Water valence shell electron pair

This simple method of deducing the structure of molecules is called Valence Shell Electron Pair Repulsion Theory (VSEPRT). It says that all electron pairs, both bonding and nonbonding, in the outer or valence shell of an atom repel each other. This simple approach predicts (more or less) the correct structures for methane, ammonia, and water with four electron pairs arranged Lctrahedrally in each case. 

Wilson and Geratt [56] discussed a pair-function model constructing geminals from non-orthogonal one-electron orbitals. Their calculations, performed on the water molecule, supported qualitative valence-shell electron-pair (VSEPR) models [57] of directed valence. 

Valence-shell electron-pair repulsion (VSEPR) theory and the concept of hybridization suggest that the water molecule has two O—H bonds and two non-bonded pairs arranged tetrahedrally. More accurate calculations show that this does not provide a true picture of the total electron density in H20. 

In Section 4.3, we learned that the shape of a molecule is an important factor in determining the properties of the substances that it composes. For example, we learned that water would boil away at room temperature if it had a straight shape instead of a bent one. We now develop a simple model called valence shell electron pair repulsion (VSEPR) theory that allows us to predict the shapes of molecules from their Lewis structures. 

The theoretical basis for a molecule possessing a particular geometric shape is the concept that electron pairs, whether they are part of a covalent bond (as in a bonding pair) or not (as in a nonbonding pair), will repel each other. In methane, water, and all other molecules, this repulsion means that the electron pairs will get as far away from each other as they can get. The general name for this theory is the valence-shell electron-pair repulsion (VSEPR) theory. 

A Lewis structure for the water molecule (which shows only the valence shell electrons) is depicted in the following structure. There are two nonbonding pairs of electrons around the oxygen as well as two bonding pairs. 

A pair of nonbonding electrons localized in the valence shell on a single atom. Examples are the lone electron pair on the nitrogen atom in ammonia and the two lone pairs on the oxygen atom in water. In these cases, the lone pairs participate in hydrogen bonding interactions. 

A Lewis base is a molecule that can donate a lone pair of electrons to fill the valence shell of a Lewis acid (Following fig.). The base can be a negatively charged group such as a halide, or a neutral molecule like water, an amine, or an ether, as long as there is an atom present with a lone pair of electrons (i.e. O, N or a halogen). 

A complete structural model for a molecule also shows the positions of electrons not involved in covalent bonding. For example, in the Lewis structures of formaldehyde and water (Figure 2.1), the oxygen atom in each carries two pairs of unshared electrons from the outer valence shell. Each of these electrons, not involved in a covalent bond, is represented by a dot. The oxygen atom in water has four nonbonding electrons, and the oxygen atom in formaldehyde carries two pairs of unshared electrons, represented by four dots on the oxygen atoms of the two molecules in the Lewis structure. 

---