Chapter 5 Atomic Models

1
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. (??)
• 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 atoms internal
  structure.

2
5.1 Models help us visualize the invisible world
of atoms

• Atom is very small.

3
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.
4
We can see atoms indirectly through scanning
tunneling microscope (STM), which was invented in
1980s.
Fig1Scanning tunneling microscope Fig2an image
of gallium and arsenic atoms obtained with an STM
fig3 an STM image of mono-layer of perylene
derivative on graphite substrate, where the
epitaxial relationship is observed between the
organic molecule and the substrate graphite
5
5.2 Light is a form of energy
6
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 elements atomic spectrum (????).
  In 1800s researchers noted the orderliness of
  elements atomic spectrum, especially hydrogen,
  but could not give the explanation.

7
?????
8
5.4 Niels Bohr used the quantum hypothesis to
explain atomic spectra

• Max Plancks quantum hypothesis (????) 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 (??). Light behaves as both a
  wave and a particle.

9
Bohrs explanation
Electron loses potential energy and moves closer
to nucleus. A photon of light is emitted
Electron gains potential energy and moves farther
from nucleus. A photon of light is absorbed
10
Bohrs 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 (????).
  The energy level with n1 has the lowest energy.

11

• 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.

12
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.  

13

• The electrons wave nature can be used to explain
  the Bohrs 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
  self-reinforcing wave, we that the circumference
  of the smallest orbit can be no smaller than a
  single wavelength.

14
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.

15
Atomic orbitals have shapes and sizes!
16

• The size of orbital is indicated by Bohrs
  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
17
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.

18

• 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 (?????), since they
are the ones in direct contact with the external
environment.
19
C 6 electrons
N 7 electrons
O 8 electrons
F 9 electrons
Ne 10 electrons
C 1s22s22p2
N 1s22s22p3
O 1s22s22p4
F 1s22s22p5
Ne 1s22s22p6
7.7
20
5.7 Orbitals of similar energies can be grouped
into shells
The seven rows correspond to 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)
21

• 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
  (???), the amount of energy needed to pull an
  electron away from an atom). The ionization
  energy also determines the atoms chemical
  behavior, which will be discussed in chapter 6.