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Periodic Trends

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Periodic Trends Electronegativity - Groups Electronegativity decreases going down a group The bonding electrons are increasingly distant from the attraction of the ... – PowerPoint PPT presentation

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Title: Periodic Trends


1
Periodic Trends
2
Periodic Trends
  • What is a trend?
  • A trend is the general direction in which
    something tends to move.

3
Periodic Trends
  • The elements on the Periodic Table of Elements
    show many trends in their physical and chemical
    properties.
  • Across the rows (periods or series)
  • Down the columns (groups or families)

4
Atomic Radius
  • ½ the distance between the nuclei of 2 like atoms
    in a diatomic molecule
  • The atoms of the 8 main groups are shown here.

5
Atomic Radius - Groups
  • Atomic radius increases as you move down a group
  • Why?
  • More electrons in more Principal Energy Levels
  • Atomic size increases

6
Atomic Radius - Periods
  • Atomic radius decreases as you move across a
    period
  • Why?
  • (-) electrons increase, but so do () protons !!!
  • Increased () nuclear charge pulls the (-)
    electrons closer to the nucleus
  • Atomic size decreases

7
Atomic Radius - Periods
  • The size trend in periods is less pronounced than
    in groups because of the electron shielding
    effect.

8
Shielding Effect
  • Reduction in effective nuclear charge on an
    electron that is caused by the repulsive forces
    of other electrons between it and the nucleus
  • In an atom with one electron, that electron
    experiences the full charge of the positive
    nucleus. However, in an atom with many electrons,
    the outer electrons are simultaneously attracted
    to the positive nucleus and repelled by the
    negatively charged electrons.

9
Atomic Radius - Graph
10
Atomic Radius
  • Click on source to see a short video
  • Source http//cwx.prenhall.com/petrucci/medialib/
    media_portfolio/text_images/046_AtomicRadii.MOV

11
Atomic Radius
12
Ionic Radius
  • What are ions?
  • Ions are charged atoms, either or -
  • Cations are positive ions
  • Cations form when atoms lose electrons
  • Anions are negative ions
  • Anions form when atoms gain electrons

13
Ionic Radius
14
Ionic Radius Cations Group
  • Cations are smaller than their parent atoms.
  • Why?
  • By losing their valence electrons, they lose
    their entire valence shell
  • Cations are formed by the metals on the left side
    of the Periodic Table

15
Ionic Radius Cations Groups
  • Ionic size increases as you move down a group for
    the same reason atomic size increases
  • Number of principal energy levels increases

16
Ionic Radius Anions Groups
  • Anions are larger than their parent ions
  • Why?
  • When extra (-) electrons are added, extra ()
    protons are NOT added to the nucleus
  • Effective nuclear attraction is less for the
    increases number of electrons

17
Ionic Radius Anions Group
  • Ionic size increases as you move down a group for
    the same reason atomic size increases
  • Number of principal energy levels increases
  • Cations are formed by nonmetals on the right side
    of the Periodic Table

18
Ionic Radius - Periods
  • Just like their parent atoms
  • Cations get smaller as you move from left to
    right
  • Anions get smaller as you move from left to right
  • Increased () nuclear charge pulls the (-)
    electrons closer to the nucleus

19
Ionic Radius
20
Ionization Energy
  • Energy is needed to remove an electron from an
    atom
  • The energy needed to overcome the attraction of
    the nuclear charge and remove an electron (from a
    gaseous atom) is called the Ionization Energy

21
1st Ionization Energy
  • The energy needed to remove the 1st electron from
    an atom is the 1st Ionization Energy

22
Factors AffectingIonization Energy
  • Atomic Radius
  • Smaller atoms hang on to valence electrons more
    tightly, and so have higher ionization energy

23
Factors AffectingIonization Energy
  • Charge
  • The higher the positive charge becomes, the
    harder it is to pull away additional electrons
  • Second ionization energy is always higher than
    the first

24
Factors AffectingIonization Energy
  • Orbital Type
  • It's easier to remove electrons from p orbitals
    than from s orbitals, which are deeper

25
Factors AffectingIonization Energy
  • Electron Pairing
  • Within a subshell, paired electrons are easier to
    remove than unpaired ones
  • Reason repulsion between electrons in the same
    orbital is higher than repulsion between
    electrons in different orbitals

26
Factors AffectingIonization Energy
  • Electron Pairing Example
  • On the basis of gross periodic trends, one might
    expect O to have a higher ionization energy than
    N. However, the ionization energy of N is 1402
    kJ/mol and the ionization energy of O is only
    1314 kJ/mol.
  • Taking away an electron from O is much easier,
    because the O contains a paired electron in its
    valence shell which is repelled by its partner.

27
1st Ionization Energy
28
2nd, 3rd, 4th Ionization Energies, etc.
  • Subsequent electrons require more energy to
    remove than the first electron
  • How much more energy is needed depends on what
    energy levels and orbitals the electrons are in

29
2nd, 3rd, 4th Ionization Energies, etc.
  • Source http//chemed.chem.purdue.edu/genchem/topi
    creview/bp/ch7/ie_ea.html

30
Ionization Energy
  • Click on source to see a short video.
  • Source http//cwx.prenhall.com/petrucci/medialib/
    media_portfolio/text_images/047_IonizationEner.MOV

31
Electronegativity
  • Ability of an atom to attract electrons toward
    itself in a chemical bond

32
Electronegativity
  • The difference between the electronegativities
    of two atoms will determine what kind of bond
    they form
  • Linus Pauling used an element's ionization energy
    and electron affinity to predict how it will
    behave in a bond.
  • The more energy it takes to pull off the outer
    electron of an atom, the less likely it is to
    allow another atom to take those electrons. The
    more energy the atom releases when it gains an
    electron, the more likely it is to take electrons
    from another atom in bonding. These two energies
    were used to compute a numerical score.

33
Electronegativity - Periods
  • Electronegativity increases going left to right
    across the periodic table.
  • Fluorine's high nuclear charge coupled with its
    small size make it hold onto bonding electrons
    more tightly than any other element. Lithium has
    a lower nuclear charge and is actually larger
    than fluorine. Its valence electron is not
    tightly held and it tends to surrender it in
    chemical bonds.

Li Be B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
34
Electronegativity - Groups
  • Electronegativity decreases going down a group
  • The bonding electrons are increasingly distant
    from the attraction of the nucleus

H 2.1
Li 1.0
Na 0.9
K 0.8
Rb 0.8
Cs 0.7
Fr 0.7
35
Electronegativity
36
Electronegativity
37
Electron Affinity
  • The electron affinity is a measure of the energy
    absorbed when an electron is added to a neutral
    atom to form a negative ion.
  • Most elements have a negative electron affinity.
    This means they do not require energy to gain an
    electron instead, they release energy.
  • Atoms more attracted to extra electrons have a
    more negative electron affinity.
  • The more negative the value, the more stable the
    ion is.

38
Electron Affinity
  • Click on source to see a short video.
  • Source http//cwx.prenhall.com/petrucci/medialib/
    media_portfolio/text_images/049_ElectrAffinity.MOV

39
Electron Affinity
  • Electron affinity is essentially the opposite of
    the ionization energy.

40
Electron Affinity - Trends
  • Click on source to see a short video.
  • Source http//cwx.prenhall.com/petrucci/medialib/
    media_portfolio/text_images/050_PeriodElectron.MOV

41
Image Sources
  • http//www.Chem4kids.com
  • http//images.encarta.msn.com
  • http//antoine.frostburg.edu
  • http//www.webelements.com
  • http//cwx.prenhall.com
  • http//www.800mainstreet.com
  • http//intro.chem.okstate.edu/1215

42
Credits
  • PowerPoint Adela J. Dziekanowski
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