Chapter 7' Periodic Properties of the Elements

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Chapter 7. Periodic Properties of the Elements

• Prof. G. Matthews

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7.1 Development of the periodic Table

• 1800 Thirty elements had been isolated and
  identified.
• 1829 Dobereiner several elements could be
  classified into triads.
• Similarity in chemical reactions and physical
  density and atomic mass.

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7.1 Development of the periodic Table

• 1860s Mendeleev suggested arranging the elements
  in a table according to increasing atomic mass.
• Predicted the existence of elements before their
  discovery ekasilicon (germanium).
• 1865 Newlands arranged the sixty two elements
  be arranged into groups of seven according to
  increasing atomic mass.

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7.1 Development of the periodic Table

• Mendeleev arranged his elements by increasing
  atomic mass across the periodic table, and by
  similar reactivity down the periodic table. This
  system is similar to the one used today.

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7.1 Development of the periodic Table

• 1913- Moseley arranged the elements in increasing
  atomic number based on his experimental work.
• Periodic Law properties of elements recur in a
  repeating pattern when arranged according to
  increasing atomic number.

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7.1 Development of the periodic Table

• Arranging the elements by atomic number
  illustrates the periodic, or repeating, pattern
  in properties that is the basis of the periodic
  table.

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7.2 Effective Nuclear Charge

• Inner-shell shielding
• Inner-shell electrons shield electrons farther
  out from some of the attractive pull exerted by
  the nucleus.
• Effective nuclear charge Z
• Actual nuclear charge experienced by an electron.

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7.2 Effective Nuclear Charge

• Both electrons in a helium atom experience the
  same attractive force from the nucleus.

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7.2 Effective Nuclear Charge

• Lithiums two first-shell electrons shield the
  second-shell electron from the nucleus.

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7.2 Effective Nuclear Charge

• Calculate the approximate value of Z on the
  outer electron by subtracting the number of
  inner electrons from the charge on the nucleus.

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7.3 Sizes of Atoms and Ions

• Atomic radius increases down a group more
  energy levels mean the atom is bigger.
• Atomic radius decreases across a period
  increase in effective nuclear charge.

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7.3 Sizes of Atoms and Ions

• A graph of atomic radius versus atomic number
  demonstrates periodic behavior.

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7.3 Sizes of Atoms and Ions

• As Zeff increases, the valence-shell electrons
  are attracted more strongly to the nucleus, and
  the atomic radius therefore decreases.

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7.3 Sizes of Atoms and Ions

• Cations have a smaller radius than their
  corresponding atoms
• Remove outer electrons.
• Effective nuclear charge pulls electrons inwards.
• Anions have a larger radius than their
  corresponding atoms
• Adding electrons to outer orbitals.
• Electrons are less tightly held.

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7.4 Ionization Energy

• Ionization Energy amount of energy needed to
  remove an electron from the atom in the gaseous
  state.
• Noble gases highest ionization energy.
• Alkali metals lowest ionization energy.

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7.4 Ionization Energy

• Sharp increase in ionization energy when an
  inner-shell electron is removed because the
  electron is strongly held by the effective
  nuclear charge.
• Decrease in ionization energy going from 2s to 2p
  orbital because 2s electrons are more effectively
  held.
• Be has a lower ionization energy than B.

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7.4 Ionization Energy

• Ionization energy decreases going down a group
  and increases from left to right across a period.
• Decreases down a group because the electrons are
  shielded and are further away from the attractive
  positive force of the nucleus.
• Increases from left to right because of the
  increase in effective nuclear charge (Z)

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7.4 Ionization Energy

• Ionization energy decreases going down a group
  and increases from left to right across a period.

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7.4 Ionization Energy

• Electronic Configuration of Ions
• Remove the electrons from the outermost orbitals
• Li 1s22s1 Li 1s2
• Fe Ar3d64s2 Fe2 Ar3d64s2
• Add the electrons to the outermost orbitals
• F1s22s22p5 F- 1s22s22p6

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7.5 Electron Affinities

• Electron Affinity
• The energy change that occurs when an electron is
  added to an atom in the gas state.
• Energy is released when an electron is added if
  the process is favored
• Cl(g) e- Cl-(g) ?E -349kJ/mol
• Ar(g) e- Ar-(g) ?E gt 0 kJ/mol

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7.5 Electron Affinities

• Electron Affinities
• Become more negative going across a period to the
  halogens.
• ?E gt0 for noble gases because adding an electron
  is unfavored.
• ?E gt0 for Group 5 elements because adding an
  electron would pair up electrons.
• ?E gt0 for Be and Mg because electrons would
  reside in an empty p subshell.
• Electron affinity for the halogens decreases
  because decreased electron repulsions.

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7.6 Metals, Nonmetals, and Metalloids

• Periodic table of the elements, showing the
  division of elements into metals, metalloids, and
  nonmetals.

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7.6 Metals, Nonmetals, and Metalloids

• Metallic character indication of an atoms
  ability to donate electrons.
• Metallic character increases down a group,
  electrons held less tightly.
• Metallic character - decreases across a period
  increasing effective nuclear charge.

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7.6 Metals, Nonmetals, and Metalloids

• Metals
• shiny substances, opaque, good conductors of heat
  and electricity, malleable, ductile i.e. Cu.
• Low ionization energies, tend to form cations.
• Tend to form ionic compounds.
• Form basic metal oxides.
• Left side of the table.

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7.6 Metals, Nonmetals, and Metalloids

• Nonmetals
• Do not conduct electricity or heat, not malleable
  or ductile, brittle i.e. C.
• Tend to form anions because of their electron
  afinities.
• Two or more nonmetals tend to form molecular
  substances.
• Form acidic nonmetal oxides
• Right side of the table.

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7.6 Metals, Nonmetals, and Metalloids

• Metalloids
• Elements that have properties of both metals and
  nonmetals.
• Weak conductors of electricity. i.e. Si.

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7.7 Group Trends for the Active Metals

• Group 1 alkali metals.
• Group 2 alkaline-earth metals.
• Group 7 halogens.
• Group 8 noble gases.
• Transition Elements IIIB-IIB.
• Inner Transition Elements Lanthanide and
  Actinide Series.

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7.7 Group Trends for the Active Metals
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7.7 Group Trends for the Active Metals

• Alkali metals
• Soft metallic solids.
• Low densities and melting points.
• Lowest ionization energy in a period.
• Exist in nature as compounds.
• React with water (Activity series)
• Reactions with oxygen depend on alkali metal and
  reaction conditions.
• Emit light of specific frequency in a flame test.

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7.7 Group Trends for the Active Metals

• Alkaline Earth Metals
• Metallic solids that are harder, more dense and
  have a higher melting point than the alkali
  metals.
• Ionization energies higher than the Alkali metals
  therefore not as reactive.
• Give off characteristic colors in flame tests.

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7.8 Group Trends for Selected Nonmetals

• Hydrogen
• Properties of metals and nonmetals, depending on
  the conditions.
• High ionization energy because its electron
  cant undergo nuclear shielding.
• Tends to form molecular compounds.
• Reacts with active metals to form metal hydrides.

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7.8 Group Trends for Selected Nonmetals

• Group 6
• Change from metallic to nonmetallic character in
  the elements.
• Oxygen forms metal oxides.
• Sulfur forms sulfides.

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7.8 Group Trends for Selected Nonmetals

• Group 7 The halogens
• Diatomic molecules.
• Gas, liquid, solid as go down the group.
• Halogens have high electron affinities.
• Form halides.
• Group 8. The Noble Gases
• Monoatomic.
• Unreactive.