Chemistry

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Group 14 (Carbon Group) Elements

Occurrence:

Carbon is found in the native form as graphite. Coal, crude oil and carbonate rocks such as calcite, magnesite etc contains large quantities of carbon in its combined form with other elements. Silicon occurs as silica (sand and quartz crystal). Silicate minerals and clay are other important sources for silicon.

Physical Properties:

Some of the physical properties of the group 14 elements are listed below

Table 2.4 Physical properties of group 14 elements

Tendency for Catenation

Catenation is an ability of an element to form chain of atoms. The following conditions are necessary for catenation.

• The valency of element is greater than or equal to two
• Element should have an ability to bond with itself
• The self bond must be as strong as its bond with other elements
• Kinetic inertness of catenated compound towards other molecules.

Carbon possesses all the above properties and forms a wide range of compounds with itself and with other elements such as H, O, N, S and halogens.

Allotropes of Carbon

Carbon exists in many allotropic forms. Graphite and diamond are the most common allotropes. Other important allotropes are graphene, fullerenes and carbon nanotubes.

Graphite is the most stable allotropic form of carbon at normal temperature and pressure. It is soft and conducts electricity. It is composed of flat two dimensional sheets of carbon atoms. Each sheet is a hexagonal net of sp² hybridised carbon atoms with a C-C bond length of 1.41 Å which is close to the C-C bond distance in benzene (1.40 Å). Each carbon atom forms three σ bonds with three neighbouring carbon atoms using three of its valence electrons and the fourth electron present in the unhybridised p orbital forms a π-bond.

These π electrons are delocalised over the entire sheet which is responsible for its electrical conductivity. The successive carbon sheets are held together by weak vander Waals forces. The distance between successive sheet is 3.40 Å. It is used as a lubricant either on its own or as a graphited oil.

Unlike graphite the other allotrope diamond is very hard. The carbon atoms in diamond are sp³ hybridised and bonded to four neighbouring carbon atoms by σ bonds with a C-C bond length of 1.54 Å. This results in a tetrahedral arrangement around each carbon atom that extends to the entire lattice as shown in figure 2.5. Since all four valance electrons of carbon are involved in bonding there is no free electrons for conductivity. Being the hardest element, it used for sharpening hard tools, cutting glasses, making bores and rock drilling.

Fullerenes are newly synthesised allotropes of carbon. Unlike graphite and diamond, these allotropes are discrete molecules such as C₃₂, C₅₀, C₆₀, C₇₀, C₇₆ etc. These molecules have cage like structures as shown in the figure. The C₆₀ molecules have a soccer ball like structure and is called buckminster fullerene or buckyballs.

It has a fused ring structure consists of 20 six membered rings and 12 five membered rings. Each carbon atom is sp² hybridised and forms three σ bonds & a delocalised π bond giving aromatic character to these molecules. The C-C bond distance is 1.44 Å and C = C distance 1.38 Å.

Carbon nanotubes, another recently discovered allotropes, have graphite like tubes with fullerene ends. Along the axis, these nanotubes are stronger than steel and conduct electricity. These have many applications in nanoscale electronics, catalysis, polymers and medicine.

Another allotrophic form of carbon is graphene. It has a single planar sheet of sp² hybridised carbon atoms that are densely packed in a honeycomb crystal lattice.

Carbon Monoxide [CO]:

Preparation:

Carbon monoxide can be prepared by the reaction of carbon with limited amount of oxygen.

2C + O₂ → 2CO

On industrial scale carbon monoxide is produced by the reaction of carbon with air. The carbon monoxide formed will contain nitrogen gas also and the mixture of nitrogen and carbon monoxide is called producer gas.

The producer gas is then passed through a solution of copper(I)chloride under pressure which results in the formation of CuCl(CO).2H₂O. At reduced pressures this solution releases the pure carbon monoxide. Pure carbon monoxide is prepared by warming methanoic acid with concentrated sulphuric acid which acts as a dehydrating agent.

HCOOH + H₂SO₄ → CO + H₂SO₄.H₂O

Properties

• It is a colourless, odourless, and poisonous gas. It is slightly soluble in water.
• It burns in air with a blue flame forming carbon dioxide.

2CO + O₂ → 2CO₂

When carbon monoxide is treated with chlorine in presence of light or charcoal, it forms a poisonous gas carbonyl chloride, which is also known as phosgene. It is used in the synthesis of isocyanates.

CO + Cl₂ → COCl₂

Carbon monoxide acts as a strong reducing agent.

3CO + Fe₂ O₃ → 2Fe + 3CO₂

Under high temperature and pressure a mixture of carbon monoxide and hydrogen (synthetic gas or syn gas) gives methanol.

CO + 2H₂ → CH₃OH

In oxo process, ethene is mixed with carbon monoxide and hydrogen gas to produce propanal.

CO + C₂H₄ + H₄ → CH₃CH₂CHO

Fischer Tropsch Synthesis:

The reaction of carbon monoxide with hydrogen at a pressure of less than 50 atm using metal catalysts at 500 – 700 K yields saturated and unsaturated hydrocarbons.

nCO + (2n+1)H₂ → C_nH(_2n+2) + nH₂O
nCO + 2nH₂ → C_nH_2n + nH₂O

Carbon monoxide forms numerous complex compounds with transition metals in which the transition meal is in zero oxidation state. These compounds are obtained by heating the metal with carbon monoxide.

Eg. Nickel tetracarbonyl [Ni(CO)₄], Iron pentacarbonyl [Fe(CO)₅], Chromium hexacarbonyl [Cr(CO)₆].

Structure:

It has a linear structure. In carbon monoxide, three electron pairs are shared between carbon and oxygen. The bonding can be explained using molecular orbital theory as discussed in XI standard. The C-O bond distance is 1.128Å. The structure can be considered as the resonance hybrid of the following two canonical forms.

Uses of Carbon Monoxide:

1. Equimolar mixture of hydrogen and carbon monoxide – water gas and the mixture of carbon monoxide and nitrogen – producer gas are important industrial fuels.
2. Carbon monoxide is a good reducing agent and can reduce many metal oxides to metals.
3. Carbon monoixde is an important ligand and forms carbonyl compound with transition metals.

Carbon Dioxide:

Carbon dioxide occurs in nature in free state as well as in the combined state. It is a constituent of air (0.03%). It occurs in rock as calcium carbonate and magnesium carbonate.

Production

On industrial scale it is produced by burning coke in excess of air.
C + O₂ → CO₂ ∆H = – 394 kJ mol^-1

Calcination of lime produces carbon dioxide as by product.
CaCO₃ → CaO + CO₂

Carbon dioxide is prepared in laboratory by the action of dilute hydrochloric acid on metal carbonates.
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

Properties

It is a colourless, nonflammable gas and is heavier than air. Its critical temperature is 31^° C and can be readily liquefied.

Carbon dioxide is a very stable compound. Even at 3100 K only 76% decomposes to form carbon monoxide and oxygen. At still higher temperature it decomposes into carbon and oxygen.

Oxidising Behaviour:

At elevated temperatures, it acts as a strong oxidising agent. For example,
CO₂ + 2Mg → 2MgO + C

Water Gas Equilibrium:

The equilibrium involved in the reaction between carbon dioxide and hydrogen, has many industrial applications and is called water gas equilibrium.

Acidic Behaviour:

The aqueous solution of carbon dioxide is slightly acidic as it forms carbonic acid.

CO₂ + H₂O ⇄ H₂CO₃ ⇄ H⁺ + HCO₃^–

Structure of Carbon Dioxide

Carbon dioxide has a linear structure with equal bond distance for the both C-O bonds. In this molecule there is two C-O sigma bond. In addition there is 3c-4e bond covering all the three atoms.

Uses of Carbon Dioxide

1. Carbon dioxide is used to produce an inert atomosphere for chemical processing.
2. Biologically, it is important for photosynthesis.
3. It is also used as fie extinguisher and as a propellent gas.
4. It is used in the production of carbonated beverages and in the production of foam.

Silicon Tetrachloride:

Preparation:

Silicon tetrachloride can be prepared by passing dry chlorine over an intimate mixture of silica and carbon by heating to 1675 K in a porcelain tube

SiO₂ + 2C + 2Cl₂ → SiCl₄ + 2CO

On commercial scale, reaction of silicon with hydrogen chloride gas occurs above 600 K

Si + 4HCl → SiCl₄ + 2H₂

Properties:

Silicon tetrachloride is a colourless fuming liquid and it freezes at – 70°C. In moist air, silicon tetrachloride is hydrolysed with water to give silica and hydrochloric acid.

SiCl₄ + 2H₂O → 4HCl + SiO₂

When silicon tetrachloride is hydrolysed with moist ether, linear perchloro siloxanes are formed [Cl-(Si Cl₂O)_nSiCl₃ where n = 1 – 6.

Alcoholysis

The chloride ion in silicon tetrachloride can be substituted by nucleophile such as OH, OR, etc using suitable reagents. For example, it forms silicic esters with alcohols.

Ammonialysis

Similarly silicon tetrachloride undergoes ammonialysis to form chlorosilazanes.

Uses:

1. Silicon tetrachloride is used in the production of semiconducting silicon.
2. It is used as a starting material in the synthesis of silica gel, silicic esters, a binder for ceramic materials.

Silcones:

Silicones or poly siloxanes are organo silicon polymers with general empirical formula (R₂SiO). Since their empirical formula is similar to that of ketone (R₂CO), they were named “silicones”. These silicones may be linear or cross linked. Because of their very high thermal stability they are called high – temperature polymers.

Preparation:

Generally silicones are prepared by the hydrolysis of dialkyldichlorosilanes (R₂SiCl₂) or diaryldichlorosilanes
Ar₂SiCl₂, which are prepared by passing vapours of RCl or ArCl over silicon at 570 K with copper as a catalyst.

The hydrolysis of dialkylchloro silanes R₂SiCl₂ yields to a straight chain polymer which grown from both the sides

The hydrolysis of monoalkylchloro silanes RSiCl₃ yields to a very complex cross linked polymer. Linear silicones can be converted into cyclic or ring silicones when water molecules is removed from the terminal – OH groups.

Types of Silicones:

(i) Liner Silicones:

They are obtained by the hydrolysis and subsequent condensation of dialkyl or diaryl silicon chlorides.

(a) Silicone Rubbers:

These silicones are bridged together by methylene or similar groups.

(b) Silicone Resins:

They are obtained by blending silicones with organic resins such as acrylic esters.

(ii) Cyclic Silicones

These are obtained by the hydrolysis of R₂SiCl₂.

(iii) Cross linked Silicones

They are obtained by hydrolysis of RSiCl₃

Properties

The extent of cross linking and nature of alkyl group determine the nature of polymer. They range from oily liquids to rubber like solids. All silicones are water repellent. This property arises due to the presence of organic side groups that surrounds the silicon which makes the molecule looks like an alkane.

They are also thermal and electrical insulators. Chemically they are inert. Lower silicones are oily liquids whereas higher silicones with long chain structure are waxy solids. The viscosity of silicon oil remains constant and doesn’t change with temperature and they don’t thicken during winter.

Uses:

1. Silicones are used for low temperature lubrication and in vacuum pumps, high temperature oil baths etc
2. They are used for making water proofing clothes
3. They are used as insulting material in electrical motor and other appliances
4. They are mixed with paints and enamels to make them resistant towards high temperature, sunlight, dampness and chemicals.

Silicates

The mineral which contains silicon and oxygen in tetrahedral [SiO₄]^2- units linked together in different patterns are called silicates. Nearly 95% of the earth crust is composed of silicate minerals and silica. The glass and ceramic industries are based on the chemistry silicates.

Types of Silicates:

Silicates are classified into various types based on the way in which the tetrahedral units, [SiO₄]^4- are
linked together.

Ortho Silicates (Neso Silicates):

The simplest silicates which contain discrete [SiO₄]^4- tetrahedral units are called ortho silicates or
neso silicates.

Examples:

Phenacite – Be₂SiO₄ (Be²⁺ ions are tetrahedrally surrounded by O^2- ions), Olivine – (Fe/Mg)₂SiO₄ (Fe²⁺ and Mg²⁺ cations are octahedrally surrounded by O^2- ions).

Pyro Silicate (or) Soro Silicates:

Silicates which contain [Si₂O₇]^6- ions are called pyro silicates (or) Soro silicates. They are formed by joining two [SiO₄]^4- tetrahedral units by sharing one oxygen atom at one corner (one oxygen is removed while joining). Example: Thortveitite – Sc₂Si₂O₇

Cyclic Silicates (or Ring Silicates)

Silicates which contain (SiO₃)_n^2n- which are formed by linking three or more tetrahedral SiO₄^2- units cyclically are called cyclic silicates. Each silicate unit shares two of its oxygen atoms with other units.

Example: Beryl [Be₃Al₂ (SiO₃)₆] (an aluminosilicate with each aluminium is surrounded by 6 oxygen atoms octahedrally)

Inosilicates:

Silicates which contain ‘n’ number of silicate units linked by sharing two or more oxygen atoms are called inosilicates. They are further classified as chain silicates and double chain silicates.

Chain Silicates (or Pyroxenes):

These silicates contain [(SiO₃)_n]^2n- ions formed by linking ‘n’ number of tetrahedral [SiO₄]^4- units linearly. Each silicate unit shares two of its oxygen atoms with other units. Example: Spodumene – LiAl(SiO₃)₂.

Double Chain Silicates (or Amphiboles):

These silicates contains [Si₄O₁₁]_n^6n- ions. In these silicates there are two different types of tetrahedra:

• These sharing 3 vertices
• Those sharing only 2 vertices.

Examples:

1. Asbestos:

These are fibrous and noncombustible silicates. Therefore they are used for thermal insulation material, brake linings, construction material and filters. Asbestos being carcinogenic silicates, their applications are restricted.

Sheet or Phyllo Silicates

Silicates which contain (Si₂O₅)_n^2n- are called sheet or phyllo silicates. In these, Each [SiO₄]^4- tetrahedron unit shares three oxygen atoms with others and thus by forming twodimensional sheets. These sheets silicates form layered structures in which silicate sheets are stacked over each other. The attractive forces between these layers are very weak, hence they can be cleaved easily just like graphite.

Example: Talc, Mica etc.

Three Dimensional Silicates (or Tecto Silicates):

Silicates in which all the oxygen atoms of [SiO₄]^2- tetrahedra are shared with other tetrahedra to form three-dimensional network are called three dimensional or tecto silicates. They have general formula (SiO₂)_n.

Examples: Quartz

These tecto silicates can be converted into three dimensional aluminosilicates by replacing [SiO₄]^4- units
by [AlO₄]^5- units. Eg. Feldspar, Zeolites etc.,

Zeolites:

Zeolites are three-dimensional crystalline solids containing aluminium, silicon, and oxygen in their regular three dimensional framework. They are hydrated sodium alumino silicates with general formula Na₂O.(Al₂O₃).x(SiO₂).yH₂O (x=2 to 10; y=2 to 6).

Zeolites have porous structure in which the monovalent sodium ions and water molecules are loosely held. The Si and Al atoms are tetrahedrally coordinated with each other through shared oxygen atoms. Zeolites are similar to clay minerals but they differ in their crystalline structure.

Zeolites have a three dimensional crystalline structure looks like a honeycomb consisting of a network of interconnected tunnels and cages. Water molecules moves freely in and out of these pores but the zeolite framework remains rigid.

Another special aspect of this structure is that the pore/channel sizes are nearly uniform, allowing the crystal to act as a molecular sieve. We have already discussed in XI standard, the removal of permanent hardness of water using zeolites.

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Group 13 (Boron Group) Elements

The boron occurs mostly as borates and its important ores are borax – Na₂[B₄O₅(OH)₄].8H₂O and kernite – Na₂[B₄O₅(OH)₄].2H₂O. Aluminium is the most abundant metal and occurs as oxides and also found in aluminosilicate rocks. Commercially it is extracted from its chief ore, bauxite (Al₂O₃.2H₂O). The other elements of this group occur only in trace amounts. The other elements Ga, In and Tl occur as their sulphides.

Physical Properties:

Some of the physical properties of the group 13 elements are listed below

Table 2.3 Physical properties of group 13 elements

Element	Most Common Allotropes
Boron	Amorphous boron, α-rhombohedral boron, β-rhombohedral boron, γ-orthorhombic boron, α-tetragonal boron, β-tetragonal boron
Carbon	Diamond, Graphite, Graphene, Fullerenes, Carbon nanotubes
Silicon	Amorphous Silicon, Crystalline Silicon
Germanium	α-germanium, β-germanium
Tin	Grey tin, white tin, rhombic tin, sigma tin
Phosphorus	White phosphorus, Red phosphorus, Scarlet phosphorus, Violet phosphorus, Black phosphorus.
Arsenic	Yellow arsenic, gray arsenic & Black arsenic
Anitimony	Blue-white antimony, Yellow, Black
Oxygen	Dioxygen, ozone
Sulphur	Rhombus sulphur, monoclinic sulphur
Selenium	Red selenium, Gray selenium, Black selenium, Monoclinic selenium
Tellurium	Amorphous & Crystalline

Chemical Properties of Boron:

Boron is the only nonmetal in this group and is less reactive. However, it shows reactivity at higher temperatures. Many of its compounds are electron deficient and has unusual type of covalent bonding which is due to its small size, high ionisation energy and similarity in electronegativity with carbon and hydrogen.

Formation of Metal Borides:

Many metals except alkali metals form borides with a general formula M_xB_y (x ranging upto 11 and y ranging upto 66 or higher) Direct combination of metals with boron:

Direct Combination of Metals With Boron:

Reduction of Boron Trihalides:

Reduction of borontrichloride with a metal assisted by dihydrogen gives metal borides.

Formation of Hydrides:

Boron does not react directly with hydrogen. However, it forms a variety of hydrides called boranes. The simplest borane is diborane – B₂H₆. Other larger boranes can be prepared from diborane. Treatment of gaseous boron triflouride with sodium hydride around 450 K gives diborane. To prevent subsequent pyrolysis, the product diborane is trapped immediately.

Formation of Boron Trihalides:

Boron combines with halogen to form boron trihalides at high temperatures.

Formation of Boron Nitride:

Boron burns with dinitrogen at high temperatures to form boron nitride.

Formation of Oxides:

When boron is heated with oxygen around 900 K, it forms its oxide.

Reaction with Acids and Alkali:

Halo acids have no reaction with boron. However, boron reacts with oxidising acids such as sulphuric acid and nitric acids and forms boric acid.

2B + 3H₂SO₄ → 2H₃BO₃ + 3SO₂
B + 3HNO₃ → H₃BO₃ + 3NO₂

Boron reacts with fused sodium hydroxide and forms sodium borate.

2B + 6NaOH → 2Na₂BO₃ + 3H₂

Uses of Boron:

1. Boron has the capacity to absorb neutrons. Hence, its isotope ¹⁰B₅ is used as moderator in nuclear reactors.
2. Amorphous boron is used as a rocket fuel igniter.
3. Boron is essential for the cell walls of plants.
4. Compounds of boron have many applications. For example eye drops, antiseptics, washing powders etc contains boric acid and borax. In the manufacture of Pyrex glass, boric oxide is used.

Borax [Na₂B₄O₇.10H₂O]:

Preparation:

Borax is a sodium salt of tetraboric acid. It is obtained from colemanite ore by boiling its solution with sodium carbonate.

Borax is normally formulated as Na₂B₄O₇.10H₂O. But it contains, tetranuclear units [B₄O₅. (OH)₄]^2-. This form is known as prismatic form. Borax also exists two other forms namely, jeweller or octahderal borax (Na₂B₄O₇.5H₂O) and borax glass (Na₂B₄O₇).

Properties

Borax is basic in nature and its solution in hot-water is alkaline as it dissociates into boric acid and sodium hydroxide.

Na₂B₄O₇ + 7H₂O → 4H₃BO₃ + 2NaOH

On heating it forms a transparent borax beads.

Borax reacts with acids to form sparingly soluble boric acid.

Na₂B₄O₇ + 2HCl + 5H₂O → 4H₃BO₃ + 2NaCl
Na₂B₄O₇ + H₂SO₄ + 5H₂O → 4H₃BO₃ + Na₂SO₄

When treated with ammonium chloride it forms boron nitride.

Na₂B₄O₇ + 2NH₄Cl → 2NaCl + 2BN + B₂O₃ + 4H₂O

Uses of Borax:

1. Borax is used for the identifiation of coloured metal ions
2. In the manufacture optical and borosilicate glass, enamels and glazes for pottery
3. It is also used as a flx in metallurgy and also acts as a preservative

Boric Acid [H₃BO₃ or B(OH)₃]:

Preparation:

Boric acid can be extracted from borax and colemanite.

Na₂B₄O₇ + H₂SO₄ → 4H₃BO₃ + Na₂SO₄
Ca₂B₆O₁₁ + 11H₂O + 4SO₂ → 2Ca(HSO₃)₂ + 6H₃BO₃

Properties:

Boric acid is a colourless transparent crystal. It is a very weak monobasic acid and, it accepts hydroxyl ion rather than donating proton.

B(OH)₃ + 2H₂O ⇄ H₃O⁺ + [B(OH)₄]^–

It reacts with sodium hydroxide to form sodium metaborate and sodium tetraborate.

H₃BO₃ + NaOH → NaBO₂ + 2H₂O
4H₃BO₃ + 2NaOH → Na₂B₄O₇ + 7H₂O

Action of Heat:

Boric acid when heated at 373 K gives metaboric acid and at 413 K, it gives tetraboric acid. When heated at red hot, it gives boric anhydride which is a glassy mass.

Action of Ammonia

Fusion of urea with B(OH)₃, in an atmosphere of ammonia at 800 – 1200 K gives boron nitride.

Ethyl Borate Test

When boric acid or borate salt is heated with ethyl alcohol in presence of conc. sulphuric acid, an ester, triethylborate is formed. The vapour of this ester burns with a green edged flame and this reaction is used to identify the presence of borate.

Note:

The trialkyl borate on reaction with sodium hydride in tetrahydrofuran to form a coordination compound Na[BH(OR)₃], which acts as a powerful reducing agent.

Formation of Boron Triflouride:

Boric acid reacts with calcium flouride in presence of conc. sulphuric acid and gives boron triflouride.

3CaF₂ + 3H₂SO₄ + 2 B(OH₃) → 3CaSO₄ + 2BF₃ + 6H₂O

Boric acid when heated with soda ash it gives borax

Na₂CO₃ + 4B(OH)₃ → Na₂B₄O₇ + CO₂ + 6H₂O

Structure of Boric Acid:

Boric acid has a two dimensional layered structure. It consists of [BO₃]^3- unit and these are linked to each other by hydrogen bonds as shown in the Figure 2.2.

Uses of Boric Acid:

1. Boric acid is used in the manufacture of pottery glases, enamels and pigments.
2. It is used as an antiseptic and as an eye lotion.
3. It is also used as a food preservative.

Diborane

Preparation:

As discussed earlier diborane can be prepared by the action of metal hydride with boron. This method is used for the industrial production. Diborane can also be obtained in small quantities by the reaction of iodine with sodium borohydride in diglyme.

2NaBH₄ + I₂ → B₂H₆ + 2NaI + H₂

On heating magnesium boride with HCl a mixture of volatile boranes are obtained.

2Mg₃B₂ + 12HCl → 6MgCl₂ + B₄H₁₀ + H₂
B₄H₁₀ + H₂ → 2B₂H₆

Properties:

Boranes are colourless diamagnetic compounds with low thermal stability. Diborane is a gas at room temperature with sweet smell and it is extremely toxic. It is also highly reactive.

At high temperatures it forms higher boranes liberating hydrogen.

Diboranes reacts with water and alkali to give boric acid and metaborates respectively.

B₂H₆ + 6H₂O → 2H₃BO₃ + 6H₂
B₂H₆ + 2NaOH + 2H₂O → 2NaBO₂ + 6H₂

Action of Air:

At room temperature pure diborane does not react with air or oxygen but in impure form it gives B₂O₃ along
with large amount of heat.

B₂H₆ + 3O₂ → B₂O₃ + 3H₂O
∆H = – 2165 KJ mol^2-

Diborane reacts with methyl alcohol to give trimethyl Borate.

B₂H₆ + 6CH₃OH → 2B(OCH₃)₃ + 6H₂

Hydroboration:

Diborane adds on to alkenes and alkynes in ether solvent at room temperature. This reaction is called hydroboration and is highly used in synthetic organic chemistry, especially for anti Markovnikov addition.

B₂H₆ + 6RCH = CHR → 2(RCH₂ – CHR)₃B

Reaction with Ionic Hydrides

When treated with metal hydrides it forms metal borohydrides

Reaction with Ammonia:

When treated with excess ammonia at low temperatures diborane gives diboranediammonate. On heating at higher temperatures it gives borazole.

Structure of Diborane:

In diborane two BH₂ units are linked by two bridged hydrogens. Therefore, it has eight B-H bonds. However, diborane has only 12 valance electrons and are not sufficient to form normal covalent bonds. The four terminal B-H bonds are normal covalent bonds (two centre – two electron bond or 2c-2e bond).

The remaining four electrons have to be used for the bridged bonds. i.e. two three centred B-H-B bonds utilise two electrons each. Hence, these bonds are three centre-two electron bonds (3c-2e). The bridging hydrogen atoms are in a plane as shown in the figure 2.3. In diborane, the boron is sp³ hybridised.

Three of the four sp³ hybridised orbitals contains single electron and the fourth orbital is empty. Two of the half filled hybridised orbitals of each boron overlap with the 1s orbitals of two hydrogens to form four terminal 2c-2e bonds, leaving one empty and one half filled hybridised orbitals on each boron. The Three centre – two electron bonds), B-H-B bond formation involves overlapping the half filled hybridised orbital of one boron, the empty hybridised orbital of the other boron and the half filled 1s orbital of hydrogen.

Uses of Diborane:

1. Diborane is used as a high energy fuel for propellant
2. It is used as a reducing agent in organic chemistry
3. It is used in welding torches

Boron Triflouride:

Preparation:

Boron triflouride is obtained by the treatment of calcium fluoride with boron trioxide in presence of conc. sulphuric acid.

It can also be obtained by treating boron trioxide with carbon and fluorine.

B₂O₃ + 3C + 3F₂ → 2BF₃ + 3CO

In the laboratory pure BF₃ is prepared by the thermal decomposition of benzene diazonium tetrafloro borate.

Properties:

Boron triflouride has a planar geometry. It is an electron deficient compound and accepts electron pairs to form coordinate covalent bonds. They form complex of the type [BX₄]^2-.

BF₃ + NH₃ → F₃B ← NH₃
BF₃ + H₂O → F₃B ← OH₃

On hydrolysis, boric acid is obtained. This then gets converted into Hydro floroboric acid.

Uses of Boron Triflouride:

1. Boron trifloride is used for preparing HBF₄, a catalyst in organic chemistry
2. It is also used as a flourinating reagent.

Aluminium Chloride:

Preparation:

When aluminium metal or aluminium hydroxide is treated with hydrochloric acid, aluminium trichloride is formed. The reaction mixture is evaporated to obtain hydrated aluminium chloride.

2Al + 6HCl → 2AlCl₃ + 3H₂
Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O

McAfee Process:

Aluminium chloride is obtained by heating a mixture of alumina and coke in a current of chlorine.

2Al₂O₃ + 3C + 6Cl₂ → 4AlCl₃ + 3CO₂

On industrial scale it is prepared by chlorinating aluminium around 1000 K

Properties:

Anhydrous aluminium chloride is a colourless, hygroscopic substance. An aqueous solution of aluminium chloride is acidic in nature. It also produces hydrogen chloride fumes in moist air.

AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl

With ammonium hydroxide it forms aluminium hydroxide.

AlCl₃ + 3NH₄OH → Al(OH)₃ + 3NH₄Cl

With excess of sodium hydroxide it produces metal aluminate

AlCl₃ + 4NaOH → NaAlO₂ + 2H₂O + 3NaCl

It behaves like a Lewis acid and forms addition compounds with ammonia, phosphine and carbonylchloride etc… Eg. AlCl₃.6NH₃.

Uses of Aluminium Chloride:

1. Anhydrous aluminium chloride is used as a catalyst in Friedels Craft reactions
2. It is used for the manufacture of petrol by cracking the mineral oils.
3. It is used as a catalyst in the manufacture on dyes, drugs and perfumes.

Alums:

The name alum is given to the double salt of potassium aluminium sulphate [K₂SO₄.Al₂(SO₄)₃.24.H₂O]. Now a days it is used for all the double salts with M^‘SO₄.M^”₂(SO₄)₃. 24H₂O, where M’ is univalent metal ion or [NH₄]⁺ and M^” is trivalent metal ion.

Examples:

Potash alum [K₂SO₄.Al₂(SO₄)₃.24.H₂O]; Sodium alum [Na₂SO₄.Al₂(SO₄)₃. 24.H₂O], Ammonium alum [(NH₄)₂(SO₄)₃.24.H₂O], Chrome alum [K₂SO₄.Cr₂(SO₄)₃.24.H₂O]. Alums in general are more soluble in hot water than in cold water and in solutions they exhibit the properties of constituent ions.

Preparation:

The alunite the alum stone is the naturally occurring form and it is K₂SO₄. Al₂(SO₄)₃. 4Al(OH)₃. When alum stone is treated with excess of sulphuric acid, the aluminium hydroxide is converted to aluminium sulphate. A calculated quantity of potassium sulphate is added and the solution is crystallised to generate potash alum. It is purified by recrystallisation.

K₂SO₄.Al₂(SO₄)₃.4Al(OH)₃ + 6H₂SO₄ → K₂SO₄ + 3Al₂(SO₄)₃ + 12H₂O
K₂SO₄ + Al₂(SO₄)₃ + 24 H₂O → K₂SO₄.Al₂(SO₄)₃. 24 H₂O

Properties

Potash alum is a white crystalline solid it is soluble in water and insoluble in alcohol. The aqueous solution is acidic due to the hydrolysis of aluminium sulphate it melts at 365 K on heating. At 475 K loses water of hydration and swells up. The swollen mass is known as burnt alum. Heating to red hot it decomposes into potassium sulphate, alumina and sulphur trioxide.

Potash alum forms aluminium hydroxide when treated with ammonium hydroxide.

K₂SO₄.Al₂(SO₄)₃.24 H₂O + 6NH₄OH → K₂SO₄ + 3(NH₄)₂SO₄ + 24 H₂O + 2Al(OH)₃

Uses of Alum:

1. It is used for purification of water
2. It is also used for water proofing and textiles
3. It is used in dyeing, paper and leather tanning industries
4. It is employed as a styptic agent to arrest bleeding.

Find free online chemistry topics covering a broad range of concepts from research institutes around the world.

General Trends In Properties Of P -Block Elements

We already learnt that the properties of elements largely depends on their electronic configuration, size,  ionisation enthalpy, electronegativity etc. Let us discuss the general trend in such properties of various p-block elements.

Electronic Configuration and Oxidation State:

The p-block elements have a general electronic configuration of ns², np^1-6. The elements of each group have similar outer shell electronic configuration and differ only in the value of n (principal quantum number). The elements of group 18 (inert gases) have completely filled p orbitals, hence they are more stable and have least reactivity.

The elements of this block show variable oxidation state and their highest oxidation state (group oxidation state) is equal to the total number of valance electrons present in them. Unlike s-block elements which show only positive oxidation state, some of the p-block elements show negative oxidation states also.

The halogens have a strong tendency to gain an electron to give a stable halide ion with completely filled electronic configuration and hence – 1 oxidation state is more common in halogens. Similarly, the other elements belonging to pnictogen and chalcogen groups also show negative oxidation states.

General Electronic Configurations and Oxidation States of P-Block Elements

The tendency of an element to form a cation by loosing electrons is known as electropositive or metallic character. This character depends on the ionisation energy. Generally on descending a group the ionisation energy decreases and hence the metallic character increases.

Figure 2.1 P-Block Elements With Their Ionisation Enthalpies, Electronegativity and Metallic Nature.

In p-block, the elements present in lower left part are metals while the elements in the upper right part are non metals. Elements of group 13 have metallic character except the first element boron which is a metalloid, having properties intermediate between the metal and nonmetals. The atomic radius of boron is very small and it has relatively high nuclear charge and these properties are responsible for its nonmetallic character.

In the subsequent groups the non-metallic character increases. In group 14 elements, carbon is a nonmetal while silicon and germanium are metalloids. In group 15, nitrogen and phosphorus are non metals and arsenic & antimony are metalloids. In group 16, oxygen, sulphur and selenium are non metals and tellurium is a metalloid. All the elements of group 17 and 18 are non metals.

Ionisation Enthalpy:

We have already learnt that as we move down a group, generally there is a steady decrease in ionisation enthalpy of elements due to increase in their atomic radius. In p-block elements, there are some minor deviations to this general trend. In group 13, from boron to aluminium the ionisation enthalpy decreases as expected. But from aluminium to thallium there is only a marginal difference.

This is due to the presence of inner d and f-electrons which has poor shielding effect compared to s and p-electrons. As a result, the effective nuclear charge on the valance electrons increases. A similar trend is also observed in group 14.

The remaining groups (15 to 18) follow the general trend. In these groups, the ionisation enthalpy decreases, as we move down the group. Here, poor shielding effect of d- and f-electrons are overcome by the increased shielding effect of the additional p-electrons. The ionisation enthalpy of elements in successive groups is higher than the corresponding elements of the previous group as expected.

Electronegativity

As we move down the 13^th group, the electronegativity first decreases from boron to aluminium and then marginally increases for Gallium, there after there is no appreciable change. Similar trend is also observed in 14^th group as well. In other groups, as we move down the group, the electro negativity decreases. This observed trend can be correlated with their atomic radius.

Anomalous Properties of the First Elements:

In p-block elements, the first member of each group differs from the other elements of the corresponding group. The following factors are responsible for this anomalous behaviour.

1. Small size of the first member
2. High ionisation enthalpy and high electronegativity
3. Absence of d orbitals in their valance shell

The first member of the group 13, boron is a metalloid while others are reactive metals. Moreover, boron shows diagonal relationship with silicon of group 14. The oxides of boron and silicon are similar in their acidic nature. Both boron and silicon form covalent hydrides that can be easily hydrolysed. Similarly, except boron trifluoride, halides of both elements are readily hydrolysed.

In group 14, the first element carbon is strictly a nonmetal while other elements are metalloids (silicon & germanium) or metals (tin & lead). Unlike other elements of the group carbon can form multiple bonds such as C=C, C=O etc. Carbon has a greater tendency to form a chain of bonds with itself or with other atoms which is known as catenation. There is considerable decrease in catenation property down the group (C>>Si>Ge≈Sn>Pb).

In group 15 also the first element nitrogen differs from the rest of the elements of the group. Like carbon, the nitrogen can from multiple bonds (N=N, C=N, N=O etc…). Nitrogen is a diatomic gas unlike the other members of the group. Similarly in group 16, the first element, oxygen also exists as a diatomic gas in that group. Due to its high electronegativity it forms hydrogen bonds.

The first element of group 17, flourine the most electronegative element, also behaves quiet differently compared to the rest of the members of group. Like oxygen it also forms hydrogen bonds. It shows only -1 oxidation state while the other halogens have +1, +3, +5 and +7 oxidation states in addition to -1 state. The flourine is the strongest oxidising agent and the most reactive element among the halogens.

Inert Pair Effect:

We have already learnt that the alkali and alkaline earth metals have an oxidation state of +1 and +2 respectively, corresponding to the total number of electrons present in them. Similarly, the elements of p block also show the oxidation states corresponding to the maximum number of valence electrons (group oxidation state).

In addition they also show variable oxidation state. In case of the heavier post-transition elements belonging to the groups (13 to 16), the most stable oxidation state is two less than the group oxidation state and there is a reluctance to exhibit the group oxidation state.

Let us consider group 13 elements. As we move from boron to heavier elements, there is an increasing tendency to have +1 oxidation state, rather than the group oxidation state, +3. For example Al⁺³ is more stable than Al⁺¹ while Tl⁺¹ is more stable than Tl⁺³. Aluminium (III) chloride is stable whereas thallium (III) chloride is highly unstable and disproportionates to thallium (I) chloride and chlorine gas.

This shows that in thallium the stable lower oxidation state corresponds to the loss of np electrons only and not ns electrons. Thus in heavier posttransition metals, the outer s electrons (ns) have a tendency to remain inert and show reluctance to take part in the bonding, which is known as inert pair effect. This effect is also observed in groups 14, 15 and 16.

Allotropism in P-Block Elements:

Some elements exist in more than one crystalline or molecular forms in the same physical state. For example, carbon exists as diamond and graphite. This phenomenon is called allotropism (in greek ‘allos’ means another and ‘trope’ means change) and the different forms of an element are called allotropes. Many p-block elements show allotropism and some of the common allotropes are listed in the table.

Table 2.2 : Some of common allotropes of p-block elements

Element	Most Common Allotropes
Boron	Amorphous boron, α-rhombohedral boron, β-rhombohedral boron, γ-orthorhombic boron, α-tetragonal boron, β-tetragonal boron
Carbon	Diamond, Graphite, Graphene, Fullerenes, Carbon nanotubes
Silicon	Amorphous Silicon, Crystalline Silicon
Germanium	α-germanium, β-germanium
Tin	Grey tin, white tin, rhombic tin, sigma tin
Phosphorus	White phosphorus, Red phosphorus, Scarlet phosphorus, Violet phosphorus, Black phosphorus.
Arsenic	Yellow arsenic, gray arsenic & Black arsenic
Anitimony	Blue-white antimony, Yellow, Black
Oxygen	Dioxygen, ozone
Sulphur	Rhombus sulphur, monoclinic sulphur
Selenium	Red selenium, Gray selenium, Black selenium, Monoclinic selenium
Tellurium	Amorphous & Crystalline

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Applications Of Metals

Applications of Al

Aluminium is the most abundant metal and is a good conductor of electricity and heat. It also resists corrosion. The following are some of its applications.

Many heat exchangers/sinks and our day to day cooking vessels are made of aluminium.

It is used as wraps (aluminium foils) and is used in packing materials for food items.

Aluminium is not very strong, However , its alloys with copper, manganese, magnesium and silicon are light weight and strong and they are used in design of aeroplanes and other forms of transport.

As Aluminium shows high resistance to corrosion, it is used in the design of chemical reactors, medical equipments, refrigeration units and gas pipelines.

Aluminium is a good electrical conductor and cheap, hence used in electrical overhead electric cables with steel core for strength.

Applications of Zn

Metallic zinc is used in galvanising metals such as iron and steel structures to protect them from rusting and corrosion. Zinc is also used to produce die-castings in the automobile, electrical and hardware industries.

Zinc oxide is used in the manufacture of many products such as paints, rubber, cosmetics, pharmaceuticals, plastics, inks, batteries, textiles and electrical equipment. Zinc sulphide is used in making luminous paints, florescent lights and x-ray screens.

Brass an alloy of zinc is used in water valves and communication equipment as it is highly resistant to corrosion.

Applications of Fe

Iron is one of the most useful metals and its alloys are used everywhere including bridges, electricity pylons, bicycle chains, cutting tools and rifle barrels. Cast iron is used to make pipes, valves and pumps stoves etc. Magnets can be made from iron and its alloys and compounds.

An important alloy of iron is stainless steel, and it is very resistant to corrosion. It is used in architecture, bearings, cutlery, surgical instruments and jewellery. Nickel steel is used for making cables, automobiles and aeroplane parts. Chrome steels are used for maufacturing cutting tools and crushing machines.

Applications of Cu

Copper is the first metal used by the human and extended use of its alloy bronze resulted in a new era,’Bronze age’. Copper is used for making coins and ornaments along with gold and other metals. Copper and its alloys are used for making wires, water pipes and other electrical parts.

Applications of Au

Gold, one of the expensive and precious metals. It is used for coinage, and has been used as standard for monetary systems in some countries.

It is used extensively in jewellery in its alloy form with copper. It is also used in electroplating to cover other metals with a thin layer of gold which are used in watches, artificial limb joints, cheap jewellery, dental filings and electrical connectors. Gold nanoparticles are also used for increasing the efficiency of solar cells and also used an catalysts.

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Refining Process

Generally the metal extracted from its ore contains some impurities such as unreacted oxide ore, other metals, nonmetals etc. Removal of such impurities associated with the isolated crude metal is called refiing process. In this section, let us discuss some of the common refining methods.

Distillation

This method is employed for low boiling volatile metals like zinc (boiling point 1180 K) and mercury (630 K). In this method, the impure metal is heated to evaporate and the vapours are condensed to get pure metal.

Liquation

This method, is employed to remove the impurities with high melting points from metals having relatively low melting points such as tin (Sn; mp = 904 K), lead (Pb; mp = 600 K), mercury (Hg; mp = 234 K), and bismuth (Bi; mp = 545 K). In this process, the crude metal is heated to form fusible liquid and allowed to flow on a sloping surface.

The impure metal is placed on sloping hearth of a reverberatory furnace and it is heated just above the melting point of the metal in the absence of air, the molten pure metal flows down and the impurities are left behind. The molten metal is collected and solidified.

Electrolytic Refining:

The crude metal is refined by electrolysis. It is carried out in an electrolytic cell containing aqueous solution of the salts of the metal of interest. The rods of impure metal are used as anode and thin strips of pure metal are used as cathode.

The metal of interest dissolves from the anode, pass into the solution while the same amount of metal ions from the solution will be deposited at the cathode. During electrolysis, the less electropositive impurities in the anode, settle down at the bottom and are removed as anode mud.

Let us understand this process by considering electrolytic refiing of silver as an example.

Cathode: Pure silver
Anode: Impure silver rods
Electrolyte: Acidified aqueous solution of silver nitrate.

When a current is passed through the electrodes the following reactions will take place

Reaction at anode Ag(s) → Ag⁺ (aq) + 1e^–
Reaction at cathode Ag⁺ (aq) + 1e^– → Ag(s)

During electrolysis, at the anode the silver atoms lose electrons and enter the solution. The positively charged silver cations migrate towards the cathode and get discharged by gaining electrons and deposited on the cathode. Other metals such as copper, zinc etc.,can also be refined by this process in a similar manner.

Zone Refining

This method is based on the principles of fractional crystallisation. When an impure metal is melted and allowed to solidify, the impurities will prefer to be in the molten region. i.e. impurities are more soluble in the melt than in the solid state metal.

In this process the impure metal is taken in the form of a rod. One end of the rod is heated using a mobile induction heater which results in melting of the metal on that portion of the rod. When the heater is slowly moved to the other end the pure metal crystallises while the impurities will move on to the adjacent molten zone formed due to the movement of the heater. As the heater moves further away, the molten zone containing impurities also moves along with it.

The process is repeated several times by moving the heater in the same direction again and again to achieve the desired purity level. This process is carried out in an inert gas atmosphere to prevent the oxidation of metals. Elements such as germanium (Ge), silicon (Si) and galium (Ga) that are used as semiconductor are refined using this process.

Vapour Phase Method

In this method, the metal is treated with a suitable reagent which can form a volatile compound with the metal. Then the volatile compound is decomposed to give the pure metal. We can understand this method by considering the following process.

Mond Process for Refining Nickel:

The impure nickel is heated in a stream of carbon monoxide at around 350 K. The nickel reacts with the CO to form a highly volatile nickel tetracarbonyl. The solid impurities are left behind.

Ni (s) + 4 CO (g) → [Ni(CO)₄] (g)

On heating the nickel tetracarbonyl around 460 K, the complex decomposes to give pure metal.

[Ni(CO)₄] (g) → Ni (s) + 4 CO (g)

Van-Arkel Method for Refining Zirconium/Titanium:

This method is based on the thermal decomposition of metal compounds which lead to the formation of pure metals. Titanium and zirconium can be purified using this method. For example, the impure titanium metal is heated in an evacuated vessel with iodine at a temperature of 550 K to form the volatile titanium tetra iodide. (TiI₄). The impurities are left behind, as they do not react with iodine.

The volatile titanium tetraiodide vapour is passed over a tungsten filament at a temperature aroud 1800 K. The titanium tetraiodide is decomposed and pure titanium is deposited on the filament. The iodine is reused.