Question

Draw the Lewis structures for each of the following molecules or ions. Identify instances where the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state how many electrons surround these atoms: \((\mathbf{a}) \mathrm{PF}_{6}^{-},(\mathbf{b}) \mathrm{BeCl}_{2},(\mathbf{c}) \mathrm{NH}_{3},(\mathbf{d}) \mathrm{XeF}_{2} \mathrm{O}\) (the Xe is the central atom), (e) \(\mathrm{SO}_{4}^{2-}\).

Short answer

In summary, the Lewis structures for the given molecules or ions are as follows: a) PF6-: All atoms obey the octet rule. b) BeCl2: The Be atom does not obey the octet rule and has only 4 electrons around it. c) NH3: All atoms obey the octet rule. d) XeF2O: The Xe atom does not fully obey the octet rule and has 12 electrons around it. e) SO4^2-: All atoms obey the octet rule.

Step-by-step solution

Drawing Lewis Structure for PF6-

PF6- has a total of 5 valence electrons on the central P atom, 6 valence electrons on each of the 6 F atoms, and an extra electron due to its -1 charge. The total number of valence electrons is thus 5 + 6*6 + 1 = 42. Place the P atom in the center, surrounded by the 6 F atoms. Then, put a single bond between each P-F pair with 2 shared electrons. Lastly, complete the octet rule by assigning the remaining 24 electrons (6 unshared pairs) to each F atom. The Lewis structure for PF6- is complete. The octet rule is obeyed in PF6-.

Drawing Lewis Structure for BeCl2

BeCl2 has a total of 2 valence electrons on the central Be atom and 7 valence electrons on each of the two Cl atoms. The total number of valence electrons is thus 2 + 7*2 = 16. Place the Be atom in the center, surrounded by the 2 Cl atoms. Put a single bond between Be-Cl pairs to share 2 electrons and assign the remaining 12 electrons (6 unshared pairs) to each Cl atom. The Lewis structure for BeCl2 is complete. The octet rule is not obeyed in BeCl2 since Be has only 4 electrons around it.

Drawing Lewis Structure for NH3

NH3 has a total of 5 valence electrons on the central N atom and 1 valence electron on each of the three H atoms. The total number of valence electrons is thus 5 + 1*3 = 8. Place the N atom in the center, surrounded by the 3 H atoms. Put a single bond between each N-H pair and share 2 electrons. Assign the remaining 2 electrons (1 unshared pair) to the N atom. The Lewis structure for NH3 is complete. The octet rule is obeyed in NH3.

Drawing Lewis Structure for XeF2O

XeF2O has a total of 8 valence electrons on the central Xe atom, 7 valence electrons on each of the two F atoms, and 6 valence electrons on the O atom. The total number of valence electrons is thus 8 + 7*2 + 6 = 28. Place the Xe atom in the center, surrounded by the 2 F atoms and the O atom. Put single bonds between each Xe-F and Xe-O pair to share 6 electrons. Complete the octet rule for each Xe-F and Xe-O atom. Also add 3 lone pairs (6 unshared electrons) to the Xe atom. The Lewis structure for XeF2O is complete. The octet rule is not fully obeyed in XeF2O because the Xe atom has 12 electrons around it.

Drawing Lewis Structure for SO4^2-

In the SO4^2- ion, there are 6 valence electrons on the central S atom, 6 valence electrons on each of the four O atoms, and two extra electrons due to the -2 charge. The total number of valence electrons is 6 + 6*4 + 2 = 32. Place the S atom in the center, surrounded by the four O atoms. Put a double bond between the S-O1 and S-O2 pair, and single bonds between the S-O3 and S-O4 pair for a total of 10 shared electrons. Assign the remaining 16 electrons (2 unpaired pairs) around each O atom except those with the double bond. The Lewis structure for SO4^2- is complete. The octet rule is obeyed in SO4^2-.

Key concepts

Octet Rule

The octet rule is a fundamental principle in chemistry that refers to the tendency of atoms to prefer having eight electrons in their valence shell. This rule is based on the idea that a full valence shell resembles the electron configuration of noble gases, which are stable and non-reactive. To achieve a full octet, atoms can form chemical bonds by sharing, gaining, or losing electrons.

In the step-by-step exercise you've seen, certain compounds like \(\mathrm{PF}_{6}^{-}\) and \(\mathrm{NH}_{3}\) adhere strictly to the octet rule, meaning all atoms involved have eight electrons in their valence shells after bonding. However, there are exceptions, such as \(\mathrm{BeCl}_{2}\) and \(\mathrm{XeF}_{2}\mathrm{O}\), where the octet rule is not fully obeyed.

• Atoms like Beryllium (Be) and Boron (B) are content with fewer than eight electrons.
• Atoms in Period 3 or higher can have expanded octets, such as Xenon (Xe) in \(\mathrm{XeF}_{2}\mathrm{O}\), due to available d orbitals.

Understanding the octet rule and its exceptions is crucial for correctly drawing Lewis structures and predicting molecular behavior.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in chemical bonding and the formation of molecules. These electrons determine the bonding behavior of an atom because they are involved in the formation of chemical bonds. When drawing Lewis structures, counting valence electrons helps in determining how atoms share or give away electrons.

Each element has a specific number of valence electrons that can be found using the periodic table:

• Elements in Group 1 have 1 valence electron.
• Elements in Group 2 have 2 valence electrons.
• Elements in Group 17 have 7 valence electrons, and so on.

In the examples you worked with, knowing that Phosphorus (P) has 5 valence electrons, each Chlorine (Cl) has 7, and Nitrogen (N) has 5 allowed you to correctly arrange these electrons in the form of bonds or lone pairs around the central atom. Valence electrons help us understand not just the structure, but also the potential reactivity and characteristics of molecules.

Chemical Bonding

Chemical bonding refers to the forces that hold atoms together in molecules and compounds. The main types of chemical bonds are ionic, covalent, and metallic, but in Lewis structures, we focus primarily on covalent bonds, where electrons are shared between atoms.

The nature of chemical bonding largely dictates the structure and stability of the molecules. For instance, in a covalent bond, involved in the molecules you analyzed, two atoms share one or more pairs of electrons:

• In \(\mathrm{PF}_{6}^{-}\), each Fluorine (F) forms a single covalent bond with Phosphorus (P).
• In \(\mathrm{NH}_{3}\), each Hydrogen (H) shares an electron with Nitrogen (N).
• Double bonds, as found in \(\mathrm{SO}_{4}^{2-}\), represent two shared pairs of electrons.

Understanding chemical bonding allows you to predict molecular shapes, understand reactivity, and explain the properties of compounds. Different types of bonds affect the physical and chemical properties like melting points, boiling points, and solubility.