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Lecture Presentation
Chapter 8
Periodic Properties
of the Elements
Catherine MacGowan
Armstrong Atlantic State University
© 2013 Pearson Education, Inc.
Periodic Table: Historic Perspective
•
Dmitri Mendeleev (1834-1907) developed
the modern periodic table.
– He argued that element properties are
periodic functions of their atomic
weights.
•
Periodic law: When the elements are
arranged in order of increasing atomic
mass, certain sets of properties recur
periodically.
– Elements having similar physical and
chemical properties fall within a
column.
•
We now know that element properties are
periodic functions of their ATOMIC
NUMBERS.
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Mendeleev’s Periodic Table
Mendeleev’s “periodic table” allowed for the prediction of
elements not yet discovered based on their physical and
chemical properties.
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Periodic Table: Organization
• The elements are arranged from left to right in increasing
ATOMIC NUMBER (number of protons an element has).
• Rows in the periodic table are referred to as PERIODS.
• Columns within the periodic table are sometimes referred to as
families.
– Families because the elements within the column have similar
physical and chemical properties
• The elements, their names, and symbols are given on the
PERIODIC TABLE.
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Electron Configuration
A map of atom’s electrons
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Electron Configuration
• According to quantum mechanics:
– Electron location around the atom’s nucleus is
described by the four quantum numbers:
• n (principle energy level)
• l (orbital type: s, p, d, f…)
• ml (orientation of orbital)
• ms (spin of electron in orbital)
– No two electrons can have the same four quantum
numbers: Pauli exclusion principle
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Writing Atomic Electron Configurations
Electron configuration for He, atomic number = 2
number of electrons
in the orbital
1 s2
orbital type (value of l)
Energy level (value of n)
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Arrangement of Electrons in Atoms
Electrons are arranged around an
atom’s nucleus according to the
Aufbau principle.
Electrons enter the lowest energy
level (n) and orbitals (l) available.
• n + l rule
•Electrons will enter lowest
energy subshell (l) that is
empty in that energy level
(n).
• Example: 2s vs. 2p
• For 2s: n = 2 and l = 0 (s)
So,
(n + l) = 2 = 0 = 2
• For 2p:
So,
n = 2 and l = 1 (p)
(n + l) = 2 + 1 = 3
The electron goes into the empty
2s orbital, not the 2p.
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General Energy Ordering of Electrons
in Atomic Orbitals
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Li: Electron Configuration
Box orbital diagram of Li
Li has an atomic
number of 3, so to
be neutral it must
have 3 electrons.
Its electron
configuration is
1s22s1.
3
3p
3s
E
2
2p
2s
1
1s
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N: Electron Configuration
N has an atomic number
of 7, so to be neutral it
must have 7 electrons.
Its electron configuration
is 1s22s22p3.
Box orbital diagram of N
3
3s
Here we see for the first
time HUND’S rule.
When placing electrons in
a set of orbitals having
the same energy
(degenerate orbitals), the
electrons are placed
singly as long as possible.
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3p
E
2
2p
2s
1
1s
Na: Electron Configuration
Na has an atomic number
of 11, so to be neutral it
must have 11 electrons.
Box orbital diagram of Na
Its electron configuration
is 1s22s22p63s1.
3
Another way to write its
electron configuration is
using “noble gas
notation”:
[Ne] 3s1
Where the first two
energy levels are filled or
the “neon core”
(1s22s22p6) + 3s1
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3p
3s
E
2
2p
2s
1
1s
Periodic Table and Electron Configuration
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Inner Core and Valence Electron
Configuration
• Inner core electrons: the filled “shells”
• Valence electrons: electrons occupying the outer
energy level
– These are the electrons that participate in:
• Bonding
• Making cations – lose
• Making anions – gain
Si:
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1s22s22p63s23p2
[Ne] 3s23p2
d and f block elements
•
Transition metals
– Their valence shell contains d
electrons
• d electrons first appear in the 4th
period
– 3d electrons in 4th period
– 4d electrons in 5th period
– 5d electrons in 6th period
– 6d electrons in 7th period
•
Inner transition metals
– Their valence shell contains f
electrons
• f electrons first appear in the 6th
period
– 4f electrons in 6th period
– 5f electrons in 7th period
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Anomalous Electron Configurations
•
The actual electron configuration of some elements, specifically those
found within the transition and inner transition metal section, have an
arrangement that is not predicted.
•
Reason: sublevel splitting
– The (n)s (where n = 4, 5, 6,…) has a lower energy than the (n-1)d
(where n = 3, 4, 5,…) when empty, and therefore electrons enter the
(n)s orbital before the (n-1)d.
• Example: 4s is filled first, then electrons enter the 3d orbital.
– The energy difference between s and d orbitals is not large.
– Also, as the (n-1)d orbitals are filled, their energy level becomes lower
than the (n)s orbitals.
•
Anomalous electron configurations occur when:
– the (n)s orbital only partially fills before electrons start entering the
(n−1)d or
– doesn’t fill at all and all electrons are in the (n-1)d orbital
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Examples of Anomalous
Electron Configurations
Predicted as per
Aufbau principle
[Ar]4s23d4
Cr =
Cu = [Ar]4s23d9
Mo = [Kr]5s24d4
Ru = [Kr]5s24d6
Pd = [Kr]5s24d8
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Actual as found
experimentally
Cr = [Ar]4s13d5
Cu = [Ar]4s13d10
Mo = [Kr]5s14d5
Ru = [Kr]5s14d7
Pd = [Kr]5s04d10
Electron configurations:
Ions
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The “Full-Shell” Look:
Eight Valence Electrons
•
Quantum-mechanical calculations show that an atom with eight valence
electrons should be unreactive – meaning the atom is very stable.
•
The noble gases have eight valence electrons and are all very stable and
unreactive.
• Exception is Helium (He), which has two valence electrons, but
that fills its valence shell.
•
If a “full-shell” look implies stability, then elements having either one
more electron (alkali metals) or one less electron (halogens) should be
predicted to be very reactive.
– Halogens, with seven valence electrons, are the most reactive of the
nonmetals.
• They gain ONE electron to form anions.
– Alkali metals, with one more electron than the preceding noble gas,
are the most reactive of the metals.
• They lose ONE electron to form cations.
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Electron Configuration for Ions
• Electrons are removed or added to the VALENCE SHELL.
– Valence shell is the outermost energy level that an atom’s
electrons occupy.
– Electrons in the valence shell are referred to as the valance
electrons.
• Cations
– Metals tend to form cations.
– Cations are elements having fewer electrons than protons.
– Cations are positively charged atoms.
• Anions
– Nonmetals tend to form anions.
– Anions are elements that have more electrons than protons.
– Anions are negatively charged atoms.
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Electron Configuration for Cations
• s column metals (1A, alkali and 2A, alkaline earths) lose their svalence electrons.
• Mg [Ne] 3s2
Mg2+ [Ne]
• p block elements (3A-7A or 13-17) lose their p-valence electrons
FIRST and then their s-valence electrons.
• Sn [Kr] 4d10 5s25p2
• Sn [Kr] 4d105s25p2
Sn2+ [Kr] 4d105s2
Sn4+ [Kr] 4d10
• Full d subshells are very stable.
– They behave as pseudo inner (full-shell) core electrons.
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Electron Configuration for Cations
• d block elements (1B-2B or 3-12) lose their s-valence electrons
FIRST and then their d-valence electrons.
• Fe [Ar] 3d64s2
• Fe [Ar] 3d64s2
Fe2+ [Ar] 3d6
Fe3+ [Ar] 3d5
*Half-filled d subshells are very stable
• f group elements (lanthanides 58-71 and actinides 90-103) lose
their s-valence electrons FIRST and then their f-valence
electrons.
• Bk [Rn] 5f97s2
• Bk [Rn] 5f97s2
Bk2+ [Rn] 5f9
Bk4+ [Rn] 5f7
*Half-filled f subshells are very stable.
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Ion Configurations: Anions
• Anions gain electrons.
• Electrons go in the available valance orbitals.
O:
O2-:
[He] 2s22p4
[He] 2s22p6 or [Ne]
I:
I-:
[Kr] 4d105s25p5
[Kr] 4d105s25p6 or [Xe]
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Magnetic Properties of
Transition Metal Atoms and Ions
• Electron configurations that result in unpaired
electrons mean that the atom or ion will have a net
magnetic field. This is called paramagnetism.
– Will be attracted to a magnetic field
• Electron configurations that result in all paired
electrons mean that the atom or ion will have no
magnetic field. This is called diamagnetism.
– Slightly repelled by a magnetic field
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Why Iron (III) Oxide Is Magnetic
• Ions WITH UNPAIRED ELECTRONS are
PARAMAGNETIC.
• Ions having NO UNPAIRED ELECTRONS are
DIAMAGNETIC.
– Fe3+ ([Ar] 3d5 )in Fe2O3 has five unpaired electrons.
4s
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4s
3d
3d
Fe
Fe3+
Periodic Trends
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Electron Configuration, Element Properties,
and the Periodic Table
• The number of valence electrons largely
determines the behavior of an element.
– Chemical and some physical
• The number of valence electrons follows
a periodic pattern; the properties of the
elements should also be periodic.
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Effective Nuclear Charge
Effective nuclear charge (Zeff):
•
It is the pull/force an electron “feels” from the nucleus (protons).
•
The closer the electrons are to the nucleus, the greater the “pull” on
the electrons.
•
The greater the Zeff, the more tightly the electrons are held and the
more energy needed to remove the electrons.
– Electrons located farthest from the nucleus (e.g., valence shell)
experience less Zeff.
•
Zeff explains the reason for the periodic properties and trends of the
elements.
•
Zeff general trend:
Zeff increases going across periods.
Zeff decreases going down periods.
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Effective Nuclear Charge, Zeff,
and Screening/Shielding Aspect
•
Zeff increases across a period owing
to incomplete shielding by inner
electrons in atomic orbitals
(subshells).
•
Shielding ability of subshells:
s>p>d>f
•
Estimate Zeff
= [Z (atomic #) - (# inner electrons)]
– Pull felt by 2s electron in Li:
Zeff = 3 - 2 = 1
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Zeff Experienced by Outermost Valence
Electron
Li
B
C
N
O
F
+1.28
+2.58
+3.22
+3.85
+4.49
+5.13
Observed
+1.00
+1.00
+1.00
+1.00
+1.00
+1.00
Predicted
3-2=1
5-4=1
6-5=1
7-6=1
8-7=1
9-8=1
•
Zeff increases as you go across a period due to the ineffective
shielding of the atomic orbitals (subshells) within the period.
•
The values observed are different than predicted values.
•
Reason: shielding abilities of the atomic orbitals (subshells)
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Atomic Radii
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Atomic Radii
•
Radii (size) increase going
down a group due to
decrease in Zeff.
– As electrons are added in
different levels (shell)
farther from the nucleus,
there is less attraction by
the protons in the nucleus
to the orbiting electrons.
•
Radii (size) decrease going
across a period due to the
increase in Zeff.
– Ineffective shielding of the
orbitals (subshells),
resulting in an increase in
attraction by protons in
the nucleus to the orbiting
electrons
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Which has the larger atomic radius?
a.
b.
c.
d.
N or F
C or Ge
B or Al
Li or K
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Which has the larger atomic radius?
ANSWER
a. N or F
N is because it is more left in the
period.
b. C or Ge
Ge is because it is farther down the
column.
c. B or Al
Al is because it farther down the
column and to the left in the period.
d. Li or K
K is because it is farther down the
column.
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Atomic Radii vs. Cation Radii:
Size Comparison
+
Li, 152 pm
(3 e- and 3 p+)
Li+, 78 pm
(2 e- and 3 p+)
• CATIONS are SMALLER than the atoms from
which they come.
• The electron/proton attraction has
INCREASED, so size DECREASES.
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Ionic Radii: Cations
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Atomic Radii vs. Anion Radii:
Size Comparison
F, 71 pm
(9 e- and 9 p+)
F-, 133 pm
(10 e- and 9 p+)
• ANIONS are LARGER than the atoms from which
they come.
• The electron/proton attraction has DECREASED,
so size INCREASES.
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Ionic Radii: Anions
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Choose the larger of each pair.
• S or S2−
• Ca or Ca2+
• Br− or Kr
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Answer: Choose the larger of each pair.
•
S or S2−
S2− is larger because there are more electrons (18 e−)
for the 16 protons to hold.
The anion is larger than the neutral atom.
•
Ca or Ca2+
Ca is larger because its valence shell electrons have
been lost to form Ca2+.
The cation is always smaller than the neutral atom.
•
Br− or Kr
Br− is larger because it has fewer protons (35 p+) to
hold the 36 electrons than does Kr (36 p+).
*For isoelectronic species, the more negative the
charge, the larger the atom or ion.
* An isoelectronic series is elements or ions having the
same electronic configuration.
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General Periodic Trend for
1st Ionization Energy
IE is the energy required to remove an
electron from an atom in the gas phase.
•
1st IE increases across a period
because Zeff increases due to poor
shielding ability of p electrons.
•
1st IE decreases going down a
column due to increase in atomic
size and decrease in Zeff.
– This is why metals lose
electrons more easily than
nonmetals and are good
reducing agents.
• Have lower 1st IE
– Nonmetals lose electrons with
difficulty.
• Have higher 1st IE
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Periodic Trend: Ionization Energy (IE)
IE is endothermic, DH (+), in nature.
1st IE: Mg(g) + 738 kJ  Mg+(g) + e2nd IE: Mg+(g) + 1451 kJ  Mg2+(g) + e-
3rd IE: Mg2+(g) + 7733 kJ  Mg3+(g) + eNOTE:
– Mg+ has 12 protons and only 11 electrons.
• Therefore, 1st IE for Mg+ > Mg.
– Mg2+ has 12 protons and only 10 electrons.
• Therefore, 2nd IE for Mg2+ > Mg+.
– Energy cost is VERY high to dip into a full
(inner core) shell (lower zx).
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Trends in 1st Ionization Energy
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Choose the atom in each set having the
higher 1st ionization energy.
a. Al or S
b. As or Sb
c. N or Si
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Answers: Choose the atom in each set
having the higher 1st ionization energy.
ANSWER
a. Al or S
b. As or Sb
sulfur
arsenic
c. N or Si
nitrogen
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General Trends in Electron Affinity (EA)
Electron affinity is the change in energy (DE) involved when a
atom gains an electron to form an anion.
A(g) + e-  A-(g)
EA is about a ∆E
•
EA can be either endothermic or exothermic in nature.
– Why either energy exchange?
• It is due to electron-electron repulsion within orbitals
and the volume of the atom.
•
General trends:
– EA increases across a period.
• EA becomes more positive due to increase in Zeff.
– EA decreases down a group.
• EA becomes less positive due to decrease in Zeff.
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General Trends in Electron Affinity (EA)
Electron affinity is the willingness to
accept electrons into the valence shell.
•
The fluoride ion has a LOWER EA
than the chloride ion because it has
a smaller volume and greater e e
than chloride.
•
EA greatest for halogens
•
EA greater for nonmetals vs. metals
•
A LARGER EA value means a very
STABLE ion.
– Larger EA is more negative energy
• More stable
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1st and 2nd Electron Affinity of Oxygen
O atom:
[He] 
 

+ electron
∆E is for the first EA is
EXOthermic because
oxygen has an affinity for
an electron (-141 kJ/mole).
1st EA = -141 kJ
O- ion:
[He] 
O2- ion:

[He]
 
 


DE for second EA is
ENDOthermic because of
increase in electronelectron repulsion (780
kJ/mol).
2nd EA = +780 kJ
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Properties of Metals and Nonmetals
Metal element
characteristics
• Malleable and ductile
• Shiny, lustrous, reflect
light
• Conduct heat and
electricity
• Most oxides are basic and
ionic
• Form cations in solution
• Lose electrons in
reactions
– oxidized
Nonmetal element
characteristics
• Brittle in solid state
• Dull, nonreflective solid
surface
• Electrical and thermal
insulators
• Most oxides are acidic
and molecular.
• Form anions and
polyatomic anions
• Gain electrons in
reactions
– reduced
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Metallic Character
• Metallic character of an element is about how closely its
properties match the ideal properties of a metal.
– The more malleable and ductile, the better the conductor
and easier to ionize.
• Metallic character decreases from left to right across a
period.
– Metals are found at the left of the period and nonmetals
are to the right.
• Metallic character increases down the column.
– Nonmetals are found at the top of the middle maingroup elements and metals are found at the bottom.
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Metallic Trend
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