Amphiprotic Intermediate Ions

Introduction

An intermediate member of a polyprotic acid family can often act both as an acid and as a base. Bicarbonate HCO₃⁻ can donate H⁺ to become CO₃²⁻ or accept H⁺ to become the carbonic-acid/CO₂ side of the family. Dihydrogen phosphate H₂PO₄⁻ similarly has two possible proton-transfer roles. Calling such a species amphiprotic identifies those roles without claiming equal strength in both directions.

Core explanation

For a generic diprotic acid H₂A, the sequence is H₂A ⇌ H⁺ + HA⁻ and HA⁻ ⇌ H⁺ + A²⁻. The middle species HA⁻ is amphiprotic. As an acid, HA⁻ donates H⁺ with acid constant Ka2. As a base, HA⁻ accepts H⁺ to become H₂A; its base constant for reaction with water is Kb = Kw/Ka1 in a consistent aqueous idealisation. Compare Ka2 and Kb to see which tendency is stronger in water; the mere word “amphiprotic” does not predict pH by itself.

In water, write the acid reaction HA⁻ + H₂O ⇌ A²⁻ + H₃O⁺ and the base reaction HA⁻ + H₂O ⇌ H₂A + OH⁻. These two equations show opposite proton movements. The first adds hydronium, while the second adds hydroxide. Depending on constants and concentration, one may dominate net solution behaviour, while both still occur at equilibrium.

HCO₃⁻ is an important example. It can be protonated toward dissolved CO₂/carbonic acid or deprotonated toward carbonate. A solution of sodium bicarbonate often behaves mildly basic because the base tendency outweighs its acid tendency under familiar conditions, but its exact pH depends on concentration, temperature, gas exchange and ionic strength. In particular, an open bicarbonate solution can exchange CO₂ with air, so a sealed-system calculation may not match it exactly.

Phosphate offers another set of intermediates. H₂PO₄⁻ can accept H⁺ to become H₃PO₄ or donate H⁺ to become HPO₄²⁻. HPO₄²⁻ can accept H⁺ to become H₂PO₄⁻ or donate H⁺ to become PO₄³⁻. Their acid/base behaviour differs because they sit between different pairs of stepwise Ka values. The fully protonated H₃PO₄ is predominantly an acid donor in the usual model, while fully deprotonated PO₄³⁻ is a base acceptor.

Charge bookkeeping is essential. HCO₃⁻ losing H⁺ gives CO₃²⁻: charge changes from −1 to −2. Accepting H⁺ gives a neutral carbonic-acid form. For phosphate, H₂PO₄⁻ losing H⁺ gives HPO₄²⁻. An incorrect charge can produce an impossible electroneutrality equation even if the species names are right.

Amphiprotic species can form buffers with either adjacent family member. H₂PO₄⁻/HPO₄²⁻ is a conjugate pair near the corresponding pKa; HCO₃⁻/CO₃²⁻ is another. A buffer pair contains two species differing by one proton, not simply any two salts with similar names.

Step-by-step reasoning

1. List the polyprotic sequence from most to least protonated. 2. Pick an intermediate species with a neighbour on both sides. 3. Write proton donation to the lower-proton species. 4. Write proton acceptance to the higher-proton species. 5. Check formulas, charges and appropriate Ka or Kb labels.

Visual explanation

Draw H₂A ⇌ HA⁻ ⇌ A²⁻ as three boxes. Above HA⁻, draw one arrow to H₂A marked “accept H⁺” and one to A²⁻ marked “donate H⁺.”

Real-world analogy

A person in the middle of a row can pass a ball to a neighbour on one side or receive one from the other. An intermediate ion can donate or accept a proton, although the two exchanges may have different preferences.

Real-world example

The sodium bicarbonate ion HCO₃⁻ can respond to added acid by accepting H⁺ and to added base by donating H⁺. Its amphiprotic nature helps explain its role in carbonate buffering systems.

Why?

Why is an intermediate ion more likely to be amphiprotic than the terminal forms of its family? It has a removable proton and also a more protonated neighbour it can reach by accepting another proton.

Common misconception

“Amphiprotic means exactly neutral in water.” It means capable of both proton donation and acceptance; relative constants determine whether the solution is acidic, basic or near neutral.

Worked example

Classify H₂PO₄⁻. It donates H⁺: H₂PO₄⁻ + H₂O ⇌ HPO₄²⁻ + H₃O⁺. It accepts H⁺: H₂PO₄⁻ + H₂O ⇌ H₃PO₄ + OH⁻. The ion is amphiprotic. If a problem supplies Ka1 and Ka2, compare Ka2 for acid behaviour with Kw/Ka1 for base behaviour rather than assuming equal tendencies.

Quick check

1. What are the conjugate acid and conjugate base of HCO₃⁻ in the simplified carbonate family? Answer: Its conjugate acid is H₂CO₃ or the appropriate CO₂ pool under a consistent convention; its conjugate base is CO₃²⁻.

Exam focus

Write both water-reaction equations and check charges. Compare relevant constants if asked to predict pH direction.

Advanced insight

An amphiprotic salt's approximate pH can be near the midpoint of neighbouring pKa values under suitable concentration assumptions. The next page derives and limits that approximation.

Summary

An amphiprotic intermediate can donate H⁺ to a less protonated family member or accept H⁺ to a more protonated one. HA⁻, HCO₃⁻ and phosphate intermediates are examples. Their solution pH depends on relative acid/base equilibria, not on the label alone.

Practice questions

1. What does HA⁻ become when it donates H⁺ in a diprotic family? Answer: A²⁻. 2. What does HA⁻ become when it accepts H⁺? Answer: H₂A. 3. Is HPO₄²⁻ amphiprotic? Answer: Yes. It can accept H⁺ to become H₂PO₄⁻ or donate H⁺ to become PO₄³⁻. 4. Does amphiprotic behaviour guarantee pH 7? Answer: No. The acid and base equilibrium strengths and conditions determine pH.