Question

Write Lewis structures that obey the octet rule for each of the following molecules and ions. (In each case the first atom listed is the central atom.) a. \(\mathrm{POCl}_{3}, \mathrm{SO}_{4}^{2-}, \mathrm{XeO}_{4}, \mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{4}^{-}\) b. \(\mathrm{NF}_{3}, \mathrm{SO}_{3}^{2-}, \mathrm{PO}_{3}^{3-}, \mathrm{ClO}_{3}^{-}\) c. \(\mathrm{ClO}_{2}-, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\) d. Considering your answers to parts a, b, and c, what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons?

Short answer

In summary, the Lewis structures of the given molecules and ions obey the octet rule. For species with the same number of atoms and valence electrons, their shapes depend mainly on the arrangement of atoms and electron pairs around the central atom. However, such species might not necessarily have the same shape, as the arrangement of electrons can be influenced by differences in electronegativity and the central atom's ability to expand its valence shell.

Step-by-step solution

a. Lewis structures for \(\mathrm{POCl}_{3}, \mathrm{SO}_{4}^{2-}, \mathrm{XeO}_{4}, \mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{4}^{-}\)

1. First, count the total number of valence electrons for each molecule or ion. 2. Place the least electronegative atom as the central atom. 3. Draw single bonds between the central atom and surrounding atoms. 4. Distribute the remaining electrons to satisfy the octet rule for all atoms. The Lewis structures for each molecule or ion are as follows: \(\mathrm{POCl}_{3}\): \[ \chemfig{P(-[:0]O)(-[:-90]Cl)(-[:90]Cl)(-[:180]Cl)} \] \(\mathrm{SO}_{4}^{2-}\): \[ \chemfig{S(=[:0]O)(=[:90]O)(=[:180]O)(=[:-90]O)}^{-2} \] \(\mathrm{XeO}_{4}\): \[ \chemfig{Xe(=[:0]O)(=[:90]O)(=[:180]O)(=[:-90]O)} \] \(\mathrm{PO}_{4}^{3-}\): \[ \chemfig{P(-[:0]O)(-[:90]O)(-[:180]O)(-[:-90]O)}^{3-} \] \(\mathrm{ClO}_{4}^{-}\): \[ \chemfig{Cl(-[:0]O)(-[:90]O)(-[:180]O)(-[:-90]O)}^{-} \]

b. Lewis structures for \(\mathrm{NF}_{3}, \mathrm{SO}_{3}^{2-}, \mathrm{PO}_{3}^{3-}, \mathrm{ClO}_{3}^{-}\)

Follow the same procedure as mentioned earlier for each molecule or ion: \(\mathrm{NF}_{3}\): \[ \chemfig{N(-[1]F)(-[3]F)(-[5]F)} \] \(\mathrm{SO}_{3}^{2-}\): \[ \chemfig{S(-[:0]O)(-[:120]O)(-[:-120]O)}^{2-} \] \(\mathrm{PO}_{3}^{3-}\): \[ \chemfig{P(-[:0]O)(-[:90]O)(-[:-90]O)}^{3-} \] \(\mathrm{ClO}_{3}^{-}\): \[ \chemfig{Cl(-[:0]O)(-[:90]O)(-[:-90]O)}^{-} \]

c. Lewis structures for \(\mathrm{ClO}_{2}-, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\)

Follow the same procedure as mentioned earlier for each molecule or ion: \(\mathrm{ClO}_{2}^{-}\): \[ \chemfig{Cl(=[:0]O)(-[:-90]O)}^{-} \] \(\mathrm{SCl}_{2}\): \[ \chemfig{S(-[:0]Cl)(-[:-90]Cl)} \] \(\mathrm{PCl}_{2}^{-}\): \[ \chemfig{P(-[:0]Cl)(-[:-90]Cl)}^{-} \]

d. Conclusions

From the Lewis structures of species containing the same number of atoms and valence electrons, we can draw the following conclusions: 1. The shape of the molecules or ions depends on the arrangement of atoms and electron pairs around the central atom. 2. Species with the same number of atoms and valence electrons might not necessarily have the same shape. 3. Species with different central atoms, even though they might have the same number of atoms and valence electrons, will have different arrangements of electrons due to differences in electronegativity and the central atom's ability to expand its valence shell.

Key concepts

Octet Rule

In chemistry, the octet rule is an essential concept that helps explain how atoms bond and form stable molecules. The rule states that atoms tend to achieve eight electrons in their valence shell, reaching a noble gas-like electron configuration. This stability arises because a filled outer shell, typically with eight electrons, is energetically favorable for most elements.

The octet rule is especially applicable to elements in the second period of the periodic table, such as carbon, nitrogen, and oxygen. For instance, in the molecule \( ext{POCl}_3\), phosphorus serves as the central atom and forms bonds with surrounding atoms to satisfy its octet.

Some atoms, like sulfur, can exceed the octet rule due to the availability of d-orbitals. Thus, they can hold more than eight electrons when forming compounds beyond the second period, like in \( ext{SO}_4^{2-}\). However, the octet rule provides an excellent starting point for drawing Lewis structures and understanding chemical bonding in molecules.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining how atoms interact to form chemical bonds. These electrons are found in the last electron shell (highest energy level) of an atom and participate directly in bonding. Understanding the number of valence electrons helps to predict how atoms will bond and organize when forming molecules.

Atoms tend to form bonds to achieve a full valence shell, often aligning with the octet rule. In crafting Lewis structures, counting valence electrons for each atom in a molecule is a fundamental step. For example, in \( ext{NF}_3\), nitrogen has five valence electrons, while each fluorine atom has seven.

When drawing Lewis structures, these valence electrons are represented as dots around the chemical symbols, indicating possibilities for forming single, double, or even triple bonds. Knowing the number of valence electrons guides us in creating accurate models of molecular structures and predicting molecule behavior.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms within a molecule. It is a key factor that affects the physical and chemical properties of substances. Predicting geometry requires understanding the Valence Shell Electron Pair Repulsion (VSEPR) theory, which states that electron pairs around a central atom will arrange themselves to minimize repulsion, resulting in various geometrical shapes.

For instance, \( ext{XeO}_4\) is an example where the central xenon atom is surrounded by four oxygen atoms, forming a tetrahedral shape. Tetrahedral geometry arises when there are four electron pairs arranged symmetrically around the central atom.

Molecular geometry significantly influences how molecules interact. The shape determines molecule polarity, bond angles, and reactivity. Recognizing molecular shapes helps predict how molecules might assemble, react, or interact with others in chemical reactions.

Electronegativity

Electronegativity is a chemical property that describes an atom's ability to attract and bind electrons in a molecule. It is a crucial factor in determining bond polarity and molecular behavior. You can think of it as a measure of how "greedy" an atom is for electrons. The more electronegative an atom, the stronger its pull on the electron cloud shared in a covalent bond.

This concept is essential when assessing the type and strength of bonds in a molecule. For example, in the molecule \( ext{PCl}_2^-\), chlorine is more electronegative than phosphorus, creating a polar bond where electron density is shifted towards chlorine.

Differences in electronegativity between bonded atoms define whether a bond is polar or nonpolar. These attributes dictate how molecules interact in larger systems such as biological processes and materials. Understanding electronegativity helps predict the behavior of molecules, an invaluable tool for chemists and researchers studying reaction mechanisms and material properties.