Transcript Electromagnetic Radiation

MODERN ATOMIC THEORY
Chapter 10
ANCIENT GREEKS’ VIEW OF
MATTER
About 400 B.C. , Aristotle thought all matter was
made of four “elements” :
•
earth
•
air
•
fire
•
water
ANCIENT GREEKS’ VIEW
OF MATTER
At about the same time another Greek
philosopher, Democritus, said that
matter was made of tiny, indivisible
particles called atoms.
Atomos is the Greek word for indivisible.
Modern View of the Atom
Tiny, dense, positively charged nucleus
made up of positive protons and neutral
neutrons.
Negatively charged electron shells enclose
the nucleus and contain negative
electrons.
Atomic Spectra and Bohr
One view of atomic structure in early 20th
century was that an electron (e-) traveled
about the nucleus in an orbit.
+
Electron
orbit
1. Any orbit should be possible and
so is any energy.
2. But a charged particle moving in
an electric field should emit
energy.
End result should be destruction!
Electromagnetic
Radiation
Radiant energy that exhibits
wave-like behavior and travels
through space at the speed of
light in a vacuum.
The electromagnetic spectrum.
Electromagnetic Radiation
wavelength
Visible light
Amplitude
wavelength
Ultraviolet radiation
Node
Waves
Waves have 3 primary characteristics:
1. Wavelength: distance between two peaks
in a wave.
2. Frequency: number of waves per second
that pass a given point in space.
3. Speed: speed of light is 2.9979  108 m/s.
As the wavelength () decreases, the frequency ()
increases.
Wavelength and frequency can be interconverted.
 = c/
 = frequency (s1, Hz, cyc/s, or waves/s )
 = wavelength (m)
c = speed of light (m/s)
Huygens thought light travels as waves, while Newton
believed it travels as particles.
Photons
Photons -- tiny particle of electromagnetic
radiation -- a bundle of light energy.
Ground state -- electrons are at their lowest
energy state in an atom.
Excited state -- electrons have absorbed
energy by jumping up to a higher energy
state in the atom.
Larger energy jumps by electrons produce
shorter wavelength (more energetic) light.
Line Spectra
of Excited Atoms
High E
Short 
High 
Low E
Long 
Low 
Visible lines in H atom spectrum are called
the BALMER series.
Line Spectra
of Excited Atoms
Excited atoms emit light of only certain
wavelengths
The wavelengths of emitted light depend
on the element.
Atomic Spectrum of Hydrogen
Continuous spectrum: Contains all the
wavelengths of light.
Bright Line (discrete) spectrum: Contains
only some of the wavelengths of light.
The diagrams above present evidence for discrete energy
levels about a nucleus. Electrons can only be found in
certain energy levels with certain energies.
Atomic Line Spectra and
Niels Bohr
Niels Bohr
(1885-1962)
Bohr’s greatest contribution to
science was in building a
simple model of the atom.
It was based on an
understanding of the
BRIGHT LINE SPECTRA
of excited atoms.
Bohr’s Model
Bohr’s Model was incorrect.
Replaced by QUANTUM or WAVE
MECHANICS MODEL.
e- can only exist in certain discrete orbitals.
e- is restricted to QUANTIZED energy states.
e- can not be exactly located--location based upon
probability.
Quantum or Wave
Mechanics
L. de Broglie
(1892-1987)
de Broglie (1924)
proposed that all
moving objects have
wave properties.
Quantum or Wave
Mechanics
E. Schrodinger
1887-1961
Schrodinger applied idea
of e- behaving as a
wave to the problem
of electrons in atoms.
Failure of the Bohr Model
The Bohr Model of the atom paved
the way for the Quantum Mechanical
Theory, but current theory is in no
way derived from the Bohr Model of
the atom. The Bohr Model of the
Atom was fundamentally incorrect-atoms do not move in circular orbits
about the nucleus.
1s Orbital
2s Orbital
p Orbitals
A p orbital
The three p
orbitals lie 90o
apart in space
2px Orbital
2py Orbital
2pz Orbital
3px Orbital
3dxy Orbital
3dxz Orbital
3dyz Orbital
3dyz Orbital
2
2
3dx - y
Orbital
Quantum Numbers (QN)
1. Principal QN (n = 1, 2, 3, . . .) - related to size
and energy of the orbital.
2. Angular Momentum QN -- l (s, p, d, & f) relates to shape of the orbital.
3. Magnetic QN -- ml (x, y, or z plane) - relates to
orientation of the orbital in space relative to
other orbitals.
4. Electron Spin QN -- ms (+1/2, 1/2) - relates to the
spin states of the electrons-- clockwise or
counterclockwise.
Electron Arrangement
Level
Sublevel # Orbitals # electrons
1-7
s
1
2
2-7
p
3
6
3-7
d
5
10
4-7
f
7
14
Energy Levels and Orbitals
•
n = the number of the energy level.
•
n2 = the number of orbitals in an energy
level.
•
2n2 = the number of electrons in an energy
level.
Pauli Exclusion Principle
In a given atom, no two electrons can have
the same set of four quantum numbers (n,
l, ml, ms).
Therefore, an orbital can hold only two
electrons, and they must have opposite
spins.
Aufbau Principle
As protons are added one by one
to the nucleus to build up the
elements, electrons are similarly
added to these hydrogen-like
orbitals.
Electron
Filling
Order
-Aufbau
Hund’s Rule
The lowest energy configuration
for an atom is the one having the
maximum number of unpaired
electrons allowed by the Pauli
principle in a particular set of
degenerate orbitals. Orbitals halffill before they completely fill.
Writing Atomic Electron
Configurations
Two ways of writing
configs. One is
called the
electron
configuration
notation.
Electron configuration notation
for H, atomic number = 1
1
1s
value of n
Electron-dot symbol is H.
no. of
electrons
value of l
Writing Atomic Electron
Configurations
Two ways of
writing
configs. Other
is called the
orbital box
notation.
ORBITAL BOX NOTATION
for He, atomic number = 2
2
1s
1s
Arrows
depict
electron
spin
Quantum numbers are an energy address
instead of a positional address.
Electron-dot symbol is He:
Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
3p
3s
Li.
2p
2s
1s
Beryllium
Group 2A
Atomic number = 4
3p
3s
1s22s2 ---> 4 total
electrons
2p
2s
1s
Be:
Boron
Group 3A
Atomic number = 5
3p
1s2 2s2 2p1 --->
3s
2p
2s
1s
5 total electrons
.
B:
Carbon
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
3p
3s
2p
2s
1s
.
.C :
Here we see for the first time
HUND’S RULE. When placing
electrons in a set of orbitals
having the same energy, we
place them singly as long as
possible.
Nitrogen
Group 5A
Atomic number = 7
1s2 2s2 2p3 --->
3p
3s
2p
2s
1s
.
7 total electrons
.N :
.
Oxygen
Group 6A
Atomic number = 8
3p
1s2 2s2 2p4 --->
3s
2p
2s
1s
..
8 total electrons
.O :
.
Fluorine
Group 7A
Atomic number = 9
3p
1s2 2s2 2p5 --->
3s
2p
2s
1s
9 total electrons
..
:F:
.
Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
3p
3s
2p
2s
1s
..
Note that we have reached
the end of the 2nd period,
and the 2nd shell is full!
: Ne :
..
Electron Dot Filling Order
63
4
7
X
58
2
1
Sodium
Group 1A
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
Na.
[Ne] 3s1 (uses rare gas notation)
Note that we have begun a new period.
Aluminum
Group 3A
Atomic number = 13
1s2 2s2 2p6 3s2 3p1
3p
3s
[Ne] 3s2 3p1
.
Al :
2p
2s
1s
Phosphorus
Group 5A
Atomic number = 15
1s2 2s2 2p6 3s2 3p3
[Ne]
3s2 3p3
3p
.
.P :
3s
2p
2s
.
1s
Calcium
Group 2A
Atomic number = 20
6 4s2
1s2 2s2 2p6 3s2 3p
Ca :
[Ar] 4s2
Valence Electrons
The electrons in the outermost principle
quantum level of an atom.
Inner electrons are called core electrons.
Relationship of Electron
Configuration and Region of the
Periodic Table
Green = s block
Yellow = p block
Lt. Blue = d block
Med. Blue = f block
Broad Periodic Table
Classifications
Representative Elements (main group):
filling s and p orbitals (Na, Al, Ne, O)
Transition Elements: filling d orbitals (Fe,
Co, Ni)
Lanthanide and Actinide Series (inner
transition elements): filling 4f and 5f
orbitals (Eu, Am, Es)
Transition Metals
All transition elements have the configuration
(n-1) d. D orbitals are always one behind
their period.
Chromium
Iron
Copper
Transition Element Configurations
3d orbitals used for
Sc - Zn
Lanthanides and Actinides
Rare earth elements always have the configuration
(n-2) f. F orbitals are always two behind their
period.
Cerium
[Xe] 6s2 5d1 4f1
Uranium
[Rn] 7s2 6d1 5f3
Lanthanide Element Configurations
4f orbitals used for
Ce - Lu and 5f for
Th - Lr
Properties of Metals
•
malleable
•
ductile
•
good conductors of heat & electricity
•
tend to lose electrons--oxidation
•
left of zigzag line on periodic table
•
most active metal in lower left corner
(Fr)
Properties of Nonmetals
•
not malleable or ductile
•
brittle
•
nonconductors of heat & electricity
•
tend to gain electrons -- reduction
•
right of zigzag line on periodic table
•
most active nonmetal in upper right
corner (F)
Properties of Metalloids
•
properties intermediate between metals
and nonmetals
•
found bordering zigzag line on periodic
table
•
B, Si, Ge, As, Sb, & Te
ATOMIC ELECTRON CONFIGURATIONS AND
PERIODICITY
Atomic Size
SIZE
Size goes UP on going down a group.
Because electrons are added further
from the nucleus, there is less
attraction.
Size goes DOWN on going across a
period, the addition of protons
pulls electrons tighter.
Atomic Radii
Trends in Atomic Size
Radius (pm)
250
K
1st transition
series
3rd period
200
Na
2nd period
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
Atomic Number
25
30
35
40
Sizes of Transition Elements
3d subshell is inside the 4s subshell.
4s electrons feel a more or less constant Z*.
Sizes stay about the same and chemistries are similar!
Ion Sizes
F,64 pm
9e and 9p
F- , 136 pm
10 e and 9 p
Forming
an anion.
ANIONS are LARGER than the atoms from
which they come.
The electron/proton attraction has gone DOWN
and so size INCREASES.
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 60 pm
2e and 3 p
Forming
a cation.
. CATIONS are SMALLER than the atoms
from which they come.
The electron/proton attraction has gone UP and
so size DECREASES.
Trends in Ion Sizes
Ionization Energy
IE = energy required to remove an
electron from an atom in the gas
phase.
Mg (g) + 738 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
Trends in Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
3
H
Li
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
35
Trends in Ionization Energy
IE increases across a period
because Z* increases.
Metals lose electrons more
easily than nonmetals.
Metals are good reducing
agents.
Nonmetals lose electrons with
difficulty.
Trends in Ionization Energy
IE decreases down a group
Because size increases.
Reducing ability generally
increases down the
periodic table.
Electron Affinity
A few elements GAIN electrons to
form anions.
Electron affinity is the energy
involved when an anion loses an
electron.
A-(g) ---> A(g) + e-
E.A. = DE
Trends in Electron Affinity
Affinity for electron
increases across a
period (EA becomes
more positive).
Affinity decreases down
a group (EA becomes
less positive).
Atom EA
F
+328 kJ
Cl +349 kJ
Br +325 kJ
I
+295 kJ
Trends in Electron Affinity
F Cl
Br
350
300
S
Si
200
Se
150
S4
100
Ge P
S3
Period
50
S2
S1
0
K
1
2
3
4
5
Group
6
7
Electron aff inity ( kJ/mol)
C
H
250
O
electronegativity, ionization energy, ionic radii, electron affinity
atomic radii
ionization energy,
electron affinity,
& electronegativity
Noble
gases
02_29
Alkaline
1 earth metals
Halogens
1A
1
Alkali metals
H
8A
2
13
14
15
16
17
2A
3A
4A
5A
6A
7A
2
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
13
14
15
16
17
18
Al
Si
P
S
Cl
Ar
11
12
Na
Mg
4
3
5
6
9
8
7
Transition metals
10
11
12
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La*
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
104
105
106
107
108
109
110
111
87
88
Fr
Ra
89
Ac†
Unq Unp Unh Uns Uno Une Uun Uuu
*Lanthanides
† Actinides
ionic & atomic
radii
18
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Increasing Periodic Trends
Similarity of Elements
Elements are grouped together in vertical
columns (Groups) that have similar
properties.
Alkali Metals -- Li, Na, K, Rb, & Cs
Halogens -- F2, Cl2, Br2, & I2
Noble Gases -- He, Ne, Ar, Kr, Xe, & Rn
Noble
gases
02_29
Alkaline
1 earth metals
Halogens
1A
1
Alkali metals
H
18
8A
2
13
14
15
16
17
2A
3A
4A
5A
6A
7A
2
He
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
3
4
5
6
7
8
9
Transition metals
10
11
12
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La*
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
104
105
106
107
108
109
110
111
87
88
Fr
Ra
89
Ac†
Unq Unp Unh Uns Uno Une Uun Uuu
*Lanthanides
† Actinides
58
59
60
61
62
63
64
65
66
67
68
69
70
71
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
Er
Tm
Yb
Lu
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
Lr
Periodic Table of the Elements