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Lesson 5
Atomic Theory
Anything in black letters = write it in
your notes (‘knowts’)
Section 1 – Defining the Atom
Section 2 – The Spectrum of Light
Section 3 – Development of Modern Atomic Theory
Section 4 – The Nucleus
Section 5 – Electrons in Atoms
5.3 – Development of Modern Atomic Theory
Wilhelm Roentgen
Discovered x-rays in 1895.
Crooke’s Tube
Roentgen noticed a glow on a nearby fluorescent screen when the Crooke’s
tube was on. Even with the tube covered in black cardboard, the glow was still
present. No known ray of light could do this – these unknown rays were called
x-rays; x for unknown.
First x-ray taken
When roentgen placed his hand between the crookes’ tube and the
fluorescent screen, he saw the bones of his own hands.
Henri Becquerel
Discovered radioactivity in 1896.
Becquerel was studying phosphorescence in uranium containing
compounds. He noticed that when these compounds were wrapped in
black paper and left in the dark they still marked photographic plates. This
indicated that the uranium compound was giving off energy at all times.
3 Types of Radioactivity
Alpha () – helium nucleus; heavy and slow
Beta () – electron; fast and light
Gamma () – high energy light wave
J. J. Thomson
Discovered the electron in 1897
Thomson performed experiments that involved passing electric current
through gases at low pressure.
The result was a glowing beam, or cathode ray, that traveled from the cathode to
the anode.
Thomson found that a cathode ray is deflected by electrically charged metal plates.
Thompson knew that opposite charges attract and like charges repel, so he
hypothesized that a cathode ray is a stream of tiny negatively charged
particles moving at high speed; now called electrons.
To test his hypothesis,
Thompson set up an
experiment to measure
the ratio of an electron’s
charge to its mass.
Also, the charge-to-mass
ratio of electrons did not
depend on the kind of gas
in the cathode-ray tube or
the type of metal used for
the electrodes.
A cathode ray can also be deflected by a magnet.
The discovery of the electron led to the “Plum
Pudding” model of the atom.
Democritus
Dalton
Thomson’s Plum
Pudding Model
In 1911, Ernest Rutherford and others performed
the Gold Foil Experiment to test the plum pudding
model
Ernest
Rutherford
The Gold Foil Experiment
The results…
It was expected that alpha
particles would pass through
the plum pudding model of the
gold atom undisturbed.
Expected
It was observed that a small
portion of the alpha particles
were deflected, indicating a
small, concentrated positive
charge (the nucleus!)
Actual
“It was quite the most incredible event that has ever happened to
me in my life. It was almost as incredible as if you fired a 15-inch
shell at a piece of tissue paper and it came back and hit you. On
consideration, I realized that this scattering backward must be the
result of a single collision, and when I made calculations I saw that
it was impossible to get anything of that order of magnitude unless
you took a system in which the greater part of the mass of the atom
was concentrated in a minute nucleus. It was then that I had the
idea of an atom with a minute massive center, carrying a charge.”
— Ernest Rutherford
Nucleus – tiny positively charge core of an atom
Rutherford’s Nuclear
Model of the Atom
Is this the current
model of the atom?
NO…
• If an atom were the size of a football stadium,
the nucleus would be about the size of a marble
Proton – positively charged particle in the
nucleus of all atoms.
Neutron – particle with no charge in the nucleus
of all atoms except 1H
Properties of Subatomic Particles
Particle
Symbol
Relative
charge
Relative mass
(mass of proton = 1)
Actual mass
(g)
Electron
e–
1–
1/1840
9.11  10–28
Proton
p+
1+
1
1.67  10–24
Neutron
n0
0
1
1.67  10–24
Lesson 5 Quick Quiz
1. Identify the idea or discovery these people are credited with.
Democritus, Dalton, Roentgen, Becquerel, Thomson, Rutherford
2. Place the above ideas or discoveries in chronological order.
3. Describe the setup of the gold foil experiment. Also describe the
expected and actual results of this experiment. What conclusion was
made about the atom from the gold-foil experiment?
4. Name the three subatomic particles. Also give their symbol and
charge. What particles are present in the nucleus of the atom?
5. What is an atom?
5.4 – The Nucleus
Atomic Number (Z) The number of protons in an atom; identifies
the element.
Atoms have no net charge: # p+ = # e-
Mass Number (A) The number of protons (p+) plus the number
of neutrons (n0) in an atom.
The mass number is NOT the atomic mass.
Element
H
O
Ca
Atomic
Protons
Number (Z)
(p+)
Electrons
(e-)
Neutrons
(n0)
1
1
1
???
8
8
8
???
20
20
20
???
The number of n0 depends on the
mass number of the isotope
Isotopes Atoms of the same element that have a different
number of neutrons.
Chemical Symbols for Isotopes
A is the superscript
20
10
Ne
21
10
Ne
22
10
Z is the subscript
Ne
Determining the Composition of an Atom
How many protons, electrons, and neutrons
are in each atom?
9
4
a. Be
p+
en0
20
10
b. Ne
23
11
c. Na
Naturally Occurring Isotopes of Neon
20
10
Ne
21
10
Ne
22
10
Ne
Percent Abundance in Nature
90.48%
0.27%
9.25%
The masses of atoms are rarely expressed in grams.
The C-12 isotope has been given a mass of exactly
12 atomic mass units (amu)
The masses of all other elements are based on the
mass of the C-12 isotope.
Atomic Mass –
Weighted average of all the naturally occurring
isotopes of the element.
12
6
C
13
6
C
12.000 amu
13.003 amu
98.93 %
1.07 %
14
6
C
14.003 amu
0.0000000001 %
(12 x 0.9893)  (13.003 x 0.0107)  12.011
Atomic Mass of Carbon = 12.011 amu
12.000 amu
13.003 amu
98.93 %
1.07 %
14.003 amu
0.0000000001 %
Atomic Mass of Carbon = 12.011 amu
No atom of carbon actually weighs 12.011 amu. But
a typical carbon atom averages 12.011 amu.
Atomic masses are weighted averages.
There are 2 stable isotopes of silver
Silver-107; 106.905097 amu; 51.84%
Silver-109; 108.904752 amu; 48.16%
Calculate the atomic weight of silver.
Atomic Weight of Silver = 107.868 amu
Weighted Averages are NOT just for Atomic Masses
Your Chemistry grade is calculated as a weighted average.
50%
20%
15%
15%
(0.50) Quizzes/Tests,
(0.20) Homework,
(0.15) Lab Reports,
(0.15) Quarter Project,
4/10 = 40%
50/50 = 100%
70/75 = 93.3%
45/50 = 90%
Calculate the un-weighted average grade
169/185 = 91.35%
Calculate the weighted average grade
67.50%
Mass Spectrometer – separates isotopes by
mass differences.
Mass Spectrum for Cadmium
Mass Spectrum for Zinc; Atomic Weight = 65.395
Despite differences in the number of neutrons,
isotopes of an element are chemically similar.
Neutrons do not determine chemical reactivity; the
electrons do.
Iso
Example…
64Zn
Calculate
65Zn
The average mass of zinc
66Zn
67Zn
68Zn
70Zn
The weighted average mass 72Zn
of zinc
NA (%)
48.6
syn
27.9
4.1
18.8
0.6
syn
Mass
63.929
64.929
65.926
66.927
67.925
69.925
71.927
Terms to Know for Section 3 & 4
Atom
Atomic Number (Z)
Mass Number (A)
Proton, Neutron, Electron
Nucleus
Isotope
Atomic Mass (a.m.u.)
Weighted Average
Natural Abundance
Things to know for Section 3 & 4
Roentgen, Becquerel, Thomson, Rutherford
3 types of radiation (, , )
Explain Gold Foil Experiment (expected and actual
results, conclusion)
How to determine the number of p+, n0, eHow to write isotope symbols (Sodium-22 or 22Na)
How to calculate weighted averages
Basically know how to do the Practice Quiz!
5.5 – Electrons in Atoms
Rutherford’s Planetary Model of the Atom
Electrons moving around a tiny nucleus
Problems with Rutherford’s Model
1.Did not explain the chemical properties
of the elements.
2.Did not explain atomic spectra (…later…)
3.e- would spiral into the nucleus, but they
don’t
Review…..
The Electromagnetic Spectrum
Low energy
( = 700 nm)
Frequency  (s-1)
3 x 106
102
High energy
( = 380 nm)
3 x 1012
3 x 1022
10-8
10-14
White light produces a rainbow of colors.
Screen
Light
bulb
Slit
Prism
A prism separates light into the colors it contains.
Light from a helium lamp produces discrete
lines.
Screen
Slit
Helium
lamp
Prism
The lines that result are unique for each element
and are called its atomic emission spectrum.
For years, nobody had an explanation for
atomic spectra…
The Bohr Model ~1913
Electrons are found only in specific
locations (or energy levels) around the
nucleus.
To move from one energy
level to another, an e- must
gain or lose a quantum of
energy.
Niels Bohr
The energy levels in atoms are
unequally spaced, like the
rungs in this unusual ladder.
The higher energy levels are
closer together.
Bohr’s Model explained the emission spectra
When an atom absorbs energy an electron
jumps to a higher energy level (excited
state).
The electron returns to the lower energy level,
emitting a photon with a definite energy.
The photon’s energy shows up as a line in the
emission spectrum.
The Quantum Mechanical Model
The modern description of e- in atoms.
Similar to Bohr Model except the exact
location of an electron is impossible to know
Electrons are likely to be
found in electron ‘clouds’
around the nucleus
We will use the Bohr
model in this class…
Electron
Atomic orbital –
Most probable place for e- to be.
Orbitals can hold 2 e- maximum.
The orbitals are named s, p, d & f
S (1 type)
p (3 types)
d (5 types)
There are 7 types of f orbitals
Don’t worry about these shapes…
Orbital
Name
s
p
d
f
Types of
Maximum Electron
Orbital
Capacity
x2
2
1
3
x2
6
x2
5
10
x2
7
14
Each orbital can hold 2
e- maximum
Summary of Principal Energy Levels and Sublevels
Energy
Level
Type of Orbitals in Energy
Level
Maximum Number of
Electrons in Energy
Level
n=1
1s (1 orbital)
2
n=2
2s (1 orbital), 2p (3 orbitals)
8
n=3
3s (1 orbital), 3p (3 orbitals),
3d (5 orbitals)
18
n=4
4s (1 orbital), 4p (3 orbitals),
4d (5 orbitals), 4f (7 orbitals)
32
Aufbau Principle –
e- occupy the orbitals of lowest energy
first.
Aufbau Diagram
6p
5d
6s
5p
4d
5s
Increasing energy
4f
4p
3d
4s
3p
3s
2p
2s
E- fill the lowest energy orbitals first
Notice the 4s fills before the 3d
1s
Another Aufbau Diagram
(write this one down!)
Orbital
Maximum Number of
e- Due to Orbitals
s
p
d
f
2
6
10
14
Aufbauprinzip,
(german)
Electron Configuration –
shows how e- are arranged in an atom
Example: Write the electron configuration for N
How many e-?
7
Use Aufbau Diagram
2
2
3
1s 2s 2p
# of e- in
orbitals
Your turn…
Write the electron configuration for
a) boron
b) silicon
c) sulfur
Quick QUIZ!!
LEFT
RIGHT
1. What is the maximum number of electrons that can go
into
p orbital
d orbital
2. What is the maximum number of electrons in the
p orbitals
d orbitals
3. Write the electron configuration for
N
Mg
Mg
K
Cl
C
Can you see how the periodic table can be used as an
Aufbau Diagram?
Lesson 5 Quick Quiz
1. Explain the main difference between the Bohr Model and th
Quantum Model of the atom.
2. An atomic orbital can hold a maximum of _____ electrons.
3. How many types of s, p, d, and f orbitals are there?
s = ______,
p = _______, d = ______, f = ______
4. What is the maximum e- capacity of the
s orbitals ____, p orbitals ____, d orbitals ____, f orbitals
____
5. Write the electron configuration for the following elements
a) Helium (Z = 2)
b) Strontium (Z = 38)
c) Aluminum (Z = 13)
d) Chlorine (Z=17)
e) Silver (Z = 47)
f) Arsenic (Z = 33)
5. Write the electron configuration for the following elements.
a) Helium (Z = 2)
b) Strontium (Z = 38)
c) Aluminum (Z = 13)
d) Chlorine (Z=17)
e) Silver (Z = 47)
f) Arsenic (Z = 33)
Rutherford  Bohr  Quantum Mechanical Models,
Energy Levels, Atomic Orbitals (s, p, d, f) ,
Aufbau Principle & Diagram,
Electron Configurations,
Explanation of Atomic Emission Spectra
Big Sale at the Nuclear Mall today!
Make sure to get the closest
parking stall.
Nuclear
Mall
1s
2p
2s
2p
3d
3p
3s
3p
3d
4d
4p
4s
4p
4d