1. 5-1 Quantum Theory of the atom
Bohr Model of the Atom
Proposed that electrons orbited the nucleus in circular paths.
Ground state- lowest allowable energy states of an atom.
Excited state- atom gains energy; H atoms can have many different excited states although it contains 1 e-.
Electrons move around a H atom in circular orbit
Orbits equal to a principal quantum number n, where n=1 is lowest energy level, closest to nucleus.

2. Bohr model of the atom nucleus
Orbits/ levels are like rungs in step ladder
Cannot stand b/w rungs, e- can’t exist b/w levels (orbits).
E- move from 1 orbit to the next emitting or absorbing certain amounts of energy (quanta).
The smaller the e- orbit, the lower the energy state/level
The larger the e- orbit, the higher the energy state/level
n =1
n =2
n =3
n =4
n =5
n =6
nucleus

3. 5-1 Quantum Theory and the atom
Quantum mechanical model is the modern atomic model and comes from
Louis De Broglie: radiation (energy) behaves like particles and vice versa.
All particles w/ a mass have wave characteristics
E- move around nucleus in a wave-like manner
Heisenberg uncertainty principle- impossible to know both the velocity and position of an e- at the same time.
Shrodinger: e-’s energy are limited to certain values (quantum) but does not predict path
Treated e-’s as waves
Created wave function = predicts probability of finding e- in a volume of space (location)

4. Hydrogen’s Atomic Orbitals
Shrodinger’s wave eqn predicts atomic orbitals
Atomic orbital - 3D regions around the nucleus that describes the e-’s probable location.
atomic orbital = fuzzy cloud
Do not have a defined size
Shape = volume that contains 90% of the probable location of e-’s inside that region.

5. Quantum mechanical model
Like Bohr, electrons occupy space surrounding the nucleus and exist in several principal energy levels = principal quantum number (n)
Relative size and energies of atomic orbital
n = 1,2, 3, etc. = period
Principal energy levels consist of energy sublevels with different energy values.
Energy sublevels – shape of the atoms’ orbitals
s = spherical
p = dumbbell
d, f= different shapes

6. Quantum Mechanical model
Principal energy levels have specific allowed sublevels - shapes.
s sublevel is lower in energy and f has higher energy 1 2 3 4
n = s p d f s p d s p s

7. Quantum Mechanical Model
Sublevels consist of orbitals of different orientation.
Orbitals in same sublevel are = in energy (no matter orientation)
Orbitals only hold 2e- maximum with opposite spins (+ or – spins).
Sublevel Orientations/ Orbitals Max # e- s p d f

8. Orientations/ orbitals per sublevel
s- spherical only 1 orbital orientation
p- dumbbell has 3 orbital orientations
d- 2dumbbells with 5 orbital orientations
f- 3dumbbells with 7 orbital orientations
dex.html

9. Bohr model of the atom Hydrogen’s Line Spectrum (AES)
At n= 1 H atom is in ground state
When energy is added, e- moves to higher energy level, n=2 (excited state).
e- drop back to lower energy level n=1 and emitts a photon equal to the difference b/w levels.
A photon is emitted with E= hυ
A photon is absorbed

10. Hydrogen’s line spectrum
Lines which show up have specific energies which correspond to a frequency of a color of light.
A photon is emitted with E= hυ for each frequency
Energy of Hydrogen Atom 1 2 3 4 5 6 n
E= 4.85 x J
E= 3.03 x J

11. 5-2 Electron configurations
Electron configuration – arrangement of e- in atoms; lower nrg arrangements
Arrangements defined by:
Aufbau principle – e- occupy lowest nrg orbital available
All orbitals in a sublevel are = in nrg (px py pz )
Sublevels within an energy level have different energies
Ex: 2s lower in nrg than 2p
Order of energy = s, p, d, f
Sublevels in one energy level can overlap with sublevels in another principal energy level.
Ex: 4s lower in nrg than 3d

12. Aufbau diagram

13. Electron Configurations
Pauli exclusion principle – a max of 2 e- may occupy a single orbital only if they have opposite spins.
Hund’s rule – energy charged e- repel each other.
All same nrg orbitals are filled first with e- containing same spin before extra e- can occupy the same orbital with opposite spins.
Ex: 3 orbitals of 2p
2px
2py
2pz

14. Filling sublevels with electrons
Energy sublevels are filled from lower energy to higher energy following the diagram.
ALWAYS start at the beginning of each level and follow it until all e- in an element have been placed. 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
Increasing Energy

15. Orbital diagram and E- Configurations
Orbital diagram for Fe:
Iron has how many e- ?
26 e-
1s 2s p s p s d
Electron configuration for Fe:
Iron has 26 e-
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Shortcut to the E- config. for Fe is Noble gas notation
Group 18 or 8A are the Nobel Gases
Argon has 18 e-
Noble gas notation: ] [
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s2 3d6
[Ar] 4s2 3d6

16. Valence electrons and Electron dot structures
Valence electrons – outer energy level/orbital electrons which are involved in bonding.
Valence electrons = groups 1A to 8A
B groups do not count
E- dot structures- consists of the element’s:
Symbol - represents the atomic nucleus & inner-level electrons
Surrounded by dots- represent the valence electrons.
Ex: O = 1s2 2s2 2p4 or [He]2s2 2p4 ve- =6 in grp 6A O

17. Periodic table shortcut
Groups (A only) = Valence e-
Energy level = n-1 for d sublevel
Periods = Energy Level
Energy level =
n-2 for f sublevel

18. 5-3 Light and Quantized Energy
Some elements emit visible light when
heated with a flame.
This chemical behavior is due to the
arrangement of e- in atoms.

19. Electromagnetic radiation
Form of energy that exhibits wave-like behavior as it travels through space.
There are many types of electromagnetic radiation and all are represented in the electromagnetic spectrum

20. Electromagnetic spectrum

21. Parts of a wave
Frequency (v, nu) –The number of complete wavelengths that pass a given point each second.
Units: wave/second = 1/s = s-1 = Hertz (Hz)
Wavelength (l, lambda) – The distance between identical points on successive waves. (crest to crest or trough to trough)
Units: meters (m)
c = l v
c = speed of light, 3.00 x 108 m/s

22. Wave nature of Light
Max Planck theorized that all matter can gain/ lose energy in small “chunks” of light (quanta).
Quantum- minimum amt of energy that can be gained or lost by an atom.
Ex: Iron when hot appears red or blue, emits energy that is quantized has a specific frequency.
Heating water – temp increases by molecules absorbing a specific amt or quanta.
Calculated as follows:
Equantum= hv
E = Energy (J)
h = Planck’s constant x (J s)
v = frequency ( Hz or s-1)

23. Particle Nature of Light
Photoelectric effect – electrons are emitted from a metal’s surface when light of a specific frequency shines on the surface.
Albert Einstein (1905) assumed that light travelled as a stream of tiny particles or packets of energy called photons.
Photons- EM radiation w/ no mass that carries a quantum of energy.
EM radiation has both wave-
like and particle- like nature.
Ephoton= hv
Photon = quantum of energy

24. Atomic emission spectra
Set of frequencies of light waves emitted by an atom of an element.
Line spectrum – consists of several individual lines of color from light energy emitted by excited unstable atoms
Only certain colors (frequencies) appear in an element’s AES & it can be used to identify the element.