Transcript Periodic Table Larry Scheffler Lincoln High School IB Chemistry 1-2

Periodic Table
Larry Scheffler
Lincoln High School
IB Chemistry 1-2
.1
The Periodic Table-Key Questions
What is the periodic table ?
What information does the table provide ?
How can one use the periodic table to predict
the properties of the elements?
.2
Periodic Table
• The development of the periodic table brought a
system of order to what was otherwise an
collection of thousands of pieces of
information.
• The periodic table is a milestone in the
development of modern chemistry. It not only
brought order to the elements but it also
enabled scientists.
to predict the existence
of elements that had
not yet been discovered .
.3
Early Attempts to Classify
Elements
Dobreiner’s Triads (1827)
• Classified elements in sets of three
having similar properties.
• Found that the properties of the middle
element were approximately an average
of the other two elements in the triad.
.4
Dobreiner’s Triads
Element Atomic
Mass
Average Density
Average
(g cm-3)
(amu)
Cl
Br
I
35.5
79.9
126.9
Ca
Sr
Ba
40.1
87.6
137.3
81.2
1.56
3.12
4.95
3.25
88.7
1.55
2.6
3.5
2.53
Note: In each case, the numerical values for the
atomic mass and density of the middle element are
close to the averages of the other two elements
.5
Newland’s Octaves -1863
John Newland attempted to
classify the then 62 known
elements of his day.
He observed that when classified
according to atomic mass, similar
properties appeared to repeat for
about every eighth element
His Attempt to correlate the
properties of elements with musical
scales subjected him to ridicule.
In the end his work was
acknowledged and he was
vindicated with the award of the
Davy Medal in 1887 for his work.
.6
Dmitri Mendeleev
Dmitri Mendeleev is
credited with creating
the modern periodic
table of the elements.
He gets the credit
because he not only
arranged the atoms,
but he also made
predictions based on
his arrangements His
predictions were later
shown to be quite
accurate.
.7
Mendeleev’s Periodic Table
• Mendeleev organized all of the elements into
one comprehensive table.
• Elements were arranged in order of
increasing mass.
• Elements with similar properties were placed
in the same row.
.8
Mendeleev’s Periodic Table
.9
Mendeleev’s Periodic Table
Mendeleev left some blank spaces in his periodic table.
At the time the elements gallium and germanium were
not known. He predicted their discovery and estimated
.10
their properties.
The Modern Periodic Table
The Periodic Table has undergone several modifications
before it evolved in its present form. The current form
is usually attributed to Glenn Seaborg in 1945
.11
Periodic Table Expanded View
The Periodic Table can be
arranged by energy sub levels
The s-block is Group IA and & IIA,
the p-block is Group IIIA - VIIIA.
The d-block is the transition
metals, and the f-block are the
Lanthanides and Actinide metals
The way the periodic table
usually shown is a compressed
view. The Lanthanides and
actinides (F block) are cut out and
placed at the bottom of the table.
.12
Periodic Table: Metallic
Arrangement
Layout of the Periodic Table: Metals vs.
nonmetals
1
IA
.
1
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
18
VIIIA
17
VIIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
Nonmetals
4
5
6
Metals
7
.13
The Three Broad Classes are the
Representative, Transition, & Rare Earth
Main (Representative),
Transition metals,
lanthanides and actinides (rare earth)
.14
Additional Groupings in the
Periodic Table
Nonmetals, Metals, Metalloids, Noble gases
.15
Periodic Table:
The electron
configurations are inherent in the
periodic table
1
IA
1
2
3
4
5
6
7
H
1s1
18
VIIIA
2
IIA
Li Be
2s1 2s2
Na Mg
3s1 3s2
13
IIIA
B
2p1
3
IIIB
4
IVB
Sc
3d1
Rb
5s1
Ca
4s2
Sr
5s2
Y
4d1
V
Ti
Cr Mn Fe Co
3d2 3d3 4s13d5 3d5 3d6 3d7
Zr Nb Mo Tc Ru Rh
4d2 4d3 5s14d5 4d5 4d6 4d7
Cs
6s1
Ba
6s2
La
5d1
Hf Ta W Re Os
5d2 5d3 6s15d5 5d5 5d6
Fr
7s1
Ra
7s2
Ac Rf
6d1 6d2
K
4s1
5
VB
6
VIB
7
VIIB
Db Sg Bh
6d3 7s16d5 6d5
8
9
VIIIB
14
IVA
15
VA
16
VIA
17
VIIA
B
1
2p
C
N
O
2
3
2p 2p 2p4
F
2p5
Ne
2p6
Cl
3p5
Ar
3p6
He
1s2
10
11
IB
12
IIB
Ni
3d8
Cu
4s13d10
Ni
4d8
5s14d10
Zn Ga Ge
3d10 4p1 4p2
Cd
In Sn
10
4d
5p1 5p2
As Se Be
4p3 4p4 4p5
I
Sb Te
5p3 5p4 5p5
Kr
4p6
Xe
5p6
Hg
Tl Pb
5d10 6p1 6p2
Bi Po At
6p3 6p4 6p5
Rn
6p6
Ir
Ni
7
5d 5d8
Hs Mt
6d6 6d7
Ag
Au
6s15d10
Al Si
3p1 3p2
S
P
3
3p 3p4
.16
Periodic Table Organization-----Groups or Families
Vertical columns in the periodic table are known as groups or
families The elements in a group have similar electron
configurations
.17
Periodic Table Organization
---- Periods
Horizontal Rows in the periodic table are known as Periods The
Elements in a period undergo a gradual change in properties as
one proceeds from left to right
.18
Periodic Properties
Elements show gradual changes in certain
physical properties as one moves across a
period or down a group in the periodic table.
These properties repeat after certain
intervals. In other words they are PERIODIC
Periodic properties
include:
-- Ionization Energy
-- Electronegativity
-- Electron Affinity
-- Atomic Radius
-- Ionic Radius
.19
Trends in Ionization Energy
Ionization energy is the energy required to
remove an electron from an atom
Metals lose electrons
more easily than
nonmetals.
Nonmetals lose electrons
with difficulty. (They like
to GAIN electrons).
Ionization energy
increases across a period
because the positive
charge increases.
.20
Trends in Ionization Energy
The ionization energy is
highest at the top of a
group. Ionization energy
decreases as the atom
size increases.
This results from an
effect known as the
Shielding Effect
.21
Ionization Energies of the
Representative Groups
.22
Ionization Energies are Periodic
.23
Electronegativity
Electronegativity
is a measure of
the ability of an
atom in a
molecule to
attract electrons
to itself.
This concept was first proposed by Linus
Pauling (1901-1994). He later won the Nobel
Prize for his efforts.
.24
Periodic Trends:
Electronegativity
In a group: Atoms with fewer
energy levels can attract
electrons better (less shielding).
So, electronegativity increases
UP a group of elements.
In a period: More protons, while
the energy levels are the same,
means atoms can better attract
electrons. So, electronegativity
increases RIGHT in a period of
elements.
.25
Trends in Electronegativity
Electronegativity increases across
a period and up a group
.26
Electronegativity
.27
Electronegativity
.28
Electron Affinities
.29
Electron Affinities Are Periodic
Electron Affinity v Atomic Number
.30
The Electron Shielding Effect
Electrons
between the
nucleus and
the valence
electrons repel
each other
making the
atom larger.
.31
Atomic
Radius
The radius increases on going down a group.
Because electrons are added further from the
nucleus, there is less attraction. This is due to
additional energy levels and the shielding
effect. Each additional energy level “shields”
the electrons from being pulled in toward the
nucleus.
The radius decreases on going across a
period.
.32
Atomic Radius
The radius decreases across a period owing
to increase in the positive charge from the
protons.
Each added electron feels a greater and
greater + charge because the protons are
pulling in the same direction, whereas the
electrons are scattered.
Large
All values are in
nanometers
Small
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Atomic Radius
.34
Atomic Radius
.35
Trends in Ion Sizes
Radius in pm
.36
Cations
Cations (positive ions) are smaller
than their corresponding atoms
.37
Ionic Radius
+
Li
0.152 nm
3e and 3p
Li + ,
0.078 nm
2e and 3 p
Forming a
cation.
CATIONS are SMALLER than the atoms from
which they come.
The electron/proton attraction has gone UP
and so the radius DECREASES.
.38
Ionic Radius for Cations
Positive ions or
cations are
smaller than the
corresponding
atoms.
Cations like
atoms increase
as one moves
from top to
bottom in a
group.
.39
Anions
Anions (negative ions) are larger
than their corresponding atoms
.40
Ionic Radius-Anions
F 0.064 nm
9e- and 9p+
F- 0.133 nm
10 e- and 9 p+
Forming an anion.
ANIONS are LARGER than the atoms
from which they come.
The electron/proton attraction has gone
DOWN and so size INCREASES.
Trends in ion sizes are the same as atom
sizes.
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Ion Sizes
Does the size go up or
down when gaining an
electron to form an
anion?
.42
Ionic Radii for Anions
Negative ions or
anions are larger
than the
corresponding
atoms.
Anions like atoms
increase as one
moves from top
to bottom in a
group.
.43
Ionic Radius for an
Isoelectronic Group
Isoelectronic ions
have the same
number of electrons.
The more negative
an ion is the larger it
is and vice versa.
.44
Summary of Periodic Trends
Properties of the Third Period
Oxides
.46
Properties of the Third Period
Chlorides
.47
The D Block Elements
The d block elements
fall between the s
block and the p block.
They share common
characteristics since
the orbitals of d
sublevel of the atom
are being filled.
.48
The D Block Elements
The D block elements include the transition
metals. The transition metals are those d block
elements with a partially filled d sublevel in one
of its oxidation states.
Since the s and d sublevels are very close in
energy, the d block elements show certain
special characteristics including:
1. Multiple oxidation states
2. The ability to form complex ions
3. Colored compounds
4. Catalytic behavior
5. Magnetic properties
.49
The D Block Elements
The d electrons are close in energy to the s
electrons.
D block elements may lose 1 or more d
electrons as well as s electrons. Hence they
often have multiple oxidation states
Some common D block oxidation states
.50
Multiple Oxidation States
There is no sudden sharp increase in ionization
energy as one proceed through the d electrons as
there would be with the s block.
D block elements can lose or share d electrons
as well as s electrons, allowing for multiple
oxidation states.
Most d Block elements have a +2 oxidation State
which corresponds to the loss of the two s
electrons.
This is especially true on the right side of the d
block, but less true on the left.
---- For example Sc+2 does not exist, and
Ti+2 is unstable, oxidizing
in the presence of any
water to the +4 state.
.51
Complex Ions
The ions of the d block and the lower p block
have unfilled d or p orbitals.
These orbitals can accept electrons either an
ion or polar molecule, to form a dative bond.
This attraction results in the formation of a
complex ion.
A complex ion is made up of two or more ions
or polar molecules joined together.
The molecules or ions that surround the metal
ion donating the electrons to form the complex
ion are called ligands.
.52
Complex Ions
Compounds that are formed with
complex ions are called coordination
compounds
Common ligands
Complex ions usually have either 4 or 6
ligands.
K3Fe(CN)6
Cu(NH3)42+
Complex Ions
The formation of complex ions stabilizes
the oxidations states of the metal ion
and they also affect the solubility of the
complex ion.
The formation of a
complex ion often has
a major effect on the
color of the solution of
a metal ion.
.54
The D Block Colored Compounds
In an isolated atom all of the d sublevel electrons
have the same energy.
When an atom is surrounded by charged ions or polar
molecules, the electric field from these ions or
molecules has a unequal effect on the energies of the
various d orbitals and d electrons.
The colors of the ions and complex ions of d block
elements depends on a variety of factors including:
– The particular element
– The oxidation state
– The kind of ligands bound to the element
Various oxidation
states of Nickel (II)
.55
Colors in the D Block
The presence of a partially filled d sublevels in a
transition element results in colored compounds.
Elements with completely full or completely empty
subshells are colorless,
– For example Zinc which has a full d subshell. Its
compounds are white
A transition metal ion exhibits color, if it absorbs
light in the visible range (400-700
nanometers)
If the compound absorbs a
particular wavelengths of light its
color is the composite of those
wavelengths that it does not absorb.
It shows the complimentary color.
.56
Colors and d Electron Transitions
The d orbitals may split into two groups so that two
orbitals are at a lower energy than the other three
The difference in energy of these orbitals varies slightly
with the nature of the ligand or ion surrounding the
metal ion
When white light passes through a compound of a
transition metal, light of a particular frequency is
absorbed as an electron is promoted from a lower
energy d orbital to a higher one.
When the energy of the transition: ∆E =hn may occur in
the visible region, the compound is colored
.57
Magnetic Properties
Paramagnetism --- Molecules with
one or more unpaired electrons are
attracted to a magnetic field. The
more unpaired electrons in the
molecule the stronger the attraction.
This type of behavior is called
Diamagnetism --- Substances with
no unpaired electrons are weakly
repelled by a magnetic field.
Transition metal complexes with
unpaired electrons exhibit simple
paramagnetism.
The degree of paramagnetism
depends on the number of unpaired
electrons
.58
Catalytic Behavior
Many D block elements are
catalysts for various reactions
Catalysts speed up the rate of a
chemical reaction with out being
consumed.
The transition metals form complex
ions with species that can donate
lone pairs of electrons.
This results in close contact
between the metal ion and the
ligand.
Transition metals also have a wide
variety of oxidation states so they
gain and lose electrons in redox
reactions
.59
Some Common D Block Catalysts
Examples of D block elements that are
used as catalysts:
1. Platnium or
rhodium is used in a
catalytic converter
2. MnO2 catalyzes the
decomposition
of hydrogen peroxide
3. V2O5 is a catalyst for
the contact process
4. Fe in Haber process
5. Ni in conversion of
alkenes to alkanes
.60
Alternate Periodic Tables
Although we are most familiar
with the periodic table that
Seaborg proposed more than
60 years ago, several alternate
designs have been proposed.
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Alternate Periodic Tables
.62
Alternate Periodic Tables II
.63
Alternate Periodic Tables III
.64
Alternate Periodic Tables IV