CHAPTER 4 - CHEMICAL BONDING

2. The polarity of molecules

For molecules that have similar molecular
mass, the dipole-dipole interaction will be
more dominant. Polar molecules have higher
boiling point due to the existence of stronger
dipole-dipole attraction.

Example: Molecular Polarity Boiling Point
Mass (oC)
F2 Non polar
HCl 38.0 Polar 85.01

36.5 188.11

251

B. Hydrogen bond

 Dipole-dipole interaction that acts between
molecules that have a hydrogen atom that is
covalently bonded to a highly electronegative
atom; F, O ,N in one molecule and F,O or N of
another molecule.

Example: + - + -

HF HF

Hydrogen bond 252

Other examples: H2O

NH3 liquid Hydrogen bond
OO
..
O
N
O Hydrogen bond
Hydrogen bond

..

N

Covalent bond

253

Effect of Hydrogen Bonding on Physical
Properties

(a) Boiling point

• Boiling point of compounds having hydrogen
bonds are relatively higher than compounds
having dipole-dipole interaction or London forces.

• It is due to Hydrogen bond is the strongest
attraction force compared to the dipole-dipole
interaction or the London forces.

254

Example:

Explain the trend of boiling point given by the
graph below:

T/oC
HF

HI

HBr
HCl

Molecular mass

255

Answer:

 HF form hydrogen bonds between molecules
while HCl, HBr and HI have van der Waals forces
acting between molecules.
Hydrogen bond is stronger that the van der
Waals forces. More energy is required to break
the Hydrogen bond. HF has the highest boiling
point.

 Boiling point increases from HCl to HI. The
strength of van der Waals forces increases with
molecular size. Since molecular size increases
from HCl to HI, thus the boiling point will also
increase in the same pattern.

256

 The boiling points of compounds with hydrogen
bonds are affected by:
a) the number of hydrogen bonds per
molecule
b) the strength of hydrogen bond which
directly depends on the polarity of the
hydrogen bond

257

Example:
The order of the increase in boiling point is:
H2O > HF > NH3 > CH4
Explain the trend of boiling points.

258

Answer:

 by looking at the polarity of the bond, we have

(order of polarity: HF > H2O > NH3)
but H2O has the highest boiling point.
For H2O, the number of hydrogen bonds per molecule
affects the boiling point.

 Each water molecule form more hydrogen bonds
with other water molecules. More energy is
required to break the Hydrogen bonds.

 HF has higher boiling point than NH3 because F is
more electronegative than Nitrogen. Therefore the
hydrogen bond in HF is stronger.

 CH4 is the lowest - it is a non polar compound and
has weak van der Waals forces acting between

molecules. 259

Hydrogen bonds between HF molecules and

H2O molecules H2O Hydrogen bond

+ - + - O O

HF HF

O

Hydrogen bond

O Hydrogen bond

260

(b) Solubility

 Molecules which form hydrogen bond with
water molecules are soluble in water.
Examples: NH3, HF, CH3OH

O

Hydrogen bond

.. ..

N N

 NH3 dissolves in water because it form
hydrogen bond with water.

261

• Organic compound that soluble in water include:
i. Amine (eg: methylamine, CH3NH2)
ii. Alcohol (eg: ethanol, CH3CH2OH)
iii. Carboxylic acid (eg: ethanoic acid, CH3COOH)

• As the relative mass increase, the non-polar
hydrocarbon portion (hydrophobic) become
larger while polar group (hydrophilic) represents
an increasingly smaller portion of the molecule.

•Therefore, the solubility decreases.

262

R OH .. H

Hydrophobic O.. H
group
Hydrogen
(water-hating) bonding

263

Name Formula Solubility
(mol/100g of
Methanol CH3OH
Ethanol C2H5OH water)
Propanol C3H7OH miscible
Butanol C4H9OH miscible
Pentanol C5H11OH miscible
Hexanol C6H13OH
Heptanol C7H15OH 0.11
0.030
0.0058
0.0008

264

Example

Solubility of butanol and hexanol in water:

butanol: CH3CH2CH2CH2 OH .. H

hexanol: CH3CH2CH2CH2CH2CH2 OH .. O.. H

H

O.. H

Larger hydrophobic group,
less soluble in water

265

(c) Density

 The density of water is relatively high
compared to other molecules with similar
molar mass.

Reason:
Hydrogen bonds are stronger than the
dipole-dipole or the London forces. Thus the
water molecules are drawn closer to one
another and occupy a smaller volume.

266

Ice (solid H2O) has lower density compared to its liquid.

• The hydrogen bonds in ice arrange the H2O molecules
in open hexagonal crystal (open structure)

Hydrogen bond takes
one of the tetrahedral
orientation and occupy
some space

267

• This arrangement leaves a relatively large
amount of empty space between them.

• As a result, ice has larger volume
compared to the liquid water.

• Thus, the density of ice is less than liquid
water.

268

• When ice melts, H2O molecules have higher
kinetic energy and can overcome the hydrogen
bond.

• The hydrogen bonds are broken and the open
structure collapses. V-shaped water molecules
slide between each other.

• This causes the volume occupied by water to
decrease causing a corresponding increase in
the density

269

270

271









4.5 Metallic Bond

Learning Outcomes:

a) Explain the formation of metallic bond by using
electron sea model.

b) Relate metallic bond to the properties of metal:
malleability, ductility, electrical conductivity and
thermal conductivity.

c) Explain the factors that affect the strength of
metallic bond.

d) Relate boiling/melting point to the molecular
structure, types of bonding and intermolecular
forces for elements of: period 3, group 1 and group
17.

277

Metallic bond

 An electrostatic force between positive charge
metallic ions and the sea of delocalised valence
electrons.

 The metallic bond can be imagined as an array
of positively charged ions immersed in a sea of
delocalized valence electron.

 In the solid state, metal atoms are closely
packed in a regular arrangement.

 The solid lattice held together by the strong
attractive forces between the delocalised
valence electrons and positively charge ions.

278

Example : Na

279

Physical properties of metals

Metal have the following physical properties:

(a) Malleability
(b) Ductility
(c) Electrical conductivity
(d) Thermal conductivity

280

Malleability and Ductility

Malleable – can be pressed into different shapes.
Ductile – can be pulled into wire

• Metal atoms are arranged closely packed
• When sufficient force is applied to the metal, one

layer of atoms can slide over another without
disrupting the metallic bonding.
• As a result, metals are malleable and can be
drawn into wires(ductile)

281

Electrical and Thermal Conductivity
• Metals have high electrical and thermal

conductivity because the sea of delocalised
electrons are free to move from one end to
another end when there are differences in
electrical potential or heat.

282

Factors that affect
the strength of the metallic bond

 The strength of metallic bonding depends on:
i. Size of positive metal ions
ii. Number of valence electrons

 The strength of metallic bonds is proportional to
the number of valence electrons and inversely
proportional to the size of the atom.

 The smaller size of positive metal ions and the

more number of valence electrons will exert

stronger attraction of metallic bond. 283

Effect of the strength of metallic bond on
boiling point

 The boiling point of a metal depends on the
strength of metallic bond.

 The stronger the metallic bond, the higher the
boiling point because high energy is required to
overcome these strong electrostatic forces
between the positive ions and the delocalised
valence electron sea in the metallic bond

284

Example:

Explain the difference in the boiling point of the two
metals given:

Boiling Point (oC)
Mg 1120
Al 2450

285

Boiling point of Al is higher than that of Mg:

 The strength of metallic bond in Aluminium is
stronger than Magnesium because Al has three
valence electrons and smaller cationic radius
while Mg has only two valence electrons and
larger cationic radius.

 The strength of metallic bond is directly
proportional to the boiling point.

 The stronger metallic bond, the higher the
boiling point.

286