Electronic configuration

• Conſegnation


  Development of Electronic Compormations
• Subatomic particles


  • More than 220 have been discovered so far
  • Electrons, neutrons and protons are of much interest to chemists
• Atomic number


  The number of electrons which is the same as the number of protons in an atom
• Atomic mass


  The sum of the number of protons and neutrons
• The number of electrons and their arrangement in an atom determines the reactivity of the element
• Electron configuration


  The arrangement of electrons in an atom of an element
• Quantum numbers


  • Principal, subsidiary, magnetic and spin quantum numbers describe the position of each electron in terms of the shell, subshell, orbital or spin
• Aufbau process


  1. Electrons fill the orbitals first, and then move up to higher energy levels

  2. Lowest energy orbitals are filled first
• There is a problem with the Aufbau process as the 3d orbitals should be filled before the 4s orbitals, because the 3d orbitals have a lower value and thus a lower energy
• Aufbau principle


  A shorthand notation for writing electronic configurations using the formula nxl, where n is the principal quantum number, x is the subshell (orbital) and l is the number of electrons in the subshell
• Three concepts in the Aufbau principle

  • Aufbau rule
  • Hund's rule
  • Pauli exclusion principle
• Aufbau rule


  Electrons will enter the subshell (orbital) of lowest energy first
• The relative energies of the orbitals are in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, etc.
• Pauli exclusion principle


  No two electrons in an atom (or molecule) can have the same set of quantum numbers
• An orbital can contain a maximum of only two electrons, and they must have opposing spins
• Hund's rule


  For a set of orbitals of equal energy, the lowest energy state is the one with the greatest electron spin multiplicity
• Hund's rule states that electrons will singly occupy orbitals of equal energy before pairing up
• Atoms acquire extra stability where degenerate orbitals are either half filled (all parallel spins) or completely filled
• The ground state of atoms follows Hund's rule, where the maximum number of unpaired electrons (and hence maximum total spin state) is assumed
• Orbitals


  The maximum number of electrons that can occupy an orbital
• Unpaired electrons


  The maximum total spin state
• In general, atoms acquire extra stability where degenerate orbitals are either half filled (all parallel spins) or completely filled
• Noble gases have half-filled (8s, 2p, 7s) or completely filled configurations (2p6, 3s2, 3p6) which show enhanced stability
• Nitrogen (Z=7) has electronic configuration 1s2 2s2 2p3, which is not half-filled or completely filled
• Chromium (Z=24) has electronic configuration 1s2 2s2 2p6 3s2 3p6 3d5 4s1, which is not half-filled or completely filled
• Oxygen (Z=8) has electronic configuration 1s2 2s2 2p4, which is not half-filled or completely filled
• Hund's rule


  The electrons tend to occupy orbitals singly with parallel spins as much as possible
• When all the orbitals are singly occupied, only then pairing of electrons commences
• In the ground state, the electrons will have their spins parallel
• Hund's rule suggests that electronic arrangement with maximum multiplicity is more stable than the one with minimum multiplicity
• Electronic configuration of ions

  The only difference is that when writing an electronic configuration for an ion, you have to remember to add or subtract electrons from the neutral atom's configuration
• The (n+l) rule is still followed when filling the orbitals for an ion
• When many orbitals have the same (n+l) value, electrons fill the orbital with the higher l value first
• (n+l) rule


  Used to determine the relative energy of the orbitals, the orbital with the lowest (n+l) value is the most stable
• Where two orbitals have the same (n+l) value, the more stable orbital is the one with the lower l value
• Quantum numbers n, l, m

  • n = 2, l = 1 can have 6 electrons
  • n = 3, l = 2 can have 10 electrons
  • n = 3, l = 3 is not possible
  • n = 4, l = 3 can have 14 electrons
• The maximum number of electrons an orbital can have is 2(2l+1)
• Paramagnetic substances are attracted to a magnetic field due to the presence of unpaired electrons
• Diamagnetic substances are weakly repelled from a magnetic field due to the absence of unpaired electrons
• Ferromagnetism overcomes diamagnetism when both occur in the same substance