Acids and bases are the foundation of chemistry — from the sour taste of lemons to the slippery feel of soap, their reactions shape our daily life. But have you ever wondered why some substances act as acids and others as bases? Scientists have explained this behavior using three famous theories: the Arrhenius, Bronsted–Lowry, and Lewis concepts.
In this detailed blog by Learn Chemistry by Inam Jazbi, you’ll learn the 3 main concepts of acids and bases with easy definitions, examples, and a comparison table — all explained in a way that’s simple, clear, and perfect for MDCAT, FSc, and O-Level students.
Different concepts
of Acids and Bases
1. The
Arrhenius concept
2. Bronsted-Lowry
concept
3. Lewis concept
Acids, bases and salts are three distinct classes in which almost all the organic and inorganic compounds are classified. A famous Muslim Chemist Jabir Bin Hayan prepared nitric acid (HNO3), hydrochloric acid (HCl) and sulphuric acid (H2SO4). In 1787, Lavoisier named binary compounds of oxygen such as CO2 and SO2 as acids which on dissolution in water gave acidic solutions. Later on in 1815, Sir Humphrey Davy discovered that there are certain acids which are without oxygen, e.g. HCl. Davy proved the presence of hydrogen as the main constituent of all acids. It was also discovered that all water soluble metallic oxides turn red litmus blue, which is a characteristics of bases. The word acid is derived from the Latin word ‘Acidus’ meaning sour. The first acid known to man was acetic acid, i.e., in the form of vinegar.
We all have
a little concentration of hydrochloric acid in our stomach, which helps to
break down the food. Sometimes, the amount of stomach acid becomes too much,
which causes ‘acidity’. This uncomfortable feeling is easily treated by taking
an alkaline medicine. The alkali neutralizes the acid, producing a harmless
chemical called a salt.
The Arrhenius Theory
Svante-Arrhenius, a Swedish scientist in 1887 first defined acids and bases on their ionic dissociation in water in his theory of ionization. According to Arrhenius theory, acids are substances that dissociates in water to give hydrogen ions H+(aq) and bases are substances that produce hydroxyl ions OH–(aq). A salt is an ionic compound that is formed by the reaction of an acid and a base. Such a reaction is called neutralization.
Examples of Some Important Arrhenius Acids and bases
Bases: Sodium hydroxide (NaOH), Potassium hydroxide, KOH, Calcium
hydroxide, Ca(OH)2, Al(OH)3
Acids: Hydrochloric acid, HCI, Nitric acid, HNO3 Sulphuric acid, H2SO4 Phosphoric acid, H3PO4
Arrhenius
Definition of Acid
An acid is a substance (such as HCl, HNO3, HCN, CH3COOH etc.), which dissociates in aqueous solution to yield hydrogen ions or protons (H+). Properties of acids are due to presence of hydrogen ions (H+)
general ionization
The general ionization of an acid HY(aq) can be represented by the following equations.
Examples of Arrhenius Acids
Halogen acids (HX; HCl, HBr, HI, HF), nitric acid, sulphuric acid, phosphoric acid, acetic acid etc. produce H+ ions when dissolved in water and thus they are typical acids (but H3BO3 is not an Arrhenius acid). [Thus all acids contain hydrogen but not all hydrogen- containing substances are acids].
(A bare proton, H+ is very
reactive and cannot exist freely in aqueous solutions. Thus, it bonds to the
oxygen atom of a solvent water molecule to give trigonal pyramidal hydronium ion, H3O+ {[H(H2O)]+}.
Hence H+(aq) and H3O+(aq)
are used interchangeably to mean the same i.e. a hydrated proton).
Arrhenius
Definition of Base
A base is a substance such as NaOH, KOH, NH4OH, Ca(OH)2 etc. that gives hydroxide (OH–) ions in aqueous solution. Properties of bases are due to presence of hydroxide ions (OH–)
general ionization
The general ionization of a base MOH(aq) (BOH) can be represented by the following equations.
Examples of Arrhenius Bases
Sodium hydroxide, Potassium hydroxide, Ammonium hydroxide are considered
as bases as they furnish OH– ion in water. (The hydroxyl ion also
exists in the hydrated form in the aqueous solution).
Limitations of Arrhenius Concept/Drawbacks of Arrhenius Concept
1. It is only applicable in aqueous solution.
One of the limitations of Arrhenius’s definition of acids and bases is that it is restricted to aqueous solution. This concept is applicable only in aqueous medium and does not explain nature of acids and bases in non-aqueous medium.
2. It also failed to account for the basicity or acidity of substances that do not contain H+ (CO2) or OH– ions (e.g. NH3) i.e. It does not explain the basicity of ammonia (NH3), acidity of carbon dioxide (CO2) & other similar compounds.
According to this concept, acids and bases are only those compounds which contain hydrogen (H+) and hydroxide (OH–) ions, respectively. It can’t explain the nature of compounds like CO2, NH3, etc. which are acid and base, respectively.
3. Hydrogen ions do not exist in water and they react with water to form hydronium ions (H3O+).
Although
this concept has limited scope yet, it led to the development of more general
theories of acid-base behaviour.
Bronsted- Lowry Theory [Proton-donor and acceptor Theory/Proton transfer theory]
Hydronium and Hydroxyl Ions
Hydrogen ion by itself is a bare proton with extremely small size (~10–15 m or 10–13 cm radius, comparing other ions which have diameters of the order of 10–8 cm) having intense electric field and very high charge density. Therefore, hydrogen ion cannot exist in water or any other solvent.
Hydrogen ion binds itself with the water molecule at one of the two available lone pairs on it giving H3O+ ion which is called hydroxonium or hydronium or oxonium ion. This species has been detected in many compounds (e.g. H3O+Cl–) in the solid state. In aqueous solution the hydronium ion is further hydrated to give species like
Similarly the hydroxyl ion is hydrated
to give several ionic species like H3O2–, H5O3–
and H7O4– etc.
Introduction
A broader, more general and Protonic definition of acids and bases were proposed independently by the Danish Chemist Johannes Bronsted and the English Chemist Thomas M. Lowry in 1923 on the basis of proton-transfer.
Definition
According to Bronsted-Lowry theory, an acid is a
substance that is capable of donating a hydrogen ion (H+) or proton to another substance
and a base is a substance capable of accepting a hydrogen ion (H+)
or proton from another substance. Hence, In short, acids are proton donors and bases are proton acceptors. For example, HCl acts as an acid while NH3 acts as a base. All
Arrhenius acids are Bronsted acids but but except OH− other Bronsted bases are not Arrhenius bases.
OR
In other words, an acid is a substance that gives
oxonium ion or hydroxonium ion (H3O+) in aqueous
solution by donating H+ ions. a
base is a substance that combines or adds with proton of oxonium ion (H3O+)
or conjugate acid accepting or removing a proton forming water water or OH-
ion i.e. accepts).
According to this concept, acid-base pairs i.e. proton donor and proton acceptor must co-exist. Thus an acid and a base always work together to transfer a proton.
Stated differently, no substance can act as an acid in solution unless a base
is present to accept a proton i.e. the reactions of acids are reaction between acids and bases. Similarly, all reactions of bases in solution are
acid-base reactions. The
products of acid-base reaction are themselves acids and base which are called
conjugate acids and conjugate base respectively.
amphoteric or amphiprotic
substances
Some substances can act as proton donor as well as proton abstractor and are called amphoteric or amphiprotic. i.e. they
can behave as an acid, as well as, a base e.g. H2O,
According to Bronsted-Lowry concept, an acid and a base always work together to transfer a proton. That means, a substance can act as an acid (proton donor) only when another substance simultaneously behaves as a base (proton acceptor). Hence, a substance can act as an acid as well as a base, depending upon the nature of the other substance.
General Example
Consider the dissociation reaction of a general monoprotic acid HA in water in which HA dissolves in water in a reversible manner by donating
a proton to water. Therefore, HA is the Bronsted-Lowry acid and H2O is the Bronsted-Lowry base. It is seen that the products of
acid-base reaction (H3O+ and A-) are
themselves acids and bases which are called conjugate acid and conjugate
base respectively.
In
each acid-base reaction an acid reacts with a base to give a conjugate base and
a conjugate acid.
Example # 1 of Bronsted-Lowry
Concept
When hydrogen chloride is dissolved in
water, a reversible reaction takes place called ionization. In this reaction,
HCl donates its one proton to water acting as a Bronsted acid giving Cl– ion while water accepts that
one proton acting as a Bronsted base forming hydronium (H3O+)
ion. H3O+ is called conjugate acid and Cl– is a conjugate base. The products of this reaction are
themselves acid and base which are called conjugate acid and conjugate base
respectively].
Example # 2 of Bronsted-Lowry
Concept
Consider the
dissolution of acetic acid (CH3COOH) in water. In the forward
reaction, acetic acid is a Bronsted acid as it donates its one proton to water
while water is a Bronsted base as it accepts a proton. Like this we have pairs
of conjugated acids-base pairs. Conjugate acid is formed by accepting a proton
by a base and conjugate base is produced by donating a proton by an acid.
Example # 3 of Bronsted-Lowry
Concept
Example # 4 of Bronsted-Lowry
Concept
Conjugate Acid-Base Pair
Acids and bases occur as conjugate acid-base pair (the word conjugate means “joined together or tie together as a pair”) which are defined as
an acid and a base that differ only in the presence or absence of a proton or pair of acid and base that are related to each other by loss or gain of a proton.
Conjugate Base
Every acid has a conjugate base which is the negatively charged or neutral specie formed by the removal or release of a proton from the acid. A conjugate base is a species that results when an acid loses a proton.
Conjugate Acid
Every base is associated with a conjugate acid which is the positively charged ion produced by the acceptance or addition of a proton by a base. The species that results when a base accepts a proton from an acid is called the conjugate acid.
1. According to this concept, proton donor and proton acceptor must co-exist. But there are many reactions in which this does not happen.
2. It also could not explain the behaviour of those acids and bases which do not contain hydrogen at all having no tendency to lose H+ ions i.e. it does not explain acidic behaviour of aprotic acids like SO2, SO3, CO2, AlCl3, SiCl4 etc. e.g. SO3 behave as an acid although it does not have the ability to donate a proton. CaO behaves as a base but it cannot accept a proton.
3. It could
not explain the basic nature compounds having OH- ions e.g. NaOH,
KOH, Ca(OH)2
Lewis Concept
(Electronic Concept)
Significance
AlCl3, BF3, BCl3, ZnCl2, FeCl3 are considered as acids although they do not have hydrogen. In 1923, an American Chemist G.N. Lewis proposed a more general and broader electronic concept of acids and bases focusing on electron transfer instead of proton transfer. Lewis concept can be applicable to non-aqueous solutions or solutions lacking hydrogen ions and reactions that do not involve hydrogen ions at all. The Lewis definition is much more useful than the others because it can be applied to all species and reaction. The other definitions are only useful for species and reactions involving H+.
Lewis Definition of Acid and Base
In the Lewis theory of
acid-base reactions, bases donate pairs of electrons and
acids accept pairs of electrons. A Lewis acid is any substance, such as the H+ ion,
that can accept a pair of nonbonding electrons i.e. a
Lewis acid is an electron-pair acceptor.
A Lewis base is
any substance, such as the OH- ion, that can donate a pair of non-bonding electrons i.e. A Lewis
base is an electron-pair donor.
A Lewis acid must have a vacant orbital into which it can accept the electron pair. H+ is a Lewis acid. A Lewis acid-Lewis base reaction gives Lewis adduct. Lewis acid-base reactions in general do not have to involve H+
Lewis Definition of Acid
An acid is any species (molecule or ion) which can accept a (lone) pair of electrons during a reaction i.e. Lewis Acids are an electron-pair acceptor. They are also called Electrophiles (meaning electron loving) and have vacant orbitals into which it can accept the electron pair.
Lewis acids include not only H+ (protons) or H3O+ (oxonium ion) but also other cations and neutral molecules having vacant valence orbitals like AlCl3, AlBr3, BF3, BCl3, ZnCl2, FeCl3 etc.
Example # 1 of Lewis Acid and Base
Consider the reaction between proton and water. The proton (H+) has a tendency to accept a pair of electrons while H2O has a tendency to donate a pair of electrons to form coordinate covalent bond or
donor-acceptor bond. Hence in Lewis concept, H+ is a Lewis acid and H2O is a Lewis base.
Example # 2 of Lewis Acid and Base
H+ ion acts as Lewis acid because it is short of two electrons for completion of its duplet and thus capable of accepting a lone pair of electrons from N of ammonia which is a Lewis base making co-ordinate covalent bond.
Example # 3 of Lewis Acid and Base
Consider the reaction between Nh3
and BF3. Here NH3 acts as an electron pair donor and
hence it is Lewis base while BF3 acts as an electron pair acceptor
and hence it is Lewis acid and they combine to form adduct through coordinate
covalent bond.
Lewis Definition
of Base
A base is any species (molecule or ion) which can donate a (lone) pair of electrons during a reaction i.e. Lewis bases are electron pair donor. They are also called Nucleophiles (meaning nucleus loving) and have lone pair of electrons.
Lewis bases include not only OH– ions but also all anions (F–, Cl– etc) and neutral molecules having lone pair of electrons (NH3, C2H5OH
etc.). Ammonia is base in all three concepts.
Example # 5 of Lewis Acid and Base
ammonia is a Lewis base as it donates its an electron pair to boron of
boron trifluoride lacking a pair of electrons to complete its octet acting as
Lewis acid to form co-ordinate covalent bond.
Example # 6 of Lewis Acid and Base
Chloride ion is a Lewis base as it donates its an electron pair to aluminium
of aluminium trichloride lacking a pair of electrons to complete its octet
acting as Lewis acid to form co-ordinate covalent bond.
Types of Lewis Acids
1. All Metals cations e.g. Li+,
Ag+, Al3+, Mg2+ etc. and proton
itself act as Lewis acids.
2.
Molecules having incomplete octet e.g. BF3,
BCl3, AlCl3, B(OH)3 or H3BO3,
etc.
3.
Molecules having multiple bonds between
atoms of different E.N e.g. CO2, SO2, SO3 etc.
4. Molecules having vacant d-orbitals e.g. FeCl3, SF4, SF6, SnCl2, SnCl4 etc.
(i) Simple cations can act as Lewis acids. All cations act as
Lewis acids since they are deficient in electrons. However, cations such as Na+,
K+, Ca2+,
(ii) Molecules in which the central atom
has incomplete octet are always Lewis acids. For example, in BF3,
AICl3, FeCl3, the central atoms have only six electrons
around them, therefore, these can accept an electron pair.
Types or Examples of Lewis Bases
1. All anions e.g. chloride (Cl-), cyanide (CN-), hydroxide
(OH-), O2-,
2. Neutral
molecules having lone
(unshared) pair of electrons e.g. NH3, H2O , R-NH2, R2NH,
ROR etc.
3. pi-electrons containing Molecules being electron rich are Lewis bases e.g. alkenes,
alkynes, benzene
Neutral species having at least one lone pair of
electrons act as Lewis bases. For example, ammonia, amines, alcohols etc. act
as Lewis bases as they contain a lone pair of electrons:
The
cations (proton itself or metal ions) act as Lewis acids.
The product
of any Lewis acid-base reaction is a single specie, called an adduct. So, a
neutralization reaction according to Lewis concept is donation and acceptance
of an electron pair to form a coordinate covalent bond in an adduct.
Acids are electron pair acceptors while bases are electron pair donors. Thus, it is evident that any substance which has an unshared pair of electrons can act as a Lewis base while a substance which has an empty orbital that can accommodate a pair of electrons acts as Lewis acid.
All the
Lewis bases are Bronsted bases but all the Lewis acids are not Bronsted acids
It may be noted that all Bronsted bases are also Lewis
bases but all Bronsted acids are not Lewis
acids.
According to Bronsted concept, a base is a substance which can accept a proton,
while according to Lewis concept, a base is a substance which can donate a pair
of electrons. Lewis bases generally contain one or more lone pair of electrons
and therefore, they can also accept a proton (Bronsted base). Thus, all Lewis
bases are also Bronsted bases. On the other hand, Bronsted acids are those
which can give a proton. For example, HCI, H2SO4 are not
capable of accepting a pair of electrons. Hence, all Bronsted acids are not
Lewis acids.
🧪 Summary of 3 Powerful Concepts of Acids and Bases Explained Simply | Learn Chemistry by Inam Jazbi
🔍 Introduction
Acids and bases are among the most fascinating topics in chemistry. From lemon juice to soaps, almost every chemical reaction around us depends on whether a substance behaves as an acid or a base. Scientists have proposed three main concepts to define acids and bases: Arrhenius, Bronsted–Lowry, and Lewis theories. Understanding these gives you a complete picture of acid–base behavior in both aqueous and non-aqueous systems.
⚗️ 1. Arrhenius Concept of Acids and Bases
📘 Definition
Proposed by Svante Arrhenius (1884), this theory applies to aqueous solutions only.
- Acid: A substance that produces H⁺ ions in water.
- Base: A substance that produces OH⁻ ions in water.
🧩 Examples
- HCl (aq) → H⁺ + Cl⁻ → acid
- NaOH (aq) → Na⁺ + OH⁻ → base
💡 Limitations
- Works only in aqueous media.
- Cannot explain acid–base behavior in non-aqueous solvents.
- Fails to explain why NH₃ acts as a base (no OH⁻ ions).
⚗️ 2. Bronsted–Lowry Concept of Acids and Bases
📘 Definition
Proposed by Johannes Bronsted and Thomas Lowry (1923), this concept defines acids and bases in terms of proton transfer.
- Acid: Proton (H⁺) donor
- Base: Proton (H⁺) acceptor
🧩 Example
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
- NH₃ → Base (accepts H⁺)
- H₂O → Acid (donates H⁺)
- NH₄⁺ and OH⁻ form a conjugate acid–base pair
💡 Key Points
- Explains acid–base behavior in non-aqueous systems.
- Introduces the idea of conjugate acid–base pairs.
- Limited to reactions involving proton transfer.
⚗️ 3. Lewis Concept of Acids and Bases
📘 Definition
Proposed by G. N. Lewis (1923), this theory focuses on electron pair transfer rather than protons.
- Lewis Acid: Electron pair acceptor
- Lewis Base: Electron pair donor
🧩 Examples
- NH₃ + BF₃ → H₃N–BF₃
- NH₃ donates an electron pair → Lewis base
- BF₃ accepts an electron pair → Lewis acid
- H⁺ + :OH⁻ → H₂O
- H⁺ = Lewis acid
- OH⁻ = Lewis base
💡 Importance
- Explains reactions that do not involve H⁺.
- Most general and widely applicable acid–base concept.
- Used in organic and coordination chemistry.
📊 Comparison Table of the Three Concepts
| Feature | Arrhenius | Bronsted–Lowry | Lewis |
|---|---|---|---|
| Basic Idea | Ionization in water | Proton transfer | Electron pair transfer |
| Medium | Only aqueous | Aqueous & non-aqueous | Any medium |
| Acid | H⁺ producer | Proton donor | Electron pair acceptor |
| Base | OH⁻ producer | Proton acceptor | Electron pair donor |
| Examples | HCl, NaOH | NH₃, H₂O | BF₃, NH₃ |
🧠 Quick Revision Tips
- Arrhenius – Ions
- Bronsted–Lowry – Proton
- Lewis – Electrons
- Arrhenius is simplest, Lewis is most general.
- Practice examples from inorganic & organic chemistry.
💧 MDCAT Style MCQs on Acids, Bases, and Salts | Learn Chemistry by Inam Jazbi
Test your knowledge of acids, bases, and salts with these MDCAT-style multiple choice questions. Each question includes a brief explanation for better understanding.
1. Which of the following is an Arrhenius acid?
A) NaOH
B) NH₃
C) HCl
D) K₂CO₃
Explanation: Arrhenius acids release H⁺ ions in water. For example, HCl → H⁺ + Cl⁻.
2. Which of the following acts as a Bronsted–Lowry base?
A) H₂SO₄
B) NH₃
C) HCl
D) H₂O
Explanation: A Bronsted–Lowry base is a proton (H⁺) acceptor. NH₃ accepts H⁺ from water to form NH₄⁺.
3. Which of the following is a Lewis acid?
A) NH₃
B) BF₃
C) OH⁻
D) H₂O
Explanation: A Lewis acid accepts an electron pair. BF₃ can accept an electron pair from NH₃.
4. Which salt is formed when HCl reacts with NaOH?
A) KCl
B) NaCl
C) NH₄Cl
D) Na₂SO₄
Explanation: HCl + NaOH → NaCl + H₂O. Neutralization produces a salt and water.
5. Which of the following pairs is a conjugate acid–base pair?
A) H₂SO₄ / SO₄²⁻
B) NH₄⁺ / NH₃
C) HCl / Cl₂
D) NaOH / H₂O
Explanation: A conjugate acid–base pair differs by one proton. NH₄⁺ donates H⁺ to form NH₃.
6. Which of the following is a basic salt?
A) NaCl
B) CH₃COONa
C) NH₄Cl
D) KNO₃
Explanation: CH₃COONa is formed from a weak acid (CH₃COOH) and a strong base (NaOH), making it basic.
7. Which solution has a pH less than 7?
A) NaOH
B) H₂SO₄
C) Na₂CO₃
D) KOH
Explanation: Acids have pH < 7. H₂SO₄ is a strong acid producing many H⁺ ions.
8. Which salt solution is acidic?
A) NH₄Cl
B) Na₂CO₃
C) KCl
D) NaCl
Explanation: NH₄Cl is formed from a weak base (NH₃) and strong acid (HCl), so its solution is acidic.
9. Which of the following acids is diprotic?
A) HCl
B) H₂SO₄
C) HNO₃
D) CH₃COOH
Explanation: Diprotic acids can donate 2 protons. H₂SO₄ → 2H⁺ + SO₄²⁻.
10. What is the product of neutralization between HNO₃ and KOH?
A) K₂SO₄
B) KNO₃
C) NH₄NO₃
D) NaNO₃
Explanation: Acid + Base → Salt + Water. HNO₃ + KOH → KNO₃ + H₂O.
✨ Quick Tip by Inam Jazbi
Remember: Acids donate H⁺, Bases accept H⁺, and Salts form from neutralization. Practice these MCQs to strengthen your MDCAT chemistry preparation.