Q.What happens when H3PO3 is heated?

First sequence: $P_2O_3 \xrightarrow{H_2O} H_3PO_3$ (phosphorous acid, $X$). Condensation of two $H_3PO_3$ units gives pyrophosphorous acid, $Y = H_4P_2O_5$, structure $(HO)(H)(O{=})P{-}O{-}P(=O)(H)(OH)$:

• $2$ $P{-}OH$ bonds, $2$ $P{-}H$ bonds, $2$ $P{=}O$ bonds, $1$ $P{-}O{-}P$ bond.

Second sequence: $P_4O_{10} \xrightarrow{H_2O} H_3PO_4$ (orthophosphoric acid, $Z$); on heating ($\Delta$) it condenses to pyrophosphoric acid, $W = H_4P_2O_7$, structure $(HO)_2(O{=})P{-}O{-}P(=O)(OH)_2$:

• $4$ $P{-}OH$ bonds, $0$ $P{-}H$ bonds, $2$ $P{=}O$ bonds, $1$ $P{-}O{-}P$ bond. …

Concept & Intuition

Oxoacids of phosphorus are formed by reacting phosphorus halides, oxides, or elemental phosphorus with water or alkali. The oxidation state of phosphorus in the product depends on the starting material’s oxidation state and whether disproportionation occurs. For phosphorus, common oxidation states are +3 (e.g., in phosphorous acid, $\mathrm{H_3PO_3}$) and +5 (e.g., in phosphoric acid, $\mathrm{H_3PO_4}$). We need to identify which reactions produce an oxoacid where P is +3.

Step-by-step reasoning

1. Set I: $\mathrm{PCl_3 + H_3PO_3}$

   • $\mathrm{PCl_3}$ contains P in +3 state (since Cl is -1, total -3, so P = +3).
   • $\mathrm{H_3PO_3}$ is phosphorous acid, where P is also +3 (check: H = +1 each, O = -2 each; 3(+1) + 3(-2) = -3, so P = +3).
   • Mixing them does not change oxidation states; the product is still an oxoacid with P(+3). So I works.

2. Set II: $\mathrm{P_2O_3 + H_2O}$

   • $\mathrm{P_2O_3}$ is phosphorus trioxide. In $\mathrm{P_2O_3}$, O is -2 each, total -6, so each P is +3.
   • Reaction: $\mathrm{P_2O_3 + 3H_2O \rightarrow 2H_3PO_3}$ (phosphorous acid).
   • The product $\mathrm{H_3PO_3}$ has P in +3 state. So II works.

3. Set III: Red $\mathrm{P_4}$ + alkali

   • Elemental phosphorus (red $\mathrm{P_4}$) has oxidation state 0.
   • With hot concentrated alkali, it disproportionates: $\mathrm{P_4 + 3OH^- + 3H_2O \rightarrow PH_3 + 3H_2PO_2^-}$ The product includes hypophosphite ($\mathrm{H_2PO_2^-}$), where P is +1 (not +3). Also, phosphine ($\mathrm{PH_3}$) has P in -3. …

Concept & Intuition

Basicity in oxoacids is not simply the total number of hydrogens in the formula. Instead, it counts only those hydrogens that can be released as $\mathrm{H^+}$ ions — which are the ones attached to oxygen atoms (O–H bonds). Hydrogens bonded directly to the central atom (like P–H bonds in phosphorus oxoacids) are not acidic and do not contribute to basicity. So to find which pair both have basicity 2, we must examine the structure of each acid, not just its molecular formula.

1. Recall the structures of the given acids

   • $\mathrm{H_3PO_3}$ (phosphorous acid): Has one P–H bond and two O–H bonds → basicity = 2.
   • $\mathrm{H_2SO_4}$ (sulfuric acid): Both hydrogens are on oxygen → basicity = 2.
   • $\mathrm{H_3PO_2}$ (hypophosphorous acid): Has two P–H bonds and only one O–H bond → basicity = 1.
   • $\mathrm{H_2SO_3}$ (sulfurous acid): Both hydrogens are on oxygen → basicity = 2.
   • $\mathrm{H_3PO_4}$ (phosphoric acid): All three hydrogens are on oxygen → basicity = 3.
   • $\mathrm{H_2S_2O_8}$ (peroxodisulfuric acid): Both hydrogens are on oxygen → basicity = 2.

2. Check each option

   • (A) $\mathrm{H_3PO_3}$ (basicity 2) and $\mathrm{H_2SO_4}$ (basicity 2) → both have basicity 2.
   • (B) $\mathrm{H_3PO_2}$ (basicity 1) and $\mathrm{H_2SO_3}$ (basicity 2) → not both 2. …

The problem asks us to identify an oxoacid of phosphorus that can reduce silver nitrate to metallic silver, and in the process gets oxidized to another phosphorus oxoacid. This is a classic redox reaction involving phosphorus in a lower oxidation state being oxidized to a higher one.

Concept and Intuition:

Phosphorus forms several oxoacids where the central phosphorus atom has different oxidation states. The reducing power of these acids depends on how easily phosphorus can lose electrons (be oxidized). Hypophosphorous acid ($\mathrm{H_3PO_2}$) has phosphorus in the +1 oxidation state — the lowest among common phosphorus oxoacids — so it is the strongest reducing agent. When it reduces silver ions ($\mathrm{Ag^+}$) to silver metal, phosphorus itself gets oxidized to the +5 state, forming phosphoric acid ($\mathrm{H_3PO_4}$). Other oxoacids like phosphorous acid ($\mathrm{H_3PO_3}$, +3) are weaker reductants, and metaphosphoric acid ($\mathrm{HPO_3}$, +5) cannot be oxidized further.

Let’s work through the reasoning step by step.

1. Determine the oxidation states of phosphorus in each candidate.

   • In $\mathrm{HPO_3}$ (metaphosphoric acid): H = +1, O = -2, so P = +5.
   • In $\mathrm{H_3PO_4}$ (orthophosphoric acid): H = +1, O = -2, so P = +5.
   • In $\mathrm{H_3PO_2}$ (hypophosphorous acid): H = +1, O = -2, so P = +1.
   • In $\mathrm{H_3PO_3}$ (phosphorous acid): H = +1, O = -2, so P = +3.
   • $\mathrm{HNO_3}$ is not a phosphorus oxoacid, so option (D) is immediately suspect.

2. Identify which phosphorus oxoacid can reduce $\mathrm{Ag^+}$.

   Silver ions ($\mathrm{Ag^+}$) are reduced to silver metal ($\mathrm{Ag}$) by gaining an electron. The reducing agent must lose electrons (be oxidized). Only a phosphorus oxoacid in a lower oxidation state can be oxidized to a higher state.

   • $\mathrm{HPO_3}$ (P = +5) is already at the maximum oxidation state for phosphorus; it cannot be oxidized further, so it cannot act as a reducing agent.
   • $\mathrm{H_3PO_4}$ (P = +5) is also at maximum; same reasoning.
   • $\mathrm{H_3PO_3}$ (P = +3) can be oxidized to +5, so it is a possible reducing agent, but it is weaker than $\mathrm{H_3PO_2}$.
   • $\mathrm{H_3PO_2}$ (P = +1) can be oxidized to +3 or +5, and it is known to be a strong reducing agent — in fact, it is used to reduce silver salts in electroless silver plating.

3. Check the oxidation product.

   The problem states that the oxoacid X gets oxidized to another compound Y. For $\mathrm{H_3PO_2}$, the common oxidation product is $\mathrm{H_3PO_4}$ (phosphoric acid), not $\mathrm{H_3PO_3}$ (which would be a partial oxidation). The reaction with silver nitrate is well-known:

   $\mathrm{H_3PO_2 + 2AgNO_3 + 2H_2O \rightarrow H_3PO_4 + 2Ag + 2HNO_3}$ …

The key to this question lies in understanding the structure of oxoacids of phosphorus. Unlike many other oxoacids, phosphorus oxoacids often have hydrogen atoms attached directly to the phosphorus atom (P–H bonds) in addition to the –OH groups. The number of P–H bonds determines the basicity (number of ionizable hydrogens) and the reducing nature of the acid.

Let’s examine each option systematically.

1. Hypophosphoric acid ($H_4P_2O_6$): This is a diphosphorus acid with a P–P bond. Its structure has each phosphorus atom bonded to two oxygen atoms (one double-bonded) and one –OH group. No hydrogen is directly attached to phosphorus; all four hydrogens are in –OH groups. So it has zero P–H bonds.

2. Orthophosphoric acid ($H_3PO_4$): The classic phosphoric acid. Its structure is a central phosphorus with one double-bonded oxygen and three –OH groups. All three hydrogens are on oxygen. Again, zero P–H bonds.

3. Pyrophosphoric acid ($H_4P_2O_7$): Formed by condensation of two orthophosphoric acid molecules (eliminating water). Each phosphorus has one double-bonded oxygen and two –OH groups, with a P–O–P bridge. All four hydrogens are on oxygen. Zero P–H bonds. …

Concept & Intuition

Phosphorus oxides react with water to give oxyacids. The number of P=O (phosphoryl) bonds in these acids is fixed by the oxidation state and structure of phosphorus. Instead of memorizing, we can deduce the formulas by balancing the reactions and then drawing the Lewis structures. The classic trap is confusing P–OH single bonds with P=O double bonds — only the latter count.

Step-by-step reasoning

1. Identify the products $X$ and $Y$
   • $P_2O_3$ is phosphorus(III) oxide. With water, it forms phosphorous acid:

 So $X = H_3PO_3$.  

• $P_4O_{10}$ is phosphorus(V) oxide. With water, it forms phosphoric acid:

 So $Y = H_3PO_4$.

2. Determine the structure of $H_3PO_3$ (X)

• In $H_3PO_3$, phosphorus is in the +3 oxidation state. The molecule has one P=O double bond and two P–OH single bonds, plus one P–H bond (yes, a direct P–H bond, not O–H).
• The structure is:

 Count of P=O bonds = **1**.

3. Determine the structure of $H_3PO_4$ (Y)

• In $H_3PO_4$, phosphorus is in the +5 oxidation state. The molecule has one P=O double bond and three P–OH single bonds.
• The structure is: $\text{O=P(OH)}_3$ …

Concept. A $\text{P–O–P}$ (oxygen-bridged) linkage forms when two phosphorus tetrahedra condense by sharing a bridging oxygen atom. A direct $\text{P–P}$ bond is a different structural feature and is not an oxygen bridge.

Each acid.

• III. $\text{H}_4\text{P}_2\text{O}_7$ — pyrophosphoric acid, obtained by condensing two $\text{H}_3\text{PO}_4$ units with loss of water; the two tetrahedra are joined through a bridging oxygen, giving one $\text{P–O–P}$ bond.
• IV. $(\text{HPO}_3)_3$ — trimetaphosphoric acid, a six-membered ring of alternating P and O atoms, giving three $\text{P–O–P}$ bonds. …

The chemistry of phosphorus oxoacids hinges on understanding their structures, not just formulas. The oxidation state of phosphorus depends on how many P–OH groups, P=O bonds, and P–H bonds are present. P–H bonds are reducing (they lower the oxidation state), while P–OH and P=O are oxidizing. Metaphosphoric acid's behavior follows from its tendency to condense.

Let me examine each statement:

Statement I: Oxidation state of P in hypophosphoric acid is +4

Hypophosphoric acid has the formula $\ce{H4P2O6}$. Its structure contains a P–P bond with each phosphorus bearing two –OH groups and one =O group:

Actually, the correct structure is:

For each phosphorus atom:

• Two P–O (in –OH groups): contributes $2 \times (+1) = +2$ to oxidation state
• One P=O bond: contributes $+2$
• One P–P bond: contributes $0$ (same element)

Total oxidation state: $+2 + 2 = +4$

Alternatively, using the formula method: if we assign oxidation state $x$ to each P,

Statement I is correct.

Statement II: P–H bonds in orthophosphoric and pyrophosphoric acid

Orthophosphoric acid is $\ce{H3PO4}$ with structure:

All hydrogen atoms are in –OH groups bonded to oxygen, not directly to phosphorus. There are no P–H bonds.

Pyrophosphoric acid is $\ce{H4P2O7}$ with structure:

Again, all hydrogens are in –OH groups. There are no P–H bonds. …

The key is to identify the compound X formed in the first reaction, then determine Y from the second, and finally check each statement about Y.

Concept and Intuition

White phosphorus ($P_4$) reacts with hot concentrated alkali to give phosphine ($PH_3$) and a salt of hypophosphorous acid. The reaction here uses NaOH and water, with $CO_2$ possibly acting as a mild acid to help drive the formation. The salt X is sodium hypophosphite ($NaH_2PO_2$). When treated with HCl, it yields Y, which is hypophosphorous acid ($H_3PO_2$). This acid has a peculiar structure: only one hydrogen is ionizable (attached to oxygen), making it monobasic, despite having three hydrogens. The central phosphorus atom has an oxidation state of +1. Let’s verify each option.

Step-by-step reasoning

1. Identify X from the first reaction The reaction of $P_4$ with NaOH and water is a disproportionation:

Here, phosphorus in $P_4$ (oxidation state 0) is both reduced to -3 in $PH_3$ and oxidized to +1 in $\text{NaH}_2\text{PO}_2$ (sodium hypophosphite). So X is $\text{NaH}_2\text{PO}_2$.

2. Identify Y from the second reaction

Thus Y is hypophosphorous acid, $\text{H}_3\text{PO}_2$.

3. Analyze each statement about Y ($H_3PO_2$)
   • (A) It is also called phosphonic acid Phosphonic acid is actually $H_3PO_3$ (phosphorous acid). Hypophosphorous acid is phosphinic acid. So this name is incorrect for Y.
   • (B) It is a monobasic acid $H_3PO_2$ has only one ionizable H (the one bonded to oxygen); the two H atoms directly bonded to P are not acidic. So it is indeed monobasic — correct statement.
   • (C) It is a reducing agent …

The question asks you to match a specific bond count to a known oxoacid of phosphorus. This is a structural chemistry problem — you need to know the molecular formulas and the bonding patterns of common phosphorus oxoacids. The trick is that phosphorus can form multiple oxoacids with different numbers of P–O–H (acidic hydroxyl groups), P=O (phosphoryl groups), and P–O–P (anhydride bridges). Let’s work through each option.

1. Orthophosphoric acid ($\mathrm{H_3PO_4}$)

   Its structure has one phosphorus atom bonded to four oxygen atoms: one P=O and three P–O–H groups. There is no P–O–P bond because there’s only one phosphorus. So it has 3 P–O–H, 1 P=O, and 0 P–O–P bonds. This doesn’t match the required 4, 2, and 1.

2. Metaphosphoric acid ($\mathrm{HPO_3}$)

   This is a cyclic or polymeric form. In its simplest cyclic trimer ($\mathrm{(HPO_3)_3}$), each phosphorus has one P=O, one P–O–H, and two P–O–P bonds (linking to neighbours). For a single $\mathrm{HPO_3}$ unit, you get 1 P–O–H, 1 P=O, and 2 P–O–P bonds (if cyclic). That’s far from the required counts.

3. Pyrophosphoric acid ($\mathrm{H_4P_2O_7}$) …

The question as written is ambiguous: "Phosphorous and phosphoric acids" — but phosphorus itself is an element, not an acid. In standard Indian exam language, "phosphorous acid" means $H_3PO_3$ and "phosphoric acid" means $H_3PO_4$. The trick is to understand basicity — the number of ionizable hydrogen atoms attached to oxygen (not directly to phosphorus).

1. Basicity is not the same as the number of hydrogens in the formula.

   $H_3PO_3$ has three hydrogens, but only two are bonded to oxygen (and thus can be released as $H^+$). The third hydrogen is directly bonded to the phosphorus atom, making it non-ionizable under normal conditions. So $H_3PO_3$ is dibasic.

2. $H_3PO_4$ has all three hydrogens attached to oxygen.

   In phosphoric acid, the structure is $OP(OH)_3$ — each hydrogen is on an $-$OH group. All three can be donated stepwise, so it is tribasic.

3. Why does this matter? …