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Atomic Structure
Atoms and their structure
1
History of the atom
Not the history of atom, but the history of the
idea of the atom
 Original idea Ancient Greece (400 B.C..)
 Democritus and Leucippus - two Greek
philosophers

2
History of Atom
Looked at beach
 Made of sand
 Cut sand - smaller
sand

Smallest
possible
piece?
Atomos - not to be cut
3
Another Greek
Aristotle - Famous philosopher
 All substances are made of 4 elements
 Fire - hot
 Air - light
 Earth - cool, heavy
 Water - wet
 Blend these in different proportions to
get all substances

4
Who Was Right?








Greek society was slave based
It was beneath the famous to work with their
hands
They did not experiment
Greeks settled disagreements by argument
Aristotle was more famous
He won
His ideas carried through to the middle ages.
Alchemists therefore tried to change lead to
gold
5
Who’s Next?
Late 1700’s - John Dalton - a famous
English chemist conducted experiments
 Summarized results of his experiments
and those of other’s
 Where?
 In Dalton’s Atomic Theory
 Combined ideas of elements with that of
atoms

6
Dalton’s Atomic Theory
 All matter is made of tiny indivisible
particles called atoms.
 Atoms of the same element are identical,
those of different atoms are different.
 Atoms of different elements combine in
whole number ratios to form compounds
 Chemical reactions involve the
rearrangement of atoms. No new atoms
are created or destroyed.
7
Law of Definite Proportions
Each compound has a specific ratio of
elements
 It is a ratio by mass
 Water is always 8 grams of oxygen for
each gram of hydrogen

8
Law of Multiple Proportions

If two elements form more that one
compound, the ratio of the second
element that combines with 1 gram of
the first element in each is always a
simple whole number.
9
Parts of Atoms
J. J. Thomson - English physicist. 1897
 Made a piece of equipment called a
cathode ray tube.
 It is a vacuum tube - all the air has been
pumped out.

10
Thomson’s Experiment
Voltage source
-
+
Vacuum tube
Metal Disks
11
Thomson’s Experiment
Voltage source
-
+
12
Thomson’s Experiment
Voltage source
-
+
13
Thomson’s Experiment
Voltage source
-
+
14
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
15
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
16
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
17
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
18
Thomson’s Experiment
Voltage source

By adding an electric field
19
Thomson’s Experiment
Voltage source
+
 By adding an electric field
20
Thomson’s Experiment
Voltage source
+
 By adding an electric field
21
Thomson’s Experiment
Voltage source
+
 By adding an electric field
22
Thomson’s Experiment
Voltage source
+
 By adding an electric field
23
Thomson’s Experiment
Voltage source
+
 By adding an electric field
24
Thomson’s Experiment
Voltage source
+
 By adding an electric field he found that the
moving pieces were negative
25
J. J. Thomsom’s Model




Discovered the
electron
Couldn’t find positive
(for a while)
Said the atom was
like plum pudding
A bunch of positive
stuff, with the
electrons able to be
removed
26
Other pieces
Proton - positively charged pieces
1840 times heavier than the electron
 Neutron - no charge but the same mass
as a proton.
 Where are the pieces?

27
Rutherford’s experiment






Ernest Rutherford - another famous English
physicist. (1910)
He believed in the plum pudding model of the
atom.
He wanted to see how big the electrons
(plums) were.
Used radioactive Uranium.
Alpha particles - positively charged pieces
given off by uranium 2He4
Shot them at gold foil which can be made a
few atoms thick
28
Rutherford’s experiment
When the alpha particles hit a florescent
screen, it glows.
 Here’s what the set up of the
experiment looked like.

29
Lead
block
Uranium
Florescent
Screen
Gold Foil
30
He Expected
The alpha particles to pass through
without changing direction very much
Because
– The positive charges were spread out
evenly. Alone they were not enough
to stop the alpha particles
31
What he expected
32
Because
33
Because, he thought the mass was
evenly distributed in the atom
34
Because, he thought
the mass was evenly
distributed in the atom
35
What he got
36
Conclusions of Rutherford's Experiment:
1.
Atom is mostly empty
Because most particles passed straight through
2.
Small dense,
positive piece
at center
+
Because Alpha particles
are deflected by the
nucleus it if they get
close enough
37
+
38
Bohr Model
Planetary Model or Heliocentric Model
 Places the nucleus in the center of the
atom like the sun in the center of the
universe
 Places the electrons in orbitals revolving
around the nucleus, like the planets
revolve around the sun. (fixed paths)

39
40
41
Modern View
The atom is mostly
empty space
 Two regions
– Nucleus:

» protons and neutrons
– Electron cloud:
» region where you might find
an electron based on
mathematical probability.
42
Density and the Atom
Since most of the particles went
through, it was mostly empty.
 Because the pieces turned so much, the
positive pieces were heavy.
 Small volume, big mass, big density
 This small dense positive area is the
nucleus
 Scale: marble in a football field.

43
Subatomic particles
Relative Actual
mass (g)
Name Symbol Charge mass
Electron
e-
-1
1/1840 9.11 x 10-28
Proton
p+
+1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
44
Structure of the Atom

There are two regions
– The nucleus
» With protons and neutrons
» Positive charge
» Almost all the mass
– Electron cloud- Most of the volume of
an atom
» The region where the electron can be found
45
Size of an atom
Atoms are small.
 Measured in picometers, 10-12 meters
 Hydrogen atom, 32 pm radius
 Nucleus tiny compared to atom (marble in a football

field)

Radius of the nucleus near 10-15m.
46
Counting the Pieces
Atomic Number = number of protons
– # of protons determines kind of atom
– The # of protons is the same as the
number of electrons in a neutral atom
 Mass Number =
the number of protons + neutrons

47
End of Part One
48
Atomic Structure - Part Two
49
Symbols

Contain the symbol of the element, the
mass number and the atomic number
50
Symbols

Contain the symbol of the element, the
mass number and the atomic number
Mass
number
Atomic
number
X
51
Symbols

Find the
– number of protons
– number of neutrons
– number of electrons
– Atomic number
– Mass Number
19
9
F
52
Symbols
 Find
the
–number of protons
–number of neutrons
–number of electrons
–Atomic number
–Mass Number
80
35
Br
53
Symbols
 if
an element has an atomic
number of 34 and a mass number
of 78 what is the
–number of protons
–number of neutrons
–number of electrons
–Complete symbol
54
Se
Symbols
 if
an element has 91 protons and
140 neutrons what is the
–Atomic number
–Mass number
–number of electrons
–Complete symbol
55
Pa
Symbols
 if
an element has 78 electrons and
117 neutrons what is the
–Atomic number
–Mass number
–number of protons
–Complete symbol
56
Pt
Naming Isotopes
Put the mass number after the name of
the element
 carbon- 12
 carbon -14
 uranium-235

57
Isotopes
Dalton was wrong.
 Atoms of the same element can have
different numbers of neutrons
 Different mass numbers
 Called isotopes (do not confuse with
allotropes)

58
Atomic Mass

How heavy is an atom of oxygen?
– There are different kinds of oxygen atoms.
– More concerned with average atomic
mass.
– Weighted average based on the
abundances of all the naturally occurring
isotopes in nature.
– Don’t use grams because the numbers
would be too small.
59
Measuring Atomic Mass

Unit is the Atomic Mass Unit (amu)
– AMU is based on the C-12 atom. It is
one twelfth the mass of a carbon-12
atom.
– Each isotope has its own atomic
mass we need the average from
percent abundance.
60
It all averages out
Test
1
2
3
4
5
6
Avg. Grade
Student A
95
74
95
95
74
95
Student B
89
88
88
87
88
88
61
Test 6 Counts 50%
Student A
 95 x 10% =
 74 x 10% =
 95 x 10% =
 95 x 10% =
 74 x 10% =
 95 x 50% =
StudentA B
Student

9589x x.110%
= =

7488x x.110%
= =

9588x x.110%
= =
 95 x .1 =
 87 x 10% =
 74 x .1 =
 88 x 10% =
 95 x .5 =
=
__________  88 x 50%__________
__________
62
Atomic Mass

Calculate the atomic mass of copper if
copper has two isotopes. 69.1% has a mass
of 62.93 amu and the rest has a mass of
64.93 amu.
63
Atomic Mass

Magnesium has three isotopes. 78.99%
magnesium 24 with a mass of 23.9850
amu, 10.00% magnesium 25 with a mass of
24.9858 amu, and the rest magnesium 26
with a mass of 25.9826 amu. What is the
atomic mass of magnesium?

If not told otherwise, the mass of the
isotope is the mass number in amu
64
Atomic Mass
Is not a whole number because it is an
average.
 The decimal numbers on the periodic
table are based on the weighted
average of all the known naturally
occurring isotopes.

65
The End
66
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
67
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
68
Bohr’s Model
Increasing energy
Fifth
Fourth
Third
Second
First
Nucleus
Further away
from the
nucleus means
more energy.
 There is no “in
between”
energy
 Energy Levels

69
The Quantum Mechanical Model
Energy is quantized. It comes in chunks.
 Quanta - the amount of energy needed to
move from one energy level to another.
 Quantum leap in energy.
 Schrödinger derived an equation that
described the energy and position of the
electrons in an atom
 Treated electrons as waves

70
The Quantum Mechanical Model
a mathematical solution
 It is not like anything you can
see.

71
The Quantum Mechanical Model
Does have energy levels
for electrons.
 Orbits are not circular.
 It can only tell us the
probability of finding
an electron a certain distance from the
nucleus.

72
The Quantum Mechanical Model
The electron is found
inside a blurry
“electron cloud”
 An area where there
is a chance of finding
an electron.
 Draw a line at 90 %

73
Atomic Orbitals
Principal Quantum Number (n) = the
energy level of the electron.
 Within each energy level the complex
math of Schrödinger's equation
describes several shapes.
 These are called atomic orbitals.

– To calculate number of orbitals : n2
– To calculate maximum electrons per energy level:
2n2

Regions where there is a high
probability of finding an electron.
74
S orbitals
1 s orbital for every energy level
 Spherical
shaped

Each s orbital can hold 2 electrons
 Called the 1s, 2s, 3s, etc.. orbitals.

75
P orbitals
Start at the second energy level
 3 different directions
 3 different shapes (dumbell)
 Each can hold 2 electrons

76
P Orbitals
77
D orbitals
Start at the third energy level
 5 different shapes
 Each can hold 2 electrons

78
F orbitals
Start at the fourth energy level
 Have seven different shapes
 2 electrons per shape

79
F orbitals
Images
J mol
80
Summary
# of
shapes
Max
electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
81
Filling order
Lowest energy fill first.
 The energy levels overlap
 The orbitals do not fill up order of energy
level.
 Counting system
– Each box is an orbital shape
– Room for two electrons

83
Increasing energy
7s
6s
5s
7p
6p
6d
5p
5d
4s
4p
4d
3s
3p
3d
2s
2p
5f
4f
1s
84
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
6d
5d
4d
5f
4f
3d
3p
3s
2p
2s
1s
85
Electron Configurations
Aufbrau principle- electrons enter the
lowest energy first.
 Pauli Exclusion Principle- at most 2
electrons per orbital - different spins
 Heisenberg’s Principle- we do not know
the precise location of an electron only a
probability of where the electron is in the
orbital cloud.

86
Electron Configuration

Hund’s Rule- When electrons occupy
orbitals of equal energy they don’t pair
up until they have to . (up, up, up, then
down, down, down.)

Phosphorus electron config.
87
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2s
1s
4f
3d
4s
3p
5f
The first 2 electrons
go into the 1s orbital
2p
 Notice the opposite
spins
 only 13 more

88
Increasing energy
7s
6s
5s
7p
6p
6d
5d
5p
4d
4p
3s
2s
4f
3d
4s
3p
5f
The next electrons
go into the 2s orbital
2p
 only 11 more

1s
89
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more, but follow
Hund’s Rule – which is
NOT shown here.
90
Increasing energy
7s
6s
5s
4s
3s
2s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
1s
91
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into each
in one spin direction
(Hund’s Rule)
3 unpaired electrons
92
1s22s22p63s23p3
The easy way to remember
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
1s
• 2 electrons
93
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
1s 2s
• 4 electrons
94
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
• 12 electrons
95
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
3p 4s
• 20 electrons
96
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2
• 38 electrons
97
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
• 56 electrons
98
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
• 88 electrons
99
Fill from the bottom up following
the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
5f14 6d10 7p6
• 118 electrons
100
Rewrite when done
•
2
2
6
2
6
2
10
6
2
1s 2s 2p 3s 3p 4s 3d 4p 5s
4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10
7p6

Group the energy levels together
• 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10
4f14 5s2 5p6 5d105f146s2 6p6 6d10 7s2
6
7p
101
The Modern Table
Elements are still grouped by properties
 Similar properties are in the same column
 Order is in increasing atomic number
 Added a column of elements Mendeleev
didn’t know about.
 The noble gases weren’t found because
they didn’t react with anything.

102
Horizontal rows are called periods or
principal energy levels.
 There are 7 periods

103
Vertical columns are called groups.
 Elements are placed in columns by
similar properties.
 Also called families

104

1A
2A
The elements in the A groups
8A
0
are called the representative
3A 4A 5A 6A 7A
elements
105
Representative Elements
Groups IA-VIIIA
 S and p Block
 Obey the octet rule

106
VIIIA
VIIA
VIA
IIIA
IIB
13 14 15 16 17
3A 4A 5A 6A 7A
IB
VIIIB
VIIB
VIB
VA
IVB
IIIB
1 2
1A 2A
VA
IIA
IA
IUPAC System for 1-18 (International Union
of Pure and Applied Chemistry)
IVA
CAS System for A and B
(Chemical Abstract Service)
Other Systems
3 4 5 6 7 8 9 10 11 12
3B 4B 5B 6B 7B 8B 8B 8B 1B 2B
107
18
8A
Metals
108
Metals
Luster – shiny.
 Ductile – drawn into wires.
 Malleable – hammered into sheets.
 Conductors of heat and electricity.
 Why? Sea of “mobile” electrons.

109
Transition metals

The Group B
elements
110
Dull
 Brittle
 Nonconductors
- insulators

Non-metals
111
Metalloids or Semimetals
Properties of both
 Semiconductors

112

These are called the inner transition elements
and they belong here
113
114
Group 1A are the alkali metals
 Group 2A are the alkaline earth metals

115
Group 7A is called the Halogens
 Group 8A are the noble gases

116
Why?
The part of the atom another atom sees
is the electron cloud.
 More importantly the outside orbitals
 The orbitals fill up in a regular pattern
 The outside orbital electron configuration
repeats
 So.. the properties of atoms repeat.

117
H
Li
1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p63d104s24p65s1
1s22s22p63s23p63d104s24p64d105s2
5p66s1
1s22s22p63s23p63d104s24p64d104f145s2
5p65d106s26p67s1
118
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar18
1s22s22p63s23p63d104s24p6 Kr
36
1s22s22p63s23p63d104s24p64d105s25p6 Xe
54
1s22s22p63s23p63d104s24p64d105s24f14 Rn
5p65d106s26p6 86
119
S- block
s1
s2
Alkali metals all end in s1
 Alkaline earth metals all end in s2
 really have to include He but it fits
better later
 He has the properties of the noble
gases

120
Transition Metals -d block
s1 d5
d1
d2
d3
d5
d6
d7
d8
s1
d9
d10
121
The P-block
p1 p2
p3
p4
p5
p6
122
F - block

inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
123
1
2
3
4
5
6
7

Each row (or period) is the energy level
for s and p orbitals
124

d orbitals fill up after previous energy level
so first d is 3d even though it’s in row 4
1
2
3
4
5
6
7
3d
125
1
2
3
4
5
6
7
4f
5f

f orbitals start filling at 4f
126
Writing Electron
configurations the easy way
Yes there is a shorthand
127
Electron Configurations repeat
The shape of the periodic table is a
representation of this repetition.
 When we get to the end of the row the
outermost energy level is full.
 This is the basis for our shorthand

128
The Shorthand
Write the symbol of the noble gas
before the element in brackets [ ]
 Then the rest of the electrons.
 Aluminum - full configuration
 1s22s22p63s23p1
 Ne is 1s22s22p6
 so Al is [Ne] 3s23p1

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More examples
Ge = 1s22s22p63s23p63d104s24p2
2 10 2
 Ge = [Ar] 4s 3d 4p
 Ge = [Ar] 3d104s24p2
 Hf=1s22s22p63s23p64s23d104p64f14
4d105s25p65d26s2
 Hf=[Xe]6s24f145d2
 Hf=[Xe]4f145d26s2

130
The Shorthand
Sn- 50 electrons
The noble gas before it
is Kr
Takes care of 36
Next 5s2
Then 4d10
Finally 5p2
[ Kr ]
5s2 4d10
5p2
131
Electron configurations and groups
Representative elements have s and p
orbitals as last filled
– Group number = number of electrons
in highest energy level
 Transition metals- d orbitals
 Inner transition- f orbitals
 Noble gases s and p orbitals full

132