Question

(a) Methane \(\left(\mathrm{CH}_{4}\right)\) and the perchlorate ion \(\left(\mathrm{ClO}_{4}-\right)\) are both described as tetrahedral. What does this indicate about their bond angles? (b) The \(\mathrm{NH}_{3}\) molecule is trigonal pyramidal, while \(\mathrm{BF}_{3}\) is trigonal planar. Which of these molecules is flat?

Short answer

(a) Both methane (CH4) and the perchlorate ion (ClO4-) have tetrahedral geometries, resulting in bond angles of \(109.5^\circ\). (b) The NH3 molecule has a trigonal pyramidal geometry and is not flat, while the BF3 molecule has a trigonal planar geometry and is flat.

Step-by-step solution

Part (a): Identifying bond angles of tetrahedral molecules

Methane (CH4) and the perchlorate ion (ClO4-) share a tetrahedral geometry. In a tetrahedral structure, there are four regions of electron density around the central atom, pushing each other as far apart as possible to minimize electron repulsion. This arrangement results in bond angles of \(109.5^\circ\). Therefore, both CH4 and ClO4- have bond angles of \(109.5^\circ\).

Part (b): Determining the flat molecule

To determine which molecule, NH3 or BF3, is flat, we need to analyze their molecular geometries. Ammonia (NH3) has a trigonal pyramidal structure. In this geometry, there are three bonding electron-pair regions and one lone pair of electrons around the central atom, resulting in a three-dimensional shape. This means that the NH3 molecule is not flat. On the other hand, boron trifluoride (BF3) has a trigonal planar structure. In this geometry, there are three bonding electron-pair regions and no lone pairs around the central atom, leading to a two-dimensional shape with bond angles of \(120^\circ\). Therefore, the BF3 molecule is flat.

Key concepts

Tetrahedral Bond Angles

Molecules with a tetrahedral geometry, like methane \( \mathrm{CH}_4 \) and the perchlorate ion \( \mathrm{ClO}_4^- \), have distinct bond angles due to the spatial arrangement of their atoms. This structure results when a central atom is surrounded by four bonded atoms or groups of electrons that orient themselves as far away from each other as possible to reduce electron-pair repulsion.
This arrangement creates bond angles of \( 109.5^\circ \). You can visualize this by imagining a three-dimensional pyramid where the central atom is at the center and the other atoms are at the corners. Such an arrangement gives the molecule a three-dimensional, non-linear geometry.

• The electron pairs spread evenly in three dimensions.
• All bond angles are equal at \( 109.5^\circ \).
• This geometry ensures maximum distance between electron pairs, minimizing repulsion.

Understanding tetrahedral bond angles helps explain why molecules like \( \mathrm{CH}_4 \) are not planar but rather exhibit a robust three-dimensional structure.

Trigonal Pyramidal

The trigonal pyramidal geometry is characterized by a central atom surrounded by three bonded atoms and one lone pair of electrons. An example of this structure is ammonia \( \mathrm{NH}_3 \). While it may seem similar to the tetrahedral arrangement, the lone pair of electrons in trigonal pyramidal molecules affects their shape and bond angles.
The lone pair exerts more repulsion on the bonded pairs compared to the repulsion between bonded atoms. This causes the bond angles to decrease slightly from the ideal tetrahedral angle of \( 109.5^\circ \) to about \( 107^\circ \).

• Presence of a lone pair affects bond angles, making it less than the tetrahedral angle.
• Molecules have a three-dimensional, non-flat shape.
• The arrangement leads to a structure resembling a pyramid with a triangular base.

In practical terms, this means that although ammonia is not flat, the distinction in its structure is predominantly due to the lone pair's interaction.

Trigonal Planar

Trigonal planar molecular geometry is seen in molecules like boron trifluoride \( \mathrm{BF}_3 \). This arrangement includes a central atom bonded to three surrounding atoms without any lone pairs on the central atom. As a result, these molecules are flat and two-dimensional.
The bond angles in a trigonal planar arrangement are \( 120^\circ \), as the atoms try to be as far apart as possible in a plane. This geometry contrasts with trigonal pyramidal geometry by being planar, demonstrating how the absence of lone pairs leads to a flat molecule.

• No lone pairs allows for equal bond distribution in a flat plane.
• Bond angles are uniform at \( 120^\circ \).
• Molecules are two-dimensional, making them inherently flat.

Understanding trigonal planar structures provides insight into why \( \mathrm{BF}_3 \) is not only flat but also exhibits consistent bond angles across its structure.