Molecular Geometry and Bonding Theories (CH.9)

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Molecular Geometry and Bonding Theories (CH.9)

• By Maggie Dang

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9.1 Molecular Shapes

• The overall shape of a molecule is determined by
  its bond angles, the angles made by the lines
  joining the nuclei of the atoms in the molecule
• Molecules with a central atom A surrounded by n
  atoms B, denoted ABn, adopt a number of different
  geometric shapes, depending o n the value of n
  and on the particular atoms involved

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9.2 The VSEPR Model

• Valence-shell electron-pair repulsion (VSEPR)
  model rationalizes molecular geometries in terms
  of the repulsions between electron domains, which
  are regions about a central atom in which
  electrons are likely to be found.
• Bonding pairs of electrons are involved in making
  bonds
• Nonbonding pairs of electrons, also called lone
  pairs, both create electron domains around an atom

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Electron Domains

• Based on the VSEPR model, electron domains orient
  themselves to minimize electrostatic repulsions
  and to remain as far apart as possible
• Electron domains from nonbonding pairs exert
  slightly greater repulsions than those from
  bonding pairs
• Electron domains from multiple bonds exert
  greater repulsions than those from single bonds
• Electron-Domain Geometry arrangement of electron
  domains around a central atom
• Molecular Geometry arrangement of atoms

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Steps to Predicting Molecular Geometries with the
VSEPR Model

• Sketch the Lewis structure of the molecule or ion
• Count the total number of electron domains around
  the central atom, and arrange them in the way
  that minimizes the repulsions among them
• Describe the molecular geometry in terms of the
  angular arrangement of the bonded atoms
• A double or triple bond is counted as one
  electron domain when predicting geometry.
• Ex CO2 has CO double bonds
• When we apply the VSEPR model to CO2, each double
  bond counts as one electron domain. The VSEPR
  model predicts that CO2 is linear.
• Because multiple bonds count as one electron
  domain, the number of electron domain can be
  counted as ( of electron domains) ( of atoms
  bonded to the central atom) ( of nonbonding
  pairs on the central atom)
• Refer to pgs 207-209 for molecular geometry
  tables

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Example

• Using the VSEPR model, predict the molecular
  geometries of O3.

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The Effect of Nonbonding Electrons and Multiple
Bonds on Bond Angles

• Electron domains for non-bonding electron pairs
  exert greater repulsive forces on adjacent
  electron domains and thus tend to compress the
  bond angles
• Electron domains for multiple bonds exert a
  greater repulsive force on adjacent
• electron domains than do single bonds.

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Molecules with Expanded Valence Shells

• These shapes generally contain axial and
  equatorial positions
• When pointing toward an axial position, an
  electron domain is situated 90 from three
  equatorial positions
• In equatorial position an electron domain is
  situated 120 from the other two equatorial
  positions and 90 from the two axial positions
• Repulsions between domains are much greater when
  they are situated 90 from each other than when
  they are at 120.
• Variations of the trigonal bipyramidal shape show
  lone electron pairs in the equatorial position
• Variations of the octahedral shape show lone
  electron pairs in the axial positions

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Molecules with More than One Central Atom

• The VSEPR theory can be used for molecules with
  more than one central atom

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9.3 Polarity of Polyatomic Molecules

• The dipole moment of a polyatomic molecule
  depends on the vector sum of the dipole moment
  due to each individual bond, called the bond
  dipole.
• Certain molecular shapes, such as linear AB2 and
  trigonal planar AB3, assure that the bond dipoles
  cancel, leading to a dipole moment of zero for
  the molecule.
• In other shapes such as bent AB2 and trigonal
  pyramidal AB3, the bond dipoles do not cancel and
  the molecule will have a nonzero dipole moment
  called a polar molecule
• One with a zero dipole moment is called nonpolar.
• Polarity is also used when talking about
  covalently bonded molecules.
• If the molecule has only 2 different atoms, such
  as, HF or CCl4 you can calculate the
  electronegativity difference and determine the
  type of covalent bond (polar or non-polar).

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Example

• Predict whether BrCl is polar or nonpolar.
• Chlorine is more electronegative than bromine.
  Consequently, BrCl will be polar with chlorine
  carrying the partial negative charge.

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Polarity and Bond Type

• Electronegativity Difference

• Bonding Type

• lt0.5
• 0.5 1.9
• gt 2.0

• Non-polar covalent
• Polar covalent
• ionic

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9.4 Covalent Bonding and Orbital Overlap

• Valence-bond theory is an extension of Lewiss
  notion of electron-pair bonds. In valence-bond
  theory, covalent bonds are formed when atomic
  orbitals on neighboring atoms overlap one
  another.
• The overlap region is a favorable one for the two
  electrons because of their attraction to two
  nuclei.
• The greater the overlap between two orbitals, the
  stronger the bond that is formed.

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9.5 Hybrid Orbitals

• To extend the ideas of valence-bond theory to
  polyatomic molecules, it is useful to envision
  the mixing of s,p, and sometimes d orbitals to
  form hybrid orbitals.
• Hybrid orbitals can overlap with orbitals on
  other atoms to make bonds, or they can
  accommodate nonbonding pairs.
• The process of hybridization leads to hybrid
  orbitals that are directed along certain definite
  directions

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sp, sp2, and sp3 Hybrid Orbitals

• One s orbital and one p orbital can hybridize to
  form two equivalent sp hybrid orbitals
• For sp2, using BF3 as an example, a 2s electron
  on the B atom can be promoted to a vacant 2p
  orbital. Mixing the 2s and two of the 2p
  orbitals yields three equivalent sp2 hybrid
  orbitals
• For sp3, using CH4 as an example, it forms four
  equivalent bonds with the four hydrogen atoms.
  This process results from the mixing of the 2s
  and all three 2p atomic orbitals of carbon to
  creat four equivalent sp3 hybrid orbitals.
• See pgs 318-320 for diagrams

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Hybridization Involving d Orbitals

• Atoms in the third period and beyond can use d
  orbitals to form hybrid orbitals. Mixing one s
  orbital, three p orbitals, and one d orbital
  leads to five sp3d hybrid orbitals. These hybrid
  orbitals are directed toward the vertices of a
  trigonal bipyramid.
• Similarly, mixing one s orbital, three p
  orbitals, and two d orbitals gives six sp3d2
  hybrid orbitals, which are directed toward the
  vertices of an octahedron.
• The use of d orbitals in constructing hybrid
  orbitals corresponds to the notion of an expanded
  valence shell.

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Example

• Predict the hybridization of SF4
• - There are five electron domains around S,
  giving rise to the trigonal bipyramidal
  electron-domain geometry. With an expanded octet
  of 10 electrons, the use of a d orbital on the
  sulfur is required. The trigonal bipyramidal
  electron-domain geometry corresponds to SP3 d
  hybridization. One of the hybrid orbitals that
  points in an equatorial direction contains a
  nonbonding pair of electrons the other four are
  used in forming the S-F bonds.

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Steps to Predict Hybrid Orbitals

1. Draw the Lewis structure for the molecule or ion.
2. Determine the electron-domain geometry using the
   VSEPR model.
3. Specify the hybrid orbitals needed to accommodate
   the electron pairs based on their geometric
   arrangement.

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9.6 Multiple Bonds

• Sigma bonds (s) covalent bonds in which the
  electron density lies along the line connecting
  the atoms
• Pi bonds (p) formed from the overlap of p
  orbitals that are oriented perpendicular to the
  internuclear axis
• A double bond, such as that in C2H4, consists of
  one sigma bond and one pi bond.
• A triple bond, such as that in C2H2, is composed
  of one sigma bond and two pi bonds.
• Ex H-H has one sigma bond, N2 has one sigma plus
  two pi bond

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Delocalized Pi Bonding

• Every pair of bonded atoms shares one or more
  pairs of electrons. In every bond at least one
  pair of electrons is localized in the space
  between the atoms, in a sigma bond.
• The electrons in sigma bonds are localized in the
  region between two bonded atoms and do not make a
  significant contribution to the bonding between
  any other two atoms.
• When atoms share more than one pair of electrons,
  the additional pairs are in pi bonds. The
  centers of charge density in a pi bond lie above
  and below the bond axis
• Molecules with two or more resonance structures
  can have pi bonds that extend over more than two
  bonded atoms. Electrons in pi bonds that extend
  over more than two atoms are said to be
  delocalized.
• Delocalized Pi bonds are spread among several
  atoms

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9.7 Molecular Orbitals

• Molecular orbital theory another model used to
  describe the bonding in molecules. In this
  model, the electrons exist in allowed energy
  states call molecular orbitals (MOs).
• A molecular orbital can be spread among all the
  atoms of a molecule, can have a definite energy,
  and can hold two electrons of opposite spin.

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The Hydrogen Molecule

• Whenever two atomic orbitals overlap, two
  moleular orbitals form. Thus, the overlap of the
  1s orbitals of two hydrogen atoms to form H2
  produces two MOs .
• The lower-energy MO of H2 concentrates electron
  density between the two hydrogen nuclei and is
  called the bonding molecular orbital.
• The higher-energy MO has very little electron
  density between the nuclei and is called the
  antibonding molecular orbital.
• The electron density in both the bonding and the
  antibonding molecular orbitals of H2 is centered
  about the internuclear axis. MOs of this type
  are called sigma molecular orbitals.
• The bonding sigma molecular orbital of H2 is
  labeled s1s, indicating that the MO is formed
  from two 1s orbitals.
• The antibonding sigma molecular orbital of H2 is
  labeled s1s, the asterisk denoting that MO is
  antibonding.
• The interaction between two 1s orbitals to form
  s1s and s1s molecular orbitals can be
  represented by an energy-level diagram(molecular
  orbital diagram). It shows the interacting
  atomic orbitals in the left and right columns and
  the MOs in the middle column.

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Bond Order

• The stability of a covalent bond is related to
  its bond order.
• Bond order ½ ( of bonding electrons - of
  antibonding electrons)
• A bond order of 1 represents a single bond, a
  bond order of 2 represents a double bond, and a
  bond order of 3 represents a triple bond.
• Because MO theory also treats molecules with an
  odd number of electrons, bond orders of ½, 3/2,
  or 5/2 are possible.
• Ex H2 has 2 bonding electrons and no antibonding
  ones. It has a bond order of ½(2-0)1.

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Example

• What is the bond order of the O2 ion?
• The O2 ion has eight bonding electrons and three
  antibonding ones. Thus, its bond order is
• Bond order ½ (8-3) 2.5

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9.8 Second-Row Diatomic Molecules

• Second-row atoms have more than one atomic
  orbital
• The way we place electrons in the orbitals
• The number of Mos formed equals the number of
  atomic orbitals combined
• Atomic orbitals combine most effectively with
  other atomic orbitals of similar energy.
• The effectiveness with which 2 atomic orbitals
  combine is proportional to their overlap with one
  another. As the overlap increases, the bonding
  MO is lowered in energy, and the antibonding MO
  is raised in energy.
• Each molecular orbital can accommodate, at most,
  two electrons, with their spins paired (Pauli
  exclusion principle)
• When Mos have the same energy, one electron
  enters each orbital ( with the same spin) before
  spin pairing occurs (Hunds rule)

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Molecular Orbitals for Li2 and Be2

• Core electrons usually do not contribute
  significantly to bonding in molecule formation.

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Molecular Orbitals from 2p Atomic Orbitals

• The p orbitals that point directly at one another
  can form sigma bonding and sigma antibonding MOs.
  
• The p orbitals that are oriented perpendicular to
  the internuclear axis combine to form pi
  molecular orbitals.
• In diatomic molecules, the pi molecular orbitals
  occur as pairs of degenerate (same energy)
  bonding and antibonding MOs.
• The s2p bonding MO is expected to be lower in
  energy (more stable) than the p2p bonding Mos
  because of larger orbital overlap. This ordering
  is reversed in B2, C2, and N2 because of
  interaction between the 2s and 2p atomic orbitals.

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Electron Configurations for B2 Through Ne2

• For B2, C2, and N2, the s2p MO is above the p2p
  molecular orbitals in energy. For O2, F2, and
  Ne2, the s2p MO is below the p2p molecular
  orbitals.

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Electron Configurations and Molecular Properties

• Molecules with 1 or more unpaired electrons are
  attracted into a magnetic field. The more
  unpaired electrons in a species, the stronger the
  force of attraction. This type of magnetic
  behavior is called paramagnetism.
• Substances with no unpaired electrons are weakly
  repelled from a magnetic filed. This property is
  called diamagnetism. It is a much weaker effect
  than paramagnetism.