1. 1 2 3 4 5 6 7 3 4 5 6 4 5 n d (l=2) p (l=1) s (l=0) f (l=3)
Saunders 7.13
f (l=3)

2. 1s
n=1, l=0, ml = 0 (ms = ±(1/2) )
2py
n=2, l=1, ml =-1 (ms= ±(1/2) )
n=2, l=1, ml =1 (ms= ±(1/2) )
2px
2pz
n=2, l=1, ml =0 (ms= ±(1/2) )

3. Electron Configuration
Oxygen Orbital Notation Simplified Notation
1s s p
O s22s22p [He]2s22p4
Unpaired
Ground State
Spin Paired
Excited State [He]2s22p33s1
Anion - O [Ne]
Aufbau Principle-add electrons to give the lowest energy
Hund’s Rule - electrons fill all orbitals of a subshell before pairing
and they will have parallel spin
Pauli Exclusion Principle - no two electrons can have the same four
quantum numbers
pg. 216

4. paramagnetic-weakly attracted into a magnetic field due to unpaired 4s
saunders 8.3sb
for magnetism
Mn [Ar] 3d54s2
K [Ar] 4s1
K [Ar]
paramagnetic-weakly attracted into a magnetic field due to unpaired
electrons 3d 4s
Zn [Ar] 3d104s2
diamagnetic-very weakly repelled by a magnetic field due to all
electrons being spin coupled
Exceptions
Cr [Ar] 3d54s Cu [Ar] 3d104s1
pg. 224