In the previous post, we talked about electron configurations, what they show, how to read, and how to write them based on energy levels and sublevels.

As a reminder, electron configurations show the number and location of electrons in an atom or an ion. For example, the electron configuration of fluorine is1s²2s²2p⁵.

The number before each letter indicates the main energy level, and this is the principal quantum number, n. The letters (s, p, d, f) tell us the type of the orbital, and the exponent is for the number of electrons.

The 1s² means there are 2 electrons in the s orbital that is in the first row. Next, we have the orbitals in the second energy level because fluorine is in the second row of the periodic table. In this level, there are 2 electrons in an s orbital, and 5 electrons in the p orbital:

This is the very least you should know in addition to the order of filling the energy levels and sublevels so please go over the detailed post on the electron configurations before getting to ions.

Electron Configuration of Ions

The first thing you need to remember here is that cations are formed by losing an electron(s), and anions are formed by gaining an electron(s).

The charge of the ion is a result of an imbalance between the number of protons and electrons. If it is a cation, then the positive indicates how many more protons it has compared to the number of electrons. For anions, the charge tells how many extra electrons there are compared to the number of protons.

Recall, this pattern for the formation of anions and cations: Metals tend to lose electron(s) and become cations (positively charged ions).

Nonmetals tend to gain an electron(s) and become anions (negatively charged ions).

Notice that the number of protons is not changed, and the ions are charged because, unlike atoms, their number of protons and electrons is not equal.

Electron Configuration of Main Group Cations

Now, how do we determine the electron configuration of an ion? If it is, for example, a +1 charged cation, that means the atom has lost one electron. This electron is going to be from the outermost valence shell as these are the electrons farthest away from the nuclei and thus not as strongly attracted to it.

Let’s see an example on Na. It is in the first group, so it loses one electron to become Na⁺.  The electron configuration of sodium is 1s^²2s^²2p^⁶3s^¹, and the electron is removed from the energy level with the greatest n value – 3s. Therefore, the electron configuration of the Na⁺ ion will be 1s^²2s^²2p^⁶:

Notice that the ion has a configuration with a complete shell of p orbitals which is characteristic of noble gases. In fact, this is the electron configuration of Ne, and we say that Na⁺ and Ne are isoelectronic (same electronic structure). The reason for this is that, remember, noble gases are very stable because of the low energy level of complete orbitals.

Na (1s²2s²2p⁶3s¹) ⟶ e⁻ + Na⁺([He]2s²2p⁶) [isoelectronic with Ne ([He] 2s²2p⁶)]

Let’s now determine the electron configuration of the Ca²⁺ ion. It is formed when an atom of Ca (1s²2s²2p⁶3s²3p⁶4s² or [Ar]4s²) loses two electrons from the outer shell 4s. 

Ca ([Ar]4s²) ⟶ 2e⁻ + Ca²⁺([Ar]) [isoelectronic with Ar]

This pattern explains why the metals in the first group become +1, the ones in the second group become +2, and Al, for example, becomes +3. It takes removing one electron from a metal in the first group to obtain the electron configuration of the previous noble gas, it takes two for the group two metals, etc.

Electron Configuration of Anions

Anions are formed when the atom gains as many electrons as necessary to attain the electron configuration of the next noble gas in the periodic table. For example, oxygen is in group 6, and therefore, it will need two electrons to attain the electron configuration of Ne:

Notice again that the two electrons go to the outermost valence orbital. Oxygen has 6 valence electrons – those in the 2s and 2p orbitals, however, since p sublevels are higher in energy, and they are the only ones capable of accepting additional electrons, the two electrons go to the 2p orbitals. The electron configuration of the oxide ion (O^2-) is therefore, 1s^²2s^²2p^⁶.

O (1s²2s²2p⁴) + 2e⁻⟶ + O^2- ([He]2s²2p⁶) [isoelectronic with Ne ([He] 2s²2p⁶)]

Another common type of monoatomic anions are the halides. Halogens are in group 7, and therefore, they only need one electron to attain the electron configuration of the noble gas following them in the periodic table. For example, bromine takes one electron and becomes isoelectronic to t Kr:

Br([Ar] 4s²3d¹⁰4p⁵) + e⁻ ⟶ Br⁻([Ar] 4s²3d¹⁰4p⁶) [isoelectronic with Kr ([Ar] 4s²3d¹⁰4p⁶)]

Electron Configuration of Transition Metal Cations

In contrast to main-group ions, transition metal ions do not usually attain a noble gas configuration. This is because the ns level is the outermost level, and the (n-1)d is considered an inner level therefore, it will take too much energy to remove those electrons and achieve a noble gas configuration. Therefore, the cation of a transition metal is formed by removing first the electrons from the ns (highest principal quantum number) orbital and then from the (n -1)d orbitals.

For example, the electron configuration of Zn is 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰ or [Ar]4s²3d¹⁰, and it loses the two electrons from the 4s orbital to become Zn²⁺[Ar]3d¹⁰. This is “not a bad” electron configuration considering the filled d orbitals.

Check Also

• Atomic Orbitals
• Electron Configurations
• Orbital Diagrams
• Aufbau’s Principle, Hund’s Rule, and Pauli’s Exclusion Principle
• Hund’s Rule
• Pauli Exclusion Principle
• Quantum Numbers (n, l, m_l, m_s)
• Bohr Model of the Hydrogen Atom
• Rydberg Formula
• The Photoelectric Effect
• Calculating The Energy of a Photon
• Ionization Energy
• Electron Affinity
• Energy, Wavelength, and Frequency Practice Problems