Chem Quantum Numbers

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Quantum Set

[3,1,0,+1/2] Describes the possible locations of e-

Principle Quantum Number

[3,1,0,+1/2] # → is energy level ‘n’ (1,2,3,4, → 7) can not be 0

QM Model

[3,1,0,+1/2] 2nd Quan # → L, describes the shape of sublevel

l=0 → s (spherical)

l=1 → p (dumbbell shaped)

l=2 → d (clover shaped)

l=3 → f (complex shape)

Magnetic Quantum Number

[3,1,0,+1/2] What orbital the e- is in the sublevel

s → 0

p → -1,0,+1

d → -2,-1,0,+1,+2

f → -3,-2,-1,0,+1,+2,+3

Electron Spin

[3,1,0,+1/2 ] shows what spin the electron is spinning, can only be -1/2 or +1/2

Orbital Occupancy

Each orbital can only have two electrons. To fill, they need to have opposite spin.

Energy Levels

n - 1 = s → 2e-

n - 2 = s, p → 8e-

n - 3 = s, p, d → 18e-

n - 4 = s, p, d, f → 32e-

4-7 = 32e-

Written Sublevel Occupancy

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, ect.

Diamagnetism

All electrons are pared, there is no net spin. All electron spin gets canciled out by eachother.

Paramagnetism

Electron spins are randomly oriented

Ferromagnetism

Electron spins are parallel

Ferrimagnetism

Electron spins are parallel and an opposite directions, but don't cancel each other out.

Aufbau’s Rule

Electrons fill atomic orbitals from the lowest energy level to the highest

Pauli Exclusion Rule

States that no two electrons in an atom can have identical spins

Hund’s Rule

electrons in a sub level will singly occupy each orbital with parallel spins before any orbital is doubly occupied