Electronic Configuration (Cambridge O Level Chemistry)

Electronic configuration

• We can represent the structure of the atom in two ways: using diagrams called electron shell diagrams or by writing out a special notation called the electronic configuration (or electronic structure or electron distribution)

Electron shell diagrams

• Electrons orbit the nucleus in shells (or energy levels) and each shell has a different amount of energy associated with it
• The further away from the nucleus, the more energy a shell has
• Electrons fill the shell closest to the nucleus 
• When a shell becomes full of electrons, additional electrons have to be added to the next shell
• The first shell can hold 2 electrons
• The second shell can hold 8 electrons 
• For this course, a simplified model is used that suggests that the third shell can hold 8 electrons
  • For the first 20 elements, once the third shell has 8 electrons, the fourth shell begins to fill
• The outermost shell of an atom is called the valence shell and an atom is much more stable if it can manage to completely fill this shell with electrons 

A simplified model showing the electron shells

• The arrangement of electrons in shells can also be explained using numbers
• Instead of drawing electron shell diagrams, the number of electrons in each electron shell can be written down, separated by commas
• This notation is called the electronic configuration (or electronic structure)
  • E.g. Carbon has 6 electrons, 2 in the 1st shell and 4 in the 2nd shell
    • Its electronic configuration is 2,4
• Electronic configurations can also be written for ions
  • E.g. A sodium atom has 11 electrons, a sodium ion has lost one electron, therefore has 10 electrons; 2 in the first shell and 8 in the 2nd shell
    • Its electronic configuration is 2,8

The Electronic Configuration of the First Twenty Elements

Element	Atomic Number	Electronic Configuration
hydrogen	1	1
helium	2	2
lithium	3	2,1
berylium	4	2,2
boron	5	2,3
carbon	6	2,4
nitrogen	7	2,5
oxygen	8	2,6
fluorine	9	2,7
neon	10	2,8
sodium	11	2,8,1
magnesium	12	2,8,2
aluminium	13	2,8,3
silicon	14	2,8,4
phosphorus	15	2,8,5
sulfur	16	2,8,6
chlorine	17	2,8,7
argon	18	2,8,8
potassium	19	2,8,8,1
calcium	20	2,8,8,2

Note: although the third shell can hold up to 18 electrons, the filling of the shells follows a more complicated pattern after potassium and calcium. For these two elements, the third shell holds 8 and the remaining electrons (for reasons of stability) occupy the fourth shell first before filling the third shell.

Electronic Configuration of Ions 

• Ions are formed when an atom loses or gains electrons to become stable 
  • Positively charged ions are formed when an atom loses electrons
  • Negatively charged ions are formed when an atom gains electrons
  • The size of the charge indicates the number of electrons that have been lost or gained
• To find the electronic configuration of an ion: 
  • Identify the electronic configuration for the atom
  • Identify whether it has lost or gained electrons from its charge
  • Add or remove electrons depending on the charge of the atom 

• There is a clear relationship between the electronic configuration and how the Periodic Table is designed
• The number of notations in the electronic configuration will show the number of occupied shells of electrons the atom has, showing the period in which that element is in
• The last notation shows the number of outer electrons the atom has, showing the group that element is in (for elements in Groups I to VII)
• Elements in the same group have the same number of outer shell electrons 

The electronic configuration for chlorine

 

Period: The red numbers at the bottom show the number of notations which is 3, showing that a chlorine atom has 3 occupied shells of electrons and is in Period 3

Group: The final notation, which is 7 in the example, shows that a chlorine atom has 7 outer electrons and is in Group VII

 

The position of chlorine on the Periodic Table

 

• In most atoms, the outermost shell is not full and therefore these atoms react with other atoms in order to achieve a full outer shell of electrons (which would make them more stable)
• In some cases, atoms lose electrons to entirely empty this shell so that the next shell below becomes a (full) outer shell
• All elements wish to fill their outer shells with electrons as this is a much more stable configuration

The noble gases

• The atoms of the Group VIII elements (the noble gases) all have a full outer shell of electrons
• All of the noble gases are unreactive as they have full outer shells and are thus very stable

The noble gases are on the Periodic Table in Group 8/0