Brønsted–Lowry Acids in Inorganic Systems

Introduction

An acid does not need to produce H₃O⁺ in water to count as a Brønsted–Lowry acid. The definition is based on proton transfer: one species donates H⁺ and another accepts it. That makes it useful in gas reactions, liquid ammonia and metal-aqua hydrolysis as well as ordinary aqueous acid-base chemistry. Identifying both conjugate pairs prevents labels from being attached to a substance without reference to its reaction partner.

Core explanation

The general reaction is HA + B ⇌ A⁻ + BH⁺. HA and A⁻ are a conjugate acid-base pair; BH⁺ and B are another. Each pair differs by one proton. For NH₄⁺ + OH⁻ → NH₃ + H₂O, ammonium is the acid and hydroxide the base. NH₃ is the conjugate base of NH₄⁺, while H₂O is the conjugate acid of OH⁻. The charges and atoms balance, allowing the roles to be checked mechanically.

In water, HCl donates a proton to H₂O, forming H₃O⁺ and Cl⁻. But gaseous HCl can also transfer a proton to gaseous NH₃ to form NH₄Cl solid under suitable conditions. Water is not part of that reaction, yet the same proton-donor/acceptor description applies. In liquid ammonia, NH₄⁺ can act as an acid and NH₂⁻ as a base; the solvent itself can autoionise as 2NH₃ ⇌ NH₄⁺ + NH₂⁻. The chemical meaning of “strong acid in this solvent” depends on that medium's proton-transfer equilibria.

Water is amphiprotic: it can accept a proton from HCl or donate one to NH₃. Hydrogencarbonate HCO₃⁻ is another amphiprotic inorganic ion. It can accept H⁺ to form H₂CO₃/CO₂-related species or donate H⁺ to form CO₃²⁻. Its role depends on the partner. Calling HCO₃⁻ “an acid” or “a base” without specifying the reaction is incomplete.

Hydrated metal ions provide subtler Brønsted acidity. A small, highly charged metal such as Al³⁺ draws electron density from coordinated water, making an O–H bond easier to deprotonate: [Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅(OH)]²⁺ + H₃O⁺. The metal centre itself is a Lewis acid, while the coordinated water molecule acts as a Brønsted acid. The two acid-base models describe connected but different electron and proton events.

Acid strength is a relative equilibrium property. HCl behaves as a strong acid in water because proton transfer to water is strongly favoured, and water levels sufficiently strong acids to H₃O⁺. In another solvent, relative strengths can differ, and the conjugate base's solvation matters. A proton is not an unbound sphere floating in bulk solution; it is accepted by a base or solvent.

Step-by-step reasoning

1. Find which species loses one proton and which gains it. 2. Pair the donor with its deprotonated conjugate base. 3. Pair the acceptor with its protonated conjugate acid. 4. Check atom and charge balance, including solvent participation. 5. State the solvent and avoid transferring an aqueous strength ranking uncritically.

Visual explanation

Draw HA above A⁻ with a downward “−H⁺” arrow and B above BH⁺ with a downward “+H⁺” arrow. Next to the diagram place H₂O in the middle of two examples: accepting H⁺ from HCl and donating H⁺ to NH₃.

Real-world analogy

A person giving a book is the donor only relative to someone receiving it. In a different exchange, the former receiver might give the book onward. Water and hydrogencarbonate likewise can be acid or base depending on their reaction partner.

Real-world example

Liquid-ammonia chemistry supports proton-transfer reactions that differ from water chemistry. An ammonium salt and an amide base can neutralise to ammonia molecules. The example demonstrates that Brønsted–Lowry analysis is broader than an aqueous pH scale.

Why?

Why can [Al(H₂O)₆]³⁺ acidify water? The high-charge-density metal pulls electron density from ligand oxygen, weakening the ligand's O–H proton attachment. A surrounding water molecule can accept that proton and become H₃O⁺.

Common misconception

“An acid must contain H₃O⁺ before reaction” is false. An acid donates a proton to a base; H₃O⁺ is a product when water is the acceptor. Gas-phase and nonaqueous proton transfers still satisfy the definition.

Worked example

Analyse HCO₃⁻ + NH₃ ⇌ CO₃²⁻ + NH₄⁺. Hydrogencarbonate loses H⁺, so it is the acid and CO₃²⁻ its conjugate base. Ammonia accepts H⁺, so it is the base and NH₄⁺ its conjugate acid. The total charge is −1 on each side. The equilibrium position depends on the medium and relative acid strengths rather than just formal labels.

Quick check

1. In NH₄⁺ + OH⁻ → NH₃ + H₂O, what is the conjugate base of NH₄⁺? Answer: NH₃, formed when ammonium donates one proton.

Exam focus

Show the transferred proton explicitly and identify both conjugate pairs. Include solvent species if they accept or donate the proton. Do not assume an aqueous pH ranking applies unchanged in liquid ammonia or a gas reaction.

Advanced insight

The leveling effect means a solvent converts acids stronger than its conjugate acid into the same protonated solvent species. This limits how well acid strengths can be distinguished within that medium. Comparing solvents can reveal differences hidden in water, but requires appropriate equilibrium and solvation data.

Summary

Brønsted–Lowry chemistry tracks proton donation and acceptance in any suitable medium. Conjugate pairs differ by one proton, and amphiprotic species switch roles with reaction partner. Solvent and coordination effects control equilibrium strength.

Practice questions

1. Identify the acid in HCl(g) + NH₃(g) → NH₄Cl(s). Answer: HCl donates a proton to NH₃, so HCl is the Brønsted acid. 2. Write the autoionisation of liquid ammonia. Answer: 2NH₃ ⇌ NH₄⁺ + NH₂⁻. 3. Can HCO₃⁻ act as a base? Give its conjugate acid. Answer: Yes; it can accept H⁺ to form H₂CO₃, represented in water by related dissolved CO₂/carbonic-acid species.