Chapter 7
Atomic Structure
and Periodicity
What is New for 2015 AP Chemistry Exam
1-Students do not need to memorize all the exceptions to the Aufbau principle.
2-Students do not need to know how to assign quantum numbers to electrons.
3-Students will not be tested on phase diagrams.
4-Colligative properties will not be tested on the exam.
5-Students are not required to solve problems involving molality, percent by mass, and percent by volume.
6-Problems involving hydrogen bonding are limited to interactions between molecules that include hydrogen atoms covalently
bonded to nitrogen, oxygen or fluorine. Students are not required to know other cases of much weaker hydrogen bonding.
7-Students do not need to memorize the different types of crystal structures.
8-Students will not be tested on the use of formal charge to explain why certain molecules do not obey the octet rule.
9-Students do not need to learn how to defend Lewis models based on the assumptions about the limitations of the models.
10-Academics are still debating whether to teach hybridization but until this issue is cleared, AP students are required to know the
difference between sp, sp2, and sp3 hybridizations and the bond angles. However, students do not need to know how these
orbitals are derived or depicted.
11-Aspects such as recall or filling of the molecular orbital diagrams related to the molecular orbital theory will not be tested.
12-Students will not be tested on the specific varieties of crystal lattices for ionic compounds.
13-Students will not be tested on Lewis acid-base concepts.
14-The terms “reducing agent” and “oxidizing agent” will not be used in this course or on the test.
15-Students do not need to label electrodes as positive or negative in electrochemistry cells.
16-Students will not be tested about the Nernst equation.
17-Students are not required to solve problems involving Arrhenius equation.
18-Collection of data involving reaction intermediates will not be covered on the test.
19-Students are not required to calculate the concentrations of each species present in the titration curve for polyprotic acids.
20-Students are not required to calculate the change in pH resulting from the addition of an acid or a base to a buffer.
21-Students do not need to know how to derive the Henderson-Hasselbalch equation.
22-Students do not need to memorize all the solubility rules. Students will only be assessed on the solubility of sodium, potassium,
ammonium, and nitrate salts.
23-Students will not be required to calculate solubility or solubility as a function of pH.
Section 5.6
The Kinetic Molecular Theory of Gases
Section 6.1
The Nature of Energy
Section 6.1
The Nature of Energy
Section 6.1
The Nature of Energy
Light
matter
Section 6.1
The Nature of Energy
Not on
AP Exam
Practice
Section 6.1
The Nature of Energy
Section 6.1
The Nature of Energy
for degenerate orbitals, the lowest energy is attained when
the number of electrons with the same spin is maximized.
Section 6.1
The Nature of Energy
& Coulomb’s Law
& Electronegativity
PES
Absorbance
Review Questions
Section 6.1
The Nature of Energy
Section 6.1
The Nature of Energy
https://www.youtube.com/watch?v=yteVW9CVuE4
Section 6.1
The Nature of Energy
https://www.youtube.com/watch?v=5V-UZ-t5AJ4
Section 6.1
The Nature of Energy
https://www.youtube.com/watch?v=NCZkCEaknxE
Electron Configuration
Section 6.1
The Nature of Energy
https://www.youtube.com/watch?v=9WKzgVf8Iyg
Section 6.1 Periodic Table Trends Questions
The Nature of Energy
https://www.youtube.com/watch?v=ELGHGflwYtw
Section 6.1
The Nature of Energy
https://www.youtube.com/watch?v=K-3jRZdscJA
Periodic Table Elements
Section 6.1
The Nature of Energy
https://www.youtube.com/watch?v=f4BWgxhrCt8
Section 7.1
Electromagnetic Radiation
Different Colored
Fireworks
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Section 7.1
Electromagnetic Radiation
Questions to Consider
 Why do we get colors?
 Why do different chemicals give us different colors?
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Section 7.1
Electromagnetic Radiation
Electromagnetic Radiation
 One of the ways that energy travels through space.
 Three characteristics:
 Wavelength
 Frequency
 Speed
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Section 7.1
Electromagnetic Radiation
Characteristics
 Wavelength () – distance between two consecutive
peaks or troughs in a wave.
 Frequency ( ) – number of waves (cycles) per
second that pass a given point in space
 Speed (c) – speed of light (2.9979×108 m/s)
c = 
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Section 7.1
Electromagnetic Radiation
The Nature
of Waves
23
Section 7.1
Electromagnetic Radiation
Classification of Electromagnetic Radiation
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Section 7.2
The Nature of Matter
Pickle Light
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Section 7.2
The Nature of Matter
 Energy can be gained or lost only in whole number
multiples of hν .
 A system can transfer energy only in whole quanta (or
“packets”).
 Energy seems to have particulate properties too.
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Section 7.2
The Nature of Matter
 Energy is quantized.
 Electromagnetic radiation is a stream of “particles”
called photons.
hc
Ephoton = hν =

 Planck’s constant = h = 6.626 × 10-34 Js
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Section 7.2
The Nature of Matter
The Photoelectric effect
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Section 7.2
The Nature of Matter
 Energy has mass E = mc2
 Dual nature of light:
 Electromagnetic radiation (and all matter) exhibits
wave properties and particulate properties.
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Section 7.3
The Atomic Spectrum of Hydrogen
 Continuous spectrum (results when white light is passed
through a prism) – contains all the wavelengths of visible
light
 Line spectrum – each line corresponds to a discrete
wavelength:
 Hydrogen emission spectrum
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Section 7.3
The Atomic Spectrum of Hydrogen
Refraction of White Light
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Section 7.3
The Atomic Spectrum of Hydrogen
The Line Spectrum of Hydrogen
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Section 7.3
The Atomic Spectrum of Hydrogen
Significance
 Only certain energies are allowed for the electron in the
hydrogen atom.
 Energy of the electron in the hydrogen atom is
quantized.
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Section 7.3
The Atomic Spectrum of Hydrogen
CONCEPT CHECK!
Why is it significant that the color emitted from
the hydrogen emission spectrum is not white?
How does the emission spectrum support the
idea of quantized energy levels?
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Section 7.4
The Bohr Model
 Electron in a hydrogen atom moves around the nucleus
only in certain allowed circular orbits.
 Bohr’s model gave hydrogen atom energy levels
consistent with the hydrogen emission spectrum.
 Ground state – lowest possible energy state (n = 1)
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Section 7.4
The Bohr Model
Electronic Transitions in
the Bohr Model for the
Hydrogen Atom
a) An Energy-Level Diagram for
Electronic Transitions
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Section 7.4
The Bohr Model
Electronic Transitions in
the Bohr Model for the
Hydrogen Atom
b) An Orbit-Transition Diagram,
Which Accounts for the
Experimental Spectrum
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Section 7.4
The Bohr Model
 For a single electron transition from one energy level to
another:
 1
1 
18
E =  2.178  10 J  2  2 
ninitial 
 nfinal
ΔE = change in energy of the atom (energy of the emitted photon)
nfinal = integer; final distance from the nucleus
ninitial = integer; initial distance from the nucleus
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Section 7.4
The Bohr Model
 The model correctly fits the quantized energy levels of
the hydrogen atom and postulates only certain allowed
circular orbits for the electron.
 As the electron becomes more tightly bound, its energy
becomes more negative relative to the zero-energy
reference state (free electron). As the electron is brought
closer to the nucleus, energy is released from the system.
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Section 7.4
The Bohr Model
 Bohr’s model is incorrect. This model only works for
hydrogen.
 Electrons move around the nucleus in circular orbits.
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Section 7.4
The Bohr Model
EXERCISE!
What color of light is emitted when an
excited electron in the hydrogen atom falls
from:
a) n = 5 to n = 2
b) n = 4 to n = 2
c) n = 3 to n = 2
blue, λ = 434 nm
green, λ = 486 nm
orange/red, λ = 657 nm
Which transition results in the longest
wavelength of light?
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Section 7.5
The Quantum Mechanical Model of the Atom
 We do not know the detailed pathway of an electron.
 Heisenberg uncertainty principle:
 There is a fundamental limitation to just how precisely
we can know both the position and momentum of a
particle at a given time.
x    m 
h

4
Δx = uncertainty in a particle’s position
Δ(mν) = uncertainty in a particle’s momentum
h = Planck’s constant
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Section 7.5
The Quantum Mechanical Model of the Atom
Physical Meaning of a Wave Function (Ψ)
 The square of the function indicates the probability of
finding an electron near a particular point in space.
 Probability distribution – intensity of color is used to
indicate the probability value near a given point in
space.
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Section 7.5
The Quantum Mechanical Model of the Atom
Probability Distribution for the
1s Wave Function
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Section 7.5
The Quantum Mechanical Model of the Atom
Radial Probability Distribution
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Section 7.5
The Quantum Mechanical Model of the Atom
Relative Orbital Size
 Difficult to define precisely.
 Orbital is a wave function.
 Picture an orbital as a three-dimensional electron density
map.
 Hydrogen 1s orbital:
 Radius of the sphere that encloses 90% of the total
electron probability.
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Section 7.6
Quantum Numbers
 Principal quantum number (n) – size and energy of the
orbital.
 Angular momentum quantum number (l) – shape of
atomic orbitals (sometimes called a subshell).
 Magnetic quantum number (ml) – orientation of the
orbital in space relative to the other orbitals in the atom.
47
Section 7.6
Quantum Numbers
Quantum Numbers for the First Four Levels of Orbitals in
the Hydrogen Atom
Section 7.6
Quantum Numbers
EXERCISE!
For principal quantum level n = 3, determine the
number of allowed subshells (different values of
l), and give the designation of each.
# of allowed subshells = 3
l = 0, 3s
l = 1, 3p
l = 2, 3d
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Section 7.6
Quantum Numbers
EXERCISE!
For l = 2, determine the magnetic quantum numbers
(ml) and the number of orbitals.
magnetic quantum numbers = –2, – 1, 0, 1, 2
number of orbitals = 5
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Section 7.7
Orbital Shapes and Energies
1s Orbital
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Section 7.7
Orbital Shapes and Energies
Three Representations
of the Hydrogen 1s, 2s,
and 3s Orbitals
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Section 7.7
Orbital Shapes and Energies
2px Orbital
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Section 7.7
Orbital Shapes and Energies
2py Orbital
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Section 7.7
Orbital Shapes and Energies
2pz Orbital
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Section 7.7
Orbital Shapes and Energies
The Boundary Surface Representations of All Three 2p Orbitals
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Section 7.7
Orbital Shapes and Energies
3dx -y Orbital
2
2
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Section 7.7
Orbital Shapes and Energies
3dxy Orbital
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Section 7.7
Orbital Shapes and Energies
3dxz Orbital
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Section 7.7
Orbital Shapes and Energies
3dyz Orbital
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Section 7.7
Orbital Shapes and Energies
3dz2
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Section 7.7
Orbital Shapes and Energies
The Boundary Surfaces of All of the 3d Orbitals
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Section 7.7
Orbital Shapes and Energies
Representation of the 4f Orbitals in Terms of Their Boundary Surfaces
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Section 7.8
Electron Spin and the Pauli Principle
Electron Spin
 Electron spin quantum number (ms) – can be +½ or -½.
 Pauli exclusion principle - in a given atom no two
electrons can have the same set of four quantum
numbers.
 An orbital can hold only two electrons, and they must
have opposite spins.
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Section 7.9
Polyelectronic Atoms
 Atoms with more than one electron.
 Electron correlation problem:
 Since the electron pathways are unknown, the
electron repulsions cannot be calculated exactly.
 When electrons are placed in a particular quantum level,
they “prefer” the orbitals in the order s, p, d, and then f.
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Section 7.9
Polyelectronic Atoms
Penetration Effect
 A 2s electron penetrates to the nucleus more than one in
the 2p orbital.
 This causes an electron in a 2s orbital to be attracted to
the nucleus more strongly than an electron in a 2p
orbital.
 Thus, the 2s orbital is lower in energy than the 2p orbitals
in a polyelectronic atom.
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Section 7.9
Polyelectronic Atoms
Orbital Energies
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Section 7.9
Polyelectronic Atoms
A Comparison of the Radial Probability Distributions of the 2s and 2p
Orbitals
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Section 7.9
Polyelectronic Atoms
The Radial Probability Distribution of the 3s Orbital
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Section 7.9
Polyelectronic Atoms
A Comparison of the Radial Probability Distributions of the 3s, 3p, and
3d Orbitals
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Section 7.10
The History of the Periodic Table
 Originally constructed to represent the patterns
observed in the chemical properties of the elements.
 Mendeleev is given the most credit for the current
version of the periodic table because he emphasized how
useful the periodic table could be in predicting the
existence and properties of still unknown elements.
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71
Section 7.11
The Aufbau Principle and the Periodic Table
Aufbau Principle
 As protons are added one by one to the nucleus to build
up the elements, electrons are similarly added to
hydrogen-like orbitals.
 An oxygen atom has an electron arrangement of two
electrons in the 1s subshell, two electrons in the 2s
subshell, and four electrons in the 2p subshell.
Oxygen: 1s22s22p4
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Section 7.11
The Aufbau Principle and the Periodic Table
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 (same energy) orbitals.
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Section 7.11
The Aufbau Principle and the Periodic Table
Orbital Diagram
 A notation that shows how many electrons an atom has
in each of its occupied electron orbitals.
Oxygen: 1s22s22p4
Oxygen:
1s
2s
2p
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Section 7.11
The Aufbau Principle and the Periodic Table
Valence Electrons
 The electrons in the outermost principal quantum level
of an atom.
1s22s22p6 (valence electrons = 8)
 The elements in the same group on the periodic table
have the same valence electron configuration.
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Section 7.11
The Aufbau Principle and the Periodic Table
The Orbitals Being Filled for Elements in Various Parts of the Periodic
Table
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Section 7.11
The Aufbau Principle and the Periodic Table
EXERCISE!
Determine the expected electron configurations for
each of the following.
a) S
1s22s22p63s23p4 or [Ne]3s23p4
b) Ba
[Xe]6s2
c) Eu
[Xe]6s24f7
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Section 7.12
Periodic Trends in Atomic Properties
Periodic Trends
 Ionization Energy
 Electron Affinity
 Atomic Radius
Section 7.12
Periodic Trends in Atomic Properties
Ionization Energy
 Energy required to remove an electron from a gaseous atom or
ion.
 X(g) → X+(g) + e–
Mg → Mg+ + e–
Mg+ → Mg2+ + e–
Mg2+ → Mg3+ + e–
I1 = 735 kJ/mol(1st IE)
I2 = 1445 kJ/mol
(2nd IE)
I3 = 7730 kJ/mol
*(3rd IE)
*Core electrons are bound much more tightly than valence
electrons.
Section 7.12
Periodic Trends in Atomic Properties
Ionization Energy
 In general, as we go across a period from left to right, the
first ionization energy increases.
 Why?
 Electrons added in the same principal quantum level
do not completely shield the increasing nuclear charge
caused by the added protons.
 Electrons in the same principal quantum level are
generally more strongly bound from left to right on
the periodic table.
Section 7.12
Periodic Trends in Atomic Properties
Ionization Energy
 In general, as we go down a group from top to bottom,
the first ionization energy decreases.
 Why?
 The electrons being removed are, on average, farther
from the nucleus.
Section 7.12
Periodic Trends in Atomic Properties
The Values of First Ionization Energy for the Elements in the First Six
Periods
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Explain why the graph of ionization energy versus
atomic number (across a row) is not linear.
electron repulsions
Where are the exceptions?
some include from Be to B and N to O
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Which atom would require more energy to remove
an electron? Why?
Na
Cl
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Which atom would require more energy to remove
an electron? Why?
Li
Cs
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Which has the larger second ionization energy? Why?
Lithium or Beryllium
Section 7.12
Periodic Trends in Atomic Properties
Successive Ionization Energies (KJ per Mole) for the Elements in
Period 3
Section 7.12
Periodic Trends in Atomic Properties
Electron Affinity
 Energy change associated with the addition of an
electron to a gaseous atom.
 X(g) + e– → X–(g)
 In general as we go across a period from left to right, the
electron affinities become more negative.
 In general electron affinity becomes more positive in
going down a group.
Section 7.12
Periodic Trends in Atomic Properties
Atomic Radius
 In general as we go across a period from left to right, the
atomic radius decreases.
 Effective nuclear charge increases, therefore the
valence electrons are drawn closer to the nucleus,
decreasing the size of the atom.
 In general atomic radius increases in going down a group.
 Orbital sizes increase in successive principal quantum
levels.
Section 7.12
Periodic Trends in Atomic Properties
Atomic Radii for
Selected Atoms
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Which should be the larger atom? Why?
Na
Cl
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Which should be the larger atom? Why?
Li
Cs
Section 7.12
Periodic Trends in Atomic Properties
CONCEPT CHECK!
Which is larger?
 The hydrogen 1s orbital
 The lithium 1s orbital
Which is lower in energy?
The hydrogen 1s orbital
The lithium 1s orbital
Section 7.12
Periodic Trends in Atomic Properties
Atomic Radius of a Metal
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Section 7.12
Periodic Trends in Atomic Properties
Atomic Radius of a Nonmetal
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Section 7.12
Periodic Trends in Atomic Properties
EXERCISE!
Arrange the elements oxygen, fluorine, and sulfur
according to increasing:
 Ionization energy
S, O, F
 Atomic size
F, O, S
Section 7.13
The Properties of a Group: The Alkali Metals
The Periodic Table – Final Thoughts
1. It is the number and type of valence electrons that
primarily determine an atom’s chemistry.
2. Electron configurations can be determined from the
organization of the periodic table.
3. Certain groups in the periodic table have special names.
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Section 7.13
The Properties of a Group: The Alkali Metals
Special Names for Groups in the Periodic Table
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Section 7.13
The Properties of a Group: The Alkali Metals
The Periodic Table – Final Thoughts
4. Basic division of the elements in the periodic table is
into metals and nonmetals.
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Section 7.13
The Properties of a Group: The Alkali Metals
Metals Versus Nonmetals
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Section 7.13
The Properties of a Group: The Alkali Metals
The Alkali Metals
 Li, Na, K, Rb, Cs, and Fr
 Most chemically reactive of the metals
 React with nonmetals to form ionic solids
 Going down group:




Ionization energy decreases
Atomic radius increases
Density increases
Melting and boiling points smoothly decrease
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101
Photoelectron
Spectroscopy
Straightforward evidence for
the validity of orbital diagram
Konsler 2013
2014 AP Exam PES
2014 AP Exam PES
Electron Configuration
Is there any direct
evidence that this
diagram is
accurately showing
potential energy of
electrons on the
atom?
Some Available Evidence
• Atomic Emission Spectra
• Successive Ionization
• Photoelectron Spectroscopy
The topic is probably best introduced in1st year
Chemistry under Atomic Electronic Structure
Atomic Emission Spectra
Atomic emission spectra are evidence that there are
specific allowed PE values by showing the energy
difference between some of them.
Limitations
Differences between PE values, not the values themselves.
Ambiguous origin and destination orbitals.
Impossible to visually compare PE of orbitals on multiple
atoms; large data sets will require many calculations.
Successive Ionization
Using successive ionization it
is usually possible to
determine the number of
valence electrons an atom
has. This is evidence that our
PE diagram is an accurate
representation
Limitations
Each ionization causes a
reorganization of the remaining
electrons, meaning the successive
ionization is not a measurement of
the characteristics of the original
atom. We infer there is a
relationship between the ions and
the parent atom.
Limitations
It’s not feasible to do more
than 5 successive ionizations,
making core electrons (and
sometimes even valence
electrons) impossible to obtain
values for.
Limitations
It’s not possible to determine if
electrons at the same energy
level on the atom have the
same PE initially, because we
remove them one at a time. As
a result, the method is really a
support only for n, not the
other 3 quantum numbers.
Photoelectron Spectroscopy
Ephoton = hv
Atom
Monochromatic
Beam of X-Rays
IEelectron = Ephoton - KE
Each event happens once for a
single atom. This is a quantum
event. Since Ephoton > IEelectron for
all the electrons on the atom, the
electron removed is random and
the KE of the electron is a
characteristic of that electron as
it exists on the atom at the
moment of the event.
KE = mv2
2
e-
Negatively Charged Hemisphere (Constant V)
Retarding Voltage is directly
proportional to KEelectron
Independent
Variable (x)
Retarding
Voltage of Lens
Negatively
Charged
Hemisphere
(Constant V)
Dependent
Variable (y)
Intensity
Beam of Atoms
Electron Detector
Electrons scan past detector
Electrostatic Lens
(Focuses and slows electrons)
The intensity is measured as the retarding voltage on
the lens is reduced at a constant rate
X-Rays
(monochromatic)
Electrons ejected
Photoelectron Spectroscopy is still limited by Ephoton. If IEelectron is
too large, the electron will not be detected. This means that
electrons with low values of n on atoms with high value of Z may
be off the left hand side of the chart.
High Ionization Energy
Low electron PE
20 MJ/mol
Low Ionization Energy
High electron PE
10 MJ/mol
IEelectron = Ephoton - KEelectron
0 MJ/mol
However, a fluorescence source emits X-Rays with energy of over
1000 eV, several times greater than for successive ionization
instrumentation. In addition, core electrons are removed from a
neutral atom, requiring less energy.
High Ionization Energy
Low electron PE
20 MJ/mol
Low Ionization Energy
High electron PE
10 MJ/mol
IEelectron = Ephoton - KEelectron
0 MJ/mol
Photoelectron Spectrum
Valence
Helium possesses the valence electron with the
lowest PE of all elements. The measured value for
He is 2.37 MJ/mol. Any values observed with greater
values (to the left of this line) must be core electrons.
Note, d-subshell core electrons will be to the right of
this line.
20 MJ/mol
10 MJ/mol
0 MJ/mol
“Valence Line”: He(1s) = 2.37 MJ/mol
Photoelectron Spectrum
Relative Intensity = 2
Since each event removes a random electron from
a separate atom, the relative intensity shows the
proportion of electrons at that PE.
Relative Intensity = 1
20 MJ/mol
10 MJ/mol
0 MJ/mol
Photoelectron Spectrum
Valence
19.3 MJ/mol
1.36 MJ/mol
RI = 2
RI = 2
0.80 MJ/mol
RI = 1
20 MJ/mol
10 MJ/mol
0 MJ/mol
Photoelectron Spectrum
Boron (Z=5)
19.3 MJ/mol
RI = 2
Analysis:
1) Valence has 2 values:
2) RI is 2 to 1 in valence:
3) Closest core has RI 2 not 6:
4) s2s2p1 must be 1s22s22p1
sp
s2p1
not pxs2p1
10 MJ/mol
RI = 2
0.80 MJ/mol
RI = 1
2s2
1s2
20 MJ/mol
1.36 MJ/mol
2p1
0 MJ/mol
Photoelectron Spectrum
2p6
3.67 MJ/mol
1s2
2s2
104 MJ/mol
Sodium (Z=11)
3s1
6.84 MJ/mol
0.50 MJ/mol
{}
8 MJ/mol
4 MJ/mol
0 MJ/mol
But there aren’t even any instruments!
•
•
•
All PES spectra will probably be “simulated” for the following reasons:
PES post-dates atomic theory.
Most PES is a solid-phase technique. Having so many nearby atoms will
cause interference. For example, an electron promoted out of a low lying
orbital will leave a vacancy which can be occupied by an outer orbital electron
with concomitant production of a photon. If that photon is of high enough
energy, it might eject a more easily ionized electron in the same atom or a
nearby atom. This will result in the production of background interference and
even spurious peaks. As a result, PES spectra derived from solids would be
impossible for a non-expert to interpret.
Why do it?
•
•
•
•
Although it requires “simulated spectra” it is superior to other methods we describe
in General Chemistry because, even in principle, atomic emission and successive
ionization do not provide data which could be used to identify an atom in the context
of an exam. At best they would allow a student to choose the most likely among a
set of possibilities.
Photoelectron spectroscopy allows questions to be written using visual as opposed
to tabular data (as is required for successive ionization).
Students can be asked to interpret a spectrum but they can also legitimately be
asked to draw one. It’s a decently easy grade if you give the axes and requires
essentially no “give away” information in the stimulus.
Therefore, while the instruments aren’t accessible, understanding the principle
allows more thoughtful questions to be asked about electronic structure than were
previously possible.
Photoelectron Spectroscopy
(PES)
Spectroscopy
• Method of analyzing matter using
electromagnetic radiation.
Photoelectron Spectroscopy
• PES apparatus:
iramis.cea.fr
Photoelectron Spectroscopy
How it works:
1. Sample is
exposed to
EM radiation
2. Electrons
jump out of
sample and
go through
analyzer
http://chemwiki.ucdavis.edu
PES Data
Electrons
generally closer
to the nucleus
Electrons
generally
farther from
the nucleus
Each peak represents the
electrons in a single
sublevel in the atom
The bigger the peak – the
more electrons
Number of electrons
Energy to remove an electron
(binding energy)
(increases to the left!)

Hydrogen vs. Helium
Helium
Hydrogen
#e-
#e-
 energy
1 electron in 1s
 energy
2 electrons in 1s
The helium peak is twice as tall because there are twice as many electrons in the 1s sublevel
Hydrogen vs. Helium
Helium
Hydrogen
#e-
#e-
 energy
1 electron in 1s
 energy
2 electrons in 1s
The helium peak is farther to the left (higher energy) thus more energy is needed to remove
the 1s electrons in helium. They must be held more tightly because there is a higher effective
nuclear charge. (Helium has 2 protons pulling on 1s but hydrogen only has 1)
Oxygen (1s22s22p4)
Number of electrons
2 electrons in 1s
2 electrons in 2s
Energy to remove an electron
(binding energy)
(increases to the left!)

4 electrons in 2p
Scandium (1s22s22p63s23p64s23d1)
Number of electrons
*Notice that it takes more
energy to remove an electron from
3d than from 4s.
This is because as electrons are
added to 3d they shield 4s thus it’s
2 in 3s
easier (takes less energy) to remove 2 in1s
4s electrons compared to 3d
2 in 2s 6 in 2p
electrons.
Remember when transition metals make
positive ions - it’s the s electrons that are
lost first!
2 in 4s
6 in 2p
1 in 3d
Energy to remove an electron
(binding energy)
(increases to the left!)

Example 1:
Number of electrons
Identify the element whose
PES data is shown to the right.
Sodium
Why is one peak much larger
Than the others?
This peak represents 6 electrons
In the 2p sublevel the other
Peaks represent only 1 or 2
electrons
In which sublevel are the electrons
Represented by peak A
3s
A
Energy

Example 2:
Oxygen
Nitrogen
#e-
#e-
 energy
 energy
The PES data above shows only the peak for the 1s electrons. Why is the peak for
Oxygen farther to the left?
It takes less energy to remove a 1s electron from nitrogen because it has a lower
Effective nuclear charge (less protons) than oxygen
Example 3:
Number of electrons
Draw the expected PES
Spectrum for the element boron
Energy

Disregard This
Hydrogen emission
spectrum
Line
Spectra
Helium was first detected in the Sun due to its
characteristic spectral lines
Greek word for the Sun, ήλιος (ílios or helios)
. Emit color light
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Spectrum
Excited Staten=4
UV
Excited Staten=3
Excited State unstable
and drops back down
Excited State
n=2
But only as far as
n = 2 this time
•Energy released as a photon
•Frequency proportional
to energy drop
V
i
s
i
b
l
e
IR
n=1
Ground State
. Emit color light
Continuous Spectrum
Emission Line Spectrum
. Emit color light
DNA
Hydrogen emission
spectrum
Electron Orbitals
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Heisenberg Uncertainty Principle
Definition: The scientific principle stating
that it is impossible to determine with
perfect accuracy both the position and
momentum of a particle at any given point in
time.
Azimuthal
Orbital
Shapes & Energies
n = number – shell
- energy level
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Orbital Energies
Electron
Configuration
Practice writing Electron Configuration
Trends
Periodic Table Trends Questions
Effective Nuclear charge
Electron Shielding
Energy/Attractive/Repulsive
forces
https://www.youtube.com/watch?v=K-3jRZdscJA
Effective Nuclear charge
Electron Shielding
Energy/Attractive/Repulsive
forces
The Values of First Ionization Energy for the Elements in the First Six
Periods
Practice Activities
Not on New AP Exam
Not on New AP Exam
Absorba
nce