1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 9 Bonding and Molecular

1

© 2006 Brooks/Cole - Thomson

Chemistry and Chemical Reactivity 6th Edition

John C. Kotz Paul M. Treichel

Gabriela C. Weaver

CHAPTER 9

Bonding and Molecular Structure:Fundamental Concepts

© 2006 Brooks/Cole Thomson

Lectures written by John Kotz

cellE

2


Chemical Bonding

Problems and questions —How is a molecule or

polyatomic ion held together?

Why are atoms distributed at strange angles?

Why are molecules not flat?Can we predict the structure?How is structure related to

chemical and physical properties?

3


Structure & Bonding

NN triple bond. Molecule is unreactive

Phosphorus is a tetrahedron of P atoms. Very reactive!

Red phosphorus, a polymer. Used in matches.

D:%5CMedia%20Kotz%206e%201%20to%2010%5CMovies%5C04M04VD1.MOV


4


Forms of Chemical Bonds

• There are 2 extreme forms of connecting or bonding atoms:

• Ionic—complete transfer of 1 or more electrons from one atom to another

• Covalent—some valence electrons shared between atoms

• Most bonds are somewhere in between.

5


Ionic CompoundsMetal

Nonmetal

2 Na(s) + Cl2(g) ---> 2 Na+ + 2 Cl-

6


Covalent BondingThe bond arises from the mutual attraction of

2 nuclei for the same electrons. Electron sharing results.

Bond is a balance of attractive and repulsive forces.

D:%5CMedia%20Kotz%206e%201%20to%2010%5CMovies%5C09M03AN1.MOV

7


Bond FormationA bond can result from a “head-to-head”

overlap of atomic orbitals on neighboring atoms.

ClH H Cl••

••

••

••

••

••

+

Overlap of H (1s) and Cl (3p)

Note that each atom has a single, unpaired electron.

8


Chemical Bonding: Objectives

Objectives are to understand:

1. valence e- distribution in molecules and ions.

2. molecular structures3. bond properties and their

effect on molecular properties.

9


Electron Distribution in Molecules

• Electron distribution is

depicted with Lewis electron dot structures

• Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS.

G. N. Lewis 1875 - 1946

10


Bond and Lone Pairs• Valence electrons are distributed as

shared or BOND PAIRS and unshared or LONE PAIRS.

•

••

•

••

H Cllone pair (LP)

shared orbond pair

This is called a LEWIS ELECTRON DOT structure.

11


Valence ElectronsElectrons are divided between core and valence

electronsB 1s2 2s2 2p1

Core = [He] , valence = 2s2 2p1

Br [Ar] 3d10 4s2 4p5

Core = [Ar] 3d10 , valence = 4s2 4p5

12


Rules of the GameNo. of valence electrons of a main

group atom = Group number

•For Groups 1A-4A, no. of bond pairs = group number.

• For Groups 5A -7A, BP’s = 8 - Grp. No.

13


Rules of the Game•Except for H (and sometimes atoms of 2nd and 3rd families and 3rd and higher periods),

BP’s + LP’s = 4

This observation is called the

OCTET RULE

14


Building a Dot StructureAmmonia, NH3

1. Decide on the central atom; never H. Central atom is atom of lowest affinity

for electrons. Therefore, N is central2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs

15


3. Form a single bond between the central atom and each surrounding atom

H H

H

N

Building a Dot Structure

H••

H

H

N4. Remaining electrons form LONE PAIRS to complete octet as needed.

3 BOND PAIRS and 1 LONE PAIR.

Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

16


10 pairs of electrons are now left.

Sulfite ion, SO32-

Step 1. Central atom = SStep 2. Count valence electrons

S = 6 3 x O = 3 x 6 = 18 Negative charge = 2

TOTAL = 26 e- or 13 pairsStep 3. Form bonds

17


Sulfite ion, SO32-

Remaining pairs become lone pairs, first on outside atoms and then on central atom.

O O

O

S••

••

••

••

•• ••

••

••

••

••

Each atom is surrounded by an octet of electrons.

18


Carbon Dioxide, CO2

1. Central atom = _______2. Valence electrons = __ or __ pairs3. Form bonds.

4. Place lone pairs on outer atoms.

This leaves 6 pairs.

19


Carbon Dioxide, CO2

4. Place lone pairs on outer atoms.

The second bonding pair forms a pi (π) bond.

5. So that C has an octet, we shall form DOUBLE BONDS between C and O.

20


Double and even triple bonds are commonly observed for C, N, P, O, and S

H2CO

SO3

C2F4

21


Sulfur Dioxide, SO2

1. Central atom = S2. Valence electrons = 18 or 9 pairs

bring inleft pair

OR bring inright pair

3. Form double bond so that S has an octet — but note that there are two ways of doing this.

•• OS

••

••

••

••••••O

22


Sulfur Dioxide, SO2

This leads to the following structures.

These equivalent structures are called RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two.

23


Formal Atom Charges

• Atoms in molecules often bear a charge (+ or -). • The predominant resonance structure of a

molecule is the one with charges as close to 0 as possible.

• Formal charge = Group number – 1/2 (no. of bonding electrons)

- (no. of LP electrons)

24


••

••O OC

••••

+4 - (1/2)(8) - 0 = 0

+6 - (1/2)(4) - 4 = 0

Carbon Dioxide, CO2

25


Thiocyanate Ion, SCN-

6 - (1/2)(2) - 6 = -1 5 - (1/2)(6) - 2 = 0

4 - (1/2)(8) - 0 = 0

••

•

••S NC

•••

26


Thiocyanate Ion, SCN-

••

•

••S NC

•••

••

•

••S NC

•••

•••

••S NC

•••

Which is the most important resonance form?

27


Calculated Partial Charges in SCN-

All atoms negative, but most on the S

••

•

••S NC

•••

28


Violations of the Octet Rule

Usually occurs with B and elements of higher periods.

BF3 SF4

29


Boron Trifluoride• Central atom = _____________• Valence electrons = __________ or

electron pairs = __________• Assemble dot structure

The B atom has a share in only 6 pairs of electrons (or 3 pairs). B atom in many molecules is electron deficient.

30


Boron Trifluoride, BF3

What if we form a B—F double bond to satisfy the B atom octet?

F••

••

••

F

F

B••

••

••

••

••

+1

-1

31


Sulfur Tetrafluoride, SF4• Central atom =

• Valence electrons = ___ or ___ pairs.• Form sigma bonds and distribute electron

pairs.

5 pairs around the S atom. A common occurrence outside the 2nd period.

MOLECULAR GEOMETRY

33


VSEPR • Valence Shell Electron

Pair Repulsion theory.• Most important factor in

determining geometry is relative repulsion between electron pairs.

Molecule adopts the shape that minimizes the electron pair repulsions.

MOLECULAR GEOMETRY



34


Electron Pair GeometriesActive Figure 9.8

35




36




37




38


Electron Pair GeometriesActive Figure 9.8

39


Structure Determination by

VSEPRAmmonia, NH3

1. Draw electron dot structure2. Count BP’s and LP’s = 4

H••

H

H

N

3. The 4 electron pairs are at the corners of a tetrahedron.

40



VSEPRAmmonia, NH3

There are 4 electron pairs at the corners of a tetrahedron.

The ELECTRON PAIR GEOMETRY is tetrahedral.

41


Ammonia, NH3

The electron pair geometry is tetrahedral.


VSEPR

The MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDAL.



42



VSEPRWater, H2O

1. Draw electron dot structure

The electron pair geometry is TETRAHEDRAL.

2. Count BP’s and LP’s = 4

3. The 4 electron pairs are at the corners of a tetrahedron.

43



VSEPRWater, H2O

The electron pair geometry is TETRAHEDRAL

The molecular geometry is BENT.

44


Geometries for Four Electron Pairs

Figure 9.9

45



VSEPRFormaldehyde, CH2O

1. Draw electron dot structure

The electron pair geometry is PLANAR TRIGONAL with 120o bond angles.

••

C HH

O••

••

CHH

O••

2. Count BP’s and LP’s at C3. There are 3 electron “lumps” around C at the corners of a planar triangle.

46



VSEPRFormaldehyde, CH2O

The electron pair geometry is PLANAR TRIGONAL

The molecular geometry is also planar trigonal.

47


H-C-H = 109o

C-O-H = 109o

In both cases the atom is surrounded by 4 electron pairs.


VSEPRH

H

••

••H—C—O—H

109˚ 109˚

Methanol, CH3OH

Define H-C-H and C-O-H bond angles



48



VSEPRAcetonitrile, CH3CN

••

H

H

H—C—C N

180˚109˚H-C-H = 109o C-C-N = 180o

One C is surrounded by 4 electron “lumps” and the other by 2 “lumps”

Define unique bond angles

49


Phenylalanine, an amino acid

50


Phenylalanine

51


Structures with Central Atoms with More Than or Less Than 4 Electron Pairs

Often occurs with Group 3A elements and with those of 3rd period and higher.

52


Boron Compounds

Consider boron trifluoride, BF3

Geometry described as

planar trigonal

The B atom is surrounded by only 3 electron pairs.

Bond angles are 120o

53


5 electron pairs

Compounds with 5 or 6 Pairs Around the Central Atom



54


Molecular Geometries for Five Electron

PairsFigure 9.11

All based on trigonal bipyramid

55


• Number of valence electrons = 34

• Central atom = S• Dot structure

Sulfur Tetrafluoride, SF4

Electron pair geometry--> trigonal bipyramid (because there are 5 pairs around the S)

56


Lone pair is in the equator because it requires more room.

Sulfur Tetrafluoride, SF4

57


Molecular Geometries

for Six Electron

PairsFigure 9.14

All are based on the 8-sided octahedron

58


6 electron pairs

Compounds with 5 or 6 Pairs Around the Central Atom



59


Bond Properties• What is the effect of bonding and structure on

molecular properties?

Free rotation around C–C single bond

No rotation around C=C double bond

60


Bond Order # of bonds between a pair of atoms

Double bondSingle bond

Triple bond

Acrylonitrile

61


Bond OrderFractional bond orders occur in molecules with resonance

structures.

Consider NO2-

The N—O bond order = 1.5

62


Bond OrderBond order is proportional to two important bond

properties:

(a) bond strength(b) bond length

745 kJ

414 kJ

110 pm

123 pm

63


Bond LengthBond length depends on

size of bonded atoms.

H—F

H—Cl

H—I

Bond distances measured in Angstrom units where 1 A = 10-2 pm.

64


Bond length depends on bond order.

Bond distances measured in Angstrom units where 1 A = 10-2 pm.

Bond Length

65


• —measured by the energy req’d to break a bond. See Table 9.10.

• BOND STRENGTH (kJ/mol) H—H 436 C—C 346 C=C 602 CC 835 NN 945

The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond.

Bond Strength

66


Bond Order Length Strength

HO—OH

O=O

1 142 pm 210 kJ/mol

2 121 498

1.5 128 ?

Bond Strength

67


Molecular Polarity

Why do ionic compounds dissolve in water?

WaterBoiling point =

100 ˚C

MethaneBoiling point

= -161 ˚C

Why do water and methane differ so much in their

boiling points?

68


Bond Polarity

HCl is POLAR because it has a positive end and a negative end.

Cl has a greater share in bonding electrons than does H.

Cl has slight negative charge (-d) and H has slight positive charge (+ d)

69


Bond Polarity

• Three molecules with polar, covalent bonds.

• Each bond has one atom with a slight negative charge (-d) and and another with a slight positive charge (+ d)

70


This model, calc’d using CAChe

software for molecular

calculations, shows that H is +

(red) and Cl is - (yellow). Calc’d

charge is + or - 0.20.

Bond Polarity

71


Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond.

BOND ENERGY“pure” bond339 kJ/mol calc’dreal bond 432 kJ/mol measured

Difference = 92 kJ. This difference is proportional to the difference in

ELECTRONEGATIVITY, .

Bond Polarity

72


Electronegativity, is a measure of the ability of an atom in a

molecule to attract electrons to itself.

Concept proposed by Linus Pauling 1901-1994

73


Linus Pauling, 1901-1994

The only person to receive two unshared Nobel prizes (for Peace and Chemistry).

Chemistry areas: bonding, electronegativity, protein structure



74


ElectronegativityFigure 9.14

75


• F has maximum .• Atom with lowest is the center atom in

most molecules.• Relative values of determine BOND

POLARITY (and point of attack on a molecule).

Electronegativity,

See Figure 9.14

76


Bond PolarityWhich bond is more polar (or DIPOLAR)? O—H O—F 3.5 - 2.1 3.5 - 4.0 1.4 0.5 OH is more polar than OF

and polarity is “reversed.”

77


Molecular PolarityMolecules—such as HI and H2O—

can be POLAR (or dipolar).

They have a DIPOLE MOMENT. The polar HCl molecule will turn to align with an electric field.

78


Molecular PolarityThe magnitude of the dipole is

given in Debye units.

Named for Peter Debye

(1884 - 1966). Rec’d 1936

Nobel prize for work on x-

ray diffraction and dipole

moments.

79


Dipole Moments

Why are some molecules polar but others are not?

80


Molecular PolarityMolecules will be polar ifa)bonds are polar ANDb) the molecule is NOT “symmetric”

All above are NOT polar

81


Polar or Nonpolar?Compare CO2 and H2O. Which one is polar?

82


Carbon Dioxide

• CO2 is NOT polar even though the CO bonds are polar.

• CO2 is symmetrical.

+1.5 -0.75-0.75

Positive C atom is reason CO2 and H2O react to give H2CO3

83


Polar or Nonpolar?• Consider AB3 molecules: BF3, Cl2CO, and NH3.

84


Molecular Polarity, BF3

B atom is positive and F atoms are negative.

B—F bonds in BF3 are polar.

But molecule is symmetrical and NOT polar

85


Molecular Polarity, HBF2

B atom is positive but H & F atoms are negative.

B—F and B—H bonds in HBF2 are polar. But molecule is NOT symmetrical and is polar.

86


Is Methane, CH4, Polar?

Methane is symmetrical and is NOT polar.

87


Is CH3F Polar?

C—F bond is very polar. Molecule is not symmetrical and so is polar.

88


CH4 … CCl4

Polar or Not?

• Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”

89


Substituted Ethylene

• C—F bonds are MUCH more polar than C—H bonds.

• Because both C—F bonds are on same side of molecule, molecule is POLAR.

90


Substituted Ethylene

• C—F bonds are MUCH more polar than C—H bonds.

• Because both C—F bonds are on opposing ends of molecule, molecule is NOT POLAR.

Documents

1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 9 Bonding and Molecular