1. Chapter 2 Notes: electron configuration
The electron configuration of an atom can be deduced from its atomic number.
Chapter 2.2

2. Important Terms for this Section
Electromagnetic spectrum -speed -wavelength -frequency Spectrum Continuous spectrum Absorption spectrum Line/emission spectrum States excited state ground state Photon Quantized Planck equation Converge
Ionization energy
Hydrogen
Paschen series
Balmer series
Lyman series
Heisenberg’s uncertainty principle
Atomic orbitals
sub-levels
Aufbau principle
electron configuration
Hund’s third rule
Valence electrons
S, p, d, f blocks

3. Light as a wave
Visible light is only a small portion of the full light spectrum All light travels at the same speed, c: 3.00 x 108 m s-1 All electromagnetic waves have: frequency and wavelength Wavelength - distance between any two corresponding points on successive waves (m) Frequency - number of waves that pass a point in space during any time interval (s-1) c = λf

5. Continuous spectrum
Absorption or emission spectra

6. Absorption or emission spectra
When electromagnetic energy passes through prism: continuous spectrum When electromag E. pass through a sample of atoms: emission or absorption spectrum This is also seen when some elements are passed through a flame - rad. Excited some of the atoms to higher E level Absorption: light passed through a cold gas sample – energy absorbed by e- so they leave blank lines in the spectrum Emission: heated gas (voltage applied) to gas – energy released by heated gas so they produce lines in spectrum Each element has a specific spectrum fingerprint

7. Evidence of bohr model
Packet of energy that is emitted when excited e- falls back down to ground state is called a PHOTON Energy of photon is proportional to the frequency of the radiation (what goes in as EM radiation, must come out as photon rad): ΔEe- = Ephoton Also frequency is related to E of photon: Ephoton = hv where h is Planck’s constant (in data booklet) Therefore, ΔEe- = hv There are steps to the energy levels that can be occupied by the electrons = quantized (discrete amounts)

8. Hydrogen spectrum
When e- falls from high to low energy, emission of photon
The energy of the drop determines what part of spectrum:
Drop from higher to 3 = infrared
Drop from higher to 2 = visible
Drop from higher to 1 = ultraviolet
Lines converge at higher energies
because energy levels are closer
At n = ∞ the atom is ionized
Problems with Bohr model:
Spectral lines of other atoms
Trajectory of e- is not known

9. Wave and particle models
Wave model - Diffraction of light Particle model - Photon emission of light Both models are needed to describe the nature of light

10. Heisenberg’s uncertainty principle
If we were to use energy to locate the precise location of an electron: - we would influence the electron and it would travel off in a random direction Because of this, we cannot know the true location of an electron at any given moment. BUT!... We can have a probability of where the electron is most likely to be!

11. SchrÖdinger model of hydrogen atom
Atomic Orbitals!!!! These are the regions around an atomic nucleus where there is a 90% probability of finding an electron with that particular energy. The shape of the orbital is also dependent on the energy of the electron The higher the energy, the farther away the electron is likely to be found

12. Atomic orbitals
1st orbital is spherical and holds 2 e- 2nd orbital has a sphere and 3 “dumbells” that can hold up to 8 e- - these are sub-levels: - sphere = 2e- - “dumbells” hold 2 e- each = 6 e- total Energy levels are said to be degenerate if it corresponds to two or more different measurable states of a quantum system - p sub-level = degenerate - have same shape and energy - but are in different spatial orientation

13. Pauli exclusion principle and sub-levels of electrons
States that no more than 2 e- can occupy any one orbital, and if there are 2 e- in one orbital, they must spin in opposite directions. Each orbital can hold 2 e- Each electron has an opposite spin – they must spin in opposite directions or they would be the same electron (mutual repulsion)

14. Aufbau principle: orbital diagrams
Aufbau is German for “building up” and refers to the orbital diagrams and the relative increase in energy moving up the diagram
This demonstrates Pauli Exclusion Principle and Hund’s Third Rule: Hund’s Rule: - Orbitals will fill 1 e- at a time (so e- have parallel spin)... This results in lower energy - ...Before additional electrons fill the orbitals (for the opposite spin)

15. Electron configuration using periodic table

16. Electron configurations
The relative energy of the orbitals is dependent on the atomic number (and valence electrons)
- Starts to show separation when there are more protons and electrons added
There are some special notes to make about configurations:
The 3d sub-level is lower than 4s for elements after Ca and should be written before the 4s level
Chromium family has a special electron configuration:
...d5 s1
Copper family has a special electron configuration:
...d10 s1

17. Electron configurations
Electron configuration for Hydrogen-1: Electron configuration for Nitrogen-7: Electron configuration for Krypton-36: Electron configuration for Silver-47: Short hand electron configuration for Cobalt-27: Short hand electron configuration for Silver-47: Short hand electron configuration for Tungsten-74: Short hand electron configuration for Iridium-77: Short hand electron configuration for Fermium-100:

18. Electron configuration for ions
Electrons are removed from the highest energy level or added to the next available orbital Therefore, transition metals: the electrons are removed from the highest level (s) before the next d level. Electron configuration for Ti2+: Electron configuration for N3-: Electron configuration for Br1-: Electron configuration for Ag1+: Short hand electron configuration for Co2+: Short hand electron configuration for Ag1+: Short hand electron configuration for Pb4+: