Electron Configuration and Energy Level Filling Explained (OCR A-Level Chemistry)
- 4 days ago
- 8 min read
Updated: 3 days ago
By Paul Morgan
Understanding Electron Configuration
One of the most common areas of confusion in OCR A-Level Chemistry is electron configuration. Many students can remember the order in which orbitals fill but become unsure when questions involve transition metals, ions or the exceptions to the usual rules.
Key Concepts
Order of Orbital Filling
Transition Metals
Ions and Their Configurations
Exceptions to the Rules
Exam questions often test much more than simply writing an electron configuration. You may be asked why the 4s orbital fills before the 3d orbital, why 4s electrons are removed first during ionisation, or why chromium and copper don't follow the expected pattern.
Fortunately, there are only a handful of rules you need to understand. Once you know them, electron configuration questions become much more straightforward.
In this guide you'll learn:
• What electron configuration means.
• The Aufbau Principle.
• Hund's Rule.
• The Pauli Exclusion Principle.
• Why the 4s orbital fills before the 3d orbital.
• Why 4s electrons are removed before 3d electrons.
• The chromium and copper exceptions.
• OCR-style worked examples.
• Common exam mistakes and how to avoid them.
What Is Electron Configuration?
Electron configuration describes how electrons are arranged in the orbitals around the nucleus of an atom. Electrons always occupy the lowest available energy orbitals first. As more electrons are added, they gradually fill orbitals with higher energy.
Understanding electron configurations helps explain many important topics in OCR A-Level Chemistry, including:
• Ion formation.
• Transition metals.
• Oxidation states.
• Periodicity.
• Chemical bonding.
Rather than memorising long strings of numbers and letters, try to understand why electrons fill orbitals in a particular order.
Once you understand the pattern, writing electron configurations becomes much easier.
The Aufbau Principle
The word Aufbau comes from the German word meaning "building up."
The Aufbau Principle states:
Electrons fill the lowest energy orbital available before occupying higher energy orbitals.
For the first twenty elements, the order is:
1s ↑↓
2s ↑↓
2p ↑↓ ↑↓ ↑↓
3s ↑↓
3p ↑↓ ↑↓ ↑↓
4s ↑↓
3d ↑
Only after the 4s orbital has filled do electrons begin to occupy the 3d orbitals.
This often surprises students because the orbital is called 4s, yet it fills before 3d.
We'll explain exactly why later in this guide.
Energy Level Filling Order
The order of increasing energy is:
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
4f
5d
6p
7s
You do not need to memorise the entire sequence for OCR A-Level Chemistry, but you should be confident with the orbitals up to krypton.
Worked Example 1
Write the electron configuration for calcium (atomic number 20).
Calcium has 20 electrons.
Fill the orbitals in order of increasing energy.
1s²
2s²
2p⁶
3s²
3p⁶
4s²
Therefore, the electron configuration of calcium is:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Worked Example 2
Write the electron configuration for scandium (atomic number 21).
Scandium has 21 electrons.
The first twenty electrons fill exactly as they do for calcium.
The twenty-first electron enters the 3d orbital.
Therefore:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
This is the first transition metal in the periodic table.
Hund's Rule
Hund's Rule explains how electrons occupy orbitals that have the same energy.
For example, the three 2p orbitals all have identical energy.
Hund's Rule states:
Electrons occupy empty orbitals singly before pairing begins.
Imagine three empty seats on a bus.
Passengers will usually choose an empty seat before sitting next to someone else.
Electrons behave in exactly the same way.
For example, nitrogen has three electrons in the 2p sub-level.
Instead of pairing immediately, they occupy separate orbitals:
Nitrogen (2p³)
2p
↑ ↑ ↑
Oxygen (2p⁴)
2p
↑↓ ↑ ↑
Only when each orbital contains one electron do additional electrons begin to pair.
Remember:
• Electrons spread out before pairing.
• This arrangement is more stable because electron repulsion is reduced.
The Pauli Exclusion Principle
The Pauli Exclusion Principle states:
An orbital can contain a maximum of two electrons.
If two electrons occupy the same orbital, they must have opposite spins.
This is usually represented as:
↑↓
Never draw two electrons with the same spin in one orbital.
Doing so breaks the Pauli Exclusion Principle and would lose marks in OCR examinations.
Why Does the 4s Orbital Fill Before the 3d Orbital?
One of the most common questions students ask is:
"If the 3d orbital comes before 4p, why does the 4s orbital fill first?"
The answer lies in the energy of the orbitals.
When electrons are first added to an atom, the 4s orbital is slightly lower in energy than the 3d orbitals.
Electrons always occupy the lowest energy orbital available.
That is why the filling order is:
1s
2s
2p
3s
3p
4s
3d
Although the 4s orbital has a higher principal energy level (4), it is actually slightly lower in energy than the empty 3d orbitals.
This is why calcium has the electron configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
rather than:
1s² 2s² 2p⁶ 3s² 3p⁶ 3d²
Why Are 4s Electrons Removed First?
This seems to contradict everything you've just learned.
If the 4s orbital fills before the 3d orbital, why are the 4s electrons removed first when transition metals form ions?
The answer is that once the 3d orbitals begin to fill, their energy changes.
The occupied 3d orbitals become slightly lower in energy than the 4s orbital.
As a result, the 4s electrons become the highest-energy electrons and are removed first during ionisation. This is an important OCR examination point.
Worked Example 3
Write the electron configuration for the calcium ion, Ca²⁺.
A calcium atom has:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Calcium loses its two outer 4s electrons.
Therefore:
Ca²⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶
Notice that the 4s orbital is now empty.
Worked Example 4
Write the electron configuration for the scandium ion, Sc³⁺.
A scandium atom has:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
During ionisation:
• The two 4s electrons are removed first.
• The remaining electron is removed from the 3d orbital.
The electron configuration becomes:
1s² 2s² 2p⁶ 3s² 3p⁶
Sc³⁺ has the same electron configuration as argon.
A Quick Shortcut for Ions
Many students find it easier to use a simple shortcut.
After ionisation, count how many electrons remain.
Then write the electron configuration for that number of electrons.
For example:
Calcium has 20 electrons.
Ca²⁺ has 18 electrons.
Simply write the electron configuration for 18 electrons:
1s² 2s² 2p⁶ 3s² 3p⁶
The answer is exactly the same.
OCR Exam Tip
There are two correct ways to write electron configurations for ions.
• Write the neutral atom first and remove electrons from the 4s orbital before the 3d orbital.
• Alternatively, count the number of electrons remaining after ionisation and write the electron configuration for the ion directly. In this method, the occupied 3d orbitals are written before the empty 4s orbital.
Both approaches give the same final electron configuration.
The Chromium Exception
Most elements follow the expected filling order.
Chromium does not.
Based on the Aufbau Principle, you might expect chromium to be:
[Ar] 4s² 3d⁴
However, its actual electron configuration is:
[Ar] 4s¹ 3d⁵
Why?
A half-filled 3d sub-level is more stable than a partially filled one.
One electron moves from the 4s orbital into the 3d orbital, producing five unpaired electrons.
You do not need to explain the detailed quantum mechanics in OCR A-Level Chemistry.
Simply remember that chromium is an exception.
The Copper Exception
Copper behaves similarly.
The expected configuration would be:
[Ar] 4s² 3d⁹
Instead, copper is:
[Ar] 4s¹ 3d¹⁰
A completely filled 3d sub-level is particularly stable.
Again, one electron moves from the 4s orbital into the 3d orbital.
OCR Exam Tip
The two transition metals you should always remember are:
• Chromium
[Ar] 4s¹ 3d⁵
• Copper
[Ar] 4s¹ 3d¹⁰
These are by far the most commonly tested exceptions.
Electron Configurations of Negative Ions
Negative ions are formed when atoms gain electrons.
Unlike positive ions, no electrons are removed, so there is no need to think about the 4s and 3d rule. Simply add the extra electrons to the next available orbital following the normal filling order.
Example: Chloride ion (Cl⁻)
Chlorine atom:
1s² 2s² 2p⁶ 3s² 3p⁵
Add one electron:
Cl⁻ = 1s² 2s² 2p⁶ 3s² 3p⁶
Example: Oxide ion (O²⁻)
Oxygen atom:
1s² 2s² 2p⁴
Add two electrons:
O²⁻ = 1s² 2s² 2p⁶
OCR Exam Tip
For negative ions, simply add electrons to the next available orbital. The rule that 4s electrons are removed before 3d electrons applies only to positive ions of transition metals.
Common OCR Exam Mistakes
Mistake 1
Writing the 3d orbitals before the 4s orbital when filling electrons.
Remember:
Electrons fill 4s before 3d.
Mistake 2
Removing electrons from the 3d orbital before the 4s orbital.
When transition metals form ions:
4s electrons are removed first.
Mistake 3
Forgetting the chromium and copper exceptions.
Learn these two electron configurations separately.
Mistake 4
Pairing electrons too early.
Apply Hund's Rule.
Fill empty orbitals singly before pairing begins.
OCR Exam Tip
Whenever you're asked to write an electron configuration:
• Count the total number of electrons.
• Fill orbitals in the correct order.
• Apply Hund's Rule.
• Check whether the element is chromium or copper.
• If writing an ion, remove electrons from the 4s orbital before the 3d orbital.
Following exactly the same method every time will prevent most common errors.
Practice Question
Write the electron configuration for each of the following:
• Potassium (K)
• Iron (Fe)
• Fe²⁺
• Fe³⁺
Pause before looking at the answers.
Answers
Potassium
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
Iron
[Ar] 4s² 3d⁶
Fe²⁺
Remove the two 4s electrons first.
[Ar] 3d⁶
Fe³⁺
Remove one additional 3d electron.
[Ar] 3d⁵
Students lose lots of marks on Fe²⁺ and Fe³⁺.
OCR love these questions.
Frequently Asked Questions
Why does 4s fill before 3d?
Because the empty 4s orbital is slightly lower in energy than the empty 3d orbitals.
Why are 4s electrons removed first?
Once the 3d orbitals contain electrons, they become lower in energy than the 4s orbital.
The 4s electrons are therefore removed first.
Do chromium and copper always appear in OCR examinations?
Not every year, but they are among the most common electron configuration questions.
They are well worth learning.
Do I need to memorise every orbital up to 7p?
No.
You should be confident with the filling order needed for the OCR A-Level Chemistry specification, particularly up to krypton and the first-row transition metals.
Key Takeaways
Remember these rules:
• Electrons occupy the lowest energy orbitals first.
• The 4s orbital fills before the 3d orbital.
• Hund's Rule states that electrons occupy empty orbitals singly before pairing.
• The Pauli Exclusion Principle allows a maximum of two electrons per orbital with opposite spins.
• During ionisation, 4s electrons are removed before 3d electrons.
• Chromium and copper are important exceptions to the usual filling order.
Understanding these rules will make electron configuration questions much easier and help you avoid some of the most common mistakes in OCR A-Level Chemistry.
Free OCR A-Level Chemistry Guides
If you're studying OCR A-Level Chemistry and want to avoid the most common mistakes that hold students back, download my free guides:
Year 12 Students:
4 Mistakes That Cause Strong GCSE Students To Struggle In Year 12 Chemistry
DOWNLOAD YEAR 12 GUIDE
Year 13 Students:
4 Mistakes Keeping Capable OCR Chemistry Students Stuck At Grade B Or Below
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Continue Your OCR Chemistry Revision
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Understanding Relative Atomic Mass
Limiting Reagents Explained
Electron Configuration and Energy Level Filling
How to Balance Redox Equations
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