Electronic Configurations

Electrons are arranged in shells (also called energy levels) around the nucleus. The shell closest to the nucleus has the lowest energy, so electrons always fill the inner shells first before starting a new one.

For the first 20 elements the maximum number of electrons each shell can hold is:

• 1st shell: holds up to 2 electrons.
• 2nd shell: holds up to 8 electrons.
• 3rd shell: holds up to 8 electrons (for the first 20 elements).
• 4th shell: just starts to fill (with the last 1 or 2 electrons) for potassium and calcium.

So you fill electrons in the order 2, 8, 8, 2. The number of electrons to place equals the atomic number (because a neutral atom has equal protons and electrons).

How to write a configuration from the atomic number. Take the atomic number, then "spend" the electrons shell by shell:

• Sodium, Na (Z = 11): 2 in shell 1, 8 in shell 2, 1 left → 2,8,1.
• Chlorine, Cl (Z = 17): 2, then 8, then 7 left → 2,8,7.
• Argon, Ar (Z = 18): 2, then 8, then 8 → 2,8,8 (full).
• Calcium, Ca (Z = 20): 2, 8, 8, then 2 left → 2,8,8,2.

Write the numbers separated by commas with the inner shell first. The numbers must add up to the atomic number — always check this.

As well as writing 2,8,1, you must be able to draw the configuration as a shell (Bohr) diagram. The rules:

• Draw the nucleus as a small circle in the centre with the element's symbol inside.
• Draw a circle (shell) for each occupied energy level.
• Put the correct number of electrons (small dots or crosses) on each circle, spacing them out neatly (often in pairs).
• The total number of electrons drawn must equal the atomic number.

Marks-saving tips.

• Count your electrons before you submit — the total must equal the atomic number.
• Don't overfill: never put more than 2 in the first circle or 8 in the second.
• A neat outer shell makes the group number obvious, which is often the next mark.

The electronic configuration is directly linked to an element's position in the periodic table — this is one of the most commonly tested ideas.

Group number = number of electrons in the OUTER shell (for Groups 1–7).

• Na = 2,8,1 → 1 outer electron → Group 1.
• Mg = 2,8,2 → 2 outer electrons → Group 2.
• Cl = 2,8,7 → 7 outer electrons → Group 7.
• Ar = 2,8,8 → full outer shell → Group 0 (the noble gases).

Period number = number of OCCUPIED shells.

• Li = 2,1 → 2 shells → Period 2.
• Na = 2,8,1 → 3 shells → Period 3.
• K = 2,8,8,1 → 4 shells → Period 4.

Working backwards. You can also use a configuration to identify an element (count the electrons → that's the atomic number → look it up on the periodic table) or to predict its position without naming it.

This pattern works because elements in the same group have the same number of outer electrons, which is why they have similar chemical properties — the outer electrons control how an atom reacts.

An atom's reactivity is decided by its outer shell. Atoms "want" the stable, full outer shell of the nearest noble gas.

• Full outer shell (2 for helium; 8 for neon, argon, etc.) → already stable → Group 0 noble gases are unreactive. They do not need to gain, lose or share electrons.
• 1 or 2 outer electrons (Groups 1 and 2) → an atom can easily lose them to leave a full shell beneath → reactive metals.
• 6 or 7 outer electrons (Groups 6 and 7) → an atom can easily gain electrons to complete the shell → reactive non-metals.

Worked logic — sodium vs argon.

• Na = 2,8,1. One lonely outer electron is easy to lose → sodium is very reactive.
• Ar = 2,8,8. A full outer shell → argon is unreactive (used to fill light bulbs because it won't react).

So when an exam asks "explain why this element is unreactive", the mark is for: it has a full outer shell (of electrons). When it asks "why is it reactive", the mark is for: it has only 1/2 outer electrons to lose or needs only 1/2 more to fill its shell.