• Electron clouds are 3D, not flat
• Electrons are spread out as much as possible,
not usually round, and are moving rapidly
• These things are hard to see in a still picture
• Electrons spread out in orbitals
• Orbitals have different SHAPES and ENERGY
(distance from nucleus)
• Quantum numbers describe orbitals
• There are four quantum numbers
• First of the four (n, #, #, #)
• Describes the distance from the nucleus to the
orbital (and therefore describes the ENERGY of
the orbital)
• Values: integers ≥ 1
• As the distance from the nucleus increases, n
increases.
• As the energy increases, n increases.
• Each period as a unique n value
• Period 1: n=1
• Period 2: n=2
•Period 3: n=3
• ETC
• The transition from n≥2 to n=1in a hydrogen
atom
• Result: ultraviolet emission lines of the hydrogen
atom
• Greater the difference in the principal quantum
numbers, the higher the energy of the
electromagnetic emission
• The transition from n≥3 to n=2 in a hydrogen atom
• Result: spectral line emissions of the hydrogen atom
• As the n value increases, the wavelength emitted
decreases (in nm)
Starting n value
3
4
5
6
7
8
Wavelength (nm)
656.3
486.1
434.1
410.2
397.0
388.9
Color
Red
Blue
Violet
Violet
Ultraviolet Ultraviolet
• The transition from n≥4 to n=3 in a
hydrogen atom
• Result: emission lines in the infrared band
• Second of the four (n, l, #, #)
• Shape of the sublevel
• Range from 0 to n-1 (we will never deal with
anything above l=3)
• l=0 = s
• l=1 = p
• l=2 = d
• l=3 = f
• Third of the four numbers (#, #, ml , #)
• Denotes the orbital sublevel that is filled
• s sublevel has ONE orbital (sphere has one orientation
in space)
• p sublevel has THREE orbitals (three orientations in
space)
• d sublevel has FIVE orbitals (five orientations in space)
• f sublevel has SEVEN orbitals (seven orientations in
space)
• s sublevel: one orbital
• One orientation in space
• P sublevels: three orbitals
• Three orientations in space
• d sublevels: five orbitals
• Five orientations in space
• f sublevel: seven orbitals
• Seven orientations in space
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Integers from -l to l
SO:
s: ml = 0 only since l= 0
p: ml = -1,0,1 since l= 1
d: ml = -2,-1,0,1,2 since l= 2
f: ml = -3,-2,-1,0,1,2,3 since l= 3
• Fourth number (#, #, #, ms)
• It is either -1/2 or ½
• Down or up
• There is a very specific order in which electrons
fill orbitals. It is not random. There are some
exceptions.
1. Aufbau (next) Principle
2. Pauli Exclusion Principle
3. Hund’s Rule
• Electrons fill the LOWEST energy
sublevel before going to the next
sublevel
• 1s fills, then 2s fills, then 2p fills, then
3s fills, then 3p fills ….
• Electrons pair according to OPPOSITE
spins
• ↑↓, not ↑↑ or ↓↓
• Electrons spread out in equal energy sublevels
before pairing electrons
• (↑ ↑ ↑ and not ↑↓ ↑ _)
• First level to fill is 1s level
• Lowest energy sublevel
• Holds two electrons
•They are oppositely paired
• A sublevel is represented by __ and holds 2
electrons
• Second sublevel is the 2s sublevel
• It holds 2 electrons because of s
• Electrons are oppositely paired
• So, we filled 1s, we filled 2s
• Now comes 2p
• Holds six electrons because p orbitals hold 6
electrons 1s22s22p6
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From 2p, 3s fills with 2 electrons
3p fills with 6 electrons
4s fills with 2 electrons
3d fills with 10 electrons
4p fills with 6 electrons
5s fills with 2 electrons
• Neutral carbon: 6 electrons
↑↓ ↑↓ ↑ ↑ _
1s 2s 2p
• Six arrows for six electrons
• 1s2 2s2 2p2
• Some energy levels are SUPER close together
• The levels are so close that electrons are able to move between
these orbitals in order to minimize repulsion…
• 4s and 3d orbitals are very close in energy
• Exceptions exist for some period 4 d block elements
• Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4
• Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5
• It actually takes LESS energy to split the electrons between the
5 sublevels than it does to put them together in the 4s and 3d