The Periodic Table

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The Periodic Table
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Dimitri Mendeleev

• Russian scientist developed the first published
  table in 1869.
• Arranged elements in order of atomic mass
• Believed most important property of elements

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• Elements with similar properties were placed in
  columns
• Showed periodicity / patterns
• left spaces for undiscovered elements

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Henri Moseley (1914)

• Used X-rays to reveal atomic numbers of several
  elements.
• Suggested that the elements should be arranged in
  order of atomic number instead of atomic mass.

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Modern Periodic Law(Moseley)

• The chemical and physical properties of the
  elements are the periodic functions of their
  atomic number.

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Reading the Periodic Table

• Groups
• Numbered 1 to 18
• Elements in groups (families) have similar
  characteristics

• Periods

.
When you write a sentence the period goes at the
end
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Classifications of Elements
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Metals

• Solids at STP (except Hg)
• good thermal (heat) conductors good electrical
  conductors
• Mobile electrons
• shiny luster
• ductile - drawn (made) into wire
• malleable - hammered into sheets
• in compounds have a positive oxidation state (,
  cation)

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• Elements to the left of the steps are metals
• except H, Ge Sb

Strongest metallic character found in group 1
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Non-metals

• Solid, liquid or gas at STP
• used as insulators
• dull
• brittle
• when combined in a compound have negative
  oxidation state (-, anion)

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Non-metals
strongest non-metal characteristics are found in
Group 18
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Metalloids

• a.k.a. Semimetals
• both metal and non-metal properties
• For example
• Boron
• Shiny poor conductor at room temp
• Arsenic
• Shiny brittle

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Metalloids

• sit on the steps
• 2 hide underneath

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Elements are found as (s), (l), (g)

• Solid
• particles have vibratory motion and are tightly
  packed
• Almost ALL of the elements
• Liquid
• particles can move throughout substance
• particles are farther apart than in solids
• conforms to shape of container
• ONLY Hg and Br

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• Gas
• H, N, O, F, Cl and Group 18
• molecules in constant, random, straight line
  motion
• molecules fill container
• large space between volumeless molecules
• conforms to the shape of the container

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Allotropes

• Different forms of an element found naturally in
  the same state
• Different molecular structures
• Different physical chemical properties

Carbon diamond, coal, graphite
buckminsterfullerene
Oxygen O2 Ozone O3
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Diatomic elements

• Elements found bonded to itself
• free state
• uncombined
• N2 O2 F2 Cl2 Br2 I2 and H2

H
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Monoatomic elements

• Dont form compounds
• Full valence
• Inert

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Electron configuration

• Periods correspond to the orbital which starts to
  be filled is the valence shell.
• Blocks refer to the sublevels being filled.
• Sublevels s, p, d, f

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Group / Family Characteristics

• Each group has characteristic properties that are
  directly related to electron configuration
  especially the number of valence electrons

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Group 1

• Alkali Metals
• Arabic al aqali the ashes
• Ash contains compounds of Na K
• all elements except hydrogen
• most reactive metal group
• SO reactive they are never found alone in nature
• always bonded to another element
• form 1 ions

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Group 2

• Alkaline Earth Metals
• second most reactive metals group
• SO reactive they dont occur alone in nature
• always bonded to another element
• form 2 ions

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d- block elements

• Transition Metals
• ions solutions are
• COLORful

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• Transition metals not as reactive as other metals
• some found in free state
• Au, Ag, Pt
• multiple oxidation states
• Iron (IV) oxide FeO2
• Iron (II) oxide FeO

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REVIEW

• Formula naming writing

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• Inner Transition Metals
• f-block
• Elements 93 and above, Pm Tc
• Not found in nature
• Man-made
• Metals in p-block
• other metals

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Group 17

• Halogens
• hals salt genesis to be born
• Halogens form many salt compounds
• Most reactive nonmetal group
• SO reactive that they dont occur by themselves
  in nature
• at least bonded to themselves (diatomic)
• form -1 ions

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Group 18

• Noble gases
• Inert gases
• dont like to combine with other elements
• Valence shells filled

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Ionization Energy
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First Ionization Energy

• Definition
• energy needed to remove the most loosely held e-

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Why 1st? Is there a 2nd?

• 2nd ionization energy
• Energy needed to remove the second electron
• May be higher or lower depending on the of
  valence electrons
• 3rd ionization energy
• Energy needed to remove the third electron
• May be higher or lower depending on the of
  valence electrons

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Whats holding the e- in place?

• Opposites attract
• There is a force of attraction between protons
  () and electrons (-)

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All elements want 8 e- (octet rule)

• Elements with only a few valence electrons will
  tend to have lower ionization E
• In other words
• It doesnt take a lot of E to remove their e-

Metals
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Most elements want 8 e- (octet rule)

• Elements with almost 8 valence electrons wont
  give them up so easy
• It takes a lot of E to remove their e-
• Non-metals have high IE

Who has the highest IEs?
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Ionization energy

• Hint energy needed to make an ion by losing
  electrons

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Ionization E trend?

• As move across a periodIE increases. WHY?

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• As you go across a period
• a large proton and a tiny electron are being
  added.
• more p hold the e- tighter when in the same
  energy level
• Increasing nuclear charge

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• As move UP a group IE increases. WHY?
• Period 1
• Period 2
• Period 3
• Period 4
• Lets figure it out . . .

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Potassium 19 p
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Sodium 11 p
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Lithium 3 p
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because . . .

• As you go up a group
• the electron shells get closer to the nucleus and
  can hold on to the e- tighter.
• ALSO.

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SHIELDING

• as you go DOWN a group
• Not only does the distance increase between the
  p and valence e-
• e- in outer shells repel each other
• For metals
• Lower i.e. is more reactive

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Electronegativity

• E-neg

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Electronegativity

• Ability of an atom to attract electrons of other
  atoms.
• In other words. . .
• Atoms with high e-neg are bullies that steal
  electrons

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Electronegativity

• Periodic trend

F
EXCEPT for Noble gases. WHY? Already have full
valence shell!
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Summary

• Atoms with
• high electronegativities
• also have
• high ionization E

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Atomic Radius

• Radius
• distance from the center of a circle (Nucleus)
  to the outermost edge (valence shell)

R
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Atomic Radius Periodicity
DECREASES
DECREASES
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Why?????

• Why does atomic radius DECREASE as you move up a
  group?

• Losing layers of e-

• Why does atomic radius DECREASE as you move
  across a period?

• Increasing the of p holds the e- in tighter

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Ionic Radius

• When elements are in their ionic state
• Full valence shell by losing or gaining e-

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NOW . . .

• Compare atomic radius to ionic radius

• Ionic radius is smaller

• Compare atomic radius to - ionic radius

• Ionic radius is larger

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Compare 2 different elements

• Compare the ionic radius of Mg2 and the atomic
  radius of Ne.

The ionic radius of Mg2 is smaller than the
atomic radius of Ne because the Mg 2 ion has
more p (12) than the Ne atom (10).
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General Formulas of compounds

• You can look at the groups on the periodic table
  and determine how they will combine with elements
  of different groups.

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• Group 1 Group 2 Group16 Group 17
• 1 2 -2
  -1

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General Formula AB
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• Group 1 Group 2 Group16 Group 17
• 1 2 -2
  -1

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General Formula AB
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• Group 1 Group 2 Group16 Group 17
• 1 2 -2
  -1

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General Formula A2B
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• Group 1 Group 2 Group16 Group 17
• 1 2 -2
  -1

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General Formula AB2
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Formula Writing Naming Review

• Lead IV oxide
• Phosphorus pentoxide
• SO3
• Oxygen
• Argon
• Aluminum oxide
• NiO

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Different Forms of the Periodic Table

• Changes were made to Mendeleevs table to look
  like the modern Periodic table we use today.
• This is not the only form of the periodic table
  that exists, however it is the most widely
  accepted.

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Stowes Physicists p.t.
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Benfey p.t.
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Zmaczynski p.t.
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Alexander Arrangement p.t.
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??????????? p.t.
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THE END