Bonds in solids

2. • In nature one comes across several types of solids.
• Many solids are aggregates of atoms.
• The arrangement of atoms in any solid material is
determined by the character, strength and
directionality of the chemical binding forces, cohesive
forces or chemical bonds. We call these binding
forces as atomic interaction forces.

3. • The atoms, molecules or ions in a solid state
are more closely packed than in the gaseous
and liquid states and are held together by
strong mutual forces of attraction and
repulsion.
• One can describe the atomic arrangement in
elements and compounds on the basis of this.
• The type of bond that appears between
atoms in crystal is determined by the
electronic structure of interacting atoms.

4. TYPES OF BOND

7. Ionic Bond
• Ionic bond is formed when the atoms with small number
of electrons on the outer shells, have a tendency to give
up electrons to the atoms with an almost filled outer
shell.
• Ionic bond is the simplest type of interatomic bond.
• By loosing electrons, metallic atoms like alkali-metals,
alkaline-earth metals etc.,
• Ionic bonding occurs between electropositive elements
(metals, i.e., those elements on the left side of the
periodic table) and the electronegative elements (non-
metals; i.e. on the right hand of the periodic table).

8. Ionic Bonds

9. Ionic Bond

10. IONIC BOND

11. CHA.IONIC SOLIDS
• Ionic solids are generally rigid and crystalline in nature.
• In ionic bonding, a metallic element loses from the outer
electron shell of its atom a number of electrons equal to
its valency (numerical). Obviously, an electrostatic
attraction between positive and negative ions occurs.
• In ionic crystal each positive ion attracts all neighbouring
negative ions and vice versa, the bond itself is non-
directional.
• Ionic solids are generally non-conductors of electricity.
• Ionic solids are highly soluble in water but insoluble in
organic solvents.

12. COVALENT BOND
• In covalent bonding, stable electron setups are expected
by the sharing of electrons between neighbouring
atoms.
• Two neighbouring atoms each having incomplete
outermost shells.
• Unlike ionic bonding, the atoms participating in the
covalent bond have such electronic configurations that
they cannot complete their octets by the actual transfer
of electrons from one atom to the other.
• Obviously, there is no charge associated with any atom of
the crystal.
• The majority of solids incorporating covalent bonds are
also bound by either ionic or Vander Waals bonds.

13. covalent bond
• A covalent bond is formed between similar or dissimilar
atoms each having a deficiency of an equal number of
electrons.
• When two atoms, each having a deficiency of one
electron, come so close that their electronic shells start
overlapping, the original atomic charge distributions of
atoms are distorted and each atom transfers its unpaired
electron to the common space between the atoms.
• Obviously, the common space contains a pair of electrons
which belongs equally to both the atoms and serves to
complete the outermost shell of each atom. This is called
sharing of electrons.

14. covalent bond
• The sharing is effective if the shared electrons have
opposite spins.
• In such a case the atoms attract each other and a
covalent bond is formed.
• As the participating atoms in the bond have the same
valence state, this bond is also called the ‘valence
bond’.

15. Covalent Bonds

16. CHA. Covalent COMPOUNDS
• Covalent compounds are mostly gases and liquids.
• They are usually electric insulators.
• They are directional in nature.
• They are insoluble in polar solvents like water but are soluble
in non-polar solvents, e.g,, benzene, chloroform, alcohol,
paraffins etc.
• Covalent compounds are homopolar, i.e. the valence electrons
are bound to individual or pairs of atoms and electrons cannot
move freely through the material as in the case of metallic
bonds.
• Covalent compounds are soft, rubbery elastomers, and form a
variety of structural materials usually termed as plastics.
• The melting and boiling points of these compounds are low.

17. Metallic bond
• Metallic bond is formed by the partial sharing of
valence electrons by the neighbouring atoms.
• It is somewhat intermediate between covalent and
ionic bonding.
• They are formed by the elements of all subgroups A
and I-III, subgroups B.
• Metallic materials have one, two, or at most, three
valence electrons.

18. Metallic bond
• With this model, these valence electrons are not
bound to a specific molecule in particular and are
pretty much allowed to float all through the whole
metal.
• They might be considered having a place with the
metal overall, or framing an “ocean of electrons” or
an “electron cloud.”
• Metallic bonds are electropositive.
• When interacting with elements of other groups,
atoms in a metallic crystal can easily give off their
valence electrons and change into positive ions.

19. Metallic bond
• When interacting with one another, the valence energy
zones of atoms overlap and form a common zone with
unoccupied sublevels.
• The valence electrons thus acquire the possibility to
move freely with the zone.
• Obviously, valence electrons are shared in the volume of
a whole crystal. Thus the valence electrons in a metal
cannot be considered lost or acquired by atoms.
• They are shared by atoms in the volume of a crystal,
unlike covalent crystals where sharing of electrons is
limited to a single pair of atoms.

20. METALLIC BOND

21. CHA. Metallic compounds
• Metallic compounds are crystalline in nature due to
the symmetrical arrangements of the positive ions in
a space lattice .
• Metallic bonding is weaker than ionic and covalent
bonding, but stronger than Vanderwaals' bonding.
• Metallic bonds being weak, metals have a melting
point moderate to high, i.e., the melting points of
metallic crystals are lower than those of
electrovalent crystals.
• Metals are opaque to light, because the free
electrons in a metal absorb light energy.

22. • Metals possess high thermal and electrical conductivities. This
is because of the free electron movement in the lattice.
Metals exhibit ductility.
• They are brittle below a certain temperature. Metals exhibit
brittle to ductile fracture transition at a certain temperature.
Many metals like tungsten are tough. But some metal like
pure silver, aluminium, gold are soft and malleable. Their
melting points vary from 64 0 C (Potassium) to 2000 0 C
[Titanium (1812 0 C), Molydenum, Tungsten].
• Metallic alloys are extremely useful for all engineering
applications. Their high conducting values (gold, platinum,
silver, copper etc.,) make them useful as electrodes in
scientific and technical devices. Metals are opaque to light
and they reflect light energy very efficiently.
CHA. Metallic compounds