6.10 Acids, Bases and Salts

Many acids, bases and salts find widespread occurrence in nature. Acetic acid is known to be the main constituent of vinegar. Lemon and orange juices contain citric and ascorbic acids, hydrochloric acid in gastric juices, and tartaric acid is found in tamarind paste.

Similarly, many bases are found such as lime water. We use many of these acids in our day-to-day life, such as vinegar or acetic acid in the kitchen, boric acid for laundry, baking soda for the purpose of cooking, washing soda for cleaning, etc.

Acids

An acid is defined as a substance whose water solution tastes sour, turns blue litmus red, and neutralizes bases. As most of the acids taste sour, the word “acid” has been derived from a latin word “acidus” meaning sour. Acids are known to turn blue litmus paper into red and liberate dihydrogen on reacting with some metals.

These dissociate in their aqueous solution to form their constituent ions, as given by the following examples.
\(
\mathrm{HCl}(\mathrm{aq}) \longrightarrow \mathrm{H}^{+}+\mathrm{Cl}^{-}
\)
\(
\mathrm{H}_2 \mathrm{SO}_4(\mathrm{aq}) \longrightarrow 2 \mathrm{H}^{+}+\mathrm{SO}_4^{2-}
\)
\(
\mathrm{CH}_3 \mathrm{CO}_2 \mathrm{H}(\mathrm{aq}) \longrightarrow \mathrm{H}^{+}+\mathrm{CH}_3 \mathrm{CO}_2^{-}
\)
\(
\text { Dissociation of Acids }
\)
Based on their occurrence, they are divided into two types- Natural and mineral acids.

Natural Acids: These are obtained from natural sources, such as fruits and animal products. For e.g. lactic, citric, and tartaric acid etc.

Mineral Acids: Mineral acids are acids prepared from minerals. Examples are Hydrochloric acid \((\mathrm{HCl})\), Sulphuric Acid \(\left(\mathrm{H}_2 \mathrm{SO}_4\right)\), and nitric acid \(\left(\mathrm{HNO}_3\right)\), etc.

Bases

A substance is called base if its aqueous solution tastes bitter, turns red litmus blue, or neutralizes acids. Bases are known to turn red litmus paper blue, taste bitter and feel soapy. A common example of a base is washing soda used for washing purposes.

The bases dissociate in their aqueous solution to form their constituent ions, given in the following examples.

\(
\begin{aligned}
& \mathrm{NaOH}(\mathrm{aq}) \longrightarrow \mathrm{Na}^{+}+\mathrm{OH}^{-} \\
& \mathrm{Ca}(\mathrm{OH})_2 \longrightarrow \mathrm{Ca}^{2+}+2 \mathrm{OH}^{-}
\end{aligned}
\)

Dissociation of Bases

Salt

Salt is a neutral substance whose aqueous solution does not affect litmus. When acids and bases are mixed in the right proportion they react with each other to give salts. Some commonly known examples of salts are sodium chloride, barium sulphate, sodium nitrate. Sodium chloride (common salt) is an important component of our diet and is formed by reaction between hydrochloric acid and sodium hydroxide. It exists in the solid state as a cluster of positively charged sodium ions and negatively charged chloride ions which are held together due to electrostatic interactions between oppositely charged species (Fig. 7.10).

The formation of salt can be seen from the chemical reactions shown in the equations below.
\(
\begin{array}{ll}
\mathrm{HCl}+\mathrm{NaOH} & \longrightarrow \mathrm{NaCl}+\mathrm{H}_2 \mathrm{O} \\
\mathrm{H}_2 \mathrm{SO}_4+\mathrm{Ca}(\mathrm{OH})_2 & \longrightarrow \mathrm{CaSO}_4+2 \mathrm{H}_2 \mathrm{O}
\end{array}
\)
Salts are Formed from the Neutralization Reactions of Acids and Bases

Note: The terms dissociation and ionization have earlier been used with different meaning. Dissociation refers to the process of separation of ions in water already existing as such in the solid state of the solute, as in sodium chloride. On the other hand, ionization corresponds to a process in which a neutral molecule splits into charged ions in the solution. Here, we shall not distinguish between the two and use the two terms interchangeably.

Arrhenius Concept of Acids and Bases

According to Arrhenius theory, acids are substances that dissociates in water to give hydrogen ions \(\mathrm{H}^{+}(\mathrm{aq})\) and bases are substances that produce hydroxyl ions \(\mathrm{OH}^{-}(\mathrm{aq})\). The ionization of an acid \(\mathrm{HX}\) (aq) can be represented by the following equations:
\(
\begin{gathered}
\mathrm{HX}(\mathrm{aq}) \rightarrow \mathrm{H}^{+}(\mathrm{aq})+\mathrm{X}^{-}(\mathrm{aq}) \\
\text { or } \\
\mathrm{HX}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{X}^{-}(\mathrm{aq})
\end{gathered}
\)
A bare proton, \(\mathrm{H}^{+}\)is very reactive and cannot exist freely in aqueous solutions. Thus, it bonds to the oxygen atom of a solvent water molecule to give trigonal pyramidal hydronium ion, \(\mathrm{H}_3 \mathrm{O}^{+}\left\{\left[\mathrm{H}\left(\mathrm{H}_2 \mathrm{O}\right)\right]^{+}\right\}\). In this chapter we shall use \(\mathrm{H}^{+}(\mathrm{aq})\) and \(\mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})\) interchangeably to mean the same i.e., a hydrated proton.
Similarly, a base molecule like \(\mathrm{MOH}\) ionizes in aqueous solution according to the equation:
\(
\mathrm{MOH}(\mathrm{aq}) \rightarrow \mathrm{M}^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})
\)

The hydroxyl ion also exists in the hydrated form in the aqueous solution. Arrhenius concept of acid and base, however, suffers from the limitation of being applicable only to aqueous solutions and also, does not account for the basicity of substances like, ammonia which do not possess a hydroxyl group.

The Brönsted-Lowry Acids and Bases

According to Brönsted-Lowry theory, acid is a substance that is capable of donating a hydrogen ion \(\mathrm{H}^{+}\)and bases are substances capable of accepting a hydrogen ion, \(\mathrm{H}^{+}\). In short, acids are proton donors and bases are proton acceptors.

The species B accepts a proton and thus behaves as a base, while the species \(\mathrm{BH}^{+}\)gives up a proton and thus acts as an acid
\(
\mathrm{B}+\quad \quad  \mathrm{H}^{+} \quad \rightleftharpoons \mathrm{BH}^{+}
\)
\(
\text { (Base) (Proton) (acid) }
\)

Consider the example of dissolution of \(\mathrm{NH}_3\) in \(\mathrm{H}_2 \mathrm{O}\) represented by the following equation:

The basic solution is formed due to the presence of hydroxyl ions. In this reaction, water molecule acts as proton donor and ammonia molecule acts as proton acceptor and are thus, called Lowry-Brönsted acid and base, respectively. In the reverse reaction, \(\mathrm{H}^{+}\) is transferred from \(\mathrm{NH}_4^{+}\)to \(\mathrm{OH}^{-}\). In this case, \(\mathrm{NH}_4^{+}\)acts as a Bronsted acid while \(\mathrm{OH}^{-}\)acted as a Brönsted base. The acid-base pair that differs only by one proton is called a conjugate acid-base pair. Therefore, \(\mathrm{OH}^{-}\)is called the conjugate base of an acid \(\mathrm{H}_2 \mathrm{O}\) and \(\mathrm{NH}_4^{+}\)is called conjugate acid of the base \(\mathrm{NH}_3\). If Brönsted acid is a strong acid then its conjugate base is a weak base and viceversa. It may be noted that conjugate acid has one extra proton and each conjugate base has one less proton.

Consider the example of ionization of hydrochloric acid in water. \(\mathrm{HCl}(\mathrm{aq})\) acts as an acid by donating a proton to \(\mathrm{H}_2 \mathrm{O}\) molecule which acts as a base.

It can be seen in the above equation, that water acts as a base because it accepts the proton. The species \(\mathrm{H}_3 \mathrm{O}^{+}\)is produced when water accepts a proton from \(\mathrm{HCl}\). Therefore, \(\mathrm{Cl}^{-}\)is a conjugate base of \(\mathrm{HCl}\) and \(\mathrm{HCl}\) is the conjugate acid of base \(\mathrm{Cl}^{-}\). Similarly, \(\mathrm{H}_2 \mathrm{O}\) is a conjugate base of an acid \(\mathrm{H}_3 \mathrm{O}^{+}[latex]and [latex]\mathrm{H}_3 \mathrm{O}^{+}\)is a conjugate acid of base \(\mathrm{H}_2 \mathrm{O}\).

It is interesting to observe the dual role of water as an acid and a base. In case of reaction with \(\mathrm{HCl}\) water acts as a base while in case of ammonia it acts as an acid by donating a proton.

Example 7.12: What will be the conjugate bases for the following Brönsted acids: \(\mathrm{HF}, \mathrm{H}_2 \mathrm{SO}_4\) and \(\mathrm{HCO}_3^{-}\)?

Answer: The conjugate bases should have one proton less in each case and therefore the corresponding conjugate bases are: \(\mathrm{F}^{-}\), \(\mathrm{HSO}_4^{-}\)and \(\mathrm{CO}_3^{2-}\) respectively.

Example 7.13: Write the conjugate acids for the following Brönsted bases: \(\mathrm{NH}_2^{-}, \mathrm{NH}_3\) and \(\mathrm{HCOO}^{-}\).

Answer: The conjugate acid should have one extra proton in each case and therefore the corresponding conjugate acids are: \(\mathrm{NH}_3\), \(\mathrm{NH}_4^{+}\)and \(\mathrm{HCOOH}\) respectively.

Example 7.14: The species: \(\mathrm{H}_2 \mathrm{O}, \mathrm{HCO}_3^{-}, \mathrm{HSO}_4^{-}\)and \(\mathrm{NH}_3\) can act both as Bronsted acids and bases. For each case give the corresponding conjugate acid and conjugate base.

Answer: The answer is given in the table below.

\(
\begin{array}{lcc}
\text { Species } & \begin{array}{c}
\text { Conjugate } \\
\text { acid }
\end{array} & \begin{array}{c}
\text { Conjugate } \\
\text { base }
\end{array} \\
\mathrm{H}_2 \mathrm{O} & \mathrm{H}_3 \mathrm{O}^{+} & \mathrm{OH}^{-} \\
\mathrm{HCO}_3^{-} & \mathrm{H}_2 \mathrm{CO}_3 & \mathrm{CO}_3^{2-} \\
\mathrm{HSO}_4^{-} & \mathrm{H}_2 \mathrm{SO}_4 & \mathrm{SO}_4^{2-} \\
\mathrm{NH}_3 & \mathrm{NH}_4^{+} & \mathrm{NH}_2^{-} \\
\end{array}
\)

Lewis Acids and Bases

G.N. Lewis in 1923 defined an acid as a species which accepts electron pair and base which donates an electron pair. As far as bases are concerned, there is not much difference between Brönsted-Lowry and Lewis concepts, as the base provides a lone pair in both cases. However, in Lewis concept many acids do not have proton. A typical example is reaction of electron-deficient species \(\mathrm{BF}_3\) with \(\mathrm{NH}_3\).
\(\mathrm{BF}_3\) does not have a proton but still acts as an acid and reacts with \(\mathrm{NH}_3\) by accepting its lone pair of electrons. The reaction can be represented by,
\(
\mathrm{BF}_3+: \mathrm{NH}_3 \rightarrow \mathrm{BF}_3: \mathrm{NH}_3
\)

Electron deficient species like \(\mathrm{AlCl}_3, \mathrm{Co}^{3+}\), \(\mathrm{Mg}^{2+}\), etc. can act as Lewis acids while species like \(\mathrm{H}_2 \mathrm{O}, \mathrm{NH}_3, \mathrm{OH}^{-}\)etc. which can donate a pair of electrons, can act as Lewis bases.

Example 7.15: Classify the following species into Lewis acids and Lewis bases and show how these act as such:
(a) \(\mathrm{HO}^{-}\)
(b) \(\mathrm{F}^{-}\)
(c) \(\mathrm{H}^{+}\)
(d) \(\mathrm{BCl}_3\)

Answer: (a) Hydroxyl ion is a Lewis base as it can donate an electron lone pair (: \(\left.\mathrm{OH}^{-}\right)\).
(b) Flouride ion acts as a Lewis base as it can donate any one of its four electron lone pairs.
(c) A proton is a Lewis acid as it can accept a lone pair of electrons from bases like hydroxyl ion and fluoride ion.
(d) \(\mathrm{BCl}_3\) acts as a Lewis acid as it can accept a lone pair of electrons from species like ammonia or amine molecules.