Bronsted Lowry Acid and Base: Videos & Practice Problems

The Brønsted–Lowry theory, established in 1923, defines acids as proton donors (H⁺) and bases as proton acceptors. This expands on the Arrhenius definition, which focuses on aqueous solutions. Brønsted–Lowry acids and bases form conjugate acid-base pairs, differing by one hydrogen atom. For example, H₃O⁺ and H₂O are conjugate pairs. Understanding these concepts is crucial for grasping acid–base reactions and their applications in various chemical processes.

Bronsted Lowry Acids & Bases Concept 1

In 1923, Bronsted and Lowry introduced a significant advancement in the understanding of acids and bases, refining earlier definitions established by Arrhenius. According to the Bronsted-Lowry theory, an acid is defined as a proton donor, specifically referring to the hydrogen ion (H⁺). This aligns with the Arrhenius definition, which states that an Arrhenius acid increases the concentration of H⁺ ions in aqueous solutions.

However, the Bronsted-Lowry theory diverges from Arrhenius in its definition of bases. While Arrhenius defined a base as a substance that produces hydroxide ions (OH^-) in water, Bronsted and Lowry proposed that a base is a proton acceptor. This means that a base can accept H⁺ ions, which is often facilitated by the presence of lone pairs of electrons or a negative charge. This fundamental shift allows for the consideration of acid-base reactions in non-aqueous solutions, broadening the scope of acid-base chemistry.

Both theories agree on the behavior of acids: Bronsted-Lowry acids donate H⁺ ions, which is consistent with the Arrhenius definition. However, the Bronsted-Lowry approach emphasizes the interaction between acids and bases, highlighting that they always exist in pairs known as conjugate acid-base pairs. These pairs differ by a single hydrogen ion. For example, in the pair consisting of H₂O and OH^-, water (H₂O) can donate an H⁺ to become OH^-, while H₃O⁺ (hydronium ion) and H₂O also form a conjugate pair, differing by one H⁺ ion.

Understanding these concepts is crucial for grasping the dynamics of acid-base reactions and their applications in various chemical contexts.

Bronsted Lowry Acids & Bases Example 1

To determine the conjugate base of a compound, it is essential to understand the concept of proton transfer. In this case, we are examining the compound HSO₄^-. The conjugate base is formed by removing a hydrogen ion (H⁺), which is a positively charged proton. This process results in a change in the overall charge of the compound.

Starting with HSO₄^-, we note its initial charge of -1. When we remove H⁺, we are effectively taking away a positive charge. This action leads to a decrease in the overall positive charge, making the compound more negative. To visualize this, we can think of a number line where removing a positive value shifts the charge further into the negative range.

Thus, after removing H⁺, the remaining species is SO₄^2-. Therefore, the conjugate base of HSO₄^- is SO₄^2-, which carries a charge of -2. This understanding of conjugate acids and bases is crucial in acid-base chemistry, as it helps predict the behavior of substances in chemical reactions.

Bronsted Lowry Acids & Bases Example 2

In the context of acid-base chemistry, understanding conjugate acids and bases is essential. A conjugate acid is formed when a base gains a proton (H⁺). For instance, when considering the compound \( \text{V}_2\text{O}_5 \), adding a proton results in the formation of \( \text{H}_2\text{V}_2\text{O}_5 \). This process effectively increases the positive charge of the compound, shifting its charge from -2 to -1.To visualize this, think of a number line where the addition of a positive charge (H⁺) moves the compound closer to neutrality. Thus, the conjugate acid of \( \text{V}_2\text{O}_5 \) is represented as \( \text{H}_2\text{V}_2\text{O}_5^{-1} \). Practicing with additional questions will help solidify your understanding of how to identify and create conjugate acids and bases.

Write the formula of the conjugate base for the following compound:
H₂Se

• A.
• B.
• C.
• D.

Answer: (A)

Write the formula of the conjugate for the following compound:
NH₂NH₂

• A.
• B.
• C.
• D.

Answer: (D)

Bronsted Lowry Acids & Bases Example 3

In the context of Bronsted-Lowry acids and bases, understanding the transfer of protons (H⁺) is crucial. In the reaction between hydrofluoric acid (HF) and water (H₂O), we can identify the roles of each component based on their behavior during the reaction.

In this example, HF acts as the Bronsted-Lowry acid because it donates a proton (H⁺) to water. As a result of this donation, HF transforms into its conjugate base, fluoride ion (F^-). The water molecule, which accepts the proton, is classified as the Bronsted-Lowry base and is converted into hydronium ion (H₃O⁺), which serves as the conjugate acid.

To summarize, the acid-base pairs in this reaction are as follows: HF is the acid, and F^- is its conjugate base. Conversely, H₂O is the base, and H₃O⁺ is its conjugate acid. These pairs are linked by the transfer of a single proton, highlighting the fundamental principle that conjugate acids and bases differ by one hydrogen ion.

Bronsted Lowry Acids & Bases Example 4

In the context of acid-base chemistry, understanding the behavior of species like cyanide (CN^-) and water (H₂O) is crucial. When CN^- accepts a proton (H⁺), it transforms into hydrogen cyanide (HCN), demonstrating its role as a base. The proton donor in this reaction is water, which acts as an acid by releasing H⁺. Consequently, water becomes hydroxide (OH^-), which is identified as the conjugate base of the acid. This relationship highlights the concept of conjugate acid-base pairs, where the base and its conjugate acid are related through the transfer of a proton.

Water is a prime example of an amphoteric substance, meaning it can function as either an acid or a base depending on the surrounding chemical environment. In the first example, water acts as a base, while in the second, it serves as an acid. This duality is essential for understanding various chemical reactions.

When comparing acids, such as hydrofluoric acid (HF) and water, it is important to consider electronegativity. Both HF and H₂O contain hydrogen bonded to electronegative elements (fluorine and oxygen, respectively). However, HF is a stronger acid than H₂O due to the greater electronegativity of fluorine compared to oxygen. This principle helps determine which species will act as an acid and which will act as a base in a reaction.

To identify a Brønsted-Lowry acid, remember that it must contain hydrogen (H⁺) and be bonded to an electronegative element, creating a polar bond that can easily release the proton. For instance, diatomic hydrogen (H₂) does not qualify as an acid because it forms a nonpolar bond. Similarly, hydrogen bonded to carbon (C) is also unlikely to exhibit acidic behavior due to the similar electronegativities of carbon and hydrogen, resulting in a nonpolar bond.

Utilizing these guidelines will aid in recognizing acids and bases in various chemical contexts, enhancing your understanding of acid-base reactions and their underlying principles.

Which of the following is a Bronsted-Lowry acid?

• A. CH₄
• B. HCN
• C. NH₃
• D. Br₂

Answer: (B)

Determine the chemical equation that would result when carbonate, CO3^2-, reacts with water.

• A.
• B.
• C.
• D.

Answer: (B)

Here’s what students ask on this topic:

What is the Brønsted–Lowry definition of acids and bases?

The Brønsted–Lowry theory, established in 1923, defines acids as proton donors (H⁺) and bases as proton acceptors. This definition expands on the Arrhenius concept, which focuses on aqueous solutions. According to Brønsted–Lowry, an acid donates a proton (H⁺), while a base accepts a proton. This theory is crucial for understanding acid-base reactions beyond just aqueous environments, making it more versatile in various chemical processes.

How do Brønsted–Lowry acids and bases differ from Arrhenius acids and bases?

Brønsted–Lowry acids and bases differ from Arrhenius acids and bases primarily in their definitions and applicability. Arrhenius acids increase the concentration of H⁺ ions in aqueous solutions, and Arrhenius bases increase the concentration of OH⁻ ions. In contrast, Brønsted–Lowry acids are proton donors (H⁺), and Brønsted–Lowry bases are proton acceptors, applicable in both aqueous and non-aqueous solutions. This broader definition allows for a more comprehensive understanding of acid-base reactions.

What are conjugate acid-base pairs in the Brønsted–Lowry theory?

In the Brønsted–Lowry theory, conjugate acid-base pairs are two species that differ by only one proton (H⁺). When an acid donates a proton, it becomes its conjugate base, and when a base accepts a proton, it becomes its conjugate acid. For example, H₂O and OH⁻ form a conjugate acid-base pair, as do H₃O⁺ and H₂O. Understanding these pairs is essential for analyzing acid-base reactions and predicting the behavior of substances in different chemical environments.

Why is the Brønsted–Lowry theory important in chemistry?

The Brønsted–Lowry theory is important in chemistry because it provides a more comprehensive framework for understanding acid-base reactions. Unlike the Arrhenius definition, which is limited to aqueous solutions, the Brønsted–Lowry theory applies to a broader range of environments, including non-aqueous solutions. This theory helps chemists predict the behavior of acids and bases in various chemical processes, making it a fundamental concept in both academic and practical applications of chemistry.

Can you give an example of a Brønsted–Lowry acid-base reaction?

An example of a Brønsted–Lowry acid-base reaction is the interaction between hydrochloric acid (HCl) and ammonia (NH₃). In this reaction, HCl donates a proton (H⁺) to NH₃, forming NH₄⁺ (ammonium ion) and Cl⁻ (chloride ion). The reaction can be represented as:

${\mathrm{HCl}}_{}+{\mathrm{NH}3}_{}\to {\mathrm{NH}4}_{}+{\mathrm{Cl}}_{}$

In this reaction, HCl is the Brønsted–Lowry acid, and NH₃ is the Brønsted–Lowry base.