CBSE Class 11 Chemistry Notes For Some P Block Elements Group-14 Elements

Class 11 Chemistry Some P Block Elements Group-14 Elements (Carbon Family) Introduction

The valence shell electronic configuration of the elements of group 14 is ns²np², where n = 2-6.

It becomes clear from these electronic configurations (given in the table below) that carbon and silicon have noble gas cores, germanium and tin have noble gas plus 10 d-electron cores and lead has a noble gas core in addition to 14 f and 10 d-electron cores.

Thus, the electronic configurations of group-14 elements are similar to that of group-13 elements. However, they contain one more p-electron as compared to group-13 elements.

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Electronic Configurations Of The Group- 14 Elements:

Occurrence Of Group-14 Elements

The members of group 14 are carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb).

Carbon is the seventeenth most abundant element by mass in the earth’s crust It is widely distributed in nature in free as well as in combined states. In a free state, it is found to occur in coal, graphite and diamond. These are the main allotropes of carbon.

• Carbon in the form of coal and coke is used mainly as fuel. In a combined state, it is present widely as metal carbonates, hydrocarbons (petroleum), carbohydrates and carbon dioxide (0.03%) in air. Gases like propane and butane are the major constituents of LPG.
• Moreover, the main constituent of all organic compounds is carbon. Two stable isotopes of carbon are present in nature namely ¹²C₆  and ¹³C₆. Another isotope of carbon (¹⁴C₆) is radioactive. The age of antique articles is determined by the ratio of ¹²C₆  and ¹³C₆. present in them. This process is called radiocarbon dating.
• Silicon is the second (about 27.7% by mass) most abundant element (next to oxygen) in the earth’s crust and is present in the form of silica and silicates. Germanium occurs only in traces (1.5 ppm). Both germanium and silicon in very pure form find applications as semiconductors.
• The natural abundance of tin and lead is very low (2 and 13ppm respectively). The principal ore of tin is tinstone or cassiterite (SnO₂) and that of lead is galena (PbS). Both tin and lead form several alloys. Tin is also used for tin plating while some lead-containing compounds are used as the constituents of paints.
• The first two elements of this group, carbon and silicon are non-metals, germanium is a semi-metal (metalloid) while tin and lead are metals

Allotropic Forms Of Carbon

The phenomenon of the existence of an element in two or more forms having different physical but similar chemical properties i? called allotropy and the different forms are called allotropes. Carbon exists in some allotropic forms which may be classified as

1. Crystalline and
2. Amorphous: ‘

The Four Crystalline Allotropic Forms Of Carbon Are:

• Diamond
• Graphite
• Fullerene
• Carbon nanotubes.

The Four Amorphous Allotropic Forms Of Carbon Are:

1. Charcoal
2. Soot or lamp black,
3. Coke and
4. Gas carbon.

Amorphous carbon is not pure and remains mixed with various elements and compounds. Finer X-ray studies have shown that the amorphous varieties ofc arbon are composed of very minute crystalline units like graphite which are distributed throughout their masses in a most disordered fashion. A synopsis of various is allotropes of carbon is given in the following chart

Crystalline Carbon

1. Diamond

Diamond is a very precious substance which is available in South Africa, New South Wales, Brazil, the Ural mountains and at Golconda in India. Two varieties of natural diamonds are available. One is the lustrous and colourless (or slightly coloured) variety which is generally used as precious gemstones and the other is the black or deep-coloured opaque variety, known as carbonado orbort.

The weight of precious diamond is expressed in carats (1 carat = 200mg). Some well-known diamonds are:

• Cullinan (3032 carats), Kohinoor (present weight 106 carats),
• Pitt (136.25 carats), Regent (193 carats),
• Orloff (193 carats) and Great Mogul (186 carats).
• Diamond can also be prepared artificially but because of high cost and poor quality is seldom made artificially

Diamond Physical Properties

• Diamond is transparent, lustrous and crystalline. It may be colourless or slightly yellow coloured although some black or dark-coloured varieties are also available.
• It is the hardest naturally occurring substance known and it has a very high melting point (3843 K).
• It is the heaviest among all the allotropic forms of carbon; its density is 3.51 gem-3.
• Diamond has a very high refractive index (2.417) and thus light passing through it, suffers total internal reflection innumerable times.
• For this reason, the diamond appears to be extremely bright and lustrous.0It is transparent to X-rays and this property helps to distinguish a real diamond from an artificial one (made of glass).
• Diamond is a non-conductor of electricity but a good conductor of heat

Diamond Chemical Properties:

At ordinary temperatures, diamond is chemically inert. It does not react with acids, alkalies, chlorine, potassium chlorate etc. However, it reacts with certain substances at much higher temperatures.

1. It is oxidised by oxygen at 800-900°C to produce pure carbon dioxide:

C + O₂  → CO₂

2.  Diamond is converted into graphite in the absence of air at much higher temperatures

Molten diamond can be converted into graphite by applying heat but graphite cannot be converted into diamond by heating to a very high temperature. The change is, therefore, unidirectional and this is because graphite is thermodynamically more stable than diamond. This type of allotropy is known as monotropy.

3. When a diamond is reacted with molten sodium carbonate, sodium monoxide and carbon monoxide are produced

C + Na₂CO₃→ Na₂O + 2CO

4. It undergoes oxidation by fluorine at 700°C to form carbon tetrafluoride C+ 2F CF

5. At about 250°C, diamond gets oxidised by a mixture of K₂Cr₂O₇ and concentrated sulphuric acid (i.e., chromic acid) to CO₂

6. Diamond reacts with sulphur vapour at 1000°C to form carbon disulphide: C + 2S→ CS₂

Structure Of Diamond

• In diamond, each C-atom is sp³ – hybridised and linked to four other C-atoms tetrahedrally by covalent bonds.
• The value of each C—C—C bond angle is 109°28′ and each C —C bond distance is 1.54A
• An innumerable number of such tetrahedral units are linked together to form a three-dimensional giant molecule containing very strong bonds extended in all directions. Because of such a three-dimensional network of strong covalent bonds, diamond is extremely hard.
• Since a huge amount of thermal energy is required to break a large number of strong covalent bonds, its melting point is very high.
• All 4 valence electrons of each sp³ -hybridised C-atom in a diamond crystal participate in forming covalent bonds and there is no free electron on any carbon atom. Thus, a diamond is an anon-conductor of electricity.
• Diamond has the highest known thermal conductivity because its structure distributes thermal motion in three dimensions very effectively.
• Unlike graphite in which the C-atoms are arranged in different distant layers, the C-atoms in diamond are placed at a covalent bond distance (1.54A). Because of this, the density of diamond is higher than that of graphite

Diamond Uses:

• Because of its transparency, dazzling lustre and beauty, diamond is extensively used as precious gemstone.
• Because of its extreme hardness,it is used for cutting glass, polishing hard surfaces and drilling purposes. Black or dark-coloured diamonds are generally used for this purpose.

2. Graphite

Graphite is available as minerals in Sri Lanka, Mexico, Italy, California (U.S.A), Siberia, Korea, Spain and India. The word graphite originates from the Greek word ‘graphic’ which

Preparation Of Artificial Graphite:

Acheson process in this process, coke dust mixed with silica is heated to a temperature of 3000-3500°C with the help of electrodes in an electric furnace made offire-bricks for 25-30 hours. The mixture is kept covered by sand.

• In the first stage of the reaction, silica reacts with carbon to form silicon carbide (SiC) and carbon monoxide (CO).
• Silicon carbide thus formed decomposes to yield graphite and silicon. Q At higher temperatures, silicon, on being vapourised, escapes from the furnace and graphite is left.

SiO₂ + 3C→SiC + 2CO↑

SiC→Si + C [graphite]

Graphite Physical Properties

•  Graphite is a dark greyish-coloured opaque, soft and slippery crystalline substance possessing metallic lustre.
• It is lighter (density 2.25 g-cm-3 ) than diamond.
• It is a good conductor of heat and electricity.

Graphite Chemical properties

Graphite is more reactive than diamond.

1. When graphite is heated in air at 700°C, it is oxidised to carbon dioxide:  C + O₂→CO₂

2. At 500°C, fluorine reacts with graphite to produce carbon tetrafluoride (CF₄).  The compound is a non-conductor of electricity.It is also called graphite fluoride.

C+ 2F₂ → CF₄

3. Graphite is not attacked by dilute acid or alkali. However, when it is subjected to react with molten sodium carbonate, carbon monoxide is formed.

4. Graphite is oxidised to carbon dioxide with a mixture of K₂Cr₂O and concentrated H₂SO₄ (chromic acid).

When graphite is heated in the presence of a mixture of cones. nitric acid and sulphuric acid containing a small amount of potassium chlorate, greenish-yellow-coloured solid graphitic acid (CnH₄O₅) is obtained.

Its exact structural formula is not known yet. Graphite can be identified by this test (diamond does not respond to this test).

On complete combustion, graphite produces mellitic acid [C₆(COOH)₆].

Structure Of Graphite

•  Each carbon atom in graphite is sp² -hybridised and is linked to three other carbon atoms directly in the same plane forming a network of planar hexagons and these two-dimensional layers exist in different parallel planes.
• In each layer, the C—C bond length is 1.42A and the distance between two adjacent layers is 3.35A which is greater than the C—C covalent bond distance. So, the layers are supposed to be held together by relatively weak van der Waals forces of attraction
• As the distance between two parallel layers is sufficiently large, graphite is less dense than diamond.
• Since the layers are weakly held together, on the application of pressure, one layer easily slides over the other. Thus, graphite is found to be soft and lubricating in nature.
• In the formation of hexagons in a layer of graphite, only three of its four valence electrons are used to form three sigma bonds (C_sp²-C_sp²). The remaining electrons of each carbon atom present in an unhybridised p -orbital is utilised to form n -n-bonds.
• The -π-electrons are mobile and can move freely through the graphite crystal. Because of the presence of free mobile electrons, graphite is a good conductor of electricity and heat. Of all the crystalline allotropes of carbon, graphite is thermodynamically the most stable one. Its standard enthalpy of formation (AfH°) is taken as zero.

Graphite Uses

• Graphite is largely used for lining and making electrodes for electric furnaces.
• When mixed with oil and water, graphite is used as a lubricant in machinery.
• It finds use in making crucibles resistant to high temperatures.
• By mixing with desired quantities of wax or clay, graphite is used for making cores of lead pencils.
• Graphite is used as a moderator in nuclear reactors

3. Fullerens

Fullerenes or Buckminsterfullerene (named after the famous American designer of the geodesic dome, Robert Buckminster Fuller) is the latest allotrope of carbon discovered in 1985 collectively by three scientists namely R. E. Smalley, R. F. Curl and H. W. Kroto.It is a crystalline allotrope of carbon in which the carbon atoms exist in a cluster form.

It is also known to be the purest form of carbon because unlike diamond and graphite does not have surface bonds that are to be attracted by other atoms.

1. Structure Of Fullerenes C₆₀:

• Fullerenes are expressed by the general formula Cn, where n is an even number between 30-600, for example, C₆₀, C₇₀, C₈₄ ….etc.
• All these are cage-like spheroidal molecules having polyhedral geometry containing pentagonal and hexagonal planes. Number of hexagons in a C_n molecule =(n/2- 10),
• For example,  in fullerene C₆₀ molecule, number of hexagons =(60/2-10) = 20. Structure of fullerene C₆₀; The C₆₀ molecule consists of twenty-six-membered rings and twelve five-membered rings of
• sp² -hybridised carbon atoms fused into each other. Each carbon atom forms three or -bonds with the other three carbon atoms and the remaining electron on each carbon is involved in the formation of n -bond and as a result, the system is expected to be aromatic.
• However, it is not aromatic because the molecule is not planar and it does not have (4n + 2) electrons. It is a non¬ aromatic system. This fusion pattern provides a marvellous symmetry to the structure in which the fused ring system bends around and closes to form a soccer ball-shaped molecule (“buckyball”). Of all the fullerenes, C₆₀ is the most stable one.

2. Structure Of Fullerenes C7₀:

The molecule acquires the shape of a rugby ball. It consists of twelve five-membered rings and twenty-five six-membered rings and their arrangement is the same as that of a C₆₀ molecule.

Fullerens Preparation

• The preparation of fullerenes involves heating of graphite in an electric arc in the presence of an inert gas such as He or Ar.
• A sooty material is recovered which consists mainly of C₆₀ with a small amount of C₇₀ and traces of other with an even number of C atoms up to 350 and above. The C₆₀ and C₇₀ fullerenes can further be separated from the sooty material by extraction with benzene or toluene followed by chromatographic separation using alumina (Al₂O₃) as the adsorbent.
• In Russia, America, Canada and New Zealand C₆₀ and C₇₀ fullerenes are isolated from natural sources. Fullerenes of this type are formed in the red giant star Antares.

Fullerenes Properties And Applications:

• Fullerenes are solids with high melting points.
• Being covalent, they are soluble in organic solvents.
• They react with alkali metals to form solid compounds such as K₃C₆₀.
• This compound acts as a superconductor even at temperatures of the order of 10-40K.
• Because of their spherical shapes, they exhibit wonderful lubricating property

4. Carbon Nanotubes

Carbon nanotubes are crystalline allotropes of carbon with cylindrical nanostructure. This allotrope was discovered by Sumio Iijima (Japan) in 1991. A carbon nanotube consists of a two-dimensional array of hexagonal rings of carbon just as in a layer of graphite or a chicken wire. The layer is then rolled. Cylinder and capped at each end with half of a C₆₀. fullerene

Carbon nanotubes Properties and applications:

• The nanotubes are approximately 50000 times thinner than a human hair. These are very tough, about 100 times as strong as steel. They are electrically conducting along the length of the tube.
• These cylindrical carbon molecules having unusual properties are valuable for nanotechnology, electronics, optics and other fields of material science and technology.
• They are also being used as probe tips for the analysis of DNA and proteins by atomic force microscopy (AFM). Many other applications have been envisioned for them as well, including molecular-size test tubes or capsules for drug delivery