Lesson 8.1: Acid–Base Theories and Equilibrium: The Nature of Acids and Bases

8.1 The Nature of Acids and Bases

Introduction to Acids and Bases

Acids and bases are two fundamental classes of substances in chemistry, recognized for their distinct properties and behaviors. Acids typically taste sour (e.g., vinegar contains acetic acid, lemons contain citric acid), while bases are characterized by a bitter taste and slippery feel (e.g., substances used in drain cleaners). Over time, scientists have developed several theories to explain the nature of acids and bases, including the Arrhenius and Brønsted–Lowry theories.

The Arrhenius Theory of Acids and Bases

The Arrhenius theory was the first to provide a scientific explanation for acids and bases in aqueous solutions:

• Arrhenius Acid: A substance that produces hydrogen ions, H+(aq), when dissolved in water.

• Arrhenius Base: A substance that produces hydroxide ions, OH–(aq), when dissolved in water.

For example:

• Hydrochloric acid: $\mathrm{HCl}(g) \xrightarrow{\text{water}} \mathrm{H}^+(aq) + \mathrm{Cl}^-(aq)$

• Sodium hydroxide: $\mathrm{NaOH}(s) \xrightarrow{\text{water}} \mathrm{Na}^+(aq) + \mathrm{OH}^-(aq)$

The Arrhenius theory is limited because it only recognizes bases that contain hydroxide ions and only applies to aqueous solutions.

The Brønsted–Lowry Theory of Acids and Bases

The Brønsted–Lowry theory provides a broader definition:

• Brønsted–Lowry Acid: A proton (hydrogen ion) donor.

• Brønsted–Lowry Base: A proton (hydrogen ion) acceptor.

This theory includes all Arrhenius acids and bases, as well as substances like ammonia, NH3, which do not contain hydroxide ions but can still act as bases. (This theory is not limited to bases that only contain hydroxide ions, or limited to aqueous solutions , unlike the Arrhenius theory.)

The Brønsted–Lowry Theory in Action

Acidic Solutions

When hydrogen fluoride, HF, dissolves in water, the hydrogen ion does not exist freely but forms a hydronium ion, H3O+:

• $\mathrm{HF}(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{F}^-(aq) + \mathrm{H}_3\mathrm{O}^+(aq)$

Here, HF is the acid (proton donor), and H2O is the base (proton acceptor).

Basic Solutions

Ammonia reacts with water to produce ammonium and hydroxide ions:

• $\mathrm{NH}_3(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{NH}_4^+(aq) + \mathrm{OH}^-(aq)$

Water acts as an acid (proton donor), and ammonia acts as a base (proton acceptor).

Conjugate Acid–Base Pairs

Acid–base reactions are reversible, involving pairs of substances that differ by a single proton:

• Conjugate Acid: The species formed when a base gains a proton.

• Conjugate Base: The species formed when an acid loses a proton.

General equations:

• $\mathrm{HA}(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{A}^-(aq) + \mathrm{H}_3\mathrm{O}^+(aq)$

• $\mathrm{B}(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{HB}^+(aq) + \mathrm{OH}^-(aq)$

Each acid–base reaction involves two conjugate acid–base pairs, differing only by a hydrogen ion.

The Brønsted–Lowry Theory and Non-aqueous Reactions

The Brønsted–Lowry theory is not limited to aqueous solutions. For example, gaseous hydrogen chloride reacts with gaseous ammonia to form solid ammonium chloride:

• $\mathrm{HCl}(g) + \mathrm{NH}_3(g) \rightarrow \mathrm{NH}_4\mathrm{Cl}(s)$

Here, HCl is the acid, NH3 is the base, NH4+ is the conjugate acid, and Cl– is the conjugate base.

Amphiprotic (Amphoteric) Substances

Some substances can act as either acids or bases depending on the reaction. These are called amphiprotic (or amphoteric) substances. For example, water can act as a base in one reaction and as an acid in another. The hydrogen carbonate ion, HCO3–, is also amphiprotic:

• $\mathrm{HCO}_3^-(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{H}_2\mathrm{CO}_3(aq) + \mathrm{OH}^-(aq)$

• $\mathrm{HCO}_3^-(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{CO}_3^{2-}(aq) + \mathrm{H}_3\mathrm{O}^+(aq)$

The Acid Ionization Constant, Ka

The acid ionization constant, Ka, quantifies the extent to which an acid ionizes in water. For a generic weak acid, HA:

• $\mathrm{HA}(aq) + \mathrm{H}_2\mathrm{O}(l) \rightleftharpoons \mathrm{H}_3\mathrm{O}^+(aq) + \mathrm{A}^-(aq)$

The equilibrium expression is:

• $K_a = \frac{[\mathrm{H}_3\mathrm{O}^+][\mathrm{A}^-]}{[\mathrm{HA}]}$

Alternatively, since $[\mathrm{H}^+]$ is equivalent to $[\mathrm{H}_3\mathrm{O}^+]$:

• $K_a = \frac{[\mathrm{H}^+][\mathrm{A}^-]}{[\mathrm{HA}]}$

Values of Ka indicate acid strength; larger values mean stronger acids.

Acid	Acid Ionization Constant, Ka
Hydrocyanic, HCN(aq)	6.2 × 10–10
Benzoic, HC6H5CO2(aq)	6.3 × 10–5
Propanoic, HC3H5O2(aq)	1.3 × 10–5
Ethanoic (acetic), HC2H3O2(aq)	1.8 × 10–5
Hydrofluoric, HF(aq)	6.6 × 10–4
Nitrous, HNO2(aq)	4.6 × 10–4
Methanoic (formic), HCHO2(aq)	1.8 × 10–4

Summary of Key Concepts

• Arrhenius acids produce H+ in water; Arrhenius bases produce OH– in water.

• Brønsted–Lowry acids are proton donors; Brønsted–Lowry bases are proton acceptors.

  • *Think - Lowry = dowry = relates to the donation and acceptance of H+ ions

• Acid–base reactions involve conjugate acid–base pairs, differing by one proton.

• Amphiprotic substances can act as either acids or bases depending on the reaction.

• The acid ionization constant, Ka, measures the strength of an acid in water.