Question

The Group \(5 \mathrm{~A}\) elements can form molecules or ions that involve three, five, or six covalent bonds; \(\mathrm{NH}_{3}, \mathrm{AsCl}_{5}\), and \(\mathrm{PF}_{6}^{-}\) are examples. Draw the Lewis structure for each of these substances, and predict the molecular structure and hybridization for each. Why doesn't \(\mathrm{NF}_{5}\) or \(\mathrm{NCl}_{6}^{-}\) form?

Short answer

The Lewis structures, molecular structures, and hybridizations for the given species are: - NH₃: Lewis structure shows 3 N-H bonds and 1 lone pair on N. Molecular shape is trigonal pyramidal, and hybridization is sp³. - AsCl₅: Lewis structure shows 5 As-Cl bonds. Molecular shape is trigonal bipyramidal, and hybridization is sp³d. - PF₆⁻: Lewis structure shows 6 P-F bonds. Molecular shape is octahedral, and hybridization is sp³d². NF₅ and NCl₆⁻ don't form because Nitrogen lacks enough available orbitals to form the required number of bonds with Fluorine and Chlorine.

Step-by-step solution

Drawing Lewis structures

To draw the Lewis structures of these three species, we need to follow some basic steps: - Count the total available valence electrons. - Assign pairs of electrons to the central atom to fulfill the octet rule. - Arrange the remaining electrons around the outer atoms to also satisfy the octet rule.

NH₃ Lewis structure

1. The total valence electrons are \[5 (N) + 3 \times 1 (H) = 8\text{ electrons}\]. 2. Nitrogen is the central atom, surrounded by three Hydrogen atoms. 3. Distribute the remaining valence electrons: Nitrogen shares three pairs of electrons with three Hydrogen atoms, forming three single bonds (N-H). Nitrogen also has one lone pair of electrons left. The structure is: ``` H | H--N--H ```

AsCl₅ Lewis structure

1. The total valence electrons are \[5 (As) + 5 \times 7 (Cl) = 40\text{ electrons}\]. 2. Arsenic (As) is the central atom, surrounded by five Chlorine (Cl) atoms. 3. Distribute the remaining valence electrons: Each Chlorine atom shares a pair of electrons with Arsenic, forming five single bonds (As-Cl). Each Chlorine atom also has three lone pairs of electrons. The structure is: ``` Cl | Cl--As--Cl / \ Cl Cl ```

PF₆⁻ Lewis structure

1. The total valence electrons are \[5 (P) + 6 \times 7 (F) + 1 (\text{extra e⁻}) = 48\text{ electrons}\]. 2. Phosphorus (P) is the central atom, surrounded by six Fluorine (F) atoms. 3. Distribute the remaining valence electrons: Each Fluorine shares one pair of electrons with Phosphorus, forming six single bonds (P-F). Each Fluorine atom also has three lone pairs of electrons. The structure is: ``` F | F--P--F / \ F F \ F ```

Predicting molecular structures and hybridization

Now that we have the Lewis structures, we can predict the molecular geometries using the VSEPR theory, which states that electron pairs repel each other to get their geometry. - NH₃: Nitrogen is surrounded by three bonding pairs and one lone pair, which gives it an electron geometry of tetrahedral. However, considering only bond pairs, the molecular shape is trigonal pyramidal. The hybridization would be sp³ since there are four electron pairs. - AsCl₅: Arsenic is surrounded by five bonding pairs. Thus, the electron geometry is trigonal bipyramidal. The molecular shape is also trigonal bipyramidal, and the hybridization would be sp³d. - PF₆⁻: Phosphorus is surrounded by six bonding pairs, giving it an electron geometry of octahedral. The molecular shape is octahedral as well, and the hybridization would be sp³d².

Explaining why NF₅ and NCl₆⁻ don't form

- NF₅: Nitrogen is a very small atom, with a valence electron shell (n=2) that can accommodate only four orbitals: one s and three p orbitals. In order to form five bonds with Fluorine (F), Nitrogen would need five orbitals (sp³d), which is not possible for Nitrogen since it lacks available d orbitals. Furthermore, F is highly electronegative, which would lead to increased electron repulsion in such a compact structure. - NCl₆⁻: Similar to NF₅, Nitrogen's valence shell doesn't have enough orbitals to form six bonds with Chlorine (Cl) atoms, as it would require six orbitals (sp³d²) and Nitrogen lacks d orbitals.

Key concepts

Covalent Bonds

Covalent bonds are a type of chemical bond where atoms share pairs of electrons. This sharing allows each atom to achieve a stable electronic configuration, similar to that of the noble gases. In molecules like

• NH₃ (Ammonia): Three hydrogen atoms share their electrons with nitrogen, forming three single covalent bonds.
• AsCl₅ (Arsenic Pentachloride): Arsenic shares its electrons with five chlorine atoms, forming five single covalent bonds.
• PF₆⁻: Phosphorus forms six covalent bonds with fluorine atoms.

Each bond is crucial as it helps stabilize the molecule by filling the outer electron shells of the involved atoms. Covalent bonds can be single, double, or triple, though only single bonds appear in the given examples. Understanding these bonds is fundamental to predicting molecular geometry and chemical reactivity.

Hybridization

Hybridization is a concept that explains the mixing of atomic orbitals in an atom to form new hybrid orbitals. This is essential for understanding molecular geometry. In the given examples, hybridization helps explain the shape and bond angles in:

• NH₃: Nitrogen undergoes sp³ hybridization, combining one s orbital and three p orbitals to form four equivalent sp³ orbitals. This allows for the trigonal pyramidal shape with bond angles slightly less than 109.5° due to the lone pair of electrons.
• AsCl₅: The arsenic atom uses sp³d hybridization, which involves the s, p, and d orbitals to create five hybrid orbitals. This results in a trigonal bipyramidal shape, accommodating five bonds around the central atom.
• PF₆⁻: Phosphorus uses sp³d² hybridization to form six hybrid orbitals, arranging themselves to form an octahedral shape. This allows equal spacing and minimizes repulsion.

Hybridization is a useful tool to predict and explain molecular structures and their chemical behavior.

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the geometry of molecules based on electron pair repulsion. According to this theory, electron pairs around a central atom will arrange themselves to minimize repulsion. Here's how it applies:

• NH₃: With four valence electron groups (three bonding pairs, one lone pair), the electron geometry is tetrahedral, but the molecular geometry is trigonal pyramidal due to the lone pair.
• AsCl₅: Arsenic has five bonding pairs and no lone pairs, resulting in a trigonal bipyramidal shape. This arrangement allows for minimized repulsion among the bonds.
• PF₆⁻: With six bonding pairs, the VSEPR theory predicts an octahedral shape, where all positions are occupied equally, leading to minimized electron pair repulsion.

By using VSEPR theory, we can predict the ideal bond angles and molecular shapes, helping to understand the physical properties and potential reactions of these molecules.