Conjugate (acid-base theory)

Aconjugate acid,within theBrønsted–Lowry acid–base theory,is achemical compoundformed when an acidgives a proton(H⁺) to abase—in other words, it is a base with ahydrogen ionadded to it, as it loses a hydrogen ion in the reverse reaction. On the other hand, aconjugate baseis what remains after an acid has donated a proton during a chemical reaction. Hence, a conjugate base is a substance formed by theremoval of a protonfrom an acid, as it can gain a hydrogen ion in the reverse reaction.^[1]Becausesome acidscan give multiple protons, the conjugate base of an acid may itself be acidic.

In summary, this can be represented as the followingchemical reaction: ${\text{acid}}+{\text{base}}\;{\ce {<=>}}\;{\text{conjugate base}}+{\text{conjugate acid}}$

Johannes Nicolaus Brønsted(left) andMartin Lowry(right).

Johannes Nicolaus BrønstedandMartin Lowryintroduced 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 symbolH⁺because it has thenucleusof a hydrogenatom,^[2]that is, ahydrogen cation.

Acationcan be a conjugate acid, and ananioncan be a conjugate base, depending on whichsubstanceis involved and whichacid–base theoryis used. The simplest anion which can be a conjugate base is thefree electron in a solutionwhose conjugate acid is the atomic hydrogen.

Acid–base reactions

In anacid–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 anelectron pairfrom the base and ends at the hydrogen ion (proton) that will be transferred: In this case, the water molecule is the conjugate acid of the basic hydroxide ion after the latter received the hydrogen ion fromammonium.On the other hand,ammoniais 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 ofH
₂O,since the water molecule donates a proton to giveNH⁺
₄in 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.

Strength of conjugates

The strength of a conjugate acid is proportional to itssplitting constant.A stronger conjugate acid will split more easily into its products, "push" hydrogen protons away and have a higherequilibrium 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.

If a chemical is a strong acid, its conjugate base will be weak.^[3]An example of this case would be the splitting ofhydrochloric acidHClin water. SinceHClis a strong acid (it splits up to a large extent), its conjugate base (Cl⁻
) will be weak. Therefore, in this system, mostH⁺
will behydroniumionsH
₃O⁺
instead of attached to Cl⁻anions and the conjugate bases will be weaker than water molecules.

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 abase splitting constant(Kb) of about5.6×10⁻¹⁰,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.

Identifying conjugate acid–base pairs

To identify the conjugate acid, look for the pair of compounds that are related. Theacid–base reactioncan 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 anacidbecause it donates a proton to the water molecule and itsconjugate baseisnitrate(NO⁻
₃). The water molecule acts as a base because it receives the hydrogen cation (proton) and its conjugate acid is thehydroniumion (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 abuffer 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 thecarbonic acid-bicarbonate buffer,which prevents drastic pH changes whenCO
₂is introduced. This functions as such: ${\ce {CO2 + H2O <=> H2CO3 <=> HCO3^- + H+}}$

Furthermore, here is a table of common buffers.

Buffering agent	pK_a	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 formulaCH
₃COOH,was made into a buffer solution, it would need to be combined with its conjugate baseCH
₃COO⁻
in the form of a salt. The resulting mixture is called an acetate buffer, consisting of aqueousCH
₃COOHand aqueousCH
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 solutionis an example where the conjugate base of an organic acid,lactic acid,CH
₃CH(OH)CO⁻
₂is combined with sodium, calcium and potassium cations and chloride anions in distilled water^[4]which together form a fluid which isisotonicin relation to human blood and is used forfluid resuscitationafterblood lossdue totrauma,surgery,or aburn injury.^[5]

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⁺ion). Acid strength decreases and conjugate base strength increases down the table.

Acid	Conjugate base
H ₂F⁺ Fluoroniumion	HFHydrogen fluoride
HClHydrochloric acid	Cl⁻Chlorideion
H₂SO₄Sulfuric acid	HSO⁻ ₄Hydrogen sulfateion (bisulfateion)
HNO₃Nitric acid	NO⁻ ₃Nitrateion
H₃O⁺Hydroniumion	H₂OWater
HSO⁻ ₄Hydrogen sulfateion	SO²⁻ ₄Sulfateion
H₃PO₄Phosphoric acid	H₂PO⁻ ₄Dihydrogen phosphateion
CH₃COOHAcetic acid	CH₃COO⁻Acetateion
HFHydrofluoric acid	F⁻Fluorideion
H₂CO₃Carbonic acid	HCO⁻ ₃Hydrogen carbonateion
H₂SHydrosulfuric acid	HS⁻Hydrosulfideion
H₂PO⁻ ₄Dihydrogen phosphateion	HPO²⁻ ₄Hydrogen phosphateion
NH⁺ ₄Ammoniumion	NH₃Ammonia
H₂O Water (pH=7)	OH⁻Hydroxideion
HCO⁻ ₃Hydrogencarbonate(bicarbonate)ion	CO²⁻ ₃Carbonateion

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 ₂Ethylamine	C ₂H ₅NH⁺ ₃Ethylammoniumion
CH ₃NH ₂Methylamine	CH ₃NH⁺ ₃Methylammoniumion
NH ₃Ammonia	NH⁺ ₄Ammoniumion
C ₅H ₅NPyridine	C ₅H ₆N⁺ Pyridinium
C ₆H ₅NH ₂Aniline	C ₆H ₅NH⁺ ₃Phenylammoniumion
C ₆H ₅CO⁻ ₂Benzoateion	C ₆H ₆CO ₂Benzoic acid
F⁻ Fluorideion	HFHydrogen fluoride
PO³⁻ ₄Phosphateion	HPO²⁻ ₄Hydrogen phosphateion
OH⁻Hydroxideion	H₂OWater(neutral,pH7)
HCO⁻ ₃Bicarbonate	H ₂CO ₃Carbonic acid
CO²⁻ ₃Carbonate ion	HCO⁻ ₃Bicarbonate
Br⁻ Bromideion	HBrHydrogen bromide
HPO²⁻ ₄Hydrogen phosphate	H ₂PO⁻ ₄Dihydrogen phosphateion
Cl⁻ Chlorideion	HClHydrogen chloride
H ₂OWater	H ₃O⁺ Hydroniumion
${\ce {NO2-}}$ Nitriteion	${\ce {HNO2}}$ Nitrous acid

References

^Zumdahl, Stephen S., & Zumdahl, Susan A.Chemistry.Houghton Mifflin, 2007,ISBN 0618713700
^"Brønsted–Lowry theory | chemistry".Encyclopedia Britannica.Retrieved25 February2020.
^"Strength of Conjugate Acids and Bases Chemistry Tutorial".ausetute.au.Retrieved25 February2020.
^British national formulary: BNF 69(69 ed.). British Medical Association. 2015. p. 683.ISBN 9780857111562.
^Pestana, Carlos (7 April 2020).Pestana's Surgery Notes(Fifth ed.). Kaplan Medical Test Prep. pp. 4–5.ISBN 978-1506254340.

External links

[1] Zumdahl, Stephen S., & Zumdahl, Susan A.Chemistry.Houghton Mifflin, 2007,ISBN 0618713700

[2] "Brønsted–Lowry theory | chemistry".Encyclopedia Britannica.Retrieved25 February2020.

[Conjugate_Strength-3] "Strength of Conjugate Acids and Bases Chemistry Tutorial".ausetute.au.Retrieved25 February2020.

[BNF69-4] British national formulary: BNF 69(69 ed.). British Medical Association. 2015. p. 683.ISBN 9780857111562.

[PestanaFifth-5] Pestana, Carlos (7 April 2020).Pestana's Surgery Notes(Fifth ed.). Kaplan Medical Test Prep. pp. 4–5.ISBN 978-1506254340.

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