Question

(a) What is the difference between the Arrhenius and the Bronsted-Lowry definitions of a base? (b) Can a substance behave as an Arrhenius base if it does not contain an OH group? Explain.

Short answer

In conclusion, the main difference between the Arrhenius and Bronsted-Lowry definitions of a base is that Arrhenius bases increase the concentration of hydroxide ions (\( OH^- \)) in water, whereas Bronsted-Lowry bases accept protons (\( H^+ \)). A substance without an OH group, such as ammonia (\( NH_3 \)), can still act as an Arrhenius base if it reacts with water to produce hydroxide ions.

Step-by-step solution

Comparing Arrhenius and Bronsted-Lowry definitions of a base

The Arrhenius definition of a base states that a base is a substance that increases the concentration of hydroxide ions (\( OH^- \)) when dissolved in water. On the other hand, the Bronsted-Lowry definition states that a base is a substance that can accept a proton (\( H^+ \)).

Explaining the difference

The main difference between these two definitions is that the Arrhenius definition is limited to the substances that increase the concentration of hydroxide ions in water while the Bronsted-Lowry definition is applicable to a broader range of substances which can accept protons, not necessarily increasing the concentration of hydroxide ions.

Substance behaving as an Arrhenius base without an OH group

To answer this question, let's recall the Arrhenius definition: a substance is considered an Arrhenius base if it increases the concentration of hydroxide ions (\( OH^- \)) in water. If a substance does not contain an OH group, it may still be considered an Arrhenius base if it can react with water to produce hydroxide ions.

Example of a substance without an OH group acting as an Arrhenius base

An example of such a substance is ammonia (\( NH_3 \)). Although it does not contain an OH group, ammonia can act as an Arrhenius base because when it is dissolved in water, it reacts with water molecules to form ammonium ions (\( NH_4^+ \)) and hydroxide ions (\( OH^- \)): \[ NH_3(g) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq) \] In this reaction, ammonia accepts a proton from water, creating hydroxide ions that increase the concentration of hydroxide ions in the solution.

Conclusion

In conclusion, the Arrhenius definition of a base involves increasing the concentration of hydroxide ions in water, while the Bronsted-Lowry definition involves accepting protons. A substance can behave as an Arrhenius base even if it does not contain an OH group, as long as it can react with water or any other substance to produce hydroxide ions.

Key concepts

Definitions of Acids and Bases

Understanding the nature of acids and bases is fundamental in chemistry. An acid can be defined as a substance that donates a proton (\( H^+ \)), according to the Bronsted-Lowry theory, or one that increases the concentration of hydrogen ions (\( H^+ \)) in an aqueous solution, as per the Arrhenius definition. On the contrary, a base is characterized by its ability to accept a proton (Bronsted-Lowry) or increase the concentration of hydroxide ions (\( OH^- \)) in water (Arrhenius).

These two definitions highlight the conditions under which substances behave as acids or bases. While Arrhenius's definition is restricted to aqueous solutions and focuses on the presence of ions, Bronsted-Lowry's definition can be applied to more general scenarios, including non-aqueous environments. It's a more inclusive approach that accommodates substances that don't necessarily produce hydroxide ions but can still accept protons, aligning with a broader range of observed chemical behaviors.

Proton Acceptor

The concept of a proton acceptor is central to the Bronsted-Lowry definition of bases. In this context, the ability of a substance to accept a proton (\( H^+ \)) is what qualifies it as a base. This definition expands the horizon beyond just substances producing hydroxide ions in water.

For instance, ammonia (\( NH_3 \)) is a classic example of a Bronsted-Lowry base. It does not contain a hydroxide group but readily accepts a proton from a water molecule to form ammonium (\( NH_4^+ \)) in a reversible reaction, exemplifying its base-like behavior without the presence of \( OH^- \) ions from the molecule itself. This perspective allows us to categorize a greater variety of substances as bases, such as those that might donate pairs of electrons to accept protons in chemical reactions.

Hydroxide Ion Concentration

The hydroxide ion concentration in an aqueous solution is pivotal for the Arrhenius definition of a base. In this framework, a base is defined by its ability to increase the concentration of \( OH^- \) ions when dissolved in water. The more \( OH^- \) ions present, the more basic the solution. This ion is what gives bases their characteristic properties, such as a slippery feel and the ability to neutralize acids.

Common substances like sodium hydroxide (\( NaOH \)) raise the hydroxide ion concentration directly upon dissolution. However, other compounds, like the aforementioned ammonia, increase \( OH^- \) via chemical reactions with water molecules, indicating that even without an \( OH \) group, a substance can still act as an Arrhenius base.

Chemical Reaction in Aqueous Solution

A chemical reaction in aqueous solution involves solutes reacting with water or with each other while dissolved in water. This medium is essential for the behavior of acids and bases under both Arrhenius and Bronsted-Lowry theories. The water serves as a reactant itself in many base-generating reactions, exemplified by ammonia's reaction to form hydroxide ions.

In such reactions, water not only dissolves substances but also interacts with them, altering the solution's chemistry - like becoming more acidic or basic. These reactions are an integral part of understanding how acids and bases behave and how they are characterized in both theoretical frameworks.