Question

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

Short answer

Here's a summary of the balanced net ionic equations for each reaction: a) \(2 \mathrm{H}_{2} \mathrm{O}_{2} + \mathrm{N}_{2} \mathrm{H}_{4} \rightarrow \mathrm{N}_{2} + 4 \mathrm{H}_{2} \mathrm{O}\) b) \(2 \mathrm{H}_{2} \mathrm{O}_{2} + 2 \mathrm{SO}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O} + \mathrm{SO}_{4}^{2-}\) c) \(\mathrm{H}_{2} \mathrm{O}_{2} + \mathrm{NO}_{2}^{-} \rightarrow \mathrm{NO}_{3}^{-} + \mathrm{H}_{2} \mathrm{O}\) d) \(2 \mathrm{H}_{2} \mathrm{O}_{2} + 2 \mathrm{H}_{2} \mathrm{S} \rightarrow 4 \mathrm{H}_{2} \mathrm{O} + 2 \mathrm{S}\) e) \(\mathrm{H}_{2} \mathrm{O}_{2} + \mathrm{Fe}^{2+} \rightarrow \mathrm{Fe}^{3+} + 2 \mathrm{H}_{2} \mathrm{O}\)

Step-by-step solution

a) H2O2 + H2O2 → N2 + H2O + H2O

Step 1: Determine the oxidation states - H2O2: H is +1, O is -1. - N2H4: N is -2, H is +1. - N2: N is 0. - H2O: H is +1, O is -2. Step 2: Identify oxidation and reduction - Hydrazine (N2H4) is oxidized (from -2 to 0). - Hydrogen peroxide (H2O2) is reduced (from -1 to -2). Step 3: Balance atoms and charge In order to balance the redox reaction, we need to balance the atoms and charge: \(2 \mathrm{H}_{2} \mathrm{O}_{2} + \mathrm{N}_{2} \mathrm{H}_{4} \rightarrow \mathrm{N}_{2} + 4 \mathrm{H}_{2} \mathrm{O} \)

b) H2O2 + 2SO2 → H2O + H2O + SO42-

Step 1: Determine the oxidation states - H2O2: H is +1, O is -1. - SO2: S is +4, O is -2. - SO42-: S is +6, O is -2. - H2O: H is +1, O is -2. Step 2: Identify oxidation and reduction - SO2 is oxidized (from +4 to +6). - H2O2 is reduced (from -1 to -2). Step 3: Balance atoms and charge - \(2 \mathrm{H}_{2} \mathrm{O}_{2} + 2 \mathrm{SO}_{2} \rightarrow 2 \mathrm{H}_{2} \mathrm{O} + \mathrm{SO}_{4}^{2-} \)

c) H2O2 + NO2- → NO3- + H2O

Step 1: Determine the oxidation states - H2O2: H is +1, O is -1. - NO2-: N is +3, O is -2. - NO3-: N is +5, O is -2. - H2O: H is +1, O is -2. Step 2: Identify oxidation and reduction - NO2- is oxidized (from +3 to +5). - H2O2 is reduced (from -1 to -2). Step 3: Balance atoms and charge - \(\mathrm{H}_{2} \mathrm{O}_{2} + \mathrm{NO}_{2}^{-} \rightarrow \mathrm{NO}_{3}^{-} + \mathrm{H}_{2} \mathrm{O} \)

d) H2O2 + H2S -> H2O + H2O + S

Step 1: Determine the oxidation states - H2O2: H is +1, O is -1. - H2S: H is +1, S is -2. - H2O: H is +1, O is -2. - S: S is 0. Step 2: Identify oxidation and reduction - H2S is oxidized (from -2 to 0). - H2O2 is reduced (from -1 to -2). d) Step 3: Balance atoms and charge - \(2 \mathrm{H}_{2} \mathrm{O}_{2} + 2 \mathrm{H}_{2} \mathrm{S} \rightarrow 4 \mathrm{H}_{2} \mathrm{O} + 2 \mathrm{S} \)

e) H2O2 + Fe2+ → Fe3+ + H2O + H2O

Step 1: Determine the oxidation states - H2O2: H is +1, O is -1. - Fe2+: Fe is +2. - Fe3+: Fe is +3. - H2O: H is +1, O is -2. Step 2: Identify oxidation and reduction - Fe2+ is oxidized (from +2 to +3). - H2O2 is reduced (from -1 to -2). Step 3: Balance atoms and charge - \(\mathrm{H}_{2} \mathrm{O}_{2} + \mathrm{Fe}^{2+} \rightarrow \mathrm{Fe}^{3+} + 2 \mathrm{H}_{2} \mathrm{O} \)

Key concepts

Balancing Redox Equations

Balancing redox equations is an essential skill in chemistry that ensures both the number of atoms and the charge are equal on both sides of a reaction. This is done by following three main steps: identifying oxidation states, recognizing which species are oxidized and reduced, and finally, balancing the equation by adding coefficients.

For example, in the oxidation of hydrazine to nitrogen gas using hydrogen peroxide, we start by determining the oxidation states. Hydrazine has nitrogen in the -2 state, which goes to 0 in nitrogen gas, indicating oxidation. Hydrogen peroxide, on the other hand, goes from having oxygen in the -1 state to -2 in water, indicating reduction.

Once we've identified these changes, we balance the atoms that have changed oxidation state, followed by ensuring that the overall charge is the same on both sides. Any imbalance is corrected by adding appropriate coefficients before the compounds in the equation. This balanced approach not only conserves mass but also charge, adhering to the foundational laws of chemistry.

Oxidation States

Understanding oxidation states is key to working with redox reactions. An oxidation state is a number assigned to an element in a chemical compound that represents the number of electrons lost or gained by an atom of that element in the compound. The rules for determining oxidation states include: elements in their elemental form have an oxidation state of 0, oxygen typically has an oxidation state of -2 (except in peroxides where it's -1 or when bonded to fluorine), and hydrogen is usually +1 (except when bonded to metals where it's -1).

Using hydrogen peroxide as an example, the oxygen has an unusual oxidation state of -1 because it's a peroxide. In the oxidation of sulfur dioxide to sulfate, the oxidation state of sulfur increases from +4 to +6, reflecting its loss of electrons and subsequent oxidation. These oxidation numbers are like the 'accounting system' of electrons in chemical reactions, providing invaluable insight into the transfer of electrons that occurs in redox processes.

Chemical Oxidizing Agents

In redox reactions, a chemical oxidizing agent is responsible for accepting electrons from another species, causing the other species to oxidize. The oxidizing agent itself gets reduced in the process. Hydrogen peroxide is a powerful oxidizing agent, which can take part in various redox reactions, as seen in the textbook examples.

In the case of the conversion of hydrazine to nitrogen, hydrogen peroxide receives electrons from hydrazine, allowing for the breakdown of the N-N bond to form nitrogen gas. The ability of an oxidizing agent to accept electrons is a key characteristic that determines its strength and usefulness in chemical reactions. A robust oxidizing agent like hydrogen peroxide is versatile and can be used to drive many different types of redox reactions, evidenced by its ability to oxidize a wide range of substances from sulfur dioxide to iron ions.