The elecrtrons in atoms can be split into two main categories. First, the core electrons which do not participate in forming bonds to other atoms. Second, the valence electrons, which are involved in bonding. The number of valence electrons, which equals the number in the outermost shell of an atom, determines how an atom bonds to other atoms.
The number of valence electrons an atom has is determined by the arrangement of the electrons shells and the total number of electrons found in a neutral atom. The complete description of how this is determined is unecessary when the aim is to be able to draw molecular structures. A simple way to count valence electrons is to look at the group number. In the figure at right the old group numbers are given as roman numerals. The number of dots arranged around each atomic symbol gives the number of valence electrons. Note that the value of the group number and the number of valence electrons is equal. Modern periodic tables have changed group number labels so that groups III – VIII (zero in this figure) are now labeled as groups 13 – 18. Still, the pattern— once seen— is easy to remember.
Atoms form bonds with one another based on the number of valence electrons they have. The general rule is known as the ‘octet rule’. This rule governs the number of electrons an atom needs to borrow from other atoms in order to have a full complement of eight electrons in its outermost shell. The reason for this rule is that eight electrons is a sort of magic number. When an atom has eight valence electrons it is in a lower energy state than when it does not. An atom in a Lewis dot structure has an octet when the sum of shared and un-shared electrons equals eight.
Conventionally, electrons are considered in pairs (an exception is a molecular radical with an odd number of electrons). Electron pairs in a bond are represented as a line between the atomic symbols: H—O—H. Non-bonding pairs (also called lone pairs) are represented by a pair of dots near the atomic symbol they belong to (for example, :N—F3). Lone pair (also called non-bonding pair) electrons are valence shell electrons that do not participate in bonding.
An easy way to find formal charge is to draw a circle around an atom in a Lewis Structure and count the electrons inside the circle (counting only half of each bond as being inside the circle). Subtract the result from the number of valence electrons to find formal charge.
Note formal charges next to atoms by using a small + or – number in a circle.
Total formal charge for a molecule must equal the overall charge of the molecule.
Minimize formal charges so that they are all as small as possible and so that as few atoms as possible have formal charges. For central atoms from period 3 or higher it is permitted to allow further bonds that exceed an octet as long as the formal charge is zero or at most +1.
Formal charge is simply a way to decide whether electrons are equally or unequally shared in a molecule. When formal charge is zero it means that the number of electrons shared or owned by an atom equals the number of valence electrons. When an atom has a formal charge that is negative it means that the atom has taken electrons away from its neighbors and is sharing unfairly by taking more electrons. Usually atoms with high electronegativity will do this. In other words, non-metals and especially non-metals at the furthest right-hand extreme of the periodic table. When an atom has a formal charge that is positive it means that the atom has given away electrons to its neighbors and is sharing more generously. Usually atoms with low electronegativity will do this. In other words, metals may be found to occur in some molecules with a positive formal charge.
Not all molecular model drawings (Lewis diagrams) require the calculation of formal charge. You only need to calculate formal charge under any one of the following conditions:
|Usual Bonding Patterns|
|1||1||H Li Na K||1||0|
|2||2||Be Mg Ca||2||0|
|15||5||N P As||3||1|
|16||6||O S Se||2||2|
|17||7||F Cl Br I||1||3|