Transcript (FOR STUDENTS 2015)

The Periodic Table
(very) Brief History
• 1869 Mendeleev* & Meyer published similar
tables
*First to be recognized at international convention
– Elements were ordered based on atomic mass
• Henry Moseley developed the atomic # concept
– Proved more accurate than Mendeleev’s atomic mass
method
The Periodic Law
• When elements are arranged by increasing atomic #
their chemical & physical properties show a periodic
pattern.
Review of Element Symbol
• Atomic #
– #protons
• Element Symbol
• Atomic Mass
54.938049
54.938049
– Weighted average of all
isotopes’ mass #
– Listed in AMUs
– Equal to MM (g/mol)
• Element Name
Periodic Table Layout
• Groups or Families
– Columns, Vertical
– 18 groups
• Periods
Groups to know
–
–
–
–
Group 1 - Alkali Metals
Group 2 - Alkaline Earth Metals
Group 17 Halogens
Group 18 - Noble Gases
– Rows, Horizontal
– 7 periods
• Kinds of Elements
– Metals, Nonmetals, Semi-metals
– Varying Properties
Electron Configurations
Using the Periodic Table
s-block elements:
1
Group 1: Alkali
H-
1s1
Li - 1s22s1
H
2
IIA
3
4
Li
Be
1.00797
2
3
Na- 1s22s22p63s1
K - 1s22s22p63s23p64s1
1
IA
1
4
5
6
7
6.939
11
9.0122
12
Na
Mg
22.9898 24.305
19
20
K
39.102
37
Rb
Ca
40.08
38
Sr
85.47
55
87.62
56
Cs
Ba
Fr
Ra
132.905 137.34
87
88
[223]
[226]
s-block elements:
Group 2: Alkaline Earth
Be - 1s22s2
Mg- 1s22s22p63s2
Ca - 1s22s22p63s23p64s2
Electron Configurations (con’t)
p-block elements:
Group 13: B -1s22s22p1
Group 14: Si
-1s22s22p63s23p2
Group 15: As-1s22s22p63s23p64s23d104p3
Group 16: S 1s22s22p63s23p4
Group 17: Br1s22s22p63s23p64s23d104p5
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
18
VIIIA
2
He
5
6
7
8
9
4.0026
10
B
C
N
O
F
Ne
Al
Si
P
S
Cl
Ar
Ge
As
Se
Br
Kr
In
Sn
Sb
Te
I
Xe
Tl
Pb
Bi
Po
At
Rn
10.811 12.0112 14.0067 15.9994 18.9984 20.179
13
14
15
16
17
18
26.9815 28.086 30.9738 32.064 35.453 39.948
31
32
33
34
35
36
Ga
7 65.37
49
72.59 74.9216 78.96 79.909 83.80
50
51
52
53
54
40 114.82 118.69 121.75 127.60 126.904 131.30
81
82
83
84
85
86
59 204.37 207.19 208.980 [210]
[210]
[222]
Noble Gas Configuration
Choose Noble Gas Prior to Element
FOCUS ON THE VALENCE SHELL
Put Noble gas in brackets,
ex. [Ne]
This represents the “inner
shell” of the element
Add the outer level electrons by
row # (2-7) s & p
**d level starts on 4th row**
4f & 5f start on 6th & 7th row
Examples:
Ni: [Ar] 4s23d8 **remember that 3d fills after 4s and before 4p
Sb: [Kr] 5s24d105p3
Trends in the Periodic Table
Definition: Predictable changes in properties of the
elements as you move through the table. (Realize
there are exceptions)
Decrease
Increase
Increase
Increas
e
Atomic Radii
1
1
H
2
Li
Be
3
Na Mg
11
12
4
K
5
Rb
6
Cs
7
Fr
3
19
37
55
87
21
Sc
38
39
56
Ba
88
Ra
He
5
B
- Ions are larger
Ca
Sr
2
Ionic Size
+Ions are Smaller
4
20
Ionization Energy
Electron Affinity
Y
57
22
Ti
40
Zr
72
La
Hf
89
104
Ac Rf
23
24
V
Cr
41
42
25
43
Nb Mo Tc
73
44
27
Co
45
Ru Rh
Ta
W
Re Os
75
76
Ir
105
106
107
108
109
Db
74
26
Mn Fe
Sg
77
13
28
29
Ni
Cu
46
47
Pd Ag
78
Pt
79
30
Zn
48
Cd
80
Au Hg
6
7
N
O
16
Al
Si
14
15
31
32
33
P
Ga Ge As
49
In
81
Tl
8
C
50
Sn
82
Pb
51
Sb
83
Bi
S
34
9
F
17
Cl
35
Se
Br
52
53
Te
84
Po
Bh Hs Mt
Atomic Radii - distance from the nucleus to outermost electron.
Ionization Energy - energy required to remove an electron (kJ/mol)
Electron Affinity – energy change when neutral atom gains electron
I
85
At
10
Ne
18
Ar
36
Kr
54
Xe
86
Rn
Ionization Energy
• In general, ionization energies of the main-group
elements increase across each period.
• This increase is caused by increasing nuclear charge.
• A higher charge more strongly attracts electrons in the same
energy level.
• Among the main-group elements, ionization energies
generally decrease down the groups.
• Electrons removed from atoms of each succeeding element
in a group are in higher energy levels, farther from the
nucleus.(electron shielding)
• The electrons are removed more easily.
Additional Trend Information
Electron Affinity
The energy associated with an atom gaining or losing an
electron. (kJ/mol)
• + Energy means it requires energy …not favorable
• - Energy means it gives up energy…favorable
Electronegativity
The ability to attract an electron during bonding
• Increases up a group and across a period
Electron Affinity
• The energy change that occurs when an electron is
acquired by a neutral atom is called the atom’s
electron affinity.
• Electron affinity generally increases across periods.
• Increasing nuclear charge along the same
sublevel attracts electrons more strongly
• Electron affinity generally decreases down groups.
• The larger an atom’s electron cloud is, the farther
away its outer electrons are from its nucleus.
Ionic Radii
• A positive ion is known as a cation.
• The formation of a cation by the loss of one or more
electrons always leads to a decrease in atomic radius.
• The electron cloud becomes smaller.
• The remaining electrons are drawn closer to the nucleus by its
unbalanced positive charge.
• A negative ion is known as an anion.
• The formation of an anion by the addition of one or
more electrons always leads to an increase in
atomic radius.
Ionic Radii, continued
• Cationic and anionic radii decrease across a period.
• The electron cloud shrinks due to the increasing
nuclear charge acting on the electrons in the same
main energy level.
• The outer electrons in both cations and anions are in
higher energy levels as one reads down a group.
• There is a gradual increase of ionic radii down a
group.
Electronegativity
• Valence electrons hold atoms together in chemical
compounds.
• In many compounds, the negative charge of the
valence electrons is concentrated closer to one
atom than to another.
• Electronegativity is a measure of the ability of an
atom in a chemical compound to attract electrons
from another atom in the compound.
• Electronegativities tend to increase across periods,
and decrease or remain about the same down a
group.
The Periodic Table
(very) Brief History
• 1800’s - Dobereiner introduced “Triads”
– 3 elements with similar properties
• 1865 - Newlands introduced “Law of Octaves”
– At this time 62 known elements
– 1st behaved like 8th, 2nd like 9th…
• 1869 Mendeleev* & Meyer published similar
tables
*First to be recognized at international convention
– Elements were ordered based on atomic mass
The Periodic Law
• When elements are arranged by increasing atomic #
their chemical & physical properties show a periodic
pattern.
1913 - Henry Moseley developed the atomic # concept.
– Proved more accurate than Mendeleev’s atomic mass
method
Review of Element Symbol
• Atomic #
– #protons
• Element Symbol
• Atomic Mass
54.938049
– Weighted average of all
isotopes’ mass #
– Listed in AMUs
– Equal to MM (g/mol)
• Element Name
Periodic Table Layout
• Groups or Families
– Columns, Vertical
– 18 groups
• Periods
Groups to know
–
–
–
–
Group 1 - Alkali Metals
Group 2 - Alkaline Earth Metals
Group 17 Halogens
Group 18 - Noble Gases
– Rows, Horizontal
– 7 periods
• Kinds of Elements
– Metals, Nonmetals, Semi-metals
– Varying Properties
Electron Configurations
Using the Periodic Table
s-block elements:
Group 1: Alkali
H - 1s1
Li - 1s22s1
Na- 1s22s22p63s1
K-
1
1
IA
1
H
2
IIA
3
4
1.00797
2
3
4
5
1s22s22p63s23p64s1
6
7
Li
6.939
11
Na
Be
9.0122
12
Mg
s-block elements:
Group 2: Alkaline Earth
Be - 1s22s2
22.9898 24.305
19
20
Mg- 1s22s22p63s2
39.102
37
40.08
38
Ca - 1s22s22p63s23p64s2
85.47
55
87.62
56
K
Rb
Cs
Ca
Sr
Ba
132.905 137.34
87
88
Fr
[223]
Ra
[226]
Electron Configurations (con’t)
p-block elements:
Group 13: B -1s22s22p1
13
IIIA
5
14
IVA
6
15
VA
7
16
17
VIA VIIA
8
9
18
VIIIA
2
He
4.0026
10
B
C
N
O
F
Ne
Al
Si
P
S
Cl
Ar
Ge
As
Se
Br
Kr
Sn
Sb
Te
I
Xe
Pb
Bi
Po
At
Rn
Group 14: Si -1s22s22p63s23p2
10.811 12.0112 14.0067 15.9994 18.9984 20.179
13
14
15
16
17
18
Group 15: As-1s22s22p63s23p64s23d104p3
26.9815 28.086 30.9738 32.064 35.453 39.948
31
32
33
34
35
36
Ga
7 65.37
49
Group 16: S 1s22s22p63s23p4
In
40 114.82 118.69 121.75 127.60 126.904 131.30
81
82
83
84
85
86
Tl
Group 17: Br-
72.59 74.9216 78.96 79.909 83.80
50
51
52
53
54
59 204.37 207.19 208.980 [210]
1s22s22p63s23p64s23d104p5
[210]
[222]
Noble Gas Configuration
Choose Noble Gas Prior to Element
FOCUS ON THE VALENCE SHELL
Put Noble gas in brackets,
ex. [Ne]
This represents the “inner
shell” of the element
Add the outer level electrons by
row # (2-7) s & p
**d level starts on 4th row**
4f & 5f start on 6th & 7th row
Examples:
Ni: [Ar] 4s23d8 **remember that 3d fills after 4s and before 4p
Sb: [Kr] 5s24d105p3
Trends in the Periodic Table
Definition: Predictable changes in properties of the
elements as you move through the table.
Decrease
Increase
Increas
e
Atomic Radii
1
Ionization Energy
2
1
H
2
Li
Be
4
+Ions are Smaller
3
11
12
Na Mg
- Ions are larger
4
K
5
Rb Sr
6
Cs Ba La Hf Ta
W Re Os
Ir
7
Fr Ra Ac Rf
109
3
19
Ionic Size
20
21
37
38
39
Y
Zr Nb Mo Tc
55
56
87
88
He
5
B
13
6
7
8
C
N
O
Al
Si
14
15
16
32
33
34
50
51
52
30
31
44
45
46
47
48
49
78
79
80
81
35
36
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
24
25
Br Kr
40
41
42
43
53
57
72
73
74
75
76
89
104
106
107
108
105
29
18
Ar
V
Db
28
17
Cl
23
Ti
27
S
10
Ne
22
Ca Sc
26
P
9
F
Ru Rh Pd Ag Cd In
77
Pt Au Hg Tl
Sn Sb Te
82
Pb
83
Bi
84
Po
Sg Bh Hs Mt
Atomic Radii - distance from the nucleus to outermost electron.
Ionization Energy - energy required to remove an electron (kJ/mol)
54
I
Xe
85
86
At Rn
Additional Trend Information
Electron Affinity
The energy associated with an atom gaining or losing an
electron. (kJ/mol)
• + Energy means it requires energy …not favorable
• - Energy means it gives up energy…favorable
Electronegativity
The ability to attract an electron during bonding
• Increases up a group and across a period