Isoelectronic Species

Key Concepts

• Species with the same electronic configuration are said to be isoelectronic.¹.

  Name	Symbol	Electronic Configuration
  		shells	subshells
  Helium	He	2	1s2
  Hydride ion	H-	2	1s2
  Lithium ion	Li+	2	1s2
  He, H- and Li+ are all isoelectronic species because they have the same electronic configuration.

• Isoelectronic species must have the same number of electrons in total.

  Name	Symbol	Electronic Configuration		Total Number of Electrons
  		shells	subshells
  Helium atom	He	2	1s2	2
  Hydride ion	H-	2	1s2	2
  Lithium ion	Li+	2	1s2	2
  He, H- and Li+ are all isoelectronic species and have the same total number of electrons.

• To decide if two or more species are isoelectronic:

  (a) Write the electronic configuration for electrically neutral² atoms of the elements.

  (b) Write the electronic configuration of any ions:

  ⚛ For positively charged ions (cations) remove electrons from the atom's highest energy level

  ⚛ For negatively charged ions (anions) add electrons to the atom's highest occupied energy level

  (c) If two or more species have the same electronic configuration they are said to be isoelectronic.

• Note that isoelectronic species tend to have very similar chemical properties.

  Helium atom, hydride ion and lithium ion are all isoelectronic species:

  They all have the same ground-state electronic configuration as the Noble Gas helium: 1s²

  Just like helium, hydride ions and lithium ions are considered unreactive.

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Isoelectronic Configurations of Atoms and Ions

Consider an atom of the Nobel Gas (Group 18 element) argon, Ar.
Argon has atomic number of 18 (Z = 18).
An atom of argon has 18 positively charged protons in its nucleus and 18 negatively charged electrons "orbiting"³ the nucleus in the various energy levels:

First energy level: 2 electrons in the s subshell
Second energy level: 2 electrons in the s subshell and 6 electrons in the p subshell
Third energy level: 2 electrons in the s subshell and 6 electrons in the p subshell

Argon is the ONLY element whose atoms have the ground-state electronic configuration of 1s²2s²2p⁶3s²3p⁶

But there are ions of other elements that can have the ground-state electronic configuration 1s²2s²2p⁶3s²3p⁶

Consider an atom of chlorine, Cl.
Chlorine has an atomic number of 17 (Z = 17).
An atom of chlorine has 17 positively charged protons in its nucleus and 17 negatively charged electrons "orbiting" the nucleus in various energy levels:

First energy level: 2 electrons in the s subshell
Second energy level: 2 electrons in the s subshell and 6 electrons in the p subshell
Third energy level: 2 electrons in the s subshell and 5 electrons in the p subshell

Like all Group 17 (halogen) elements, atoms of chlorine can gain an electron to form an anion (negatively charged ion) with a charge of 1-.

Now, let's consider what happens if this chlorine atom gains an electron to form the chloride ion.

Where will this "extra" electron go?
It will enter the 3p subshell to complete this subshell (and also complete this energy level).

The electronic configuration of the chloride ion, Cl^-, is 1s²2s²2p⁶3s²3p⁶.
The electronic configuration of an argon atom in the ground-state is also 1s²2s²2p⁶3s²3p⁶.
We say that that the chloride ion is isoelectronic with the argon atom.

It is also possible for cations, positively charged ions, to be isoelectronic with the argon atom in its ground state.

Consider an atom of potassium, K, in its ground state.
Potassium has an atomic number of 19 (Z = 19).
There are 19 positively charged protons in the nucleus of an atom of potassium.
There are 19 negatively charged electrons "orbiting" the nucleus of a potassium atom in the ground state.

First energy level: 2 electrons in the s subshell
Second energy level: 2 electrons in the s subshell and 6 electrons in the p subshell
Third energy level: 2 electrons in the s subshell and 6 electrons in the p subshell
Fourth energy level: 1 electron in the s subshell

Like all Group 1 (alkali metal) elements, potassium will readily lose an electron to form a cation with a charge of +1.

The electron that is lost will come from the highest energy level, the fourth energy level:

And we can see that the potassium ion, K⁺, has the same electronic configuration as the chloride ion, Cl^-, and the same electronic configuration as an atom of argon, Ar.
Therefore, Ar, Cl^-, and K⁺ are said to be isoelectronic species.

Similary, we can see that an atom of calcium, Ca, (atomic number = 20) has en electronic configuration of 1s²2s²2p⁶3s²3p⁶4s²
Like all Group 2 (alkali-earth) metals, calcium will lose 2 electrons from its highest energy level to form a cation with a charge of 2+.
The calcium ion, Ca²⁺, will have the electronic configuration 1s²2s²2p⁶3s²3p⁶
Ca²⁺ is said to be isoelectronic with Ar, Cl^- and K⁺

Summary Table of Some Isoelectronic Species

If we consider the possible ions of the first 20 elements of the Periodic Table, we can draw up a table summarising which of the species are isoelectronic with atoms of a Group 18 (Noble Gas) element:

Worked Example: Determining Which Species are Isoelectronic

Question : Which of the species below is NOT isoelectronic with Al³⁺ ?
    ⚛ Na⁺
    ⚛ O^2-
    ⚛ B³⁺
Justify your answer.

Solution:

(Based on the StoPGoPS approach to problem solving.)

1. What is the question asking you to do?

   (i) Determine which species is not isoelectronic with the others.
   (ii) Justify your answer.

2. What data (information) have you been given in the question?

   Extract the data from the question:

   Formula of the species:
       Al³⁺ (reference)
       ⚛ Na⁺
       ⚛ O^2-
       ⚛ B³⁺

3. What is the relationship between what you know and what you need to find out?

   Definition: isoelectronic species have the same electronic configuration.

   Write the electronic configuration of each species:

   Al³⁺ (reference)
       Al, Z = 13, ground-state electronic configuration 1s²2s²2p⁶3s²3p¹
       Atom of Al loses 3 electrons from highest energy level (third energy level) to form Al³⁺:
       Al³⁺ ground-state electronic configuration: 1s²2s²2p⁶

   Na⁺
       Na, Z = 11, ground-state electronic configuration 1s²2s²2p⁶3s¹
       Atom of Na loses 1 electron from highest energy level (third energy level) to form Na⁺:
       Na⁺ ground-state electronic configuration: 1s²2s²2p⁶

   O^2-
       O, Z = 8, ground-state electronic configuration 1s²2s²2p⁴
       Atom of O gains 2 electrons to form O^2-:
       O^2- ground-state electronic configuration: 1s²2s²2p⁶

   B³⁺
       B, Z = 5, ground-state electronic configuration 1s²2s²2p¹
       Atom of B loses 3 electrons from second energy level to form B³⁺:
       B³⁺ ground-state electronic configuration: 1s²

4. Decide which species is NOT isoelectronic with Al³⁺

   All the species (Al³⁺, Na⁺ and O^2-) have the same electronic configuration (1s²2s²2p⁶) EXCEPT B³⁺ which has the electronic configuation 1s².
   B³⁺ is therefore NOT isoelectonic with Al³⁺

5. Is your answer plausible?

   Check that Al³⁺, Na⁺ and O^2- all have the same total number of electrons and that B³⁺ has a different number of electrons.
   no. electrons(Al³⁺) = Z(Al) -3 = 13 - 3 = 10
   no. electrons(Na⁺) = Z(Na) -1 = 11 - 1 = 10
   no. electrons(O^2-) = Z(O) +2 = 8 + 2 = 10
   no. electrons(B³⁺) = Z(B) -3 = 5 - 3 = 2

   All the species have 10 electrons in total EXCEPT B³⁺, so B³⁺ can not be isoelectronic with the other species.

6. State your solution to the problem "species that is not isoelectronic and justify answer":

   (i) B³⁺ is not isoelectronic with Al³⁺
   (ii) because the electronic configuration of B³⁺ (1s²) is NOT the same as the electronic configuration of Al³⁺ (1s²2s²2p⁶).