Transcript Periodic Table HTHS

Periodic Table
Dmitri Mendeleev
• Russian chemist
• Organized elements by
properties
• Arranged elements by
atomic mass
• Predicted existence of
several unknown elements
• Element 101 Mendeleevium (Md)
Dmitri Mendeleev
Mendeleev’s Periodic
Table
Group I
II
III
IV
V
VI
VII
Period
1
H=1
2
Li = 7
Be= 9.4
B = 11
C = 12
N = 14
O = 16
F = 19
3
Na = 23
Mg = 24
Al = 27.3
Si = 28
P = 31
S = 32
Cl = 35.5
4
K = 39
Ca = 40
? = 44
Ti = 48
V = 51
Cr = 52
Mn = 55
5
Cu = 63
Zn = 65
? = 68
? = 72
As = 75
Se = 78
Br = 80
6
Rb = 85
Sr = 87
? Yt = 88
Zr = 90
Nb = 94
Mo = 96
? = 100
7
Ag = 108
Cd = 112
In = 113
Sn = 118
Sb = 122
Te = 125
I = 127
8
Cs = 133
Ba = 137
?Di = 138
?Ce = 140
?Er = 178
?La = 180
Ta = 182
W = 184
Tl = 204
Pb = 207
Bi = 208
VIII
Fe =56, Co = 59,
Ni = 59
Ru= 104, Rh = 104,
Pd = 106
9
10
11
12
Au = 199
Hg = 200
Th = 231
U = 240
Os = 195, Ir = 197,
Pt = 198
Elements Properties are Predicted
Property
Mendeleev’s Predictions in 1871
Observed Properties
Scandium (Discovered in 1877)
Molar Mass
Oxide formula
Density of oxide
Solubility of oxide
44 g
M2O3
3.5 g / ml
Dissolves in acids
43.7 g
Sc2O3
3.86 g / ml
Dissolves in acids
Gallium (Discovered in 1875)
Molar mass
Density of metal
Melting temperature
Oxide formula
Solubility of oxide
68 g
6.0 g / ml
Low
M2O3
Dissolves in ammonia solution
69.4 g
5.96 g / ml
30 0C
Ga2O3
Dissolves in ammonia
Germanium (Discovered in 1886)
Molar mass
Density of metal
Color of metal
Melting temperature
Oxide formula
Density of oxide
Chloride formula
Density of chloride
Boiling temperature
of chloride
72 g
5.5 g / ml
Dark gray
High
MO2
4.7 g / ml
MCl4
1.9 g / ml
Below 100 oC
71.9 g
5.47 g / ml
Grayish, white
900 0C
GeO2
4.70 g / ml
GeCl4
1.89 g / ml
86 0C
Modern Periodic Table
•
•
•
•
•
Based on Mendeleev’s system
Arrange elements by atomic number
Elements organized by properties
Columns are groups or families(1 to 18)
Rows are periods (1 to 7) showing the
variation of chemical and physical
properties.
Groups of Elements
1A
1
H
1
2
3
Be
3
4
7
2A
Alkali earth metals
6A
Oxygen group
Transition metals
7A
Halogens
3A
Boron group
8A
Noble gases
4A
Carbon group
8A
He
3A 4A
B C
5A 6A 7A 2
N O F Ne
Hydrogen
Inner transition metals
5
6
7
8
9
10
Al
Si
P
S
Cl
Ar
8B
K
3B 4B 5B 6B 7B
1B 2B 13 14 15 16 17
Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
12
20
21
22
Rb Sr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd
In
39
40
41
42
49
Hf
Ta
W
72
73
74
37
6
Nitrogen group
Na Mg
19
5
5A
2A
Li
11
4
1A Alkali metals
38
Cs Ba
55
56
Fr
Ra
87
88
*
W
25
43
26
44
Re Os
75
76
27
28
29
47
30
45
46
Ir
Pt Au Hg
Tl
77
78
81
79
48
31
80
32
33
34
Sn Sb Te
50
51
Pb Bi
82
83
52
35
36
I
Xe
53
54
Po At Rn
84
85
86
105
106
107
108
109
La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
57
W
24
Kr
Rf Db Sg Bh Hs Mt
104
*
23
18
58
59
Ac Th Pa
89
90
91
60
U
92
61
62
63
64
65
66
Np Pu Am Cm Bk Cf
93
94
95
96
97
98
67
68
69
70
71
Es Fm Md No Lr
99
100
101
102
103
Metals and Nonmetals
1
2
3
H
He
1
2
Li
Be
B
C
3
4
5
Na Mg
11
4
K
19
5
7
Ca Sc
O
F
Ne
6
7
8
9
10
Al
Si
P
S
Cl
Ar
13
14
15
16
17
18
Ti
V
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
Kr
23
24
35
36
I
Xe
53
54
20
21
22
Rb Sr
Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd
In
39
40
41
42
49
Hf
Ta
W
72
73
74
37
6
12
N
38
Cs Ba
55
56
Fr
Ra
87
88
*
W
Nonmetals
25
26
27
28
29
30
METALS
43
44
Re Os
75
76
47
45
46
Ir
Pt Au Hg
Tl
77
78
81
79
48
31
80
32
33
34
Sn Sb Te
50
51
Pb Bi
82
83
52
Po At Rn
84
85
86
Rf Db Sg Bh Hs Mt
104
105
106
107
108
Metalloids
109
La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
57
58
59
Ac Th Pa
89
90
91
60
U
92
61
62
63
64
65
66
Np Pu Am Cm Bk Cf
93
94
95
96
97
98
67
68
69
70
71
Es Fm Md No Lr
99
100
101
102
103
Properties of Metals,
Nonmetals, and Metalloids
METALS
malleable, lustrous, ductile, good conductors of heat
and electricity
NONMETALS
gases or brittle solids at room temperature, poor
conductors of heat and electricity (insulators)
METALLOIDS (Semi-metals)
dull, brittle, semi-conductors (used in computer chips)
exhibit properties of both metals and nonmetals
Electrons filling the orbitals shown in
the Periodic Table
1
Periods
1
1s
8
Groups
2
3
4
5
2
2s
2p
3
3s
3p
4
4s
3d
4p
5
5s
4d
5p
6
6s
La
5d
6p
7
7s
Ac
6d
6
7 1s
4f
Lanthanide series
5f
Actinide series
Size of Atoms - Trends
Periodic Trends in Atomic Radii
Relative Size of Atoms
Shielding Effect
Valence
+
nucleus
Kernel electrons
block the attractive
force of the
nucleus from the
valence electrons
-
-
Electron
Shield
“kernel”
electrons
Electrons
Atomic Radius vs. Atomic Number
Radii vs atomic # up to 53
0.23
0.21
0.19
Radii (nonometers)
0.17
0.15
0.13
0.11
0.09
0.07
0.05
3
4
5
6
7
8
9
11
12
13
14
15
16
17
19
Atomic #
20
31
32
33
34
35
37
38
49
50
51
52
Atomic Radii trend explained
• As you go across the period the number of shielding electrons are the
same.
• The nuclear charge is increasing (adding protons as you go across).
• The electrons added are in the same valence shell – same distance
from the nucleus.
• More + nuclear charge gets out to the valence electrons, pulling the
valence electrons in closer (stronger attraction).
Ne
Li
As we go down a group each atom has another
energy level, so the atoms get bigger.
There are more levels in the kernel and therefore
greater shielding of valence electrons (weaker
attraction).
Li
Cs
Why do elements react?
• Atoms with filled valence shells are stable
– low in energy. Atoms attain a full valence
shell by losing, gaining or sharing valence
electrons. The result is a particle which is
isoelectronic with a noble gas (has the
same electron configuration as a noble
gas).
• Na 1+ =1s2 2s2 2p6 (isoelectronic Ne).
• F1- = 1s2 2s2 2p6
Atomic and Ionic Radii vs Atomic Number
Radii period trends explained
• Metals are “born losers”, the atoms lose their valence
electrons to form cations (+ ions). The kernel is smaller.
The remaining electrons are more strongly attracted to the
nucleus. There are more protons than electrons so the
remaining electrons are pulled in closer (+ ion is smaller).
• Nonmetals will gain valence electrons to fill the valence
shell to form anions (- ions). There are more electrons to
share the nuclear charge so there is a weaker attraction
between the electrons and the nucleus. The weaker
attraction leads to the valence electrons being further
away (- ion is larger)
N 3Li1+
B3+
Be2+
C4+
O2-
F1-
Radii group trends explained
Li+1
• As you go down the
group you are adding
energy levels so the
cations and anions
get bigger.
Na+1
K+1
Rb+1
Cs+1
First Ionization Energy Plot
First ionization energy trend explained
• 1st ionization energy is the energy required to
remove one electron from the gaseous atom of
an element.
• 1st ionization energy is increased by strong
attraction to valence electrons and a stable
electron configuration.
• As you go across the period the attraction to
valence electrons increases so ionization energy
increases.
• Peaks in energy occur at Be (full s sublevel), N
(½ full p sublevel) and Ne (full s & p sublevels –
full valence shell).
Electronegativity values
• Electronegativity is the tendency for an atom to attract electrons to
itself when it is chemically combined with another element.
• Big electronegativity values means the atom pulls the valence
electrons toward the nucleus, strong attraction to valence electrons.
• Atoms with small electronegativity values have weak attraction to
valence electrons and these electrons are drawn away from this
atom.
2.1
H
1.0
Li
1.5
Be
2.0
B
2.5
C
3.0
N
3.5
O
4.0
F
1.0
Na
1.2
Mg
1.5
Al
1.8
Si
2.1
P
2.5
S
3.0
Cl
0.9
K
1.0
Ca
1.3
Sc
1.4
Ti
1.5
V
1.6
Cr
1.6
Mn
1.7
Fe
1.7
Co
1.8
Ni
1.8
Cu
1.6
Zn
1.7
Ga
1.9
Ge
2.1
As
2.4
Se
2.8
Br
0.9
Rb
1.0
Sr
1.2
Y
1.3
Zr
1.5
Nb
1.6
Mo
1.7
Tc
1.8
Ru
1.8
Rh
1.8
Pd
1.6
Ag
1.6
Cd
1.6
In
1.8
Sn
1.9
Sb
2.1
Te
2.5
I
0.8
Cs
1.0
Ba
1.1
La
1.3
Hf
1.4
Ta
1.5
W
1.7
Re
1.9
Os
1.9
Ir
1.8
Pt
1.9
Au
1.7
Hg
1.6
Tl
1.7
Pb
1.8
Bi
1.9
Po
2.1
At
Electronegativity tend explained
• Group Trend: The further down a group the
farther the electron is away and the more
shielding electrons an atom has – the lower the
electronegativity value.
Period Trend: Metals on the left lose electrons
easily (weak attraction for valence electrons).
Metals have low electronegativity values.
• Nonmetals on the right need more electrons to
complete the valence shell and have strong
attraction for valence electrons. Nonmetals have
high electronegativity values.
• Electronegativity values increase as you go
across the period
Electron Affinity values
• Electron Affinity is the energy change associated with adding an electron to a
gaseous atom.
• If the atom becomes more stable (electron configuration like the noble gases)
there is a loss of energy and the energy change is shown a negative value.
• If the atom becomes less stable the energy level goes up and the energy change
is shown as a positive value.
H -73
/
Li -60
Be +240
\
B
-27
Na -53
Mg +230
/
Al -44
K -48
Ca +156
\
Rb -47
Sr +170
Cs -45
Ba +52
C
-122
N
+9
O -141
F -328
Si -134
P -72
S -200
Cl -348
Ga -30
Ge -120
As -77
Se -195
Br -325
/
In -30
Sn -121
Sb -101
Te -190
I -295
\
Tl -30
Pb -110
Bi -110
Po -183
At -270
Electron Affinity tend explained
• Across a period the electron affinity is low for
metals and high for nonmetals (electron affinity
increases from left to right). In Group 7A valence
electrons are strongly attracted to the nucleus and
an extra electron gives the atom a full valence
shell.
• Electron affinity decreases as we go down a group
because the atoms are getting bigger and the
valence electrons are not attracted as strongly to
the nucleus.
Nuclear charge increases
Shielding increases
Atomic radius increases
Ionic size increases
Ionization energy decreases
Electronegativity decreases
Electron affinity decreases
Summary of Periodic Trends
Nuclear charge increases
Shielding is constant
Atomic radius decreases
Ionization energy increases
Electronegativity increases
Electron affinity increases
1A
2A
0
3A 4A 5A 6A 7A
Ionic size (cations)
decreases
Ionic size (anions)
decreases