Lone Electron Pairs



 

Learning Objective

  • Recognize the effect of lone electron pairs on molecules’ geometries.

Key Points

    • Orbitals containing the various bonding and nonbonding pairs in the valence shell will extend out from the central atom in directions that minimize their repulsions.
    • A nonbonding orbital has no atomic nucleus at its far end to draw the electron cloud toward it; the charge in such an orbital will therefore be concentrated closer to the central atom.
    • Nonbonding orbitals exert more repulsion on other orbitals than do bonding orbitals.

Terms

  • lone paira valence set of two electrons that exists without bonding or sharing with other atoms
  • coordination numberin chemistry and crystallography, the number of a central atom’s neighbors in a molecule or crystal

Molecular Geometries with Lone Pair Electrons

So far, we have only discussed geometries without any lone pairs of electrons. As you likely noticed in the table of geometries and the AXE method, adding lone pairs changes a molecule’s shape. We mentioned before that if the central atom also contains one or more pairs of nonbonding electrons, these additional regions of negative charge will behave much like those associated with the bonded atoms. The orbitals containing the various bonding and nonbonding pairs in the valence shell will extend out from the central atom in directions that minimize their mutual repulsions.

AXE methodLone pairs change a molecule’s shape.

Coordination Number and the Central Atom

Coordination number refers to the number of electron pairs that surround a given atom, often referred to as the central atom. The geometries of molecules with lone pairs will differ from those without lone pairs, because the lone pair looks like empty space in a molecule. Both classes of geometry are named after the shapes of the imaginary geometric figures (mostly regular solid polygons) that would be centered on the central atom and have an electron pair at each vertex.

In the water molecule (AX2E2), the central atom is O, and the Lewis electron dot formula predicts that there will be two pairs of nonbonding electrons. The oxygen atom will therefore be tetrahedrally coordinated, meaning that it sits at the center of the tetrahedron. Two of the coordination positions are occupied by the shared electron-pairs that constitute the O–H bonds, and the other two by the non-bonding pairs. Therefore, although the oxygen atom is tetrahedrally coordinated, the bonding geometry (shape) of the H2O molecule is described as bent.

The effect of the lone pair on waterAlthough the oxygen atom is tetrahedrally coordinated, the bonding geometry (shape) of the H2O molecule is described as bent.

The Repulsive Effect of the Lone Pair Electrons

There is an important difference between bonding and non-bonding electron orbitals. Because a nonbonding orbital has no atomic nucleus at its far end to draw the electron cloud toward it, the charge in such an orbital will be concentrated closer to the central atom; as a consequence, nonbonding orbitals exert more repulsion on other orbitals than do bonding orbitals. In H2O, the two nonbonding orbitals push the bonding orbitals closer together, making the H–O–H angle 104.5° instead of the tetrahedral angle of 109.5°.

The electron-dot structure of NH3 places one pair of nonbonding electrons in the valence shell of the nitrogen atom. This means that there are three bonded atoms and one lone pair for a coordination number of four around the nitrogen, the same as occurs in H2O.

The Lewis dot structure for ammonia, NH3.The lone pair attached to the central nitrogen creates bond angles that differ from the tetrahedral 109.5 °.

We can therefore predict that the three hydrogen atoms will lie at the corners of a tetrahedron centered on the nitrogen atom. The lone pair orbital will point toward the fourth corner of the tetrahedron, but since that position will be vacant, the NH3 molecule itself cannot be tetrahedral; instead, it assumes a pyramidal shape, more specifically, that of a trigonal pyramid (a pyramid with a triangular base). The hydrogen atoms are all in the same plane, with the nitrogen outside of the plane. The non-bonding electrons push the bonding orbitals together slightly, making the H–N–H bond angles about 107°.

In 5-coordinated molecules containing lone pairs, these non-bonding orbitals (which are closer to the central atom and thus more likely to be repelled by other orbitals) will preferentially reside in the equatorial plane. This will place them at 90° angles with respect to no more than two axially-oriented bonding orbitals. We can therefore predict that an AX4E molecule (one in which the central atom A is coordinated to four other atoms X and to one nonbonding electron pair) such as SF4 will have a “see-saw” shape.

Example of a see-saw structureTry to imagine this molecule teetering on each end, and you will have a visual representation of a see-saw.

Substituting nonbonding pairs for bonded atoms reduces the triangular bipyramid coordination to even simpler molecular shapes.

Interactive: Unshared Electrons and the “Bent” ShapeUse the 3D model to see how unshared electrons repel those that are shared in the bonds between hydrogen and oxygen, causing the molecule to have a “bent” shape.