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Title: Chapt. 11 Atomic Structure


1
Chapt. 11 Atomic Structure
2
From macroscopic to microscopic
  • http//micro.magnet.fsu.edu/primer/java/scienceopt
    icsu/powersof10/

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History
  • Greek Philosopher Democritus (460-370 B.C.)
  • all matter composed of small atoms
  • atomos indivisible

6
  • 384-322 BC
  • Aristotle and Plato favored the earth, fire, air
    and water approach to the nature of matter

7
400 to 0, 0 to 1600 400 1600 2000 years
8
Renaissance
Medieval
  • http//www.youtube.com/watch?vwrD49Jci6h8feature
    related
  • http//www.youtube.com/watch?vrFI3UkPTpMsfeature
    related

9
Baroque
  • http//www.youtube.com/watch?v1ZhHjZLgWOs
  • Boyle
  • 1600s

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Avogadro who, in 1811,1 hypothesized that two
given samples of an ideal gas, at the same
temperature, pressure and volume, contain the
same number of molecules
Dalton 1803
Bernoulli in 1738
Boltzmann 1898
  • Boyle published it in 1662
  • about 1787,Charles had found that oxygen,
    nitrogen, hydrogen, carbon dioxide, and air
    expand to the same extent over the same 80 degree
    interval.

1687 Newton
1845, Watterson
12
  • We might as well attempt to introduce a new
    planet into the solar system, or to annihilate
    one already in existence, as to create or destroy
    a particle of hydrogen.
  • John Dalton, A New System of Chemical Philosophy,
    1808)

13
Dalton's Postulates
  • 1. All matter consists of tiny
  • particles.

14
Dalton's Postulates
  • 2. Elements are characterized by the mass of
    their atoms. All atoms of the same element have
    identical weights, Dalton asserted. Atoms of
    different elements have different weights.

15
Dalton's Postulates
  • 3. Atoms are indestructible
  • and unchangeable.

16
Daltons Postulates
  • 4. When elements react, their atoms combine in
    simple, whole-number ratios.

17
John Daltons Atomic Theory
  • Almost right. A good start.

very small
Structure of the atom after Dalton (ca. 1810)
18
our future discoveries must be looked for in the
6th decimal place, 1894, at the dedication of
Ryerson Physics Laboratory, Chicago
  • Heading into the 20th century there was a feeling
    by many in the chemistry and physics communities
    that our scientific knowledge was nearly
    complete. It was universally accepted that atoms
    were the most basic constituent of matter and
    that the behavior of all matter could be
    explained through Newtonian mechanics.
  • BUTseveral discoveries and observations
    contradicted these theories

19
  • I. Discovery of the electron

20
Cathode ray tube and the electron
  • http//videos.howstuffworks.com/science-channel/29
    292-100-greatest-discoveries-the-cathode-ray-tube-
    video.htm

21
J.J. Thomson (1897) Cathode Rays
Atoms subjected to high voltages give off cathode
rays.
22
J.J. Thomson Cathode Rays
Cathode rays can be deflected by a magnetic field.
Cathode rays are negatively charged particles
(electrons).
Electrons are in atoms.
23
Thompsons Plum Pudding Model
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400 to 0, 0 to 1600 400 1600 2000 years
25
Radioactivity
  • Radioactivity is the spontaneous emission of
    radiation by an atom.
  • First observed by Henri Becquerel
  • (1852-1908).
  • Marie and Pierre Curie also studied it.
  • Nobel Prize in 1903 (physics).

26
Radiation named by Rutherford
27
2. Discovery of the nucleus
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Discovery of the nucleus
  • http//www.youtube.com/watch?v5pZj0u_XMbc

30
  • Since the vast majority of a particles pass
    through the Au foil undeflected, the Au atoms are
    mostly ______.
  • Empty space

31
  • A very tiny percentage of a particles hit
    something massive in the atom and backscatter
    (bounce back).
  • This indicates that most of the mass of the atom
    is concentrated in a very small volume relative
    to the volume of the entire atom.
  • We now call this the NUCLEUS.

32
  • Rutherford proposed the charge on the nucleus to
    be positive, since electrons are negatively
    charged and atoms are neutral.

33
Rutherfords model
34
  • The primary limitation to Bohr Theory is that it
    was limited to a description of a one electron
    system, namely Hydrogen.
  • The description of multiple electron systems is
    much more complex, and was only adequately
    handled by the Modern Theory of Atomic Structure.

35
Bohrs Theory
  • 1. The electron travels in a circular path around
    the nucleus. This path is called an orbit.

36
  • 2. At normal living conditions, room temperature,
    the electron resides in the orbit which is
    closest to the nucleus.
  • This is the position of lowest energy content for
    the electron, and is referred to as the Ground
    State. (This statement implies that there will be
    more that one orbit available to an electron.)

37
  • 3. As long as the electron remains in a specific
    orbit, no energy is gained or lost by the system.
  • 4. If energy is added to an electron, the
    electron will move to a new orbit. This orbit
    will be farther from the nucleus, and is a
    position of higher energy content. This new
    position is known as an excited state.

38
  • 5. When an electron moves from one orbit to
    another orbit, it does so without ever passing
    through the space between the orbits. In other
    words, the electron is only allowed to exist at
    very specific distances from the nucleus, or
    positions of very specific energy content. (This
    idea is much like climbing a ladder. The foot is
    only allowed to be placed in very specific
    locations.) This idea is known as a quantum jump,
    a transition in which the electron gains or loses
    a very specific amount of energy.

39
  • Part Five off the Bohr Theory is, perhaps, the
    most controversial item. It says that an electron
    is restricted to having certain specific
    quantities of energy. The electron will never be
    allowed to have energy in between the allowed
    values. This is referred to as the quantization
    of energy. The idea was first expressed by Max
    Planck. This piece of information, when given to
    Bohr, suddenly made the ideas that he expressed
    much more meaningful.

40
  • In essence, it now becomes clear why an atom will
    only release specific colors of light, or
    specific wavelengths of electromagnetic
    radiation. Without the Planck contribution, the
    Bohr atom would release all colors of light.

41
  • 6. When an electron is in an excited state, it
    will always drop down to a lower energy state,
    ultimately returning to ground state. Each
    electron transition to a lower energy state will
    be accompanied by the simultaneous release of
    energy. This energy is released as
    electromagnetic radiation. The energy of the
    released radiation will correspond to the
    difference in energy content between the two
    levels.

42
EMR (electromagnetic radiation)
43
3 problems with the Bohr Theory
  • The theory only works for a one electron system.
    What happens when an atom has more than one
    electron?
  • The theory violates the Heisenberg Uncertainty
    Principle. Bohr Theory makes the behavior of the
    electron entirely to predictable. Bohr claims it
    is possible to know exactly where an electron is
    and what it is doing. The Heisenberg Uncertainty
    Principle says that is not possible.

44
  • The Bohr Theory will based on trying to explain
    four visible colors in the hydrogen atomic
    spectrum. He worked with a red line, blue-green
    line, blue line, and violet line. With improved
    instrumentation, it is now known that the red
    line is actually two red lines. These lines are
    extremely close together, and are referred to as
    a doublet. The instruments that were available to
    Bohr were not sophisticated enough to distinguish
    the two red lines. To him they looked as if they
    were one wide red line.

45
  • The Atomic Spectrum is a series of lines of color
    produced when light from an excited atom is
    passed through a prism. It is also known as a
    line spectrum.

46
  • 3. (Failure of) the classical description of the
    atom

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48
Water wave
Sound wave
49
EMR electromagnetic radiation
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52
Orbitals...
53
Electrons are part of what makes an atom an atom
54
Electrons are part of what makes an atom an atom
But where exactly are the electrons inside an
atom?
atom
55
Orbitals are areas within atoms where there is a
high probablility of finding electrons.
56
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
57
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
58
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
59
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
60
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
61
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
62
Knowing how electrons are arranged in an atom
is important because that governs how atoms
interact with each other
63
It has been determined where the orbitals are
inside an atom, but it is not known precisely
where the electrons are inside the orbitals
64
It has been determined where the orbitals are
inside an atom, but it is not known precisely
where the electrons are inside the orbitals (as
described by Heisenburgs Uncertainty
Principle)
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66
Hey, where am I?
67
The area where an electron can be found, the
orbital, is defined mathematically, but we can
see it as a specific shape in 3-dimensional space
68
z
y
x
69
z
y
The 3 axes represent 3-dimensional space
x
70
z
y
For this presentation, the nucleus of the atom is
at the center of the three axes.
x
71
The 1s orbital is a sphere, centered around the
nucleus (l 0)
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The 2s orbital is also a sphere.
75
The 2s electrons have a higher energy than the
1s electrons. Therefore, the 2s electrons are
generally more distant from the nucleus, making
the 2s orbital larger than the 1s orbital.
76
1s orbital
77
2s orbital
78
Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
79
Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
The s orbitals are spheres
80
There are three 2p orbitals
81
The three 2p orbitals (l 1) are
oriented perpendicular to each other
82
z
This is one 2p orbital (2py)
y
x
83
z
another 2p orbital (2px)
y
x
84
z
the third 2p orbital (2pz)
y
x
85
Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
86
Dont forget an orbital is the shape of
the space where there is a high probability of
finding electrons
This is the shape of p orbitals
87
z
y
x
88
z
2px
y
x
89
z
2px and 2pz
y
x
90
z
The three 2p orbitals, 2px, 2py, 2pz
y
x
91
z
The three 2p orbitals, 2px, 2py, 2pz (m -1,
0, 1)
y
x
92
once the 1s orbital is filled,
93
the 2s orbital begins to fill around the 1s
orbital
94
once the 2s orbital is filled,
95
the 2p orbitals begin to fill
96
each 2p orbital intersects the 2s orbital and the
1s orbital
97
each 2p orbital gets one electron before pairing
begins
98
once each 2p orbital is filled with a pair of
electrons, then
99
the 3s orbital gets the next two electrons
100
the 3s electrons have a higher energy than 1s,
2s, or 2p electrons,
101
so 3s electrons are generally found further from
the nucleus than 1s, 2s, or 2p electrons
102
What does that have to do with anything??
103
the billions of interactions of atoms constantly
going on around you depend on how the
electrons are arranged in each atom
104
the billions of interactions of atoms constantly
going on around you depend on how the
electrons are arranged in each atom
the arrangement of an atoms electrons (its
orbitals) govern how that atom will interact with
other atoms
105
the billions of interactions of atoms constantly
going on around you depend on how the
electrons are arranged in each atom
the arrangement of an atoms electrons (its
orbitals) govern how that atom will interact with
other atoms
If atoms did not interact with each other, you
would not be sitting here reading this
106
An interesting place where electrons have a
specific organization within atoms, allowing for
intersting atom interactions
107
Not an interesting place, where electrons have no
specific organization within atoms, where atoms
wander aimlessly about
An interesting place where electrons have a
specific organization within atoms, allowing for
intersting atom interactions
108
Principal Quantum Number, n
  • Indicates main energy levels
  • n 1, 2, 3, 4
  • Each main energy level has sub-levels

109
Energy Sublevels
  • s p d f g

110
  • The principle quantum number, n, determines the
    number of sublevels within the principle energy
    level.

111
Orbital Quantum Number, l(Angular Momentum
Quantum Number)
  • Indicates shape of orbital sublevels
  • l n-1
  • l sublevel
  • 0 s
  • 1 p
  • 2 d
  • 3 f
  • 4 g

112
Orbital
  • The space where there is a high probability that
    it is occupied by a pair of electrons.
  • Orbitals are solutions of Schrodingers equations.

113
Orbitals in Sublevels
  • Sublevel Orbitals electrons
  • s 1 2
  • p 3 6
  • d 5 10
  • f 7 14
  • g 9 18

114
Three rules are used to build the electron
configuration
  • Aufbau principle
  • Pauli Exclusion Principle
  • Hunds Rule

115
Aufbau Principle
  • Electrons occupy orbitals of lower energy first.

116
Aufbau Diagram
117
-Pauli Exclusion Principle(Wolfgang Pauli,
Austria, 1900-1958)-Electron Spin Quantum Number
  • An orbital can hold only two electrons and they
    must have opposite spin.
  • Electron Spin Quantum Number (ms)
  • 1/2, -1/2

118
Hunds Rule
  • In a set of orbitals, the electrons will fill the
    orbitals in a way that would give the maximum
    number of parallel spins (maximum number of
    unpaired electrons).
  • Analogy Students could fill each seat of a
    school bus, one person at a time, before doubling
    up.

119
Aufbau Diagram for Hydrogen
120
Aufbau Diagram for Helium
121
Aufbau Diagram for Lithium
122
Aufbau Diagram for Beryllium
123
Aufbau Diagram for Boron
124
Aufbau Diagram for Carbon
125
Aufbau Diagram for Nitrogen
126
Aufbau Diagram
127
Notations of Electron Configurations
  • Standard
  • Shorthand

128
Aufbau Diagram for Fluorine
129
Standard Notation of Fluorine
Number of electrons in the sub level 2,2,5
1s2 2s2 2p5
Main Energy Level Numbers 1, 2, 2
Sublevels
130
Shorthand Notation
  • Use the last noble gas that is located in the
    periodic table right before the element.
  • Write the symbol of the noble gas in brackets.
  • Write the remaining configuration after the
    brackets.
  • Ex Fluorine He 2s2 2p5

131
Blocks in the Periodic Table
132
General Periodic Trends
  • Atomic and ionic size
  • Ionization energy
  • Electron affinity

Higher effective nuclear charge.
Electrons held more tightly
133
Atomic Size
  • Size goes UP on going down a group.
  • Because electrons are added farther from the
    nucleus, there is less attraction.
  • Size goes DOWN on going across a period.

134
Atomic Radii
Figure 8.9
135
Trends in Atomic SizeSee Figures 8.9 8.10
136
Ion Sizes
Does the size go up or down when losing an
electron to form a cation?

137
Ion Sizes
Forming a cation.
Li,152 pm
3e and 3p
  • CATIONS are SMALLER than the atoms from which
    they come.
  • The electron/proton attraction has gone UP and so
    size DECREASES.

138
Ion Sizes
  • Does the size go up or down when gaining an
    electron to form an anion?

139
Ion Sizes
Forming an anion.
  • ANIONS are LARGER than the atoms from which they
    come.
  • The electron/proton attraction has gone DOWN and
    so size INCREASES.
  • Trends in ion sizes are the same as atom sizes.

140
Trends in Ion Sizes
Figure 8.13
141
Ionization EnergySee Screen 8.12
  • IE energy required to remove an electron from
    an atom in the gas phase.

Mg (g) 738 kJ ---gt Mg (g) e-
142
Ionization EnergySee Screen 8.12
  • Mg (g) 735 kJ ---gt Mg (g) e-
  • Mg (g) 1451 kJ ---gt Mg2 (g) e-

Mg2 (g) 7733 kJ ---gt Mg3 (g) e-
Energy cost is very high to dip into a shell of
lower n. This is why ox. no. Group no.
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