CBSE Class 11 Chemistry Revision Notes Chapter 11

Class 11 Chemistry Revision Notes for Chapter 11 – The p-Block Elements 

Furnished with additions from all significant elements of Class 11 Chemistry Chapter 11, these revision notes from Extramarks are available in a downloadable format on the Extramarks website. The notes have quick access to all subheadings of the chapter and a detailed explanation of the necessary concepts. These are curated for quick and easy preparation for the chemistry CBSE examinations. Subject matter experts of chemistry have created these notes for the benefit of students of any calibre. 

Access Class 11 Chemistry Chapter 11 – The p-block elements

Introduction:

Metals, metalloids and non-metals, belonging to groups 13 to 18 of the periodic table come under p-block elements. The electrical configuration of these is ns2np1-6. 

Some characteristics of p-block elements are as follows :

• They are tiny in size
• Strong electronegativity
• Lack of d-orbitals
• The first member of a group belonging to the p-block has the capacity to form pπ-pπ multiple bonds within itself.
• The group number minus 10 is the maximum oxidation 

The last electron moves into the outermost p orbital in p-block elements. The maximum number of electrons that can fit in a set of p orbitals is six because we know that there are three p orbitals. As a result, the periodic table contains six groups of p-block elements, ranging in number from 13 to 18. The groups are headed by boron, carbon, nitrogen, oxygen, fluorine, and helium. A p-block element’s maximum oxidation state corresponds to the total number of valence electrons (i.e., the sum of the sand p-electrons).

Trends in Properties of P-Block Elements:

Various trends are observed in the p-block with respect to many characteristics. 

(A) Group 13 Elements: The Boron Family:

Group 13 includes boron, a nonmetal, aluminium, and a metal showing chemical similarities to boron and gallium, indium, and thallium which are typical metals. 

1. Electronic configuration : ns2np1
2. Atomic radii: Typically, atomic radii increases in a group and so does the size. However, for Gallium, the presence of an extra 10 d-electrons gives a screening effect for the outer electron, pulling them closer due to the high nuclear charge. Hence Gallium is smaller in atomic radii than aluminium, although it comes after aluminium in the group.

B < Ga < Al <In < TlB < Ga < Al < In < Tl

• Ionisation enthalpy: The sum of the first ionisation enthalpies of each element is very high, and the values are not consistently decreasing in group 13. As size increases from boron to aluminium, the drop is justified. However, as the d and f electrons cannot shield the increased nuclear charge in gallium there is an inconsistency in values between Al and Ga as well as between In and Tl. 

B > Al > Ga > In > TlB > Al > Ga > In > Tl

• Electronegativity: As atomic size varies between elements,  a decrease in electronegativity from boron to aluminium and a sudden increase down the group is observed.

B > Tl > In > Ga > Al Tl > In > Ga > Al

Physical properties : 

• Boron is a solid black, hard non-metallic element that shows a variation in allotropic forms. It has a high melting point due to a crystalline lattice. 
• Gallium has a low melting point (303 K) and thus may be liquid during summers. It, however, has a high boiling point (2676 K) which can be used to measure high temperatures. 
• The density of the elements increases from B to Tl

Chemical Properties:

• Boron has an extremely high sum of its first three ionisation enthalpies, which makes it difficult to form +3 ions. It forms covalent compounds. For aluminium, this sum is not that high, allowing it an Al3+ configuration. 
• Down the group, the bonding of ns electrons decreases, and hence the p-orbital participates primarily in bonding. 
• Ga, In and Tl show oxidation states +1 and +3, with +1 being most common in thallium and +3 state being the most oxidising. 
• Compounds in the +1 state are more ionic than in the +3 state. 
• Electron deficient molecules act as Lewis acids, to reach a stable configuration. As the group grows smaller, the tendency to act as a Lewis acid reduces. 
• A tetrahedral compound, [M(OH)4]– is formed as a trichloride when hydrolysed in water. 
• Reactivity Towards Air: In the presence of oxygen, boron and aluminium form B2O3 and Al2O3 upon heating in the air. In its crystalline form, boron is non-reactive and a thin layer of oxide develops on aluminium. The creation of nitrides is observed when dinitrogen is heated to a raised temperature. 

2E + 3O2  → 2E2O3

2E + N2  →  2EN 

• Reactivity Towards Acids and Alkalis: Boron shows no reaction with acids and alkalis, but Al shows amphoteric property by dissolving in mineral acids and aqueous alkalies. The reaction of Al with diluted HCl and an aqueous alkali (NaOH) is given below :

2Al + 6HCl  → 2Al3+ + 6Cl– + 3H2

2Al + 2NaOH + 6 H2O  →  2Na+ [Al (OH)4 ] + 3H2

• Reactivity Towards Halogens: Reaction with halogens results in trihalide formations, except in ThI3

2E + 2X2  → 2EX3   (X = F, Cl, Br, I)

Important Trends and Anomalous Properties of Boron:

Group 13 elements’ chemical behaviour exhibits a few significant trends. Tri-chlorides, bromides, and iodides, which are covalent in nature, are hydrolyzed in water.

Species like tetrahedral [M(OH)4]– and octahedral [M(H2O)6] 3+, except in the case of boron, exist in an aqueous medium. Due to their electron deficiency, monomeric trihalides are strong Lewis acids. To complete the octet around boron, boron trifluoride easily reacts with Lewis bases such as NH3.

F3B  +  :NH3  →   F3B   ← NH3

It is due to the absence of d orbitals that the maximum covalence of B is 4. Since the d orbitals are available with Al and other elements, the maximum covalence can be expected beyond 4.

Boron (B):

Borax, orthoboric acid and diborane are some of the useful boron compounds.

Occurrence:

The following minerals may show the occurrence of boron :

An anionic complex, Borax (Na+)2B4O2−7• 10H2O. 

Boric acid H3BO3

Kernite Na2B4O7 •4H2O 

Colemanite Ca2B6O11 • 5H2O

Extraction of Boron:

1. Reducing B2O3 with Mg, Na or K in the absence of oxygen 

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

                              Δ

2H3BO3   —-→  B2O3 + 3H2O    ;    B2O3 + 3Mg —→ 2B + 3MgO

This leaves behind dark, amorphous boron powder B.

• Heating potassium fluoroborate (KBF4) with potassium metal yields boron.

KBF4 + 3K   →  4KF + B

Treating with dil. HCl after this process leaves behind pure B.

1. Crystalline boron can be formed either by reduction of BBr3 with H2 upon a heated titanium metal filament at 1275 – 1475 K or by the decomposition of BI3 vapours using tungsten electrodes and a high tension arc (80kV)

Properties:

1. 1. Boron is found in crystalline forms as well as amorphous forms. Crystalline boron is chemically inert, black in appearance and made up of hard B12 clusters. Melting points are in the range of 2300℃. Amorphous boron is chemically active, and brown. 

• Upon reaction with air, 

2E  +   3O2   →   2E2O3 ;         2E +  N2   →  2EN

• The action of alkalis and acids,

2B + 2NaOH + 2H2O →  2NaBO2    +    3H2 

          oxidation

2B  + 3H2SO4 —————→ 2H3BO3 + 3SO2

         oxidation

2B  +  6HNO3    −————-→ 2H3BO3 + 6NO2

• Reaction with Mg and Ca : 

3Mg +  2B  →  Mg3B2 

3Ca   +  2B   →  Ca3B2

On repetitive hydrolysis, Mg3B2 yields diborane. 

                                 hydrolysis

Mg3B2 +  6HCl    −−−−−−→    3MgCl2 + B2H6 ; B2H6  +  6H2O  →  2H3BO3   +  6H2 

• Reducing properties :

3SiO2 + 4B    →  2B2O3   +   3Si 

3CO2 + 4B   →  2B2O3   +    3C

It decays in steam, liberating hydrogen gas. 

2B + 3H2O (steam)  →   B2O3  + 3H2

Uses: Boron absorbs neutrons and hence is used in reactor rods for atomic reaction control. It is also used in the manufacture of high-impact steel. 

Compounds Of Boron:

• Boron trioxide (B2O3) 

Preparation:

               100℃                  160℃                       red heat

H3BO3   −−−−→   HBO2   −−−−→   H2B4O7   −−−−−→    B2O3

Properties:

1. B2O3 is a slightly acidic oxide. It forms borates with bases.

3Na2O   +  B2O3    →   2Na3BO3    (sodium orthoborate)

• It forms orthoboric acid upon slow reaction with water and forms coloured compounds with transition metal salts. 

H2O  +  B2O3  →  2HBO2 ;  HBO2  + H2O  →  H3BO3 

B2O3  +  P2O5  ⇌  2BPO4

1. Orthoboric Acid ( H3BO3 )

An oxyacid of Boron. 

Preparation : 

1. Concentrated Borax solution is treated with H2SO4  to precipitate orthoboric acid

Na2B4O7   +    H2SO4   +  5H2O  →  Na2SO4    +   4H3BO3↓

1. Powdered colemanite is suspended in water and surplus SO2 is filtered. H3BO3 is formed after chilling the filtrate. 

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

Properties :

• A weak monobasic acid and the boron atom eliminate OH– from water molecules thus completing its octet in an aqueous solution
• The properties shown are similar to lewis acids and behave as a strong acid upon the addition of glycol or glycerol to its aqueous solution. 
• Production of metaboric acid (HBO2) first upon heating followed by boron trioxide. 
• Physical properties include less solubility in cold water than in hot water and oily texture upon touch. It shows planar connecting units of BO3 in a multilayer structure with hydrogen bonds. 

Test For Borate Radical: 

Treating borate with ethyl alcohol creates a green-edged flame. 

H3BO3  + 3C2H5OH  →  B(OC2H5)3  + 3H2O

Uses:

• The water solution of ethyl borate is used to cleanse the eyes.
• Used as an antiseptic.
• Used in glass, enamel and pottery industries

1. Borax (Na2B4O7 • 10H2O)

Preparation:

1. Colemanite, along with a solution of Na2CO3, allows white crystals of borax to precipitate. NaBO2  changes to Na2B4O7  in this reaction. 
2. The action of Na2CO3 on orthoboric acid produces borax. 

Properties:

1. A white powder is more soluble in hot water than in cold water. 
2. The aqueous solution of borax is alkaline, as it can hydrolyse weak acid H3BO3 and strong alkali NaOH
3. Borax powder expands when heated as there is a loss of water as steam. Clear borax beads are formed at 740℃. 
4. The action of acids like HCl forms boric acid
5. Borax first loses its water molecules and expands when heated.
6. Further heating transforms it into a transparent liquid that eventually hardens into a glass-like substance known as a borax bead.
7. Borax bead test is performed to find the metallic ions (cations) in salts

1. Diborane (B2H6)

Binary compounds of B with H are called boranes. Two types of series of boranes are BnHn+4 and BnHn+6

Preparation : 

                                             ether

1. 4BF3 +  3LiAlH4     −−−→   2B2H6 +  3LiF  +  3AlF3 

                                           silent electric discharge

1. 2BCl3  +   6H2(excess)    −−−−−−−−−−−−−−−− −→   B2H6   +    6HCl 

                         ether

1. 8BF3  +  6LiH   −−−−−→    B2H6  +   6LiBF4 

                                      ether

1. 2NaBH4 + I2   −−−−−→  B2H6 + 2NaI + H2

                                                            450K

1. 3NaBH4  +  4BF3    −−−−−−→  3NaBF4  +  2B2H6

                                     ether

1. 2NaBH4   +  H2SO4     −−−−−→    B2H6  +   2H2 + Na2SO4

1. 2NaBH4       +   2H3PO4    −−−−− →  B2H6  +    2H2    +   2NaH2PO4

                                                       450K

1. 2BF3    +    6NaH    −−→   B2H6 +   6NaF (Industrial method)

Properties :

1. A colourless gas with a boiling point of 183K, diborane quickly decomposes by water and forms H3BO3  and H2
2. The reaction of diborane with air is a combustion reaction, with the release of heat. 
3. Diborane undergoes pyrolysis above 375K and results in a mixture of boranes. 
4. Diborane reacts with alkenes and alkynes in ether solvents to form organoboranes. This is called hydroboration. 
5. Certain cleavage reactions also take place with diborane. 

Aluminium

Extraction (Hall-Heroult Process):

Aluminium derived from the bauxite ore is purified by Bayer’s method. Combination and fusion of Al2O3 with the ore are treated with Na3AlF6 and CaF2 for an electrolytic reduction. Purification is done by Hoppe’s method. 

Properties :

• A silvery metal with 2.7 g/cc density and a melting point of 660℃ 

1. Dry air shows no effect, wet air creates a thin coating of Al2O3
2. Halogens react with the aluminium to form anhydrous AlX3
3. Reaction with concentrated NaOH releases H2 gas and sodium aluminate solution
4. Reacts with both H2SO4 and HCl  but not with HNO3
5. N2 gas is passed over Al to form AlN
6. As a result of its highly negative redox potential, aluminium produces hydrogen gas when it interacts with water.
7. Al when added to HgCl2 releases Mercury.
8. Less reactive metal oxides are heated w Aluminium to release less reactive metal. 

Uses :

Uses of aluminium include plating pipes, tanks etc to prevent corrosion and manufacturing aluminium cables etc. 

Compounds of Aluminium:

• Aluminium Oxide (Al2O3) 

Alumina is found in bauxite and corundum forms in nature and is usually found in gemstones. 

Preparation:

Alumina is obtained by igniting Al2(SO4)3  or ammonium alum. 

                 Δ

Al2(SO4)3 −→Al2O3 +  3SO3

Properties :

Alumina is an amphoteric white powder in appearance and is a covalent molecule with a positive charge. 

Uses :

1. Extraction of aluminium
2. Production of fake gems
3. Furnace linings fabrication
4. It is a refractory substance

1. Aluminium Chloride (AlCl3 . 6H20 )

 A white deliquescent solid is water soluble and has a covalent bond.

Preparation:

1. Mixing Al with dilute HCl yields AlCl3
2. The action of Cl2 on heated aluminium yields anhydrous AlCl3 
3. Heating alumina with coke and passing chlorine over it yields AlCl3

Properties:

1. Hydrated AlCl3 transforms to Al2O3
2. Upon exposure to air, fumes of HCl are released
3. Anhydrous AlCl3 absorbs NH3 because it is a Lewis acid
4. Sodium aluminate is formed ultimately when NaOH reacts with aqueous AlCl3. This reaction is used to determine the difference between an aluminium salt and salts of Mg, Ca, Sr and Ba.
5. In a reaction with NH4OH, a white precipitate of Al(OH) is obtained. This precipitate is insoluble, unlike reaction with Zn(OH)2
6. Upon dissolution in water, Al(OH)3 – weak base – and HCl – strong acid – are yielded.

Uses :

A catalyst used in Friedel Craft’s processes and petroleum cracking. 

1. Alums [ M2SO4 . M2’ (SO4)3 . 24H2O ]

Alums are double salts that produce metal ions and sulphate ions in water. 

Preparation:

Fusing M2SO4 and M’2(SO4)3 in a 1:1 molar ratio and dissolving the resultant in water, crystallises alum in solution. 

Uses:

Alum is a mordant and a germicide used for water purification, and as a coagulating agent to remove colloidal contaminants. 

(B) Group 14 Elements: The Carbon Family

Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn) and lead (Pb) are part of Group 14 of the periodic table. Carbon has two stable isotopes -12C and 13C. A radioactive isotope of carbon is 14C. 

• Electronic configuration : ns2 np2

• Covalent radius: A rise in covalent radius is followed by a minor increase in radius from Si to Pb due to completely filled d and f orbitals. 

1. 1. Ionisation Enthalpy: Ionisation enthalpy of Group 14 elements is greater than that of members of group 13, as weak shielding effects of adjacent d and f orbitals and an increase in atomic size cause decreases from Si to Ge to Sn and then a rise from Sn to Pb. 

• Electronegativity: These elements are more electronegative due to their smaller size. Nearly identical values are observed from Si to Pb. 

• Physical Properties: Melting point and Boiling point values are greater than group 13 elements. All members in group 14 are solids, with carbon and silicon being non-metals, germanium a metalloid, and tin and lead being soft metals. 

• Chemical Properties:  The most common oxidation states of Group 14 elements are +2 and +4. Due to the presence of d orbital in them, they can surpass their covalence. Valence shell electrons are ns2 and thus show an inability to engage in bonding. 

• Reactivity towards oxygen – Monoxide and Dioxide forms of compounds are seen. 
• Reactivity towards water – No effects are seen on C, Si, or Ge. Tin shows the property of decomposition of water to give H2 gas and O2. 
• Reactivity towards halogen – Halides created by these elements are MX2 and MX4 types of halides. Majorly yield MX4 and covalency, but SnF4 and PbF4 are ionic in nature.

Important Trends And Anomalous Behaviour Of Carbon

Carbon shows anomalous behaviour due to its small size, high electronegativity, higher ionisation enthalpy and its ability to form pπ-pπ multiple bonds with itself, and with other smaller atoms. Carbon atoms can build chains and rings using covalent bonding owing to its property of catenation. 

Allotropes of Carbon

1. Diamond: The hardest substance on the planet, diamond has a crystalline lattice that undergoes sp hybridisation and is connected to four other carbons in a tetrahedral arrangement. 
2. Graphite: Graphite contains each carbon atom in a hexagonal ring structure, with van der Waals forces acting upon the atoms. Electrons are somewhat mobile, allowing the conduction of electricity. 
3. Fullerenes: Heating graphite with inert gases like helium and argon creates fullerenes, which are pure forms of carbon. A common example of fullerenes is Buckminsterfullerene or C60.

Uses of Carbon:

• Structural work like tennis racquets, fishing rods etc is made out of graphite fibres. 
• Also used as electrodes in batteries, graphite is a good conductor. 
• Activated charcoal is used to absorb harmful gases. 

Properties Of Carbon:

1. Carbon forms carbon dioxide upon reacting with oxygen. Carbon monoxide may also be formed. 
2. The carbon and water react when heated sufficiently (i.e., to hundreds of degrees), producing carbon monoxide and hydrogen, also known as water gas.
3. If dilute acids like HNO3 and H2SO4  are acted upon by carbon, they get reduced to water and other gases like NO2 and CO2.

Oxides Of Carbon:

• Carbon Dioxide (CO2 )

Preparation:

• CO2 is formed by the action of weak HCl on marble chips
• In the manufacturing of quicklime and fermentation operations, it is generated as a by-product.

Properties:

• A heavy gas, at normal temperature and pressure, CO2 is odourless and colourless. Faster dissolution at higher temperatures urges the mixing of CO2 in sodas and fizzy drinks.
• At its critical temperature, 31.1℃ it is easily liquified and can be used as a fire extinguisher. 
• It is the anhydride of carbonic acid that ionises as a weak dibasic acid. H2CO3 / HCO3– is used in maintaining blood pH in the 7.26 to 7.42 range. 
• It performs the limewater test and gives a positive result 
• It is a crucial ingredient in the process of photosynthesis of green plants.
• Utilised for urea production.

• Carbon Monoxide (Co):

Preparation:

• Oxidation of carbonaceous material by air or oxygen produces CO
• CO2 reduction using red hot carbon also produces CO
• Dehydrating methanoic acid and mixing it with conc.  H2SO4
• Steam passed over heated coke also produces CO

Properties:

• A colourless, odourless gas which produces a blue flame on burning, carbon monoxide is very toxic to us. 
• It readily binds to haemoglobin and causes issues in breathing.
• It is a strong reducing agent used in iron and nickel extraction.
• It reacts with sulphur to form carbonyl sulphide, and with chlorine in the presence of light to form phosgene. 
• Many such properties of CO are seen in Class 11 Chemistry Chapter 11 based on the NCERT curriculum. 

• Carbon Suboxide (C3O2) 

• Carbon suboxide is the anhydride of malonic acid which is produced by dehydrating it with phosphorus pentoxide. Upon heating to 200℃, it decomposes into CO2 and C.
• It has a linear structure O = C =C = C = O.

• Carbonates (CO32- ) and Bicarbonates (HCO3–)

• Carbonic acid produces these salts when hydrogen atoms are replaced from H2CO3.

Preparation:

1. Addition of NaOH to CO2 leads to formation of NaHCO3
2. Precipitation of BaCO3 from a mixture of BaCl2 and Na2CO3

1. Carbides : 

Carbon-based binary compounds with elements that have less or equivalent electronegativity.

The three groups of carbides are :

1. Ionic carbides – Further classified into Methanides, Acetylides, Allylides.
2. Covalent carbides (CH4 , CO2 , CS2 )
3. Interstitial  or metallic carbides 

1. Carborundum Sic.

Preparation:

SiO2  + 3C SiC + 2CO  (At a furnace, 2000℃)

Properties: 

• It is a brittle material
• Does not melt when heated, instead it decomposes into its components.
• Acids show no effect 
• At high temperatures, each atom is sp3 hybridised thus gets surrounded by 4 atoms of opposite kind.

Silicon:

Found in nature as sand, flint, quartz etc. Silicon is the second most prevalent element on earth. 

Preparation:

1. Silica (SiO2)  is reduced with carbon in an electric furnace to form pure silicon. 
2. Zone refining process means to have very pure silicon from chemically pure silicon.

SiO2  + 2Mg  2MgO + Si

Properties:

1. Silicon is inert in its bulk state
2. In a powdered state, alkalis and halogens react with silicon. 
3. Silicon is unaffected by acids, unless it is hydrofluoric acid, to produce hexafluorosilicic acid. 

Si(s)  +  6HF(g)−→ SiF6(aq)  +  2H2(g)                                                           [In the presence of H2]

1. Na2CO3 + Si →  Na2SiO3 + C                             [In the presence of heat]

Compounds of Silicon:

• Silicon dioxide (SiO2) :A three-dimensional covalent solid network wherein each silicon atom is linked to four oxygen atoms in tetrahedral structure. 

It is resistant to halogens, dihydrogen and most acids and metals. It reacts with HF and NaOH.

SiO2 + NaOH → Na2SiO3 + H2O

SiO2 + 4HF → SiF4 + 2H2O

Silicates: These are binary silicon-oxygen compounds that also contain additional metals in their atomic structures. The linkage present in silicates can be considered 50% ionic and 50% covalent. 

Classification Of Silicates:

The classification of silicates is as follows :

1. Orthosilicates – Distinct units of [SiO4]4 without sharing of corners. 
2. Pyrosilicates – Two tetrahedral units are linked by sharing one oxygen at a single corner. Shows [Si2O7]6- units.
3. Cyclic silicates – Two oxygen atoms per tetrahedron are shared to form closed rings having general formula (SiO32-)n
4. Chain silicates – Two corners of each tetrahedron are linked together by shared oxygen to form a lengthy tetrahedron chain. 
5. Three-dimensional sheet silicates – Sharing of all four oxygen atoms by neighbouring SiO44 tetrahedral units. For Example, quartz, zeolite

Silicones: These are synthetic organosilicon compounds having an alkyl or aryl group in their general formula, (R2SiO)n . These compounds show Si -O-Si bonds. 

Silicones can be made from 

1. R3SiCl 
2. R2SiCl2
3. RSiCl3

Some uses of silicones are to make products with physicalities similar to rubbers, oils, and resins since they are thermally stable and their viscosity changes very little as a function of temperature. 

Tin and Lead:

Compounds of Tin:

• Stannous oxide (SnO)

Preparation:

1. Sn(OH)2 is heated in absence of air to give SnO
2. Heating stannous oxalate yields SnO in the absence of air.

Properties:

1. SnO is insoluble in water, and is an amphoteric dark grey or black solid oxide. 
2. Upon dissolution in acids, it forms stannous salts. 
3. SnO in a heated solution of NaOH creates sodium stannites and water. Only aqueous solutions contain stannites.

Uses :

Stannous chloride and stannous sulphate are made from it.

• Stannous Chloride (SnCl2 – H2O) 

A colourless solid whose solution becomes milky owing to the hydrolysis of Sn(OH)2 and HCl. It is soluble in ether and alcohol.

Preparation :

1. Sn + 2HCL (conc.) → SnCl2(aq) + H2 
2. SnO + 2HCl → SnCl2(aq) + H2O

Properties:

1. With the addition of SnCl2  to HgCl2, a white precipitate of Hg2Cl2 precipitates to form a black substance. 
2. FeCl3 is reduced to ferrous chloride and SnCl4.
3. Hydrolysis of SnCl2 in water gives a white precipitate of Sn(OH)2

Uses:

 It is a reducing agent in the dye industry, also necessary for mercuric salt testing.

• Stannic Oxide (SnO2)

Preparation:

1. Sn + O2  → SnO2
2. Sn + 4HNO3  → H2SnO3 + 4NO2  ↑   +H2O

Properties:

1. It is a white solid with low acidity
2. It does not dissolve in water but forms stannic sulphate when dissolved in H2SO4 
3. It dissolves in alkali metal stannate solution when mixed with concentrated alkalis. 

• Stannic Chloride (SnCl4) 

Preparation:

1. Sn  + 2Cl2 → SnCl4
2. SnCl2    +   Cl2   →  SnCl4

Properties:

1. It is a colourless flammable liquid, which acts as a Lewis acid.
2. It converts to hydrated stannic chlorides upon absorbing moisture. 
3. It generates HCl upon hydrolysis in water. 
4. The boiling point of SnCl4  is 114℃

Uses: 

It is used for preparing stannic compounds.

Compounds Of Lead:

• Litharge (PbO) 

It is an unstable yellow compound obtained after heating Pb at 180℃. It is amphoteric in nature. 

• Lead dioxide (PbO2)

This compound is a dioxide, not a peroxide. 

Preparation:

1. PbO  +  NaOCl−→ PbO2(insoluble)  +  NaCl                     [In the presence of heat]
2. Pb3O4 +  4HNO3(dilute)→2Pb(NO3)2 +  PbO2

Properties:

1. An insoluble chocolate-coloured powder in appearance 
2. It oxidises HCl to Cl2 
3. It gets dissolved in conc. NaOH solution
4. It oxidises Mn to Permanganic acid ( 2HMnO4)

Uses:

It is used to make the igniting surface of match boxes where KMnO4 is the match powder. 

• Red Lead (Pb3O4) 

Preparation:

6PbO  + O2     −−−−→   2Pb3O4    [  At 450oC ]

Properties:

1. When heated with conc. HNO3 it produces a precipitate of PbO2
2. Above 550oC, it undergoes decomposition into PbO and liberates oxygen gas.
3. Oxidation of conc. HCl to chlorine takes place.
4. Heating with conc. H2SO4 , Pb3O4 evolves oxygen. 

Uses:

It is used in red paint and for making special lead cement. 

• Lead Chloride (PbCl2)

Preparation: 

1. Pb(NO3)2 + 2HCl→PbCl2↓+2HNO3 
2. Pb(NO3)2 +2NaCl→PbCl2↓+2NaNO3
3. Pb(CH3COO)2 + 2HCl→PbCl2↓+2CH3COOH
4. PbO+2HCl→PbCl2↓+H2O
5. Pb(OH)2+2HCl→PbCl2↓+2H2O
6. Pb(OH)2⋅PbCO3+ 4HCl→2PbCl2↓+CO2↑

Properties:

It is a white crystalline substance, insoluble in cold water but forms a complex ion when added to strong HCl. 

2HCl  +  PbCl2  ⇆   H2PbCl4  (chloro plumbous acid)

Uses:

Used in paint pigment manufacturing.

• Lead Tetrachloride (PbCl4)

Preparation:

1. Adding PbO2 to cold conc. HCl
2. NH4Cl  is introduced to a chloro plumbic acid solution, to produce a yellow precipitate of ammonium chloroplumbate. 
3. By action of Cl2  on PbCl 2
4. (NH4)2PbCl6  + H2SO4    →  PbCl4    +  (NH4)2SO4    + 2H

Properties:

1. Precipitation of PbO2 occurs with rapid hydrolysis in water.
2. It freezes at -10℃ and dissolves in organic solvents.
3. It is a yellow oily liquid in physical appearance. 

About The p-Block Elements

The Chapter on p-block elements covers majorly Group-13 and Group-14 of the modern periodic table and explores the trends in these groups. The physical and chemical properties of the elements are observed which can help understand their practical applications in day-to-day chemical products. Some concepts like electronegativity, ionisation enthalpy and the shielding by d and f electrons are important from an exam point of view.  Class 11 Chemistry Revision Notes for Chapter 11 – The p-Block Elements provide a concise overview of all these important topics, which is crucial for your exam preparation strategy. 

Sub-Topics Covered Under The p-Block Elements :

1. Group 13 Elements – The Boron Family

• Ionisation Enthalpy
• Atomic Radii
• Electronegativity
• Electronic Configuration
• Physical and Chemical Properties

1. Important Trends & Anomalous Properties of Boron

1. Some Important Compounds of Boron

• Borax

• Diborane, B2H6

• Orthoboric Acid

1. Uses of Boron & Aluminium & Their Compounds

1. Group 14 Elements: The Carbon Family

• Covalent Radius

• Electronic Configuration

• Ionisation Enthalpy

• Physical and Chemical properties

• Electronegativity

1. Important Trends & Anomalous Behaviour of Carbon
2. Allotropes of Carbon

• Graphite

• Diamond

• Fullerenes

• Uses of Carbon

1. Some Important Compounds of Carbon & Silicon

• Carbon Monoxide 

• Carbon Dioxide
• Silicon Dioxide, SiO2 
• Silicates
• Silicones
• Zeolites