Oxoacids of Phosphorus

Introduction

Phosphorus oxoacids show why a molecular formula alone can mislead. H₃PO₃ and H₃PO₄ each contain three hydrogen atoms, but they do not have the same number of readily ionizable protons. The structural location of H—on oxygen or directly on phosphorus—determines acid basicity in the common aqueous model.

Core explanation

Phosphoric acid, H₃PO₄, is commonly represented OP(OH)₃. Three hydrogens lie in O–H groups, so it is triprotic: it can lose protons stepwise to form H₂PO₄⁻, HPO₄²⁻ and PO₄³⁻. These steps have different equilibria; “triprotic” does not mean all three protons leave completely at once. The first dissociation is stronger than later ones because removing a proton from an increasingly negative ion is less favorable.

Phosphorous acid, H₃PO₃, is commonly represented HPO(OH)₂. Two hydrogens are on oxygen and one is directly bonded to phosphorus. It is therefore diprotic in ordinary aqueous acid-base behavior, yielding H₂PO₃⁻ and HPO₃²⁻ through its two O–H groups. The P–H hydrogen is not normally counted as a readily ionizable acid proton. A formula beginning H₃ does not establish three acid dissociations.

Hypophosphorous acid, H₃PO₂, can be represented H₂PO(OH). Only one hydrogen is attached through oxygen, so it is monoprotic. Its two P–H bonds are important to redox chemistry rather than ordinary proton donation. The sequence H₃PO₂, H₃PO₃ and H₃PO₄ gives one, two and three ionizable O–H protons respectively, despite all formulas having three H atoms.

Assign oxidation states with H +1 and O −2 in these neutral acids. For H₃PO₂, 3(+1)+x+2(−2)=0, so P is +1. For H₃PO₃, 3+x−6=0, giving +3. For H₃PO₄, 3+x−8=0, giving +5. These formal numbers track phosphorus oxidation, while acid basicity tracks O–H groups. The two sequences happen to rise together here but answer separate questions.

The common structural P=O drawing is a useful Lewis representation; modern bonding discussions can describe electron density more subtly than a fixed localized double bond. The practical exam lesson is the position of P–H versus P–OH and charge-balanced deprotonation. It is unnecessary to invoke impossible extra proton release merely to match a formula.

Phosphate chemistry is important in fertilizers, biology and mineral structure. H₂PO₄⁻ and HPO₄²⁻ form a buffer pair in appropriate pH ranges. Their proportions depend on pH rather than on the original acid's formula alone. Phosphite and hypophosphite species differ chemically from phosphate, including reducing behavior associated with lower phosphorus oxidation states.

Step-by-step reasoning

1. Draw or obtain the structural formula of the phosphorus oxoacid. 2. Count O–H groups to determine ordinary acid basicity. 3. Exclude P–H hydrogen from that count. 4. Assign formal H +1 and O −2 to calculate P oxidation state. 5. Write stepwise conjugate-base formulas and charges.

Visual explanation

Place H₂PO(OH), HPO(OH)₂ and OP(OH)₃ in three columns. Circle only the O–H hydrogens, counting one, two and three. Below, calculate P states +1, +3 and +5 separately to emphasize that proton count and oxidation number are different calculations.

Real-world analogy

Three keys on a ring may look alike in an inventory, but only keys attached to accessible doors can be used for a particular task. Total hydrogen count is the inventory; the O–H positions identify the protons released in aqueous acid-base steps.

Real-world example

Phosphate buffers use the pair H₂PO₄⁻/HPO₄²⁻. Adding a small amount of acid or base shifts their relative amounts, illustrating that stepwise deprotonation is chemically useful rather than only a naming rule.

Why?

Why is H₃PO₃ diprotic? Its common structure has two P–OH groups and one P–H bond. The two O–H protons enter ordinary acid dissociation; the P–H hydrogen does not behave as a third such proton.

Common misconception

“Every H in an acid formula is replaceable as H⁺.” Structure decides which hydrogens are attached to electronegative oxygen and participate in the ordinary acid equilibria.

Worked example

For H₃PO₃, calculate phosphorus state and acid basicity independently. Let P be x: 3(+1)+x+3(−2)=0, so x=+3. Draw HPO(OH)₂: two O–H groups can lose protons stepwise, so the acid is diprotic. The result is not “triprotic because H₃” and not “three protons because P is +3.”

Quick check

1. How many ordinary ionizable protons does H₃PO₂ have? Answer: One, from its single P–OH group.

Exam focus

Show structures, count P–OH groups and calculate oxidation states separately. Give the correct mono-, di- and triprotic order for H₃PO₂, H₃PO₃ and H₃PO₄.

Advanced insight

Phosphorous and hypophosphorous acids can act as reducing agents because phosphorus at +3 or +1 can be oxidized toward +5. Acid dissociation and redox may therefore occur in different contexts without sharing a single “acid strength” explanation.

Summary

Phosphorus oxoacid basicity depends on O–H groups: H₃PO₂ is monoprotic, H₃PO₃ diprotic and H₃PO₄ triprotic. Phosphorus oxidation states are +1, +3 and +5 respectively. Count protons and oxidation states by different rules.

Practice questions

1. Give the common structure and basicity of H₃PO₃. Answer: HPO(OH)₂; diprotic because it has two O–H groups. 2. What is phosphorus's oxidation state in H₃PO₄? Answer: +5, from 3(+1)+x+4(−2)=0. 3. What is the conjugate base after H₃PO₄ loses one proton? Answer: H₂PO₄⁻.