Short Answer

The main limitation of the Arrhenius theory is that it explains acids and bases only in water solutions. It cannot describe how acids and bases behave in other solvents. The theory also fails to explain substances like ammonia (NH₃), which behaves like a base even though it does not release hydroxide ions directly.

Another limitation is that the theory gives a very narrow definition of acids and bases because it focuses only on the production of H⁺ and OH⁻ ions. Many acid–base reactions that do not involve these ions cannot be understood using Arrhenius theory.

Detailed Explanation :

Limitation of Arrhenius Theory

The Arrhenius theory of acids and bases, proposed in 1887, was the first scientific attempt to describe why acids and bases show their unique properties. Although it played a major role in developing modern chemistry, the theory has several limitations. These limitations arise mainly because the theory is simple and applies only to a limited number of situations. As chemistry developed further, scientists found that many substances and reactions could not be explained using Arrhenius ideas alone.

Understanding the limitations of Arrhenius theory helps us appreciate why new and more advanced theories, such as the Bronsted–Lowry and Lewis theories, were needed.

1. Applicable Only to Aqueous Solutions

The most important limitation is that Arrhenius theory applies only to substances dissolved in water.
According to Arrhenius:

• An acid must release H⁺ ions in water.
• A base must release OH⁻ ions in water.

This means that the theory cannot explain acid–base behavior in other solvents such as alcohol, ammonia, benzene, or liquid hydrogen fluoride. Many reactions take place outside water, and Arrhenius theory is unable to describe them.

For example:

• Hydrogen chloride (HCl) behaves as an acid in water but not in some other solvents.
• Some substances behave like bases in liquid ammonia even when they do not release OH⁻ ions.

This shows that acid–base behaviour is not limited to water alone.

2. Cannot Explain Behaviour of Substances Without OH⁻ Ions

According to Arrhenius, a base must release hydroxide ions.
However, some substances behave as bases even though they do not contain OH⁻ ions.

A common example is ammonia (NH₃).
Ammonia does not contain hydroxide ions, but:

• It turns red litmus blue
• It reacts with acids
• It accepts hydrogen ions (H⁺)

Arrhenius theory cannot classify ammonia as a base because it does not release OH⁻ ions directly. But in reality, ammonia is a well-known base. This is a major weakness because many important bases behave similarly.

3. Cannot Explain Acid–Base Reactions Without Ion Formation

Some acid–base reactions take place without the involvement of H⁺ or OH⁻ ions.
Arrhenius theory is unable to explain such reactions.

For example:

• In gaseous reactions, acids and bases can react even when no ions are present.
• Certain reactions between gases like hydrogen fluoride (HF) and ammonia occur without water.

Since Arrhenius theory requires ion formation in water, it fails to explain such reactions.

4. Narrow Definition of Acids and Bases

The theory defines:

• Acids only as H⁺ producers
• Bases only as OH⁻ producers

This is a very limited definition because:

• Many acids do not release free hydrogen ions directly
• Many bases do not release hydroxide ions directly
• Some substances can act as acids or bases without forming ions in water

Because of this narrow focus, the theory cannot describe amphoteric substances (like water), which can behave both as acids and bases.

5. Cannot Explain Strength of Acids and Bases Completely

Arrhenius theory only explains strength based on the number of ions released. But:

• The strength of acids and bases also depends on molecular structure
• Solvent properties affect ionisation, which the Arrhenius theory does not consider
• Acids like acetic acid ionise poorly in water but behave strongly in other solvents

Thus, the theory gives an incomplete picture of strength.

6. Does Not Consider Proton Transfer

Modern acid–base theories explain reactions using proton transfer (movement of H⁺ ions).
Arrhenius theory does not consider proton donors or acceptors. It only considers ion formation in water. Because of this, many acid–base reactions cannot be explained properly.

For example:

• Ammonia accepts a proton to form NH₄⁺, but Arrhenius cannot describe this behaviour.

Conclusion

The Arrhenius theory is simple and useful, but it has several limitations. It works only for water solutions, does not explain the behaviour of substances like ammonia, gives a narrow definition of acids and bases, and cannot describe many acid–base reactions outside aqueous systems. These limitations led to the development of more advanced theories such as the Bronsted–Lowry and Lewis theories, which offer a broader and more accurate understanding of acids and bases.