1. Chemical Bond
A chemical bond is a strong attractive force that exists between certain atoms in a substance.
There are three types of chemical bonds:
Ionic bonds
Covalent bonds
Metallic bonds
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2. An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions.
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3. Ionic Bond
An ionic bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of another atom.
The atom that transferred the electron(s) becomes a cation.
The atom that gained the electron(s) becomes an anion.
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4. A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the chemical symbol of the element.
Note – Dots are placed one to each side, until all four sides are occupied.
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5. Lewis electron-dot symbol
Table 9.1 illustrates the Lewis electron-dot symbols for second- and third-period atoms.
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6. Lewis electron-dot symbol
Use Lewis electron-dot symbols to represent the transfer of electrons from magnesium to fluorine atoms to form ions with noble-gas configurations.
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7. Ionic Bond An ionic bond between two atoms is formed in two steps:
An electron is transferred between the two separate atoms to give ions.
The ions attract one another, forming an ionic bond.
Both steps occur simultaneously.
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8. Coulomb's Law
The potential energy obtained in bringing two charges Q1 and Q2, initially apart, up to a distance r apart is directly proportional to the product of the charges and inversely proportional to the distance between them.
k is a physical constant (8.99 × 109 J·m/C2).
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9. NaCl(s) → Na+(g) + Cl-(g)
The lattice energy is the change in energy that occurs when an ionic solid is separated into gas-phase ions.
It is very difficult to measure lattice energy directly.
NaCl(s) → Na+(g) + Cl-(g)
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10. Trends in Lattice Energy: Ion Size
The force of attraction between charged particles is inversely proportional to the distance between them.
Larger ions mean the center of positive charge (nucleus of the cation) is farther away from the negative charge (electrons of the anion).
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11. Lattice Energy versus Ion Size
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12. Trends in Lattice Energy: Ion Charge
The force of attraction between oppositely charged particles is directly proportional to the product of the charges.
Larger charge means the ions are more strongly attracted.
Larger charge = stronger attraction
Stronger attraction = larger lattice energy
Of the two factors, ion charge is generally more important.
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13. Properties of Ionic Compounds
Ionic solids are high-melting substances.
This is due to the strong interactions between the ions.
Charge on the ions of MgO is double the charge on the ions of NaCl thus, the force will be four times stronger.
The size of Na+ is larger than that of Mg2+; the size of Cl- is larger than that of O2-.
The distance between Mg2+ and O2- is smaller than the distance between Na+ and Cl-, the force between Mg2+ and O2- will be greater.
Thus, MgO has a higher melting point than NaCl.
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14. Properties of Ionic Compounds
Based on the higher charge and the smaller distance for MgO, its melting point MgO should be significantly higher than the melting point of NaCl.
The actual melting point of NaCl is 801°C; the melting point of MgO is 2800°C.
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15. Properties of Ionic Compounds
Hard and brittle crystalline solids
All are solids at room temperature.
Melting points generally > 300 C
The liquid state conducts electricity.
The solid state does not conduct electricity.
Many are soluble in water.
The solution conducts electricity well.
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16. Properties of Ionic Compounds
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17. Ionic Bonding Model versus Reality
Lewis theory implies that if the ions are displaced from their positions in the crystal lattice, repulsive forces should occur.
This predicts the crystal will become unstable and break apart. Lewis theory predicts ionic solids will be brittle.
Ionic solids are brittle. When struck, they shatter.
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18. Conductivity of NaCl
In NaCl(s), the ions are stuck in position and not allowed to move to the charged rods.
In NaCl(aq), the ions are separated and allowed to move to the charged rods.
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19. Covalent Bonding in Molecules
Covalent Bond: A bond that results from the sharing of electrons between atoms
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20. Covalent Bond
To consider how a covalent bond forms, we can monitor the energy of two isolated hydrogen atoms as they move closer together.
The potential energy decreases—first gradually, and then more steeply—to a minimum. As the atoms continue to move closer, it increases dramatically.
The distance between the atoms when energy is at a minimum is called the bond length.
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21. Covalent Bonding
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22. Covalent Bonding G.N. Lewis (1916): Some atoms share e- to form bonds.
Number of bonds = Number shared e- pairs.
Molecules with shared e- are covalent compounds with covalent bonds.
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23. H H − H Single Covalent Bonds
Lewis structures: dot = 1 e-. Line = 1 pair of e- H
H − H
Single bond: one shared pair of e-.
Octet rule
To form bonds, elements gain, lose, or share e- to achieve 8 valence e-.
Doesn’t apply if Z ≤ 5 (H, He, Li, Be and B)
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24. H O H O Single Covalent Bonds
Bonding pairs = e- pairs shared between atoms.
Lone pairs = unshared e- pair.
bonding
lone pair H O H O
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25. F H O H N Single Covalent Bonds 2F = 2(7) = 14 valence e-
Share 2 e- to form octets
O + 2H = 6 + 2(1) = 8 val. e-
Two O-H bonds H O H N
N + 3H = 5 + 3(1) = 8 val. e-
3 N-H bonds
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26. A coordinate covalent bond is formed when both electrons of the bond are donated by one atom.
The two electrons forming the bond with the hydrogen on the left were both donated by the nitrogen. Once shared, they are indistinguishable from the other N—H bonds.
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27. Single Covalent Bonds
Those examples reveal a general rule for main-group elements:
Number of e- shared = 8 – (A group number)
Group Number
Number of Valence e-
Number of e- shared
Example 4A 4
8 – 4 = 4
C in CH4 5A 5 3
N in NH3 6A 6 2
O in H2O 7A 7 1
F in HF
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28. Guidelines for Writing Lewis Structures
1. Count the valence e- in the molecule.
2. Draw a skeleton structure. Join atoms with single lines (pairs of e-).
3. Add e- pairs to form octets (except H). Start with terminal atoms.
Extra e- ? Place around the central atom.
Too few e-? Convert lone pairs into multiple bonds.
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29. F P F F Writing Lewis Structures Phosphorus trifluoride, PF3
1. PF3 = (7) = 26 valence e-
3 x F (group 7A)
P (group 5A) F P F
Skeleton (X is central in XYn ). F
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30. F P Writing Lewis Structures
3. Build octets – start with terminal atoms. F P
6 e- used in 3 bonds, 20 e- remain (10 pairs)
= 26 e-
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31. O O P O O Writing Lewis Structures Phosphate ion, PO43-
P (grp 5A)
4 x O (grp 6A)
charge (-3) 3 e- O O P O
2. Skeleton; P is central O
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32. O P O P Writing Lewis Structures 3. Add e- pairs:
8 e- used in 4 bonds, 24 e- remain (12 pairs)
32 e- used. O P 3-
Add brackets and overall charge to show this is an ion.
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33. O C H Multiple Covalent Bonds
Too few “dots” to complete all the octets?
Convert lone pairs to shared pairs.
Methanal (formaldehyde) H2CO
1. Valence e- = 2(1) = 12
2. Skeleton O C H
6 e- in bonds. Add the other 3 pairs to O (outer atom).
Each H shares 2 e-
C only “has” 6 e-.
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34. O C H O C H Multiple Covalent Bonds Convert lone pairs to bond pairs.
Each H shares 2 e-
C shares 8
O “has” 8
4 shared + 2 lone pairs
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35. O C O C O Multiple Covalent Bonds Carbon dioxide CO2
2. Skeleton
4 e- in bonds. Add 3 pairs to each O.
4. Convert lone pairs to bond pairs. O C O
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36. Multiple Covalent Bonds
A single bond is a covalent bond in which one pair of electrons is shared by two atoms.
A double bond is a covalent bond in which two pairs of electrons are shared by two atoms.
ethene C2H4
(ethylene) H C
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37. Multiple Covalent Bonds
A triple bond is a covalent bond in which three pairs of electrons are shared by two atoms.
Double bonds form primarily with C, N, O, and S atoms.
Triple bonds form primarily with C and N atoms.
ethyne (acetylene)
H-C≡C-H
C2H2
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38. A polar covalent bond (or polar bond) is a covalent bond in which the bonding electrons spend more time near one atom than near the other atom.
Electronegativity, X, is a measure of the ability of an atom in a molecule to draw bonding electrons to itself.
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39. Electronegativity
Electronegativity increases from left to right and from bottom to top in the periodic table. F, O, N, and Cl have the highest electronegativity values.
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40. Polar Covalent Bond
The difference in electronegativity between the two atoms in a bond is a rough measure of bond polarity.
A small difference results in a nonpolar bond.
A large difference results in a polar bond.
A very large difference can result in an ionic bond.
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41. Polar Covalent Bond
Using electronegativities, arrange the following bonds in order by increasing polarity: P—H, H—O, C—Cl.
The difference for P—H is 0.0
The difference for H—O is 1.4
The difference for C—Cl is 0.5
The order is
P—H <
C—Cl <
H—O
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42. Polar Covalent Bond
During bond formation, electrons are piled toward the more electronegative atom. The direction of electrons can be predicted using the electronegativity scale.
In H–Cl bond, electrons are pulled toward the Cl atom (X = 3.0) rather than the H atom (X = 2.1).
The Cl atom acquires a partial negative charge.
The H atom acquires a partial positive charge. Thus, the HCl molecule is a polar molecule.
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43. Bond Properties: Electronegativity
Differences, ΔEN, determine bond polarity: EN
F–F
H–Br
H–F
Na+ F-
ΔEN
0.5
1.9
3.1
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44. Formal Charge
The charge a bonded atom would have if its bonding e- were shared equally.
Used to study charge distribution in a molecule.
Method
Find the number of e- assigned to each atom:
e- “on” an atom = (lone pair e-) + ½ (bonding e-)
Formal charge of each atom
= (# of valence e-) – (e- “on” the atom).
Note: sum of formal charges = molecular charge
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45. [ ] O C N Formal Charge – O C N Valence e- 6 4 5 Lone pair e- 6 0 2
[ ] – O C N
O C N
Valence e
Lone pair e
½ shared e
Formal Charge
Check:  (formal charges) = ion charge = -1
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46. N O N O Formal Charge If there is choice between Lewis structures:
Smaller formal charges are favored.
Negative formal charges should be on the most EN atoms
Like charges should not be on adjacent atoms
Which N2O structure is preferred? N O N O
Formal charges: -1 +1 +1 -1
Preferred. ENO > ENN
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47. Cl O Formal Charge Which ClO2- structure is preferred? Preferred
Formal charges: -1 +1
Preferred
(Smaller charges)
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48. Lewis Structures & Resonance
Ozone has 2 equivalent structures: O
Both:
obey the octet rule
have the same number and types of bonds
have the same formal charges
Experiments show that the OO bonds are identical.
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49. Resonance
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50. Lewis Structures & Resonance
Resonance structures are used to show O3 is a mixture of both: O
A resonance hybrid best represents O3. Each bond is ≈1½ bonds. Ozone does NOT “flip” back and forth.
Valid resonance structures:
Bonding and e- pair positions are changed.
Atom positions must not change.
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51. Lewis Structures & Resonance
Resonance in CO32-
Experiment: All three CO bonds = 129 pm
Typical bond lengths
C-O =143 pm; C=O = 122 pm.
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52. Lewis Structures and Resonance
Benzene, C6H6, is a ring compound best described using resonance ideas:
Experiments: All C-C bonds have identical lengths.
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53. Resonance & Structure of Benzene
Benzene is usually drawn without its hydrogen atoms shown, or with a ring representing six delocalized electrons spread evenly over the carbon atoms: or
Solid ring shows resonance
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54. Metallic Bonds
The low ionization energy of metals allows them to lose electrons easily.
The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared by all atoms/ions in the metal.
An organization of metal cation islands in a sea of electrons
Electrons delocalized throughout the metal structure
Bonding results from attraction of the cations for the delocalized electrons.
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55. Exceptions to the Octet Rule
Be and B form e- deficient compounds: H Be
2 + 2(1) = 4 valence e-
3 + 3(7) = 24 valence e- B F
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56. Fewer than Eight Valence Electrons
Often very reactive B F N H B F N H +
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57. O N O N Odd Number of Valence Electrons
Some stable molecules have an odd number of e-.
Examples
NO = 11 valence e- O N O N
NO (6) = 17 valence e-
Free radical = atom or molecule with unpaired e-.
Very reactive. Most stable molecules have paired e-.
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58. More Than Eight Valence Electrons
“Expanded octets” are relatively common.
Low-lying d orbitals can accept extra e-
(only 3rd period and beyond).
Examples
5 bonds (5 e- pairs) around P in PF5
NF5 does not exist
4 bonds and 1 lone-pair (5 e- pairs) around S in SF4
OF4 does not exist.
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59. More Than Eight Valence ElectronsCl F
ClF (7) = 28 val. e-
Make octets on F
24 e- used, 4 remain
[bonds (3 x 2); Lone pairs (3 x 6)]
Add 2 lone pairs to Cl – the 3rd period element
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60. Bond Properties: Bond Length
Atom size is important.
Remember the periodic table trends:
increasing
size
< <
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61. Bond Length Average single bond lengths (pm) I Br Cl S P Si F O N C H
161
142
127
132
138
145 92 94 98
110 74
210
191
176
181
187
194
141
143
147
154
203
184
169
174
180
134
136
140
199
165
170
173
130
197
178
163
168
128
250
231
216
221
227
234
243
224
209
214
220
237
218
208
232
213
200
247
228
266
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62. Bond Length More bonds = greater e- density between atoms
= atoms pulled together more strongly
Multiple bonds (pm): N-N 140 N=N 120 N≡N 110
C-C 154 C=C 134 C≡C 121
C-N 147 C=N 127 C≡N 115
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63. Bond Enthalpy Average bond enthalpies (kJ/mol)Br Cl S P Si F O N C H
299
366
431
347
322
323
566
467
391
416
436
213
285
327
272
264
301
486
336
356
193
~200
335
201
160
205
~340
173
190
146
255
326
490
582
158
234
310
226
184
319
209
217
242
180
151
Shorter bond = stronger bond.
Multiple bonds: N-N 160 N=N 418 N≡N 946
C-C 356 C=C 598 C≡C 813
C-N 285 C=N 616 C≡N 866
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64. Molecular Geometry
Diatomic molecules are the easiest to visualize in three dimensions
HCl
Cl2
Diatomic molecules are linear
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65. Two Electron Pairs Linear Bent Bond angles
Always 180° in a linear molecule
Always less than 180° in a bent molecule
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66. Valence Shell Electron Pair Repulsion Theory
Ideal geometry of a molecule is determined by the way the electron pairs orient themselves in space
Orientation of electron pairs arises from electron repulsions
Electron pairs spread out so as to minimize repulsion
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67. VSEPR Electron-Pair Geometries
Fundamental geometry exists that corresponds to the total number of electron pairs around the central atom: 2, 3, 4, 5, and 6
linear
trigonal
planar
tetrahedral
trigonal
bipyramidal
octahedral
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68. Molecular Geometries Trigonal planar Tetrahedral
Electron pairs form an equilateral triangle around the central atom
Bond angles are 120°
Tetrahedral
Bond angles are 109.5°
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69. Molecular Geometries Trigonal bipyramid Bond angles vary
In the trigonal plane, 120°
Between the plane and the apexes, 90°
Between the central atom and both apexes, 180°
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70. Molecular Geometries Octahedron Is a square bipyramid Bond angles vary
90° in and out of plane
180° between diametrically opposite atoms and the central atom
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71. Molecular Geometry Summarized
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72. Molecular Geometry Summarized

73. Unshared Pairs and Geometry
Electron-pair geometry
Same as that when single bonds are involved
Bond angles are a smaller than the ideal values
Molecular geometry
Quite different when one or more unshared pairs are present
While describing, only the positions of the bonded atoms are referred
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74. Two Ways of Showing the Geometry of the NH3 Molecule
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75. Geometries with Two, Three, or Four Electron Pairs Around a Central Atom
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76. Nonbonding Pairs and Bond Angle
Nonbonding pairs are physically larger than bonding pairs.
Therefore, their repulsions are greater; this tends to compress bond angles.
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77. Multiple Bonds and Bond Angles
Double and triple bonds have larger electron domains than single bonds.
They exert a greater repulsive force than single bonds, making their bond angles greater.
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78. Example
Predict the geometry of NH4+, BF3, and PCl3
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79. Example Strategy: Start by writing Lewis structures for each species
Focus on the central atom, then decide what species type (AX2, AX3, ) the molecule or ion is
A represents the central atom
X represents the terminal atoms
E represents the unshared electron pairs
Match the species type with the molecular geometry and ideal bond angles for the species
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80. Example Solution: Lewis structure for NH4+
Species type A = N, X = H(4), no E→ AX4
Geometry tetrahedral, 109.5° bond angles
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81. Example Lewis structure for BF3
Species type A = B, X = F(3), no E→ AX3
Geometry trigonal planar, 120° bond angles
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82. Example Lewis structure for PCl3
Species type A = P, X = Cl(3), no E = 1→ AX3E
Geometry trigonal pyramid (The ideal bond angles are 109.5° but actually are 104°)
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83. The A-X-E Notation A denotes a central atom X denotes a terminal atom
E denotes a lone pair
Example
Water
H2O
O is central
Two lone pairs
Two hydrogens
AX2E2
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84. Lone Pairs and Expanded Octets
Where expanded octets are possible, place the extra lone pairs on the central atom
Example: XeF4
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85. Multiple Bonds
For molecular geometry purposes, multiple bonds behave the same as single bonds
Electron pairs are located in the same place (between the nuclei)
Geometry of the molecule is determined by the number of terminal atoms
Not affected by the presence of a double or triple bond
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86. Molecular Geometries for Molecules with Expanded Octets
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87. Molecular Geometries for Molecules with Expanded Octets
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88. Example
Predict the geometries of the ClO3– ion, the NO3– ion, and the N2O molecule, which have the following Lewis structures, respectively
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89. Example Strategy Classify each species as AXmEn
Multiple bonds count as single bonds. It is the number of terminal atoms (X) that are counted, not the number of bonds
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90. Example Solution For ClO3–:
Species type A = Cl, X = O =3, E = 1 → AX3E
Geometry trigonal pyramid, ideal bond angles are 109.5°
For NO3–:
Species type A = N, X = O = 3, E = 0 → AX3
Geometry trigonal planar, ideal bond angles are 120°
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91. Example For N2O: Species type A = N, X = N and O = 2, E = 0 → AX2
Geometry linear, ideal bond angles are 180°
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92. Describing the Geometry of Glycine
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93. Dipoles
When two equal, but opposite, charges are separated by a distance, a dipole forms.
A dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:
 = Q x r
It is measured in debyes (D).
1D = 3.34X10-30 Cm. (coulomb-meters)
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94. In part A, there is no electric field; molecules are oriented randomly.
Dipoles
In part B, there is an electric field; molecules align themselves against the field. A B
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95. Molecular 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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96. Molecular 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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97. Molecular Polarity
The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. Because the molecule is bent, not linear, the net result is a polar molecule.
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98. Vector Addition
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99. 8-קשר כימי

100. Polarity of Molecules
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101. Bond and Molecular Polarity
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102. The dipole moment of a molecule can affect its properties
cis-1,2-dichloroethene
trans-1,2-dichloroethene
There is no net polarity; this is a nonpolar molecule.
The net polarity is down; this is a polar molecule.
Boiling point 48°C.
Boiling point 60°C.
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103. Valence Bond Theory
Valence Bond Theory: A quantum mechanical model that shows how electron pairs are shared in a covalent bond
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104. Valence Bond Theory
Valence Bond Theory: A quantum mechanical model that shows how electron pairs are shared in a covalent bond
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105. Valence Bond Theory
Covalent bonds are formed by overlap of atomic orbitals, each of which contains one electron of opposite spin.
Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms.
The greater the amount of overlap, the stronger the bond.
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106. Hybridization and sp3 Hybrid Orbitals
How can the bonding in CH4 be explained?
4 valence electrons
2 unpaired electrons
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107. Hybridization and sp3 Hybrid Orbitals
How can the bonding in CH4 be explained?
4 valence electrons
2 unpaired electrons
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108. Hybridization and sp3 Hybrid Orbitals
How can the bonding in CH4 be explained?
4 nonequivalent orbitals
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109. Hybridization and sp3 Hybrid Orbitals
How can the bonding in CH4 be explained?
4 equivalent orbitals
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110. 8-קשר כימי

111. Hybridization and sp3 Hybrid Orbitals
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112. Other Kinds of Hybrid Orbitals-sp2
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113. Other Kinds of Hybrid Orbitals-sp2
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114. Bond Rotation
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115. Double Bonds & Isomerism
Bond rotation:
CHCl=CHCl is “locked” - 2 isomers are possible.
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116. Double Bonds & Isomerism
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117. Other Kinds of Hybrid Orbitals-sp
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118. Other Kinds of Hybrid Orbitals-sp
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119. Consider a Molecule Like PCl5
Phosphorous:valence electron configuration
of 3s23p3. (5 electrons)
Each of the five electrons forms a single bond with a chlorine atom.
This means that central atom in the molecule needs 5 bonding orbitals to achieve the trigonal bipyramidal electronic geometry.
This cannot happen with sp3 hybridization… 3s 3p  
mix or
“hybridize”
sp3 hybrid valence bond orbitals
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120. mix the 3s, the three 3p and one 3d
Valence Bond Theory (2):
Expanded Valence
The only way to produce 5 half-filled orbitals on phosphorous is by adding a fifth atomic orbital… 3d 3d        
new sp3d hybrid valence bond orbitals 3p
mix the 3s, the three 3p and one 3d  3s
Each of these half–filled sp3d orbitals can form a –bond with a chlorine atom in PCl5.
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121. Additional Example of sp3d Hybrid Molecules
SF4 (sulfur tetrafluoride)
ClF3 (chlorine trifluoride)
EPG: Trigonal Bipyramidal
MG: T-shape
EPG: Trigonal Bipyramidal
MG: See Saw
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122. mix the 3s, the three 3p and two 3d’s
Valence Bond Theory (2): Expanded Valence
What about molecules with 12 electrons in the valence?
In order to achieve an expanded valence that can hold six electron pairs (bp & lp) we need to form 6 new hybrid orbitals.
This requires the mixing of an s, three p’s and two d–atomic orbitals. 3d 3d
new sp3d2 hybrid valence bond orbitals 3p
mix the 3s, the three 3p and two 3d’s 3s
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123. mix the 3s, the three 3p and two 3d’s
sp3d2 Hybridization: SF6 3s 3p 3d
mix the 3s, the three 3p and two 3d’s  
six new sp3d2 hybrid valence bond orbitals
Each of these half–filled sp3d2 orbitals can form a –bond with a fluorine atom.
sulfur
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124. sp3d2 Hybridization: SF6 SF6
Sulfur has a valence electron configuration of 3s23p4 (6 electrons).
Each of the six electrons forms a single bond with a fluorine atom forming an octahedral MG and EPG.
The bonding can be described in terms of sp3d2 hybrid orbitals.
SF6
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125. Molecular Orbital Theory: The Hydrogen Molecule
Atomic Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in an atom
Molecular Orbital: A wave function whose square gives the probability of finding an electron within a given region of space in a molecule
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126. Molecular Orbital Theory: The Hydrogen Molecule
σ bonding orbital
σ* antibonding orbital
Bond order =
(# bonding e– – # antibonding e–) 2
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127. Molecular Orbital Theory: The Hydrogen Molecule
= 1 2
2 – 0
Bond order =
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128. Molecular Orbital Theory: The Hydrogen Molecule
= 1/2 2
2 – 1
Bond order:
= 0 2
2 – 2
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129. Interaction of p Orbitals
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130. Interaction of p Orbitals
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131. Interaction of p Orbitals
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132. 8-קשר כימי

133. O2 Dioxygen is paramagnetic.
Paramagnetic material has unpaired electrons.
Neither Lewis theory nor valence bond theory predict this result.
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134. O2 as Described by Lewis and VB Theory
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135. Molecular Orbital Theory: Other Diatomic Molecules
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136. s and p Orbital Interactions
In some cases, s orbitals can interact wit the pz orbitals more than the px and py orbitals.
It raises the energy of the pz orbital and lowers the energy of the s orbital.
The px and py orbitals are degenerate orbitals.
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137. Heteronuclear Diatomic Molecules
Diatomic molecules can consist of atoms from different elements.
How does a MO diagram reflect differences?
The atomic orbitals have different energy, so the interactions change slightly.
The more electronegative atom has orbitals lower in energy, so the bonding orbitals will more resemble them in energy.
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138. Benzene
The organic molecule benzene (C6H6) has six -bonds and a p orbital on each C atom, which form delocalized bonds using one electron from each p orbital.
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