Periodic Trends

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Periodic Trends
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Periodic Trends

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

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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)

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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.

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Atomic Radius - Groups

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

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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

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Atomic Radius - Periods

• The size trend in periods is less pronounced than
  in groups because of the electron shielding
  effect.

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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.

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Atomic Radius - Graph
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Atomic Radius

• Click on source to see a short video
• Source http//cwx.prenhall.com/petrucci/medialib/
  media_portfolio/text_images/046_AtomicRadii.MOV

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Atomic Radius
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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

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Ionic Radius
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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

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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

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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

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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

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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

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Ionic Radius
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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

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1st Ionization Energy

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

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Factors AffectingIonization Energy

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

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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

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Factors AffectingIonization Energy

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

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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

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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.

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1st Ionization Energy
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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

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2nd, 3rd, 4th Ionization Energies, etc.

• Source http//chemed.chem.purdue.edu/genchem/topi
  creview/bp/ch7/ie_ea.html

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Ionization Energy

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

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Electronegativity

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

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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.

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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
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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
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Electronegativity
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Electronegativity
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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.

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Electron Affinity

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

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Electron Affinity

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

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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
  

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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

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Credits

• PowerPoint Adela J. Dziekanowski