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

The given reaction consists of the following species: \(\mathrm{CO}_{2}\), \(\mathrm{H}_{2} \mathrm{O}\), and \(\mathrm{H}_{2}\mathrm{CO}_{3}\). First, let's examine the structures of these molecules to help us identify the electron-pair donor and acceptor. The structure of \(\mathrm{CO}_{2}\) is linear, with carbon in the middle and a double bond to each oxygen atom. Oxygen is more electronegative than carbon, resulting in a partial positive charge on the carbon atom. This makes the central carbon atom in \(\mathrm{CO}_{2}\) an electron-pair acceptor, therefore a Lewis acid. The structure of \(\mathrm{H}_{2} \mathrm{O}\) is bent, with oxygen in the middle and a single bond to each hydrogen atom. Oxygen has two lone pairs of electrons and can donate an electron pair, making \(\mathrm{H}_{2} \mathrm{O}\) a Lewis base. Now let's analyze the reaction step by step.

In this reaction, the Lewis base \(\mathrm{H}_{2} \mathrm{O}\) donates an electron pair to the Lewis acid \(\mathrm{CO}_{2}\), forming a new bond between the oxygen atom of \(\mathrm{H}_{2} \mathrm{O}\) and the carbon atom of \(\mathrm{CO}_{2}\), leading to the formation of \(\mathrm{H}_{2}\mathrm{CO}_{3}\). $$ \mathrm{CO}_{2}(g) + \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{2}\mathrm{CO}_{3}(a q) $$ To summarize, according to the Lewis acid-base model, \(\mathrm{CO}_{2}\) acts as the Lewis acid (electron-pair acceptor) and \(\mathrm{H}_{2} \mathrm{O}\) acts as the Lewis base (electron-pair donor) in this reaction. The products are formed after the electron-pair donation from the oxygen atom of \(\mathrm{H}_{2} \mathrm{O}\) to the central carbon atom of \(\mathrm{CO}_{2}\), which forms a new bond and produces \(\mathrm{H}_{2} \mathrm{CO}_{3}\).