Molecules lone pairs count

Each atom in a polyatomic molecule completes its octet (or duplet for hydrogen) by sharing pairs of electrons with its immediate neighbors. Each shared pair counts as one covalent bond and is represented by a line between the two atoms. A Lewis structure does not portray the shape of a polyatomic molecule it simply displays which atoms are bonded together and which atoms have lone pairs. 

Count the number of lone pairs and bonding partners an atom actually has within a molecule. You can do this by looking at the Lewis structure. 

As is obvious from the drawings, the E atoms in these simple molecules retain a lone pair of electrons may be used for a dative bond to another ML fragment (10-13). Data on complexes of types 12 and 13 are summarized in Tables III and IV, respectively. The dilemma that arises is that from the electron counting procedures one would expect two types of bonds— covalent and dative (see Ref. 62 for a discussion of the differences between dative and covalent bonding)—for complexes 11-13, but the ML fragments... 

To determine the molecular structure, we must count the electron pairs around the sulfur atom. In each resonance structure the sulfur has one lone pair, one pair in a single bond, and one double bond. Counting the double bond as one pair yields three effective pairs around the sulfur. According to Table 13.8, a trigonal planar arrangement is required, yielding a V-shaped molecule ... 

Formal charge is the apparent electronic charge of each atom in a molecule, based on the electron-dot structure. The number of valence electrons available in a free atom of an element minus the total for that atom in the molecule (determined by counting lone pairs as two electrons and bonding pairs as one assigned to each atom) is the formal charge on the atom ... 

There are two alternative approaches to hybridization for the water molecule. For example, the electron pairs around the oxygen atom in water can be considered as having nearly tetrahedral symmetry (counting the two lone pairs and the two bonds equally). All four valence orbitals of oxygen are used, and the hybrid orbitals are sp. The predicted bond angle is then the tetrahedral angle of 109.5° compared with the experimental value of 104.5°. Repulsion by the lone pairs, as described in the VSEPR section of Chapter 3, is one explanation for this smaller angle. 

A measure of the hydrogen-bonding ability of a molecule expressed in terms of number of possible hydrogen-bond acceptors. In particular it is calculated as the count of lone pairs on oxygen and nitrogen atoms in the molecule. 

The formal charge of an atom is computed as the difference between the number of valence electrons that a neutral atom would have and the number of electrons that belong to it in the Lewis structure when one counts lone pair electrons as belonging fully to the atom, while electrons in covalent bonds are split equally between the atoms involved in the bond. Tlie total of the formal charges on an ion should be equal to the charge on the ion, and the total of the formal charges on a neutral molecule should be equal to zero. 

When we are to determine how many electron groups that surround an atom, the Lewis structure can be of great help (see the previous section 2.23 Lewis structure). From the Lewis structure of a given molecule you can simply count how many bonds and lone pairs that surround an atom. That way you have the number of electron groups. The VSEPR theoiy tells us that these electron groups will be placed as far apart as possible. In the following example we will use the VSEPR theory to predict the molecular geometries of a water molecule and a carbon dioxide molecule. That way we will discover why a carbon dioxide molecule is linear and why a water molecule is V-shaped. 

From the Lewis structure the number of electron groups surrounding the central atoms is counted. The oxygen atoms in the water molecule are surrounded by four electron groups (two lone pairs and two single bonds). The carbon atom in the carbon dioxide molecule is surrounded by two electron groups (two double bonds). 

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