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Lewis Structures

The world of molecules is built on a foundation of shared electrons. Understanding how these electrons arrange themselves is crucial for predicting molecular behavior and properties. Lewis structures, also known as Lewis dot structures, provide a simple yet powerful tool for visualizing the distribution of valence electrons around atoms in a molecule. By depicting atoms as symbols and their valence electrons as dots, Lewis structures offer a clear picture of bonding and lone pairs, laying the groundwork for further exploration of molecular structure, reactivity, and function.

Let's start by looking at Lewis dot structures of atoms.

Lewis Dot Structures of Atoms

To write an element’s Lewis dot structure, we place dots representing its valence electrons, one at a time, around the element’s chemical symbol.  Since most atoms we deal with in organic chemistry have at most four orbitals (i.e. s, px, py and pz) simply draw the atom with 4 orbitals on it.

Let's try a lithium atom.  The symbol for lithium is Li so write the symbol and draw 4 orbitals, "hoopie things", around the symbol.  Looking at the periodic table we see that lithium has one valence electron (1s22s1), so we add one electron "dot" in one of the orbitals.

Since lithium has one unpaired electron it likes to have one bond to it.  For example, Li-Cl is a common compound of Lithium.

Up to four dots are placed around each atom (in any order, as long as elements with four or fewer valence electrons have no more than one dot in each orbital).  It's Hund's rule.  For example, the Lewis dot diagrams for beryllium (Be), boron (B), and carbon (C) are as follows.  Be, B, and C like to have two, three and four bonds respectively (e.g. BeCl2, BH3, CH4) because each has two, three, and four unpaired electrons.

The next dots, for elements with more than four valence electrons, are again distributed one at a time, each paired with one of the first four. For example, the electron configurations for atomic nitrogen, oxygen, and fluorine are as follows.  Electrons in excess of four become lone pairs.  Nitrogen has three bonds and one lone pair of electrons (e.g. NH3).  Oxygen atoms can have two bonds and two lone pairs (e.g. H2O).  Fluorine can have one bond and has three lone pairs (H-F).

You could envision making bonds with other atoms to the orbitals with only one electron.

Lewis Structures of Molecules

A hydrogen atom only has one orbital (i.e. 1s), so it only needs one orbital on it.

If we combine four hydrogen atoms with one carbon atom we make the Lewis structure for methane (CH4). Each bond is made up of one hydrogen electron and one carbon electron.

We can drop the orbitals, "whoopie things", to simplify the structures - just remember only two electrons per orbital or bond.

 

 

Take Note
  • All valence electrons must be shown for each element in the molecule.
  • Generally, electrons are paired. Unpaired electrons are observed in odd electron molecules such as NO and NO2.
  • Generally, each atom acquires eight electrons in its valence shell except hydrogen acquires only two electrons.
  • Multiple bonds (double and triple bonds) can be formed by C, N, O, P, and S. 
  • Hydrogen atoms are terminal atoms.
  • Central atoms are generally less electronegative than terminal atoms.

In drawing Lewis structures for relatively small molecules and polyatomic ions, the structures tend to be more stable
when they are compact and symmetrical rather than extended chains of atoms.

Practice Time - Lewis Structure Tool

Action
  1. Click the "+" icon in the app below and then click on the H (hydrogen) atom.
    • Notice there is one unpaired electron.  This means it can form one bond.
  2. Click the "+" icon and add another H atom.
  3. Click the unpaired electron on one H atom, then click the unpaired electron on the other H atom to form a bond.

That's it! You've made your first Lewis structure.

 

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The Octet Rule

Recall from general chemistry and our discussions of electron configurations of atoms, that atoms tend to lose, gain, or share electrons to reach a total of eight valence electrons, called an octet.  We now know from quantum mechanics that the number eight corresponds to one s and three p valence orbitals, which together can accommodate a total of eight electrons.  An exception to the octet rule is hydrogen and helium, whose 1s2 electron configurations give a full n = 1 shell with a maximum of only two electrons.  Hence hydrogen abides by the duet rule.