ATOMIC STRUCTURE &
PERIODICITY
Chapter 7
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.
Electromagnetic
Radiation
Radiant energy that exhibits
wavelength-like behavior and
travels through space at the speed
of light in a vacuum.
Electromagnetic Radiation
wavelength
Visible light
Amplitude
wavelength
Ultaviolet 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
.
Wavelength and frequency can be interconverted.
c=
 = frequency (s1, Hz, cyc/s, or waves/s )
 = wavelength (m)
c = speed of light (m/s)
Quantization of Energy
Max Planck (1858-1947)
An object can gain or lose energy by absorbing
or emitting radiant energy in QUANTA.
Planck’s Constant
Transfer of energy is quantized, and can only
occur in discrete units, called quanta.
E = h =
hc

E = change in energy, in J
h = Planck’s constant, 6.626  1034 J s
 = frequency, in s1
 = wavelength, in m
Photoelectric Effect
A. Einstein (1879-1955)
Experiment demonstrates the particle nature of
light.
Classical theory said that E of ejected electron
should increase with increase in light
intensity—not observed!
No e- observed until light of a certain minimum
E is used.
Number of e- ejected depends on light intensity.
Energy and Mass
Energy has mass
E = mc2
E = energy
m = mass
c = speed of light
Energy and Mass
Ephoton =
mphoton
hc

h
=
c
(Hence the dual nature of light or
wave-particle duality.)
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.
Line (discrete) spectrum: Contains only
some of the wavelengths of 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.
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
SHARP LINE SPECTRA
of excited atoms.
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!
Atomic Spectra and Bohr
Bohr said classical view is wrong.
Need a new theory — now called QUANTUM or WAVE
MECHANICS.
e- can only exist in certain discrete orbits — called
stationary states.
e- is restricted to QUANTIZED energy states.
Energy of state = - C2
n
where C is a constant &
where n = quantum no. = 1, 2, 3, 4, ....
The Bohr Model
The electron in a hydrogen atom moves around the
nucleus only in certain allowed circular orbits.
2

z 
-18
E = -2.178 × 10 J  2 
n 
E = energy of the levels in the H-atom
z = nuclear charge (for H, z = 1)
n = an integer
E becomes more negative as electrons move closer
to the nucleus.
The Bohr Model
Ground State: The lowest
possible energy state for an
atom (n = 1).
Higher energy electrons are in outer shells, and
are easier to remove because they are more
shielded from the positive nucleus by inner shell
electrons.
Energy Changes in the
Hydrogen Atom
E = Efinal state  Einitial state
hc
 =
E
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-electrons do not move in circular orbits
about the nucleus.
Quantum or Wave
Mechanics
L. de Broglie
(1892-1987)
de Broglie (1924)
proposed that all
moving objects have
wave properties.
Wavelength and Mass
de Broglie’s Equation
h
 =
m
 = wavelength, in m
h = Planck’s constant, 6.626  1034 J s =
kg m2 s1
m = mass, in kg
v = velocity in m/s
Quantum or Wave
Mechanics
See Sample Exercise 7.3 on page 298.
Baseball (0.100 kg) at 35
m
s
 = 1.9 x 10-34 m
e- with velocity = 1.0 x
 = 7.27 x 10-11 m
m
7
10
s
Quantum or Wave
Mechanics
E. Schrodinger
1887-1961
Schrodinger applied idea
of e- behaving as a
wave to the problem
of electrons in atoms.
Quantum Mechanics
The Wave Equation
Based on the wave properties of the atom
H  = E
 = wave function
H = mathematical operator
E = total energy of the atom
A specific wave function is often called an
orbital. Electrons move in orbitals not
Bohr orbits.
Heisenberg Uncertainty
Principle
Problem of defining nature of
electrons in atoms solved by
W. Heisenberg.
W. Heisenberg
1901-1976
Cannot simultaneously define the
position and momentum (=
m•v) of an electron.
We define e- energy exactly but
accept limitation that we do
not know exact position.
Heisenberg Uncertainty
Principle
 x   mv  
h
4
x = position
mv = momentum
h = Planck’s constant
The more accurately we know a
particle’s position, the less accurately we can
know its momentum.
Probability Distribution
square of the wave function
probability of finding an electron at a
given position
Radial probability distribution is the
probability distribution in each spherical
shell.
Quantum Numbers (QN)
1. Principal QN (n = 1, 2, 3, . . .) - related to size
and energy of the orbital.
2. Angular Momentum QN (l = 0 to n  1) - relates
to shape of the orbital.
3. Magnetic QN (ml = l to l) - 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.
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.
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.
1s Orbital
2s Orbital
3s Orbital
p Orbitals
Typical p orbital
When n = 2, then l = 0 and 1
Therefore, in n = 2 shell there are 2
types of orbitals — 2 subshells
For l = 0
ml = 0
this is a s subshell
For l = 1
ml = -1, 0, +1
this is a p subshell
with 3 orbitals
planar node
When l = 1, there is
a
PLANAR NODE
thru
the nucleus.
p Orbitals
A p orbital
The three p
orbitals lie 90o
apart in space
2px Orbital
2py Orbital
2pz Orbital
3px Orbital
3py Orbital
3pz Orbital
d Orbitals
When n = 3, what are the values of l?
l = 0, 1, 2
and so there are 3 subshells in the shell.
For l = 0, ml = 0
---> s subshell with single orbital
For l = 1, ml = -1, 0, +1
---> p subshell with 3 orbitals
For l = 2, ml = -2, -1, 0, +1, +2
---> d subshell with 5 orbitals
d Orbitals
typical d orbital
s orbitals have no planar node (l
= 0) and so are spherical.
p orbitals have l = 1, and have 1
planar node,
and so are “dumbbell” shaped.
This means d orbitals (with l = 2)
have
2 planar nodes
planar node
planar node
3dxy Orbital
3dxz Orbital
3dyz Orbital
3dyz Orbital
2
2
3dx - y
Orbital
f Orbitals
When n = 4, l = 0, 1, 2, 3 so there are 4 subshells in the
shell.
For l = 0, ml = 0
---> s subshell with single orbital
For l = 1, ml = -1, 0, +1
---> p subshell with 3 orbitals
For l = 2, ml = -2, -1, 0, +1, +2
---> d subshell with 5 orbitals
For l = 3, ml = -3, -2, -1, 0, +1, +2, +3
---> f subshell with 7 orbitals
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.
Valence Electrons
The electrons in the outermost principle
quantum level of an atom.
Atom
Valence Electrons
Ca
2
N
5
Br
7
Inner electrons are called core electrons.
ATOMIC ELECTRON CONFIGURATIONS AND
PERIODICITY
Arrangement of
Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
Assigning Electrons to Atoms
Electrons generally assigned to orbitals of
successively higher energy.
For H atoms,
E=-
C
n2
. E depends only on n.
For many-electron atoms, energy depends on both n
and l.
Effective Nuclear Charge, Zeff
Zeff is the nuclear charge experienced by the outermost
electrons. Explains why E(2s) < E(2p)
Zeff increases across a period owing to incomplete
shielding by inner electrons.
Estimate Zeff by --> [ Z - (no. inner electrons) ]
Charge felt by 2s e- in Li
Be
Zeff = 4 - 2 = 2
B
Zeff = 5 - 2 = 3
Zeff = 3 - 2 = 1
and so on!
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
Arrows
2
depict
electron
spin
1s
1s
One electron has n = 1, l = 0, ml = 0, ms = + 1/2
Other electron has n = 1, l = 0, ml = 0, ms = - 1/2
Lithium
Group 1A
Atomic number = 3
3p
1s22s1 ---> 3 total
electrons
3s
2p
2s
1s
Li 
Beryllium
Group 2A
Atomic number = 4
3p
3s
1s22s2 ---> 4 total
electrons
2p
2s
1s
Be 
Boron
Group 3A
Atomic number = 5
3p
3s
1s2 2s2 2p1 --->
2p
2s
1s

B 
5 total
electrons
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
3p
1s2 2s2 2p3 --->
3s
2p
2s
1s


N
 

7 total electrons
Oxygen
Group 6A
Atomic number = 8
3p
1s2 2s2 2p4 --->
2p
8 total
electrons
3s
2s
1s

 O 

Fluorine
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
3p
3s
2p
2s
1s

 F 

9 total
electrons
Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
3p
3s
2p
2s
1s

 Ne 

Note that we have reached
the end of the 2nd period,
and the 2nd shell is full!
Electron Dot Filling Order
6
3
1
4
2
7
58
X
Sodium
Group 1A
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Note that we have begun a new period.
All Group 1A elements have [core]ns1
configurations.
Aluminum
Group 3A
Atomic number = 13
1s2 2s2 2p6 3s2 3p1
[Ne] 3s2 3p1
All Group 3A elements
have
[core] ns2 np1
configurations where n is
the period number.
3p
3s
2p
2s
1s
Phosphorus
Group 5A
Atomic number = 15
1s2 2s2 2p6 3s2 3p3
[Ne] 3s2 3p3
All Group 5A
elements have
[core ] ns2 np3
configurations
where n is the period
number.
3p
3s
2p
2s
1s
Calcium
Group 2A
Atomic number = 20
1s2 2s2 2p6 3s2 3p6 4s2
[Ar] 4s2
All Group 2A elements have [core]ns2
configurations where n is the period
number.
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 4th period elements have the configuration
[argon] nsx (n - 1)dy and so are “d-block”
elements.
Chromium
Iron
Copper
Transition Element Configurations
3d orbitals used for
Sc - Zn
Lanthanides and Actinides
All these elements have the configuration [core] nsx (n 1)dy (n - 2)fz and so are “f-block” elements.
Cerium
[Xe] 6s2 5d1 4f1
Uranium
[Rn] 7s2 6d1 5f3
Lanthanide Element Configurations
4f orbitals used for
Ce - Lu and 5f for
Th - Lr
Ion Configurations
To form cations from elements remove 1 or
more e- from subshell of highest n [or
highest (n + l)].
P [Ne] 3s2 3p3 - 3e- ---> P3+ [Ne] 3s2 3p0
Ion Configurations
To form cations from elements remove 1 or more efrom subshell of highest n [or highest (n + l)].
P [Ne] 3s2 3p3 - 3e- ---> P3+ [Ne] 3s2 3p0
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
Ion Configurations
For transition metals, remove ns electrons and
then (n - 1) electrons.
Fe [Ar] 4s2 3d6
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
loses 3 electrons ---> Fe3+ [Ar] 4so 3d5
Information Contained in the
Periodic Table
1. Each group member has the same valence
electron configuration (these electrons primarily
determine an atom’s chemistry).
2. The electron configuration of any representative
element.
3. Certain groups have special names (alkali
metals, halogens, etc).
4. Metals and nonmetals are characterized by their
chemical and physical properties.
General Periodic Trends
Higher Z*.
Electrons held
more tightly.
Larger orbitals.
Electrons held less
tightly.
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 UP on going across a
period.
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
25
Atomic Number
30
35
40
Sizes of Transition Elements
3d subshell is inside the 4s subshell.
4s electrons feel a more or less constant Zeff.
Sizes stay about the same and chemistries are similar!
Ion Sizes
F,64 pm
9e and 9p
Does -the size go up or
down when gaining
F , 136 pm
electron
to form an
10 ean
and
9p
anion?
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.
Trends in ion sizes are the same as atom sizes.
Ion Sizes
Li,152 pm
3e and 3p
Does+ the size go
up+ or down
Li , 60 pm
when
an
2e and 3losing
p
electron to form
a cation?
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
Redox
Reactions
Why do metals lose
electrons in their
reactions?
Why does Mg form Mg2+
ions and not Mg3+?
Why do nonmetals take on
electrons?
Ionization Energy
IE = energy required to remove an
electron from an atom in the gas
phase.
Mg (g) + 738 kJ → Mg+ (g) + e-
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 Zeff 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.
2nd IE / 1st IE
Li
Na
K
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. = E
Electron Affinity of Oxygen
E is ENDOthermic
because O has an
affinity for an e-.
O- ion [He] 
 

- electron
O atom [He] 
 
EA = + 141 kJ

Electron Affinity of Nitrogen
N- ion [He] 
E is zero for N-
due to
electronelectron
repulsions.
 

- electron
N atom [He] 


EA = 0 kJ

Trends in Electron Affinity
Affinity for electron
increases across a
period (EA becomes
more negative).
Affinity decreases down a
group (EA becomes
less negative).
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
35 0
30 0
S
Si
20 0
Se
15 0
S4
10 0
Ge P
S3
Period
50
S2
S1
0
K
1
2
3
4
5
Group
6
7
Elec tron affinity (kJ/m ol)
C
H
25 0
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
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