Lecture 24 © slgCHM 151

TOPICS:

QUIZ 6

1. Bond and Molecular Polarity

Key, Q 6 ClO 2-

7 + 12 + 1= 20/2 = 10 pr e's

O Cl O

-

O

Cl

O

Tetrahedral~109.5o

“bent”

Same as BrO2-

Lewis Structure

Shape

Key, Q 6

Trigonal bipyramidal90, 180o

“T shaped”

Same as IF3

Lewis Structure

Shape

BrC l37 + 21 = 28/2= 14 pr e's

Cl Br Cl

Cl

Br Cl

Cl

Cl

Oxidation Numbers (from formula)

Formal charge (from Lewis structure)

O 2[ 6- (1+6)] = -2

Cl 7- (2+4) = +1

Sum, all FC = -1

O Cl O

-

Cl O 2 -

-2

-4

?=+3

+3

(?) + (-4) = -1

(?) = 4 - 1=+3

Oxidation Numbers (from formula)

Formal charge (from Lewis structure)

Br Cl 3

-1

-3?=+3

?=+3 (?)+(-3) = 0

(?) = 0 +3= +3

Cl Br Cl

ClBr 7- (3 + 4) = 0

Cl 3[7- (1 + 6)] = 0

Sum, all FC = 0

UNIT FIVE

Bond and Molecular Polarity (Test 5)

Gas Laws (Test 5)

Intermolecular Attractions (ACS test only)

We have classified bonds “ionic” and “covalent”,depending on whether electron pairs are shared orelectrons are completely transferred from one atom to another.

In actuality, there is no sharp dividing line between thetwo types but rather a continuum:

Evenly shared electrons

Unevenly shared electrons

Transferredelectrons

To determine where a bond lies in this “continuum”, it is useful to consider the difference in electronegativity ( X) between the two atoms making up the bond:

When the difference ( X) is less than 0.5, sharing is fairlyeven and electrons are not much closer to one atom than the other. The bonds are considered “non-polar.”

When the difference is between 0.5 and about 1.7, theelectrons are closer to the more electronegative atom and partial charge buildup, polarization, develops.

When ( X) is greater than 1.7 or so, ionic bonding becomes the more likely bond type and valence electrons are transferred to the more electronegative atom.

Cl Cl H H H Cl Na+ Cl-

X 3.0 3 .0 2.1 2 .1 2.1 3 .0 1 .0 3 .0

X 0 0 .9 2.0

BONDTYPE:

(NONPOLAR)COVALENT

POLARCOVALENT IONIC

So, we need to consider a third more specializedtype of bond, “the polar covalent bond:”

This type of bond will be the important factor tobe considered when we look at molecular polarity, which arises from molecular shape and bondpolarity.

The polar molecular in turn will exhibit different solubilities and boiling points than non polar molecules.

Let us consider the bond between H and Cl in amolecule of hydrogen chloride (only hydrochloricacid when in water!):

E pair closer to Cl,more electronegative

Orbital between H and Cl

ClH + ClH ClH

The electron cloud from the pair of sharedelectrons is more dense closer to the chlorine,and much less dense closer to the hydrogen.

The bond has become “polarized”: it has developed a region (or “pole”) of partial positive charge buildup and a region (or “pole”) of partial negative buildup.

ClH

Major portion of electron density

H Cl~ + ~ -

H Cl~ + ~ -

or

Arrow to indicate polarbond, pointing to more(-) atom

“partially positive”

“partiallynegative”

H Cl

The molecule has only one bond, and it is polar.This makes the entire molecule a “dipole”, one whichhas a positive and negative pole and will align in anelectrical or magnetic field:

All diatomicmolecules withpolarized bondingbetween the two atoms are DIPOLES.

Other examples of diatomic dipoles:

C O

C O

X 2.5 3.5

X 1.0

CARBON MONOXIDE

I F

X 2.5 4.0

X 1.5

IODINE FLUORIDE

I F

MOLECULAR POLARITY, LARGER MOLECULES

All diatomic molecules with polar bond(s) are dipoles,but the situation is not so simple for larger molecules.

There are two factors to consider:

• Are the bonds polar?

• Are they arranged so that the center of positive charge and the center of negative charge do not “coincide”?

NH 3 NH H

H

X 3.0 2.1

X 0.9

BOND POLARITY MOLECULARPOLARITY

NH H

H

NH H

H

Center of +,-charges coincide,center of molecule

CCl Cl

Cl

Cl

CH H

H

H

X 3.0 2.5

X 0.5

X 2.1

X 0.4

CCl Cl

Cl

Cl

CH H

H

H

2.5

Polar nonpolar

NONPOLARMOLECULES

POLAR BONDS, NON-POLAR MOLECULE

Center of ChargeCenter of Charge

CH Cl

H

Cl

X 2.1

X 0.4

H

HCl

Cl

CH 2Cl 2

C H

2.5 X 3.0

X 0.5

C Cl

2.5

DIPOLE

CO O

O2-

CO32-

X 3.5 2.5

X 1.0

CO O

O2-

C

O

O

~-

~-2-

O~-

Centers Coincide,no dipole

CH 2O

CH H

O

X 2.5 3.5

X 1.0

C O

X 2.1

X 0.4

C H

2.5

O

HH

DIPOLE

C

H2O

H

OH H

H

X 3.5 2.1

X 1.4

O

DIPOLE

In conclusion, to be a dipole, a polar molecule (or polyatomic ion), the presence of polar bonds is required.

However, in addition, the polar bonds must be arranged so that they are not canceling.

Molecular shape must be such that the center of the negative charge buildup does not coincide with thecenter of positive charge buildup.

Group Work Determine the polarity of each of the below:

NO 2+

NO O+

NO 3-

NO O

O-

NO 2-

NO O

-•Draw to shape•Check en•Draw arrow, if dipole

NO 2-

NO O

-

X 3.5

X 0.5

3.0

NO O

-

NO O

dipole

NO 2+

NO O+

N

NO DIPOLE : CENTERS COINCIDE

2 regions : s hape LINEAR

ELECTRONEGATIVITY:

O 3.5 N 3.0

0.5 BONDS POLAR

NO 3-

NO O

O-

N NO DIPOLE : CENTERS COINCIDE

NO O

O-

Hybrid structure

Importance of Polarity

As it turns out, it is the difference in polarity whichdetermines, for the same sized species, whetherit is soluble in water and whether it is a gas atroom temperature or a volatile liquid whichevaporates quickly or a high boiling liquid whichdoes not evaporate at all.

The attractions between molecules which causes these differences all arise from increasing polarity or its complete lack...