1.6:
Lewis Structures and Formal Charges
Lewis structures are simplified representations of chemical bonds between atoms. Consider ethene, where two carbon atoms with four valence electrons each are surrounded by four hydrogen atoms having one valence electron each. Hence, ethene has a total of twelve valence electrons.
It takes twenty-four electrons to satisfy the octets for two carbon atoms and the duets for four hydrogen atoms. Thus, twelve electrons must be shared between the atoms. With only twelve electrons available, they must all be bonding electrons to allow all atoms to reach a stable electronic configuration.
Sometimes, atoms in polyatomic structures do not exhibit the standard number of valence electrons. In such cases, a nonzero formal charge, F, is associated with the anomalous atom. The formal charge is equal to the number of valence electrons on the neutral atom minus the number of bonds and unshared electrons on that atom.
For example, in the ammonium ion, each hydrogen atom has one valence electron in its neutral form and is bonded by a single bond, yielding a formal charge of zero.
This leaves the nitrogen atom with four valence electrons instead of its usual five valence electrons. Since it is missing one electron, it bears a formal charge of one-plus. Thus, the overall charge on the ammonium ion is one-plus.
Stable configurations generally minimize formal charges. For some larger elements, the octet may be expanded to share more than eight electrons to reach a configuration with fewer formal charges.
For example, the sulfur atom in sulfur trioxide can form both single and double bonds with the surrounding oxygen atoms. However, the double bond arrangement is preferred, as it reduces the overall formal charge on the atom.
Other exceptions to the octet rule include boron-bearing compounds, which have a target of six shared electrons, rather than eight, because achieving an octet requires a formal charge.
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:
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).
A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO).
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.