1. VSEPR theory Molecular Polarity

2. Taste
The taste of a food: Interaction between food molecules and taste cells
Factors: Shape of the molecule and charge distribution within the molecule
Food or “spicy” molecule fit snugly into the active site of specialized proteins on the surface of taste cells
When this happens, changes in the protein structure cause a nerve signal
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3. Sweet Taste: Key/Lock model
Sugar molecules (“Key”) fit into the active site of taste cell receptors (“Lock”):
When the sugar molecule enters the active site, parts of the taste cell receptor split apart
ion channels in the cell membrane to open
resulting in nerve signal transmission
Artificial sweeteners also fit into the same receptor, sometimes binding even stronger than sugar (making them “sweeter” than sugar)
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4. Structure Determines Properties!
Properties of molecular substances depend on the structure of the molecule
The structure includes many factors, such as:
the skeletal arrangement of the atoms
the kind of bonding between the atoms
ionic, polar covalent, or covalent
the shape of the molecule
Bonding theory should allow you to predict the shapes of molecules
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5. Molecular Geometry Molecules: 3-dimensional objects
Molecular Geometry: Shape of a molecule
Two important factors in Molecular Geometry:
Bond Angle
Bond Length
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6. Lewis Theory Predicts Electron Groups
Lewis theory predicts there are regions of electrons in an atom
Electron groups (= VSEPR group): Bonding pairs and Lone pairs
Bonding pairs: Shared pairs of valence electrons between bonding nuclei
Nonbonding pairs (= Lone Pairs): Unshared valence electrons on a single nuclei
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7. O N •
VSEPR Theory
Electron groups around the central atom repel each other (same “-” charge)
Electron groups are spaced away to minimize the repulsion –valence shell electron pair repulsion theory (VSEPR theory)
VSEPR theory can predict the shapes and bond angles in the molecule
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8. Identify Electron Groups
VSEPR groups: Regions of Valence Electrons
Each lone pair of electrons constitutes one electron group on a central atom
Each bonding region constitutes one electron group on a central atom
regardless of whether it is single, double, or triple O N •
____ electron groups on N
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9. Electron Group Geometry
#Electron groups on central atom = 2, 3, 4, 5, 6
5 basic arrangements of electron groups around a central atom
Rarely >6, like IF7
e-group arrangements leads to different geometries
Lone pair causes slight different geometry than bonding pair
Resonance does NOT affect the electron geometry
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10. Linear Electron Geometry
Two electron groups around the central atom: electron groups occupy positions on opposite sides of the central atom
Linear geometry
Bond angle = 180°
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11. Trigonal Planar Electron Geometry
Three electron groups around the central atom: electron groups taking a trigonal planar geometry
bond angle = 120°
Examples: COCl2, CH2O, SO3, BF3
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12. Trigonal Geometry
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13. 3D Animation: Three Electron Groups on Central Atom (bond angle  _________)
Three terminal atoms
(no lone pair electrons)
Two terminal atoms
(ONE lone pair electrons)

14. Tetrahedral Electron Geometry
Four electron groups around the central atom: Tetrahedral geometry
bond angle = 109.5°
Examples: CF4, SiCl4, CH2Cl2, CHCl3
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15. Tetrahedral Geometry
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16. Non-Octet: Trigonal Bipyramidal Electron Geometry
5 electron groups around the central atom: Trigonal Bipyramidal geometry
Axial positions are above and below the central atom
Equatorial positions are in the same base plane as the central atom
<Eq-Central-Eq = 120°
<Ax-Central-Eq = 90°
Examples: SF6, PF6-, SiF62-
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17. Trigonal Bipyramid
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18. Non-Octet: Octahedral Electron Geometry
6 electron groups around the central atom: Two square-base pyramids that are base-to-base with the central atom in the center of the shared bases: Octahedral geometry
eight sides
All positions are equivalent
The bond angle is 90°
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19. Octahedral Geometry
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20. Molecular Geometry vs. Electron Geometry
The actual geometry of the molecule may be different from the electron geometry when:
The electron groups are attached to atoms of different size, or when the bonding to one atom is different than the bonding to another
Lone pairs occupy more space on the central atom than bonding pairs
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21. Not Quite Perfect Geometry
Because the bonds and atom sizes are not identical in formaldehyde, the angles are slightly off
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22. The Effect of Lone Pairs
<H-O-H in water = 104.5° < 109.5°, why?
Lone pair (LP) groups “occupy more space” bonding pair (BP).
electron density in LP is on the central atom only, not shared like bonding electron groups
Ranking the repulsive force interactions:
LP-LP > LP-BP > BP-BP
Stronger repulsive force from LP  ____ angles from LP  _____ BP-BP angle
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23. Repulsive force: BP vs. LP
BP electrons are shared by two atoms
Negative charge in BP partially “neutralized” by nuclear charge
less repulsive
LP electrons are localized on the central atom
 negative charge from LP takes more space
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24. Bond Angle Distortion from Lone Pairs
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25. Bond Angle Distortion from Lone Pairs
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26. Bent Molecular Geometry: Derivative of Trigonal Planar
3 electron groups around the central atom (2 BP + 1 LP): trigonal planar — bent shape
The bond angle is less than 120°
because the lone pair takes up more space, “pushing” bonding electrons closer to each other.
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27. Molecular Geometry: Pyramidal & Bent
Tetrahedral Electron Geometry (1 LP + 3 BPs): pyramidal shape
Larger LP-BP repulsion  Bond angle ___109.5°
Example: Ammonia, PCl3
Tetrahedral Electron Geometry (2 BPs + 2 LPs): tetrahedral—bent shape
it is planar
Bond angle ____ 109.5°
Example: SCl2 , H2O
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28. Pyramidal Shape
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29. Pyramidal Shape
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30. Tetrahedral–Bent Shape
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31. Tetrahedral–Bent Shape
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32. 3D Animation: Four Electron Groups on Central Atom (bond angle  _________)
Three Terminal Atoms
(One lone pair electrons)
Four Terminal Atoms
(no lone pair electrons)
Two Terminal Atoms
(Two lone pair electrons)

33. Non-Octet: Trigonal Bipyramidal
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34. Lone Pair in Trigonal Bipyramidal Electron Geometry
Lone Pair(s) has priority to occupy the equatorial positions Because larger angle (<eq-Cen-eq =120°) gives less repulsion than LP at axial position (90°)
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35. Derivatives of the Trigonal Bipyramidal Electron Geometry
Since Lone Pair(s) occupy equatorial positions
1 LP: seesaw shape
aka distorted tetrahedron
2 LP: T-shaped
3 LP: linear shape
Note: Distortion due to LP
<Eq-Central-Eq < 120°
<Axial-Central-Eq < 90°
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36. 1x LP in Trigonal Bipyrimidal e-Geometry: Seesaw Shape
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37. 2x LPs in Trigonal Bipyrimidal e-Geometry: T–Shape
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38. 3x LPs in Trigonal Bipyrimidal e-Geometry: Linear Shape
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39. Non-Octet: OctahedralS F
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40. 1x LP Derivatives of the Octahedral Geometry: square pyramid shape
the bond angles between axial and equatorial positions is less than 90°

41. 2x LP Derivatives of the Octahedral Geometry: square pyramid shape
When more than ONE LP in Octahedral electron groups geometry, LPs are positioned opposite to each other (LP-LP repulsion is the strongest)
the bond angles between equatorial positions is 90°

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43. Predicting the Shapes Around Central Atoms
1. Draw the Lewis structure
2. Determine the number of electron groups around the central atom
3. Classify each electron group as bonding or lone pair, and count each type
remember, multiple bonds count as one group
4. Determine the shape and bond angles
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44. Example: Predict the geometry and bond angles of PCl3
26 valence electrons
Four electron groups on P atom: 1 LP, 3 BP
Electron group geometry = __________
Bond angle =
Molecular geometry ________
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45. Example: Predict the geometry and bond angles of SiF5−
40 valence electrons
____ electron groups on P atom: ___ LP, ___ BP
Electron group geometry = __________
Bond angle =
Molecular geometry ________
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46. Example: Predict the geometry and bond angles of ClO2F
26 valence electrons
____ electron groups on P atom: ___ LP, ___ BP
Electron group geometry = __________
Bond angle =
Molecular geometry ________
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47. Representing 3-Dimensional Shapes on a 2-Dimensional Surface
General guideline:
The central atom is put in the plane of the paper
Put as many other atoms as possible in the same plane and indicate with a straight line
Atoms in front of the plane: solid wedge
Atoms behind the plane: hashed wedge
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49. SF6 S F
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50. Multiple Central Atoms
Many molecules have larger structures with many interior atoms
They have multiple central atoms
Consider each central atom in sequence
shape around left C is tetrahedral
shape around center C is trigonal planar
shape around right O is tetrahedral-bent
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51. Geometry of Methanol: Smallest Alcohol molecule
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52. Geometry of Glycine: Smallest Amino Acid (units of Protein)
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53. Practice – Predict the molecular geometries in H3BO3
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54. Practice – Predict the molecular geometries in H3BO3
oxyacid, so H attached to O
3 electron groups on B
4 electron groups on O
B least electronegative
O has
2 bonding groups
2 lone pairs
B has
3 bonding groups
0 pone pairs
B Is Central Atom
Total = 24e─
Shape on B = trigonal planar
Shape on O = tetrahedral bent
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55. Polarity of Molecules For a molecule to be polar it must
have polar bonds
electronegativity difference - theory
bond dipole moments - measured
have an unsymmetrical shape
vector addition
Polarity affects the intermolecular forces of attraction
therefore boiling points and solubilities
like dissolves like
Nonbonding pairs affect molecular polarity, strong pull in its direction
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56. Molecule Polarity
The H─Cl bond is polar. The bonding electrons are pulled toward the Cl end of the molecule. The net result is a polar molecule.
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57. Vector Addition: A. Simple Addition/Cancellation
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58. Vector Addition: B. Addition using Parallelogram
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59. Vector Addition C. Addition of Multiple vectors
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61. Molecule Polarity
The O─C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule.
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62. Molecule Polarity
The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule.
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63. Predicting Polarity of Molecules
1. Draw the Lewis structure and determine the molecular geometry. Bent, Trig. Pyr., Seesaw, T-shape: always polar!!!
2. Determine if bonds are polar
if there are no polar bonds (DEN = 0), the molecule is nonpolar
3. Determine whether the polar bonds add together to give a net dipole moment
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64. Practice – Decide whether the following molecules are polarEN
O = 3.5
N = 3.0
Cl = 3.0
S = 2.5
H = 2.1
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65. Molecular Polarity Affects Solubility in Water
Polar molecules are attracted to other polar molecules
Because water is a polar molecule, other polar molecules dissolve well in water
and ionic compounds as well
Some molecules have both polar and nonpolar parts
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66. Example: Predict the geometry and bond angles of PCl3
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67. Practice – Predict the molecular geometry and bond angles in SiF5─
Si least electronegative
5 electron groups on Si
Si is central atom
5 bonding groups
0 lone pairs
Si = 4e─
F5 = 5(7e─) = 35e─
(─) = 1e─
total = 40e─
Shape = trigonal bipyramid
Bond angles
Feq–Si–Feq = 120°
Feq–Si–Fax = 90°
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68. Practice – Predict the molecular geometry and bond angles in ClO2F
Cl least electronegative
4 electron groups on Cl
Cl is central atom
3 bonding groups
1 lone pair
Cl = 7e─
O2 = 2(6e─) = 12e─
F = 7e─
Total = 26e─
Shape = trigonal pyramidal
Bond angles
O–Cl–O < 109.5°
O–Cl–F < 109.5°

69. Example: Predict whether NH3 is a polar molecule
1. Draw the Lewis structure and determine the molecular geometry
a) eight valence electrons
b) three bonding + one lone pair = trigonal pyramidal molecular geometry
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70. Example: Predict whether NH3 is a polar molecule
2. Determine if the bonds are polar
a) electronegativity difference
b) if the bonds are not polar, we can stop here and declare the molecule will be nonpolar
ENN = 3.0
ENH = 2.1
3.0 − 2.1 = 0.9
therefore the bonds are polar covalent
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71. Example: Predict whether NH3 is a polar molecule
3) Determine whether the polar bonds add together to give a net dipole moment
a) vector addition
b) generally, asymmetric shapes result in uncompensated polarities and a net dipole moment
The H─N bond is polar. All the sets of bonding electrons are pulled toward the N end of the molecule. The net result is a polar molecule.
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72. Practice – Decide whether the following molecules Are polar
Trigonal
Bent
Trigonal
Planar
2.5
1. polar bonds, N-O
2. asymmetrical shape
1. polar bonds, all S-O
2. symmetrical shape
polar
nonpolar
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