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AP Notes Chapter 8. Bonding and Molecular Structure: Fundamental Concepts Valence e- and Bonding Covalent Ionic Resonance & Exceptions to Octet Rule Bond Energy & Length Structure, Shape & Polarity of Compounds. What is a Bond?. A force that holds atoms together. Why?

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slide1

AP Notes Chapter 8

Bonding and Molecular Structure: Fundamental Concepts

Valence e- and Bonding

Covalent

Ionic

Resonance & Exceptions to Octet Rule

Bond Energy & Length

Structure, Shape & Polarity of Compounds

what is a bond
What is a Bond?
  • A force that holds atoms together.
  • Why?
  • We will look at it in terms of energy.
  • Bond energy the energy required to break a bond.
  • Why are compounds formed?
  • Because it gives the system the lowest energy.
covalent compounds
Covalent compounds?
  • The electrons in each atom are attracted to the nucleus of the other.
  • The electrons repel each other,
  • The nuclei repel each other.
  • The reach a distance with the lowest possible energy.
  • The distance between is the bond length.
why isn t helium diatomic

He2

.

.

E

He + He

.

.

Inter-nuclear Distance

Why Isn’t Helium Diatomic?
slide7
F + F F2

2p____ ____ ___ ___ ____ ____ 2p2s ____ ____ 2s

F F

ionic bonding
Ionic Bonding
  • An atom with a low ionization energy reacts with an atom with high electron affinity.
  • The electron moves.
  • Opposite charges hold the atoms together.
slide9
Li + Cl1s22s1 [Ne] 3s23p52s ___ 3p _____ _____ ___1s _____ 3s _____[Ne]
electronegativity
Electronegativity

Describes the relative ability of an atom within a molecule to attract a shared pair of electrons to itself.

electronegativity1
Electronegativity

Pauling electronegativity values, which are unit-less, are the norm.

electronegativity range from 0 7 to 4 0
ElectronegativityRange from 0.7 to 4.0

Figure 9.9 – Kotz & Treichel

bond a b
Bond: A - B

DEN = | ENA - ENB |

bond character
Bond Character

“Ionic Bond” - Principally Ionic Character

“Covalent Bond” - Principally Covalent Character

slide18

F - F EN = 0

Non-polar

slide19

N - O EN = |3.0 - 3.5| = 0.5

O

N

Slightly polar

slide21

Ca - O  EN = |1.0 - 3.5| = 2.5

Ca

O

Ionic Bond with somecovalent character

electronegativity2
Electronegativity
  • The ability of an electron to attract shared electrons to itself.
  • Pauling method
  • Imaginary molecule HX
  • Expected H-X energy = H-H energy + X-X energy 2
  • D = (H-X) actual - (H-X)expected
electronegativity3
Electronegativity
  • D is known for almost every element
  • Gives us relative electronegativities of all elements.
  • Tends to increase left to right.
  • decreases as you go down a group.
  • Noble gases aren’t discussed.
  • Difference in electronegativity between atoms tells us how polar.
slide25

Polar

Covalent

Ionic

Electronegativity

difference

Bond

Type

Zero

Covalent

Covalent Character

decreases

Ionic Character increases

Intermediate

Large

dipole moments
Dipole Moments
  • A molecule with a center of negative charge and a center of positive charge is dipolar (two poles),
  • or has a dipole moment.
  • Center of charge doesn’t have to be on an atom.
  • Will line up in the presence of an electric field.
which molecules have them
Which Molecules Have Them?
  • Any two atom molecule with a polar bond.
  • With three or more atoms there are two considerations.
  • There must be a polar bond.
  • Geometry can’t cancel it out.
slide34

Molecular Polarity

Vector Sum of Bond Polarities

slide35

MgBr2

Mg - Br EN = |1.2 - 2.8| = 1.6

Mg

Br

Br

Covalent BOND w/much ionic character, BUT NON-POLAR molecule

slide39

Lewis

Structures

slide40

The most important requirement for the formation of a stable compound is that the atoms achieve noble gas e- configuration

valence shell electron pair repulsion model vsepr
Valence Shell ElectronPair Repulsion Model(VSEPR)

The structure around a given atom is determined principally by minimizing electron-pair repulsions

vsepr

Electron

pairs

Bond

Angles

Underlying

Shape

2

180°

Linear

3

120°

Trigonal Planar

4

109.5°

Tetrahedral

90° &

120°

Trigonal Bipyramidal

5

6

90°

Octagonal

VSEPR
lewis structures
LEWIS STRUCTURES
  • : draw skeleton of species
  • : count e- in species
  • : subtract 2 e- for each bond in skeleton
  • : distribute remaining e-
slide44

Distinguish Between

ELECTRONIC

Geometry

&

MOLECULAR

Geometry

slide45

CH4

Bond angle = 109.50

Electronic geometry: tetrahedral

Molecular geometry: tetrahedral

slide46

H3O+

Bond angle ~ 1070

Electronic geometry: tetrahedral

Molecular geometry: trigonal pyramidal

slide47

H2O

Bond angle ~ 104.50

Electronic geometry: tetrahedral

Molecular geometry: bent

slide48

NH2-

Bond angle ~ 104.50

Electronic geometry: tetrahedral

Molecular geometry: bent

slide51

BaI2

Bond angle =1800

Electronic geometry: linear

Molecular geometry: linear

slide52

BF3

Bond angle =1200

Electronic geometry: trigonal planar

Molecular geometry: trigonal planar

slide53

PF5

Bond angle = 1200 & 900

Electronic geometry: trigonal bipyramidal

Molecular geometry: trigonal bipyramidal

slide54

SF4

Bond angle = 1200 & 900

Electronic geometry: trigonal bipyramidal

Molecular geometry: see-saw

slide55

ICl3

Bond angle <= 900

Electronic geometry: trigonal bipyramidal

Molecular geometry: T-shape

slide56

I3-

Bond angle = 1800

Electronic geometry: trigonal bipyramid

Molecular geometry: linear

slide58

PCl6-

Bond angle = 900

Electronic geometry: octahedral

Molecular geometry: octahedral

slide59

BrF5

Bond angle ~ 900

Electronic geometry: octahedral

Molecular geometry: square pyramidal

slide60

ICl4-

Bond angle = 900

Electronic geometry: octahedral

Molecular geometry: square planar

actual shape
Actual shape

Non-BondingPairs

ElectronPairs

BondingPairs

Shape

2

2

0

linear

3

3

0

trigonal planar

3

2

1

bent

4

4

0

tetrahedral

4

3

1

trigonal pyramidal

4

2

2

bent

actual shape1
Actual Shape

Non-BondingPairs

ElectronPairs

BondingPairs

Shape

5

5

0

trigonal bipyrimidal

5

4

1

See-saw

5

3

2

T-shaped

5

2

3

linear

slide63

Actual Shape

Non-BondingPairs

ElectronPairs

BondingPairs

Shape

6

6

0

Octahedral

6

5

1

Square Pyramidal

6

4

2

Square Planar

6

3

3

T-shaped

6

2

1

linear

examples
EXAMPLES

Ethene

Acetic Acid

Oxygen

Nitrogen

slide66

RESONANCE

&

FORMAL CHARGE

resonance
Resonance
  • Sometimes there is more than one valid structure for an molecule or ion.
  • NO3-
  • Use double arrows to indicate it is the “average” of the structures.
  • It doesn’t switch between them.
  • NO2-
  • Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.
examples1
EXAMPLES

Nitrate ion

Ozone

formal charge
FORMAL CHARGE

the charge assigned to an atom in a molecule or polyatomic ion

FC atom = Family# - [LPE + ½(BE)]

Sum FC’s atoms = ion charge

Closer sum FC’s is to zero more stable

formal charge1
Formal Charge
  • For molecules and polyatomic ions that exceed the octet there are several different structures.
  • Use charges on atoms to help decide which.
  • Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.
formal charge2
Formal Charge
  • The difference between the number of valence electrons on the free atom and that assigned in the molecule.
  • We count half the electrons in each bond as “belonging” to the atom.
  • SO4-2
  • Molecules try to achieve as low a formal charge as possible.
  • Negative formal charges should be on electronegative elements.
assignment of e
Assignment of e-

1. Lone pairs belong entirely to atom in question

2. Shared e- are divided equally between the two sharing atoms

slide73

The sum of the formal charges of all atoms in a species must equal the overall charge on the species.

a useful equation
A useful equation
  • (happy-have) / 2 = bonds
  • POCl3P is central atom
  • SO4-2S is central atom
  • SO3-2S is central atom
  • PO4-2P is central atom
  • SCl2 S is central atom
exceptions to the octet
Exceptions to the octet
  • BH3
  • Be and B often do not achieve octet
  • Have less than and octet, for electron deficient molecules.
  • SF6
  • Third row and larger elements can exceed the octet
  • Use 3d orbitals?
  • I3-
exceptions to the octet1
Exceptions to the octet
  • When we must exceed the octet, extra electrons go on central atom.
  • ClF3
  • XeO3
  • ICl4-
  • BeCl2
slide77

If nonequivalent Lewis structures exist, the one(s) that best describe the bonding in the species has...

favored lewis structures
FAVORED LEWIS STRUCTURES

1. formal charges closest to zero

2. negative formal charge

is on the most electronegative atom

examples2
EXAMPLES

Carbon dioxide

Thiocyanate ion

Sulfate ion

slide80

BOND

ENERGY & LENGTH

slide82

Bonds form between atoms because bonded atoms exhibit a lower energy.

Thus, energy is required to break bonds and energy is released when bonds are formed.

slide83
Bond Order = # bonds to a specific set of elementsC-C the BO=1C=C the BO=2C C the BO=3Fractions are possible
covalent bonds
COVALENT BONDS

Bond Dissociation

Energy

Table 9.9 (text)

slide85

Bond Energy

(kJ/mol)

H-F 565

H-Cl 432

H-Br 366

H-I 299

slide86

Bond Energy

(kJ/mol)

Cl-Cl 242

Br-Br 193

I-I 151

slide87

Bond Energy

(kJ/mol)

slide88

Bond Energy

(kJ/mol)

slide90

Energy

0

Internuclear Distance

slide91

Energy

0

Internuclear Distance

slide92

Energy

0

Internuclear Distance

slide93

Energy

0

Internuclear Distance

slide94

Energy

0

Bond Length

Internuclear Distance

slide95

Energy

Bond Energy

0

Internuclear Distance

slide96

Bond: Energy Length

(kJ/mol) (pm)

slide97

Bond: Energy Length

(kJ/mol) (pm)

binary ionic compounds
Binary Ionic Compounds

metal(s) + non-metal (g) ---> salt(s)

lattice energy
Lattice Energy

Energy change occurring when separated gaseous ions are packed together to form an ionic solid

M+(g) + NM-(g) --> M-NM

slide100

What is the lattice energy of NaCl(s)?

Na+(g) + Cl-(g) ---> NaCl(s)

lattice energies
Lattice Energies

LiCl 834 : BeCl2 3004

NaCl 769 : MgCl2 2326

KCl 701 : CaCl2 2223

Li2O 2799 : BeO 4293

Na2O 2481 : MgO 3795

K2O 2238 : CaO 3414

slide102

LE = Lattice Energy

Where: k = proportionality constant dependent on structure of solid and on electron configuration of the ions

Where: Q1 & Q2 = charges on the ions

Where: r = the shortest distance

between the centers of the cation and anion