Electronic Configuration Exceptions: Cr, Cu, Mo, Ag, Pd and Au

The electronic configuration of an atom describes how its electrons are arranged in different shells, subshells and orbitals. The Aufbau principle provides a useful method for predicting this arrangement by filling lower-energy orbitals before higher-energy orbitals.

However, some elements do not follow the simple filling order expected from this principle. Chromium, copper, molybdenum, silver, palladium and gold are important examples.

These are commonly studied as electronic configuration exceptions in transition elements. Their actual ground-state configurations depend on the relative energies of orbitals and the total energy of the atom.

Understanding these configurations is important for learning transition-metal chemistry, predicting common ions and solving inorganic chemistry questions in JEE and NEET.

Also Read: Chemical Exceptions in Inorganic Chemistry

What Is Electronic Configuration?

Electronic configuration is the arrangement of electrons in the orbitals of an atom or ion.

For example, the electronic configuration of sodium is 1s² 2s² 2p⁶ 3s¹. This means that sodium has one electron in its outermost 3s orbital.

The electronic configuration of an atom helps explain:

• Its position in the periodic table.

• The number of valence electrons.

• Its common oxidation states.

• Its tendency to form ions.

• Some of its magnetic and chemical properties.

To understand exceptional configurations, students should first learn three basic principles.

Aufbau principle

The Aufbau principle states that electrons generally occupy orbitals in order of increasing energy.

For example, the 4s orbital is generally filled before the 3d orbitals in the simplified filling sequence used to predict the configurations of many elements.

However, orbital energies can change as electron occupancy changes. Therefore, the Aufbau principle is a useful guide, not an unbreakable rule.

Pauli exclusion principle

The Pauli exclusion principle states that an orbital can contain a maximum of two electrons, and those electrons must have opposite spins.

Hund's rule of maximum multiplicity

Hund's rule states that electrons occupy orbitals of the same energy individually, with parallel spins, before pairing begins.

For example, three electrons in a p subshell occupy the three p orbitals separately before any pairing takes place.

These principles help explain electron arrangements, but they do not independently predict every exceptional ground-state configuration.

Why Do Electronic Configuration Exceptions Occur?

The simple Aufbau sequence does not always predict the actual ground-state configuration of an atom because the energies of orbitals can be very close.

In transition elements, the energy difference between the outer s and d orbitals is often small. Electron-electron interactions and the overall energy of the atom also influence the preferred arrangement.

Some configurations involving half-filled or completely filled d subshells are observed because the complete arrangement can be energetically favourable.

However, students should not assume that half-filled or fully filled subshell stability alone explains every exception.

The important factors include:

• Relative energies of s and d orbitals.

• Electron-electron repulsion.

• Exchange interactions between electrons.

• The overall energy of the atom.

• Changes in orbital energies as electrons are added or rearranged.

1. Electronic Configuration of Chromium (Cr)

Chromium has the atomic number 24.

According to a simple Aufbau-based prediction, its electronic configuration would be:

[Ar] 3d⁴ 4s²

However, the experimentally established ground-state configuration of chromium is: [Ar] 3d⁵ 4s¹

Why does chromium show an exceptional configuration?

Instead of having four electrons in the 3d subshell and two in the 4s orbital, chromium has five electrons in the 3d subshell and one in the 4s orbital.

This arrangement gives chromium a half-filled 3d subshell.

The observed configuration is energetically more favourable than the simple predicted arrangement. The relative energies of the 3d and 4s orbitals, together with electron interactions, contribute to this result.

Key point for exams: The ground-state electronic configuration of chromium is [Ar] 3d⁵ 4s¹, not [Ar] 3d⁴ 4s².

2. Electronic Configuration of Copper (Cu)

Copper has the atomic number 29.

A simple Aufbau-based prediction gives: [Ar] 3d⁹ 4s²

However, the actual ground-state configuration of copper is: [Ar] 3d¹⁰ 4s¹

Why does copper show an exceptional configuration?

Copper has a completely filled 3d subshell and one electron in the 4s orbital.

The observed arrangement has a lower total energy than the simple predicted configuration. Relative orbital energies and electron interactions contribute to this preference.

A completely filled d subshell is therefore associated with the observed configuration of copper, but it should not be treated as the only explanation.

Key point for exams: Copper has the ground-state configuration [Ar] 3d¹⁰ 4s¹.

3. Electronic Configuration of Molybdenum (Mo)

Molybdenum has the atomic number 42.

Its simple predicted configuration would be: [Kr] 4d⁴ 5s²

However, its actual ground-state configuration is: [Kr] 4d⁵ 5s¹

Molybdenum shows a configuration similar to chromium. One electron occupies the 5s orbital, while the 4d subshell becomes half-filled.

The observed configuration is favoured by the overall energy of the atom, including the relative energies of the 4d and 5s orbitals and electron interactions.

Key point for exams: Molybdenum has the ground-state configuration [Kr] 4d⁵ 5s¹.

4. Electronic Configuration of Silver (Ag)

Silver has the atomic number 47.

A simple predicted configuration would be: [Kr] 4d⁹ 5s²

However, its actual ground-state configuration is: [Kr] 4d¹⁰ 5s¹

Silver has a completely filled 4d subshell and one electron in the 5s orbital.

As with copper, the actual arrangement is determined by the overall energy of the atom rather than a rigid application of the Aufbau filling sequence.

Key point for exams: Silver has the ground-state configuration [Kr] 4d¹⁰ 5s¹.

5. Electronic Configuration of Palladium (Pd)

Palladium has the atomic number 46.

Its electronic configuration is particularly important because it differs from the simple configuration predicted by the usual filling sequence.

A simple prediction would give: [Kr] 4d⁸ 5s²

However, the actual ground-state configuration is: [Kr] 4d¹⁰

The 5s orbital is unoccupied in the ground-state atom, while the 4d subshell is completely filled.

Palladium is therefore an important reminder that the relative energies of the 4d and 5s orbitals can produce configurations that are not predicted by a rigid filling order.

Key point for exams: Palladium has the ground-state configuration [Kr] 4d¹⁰, with no 5s electrons.

6. Electronic Configuration of Gold (Au)

Gold has the atomic number 79.

Its actual ground-state electronic configuration is: [Xe] 4f¹⁴ 5d¹⁰ 6s¹

Gold has a completely filled 5d subshell and one electron in the 6s orbital.

Its configuration is influenced by the complex energy relationships in a heavy atom. Relativistic effects also contribute to the electronic structure of heavy elements such as gold.

This is a more advanced consideration than the simple half-filled or fully filled subshell explanation.

Key point for exams: Gold has the ground-state configuration [Xe] 4f¹⁴ 5d¹⁰ 6s¹.

7. Comparison of Important Electronic Configuration Exceptions

The following table is useful for quick revision.

Remember: The predicted configurations in this table are simplified Aufbau-based predictions used for comparison. The actual configurations are the established ground-state arrangements.

8. Why Are Half-Filled and Fully Filled Subshells Important?

Students often learn that half-filled and fully filled subshells are especially stable.

This is a useful introductory idea, but it needs a more complete explanation.

A half-filled or completely filled subshell can be associated with favourable electron arrangements. Exchange interactions and electron-electron interactions contribute to the energy of the atom.

However, the actual configuration depends on the total energy of the atom, including the relative orbital energies and electron interactions.

For example:

• Chromium has a half-filled 3d subshell in its ground-state configuration.

• Molybdenum has a half-filled 4d subshell.

• Copper and silver have completely filled d subshells.

• Palladium has a completely filled 4d subshell with an unoccupied 5s orbital.

These examples should not be reduced to a single rule that every atom will rearrange its electrons to obtain a half-filled or completely filled subshell.

9. Why Are 4s Electrons Removed Before 3d Electrons?

One of the most important concepts in transition-element chemistry is the order of electron removal during ion formation.

Students learn that the 4s orbital is generally filled before the 3d orbitals in the simplified Aufbau sequence. This can create confusion when transition-metal atoms form positive ions.

For common transition-metal cations, the 4s electrons are generally removed before the 3d electrons.

Example: Formation of iron ions

The ground-state electronic configuration of neutral iron is: Fe: [Ar] 3d⁶ 4s²

When iron forms an Fe²⁺ ion, it loses two electrons from the 4s orbital. Fe²⁺: [Ar] 3d⁶

When iron forms an Fe³⁺ ion, one additional electron is removed from the 3d subshell. Fe³⁺: [Ar] 3d⁵

Why are 4s electrons removed first?

The relative energies of the 4s and 3d orbitals change when electrons occupy them and when the atom forms an ion.

Although the 4s orbital is generally filled before 3d in the simplified Aufbau sequence, the energy relationship changes in the transition-metal atom.

During ion formation, the 4s electrons are generally removed first in common transition-metal cations.

Important: Do not apply the neutral-atom filling order directly to electron removal. Electron filling and electron removal are different processes.

10. Common Transition-Metal Ion Examples

The following examples help students practise the removal of 4s electrons before 3d electrons.

These are representative common ions. The table illustrates electron removal and does not imply that every listed oxidation state has the same stability or chemical importance.

11. Important Scientific Principles Behind Electronic Configuration Exceptions

12. Quick Revision Table

13. How to Prepare Electronic Configuration Exceptions for JEE and NEET

Use the following approach to revise this topic effectively.

1. Learn the Aufbau principle first. Understand the general orbital filling sequence before studying its exceptions.

2. Practise the three basic rules. Revise the Aufbau principle, Pauli exclusion principle and Hund's rule.

3. Memorise the six important configurations. Focus on Cr, Cu, Mo, Ag, Pd and Au.

4. Understand the scientific reason. Remember that total atomic energy and relative orbital energies matter. Do not rely only on the phrase half-filled or fully filled stability.

5. Practise ion formation. Write the configurations of neutral atoms first, then remove electrons from the appropriate orbitals.

6. Pay attention to 4s electron removal. Common transition-metal cations generally lose 4s electrons before 3d electrons.

7. Practise oxidation-state questions. Connect the number of electrons removed with the resulting d-electron configuration.

8. Make a revision chart. Keep the six exceptional configurations and common ion examples together for quick practice.

For JEE Main and NEET, focus on correct ground-state configurations, the six important examples and common ion formation. For JEE Advanced, develop a stronger understanding of orbital energy, electron interactions and the limitations of simplified filling rules.

Conclusion

Electronic configuration exceptions show that the Aufbau principle is a useful prediction method, but it does not describe every atom perfectly. The actual ground-state arrangement depends on orbital energies, electron interactions and the total energy of the atom.

Remember the configurations of chromium, copper, molybdenum, silver, palladium and gold. Also practise the removal of 4s electrons before 3d electrons when common transition-metal cations form.

Continue with the other guides in the Chemical Exceptions in Inorganic Chemistry series to understand unusual bonding behaviour, block-wise anomalies and coordination chemistry.