Chapter 6 The Periodic Table and Periodic Law

1
Chapter 6 The Periodic Table and Periodic Law

• The elements, which make up all living and
  non-living matter, fit into a orderly table.
  When interpreted properly, the table describes
  much of the elements physical and chemical
  properties.

2
What is the Periodic Law and how was it
formulated?

• Demitri Mendeleev is known as the father of the
  periodic table
• He arranged the elements in families (groups) and
  periods (rows,series) according to atomic mass
  and properties

3

• Mendeleev noted that the chemical and physical
  values for elemental properties would either be
  high or low depending upon the group under
  observation.
• He proposed the first Periodic Law "The
  properties of the elements are a periodic
  function of their atomic masses"
• Left blanks in his table for undiscovered
  elements

4
Moseleys Modern-Periodic Law

• There was some inconsistencies with Mendeleevs
  table
• In the early 1900s Moseley was able to
  experimentally determine the atomic number of all
  known elements

5

• Moseley then proceeded to rearrange the elements
  according to increasing atomic numbers.
• New/Modern Periodic law states that the
  properties of elements are a periodic function of
  their atomic number

6
The Modern Periodic Table

• Glenn T. Seaborg won the Nobel Prize for his work
  in nuclear chemistry
• In 1944, formulated the actinide concept of
  heavy element electronic structure. This concept
  predicted that the fourteen actinides

7
Some Characteristics of Groups
8
Group 1 (IA) - Alkali Metals

• ( metal characteristics - shiny, malleable,
  ductile, good conductors)
• Very active metals - activity increases as you go
  down group
• All have one valence electron - form 1 cations
  by losing an electron
• react violently with water

9
Group 2 (IIA) - Alkaline Earth Metals

• activity increases as you move down the column
  not as reactive as alkali metals
• Ca, Sr, and Ba react violently when they come
  into contact with water
• All have two valence electrons
• Form 2 cations by losing 2 electrons

10
Group 17 (VIIA) - Halogens

• All gain one electron to form anions with a
  charge of -1.
• All are nonmetals except for At which is a
  semimetal
• All are diatomic in their elemental form

11
Group 18 (VIIIA) - Noble (Rare) Gases

• Mistakenly labeled as "inert gases" until about
  30 years ago because it was thought that these
  gases did not react with anything.
• Noble gases have filled valence (outermost)
  shells.

12
Periodic Trends

• Atomic Radii
• 1) As you move down a group, atomic radius
  increases.
• 2) WHY? - The number of energy levels
  increases as you move down a group as the number
  of electrons increases.

13

• As you move across a period, atomic radius
  decreases.
• WHY? - As you go across a period, electrons are
  added to the same energy level. At the same time,
  protons are being added to the nucleus. The
  concentration of more protons in the nucleus
  creates a "higher effective nuclear charge."

14
First Ionization Energy

• Definition The energy required to remove the
  outermost (highest energy) electron from a
  neutral atom in its ground state.
• 1) As you move down a group, first ionization
  energy decreases.
• 2) WHY? Electrons are further from the nucleus
  and thus easier to remove the outermost one -
  shielding effect of other electrons

15

• 3) As you move across a period, first
  ionization energy increases.
• 4) WHY? - As you move across a period, the
  atomic radius decreases, the attraction from the
  positive nucleus gets larger

16
Electronegativity

• Definition The ability of an element to attract
  electrons in a chemical bond
• 1) As you move down a group,
• electronegativaty decreases
• 2) As you move across a period, it
• increases.

17
(No Transcript)