Select the nature or type of redox change in the following reactions: (ii) \(2 \mathrm{Cu}^{+} \longrightarrow \mathrm{Cu}^{2+}+\mathrm{Cu}^{0}\) (b) \(\quad \mathrm{Cl}_{2} \longrightarrow \mathrm{ClO}^{-}+\mathrm{Cl}^{-}\) (c) \(2 \mathrm{KClO}_{3} \stackrel{\Delta}{\longrightarrow} 2 \mathrm{KCl}+3 \mathrm{O}_{2}\) (d) \(\quad\left(\mathrm{NH}_{4}\right)_{2} \mathrm{Cr}_{2} \mathrm{O}_{7} \longrightarrow \mathrm{N}_{2}+\mathrm{Cr}_{2} \mathrm{O}_{3}+4 \mathrm{H}_{2} \mathrm{O}\) (e) \(10 \mathrm{FeSO}_{4}+2 \mathrm{KMnO}_{4}+8 \mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow 2 \mathrm{MnSO}_{4}+5 \mathrm{Fe}_{2}\left(\mathrm{SO}_{4}\right)_{3}+\) \(\mathrm{K}_{2} \mathrm{SO}_{4}+8 \mathrm{H}_{2} \mathrm{O}\) (f) \(5 \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}+2 \mathrm{KMnO}_{4}+3 \mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{K}_{2} \mathrm{SO}_{4}+2 \mathrm{MnSO}_{4}+\) \(10 \mathrm{CO}_{2}+8 \mathrm{H}_{2} \mathrm{O}\)

Step by step solution

01

- Determine Oxidation States for Cu

In the reaction (ii) 2 Cu+ → Cu2+ + Cu0, identify the oxidation states of copper before and after the reaction. Before the reaction, copper has an oxidation state of +1 (Cu+). After the reaction, copper appears in two different forms: Cu2+ has an oxidation state of +2, and Cu0 has an oxidation state of 0.

02

- Analyze the Redox Process for Cu

Since one mole of Cu+ is losing one electron to become Cu2+ (oxidation process) while the other mole of Cu+ is gaining one electron to become Cu0 (reduction process), both oxidation and reduction are taking place in the same reaction. Therefore, it is a disproportionation redox reaction.

03

- Determine Oxidation States for Chlorine in Cl2

In the reaction (b) Cl2 → ClO- + Cl-, identify the oxidation states. Elemental chlorine (Cl2) has an oxidation state of 0. The chlorine in ClO- (hypochlorite ion) has an oxidation state of +1, and the chlorine in Cl- (chloride ion) has an oxidation state of -1.

04

- Analyze the Redox Process for Cl2

In the reaction (b) since chlorine undergoes an increase in oxidation state from 0 to +1 in ClO-, and a decrease from 0 to -1 in Cl-, this is also a disproportionation reaction, where the same element is both oxidized and reduced.

05

- Identify the Decomposition Reaction in (c)

Reaction (c) 2 KClO3 → 2 KCl + 3 O2 is a decomposition reaction. Upon heating, potassium chlorate decomposes into potassium chloride and oxygen gas. This reaction does not fit the typical pattern of a redox reaction and is not typically classified as such.

06

- Identify Products of Decomposition

In reaction (c), KClO3 produces KCl and O2. Oxygen starts with an oxidation state of -2 in KClO3 and ends up as O2 with an oxidation state of 0. This implies a redox change for oxygen in which it is being reduced. Potassium and chlorine maintain a constant oxidation state. Therefore, the reaction involves a redox process.

07

- Identify Redox Reaction in (d)

Reaction (d) (NH4)2Cr2O7 → N2 + Cr2O3 + 4 H2O is a redox reaction. Ammonium dichromate undergoes thermal decomposition to form nitrogen gas, chromium(III) oxide, and water. This reaction involves changes in oxidation states for both nitrogen and chromium, and thus it's a redox reaction.

08

- Analyze Complex Redox Reaction in (e)

In reaction (e) 10 FeSO4 + 2 KMnO4 + 8 H2SO4 → 5 Fe2(SO4)3 + 2 MnSO4 + K2SO4 + 8 H2O, iron is oxidized from Fe2+ in FeSO4 to Fe3+ in Fe2(SO4)3, and manganese is reduced from Mn7+ in KMnO4 to Mn2+ in MnSO4. This is a complex redox reaction involving electron transfer between different species.

09

- Analyze Complex Redox Reaction in (f)

In reaction (f) 5 H2C2O4 + 2 KMnO4 + 3 H2SO4 → K2SO4 + 2 MnSO4 + 10 CO2 + 8 H2O, oxalic acid (H2C2O4) is oxidized to CO2, and KMnO4 is reduced to MnSO4. This is another complex redox reaction where electron transfer occurs between reactants.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Disproportionation Reaction

In the world of chemistry, a disproportionation reaction is a fascinating and unique process in which a single element simultaneously undergoes oxidation and reduction, transforming into at least two different products with distinct oxidation states. This transformation can be envisioned as a balancing act, where the element donates and accepts electrons within the same reaction.

For example, take the transformation of copper(I) ion, where the reaction is described as follows: \(2 \text{Cu}^+ \longrightarrow \text{Cu}^{2+}+\text{Cu}^{0}\). Here, the copper in \(\text{Cu}^+\) has an oxidation state of +1 but emerges in two forms: \(\text{Cu}^{2+}\) with an oxidation state of +2 (oxidation) and \(\text{Cu}^{0}\) with an oxidation state of 0 (reduction). This reaction is a classic case of disproportionation as copper is both losing and gaining electrons simultaneously.

Oxidation State Determination

Determining oxidation states is key to understanding redox reactions. Oxidation state, sometimes referred to as oxidation number, is a conceptual charge assigned to an atom assuming complete transfer of electrons. It helps us keep track of electron movement and understand which substances are undergoing oxidation (loss of electrons) or reduction (gain of electrons).

Let's demystify this with an example from a common disinfectant, chlorine. In the transformation \(\text{Cl}_2 \longrightarrow \text{ClO}^-+\text{Cl}^-\), elemental chlorine (\text{Cl}_2) has an oxidation state of 0. In the products, the chlorine in \(\text{ClO}^-\) (hypochlorite ion) has an oxidation state of +1, indicating it has accepted an electron, and chlorine in \(\text{Cl}^-\) has an oxidation state of -1, indicating it has donated an electron, this variation is definitive evidence of a redox process.

Decomposition Reaction

A decomposition reaction is akin to a magic trick in chemistry, where a compound splits into two or more simpler substances. It often applies heat, electricity, or another form of energy to 'break the spell' that holds together the ingredients in a single compound. For instance, when potassium chlorate (\text{KClO}_3) decomposes upon heating, it divides into potassium chloride (KCl) and free oxygen (O_2) as in: \(2 \text{KClO}_3 \rightarrow 2 \text{KCl} + 3 \text{O}_2\). While decomposition doesn't always involve redox changes, this particular reaction does feature a redox process—oxygen's oxidation state decreases from -2 in \text{KClO}_3 to 0 in \text{O}_2, indicative of reduction. Potassium and chlorine, however, steadfastly maintain their oxidation states throughout this transformation, showcasing that not all elements in a decomposition reaction undergo redox changes.

Complex Redox Processes

Complex redox processes are essentially intricate dance routines of electrons among participating species. They often involve multiple oxidations and reductions, and typically, reactions in this category are a bit more elaborate and challenging to balance. An illustrative example of this complexity can be seen in the reaction between ferrous sulfate (\text{FeSO}_4), potassium permanganate (\text{KMnO}_4), and sulfuric acid (\text{H}_2\text{SO}_4), which is represented by the equation: \(10 \text{FeSO}_4 + 2 \text{KMnO}_4 + 8 \text{H}_2\text{SO}_4 \longrightarrow 5 \text{Fe}_2(\text{SO}_4)_3 + 2 \text{MnSO}_4 + \text{K}_2\text{SO}_4 + 8 \text{H}_2\text{O}\). Here, iron transitions from an oxidation state of +2 to +3 (oxidation), while manganese changes from +7 to +2 (reduction). Such reactions showcase the interplay of different elements changing their oxidation states by giving and accepting electrons in a network of chemical shifts.