Molecular Geometry Practice Quiz

A molecule with two electron pairs arranges them on opposite sides to minimize repulsion, resulting in a 180° bond angle. Thus, the molecule is linear.

Three bonding pairs evenly spaced around the central atom minimize repulsion by forming 120° angles. This arrangement gives a trigonal planar shape.

When a molecule has four bonding pairs with no lone pairs, the electron domains arrange themselves in a tetrahedral geometry. This leads to bond angles close to 109.5°.

A molecule with five regions of electron density arranges them in a trigonal bipyramidal geometry to maximize separation. All bonding pairs contribute to this idealized arrangement.

When there are six regions of electron density, they arrange themselves in an octahedral geometry. This configuration minimizes repulsions by having bonds at 90° and 180°.

Although the electron geometry is tetrahedral, the presence of one lone pair reduces the symmetry. This results in a trigonal pyramidal molecular geometry with bond angles slightly less than 109.5°.

The ideal electron geometry for five regions is trigonal bipyramidal, but the presence of one lone pair distorts the shape. This distortion leads to a see-saw molecular geometry.

With three bonding pairs and two lone pairs in a trigonal bipyramidal arrangement, the two lone pairs typically occupy equatorial positions. This results in a T-shaped molecular geometry.

An octahedral geometry arranges six bonds with each adjacent bond at a 90° angle. This uniform angle distribution minimizes electron repulsions across the molecule.

A tetrahedral molecule with identical bonds has a symmetrical arrangement that causes the dipole moments to cancel. This results in a nonpolar molecule despite the polar nature of individual bonds.

Tetrahedral molecules ideally have bond angles of approximately 109.5°. This arrangement minimizes electron pair repulsions and is a fundamental concept in VSEPR theory.

A trigonal planar geometry is produced by sp2 hybridization, which forms three hybrid orbitals arranged at 120° to one another. This hybridization is essential for achieving the planar structure.

When two of the four electron pairs around a central atom are lone pairs, the molecular geometry deviates from the ideal tetrahedral, resulting in a bent shape. The lone pairs exert greater repulsion, compressing the bond angle.

Lone pairs occupy more space than bonding pairs because of their higher electron density. This increased repulsion among lone pairs reduces the bond angles between adjacent bonding pairs.

SF4 has four bonding pairs and one lone pair around the central sulfur atom. This combination leads to a distorted trigonal bipyramidal electron geometry, resulting in a see-saw molecular shape.

ClF3 has five domains of electron density with three bonding pairs and two lone pairs. The lone pairs occupy equatorial positions, resulting in a T-shaped molecular geometry.

A trigonal bipyramidal geometry arises from sp3d hybridization, where one s orbital, three p orbitals, and one d orbital mix to form five equivalent orbitals. This hybridization supports the five regions of electron density required for the structure.

BF3 has three bonding pairs and no lone pairs, which forces the central boron atom into sp2 hybridization. This results in a trigonal planar shape with bond angles of approximately 120°.

XeF4 exhibits a square planar geometry because it has four bonding pairs and two lone pairs arranged around the central xenon. The lone pairs are positioned to minimize repulsions, resulting in a flat, square structure.

The spatial arrangement of electron domains determines whether the bond dipoles cancel out. In symmetrical molecules, the dipole moments balance, producing a nonpolar molecule, whereas an asymmetrical arrangement yields a net dipole moment and polarity.