Shapes of Molecules

Shapes of Molecules

Molecular Shapes

  • Valence Shell Electron Pair Repulsion (VSEPR) Theory: This theory is used to explain the shapes of molecules. It postulates that electron pairs around a central atom orient themselves to minimize repulsion.

  • Bond Pairs and Lone Pairs: A bond pair is shared by two atoms while a lone pair belongs to a single atom. Lone pairs repel more than bond pairs due to their closer proximity to the nucleus, making the lone pair occupy more space than bond pairs.

  • Linear Shape: A molecule without any lone pairs and with two areas of electron density (such as BeCl2) is linear. The bond angle is 180°.

  • Trigonal Planar Shape: Molecules with three areas of electron density and no lone pairs (e.g., BCl3) have a trigonal planar shape. The bond angle is 120°.

Alteration of Bond Angles

  • Lone Pair/Lone Pair Repulsion: The repulsion between two lone pairs of electrons is greater than between two bonding pairs or a bonding pair and a lone pair. This causes deviation from the expected bond angles.

  • Trigonal Pyramidal Shape: Molecules with three bonding pairs and one lone pair (NH3, for example) adopt a trigonal pyramidal shape, with bond angles slightly less than 109.5° due to lone pair/lone pair repulsion.

  • Bent Shape: Molecules with two bonding pairs and two lone pairs (like H2O) have a bent or V shape, with bond angles less than 109.5°, again due to lone pair/lone pair repulsion.

Unusual Shapes

  • Octahedral Shape: SF6 is an example of a molecule with an octahedral shape due to six areas of electron density from bond pairs, with bond angles of 90°.

  • Square Planar Shape: Examples of molecules with a square planar shape include XeF4, which has four bond pairs and two lone pairs, reducing the bond angles to be 90°.

Keep in mind that the presence of lone pairs of electrons can greatly influence the shape of molecules and their bond angles due to increased repulsion. The VSEPR theory provides a basis for predicting the shapes of molecules by considering the electron density around a central atom.