Periodic Table

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Periodic Table

• Chapter 6

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Periodic Table

• Many different versions of the Periodic Table
  exist
• All try to arrange the known elements into an
  organized table

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Alternate Periodic Tables
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Alternate Periodic Tables
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Alternate Periodic Tables
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Alternate Periodic Tables
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Alternate Periodic Tables
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Alternate Periodic Tables
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Alternate Periodic Tables
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History

• 1869 - Russian chemist and teacher, Dmitri
  Mendeleev proposed a table for organizing
  elements
• Mendeleev arranged the elements in a table based
  on increasing atomic mass.

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History

• Mendeleev placed elements next to each other with
  similar chemical properties
• He would leave elements out of order based on
  atomic mass if they lined up better based on
  chemical properties

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History

• Mendeleev left spaces for elements not yet
  discovered
• He predicted properties of elements that would
  fit in those spots
• He predicted very closely the properties of Ge,
  Ga, Sc, and 5 others

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History

• 1913 - British physicist, Henry Moseley,
  determined the atomic numbers for the elements
• The modern periodic table is arranged in order of
  increasing atomic number.

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Periodic Table
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Arrangement

• Columns are called Groups
• Numbered 1-18
• Rows are called Periods
• Elements in the same group have similar properties

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Group Names

• Group 1 - Alkali Metals
• Group 2 - Alkaline earth metals
• Group 17 - Halogens
• Group 18 - Inert or Noble gases.

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Group Names

• Groups 3-11 Transition Metals
• Bottom 2 rows Inner Transition

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Phases at STP

• Most elements are solids at STP
• Hg and Br are liquids at STP
• H, N, O, F, Cl and Noble Gases are all gases at
  STP

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Periodic Law

• 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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Valence Electrons

• Electrons in outermost occupied energy level
• Valence Electrons are responsible for most
  chemical properties
• Elements in the same group have similar
  properties because they have the same number of
  valence electrons

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Classifying Elements

• Elements are classified into 3 groups based on
  their properties
• Metals Left and Middle
• Nonmetals Right
• Metalloids - Staircase

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Metals

• Good conductors of heat and electrical current
• High luster or sheen
• Many are ductile, meaning they can be drawn into
  wires
• Most are malleable, meaning they can be hammered
  into thin sheets

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Metals

• Metallic Character increases as you move towards
  the lower left
• Most Metallic Element is Francium, Fr

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Nonmetals

• Most are gases at room temperature, some are
  solids, and one is liquid
• Most are poor conductors
• Most solids are brittle

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Nonmetals

• Non-Metallic Character increases as you move
  towards upper right
• Most nonmetallic element is Fluorine, F

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Metalloids

• B, Si, Ge, As, Sb, Te
• Have properties of both metals and nonmetals,
  based on conditions
• Exceptions
• Al and Po are metals
• At is a nonmetal

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Group Characteristics

• Alkali Metals (Group 1)
• H, Li, Na, K, Rb, Cs, Fr
• All have 1 valence electron, tend to form 1 ions
• Most reactive metals
• Not found in nature by themselves, always
  combined with someone else
• Have properties of metals but are softer and less
  dense

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Group Characteristics (cont)

• Alkaline Earth Metals (Group 2)
• Be, Mg, Ca, Sr, Ba, Ra
• All have 2 valence electrons, tend to form 2
  ions
• Harder and more dense than alkali metals, but
  also have higher melting and boiling points
• Highly reactive, but not as much as alkali metals
• Not found by themselves in nature

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Group Characteristics (cont)

• Halogens (Group 17)
• F, Cl, Br, I, At
• All have 7 valence electrons, tend to form -1
  ions
• Strongly non-metallic
• Most active nonmetals
• Have low melting and boiling points
• Combine readily with metals to form salts

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Group Characteristics (cont)

• Noble Gases (Group 18)
• He, Ne, Ar, Kr, Xe, Rn
• Colorless gases that are extremely non-reactive
• Full valence shell, non-reactive
• All are found in small amounts in our atmosphere

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Group Characteristics (cont)

• Transition Metals (Groups 3-11)
• Most are excellent heat and electrical conductors
• Most have high melting points and are hard,
  except Hg
• Less active than group 1 and 2 metals
• Many combine with Oxygen to form oxides (Chemical
  property)
• Many have more than one oxidation number
• Form compounds that are colorful

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Reminder

• STP
• Standard Temperature and Pressure
• 1 atm, 0C
• Reference Point for most measurements

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Diatomics

• Eight elements are diatomic molecules when alone
  in nature (exist as two atoms bonded together)
• H2, N2, O2, F2, Cl2, Br2, I2, At2

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Diatomics

• Hydrogen and the Magic 7

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Coloring

• Color in the specific groups with your own color
  choices

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Coloring

• Color in the different classifications with your
  own color choices
• Metals, Nonmetals, Metalloids

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Orbital Blocks
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Periodic Trends

• How a property changes either across a period or
  down a group
• Atomic Number
• Atomic Mass
• Atomic Radius
• Ionic Radius
• Ionization Energy
• Electronegativity

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Trends

• Atomic number increases across a period.
• Increasing number of protons
• Atomic number increases down a group
• Increasing number of protons

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Trends

• Atomic mass generally increases across a period.
• Increasing protons, neutrons, and electrons.
• Atomic mass increases down a group.
• Increasing protons, neutrons, and electrons.

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Radius

• Atomic Radius measure of the size of the atom
• Half the distance between two nuclei
• Ionic Radius measure of the size of an ion

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Trends

• Atomic Radius decreases across a period
• More protons to pull on the electrons
• Atomic Radius increases down a group
• Increasing electrons into more energy levels
  (more shells)

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Ions

• Atom, or group of atoms, that has gained or lost
  electrons
• Cation positive ion
• Anion negative ion

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Ions

• When an atom loses an electron, it becomes
  positively charged
• The radius becomes smaller
• Metals tend to lose electrons

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Ions

• When an atom gains an electron, it becomes
  negatively charged
• The radius becomes larger
• Nonmetals tend to gain electrons

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Trends

• Ionic Radius decreases for positive ions across a
  period
• More protons to pull on the electrons
• Ionic Radius decreases for negative ions across a
  period
• More protons to pull on the electrons

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Ionic Radius
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Trends

• Ionic Radius increases down a group
• Increasing electrons into more energy levels
  (more shells)

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Ionization Energy (IE)

• Amount of energy required to remove an electron
  from an atom
• Ca ? Ca e- 590kJ/mol
• First ionization energy is removing the first
  electron
• Second Ionization energy is removing the second
  electron after having the first removed
• Ca ? Ca2 e- 1145kJ/mol

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IE Trends

• Ionization energy tends to increase across a
  period
• More protons are able to hold on tighter to
  electrons
• Ionization energy tends to decrease down a group
• Electrons are farther away from the protons (more
  shells)

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Electronegativity (EN)

• Ability of an atom to attract an electron from
  another atom when in a compound.
• Noble gases are usually omitted since they dont
  form compounds
• Fluorine, F, is the most electronegative element
  with a value of 4.0
• Francium, Fr, is the least electronegative
  element with a value of 0.7

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EN Trends

• Electronegativity tends to increase across a
  period
• More protons are able to attract electrons better
• Electronegativity tends to decrease down a group
• Electrons are farther away from the protons (more
  shells)

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Trends Summary
Property Period (L?R) Group (T?B)
Atomic Number
Atomic Mass
Atomic Radius
Ionic Radius
Ionization Energy
Electronegativity
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Reactivity

• Elements that are more reactive tend to either
  gain or lose electrons very easily
• Elements that lose electrons easily have low IE
  and low EN
• Lower left, Fr
• Elements that gain electrons easily have high IE
  and high EN
• Upper right, F