Question

Write the electron dot formula and draw the structural formula for each of the following polyatomic ions: (a) \(\mathrm{SO}_{4}^{2-}\) (b) \(\mathrm{HSO}_{4}\) (c) \(\mathrm{SO}_{3}^{2-}\) (d) \(\mathrm{HSO}_{3}^{2-}\)

Short answer

Draw Lewis structures using electron count for each ion and depict bonds among atoms.

Step-by-step solution

Understanding Electron Dot Formula

The electron dot formula, or Lewis structure, represents the valence electrons of atoms within a molecule or ion using dots. We'll start by determining the total number of valence electrons for each given polyatomic ion.

Group Valence Electrons for SO4^2-

In \( \mathrm{SO}_{4}^{2-} \), sulfur (S) has 6 valence electrons, each oxygen (O) has 6 valence electrons, and the ion has a 2- charge, adding 2 more electrons. Total = 6 + (6×4) + 2 = 32 electrons.

Draw Lewis Structure for SO4^2-

Place sulfur in the center and surround it with four oxygens. Create single bonds between S and each O (using 8 electrons). Then, distribute the remaining 24 electrons to satisfy the octet rule, resulting in double bonds with two oxygens (forming resonance structures) to account for full octets.

Group Valence Electrons for HSO4^-

In \( \mathrm{HSO}_{4}^{-} \), add 1 electron for hydrogen, maintaining the \( \mathrm{SO}_{4}^{2-} \) foundation. Total = 32 (from \( \mathrm{SO}_{4}^{2-} \)) + 1 - 1 (lost for H bonding) = 32 electrons.

Draw Lewis Structure for HSO4^-

Assign hydrogen to one of the oxygens of \( \mathrm{SO}_{4}^{2-} \) through a single bond (using one electron), then adjust electron pairs similar to that of \( \mathrm{SO}_{4}^{2-} \) to maintain octet configuration.

Group Valence Electrons for SO3^2-

For \( \mathrm{SO}_{3}^{2-} \), sulfur has 6 electrons, each oxygen supplies 6 electrons, and the ion's charge adds 2 electrons. Total = 6 + (3×6) + 2 = 26 electrons.

Draw Lewis Structure for SO3^2-

Place sulfur in the center with three surrounding oxygens. Connect with single bonds (6 electrons used) and place remaining 20 electrons to form double bonds (with possible resonance structures) to satisfy the octet rule.

Group Valence Electrons for HSO3^2-

In \( \mathrm{HSO}_{3}^{2-} \), repeat SO3 counting but add H's 1 valence without losing any due to charge (-2). Total = 26 (from \( \mathrm{SO}_{3}^{2-} \)) + 1 = 27 electrons.

Draw Lewis Structure for HSO3^2-

Attach hydrogen to one oxygen, completing the structure by arranging remaining electrons around sulfur and oxygens to fulfill octet requirements with single/double bonds.

Key concepts

Electron Dot Formulas

Electron dot formulas, also known as Lewis structures, are a visual way to represent the valence electrons of atoms within a molecule or ion. These diagrams use dots placed around the chemical symbols to illustrate each valence electron involved in the formation of chemical bonds. This approach helps in visualizing the arrangement of electrons around each atom, and predicting how bonds form in a molecule.
For example, consider the polyatomic ion sulfate, \( \mathrm{SO}_{4}^{2-} \). The sulfur atom, being the central atom, is surrounded by four oxygen atoms. The electron dot formula starts by calculating the total number of valence electrons available. This includes accounting for the ion charge. Applying this to \( \mathrm{SO}_{4}^{2-} \), we sum up the valence electrons from sulfur, the oxygen atoms, and the additional electrons from the overall negative charge. Understanding where each electron goes helps in correctly representing the molecule's electronic structure, following steps to show bonding and lone pairs around each element.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding. They are primarily responsible for the chemical properties of elements because these are the electrons that participate in forming bonds. While drawing electron dot formulas, it’s essential to correctly identify and account for these electrons to accurately depict molecule structure.
In the sulfate ion \( \mathrm{SO}_{4}^{2-} \), each of the atoms has a certain number of valence electrons: sulfur has 6, and each oxygen also has 6. The 2- charge on the ion adds two more electrons to the total. Thus, understanding the count and distribution of valence electrons allows us to illustrate stable structures that follow chemical rules. Proper allocation of valence electrons is crucial, as it dictates the configuration of single and double bonds in resonance forms of the molecule.

Polyatomic Ions

Polyatomic ions are charged species composed of two or more atoms covalently bonded, or of a metal complex that acts as a single unit. These ions carry a net charge because their total number of protons does not equal the total number of electrons. When writing Lewis structures for polyatomic ions, it's important to consider their overall charge, which affects electron count and distribution.
For instance, in drawing the electron dot formula for the ion \( \mathrm{HSO}_{4}^{-} \), we begin by calculating total valence electrons including those from the negative charge. This calculation influences the connections between atoms, helping us determine the placement of hydrogen and the distribution of charge over the structurally important oxygens. By accurately representing polyatomic ions, we can predict how they might interact in reactions and form meaningful chemical insights.

Octet Rule

The octet rule is a cornerstone principle in chemistry stating that atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell, achieving the electronic configuration of a noble gas. This rule acts as a guideline to ensure stability in molecules and ions. In practice, the octet rule helps dictate how atoms share electrons to form single or multiple bonds.
In ions like \( \mathrm{SO}_{3}^{2-} \) and \( \mathrm{HSO}_{3}^{2-} \), the distribution of electrons adheres to the octet rule to ensure all atoms reach a stable electron configuration. By following the octet rule, we can predict the formation of double bonds in resonance structures or the need to rearrange electrons around atoms that are close to fulfilling their octet. Although there are exceptions, such as elements from the third period and beyond, the octet rule remains a reliable tool for understanding and predicting molecular structure and reactivity.