2. Chemical Bond 1

18, 20 Oct 97 Bonding and Structure 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

polarity, bond order, bond strength


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?


Most bonds are somewhere in between.

Forms of Chemical Bonds• There are 2 extreme forms of connecting

or bonding atoms:

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

•Covalent—electrons shared between atoms


Ionic Ionic BondsBonds

Ionic compounds

- essentially complete electron transfer from an element of low IE (metal) to an element of high electron affinity (EA) (nonmetal)

Na(s) + 1/2 Cl2(g) Na+ + Cl-

NaCl (s)

- NON-DIRECTIONAL bonding via Coulomb (charge) interaction

- primarily between metals (Grps 1A, 2A and transition metals) and nonmetals (esp O and halogens)


Covalent Bonding

Covalent bond is the sharing of the VALENCE ELECTRONS of each atom in a bond

Recall: Electrons are divided between core

and valence electrons.

ATOM core valenceNa 1s2 2s2 2p6 3s1 [Ne] 3s1

Br [Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5

Br Br


Valence Electrons1A1A

2A2A 3A3A 4A4A 5A5A 6A6A 7A7A

8A8A

Number of valence electrons is equal to the Group number.


Covalent BondingThe bond arises from the mutual attraction of

2 nuclei for the same electrons.

HB+ HAHBHA

A covalent bond is a balanceof attractive and repulsive forces.

6_H2bond.mov


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

of atomic orbitals on neighboring atoms.

H H Cl••

••

•• Cl

••

••

••

+

Overlap of H (1s) and Cl (2p)

This type of overlap places bonding electrons in a

MOLECULAR ORBITAL along the line between

the two atoms and forms a SIGMA BOND ().


Sigma Bond Formation by Orbital Overlap

•• ••

sigma bond ( )

+HH

Two s Atomic Orbitals (A.O.s) overlap to form an s (sigma) Molecular Orbital (M.O.)

6_H2pot.mov


Sigma Bond Formation by Orbital Overlap

•• ••

sigma bond ( )

+HH

Two s A.O.s overlap to from an s M.O.

Similarly, two p A.O.s can overlap end-on to from a p M.O.

e.g.F2


Electron Electron Distribution in Distribution in

MoleculesMolecules• Electron distribution

is depicted with

Lewis electron dot structures

• Electrons are distributed as:

• shared or BOND PAIRS and

• unshared or LONE PAIRS.

G. N. Lewis 1875 - 1946


Bond and Lone Pairs• Electrons are distributed as shared or BOND

PAIRS and unshared or LONE PAIRS.

••H Cl

••••

This is a LEWIS ELECTRON DOT structure.

shared or bond pair

Unshared orlone pair (LP)


• This observation is called the OCTET RULE

Rules of Lewis StructuresRules of Lewis Structures• No. of valence electrons of an atom =

Group number

• Except for H (and atoms of 3rd and higher periods),

#Bond Pairs + #Lone Pairs = 4

• For Groups 5A-7A (N - F), no. of BOND PAIRS = 8 - group No.

• For Groups 1A-4A (Li - C), no. of BOND PAIRS = group number


2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons or

1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. In ammonia, N is central

Building a Dot Structure

Ammonia, NH3

4 pairs


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

3. Form a sigma bond between the central atom and surrounding atoms.

H H

H

N

Building a Dot Structure

H••

H

H

N

3 BOND PAIRS and 1 LONE PAIR.

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


Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2

TOTAL = 6 + 18 + 2 = 26 e- or 13 pairs

Step 1. Central atom = S

10 pairs of electrons are left.

Sulfite ion, SO32-

Step 3. Form sigma bonds

O O

O

S


Remaining pairs become lone pairs, first on outside atoms

then on central atom.

Sulfite ion, SO32- (2)

Each atom is surrounded by an octet of electrons.

••O O

O

S

••

••

•• ••

••

••

••

••

••

NOTE - must add formal charges (O-, S+) for complete dot diagram


Carbon Dioxide, CO2

1. Central atom = __C____

2. Valence electrons = _16_ or _8_ pairs

3. Form sigma bonds.

O OC

••O OC

•• ••

••••••

This leaves __6__ pairs.4. Place lone pairs on outer atoms.


••O OC

•• ••

••••••

Carbon Dioxide, CO2 (2)4. Place lone pairs on outer atoms.

••O OC

•• ••

••••••

••O OC

•• ••

••

The second bonding pair forms a pi () bond.

5. To give C an octet, form DOUBLE BONDS between C and O.


SO3

H2CODouble and even triple bonds are commonly observed for C, N, P, O, and S

••O OC

•• ••

••

C2F4


Sulfur Dioxide, SO2

1. Central atom = S

2. Valence electrons = 6 + 2*6 = 18 electrons

or 9 pairs••O OS

••

••

••

••••••

••O OS

••

••

••

••••••

bring inleft pair

OR bring inright pair

3. Form pi () bond so that S has an octet — note that there are two ways of doing this.


Sulfur Dioxide, SO2

••O OS

••

••

••

••••••

bring inleft pair

OR bring inright pair

••O OS

••

••

••

••••

••O OS••

••

••

••

••

Equivalent structurescalled:

RESONANCE STRUCTURES

The proper Lewis structure is a HYBRID of the two.

A BETTER representation of SO2 is made by forming 2 double bonds

O = S = OEach atom has - OCTET - formal charge = 0


Urea (NHUrea (NH22))22COCO1. Number of valence electrons = 24 e-

2. Draw sigma bonds.CN N H

HH

H

O

4. Place remaining electron pairs on oxygen

3. Complete C atom octet with double bond.

CN N H

HH

H

O

Leaves 24 - 14 = 10 e- pairs.

and nitrogen atoms.


Violations of the Octet RuleUsually occurs with:

Boron

BF3 SF4

elements of higher periods.


Boron Trifluoride

• Central atom = B

• Valence electrons = 3 + 3*7 = 24

or electron pairs = 12

• Assemble dot structure

F••

••

••

F

F

B••

••

••

••

••

••

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


Sulfur Tetrafluoride, SF4

• Central atom = S

• Valence electrons = 6 + 4*7 = 34 e-

or 17 pairs.

• Form sigma bonds and distribute electron pairs.

F

••

••

••

F

F

S••

••

••

••

•• F

••

••

••

••

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


Formal charge = Group no. - 1/2 (no. bond electrons)

- (no. of LP electrons)

Formal Atom Charges

• Atoms in molecules often bear a charge (+ or -).

• The most important dominant resonance structure of a molecule is the one with formal charges as close to 0 as possible.


04 - (1/2)(8) - 0 =

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

Carbon Dioxide, CO2

At OXYGEN

O C O••

• •••

• •

At CARBON


C atom charge is 0.

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

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

Carbon Dioxide, CO2 (2)

O C O••

• •••

• •

An alternate Lewis structure is:

AND the corresponding resonance form

+

O C O••

• •••

• •

+


• REALITY: Partial charges calculated

by CAChe molecular modeling system (on CD-ROM).

+1.46-0.73 -0.73

Carbon Dioxide, CO2 (3)

Which is the predominant resonance structure?

O C O••

• •••

• •

ORO C O•

•

• •••

• •

O C O••

• •••

• •

+

+

Answer ?Form without formal charges isBETTER - no +ve charge on O


Boron Trifluoride, BF3

F••

••

••

F

F

B••

••

••

••

••

••

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


Boron Trifluoride, BF3 (2)

• To have +1 charge on F, with its very high electron affinity is not good. -ve charges best placed on atoms with high EA.

• Similarly -1 charge on B is bad

• NOT important Lewis structure

fc = 7 - 2 - 4 = +1 Fluorine

F••

••

F

F

B••

••

••

••

••

••

fc = 3 - 4 - 0 = -1 Boron

+


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

2. Chemical Bond 1

Documents

Transcript of 2. Chemical Bond 1