Molecular Structure

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Molecular Structure

• Chapter 9

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VSEPR MODEL

• The shape of a molecule is described by reporting
  the locations of its atoms.

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• In the VSEPR model, we focus on one atom at a
  time, examining the bonds it forms and any lone
  pairs it may possess. In order to understand and
  study the properties of molecules, we must be
  able to recognize and understand the orientation
  of atoms within a molecule. 

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• The atoms in a molecule that are bonded to each
  other share a pair of valence electrons.  Some
  molecules also have atoms with nonbonding
  electron pairs.  Electron pairs repel each other,
  and they want to take up a position in space that
  will minimize their interactions with other
  electron pairs, bonded or non bonded. 
•   The goal of the VSEPR model is to arrange the
  electron pairs around the central atom so that
  there is the least amount of repulsions among
  them.  This occurs when the electron pairs are as
  far away from each other as possible. 

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• To report shapes more precisely, we give the bond
  angles, which are the angles between bonds
  visualized as straight lines joining the atom
  centers.

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Molecules with multiple bonds

• The VSEPR model does not distinguish between a
  single and multiple bonds. A multiple bond is
  treated just like other region of high electron
  concentration.

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Class Practice

• Suggest a shape for the ethyne molecule,

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• To understand any molecule, one must first
  complete a Lewis dot structure. It is then
  possible to predict the molecular shape using Two
  basic Principles
• 1. The shapes of molecules are determined by the
  repulsion between electron pairs in the outer
  shell of the central atom. Both bond pairs
  (electron pairs shared by two atoms) and lone
  pairs (those located on a central atom but not
  shared) must be considered.
• 2. Lone pairs repel more than bond pairs.

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• The molecule BF3 has a dot symbol as shown below.
  Here the B atom has 3 bonded pairs in its outer
  shell. Minimizing the repulsion causes this
  molecule to have a trigonal planar shape, with
  the F atoms forming an equilateral triangle about
  the B atom. The F-B-F bond angles are all 120,
  and all the atoms are in the same plane. Central
  atoms with octet configurations

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• The molecule NH3 has a dot symbol much like that
  for BF3. However, there is a lone pair in the
  outer shell of the central N atom. 
• In NH3 the N has 3 bond pairs and 1 lone pair,
  (4 total pairs). The shape is called trigonal
  pyramidal (approximately tetrahedral minus one
  atom).

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• 3 groups of electrons (3 bonds around a central
  atom, no lone pairs)
• BF3 Trigonal Planar
• Molecule
• 4 groups of electrons (4 bonds around a central
  atom, no lone pairs) CF4 Tetrahedral
  
• Molecule 

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• 4 groups of electrons (3 bonds around a central
  atom, 1 lone pair) NH3 Trigonal pyramid 
• 4 groups of electrons (2 bonds around a central
  atom, 2 lone pairs) H2O Bent Molecule 

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• 4 groups of electrons (1 bond around a central
  atom, 3 lone pairs) HF Linear Molecule 
• 5 groups of electrons (5 bonds around a central
  atom, no lone pairs) PF5 Trigonal Bipyramid 

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• 5 groups of electrons (4 bonds around a central
  atom, 1 lone pair) SF4 See-saw Molecule
• 5 groups of electrons (3 bonds around a central
  atom, 2 lone pairs) ClF3 T-shaped Molecule 

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• 5 groups of electrons (2 bonds around a central
  atom, 3 lone pairs) XeF2 Linear Molecule 
• 6 groups of electrons (6 bonds around a central
  atom, no lone pairs) SF6 Octahedral Molecule 

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• 6 groups of electrons (5 bonds around a central
  atom, 1 lone pair) BrF5 Square Pyramid 
• 6 groups of electrons
• (4 bonds around a central
• atom, 2 lone pairs) XeF4 Square Planar

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Polar bonds

• A covalent bond in which the electron pair is
  shared unequally has partial ionic character and
  is called a polar covalent bond.
• A polar covalent bond is a bond between two atoms
  that have partial electric charges arising from
  their difference in electronegativity. Partial
  charges give rise to an electric dipole moment.

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Polar molecules

• A diatomic molecule is polar if its bond is
  polar. A polyatomic molecule is polar if it has
  polar bonds arranged in space in such a way that
  their dipoles do not cancel.
• Water is a polar molecule because of the way the
  atoms bind in the molecule such that there are
  excess electrons on the oxygen side and a lack or
  excess of positive charges on the Hydrogen side
  of the molecule.

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Water is a polar molecule with positive
chargeson one side and negative on the other.
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• Examples of polar molecules of materials that are
  gases under standard conditions are Ammonia
  (NH3), Sulfur Dioxide (SO2) and Hydrogen Sulfide
  (H2S).

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Home work

• Page 413
• 9.52,9.58,9.60,

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Bond Strengths in diatomic molecules

• The strength of a bond between two atoms is
  measured by the bond enthalpy . The bond enthalpy
  typically increases as the order of the bond
  increases, decreases as the number of lone pairs
  on neighboring atoms increases, and decreases as
  the atomic radius increases.
• The greater the bond strength the harder it will
  be to break the bond.

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Bond strengths in polyatomic molecules

• The first step in predicting the strength of a
  bond in polyatomic molecule is to identify the
  atoms and the bond order. An electronegative
  atom can pull electrons toward itself from more
  distant parts of the molecule. So all the atoms
  in the molecule experience a slight pull. Because
  these variations in strength are not very great,
  the average bond enthalpy ?HB is a guide to the
  strength of a bond in any molecule.

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• The average bond enthalpies can be used to
  estimate reaction enthalpies and to predict the
  stability of a molecule.

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Class Practice

• Decide whether the reaction
• CH3CH2I(g) H2O-? CH3CH2OH(g) HI(g)
• is endothermic or exothermic.

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Bond lengths

• A bond length is the distance between the centers
  of two atoms joined by a chemical bond. Bond
  lengths affect the overall size of a molecule.
• Bond length is directly related to bond order,
  when more electrons participate in bond formation
  the bond will get shorter. Bond length is also
  inversely related to bond strength and the bond
• dissociation energy, as a stronger bond is also
  a shorter bond. In a bond between two identical
  atoms half the bond distance is equal to the
  covalent radius.

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• The covalent radius of an atom is the
  contribution it makes to the length of the
  covalent bond. Covalent radii are added together
  to estimate the lengths of bonds in molecules.

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Sigma and Pi bonds

• Pi bonds involve the electrons in the leftover
• p orbital for each carbon atom. Those p orbitals
  are the electron clouds or orbitals that are
  shown going up above and below each carbon atom

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• The distinction between a sigma bond and a pi
  bond is shown below. The sigma bond has orbital
  overlap directly between the two nuclei. The pi
  bond has orbital overlap off to the sides of the
  line joining the two nuclei.

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• A single bond is a s- bond.
• A double bond is a s bond plus one p bond.
• A triple bond is a s bond and two p bonds.
• In valence bond theory a bond forms when unpaired
  electrons in valence shell atomic orbitals
• on two atoms pair. The atomic orbitals they
  occupy overlap end to end to form s bonds or side
  by side to form p bonds.

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Hybridization

• Valence Bond theory cannot account for bonding in
  polyatomic molecules like CH4.
• The electronic configuration of carbon is
• He 2s²2px¹ 2py¹.The 2s orbital is filled in
  the ground state of the carbon atom. Consider
  forming hybrids by mixing the 2s orbital with the
  2pz orbital.
• The configuration of carbon would be He 2s¹2px¹
  2py¹2pz¹

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• Without promotion , a carbon atom has only two
  unpaired electrons and so can form only two bonds
  after promotion it has four unpaired electrons
  and can form four bonds. Each bond releases
  energy as it forms. Despite the energy cost of
  promoting the electron, the overall energy of
  methane molecule is lower than it would be if
  carbon formed only two
• C-H bonds.

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• The carbon in methane is now sp³ hybridised. Four
  hydrogen should form 4 bonds, one with the 2s
  orbital of carbon and three with the 2p orbitals
  in carbon.

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• The four half-filled orbitals have equal energy
  and are called sp3 hybrid orbitals. The s has a
  superscript of 1 and the p has a 3, thus meaning
  that there are four sp3 orbitals, all with the
  same amount of energy. The hydrogen atoms can now
  bond to these orbitals. This is how methane
  forms.

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• Hybridization also occurs in boron atoms, where
  it forms sp2 orbitals. One electron from the 2s
  orbital gets promoted to an empty 2p orbital,
  thus forming three sp2 hybrid orbitals.
• .

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• The sp hybridization, occurs in beryllium. In
  this example, one of the two electrons in the 2s
  orbital gets 'promoted' to an empty 2p orbital,
  thus forming two sp hybrid orbitals.

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Hybrids including d-orbitals

• This usually occurs when the central atom is an
  element in period 3 or later, it uses its d
  orbital to expand the valence shell to
  accommodate more than 4 electron pairs.
• The atom now uses one d orbital in addition to
  four s and p orbitals of the valence shell. The
  resulting orbital is called as sp³d hybrid
  orbital. This alignment is Trigonal Bipyramidal
  because the five outer atoms form two pyramids
  (one on top, the other upside down on the bottom)
  around the central atom. In this model, all atoms
  are 90 degrees or 120 degrees apart and have the
  orbital hybridization of "sp3d"

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• sp3d2
• The final alignment is octahedral, not because it
  has eight outer atoms, it actually has six, but
  because it forms eight 90 degree angles. The
  orbital hybridization is "sp3d2

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Paramagnetic and dimagnetic substances

• Paramagnetic substance means that it is attracted
  to a magnetic field. It is a property of unpaired
  electrons which act as tiny magnets. Paramagnetic
  substances tend to move into a magnetic field
• Most substances contain fully paired electrons
  and are repelled by a magnetic field such
  substances are diamagnetic.

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Limitations of Lewis Theory

• Lewis theory cannot account for electron
  deficient compounds or the paramagnetism of
  oxygen. In molecular orbital theory, electrons
  occupy orbitals that spread throughout the
  molecule. The Pauli exclusion principle allows no
  more than two electrons to occupy each molecular
  orbital.

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Bonding and antibonding orbitals

• Constructive interference produces MO's that
  places the bonding electrons between the nuclei
  of the bonding atoms. Electrons are
  simultaneously attracted by both nuclei which
  lowers their energies. The MO's produced are
  called bonding orbitals.
• Destructive interference produces MO's that
  places the electrons away from the space between
  the nuclei. The nuclei repel which makes the bond
  less stable (i.e. higher energy). These MO's are
  called antibonding orbitals.

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Molecular orbitals in period 2 diatomic molecules

• Consider two helium atoms approaching. Two
  electrons go into the sigma 1s bonding MO, and
  the next two into the sigma star antibonding MO.
  Bond order(of electrons in bonding orbital-of
  electrons in antibonding orbitals)/2

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• As antibonding MOs are more antibonding than
  bonding MOs are bonding, He2 (dihelium), is not
  expected to exist
• And dihelium has a bond order of zero.

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Home work

• Page 412
• 9.50,9.52,9.58