Problem 121 - Introductory Chemistry Atoms First

Question

Which of the following are electron-transfer reactions? For those that are, indicate which reactant is the reducing agent and which reactant is the oxidizing agent. (a) \(\mathrm{PF}_{3}+3 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{H}_{3} \mathrm{PO}_{3}+3 \mathrm{HF}\) (b) \(\mathrm{H}_{2}+\mathrm{Cl}_{2} \rightarrow 2 \mathrm{HCl}\) (c) \(2 \mathrm{Cr}_{2} \mathrm{O}_{3}+3 \mathrm{Si} \rightarrow 4 \mathrm{Cr}+3 \mathrm{SiO}_{2}\) (d) \(\mathrm{HCl}+\mathrm{NaOH} \rightarrow \mathrm{NaCl}+\mathrm{H}_{2} \mathrm{O}\)

Step-by-Step Solution

Answer

(a) Not an electron-transfer reaction. (b) Electron-transfer reaction; \(\mathrm{H_{2}}\) is the reducing agent and \(\mathrm{Cl_{2}}\) is the oxidizing agent. (c) Electron-transfer reaction; \(\mathrm{Si}\) is the reducing agent and \(\mathrm{Cr_{2}O_{3}}\) is the oxidizing agent. (d) Not an electron-transfer reaction.

Step 1: Identify the oxidation state of elements in each compound

For each given reaction, we should first determine the oxidation state of the elements in each compound both in the reactants and products. This will help identify any changes in oxidation state and classify the reactions as electron-transfer reactions or not. (a) \(\mathrm{PF_{3}}+3 \mathrm{H_{2}O} \rightarrow \mathrm{H_{3}PO_{3}}+3 \mathrm{HF}\) Oxidation states: P in \(\mathrm{PF_{3}}\): +3 F in \(\mathrm{PF_{3}}\): -1 P in \(\mathrm{H_{3}PO_{3}}\): +3 F in \(\mathrm{HF}\): -1 No change in the oxidation state of any element, so it's not an electron-transfer reaction. (b) \(\mathrm{H_{2}}+\mathrm{Cl_{2}} \rightarrow 2 \mathrm{HCl}\) Oxidation states: H in \(\mathrm{H_{2}}\): 0 Cl in \(\mathrm{Cl_{2}}\): 0 H in \(\mathrm{HCl}\): +1 Cl in \(\mathrm{HCl}\): -1 There are changes in the oxidation states, so it's an electron-transfer reaction. (c) \(2 \mathrm{Cr_{2}O_{3}}+3 \mathrm{Si} \rightarrow 4 \mathrm{Cr}+3 \mathrm{SiO_{2}}\) Oxidation states: Cr in \(\mathrm{Cr_{2}O_{3}}\): +3 O in \(\mathrm{Cr_{2}O_{3}}\): -2 Si in \(\mathrm{Si}\): 0 Cr in \(\mathrm{Cr}\): 0 O in \(\mathrm{SiO_{2}}\): -2 Si in \(\mathrm{SiO_{2}}\): +4 There are changes in the oxidation states, so it's an electron-transfer reaction. (d) \(\mathrm{HCl}+\mathrm{NaOH} \rightarrow \mathrm{NaCl}+\mathrm{H_{2}O}\) Oxidation states: H in \(\mathrm{HCl}\): +1 Cl in \(\mathrm{HCl}\): -1 Na in \(\mathrm{NaOH}\): +1 O in \(\mathrm{NaOH}\): -2 H in \(\mathrm{H_{2}O}\): +1 O in \(\mathrm{H_{2}O}\): -2 Na in \(\mathrm{NaCl}\): +1 Cl in \(\mathrm{NaCl}\): -1 No change in the oxidation state of any element, so it's not an electron-transfer reaction.

Step 2: Identify the reducing and oxidizing agents

For the electron-transfer reactions identified in Step 1, we should now indicate which reactant is the reducing agent and which reactant is the oxidizing agent. (b) \(\mathrm{H_{2}}+\mathrm{Cl_{2}} \rightarrow 2 \mathrm{HCl}\) Oxidation: H in \(\mathrm{H_{2}}\) goes from 0 to +1, losing electrons - reducing agent Reduction: Cl in \(\mathrm{Cl_{2}}\) goes from 0 to -1, gaining electrons - oxidizing agent (c) \(2 \mathrm{Cr_{2}O_{3}}+3 \mathrm{Si} \rightarrow 4 \mathrm{Cr}+3 \mathrm{SiO_{2}}\) Oxidation: Si in \(\mathrm{Si}\) goes from 0 to +4, losing electrons - reducing agent Reduction: Cr in \(\mathrm{Cr_{2}O_{3}}\) goes from +3 to 0, gaining electrons - oxidizing agent So, the final solutions are: (a) Not an electron-transfer reaction (b) Electron-transfer reaction; \(\mathrm{H_{2}}\) is the reducing agent and \(\mathrm{Cl_{2}}\) is the oxidizing agent. (c) Electron-transfer reaction; \(\mathrm{Si}\) is the reducing agent and \(\mathrm{Cr_{2}O_{3}}\) is the oxidizing agent. (d) Not an electron-transfer reaction

Key Concepts

Oxidation States

Oxidation states, also known as oxidation numbers, are essential in understanding electron-transfer reactions. They represent the hypothetical charge an atom would have if all bonds to atoms of different elements were completely ionic. This aids in keeping track of electron transfer. For instance, hydrogen in its elemental form, \(\mathrm{H_{2}}\), has an oxidation state of 0, meaning it neither gains nor loses electrons. However, when hydrogen reacts to form \(\mathrm{HCl}\), its oxidation state changes to +1, signifying the loss of electrons.
Understanding oxidation states is crucial because:

• They help identify the flow of electrons during a chemical reaction.
• They allow us to determine whether reactions involve electron transfer, making them potential redox (reduction-oxidation) reactions.

By observing changes in oxidation states, we can ascertain whether a chemical reaction involves electron transfer and, consequently, identify which species are oxidized and which are reduced.

Reducing Agents

A reducing agent, or reductant, is a species that donates electrons to another species in a chemical reaction. It "reduces" the other compound by giving it electrons. An excellent example is the molecule \(\mathrm{H_2}\) in the reaction \(\mathrm{H_{2} + Cl_{2} \rightarrow 2 \, HCl}\). Here, hydrogen loses electrons, increasing its oxidation state from 0 in \(\mathrm{H_2}\) to +1 in \(\mathrm{HCl}\). By losing these electrons, hydrogen reduces chlorine (\(\mathrm{Cl_2}\)).
The key properties of reducing agents include:

• Undergoing oxidation themselves as they donate electrons.
• Having a low electronegativity and a high tendency to release electrons.

A solid grasp of reducing agents helps in identifying which reactants provide electrons in a redox process, showcasing their role in balancing such reactions.

Oxidizing Agents

Oxidizing agents, or oxidants, are substances that accept electrons from other substances during a chemical reaction. They "oxidize" the other component by taking electrons. As an example, consider \(\mathrm{Cl_2}\) in the reaction \(\mathrm{H_{2} + Cl_{2} \rightarrow 2 \, HCl}\).Here, chlorine undergoes a reduction in its oxidation state from 0 to -1, symbolizing a gain of electrons as it forms \(\mathrm{HCl}\). In this process, chlorine oxidizes hydrogen.
Characteristics of oxidizing agents include:

• Undergoing a reduction themselves as they gain electrons.
• Possessing a high electronegativity and a strong attraction for electrons.

Knowing about oxidizing agents is crucial to determining which components in a chemical reaction act as electron acceptors. This comprehension is vital for balancing both simple and complex redox equations.

Chemical Reactions Analysis

Analyzing chemical reactions involves identifying the reactants, products, and the changes they undergo. A key aspect is determining whether these reactions involve electron transfers, which are termed redox reactions.
To analyze a chemical reaction effectively, one should:

• Identify changes in oxidation states of elements between the reactants and products.
• Determine which species are oxidized (lose electrons) and which are reduced (gain electrons).
• Pinpoint reducing and oxidizing agents based on electron transfers.

For example, inspecting the reaction \(\mathrm{2 \, Cr_{2}O_{3} + 3 \, Si \rightarrow 4 \, Cr + 3 \, SiO_{2}}\) reveals a transfer of electrons:- Silicon's oxidation state increases from 0 to +4, acting as a reducing agent.- Chromium's state decreases from +3 to 0, indicating it's the oxidizing agent.
This method of analysis is invaluable in understanding the mechanism and the balancing of chemical equations. Spotting these electron transfers clarifies the underlying process of many reactions, both in a lab setting and in natural phenomena.