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AP Chemistry
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CHAPTER 7
ATOMIC STRUCTURE and PERIODICITY
HW: Warning: If you do not normally read the chapter, please start NOW with chapter 7. This
chapter is very conceptual (vs. plug and chug), so complete understanding will require more than
simply completing the HW assignment.
p.336 # 12, 14, 18, 24, 26, 27, 31, 40, 44, 62, 64, 68, 74, 78, 80, 84, 86, 90, 94, 96, 98, 106, 114,
118, 124
BOHR MODEL LAB
PERIODIC TABLE ASSIGNMENT #2
Consult chapter 7, 10, 19, 20, 21
Add the following information to your large periodic table.
1. Below the appropriate groups, label the following electron configuration endings: s 1, s2, p1,
p2,p3, p4,p5,p6,d1-d10, f1-f14
2. Write chromium’s electron configuration below the element symbol.
3. Write copper’s electron configuration below the element symbol.
4. Draw the trend for increasing ionization energy on your table (or below it in a small box).
5. Draw the trend for increasing electron affinity on your table (or below it in a small box).
6. Draw the trend for increasing electronegativity on your table.
7. Find the most electronegative atom on the table and write its electronegativity value (from the
Pauling scale) in its box.
8. Find the least electronegative atom on the table and do the same as in the last exercise.
9. Find an element that is a yellow crystalline solid and color its box yellow.
10. Find the element that is gold and color its box gold.
11. Find the element known as “quicksilver” and write its name and symbol and color it silver.
12. Find the element that is a violet – black crystal and color its box purple.
13. Find the element that is a red-brown liquid and color its box accordingly.
14. Find the element that is a yellow-green poisonous gas and color its box accordingly.
15. Find three noble gases that will actually form compounds. List a few of their compounds in
their boxes.
16. Find a noble gas that is a major component of the sun and whose nucleus is known as an
alpha particle. Write and alpha symbol in its box.
17. Find an element that can form a peroxide or a superoxide and write the general formulas for
each in its box.
18. This metallic element is prepared from bauxite. Find it and write the formula for bauxite in its
box.
19. This metallic element melts at 30ºC. Write this fact in its box.
20. This metallic element is prepared from cassiterite. Find it and write the formula for cassiterite
in its box.
21. This metallic element can be prepared from galena. Find it and write the formula for galena in
its box.
22. This element is the most abundant in the earth’s crust. Write crust in its box.
23. This element is the most abundant in the atmosphere. Write atmosphere in its box.
24. This element comes in black, white, or red. Color it accordingly.
25. This element has forms called monoclinic and rhombic. List these terms in its box. Color the
box appropriately.
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AP Chemistry
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CHAPTER OVERVIEW
(7.1-7.2) Electromagnetic radiation, solving for wavelength and frequency, quantized energy,
Debroglie relationship for calculating wavelength of a particle, particle wave duality and continuous
vs. discrete line spectrum.
(7.3-7.4) The Bohr model of the atom, energy calculations for electron transitions. (see bohr lab)
(7.5) Quantum mechanical view of the atom, Heisenberg uncertainty principle, electron probability
distributions.
(7.6-7.7) Quantum numbers n, l, ml and ms and orbital shapes and energies.
(7.8-7.11) Pauli exclusion, Hund’s rule and aufbau principle.
(7.12) Shielding effect, Z effective (eff. Nuclear charge), ionization energy, orbital filling across a
period.
(7.13) Trends of the first and second ionization energy, electron affinity, electronegativity, and atomic
radius.
UNDERSTANDING BOHR’s MODEL of the ATOM
To understand the ring model that Bohr proposed, we have to
understand how an electron is moving.
PARTICLE WAVE DUALITY
Who? Louis DeBroglie
What? The electron can travel as a particle or as a wave.
All matter has particle/ wave like properties. Some have such a
small wavelength that we don’t notice.
When? 1923
WAVE MOTION
 Characteristic of all electromagnetic radiation (EMR)
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ROYGBIV
 Wave motion is described by:
1. Wavelength
Defined as: the distance between two crests of a wave
Symbol:  (lambda)
Units: m or nanometers
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2. Amplitude
Defined as: height of the wave (from rest to crest)
Symbol: A
Units: m
3. Frequency Rays
Defined as: the number of waves that pass per second
Symbol:  (nu)
Units: hertz
1 hz = 1/ second
106 hz = 1 Megahertz
 All EMR travels at the speed of light.
c = 2.9979 x 108 m/s
we will use 3.00 x 108 m/s
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 Relationship between c, , and 

c = 
Ex. What is the wavelength of light that has a frequency of 5 Hz?
Is this visible?
Ex. What is the frequency of blue light with a wavelength of 484
nm?
PRIOR to DeBROGLIE:
Matter and energy were seen as different from each other in
fundamental ways. Matter was treated as a particle. Energy could
come in waves, with any frequency. Until,
Who? Max Planck
What? Found that the colors of light emitted from hot objects
(heated to incandescence) couldn’t be explained by viewing energy
as a wave.
Instead, he proposed that light was given off in the form of photons
with a discrete amount of energy called quanta.
When? 1900
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HOW? can the energy of a photon can be calculated?
E = h
Where:
E is energy
h is Planck’s constant = 6.626 x 10-34 Joule seconds
 is frequency
Ex. What is the energy associated with light with a frequency of
6.65 x 108 / second?
Who? Einstein
What? Said electromagnetic radiation is quantized in particles
called photons.
When? 1905
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AP Chemistry
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Each photon has energy E = h = hc/
Combine this with E = mc2
m = h / (c)
or
 = h/ mv (careful – velocity)
Which is it? Particle Wave Duality
 Is energy a wave like light, or a particle ?
 Does matter a wavelength?
Treating matter as a wave:
 Use the velocity v to find wavelength
DeBroglie’s equation:  = h/ mv
Ex. Sodium atoms have a characteristic color when excited in a
flame. The color comes from the emission of light of 589.0 nm.
What is the frequency of this light ?
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AP Chemistry
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What is the energy of a photon of this light ?
What is the apparent mass of a photon of this light ?
What is the energy of a mole of these photons?
What is the wavelength of an electron travelling at 1.0 x 107 m/s?
Mass of e-1 = 9.11 x 10-31 kg
What is the wavelength of a softball with a mass of 0.10 kg
moving at 99 mi/hr?
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Continuous Spectrum
 The range of frequencies present in light.
 White light has a continuous spectrum.
 All the colors are possible.
 A rainbow can be seen through a spectroscope or prism.
Hydrogen spectrum
 Emission spectrum because these are the colors it gives off or
emits.
 Called a line spectrum.
 There are just a few discrete lines showing. What this means:
o Only certain energies are allowed for the hydrogen atom.
o Can only give off certain energies.
o Energy in the in the atom is quantized.
 Use E = h = hc /  486 nm
656 nm
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Who? Niels Bohr
What? Developed the quantum model of the hydrogen atom.
He said the atom was like a solar system.
The electrons were attracted to the nucleus because of opposite
charges.
Didn’t fall in to the nucleus because it was moving around.
The Bohr Ring Atom
He didn’t know why but only certain energies were allowed.
He called these allowed energies energy levels.
Putting Energy into the atom moved the electron away from the
nucleus from ground state to excited state.
When it returns to ground state it gives off light of a certain energy.
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The Bohr Model
 n is the energy level
 n = 1 is called the ground state
 Z is the nuclear charge, which is +1 for hydrogen.
 For each energy level the energy is:
E = -2.178 x 10-18 J (Z2 / n2)
 When the electron is removed, n =
,E=0
 We are worried about the change when the electron moves
from one energy level to another.
ΔE = E final – E initial
ΔE = -2.178 x 10-18J Z2 (1/ nf2 - 1/ ni2)
\\\\
2)
Ex. Calculate the energy need to move an electron from its ground
state to the third energy level.
Ex. Calculate the energy released when an electron moves from n=
4 to n=2 in a hydrogen atom.
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Ex. Calculate the energy released when an electron moves from n=
5 to n=3 in a He+1 ion
When is it true?
 Only for hydrogen atoms and other monoelectronic species.
 Why the negative sign?
o To increase the energy of the electron you make it closer
to the nucleus.
o n the maximum energy an electron can have is zero, at an
infinite distance.
QUANTUM MECHANICS
The mathematical relationships predicted by BOHR (and
demonstrated in our investigation)




successfully predict wavelengths of light emitted for an electron
transitioning between two energy levels within the hydrogen atom
predict the most probable radius of the energy levels from
nucleus
This model fails when applied to POLYELECTRONIC systems
(atoms with more than one e-).
e- interactions and Z (the nuclear charge) make it impossible to
apply BOHR’s relationship
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QUANTUM MECHANICAL VIEW OF THE ATOM
Also known as the WAVE MECHANICAL view.
Predicted by:



HEISENBERG
DeBROGLIE
SCHRODINGER
Premises:


e- is a particle that can travel as a WAVE (DeBroglie relationship
= h/mv)
waves have only some allowable energy levels (corresponding to
n= 1, 2, etc. in the H atom)…these allowable energy levels are
called QUANTUM LEVELS.
Let’s look at a wave pattern between two fixed points (like an
electron traveling between two walls) or a guitar string. This is
known as a STANDING wave.
There are only certain frequencies at which the wave can travel
because the ends are fixed.
Set frequencies mean set
WAVELENGTHS.
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


When a wave is set up, it can be defined by it’s number of NODES.
These are areas when the wave goes from + to – in value.
 0 nodes = the LONGEST wavelength. It has the lowest
FREQUENCY and the longest wavelength. This is known as
the GROUND STATE in the atom. This is the n=1 level.
All other frequencies will be a multiple of the fundamental
frequency.
 1 node = the first HARMONIC.
ENERGY LEVEL.
Also known as the n= 2
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AP Chemistry
 2 nodes =
And so on…
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the second HARMONIC.
So, both the WAVELENGTH and the FREQUENCY of the trapped
electron are discrete or QUANTIZED: meaning there are only
certain allowable energy states and nothing in between.
HOW CAN THE ELECTRON’s POSTION or MOTION BE
DESCRIBED?
 = the WAVE FUNCTION, which tells the 3-D coordinates of the
e- position.
 is part of the SCHRODINGER equation.
- h 2 d2 = E 
2 m dx2
where h is a modification of Planck’s constant = h / 2 = 1.05457 x
10-34 Js
m = the mass of the particle
E is the energy of the wave function
and the d2 term means to take the 2nd derivative of the function
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The solution of the calculus based equation results in 4 QUANTUM
numbers, which tell us something about the electron’s behavior.
More on these later.
2 = the probability of finding an electron in a particular point in
space called an ORBITAL.
Can be shown as a PROBABILITY distribution. Where the highest
point is the most likely distance from the nucleus to find the
electron. When n= 1 this distance also coincides with the first
“orbit” predicted by BOHR.
Unfortunately, the electron’s position and momentum cannot be
known at the same time.
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This is the HEISENBERG UNCERTAINTY principle.
 x   mv = h / 4
Where x is the uncertainty about the POSITION.
 mv is the uncertainty about the MOMENTUM.
So, the more you know about the electron’s position, the less you
can know about its movement (momentum). In macroscopic
systems, this uncertainty is negligible.
QUANTUM NUMBERS
Principle quantum number
Symbol:
What does it tell about the electron?
Energy level
Values
Angular quantum number
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Symbol:
What does it tell about the electron?
Orbital shape
Values:
Magnetic
Symbol:
What does it tell about the electron?
Values:
Spin
Symbol:
What does it tell about the electron?
Values:
ENERGY LEVEL DIAGRAMS
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RULES for FILLING the DIAGRAM
 Aufbau – fill orbitals in lower energy levels before proceeding
to the next level.
 Hund’s Rule- Place electrons in separate orbitals before
pairing them within the same energy level.
 Pauli exclusion principle – every electron must have a different
set of quantum numbers. Electrons in the same orbital must
have opposite spins.
(insert sample energy level diagrams and quantum number determinations here)
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ELECTRON CONFIGURATION
 Shows the filled orbitals in short hand notation.
Ex. Mg
Ex. Cl
NOBLE GAS electron configuration: shows the noble gas core to
simplify electron configuration.
 Focuses on valence electrons - the electrons in the outermost
energy levels (not including d).
Ex. O
Ex. Br
Ex. U
How is electron configuration related to the periodic table?.
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 Elements in the same column have the same electron
configuration.
 Put in columns because of similar properties.
 Similar properties because of electron configuration.
 Noble gases have filled energy levels.
 Transition metals are filling the d orbitals
 Exceptions to filling rules:
Ti = [Ar] 4s2 3d2
V = [Ar] 4s2 3d3
Cr = [Ar] 4s13d5
Cu=[Ar] 4s13d10
Mn = [Ar] 4s23d5
 These have half filled orbitals.
 Scientists aren’t sure of why it happens. Leads to stability
due to minimizing electron repulsions.
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We can use Zeff to predict properties, if we determine its pattern on
the periodic table. Can use the amount of energy it takes to
remove an electron for this.
Ionization Energy- The energy necessary to remove an electron
from a gaseous atom.
Remember this:
E = -2.18 x 10-18 J(Z2/n2) was true for Bohr atom.
Can be derived from quantum mechanical model as well for a mole
of electrons being removed
E =(6.02 x 1023/mol) x 2.18 x 10-18 J(Z2/n2)
E= 1.13 x 106 J/mol(Z2/n2)
E= 1310 kJ/mol(Z2/n2)
Example
Calculate the ionization energy of B+4
Ionization energy =
1310 kJ/mol(Zeff 2/n2)
So we can measure Zeff
 The ionization energy for a 1s electron from sodium is 1.39 x
105 kJ/mol .
 The ionization energy for a 3s electron from sodium is 4.95 x
102 kJ/mol .
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Why?
SHIELDING
 Electrons on the higher energy levels tend to be farther out.
 Have to “look through” the other electrons to see the nucleus.
 They are less affected by the nucleus.
 Lower effective nuclear charge (Z eff).
If shielding were completely effective,
Zeff = 1
Why isn’t it?
Penetration
There are levels to the electron distribution for each
orbital.2s
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PENETRATION EFFECT
 The outer energy levels penetrate the inner levels so the
shielding of the core electrons is not totally effective.
 From most penetration to least penetration the order is
ns > np > nd > nf (within the same energy level).
This is what gives us our order of filling, electrons prefer s and p.
How do orbitals differ?
 The more positive the nucleus, the smaller the orbital.
A sodium 1s orbital is the same shape as a hydrogen 1s orbital, but
it is smaller because the electron is more strongly attracted to the
nucleus.
The helium 1s is smaller as well.
This provides for better shielding.
PERIODIC TRENDS
IONIZATION ENERGY, ATOMIC RADIUS,
ELECTRON AFFINITY, ELECTRONEGATIVITY
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IONIZATION ENERGY
Defined as: The energy required to remove an electron form a
gaseous atom.
Highest energy electron removed first.
 First ionization energy (I1) is that required to remove the first
electron.
 Second ionization energy (I2) – the second electron
Trends in ionization energy
For Mg
• I1 = 735 kJ/mole
• I2 = 1445 kJ/mole
• I3 = 7730 kJ/mole
The effective nuclear charge increases as you remove electrons.
It takes much more energy to remove a core electron than a
valence electron because there is less shielding.
Ex. Explain this trend
For Al
• I1 = 580 kJ/mole
• I2 = 1815 kJ/mole
• I3 = 2740 kJ/mole
• I4 = 11,600 kJ/mole
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Across a Period
Generally from left to right, IE increases because there is a greater
nuclear charge with the same shielding.
Down a Group
As you go down a group IE decreases because electrons are
farther away.
It is not that simple
Zeff changes as you go across a period, so will IE
Half filled and filled orbitals are harder to remove electrons from.
Here’s what it looks like graphically.
Radial Probability Distance from nucleus
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ATOMIC RADIUS
Defined as: ½ the distance between nuclei of 2 identical atoms
Across a Period
Decreases due to electrons being added in the same energy
level and the number of protons increasing. Shielding is not
as effective and higher Zeff causes e to be pulled closer to the
nucleus resulting in a smaller atomic radius.
Down a Group
Increases. Electrons are added in higher energy levels farther
from the nucleus. Core electrons shield the nuclear charge so
a lower Zeff is not as effective at pulling the electrons, so the
atomic radius increases.
IONIC RADIUS
Measured relative to the parent atom.
Cations: always smaller than the parent atom since electrons
are lost. Higher p+ to e- ratio causes the remaining e to be
pulled closer.
Anions: always larger than the parent atom since electrons are
gained. Inner electrons shield the added e and the size of the
cloud increases.
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Isoelectronic species: atoms or ions with the same number of
electrons.
Ex. Compare the size of elements that are isoelectronic with
argon
ELECTRON AFFINITY
Defined as: the amount of energy needed to add an electron to
a gaseous atom (usually in kj/mole)
(+) EA – metals – hard to add an e-, energy is required,
endothermic
(-) EA – non-metals –easy to add an e-, energy is released,
exothermic
(0) EA – noble gases – no reason to test their affinity, as they
have no reason to gain an e.
Across a period: (+) to (-) becomes more favorable (except for
noble gases)
Down a group:
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ELECTRONEGATIVITY
Defined as:
Highest electronegativity:
Across a period:
Down a group:
PLACE THE FOLLOWING IN ORDER of INCREASING
AR, IE, EA, and EN
K Ca Cr Kr
AR:
IE:
EA:
EN:
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Cs Ag Si F
AR:
IE:
EA:
EN:
O S Se Te
AR:
IE:
EA:
EN:
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