Lewis Structure

To show how atoms connect together to form a molecule, you draw a Lewis structure. The character of each molecule depends on the exact way elements attach to each other. Chemists use a standard way to depict the structure of molecules.

Only Some Elements

Figure 1: Simplified periodic table shows elements most often found in Lewis structure problems

Molecular structure problems become easier when you only consider the nonmetal elements which most often make molecules, Figure 1.

Even though other elements do form molecules, ninety-eight percent of molecules you build come from this smaller number of elements. This certainly accounts for the elements you find the most often in organic chemistry.

This simplifies how much information  when you begin the problems of molecular structure.

Elements and Valence Electrons

Valence Electrons

Figure 2: Valence electrons important in molecular structure.

Recall valance electrons consist outer electrons from the outer most electron shell. Those electrons act in chemical reactions.

Core electrons reside inside the outer electrons. All core electrons have entirely filled electron shells. This makes them chemically inert.

Lewis Dot Structure of Elements

To represent an individual element as a Lewis structure, write the element’s symbol and show the number of valence electrons with dots around the symbol.

Dot Rules

Orbitals Around Elements

In order to write a Lewis dot structure, imagine the electrons around an element, Figure 3. 

Figure 3: Vacant electron orbitals surround an elemental symbol

Electron orbitals start life as a smudgy cloud. Each elemental symbol has four sides. The four sides correspond to four orbitals. Each orbital houses two electrons. This gives a total of eight electrons around around each element.

How Electrons Fill Orbitals

Each electron gets a dot. Electrons go around an element with the following rules, Figure 4:

Figure 4: Order of electron electrons filling orbitals for Lewis dot structure of an element

Electrons singly occupy orbitals before they can doubly fill an orbital. Nonmetal elements can have no more than eight electrons in any valence shell.

Figure 5 shows what the Lewis structure of each element looks like.

Figure 5: Lewis structure of important nonmetal elements

You can see a step-wise progression across each row. Every column has the same number of valence electrons.

 Each one has the same number of dots with the same orientation. They differ only between their elemental symbol.

This simplifies deducing the correct Lewis structure. It also leads to simplifying how to write a Lewis structure.

Single and Paired Electrons

Pay special attention to the fact there are two ways electron dots are placed on the side of an element. One electron or two electrons occupies a side. The difference between these situations dictates what happens next.

This puts an electron into one of two possible categories:

1. lone unpaired electron
2. pair of electrons

Figure 6 shows how a  Lewis structure of an element converts to an element ready to attach to other elements. 

Figure 6: single dots convert to bonds, paired electrons are left alone

Single Unpaired Electron

Every single lone dot converts to a chemical bond connector. This happens when another single dot from another element fills the side with two dots.

Pairs of Electrons

In contrast, if the side of an element has two dots then the two dots are left alone.

Periodic Table

Figure 7: Nonmetals with lone pairs and spurs  used to make covalent bonds

The single electrons have a “spur” where they can bond to other nonmetals, Figure 7.

This makes building blocks out of elements you can assemble in the same specific way.

Simple Molecules

Figure 8: Water constructed from two hydrogens and one oxygen

Figure 8 shows how this works to construct Lewis structures. 

Water has the formula H₂O. Each hydrogen has one spar. Oxygen has two spars and two nonbonding lone pairs. 

Elements fit together to satisfies the needs of each atom connected in the way shown in Figure 8.

Multiple Bonds

Sometimes molecules have more spars to connect than elements . To account for unused connections, elements  form multiple bonds.

Double Bonds

Formaldehyde, CH₂O, provides an example of when elements form more than one bond , Figure 9. Each hydrogen has one connection. Carbon possesses four connections. Finally, oxygen has two connections and two pairs of nonbonding electron pairs.

Figure 9: Formaldehyde with the formula CH₂O assembled from its elements. Unused connections form an additional bond

Three of the available connections on carbon find connections: one to each hydrogen and one to oxygen. This leaves one unused connection on carbon with another unused connection on oxygen (blue).

Each of these vacant connections satisfy their need to form a bond by forming a connection with one another. That forms a double bond between carbon and oxygen.

Triple  Bonds

In addition, sometimes more than one connection remains unused. Hydrogen cyanide gives example of this kind of molecule, Figure 10.

Figure 10: carbon and nitrogen form a triple bond 

Hydrogen has one space for connection. Carbon possesses four. Nitrogen has three  connection and one set of nonbonding paired electrons.

After the elements connect through their open spots, carbon has two unused connectors. Nitrogen has two unused connections (blue). The two unused connections on each element form one chemical bond. This creates a triple bond between carbon and nitrogen.