Download Periodic Table - Derry Area School District

Periodic Table
Chapter 6
What do I know?
• On the back of the blank periodic table
write down at least 3 pieces of
information you can get from the
periodic table.
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A Brief History…
• Joseph Proust
• Law of Definite Composition
– elements combine in definite proportions
by weight
• The weight of one element that
combines with the weight of another
element = combining weight
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• Joseph Berzelius
• [1807 - 1818 ]
• Determined the combining weights of 43
elements with oxygen.
• Recognized similarities of certain elements...
– similar metallic properties
– similar reactive
properties
– Li, Na, K
– similar nonmetals
– Cl, Br, I,
• “TRIPLETS”
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• Johann Wolfgang Dobereiner
• 1829
• mathematician
• discovered that combining weight of
middle triplet is the average [or near
average]of the combining weights of
the other two
• Li, Na, K
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• Jean Stas
• 1860
• confirmed Proust theory of definite
composition
• established accurate atomic weight of
the known elements
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Was there a relationship
between the weight of an
element and its properties?
• John A.R. Newlands
• 1865
• arranged elements in order of atomic
weight
• elements with similar properties were 7,
or multiple of 7 apart
• Law of Octaves
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Dimitri Mendeleev 1869
• developed a chart -listed elements by
increasing atomic weight
• grouped elements with similar
properties in the same row
• Left Gaps where no element fit the
pattern.
• Predicted discovery of new elements
• Predicted properties of new elements
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Mendeleev’s Table
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Old Periodic Law
.
“The properties
of elements are
in periodic
dependence of
their atomic
weights.”
Dimitri Mendeleev
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ALTERATIONS and
ADDITIONS
• Sir Wm.Ramsay
• 1890’s
• Discovered Ne, Ar, Kr, Xe
• Helium and Radon disc. Previously
• New row added to Periodic Table
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•
•
•
Henry Gwyn-Jeffreys Mosley
1914-1915
Number of protons determined
– atomic number - identifies what an
element is
•
Periodic Table Rearranged
– elements arranged by increasing atomic
number
– similar elements put in columns instead
of rows
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Modern Periodic Law
• “The properties of elements are in
periodic dependence of their atomic
numbers.”
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ARRANGEMENT OF THE MODERN
PERIODIC TABLE
• A horizonal row on the periodic chart is
refered to as either a period, or a
series.
• A vertical column on the periodic chart
is refered to as either a group, or a
family.
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Element
• H
• Li
Location
on Chart
Grp 1
Grp 1
• Na
Grp 1
Na .
• K
Grp 1
K.
• Rb
Grp 1
Rb .
• Cs
Grp 1
Cs .
• Fr
•
Grp 1
Fr .
Page 16
Electron Dot
Notation
H.
Li .
Alkali Metals
Properties
Colorless gas
Soft; silver
highly reactive
Soft; silver
highly reactive
Soft; silver
highly reactive
Soft; silver
highly reactive
Soft; silver
highly reactive
Soft; silver
most reactive metal
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Element
• Be
• Mg
• Ca
• Sr
• Ba
• Ra
Page 17
Location
on Chart
Grp 2
Electron Dot
Notationc
Be :
Properties
Reactive metal
Grp 2
Mg :
Reactive metal
Grp 2
Ca :
Reactive metal
Grp 2
Sr :
Reactive metal
Grp 2
Ba :
Reactive metal
Grp 2
Ra :
Most reactive
metal of group
Alkaline Earth Metals
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What pattern(s) do we see?
• All elements in groups have same
electron dot structure.
• Group placement predicts valence.
• Groups usually have similar properties.
• Most reactive metals at the bottom of
the group.
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Element
• B
• Al
• Ga
• In
• Tl
Page 19
Location
on Chart
Electron Dot
Notation
Properties
Grp 3
B:
nonmetal; black
solid
Grp 3
Al:
Metal
Grp 3
Ga :
Metal
Grp 3
In :
Metal
Grp 3
Tl :
Most reactive metal
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Element
Location
on Chart
Electron Dot
Notation
Properties
• C
Grp 4
C:
black→clear
solid
Grp 4
Si :
Metalloid
Grp 4
Ge :
Metal
Grp 4
Sn :
Metal
Grp 4
Pb :
Most reactive metal
• Si
• Ge
• Sn
• Pb
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Element
•
•
•
•
•
•
•
•
•
•
Location
on Chart
Electron Dot
Notation
Properties
N
Grp 5
N:
gas; nonmetal
Grp 5
P:
nonmetal
Grp 5
As :
Metalloid
Grp 5
Sb :
Metalloid
Grp 5
Bi :
Metal
P
As
Sb
Bi
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Element
• O
•
•
•
•
•
•
•
•
•
S
Se
Te
Po
Page 22
Location
on Chart
Electron Dot
Notation
Properties
Grp 6
O:
gas; nonmetal
reactive
Grp 6
S:
Nonmetal
Grp 6
Se :
Nonmetal
Grp 6
Te :
Nonmetal
Grp 6
Po :
Metal
Chalcogen Family
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Element
•
•
F
•
•
Cl
•
•
Br
Location
on Chart
Grp 7
•
•
I
•
•
At
Electron Dot
Notation
:F:
gas; most reactive
nonmetal
Grp 7
:Cl :
gas; reactive
nonmetal
Grp 7
:Br :
liquid; reactive
Grp 7
:I:
solid; reactive
nonmetal
Grp 7
:At :
solid; reactive
nonmetal
Halogen Family
Page 23
Properties
nonmetal
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Element
•
•
•
•
•
•
•
•
•
•
•
He
Ne
Ar
Kr
Xe
Rn
Page 24
Location
on Chart
Electron Dot
Notation
Properties
Grp 8
He :
inert; nonmetal
Grp8
:Ne :
inert; nonmetal
Grp 8
:Ar :
inert; nonmetal
Grp 8
:Kr :
inert; nonmetal
Grp 8
:Xe :
inert; nonmetal
Grp 8
:Rn :
inert; nonmetal
Noble Gases / Inerts
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What pattern(s) do we see?
• All elements in groups have same
•
•
•
•
electron dot structure.
Group placement predicts valence.
Groups usually have similar properties –
(exception: steps)
Most reactive nonmetals at the top of
the group.
Most reactive metals at the bottom of
the group.
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I spy with my little eye an element with…
• 3 energy levels and 2 valence electrons
• Mg
• 5 energy levels and 4 valence electrons
• Sn
• 2 energy levels and 8 valence electrons
• Ne
• 1 valence electron and 5 energy levels
• Rb
• 1 valence electron and 7 energy levels
• Fr
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I spy with my little eye an element with…
• 4 energy levels and 7 valence electrons
• Br
• 3 energy levels and 5 valence electrons
• P
• 2 valence electrons and 4 energy levels
• Ca
• 3 valence electrons and 2 energy levels
• B
• 8 valence electrons and 5 energy levels
• Xe
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I spy with my little eye an element with…
• The heaviest halogen…
• At (astatine)
• The triplet with the average atomic weight of 35.5…
• Cl
• The least reactive Chalcogen
• Po (polonium)
• The group that fills the s2 valence orbital
• Alkaline Earth Metals
• A third period metalloid
• Si
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Bonding
• See interactive
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Types of Bonding
• Ionic
– Electrons transfer from one atom to
another creating + and – ions.
• Covalent
– Atoms share electrons to create a
molecule.
• Metallic
– Many atoms share electrons
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Types of Bonding
• Ionic
– Electrons transfer from one atom to another
– creating + and – ions.
e+
Page 31
-
+ energy
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Ionization Energy
• The energy required to remove the outermost
e- in an atom.
Ionization Energy
Ionization Energy (kJ/moL)
Helium
2500
Neon
2000
Argon
1500
Hydrogen
Series1
1000
500
0
Lithium
Sodium
0
10
20
30
40
Atomic Number
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• Why are some e- removed more easily?
– Electrons that are farther away from the nucleus and that
have more E levels between them and the nucleus
• Low ionization energy
– characteristic of METALS.
• High ionization energy
– characteristic of NONMETALS.
• Removing successive electrons is more difficult, but
follows the same overall pattern.
•
Na + Energy  Na+ + e•
Na+ + Energy  Na++ + e•
Na+++ Energy  Na++++ e-
Page 34
119 Kcal / mol
1090 Kcal/ mol
1652 Kcal/ mol
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Electron Affinity
The energy released / absorbed when an
electron is accepted by a neutral atom
e+
Ionization E removes eand forms + ion
Page 35
linked
-
+ energy
Electron affinity is the E
released when the neutral
atom accepts the freed eand becomes -
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Electron Affinity
Electron Affinity (kJ/moL)
Electron Affinity
100
Lithium
0
-100
0
Sodium
10
20
30
40
Series1
-200
-300
-400
Fluorine
Chlorine
Atomic Number
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Electron Affinity
Increases across a period
D
e
c
r
e
a
s
e
s
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• For atoms that have - valences:
Atom + e-  A- + E
exothermic - energy released
(electron affinity)
stable product
Atom + e- + E  Aendothermic - energy required
unstable product
•
•
•
•
•
•
•
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• Covalent
– Atoms share electrons to create a
molecule.
-’
Shared e s
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Electronegativity
• the attraction of an atom for a shared
pair of electrons
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Electronegativity
Electronegativity
Electronegativity
4.5
4
3.5
3
2.5
2
1.5
1
0.5
0
Fluorine
Chlorine
Series1
Lithium
0
Sodium
10
20
30
40
Atomic Number
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Table of Electronegativities
Page 42
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Electronegativity
• Types of Covalent Bonds:
• pure covalent - relatively even sharing of e• polar covalent - uneven sharing of e•
0 - .5 ....... pure covalent
.5 - 1.7..... polar covalent
> 1.7 ....... ionic bond
•
•
•
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Atomic Radius [size]
Atomic Radius
Atomic Radius (pm)
250
200
Lithium
150
Sodium
100
Series1
Chlorine
Fluorine
50
0
0
10
20
30
40
Atomic Number
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Atomic Radius [size]
Page 45
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•
Down a group
–
•
Across a period
–
•
•
E levels are added.
Increased attraction between the E levels and
the nucleus causes the size to decrease.
Pauli Repulsion Theory
As the number of electrons increases so
does the repulsion between the electrons;
this may help account for the irregular
increase in the radii.
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Ions [size]
• Increases down a Group
• Decreases across a Period
• Metal atoms lose electrons
– become positive (cation)
– Cations are SMALLER than the atoms from
which they come.
• Nonmetal atoms gain electrons
– become negative (anion)
– Anions are LARGER than the atoms from
which they come.
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Density
Density (g/cm^3)
Density
10
9
8
7
6
5
4
3
2
1
0
-1 0
Boron
Lithium
Series1
Aluminum
Sodium
Chlorine
10
Fluorine20
30
40
Atomic Number
Page 48
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Density
•
Here the
density of each
period is
graphed
individually
Page 49
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Page 50
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Density
• Generally, density
– increases down a group
– Increases across the metals in a period,
and then decreases across the nonmetals
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M.P and B.P.
Melting Point/Boiling Point
3500
3000
Temperature °C
2500
2000
Melting Point
1500
Boron
1000
Boiling Point
Aluminum
500
0
0
Lithium
-500
5
10
Neon
15
Sodium
20
Argon25
30
35
40
Atomic Number
Page 52
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M.P. and B.P.
• Generally, like density, M.P. and B.P.
– increases down a group
– Increases across the metals in a period,
and then decreases across the nonmetals
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Configurations and the
Periodic Table
• Electrons that reside in the
outermost shell of an atom are called
valence electrons.
– These electrons are primarily involved in chemical
reactions.
– Elements within a given group have the same
“valence shell configuration.”
– This accounts for the similarity of the chemical
properties among groups of elements.
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Configurations and the
Periodic Table
• The following slide illustrates how the
periodic table provides a sound way to
remember the Aufbau sequence.
– In many cases you need only the configuration of
the outer electrons.
– You can determine this from their position on the
periodic table.
– The total number of valence electrons for an
atom equals its group number.
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Configurations and the
Periodic Table
s2
p6
s1
1
s2
p1
p2
p3
p4
p5
2
3
S
4
filling
d1
d2
d3
d4
d5
d6
d7
d8
d9
d10
3d
P
filling
4d
5
6
*
7
**
5d
6d
f filling
Page 56
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The Elements
• The Elements in Song
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