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Module 2: Foundations in chemistry›Electron structure (2.2.1)

Electron structure (2.2.1)

An overview of how electrons are arranged within shells, orbitals and suborbitals within an atom.

6 min

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Electrons occupy specific energy levels outside the nucleus of an atom.

The arrangement of electrons in these quantised energy levels in an atom is referred to as its electron structure or electronic configuration.

Principal energy levels or principal quantum shells (n) are numbered based on their distance from the nucleus.

The first principle quantum shell is closest to the nucleus and described as n = 1.

A diagram illustrating the atomic structure, showing a blue nucleus at the center with three concentric circles representing electron orbits. The orbits are labeled with increasing energy levels: n=1, n=2, and n=3. An arrow points outward, indicating increasing energy and distance from the nucleus.

n is an integer value which increases as the shells move further from the nucleus.

The maximum number of electrons that can fill a quantum shell is given by the formula e = 2n²

The number of electrons needed to fully occupy the first four quantum shells is shown below.

A table displaying the relationship between quantum shell numbers and the corresponding number of electrons. The first column lists quantum shell numbers (n) from 1 to 4, and the second column shows the number of electrons that can occupy each shell: 2 for n=1, 8 for n=2, 18 for n=3, and 32 for n=4.

Shells are divided into subshells which can be further divided into atomic orbitals.

The atomic orbital is a three-dimensional region around the nucleus of an atom, where there is a maximum probability of finding an electron.

Each orbital can have a maximum of 2 electrons, with opposite spins.

A diagram illustrating atomic structure, featuring a central sub-shell represented by a pink sphere, surrounded by various orbitals depicted as shapes in different orientations. The orbitals are labeled with axes (X, Y, Z) and categorized into shells and sub-shells.

Within a shell, the atomic orbitals are grouped as subshells, distinguished by the letters s,p,d and f.

The s subshell has one atomic orbital, the p subshell has three orbitals, the d subshell has five orbitals while the f subshell has a total of seven atomic orbitals.

A diagram illustrating the shapes of atomic orbitals: s, p, d, and f. The s orbital is shown as a spherical shape at the top, followed by three p orbitals represented as two-lobed shapes. Below are five d orbitals with various lobular configurations, and at the bottom, seven f orbitals with more complex shapes. Each orbital is labeled with its corresponding quantum number and axes are indicated.

As the energy level of the shell increases, the size of the atomic orbitals within the subshells increases but their shapes remain similar.

The number of subshells per shell in an atom increases with the increase in principal quantum number.

The number of electrons per shell can be calculated by considering full occupancy of all the atomic orbitals present.

A table displaying the principal quantum numbers, corresponding sub-shells, and the maximum number of electrons that can occupy each shell. The table includes four rows for quantum numbers 1 to 4, detailing the sub-shells (1s, 2s, 2p, etc.) and their respective electron capacities.

An atomic orbital has a three-dimensional shape indicating the area in which an electron is likely to be found.

The s orbital has a symmetrical, spherical shape, with the nucleus located at the centre.

Illustration showing the 3D shape of an s-orbital, represented as a pink sphere with labeled axes (X, Y, Z) and a nucleus at the center. To the right, a simplified 2D representation of the s-orbital is shown as a plain circle.

The s orbitals of all shells have the same spherical shape but differ in the size of their radius.

The p orbitals are described as dumb-bell shaped, with the nucleus located between the two halves.

There are three degenerate p orbitals in each p subshell. These orbitals have an identical shape and energy, however, their orientation in space differs. They lie at right angles to one another along the X, Y, and Z axes of a Cartesian plane.

The three p orbitals are thus distinguished as p_x, p_y, and p_z.

Illustration showing three 2p orbitals on the Cartesian plane, labeled X, Y, and Z. The top section displays a combined view of the orbitals, while the bottom section presents the individual shapes of the 2px, 2py, and 2pz orbitals.

Electrons always fill up atomic orbitals in an order of increasing orbital energy. This leads to the most stable electronic configuration of the atom; the one that has the lowest overall energy.

Orbital energy increases with increasing principle quantum number (shell number). 4 > 3 > 2 > 1.

The subshells within a shell have slightly different energy levels with the trend f > d > p > s.

All orbitals in the same subshell are degenerate; they have an equal energy. Electrons-electron repulsion means all degenerate orbitals are filled singly before any electrons are paired.

The first electron is always placed in the 1s orbital, the lowest energy orbital.

Subshell notation describes the electrons contained in each subshell of an atom:

X denotes the principal quantum number (shell number)
y shows the subshell type (s, p, d, f)
z represents the number of electrons occupying the subshell

Using the subshell notation the electronic configuration of carbon (Z = 6) is:

1s² 2s² 2p²

Make use of your periodic table to remember the order in which the subshells are filled.

List the subshells in order as you move in increasing atomic number from hydrogen to the element of interest.

A periodic table of elements displaying the arrangement of elements in rows and columns, with color-coded sections for different electron configurations. The table includes labels for s, p, d, and f orbitals, along with the atomic numbers and symbols for each element.

The energy levels of 3d and 4s subshells overlap.

The energy of 4s lies slightly below that of 3d therefore 4s is filled before 3d.

A diagram illustrating the energy levels of electron shells and sub-shells in an atom. The vertical axis represents increasing energy, while the horizontal axis shows shell numbers (n = 1 to n = 4) with corresponding sub-shells labeled (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d).

However, once filled, the energy of 3d falls below the 4s energy level.

When a transition metal is ionised, the 4s electrons are easier to remove than the 3d electrons.

Box notation goes beyond subshell notation as it describes which atomic orbitals contain electrons. Each atomic orbital is drawn as a box and each electron as an arrow.

A diagram illustrating the electron configuration in atomic orbitals. The top section shows all orbitals singly occupied before pairing, with arrows indicating electron spins. Subsequent sections depict various configurations of electrons in atomic orbitals, including paired and unpaired electrons, with labels for 'Atomic orbital' and 'Electron'.

All electrons are identical negatively charged particles. Two electrons close to each other experience strong electrostatic repulsions.

To minimise electron–electron repulsion, when filling degenerate orbitals (such as three 2p orbitals), electrons first singly occupy each orbital before pairing occurs.

Two electrons in the same orbital must be in opposite spin (clockwise and anticlockwise) denoted by an upwards (↑) and a downwards (↓) pointing arrow.

A simple black and white illustration showing two vertical arrows, one pointing upwards and the other pointing downwards, enclosed in a square.

This minimises the repulsive effect between the two electrons.

The electronic configuration of oxygen is 1s² 2s² 2p⁴.

Common mistakes to avoid using box notation are:

doubly filling degenerate orbitals from left to right without singly filling first
pairing electrons with the same spin

An educational diagram illustrating the rules of electron configuration in atomic orbitals. The top section shows incorrect configurations with electrons in the same orbital having the same spin, marked with a red 'X'. The bottom section displays correct configurations with paired electrons in opposite spins, marked with a green check. Labels indicate the different orbitals: 1s, 2s, 2px, 2py, and 2pz.

Ions are formed by the loss or gain of electrons by the atom of an element.

The electronic configuration of an ion can be written using the same set of rules as that for the element and then adjusting for the increase or decrease in electron count.

For ionisation, it is important to adjust from the element rather than filling the atomic orbitals directly using the ions’ electron count. This is because, when filled, the 4s orbital becomes higher in energy than the 3d orbital.

Electrons must be taken from the 4s subshell before being removed from the 3d subshell.

A diagram illustrating the electron configuration for Scandium (Sc) and Vanadium ion (V2+). The top section shows electrons filled into the 4s orbital before the 3d orbital for Sc, while the bottom section shows electrons removed from the 4s orbital before the 3d orbital for V2+. Each orbital is represented with arrows indicating electron spins.

Vanadium 2+, , and scandium, , both have an electron count of 21 but their electron configuration is different.