Transcript Modern Atomic Theory and the Periodic Table Chapter 10 Chapter 10 - Modern Atomic Theory and the Periodic Table 10.1 A Brief History 10.2 Electromagnetic Radiation 10.3

Modern Atomic Theory
and the Periodic Table
Chapter 10
1
Chapter 10 -
Modern Atomic
Theory and the Periodic Table
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
A Brief History
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.
3
Electromagnetic Radiation
Examples
 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
4
Characteristics of a Wave
Wavelength (λ)
Light has the properties of a wave.
wavelength
wavelength
(measured
from
(measured
from
peak totrough
peak) to trough)
5
10.1
Frequency (n) is the number of wavelengths
that pass a particular point per second.
6
10.1
Speed (v) is how fast a wave moves
through space.
7
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.
8
The Electromagnetic Spectrum
X-rays are part visible
of the light
Infrared
is partlight
of is
electromagnetic
the electromagnetic
part of the
spectrum
spectrum
electromagnetic
spectrum
9
10.2
The Bohr Atom
• 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.
10
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.
11
10.3
Niels Bohr
Niels Bohr, a Danish physicist,
in 1912-1913 carried out research
on the hydrogen atom.
12
The Bohr Atom
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.
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10.4
The Bohr Atom
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.
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10.4
The Bohr Atom
Different lines of the
hydrogen spectrum
correspond to different
electron energy level
shifts.
15
10.4
The Bohr Atom
Light is not emitted
continuously. It is
emitted in discrete
packets called quanta.
16
10.4
The Bohr Atom
E1
E2
E3
An electron can have
one of several possible
energies depending on
its orbit.
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10.4
The Bohr Atom
 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.
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 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.
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 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
an atom cannot be determined.
20
 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.
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Energy Levels of Electrons
The
wave-mechanical
model
of
the
atom
According to Bohr the energies of electrons
also predicts discrete principal energy
in an atom are quantized.
levels within the atom
23
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.
24
10.7
10.7, 10.8
Each principal energy level
is subdivided into sublevels.
25
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.
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10.10
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.
27
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.
10.10
A p sublevel can hold a maximum of 6
electrons.
28
The P orbitals
pz
The three p orbitals share
a common center.
py
px
The three p orbitals point in
different directions.
29
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.
10.11
30
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
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The Hydrogen Atom

thediameter
ground state
• In
The
of
hydrogen’s single
hydrogen’s
nucleus
electron lies inelectron
the 1s
-13 cm.
is
cloud
about
is 10
about
orbital.

Hydrogen
can
absorb
times
• 100,000
The diameter
of
energy
and
thethe
electron
greater
than
hydrogen’s
electron
will move to excited
diameter
of its 10-8
cloud
states. is about
nucleus.
cm.
33
10.12
Atomic Structure of the
First 18 Elements
To determine the electronic structures of
atoms, the following guidelines are used.
34
Pauli exclusion principle
1.
No more than two
electrons can
occupy one orbital
35
10.10
1 s orbital
2.
3.
2 s orbital
Electrons occupy the lowest energy orbitals available.
They enter a higher energy orbital only after the lower
orbitals are filled.
For the atoms beyond hydrogen, orbital energies vary
as s<p<d<f for a given value of n.
36
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.
37
10.10
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
38
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
39
Orbital Filling
 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.
40
Filling the 1s Sublevel
↑
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.
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Filling the 2s Sublevel
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
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fourth electrons.
Filling the 2p Sublevel
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.
43
N ↑↓
↑↓
1s
2s
↑
↑
↑
1s22s22p3
2p
The third p electron of nitrogen enters a different p orbital than its
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.
44
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.
Ne ↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
2s
2p
1s22s22p6
There are 6 electrons in the 2p sublevel of neon, which fills the
45
sublevel.
Filling the 3s Sublevel
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.
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47
48
Electron Filling Order
49
Sublevel energy level order:
1s < 2s < 2p < 3s < 3p <
4s < 3d < 4p < 5s < 4d <
5p < 6s < 4f < 5d < 6p <
7s < 5f < 6d
You can memorize this sequence or....
Electron Structures and the
Periodic Table
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.
52
Period numbers correspond
Horizontal rows are
to the highest occupied
called periods.
energy level.
53
10.14
Elements
with
similar
Elements
in
the
B groups
groups
in
A
Groups are
numbered
properties
arenumerals.
organized
are
designated
transition
are
designated
with
Roman
in groups or families.
elements.
representative
elements.
54
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.
57
10.15
With the exception of helium which has a filled s
orbital, the nobles gases have filled p orbitals.
58
10.15
To write an electron configuration
using a noble-gas core:
1. Find the highest atomic-numbered
noble gas (Group 8A element)
less than the atomic number
of the element for which the
configuration is being written
2. Write the elemental symbol of the
noble gas in square brackets, followed
by the remaining configuration
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
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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
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Exceptions to the conventional filling order:
1. d4 configurations generally do not exist
Chromium (Z = 24):
Systematic prediction:
Cr: [Ar]4s23d4
But d4 is not likely,
so promote an electron from the 4s sublevel:
Cr: [Ar]4s13d5
2. d9 configurations generally do not exist
Copper (Z = 29):
Systematic prediction:
Cu: [Ar]4s23d9
But d9 is not likely,
so promote an electron from the 4s sublevel:
Cu: [Ar]4s13d10
d orbital numbers are 1 less
than dthe
period
number
orbital
filling
Arrangement of electrons
according to sublevel being filled.
10.16
64
f orbital numbers are 2 less
than the
period
number
f orbital
filling
Arrangement of electrons
according to sublevel being filled.
10.16
65
Period number corresponds with the
highest energy level occupied by
electrons in that period.
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10.17
The
Thegroup
elements
numbers
of a family
for thehave
representative
the same
outermost
elements electron
are equal
configuration
to the total number
except that
of
outermost
the electrons
electrons
are inindifferent
the atoms
energy
of the
levels.
group.
67
10.17
Chapter 10 -
Modern Atomic
Theory and the Periodic Table
10.1 A Brief History
10.2 Electromagnetic Radiation
10.3 The Bohr Atom – Niels Bohr description of the atom
(electron orbitals).
10.4 Energy Levels of Electrons – Electron configuration
(from the periodic table), s, p, d, and f orbitals.
10.5 Atomic structures of the First 18 Elements – Valence
electrons, Representatives and Transition elements,
Families names.
10.6 Electron Structures and the Periodic table –
68
Relationship between group number and valence electrons.