Electron Configuration

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Electron configuration is the arrangement of electrons in an atom's orbitals. It describes the specific distribution of electrons among the available energy levels and sublevels, following certain rules that determine how atoms behave chemically.

Understanding Atomic Orbitals

Electrons occupy regions of space called orbitals. Each orbital can hold a maximum of 2 electrons. Orbitals are organized into energy levels (shells) and sublevels (subshells):

Sublevel	Shape	Number of Orbitals	Maximum Electrons
s	Spherical	1	2
p	Dumbbell	3	6
d	Cloverleaf	5	10
f	Complex	7	14

Three Fundamental Rules

1. Aufbau Principle

Electrons fill orbitals starting from the lowest energy level to the highest. The filling order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

2. Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers. This means each orbital can hold a maximum of 2 electrons, and they must have opposite spins.

3. Hund's Rule

When electrons occupy orbitals of equal energy (degenerate orbitals), one electron enters each orbital before any orbital receives a second electron. All single electrons must have the same spin.

Energy Level Diagram (Aufbau Diagram)

The diagonal rule helps remember the filling order:

Energy Level	Sublevels Available	Order Filled
n = 1	1s	1st
n = 2	2s, 2p	2nd, 3rd
n = 3	3s, 3p, 3d	4th, 5th, 7th
n = 4	4s, 4p, 4d, 4f	6th, 8th, 10th, 13th

Writing Electron Configurations

Electron configurations are written using the notation: nlx where:

n = principal energy level (1, 2, 3, etc.)
l = sublevel (s, p, d, f)
x = number of electrons in that sublevel (superscript)

Shorthand (Noble Gas) Notation

For elements with many electrons, use the symbol of the previous noble gas in brackets, then continue with remaining electrons. For example, potassium (K): [Ar] 4s1 instead of writing 1s2 2s2 2p6 3s2 3p6 4s1.

Electron Configurations of Common Elements

Element	Atomic Number	Full Configuration	Noble Gas Notation
Hydrogen (H)	1	1s1	1s1
Carbon (C)	6	1s2 2s2 2p2	[He] 2s2 2p2
Oxygen (O)	8	1s2 2s2 2p4	[He] 2s2 2p4
Sodium (Na)	11	1s2 2s2 2p6 3s1	[Ne] 3s1
Iron (Fe)	26	1s2 2s2 2p6 3s2 3p6 4s2 3d6	[Ar] 4s2 3d6

Exceptions to the Aufbau Principle

Some elements have configurations that differ from predictions due to the extra stability of half-filled and fully-filled d sublevels. Notable exceptions include:

Chromium (Cr): [Ar] 4s1 3d5 (not 4s2 3d4)
Copper (Cu): [Ar] 4s1 3d10 (not 4s2 3d9)

SAT/ACT Connection

Science reasoning sections may ask you to interpret electron configurations to predict chemical properties. Understanding that valence electrons (outermost electrons) determine reactivity is key for answering these questions.

💡 Examples

Example 1: Writing Full Electron Configuration

Problem: Write the full electron configuration for chlorine (Cl, atomic number 17).

Solution:

Step 1: Chlorine has 17 electrons to place.

Step 2: Fill orbitals following the Aufbau principle:

1s2 (2 electrons, 15 remaining)
2s2 (2 electrons, 13 remaining)
2p6 (6 electrons, 7 remaining)
3s2 (2 electrons, 5 remaining)
3p5 (5 electrons, 0 remaining)

Answer: 1s2 2s2 2p6 3s2 3p5

Example 2: Noble Gas Notation

Problem: Write the electron configuration for calcium (Ca, atomic number 20) using noble gas notation.

Solution:

Step 1: Find the nearest noble gas with fewer electrons than calcium. Argon (Ar) has 18 electrons.

Step 2: Write [Ar] to represent the first 18 electrons.

Step 3: Calcium has 20 - 18 = 2 remaining electrons, which go into 4s.

Answer: [Ar] 4s2

Example 3: Identifying Valence Electrons

Problem: How many valence electrons does sulfur (S) have? Its configuration is 1s2 2s2 2p6 3s2 3p4.

Solution:

Step 1: Valence electrons are in the outermost energy level (highest n value).

Step 2: The highest energy level for sulfur is n = 3.

Step 3: Count electrons in the third energy level: 3s2 3p4 = 2 + 4 = 6 electrons.

Answer: Sulfur has 6 valence electrons.

Example 4: Electron Configuration from Position in Periodic Table

Problem: Without looking at a chart, determine the electron configuration of selenium (Se), which is in Period 4, Group 16.

Solution:

Step 1: Period 4 means the valence electrons are in the 4th energy level.

Step 2: Group 16 (or 6A) means 6 valence electrons.

Step 3: The noble gas before Period 4 is Argon [Ar].

Step 4: After Ar, fill 4s2, then 3d10, then 4p4 (to get 6 valence electrons total).

Answer: [Ar] 4s2 3d10 4p4 or [Ar] 3d10 4s2 4p4

Example 5: Configuration of an Ion

Problem: Write the electron configuration for Fe2+ (iron with a 2+ charge).

Solution:

Step 1: Iron (Fe) has 26 electrons. Neutral Fe: [Ar] 4s2 3d6

Step 2: Fe2+ means iron has lost 2 electrons (26 - 2 = 24 electrons).

Step 3: Important: Electrons are lost from the highest n value first. Electrons are removed from 4s before 3d.

Step 4: Remove both 4s electrons.

Answer: Fe2+: [Ar] 3d6

✏️ Practice

1. What is the electron configuration of nitrogen (N, atomic number 7)?

A) 1s2 2s2 2p2

B) 1s2 2s2 2p3

C) 1s2 2s2 2p5

D) 1s2 2s1 2p4

2. Which noble gas notation correctly represents potassium (K, atomic number 19)?

A) [Ar] 3d1

B) [Ne] 3s2 3p6 4s1

C) [Ar] 4s1

D) [Kr] 4s1

3. How many valence electrons does phosphorus (P) have?

A) 3

B) 5

C) 15

D) 10

4. According to the Aufbau principle, which sublevel is filled immediately after 3p?

A) 3d

B) 4s

C) 4p

D) 3s

5. What is the electron configuration of a Mg2+ ion?

A) 1s2 2s2 2p6 3s2

B) 1s2 2s2 2p6

C) 1s2 2s2 2p4

D) 1s2 2s2 2p6 3s2 3p2

6. Which element has the electron configuration [Ne] 3s2 3p4?

A) Silicon

B) Sulfur

C) Phosphorus

D) Chlorine

7. What is the maximum number of electrons that can occupy the 3d sublevel?

A) 2

B) 6

C) 10

D) 14

8. Which rule states that electrons fill orbitals of equal energy singly before pairing up?

A) Aufbau Principle

B) Pauli Exclusion Principle

C) Hund's Rule

D) Octet Rule

9. The actual electron configuration of chromium is [Ar] 4s1 3d5 instead of [Ar] 4s2 3d4. This exception occurs because:

A) The 4s orbital is higher in energy than 3d

B) Half-filled d sublevels are especially stable

C) Chromium violates the Pauli Exclusion Principle

D) The Aufbau principle doesn't apply to transition metals

10. An element has the configuration [Kr] 5s2 4d10 5p3. In which group of the periodic table is this element?

A) Group 3

B) Group 13

C) Group 15

D) Group 18

Click to reveal answers

B - Nitrogen has 7 electrons: 1s2 2s2 2p3
C - Potassium (19 electrons) uses Ar (18 electrons) as its core, with 1 electron in 4s
B - Phosphorus is in Group 15 and has 5 valence electrons (3s2 3p3)
B - Following the Aufbau principle diagonal rule, 4s fills before 3d
B - Mg has 12 electrons, loses 2 to become Mg2+ with 10 electrons (same as Ne)
B - Sulfur (S) has the configuration [Ne] 3s2 3p4
C - The d sublevel has 5 orbitals, each holding 2 electrons, for a maximum of 10
C - Hund's Rule describes the single filling of degenerate orbitals before pairing
B - Half-filled and fully-filled sublevels provide extra stability, causing exceptions
C - With 5 valence electrons (5s2 5p3), the element is in Group 15

✅ Check Your Understanding

Question 1: Why do electrons fill the 4s orbital before the 3d orbital, even though 4 is a higher number than 3?

Reveal Answer

The 4s orbital is actually lower in energy than the 3d orbital for most atoms. Energy levels overlap, and the filling order is determined by energy, not just the principal quantum number. The 4s orbital's energy is lower due to its greater penetration toward the nucleus. This is why we follow the Aufbau principle's specific filling order rather than simply filling by increasing n values.

Question 2: How does electron configuration explain why elements in the same group of the periodic table have similar chemical properties?

Reveal Answer

Elements in the same group have the same number of valence electrons, which are the electrons involved in chemical bonding and reactions. For example, all alkali metals (Group 1) have one valence electron in their outermost s orbital, making them all highly reactive and likely to form +1 ions. The similar valence configurations lead to similar chemical behaviors.

Question 3: When transition metal atoms form ions, why are electrons removed from the 4s orbital before the 3d orbital, even though 4s was filled first?

Reveal Answer

Although 4s fills before 3d in neutral atoms (because it has lower energy in that context), the situation changes when ionization occurs. In transition metal ions, the 3d electrons are actually lower in energy and closer to the nucleus than 4s electrons. The 4s orbital is larger and more exposed, making those electrons easier to remove. This is why Fe loses its 4s electrons first to form Fe2+ ([Ar] 3d6) rather than losing 3d electrons.

Question 4: Explain how you would determine the electron configuration of an element you've never seen before, using only its position in the periodic table.

Reveal Answer

The periodic table is organized by electron configuration. To determine configuration: (1) Find the element's period number, which tells you the highest energy level. (2) Find its group, which indicates valence electrons. (3) Identify the block (s, p, d, or f) where the element is located. (4) Start with the noble gas from the previous period. (5) Add electrons for the current period following the pattern: s-block elements fill the ns orbital, d-block fills (n-1)d, p-block fills np. For example, an element in Period 5, Group 16 would be [Kr] 5s2 4d10 5p4.

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