Write balanced net ionic equations for the following...

Write balanced net ionic equations for the following reactions in acidic solution: (a) $\mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{Al}(s) \longrightarrow \mathrm{H}_{2} \mathrm{S}(a q)+\mathrm{A}^{3+}(a q)$ (b) $\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \longrightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{Cr}^{3+}(a q)$ (c) $\operatorname{ClO}_{3}^{-}(a q)+\operatorname{As}_{2} S_{3}(s) \longrightarrow_{\mathrm{Cl}^{-}(a q)+\mathrm{H}_{2} \mathrm{AsO}_{4}^{-}(a q)+\mathrm{HSO}_{4}^{-}(a q)}$ (d) $\mathrm{IO}_{3}^{-}(a q)+\operatorname{Re}(s) \longrightarrow \operatorname{ReO}_{4}^{-}(a q)+\mathrm{I}^{-}(a q)$ (e) $\mathrm{HSO}_{4}^{-}(a q)+\mathrm{As}_{4}(s)+\mathrm{Pb}_{3} \mathrm{O}_{4}(s) \rightarrow _{\mathrm{PbSO}_{4}(s)+\mathrm{H}_{2} \mathrm{AsO}_{4}^{-}(a q)}$ (f) $\mathrm{HNO}_{2}(a q) \longrightarrow \mathrm{NO}_{3}^{-}(a q)+\mathrm{NO}(g)$

Write balanced net ionic equations for the following reactions in acidic solution:
(a) $\mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{Al}(s) \longrightarrow \mathrm{H}_{2} \mathrm{S}(a q)+\mathrm{A}^{3+}(a q)$
(b) $\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \longrightarrow \mathrm{S}_{4} \mathrm{O}_{6}^{2-}(a q)+\mathrm{Cr}^{3+}(a q)$
(c) $\operatorname{ClO}_{3}^{-}(a q)+\operatorname{As}_{2} S_{3}(s) \longrightarrow_{\mathrm{Cl}^{-}(a q)+\mathrm{H}_{2} \mathrm{AsO}_{4}^{-}(a q)+\mathrm{HSO}_{4}^{-}(a q)}$
(d) $\mathrm{IO}_{3}^{-}(a q)+\operatorname{Re}(s) \longrightarrow \operatorname{ReO}_{4}^{-}(a q)+\mathrm{I}^{-}(a q)$
(e) $\mathrm{HSO}_{4}^{-}(a q)+\mathrm{As}_{4}(s)+\mathrm{Pb}_{3} \mathrm{O}_{4}(s) \rightarrow _{\mathrm{PbSO}_{4}(s)+\mathrm{H}_{2} \mathrm{AsO}_{4}^{-}(a q)}$
(f) $\mathrm{HNO}_{2}(a q) \longrightarrow \mathrm{NO}_{3}^{-}(a q)+\mathrm{NO}(g)$

Key Concepts

Redox Reactions

Redox reactions involve the transfer of electrons between species, which results in changes in oxidation states. In these reactions, one species undergoes oxidation (loss of electrons) and another undergoes reduction (gain of electrons). Understanding redox principles is essential for identifying the processes occurring in the chemical equation and for applying appropriate balancing techniques.

Oxidation States

Oxidation states, or oxidation numbers, help determine which atoms undergo oxidation or reduction by tracking the loss or gain of electrons. Assigning correct oxidation numbers is a fundamental step in identifying the electron transfer and, consequently, in setting up balanced half-reactions for a redox process.

Net Ionic Equations

Net ionic equations represent the essential chemical change by omitting spectator ions and focusing on the species that directly participate in the reaction. This simplification highlights the fundamental redox changes and is particularly useful in aqueous reactions where many ions may be present but do not change oxidation state.

Half-Reaction Method

The half-reaction method is a technique for balancing redox reactions by splitting the overall reaction into two parts: one for oxidation and one for reduction. Each half-reaction is balanced separately for mass and charge, including the addition of water, hydrogen ions, and electrons as needed in acidic solutions, before recombining them to form the balanced overall equation.

Balancing in Acidic Medium

Balancing redox reactions in an acidic solution requires the addition of H+ ions and water molecules to achieve balance for both mass and charge. The acidic medium provides a source of protons (H+) that can be used to balance oxygen and hydrogen atoms, making it a distinctive aspect of balancing equations under these conditions.

Transcript

Please give Ace some feedback