The Periodic Table

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The Periodic Table

• Early attempts to classify the elements
• Dobereiner (1817) triads - group of three
  elements in which the middle one has properties
  half way between the others.

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The Periodic Table

• John Newlands (1863) law of octaves - the 49
  known elements divided into 7 groups of 7.
  Chemical properties repeated every 8th element.
  (based on music)

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The Periodic Table

• Mendeleev (1860s) - grouped elements in
  accordance to repeating properties, left blank
  spaces for elements not yet discovered.

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The Periodic Table

• Mendeleev (continued) - arranged his table in
  accordance to increasing atomic mass.
• Stated the Periodic Law

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The Periodic Table

• Periodic Law - the properties of the elements are
  a periodic function of their atomic number.

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The Periodic Table

• Elements were discovered to support Mendeleev,
  but a couple were out of sequence by mass.

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The Periodic Table

• Moseley - refined Mendeleevs table. Use X-rays
  experiments to show that the nucleus had an
  integral positive charge.

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The Periodic Table

• Mendeleevs table was then supported based on
  atomic number, not atomic mass (the modern
  periodic table).

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The Periodic Table

• s-block, p-block, d-block, and f-block review?
• Noble gas configurations review?
• Z atomic number review?

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The Transition Metals

• d-block elements (groups 3-12).
• Energy sublevel overlap ex - 4s vs. 3d
• Multiple valence

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The Periodic Table

• Brightly colored compounds and solutions.

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The Periodic Table

• Lanthanoids elements 57-70, begins 4f block.
• Actinoids elements 89-102, begins 5f block.
• Rare earth metals.

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The Periodic Table

• Nature tends towards stability.
• Atoms seek bonding situations that result in
  stable electron configurations (ex. Share
  electrons).

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The Periodic Table

• Octet rule eight electrons in the outer level (s
  ps?) render an atom unreactive.
• Atoms seek to lose, gain, or share electrons to
  seek a stable octet of electrons.

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The Periodic Table

• An atom having a filled or half-filled sublevel
  is slightly more stable than an atom without.
• Full sublevels are more stable than half-filled.

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The Periodic Table

• Full outer levels are more stable than full
  sublevels.

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The Periodic Table

• Electron promotion one electron can be promoted
  to a slightly higher sublevel in order to produce
  a full or half filled sublevel.
• Ex - Cr and Cu (p. 154)

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The Periodic Table

• Metals vs Nonmetal vs Metalloid.
• Groups vs families.
• Survey of families in the table.

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The Periodic Table

• Metals
• 1. Have luster
• 2. Good conductors of heat and electricity.

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The Periodic Table

• Metals
• 3. Maleability can be hammered into a flat sheet
• 4. Ductility can be drawn out into a wire.

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The Periodic Table

• Nonmetals
• 1. Dull
• 2. Brittle
• 3. Poor conductors of heat and electricity
  (insulators)

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The Periodic Table

• Metalloids (semi-metals) - elements having a
  mixture of both metal and nonmetal
  characteristics.
• Located along the metal/nonmetal line.

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The Periodic Table

• Dull
• Maleable but not ductile
• Poor conductors except when combined with certain
  compounds can become superconductors.

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The Periodic Table

• Periods rows across the periodic table. Refers
  to the highest principle energy level. See p.
  152.
• Ex. K is in period 4 it has a 4s electron.

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The Periodic Table

• Groups families columns of related elements.
  Each family has a similar outer electron
  configuration. Ex. s2 and p5 elements.

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The Periodic Table

• On the periodic table, be able to identify the
  group location of the alkali, alkaline earth,
  and transition metals.

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The Periodic Table

• On the periodic table, be able to identify the
  group location of the pnicogens, chalcogen,
  halogen, and noble gas elements.

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Periodic Trends (Ch10)

• Both an elements position in the table, and its
  properties, are a direct result of its electronic
  configuration.

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The Periodic Table

• The properties are predictable and repeat
  themselves. They are based on electronic
  configuration.

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The Periodic Table

• Density increases then decreases across a
  period.
• Increases from metal to metal. Decreases as
  metalloids and nonmetals are approached.

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Atomic Radii

• The distance from the nucleus to the outermost
  orbital.
• Increases as you move down a group increased
  principle energy level.

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Atomic Radii

• Decreases as you move across a period.
• Z effective - the larger the charge density of a
  nucleus, the greater the pull on the electrons.

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Atomic Radii

• Z effective - this greater attractive force pulls
  the electrons slightly closer to the nucleus and
  accounts for the trend in atomic radii across a
  period.

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Radii of Ions

• Ions charged particles which are the result of
  adding or subtracting electrons from a neutral
  atom.

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Radii of Ions

• Cations ions with a positive charge (metals).
• Anions ions with a negative charge (non-metals).

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Radii of Ions

• Atoms will add or subtract electrons to complete
  their outermost energy level (they seek
  stability a full octet of electrons).

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Radii of Ions

• Nature does whatever is easiest.
• Ex. It is easier for potassium to lose 1 electron
  rather than gain 7 to complete its octet.

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Radii of Ions

• Ionic radii is based on whether an atom will add
  or subtract electrons when it ionizes.

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Radii of Ions

• Cations are smaller than their neutral atoms.
• Anions are larger than their neutral atoms.

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Radii of Ions

• Ionic radii increase down a group and decrease
  across a period. p. 252 (notice the group 18
  elements)

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Oxidation Numbers

• Predicting oxidation states for the main group
  elements (groups 1-2, 13-18) this will be based
  on the tendency of the group towards stability.

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Oxidation Numbers

• Predicting oxidation states for the transition
  elements - theory vs reality
• Ex. Zn and Ag

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Oxidation Numbers

• Transition elements
• 1. Can lose one or both of its two s-shell
  electrons first.
• 2. Can lose each individual d-shell electrons
  only after the outer s-shell is empty

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Oxidation Numbers

• Transition elements
• 3. Will not lose d-electrons if that shell is
  half-filled.
• Ex. Scandium (2, 3)
• vs Titanium (2, 3, 4)
• Vs Copper (1, 2)

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Ionization Energy

• Ionization energy the energy required to remove
  one electron from a gaseous atom.

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Ionization Energy

• Trends decreases down a group, increases across
  a period ( takes more energy to remove an
  electron from an element that usually forms an
  anion).

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Ionization Energy

• Factors affecting ionization energy
• Nuclear charge
• Shielding effect
• Radius
• Sublevel

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Ionization Energy

• Nuclear charge the larger the Z effective, the
  greater the force attracting the electrons
  greater ionization energy.

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Ionization Energy

• Shielding effect - core electrons shield the
  attractive nuclear force from outer electrons,
  lessening ionization energy. Electron/electron
  repulsion can increase the effect.

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Ionization Energy

• Radius- the greater the atomic radii, the lower
  the ionization energy.
• Sublevels- removing an electron from a
  half-filled or full sublevel requires more energy.

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Electron Affinity

• Electron Affinity an atoms ability to attract
  additional electrons.

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Electron Affinity

• Metals have low electron affinities. Non-metals
  have high electron affinities. The trend
  increase across a period and decreases down a
  group. Why?

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Ionization Energy

• Multiple ionization energies the second and
  third ionization energies can give clues as to
  atomic structure. Ex. Al vs Mg vs Na