Charge distribution in molecules

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Bonding Molecular Structure Fundamental
Concepts

• Charge distribution in molecules
• Lewis dot structures may lead one to believe that
  the distribution of electrons in a molecule is
  fairly even, further examination reveals this is
  not quite so
• Some atoms will have a slight negative charge and
  other atoms may have a slight positive charge in
  the same molecule
• Two reasons for this
• The formal charges may be different
• Some atoms contribute more electrons to a bond
  than they get back
• The bonding electrons shared between atoms may be
  attracted to one of the atoms more than the other

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Bonding Molecular Structure Fundamental
Concepts

• Charge distribution in molecules
• The result is that some atoms have partial
  charges ( or -) and the way these charges are
  organized in a molecule is its charge
  distribution
• If a molecule has a partial positive charge at
  one end and a partial negative charge at the
  other end, in the condensed phases, the molecules
  will line up so that the partial positive charge
  in one molecule will be attracted to the partial
  negative charge in another molecule
• The strength of these intermolecular interactions
  affect such properties as boiling temperature
  or melting temperature

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Bonding Molecular Structure Fundamental
Concepts

• Charge distribution in molecules
• If a molecule or ion has a negative partial
  charge on one of its atoms, a cation would
  probably attempt to associate with the negatively
  charged atom
• Examples

An H ion will be more strongly attracted to O
atom then the N atom
4
Bonding Molecular Structure Fundamental
Concepts

• Bond polarity and electronegativity
• The electrons in the bonds between the same kind
  of atoms are shared equally
• The electrons in the bonds between the different
  atoms are not shared equally
• One of the atoms will attract electrons to itself
  better than the other atom
• This produces a polar covalent bond
• The bonded atoms will each have residual partial
  charges

• The observation that atoms in a molecule either
  attract electrons better than another
  atom in a molecule is consistent with our
  observations regarding electron affinities and
  ionization energies
• The bond has a end and a - end producing a
  dipole moment giving rise to a polar bond

• The symbols d and d- are written near the atoms
  where the respective partial charges lie

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Bonding Molecular Structure Fundamental
Concepts

• Bond polarity and electronegativity
• Linus Pauling proposed the concept of
  electronegativity to quantitate the ability of
  an atom in a molecule to attract electrons to
  itself.
• Values of electronegativity, c, are given on the
  next slide and are also given on the VWR
  periodic table
• F has the largest electronegativity, c4.0
• Cs and Fr have the smallest, c0.8
• Electronegativities increase from left to right
  in the periodic table
• Electronegativities decrease from top to bottom
  in the periodic table
• Metals have low electronegativity, c lt 2
• Metalloids have electronegativities around 2

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Bonding Molecular Structure Fundamental
Concepts

• Bond polarity and electronegativity
• Ionic compounds have large differences in
  electronegativities
• For CsF, Dc 4.0-0.8 3.2
• For NaCl, Dc 3.0-1.0 2.0
• Covalent compounds have low differences in
  electronegativities
• For HCl, Dc 3.0-2.1 0.9
• For CH4, Dc 2.5-2.1 0.4
• For covalent compounds, the greater the Dc, the
  greater the bond polarity
• HF, Dc1.9 HCl, Dc 0.9 HBr, Dc 0.7 HI, Dc
  0.4

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Bonding Molecular Structure Fundamental
Concepts

• Combining formal charge and electronegativity
• It is reasonable that the electronegative atoms
  in a molecule should take on more negative
  charge than nonelectronegative atoms
• Example BF4- ion

• According to formal charge considerations, the
  formal charge of B is -1 and each F is 0
• This is not reasonable the Fs are more
  electronegative than the B
• The electroneutrality principle says that the
  electrons in a molecule are spread out in such a
  way that the charges on the atoms are as close
  to zero as possible

• The negative charge must be spread out over the F
  atoms and not on B

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Bonding Molecular Structure Fundamental
Concepts

• Combining formal charge and electronegativity
• Example OCN-
• The protonated form of this ion would be HOCN
• Check out Exercise 9.13 and look at the answer. I
  believe there is a better answer to this problem
  than given in the back of the book. What is it?

10
Molecular Geometry

• Molecular shape is determined by the number of
  atoms in a molecule, how the atoms are
  connected, the separation between atoms and the
  number of non-bonded electron pairs on those
  atoms.
• These factors determine the arrangement of the
  atoms in three space that make up molecules.
• Lewis dot structures do not generally indicate
  molecular geometry only how the atoms in a
  molecule are connected and indicates something
  about the nature of the chemical bonds between
  those connected atoms.

11
Molecular Geometry

• VSEPR theory Valence Shell Electron Pair
  Repulsion theory
• The basis of this theory comes from the Pauli
  Exclusion Principle
• In simple terms electron pairs on an atom,
  whether they are bonding or non-bonding, will
  arrange themselves so as to minimize
  electrostatic repulsion.
• 2 electron pairs linear
• 3 electron pairs trigonal planar
• 4 electron pairs tetrahedral
• 5 electron pairs trigonal bipyramidal
• 6 electron pairs octaheral
• Figs. 9.11 9.12, p 414, Kotz and Treichel,
  shows the arrangements of electron pairs with
  formula ABn, n2 to 6

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Molecular Geometry

• Predicting Molecular Geometries
• CH4 NH3
• Both molecules have 4 electron pairs attached to
  the central atom
• The arrangement of these 4 regions of electron
  density is such that they point to the corners
  of a regular tetrahedron
• 2 electron pairs - linear, 180o
• 3 electron pairs - trigonal planar, 120o
• 4 electron pairs - tetrahedral, 109.5o
• 5 electron pairs - trigonal bipyramidal, 90o and
  120o
• 6 electron pairs - octahedral, 90o
• The molecular geometry gives the arrangement of
  the atoms in three space but does not include
  non-bonded electron pairs.
• The non-bonded electron pairs are important in
  determining the molecular geometry.

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Molecular Geometry

• For CH4 and NH3, both will have a tetrahedral
  electron-pair geometry
• However, CH4 has a tetrahedral molecular geometry
  whereas NH3 has a trigonal pyramidal molecular
  geometry.

15
Molecular Geometry

• Steps in using VSEPR theory to predict
  structures
• Draw the Lewis dot structure for the molecule
• Count the number of electron pairs - both bonding
  and non-bonding - attached to the central atom.
• Arrange the atoms and non-bonding electron pairs
  so as to minimize electron-pair repulsions.
• Describe the molecular geometry in terms of the
  arrangement of the atoms attached to the
  central atom. The non-bonded electron pairs are
  not counted in stating the molecular geometry,
  even though they are important in determining
  the molecular geometry.
• Multiple bonds have the same effect as that of an
  electron pair bond in determining the basic
  molecular geometry.
• HCN
• Its really the number of regions of electron
  density on the central atom that is important
  in VSEPR theory.

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Molecular Geometry
Electron pair geometries and molecular shapes
e- e-Pair Bonding nonbonding
Molecular Pairs Geometry Pairs
Pairs Type Geometry
Example 2 2
0 AX2E0 linear

linear 3 trigonal
3 0 AX3E0 trigonal
planar
planar
2 1 AX2E bent
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Molecular Geometry
Electron pair geometries and molecular shapes
e- e-Pair Bonding nonbonding
Molecular Pairs Geometry Pairs
Pairs Type Geometry Example
4 tetrahedral 4
0 AX4E0 tetrahedral

3 1
AX3E1 trigonal pryamidal
2
2 AX2E2 bent
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Molecular Geometry
Electron pair geometries and molecular shapes
e- e-Pair Bonding nonbonding
Molecular Pairs Geometry Pairs
Pairs Type Geometry Example
5 trigonal 5
0 AX5E0 trigonal
bipyramidal bipyramidal

4
1 AX4E1 seesaw
3
2 AX3E2 T-shaped
2
3 AX2E3 linear
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Molecular Geometry
Electron pair geometries and molecular shapes
e- e-Pair Bonding nonbonding
Molecular Pairs Geometry Pairs
Pairs Type Geometry Example
6 octahedral 6
0 AX6 octahedral
5
1 AX5E1 square
pyramidal
4 2 AX4E2
square planar