Chapter 9Chapter 9Chemical Chemical Bonding I:Bonding I:Lewis Lewis TheoryTheory

Determining the Number of Determining the Number of Valence Electrons in an AtomValence Electrons in an Atom

the column number on the Periodic Table will tell you how many valence electrons a main group atom has◦ Transition Elements all have 2 valence

electrons; Why?

2

1A 2A 3A 4A 5A 6A 7A 8A

Li Be B C N O F Ne

1 e-1 2 e-1 3 e-1 4 e-1 5 e-1 6 e-1 7 e-1 8 e-1

Lewis Symbols of AtomsLewis Symbols of Atomsaka electron dot symbolsuse symbol of element to represent nucleus

and inner electronsuse dots around the symbol to represent

valence electrons◦ pair first two electrons for the s orbital◦ put one electron on each open side for p

electrons ◦ then pair rest of the p electrons

3

Lewis Symbols of IonsLewis Symbols of IonsCations have Lewis symbols without valence

electrons◦ Lost in the cation formation

Anions have Lewis symbols with 8 valence electrons◦ Electrons gained in the formation of the

anion

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Li• Li+1

F

1

F

Stable Electron ArrangementsStable Electron ArrangementsAnd Ion ChargeAnd Ion Charge Metals form cations by

losing enough electrons to get the same electron configuration as the previous noble gas

Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas

The noble gas electron configuration must be very stable

5

Atom Atom’s Electron Config

Ion Ion’s Electron Config

Na [Ne]3s1 Na+1 [Ne]

Mg [Ne]3s2 Mg+2 [Ne]

Al [Ne]3s23p1 Al+3 [Ne]

O [He]2s22p4 O-2 [Ne]

F [He]2s22p5 F-1 [Ne]

RulesRules when atoms bond, they tend to gain, lose, or share

electrons to result in 8 valence electrons ns2np6

◦ noble gas configuration Duet Rule: sharing of 2 electrons

◦ E.g H2 H : H

Octet Rule: sharing of 8 electrons◦ Carbon, oxygen, nitrogen and fluorine always obey

this rule in a stable molecule◦ E.g F2, O2

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ExceptionsExceptions many exceptions

◦ H, Li, Be, B attain an electron configuration like He He = 2 valence electrons Li loses its one valence electron H shares or gains one electron

though it commonly loses its one electron to become H+

Be loses 2 electrons to become Be2+

though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons

B loses 3 electrons to become B3+

though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons

◦ expanded octets for elements in Period 3 or below using empty valence d orbitals

Lewis TheoryLewis Theory the basis of Lewis Theory is that there are certain

electron arrangements in the atom that are more stable◦ octet rule

bonding occurs so atoms attain a more stable electron configuration◦ more stable = lower potential energy◦ no attempt to quantify the energy as the calculation

is extremely complex Bonding pair: two of which are shared with other

atoms Lone pair or nonbonding pair: those that are not

used for bonding

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Electron-Dot StructuresElectron-Dot Structures

HH O•••••

•

••H•H• O

•••

•

••

Electron-Dot StructuresElectron-Dot Structures

Covalent BondingCovalent BondingPredictions from Lewis TheoryPredictions from Lewis Theory Lewis theory allows us to predict the formulas of

molecules Lewis theory predicts that some combinations should be

stable, while others should not◦ because the stable combinations result in “octets”

Lewis theory predicts in covalent bonding that the attractions between atoms are directional◦ the shared electrons are most stable between the

bonding atoms◦ resulting in molecules rather than an array

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ElectronegativityElectronegativity measure of the pull an atom has on bonding

electrons increases across period (left to right) and decreases down group (top to bottom)

◦ fluorine is the most electronegative element◦ francium is the least electronegative element

the larger the difference in electronegativity, the more polar the bond◦ negative end toward more electronegative atom

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Bond PolarityBond Polarity covalent bonding between unlike atoms results in

unequal sharing of the electrons◦ one atom pulls the electrons in the bond closer to its

side◦ one end of the bond has larger electron density than

the other the result is a polar covalent bond

◦ bond polarity◦ the end with the larger electron density gets a partial

negative charge◦ the end that is electron deficient gets a partial positive

charge

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Polar Covalent Bonds: Polar Covalent Bonds: ElectronegativityElectronegativity

NaCl

HCl

Cl2

Electronegativity and Bond Electronegativity and Bond PolarityPolarity

If difference in electronegativity between bonded atoms is 0, the bond is pure covalent◦ equal sharing

If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent

If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent

If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic

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“100%”0 0.4 2.0 4.0

4% 51%Percent Ionic Character

Electronegativity Difference

Bond Dipole MomentsBond Dipole Moments the dipole moment is a quantitative way of describing the

polarity of a bonda dipole is a material with positively and negatively

charged endsmeasured

dipole moment, m, is a measure of bond polarityit is directly proportional to the size of the partial

charges and directly proportional to the distance between them m = (q)(r)

r = radius q = 1.6 x 10-19 C

1 D = 3.34 x 10-30 C•m

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% ionic character =

actual dipole moment

dipole moment if electrons were completely transferred

x 100

Polarity and Dipole MomentPolarity and Dipole Moment Dipole moment:

◦ a vector quantity from the center of the positive charge to the center of negative charge

◦ Represents with an arrow

ExampleExample Determine whether bond formed between the following

pair is ionic, covalent, or polar covalent◦ N and O◦ Sr and F◦ N and Cl

E.g Draw the dipole moment for HF HCl OF

Lewis Structures of Lewis Structures of MoleculesMolecules shows pattern of valence electron distribution in the

molecule useful for understanding the bonding in many

compounds allows us to predict shapes of molecules allows us to predict properties of molecules and how

they will interact together

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Rules for writing Dots Lewis Rules for writing Dots Lewis structuresstructures

Write the correct skeletal structure for molecule◦ Least electronegative atom will be in the center◦ Hydrogen will always be the terminal

Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule◦ If polyatomic ions, charges must be considered when

calculating the total valence electrons Distribute the electrons among the atoms, giving octets

(or duet for hydrogen) to as many atoms as possible If any atoms lack an octet, form double or triple bonds

as necessary to give them octets.

Lewis StructuresLewis 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 off their presence is assumed from common bonding

patterns structures which result in bonding patterns different

from common have formal charges

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B C N O F

Formal ChargeFormal 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-

sum of all the formal charges in a molecule = 0◦ in an ion, total equals the charge

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ExamplesExamples

Draw a Lewis formula then assign formal charge for the following molecules and/or ions

HBr CO2

NH4+ SO3

2-

ResonanceResonance 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 hybrid◦ it does not resonate between the two forms, though

we often draw it that way look for multiple bonds or lone pairs

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Rules of Resonance Rules of Resonance StructuresStructures Resonance structures must have the same

connectivity ◦ only electron positions can change

Resonance structures must have the same number of electrons

Second row elements have a maximum of 8 electrons◦ bonding and nonbonding◦ third row can have expanded octet

Formal charges must total same Better structures have fewer formal charges Better structures have smaller formal charges Better structures have − formal charge on more

electronegative atom

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Drawing resonanceDrawing resonance Any compound for which more than one Lewis structure

may be written is accurately described by no single structure. The actual structure is a resonance hydrid of them all (NOT “flipping back and forth” between resonance forms). The various structures are called contributing structure or resonance forms

C

H

H

N

H

H

C N

H

H

H

H

resonance forms = resonance contributors

C N

H

H

H

H

combined representationresonance hybrid

Drawing Resonance Drawing Resonance StructuresStructures

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O N

O

O·· ··

········

··

··

1. draw first Lewis structure that maximizes octets

2. assign formal charges3. move electron pairs from

atoms with (-) formal charge toward atoms with (+) formal charge

4. if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

5. if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet.

-1

-1

+1

O N

O

O

·· ····

····

······

-1

-1 +1

Exceptions to the Octet RuleExceptions to the Octet Rule expanded octets

◦ elements with empty d orbitals can have more than 8 electrons

odd number electron species e.g., NO◦ will have 1 unpaired electron◦ free-radical◦ very reactive

incomplete octets◦ B, Al

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Drawing Resonance Drawing Resonance StructuresStructures

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1. draw first Lewis structure that maximizes octets

2. assign formal charges3. move electron pairs from atoms

with (-) formal charge toward atoms with (+) formal charge

4. if (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

5. if (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet.

O S

O

O

O

HH

·· ··

········

··

······

-1

-1

+2

O S

O

O

O

HH

··

······

··

······

0

0

0

ExamplesExamplesIdentify Structures with Better or Equal

Resonance Forms and Draw Them◦O3

◦NO2-

◦PO43-