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Chemical Bonding
Lewis TheoryLewis TheoryValence BondValence Bond‐‐VSEPRVSEPRMolecular Orbital TheoryMolecular Orbital Theory
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"...he [his father] knew the difference "...he [his father] knew the difference between knowing the name of something between knowing the name of something and knowing something"and knowing something"
Richard Philips Feynman, Nobel Laureate in Physics(1918‐1988)
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BONDING MODELSBONDING MODELS
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Formation of dihydrogen, H2
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nonpolar covalent bond
polar covalent bond
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• The element with the larger electronegativity will carry the partial negative charge
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The degree of polarity, or ionic character, varies continuously with the electronegativity difference
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Figure 9.19
Percent ionic character of electronegativity difference (EN).
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BEGIN 11/18
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Bond Type and Electronegativity
• ΔEN = 0.0 – 0.4 nonpolar covalent
• ΔEN = 0.4 – 2.0 polar covalent
• ΔEN = 2.0 – 3.3 ionic
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HF
FH
EN 2.1 EN 4.0
H F••
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Bond Polarity
ENCl = 3.03.0 ‐ 3.0 = 0Pure Covalent
ENCl = 3.0ENH = 2.1
3.0 – 2.1 = 0.9Polar Covalent
ENCl = 3.0ENNa = 1.0
3.0 – 0.9 = 2.1Ionic
Bond Dipole Moments• the dipole moment is a quantitative way of describing the polarity of a bond
– a dipole is a material with positively and negatively charged ends
– measured
• dipole moment, , is a measure of bond polarity
– it is directly proportional to the size of the partial charges and directly proportional to the distance between them
• = (q)(r)
• measured in Debyes, D
• the percent ionic character is the percentage of a bond’s measured dipole moment to what it would be if full ions
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Dipole Moments
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Polarity of Molecules
• in order for a molecule to be polar it must
1) have polar bonds• electronegativity difference ‐ theory
• bond dipole moments ‐measured
2) have an unsymmetrical shape
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2) 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
Molecule Polarity
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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.
Vector Addition
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Molecule Polarity
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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.
Molecule Polarity
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The H‐O bond is polar. The 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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Water – a VERY Polar Molecule
stream of water attracted to a charged glass
stream of hexane not attracted to a charged glass
rod rod
Molecule Polarity
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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.
Molecular Polarity Affects Solubility in Water
• polar molecules are attracted to other polar molecules
• since water is a polar molecule,
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other polar molecules dissolve well in water– and ionic compounds as well
• some molecules have both polar and nonpolar parts
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A Soap MoleculeSodium Stearate
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Practice ‐ Decide Whether the Following Are Polar
O N Cl ••
••
••
••
••••
O S
O
O
••
••
•• •
•••••
••
••ENO = 3.5N = 3.0Cl = 3.0S = 2.5
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Practice ‐ Decide Whether the Following Are Polar
O N Cl ••
••
••
••
••••
O S
O
O
••
••
•• •
•••••
••
••
TrigonalBent
TrigonalPlanar
N
3.0
3 0 O3.5
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polarnonpolar
1) polar bonds, N‐O2) asymmetrical shape 1) polar bonds, all S‐O
2) symmetrical shape
Planar
Cl O
3.0
3.5
O
O
OS
3.5 3.5
2.5
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VIDEO
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Lewis Structures of Molecules
• shows pattern of valence electron distribution in the molecule
• useful for understanding the bonding in many compoundscompounds
• allows us to predict shapes of molecules
• allows us to predict properties of molecules and how they will interact together
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Lewis Structures• use common bonding patterns
– C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs
– often Lewis structures with line bonds have the lone pairs left offpairs left off
• their presence is assumed from common bonding patterns
• structures which result in bonding patterns different from common have formal charges
B C N O F
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Practice ‐ Lewis Structures
• CO2
• SeOF2
• H3PO4
• SO3‐2
:O::C::O:O P
O
O
O
HH
H
••
••
••
••
••
••
••
••
••
O ••••
O ••••
16 e‐
32 e‐
• NO2‐1 • P2H4
F Se
O
F
••
••
•• •
•••
••
••
••
••O S
O
O
••
•• •
•••
••
••
••
••
O N O ••
••
••
••
••••
26 e‐
18 e‐
26 e‐
14 e‐ H P P H
HH
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Formal Charge• during bonding, atoms may wind up with more or less electrons in order to fulfill octets ‐ this results in atoms having a formal charge
FC = valence e‐ ‐ nonbonding e‐ ‐ ½ bonding e‐
left OFC = 6 4 ½ (4) = 0 ••••••left OFC = 6 ‐ 4 ‐ ½ (4) = 0
S FC = 6 ‐ 2 ‐ ½ (6) = +1
right O FC = 6 ‐ 6 ‐ ½ (2) = ‐1
• sum of all the formal charges in a molecule = 0
– in an ion, total equals the charge
•• •• ••••••••
••O S O••••
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Common Bonding Patterns
B C N O
C+
N+
O+
F
F+
C N O
C-
N-
O-
B-
F
‐F
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Resonance• when there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures
• the actual molecule is a combination of the resonance forms – a resonance hybridy
– it does not resonate between the two forms, though we often draw it that way
• look for multiple bonds or lone pairs
•••• •• ••••••••
•• ••O S O O S O•••••• ••••
••••
••••
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Dinitrogen monoxide
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HO
:O:
:O:
SOH HO
:O:
:O:
SOH||
||
|
|
II Structure I Structure
Structure I obeys the octet rule,
but is not consistent with experiment
Structure II violates the octet rule,
but is consistent with experiment
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