What are Orbital Diagrams?
Electron orbital diagrams are diagrams used to show the energy of electrons within the sublevels of an atom or atoms when used in bonding. Single atom diagrams (atomic orbital diagrams) consist of horizontal lines or boxes for each sublevel. Within orbitals, arrows indicate the spin direction of the occupant electrons. Multi-atom diagrams (molecular orbital diagrams) show the energy of electrons in molecular orbitals. Typically, they only show the outermost electrons. This article will explore the basics of how to draw each type of diagram, and important rules to follow in their construction.
Orbital Diagram Basics
As mentioned in the introduction, diagrams make use of horizontal lines which are filled with arrows to represent the spin direction of electrons. When constructing an orbital diagram for either a singular atom or two atoms, one must first begin with the electron configurations for the atom(s) involved. To learn more about electron configurations, see our article on writing electron configurations.
Let’s begin with looking at the atomic and molecular orbital diagrams for Hydrogen. As the first atom in the periodic table, Hydrogen has one electron in the 1s shell. When Hydrogen forms a covalent bond with another Hydrogen atom, H2 is formed. Below are the corresponding atomic and molecular orbital diagrams for this element.
To learn more about how to construct these diagrams for other elements, follow along with the steps below.
Steps for Constructing an Orbital Diagram
Atomic Orbital Diagrams
- Beginning with your selected element, determine the atomic number. Once the atomic number has been identified, write the electron configuration. As an example, we will use Argon, whose atomic number is 18 and electron configuration is 1s22s22p63s23p6.
- Next, moving from the bottom up, we will draw the sublevels for each principal energy level. Each orbital follows a pattern – one line for the s orbital, three lines for the p orbital, five lines for the d orbital, and seven lines for the f orbital. Notice that the number of lines we draw for each orbital equals half the number of electrons each orbital can hold! Complete this process up to the highest principal energy level and orbital. See the image below to see this process broken down.
- Fill each orbital (horizontal line) with the number of occupant electrons. To do so, start with the lowest energy level.
- For the s orbitals, first draw an arrow pointing up and an arrow pointing down.
- For the p orbitals, draw one arrow pointing up on each of the lines first. Then, fill the lines with an arrow pointing down, until the number of arrows drawn is equal to the electron occupancy. This process is the same for the d and f orbitals. The following image demonstrates the order for filling, which should always follow the electron configuration.
- Finally, count the number of arrows on your diagram – the total should be equal to your element’s atomic number! Below is the completed argon orbital diagram.
Molecular Orbital Diagrams
- Draw two lines to create three columns.
- In the first column, draw the atomic diagram for your first element. In the third column, draw the atomic diagram for your second element. For this example, we will be using Chlorine. To allow space for the next step, draw the orbitals further apart than you would for an atomic orbital diagram.
- In the second (center) column, draw lines above and below each orbital. The bottom line(s) represent the bonding orbital(s), and the top line(s) represent the antibonding orbital(s). See below for a quick example of the s and p orbitals.
When we put the two processes together, we get a product that looks like the image below.
- Fill the middle column using the atomic orbital diagrams, starting with the bonding orbitals of each orbital, then filling the antibonding orbitals. These diagrams follow the same process as the atomic orbitals, so be sure not to exceed the designated number of occupant electrons.
- Count the number of arrows in the middle column – the total should equal the number of arrows in both the first and third columns. If equal, you are done!
Rules for Construction
When creating our diagrams, there are a few main steps we must follow. All rules are equally important, so be sure not to miss any!
- Reference electron configurations for creating diagrams, being sure not to exceed the total number of electrons.
- Aufbau’s Principle – orbitals must be filled from lowest energy to highest energy. To ensure that we maintain proper filling order, simply follow the electron configuration from start to finish.
- Hund’s Rule – when filling orbitals, there must be one electron in each orbital of the sublevel before adding another electron. This means we will draw one up-pointing arrow in each p, d, or f orbital before we draw the down-pointing arrows.
- Pauli Exclusion Principle – the electrons in each orbital must have opposite spin. This means the orbital filling diagram will only have arrows pointing in the opposite direction.
- As with all rules, there are always some exceptions. Copper is such an element who prefers to fill its 3d orbital fully, leaving the 4s orbital with only one electron.