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
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
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.
Orbitals Around Elements
In order to write a Lewis dot structure, imagine the electrons around an element, Figure 3.
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:
- only one electron per box
- each side must be individually filled before any side can be doubly filled
- There can be no more than a total of eight electrons per 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.
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:
- lone unpaired electron
- pair of electrons
Figure 6 shows how a Lewis structure of an element converts to an element ready to attach to other elements.
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.
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.
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.
Sometimes molecules have more spars to connect than elements . To account for unused connections, elements form multiple 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.
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.
In addition, sometimes more than one connection remains unused. Hydrogen cyanide gives example of this kind of molecule, Figure 10.
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.