Question

Draw the Lewis structures of \(\mathrm{NOF}_{3}\) and \(\mathrm{POF}_{3}\) in which the group 15 element is the central atom and the other atoms are bonded to it. What differences are there in the types of bonding in the molecules?

Short answer

Answer: The main difference in bonding types between the Lewis structures of NOF3 and POF3 is that in NOF3, the nitrogen-oxygen bond consists of a double bond, while in POF3, the phosphorus-oxygen bond consists of a single bond.

Step-by-step solution

Determine the number of valence electrons

For \(\mathrm{NOF}_{3}\): - Nitrogen (N) has 5 valence electrons. - Oxygen (O) has 6 valence electrons. - Fluorine (F) has 7 valence electrons and there are 3 F atoms. Total valence electrons in \(\mathrm{NOF}_{3}\) = 5 + 6 + (3 × 7) = 5 + 6 + 21 = 32 For \(\mathrm{POF}_{3}\): - Phosphorus (P) has 5 valence electrons. - Oxygen (O) has 6 valence electrons. - Fluorine (F) has 7 valence electrons and there are 3 F atoms. Total valence electrons in \(\mathrm{POF}_{3}\) = 5 + 6 + (3 × 7) = 5 + 6 + 21 = 32

Place the central atom

For both molecules, the group 15 element (N for \(\mathrm{NOF}_{3}\) and P for \(\mathrm{POF}_{3}\)) is specified as the central atom.

Add octets to the outer atoms

Add single bonds between the central atoms and the other atoms (1 O and 3 F atoms in both molecules) to fill their octets. \(\mathrm{NOF}_{3}\): N - O - F | | F F \(\mathrm{POF}_{3}\): P - O - F | | F F

Add remaining electrons to the central atom

For both molecules, the central atom has consecutive single bonds to the other atoms and therefore, has 4 single bonds in total. Since N and P have 5 valence electrons each, we need to add one more electron to each central atom to satisfy their octets. As a result, nitrogen in \(\mathrm{NOF}_{3}\) will carry one lone pair, and phosphorus in \(\mathrm{POF}_{3}\) will carry one lone pair.

Adjust the bonds

Since the octet rule is already satisfied for all atoms in \(\mathrm{NOF}_{3}\) and \(\mathrm{POF}_{3}\), there is no need to adjust the bonds further. The final Lewis structures for the molecules are: \(\mathrm{NOF}_{3}\): O || N - O++ - F ^ || ^ | F | ------------- \(\mathrm{POF}_{3}\): P - O - F ^ | | F | F As for the difference in bonding types: In \(\mathrm{NOF}_{3}\), the nitrogen-oxygen bond consists of a double bond, while in \(\mathrm{POF}_{3}\), the phosphorus-oxygen bond consists of a single bond.

Key concepts

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. They play a crucial role in chemical bonding because they determine how an atom interacts with other atoms. In Lewis structures, these electrons are typically represented as dots around atomic symbols, visually demonstrating how atoms share or transfer electrons to form bonds.
To find the number of valence electrons for an atom, you look at its group number in the periodic table. For example:

• Nitrogen (N) belongs to group 15, so it has 5 valence electrons.
• Oxygen (O) is in group 16, giving it 6 valence electrons.
• Fluorine (F), found in group 17, has 7 valence electrons.

This understanding is essential because the total number of valence electrons in a molecule determines how the atoms within that molecule bond together. To properly draw a Lewis structure, you must account for all these electrons. For example, both NOF₃ and POF₃ have a total of 32 valence electrons that need to be arranged to obey bonding rules.

Bonding Types

Different types of bonds are formed based on how atoms share or transfer their valence electrons. Generally, there are two main types of bonding found in Lewis structures: covalent and ionic bonding. In covalent bonds, atoms share electrons to attain full outer shells, whereas in ionic bonds, electrons are transferred from one atom to another.
In the molecules NOF₃ and POF₃, covalent bonding dominates.

• In NOF₃, nitrogen forms a double bond with oxygen, sharing two pairs of electrons, and single bonds with each fluorine atom.
• This difference marks another type of interaction with oxygen in POF₃, where phosphorus forms a single bond with oxygen.

The sharing of electrons between atoms helps achieve stable electron arrangements, usually aligning with the octet rule. Understanding these bonding types is crucial in predicting the behavior and reactivity of molecules.

Octet Rule

The octet rule is a chemical rule of thumb that suggests atoms tend to form bonds until they are surrounded by eight valence electrons, much like a noble gas. This rule is key to understanding how atoms achieve stability through bonding. When drawing Lewis structures, the octet rule guides how electrons are arranged around atoms.
For the molecules NOF₃ and POF₃:

• Nitrogen and phosphorus are central atoms that initially have fewer than eight electrons, encouraging them to form multiple bonds to fulfill the octet requirement.
• Oxygen and fluorine, already electronegative elements, typically form one or two bonds to complete their octet.

Each atom in the structures strives to achieve the closest version of the octet configuration, whether by sharing electrons in a covalent bond or by completing outer shell pairs through single, double, or even triple bonds. While the rule helps predict molecular structure, there are exceptions, especially with atoms having less than or more than eight valence electrons. However, for common cases like these, it remains a useful guideline.