Chapter 13: Electrons in Atoms
Waves
Properties() Frequency-number of waves that
pass a given point per unit of time.
() Wavelength- distance between
similar points in a set of waves (crest to
crest.)
Properties of waves, cont.
Amplitude- distance from crest or
trough to the normal
Energy- Waves do not carry energy,
they transmit energy. The amount of
energy determines the amplitude and
the frequency.
Energy vs. Amplitude
Energy and amplitude are DIRECTLY
related. As energy increases, amplitude
increases. As energy decreases,
amplitude decreases
Review waves at home:
http://www.glenbrook.k12.il.us/gbssci/p
hys/Class/waves/u10l2a.html
Speed of the wave
Speed = frequency x wavelength= 
S= as frequency increases wavelength
decreases
inverse relationship- what does inverse mean?
c = Speed of light 3.00 x 108 m/s
(in a vacuum)
Practice Problems
 What is the wavelength of a radio wave
whose frequency is 1.01 x 108 Hz? Remember
scientific notation-here it is again…
Speed = wavelength x frequency
3.0 x 10 m
8

s
8
1.01 x 10 Hz

speed

  2.97 m
Practice Problems
 What is the frequency of a green light which
has a wavelength of 4.90 x 10-7 m?

Speed = wavelength x frequency
3.0 x 10 m
8
s

-7
4.9 x 10 m
speed

3.0
8  ( 7 )

 10
Hz
4.9
  6.12 x 10 Hz
14
Practice Problems
An X-ray has a wavelength of 1.15 x 10-10
m. What is its frequency?
Check your answer!
Light as Particles
 quantum – term coined by Max Planck in the
early 1900’s; it is the minimum amount of
energy that can be gained or lost by an atom.
1. Planck’s constant (h) relates energy and
frequency.
2. h = 6.626 x 10-34 Js
3. E = h
4. Frequency and energy are DIRECTLY related.
Electromagnetic Spectrum
Page 120
Frequency and energy have a direct relationship.
About Light
 Light- Characteristics of light waves
Visible light when bent (refracted) gives
different colors, each color corresponds to a
different frequency. (pg 120)


ROYGBIV
(E)
(E)
 Light travels in bundles called photons.
 Energy = Planck's constant x frequency
 E=hv (unit for energy is the joule)
Photon
 photon – a particle of electromagnetic
radiation with no mass that carries a
quantum of energy; a bundle of energy; a
stream of tiny particles
 the idea of a photon was developed when
Einstein, in 1905, said that electromagnetic
radiation has both wavelike and particlelike
natures
Quantum leap
 When an electron is excited, by heat or
electricity, it absorbs energy and jumps to a
higher energy level. When the electron
jumps down to a lower energy level it gives
off energy in the form of light. Each jump is
a different color.
 The colors given off are called the bright line
spectra; it acts as a finger print for the
element.
 Remember the flame tests we did?
Atomic Spectra
http://www.physics.lsa.umich.edu/demolab/graphics2/7b10_u2a.jpg
http://www.physics.lsa.umich.edu/demolab/graphics2/7b10_u2b.jpg
Flame tests
 Show colors unique to each element
http://www.unit5.org/christjs/flam
e%20tests.htm
Ba2+
greenish-yellow
Ca2+
brick-red
Cu2+
green
Sr2+
crimson
Na1+
yellow
K1+
Violet
Bohr Models
 Because of the quantum leap Bohr theorized
that the electron traveled in different energy
levels.
Bohr video Must watch at school!
Try this video!
A bohr diagram gives us a fair idea about
where electrons are in their clouds
The shells are labeled K, L, and M
K has 2 e-, L has 8 e-, M has 18 e-
Bohr Diagrams
Try this practice site
Draw a diagram of Carbon, Magnesium,
and Silicon
Modern Concept:
Quantum Mechanical Model
Bohr’s Model had two problems:
Electrons are always moving
Tests to “see” electrons move electrons
Bohr's theory only works well in explaining
one or two electrons.
Modern theory is based on...
Origins of Quantum Model
1. de Broglie- Dual nature of matter which says
that that light can act like a particle and that
particles can act like light waves.
2. Heisenberg's uncertainty principle - It is
impossible to know both the speed and location
of an electron at the same time.
3. Schodinger wave equation- predicts the
location of the electron most of the time. Dealt
with the probability of finding a blurry electron
cloud which he called an orbital.
4. Orbital- a region of space in which the electron
can be found 90% of the time.
Quantum numbers
 From Schrodinger's equation each
electron is assigned a set of quantum
numbers to distinguish each from all
other electrons (address)
no 2 electrons have the same set of
quantum numbers
 Pauli's Exclusion Principle Electrons in the same orbital must have
opposite spin. (an orbital will hold only
____ electrons.)
http://nobelprize.org/physics/laureates/1933/schrodinger-bio.html
Quantum Numbers Decoded
"Never say, 'I tried it once and it did not work.'" -- Ernest Rutherford
Energy Level
Quantum #n
Sublevel type
quantum # l
s sublevel
Spherical shaped
1 orbital
Sublevel p
dumbell shaped
3 orbitals
Sublevel d
Sublevel f
5 orbitals
Remember an orbital can only contain 2 electrons!
7 orbitals
Pictures of Orbital shapes
P orbitals (2
of3)
S orbital
D orbitals (1 of
5)
http://www.colorado.edu/physics/2000/elements_as_atoms/orbitals.html
Labeling Electrons
 Orbital filling diagram and electron
configuration. Both methods use the following
rules.
 Aufbau Principle- fill lowest energy level first
 Hund's Rule- fill equal orbitals before pairing
up.
 http://intro.chem.okstate.edu/WorkshopFolder/El
ectronconfnew.html
 http://tinyurl.com/9v6zu
Electron Configurations
The Electron Dorm
Always follow the order
1s 2s 2p
3s 3p 4s 3d 4p
5s 4d 5p
6s 4f 5d 6p
7s 5f 6d 7p
Orbital Diagrams
and Electron Configurations
Let’s Practice
Do the electron Configurations of elements
11-19
Lewis Dot Structures
The dots around the atom’s symbol
represent the electrons in the atoms
outermost energy level.
There will be no more than 8
Note how these dot structures follow the
main group elements (A columns) of the
periodic table.
Lewis Dot Structures
IA
VA
See your teacher to learn what the
IIA
dot structures look VIA
like or look them
up in your book!
VIIA
IIIA
VIIIA
IV A
Elements with fewer than 4 electrons in the outer
shell tend to lose electrons whereas elements
with more than 4 outershell electrons tend to
gain electrons!
Practice Problem 1:
 Write the orbital filling diagram
 Write the electron
configuration
 Write the noble gas notation.
 How many electrons in the outer shell
 In what period and group does this element appear on
the periodic table?
 Metal, nonmetal or metalloid?
 Electron dot structure?
 Gain or lose electrons?
 Highest energy level filled?
80
35
Br
Practice Problem 2:
 Write the orbital filling diagram
 Write the electron
configuration
 Write the noble gas notation.
 How many electrons in the outer shell
 In what period and group does this element appear on
the periodic table?
 Metal, nonmetal or metalloid?
 Electron dot structure?
 Gain or lose electrons?
 Highest energy level filled?
128
52
Te
Practice Problem 3:
 Write the orbital filling diagram
 Write the electron
configuration
 Write the noble gas notation.
 How many electrons in the outer shell
 In what period and group does this element appear on
the periodic table?
 Metal, nonmetal or metalloid?
 Electron dot structure?
 Gain or lose electrons?
 Highest energy level filled?
133
55
Cs
Answer to Practice Problem
 An X-ray has a wavelength of 1.15 x 10-10 m.
What is its frequency?
3.0 x 10 m
8
18
s

 2.73 x 10 Hz
-10
1.10 x 10 Hz