Question

Draw the following Lewis structures, showing any full charges where they exist on ions: (a) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (d) \(\mathrm{HCN}\) (b) \(\mathrm{NaNO}_{3}\) (e) \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\) (c) \(\mathrm{K}_{2} \mathrm{CO}_{3}\) (f) \(\mathrm{CH}_{3} \mathrm{COOH}\)

Short answer

Lewis structures show the arrangement of electrons in molecular and ionic compounds, with particular attention to valence electrons and charge distribution. Sulfuric acid is a molecule with sulfur in the center, HCN is a linear molecule with triple bond between carbon and nitrogen, NaNO3 consists of Na+ and NO3- ions, Al2(SO4)3 consists of Al3+ ions coordinated with sulfate ions, K2CO3 consists of K+ ions and a CO3^2- ion, and CH3COOH is acetic acid with two carbons, one double-bonded to an oxygen and one single-bonded to a hydroxyl group.

Step-by-step solution

Draw Lewis Structure for \(\mathrm{H}_{2} \mathrm{SO}_{4}\)

Start by drawing a central sulfur atom with double bonds to two oxygen atoms and single bonds to two hydroxyl groups (OH). Each oxygen atom not bonded to hydrogen has two lone pairs, while each bonded to hydrogen has one lone pair. Sulfur has no lone pairs.

Draw Lewis Structure for \(\mathrm{HCN}\)

Place the carbon atom in the center with a triple bond to the nitrogen atom, which has one lone pair, and a single bond to the hydrogen atom. Carbon has no lone pairs.

Draw Lewis Structure for \(\mathrm{NaNO}_{3}\)

Sodium ion (Na+) is separate from the nitrate ion (NO3-). Nitrogen is in the center with double bond to one oxygen and single bonds to the other two, each of which bears a lone pair. One of the singly bonded oxygen atoms has a negative charge. The nitrogen atom has one lone pair.

Draw Lewis Structure for \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}\)

Each aluminum ion (Al3+) coordinates with three sulfate ions (SO4^-2). Each sulfate has a central sulfur atom with double bonds to two oxygen atoms and single bonds to two oxygen atoms, each of these having two lone pairs and one with a negative charge.

Draw Lewis Structure for \(\mathrm{K}_{2} \mathrm{CO}_{3}\)

Potassium ions (K+) are separate from the carbonate ion (CO3^-2). Carbon is central with double bond to one oxygen atom, and single bonds to two oxygen atoms, each of which bears a lone pair. One oxygen atom has a negative charge and carbon has no lone pairs.

Draw Lewis Structure for \(\mathrm{CH}_{3} \mathrm{COOH}\)

Draw the central carbon with three single bonds to hydrogen atoms and one to another carbon. The second carbon is double-bonded to an oxygen atom and single-bonded to a hydroxyl group (OH). The oxygen double bonded to carbon has two lone pairs, and the hydroxyl oxygen has three lone pairs with a bond to hydrogen.

Key concepts

Chemical Bonding

Chemical bonding refers to the forces that hold atoms together in compounds, which become stable through the sharing or transfer of electrons. Bonds are often formed to allow atoms to reach a noble gas electron configuration, thereby achieving lower energy states and greater stability. Chemical bonds can be broadly classified into three types: ionic, covalent, and metallic.

In the provided exercise, we see examples of both ionic and covalent bonding. Ionic bonds occur in compounds such as H_{2}SO_{4}, NaNO3, Al_{2}(SO_{4})_{3}, and K_{2}CO_{3}, where a metal cation donates electrons to a non-metal anion, resulting in oppositely charged ions that attract each other. Covalent bonding, on the other hand, involves the sharing of electrons, as seen in molecules like HCN and CH3COOH, where non-metal atoms share electron pairs to complete their valence shells. Understanding the differences between these types of bonds is essential, as they significantly affect the properties of the resulting compounds. Establishing a strong grasp on these concepts allows one to predict the structures and behavior of various substances.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms within a molecule, which in turn influences the molecule's reactivity, polarity, phase of matter, color, and magnetism. The geometry is determined by the valence shell electron pair repulsion (VSEPR) theory, which posits that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion.

When we consider the molecular geometry of substances in the exercise such as H2SO4 and CH3COOH, we can see that each molecule has a different shape dictated by the number of bonds and the presence of lone electron pairs. For example, H2SO4 has a tetrahedral geometry around the sulfur atom, while HCN, being a linear molecule, has a completely different shape. Students should visualize these structures to understand how molecular shape affects properties and reactions. Connecting the Lewis structures to their molecular geometries can also help in predicting how these molecules interact with other substances and the angles at which they will bond.

Valence Electrons

Valence electrons are the outermost electrons of an atom that participate in the formation of chemical bonds. The number of valence electrons an element has can be typically found from its group number on the periodic table, which dictates its bonding behavior.

For instance, in the Lewis structure steps for H2SO4, sulfur has six valence electrons, while oxygen has six as well. Understanding how to count and distribute valence electrons is crucial in forming correct Lewis structures. The electron configuration of the valence shell is what allows atoms to bond together or stay as ions. In the given exercises, students learn to identify valence electrons to predict how an atom will bond: Sodium (Na) tends to lose one electron forming Na+, while nitrogen (N) tends to share electrons to complete its octet. The control over valence electron distribution is a key factor in understanding chemical reactivity and the physical properties of the substances involved.