Electronic configuration
Learn more about electrons and electron configurations using this helpful resource!
The electron
Electrons are fundamental subatomic particles with a negative electric charge. They play a crucial role in many aspects of science and technology.
Bohr planetary (particle) model and Schrödinger's wave mechanical model
Electrons have historically been described as tiny particles that revolve around the nucleus in circular orbits similar to the solar system. This is Bohr's particle model. However, this model overly simplifies the path of electrons within atoms and was eventually rejected.
The modern understanding of electron behaviour is based on Schrödinger's wave mechanical model.
The wave model shows that electrons don't have exact paths. Instead, they exist in regions of space where they are likely to be found. These regions are called orbitals. In this model, electrons occupy a three-dimensional space around the nucleus.
Bohr and Schrödinger models
The Bohr model
- A central black dot labeled "p+" representing the nucleus.
- A red oval path around the nucleus labeled "e-" representing an electron orbit.
The Schrödinger model
- A central black dot labeled "p+" representing the nucleus.
- A complex, cloud-like blue shape around the nucleus labeled "e-" representing an electron cloud.
When an electron's wave fits perfectly within an orbital, it creates a stable energy level, known as a "stationary state." This means the electron's energy is quantised, meaning that it can exist only in certain allowed states. These stable energy levels are crucial for understanding how electrons behave in atoms, influencing how atoms absorb and emit energy, and determining the chemical properties of elements. Orbitals, therefore, provide a framework for predicting where electrons are likely to be and how they will interact with other atoms in chemical reactions.
Electron energy levels
The space where electrons move around the nucleus is divided into subspaces known as shells, subshells and orbitals.
Electron shells
Electron shells are named from the nucleus outwards, as \(1\), \(2\), \(3\), \(4\) and so on. The energy of the electron shells increases with the distance from the nucleus. Therefore, the electrons closest to the nucleus have lower energy than the electrons located further away. Each electron shell can contain a maximum number of electrons. These numbers can be calculated using \(2n^2\). For example, shell \(n=2\) contains \(2\times2^{2}=8\textrm{ electrons}.\)
Electron shell diagram
- A central black dot representing the nucleus.
- Three concentric circles around the nucleus, labeled:
- n=1 (innermost circle)
- n=2 (middle circle)
- n=3 (outermost circle)
- A red arrow pointing outward from the nucleus through the circles, indicating "Increasing energy and distance from the nucleus".
Electron subshells
Within each electron shell are subshells. The number of subshells within a shell is the same as the shell number. For example, the \(n=3\) shell contains \(3\textrm{ subshells}\) and the \(n=6\) shell contains \(6\textrm{ subshells}\). Subshells are labeled as \(s,p,d\) and \(f\). Each subshell contains a certain number of electrons:
- An \(s\) subshell contains up to \(2\) electrons.
- A \(p\) subshell contains up to \(6\) electrons.
- A \(d\) subshell contains up to \(10\) electrons.
- A \(f\) subshell contains up to \(14\) electrons.
Let's consider the \(n=4\) shell. This shell contains \(4\) subshells, which are labeled as \(4s\), \(4p\), \(4d\), and \(4f\). Remember, the \(4s\) subshell can hold up to \(2\) electrons, the \(4p\) subshell can hold up to \(6\) electrons, the \(4d\) subshell can accommodate up to \(10\) electrons, and the \(4f\) subshell can contain up to \(14\) electrons. Altogether, the \(n=4\) shell can hold a maximum of \(2+6+10+14=32\) electrons.
Electron subshells are filled in order of increasing energy. The figure shows the electron filling order. Generally, shells fill from \(1\) to \(2\) to \(3\) and subshells, from \(s\) to \(p\) to \(d\), but there is an exception. Subshell \(4s\) has lower energy than subshell \(3d\); therefore, \(4s\) fills before \(3d\).
Electron shell filling diagram
- Left side: A vertical arrow labeled "Energy" pointing upwards.
-
Top labels:
- "Shell number" on the left
- "Filling order" on the right
-
Energy levels (from bottom to top):
- Shell 1: 1s (blue line)
- Shell 2:
- 2s (red line)
- 2p (red line)
- Shell 3:
- 3s (yellow line)
- 3p (yellow line)
- 3d (yellow line)
- Shell 4:
- 4s (blue line)
- 4p (blue line)
- 4d (blue line)
- Shell 5:
- 5s (red line)
- 5p (red line)
- Connecting lines: Lines indicate the order of filling, showing connections between shells such as from 3p to 4s to 3d.
Electron orbitals
Within each electron subshell are electron orbitals. Each electron orbital can hold up to two electrons, where one spin upwards and the other spins downwards. This means that:
- An \(s\) subshell, which contains up to \(2\) electrons, has one orbital.
- A \(p\) subshell, which contains up to \(6\) electrons, has three orbitals.
- A \(d\) subshell, which contains up to \(10\) electrons, has five orbitals.
- An \(f\) subshell, which contains up to \(14\) electrons, has seven orbitals.
You can see that the number of orbitals is half of the number of maximum electrons.
Orbitals have distinct shapes. For example, \(s\) orbitals are spherical and \(p\) orbitals are dumbbell-shaped. Orbitals within the same subshell that have the same shape are generally of equal energy but have a different orientation in space, such as \(p_{x}\), \(p_{y}\) and \(p_{z}\), where \(x\), \(y\), and \(z\) denote the orientation in space.
Electron orbitals
s orbital
- A red spherical shape labeled "s orbital".
- Axes labeled x, y, and z around the sphere.
p orbitals
-
px orbital:
- Two blue lobes along the x-axis labeled "px orbital".
- Axes labeled x, y, and z.
-
py orbital:
- Two blue lobes along the y-axis labeled "py orbital".
- Axes labeled x, y, and z.
-
pz orbital:
- Two blue lobes along the z-axis labeled "pz orbital".
- Axes labeled x, y, and z.
Like shells and subshells, electron fill orbitals from lowest to highest energy level. This is called the Aufbau principle. Electrons also fill orbitals of a subshell such that each orbital acquires one electron before any orbital acquires a second electron. All single electrons must have the same spin.
Aufbau principle diagram
- Energy levels labeled n=1 to n=7 on the left.
-
Diagonal lines indicate the filling order for orbitals:
- n=1: 1s
- n=2: 2s, 2p
- n=3: 3s, 3p, 3d
- n=4: 4s, 4p, 4d, 4f
- n=5: 5s, 5p, 5d, 5f
- n=6: 6s, 6p, 6d, 6f
- n=7: 7s, 7p, 7d, 7f
- Black and red arrows guide the sequence in which orbitals are filled.
Electron configurations
Electron configurations show how the electrons are distributed within an atom. They are written from lowest energy subshell to the highest energy subshell. Each subshell is written with the number of electrons in each subshell shown as a superscript. The superscripts in an electron configuration will add up to the total number of electrons for that atom.
Full electron configurations
Full electron configurations show the location of all electrons in an atom. The full electron configurations of the first 10 elements are shown in the table.
| Element | Atomic number | Electron configuration |
|---|---|---|
| \(\ce{H}\) | \(1\) | \(1s^{1}\) |
| \(\ce{He}\) | \(2\) | \(1s^{2}\) |
| \(\ce{Li}\) | \(3\) | \(1s^{2}\,2s^{1}\) |
| \(\ce{Be}\) | \(4\) | \(1s^{2}\,2s^{2}\) |
| \(\ce{B}\) | \(5\) | \(1s^{2}\,2s^{2}2p^{1}\) |
| \(\ce{C}\) | \(6\) | \(1s^{2}\,2s^{2}2p^{2}\) |
| \(\ce{N}\) | \(7\) | \(1s^{2}\,2s^{2}2p^{3}\) |
| \(\ce{O}\) | \(8\) | \(1s^{2}\,2s^{2}2p^{4}\) |
| \(\ce{F}\) | \(9\) | \(1s^{2}\,2s^{2}2p^{5}\) |
| \(\ce{Ne}\) | \(10\) | \(1s^{2}\,2s^{2}2p^{6}\) |
Condensed electron configuration
Condensed electron configurations are a short way of showing electron configurations. They are often also called noble gas configurations and are handy to avoid writing out all of the subshells. This way of representing electron configurations states the previous noble gas in square brackets, followed by the remaining electrons.
For example, let's look at the magnesium atom, which has a full electron configuration of \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\). The previous noble gas, as shown on the periodic table, is neon. Neon has an electron configuration of \(1s^{2}\,2s^{2}\,2p^{6}\). You can see that the location of the first 10 electrons is the same as in magnesium, so we can replace this part of the electron configuration with the chemical symbol for the noble gas in square brackets: \(\textrm{[Ne]}\). This makes the condensed electron configuration for magnesium: \(\textrm{[Ne]}\,3s^{2}\).
Orbital diagrams
Orbital diagrams depict spin of individual electrons within each subshell. The orbital diagrams for the first 10 elements are also shown.
Electron configurations and corresponding orbital diagrams
| Element | Electron configuration | Orbital diagram |
|---|---|---|
| H | 1s1 | 1s: ↑ |
| He | 1s2 | 1s: ↑↓ |
| Li | 1s2 2s1 | 1s: ↑↓ 2s: ↑ |
| C | 1s2 2s2 2p2 | 1s: ↑↓ 2s: ↑↓ 2p: ↑ ↑ |
| N | 1s2 2s2 2p3 | 1s: ↑↓ 2s: ↑↓ 2p: ↑ ↑ ↑ |
| O | 1s2 2s2 2p4 | 1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑ ↑ |
Example 1 – writing full electron configurations
Write the full electron configuration for \(\ce{Na}\).
Step 1: Determine the number of electrons. This is the same as the atomic number (Z).
\[Z=11\]
Therefore, \(\ce{Na}\) has \(11\textrm{ electrons}\).
Step 2: Place the electrons into orbitals according to the Aufbau principle. The first two electrons go into the \(1s\) orbital as a pair. The next two electrons are placed in the \(2s\) orbital. The next six electrons are placed in \(2p\) orbitals as three pairs. The remaining electron is placed in \(3s\) orbital.
Orbital diagram for sodium
| Element | Electron configuration | Orbital diagram |
|---|---|---|
| Na | 1s2 2s2 2p6 3s1 | 1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑↓ ↑↓ 3s: ↑ |
Step 3: Convert the orbital diagram into an electron configuration, listing the subshells in order of increasing energy. Write the total number of electrons in each subshell as a superscript.
\(\ce{Na}\) has the full electron configuration \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{1}\).
Write the full electron configuration for \(\ce{Si}\) using an orbital diagram.
Step 1: Determine the number of electrons. This is the same as the atomic number (Z).
\[Z=14\]
Therefore, \(\ce{Si}\) has \(14\textrm{ electrons}\).
Step 2: Place the electrons into orbitals according to the Aufbau principle. The first two electrons go to \(1s\) orbital as a pair. The next eight electrons are placed in the second shell (\(n=2)\) as pairs. The last four electrons enter into the third shell, where the first two occupy the \(s\) orbital as a pair and the last two fill two \(p\) orbitals as single electrons.
Step 3: Convert the orbital diagram into an electron configuration, listing the subshells in order of increasing energy. Write the total number of electrons in each subshell as a superscript.
\(\ce{Si}\) has the full electron configuration \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{2}\).
Example 3 – writing condensed electron configurations
Write the condensed electron configuration for \(\ce{Ca}\).
Step 1: Use the steps in Example 1 – writing full electron configurations to write the full electron configuration for \(\ce{Ca}\).
\(\ce{Ca}\) has the full electron configuration \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,4s^{2}\).
Step 2: Determine the full electron configuration of the previous noble gas, according to the periodic table. For \(\ce{Ca}\), the previous noble gas is \(\ce{Ar}\).
\(\ce{Ar}\) has the full electron configuration \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\).
Step 3: Write the condensed electron configuration by replacing the electron configuration of \(\ce{Ar}\) within the electron configuration of \(\ce{Ca}\) with \(\textrm{[Ar]}\).
\(\ce{Ca}\) has the condensed electron configuration \(\textrm{[Ar]}\,4s^{2}\).
Your turn – the electron
Test yourself on your understanding of electrons.
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