P – Block Elements Class 11 Notes

P – Block Elements Class 11th Chemistry Notes

Elements in which distinguishing or valence electron enters in the p – orbital are known as P – block elements. The only exceptional element in this block is helium (atomic number=2) that has the electronic configuration 1s². The general valence shell electronic configuration of P–block elements is ns² np^1-6. They are placed to the right hand side of long form of the periodic table and includes elements belonging to group 13 to 18. They consist of metals, metalloids, non – metals and noble gases. There are six families in the p – block elements.

Group	Name of family	Outermost electronic configuration	Elements
13	Boron family	ns2 np1	B, Al, Ga, In, Ti
14	Carbon family	ns2 np2	C, Si, Ge, Sn, Pb
15	Nitrogen family	ns2 np3	N, P, As, Sb, Bi
16	Chalcogens i.e. oxygen family	ns2 np4	O, S, Se, Te, Po
17	Halogens i.e. fluorine family	ns2 np5	F, Cl, Br, I, At
18	Noble gases	ns2 np6	He, Ne, Ar, Kr, Xe, Rn

GROUP-13 ELEMENTS: BORON FAMILY

The group – 13 elements contain five elements namely boron (B), Aluminium (Al), gallium (Ga), indium (In) and thallium (Tl). The outer electronic configuration of group -13 element is ns²np¹. Boron is a semi – metal while aluminium and rest of the elements are metallic in nature.

OCCURRENCE

Boron does not occur in the Free State in the nature. It is fairly rare element. It mainly occurs as borates, orthoboric acid, borax and kernite. It shows two isotopic forms as ¹⁰B (19%) and ¹¹B (81%). Aluminium also does not occur in free state. It is forth most abundant element in the earth crust. Important minerals of aluminium are bauxite and cryolite. Gallium, indium and thallium are less abundant elements in nature. Thallium is highly toxic amongst the group.

PHYSICAL PROPERTIES

1. Atomic radii

 On moving down the group, atomic radius increases due increase in extra shell of electrons. But atomic radius of gallium is less than that of aluminium. This is because of variation in the inner core of the electronic configuration.

2. Ionization enthalpy

On moving down the group, regular decreasing trend (expected) in the ionization enthalpy value is not observed. It decreases from B to Al, increases in Ga and again decreases in In then increases in Tl.

IE₁ : B > Al ~ Ga > In < Tl

3. Electropositive or Metallic character

The elements of group – 13 are less electropositive than S – block elements. The electropositive character increases from B to Al and then decreases from Al to Tl. Aluminium is metal and is most electropositive.

4. Lewis acids

Trivalent compounds of group – 13 such as BCl₃, Bf₃, AlCl₃ etc., have six electrons around the central atom and thus act as Lewis acids by accepting a lone pair of electron to achieve stable electronic configuration. Lewis acid character decreases down the group.

5. Oxidation states

Boron shows only +3 oxidation state while other elements show +1 as well as +3 oxidation states.

CHEMICAL REACTIVITY

1. Action of air (oxygen)

Boron is unreactive towards air in crystalline form. While amorphous boron and aluminium metal on heating in air form B₂O₃ and Al₂O₃ also with nitrogen at high temperature they form nitrides. All the members of this group form oxides of the type E₂O₃.

4E   +   3O        ^Δ →    2E₂O₃

2E   +     N₂     →      2EN    (E = B, Al, Ga, In, Tl)

Oxides of B is acidic, Al and Ga oxides are amphoteric and those of In and TI are basic in their properties.

2. Reactivity towards water

Pure boron does not react with water while it reacts with steam at red heat. Aluminium decomposes boiling water evolving hydrogen. Gallium and indium are not attacked by water.

2Al   +   3H₂O   →   Al₂O₃   +    3H₂ ↑

3. Reactivity towards halogens

All the elements of group – 13 react with halogen to form trihalides (except TlI₃)

2E   +   3X₂   →   2EX₃

Where   X = F, Cl, Br, I  and  E =B, Al, Ga, In, Tl

4. Reactivity towards acids

Boron does not reacts with non – oxidizing acids (HCl and dil. H₂SO₄) but with oxidizing acids (conc. H₂SO₄ and HNO₃) boron forms boric acid.

B   +   3HNO₃   →   H₃BO₃   +   3NO₂

2B   +   3H₂SO₄   →   2H₃BO₃   +   3SO₂

While other elements react with any acid to form trivalent salt.

2E   +   6HCl   →   2ECl₃   +   3H₂ ↑

2E   +   3H₂SO₄   →   E₂(SO₄)₃   +   3H₂ ↑

5. Reactivity towards alkalies

Boron and aluminium react with alkalies to give borates and meta – aluminate with the evolution of H₂ gas.

2B   +   6NaOH       ^Fuse →   2Na₃BO₃   +   3H₂ ↑

2Al   +   2NaOH   +   2H₂O       ^Fuse →   2NaAlO₂   +   3H₂ ↑

Indium and thallium are not affected by alkalis.

6. Reducing action of boron

Boron element is a powerful reducing agent and it reduces CO₂ and SiO₂ to C and Si respectively.

4B   +   3CO₂    →   2B₂O₃   +   3C

4B   +   3SiO₂    →   2B₂O₃   +   3Si

ANOMALOUS PROPERTIES OF FIRST ELEMENT BORON

Boron differs from other members of group – 13 elements because of its small size (88pm), higher electronegativity (2.0) and higher value of ionization enthalpy (801 KJmol^-1). It shows following different properties.

·         Boron is a non – metal while other members are metals.

·         Boron shows allotropy while other members do not.

·         Boron has highest melting and boiling point amongst the elements of group – 13.

·         Oxides and hydroxides of boron are weakly acidic, of aluminium are amphoteric while those of rest of the members are weakly basic.

·         Boron hydride is quite stable while hydrides of other elements are less stable.

IMPORTANT COMPOUNDS OF BORON

1) Borax

It is also known as sodium tetraborate decahydrate (Na₂, B₄O₇.

10H₂O) Borax occurs naturally as tincal. It is obtained from the mineral colemanite by boiling it with a solution of Na₂CO₃.

Ca₂B₆O₁₁   +   2Na₂CO₃   →   Na₂B₄O₇   +   2NaBO₂   +   2CaCO₃

Colemanite                            Borax

Properties

        i.            Aqueous solution of boron is alkaline due to hydrolysis.

Na₂B₄O₇   +   7H₂O   →    2NaOH     +     4H₃BO₃

                                     Strong alkali      weak acid

      ii.            Boron when heated with ethyl alcohol and concentrated H₂SO₄, it gives volatile vapors of triethyl borate which burn with green edged flame.

This is used as test for borate radical in quantitative analysis.

    iii.            The borax bead is used to detect colored radicals. It is transparent glassy mass.

2]  Orthoboric acid (H₃BO₃)

Orthoboric acid is a white crystalline solid, with a soapy touch. It is obtained from borax by treading with dil. HCl or dil. H₂SO₄.

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

It can also be obtained from mineral colemanite by passing SO₂ through mixture of powdered mineral and boiling water.

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

Properties

        i.            It is a weak monobasic acid and behaves as a Lewis acid it accepts pair of electron from OH^- ion of H₂O.

H₃BO₃   +    H₂O    →    [B(OH)₄]^-   +    H⁺

      ii.            On heating it at 370K, it forms metaboric acid (HBO₂) which on further heating gives boric oxide (B₂O₃).

H₃BO₃     ^Δ   →    HBO₂      ^Δ  →    B₂O₃

    iii.            With NaOH it forms sodium metaborate.

H₃BO₃   +    NaOH    →   NaBO₂   +   2H₂O

     iv.            With ethyl alcohol and conc. H₂SO₄ it gives triethylborate which burns with green edged flame.

H₃BO₃   +   3C₂H₅OH      ^{conc. H2SO4}  →    B(OC₂H₅)₃   +    3H₂O

3] Diborane (B₂H₆)

It is the simplest boron hydride. It is prepared by treating boron trifluoride with LiAlH₄ in diethyl ether.

4BF₃   +   3LiAH₄   →   2B₂H₆   +   3LiF   +   3AlF₃

In laboratory the diboron is prepared by the oxidation of sodium borohydride wit h iodine.

2NaBH₄   +   I₂    →    B₂H₆   +   3NaI   +   H₂

In industry it is prepared by the reaction of BF₃ with NaH.

2BF₃   +   6NaH   →   B₂H₆   +   6NaF

Properties

        i.            Diborone is colorless, highly toxic gas. It has boiling point of 180K. It catches fire spontaneously when it is exposed to atmospheric air. It burns in oxygen. The reaction is highly exothermic.

B₂H₆   +    3O₂  →  B₂O₃  +   3H₂O

ΔH   =   -1976KJmol^-1

      ii.            It is readily hydrolysed by water to form boric acid.

B₂H₆   +   6H₂O  →  2H₃BO₃  +   6H₂

    iii.            With lewis bases, diborane first undergoes cleavage to form borane (BH₃) which then reacts to form adducts.

B₂H₆   +   2NMe₃   →   2BH₃.NMe₃ 

B₂H₆   +   2CO   →   2BH₃.Cn 

     iv.            It reacts with strong alkalies to form metaborates.

B₂H₆   +   2KOH   +   2H₂O  →  2KBO₂  +   6H₂

       v.            When diborane reacts with LiH in presence of ether it farms lithium borohydride.

2LiH   +   B₂H₆      ^Ether →    2LiBH₄

GROUP – 14 ELEMENTS  :  CARBON FAMILY

The outer electronic configuration of group – 14 elements is ns²np². This group includes carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). C and Si are non metals, Ge is a metalloid where Pb and Sn are metals.

OCCURRENCE

Carbon is widely distributed in nature in free state as well as in the combined form. In elemental form it occurs as coal, graphite, diamond and in combined form it occurs as metal carbonates, hydrocarbons and carbon dioxide. Silicon is the second (27.7% by mass) most abundant element in the earth crust. It is present in nature as silicon (SiO₂) and Silicates. Germanium exists only in traces. Tin occurs as SnO₂ and lead as galena (PbS).

PHYSICAL PROPERTIES

1. Atomic radii

Atomic radii of group – 14 elements regularly increase down the group. There is large increase in atomic radius from C to Si but from Si inwards there is small increase due to ineffective shielding of the valence electrons by intervening d – and f – electrons.

2. Ionization enthalpy

In general ionization enthalpy decrease down the group. From C to Si there is large decrease in ionization enthalpy, but from Sn to Pb there is  slight increases.

3. Electronegativity

Group – 14 are slightly more electronegative than corresponding group – 13 elements. Carbon (2.5) is most electronegative elment in this group. Electronegativity values from Si to Pb are almost same.

4. Oxidation states

Carbon shows only +4 oxidation state while other elements of the group show +4 as well as +2 oxidation states.

CHEMICAL REACTIVITY

1. Reactivity towards oxygen

All elements of group – 14 when heated in oxygen forms monoxide and dioxide.

2M   +   O₂   →   2MO

M   +   O₂    →    MO₂

Where, M = C, Si, Ge, Sn, Pb.

The dioxides of C, Si and Ge are acidic in nature and that of Sn and Pb are amphoteric in nature. The monoxides of C is neutral, Ge is distinctly acidic while Sn and Pb are amphoteric.

2. Reactivity towards water

Carbon, Silicon , germanium are not affected by water but react on red heating to form oxides. Lead is unaffected by water. Tin decomposes steam to form dioxide and hydrogen gas.

Sn   +  2H₂O     ^Δ   →    SnO₂   +    2H₂ ↑

3. Reactivity towards halogens

Carbon forms only tetrahalide while other elements of group react directly with halogens to form dihalides (MX₂) as well as tetrahalides (MX₃).

M   +   X₂   →   MX₂   (except carbon)

M   +   2X₂    →    MX₄

Stability of dihalides increases down the group while that of tetrahalides decreases down the group.

4. Reactivity towards alkalies

Carbon is unaffected by alkalies while Si, Sn and Pb are Slowly attacked by cold alkali to form silicate, stannates and plumbates respectively.

M   +   2NaOH  +  H₂O  →   Na₂MO₃  +  2H₂

(where, M = Si, Sn, Pb)

ANOMALOUS PROPERTIES OF CARBON

Carbon differs from rest of the element of group – 14 due to its smaller size, higher electronegativity and high ionization enthalpy.

Some points of difference are –

·         It is the hardest element of the group – 14. It has the highest ionization enthalpy and is most electronegative element of this group.

·         It’s maximum covalency is 4 as d – electrons are absent in the valence shell.

·         It has maximum property of catenation. It can form multiple bonds. As result carbon forms large numbers of compounds.

·         Carbon dioxides (CO₂) is a gas while the oxides of other elements are solids.

CATENATION

The property of self – linking of atoms of an element to form straight or unbranched chains and rings of different sizes by covalent bonds is called catenation. The tendency of an element to form a chains depends upon the strength of the element – element bond. Among the group  14 elements, C – C bond strength (348 KJmol^-1) is maximum, therefore carbon shows maximum tendency for catenation. The order of catenation is C >> Si > Ge ~ Sn. Lead does not show catenation.

ALLOTROPES OF CARBON

Allotropy is the phenomenon in which on element exist in a different physical forms having different physical properties but chemical properties are same or similar. Among the elements of group – 14, carbon, silicon and tin have allotropic modifications. Carbon has the largest number of allotropic modifications. 

Diamond

Diamond has three dimensional structur each carbon atom is bonded to four other carbon atoms which occupy four corners of a tetrahedron and each of these to four more carbon atoms. In diamond carbon atom is SP³ hydridized. C – C bond length is 154pm. Diamond has high melting point (3873K) and high density (3.51gmCm^-1). It is non conductor of electricity and is non conductor of electricity and is the hardest on the earth.

Graphite

Graphite has two dimensional sheet like structure and these are fused system of planner hexagonal ring. In the planner hexagonal rings, each carbon atom is bonded to three other carbon atoms. In graphite carbon atom undergoes SP² hybridization. In graphite the distance between two successive layers is 340 pm while bond length is 141.5pm. It conducts electricity.

Fullerenes or Bucky balls

Fullerenes is newly discovered allotrope of carbon in 1985. Fullerenes are made by heating of graphite in an electric arc in presence of inert gases such as helium or organ. It is cage like molecule with chemical formula C₆₀. It has a shape like soccer ball and called Buckminster – fullerene. Each carbon atom is SP² hybridized the average C – C distance is 144pm.

IMPORTANT COMPOUNDS OF CARBON : OXIDES OF CARBON

1] Carbon monoxide (CO)

        i.            It is prepared by direct combination of carbon in limited supply of oxygen.

2C   +   O₂     ^Δ   →   2CO

      ii.            By laboratory method pure CO is prepared by dehydration of formic acid with concentrated H₂SO₄ at 373K.

HCOOH     ^{H2SO4 373K}   →   CO   +   H₂O

    iii.            On commercial scale it is prepared by the passage of steam over hot coke.

C   +   H₂O   →   CO   +   H₂

                          Water gas

     iv.            When air is used instead of steam a mixture of CO and N₂ is produced, which is called producer gas.

2C   +   O₂   +   4N₂     ^1273K  →   2CO   +   4N₂

Properties

It is colorless and odorless gas and almost insoluble in water. It is powerful reducing in water. It is powerful reducing agent and reduces almost all the metal oxides. CO is used in the extraction of many metals from their oxides.

Fe₂O₃   +   3CO       ^Δ →   2Fe   +    3CO₂

CuO   +    CO        ^Δ →   Cu   +    CO₂

ZnO   +    CO     →   Zn   +    CO₂

It is a poisonous gas, burns with blue flame and when inhaled produces suffocation and finally death.

2] Carbon dioxide

        i.            It is prepared by complete combustion of carbon and carbon containing fuels in excess of air.

C   +   O₂      ^Δ →   CO₂

CH₄   +   2O₂      ^Δ →   CO₂   +   2H₂O

      ii.            It is prepared in laboratory by the action of dilute HCl on CaCO₃.

CaCO₃   +   2HCl   →   CaCl₂   +   CO₂   +   H₂O

    iii.            On commercial scale it is obtained by heating limestone.

CaCO₃       ^1600K →   CaO   +  CO₂  

     iv.            It is obtained as a byproduct in manufacture of ethyl alcohol by fermentation of glucose or fructose.

C₆H₁₂O₆    ^zymase  →  2C₂H₅OH   +   2CO₂

Properties

Carbon dioxide is colorless, Odorless and tasteless gas. It is slightly soluble in water and is heavier than air. CO₂ is acidic in nature it combines with alkalies to form metal carbonates.

CO₂ acts as an oxidizing agent. When heated with Zn, Fe or C these are oxidized.

Zn   +   CO₂     ^Δ  →   ZnO   +   CO

CO₂   +   C      ^Δ →    2CO

In photosynthesis, CO₂ is converted to glucose and oxygen.

6CO₂   +   12H₂O   →   C₆H₁₂O₆   +   6H₂O   +   6O₂

CO₂ is reacts with NH₃ at 453K – 473K under a pressure of 220 atmosphere to give urea as the final product.

2NH₃   +   CO₂    ^{453K – 473K , 220atm} →  NH₂CONH₂  +  H₂O

                                                              Urea

COMPOUNDS OF SILICON

Silicon dioxide (SiO₂)

Silicon dioxide is commonly known as silica. 95% of the earth crust is made up of silica and silicates. It occurs as quartz, cristabalite, and tridymite which are crystalline.SiO₂ is covalent in nature and in normal form it is non reactive because of very high Si – O bond enthalpy. However it is attacked by HF and NaOH.

SiO₂   +   4HF  →  SiF₄   +   2H₂O

SiO₂   +    2NaOH   →    Na₂SiO₃   +   H₂O

Silicon tetrachloride (SiCl₄)

It is prepared by action of chlorine on silicon

Si   +   2Cl₂   →    SiCl₄

This tetrahalide is covalent in nature and possesses tetrahedral geometry. CCl₄ is not hydrolysed by water but SiCl₄ gets easily hydrolysed.

SiCl₄   +   4H₂O   →   Si(OH)₄   +  4HCl

Silicones

Silicones are synthetic organosilicon compounds containing repeated R₂SiO units held by Si – O – Si linkage. These are formed by the hydrolysis of alkyl or aryl substituted chlorosilanes and their subsequent polymerization.

Silicates

Silicates are metal derivatives of silicic acid, H₄SiO₄ or Si(OH)₄.

Silicates are formed by heating metal oxide or metal carbonates with sand.

e.g. Na₂CO₃    ^{fused with sand, SiO2}  →   Na₄SiO₄

Silicates can also be termed as metallic salts of silicon – oxygen anions. Silicates have basic unit of SiO₄^4-, each silicon atom is bounded with four oxide ions tetrahedrally. Rocks, clays, soils are made up of silicates of aluminium, iron, magnesium and other metals.