### Lewis Electron Dot Structures

Lewis electron dot structures are an easy way both to predict what sorts of covalent bonds will be formed between atoms, and to visualize what these bonds look like. They are very useful and very important, so be sure you have a good grasp on this concept.

Lewis dot structures are only applicable to covalent bonds, so only the right portion of the main group elements (the groups of hydrogen, beryllium, boron, carbon, nitrogen, oxygen, fluorine, and helium) can be drawn this way--no more pesky transition metals! Start by drawing the atomic symbols of the bonding elements. Let's use CO2 as an example. In simple molecules, there is a rule that the atom that appears once is the central atom, so we'll draw CO2 like this:

Not too revolutionary yet, right? Now, using little dots, two on each side of an atom, draw in the valence electrons for each element (carbon has four and oxygen has six, to refresh your memory).

Uisng the octet rule, we know that each atom wants to share until it has eight electrons. Looking at this Lewis structure, it appears that all the atoms could be satisfied if two pairs of electrons were shared between each oxygen and the carbon, as shown below:

Now that the bonding electrons have been established, each pair is typically represented by a line, as shown here:

We can now see that carbon dioxide consists of two covalent bonds between each oxygen and the carbon atom. Each atom is happy, with a full octet of electrons. The pairs of electrons that are shared in the covalent bond are called, surprisingly enough, bonding pairs, and those that did not (two pairs on each oxygen) are the non-bonding pairs.

For some molecules, one can draw several plausible Lewis structures, such as the ones below for the NOCl molecule:

So which one is the most likely? We return to formal charge to find out which structure is the most appropriate. The arrangement that keeps the formal charge on each atom to a minimum is the best structure. In the first possibility, the formal charges are -1 on the oxygen, 0 on the nitrogen, and +1 on the chlorine. On the second, the formal charge on all atoms is zero, making it the likely winner. The third has formal charges of -2 on the nitrogen and +1 on the oxygen and chlorine atoms. The final possibility has a formal charge of -1 on the nitrogen, +1 on the oxygen, and 0 on the chlorine. By comparing charges, we see that the second possibility is the best.

Lewis structures can be drawn for any covalently-bonded atoms or ions, but in the case of an ion, the extra or missing electrons must be remembered. For example, the molecule CH3COOH, acetic acid, can be drawn using Lewis structures. We must remove the final H, which is really an H+ ion ionically bonded onto the rest of the molecule. Given that carbon atoms always want to make four bonds, they should be the central atoms of the molecule. The Lewis structure becomes:

The blue-tinged electron on the final oxygen indicates that it was taken from the hydrogen atom; without this stolen electron, the oxygen would not be content with just one bond. The entire structure is placed in brackets because it is an ion, with the charge noted above.

An other interesting configuration of electrons can result when the total number of valence electrons in a molecule is odd. In these cases, the octet rule cannot be obeyed--one atom must have too many or too few electrons. There simply aren't the right number of electrons to make everybody happy. A common example is nitrogen dioxide, NO2, since N has five valence electrons and the oxygen atoms each have six, giving a total of 17. The Lewis structure for this molecule is shown below:

Nitrogen here has too many electrons, because its gets two for each of the four shared bonds, plus the extra electron sticking out on top. Note that we could have single-bonded one oxygen, giving one electron less than it would like, to give nitrogen eight, but since oxygen is more electronegative than nitrogen, it will take all the electrons it can get (up to eight, of course). Molecules with one unpaired electron are called free radicals, and they are very reactive, because they either want to get rid of or find an extra electron. Nitrogen dioxide, for instance, often doubles up with another NO2 molecule to share the extra electron and stabilize the molecule, forming N2O4. Such pairs of free radicals are called dimers:

In this case, the nitrogen atoms now have way too many electrons (10 each, with their 5 shared pairs), but the increased stability of the molecule overrules the nitrogens' discomfort.

Finally, atoms past phosphorus in the Periodic Table can "expand their octets" and bond with more atoms (or higher-order bonds than usual) through additional bonding in the d subshell. For example, XeCl4 is one of the few compounds that can be made using noble gases. Because xenon's octet is already full, it must expand its octet to produce the following Lewis structure:

There are a few more plot twists involving Lewis dot structures, but we'll discuss these on the next page, which will describe different types of covalent bonds.

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