Chapter 11 The p Block Elements

(i) B to Tl- they belong to group 13 elements with electronic configuration as ns² np¹. The atoms of these elements have 3 valence electrons, two in s subshell and one in p subshell, therefore all these elements show maximum of + 3 oxidation state. Boron shows only +3 oxidation state in its compound and other elements also show +1 oxidation state. The +1 oxidation state becomes more stable as we move down the group from boron to thallium. The +1 oxidation state is more stable than +3 oxidation state because of inert pair effect .In case of last element, after removal of one electron from p orbital, the remaining ns² electrons behave like stable noble gases and do not take part in compound formation. This reluctance of the s electron pair to take part in chemical combination is called inert pair effect. The two electrons present in the s-shell are strongly attracted by the nucleus and do not participate in bonding. This inert pair effect becomes more and more prominent on moving down the group. Hence, Ga (+1) is unstable, In (+1) is fairly stable and Tl (+1) is very stable.

Group 13 element	Oxidation state
B	+3
Al	+3
Ga, In, Tl	+1, +3

The stability of the +3 oxidation state decreases on moving down the group.

(ii) C to Pb- they belongs to group 14 in periodic table and are also called carbon family with electronic configuration of ns² np². Therefore, the most common oxidation state exhibited by them should be +4. However, the +2 oxidation state becomes more and more common on moving down the group. C and Si mostly show the +4 oxidation state. On moving down the group, the higher oxidation state becomes less stable. This is because of the inert pair effect. Thus, although Ge, Sn, and Pb show both the +2 and + 4 states, the stability of the lower oxidation state increases and that of the higher oxidation state decreases on moving down the group.

Group14 element	Oxidation state
C	+4
Si	+4
Ge, Sn, Pb	+2, +4

 

Boron and thallium belong to group 13 of the periodic table. In this group, the +1 oxidation state becomes more stable on moving down the group. BCl₃ is more stable than TlCl₃ because the +3 oxidation state of B is more stable than the +3 oxidation state of Tl. In Tl, the +3 state is highly oxidising and it reverts back to more stable +1 state.

The electronic configuration of boron is ns² np¹. It has three electrons in its valence shell. Thus, it can form only three covalent bonds. This means that there are only six electrons around boron and its octet remains incomplete. When one atom of boron combines with three fluorine atoms, its octet remains incomplete. Hence, boron trifluoride remains electron-deficient and acts as a Lewis acid.

Being a Lewis acid, BCl₃ readily undergoes hydrolysis. Boric acid is formed as a result.

BCl₃  +  3H₂O   →    3HCl   +  B(OH)₃

CCl₄ completely resists hydrolysis. Carbon does not have any vacant orbital. Hence, it cannot accept electrons from water to form an intermediate. When CCl₄ and water are mixed, they form separate layers.

CCl4   +  H₂O   →   No reaction

Boric acid is not a protic acid. It is a weak monobasic acid, behaving as a Lewis acid.

B(OH)₃   +   2HOH   →   [B(OH)₄]^-  +  H₃O⁺

It behaves as an acid by accepting a pair of electrons from ^-OH ion.

On heating orthoboric acid (H₃BO₃) at 370 K or above, it changes to metaboric acid (HBO₂). On further heating, this yields boric oxide B₂O₃.

(i) BF₃

As a result of its small size and high electronegativity, boron tends to form monomeric covalent halides. These halides have a planar triangular geometry. This triangular shape is formed by the overlapping of three sp² hybridised orbitals of boron with sp orbitals of three halogen atoms. Boron is sp² hybridised in BF₃.

(ii) BH₄^-

Boron-hydride ion (BH₄^-) is formed by the sp³ hybridisation of boron orbitals. Therefore, it is a tetrahedral structure.

A substance is called amphoteric if it displays characteristics of both acids and bases. Aluminium dissolves in both acids and bases, showing its amphoteric behaviour.

(i) 2AI_(s) + 6HCI(aq)  →   2AI³⁺_(aq)   +   6Cl^-_(aq)  +  3H_2(g)

(ii) 2AI_(s) + 2NaOH(aq) + 6H₂O_(l)   →  2Na⁺[Al(OH)₄]^-_(aq)  + 3H_2(g)

In an electron-deficient compound, the octet of electrons is not complete, i.e., the central metal atom has an incomplete octet. Therefore, it needs electrons to complete its octet.

(i) BCl₃

BCl₃ is an appropriate example of an electron-deficient compound. B has 3 valence electrons. After forming three covalent bonds with chlorine, number of electrons around it increases to 6. However, it still lacks two electrons to complete its octet.

(ii) SiCl₄

The electronic configuration of silicon is ns² np² . This indicates that it has four valence electrons. After it forms four covalent bonds with four chlorine atoms, its electron count increases to eight. Thus, SiCl₄ is not an electron-deficient compound.

There are only two resonating structures for the bicarbonate ion.

The state of hybridisation of carbon in:

(a) CO^2-₃

C in CO^2-₃  is sp² hybridised and is bonded to three oxygen atoms.

(b) Diamond

Each carbon atom in diamond is sp3 hybridised and is bound to four other carbon atoms.

(c) Graphite

Each carbon atom in graphite is sp2 hybridised and is bound to three other carbon atoms.

 

Diamond	Graphite
It has a crystalline lattice.	It has a layered structure.
In diamond, each carbon atom is sp3 hybridised and is bonded to four other carbon atoms through a   σ  bond.	In graphite, each carbon atom is sp2 hybridised and is bonded to three other carbon atoms through a σ bond. The fourth electron forms a π bond.
It is made up of tetrahedral units.	It has a planar geometry.
The C-C bond length in diamond is 154 pm.	The C-C bond length in graphite is 141.5 pm.
It has a rigid covalent bond network which is difficult to break.	It is quite soft and its layers can be separated easily.
It acts as an electrical insulator.	It is a good conductor of electricity.

 

(a) Lead belongs to group 14 of the periodic table. The two oxidation states displayed by this group is +2 and +4. On moving down the group, +2 oxidation state becomes more stable and +4 oxidation state becomes less stable. This is because of the inert pair effect. Hence, PbCl₄ is much less stable than PbCl₂. However, the formation of PbCl₄ takes place when chlorine gas is bubbled through a saturated solution of PlCl₂.

PbCl _2(s) + Cl_2(g)  →  PbCl_4(l)

(b) On moving down group IV, the higher oxidation state becomes unstable because of the inert pair effect. Pb (IV) is highly unstable and when heated, it reduces to Pb (II).

PbCl_4(l)   →  PbCl _2(s) + Cl_2(g) 

(c) Lead is known not to form PbI₄. Pb (+4) is oxidising in nature and I^- is reducing in nature. A combination of Pb(IV) and iodide ion is not stable. Iodide ion is strongly reducing in nature. Pb(IV) oxidises I^- to I²and itself gets reduced to Pb(II).

Pbl₄ →  Pbl ₂  + l₂

The B-F bond length in BF₃ is shorter than the B-F bond length in BF^-₄ . BF₃ BF3 is an electron-deficient species. With a vacant p-orbital on boron, the fluorine and boron atoms undergo pπ-pπ back bonding to remove this deficiency. This imparts a double-bond character to the B-F bond.

This double-bond character causes the bond length to shorten in BF₃(130 pm). However, when BF₃ coordinates with the fluoride ion, a change in hybridisation from sp² (in BF₃) to sp³ (in BF^-₄) occurs. Boron now forms 4 σ  bonds and the double-bond character is lost. This accounts for a B-F bond length of 143 pm in BF^-₄ ion.

As a result, the difference in the electronegativities of B and Cl, the B-Cl bond is polar in nature. However, the BCl₃ molecule is non-polar. This is because BCl₃ is trigonal planar in shape. It is a symmetrical molecule. Hence, the respective dipole-moments of the B-Cl bond cancel each other, therefore causing a zero-dipole moment.

        

Hydrogen fluoride (HF) is a covalent compound and has very strong intermolecular hydrogen-bonding. Thus, it does not provide ions and aluminium fluoride (AlF) does not dissolve in it. Sodium fluoride (NaF) is an ionic compound and when it is added to the mixture, AlF dissolves. This is because of the availability of free F^-. The reaction involved in the process is:

AlF₃  +  3 NaF  →   Na₃[AIF₆]

                               Sodium hexafluoroaluminate(lll)

When boron trifluoride (BF₃) is added to the solution, aluminium fluoride precipitates out of the solution. This happens because tendency of boron to form complexes is much more than that of aluminium. Therefore, when BF₃ is added to the solution, B replaces Al from the complexes according to the following reaction:

Na₃[AIF₆]   +  3 BF₃   →  3 Na[BF₄]  + AlF₃ 

Carbon monoxide is highly-poisonous because of its ability to form a complex with haemoglobin. The CO-Hb complex is more stable than the O_2-Hb complex. The former prevents Hb from binding with oxygen. Thus, a person dies because of suffocation on not receiving oxygen. It is found that the CO-Hb complex is about 300 times more stable than the O_2-Hb complex.

Carbon dioxide is a very essential gas for our survival. However, an increased content of CO₂ in atmosphere causes a serious threat. An increment in the combustion of fossil fuels, decomposition of limestone, and decrease in the number of trees has led to greater levels of carbon dioxide. Carbon dioxide has property of trapping the heat provided by sunrays. Higher the level of carbon dioxide, higher is the amount of heat trapped. This results in an increase in the atmospheric temperature, therefore causing global warming.

(a) Diborane B₂H₆ is an electron-deficient compound. B₂H₆ has only 12 electrons - 6 e^- from 6 H atoms and 3 e-each from 2 B atoms. Thus, after combining with 3 H atoms, none of the boron atoms has any electrons left. X-ray diffraction studies have shown the structure of diborane as:

 

2 boron and 4 terminal hydrogen atoms (Ht) lie in one plane, while the other two bridging hydrogen atoms (Hb) lie in a plane perpendicular to the plane of boron atoms. Again, of the two bridging hydrogen atoms, one H atom lies above the plane and the other lies below the plane. The terminal bonds are regular two-centre two-electron (2c - 2e^-) bonds, while the two bridging (B-H-B) bonds are three-centre two-electron (3c - 2e^-) bonds.

 

(b) Boric acid Boric acid has a layered structure. Each planar BO₃ unit is linked to one another through H atoms. The H atoms form a covalent bond with a BO3 unit, while a hydrogen bond is formed with another BO₃ unit. In the given figure, the dotted lines represent hydrogen bonds.

(a) When heated, borax undergoes various transitions. It first loses water molecules and swells. Then, it turns into a transparent liquid, solidifying to form a glass-like material called borax bead.

Na₂B₄O₇.10H₂O   →   Na₂B₄O₇   →   2NaBO₂  +  B₂O₃

   Borax                      Sodium Metaborate     Boric anhydride

 

(b) When boric acid is added to water, it accepts electrons from -OH ion.

B(OH)₃ + 2HOH   →   [B(OH)₄]^-   +   H₃O⁺

 

(c) Al reacts with dilute NaOH to form sodium tetrahydroxoaluminate(III). Hydrogen gas is liberated in the process.

2AI_(s) + 2NaOH(aq) + 6H₂O_(l)   →  2Na⁺[Al(OH)₄]^-_(aq)  + 3H_2(g)

 

(d) BF₃ (a Lewis acid) reacts with NH₃ (a Lewis base) to form an adduct. This results in a complete octet around B in BF₃.

F₃B   +  :NH₃   →   F₃B   ←:NH₃

 

(a) When silicon reacts with methyl chloride in the presence of copper (catalyst) and at a temperature of about 537 K, a class of organo silicon polymers called methyl-substituted chlorosilanes (MeSiCl₃, Me₂SiCl₂, Me₃SiCl, and Me₄Si) are formed.

 

(b) When silicon dioxide (SiO₂) is heated with hydrogen fluoride (HF), it forms silicon tetrafluoride (SiF₄). Usually, the Si-O bond is a strong bond and it resists any attack by halogens and most acids, even at a high temperature. However, it is attacked by HF.

SiO₂   +   4HF   →   SiF₄   +   2H₂O

The SiF₄ formed in this reaction can further react with HF to form hydrofluorosilicic acid.

SiF₄   +   2HF   →    H₂SiF₆

 

(c) When CO reacts with ZnO, it reduces ZnO to Zn. CO acts as a reducing agent.

ZnO_(s)  +  CO_(g)  →  Zn_(s)   +   CO_2(g)

 

(d) When hydrated alumina is added to sodium hydroxide, the former dissolves in the latter because of the formation of sodium meta-aluminate.

Al₂O₃.2H₂O   +   2NaOH    →    2NaAlO₂   +   3H₂O

(i) Concentrated HNO₃ can be stored and transported in aluminium containers as it reacts with aluminium to form a thin protective oxide layer on the aluminium surface. This oxide layer renders aluminium passive.

(ii) Sodium hydroxide and aluminium react to form sodium tetrahydroxoaluminate(III) and hydrogen gas. The pressure of the produced hydrogen gas is used to open blocked drains.

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

(iii) Graphite has layered structure and different layers of graphite are bonded to each other by weak van der Waals' forces. These layers can slide over each other. Graphite is soft and slippery. Therefore, graphite can be used as a lubricant.

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

(v) Aluminium has a high tensile strength and is very light in weight. It can also be alloyed with various metals such as Cu, Mn, Mg, Si, and Zn. It is very malleable and ductile. Therefore, it is used in making aircraft bodies.

(vi) The oxygen present in water reacts with aluminium to form thin layer of aluminium oxide. This layer prevents aluminium from further reaction. However, when water is kept in an aluminium vessel for long periods of time, some amount of aluminium oxide may dissolve in water. As aluminium ions are harmful, water should not be stored in aluminium vessels overnight.

(vii) Silver, copper, and aluminium are among the best conductors of electricity. Silver is an expensive metal and silver wires are very expensive. Copper is quite expensive and is also very heavy. Aluminium is very ductile metal. Thus, aluminium is used in making wires for electrical conduction.

Ionisation enthalpy of carbon (the first element of group 14) is very high (1086 kJ/mol). This is expected because of its small size. However, on moving down the group to silicon, there is sharp decrease in the enthalpy (786 kJ). This is because of an appreciable increase in the atomic sizes of elements on moving down the group.

                                                     Atomic Radius (in pm)

Aluminium	143
Gallium	135

Although Ga has one shell more than Al, its size is lesser than Al. This is because of the poor shielding effect of the 3d-electrons. The shielding effect of d-electrons is very poor and effective nuclear charge experienced by the valence electrons in gallium is much more than that of Al.

Allotropy is the existence of an element in more than one form, having the same chemical properties but different physical properties. The various forms of an element are called allotropes.

The rigid 3-D structure of diamond makes it a very hard substance. In fact, diamond is one of the hardest naturally-occurring substances. It is used as an abrasive and for cutting tools.

 

Graphite:

It has sp² hybridised carbon, arranged in the form of layers. These layers are held together by weak van der Walls' forces. These layers can slide over each other, making graphite soft and slippery. Therefore, it is used as a lubricant.

(1) CO = Neutral

 

(2) B₂O₃ = Acidic

Being acidic, it reacts with bases to form salts. It reacts with NaOH to form sodium metaborate.

B₂O₃  +  2NaOH   →   2NaBO₂  +   H₂O

 

(3) SiO₂ = Acidic

Being acidic, it reacts with bases to form salts. It reacts with NaOH to form sodium silicate.

SiO₂  +  2NaOH   →   2Na₂SiO₃  +   H₂O

 

(4) CO₂ = Acidic

Being acidic, it reacts with bases to form salts. It reacts with NaOH to form sodium carbonate.

CO₂  +  2NaOH   →   2Na₂CO₃  +   H₂O

 

(5) Al₂O₃ = Amphoteric

Amphoteric substances react with both acids and bases. Al₂O₃ reacts with both NaOH and H₂SO₄.

Al₂O₃ +  2NaOH   →   NaAIO₂ 

Al₂O₃ + 3H₂SO₄    →   Al₂(SO₄)₃ +  3H₂O

 

(6) PbO₂ = Amphoteric

Amphoteric substances react with both acids and bases. PbO₂ reacts with both NaOH and H₂SO₄.

PbO₂  +  2NaOH   →   Na₂PbO₃  +   H₂O

PbO₂ + 2H₂SO₄    →   2PbSO₄ +  2H₂O +O₂

 

(7) Tl₂O₃ = Basic

Being basic, it reacts with acids to form salts. It reacts with HCl to form thallium chloride.

Tl₂O₃  +  6HCl  →   2TlCl₃ + 3H₂O

Thallium belongs to group 13 of periodic table. The most common oxidation state for this group is +3. However, heavier members of this group also display the +1 oxidation state. This happens because of the inert pair effect. Aluminium displays the +3 oxidation state and alkali metals display the +1 oxidation state. Thallium displays both the oxidation states.

Thallium, like aluminium, forms compounds such as TlCl₃ and Tl₂O₃. It resembles alkali metals in compounds Tl₂O and TlCl.

The given metal X gives a white precipitate with sodium hydroxide and this precipitate dissolves in excess of sodium hydroxide. Hence, X must be aluminium.

The white precipitate (compound A) obtained is aluminium hydroxide. The compound B formed when an excess of the base is added is sodium tetrahydroxoaluminate(III).

2Al  +  3NaOH  →  Al (OH)₃ ↓  +  3Na⁺

Al(OH)₃  +  NaOH   →  Na⁺ [Al(OH)4]^-

^(A)

Now, when dilute hydrochloric acid is added to aluminium hydroxide, aluminium chloride (compound C) is obtained.

Al(OH)₃  +  3HCl   →   AlCl₃  +  3H₂O

(A)                              (C)

Also, when compound A is heated strongly, it gives compound D. This compound is used to extract metal X. Aluminium metal is extracted from alumina. Hence, compound D must be alumina.

2Al(OH)₃   →  Al₂O₃  +  3H₂O

(A)                    (D)

(a) Inert pair effect: Down the group, the tendency of s-block electrons to participate in chemical bonding decreases. This effect is known as inert pair effect. In case of group 13 elements, the electronic configuration is ns² np¹ and their group valency is +3. However, on moving down the group, the +1 oxidation state becomes more stable. This happens because of the poor shielding of the ns² electrons by the d- and f- electrons. As a result of poor shielding, the ns² electrons are held tightly by the nucleus and so they cannot participate in chemical bonding.

(b) Allotropy: Allotropy is existence of an element in more than one form, having the same chemical properties but different physical properties. The various forms of an element are called allotropes. For example, carbon exists in three allotropic forms: diamond, graphite, and fullerenes.

(c) Catenation: The atoms of some elements (such as carbon) can link with one another through strong covalent bonds to form long chains. This property is known as catenation. It is most common in carbon and quite significant in Si and S.

The given salt is alkaline to litmus. Therefore, Xis a salt of a strong base and a weak acid. Also, when Xis strongly heated, it swells to form substance Y. Therefore, Xmust be borax.

When borax is heated, it loses water and swells to form sodium metaborate. When heating is continued, it solidifies to form a glassy material Y. Hence, Ymust be a mixture of sodium metaborate and boric anhydride.

 

Caron dioxide

In the laboratory, CO₂ can be prepared by the action of dilute hydrochloric acid on calcium carbonate. The reaction involved is as follows:

CaCO₃ + 2HCl_(aq)   →   CaCl_2(aq)  +  CO _2(g)   +  H2O _(l)

CO₂ is commercially prepared by heating limestone. The reaction involved is as follows:

CaCO₃     CaO  + CO₂ ↑

Caron monoxide

In the laboratory, CO is prepared by the dehydration of formic acid with conc. H₂SO₄, at 373 K. The reaction involved is as follows:

(c) Borax is a salt of a strong base (NaOH) and a weak acid (H₃BO₃). It is, therefore, basic in nature.

(b) Boric acid is polymeric because of the presence of hydrogen bonds. In the given figure, the dotted lines represent hydrogen bonds.

(c) Boron in diborane is sp³ hybridised.

(b) Graphite is thermodynamically the most stable form of carbon.

(b)The elements of group 14 have 4 valence electrons. Therefore, the oxidation state of group is +4. However, as a result of the inert pair effect, the lower oxidation state becomes more stable and the higher oxidation state becomes less stable. Therefore, this group exhibits +4 and +2 oxidation states.

Group 14 element	Oxidation state
C	+4
Si	+4
Ge, Sn , Pb	+2 , +4