The Periodic Table

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The Periodic Table
Basic Concepts
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• 1800ish--Johann Dobereiner -- triads
• 1864 -- John Newlands -- octaves
• 1870--Dmitrii Mendeleev Julius
• Lothar Meyer--by mass
• 1913 -- Mosley--by number of protons

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Periodic Law

• When elements are arranged in order of increasing
  atomic number, there is a periodic repetition of
  their physical and chemical properties.

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Periods

• Horizontal rows on the table
• Correspond to the outermost energy level being
  filled.

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Groups

• Vertical columns on the table
• All groups have number designations
• Are also called families
• Same/similar physical and chemical properties due
  to VALENCE ELECTRONS!
• Some groups have special family names based upon
  characteristics of elements in that group

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Group Numbering Systems

• American Method
• IUPAC Method
• European Method

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Family Names

• Alkali Metals (Group 1)
• Alkaline Earth Metals (Group 2)
• Halogens (Group 17)
• Noble gases (Group 18)

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Alkali Metals (Group 1)

• Form metal hydroxides (strong bases) when
  reacting in water
• 2 Na 2 HOH ? 2 NaOH H2
• Are generally very reactive compared to other
  groups of metals
• Have one valence electron
• Form cations with a 1 charge

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Alkali metals
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Alkaline Earth Metals (Group 2)

• Form metal hydroxides (strong bases) when
  reacting in water
• Ca 2 HOH ? Ca(OH)2 H2
• Are not as reactive as alkali metals but are
  generally more reactive than transition elements
• Have two valence electrons
• Form cations with a 2 charge

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Alkaline earth metals
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Halogens (Group 17)

• Form a multitude of salts
• Are generally very reactive when compared to
  other nonmetals
• Have seven valence electrons
• Form anions with a -1 charge

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Halogens
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Noble Gases (Group 18)

• Are generally unreactive (inert)
• Have eight valence electrons
• Some compounds with xenon and krypton have been
  synthesized

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Regions of the Periodic Table

• Metals
• Nonmetals
• Metalloids (semi-metals)
• Transition metals
• Inner-transition metals
• Lanthanide Series
• Actinide Series
• Transuranium elements

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Metals
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• 1) Metals- high luster, good conductors, ductile,
  malleable, most are solid at room temp (except Hg
  is liquid)
• 2)Nonmetals- low luster, poor conductors, very
  brittle, various states of matter at room
  temperature (ex S is solid, O is gas, Br is
  liquid)
• 3) Metalloids- sit on stair-step line b/w metals
  and non-metals have properties between metals
  and non-metals

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Nonmetals
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Metalloids
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Transition Metals
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Inner-Transition Metals
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Regions of the Periodic Table Continued

• Radioactive elements ( Z gt 83)
• s block (Groups 1 2)
• p block (Groups 13-18)
• (n-1) d block (Groups 3-12)
• (n-2) f block (Lanthanide Actinide Series)
• Most active metals
• Most active nonmetals

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Orbital Blocks
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Periodic Table
Properties and Trends
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Common Periodic Properties

• Atomic Radius
• Ionization Energy
• Electron Affinity
• Ionic Radius
• Electronegativity
• Metallic Character
• Nonmetallic Character

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Periodic Trends

• A trend is NOT an EXPLANATION!

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Atomic Radii Trends on the Periodic Table

• For the main group elements
• atomic radii increase going down a group
• decrease going across a period.

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Going down a group radii increases

• Energy level is added for each successive period
• Each energy level shields (blocks) the influence
  of the nucleus

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Periodic table trends

• 1) Atomic Radii As you move down a group, atomic
  radius increases.     
• WHY? - The number of energy levels increases as
  you move down a group
•   Each subsequent energy level is further from
  the nucleus than the last. 

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• 2) 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."  In
  other words, there is a stronger force of
  attraction pulling the electrons closer to the
  nucleus resulting in a smaller atomic radius.

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Ionization Energy

• Is defined as the energy required to remove an
  electron from an atom in the gas phase.
• Ao(g) energy gt A(g) e -
• Each atom can have a series of ionization
  energies, since more than one electron can always
  be removed (except H).

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First Ionization Energy Trends on the Periodic
Table

• First ionization energies generally increase
  across a period and decrease down a group.
• Generally, the larger the atom the easier it is
  to remove an electron and the less ionization
  energy required.

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• group trends decreases as you move down a group
• why? the outermost electrons are found in
  higher energy levels as one goes down the group.
  Since the electrons are farther from the
  nucleus's pull the electrons are more easily
  removed.

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periodic trends increase from left to right

• why?
• as the atomic number increases in a period the
  nucleus is becoming stronger (more protons) but
  no new energy levels are being added. So atoms
  with larger atomic numbers have nucleus's that
  hold onto their electrons harder.

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• Low ionization energies are typical of active
  metals.
• High ionization energies are typical of active
  nonmetals.
• Very high ionization energies are found with the
  Noble Gases

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Plot of First Ionization Energies For Periods 1-4
Ionization Energy (kJ/mol)
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Ionic Radii

• Cations have lost one or more electrons and are
  smaller than the atoms from which they were
  derived.
• Anions have gained one or more electrons and are
  larger than the atoms from which they were
  derived.

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• 3) Anions (negative ions) are larger than their
  respective atoms.
• WHY? Electron-electron repulsion forces them to
  spread further apart. Electrons outnumber
  protons the protons cannot pull the extra
  electrons as tightly toward the nucleus.

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• 4) Cations (positive ions) are smaller than their
  respective atoms.
• WHY? Protons outnumber electrons the protons
  can pull the fewer electrons toward the nucleus
  more tightly. If the electron that is lost is the
  only valence electron so that the electron
  configuration of the cation is like that of a
  noble gas, then an entire energy level is lost. 
  In this case, the radius of the cation is much
  smaller than its respective atom.

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Relative Sizes of Some Common Ions
in picometers
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Metallic Character

• Is associated with
• Larger atomic radii
• Lower ionization energies
• Lower electron affinities
• Lower electronegativities
• The most active metals (ones with the most
  metallic character) are located in the lower left
  corner of the table.

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Nonmetallic Character

• Is associated with
• Smaller atomic radii
• Higher ionization energies
• Higher electron affinities
• Higher electronegativities
• The most active nonmetals are found in the upper
  right hand corner of the periodic table
  (excluding the Noble gases).

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What About the Noble Gases?

• do not behave as metals or nonmetals.
• very high ionization energies and positive
  electron affinity values.
• Noble gases usually are not assigned
  electronegativities due to their tendency to not
  form chemical bonds.