3.06B and 3.07

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3.06B and 3.07

• Concepts

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3.06 B Lewis Structures

• Just as we represented valence electrons around a
  symbol of an element in a Lewis structure of an
  individual atom, we can also represent bonded and
  non-bonded valence electrons around the atoms
  within a molecule.
• Ex

N
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3.06 B Lewis Structures

• In a Lewis dot, or electron dot, structure of a
  covalent compound, chemists usually use a
  straight line to represent the two electrons
  shared in a covalent bond.

N
N

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3.06 B Lewis Structures

• The other valence electrons not involved in
  bonding are represented by dots around the symbol
  of the element. These other valence electrons are
  called unshared, or lone pair, electrons.

N
N

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3.06 B Steps for drawing Lewis Stuctures

• Place the least electronegative element in the
  center of the molecule. 
• The central atom is usually the first element
  written in the molecular formula, except H, which
  cannot be a central atom because it only bonds
  once and then its valence is full.

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3.06 B Steps for drawing Lewis Stuctures

• Calculate the total number of valence electrons.
• Add up the total number of valence electrons
  thatshould be in the picture according to the
  periodic table. Make a list of the number of
  valence electrons (determined by the location on
  the periodic table) for each atom in the
  molecule, and then add them together. We will
  call this the reality number because this is
  the number of valence electrons that must be
  present in the molecule.

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3.06 B Steps for drawing Lewis Stuctures

• Write the skeleton structure.
• Attach all other atoms to the center with single
  bonds. You may change this later, but we know
  that each atom in the molecule must be attached
  with at least one set of shared electrons (a
  single covalent bond).

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3.06 B Steps for drawing Lewis Stuctures

• Complete the valence electrons.
• Fill every atoms valence by adding electron
  pairs until they are all full. Remember that most
  atoms need eight valence electrons to be full,
  except H, which is full with two electrons. Also,
  remember that a single bond (each line drawn in
  the model) represents two electrons being shared
  by both atoms involved in the bond. This means
  that one line counts as two electrons in the
  valence of each atom touching that line.

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3.06 B Steps for drawing Lewis Stuctures

• Tally the totals.
• Count up the number of electrons represented in
  your drawing. Remember that each line represents
  two shared electrons and each pair of dots
  represents two unshared electrons when you are
  counting the electrons in the drawing. We will
  call this counted number the picture number in
  our examples.

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3.06 B Steps for drawing Lewis Stuctures

• Perform a comparison.
• Compare the picture number to the reality number
  to see if you need to change your Lewis structure
  drawing.
• If picture Reality , then the Lewis
  structure drawing is a good representation of the
  molecule and you do not need to change anything.
• If picture gt Reality , you must fix it by
  adding multiple bonds where appropriate in order
  to reduce the Picture to equal Reality .

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3.06 B Steps for drawing Lewis Stuctures

• Distribute.
• Distribute electrons to atoms surrounding the
  central atom to satisfy octet rule
• Atoms that form multiple bonds are C, N, O, S.
  Oxygen atoms do not bond to each other (except in
  O2 O3, H2O2, peroxides, superoxides).
• All atoms must have eight electrons (octet rule),
  except hydrogen whose octet is two.

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3.06 B Practice

• Share your desktop and show the 3.06 Lewis
  Structures examples and practice.

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3.06 B What to do!

• Share your desktop and show the 3.06B Lab and
  where to complete it and submit it.

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3.07 Intermolecular Forces

• There are also forces of attraction between
  separate molecules, called intermolecular forces.
  The prefix inter- comes from the Latin stem
  meaning between, such as in words like
  Internet, interface, and international.
  Intermolecular forces, sometimes called van der
  Waals forces, vary in strength, but they are
  generally weaker than the ionic and covalent
  bonds found within compounds.

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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Molecular Geometry
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3.07 Predicting Polarity

• Share your desktop and show the steps and
  examples.

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3.07 Electronegativity and polarity

• The difference in electronegativity values will
  affect the polarity (dipole moment) of the
  molecule.

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3.07 Dipoles

•  Many molecules have dipole moments due to
  non-uniform distributions of positive and
  negative charges on the various atoms. Partial
  charges are denoted as d (delta plus)
  and d- (delta minus).
• There are three types of dipoles
• Permanent dipoles
• Instantaneous dipoles
• Induced dipoles

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3.07 Dipoles

• Instantaneous dipoles These occur due to chance
  when electrons happen to be more concentrated in
  one place than another in a molecule, creating a
  temporary dipole. A molecule is polarized when it
  carries an instantaneous or an induced dipole.

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3.07 Dipoles

• Induced dipoles These can occur when one
  molecule with a permanent dipole repels another
  molecule's electrons, "inducing" a dipole moment
  in that molecule temporarily.

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3.07 Intermolecular forces

• Intermolecular forces are attractive forces
  between molecules. These forces exist between
  molecules when they are sufficiently close to
  each other. They are responsible for the
  non-ideal behavior of gases and for properties of
  matter such as boiling point and melting point.
• There are four different types of intermolecular
  forces, depending on the polarity of the
  molecules involved.

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3.07 Intermolecular forces

• In order of increasing strength these inter
  molecular forces are
• London Dispersion
• Dipole-Dipole
• Hydrogen Bonding
• Ion-Dipole
• Share you desktop and show the different forces.

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3.07 What to do!

• Share your desktop and show the 3.07 Lab and
  where to complete it and submit it.