Theories of Acids and Bases: Arrhenius, Brønsted-Lowry, and Neutralization
Theory of Ionic Dissociation of Arrhenius
This theory postulates the existence of positive and negative ions in aqueous solutions of acids, bases, and salts (electrolytes) to explain their electrical conductivity.
Acid
An electrically neutral substance that, in aqueous solution, dissociates into H+ ions (protons) and negative ions.
Base
An electrically neutral substance that, in aqueous solution, dissociates into OH- ions (hydroxyl or hydroxide) and positive ions.
According to this theory, a neutralization reaction occurs between an acid and a base, obtaining salt and water.
Benefits:
- Defines what is a base and provides justification for the neutralization process.
- Introduces the concept of the degree of dissociation, which allows for comparison between the strength of an acid and a base. The higher the value, the more dissociated a substance is, indicating a stronger acid or base.
- Provides a quantitative notion of an acid or base.
Limitations:
- Limits the definitions to aqueous solutions.
- Considers OH- ions as solely responsible for basicity, but does not justify the basicity of substances like ammonia (NH3), the carbonate ion (CO3 2-), and oxides, which lack these ions.
- Does not account for the small size of the H+ ion and its reaction with the negative dipole of water, forming the hydronium ion: H+ + H2O -> H3O+ (hydronium ion).
Brønsted-Lowry or Conjugate Acid-Base Theory
This theory does not consider acids and bases in isolation but as interrelated entities. Reactions are viewed as proton transfer reactions.
Acid
A substance (molecule or ion) capable of donating a proton (H+) to another substance (acid). Example: HCl + H2O -> Cl- + H3O+.
Base
A substance (molecule or ion) capable of accepting a proton (H+) from another substance (base). Example: NH3 + H2O -> NH4+ + OH- (ammonium ion).
The concepts of acid and base are complementary. An acid acts as a proton donor only in the presence of a base, which accepts the proton. Conversely, a base can only accept a proton if it reacts with an acid that donates it.
Neutralization, in this context, is the transfer of a proton from an acid to a base. This can be represented as an equilibrium where substances can transfer protons among themselves.
Acid + Base Conjugate Base + Conjugate Acid
A conjugate pair (or conjugate acid-base pair) consists of an acid and its conjugate base or a base and its conjugate acid. Both pairs are interconvertible through the gain or loss of a proton.
HA + B A- + BH+
- A- is the conjugate base of acid HA -> HA/A- is a conjugate acid-base pair.
- HB+ is the conjugate acid of base B -> B/BH+ is a conjugate acid-base pair.
Strengths of Acids and Bases
The tendency to donate or accept protons is relative and depends on the substance involved in the reaction. Water is often used as a reference substance.
Strong Acids
These substances completely dissociate in water due to their strong tendency to donate protons to a base. They have a degree of dissociation (α) ≈ 1. This represents an irreversible reaction.
(Strong Acid) HA + H2O -> A- + H3O+ (Weak Conjugate Base)
Weak Acids
These substances only partially dissociate in water due to their low tendency to donate protons to a base. They have a very small degree of dissociation (α
The higher the value of Ka (acid dissociation constant), the more the equilibrium shifts to the right, favoring dissociation. This results in a higher concentration of hydronium ions (H3O+), a higher degree of dissociation, and greater acid strength. Consequently, the pH is lower.
Strong Bases
These substances completely ionize in water due to their strong tendency to accept protons from an acid. Their degree of dissociation is practically 1.
(Strong Base) B + H2O -> BH+ + OH- (Weak Conjugate Acid)
Weak Bases
These substances partially ionize in water due to their low tendency to accept protons. Their degree of dissociation is very small (α
The higher the value of Kb (base dissociation constant), the more the equilibrium shifts towards the formation of ions. This leads to a higher concentration of OH-, a higher degree of dissociation, greater base strength, and consequently, a higher pH.
α = Quantity of Substance Dissociated / Initial Quantity of Reactant
When Ka or Kb values are small (
Neutralization Reaction
Excess Acid
If there is an excess of acid, all the base reacts with only a part of the acid. The final solution will be acidic (pH
Excess Base
If there is an excess of base, all the acid reacts with only a part of the base. The final solution will be basic (pH > 7).
Stoichiometric Amounts
of acid and base: All this and all acid base react with each other, leaving no excess of any of them. In this case, the neutralization is complete. The pH of the resulting solution depends on the concentration of hydronium ion with respect to hydroxyl ions. 2