Unit 5
Atomic Structure &
The Periodic Table
AP Chemistry
Chamberlain
1
Chapter 7
Atomic Structure
•
Scientists use experimental results to test scientific models.
One such model is that of the atom.
•
When experimental results are not consistent with the
predictions of a scientific model, the model must be revised or
replaced with a new model that is able to predict or explain
the _______________________________.
•
A robust scientific model is one that can be used to
explain/predict numerous results over a wide range of
experimental circumstances.
Coulomb’s Law
• The force between two charged particles is proportional to
the magnitude of each of the two charges (q1 and q2), and
inversely proportional to the square of the distance, r,
between them.
2
•
If the two charges are of _______________________, the
force between them is attractive; if they are the same sign,
the force is _________________________.
Light & the Quantum Theory
 Light is very peculiar in that it is a _____________ and a
_____________ at the same time.
The Wave Nature of Light
 Light moves at a constant speed through a particular medium
o c=
o c=

Wavelength (λ) is

Frequency (ν) of a wave is the

Relationship between wavelength & frequency:
3
Example 7.1 Calculate the frequency of an X ray that has a
wavelength of 8.21 nm.
Try it Out– Calculate the wavelength, in nanometers, of infrared
radiation that has a frequency of 9.76 x 1013 Hz.
The Particle Nature of Light
 Light comes as particles called _______________

The energy of a photon is proportional to frequency:
o h=
Example 7.3 Calculate the energy, in joules, of a photon of violet
light that has a frequency of 6.15 × 1014 s–1.
4
Example 7.4 A laser produces red light of wavelength 632.8 nm.
Calculate the energy, in kilojoules, of 1 mol of photons of this red
light.
Try it Out - Calculate the energy, in joules per photon, of
ultraviolet light with a wavelength of 235 nm.
The Photoelectric Effect
 Light shining on a metal surface may cause electrons to be
ejected from the surface.

The __________________ of the light needs to be above
threshold frequency to induce _____________
_______________.

Kinetic energy of electrons is _______________ to frequency
of incident radiation.

Kinetic energy of electrons is ___________________ of light
intensity.

This independence contradicts the wave nature of light.
o According to wave nature, energy of electrons should
be __________________ to the __________________.
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The Electromagnetic Spectrum

In order of increasing wavelength (decreasing frequency):
1.
2.
3.
4.
5.
6.

Visible light is the part of ultraviolet light that we can
______________.

To remember the colors in order of increasing frequency
remember _________________________.
Example 7.2 Which light has the higher frequency: the bright red
brake light of an automobile or the faint green light of a distant
traffic signal?
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Line Spectra of the Elements
 The light emitted by pure elements has specific
_________________.
o Therefore, light of only specific wavelengths can be
seen
o This emitted light is called a ____________
__________________

Line spectra tell us that atoms can only have certain
____________________ ____________________
o The atoms cannot have any arbitrary value of energy

The line emission spectrum of an element is a
______________________ for that element, and can be used
to _________________ the element.
Bohr’s Hydrogen Atom
 Bohr assumed….

Orbits are _________________. Electrons cannot exist
between defined orbits.

The equation:
7
Example 7.5 Calculate the energy of an electron in the second
energy level of a hydrogen atom.
Ground State and Excited State
 Ground State –

Excited State –

Electrons are promoted to higher levels through an electric
discharge, heat, or some other source of energy.

An atom in an excited state eventually emits a photon (or
several) as the electron drops back down to the ground state.
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Emission of Photons

Different types of ____________________________ lead to
absorption or emission of photons in different spectral
regions.

_____________________ radiation is associated with
transitions in molecular vibrations and so can be used to
detect the presence of different types of bonds.

_________________________________ radiation is
associated with transitions in electronic energy levels and so
can be used to probe electronic structure.
Bohr’s Equation:
Example 7.6 Calculate the energy change, in joules, that occurs
when an electron falls from the ni = 5 to the nf = 3 energy level in
a hydrogen atom.
Example 7.7 Calculate the frequency of the radiation released by
the transition of an electron in a hydrogen atom from the n = 5
level to the n = 3 level, the transition we looked at in Example 7.6.
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Try it Out - Calculate the energy change that occurs when an
electron is raised from the n=2 energy level to the n=4 energy
level of a hydrogen atom.
DeBroglie’s Equation:
Example 7.9 Calculate the wavelength, in meters and
nanometers, of an electron moving at a speed of 2.74 × 10 6 m/s.
The mass of an electron is 9.11 × 10–31 kg, and 1 J = 1 kg m2 s–2.
Try it Out - Calculate the wavelength, in nanometers, of a proton
moving at a speed of 3.79 x 103 m/s. The mass of a proton is 1.67
x 10-27 kg.
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Quantum Mechanical Model
• Thus far, we’ve discussed the electrons in orbits around the
atom (the Bohr model). That is a useful way of thinking of the
atom, but unfortunately it is __________________.
•
In the 1920s, __________________________ used
mathematics and probabilities to determine the location of
the electrons.
•
We can describe the electron ______________________ using
_____________________________.
•
Schrödinger developed a ___________________________ to
describe the hydrogen atom.
•
An acceptable solution to Schrödinger’s wave equation is
called a _______________________________.
•
A wave function represents an energy state of the atom.
•
Today, we refer to this model as the quantum mechanical
(QM) model. It addresses known problems with the classical
shell model and is also consistent with atomic electronic
structures that correspond to the periodic table.
•
The QM model can be approximately solved using
________________________and serves as the basis for
software that calculates the structure and reactivity of
molecules.
Heisenberg’s Uncertainty Principle:
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Quantum Numbers and Orbitals
The ______________________ quantum number (n):

Can only be a positive integer (n = ______________)

It is related to the distance away from the ________________

The size of an orbital and its electron energy depend on the
value of n.

Orbitals with the same value of n are said to be in the same
____________________________________.
The ______________/angular momentum quantum number (l):

Determines the ___________________of the orbital

Can have positive integral values from __________________

Orbitals having the same values of n and of l are said to be in
the same _________________________.
Value of l
Subshell

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Each orbital designation represents a different region of space
and a different ___________________.
The ____________________ quantum number (ml):

Determines the___________________ in space of the orbitals
of any given type in a subshell.

Can be any integer from __________________________

The number of possible values for ml is (____________), and
this determines the number of orbitals in a subshell.
The ____________________ quantum number

Values can be ________________

Clockwise or counterclockwise spin on the electron

The spin refers to a magnetic field induced by the moving
electric charge of the electron as it spins.

The magnetic fields of two electrons with opposite spins
cancel one another; there is no net magnetic field for the pair.
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Example 7.11
Consider the relationship among quantum numbers and orbitals,
subshells, and principal shells to answer the following. (a) How
many orbitals are there in the 4d subshell? (b) What is the first
principal shell in which f orbitals can be found? (c) Can an atom
have a 2d subshell? (d) Can a hydrogen atom have a 3p subshell?
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Chapter 8
The Periodic Table
 Pauli Exclusion Principle:
 In an atom, electrons can’t share same set of quantum
numbers


15

Two electrons can’t be in the same place at the same time

This may seem obvious, however two photons can be in
the same place at the same time
Aufbau Principle:
 Electrons in ground state atom are in _______________
possible energy

Electrons “fill” into orbitals from lower energy to _______
energy

This is why when we start an electron configuration we
always start at 1s (hydrogen)
Hund’s Rule:
 Electrons fill orbitals singly before doubling up.

Electrons in singly occupied orbitals have the
____________ spin.

Electrons in doubly occupied orbitals have
_______________ spin.
Orbital Diagrams & Electron Configuration
 An ______________________________________ describes
the distribution of electrons among the various orbitals in the
atom

It can be represented in 2 ways:
o spdf notation

Magnesium

Molybdenum

Cobalt
o orbital diagram

Nitrogen
Example 8.1 Write electron configurations for sulfur, using both
the spdf notation and an orbital diagram.
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Try it Out – Write electron configurations for calcium using both
the spdf notation and an orbital diagram.

Noble gas-core abbreviation:
o Chromium
o Uranium
o Iodine
o Zinc (orbital notation)
Main Group and Transition Elements
17
Example 8.2 Give the complete ground-state electron
configuration of a strontium atom (a) in the spdf notation and (b)
in the noble-gas-core abbreviated notation.

Exceptions to the Aufbau Principle:
o _______________ d subshell plus ________________ s
subshell has slightly lower energy than s2d4.
o Filled ______ subshell plus half-filled _______ subshell
has slightly lower energy than s2d9.
o Silver
Valence Electrons and Core Electrons
 The ________________________ is the outermost occupied
principal shell. The valence shell contains the valence
electrons.

18
For main group elements, the number of valence shell
electrons is the same as…

Electrons in inner shells are called ______________
__________________
Electron Configuration of Ions
 To obtain the electron configuration of an _________ by the
aufbau process, we simply __________ the additional
electrons to the valence shell of the neutral nonmetal atom.

The number added usually completes the _________

A nonmetal monatomic ion usually attains the electron
configuration of a _____________________.
o O-2
o Br-1

A metal atom loses electrons to form ____________.

Electrons are ________________ from the configuration of
the atom.

The first electrons lost are those of the ___________
____________ quantum number.

If there are two subshells with the highest principal quantum
number, electrons are lost from the subshell with the higher l.
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Example 8.3 Write the electron configuration of the Co3+ ion in a
noble-gas-core abbreviated spdf notation.

Try it Out:
o
Al+3
o Ni+2
o S-2
Isoelectronic:

What are three ions that are isoelectronic with Xe?
Magnetic Properties
 Diamagnetism –

Paramagnetism –

Ferromagnetism –
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
Paramagnetic =
o The more unpaired electrons, the more attracted the
compound will be

Diamagnetic =

Examples:
o Zinc
o Manganese
Example 8.4 A sample of chlorine gas is found to be
diamagnetic. Can this gaseous sample be composed of individual
Cl atoms?
Spectrophotometry

21
A spectrophotometer measures the ____________________
or transmittance of light in a sample.

Beer’s Law –

The formula:
Periodic Properties

Certain physical and chemical properties recur at regular
intervals, and/or vary in regular fashion, when the elements
are arranged according to increasing atomic number.

Melting point, boiling point, hardness, density, physical state,
and chemical reactivity are periodic properties.

Periodicity is a useful tool when designing new molecules or
materials, since replacing an element of one group with
another of the same group may lead to a new substance with
similar properties.

For instance,

Atomic Radius
o How it’s calculated –
o Atomic radius _________________ as you go left to
right across a period.

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Why?
o Atomic radius _________________ as you go down a
group.
 Why?
Example 8.5 With reference only to a periodic table, arrange
each set of elements in order of increasing atomic radius:
(a) Mg, S, Si
(b) As, N, P
(c) As, Sb, Se

Ionic Radii
o How it’s calculated –
o Cations are ____________ than the atoms from which
they are formed.
 Why?
o Anions are ____________ than the atoms from which
they are formed.
 Why?
o Isoelectronic species have the same electron
configuration; size decreases with effective nuclear
charge
23
Example 8.6 Refer to a periodic table but not to Figure 8.14, and
arrange the following species in the expected order of increasing
radius:
Ca2+, Fe3+, K+, S2–, Se2–

Ionization energy –
o Expressed in _______________
o The first ionization energy is the minimum energy
needed to remove the least tightly held electron from
an atom or ion.
o The relative magnitude of the ionization energy can be
estimated through qualitative application of Coulomb’s
law.
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o The farther an electron is from the nucleus, the lower
its ionization energy.
o When comparing two species with the same
arrangement of electrons, the higher the
____________ charge, the higher the ionization energy
of an electron in a given subshell.
o I1<I2<I3
 Removing an electron from a ______________
ion is more difficult than removing it from a
_________________________.

A large jump in I occurs after valence electrons
are completely removed. Why?

How many valence electrons does this element
have if these are its first 4 ionization energies?
[1st IE = 578 kJ/mol]
[2nd IE = 1820 kJ/mol]
[3rd IE = 2750 kJ/mol]
[4th IE = 11,600 kJ/mol]
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AP Exam Question
Ionization Energies for Element X (kJ/mole)
First
Second
Third
Fourth
Fifth
580
1815
2760
11600
14800
The ionizations for element X are listed in the table. On the basis
of the data, element X is most likely to be:
a) Na
b) Mg
c) Al
d) Si
e) P
o Ionization energy _________________ as you go from
left to right across a period.
 Why?
o Ionization energy _________________ as you go down
a group.
 Why?
 Shielding Effect - Caused by an
increasing number of electrons between
the outer level and the nucleus
Example 8.7 Without reference to Figure 8.15, arrange each set of
elements in the expected order of increasing first ionization
energy.
(a) Mg, S, Si
(b) As, N, P
(c) As, Ge, P
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
Electronegativity –
o Electronegativity _________________ as you go from
left to right across a period.
 Why?
o Electronegativity _________________ as you go down
a group.
 Why?
Metals
 Metals have a small number of electrons in their valence shells
and tend to form _________________ ions.
Summary of
Trends:
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