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

Acid Base Theory


Theories on:

• Acids

• Bases and

• Salts


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


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

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

• 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

Lowry Theory (1923)

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

“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+

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

The Cl- is called the conjugate
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
H+, Na+, BF3,

Lewis bases are
NH3, H2O, PH3

Acid base reactions:

BF3 + :NH3 à

Lewis acids and bases

In general LA + :LB à

Lux flood

• It is applicable only for oxide system

According to this

• A base is an oxide ion donor and

• An acid is an oxide ion acceptor

Example: CaO + SiO2 → CaSiO3


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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