Electronic structure · GCSE Chemistry

Electron shells

Electrons occupy the lowest available energy levels (shells), starting nearest the nucleus.
For the first 20 elements, the first shell holds up to 2 electrons, the second up to 8, and the third up to 8 (the fourth shell starts at potassium).
Electronic structures are written as numbers, for example sodium (11 electrons) is 2,8,1.

The number of electrons equals the atomic number (for a neutral atom).
Fill the shells in order: 2, then 8, then 8, then the rest.
Examples: carbon (6) is 2,4; oxygen (8) is 2,6; chlorine (17) is 2,8,7; calcium (20) is 2,8,8,2.

The number of shells tells you the period (row). The number of electrons in the outer shell tells you the group (for groups 1 to 7).
For example, magnesium is 2,8,2: period 3, group 2.
Elements in the same group have the same number of outer electrons, so they react in similar ways.
Group 0 elements have full outer shells, which is why they are unreactive.

Metal atoms lose outer electrons to form positive ions: Na (2,8,1) becomes Na⁺ (2,8).
Non-metal atoms gain electrons to form negative ions: Cl (2,8,7) becomes Cl⁻ (2,8,8).
The ions have the same electronic structure as a noble gas.

Electronic structure: The arrangement of electrons in shells, such as 2,8,1.
Energy level: A shell in which electrons are found, at a particular distance from the nucleus.
Outer shell: The highest occupied energy level; its electrons decide how an element reacts.
Period: A horizontal row in the periodic table; its number equals the number of occupied shells.
Group: A vertical column in the periodic table; for groups 1 to 7 it matches the number of outer electrons.
Ion: An atom or group of atoms that has gained or lost electrons, so it has a charge.
Noble gas structure: A full outer shell, which is very stable.