The Periodic Table- Topic 5

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• The Periodic Table- Topic 5

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A. Dmitri Mendeleev (1869, Russian)
I. HISTORY

• Organized elements by
• increasing ATOMIC MASS.
• Elements with similar chemical properties were
  grouped together.
• There were some discrepancies.

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B. Henry Moseley

• ORGANIZED ELEMENTS BY INCREASING ATOMIC NUMBER.
• Resolved discrepancies in Mendeleevs
  arrangement.

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• When elements are arranged in order of INCREASING
  ATOMIC , elements with similar chemical
  properties appear at regular intervals.

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II. ORGANIZATION OF THE ELEMENTS
A. Arrangement of Table

• 1. Horizontal rows
• Called PERIODS
• All elements in the same period have the same
  number of ENERGY LEVELS in their atomic structure

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2. Vertical Columns

• Called GROUPS OR FAMILIES
• All elements in the same group have the
• same number of VALENCE ELECTRONS, therefore
  lose or gain the SAME number of electrons, form
  similar CHEMICAL FORMULAS and have similar
  CHEMICAL PROPERTIES
• ex. XCl2 Group 2
• Be 2 Cl -1 BeCl2
• Mg 2 Cl -1 MgCl2

• Remember When writing formulas, use the
  criss-cross rule to cancel out oxidation states

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III. Comparing Metals, Nonmetals Metalloids
Elements on the Periodic Table are divided into
three subgroups called METALS, NONMETALS and
METALLOIDS (semimetals).

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Increase nonmetallic properties
Increase metallic properties
Decrease metallic properties
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METALS located on the LEFT SIDE of the
periodic table (except H) MORE THAN 2/3 of all
elements

• 1. Chemical properties
• tend to LOSE ELECTRONS EASILY
• have LOW IONIZATION ENERGY (energy needed to
  remove electrons)
• Metallic character INCREASES as ionization energy
  decreases.
• have LOW ELECTRON AFFINITY (attraction for
  electrons)
• form POSITIVE IONS when combining with other
  atoms
• FRANCIUM most reactive metal See Table J
  http//castlelearning.com/review/reference/chem20
  table20j.htm

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2. Metals Physical Properties

• good conductors of heat and electricity
• LUSTROUS - reflect light, shine when they are
  polished
• MALLEABLE - can be rolled or hammered into
  sheets
• DUCTILE - can be drawn into wires
• are SOLIDS at room temperature except for
  MERCURY (liquid)

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B. NONMETALS
located on the right side of the periodic table
(except for Noble gases)

• Chemical properties
• tend to GAIN electrons to form NEGATIVE IONS
• have high electron affinities (electronegativity)
• produce COVALENT bonds by SHARING electrons with
  other nonmetals
• FLUORINE most reactive nonmetal see Table J

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• 2. Nonmetals Physical Properties
• exist as gases, molecular solids, or network
  solids at room temperature except BROMINE
  (liquid)
• BRITTLE - (shatters when struck)
• DULL - does not reflect light even when polished
• POOR CONDUCTORS of heat and electricity
• Allotropes Different SHAPE PROPERTIES forms
  from the same element.
• CARBON coal diamond, graphite
• OXYGEN O2 O3 (OZONE)

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C. METALLOIDS

• Found lying on the jagged line between metals and
  nonmetals flatly touching the line (except Al and
  Po).
• B,Si,Ge,As,Sb,Te At
• Exhibit properties of both metals and nonmetals
• Behave as nonmetals but their conductivity is
  like metals
• SEMICONDUCTORS Si and Ge

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IV. Periodic Trends use Table S

• A. Periodic Law
• When elements are arranged in order of increasing
  atomic , elements with similar properties appear
  at regular intervals.

• http//castlelearning.com/review/reference/chem2
  0table20s.htm

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1) Ionization Energy

• Energy needed to remove the most loosely bound
  electron from a neutral gaseous atom
• X energy X e-

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Trends in Ionization Energy
6.3
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Trends in Ionization Energy

• IE increases as you move across a period
• Why?
• The nuclear charge (atomic ) is increasing
  therefore greater attraction of the nucleus for
  electrons hence harder to remove an electron

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Trends in Ionization Energy

• IE decreases as you move down a group
• Why?
• Atom size increases making the outermost electron
  farther away from the nucleus therefore making it
  easier to remove
• Shielding increases

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• First Ionization Energy

• Increases UP and to the RIGHT

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Ionization Energy cont.

• Successive Ionization Energies

• Large jump in I.E. occurs when a CORE e- is
  removed.

• Mg 1st I.E. 736 kJ
• 2nd I.E. 1,445 kJ
• Core e- 3rd I.E. 7,730 kJ

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2. Atomic Radius

• ½ the distance between nuclei

• Decreases to the LEFT and Increases as you go DOWN

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Atomic Radius cont.

• Why is it larger going down?
• Higher energy levels have larger orbitals
• Shielding - core e- block the attraction between
  the nucleus and the valence e-
• Why is it smaller to the right?
• Increased nuclear charge without additional
  shielding pulls e- in tighter

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Comparison???

• Why is the ionization energy opposite that of
  atomic radius?
• In small atoms, e- are close to the nucleus where
  the attraction is stronger
• Why small jumps within each group?
• Stable e- configurations dont want to
• lose e-

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3. Electronegativity
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• the ability for an atom to attract electrons
  (electron affinity)
• Based on a scale of 4, Fluorine having the
  greatest EN

• A. Metals
• lose e-
• Form Cations ()
• get smaller

• B. Nonmetals
• gain e-
• Form Anions ()
• Get larger

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Trends in Electronegativity
6.3

• Representative Elements in Groups 1A through 7A

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4. Melting/Boiling Point Table S

• Melting/Boiling Point

• Highest in the middle of a period.

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Periodic Trends Summary(use reference Table S
for data comparison)
Trend Across a period (L to R) Down a group
Ionization energy increases decreases
Electronegativity increases decreases
Atomic radii decreases increases
Metallic properties decreases increases

• Click on for video clip 441

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IV. Classification

• Alkali Metals
• Alkaline Earth Metals
• Transition Metals
• Halogens
• Noble Gases

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Group 1 Alkali Metals

• extremely reactive (not found free in nature)
• form stable ionic compounds
• react with water to form a base
• react with air to form oxides
• react with acids to form salts

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Group 2 Alkaline Earth Metals

• reactive (not found free in nature) - form stable
  ionic compounds
• react with water to form a base
• react with air to form oxides
• react with acids to form salts

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Groups 3-11 Transition Metals

• multiple positive oxidation states
• Lose electrons from two outermost energy levels
• Ions form colored solutions

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Group 15 unique features

• Members range from typical nonmetals (nitrogen
  and phosphorus) through metalloids (arsenic and
  antimony) to metals (bismuth)
• Nitrogen
• Forms stable diatomic molecules with a triple
  bond
• Component of protein
• Forms some unstable compounds that are used as
  explosives
• Phosphorus
• Component of nucleic acids (DNA, RNA)
• More reactive than nitrogen at room temperature

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Group 16 unique Features

• Members range from typical nonmetals (oxygen and
  sulfur) through metalloids (selenium and
  tellurium) to metals (polonium)
• Solids except oxygen
• Oxygen can exist as O2 and O3 (it is an
  allotrope)
• Polonium is radioactive

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Group 17 Halogens

• very reactive nonmetals - high electronegativity
• not found free in nature
• form diatomic molecules when free
• react with metals to form salts (halides)
• Found in all three phases (s, l, g) due to
  differences in Van der Waals forces (these are
  weak)

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Group 18 Noble Gases

• Have complete outer shells
• Almost inert (not reactive) stable
• Krypton, xenon, and radon form compounds with
  oxygen and fluorine
• Referred to as monatomic gases