Transcript Periodic Table

Main Group Elements
Atomic radius is defined as being one half the
distance between identical nuclei bonded in a
molecule.
Atoms get smaller as you proceed from left to
right across a period (series).
The nucleus contains more protons and
the electron cloud contains more
electrons.
The increased charge results in a greater
attraction making the atom smaller.
8-1
Atomic Radius Across A Period
.
Na Mg Al
Si
P
S
Cl
Ar
Atomic Radius Decreases
8-2
Atomic Radius
Atomic Radius (pm)
Atomic Radius vs Atomic Number
250
Cs
Rb
K
200
Na
Li
150
100
Ar
50
He
Xe
Kr
Ne
0
0
10
20
30
40
50
60
Atomic Number
8-3
Main Group Elements
Atoms get larger as you proceed from top to
bottom down a group (family).
There is one more principal energy level
each time you go down a period.
The valence electrons get further from the
nucleus and feel less of an attraction by the
positive nucleus.
8-4
Atomic Radius Down A Group
.
Li
Na
Atomic
Radius
Increases
K
Rb
Cs
8-5
Atomic Radius
Atomic Radius (pm)
Atomic Radius vs Atomic Number
250
Cs
Rb
K
200
Na
Li
150
100
Ar
50
He
Xe
Kr
Ne
0
0
10
20
30
40
50
60
Atomic Number
8-6
Ionic Radius for Cations in a Period
Cations are smaller than the atoms from
which there are formed.
Metals give off their valence electrons, i.e.
Ca → Ca2+ + 2e-
1s22s22p63s23p64s2 → 1s22s22p63s23p6 + 2eThe entire highest numbered principal energy
is lost (n = 4) which also decreases the
number of electrons which decreases the
repulsion.
8-7
Ionic Radius Across A Period
.
Na Mg Al Si
P
S
Cl
Ar
Atomic radius decreases
for both cations and anions.
8-8
Ionic Radii for Anions in a Period
Anions are larger than the atoms from which
there are formed.
Nonmetals take in valence electrons to form a
complete octet, i.e.
N + 3e- → N31s22s22p3 + 3e- → 1s22s22p6
Adding electrons increases the repulsion
between electrons residing in the same
sublevel.
8-9
Ionic Radii for Anions in a Period
Compare nitrogen to oxygen:
N + 3e- → N31s22s22p3 + 3e- → 1s22s22p6
O + 2e- → O21s22s22p4 + 2e- → 1s22s22p6
N3- and O2- are isoelectronic but O2- has the
greater nuclear charge making it the smaller
anion.
8 - 10
Ionic Radius Across A Period
.
Na Mg Al Si
P
S
Cl
Ar
Atomic radius decreases
for both cations and anions.
8 - 11
Transition Elements
.
Cr2+
Cr3+
Mn2+
Fe2+
Fe3+
Co2+
Co3+
Ni2+
Cu+
Cu2+
Transition elements tend to have multiple
valence numbers.
Almost all the transition elements of the fourth
period form monatomic ions with a charge of +2.
8 - 12
Transition Elements
Because transition elements in the
d-block have 2 electrons in their valence
shell, they tend to react chemically the same.
Both lanthanoids and actinoids are found in
the f-block.
Lanthanoids occur in trace amounts in nature
and are called the rare earth elements.
Actinoids usually have large unstable nuclei
that undergo spontaneous radioactive decay.
8 - 13
Ionization Energy
First Ionization Energy
Ionization energy or ionization potential is the
minimum amount of energy needed to remove
an electron from the valence shell of a gaseous
atom.
Na(g) + IE1
Na+(g) + e-
IE indicates how easy it is for a metal to form a
cation.
8 - 14
Ionization Energy
Generally, IE increases across a period
(series) because of an increase in nuclear
charge.
IE increases as the size of the atom
decreases.
Nonmetals easily accept electrons causing
them to have a high IE.
8 - 15
Ionization Energy
As you move from left to right in a period, the
IE also depends on half-filled and
completely-filled orbitals.
 When an s sublevel is filled with 2
electrons, there is an increase in its
stability.
 When a p sublevel is half-filled with 3
electrons, there is an increase in its
stability.
8 - 16
Ionization Energy
 When a p sublevel is completely filled
with 6 electrons, there is an even greater
increase in its stability.
 The same is true for a d sublevel except
it is for 5 electrons and 10 electrons.
 Remember the two exceptions in
Period 4, Cr (3d5) and Cu (3d10) which
follows this same tendency.
8 - 17
Ionization Energy
The highest IE occurs for the noble gases
because they have a complete octet.
Generally, IE decreases from the top to the
bottom in a group or family because of the
addition of a principal energy level.
 IE decreases as the size of the atom
increases.
8 - 18
Ionization Energy
The second IE is larger than the first because
the second electron is being removed from a
cation rather than a neutral atom.
The third IE is larger than the second
because the third electron is being removed
from a cation with a +2 charge.
8 - 19
First Ionization Energy
First Ionization Energy
(kJ/mol)
First Ionization Energy vs Atomic Number
He
2500
Ne
2000
Ar
1500
Kr
Xe
1000
Rn
500
0
0
20
40
60
80
100
Atomic Number
8 - 20
Electronegativity
Electronegativity is a measure of the
attraction of an element for a shared pair of
electrons.
..
..
.. δ
H Cl
..
δ+
Comparing the electronegativity values of
hydrogen and chlorine, chlorine has a value
of 3.2 and that of hydrogen is 2.2.
8 - 21
Electronegativity
The origin of these values is unimportant and
the atom with the higher value is more
electronegative.
The most electronegative element is 9F
because it has the smallest atomic radius
with very few of it electrons shielding the
nucleus.
The least electronegative is 87Fr.
8 - 22
Electronegativity
Generally, electronegativity increases from
left to right within a period or series.
Generally, electronegativity decreases from
top to bottom within a group or family.
Electronegativity values are not assigned to
the noble gases because they are inactive.
The explanation for the trends in
electronegavity is the same as for ionization
energy.
8 - 23
Electronegativity
Electronegativity
Electronegativity vs Atomic Number
4.5
4.0
3.5
3.0
2.5
2.0
1.5
1.0
0.5
0.0
F
Cl
0
20
Br
40
I
At
60
80
100
Atomic Number
8 - 24
Electron Affinity
A measure of an atom’s tendency to gain
electrons in the gas phase.
A(g) + e-
A-(g) + thermal energy
Electron affinity is an irregular periodic
function of atomic number. In general, it
increases from left to right.
Noble gases are not included since they have
little or no tendency to gain electrons.
8 - 25
Periodic Trends in Density
Generally, the density increases as you
proceed from top to bottom in a group of
metals or nonmetals.
The atomic mass increases more rapidly
than the atomic radius.
8 - 26
Density (g/cm3)
.
Element
Density
Element
Density
Li
0.53
F
1.31
Na
0.97
Cl
1.56
K
0.86
Br
3.12
Rb
1.53
I
4.92
Cs
1.90
8 - 27
Periodic Trends in Density
Generally, the density increases as you
proceed from left to right in a period until you
reach the metalloids.
There is a big drop off in density in the
nonmetals (gases) but then starts to
increase.
8 - 28
Trends in Boiling and Melting Points
The boiling and melting points generally
decrease as you proceed from top to bottom
in the metals.
 This results from metallic bonding.
The boiling and melting points generally
increase as you proceed from top to bottom
in the nonmetals.
This results from Van der Waals forces.
8 - 29
Alkali Metals
.
Element
BP (°C)
MP (°C)
Li
1372
179
Na
892
98
K
774
64
Rb
679
39
Cs
690
28
Metallic Bonds
Boiling and Melting Points
8 - 30
Halogens
.
Element
BP (°C)
MP (°C)
F
-187
-223
Cl
-35
-101
Br
59
-72
I
185
114
Van der Waal Forces
Boiling and Melting Points
8 - 31
Metallic Characteristics
Metals are good conductors of heat and
electricity due to their “sea of electrons”.
Metals have shiny surfaces that are both
malleable and ductile due to their d-electrons.
Metals are malleable because they can be
hammered into a thin foil without breaking.
Metals are ductile because they can be
stretched into a thin wire without breaking.
8 - 32
Metallic Characteristics
Metals have three or fewer valence electrons
which they donate during chemical reactions.
Generally, metallic character increases as
you go down a group or family.
 The valence electrons are further from
the nucleus and are more shielded from
the nucleus.
8 - 33
Metallic Characteristics
Generally, metallic character decreases as
you proceed from left to right in a period
(series).
 Metals → Metalloids → Nonmetals
8 - 34
Metallic Characteristics
When a metal donates electrons it is said to
undergo oxidation.
Metals that are more easily oxidized will react
more readily in the presence of a nonmetal.
Group I and Group II are very active metals.
 Because they have such a strong tendency
to form compounds they are not found in
their elemental or free state.
8 - 35
Alkali Metals
The Group IA metals have an outer electron
configuration of ns1.
The loss of an electron to form a 1+ ion is the
basis of almost all reactions of the alkali
metals.
M → M+ + e-
8 - 36
Alkaline Earth Metals
The Group IIA metals have an outer electron
configuration of ns2.
The Group II metals are not as reactive as
the alkali metals because they need to lose
two electrons from a completely filled
s-sublevel in order to achieve a noble gas
configuration.
M → M2+ + 2e-
8 - 37
Nonmetallic Characteristics
Nonmetals are poor conductors of heat and
electricity.
Nonmetals have dull surfaces and are brittle.
Nonmetals have 4-7 valence electrons,
therefore they gain electrons.
Generally, nonmetallic character decreases
from top to bottom within a group or family.
8 - 38
Nonmetallic Characteristics
The atomic radius is a very important factor
in determining the reactivity in nonmetals.
F and Cl are the smallest halogens and will
more readily accept electrons in their
valence shell.
According to Coulomb’s Law, the positive
nucleus will attract valence electrons more in
a smaller atom.
8 - 39
Nonmetallic Characteristics
Generally, nonmetallic character increases
as you proceed from left to right in a period
(series).
 Metals → Metalloids → Nonmetals
8 - 40
Nonmetallic Characteristics
When a nonmetal accepts electrons it is said
to undergo oxidation.
Nonmetals that are more easily oxidized will
react more readily in the presence of a
nonmetal.
Group VII are very reactive nonmetals.
 Because they have such a strong tendency
to form compounds they are not found in
their elemental or free state.
8 - 41
Halogens
The common group VIIA elements are all
nonmetals. Each only needs a single electron
to achieve a noble gas configuration.
When reacting with metals, they form 1- ions.
2Na(s) + Cl2(g)
2NaCl(s)
When they have no other elements to react
with, they are found as diatomic molecules.
2F(g)
F2(g)
8 - 42
Noble Gases
Each noble gas has filled s and p sublevels
except for helium (1s2).
All are very unreactive.
A limited number of compounds have been
produced using xenon and krypton.
Xe(g) + F2(g)
XeF2(g)
8 - 43
The Anomoly of Hydrogen
Hydrogen is a nonmetallic gas at room
temperature.
While it may lose an electron to form H+, it
also can gain an electron to form H- (hydride).
2Na(l) + H2(g)
2NaH(s)
Hydrogen is placed in Group IA. Where else
could it go?
8 - 44
Semimetals (Metalloids)
These elements (B, Si, Ge, As, Sb, Te, and At)
along the “stairway” exhibit properties of
both metals and nonmetals.
They have some similarities with metals
because they are shiny and conduct
electricity.
They are similar to nonmetals because they
are brittle.
8 - 45
Allotropes
Allotropes are elements having more than
one form because of structural differences
(the way in which their atoms or molecules
are arranged).
The element oxygen has 3 forms:
 O – monatomic oxygen
O2 – molecular or diatomic oxygen
O3 - ozone
8 - 46