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Periodic Trends Section 6.3

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Periodic Trends Section 6.3 Objectives: AOD C.3.1 Define atomic radii, ionization energy, electronegativity, and energy levels. AOD C.3.2 Recognize periodic trends of ... – PowerPoint PPT presentation

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Title: Periodic Trends Section 6.3


1
Periodic TrendsSection 6.3
  • Objectives
  • AOD C.3.1 Define atomic radii, ionization energy,
    electronegativity, and energy levels.
  • AOD C.3.2 Recognize periodic trends of elements,
    including the number of valence electrons, atomic
    size, and reactivity.

2
Atomic Radius
  • Def The distance between an atoms nucleus and
    its valence electrons.
  • Size of the atom varies from substance to
    substance

3
Trends within Periods
  • Atomic radius INCREASES as you move to the left
    across a period.
  • Due to decreased nuclear charge (total charge in
    the nucleus --- which would be based on WHAT?)
  • Decreased nuclear charge has less pull on the
    outermost (highest energy, valence) electrons so
    they can move further away from each other (and
    the nucleus).

4
Trends within Groups
  • Atomic radius INCREASES as you move down a group.
  • Nuclear charge increases, but as electrons are
    added to higher energy levels, they are further
    from the pull of the nuclear protons, AND they
    are shielded by the electrons between them and
    the nucleus.
  • Summary Increases left and down!

5
Examples
  • Which element has the smallest atomic radius?
    Largest atomic radius?
  • Iodine (I)
  • Bromine (Br)
  • Fluorine (F)
  • Chlorine (Cl)
  • Fluorine
  • Iodine

6
Ionic Radius
  • Ion an atom that gains or loses electrons
  • When atoms lose electrons they form positive ions
    and become smaller.
  • The electron lost will always be a valence
    electron.
  • Loss of valence electrons may leave an empty
    outer orbital which results in a smaller radius.
    (Example Na1)
  • Even if the orbital is not emptied, repulsion
    between fewer electrons decreases allowing them
    to be pulled closer to the nucleus.

7
Ionic Radius, continued..
  • When atoms GAIN electrons they form negative ions
    and they always become larger.
  • The addition of an electron to an atom increases
    the repulsion between the valence electrons,
    forcing them to move farther apart (Example
    O-2).
  • The result is a larger radius.

8
Ionic Radius within Periods
  • Size of the positive ions gradually increases
    from right to left across a period.
  • The size of the much larger negative ions also
    gradually increases from right to left, until you
    reach the smaller positive ions.
  • Ionic radius GENERALLY INCREASES to the left
    across periods.

9
Ionic Radius within Groups
  • As you move down a group an ions outer electrons
    are in higher principal energy levels resulting
    in a gradual increase in ionic size.
  • Ionic radius INCREASES as you move down a group.
  • Summary Increases left and down.

10
Ionization Energy
  • Def the energy required to remove an electron
    from an atom in the gaseous state
  • How strongly an atoms nucleus holds on to its
    valence electrons
  • High IE indicates atom has a strong hold on its
    electrons
  • Low IE indicates an atom loses its outer electron
    more.easily

11
Ionization Energy
  • Energy required to remove the 1st electron is the
    first ionization energy.
  • Energy required to remove the 2nd electron is
    the second ionization energy.
  • 1st IE is ALWAYS highest.

12
Ionization Energy within Periods
  • INCREASES as you move from left to right across a
    period
  • The increased nuclear charge of each successive
    element produces an increased hold on the valence
    electrons, as they are all in the same principal
    energy level.

13
Ionization Energy within Groups
  • INCREASES as you move up a group because the
    valence electrons are closer to the nucleus.
  • Summary Increases up and right.

14
Octet Rule
  • Atoms tend to gain, lose, or share electrons in
    order to acquire a full set of eight valence
    electrons (resembling which elements?).
  • Elements on the right side of the periodic table
    (nonmetals) tend to gain electrons in order to
    acquire the 8 valence electrons, forming negative
    ions)
  • Elements on the left side of the periodic table
    (metals) tend to lose electrons and form positive
    ions.

15
Electronegativity of an Element
  • Indicates the relative ability of an atom to
    attract electrons in a chemical bond.
  • Noble Gases are not assigned values
  • Fluorine is the most electronegative element.
  • Fr Cs are the least.
  • In a chemical bond the atom with the greater
    electronegativity more strongly attracts the
    shared electrons.

16
Electronegativity Trends within Periods Groups
  • INCREASES as you move up a group and across a
    period (up and right).
  • The lowest electronegativities are found at the
    lower left side
  • Highest are found at the upper right side

17
Example Problems
  • Which element has the highest electronegativity?
    Lowest?
  • N- Nitrogen
  • P- Phosphorus
  • As-Arsenic
  • Sb-Antimony
  • Bi- bismuth
  • Nhighest
  • Bi Lowest

18
Homework Problems
  • Pg. 175
  • 56, 57, 59, 62, 63, 65-67
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