Shapes of Orbitals (Edexcel International A Level Chemistry)

Orbitals

• We have seen before that the total number of electrons that each subshell in each orbital can hold are:
  • s : one orbital (1 x 2 = total of 2 electrons)
  • p : three orbitals ( 3 x 2 = total of 6 electrons)
  • d : five orbitals (5 x 2 = total of 10 electrons)
  • f : seven orbitals (7 x 2 = total of 14 electrons)

• • The orbitals have specific 3-D shapes

s orbital shape

• The s orbitals are spherical in shape
• The size of the s orbitals increases with increasing shell number
  • E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1)

p orbital shape

• The p orbitals have a dumbbell shape
• Every shell has three p orbitals except for the first one (n = 1)
• The p orbitals occupy the x, y and z axes and point at right angles to each other, so are oriented perpendicular to one another
• The lobes of the p orbitals become larger and longer with increasing shell number

Representation of orbitals (the dot represents the nucleus of the atom) showing spherical s orbitals (a), p orbitals containing ‘lobes’ along the x, y and z axis 

• Note that the shape of the d orbitals is not required

 An overview of the shells, subshells and orbitals in an atom

• Electrons can be imagined as small spinning charges which rotate around their own axis in either a clockwise or anticlockwise direction
  • The spin of the electron is represented by its direction
  • The spin creates a tiny magnetic field with N-S pole pointing up or down

Electrons can spin either in a clockwise or anticlockwise direction around their own axis

• Electrons with the same spin repel each other which is also called spin-pair repulsion
  • Therefore, electrons will occupy separate orbitals in the same subshell first to minimise this repulsion and have their spin in the same direction
  • They will then pair up, with a second electron being added to the first p orbital, with its spin in the opposite direction

• This is known as Hund's Rule
  • E.g. if there are three electrons in a p subshell, one electron will go into each p_x, p_y and p_z orbital

  

Electron configuration: three electrons in a p subshell

• The principal quantum number indicates the energy level of a particular shell but also indicates the energy of the electrons in that shell
  • A 2p electron is in the second shell and therefore has an energy corresponding to n = 2

• Even though there is repulsion between negatively charged electrons, they occupy the same region of space in orbitals
• An orbital can only hold two electrons and they must have opposite spin - the is known as the Pauli Exclusion Principle
• This is because the energy required to jump to a higher empty orbital is greater than the inter-electron repulsion
• For this reason, they pair up and occupy the lower energy levels first