Chemical Bonding and Molecular Structure (Chapter 9)

20 Oct 97 Bonding and structure (2) 1

Chemical Bonding and

Molecular Structure (Chapter 9)

• Ionic vs. covalent bonding• Molecular orbitals and the covalent bond (Ch. 10)• Valence electron Lewis dot structures

octet vs. non-octetresonance structuresformal charges

• VSEPR - predicting shapes of molecules• Bond properties

electronegativitypolarity, bond order, bond strength


•OCTET RULE: #Bond Pairs + #Lone Pairs = 4 (except for H and atoms of 3rd and higher periods)

Rules for making Lewis dot structures

2. Place a bond pair (BP) between connected atoms3. Complete octets by using rest of e- as lone pairs (LP)4. For atoms with <8 e-, make multiple bonds to complete octets5. Assign formal charges : fc = Z - (#BP/2) - (#LP) Indicate equivalent (RESONANCE) structures6. Structures with smaller formal charges are preferred - consider non-octet alternatives (esp. for 3rd, 4th row)

— 2 for # of PAIRS# of PAIRS

1. Count no. of valence electrons (- don’t forget to include the charge on molecular ions!)

#lone pairs at central atom in AXn = {(#e-) - 8*n}/2


Sulfur Dioxide, SO2

••O OS

••

••

••

••••••

bring inleft pair

OR bring inright pair

These equivalent structuresare called:

RESONANCE STRUCTURES. The proper Lewis structure

is a HYBRID of the two.

Each atom has OCTET . . . . . BUT there is a +1 and -1 formal charge

••O OS

••

••

••

••••

••O OS••

••

••

••

••

+— —+

Rules 1-3 O—S —O


SO2 (2)

Alternate Lewis structure for SO2 uses 2 double bonds

NB: # of central atom lone pairs = (3*6 -8*2)/2 = 1 in both O=S+-O- and O=S=O structures

O = S = O Sulfur does not obey OCTET ruleBUT the formal charge = 0

This is better structure than O=S+-O- since it reduces formal charge (rule 6). 3rd row S atom can have 5 or 6 electron pairs


A. S=C=N

Thiocyanate ion, (SCN)-

Which of three possible resonance structuresis most important?

-0.52 -0.32-0.16

Calculated partial charges

ANSWER:

C > A > B

C. S-C N

B. S=C - N


VSEPR • Valence Shell Electron Pair

Repulsion theory.

• Most important factor in determining geometry is relative repulsion between electron pairs.

MOLECULAR GEOMETRY

Molecule adopts the shape that minimizes the electron pair repulsions.

6_VSEPR.mov


H

HH

H

tetrahedral

109o

C4

F

120o

planar trigonalFF

B3

GeometryExample

No. of e- PairsAround CentralAtom

180o

linear2 F—Be—F

CACheimage


Structure Determination by VSEPR

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

lone pair of electronsin tetrahedral position

H

H HNH

••

H

H

N

The ELECTRON PAIR GEOMETRY is tetrahedral.

Ammonia, NH3


Although the electron pair geometry is tetrahedral . . .

VSEPR - ammonia

Ammonia, NH3

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

lone pair of electronsin tetrahedral position

H

H HN


AXnEm notation

• a good way to distinguish between electron pair and molecular geometries is the AXnEm notationwhere: A - atom whose local geometry is of interest

(typically the CENTRAL ATOM) Xn - n atoms bonded to A Em - m lone pair electrons at A

NH3 is AX3E system pyramidal

(NB this notation not used by Kotz)


2. Count BP’s and LP’s = 4

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

H

HO

The electron pair geometry is TETRAHEDRAL.

H - O - H••

••

VSEPR - water

Water, H2O

1. Draw electron dot structure


H

HO

. . . the molecular . . . the molecular geometry is bent.geometry is bent.

VSEPR - water (2)

Although the electron pair geometry is TETRAHEDRAL . . .

H - O - H••

••

H2O - AX2E2 system - angular geometry


2. Count BP’s and LP’s: At Carbon there are 4 BP but . . .

3. These are distributed in ONLY 3 regions. Double bond electron pairs are in same region.There are 3 regions of electron densityElectron repulsion places them at the corners of a planar triangle.

Both the electron pair geometry and the molecular geometry are PLANAR TRIGONAL 120o bond angles.

••

C HH

••OVSEPR - formaldehyde Formaldehyde, CH2O

1. Draw electron dot structure

• ••

CHH

•O

H2CO at the C atom is an AXAX33 species species


AXnEm designation ?

at C

at O

Define bond angles 1 and 2

Angle 1 = H-C-H = ?

Angle 2 = H-O-C = ?

Answer:

VSEPR - Bond Angles

H

HAngle 2

••H—C—O—H

Angle 1

6_CH3OH.mov

Methanol, CH3OH

109o because both the C and O atoms are surrounded by 4 electron pairs.

••

AX4 = tetrahedral

AX2E2 = bent


AXnEm designation ?

at CH3 carbon

at CN carbon

Define bond angles 1 and 2

VSEPR - bond angles (2) Acetonitrile, CH3CN

Angle 1 = ?

H 1

H—C—C

2

H

••N

Why ? : The CH3 carbon is surrounded by 4 bond charges

The CN carbon is surrounded by 2 bond charges

Angle 2 = ?

AX4 = tetrahedral

AX2 = linear

109o

180o


What about:STRUCTURES WITH CENTRAL ATOMS

THAT DO NOT OBEY THE OCTET RULE ?

BF3

SF4

PF5


Geometry for non-octet species also obey VSEPR rules

The B atom is surrounded by only 3 electron pairs.

Bond angles are 120o F••

••

••

F

F

B••

••

••

••

••

••

Molecular Geometry is planar trigonalBF3 is an AX3 species

Consider boron trifluoride, BF3


FF

F

FF

Trigonal bipyramid

120

90

P 5 electron pairs

FF F

Octahedron90 F

FF90S6 electron pairs

Compounds with 5 or 6 Pairs Around the Central Atom 6_VSEPR.mov

AX5 system

AX6 system


There are 5 (BP + LP)e- pairs around the STHEREFORE:electron pair geometry ?

F

F

F

F

••• •••

••••••

•••

••

••••

••

••••••

••

S

Sulfur Tetrafluoride, SF4

F

FFF•• S

Number of valence e- = 34No. of S lone pairs =

{17 - 4 b.p. - 3x4 l.p.(F)}= 1 lone pair on S

= trigonal bipyramid

AX4E system. Molecular geometry ?

F

FFF

••

SOR


120

90

F

F

F

F•• S

Sulfur Tetrafluoride, SF4 (2)

axial

equatorial

Molecular geometry of SF4 is “see-saw”

Q: What is molecular geometry of SO2 ?

Lone pair is in the equatorial position because it requires more room than a bond pair.


Bonding with Hybrid Atomic Orbitals

4 C atom orbitals hybridize to form four4 C atom orbitals hybridize to form four

equivalent spequivalent sp33 hybrid atomic orbitals. hybrid atomic orbitals.

6_CH4.mov

But atomic carbon has an s2p2 configurationWhy can it make more than 2 bonds ?

- Carbon prefers to make 4 bonds as in CH4


Orbital Hybridization

BONDS SHAPE HYBRID REMAIN e.g.

2 linear {2 x sp &2 p’s} C2H2

3 trigonal {3 x sp2 & 1 p} C2H4

planar

4 tetrahedral {4 xsp3 } CH4

s2p2


Multiple Bonds and Bonding in C2H4

• The extra p orbital electron on each C atom overlaps the p orbital on the neighboring atom to form the bond.

p orbital

3 sp2

hybrid orbitals

2p2s

C atom orbitals are COMBINED (= re-hybridized) to form orbitalsbetter suited for BONDING

• The 3 sp2 hybrid orbitals are used to make the C-C and two C-H bonds

6_C2H4-sg.mov

6_C2H4.mov

H HC

H Hsp2120 C6_C2H4-pi.mov


Consequences of Multiple Consequences of Multiple BondingBonding

Restricted rotation around C=C bond in1-butene = CH2=CH-CH2-CH3.

27

233

E (

kJ/m

ol)

-180 0 180C-C=C angle (o)

P. 475 - Photo-rotation about double bonds lets us see !!

See Butene.Map in ENER_MAP in CAChe models.


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

molecular properties ?

Buckyball in HIV-protease, see page 107

- bond order - bond length - bond strength - bond polarity - MOLECULAR polarity


H

H

H

C C NC

Bond Order

• the number of bonds between a pair of atoms.

singleBO = 11

triple, BO = 31 and 2

double, BO = 21 and 1

CH2CHCNAcrylonitrile


Bond order = Total # of e - pairs used for a type of bond

Total # of bonds of that type

Bond Order (2)

Fractional bond orders occur in molecules with resonance structures.

Consider NO2-

Bond order in NO2- = 3 (e - pairs in N-O bonds)

2 (N - O bonds)N-O bond order in NO2

- = 1.5

O O O O

N••

••••

••••

••••••••••

••••

••N


Bond Order and Bond Length

Bond order is related to two important bond properties:

(a) bond strength

as given by DE

(b) Bond length - the distance between the nuclei of two bonded atoms.

745 kJ745 kJ

414 kJ414 kJ 123 pm123 pm

110 pm110 pm

Formaldehye


Bond Length

- depends on size of bonded atoms:

Molecule R(H-X)H- F 104 pmH- Cl 131 pmH- I 165 pm

- depends on bond order.

Molecule R(C-O)CH3C- OH 141 pmO=C=O 132 pm

C O 119 pm


Bond Strength• Bond Dissociation energy (DE) - energy required to

break a bond in gas phase. • See Table 9.5

BOND STRENGTH (kJ/mol) LENGTH (pm)H—H 436 74

C—C 347 154C=C 611 134CC 837 121

NN 946 110

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


Bond Strength (2)

Bond Order Length Strength

HO—OH 1 149 pm 210 kJ/mol

O=O 2 121 498 kJ/mol

O O•••••••

••••

••O 1.5 128 ?

HOW TO CALCULATE ?

Hrxn = {3xHf(O) - Hf(O3)} = {3x249.2 - 142.7} = 605 kJ/mol 2 O-O bonds in O3 DE (O3) = 605/2 = 302.5 kJ/mol

O3 (g) 3 O(g)

303 kJ/mol


Bond Polarity

HCl is POLAR because it has a positive end and a negative end (partly ionic).

Polarity arises because Cl has a greater share of the bonding electrons than H.

Cl

-+

•••H••

••

Calculated charge by CAChe:

H (red) is +ve (+0.20 e-)

Cl (yellow) is -ve (-0.20 e-).

(See PARTCHRG folder in MODELS.)


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

Cl

-+

•••H••

••

Bond Polarity (2)

BOND ENERGY

“pure” bond 339 kJ/mol calculated

real bond 432 kJ/mol measured

ELECTRONEGATIVITY, .

Difference 92 kJ/mol.

This difference is the contribution of IONIC bondingIt is proportional to the difference in


Electronegativity,

is a measure of the ability of an atom in a molecule to attract electrons to itself.

Concept proposed byLinus Pauling (1901-94)Nobel prizes:Chemistry (54), Peace (63)

See p. 425; 008vd3.mov (CD)


• F has maximum .

• Atom with lowest is the center atom in most molecules.

• Relative values of determines BOND POLARITY (and point of attack on a molecule).

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 180

0.5

1

1.5

2

2.5

3

3.5

4

H

FCl

CN

O

SP

Si

Electronegativity,

Figure 9.7


Bond Polarity

(A) - (B) 3.5 - 2.1

1.4

+ -+-O—FO—H

H 2.1O F3.5 4.0

Also note that polarity is “reversed.”

Which bond is more polar ? (has larger bond DIPOLE)

O—H O—F

3.5 - 4.0 0.5

(O-H) > (O-F) Therefore OH is more polar than OF


Molecular Polarity

• Molecules—such as HCl and H2O—

can be POLAR (or dipolar).

• They have a DIPOLE MOMENT.

• Polar molecules turn to align their dipole with an electric field.

POSITIVE

NEGATIVE

H—Cl

POSITIVE

NEGATIVE

H—Cl


Predicting molecular polarity

A molecule will be polar ONLY if

a) it contains polar bonds

AND

b) the molecule is NOT “symmetric”

Symmetric molecules


H

HH H

O

••

••

O

+polar

Molecular Polarity: H2O

Water is polar because:

a) O-H bond is polarb) water is non-symmetric

The dipole associated with polar H2O is the basis for absorption of microwaves used in cooking with a microwave oven


Carbon Dioxide

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

• Because CO2 is symmetrical the BOND polarity cancels

The positive C atom is why water attaches to CO2

CO2 + H2O H2CO3

-0.73 +1.46 -0.73


B—F, B—H bonds polarmolecule is NOT symmetric

B—F bonds are polarmolecule is symmetric

Molecular Polarity in NON-symmetric molecules

F

F FB

B +ve F -ve

H

F FB

Atom Chg. B +ve 2.0H +ve 2.1F -ve 4.0

BF3 is NOT polar HBF2 is polar


Fluorine-substituted Ethylene: C2H2F2

CIS isomer

• both C—F bonds on same side

molecule is POLAR.

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

TRANS isomer

• both C—F bonds on opposite side

molecule is NOT POLAR.

(C-F) = 1.5, (C-H) = 0.4


Chemical Bonding and

Molecular Structure (Chapter 9)

• Ionic vs. covalent bonding• Molecular orbitals and the covalent bond (Ch. 10)• Valence electron Lewis dot structures

octet vs. non-octetresonance structuresformal charges

• VSEPR - predicting shapes of molecules• Bond properties

electronegativitypolarity, bond order, bond strength

Documents

Chemical Bonding and Molecular Structure (Chapter 9)