Conjugate (acid-base theory)

In an acid–base reaction, an acid and a base react to form a conjugate base and a conjugate acid respectively. The acid loses a proton and the base gains a proton. In diagrams which indicate this, the new bond formed between the base and the proton is shown by an arrow that starts on an electron pair from the base and ends at the hydrogen ion (proton) that will be transferred:File:Conjugate base reaction.svg

In this case, the water molecule is the conjugate acid of the basic hydroxide ion after the latter received the hydrogen ion from ammonium. On the other hand, ammonia is the conjugate base for the acidic ammonium after ammonium has donated a hydrogen ion to produce the water molecule. Also, OH⁻ can be considered as the conjugate base of {{Chem|H|2|O}}, since the water molecule donates a proton to give {{Chem|NH|4|+}} in the reverse reaction. The terms "acid", "base", "conjugate acid", and "conjugate base" are not fixed for a certain chemical substance but can be swapped if the reaction taking place is reversed.{{cn|date=May 2025}}

The strength of a conjugate acid is proportional to its splitting constant. A stronger conjugate acid will split more easily into its products, "push" hydrogen protons away and have a higher equilibrium constant. The strength of a conjugate base can be seen as its tendency to "pull" hydrogen protons towards itself. If a conjugate base is classified as strong, it will "hold on" to the hydrogen proton when dissolved and its acid will not split.{{cn|date=May 2025}}

If a chemical is a strong acid, its conjugate base will be weak.{{Cite web| url=https://www.ausetute.com.au/cabstrength.html|title=Strength of Conjugate Acids and Bases Chemistry Tutorial| website=www.ausetute.com.au| access-date=2020-02-25}} An example of this case would be the splitting of hydrochloric acid {{Chem|H|Cl}} in water. Since {{Chem|H|Cl}} is a strong acid (it splits up to a large extent), its conjugate base ({{Chem|Cl|-}}) will be weak. Therefore, in this system, most {{Chem|H|+}} will be hydronium ions {{Chem|H|3|O|+}} instead of attached to Cl⁻ anions and the conjugate bases will be weaker than water molecules.{{cn|date=May 2025}}

On the other hand, if a chemical is a weak acid its conjugate base will not necessarily be strong. Consider that ethanoate, the conjugate base of ethanoic acid, has a base splitting constant (Kb) of about {{val|5.6e-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.{{cn|date=May 2025}}

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:

{{block indent | em = 1.5 | text = {{chem|HNO|3}} + {{chem|H|2|O}} → {{chem|H|3|O|+}} + {{chem|NO|3|-}}}}

Nitric acid ({{chem|HNO|3}}) is an acid because it donates a proton to the water molecule and its conjugate base is nitrate ({{chem|NO|3|-}}). The water molecule acts as a base because it receives the hydrogen cation (proton) and its conjugate acid is the hydronium ion ({{chem|H|3|O|+}}).

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 {{Chem|CO|2}} is introduced. This functions as such:{{cn|date=May 2025}}

CO2 + H2O <=> H2CO3 <=> HCO3^- + H+

Furthermore, here is a table of common buffers.

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 {{Chem|CH|3|COOH}}, was made into a buffer solution, it would need to be combined with its conjugate base {{Chem|CH|3|COO|-}} in the form of a salt. The resulting mixture is called an acetate buffer, consisting of aqueous {{Chem|CH|3|COOH}} and aqueous {{Chem|CH|3|COO|Na}}. 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.{{cn|date=May 2025}}

Ringer's lactate solution is an example where the conjugate base of an organic acid, lactic acid, {{chem|CH|3|CH(OH)CO|2|−}} is combined with sodium, calcium and potassium cations and chloride anions in distilled water{{cite book|title=British national formulary: BNF 69|date=2015|publisher=British Medical Association|isbn=9780857111562|page=683|edition=69}} 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.{{cite book |last1=Pestana |first1=Carlos |title=Pestana's Surgery Notes |date=7 April 2020 |publisher=Kaplan Medical Test Prep |isbn=978-1506254340 |pages=4–5 |edition=Fifth }}

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

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

| {{Chem|N|H|4|+}} Ammonium ion

|-

| {{Chem|C|5|H|5|N}} Pyridine

| {{Chem|C|5|H|6|N|+}} Pyridinium

|-

| {{Chem|C|6|H|5|N|H|2}} Aniline

| {{Chem|C|6|H|5|N|H|3|+}} Phenylammonium ion

|-

| {{Chem|C|6|H|5|C|O|2|-}} Benzoate ion

| {{Chem|C|6|H|6|C|O|2}} Benzoic acid

|-

| {{Chem|F|-}} Fluoride ion

| {{Chem|H|F}} Hydrogen fluoride

|-

| PO{{su|b=4|p=3−}} Phosphate ion

| HPO{{su|b=4|p=2−}} Hydrogen phosphate ion

|-

| OH⁻ Hydroxide ion

| H₂O Water (neutral, pH 7)

|-

| {{Chem|H|C|O|3|-}} Bicarbonate

| {{Chem|H|2|C|O|3}} Carbonic acid

|-

| {{Chem|C|O|3|2-}} Carbonate ion

| {{Chem|H|C|O|3|-}} Bicarbonate

|-

| {{Chem|Br|-}} Bromide ion

| {{Chem|H|Br}} Hydrogen bromide

|-

| {{Chem|H|P|O|4|2-}} Hydrogen phosphate

| {{Chem|H|2|P|O|4|-}} Dihydrogen phosphate ion

|-

| {{Chem|Cl|-}} Chloride ion

| {{Chem|H|Cl}} Hydrogen chloride

|-

| {{Chem|H|2|O}} Water

| {{Chem|H|3|O|+}} Hydronium ion

|-

| {{Chem|N|O|2|-}} Nitrite ion

| {{Chem|H|N|O|2}} Nitrous acid

|}

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

{{reflist}}

• [http://schoolbag.info/chemistry/mcat_2/78.html MCAT General Chemistry Review - 10.4 Titration and Buffers]
• [https://pharmlabs.unc.edu/labs/ophthalmics/buffers.htm The Pharmaceutics and Compounding Laboratory - Buffers and Buffer Capacity.] {{Webarchive|url=https://web.archive.org/web/20210428020925/https://pharmlabs.unc.edu/labs/ophthalmics/buffers.htm |date=28 April 2021 }}

Category:Acid–base chemistry