1. CH 5 ELECTRONS IN
ATOMS
5-1 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.

2. 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

3. ELECTROMAGNETIC SPECTRUM

4. 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 (λ, lambda) – The distance between
identical points on successive waves. (crest to
crest or trough to trough)
 Units: meters (m)
c = λ v
c = speed of light, 3.00 x 108
m/s

5. 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.
o Ex: Iron when hot appears red or blue, emits
nrg that is quantized has a specific frequency.
o Heating water – temp increases by molecules
absorbing a specific amt or quanta.
 Calculated as follows:
Equantum= hv
o E = Energy (J)
o h = Planck’s constant 6.626 x 10-34
(J s)
o v = frequency ( Hz or s-1
)

6. 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

7. 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.

8. 5-2 QUANTUM THEORY OF THE
ATOM
 Bohr Model of the Atom
 Used to explain why AES was set of discontinuous
lines of specific frequencies (color).
 Proposed that Hydrogen atoms have only certain
allowable energy states based on Planck’s and
Einstein’s quantized energy.
 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 nrg level, closest to nucleus.

9. BOHR MODEL OF THE ATOM
 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 amts of nrg (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

10. BOHR MODEL OF THE ATOM
 Hydrogen’s Line Spectrum (AES)
 At n= 1 H atom is in ground state
 When nrg 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
absorbed
A photon is emitted
with E= hυ

11. HYDROGEN’S LINE SPECTRUM
 Lines which show up have specific energies
which correspond to a frequency of a color of
light.
EnergyofHydrogenAtom
1
2
3
4
5
6
n
A photon is
emitted with
E= hυ for each
frequency
E= 4.85 x 10-19
J
E= 3.03 x 10-19
J

12. 5-2 QUANTUM THEORY AND THE
ATOM
 Quantum mechanical model is the modern
atomic model and comes from
A. Louis De Broglie: radiation (energy) behaves like
particles and vice versa.
1. All particles w/ a mass have wave characteristics
2. E- move around nucleus in a wave-like manner
B. Heisenberg uncertainty principle- impossible
to know both the velocity and position of an e- at
the same time.
C. Shrodinger: e-’s energy are limited to certain
values (quantum) but does not predict path
1. Treated e-’s as waves
2. Created wave function = predicts probability of finding
e- in a volume of space (location)
Moving
Electron
Photon
Before
Electron
velocity changes
Photon
wavelength
changes
After

13.  Shrodinger’s wave eqn predicts atomic orbitals
 Atomic orbital - 3D regions around the nucleus
that describes the e-’s probable location.
a. atomic orbital = fuzzy cloud
b. Do not have a defined size
c. Shape = volume that contains 90% of the probable
location of e-’s inside that region.
HYDROGEN’S ATOMIC ORBITALS

14. QUANTUM MECHANICAL MODEL
 Like Bohr, electrons occupy space surrounding
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 nrg levels consist of energy sublevels
with different nrg values.
 Energy sublevels – shape of the atoms’ orbitals
s = spherical
p = dumbbell
d, f= different shapes

15. 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
s p
s p d
s p d f

16. 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 1 2
p 3 6
d 5 10
f 7 14

17. 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
 http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html

18. 5-3 ELECTRON CONFIGURATIONS
 Electron configuration – arrangement of e- in
atoms; lower nrg arrangements
 Arrangements defined by:
1. Aufbau principle – e- occupy lowest nrg orbital
available
a. All orbitals in a sublevel are = in nrg (px py pz )
b. Sublevels within an energy level have different energies
 Ex: 2s lower in nrg than 2p
a. Order of energy = s, p, d, f
b. Sublevels in one energy level can overlap with sublevels
in another principal energy level.
a. Ex: 4s lower in nrg than 3d

19. AUFBAU DIAGRAM

20. ELECTRON CONFIGURATIONS
2. Pauli exclusion principle – a max of 2 e-
may occupy a single orbital only if they have
opposite spins.
3. 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

21. 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

22.  Orbital diagram for Fe:
 Iron has how many e- ?
 26 e-
1s 2s 2p 3s 3p 4s 3d
 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-
 Iron has 26 e-
 Noble gas notation:
1s2
2s2
2p6
3s2
3p6
ORBITAL DIAGRAM AND E-
CONFIGURATIONS
][1s2
2s2
2p6
3s2
3p6
4s2
3d6
[Ar] 4s2
3d6

23. 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:
a. Symbol - represents the atomic nucleus & inner-level
electrons
b. Surrounded by dots- represent the valence electrons.
c. Ex: O = 1s2
2s2
2p4
or [He]2s2
2p4
ve- =6 in grp 6A
O

24. PERIODIC TABLE SHORTCUT
Periods=EnergyLevel
Groups (A only) = Valence e-
1A
2A 3A 4A 5A 6A 7A
8A
Energy level = n-1 for d sublevel
Energy level =
n-2 for f sublevel