Question

Identify correct statements for the following incomplete reactions. (i) \(\mathrm{KI}+\mathrm{Cl}_{2} \rightarrow\) (ii) \(\mathrm{KClO}_{3}+\mathrm{I}_{2} \rightarrow\) (iii) \(\mathrm{KNO}_{3(3)}+{ }^{3} \mathrm{~K}_{(s)} \rightarrow\) (iv) \(\mathrm{P}+\mathrm{I}_{2}+\mathrm{H}_{2} \mathrm{O} \rightarrow\) (Red) (a) All are redox reactions (b) Only I, II and III have one of the product in its element form (c) I and II are displacement reactions (d) III is disproportionation reaction

Short answer

Statements (a), (b) are likely correct upon detailed reaction review and completion.

Step-by-step solution

Analyze Reaction (i)

For the reaction \( \mathrm{KI} + \mathrm{Cl}_{2} \rightarrow \), potassium iodide (KI) reacts with chlorine gas \( \mathrm{Cl}_{2} \). This is a type of halogen displacement reaction where \( \mathrm{Cl}_{2} \) displaces iodine from \( \mathrm{KI} \) to form \( \mathrm{KCl} \) and \( \mathrm{I}_{2} \). Therefore, the reaction is \( \mathrm{KI} + \mathrm{Cl}_{2} \rightarrow \mathrm{KCl} + \mathrm{I}_{2} \).

Analyze Reaction (ii)

In the reaction \( \mathrm{KClO}_{3} + \mathrm{I}_{2} \rightarrow \), usually iodine can be reduced. However, the reaction setup is missing specifics on its products. In a typical redox reaction, \( \mathrm{KClO}_{3} \) acts as an oxidizer. This is more complex but can be redox.

Analyze Reaction (iii)

For the reaction \( \mathrm{KNO}_{3} + \mathrm{K}_{(s)} \rightarrow \), this seems erroneous due to notation \( \mathrm{KNO}_{3(3)} + { }^{3} \) but considering a similar correct form, it typically involves redox where \( \mathrm{KNO}_{3} \) undergoes reduction. Possible reaction results in products like \( \mathrm{N}_{2} \) (an element form).

Analyze Reaction (iv)

In the reaction \( \mathrm{P} + \mathrm{I}_{2} + \mathrm{H}_{2} \mathrm{O} \rightarrow \), phosphorus reacts with iodine and water. Here, iodine could be reduced and depending on conditions, phosphorus might be oxidized - indicating redox potential. Reaction specifics and media determine exact products.

Evaluate Statements

(a) True: All involve potential redox character where an element may experience oxidization or reduction. (b) True: In reaction (i), (ii), (iii) - each has iodine or nitrogen respectively potentially in elemental form. (c) Partially True: Indeed, reaction (i) is displacement but (ii) lacks details to confirm fully. (d) Unlikely: From setup, (iii) more likely stays redox, rapid validation needed.

Key concepts

Halogen Displacement Reaction

Halogen displacement reactions occur when a more reactive halogen displaces a less reactive halogen from a compound. This concept is demonstrated in the first reaction example, where chlorine gas \( \mathrm{Cl}_{2} \) displaces iodine from potassium iodide \( \mathrm{KI} \). Chlorine is more reactive than iodine, allowing it to replace iodine, resulting in potassium chloride \( \mathrm{KCl} \) and free iodine \( \mathrm{I}_{2} \).

• Chlorine displaces iodine due to its higher reactivity.
• The resulting products are \( \mathrm{KCl} \) and \( \mathrm{I}_{2} \).

This type of reaction is a common example in chemistry to illustrate reactivity trends in halogens. Displacement reactions highlight the ability of elements higher in the periodic table (within the same group) to displace those below them.

Disproportionation Reaction

A disproportionation reaction is a specific type of redox reaction where a single element is simultaneously oxidized and reduced, forming two different products. While the initial exercise considers if Reaction (iii) might be disproportionation, the notation suggests otherwise. It's crucial to note that true disproportionation involves one element changing to two oxidation states.
In possible scenarios, if such a reaction existed, you would observe:

• A compound containing one element in an intermediate oxidation state.
• The formation of two compounds, where the element is both oxidized and reduced.

For example, in a classical disproportionation of chlorine in water, \( \mathrm{Cl}_{2} + \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{HCl} + \mathrm{HClO} \), chlorine undergoes both oxidation to hypochlorous acid and reduction to hydrochloric acid.

Oxidation-Reduction (Redox)

Redox reactions involve the transfer of electrons between substances, whereby one element is oxidized (loses electrons) and another is reduced (gains electrons). In the context of the given exercises, all potential reactions suggested elements that could undergo oxidation or reduction, re-affirming their redox nature.
Key points to understand about redox reactions are:

• Reduction is indicated by a gain in electrons, commonly seen as a decrease in oxidation state.
• Oxidation occurs when an element loses electrons, leading to an increase in oxidation state.
• All reactions where electrons transfer between reactants are categorized as redox.

For instance, in reaction (i), iodine is reduced to free iodine \( \mathrm{I}_{2} \), while chlorine displaces it, which can be perceived as a classic outcome in redox reactions.

Chemistry Reaction Analysis

Chemistry reaction analysis involves breaking down a chemical reaction into its fundamental parts to understand the processes involved. By examining each reactant and product, we glean insights into reaction types and mechanisms. Every analysis should include:

• Determining the reactants and products to identify transformation.
• Assessing changes in oxidation states to discern redox events.
• Recognizing the nature of the reaction, such as displacement or combination.
• Considering conditions like pH and temperature, which affect the reaction pathway.

For example, scrutinizing the halogen displacement and possible redox processes in the given exercise helps in predicting the resulting products and understanding the reactants' reactivity. Chemistry reaction analysis aids in building predictive models in chemical reactivity and reaction feasibility.