Principal Energy Levels & Sub-Levels

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Energy levels / shells numbered 1, 2, 3, 4, etc. - called principal energy levels.

Number of principal energy levels matches the period number of the element.

First energy level accommodates two electrons - filled first, closest to nucleus.

Second energy level holds up to eight electrons - filled after first energy level but before third.

Third energy level accommodates up to 18 electrons.

Fourth energy level accommodates up to 32 electrons.

Principal energy levels divided into a number of sub-levels. Sub-levels assigned the letters s, p, d and f and contain orbitals.

The s sub-level contains one orbital - holds two electrons max.

The p sub-level contains three orbitals - each orbital holds two electrons max. so six electrons in total.

The d sub-level contains five orbitals - each one can hold two electrons max. so ten electrons in total.

The f sub-level contains seven orbitals - each one can hold two electrons max. so fourteen electrons in total.

Electron orbitals represent a region of space around atom's nucleus where there is a high probability of finding a particular electron.

Full electronic configuration of element gives information about number of electrons in each occupied principal energy level and sub-level.

Filling order follows a pattern which takes into account principal energy level, sub-level and maximum number of electrons in each energy level.

Electronic configuration of an element is such that the lowest possible energy state is obtained so electrons placed in lowest energy orbitals first.

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Break in logical filling order after 3p as 4s orbital filled before 3d.

Simply because 4s orbital is of lower energy than 3d orbital.

Once 3d orbitals filled, 4p orbitals filled - known as Aufbau Principle.

Chromium and copper have one electron in their 4s orbital as they have half-filled and filled d-orbitals respectively.

These are the lowest energy configurations as half full and full orbitals have the most stability.

Cr [Ar] 3d⁵ 4s¹ is more stable than Cr [Ar] 4s² 3d⁴

Cu [Ar] 3d¹⁰ 4s¹ is more stable than Cu [Ar] 4s² 3d⁹

Positive ions (cations) formed when electrons lost = positive charge.

Negative ions (anions) formed when electrons gained = negative charge.

Isoelectronic where two ions have the same electronic structure.

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Pauli Exclusion Principle - any electron orbital can hold one or two electrons but no more.

Electrons carry charge but they also carry spin.

Electrons can exist in the same orbital as a pair because they have opposite spin.

Reduces mutual repulsion so one electron goes clockwise, the other anti-clockwise.

Anticlockwise represented by a down arrow.

Clockwise represented by an up arrow.

Electrons will occupy orbitals as unpaired electrons due to like charge repulsion but only until orbitals of equal energy have one pair each.

Key ideas:

Electrons fill lowest available orbitals first.

4s orbital filled first before the 3d orbital as it is of lower energy.

Sub-levels split into orbitals which can carry a maximum of two electrons each.

s, p, d and f orbitals have different shapes.

Electron-electron repulsion reduced in sub-levels by having unpaired electrons in preference to paired electrons.

In transition metals, electrons lost from the 4s orbital first to form ions.