# Acid - base theories

Basic definitions of acid and base

• According to Arrhenius

Are acidic all those substances which when placed in water are able to dissociate, freeing H+ ions in solution. For example:

-  $HCl&space;\xrightarrow[]{H_2O}&space;H^+&space;+&space;Cl^-$

-  $CH_3COOH&space;\xrightarrow[]{H_2O}&space;CH_3COO^-&space;+&space;H^+$

Vice versa, are bases all those substances that once placed in water are able to dissociate, freeing OH- ions. For example:

-  $NaOH&space;\xrightarrow[]{H_2O}&space;Na^+&space;+&space;OH^-$

limits of the TheoryIt has been verified experimentally that in solution may not exist the free proton (H+). This has a large charge density and generates a very strong electron field. Therefore, is always associated with one or more molecules of water, H3O+ , H9O4+

Although for simplicity we'll continue writing H+, in reality this ion is always in his hydrated form. Then an acid cannot be a substance that simply free H+ ions; there must be another substance that accepts the proton.

The basic concept should be extended. In fact there are bases that send in solution hydroxyl ions (OH-) and that therefore able to neutralize the acidity. However there are many substances that are able to neutralize the acids even if they do not possess OH- functions.

• According to Bronsted and Lowry

- An acid is a species that has the tendency to lose (or donate ) a proton.

- a base is a species that has the tendency to buy (or accept) a proton.

Implications → Given that a proton cannot exist as free in aqueous solutions , it follows that " a substance can act by acid only in presence of a base" and vice versa "a substance can behave as a base only in presence of an acid"

In terms of the Bronsted and Lowry theory, a substance can donate a proton only if there is another substance ready to take it and vice versa.

HA + B  $\rightleftharpoons$ A- + HB+

Example:

HI + H2O  $\rightleftharpoons$ I+ H3O+

There are substances that can behave as bases in respect of certain substances and as acids in respect of others → amphoterism

H2O + NH3  $\rightleftharpoons$ OH- + NH4   (water behaves as acid)

H2O + HCl  $\rightleftharpoons$ H3O+ + Cl-   (water behaves as base)

Given that these reactions have the sign of reversibility ($\rightleftharpoons$) we deduce that they can also go in the opposite direction; for example, Cl- could take a proton from H3O+ giving back HCl.  Actually this happens but to an extent totally negligible. Species that differ for 1 proton are called conjugate acid - base pairs (es: HCl/Cl- , H3O+/H2O).

In an acid-base reactions we can deduce the rcelative concentrations of the two species of the pair. It is however necessary to know the concept of acidic and basic strength, described below.

Limits of the theory →  The Bronsted-Lowry theory does not justify some reactions that occur without transfer of protons.

• According to Lewis

- Are acidic those substances that are able to accept a pair of non-bonding electrons.

- Are basic those substances that can donate a pair of non-bonding electrons.

AlCl+ Cl-  $\rightleftharpoons$  AlCl4-

For example, this reaction can be considered an acid-base. The aluminum is able to accept a pair of electrons, and therefore behaves as a Lewis acid, while the Cl- donates a pair of non-bonding electrons and therefore behaves as a Lewis base.

It should be noticed that a base according to Bronsted is also a base according to Lewis. In fact, a Lewis base can always associate a proton through a pair of non-bonding electrons.

A Bronsted acid requires a proton to transfer. This condition is not required instead with a Lewis acid, that only needs an available orbital to associate an electronic doublet.