Q.The compound AgF2 is unstable compound. However, if formed, the compound acts as a very strong oxidising agent. Why ?

Concept & Intuition

A substance acts as an oxidising agent if it can accept electrons (its own oxidation number decreases), and as a reducing agent if it can donate electrons (its oxidation number increases). Whether a compound can do one or both depends on the range of stable oxidation states of the element that changes oxidation number. If an element is already in its highest possible oxidation state, it can only be reduced (oxidising agent only). If it is in its lowest, it can only be oxidised (reducing agent only). If it has intermediate states, it can do both.

Step-by-step reasoning

1. Examine the assertion

   • Sulphur dioxide (SO₂): Sulphur in SO₂ has oxidation state +4. Sulphur can exist in states from –2 (e.g., H₂S) to +6 (e.g., H₂SO₄). Since +4 is intermediate, SO₂ can be oxidised to +6 (acts as reducing agent) or reduced to lower states (acts as oxidising agent).
   • Hydrogen peroxide (H₂O₂): Oxygen here is –1. Oxygen normally has states –2 (in H₂O) and 0 (in O₂). The –1 state is intermediate, so H₂O₂ can be oxidised to O₂ (reducing agent) or reduced to H₂O (oxidising agent).
   • Nitric acid (HNO₃): Nitrogen in HNO₃ is +5, the highest stable oxidation state for nitrogen (compare: NH₃ is –3, N₂ is 0, NO is +2, NO₂ is +4). Being at the maximum, nitrogen cannot be further oxidised; it can only be reduced. Hence HNO₃ acts only as an oxidising agent. → The assertion is correct.

2. Examine the reason

   • The reason states: “Sulphur and Oxygen can exhibit more than one stable oxidation state while Nitrogen does not.”
   • Sulphur: –2, 0, +4, +6 (many stable states).
   • Oxygen: –2, –1, 0 (stable states, e.g., in peroxides and superoxides). …

1. Decode the name

Pyridinium Chlorochromate:

• pyridinium — the protonated pyridine cation, $\mathrm{C_5H_5NH^{+}}$ (the source of both the pyridine and the proton, from HCl);
• chloro-chromate — the anion $[\mathrm{CrO_3Cl}]^{-}$ (the chromium is supplied by $\mathrm{CrO_3}$, the chlorine by HCl).

So all three components — chromium trioxide, pyridine, and HCl — are required, which is exactly what option (D) says.

2. Why PCC matters synthetically

PCC is a mild oxidising agent used in a non-aqueous solvent (dichloromethane, $\mathrm{CH_2Cl_2}$). Its defining property:

The absence of water is the key — a strong aqueous oxidant like $\mathrm{K_2Cr_2O_7/H_2SO_4}$ would hydrate the aldehyde and over-oxidise it to the carboxylic acid.

3. Eliminating the distractors …

The question is about controlled oxidation of primary alcohols. A primary alcohol ($RCH_2OH$) can be oxidised first to an aldehyde ($RCHO$) and then further to a carboxylic acid ($RCOOH$). The challenge is to stop the reaction at the aldehyde stage — which requires a mild oxidising agent that works in anhydrous conditions.

Let’s examine each option:

1. PCC (pyridinium chlorochromate) — This is a mild, anhydrous reagent specifically designed for this purpose. It is prepared by dissolving chromium trioxide in pyridine and hydrochloric acid. PCC in dichloromethane ($CH_2Cl_2$) cleanly oxidises primary alcohols to aldehydes without further oxidation to acids. This is the go-to reagent in organic chemistry for aldehyde synthesis.

2. Alkaline KMnO$_4$ — This is a very strong oxidising agent. Under alkaline conditions, it oxidises primary alcohols all the way to carboxylic acids (via the aldehyde, which is rapidly consumed). It cannot stop at the aldehyde stage.

3. Acidified K$_2$Cr$_2$O$_7$ — This is also a strong oxidising agent (the familiar "chromic acid" mixture). It oxidises primary alcohols to carboxylic acids, not aldehydes. The aldehyde, if formed, is quickly further oxidised under the acidic aqueous conditions. …

The key to this question lies in understanding what makes a compound a good reducing agent. A reducing agent donates electrons (or hydrogen atoms) to another substance, getting oxidized itself in the process. So we need to look at the structure of hypophosphorous acid and see what part of it is easily given away.

Hypophosphorous acid has the formula $H_3PO_2$, but its structure is not what you might guess from the formula alone. It is a monoprotic acid, meaning only one of its three hydrogens is acidic (ionizable as $H^+$). The other two hydrogens are directly bonded to the phosphorus atom. The structure is best written as $H_2P(O)OH$ — a central phosphorus atom with one double-bonded oxygen, one hydroxyl group ($-OH$), and two hydrogens directly attached to the phosphorus.

Now, the P–H bond is relatively weak and polarizable. In a redox reaction, these two hydrogen atoms (each with an oxidation state of +1 when bonded to P) can be readily released as hydrogen atoms ($H^\cdot$) or as hydride ions ($H^-$), effectively reducing another species. This is the source of the strong reducing power.

Let's examine each option carefully.

1. Option (A): Two P–H bonds. This is the correct reason. The two hydrogen atoms directly attached to phosphorus are the ones that get donated in reduction reactions. For example, hypophosphorous acid reduces copper(II) salts to copper metal, and it itself gets oxidized to phosphoric acid ($H_3PO_4$). The two P–H bonds are the "fuel" for this reduction.

2. Option (B): Presence of phosphorus in its highest oxidation state. This is false. In $H_3PO_2$, the oxidation state of phosphorus is +1 (calculate: $3(+1) + x + 2(-2) = 0 \Rightarrow x = +1$). The highest oxidation state of phosphorus is +5 (as in $H_3PO_4$). A species in its highest oxidation state can only be reduced, not oxidized — it would be an oxidizing agent, not a reducing agent. So this option is the opposite of the truth.

3. Option (C): Its concentration. Concentration affects the rate of a reaction and the amount of reducing agent available, but it does not determine the inherent reducing property of the substance. A dilute solution of a strong reducing agent is still a strong reducing agent; a concentrated solution of a weak reducing agent is still weak. The property is intrinsic to the molecule, not its concentration. …

The Concept: What Makes an Antioxidant?

An antioxidant is a substance that prevents or delays the oxidation of other molecules. In food chemistry, oxidation is the enemy — it causes fats and oils to turn rancid (developing a bad smell and taste), and can also cause discoloration in fruits and vegetables. Antioxidants work by reacting with free radicals (highly reactive molecules that start the oxidation chain reaction) and neutralizing them, effectively breaking the chain.

Common food antioxidants include natural ones like Vitamin C (ascorbic acid) and Vitamin E (tocopherols), and synthetic ones like BHA (Butylated Hydroxy Anisole) and BHT (Butylated Hydroxy Toluene). The question asks you to pick the one that acts as an antioxidant from a list of common food additives.

Step-by-Step Reasoning

1. Identify the role of each option.

   Let’s look at what each substance is primarily used for in food:

   • Salt (NaCl): A preservative and flavor enhancer. It works by drawing out water from microbes, not by preventing oxidation.
   • BHA (Butylated Hydroxy Anisole): A synthetic phenolic compound. It is specifically added to foods (like butter, lard, and packaged snacks) to prevent the oxidation of fats and oils.
   • Saccharin: An artificial sweetener. It provides sweetness without calories, but has no role in preventing oxidation.
   • Sugar syrup: A sweetener and preservative (by creating a high osmotic pressure), but it does not act as an antioxidant.

2. Match the function to the definition.

   The question asks for an "anti-oxidant". Only BHA fits this description. Salt and sugar syrup are preservatives through different mechanisms (osmotic effect), and saccharin is a sweetener. …

The key idea here is phase uniformity. In homogeneous catalysis, the catalyst and the reactants form a single phase — all gases, all liquids, or all solids. If the catalyst is in a different phase from the reactants, it's heterogeneous catalysis.

Let’s examine each process:

1. Oxidation of NH₃ in Ostwald’s process

   This is the first step in making nitric acid. Ammonia is oxidised over a platinum–rhodium gauze (a solid catalyst). The reactants are gases, the catalyst is solid — different phases. This is heterogeneous catalysis.

2. Oxidation of SO₂ in the lead chamber process

   Here, sulphur dioxide is oxidised to sulphur trioxide using nitric oxide (NO) as a catalyst. Both SO₂ (reactant) and NO (catalyst) are gases. They are in the same phase (gas). This is a classic example of homogeneous catalysis.

3. Oxidation of SO₂ in the contact process

   In the modern contact process, SO₂ is oxidised over vanadium pentoxide (V₂O₅) — a solid catalyst. Again, the reactants are gases, the catalyst is solid — heterogeneous.

4. Manufacture of NH₃ by Haber’s process

   Nitrogen and hydrogen react over finely divided iron (with promoters like K₂O, Al₂O₃). Solid catalyst, gaseous reactants — heterogeneous. …