Problem 139 - Chemistry

Question

Most of the hydrides of Group \(1 \mathrm{~A}\) and Group \(2 \mathrm{~A}\) metals are ionic (the exceptions are \(\mathrm{BeH}_{2}\) and \(\mathrm{MgH}_{2}\), which are covalent compounds). (a) Describe the reaction between the hydride ion \(\left(\mathrm{H}^{-}\right)\) and water in terms of a Brønsted acid-base reaction. (b) The same reaction can also be classified as a redox reaction. Identify the oxidizing and reducing agents.

Step-by-Step Solution

Answer

(a) H⁻ is the base, H₂O is the acid. (b) H⁻ is reducing agent, H₂O is oxidizing agent.

Step 1: Define Brønsted Acid-Base Reaction

In a Brønsted acid-base reaction, the acid donates a proton (H⁺) and the base accepts a proton. In the case of the hydride ion (H⁻) reacting with water (H₂O), the water acts as the acid and the hydride ion acts as the base. The reaction is: \[ \text{H}^- + \text{H}_2\text{O} \rightarrow \text{OH}^- + \text{H}_2 \] Here, H⁻ accepts a proton from H₂O, forming OH⁻ and H₂.

Step 2: Identify Oxidation and Reduction in Redox Reaction

In a redox reaction, the oxidizing agent gains electrons and is reduced, while the reducing agent loses electrons and is oxidized. In the hydride ion and water reaction: \[ \text{H}^- + \text{H}_2\text{O} \rightarrow \text{OH}^- + \text{H}_2 \] The hydride ion \( \text{H}^- \) donates an electron and is oxidized to form H₂. Thus, \( \text{H}^- \) is the reducing agent. Conversely, in H₂O, the oxygen accepts an electron to form OH⁻, so H₂O is the oxidizing agent.

Key Concepts

Brønsted Acid-Base Theory

The Brønsted acid-base theory is a fundamental concept in chemistry that describes how substances behave as acids or bases during a chemical reaction. According to this theory, an acid is a substance that donates a proton (\(\text{H}^+\)), while a base is a substance that accepts a proton. In the context of the reaction between the hydride ion (\(\text{H}^-\)) and water (\(\text{H}_2\text{O}\)), we can analyze how this theory applies. * In this specific reaction: - Water (\(\text{H}_2\text{O}\)) acts as the Brønsted acid because it donates a proton to the hydride ion. - The hydride ion (\(\text{H}^-\)) behaves as the Brønsted base as it accepts the proton from water.The reaction can be represented as follows: \[\text{H}^- + \text{H}_2\text{O} \rightarrow \text{OH}^- + \text{H}_2\] From this equation, the hydride ion accepts a proton, resulting in the formation of a hydroxide ion (\(\text{OH}^-\)) and hydrogen gas (\(\text{H}_2\)). This showcases a classic acid-base interaction where roles are clearly defined.

Ionic Hydrides

Ionic hydrides are compounds formed between hydrogen and metals from Groups 1 and 2 of the periodic table. These hydrides generally contain the hydride ion (\(\text{H}^-\)), which has one more electron than a neutral hydrogen atom. This gives it a negative charge. They are typically formed by the reaction of hydrogen gas with a metal.Some characteristics of ionic hydrides include:

• The hydride ion is highly reactive, especially with polar compounds like water. This reactivity is due to its desire to donate its excess electron and form neutral hydrogen (\(\text{H}_2\)).
• Ionic hydrides are generally solid at room temperature and are known for their strong reducing properties.
• They conduct electricity when melted or dissolved in water, indicating the presence of free-moving ions.

The provided exercise mentions they mostly include Group 1A and Group 2A metals, excluding a few covalent hydrides like \(\text{BeH}_2\) and \(\text{MgH}_2\). In aqueous reactions, their behavior is crucial to understanding their role in acid-base and redox reactions.

Redox Reactions

Redox reactions are chemical processes where the oxidation state of atoms is changed, typically through the transfer of electrons. In the case of the reaction between the hydride ion (\(\text{H}^-\)) and water, we can further understand how these reactions operate.In every redox reaction:

• There is an oxidizing agent that gains electrons and, in doing so, becomes reduced.
• Conversely, there is a reducing agent that loses electrons and is oxidized in the process.

For the reaction \[\text{H}^- + \text{H}_2\text{O} \rightarrow \text{OH}^- + \text{H}_2\]:

• The hydride ion (\(\text{H}^-\)) acts as a reducing agent because it donates its electron, transforming into molecular hydrogen (\(\text{H}_2\)).
• The water (\(\text{H}_2\text{O}\)), on the other hand, serves as an oxidizing agent. It accepts an electron from the hydride ion, forming a hydroxide ion (\(\text{OH}^-\)).

This transformation highlights the dual nature of the reaction, serving as both an acid-base and a redox reaction, where electron exchange and proton transfer occur simultaneously.