Ch. 9: Electrons in Atoms and
the Periodic Table
Dr. Namphol Sinkaset
Chem 152: Introduction to General
Chemistry
I. Chapter Outline
I.
II.
III.
IV.
V.
VI.
Introduction
Electromagnetic Radiation
The Bohr Model of the Atom
The Quantum-Mechanical Atom
Electrons and the Periodic Table
Periodic Trends
I. Hydrogen vs. Helium
I. Electrons and Chemistry
• Chemistry is all about electrons.
• Therefore, how electrons are organized
in the atom is an important concept.
• We will see that reactivity and the
arrangement of the periodic table are
both related to electrons.
I. Stranger Than Anyone
Thought
• Theories used to explain how electrons
are organized in the atom were devised
by scientists like Bohr, Schrödinger,
Planck, and Einstein.
• “I don’t like it, and I am sorry I ever had
anything to do with it.”
• “God does not play dice with the
universe.”
II. It Started With Light
• Light can interact with atoms, and
studying the interactions led to
understanding how electrons are
organized in the atom.
• Light is not a form of matter; it is
electromagnetic radiation.
• EM radiation is a type of energy that
travels at 3.0 x 108 m/s.
II. Classical Wave
• Light was
considered a
classical wave
phenomenon – like
ripples in water or a
moving rope.
• But that wasn’t quite
right…
II. Describing Waves
• EM waves are characterized by their
wavelength (λ) and their frequency (ν).
II. White Light
• White light can
be separated.
• Note that short
wavelength =
high frequency
and vice versa.
II. Light As Particles
• In the early 20th century, scientists like
Einstein saw that light was not a
classical wave.
• They explained that light acted like
particles, which were called photons.
• Thus, when we are under a light, we are
being showered with light particles.
II. The EM Spectrum
III. Atoms Can Emit Light
III. Atomic Emission Spectra
III. Atomic LINE Spectra
• The individual lines were key to
formulation of Bohr’s atomic model.
• The movement of electrons were the
reason for atoms emitting light.
• But why lines? Lines meant only
specific wavelengths (colors) were
allowed.
III. The Bohr Model
• Bohr reasoned that
electrons were only
allowed to have
certain energies.
• Strange – it’s like
someone telling you
were only allowed to
have certain
amounts of money.
III. Bohr Energies
• The electron orbits
in the Bohr model
are like rungs on a
ladder.
• You can stand on
one step or another,
but never in
between.
III. Bohr Orbits
• Each orbit in the Bohr model has a
specific energy that is specified by a
quantum number, n.
• When an electron moves to a higher
orbit, it must absorb a quantum of
energy.
• When an electron moves to a lower
orbit, it must emit a quantum of energy.
III. Moving Between Orbits
III. Line Spectra Correspond
to Electron Transitions
III. Works Great for Hydrogen
• Originally, Bohr set out to model only
the hydrogen atom.
• When people tried to extend it atoms
with more than one electron, it didn’t
work!
• A new model that worked for all atoms
was needed.
IV. Wavy Electrons
• Experiments found that electrons don’t
always act like particles – they
sometimes act like waves!
• Electrons phase in and out (similar to
how waves oscillate), so we don’t know
exactly where they are.
• The best we can do is plot probability
maps.
IV. Baseballs vs. Electrons
IV. Orbits to Orbitals
• In the Bohr model, electrons were in
well-defined orbits like planets around
the sun.
• In the new quantum-mechanical model,
orbits are replaced by orbitals, which
are probability maps of where an
electron could be found.
IV. Orbitals
• In the Bohr model, each orbit was
labeled with a single quantum number.
• For orbitals, it’s more complex, so we
need something else.
• Orbitals are labeled with a principal
quantum number (n)and a subshell
letter designation (s, p, d, f).
IV. Principal Quantum Number
• The principal
quantum number
specifies the
principal shell of the
orbital.
• Higher n means
higher energy.
IV. Subshells
• Each principal shell has one or more
subshells.
• Each subshell has a different “shape.”
IV. The s Orbitals
• s orbitals are
spherical.
• The 1s orbital is the
lowest possible
energy for an
electron; it is the
ground state.
IV. 1s vs. 2s
• A 2s orbital is bigger
and has more energy
than a 1s orbital.
• If a 1s electron in
hydrogen transitions
to 2s, then the
hydrogen atom is
now in an excited
state.
IV. The p Orbitals
• There are three p orbitals, each with a
different orientation.
IV. The d Orbitals
• There are five d orbitals.
IV. The f Orbitals
• There are seven f orbitals.
IV. Energy Order of Orbitals
IV. Electrons and Orbitals
• Electrons have an intrinsic spin
property. They can spin up or down.
• According to the Pauli exclusion
principle, only two electrons with
opposite spin can “occupy” an orbital.
• We use electron configurations or
orbital diagrams to show electrons in
atoms.
IV. The Hydrogen Atom
IV. Energy Order of Orbitals
• Order of the orbitals
can be obtained
from the periodic
table.
• But, if you don’t
have one, you can
remember the
orbitals with a
simple diagram.
IV. Hund’s Rule
• We know there are only two electrons
w/ opposite spin per orbital.
• What about when more than one of the
same type is available (for > s)?
IV. Sample Problem
• Write electron configurations and orbital
diagrams for the following. Write
condensed forms for the last two.




Al
Mn
Sr
Br
V. Blocks on the Periodic Table
• We can organize the Periodic Table into
blocks in which s, p, d, or f orbitals are
being filled.
• This allows us to easily write electron
configurations or orbital diagrams based
on an element’s location.
• Note there are 2 columns for s-block, 6
columns for p-block, 10 columns for dblock, and 14 columns for f-block. Why?
V. e- Config Based on Location
V. Using the Periodic Table
V. Two Types of Electrons
• Core electrons are those that are not in
the outermost principal shell.
• Valence electrons are those in the
outermost principal shell.
V. Valence Electrons
• The reactivity of an atom is determined
by its valence electrons.
• The valence electrons are loosely held
by the atom because they are the
furthest away from the nucleus.
• Thus, they can be easily lost or gained,
which leads to chemical
reactions/properties.
V. Valence Electrons and
Element Families
• Elements in same family have similar
reactivity because they have the same
valence electron configuration.
V. The Periodic Table
V. Why Group
17’s Form 1Anions
V. Why Group 1’s
Form 1+ Cations
VI. Periodic Trends
• The quantum mechanical model of the
atom allows prediction of some periodic
trends.
• We will examine the trends of atomic
size, ionization energy, and metallic
character.
VI. Shells and Subshells
• Shells are like layers of groups of
electrons around the nucleus.
 More shells = larger size.
• Subshells rest inside a shell and don’t
add any thickness to the shell.
• As we go across a period, we add
electrons AND protons.
• This knowledge will help us understand
the trends.
VI. Main Group Atomic Size
VI. Explaining the Trend
• Down a family is easy: opening more
shells, so atomic size must increase.
• Across a period, the principal shell stays
the same; electrons are just filling
subshells.
 However, each additional proton pulls
everything in closer.
VI. Ionization Energy
• Ionization energy is the energy needed
to take away an electron from an atom
in the gas phase.
 Na + ionization energy  Na+ + 1e-
• The ionization energy follows a clear
trend.
VI. Ionization Energy Trend
VI. Explaining the Trend
• The trend can be correlated with atomic
size.
• An electron is easier to remove if it is
further away from the nucleus.
• Thus, larger atoms will have LOWER
ionization energies than smaller atoms.
VI. Metallic Character
• One characteristic of metals is that they
tend to lose electrons.
• If we use this as a criteria for metallic
character, then atoms with low
ionization energies are more metallic
than those with high ionization energies.
• Thus, the trend in metallic character is
the opposite of the trend in ionization
energy.
VI. Metallic Character Trend
VI. Sample Problem
• Choose the appropriate atom in each
pair.






Larger atom: Pb or Po
Larger atom: Cl or Se
Higher ionization energy: Mg or Sr
Higher ionization energy: Cu or P
More metallic: Au or Cu
More metallic: N or S