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.

Electron Configurations

Review of an Atom

This isn’t an accurate depiction.

• 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

Orbitals

• Electrons spread out in orbitals• Orbitals have different SHAPES and

ENERGY (distance from nucleus)• Quantum numbers describe orbitals• There are four quantum numbers

Principal Quantum Number (n)

• 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.

Principal Quantum Number

n relates to the Periodic Table!

• Each period as a unique n value•Period 1: n=1•Period 2: n=2•Period 3: n=3•ETC

Lyman Series

• 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

Lyman Series

Balmer Series• 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

Balmer Series

Paschen Series

• The transition from n≥4 to n=3 in a hydrogen atom• Result: emission lines in the infrared band

Paschen Series

Angular Momentum (l)

• 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

s sublevel

p sublevel

d sublevel

f orbitals

Magnetic Number (ml)

• 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)

Magnetic Number

• s sublevel: one orbital • One orientation in space

Magnetic Number

• P sublevels: three orbitals• Three orientations in space

Magnetic Number

• d sublevels: five orbitals• Five orientations in space

Magnetic Number

• f sublevel: seven orbitals• Seven orientations in

space

Magnetic Number

• 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

Spin (ms)

• Fourth number (#, #, #, ms)

• It is either -1/2 or ½• Down or up

What does all of this tell us?

• There is a very specific order in which electrons fill orbitals. It is not random.

There are some exceptions.

Orbital Filling Order

Three Important Principles

1.Aufbau (next) Principle2.Pauli Exclusion Principle

3.Hund’s Rule

Aufbau

• 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 ….

Pauli Exclusion Principle

• Electrons pair according to OPPOSITE spins•↑↓, not ↑↑ or ↓↓

Hund’s Rule

• Electrons spread out in equal energy sublevels before pairing electrons• (↑ ↑ ↑ and not ↑↓ ↑ _)

Step by Step…

• 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

Finally..

• 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

Example: Carbon

• Neutral carbon: 6 electrons↑↓ ↑↓ ↑ ↑ _1s 2s 2p

• Six arrows for six electrons• 1s2 2s2 2p2

Exceptions…• 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