Ch. 7 Atomic Structure
and Periodicity
7.1 Electromagnetic
Radiation
I. Waves
• A. Energy travels via electromagnetic radiation
(light) with wave-like nature
• B. Wavelength (): distance between two
crests (highest point) or troughs (lowest point), in
meters
II. Frequency
• A. Frequency (): # of wave cycles per second
passing given point (units are Hertz (1/second))
2 cycles per second =
2 Hertz (2 Hz)
3 cycles per second =
3 Hertz (3 Hz)
Longer wavelength =
lower frequency
III. Types
of E.R.
• A. Speed of light is a constant (C) =  • 
• B. C = 3.00 x 108 m/s
• C. Different forms of light have different energy,
wavelength
7.2: I. Nature of Matter
• A. Was thought light only moved in waves so they
have continuous amounts of energy
• B. Found that energy can be lost in whole number
multiples (h).
• C. Planck’s constant (h) = 6.626x10-34 J•sec
• D. “Quantized”: light travels in packets of energy
(“Quantum”)
II. Photons of Energy
• A. Photons: electromagnetic particles
• B. Ephoton = h = hC/
• C. E = mC2
III. Duality of Light
• A. Light can exist as wave and particle simultaneously
• B. Determined through a double slit experiment
IV. De Broglie
• A. De Broglie’s Eqn: found that particles of all
material can exhibit wave-like properties
• B.  = h/mV, where V is velocity
• C. Large objects act as mostly particles
• D. Small objects, mostly waves (light)
• E. Intermediate objects exhibit equal amounts of both
(ex. Electrons)
V. Diffraction
• A. Diffraction: scattering of light when light is
bent by a medium (ex. Through prism)
• B. Diffraction pattern: light scattered as it’s
passed through a crystalline molecule due to light
interference (ex. NaCl)
7.3: I. Atomic Spectrum of Hydrogen
• A. H2 gives off light when given energy
• B. Line spectrum: only color wavelengths that are
emitted are visible
• C. ***Shows that only certain wavelengths of
energy are present in certain elements***
• D. Reinforces “quantized” energy of light
7.4: I. The Bohr Model
• A. Bohr thought that electrons move around
the nucleus in certain allowed circular orbits
• B. Believed that electrons
could only inhabit these fixed
orbits when gaining or losing
energy (cannot be in-between)
• C. Derived calculation to
determine energy levels
• D. E = 2.18x10-18 J (Z2/n2)
• E. Z = nuclear charge, n = integer for orbit #
• F. If we calculate the energy associated with
different orbits for Hydrogen e-s, we get this chart
• G. Notice that higher
orbital levels have more
energy (Potential)
• H. The “ground state” is
the most stable
• I. Returning to the
ground state from higher
levels releases extra
energy as photons of
light
II. In Reality…
• A. Bohr’s model only works for Hydrogen
• B. In fact, electron orbits are not all circular
• C. This model was important though in the
formulation of further theories
7.5: I. Quantum Mechanical Model
• A. Schrödinger and DeBroglie focused on wave
nature of electrons
• B. Related motion of electrons to standing waves
• C. Nodes: fixed points of waves
• D. Antinodes: area of maximum movement
• E. Wave function: represents location of electron
in 3-D space (“Schrödinger equation”)
• F. Orbital: any possible solution of the wave
function, where electrons can actually exist
II. Heisenberg Uncertainty Principle
• A. States that we cannot know the position and
the momentum of a particle at a given time
• B. We cannot figure out how the electron is
moving around the nucleus in its possible
positions, but we can mathematically determine
that they are
• C. ∆x • ∆(mV) ≥ h/4
• D. ∆x = a particles uncertainty, V is velocity, m in
kilograms
• E. The more accurately
we know an electron’s
position, the less
accurately we know its
momentum
• F. Only significant for
small particles
• G. Can’t assume we
know how the
electrons are traveling
around the nucleus
III. Physical Meaning of Wave
Function
• A. The square of the
wave function is the
probability of finding an
electron near a point in
space
• B. We determine electron
locations by probability
distribution or electron
density mapping
• C. This is the first orbital
of Hydrogen called “1s”
7.6: I. Quantum Numbers
• A. Describe various properties of the orbitals
electrons can inhabit
• B. Principal Quantum Number (n): whole
number values, relate to size and energy of
orbital
• C. Higher n values mean larger orbitals (farther
from nucleus), higher energy
• D. Angular Momentum Quantum Number (ℓ):
has values from 0 to n-1 for each Principal
Quantum #, describes shape of orbital
(represented by letters for different shapes)
• E. Magnetic Quantum Number (mℓ): values
between ℓ and -ℓ including 0
• F. Electron Spin Quantum Number (ms):
indicate two directions electrons can spin (+1/2,
-1/2), both Quantum
exist forNumbers
each orbital
n
1
ℓ
0
Orbital type
1s
mℓ
0
# of orbitals
1
2
0
1
2s
2p
0
-1,0,1
1
3
3
0
1
2
3s
3p
3d
0
-1,0,1
-2,-1,0,1,2
1
3
5
4
0
1
2
3
4s
4p
4d
4f
0
-1,0,1
-2,-1,0,1,2
-3,-2,-1,0,1,2,3
1
3
5
7
7.7: I. Orbital Shapes and Energies
• A. Shapes of s,p,d,f
orbitals are best
illustrated by electron
probability diagram
• B. Like a standing wave, orbitals have nodes
and antinodes, nodes are fixed parts (no
electrons), antinodes are where electrons can
exist
II. S Orbitals
• A. Each level up has
one extra area for
electrons to inhabit
• B. Each antinode of
this orbital can hold up
to 2 electrons
III. P Orbitals
• A. Figure-8 shaped
along x, y, and z
axis
• B. Node exists
between two lobes
• C. Each orbital
contains two
electrons (6 total)
• D. Starts at
Quantum level 2
• E. Each orientation
named after axis it goes
along
Ex. 2px, 2py, 2pz
IV. D Orbitals
• A. Has 5 orientations (“clover-leaf”) each
holding 2 electrons (10 total)
• B. Starts at n = 3
V. F Orbitals
• A. Start at n=4, contain 7 orientations, each
holding two electrons (14 total)
7.8: I. Electron Spin and Pauli Principle
• A. Magnetism is caused by moving charges
• B. It was observed that electrons can have two
possible magnetic states meaning that
electrons can spin in one of two directions
C. Pauli Exclusion
Principle says since all
quantum states cannot be
identical for two electrons,
spin state must be
different for electrons in
same orbital
7.9: I. Polyelectronic Atoms
• A. In atoms with more than one electron (all but
Hydrogen) need to consider repulsion between
electrons to determine energy of electrons
• B. Need to treat each electron like it is moving
as a result of nuclear charge and the average
repulsion of other electrons
• C. Electrons are easier to remove if there is
electron repulsion because the positive nuclear
charge won’t have as much influence on
attraction
II. “Effective Nuclear Charge”
• A. Zeff = Zactual – effect of electron repulsion
• B. Z is the number of protons
• C. Zeff can be plugged into the Bohr equation
replacing Z
• D. We need to treat each electron separately to
determine overall effect of nuclear charge
7.10: I. History of the Periodic Table
• A. Mendeleev is one of two people to
independently devise early periodic table
• B. Was organized based on element properties
and atomic mass
• C. Predicted placement of some unknown
elements
• D. Current P.T. based on properties and atomic
number
7.11: I. Aufbau Principle
• A. Electrons are added to different electron
orbitals as they are added to atoms
• B. Orbital Diagram: shows placement of
electrons in orbitals
• C. Rules: 1. Electron order based on energy of
orbitals (relates to order of periodic table)
2. S orbitals have one orbital (2 e-s)
P orbitals have 3 orbitals (6 e-s)
D orbitals have 5 orbitals (10 e-s)
F orbitals have 7 orbitals (14 e-s)
Orbital Order Cheat Sheet
• 7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
3p
2p
7d
6d
5d
4d
3d
7f
6f
5f
4f
Or follow order
from P.T.
Orbital Diagram Rules Continued…
• 3. “Hund’s Rule”: lowest energy state is when
electrons are unpaired in orbital groups before
paired
4. Different electron spins are shown by
up/down arrows
 

• D. Ex. Carbon 1s
2s
2p
• E. Also can write as: 1s2 2s2 2p2
• F. Noble Gas Shortcut: [He] 2s2 2p2
• G. Valence electrons: electrons in outermost
principle quantum level
• H. Ex. Carbon valence e-s are 2s and 2p (4
valence e-s)
• I. Core electrons: innermost electrons
• J. Important Trend of P.T.: elements in same
group have same # of valence e-s
• K. Unique e- configurations: Cr, Cu
• L. Chromium should be: [Ar] 4s2 3d4
Actually: [Ar] 4s1 3d5
• M. Copper should be: [Ar] 4s2 3d9
Actually: [Ar] 4s1 3d10
7.12: I. Development of Polyelectronic Model
• A. Ionization energy equation based on Bohr’s
model multiplied by Avogadro’s number to
determine J/mole or KJ/mole
• B. Eionization = 1.31x106J/mole(z2/n2) or
= 1310 kJ/mole(z2/n2)
• C. We can also calculate the value for the
effective nuclear charge Zeff by plugging in
values and solving for it in Bohr’s equation or
ionization equation
Sample Calc. of Zeff
• Eion for Na = 1.39x105 KJ/mole
• Plug into Bohr or Eion equation, to solve for Zeff
for n = 1
• Eion = 1310 KJ/mole (Zeff2/n2)
• 1.39x105 KJ/mole = 1310 KJ/mole (Zeff2/12)
• Zeff2 = 1.39x105 KJ/mole / 1310 KJ/mole
• Zeff2 = 1.06x102
• Zeff = 10.3
II. Interpretation of Zeff values
• A. For Sodium, the actual Z value is 11
because of the number of protons, yet there is
less of an effective charge (10.3) because of
the electron repulsion
• B. If we calculate the n=3 Zeff for Sodium you
get 1.84 meaning that the farther you get from
the nucleus, the less the effect of the charge of
the nucleus on electrons (which should make
sense)
• C. The inner electrons “Shield” the outer
electrons from the effect of the nuclear charge
7.13: I. Periodic Trends in Atomic
Properties
• A. Ionization Energy: energy required to
remove an electron from an atom or ion that are
in their ground state
• B. First electron to be removed is highest
energy one because it needs least energy to
escape
• C. First ionization energy: energy to remove
first electron
• D. Second ionization energy is always higher
II. Eionization Rationalization
• A. As e- removed atomic charge increases and
holds e- more strongly
• B. Valence e- easier to remove than core ebecause of shielding
• C. On P.T.: 1st Eionization increases as you go left
to right and Bottom to top
• D. Left to right because increasing protons have
more attractive force than e- shielding
• E. Bottom to top because bottom atoms have
farther out orbitals, less nuclear attraction
III. Electron Affinity
• A. Energy change when
adding an e• B. Increase left to right due to
repulsion of e- by pairing them
• C. Ex. Carbon has 2 unpaired 2p e-s, so it can take
in one more without pairing them, but Nitrogen has
3 unpaired 2p e-s and would have to pair them to
become -1 so it releases more energy
• D. No particular trend seen up or down because
you still have similar repulsive forces for groups
IV. Atomic Radius
• A. Increase from right to left (increased nuclear
charge pulls electrons closer)
• B. Increase top to bottom (adding more orbital
shells)
V. Summary of Periodic Trends
Increasing ionization energy, electron affinity
Increasing atomic radii
Increasing
atomic
radii,
decreasing
ionization
energy
VI. Properties of Metals
• A. Metals found on left of table, low ionization
energies, form positive ions by giving up e• B. Non-metals on right of table, high ionization
energies, form anions when adding electrons
• C. Metalloids or semi-metals: have both
metallic and non-metallic properties based on
conditions
VII. Alkali Metals
• A. Li, Na, K, Rb, Cs, and Fr most chemically
reactive of all metals
• B. Hydrogen too small to act as metal because
single electron bound so tightly to nucleus
• C. Low ionization energies allow e- to be lost easily
• D. Hydration energy: energy released when water
attached
• E. Lithium has most because it is small; less
shielding effect of e-; more effective nuclear charge
for water to grab onto (called “charge density”)