Modern Atomic Theory
and the Periodic Table
Chapter 10
Hein and Arena
Version 2.0
12th Edition
Eugene Passer
Chemistry Department
Bronx Community
1 College
© John Wiley and Sons, Inc
Chapter Outline
10.1 A Brief History
10.2 Electromagnetic Radiation
10.3 The Bohr Atom
10.5 Atomic Structures of the First
18 Elements
10.6 Electron Structures and the
Periodic Table
10.4 Energy Levels of Electrons
2
10.1
A Brief History
3
A Brief History of Atomic Theory
Greeks were the first to suggest
that matter is made up of atoms
Early chemists
performed experiments
Their experiments led to
Dalton's Atomic Theory
Limitations of Dalton's model
led to the Thompson and Rutherford
models of the atom.
While these models work reasonably well
their limitatons have led to more modern theories
as to the nature of the atom.
4
10.2
Electromagnetic Radiation
5
Energy can travel through space as
Examples
electromagnetic radiation.
6
•
•
•
•
•
•
light from the sun
x-rays
microwaves
radio waves
television waves
radiant heat
All show wavelike
behavior.
Each travels at
the same speed
in a vacuum.
3.00 x 108 m/s
7
Characteristics of a Wave
8
Wavelength (λ)
9
Light has the properties of a wave.
wavelength
wavelength
(measured
from
(measured
from
peak totrough
peak) to trough)
10
10.1
Frequency (ν)
11
Frequency is the number of wavelengths
that pass a particular point per second.
12
10.1
Speed (v)
13
Speed is how fast a wave moves through
space.
14
10.1
• Light also exhibits the properties of a
particle. Light particles are called
photons.
• Both the wave model and the particle
model are used to explain the
properties of light.
15
The Electromagnetic Spectrum
16
X-rays are part visible
of the light
Infrared
is partlight
of is
electromagnetic
the electromagnetic
part of the
spectrum
spectrum
electromagnetic
spectrum
17
10.2
10.3
The Bohr Atom
18
• At high temperatures or voltages,
elements in the gaseous state emit light
of different colors.
• When the light is passed through a
prism or diffraction grating a line
spectrum results.
19
Each element has its own
unique set of spectral emission
lines that distinguish it from
other elements.
These colored lines
indicate that light is
being emitted only at
certain wavelengths.
Line spectrum of hydrogen. Each line corresponds
to the wavelength of the energy emitted when the
electron of a hydrogen atom, which has absorbed
energy falls back to a lower principal energy level.
20
10.3
Niels Bohr
21
Niels Bohr, a Danish physicist,
in 1912-1913 carried out research
on the hydrogen atom.
22
The Bohr Atom
23
Electrons
revolve
An
electron
has a
around the
nucleus
in it
discrete
energy
when
orbits thatan
areorbit.
located
occupies
at fixed distances from
the nucleus.
24
10.4
Whencolor
an electron
fallslight
The
of the
from a higher
energy level
emitted
corresponds
to
to a lower
energy
a
one
of the
lines level
of the
quantum of
energy in the
hydrogen
spectrum.
form of light is emitted by
the atom.
25
10.4
Different lines of the
hydrogen spectrum
correspond to different
electron energy level
shifts.
26
10.4
Light is not emitted
continuously. It is
emitted in discrete
packets called quanta.
27
10.4
E1
E2
E3
An electron can have
one of several possible
energies depending on
its orbit.
28
10.4
• Bohr’s calculations succeeded very well
in
correlating
the
experimentally
observed spectral lines with electron
energy levels for the hydrogen atom.
• Bohr’s methods did not succeed for
heavier atoms.
• More theoretical work on atomic structure
was needed.
29
• In 1924 Louis De Broglie suggested
that all objects have wave properties.
– De Broglie showed that the wavelength
of ordinary sized objects, such as a
baseball, are too small to be observed.
– For objects the size of an electron the
wavelength can be detected.
30
• In 1926 Erwin Schröedinger created a
mathematical model that showed
electrons as waves.
– Schröedinger’s work led to a new branch
of physics called wave or quantum
mechanics.
– Using Schröedinger’s wave mechanics,
the probability of finding an electron in a
certain region around the atom can be
determined.
– The actual location of an electron within
31
an atom cannot be determined.
• Based on wave mechanics it is clear that
electrons are not revolving around the
nucleus in orbits.
• Instead of being located in orbits, the
electrons are located in orbitals.
• An orbital is a region around the nucleus
where there is a high probability of
finding an electron.
32
10.4
Energy Levels
of Electrons
33
TheAccording
wave-mechanical
model
of
the
atom
to Bohr the energies of
also predicts discrete principal energy
electrons in an atom are quantized.
levels within the atom
34
As n increases, the
energy of the electron
increases.
The first four
principal energy
levels of the
hydrogen atom.
Each level is
assigned a principal
quantum number n.
35
10.7
10.7, 10.8
Each principal energy level
is subdivided into sublevels.
36
Within sublevels the electrons are found in
orbitals.
An s orbital is spherical in
shape.
The
spherical
surface
encloses a space where
there is a 90% probability
that the electron may be
found.
37
10.9
An atomic orbital can hold a maximum of two
electrons.
An electron can spin in one
of two possible directions
represented by ↑ or ↓.
The two electrons that
occupy an atomic orbital
must have opposite spins.
This is known as the Pauli
Exclusion Principal.
38
10.10
A p sublevel is made up of three orbitals.
Each p orbital has two lobes.
Each p orbital can hold a maximum of two
electrons.
10.9
A p sublevel can hold a maximum of 6
electrons.
39
pz
The three p orbitals share
a common center.
py
px
The three p orbitals point in
different directions.
40
10.10
A d sublevel is made up of five orbitals.
The five d orbitals all point in different directions.
Each d orbital can hold a maximum of two
electrons.
A d sublevel can hold a maximum of 10 electrons.
41
10.11
Number of Orbitals in a Sublevel
10.8 10.10 10.11
42
Distribution of Subshells by
Principal Energy Level
n=1
1s
n=2
2s
2p 2p 2p
n = 3 3s
3p 3p 3p
3d 3d 3d 3d 3d
n = 4 4s
4p 4p 4p
4d 4d 4d 4d 4d
4f 4f 4f 4f 4f 4f 4f
43
The Hydrogen Atom
• The
In the
diameter
ground of
state
hydrogen’s electron
single
nucleus
electron
is
cloud
about
is 10
about
lies-13incm.
the
1s
timesof
• 100,000
Theorbital.
diameter
than
the
• greater
Hydrogen
hydrogen’s
can
electron
diameter
of its 10
absorbisenergy
cloud
about
and-8
nucleus.
the electron will
cm.
move to excited
states.
10.12
44
10.5
Atomic Structure of the
First 18 Elements
45
To determine the electronic structures of
atoms, the following guidelines are used.
46
1. No more than
two electrons
can occupy one
orbital
47
10.10
1 s orbital
2 s orbital
2. Electrons occupy the lowest energy orbitals
available. They enter a higher energy orbital
only after the lower orbitals are filled.
3. For the atoms beyond hydrogen, orbital
energies vary as s<p<d<f for a given value
of n.
48
10.10
4. Each orbital in a sublevel is occupied
by a single electron before a second
electron enters. For example, all three
p orbitals must contain one electron
before a second electron enters a p
orbital.
10.10
49
Nuclear makeup and electronic structure of
each principal energy level of an atom.
number of protons
and of electrons
number
neutrons in the nucleus
in each sublevel
50
10.13
Electron Configuration
Number of
electrons in
sublevel orbitals
Arrangement of
electrons within their
respective sublevels.
6
2p
Principal
Type of orbital
energy level
51
Orbital Filling
52
• In the following diagrams boxes
represent orbitals.
• Electrons are indicated by arrows: ↑ or
↓.
– Each arrow direction represents one of
the two possible electron spin states.
53
Filling the 1s Sublevel
54
H
↑
1s1
Hydrogen has 1 electron. It will occupy the orbital of lowest
energy which is the 1s.
He
↑↓
1s2
Helium has two electrons. Both helium electrons occupy the
1s orbital with opposite spins.
55
Filling the 2s Sublevel
56
Li
↑↓
↑
1s
2s
1s22s1
The 1s orbital is filled. Lithium’s third electron will enter the
2s orbital.
Be
↑↓
↑↓
1s
2s
1s22s2
The 2s orbital fills upon the addition of beryllium’s third and
fourth electrons.
57
Filling the 2p Sublevel
58
B
↑↓
↑↓
1s
2s
↑
1s22s22p1
2p
Boron has the first p electron. The three 2p orbitals have the same
energy. It does not matter which orbital fills first.
C ↑↓
↑↓
1s
2s
↑
↑
1s22s22p2
2p
The second p electron of carbon enters a different p orbital than the
first p electron so as to give carbon the lowest possible energy.
N ↑↓
↑↓
1s
2s
↑
↑
↑
1s22s22p3
2p
The third p electron of nitrogen enters a different p orbital than its
59
first two p electrons to give nitrogen the lowest possible energy
.
O ↑↓
↑↓
1s
2s
↑↓ ↑
↑
1s22s22p4
2p
There are four electrons in the 2p sublevel of oxygen. One of the
2p orbitals is now occupied by a second electron, which has a spin
opposite to that of the first electron already in the orbital.
F
↑↓
↑↓
↑↓ ↑↓ ↑
1s
2s
2p
1s22s22p5
There are five electrons in the 2p sublevel of fluorine. Two of the 2p
orbitals are now occupied by a second electron, which has a spin
opposite to that of the first electron already in the orbital.
60
Ne ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
1s22s22p6
There are 6 electrons in the 2p sublevel of neon, which fills the
sublevel.
61
Filling the 3s Sublevel
62
Na ↑↓
↑↓
↑↓ ↑↓ ↑↓
↑
1s
2s
2p
3s
1s22s22p63s1
The 2s and 2p sublevels are filled. The next electron enters the
3s sublevel of sodium.
Mg ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
↑↓ 1s22s22p63s2
3s
The 3s orbital fills upon the addition of magnesium’s twelfth
electron.
63
64
65
10.6
Electron Structures and
the Periodic Table
66
In 1869 Dimitri Mendeleev of Russia and
Lothar Meyer of Germany independently
published periodic arrangements of the
elements based on increasing atomic
masses.
Mendeleev’s arrangement is the precursor
to the modern periodic table.
67
Period numbers correspond
Horizontal rows are
to the highest occupied
called periods
energy level.
68
10.14
Elements with similar
properties are organized
in groups or families.
69
10.14
Elements in the A groups
are designated
representative elements
70
10.14
Elements in the B groups
are designated
transition elements
71
10.14
TheForchemical
A family elements
behaviortheand
valence
properties
electron of
elements
configuration
in a family
is the same
are associated
in each column.
with the
electron configuration of its elements.
72
10.15
With the exception of helium which has a filled s
orbital, the nobles gases have filled p orbitals.
73
10.15
The electron configuration of any of the
noble gas elements can be represented by
the symbol of the element enclosed in
square brackets.
1s22s22p1
[He]2s22p1
Na
1s22s22p63s1
[Ne]3s1
Cl
1s22s22p63s23p5
[Ne]3s23p5
B
74
The electron configuration of argon is
Ar 1s22s22p63s23p6
The elements after argon are potassium
and calcium Instead of entering a 3d
orbital, the valence electrons of these
elements enter the 4s orbital.
K
1s22s22p63s23p64s1
[Ar]4s1
Ca
1s22s22p63s23p6 4s2
[Ar]4s2
75
The number of a d orbital is 1
less than
its period
d orbital
fillingnumber
Arrangement of electrons
according to sublevel being filled.
10.16
76
The number of an f orbital is 2
less than
its period
f orbital
fillingnumber
Arrangement of electrons
according to sublevel being filled.
10.16
77
A period number corresponds to the
highest energy level occupied by
electrons in the period.
78
10.17
The elements
group numbers
of a family
for thehave
representative
the same
elements are
outermost
electron
equalconfiguration
to the total except
numberthat
of
outermost
the
electrons
electrons
are in different
in the atoms
energy
of the
levels.
group.
79
10.17
80