Electron Configurations

The electron configuration is the distribution of the electrons in an atom. It is how the electrons are arranged that helps us understand chemical bonding and chemical reactions.

So, if you know how electrons are arranged, especially the valence electrons, you will better understand chemical formulas, equations and reactions. But, first a few rules need to be stated concerning the placement of electrons into energy levels.

Electron Arrangement Rules

Now that the quantum numbers have been introduced, the possible electron arrangements for the elements in an atom can be described.

1. The number of electrons equals the number of protons in a stable atom.
2. The number of electrons each energy level can hold is 2n².
3. The number of sub shells in an energy level is equal to n.
4. The s sub shell has only one possible position, the p sub shell has three, the d sub shell has five and the f sub shell has seven. Each possible position is an orbital.
5. Only two electrons can occupy each orbital.
6. Hund's Rule: When a p, d, or f sublevel is being filled, one electron will occupy each orbital before pairing.
7. The maximum number of electrons is two in any s sub shell, six in any p sub shell, ten in any d subshell, and fourteen in any f sub shell.
8. Pauli Exclusion Principle: No two electrons in an atom have the same four quantum numbers.
9. Aufbau Principle: An electron occupies the lowest energy level available, filling in orbitals of higher energy levels until all electrons are distributed.

Using the above rules, you can easily diagram the distribution of the electrons in an atom. Just determine how many:

• electrons the atom has
• energy levels to be used
• sub energy levels to be used
• orbitals there will be in each sub energy level

Let's look at carbon as an example.

According to the first rule, carbon has six electrons and according to the second rule would use two energy levels. Using rule three, you can determine that the first energy level will have only one sub energy level and only one orbital (s). The second energy level will have two sub energy levels and four orbitals ( one s and three p). Now, using Hund's rule and Aufbau's principle distribute carbon's six electrons:

• two electrons go into the first energy level's only orbital (s orbital)
• then two go into the second energy level's first orbital (s orbital)
• then one in the second energy level's first p orbital
• finally one in the second energy level's second p orbital

carbon = 1s² 2s² 2p¹ 2p ¹ 2p⁰

Electron and Orbital Notations

Using our knowledge of quantum numbers and the distribution rules, let's see how electron configurations and orbital notations of the elements are represented.

Complications

If you’re thinking this is too easy to be true, you’re right. There are a few complications as the atoms get larger. As the energy levels get farther from the nucleus, the distance between the energy levels decreases.

As a matter of fact, it is believed that the energy levels actually overlap. Therefore, some energy levels start filling orbitals before the previous energy level is finished filling its sublevels.

The first time this is encountered is with potassium, in which the 4s starts to fill before the 3d.

There's More ...

The second complication has to do with a variation of Hund’s Rule that takes into account the minimizing of the electron-electron repulsion.

It states, the most stable arrangement of electrons is the arrangement with the maximum number of unpaired electrons. So, when the transition metals’ orbitals are filling with electrons, at d⁴ and d⁹, an electron from the s jumps up into the d⁵ or d¹⁰.

A Shortcut

Writing out electron configurations and orbital notations can become awkward as the atoms increase in the number of electrons. So, scientists have agreed on a type of shorthand to help make writing electron configurations and orbital notations less cumbersome.

The shorthand involves using the abbreviation of the last noble gas (placed in brackets) to indicate that all the orbitals to that point are full. Then the notation is continued as usual.