The Periodic Table - PowerPoint 4 by Ax2YAZF


									The Periodic Table
     The Periodic Table
is a chart that organizes all the
elements in a way that shows patterns –
both similarities and differences among
the elements.
Elements are listed by increasing
atomic number.
The horizontal rows are called periods.
There are 7 periods.
The vertical columns are called groups.
There are 18 groups.
Elements in the same group have
similar chemical properties.
               Group       18
Period 1       2

Period 2
Period 3           3 4 5

Period 7
         Special Groups
Group 1 elements are known as the alkali
Group 2 elements are known as the alkaline
earth metals.
Group 17 elements are called halogens.
Group 18 elements are called noble gases.
Groups 3-12 are called transition metals.
The Lanthanides and Actinides are the
excerpts at the bottom.
There are 3 major sections of elements
in the periodic table:

 1.   Metals
 2.   Non-metals
 3.   Metalloids

Metals are located to the left of the
“staircase line” and non-metals are to
the right (except H). Metalloids border
the line.
     Properties of Metals
solids at room temperature (except Hg)
give up (lose) electrons
good conductors of electricity
  Properties of Non-metals
solids, liquids, and gases
gain electrons
not shiny
not bendable
poor conductors of electricity
acidic in water
  Properties of Metalloids
Metalloids border the “staircase” and
tend to have some properties of both
metals and non-metals.
can lose or gain electrons
semi-conductors of electricity
Within a period (row), elements
gradually change from metallic to non-
metallic as you go from left to right
across the table.
Periodic Trends
        Periodic Trends
As Mentioned before, the Periodic Table
is a chart that organizes all the
elements in a way that shows patterns –
both similarities and differences among
the elements.
Trends can be predicted using the
periodic table and explained by electron
configurations of the elements.
        Periodic Trends
Elements will gain, lose, or share
valence electrons (electrons on the
outermost shell) in order to achieve
stable octet formation.
Examples of natural stable octets are
the inert gases also known as the noble
gases located in group 18
    1 of 2 Important Trends
Electrons are added one at a time moving
from left to right across a period.
   For example
     N has 7 electrons, O has 8 electrons, and F has 9
       electrons (balanced atoms)
As a result electrons of the outermost shell
have an increasingly strong attraction to the
nucleus, causing the electrons to become
closer to the nucleus and more tightly bound
to it.
    2 of 2 Important Trends
moving down a column in the periodic
table, the outermost electrons become
less tightly bound to the nucleus.
 This happens because the number of filled
  energy levels increases downward within
  each group.
 Each shell is like a shield to the outermost
  electrons, blocking some of the attraction
  to the nucleus.
Periodicity: Commonly Referred
     to as Periodic Trends
These 2 trends explain the periodicity
observed in the elemental properties of:
   atomic radius
  ionization energy

  electron affinity

  electronegativity.
          Atomic Radius
half of the distance between the centers of
two atoms of that element that are just
touching each other
atomic radius decreases across a period from
left to right and increases down a given
The atoms with the largest atomic radii are
located in Group I and at the bottom of
Why Does the Radius Have This
Moving from left to right across a period,
electrons are added one at a time to the
same outermost energy shell.
  Electrons within a shell cannot shield
   each other from the attraction to
   protons. Since the number of protons
   is also increasing, the effective
   nuclear charge increases across a
   period. This causes the atomic radius
   to decrease.
Why Does the Radius Have This
Moving down a group in the periodic
table, the number of electrons and filled
electron shells increases, but the
number of valence electrons remains the
  The outermost electrons in a group are
   exposed to the same effective nuclear
   charge, but electrons are found farther
   from the nucleus as the number of
   filled energy shells increases.
   Therefore, the atomic radii increase.
          Ionization Energy
The ionization energy also known as the
ionization potential, is the energy required to
completely remove an electron from a
gaseous atom or ion.
The closer and more tightly bound an
electron is to the nucleus, the more difficult it
will be to remove, and the higher its ionization
energy will be.
   The first ionization energy is the energy required
    to remove one electron from the parent atom.
    Successive ionization energies increase. The
    second ionization energy is always greater than
    the first ionization energy.
  Why is there an Ionization
       Energy Pattern
Ionization energies increase moving
from left to right across a period
(decreasing atomic radius).
Ionization energy decreases moving
down a group (increasing atomic
Group I elements have low ionization
energies because the loss of an
electron forms a stable octet.
         Electron Affinity
Electron affinity reflects the ability of an
atom to accept an electron.
It is the energy change that occurs
when an electron is added to a gaseous
Atoms with stronger effective nuclear
charge have greater electron affinity.
           Electron Affinity
The Group 2 elements, alkaline earths, have low
electron affinity values. These elements are relatively
stable because they have filled s subshells.
Group 17 elements, the halogens, have high electron
affinities because the addition of an electron to an
atom results in a completely filled outermost shell.
Group 18 elements, noble gases, have electron
affinities near zero, since each atom possesses a
stable octet and will not accept an electron readily.
Elements of other groups have low electron affinities
Electronegativity is a measure of the
attraction of an atom for the electrons in
a chemical bond.
The higher the electronegativity of an
atom, the greater its attraction for
bonding electrons.
     Why Electronegativity?
Electronegativity is related to ionization energy.
Electrons with low ionization energies have low
electronegativities because their nuclei do not exert a
strong attractive force on electrons.
Elements with high ionization energies have high
electronegativities due to the strong pull exerted on
electrons by the nucleus.
In a group, the electronegativity decreases as atomic
number increases, as a result of increased distance
between the valence electron and nucleus (greater
atomic radius).
   An example of an electropositive (i.e., low electronegativity)
    element is cesium; an example of a highly electronegative
    element is fluorine.
Moving Left → Right
 Atomic Radius Decreases
 Ionization Energy Increases
 Electronegativity Increases
Moving Top → Bottom
 Atomic Radius Increases
 Ionization Energy Decreases
 Electronegativity Decreases

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