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Chapter 14: Problem 142

Use the Lewis acid-base model to explain the following reaction. $$ \mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2} \mathrm{CO}_{3}(a q) $$

Short Answer

Expert verified
In the given reaction, CO2 acts as the Lewis acid, accepting a lone pair of electrons from H2O, which is the Lewis base. The electron pair donation from H2O forms a covalent bond between the carbon atom in CO2 and an oxygen atom in H2O. An unstable molecule is formed, which undergoes a proton transfer reaction, resulting in the formation of carbonic acid (H2CO3). Thus, the reaction demonstrates the Lewis acid-base model with CO2 accepting an electron pair from H2O: \[ \mathrm{CO}_{2}(g) + \mathrm{H}_{2}\mathrm{O}(l) \longrightarrow \mathrm{H}_{2}\mathrm{CO}_{3}(aq) \]

Step by step solution

01

Identify the Lewis acid and base in the reaction

According to the Lewis acid-base model, a Lewis acid is a species that can accept an electron pair, while a Lewis base is a species that can donate an electron pair. In this reaction, CO2 is the Lewis acid, as it accepts a lone pair of electrons from H2O. The H2O molecule is the Lewis base, as it donates a lone pair of electrons to CO2.
02

Explain the interaction between the Lewis acid and base

The lone pair of electrons on one of the oxygen atoms in the H2O molecule forms a covalent bond with the carbon atom in the CO2 molecule. This creates a new molecule with a carbon atom bonded to three oxygen atoms and one hydrogen atom.
03

Explain the formation of the other hydrogen bond

The new molecule formed in Step 2 is unstable, and it undergoes a proton transfer reaction. The remaining hydrogen atom from the original water molecule gets attracted to the negatively charged oxygen atom within the new molecule, forming another covalent bond. This results in the formation of the carbonic acid (H2CO3) molecule.
04

Summarize the reaction according to the Lewis acid-base model

In this reaction, CO2 (the Lewis acid) accepts an electron pair from H2O (the Lewis base). This forms an unstable molecule, which undergoes a proton transfer reaction to produce the final product H2CO3. The interaction between the electron pair donor (H2O) and electron pair acceptor (CO2) demonstrates the Lewis acid-base model in the given reaction: $$ \mathrm{CO}_{2}(g) + \mathrm{H}_{2}\mathrm{O}(l) \longrightarrow \mathrm{H}_{2}\mathrm{CO}_{3}(aq) $$

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Most popular questions from this chapter

The pH of human blood is steady at a value of approximately 7.4 owing to the following equilibrium reactions: $$ \mathrm{CO}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \rightleftharpoons \mathrm{HCO}_{3}^{-}(a q)+\mathrm{H}^{+}(a q) $$ Acids formed during normal cellular respiration react with the \(\mathrm{HCO}_{3}^{-}\) to form carbonic acid, which is in equilibrium with \(\mathrm{CO}_{2}(a q)\) and \(\mathrm{H}_{2} \mathrm{O}(l) .\) During vigorous exercise, a person's \(\mathrm{H}_{2} \mathrm{CO}_{3}\) blood levels were \(26.3 \mathrm{mM},\) whereas his \(\mathrm{CO}_{2}\) levels were 1.63 \(\mathrm{mM}\) . On resting, the \(\mathrm{H}_{2} \mathrm{CO}_{3}\) levels declined to 24.9 \(\mathrm{m} M\) . What was the \(\mathrm{CO}_{2}\) blood level at rest?

Will the following oxides give acidic, basic, or neutral solutions when dissolved in water? Write reactions to justify your answers. a. \(\mathrm{CaO}\) b. \(\mathrm{SO}_{2}\) c. \(\mathrm{Cl}_{2} \mathrm{O}\)

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