Editing Conjugate (acid-base theory)

{{Short description|Chemical compound formed when an acid donates a proton to a base}}
{{use dmy dates |date=April 2021}}
{{Acids and bases}}

A '''conjugate acid''', within the [[Brønsted–Lowry acid–base theory]], is a [[chemical compound]] formed when an acid [[protonation|gives a proton]] ({{chem2|[[Hydron (chemistry)|H+]]}}) to a [[base (chemistry)|base]]—in other words, it is a base with a [[hydrogen  ion]] added to it, as it loses a hydrogen ion in the reverse reaction. On the other hand, a '''conjugate base''' is what remains after an acid has donated a proton during a chemical reaction. Hence, a conjugate base is a substance formed by the [[deprotonation|removal of a proton]] from an acid, as it can gain a hydrogen ion in the reverse reaction. <ref>Zumdahl, Stephen S., & Zumdahl, Susan A. ''Chemistry''. Houghton Mifflin, 2007, {{ISBN|0618713700}}</ref> Because [[Polyprotic acid|some acids]] can give multiple protons, the conjugate base of an acid may itself be acidic.

In summary, this can be represented as the following [[chemical reaction]]:
<math chem display="block">\text{acid} + \text{base} \; \ce{<=>} \; \text{conjugate base} + \text{conjugate acid}</math>

{{multiple image|left|
 | width = 100
 | footer = [[Johannes Nicolaus Brønsted]] (left) and [[Martin Lowry]] (right).
 | image1 = Johannes Brønsted.jpg
 | image2 =Thomas Martin Lowry2.jpg
 }}
[[Johannes Nicolaus Brønsted]] and [[Martin Lowry]] introduced the Brønsted–Lowry theory, which said that any compound that can give a proton to another compound is an acid, and the compound that receives the proton is a base. A proton is a subatomic particle in the nucleus with a unit positive electrical charge. It is represented by the symbol {{chem2|H+}} because it has the [[atomic nucleus|nucleus]] of a hydrogen [[atom]],<ref>{{Cite web|url=https://www.britannica.com/science/Bronsted-Lowry-theory|title=Brønsted–Lowry theory {{!}} chemistry|website=Encyclopedia Britannica|language=en|access-date=2020-02-25}}</ref> that is, a [[hydron (chemistry)|hydrogen cation]].

A [[cation]] can be a conjugate acid, and an [[anion]] can be a conjugate base, depending on which [[chemical substance|substance]] is involved and which [[acid–base reaction|acid–base theory]] is used. The simplest anion which can be a conjugate base is the [[solvated electron|free electron in a solution]] whose conjugate acid is the atomic hydrogen.

==Acid–base reactions==
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<sup>−</sup> 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}}

==Strength of conjugates==
The strength of a conjugate acid is proportional to its [[dissociation constant|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.<ref name="Conjugate Strength">{{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}}</ref> 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<sup>−</sup> 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 dissociation constant|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}}

==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:

{{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|+}}).

{| class="wikitable" style="text-align:center"
|-
! Equation !! Acid !! Base !! Conjugate base !! Conjugate acid
|-
| {{chem|HClO|2}} + {{chem|H|2|O}} → {{chem|ClO|2|-}} + {{chem|H|3|O|+}} || [[chlorous acid|{{chem|HClO|2}}]] || [[Properties of water|{{chem|H|2|O}}]] || [[chlorite|{{chem|ClO|2|-}}]] || [[hydronium|{{chem|H|3|O|+}}]]
|-
| {{chem|ClO|-}} + {{chem|H|2|O}} → {{chem|HClO}} + {{chem|OH|-}} || {{chem|H|2|O}} || [[hypochlorite|{{chem|ClO|-}}]] || [[hydroxide|{{chem|OH|-}}]] || [[Hypochlorous acid|{{chem|HClO}}]]
|-
| {{chem|HCl}} + {{chem|H|2|PO|4|-}} → {{chem|Cl|-}} + {{chem|H|3|PO|4}} || [[hydrochloric acid|{{chem|HCl}}]] || [[phosphate#Chemical properties|{{chem|H|2|PO|4|-}}]] || [[chloride|{{chem|Cl|-}}]] || [[phosphoric acid|{{chem|H|3|PO|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 [[Bicarbonate buffering system|carbonic acid-bicarbonate buffer]], which prevents drastic pH changes when {{Chem|CO|2}} is introduced. This functions as such:{{cn|date=May 2025}}
<chem display="block">CO2 + H2O <=> H2CO3 <=> HCO3^- + H+</chem>

Furthermore, here is a table of common buffers.
{| class="wikitable" style="text-align:center"
!Buffering agent!!pK<sub>a</sub>!!Useful pH range
|-
|[[Citric acid]]||3.13, 4.76, 6.40||2.1 - 7.4
|-
|[[Acetic acid]]||4.8||3.8 - 5.8
|-
|[[potassium dihydrogenphosphate|KH<sub>2</sub>PO<sub>4</sub>]]||7.2|| 6.2 - 8.2
|-
|[[N-Cyclohexyl-2-aminoethanesulfonic acid|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 {{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<ref name=BNF69>{{cite book|title=British national formulary: BNF 69|date=2015|publisher=British Medical Association|isbn=9780857111562|page=683|edition=69}}</ref> which together form a fluid which is [[Tonicity#Isotonicity|isotonic]] in relation to human blood and is used for [[Fluid replacement|fluid resuscitation]] after [[blood loss]] due to [[Physical trauma|trauma]], [[surgery]], or a [[burn injury]].<ref name="PestanaFifth">{{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 }}</ref>

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

{| class="wikitable"
|-
! Acid
! Conjugate base
|-
| {{Chem|H|2|F|+}} [[Fluoronium]] ion
| HF [[Hydrogen fluoride]]
|-
| HCl [[Hydrochloric acid]]
| Cl<sup>−</sup> [[Chloride]] ion
|-
| H<sub>2</sub>SO<sub>4</sub> [[Sulfuric acid]]
| HSO{{su|b=4|p=−}} [[Hydrogen sulfate]] ion (''bisulfate'' ion)
|-
| HNO<sub>3</sub> [[Nitric acid]]
| NO{{su|b=3|p=−}} [[Nitrate]] ion
|-
| H<sub>3</sub>O<sup>+</sup> [[Hydronium]] ion
| H<sub>2</sub>O [[water (molecule)|Water]]
|-
| HSO{{su|b=4|p=−}} [[Hydrogen sulfate]] ion
| SO{{su|b=4|p=2−}} [[Sulfate]] ion
|-
| H<sub>3</sub>PO<sub>4</sub> [[Phosphoric acid]]
| H<sub>2</sub>PO{{su|b=4|p=−}} [[Dihydrogen phosphate]] ion
|-
| CH<sub>3</sub>COOH [[Acetic acid]]
| CH<sub>3</sub>COO<sup>−</sup> [[Acetate]] ion
|-
| HF [[Hydrofluoric acid]]
| F<sup>−</sup> [[Fluoride]] ion
|-
| H<sub>2</sub>CO<sub>3</sub> [[Carbonic acid]]
| HCO{{su|b=3|p=−}} [[Hydrogen carbonate]] ion
|-
| H<sub>2</sub>S [[Hydrosulfuric acid]]
| HS<sup>−</sup> [[Hydrosulfide]] ion
|-
| H<sub>2</sub>PO{{su|b=4|p=−}} [[Dihydrogen phosphate]] ion
| HPO{{su|b=4|p=2−}} [[Monohydrogen phosphate|Hydrogen phosphate]] ion
|-
| NH{{su|b=4|p=+}} [[Ammonium]] ion
| NH<sub>3</sub> [[Ammonia]]
|-
| H<sub>2</sub>O Water ([[pH]]=7)
| OH<sup>−</sup> [[Hydroxide]] ion
|-
| HCO{{su|b=3|p=−}} [[Hydrogencarbonate]] ''(bicarbonate)'' ion
| CO{{su|b=3|p=2−}} [[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.
{| class="wikitable"
|-
! Base
! Conjugate acid
|-
| {{Chem|C|2|H|5|N|H|2}} [[Ethylamine]]
| {{Chem|C|2|H|5|N|H|3|+}} [[Ethylammonium]] ion
|-
| {{Chem|C|H|3|N|H|2}} [[Methylamine]]
| {{Chem|C|H|3|N|H|3|+}} [[Methylammonium]] ion
|-
| {{Chem|N|H|3|}} [[Ammonia]]
| {{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−}} [[Monohydrogen phosphate|Hydrogen phosphate]] ion
|-
| OH<sup>−</sup> [[Hydroxide]] ion
| H<sub>2</sub>O [[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-}} [[Monohydrogen phosphate|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]]
|}

==See also==
* [[Buffer solution]]
* [[Deprotonation]]
* [[Protonation]]
* [[Salt (chemistry)]]
* [[Carboxylate]]

==References==
{{reflist}}

==External links==
* [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]]