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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
Democritus
 First to suggest the existence of atoms
(from the Greek word “atomos”)

 He
believed that atoms were
indivisible and
indestructible.
Dalton’s Atomic Theory
1) All elements are composed of
tiny indivisible particles
called atoms.
John Dalton
(1766 – 1844)
2) Atoms of the same element are
identical.
--Atoms of any one element are different
from those of any other element.
Dalton’s Atomic Theory
3) Atoms of different elements
combine in whole-number
ratios to form compounds.
4) In chemical reactions, atoms are combined,
separated, or rearranged – but never
changed into atoms of another element.
Sizing up the Atom
100,000,000 atoms = 1 cm
1,000,000 atoms = width of hair
Can be observed with scanning
tunneling (electron) microscopes
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 three
subatomic particles:
 Electrons
 Protons
 Neutrons
Discovery of the Electron
J.J. Thomson used a cathode ray tube to
discover the negatively charged electron
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
Robert Millikan determined the mass of the
electron: 1/1840 the mass of a hydrogen
atom
Conclusions from the Study
of the Electron:
a) Atoms have no charge, so there must be
positive particles to balance the negative
charge of the electrons
b) Electrons have so little mass that other
particles must account for most of the
mass
Conclusions from the Study
of the Electron:
 Eugen Goldstein observed positive
proton
 Mass of 1 (or 1840 times that of an
electron)
 James Chadwick confirmed the neutral
neutron
 Mass nearly equal to a proton
Subatomic Particles
Particle
Charge
Mass (g)
Location
Electron
(e-)
-1
9.11 x 10-28
Electron
cloud
Proton
(p+)
+1
1.67 x 10-24
Nucleus
Neutron
(no)
0
1.67 x 10-24
Nucleus
Thomson’s Atomic Model
J. J. Thomson
Thomson - plum pudding model.
Electrons were like plums embedded in
a positively charged pudding.
Ernest Rutherford’s
Gold Foil Experiment - 1911
 Alpha
particles (helium nuclei) fired at
a thin gold foil.
 Particles that hit on the detecting
screen are recorded
Rutherford’s Findings
 Most
of the particles passed right through
 A few particles were deflected.
Conclusions:
a) The nucleus is small,
dense, and, positively
charged
The Rutherford Atomic Model

Based on his experimental evidence:
• Atom is mostly empty space.
•
All the positive charge, and almost all the
mass is in the center at the nucleus.
The Rutherford Atomic Model
• Nucleus is made of protons and neutrons
• Electrons surround the nucleus.
• Called the “nuclear model”
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 identical
protons, neutrons, and electrons
•
How then are atoms of one element
different from another element?
Atomic Number



Elements are different because they contain
different numbers of PROTONS
Atomic number - number of protons in the
nucleus (smaller #)
# protons = # electrons
Atomic Number: 35
# p+ : 35
# e- : 35
Atomic Number: 53
# p+ : 53
# e- : 53
Atomic Number
Atomic number (Z) of an element is the
number of protons in the nucleus of
each atom of that element.
Element
# of protons
Atomic # (Z)
Carbon (C)
6
Phosphorus (P)
15
15
Gold (Au)
79
79
6
Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an isotope:
Mass # = p+ + n0
Atomic Number: 35
Mass Number: 79.9
# p+ : 35
# e- : 35 #n0 : 45
Atomic Number: 53
Mass Number: 127
# p+ : 53
# e- : 53 #n0 : 74
Mass Number Practice
Atom
p+
n0
e- Mass #
Oxygen
8
8
8
Arsenic
33
Phosphorus
15
41
16
16
33
74
15
31
Complete Symbols
Contain the symbol of the element,
the mass number and the atomic
number.
Atomic
Superscript →
number

Subscript →
X
Mass
number
Symbols

Find each of these:
a) number of protons 11
b) number of
11
12
neutrons
23
c) number of
11
electrons
d) Atomic number 11
e) Mass Number 23
Na
Symbols

If an element has an atomic
number of 34 and a mass
number of 78, what is the:
a) number of protons = 34
b) number of neutrons = 43
c) number of electrons = 34
d) complete symbol 34 X
78
Symbols
 If an element has 91
protons and 140 neutrons
what is the
a) Atomic number = 91
b) Mass number = 131
c) number of electrons = 91
d) complete symbol
Symbols
 If an element has 78
electrons and 117 neutrons
what is the
78
a) Atomic number
195
b) Mass number
c) number of protons 78
d) complete symbol
Isotopes
Dalton was wrong about all elements of
the same type being identical
 Atoms of the same element can have
different numbers of neutrons.
 Thus, different mass numbers.
 These are called isotopes.

Atomic #: 6
Mass #: 12
# p+: 6
#n0: 6
Atomic #: 6
Mass #: 13
# p+: 6
#n0: 7
Atomic #: 6
Mass #: 14
# p+: 6
#n0: 8
Isotopes

Isotopes are atoms of the
same element with different masses,
due to varying numbers of neutrons.
Naming Isotopes
 We
can also put the mass
number after the name of the
element:
carbon-12
• carbon-14
• uranium-235
•
Mass: 12
Mass: 14
Mass: 235
Isotopes 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
Nucleus
What’s the only thing that changes? # of neutrons
Atomic Mass



How heavy is an atom of oxygen?
 Depends - there are different
masses of oxygen atoms.
We want the average atomic mass.
Based on abundance (percentage)
of each variety of that element in nature.
Measuring Atomic Mass



Measure atomic mass with the
Atomic Mass Unit (amu)
Defined as one-twelfth the mass of a
carbon-12 atom.
Each isotope has its own atomic mass, thus
we determine the average from
percent abundance.
To calculate the average:
Multiply the atomic mass of each
isotope by it’s abundance (expressed as
a decimal), then add the results.
 Expressed as amu.


C-12 = 12 amu.
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 =

(10.0)(.198) + (11.0)(.802) = 10.8 amu
The Periodic Table:
A Preview
 Periodic table - arrangement of elements in
which the elements are separated into
groups based on a set of
repeating properties.
Allows easy comparison of the
properties of different elements
The Periodic Table:
A Preview
 Period - horizontal row
(there are 7 of them)
Group - vertical column
Also called a family
Elements in a group have similar
chemical and physical properties
Identified with number and “A” or “B”
Draw an arrow and label a period and a group.
Group
Period
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 dense
positive “nucleus”
Electrons surround and
orbit nucleus
 Like planets around the sun
Atom is mostly
empty space
Ernest Rutherford’s Model

Did not explain
chemical
properties

of elements
Better description of the
electron
behavior
was needed
Niels Bohr’s Model
 Electrons
move in
specific circular
paths, or orbits,
at different levels.
 Amount of fixed energy,
(energy levels,)
separates one level from
another.
 Electrons can jump
from one level to another.
Bohr’s model
Energy level analogous to the
rungs of a ladder
 Electrons cannot
exist between
energy levels, just like you can’t
stand between rungs on a ladder

Bohr’s model
Quantum –
amount of energy required to
move an electron
from one energy level to another
 The further away
from the nucleus,
the more energy
the electron has.

The Quantum Mechanical
Model



In 1926, Erwin Schrodinger derived the quantum
mechanical model
Determines the energy
of an electron
States the probability
of finding an electron a
certain distance from the nucleus.
The Quantum Mechanical
Model



The atom is found
inside a blurry
“electron cloud”
An area where there is a
chance of
finding an electron.
Think of fan blades
Atomic Orbitals



Atomic orbitals –
regions where there is a
high probability
of finding an electron.
Sublevels- arranged in sections:
 letters s, p, d, and f
Each sublevel
corresponds to a
different shape.
Principal Quantum Number
“n”- it denotes the energy level
in which the electron is located.
~ 1, 2, 3, etc.
Maximum number of electrons that can
fit in an energy level is:
2n2
How many e- in level 2? 3?
# 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
S-orbital
P-orbitals
D-orbitals
F-orbitals
By Energy Level




First Energy Level
Has only s orbital
only 2 electrons
1s2





Second Energy Level
Has s and p orbitals
Electrons:
2 in s, 6 in p
2s22p6
8 total electrons
By Energy Level





Third energy level
Has s, p, and d
orbitals
Electrons: 2 in s, 6 in p,
and 10 in d
3s23p63d10
18 total electrons





Fourth energy level
Has s, p, d, and f
orbitals
Electrons: 2 in s, 6 in p,
10 in d, and 14 in f
4s24p64d104f14
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 = s level – 1
f = s level – 2
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
electrons are
arranged in various
orbitals around the nuclei of atoms. Three
rules tell us how:
Rule #1 - Aufbau Principle

Electrons must occupy
the orbital with the
lowest energy first

Example: Oxygen 1s22s22p4
1s
2s
2p
1s
2s
2p
Rule #2 - Pauli Exclusion Principle

Orbitals can only have two electrons max
The 2 electrons must have
opposite spins

Example: Oxygen 1s22s22p4

1s
2s
2p
1s
2s
2p
Rule #3 - Hund’s Rule


Orbitals of equal energy are each occupied by one
electron before 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
Lowest energy to
higher energy.
 Adding electrons can change the energy of the
orbital.
Full orbitals
are the absolute best situation.

Orbitals fill in an order
 However,
•
•
half filled
orbitals have a lower energy, and are
next best.
Makes them more stable.
Changes the filling order
Practice Problems
Write electron configurations for the following atoms
1.
2.
3.
4.
Li
N
Be
C
1s22s1
1s22s22p3
1s22s2
1s22s22p2
5.
6.
7.
8.
P
Si
Mg
Al
1s22s22p63s23p3
1s22s22p63s23p2
1s22s22p63s2
1s22s22p63s1
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


1s22s22p63s23p64s23d2
1s22s22p63s23p64s23d3
Chromium - 24 electrons
2 2 6 2 6 2 4
 1s 2s 2p 3s 3p 4s 3d (expected)
 But
this is not what happens!!
Chromium is actually:
1s22s22p63s23p64s13d5
 Why?
 This gives us two
half filled orbitals
(the others are all still full)
 Half full is slightly
lower in energy.
 The same principal applies to copper.

Copper’s electron configuration





Copper has 29 electrons so we expect:
1s22s22p63s23p64s23d9
But the actual configuration is:
1s22s22p63s23p64s13d10
This change gives one more filled orbital and one
that is half filled.
Remember these exceptions: d4, d9
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
magnetic field.
s describes the direction an
electron spins

They must spin in
opposite directions


Spin= up
down

There are two values of s: +1/2 and -1/2
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
Wavelength
Origin
Amplitude
Trough
Low
Energy
High
Energy
Radio Micro Infrared
waves waves .
Low
Frequency
Long
Wavelength
Ultra- Xviolet Rays
Visible Light
Gamma
Rays
High
Frequency
Short
Wavelength
Long
Wavelength
=
Low
Frequency
=
Low
ENERGY
Short
Wavelength
=
High
Frequency
=
High
ENERGY
Wavelength Table
Atomic Spectrum

•
•
Atomic Emission
Spectrum –
the discrete lines
representing the
frequencies of light
emitted by an element
Unique to each
element, like fingerprints!
Very useful for identifying
elements

When these
electrons return
to their
lower
energy levels,
they lose
energy by
emitting
light
Changing the energy

Let’s look at a hydrogen atom, with only
one electron, 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
“excited”
Changing the energy

As the electron falls back to the
ground state, it gives the energy back as light
Changing the energy


Fall down in specific steps
Each step has a different energy
Ultraviolet

Visible
The further they fall,
more energy is released
and the higher
the frequency.
Infrared