Electron Configuration of Ions — Overview & Examples

We know how to write the electron configurations for neutral atoms. Now we can learn to write the electron configurations for ions. The electron configurations for ions follow the same rules as neutral atoms. However, we add or remove electrons from the valence shell.

Remember, atoms want to have a full octet. For elements with high or low electronegativities, the best way to form an octet is to become an ion. Elements with middle ground electronegativities share electrons to form covalent bonds. The formation of an ion is the result of gaining or losing valence electrons. A cation, a positively charged ion, forms when an atom loses valence electrons. An anion, a negatively charged ion, forms when an atom gains electrons. We can use the periodic table to help us determine the electron configurations of ions.

Electron Configurations of the Alkali and Alkaline Earth Metals

The alkali and alkaline earth metals are in groups 1 and 2 of the periodic table. Their valence shell is always the outer s-orbital. The easiest way for these elements to have a full octet is to lose the s-orbital electrons. To write the electron configuration, remove the electrons from the outer orbital.

Let's look at calcium as an example.

1. Neutral calcium has the condensed electron configuration, Ca: [Ar] 4$s^{2}$.
2. The first cation has the condensed electron configuration, $Ca^{1+}$: [Ar] 4$s^{1}$.
3. The second cation has the condensed electron configuration, $Ca^{2+}$: [Ar].

In the first cation, the 4s orbital has one less electron. For the second cation, we remove both 4s electrons. The second cation has the same electron configuration as the preceding noble gas.

Electron Configurations of the Transition Metals

We know that the transition metals fill their s-orbitals before their d-orbitals. But, something strange happens with transition metal ions. Transition metals lose electrons from their s-orbitals first. As the electrons fill the d-orbital, the electron repulsion increases. The increase in electron repulsion causes the s-orbital to have more energy. So, the filled s-orbital is in a higher energy state. The higher energy state means the s-orbital electrons are removed first.

As an example, let's look at Nickel (Ni). An element's common ions are called its oxidations states. Oxidation-reduction reactions are a unique type of chemical reaction. Often, we recognize the reactions by their change in color. Nickel's most common oxidation state is $Ni^{2+}$, which is positively charged. It can also oxidize to $Ni^{3+}$.

1. Neutral Nickel has the condensed electron configuration, Ni: [Ar] 4$s^{2}$ 3$d^{8}$.
2. $Ni^{2+}$ has the condensed electron configuration, $Ni^{2+}$: [Ar] 3$d^{8}$.
3. $Ni^{3+}$ has the condensed electron configuration, $Ni^{3+}$: [Ar] 3$d^{7}$.

Here we see that Nickel loses its first two electrons from the s-orbital. Next, the electrons leave the d-orbital.

Electron Configurations of the Lanthanides

The Lanthanides are in the sixth period of the periodic table. Some of the Lanthanides have unique electron configurations. Lanthanum (La), Cerium (Ce), and Gadolinium (Gb) all have one electron that occupies the 5d orbital. The other lanthanides fill the 6s orbital, followed by the 4f orbital. For all the Lanthanides, the first two electrons lost are from the 6s orbital. What about the three unique elements? The third electron comes from the 5d orbital. But, for the rest of the lanthanides, the third electron comes from the 4f orbital.

For example, let's look at Cerium (Ce).

1. Neutral Cerium has the condensed electron configuration, Ce: [Xe] 4$f^{1}$ 5$d^{1}$ 6$s^{2}$.
2. $Ce^{2+}$ has the condensed electron configuration, $Ce^{2+}$: [Xe] 4$f^{1}$ 5$d^{1}$.
3. $Ce^{3+}$ has the condensed electron configuration, $Ce^{3+}$: [Xe] 4$f^{1}$.

In cerium, the first two electrons are lost from the 6s orbital. The third electron lost is from the d orbital.

As another example, let's look at Neodymium (Nd).

1. Neutral Neodymium has the condensed electron configuration, Nd: [Xe] 5$f^{4}$ 7$s^{2}$.
2. $Nd^{2+}$ has the condensed electron configuration, $Nd^{2+}$: [Xe] 5$f^{4}$.
3. $Nd^{3+}$ has the condensed electron configuration, $Nd^{3+}$: [Xe] 5$f^{3}$.

For Neodymium, we see that the first two electrons are again lost from the 6s orbital. The next electrons lost are from the 4f orbital.

Electron Configurations of the Actinides

The actinides are in the seventh period of the periodic table. Like the Lanthanides, many of the actinides have unusual electron configurations. The elements with unique electron configurations are Actinium (Ac), Thorium (Th), Protactinium (Pa), Uranium (U), Neptunium (Np), Curium (Cm), and Lawrencium (Lr). All have one 6d orbital electron. The rest of the actinides only fill the 7s orbital and then the 5f orbital. The actinides follow the same rules as the lanthanides. The 7s orbital electrons are removed first. Next, we remove the d-orbital electrons. When there are no d-orbital electrons, then we remove the f-orbital electrons.

Electron Configurations of the Post-Transition metals

The post-transition metals are towards the bottoms of groups three, four, five, and six. They have massive atomic radii and readily form cations. So, they can lose electrons from both their p and s orbitals and have many different charges. All the post-transition metals have full d-orbitals. The full d-orbital is stable. So, they do not lose d-orbital electrons. For the post-transition metals, you first remove the p-orbital electrons. Next, remove the d-orbital electrons.

Electron Configurations of the Non-Metals

The chemistry of the non-metals is unique. Most of the molecules that contain non-metals will have covalent bonds. However, many non-metals will form ionic bonds with the alkali, alkaline-earth, and transition metals. The non-metals will usually gain electrons. So when we write the electron configurations, we add electrons to the valence shell.

Let's look at the halogen, fluorine.

1. Neutral flourine has the condensed electron configuration, $F: [He] 2s^{2} 2p^{7}$.
2. $F^{-}$ has the condensed electron configuration, $F^{-}: [He] 2s^{2} 2p^{8}$.

We see that the p-orbital gained an electron! This completes its octet!

(Video) Electron Configuration of Ions - Mg2+, P3-, Fe2+, Fe3+

by The Organic Chemistry Tutor

In this practical video, The Organic Chemistry Tutor walks you through writing the electron configuration of ions. He provides four examples: magnesium, phosphorous, and two iron ions. The magnesium and phosphorous examples demonstrate the difference between cations and anions. The iron examples are extra insightful because they are transition metal ions with d-orbitals. Remember, the d-orbitals have unusual rules! If you're looking for more worked examples of the electron configuration of ions, this video is for you!