The P -block Elements

The melting point of Gallium is 30°C and the boiling point is 2240°C.

Thus, the element Ga exists in the liquid state for a wide range of temperaturesand also due to high boiling point, it can be used for measuring high temperature.

Lewis acid is defined as a compound that has the ability to accept at least one lone pair.

The p-Orbital of Al is empty in compound AlCl₃.

Thus, it can accept a lone pair.

Whereas in all the other compounds the p-orbital is completely filled with electrons.

Boron has an electronic configuration:
In the excited state, 2s-orbital electrons are impaired and one electron is shifted to a p-orbital.

Now, hybridization occurs between one sand three p-orbitals to give sp³ hybridization and tetrahedral geometry.

B₂O₃ is acidic in nature. It reacts with basic oxides to form metal borates. Acidic nature decreases on moving down the group.

Thus, Boron being on the top of its group has the maximum acidic nature.

The element M in the complex-ion MF₆^3- has a coordination number of six.

Boron has only s- and p-orbitals and no d – orbitals.

Therefore, at the maximum, it can show a coordination number of 4. Thus, B cannot form a complex of the type MF₆^3-.

Because of the small size of boron atom and presence of only six electrons in its valence shell, B(OH)₃ accepts a pair of electrons from OH^- ion of H₂O, releasing a proton.

The reaction mechanism is:

The decrease in catenation property is linked with M – M bond energy which decreases from carbon to tin.

The greater the M-M bond energy the greater is the tendency to form catenation.

The energy difference between the Carbon-Carbon and silicon-silicon bond energy is very high and that between Ge-Ge and Sn-Sn is very less.

Me₃SiCl is added to control the chain length of a silicone polymer.

On moving down the group from B to Tl, a regular decreasing trend in the ionisation energy values is not observed.

This is due to Ga, there are 10 d-electrons in the penultimate shell which screen the nuclear charge less effectively.

Thus, outer electronsare held firmly. As a result, the ionization energy of both Al and Ga is nearly the same.

The increase in ionization energy from In to Tl is due to poor screening effect of 14f electrons present in the inner shell.

The 2 boron atoms and 4 terminal hydrogen atoms in diborane lie in the same plane and the 2 bridging hydrogen atoms lie in the perpendicular plane.

This can be seen in the structure of diborane:

The Compound X is B₂H₆ which reacts with the NH₃ give Y which is Borazine (also known as inorganic benzene).

Quartz behaves as a piezoelectric material as it can allow the conduction of electricity when subjected to mechanical stress.

The quartz is made up of crystalline SiO₄.

Thus, it contains Silicon Si.

The +4oxidation state of Sn is more stable than +2 oxidation state. Therefore, Sn²⁺ can be easily oxidized to Sn⁴⁺ and hence SnCl₂ acts as a reducing agent.
SnCl₂ + 2Cl → SnCl₄ + 2e^-

The other compound elements are in their most stable electronic state.

The solid-state of carbon dioxide gas is called dry ice.

It is given this name because it does not melt into a liquid state on heating rather changes to gaseous CO₂.

This process is called sublimation.

Cement is a product obtained by combining a material rich in lime. Calcium Oxide with other materials like clay which contains silica (SiO₂) along with the oxides of aluminium, iron, and magnesium

The average composition of Portland cement is:
CaO(50−60%),SiO₂(20−25%),Al₂O₃(5−10%),Fe₂O₃(1−2%),SO₂(1−2%) and MgO(2−3%).
Thus, it contains elements of group 2 (Ca), group 13 (Al) and group 14 (Si).

The atomic radius of Ga is less than Al because of poor screening effect.

The atomic radius of Ga is slightly lesser than of Al because in going from Al to Ga, the additional 10 electrons in 3-d orbital offer poor screening effect for the valence electrons.

The net nuclear charge increases on them and the size of Ga reduces.

The linear shape of CO₂ is due to the pπ-pπ bonding between carbon and oxygen and sp hybridization of carbon.

The chain length of the polymer can be controlled by adding Me₃SiCl, which blocks the ends of the silicone polymer.

Thus, by blocking the end of a polymer chain the polymerization of the organo-silicones is stopped and the chain length can be maintained.

Fullerenes are cage-like (soccer or rugby ball) molecules and graphite is thermodynamically most stable allotrope of carbon.

Thus, fullerenes have no dangling bonds and graphite is slippery and soft and is used as a lubricant.

The 2 boron atoms and 4 terminal hydrogen atoms in diborane lie in the same plane and the 2 bridging hydrogen atoms lie in the perpendicular plane.

Each of the two boron atoms is in the sp³-hybrid state.

Two of the four orbitals of each, of the boron atom overlap with two terminal hydrogen atoms forming two normal B - H σ-bonds. One of the remaining hybrid orbitals (either empty or singly occupied) of one of the boron atoms, 15-orbital of H (bridge atom) and one of the hybrid orbitals of the other boron atom overlap to form a delocalized orbital covering the three nuclei with a pair of electrons. This is a three-centre two-electron bond.

Similar overlapping occurs with the second hydrogen atom (bridging) forming three centre two electrons bond.

Carbon dioxide, or CO2, has three resonance structures, out of which one is a major contributor.

The CO₂ molecule has a total of 16 valence electrons.

The three resonance structure is:

The Lewis acid (BCl₃) and the Lewis base (NH₃) combine together to form a compound.

The structures of the BCl₃.NH₃is:

In AlCl₃, Al has six electrons in the valence shell. Therefore, it is an electron-deficient molecule and needs two more electrons to complete its octet. The Cl atom has a lone pair of electron.

Thus, a molecule of AlCl₃ combines with the other molecule of AlCl₃ to form a dimer in order to complete its octet.

Boric acid is a weak monobasic acid and acts as a Lewis acid by accepting electrons from a hydroxyl ion.

In the reaction, it is clearly shown that the OH^- is accepted by the Boric acid and the formation of hydroxyl ion takes place.

Thus, Boric acid act as a Lewis acid in water.

The central boron atom is connected to three hydroxyls (-OH) groups, which are capable of strong hydrogen bonding.

Its solid crystalline structure consists of parallel layers of boric acid held together in place by hydrogen bonds.

Hydrogen bonding in Boric acid

When boric acid reacts with water it forms species.

The hybridization of Boron in this species is sp³.

(i) InBCl₃, Boron has 6 electrons in it is outermost orbital and has a vacant p orbital. Thus, it is an electron-deficient compound.

Hence, it acts as Lewis acid and accepts a lone pair of electrons.

(ii) AlCl3 is an electron-deficient compound. Aluminum has three electrons in its valence shell so it forms a covalent compound with chlorine i.e. it forms three single bonds with chlorine.

The octet of Al is not complete by forming such bonds. Thus it can form a coordinate bondto complete its octet.

Hence, AlCl₃ acts as a Lewis acid.

(i); CCl₄ is non-polar and covalent in nature and water is polar in nature. Thus,CCl₄ is insoluble in water and secondly the carbon atom doesn’t have empty d-orbital to accommodate the electrons donated by the oxygen atom of the water.On another hand SiCl₄ is soluble in water because the Si atom has empty d-orbital to accommodate the electrons from oxygen atom.

Thus, SiCl₄is easily hydrolyzed whereas CCl₄ is insoluble in water.

(ii); The decrease in catenation property is linked with M – M bond energy which decreases from carbon to tin in group 14.

The greater the M-M bond energy the greater is the tendency to form catenation.

The energy difference between the Carbon-Carbon and silicon-silicon bond energy is very high i.e. the tendency of carbon to form catenation is greater than the silicon.

(i) Carbon has a much smaller size compared to silicon.

Due to its small size it forms strong double bonds with oxygen atom (due to ) to give discrete CO₂ molecules.

Silicon atom has a large size and does not form good with other atoms. It forms four single covalent bonds with oxygen atoms. Each oxygen atom is linked with two Si atoms. Thus, a large giant molecule of 3-d structure comes into existence.

Thus, CO₂ is a gas whereas SiO₂ is solid.

(ii); Silicon has lower energy 3-d orbital so it can expand its octet givingsp³d² hybridization while d-orbitals are not present in the valence shell of carbon. It can undergo sp³-hybridisation only. The size of carbon atom is very small to accommodate six F^- anions

Thus, carbon is unable to form CF^6- anion.

This property is due to an effect called the inner pair effect.

In groups 13 and 14 as we move down the group the tendency of the s-orbital electrons to participate in the bond formation decreases this is due to the poor shielding of s-orbital electrons by the d and f orbitals. This is called an inner pair effect.

Thus,the +1 oxidation state in group 13 and +2 oxidation state in group 14 becomes more and more stable with increasing atomic number.

Carbon has ability to form stable bonding with itself and other small atoms like oxygen and nitrogen due to its small size.it carbon dioxide each oxygen atom is double-bonded with the carbon atom with overlapping.

The silicon cannot form stable bonding due to its large size. It forms four single covalent bonds with oxygen atoms. Each oxygen atom is linked with two Si atoms. Thus, a large giant molecule of 3-d tetrahedral structure comes into existence.

If a few tetrahedral Si atoms in a three-dimensional network structure of SiO2 are replaced by an equal number of trivalent atoms, then one valence electron of each Si atom will become free.

As a result, the substitution of each Si atom by a trivalent atom will introduce one unit negative charge into the three-dimensional network structure of the SiO₂.

Thus, the SiO₂ structure will become negatively charged.

The reaction of the BCl₃ with water occurs in two steps:

The Boron in the ion involves one 2s orbital and 3 2p orbitals.

Thus, the hybridization of the Boron atom is sp³.

The 6 H₂O molecules get attach with Al i.e. donate 6 electron pair to the 3s,3p and 3d orbital of Al³⁺ ion.

Thus, the hybridization of the Al atom in species is sp³d².

Aluminum is an amphoteric metal i.e it reacts with both acid and base to give hydrogen gas which burns in air with a pop sound.

Nitric acid is a very strong oxidizing agent it leads to the formation of a thin layer of Al₂O₃(aluminum oxide) on the surface of the aluminum.

Thus, the further reaction is prevented and no hydrogen is liberated.

(i); Due to the ineffective shielding of valence electrons by the intervening 3d electrons, the effective nuclear charge on Ga is slightly higher than that on Al.

Hence the ΔH_ionization of gallium is slightly higher than that of Al.

(ii); Boron has three electrons in the valence shell. Because of its small size and a high sum of the first three ionization enthalpies, it doesn’t form B³⁺ ion.

(iii); Aluminum has empty d-orbital to accommodate the electrons from the fluorine atom but boron has no empty d-orbital to accommodate the electrons.

Thus,aluminum forms [AlF₆]^3- ion but boron does not form [BF₆]^3- ion.

(iv); Pb is the member of the group 14 of the periodic table (carbon family).The valence shell electronic configuration of this element is ns² np² type. Pb can show variable oxidation states of +2 and +4.

On moving top to bottom in the group, the lower oxidation state i.e +2 is more stable than the higher one due to the inert pair effect.

Thus, Pbdue to the inert pair effect (poor shielding of the inner electronic orbitals) shows +2 as a stable oxidation state rather than +4.

Thus, PbX₂ is more stable than PbX₄.

(v); Sn and Pb are the members of the group 14 of the periodic table (carbon family).The valence shell electronic configuration of these elements is ns² np²type.All these elements contain four electrons in the valence shell.These elements show variable oxidation states of +2 and +4.

On moving top to bottom in the group, the lower oxidation state i.e +2 is more stable than the higher one due to the inert pair effect.

Thus, Pbdue to the inert pair effectshows +2 as stable oxidation state rather than +4. So, by accepting two electrons Pb⁴⁺ will get converted into Pb²⁺.

Thus, Pb⁴⁺ by undergoing self-reductionacts as an oxidizing agent.

In the case of Sn, the inert pair effect is very less,the +4 oxidation state is more stable than +2.

Thus,by losing two electrons,Sn²⁺ gets converted into Sn⁴⁺ and can act as reducing agent.

(vi); Fluorine has a very small size thus the electrons of the 2p orbital experiences the interelectronic repulsion.

Thus, the electron coming out of the orbital does not experience much of the attraction from the nucleus hence the negative electron gain enthalpy of fluorine is less than that of chlorine.

(vii); Due to the poor shielding effect of the inner electronic orbitals, the +3 oxidation state of Tl is less stable than its +1 oxidation state.

Since in Tl(NO₃)₃, the oxidation state of Tl is +3, therefore, it can easily gain two electrons to form TlNO₃ in which the oxidation state of Tl is +1.

Consequently, Tl(NO₃)₃ acts as an oxidizing agent.

(viii);The catenation property depends on two factorsone is atom size and second is the M-M bond energy. Smaller the atomic radii and the greater the M-M bond energy, the greater is the tendency to show catenation.

On moving down a group the atomic size increases and the M-M bond energy also reduces.

Thus, Carbon shows the catenation whereas lead does not.

(ix);BF₃ does not hydrolyze completely. Instead, it hydrolyzes incompletely to form boric acid and fluoroboric acid. This is because the HF first formed reacts with H₃BO₃.

These equations can be written as:

Thus, BF₃ does not hydrolyze completely.

(x); Graphite is a macromolecule that is made up of hexagonal carbon rings.

This property is due to the catenation of the carbon atoms. Catenation is the tendency of an element to form covalent bonds between its atoms.

The catenation property depends on two factorsone is atom size and second is the M-M bond energy. Smaller the atomic radii and the greater the M-M bond energy, the greater is the tendency to show catenation.

On moving down a group the atomic size increases and the M-M bond energy also reduces.

Thus, Carbon shows the catenation whereas silica does not.

A is Borax which reacts with HCl in the presence of water to give Orthoboric acid(X).

Orthoboric acid onstrongly heating(up to 370K) gives metaboric acid which on further heating (T>370K) yields Boron trioxide (Z).

Z is BF₃ which on reaction with a strong reducing agent like LiAlH₄ gives B₂H₆ (X) which on oxidation yields borate.

Thus, the reactions are:

(i) BF^-₄- d) Can be further oxidised

Oxidation is a process which involves the gaining/ addition of Oxygen or removal of Hydrogen. It can also be stated as loss of electrons from any given ionic species or molecule. In the case of BF^-_4, it contains 4 Fluorine atoms, where 3 atoms covalently bind to Boron atom and the remaining Fluorine atom donates its lone pair of electron and forms BF^-_4. Now, the extra Fluorine atom if gets removed, it can be oxidised to form BF₃ by losing extra electrons.

ii) AlCl₃– c) Lewis acid.

The species which are electron deficit or electrophilic in nature and are capable of accepting electrons are referred to as Lewis acids.In other words, any substance can act as Lewis acid which can accept lone pair or non-bonding electrons. In the case of AlCl₃, Al has empty p- orbitals, thus, it can accept electrons in order to complete its octet. Therefore, it acts as a Lewis acid.

ii) AlCl₃– e) Tetrahedral shape

At normal temperature, Aluminium Chloride exists in solid lattice structure. But as soon the temperature increases (180^oC - 190^oC), Aluminium chloride starts converting into liquid state by forming a more stable of dimer of AlCl₃ and exhibits in tetrahedral structure.

iii) SnO – d) Can be further oxidised

Oxidation is a process which involves the gaining/ addition of Oxygen or removal of Hydrogen. It can also be stated as loss of electrons from any given ionic species or molecule. SnO can be further oxidised to form SnO₂by addition of another Oxygen atom. Though, Stannous oxide is much more stable than Stannous dioxide.

iv) – b) stronger oxidising agent

is a stronger oxidising agent. Inert pair effect is defined as the inertness caused by ns electrons of outer shell during bonding because of poor shielding effect. Due to dominating effect of inert pair effect, +2 oxidation state is more stable, thereby making the conversion of to easily making it a more stronger oxidising agent.

iv) - (a) Oxidation state of central atom is +4

Lead has the oxidation number of +4 in Lead oxide.

Oxidation number is calculated as sum of oxidation number (ON) of all atoms in a polyatomic ion which is equated with the total charge of the molecule or ion.

Let ON of Pb be x

So, ON of each O = -2 and total charge = 0

So,

i) Diborane - (c) Banana bonds

The formation of diborane leads to the creation of new types of bonding with hydrogen atoms a) 4 terminal – 2 centered – 2 electrons and b) 2 bridging – 3 centered – 2 electrons, leading to the formation of banana bonds. In other words, in total there are 12 valence electrons, (3 from each Boron and 6 from 6 6 H-atoms), now, out of these 4 H-atoms bind covalently with the 2 Boron atoms (8 used), while the remaining 4 electrons are shared by 2 Hydrogen – Boron atoms, form B-H-B bridge. This bridging is referred to as Banana bond.

(ii) Gallium - (d) Low melting, high boiling, useful for measuring high temperatures

Due to greater cohesive forces which held the structure together, Gallium exhibits high boiling point while its melting point is low. Therefore, used to measure high temperatures.

(iii) Borax - (a) Used as a flux for soldering metals

Borax is used for soldering metals. Borax is used

along with (flux) for welding purposes. Sometimes, it is also

mixed with water to solder jewellery and gold.As, it lowers the

melting pointand allows themolten mixture to flow easily in the

shapes. Thus, helps used in soldering.

(iv) Aluminosilicate - (e) Used as catalyst in petrochemical industries

Aluminosilicate is a compound which is composed of Aluminium and Silicon mainly, also called as Zeolites. Nowadays, aluminosilicates are commonly used in petrochemical industries as a catalyst used to synthesis ethylene and propene. These are the compounds which can accelerates the chemical reactions, called as catalysts. They are usually used in the preparation of gasoline from crude oil and in the petrochemical industries for cracking of hydrocarbons.

(v) Quartz - (b) Crystalline form of silica

Quartz is composed up of elements - Si and Oxygen and having formula . It is a crystalline structure formed by the joining of repeating units of. It is available in several colorsand varieties. They used in jewellery making and stone carving.

(a part showing crystalline form of Quartz)

i) hybridisation of Boron in - b) sp³

The EC of B is 1s² 2s² 2px¹ 2py⁰ 2pz⁰. In the excited state, one of the electrons from 2s orbitals gets shifted to a p-orbital. The tetrahedral geometry is formed when hybridisation occurs between one s and 3 p-orbitals to give sp³ hybridisation.

Here, in this case 3 molecules bind to the 3 orbitals of the Boron, but still the octet of the Boron is not yet completely filled. Therefore, an extra molecule of binds to it. Thereby involving one s-orbital and 3 p-orbitals in the hybridization process. Thus, it is hybridised.

ii) Aluminium in [Al(H₂O)₆]³⁺ - c) sp³d²

The EC of Al is 1s² 2s² 2px² 2py² 2pz² 3s² 3px¹. While in the excited state, one of the electrons from 2p and 3s orbitals gets shifted to 4s and 3d orbitals.

In other words, it can be said that in the excited state Al loses its 3 valence electrons (from the s and p-orbitals of the valence shell). Now, it accepts 6 electron pairs from water molecule and extended its orbital to d. Thus, forming sp³d² hybridisation.

iii) Boron in - b) sp³

The EC of B is 1s² 2s² 2px¹ 2py⁰ 2pz⁰. In the excited state, one of the electrons from 2s orbitals gets shifted to a p-orbital. When hybridisation occurs between one s and 3 p-orbitals it gives sp³ hybridisation.

Normal state EC of B: 1s² 2s² 2px¹ 2py⁰ 2pz⁰

Here, in this case 3 molecules bind to the 3 orbitals of one Boron. Similarly, 3 hydrogen ions are bonded with another Boron atom by making 1 s- and 3 p- orbitals. Thereby involving one s-orbital and 3 p-orbitals in the hybridization process in each case. Thus, it is hybridised.

iv) Carbon in Buckminsterfullerene – a) sp²

In this case, each carbon atom is sp² hybridised while the remaining p- orbital available for bonding binds with another C-atom with sp² hybridisation.

Normal state EC of C: 1s² 2s² 2px² 2py⁰ 2pz⁰

In the case of Buckminsterfullerene, it is composed of 60 atoms of carbon (allotrope of Carbon), here, the hydrogen atoms hybridize with Carbon atoms while, the last p-orbital participates in hybridisation with the adjacent Carbon atoms. Thus, sp² hybridisation.

v) Silicon in SiO^4-₄ - b) sp³

The EC of Si is 1s² 2s² 2p⁶ 3s² 3p². In the excited state, one of the electrons from 3s and 3p orbitals gets shifted to a p-orbitals. When hybridisation occurs between one s and 3 p-orbitals it gives sp³ hybridisation.

Normal state EC of Si: 1s² 2s² 2p⁶ 3s² 3p²

In this case, due to the presence of vacant d-orbitals Si extends its coordination number and forms bonding and accommodates four atoms of oxygen.

(vi) Germanium in [GeCl₆]^2-- c) sp³d²

The EC of Ge = [Ar] 3d¹⁰ 4s² 4p² 4d⁰

EC of Ge (IV) [Ar] 3d¹⁰ 4s⁰ 4p⁰ 4d⁰

Now, 6 Cl atoms donate 6 lone pairs of electrons to Ge which gets accommodated in 1 4s orbitals, 3 4p orbitalswhile the remaining 2 Cl atoms gets into the 4d orbitals thereby forming sp³d² hybridisation.

If aluminium atoms replace a few silicon atoms in 3D network of silicon dioxide, the overall structure acquires a negative charge. Then, Aluminium is trivalent while silicon is tetravalent. It is so because as the equivalent number of trivalent ions like (Al⁺³) replace few tetrahedral Si atoms in the 3D structure of Silicon dioxide, the valence electron of each Si atom gets freed, thereby resulting in the formation of a negatively charged structure. Thus, SiO₂ acquires negative charge.

Therefore, both Assertion and Reason are correct and Reason is the correct explanation of Assertion.

So, option (i) is correct.

Silicones are made up of alkyl groups having general formula like which are non-polar in nature. Thus, they show water-repellent nature.

Silicones are organosilicon polymers, which have (–R₂SiO–) as repeating unit. As they are made up of Silicon and Oxygen, therefore, they act as polar compound. So, they show hydrophilic nature.

Therefore, both assertion and reason are correct statements but reason is not correct explanation of assertion.

So, option (ii) is correct.

(i) ATOMIC SIZE

In group 13

Down the group, atomic radius increases and every successive member increases 1 extra shell of the electron while going down the group, but some deviation is seen where atomic radius of gallium is less than that of aluminium, this can be understood by inner core of electronic configuration which says d orbital does not screen effectively whereas s and p orbital can. Gallium has 10 electrons in 3d subshell and each shell of d orbit does the poor screening of nuclear charge than s and p orbital. Therefore nuclear charge increase and atomic radius of gallium which is 135 pm less than that of aluminium which is 143 pm

B < Ga < Al < In < Tl

In group 14

Down the group, covalent radius increases from carbon to silicon and then from Silicon to lead a small increase is observed due to presence of completely filled d and f orbital in heavy members.

C < Si < Ge < Sn < Pb

(ii) IONISATION ENTHALPY

In group 13

Ionisation enthalpy does not decrease down the group as expected. There is a decreasing trend from B to Al with increase in size. The discontinuity occur in Al to Ga and In to Tl due to inability of d and f orbital to screen the nuclear charge. They have low screening effect.

B > TI > Ga > Al > In

In group 14

First Ionisation enthalpy decreases from silicon to tin then show a slight increase in ionisation enthalpy from tin to lead due to the poor shielding effect of d and f orbital as there is a poor screening of nuclear charge in d and f orbital in comparison to s and p orbital. C > Si > Ge > Sn < Pb

(iii) METALLIC CHARACTER

In group 13

Boron is non-metallic in nature and is very hard and is black in colour. Boron also consists of allotropes like carbon. Due to very strong crystal lattice boron has a high melting point, rest are soft metals with less melting point.

In group 14

Carbon and silicons are non-metals, germanium is metalloid and tin and lead are soft metals with less melting point

(iv) OXIDATION STATE

In group 13

Boron has a small size and first 3 ionisation enthalpy are very high, which prevent boron to form +3 oxidation state and force boron to form only covalent compounds. Now when we move down the group ionisation enthalpy drastically decrease and aluminium forms Al3+ and due to poor shielding effect of d and f orbitals, the increase in nuclear charge hold ns electron tightly which shows inert pair effect (effect which shows oxidation state reduce by 2, which is more stable than other oxidation states) and restrict them to participate in bonding and as a result p orbital electron participate in bonding. Ga, In, and Tl shows both +1 and +3 oxidation state. The compound in +1 oxidation state is more ionic than +3 oxidation state.

In group 14

Oxidation state is +2 and +4. +2 oxidation state is common in heavy elements Ge< Sn<Pb because the ns² electrons of valence shells are unable to participate in bonding. Carbon and Silicon consist of a +4 oxidation state. Germanium is stable in +4 oxidation state and some compound of germanium in +2 states. Tin form compound in both +2 and +4 state and Tin is a reducing agent in +2 states. Lead is a strong oxidising agent in +4 oxidation state and very stable in +2 oxidation state.

(v) NATURE OF HALIDES

In group 13

Element react with halogens to form trihalides except TlI3

2E + 3X₂ →2EX₃

X can be F, Cl, Br, I

The trichloride, tribromide, triiodide are covalent in nature and hydrolysed in water. The monomeric trihalides are electron deficient so-called as strong Lewis acid.

In group 14

The element can form halides of MX₂ and MX₄, X can be F, Cl, Br, I.Except carbon all other member react directly with halogen under some condition forming halides. Mostly MX₄ is covalent in nature. Central metal goes in sp3 hybridisation forming the tetrahedral shape. SnF₄ and PbF₄ are exceptions as they are ionic in nature. PbI₄ does not exist as it form initially and does not releaseenergy for 6s² electron and excite one of them to higher orbital so that lead has 4 unpaired electrons around it. Ge and Pb make the formula of MX₂. Down the group stability of dihalides increases.

(i) Aluminium has 3 electrons in its valence shell ([Ne]3s²3p¹) which are incomplete and chlorine has electronic configuration of ([Ne3s²3p⁵]),In trivalent state, the central atom in a molecule of a compound has only 6 electrons and such deficiency of electron have a tendency to accept the pair of electron to achieve stable electronic configuration and behave as a Lewis acid.

(ii) Boron, ([He]2s²2p¹) in BF₃has incomplete 2p orbital whereas fluorine ([He]2s²2p⁵) has completely filled 2p orbital. BF₃ being electron deficient are strong Lewis acid and also BF₃ reacts easily with Lewis base like NH₃ to complete the octet around boron but the aptness to behave as Lewis acid decrease with increase in size down the group.BCl₃ accept lone pair of ammonia to form BCl₃.NH₃.

(iii) Sn(tin) decompose steam to form SnO₂ and hydrogen

Sn + 2H₂O → SnO₂ + 2H₂

But lead forms protective oxide film when reacting with water.

That is why PbO₂ is more oxidising agent than SnO₂.

(iv)Thallium, with electronic configuration of [Xe]4f¹⁴5d¹⁰6s²6p¹ and has both +1 and +3 oxidation state. +1 oxidation state is more dominant but +3 oxidation sate is highly oxidising in nature. The stability of +1 oxidation state increases for heavy element, thus +1 oxidation state is more stable than +3 oxidation state.

Al <Ga< In <Tl

Orthoboric acid is formed when the aqueous solution is acidified with borax. Orthoboric acid is crystalline and soapy to touch. It is soluble in hot water.

Na₂B₄O₇ + 2HCl + 5H₂O → 2NaCl + 4B(OH)₃

Boric acid is a weak monobasic acid. It acts a lewis acid by taking electrons from the hydroxyl group.

B(OH)₃ + 2H₂O → [B[OH]₄]^- + H₃O⁺

Heating orthoboric acid at 370K gives metaboric acid and on further heating give Boric acid

H₃BO₃ → HBO₂ → B₂O₃

(i) Thallium, with electronic configuration of [Xe]4f¹⁴5d¹⁰6s²6p¹ and has both +1 and+3 oxidation state. +1 oxidation state is more dominant but +3 oxidation sate is highly oxidising in nature. The stability of +1 oxidation state increases for heavy element, therefore TlCl will be more stable than TlCl₃.

Al < Ga < In < Tl

(ii) In aluminium, the sum of the first 3 ionisation enthalpy decreases, thus aluminium form Al³⁺. Also, there are No d and f orbital in aluminium which cause poor shielding and responsible for inert pair effect (an effect which shows oxidation state reduce by 2, whichis more stable than other oxidation states). Therefore AlCl₃ is more stable than AlCl.

(iii)Indium with electronic configuration [Kr]4d¹⁰5s²5p¹and has both +1 and +3 oxidation state. Due to inert pair effect(an effect which shows oxidation state reduce by 2, which is more stable than other oxidation states), in heavy elements, the stability of +1 oxidation state increases, therefore, InCl will be more stable than InCl₃.

Al < Ga < In < Tl

BCl₃ is in trivalent state and has 6 electrons in its valence shell and need two more electrons to complete the octet. BCl₃ act as a lewis acid and NH₃ can donate its electron. BCl₃readily accepts the lone pair of an electron from NH₃ and form BCl₃.NH₃.

BCl₃ + NH₃ → BCl₃.NH₃.

AlCl₃ forms a stable molecule by making a Dimer. Aluminium has 6 electrons in the valence shell and need 2 more electrons to complete the octet whereas chlorine has 3 lone pair of electron. Aluminium takes up 1 lone pair and forms Dimer with chlorine.

pπ−pπ back bonding occurs in BF₃.Boron has vacant 2p orbital and fluorine has completely filled 2p orbital so fluorine transfer 2p electron to boron. Now Boron’s deficiency has been filled making BF₃ stable.

In BH₃ hydrogen has lone pair and cannot fulfil the deficiency of boron so dimerise to form B₂H₆ which is in the shape of a banana and also known as a banana bond.

(i) Silicon are the group of organosilicon polymers having (R₂SiO) repeating unit. For manufacturing silicone, starting material is alkyl or aryl substituted silicon chloride.

Methyl chloride react with silicon in the presence of copper powder as a catalyst at 573K, many types of methyl chlorosilanes like MeSiCl₃, Me₂SiCl₂, Me₃SiCl with small amount of Me₄Si are formed. Further addition of water form the straight-chain polymer.

Uses:

They are used as greases, sealant, for waterproofing fabric, electrical insulators.

Also used as a surgical and cosmetic plant.

(ii) Borane is a simple form of boron hydride. It is prepared by treating lithium aluminium hydride with trifluoride in diethyl ether.

4BF₃ + 3 LiAlH₄ → 2B₂H₆ + 3LIF + 3AlF₃

It is also prepared by oxidation of sodium borohydride with iodine. For example:

2NaBH₄ + I₂ → B₂H₆ + 2NaI + H₂

Diborane is also produced on an industrial scale by reacting with BF₃ with NaH

2BF₃ + 6NaH → B₂H₆ + 6NaF

Diborane is colourless gas and is highly toxic. It has boiling point 180K.

B₂H₆ + 2NMe → 2BH_3.NMe₃

A B C

3 B₂H₆ + 6NH₃ → 3[BH₃(NH₃)₂]⁺[BH₄]⁺ → 2B₃N₃H₆ + 12H₂

A = B₂H₆ (DIBORANE)

B = 2BH_3.NMe₃ (ADDUCT)

C = 2B₃N₃H₆ (INORGANIC BORAZINE)

Element is Boron which is very hard and black in colour. It's melting point is very high.

BF₃ + NH₃ → BF₃ß NH₃

Monomeric trihalides have deficient electrons so they act as strong lewis acid. To complete the octet of boron, trifluoride of boron easily react with Lewis bases such as NH₃. It is due to the absence of d electrons in boron, it cannot exceed the covalence above 4.

2C + O₂ →2CO (direct oxidation of carbon)

_373K

HCOOH → H₂O + CO (DEHYDRATION OF FORMIC ACID AT 373K WITH CONCENTRATED H₂SO₄)

C + H₂O → CO + H₂(a mixture of CO and H₂ is called water gas)

2C + O₂ + 4N₂ → 2CO + 4N₂ (a mixture of CO and N₂is producer gas)

REDUCTION OF FERRIC OXIDE WITH MONOXIDE

Fe₂O₃ + 3CO → → 2Fe + 3CO₂