DISCOVER
×

How to calculate bond angles

Updated April 17, 2017

Bond angles refer to the angles created in the formation of molecular compounds. When molecular compounds are created, a variety of bond angles are possible based on the orientation of the atoms and electrons. Nevertheless, molecular geometry is important to the value of the bond angles. The standard types of molecular shapes created when lone pairs of electrons are not present on the centre atom are the linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. When lone pairs are present on the central atom, then the possible shapes are trigonal planar bent, tetrahedral trigonal pyramidal, tetrahedral bent, trigonal bipyramidal distorted tetrahedron, trigonal bipyramidal T-shaped, trigonal bipyramidal linear, octahedral square pyramidal and octahedral square planar. You can calculate the bond angle based on the shape of the molecule.

Analyse the formula of the molecular compound and draw the corresponding Lewis structure. You can draw the Lewis structure by placing the proper atom at the centre and counting the number of valence electrons. You can make the proper double, single or triple bonds depending on the number of electrons that you need to place in the structure.

Apply the VSEPR method to the Lewis structure. The VSEPR method is the Valence Shell Electron Pair Repulsion Method. To use this principle, you must count the number of electron pairs. This includes both the lone pairs and the bonded pairs. In the VSEPR model, double bonds and triple bonds can be treated as single bonds.

Distinguish the number of bonding pairs from the number of lone pairs. If the molecular structure has no lone pairs, then it is either linear, trigonal planar, tetrahedral, trigonal bipyramidal or octahedral. A linear shape has two bonding pairs and a bond angle of 180 degrees. A trigonal planar shape has three bonding pairs and a corresponding bond angle of 120 degrees. A tetrahedral shape has four bonding pairs and a bond angle of 109.5 degrees. A trigonal bipyramidal shape has five bonding pairs and a bond angle of 120 degrees between the equatorial bonds, and 90 degrees between the axial bonds. An octahedral shape has six bonding pairs and bond angles of 90 degrees.

Determine whether or not the molecule has lone pairs. If the molecule has lone pairs, first, count the number of bonding pairs. The number of bonding pairs will indicate the general geometry, while the number of lone pairs will indicate the molecular geometry. Nevertheless, a molecule with two bonding pairs and one lone pair has a trigonal planar bent shape. It has bond angles of less than 120 degrees. A molecule with three bonding pairs and one lone pair is tetrahedral trigonal pyramidal, and a molecule with two bonding pairs and two lone pairs is tetrahedral bent. Both these shapes have bond angles of less than 109.5 degrees. A molecule with four bonding pairs and one lone pair is trigonal bipyramidal distorted tetrahedron and has bond angles of less than 120 degrees in the equatorial plane and less than 90 degrees in the axial plane. A molecule with three bonding pairs and two lone pairs has a trigonal bipyramidal T-shaped geometry and bond angles of less than 90 degrees. A molecule with two bonding pairs and three lone pairs has a trigonal bipyramidal linear shape and a bond angle of 180 degrees in the axial plane. A molecule with five bonding pairs and one lone pair is octahedral square pyramidal and has bond angles of less than 90 degrees. A molecule with four bonding pairs and two lone pairs is octahedral square planar and has bond angles 90 degrees.

Tip

The value of the bond angles corresponds to the molecular and electron pair geometry.

Things You'll Need

  • Molecular compound formula
  • Lewis structure
Cite this Article A tool to create a citation to reference this article Cite this Article

About the Author

Mara Pesacreta has been writing for over seven years. She has been published on various websites and currently attends the Polytechnic Institute of New York University.