Chemical Bonding and Molecular Structure

Bonding and Structure 1

Chemical Bonding and

Molecular Structure

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

octet vs. non-octetresonance structuresformal charges

• VSEPR (Valence shell electron pair repulsion) - predicting shapes of molecules• Bond properties

polarity, bond order, bond strength

Bonding and Structure 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?

Bonding and Structure 3

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

Bonding and Structure 4

Ionic Ionic BondsBonds

Ionic compounds - essentially complete electron

transfer from an element of low (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 (Groups 1A, 2A and transition metals) and nonmetals (esp O and halogens)

Bonding and Structure 5

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

Bonding and Structure 6

Valence Electrons1A1A

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

8A8A

Number of valence electrons is equal to the Group number.

Bonding and Structure 7

Covalent BondingThe bond arises from the mutual attraction of

2 nuclei for the same electrons.

HB+ HA HBHA

A covalent bond is a balanceof attractive and repulsive forces.

Bonding and Structure 8

←Fig. (a) In the nonpolar covalent bond present, there is a symmetrical distribution of electron density. (b) In the polar covalent bond present, electron density is displaced because of its electronegativity.

Polar and non-polar covalent bond

Dipole moment, µ= e (esu) x d (angstrom)

Greater the DM greater the polarity

Bonding and Structure 9

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

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 ().

S-s, s-p and p-p orbitals form sigma bond.

Bonding and Structure 10

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

Bonding and Structure 11

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

Bonding and Structure 12

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

Bonding and Structure 13

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)

Bonding and Structure 14

• 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

Bonding and Structure 15

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 StructureAmmonia, NH3

4 pairs

Bonding and Structure 16

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

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

H HHN

Building a Dot Structure

H•• HHN

3 BOND PAIRS and 1 LONE PAIR.

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

Bonding and Structure 17

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

Bonding and Structure 18

Thanks to all

Bonding and Structure 19

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

Bonding and Structure 20

Carbon Dioxide, CO2

1. Central atom = __C____2. Valence electrons = _16_ or _8_ pairs3. Form sigma bonds.

O OC

••O OC

•• ••

••••••

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

Bonding and Structure 21

••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.

Bonding and Structure 22

SO3

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

••O OC

•• ••

••C2F4

Bonding and Structure 23

Sulfur Dioxide, SO21. Central atom = S2. 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.

Bonding and Structure 24

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

Bonding and Structure 25

Urea (NHUrea (NH22))22COCO1. Number of valence electrons = 24 e-2. Draw sigma bonds.

CN N HHH

H

O

4. Place remaining electron pairs on oxygen3. Complete C atom octet with double bond.

CN N HHH

H

O

Leaves 24 - 14 = 10 e- pairs.

and nitrogen atoms.

Bonding and Structure 26

Violations of the Octet RuleUsually occurs with:

Boron

BF3 SF4

elements of higher periods.

Bonding and Structure 27

Boron Trifluoride• Central atom = B• Valence electrons = 3 + 3*7 = 24 or electron pairs = 12• Assemble dot structure

F••

••

••

F

FB

••

••

••

••

••

••

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

Bonding and Structure 28

Sulfur Tetrafluoride, SF4

• Central atom = S• Valence electrons = 6 + 4*7 = 34 e-

or 17 pairs.• Form sigma bonds and distribute

electron pairs.

F

••

••

••

FF

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.

Bonding and Structure 29

Explanation of the failure of octet rule

Sidgwick’s rule of maximum covalency

• It is not essential to have 8 electron surrounding in an atom

• Maximum covalency depends on its period

period Maximum shared electron

Atom

n=1 2 H

n=2 4 Li, F

n=3 6 Na, Mg

n=4 8 k

Bonding and Structure 30

Sugden’s view of singlet linkages

Octet rule never violated

In PCl5, 3 Cl atom linked by normal covalent bond and others 2 by sharing only 1 electron called singlet linkage

Formal Charges• Formal charge is the charge calculated

for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs.

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• We will calculate the formal charge for each atom in this Lewis structure.

Formal charge of HFormal charge of H

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• Hydrogen shares 2 electrons with oxygen.• Assign 1 electron to H and 1 to O.• A neutral hydrogen atom has 1 electron.• Therefore, the formal charge of H in nitric

acid is 0.

Formal charge of HFormal charge of H

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• Oxygen has 4 electrons in covalent bonds.• Assign 2 of these 4 electrons to O.• Oxygen has 2 unshared pairs. Assign all 4 of

these electrons to O.• Therefore, the total number of electrons assigned

to O is 2 + 4 = 6.

Formal charge of OFormal charge of O

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• Electron count of O is 6.• A neutral oxygen has 6 electrons.• Therefore, the formal charge of O is 0.

Formal charge of OFormal charge of O

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons).

• A neutral oxygen has 6 electrons.• Therefore, the formal charge of O is 0.

Formal charge of OFormal charge of O

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons).

• A neutral oxygen has 6 electrons.• Therefore, the formal charge of O is -1.

Formal charge of OFormal charge of O

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• Electron count of N is 4 (half of 8 electrons in covalent bonds).

• A neutral nitrogen has 5 electrons.• Therefore, the formal charge of N is +1.

Formal charge of NFormal charge of N

––

Nitric acid

.... ::

....HH OO

OO

OO

NN

::

::....

....

• A Lewis structure is not complete unless formal charges (if any) are shown.

Formal chargesFormal charges

––++

Formal ChargeFormal Charge

Formal charge = Formal charge =

group numbergroup numberin periodic tablein periodic table

number ofnumber ofbondsbonds

number ofnumber ofunshared electronsunshared electrons–– ––

An arithmetic formula for calculating formal charge.An arithmetic formula for calculating formal charge.

"Electron counts""Electron counts" and formal and formal charges in NHcharges in NH44

+ + and BFand BF44--

11

44

NN

HH

HH HH

HH

++77

44

....

BBFF

FF

FF

FF

....

............:: ::

:: ::

:: ::

....

––

Resonance

Resonance

In chemistry, resonance or mesomerism is a way of

describing delocalized electrons within certain molecules

or polyatomic ions where the bonding cannot be

expressed by one single Lewis formula. A molecule or ion

with such delocalized electrons is represented by several

contributing structures (also called resonance structures

or canonical forms).

Bonding and Structure 44

General characteristics of resonanceMolecules and ions with resonance (also called

mesomerism) have the following basic characteristics:•They can be represented by several correct Lewis formulas, called

"contributing structures", "resonance structures“ or "canonical forms".

However, the real structure is not a rapid interconversion of contributing

structures. Several Lewis structures are used together, because none of

them exactly represents the actual structure. To represent the

intermediate, a resonance hybrid is used instead.

•The contributing structures are not isomers. They differ only in the

position of electrons, not in the position of nuclei.

Bonding and Structure 45

• Each Lewis formula must have the same number of valence electrons (and thus the same total charge), and the same number of unpaired electrons, if any.

• Bonds that have different bond orders in different

contributing structures do not have typical bond lengths. Measurements reveal intermediate bond lengths.

• The real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be.

General characteristics of resonance

Bonding and Structure 46

Contributing structures of some ion

•same atomic positions

•differ in electron positions

more stable more stable Lewis Lewis

structurestructure

less stable less stable Lewis Lewis

structurestructure

........

CC OO NN OOHH

HH

HH.... ::....

++ ––....

....CC OO NN OOHH

HH

HH

....::....

Resonance Structures of Methyl Nitrite

•same atomic positions

•differ in electron positions

more stable more stable Lewis Lewis

structurestructure

less stable less stable Lewis Lewis

structurestructure

........

CC OO NN OOHH

HH

HH.... ::....

++ ––....

....CC OO NN OOHH

HH

HH

....::....

Resonance Structures of Methyl Nitrite

• Electrons in molecules are often delocalizedbetween two or more atoms.

• Electrons in a single Lewis structure are assigned to specific atoms-a single Lewis structure is insufficient to show electron delocalization.

• Composite of resonance forms more accurately depicts electron distribution.

Why Write Resonance Structures?

•Ozone (O3)

–Lewis structure of

ozone shows one

double bond and

one single bond

Expect: one short bond and one Expect: one short bond and one long bondlong bond

Reality: bonds are of equal length Reality: bonds are of equal length (128 pm)(128 pm)

Example

OO OO••••OO••••••••••••

••••••••––++

•Ozone (O3)

–Lewis structure of

ozone shows one

double bond and

one single bond

Resonance:Resonance:

Example

OO OO••••OO••••••••••••

••••••••––++

OO OO••••OO••••••••••••

••••••••––++

OO OOOO••••••••••••

••••••••–– ++

••••

Bonding and Structure 52

Bonding and Structure 53

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.

Bonding and Structure 54

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

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

Carbon Dioxide, CO2

At OXYGEN

O C O••

• •••

• •

At CARBON

Bonding and Structure 55

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••

• •••

• •

+

Bonding and Structure 56

• 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

Bonding and Structure 57

Boron Trifluoride, BF3

F••

••

••

F

FB

••

••

••

••

••

••

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

Bonding and Structure 58

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

FB

••

••

••

••

••

••

fc = 3 - 4 - 0 = -1 Boron

+

Bonding and Structure 59

A. S=C=N

Thiocyanate ion, (SCN)-Which of three possible resonance structuresis most important?

-0.52 -0.32-0.16

Calculated partial chargesANSWER:C > A > B

C. S-C N

B. S=C - N