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 np1-6. 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.
Name of family
Outermost electronic configuration
Boron family
ns2 np1
B, Al, Ga, In, Ti
Carbon family
ns2 np2
C, Si, Ge, Sn, Pb
ns2 np3
N, P, As, Sb, Bi
Chalcogens i.e. oxygen family
ns2 np4
O, S, Se, Te, Po
Halogens i.e. fluorine family
ns2 np5
F, Cl, Br, I, At
Noble gases
ns2 np6
He, Ne, Ar, Kr, Xe, Rn

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.

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.

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.

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

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.

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
        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

        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

        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

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.

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).

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.

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)

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.

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.

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 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 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.

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

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

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

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 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 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 SiO44-, 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.

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