In this article, we will be discussing organic acids and bases. Different scientists have proposed different theories for acids and bases. So in this article, we will be looking at the different theories proposed by them one by one.
As the Lowry-Bronsted theory was considered to be somewhat restrictive, Lewis defined acids as substances capable of accepting electron pairs, and bases as substances having available unshared electron pair. It should be noted that a base in Lewis theory is the same as in the Lowry-Bronsted theory; both are substances with an available pair of electrons. However, Lewis definition of acids includes many substances, such as boron trifluoride, zinc chloride, and aluminium chloride, which have only six electrons in the outer shell, whereas they can accommodate 8 electrons, and consequently react with bases.
Nucleophiles and electrophiles may also be looked upon as bases and acids, respectively. By Lewis definition, an acid is a substance capable of accepting an electron pair and a base is an electron pair donor. Since electrophiles are electron seeking reagent they may be called Lewis acids whereas electron donating nature nucleophile may be considered as Lewis bases. However, the terms acids and bases are commonly confined to their capacity for breaking and making of bonds to hydrogen while electrophile and nucleophile refer to the reactivity at carbon. Thus, Hydroxide ion is called a base when it attacks a Proton and a nucleophile when it attacks the carbon atom of methyl chloride.
Strength of Acids and Bases:
All protonic acids ionize in aqueous solution by transferring protons to the solvent molecules; stronger acids give up most of their ionizable protons, while weaker acids transfer only a smaller proportion. The equilibrium constant (Ka) of the process is given by the following expression:
Water is present in excess, and its concentration has been taken as a constant. A large value of Ka signifies a high degree of ionization, and therefore, a stronger acid. Since the position of equilibrium depends on the relative stabilities of reactants and products, the acidity constant (Ka) of acid is related to standard free energy change (ΔGo) by the following expression:
ΔGo = -2.303 RT log Ka
It is more convenient to express the strength of acid by pKa which is the negative logarithm of Ka. A stronger acid has larger Ka, but smaller pKa value. The strength of a base is expressed by an equilibrium constant, Kb. The relationship between pKa and pKb in water is given by the expression, pKa + pKb = 14. It has been found more convenient to express the base strength by Ka of the conjugate acid.
The weaker is base, the stronger it’s conjugate acid. Thus, the strength of acids and bases may be compared by pKa values, which express the relative acidity of the conjugate acid for any conjugate pair. Larger the value of pKa, stronger the base, the pKa values are not related to the degree of ionization beyond 0 to 14 range. Values for such acids and bases are related to water by indirect methods and are often less accurate than those between 0 to 14.