Molecular Geometry | General Chemistry 1

Importance of molecular shape:

The shape of a molecule is critical for determining its physical and chemical properties, such as reactivity, polarity, phase of matter, color, magnetism, and biological activity.

Molecular shape from Lewis structures:

Lewis structures show how atoms are connected in a covalent molecule or ion and provide a visual representation of the arrangement of electrons in a molecule. However Lewis structures do not provide information about the orientation of the bonds in space and the molecular geometry.

Limitations of Lewis structures: 

• The Lewis structure of dichloromethane (CH₂Cl₂) shows a carbon atom bonded to 2 hydrogen atoms and 2 chlorine atoms. 2 Lewis structures can be drawn by rotating the positions of hydrogen and chlorine atoms around the carbon atom. 
• However, these 2-D representations do not convey the actual 3-D shape of the molecule and contradict the fact that only one molecule of dichloromethane exists. The Lewis structures suggest possible isomerism, which is not observed experimentally.
• Tetrahedral shape explains why dichloromethane has only one isomer:
   

Steps for determining shape:

Despite their limitations, Lewis structures are the first step in predicting molecular geometries. By identifying the number of bonding pairs and lone pairs of electrons around the central atom, we can apply VSEPR theory to predict the shape of the molecule.

1. Draw the Lewis structure.
2. Count all regions of electron density (bonds and lone pairs) around the central atom.
3. Use the VSEPR theory to predict the arrangement of electron domains and the resulting molecular geometry.

VSEPR model: 

The VSEPR (Valence-Shell Electron-Pair Repulsion) model predicts the shape of molecules based on the repulsion between electron pairs (bonding and lone pairs) in the valence shell of the central atom. These electron pairs arrange themselves as far apart as possible to minimize repulsive interactions.
 

VSEPR notation:

BeH₂: AX₂,   H₂O: AX₂E₂

NH₃: AX₃E₁,   CO₂: AX₂

Electron-domain geometry vs. molecular geometry:

• Electron-domain geometry: The arrangement of all electron domains (both bonding pairs and lone pairs) around the central atom.
• Molecular geometry: The arrangement of only the bonded atoms around the central atom, ignoring lone pairs.

Predicting molecular geometry:

1. Draw the Lewis structure: Identify the total number of valence electrons and arrange them to show bonding pairs and lone pairs.
2. Count the number of electron domains around the central atom.
3. Determine the electron geometry according to the VSEPR theory.
4. Determine the molecular geometry based on the positions of only the bonded atoms.

Common geometries:

2 electron domains:

• Electron-domain geometry: linear (angle = 180°)
• 1 Molecular geometry: linear (AX₂)

3 electron domains:

• Electron-domain geometry: trigonal planar (angle = 120°)
• 2 Molecular geometries: trigonal planar (AX₃), bent (AX₂E)

4 electron domains:

• Electron-domain geometry: tetrahedral (angle = 109.5°)
• 3 Molecular geometry: tetrahedral (AX₄), trigonal pyramidal (AX₃E), bent (AX₂E₂)

5 electron domains:

• Electron-domain geometry: trigonal bipyramidal (angles = 90° and 120°)
• 4 Molecular geometries: trigonal bipyramidal (AX₅), seesaw (AX₄E), T-shaped (AX₃E₂), linear (AX₂E₃)

6 electron domains:

• Electron-domain geometry: octahedral (angle = 90°)
• 3 Molecular geometries: octahedral (AX₆), square pyramidal (AX₅E), square planar (AX₄E₂)

Repulsions between electron domains:

Electron domain repulsions can be categorized based on the type of pairs involved. Lone pairs exert the greatest repelling effect because they are closer to the nucleus of the central atom and the unshared electrons can spread out more:

lp-lp (lone pair-lone pair) > lp-bp (lone pair-bonding pair) > bp-bp (bonding pair-bonding pair)

Position of lone pairs:

Lone pairs will occupy positions that minimize their interaction with other electron domains ⇒ Lone pairs significantly influence the shape and bond angles of a molecule.

Trigonal bipyramidal positions:

• Axial Positions: Have three strong 90° interactions.
• Equatorial Positions: Have two strong 90° interactions.

In a trigonal bipyramidal geometry, lone pairs will prefer equatorial positions to reduce the number of 90° repulsions. This minimizes the overall repulsion within the molecule.

For example, the lone pair of SF₄ occupies an equatorial position, resulting in a see-saw molecular geometry:
 

Lone pairs effect

Lone pairs have more freedom to spread out than bonding pairs because they are not shared between atoms. This increased electron density allows lone pairs to exert a greater repulsive force on other electron domains, causing bond angles to deviate from ideal values ⇒ The angle between a lone pair and a single bond is greater than the angle between 2 single bonds.
 

In CCl₄, all bond angles are 109.5°. However, in NH₃, which also has a tetrahedral electron-domain geometry, the presence of one lone pair reduces the H-N-H bond angles to approximately 107°:
 

Multiple bonds effect

Multiple bonds (double and triple bonds) contain more electron density than single bonds, leading to greater repulsive forces between electron domains. This causes bond angles involving multiple bonds to be larger than those involving only single bonds.

Overall dipole moment:

The dipole moment of a molecule is determined by the vector addition of the individual bond dipoles. A dipole moment is a measure of the separation of positive and negative charges in a molecule.

Conditions for a molecule to be polar:

• The molecule must contain polar bonds, which occur when there is a significant difference in electronegativity between the bonded atoms, resulting in partial positive and negative charges.
• The shape of the molecule must not cancel out the effects of the polar bonds. This means that the molecular geometry must be such that the dipole moments do not cancel each other out through symmetry. It is possible for a molecule to contain polar bonds, but not be polar. Its overall dipole moment is equal to 0.

Determining Polarity:

1. Check bond polarity: Identify if the bonds within the molecule are polar by comparing the electronegativities of the bonded atoms.
2. Assess molecular shape: Use VSEPR theory to determine the molecular geometry.
3. Calculate net dipole Moment: Evaluate the vector sum of all bond dipoles. If the vectors add up to a nonzero value, the molecule is polar. If they cancel out, the molecule is nonpolar.

Carbon Dioxide (CO₂):

• CO₂ is linear ⇒ bond dipoles in CO₂ cancel each other
• overall dipole moment of CO₂ = 0 ⇒ CO₂ is not polar

Water (H₂O):

• H₂O is bent ⇒ bond dipoles do not cancel each other
• overall dipole moment of H₂O ≠ 0 ⇒ water is a polar molecule