Electron arrangements

Electron arrangements

1. Energy levels

   • Electrons in different shells have differing amounts of energy
   • Shells are called main energy levels
     • Divided into sub-levels
   • S sublevel
     • Can hold 2 electrons
     • Spherical
   • P sublevels
     • Can hold 6 electrons
     • 3 dumbell shape
   • D sublevel
     • Can hold 10 electrons
     • 5 dumbell shape
   • F sublevel

2. Quantum mechanics

   • Developed 1920's
   • Schrodinger equation
   • Give the probability of finding an electron in a given volume of space called an atomic orbital

3. Atomic orbitals

   • Electron
     • Cloud of negative charge
     • Fills volume known as its atomic orbital
   • Concept
     • Different orbitals = different energies
       • Number tells us which energy level it corresponds to
     • Atomic orbitals of each main level have different shapes which have slightly different energies
     • Shapes of orbitals represent a volume of space in which there is a 95% probability of finding an electron
     • First main energy level consists of just one s orbital, the second has an s and a p,
     • Any atomic orbital can hold 2 electrons
       • S orbitals can hold up to two electrons
       • P orbitals can hold 2 each, but come in groups of 3
       • D orbitals can hold 2 each, but come in groups of 5

4. Spin

   • Electrons have a property called spin
   • 2 electrons in the same orbital must have opposite spin
     • This creates a magnetic attraction between the two, allowing them to sit in the orbital without repulsion
   • Electrons are usually represented by arrows pointing up or down to show different directions of spin
   • Repulsion
     • Has to be some sort of attraction to hold e- in an orbital

5. Order of filling

   • Aufbau principle
     • Lower energy orbitals filled first
     • No atomic orbital can hold more than 2 electrons
     • 4s fills before 3d
   • Hunts rule
     • FIll singly before pairing
     • Opposite number of spins minimised