# A Guide to Write Electron Configuration for Every Element

Written by: Kaushik Jethva

Last Updated: May 25, 2023

Zn Electron configuration is a fundamental concept in understanding the arrangement of electrons around an atom’s nucleus. This information is crucial for comprehending an element’s chemical properties, reactivity, and bonding behaviour. In this blog post, we will dive into the process of writing electron configurations for atoms of any element. We’ll begin with a brief overview of electron behaviour principles and then dive into the step-by-step process for determining an element’s electron configuration.

## 1 Understanding the Basics

### Quantum Numbers

To understand electron configuration, it is essential to become familiar with the concept of quantum numbers. Quantum numbers are a set of four numerical values that describe the distribution and behaviour of electrons in an atom. The four quantum numbers are:

• Principal Quantum Number (n): Represents the energy level (shell) of an electron, with values ranging from 1 to 7.
• Azimuthal Quantum Number (l): Represents the subshell within the energy level, with values ranging from 0 to (n-1).
• Magnetic Quantum Number (m_l): Represents the orbital within the subshell, with values ranging from -l to +l.
• Spin Quantum Number (m_s): Represents the electron’s spin within the orbital, with values of -½ or +½.

### Pauli Exclusion Principle

The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of quantum numbers. This principle is crucial in determining electron configuration, as it helps dictate the number of electrons that can occupy a specific orbital.

### Aufbau Principle

The Aufbau Principle (German for “building up”) is a guideline for filling the orbitals in an atom, stating that electrons will fill the lowest energy orbitals first. The order of filling is determined by increasing the values of the sum (n + l).

### Hund's Rule

Hund’s Rule states that when electrons fill degenerate orbitals (orbitals with the same energy), they will occupy separate orbitals with parallel spins until all the orbitals are half-filled. This rule helps maximise the stability of an atom’s electron configuration.

## 2 Writing Electron Configurations

### Step 1: Identify the Element's Atomic Number

The atomic number (Z) of an element corresponds to the number of protons in the nucleus and the number of electrons orbiting it. This number is essential for determining the electron configuration.

### Step 2: Determine the Order of Orbital Filling

Use the Aufbau Principle to determine the order in which orbitals are filled. A mnemonic device known as the “diagonal rule” can be helpful:
1s -> 2s -> 2p -> 3s -> 3p -> 4s -> 3d -> 4p -> 5s -> 4d -> 5p -> 6s -> 4f -> 5d -> 6p -> 7s -> 5f -> 6d -> 7p

### Step 3: Distribute Electrons

Starting with the lowest energy orbital, add electrons until the atomic number is reached. Remember to follow Hund’s Rule when filling degenerate orbitals.

### Step 4: Write the Configuration

Express the electron configuration using the following notation: [energy level] [subshell] [number of electrons]. For example, the configuration for hydrogen would be written as 1s¹.

## 3 Example 1: Writing the Electron Configuration for Oxygen (Z = 8)

1. Identify the Atomic Number: Oxygen has an atomic number of 8.
2. Determine the Order of Orbital Filling: Following the diagonal rule, the order is
1s -> 2s -> 2p -> 3s -> 3p -> 4s -> 3d -> 4p.
3. Distribute Electrons: Allocate the 8 electrons of oxygen across the orbitals.
• 1s orbital can hold a maximum of 2 electrons: 1s²
• 2s orbital can hold a maximum of 2 electrons: 2s²
• 2p orbital can hold a maximum of 6 electrons: 2p⁴ (since only 4 electrons remain)
4. Write the Configuration: Combining the results from step 3, the electron configuration for oxygen is written as 1s² 2s² 2p⁴.

## 4 Example 2: Writing the Electron Configuration for Zinc (Z = 30)

1. Identify the Atomic Number: Zinc has an atomic number of 30.
2. Determine the Order of Orbital Filling: Following the diagonal rule, the order is 1s -> 2s -> 2p -> 3s -> 3p -> 4s -> 3d -> 4p -> 5s -> 4d -> 5p -> 6s.
3. Distribute Electrons: Allocate the 30 electrons of zinc across the orbitals.
• 1s orbital can hold a maximum of 2 electrons: 1s²
• 2s orbital can hold a maximum of 2 electrons: 2s²
• 2p orbital can hold a maximum of 6 electrons: 2p⁶
• 3s orbital can hold a maximum of 2 electrons: 3s²
• 3p orbital can hold a maximum of 6 electrons: 3p⁶
• 4s orbital can hold a maximum of 2 electrons: 4s²
• 3d orbital can hold a maximum of 10 electrons: 3d¹⁰ (since only 10 electrons remain)
4. Write the Configuration: Combining the results from step 3, the electron configuration for zinc is written as 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰.

## 5 Utilising the Noble Gas Shorthand

For elements with a large atomic number, electron configurations can become lengthy. To simplify these configurations, the noble gas shorthand method is used. This method involves representing the electron configuration of the previous noble gas in square brackets, followed by the remaining electron configuration. For example, the electron configuration for chlorine (Z = 17) is:

1s² 2s² 2p⁶ 3s² 3p⁵

Since the electron configuration of the noble gas neon (Z = 10) is 1s² 2s² 2p⁶, we can rewrite chlorine’s configuration as:

[Ne] 3s² 3p⁵

## 6 Wrapping Up

Understanding electron configurations is fundamental to the study of chemistry, as it provides insight into the behaviour and properties of elements. By mastering the principles and rules outlined in this blog post, you can determine the electron configuration for any element on the periodic table. Remember to consider quantum numbers, the Pauli Exclusion Principle, the Aufbau Principle, and Hund’s Rule when building configurations, and don’t forget the noble gas shorthand for simplifying larger atomic numbers.

Also, do not forget to check our blog on how to memorise trigonometric functions.

## Community Q&A 