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1.6: Lewis Structures and Formal Charges
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Organic Chemistry

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Lewis Structures and Formal Charges
 
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1.6: Lewis Structures and Formal Charges

Lewis symbols can be used to indicate the formation of covalent bonds, which are shown in Lewis structures—drawings that describe the bonding in molecules and polyatomic ions. The periodic table can be used to predict the number of valence electrons in an atom and the number of bonds that will be formed to reach an octet. Group 18 elements, such as argon and helium, have filled electron configurations and thus rarely participate in chemical bonding. However, atoms from group 17, such as bromine or iodine, need only one electron to reach the octet. Hence, atoms belonging to group 17 can form a single covalent bond. The atoms of group 16 need two electrons to reach an octet; hence, they can form two covalent bonds. Similarly, carbon, which belongs to group 14, needs four more electrons to reach an octet; thus, carbon can form four covalent bonds.

Consider the Lewis structure of the chlorine molecule:

Figure1

The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called lone pairs) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:  Cl–Cl

A single shared pair of electrons is called a single bond. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet.

A double bond forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH2O (formaldehyde).

Figure2

A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO).

Figure3

The formal charge of an atom in a molecule is the hypothetical charge the atom would have if the electrons in the bonds are evenly distributed between the atoms. The formal charge can be calculated by subtracting the sum of the number of nonbonding electrons and the number of bonds on an atom (or half the number of bonding electrons) from the number of valence electrons of the neutral atom:

Formal charge = # valence shell electrons (free atom) − # lone pair electrons − # bonds

The formal charge calculations can be double-checked by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a neutral molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion. Remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. The formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.


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