Bronsted-Lowry and Lewis Acid-Base Concepts

Understanding Acid‑Base Theories: Brønsted‑Lowry and Lewis Concepts

Acid‑base chemistry is a cornerstone of modern chemistry, influencing everything from biological systems to industrial processes. This course explores the two most widely taught definitions – the Brønsted‑Lowry and Lewis concepts – and applies them to common reactions, equilibrium constants, and special species such as amphiprotic molecules.

1. The Brønsted‑Lowry Definition of Acids and Bases

The Brønsted‑Lowry theory, introduced in the early 20th century, defines:

• Acids as proton (H⁺) donors.
• Bases as proton acceptors.

In practice, a base acts like a “proton sponge,” seeking out H⁺ ions to form its conjugate acid. This simple definition explains many aqueous reactions and provides a clear way to identify conjugate pairs.

Key takeaway: The substance that accepts a proton (H⁺) is the Brønsted‑Lowry base.

2. Conjugate Acid–Base Pairs in Reversible Reactions

When an acid donates a proton, it becomes its conjugate base; the base that accepts the proton becomes its conjugate acid. Consider the classic equilibrium:

HCl + NH₃ ⇌ NH₄⁺ + Cl⁻

Here, HCl loses a proton to become Cl⁻, forming the HCl/Cl⁻ conjugate pair. Simultaneously, NH₃ gains a proton to become NH₄⁺, creating the NH₃/NH₄⁺ pair. Recognizing these pairs helps predict reaction direction and calculate equilibrium constants.

3. Lewis Acid‑Base Theory: Beyond Protons

Gilbert N. Lewis expanded the definition in 1923 to include reactions that involve the transfer of electron pairs rather than protons. According to Lewis:

• Lewis acids are electron‑pair acceptors.
• Lewis bases are electron‑pair donors.

Typical Lewis acids lack a full valence shell and possess an empty orbital ready to receive electrons. A classic example is BF₃, which has an empty p orbital that can accept a lone pair from a donor such as NH₃, forming a coordinate covalent bond.

Mnemonic: Think of BF₃ as an “electron‑hungry sponge” that eagerly takes in a pair of electrons.

4. Comparing Acid Strengths Using the Acid‑Dissociation Constant (Ka)

The acid‑dissociation constant, Ka, quantifies the extent to which an acid dissociates in water. Larger Ka values indicate stronger acids because a greater fraction of the acid molecules release H⁺ ions.

For example:

• Acetic acid: Ka = 1.8 × 10⁻⁵
• Hydrogen cyanide (HCN): Ka = 7.1 × 10⁻¹⁰

Since 1.8 × 10⁻⁵ > 7.1 × 10⁻¹⁰, acetic acid is considerably stronger than HCN. The larger Ka reflects a higher degree of ionisation, meaning more H⁺ is released at equilibrium.

Visual analogy: A larger faucet leak (higher Ka) pours more water (protons) into the sink (solution).

5. Deriving Ka from Concentration and Degree of Ionisation

For a monoprotic weak acid HA with an initial concentration c and degree of ionisation α (the fraction of molecules that dissociate), the equilibrium concentrations are:

• [H⁺] = cα
• [A⁻] = cα
• [HA] = c(1‑α)

Substituting these into the expression Ka = [H⁺][A⁻]/[HA] yields:

Ka = c·α²

This relationship is especially useful when α is small, allowing the approximation α ≈ √(Ka/c) for quick calculations.

6. Amphiprotic Species: Acting as Both Acid and Base

Some molecules can both donate and accept protons; they are termed amphiprotic. Water (H₂O) is the quintessential example:

• As a base: H₂O + H⁺ → H₃O⁺
• As an acid: H₂O → OH⁻ + H⁺

This dual behaviour is crucial in many buffer systems and in the self‑ionisation of water, which underpins the pH scale.

7. Applying the Lewis Concept to BF₃

Why does BF₃ act as a Lewis acid? The molecule has an incomplete octet on boron, leaving an empty p orbital. This vacant orbital can accept a lone pair from a Lewis base such as NH₃, forming the adduct BF₃·NH₃. The reaction does not involve proton transfer, highlighting the broader applicability of the Lewis definition.

Remember: An empty orbital → electron‑pair acceptor → Lewis acid.

8. Calculating the Degree of Dissociation (α) for Weak Acids

When the Ka value and initial concentration are known, the degree of dissociation can be estimated using the approximation:

α ≈ √(Ka / c)

For a 0.02 M solution of HF with Ka = 3.2 × 10⁻⁴:

• α ≈ √(3.2×10⁻⁴ / 0.02) ≈ √(0.016) ≈ 0.04

Thus, about 4 % of the HF molecules dissociate, providing a practical example of how equilibrium constants translate into observable solution behaviour.

9. Summary of Key Concepts

• Brønsted‑Lowry bases accept protons; acids donate them.
• Conjugate acid–base pairs are formed by the transfer of a single proton.
• Lewis acids accept electron pairs; Lewis bases donate them.
• Acid strength correlates directly with the magnitude of Ka.
• For a monoprotic weak acid, Ka = c·α², and α ≈ √(Ka/c) when α is small.
• Water is amphiprotic, capable of acting as both acid and base.
• BF₃ exemplifies a Lewis acid due to its empty orbital.
• Degree of dissociation calculations help predict the extent of ionisation in weak‑acid solutions.

Mastering these principles equips you to analyse a wide range of chemical reactions, design buffer systems, and understand the underlying thermodynamics of acid‑base equilibria.