Periodic Table

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Periodic Table

• Larry Scheffler
• Lincoln High School
• IB Chemistry 1-2

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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?

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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 .

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Early Attempts to Classify Elements

• Dobreiners 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.

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Dobreiners Triads
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
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Newlands 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.

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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.

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Mendeleevs 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.

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Mendeleevs Periodic Table
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Mendeleevs 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 their properties.
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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

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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.

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Periodic Table Metallic Arrangement

• Layout of the Periodic Table Metals vs.
  nonmetals
• .

Nonmetals
Metals
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The Three Broad Classes are the Representative,
Transition, Rare Earth

• Main (Representative),
• Transition metals,
• lanthanides and actinides (rare earth)

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Additional Groupings in the Periodic Table

• Nonmetals, Metals, Metalloids, Noble gases

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Periodic Table The electron configurations are
inherent in the periodic table
H 1s1
He 1s2

• B
• 2p1

F 2p5
Be 2s2
B 2p1
C 2p2
N 2p3
Ne 2p6
Li 2s1
O 2p4
Na 3s1
Mg 3s2
Cl 3p5
Al 3p1
Si 3p2
P 3p3
S 3p4
Ar 3p6
K 4s1
Ca 4s2
Zn 3d10
As 4p3
Be 4p5
Sc 3d1
Ti 3d2
V 3d3
Cr 4s13d5
Mn 3d5
Fe 3d6
Co 3d7
Ga 4p1
Ge 4p2
Se 4p4
Kr 4p6
Ni 3d8
Cu 4s13d10
Sr 5s2
Rb 5s1
Nb 4d3
Mo 5s14d5
Ru 4d6
Rh 4d7
Sn 5p2
I 5p5
Xe 5p6
Cd 4d10
Zr 4d2
Tc 4d5
In 5p1
Sb 5p3
Te 5p4
Y 4d1
Ni 4d8
Ag 5s14d10
Cs 6s1
Hf 5d2
Ta 5d3
W 6s15d5
Re 5d5
Os 5d6
Ir 5d7
At 6p5
Rn 6p6
La 5d1
Ni 5d8
Ba 6s2
Tl 6p1
Pb 6p2
Bi 6p3
Po 6p4
Hg 5d10
Au 6s15d10
Mt 6d7
Fr 7s1
Bh 6d5
Hs 6d6
Ra 7s2
Rf 6d2
Db 6d3
Sg 7s16d5
Ac 6d1
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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
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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

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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
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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.

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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

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Ionization Energies of the Representative Groups

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Ionization Energies are Periodic
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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.
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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.

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Trends in Electronegativity
Electronegativity increases across a period and
up a group
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Electronegativity
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Electronegativity
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Electron Affinities
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Electron Affinities Are Periodic

• Electron Affinity v Atomic Number

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The Electron Shielding Effect

• Electrons between the nucleus and the valence
  electrons repel each other making the atom larger.

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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.

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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.

All values are in nanometers
Large
Small
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Atomic Radius

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Atomic Radius
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Trends in Ion Sizes
Radius in pm
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Cations

• Cations (positive ions) are smaller than their
  corresponding atoms

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Ionic Radius
Forming a cation.
Li 0.152 nm
3e and 3p

• CATIONS are SMALLER than the atoms from which
  they come.
• The electron/proton attraction has gone UP and so
  the radius DECREASES.

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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.
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Anions

• Anions (negative ions) are larger than their
  corresponding atoms

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Ionic Radius-Anions
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?

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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.
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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.
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Summary of Periodic Trends
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Properties of the Third Period Oxides
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Properties of the Third Period Chlorides
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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.

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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
• Multiple oxidation states
• The ability to form complex ions
• Colored compounds
• Catalytic behavior
• Magnetic properties

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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

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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 Sc2 does not exist, and
• Ti2 is unstable, oxidizing
• in the presence of any
• water to the 4 state.

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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.

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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

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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.

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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)
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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.

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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

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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

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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

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Some Common D Block Catalysts

• Examples of D block elements that are used as
  catalysts

• Platnium or
• rhodium is used in a
• catalytic converter
• MnO2 catalyzes the decomposition
• of hydrogen peroxide
• V2O5 is a catalyst for the contact process
• Fe in Haber process
• Ni in conversion of
• alkenes to alkanes

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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
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Alternate Periodic Tables II
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Alternate Periodic Tables III
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Alternate Periodic Tables IV