Problem 101 - Chemistry Matter and Change

Question

Identify whether each of the reactants in these reactions is acting as an acid or a base. (Chapter 18) a. \(HBr+H_{2} O \rightarrow H_{3} O^{+}+Br^{-}\) b. \(NH_{3}+HCOOH \rightarrow NH_{4}^{+}+HCOO^{-}\) c. \(HCO_{3}^{-}+H_{2} O \rightarrow CO_{3}^{-}+H_{3} O^{+}\)

Step-by-Step Solution

Answer

In summary: a. In the reaction \(HBr + H_{2}O \rightarrow H_{3}O^{+} + Br^{-}\), HBr is acting as an acid and \(H_{2}O\) is acting as a base. b. In the reaction \(NH_{3}+HCOOH \rightarrow NH_{4}^{+}+HCOO^{-}\), HCOOH is acting as an acid and \(NH_{3}\) is acting as a base. c. In the reaction \(HCO_{3}^{-}+H_{2}O \rightarrow CO_{3}^{-}+H_{3}O^{+}\), \(HCO_{3}^{-}\) is acting as an acid and \(H_{2}O\) is acting as a base.

Step 1: Reaction a: \((\)HBr+H_{2}O \rightarrow H_{3}O^{+}+Br^{-}\()\)

In this reaction, the reactant HBr donates a proton (H⁺) to H₂O, resulting in the formation of H₃O⁺ and Br⁻. According to the Brønsted-Lowry definition, HBr acts as an acid, while H₂O acts as a base.

Step 2: Reaction b: \((\)NH_{3}+HCOOH \rightarrow NH_{4}^{+}+HCOO^{-}\()\)

In this reaction, the reactant HCOOH donates a proton (H⁺) to NH₃, resulting in the formation of NH₄⁺ and HCOO⁻. According to the Brønsted-Lowry definition, HCOOH acts as an acid, while NH₃ acts as a base.

Step 3: Reaction c: \((\)HCO_{3}^{-}+H_{2}O \rightarrow CO_{3}^{-}+H_{3}O^{+}\()\)

In this reaction, the reactant HCO₃⁻ donates a proton (H⁺) to H₂O, resulting in the formation of CO₃²⁻ and H₃O⁺. According to the Brønsted-Lowry definition, HCO₃⁻ acts as an acid, while H₂O acts as a base. In summary: a. HBr is an acid, H₂O is a base b. HCOOH is an acid, NH₃ is a base c. HCO₃⁻ is an acid, H₂O is a base

Key Concepts

Proton Transfer

Proton transfer is a key process in many chemical reactions, especially in the context of acid-base chemistry. When a proton ( H⁺ ) is transferred from one molecule to another, it often results in the formation of a new substance and changes the nature of the reactants. This simple movement can determine whether a molecule behaves as an acid or a base according to the Brønsted-Lowry theory.

This theory defines acids as proton donors and bases as proton acceptors.

• During proton transfer, the acid loses a proton.
• The base gains a proton.

This interchange is crucial for the understanding of many biochemical pathways and industrial processes. It's the basic principle behind the behavior of acids and bases.

Acid-Base Reactions

Acid-base reactions involve the exchange of protons between reactants. These reactions are essential to understanding numerous natural and industrial processes. The Brønsted-Lowry theory provides a framework for identifying acids and bases in these reactions.

In a typical acid-base reaction:

• The acid gives up a proton, becoming its conjugate base.
• The base accepts the proton, becoming its conjugate acid.

For example, in the reaction NH₃ + HCOOH → NH₄⁺ + HCOO⁻, HCHO relinquishes a proton to NH₃. Therefore, HCOOH acts as an acid, whereas NH₃ is the base. Understanding these dynamics helps predict the outcome of the reaction and the nature of the resulting compounds.

HBr as an Acid

Hydrobromic acid ( HBr ) is a classic example of a strong acid in Brønsted-Lowry theory. It effortlessly donates its proton to nearby water molecules, demonstrating its strength as an acid. This exchange is seen in the reaction: HBr + H₂O → H₃O⁺ + Br⁻.

Here, HBr serves as a proton donor, which defines it as an acid. The resulting hydronium ion ( H₃O⁺ ) exhibits increased acidity in water, with bromide ion ( Br⁻ ) as the resultant conjugate base. Knowing HBr's behavior is essential in understanding how strong acids fully ionize in water, impacting the pH level significantly.

HCOOH as an Acid

Formic acid ( HCOOH ) is a relatively weak acid compared to strong acids like HBr . However, it still functions as an acid by donating a proton in acid-base reactions. In the context of the reaction: HCOOH + NH₃ → NH₄⁺ + HCOO⁻, HCOOH is the acid because it donates a proton to ammonia ( NH₃ ), forming an ammonium ion ( NH₄⁺ ) and the conjugate base, formate ion ( HCOO⁻ ).

This illustrates how even weak acids like HCOOH can engage in proton transfer to form new compounds, emphasizing their role in chemical processes despite their lower acidity.

HCO₃⁻ as an Acid

Bicarbonate ions ( HCO₃⁻ ) highlight an interesting aspect of acid-base chemistry. Though often considered basic, HCO₃⁻ can act as an acid, donating a proton to a base such as water. This behavior is seen in the reaction: HCO₃⁻ + H₂O → CO₃²⁻ + H₃O⁺.

In this case, HCO₃⁻ donates a proton to water, forming a carbonate ion ( CO₃²⁻ ) and hydronium ion ( H₃O⁺ ). It shows HCO₃⁻'s dual role as both an acid and a conjugate base, which is central to many natural processes, like blood buffering. This dual capability makes it significant in maintaining pH stability in various chemical and biological systems.