Periodic Table - PowerPoint PPT Presentation

1 / 64
About This Presentation

Periodic Table


The development of the periodic table brought a system of order to what was ... The Periodic Table can be arranged by energy sub levels The s-block is Group IA ... – PowerPoint PPT presentation

Number of Views:151
Avg rating:3.0/5.0
Slides: 65
Provided by: lincolnP
Tags: periodic | table


Transcript and Presenter's Notes

Title: Periodic Table

Periodic Table
  • Larry Scheffler
  • Lincoln High School

The Periodic Table-Key Questions
  • What is the periodic table ?
  • What information is obtained from the table ?
  • How can elemental properties be predicted based
    on the Periodic Table?

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 yet
  • elements.

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.

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
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
  • In the end his work was acknowledged and he was
    vindicated with the award of the Davy Medal in
    1887 for his work.

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 made predictions based
    on his arrangement which were shown to be quite

Mendeleevs Periodic Table
  • Mendeleev organized all of the elements into one
    comprehensive table.
  • Elements were arranged in order of increasing
  • Elements with similar properties were placed in
    the same row.

Mendeleevs Periodic Table
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
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

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

Periodic Table Metallic Arrangement
  • Layout of the Periodic Table Metals vs.

The Three Broad Classes Are Main, Transition,
Rare Earth
  • Main (Representative), Transition metals,
  • lanthanides and actinides (rare earth)

Reading the Periodic Table Classification
  • Nonmetals, Metals, Metalloids, Noble gases

(No Transcript)
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
Periodic Table Organization------ Groups or
Vertical columns in the periodic table are known
as groups or families The elements in a group
have similar electron configurations
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

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
Trends in Ionization Energy
Ionization energy is the energy required
to Remove an electron from an atom
  • Ionization energy increases across a period
    because the positive charge increases.
  • Metals lose electrons more easily than nonmetals.
  • Nonmetals lose electrons with difficulty (they
    like to GAIN electrons).

Trends in Ionization Energy
  • The ionization energy increases UP a group
  • Because size increases due to an effect known as
    the Shielding Effect

Ionization Energies

Ionization Energies are Periodic
  • Electronegativity is a measure of the ability
    of an atom in a molecule to attract electrons to

This concept was first proposed by Linus Pauling
(1901-1994). He later won the Nobel Prize for
his efforts
Periodic Trends Electronegativity
  • In a group Atoms with fewer energy levels can
    attract electrons better (less shielding). So,
    electronegativity increases UP a group of
  • 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.

Trends in Electronegativity
Electron Affinities
Electron Affinities Are Periodic
  • Electron Affinity v Atomic Number

The Electron Shielding Effect
  • Electrons between the nucleus and the valence
    electrons repel each other making the atom larger.

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
  • The radius decreases on going across a period.

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, where the electrons are scattered.

Atomic Radius

Atomic Radius
Trends in Ion Sizes
Radius in pm
  • Cations (positive ions) are smaller than their
    corresponding atoms

Ion Sizes
  • Does the size go up or down when gaining an
    electron to form an anion?

Ionic Radius
Forming a cation.
Li,152 pm
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.

Ionic Radius for Cations
Positve ions or cations are smaller than the
corresponding atoms. Cations like atoms increase
as one moves from top to bottom in a group.
  • Anions (negative ions) are larger than their
    corresponding atoms

Ionic Radius-Anions
Forming an anion.
  • ANIONS are LARGER than the atoms from which they
  • The electron/proton attraction has gone DOWN and
    so size INCREASES.
  • Trends in ion sizes are the same as atom sizes.

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

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

The D Block Elements
  • The d electrons are close in energy to the s
  • 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

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
  • 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.

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
  • 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.

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.

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
  • The colors of the ions and complex ions of d
    block elements depends on a variety of factors
  • The particular element
  • The oxidation state
  • The kind of ligands bound to the element

Various oxidation states of Nickel (II)
Colors in the D Block
  • The presence of a partially filled d sublevels in
    a transition element results in colored
  • 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 is colored, if it absorbs
    light in the visible range (400-700
  • nanometers).
  • If the compound absorbs a
  • particular wavelength of light its
  • color will be the composite of those
  • wavelengths that it does not absorb.
  • In other words it shows its
  • complimentary color.

Colors and d Electron Transitions
  • When ligands are attached to transition metal
    ions, the d orbitals may split into two groups.
    Some of the orbitals are at a lower energy than
    the others
  • The difference in energy of these orbitals varies
    slightly with the nature of the ligand or ion
    surrounding the metal ion
  • The energy of the transition ?E hn may occur
    in the visible region.
  • 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.
  • The result is a colored compound

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
  • Transition metal complexes with unpaired
    electrons exhibit simple paramagnetism.
  • The degree of paramagnetism depends on the
    number of unpaired electrons

Catalytic Behavior
  • Many D block elements are catalysts for various
    chemical reactions
  • Catalysts speed up the rate of a reaction with
    out being consumed.
  • The transition metals form complex ions with
    ligands 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 oxidation- reduction reactions

Some Common D Block Catalysts
  • Examples of D block elements that are used as
  • Platnium or
  • rhodium in a
  • catalytic converter
  • MnO2 decomposition
  • of hydrogen peroxide
  • V2O5 in the contact
  • process
  • Fe in Haber process
  • Ni in conversion of
  • alkenes to alkanes

The Periodic Table--Summary
The periodic table is a classification system.
Although we are most familiar with the periodic
table that Seaborg proposed more than 60 years
ago, several alternate designs have been proposed.
Alternate Periodic Tables
Alternate Periodic Tables II
Write a Comment
User Comments (0)