Orbitals

Orbitals

1. Complex Model

   • Energy levels of principal quantum numbers n=1,2,3,4 etc. don't have precise energy values. Instead they each consist of set of subshells - contain orbitals with different energy levels.
   • Sublevels are of diff types labelled s,p,d + f. S sublevel has 1 orbital, p has 3, d has 5 + f has 7.
   • Electron orbital represents region of space around nucleus of atom, capable of accommodating 2 electrons. Therefore, s sublevel can accommodate 2 electrons, p - 6, d - 10 + f - 14.
   • Each orbital has its own approx, 3D shape. Not possible to draw shape of orbitals precisely. Don't have exact boundaries but are fuzzy, often called charge-clouds.

2. Spin-pairing

   • Shell n=1 consists of 1 'sublevel' containing single s orbital called 1s; n=2 consists of s + p orbitals in sublevels called 2s + 2p; n=3 consists s, p + d orbitals in sublevels called 3s, 3p + 3d.
   • Important principle concerning orbitals that affects all electron configurations. Theory that any individual orbital can hold 1/or 2 electrons but not more.
   • Along with charge, electrons have 'spin' property. Can visualise spin as electron rotating at fixed rate. 2 electrons can exist as pair in orbital by each having opposite spins - reduces effect of repulsion - why orbital can hold 2 electrons with neg charges that repel strongly.
   • Scientists now believe:
     • shell n=1 contains up to 2 electrons in s orbital.
     • shell n=2 contains up to 8 electrons, 2 in a orbital + 6 in p sublevel, with 2 in each of px, py + pz orbitals.
     • shell n=3 contains up to 18 electrons, 2 in s orbital, 6 in p sublevel + 10 in d sublevel, with 2 in each of 5 orbitals.

3. Order of filling shells + orbitals

   • Each successive element of Periodic Table, order of filling shells + orbitals is order of relative energy. Electronic configuration of each atom is 1 that gives as low energy state as possible to atom as whole.
   • Means lowest-energy orbitals filled first. Order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...
   • Expected order followed up to 3p sublevel, but there's variation, as 4s if filled before 3d. Variation + other variations further along in order caused by increasingly complex influences of nuclear attractions + electron repulsions upon individual electrons.

4. Filling Orbitals

   • Further Points
     • When 4s orbital is filled, next electron goes into 3d orbital. Begins pattern of filling up 3d sublevel, which finishes at zinc. Elements that add electrons to do sublevels are d-block elements; subset of these is transition elements.
     • Variations in pattern filling of d sublevel at elements 24 (chromium) + 29 (copper). Elements have only 1 electron in 4s orbital. Chromium has 5 d electrons, rather than expected 4; copper has 10 d rather than 9. Outcome of complex of complex interactions of attractions + repulsions in atoms.
     • From element 31 (gallium) to 36 (krypton) electrons add to 4p sublevel. Similar to pattern of filling 3p sublevel from elements 13 (aluminium) to 18 (argon) in Period 3.
   • Whenever possible, electrons will occupy orbitals singly - due to repulsion of neg charges. Electrons remain unpaired until avaiable orbitals of equal energy have 1 electron each.
   • When there're more electrons than orbitals can hold as singles - pair up by spin-pairing. If there are 3 electrons occupying p sublevel, 1 each will go to px, py, + pz orbitals rather than 2 in px, 1 in py + 0 in pz.
     • E.g. carbon (6 electrons) - 1s²2s²2p²