PERIODICITY
Periodic Table
• Dmitri Mendeleev developed
the modern periodic table.
Argued that element
properties are periodic
functions of their atomic
weights.
• We now know that element
properties are periodic
functions of their ATOMIC
NUMBERS.
Periods in the Periodic Table
Groups in the Periodic Table
Regions of the Periodic
Table
6.1
Metals, Nonmetals, and
Metalloids
– Three classes of elements are metals,
nonmetals, and metalloids.
– Across a period, the properties of elements
become less metallic and more nonmetallic.
6.1
Metals, Nonmetals, and
Metalloids
»Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
Metals, Nonmetals, and
Metalloids
»Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
Metals, Nonmetals, and
Metalloids
»Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
Metals, Nonmetals, and
Metalloids
»Metals, Metalloids, and Nonmetals in the Periodic Table
6.1
Metals, Nonmetals, and
Metalloids
– Metals
• Metals are good conductors of heat and electric
current.
– 80% of elements are metals.
– Metals have a high luster, are ductile, and are malleable.
6.1
Metals, Nonmetals, and
Metalloids
» Uses of Iron, Copper, and Aluminum
6.1
Metals, Nonmetals, and
Metalloids
» Uses of Iron, Copper, and Aluminum
6.1
Metals, Nonmetals, and
Metalloids
» Uses of Iron, Copper, and Aluminum
6.1
Metals, Nonmetals, and
Metalloids
– Nonmetals
• In general, nonmetals are poor conductors of heat
and electric current.
– Most nonmetals are gases at room temperature.
– A few nonmetals are solids, such as sulfur and
phosphorus.
– One nonmetal, bromine, is a dark-red liquid.
6.1
Metals, Nonmetals, and
Metalloids
– Metalloids
• A metalloid generally has properties that are similar
to those of metals and nonmetals.
• The behavior of a metalloid can be controlled by
changing conditions.
Element Abundance
C
O
Al
Si
Fe
http://www.webelements.com/webelements/elements/text/Si/geol.html
Hydrogen
Shuttle main engines
use H2 and O2
The Hindenburg crash,
May 1939.
Group 1A: Alkali Metals
Reaction of
potassium + H2O
Cutting sodium metal
Group 2A: Alkaline Earth Metals
Magnesium
Magnesium
oxide
Calcium Carbonate—
Limestone
The Appian Way, Italy
Champagne cave carved into
chalk in France
Group 3A: B, Al, Ga, In, Tl
Aluminum
Boron halides
BF3 & BI3
Gems & Minerals
• Sapphire: Al2O3
with Fe3+ or Ti3+
impurity gives
blue whereas V3+
gives violet.
• Ruby: Al2O3
with Cr3+
impurity
Group 4A: C, Si, Ge, Sn, Pb
Quartz, SiO2
Diamond
Group 5A: N, P, As, Sb, Bi
Ammonia, NH3
White and red
phosphorus
Phosphorus
• Phosphorus first
isolated by Brandt
from urine, 1669
Group 6A: O, S, Se, Te, Po
Sulfuric acid dripping
from snot-tite in cave
in Mexico
Sulfur from
a volcano
Group 7A:
F, Cl, Br, I, At
Group 8A:
He, Ne, Ar, Kr, Xe, Rn
• Lighter than air balloons
• “Neon” signs
XeOF4
Transition Elements
Lanthanides and actinides
Iron in air gives
iron(III) oxide
Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total
electrons
3p
3s
2p
2s
1s
Beryllium
3p
3s
2p
2s
1s
Group 2A
Atomic number = 4
1s22s2 ---> 4 total
electrons
Boron
3p
3s
2p
2s
1s
Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons
Carbon
3p
3s
2p
2s
1s
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
Nitrogen
3p
3s
2p
2s
1s
Group 5A
Atomic number = 7
1s2 2s2 2p3 --->
7 total electrons
Oxygen
3p
3s
2p
2s
1s
Group 6A
Atomic number = 8
1s2 2s2 2p4 --->
8 total electrons
Fluorine
3p
3s
2p
2s
1s
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
9 total
electrons
Neon
3p
3s
2p
2s
1s
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
Colors of Transition Metal
Compounds
Iron
Cobalt
Nickel
Copper
Zinc
General Periodic Trends
•
•
•
•
Atomic and ionic size
Ionization energy
Electron affinity
Electronegativity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Effective Nuclear Charge, Z*
• Explains why E(2s) < E(2p)
• Z* is the nuclear charge experienced by the outermost
electrons. Is the result of the nuclear attraction being blocked
by the core electrons. Nuclear attraction increases with an
increase in protons
• Estimate Z* by --> [ Z - (no. core electrons) ]
• Charge felt by 2s e- in Li Z* = 3 - 2 = 1
• Be Z* = 4 - 2 = 2
• B
Z* = 5 - 2 = 3
and so on!
Effective Nuclear Charge, Z*
• Shielding effect remains constant across a period.
As the nuclear attraction increases across the
shielding effect is less effective.
• Shielding effect increases down a group thus
effectively blocking any increase in nuclear
attraction.
• Electrons with a higher quantum number have more
kinetic energy and thus are less affected by the
nuclear charge.
Each of these forces need to be accounted for in
each trend.
Effective Nuclear Charge,
Z*
• Atom
•
•
•
•
•
•
•
Li
Be
B
C
N
O
F
Z* Experienced by Electrons in
Valence Orbitals
+1.28
------Increase in
+2.58
Z* across a
+3.22
+3.85
period
+4.49
+5.13
Lithium
Periodic Trend in
the Reactivity of
Alkali Metals with
Water
Sodium
Potassium
Atomic Size
• Size goes UP on going down a
group.
• Because electrons are added
further from the nucleus, there is
less attraction, due to an increase
in sheilding effectiveness and in
increase in kinetic energy.
Atomic Size
• Size goes UP on going down a
group.
• Because electrons are added
further from the nucleus, there is
less attraction, due to an increase
in sheilding effectiveness and in
increase in kinetic energy.
General Outline for Trends
• Across a period
• Trend-define
• Nuclear attraction
• Down a group
– Trend, effect
• Nuclear attraction-define
– Shielding effect
once
– Trend, effect
– Shielding effect-define once
• Trend, effect
– Kinetic energy-define once
• Trend, effect
• Trend, effect
– Kinetic energy
• Trend, effect
Atomic Radius
• Atomic radius is the distance from the nucleus to
the valance electrons.
– Nuclear attraction (the attraction of the protons in the
nucleus on valance electrons) increases going down a
group. This should pull the electrons in closer to the
nucleus.
– Shielding effect (the blocking of nuclear attractions by
core electrons) Shielding effect increases down a group
offsetting the increase in nuclear attraction.
– Kinetic energy (the energy of valance electrons
associated with principle energy levels) increases down
a group allowing the valance electrons to orbit farther
from the nucleus increasing atomic radius.
Atomic Radius
– Nuclear attraction increases across a period.
This should pull the electrons in closer to the
nucleus decreasing atomic radius.
– Shielding effect remains constant across a
period not offsetting nuclear attraction.
– Kinetic energy remains constant across a period
so effective nuclear attraction is greater and the
atomic radius decreases.
Atomic Radii
Figure 8.9
Atomic Size
Size decreases across a period owing
to increase in Z*. Each added electron
feels a greater and greater + charge.
Large
Small
Ion Sizes
Li,152 pm
3e and 3p
Does+ the size go
up+ or down
Li , 60 pm
when
an
2e and 3losing
p
electron to form
a cation?
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming
a cation.
• CATIONS are SMALLER than the
atoms from which they come.
• The electron/proton attraction has
gone UP and so size DECREASES.
Ion Sizes
Does the size go up or
down when gaining an
electron to form an
anion?
Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming
an anion.
• ANIONS are LARGER than the atoms
from which they come.
• The electron/proton attraction has gone
DOWN and so size INCREASES.
• Trends in ion sizes are the same as atom
sizes.
Trends in Ion Sizes
Figure 8.13
Ionic Size
• Ionic size is the distance from the nucleus to the
valence electrons after an atom has lost or gained
electrons.
• Cations form when an atom loses one or more
electrons.
• Cations are smaller than the atoms from which
they form
• Ionic size - Cations
– Effective nuclear charge increases dramatically when
electrons are removed.
– Shielding effect decreases compared to the atom
because valence electrons are lost and some or all of the
core electrons become valence electrons.
– Kinetic energy the new valence electrons have less
kinetic energy to resist the pull of the nucleus.
Ionic Size
• Anions form when an atom gains one or more
electrons
• Ionic size - Anions
– After the addition of valence electron(s) the nuclear
attraction is diluted.
– Shielding effect remains still resisting the nuclear
attraction.
– Kinetic energy increases because of additional
repulsion due to more electrons in the valence shell
increasing the anions size.
– Anions are larger than the atoms from which they form
Ionization Energy
IE = energy required to remove an electron
from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Ionization Energy
IE = energy required to remove an
electron from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Mg+ (g) + 1451 kJ ---> Mg2+ (g) +
e- only 11 electrons.
Mg+ has 12 protons and
Therefore, IE for Mg+ > Mg.
Ionization Energy
Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a
shell of lower n.
General Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
Higher Z*.
Electrons held
more tightly.
Larger orbitals.
Electrons held less
tightly.
Trends in Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
3
Li
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
35
Trends in Ionization Energy
• IE increases across a period
because Z* increases.
• Metals lose electrons more
easily than nonmetals.
• Metals are good reducing
agents.
• Nonmetals lose electrons with
difficulty.
Trends in Ionization Energy
• IE decreases down a group
• Because size increases.
• Reducing ability generally
increases down the periodic
table.
• See reactions of Li, Na, K
Ionization Energy
See Screen 8.12
Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a
shell of lower n.
This is why ox. no. = Group no.
Ionization Energy
• Ionization energy is the energy
needed to remove an electron
from an atom.
• Nuclear attraction increases across
a period holding the electrons
tighter.
• Shielding effect is constant across
• Nuclear attraction
and does not offset the increase in
– Increases down a group holding
nuclear attraction.
the electrons tighter.
– Shielding effect increases down • Kinetic energy remains constant
across. With the same initial
a group offsetting the increase
energy valence electrons are
in nuclear attraction.
increasingly harder to remove due
– Kinetic energy increases down a
to the greater effective nuclear
group giving the electrons
charge.
greater initial energy. This
reduces the additional energy
needed to remove an electron.
Electron Affinity
A few elements GAIN electrons to
form anions.
Electron affinity is the energy
involved when an atom gains an
electron to form an anion.
A(g) + e- ---> A-(g) E.A. =
∆E
Electron Affinity of Oxygen
O atom [He] 
 

+ electron
O- ion [He] 
 
EA = - 141 kJ

∆E is EXOthermic
because O has
an affinity for an
e-.
Electron Affinity of
Nitrogen
N atom [He] 
 

+ electron
N- ion
[He] 


EA = 0 kJ

∆E is zero for Ndue to electronelectron
repulsions.
Trends in Electron Affinity
• See Figure 8.12 and
Appendix F
• Affinity for electron
increases across a period
(EA becomes more
positive).
• Affinity decreases down
a group (EA becomes
less positive).
Atom EA
F
+328 kJ
Cl +349 kJ
Br +325 kJ
I
+295 kJ
Trends in Electron Affinity
Electronegativity Values
See page 177 in text
Electronegativity
• Electronegativity is the tendency •
of an atom to remove an electron
from another atom when
•
forming a compound.
• Nuclear attraction
– Increases down a group attracting
the electrons more.
– Shielding effect increases down a
group offsetting the increase in
nuclear attraction.
– Kinetic energy increases down a
group giving the atom a larger
radius and increasing the proximity
of the nucleus to adjacent electrons
decreasing electronegativity
Nuclear attraction increases
across a period attracting
electrons more.
Shielding effect and kinetic
energy are constant across.
This increases effective nuclear
charge allowing the atom to
remove electron from other
atoms with lesser
electronegativity.