Question

Examine the following structures and predict the most likely geometry, using concepts of orbital hybridization. State whether the molecule should be planar or nonplanar, and list the approximate values expected for the bond angles. a. \(\mathrm{SiCl}_{4}\) b. \(\mathrm{HCOO}^{\ominus}\) c. \(\mathrm{CH}_{3}-\mathrm{C} \equiv \mathrm{CH}\) d. \(\mathrm{F}_{2} \mathrm{C}=\mathrm{C}=\mathrm{CF}_{2}\) e. \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{O}^{\oplus}\)

Short answer

a. Tetrahedral, nonplanar, \( 109.5^\circ \). b. Trigonal planar, planar, \( 120^\circ \). c. Linear/tetrahedral, nonplanar, \( 180^\circ \) or \( 109.5^\circ \). d. Planar, \( 120^\circ \) or \( 180^\circ \). e. Trigonal pyramidal, nonplanar, slightly less than \( 109.5^\circ \).

Step-by-step solution

Analyze SiCl4

The molecule \( \mathrm{SiCl}_4 \) consists of silicon (Si) surrounded by four chlorine (Cl) atoms. Silicon typically undergoes \( sp^3 \) hybridization when forming tetrahedral compounds. ### Geometry: Tetrahedral### Planarity: Nonplanar### Bond Angles: Approximately \( 109.5^\circ \)

Analyze HCOO egative

In the \( \mathrm{HCOO}^{\ominus} \) ion, the central carbon atom is bonded to two oxygens and one hydrogen, with resonant structures allowing delocalization of the negative charge. The carbon is \( sp^2 \) hybridized.### Geometry: Trigonal Planar### Planarity: Planar### Bond Angles: Approximately \( 120^\circ \)

Analyze CH3-C≡CH

The molecule \( \mathrm{CH_3-C \equiv CH} \) contains a central carbon-carbon triple bond, where the carbon atoms on the triple bond typically undergo \( sp \) hybridization. The terminal carbon in \( \mathrm{CH_3} \) is \( sp^3 \) hybridized, and the central carbon is \( sp \) hybridized.### Geometry: Linear (around \( C \equiv C \) bond), Tetrahedral (around \( \mathrm{CH_3} \))### Planarity: Nonplanar (due to different geometries)### Bond Angles: \( 180^\circ \) around \( C \equiv C \), \( 109.5^\circ \) in \( \mathrm{CH_3} \)

Analyze F2C=C=CF2

In \( \mathrm{F_2C=C=C=CF_2} \), the carbon atoms involved in double bonds are \( sp^2 \) hybridized, with each bond plane aligned, creating a planar structure.### Geometry: Linear for \( C=C=C \) segments### Planarity: Planar### Bond Angles: \( 120^\circ \) in \( \mathrm{F_2C} \) and linear segments are \( 180^\circ \)

Analyze (CH3)3O+

In \( \left( \mathrm{CH}_3 \right)_{3} \mathrm{O}^{\oplus} \), the oxygen atom is bonded to three methyl groups. As a positively charged species, it adopts \( sp^3 \) hybridization with a slightly distorted tetrahedral geometry.### Geometry: Trigonal Pyramidal### Planarity: Nonplanar### Bond Angles: Slightly less than \( 109.5^\circ \) (due to positive charge and lone pair repulsion)

Key concepts

Orbital Hybridization

Orbital hybridization is a key concept in chemistry that explains the mixing of atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds. It helps predict the geometry of molecules based on the type and arrangement of orbitals that electrons occupy during bond formation.
\( \mathrm{SiCl}_4 \) is a classic example where the silicon atom undergoes \( sp^3 \) hybridization. This means one \( s \) and three \( p \) orbitals combine to form four equivalent \( sp^3 \) hybrid orbitals. As a result, this molecule adopts a tetrahedral shape.
In the case of \( \mathrm{HCOO}^{\ominus} \), the central carbon is \( sp^2 \) hybridized, which involves mixing one \( s \) and two \( p \) orbitals to form three equivalent \( sp^2 \) hybrid orbitals. This results in a trigonal planar geometry.

• \( sp^3 \) hybridization leads to a tetrahedral geometry.
• \( sp^2 \) hybridization results in a planar geometry.
• \( sp \) hybridization is associated with linear geometry.

Tetrahedral Geometry

Tetrahedral geometry occurs when a central atom is bonded to four other atoms, and all the bond angles are approximately \( 109.5^\circ \). This three-dimensional shape is symmetrical, distributing atoms as far apart as possible from each other, minimizing repulsion.
For instance, \( \mathrm{SiCl}_4 \) exhibits a perfect tetrahedral geometry because of the \( sp^3 \) hybridization of silicon. The four chlorine atoms are positioned at the corners of a tetrahedron, creating equal angles between bonds.
Another example is the methyl group in \( \mathrm{CH_3-C \equiv CH} \), where the terminal carbon in the \( \mathrm{CH_3} \) group is \( sp^3 \) hybridized, forming tetrahedral geometry around it. This creates bond angles near \( 109.5^\circ \).

Planar Molecules

Planar molecules are characterized by all of their atoms lying in the same plane. A molecule is often planar when the central atom is \( sp^2 \) hybridized, forming a flat, triangular shape called trigonal planar.
For the \( \mathrm{HCOO}^{\ominus} \) ion, the central carbon's \( sp^2 \) hybridization results in three \( sp^2 \) orbitals in a plane, with bond angles of approximately \( 120^\circ \). The delocalization of negative charge across the oxygen atoms contributes to its planar shape.
In the compound \( \mathrm{F_2C=C=C=CF_2} \), each carbon atom is \( sp^2 \) hybridized, resulting in a configuration that allows for planar alignment through resonation of \( \pi \) bonds. This makes the entire molecule flat.

• \( \mathrm{HCOO}^{\ominus} \) and \( \mathrm{F_2C=C=C=CF_2} \) are examples of planar molecules.
• Planarity in molecules is often due to \( sp^2 \) hybridized central atoms.

Bond Angles

Bond angles refer to the angles formed between adjacent bonds at an atom in a molecule, offering insight into its geometry.
In a tetrahedral geometry like \( \mathrm{SiCl}_4 \), the bond angles are about \( 109.5^\circ \). This geometry arises naturally to minimize electron pair repulsions between the \( sp^3 \) orbitals.
In trigonal planar molecules like \( \mathrm{HCOO}^{\ominus} \), the bond angles are approximately \( 120^\circ \). The \( sp^2 \) hybridized carbon forms three bonds in one plane, evenly spaced due to electron repulsion.
Linear geometries, seen in molecules with \( sp \) hybridization like segments of \( \mathrm{F_2C=C=C=CF_2} \), have bond angles of \( 180^\circ \).

• \( 109.5^\circ \) is typical for tetrahedral molecules.
• \( 120^\circ \) is typical for trigonal planar structures.
• \( 180^\circ \) is typical for linear molecules.

Understanding these bond angles helps predict molecular geometry and chemical behavior.