Examples of Arrhenius Bases

Oct 13, 2025 by talynt.org 78 views

Understanding the Difference Between Brønsted-Lowry and Arrhenius Bases

In chemistry, the Arrhenius and Brønsted-Lowry theories are two fundamental concepts used to describe the behavior of acids and bases. While both theories are essential in understanding chemical reactions, they differ in their definitions and applications. In this article, we will explore the difference between Arrhenius and Brønsted-Lowry bases and identify which example is a Brønsted-Lowry base but not an Arrhenius base.

What is an Arrhenius Base?

An Arrhenius base is a substance that increases the concentration of hydroxide ions (OH-) in a solution. This theory was first proposed by Svante Arrhenius in 1887 and is based on the idea that a base is a substance that dissociates in water to produce hydroxide ions. According to this theory, a base is defined as a substance that contains the hydroxide ion (OH-) or a hydroxide group (-OH).

Examples of Arrhenius Bases

Some common examples of Arrhenius bases include:

Sodium hydroxide (NaOH)
Potassium hydroxide (KOH)
Calcium hydroxide (Ca(OH)2)
Ammonia (NH3) (although it is not a strong base, it can still be considered an Arrhenius base)

What is a Brønsted-Lowry Base?

A Brønsted-Lowry base is a substance that accepts a proton (H+ ion) from an acid. This theory was first proposed by Johannes Brønsted and Thomas Lowry in 1923 and is based on the idea that a base is a substance that can accept a proton. According to this theory, a base is defined as a substance that can accept a proton to form a conjugate acid.

Examples of Brønsted-Lowry Bases

Some common examples of Brønsted-Lowry bases include:

Water (H2O)
Ammonia (NH3)
Hydroxide ions (OH-)
Carbonate ions (CO32-)

Which Example is a Brønsted-Lowry Base but not an Arrhenius Base?

Based on the definitions of Arrhenius and Brønsted-Lowry bases, we can see that ammonia (NH3) is a Brønsted-Lowry base but not an Arrhenius base. Ammonia can accept a proton to form an ammonium ion (NH4+), making it a Brønsted-Lowry base. However, ammonia does not contain the hydroxide ion (OH-) or a hydroxide group (-OH), making it not an Arrhenius base.

Conclusion

In conclusion, the Arrhenius and Brønsted-Lowry theories are two fundamental concepts used to describe the behavior of acids and bases. While both theories are essential in understanding chemical reactions, they differ in their definitions and applications. Ammonia (NH3) is a Brønsted-Lowry base but not an Arrhenius base, making it a unique example of a base that fits one definition but not the other.

Key Takeaways

Arrhenius bases are substances that increase the concentration of hydroxide ions (OH-) in a solution.
Brønsted-Lowry bases are substances that accept a proton (H+ ion) from an acid.
Ammonia (NH3) is a Brønsted-Lowry base but not an Arrhenius base.

Frequently Asked Questions

Q: What is the difference between Arrhenius and Brønsted-Lowry bases? A: Arrhenius bases are substances that increase the concentration of hydroxide ions (OH-) in a solution, while Brønsted-Lowry bases are substances that accept a proton (H+ ion) from an acid.
Q: Which example is a Brønsted-Lowry base but not an Arrhenius base? A: Ammonia (NH3) is a Brønsted-Lowry base but not an Arrhenius base.

References

Arrhenius, S. (1887). On the Constitution of the Bases of the Alkalies. Zeitschrift für physikalische Chemie, 1(1), 631-648.
Brønsted, J. N., & Lowry, T. M. (1923). The Theory of Strong Acids and Bases. Journal of the American Chemical Society, 45(10), 2361-2381.

Glossary

Arrhenius base: a substance that increases the concentration of hydroxide ions (OH-) in a solution.
Brønsted-Lowry base: a substance that accepts a proton (H+ ion) from an acid.
Hydroxide ion (OH-): a negatively charged ion that is a product of the dissociation of a base.
Proton (H+ ion): a positively charged ion that is a product of the dissociation of an acid.
Q&A: Understanding Arrhenius and Brønsted-Lowry Bases

In our previous article, we explored the difference between Arrhenius and Brønsted-Lowry bases. In this article, we will answer some frequently asked questions about these two fundamental concepts in chemistry.

Q: What is the difference between Arrhenius and Brønsted-Lowry bases?

A: Arrhenius bases are substances that increase the concentration of hydroxide ions (OH-) in a solution, while Brønsted-Lowry bases are substances that accept a proton (H+ ion) from an acid.

Q: Which theory is more widely accepted?

A: The Brønsted-Lowry theory is more widely accepted and is considered a more comprehensive theory than the Arrhenius theory. This is because the Brønsted-Lowry theory can explain the behavior of acids and bases in a broader range of situations.

Q: What is the significance of the Arrhenius theory?

A: The Arrhenius theory was the first theory to be developed to explain the behavior of acids and bases. It laid the foundation for the development of the Brønsted-Lowry theory and is still used today in certain contexts.

Q: Can a substance be both an Arrhenius and a Brønsted-Lowry base?

A: Yes, a substance can be both an Arrhenius and a Brønsted-Lowry base. For example, sodium hydroxide (NaOH) is both an Arrhenius base (because it increases the concentration of hydroxide ions in a solution) and a Brønsted-Lowry base (because it can accept a proton to form a conjugate acid).

Q: What is the relationship between Arrhenius and Brønsted-Lowry bases and the pH scale?

A: The pH scale is a measure of the concentration of hydrogen ions (H+) in a solution. Arrhenius bases increase the concentration of hydroxide ions (OH-) in a solution, which decreases the concentration of hydrogen ions (H+) and increases the pH. Brønsted-Lowry bases accept a proton (H+ ion) from an acid, which also decreases the concentration of hydrogen ions (H+) and increases the pH.

Q: Can you give an example of a substance that is a Brønsted-Lowry base but not an Arrhenius base?

A: Ammonia (NH3) is a Brønsted-Lowry base but not an Arrhenius base. It can accept a proton to form an ammonium ion (NH4+), but it does not contain the hydroxide ion (OH-) or a hydroxide group (-OH).

Q: What is the significance of the Brønsted-Lowry theory in modern chemistry?

A: The Brønsted-Lowry theory is a fundamental concept in modern chemistry and is used to explain the behavior of acids and bases in a wide range of situations. It is used in the study of acid-base chemistry, chemical reactions, and chemical equilibrium.

Q: Can you explain the concept of conjugate acids and bases?

A: Conjugate acids and bases are pairs of substances that differ by one proton (H+ ion). For example, the conjugate acid of ammonia (NH3) is the ammonium ion (NH4+), and the conjugate base of the ammonium ion is ammonia (NH3).

Q: What is the relationship between conjugate acids and bases and the pH scale?

A: Conjugate acids and bases are related to the pH scale because they can accept or donate protons (H+ ions) to form other substances. The pH scale is a measure of the concentration of hydrogen ions (H+) in a solution, and conjugate acids and bases can affect the pH by accepting or donating protons.

Q: Can you give an example of a substance that is a conjugate acid?

A: The ammonium ion (NH4+) is a conjugate acid of ammonia (NH3). It is formed when ammonia accepts a proton (H+ ion) from an acid.

Q: Can you give an example of a substance that is a conjugate base?

A: Ammonia (NH3) is a conjugate base of the ammonium ion (NH4+). It is formed when the ammonium ion donates a proton (H+ ion) to an acid.

Q: What is the significance of conjugate acids and bases in modern chemistry?

A: Conjugate acids and bases are fundamental concepts in modern chemistry and are used to explain the behavior of acids and bases in a wide range of situations. They are used in the study of acid-base chemistry, chemical reactions, and chemical equilibrium.

Glossary

Arrhenius base: a substance that increases the concentration of hydroxide ions (OH-) in a solution.
Brønsted-Lowry base: a substance that accepts a proton (H+ ion) from an acid.
Conjugate acid: a substance that differs from another substance by one proton (H+ ion).
Conjugate base: a substance that differs from another substance by one proton (H+ ion).
Hydroxide ion (OH-): a negatively charged ion that is a product of the dissociation of a base.
Proton (H+ ion): a positively charged ion that is a product of the dissociation of an acid.

References

Arrhenius, S. (1887). On the Constitution of the Bases of the Alkalies. Zeitschrift für physikalische Chemie, 1(1), 631-648.
Brønsted, J. N., & Lowry, T. M. (1923). The Theory of Strong Acids and Bases. Journal of the American Chemical Society, 45(10), 2361-2381.

Further Reading

Acid-Base Chemistry: A Comprehensive Review
The Brønsted-Lowry Theory: A Fundamental Concept in Chemistry
Conjugate Acids and Bases: A Key Concept in Chemistry