Acid-base reactions

What are acids and bases?

In the past, scientists classified the then known substances into acids, bases and salts. In principle, they classified them according to a series of similar properties. In 1663 Robert Boyle established a series of properties common to all acids. Thus, an acid was defined as a substance whose aqueous solution:

  • React with active metals (zinc) releasing H2.
  • Neutralize bases forming salts (they lose their properties).
  • Discolor the violet phenolphthalein solution and colors the litmus red.
  • Dissolve many substances.
  • In dilute solutions, they have a characteristic sour taste (called acid taste).

a base would be defined as the substance that in aqueous solution:

  • Color the phenolphthalein solutions violet red and the litmus paper blue.
  • Neutralizes acids forming salts (they lose their characteristic properties when reacting with acids).
  • For its softness to the touch (hydrolyzes skin proteins).
  • Precipitates many substances that are soluble in acids.
  • Characteristic bitter taste.

A misconception put forward by Lavoisier was that “oxygen is an element common to all acids and it is the presence of oxygen that constitutes acidity”. However, it was later shown that there are acids that do not have oxygen in their molecule (HCl, HCN, H2S, etc.).

In 1838, Liebig defined acids as “compounds containing atomic hydrogen, in which this can be replaced by metals”.

Acids with oxygen and a non-metal element are called oxyacids and are obtained by adding water to an anhydride:

N2O5 + H2O -> 2HNO3

SO3 + H2O -> H2SO4

Cl2O7 + H2O -> 2HClO4

Acids formed by a nonmetal and hydrogen are called hydrazides:

Cl2 + H2 -> 2HCl

S + H2 -> H2S

Bases are compounds made up of oxygen, hydrogen and a metallic element, in which the oxygen and hydrogen meet to form an OH hydroxyl anion. They are obtained by adding water to the oxides:

BaO + H2O -> Ba(OH)2

K2O + H2O -> 2KOH

Acid-base theories

In the course of history, different scientific theories have been proposed for the definition of acidic and basic compounds, which are listed below:

Arrhenius acid-base theory

Brønsted-Löwry acid-base theory

Lewis acid-base theory


Strength of acids and bases. Dissociation constant and degree of dissociation

Not all substances capable of acting as acids give up a proton with the same ease. Acids are all the stronger the greater their tendency to give up a proton.

Similarly, a base will be stronger the greater its tendency to accept protons. The strength of acids and bases can be measured by the magnitude of their reaction with a reference substance (the H2O in general).

The acids HCl, HNO3, H2SO4, HClO4, etc. They are strong acids because they are totally dissociated, i.e. the dissociation equilibrium is completely shifted (to the right) and the values of their constants will therefore be very high.

AH + H2O -> A + H3O+

However, its conjugate bases, Cl, NO3, SO42-, ClO4, etc. are very weak.

On the other hand, the acids CH3-COOH, CNH, CO3H2, etc, are weak acids and their dissociation equilibria are not so shifted to the right as a consequence their dissociation constants will be lower.

Their conjugate bases, CH3-COO, CN, CO3=, etc., will generally be strong bases.

Thus the measure of the strength of an acid is given by the value of its dissociation constant which is obtained by applying the law of mass action (LAM) to the dissociation equilibrium:

According to Arrhenius:

According to Brønsted-Löwry:

Considering that [H2O] = cte, we have that [H3O+] = [H+]

The same treatment can be done for the bases:

The degree of dissociation (α) of an acid or a base is the percent per one of acid AH that is in dissociated form. Thus, in the case of an acid, α is the percent per one of acid AH that is in dissociated form A y H+.

If the acid is strong, it will be fully dissociated, so α=1.

If the acid is weak, it will be partially dissociated with values 0< α< 1.

Suppose an acid that is at an initial concentration C:

AH <–> A + H+

Initial C 0 + 0

Equilibrium C(1-α) Cα + Cα

Ka = [A]·[H+]/[AH] = c2α2/c(1-α) = cα2/(1-α)

In the case of polyprotic acids, dissociation occurs gradually, as in the case of several acids, and each equilibrium process has a corresponding ionization constant.


H3PO4 <–> H2PO4 + H+

Ka1 = [H2PO4]·[H+]/[H3PO4] = 1.1·10-2 mol/l

H2PO4 <–> HPO4= + H+

Ka1 = [HPO4=]·[H+]/[H2PO4] = 7.5·10-8 mol/l

HPO4= <–> PO43- + H+

Ka1 = [PO43-]·[H+]/[HPO4=] = 5.0·10-13 mol/l

Water dissociation constant

As already mentioned, there are substances that can behave as acids or bases, depending on against whom they act, these substances are called amphoteric, amphiprotic or ampholytic substances.


Al(OH)3  AlO2 + H3O+ (acid)

Al(OH)3  Al3+ + 3OH (base)

In the case of water, the autoprotolysis reaction occurs.

H2O + H2O  H3O+ + OH

A1 + B A2 + B1

With very sensitive apparatus, they show that even pure water has a small conductivity.

H2O + H2O  H3O+ + OH

Kw cte of ionic product of water

The value of Kw at 25ºC is 10-14.

In pure water, each H3O+ ion formed is accompanied by an [H3O+] = [OH] y [H3O+] = [OH] = = 10-7 ion.


Neutral solutions:

[H3O+] = [OH] = 10-7

Acid solutions:

[H3O+] > 10-7 y [OH] < 10-7

[H3O+] > [OH]

Basic solutions:

[H3O+] < 10-7 y [OH] > 10-7

[OH] > [H3O+]

Concept and pH scale

Leveling and differentiating effect of a solvent

Acids such as HI, HBr, HCl, HClO4, HNO3, H2SO4, etc. are practically 100% dissociated in aqueous solutions that are not too concentrated. We could conclude that all these acids would have the same strength, which does not seem very reasonable if we consider the structural differences between them. The reason lies in the fact that water behaves as a strong base against these acids, retaining with intensity the yielded proton.

Water, in this sense, levels the forces of all these strong acids, it is therefore a leveling solvent.

We define the intrinsic acidity constant as:

AH <–> A + H+

KAHa = [A]·[H+]/[AH]

The solvent

DH + H+ DH2+

KDHb = [DH2+]/[DH]·[H+]

La constante de acidez real será:

AH +DH  A + DH2+

Ka = [A]·[DH2+]/[AH]·[DH] = KAHa·KDHb

In the same way, the following is obtained: Kb = KBb·KDHa


A1H ; KA1Ha = 105

A2H ; KA2Ha = 10-2

A3H ; KA3Ha = 10-4

DH ; KDHb = 1010 (water)


Ka1 = 1015

Ka2 = 108

Ka3 = 106

All are strong acids

It is the leveling effect of the solvent.


A1H ; KA1Ha = 105

A2H ; KA2Ha = 10-2

A3H ; KA3Ha = 10-4

DH ; KDHb = 102 (methanol)


Ka1 = 107 strong acid

Ka2 = 1 normal acid

Ka3 = 10-2 weak acid

This is the differentiating effect of a solvent such as methanol compared to these acids.


pH buffer solutions

Acid-base indicators

Acid-base titrations