Solids, Chemistry tutorial

Introduction:

The general characteristics of the solid state are highlighted and illustrated by employing the kinetic theory. We are very much familiar that the closeness and low average kinetic energy of solid particles are necessary for strong attractive forces by keeping the particles in fixed positions. Solid matter thus takes place naturally in different shapes and sizes example: small stones, big rocks and mountains.

Scientists have as well synthesized, modified, shaped and reshaped solid mater example: polymetric materials like fibres, resins, rubbers and plastics are available in different shapes, sizes and colors: plastic buckets, cups and so on, textile materials, electrical parts: plugs; cables and sockets, motor vehicle and machinery parts are just few examples. Unlike the gas and liquid matter that require containers for storage and moving them from place to place, solid matter is the simplest to handle and move around. The given are some of the very significant use of solids; shelter, support, for movement and for food.

Solid Classification:

Solids are categorized as crystalline or amorphous by virtue of the order in particle arrangement. Crystalline solids are basically built from simple structural units that repeat themselves in (3-D) three dimensions to provide the solid its characteristic geometric form. Most of the metal and ionic solids are crystalline in nature.

Other solids like glass, rubber and plastics don't encompass definite geometrical forms and are known as amorphous. The regular and repeated arrangement of particles in the crystalline form a pattern known as the crystal lattice.

Solid Properties:

Solids are usually characterized through hardness or fineness and high mechanical strength. The mechanical strength of any solid material is found out by the kind of bonds holding the particles in position in the solid structure.

The kind of force in turn based on the chemical nature of the units which occupy the fixed position in the lattice. Solids are grouped into four groups based on the nature of the intermolecular forces between the particles of the solid matter.

The molecular solid:

The lattice points in molecular solids are occupied via molecules that don't carry any charge. The forces holding the molecules altogether are of two kinds, the dipole-dipole interaction for polar molecules and Van der waals forces that are the non-polar molecules. We are familiar that these forces that is, the dipole-dipole and Van der Waals forces are somewhat weak forces. This is why molecular solids are generally soft and of extremely low melting points. The lack of electrical charge in the structure makes them non-conductors. They are employed mainly as insulators. Illustration of compounds that form molecular solids are ammonia, ice-water, candle wax and carbon (iv) oxide. Most of the molecular solids are liquids or gases at normal laboratory pressures and temperatures. Water for instance is a liquid and carbon(iv) oxide is a gas at normal lab temperature.

The metallic solid:

The lattice points in the metallic solid are occupied via the positive core of the metallic atom surrounded through an electron cloud. The force of attraction here is electrostatic among the positive metal cores and the electron cloud. This electrostatic force is strong and is accountable for the compact structure of metals and their very high boiling and melting points. The mobility of the electron cloud beneath the influence of heat and electricity accounts for the high thermal and electrical conductivities of metals. Illustrations of metallic solids are iron, copper, lead, aluminium and zinc.

The ionic solid:

The units engaging the lattice points in ionic solids are negative and positive ions. There is extremely strong attraction between the negative and positive ions. In the solid structure each in is surrounded via as many ions as possible of the opposite charge leading to a giant assembly of ions held altogether in a rigid structure.

Ionic solids are featured via the given properties:

1) High boiling and melting points.

2) They are non-conductors of electricity in the solid state however conduct whenever molten or in solution.

3) As the energy to separate the ions is relatively high, ionic solids are hard and fragile. Illustrations of ionic solids are NaCl, KCl, KNO3 and CuSO4.

The covalent solid:

The lattice points in covalent solids are occupied through atoms linked altogether via a Continuous system of covalent bonds. The shared electrons among the atoms in the lattice yield in strong bonding between the atomic nuclei. Illustrations of covalent solids are diamond, graphite and quartz (SiO2). If all the valence electrons are comprised in bonding as in diamond and quartz, the solid will not conduct electricity. In a few arrangements where all valence electrons are not employed in bonding as in graphite, electrical conductivity is possible as the free electrons move beneath the affect of applied electrical potential.

The Structure of Diamond and Graphite:

We are familiar that diamond and graphite are allotropes of carbon.

Both of them are as well covalent solids however while diamond is hard, encompass very high density and melting point, graphite is soft, less dense and encompass a much lower melting point. In both diamond and graphite, the lattice points are occupied through carbon atoms. In diamond, carbon atoms are tetrahedral bonded by employing Sp3 hybridized orbitals. All the four valence electrons of carbon are employed in bonding resultant in a rigid (3-D) three dimensional network. This accounts for the extremely hard nature of diamond. Diamond is the hardest substance acknowledged. It can't be cut through any other substance; therefore it is employed in cutting glass and metals.

In the structure of graphite though, each and every carbon atom is linked to three others by employing Sp2 hybridized orbitals. The carbon atoms are arranged in layers one above the other. The layers are held altogether by weak van der waal forces. Graphite consists of a relatively high melting point and it is less dense than diamond. The presence of mobile electrons in the crystal lattice makes graphite a very good conductor of electricity.

It is well known that carbon consists of four valence electrons and only three are employed in bonding in graphite however all the our electrons are employed in diamond. Graphite is employed as an inert electrode in electrolysis. The layered structure of graphite allows one layer to slide over the other easily accounting for the lubricating property of graphite. Dissimilar oil, it is non-sticky and is generally employed on bicycle chains and for bearings of a few motor cars.

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Fig: Structure of diamond and Hexagonal rings in graphite

A few properties of diamond and graphite are summarized in the table illustrated below:

Property

Diamond

Graphite

1. Appearance

Colorless, transparent solid, sparks when cut

Black opaque with metallic lustre

2. Crystalline Structure

Three dimensional network

Two dimensional network

3. Density

Very dense

Less dense

4. Hardness

Hardest natural substance known

Soft

5. Electrical conductivity

Non conductor

Good conductor

6. Chemical activity

Inert unreactive

More reactive

7. Uses

- Employed as a precious jewel

- Drilling rocks and cutting glass

- To sharpen tools

- Pivot supports in precision instruments

- Dies for drawing wires

- Dry lubricant

- Electrodes in electrolysis

- Employed in lead pencil

- Black pigment in paints

- Neutron moderator in atomic reactors

Melting and Melting Point:

Heating of solid matter raises the average kinetic energy of the particles. We are familiar that at a specific temperature feature of the solid particles will get adequate energy and overcome the attractive forces holding them in the fixed positions. At this temperature the solid structure collapses and a liquid is produced. This is known as the melting point. We are as familiar that melting of a pure solid sample takes place at a constant temperature and is a test of purity.

The melting point of a solid sample is found out through the capillary tube process. A small amount of the solid sample is put in a glass capillary tube. The capillary tube is joined to a thermometer by employing a rubber band. The capillary tube is heated in a transparent liquid bath and the temperature at which the solid begins melting is noted. For a pure solid sample, this temperature remains constant till the entire solid is melted.

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Fig: Melting point of naphthalene

The electrically operated melting point container can as well be employed however the equipment is fairly costly. The apparatus or container is fitted by a thermometer. It as well consists of apartments for placing the capillary tube. After the capillary tube having the solid sample is put into the apparatus, the apparatus is switched on. As the solid heats up, the procedure is monitored via a magnifying glass in front of the apparatus till melting takes place.

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