Which one of the following reducing agents can also act as an oxidizing agent ?

1. Redox Reactions: Basics of Oxidation and Reduction (basic)

Welcome to your first step in mastering chemical principles! To understand how batteries work, how metals are extracted from the earth, or even how our bodies produce energy, we must first master Redox Reactions. The term 'Redox' is a portmanteau of Reduction and Oxidation. These two processes are like two sides of a coin; they always occur simultaneously because if one substance loses something, another must be there to gain it.

At its most basic level, we define these processes by the movement of oxygen and hydrogen. If a substance gains oxygen or loses hydrogen during a reaction, it is said to be oxidized. Conversely, if a substance loses oxygen or gains hydrogen, it is reduced Science, Chemical Reactions and Equations, p.12-13. For example, when Copper Oxide (CuO) reacts with Hydrogen (H₂), the CuO loses oxygen to become metallic copper (reduction), while the H₂ gains that oxygen to become water (oxidation). In the industrial world, we use this principle to obtain pure metals from their ores—a process that is essentially a large-scale reduction Science, Metals and Non-metals, p.51.

To identify the players in this 'chemical exchange,' we use the terms Oxidizing Agent and Reducing Agent. It can be a bit counter-intuitive at first:

• The Reducing Agent is the substance that gets oxidized (it 'donates' the ability to reduce the other reactant).
• The Oxidizing Agent is the substance that gets reduced (it 'takes' the electrons or oxygen, causing the other to be oxidized).

While some substances like Hydrogen (H₂) or Carbon (C) are classic reducing agents, others are versatile. Sulfur dioxide (SO₂) is a fascinating example because it can act as both an oxidizing and a reducing agent depending on what it reacts with. This is because the sulfur in SO₂ is in an intermediate state; it can either lose more electrons to reach a higher state (acting as a reducer) or gain electrons to drop to a lower state (acting as an oxidizer).

Process	Oxygen Transfer	Hydrogen Transfer	Electron Transfer (Advanced)
Oxidation	Gain of Oxygen	Loss of Hydrogen	Loss of Electrons
Reduction	Loss of Oxygen	Gain of Hydrogen	Gain of Electrons

Sources: Science, Chemical Reactions and Equations, p.12; Science, Chemical Reactions and Equations, p.13; Science, Metals and Non-metals, p.51

2. Understanding Oxidation States and Numbers (basic)

Think of oxidation states (or oxidation numbers) as a chemical bookkeeping system. While we often define oxidation simply as the gain of oxygen or loss of hydrogen Science, Class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12, a more universal way to understand it is through the movement of electrons. An oxidation state represents the "apparent charge" an atom carries. By tracking these numbers, we can predict how a substance will behave in a redox reaction—specifically, whether it will act as an oxidizing agent (taking electrons) or a reducing agent (giving them away).

Most elements have a specific range of possible oxidation states, effectively a "floor" (lowest) and a "ceiling" (highest). If an element is already at its floor, it cannot gain any more electrons; it can only lose them (be oxidized). Therefore, it can only act as a reducing agent. Conversely, if it is at its ceiling, it can only act as an oxidizing agent. However, some substances contain an element in an intermediate oxidation state. These "middle-ground" substances are chemical chameleons—they can either climb higher or drop lower depending on who they are reacting with.

A classic example of this dual behavior is Sulfur Dioxide (SO₂). In SO₂, Sulfur has an oxidation state of +4. Because Sulfur's possible states range from -2 up to +6, the +4 state is intermediate. This allows SO₂ to be oxidized to +6 (acting as a reducing agent) or reduced to 0 or -2 (acting as an oxidizing agent) Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.56. This contrasts with H₂S, where Sulfur is at its absolute minimum of -2 and can only serve as a reducing agent.

Substance	Oxidation State of Sulfur	Redox Role
Hydrogen Sulfide (H₂S)	-2 (Minimum)	Reducing Agent only
Sulfur Dioxide (SO₂)	+4 (Intermediate)	Both Oxidizing & Reducing Agent
Sulfuric Acid (H₂SO₄)	+6 (Maximum)	Oxidizing Agent only

Sources: Science, Class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.56

3. Identifying Oxidizing and Reducing Agents (intermediate)

In any chemical reaction, oxidation and reduction are two sides of the same coin. To master this, you must distinguish between the process and the 'agent.' An oxidizing agent is a substance that facilitates oxidation by providing oxygen or removing hydrogen from another substance Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.71. Crucially, because it takes away electrons or adds oxygen, the oxidizing agent itself gets reduced. Conversely, a reducing agent is the one that loses oxygen or adds hydrogen to another substance, thereby getting oxidized in the process Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.51.

Identifying these agents becomes easier when you look at the oxidation state of the central atom. Think of it like a ladder: if an element is already at its highest possible 'rung' (oxidation state), it can only go down, meaning it can only be reduced and thus acts only as an oxidizing agent (e.g., Potassium Permanganate). If it is at the bottom floor, like Sulfur in H₂S (-2 state), it can only go up, acting only as a reducing agent. However, substances like Sulfur Dioxide (SO₂) are unique because they sit on a middle rung (+4 state). Depending on the partner they react with, they can climb higher to +6 (acting as a reducing agent) or drop down to 0 or -2 (acting as an oxidizing agent).

Role	Action on Others	Self-Transformation
Oxidizing Agent	Gains electrons / Adds Oxygen	Gets Reduced
Reducing Agent	Loses electrons / Removes Oxygen	Gets Oxidized

In industrial applications, we use these principles to extract pure metals. For instance, carbon is frequently used as a reducing agent to strip oxygen away from metal oxides like Zinc Oxide (ZnO), leaving behind pure metallic Zinc Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.51. Understanding this dual behavior is essential for predicting how pollutants like SO₂ will behave in the atmosphere compared to stable compounds.

Sources: Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.71; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.51; Science, Class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12

4. Environmental Chemistry: Oxides of Sulfur and Nitrogen (intermediate)

In the study of air pollution, Oxides of Sulfur (SOₓ) and Nitrogen (NOₓ) are two of the most critical gaseous pollutants. Most of these gases enter our atmosphere through the combustion of fossil fuels in motor vehicles and industrial processes Majid Hussain, Environment and Ecology, Environmental Degradation and Management, p.8. From a chemical perspective, Sulfur Dioxide (SO₂) is particularly fascinating because of its "dual personality." In SO₂, sulfur exists in an intermediate oxidation state (+4). Because sulfur can range from an oxidation state of -2 (as in H₂S) to +6 (as in sulfates), SO₂ can either lose electrons to become +6 (acting as a reducing agent) or gain electrons to move toward a lower state (acting as an oxidizing agent). This versatility makes it a central player in various atmospheric chemical reactions.

When these oxides are released, they don't just stay in their original form; they undergo a transformation known as secondary pollutant formation. In the presence of sunlight, which stimulates the creation of photo-oxidants like ozone, SO₂ and NO₂ react with moisture and oxygen to form Sulfuric Acid (H₂SO₄) and Nitric Acid (HNO₃) Shankar IAS Academy, Environment, Environmental Pollution, p.103. These acids then fall to the earth as Acid Rain, which can be deposited either as "wet deposition" (rain, snow, fog) or "dry deposition" (particles and gases sticking to surfaces).

The human health impact of these gases is profound. High concentrations of SOₓ and NOₓ are directly linked to respiratory ailments, including severe asthma attacks and permanent lung damage. In major urban centers like Delhi, these pollutants are significant contributors to premature mortality rates Majid Hussain, Environment and Ecology, Environmental Degradation and Management, p.39. To manage this, the National Air Quality Monitoring Programme (NAMP) tracks these pollutants across India to determine trends and ensure compliance with air quality standards Shankar IAS Academy, Environment, Environmental Pollution, p.69.

Pollutant	Primary Source	Atmospheric End-Product
Sulfur Dioxide (SO₂)	Coal-burning power plants, refineries	Sulfuric Acid (H₂SO₄)
Nitrogen Oxides (NOₓ)	Vehicle exhaust, high-temp combustion	Nitric Acid (HNO₃)

Sources: Environment and Ecology, Majid Hussain, Environmental Degradation and Management, p.8, 39; Environment, Shankar IAS Academy, Environmental Pollution, p.69, 103

5. Periodic Table Trends: Electronegativity and Reactivity (intermediate)

In our journey through the Periodic Table, we move from simple atomic structure to understanding the chemical personality of elements. This personality is largely defined by two concepts: Electronegativity and Reactivity. Think of Electronegativity as an atom's "hunger" for electrons. It is the measure of how strongly an atom can pull a shared pair of electrons toward itself in a chemical bond. As we move from left to right across a period, the nuclear charge increases, pulling electrons more tightly. Consequently, Electronegativity increases across a period and decreases as we go down a group Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.40.

Reactivity, on the other hand, is the actual speed or ease with which an element undergoes a chemical change. However, reactivity follows different "rules" depending on whether we are looking at metals or non-metals:

• Metals: These elements react by losing electrons. The further an electron is from the nucleus, the easier it is to lose. Therefore, metallic reactivity increases as you move down a group. This is why Potassium (K) and Sodium (Na) are so reactive they must be stored in kerosene to prevent them from catching fire in the air Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.42.
• Non-metals: These react by gaining electrons. Because smaller atoms with higher electronegativity (like Fluorine) pull electrons in more effectively, non-metallic reactivity decreases as you move down a group.

Trend	Across a Period (Left to Right)	Down a Group (Top to Bottom)
Electronegativity	Increases	Decreases
Metallic Reactivity	Decreases	Increases
Non-metallic Reactivity	Increases	Decreases

To help scientists and students predict how these elements will behave in the real world, we use the Activity Series. This is a list where metals are ranked in order of their decreasing reactivity. At the top, you find the "hyper-reactive" metals like Potassium, while the "noble" metals like Gold (Au) sit at the very bottom, rarely reacting with anything Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.45.

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.40; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.42; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.45

6. Dual Redox Behavior and Intermediate Oxidation States (exam-level)

In the study of chemistry, we often classify substances strictly as either oxidizing agents (which add oxygen or remove hydrogen/electrons) or reducing agents (which remove oxygen or add hydrogen/electrons) Science, Chemical Reactions and Equations, p.12-13. However, some unique substances exhibit Dual Redox Behavior. To understand why, we must look at the oxidation state—essentially the 'electronic status'—of the central atom. If an element is in its intermediate oxidation state, it possesses the flexibility to either lose electrons (be oxidized) or gain electrons (be reduced).

Sulfur Dioxide (SO₂) is the classic example of this chemical versatility. In SO₂, sulfur is in a +4 oxidation state. Since sulfur's range of common oxidation states spans from -2 to +6, the +4 state is right in the middle.

• As a Reducing Agent: SO₂ can be oxidized to the +6 state (like in sulfate ions or SO₃). By being oxidized itself, it reduces the other reactant.
• As an Oxidizing Agent: SO₂ can be reduced down to the 0 state (elemental sulfur) or even to -2 (H₂S). By being reduced, it oxidizes the other reactant.

This is why SO₂ is such a vital reagent in industrial processes; it changes its behavior based on the 'strength' of the substance it encounters.

In contrast, substances like Hydrogen Sulfide (H₂S) and Hydriodic Acid (HI) typically behave only as reducing agents. This is because sulfur in H₂S is at its lowest state (-2) and iodine in HI is at its lowest state (-1). Since they are already at the 'bottom' of their electronic ladder, they can only be oxidized (move up), meaning they can only act as reducing agents. While some elements like Hydrogen (H₂) can occasionally show dual behavior (acting as an oxidizing agent when reacting with active metals), SO₂ remains the primary example of a substance that switches roles based on its intermediate chemical environment Science, Carbon and its Compounds, p.71.

Substance	Oxidation State of Central Atom	Redox Role
SO₂	+4 (Intermediate)	Both Oxidizing & Reducing Agent
H₂S	-2 (Minimum)	Reducing Agent only
KMnO₄	+7 (Maximum)	Oxidizing Agent only

Sources: Science, class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12-13; Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.71; Science-Class VII (NCERT 2025 ed.), The World of Metals and Non-metals, p.53

7. Solving the Original PYQ (exam-level)

Now that you have mastered the fundamentals of oxidation states and redox reactions, this question tests your ability to apply those building blocks to real-world chemical behavior. The core principle at play here is the range of oxidation states a central atom can occupy. For a molecule to exhibit dual behavior—acting as both an oxidizing and a reducing agent—the central atom must sit in an intermediate oxidation state, allowing it to either lose or gain electrons depending on its reaction partner.

Let’s walk through the reasoning for the correct answer, (C) SO₂. In sulfur dioxide, the sulfur atom is in a +4 oxidation state. Because sulfur’s common oxidation states range from -2 to +6, it is perfectly positioned in the middle. It can be oxidized to +6 (acting as a reducing agent) or reduced to 0 or -2 (acting as an oxidizing agent). This versatility is a classic concept highlighted in Science, class X (NCERT 2025 ed.), where substances in intermediate states are identified by their ability to function in both capacities.

UPSC often uses options like H₂S and HI as traps because they are chemically active, but they are at their lowest possible oxidation states (-2 for Sulfur and -1 for Iodine). Since they cannot gain more electrons, they can only be oxidized, making them strictly reducing agents. While H₂ can act as an oxidizing agent when forming metal hydrides, SO₂ is the superior answer because its dual redox nature is its defining characteristic in the standard chemical contexts typically tested in the Civil Services Examination.