Hello future chemists! Welcome to this exciting journey into the world of Sulphur and its fascinating compounds. Today, we're going to unravel some fundamental concepts about Sulphur's oxoacids, and two very important compounds: Hydrogen Sulphide (H2S) and Sulphur Dioxide (SO2). Don't worry, we'll start from the very beginning and build up our understanding step-by-step!

### 1. Understanding Oxoacids: Sulphur's Acidic Family

First things first, what exactly is an oxoacid? It sounds fancy, but it's quite simple! An oxoacid is basically an acid that contains

• at least one oxygen atom
• at least one hydrogen atom (which is acidic and can be released as H+)
• and at least one other element (in our case, Sulphur)

So, in an oxoacid of sulphur, you'll always find sulphur, oxygen, and hydrogen bonded together.

Think of Sulphur as a versatile chef in the chemical kitchen. It can combine hydrogen and oxygen in various proportions and arrangements to create a whole menu of acidic "dishes" – these are its oxoacids! The most famous one, of course, is Sulphuric Acid (H₂SO₄), which you might know as the "King of Chemicals." But Sulphur has many other relatives in this acidic family.

Why are they important? Well, sulphuric acid alone is vital for countless industrial processes, from fertilizers to detergents to car batteries. Understanding these acids helps us grasp their reactivity and applications.

The general structure of an oxoacid of sulphur usually involves the sulphur atom at the center, bonded to oxygen atoms. Some oxygen atoms will also be bonded to hydrogen atoms, forming -OH groups. These -OH groups are where the acidic hydrogen comes from! Other oxygen atoms might be double-bonded to sulphur, forming =O groups.

When we talk about oxoacids, we often consider the oxidation state of the central sulphur atom. This tells us how many electrons sulphur has "lost" or "gained" (conceptually) in forming bonds. For example:

• In Sulphurous Acid (H₂SO₃), Sulphur has an oxidation state of +4.


• In Sulphuric Acid (H₂SO₄), Sulphur has an oxidation state of +6.

This difference in oxidation state is key to understanding their different chemical behaviors, especially in redox reactions!

CBSE vs. JEE Focus: For CBSE, knowing the common oxoacids like H₂SO₃ and H₂SO₄, their formulas, and basic properties (like being acids) is important. For JEE, you'll need to delve deeper into their structures, the oxidation states of sulphur, their relative acid strengths, and their redox properties.

### 2. Hydrogen Sulphide (H₂S): The Rotten Egg Gas

Imagine a colorless gas that smells exactly like rotten eggs. That's Hydrogen Sulphide (H₂S)! It's famously known for this pungent odor.

Structure and Bonding:
Just like water (H₂O), H₂S has a bent molecular geometry. The central sulphur atom is bonded to two hydrogen atoms. Due to the lone pairs on sulphur, the molecule isn't linear but forms an angle.

Physical Properties - The Basics:

• Colorless gas: You can't see it, only smell it.


• Rotten egg smell: This is its defining characteristic! Be warned, it's highly poisonous even in low concentrations. Your nose gets fatigued quickly, so if you stop smelling it, it doesn't mean it's gone!


• Slightly heavier than air: It tends to accumulate in low-lying areas.


• Soluble in water: It dissolves to form hydrosulphuric acid, which is a weak acid.

Chemical Properties - Getting Reactive!

Let's look at how H₂S behaves chemically:

1. Acidic Nature (Weak Acid):
   When H₂S dissolves in water, it forms an acidic solution (hydrosulphuric acid). It's a diprotic acid, meaning it can donate two protons (H⁺ ions) in two stages, but it's a very weak acid compared to something like HCl or H₂SO₄.
   
   
   H₂S (aq) + H₂O (l) ⇌ H₃O⁺ (aq) + HS⁻ (aq)
   
   
   HS⁻ (aq) + H₂O (l) ⇌ H₃O⁺ (aq) + S²⁻ (aq)
   
   
   This weak acidic nature allows it to react with bases.
   


2. Strong Reducing Agent:
   This is a super important property! In H₂S, sulphur is in its lowest possible oxidation state: -2. Since it can't go any lower, it loves to lose electrons and get oxidized to a higher oxidation state (like 0 in S, +4 in SO₂, or +6 in SO₄²⁻). This makes it a powerful reducing agent.
   
   
   

   Example 1: Reaction with halogens (Oxidation to Sulphur)

   
   H₂S reduces halogens (like chlorine, Cl₂) to halide ions while itself getting oxidized to elemental sulphur.
   
   
   H₂S (g) + Cl₂ (g) → 2HCl (g) + S (s)
   
   
   (Here, Sulphur goes from -2 to 0)
   
   
   

   Example 2: Reaction with SO₂ (Redox Reaction)

   
   Yes, H₂S can even react with SO₂!
   
   
   2H₂S (g) + SO₂ (g) → 3S (s) + 2H₂O (l)
   
   
   (H₂S is oxidized, SO₂ is reduced. This is an important reaction for sulphur recovery!)
   


3. Combustion:
   H₂S is flammable and burns in air (oxygen) to produce sulphur dioxide and water. If there isn't enough oxygen, it might just produce elemental sulphur.
   
   
   2H₂S (g) + 3O₂ (g) → 2SO₂ (g) + 2H₂O (l) (Complete combustion)
   
   
   2H₂S (g) + O₂ (g) → 2S (s) + 2H₂O (l) (Incomplete combustion)
   


4. Precipitation of Metal Sulphides:
   This is a classic reaction in qualitative analysis! H₂S is often used to precipitate metal sulphides from solutions, which are often brightly colored and insoluble.
   
   
   

   Example: Reaction with Lead acetate

   
   Pb(CH₃COO)₂ (aq) + H₂S (g) → PbS (s) ↓ + 2CH₃COOH (aq)
   
   
   (Lead sulphide, PbS, is a black precipitate)
   
   
   This property is used to identify and separate different metal ions in solution.
   

CBSE vs. JEE Focus: For CBSE, remember its smell, weak acidic nature, and its role as a reducing agent and in precipitating metal sulphides. For JEE, a deeper understanding of its redox chemistry (how its oxidation state changes) and its quantitative use in analytical chemistry is crucial.

### 3. Sulphur Dioxide (SO₂): The Choking Gas

Next up, we have Sulphur Dioxide (SO₂). This gas is often associated with volcanic eruptions and industrial pollution. It has a characteristic pungent, choking smell.

Structure and Bonding:
SO₂ also has a bent molecular geometry, similar to H₂S and H₂O. The central sulphur atom is double-bonded to two oxygen atoms, and it also has a lone pair of electrons, which contributes to its bent shape. It exhibits resonance, meaning its actual structure is an average of two contributing structures, giving equal bond lengths to both S-O bonds.

Physical Properties - The Basics:

• Colorless gas: Just like H₂S.


• Pungent, choking smell: Very irritating to the respiratory system.


• Heavier than air: Tends to settle.


• Highly soluble in water: Much more soluble than H₂S. It reacts with water to form sulphurous acid (H₂SO₃).

Chemical Properties - The Chameleon!

SO₂ is quite interesting because it can act as both a reducing and, in some cases, an oxidizing agent. This is because sulphur in SO₂ has an oxidation state of +4. This means it can:

• Be oxidized to +6 (e.g., in H₂SO₄) – acting as a reducing agent.


• Be reduced to 0 (elemental S) or -2 (sulphide) – acting as an oxidizing agent.

Let's explore its typical reactions:

1. Acidic Nature:
   When SO₂ dissolves in water, it forms sulphurous acid, a weak acid.
   
   
   SO₂ (g) + H₂O (l) ⇌ H₂SO₃ (aq)
   
   
   This H₂SO₃ can then react with bases.
   


2. Strong Reducing Agent (Most Common Behavior):
   Because sulphur in SO₂ is at +4, it often gets oxidized to +6, making it a powerful reducing agent, especially in the presence of strong oxidizing agents.
   
   
   

   Example 1: Reduction of acidified Potassium Permanganate (KMnO₄)

   
   SO₂ decolorize pink/purple acidified KMnO₄ solution, which is a classic test for SO₂.
   
   
   5SO₂ (g) + 2KMnO₄ (aq) + 2H₂O (l) → K₂SO₄ (aq) + 2MnSO₄ (aq) + 2H₂SO₄ (aq)
   
   
   (Sulphur goes from +4 to +6, Manganese goes from +7 to +2)
   
   
   

   Example 2: Reduction of Potassium Dichromate (K₂Cr₂O₇)

   
   SO₂ turns orange acidified K₂Cr₂O₇ solution green.
   
   
   3SO₂ (g) + K₂Cr₂O₇ (aq) + H₂SO₄ (aq) → K₂SO₄ (aq) + Cr₂(SO₄)₃ (aq) + H₂O (l)
   
   
   (Sulphur goes from +4 to +6, Chromium goes from +6 to +3)
   


3. Bleaching Action:
   SO₂ acts as a bleaching agent. It usually bleaches by reduction (unlike Cl₂, which bleaches by oxidation). Its bleaching effect is temporary and reversed by air. It's often used for delicate materials like silk, wool, and straw.
   
   
   Colored matter + SO₂ + H₂O → Colorless reduced matter + H₂SO₄
   
   
   (The colored substance is reduced and loses its color.)
   


4. Oxidizing Agent (Less Common):
   SO₂ can also act as an oxidizing agent, though it's less common than its reducing nature. It gets reduced to a lower oxidation state, typically elemental sulphur (0) or even sulphide (-2).
   
   
   

   Example: Reaction with H₂S (as seen before)

   
   2H₂S (g) + SO₂ (g) → 3S (s) + 2H₂O (l)
   
   
   (Here, SO₂ is reduced from +4 to 0, while H₂S is oxidized.)
   

CBSE vs. JEE Focus: For CBSE, focus on SO₂'s acidic nature, its common role as a reducing agent (decolorizing KMnO₄, turning dichromate green), and its temporary bleaching action. For JEE, you'll need to understand the nuances of its dual nature (both reducing and oxidizing), its resonance structure, and applications in pollution control and industrial chemistry (e.g., Contact process for H₂SO₄ where SO₂ is oxidized).

So there you have it – a foundational look at oxoacids of sulphur, H₂S, and SO₂! Keep these basics strong, and you'll be well-prepared to tackle more complex aspects of these fascinating compounds. Keep practicing and stay curious!

Welcome back, future IITians! Today, we're diving deep into the fascinating world of sulphur and its compounds, specifically focusing on hydrogen sulphide (H₂S), sulphur dioxide (SO₂), and the diverse family of its oxoacids. These compounds are crucial from both theoretical and practical standpoints, frequently appearing in JEE questions due to their unique properties and varied oxidation states.

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### 1. Hydrogen Sulphide (H₂S): The Rotten Egg Gas

Hydrogen sulphide, commonly known as rotten egg gas, is a compound of sulphur in its lowest possible oxidation state of -2. This low oxidation state is key to understanding its primary chemical characteristic: its strong reducing power.

#### 1.1 Structure of H₂S

H₂S has a bent or V-shaped molecular geometry, similar to water (H₂O). The central sulphur atom undergoes sp³ hybridization, leading to a tetrahedral electron geometry, but two of these hybrid orbitals are occupied by lone pairs of electrons.
* Bond Angle: Approximately 92°, which is less than the ideal tetrahedral angle (109.5°) due to the repulsion from the two lone pairs on sulphur.
* Polarity: H₂S is a polar molecule due to the bent shape and the difference in electronegativity between sulphur and hydrogen.

#### 1.2 Preparation of H₂S

1. Laboratory Method: By the action of dilute non-oxidizing acids on metal sulphides.


Example: Reaction of iron(II) sulphide with dilute sulphuric acid.

FeS (s) + H₂SO₄ (aq) → FeSO₄ (aq) + H₂S (g)

ZnS (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂S (g)

2. Industrial Method:
* From natural gas: H₂S is a significant component of "sour" natural gas and is removed before distribution.
* From the reaction of hydrogen with molten sulphur:

H₂ (g) + S (l) → H₂S (g) (at high temperatures)

#### 1.3 Properties of H₂S

##### 1.3.1 Physical Properties:
* Colorless gas with a characteristic rotten egg smell.
* Highly poisonous, even in low concentrations. Prolonged exposure can lead to paralysis of the olfactory nerves, making its detection by smell unreliable.
* Slightly heavier than air.
* Moderately soluble in water, forming a weak acidic solution called hydrosulphuric acid.

##### 1.3.2 Chemical Properties:

1. Acidic Nature:
H₂S is a weak dibasic acid (can donate two protons). It dissociates in two steps in water:

H₂S (aq) ⇌ H⁺ (aq) + HS⁻ (aq)     (Ka1 = 1.0 × 10⁻⁷)

HS⁻ (aq) ⇌ H⁺ (aq) + S²⁻ (aq)       (Ka2 = 1.2 × 10⁻¹³)

It reacts with bases to form salts (sulphides or hydrosulphides).


Example:

H₂S + NaOH → NaHS (sodium hydrosulphide) + H₂O

H₂S + 2NaOH → Na₂S (sodium sulphide) + 2H₂O

JEE Tip: The K_a values indicate that H₂S is a very weak acid, much weaker than common mineral acids. Its second dissociation constant (K_a2) is extremely small, meaning S²⁻ ions are present in very low concentrations in aqueous solutions of H₂S.

2. Strong Reducing Agent:
This is the most significant chemical property of H₂S. Sulphur in H₂S is in its lowest oxidation state (-2) and can only increase its oxidation state (e.g., to 0 in S, +4 in SO₂, or +6 in H₂SO₄).


Examples:
* Reaction with Halogens: H₂S reduces halogens to halide ions, getting oxidized to free sulphur.

H₂S (g) + Cl₂ (g) → 2HCl (g) + S (s)  

H₂S (g) + Br₂ (aq) → 2HBr (aq) + S (s)

* Reaction with Sulphur Dioxide (SO₂): An important reaction where both act as redox agents.

2H₂S (g) + SO₂ (g) → 3S (s) + 2H₂O (l)

(Here, S in H₂S goes from -2 to 0, and S in SO₂ goes from +4 to 0).
* Reaction with Oxidizing Agents: H₂S readily reduces strong oxidizing agents like potassium permanganate (KMnO₄), potassium dichromate (K₂Cr₂O₇), nitric acid (HNO₃), and iron(III) salts.

5H₂S + 2KMnO₄ + 3H₂SO₄ → 5S + 2MnSO₄ + K₂SO₄ + 8H₂O

(Mn goes from +7 to +2, S from -2 to 0)

H₂S + 2FeCl₃ → 2FeCl₂ + S + 2HCl

(Fe goes from +3 to +2, S from -2 to 0)

3. Combustion:
* In sufficient oxygen: Burns with a blue flame to form sulphur dioxide and water.

2H₂S (g) + 3O₂ (g) → 2SO₂ (g) + 2H₂O (l)

* In insufficient oxygen: Forms sulphur and water.

2H₂S (g) + O₂ (g) → 2S (s) + 2H₂O (l)

4. Precipitation of Metal Sulphides:
H₂S is widely used in qualitative analysis for the precipitation of metal sulphides. The selective precipitation of various metal ions as their sulphides, depending on the pH of the solution, is a key concept.


Example:
* In acidic medium, only highly insoluble sulphides (e.g., HgS, PbS, CdS, Bi₂S₃, CuS, As₂S₃, Sb₂S₃, SnS₂) precipitate because the concentration of S²⁻ ions is very low.
* In alkaline medium, the S²⁻ concentration is higher, leading to precipitation of less insoluble sulphides (e.g., FeS, NiS, CoS, MnS, ZnS).


Remember: The solubility product (Ksp) and the common ion effect play a crucial role here.

#### 1.4 Uses of H₂S
* Chiefly used as an analytical reagent in qualitative analysis for the precipitation of metal sulphides.

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### 2. Sulphur Dioxide (SO₂): The Pungent Gas

Sulphur dioxide is another important oxide of sulphur, where sulphur exhibits a +4 oxidation state. This intermediate oxidation state allows SO₂ to act as both an oxidizing and a reducing agent.

#### 2.1 Structure of SO₂

SO₂ has a bent or V-shaped geometry. The central sulphur atom is sp² hybridized. It exhibits resonance between two equivalent structures:

O=S-O⁻ ↔ ⁻O-S=O

Both S-O bond lengths are identical and intermediate between a single and a double bond. There is one lone pair on the sulphur atom.
* Bond Angle: Approximately 119.5°.
* Polarity: SO₂ is a polar molecule.

#### 2.2 Preparation of SO₂

1. Laboratory Method:
* By burning sulphur in air:

S (s) + O₂ (g) → SO₂ (g)

* By the action of dilute sulphuric acid on sulphites:

Na₂SO₃ (s) + H₂SO₄ (aq) → Na₂SO₄ (aq) + H₂O (l) + SO₂ (g)

(Or from NaHSO₃ with HCl)

2. Industrial Method:
* Roasting of sulphide ores: This is the most common industrial method. Metal sulphides are heated in air to convert them into metal oxides and sulphur dioxide.

4FeS₂ (s) + 11O₂ (g) → 2Fe₂O₃ (s) + 8SO₂ (g)

2ZnS (s) + 3O₂ (g) → 2ZnO (s) + 2SO₂ (g)

* By-product of power generation: Burning of fossil fuels containing sulphur.

#### 2.3 Properties of SO₂

##### 2.3.1 Physical Properties:
* Colorless gas with a pungent and suffocating smell.
* Highly soluble in water.
* Easily liquefiable at room temperature under pressure (boiling point -10°C).
* Heavier than air.

##### 2.3.2 Chemical Properties:

1. Acidic Nature:
SO₂ is an acidic oxide. It dissolves in water to form sulphurous acid (H₂SO₃), a weak dibasic acid. Note that H₂SO₃ is an unstable acid and exists only in aqueous solution.

SO₂ (g) + H₂O (l) ⇌ H₂SO₃ (aq)

Sulphurous acid reacts with bases to form sulphites and bisulphites.


Example:

SO₂ + 2NaOH → Na₂SO₃ (sodium sulphite) + H₂O

SO₂ + NaOH → NaHSO₃ (sodium bisulphite) + H₂O

2. Reducing Agent:
Since sulphur is in the +4 oxidation state, it can be oxidized further to +6. This makes SO₂ a strong reducing agent, especially in the presence of water.


Examples:
* Reaction with Halogens: Reduces halogens to hydrohalic acids, forming sulphuric acid.

SO₂ (g) + Cl₂ (g) + 2H₂O (l) → H₂SO₄ (aq) + 2HCl (aq)

(S goes from +4 to +6, Cl from 0 to -1)
* Reaction with KMnO₄: Decolorizes acidified potassium permanganate solution (purple to colorless).

5SO₂ + 2KMnO₄ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄

(Mn goes from +7 to +2, S from +4 to +6)
* Reaction with K₂Cr₂O₇: Changes acidified potassium dichromate solution from orange to green.

3SO₂ + K₂Cr₂O₇ + H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + H₂O

(Cr goes from +6 to +3, S from +4 to +6)
* Bleaching Action: SO₂ acts as a temporary bleaching agent for delicate articles like silk, wool, and straw. It bleaches by reduction (formation of colorless compounds). The bleached color can be restored on exposure to air (oxidation).

Colored substance + SO₂ + H₂O → Reduced colorless product + H₂SO₄

Contrast with Cl₂: Chlorine bleaches by oxidation, which is permanent.

3. Oxidizing Agent:
Sulphur in SO₂ (+4 oxidation state) can also be reduced to lower oxidation states (0 in S, or -2 in H₂S). Thus, SO₂ also acts as an oxidizing agent, especially with strong reducing agents.


Examples:
* Reaction with H₂S:

SO₂ (g) + 2H₂S (g) → 3S (s) + 2H₂O (l)

(S in SO₂ goes from +4 to 0, S in H₂S from -2 to 0)
* Reaction with Magnesium:

2Mg (s) + SO₂ (g) → 2MgO (s) + S (s)

(Mg goes from 0 to +2, S from +4 to 0)
* Reaction with Carbon:

C (s) + 2SO₂ (g) → CO₂ (g) + 2S (s)

(C goes from 0 to +4, S from +4 to 0)

4. Addition Reactions:
SO₂ can form addition compounds, e.g., with unsaturated hydrocarbons (in specific conditions) or with sodium sulphite (forming sodium thiosulphate).

#### 2.4 Uses of SO₂
* Manufacture of sulphuric acid (Contact Process).
* Bleaching agent for wool, silk, straw, etc.
* Antichlor (removes excess chlorine) and disinfectant.
* Preservative for food (e.g., fruit juices).
* Refrigerant (liquid SO₂).

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### 3. Oxoacids of Sulphur: A Diverse Family

Sulphur, being a versatile element, forms a large number of oxoacids (acids containing oxygen, hydrogen, and sulphur). The ability of sulphur to exist in various oxidation states (+2, +4, +5, +6) and to form both S-S bonds and S-O-S linkages contributes to this diversity.

#### 3.1 General Features of Oxoacids of Sulphur

* All oxoacids of sulphur contain at least one S=O (sulphuryl) group and at least one S-OH (hydroxyl) group. The hydrogen atom in the S-OH group is acidic.
* Some oxoacids contain S-S linkages (e.g., polythionic acids, dithionic acid), while others contain S-O-S (pyro-series) or S-O-O-H (peroxo-series) linkages.
* The oxidation state of sulphur is crucial for characterizing these acids.

#### 3.2 Important Oxoacids of Sulphur (with structures & oxidation states)

Let's explore the key oxoacids, focusing on their structures, which are frequently asked in JEE exams. To determine the oxidation state of sulphur, we generally assign -2 to oxygen and +1 to hydrogen, then balance the charge. If an S-S bond is present, the oxidation state calculation needs careful consideration (divide the total charge contribution equally if the sulphurs are identical, or calculate individually if different).

Name	Formula	Oxidation State of S	Structure (Simplified)	Key Features / JEE Focus
Sulphurous Acid	H₂SO₃	+4	O=S(OH)₂ (Exists in solution only)	Weak dibasic acid. Sulphur is sp³ hybridized. Not isolable.
Thiosulphuric Acid	H₂S₂O₃	+2 (average) (+6, -2 for the two S atoms respectively)	(One S is like central S of H2SO4, other S replaces an O)	Contains an S-S bond. Unstable, disproportionates. Key in photography (hypo).
Sulphuric Acid	H₂SO₄	+6	O₂S(OH)₂ (S has 4 bonds to O: 2=O, 2-OH)	Strong dibasic acid. The most important oxoacid. Sulphur is sp³ hybridized.
Pyrosulphuric Acid (Oleum)	H₂S₂O₇	+6	(HO)O₂S-O-SO₂OH (Contains an S-O-S linkage)	Also known as fuming sulphuric acid. Produced by dissolving SO₃ in H₂SO₄. Key intermediate in contact process.
Peroxomonosulphuric Acid (Caro's Acid)	H₂SO₅	+6	HO-O-SO₂OH (Contains one S-O-O-H peroxo linkage)	Strong oxidizing agent due to the peroxo bond. Sulphur is sp³ hybridized.
Peroxodisulphuric Acid (Marshall's Acid)	H₂S₂O₈	+6	HO₃S-O-O-SO₃H (Contains one S-O-O-S peroxo linkage)	Very strong oxidizing agent. Used in the synthesis of H₂O₂. Both S atoms are sp³ hybridized.
Dithionic Acid	H₂S₂O₆	+5	HO₃S-SO₃H (Contains an S-S bond)	S-S bond is directly between the two sulphur atoms. The two S atoms are equivalent, each in +5 oxidation state.
Polythionic Acids	H₂SnO₆ (n = 2-6)	Variable (e.g., in H₂S₄O₆, two terminal S are +5, two central S are 0)	HO₃S-(S)n-2-SO₃H (Contains a chain of 'n-2' sulphur atoms between two -SO₃H groups)	Family of acids with variable number of sulphur atoms linked in a chain. The central S atoms are in zero oxidation state.

#### 3.3 Drawing Structures and Determining Oxidation States (JEE Focus)

For JEE, being able to draw the structure and assign oxidation states is vital.
Key Rules:
1. Sulphur usually forms four or six bonds.
2. Oxygen usually forms two bonds (S=O or S-O-H).
3. Hydrogen forms one bond (O-H).
4. Each S-OH group contributes one acidic proton.
5. Oxidation State Calculation:
* Assign O as -2 (unless in a peroxo linkage, where it's -1).
* Assign H as +1.
* Sum the charges and equate to the overall charge of the molecule (0 for neutral molecules).
* For S-S bonds, consider the electron sharing. If two S atoms are bonded, they contribute equally to the bond, so each S-S bond doesn't directly change the individual oxidation state. However, in polythionic acids, the central 'n-2' S atoms are often considered to be in 0 oxidation state, while the terminal S atoms in the -SO₃H groups are in +5.

Example: Let's determine the oxidation state of S in H₂SO₅ (Caro's Acid):
The structure is HO-O-SO₂OH.
* One H is +1.
* The O in the S-OH group is -2.
* The two O atoms double-bonded to S are -2 each.
* The O-O (peroxo) linkage has each O as -1.
Let x be the oxidation state of S.
Equation: (2 * +1 for H) + (1 * -1 for peroxo O) + (1 * -1 for peroxo O) + (1 * -2 for OH O) + (2 * -2 for S=O) + x = 0
2 - 1 - 1 - 2 - 4 + x = 0
x - 6 = 0
x = +6

#### 3.4 Acidity of Oxoacids

The acidity of oxoacids generally increases with:
1. Increase in the number of non-hydroxyl oxygen atoms: These oxygen atoms are double-bonded to sulphur (S=O), making sulphur more electronegative, thus pulling electron density from the S-OH bond and making the H more easily dissociable.


Example: H₂SO₄ (two S=O) is a much stronger acid than H₂SO₃ (one S=O).
2. Increase in the oxidation state of the central atom: A higher positive oxidation state means greater electron-withdrawing capacity, strengthening the acid.


Example: S in H₂SO₄ (+6) vs. S in H₂SO₃ (+4).

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By thoroughly understanding the structure, preparation, and properties of H₂S and SO₂, along with the various oxoacids of sulphur, you'll be well-equipped to tackle the advanced problems in JEE. Pay close attention to the redox reactions of H₂S and SO₂, and master the structural aspects and oxidation state determination for the oxoacids. Good luck!

Here are some mnemonics and shortcuts to help you remember the key aspects of Oxoacids of Sulphur, H2S, and SO2 properties for your JEE and CBSE exams.

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1. Oxoacids of Sulphur: Names, Formulas, and Oxidation States

Remembering the common oxoacids and their properties can be tricky. Use these shortcuts:

• 
  "OUS" vs. "IC" acids:
  
  
  
  • Shortcut: "OUS is Less Oxygen, IC is More Oxygen."
  
  
  • Example:
    
    H₂SO₃ (SulphurOUS acid, S = +4)
    
    H₂SO₄ (SulphurIC acid, S = +6)
    
    This rule generally applies to common oxyacids.
    
  
  
  
  


• 
  Peroxoacids (Caro's and Marshall's):
  
  
  
  • Mnemonic: "CArO's is one O, MARshAll's has two O's." (Focus on the 'O' for peroxo-linkage)
  
  
  • Formulas:
    
    
    
    • Caro's Acid (Peroxomonosulphuric acid): H₂SO₅ (one peroxo linkage -O-O-)
    
    
    • Marshall's Acid (Peroxodisulphuric acid): H₂S₂O₈ (one peroxo linkage -O-O-)
    
    
    
    
  
  
  
  


• 
  Pyrosulphuric Acid (Oleum):
  
  
  
  • Shortcut: "Oleum is two Sulphuric acids, less one water."
  
  
  • Formula: 2H₂SO₄ - H₂O = H₂S₂O₇
  
  
  
  


• 
  Thiosulphuric Acid:
  
  
  
  • Mnemonic: "THIO means 'S' instead of 'O'."
  
  
  • Formula: From H₂SO₄, replace one oxygen with sulphur to get H₂S₂O₃.
  
  
  
  

---

2. Hydrogen Sulphide (H₂S) Properties

Remember the three main characteristics of H₂S using this simple phrase:

• 
  Mnemonic: "H₂S: Rotten, Reducing, Precipitates."
  
  
  
  • Rotten: Distinctive smell of rotten eggs. (Helps identify the gas)
  
  
  • Reducing: H₂S is a strong reducing agent. The sulfur in H₂S is in its lowest oxidation state (-2) and readily gets oxidized (e.g., to S, SO₂, SO₄^2-).
    
    
    
    • JEE Tip: Remember it reduces FeCl₃ to FeCl₂, and oxidises to colloidal sulfur with oxidising agents like HNO₃, KMnO₄, etc.
    
    
    
    
  
  
  • Precipitates: H₂S precipitates heavy metal ions as insoluble metal sulfides (e.g., CuS, PbS, CdS). This is a crucial analytical test.
  
  
  
  

---

3. Sulphur Dioxide (SO₂) Properties

SO₂ is versatile! Remember its properties with:

• 
  Mnemonic: "SO₂ is an Acidic, Bleaching RedOX agent."
  
  
  
  • Acidic: It's an acidic oxide. It dissolves in water to form sulphurous acid (H₂SO₃).
    
    
    
    • SO₂(g) + H₂O(l) → H₂SO₃(aq)
    
    
    
    
  
  
  • Bleaching: SO₂ acts as a bleaching agent.
  
  
  
  
  • Shortcut: "SO₂'s bleaching is Temporary and Reductive." (Turns back to original color on exposure to air/oxidation)
  
  
  • It bleaches by reduction (forming colorless addition products), not by oxidation.
  
  
  
  
  • RedOX Agent: Sulfur in SO₂ is in the +4 oxidation state. This intermediate state allows it to act as both:
    
    
    
    • Reducing Agent: S(+4) → S(+6). It reduces strong oxidising agents like KMnO₄ (purple to colorless), K₂Cr₂O₇ (orange to green), halogens (Cl₂, Br₂, I₂).
      
      
      
      • SO₂ + Cl₂ + 2H₂O → H₂SO₄ + 2HCl
      
      
      
      
    
    
    • Oxidizing Agent: S(+4) → S(0) or S(-2). It oxidizes strong reducing agents like H₂S and Mg.
      
      
      
      • SO₂ + 2H₂S → 3S + 2H₂O (Forms elemental sulphur)
      
      
      • SO₂ + 2Mg → 2MgO + S
      
      
      
      
    
    
    
    
  
  
  
  

Keep practicing these mnemonics to solidify your understanding for exam success!

💡 Quick Tips for Oxoacids of Sulphur; H₂S and SO₂ (Properties)

Mastering the properties and structures related to sulphur compounds is crucial for both JEE Main and CBSE exams. These quick tips focus on high-yield information.

1. Oxoacids of Sulphur

• Oxidation State Calculation: Always determine the oxidation state of sulphur in oxoacids. For example:
  
  
  
  • H₂SO₃ (Sulphurous acid): S is +4
  
  
  • H₂SO₄ (Sulphuric acid): S is +6
  
  
  • H₂S₂O₇ (Pyrosulphuric acid/Oleum): S is +6
  
  
  • H₂SO₅ (Peroxomonosulphuric acid/Caro's acid): S is +6 (contains a peroxide linkage, S-O-O-H)
  
  
  • H₂S₂O₃ (Thiosulphuric acid): Average S is +2 (one S is +6, one is -2 in the S-S linkage)
  
  
  
  


• Structure Matters (JEE Focus): For JEE, understanding the structure of oxoacids is vital. Key features:
  
  
  
  • Sulphur generally exhibits sp³ hybridization.
  
  
  • S=O bonds: Always count the number of S=O and S-OH bonds to determine basicity.
  
  
  • Special Linkages: Identify S-S (e.g., polythionic acids, thiosulphuric acid) or S-O-O-S/S-O-O-H (peroxoacids like H₂SO₅ or H₂S₂O₈) linkages, as they indicate unusual oxidation states or peroxide bonds.
  
  
  • Basicity: The number of -OH groups directly attached to sulphur determines the basicity. E.g., H₂SO₄ is dibasic.
  
  
  
  


• Stability: Higher oxidation states are generally more stable for sulphur, especially H₂SO₄.

2. Hydrogen Sulphide (H₂S)

• Nature: H₂S is a highly poisonous gas with a characteristic rotten egg smell. It is a weak diprotic acid (H₂S ⇌ H⁺ + HS⁻; HS⁻ ⇌ H⁺ + S²⁻).


• Reducing Agent: H₂S is a strong reducing agent. It gets oxidized to sulphur, SO₂, or H₂SO₄ depending on the oxidizing agent and conditions.
  
  
  
  • Example: H₂S + Cl₂ → 2HCl + S (H₂S reduces Cl₂ to HCl, itself oxidized to S)
  
  
  • Used to reduce Fe³⁺ to Fe²⁺ in acidic medium.
  
  
  
  


• Precipitation of Metal Sulphides: Crucial for qualitative analysis (Group II and Group IV metal ions). H₂S gas passed through acidic/basic solutions precipitates metal sulphides.
  
  
  
  • Example: CuSO₄ + H₂S → CuS (black ppt) + H₂SO₄ (in acidic medium, CuS is less soluble)
  
  
  
  


• Formation: Often formed by the action of dilute acids on metal sulphides (e.g., FeS + 2HCl → FeCl₂ + H₂S).

3. Sulphur Dioxide (SO₂)

• Physical Properties: Pungent smelling gas, highly soluble in water forming sulphurous acid (H₂SO₃).


• Acidic Oxide: SO₂ is an acidic oxide. It reacts with water to form sulphurous acid (H₂SO₃) and with bases to form sulphites.
  
  
  
  • SO₂ + H₂O ⇌ H₂SO₃
  
  
  • SO₂ + 2NaOH → Na₂SO₃ + H₂O
  
  
  
  


• Bleaching Action: SO₂ acts as a temporary bleaching agent by reduction. It bleaches delicate articles like silk, wool, straw by reducing their color to a colorless compound. The bleaching effect is temporary as the color can be restored upon exposure to air (oxidation).
  
  
  
  • Mechanism: Colored substance + SO₂ + 2H₂O → Colorless reduced product + H₂SO₄
  
  
  
  


• Redox Properties (Dual Nature):
  
  
  
  • Reducing Agent: SO₂ acts as a strong reducing agent (S goes from +4 to +6).
    
    
    
    • Example: SO₂ + 2H₂O + Cl₂ → H₂SO₄ + 2HCl (SO₂ reduces Cl₂ to HCl)
    
    
    • It reduces Fe³⁺ to Fe²⁺, acidified KMnO₄ (purple) to colorless Mn²⁺, and K₂Cr₂O₇ (orange) to green Cr³⁺. These are important diagnostic tests.
    
    
    
    
  
  
  • Oxidizing Agent: SO₂ also acts as an oxidizing agent (S goes from +4 to 0 or -2) with strong reducing agents like H₂S or Mg.
    
    
    
    • Example: 2H₂S + SO₂ → 2H₂O + 3S (SO₂ oxidizes H₂S to S, itself reduced to S)
    
    
    
    
  
  
  
  


• Structure: SO₂ has a bent shape (V-shape) due to sp² hybridization of sulphur and one lone pair. It exhibits resonance.

Stay focused on these core concepts, and remember to practice reactions and structural aspects for better retention. Good luck!

Intuitive Understanding: Oxoacids of Sulphur; H₂S and SO₂ (Properties)

Understanding the chemistry of sulphur compounds, especially its oxoacids, hydrogen sulphide (H₂S), and sulphur dioxide (SO₂), is crucial for both JEE and board exams. The key to intuitive understanding lies in connecting their properties to the fundamental concepts of oxidation states, electronegativity, and molecular structure.

Oxoacids of Sulphur: The Oxidation State Dictates Character

Sulphur, being a versatile element in Group 16, forms numerous oxoacids where it acts as the central atom bonded to oxygen and hydrogen. The most important ones you'll encounter are sulphuric acid (H₂SO₄) and sulphurous acid (H₂SO₃), along with pyrosulphuric acid (H₂S₂O₇, also known as oleum).

* Conceptual Core: The properties of these oxoacids are largely determined by the oxidation state of sulphur and the number of terminal oxygen atoms.
* In H₂SO₄, sulphur is in its highest stable oxidation state of +6. This means sulphur has largely "given up" its electrons to highly electronegative oxygen atoms. This makes H₂SO₄ a very strong acid (easy to lose H⁺) and a powerful oxidizing agent (sulphur wants to gain electrons to reduce its high oxidation state) and a potent dehydrating agent (it strongly attracts water molecules).
* In H₂SO₃, sulphur is in the +4 oxidation state. This is an intermediate oxidation state, meaning H₂SO₃ can act as both an oxidizing agent (getting reduced to +2, 0, or -2) and, more commonly, as a reducing agent (getting oxidized to +6, as in H₂SO₄). This explains its mild bleaching action.
* Structure Connection: In H₂SO₄, sulphur is sp³ hybridized, forming a tetrahedral geometry. The two -OH groups allow for the release of two protons, making it a diprotic acid.

Hydrogen Sulphide (H₂S): The Rotten Egg Gas

H₂S is a highly toxic, colourless gas with a characteristic "rotten egg" smell.

* Size and Acidity: Compare H₂S with H₂O. Although both are hydrides of Group 16, H₂S is a gas at room temperature, while H₂O is a liquid. This is because sulphur is much larger than oxygen, leading to weaker intermolecular hydrogen bonding in H₂S.
* Acidity: H₂S is significantly more acidic than H₂O. This is because the S-H bond is longer and weaker than the O-H bond due to the larger size of sulphur. A weaker bond means H⁺ is more easily released. H₂S is a weak diprotic acid.
* Reducing Nature: Sulphur in H₂S is in its lowest possible oxidation state, -2. This means it has a strong tendency to lose electrons and get oxidized to a higher oxidation state (like 0 in S, or +4 in SO₂, or +6 in SO₄²⁻). Therefore, H₂S is a strong reducing agent. This property is often tested in qualitative analysis reactions where it precipitates metal sulphides.

Sulphur Dioxide (SO₂): The Dual Nature Gas

SO₂ is a pungent, colourless gas formed during the burning of sulphur-containing fuels. It is a major air pollutant.

* Acidic Oxide: SO₂ is an acidic oxide. When dissolved in water, it forms sulphurous acid (H₂SO₃): SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq). This contributes to acid rain.
* Redox Dualism: Sulphur in SO₂ is in the +4 oxidation state. This intermediate oxidation state gives SO₂ a unique dual character:
* Reducing Agent: It can be oxidized to +6 (e.g., to SO₃ or SO₄²⁻). This is its more prominent role, especially in the presence of strong oxidizing agents (e.g., reacting with halogens, KMnO₄, K₂Cr₂O₇). Its temporary bleaching action is due to its reducing property (it reduces coloured substances by forming addition compounds, which fade on standing in air as SO₂ gets oxidized).
* Oxidizing Agent: It can also be reduced to lower oxidation states like 0 (S) or -2 (H₂S), especially when reacting with strong reducing agents (e.g., H₂S).
* Molecular Geometry: SO₂ has a bent molecular geometry due to the presence of one lone pair on the central sulphur atom, which also undergoes sp² hybridization. This gives it a dipole moment.

By understanding these core concepts – oxidation states, size effects, and molecular structure – you can predict and explain the properties of these crucial sulphur compounds, making complex reactions easier to grasp for both JEE and CBSE exams.

Real World Applications: Oxoacids of Sulphur; H₂S and SO₂ (Properties)

Understanding the properties of oxoacids of sulphur, hydrogen sulphide (H₂S), and sulphur dioxide (SO₂) is crucial not just for exams, but also to appreciate their widespread applications and environmental impact in various industries and daily life.

1. Oxoacids of Sulphur (Primarily Sulphuric Acid, H₂SO₄)

Sulphuric acid, often called the "King of Chemicals," is the most important oxoacid of sulphur and its applications are a cornerstone of modern industry.

• 
  Fertilizer Industry: H₂SO₄ is essential for producing phosphate fertilizers (e.g., superphosphate, triple superphosphate) and ammonium sulphate, both vital for agriculture to enhance soil fertility and crop yield.
  


• 
  Chemical Synthesis: It's widely used as a dehydrating agent (e.g., in the nitration of benzene to nitrobenzene) and as a strong acid in the production of various chemicals, including nitric acid, hydrochloric acid, and many organic compounds.
  


• 
  Petroleum Refining: Sulphuric acid is employed to remove impurities (like mercaptans and other sulfur compounds) from gasoline and other petroleum products, improving their quality and reducing harmful emissions.
  


• 
  Detergents and Dyes: It plays a key role in the manufacture of synthetic detergents and a vast array of dyes and pigments, which are used across textiles, paints, and plastics industries.
  


• 
  Metallurgy: Used for pickling (removing rust and scale) from steel and other metals before plating or galvanizing, ensuring better surface adhesion and finish.
  


• 
  Lead-Acid Batteries: A diluted solution of sulphuric acid serves as the electrolyte in lead-acid batteries, commonly found in automobiles, providing the medium for electrochemical reactions.
  


• 
  Pharmaceuticals: Used in the synthesis of various drugs and medicinal compounds.
  

2. Hydrogen Sulphide (H₂S)

While known for its toxicity and "rotten egg" smell, H₂S has niche applications and is also an important environmental and biological molecule.

• 
  Analytical Chemistry: In qualitative analysis, H₂S is a traditional reagent for precipitating metal sulphides from aqueous solutions. Its selective precipitation based on pH helps in identifying and separating various metal ions (e.g., Pb²⁺, Cu²⁺, Cd²⁺, Zn²⁺, Ni²⁺, Fe²⁺).
  


• 
  Heavy Water Production: The Girdler Sulphide (GS) process, a common method for heavy water production, involves the exchange of hydrogen isotopes between H₂S and water.
  


• 
  Biological Signaling: In very low concentrations, H₂S acts as a gasotransmitter in the body, playing roles in vasodilation, neuroprotection, and inflammation modulation.
  


• 
  Geothermal Energy: H₂S is a naturally occurring component in geothermal steam and gases. Its presence needs careful management due to its corrosive and toxic nature.
  

3. Sulphur Dioxide (SO₂)

Sulphur dioxide, a pungent gas, has several industrial uses, despite being a major air pollutant.

• 
  Production of Sulphuric Acid: SO₂ is the primary intermediate in the Contact Process for the industrial manufacture of H₂SO₄, where it is catalytically oxidized to SO₃.
  


• 
  Food Preservative: It is widely used as a preservative for dried fruits, wine, and fruit juices. SO₂ prevents bacterial growth, inhibits enzymatic browning, and preserves the color and flavor of food products.
  


• 
  Bleaching Agent: Due to its reducing properties, SO₂ is used to bleach delicate materials like wool, silk, straw, and paper pulp, where chlorine would be too harsh.
  


• 
  Refrigerant: Historically, SO₂ was used as a refrigerant, though it has largely been replaced by other compounds due to its toxicity and corrosive nature.
  


• 
  Antichlor: Used to remove excess chlorine from textiles or paper after bleaching, preventing damage to the material.
  


• 
  Solvent: SO₂ can act as a non-aqueous solvent for certain reactions.
  

These applications highlight the diverse roles these sulphur compounds play in supporting various industries and our daily lives, while also underscoring the importance of managing their environmental impact.

Analogies can simplify complex chemical concepts, making them easier to understand and recall. Here are some common analogies related to Oxoacids of Sulfur, H₂S, and SO₂:

• 
  Oxoacids of Sulfur – The "Versatile Chef" Sulfur:
  

  Imagine Sulfur (S) as a versatile chef who can create various dishes (oxoacids) by combining with different numbers of oxygen atoms and hydroxyl groups. The "ingredients" and "cooking conditions" (oxidation state) determine the final "dish."

  
  
  
  
  • Sulfuric Acid (H₂SO₄): This is the chef's most robust and stable creation, like a perfectly cooked, long-lasting meal. Its high oxidation state (+6) and strong structure make it very stable.
  
  
  • Sulfurous Acid (H₂SO₃): This is a less stable, intermediate dish, like a meal that can easily change or spoil. Its lower oxidation state (+4) makes it prone to further oxidation.
  
  
  
  


• 
  H₂S – The "Rotten Egg Alarm":
  

  The characteristic pungent smell of Hydrogen Sulfide (H₂S) is often compared to rotten eggs. This distinct smell acts like a "chemical alarm bell", signaling its presence even in very low concentrations. This analogy emphasizes its immediate and unmistakable sensory detection.

  
  


• 
  H₂S – The "Generous Electron Donor" (Reducing Agent):
  

  Hydrogen Sulfide (H₂S) is a strong reducing agent. Think of it as a "generous giver" of electrons. Sulfur in H₂S is in its lowest possible oxidation state (-2) and is eager to give up electrons to achieve a more stable, higher oxidation state (like 0 in S or +4 in SO₂, or +6 in SO₄²⁻). It's like someone who has excess and readily shares with others.

  
  


• 
  SO₂ – The "Chameleon" Gas:
  

  Sulfur Dioxide (SO₂) is known for its dual nature; it can act as both a reducing agent and an oxidizing agent (though primarily a reducing agent). This is like a "chameleon" that changes its properties depending on the chemical environment.

  
  
  
  
  • As a reducing agent: Sulfur in SO₂ is in the +4 oxidation state. It can donate electrons to go to +6 (e.g., SO₄²⁻). This is its more common role, like a chameleon blending into a bright, colorful background.
  
  
  • As an oxidizing agent: It can also accept electrons to go to a lower oxidation state, like 0 (S) or -2 (S²⁻), but this is less frequent. This is like a chameleon adopting a darker, less vibrant camouflage.
  
  
  
  


• 
  SO₂ Bleaching – The "Temporary Eraser":
  

  The bleaching action of SO₂ is often temporary, unlike the permanent bleaching by chlorine. Imagine SO₂ as a "chemical eraser" that temporarily fades a drawing, rather than tearing up the paper. The color is restored when the bleached substance is exposed to air, as atmospheric oxygen oxidizes the addition product formed with SO₂.

  
  

To effectively understand Oxoacids of Sulfur, H₂S, and SO₂, a strong foundation in the following core chemistry concepts is essential. Mastering these prerequisites will ensure clarity and ease in grasping the properties and reactions of these important compounds.

• 
  Atomic Structure and Periodicity:
  
  
  
  • Electron Configuration: Knowledge of electron configurations, particularly for Sulfur (S), Oxygen (O), and Hydrogen (H). This helps understand valence electrons and bonding capacity.
  
  
  • Electronegativity: Understanding electronegativity differences is crucial for predicting bond polarity and the nature of bonds in H₂S, SO₂, and oxoacids.
  
  
  • Oxidation States: The ability to accurately assign and understand oxidation states for elements in compounds is paramount, especially for sulfur in its various oxoacids and compounds like H₂S (-2) and SO₂ (+4).
  
  
  • Periodic Trends: Basic understanding of Group 16 elements (Chalcogens) and their general properties, particularly how sulfur's properties compare to oxygen and selenium.
  
  
  
  


• 
  Chemical Bonding:
  
  
  
  • Lewis Structures: Proficiency in drawing Lewis dot structures for molecules like H₂S, SO₂, SO₃, and various oxoacids (e.g., H₂SO₄, H₂SO₃). This helps visualize lone pairs and bonding patterns.
  
  
  • VSEPR Theory: Applying VSEPR theory to predict molecular geometries (e.g., bent shape of H₂S and SO₂, tetrahedral around sulfur in H₂SO₄).
  
  
  • Hybridization: Basic understanding of hybridization (e.g., sp³ in H₂S, sp² in SO₂) to explain bond angles and molecular shapes.
  
  
  • Types of Covalent Bonds: Familiarity with single, double, and dative (coordinate) bonds, which are prevalent in sulfur oxoacids.
  
  
  
  


• 
  Redox Reactions:
  
  
  
  • Definition of Oxidation and Reduction: Clear understanding of electron transfer concepts.
  
  
  • Identifying Oxidizing and Reducing Agents: Ability to identify which species is oxidized/reduced and which acts as an oxidizing/reducing agent (e.g., H₂S is a strong reducing agent, SO₂ can be both).
  
  
  • Balancing Redox Reactions: Essential for understanding the chemical equations involving H₂S and SO₂.
  
  
  • Disproportionation: Understanding this type of redox reaction where an element in one oxidation state is simultaneously oxidized and reduced, as sulfur can exhibit this property.
  
  
  
  


• 
  Acid-Base Concepts:
  
  
  
  • Brønsted-Lowry Theory: Understanding the concepts of proton donors (acids) and acceptors (bases).
  
  
  • Acid Strength: Factors affecting the strength of oxoacids (e.g., electronegativity of the central atom, number of oxygen atoms, inductive effect). This is crucial for comparing the acidity of different sulfur oxoacids.
  
  
  • Nature of Oxides: Knowing that non-metal oxides like SO₂ are acidic in nature.
  
  
  
  

Related Topics: Building Foundational Understanding

To thoroughly grasp the chemistry of oxoacids of sulfur, H₂S, and SO₂, it is crucial to understand several related concepts and compounds. These topics provide the necessary context and foundational knowledge for a deeper, exam-oriented understanding.

• 
  P-Block Elements (Group 16 - Chalcogens):
  
  
  
  • Understanding the general characteristics, electronic configuration, oxidation states, and anomalous behavior of oxygen compared to sulfur.
  
  
  • Trends in atomic size, electronegativity, ionization enthalpy, and metallic character within Group 16 help explain sulfur's diverse reactivity.
  
  
  • JEE Focus: Be aware of general trends and exceptions, as these are frequently tested in objective questions.
  
  
  
  


• 
  Chemical Bonding and Molecular Structure:
  
  
  
  • Concepts like VSEPR theory and hybridization are fundamental for predicting the shapes of SO₂, H₂S, and various sulfur oxoacids (e.g., trigonal pyramidal for H₂SO₃, tetrahedral for H₂SO₄).
  
  
  • Understanding polarity and intermolecular forces helps explain physical properties.
  
  
  • JEE Focus: Expect questions on predicting hybridization and molecular geometry for specific sulfur compounds.
  
  
  
  


• 
  Redox Reactions:
  
  
  
  • Sulfur exhibits various oxidation states (-2 in H₂S, +4 in SO₂ and H₂SO₃, +6 in H₂SO₄).
  
  
  • Understanding how to assign oxidation numbers and balance redox reactions is essential, as H₂S is a strong reducing agent and SO₂ can act as both an oxidizing and reducing agent depending on the reaction conditions.
  
  
  • Common Mistake: Incorrectly assigning oxidation states, especially in compounds with peroxide linkages (though less common in simple oxoacids of sulfur).
  
  
  
  


• 
  Acidity and Basicity:
  
  
  
  • Knowledge of acid strength, acid-base reactions, and concepts like Lewis acids/bases (SO₂ can act as a Lewis acid) is crucial for understanding the properties of oxoacids.
  
  
  • Factors affecting the acidity of oxoacids (e.g., number of non-hydroxylated oxygen atoms) are directly applicable.
  
  
  
  


• 
  Industrial Preparation of Sulfuric Acid (Contact Process):
  
  
  
  • This is a very important industrial application involving SO₂ as a key intermediate.
  
  
  • The entire process, including the oxidation of SO₂ to SO₃ and subsequent absorption, is a significant topic directly related to the chemistry of SO₂.
  
  
  • JEE Focus: Details of catalysts, optimum conditions, and reactions involved in the Contact Process are frequently asked.
  
  
  
  


• 
  Environmental Chemistry:
  
  
  
  • The role of SO₂ in air pollution and the formation of acid rain is an important related concept, linking chemistry to real-world environmental issues.
  
  
  
  

Mastering these related topics will not only strengthen your understanding of sulfur chemistry but also enhance your problem-solving skills for both board and JEE examinations.

Navigating the P-block elements, especially sulfur compounds, can be tricky. Many students fall into predictable traps. Be vigilant and avoid these common pitfalls in exams for Oxoacids of Sulfur, H₂S, and SO₂ properties.

• 
  Trap 1: Incorrect Oxidation State Calculation in Oxoacids
  
  
  
  • The Mistake: Assuming all oxygen atoms are -2 and hydrogen atoms are +1, especially in peroxoacids like H₂SO₅ (Caro's acid) or H₂S₂O₈ (Marshall's acid). This leads to an incorrectly high oxidation state for sulfur.
  
  
  • The Fix: Always draw the structure for peroxoacids. Remember that an O-O bond exists, where each oxygen in the peroxo linkage contributes -1, not -2. For H₂SO₅, the oxidation state of S is +6 (not +8). For H₂S₂O₈, each S is +6 (not +7).
  
  
  • JEE Focus: Questions frequently test your understanding of peroxo linkages to calculate the correct oxidation state.
  
  
  
  


• 
  Trap 2: Confusing Basicity of Oxoacids
  
  
  
  • The Mistake: Assuming all hydrogen atoms in an oxoacid are acidic and contribute to its basicity.
  
  
  • The Fix: Only hydrogen atoms attached to oxygen are acidic. Hydrogen atoms directly attached to the central sulfur atom are generally not acidic. For example, H₂SO₃ (sulfurous acid) is diprotic (2 acidic H), and H₂SO₄ (sulfuric acid) is also diprotic (2 acidic H).
  
  
  • CBSE Focus: While less common for sulfur, this is a significant trap for phosphorus oxoacids (e.g., H₃PO₂ is monobasic). Be aware of the principle.
  
  
  
  


• 
  Trap 3: Underestimating the Reducing Power of H₂S
  
  
  
  • The Mistake: Forgetting that H₂S, with sulfur in its lowest oxidation state (-2), is a strong reducing agent.
  
  
  • The Fix: H₂S will readily be oxidized. Products depend on the strength of the oxidizing agent:
    
    
    
    • Weak oxidizers (e.g., SO₂, Fe³⁺): H₂S → S (elemental sulfur)
    
    
    • Strong oxidizers (e.g., HNO₃, KMnO₄): H₂S → SO₂ or H₂SO₄
    
    
    
    Remember its use in qualitative analysis to precipitate metal sulfides.
    
  
  
  
  


• 
  Trap 4: Missing the Dual Nature of SO₂
  
  
  
  • The Mistake: Viewing SO₂ (sulfur dioxide) as only an oxidizing agent or only a reducing agent.
  
  
  • The Fix: Sulfur in SO₂ is in the +4 oxidation state, which is an intermediate state.
    
    
    
    • As a reducing agent: It gets oxidized to +6 (e.g., H₂SO₄ or SO₃). This happens with strong oxidizing agents like KMnO₄ (decolorizes it) or K₂Cr₂O₇ (turns green).
    
    
    • As an oxidizing agent: It gets reduced to 0 (S) or -2 (H₂S). This happens with strong reducing agents like H₂S.
    
    
    
    
  
  
  • JEE & CBSE Focus: Questions often involve reactions where you need to identify its role based on the other reactant.
  
  
  
  


• 
  Trap 5: Confusing SO₂'s Bleaching Action
  
  
  
  • The Mistake: Believing SO₂'s bleaching action is permanent or occurs via oxidation, similar to Cl₂.
  
  
  • The Fix: SO₂ bleaches by reduction (removing oxygen, converting color to colorless). This bleaching action is temporary; the bleached substance can regain its color on exposure to air (oxygen). Chlorine, on the other hand, bleaches by oxidation and is permanent.
  
  
  
  

By understanding these common traps, you can approach questions on sulfur compounds with greater precision and avoid losing easy marks.

Here are the key takeaways for Oxoacids of Sulphur, H₂S, and SO₂ properties, essential for both JEE Main and CBSE Board exams.

Key Takeaways: Oxoacids of Sulphur; H₂S and SO₂

• 
  Oxoacids of Sulphur:
  
  
  
  • Common Oxoacids: Important oxoacids include Sulfurous acid (H₂SO₃), Sulfuric acid (H₂SO₄), Thiosulfuric acid (H₂S₂O₃), Pyrosulfuric acid (H₂S₂O₇, Oleum), Peroxomonosulfuric acid (H₂SO₅, Caro's acid), and Peroxodisulfuric acid (H₂S₂O₈, Marshall's acid).
  
  
  • Oxidation States: Sulphur exhibits various oxidation states. For example, in H₂SO₃, S is +4; in H₂SO₄, S is +6; in H₂S₂O₃, the average is +2 (but one S is +6, other is -2); in H₂SO₅, S is +6 (contains a peroxide linkage).
  
  
  • Structure and Bonding:
    
    
    
    • All oxoacids of sulphur contain at least one S=O bond and one S-OH bond.
    
    
    • S-S bonds are present in compounds like H₂S₂O₃ (thiosulfuric acid).
    
    
    • S-O-S (ether-like) linkage is found in H₂S₂O₇ (pyrosulfuric acid).
    
    
    • Peroxide (O-O) linkages are present in H₂SO₅ and H₂S₂O₈. These are crucial for their strong oxidizing properties.
    
    
    
    
  
  
  • Acidity: Generally, acidity increases with increasing oxidation state of sulphur. H₂SO₄ is a strong diprotic acid.
  
  
  • JEE Focus: Pay close attention to structures, oxidation states, and the presence of S-S, S-O-S, and O-O bonds, as these are frequently tested.
  
  
  
  


• 
  Hydrogen Sulfide (H₂S):
  
  
  
  • Preparation: Typically prepared by reacting dilute acids with metal sulfides (e.g., FeS + 2HCl → FeCl₂ + H₂S).
  
  
  • Physical Properties: Colorless gas with a characteristic rotten egg smell. Highly poisonous.
  
  
  • Chemical Properties:
    
    
    
    • Weak Acidic Nature: H₂S is a weak diprotic acid (H₂S ⇌ H⁺ + HS⁻; HS⁻ ⇌ H⁺ + S²⁻). It reacts with bases to form sulfides and hydrosulfides.
    
    
    • Strong Reducing Agent: Due to sulphur being in its lowest oxidation state (-2), H₂S is a powerful reducing agent. It readily gets oxidized to elemental sulphur, SO₂, or SO₃ depending on the reaction conditions and the oxidizing agent.
      
      
      
      • Example: H₂S reduces Fe³⁺ to Fe²⁺, halogens to halide ions, and KMnO₄ to Mn²⁺.
      
      
      
      
    
    
    • Qualitative Analysis: Used extensively for the precipitation of metal sulfides (e.g., CdS, PbS, CuS) from their salt solutions, often under controlled pH conditions.
    
    
    
    
  
  
  
  


• 
  Sulphur Dioxide (SO₂):
  
  
  
  • Preparation: Formed by burning sulphur in air (S + O₂ → SO₂), roasting sulphide ores (e.g., 2ZnS + 3O₂ → 2ZnO + 2SO₂), or in the lab by reacting sulphites with dilute acids (e.g., Na₂SO₃ + H₂SO₄ → Na₂SO₄ + H₂O + SO₂).
  
  
  • Physical Properties: Colorless gas with a pungent, suffocating smell. Highly soluble in water.
  
  
  • Chemical Properties:
    
    
    
    • Acidic Oxide: SO₂ dissolves in water to form sulfurous acid (H₂SO₃), a weak acid. It reacts with alkalis to form sulfites (Na₂SO₃) and bisulfites (NaHSO₃).
    
    
    • Reducing Agent: Sulphur in SO₂ is in the +4 oxidation state, so it can be oxidized to +6. It is a good reducing agent, especially in the presence of water.
      
      
      
      • Example: Decolorizes acidic KMnO₄ solution (purple to colorless), reduces dichromate ions (orange to green), and reduces Fe³⁺ to Fe²⁺.
      
      
      
      
    
    
    • Oxidizing Agent: Less common, but SO₂ can act as an oxidizing agent, getting reduced to elemental sulphur or sulphide.
      
      
      
      • Example: 2H₂S + SO₂ → 3S + 2H₂O.
      
      
      
      
    
    
    • Bleaching Action: SO₂ acts as a bleaching agent in the presence of moisture, by reduction. This bleaching action is temporary, as the bleached color is restored on exposure to air. It is particularly used for delicate materials like silk and wool.
    
    
    • Industrial Importance: A key intermediate in the manufacture of sulphuric acid (Contact Process).
    
    
    
    
  
  
  • CBSE & JEE: Differentiating the redox behavior (oxidizing vs. reducing) of H₂S and SO₂ is a common question. H₂S is *always* a reducing agent, while SO₂ can be *both* reducing and oxidizing, but predominantly reducing.
  
  
  
  

Keep these points in mind for quick revision. Understanding the oxidation states is crucial for predicting the chemical behavior of these compounds.

Problem Solving Approach for Oxoacids of Sulfur, H₂S, and SO₂

To effectively solve problems related to oxoacids of sulfur, H₂S, and SO₂, a systematic approach focusing on oxidation states, structures, and characteristic reactions is essential.

1. Understanding the Core Concepts

Before attempting any problem, ensure a clear understanding of:

• Oxidation State of Sulfur: This is paramount. Sulfur exhibits a wide range of oxidation states (-2 to +6). Knowing the oxidation state in a given compound helps predict its redox behavior. For oxoacids, the sum of oxidation states must be zero.


• Molecular Structures: For oxoacids, H₂S, and SO₂, knowledge of their geometries (VSEPR theory), hybridization, and specific linkages (e.g., S-O-H, S=O, S-O-S, S-S, S-O-O-S) is critical.


• Acidic/Basic Nature: Identify whether the compound acts as an acid, base, or an amphoteric species.


• Redox Properties: Determine if the compound is primarily an oxidizing agent, a reducing agent, or can exhibit both.

2. Step-by-Step Problem Solving Strategy

Follow these steps when tackling problems:

1. Analyze the Question Carefully:
   
   
   
   • Identify what is being asked: Structure, properties (acidity, redox), reaction products, comparison between compounds, or a specific application.
   
   
   • Note all given compounds and conditions (e.g., presence of oxidizing/reducing agents, pH).
   
   
   
   


2. Determine Oxidation States of Sulfur:
   
   
   
   • For any sulfur compound, calculate the oxidation state of sulfur. This is the most crucial first step for predicting redox behavior and often, stability.
   
   
   • Example: In H₂SO₄, S is +6. In H₂SO₃, S is +4. In H₂S, S is -2. In SO₂, S is +4.
   
   
   
   


3. Recall Structural Features (for structural questions):
   
   
   
   • Based on the oxidation state and formula, draw the most plausible Lewis structure.
   
   
   • Identify the hybridization of the sulfur atom (e.g., sp³ for H₂SO₄, sp² for SO₂).
   
   
   • Look for specific bonds/linkages: S-S bonds (e.g., H₂S₂O₃), S-O-S linkages (e.g., H₂S₂O₇), or peroxo linkages (S-O-O-S).
   
   
   • Count sigma and pi bonds, lone pairs, and determine geometry.
   
   
   
   


4. Predict Chemical Properties (for property/reaction questions):
   
   
   
   • Redox Behavior:
     
     
     
     • If S is in its lowest oxidation state (-2 in H₂S), it will be a strong reducing agent (gets oxidized).
     
     
     • If S is in its highest oxidation state (+6 in H₂SO₄), it will be an oxidizing agent (gets reduced).
     
     
     • If S is in an intermediate oxidation state (+4 in SO₂, H₂SO₃), it can act as both an oxidizing and a reducing agent (disproportionation is possible).
     
     
     
     
   
   
   • Acidic/Basic Nature:
     
     
     
     • SO₂ is an acidic oxide, forming H₂SO₃ in water and reacting with bases.
     
     
     • H₂S is a weak diprotic acid.
     
     
     • Oxoacids of sulfur are acids, with acidity generally increasing with the oxidation state of sulfur (H₂SO₄ > H₂SO₃) due to increased electronegativity of S and better stabilization of the conjugate base.
     
     
     
     
   
   
   
   


5. Identify Products of Reactions:
   
   
   
   • Based on the redox and acid-base properties, predict the products. For instance, a strong reducing agent like H₂S will reduce oxidizing agents. SO₂ (a reducing agent) will typically be oxidized to SO₄²⁻ or H₂SO₄.
   
   
   • Consider common reactions: e.g., H₂S reacting with metal ions to form insoluble sulfides (qualitative analysis). SO₂'s bleaching action (due to reduction).
   
   
   
   

3. JEE Specific Pointers

• Complex Oxoacids: Be prepared for questions on less common oxoacids like pyrosulfuric acid (H₂S₂O₇), thiosulfuric acid (H₂S₂O₃), and various peroxoacids of sulfur (e.g., H₂SO₅, H₂S₂O₈). Focus on their structures and specific linkages.


• Balancing Redox Reactions: Questions often involve balancing redox reactions where H₂S or SO₂ are reactants or products, particularly with strong oxidizing agents like KMnO₄ or K₂Cr₂O₇.


• Qualitative Analysis: The use of H₂S in qualitative analysis for precipitating metal sulfides in specific groups is a common application.


• Comparison Questions: Be ready to compare properties like acidity, reducing/oxidizing strength, and thermal stability among different oxoacids or sulfur compounds.

By systematically applying these steps, you can confidently approach and solve a wide range of problems related to oxoacids of sulfur, H₂S, and SO₂.

The P-block elements are crucial for CBSE board examinations, with oxoacids of sulfur, H₂S, and SO₂ being frequently tested topics. Focus areas typically include structural aspects, oxidation states, and key chemical properties, especially their redox behavior.

Oxoacids of Sulphur (CBSE Focus)

For CBSE, the primary emphasis on oxoacids of sulphur revolves around their structures, oxidation states, and the presence of specific linkages. You should be able to draw the structures and identify key features.

• Important Oxoacids:
  
  
  
  • Sulphurous acid (H₂SO₃): Planar, S in +4 oxidation state.
  
  
  • Sulphuric acid (H₂SO₄): Tetrahedral, S in +6 oxidation state. Two -OH groups.
  
  
  • Thiosulphuric acid (H₂S₂O₃): Contains S=S bond, S has average +2 oxidation state.
  
  
  • Pyrosulphuric acid (H₂S₂O₇ - Oleum): Contains S-O-S (ether-like) linkage. S in +6 oxidation state.
  
  
  • Peroxomonosulphuric acid (H₂SO₅ - Caro's acid): Contains one S-O-O-H (peroxo) linkage. S in +6 oxidation state.
  
  
  • Peroxodisulphuric acid (H₂S₂O₈ - Marshall's acid): Contains one S-O-O-S (peroxo) linkage. S in +6 oxidation state.
  
  
  
  


• Key Structural Aspects:
  
  
  
  • Drawing structures: Practice drawing the structures for all listed oxoacids, clearly showing S-O-H, S=O, S-S, S-O-S, and S-O-O-S linkages.
  
  
  • Oxidation States: Be able to calculate the oxidation state of sulfur in each oxoacid.
  
  
  • Bonding: Note that sulfur exhibits dπ-pπ bonding in its oxoacids, leading to higher coordination.
  
  
  
  

Hydrogen Sulphide (H₂S) Properties (CBSE Focus)

H₂S is a crucial compound, particularly for its reducing properties and analytical applications.

• Preparation:
  
  
  
  • Lab Method: By the action of dilute non-oxidizing acids on metal sulfides.
    
    FeS(s) + H₂SO₄(dil.) → FeSO₄(aq) + H₂S(g)
  
  
  
  


• Physical Properties:
  
  
  
  • Colourless gas with a characteristic rotten egg smell.
  
  
  • Highly poisonous, should be handled with care in labs.
  
  
  • Slightly soluble in water.
  
  
  
  


• Chemical Properties:
  
  
  
  • Acidic Nature: A weak dibasic acid in aqueous solution. It dissociates in two steps to give H⁺ ions.
    
    H₂S ⇌ H⁺ + HS⁻; HS⁻ ⇌ H⁺ + S²⁻
  
  
  • Reducing Agent: This is its most important chemical property. Sulfur in H₂S is in its lowest oxidation state (-2) and thus readily gets oxidized.
    
    
    
    • Reduces halogens (Cl₂, Br₂, I₂) to respective hydrohalic acids, itself oxidizing to sulfur.
      
      H₂S + Br₂ → 2HBr + S
    
    
    • Reduces Fe³⁺ to Fe²⁺.
    
    
    • Reduces oxidizing agents like KMnO₄ (purple to colorless) and K₂Cr₂O₇ (orange to green).
    
    
    
    
  
  
  • Precipitation of Metal Sulphides: Used in qualitative analysis to precipitate metal sulfides (e.g., PbS, CdS, CuS) from acidic solutions.
  
  
  • Combustion: Burns in air/oxygen to form SO₂ and H₂O.
  
  
  
  

Sulphur Dioxide (SO₂) Properties (CBSE Focus)

SO₂ is an important industrial chemical and an air pollutant. Its properties, especially redox behavior, are often tested.

• Preparation:
  
  
  
  • From Sulphur: S(s) + O₂(g) → SO₂(g)
  
  
  • From Sulphide Ores (Roasting): 2FeS₂(s) + 11O₂(g) → 2Fe₂O₃(s) + 8SO₂(g)
  
  
  • From Sulphites: Na₂SO₃(s) + H₂SO₄(dil.) → Na₂SO₄(aq) + H₂O(l) + SO₂(g)
  
  
  
  


• Physical Properties:
  
  
  
  • Colourless gas with a pungent smell.
  
  
  • Highly soluble in water, forming sulphurous acid (H₂SO₃).
  
  
  • Easily liquefied at room temperature under pressure.
  
  
  
  


• Chemical Properties:
  
  
  
  • Acidic Nature: An acidic oxide. Reacts with water to form sulphurous acid and with alkalis to form sulphites.
    
    SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq)
    
    SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l)
  
  
  • Reducing Agent: Sulphur in SO₂ (+4 oxidation state) can be oxidized to +6.
    
    
    
    • Decolorizes acidified KMnO₄ solution (purple to colorless).
      
      5SO₂ + 2KMnO₄ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄
    
    
    • Reduces Fe³⁺ to Fe²⁺.
    
    
    • Bleaching action: Temporary bleaching action on moist delicate fabrics (due to reduction to colorless compounds). The color is restored on exposure to air.
    
    
    
    
  
  
  • Oxidizing Agent: Sulphur in SO₂ (+4 oxidation state) can also be reduced to 0 or -2. This property is less common but important.
    
    
    
    • With strong reducing agents like H₂S:
      
      SO₂(g) + 2H₂S(g) → 2H₂O(l) + 3S(s)
    
    
    
    
  
  
  • Addition Reactions: Forms addition products, e.g., with chlorine in the presence of charcoal (SO₂Cl₂).
  
  
  
  

Mastering these properties and structures will ensure a strong performance in the CBSE board exams for this topic. Pay special attention to the redox behavior of H₂S and SO₂.

JEE Focus Areas: Oxoacids of Sulfur, H₂S, and SO₂

This section outlines the most crucial concepts and frequently tested aspects related to oxoacids of sulfur, H₂S, and SO₂ for JEE Main.

• 
  Oxoacids of Sulfur: Structural Insights and Oxidation States
  
  
  
  • Key Structures: Focus on drawing and understanding the structures of common oxoacids like Sulfuric Acid (H₂SO₄), Sulfurous Acid (H₂SO₃), Pyrosulfuric Acid (H₂S₂O₇, Oleum), Peroxodisulfuric Acid (H₂S₂O₈, Marshall's Acid), and Thiosulfuric Acid (H₂S₂O₃).
  
  
  • Critical Aspects:
    
    
    
    • Oxidation State: Be able to calculate the oxidation state of sulfur in each oxoacid.
    
    
    • Linkages: Identify the presence of S-S (e.g., H₂S₂O₃) or S-O-O-S (peroxide linkage in H₂S₂O₈) bonds.
    
    
    • Hybridization & Geometry: Sulfur is typically sp³ hybridized in these oxoacids, leading to tetrahedral geometry around the sulfur atom(s).
    
    
    
    
  
  
  • JEE Tip: Questions often involve identifying the number of S=O and S-OH bonds, or the presence of specific linkages. Understanding the structure is key to predicting properties.
  
  
  
  


• 
  Hydrogen Sulfide (H₂S): Acidic and Reducing Properties
  
  
  
  • Preparation: Know common lab preparations, e.g., action of dilute acids on metal sulfides (FeS + H₂SO₄(dil) → FeSO₄ + H₂S).
  
  
  • Acidic Nature: H₂S is a weak diprotic acid. Its dissociation constants (Ka₁ and Ka₂) are important, especially Ka₁ >> Ka₂.
    
    
    
    • H₂S ⇌ H⁺ + HS⁻ (Ka₁)
    
    
    • HS⁻ ⇌ H⁺ + S²⁻ (Ka₂)
    
    
    
    
  
  
  • Reducing Nature: H₂S is a strong reducing agent. It is readily oxidized to elemental sulfur (S), sulfur dioxide (SO₂), or sulfate (SO₄²⁻) depending on the oxidizing agent.
    
    
    
    • Example: H₂S + Cl₂ → 2HCl + S (H₂S reduces chlorine)
    
    
    
    
  
  
  • JEE Significance (Qualitative Analysis): H₂S is crucial in qualitative analysis for precipitating metal sulfides in specific groups (e.g., Group II in acidic medium, Group IV in basic medium). Be familiar with the characteristic colors of common metal sulfides (e.g., CdS yellow, CuS black, ZnS white).
  
  
  
  


• 
  Sulfur Dioxide (SO₂): Versatile Redox and Bleaching Actions
  
  
  
  • Structure: SO₂ has a bent molecular geometry with sp² hybridization on sulfur, involving resonance structures.
  
  
  • Acidic Nature: It is an acidic oxide, dissolving in water to form sulfurous acid (H₂SO₃), a weak acid.
    
    
    
    • SO₂ + H₂O ⇌ H₂SO₃
    
    
    
    
  
  
  • Reducing Properties: SO₂ acts as a strong reducing agent, especially in aqueous solutions, getting oxidized to SO₃ (or H₂SO₄).
    
    
    
    • Example: 5SO₂ + 2KMnO₄ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄ (decolorizes pink KMnO₄ solution).
    
    
    
    
  
  
  • Oxidizing Properties: It can also act as an oxidizing agent with stronger reducing agents (e.g., H₂S).
    
    
    
    • Example: SO₂ + 2H₂S → 3S + 2H₂O
    
    
    
    
  
  
  • Bleaching Action: SO₂ bleaches by reduction, which is temporary. The color can be restored upon exposure to air (re-oxidation). This distinguishes it from chlorine's permanent bleaching action (by oxidation).
    
    
    
    • Colored matter + SO₂ + H₂O → Colorless reduced product
    
    
    
    
  
  
  
  

Pro Tip: Master the unique properties and characteristic reactions of each compound. Pay special attention to redox reactions and the conditions under which H₂S and SO₂ act as oxidizing or reducing agents. Good luck!