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Chapter 5 Atomic Models
Different metal containing compounds emit colored light
when fireworks burn. This colored light is a combination
of a series of colors called a spectral pattern. (谱带)
Each element emits its own characteristic spectral pattern,
which can be used to identify the element. This allowed the
scientists in 1900s to develop models of the atom’s internal
structure.
5.1 Models help us visualize the invisible
world of atoms
Atom is very small.
We can not see them in the usual sense. This is because light
travels in waves and atoms are smaller than the wavelengths
of visible light, which is the light that allows the human eye
to see things. So we can not see atoms through the media of
light, even with a microscope.
We can see atoms indirectly through scanning
tunneling microscope (STM), which was invented in
1980s.
Fig1:Scanning tunneling
microscope
Fig2:an image of gallium
and arsenic atoms
obtained with an STM
fig3 : an STM image of monolayer of perylene derivative on
graphite substrate, where the
epitaxial relationship is observed
between the organic molecule and
the substrate graphite
5.2 Light is a form of energy
5.3 Atoms can be identified by the light
they emit
When we view the light from glowing atoms, we
see that the light consists of a number of discrete
frequencies rather than a continuous spectrum.
This is called element’s atomic spectrum (原子
光谱). In 1800s researchers noted the orderliness
of element’s atomic spectrum, especially
hydrogen, but could not give the explanation.
氢原子光谱
5.4 Niels Bohr used the quantum
hypothesis to explain atomic spectra
Max Planck’s quantum hypothesis (量子假设): a beam of
light energy is not the continuous stream of energy, but
consists of small, discrete packets of energy. Each packet
was called a quantum. In 1905, Einstein recognized that
these quanta of light behave like particles. Each quantum
was called a photon (光子). Light behaves as both a wave
and a particle.
Bohr’s explanation
Electron gains
potential
energy and
moves farther
from nucleus.
A photon of
light is
absorbed
Electron
loses
potential
energy and
moves closer
to nucleus. A
photon of
light is
emitted
Bohr’s planetary model of atom
There are only a limited
number
of
permitted
energy levels in an atom,
and an electron can only
stay in these energy levels.
Each energy level has a
principal
quantum
number n ( 主 量 子 数 ).
The energy level with n=1
has the lowest energy.
Photons are emitted by atoms
as electrons move from
higher-energy outer orbits
to lower-energy inner
orbits. The energy of an
emitted photon is equal to
the difference in energy
between the two orbits.
Because an electron is
restricted to discrete orbits,
only
particular
light
frequencies are emitted.
5.4 Electrons exhibit wave properties
Question by Louis de Groglie: If light
has both wave properties and
particle properties, why can not
material particle, such as electron,
also have both?
Answer: Every particle of matter is
somehow endowed with a wave to
guide it as it travels. The more
slowly an electron moves, the more
its bahvior is that of a particle with
mass. The more quickly it moves,
the more its behavior is that of a
wave of energy.
In an atom, an electron moves at very
high speeds: on the order of
2000000 meters per second.
Practical application of the wave
behavior of electrons: electron
microscope.
The electron’s wave nature can be used to explain
the Bohr’s planetary model: (1) Permitted energy
levels are a natural consequence of electron waves
closing in on themselves in a synchronized manner.
(2) By viewing each electron orbit as a selfreinforcing wave, we that the circumference of the
smallest orbit can be no smaller than a single
wavelength.
How to visualize electron waves?
Probability clouds and atomic orbitals
In 1926, Erwin Schrodinger
formulated a equation
from which the intensities
of electron waves in an
atom could be calculated.
It was found that the
intensity at any given
location determined the
probability of finding the
electron at that location.
The electron was mostly
likely to be found where
its wave intensity was
grestest.
Where 90% of the
e- density is found
for the 1s orbital
e- density (1s orbital) falls
off rapidly as distance from
nucleus increases
Atomic orbitals have shapes and sizes!
l = 0 (s orbitals)
l = 1 (p orbitals)
l = 2 (d orbitals)
The size of orbital is indicated by Bohr’s principal quantum
number n = 1, 2, 3, 4, 5, 6, 7…
Fig5.19 The fluorine atom has five
overlapping atomic orbitals that contain
its nine electrons, which are not shown
5.6 Energy-level diagrams describe how
orbitals are occupied
Each orbital has a capacity of two,
but no more than two, electrons.
They spin in opposite directions,
which generates two oppositely
oriented magnetic fields that are
attractive
and
partly
compensate for the electrical
repulsion between the electrons.
Example:
Lithium (3): 1s22s1
Boron (5): 1s22s22p1
Carbon (6): 1s22s22p2
Nitrogen (7): 1s22s22p3
Oxygen (8): 1s22s22p4
The properties of an atom are determined mostly by its
outermost electrons (最外层电子), since they are the
ones in direct contact with the external environment.
The most stable arrangement of electrons
in subshells is the one with the greatest
number of parallel spins (Hund’s rule).
Ne97
C
N
O
F
6
810
electrons
electrons
electrons
22s
222p
22p
5
246
3
Ne
C
N
O
F 1s
1s222s
7.7
5.7 Orbitals of similar energies can be
The
seven
correspond to
grouped into shells
rows
the
seven periods in the
periodic table.
The maximum number
of electrons each row
can hold is equal to
the number of
elements in the
corresponding period.
(2, 8, 8, 18, 18, 32, 32)
Increasing First Ionization Energy
From electron configuration, we can predict:
(1)
The smallest atoms are at the upper right of the periodic
table.
The smallest atoms have the most strongly held electrons.
(Represented by ionization energy (电离能), the amount of
energy needed to pull an electron away from an atom). The
ionization energy also determines the atom’s chemical behavior,
which will be discussed in chapter 6.
Increasing First Ionization Energy