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Chapter 4
“Atomic Structure”
Section 4.1 Defining the Atom

OBJECTIVES:
 Describe Democritus’s ideas about atoms.
 Explain Dalton’s atomic theory.
 Identify what instrument is used to observe
individual atoms.
Section 4.1 Defining the Atom
_
 First to suggest the existence of atoms
(from the Greek word “atomos”)

 He
believed that atoms were
_
.
Dalton’s Atomic Theory
1)
are
composed of tiny
called
.
John Dalton
(1766 – 1844)
2) Atoms of the
.
--Atoms of any
_ are
are
from those of
any
.
Dalton’s Atomic Theory
3) Atoms of
elements
in
_
ratios to form
.
4) In chemical reactions, atoms are
– but
of
_
into atoms
.
Sizing up the Atom
100,000,000 atoms = 1 cm
1,000,000 atoms = width of hair
Can be observed with
_
.
Section 4.2
Structure of the Nuclear Atom
 OBJECTIVES:
Identify
three types of
subatomic particles.
Describe the structure of
atoms, according to the
Rutherford atomic model.
Section 4.2
Structure of the Nuclear Atom
 Atoms



are divisible into
particles:
_
_
_
_
Discovery of the Electron
used a
to discover the
_
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
determined the
of the
the mass of a
_
atom
Conclusions from the Study
of the Electron:
a) Atoms have
must be
the
, so there
to
_ of
the electrons
b)
that
for
have so little
_
must account
of the
_
Conclusions from the Study
of the Electron:
 Eugen Goldstein observed

_
(or 1840 times
that of an electron)
 J. Chadwick confirmed the
_

_
nearly
to a
Subatomic Particles
Particle
Charge
Mass (g)
Electron
(e-)
-1
9.11 x 10-28
Proton
(p+)
_
1.67 x 10-24
Neutron
(no)
_
1.67 x 10-24
Location
_
_
Thomson’s Atomic Model
J. J. Thomson
Thomson -
.
were like
embedded in a
charged
.
Ernest Rutherford’s
Gold Foil Experiment - 1911
particles (
fired at a thin
 Particles that
are
)

.
on the detecting
_
Rutherford’s Findings
 Most
A
of the particles
particles were
.
Conclusions:
a) The
is
,
.
,
_
The Rutherford Atomic Model

Based on his experimental evidence:
• Atom is mostly
the
•
the
, and
is in the center at the
.
all
.
The Rutherford Atomic Model
•
is made of
and
•
• Called the “
.
the nucleus.
”
Section 4.3
Distinguishing Among Atoms
OBJECTIVES:
 Explain what makes elements and isotopes
different from each other.
 Calculate the number of neutrons in an atom.
 Calculate the atomic mass of an element.

 Explain
table.
why chemists use the periodic
Atomic Number

Atoms are composed of
protons, neutrons, and electrons
•
How then are atoms of one
from
element?
Atomic Number

Elements are
contain different

in the nucleus (smaller #)
# protons = # electrons

because they
of
_
_
Atomic Number:
_
# p+ :
# e- :
_
Atomic Number:
_
# p+ :
# e- :
_
_
_
Atomic Number
of an element is
the
of
_
in the nucleus of each atom of that
element.
Element
Carbon (C)
# of protons
_
Atomic # (Z)
_
Phosphorus (P)
_
_
Gold (Au)
_
_
Mass Number
and
nucleus of an isotope:
Mass # =
is the number of
in the
_
Atomic Number:
_ Mass Number:
# p+ :
# e- :
_
_ #n0 :
_
_
Atomic Number: _
Mass Number:
# p+ : _
# e- : _ #n0 : _
_
Mass Number Practice
Atom
p+
Oxygen
8
n0
e- Mass #
_
_
_
74
Arsenic
_
41
_
Phosphorus
_
_
15
_
Complete Symbols

Contain the symbol of the element,
the mass number and the atomic
number.
Superscript →
Subscript →
X
Symbols

Find each of these:
a) number of protons _
b) number of
11
_
neutrons
23
c) number of
_
electrons
_
d) Atomic number
e) Mass Number _
Na
Symbols

If an element has an atomic
number of 34 and a mass
number of 78, what is the:
a) number of protons =
b) number of neutrons =
c) number of electrons =
d) complete symbol
Symbols
 If an element has 91
protons and 140 neutrons
what is the
a) Atomic number =
b) Mass number =
c) number of electrons =
d) complete symbol
Symbols
 If an element has 78
electrons and 117 neutrons
what is the
a) Atomic number
b) Mass number
c) number of protons
d) complete symbol
Isotopes




Dalton was
about
of the
type being
.
Atoms of the same element
numbers of
Thus, different
These are called
.
Atomic #:
Mass #:
# p+:
#n0:
Atomic #:
Mass #:
# p+:
#n0:
_
have
.
.
Atomic #:
Mass #:
# p+:
#n0:
Isotopes

due to
are atoms of the
with
numbers of
,
.
Naming Isotopes
 We
can also put the mass
number after the name of the
element:
carbon-12
• carbon-14
• uranium-235
•
Mass:
Mass:
Mass:
are atoms of the same
element having different masses, due to varying
numbers of neutrons.
Isotope
Proton
s
Hydrogen–1
Electrons
Neutrons
(protium)
Hydrogen-2
(deuterium)
1
1
0
1
1
1
Hydrogen-3
(tritium)
1
1
2
What’s the only thing that changes?
Nucleus
Atomic Mass



How heavy is an atom of oxygen?

- there are different
of oxygen atoms.
We want the
.
Based on
_
of
of that element in nature.
Measuring Atomic Mass

Measure atomic mass with the
_
the mass of a

Defined as

atom.
Each isotope has its own atomic
thus we determine the
,
from
.
To calculate the average:
the atomic
of each isotope by it’s
(expressed as a
then
the results.
 Expressed as


C-12 =
.
_
),
Atomic Masses
Atomic mass is the average of all the
naturally occurring isotopes of that element.
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition of
the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
Carbon = 12.011
% in nature
98.89%
1.11%
<0.01%
Atomic Mass Example

B-10 = 19.8%
B-11 = 80.2%

At. Mass =

The Periodic Table:
A Preview

elements in which the
are
based on a set of
Allows easy
- arrangement of
_
into
.
of the
of different elements
The Periodic Table:
A Preview

row
(there are 7 of them)

column
Also called a
_
Elements in a
have similar
and physical
.
Identified with number and “A” or “B”
Draw an arrow and label a period and a group.
Chapter 5
“Electrons in Atoms”
Section 5.1
Models of the Atom

OBJECTIVES:
• Identify the inadequacies in the Rutherford atomic
model.
• Identify the new proposal in the Bohr model of the
atom.
• Describe the energies and positions of electrons
according to the quantum mechanical model.
 Describe how the shapes of orbitals related to different
sublevels differ.
Ernest Rutherford’s Model



Discovered
_
“nucleus”
Electrons surround and
nucleus
 Like planets around the sun
Atom is mostly
_
Ernest Rutherford’s Model

Did not explain

of elements
Better description of the
was needed
Niels Bohr’s Model
 Electrons
move in
circular
,
paths, or
at different levels.
 Amount of fixed energy,
(
separates one level from
another.
 Electrons can
from one level to another.
,)
Bohr’s model

analogous to the rungs of a ladder
 Electrons
exist
energy levels, just like you can’t
stand between rungs on a ladder
Bohr’s model

–
amount of energy required to
an electron
from one energy level to another
 The
away
from the
,
the more
the electron has.
The Quantum Mechanical
Model

In 1926, Erwin Schrodinger derived the

Determines the
of an
States the
of
certain distance from the

an electron a
.
The Quantum Mechanical
Model



The atom is found
inside a blurry
“
An area where there is a
of
finding an electron.
Think of fan blades
”
Atomic Orbitals

regions where there is a


of finding an electron.
- arranged in sections:
 letters
Each
corresponds to a
different
.
–
Principal Quantum Number
“
”- it denotes the
in which the electron is located.
~
Maximum number of electrons that can
fit in an energy level is:
_
How many e- in level 2? 3?
level
# of shapes
(orbitals)
Maximum electrons
Starts at energy level
s
1
2
1
p
3
d
5
6
10
2
f
7
14
4
3
Number of sublevels due to
number of different shapes of orbitals
_
_
_
_
By Energy Level




First Energy Level
Has
orbital
only
electrons
1s2



Second Energy Level
Has
orbitals
Electrons:
2 in s, 6 in p


8 total electrons
By Energy Level





Third energy level
Has
_
orbitals
Electrons: 2 in s, 6 in p,
and 10 in d
_
18 total electrons





Fourth energy level
Has
_
orbitals
Electrons: 2 in s, 6 in p,
10 in d, and 14 in f
_
32 total electrons
Electron Configuration
Sublevel
# of Orbitals
Available
# of Electrons
Available
s
p
d
1
3
5
2
6
10
f
7
14
d=
f=
Section 5.2
Electron Arrangement in Atoms

OBJECTIVES:
• Describe how to write the electron
configuration for an atom.
• Explain why the actual electron
configurations for some elements differ from
those predicted by the aufbau principle.
7p
7s
6s
6p
5p
6d
5f
5d
4f
4d
Increasing energy
5s
4p
3d
4s
3p
3s
2p
2s
aufbau diagram - page 133
1s
Aufbau is German for “building up”
Electron Configurations…

…are the way
are
in various
orbitals around the nuclei of atoms. Three
rules tell us how:
Rule #1 - Aufbau Principle

Electrons must
the orbital with the
energy

_
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Rule #2 - Pauli Exclusion Principle


Orbitals can only have
The 2 electrons must have
electrons max
spins

Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Rule #3 - Hund’s Rule


Orbitals of
energy are each
by one
electron
any pairing occurs
Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
7p
7s
6s
6d
5f
6p
5d
4f
5p
4d
Increasing energy
5s
4p
3d
4s
3p
2p
2s
1s
Elec. Conf. of P?
 The first two electrons go
into the 1s orbital
Notice the opposite direction
of the spins
 only 13 more to go...

3s
7p
7s
6s
6d
5f
6p
5d
4f
5p
4d
Increasing energy
5s
4p
3d
4s
3p
3s
2p

2s

1s
The next electrons go
into the 2s orbital
only 11 more...
7p
7s
6s
6p
5p
6d
5f
5d
4f
4d
Increasing energy
5s
4p
3d
4s
3p
3s
2p
2s
1s
• The next electrons go
into the 2p orbital
• only 5 more...
7p
7s
6s
6p
5p
6d
5f
5d
4f
4d
Increasing energy
5s
4p
3d
4s
3p
3s
2p
2s
1s
• The next electrons go
into the 3s orbital
• only 3 more...
7p
7s
6s
6p
5p
6d
5f
5d
4f
4d
Increasing energy
5s
4p
3d
4s
3p
3s
2p
2s
1s
Orbital
notation
• The last three electrons go
into the 3p orbitals.
They each go into separate
shapes (Hund’s)
• 3 unpaired electrons
= 1s22s22p63s23p3
Orbitals fill in an order
energy to
energy.
 Adding electrons can change the energy of the
orbital.

are the absolute best situation.
Orbitals fill in an order
 However,
•
•
orbitals have a lower energy, and are
next
.
Makes them more
.
Changes the filling order
Practice Problems
Write electron configurations for the following atoms
1.
Li
5. P
2.
N
6. Si
3.
Be
7. Mg
4.
C
8. Al
Electron Configurations can be written in terms of
noble gases
To save space, configurations can be written in terms of
noble gases
Elec. Conf. for S? Look at noble gas before it.
Example 1: Ne = 1s22s22p6
S = 1s22s22p63s23p4
Or
S=
[Ne] 3s23p4

Elec.
Conf. for Mn? Look at noble gas before it.
Example 2: Ar = 1s22s22p63s23p6
Mn = 1s22s22p63s23p64s23d5
Mn =
[Ar]
4s23d5
Write the electron configurations
for these elements:

Titanium - 22 electrons


Vanadium - 23 electrons


Chromium - 24 electrons

 But
this is not what happens!!
Chromium is actually:
1s22s22p63s23p6
 Why?
 This gives us two

orbitals
(the others are all still full)
 Half full is slightly
in energy.
 The same principal applies to copper.
Copper’s electron configuration





Copper has 29 electrons so we expect:
1s22s22p63s23p6
But the actual configuration is:
1s22s22p63s23p6
This change gives one more filled orbital and one
that is half filled.
Remember these exceptions:
Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make
its 3d sublevel HALF FULL
Copper steals a 4s electron to
FILL its 3d sublevel
s = spin

When an electron moves, it generates a
.

s describes the direction an

They must spin in

Spin= up
down

There are two values of s:
directions
Section 5.3
Physics and the Quantum
Mechanical Model

OBJECTIVES:
• Describe the relationship between the wavelength and
frequency of light.
• Identify the source of atomic emission spectra.
• Explain how the frequencies of emitted light are related
to changes in electron energies.
• Distinguish between quantum mechanics and classical
mechanics.
- Page 139
“R O
Frequency Increases
Wavelength Longer
Y
G
B I V”
Parts of a wave
Crest
Origin
____________
Trough
_______
Energy
______
Energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
__________
__________
Frequency
Frequency
___________
___________
Wavelength
Wavelength
Visible Light
Wavelength
=
Frequency
=
ENERGY
Wavelength
=
Frequency
=
ENERGY
_
_
_
_
_
_
Wavelength Table
Atomic Spectrum

•
•
–
the discrete
representing the
frequencies of
emitted by an element
to each
element, like fingerprints!
Very useful for identifying
elements

When these
electrons return
to their
energy levels,
they lose
energy by
emitting
Changing the energy

Let’s look at a hydrogen atom, with only
, and in the first energy
level.
Changing the energy

Heat, electricity, or light can move the electron
up to different energy levels. The electron is
now said to be
“
”
Changing the energy

As the electron
back to the
ground state, it gives the energy back as
Changing the energy


Fall down in
Each step has a different
steps
Ultraviolet

The
more energy is
and the higher
the
Visible
Infrared
they fall,
.