Transcript Chapter 7

Chapter 7
Atomic Energies and
Periodicity
Department of Chemistry
and Biochemistry
Seton Hall University
Nuclear Charge
• n - influences orbital energy
• Z - nuclear charge also has a
large effect
• We can measure this by
ionization energies (IE)
– A  A + e-
• Consider H and He+
– H  H+ + e– He  He+ + e-
2.18  10-18 J
8.72  10-18 J
• Orbital stability increases with
Z2
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Electron-electron
Repulsion
• Negatively charged electron is
attracted to the positively
charged nucleus but repelled by
negatively charged electrons
• Screening, , is a measure of
the extent to which some of the
attraction of an electron to the
nucleus is cancelled out by the
other electrons
• Effective nuclear charge
– Zeff = Z - 
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Screening
• Complete screening would
mean that each electron would
experience a charge of +1
• Consider He
– w/o screening the IE would be the
same as for He+
– Complete screening the IE would
be the same as for H
– Actual IE is between the two
values
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Screening
• Screening is incomplete because
both electrons occupy an extended
region of space, so neither is
completely effective at screening the
other from the He2+ nucleus
• Compact orbitals (low values of n)
are more effective as screening since
they are packed tightly around the
nucleus
• Therefore,  decreases with orbital
size (as n increases)
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Screening
• Electrons in orbitals of a given value
n screen the electrons in orbitals
with larger values of n
• Screening also depends on orbital
shape (electron density plots, 2 vs r,
help show this)
• Generally, the larger the value of l,
the more that orbital is screened by
smaller, more compact orbitals
• Quantitative information about this
can be obtained from photoelectron
spectroscopy
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Structure of the periodic
table
• The periodic table is arranged
the way it is because the
properties of the elements
follow periodic trends
• Elements in the same column
have similar properties
• Elemental properties change
across a row (period)
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Electron configurations
• The Pauli Exclusion Principle
– No two electrons can have the
same four quantum numbers
• Hund’s rule
– The most stable configuration is
the one with the most unpaired
electrons
• The aufbau principle
– each successive electron is placed
in the most stable orbital whose
quantum numbers are not already
assigned to another electron
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Orbital diagrams and
rules
• The Pauli Exclusion Principle - no
two electrons may have the same
four quantum numbers.
• Practically, if two electrons are in the
same orbital, they have opposite
spins
• Hund’s Rule - when filling a
subshell, electrons will avoid
entering an orbital that already has
an electronic in it until there is no
other alternative
• Consider the dorm room analogy (I
suggested this to the author!!!)
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Summary of the rules
• Each electron in an atom
occupies the most stable orbital
available
• No two electrons can have the
same four quantum numbers
• The higher the value of n, the
less stable the orbital
• For equal values of n, the higher
value of l, the less stable the
orbital
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Shell designation
• The shell is indicated by the
principle quantum number n
• The subshell is indicated by the
letter appropriate to the value of
l
• The number of electrons in the
subshell is indicated by a right
superscript
• For example, 4p3
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Electronic
configurations
• We use only as many subshells
and shells as are needed for the
number of electrons
• The number of available
subshells depends on the shell
that is being filled
– n = 1 only has an s subshell
– n = 2 has s and p subshells
– n = 3 has s, p and d subshells
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Example
• Consider S
• Sulfur has 16 electrons
• Electronic configuration is
therefore
1s22s22p63s23p4
• d and f subshells are used for
heavier elements
• You are expected to do this for
any element up to Ar
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Core and valence shells
• Chemically, we find that the
electrons in the shell with the highest
value of n are the ones involved in
chemical reactions
• This shell is termed the valence shell
• Electrons in shells with lower n
values are chemically unreactive
because they are of such low energy.
• These shells are grouped together as
the core
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Electron configurations and the
periodic table
• We develop a shorthand for the
electron configuration by noting
that the core is really the same
as the electron configuration for
the noble gas that occurs earlier
in the periodic table
• E.g. for S (1s22s22p63s23p4), the
core is 1s22s22p6 which is the
same as the electron
configuration for Ne
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Atomic properties
• Ionization energy (IE)
A(g)  A+(g) + e• Electron affinity (EA)
A(g) + e-  A-(g)
• Ion sizes
– Cations are smaller than the
neutral atom
– Anions are larger that the neutral
atom
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Electron configuration
shorthand
• We can then write the electron
configuration of S as [Ne]3s23p4
• We note that the valence shell
electron configuration has the
same pattern for elements in the
same group
• For S (a chalcogen) all the
elements have the valence
electron configuration
[core]ns2np4
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Periodic trends
• Atomic radii decrease across a
period
• Atomic radii increase down a
group
• Ionization energies increase
across a period
• Ionization energies decrease
down a group
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Near degenerate orbitals
• degenerate orbitals are those
that have the same energy
• normally, certain orbitals will be
degenerate for quantum
mechanical reasons
• near degenerate orbitals have
close to the same energy for a
variety of reasons
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Ion electronic
configurations
• Electronic configurations for
ions involves adding or
subtracting electrons from the
appropriate atomic
configuration
• Example: Na  Na+
– 1s22s22p63s1  1s22s22p6
• Example: Cl  Cl– 1s22s22p63s23p5 
1s22s22p63s23p6
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Magnetic properties
• The spin of electrons generates a
magnetic field
• Two types of magnetism
• Diamagnetism - all electrons are
paired
• Paramagnetism - one or more
electrons are unpaired
• In solids, two types of condensed
phase magnetism results in bulk
magnetic properties ferromagnetism and
antiferromagnetism
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Energetics of ionic
compounds
• Ions in solids have very strong
attractions (ionic bonding)
• Due mostly to cation-anion
attraction, and includes a
component termed lattice
energy
• We can calculate this energy
from a Born Haber cycle
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Path yielding a net
reaction
• Vaporization Evaporization = 108 kJ/mol
• Ionization E = IE = 495.5 kJ/mol
• Bond breakage E = ½(bond energy) =
120 kJ/mol
• Ionization E = EA = -348.5 kJ/mol
• Condensation - includes all ion-ion
attractive and repulsive interactions (the
lattice energy)
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The Born-Haber Cycle
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Calculating the lattice
energy
• Coulomb’s law allows us to calculate the
electrical force between charged particles
• q1,q2 are the electrical charges of the
particles
• k = 1.389  105 kJ pm/mol
• r = interionic distance in pm
Ecoulomb
(q1 )( q2 )
k
r
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Calculating the lattice
energy
• Result of calculation yields a
value of -444 kJ/mol
• This includes only part of the
lattice energy, since the
coulombic interactions do not
stop at the individual ions pairs.
• An expansion of Coulomb’s law
to include the three dimensional
ion interactions yields a value
for the lattice energy of -781
kJ/mol
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3 D interaction in crystal
• Note that NaCl
extends in all
directions
• Each ion
experiences
attractions and
repulsions from
other ions past
the ones
directly in
contact
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The overall ionic
bonding energy
• The energy for the overall
process:
Na(s) + ½Cl2 (g)  NaCl(s)
• Calculated = -406 kJ/mol
• Actual = -411 kJ/mol
• This treatment assumes the
interaction between Na+ and Clis only ionic. The slight
discrepancy is ascribed to a
small degree of electron sharing
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Why not
2+
2Na Cl ?
• Main reason is the very large
ionization energy of the core of
Na
Na  Na+ IE1 = 495.5 kJ/mol
Na+  Na2+ IE2 = 4562 kJ/mol
• EA2 for Cl is expected to be
large and positive
• Basic point is that it costs way
too much energy to ionize the
core of Na
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Ion stability
• Group 1 and 2 ions will lose all
of their valence electrons
• Above Group 2, removal of all
valence electrons is generally
not observed
• Anions will generally add
enough valence electrons to fill
the valence shell
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