Electronegativity, Oxidation Numbers, and Formal Charge
Even in covalent bonds, the allocation of electrons between atoms is often not uniform: one atom will usually take more of the electron than another. In fact, the only situation where the covalent bond is not slightly asymmetrical is when two atoms of the same element join together. Chemists have devised a measure of how much atoms want their electrons called electronegativity: the lower (more negative) the number, the more an atom wants an electron. Electronegativities for all the elements are listed in our "Reference" section.
Electronegativities are useful for predicting dipoles in molecules. Dipoles occur when the molecule is not electrically symmetrical, either because of its shape or because one atom is more electronegative than others. In a molecule like ClF, for example, fluorine has the highest electonegativity, meaning that the shared electrons will gravitate toward it, making one side of the molecule (fluorine's) negatively charged, and the other end (chlorine's) will be positively charged. The symbol for a dipole is the small-case Greek letter delta (δ), followed by either a plus or minus sign, indicating the charge. Dipoles are typically written in superscript when appearing along with the molecules they describe. For example, the ClF molecule would be written as δ+Cl-Fδ- using dipole notation. The numerical representation of a dipole's charge is called the dipole moment.
Another measurement similar to electronegativity should be mentioned, called electron affinity. This value is the energy released when an atom captures a free electron. Many atoms (especially those on the right of the periodic table) will acquire an electron readily, releasing a great deal of energy, but others must actually absorb energy to accept the electron (such as the noble gases, who are perfectly happy with filled s and p subshells).
To map how electrons behave in covalent bonds, two measures can be used. The first is called an oxidation state, whereby electrons in a covalent bond are allocated to one electron or another; no sharing is indicated. The atom with the strongest electronegativity will take as many atoms as it needs to fill its s and p subshells, giving it a negative charge. If there is only one other atom in the molecule, it will have a positive charge as needed to give up its electrons. If there are three or more molecules, the atom with the second-highest electronegativity will try to give up as few electrons as possible (the most electronegative atom will "steal" electrons from the least electronegative atom first). In ions, the sum of all the molecule's oxidation states must equal the charge on the ion.
For example, what are the oxidation states in the molecule NH3. Using an electronegativity table, we see that nitrogen is the most electronegative element in the molecule. Since it needs three electrons to reach a noble gas configuration, it will take one electron from each hydrogen atom. Therefore, the nitrogen has a -3 oxidation state and each of the hydrogen atoms has a +1. The oxidation states all add up to zero, the overall charge of the molecule.
Let's try an ion: SO42-. Looking in our electronegativity chart, we see that oxygen is the most electronegative atom, so each oxygen will take 2 electrons to fill its s and p subshells. The total charge of the four oxygens is -8, but the molecule has a -2 charge, meaning that the sulphur only has to surrender six electrons. The fact that sulphur has such a positive charge is unrealistic, so another representation of atomic charge was developed.
Because oxidation numbers do not really represent the sharing of electrons that occurs in a covalent bond, another measure of how electrons are allocated has been devised, called formal charge. Formal charge is not completely accurate, either, because it assumes the boding pair of electrons is shared equally between atoms, but it is more reasonable than oxidation numbers. Formal charge is calculated by taking the number of valence electrons in an atom, subtracting the number of its electrons that do not participate in a bond, and subtracting half the electrons in any covalent bonds.
The CO2 molecule will be our example for formal charge. In this bond, carbon wants four more electrons, while each oxygen wants two. Therefore, two pairs of electrons are shared between the carbon and each oxygen atom. Carbon has 4 valence electrons, and all of its electrons participate in the bonding process. It takes half of the four pairs of electrons being shared, giving it a formal charge of zero. The oxygens each have 6 valence electrons, and four do not participate in the bond (since two electrons are being shared). Each oxygen takes half of the two pairs of electrons being shared (or two electrons), so the formal charge is 6-4-2, or zero as well. As with oxidation states, the sum of all formal charges must equal the charge on the molecule; in carbon dioxide, the total molecular charge is zero, as expected.
Formal charge is extremely useful in determining which configuration of a molecule is most stable; we will return to oxidation states in the "Redox Reactions" chapter.