Use Lewis structures to show the Lewis acid-base reaction between \(\mathrm{SO}_{3}\) and \(\mathrm{H}_{2} \mathrm{O}\) to give \(\mathrm{H}_{2} \mathrm{SO}_{4}\). Identify the Lewis acid and the Lewis base in the reaction.

Lewis structures, also known as electron dot structures, are diagrams that represent the valence electrons of atoms within a molecule. They are an invaluable tool in visualizing the distribution of electrons and the arrangement of atoms. Understanding Lewis structures begins with identifying the total number of valence electrons in the molecule. Each dot in a Lewis structure represents a valence electron, while lines symbolize covalent bonds between atoms.

When drawing Lewis structures, it's essential to follow a sequence of steps to ensure accuracy. First, count the valence electrons from all atoms in the molecule. Next, choose a central atom (usually the least electronegative one) and arrange the other atoms around it. After that, place a single bond (two electrons) between the central atom and the surrounding atoms. Distribute the remaining electrons to complete the octets (or duets for hydrogen) for each atom. If there are still electrons left or if there are atoms with an incomplete octet, you may need to form double or triple bonds.

In our exercise, for the Lewis structure of \texttt{SO\(_3\)}, sulfur is the central atom, double-bonded to three oxygen atoms. Since sulfur can have an expanded octet, this molecule does not require all atoms to follow the octet rule strictly. For \texttt{H\(_2\)O}, the oxygen atom is central with two lone pairs and forms single bonds with hydrogen atoms.

The concept of electron pair acceptor and donor lays the groundwork for understanding a Lewis acid-base reaction. An electron pair donor, also known as a Lewis base, has a lone pair of electrons and is willing to share these with another atom or molecule. Conversely, a Lewis acid is an electron pair acceptor and can accept an additional pair of electrons to form a new chemical bond.

In the given exercise, \texttt{H\(_2\)O} functions as a Lewis base because it has lone pairs of electrons on the oxygen atom that are not shared or bonded with any other atom. These lone pairs are the 'electron currency' that the base offers in the chemical transaction. On the other hand, \texttt{SO\(_3\)} is a Lewis acid; despite having a double bond structure, it has an empty orbital that can accommodate another electron pair, effectively allowing it to accept electrons.

This interaction between the electron pair donor (\texttt{H\(_2\)O}) and the electron pair acceptor (\texttt{SO\(_3\)}) leads to the formation of \texttt{H\(_2\)SO\(_4\)} as the lone pair from the oxygen atom in water is shared with the sulfur atom in \texttt{SO\(_3\)}, completing the acid-base reaction.

Chemical bonding is the physical phenomenon of chemical substances being held together by the attraction of atoms to each other through sharing, as well as exchanging, of electrons - known as covalent and ionic bonds, respectively. The Lewis acid-base reaction can be seen as a special case of covalent bonding where an electron pair is donated by one species (the base) and accepted by another (the acid) to form a covalent bond.

This type of bonding is central to chemistry because it determines the structure and properties of molecules. For example, the difference in electron sharing leads to the distinction between polar and nonpolar molecules. Covalent bonds can also have different bond strengths, which are reflected by single, double, and triple bond formations.

In the reaction between \texttt{SO\(_3\)} and \texttt{H\(_2\)O}, the electron pair from \texttt{H\(_2\)O} is shared with \texttt{SO\(_3\)} to form \texttt{H\(_2\)SO\(_4\)}. This sharing of electrons creates a new covalent bond, illustrating the dynamic nature of chemical bonding. The formation of \texttt{H\(_2\)SO\(_4\)} shows the conversion of a pi bond in \texttt{SO\(_3\)} to a sigma bond in \texttt{H\(_2\)SO\(_4\)}, changing the bonding structure and consequently the properties of the molecule.