Conjugate acid

Acids and bases
Acid Acid–base reaction Acid–base homeostasis Acid strength Acidity function Amphoterism Base Buffer solutions Dissociation constant Equilibrium chemistry Extraction Hammett acidity function pH Proton affinity Self-ionization of water Titration Lewis acid catalysis Frustrated Lewis pair Chiral Lewis acid
Acid types
Brønsted–Lowry Lewis Acceptor Mineral Organic Oxide Strong Superacids Weak Solid
Base types
Brønsted–Lowry Lewis Donor Organic Oxide Strong Superbases Non-nucleophilic Weak
v t

A conjugate acid, within the Brønsted–Lowry acid–base theory, is a chemical compound formed when an acid donates a proton (H⁺) to a base—in other words, it is a base with a hydrogen ion added to it, as in the reverse reaction it loses a hydrogen ion. On the other hand, a conjugate base is what is left over after an acid has donated a proton during a chemical reaction. Hence, a conjugate base is a species formed by the removal of a proton from an acid, as in the reverse reaction it is able to gain a hydrogen ion.^[1] Because some acids are capable of releasing multiple protons, the conjugate base of an acid may itself be acidic.

In summary, this can be represented as the following chemical reaction:

Acid + Base ⇌ Conjugate Base + Conjugate Acid

Johannes Nicolaus Brønsted and Martin Lowry introduced the Brønsted–Lowry theory, which proposed that any compound that can transfer a proton to any other compound is an acid, and the compound that accepts the proton is a base. A proton is a nuclear particle with a unit positive electrical charge; it is represented by the symbol H⁺ because it constitutes the nucleus of a hydrogen atom,^[2] that is, a hydrogen cation.

A cation can be a conjugate acid, and an anion can be a conjugate base, depending on which substance is involved and which acid–base theory is the viewpoint. The simplest anion which can be a conjugate base is the solvated electron whose conjugate acid is the atomic hydrogen.

Acid–base reactions[]

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

Conjugates are formed when an acid loses a hydrogen proton or a base gains a hydrogen proton. Refer to the following figure:

We say that the water molecule is the conjugate acid of the hydroxide ion after the latter received the hydrogen proton donated by ammonium. On the other hand, ammonia is the conjugate base for the acid ammonium after ammonium has donated a hydrogen ion towards the production of the water molecule. We can also refer to OH- as a conjugate base of H
₂O, since the water molecule donates a proton towards the production of NH⁺
₄ in the reverse reaction, which is the predominating process in nature due to the strength of the base NH
₃ over the hydroxide ion. Based on this information, it is clear that the terms "Acid", "Base", "conjugate acid", and "conjugate base" are not fixed for a certain chemical species; but are interchangeable according to the reaction taking place.

Strength of conjugates[]

The strength of a conjugate acid is directly proportional to its dissociation constant. If a conjugate acid is strong, its dissociation will have a higher equilibrium constant and the products of the reaction will be favored. The strength of a conjugate base can be seen as the tendency of the species to "pull" hydrogen protons towards itself. If a conjugate base is classified as strong, it will "hold on" to the hydrogen proton when in solution and its acid will not dissociate.

If a species is classified as a strong acid, its conjugate base will be weak.^[3] An example of this case would be the dissociation of hydrochloric acid HCl in water. Since HCl is a strong acid (it dissociates to a great extent), its conjugate base (Cl⁻
) will be a weak conjugate base. Therefore, in this system, most H⁺
will be in the form of a hydronium ion H
₃O⁺
instead of attached to a Cl⁻ anion and the conjugate base will be weaker than a water molecule.

On the other hand, if a species is classified as a weak acid its conjugate base will not necessarily be a strong base. Consider that acetate, the conjugate base of acetic acid, has a base dissociation constant (Kb) of approximately 5.6×10⁻¹⁰, making it a weak base. In order for a species to have a strong conjugate base it has to be a very weak acid, like water for example.

Identifying conjugate acid–base pairs[]

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 the water molecule and its conjugate base is nitrate (NO⁻
₃). The water molecule acts as a base because it receives the hydrogen cation (proton) and its conjugate acid is the hydronium ion (H
₃O⁺
).

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

Applications[]

One use of conjugate acids and bases lies in buffering systems, which include a buffer solution. In a buffer, a weak acid and its conjugate base (in the form of a salt), or a weak base and its conjugate acid, are used in order to limit the pH change during a titration process. Buffers have both organic and non-organic chemical applications. For example, besides buffers being used in lab processes, human blood acts as a buffer to maintain pH. The most important buffer in our bloodstream is the carbonic acid-bicarbonate buffer, which prevents drastic pH changes when CO
₂ is introduced. This functions as such:

${\displaystyle {\ce {CO2 + H2O <=> H2CO3 <=> HCO3^- + H+}}}$

Furthermore, here is a table of common buffers.

Buffering agent	pKa	Useful pH range
Citric acid	3.13, 4.76, 6.40	2.1 - 7.4
Acetic acid	4.8	3.8 - 5.8
KH2PO4	7.2	6.2 - 8.2
CHES	9.3	8.3–10.3
Borate	9.24	8.25 - 10.25

A second common application with an organic compound would be the production of a buffer with acetic acid. If acetic acid, a weak acid with the formula CH
₃COOH, was made into a buffer solution, it would need to be combined with its conjugate base CH
₃COO⁻
in the form of a salt. The resulting mixture is called an acetate buffer, consisting of aqueous CH
₃COOH and aqueous CH
₃COONa. Acetic acid, along with many other weak acids, serve as useful components of buffers in different lab settings, each useful within their own pH range.

Ringer's lactate solution is an example where the conjugate base of an organic acid, lactic acid, CH
₃CH(OH)CO⁻
₂ is combined with sodium, calcium and potassium cations and chloride anions in distilled water^[4] which together form a fluid which is isotonic in relation to human blood and is used for fluid resuscitation after blood loss due to trauma, surgery, or a burn injury.^[5]

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 conjugate base strength increases down the table.

Acid	Conjugate base
H 2F+ Fluoronium ion	HF Hydrogen fluoride
HCl Hydrochloric acid	Cl− Chloride ion
H2SO4 Sulfuric acid	HSO− 4 Hydrogen sulfate ion
HNO3 Nitric acid	NO− 3 Nitrate ion
H3O+ Hydronium ion	H2O Water
HSO− 4 Hydrogen sulfate ion	SO2− 4 Sulfate ion
H3PO4 Phosphoric acid	H2PO− 4 Dihydrogen phosphate ion
CH3COOH Acetic acid	CH3COO− Acetate ion
HF Hydrofluoric acid	F− Fluoride ion
H2CO3 Carbonic acid	HCO− 3 Hydrogen carbonate ion
H2S Hydrosulfuric acid	HS− Hydrogen sulfide ion
H2PO− 4 Dihydrogen phosphate ion	HPO2− 4 Hydrogen phosphate ion
NH+ 4 Ammonium ion	NH3 Ammonia
H2O Water (pH=7)	OH− Hydroxide ion
HCO− 3 Hydrogencarbonate (bicarbonate) ion	CO2− 3 Carbonate ion

Table of bases and their conjugate acids[]

In contrast, here is a table of bases and their conjugate acids. Similarly, base strength decreases and conjugate acid strength increases down the table.

Base	Conjugate acid
C 2H 5NH 2 Ethylamine	C 2H 5NH+ 3 Ethylammonium ion
CH 3NH 2 Methylamine	CH 3NH+ 3 Methylammonium ion
NH 3 Ammonia	NH+ 4 Ammonium ion
C 5H 5N Pyridine	C 5H 6N+ Pyridinium
C 6H 5NH 2 Aniline	C 6H 5NH+ 3 Phenylammonium ion
C 6H 5CO− 2 Benzoate ion	C 6H 6CO 2 Benzoic acid
F− Fluoride ion	HF Hydrogen fluoride
PO3− 4 Phosphate ion	HPO2− 4 Hydrogen phosphate ion
OH− Hydroxide ion	H2O Water (neutral, pH 7)
HCO− 3 Bicarbonate	H 2CO 3 Carbonic acid
CO2− 3 Carbonate ion	HCO− 3 Bicarbonate
Br− Bromine	HBr Hydrogen bromide
HPO2− 4 Hydrogen Phosphate	H 2PO− 4
Cl− Chloride ion	HCl Hydrogen chloride
H 2O Water	H 3O+ Hydronium ion

See also[]

• Buffer solution
• Deprotonation
• Protonation
• Salt (chemistry)

References[]

External links[]

• MCAT General Chemistry Review - 10.4 Titration and Buffers
• The Pharmaceutics and Compounding Laboratory - Buffers and Buffer Capacity.