Conjugate acid

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Acids and bases

Acid dissociation constant Acid-base extraction Acid–base reaction Acid–base titration Dissociation constant Acidity function Buffer solutions pH Proton affinity Amphoterism Self-ionization of water Acid strength
Acid types
Brønsted Lewis Mineral Organic Strong Superacids Weak
Base types
Brønsted Lewis Organic Strong Superbases Non-nucleophilic Weak
v t e

A conjugate acid, within the Brønsted–Lowry theory, is a species formed by the reception of a proton (H⁺), by a base - in other words, the base with a hydrogen ion added to it - while a conjugate base is formed by the removal of a proton from an acid: the conjugate base of an acid is that acid with a hydrogen ion removed.^[1]

Visually, this can be represented as:

A conjugate acid of the base → Base + H+

A conjugate base of the acid → Acid − H+

The Brønsted-Lowry model is based on the idea that acids are proton donors and bases are proton acceptors; the conjugate base or conjugate acid is merely what is left after an acid has lost a proton or a base has gained a proton, respectively.

Acid-base reactions

In an acid-base reaction, an acid plus a base reacts to form a conjugate base plus a conjugate acid:

acid + base

conjugate base + conjugate acid

The conjugate acid of a base is formed when the base gains a proton. Refer to the following equation:

NH
3(g) + H
2O(l)

NH+
4(aq) + OH−
(aq)

We say that ammonium (NH+
4) is the conjugate acid to the base ammonia (NH
3), because NH
3 gained a hydrogen ion to form NH+
4, the conjugate acid. The conjugate base of an acid is formed when the acid donates a proton. In the equation, we say that hydroxide (OH−
) is the conjugate base to the acid water (H
2O), because H
2O donates a hydrogen ion to form OH−
, the conjugate base.

Strength of conjugates

The stronger the acid or base, the weaker the conjugate. The weaker the acid or base, the stronger the conjugate. However, a weak acid or base will not necessarily have a strong conjugate base or acid; there are a number of pairs of weak conjugates. For example, acetic acid (CH₃COOH) and the acetate ion (CH₃COO^-) are both weak.

This principle is discussed in detail in the article on acid-base reaction theories. The position of the equilibrium is measured as an acid dissociation constant (K_a)

Identifying conjugate acid-base pairs

Example of a Bronsted-Lowry Reaction

The acid and conjugate base as well as the base and conjugate acid are conjugate pairs. When finding a conjugate acid or base, it is important to look at the reactants. The reactants are the acids and bases, and the acid corresponds to the conjugate base on the product side of the chemical equation. This goes for the base too; the base corresponds to the conjugate acid on the product side of the equation.

To identify the conjugate acid, look for the pair of compounds that are related. The acid-base reaction can be viewed in a before and after sense. The before is the reactant side of the equation, the after is the product side of the equation. The conjugate acid in the after side of an equation gains a hydrogen ion, so in the before side of the equation the compound that has one less hydrogen ion of the conjugate acid is the base. The conjugate base in the after side of the equation lost a hydrogen ion, so in the before side of the equation the compound that has one more hydrogen ion of the conjugate base is the acid.

Consider the following acid-base reaction:

HNO
3 + H
2O → H
3O+
+ NO−
3

Nitric acid (HNO
3) is an acid because it donates a proton to water and its conjugate base is nitrate (NO−
3). An easy way to identify the conjugate base is that it differs from the acid by one proton. Water (H
2O) is a base because it accepts a proton from HNO
3 and its conjugate acid is hydronium (H
3O+
). Again to identify the conjugate acid (or any conjugate pair) is that it differs from the base by one proton.

Equation	Acid	Base	Conjugate Base	Conjugate Acid
HClO 2 + H 2O → ClO− 2 + H 3O+	HClO 2	H 2O	ClO− 2	H 3O+
ClO− + H 2O → HClO + OH−	H 2O	ClO−	OH−	HClO
HCl + H 2PO− 4 → Cl− + H 3PO 4	HCl	H 2PO− 4	Cl−	H 3PO 4

Table of Acids and their Conjugate Bases

Tabulated below are several examples of acids and their conjugate bases; notice how they differ by just one proton (H⁺ ion). Acid strength decreases and base strength increases down the table.

Acid	Base
HFSbF₅ Fluoroantimonic acid	SbF₆⁻ Hexafluoroantimonate ion
HCl Hydrochloric acid	Cl⁻ Chloride ion
H₂SO₄ Sulfuric acid	HSO₄⁻ Hydrogen sulfate ion
HNO₃ Nitric acid	NO₃⁻ Nitrate ion
H₃O⁺ Hydronium ion	H₂O Water
HSO₄⁻ Hydrogen sulfate ion	SO₄²⁻ Sulfate ion
H₃PO₄ Phosphoric acid	H₂PO₄⁻ Dihydrogen phosphate ion
CH₃COOH Acetic acid	CH₃COO⁻ Acetate ion
H₂CO₃ Carbonic acid	HCO₃⁻ Hydrogen carbonate ion
H₂S Hydrosulfuric acid	HS⁻ Hydrogen sulfide ion
H₂PO₄⁻ Dihydrogen phosphate ion	HPO₄²⁻ Hydrogen phosphate ion
NH₄⁺ Ammonium ion	NH₃ Ammonia
HCO₃⁻ Hydrogencarbonate (bicarbonate) ion	CO₃²⁻ Carbonate ion
HPO₄²⁻ Hydrogen phosphate ion	PO₄³⁻ Phosphate ion
H₂O Water (neutral, pH 7)	OH⁻ Hydroxide ion

References

↑ Zumdahl, Stephen S., & Zumdahl, Susan A. Chemistry. Houghton Mifflin, 2007, ISBN 0618713700

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