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**Writing Electron Configurations****Dr. MJ Patterson**

The main goal of this module is to be able to write the electron configurations for atoms and ions. This web page will provide an outline of how to write these configurations.

First, a couple of definitions:

- paired spins = two electrons who have opposite spins, one up (or +1/2) and one down (or -1/2). Frequently the two electrons will be in the same orbital.
- parallel spins = two electrons who have the same spin, both up or both down

**Electron Configurations**

An electron configuration shows how the electrons are divided up among the orbitals in an atom. There are three common notations that you should be comfortable with - box, spectroscopic and noble gas.

To begin with, you need to know how many electrons you will be working with in a particular atom or ion. To review the number of electrons in an atom or ion, see the page in this module's web site called Counting Electrons in an Atom or Ion.

The overall guiding principle behind the rules which will follow is that we want to take the atom's or ion's electrons, and arrange them among the orbitals so that the energy of the entire arrangement is the smallest possible amount. All of these rules are derived from that principle.

The next several section will outline these rules, and then we will put them all together to write the configurations in several examples.

**The Aufbau Principle**

This principle simply states that we can build up an electronic configuration one electron at a time by putting each electron in the lowest energy orbital available. The energy ordering of the orbitals can be remembered from this diagram which we first saw in the last module:

After you draw the diagram, connect the orbitals in a diagonal fashion as follows. Start with the 1s subshell and draw a diagonal line through it from the lower right to upper left corner. Then, snake that line back between the 1s and 2s subshells, parallel to the first. Turn this line around and go through the 2s. Turn around again and come back between 2s and 2p. Turn around and go through 2p and 2s. Continue in this fashion until you have worked your way through the diagram.

If you follow this line through the diagram, it traces out the subshells in this order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p ...

This is the order of the subshells from lowest energy to highest. When we arrange electrons among orbitals, we will start with the lowest energy subshell, the 1s, and add electrons until it is filled. Then, we will move on to the next subshell, the 2s, until it is filled. We will continue in this manner until we run out of electrons.

**The Pauli Exclusion Principle**

This principle says that an orbital is filled when it contains 2 electrons. After that, you have to put the electrons in a different orbital.

Let's look at how this translates into subshells. Each type of subshell contains a different number of orbitals. And, each orbital can hold at most 2 electrons. The following table shows how many electrons each type of subshell can hold.

Subshell Type | # of Orbitals | Maximum # of Electrons |

s | 1 | 2 |

p | 3 | 6 |

d | 5 | 10 |

f | 7 | 14 |

**Hund's** **Rule**

Within a subshell, the electrons will occupy the orbitals singly first, and will only pair up when there are no longer any empty orbitals available in that subshell.

**Magnetism**

Individual atoms and ions are either paramagnetic or diamagnetic (ferromagnetism is a property of bulk matter, not just an individual atom or ion). To determine whether an atom is paramagnetic or diamagnetic, look at the box notation for the electron configuration. If all of the electrons are paired, it is diamagnetic. If one or more electrons are unpaired, it is paramagnetic.

(Remember that paired spins mean that for every electron with spin up, there is one electron with spin down.)

**Examples:**

Write the electronic configuration for the following atoms or ions. Use the box notation, spectroscopic notation and noble gas notation. Are they paramagnetic or diamagnetic?

1. H

2. He

3. C

4. Ne

5. Cl^{-}

Solutions:**1. H** = 1 electron (atomic number from the periodic table = 1)

We will start with the lowest energy subshell, the 1s, which has 1 orbital. The single electron is placed in this orbital.

Box notation: We can explicitly show each electron in each orbital. Each orbital is shown as a box. Each electron is shown as an arrow. An arrow point up indicates spin up, and an arrow point down indicates spin down. Each subshell of orbitals is labeled underneath the grouping of orbitals for that subshell.

Spectroscopic notation: The orbital is named, and the number of electrons inside the orbital is shown as a superscript.

1s^{1}

Hydrogen atoms are paramagnetic with one unpaired electron.

**Helium:**

The atomic number is 2, so He has 2 electrons to place.

Spectroscopic Notation: He = 1s^{2}

Box notation:

Note that helium is diamagnetic since all of the electrons are paired. Note, also, that the 1s subshell is now filled since it only contains one orbital and can hold at most 2 electrons.

**Carbon:**

Atomic number = 6, so C has 6 electrons to place.

The 1s subshell is filled with just 2 electrons. From the Aufbau diagram above, we see that the next subshell to fill is the 2s. It will also be filled with 2 electrons, so we will have to move to the 2p subshell for the remaining 2 electrons.

spectroscopic notation: C = 1s^{2}2s^{2}2p^{2}

Noble gas core notation: Instead of writing out all of the electrons in the configuration, we can write out just the ones since the last noble gas. Find carbon on the periodic table, and then go backwards until you reach a noble gas. In this case, it is helium. We can use a shorthand to indicate all of the electrons that are identical to helium's configuration by putting He in square brackets, and substituting it for those electrons.

Noble Gas Core Notation: C = [He]2s^{2}2p^{2}

In this case, we did not really save any effort, but for much bigger atoms, the noble gas core notation can be very convenient.

Box notation: Each s subshell has one orbital, but the p subshell has 3 orbitals. Hund's rule tells us that we have to put the 2p

electrons in separate orbitals since there is room to do so.

Thus, carbon is paramagnetic with two unpaired electrons.

**Neon:**

Atomic number = 10, so there are 10 electrons to place.

Spectroscopic Notation: Ne = 1s^{2}2s^{2}2p^{6}

Noble gas core notation: [Ne]

Box notation:

Since all of neon's electrons are paired, it is diamagnetic. The six electrons in the p subshell completely fill it. If you needed to place 11 electrons (for sodium) you would have to go to the next subshell, 3s.

**Cl**^{-}

Atomic number = 17 for Cl; Add 1 for the anion = 18 electrons in Cl^{-}

Spectroscopic Notation: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}

Noble gas core notation: [Ar]

Box notation: