The Periodic Table

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The Periodic Table
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Dmitri Mendeleev(1834 1907)

• He organized elements into the first periodic
  table
• He arranged elements by increasing atomic mass

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Henry Moseley(1913)

• He arranged elements according to atomic number
  rather than atomic mass
• The modern periodic table is arranged by atomic
  number

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Periodic Law

• The periodic law states that
• there is periodic repetition
• of chemical and physical
• properties of elements

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The Modern Periodic Table

• There are 18 groups (columns up and down)
• The group As (the tall columns) are called
  representative elements

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• The group Bs (the middle columns) are called
  transition metals

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• There are seven periods
• (rows across the periodic
• table)

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• Metals are to the LEFT of the
• zig-zag line (except hydrogen!)

Metals in yellow
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• Nonmetals are to the RIGHT of
• the zig-zag line

nonmetals in red
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Metalloids

• Metalloids are those elements ON the zig-zag line

Metalloids border the zig-zag line
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• Now . . .YOU fill in the
• chart using your book!

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Metals

• solid at room temperature
• shiny (have luster) and smooth
• good conductors of heat and electricity

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Metals

• malleable bendable (can be pounded into
  sheets)
• ductile - can be pulled
• into wires

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Metals

• react with acids
• mercury (Hg) is the only LIQUID metal

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Nonmetals

• generally gases or brittle, dull looking solids
  at room temperature
• poor conductors of heat and electricity
• Bromine (Br) is the only LIQUID nonmetal

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Metalloids

• sometimes called semimetals
• metalloids have properties of both metals and
  nonmetals

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Metalloids

• silicon and germanium are two of the most
  important metalloids (theyre used in computer
  chips and solar cells)

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Trends of the Periodic Table
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Periodic Law

• If elements are organized
• according to atomic number,
• their properties will repeat
• periodically

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Four Periodic Trends

• Atomic radii
• Ionic radii
• Electronegativity
• 4.Ionization energy

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Atomic Radius

• The atomic radius basically
• tells you the size of the
• atom. It is half the distance between two nuclei
  of identical atoms bonded together.
• 
  

radius
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The trend atomic radii DECREASE across a period

• Why?
• Each time a positive proton is added to the
  nucleus, the negative electrons feel a greater
  attraction to the positively-charged nucleus and
  get pulled in tighter

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Decreasing getting smaller!
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Trend atomic radiiINCREASE down a group

• Why?
• electrons are added to higher and higher energy
  levels as you go down

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Atomic radii DECREASE down a group!
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• The farther the electrons
• from the nucleus, the larger
• the atomic radii!!!!!

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Try these . . .

• Which element has the
• larger atomic radius C or F?
• ? carbon
• 2.Which element has the smaller atomic radius Ar
  or Kr?
• ? argon

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Ionic Radii

• Is basically the size of an ion or half the
  distance between the nuclei of two ions bonded
  together
• What is ion???

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Ionic Radii

• Ion an atom with a charge ( or - )
• An ion is formed when atoms lose or gain electrons

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• What happens to an atom if it LOSES an electron?
• ? it loses a negative charge so it becomes
  POSITIVE

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• A positively charged ion is called a cation

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• Positive ions (cations) are
• smaller than the atoms they
• come from because they lose
• electrons making the atom
• smaller.

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• What happens to an atom if it GAINS an electron?
• ? It gets more negative
• (so it has a negative charge)

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• A negatively charged ion is
• called an anion.

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Negative ions (anions) are LARGER than the atoms
they come from because they gain electrons
making the atom LARGER!
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The trend

• See the board

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Remember . . .

• all atoms want a full octet
• (8 valence electrons)
• atoms with 1 valence electron will give up that
  electron VERY QUICKLY to become stable

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example

• sodium has one valence
• electron 1s22s22p63s1
• if sodium gives it away,
• then the configuration will be
• 1s22s22p6
• sodium will have a full octet

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• atoms with 7 valence electrons will hold on to
  those electrons VERY TIGHTLY
• they try to get one more and become stable

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Ionization Energy

• The amount of energy needed to remove an electron
• think of it as how tightly an atom holds on to
  its electrons

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The trend

• ionization energy INCREASES across a period

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• Why?
• the more valence electrons an element has, the
  more difficult it is to remove them!

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The trend

• Ionization energy DECREASES
• down a group

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• Valence electrons in higher energy levels are NOT
  held as tightly because they are farther from the
  nucleus
• Therefore, it is easier to remove an electron
  that is farther from the nucleus

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Try these . . .

• Which has a higher ionization energy Na or Cl
• ? Chlorine
• Which has a lower ionization energy Li or O
• ? Lithium

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Electronegativity

• The ability of an atom to attract electrons to
  itself

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The Trend electronegativity INCREASES across a
period

• Why?
• atoms are trying harder to attract electrons to
  get a full octet

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The trend electronegativity DECREASES down a
group

• Why?
• it is harder to hold on to the electrons that are
  farther away from the nucleus

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Try these . . .

• Which element is more electronegative? F or Br
• ? Fluorine
• Which element is more electronegative? B or Ca
• ? Boron

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Finished!