Which one of the following is the correct order of oxidation number of Iodine (I) in I2, HI, HIO4 and ICI?

1. Basics of Redox Reactions (basic)

In the world of chemistry, reactions rarely happen in isolation. The term Redox is actually a portmanteau of two simultaneous processes: Reduction and Oxidation. At its simplest level, oxidation is defined as the gain of oxygen or the loss of hydrogen by a substance. Conversely, reduction is the loss of oxygen or the gain of hydrogen Science, Class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12. Because oxygen is highly reactive, these reactions are everywhere—from the respiration that powers your body to the rusting of iron in red soil Physical Geography by PMF IAS, Geomorphic Movements, p.91.

To master this concept for competitive exams, we must look beyond just oxygen. Chemically, redox reactions are about the transfer of electrons. When an atom loses an electron, it is 'oxidized'; when it gains an electron, it is 'reduced.' This transfer changes the 'Oxidation Number' of the elements involved—a bookkeeping system scientists use to track where electrons are going. For example, when a reactive metal displaces a less reactive one from a solution, electrons are being swapped, changing the oxidation states of the metals involved Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.45.

The table below summarizes these two sides of the same coin:

Feature	Oxidation	Reduction
Oxygen	Gain of Oxygen	Loss of Oxygen
Hydrogen	Loss of Hydrogen	Gain of Hydrogen
Electrons	Loss of Electrons	Gain of Electrons
Oxidation Number	Increases	Decreases

Sources: Science, Class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12; Physical Geography by PMF IAS, Geomorphic Movements, p.91; Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.45

2. Electronic Concept of Oxidation and Reduction (basic)

In our previous look at chemical reactions, we often defined oxidation as the gain of oxygen and reduction as the loss of oxygen Science, Class X, Chemical Reactions and Equations, p.12. However, to truly master chemistry for the UPSC, we must look deeper into the Electronic Concept. At its core, every chemical reaction is a redistribution of electrons. Atoms react because they have a drive to achieve a stable, filled outermost shell—similar to the stable electronic configuration of noble gases Science, Class X, Metals and Non-metals, p.46.

Under the electronic concept, Oxidation is defined as the process involving the loss of electrons by an atom or ion. When an atom loses a negatively charged electron, its overall positive character increases. Conversely, Reduction is the process involving the gain of electrons. For example, when Magnesium reacts with Hydrochloric acid Science, Class X, Chemical Reactions and Equations, p.15, the Magnesium atom (Mg) loses two electrons to become a Magnesium ion (Mg²⁺). This loss makes the Magnesium "oxidized."

To keep track of these migrating electrons, scientists use a bookkeeping tool called the Oxidation Number (or State). This represents the total number of electrons that an atom either gains or loses in order to form a chemical bond with another atom. In a neutral element like pure Iodine (I₂), the oxidation number is 0 because no electrons have been shifted. However, in a compound, the more "greedy" (electronegative) atom pulls electrons toward itself, taking a negative oxidation state, while the atom that loses its grip on electrons takes a positive state.

Process	Electron Movement	Oxidation Number Change
Oxidation	Loss of Electrons	Increases (becomes more positive)
Reduction	Gain of Electrons	Decreases (becomes more negative/less positive)

Sources: Science, Class X, Chemical Reactions and Equations, p.12, 15; Science, Class X, Metals and Non-metals, p.46

3. Introduction to Oxidation Numbers (basic)

Oxidation Number (also called oxidation state) is a fundamental concept used by chemists to keep track of electrons during chemical reactions. Think of it as a "bookkeeping" system: it represents the hypothetical charge an atom would carry if all its bonds were purely ionic. While in reality electrons are often shared, assigning these numbers helps us understand which substance is gaining or losing electrons in a reaction, such as when metals like copper or aluminium react with oxygen to form oxides Science, Class X, Metals and Non-metals, p.41.

To determine these numbers, we follow a set of logical rules based on the nature of the elements involved:

• Elemental Rule: Any element in its free or uncombined state (like O₂, I₂, or solid Fe) has an oxidation number of 0 Science, Class VIII, Nature of Matter, p.123.
• Hydrogen & Oxygen Rule: In most compounds, Hydrogen is assigned +1 and Oxygen is assigned -2 (as seen in metal oxides like CuO or Al₂O₃) Science, Class X, Metals and Non-metals, p.41.
• Electronegativity Rule: When two different non-metals (heteroatoms) bond, the more "electron-greedy" (electronegative) atom is assigned the negative number Science, Class X, Carbon and its Compounds, p.66. For example, in a bond between Iodine and Chlorine (ICl), Chlorine is more electronegative and takes the -1 state, leaving Iodine at +1.
• Summation Rule: For a neutral molecule, the sum of all oxidation numbers must equal zero.

Let's look at a practical calculation using Periodic Acid (HIO₄). To find the state of Iodine (I), we use the knowns: Hydrogen is +1 and each of the four Oxygen atoms is -2. Calculating the total: (+1) + I + 4(-2) = 0. This simplifies to: 1 + I - 8 = 0, which means I = +7. By applying these simple steps, we can determine the oxidation state of any element in any compound!

Substance Type	Oxidation Number Rule	Example
Pure Element	Always 0	I₂ = 0
Hydrogen (with non-metals)	Always +1	HI (H is +1)
Neutral Compound	Sum must be 0	H₂O (2(+1) + (-2) = 0)

Sources: Science, Class X, Metals and Non-metals, p.41; Science, Class VIII, Nature of Matter, p.123; Science, Class X, Carbon and its Compounds, p.66

4. Electronegativity and Periodic Trends (intermediate)

Electronegativity is a fundamental concept in chemistry that describes the 'tug-of-war' for electrons between atoms. When two atoms form a covalent bond, they share a pair of electrons to achieve a stable electronic configuration, often resembling the nearest noble gas Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60. However, this sharing is rarely equal. Electronegativity is the measure of an atom's ability to attract that shared pair of electrons toward itself. Think of it as the 'greediness' of an atom for electrons; the more electronegative an atom is, the more strongly it pulls electrons away from its partner. To understand why some atoms are 'greedier' than others, we must look at the Periodic Trends. These trends are driven by two factors: the positive charge of the nucleus (protons) and the distance of the outer electrons from that nucleus. As we move across a period (left to right), the number of protons increases, which exerts a stronger pull on the electrons. Simultaneously, the atoms are not adding new shells, so the size remains relatively compact, making the pull very effective. This is why non-metals like Fluorine, Oxygen, and Chlorine have such high electronegativity Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.47. Conversely, as we move down a group, electronegativity decreases. Although the number of protons increases, new electron shells are being added, which places the outer electrons much further from the nucleus. This 'shielding effect' by inner shells weakens the nucleus's grip on bonding electrons. In chemical reactions, this concept is vital because it determines the oxidation state: the more electronegative atom in a bond is assigned a negative oxidation number because it 'claims' the electrons, while the less electronegative atom is assigned a positive one.

Direction	Electronegativity Trend	Reasoning
Across a Period (→)	Increases	Higher nuclear charge (more protons) pulls electrons closer.
Down a Group (↓)	Decreases	Increased atomic radius and shielding make the nuclear pull weaker.

Sources: Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.60; Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.47

5. Group 17 Elements: The Halogens (intermediate)

The Halogens (Group 17) consist of Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). Their name literally translates to "salt-formers," reflecting their high reactivity with metals to form ionic salts. From a chemical perspective, they all possess seven electrons in their outermost shell (ns²np⁵), needing just one more electron to reach a stable noble gas configuration. This makes them highly electronegative—with Fluorine being the most electronegative element in the entire periodic table. In organic chemistry, they often act as heteroatoms, replacing hydrogen in hydrocarbons to form "haloalkanes" like chloropropane or bromopropane, thereby conferring specific chemical properties to the molecule Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.66-68. While the "default" oxidation state for halogens is -1 (as seen in Hydrogen Iodide, HI), their chemistry is much more versatile. Except for Fluorine, halogens can exhibit positive oxidation states like +1, +3, +5, and +7 because they possess vacant d-orbitals that allow them to expand their octet. When a halogen bonds with a more electronegative atom, it loses its negative status. For instance, in Iodine Monochloride (ICl), Chlorine is more electronegative than Iodine; thus, Chlorine takes the -1 state, forcing Iodine into a +1 state. This hierarchy of electronegativity is a fundamental principle in predicting how these elements will behave in complex molecules. In high-oxygen environments or oxyacids, halogens can reach their maximum oxidation state of +7. A classic example is Periodic acid (HIO₄), where Iodine is surrounded by four oxygen atoms. Since oxygen is much more electronegative and usually carries a -2 charge, the central Iodine must balance the charge to maintain neutrality, resulting in a +7 state. Understanding these shifts is vital for environmental science too; for example, bromine atoms are significantly more efficient than chlorine at destroying ozone molecules because of the specific stability and reactivity of their oxide intermediates Environment, Shankar IAS Academy, Ozone Depletion, p.269.

Molecule	Iodine Oxidation State	Reasoning
HI	-1	Iodine is more electronegative than Hydrogen (+1).
I₂	0	Elemental form: atoms share electrons equally.
ICl	+1	Chlorine is more electronegative than Iodine.
HIO₄	+7	High number of Oxygen atoms (-2 each) pull electrons away.

Sources: Science, class X (NCERT 2025 ed.), Carbon and its Compounds, p.66-68; Environment, Shankar IAS Academy, Ozone Depletion, p.269

6. Rules for Calculating Oxidation States (exam-level)

To understand chemical reactions deeply, we move beyond the simple idea of losing or gaining oxygen Science, class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12 and use Oxidation States. This is a bookkeeping system where we assign a formal charge to atoms to track electron movement. Think of it as a set of 'priority rules' used to determine which atom 'owns' the electrons in a bond. The fundamental rule is the Sum Rule: in a neutral compound, the sum of all oxidation states must be zero, while in a polyatomic ion, the sum must equal the ion's charge.

We start with fixed anchors. For instance, atoms in their elemental state (like pure Iron, Fe, or Nitrogen gas, N₂) always have an oxidation state of 0. When compounds form, we look to reliable elements like Oxygen, which is almost always assigned -2 Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.41, and Hydrogen, which is +1 when bonded to non-metals. For example, in water (H₂O), two +1 hydrogens (+2 total) perfectly balance one -2 oxygen to reach zero.

When two non-metals bond, the more electronegative element (the one hungrier for electrons) gets the negative state. Consider the rules applied to Iodine in different environments:

• I₂ (Elemental): Since it is a pure element, its state is 0.
• HI (Hydrogen Iodide): H is +1, so Iodine must be -1 to balance the molecule.
• ICl (Iodine Monochloride): Chlorine is more electronegative than Iodine. Thus, Cl is assigned -1, forcing Iodine to be +1.
• HIO₄ (Periodic Acid): Using our anchors (H = +1, O = -2), the calculation is: (+1) + (I) + 4(-2) = 0. This simplifies to 1 + I - 8 = 0, meaning Iodine has a high oxidation state of +7.

Sources: Science, class X (NCERT 2025 ed.), Chemical Reactions and Equations, p.12; Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.41

7. Solving the Original PYQ (exam-level)

Great job completing the foundational modules! This question is a perfect synthesis of the rules for assigning oxidation numbers and electronegativity trends that you just mastered. To solve this, you need to bring together the rules for elemental states, the hydrogen-halogen convention, and the algebraic sum of oxidation states in neutral molecules. By applying these building blocks systematically, you can transform a list of chemical formulas into a simple numerical sequence that reveals the correct answer.

Let’s walk through the reasoning as a coach would. First, identify the elemental state: in I2, the oxidation number is always 0. Next, look at HI; since hydrogen is +1, Iodine must be -1 to maintain neutrality. For ICl, remember your electronegativity lessons—Chlorine is more electronegative than Iodine, so Cl is assigned -1, meaning Iodine must be +1. Finally, in HIO4, we calculate the balance: [+1 (H) + x (I) + (-8 from four Oxygens) = 0], which gives Iodine a high state of +7. Arranging these— -1, 0, +1, and +7—leads us directly to the correct sequence in (B) HI < I2 < ICI < HI04.

UPSC often uses directional traps and relative electronegativity to test your precision. Option (A) is a classic "reverse trap" where the values are correct but the order is decreasing rather than increasing. Options (C) and (D) are distractors designed to catch students who fail to read the full requirement of the question. A common mistake is assuming Iodine is always -1 because it is a halogen; however, as NCERT Chemistry Class XI emphasizes, when a halogen bonds with a more electronegative element like Oxygen or Chlorine, it must take a positive oxidation state. Recognizing this distinction is the key to avoiding the most common pitfalls in these types of chemistry-based GS questions.