Acids and Bases

Definitions

Venn diagram showing the relationship between three acid-base definitions: Arrhenius acids and bases are a subset of Brønsted–Lowry acids and bases, which in turn are a subset of Lewis acids and bases.


Arrhenius acid (green) – a substance that, when added to water, causes an increase in the molar concentration of hydronium ions ([H₃O⁺]) and a decrease in the molar concentration of hydroxide ions ([OH⁻]).

\[\begin{aligned} {\color{green}\mathrm{HA}}\mathrm{(aq)} + \mathrm{H_2O(l)} &\rightleftharpoons \mathrm{H_3O^+(aq)} + {\color{red}\mathrm{A^-}}\mathrm{(aq)} \end{aligned}\]

Arrhenius base (red) – a substance that, when added to water, causes an increase in the molar concentration of hydroxide ions ([OH⁻]) and a decrease in the molar concentration of hydronium ions ([H₃O⁺]).

\[\begin{aligned} \textrm{(by dissociation)} \quad \mathrm{NaOH}\mathrm{(s)} &\rightarrow \mathrm{Na^+(aq)} + {\color{red}\mathrm{OH^-}}\mathrm{(aq)} \\ \textrm{(by reaction)} \quad {\color{red}\mathrm{B}}\mathrm{(aq)} + \mathrm{H_2O(l)} &\rightleftharpoons \mathrm{OH^-(aq)} + {\color{green}\mathrm{HB^+}}\mathrm{(aq)} \end{aligned}\]


Brønsted-Lowry acid (green) – a molecular entity capable of donating a proton (hydron) to a base (a proton donor).

Brønsted-Lowry base (red) – a molecular entity capable of accepting a proton (hydron) from an acid (a proton acceptor).

A B-L reaction is a proton transfer. Water can act as either an acid or a base (water is amphoteric).

\[\begin{array}{ccccccc} {\color{green}\mathrm{HA}} & \!\!\!+\!\!\! & {\color{red}\mathrm{H_2O}} & \!\!\!\rightleftharpoons\!\!\! & {\color{green}\mathrm{H_3O^+}} & \!\!\!+\!\!\! & {\color{red}\mathrm{A^-}} \\ \small\text{B-L acid} & & \small\text{B-L base} & & \small\text{conj. acid} & & \small\text{conj. base} \end{array}\] \[\begin{array}{ccccccc} {\color{green}\mathrm{H_2O}} & \!\!\!+\!\!\! & {\color{red}\mathrm{B}} & \!\!\!\rightleftharpoons\!\!\! & {\color{green}\mathrm{HB^+}} & \!\!\!+\!\!\! & {\color{red}\mathrm{OH^-}} \\ \small\text{B-L acid} & & \small\text{B-L base} & & \small\text{conj. acid} & & \small\text{conj. base} \end{array}\]

The B-L theory also applies to non-aqueous systems:

\[\begin{array}{ccccc} {\color{green}\mathrm{HCl}}(\mathrm{g}) & \!\!\!+\!\!\! & {\color{red}\mathrm{NH_3}}(\mathrm{g}) & \!\!\!\rightleftharpoons\!\!\! & {\color{green}\mathrm{NH_4Cl}}(\mathrm{s}) \\ \small\text{B-L acid} & & \small\text{B-L base} & & \small\text{salt} \end{array}\]


Lewis acid (green) – a molecular entity that is an electron-pair acceptor.

Lewis base (red) – a molecular entity that is an electron-pair donor.

The Lewis theory is the most general acid-base theory because it defines reactions in terms of electron-pair transfer, allowing it to encompass reactions that do not involve protons.

\[\begin{array}{ccccc} {\color{green}\mathrm{BF_3}} & \!\!\!+\!\!\! & {\color{red}\mathrm{:NH_3}} & \!\!\!\rightarrow\!\!\! & \mathrm{F_3B{-}NH_3} \\ \small\text{Lewis acid} & & \small\text{Lewis base} & & \small\text{adduct} \end{array}\]

Here, the lone pair on nitrogen is donated to the electron-deficient boron atom.


Conjugate acid-base pairs

A deprotonated (−H+) acid results in a conjugate base. A protonated (+H+) base results in a conjugate acid.

\[{\color{green}\mathrm{HA}}\underset{+\mathrm{H}^+}{\overset{-\mathrm{H}^+}{\rightleftharpoons}} {\color{red}\mathrm{A^-}}\]

A protonated (+H+) base results in a conjugate acid. A deprotonated (−H+) acid results in a conjugate base.

\[{\color{red}\mathrm{B}} \underset{-\mathrm{H}^+}{\overset{+\mathrm{H}^+}{\rightleftharpoons}} {\color{green}\mathrm{HB^+}}\]

Classifications

Table 1: Some common acids and bases


Table 2: Common anions and their corresponding acids

Dissociation Constants

Notes
Table 3: Dissociation constants for inorganic acids and bases


Table 4: Dissociation constants for organic acids and bases