Brønsted Acid–Base Theory

Fundamental Idea of the Brønsted Concept

The Brønsted–Lowry acid–base theory defines acids and bases through the transfer of protons (hydrogen ions, $ \text{H}^+ $):

• A Brønsted acid is a proton donor.
• A Brønsted base is a proton acceptor.

This definition focuses on what particles do during a reaction, not on how they behave in isolation. It is especially useful because it:

• Works in many solvents, not just water.
• Treats acids and bases symmetrically (every acid has a related base and vice versa).
• Connects directly to proton transfer reactions, which are a major class of chemical reactions.

Proton Transfer and Conjugate Pairs

Proton Transfer as the Central Process

In any Brønsted acid–base reaction, a proton moves from an acid to a base. Symbolically:

$$
\text{Acid}_1 + \text{Base}_2 \rightleftharpoons \text{Base}_1 + \text{Acid}_2
$$

Here:

• $ \text{Acid}_1 $ donates $ \text{H}^+ $, becoming $ \text{Base}_1 $.
• $ \text{Base}_2 $ accepts $ \text{H}^+ $, becoming $ \text{Acid}_2 $.

This leads directly to the idea of conjugate acid–base pairs.

Conjugate Acid–Base Pairs

A conjugate acid–base pair consists of two species that differ by one proton:

• The acid form has one more $ \text{H}^+ $.
• The base form has one fewer $ \text{H}^+ $.

Examples:

• $ \text{HCl} / \text{Cl}^- $
  $ \text{HCl} $ (acid) $ \rightarrow $ donates $ \text{H}^+ $ to become $ \text{Cl}^- $ (conjugate base).
• $ \text{NH}_4^+ / \text{NH}_3 $
  $ \text{NH}_4^+ $ (acid) $ \rightarrow $ donates $ \text{H}^+ $ to become $ \text{NH}_3 $ (conjugate base).
• $ \text{H}_2\text{O} / \text{OH}^- $
  $ \text{H}_2\text{O} $ (acid) $ \rightarrow $ donates $ \text{H}^+ $ to become $ \text{OH}^- $ (conjugate base).
• $ \text{H}_2\text{CO}_3 / \text{HCO}_3^- / \text{CO}_3^{2-} $
  Stepwise:
• $ \text{H}_2\text{CO}_3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+ $
  $ \text{H}_2\text{CO}_3 $ and $ \text{HCO}_3^- $ are a conjugate pair.
• $ \text{HCO}_3^- \rightleftharpoons \text{CO}_3^{2-} + \text{H}^+ $
  $ \text{HCO}_3^- $ and $ \text{CO}_3^{2-} $ are another pair.

In any acid–base equation, you can identify two conjugate pairs, one on each “side” of the proton transfer.

Example in water:

$$
\text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^-
$$

Pairs:

• $ \text{HCl} $ / $ \text{Cl}^- $
• $ \text{H}_2\text{O} $ / $ \text{H}_3\text{O}^+ $

Brønsted Acids and Bases in Different Environments

In Aqueous Solution

In water, the hydronium ion $ \text{H}_3\text{O}^+ $ represents a proton attached to a water molecule. Protons do not exist “naked” in solution; they are always bound to a solvent.

Typical aqueous Brønsted acids:

• $ \text{HCl}, \text{HNO}_3, \text{H}_2\text{SO}_4 $ (strong acids, donate protons readily)
• $ \text{CH}_3\text{COOH} $ (acetic acid, weaker acid)

Typical aqueous Brønsted bases:

• $ \text{OH}^- $ (from many metal hydroxides)
• $ \text{NH}_3 $ (ammonia)
• $ \text{HCO}_3^- $ (bicarbonate)

Brønsted theory does not require water, but water is a common and important example.

In Non-Aqueous and General Solvents

Because the definition is not tied to water, many species that do not look “basic” in the simple sense can act as Brønsted bases by accepting a proton.

Examples:

• In liquid ammonia as solvent:
• $ \text{NH}_4^+ $ is an acid (can donate $ \text{H}^+ $).
• $ \text{NH}_2^- $ is a very strong base (can accept $ \text{H}^+ $).
• In acetic acid as solvent:
• $ \text{CH}_3\text{COOH} $ can be the “solvent base”.
• $ \text{CH}_3\text{COOH}_2^+ $ is then a conjugate acid.

The Brønsted concept thus allows proton transfer chemistry to be described in any medium where proton transfer is meaningful.

Ampholytes (Amphiprotic Substances)

A substance that can act both as a Brønsted acid and as a Brønsted base is called amphiprotic (or, more generally: ampholyte).

An amphiprotic species:

• Donates $ \text{H}^+ $ in the presence of a stronger base.
• Accepts $ \text{H}^+ $ in the presence of a stronger acid.

Key examples:

• Water, $ \text{H}_2\text{O} $:
• As an acid:
  $$
  \text{H}_2\text{O} + \text{NH}_3 \rightarrow \text{NH}_4^+ + \text{OH}^-
  $$
• As a base:
  $$
  \text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^-
  $$
• Hydrogen carbonate, $ \text{HCO}_3^- $:
• As an acid:
  $$
  \text{HCO}_3^- \rightleftharpoons \text{CO}_3^{2-} + \text{H}^+
  $$
• As a base:
  $$
  \text{HCO}_3^- + \text{H}^+ \rightleftharpoons \text{H}_2\text{CO}_3
  $$
• Dihydrogen phosphate, $ \text{H}_2\text{PO}_4^- $:
• As an acid:
  $$
  \text{H}_2\text{PO}_4^- \rightleftharpoons \text{HPO}_4^{2-} + \text{H}^+
  $$
• As a base:
  $$
  \text{H}_2\text{PO}_4^- + \text{H}^+ \rightleftharpoons \text{H}_3\text{PO}_4
  $$

Amphiprotic substances are central in buffer systems and many acid–base equilibria.

Polyprotic (Polyprotic) Brønsted Acids and Their Conjugate Bases

A polyprotic acid (also called polyprotic) can donate more than one proton per molecule:

• Diprotic acid: can donate 2 protons (e.g. $ \text{H}_2\text{SO}_4, \text{H}_2\text{CO}_3 $).
• Triprotic acid: can donate 3 protons (e.g. $ \text{H}_3\text{PO}_4 $).

Each deprotonation step gives a new conjugate base:

Example: Carbonic acid, $ \text{H}_2\text{CO}_3 $:

1. First proton donation:
   $$
   \text{H}_2\text{CO}_3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+
   $$
   Pair: $ \text{H}_2\text{CO}_3 / \text{HCO}_3^- $
2. Second proton donation:
   $$
   \text{HCO}_3^- \rightleftharpoons \text{CO}_3^{2-} + \text{H}^+
   $$
   Pair: $ \text{HCO}_3^- / \text{CO}_3^{2-} $

Each step is a separate Brønsted acid–base equilibrium, with its own conjugate pair.

In general, for a polyprotic acid $ \text{H}_n\text{A} $:

• First conjugate base: $ \text{H}_{n-1}\text{A}^- $
• Then $ \text{H}_{n-2}\text{A}^{2-} $, etc., down to $ \text{A}^{n-} $.

Relative Strength: Stronger Acids and Stronger Bases

Acid Strength and Conjugate Base Strength

Within Brønsted theory, the strength of an acid (how readily it donates a proton) is closely related to the strength of its conjugate base (how readily that base would re-accept the proton):

• A strong acid has a weak conjugate base.
• A weak acid has a stronger conjugate base (relative to the conjugate base of a strong acid).

Qualitatively:

• If acid HA donates a proton almost completely to a given base (e.g. water), HA is called strong in that medium.

Examples in water:

• Strong acids (almost complete proton donation to water):
• $ \text{HCl}, \text{HBr}, \text{HI}, \text{HNO}_3, \text{HClO}_4 $
• Their conjugate bases ($ \text{Cl}^-, \text{Br}^-, \text{I}^-, \text{NO}_3^-, \text{ClO}_4^- $) are very weak bases.
• Weak acids (only partial proton donation):
• $ \text{CH}_3\text{COOH}, \text{H}_2\text{CO}_3, \text{H}_2\text{S}, \text{HF} $
• Their conjugate bases ($ \text{CH}_3\text{COO}^-, \text{HCO}_3^-, \text{HS}^-, \text{F}^- $) are relatively stronger bases.

Similarly, for bases in water:

• Strong base: $ \text{OH}^- $ (conjugate base of the very weak acid $ \text{H}_2\text{O} $).
• Weaker bases: $ \text{NH}_3, \text{CH}_3\text{COO}^-, \text{F}^- $, etc.

The detailed, quantitative treatment of acid and base strength (e.g. using equilibrium constants and $pK_\text{a}$ values) belongs to acid–base equilibrium discussions, but the qualitative idea already appears directly from the Brønsted concept.

Direction of Proton Transfer

Proton transfer “prefers” the side where the weaker acid and weaker base are present. In many simple reactions in water:

• A stronger acid will donate a proton to water, forming $ \text{H}_3\text{O}^+ $.
• A stronger base will take a proton from water, forming $ \text{OH}^- $.

Thus, identifying relative strengths of acids and their conjugate bases helps predict the likely direction of a Brønsted acid–base reaction.

Writing and Analyzing Brønsted Acid–Base Equations

Identifying Acids and Bases in a Reaction

When given an equation:

1. Locate the proton(s) that move: what species lose or gain $ \text{H}^+ $?
2. The species losing $ \text{H}^+ $ is the Brønsted acid; the product that results is its conjugate base.
3. The species gaining $ \text{H}^+ $ is the Brønsted base; the product that results is its conjugate acid.

Example:

$$
\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-
$$

• $ \text{NH}_3 $ gains $ \text{H}^+ \rightarrow \text{NH}_4^+ $:
  $ \text{NH}_3 $ = base, $ \text{NH}_4^+ $ = conjugate acid.
• $ \text{H}_2\text{O} $ loses $ \text{H}^+ \rightarrow \text{OH}^- $:
  $ \text{H}_2\text{O} $ = acid, $ \text{OH}^- $ = conjugate base.

Conjugate pairs:
$ \text{NH}_3 / \text{NH}_4^+ $ and $ \text{H}_2\text{O} / \text{OH}^- $.

Net Proton Transfer View

Even when an equation is given in ionic form, the same procedure applies.

Example:

$$
\text{HS}^- + \text{HCO}_3^- \rightleftharpoons \text{H}_2\text{S} + \text{CO}_3^{2-}
$$

Track protons:

• $ \text{HS}^- $ gains $ \text{H}^+ $ to form $ \text{H}_2\text{S} $
  $ \Rightarrow \text{HS}^- $ is base, $ \text{H}_2\text{S} $ is conjugate acid.
• $ \text{HCO}_3^- $ loses $ \text{H}^+ $ to form $ \text{CO}_3^{2-} $
  $ \Rightarrow \text{HCO}_3^- $ is acid, $ \text{CO}_3^{2-} $ is conjugate base.

Autoionization (Self-Protolysis) as a Brønsted Process

Some pure substances undergo a self-proton-transfer reaction, in which one molecule acts as a Brønsted acid and another molecule of the same substance acts as a Brønsted base. This is called autoionization or self-protolysis.

For water:

$$
2 \, \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-
$$

• One $ \text{H}_2\text{O} $ donates $ \text{H}^+ $ (acid) and becomes $ \text{OH}^- $ (conjugate base).
• Another $ \text{H}_2\text{O} $ accepts $ \text{H}^+ $ (base) and becomes $ \text{H}_3\text{O}^+ $ (conjugate acid).

Similar self-protolysis reactions exist for other amphiprotic solvents (e.g. ammonia, acetic acid), always describable in Brønsted terms.

Comparison with Other Acid–Base Concepts (Qualitative)

The Brønsted concept:

• Focuses specifically on proton transfer.
• Requires that an acid possess at least one removable $ \text{H}^+ $.
• Requires that a base have a site that can accept a proton (often a lone pair of electrons, but the concept itself does not explicitly mention electrons).

It is more general than the simple “Arrhenius” definitions tied to $ \text{H}_3\text{O}^+ $ and $ \text{OH}^- $, yet more specific than theories that interpret acids and bases in terms of electron pairs or more abstract criteria.

The broader, more electron-focused description of acids and bases is handled by other concepts treated elsewhere, while the core of the Brønsted theory remains:

Acid–base reactions are proton transfer reactions between conjugate acid–base pairs.