Acid Base Theory - Pharmaceutical Inorganic Chemistry B. Pharma 1st Semester

Acid Base Theory


Theories on:

• Acids

• Bases and

• Salts

Learning Objectives

At the end of this lecture, the student will be able to:


• Acids

• Bases and

• Salts based on various theories

History of Acids and Bases

In the early days of chemistry chemists were organizing physical and chemical properties of substances. They discovered that many substances could be placed in two different property categories:

Substance A

Substance B

1. Sour taste

1. Bitter taste

2. Reacts with carbonates to make CO2

2. Reacts with fats to make soaps

3. Reacts with metals to produce H2

3. Do not react with metals

4. Turns blue litmus pink

4. Turns red litmus blue

5. Reacts with B substances to make salt and water

5. Reacts with A substances to make salt and  water

6. pH < 7

6. pH >7


Arrhenius was the first person to suggest a reason why substances are in A or B due to their ionization in water

Arrhenius Theory

The Swedish chemist Svante Arrhenius proposed the first definition of acids and bases (1887)

According to the Arrhenius model:

“Acids are substances that dissociate in water to produce H+ ions and bases are substances that dissociate in water to produce OH- ions”

NaOH (aq) à Na+ (aq) + OH-(aq)  Base

HCl (aq) à H+ (aq) + Cl-(aq)  Acid

Arrhenius theory: Neutralization reactions

• Arrhenius acids and bases react with each other to form water and aqueous salts in neutralization reactions

H+ (aq) + A-(aq) + M+ (aq) + OH-(aq) à H2O (l) + M+ (aq) + A-(aq)

• The net ionic equation is

H+ (aq) + OH- (aq) à H2O (l)

Classification of acid and base based on an Arrhenius concept:



Strong acid

Strong base

weak acid

Weak base

Mono basic acid

Mono acidic base

Dibasic acid

Di acidic base

Tribasic acid

Tribasic base


Limitations of Arrhenius concept

• The definitions are only in terms of aqueous solution and not in terms of substance

• The theory is not able to explain acidic and basic properties of substances in non-aqueous solvents

Example: Ammonium nitrate in liquid ammonia acts as an acid though it does not give H+ ions

• The basic nature of substances like ammonia or sodium carbonate which do not contain OH- ions was not explained by this concept

• The acidic nature of carbon di oxide, sulphur di oxide which do not contain H+ ions was not explained by this concept

• The neutralisation of acid and base in absence of solvent could not be explained

Bronsted Lowry Theory (1923)

Johannes Bronsted and Thomas Lowry revised Arrhenius’s acid-base theory to include this behavior. They defined acids and bases as follows:

“An acid is a hydrogen containing species that donates a proton. A base is any substance that accepts a proton”

HCl (aq) + H2O (l) à Cl-(aq) + H3O+(aq)

In the above example what is the Brønsted acid? What is the Brønsted base?

In reality, the reaction of HCl with H2O is an equilibrium and occurs in both directions, although in this case the equilibrium lies far to the right.

HCl (aq) + H2O (l)   à Cl-( aq) +   H3O+ (aq)

For the reverse reaction Cl- behaves as a Bronsted base and H3O+ behaves as a Bronsted acid.

The Cl- is called the conjugate base of HCl. Bronsted acids and bases always exist as conjugate acid-base pairs.

Bronsted-Lowry Theory of Acids & Bases Conjugate Acid-Base Pairs

General Equation

Bronsted-Lowry Theory of Acids & Bases: Example

Notice that water is both an acid & a base = amphoteric

Bronsted-Lowry Theory of Acids & Bases Proton transfer reactions

• Pairs of compounds are related to each other through Bronsted-Lowry acid-base reactions. These are conjugate acid-base pairs.

• Generally, an acid HA has a conjugate base A- (a proton hastransferred away from the acid). Conversely, a base B has a conjugate acid BH+ (a proton has transferred toward the base).


Bronsted acid:

Mono protic acid: Capable of donating one proton only

Example: HF, CH3COOH

Poly protic acid: Capable of donating more than one proton

Example: H2S, H2O

Bronsted Base:

Mono protic base: Which can accept one proton

Example: water

Poly protic base: Which can accept two or more proton

Example: Sulphate ion, Phosphate ion

Lewis acids and bases

Gilbert Newton Lewis (1875-1946) influential American chemist. His theories include the Lewis dot structure taught in Chem120 and covalent bond theories.

Lewis acids are electrophils: H+, Na+, BF3,

Lewis bases are nucleophils: NH3, H2O, PH3

Acid base reactions:

BF3 + :NH3 à F3B:NH3

Lewis acids and bases

In general LA + :LB à LA-LB

Lux flood theory

• It is applicable only for oxide system

According to this theory:

• A base is an oxide ion donor and

• An acid is an oxide ion acceptor

Example: CaO + SiO2 → CaSiO3

Uranvonish theory

This theory is given by Russian chemist: Uranvonish in 1939

• An acid is a substance which reacts with a base to give up a cation

• A base is a substance that react with an acid giving up an anion and accepts cation or electron

Example: SO3-- + Na2O → Na2SO4



• Acids are sour tasting

• Arrhenius acid:  Any substance that, when dissolved in water, increases the concentration of hydronium ion (H3O+)

• Bronsted-Lowry acid:  A proton donor

• Lewis acid:   An electron acceptor


• Bases are bitter tasting and slippery

• Arrhenius base:  Any substance that, when dissolved in water, increases the concentration of hydroxide ion (OH-)

• Bronsted-Lowery base:  A proton acceptor

• Lewis acid:   An electron donor

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