Question

Write balanced net ionic equations for the following reactions in basic solution: (a) \(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q)+\mathrm{I}_{2}(a q) \rightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{I}^{-}(a q)\) (b) \(\mathrm{Mn}^{2+}(a q)+\mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow \mathrm{MnO}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) (c) \(\mathrm{Zn}(s)+\mathrm{NO}_{3}^{-}(a q) \rightarrow \mathrm{NH}_{3}(a q)+\mathrm{Zn}(\mathrm{OH})_{4}^{2-}(a q)\) (d) \(\mathrm{Bi}(\mathrm{OH})_{3}(s)+\mathrm{Sn}(\mathrm{OH})_{3}^{-}(a q) \rightarrow \mathrm{Bi}(s)+\mathrm{Sn}(\mathrm{OH})_{6}^{2-}(a q)\)

Short answer

Write half-reactions, balance atoms and charges, then combine. Confirm the net ionic equations.

Step-by-step solution

Write the complete reaction equation

For each of the given reactions, first write down the complete, unbalanced reaction equations with all species in their respective states.

Write the half-reactions

Identify the oxidation and reduction half-reactions for each equation. For (a), consider the thiosulfate to tetrathionate and iodine to iodide reactions. For (b), involve Mn and H2O2 in separate half-reactions. For (c), convert Zn and NO3^- to their respective products. Finally, for (d), consider Bi(OH)3 and Sn(OH)3^- as separate half-reactions.

Balance atoms other than O and H

For each half-reaction, begin by balancing the atoms that are not oxygen or hydrogen. Focus on the central metal or nonmetal atom involved in redox changes.

Balance oxygen atoms by adding water

Add H2O molecules to one side of each half-reaction to balance the oxygen atoms.

Balance hydrogen atoms by adding OH⁻

In each half-reaction, add OH⁻ ions to balance the hydrogen atoms. This is necessary because the reaction takes place in a basic solution.

Balance the charge with electrons

Add electrons to one side of each half-reaction to balance the charges, ensuring that both sides of the equation have equal net charges.

Equalize electron transfer between reactions

Multiply the half-reactions by appropriate coefficients to ensure the number of electrons transferred in the oxidation equals those gained in the reduction process.

Add half-reactions to get the net ionic equation

Combine the balanced half-reactions, canceling out identical species on both sides (including electrons), to form the final balanced net ionic equation.

Verify atom and charge balance

Confirm that all atoms and charges are balanced in the final equation for consistency and correctness.

Key concepts

Redox Reactions

Redox reactions, also known as oxidation-reduction reactions, involve the transfer of electrons between two species. One species loses electrons (oxidation), and another gains electrons (reduction).
Such reactions are a crucial component of many chemical processes, including those occurring in batteries and metabolism.
Understanding redox reactions involves identifying the oxidizing and reducing agents:

• **Oxidizing agent:** Gains electrons and is reduced.
• **Reducing agent:** Loses electrons and is oxidized.

In the context of writing net ionic equations in basic solution, each reaction can be divided into two half-reactions: one showing oxidation and the other showing reduction. Each half-reaction is balanced individually before they are combined to form the balanced net ionic equation. This means identifying the substances undergoing oxidation and reduction to properly balance electron transfer.
By methodically following steps like balancing atoms and charges, and then combining the half-reactions, students can achieve the correct balanced net ionic equation for a given redox process.

Basic Solution

When balancing chemical equations in a basic solution, the chemical context changes slightly from the more common acidic conditions. Instead of using hydrogen ions ( $H^+$ ), hydroxide ions ( $OH^{-}$ ) and water ( $H_2O$ ) are used to help balance hydrogen and oxygen atoms in the equation.
A basic solution contains an excess of hydroxide ions. This requires special consideration when dealing with redox reactions, as both oxygen and hydrogen balance may be affected differently compared to an acidic environment.
Key steps in dealing with basic solutions involve:

• Balancing oxygen atoms by adding water ( $H_2O$ ) molecules to one side of the half-reaction.
• Balancing hydrogen by adding hydroxide ions ( $OH^-${} ) to the opposite side.
• Ensuring that after these adjustments, the overall charge is balanced by adding electrons as necessary.

Following these steps helps in properly accounting for the chemical environment specific to basic solutions, which is crucial for predicting the result of the reaction accurately.

Balancing Chemical Equations

Balancing chemical equations is a fundamental skill in chemistry, particularly for reactions occurring in specific environments like basic solutions. The goal is to have equal numbers of each type of atom on both sides of the equation and to ensure that the overall charge is balanced.
The process usually involves several systematic steps:

• **Identify and Separate the Half-Reactions:** Break down the complete equation into oxidation and reduction half-reactions. This simplifies the balancing process.
• **Balance Atoms Other Than O and H:** This often involves the key elements undergoing redox changes.
• **Balance Oxygen Atoms by Adding Water:** Each half-reaction should be adjusted so the number of oxygen atoms is equal on both sides.
• **Balance Hydrogen Atoms by Adding Hydroxide Ions:** Adjust for hydrogen atoms next. Use $OH^-${} for reactions in a basic solution.
• **Balance Charges by Adding Electrons:** After atoms are balanced, check the electric charges on both sides and add electrons as needed.
• **Combine and Simplify the Half-Reactions:** Multiply each half-reaction by appropriate coefficients to ensure that electrons lost and gained are equal, then combine and cancel out identical terms.

Careful attention to these steps ensures a properly balanced chemical equation, maintaining consistency and adhering to the law of conservation of mass.