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Arrhenius Acids and Bases

Arrhenius Acids and Bases: Definition, Examples & Facts

The Arrhenius definition of acids and bases, developed by Swedish chemist Svante Arrhenius, is a fundamental concept in chemistry.

This early theory provides insights into the behavior of acids and bases in aqueous solutions. Understanding the Arrhenius definition lays the groundwork for comprehending acid-base reactions and their impact on various chemical processes.

Arrhenius Acids and Bases

  • Acids: Acids are substances that, when dissolved in water, increase the concentration of hydrogen ions (H+) in the solution. In simpler terms, acids donate protons to the solution.
  • Bases: Bases are substances that, when dissolved in water, increase the concentration of hydroxide ions (OH−) in the solution. Bases accept protons from the solution.

By exploring this concept, we can delve deeper into the fascinating world of chemistry and gain a better understanding of how substances interact in solution.

Examples of Arrhenius acids and bases:

Hydrochloric Acid (HCl)

Hydrochloric acid is a common example of an Arrhenius acid. When it is dissolved in water, it breaks apart into hydrogen ions (H+) and chloride ions (Cl-). This dissociation allows the acid to conduct electricity in solution.

Sodium Hydroxide (NaOH)

Sodium hydroxide is an example of an Arrhenius base. When sodium hydroxide dissolves in water, it dissociates into sodium ions (Na+) and hydroxide ions (OH-). The presence of hydroxide ions makes the solution basic.

Sodium Hydroxide (NaOH)

Ammonia (NH3)

Ammonia can act as both an Arrhenius base and a weak acid depending on the reaction conditions.

In the presence of strong acids, ammonia acts as a base by accepting a proton to form ammonium ions (NH4+).

On the other hand, in the presence of strong bases, ammonia can donate a proton to form amide ions (NH2-).

These examples demonstrate how different substances can exhibit acidic or basic properties according to the Arrhenius definition. It’s important to note that there are many other examples of Arrhenius acids and bases beyond those mentioned here.

Arrhenius theory: acids and bases explained

The Arrhenius theory provides an explanation for the behavior of acids and bases.

According to this theory, acids are substances that release hydrogen ions (H+) when dissolved in water.

On the other hand, bases are substances that release hydroxide ions (OH-) when dissolved in water.

The presence of excess H+ or OH- ions determines whether a solution is acidic, basic, or neutral. When an acid dissolves in water, it dissociates into its constituent ions, releasing H+ ions into the solution. This process is known as acid dissociation.

Similarly, when a base dissolves in water, it undergoes a reaction that generates hydroxide ions (OH-). This reaction is called base ionization or base hydrolysis. The released OH- ions contribute to the alkaline nature of the solution.

To illustrate this concept further, let’s consider the example of hydrochloric acid (HCl), which is a strong acid. When HCl dissolves in water, it completely dissociates into H+ and Cl- ions according to the following balanced chemical equation:

HCl(aq) → H+(aq) + Cl-(aq)

On the other hand, sodium hydroxide (NaOH), a strong base, releases OH- ions upon dissolution:

NaOH(aq) → Na+(aq) + OH-(aq)

When an acid and a base react with each other in an aqueous solution, they undergo a neutralization reaction. During this process, H+ from the acid combines with OH- from the base to form water molecules:

H+(aq) + OH-(aq) → H2O(l)

This reaction results in the formation of salts as well. Salts are ionic compounds composed of positive and negative ions derived from both the acid and base involved.

Limitations of the Arrhenius theory:

Only Applicable to Aqueous Solutions

The Arrhenius theory, although widely accepted, has its limitations.

One major drawback is that it only applies to aqueous solutions, meaning solutions in which water is the solvent.

This theory fails to account for non-aqueous acid-base reactions, where other solvents are used instead of water. So if you’re dealing with acids and bases outside of a watery environment, the Arrhenius theory might not be your go-to.

Fails to Explain Non-Ionic Reactions

Another limitation of the Arrhenius theory is its inability to explain why some substances can exhibit acidic or basic properties without releasing H+ or OH- ions.

In other words, this theory doesn’t cover cases where substances show acid or base behavior without any presence of traditional acid or base ions.

So if you come across a substance that seems acidic or basic but doesn’t fit the Arrhenius definition, don’t scratch your head too hard – it’s just one of those things this theory can’t explain.

Ignores Proton Transfer

Proton transfer plays a crucial role in many acid-base reactions.

However, the Arrhenius theory fails to address this concept adequately.

It focuses primarily on the release and acceptance of H+ and OH- ions but overlooks the actual transfer of protons between species involved in an acid-base reaction. This oversight limits our understanding of certain complex reactions where proton transfer is key.

Comparing Arrhenius, Bronsted-Lowry, and Lewis theories






 According to Arrhenius :Acid produces H+ ions in aqueous solution Acid: donates a proton (H+) to a base Acid: accepts an electron pair from a base
Base: produces OH- ions in aqueous solution According to Lowry Base accepts a proton (H+) from an acid Base: donates an electron pair to an acid


HCl -> H+ + Cl- HCl + H2O -> H3O+ + Cl- BF3 + NH3 -> BF3NH3


According to this theory Acid donates H+ : Base: OH- donor Acid: H+ donor, Base: H+ acceptor Acid: electron pair acceptor, Base: electron


pair donor


It applies  only to aqueous solutions Only applies to proton transfer reactions It applies to electron pair transfer
reactions reactions


Lays foundation for modern acid-base theories Explains acid-base reactions in solvents other  It Explains acid-base reactions in non-aqueous
than water media

The Arrhenius, Bronsted-Lowry, and Lewis theories are all important in understanding acid-base chemistry. However, they differ in their scope and focus. Let’s take a closer look at how these theories compare.

Arrhenius Theory

  • The Arrhenius theory mainly focuses on ionization in aqueous solutions.
  • It describes acids as substances that produce hydrogen ions (H+) when dissolved in water, and bases as substances that produce hydroxide ions (OH-) when dissolved in water.
  • This theory is limited to reactions that occur specifically in water.

Bronsted-Lowry Theory

  • Unlike the Arrhenius theory, the Bronsted-Lowry theory provides a broader understanding of acid-base chemistry beyond just aqueous solutions.
  • According to this theory, an acid is a substance that donates a proton (H+), while a base is a substance that accepts a proton.
  • In this theory, acids and bases can exist outside of water and can react with each other through proton transfer.

Lewis Theory

  • The Lewis theory expands upon the concepts of acids and bases even further.
  • According to this theory, an acid is defined as an electron pair acceptor, while a base is defined as an electron pair donor.
  • Unlike the previous two theories, the Lewis theory does not require the presence of protons or hydroxide ions for acid-base reactions to occur.

Advantages of the Bronsted-Lowry theory

The Bronsted-Lowry theory takes the Arrhenius definition a step further by considering the transfer of protons between acids and bases.

This expanded perspective provides us with a more comprehensive understanding of acid-base reactions in various environments, whether they are aqueous or non-aqueous.

Comprehensive Understanding of Acid-Base Reactions:

Unlike the Arrhenius definition, which focuses solely on aqueous solutions, the Bronsted-Lowry theory allows us to grasp acid-base reactions in both aqueous and non-aqueous systems.

It recognizes that proton transfer can occur not only between substances dissolved in water but also within other solvents or even gases.

This broader scope enables scientists to analyze a wider range of chemical reactions and their underlying principles.

Conjugate Acid-Base Pairs:

One key concept introduced by the Bronsted-Lowry theory is that of conjugate acid-base pairs. According to this theory, an acid donates a proton (a hydrogen ion) to a base, forming its conjugate base.

Similarly, a base accepts a proton, becoming its conjugate acid. This notion helps explain important phenomena such as buffer systems and equilibrium reactions.

By understanding how acids and bases can transform into their respective conjugates through proton transfer, we gain insight into how these systems maintain stability and balance.

Buffer solutions, for example, consist of a weak acid and its conjugate base (or vice versa), allowing them to resist changes in pH when small amounts of acid or base are added.

Relationship between Arrhenius and Lewis acid-base theories

The Arrhenius theory and the Lewis theory provide different perspectives on acid-base behavior. The Arrhenius theory focuses on ionization, whereas the Lewis theory defines acids as electron pair acceptors and bases as electron pair donors.

Both theories can be applied to non-aqueous systems, offering complementary insights into acid-base interactions. While the Arrhenius theory primarily applies to aqueous solutions, the Lewis theory encompasses a wider range of substances that can exhibit acidic or basic properties.

Here’s a breakdown of how these two theories differ:

Arrhenius Theory:

  • Defines an acid as a substance that produces hydrogen ions (H+) in water.
  • Defines a base as a substance that produces hydroxide ions (OH-) in water.
  • Limited to aqueous solutions and does not consider other solvents or non-aqueous systems.

Lewis Theory:

  • Defines an acid as a species that accepts an electron pair.
  • Defines a base as a species that donates an electron pair.
  • Applicable to both aqueous and non-aqueous systems, providing a broader understanding of acid-base reactions.


The Arrhenius definition of acids and bases is a cornerstone concept in chemistry, offering valuable insights into the behavior of substances in solution.

By understanding the characteristics and properties of acids and bases, scientists and researchers can make significant advancements in various fields, contributing to the betterment of society and the environment.

As we continue to delve deeper into the realms of chemistry, the Arrhenius definition remains a guiding principle, enlightening our path toward new discoveries and innovations.


What are some common examples of Arrhenius acids?

Arrhenius acids include hydrochloric acid (HCl), sulfuric acid (H2SO4), and nitric acid (HNO3). These are all strong acids that readily dissociate in water to release hydrogen ions (H+).

Can you give an example of an Arrhenius base?

Sodium hydroxide (NaOH) is a common example of an Arrhenius base. When dissolved in water, it dissociates into hydroxide ions (OH-) which can accept protons.

What are the limitations of the Arrhenius theory?

The primary limitation of the Arrhenius theory is that it only defines acids as substances that release hydrogen ions (H+) in water and bases as substances that release hydroxide ions (OH-) in water. It does not account for non-aqueous solutions or reactions where no water is involved.

How does Bronsted-Lowry theory differ from Arrhenius theory?

While both theories define acids and bases, the Bronsted-Lowry theory focuses on proton transfer. According to this theory, an acid is a proton donor, and a base is a proton acceptor. Unlike the Arrhenius theory, Bronsted-Lowry theory extends beyond aqueous solutions.

What is the relationship between Arrhenius and Lewis acid-base theories?

The Lewis acid-base theory expands upon the Arrhenius definition by considering electron pair donation and acceptance. According to the Lewis theory, an acid is an electron pair acceptor, while a base is an electron pair donor. This broader perspective allows for the classification of substances that do not fit into the Arrhenius or Bronsted-Lowry definitions.


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