Lewis Dot Structures

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During chemical bonding it is the valence electrons which move amongst different atoms. In order to keep track of the valence electrons for each atom and how they may be shared in bonding we use the Lewis Dot Structure for atoms and molecules. In this approach we represent the valence electrons as dots around the element symbol. For example, oxygen has 6 valence electrons, so we write the symbol O for oxygen and surround it with 6 dots:

Oxygen

The unpaired electrons are represented as single dots, and the paired electrons as double dots. The placement of the single or double dots around the symbol is not critical. Alternatively, we can represent the paired electrons as a line. That is, we replace the double dots as shown below:

Oxygen

Let's consider other examples. A sodium atom has 11 electrons, but only one is a valence electron. The other 10 are inside a closed shell with a Neon electron configuration. Thus, we draw the Lewis structure for a sodium atom as the symbol Na with a single dot:

LewisDotSodium

A chlorine atom has 17 electrons, but only 7 of these are valence electrons. Thus, we draw the Lewis structure as:

LewisDotCl

In Ionic Bonds valence electrons are completely transferred (not shared). Thus, we write the Lewis structure for NaCl as:

NaClLewisDot

As you can see Chlorine is now surrounded by 8 electrons in the n=3 shell and Sodium has lost its one valence electron in the n=3 shell. Of course, Sodium, is still surrounded by the 8 electrons of the n=2 shell, but we do not show electrons in the inner closed shells.

For period 2 elements, where all the valence electrons of an atom are in s and p orbitals, we find that the Lewis dot structure of molecules will often follow the Octet Rule:

Octet Rule - Atoms tend to gain, lose, or share electrons until they are surrounded by eight electrons (4 electron pairs).

Using Lewis dot structures and the octet rule, we can predict and represent the electronic structure of covalently bonded molecules. For example, when two chlorine atoms, each with 7 valence electrons, come together to form a diatomic chlorine molecule, the Lewis structure shows that there will be a sharing of two electrons between the two chlorine atoms which allows both chlorine to be surrounded by 8 electrons.

ChlorineMolecule

Of course, hydrogen is a period 1 element, with only has a 1s orbital, so it has a maximum of two electrons allowed in its valence shell. When two hydrogen atoms come together into a diatomic H2 molecule the Lewis structure shows that there will be a sharing of two electrons between the two hydrogen, allowing both hydrogen to be surrounded by a closed n=1 shell of 2 electrons:

HydrogenMolecule

We can represent the electronic structure and reaction of hydrogen and chlorine atoms to form HCl with Lewis structures:

PolarHCl

For diatomic oxygen, the Lewis dot structure predicts a double bond.

O2

While the Lewis diagram correctly predict that there is a double bond between O atoms, it incorrectly predicts that all the valence electrons are paired (i.e., it predicts that each valence electron is in an orbital with another electron of opposite spin). Later we will examine a more advanced theoretical approach called Molecular Orbital Theory, which correctly predicts both the double bond of O2 and its unpaired valence electrons. Generally, Lewis-dot structures have the advantage that they are simple to work with, and often present a good picture of the electronic structure. Let's consider another example. For diatomic nitrogen, the Lewis-dot structure correctly predicts that there will be a triple bond between nitrogen atoms:

N2

This triple bond is very strong. The strength of the triple bond makes the N2 molecule very stable against chemical change, and, in fact, N2 is considered to be a chemically inert gas.There is a relationship between the number of shared electron pairs and the bond length.

BondBond Length
N-N1.47 Å
N=N1.24 Å
N≡N1.10 Å

The distance between bonded atoms decrease as the number of shared electron pairs increase.

Demo:
  1. Ball and Stick Models of NH2-NH2, NH=NH, N≡N Structures