Chemistry Part Ii Chapter 11 The P-Block Elements

Both Aluminium and Gallium  forms amphoteric hydroxides.

Thallium(Tl). This is due to inert pair effect because of which it forms the most stable +1 oxidation state. 

It is an electron deficient compound. 

No, because borax bead test is only performed for coloured salts.

Boron nitride is the hardest compound of boron.

Borazine is called inorganic benzene since it has a structure similar to benzene. Its formula is B₃N₃H₆.

Structure of  ion is triangular planar.

Sodium borohydride (NaBH₄).

No, it does not exist.

The relative stability of +1 oxidation state progressively increases for heavier elements: Al<Ga<In<Tl.

No, there are two types of bonds in diborane, normal covalent bonds and three centred electron pair bond. 

It is ore of boron. It is a crude form of boron Na₂B₄O₇.10H₂O.

The B.........H.........B bridges in diborane because of distorted sp³ hybridization, involving the empty 2p_z orbital,thus resulting in the formation of a tau bond or a banana bond. It is a typical case of dual hybridization in the boron atom.

Boron and silicon

boron does not form B⁺³ because of its small size and the high sum of first three ionisation energies it does not lose its three valence electrons to form B⁺³ ions.

In BF₃ molecules three electrons of Boron and three electrons of fluorine combine to forms the compound. Since boron in its compounds does not have a complete octet (8 electrons).

C.

basic

Basic. Since borax is a salt of strong base (NaOH) and a weak acid (H₃BO₃), therefore it is basic in nature.

B.

the presence of hydrogen bonds

Since boric acid is polymeric due to the presence of hydrogen bonds.

C.

Graphite.

Peat, Lignite, Coke, Anthracite are different varieties of coal. Anthracite has the maximum percentage of carbon (90- 95%)

Among the elements of group 14, C -C bond has the maximum bond strength and hence carbon shows maximum tendency for catenation. 

Due to a smaller size and higher electronegativity, carbon undergoes  overlap to form multiple bonds but silicon does not.

Diamond has a three-dimensional network involving strong C-C bonds, which are very difficult to break and hence it is the high melting point.

Two oxides of carbon:
(i) Carbon monoxide (CO)
(ii) Carbon dioxide (CO₂)

This is due to its combination with haemoglobin in the red blood cells to form carboxyhaemoglobin. As a result, blood cannot absorb oxygen and supply it to the body.

By heating sodium formate with conc. H₂SO_4.

Two uses of carbon monoxide:
(i) It is a reducing agent in the preparation of metals. 
(ii) It helps in the preparation of a number of useful compounds such as metal carbonyls.

Carbon dioxide dissolves in water to form carbonic acid. Thus, it is called carbonic anhydride.

Graphite is used as a moderator in a nuclear reactor.

(i) Washing soda does not decompose on heating. 
(ii)Baking soda decomposes on heating to evolve carbon dioxide. 

B.

graphite

Graphite. Because thermodynamically, the most stable form of carbon is graphite.

B.

(i) Pyrene: CCl₄
(ii) Freon: CF₂Cl₂

(i) Carborundum: SiC
(ii) Dry ice: Solid CO₂.

CCl₂F₂ (Freon-12).

Graphite.

carbon compounds are relatively inert because carbon-carbon bond dissociation energies are quite high. 

Covalent.

(i) Graphite - sp²
(ii) Diamond - sp³

Diamond.

Carbon monoxide.

Silicones are synthetic organosilicon compounds containing repeated R₂SiO units held by Si–O–Si linkages.

Silicones are heat resistant and water repellant which make them useful in electronic and electrical appliances.

Glass and cement.

(a) Occurence of boron and uses:
Boron is not found free in nature. It mainly occurs as bortes and orthoboric acid. The important minerals of boron are:
1. Borax (Tincal) Na₂B₄O₇ . 10H₂O
2. Colemanite Ca₂B₆O₁₁ . 5H₂O
3. Kernite Na₂B₄O₇.2H₂O
4. Boric acid H₃BO₃
In India borax occurs in Puga valley (Ladakh) and Sambhar Lake (Rajasthan). The abundance of boron in the earth crust is less than 0.0001% by mass. There are two isotopic forms of boron ^l0B(19%) and ¹¹B(81 %).
(b) Occurrence of aluminium and uses:
Aluminium does not occur freely in nature, but its compounds (minerals) are widely distributed in nature. It is the third most abundant element (next to oxygen and silicon). Some minerals of aluminium are
1. Oxides.
(i) Corundum (Al₂O₃)
(ii) Diaspore (Al₂O₃H₂O)
(iii) Bauxite (Al₂O₃.2H₂O)

2. Fluoride, Cryolite Na₃AlF₆

3. Silicate.
(i) Feldspar (KA1 Si₃O₈)
(ii) Kaolin or slate (K₂O.2H₂O.3Al₂O₃.6SiO₂)
4. Basic sulphate Alunite or Alum stone K₂SO₄ Al₂(SO₄)₃.4Al(OH)₃
The important ore of aluminium is bauxite. In India aluminium is found as mica in Madhya Pradesh, Karnataka, Orissa and Jammu.

(c) Occurrence of gallium, indium and thallium These are less abundant elements in nature.

(i) Atomic and ionic radii. Atomic and ionic radii of group 13 elements increase from top to bottom in the group. This is due to increase in the number of energy shells in each succeeding element. However, the atomic radius of gallium (Ga) is less than that of aluminium (Al).
(ii) It should be noted that there is a very small increase in the atomic radii of gallium and indium. This is due to inert pair effect i.e. ineffective shielding of the valence shell by the intervening d and f-electrons.

(iii) Ionisation enthalpy. As we move down the group from B to Al, the ionisation enthalpy decreases due to the increase in atomic size and screening effect. However, gallium has higher ionisation enthalpy than aluminium.

(iii) Density, melting point and boiling point. The density of these elements increases as we move from top to bottom in the group. This is due to increase in the atomic mass of the element which outweighs the effect of an increase in atomic size.
The melting points in this group decrease considerably on moving down the group up to gallium and then increase in case of indium and thallium. Boiling points decrease regularly from boron to thallium.

It is due to the poor shielding of the valence electrons of Ga by the inner 3d-electrons. As a result, the effective nuclear charge of Ga is somewhat greater in magnitude than that of Al. Thus, the electrons in gallium experience the greater force of attraction by the nucleus than in aluminium. Hence the atomic size of Ga(135 pm) is slightly less than that of Al(143 pm).

(i) Electronegativity: There is no regular change in electronegativity on moving down the group. Down the group, electronegativity first decreases from B to Al and then increases marginally. This is because of the discrepancies in atomic size of the elements.

(ii) Nature of bonds: According to Fajan’s rule, the smaller the cation, the greater is its tendency to form covalent compounds. Thus, with the increase of M³⁺ ionic radii down the group from B³⁺ to Tl³⁺, the tendency of these ions to form covalent compounds decreases. Thus boron forms mostly covalent compounds while other members form ionic compounds also.

(iii) Oxidation states: The general electronic configuration of group 13 elements [Noble gas] is ns² np¹, so these elements are expected to show a uniform oxidation state of + 3. Boron and aluminium show an oxidation state of +3 only, but gallium, indium and thallium, however, show oxidation states of + 1 and +3.

(i) The electronic configuration of ₅B is 1s² 2s² 2p¹. Its tri valency is explained by promoting the electron from 2s to 2p so that number of unpaired electrons become three. One 2s and two 2p orbitals are hybridised to form a set of three equivalent sp² hybrid orbitals which are triangular planar with a bond angle of 120°.

(ii) This is because they have small atomic sizes and hence higher ionisation energies due to which electrons can not be lost. Thus they also not form trivalent cations and their compounds are generally covalent in nature which is formed by the sharing of valence electrons with the valence electrons of the other atoms. 

Inert pair effect:The inert pair effect represents the reluctance of the valence electrons to take part in the chemical combination due to their penetration in the nucleus of heavy elements.
B and Al do not exhibit inert pair effect due to the absence of d – or f-electrons. As a result, they show an oxidation state of +3 only due to the presence of two electrons in the s– and one electron in the p-orbital of the valence shell.

On the other hand, the elements from Ga to Tl contain only d and f-electrons and hence show oxidation states of +1 and  +3 due to inert pair effect. 
As we move down the group, the stability of +3 oxidation state decreases and that of +1 oxidation state increases. This means that as we move down the group, the tendency of the electrons of the valence shell to participate in bond formation decreases. In other words, ns² electron pair in Ga, In and Tl tends to remain paired. This is called inert pair effect. Because of inert pair effect, only the electron of thallium takes parts in bonding with the atoms of the other elements. Thus, monovalent compounds of thallium are stable.

This is because inert pair effect is maximum in thallium in which poor shielding of the s-electrons of the valence shell (6s²) by the 3d, -4d, -5d and 4f-electron occurs. As a result, only 6p¹ electron participates in bond formation and thus the most stable oxidation state of Tl is +1 and not +3. Therefore, TlCl is stable while TlCl₃ is unstable.
On the other hand, all the three valence electrons (2s²2p¹) of boron take part in the bond formation due to the absence of d–and f-electrons in B(no inert pair effects). Hence B shows an oxidation state of +3 and thus forms BCl₃ easily. Thus BCl₃ is more stable than TlCl₃.

Standard electrode potential values for two half-cell reactions indicate that aluminium has a high tendency to form Al³⁺ (aq) ions, while Tl³⁺ is not only unstable in solution but is also a powerful oxidising agent. Thus Tl⁺ is more stable in solution than Tl³⁺. Since aluminium being able to form +3 ions easily, therefore it is more electropositive than thallium. 

Species in which the central atom either does not have eight electrons in the valence shell or those which have 8 electrons in the valence shell but can expand their covalency beyond 4 due to the presence of d-orbitals are called electron deficient molecules.
BCl₃ is an electron deficient compound because the central boron atom has only six electrons. As a result, it accepts a pair of electrons from NH₃ to form an adduct.

(i) Due to small atomic size boron forms strong covalent bonds with the neighbouring atoms. Thus boron atoms are closely packed in its solid, thus boron has high melting and boiling points. 
(ii) Because of its high affinity for oxygen, it reduces many metallic oxides to metals. 
For example.

Boron has only three electrons in the outermost shell which it can share with other atoms. Hence in their compounds, there are only six electrons present around B-atom i.e. octet is not complete. Hence, boron forms electron deficient compounds.

(i) Metallic character: The metallic nature i.e. electropositive character of elements increases from boron to aluminium (B<Al). It is due to a smaller size and higher ionisation potential of B than Al.
However, the metallic nature (electropositive character) decreases slowly from Gallium to Thallium (Al > Ga > In > Tl). This is due to the fact that the extra d-electrons in the atoms of these elements exert a very little shielding effect on outer electrons. Therefore, these electrons are more firmly held by the nucleus and hence there is a decrease in the metallic (or electropositive) character.
(ii) The tendency to exhibit the inert-pair effect. Inert pair effect means that the two s-electrons of the valence shell of heavier p-block elements form an inert pair and do not participate in bond formation. The tendency to exhibit the inert pair effect increases as we go towards the bottom of the group. For example, aluminium gives Al³⁺ in solution. Gallium, indium and thallium show +1 and +3 states. The stable state of thallium (Tl) is +1. This is due to the inert pair effect as 6s² electrons do not participate in bonding and only 6p¹ electron takes part in bonding.

Boron trichloride is a planar molecule and three covalent bond results due to sp²- hybridisation. 

BCl₃ does not form a dimer. On the other hand, aluminium trichloride exists in the dimeric state (Al₂Cl₆).

In the dimeric state, each aluminium atom accepts a pair of electrons from the chlorine atom of another aluminium chloride molecule and thereby acquires an octet of electrons. In other words AlCl₃ achieves stability by forming a dimer.

The shape of BF₃ is a planar molecule in which the central boron atom is sp² hybridised. A sp² hybridised boron atom has a vacant p-orbital. 

In BF₃, boron is sp² hybridised and therefore BF₃ is a planar molecule. It has a vacant 2p-orbital. F-atom has three lone pairs of electrons. In BF₃ molecule, one 2p-orbital of fluorine atom overlaps sidewise with empty 2p-orbtial of boron to form  back bonding (back donation) in which the lone pair is transferred from F to B as shown.

As a result of this back bonding (or black donation), the B-F bond acquires some double bond character. 
On the other hand in  ion,  boron is sp³ hybridised and therefore  is a tetrahedral molecule. B in  ion does not have vacant p-orbital available to accept the electrons donated by the F atom. Hence  ion, B -F is a purely single bond. Since double bonds are shorter than single bonds, therefore B-F bond length in BF₃ is shorter (130 pm) than B-F bond length (143 pm) in  [BF₄]^–.

B and Cl have different electronegativities and chlorine (E.N. = 3.00) is more electronegative than B(E.N.  = 2.00). As a result B – Cl bond is polar and hence has a finite dipole moment. Now BCl₃ is a planar molecule in which the three B – Cl bonds are inclined at an angle of 120°. Therefore, the resultant of two B – Cl bonds in cancelled by equal and opposite dipole moment of the bond B – Cl bond as shown.

Hence overall dipole moment of BCl₃ is zero.

Aluminium reacts with dilute NaOH and liberates dihydrogen.
  

Alumina dissolves in aqueous NaOH solution to form sodium meta-aluminate.

Aluminium dissolves both in acids and alkalies evolving dihydrogen.
           

(i) Action with boiling water. It decomposes boiling water to liberate hydrogen.
                
(ii) Action with concentrated H₂SO₄. It reacts with hot concentrated sulphuric acid to liberate sulphur dioxide.
               
(iii) Action with dilute nitric acid. Aluminium reacts with dilute nitric acid to form aluminium nitrate and ammonium nitrate.
        
(iv) Action with concentrated HCl. Aluminium dissolves in moderately concentrated hydrochloric acid to form (hydrated) chloride. 
      
(v) Action with NaOH. It dissolves in hot sodium hydroxide solution to form sodium meta-aluminate. 
   

The strong affinity of aluminium for oxygen i.e. its reducing character is shown by the following reactions:
     

When a strip of aluminium is dipped in concentrated nitric acid, it is oxidised to aluminium oxide (Al₂O₃). It forms a protective coating on the surface of the metal and checks further attack by the acid. Therefore, aluminium becomes passive.

2Al (s) + 2NaOH(aq) + 6H₂O(l)
              ↓
2 Na⁺[Al(OH)₄] ^–(aq) + 3H₂(g)
Sodium
tetra hydroxo aluminate(III)

Aluminium is cheply available on a weight-to weight basis,the electrical conductivity of aluminium is twice that of copper. Hence, aluminium wire is used to make transmission cables.

(i) Anhydrous HF, being a covalent compound and is strongly H-bonded, therefore it does not give ions. Hence AlF₃ does not dissolve in HF. On the other hand, NaF is an ionic compound and gives F^– ions and hence AlF₃ dissolves in NaF forming soluble complex
           
(ii) Borax has much higher tendency to form complexes than aluminium because of its smaller size and higher electronegativity. Hence when gaseous BF₃ is bubbled through the resulting solution. AlF₃ gets precipitated. 

(i) All 13 group elements have three electrons such as   in the outermost orbit. 
(ii) All these elements show a group valency of three. The sum of the first three ionisation energies of these elements is very high. Thus the formation of their trivalent ions (M³⁺) would require very high energy. Hence, their compounds when anhydrous are either essentially covalent or contain an appreciable amount of covalent character.

(iii) These elements are not strongly electropositive.
(iv) They readily form tri-covalent compounds, which are usually trigonal planar in shape.
(v) These elements have only three valence electrons, thus even after the formation of three covalent bonds, they possess only six valence electrons i.e. they are electron deficient compounds. For example.

(vi) In their compounds they have only six valency electrons i.e. they require a pair of electrons to complete their octets. Thus their compounds act as Lewis acids or electron pair acceptors. For example,

(vii) Many of these ‘donor-acceptor’ compounds are known: they are sometimes known as ‘adducts.’
(viii) All the metal ions exist in the hydrated state.
(ix) They form complexes, 4-coordinate and tetrahedral for boron [e.g. (BF₄)^–] and often 60-coordinate and octahedral for the other elements [e.g.(AlF₆)^–].

Boron differs from other elements of group 13 due to its:
(i) small size
(ii) high electronegativity and high ionisation energy
(iii) non-availability of d-orbitals.

Main points of difference:
(i) Boron is non-metal while other members of the group are metals.
(ii) Boron is a bad conductor of electricity while other members are good conductors of electricity.
(iii) The melting and boiling points of boron are much higher than those of other members of the group.
(iv) Boron exhibits maximum covalency of four as in [BH₄]^– ion while other members exhibit a covalency of six as in [Al (OH)₆]^3–.
(v) Boron forms a large number of hydrides called boranes while other members do not do so.
(vi) B(OH)₃ i.e. H₃BO₃ is acidic while Al(OH)₃ as well as Ga(OH)₃ is amphoteric; indium and thallium hydroxides are basic.

The diagonal relationship is due to the fact that both these elements have almost similar values of atomic radii. electronegativity and ionisation energies. Some important points of similarities are:
(i) Both boron and silicon exhibit typical properties of non-metals i.e. possess high melting and boiling points and are non-conductors of electricity.
(ii) Both elements are semiconductors.
(iii) Both exist in amorphous and crystalline forms.
(iv) Both boron and silicon do not form cations.
(v) Both form weak acids as H3BO3 and H₄SiO₄.
(vi) Chlorides of both elements get hydrolysed easily.

(vii) Both form metallic compounds known as borides and silicides as:
          

Thallium resembles aluminium:
(i) Al shows a uniform oxidation state of + 3 and Tl also shows +3 oxidation states in its certain compounds such as TlCl₃, Tl₂O₃ etc.
(ii) Both Al and Tl form octahedral complexes [AlF₆]^3– and [TlF₆]^3– respectively.
Thallium resembles group 1 metals:
(i) Thallium (Tl) and group 1 metals show an oxidation state of + 1. Thallium shows an oxidation state of + 1 due to inert pair effect in some of its compounds like Tl₂O, TICl₆ TIClO₄ etc.
(ii) Like group 1 oxides, Tl₂O is strongly basic.

(i) Boron does not directly combine with hydrogen but a number of interesting hydrides are known. Boron hydrides such as B₂H₆ (diborane), B₄H₁₀ (tetraborane), B₅H₉ (pentaborane), B₆H₁₀ (hexaborane), B₁₀H₁₄ (decaborane) are collectively called boranes.
(ii) Preparation of diborane:
(a) By reducing boron trifluoride with lithium hybrid. 

(b) By reducing boron trichloride with hydrogen or lithium aluminium hydride. 
          
Uses of diborane: It is used:
(i) as fuel for supersonic rockets.
(ii) as a reducing agent in organic reactions.
(iii) for preparing LiBH₄, NaBH₄ etc.

(a) Action with water: Diborane is readily hydrolysed by water liberating H₂ gas.

(b) Action with LiH: Diborane reacts with Lithium hydride in diethyl ether forming lithium borohydride.
  

3B₂H₆ +6NH₃ → 2B₃N₃H₆ +12H₂

Preparation of borax:
(i) From tincal: The natural occurring borax (N₂B₄O₇.10H₂O) is called tincal. Tincal contains about 45% of boron. The natural tincal is dissolved in hot water and insoluble impurities are filtered off. The solution is concentrated and cooled when crystals of borax are obtained.
(ii) From colemanite: Te powdered mineral is heated with a slight excess of aqueous sodium carbonate.

The precipitates of calcium carbonate thus formed are removed by filtration. The filtrate (clear solution) is concentrated and cooled when crystals of borax separate out. Crystals of borax are separated and CO₂ gas is passed through the mother liquor when sodium metaborate is also converted into borax. 

Sodium carbonate is again used.
Uses of borax:
It is used:
(i) In the manufacture of enamels, glazes and optical glass.
(ii) In the manufacture of soap and of drying oils.
(iii) As a flux in welding and soldering.
(iv) As an antiseptic.
(v) In the laboratory for borax bead test.

When borax is heated strongly, a transparent glass bead which consists of sodium metaborate (NaBO₂) and boric anhydride is formed.

When concentrated H₂SO₄ is added to a hot solution of X (Borax), white crystals of an acid Z separate out, It shows that Z must be an orthoboric acid.
Chemistry of the reactions:

Boric acid a weakly acid crystalline compound derived from borax used as a mild antiseptic and in the manufacture of heat-resistant  glass and enamels. several boric acids are known, all derived from boron trioxide with varying amount of water. For example,
Orthoboric acid H₃BO₃ or B₂O₃ . 3H₂O
Metaboric acid HBO₂ or B₂O₃ . H₂O
Pyroboric acid H₆B₄O₉ or 2B₂O₃ . 3H₂O
Tetraboric acid H₂B₄O₇ or 2B₂O₃ . H₂O
of these, orthoboric acid is important.

Orthoboric acid is also called boric acid (H₃BO₃) and prepared as follows:
(i) From borax: It may be prepared by adding an excess of sulphuric acid or hydrochloric acid to a hot saturated solution of borax.

On cooling the reaction mixture, crystals of boric acid separate out.
(ii) From colemanite:  The mineral is powdered, mixed with boiling water and sulphur dioxide is passed through the mixture. 
On cooling, the reaction mixture, crystals of boric acid separate out. 
(iii) From colemanite: The mineral is powdered, mixed with boiling water and sulphur dioxide is passed through the mixture. 

On crystallisation, boric acid separates out while calcium sulphite remains in solution. 
Boric acid can be purified by recrystallization.
(iv) From boron halides: When boron halides are hydrolysed, orthoboric acid is formed. 

Boric acid on heating forms a number of products depending upon the temperature. For example. 

Boric acid acts as a weak Lewis acid by accepting a hydroxide ion of water and releasing a proton into the solution. 

Boric acid has a layer structure in which

the planar trigonal 
H-atoms form a covalent bond with one BO₃ unit and a hydrogen bond with another BO₃ unit. 

The elements carbon (C), silicon (Si), germanium(Ge), tin(Sn) and lead (Pb) constitute group 14 of the periodic table. Carbon and silicon are typical non-metals, germanium is a semi-metal with some metallic characteristics while tin and lead are typical metals.
Electronic configuration: There are four electrons in the valence shell of the atoms of carbon family. The general electronic configuration may be expressed as [Noble gas] ns²np².

Occurrence of Group 14 elements:
Occurrence of Carbon and its uses. Carbon is one of the most important element since it is an essential constituent of most of the compounds found in plants and animals. It occurs:
(i) In native state. Carbon occurs in native state as diamond, graphite and coal.
(ii) In combined state.
(a) It occurs in all living organisms, plants and animal body as proteins, carbohydrates, fats and other complex compounds.
(b) Atmosphere. It is present as CO₂ to the extent of 0.03 percent.
(c) Hydrocarbon. It occurs as hydrocarbons in marsh gas, petroleum, paraffin, vaseline, paper, cloth etc.
(d) Minerals. It occurs widely in minerals such as chalk, marble, kanker, lime stone which are different varieties of calcium carbonate ; magnesite, MgCO₃ ; dolomite, MgCO₃. CaCO₃ ; spathic ore of iron, FeCO₃ ; zinc carbonate ZnCO₃, etc.

C - C bonds are very strong in diamond. Thus, Diamond acts as an abrasive. The various layers of graphite can slide over each other so it is slippery and used as a lubricant. 

(i) Atomic radius: Atomic radius of these elements show a regular increase from carbon to lead. This is due to the addition of an extra shell in each succeeding element. However, the increase in the atomic radii from Si onwards is small. This is due to the ineffective shielding of valence electrons by the intervening d and f-orbitals.
(ii) Ionisation enthalpy: The ionisation enthalpy of these elements decrease as we move from carbon to lead. This is due to the increase in atomic size because of which nuclear pull on the outermost electrons decreases down the group.
However, the first ionisation enthalpy of these elements are higher than the corresponding elements of group 13. This is due to the higher nuclear charge and smaller size of atoms of group 14 elements.

(iii) Electropositive character: Due to high ionisation energies, the elements of group 14 are less electropositive as compared to group 13 elements. On moving down the group, the electropositive character increases due to the decrease in ionisation energies.

(iv) Melting and boiling points: There is a regular decrease in the melting and boiling points as we move down the group i.e. from carbon to lead. This is because as we move down the group, the size of the atoms increases and hence, interatomic forces of attraction decrease. As a result, lesser amount of energy is needed to melt or boil them.

(i) Oxidation states: All the elements of group 14 elements show +4 oxidation state. However as we go down the group, the stability of +4 oxidation state decreases while that of +2 oxidation state increases. This is due to the inert pair effect.
(ii) Metallic character:The elements of group 14 are less metallic than the elements of group 13. This is due to the less electropositive character of group 14 elements. Carbon and silicon are non-metallic, germanium is a metalloid while tin and lead are metallic. Thus the metallic character increases down the group.

The general electronic configuration of group 14 elements is [Noble gas] ns²np², so these elements are expected to show a uniform oxidation state of +4. Carbon and silicon show an oxidation state of +4 only, but tin and lead, however, show oxidation states of +2 and +4.
As we move down the group, the stability of +4 oxidation state decreases and that of+2 oxidation state increases. This means that as we move down the group, the tendency of the s- electron of the valence shell to precipitate in a bond formation decreases. In other words, the pair of s-electrons, 5s in tin and 6s in lead behaves as inert. This is called inert pair effect. Because of inert pair effect, only the np² electron of tin and lead take part in bonding with the atoms of the other elements. Thus, divalent compounds of tin and lead are stable.

C and Si do not exhibit inert pair effect due to the absence of d-or f-electrons. As a result, they show an oxidation state of +4 only due to the presence of two electrons in the s-and two electrons in the p-orbital of the valence shell.

On the other hand, the elements from Ge to Pb contain the only d- and f-electrons and hence show oxidation state of +2 and +4 due to inert pair effect. Further, as we move down the group, the number of d-and f -electrons increases and consequently the inert pair effect becomes more and more pronounced. In other words ns² electron pair in Ge, Sn and Pb tends to remain paired, thus stability of +2 oxidation state increases while that of +4 oxidation state decreases. Thus +2 oxidation state of Pb is more stable than its +4 oxidation state.

Carbon cannot have a valency beyond four
 
This is because of the absence of d-orbitals in the valence shell of carbon and presence of 3d-orbtials in the valence shell of silicon. Since silicon has vacant 3d- orbitals, it can accept electrons from electron donating species such as F^- ion to form complex-ion 

In graphite, carbon is sp² hybridised and each carbon is thus linked to three other carbon atoms forming hexagonal rings and also it involves the formation of pi-pi double bonds. On the other hand, the size of a silicon atom is larger than carbon and hence silicon has no tendency to undergo sp² or sp hybridization to form multiple bonds between themselves or with other atoms such as oxygen, nitrogen (for example Si = 0, Si = N). Therefore, silicon always undergoes sp² hybridization and hence all the silicon compounds have tetrahedral geometry. Also, silicon has a lesser tendency for catenation than carbon because Si – Si bonds are much weaker than C – C bonds.

(i) Carbon    (ii) lead     (iii) silicon and germanium.

The main reasons are:
(i) Six large chloride ions cannot be accommodated around Si⁴⁺ due to limitation of its size.
(ii) Interaction between lone pair of chloride ion and Si⁴⁺ is not very strong.

It is because of the presence of d-orbital in Si, Ge and Sn. Due to this,they have the tendency to form complexes by accepting electron pairs from donor species.
 

Allotropy: Allotropes or allotropic forms are defined as the different forms of an element which have different physical but similar chemical properties. The phenomenon of existence of allotropic forms is known as allotropy.
Allotropic forms of carbon:
Carbon exists in two types of allotropic forms:
(i) Crystalline forms: Diamond and graphite.
(ii) Amorphous forms: Coal, coke, charcoal, lamp black.

(i) Diamond
(ii) Graphite
(iii) Charcoal
(iv) Graphite
(v) Animal charcoal
(vi) Lamp black

Allotropes. Allotropes are defined as the different forms of an element which have different physical but similar chemical properties.

Both diamond and graphite are network covalent solids in which the carbon atoms are linked by covalent bonds.

In diamond: Each carbon atom is sp³ hybridised. It is bonded tetrahedrally to four other carbon atoms by simple covalent bonds. Thus, it has a three-dimensional network of strong covalent bonds in which C–C bond length is 154 pm and each bond angle is 109° – 28'. Due to sp³ hybridization, the carbon atoms in diamond are closely packed and all the valence electrons are involved in bonding, leaving no free electrons, so, diamond is very hard, bad conductor of electricity and has high melting point.

In graphite: Each carbon atom is sp² hybridised and is bonded to three other carbon atoms through covalent bonds forming hexagonal planar rings. The C–C covalent bond distance in rings is 142 pm. Graphite has a two-dimensional sheet-like structure. Due to sp² hybridization, one p-unhybridised orbital of the carbon atom in graphite

is not involved in bond formation. Thus, one valence electron of each carbon atom is free to move (mobile electron) from one point to other and this accounts for the electrical conductivity and soft nature of graphite.

In diamond, the carbon is sp³ hybridised. Each carbon atom is bonded to four other carbon atoms with the help of strong covalent bonds. These covalent bonds are present throughout the surface, giving it a very rigid 3-d structure. It is very difficult to break this extended covalent bonding and for this reason, diamond is the hardest substance known. Thus, it is used as an abrasive and for cutting tools.