The Periodic Table of the Elements

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The Periodic Table of the Elements

• Unit III

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Figure 1 The Periodic Table of the
Elements From http//www.vcs.ethz.ch/chemglobe/pt
oe/periodic.gif
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• The Periodic Table Historical Development

a) Dimitri Mendeleev Russian chemist
developed early periodic law organized elements
of the periodic table based upon atomic mass.

• Problem gaps existed between elements and same
  elements were found to have different atomic
  masses (isotopes).

b) Henry Moseley British physicist revised
Mendeleevs periodic table to produce todays
version organized elements based upon atomic
numbers.
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II. Structure of the Periodic Table

• The periodic table is organized into groups and
  periods.

• Groups Characteristics

• Groups are vertical families of elements.

• The elements within groups share similar
  chemical properties.

• Elements within groups have the same number of
  valence electrons.

• The groups of the periodic table range from 1-18.

Na.
Li.
Be Mg
Group 1 Elements
Group 2 Elements
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b) Periods Characteristics

• Periods are horizontal families of elements.

• The elements within periods have electrons that
  occupy the same energy levels.

• Elements within periods do not have the same
  number of valence electrons.

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Groups (1-18)
Periods 1-7
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1. Group 1 Alkali Metals

• The alkali metals are the most active metals.

• They each contain one valence electron.

• The alkali metals have oxidation numbers of 1.

• They will form positively charged ions (cations).

• They will have a tendency to lose electrons.

• Francium and Cesium are most active.

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2. Group 2 Alkaline Earth Metals

• The Alkaline Earth Metals are active metals.

• They contain two valence electrons.

• The Alkaline Earth Metals have oxidation numbers
  of 2.

• They will form positively charged ions (cations).

• The Alkaline Earth Metals will have a tendency
  to lose electrons.

• Radium and Barium are most active.

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3. Groups 3-12 Transition Elements
a) Characteristics

• Transition elements are all metals.

• multiple positive oxidation numbers.

• form colored aqueous solutions.

• have incomplete inner d-sublevels.

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4. Group 17 Halogens
Characteristics

• contain four diatoms (F2 (g), Cl2 (g), Br2 (l),
  I2 (s)).

• most active nonmetals.

• have elements that exist in all three states of
  matter.

• all have primary oxidation numbers of -1.

• contain seven valence electrons.

• will have a tendency to gain electrons.

• will have a tendency to form negatively charged
  ions (anions).

• Fluorine is the most active.

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5. Group 18 The Noble (Inert) Gases
Monoatomic Molecules
Characteristics

• most stable elements.

• all have complete valence shells.

• are extraordinarily inactive.

• are all gases at STP.

• are also known as the monatomic molecules.

NOTE Bromine (Br) is the only nonmetal that is
a liquid at STP. Mercury (Hg) is the only metal
that is a liquid at STP.
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III. Electronegativity, Ionization Energy,
Electron Affinity, and Atomic Radii
a) Electronegativity

• Electronegativity is a measure of the attraction
  that an atom has for electrons in a covalent bond.

• The force of attraction that an atoms nucleus
  has on its own valence electrons and those of
  other elements.

• The difference in electronegativity values
  between two elements enables chemists to predict
  the type of bonds in chemical compounds.

• Two scales of electronegativity are in common
  use the Pauling scale (proposed in 1932) and the
  Mulliken scale (proposed in 1934).

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1) The Pauling scale

• was devised in 1932.

• The most electronegative element (fluorine) is
  given an electronegativity value of 4.0

• The least electronegative element (francium) has
  a value of 0.7, and the remaining elements have
  values in between.

• Electronegativity values for all elements are
  found on Table S of The Physical Setting
  Reference Tables.

2) Mulliken Scale

• On the Mulliken scale, numbers are obtained by
  averaging ionization potential and electron
  affinity.

• Consequently, the Mulliken electronegativities
  are expressed directly in energy units, usually
  electron volts (eV).

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b) Ionization Energy

• The amount of energy needed to remove the most
  loosely bonded electron from an atom.

• Metals will generally have small ionization
  energies. This is due to the fact that metals
  will easily loose electrons due to their low
  electronegativity values.

• Nonmetals will generally have large ionization
  energies. This is due to the fact that nonmetals
  will NOT easily loose electrons due to their
  high electronegativity values.

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c) Atomic Radii
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IV. Classes of Elements Metals, Nonmetals, and
Metalloids
a) Metals

• account for 2/3 of the elements on the periodic
  table.

• are found on the left side of the periodic table
  (left of the steps).

• have a tendency to lose electrons.

• form positively charges ions (cations).

• are excellent conductors of heat and electricity.

• exist as solids at STP (except Hg).

• are malleable (can be flattened) and ductile
  (can be stretched into a wire) .

• are lustrous (shiny).

• have low electronegativities and ionization
  energies.

• have high atomic radii.

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b) Nonmetals

• account for 1/3 the elements on the periodic
  table.

• are found on the right side of the periodic
  table (right of the steps).

• have a tendency to gain electrons.

• have a tendency to form negatively charged ions
  (anions).

• are poor conductors of heat and electricity
  (good insulators).

• exist in all three states of matter.

• are brittle as solids.

• are dull as solids.

• have high electronegativities and ionization
  energies.

• have low atomic radii.

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c) Metalloids

• are found along side of the steps.

• have properties intermediate of metals and
  nonmetals.

• Examples include boron (B), arsenic (As),
  silicon (Si), and tellurium (Te).

NOTE Hydrogen is the only nonmetal that is
found on the left side of the periodic table.
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V. Trends of the Periodic Table
a) Electronegativity the force of attraction
that an atoms nucleus has on its own valence
electrons.

• Across a period Increases due to increased
  nuclear charge (number of positively charged
  protons).
• Down a group Decreases due to a greater
  distance between the positively charged nucleus
  and outer electrons and shielding.

b) Ionization Energy the amount of energy
needed to remove the most loosely bound valence
electron from a given atom.

• Across a period Increases due to the stronger
  force of attraction that atoms positively
  charged nucleus has on the negatively charged
  electrons. Meaning that it take more energy to
  remove an electron from that atom.

• Down a group Decreases due to the weaker force
  of attraction that atoms positively charged
  nucleus has on the negatively charged electrons.
  Nucleus and electrons are further apart.

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c) Atomic Radius

• ½ of the distance between the nuclei of two
  adjacently bonded atoms.

• Across a period Decreases due to the greater
  electronegativity. The nuclei of atoms are
  closer together.
• Down a group Increases due to the lower
  electronegativity. The nuclei of atoms are
  further apart.