Chapter 13:
Electrons in the
Atom
College Prep Chemistry
Orbital
Interactive
Evolution of Atomic Models
1. John Dalton
–
indestructible mass
 - no subatomic particles
Evolution of Atomic Models
2. J.J. Thomson – plum-pudding model
 Discovered
electrons
 electrons stuck into a lump of positively charged
material
 “mint chocolate chip ice cream model”
Evolution of Atomic Models
3. Ernest Rutherford
 nucleus
of the atom is positively charged
 Evidence – gold foil experiment
Evolution of Atomic Models
4. Niels Bohr
 electrons
travel in definite orbitals around the
nucleus
 Energy level = the region around the nucleus
where the electron is likely to be moving

Fixed energy levels analogous to the rungs of a
ladder
Evolution of Atomic Models
Quantum Mechanical Model
5.
Describes the probability of finding an electron in a
region of space around the nucleus



It is impossible to know the exact position and
momentum of an electron at the same time
Based on probability rather than certainty
Instead of traveling in defined orbits as Bohr
proposed, electrons actually travel in diffuse clouds
around the nucleus
Principal Energy Levels (n)
 Principal
(main) energy levels are assigned
numbers according to their energy



n=1, 2, 3, 4 …
Generally, energy increases with increasing n
Distance of the electron from the nucleus increases
with increasing n
Sublevels
 For
each principal energy level, there are one or
more sublevels
 s, p, d, f
 # of sublevels = the principal energy level (n)

For example,
 1st
principal energy level has 1 sublevel (s)
 2nd principal energy level has 2 sublevel (s,p)
 3rd principal energy level has 3 sublevels (s,p,d)
Sublevels and Orbitals
 Each
sublevel has a specific shape
 Each
sublevel houses a specific # of electron
orbitals (“where the electrons live”)
 Each
orbital can contain 2 electrons
Sublevels and Orbitals
 s-sublevel:



Spherical in shape
One orbital
Contains 2e- maximum (2e- per each s orbital)
Sublevels and Orbitals
 p-sublevel:



dumbbell in shape
3 orbitals
Holds 6 e- maximum (2e- per each p orbital)
Sublevels and Orbitals
 d-sublevel:



Double dumbbell in shape
5 orbitals
Holds 10 e- maximum (2e- per each d orbital)
Sublevels and Orbitals
 f-sublevel:



complex in shape
7 types of f-orbital
Holds 14 e- maximum in each orbital (2e- per each
f orbital)
Atomic Orbital Chart
Energy
# of
Level Sublevels
Sublevels
# of Orbitals in each Sublevel
1
3
5
n=1
1
s
n=2
2
s
p
n=3
n=4
3
s
p
d
4
s
p
d
Total # of
Electrons
7
2
8
18
f
32
Electron Arrangement in Atoms
Electron Configuration: the ways in which
electrons are arranged around the nuclei of
atoms


Gives information about principal energy levels,
sublevels, orbitals
Three Rules determine
electron configurations
1.
2.
3.
The Aufbau Principal
Hund’s Rule
Pauli Exclusion Principle
Rules for Electron Configurations
 Aufbau principle:


Electrons enter orbitals of
the lowest energy first
The s sublevel is always
the lowest in energy
Rules for Electron Configurations
1. Aufbau principle: Draw your own
Aufbau filling diagram
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f
6f
5f
4f
(the #4 spelled out has an f)
(we see in 3 dimensions!)
(two “peas” in a pod)
(the #7 spelled out has an s)
Rules for Electron Configurations
1. Aufbau principle: Draw your own
Aufbau filling diagram
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f
6f
5f
4f
Rules for Electron Configurations
2. Hund’s Rule: (“hogs don’t like each other”)
 Every
orbital in a sublevel is singly occupied
before any orbital is doubly occupied
 All
of the electrons in singly occupied orbitals
have the same spin
↑ ↑ ↑_
Rules for Electron Configurations
3. Pauli Exclusion Principle:
 an atomic orbital may have a
maximum of two electrons
 Two electrons that occupy the
same orbital must have opposite
spins
 designated with 
Orbital Notation Examples
Li
B
C
↑↓
1s
2s
↑↓
1s
↑↓
2s
↑
↑↓
1s
↑↓
2s
↑
2p
↑
2p
Hund’s Rule
Orbital Notation practice
 Elements
#1-20
 Use the Aufbau
Diagram Provided
 Remember all 3 rules!
Electron Configurations
A
shorthand way of identifying the location
of electrons
C
↑↓
1s
↑↓
2s
↑
C
1s2
2s2
2p2
↑
2p
Electron Configurations
Rearranging the e- configurations
 Scientists
rearrange the configurations so
that all similar energy levels stay together
#21 = Sc (fill with Aufbau first!)
1s22s22p63s23p64s23d1
Then Switch!
1s22s22p63s23p63d14s2
Noble Gas electron configurations
Noble
gas configurations – look on
the periodic table!
#11 = Na
 The
noble gas that precedes Na is Ne
1s22s22p6
So
instead of 1s22s22p63s1 use
[Ne]3s1
Orbital blocks on the Periodic
Table
Exceptions to electron
configurations
24 = Cr
#
1s22s22p63s23p64s23d4
↑↓ ↑ ↑ ↑ ↑ __
 Takes
less energy to half fill all
orbitals
1s22s22p63s23p64s13d5
↑_ ↑ ↑ ↑ ↑ ↑_
Exceptions to the electron
configurations
 Also
true for the rest of column 6 & 11
#29 = Cu
 Following

1s22s22p63s23p64s23d9
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_
 Actual

the rules
configuration
1s22s22p63s23p64s13d10
↑_ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
Section 13.3
Physics and the Quantum
Mechanical Model…
Electrons and Light
Back to Bohr
 Bohr’s
model is based on atomic emission
spectra
 Atoms only give off light of certain colors
(wavelengths)
Wave model of light
 Light
is a wave with a frequency, speed and
wavelength
 Emission of light is related to the behavior of
electrons in an atom
Parts of a Wave (1 of 2)
Origin
– center line
Wavelength (l) – distance
from crest to crest
Amplitude (A) – distance
from origin to crest
origin
amplitude
Parts of a Wave (2 of 2)
Frequency
SI
(f) - # of waves per second (s-1)
Unit (Hertz)
Hz = 1 s-1
inversely related to wavelength
1
Electromagnetic Radiation (EMR)
a
form of energy that exhibits wave-like behavior
as it travels through space


Includes radio waves, microwaves, infrared waves,
Visible light, ultraviolet waves, X-rays and gamma
rays
All waves travel at the speed of light
 3.0
X 108 m/s
Visible Light
A
prism separates sunlight into a spectrum of
colors
 Sunlight consists of a continuous range of colors


Each color has a specific wavelength and f
Red light = lowest f, longest wavelength
Electromagnetic Spectrum
Wave model of light
c=lxf
Speed
Wavelength
Frequency
m/s
m
Hz (hertz)
1/s
c = Speed of light = 3.00 x 108 m/s
Calculating Frequency
Determine the frequency of light
with a wavelength of 500 nm?
c=lxf
l = 500 nm =5 x 10-7 m
c = 3.00 x 108 m/s
3.00 x 108 m/s = (5 x 10-7 m) f
f = 6.00 x 1014 Hz
Calculating Wavelength
What
is the wavelength of the
yellow light emitted by a sodium
lamp if the frequency of the
radiation is 5.10 x 1014 s-1?
Sample Problem Answer
 Given:
f = 5.10 x 1014 s-1
c = 3.00 x 108 m/s
 Unknown: l
 Parent Equation: c = f x l
 Answer:
5.88 x 10-7 m = 588 nm
Your turn #1
What
is the wavelength of radiation
with a frequency of is 1.50 x 1013 s-1?
2.00 x 10 -5 m
Active Inspire
Your turn #2
What
is the frequency of radiation
with a wavelength of 5.00 x 10 -6 m?
6.00 x 10 13 s -1
Active Inspire
Physics and the Quantum
Mechanical Model
Max
Planck (1858-1947)
Discovered a direct relationship
between frequency and energy
 Higher
frequency = higher energy
E = h x f
h
= Planck’s Constant = 6.63 x 10-34 J·s
Sample Problem
What
is the energy of a
photon with a frequency of
5.00 x 1015 s-1?
Sample Problem Answer
 Given:
f = 5.00 x 1015 s-1
h = 6.63 x 10-34 J-s
 Unknown: E = ?
 Parent Equation: E = h x f
 Answer:
3.32 x 10-18 J
Your turn #1
What
is the energy of radiation with a
frequency of is 5.50 x 1014 Hz?
Active Inspire
Photoelectric Effect
 Photon
- electron in the form of light
 Ground state – normal position of electron
 Excited state – electron jumps to higher energy level
when energy is applied
 Photoelectric effect – release of energy in the form of
light when electron falls back to ground state


Color of light emitted depends on the energy level of the
electrons
Blue light has a higher energy than red light
Photoelectric Effect
Atomic Emission Spectrum
 Elements
emit light when electrocuted in gaseous
form
 The light is passed through a prism and an atomic
emission spectrum is obtained
 Atomic emission spectra are NOT continuous like
sunlight


Each line represents one distinct wavelength and
frequency
Every element has a unique emission spectra
Atomic Emission Spectrum
Spectrums
 When
a narrow beam of emitted light is shined
through a prism, it is separated into colors of the
visible SPECTRUM
Types of Spectrums
 Visible

light Spectrum:
Continuous range
 Atomic




emission spectrum
Light emitted by an element through a prism
Every element has a distinct emission spectrum
Discontinuous
Each line = one frequency
Different Elements Have Different Spectrums
Video of Spectra via Spectroscopes
 http://www.mhhe.com/physsci/chemistry/essent
ialchemistry/flash/linesp16.swf
 http://www.flinnsci.com/atomicspectrum
More practice problems
with energy, frequency,
and wavelength
 Last
page of your chapter 13 packet
Chapter 13 Review Problems
p.
386 #22, 27,
28, 33, 34, 36, 38,
45, 47, 51, 60, 63
Evolution of Atomic Models
5. Quantum Mechanical Model (1 of 2)
Quantum of energy = amount of energy required
to move an electron from one energy level to
the next higher level




Energy levels are not equally spaced
Levels get closer together the further from the
nucleus
The further away from the nucleus, the less energy
is required for an electron to escape