1.16-1.22 - Orbitals And Electronic Configurations

Shells, Subshells And Electron Addresses

A chemist can often tell where an element belongs from a compact electron address such as $\mathrm{1s}^{2}\,\mathrm{2s}^{2}\,\mathrm{2p}^{6}\,\mathrm{3s}^{2}$. That address has three layers:

• the number, such as $3$, names the shell or principal energy level
• the letter, such as $\mathrm{s}$ or $\mathrm{p}$, names the subshell
• the superscript, such as 2, counts electrons in that subshell

The first four quantum shells can hold these maximum numbers of electrons:

The maximum capacity is not the same as the order electrons fill in real ground-state atoms. A GCSE-style shell diagram for calcium may look like 2,8,8,2, but that does not mean the third shell can only hold 8 electrons. It means that, for calcium, electrons enter $\mathrm{4s}$ before the $\mathrm{3d}$ subshell starts filling.

Orbitals And Subshell Capacities

Each subshell contains one or more orbitals. The number of orbitals sets its electron capacity.

Opposite spins are often shown as one upward arrow and one downward arrow in a box.

[DIAGRAM: asset_name: S And P Orbitals; asset_slug: edexcel_a_level_chemistry_l006_s_and_p_orbitals; recommended_method: matplotlib; description: Pale-colour schematic showing a spherical s orbital and a dumbbell-shaped p orbital, each with a central nucleus point.]

In the diagram, the sphere on the left represents an s orbital; the two-lobed shape on the right represents a p orbital. They show orbital regions, not paths traced by individual electrons.

For this specification, you need the shapes of s and p orbitals:

• an s orbital is spherical
• a p orbital is dumbbell-shaped, with two lobes on opposite sides of the nucleus

You also need the number of electrons that occupy each type of subshell:

The capacity follows from "one orbital holds up to two electrons". A p subshell has three p orbitals, so it can hold $3\times 2 = 6$ electrons. A d subshell has five d orbitals, so it can hold $5\times 2 = 10$ electrons. You do not need d-orbital shapes in this lesson, but you do need the d-subshell capacity because configurations up to krypton include $\mathrm{3d}$.

Filling Orbitals Without Losing Electrons

For ground-state configurations up to Z = 36, use this filling order:

$\mathrm{1s}$ → $\mathrm{2s}$ → $\mathrm{2p}$ → $\mathrm{3s}$ → $\mathrm{3p}$ → $\mathrm{4s}$ → $\mathrm{3d}$ → $\mathrm{4p}$

Two filling rules do a lot of the work:

• electrons fill orbitals of equal energy in the same subshell singly with parallel spins before pairing; this is the lowest-energy distribution
• two electrons in the same orbital must have opposite spins

In electron-in-box notation, each box is one orbital. The three boxes of a p subshell have the same energy, so electrons spread out before they pair.

For nitrogen, $Z = 7$, so there are 7 electrons:

1s  [↑↓]
2s  [↑↓]
2p  [↑] [↑] [↑]

This is $\mathrm{1s}^{2}\,\mathrm{2s}^{2}\,\mathrm{2p}^{3}$. The $\mathrm{2p}$ electrons are not drawn as [↑↓] [↑] [ ], because pairing starts only after each orbital in that subshell has one electron.

For oxygen, $Z = 8$, one more electron is added:

1s  [↑↓]
2s  [↑↓]
2p  [↑↓] [↑] [↑]

This is $\mathrm{1s}^{2}\,\mathrm{2s}^{2}\,\mathrm{2p}^{4}$. The paired electrons in the first $\mathrm{2p}$ box have opposite spins.

Worked example: draw the box notation for sulfur

Sulfur has $Z = 16$, so a sulfur atom has 16 electrons. Fill in order:

The boxes are:

1s  [↑↓]
2s  [↑↓]
2p  [↑↓] [↑↓] [↑↓]
3s  [↑↓]
3p  [↑↓] [↑] [↑]

The chemical interpretation is that sulfur has six outer-shell electrons in shell 3: two in $\mathrm{3s}$ and four in $\mathrm{3p}$. That outer-shell pattern will matter later when explaining periodic properties and bonding, but here the key point is the correct filling and pairing.

Predicting Atom Configurations Up To Krypton

For a neutral atom, the number of electrons equals the atomic number, $Z$. The method is mechanical, but the checking is chemical:

1. Count electrons from $Z$.
2. Fill subshells in the order $\mathrm{1s}$, $\mathrm{2s}$, $\mathrm{2p}$, $\mathrm{3s}$, $\mathrm{3p}$, $\mathrm{4s}$, $\mathrm{3d}$, $\mathrm{4p}$.
3. Stop when the electron total equals $Z$.
4. Check the last occupied subshell against the element's periodic-table position.

Worked example: selenium, Z = 34

Start with 34 electrons and fill:

Count the electrons:

$2 + 2 + 6 + 2 + 6 + 2 + 10 + 4 = 34$

Selenium is in the p-block, and the configuration ends in $\mathrm{4p}^{4}$, so the periodic-table check agrees.

For longer configurations, noble-gas shorthand is a useful working check:

When a question asks for $\mathrm{1s}$ notation, the full form is safest unless the question or mark scheme clearly allows shorthand.

Chromium and copper caution

The simple filling pattern predicts most atoms up to krypton, but chromium and copper are the important period-4 exceptions:

Cr: $[\mathrm{Ar}]\,\mathrm{4s}^{1}\,\mathrm{3d}^{5}$

Cu: $[\mathrm{Ar}]\,\mathrm{4s}^{1}\,\mathrm{3d}^{10}$

The small energy difference between the 4s and 3d subshells means electron interactions favour these half-filled or full $\mathrm{3d}$ subshells. The copper exception is commonly tested, so do not write copper as $[\mathrm{Ar}]\,\mathrm{4s}^{2}\,\mathrm{3d}^{9}$.

Ions And Blocks

For ions, change the electron count before writing the configuration.

• A positive ion has lost electrons.
• A negative ion has gained electrons.
• The specification asks for ions only for s- and p-block elements up to $Z = 36$.

For s- and p-block ions, electrons are removed from or added to the outer shell. That usually means the highest principal shell number, such as shell 3 before shell 2.

Worked example: magnesium ion, $\ce{Mg^2+}$

Magnesium has $Z = 12$, so neutral magnesium is:

$\ce{Mg^2+}$ has lost two electrons, so it has 10 electrons:

The two lost electrons came from the $\mathrm{3s}$ subshell, because that is the outer occupied subshell in magnesium.

Worked example: chloride ion, $\ce{Cl-}$

Chlorine has $Z = 17$, so neutral chlorine is:

$\ce{Cl-}$ has gained one electron, so it has 18 electrons:

This is the same electron configuration as argon. It is not the same atom: chloride has 17 protons and a $1-$ charge.

Block classification is based on the subshell being filled in the atom:

Helium is displayed with the noble gases because of its chemical behaviour, but its electron configuration is $\mathrm{1s}^{2}$. The block label follows configuration, while group placement also reflects chemical properties.

For transition-metal ions, such as $\ce{Fe^2+}$, the specification revisits the issue in Topic 15A. The key later caution is that period-4 transition-metal ions lose $\mathrm{4s}$ electrons before $\mathrm{3d}$ electrons; this lesson's ion-prediction requirement is only for s- and p-block ions.

Recap

A complete configuration shows the correct electron count, the correct filling order, the correct pairing in boxes, and the correct ion charge handling. Once the notation is right, block classification is quick: look at the subshell being filled, not the total number of shells in the atom.