Molecular geometry

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Molecular geometry

• According to VSEPR, shared pairs of electrons are
  oriencted as far away from each other as possible
• Models AB2, such as BeF2 linear
• AB3, such as BF3 trigonal planar

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Molecular geometry

• AB4, such as CH4 tetrahedral

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Valence Shell Electron Pair Repulsion

• An electron group a lone pair, or a group of
  bonding e- (all the e- in a multiple bond are
  considered as one group).
• Distinct electron groups arranged around a
  central atom repel one another (Coulombs law).
• these repulsive forces result in the arrangement
  of electron groups around the central atom such
  that the distance between the groups is
  maximized.
• for the case of 4 groups the maximum distance is
  created in a tetrahedron (109.5 angles)

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Sample problemsUse VSEPR theory to predict the
molecular geometry of aluminum trichloride,
AlCl3Predict the shape of a molecule of Carbon
DioxidePredict the shape of a chlorate ion,
ClO-3
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Lone pair (and multiple bonds) occupy more space
than a single bonding pair of electrons
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polar vs. non-polar molecules

• A molecule is considered polar if it has a net
  dipole moment. A molecule is non-polar if it
  has no net dipole moment (or if the dipole moment
  is negligibly small)
• Characteristics of a polar molecule
• Polar covalent bond(s)
• Molecular geometry that gives
• rise to a net dipole moment

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Covalent Bonding

• Multiple Bonds
• It is possible for more than one pair of
  electrons to be shared between two atoms
  (multiple bonds)
• One shared pair of electrons single bond (e.g.
  H2)
• Two shared pairs of electrons double bond (e.g.
  O2)
• Three shared pairs of electrons triple bond
  (e.g. N2).
• Generally, bond distances decrease as we move
  from single through double to triple bonds.

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Bond Polarity and Electronegativity

• In a covalent bond, electrons are shared.
• Sharing of electrons to form a covalent bond does
  not imply equal sharing of those electrons.
• There are some covalent bonds in which the
  electrons are located closer to one atom than the
  other.
• Unequal sharing of electrons results in polar
  bonds.

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Bond Polarity and Electronegativity

• Bond polarity helps describe the sharing of
  electrons between atoms.
• Nonpolar covalent bond one in which the
  electrons are shared equally between two atoms
• Polar covalent bond one of the atoms exerts a
  greater attraction for the bonding electrons than
  the other. If the difference in ability to
  attract electrons is large enough, an ionic bond
  is formed.

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Bond Polarity and Electronegativity

• Electronegativity
• Electronegativity The ability of one atoms in a
  molecule to attract electrons to itself. i.e. the
  greater the electronegativity, the greater the
  ability to attract electrons to itself

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Bond Polarity and Electronegativity
Ionization Energy Measures how strongly an atom
holds on to its electrons (section 7.4) Electron
affinity A measure of how strongly an atom
attracts additional electrons (section 7.5) (the
greater the attraction between a given atom and
an added electron, the more negative the atoms
electron affinity will be the greater the
affinity)
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Bond Polarity and Electronegativity
Electron affinities
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Bond Polarity and Electronegativity

• Electronegativity
• Electronegativity The ability of one atoms in a
  molecule to attract electrons to itself. i.e. the
  greater the electronegativity, the greater the
  ability to attract electrons to itself
• Pauling set electronegativities on a scale from
  0.7 (Cs) to 4.0 (F).
• Electronegativity increases
• across a period and
• down a group.

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Bond Polarity and Electronegativity
Electronegativity
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Bond Polarity and Electronegativity

• Electronegativity and Bond Polarity
• Difference in electronegativity is a gauge of
  bond polarity
• electronegativity differences around 0 result in
  non-polar covalent bonds (equal or almost equal
  sharing of electrons)
• electronegativity differences around 2 result in
  polar covalent bonds (unequal sharing of
  electrons)
• electronegativity differences around 3 result in
  ionic bonds (transfer of electrons).

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Bond Polarity and Electronegativity

• Electronegativity and Bond Polarity
• There is no sharp distinction between bonding
  types.
• The positive end (or pole) in a polar bond is
  represented ? and the negative pole ?-.

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Bond Polarity and Electronegativity

• Dipole Moments
• Consider HF
• The difference in electronegativity leads to a
  polar bond.
• There is more electron density on F than on H.
• Since there are two different ends of the
  molecule, we call HF a dipole.
• Dipole moment, m, is the magnitude of the dipole
• where Q is the magnitude of the charges.
• Dipole moments are measured in debyes, D.

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Bond Polarity and Electronegativity

• Bond Types and Nomenclature
• In general, the least electronegative element is
  named first.
• The name of the more electronegative element ends
  in ide.
• Ionic compounds are named according to their
  ions, including the charge on the cation if it is
  variable.
• Molecular compounds are named with prefixes.

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Dipole-dipole forces

• Dipole-dipole forces the forces of attraction
  between polar molecules

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Hydrogen Bonding

• The intermolecular force in which a hydrogen atom
  that is bonded to a highly electronegative atom
  is attracted to an unshared pair of electrons of
  an electronegative atom in a nearby molecule

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Homework due tomorrow

• Problems 43-49, pg 178

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Dipoles do not cancel thus, there is a net
dipole moment
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Dipoles cancel thus, there is no net dipole
moment
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Increasing p character
Increasing s character
sp2
sp3
sp
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sp, sp2, sp3 orbitals describe molecular geometry
in organic chemistry (see Table 3.2)
two p
one p