation number of O in compounds of fluorine is positive as F is the most electronegative element.
➡️10. Electronegativity values of no two elements are same
P > H, C > H, S > C, Cl > N
➡️11. Oxidation state of same element can be different in same or different compounds
O.N of S in H₂S = ‒ 2
O.N of S in H₂SO₃ = +4
O.N of S in H₂SO₄ = +6
🔍Details of Important Points of Oxidation Number
1. Positive and Negative Oxidation Numbers
2. Fractional Oxidation Number
3. Range of Oxidation Number
4. Maximum Oxidation Numbers
5. Oxidation Number measures covalent and ionic character
7. Oxidation number and group number
➡️1. Positive and Negative Oxidation Numbers
In general metallic elements (present at the farther left and middle of the periodic table) have only positive oxidation numbers (e.g. Li, Na, K, Mg, Ca, Ba, Al, Pb, Sn, Fe, Cu, etc.) whereas non-metallic elements may have either positive or negative oxidation numbers. But usually non-metals such as F, O, N and other halogens (Cl, Br, and I) have negative oxidation numbers.
➡️2. Fractional Oxidation Number
Elements as such do not have any fractional oxidation numbers. In reality no element can have a fractional oxidation state as electrons cannot be transferred in fraction.
When two or more atoms of an element are present in different oxidation states, then calculated oxidation number in a compound or ion may come out as fractional due to average of all the different oxidation states.
Fractional oxidation number is the average oxidation number. Fractional oxidation number of a particular element can be calculated only if we know about the structure of the compound in which it is present.
e.g.
⚡(i) In tetrathionate (S₄O₆²⁻) ion, the oxidation number of end S atoms is +5 each and that of the middle S atoms is 0 each. The total oxidation number of 4 S atoms is 5+0+0+5=+10 and the average oxidation number is 10 ÷ 4 = 2.5.
⚡(ii) In sodium tetrathionate (Na₂S₄O₆), the oxidation number of both S⁺ is equal to 0 (pure covalent bond) and other two terminal sulphur atoms have oxidation number = +5 having the structure
⚡(iii) in C₃O₂ the oxidation number of C is +4/3 or + 1.33.
➡️3. Range of Oxidation Number
Oxidation number of an atom in a molecule may be positive (metallic elements), negative (non-metallic elements) or any value ranges –4 to zero to +7 (or +8 in Os⁺⁸O₄, Ru⁺⁸O₄ etc.) or even fractional value e.g. Fe₃O₄ (Fe = +2.6), C₃O₄ (C = +2.6), C₃O₂ (C = +1.33).
Oxidation number of an atom in a molecule may have zero value. e.g. Ni(CO)₄ (Ni = 0), Fe(CO) ₅ (Fe = 0), C₆H₁₂O₆ (C = 0), C₁₂H₂₂O₁₁ (C = 0), CH₂O (C =0) etc.
The highest known oxidation state is +8 in the tetroxides of ruthenium, xenon, osmium, iridium, hassium, and some complexes involving plutonium; the lowest known oxidation state is −4 for some elements in the carbon group.
➡️4. Variable Oxidation Numbers
The transition metals of group B of the periodic table usually have several possible oxidation states except IIIB group (Sc, Y, La and Ac = +3) and IIB group (Zn, Cd = +2). The variable oxidation state of d-block elements is due to involvement of unpaired electrons of d-subshell. Similarly p-block elements exhibit variable oxidation states which is due to inert pair effect.
➡️5. Maximum Oxidation Numbers
Maximum oxidation number is always positive and maximum oxidation numbers of an atom in a molecule is equal to its group number. e.g. maximum oxidation number of Cl of group VIIA is +7 and that of Cr of group VIB is +6.
⚡ Os, Ru, Xe show maximum oxidation number i.e. +8
⚡ In 3d-series of transition metals, Mn shows maximum oxidation number of +7.
⚡ Maximum oxidation number of element = Group number in periodic table (applicable only for metals).
⚡ Minimum oxidation number of element = Group number – 8 (applicable only for non-metals).
Maximum or highest oxidation state is not stable. (Thus compounds containing central atom with its highest oxidation state are unstable and tend to decompose to reduce oxidation number
e.g.
KMnO₄ (Mn = +7)
AgNO₃ (N = +5)
Mn₂O₇ (Mn = +7)
HClO₄ (Cl = +7) etc.
⚡ For p-block elements (except F and O), the highest oxidation number is equal to their group number and lowest oxidation number is equal to the group number minus eight.
⚡ In transition elements the lowest oxidation number is equal to the number of ns electrons and highest oxidation number is equal to number of ‘ns’ and (n–1)d unpaired electrons. Maximum oxidation number of transition elements is given by:
Maximum oxidation number of atom = Number of ‘s’ electrons + Number of unpaired ‘d’ electrons
e.g.
Maximum oxidation number of Mn = 2 + 5 = +7 (group VIIA)
Maximum oxidation number of Fe = 2 + 4 = +6 (group VIIIA)
➡️6. Oxidation Number measures covalent and ionic character
A high oxidation number usually indicates significant (more) covalent character in the bonding of that compound
e.g. Mn₂O₇ (Mn = +7) and MnO₄⁻ (Mn = +7) ion have more covalent character.
Compounds with lower oxidation states have more ionic character
e.g. MnO₂ (Mn = +4) and Mn₂O₃ (Mn = +3) have significant ionic character.
➡️7. Oxidation number and group number
Oxidation number is directly related to the group number to which the element belongs.
e.g. the oxidation number of group IA is +1 and that of IIA is +2. Similarly zero group shows zero oxidation state.
Oxidation number of p-block elements is the number of electrons in the valence shell or deficiency of electrons in the valence shell.
➡️8. Oxidation number of two or more atoms of same elements may be different
If a compound contains two or more atoms of the same element, all of them may or may not have same oxidation number
e.g.
⚡(i) In Na₂S₂O₃, one S-atom has oxidation number = -2 while the other has oxidation number = +6.
⚡(ii) In bleaching powder; CaOCl₂ or Ca(OCl)Cl, oxidation number of one Cl = ‒1 while oxidation number of other Cl = +1.
⚡(iii) In Fe₃O₄ or FeO.Fe₂O₃, oxidaiton number of one Fe = +2 while that of each of the other two = +3.
⚡(iv) In NH₄NO₃, oxidation number of N of NH₄⁺ = ‒3 while that of N in NO₃⁻ = +5.
🔍How to get Oxidation Number
➡️1) Given a compound, write its Lewis dot structure.
➡️2) For each separate bond decide which element is most electronegative (EN).
➡️3) Give the most electronegative element all the electrons of that bond. If the atoms in the bond are the same, give each element half of the electrons.
➡️4) When all electrons have been assigned subtract the number of electrons on each atom from the valence of each element to get the oxidation state (number).
🔍Significance of Oxidation number
Oxidation number provides a measure of whether the atom in a molecule is neutral, electron rich or electron-poor. It guides us to identify elements that are oxidized (oxidation number increases) and reduced (oxidation number decreases) at a glance by comparing its oxidation number before and after the reaction.
Oxidation numbers are used:
➡️1. In nomenclature (naming) of compounds.
➡️2. In classifying types of reactions (as redox, non-redox or auto-redox).
➡️3. In balancing of equations of redox reactions.
➡️4. In examining trends in chemical reactivity across the periodic table.
➡️5. In exploring the systematic chemistry of elements.
➡️6. In identifying redox (oxidation-reduction) reactions.
➡️7. In determining the Equivalent weights
➡️8. in comparing the strength of acid and base
⚡(a) Strength of acids increases with increase in oxidation number.
⚡(b) Strength of base decreases with increase in oxidation number.
➡️9. In determining the oxidizing and reducing nature of compounds
⚡(a) If any compound is in maximum oxidation state, then it will act as oxidant only.
⚡(b) If any compound is in minimum oxidation state, then it will act as reductant only.
⚡(c) If the oxidation state is intermediate, then compound can act as both reductant as well as oxidant.
➡️10. To determine possible molecular formula of any compound
Suppose that there are three atoms A, B, C and their oxidation number are +6, ‒1, ‒2 respectively. Then the molecular formula of compound formed by them will be AB₄C because
+6 = (‒1×4)+( ‒2)
or +6 = ‒6
🔍Details
➡️In comparing the strength of acid and base
⚡Example 1
Order of acidic strength of oxyacids of chlorine in HClO, HClO₂, HClO₃ and HClO₄ will be HClO< HClO₂< HClO₃ < HClO₄ or HClO₄ > HClO₃ > HClO₂ > HClO. It is due to increasing oxidation number of Cl from +1 to +7 with incrasing number of oxygen.
⚡Example 2
Order of acidic strenght of oxyacids of nitrogen in H₂N₂O₂ (hyponitrous acid), HNO₂ (nitrous acid), and HNO₃ (nitric acid) will be
⚡Example 3
➡️Reason
H₃PO₂ (Hypophosphorous acid) = Ka ↑↑↑
H₃PO₃ (Phosphorous acid) = Ka ↑↑
H₃PO₄ (Ortho phosphoric acid = Ka ↑
➡️In determining the oxidizing and reducing nature of compounds
⚡(a) If central element in a compound is in maximum oxidation state, then compound acts as oxidant only.
⚡(b)If central element in a compound is in minimum oxidation state, then compound acts as reductant only.
⚡(c) If central element in a compound has intermediate oxidation state, then compound can act as both reductant as well as oxidant.
⚡(d) To calculate the equivalent weight of compounds
The equivalent weight of an oxidizing or reducing agent is that weight which accepts or loses one mole of electrons in a chemical reaction.
Equivalent weight of an oxidizing agent = Molecular weight/no. of electrons gained by one mole
⚡Example
In acidic medium, the oxidation number of dichromate is M/6 as central chromium atom undergoes reduction by gaining six electrons decreasing its oxidation state from +6 to +3.
Cr₂O₇²⁻ + 14H⁺ + 6ē → 2Cr³⁺ + 7H₂O
Equivalent weight of K₂Cr₂O₇ = MW/3 x 2 = M/6
Equivalent weight of a reductant = Molecular weight/no. of electrons lost by one mole
In acidic medium, the oxidation number of oxalate ion is M/2 as central carbon atom undergoes reduction by losing two electrons decreasing its oxidation state from +3 to +4.
Equivalent weight of C₂O₄²⁻ = MW/2 = M/2
In different conditions, a compound may have different equivalent weight because it depends upon the number of electrons gained or lost by that compound in that reaction.
⚡Example
MnO₄⁻(+7) → Mn²⁺ (+2) (Acidic medium)
Here five electrons are taken by MnO₄⁻, so its equivalent weight = M/5 = 158/5 = 31.6
MnO₄⁻(+7) → MnO₂ (+4) (Neutral or weak alkaline medium)
Here only three electrons are taken by MnO₄⁻, so its equivalent weight = M/3 = 158/3 = 52.7
(When only alkaline medium is given, consider it as weak alkaline medium)
MnO₄⁻(+7) → MnO₄²⁻(+6) (Strong alkaline medium)
Here only one electron is taken by MnO₄⁻, so its equivalent weight = M/1 = 158/1 = 158
Equivalent weight = Molecular weight/valency factor (General formula)
Valency factor = (Final ON – Initial ON) x No. of atms oxidized or reduced in 1 mole
Or
Valency factor = Change in ON x No. of atms oxidized or reduced in 1 mole
Note: KMnO₄ acts as an oxidant in every medium although with different strength which follows the order
Acidic medium > neutral medium > alkaline medium
Contrarily, K₂Cr₂O₇ acts as an oxidant only in acidic medium as follows
Cr₂O₇²⁻ → 2Cr³⁺
(2×+6) (2×+3)
Here, six electrons are gained by K₂Cr₂O₇, equivalent weight = M/6 = 294/6 = 49
🔍Summary of Rules for Assigning Oxidation State
🔍Rules for Finding Oxidation Number
➡️1. Oxidation Number of Free Elements is zero
The oxidation number of an atom in its elemental form or uncombined state is always zero i.e. the oxidation number of an element in a free atomic state (Na, H, Cl, O, P etc.) or in its poly-atomic state (graphite, H₂, O₂, P₄, S₈ etc.) or alloy form (Na/Hg) is always zero. E.g. K⁰, Cu⁰, H₂⁰, Cl₂⁰, O₂⁰. (Oxidation number is zero for any elemental substance, which occurs in diatomic, triatomic or polyatomic forms or allotropic forms (diamond, graphite) or alloy form (Na/Hg). (Free state = most stable state, uncombined state).
e.g. each atom in Na, Mg, C, O₂, N₂, H₂, Br₂, F₂, I₂, O₃, P₄, S₈, has an oxidation number of zero.
⚡(i) Oxidation state of atoms present in homoatomic molecules is zero. e.g. H₂, O₂, N₂, P₄, S₈ = zero
⚡(ii) Oxidation state of an element in any of its allotropic form is zero.
Cdiamond = 0, Cgraphite = 0, Smonoclinic = 0, Srhombic = 0
⚡(iii) Oxidation state of all the components of any an alloy are 0 e.g. Na⁰/Hg⁰
⚡(iv) In complex compounds, oxidation state of some neutral ligands is zero. e.g. CO, NO, H₂O, NH₃
➡️2. Oxidation Number of Monoatomic Ion equal to its charge
The oxidation number of atom in monoatomic ion (composed of only one atom) is equal to its charge.
e.g.
Oxidation number of Na in Na¹⁺ is +1, that of Ba in Ba²⁺ ion is +2, that of Al in Al³⁺ ion is +3, that of Ca in Ca²⁺ is +2, that of Cl in Cl⁻ion is –1, that of O in O²⁻ ion is –2, that of P in P³⁻ ion is –3 and so on.
➡️3. Oxidation Number of atoms in Polyatomic Ion
The oxidation number an atom in a polyatomic ion is usually equal to its oxidation number that it would have if it were a monoatomic ion.
For example;
In hydroxide ion (OH⁻), the oxygen atom has an oxidation number of –2 as if it were a monatomic oxide (O²⁻) ion and the hydrogen atom has an oxidation number of +1 as if it were simple H⁺ ion.
[In oxyanions, the oxidation number of central atom is always positive which is usually equal to its highest oxidation number e.g. in CO₃²⁻ ion, the oxidation number of carbon is +4 which is its highest oxidation state].
➡️4. Sum of Oxidation Numbers of all atoms in Polyatomic Ion equals net ionic charge
In polyatomic ions (or compound radical), the sum of oxidation numbers of its all atoms is equal to overall (net) charge of the ion.
e.g.
In the ammonium ion (NH₄⁺), the oxidation number of each H is +1 and that of N is –3. Thus the sum of the oxidation numbers is –3 + 4(+1) = +1, which is equal to net charge of the ion.
➡️5. Sum of Oxidation Numbers of all atoms in a molecule is always zero
In a neutral species (molecule of a compound), the sum of the oxidation numbers of all elements is always zero to comply with law of charge conservation. e.g.
[This rule is particularly useful for finding the oxidation number of an atom in difficult cases by assigning oxidation numbers to the ‘Easy’ atoms first and then find the oxidation number of the ‘Difficult‘ atom by subtraction].
➡️6. Oxidation Number of Atoms in Binary Polar Compounds
In binary polar compounds (those with two different elements), more electronegative element has negative oxidation number (equal to its charge in simple ionic compounds of the element) while less electronegative element has positive oxidation number. e.g.
➡️7. Oxidation Number of Atoms in Ternary Compounds
In ternary polar compounds (those with three or more different elements), only more electronegative element has negative oxidation number (equal to its normal oxidation number) while all other elements have positive oxidation numbers
⚡Exceptions
In few organic compounds and in few inorganic compounds where two more electronegative elements are present, more than one element has negative oxidation number while only one element has positive oxidation number.
➡️8. Oxidation Number of Fluorine is always –1
The oxidation number of Fluorine in its compounds is always –1.
(Due to restriction of negative oxidation number, F cannot form oxyacids or oxyanions for that it has to assign positive oxidation number)
➡️9. Oxidation Number of Other Halogens is usually –1
The oxidation number of other halogens (Cl, Br and I) in binary compounds where they occur as halide ion (X⁻) is usually –1. e.g.
The major exception is in compounds or ions of Cl, Br and I where they are bonded to oxygen atom e.g. in oxyacids, oxysalts and oxyanions of halogens like H⁺¹Cl⁺⁷O₄⁻⁸, Na⁺¹Cl⁺⁵O₃⁻⁶ etc.
➡️10. Oxidation Number of Hydrogen is mostly +1
➡️11. Oxidation Number of Oxygen
➡️12. Oxidation Number of Elements in Groups of the Periodic Table
The oxidation number of each element of Group IA (Li, Na, K, Rb, Cs), IIA (Be, Mg, Ca, Sr, Ba), IIIA (B,Al), IVA (C, Si, Ge, Sn, Pb), VA (N, P, As, Sb, Bi) and VIA (O, S, Se, Te), in their compounds is +1, +2, +3, –4/+4, –3/+5, –2, –1, +1 to +6 respectively. (The most common oxidation state of group IIIA is +3 but it can also show +1 oxidation state due to inert pair effect)
🔍Some helping rules for calculating oxidation number
🔍Quick Recap
➡️What is Oxidation Number? 🧐
Oxidation Number (ON) or oxidation state is the charge an atom would have if electrons were transferred completely in a compound.
It helps in understanding redox (oxidation & reduction) reactions by tracking electron loss (oxidation) and electron gain (reduction) during reactions 🔄.
💡 Rules for Assigning Oxidation Numbers 📝:
➡️Free/uncombined Pure Elements (like H₂, O₂, N₂, Cl₂, Na, Zn) → Oxidation number = 0 🏔️
➡️Monatomic ions → Oxidation number = Ion Charge (e.g., Na⁺ = +1, Cl⁻ = -1) ⚡
➡️Molecule of a compound → Sum of oxidation numbers = 0 (e.g., in H₂O: (2×+1) + (-2) = 0)🏁
➡️Polyatomic Ions → Sum of oxidation numbers = Ion Charge (e.g., in SO₄²⁻: (S = +6, O = -2) → (+6) + (4×-2) = -2) 💥
➡️Hydrogen (H) = +1 when bonded to non‑metals (HCl, H₂O) (except with metals in hydrides like NaH, CaH₂, where it’s -1) 🔋
➡️Oxygen (O) = -2 (except in peroxides like H₂O₂, Na₂O₂, where it’s -1, in superoxides where it is -½, +2 in OF₂ ) 🧪
➡️Fluorine (F) = -1 always! (most electronegative).🦷
➡️Group 1 Elements (Li, Na, K, etc.) = +1 🤖
➡️Group 2 Elements (Mg, Ca, etc.) = +2 ⚙️
🔥⚡ Identifying Oxidation & Reduction
⚡Oxidation = increase in oxidation number (loss of electrons) ➡️ 🧑🔬
⚡Reduction = decrease in oxidation number (gain of electrons) ➡️ ✨
⚡Oxidizing agent = species reduced, decrease in oxidation number (gains electrons) ➡️ 💧
⚡Reducing agent = species oxidized, increase in oxidation number (loses electrons) ➡️ 🤖
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