Electron Configuration

How to Master Electron Configuration and Finally Understand the Periodic Table

Why do some elements form bonds instantly, while others barely react at all?

Is it random… or is there a pattern behind it?

The answer lies in electron configuration—the way electrons are arranged inside an atom.

Think of it this way: if atoms were people, electrons would determine their personality. How they interact, what they “prefer,” and how they react all depend on where their electrons are.

But here’s the problem: Electron configuration can feel overwhelming.

What does 1s²2s²2p⁶ even mean?

Why does 4s fill before 3d?

Why don’t electrons just pair up immediately?

If you’ve ever felt like you’re just memorizing rules without understanding them—you’re not alone. This guide will help change that.

By the end, electron configuration won’t feel like memorization—it’ll feel like a system you can actually understand.

What You’ll Learn:

• How electron configurations actually work (without the confusion)
• The simple rules that govern electron placement
• How to use the periodic table as a shortcut
• Why electron arrangement explains chemical behavior

Key Takeaways

• Electron configuration determines chemical behavior, explaining how atoms bond, react, and why periodic trends exist.
• Electrons fill orbitals following three key rules: Pauli Exclusion Principle (max 2 per orbital), Aufbau Principle (lowest energy first), and Hund’s Rule (singly with parallel spins before pairing).
• Electron configurations follow a predictable filling order based on the Aufbau Principle, with fixed subshell capacities (s = 2, p = 6, d = 10, f = 14).
• The periodic table mirrors electron configuration patterns, with s, p, d, and f blocks showing which subshell is being filled.
• Valence electrons are key to reactivity, while concepts like noble gas notation and exceptions (e.g., Cr, Cu) refine how configurations are written and understood.

Electron Configuration

When we talk about electron configuration, we are describing how electrons are arranged within an atom—specifically, how they occupy different energy levels, subshells, and orbitals. But why should we care about this arrangement?

The answer lies in chemical behavior. The way electrons are distributed determines how atoms interact, bond, and react. In fact, one of the central ideas in chemistry is that elements with similar electron configurations exhibit similar chemical properties. This is why the periodic table shows predictable patterns: it reflects repeating patterns in electron arrangement.

Electron Configuration Notation

The electron configuration notation or spdf notation shows the energy level (n), the subshell (s, p, d, f), and the number of electrons (as a superscript).

For example, a carbon atom has 6 electrons and the electron configuration 1s²2s²2p². This means that carbon has 2 electrons in the 1s subshell, another 2 in the 2s subshell, and another 2s in the 2p subshell.

As such, this compact notation answers the question: Where are the electrons located in the atom?

Orbital Diagrams

Electron configurations can also be shown through orbital diagrams, also known as box notation. This uses boxes or lines to represent each orbital and arrows to show electrons. The opposite spins of the electrons are shown through arrows pointing up and pointing down.

Recall that the magnetic quantum number m_l describes orbital orientation with values ranging from –l to l. Each box in an orbital diagram corresponds to an m_l value from this range.

Magnetic quantum number (m_l) in orbital diagrams. Boxes are often labeled from –l to +l, but this is purely representational. The positions don’t matter—what matters is that each orbital in a subshell has a different m_l value.

The corresponding orbital diagram of the electron configuration of carbon is:

You might now be wondering how and why carbon has the given electron configuration. Let’s clear that up and answer the question: how do we determine the electron configuration?

Rules for Electron Arrangement

Before assigning electrons to orbitals, we need to understand the rules they follow. These rules arise from quantum mechanics and ensure that electrons occupy the lowest energy arrangement possible.

The Pauli Exclusion Principle: Why Can Only Two Electrons Fit in an Orbital?

If we imagine placing electrons into orbitals, a natural question arises: Can we keep adding electrons indefinitely into the same orbital?

The answer is no. According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers. In simpler terms, an orbital can hold a maximum of two electrons, and these must have opposite spins.

For example, helium has two electrons in the same 1s orbital. One electron is assigned a spin of +½, and the other −½. This pairing ensures that each electron remains unique in its quantum state.

This principle has a powerful consequence: it sets strict limits on how many electrons each subshell can hold. Since:

• an s subshell has 1 orbital → it holds 2 electrons
• a p subshell has 3 orbitals → it holds 6 electrons
• a d subshell has 5 orbitals → it holds 10 electrons
• an f subshell has 7 orbitals → it holds 14 electrons

Additionally, since there are n² orbitals in a given shell n, the maximum number of electrons in any shell is 2n².

The Aufbau Principle: Which Orbitals are Filled First?

Take a look back at the electron configuration of carbon. The 1s orbital is filled first, followed by the 2s orbital. The 2p orbitals are filled last.

This might be making you wonder: How do we systematically assign electrons to orbitals?

The answer lies in the Aufbau principle, which means “building up.” Electrons are added one by one to the lowest-energy orbitals available, ensuring that the atom remains in its lowest energy, known as the ground state.

In other words, sublevels fill in order of increasing energy, which follows the rule:

• Orbitals with lower n + l values fill first
• If two orbitals have the same value, the one with lower n fills first

Aufbau principle. The order of filling in orbitals based on the Aufbau principle is summarized in the given diagram: the 1s subshell is filled first, followed by 2s, then 2p, 3s, 3p, 4s, and so on as indicated by the arrows.

This explains patterns such as:

• 2s filling before 2p
• 4s filling before 3d

Hund’s Rule: Why Don’t Electrons Pair Immediately?

Looking back at the orbital diagram of carbon, you might notice that the 1s and 2s orbitals are both filled, but the two electrons in the 2p subshell are placed in different orbitals (boxes).

This naturally leads to the question: If orbitals can hold two electrons, why don’t electrons just pair up right away?

The answer is given by Hund’s rule, which states that electrons occupy orbitals of equal energy singly first, with parallel spins, before pairing occurs.

This happens because electrons repel each other. By spreading out into separate orbitals, they minimize repulsion and lower the atom’s energy.

This is why the two electrons in the 2p subshell of carbon occupy two different orbitals rather than pairing in one. Only after each orbital has one electron does pairing begin.

Hence, orbital diagrams are especially useful for visualizing electron pairing and Hund’s rule in action. The boxes are typically filled with the upward arrow first, which is typically attributed to the +½ spin value, but these are only representational values. As long as paired electrons are shown to have opposite spins and unpaired electrons in the same sublevel show parallel spins, the orbital diagram is technically correct.

Hund’s rule and Pauli exclusion principle. Electrons fill orbitals singly with parallel spins before pairing (Hund’s rule). In a p subshell, the first three electrons occupy separate orbitals; the fourth begins pairing. Each orbital holds at most two electrons with opposite spins (Pauli exclusion principle).

The following question might then be helpful when thinking about electron placement: Is there an empty orbital available before pairing electrons?

Patterns in Electron Configurations Across Periods

Let’s apply all three rules to determine the electron arrangement of the periods and sections in the periodic table. As we go along, we will be observing electron configuration patterns in the periodic table.

Period 1

The only elements in Period 1 are hydrogen and helium. Since hydrogen only has 1 electron, it has the following electron configuration and orbital diagram:

Consequently, this electron has the following quantum numbers: n = 1, l = 0, m_l = 0, and m_s = +½ or –½.

Helium has 2 electrons. One electron has the same set of quantum numbers as the lone electron of hydrogen, but the second electron in the orbital should have an opposite spin, following the Pauli exclusion principle. This gives the following electron configuration and orbital diagram:

Hence, the Period 1 elements only have electrons in the 1s orbital.

Period 2

By the end of Period 1, the 1s orbital is fully filled, so in Period 2 electrons begin filling the 2s and then the 2p subshells, following the Aufbau principle.

The first element in Period 2 is lithium, which has 3 electrons. The third electron in lithium is the first electron in the 2s orbital, and thus have the following quantum numbers: n = 2, l = 0, m_l = 0, and m_s = +½ or –½. Its electron arrangement is shown below:

Following lithium, beryllium has 4 electrons. With an opposite spin, the fourth electron pairs up with the electron in the 2s orbital:

Boron, with 5 electrons, is the next element. The 2s orbital is completely filled now, so the fifth electron becomes the first to occupy the 2p subshell. Since all three orbitals of the 2p subshell have equal energies, the fifth electron can go into any of the orbitals although orbital filling is conventionally done from left to right. This means the fifth electron of boron would have the following quantum numbers: n = 2, l = 1, m_l = –1, 0, 1, and m_s = +½ or –½.

Carbon then follows. It has 6 electrons whose arrangement has previously been explained. The n and l values of the last electron of carbon are the same as that of the fifth electron of boron. The m_s value is also the same since the two unpaired 2p electrons of carbon have parallel spins.

What differs is the m_l value, which can be either of the values that is not the m_l of the fifth electron of boron. For instance, if the fifth electron of boron has an m_l = –1, then the last electron of carbon would have m_l = 0 or 1.

The next element is nitrogen, which has 7 electrons. The seventh electron of nitrogen goes into the remaining unoccupied 2p orbital, following Hund’s rule. This electron would have the same n, l, and m_s values as the electrons in the other 2p orbitals, but its m_l would be whichever is not yet taken by the other 2p electrons.

Next is oxygen, which has 8 electrons. The eighth electron of oxygen would be the first to pair up with one of the electrons in the 2p sublevel. Again, pairing up is conventionally done from left to right although the electron can be paired with any of the unpaired electrons.

Fluorine and neon, which have 9 and 10 electrons respectively, are next. Their last electrons pair up with the remaining unpaired 2p electrons as shown below:

By the end of Period 2, all the orbitals with n = 2 have been completely filled.

Period 3 and Other Notations

Period 3 follows the same logic as Period 2, but with the 3s and 3p subshells. This repetition reinforces the idea that electron configuration patterns are periodic, which is one of the foundations of the periodic table itself.

However, writing entire electron configurations and drawing entire orbital diagrams can become too lengthy at this point, so they are typically written more concisely using two ways: partial orbital diagrams and noble gas notation, also known as the condensed electron configuration.

Partial orbital diagrams draw only the sublevels being filled, while the noble gas notation puts the element symbol of the previous noble gas in brackets and writes out the remainder of the electron configuration. Notice that noble gases like helium and neon have all of their orbitals completely filled, so the noble gas in brackets is a shorthand representing the entire electron configuration of that noble gas.

For Period 3 elements, the preceding noble gas is neon, which has 10 electrons and the electron configuration 1s²2s²2p⁶. As such, the noble gas notation of silicon (14 electrons), for instance, is [Ne] 3s²2p², with [Ne] representing 1s²2s²2p⁶.

Period 4 and Transition Elements

By the end of Period 3, the 3s and 3p orbitals are completely occupied. Looking back at the Aufbau principle, the next lowest-energy orbital is 4s, followed by 3d and 4p respectively.

Shielding and Penetration

You might be wondering: what makes the 4s orbital lower in energy than the 3d orbitals? In other words, why is orbital energy determined by n + l?

In atoms with many electrons, energy levels are affected not only by distance from the nucleus but also by electron repulsion and orbital shape. Two simple ideas help explain this:

• Shielding: Inner electrons block some of the nucleus’s pull from reaching outer electrons. This makes outer electrons easier to remove and higher in energy. This lowers the effective nuclear charge, which is the attraction 
• Penetration: Due to their shape, some orbitals allow electrons to get closer to the nucleus than others. The closer an electron gets, the stronger the attraction, and the lower its energy.

A simple pattern emerges: s orbitals penetrate closest, then p, then d, then f. So for the same energy level n, the sublevel energy order is s < p < d < f.

Now we can put everything together. Even though the 4s orbital is in a higher energy level (n = 4), its electrons can get closer to the nucleus than electrons in the 3d orbital. Because of this:

• 4s electrons feel a stronger attraction to the nucleus
• They are less affected by shielding

In contrast, 3d electrons:

• Stay farther from the nucleus
• Are more shielded by inner electrons

As a result, the 4s orbital is actually slightly lower in energy than 3d, so it fills first.

Exceptions in Electron Configuration

Take a look at the electron configurations and orbital diagrams of chromium and copper. You would notice something unusual:

• Chromium has the electron configuration [Ar] 4s¹3d⁵ instead of [Ar] 4s²3d⁴. What happens here is that one electron “moves” from 4s to 3d, giving both a half-filled status.
• Copper has the electron configuration [Ar] 4s¹3d¹⁰ instead of [Ar]4s²3d⁹. One electron “moves” from 4s to 3d, giving a completely filled 3d subshell.

Why do these exceptions in electron configuration occur? These exceptions arise because half-filled and fully filled subshells are especially stable, which lowers the overall energy of the atom.

As the 3d sublevel begins to fill, its energy becomes very close to that of the 4s sublevel. When sublevel energies are this close—especially in larger atoms—an electron may shift from one subshell to another to achieve a more stable arrangement. This behavior is also observed in elements such as molybdenum (Mo) and silver (Ag), among other transition metals.

This leads to an important insight: Electron configurations are guided by energy minimization, not strict rules alone.

Periods 6 and 7: Lanthanides and Actinides

Periods 6 and 7 include the lanthanides (Ce-Lu) and actinides (Th-Lr), also known as the inner transition metals. These fill in the f orbitals. Specifically, the lanthanides occupy the 4f orbitals while actinides occupy the 5f, so both have l = 3 and m_l = –3, –2, –1, 0, +1, +2, or +3.

Types of Electrons

Not all electrons play the same role in chemistry, so it is useful to distinguish between them. Chemists distinguish three categories based on where electrons are and what they do:

Core (inner) electrons are found in filled inner shells. This includes electrons shared with the preceding noble gas, as well as electrons in completed transition series. These are not involved in chemical reactions.

Outer electrons are those in the highest energy level (n). These spend the most time farthest from the nucleus. For main-group elements, the group number indicates the number of outer electrons: Group 1 (1a) elements have 1 outer electron, Group 2 (2a) elements have 2 outer electrons, Group 13 (3a) elements have 3, Group 14 (4a) elements have 4, and so on.

• Valence electrons are the ones involved in chemical bonding. They are especially important as they determine how atoms bond, how reactive they are, and what compounds they form. 

For main-group elements, outer electrons and valence electrons are the same, so the group number also indicates the number of valence electrons. However, for transition metals, d electrons can also act as valence electrons, which adds complexity to their chemistry.

Determining Electron Configuration from the Periodic Table

You might be thinking that it’s difficult to determine and memorize the order of orbital filling, but here’s the beautiful payoff: the periodic table is essentially an electron configuration chart in disguise.

The period number indicates the highest energy level (or shell number, n) occupied by an atom. In addition, the periodic table can be divided into four blocks, based on which type of sublevel is being filled:

• s-block: Includes Groups 1-2 and helium. These elements fill the ns orbitals, and their period number (except for helium) corresponds to the n value of the outermost s sublevel. Helium is placed in Group 18 due to its filled 1s orbital, even though it is technically part of the s-block.
• p-block: Includes Groups 13-18 (except helium). These elements fill the np orbitals, and their period number gives the n value of the outermost p sublevel. Their valence electron configuration follows the pattern ns²np^x, where x = 1 to 6 across the block.
• d-block: Composed of the transition metals. These elements fill the (n − 1)d orbitals, meaning the d sublevel being filled belongs to one energy level lower than the period number.
• f-block: Composed of the inner transition elements. These elements fill the (n − 2)f orbitals, meaning the f sublevel being filled is two energy levels lower than the period number.

Orbital map of the periodic table. The periodic table is organized into s, p, d, and f blocks, which follow the order of electron filling. The energy level or shell (n) being filled can also be determined from the period number.

Because electron configurations repeat periodically, chemical properties also repeat periodically, especially for main-group elements. As a result, main-group elements in the same group have the same number of valence electrons, which explains why they behave similarly in chemical reactions. Let’s take a look at some examples:

• Group 1 elements (ns¹): Highly reactive metals that tend to lose one electron in reactions.
• Group 17 elements (ns²np⁵): Reactive nonmetals that tend to gain one electron.
• Group 18 elements (ns²np⁶): With filled energy levels, these elements are very stable and largely unreactive.

Let’s help you remember the key topics with these memory tricks or mnemonics

Why Electron Configuration Matters

“Configuration = Chemistry”

• How electrons are arranged determines how atoms behave.

“Same Setup → Same Behavior”

• Elements with similar configurations act alike

What the notation means (e.g., 2p²)

“State–City–Population”

• State → energy level/shell (n)
• City (in the state) → subshell (s, p, d, f) in the shell
• Population (of the city) → electrons

Pauli Exclusion Principle

“2 per room, opposite moods”

• Max 2 electrons per orbital (“room”)
• Opposite spins

Aufbau Principle

“Build from the basement up”

• Follow the electron configuration diagram or use the periodic table

Hund’s Rule

“Spread out before the pair up”

• Each orbital must be filled with electrons of parallel spins before pairing up occurs

Why ns is filled before (n – 1)d

“Closer beats lower”

• 4s penetrates closer to the nucleus → lower energy

Shielding

“Inner blocks, outer feels less”

• Inner (core) electrons reduce the attraction between the nucleus and outer electrons

Penetration order

“s, p, d, f → s best, f worst”

• The s orbitals can penetrate closest to the nucleus, followed by p, d, and f respectively

Types of Electrons

• “Core = Ignore” – core electrons don’t typically participate in chemistry
• “Outer = Farthest” – electrons in the highest n
• “Valence = Violence” – valence electrons are the ones that react and bond

Periodic Table as a Map

“The table IS the configuration”

• s-block: “2 groups wide” → s orbital = max 2 electrons
• p-block: “6 groups wide” → p orbital = max 6 electrons
• d-block: “10 groups wide (n – 1)” → d orbital = max 10 electrons
• f-block: “14 groups wide (n – 2)” → f orbital = max 14 electrons

Conclusion

Electron configuration is more than just a notation system—it is the bridge between quantum mechanics and observable chemical behavior.

By applying a few key principles—the Aufbau principle, the Pauli exclusion principle, and Hund’s rule—we can predict how electrons arrange themselves within an atom. From this arrangement, we can make sense of periodic trends, bonding behavior, and overall chemical reactivity.

Instead of relying on memorization, you now have a framework that explains why elements behave the way they do.

Ultimately, the guiding idea is simple but powerful: the structure of the periodic table—and the behavior of every element within it—is a direct consequence of how electrons are arranged in atoms.