Ch. 10a: Chemical Bonding II: Molecular Shapes

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Ch. 10a Chemical Bonding II Molecular Shapes

• Dr. Namphol Sinkaset
• Chem 200 General Chemistry I

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I. Chapter Outline

• Introduction
• Lewis Structures
• Resonance
• Exceptions
• VSEPR Theory
• Molecular Polarity

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I. Importance of Shape

• In condensed phases (liquids/solids), molecules
  are in close proximity, so they interact
  constantly.
• The 3-D shape of a molecule determines many of
  its physical properties.
• We want to be able to predict 3-D shape starting
  from just a formula of a covalent compound.

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II. Lewis Structures

• The first step to getting the 3-D shape of a
  molecule is getting the correct 2-D structure.
• The 2-D structure will be the basis of our 3-D
  shape assignment.
• A 2-D representation of the bonding in molecule
  is known as a Lewis structure.

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II. Steps for Drawing Lewis Structures

• Determine total of valence e-.
• Place atom w/ lower Group (lower
  electronegativity) as the central atom.
• Attach other atoms to central atom with single
  bonds.
• Fill octet of outer atoms. (Why?)
• Count of e- used so far. Place remaining e- on
  central atom in pairs.
• If necessary, form higher order bonds to satisfy
  octet rule of central atom.
• Allow expanded octet for central atoms from
  Period 3 or lower.

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II. Lewis Structure Practice

• Draw correct Lewis structures for NF3, CO2,
  SeCl2, PF6-, PI5, IF2-, IF6, and H2CO.

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III. Multiple Valid Lewis Structures

• Sometimes more than one Lewis structure can be
  drawn for the same molecule.
• For example, ozone (O3).

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III. Resonance Forms

• Resonance forms are also known as resonance
  structures.
• Resonance forms have the same relative placement
  of atoms, but different locations of bonding and
  lone e- pairs.

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III. Resonance Hybrid

• Neither resonance form is a true picture of the
  molecule.
• The molecule exists as a resonance hybrid, which
  is an average of all resonance forms.
• In a resonance hybrid, e- are delocalized over
  the entire molecule.

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III. Sample Problem

• Draw the resonance forms of the carbonate anion.

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III. Important Resonance Forms

• If all resonance forms have the same surrounding
  atoms, then each contributes equally to the
  resonance hybrid.
• If this is not the case, then one or more
  resonance forms will dominate the resonance
  hybrid.
• How can we determine which forms will dominate?

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III. Formal Charge

• formal charge the charge an atom would have if
  bonding e- were shared equally

formal charge ( valence e-) (unshared e- ½
shared e-)
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III. Formal Charges in O3

• We calculate formal charge for each atom in the
  molecule.
• For oxygen atom A (on the right), there are 6
  valence e-, 4 unshared e-, and 4 shared e-. The
  formal charge for this O atom is 0.
• NOTE sum of all formal charges must equal the
  overall charge of the molecule!

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III. Using Formal Charges

• Formal charges help us decide the most important
  resonance forms when we consider to the following
  guidelines
• Smaller f.c.s are better than larger f.c.s.
• Same sign f.c.s on adjacent atoms is
  undesirable.
• Electronegative atoms should carry higher
  negative f.c.s.

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III. Sample Problem

• Find the dominant resonance structures for the
  sulfate anion.

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IV. Exceptions to the Octet Rule

• Weve already discussed expanded valence cases,
  but there are other exceptions as well.
• e- deficient atoms like Be and B, e.g. BeCl2 and
  BF3.
• Compounds w/ odd of e-s free radicals.
  Examples include NO and NO2.
• Expanded valence when d orbitals are used to
  accommodate more than an octet.

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V. VSEPR Theory

• From a correct Lewis structure, we can get to the
  3-D shape using this theory.
• VSEPR stands for valence shell electron pair
  repulsion.
• The theory is based on the idea that e- pairs
  want to get as far away from each other as
  possible!

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V. VSEPR Categories

• There are 5 categories from which all molecular
  shapes derive.

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V. Drawing w/ Perspective

• We use the conventions below to depict a 3-D
  object on a 2-D surface.

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V. Determining 3-D Shape

• The 5 categories are a starting point.
• To determine the 3-D shape of a molecule, we
  consider the of atoms and the of e- pairs
  that are associated w/ the central atom.
• All the possibilities for molecular geometry can
  be listed in a classification chart.

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V. Linear/Trigonal Planar Geometries

• First, we have the linear and trigonal planar
  categories.

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V. Tetrahedral Geometries
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V. Trigonal Bipyramidal Geometries
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V. Octahedral Geometries
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V. Steps to Determine Molecular Shape

• Draw Lewis structure.
• Count of bonds and lone pair e-s on the
  central atom.
• Select geometric category.
• Place e-s and atoms that lead to most stable
  arrangement (minimize e- repulsions).
• Determine 3-D shape.

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V. Trig Bipy is Special

• In other categories, all positions are
  equivalent.
• In trig bipy, lone pairs always choose to go
  equatorial first.
• Why?

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V. Distortion of Angles

• Lone pair e-s take up a lot of room, and they
  distort the optimum angles seen in the geometric
  categories.

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V. Some Practice

• Draw the molecular shapes for SF4, BeCl2, ClO2-,
  TeF5-, ClF3, NF3.

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VI. Molecular Polarity

• Individual bonds tend to be polar, but that
  doesnt mean that a molecule will be polar
  overall.
• To determine molecular polarity, you need to
  consider the 3-D shape and see if polarity arrows
  cancel or not.

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VI. Sample Problem

• Determine the molecular geometry of IF2- and
  state whether it is polar or nonpolar.