Chemical Bonding and Molecul

7/27/2019 Chemical Bonding and Molecul

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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-octet

resonance structures

formal charges

• VSEPR - predicting shapes of molecules• Bond properties

electronegativity

polarity, 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 atoms

3. Complete octets by using rest of e- as lone pairs (LP)

4. For atoms with <8 e-, make multiple bonds to complete octets

5. Assign formal charges : fc = Z - (#BP/2) - (#LP)I ndicate equivalent (RESONANCE) structures

6. Structures with smaller formal charges are preferred

- consider non-octet alternatives (esp. for 3rd, 4th row)

—

2 for # 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 in

left pair

OR bring in

right 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 = OSulfur does not obey OCTET rule

BUT 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 structures

is 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

R epulsion 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

H H

H

tetrahedral

109o

C 4

F

120o

planar trigonal F F

B 3

Geometry Example

No. of e- Pairs Around Central Atom

180o linear 2 F

— Be

— F

CAChe

image




Structure Determination by VSEPR

There are 4 electron pairs at the corners of a

tetrahedron.

lone pair of electrons

in tetrahedral position

H

H H

N H

••

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 electrons

in tetrahedral position

H

H H

N




AXnEm notation

•

a good way to distinguish betweenelectron pair and molecular geometries

is the AXnEm notation

where:

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

H

O The electron pair geometry is

TETRAHEDRAL.

H - O - H••

••

VSEPR - water

Water, H2O

1. Draw electron dot structure




H

H

O

. . . the molecular geometry is bent.

VSEPR - water (2)

Although the electron

pair geometry isTETRAHEDRAL . . .

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 aplanar triangle.

Both the electron pair geometry and the

molecular geometry are PLANAR TRIGONAL 120o bond angles.

• •

C HH

• • O

VSEPR - formaldehyde Formaldehyde, CH2O

1. Draw electron dot structure

• • •

CHH

• O

H2CO at the C atom is an AX3 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

H

Angle 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

http://c/chem_eve/6_CH3OH.mov




AXnEm designation ?

at CH3 carbon

at CN carbon

Define bond angles 1 and 2

VSEPR - bond angles (2) Acetonitrile, CH3CN

Angle 1 = ?

H1

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 ATOMSTHAT 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

FB

••

••

•

•

•

•

•

•

•

•

Molecular Geometry is

planar trigonalBF3 is an AX3 species

Consider boron trifluoride, BF3




FF

F

FF

Trigonal bipyramid

120

90

P 5 electron pairs

F

F F

Octahedron90 F

FF 90S6 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 S

THEREFORE:

electron pair geometry ?

F F

F

F

• • • • • •

• • • •

• •

•

• •

••

•• ••

••

•• •• ••

••

S

Sulfur Tetrafluoride, SF4

F

FF

F• •

S

Number of valence e- = 34

No. 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 four

equivalent sp3 hybrid atomic orbitals.

6_CH4.mov

But atomic carbon has an s2p2 configuration

Why can it make more than 2 bonds ?

- Carbon prefers to make 4 bonds as in CH4

http://c/chem_eve/6_CH4.mov




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 s and p Bonding in C2H4

• The extra p orbitalelectron on each C atomoverlaps the p orbital onthe neighboring atom to

form the p bond.

porbital

3 sp2

hybrid

orbitals

2p 2s

C atom orbitals are COMBINED

(= re-hybridized) to form orbitals

better suited for BONDING

• The 3 sp2 hybrid orbitals

are used to make the C-C

and two C-H s bonds6_C2H4-sg.mov

6_C2H4.mov

H HC

H H

sp2 120 C6_C2H4-pi.mov

http://c/chem_eve/6_C2H4.mov

http://c/chem_eve/6_C2H4-pi.mov

http://c/chem_eve/6_C2H4-sg.mov




Consequences of Multiple Bonding

Restricted rotation around C=C bond in

1-butene = CH2=CH-CH2-CH3.

27

233

E ( k J / m o

l )

-180 0 180

C-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.

single

BO = 1

1 s

triple, BO = 31 s and 2 p

double, BO = 2

1 s and 1 p

CH2CHCN

Acrylonitrile




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 bondedatoms.

745 kJ

414 kJ 123 pm

110 pm

Formaldehye




Bond Length

- depends on size of bonded atoms:

Molecule R(H-X)H- F 104 pm

H- Cl 131 pm

H- 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 611 134

CC 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 StrengthHO — 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 positiveend 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 calculatedreal bond 432 kJ/mol measured

ELECTRONEGATIVITY, c.

Difference 92 kJ/mol.

This difference is the contribution of IONIC bonding

It is proportional to the difference in




Electronegativity, c

c is a measure of the ability of an atom in a

molecule to attract electrons to itself.

Concept proposed by

Linus Pauling (1901-94)

Nobel prizes:

Chemistry (54), Peace (63)

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




• F has maximum c.

• Atom with lowest c is the center atom in most

molecules.

• Relative values of c 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

C

NO

SP

Si

Electronegativity, c

Figure 9.7




Bond Polarity

c(A) - c(B) 3.5 - 2.1

c 1.4

+ -+-O—FO—H

c H

2.1

O F

3.5 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

c(O-H) > c(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




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 polar

b) water is non-symmetric

The dipole associated with polar H2Ois 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 whywater attaches to CO2

CO2 + H2O H2CO3

-0.73 +1.46 -0.73




B — F, B — H bonds polar molecule is NOT symmetric

B — F bonds are polar molecule is symmetric

Molecular Polarity inNON-symmetric molecules

F

F F

B

B +ve

F -ve

H

F F

B

Atom Chg. c B +ve 2.0

H +ve 2.1

F -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(C-F) = 1.5, c(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 structuresoctet vs. non-octet

resonance structures

formal charges

•

VSEPR - predicting shapes of molecules• Bond properties

electronegativity

polarity, bond order, bond strength

Chemical Bonding and Molecul

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