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ATOMIC ELECTRON
CONFIGURATIONS
AND CHEMICAL
PERIODICITY
CHAPTER 7
DMITRI MENDELEEV
• In 1870, Dmitri Mendeleev began to organize the
periodic table due to repeating patterns - mass
• one of the founders of the modern periodic table
• 1913 Mosley arranged it by number of protons (atomic
number)
7.1 PAULI EXCLUSION PRINCIPLE
• No atomic orbital can contain more than two
electrons and they must be opposite spin!
PAULI EXCLUSION PRINCIPLE
• Electrons can be
represented by
arrows in a box; this is
referred to as an
orbital box diagram.
• Figure 7.2 describes
the subshell filling
order, or the order
can be determined
by using the periodic
table.
7.2 ATOMIC SUBSHELL ENERGIES
• Electrons are assigned to subshells according to the
aufbau principle.
• assigned in order of “n + l” values
• 2 subshells with same “n + l” value  electrons assigned to
lower “n” first
EFFECTIVE NUCLEAR CHARGE, Z*
• Outer electrons may penetrate the inner electron
region. These core electrons screen the positive
nuclear charge from the outer (valence) electrons
and so the outer electrons experience an average
nuclear charge.
• Electrons are found in configurations that result in
the lowest energy for the atom.
7.3 ELECTRON CONFIGURATIONS
• The use of orbital notations is referred to as the
“spdf” notation.
• The core electrons can be summed up in the
“noble gas” notation.
• What is the spdf for the lithium atom?
• What is the orbital box diagram for the boron atom?
• What is the noble gas notation for potassium?
HUND’S RULE
• The most stable arrangement of electrons is that
with the maximum number of unpaired electrons;
minimizes electron-electron repulsions
• All single electrons must have parallel spins to reduce
repulsion.
ELECTRON CONFIGURATIONS
• Give the electron configuration of chlorine using the
spdf, noble gas, and orbital box notations.
• Write the electron configuration for Al using the
noble gas notation and give a set of quantum
numbers for each of the electrons with n = 3 (the
valence electrons).
ELECTRON CONFIGURATIONS
• Transition Elements
• Electrons may be found in s, p, and d sublevels
• Most of the time, the periodic table can be used to
determine electron filling
• Differences may occur between the expected and the
actual configurations… (not tested on the AP exam)
• Chromium is expected to be [Ar]3d44s2; however due to the
fact that the 3d and 4s are so similar in energy, each of the six
valence electrons is assigned to a different orbital and therefore
the actual configuration is [Ar]3d54s1
• Copper also has an unusual configuration of [Ar]3d104s1
ELECTRON CONFIGURATIONS
• Lathanides and Actinides
• f subshells are filled or partially filled in the inner transition
metals
• La (lanthanum) [Xe]5d16s2
• Ce (cerium) [Xe]4f15d16s2
PRACTICE PROBLEMS
• What element is 1s22s22p63s23p5?
• Do the spdf and orbital box diagram for
phosphorus.
• Do the spdf and noble gas for technetium (Tc) and
osmium (Os).
HOMEWORK
• After reading sections 7.1-7.3, you should be able to
do the following…
• p. 332 (11-21)
7.4 ELECTRON CONFIGURATIONS
OF IONS
• Electrons are removed from the outermost energy
level (shell of highest n).
• If there are more than one subshell in the outermost
level, the electrons are removed from maximum l
• Na+ – the 3s1 electron is removed [1s22s22p6]
• Fe2+ - [Ar]3d6
PRACTICE PROBLEM
• What is the electron configuration of V2+, V3+, and
Co3+? Are any of the ions paramagnetic? If so,
how many unpaired electrons are there?
7.5 ATOMIC PROPERTIES
• The similarities in properties of the elements are
the result of similar valence shell electron
configurations.
• Atomic Size
• Atomic radius is ½ the experimentally determined
distance between the centers of the two atoms.
• For the main group elements, atomic radius increases
going down a group due to the fact that the
outermost electrons have a higher n value.
• Atomic radius decreases going across a period due to
the fact that effective nuclear charge increases as
protons are added.
• Transition metals are different due to the large filled dsublevels; electron repulsion increases size toward the
right.
PRACTICE PROBLEMS
• Place the three elements Al, C, and Si in order of
increasing atomic radius.
• If the interatomic distance in Br2 is 228 pm, what is
the radius of Br? Using this value, and that for Cl (99
pm), estimate the distance between atoms in BrCl.
IONIZATION ENERGY
• The energy required to remove an electron from an
atom in the gas phase is referred to as ionization
energy.
• Excluding hydrogen, each atom can lose more
than one electron and therefore has a series of
ionization energies.
• Each successive electron removal requires more
energy because electrons are being removed from
an increasingly positive ion.
IONIZATION ENERGY
• For main group elements, first ionization energies
increase across a period due to the increase in
effective nuclear charge (increasing atomic
number).
• First ionization energies decrease down a group
occurs because the electron removed is farther
from the nucleus and therefore the nucleuselectron attraction is reduced.
ELECTRON AFFINITY
• The electron affinity, EA, of an atom is defined as
the energy of a process in which an electron is
acquired by the atom in the gas phase.
• Both electron affinity and ionization energy
represent the energy involved in the gain or loss of
an electron by an atom.
• An element with a high ionization energy generally has a
high affinity for an electron.
PRACTICE PROBLEMS
• Compare the three elements B, Al, and C.
• Place the three elements in order of increasing atomic
radius.
• Rank the elements in order of increasing ionization energy.
• Which element is expected to have the most negative
electron affinity value?
ION SIZES
• The radius of a cation is always smaller than that of
the atom from which it is derived. Once an
electron has been removed, the attractive force of
the protons are exerted over fewer electrons.
• Anions are always larger than the atoms from which
they are derived due to the addition of electron(s)
and an increase in electron-electron repulsions.
ION SIZES
7.6 PERIODIC TRENDS AND
CHEMICAL PROPERTIES
• Main group metals generally form cations with an
electron configuration equivalent to that of the
nearest noble gas.
• Non-metals generally acquire enough electrons to
form an anion with the electron configuration of the
next, higher noble gas.
HOMEWORK
• After reading sections 7.4-7.6, you should be able to
do the following…
• p. 333 (23-32)