Chapter 8
Electron
Configurations
and Periodicity
Contents and Concepts
Electronic Structure of Atoms
In the previous chapter, you learned that we
characterize an atomic orbital by four quantum
numbers: n, l, ml, and ms. In the first section, we
look further at electron spin; then we discuss how
electrons are distributed among the possible orbitals
of an atom.
1. Electron Spin and the Pauli Exclusion Principle
2. Building-Up Principle and the Periodic Table
3. Writing Electron Configurations Using the
Periodic Table
4. Orbital Diagrams of Atoms; Hund’s Rule
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Periodicity of the Elements
You learned how the periodic table can be
explained by the periodicity of the ground-state
configurations of the elements. Now we will look at
various aspects of the periodicity of the elements.
5. Mendeleev’s Predictions from the Periodic Table
6. Some Periodic Properties
7. Periodicity in the Main-Group Elements
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Learning Objectives
Electronic Structure of Atoms
1. Electron Spin and the Pauli Exclusion Principle
a. Define electron configuration and orbital
diagram.
b. State the Pauli exclusion principle.
c. Apply the Pauli exclusion principle.
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2. Building-Up Principle and the Periodic Table
a. Define building-up principle.
b. Define noble-gas core, pseudo-noble-gas
core, and valence electron.
c. Define main-group element and (d-block
and f-block) transition element.
3. Writing Electron Configurations Using the
Periodic Table
a. Determine the configuration of an atom
using the building-up principle.
b. Determine the configuration of an atom
using the period and group number.
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4. Orbital Diagrams of Atoms; Hund’s Rule
a. State Hund’s rule.
b. Apply Hund’s rule.
c. Define paramagnetic substance and
diamagnetic substance.
Periodicity of the Elements
5. Mendeleev’s Predictions from the Periodic
Table
a. Describe how Mendeleev predicted the
properties of undiscovered elements.
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6. Some Periodic Properties
a. State the periodic law.
b. State the general periodic trends in size of
atomic radii.
c. Define effective nuclear charge.
d. Determine relative atomic sizes from
periodic trends.
e. State the general periodic trends in
ionization energy.
f. Define first ionization energy.
g. Determine relative ionization energies from
periodic trends.
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h. Define electron affinity.
i. State the broad general trend in electron
affinity across any period.
7. Periodicity in the Main-Group Elements
a. Define basic oxide, acidic oxide, and
amphoteric oxide.
b. State the main group corresponding to an
alkali metal, an alkaline earth metal, a
chalcogen, a halogen, and a noble gas.
c. Describe the change in metallic/nonmetallic
character (or reactivities) in going through
any main group of elements.
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In 1921, Otto Stern and Walther Gerlach first
observed electron spin magnetism. In the diagram
below, a beam of hydrogen atoms divides in two
while passing through a magnetic field. This
correlates with the two values of ms: +½ and -½.
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The two possible
spin orientations
of an electron and
the conventions
for ms are
illustrated here.
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An electron configuration of an atom is a
particular distribution of electrons among available
subshells.
An orbital diagram of an atom shows how the
orbitals of a subshell are occupied by electrons.
Orbitals are represented with a circle; electrons
are represented with arrows up for ms= +½ or
down for ms= -½.
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The Pauli exclusion principle summarizes
experimental observations that no two electrons in
one atom can have the same four quantum
numbers.
That means that within one orbital, electrons must
have opposite spin. It also means that one orbital
can hold a maximum of two electrons (with
opposite spin).
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An s subshell, with one orbital, can hold a
maximum of 2 electrons.
A p subshell, with three orbitals, can hold a
maximum of 6 electrons.
A d subshell, with five orbitals, can hold a
maximum of 10 electrons.
An f subshell, with seven orbitals, can hold a
maximum of 14 electrons.
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The lowest-energy configuration of an atom is
called its ground state.
Any other configuration represents an excited
state.
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The building-up principle (or aufbau principle) is
a scheme used to reproduce the ground-state
electron configurations by successively filling
subshells with electrons in a specific order (the
building-up order).
This order generally corresponds to filling the
orbitals from lowest to highest energy. Note that
these energies are the total energy of the atom
rather than the energy of the subshells alone.
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1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
4f
5f
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This results in the following order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p,
6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
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Another way to learn the building-up order is to
correlate each subshell with a position on the
periodic table.
The principal quantum number, n, correlates with
the period number.
Groups IA and IIA correspond to the s subshell;
Groups IIIA through VIIIA correspond to the p
subshell; the “B” groups correspond to the d
subshell; and the bottom two rows correspond to
the f subshell. This is shown on the next slide.
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There are a few exceptions to the building-up
order prediction for the ground state.
Chromium (Z=24) and copper (Z=29) have been
found by experiment to have the following groundstate electron configurations:
Cr:
Cu:
1s2 2s2 2p6 3s2 3p6 3d5 4s1
1s2 2s2 2p6 3s2 3p6 3d10 4s1
In each case, the difference is in the 3d and 4s
subshells.
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There are several terms describing electron
configurations that are important.
The complete electron configuration shows every
subshell explicitly.
Br:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5
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The noble-gas configuration substitutes the
preceding noble gas for the core configuration and
explicitly shows subshells beyond that.
Br:
[Ar]3d104s24p5
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The pseudo-noble-gas core includes the noblegas subshells and the filled inner, (n – 1), d
subshell.
For bromine, the pseudo-noble-gas core is
[Ar]3d10
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The valence configuration consists of the
electrons outside the noble-gas or pseudo-noblegas core.
Br: 4s24p5
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For main-group (representative) elements, an s or
a p subshell is being filled.
For d-block transition elements, a d subshell is
being filled.
For f-block transition elements, an f subshell is
being filled.
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For main-group elements, the valence
configuration is in the form
nsAnpB
The sum of A and B is equal to the group number.
So, for an element in Group VA of the third period,
the valence configuration is
3s23p3
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?
Write the complete electron
configuration of the arsenic atom, As,
using the building-up principle.
For arsenic, As, Z = 33.
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3
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?
What are the electron configurations for
the valence electrons of arsenic and
zinc?
Arsenic is in period 4, Group VA.
Its valence configuration is 4s24p3.
Zinc, Z = 30, is a transition metal in
the first transition series.
Its noble-gas core is Ar, Z = 18.
Its valence configuration is 4s23d10.
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When n = 2, there are two subshells.
The s subshell has one orbital, which could hold
one electron.
The p subshell has three orbitals, which could hold
three electrons.
This would give a total of
four elements for the second period.
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In 1927, Friedrich Hund discovered, by
experiment, a rule for determining the lowestenergy configuration of electrons in orbitals of a
subshell.
Hund’s rule states that the lowest-energy
arrangement of electrons in a subshell is obtained
by putting electrons into separate orbitals of the
subshell with the same spin before pairing
electrons.
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For nitrogen, the orbital diagram would be
1s
2s
2p
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?
Write an orbital diagram for the ground
state of the nickel atom.
For nickel, Z = 28.
1s
3s
4s
2s
2p
3p
3d
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?
Which of the following electron
configurations or orbital diagrams are
allowed and which are not allowed by
the Pauli exclusion principle? If they
are not allowed, explain why?
a. 1s22s12p3
b. 1s22s12p8
c. 1s22s22p63s23p63d8
d. 1s22s22p63s23p63d11
e.
1s
2s
a. Allowed; excited.
b. p8 is not allowed.
c. Allowed; excited.
d. d11 is not allowed.
e. Not allowed;
electrons in one
orbital must have
opposite spins.
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Magnetic Properties of Atoms
Although an electron behaves like a tiny magnet,
two electrons that are opposite in spin cancel each
other. Only atoms with unpaired electrons exhibit
magnetic susceptibility.
This allows us to classify atoms based on their
behavior in a magnetic field.
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A paramagnetic substance is one that is weakly
attracted by a magnetic field, usually as the result
of unpaired electrons.
A diamagnetic substance is not attracted by a
magnetic field generally because it has only paired
electrons.
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You learned how the organization of the periodic
table can be explained by the periodicity of the
ground-state configurations of the elements. Now
we will look at various aspects of the periodicity of
the elements.
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Mendeleev’s periodic table generally organized
elements by increasing atomic mass and with
similar properties in columns. In some places,
there were missing elements whose properties he
predicted.
When gallium, scandium, and germanium were
isolated and characterized, their properties were
almost identical to those predicted by Mendeleev
for eka-aluminum, eka-boron, and eka-silicon,
respectively.
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Periodic law states that when the elements are
arranged by atomic number, their physical and
chemical properties vary periodically.
We will look in more detail at three periodic
properties: atomic radius, ionization energy, and
electron affinity.
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Atomic Radius
While an atom does not have a definite size, we
can define it in terms of covalent radii (the radius in
covalent compounds).
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Trends
Within each group (vertical column), the atomic
radius increases with the period number.
This trend is explained by the fact that each
successive shell is larger than the previous shell.
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Within each period (horizontal row), the atomic
radius tends to decrease with increasing atomic
number (nuclear charge).
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Effective Nuclear Charge
Effective nuclear charge is the positive charge that
an electron experiences from the nucleus. It is
equal to the nuclear charge, but is reduced by
shielding or screening from any intervening
electron distribution (inner shell electrons).
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Effective nuclear charge increases across a
period. Because the shell number (n) is the same
across a period, each successive atom
experiences a stronger nuclear charge. As a
result, the atomic size decreases across a period.
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Atomic radius is plotted against atomic number in the
graph below. Note the regular (periodic) variation.
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A representation of atomic radii is shown below.
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?
Refer to a periodic table and arrange
the following elements in order of
increasing atomic radius: Br, Se, Te.
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35
Se
Br
Te is larger than Se.
Se is larger than Br.
52
Te
Br < Se < Te
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First Ionization Energy (first ionization potential)
The minimum energy needed to remove the
highest-energy (outermost) electron from a neutral
atom in the gaseous state, thereby forming a
positive ion
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Trends
Going down a group, first ionization energy
decreases.
This trend is explained by understanding that the
smaller an atom, the harder it is to remove an
electron, so the larger the ionization energy.
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Generally, ionization energy increases with atomic
number.
Ionization energy is proportional to the effective
nuclear charge divided by the average distance
between the electron and the nucleus. Because
the distance between the electron and the nucleus
is inversely proportional to the effective nuclear
charge, ionization energy is inversely proportional
to the square of the effective nuclear charge.
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Small deviations occur between Groups IIA and
IIIA and between Groups VA and VIA.
Examining the valence configurations for these
groups helps us to understand these deviations:
IIA
IIIA
ns2
ns2np1
It takes less energy to remove the
np1 electron than the ns2 electron.
VA
VIA
ns2np3
ns2np4
It takes less energy to remove the
np4 electron than the np3 electron.
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These trends and reversals are visible in the graph
of ionization energy versus atomic number.
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The size of each sphere indicates the size of the
ionization energy in the figure below.
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Electrons can be successively removed from an
atom. Each successive ionization energy
increases, because the electron is removed from a
positive ion of increasing charge.
A dramatic increase occurs when the first electron
from the noble-gas core is removed.
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Left of the line, valence shell electrons are being
removed. Right of the line, noble-gas core
electrons are being removed.
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?
Refer to a periodic table and arrange
the following elements in order of
increasing ionization energy: As, Br,
Sb.
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35
As
Br
Sb is larger than As.
As is larger than Br.
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Sb
Ionization energies:
Sb < As < Br
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Electron affinity (E.A.)
The energy change for the process of adding an
electron to a neutral atom in the gaseous state to
form a negative ion
A negative energy change (exothermic) indicates a
stable anion is formed. The larger the negative
number, the more stable the anion. Small negative
energies indicate a less stable anion.
A positive energy change (endothermic) indicates
the anion is unstable.
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The electron affinity is > 0, so the element must be
in Group IIA or VIIIA.
The dramatic difference in ionization energies is at
the third ionization.
The element is in Group IIA.
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Broadly speaking, the trend is toward more
negative electron affinities going from left to right in
a period.
Let’s explore the periodic table by group.
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Groups IIA and VIIIA do not form stable anions;
their electron affinities are positive.
Group Valence
IA
ns1
IIIA ns2np1
IVA ns2np2
VA ns2np3
VIA ns2np4
VIIA ns2np5
Anion Valence
ns2
stable
ns2np2
stable
ns2np3
stable
ns2np4
not so stable
ns2np5
very stable
ns2np6
very stable
Except for the members of Group VA, these values
become increasingly negative with group number.
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Metallic Character
Elements with low ionization energies tend to be
metals. Those with high ionization energies tend to
be nonmetals. This can vary within a group as well
as within a period.
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Oxides
A basic oxide reacts with acids. Most metal
oxides are basic. If soluble, their water solutions
are basic.
An acidic oxide reacts with bases. Most nonmetal
oxides are acidic. If soluble, their water solutions
are acidic.
An amphoteric oxide reacts with both acids and
bases.
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Group IA, Alkali Metals (ns1)
These elements are metals; their reactivity
increases down the group.
The oxides have the formula M2O.
Hydrogen is a special case. It usually behaves as
a nonmetal, but at very high pressures it can
exhibit metallic properties.
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Group IIA, Alkaline Earth Metals (ns2)
These elements are metals; their reactivity
increases down the group.
The oxides have the formula MO.
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Group IIIA (ns2np1)
Boron is a metalloid; all other members of Group
IIIA are metals.
The oxide formula is R2O3.
B2O3 is acidic; Al2O3 and Ga2O3 are amphoteric;
the others are basic.
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Group IVA (ns2np2)
Carbon is a nonmetal; silicon and germanium are
metalloids; tin and lead are metals.
The oxide formula is RO2 and, for carbon and lead,
RO.
CO2, SiO2, and GeO2 are acidic (decreasingly so).
SnO2 and PbO2 are amphoteric.
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Some oxides of Group IVA
PbO
(yellow)
SiO2
(crystalline solid quartz)
PbO2
(dark brown)
SnO2 (white)
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Group VA (ns2np3)
Nitrogen and phosphorus are nonmetals; arsenic
and antimony are metalloids; bismuth is a metal.
The oxide formulas are R2O3 and R2O5, with some
molecular formulas being double these.
Nitrogen, phosphorus, and arsenic oxides are
acidic; antimony oxides are amphoteric; bismuth
oxide is basic.
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Group VIA, Chalcogens (ns2np4)
Oxygen, sulfur, and selenium are nonmetals;
tellurium is a metalloid; polonium is a metal.
The oxide formulas are RO2 and RO3.
Sulfur, selenium, and tellurium oxides are acidic
except for TeO2, which is amphoteric. PoO2 is also
amphoteric.
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Group VIIA, Halogens (ns2np5)
These elements are reactive nonmetals, with the
general molecular formula being X2. All isotopes of
astatine are radioactive with short half-lives. This
element might be expected to be a metalloid.
Each halogen forms several acidic oxides that are
generally unstable.
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Group VIIIA, Noble Gases (ns2np6)
These elements are generally unreactive, with only
the heavier elements forming unstable
compounds. They exist as gaseous atoms.
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For R2O5 oxides, R must be in Group VA.
R is a metalloid, so R could be As or Sb.
The oxide is acidic, so
R is arsenic, As.
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