Chapter 9


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



In the VSEPR model, we focus on one atom at a
time, examining the bonds it forms and any
lone pairs it may possess. ...




The atoms in a molecule that are bonded to each
other share a pair of valence electrons.  Some
molecules also have a...


To report shapes more precisely, we give the
bond angles, which are the angles between
bonds visualized as straight lin...


The VSEPR model does not distinguish
between a single and multiple bonds. A
multiple bond is treated just like other re...


Suggest a shape for the ethyne molecule,






To understand any molecule, one must first
complete a Lewis dot structure. It is then
possible to predict the mol...


The molecule BF3 has a dot symbol as shown
below. Here the B atom has 3 bonded pairs in
its outer shell. Minimizing the...




The molecule NH3 has a dot symbol much like
that for BF3. However, there is a lone pair in the
outer shell of the ce...




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



4 groups of el...


4 groups of electrons (3 bonds around a
central atom, 1 lone pair)
NH3 Trigonal pyramid 



4 groups of electrons (2 b...


4 groups of electrons (1 bond around a central
atom, 3 lone pairs)
HF Linear Molecule 



5 groups of electrons (5 bon...


5 groups of electrons (4 bonds around a
central atom, 1 lone pair)
SF4 See-saw Molecule



5 groups of electrons (3 bo...


5 groups of electrons (2 bonds around a
central atom, 3 lone pairs)
XeF2 Linear Molecule 



6 groups of electrons (6 ...


6 groups of electrons (5 bonds around a
central atom, 1 lone pair)
BrF5 Square Pyramid 



6 groups of electrons
(4 bo...


A covalent bond in which the electron pair is
shared unequally has partial ionic character and is
called a polar covale...




A diatomic molecule is polar if its bond is
polar. A polyatomic molecule is polar if it has
polar bonds arranged in ...
Water is a polar molecule with positive charges
on one side and negative on the other.


Examples of polar molecules of materials that
are gases under standard conditions are:
Ammonia (NH3), Sulfur Dioxide (S...



Page 413
9.52,9.58,9.60,




The strength of a bond between two atoms is
measured by the bond enthalpy . The bond
enthalpy typically increases as...


The first step in predicting the strength of a
bond in polyatomic molecule is to identify the
atoms and the bond order....


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

Decide whether the reaction :
CH3CH2I(g) +H2O- CH3CH2OH(g) +HI(g)

is endothermic or exothermic.




A bond length is the distance between the centers of
two atoms joined by a chemical bond. Bond lengths
affect the ov...


The covalent radius of an atom is the
contribution it makes to the length of the
covalent bond. Covalent radii are adde...


Pi bonds involve the electrons in the leftover
p orbital for each carbon atom. Those p orbitals
are the electron clouds...


The distinction between a sigma bond and a pi
bond is shown below. The sigma bond has
orbital overlap directly between ...





A single bond is a σ- bond.
A double bond is a σ bond plus one π bond.
A triple bond is a σ bond and two π bonds....






Valence Bond theory cannot account for
bonding in polyatomic molecules like CH₄.
The electronic configuration of ...


Without promotion , a carbon atom has only
two unpaired electrons and so can form only
two bonds ; after promotion it h...


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

The four half-filled orbitals have equal energy
and are called sp3 hybrid orbitals. The s has a
superscript of 1 and th...




Hybridization also occurs in boron atoms,
where it forms sp2 orbitals. One electron from
the 2s orbital gets promote...




The sp hybridization, occurs in beryllium. In
this example, one of the two electrons in the 2s
orbital gets 'promote...




This usually occurs when the central atom is an
element in period 3 or later, it uses its d orbital to
expand the va...



sp3d2
The final alignment is octahedral, not because
it has eight outer atoms, it actually has six, but
because it fo...




Paramagnetic substance means that it is
attracted to a magnetic field. It is a property of
unpaired electrons which ...


Lewis theory cannot account for electron
deficient compounds or the paramagnetism of
oxygen. In molecular orbital theor...
Constructive interference produces MO's that places
the bonding electrons between the nuclei of the
bonding atoms. Electro...


Consider two helium atoms approaching. Two
electrons go into the sigma 1s bonding MO,
and the next two into the sigma s...




As antibonding MOs are more antibonding
than bonding MOs are bonding, He2
(dihelium), is not expected to exist
And d...



Page 412
9.50,9.52,9.58
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Ap chapter 9

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Ap chapter 9

  1. 1. Chapter 9
  2. 2.  The shape of a molecule is described by reporting the locations of its atoms.
  3. 3.   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. 
  4. 4.   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. 
  5. 5.  To report shapes more precisely, we give the bond angles, which are the angles between bonds visualized as straight lines joining the atom centers.
  6. 6.  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.
  7. 7.  Suggest a shape for the ethyne molecule,
  8. 8.    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.
  9. 9.  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:
  10. 10.   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).
  11. 11.   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 
  12. 12.  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 
  13. 13.  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 
  14. 14.  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 
  15. 15.  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 
  16. 16.  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
  17. 17.  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.
  18. 18.   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.
  19. 19. Water is a polar molecule with positive charges on one side and negative on the other.
  20. 20.  Examples of polar molecules of materials that are gases under standard conditions are: Ammonia (NH3), Sulfur Dioxide (SO2) and Hydrogen Sulfide (H2S).
  21. 21.   Page 413 9.52,9.58,9.60,
  22. 22.   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.
  23. 23.  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.
  24. 24.  The average bond enthalpies can be used to estimate reaction enthalpies and to predict the stability of a molecule.
  25. 25.  Decide whether the reaction : CH3CH2I(g) +H2O- CH3CH2OH(g) +HI(g) is endothermic or exothermic.
  26. 26.   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.
  27. 27.  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.
  28. 28.  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
  29. 29.  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.
  30. 30.     A single bond is a σ- bond. A double bond is a σ bond plus one π bond. A triple bond is a σ bond and two π 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 σ bonds or side by side to form π bonds.
  31. 31.    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
  32. 32.  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.
  33. 33.  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.
  34. 34.  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.
  35. 35.   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. .
  36. 36.   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.
  37. 37.   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"
  38. 38.   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
  39. 39.   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.
  40. 40.  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.
  41. 41. 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. 
  42. 42.  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
  43. 43.   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.
  44. 44.   Page 412 9.50,9.52,9.58

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