6.2: Brønsted–Lowry acids and bases
Some bases do not have hydroxide ions in their formula, yet they act as bases and neutralize acids. For example, ammonia (NH 3 ) and calcium carbonate (CaCO 3 ) do not contain hydroxide ions, but they neutralize acids. Further, Arrhenius's definition limits the acid-base reactions in the water medium. The acid-base reactions can take place in other mediums also, e.g., HCl –an acid, and NH 3 –a base can react with and neutralize each other in the gas phase also. The Brønsted–Lowry bordered the definition of acids and bases by including the bases mentioned above and also by including acid-base reactions in a non-aqueous medium.
Brønsted–Lowry's definition of acids and bases
Brønsted–Lowry's definition states that:
- An acid is a proton donor, and
- A base is a proton acceptor.
or example, HCl is an acid because it donates a proton to the water solvent.
\[\mathrm{HCl}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})\nonumber\]
Water is a base in the above reaction because it accepts a proton from the acid. In a reaction between HCl and NH 3 :
\[\mathrm{HCl}(\mathrm{g})+\mathrm{NH}_{3}(\mathrm{~g}) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(\mathrm{s})\nonumber\]
HCl is an acid because it donates its proton to NH 3 , and the NH 3 is a base because it accepts a proton, as shown in Fig. 6.2.1.
Brønsted–Lowry's acids have ionizable protons that they donate to bases. Therefore, Brønsted–Lowry's acid is generally written as HA, where H + is the donatable proton, and A - is the anion of the acid. Examples of acids are HCl, H 2 SO 4 , HNO 3 , and CH 3 COOH. Note that acetic acid has only one acidic proton that is attached to the O atom in the carboxylic acid group (–COOH). The rest of the protons attached to carbon atoms are not acidic. All organic acids have a carboxylic acid group (–COOH). Brønsted–Lowry's acid may have net +ve charge, no charge, or net –ve charge on it. For example, H 3 O + , HCl, and HSO 4 - are all acids because they can donate a proton to a base.
Mono-, di- and tri-protic acids
Acids that have only one acidic proton are mono- protic , e.g., H Cl , H NO 3 , CH 3 COO H , are mono- protic acids where the acid proton is shown in bold font. Some acids have two acidic protons –they are di- protic , e.g., H 2 SO 4 and H 2 CO 3 are di- protic . The acids with three acidic protons are tri- protic , e.g., H 3 PO 3 is a tri- protic acid. For example, phosphoric acid ( H 3 PO 4 ) can dissociate and donate three protons, as shown in chemical reactions below:
\begin{aligned}
&\mathrm{H}_{3} \mathrm{PO}_{4}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{H}_{2} \mathrm{PO}_{4}{ }^{-} \\
&\mathrm{H}_{2} \mathrm{PO}_{4}{ }^{-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{HPO}_{4}{ }^{2-} \\
&\mathrm{HPO}_{4}{ }^{2-}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}+\mathrm{PO}_{4}{ }^{3-}
\end{aligned}
The base accepts proton by making a bond with it. The bond is a pair of bonded electrons. Since the proton is a hydrogen atom without an electron, both electrons in the bond come from the base. The base must have a lone pair of electrons on it. The base is usually represented as to emphasize a lone pair of electrons on it that is shown as a pair of dots. For example, ammonia, water, and hydroxide ion ( , , and ) are Brønsted–Lowry's bases, because each of these has an atom with lone pair or lone pairs of electrons on them.
Conjugate acid-base pairs
The acid-base reactions described above are one-way reactions, i.e., reactants go-to products almost 100%. However, the majority of the acid-base reactions are two ways, i.e., reactants form the products and the products react with each other and re-form the reactants. Double arrows between reactants and products represent the two ways reactions. For example, hydrofluoric acid (HF) is a weak electrolyte; it partially dissociates in water to form F - and H 3 O + , and the products react to re-form the reactants, as shown in Fig. 6.2.2. In the reverse reaction, H 3 O + is acting as an acid, and F - is acting as a base. The acid and the base in the products are called conjugate acid and conjugate base, respectively. The acid HF becomes conjugate base F - after removal of a proton, and the base H 2 O becomes conjugate acid H 3 O + after accepting a proton.
The conjugate acid-base pair is related to the loss and gain of H + . For example, HF/F - is a conjugate acid-base pair, and H 3 O + /H 2 O is also a conjugate acid-base pair.
In other words, remove the acidic proton from an acid to get its conjugate base and add a proton to a base to get its conjugate acid.
Another example is ammonia NH 3 which dissolves in water and accepts a proton to form its conjugate acid NH 4 + , as shown in Fig. 6.2.3. Water H 2 O acts as an acid by donating a proton and forming its conjugate base OH - . The two conjugate acid-base pairs in this reaction are NH 4 + /NH 3 and H 2 O/OH - that are related by loss and gain or an H + .
Identify the conjugate acid-base pairs in the following reaction?
\[\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq})+\mathrm{NH}_{3}(\mathrm{aq}) \rightleftarrows \mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{NH}_{4}^{+}\nonumber\]
Solution
- Identify the substance that has donated a proton in the reactants –it is an acid.
- Remove a proton from the acid to form its conjugate base: H 3 PO 4 /H 2 PO 4 - .
- Identify the substance that has accepted a proton in the reactants –it is a base.
- Add a proton to the base to form its conjugate acid: NH 3 /NH 4 + .
Note that loss of a proton from an acid forms its conjugated base with the charge decreased by one, e.g., H 3 PO 4 /H 2 PO 4 - , HSO 4 - /SO 4 2 - , and NH 4 + /NH 3 . Similarly, the gain of a proton by a base forms its conjugate acid with the charge increased by one, e.g., HPO 4 2 - /H 2 PO 4 - , HCO 3 - /H 2 CO 3 , and NH 3 /NH 4 + .
Amphoteric substances
Water acts as a base in some reactions, e.g., with HF, and as an acid in some reactions, e.g., with NH 3 .
Substances like water that can act as an acid and also as a base are called amphoteric substances.
Other examples of amphoteric substances include HSO 4 - , HCO 3 - , and NH 3 , as illustrated in Fig. 6.2.4.