Lec.

Rana Hassan Tariq

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Chapter #4
LIQUIDS AND SOLIDS

Differentiate between intramolecular forces and intermolecular forces

Intramolecular forces Intermolecular forces
 Those forces which are present within the  Those forces, which are present between two
molecules between atoms are called molecules, are called intermolecular forces.
intramolecular forces.
 Valence electrons are involved in their  No involvement of valence electrons.
formation.
 Nature of substance is determined by  State of a substance (solid, liquid and gas) is
intramolecular forces. determined by intermolecular forces.
 Chemical properties of a substance are related to  The physical properties of a substance are related
its intramolecular forces. to its intermolecular forces.
 These are comparatively strong forces.  These are relatively weak forces.
 Examples:  Examples:
i. Covalent bond i. Dipole-Dipole forces
ii. Ionic bond ii. Dipole induced dipole forces
iii. Coordinate covalent bond iii. London dispersion forces
iv. Hydrogen bonding

DIPOLE-DIPOLE FORCES

Definition:
The attractive forces present between positive end of polar molecule and the negative end of another
polar molecule are called dipole dipole forces.

Examples:
HCl:
In case of HCI molecule, both atoms differ in electronegativity. Chlorine being more electronegative
develops the partial negative charge and hydrogen develops partial positive charge. So, whenever, molecules
are close to each other, they tend to line up.

Chloroform:
In CHCl3 molecule, there is force of attraction between hydrogen of one molecule and chlorine of
other molecule.
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Factors affecting dipole-dipole forces:

(i) Electronegativity difference:
The strength of these forces depends upon electronegativity difference between the bonded
atoms and distance between molecules.
In case of HCI, both atoms differ in electronegativity difference. Chlorine being more
electronegative develops partial positive charge and hydrogen develops partial negative charge.
Greater is the electronegativity difference greater will be the dipole-dipole forces.

(ii) Intermolecular Distance:

The distances between molecules in gas phase are greater so these forces are very, weak in
this phase. In liquids, these forces are reasonably strong.

Important points:
1. Dipole-dipole forces are 1% as effective as a covalent bond.
2. The values of thermodynamic parameters such as melting points, boiling points, heat of
vapourization and heat of fusion depends upon the strength of dipole-dipole forces.

DIPOLE INDUCED DIPOLE FORCES OR DEBYE FORCES

Sometimes, we have a substance containing polar and non-polar molecules. The positive end of polar
molecule attracts the mobile electrons of the nearby non-polar molecule. In this way, polarity is in non-polar
molecule and both the molecules become polar. These forces are called induced dipole or Debye forces.

Dipole induced dipole interaction

LONDON DESPERSION FORCES OR INSTANTANEOUS DIPOLE INDUCED

DIPOLE FORCES OR SHORT RANGE FORCES

Definition:
The momentary force of attraction between instantaneous dipole and induced dipole is called dipole-
induced dipole forces.

Explanation:
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In a gas, the electrons of one atom influence the morning electrons of the other tend to say as far
apart as possible when electrons of one atom come closer to electrons of other atom, they are pushed away
from each other. In this way, a temporary dipole is created in atom. The result is that at any moment, the
electron density of atom is no more symmetrical. It has a negative charge on one side than on the other.

Instantaneous dipole induced dipole attractions between helium atoms

At that particular instant, three helium atoms becomes a dipole. This is called instantaneous dipole. The
instantaneous dipole then disturbs the other electronic cloud of the other nearby atom. So a dipole is induced
in the second atom. This is called induced dipole. It is a very short lived because electrons keep on moving.
This movement of electron causes the dipole to vanish as quickly as they are formed. The dipoles appear
again in different orientation and again weak attractions are developed.

Define polarizability:
“The quantitative measurement of the extent to which the electronic cloud can be polarized or
distorted is called polarizability.”
As the size of the molecules increases the distortion also increases. This increased distortion of electronic
cloud creates stronger London dispersion forces and hence the values of thermodynamic parameter increase.

Occurrence:
It is present in all types of molecules whether polar or non polar but they are very significant in non
polar molecules e.g., H2, Cl2, O2, N2 noble gases.

Discovery:
This force of attraction was identified by Fritz London in 1930. Fritz London was a German
physicist.

Factors affecting on London forces:

(i) Size of electronic cloud:
London forces are weaker than dipole-dipole interaction. The strength of these forces
depends upon size of electronic cloud of the atom and molecules. When the size of atom or
molecule is larger than dispersion becomes easy and these forces become more prominent.

Example 1:
The elements of the zero groups in the periodic table are all mono atomic gases. Their boiling
points increase down the group form Helium (-268.6oC) to Radon (-61.8 oC). The atomic number
increases down the group and the outermost electrons move away from nuclei. The dispersion of
electronic clouds becomes more and more easy. So the polarizability of these atoms goes on
increasing as a result, London Dispersion Forces become stronger.

Example 2:
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Example 2:
All the halogens are non-polar diatomic molecules, but there is a big difference in their
physical state at room temperature Fluorine is a gas and boils at -188.1°C while iodine is a solid
at room temperature which boils at +184.4°C. The polarizability of iodine molecule is much
greater than that of fluorine.

(ii) No. of atoms in a molecule (atomicity):

Another important factor that affects the strength of London forces is the no. of atoms in a
non-polar molecule Greater the no. of atoms in a molecule, greater is the polarizability and hence
stronger will be London dispersion forces. (LDF)

Example:
Compare the length of chain for ethane (C6H6) and hexane (C6H14). They have the boiling
points -88.6°C and 68.7°C respectively. This means that a molecule with a longer chain length
experiences stronger attraction forces. The reason is that longer molecules have more places
along its length, which they can be attracted to other molecule.

HYDROGEN BONDING

Definition:
The electrostatic force of attraction between hydrogen atom
(bonded to a small highly electronegative atom) and the
electronegative atom (F, O, N) of another molecule is called hydrogen
bonding.

Symbol:
It is represented by dotted line (....).

Examples:
(i) Water (H2O):
In case of H20, oxygen is a more electronegative element as compared to hydrogen, so water
is a polar molecule. Hence, there will be dipole-dipole interactions between partial positively
charged hydrogen and partially negatively charged oxygen atoms. Actually, hydrogen bonding is
something more than dipole-dipole interaction. Firstly, oxygen atom has two lone pairs. Secondly
hydrogen has sufficient partial positive charge. Both the hydrogen atoms of water molecule
create strong electrical field due to their small sizes. The oxygen atom of the other molecule links
to form a coordinate covalent bond with hydrogen using one of its lone pairs of electrons. Thus,
loose bond formed is definitely stronger than simple dipole- dipole interaction. Because of the
small size of hydrogen atom, it can take part in this type of bonding. This bonding acts as a
bridge between two electronegative oxygen atoms.
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(ii) Hydrofluoric Acid (Ill):

The molecules of HF join with each other in a zigzag manner.

(iii) Chloroform and Acetone:

It is not advisable to limit the hydrogen bonding to fluorine, oxygen and nitrogen only. In
case of chloroform three chlorine atoms are responsible for H-bonding with other molecules,
these atoms deprive the carbon atom of its electrons and the partial positively charged hydrogen
can form a strong hydrogen bond with oxygen atom of acetone.

Why HF is a weak and as compared to HCl, HBr and HI?

The exceptional, low acidic strength of HF molecule as compared to HCI, HBr and HI is due
to the strong hydrogen bonding in HF molecules as the partial positive hydrogen is entrapped
between two highly electronegative atoms so this acid do not ionizes completely in water and
hence HF becomes a weak.

PROPERTIES AND APPLICATIONS OF COMPOUNDS CONTAINING

HYDROGER BONDING

1. Thermodynamic properties of covalent hydride:

“The binary compounds of hydrogen are called hydrides”.
Hydrogen bonding exists in the compounds having partial positively charged hydrogen and highly
electro-negative atoms bearing partial negative charge. Obviously such intermolecular forces will
influence the physical properties like melting point and boiling point. Let us compare the physical
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properties of hydrides of group IV-A, V-A, VI-A, and VII-A. The graphs are plotted between the
period number of periodic table on x-axis and boiling points in Kelvin on y-axis. Boiling points of
hydrides of IV-A group here low boiling points as compared to those of group V-A, VI-A and VII-A.
The reason is that these elements are least electronegative. CH4 has lowest boiling point because it is
a very small molecule and its polarizability is least. When we consider hydrides of group V-A, V1-A
and VII-A then NH3, H2O and HF show maximum boiling points in respective series. The reason is
the entrained electronegative character of N, O and F. That is why water is liquid at room
temperature but H2Se are gases. It is interesting to know that boiling point of water (100°C) seems to
be more effected by hydrogen bonding than that of HF (19.9°C). Fluorine is more electronegative
than oxygen. So, we should expect hydrogen bonding and as a result the boiling point of HF should
be higher than that of H2O. However, it is lower the reason is that fluorine atom can, make only one
hydrogen bond with electropositive hydrogen in neighboring molecule water can form two hydrogen
bonds per molecule as it has two hydrogen atoms and two lone pair on oxygen atoms. Ammonia can
form only one hydrogen bonding per molecule as it has only one lone pair. The boiling point of HBr
is slightly higher than that of HCl. It means that chlorine is electronegative enough to form a
hydrogen bond. Sometimes it is thought that HCl has a strong dipole-dipole force but in reality, it is a
borderline case. The hydrides of fourth group GeH2, ASH3, H2Se, HBr show greater boiling points
than those of third period due to greater size and enhanced polarizibilities.

2. Solubility of hydrogen bonded molecules:

Alcohol:
A group of compounds containing OH (Hydroxyl) functional group OH are called alcohols.

Water and ethyl alcohol:

Water is the best example of hydrogen bonded system. Ethyl alcohol (C2H2OH) also has the
tendency to form hydrogen bonds. So ethyl alcohol can dissolve in water because both can form
hydrogen bonds with each other.

Carboxylic acid:
“The compound containing –COOH (carboxylic group) are called carboxylic acids.”
Carboxylic acids are soluble in water which is small sized. They form hydrogen bonding with water
molecules and become dissolving. Carboxylic acids like HCOOH (formic acid) and (CH3COOH)
acetic acid are soluble in water.

Hydrocarbons:
“The compounds containing carbon and hydrogen only are called hydrocarbons.”
Hydrocarbons are not soluble in water at all because they are non-polar compounds and there are no
chances of hydrogen bonding between water and hydrocarbons molecules.
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3. Structure of ice:
The molecules of water have tetrahedral structure. Two
lone pairs of electrons on oxygen atom occupy two others corners
of tetrahedron.

Water molecule in liquid state:

In the liquid state, water molecules are extensively
associated, with each other and these associations break and are
reformed because molecules of water are mobile.

Water molecule in ice:

When temperature of water is decreased and ice is formed
then the molecules of water are mobile become more regular and this regularity extends throughout
the structure. The corners of tetrahedrons having lone pairs are linked with partial positive H atoms
of their.

Density of ice:
When ice is formed, there remain empty spaces-in the
structure of ice. These spaces are due to hydrogen bonding. Due to
this reason, when water freezes it occupies 9% more space and its
density decreases. The result is that ice floats on the surface of
water.

Diamond like structure:

The structure of ice is just like that of a diamond because
each atom of carbon in diamond is at centre of tetrahedron just like
the oxygen of water molecule in ice.

Significance of low density of ice:

The lower density of ice than the liquid water at 0°C causes water in ponds and lakes to
freeze from surface to the downward direction. Water attains the temperature of 4°C by the fall of
temperature in the surrounding. As the outer atmosphere becomes further cold, the water at the
surface becomes less dense. This less dense water below 4°C stays on the top of slightly warm water
underneath. A stage reaches when it freezes. This layer of ice insolates the water underneath for
further heat loss. Fish and plants survive under this blanket of ice for months. “The patterns of life
for plants and animals would have been totally different in the absence of hydrogen bonding in
water.”

4. Cleansing action of soaps and detergents:

The cleansing action of soaps and detergents is due to hydrogen, bonding. Reason is that
polar parts of their molecules are water soluble due to hydrogen bonding and non polar parts remain
outside the water because they are alkyl or benzyl portions and are insoluble in water.

5. Hydrogen bonding in biological compounds and food materials:

Hydrogen bonding exists in the molecules of living system. Proteins are the important part of
living organisms. Fibers like those found in the hair, silk and muscles’ consist of long chains of
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amino acids, these long chains are coiled about one another in a spiral.
This spiral is called a helix. Such a helix be either right handed or left
handed.

H-Bonding in single helix:

In the case of right, handed helix the groups like >NH and >CO are vertically adjacent to one
another and they are linked together by hydrogen bonds. These H bonds like one spiral to other. X-
ray analysis has shown that there are 27 amino acid units for each turn of helix.

Hydrogen bonding in double helix DNA:

There are two spiral chains in DNA (Deoxyribonucleic acid) which are coiled about each
other on a common axis. In this way they give a double helix. This is 18-20A in diameter. They are
iii led together by H-bonding between their submits.

H-bonding in carbohydrates:
The food materials like carbohydrates include glucose fructose and sucrose. They all have –
OH groups due to which hydrogen bonding is present.

6. Hydrogen bonding in points, dyes and textile materials:

One of the most important properties of paints and dyes is their adhesive action. This
property is developed due to hydrogen bonding. Similar type of hydrogen bonding makes glue and
honey as sticky substances.

Hydrogen bonding in clothing material:

Hydrogen bonding is very important in thread making materials like cotton, silk and synthetic
fibers for clothing. This hydrogen bonding is responsible for heir rigidity and the tensile strength.

EVAPORATION

Definition:
The spontaneous change of a liquid into vapours is called evaporation and it continues at all
temperature.
Characteristics:
1. Surface phenomenon:
The molecules of liquid are not motionless. The energy of molecules is not equally
distributed. The molecules which have low K. Energy move slowly while others with high of high
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speed molecules reaches the surface it may escape attractions of neighboring molecules and leaves
bulk of liquid.

2. Endothermic process:
As energy is required for evaporation, which is absorbed from surrounding and it is called
endothermic process.

3. Continuous at all temperature:

The process of evaporation continuous at all temperatures and liquid continuously changes in
vapours spontaneously.

4. Cooling process:
Evaporation causes cooling. The reason is that high energy molecules leave the liquid and
low energy molecules left behind, the temperature of liquid falls and heat moves fim sunounding to
liquid and then temperature of surrounding also falls.

Factors affecting rate of evaporation:

1. Temperature:
Higher temperature higher will be rate of evaporation because it increases K.E of molecules.
R∝T
2. Intermolecular forces:
Stronger the intermolecular forces, slower will rate of evaporation.

3. Surface area:
Greater surface area, greater will be rate of evaporation.
R Surface area
4. Speed of wind:
Greater speed of wind greater will be evaporation rate.
R Speed of wind
For example:
Gasoline has weak London dispersion forces than water. So rate of evaporation of gasoline is
faster than that of water.

VAPOUR PRESSURE

Definition:
Vapour pressure of a liquid is pressure exerted by the vapours of liquid in equilibrium with the liquid
at a given temperature.

Explanation:
When the molecules of liquid leave the open surface, they are mixed up with air above liquid. If the
vessel is open these molecules go on leaving the surface. But if we close the system the molecules of liquid
start gathering above the surface. These molecules not only collide with the walls of container but also with
surface of liquid as well. There are chances that these molecules are recaptured by surface of liquid. This
process is called condensation. The two process i.e., evaporation and condensation continue till a stage
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reaches when the rate of evaporation becomes equal to rate of condensation. This is called state of dynamic
equilibrium.

The no. of molecules leaving the surface is just equal to the no. of molecules coming bank to it at a constant
temperature. The molecules which are in a liquid state at any moment may be in vapour state in the next
moment.

Factors on which vapour pressure does not depend:

Vapour pressure does not depend upon:
1. Amount of liquid.
2. Volume of container.
3. Surface area of liquid.
The larger surface area also presents a larger target for returning molecules. So rate of condensation also
increases.

Factors on which vapour pressure depends:

1. Temperature
2. Inter molecular forces

Temperature:
The values of vapour pressure of vapours liquids depend fairly on the nature of liquids i.e., on the
sizes of molecules and intermolecular forces but most important parameter which controls the vapour
pressure of a liquid is its temperature. At an elevated temperature the kinetic energy of molecule is enhanced
and capability to leave the surface increases. It causes the increase of vapour pressure. Increase of vapour
pressure goes on increasing for the same temperature of from 0oC to 100oC for water. There is increase of
vapour pressure from 4.579 torr tto 9.209
.209 torr for change of temperature from O°C to 10°C. But increase is
from 527.8 torr to 760 torr when temperature changes from 90°C to 100°C.

Strength of intermolecular forces:

The difference in the strength of intermolecular forces is different liquids is directly related to their
vapour pressure at a particular temperature. The stronger the intermolecular forces the lower the vapour
pressure. At 20°C isopentane has highest vapour pressure and glycerol has lowest.
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Measurement of vapour pressure:

There are many methods for measurement of vapour pressure of a liquid. One of the most important
method is given below:

Manometric method:
It is comparatively an accurate method. The liquid whose vapour pressure is to be determined is
taken in a flask placed in a thermostat. One end of tube is connected to a manometer and other is connected
to a vacuum pump. The liquid is frozen with the help of freezing mixture and the space above the liquid is
evacuated. In this way, the air is removed from the surface of liquid along with the vapours of that liquid.
The frozen liquid is then melted to release an entrapped air liquid is again frozen and released air is
evacuated. This process is repeated many times till almost all air is removed. Now the liquid is warmed in a
thermostat to that temperature at which its vapour pressure of liquid in the flask is to be determined.
Difference in the height of columns of Hg in two limbs of manometer determines vapour pressure of liquid.
The column of mercury in the manometer being facing the vapour of the liquid is depressed. The other
column which faces atmospheric pressure rises. Actually, the pressure on the surface of liquid in flask is
equal to sum of the atmosphere pressure and vapour pressure of liquid. For this reason the column of
manometer facing the liquid is depressed then facing the atmosphere and it is given by the following
equation.
P = Pa + ∆H
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BOILING POINT

Definition:
The temperature at which the vapour pressure of a liquid becomes equal to atmospheric pressure
external pressure of any subjected pressure is called boiling point.

Factors affecting boiling point:

1. Nature of liquid.
2. Intermolecular attraction.
3. External pressure.

Nature of liquid:
The different liquids have different boiling points at same pressure.

Inter-molecular attraction:
The boiling point is directly proportional to the inter-molecular attraction, greater the inter-
molecular attraction, greater will be the boiling point and vice versa.

External pressure:
The boiling point is directly proportional to external pressure i.e greater the external pressers greater
will be the boiling point and vice versa e.g for H2O (Water).
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DISTILLATION

Definition:
A process which we vaporize a substance, withdraw condense is called distillation.

Types:
There are three types of distillation
1. Fractional distillation.
2. Vacuum distillation.
3. Destructive distillation.
Applications of boiling point:
1. Pressure cooker
2. Vacuum distillation

1. Pressure cooker:
Principle:
The working principle of pressure cooker is
“The boiling point of a liquid increases with increased external pressure”

Working:
(i) Formation of vapours:
When a liquid is heated in a pressure cooker, which is a closed container, more and more
vapours are formed over the surface of liquid, exerting more pressure

(ii) Increase in pressure:

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These vapours are not allowed to escape. In this way, they develop more pressure in the
cooker and the boiling point of H2O increases. As more heat as absorbed in water. So, food as
cooked quickly under increased pressure.

2. Vacuum distillation:
Definition:
The distillation process that is carried out under reduced pressure or an vacuum is called vacuum
distillation.

Principle:
Boiling points are lowered at lower external pressure.

Working:
Some liquids with high boiling points can decompose of distilled. In order to boil or distill them at
lower temperature, pressure is lowered or distillation is carried out under vacuum.

Example of glycerin:
Glycerin boils at 290°C at 760 torr (1 atm) pressure but decomposes at this temperature. Hence,
glycerin cannot be distilled at 290°C. Under vacuum, the boiling point of glycerin decreases to 120°C at 50
torr. It is distilled at this temperature without decomposition and hence, can be purified easily.

Advantages:
 It decreases the time, for distillation process.
 It is economical because less fuel is required.
 The decomposition of many compounds can be avoided.

PHASE CHANGE

During the phase change, there is certainly energy change but temperature does not change.
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HEAT OF VAPOURIZATION

Definition:
The amount of heat energy required to vaporize a given liquid is called vaporization.
Molar heat of vaporization (∆Hv):
The amount of heat energy required to convert one mole of liquid into is vapours at 1 atm at its
boiling points is called molar heat of vaporization. E.g, H2O
Molar heat of vaporization =40.6 KJ/mol
Heat of vaporization is inversely proportional to temperature.
Hv ∝ 1/T

Molar heat of fusion (∆HF):

Definition:
The amount of heat required to convert one mole of a solid into its vapours at particular temperature
and1atm pressure is called molar heat of sublimation.
All these are called enthalpy changes. These are positive because they are endothermic process.

Enthalpy changes:
Definition:
The heat change during a physical or chemical change at constant pressure is, called enthalpy
change.
LIQUID CRYSTALS

Discovery:
In 1888, Frederick Reinitzer is Austrian botanist discovered the liquid crystal.
Definition:
The turbid liquid phase of a solid that exists in between the melting and clearing temperature is
called liquid crystal.

Formation:
An organic compound cholesterol benzoate was studied. It twins milky liquid at 145°C and becomes
a clear liquid at 179°C. When the substance is cooled, the reverse process occurs. This turbid liquid phase is
called liquid crystal.

Characteristics:
Liquid crystal has both properties of liquids and crystals (solid).
 Like liquids: Viscosity, surface tension, fluidity.
 Like crystals: Optical properties i.e, arrangement of molecules is not damaged and they move in
a group.

Isotropic nature:
A crystalline solid may be isotropic or an isotropic but liquid crystal is always isotropic.
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Applications:
From 1888 to until about 30 years ago, liquid crystals were largely a laboratory curiosity but now
they have found a large number of applications.

Composition:
The substances which make the liquid crystals are often composed of long rod like molecules.
 In normal liquid phase these molecules are oriented in random directions.
 In crystalline liquid phase, they develop some ordering of molecules.

Types:
Depending upon the nature of ordering, liquid crystal can be divided into:
 Nematic
 Smectic
 Cholesteric

USES OF LIQUID CRYSTALS

Due to remarkable optical and electrical properties, liquid crystals find may practical application. Many
organic compounds and biological tissues behave as liquid crystals. The unique properties of liquid crystals
have intrigued the scientists since their discovery nearly hundred years ago. Some of their important uses are
as follows:
1. As temperature sensor:
Like solid crystals, liquid crystals can diffract light. When one wavelength of white Light is
reflected, from a liquid crystal it appears coloured. As the temperature charges, the distances
between layers of molecules of liquid crystal change. Therefore the reflected light changes
accordingly. Thus liquid crystals can be used as temperature sensors.

2. To find out potential failure/room thermometers:

Liquid crystals are used to find the point of potential failure in electrical circuits. Room
thermometers also contain liquid crystals with a suitable temperature range.

3. Location of infected parts and breast cancer:

Liquid crystalline substances are used to locate veins, arteries, infections and tumors. The
reason is that these parts of the body are warmer than surrounding tissues. Specialists can use the
techniques of skin thermography to detect blockages in veins and arteries. When a layer of liquid
crystal is painted other surface of the breast, a tumor shows up as a hot area which is coloured blue.
This technique has been successful in the early diagnosis of breast cancer.

4. Electrical devices:
Liquid crystals arc used in the display of electrical devices such as digital watches,
calculators and laptop computers. These devices operate due to fact that temperature, pressure and
electromagnetic fields easily affect the weak bonds, which hold molecules together in liquid crystals.

5. Solvents in chromatography:
In chromatographic separations, liquid crystals are used as solvent.
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6. Oscillograph and TV displays:

Oscillograph and TV display also use.

SOLIDS
Definition:
Those substances which are rigid, hard, having definite shape definite volume and cannot flow are
called solids.

Important points:
 The atoms, ions or molecules that make up a solid are closely packed.
 They are held together by strong cohesive forces.
 The constituent particles of solids cannot move at random. They show only vibrational motion.
 There exists a well ordered arrangement in solids.

Difference between crystalline and amorphous solid:

Crystalline solid Amorphous solid
 Those solids in which atom, ions or molecules  Those solids in which atom, ions or molecules do
are arranged in a definite three dimensional not possess a regular orderly arrangement are
pattern are called crystalline solid. called amorphous solids.
 These solids have a sharp melting point.  These solids do not have a sharp melting point.
 They are also called true solid.  They are also called:
(i) Super cooled liquid
(ii) Highly viscous liquid
(iii) Pseudo solids
 They have definite heat of fusion  They do not have definite heat of fusion.
 Examples  Examples:
(i) NaC1 (i) Rubber
(ii) Sugar (ii) Plastic
(iii) Ice (iii) Glass
(iv) Diamond etc (iv) Glue etc

How can you convert crystalline solid into amorphous solid:

Many crystalline solids can be changed into amorphous solids by melting them and then cooling the molten
mass rapidly. In this way, the constituent particles do not find time to arrange themselves. Thus a crystalline,
solid is changed into amorphous solid.

What are crystallites?

The small parts of amorphous solids which possess orderly arrangements of constituent particles are called
crystallites.

Properties of crystalline solids:

(i) Geometric shape:
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All the crystalline solids have a definite distinctive geometrical shape due to definite and
orderly arrangement of atoms, ions or molecules in three dimensional spaces. For a given
crystals, the interfacial angles, at winch the surfaces intersect are always the same no matter in
which shape, they are grown. The faces and angles remain characteristic even when the material
is ground to a fine powder.
(ii) Melting points:
Crystalline solids have sharp melting points and can be identified from their definite
melting points.

(iii) Cleavage planes:

Whenever the crystalline solids are broken they do so along definite planes. These planes
are called the cleavage planes. They are inclined to one another at a particular angle for a given
crystalline solid. The value of this angle varies from one solid to another solid.

(iv) Anisotropy:
Definition:
The phenomenon in which a crystalline solid shows variation in certain phyical properties
depending upon the direction is called anisotropy.

Anisotropy properties:
Anisotropic properties are:
 Refractive index.
 Thermal and electrical conductivities.
 Co-efficient of thermal expansion.
 Cleavage planes.

Reason:
The variation in anisotropic properties with direction is due to the fact that the orderly
arrangement of particles in crystalline solids is different indifferent directions.

Examples:
(i) Electrical conductivity of graphite is greater in one direction than in another. Actually,
electrons in graphite are mobile for electrical conduction parallel to the layers only.
Therefore, its conductivity in this direction is far better than perpendicular to their
direction.
(ii) Cleavage itself is an isotropic behavior.

(v) Symmetry:
Definition:
The repetition of angles, edges or faces when a crystal is rotated about 360o along its axis is
called symmetry.

Symmetry elements:
Following are the symmetry elements of a crystal:
(i) Plane of symmetry.
(ii) Axis of symmetry.
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(iii) Centre of symmetry.

(vi) Habit of a crystal:

Definition:
The shape of a crystal in which it usually grows is called habit of a crystal.

Preparation:
A crystal can be prepared:
(i) By moderate cooling of a saturate solution.
(ii) By slow cooling of a liquid.
(iii) By growing indifferent directions.
If the conditions for growing a crystal are maintained then the shape of the crystal always
remains the same. If the conditions are changed the shape of the crystal may change.
For example: A cubic crystal of NaCl becomes needle like when 10% urea is present in its
solution as an impurity.

(vii) Isomorphism:
Definition:
The phenomenon in which two different substances exist in the same crystalline form is
called isomorphism.

Isomorphs:
Those crystalline substances which’ have same crystalline, form are called isomorphs to each
other.
Isomorphs may be compounds or elements.
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(viii) Polymorphism:

Definition:
The phenomenon in which a compound exists in more than one crystalline form is called
polymorphism.

Polymorphs:
The compound which exists in more than one crystalline form is called polymorphic
compound and these forms are called polymorphs of each other.

Important points:
 It is a compound phenomenon.
 Polymorphs have same chemical properties.
 Polymorphs have different physical properties due -to the different structual arrangement
of their particles.

(ix) Allotropy:
Definition:
The phenomenon in which an element exists in more than one crystalline forms is called
allotropy and these forms of the element are, called allotropes or allotropic forms.

Important points:
 It is an elemental phenomenon.
 Allotropes of an element have same chemical but different physical properties.

(x) Transition temperature:

Definition:
The temperature at which two crystalline forms of the same substance can co-exist in
equilibrium with each other is called transition temperature.
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Important points:
 At transition temperature one crystalline form of a substance changes to another.
 Above and below this temperature only one form exists.
 The transition temperature, of allotropic forms of an element always less than its melting
point.

CRYSTAL LATTICE

Definition:
An array of points representing atoms, ions or molecules of, a crystal arranged at different sites in
three-dimensional space is called crystal lattice.

Lattice sites:
The points or sites representing atoms ions or molecules in a crystal lattice are called lattice point or
lattice sites.

Cubic crystal lattice

Cubic crystal lattice

Unit cell:
Definition:
The smallest part of the crystal lattice that has the entire characteristic features of the entire crystal is
called unit cell.

Important points:
 A unit cell shows the structural properties of a given crystal.
 When a unit cell is repeated in three-dimensions, it gives the entire crystal.
 The complete information about crystalline structure is present within its unit cell.
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Unit cell dimensions/crystallographic elements:

There are six (6) unit cell dimensions
(A) Unit cell lengths:
Length along x-axis = ‘a’
Length along y-axis = ‘b’
Length along z-axis = ‘c’

(B) Unit cell angles:

Angle between lengths ‘a’ and ‘b’ = γ
Angie between lengths ‘b’ and ‘c’ = α
Angie between lengths ‘c’ and ‘a’ = β
The unit cell lengths a, b, c may be assigned along x, y and z axis, respectively & it angles α, β and γ
have to be decided accordingly.

CRYSTALS AND THEIR CLASSIFICATION

Crystal system:
A crystal system may be identified by the dimensions of its unit cell along. It’s three edges or axes a,
b, and c and three angles between the axis α, β and γ.
There are seven crystal systems. These seven crystals systems are described as follow:

1. Cubic system:
Lengths:
In this system, all the three axes are of equal lengths.
a=b=c
Angles:
All axes are at right angles to each other.
α = β = γ = 900
Examples:
Fe, Cr, Ag, Au, NaCl, NaBr, Diamond.

2. Tetragonal system:
Lengths:
In this system two axes are of equal lengths and the third axis is either shorter or larger than the other
two.
a=b≠c
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Angles:
All the angles are of 900.
α = β = γ = 90o
Examples:
Sn, SnO2, MnO2, NH4Br

3. Orthorhombic or rhombic system:

Lengths:
All the three axes are of unequal length.
a≠b≠c
Angles:
All axes are at a right angle to each other.
α = β = γ = 90o
Examples:
Iodine, Rhombic, Sulphur, BaSO4, K2SO4

4. Rhombohedral or trigonal system:

Lengths:
All the thee axes are of equal length.
a=b=c
Angles:
Three angles are not equal and lie between 900 and 1200.
α ≠ β ≠ γ ≠ 90o
Examples:
Bi, Al2O3, NaNO3, KNO3

5. Hexagonal system:
Lengths:
In this system two axes are of equal length and the third axis is of different length.
a=b≠c
Angles:
Two axes are join one plane making an angle of 120° with each other and the third axis is at right
angle to these two axes.
α = β = 90o, γ = 120
Examples:
Grapite, ZnO, CdS, Ice, Zn, Cd

6. Monoclinic system:
Lengths:
In this system, all the three axes are of unequal length.
a≠b≠c
Angles:
Two axes are at right angle to each other while the third angle is, greater than 900.
α = γ = 90o, β ≠ 900
Examples:
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Sugar, sulphur, borax, Na2SO4, 10H2O

7. Triclinic system:
Lengths:
All the three axes are of unequal length.
a≠b≠c
Angles:
All the three angles are unequal and none of these angles is 90°.
α ≠ β ≠ γ ≠ 90o
H3PO3, K2Cr2O7, CuSO4.5H2O

CLASSIFICATION OF SOLDS

There are four types of crystalline solids depending upon type of bonds present in them:
1. Ionic Solids
2. Covalent Solids
3. Metallic Solids
4. Molecular Solids

Ionic Solids:
Definition:
The crystalline solids in which the particles forming the crystal are positively and negatively charged
ions which are held together by strong electrostatic forces of attraction (Ionic bond) are called ionic solids.

Examples:
The crystals of NaCl, KBr etc are ionic solid

Properties of ionic solids:

1. Physical State:
The cations and anions are arranged in a well defined geometrical pattern, so they are
crystalline solids at room temperature. Under ordinary conditions of temperature and pressure they
ever exist in the form of liquids or gases.
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2. Hardness, volatility and melting, boiling point:

Ionic crystals are very stable compounds. Very high energy is required to separate the cations
and anions from each other against the forces of attraction. That’s why ionic solids are very hard,
have low volatility and high melting and boiling points.

3. Nature of ionic solids:

Ionic solids do not exist as individual neutral independent molecules. Their cations and
anions attract each other and these forces are non directional. The close, packing of the ions enables
them to occupy minimum space. A crystal lattice is developed when the ions arrange themselves
systematically in an alternate manner.

4. Radius ratio:
The structure of ionic crystals depends upon the radius ratio of cations and onions.
Radius ratio =
In NaCl,
Na+ =95 pm
-
Cl =181pm
Radius ratio = = 0.525
NaCl and CsF have the same geometry because the radius ratio in both the cases as same.

5. Formula mass of ionic solids:

In the case of ionic crystals, we always talk about the formula mass of these substances and
not the molecular mass because they do not exist in the form of molecules.

6. Conduction of electricity:
(a) In solid state:
In solid state, ionic crystals do not conduct electricity because on account of electrostatic
force existing between them the cations and anions remains tightly held together and hence
occupy fixed positions.

(b) In solution or in molten state:

Ionic crystals conduct electricity when they are in the solution or in molten state. In both
cases, ions become free.

7. Brittleness:
Ionic solids are highly brittle because ionic solids are composed of parallel layers, which
contain cations and anions, in alternate positions, so that the opposite ions in the various parallel
layers lie over each other. When an external force is applied, one layer of ions slides a bit over the
other layer along a plane. In this way, the like, ions come in front of each, other and hence begin to
repel. So, the application of little external force develops repulsion between two layers causing
brittleness.
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8. Density:
Ionic solids are mostly of high density due to close packing of ions.

9. Ionic reactions:
Such compounds having the ionic crystals give ionic reactions in polar solvents and these are
very fast reactions.

10. Isomorphism and polymorphism:

The properties like isomorphism and polymorphism are also associated with the ionic
crystals.

Structure of sodium chloride:

The structure of ionic crystals depends upon the structure arid the size of ions. Each ion is
surrounded by a certain number of ions of opposite charge. In the structure of NaC1, each Na+ ion is
-
surrounded by six chloride Cl ions. These ions are arranged in a crystal lattice. Na+ has ten electrons while
- -
Cl has total eighteen electrons. The size of Cl is bigger than that of Na+.

Distance between ions:

-
The distance between two nearest ions of same kind i.e., Cl ions is 5.63Ao. The distance between
two adjacent ions of different kind is 5.63/2 =2.815 A°.

Location of ions:
The location of Na+ and Cl- ions is such that each Na+ is surrounded by six Cl- placed at the corners
of regular octahedron.

Coordination number:
Coordination number of an ion is equal to number of ions surrounding that ion.
Coordination number of each Na+ is six similarly each Cl- ion is also surrounded by six Na+ ions. Na+ and
Cl- are not connected to each other by pairs because all six Cl- ions are at same distance away from one Na+.
The independent moloeules of NaCl do exist in vapour phase. Anyhow in solid NaCI, there are no
independent molecules of NaCI. That’s why sodium chloride is said to have formula unit of NaCI.

Cubic structure:
There are eight Cl- ions at the corners of cube and each is shared by among eight cubes. l/8th part of
each Cl- ion is considered for this unit cell. So one complete Cl- is contributed by eight corners. Similarly,
six chloride ions are present at the face centers and each is being shared by two cells.
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No. of ions per unit cell:

Thus, per unit cell, there are 8/8 + 6/2 = 4C l- ions. If we take a unit cell having 8Na+ at eight corners
and 6Na+ at faces so there are equal number of Na+ ions and therefore 4NaCI formula units are present per
unit cell.

Covalent solids:
Those crystalline solids which consists of atoms of same or different elements held together through
covalent bonds are called covalent solids.
Covalent solids are also called atomic solids.

Examples:
Diamond, graphite SiC, BN, etc

Types of covalent solids:

Covalent solids are of two types:
(i) Giant structure covalent solids:
When the covalent bonds join to form giant molecules e.g., diamond, silicon carbide,
aluminium nitride

(ii) Layered structure covalent solids:

When the atoms join to form the covalent bonds and separate layers are produced.
e.g. cadmium iodide, graphite, boron nitride.

Properties of covalent crystals:

1. Three dimensional open structure:
The bonding in covalent crystals extends in three dimensions. They contain a network of
atoms. The valencies of atoms are directed in definite directions. So the packing of atoms in these
crystals is looser than those of ionic and metallic crystals. Thus covalent crystals have open structure.

2. Hardness, volatility, melting and boiling point:

These crystals are very hard and considerable amount of energy is required to break them.
They have high melting and boiling points and their volatility is very low.

3. Conduction:
Due to absence of free electrons and ions they are bad conductors of electricity.
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Graphite is a conductor:
In graphite, each carbon atom is sp2 hybridized. Three out
of four electrons of a carbon atom form covalent bond with
neighboring atoms. Thus a hexagonal layered structure of graphite
is formed and the electrons are available between the layers. These
electrons are delocalized and conductivity becomes possible.
Graphite is not a conductor perpendicular to the layers.

4. Solubility:
 Mostly covalent crystalline solids are insoluble in polar
solvents like water.
 They are readily soluble in non-polar solvents like benzene
and carbon tetrachloride.
 The covalent crystals having giant molecules like diamond
and silicon carbide are insoluble, in all the solvents.
Because of their big, size, they do not interact with the
solvent molecules. The chemical reactions of such
crystalline solids are very slow.

Structure of diamond:
Diamond is one of the allotropic modifications of carbon. Carbon has four electrons in its outermost
shell. The four atomic orbitals (one 2s and three 2p) undergo sp3 hybridization to give four sp3 hybridized
orbitals. They are directed in space along the four corners of tetrahedron. This is the unit cell of diamond
and a larger number of such unit cell undergo sp3— sp3 overlapping to form a huge structure. Each carbon
atom is linked with four other carbon atoms. The bonds between carbon atoms are covalent which run
through the crystal in three dimensions. All the bond angles are 109.5° and the bond length are 154pm. The
whole lattice is, therefore, continuous and because of continuity of C— C covalent bonding. The entire
diamond crystal behaves as a huge or giant three dimensional carbon molecule. This is also called
“macromolecule”. The overall structure of diamond looks face-centered cubic.

Molecular solids:
Those solid substances in which the particles forming the crystals are polar or non-polar molecule or
atoms of a substance are called molecular solids.

Types of attractive forces:

Two types of intermolecular forces hold them together:
(i) Dipole-dipole forces
(ii) Van der Waal’s forces
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These intermolecular forces are much weaker than the forces of attraction between the cations and the
anions in ionic crystals and between the atoms in covalent crystals.

Types of molecular solids:

1. Polar molecules containing molecular solids.
They have high melting and boiling points and less volatile and are soluble in water. e.g., Ice,
sugar
2. Non-polar molecules containing molecular solids.
They have low melting and boiling points. They are soluble in nonpolar solvents and are
more volatile. e.g, Iodine (I2), Phosphorous (P4), Carbon dioxide (CO2), Sulphur (S8)

Properties of molecular solids:

Regular arrangement:
X-rays analysis has shown the regular arrangements of atoms in the constituent molecules of these
solids and we get the exact position of all the atoms.

Softness:
The forces, which hold the molecules together in the molecular crystals, are very weak, so they are
soft and easily compressible.

Volatility, melting and boiling points:

They are mostly volatile and have low melting and boiling points.

Conduction, solubility and density:

They are bad conductors of electricity, have low densities and sometime transparent to light Polar
molecular crystals are mostly soluble in polar solvents while non-polar molecular crystals are usually
soluble in non-polar solvents.

Structure of solid iodine:

In solid state, the molecules of iodine align in the form of layer lattice I - I. bond distance is 271.5
pm and is appreciably longer than in gaseous state. As expected from its structure, iodine is a poor conductor
of electricity.

Metallic solids:
Those crystals in which metal atoms are held together by the attractive forces between a metal cation
and mobile electrons with its sphere of influence are called metallic crystals.

Metallic bond:
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The attractive force which binds a metal cation to a no. of electrons within its sphere of influence is
called metallic bond.

Theories about the metallic bond:

1. Electron pool or electron gas theory:
This was proposed by Drude and extended by Loren in 1923.

Definition:
Each atom in a metal crystal loses all of its valence electrons. This valence electron is form a
pool or a gas. The positively charged metal ions are held together by electron pool or gas.

Explanation:
Due to low ionization energy, metal loses electrons and forms positively charged ions. These
positively charged ions occupy definite positions at measurable distances from each other in the
crystal lattice. Valence electrons are not attracted to any individual ion or a pair of ions rather
belongs to crystal as a whole. These electrons are free to move about from one part of the crystal to
other. The force which binds a metal cation to a number of electrons within its sphere of influence is
called metallic bond.

2. Valence bond theory:

Louis Pauling tried to explain metallic bond according to valence bond theory. According to
this theory, The metallic bond is treated essentially as covalent in character. However, it is assumed
that, the covalent bonds are not localized but are highly delocalized in metal structure.

3. Molecular orbital theory:

Recently, molecular orbital theory was applied to explain the characteristics of metallic solids
According to this theory, it is assumed that electrons in the completely filled orbitals are essentially
localized, while atomic orbitals containing valence electrons interact or overlap to form a set of
delocalized orbitals. These delocalized orbitals are the molecular orbitals which extend over the
entire crystal lattice. Such a combination of atomic orbitals produces as a large number of closely
spaced states. These states of energy are also known as bonds of energy. That is why it is also called
a band theory. The energy gap between two bands determines the properties of the metallic solids.

Properties of metallic crystals:

Electrical conductivities of metals:
Metals are good conductor of electricity. When electric field is applied between two ends of a metal
then mobile electrons begin to move towards the positive pole and the new electrons from the negative pole
take their place.
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Effect of temperature on electrical conductivities of metals:

The electrical conductivities of metals decrease with the increase in temperature. The reason is that
with the increase in temperature the positive metal ions begin to oscillate, their amplitude increases and the
motion hinders the free movement of mobile electrons between the positive ions. This hindrance decreases
the electrical conductivity.

Thermal conductivity:
It is another property associated with metallic solids. When a piece of metal is heated at one end, the
mobile electrons at this end absorb heat energy and move very rapidly through metallic lattice towards the
cooler end. During the process, they collide with adjacent electrons and transfer their beat energy to them.

Metallic luster:
Whenever, the metals are freshly cut, most of them possess metallic luster which means that they
have a shinning surface. When light falls on the metallic surface, the incident light collide with the mobile
electrons and they are excited. These e1ectrons when de-excited give off some energy in the form of light.
This light appears to be reflected from the, surface of the metal which gives a shining look.

MALLEABILITY AND DUCTILITY

Metals are malleable and ductile whenever stress is applied on them. Their layers slip passes each
other. The structure of the metal changes without fracturing.
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STRUCTURE OF MATELS

The spaces during the packing are, larger when the box is shaken, the balls rearrange,

The arrangements of these balls are now stable and more completely packed. It is the natural tendency of the
balls to have closely packed arrangement eleven spheres.

In order to understand, how various unit cell of the crystal lattice are developed, consider three balls which
join together in one plane. The forth ball is inserted in the space created by other three as second layer. In
this way, a tetrahedral structure is obtained. Actually fourth ball of the second layer is placed in the
depression created by first three balls. These depressions are also, called “interstices or crevices or voids”

Consider eleven balls are present in first layer. The balls of second layer can fit into the depressions or
interstices created by the first layer. When the balls of the second layer are arranged, then all the depressions
of first layer are not occupied. There are two types of depressions as ‘a’ and ‘b’. The depressions marked ‘b’
is not occupied by the second layer and one can see the ground from looking at the top through depressions
‘b’. The new depressions marked ‘a’ are created by second layer. Through the depressions ‘a’, we cannot see
the ground but falls of the first layer.
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Now arrange the balls of third layer in the depressions of second layer. When the balls of the third layer are
placed above the second layer then there are two possibilities. The third layer balls may be accommodated in
‘a’ type or ‘b’ type interstices or depressions.

Cubic close packing:

When the atoms of third layer fit into interstices marked ‘b’ then atoms of the third layer will not be
lie directly above those of the atoms of first layer. This pattern of arrangement is called ABC ABC……… or
123 123…….. It is named as face centered cubic arrangement. The balls of fourth, seventh and tenth layers
will be in front of each other.

Hexagonal close packing:

When the atoms of third layer are arranged in such a way that they occupy the depressions created by
second layer i.e., in the ‘a’ types crevices then those atoms will directly lie above the atoms of first layer.
This pattern of arrangement is usually, written as ABAB….. or l212.....
This pattern has been named as hexagonal close packing. The balls of third, fifth and seventh layers will be
in front of each other.
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KEY POINTS

 Dipole-dipole forces, Debye’s forces, London’s forces and Hydrogen bonding are intermolecular
forces.
 Inter-molecular forces depend upon (polarity, state of matter, polarizability, shape and size of
molecules).
 Hydrogen Bonding is responsible for low density of ice and survival of aquatic life in water in
winter. It is also present in biological compounds e.g., DNA, proteins etc.
 Evaporation is a cooling process.
 When vapour pressure of a liquid becomes equal to external pressure, it starts boiling.
 Boiling Point of a liquid can be increased or decreased by varying external pressure.
 Liquid crystals are a mesomorphic state of matter existable between melting and clearing
temperatures.
 They have fluidity like liquids and optical properties like crystals.
 They find major applications is thermal sensors, clinical diagnosis and liquid crystal display (LCD)
and as solvent in chromatography.
 Solids are either crystalline or amorphous.
 Crystalline solids have three dimensional arrays of points where atoms, ions or molecules are
present.
 Unit cell is the smallest part of crystalline solid showing all properties of chunk of crystal.
 Seven crystal systems are there. Simplest one is cubic.
 Crystalline solids are of four type (a) ionic (b) covalent (c) molecular (d) metallic.
 Ionic solids are brittle, have high M.P°, B. P°, densities, are non- conductor in solid state and
electrolytes in solution or molten form. e.g., NaCI(aq).
 Covalent solids are macro molecules held together by a net work of covalent bond. They are hard
and non conductor e.g., diamond, silica.
 Molecular solids are either polar e.g., ice or non-polar e.g., I2 crystal held together by weak
intermolecular forces.
 Metallic solid are joined by metallic bond. Here cations are immersed in a pool of electrons e.g.,
copper, gold, silver, etc
 Metals are conductors, malleable and ductile.
 Metals have cubic or hexagonal closed packing of atoms in crystals.