Which of the following reactions would you expect to proceed at a faster rate at room temperature? Why? (Hint: Think about which reaction would have the lower activation energy.) $$2 \mathrm{Ce}^{4+}(a q)+\mathrm{Hg}_{2}^{2+}(a q) \longrightarrow 2 \mathrm{Ce}^{3+}(a q)+2 \mathrm{Hg}^{2+}(a q)$$ $$\mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{OH}^{-}(a q) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)$$

Short Answer

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The second reaction, H₃O⁺(aq) + OH⁻(aq) → 2 H₂O(l), is expected to proceed at a faster rate at room temperature, as it is an acid-base neutralization reaction that generally has a lower activation energy compared to the heavy metal redox reaction involving cerium and mercury ions in the first reaction.

Step by step solution

01

Reaction 1: Cerium and Mercury Reaction

In this reaction, we have cerium ions with a +4 charge reacting with a mercury ion with a +2 charge. The result of the reaction will be cerium ions with a +3 charge and mercury ions with a +2 charge. This reaction involves the change of oxidation states of both cerium and mercury ions.
02

Reaction 2: Hydronium and Hydroxide Reaction

In this reaction, hydronium ions (H_3O^+), which are essentially water molecules with an extra H^+ attached, react with hydroxide ions (OH^-), forming two water molecules (H_2O). This reaction is actually a classic example of an acid-base neutralization reaction, as hydronium ions are acidic and hydroxide ions are basic.
03

Comparing Activation Energies

In general, neutralization reactions, such as Reaction 2, have relatively low activation energies because acid-base reactions tend to be fast. On the other hand, Reaction 1 involves the change of oxidation states of heavy metal ions, such as cerium and mercury, which may require a higher activation energy for the reaction to proceed.
04

Conclusion

Considering the two reactions and the general trend of activation energies for acid-base reactions and heavy metal redox reactions, we can conclude that Reaction 2 (H_3O^+ + OH^- -> 2 H_2O) is expected to have a lower activation energy and therefore proceed at a faster rate at room temperature.

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Most popular questions from this chapter

A proposed mechanism for a reaction is $$\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{Br} \longrightarrow \mathrm{C}_{4} \mathrm{H}_{9}^{+}+\mathrm{Br}^{-} \quad \text { Slow }$$ $$\mathrm{C}_{4} \mathrm{H}_{9}^{+}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}_{2}^{+} \quad \text { Fast }$$ $$\mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}_{2}^{+}+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{C}_{4} \mathrm{H}_{9} \mathrm{OH}+\mathrm{H}_{3} \mathrm{O}^{+}\quad \text { Fast }$$ Write the rate law expected for this mechanism. What is the overall balanced equation for the reaction? What are the intermediates in the proposed mechanism?

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