VSEPR Theory (College Board AP® Chemistry)

VSEPR Theory

• The valence shell electron pair repulsion (VSEPR) theory is based on the principle that electron groups such as single bonds, double bonds, triple bonds, or single electrons repel one another through Coulombic forces

• The minimize this repulsion, the valence shell electrons should be placed as far apart in the space as possible

• VSEPR theory consist of three rules:

1. Bonding pairs and lone pairs should be arranged as far apart in space as possible

2. Double and triple bonds are one electron group

3. The Coulombic repulsion forces from lone pairs are stronger than the bonding pairs

• The VSEPR theory and the number of electron groups in an atoms is a useful tool that can be used to predict the structural and electron properties of the molecules

Two electron groups

• If there are two electron groups around the central atom, the angle between the bond is 180° which maximize their separation

• The geometry of this molecules is LINEAR

  • g. BeCl₂, CO₂, and ethyne (HC≡CH)

Examples of Linear Molecules

Beryllium chloride, carbon dioxide and ethyne all have two electron groups

Three electron groups

• If there are three electron groups around the central atom, the angle between the bonds is 120° which maximize their separation

• The geometry of this molecule is TRIGONAL PLANAR

  • E.g. BF₃, CH₂CH₂ and CH₂O

Examples of Trigonal Planar Molecules

Boron trifluoride, ethene and methanal all have three electron groups

Four electron groups

• Molecules with four electron groups have three dimensional geometries

• If there are four electron groups around the central atom, there are three case scenarios:

  • If the four electron groups are bonding groups, the angle between the bonds is 109.5° which maximize their separation

    • The geometry of this molecules is TETRAHEDRAL

    • E.g. CH₄, NH₄⁺ 

Examples of Tetrahedral Molecules

Methane and ammonium ions have four electron domains

• If three electron groups are bonding groups and one is a lone pair, the angle is approximately 107° which is slightly less than 109.5° due to the Coulombic force of repulsion generated by the lone pair

  • The geometry of this molecules is TRIGONAL PYRAMIDAL

  • E.g. NH₃ 

NH₃ has a Trigonal Pyramidal Geometry

The molecular geometry of ammonia

• If two electron groups are bonding groups and two are lone pairs, the angle is approximately 104.5° which is less than 109.5° due to the Coulombic force of repulsion generated by the two lone pairs

Bond Angles in Water

The order of electron pair repulsion is lone pairs > lone pair: bonding pair > bonding pairs

• The geometry of this molecules is BENT or ANGULAR

• E.g. H₂O

Water Has a Bent Geometry

The molecular geometry of water

Five electron groups

• Molecules with five electron groups have three dimensional geometries

• If there are five electron groups around the central atom, there are four case scenarios:

  • If the five electron groups are bonding groups, the angle between the equatorial bonds is 120° and the angle between the axial bonds is 90° which maximize their separation

    • The geometry of this molecules is TRIGONAL BIPYRAMIDAL

    • E.g. PCl₅ 

PCl₅ Has a Trigonal Bipyramidal Geometry

The molecular geometry of phosphorus pentachloride

• If four electron groups are bonding groups and one of them is a lone pair, the angle between the equatorial bonds is slightly less than 120° and the angle between the axial bonds is slightly less than 90° due to the Coulombic force of repulsion generated by the lone pair

  • The geometry of this molecules is SEESAW

  • E.g. SF₄

SF₄ Has a Seesaw Geometry

The molecular geometry sulfur tetrafluoride

• If three electron groups are bonding groups and two of them are lone pairs, it is a flat molecule

• The angle between bonds is slightly less than 90° due to the Coulombic force of repulsion generated by the lone pairs

  • The geometry of this molecules is T-SHAPED

  • E.g. ClF₃

ClF₃ Has a T-shaped Geometry

The molecular geometry of chlorine trifluoride

• If two electron groups are bonding groups and three of them are lone pairs, it is a flat molecule

• The angle between bonds is slightly 180° due to the Coulombic force of repulsion generated by the lone pairs

  • The geometry of this molecules is LINEAR

  • E.g. I₃^-

I₃^- Has a Linear Geometry

The molecular geometry of the triiodide ion

Six electron groups

• Molecules with five electron groups have three dimensional geometries

• If there are six electron groups around the central atom, there are three case scenarios:

  • If the six electron groups are bonding groups, the angle between the bonds is 90° which maximize their separation

    • The geometry of this molecules is OCTAHEDRAL

    • E.g. SF₆

SF₆ Has an Octahedral Geometry

The molecular geometry of sulfur hexafluoride

• If five electron groups are bonding groups and one of them is a lone pair, the angle between the bonds is slightly less than 90° due to the Coulombic force of repulsion generated by the lone pair

  • The geometry of this molecules is SQUARE PYRAMIDAL

  • E.g. BrF₅

BrF₅ Has a Square Pyramidal Geometry

The molecular geometry of bromine pentafluoride

• If four electron groups are bonding groups and two of them are lone pairs, it is a flat molecule

• The angle between the bonds is 90° and the lone pairs are place at 180° minimizing the repulsion interaction between them

  • The geometry of this molecules is SQUARE PLANAR

  • E.g. XeF₄

XeF₄ Has a Square Planar Geometry

The molecular geometry of xenon tetrafluoride

In the table below it is a summary of the molecules that are part of the VSEPR

Summary of the VSEPR theory

Molecular Geometry

Predicting molecular geometry

• The molecular geometry of any molecule can be determined following simple steps:

1. Draw an accurate Lewis structure for the molecule or ion

2. Count the number of electron groups

3. Count the number of bonding groups and the number of lone pairs

4. Apply the VSEPR theory to predict the molecular geometry

Bond Angles

Predicting bond angles

• The bond angle of any molecule can be determined following simple steps:

1. Draw an accurate Lewis structure for the molecule or ion

2. Count the number of electron groups and place them as far as possible

3. Count the number of bonding groups and the number of lone pairs

4. Apply the VSEPR theory to predict the bond angle

Relative Bond Energies

Bond energy

• The bond energy is the amount of energy required to break 1 mol of bonds in the gas phase

  • It reflects the strength of a chemical bond between two atoms. The higher the energy, the stronger the bond

• It is influenced by factors such as bond type (ionic or covalent), bond length, and the presence of multiple bonds

• There are some important rules to follow when predicting the relative bond energy:

  • Single bonds are generally weaker than double bonds, and double bonds are weaker than triple bonds. Therefore, single bonds have smaller bond energy

  • The presence of multiple bonds increases bond strength because there is a greater Coulombic attraction between the electrons and the nuclei of the atoms

Bond lengths

Triple bonds has the highest bond energy (in red) and therefore they are the strongest ones

• Bond order is also a great chemical tool used to compare the bond energies between molecules

Bond order

• The bond order is the amount of bonding pairs of electrons between atoms

• Higher bond orders indicates greater stability and higher bond energy

• Bond order is usually an integer

• When a molecule has resonance, the bond order does not need to be an integer

• The bond order can be calculated by using three simple steps

1. Draw an accurate Lewis structures

2. Count the number of bonding pairs in the molecule

3. Count the number of bonding groups

4. Divide the number of bonding pairs by the number of bonding groups

Relative Bond Lengths

Bond length

• The bond energy refers to the distance between the nuclei of two atoms that are chemical bonded

• It is influenced by factors such as bond type (ionic or covalent), bond length, the presence of multiple bonds, and the atomic radius of the atoms

• There are some important rules to follow when predicting the relative bond lengths:

  • If the atomic radius of the atoms involved increases, the bond length will increase

    • This occurs because the valence electrons are pulled with less Coulombic force of attraction by the nuclei of the atoms

  • The presence of multiple bonds decreased the bond length because there is a greater Coulombic attraction between the electrons and the nuclei of the atoms

Bond Lengths

Triple bonds have the smallest bond length (in blue) because of the greater Coulombic forces of attraction

• Bond order is also a great chemical tool used to compare the bond lengths between molecules

  • If you want to know how to calculate the bond order, check the previous section about Relative Bond Energy

Dipole Moments

• Dipole moment measures the separation of positive and negative charge within a molecule

• If a molecule exhibits dipole moments, it is said that the molecule is polar

• There are two important factors to consider if a molecule is polar or not:

  • The presence of polar covalent bonds in the molecule

  • The molecular geometry

• There are cases of molecules that have polar covalent bond, but their polarity its neutralized by the molecular geometry

  • This occurs because the atoms are arranged in a particular way, so the individual dipole moments cancel each other out

  • You can think of a dipole moment as a “tug of war game”. If the atoms pull the electrons with the same force in opposite directions, the net dipole moment of the molecule is zero

  • E.g. The molecule below CH₃Cl is a polar molecule. The net dipole moment of the molecule points towards the chlorine atom which attracts the electrons with the strongest Coulombic force of attraction

The CH₃Cl Molecule

There are four polar covalent bonds in CH₃Cl which do not cancel each other out causing CH₃Cl to be a polar molecule; the overall dipole is towards the electronegative chlorine atom

• E.g. The molecule below CCl₄ is a nonpolar molecule. Even if the C-Cl bonds are highly polar due to the difference in electronegativity between C and Cl, the net dipole moment cancels out because all the chlorine nuclei are pulling the electrons with the same Coulombic force of attraction

The CCl₄ molecule

Though CCl₄ has four polar covalent bonds, the individual dipole moments cancel each other out causing CCl₄ to be a nonpolar molecule