Question

Identify the following reactions as redox or nonredox: a) \(\mathrm{Mg}(s)+\mathrm{ZnCl}_{2}(a q) \longrightarrow \mathrm{Zn}(s)+\mathrm{MgCl}_{2}(a q)\) b) \(2 \mathrm{H}_{2}(g)+\mathrm{OF}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g)+2 \mathrm{HF}(g)\) c) \(2 \mathrm{KI}(a q)+\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(a q) \longrightarrow\) \(\mathrm{Pbl}_{2}(s)+2 \mathrm{KNO}_{3}(a q)\) d) \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{Ca}(\mathrm{OH})_{2}(a q)\) e) \(3 \mathrm{CuCl}_{2}(a q)+2\left(\mathrm{NH}_{4}\right)_{3} \mathrm{PO}_{4}(a q) \longrightarrow\) \(6 \mathrm{NH}_{4} \mathrm{Cl}(a q)+\mathrm{Cu}_{3}\left(\mathrm{PO}_{4}\right)_{2}(s)\) f) \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)\)

Short answer

a) Redox, b) Redox, c) Nonredox, d) Nonredox, e) Nonredox, f) Redox

Step-by-step solution

Understand Redox Reactions

Redox reactions involve the transfer of electrons between two species. A substance that loses electrons is oxidized and the substance that gains electrons is reduced.

Identify Oxidation States

Assign oxidation states to each element in the reactants and products for each reaction. Consider the common oxidation states or use the rules for determining oxidation states.

Analyze Reaction (a)

For the reaction \(\mathrm{Mg}(s)+\mathrm{ZnCl}_{2}(a q) \longrightarrow \mathrm{Zn}(s)+\mathrm{MgCl}_{2}(a q)\)Assign oxidation states:- \(\mathrm{Mg}(s)\): 0- \(\mathrm{ZnCl}_{2}(aq)\): \(\mathrm{Zn}^{2+}\) (with \(+2\)), \(\mathrm{Cl}^{-}\) (each \(-1\))- \(\mathrm{Zn}(s)\): 0- \(\mathrm{MgCl}_{2}(aq)\): \(\mathrm{Mg}^{2+}\) (with \(+2\)), \(\mathrm{Cl}^{-}\) (each \(-1\))Identify changes:- \(\mathrm{Mg}\) is oxidized from 0 to \(+2\)- \(\mathrm{Zn}^{2+}\) is reduced from \(+2\) to 0Conclusion: This is a redox reaction.

Analyze Reaction (b)

For the reaction \(2 \mathrm{H}_{2}(g)+\mathrm{OF}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(g)+2 \mathrm{HF}(g)\)Assign oxidation states:- \(\mathrm{H}_{2}\): 0- \(\mathrm{OF}_{2}\): \(\mathrm{O}(+2)\), \(\mathrm{F}(-1)\)- \(\mathrm{H}_{2} \mathrm{O}\): \(\mathrm{H}(+1)\), \(\mathrm{O}(-2)\)- \(\mathrm{HF}\): \(\mathrm{H}(+1)\), \(\mathrm{F}(-1)\)Identify changes:- \(\mathrm{H}\) is oxidized from 0 to \(+1\)- \(\mathrm{O}\) is reduced from \(+2\) to \(-2\)Conclusion: This is a redox reaction.

Analyze Reaction (c)

For the reaction \(2 \mathrm{KI}(a q)+\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(a q) \longrightarrow\mathrm{Pbl}_{2}(s)+2 \mathrm{KNO}_{3}(a q)\)Assign oxidation states:- \(\mathrm{KI}\): \(\mathrm{K}(+1)\), \(\mathrm{I}(-1)\)- \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\): \(\mathrm{Pb}^{2+}\) (with \(+2\)), \(\mathrm{NO}_{3}^{-}\) (each \(-1\) group)- \(\mathrm{Pbl}_{2}\): \(\mathrm{Pb}(+2)\), \(\mathrm{I}^{-}\) (each \(-1\))- \(\mathrm{KNO}_{3}\): \(\mathrm{K}(+1)\), \(\mathrm{NO}_{3}^{-}\) (each \(-1\) group)No changes in oxidation states are occurring.Conclusion: This is a nonredox reaction.

Analyze Reaction (d)

For the reaction \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\mathrm{Ca}(\mathrm{OH})_{2}(a q)\)Assign oxidation states:- \(\mathrm{CaO}\): \(\mathrm{Ca}(+2)\), \(\mathrm{O}(-2)\)- \(\mathrm{H}_{2} \mathrm{O}\): \(\mathrm{H}(+1)\), \(\mathrm{O}(-2)\)- \(\mathrm{Ca}(\mathrm{OH})_{2}\): \(\mathrm{Ca}(+2)\), \(\mathrm{OH}^{-}\) (each \(-1\) group)No changes in oxidation states are occurring.Conclusion: This is a nonredox reaction.

Analyze Reaction (e)

For the reaction \(3 \mathrm{CuCl}_{2}(a q)+2\left(\mathrm{NH}_{4}\right)_{3}\mathrm{PO}_{4}(a q) \longrightarrow6 \mathrm{NH}_{4} \mathrm{Cl}(a q)+\mathrm{Cu}_{3}\left(\mathrm{PO}_{4}\right)_{2}(s)\)Assign oxidation states:- \(\mathrm{CuCl}_{2}\): \(\mathrm{Cu}(+2)\), \(\mathrm{Cl}(-1)\)- \(\left(\mathrm{NH}_{4}\right)_{3}\mathrm{PO}_{4}\): \(\mathrm{NH}_{4}(+1)\), \(\mathrm{PO}_{4}(-3)\)- \(\mathrm{NH}_{4} \mathrm{Cl}\): \(\mathrm{NH}_{4}(+1)\), \(\mathrm{Cl}(-1)\)- \(\mathrm{Cu}_{3}\left(\mathrm{PO}_{4}\right)_{2}\): \(\mathrm{Cu}(+2)\), \(\mathrm{PO}_{4}(-3)\)No changes in oxidation states are occurring.Conclusion: This is a nonredox reaction.

Analyze Reaction (f)

For the reaction \(\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(g)\)Assign oxidation states:- \(\mathrm{CH}_{4}\): \(\mathrm{C}(-4)\), \(\mathrm{H}(+1)\)- \(\mathrm{O}_{2}\): 0- \(\mathrm{CO}_{2}\): \(\mathrm{C}(+4)\), \(\mathrm{O}(-2)\)- \(\mathrm{H}_{2} \mathrm{O}\): \(\mathrm{H}(+1)\), \(\mathrm{O}(-2)\)Identify changes:- \(\mathrm{C}\) is oxidized from \(-4\) to \(+4\)- \(\mathrm{O}\) is reduced from 0 to \(-2\)Conclusion: This is a redox reaction.

Key concepts

Oxidation States

Understanding oxidation states is crucial to determining whether a reaction is a redox reaction or not. Each element in a compound or molecule has an oxidation state, which helps in tracking electron transfer.
Oxidation states are typically assigned based on a set of rules:

• Elements in their pure form have oxidation states of 0. For example, \(\text{H}_2, \text{O}_2 \) and metallic elements.
• Hydrogen generally has an oxidation state of \(+1\) and oxygen \(-2\) in most compounds.
• In ions, the oxidation state is equal to the charge of the ion. For example, in NaCl, Na has \(+1\) and Cl has \(-1\).
• For polyatomic ions, the sum of the oxidation states must equal the ion's charge.

Knowing these oxidation states allows you to see the gain and loss of electrons in chemical reactions. Clearly see which atoms are oxidized (increase in oxidation state) and which are reduced (decrease in oxidation state). This helps you identify redox reactions accurately.

Electron Transfer

Redox reactions are driven by electron transfer between species. One element loses electrons (oxidation), and another gains electrons (reduction). Understanding electron transfer can help you grasp how chemical reactions occur on a microscopic level.

• Oxidation is the loss of electrons. For example, in \(Mg(s) \rightarrow Mg^{2+}(aq) + 2e^-\), magnesium is oxidized.
• Reduction is the gain of electrons. For example, in \(Zn^{2+}(aq) + 2e^- \rightarrow Zn(s)\), zinc is reduced.

An easy way to remember this is the acronym 'OIL RIG' (Oxidation Is Loss, Reduction Is Gain). By using the rules for oxidation states, you can determine how electrons are moving in a reaction. This movement tells you whether a reaction involves electron transfer, hence identifying it as a redox reaction.

Chemical Reactions

Chemical reactions can be categorized into redox and nonredox reactions. Understanding this distinction is essential for analyzing reactions correctly.

Redox Reactions

Redox reactions involve changes in oxidation states and electron transfer. For example, \(Mg(s) + ZnCl_2(aq) \rightarrow Zn(s) + MgCl_2(aq)\) is a redox reaction. Magnesium is oxidized and zinc is reduced.

Nonredox Reactions

Nonredox reactions, on the other hand, do not involve changes in oxidation states. For instance, the reaction \(2 KI(aq) + Pb(NO_3)_2(aq) \rightarrow PbI_2(s) + 2 KNO_3(aq)\) is a nonredox reaction. Here, no electrons are transferred, and oxidation states remain the same.

Understanding the nature of chemical reactions helps you predict product formation, energy changes, and reaction mechanisms. It is fundamental for mastering chemistry concepts and solving complex problems with ease.