1

A molecule of ammonia NH3 is made up of one nitrogen and

three hydrogen atoms:

Coordinate bond

The nitrogen atom forms three bonds and the hydrogen

atoms one bond each. In this case, one pair of electrons is

not involved in bond formation and this is called a lone pair

of electrons.

2

It is possible to have a shared electron pair in which

the pair of electrons comes just from one atom and not from

both. Such bond is called coordinate covalent bond.

Even though the ammonia molecule has a stable

configuration, it can react with hydrogen H+ by donating the

lone pair of electrons, forming the ammonium ion NH4+:

3

In the chlorine molecule Cl – Cl the pair of electrons

of the covalent bond is shared equally between both

chlorine atom. Because there is not a charge separation

between the chlorine atoms, Cl2 molecule is nonpolar.

Partial ionic character of covalent bonds

4

On the contrary, in HCl molecule, there is a shift of

electrons toward the chlorine atom which is more

electronegative than hydrogen. Such molecule, in which a

charge separation exists is called a polar molecule or

dipole molecule.

The polar molecule of hydrochloric acid

+ +e-

H Cld

+ -

5

The separation between the positive and negative

charges is given by the dipole moment μ. The dipole

moment is the product between the magnitude of the charges

(δ) and the distance separating them (d):

μ = δ · d

The symbol δ suggests small magnitude of charge,

less than the charge of an electron (1.602 · 10-19 C).

The unit for the dipole moment is Debye (D):

1D = 3.34 · 10-30 C m

6

C...

d20

12

30

1053210136

10343031

The charge δ for HCl molecule represents about 16%

of the electron charge (1.602 10-19 C). We can say that the

covalent H – Cl bond has about 16% ionic character.

Dipole moment values for some molecules:

Carbon dioxide CO2 μ = 0 D

Carbon monoxide CO μ = 0.112 D

Water H2O μ = 1.85 D

Hydrochloric acid HCl μ = 1.03 D

dist. between H and Cl atoms is d = 136 pm (136 10-12 m)

We can calculate the charge δ for HCl molecule:

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c) Metallic bond

The metallic bond represents the electromagnetic

attraction forces between delocalized electrons and the metal

nuclei. The metallic bond is a strong chemical bond, as

indicated by the high melting and boiling points of metals.

A metal can be regarded as a lattice of positive metal

“ions” in a “sea” of delocalised electrons.

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Metal atoms contain few electrons in their outer shells.

Metals cannot form ionic or covalent bonds.

Sodium has the electronic structure 1s22s22p63s1.

When sodium atoms come together, the electron from the 3s

atomic orbital of one sodium atom shares space with the

corresponding electron of a neighbouring atom to form a

molecular orbital. All the 3s orbitals of all the atoms overlap

to give a vast number of molecular orbitals which extend over

the whole piece of metal. There is a huge numbers of

molecular orbitals because any orbital can only hold two

electrons.

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The electrons can move freely within these molecular orbitals and so each

electron becomes detached from its parent atom. The electrons are called

delocalized electrons. The “free“ electrons of the metal are responsible

for the characteristic metallic properties: ability to conduct electricity and

heat, malleability (ability to be flattened into sheets), ductility (ability to be

drawn into wires) and lustrous appearance.

Crystal structure of sodium

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Weak chemical bonds – Intermolecular bonds

1. Van der Waals forces

Intermolecular forces are attractions between one

molecule and neighboring molecules. All molecules are under

the influence of intermolecular attractions, although in some

cases those attractions are very weak. These intermolecular

interactions are known as van der Waals forces.

Even in a gas like hydrogen (H2), if the molecules are

slow down by cooling the gas, the attractions become large

enough so the molecules will stick together and form a liquid

and then a solid.

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The attractions between the H2 molecules are so

weak that the molecules have to be cooled to 21 K (-252C)

before the attractions are enough to form liquid hydrogen.

Helium’ s intermolecular attractions are even weaker

– the molecules won’t stick together to form a liquid until the

temperature drops to 4 K ( -269 C).

Attractions are electrical in nature. In a symmetrical

molecule like hydrogen, however, it doesn’t seems to be any

electrical distortion to produce positive or negative parts.

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H2 symmetrical molecule

++-

H H

-

But that’ s only true in average. In the next figure the

symmetrical molecule of H2 is represented.

On average there is no electrical distortion.

13

But the electrons are mobile and at any one instant

they might find themselves towards one end of the

molecule. This end of the molecule becomes slightly

negative (charge -). The other end will be temporarily short

of electrons and so becomes slightly positive (+).

An instant later the electrons may move to the other end,

reversing the polarity of the temporary dipole of molecule.

Temporary dipole of H2 +

+ -

+ --

H H

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This phenomena even happens in monoatomic

molecules of rare gases, like helium, which consists of a

simple atom. If both the helium electrons happen to be on

one side of the atom at the same time, the nucleus is no

longer properly covered by electrons for that instant.

Temporary dipole of He

+-

-

+-

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How temporary dipoles give intermolecular bonds?

A molecule which has a temporary polarity approaches

another molecule which happens to be non-polar at that

moment. The electrons form the non-polar molecule will be

attracted by the slightly positive end of the polar molecule.

This is how an induced dipole is forming.

+-

+- +-

original temporary dipole

original temporary dipole

non-polar molecule

induced dipole

Dipole-dipole attraction (van der Waals forces)

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An instant later the electrons in the left hand molecule

can move up to the other end. So, they will repel the

electrons in the right hand molecule. The polarity of both

molecules reverses, but there is still attraction between - end

and + end. As long as the molecules stay close to each other,

the polarities will continue to fluctuate in synchronization so that

the attraction is always maintained.

This phenomena can occur over huge numbers of

molecules. The following diagram shows how a whole lattice of

molecules could be held together in a solid.

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Molecular distribution in a solid

The interactions between temporary dipoles and induced

dipoles are known as van der Waals dispersion forces .

+- +- +-

+- +- +-

+- +- +-

+- +- +-

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Hydrogen bond If we plot the boiling points of the hydrides of the

elements of groups 15, 16 end 17 we find that the boiling point

of the first elements in each group is abnormally high.

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In case of ammonia NH3, water H2O and hydrofluoric acid

HF, there must be some additional intermolecular forces of

attraction, requiring significantly more heat energy to break

them. These relatively powerful intermolecular forces are

called hydrogen bonds.

Hydrogen bonds are stronger than van der Waal dispersion

forces, but weaker than covalent or ionic bonds.

Hydrogen bonds can form if:

hydrogen is attached directly to one electronegative

element (F, O, N)

each of the elements to which the hydrogen atom is

attached have one “active“ lone pair of electrons.

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Let’s consider two water molecules coming close together:

The slightly + charge of hydrogen is strongly attracted to the

lone pair of electrons; as a result a coordinate bond is formed.

This is a hydrogen bond.

In liquid water, hydrogen bonds are constantly broken and

reformed.

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In solid water each

water molecule can

form hydrogen

bonds with other 4

surrounding water

molecules, creating

a 3-D structure of

ice. As a result, the

boiling point of H2O

is higher than that

of NH3 or HF.