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.

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 (i.e. only 2 electrons around it).

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, p_x, p_y and p_z) 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.  Notice that the C now has a full octet of electrons.

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

Chirality Note

Organic molecules have a three-dimensional (3D) structure that is crucial for understanding their chemical behavior and reactivity, especially in biological systems. While Lewis structures are helpful for showing connectivity and electron distribution, they fail to represent the spatial arrangement of atoms (i.e. they are inherently 2-dimensional). For example, methane (CH₄) is often drawn as a flat cross in a Lewis structure, but in reality, it has a tetrahedral shape with bond angles of 109.5°. This 3D arrangement influences how molecules interact with each other, making it essential to visualize structures beyond simple 2D representations.  Because of the tetrahedral geometry of sp³ carbon atoms, molecules can have handedness.  Just like your left and right hands are similar, but different, molecules can be left and right-handed. 

Number of Bonds Neutral Atom Like to Have

The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to achieve a noble-gas configuration. Hydrogen, with one valence electron in the 1s orbital, needs only one more electron to reach the helium configuration (1s²), so it forms one bond. Carbon has four valence electrons (2s² 2p²) and requires four more to reach the neon configuration (2s² 2p⁶), resulting in four bonds. Nitrogen, with five valence electrons (2s² 2p³), needs three additional electrons and forms three bonds. Oxygen, which has six valence electrons (2s² 2p⁴), needs two more and forms two bonds. Finally, the halogens have seven valence electrons, require one more to complete their valence shell, and form one bond.

The problem with Lewis Structure.

The problem with Lewss structures is they don't show the 3D nature of carbon atoms and other organic structures.Organic molecules have a three-dimensional (3D) structure that is crucial for understanding their chemical behavior and reactivity. While Lewis structures are helpful for showing connectivity and electron distribution, they fail to represent the spatial arrangement of atoms. For example, methane (CH₄) is often drawn as a flat cross in a Lewis structure, but in reality, it has a tetrahedral shape with bond angles of 109.5°. This 3D arrangement influences how molecules interact with each other, making it essential to visualize structures beyond simple 2D representations.

Practice Time - Lewis Structure Tool