Chapter
5
Periodicity & Atomic Structure
Chemistry, 4th Edition
McMurry/Fay
The Periodic Table
•
The periodic table is the most important organizing
principle in chemistry.
•
Chemical and physical properties of elements in
the same group are similar.
•
All chemical and physical properties vary in a
periodic manner, hence the name periodic table.
2
The Periodic Table
3
The Periodic Table
4
The Periodic Table
5
Electromagnetic Radiation
Electromagnetic Radiation:
Energy propagated by an electromagnetic field.
Electromagnetic radiation has both particle and
wave nature.
6
Electromagnetic Radiation
Spectroscopy:
Branch of physical science that deals with the interaction
of electromagnetic radiation with matter
Spectrometry:
The quantitative measurement of the intensity of
radiation at a particular wavelength of light.
7
Wave-Like Nature of Light
Frequency (, Greek nu): Number of peaks that pass a
given point per unit time.
Wavelength (, Greek lambda): Distance from one
wave peak to the next.
Amplitude: Height measured from the center of the wave.
The square of the amplitude gives intensity.
8
Wave-Like Nature of Light
9
Wave-Like Nature of Light
•
Speed of a wave is the wavelength (in meters)
multiplied by its frequency in reciprocal seconds.
Wavelength x Frequency = Speed
 (m)
x
 (s–1) = c (m/s–1)
10
Wave-Like Nature of Light
11
Particle-Like Nature of Light
Electromagnetic radiation can be described as a
stream of tiny particles, called photons, with a very
small mass and a very large velocity.
The velocity of photons traveling in a vacuum is:
c = 3.00 x 108 m/s
12
Particle-Like Nature of Light
Where does a photon come from?
One photon is emitted when one atom or molecule in
an excited state relaxes to the ground state via the
emission of radiation.
E=hν
13
Atomic Spectra
•
Atomic spectra:
Result from excited
atoms emitting light.
•
Line spectra: Result
from electron
transitions between
specific energy
levels.
14
Atomic Spectra
15
Atomic Spectra
•
Blackbody radiation is the visible glow that solid
objects emit when heated.
•
Max Planck (1858–1947): proposed the energy is
only emitted in discrete packets called quanta.
•
The amount of energy depends on the frequency:
E  h 
hc

h  6.626  10 34 J  s
16
Atomic Spectra
Albert Einstein (1879–1955):
• Used the idea of
quanta to explain the
photoelectric effect.
•
•
He proposed that
light behaves as a
stream of particles
called photons.
17
Atomic Spectra
•
A photon’s energy
must exceed a
minimum threshold
for electrons to be
ejected.
•
Energy of a photon
depends only on
the frequency.
18
Atomic Spectra
•
For red light with a wavelength of about 630 nm,
what is the energy of a single photon and one mole
of photons?
E  h 
hc

h  6.626  10 34 J  s
19
Wave–Particle Duality
•
Louis de Broglie (1892–1987): Suggested waves
can behave as particles and particles can behave
as waves. This is called wave–particle duality.
For Light :  
For a Particle :  
h
mc
h
mv


h
p
h
p
20
Quantum Mechanics
•
Niels Bohr (1885–1962): Described atom as
electrons circling around a nucleus and concluded
that electrons have specific energy levels.
•
Erwin Schrödinger (1887–1961): Proposed
quantum mechanical model of atom, which focuses
on wavelike properties of electrons.
21
Quantum Mechanics
•
Werner Heisenberg (1901–1976): Showed that it
is impossible to know (or measure) precisely both
the position and velocity (or the momentum) at the
same time.
•
The simple act of “seeing” an electron would
change its energy and therefore its position.
22
Quantum Mechanics
h
Heisenberg Uncertainty P rinciple: (x)(m ) 
4
h
Uncertainty in electron's position: (x) 
(4 )(m )
23
Quantum Mechanics
•
Erwin Schrödinger (1887–1961): Developed a
compromise which calculates both the energy of an
electron and the probability of finding an electron at
any point in the molecule.
•
This is accomplished by solving the Schrödinger
equation, resulting in the wave function, .
24
Quantum Numbers
•
Wave functions describe the behavior of electrons.
•
Each wave function contains three variables called
quantum numbers:
• Principal Quantum Number (n)
• Angular-Momentum Quantum Number (l)
• Magnetic Quantum Number (ml)
25
Quantum Numbers
•
Principal Quantum Number (n): Defines the size
and energy level of the orbital. n = 1, 2, 3, 
•
As n increases, the electrons get farther from the
nucleus.
•
As n increases, the electrons’ energy increases.
•
Each value of n is generally called a shell.
26
Quantum Numbers
•
Angular-Momentum Quantum Number (l):
Defines the three-dimensional shape of the orbital.
•
For an orbital of principal quantum number n, the
value of l can have an integer value from 0 to n – 1.
•
This gives the subshell notation:
l = 0 = s orbital
l = 1 = p orbital
l = 2 = d orbital
l = 3 = f orbital
l = 4 = g orbital
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Quantum Numbers
•
Magnetic Quantum Number (ml): Defines the
spatial orientation of the orbital.
•
For orbital of angular-momentum quantum number, l,
the value of ml has integer values from –l to +l.
•
This gives a spatial orientation of:
l = 0 giving ml = 0
l = 1 giving ml = –1, 0, +1
l = 2 giving ml = –2, –1, 0, 1, 2,
and so on…...
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Quantum Numbers
•
Spin Quantum Number:
•
The Pauli Exclusion
Principle states that no
two electrons can have
the same four quantum
numbers.
29
Quantum Numbers
30
Electron Radial Distribution
31
Electron Radial Distribution
•
s Orbital Shapes:
32
Electron Radial Distribution
•
p Orbital Shapes:
33
Electron Radial Distribution
•
d and f Orbital Shapes:
34
Effective Nuclear Charge
•
Electron shielding
leads to energy
differences among
orbitals within a shell.
•
Net nuclear charge
felt by an electron is
called the effective
nuclear charge (Zeff).
35
Effective Nuclear Charge
•
Zeff is lower than actual
nuclear charge.
•
Zeff increases
toward nucleus
ns > np > nd > nf
•
This explains certain periodic
changes observed.
36
Effective Nuclear Charge
37
Electron Configuration of Atoms
•
Pauli Exclusion Principle: No two electrons in an
atom can have the same quantum numbers (n, l,
ml, ms).
•
Hund’s Rule: When filling orbitals in the same
subshell, maximize the number of parallel spins.
38
Electron Configuration of Atoms
•
Rules of Aufbau Principle:
1.
Lower n orbitals fill first.
2.
Each orbital holds
two electrons; each
with different ms.
3.
Half-fill degenerate
orbitals before pairing
electrons.
39
Electron Configuration of Atoms
Assigning Electrons to Atomic Orbitals
1. The number of electrons in an atom is equal to the atomic
number.
2. Assign electrons to the lowest energy orbitals first, then
build up.
40
Electron Configuration of Atoms
Assigning Electrons to Atomic Orbitals
3. No more than 2 electrons can occupy a single orbital:
their spins must be paired.
4. If more than one orbital is available at the same energy,
add single electrons with the same spin to each orbital
before adding two electrons to one orbital.
5. Use the periodic table as a guide.
41
Electron Configuration of Atoms
Writing Electron Configurations
Name the occupied atomic orbitals in the atom with the
number of electrons in each orbital written as a superscript.
Li: 1s22s1
Na: 1s22s22p63s1
Fe: 1s22s22p63s23p64s23d6
42
Electron Configuration of Atoms
Writing Electron Configurations
One may also write the configuation as a noble gas closed
shell plus the valence electrons present in the atom.
Li: [He]2s1
Na: [Ne]3s1
Fe: [Ar]4s23d6
43
Electron Configuration of Atoms
Increasing Energy
Core
[He]
[Ne]
[Ar]
[Kr]
[Xe]
[Rn]
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d 4f
5d 5f
6d
44
Electron Configuration of Atoms

1s

2s
1s2 2s1
Be 
1s

2s
1s2 2s2
Li
B
  
1s 2s 2px 2py 2pz
1s2 2s2 2p1
C
  

1s 2s 2px 2py 2pz
1s2 2s2 2p2
45
Electron Configuration of Atoms
N
    
1s 2s 2px 2py 2pz
1s2 2s2 2p3
O
    
1s 2s 2px 2py 2pz
1s2 2s2 2p4
Ne     
1s 2s 2px 2py 2pz
S
[Ne]    
3s 3px 3py 3pz
1s2 2s2 2p5
[Ne] 3s2 3p4
46
Electron Configuration of Atoms
47
Electron Configuration of Atoms
48
Electron Configuration of Atoms
•
Anomalous Electron Configurations: Result from
unusual stability of half-filled & full-filled subshells.
•
Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5
•
Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10
•
In the second transition series this is even more
pronounced, with Nb, Mo, Ru, Rh, Pd, and Ag having
anomalous configurations (Figure 5.20).
49
Periodic Properties
50
Electron Configuration of Atoms
Metallic Radius: One half of the distance
between neighboring atoms in a solid sample.
Predicting Relative Atomic Radii:
1. The atom with the largest n is largest.
2. If n is equal, then the atom with the largest
nuclear charge is smallest.
51
Atomic Radii
52
Atomic Radii
53
Atomic Radii
54