Download Periodic Table

Unit 5.2 Periodic Trends
Teacher: Dr. Van Der Sluys
Objectives
• Trends in
– Electron configuration
– Atomic Radii
– Electronegativity
– Ionization Energy
– Electron Affinity
Periodic Table
e-
configuration from the periodic table
1
IA
18
VIIIA
1
H
1s1
2
IIA
13
IIIA
2
Li
2s1
Na
3s1
Be
2s2
B
2p1
Mg
3s2
3
IIIB
4
IVB
5
VB
K
4s1
Rb
5s1
Ca
4s2
Sr
5s2
Sc
3d1
V
Cr Mn Fe Co
3d3 4s13d5 3d5 3d6 3d7
Nb Mo Tc Ru Rh
4d3 5s14d5 4d5 4d6 4d7
6
Cs
6s1
7
Fr
7s1
Ba
6s2
Ra
7s2
Ti
3d2
Zr
4d2
Hf
5d2
3
4
5
Y
4d1
La
5d1
Ac Rf
6d1 6d2
6
VIB
7
VIIB
8
Ta W Re Os
5d3 6s15d5 5d5 5d6
Db Sg Bh
6d3 7s16d5 6d5
9
VIIIB
Ir
5d7
Hs Mt
6
6d 6d7
10
11
IB
12
IIB
Ni
Cu Zn
3d8 4s13d10 3d10
Ni
Ag Cd
4d8 5s14d10 4d10
Ni Au Hg
5d8 6s15d10 5d10
Al
3p1
Ga
4p1
In
5p1
Tl
6p1
14
IVA
15
VA
16
VIA
17
VIIA
He
1s2
F
•B
C N O
Ne
•2p2 1 2p3 2p4 2p5 2p6
2p
Cl Ar
Si
P
S
3p2 3p3 3p4 3p5 3p6
Ge As Se Be Kr
4p2 4p3 4p4 4p5 4p6
Sn Sb Te
I
Xe
5p2 5p3 5p4 5p5 5p6
Pb Bi Po At Rn
6p2 6p3 6p4 6p5 6p6
1
Periodic Table and Subshells
Metals, Nonmetals and
Metalloids
• Metals are on the bottom left side of
the table and represent the vast
majority of the total number of
elements.
• Nonmetals are on the top right side of
the table.
• Metalloids are along the stair case,
except aluminum, which is a metal.
Periodic Table as a Map
West (South)
Mid-plains
East (North)
METALS
Alkali
Alkaline
Transition
These elements
tend to give up
e - and form
CATIONS
METALLOID
NON-METALS
Noble gas
Halogens
Calcogens
These elements
tend to accept
e - and form
ANIONS
These elements
will give up e- or
accept e-
• The periodic
table can be
classified by
the behavior
of their
electrons
1
IA
1
18
VIIIA
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
4
5
6
7
2
Atomic Radius
• The atomic radius of an element is half of
the distance between the centers of two
atoms of that element that are just touching
each other. Generally, the atomic radius
decreases across a period from left to right
and increases down a given group. The
atoms with the largest atomic radii are
located in Group I and at the bottom of
groups.
Electronegativity
•
Electronegativity is a measure of the attraction of an atom for the
electrons in a chemical bond. The higher the electronegativity of an atom,
the greater its attraction for bonding electrons. Electronegativity is related
to ionization energy (IE) and electron affinity.
•
Electrons with low ionization energies have low electronegativities
because their nuclei do not exert a strong attractive force on electrons.
Elements with high ionization energies have high electronegativities due
to the strong pull exerted on electrons by the nucleus. In a group, the
electronegativity decreases as atomic number increases, as a result of
increased distance between the valence electron and nucleus (greater
atomic radius). An example of an electropositive (i.e., low
electronegativity) element is cesium; an example of a highly
electronegative element is fluorine .
Ionization Energy
•
The ionization energy, or ionization potential, is the energy
required to completely remove an electron from a gaseous atom
or ion. The closer and more tightly bound an electron is to the
nucleus, the more difficult it will be to remove, and the higher its
ionization energy will be. The first ionization energy is the energy
required to remove one electron from the parent atom. The
second ionization energy is the energy required to remove a
second valence electron from the univalent ion to form the
divalent ion, and so on. Successive ionization energies increase.
The second ionization energy is always greater than the first
ionization energy. Ionization energies increase moving from left
to right across a period (decreasing atomic radius). Ionization
energy decreases moving down a group (increasing atomic
radius). Group I elements have low ionization energies because
the loss of an electron forms a stable octet.
3
Electron Affinity
•
Electron affinity reflects the ability of an atom to accept an
electron. It is the energy change that occurs when an electron is
added to a gaseous atom. Atoms with stronger effective nuclear
charge have greater electron affinity. Some generalizations can
be made about the electron affinities of certain groups in the
periodic table. The Group IIA elements, the alkaline earths, have
low electron affinity values. These elements are relatively stable
because they have filled s subshells. Group VIIA elements, the
halogens, have high electron affinities because the addition of an
electron to an atom results in a completely filled shell. Group VIII
elements, noble gases, have electron affinities near zero, since
each atom possesses a stable octet and will not accept an
electron readily. Elements of other groups have low electron
affinities.
1. Electronegativity Trend
Metals < 1.6
Nonmetals > 2.0
http://library.kcc.hawaii.edu/external/chemistry/electronegativity.html
2. Trend in Atomic Radius
4
3. Trend in Ionization Potential
Ionization potential:
The energy required to remove the valence (outer most) electron from
an atomic species. Largest toward NE corner of PT since these atoms
hold on to their valence e- the tightest.
4. Trend in Electron Affinity
Electron Affinity:
The energy release
when an electron is
added to an atom.
Most favorable
toward NE corner of
PT since these atoms
have a great affinity
for e-.
Trends in Groups: Shells
• The trends down a group can best be
explained by the increasing number of
shells of electrons.
• Each element in a group has the same
number of valence electrons, but they
are always in shells that have
increasing n values and are further
from the nucleus.
5
Core Charge
• The trends in periods from left to right
can best be explained by
– the increasing core charge
• Core charge = # protons - inner shell electrons
Core Charge
Group
1
2
H
1 - 0 = +1
Li
3 - 2 = +1
Na
11 - 10 = +1
K
19 - 18 = +1
Rb
37 - 36 = +1
Cs
55 - 54 = +1
He
2 - 0 = +2
Be
4 - 2 = +2
Mg
12
Ca
20
Sr
38
Ba
56
13
B
5 - 2 = +3
Al
13
Ga
31
In
49
Tl
81
14
C
5 - 2 = +4
Si
14
Ge
32
Sn
50
Pb
82
15
N
7 - 2 = +5
P
15
As
33
Sb
51
Bi
83
16
O
8 - 2 = +6
S
16
Se
34
Te
52
Po
84
17
F
9 - 2 = +7
Cl
17
Br
35
I
53
At
85
18
Ne
10 - 2 = +8
Ar
18
Kr
36
Xe
54
Rn
86
Core charge = atomic number - # core electrons
Summary of Trend
1. Electronegativity
Electronegativity:: Largest toward NE corner of PT
3. Ionization energy: Largest toward NE corner of PT
2. Atomic Radius: Largest toward SW corner of PT
6