Orbitals and Sublevels (HSC SSCE Chemistry): Revision Notes

Orbitals and Sublevels

Introduction to electron orbitals

You may have learned a simple way to describe electron arrangement in atoms based on Bohr's theory, where electrons occupy specific energy levels or "shells". However, this simple model cannot explain some important observations. For example, why does an energy level sometimes get "half-filled" with eight electrons before the next level starts filling? Why are the noble gases so stable when their outer level has only eight electrons?

Scientists developed a more detailed picture of electron arrangement that explains these puzzling observations. This advanced model introduces the concept of orbitals and sublevels.

What is an orbital?

An orbital is a three-dimensional region of space around the nucleus where we are most likely to find one or two electrons moving randomly. Think of it as a "cloud" or zone where electrons spend most of their time.

Key points about orbitals:

• Each orbital can hold a maximum of two electrons
• Orbitals have different shapes and sizes
• Different energy levels contain different numbers and types of orbitals
• Orbitals in higher energy levels are generally larger than those in lower levels
• Orbitals can overlap with each other in space

Types and shapes of orbitals

There are four main types of orbitals, each with a characteristic shape:

s orbitals

s orbitals are spherical in shape, like a ball centred on the nucleus. Different s orbitals (1s, 2s, 3s, etc.) have the same shape but different sizes. The higher the number, the larger the sphere.

The diagram above shows how the 1s orbital (smaller sphere) compares with the 2s orbital (larger sphere). Both are perfectly spherical but the 2s orbital extends further from the nucleus.

p orbitals

p orbitals have a distinctive dumbbell shape, like two teardrops or pears joined at their narrow ends. Each p sublevel contains three p orbitals oriented at right angles to each other along the x, y, and z axes.

The figure shows the three p orbitals and how they are arranged perpendicular to one another in three-dimensional space. This arrangement ensures the orbitals cover space efficiently around the nucleus.

d orbitals

d orbitals have more complex shapes than s or p orbitals. Each d sublevel contains five d orbitals, meaning it can hold up to 10 electrons in total.

f orbitals

f orbitals have even more complicated shapes. Each f sublevel contains seven f orbitals, allowing a maximum of 14 electrons.

Energy levels and sublevels

The main energy levels (often called shells) that you learned about in simpler models are actually made up of smaller groups called sublevels or subshells.

Structure of energy levels

Each main energy level contains one or more sublevels:

First energy level:

• Contains only 1 sublevel: the 1s sublevel
• Maximum capacity: 2 electrons

Second energy level:

• Contains 2 sublevels: 2s and 2p
• Maximum capacity: 8 electrons (2 in 2s, 6 in 2p)

Third energy level:

• Contains 3 sublevels: 3s, 3p, and 3d
• Maximum capacity: 18 electrons (2 in 3s, 6 in 3p, 10 in 3d)

Fourth energy level:

• Contains 4 sublevels: 4s, 4p, 4d, and 4f
• Maximum capacity: 32 electrons (2 in 4s, 6 in 4p, 10 in 4d, 14 in 4f)

Notice the pattern: the first level has 1 sublevel, the second has 2 sublevels, the third has 3 sublevels, and so on.

This comprehensive diagram shows three related views: (a) the main energy levels with their total electron capacity, (b) how these levels split into sublevels, and (c) the individual orbitals within each sublevel. Energy increases as you move upward in the diagram.

Electron capacity of sublevels

Each type of sublevel can hold a specific maximum number of electrons:

• s sublevel: 1 orbital → maximum 2 electrons
• p sublevel: 3 orbitals → maximum 6 electrons
• d sublevel: 5 orbitals → maximum 10 electrons
• f sublevel: 7 orbitals → maximum 14 electrons

To remember the capacity: 2, 6, 10, 14 (each doubles the previous increase).

The table summarises how the sublevels within each main energy level add up to give the total electron capacity. This explains why the main energy levels can hold 2, 8, 18, and 32 electrons respectively.

Important energy consideration

Orbital notation

There is a specific way to write electron configurations that shows exactly which sublevels contain electrons. This method is called orbital notation or writing in terms of shells and subshells.

How orbital notation works

In orbital notation, we write:

• The number of the main energy level
• The letter of the sublevel type (s, p, d, or f)
• A superscript number showing how many electrons are in that sublevel

For example: $1s^2$ means 2 electrons in the 1s sublevel.

Examples of orbital notation

Electron configuration diagrams

We can also represent electrons in orbitals using a box diagram, where each box represents one orbital and arrows represent electrons:

These tables show three ways to represent electron configuration for elements 1-20:

1. Box diagrams: Arrows in boxes (↑↓ means two electrons with opposite spins in one orbital)
2. Orbital notation: Using the $1s^2 2s^2$ format
3. Simple form: Using comma-separated numbers like 2,8,1

Why this matters

Understanding orbitals and sublevels helps explain:

• Why noble gases are so stable (they have completely filled s and p sublevels in their outer level)
• Why the periodic table has its particular structure (elements in the same group have similar outer sublevel configurations)
• Why there are 10 transition metals in each row (they correspond to filling a d sublevel with 10 electrons)
• The chemical behaviour of elements (which depends largely on outer electron arrangement)

Exam tips

Remember!