Problem 85 - Chemistry: The Central Science

Question

Hydrogen peroxide is capable of oxidizing (a) hydrazine to \(\mathrm{N}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\), (b) \(\mathrm{SO}_{2}\) to \(\mathrm{SO}_{4}^{2-}\), (c) \(\mathrm{NO}_{2}^{-}\)to \(\mathrm{NO}_{3}^{-}\), (d) \(\mathrm{H}_{2} \mathrm{~S}(\mathrm{~g})\) to \(\mathrm{S}(s)\), (e) \(\mathrm{Fe}^{2+}\) to \(\mathrm{Fe}^{3+}\). Write a balanced net ionic equation for each of these redox reactions.

Step-by-Step Solution

Answer

The balanced net ionic equations for the given reactions involving hydrogen peroxide are: (a) \(2\mathrm{N}(-2) + 2\mathrm{O}(-2) + 4\mathrm{H}^{+} \to \mathrm{N}_{2} + 2\mathrm{O}_{2}(-2) + 2\mathrm{H}_{2}\mathrm{O}\) (b) \(2\mathrm{SO}_{2}(g) + 2\mathrm{H}_{2}\mathrm{O}_{2}(aq) \to 2\mathrm{SO}_{4}^{2-}(aq) + 2\mathrm{H}_{2}\mathrm{O}(l)\) (c) \(\mathrm{H}_{2}\mathrm{O}_{2}(aq) + 2\mathrm{NO}_{2}^{-}(aq) \to 2\mathrm{NO}_{3}^{-}(aq) + 2\mathrm{H}^{+}(aq)\) (d) \(\mathrm{H}_{2}\mathrm{S}(g) + 4\mathrm{H}_{2}\mathrm{O}_{2}(aq) \to \mathrm{S}(s) + 4\mathrm{H}_{2}\mathrm{O}(l) + 2\mathrm{H}^{+}(aq)\) (e) \(\mathrm{Fe}^{2+}(aq) + \mathrm{H}_{2}\mathrm{O}_{2}(aq) \to \mathrm{Fe}^{3+}(aq) + 2\mathrm{H}_{2}\mathrm{O}(l)\)

Step 1: Identify the oxidation states

We have: NH2NH2 (oxidation state of N: -2) H2O2 (oxidation state of O: -2, H: +1) N2 (oxidation state of N: 0) H2O (oxidation state of O: -2, H: +1)

Step 2: Determine the oxidation and reduction half-reactions

N in NH2NH2 is reduced from -2 to 0 (N2), while O in H2O2 is oxidized from -2 to -2 (H2O). The half-reactions are: 1. 2N(-2) -> N2 2. 2O(-2) -> O2(-2) + O2(-2)

Step 3: Balance the atoms

No further balancing required for N. For O atoms, we add H2O to the half-reactions: 1. 2N(-2) -> N2 2. 2O(-2) -> O2(-2) + O2(-2) + 2H2O

Step 4: Balance the hydrogen atoms

Add H+ to balance the hydrogen atoms in the second half-reaction: 1. 2N(-2) -> N2 2. 2O(-2) + 4H+ -> O2(-2) + O2(-2) + 2H2O

Step 5: Balance the charges

Balance charges by adding electrons: 1. 2N(-2) -> N2 + 4e- 2. 2O(-2) + 4H+ + 4e- -> O2(-2) + O2(-2) + 2H2O

Step 6: Combine the half-reactions and cancel out common ions

Adding the two half-reactions, we get the balanced net ionic equation: 2N(-2) + 2O(-2) + 4H+ -> N2 + O2(-2) + O2(-2) + 2H2O Now, let's write the balanced net ionic equations for the other reactions: (b) H2O2 oxidizing SO2 to SO4^2-: \(2\mathrm{SO}_{2}(g) + 2\mathrm{H}_{2}\mathrm{O}_{2}(aq) \to 2\mathrm{SO}_{4}^{2-}(aq) + 2\mathrm{H}_{2}\mathrm{O}(l)\) (c) H2O2 oxidizing NO2^- to NO3^-: \(\mathrm{H}_{2}\mathrm{O}_{2}(aq) + 2\mathrm{NO}_{2}^{-}(aq) \to 2\mathrm{NO}_{3}^{-}(aq) + 2\mathrm{H}^{+}(aq)\) (d) H2O2 oxidizing H2S(g) to S(s): \(\mathrm{H}_{2}\mathrm{S}(g) + 4\mathrm{H}_{2}\mathrm{O}_{2}(aq) \to \mathrm{S}(s) + 4\mathrm{H}_{2}\mathrm{O}(l) + 2\mathrm{H}^{+}(aq)\) (e) H2O2 oxidizing Fe^2+ to Fe^3+: \(\mathrm{Fe}^{2+}(aq) + \mathrm{H}_{2}\mathrm{O}_{2}(aq) \to \mathrm{Fe}^{3+}(aq) + 2\mathrm{H}_{2}\mathrm{O}(l)\)

Key Concepts

Oxidation and Reduction

Oxidation and reduction are two fundamental concepts in chemistry that describe the transfer of electrons between substances. During a redox reaction, one substance (the oxidizing agent) gains electrons, undergoing a reduction, while the other (the reducing agent) loses electrons, undergoing oxidation.

For instance, when hydrogen peroxide ((H_₂O_₂)) acts as an oxidizing agent, it causes other substances to lose electrons. Taking hydrazine ((NH_₂NH_₂)) as an example, each nitrogen atom in hydrazine has an oxidation state of -2. Through the reaction, the nitrogen is oxidized to nitrogen gas ((N_₂)) where nitrogen is in its elemental state with an oxidation state of 0. This increase in oxidation state indicates that nitrogen has lost electrons.

Conversely, the oxygen in hydrogen peroxide undergoes reduction. Since there is no change in the oxidation state of oxygen here, it might seem that no reduction occurs. However, the accompanying change that actually occurs is the transfer of electrons to oxygen, which are then immediately used to convert hydrogen ions ((H^+)) to water ((H_₂O)), illustrating the fundamental relationship between oxidation and reduction in redox reactions.

Balancing Net Ionic Equations

Balancing net ionic equations is a crucial step in describing redox processes accurately. These equations only include the particles that participate in the reaction, omitting the spectator ions.

The process to balance them involves several key steps: identifying the half-reactions for oxidation and reduction, balancing the number of atoms, then balancing the number of charges by adding electrons. It's also necessary to equalize the number of electrons lost in the oxidation half-reaction with those gained in the reduction half-reaction.

For our exercise, balancing begins with identifying that hydrazine loses electrons to form nitrogen gas, and hydrogen peroxide gains electrons to form water. This implies writing two half-reactions, one for oxidation and one for reduction, and then carefully adding water molecules, hydrogen ions, or electrons to bring balance to each side of the equations. Finally, by combining these half-reactions and canceling out the common species, a balanced net ionic equation is formed.

Oxidation States

Knowing the oxidation states of atoms in a molecule is key to understanding redox reactions. The oxidation state, or oxidation number, signifies the theoretical charge an atom would have if electrons were distributed according to certain rules.

In the given exercise, for example, the oxidation state of nitrogen in hydrazine is -2, implying it is in a relatively reduced form. During the reaction, the oxidation state of nitrogen increases from -2 to 0. This change reflects the loss of electrons, which is the essence of oxidation.

To determine the oxidation states, we often follow a set of rules that prioritize certain elements over others in terms of assigning electron ownership. For instance, hydrogen is usually assigned a +1 oxidation state and oxygen a -2, except for peroxides like hydrogen peroxide where oxygen has a -1 oxidation state. Assigning oxidation states correctly is fundamental for balancing redox reactions and for identifying the roles of different chemicals as oxidizing or reducing agents.

Chemical Oxidizing Agents

Chemical oxidizing agents are substances that accept electrons from other compounds, thereby causing the other compounds to be oxidized. In the examples provided in the exercise, hydrogen peroxide ((H_₂O_₂)) acts as an oxidizing agent.

Oxidizing agents feature in a wide range of chemical reactions, from industrial processes to biological systems. A strong oxidizing agent, like hydrogen peroxide, has a high affinity for electrons and will often have relatively high positive oxidation states after the reaction.

In each of the given reactions, hydrogen peroxide takes electrons away from the other substances like hydrazine or sulfur dioxide ((SO_₂)), which gets oxidized to nitrogen gas or sulfate ions ((SO_₄^2-)), respectively. Recognizing an oxidizing agent is valuable for predicting the course of chemical reactions and for crafting strategies to control and utilize these reactions for a variety of purposes including synthesis, energy storage, and environmental remediation.