CBSE Class 11 Chemistry Notes For Chapter 11 Some P Block Elements

CBSE Class 11 Chemistry Notes For Chapter 11 Some P Block Elements Group-13 Elements (Boron Family) Introduction

The valence shell electronic configuration of the elements of group-13 is ns²np¹ where n = 2-6.  It becomes clear from the electronic configurations that boron (B) and aluminium (Al) have noble gas cores, gallium (Ga) and indium (In) have noble gas cores plus 10 d-electrons and thallium (Tl) have noble gas cores plus 14 F -electrons and 10 d-electrons. Thus electronic configuration of the elements of group 13 is more complex compared to those of groups and 2.

This difference in electronic configuration affects the chemistry of the elements of this group.

Electronic configuration of group 13 – elements:

Occurrence Of Group-13 Elements

1. The elements present in group 13 of the periodic table are boron (B), aluminium (Al), gallium (Ga), indium (In) and thallium (Tl). Except for horon, which is a non-metal, all other elements of this group are metals. The non-metallic character of boron is due to its small atomic size, high ionization enthalpy and comparatively high electronegativity.

2. Boron is a fairly rare element which occurs to a very small extent (0.0001% by mass) in the earth’s crust Natural boron consists of two isotopes: ¹⁰B (19%) and ¹¹B (81%). Boron does not occur in a free state as it is highly reactive. It occurs mainly as orthoboric acid & as minerals like

Borax –  Na₂(B₄O₅(OH₄).8H₂O -8H₂O

Kernite – Na₂(B₄O₅(OH₄).2H₂O

Colemanite – Na₂(B₃O₄(OH₃)₂.2H₂O

3. Aluminium is the most abundant metal, and the third most abundant element (8.3%by mass) in the earth’s crust after oxygen (45.5%) and silicon (27.7%). The important minerals of aluminium are:

Bauxite – (Al₂O₃ – 2H₂O)

Cryolite -(Na₃AlF₆)

Orthoclase (feldspar) – KAlSi₃O₈, 

Mica (Muscovite)-  KAl ₂(AlSi₃O₁₀)(F, OH)₂ etc.

4. Gallium, indium and thallium are quite less abundant and occur in traces in sulphide minerals.

5. The highest concentration of Ga (0.1-1%) is found in a rare mineral known as germanite (a sulphide complex of Zn, Cu, Ge and As).

6. Traces of and Tl are available in sulphide ores of zinc and lead respectively

General Trends In Atomic And Physical Properties Of Group-13 Elements

Some important atomic and physical properties of group-13 most elements are given in the following table

Some atomic and physical properties of group-13 elements:

Trends in different atomic and physical properties of group-13 elements with explanations:

Atomic and ionic radii

1. Atomic and ionic radii of group-13 elements are smaller as compared to the corresponding elements of group-2.

Atomic and ionic radii Explanation 1:

On moving from left to right in the periodic table, i.e., on moving from group-2 to group-13in a given period, the magnitude of nuclear charge increases but the new electron is added to the same shell. Since the electrons in the same shell do not screen each other and the effective nuclear charge increases, the outermost electrons experience greater nuclear charge and are pulled more strongly towards the nucleus. As a result, atomic size decreases. The same is true in the case of ionic radius.

2. On moving down the group, both atomic and ionic radii are expected to increase due to the addition of new electronic shells. However, the observed atomic radius of Ga = 135pm is slightly lesser than that of Al = 143pm.

Atomic and ionic radii Explanation 2:

On moving from Al (Z = 13) to Ga (Z = 31), the d-orbitals are filled by electrons. Since the d-orbitals are larger, these intervening electrons in d-orbitals do not screen the nucleus effectively. As a result, the effective nuclear charge experienced by the electrons in Ga is greater than that experienced by the electrons in Al. Hence, the atomic radius of Ga is slightly less than that of Al. The Ionic radii, however, increase regularly on moving down the group

Ionisation enthalpy

1. First ionisation enthalpies (Δ_iH₁) of group-13 elements are lower than the corresponding elements ofgroup-2.

Ionisation enthalpy Explanation 1:

• The first electron, in the case of group-13 elements (ns²np¹), is to be removed from a p-orbital, while in the case of group-2 elements, it is to be removed from an s-orbital.
• Due to greater penetration of the s-orbital, the s-electron is nearer to the nucleus and is more tightly held by the nucleus than a p-electron of the same principal shell.
• The removal of the s-electron requires a greater amount of energy compared to p- the electron and because of this, the values of first ionisation enthalpies (Δ_iH_i) of the elements of group-13 are low as compared to the corresponding elements of group-2.
• The second and third ionisation enthalpies of these elements are, however, quite high because the second and third electrons are to be removed from ns-orbital.

2. On moving down the group from B to Al, the first ionisation enthalpy, (Aÿ) decreases sharply. However, the value of (Δ_iH₁) of Ga is slightly higher than that ofAl, while that ofTl is much higher than that of.

Ionisation enthalpy Explanation 2:

1. The sharp decrease in (Δ_iH₁) value from B to Al is expected because an increase in atomic size and screening effect (caused by to addition of a new shell) outweighs the effect of ‘increased nuclear charge.

2. The element Ga has ten 3d-electrons which do not screen as  N, much as s- and p-electrons. Therefore, due to poor shielding of 3d-electrons, the effective nuclear charge acting on Ga is slightly higher than that on Al.  Due to this, the (Δ_iH₁) value of Ga is slightly higher than that of Al, even though a new shell has been added on going from Al to Ga.

3. The same explanation can be offered on going from Into Tl. Tl has fourteen 4 f-electrons having a very poor screening effect and because of this, there occurs an unexpected increase in tyre effective nuclear charge, for which (Δ_iH₁) of Tl becomes much higher than that of In.

4. The order of (Δ_iH₁) values of group-13 elements is B > Al < Ga > In < Tl. However, this trend is not observed in the (Δ_iH₂) and (Δ_iH₃) values of these elements and this is because once the outermost p-electron is removed, it is not easy to remove the second and third electrons due to a large increase in effective nuclear charge. As expected, the first three ionisation enthalpies of these elements follow the order: Δ_iH₁< Δ_iH₂ < Δ_iH₃

Oxidation states

The atoms of group-13 elements have three valence electrons, two in the s -s-subshell and one in the p -subshell. Therefore, it becomes clear from their electronic configurations that +3 is expected to be the most common oxidation state of these elements. Therefore, the group oxidation state of the group-13 element is +3.

Due to the small size of the boron, the sum of its first three ionisation enthalpies is very high. Therefore, it cannot lose its valence electrons to form B³⁺ ion rather it forms covalent bonds with other atoms, for example, BH₃ or B₂H₆.

The sum of the first three ionisation enthalpies of Al is much lower than that of B so, it can form an Al³⁺ ion, for example, AlCl₃. Al is a highly electropositive metal.

B and Al exhibit only a +3 oxidation state but Ga, In and Tl show a +3 as well as a +1 oxidation state and on moving down the group the stability of the +3 oxidation state decreases while that of the +1 oxidation state progressively increases.

The stability of the +1 oxidation state follows the order:

Al < Ga <In < Tl. In the case of Tl, the +1 oxidation state is very much more stable than the +3 oxidation state.

Oxidation states Explanation:

The stability of the +1 oxidation state that increases down the group can be explained in terms of the inert pair effect. On moving down the group, the tendency of electrons of the valence shell to participate in bond formation decreases due to poor shielding of these electrons from the attraction of the nucleus by the intervening d- and f- electrons.

This reluctance or inertness ofthe s-electrons to participate in bond formation is called the inert pair effect. Since the magnitude of this effect increases down the group, the +1 oxidation state becomes more and more stable down the group as compared to the +3 oxidation state. The inert pair effect is maximum in the case of Tl and therefore, it shows mainly a +1 oxidation state. Due to lesser stability, Tl3+ salts act as strong oxidising agents. This is evident from its electrode potential data:

Tl³⁺(aq) + 2e → Tl+(aq);E° = +1.25V

The inert pair effect may also be explained by the fact that as the size of the atom increases from Al to Tl, the energy required to unpair the ns² -electrons is not compensated by the energy released due to the formation of two additional bonds.

Electropositive and metallic character

The elements of group- 13 are less electropositive or metallic as compared to the elements of group 2. On moving down the group, the electropositive character of the elements first increases from B to Al and then decreases from Al to Tl.

Electropositive and metallic Explanation:

1. Elements of group-13 are smaller in size and the sum of the three ionisation enthalpies Δ_iH₁+ Δ_iH₂ + Δ_iH₃ needed to form M³⁺ ions is much higher than the sum of two ionisation enthalpies, Δ_iH₁+ Δ_iH₂ < Δ_iH₃ for the corresponding bigger-sized elements belonging to alkaline earth metals needed to form M²⁺ ions. For this reason, the elements of group 13 are less electropositive than the elements of group 2.

2. Boron has the highest sum of the first three ionisation enthalpies among the elements of group 13. Because of this, it has very little tendency to lose electrons and hence it is the least electropositive among group-13 elements. It is a non-metal and a poor conductor of electricity.

3. On moving from B to Al, the sum of the first three ionisation enthalpies decreases considerably (6887 to 5137kJ-mol^-1 ) due to an increase in atomic size and hence, Al has a much higher tendency to lose electrons, i.e., Al is sufficiently electropositive. All is a metal and a good conductor of electricity.

4. Because of the increasingly poor shielding effect of 3d -electrons in Ga, 4d -electrons in and 4 f-electrons in Tl, the effective nuclear charge gradually increases and as a consequence, they exhibit lesser electropositive and metallic character.

Density

Because of smaller atomic and ionic radii, the elements of group 13 have a higher density as compared to the elements of group 2. On moving down the group, density increases.

Density Explanation:

On moving down the group, the density of these elements increases because the extent of the increase in atomic mass is greater than the extent of the increase in atomic size.0 On moving from B to Tl, both atomic mass and no. of electrons in the inner d- and f- subshell increases. Due to the lower shielding effect of d-and f— electrons, the effective nuclear charge increases from B to Tl. As a result, from B to Tl, the atomic size does not increase much.

Melting and boiling points

Elements of group 13 do not show a regular trend in their melting points. The melting points decrease from B to Ga and then increase from Ga to Tl.

Melting and boiling points Explanation:

This irregular trend is probably due to unusual crystal structures of B and Ga. The much higher melting point of B is due to its giant covalent polymeric crystal structure consisting of icosahedral units with B-atoms at all 12 corners and each B-atom is bonded to five equidistant neighbours resulting in much stronger attractive forces. In contrast, Ga consists of discrete Ga₂ molecules so its melting point is exceptionally low (303K). However, the boiling points of these elements decrease regularly on moving down the group.

Gallium remains liquid over a vast range of temperatures and no other low-melting metal can compare with it. Molten Ga begins to boil only when heated to a temperature of 2276K. Due to this unusual property, gallium is used in thermometers required for measuring very high temperatures (>1000°C).

Electronegativity

Elements of group-13 are more electronegative than the elements of group-1 (alkali metals) and group-2 (alkaline earth metals). On moving down the group, the electronegativity first decreases from B to Al and then increases marginally.

Electronegativity Explanation:

•  Because of the smaller atomic size and higher nuclear charge, the electronegativities of group-13 elements are higher than the corresponding elements of group-1 and 2.0 On moving down the group from B to Al, the atomic size increases considerably and as a result, the attraction of the nucleus for the electrons decreases and hence the electronegativity decreases.
• On moving from Al to Tl, the atomic size increases but at the same time effective nuclear charge increases due to poor shielding of the inner d and f-electrons. As a result, the attractive force of the nucleus for the electrons increases and hence the electronegativity increases

Chemical Elements Properties Of Group-13

The members of group-13 elements have three electrons in their valence shells. Except for the last member Tl, all other members use these electrons to form three bonds and thus exhibit an oxidation state of +3. In the +3 oxidation state, the members of the boron family are expected to form covalent bonds for the following reasons:

• Small size and high charge (+3) cause high polarisation of the anions leading to the formation of covalent bonds.
• The large value of the sum of the first three ionisation enthalpies, Δ_iH₁ + Δ_iH₂ + Δ_iH₃ of these elements also suggests that the bonds will be largely covalent.
• The difference in electronegativity between the elements
  of group 13 and those ofthe higher groups is not very high.

This fact also agrees with the formation of covalent bonds. Because of its small size and high ionisation enthalpies, it is not possible for boron to form B3+ ions by losing its three valence electrons.

Therefore, 113 Boron does not form ionic compounds. It always forms covalent compounds by showing its valence electrons. The sum of the first three ionisation enthalpies, Δ_iH₁ + Δ_iH₂ + Δ_iH₃ is also higher but less than that of B. So A1 also has a strong tendency to form covalent compounds,

Example:  AlCl₃ , AlBr3 and AlCl₃. Like Al, compounds of the rest of the members such as GaCl₃, In Cl₃ etc. are covalent when anhydrous.

However, all the members except B form metal ions in solution. This change from covalent to ionic nature may be explained by the fact that in aqueous solutions these ions undergo hydration and the amount of hydration enthalpy exceeds the ionisation enthalpy.

Ga, In and TI show two oxidation states of +1 and +3 due to the inert pair effect. The compounds in the +1 oxidation state are more ionic than the compounds in the +3 oxidation state. In a trivalent state, the number of electrons in the valence shell of the central atom in a molecule of these elements is only six (two electrons less than the octet) and therefore, such electron-deficient molecules behave as Lewis acids. For example, BCl3 (Lewis acid) readily accepts an unshared pair of electrons from ammonia (Lewis base) to form the adduct, BCl₃-NH₃.

Reaction with dioxygen or air

1. All the members of group 13 react with dioxygen at higher temperatures to form trioxides of the general formula M₂O₃. T1 forms both T1203 and some amount of Tl₂O

The reaction of Al with 02 is known as a thermite reaction which is highly exothermic (AH0 =-1670kJ mol-1 ). A very strong affinity of Al for oxygen is used in the extraction of other metals from their oxides (thermite process). For example, Mn and Cr can be extracted from Mn₃O₄ and Cr₂O₃ respectively by this process.

2. The reactivity of group-13 elements towards dioxygen increases on moving down the group. Pure crystalline boron is almost unreactive towards air at ordinary temperature. Al does not react with dry air.

However, it gets tarnished readily in moist air even at ordinary temperatures due to the formation of a thin oxide (Al₂O₃) layer on the surface which prevents the metal
from further reaction. When amorphous boron and aluminium metal are heated in air, they form boron trioxide and aluminium trioxide (Al₂O₃) respectively.

Ga and In are not affected by air but T1 forms an oxide on
its surface in the presence of air.

3. B and Al react with dinitrogen at high temperatures to form
the corresponding nitrides.

Ga, In and T₁ do not react with N₂ to form the corresponding nitrides.

Boron nitride is a white slippery solid which melts under pressure at 3246K. It is chemically inert towards the air, oxygen, hydrogen, chlorine, etc. even on heating. The total number of valence electrons of one B and one N- -atom is equal to the number of valence electrons of two C-atoms.

Therefore, the structure of boron nitride is almost the same as that of graphite having a layer lattice. In each layer, alternate B and N-atoms (both sp2 -hybridised) form
a planar hexagon. The layers are stacked over one another in such a way that the N-atom of one layer is directly over the B-atom of another layer. Because of its structural similarity with graphite, boron nitride is also called inorganic graphite.

When boron nitride is heated at 1800°C under very high pressure, it gets converted to a cubic form comparable to diamond. This extremely hard variety known as borazon is used for cutting diamonds.

The acid-base character of oxides and hydroxides:

1. Trioxides of the elements of the boron family react with water to form their corresponding hydroxides.

M₂O₃ + 3H₂O→  2M(OH)₃

2. The nature of these oxides and hydroxides changes on moving down the group. Both B203 and B(OH)3 are weakly acidic. They dissolve in alkali to form metal borates.

B₂O₃ + 2NaOH → 2NaBO₂(sodium metaborate) + H₂O

B(OH)₃ + 3NaOH → Na₃BO₃(sodium borate) + 3H₂O

Aluminium oxide and hydroxide are amphoteric. Both of them dissolve in alkalies as well as acids.

1. Al₂O₃(s) + 3H₂SO₄(s)→ Al₂(SO₄)₃(aq) + 3H₂O(l)

2.

Similarly, Al(OH)₃(s) + NaOH(s)→Na[Al(OH)₄](aq)

Al(OH)₃(aq) + 3HCl(aq)→ AlCl₃(aq) + 3H₂O(l)

The oxide and hydroxide of Ga are also amphoteric while those of In and TI are basic.

Therefore, the basic character of oxides and hydroxides increases down the group

3. Thallium forms two types of hydroxides:

Thallic hydroxide [Tl(OH)3] and thallous hydroxide (TlOH). Al(OH)3 is insoluble in H20 but TlOH is soluble and is a
strong base like alkali metal hydroxides.

Thallium Explanation:

On moving down the group, the magnitude of ionisation enthalpy decreases. As a result, the strength of the M— O bond also decreases and therefore, its cleavage becomes progressively easier resulting in the increased basic strength down the group

An extremely hard crystalline form of aluminium oxide called corrundum is used as an abrasive. It can be made by heating amorphous aluminium oxide at about 2000K. Aluminium forms a series of mixed oxides with other j metals, some of them occurring naturally as semi-precious stones. These include ruby (Cr³⁺) and blue sapphire
(CO²⁺, Fe²⁺, Tl⁴⁺).

Reaction with hydrogen

1.  Group-13 elements form hydrides ofthe type MH₃. The members of the boron family do not combine directly with hydrogen. However, several hydrides are known
which can be prepared indirectly. Boron forms several stable covalent hydrides which are collectively called boranes.

The two most important types of boranes are as follows:

1. Boraness with general formula B_nH_n+4 are called nido-boranes 
   • Examples: Diborane (B₂H₆) pentaborane-9 (B₅H₉).
2. Boranes with the general formula B_nH_n+6 are called arachnoid-boranes
   • Examples:  Tetraborane (B₄H₁₀), and pentaborane-11 (B₅H₁₁).

2. The simplest hydride is diborane (B₂H₆) which is prepared by the reaction of BF₃ with lithium hydride industrially.

3. The other members of group 13 also form several hydrides which are polymeric,

Example:

(AlH₁)_n, (GaH₃)_n, (InH₃)_n–

The stability of these hydrides decreases down the group and thallium hydride is quite unstable. Boron, aluminium and gallium also form complex anionic hydrides such as NaBH₄ (sodium borohydride), LiAlH₄ (lithium aluminium hydride) and LiGaH₄ (lithium gallium hydride). These complex hydrides act as powerful reducing agents

4. The hydrides are weak Lewis acids and readily form adducts with strong Lewis bases to form compounds of the type MH₃:B (B = base).

For example, AlH₃:NMe₃, GaH₃:NMe₃ etc.

NMe₃ + AlH₃→[Me₃N:→AlH₃]

Reaction With Acids And Alkalies

The action of acids:

1. Boron remains inert in the presence of non-oxidising acids such as HCl. However, it undergoes oxidation by strong oxidising acids such as a mixture of hot concentrated H₂SO₄ and HNO₃ (2: 1) to form boric acid (H₃BO₃) at very high temperatures.

2.  The remaining elements of this group react with both oxidising and non-oxidising acids. For example, Al dissolves in dilute HCl and liberates dihydrogen.

2Al(s) + 6HCl(aq)→ 2Al3+(a<7) + 6Cr(a<jr) + 3H2(g)

3. Concentrated nitric acid renders aluminium passive by forming a protective layer of its oxide (AlO₃) on the surface ofthe metal. Thus aluminium vessels can be used
to store concentrated HNO₃.

2Al + 6HNO₃→Al₂O₃+ 6NO₂ + 2H₂O

Ga, In and TI react with dilute acids to liberate H₂

Action of alkalies:

1. When boron is fused with alkalies (NaOH or KOH) at a temperature greater than 775K, it forms borates and liberates dihydrogen

2. Boron dissolves in a fused mixture of Na₂CO₃ and NaNO₃ at 1123K to produce borate and nitrite salt and liberate carbon dioxide.

Al and Ga also react with aqueous alkalies with the evolution of dihydrogen

In and TI does not react with alkalis.

Reaction with halogens

Elements of group-13 react with halogens at high temperatures to produce trihalides ofthe general formula, MX₃. However, thallium (III) iodide does not exist.

Trihalides of boron:

Due to its small atomic size and high ionisation enthalpy, boron forms covalent trihalides

BX₃. BF₃ is a gas, BCl₃ and BBr₃ are liquids and BI3 is a solid. All these are trigonal planar molecules in which the central B -atom is sp² -hybridised. The three unpaired electrons of p -orbitals of three halogen atoms overlap with the three sp² -orbitals of boron to form three sp²-p, B—X cr -bonds. The unhybridised empty p-orbital remains perpendicular to the plane of the molecule.

Since there are only six electrons in the valence shell of the central boron atom in boron trihalides, they can accept two more electrons to acquire a stable octet, i.e., boron trihalides can behave as Lewis acids.

The Lewis acid character, however, decreases in the order:

BI₃ > BBr₃ > BCl₃ > BF₃

Explanation: This order of relative Lewis acid strength of boron trihalides, is just the reverse of what may be expected based on the electronegativities of the
halogen atoms can well be explained based on the tendency of the halogen atom to donate its lone pair of electrons to the boron atom through pn-pn back bonding.

Since the vacant 2p -orbital of B and the 2p-orbital of Fatom containing a lone pair of electrons are equal in size, therefore, the tendency of the F -atom to donate the unshared
pair by pπ-pπ back bonding is maximum. BF₃ can well be represented as a resonance hybrid of four resonating structures. As a result of resonance involving pn-pn back bonding, the electron density on the boron atom increases effectively and so its strength as a Lewis acid decreases considerably.

As the size ofthe halogen atom increases on going from Cl to I, the extent of overlap between the 2p -orbital of boron and a large p -orbital of halogen (3p of Cl, 4p of Br and 5p of I] decreases. As a consequence, the electron deficiency of boron increases and thus, the Lewis acid strength decreases on going from BF₃ to BI₃.

Halides of aluminium:

The halides of aluminium in the vapour state as well as in an inert solvent such as benzene exist ns dimers.

For example, Al₂Cl₃ exists as Al₂Cl₃.

Halides of aluminium Explanation:

In AlCl₃, there are six electrons (two electrons less than tyre octet) around the central Al-atom. In the dimeric structure, each Al completes its octet by accepting a lone
pair of electrons from the Cl-atom of another AlCl₃ molecule. The dimeric form exists lit vapour state at < 473K.

However, at higher temperatures, it dissociates to trigonal planar AlCl₃ molecule. In polar solvents such as water, the dimer dissociates and it is the high hydration enthalpy which helps this dissociation leading to the formation of Al³⁺ ion

Al₂Clg + 6H₂O ⇌ 2[Al(H₂O)6]³⁺(aq) + 6Cl^–(aq)

Therefore, anhydrous AlCl₃ is covalent but, hydrated aluminium chloride is ionic. Some important points of distinction between boron and
the other members of its family (especially the next member Al) are discussed in the given table. Anomalous properties of boron

Unlike aluminium halides, boron halides exist as monomers and this is because the boron atom is so small that it cannot accommodate four large-sized halogen atoms around it.

Boron

• Boron is the first member of group 13 of the periodic table. There are three electrons in its valence shell (ls²2s²2p¹).
• It exhibits anomalous behaviour and differs from the other members of its family. The reasons behind its exceptional behaviour can be attributed to the
• Exceptionally small atomic size as compared to other elements of its group,
• Much higher ionisation enthalpy and absence of d orbitals in its outermost or valence shell. Boron forms electron-deficient compounds which act as Lewis acids

Anomalous behaviour of boron:

Some important points of distinction between boron and the other members of its family (especially the next member Al) are discussed in the given table

Anomalous properties of boron:

1. Occurrence

Boron does not exist in a free state in nature. It is always found in the combined state as boric acid and borates. Boron occurs in two isotopic forms, ¹⁰B (19%) and ¹¹B (81%). Its abundance in the earth’s crust is very low (about 0.001%)

Important minerals of boron:

Boron may be obtained from the jets of steam which erupt from the volcano as boric acid and also from the water of the hot spring of Tuscany in small amounts as boric acid.

2. Properties of boron

Boron Physical properties:

1. Boron is an extremely hard solid (next to diamond) having a much higher melting point (2450K) and this is because of its three-dimensional network structure. Its boiling point is 3923 K.

2. Boron exists in two allotropic forms namely:

1. Amorphous,
2. Crystalline.

Crystalline boron is of three types:

1. α – Rhombohedral
2. β – Rhombohedral and
3. ϒ – Tetragonal.

The building units of all these forms are B₁₂ icosahedral units with 20 faces and boron atoms at all the 12 comers or vertices.

The melting and boiling point of boron is 2450K and 3923K respectively. The reason behind such high melting and boiling points is attributed to very strong attractive forces among the B₁₂ units as well as its closely packed stable crystal structure.

Chemical properties of crystalline boron

• Crystalline boron is chemically very inert. It is not oxidised even when heated with oxygen Crystalline boron is not attacked by HCl or HF. It is not affected by various oxidising acids such as shot and cone. HNO₃, H₂SO₄ etc.
• When it is fused with Na₂O₂ or Na₂CO₃ and KNO₃ at high temperatures, sodium borate (Na₃BO₃) is obtained.

Chemical properties of amorphous boron

1. Reaction with air:

When amorphous boron is heated in air at 700°C, it bums with a red flame and undergoes oxidation to form boron trioxide. Boron nitride is also formed by its reaction with N₂ gas of air

4B + 3O₂ → 2B₂O₃; 2B + N₂ → 2BN

2. Reaction with strong alkali:

When amorphous boron is fused with NaOH or KOH at a temperature greater than 773K, it forms borate salts and liberates H₂ gas

3. Reaction with oxidising adds:

Boron is not affected by non-oxidising acids such as hydrochloric add. However, it reacts with oxidising adds like cone. H₂SO₄ and HNO₃ form boric add.

3H₂SO₄ + 2B → 2H₂BO₂ + 3SO₂

6HNO₃ + 2B→2H₃BO₃ + 6NO₂

4. Reaction with halogens:

Boron bums in fluorine gas to form boron trifluoride. Boron reacts with chlorine at high temperatures to form boron trichloride

2B + 3F₂→ 2BF₃ ;2B + 3Cl₂→2BCl₃

An aqueous solution of BC13 is addicting because it undergoes hydrolysis to form a mixture of HCl and boric acid.

BCl₃+ 3H₂O →  3HCl + H₃BO₃

5. Reaction with metals (oxidising property):

The binary compounds of boron with elements having electronegativity lower than boron itself (For example: Metal) are called borides. When boron is heated with a metal at high temperature in an electric arc furnace, borides are obtained (B acts as an oxidant). Borides are hard, inert and have special properties

3Mg + 2B→ Mg₃B₂; 3Ca + 2B→ Ca₃ B₂

6. Reaction with carbon:

When boron is heated with carbon having comparable size and electronegativity at high temperatures in an electric arc furnace, very hard covalent boron carbide (B₄C) is obtained. It is even harder than diamonds. It is used as an abrasive. 4B + C→ B₄C It can also be prepared by reducing B₂O₃with coke at high temperature (2500°C) in an electric furnace

2B₃O₃ → B4C + 6CO↑

Reducing property: When boron is heated strongly with SiO₂ and CO₂, it reduces these oxides to give Si and C respectively.

2B₂O₃ + 7C →  B₄C + 6CO↑

7. Reducing property:

When boron is heated strongly with SiO₂ and CO₂, it reduces these oxides to give Si and C respectively.

8. Reaction with water: Red hot boron reduces steam to yield B₂O₃ and dihydrogen

Uses of boron

• Boron, an extremely hard refractory solid with a high melting point, low density and very low electrical conductivity, finds many applications which are as follows:
• Boron fibres having enormous tensile strength are used in making bullet-proof vests and as reinforcement materials in space shuttles and aircraft.
• Because of the high tendency of isotopes to absorb neutrons, metal borides are used in nuclear reactors as protective shields and control rods.
• It is used in the steel industry (instead of using expensive metals like Mo, Cr and W) for manufacturing special types of hard steel
• Its compounds such as borax and boric add are used for making heat-resistant glass (i.e., p a mild antiseptic.
• Boron compounds are used as rocket fuels because of their high energy/mass ratio.
• Boron carbide (B₄C) is used as an abrasive for polishing or grinding.
• Boron is used as a semiconductor for making electronic devices glass), glass-wool and gÿre glass
• An aqueous solution of orthoboric acid is used as

Some Important Compounds Of Boron

1. Borax, Na₂B₄O₇.10H₂O or Na₂[B₄O₅(OH)₄] 8H₂O

Borax or sodium tetraborate decahydrate which occurs naturally as tincal (sugar) in certain dried-up lakes is the most important compound of boron. Borax contains the tetranuclear units

Therefore, its correct formula is Na₂[B₄O5(OH)₄]-8H₂O.

Borax Preparation

1. From tincal:

Naturally occurring borax or tincal, which contains about 50% borax is boiled with water, concentrated and then filtered to remove insoluble impurities. The filtrate is then concentrated and cooled when crystals of borax separate.

2. From colemanite:

Finely powdered mineral, colemanite (Ca₂ B₆O₁₁) is boiled with sodium carbonate solution and CaCO₃, Na₂B₄O₇ and NaBO₂ are obtained

Precipitate of CaCO₃ is filtered off and the filtrate is then concentrated and cooled to get the crystals of borax. A current of CO₂ is passed through the mother liquor when sodium metaborate presentient gets converted into borax

4NaBO₂ + CO₂→ Na₂B₄O₇(Borax) + Na₂CO₃

3. From boric acid:

Borax may also be obtained by neutralising boric acid with sodium carbonate. Crystals of Na₂B₄O₇.10H₂O separate on cooling.

4H₃BO₃ + Na₂CO₃ →  Na₂B₄O₇ + 6H₂O + CO₂↑

Borax Physical properties:

1. It is a white crystalline solid.
2. It is less soluble in cold water but more soluble in hot water.
3. When ordinary borax is recrystallised from water at a higher temperature (≈60°C), crystals of sodium tetraborate pentahydrate (Na₂B₄O₇.5H₂O) separate. This is called ‘goldsmith’s sugar.

Borax Chemical properties

1. Nature of aqueous solution:

The aqueous solution of borax is alkaline in nature and this is because borax undergoes hydrolysis to form the strong alkali, NaOH and the weak acid, boric acid. It acts as a buffer.

As the aqueous solution of borax is alkaline, it can be titrated against an acid using an orange indicator

Na₂B₄O₇ + 2HCl + 5H₂O→ 4H₃BO₃ + 2NaCl

When phenolphthalein is added to an aqueous solution of borax, the solution becomes pink in colour. However, when glycerol (a polyhydroxy compound) is added to the solution, it becomes colourless again.

Aqueous solution Explanation:

Since the aqueous solution of borax is alkaline in nature, it becomes pink when phenolphthalein is added to it. When glycerol is added to that solution, it combines with B(OH)₄ and removes it from the equilibrium by forming a stable chelate complex. As a consequence, the equilibrium shifts to the right making boric acid a strong acid. Because of the increased concentration of H⁺ ions, complete neutralisation of OH^– ions occurs and the solution becomes colourless again.

H₃BO₃ + H₂O  ⇌  H⁺ B(OH)₄^–; K_a = 6 × 10^-10

2. Reaction with caustic soda:

When a calculated amount of NaOH is added to borax, sodium metaborate is obtained.

Na₂B₄O₇ + 2NaOH→4NaBO₂ + H₂O

Reaction with sulphuric acid: When a calculated quantity of concentrated sulphuric acid is added to a hot concentrated solution of borax, boric acid is produced

3. Reaction with ethanol and sulphuric acid:

When borax is heated with ethanol and concentrated H₂SO₄, vapours of triethyl borate are formed which on ignition bum with a green-edged flame.

Na₂B₄O₇+ H₂SO₄ + 5H₂O →  Na₂SO₄ + 4H₃BO₃

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

This reaction is used as a test for the detection of borate ion (BO^3-₃) in qualitative analysis.

4. Action of heat:

When borax is heated strongly in the flame of a Bunsen burner, it loses its water of crystallisation and swells up to form a puffy mass. On further heating, the

mass turns into a transparent liquid which Nolldllics to form a bead that consists of sodium metaborate (NaBO₂) and boric anhydride (B₂O₃)

Preparation til boron from borax

A hot and concentrated solution of borax reacts with concentrated H₂SO₄ to form I boric acid (H₃BO₃). Boric added when heated tit high temperature, successively dissociates to form boron d ioxide, (B₂O₃). Boron trioxide when heated with Mg-powdor produces boron.

Na₂B₄O₇ + H₂SO₄+ 5H₂O→Na₂SO₄ + 4HaBO₃

Borax bead test

The borax bead test is very Important In a qualitative analysis for the detection of coloured metal ions like Cu²⁺, Ni²⁺, Co²⁺, Cr²⁺ etc.

• At first, a hot platinum loop is touched with borax and then heated in a Bunsen burner’s flame.
• Borax at first swells up and finally melts to form a colourless bead in the loop.
• The hot loop is touched with the salt under investigation and heated at first in the oxidising flame and then in the reducing flame.
• The metal ion is identified from the colour of the bead. This test is called the borax bead test.

Borax Bead Test:

Reactions:

1. Metallic compounds undergo decomposition on heating to form metallic oxides.

2M(NO₃)₂ → 2MO + 4NO₂ + O₂

2MSO₄ → 2MO + 2SO₂ + O₂

(M = Cu, Fe, Co, Ni, Mn, Cr)

2. The basic metallic oxides dissolve in the acidic diborane trioxide(BO) of the borax bead and form coloured metal metaborate salts

MO + BoO₃ → M(BO₂)₂

3. Copper Iron and other metallic salts form -ic metaborate in oxidising flame and -metaborates in reducing flame

Examples:

1. Reactions with copper salt:

In oxidising flame, cupric metaborate (blue) is formed

In reducing flame cuprous metaborate is formed

1.

2.

2. Reactions with iron salt:

In oxidising flame, ferric metaborate (yellow) is formed

In reducing flame ferrous metaborate (green) is formed

3. Reactions with cobalt salt: Both in oxidising and reducing flame, cobalt metaborate (blue) is formed

As the oxidation states of Co (+2) and Cr (+3) remain unchanged, the colour of the bead obtained from them is the same for both reducing and oxidising flame.

Uses of borax

Borax is used:

• In the manufacture of heat-resistant borosilicate (pyrex) glass,
• For preparing medicinal soaps,
• As a flux in soldering metals.
• In the candle industry as a stiffening agent,
• In softening water
• For the borax bead test,
• In the manufacture of perborate
• Na₂(OH)₂B(O — O)2B(OH)₂ 6H₂O, is an important cleansing and bleaching agent used in washing powders.

Orthoboric acid or boric acid, H₂BO₃o r B(OH)₃

The trivial name of orthoboric acid is boric acid

Orthoboric acid Preparation:

1. From colemanite:

Sulphur dioxide is passed through a hot concentrated solution of the mineral cole, Win’ll the resulting solution is concentrated and cooled, and crystals of boric add separate out. Calcium bisulphite being highly soluble In water remains dissolved In the mother liquor.

Ca₂ B₆O₁₁  (Colemanite)+ 4SO₂ + 11H₂ O →  2Ca(HSO₃)₂ (Calcium bisulphite)+ 6H₃BO₃ ( Boric acid)

2. From borax:

When a hot concentrated solution of borax is treated with hydrochloric acid or sulphuric acid, boric acid Is obtained. The resulting solution is concentrated and then cooled when crystals of boric acid separate out

Na₂B₄O₇ + 2HCl + 5H₂O →4H₃BO₃ + 2NaCI

Na₂B₄O₇ + H₂SO₄ + 5H₂O → 4H₃BO₃ + Na₂SO₄

3. From boron compounds by hydrolysis:

Certain boron compounds such as halides, hydrides and nitrides on boiling with water (hydrolysis) produce boric add.

BCI₂ + 3H₂O → H₃BO₃+ 3HCl

B₂H₆(dlborane) + 6H₂O→2H₃BO₃ + 6H₂↑

BN(boron nitride) + 3H₂O→ H₃BO₃ + NH₃

Orthoboric acid’s Physical properties

• It Is a white needle-like crystalline solid with a soft soapy touch.
• It Is sparingly soluble in cold water but highly soluble in hot water.
• It Is steam volatile.

Orthoboric acid Chemical properties

1. Addle nature:

It Is a very weak monobasic acid (K_a = 6 × 10^-10). It does not donate protons like most protonic acids. In fact, due to the small size of B and the presence of only six electrons Its valence shell B(OH)₃ behaves as a Lewis acid and accepts a pair of electrons from OH” Ion water thereby releasing a proton.

H(OH)₃ + 2H₂O ⇌  [B(OH)₄]^–+ H₃O⁺

B(OH)₃ behaves as a very weak acid (pK_a  = 9.2) because it only partially reacts with water to form [B(OH)₄]^–and H₃O⁺ Ions. So, B(OH)₃ or H₃BO₃ cannot be titrated satisfactorily with NaOH solution because no sharp end point Is obtained.

If some polyhydroxy compound such as glycerol, mannitol or catechol is added to the titration mixture, then boric acid behaves as a strong monobasic add.  This occurs due to the Removal of [B(OH)₄]^– Ion from the equilibrium mixture by the formation of a stable complex with the polyhydroxy compound.

It can then be titrated with NaOH solution and the endpoint can be detected using phenolphthalein as an indicator. ‘

2. Action of heats:

When orthoboric acid is heated, it loses molecules of water in three stages at different temperatures thus forming different products

3. Reaction with ethyl alcohol:

Boric acid reacts with ethyl alcohol In the presence of concentrated sulphuric acid to form triethyl borate

Vapours of triethyl borate burn with a green-edged flame. This test is used for detecting boric acid in qualitative analysis.

It is to be noted that this test can also be performed without using H₂SO₄ However, for detecting borate ions, the presence of H₂SO₄ is required. Therefore, boric acid and borate ions can be distinguished by this test.

4. Reaction with fluoride salt:

Boric acid reacts with fluoride salt in the presence of concentrated H₂SO₄ to form volatile boron trifluoride (BF₃). This compound burns with a green-edged flame.

2H₃BO₃ + 3CaF₂ + 3 H₂SO₄ → 3CaSO₄ + 2BF₃ + 6H₂O

5. Reaction with ammonium bi fluoride:

When boric acid is heated with ammonium bi fluoride, no residue is obtained because all the resulting compounds are gaseous

B₂O₃3(s) + 6NH₄BF₄(S)→  8BF₃(G) + 6NH₃(g) + 3H₂O(g)

6. Reaction with potassium fluoride:

When the aqueous solutions of two acidic compounds, boric acid and potassium fluoride (KHF2) are mixed, an alkaline solution is obtained due to the formation of potassium tetrafluoroborate (KBF₄) and potassium hydroxide (KOH).

B(OH)₃+ 2KHF₂→  KBF₄ + KOH + 2H₂O

Being a Lewis acid, B(OH)₃ has a strong tendency to combine with relatively smaller (F^–) ions to form fluoroborate ion (BF^–) and for this reason, this unbelievable reaction takes place.

Uses of boric acid

Boric acid is used

• As a mild antiseptic for washing eyes under the name Boric lotion
• In the manufacture of heat-resistant borosilicate glass,
• As a preservative for milk and foodstuffs,
• In the manufacture of enamels and glazes for pottery.

Structure Of boric acid:

The shining white crystals of boric acid contain B(OH)₃ units linked by H -bonds in infinite layers of nearly hexagonal symmetry. Since the adjacent layers in the boric acid crystal are held together with weak attractive forces, one layer can easily slide over the other and hence, boric acid is soft and slippery touch.

Diborane, B₂H₆

Boron hydrides are binary compounds of B and H. Although boron does not combine directly with hydrogen, several boron hydrides collectively called boranes, (in analogy with alkanes) are known. Depending upon their general formulae, these hydrides

Can be divided into several categories of which the following two are the most important:

1. Nido-boranes: General formula: \(\mathrm{B}_n \mathrm{H}_{n+4} \text {, example} \mathrm{B}_2 \mathrm{H}_6\) (diborane), B₅H₉ (pentaborane-9), B₆H₁₀ (hexaborane- 10), B₈H₁₂(octaborane-12),B₁₀H₁₄ (decaborane-14) etc. 
2. Arachno-boranes: General formula: \(\mathrm{B}_n \mathrm{H}_{n+6} \text {, example }\) B₅H₁₄ (pentaborane-11), BgH₁₂ (hexa-borane-12), BgH₁₄(octaborane-14) etc. The mostimportant hydride ofboron is diborane (B₂Hg).

Preparation of diborane

1. Laboratory preparation:

Diborane is prepared by the oxidation of sodium borohydride (NaBH₄) with I₂ in a diglyme solution.

Diglyme is a polyether whose formula is CH₃OCH₂CH₂OCH₂CH₂OCH₃

2. From boron trifluoride etherate:

It may be prepared by the reduction of boron trifluoride etherate with lithium aluminium hydride (LiAlH₄) in diethyl ether or sodium borohydride (NaBH4) in diglyme

3. Industrial preparation:

On an industrial scale, diborane is prepared by reducing BF₃ with LiH or NaH. 450K

Diborane Physical properties

Diborane is a colourless, foul-smelling, highly toxic gas having a boiling point of 180K.

Diborane Chemical properties

1.  Thermal stability:

It is stable only at low temperatures. When it is heated at 373-523K in a sealed tube, several higher boranes are obtained

However, by controlling the temperature, pressure and reaction time, various individual boranes can be prepared.

2. Combustibility:

When it is exposed to air, it spontaneously catches fire because of the strong affinity of boron towards oxygen. This reaction forming boric anhydride and water is highly exothermic.

B₂H₆ + 3O₂ →B₂O₃ + 3H₂O; ΔH = – 1976 kJ -mol^-1

The higher boranes also spontaneously in the air.

3. Hydrolysis:

It undergoes ready hydrolysis to produce boric acid.

B₂H₆(g) + 6H₂O(aq)→2H₃BO₃(aq) + 6H₂(g)↑

It reacts with methanol to form trimethyl borate.

B₂H₆ + 6CH₃OH → 2B(OCH₃)₃ + 6H₂↑

4. Reaction with Lewis bases:

When diborane is treated with Lewis base, it undergoes cleavage to form monoborane which then reacts with Lewis base to form an adduct.

B₂H₆ + 2NMe₃→ 2BH₃-NMe₃; B₂H₆ + 2CO → 2BH₃-CO

5. Reaction with ammonia:

When diborane is treated with ammonia, an additional compound is formed. The compound on further heating at about 473 K decomposes to give a volatile compound called borazine (or borazole).

Borazine is isosteric (i.e., the same number of atoms) and isoelectronic (i.e., the same number of electrons) with benzene and its structure is similar to that of benzene. Like benzene, all the atoms in borazine are sp² -hybridised. The n -n-bonding of borazine is dative and it arises due to sideways overlapping of filled p-orbitals of N and empty p-orbitals of B. Because of its similarity with benzene, borazine is also called inorganic benzene.

6. Formation of complex borohydrides:

Diborane reacts with several metal hydrides to form borohydrides containing tetrahedral [BH₄]_ ion.

Both sodium and lithium borohydrides are used as very good reducing agents in the synthesis of organic compounds in the laboratory. These two compounds may also be used as starting material for the preparation of other borohydride compounds.

7. Reaction with alkalis:

Diborane dissolves in strong alkalies such as NaOH or KOH solution to form metaborates and H₂ gas.

B₂H₆ + 2KOH + 2H₂O→ 2KBO₂ + 6H₂(Potassium Metaborate)↑

8. Reaction with halogen acids:

Diborane reacts with halogen acids to form halodiborancs and hydrogen gas. The order of reactivity of halogen acids is: HI > HBr > HCl

9. Reaction with halogens:

Diborane reacts with halogens to form corresponding halodiboranes. The order of reactivity of halogens is Cl₂ > Br₂ > I₂. Thus, chlorine reacts with diborane explosively at room temperature, bromine reacts rapidly at 373 K but iodine reacts slowly at higher temperatures

Uses of diborane:

1. Diborane is used in the preparation of several borohydrides such as LiBH₄ NaBH₄, etc.
2. It is used as a reducing agent in organic synthesis.
3. It is also used as a fuel for supersonic rockets.

Structure Of diborane, B₂H₆:

The structure and bonding of diborane seem to be very interesting. In the excited state, the B atom has the electronic configuration 2s¹2p_x¹ 2p_y¹ and therefore, it has only three electrons available for sharing. Now, 14 electrons (for six B—H and one B—B bond) are required if boron forms all conventional covalent bonds in ethane (C₂H₆)

But there are only 12 electrons (six from two B atoms and six from six H-atoms). Thus, the molecule is short of two electrons and its structure can not be similar to that of ethane (C₂H₆)

Based on electron diffraction study:

• Diborane has a bridged structure as given in There are two types of hydrogen atoms in this bridged structure. The two boron atoms and four terminal hydrogen atoms (shown by thick lines) lie in the same plane, while the remaining two hydrogen atoms (shown by dotted lines) lying above and below the plane form bridges and these are called bridge hydrogen atoms.
• The two B-H-B bridges lie in a plane which is nearly perpendicular to the plane containing the terminal B—H bonds.
• There are two bonds in the diborane molecule:
  1. The four terminal B—H bonds are normal covalent bonds, each Being formed by sharing a pair of electrons between boron and hydrogen atoms. These are quite strong bonds and called two-centre electron pair bonds or two-centre two-electron bonds (2c-2e bonds),
  2. The two bridge bonds B …. H……B are quite different from the normal electron pair bonds. Each bridge H-atom is bonded to two boron atoms by sharing only one pair of electrons.
• Such bridge bonds are called three centre electron pair bonds or three centres two-electron bonds (3c-2e bonds). Three-centre electron pair bonds or three-centre two-electron bonds are very weak bonds and are often called banana bonds as they resemble bananas in shape.
• Molecules like diborane (B₂H₆) which do not have a sufficient number of electrons to form normal covalent bonds are called electron-deficient molecules.

Based on hybridisation:

Boron atoms (excited state electronic configuration: 2s¹2p_x¹ 2p_y¹ in diborane undergo sp³ -hybridisation.

1. The two half-filled sp³ -hybrid orbitals of each boron atom overlap with the half-filled orbitals of hydrogen atoms to form normal covalent bonds.

2. The third half-filled hybrid orbital of one of the two boron atoms and the vacant orbital of the remaining boron atom overlap simultaneously with the half-filled Is -orbital of a hydrogen atom to form a B……H….B bridge bond

3. This bond involves three atoms (two boron atoms and one hydrogen atom) and contains only two electrons because one overlapping orbital of boron is empty. Hence, this B–‘H-‘-B bond is called three centre electron pair (3c-2e) bonds. Because of its typical shape resembling a banana, it is also called a banana bond

Aluminium

Aluminium, the second member of the boron family (group-13), is the most abundant metallic element in the earth’s crust. It is found in a variety of aluminosilicate compounds such as clay, mica and feldspar. The only ore of aluminium from which the metallic aluminium can be extracted profitably (in industry) is bauxite. Bauxite is hydrated aluminium oxide whose molecular formula is Al₂O₃-2H₂O

Aluminium Physical properties:

• Aluminium is a bluish-white metal with a brilliant lustre. But aluminium easily gets tarnished by the formation of a thin layer of oxide on the surface.
• It is a light metal whose density is 2.73g-cm^-1. Aluminium possesses high tensile strength, yet it is malleable and ductile.

Aluminium is a very good conductor of heat and electricity.

Aluminium Chemical properties:

It is not as reactive as its high negative electrode potential (E° = -1.66V) would imply and this is because there is a very thin layer of oxide on its surface.

1. Action of air:

Al remains unaffected in dry air but in the presence of moist air, a thin film of oxide is formed over its surface. Hence, the metallic of disappears. When burnt in oxygen it produces brilliant light.

4Al + 3O₂ → 2Al₂ O₃+ 772 kcal

The reaction is highly exothermic and the heat evolved is used for the reduction of oxides of Cr, Fe, Mn etc. (known as the thermite process).

2. Action of water: Aluminium decomposes boiling water thereby evolving hydrogen gas.

2Al + 6H₂O→2Al(OH)₃ + 3H₂↑

3. The action of non-metals:

Besides oxygen, aluminium reacts with other non-metals such as nitrogen, sulphur and halogens to form nitride, sulphide and halides respectively

4. Action of acids:

It dissolves both in dilute and concentrated hydrochloric acid and dilute sulphuric acid along with the evolution of hydrogen gas.

2Al + 6HCl→ 2AlCl₃ + 3H₂↑

2Al + 3H₂SO₄→Al₂(SO₄)₃ + 3H₃↑

The reaction with dilute sulphuric acid is very slow probably due to the insolubility of tyre oxide film over the metal in the acid. Hot and concentrated sulphuric acid dissolves aluminium with the evolution of sulphur dioxide (SO₂) gas.

2Al + 6H₂SO₄→Al₂(SO₄)₃ + 3SO₂ + 6H₂O

Dilute and concentrated nitric acid have no action on aluminium and this is due to the formation of an impenetrable oxide layer on its surface. Nitric acid may, therefore, be kept in the aluminium vessel.

5. Action of alkalis:

Aluminium dissolves in hot and cone. NaOH or KOH solutions form sodium aluminate with the evolution of hydrogen gas.

2Al + 2NaOH +2H₂O → 2NaAlO₂  (Sodium aluminate) (Soluble) +  3H₂ ↑

Aluminium reacts with hot and concentrated sodium carbonate (Na2C03) solution to form sodium aluminate, carbon dioxide and hydrogen.

2Al + Na₂CO₂ + 3H₂O  → 2NaAlO₂ (Sodium aluminate) (Soluble) + CO₂ + 3H₂↑

Uses of aluminium:

• Aluminium alloys (duralumin: Al, Mg, Cu and magnalium: Al, Mg) are light and strong and thus, are used in the construction of aircraft, ships and cars.
• It is a better conductor than copper and is used for making electric power cables.
• It is used for making doors, windows, building panels, mobile homes and household utensils.
• Finely divided Al powder is used in preparing aluminium paint and as an ingredient in solid fuels in rockets.
• Aluminium foils are used In wrapping soaps, cigarettes and confectioneries.
• Al is used to extract metals such as Cr, Mn etc., from their ores (thermite process).
• A mixture of ammonium nitrate and Al dust (ammonal) is used to make bombs and crackers.

CBSE Class 11 Chemistry Notes For Chapter 11 Some P Block Elements Group-14 Elements (Carbon Family)

1. The valence shell electronic configuration of the elements of group 14 is ns²np², where n = 2-6.
2. It becomes clear from these electronic configurations (given in the table below) that carbon and silicon have noble gas cores, germanium and tin have noble gas plus 10 d-electron cores and lead has noble gas cores in addition to 14/ and lOd -electron cores.
3. Thus, the electronic configurations of group-14 elements are similar to that of group-13 elements. However, they contain one more p-electron as compared to group-13 elements.

Electronic configurations of a group of  elements

Occurrence Of Group-14 Elements

1. The members of group 14 are carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb).

2. Carbon is the seventeenth most abundant element by mass in the earth’s crust It is widely distributed in nature in free as well as in combined states. In a free state, it occurs in coal, graphite and diamond. These are the main allotropes of carbon.

Carbon in the form of coal and coke is used mainly as fuel. In a combined state, it is present widely as metal carbonates, hydrocarbons (petroleum), carbohydrates and carbon dioxide (0.03%) in the air. Gases like propane and butane are the major constituents of LPG.

Moreover, the main constituent of all organic compounds is carbon. Two stable isotopes of carbon are present in nature namely ₆C¹² and ₆C¹³ Another isotope of carbon (₆C¹⁴) is radioactive. The age of antique articles is determined by the ratio of ₆C¹² and ₆C¹⁴ present in them. This process is called radiocarbon dating.

3. Silicon is the second (about 27.7% by mass) most abundant element (next to oxygen) in the earth’s crust and is present in the form of silica and silicates. Germanium occurs only in traces (1.5 ppm). Both germanium and silicon in very pure form find applications as semiconductors.

4. The natural abundance of tin and lead is very low (2 and 13 ppm respectively). The principal ore of tin is tinstone or cassiterite (Sn02) and that of lead is galena (PbS). Both tin and lead form several alloys. Tin is also used for tin plating while some lead-containing compounds are used as the constituents of paints.

5. The first two elements of this group, carbon and silicon are non-metals, germanium is a semi-metal (metalloid) while tin and lead are metals

General Trends In Atomic And Physical

Properties Of Group 14 Elements

Some important physical and chemical properties of group-14 elements are given in the following table. The trend in properties may be largely understood from their electronic configurations. The effect of inadequate shielding by d – and f- electrons is prominent in the case of Ge and Pb.

Some atomic and physical properties of group-14 elements

Trends in various atomic and physical properties of group -14 elements and their explanation

1. Atomic radii

1. The atomic radii of group-14 elements are smaller than the corresponding elements of group-13.

Atomic radii  Explanation 1:

On moving from a group-13 to a group-14 element within the same period, the magnitude of nuclear charge increases. As a result, the outermost electrons experience greater nuclear charge and are pulled more strongly towards the nucleus. Consequently, the atomic radius decreases.

2. On moving down the group, the atomic radii of the group- 14 elements increase regularly.

Atomic radii  Explanation 2:

The increase in atomic radii down the group is due to the addition of new electron shells. However, the increase in atomic radii from Si to Pb is small and this is because the effective nuclear charge somewhat increases due to ineffective shielding ofthe intervening d and /-electrons.

2. Ionisation enthalpy

1. First ionisation enthalpies of group-14 elements are higher than those of corresponding members of group-13.

Ionisation enthalpy Explanation 1:

This is because of the greater nuclear charge and smaller atomic size of group-14 elements as compared to the corresponding members of group-13.

2. In group 14, ionisation enthalpy decreases from C to Sn and then increases from Sn to Pb. The overall order is

Ionisation enthalpy Explanation 2:

Because of the increase in atomic size and the screening effect of inner electrons which outweigh the impact of increased nuclear charge, the ionisation enthalpies decrease down the group from C to Sn. However, a small increase in ionisation enthalpy from Sn to Pb is because the effect of increased nuclear charge (82-50 = 32 units), in this case, outweighs the inadequate shielding effect of intervening d and /-electrons.

3. Electropositive or metallic character

The group-14 elements are less electropositive and hence less metallic than the elements of group-13. Again, on moving down the group, the metallic character increases from C to Pb.

Electropositive Explanation:

The less electropositive character of group-14 elements is due to the smaller size of their atoms and higher ionisation enthalpies as compared to those of the corresponding group-13 elements. On moving down the group, the electropositive character increases as the size of the atom increases and ionisation enthalpy decreases.

4. Electronegativity

Because of the smaller size, the elements of group 14 are more electronegative than the corresponding elements of group 13. Values of electronegativity, however, do not decrease regularly down the group.

Carbon is the most electronegative element (2.5) while the remaining elements possess almost die same value (1.8 to 1.9) because of poor shielding by d -and f-electrons.

5. Melting and boiling points

The elements of group 14 have higher melting and boiling points as compared to the corresponding elements of group 13.

This is because the atoms of group-14 elements form a greater number of covalent bonds (four) with each other as compared to group-13 elements (which form only three bonds) and hence there exist strong binding forces between their atoms both in the solid as well as in the liquid states. Further, the melting and boiling points decrease down the group due to an increase in atomic size and consequent decrease in the interatomic forces of attraction. However, the melting point of Pb is higher than that of Sn.

6. Catenation

The self-inking property of the elements by which their atoms mutually combine to form long open chains (straight or branched) and rings of different sizes is called catenation. Carbon has the maximum tendency to catenate and because of this carbon forms a vast number of organic compounds. The catenation property of group- 14 elements follows the order: C>>Si > Ge = Sn. Lead does not exhibit the property of catenation.

Catenation Explanation:

The tendency of an element to undergo catenation increases with an increase in the strength of the M—M bond, where M represents an atom of the given element. As the size of the carbon atom is very small, the C—C bond is quite strong and the formation of this bond is thermodynamically very favourable.

However, on moving down the group, the atomic size increases and the strength of the bond decreases and consequently, the tendency to exhibit the property of catenation decreases

C —C (348 kj-mol^-1) > Si—Si  (297kJ-mol^-1) > Ge —Ge  (260kJ-mol^-1)> Sn—Sn (240kJ-mol^-1)

7. Oxidation state

1. Carbon and silicon exhibit a +4 oxidation state. Due to the inert pair effect, the other three elements ofthe group (Ge, Sn and Pb) show +2 and +4 oxidation states. This is due to the weak shielding effect of- and- electrons.

2. On moving down the group from Ge to Pb, the number of d- and f- electrons increases due to which the inert pair effect increases gradually. As a result, the stability of the +4 oxidation state decreases. The increasing order of the +2 oxidation state is Ge < Sn < Pb. Lead is most stable at +2 oxidation state.

3. The compounds in which the group-14 elements are in a +4 oxidation state, are covalent since the charge on the element is high and its size is small. However, those compounds in which the oxidation state is +2, are ionic, since the charge of the element is small and its size is large. For example, SnCl₂ is ionic whereas SnCl₄ is a covalent liquid. On moving down the group, the tendency of forming covalent compounds decreases and that of forming ionic compounds increases.

8. Multiple bonding

Carbon, the first member of the carbon family, forms multiple bonds with carbon (C=C, C ≡ C), oxygen (C=O) and nitrogen (C=N, C≡ N ) and this occurs because of comparable sizes and energies of overlapping orbitals (2p ). However, Si and other members possess p -orbitals having comparatively bigger sizes and higher energies (3/7, 4p, 5p etc.) as compared to that of carbon.

Therefore, sideways overlap is not effective and consequently, multiple bonding does not take place. However, in some compounds, silicon becomes involved in multiple bonds with oxygen and nitrogen atoms and these are called pn-in bonds.

Example:

This can be illustrated by the structure of trisilylamlne, N(SiH3)3. The central N-atom of this molecule is sp² – hybridised. It is, therefore, a trigonal planar molecule. A lone pair of electrons present in the 2p -orbital of the N-atom becomes involved in sideways overlap with the vacant 3d orbitals of Si leading to the formation of pπ-dπ back-bonding (shown by dotted lines). Since the availability of the unshared pair of electrons on N (engaged in pn-dn backbonding) is reduced, the compound becomes weakly basic.

The central N-atom of trimethylamine, N(CH₃)₃, on the other hand, is sp3 -hybridised. The three hybrid orbitals are involved in the formation of cr -bonds with three carbon atoms and the unshared electron pair exists in the fourth sp³ -hybrid orbital. Carbon does not possess any low-energy vacant d -d-orbital suitable for -bonding. Therefore, the shape of the molecule is pyramidal. Since the availability of the unshared electron pair on N is much higher, the molecule is strongly basic

Chemical Properties Of Carbon Family

1. Nature of compounds

1. The general valence shell electronic configuration of the elements of the carbon family is ns²np². This indicates that their atoms have four valence electrons.

2. Compounds of these elements are normallynotionicin nature as the formation of both M⁴⁺ and M^4- ions require a large amount of energy, thus making it unfavourable.

3. Compounds formed by the elements of this family are generally covalent, in which they complete their octets by sharing valence electrons with atoms of other elements.

4. Carbon and silicon exhibit a +4 oxidation state. The remaining three elements of this group, i.e., Ge, Sn and Pb, however, exhibit +4 as well as +2 oxidation states. This is due to the inert pair effect arising out of poor shielding of valence s -electrons by intervening d – or f-electrons.

5. As the number of d or f-electrons increases down the group from Ge to Pb, the inert pair effect becomes gradually more prominent. As a result, the stability of the +4 oxidation state decreases and the +2 oxidation state increases.

The stability of the +2 oxidation state follows the order:

1. Ge < Sn < Pb. Lead is most stable in the +2 oxidation state. vii] The compounds of group-14 elements in which these elements exhibit a +4 oxidation state are expected to be covalent because of their extremely high charge. However, the compound in which these elements exhibit an oxidation state of +2 is expected to be ionic because of its large size and a small charge

2. SnCl₂, for example, is anionic compound while SnCl4 is a covalent compound. Again, on moving down the group, the tendency of the elements to form covalent compounds decreases whereas the tendency to form ionic compounds increases

3. Since carbon has no d -d-orbital, it cannot expand its valence shell and hence its maximum covalency or coordination number is four. However, due to the availability of vacant d -d-orbitals, Si, Ge, Sn and Pb can form hexa-coordinated complexes by increasing their coordination number from 4 to 6. For example, [SiF₆]^2-, [GeF₆]^2-, [Sn(OH)₆]^2-, [PbCl₆]^2-etc

2. Trends in oxidising and reducing properties

1. Because of the inert pair effect, elements like Ge, Sn and Pb exhibit +2 and +4 oxidation states. Therefore, these elements in the +2 oxidation state act as reducing agents (M→ M⁴⁺+2e) while in the +4 oxidation state, they act as oxidation agents (M⁴⁺+2e →M²⁺).

2.  The +4 oxidation state of get is most stable followed by Sn and Pb. It thus follows that in group -14, Ge(2) salts are the strongest reducing agents followed by Sn(2) salts. However, germanium is much less abundant in nature than tin and therefore, Sn(II) salts such as SnCl₂ are largely used as reducing agents.

For example:

⇒  \(2 \mathrm{FeCl}_3+\mathrm{SnCl}_2\rightarrow 2 \mathrm{FeCl}_2+\mathrm{SnCl}_4 \)

⇒  \(2 \mathrm{HgCl}_2+\mathrm{SnCl}_2\rightarrow \mathrm{Hg}_2 \mathrm{Cl}_2+\mathrm{SnCl}_4 \)

⇒ \(\mathrm{Hg}_2 \mathrm{Cl}_2+\mathrm{SnCl}_2\rightarrow 2 \mathrm{Hg}+\mathrm{SnCl}_4\)

3. The +2 oxidation state of Pb is most stable followed by Sn and Ge. Therefore, Pb(IV) salts such as lead tetraacetate, Pb(OCOCH3)4 and Pb02 are largely used as oxidising agents.

For example:

Pb02 + 4HC1 (cone.)- PbCl2 + Cl2 + 2H20

Formation of oxides

When the elements of group 14 are heated in oxygen,

They form two types of oxides:

1. Monoxides of the formula MO and
2. Dioxides ofthe formula MO.

Lead also forms another oxide called trialled tetraoxide or red lead (Pb304) which can be obtained by heating litharge (PbO) more than O₂ or air at 673 K

The acid-base character of the oxides:

1.  The oxides of these elements in higher oxidation states are generally more acidic than those in lower oxidation states.

2. Again, the acidic character of the oxides decreases on moving down the group. The dioxides such as CO₂ and SiO₂ are acidic, GeO₂ is less acidic than SiO₂ while SnO₂ and PbO₂ are amphoteric. Being acidic, the dioxides of mC, Si and Ge react only with bases.

CO₂ + 2NaOH→Na₂CO₃+ H₂O

CO₂ + Ca(OH)₂→CaCO₃ + H₂O

SiO₂ + 2NaOH→Na₂SiO₃ + H₂O

GeO₂+ 2NaOH→Na₂GeO₃ + H₂O

3. Being amphoteric, the dioxides of Sn and Pb react with both acids and bases

SnO₂+ 2NaOH→Na₂SnO₃ (Sodium stannate)+ H₂O

SnO₂ + 4HCl→ SnCl₄ + 2H₂O

PbO₂ + 2NaOH→Na2PbO₃ (Sodium plumbate)+ H₂O

4. Among the monoxides of these elements, CO is neutral, GeO is acidic while SnO and PbO are amphoteric.

Oxidising and reducing the power of oxides:

1. Among the monoxides of group-14 elements, CO is the strongest reducing agent because the +4 oxidation state is the most stable for carbon. CO is used in the extraction of many metals from their oxides.

For example:

2. Among the dioxides of group-14 elements, Pb02 is the strongest oxidising agent because of the inert pair effect. It oxidises HC1 to Cl2 and reacts with concentrated H2S04 or HN03 to liberate dioxygen.

3. Formation of halides

The elements of group-14 combine with halogens to form halides of the formula MX4 and MX2 (X = F, Cl, Br or I). All the elements of this except carbon react directly with halogens under suitable conditions.

Tetrahalides ofgroup-14 elements:

1. All the elements of this group form tetrahalides of the formula MX4.

2. These tetrahalides are mostly covalent. The central atom of these halides is sp3 -hybridised and the molecules are tetrahedral shape. Thein two exceptions are SnF4 and PbF4 which are more.

3. IoThe ionic character and thermal stability of these halides decrease with an increase in n atomicnumbatomic number of halogen atatomsExample:

PbCl₄ is a stable compound but PbBr₄ is an unstable compound while Pbl4 is not known. Because of the strong oxidising power of Pb(+4) and the strong reducing power of I^–, Pbl₄ does not exist. The unstable nature of PbBr₄ is because Pb(+4) is strongly oxidising and Br is weakly reducing.In both the cases, the halide ion(I^– or Br^– ) reduces Pb⁴⁺ to Pb²⁺ (Pb⁴⁺ + 2X^–→Pb²⁺+ X₂)

Then on-existence of Pbl₄ can be explained alternatively as follows— The amount of energy released by the initial formation ofthe Pb—I bond is not sufficient to unpair 6s² electrons and excite one of them to 6p -orbital to have four unpaired electrons around the Pb-atom

4.  Carbon tetrachloride (CCl₄) does notundergo hydrolysis. However, the tetrahalides of other elements undergo ready hydrolysis. For example:

SiCl₄ + 4H₂O→Si(OH)₄ + 4HCl

SnCl₄+ 2H₂O→SnO₂ + 4HCI

Explanation:

That CCl₄ does not undergo hydrolysis can be explained by the fact that carbon having no vacant d -d-orbital cannot expand its coordination number beyond 4. However, silicon can undergo ready hydrolysis because it can expand its octet (i.e., coordination number beyond four) due to the availability ofvacant d -orbitals. reaction

CCl₄+ 2H₂O → Noreaction

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

The reaction proceeds through the following steps:

1. In the first step, a lone pair of electrons from the oxygen atom of the H20 molecule is donated to an empty d -d-orbital of Si, forming a five-coordinate intermediate which has a trigonal bipyramidal shape.
2. In the second step, the intermediate loses a molecule of HCl and in this way, one Cl-atom of SiCl₄ is displaced by one OH group.
3. The remaining three Cl-atoms are displaced successively by OH groups yielding silicic acid.

5. Since carbon has no vacant d -orbitals in its valence shell, its tetrahalides do not form complexes. However, tetrahalides of other elements of this group form complexes due to the availability of vacant d -d-orbitals in their valence shells. These can, therefore, increase their coordination number up to six.

In other words, the tetrahalides of carbon do not act as Lewis acids but the tetrahalides of the other elements act as strong Lewis acids. The tetrahalides of Si, Ge, Sn and Pb, for example, can form hexahalo complexes like [SiF₆]^2-_, [GeF₆]^2-, [SnCl₆]^2-_ and[PbCl₆]^2- with corresponding halide ions.

SiF₄+ 2HF→H₂SiF₆ (hydrofluorosilicic acid) (or hexafluorosilicic acid)

SnCl₄ + 2Cl→SnCl^2-₆ [Hexachlorostannate (IV) ion]

Dihalldes of group-14 elements:

All the elements of group-14 except carbon and silicon form dihalides of the formula MX2. On moving down the group, the stability of these dihalides increases steadily due to the inert pair effect, i.e., the stability follows the order:

GeX₂<<SnX₂ < PbX₂

Anomalous Behaviour Of Carbon

Carbon, die first member of group 14, differs from the remaining members of its family in many properties.

Reasons for anomalous behaviour of carbon: 

1. Its atomic size is exceptionally small.
2. It has higher ionisation enthalpy and electronegativity.
3. It has no vacant dorbital in its valence shell.
4. It tends to form multiple bonds

Anomalous properties of carbon:

Allotropic Forms Of Carbon

The phenomenon of the existence of an element in two or more forms having different physical but similar chemical properties i? called allotropy and the different forms are called allotropes.

Carbon exists in several allotropic forms which may be classified as:

1. Crystalline and
2. Amorphous.

The four crystalline allotropic forms of carbon are:

1. Diamond
2. Graphite,
3. Fullerene
4. Carbon nanotubes.

The four amorphous allotropic forms of carbon are:

1. Charcoal
2. Soot or lamp black
3. Coke and
4. Gas carbon.

Amorphous carbon is not pure and remains mixed with various elements and compounds. Finer X-ray studies have shown that the amorphous varieties of carbon are composed of very minute crystalline units like graphite which are distributed throughout their masses in a most disordered fashion.

A synopsis of various allotropes of carbon is given in the following chart:

Crystalline Carbon Diamond

Diamond is a very precious substance which is availablein South Africa, New South Wales, Brazil, the Ural mountains and at Golconda in India. Two varieties of natural diamonds are available. One is the lustrous and colourless (or slightly coloured) variety which is generally used as precious gem¬ stones and the other is the black or deep-coloured opaque variety, known as the carbonado robot.

The weight of precious diamond is expressed in carats (1 carat = 200mg).

Some well-known diamonds are:

Cullinan (3032 carats), Kohinoor (present weight 106 carats), Pitt (136.25 carats), Regent (193 carats), Orloff (193 carats) and Great Mogul (186 carats). Diamond can also be prepared artificially but because of high cost and poor quality is seldom made artificially

Diamond Physical properties

1. Diamond is transparent, lustrous and crystalline. It may be colourless or slightly yellow coloured although some black or dark coloured varieties are also available.
2. It is the hardest naturally occurring substance known and it has a very high melting point (3843 K).
3. It is the heaviest among all the allotropic forms of carbon; its density is 3.51 gem^-3
4. Diamond has a very high refractive index (2.417) and thus light passing through it, suffers total internal reflection innumerable times. For this reason, diamonds appear to be extremely bright and lustrous.
5. It is transparent to X-rays and this property helps to distinguish a real diamond from an artificial one (made of glass).
6. Diamond is a non-conductor of electricity but good conductor of heat

Chemical properties:

At ordinary temperatures, diamond is chemically inert. It does not react with acids, alkalies, chlorine, potassium chlorate etc. However, it reacts with certain substances at much higher temperatures.

1. It is oxidised by oxygen at 800-900°C to produce pure carbon dioxide:

C + O₂ → CO₂

2. Diamond is converted into graphite in the absence ofair at much higher temperatures

Molten diamond can be converted into graphite by applying heat but graphite cannot be converted into diamond by heating to a very high temperature. The change is, therefore, unidirectional and this is because graphite is thermodynamically more stable than diamond. This type of allotropy is known as monotropy.

3. When a diamond is reacted with molten sodium carbonate, sodium monoxide and carbon monoxide are produced

C + Na₂CO₃ → Na₂O+2CO

4. It undergoes oxidation by fluorine at 700°C to form carbon tetrafluoride

C+ 2F₂→ CF₄

5. At about 250°C, diamond gets oxidised by a mixture of K₂Cr₂O₇ and concentrated sulphuric acid (i.e., chromic acid) to CO₂

6. Diamond reacts with sulphur vapour at 1000°C to form carbon disulphide: C + 2S→ CS₂

Structure of diamond:

1. In diamond, each C-atom is sp³ – hybridised and linked to four other C-atoms tetrahedrally by covalent bonds.

2. The value of each C—C—C bond angle is 109°28′ and each C —C bond distance is 1.54Å

3. An innumerable number of such tetrahedral units are linked together to form a three-dimensional giant molecule containing very strong bonds extended in all directions. Because of such a three-dimensional network of strong covalent bonds, diamond is extremely hard.

4. Since a huge amount of thermal energy is required to break a large number of strong covalent bonds, its melting point is very high.

5. All 4 valence electrons of each sp3 -hybridised C-atom in diamond crystal participate in forming covalent bonds and there is no free electron on any carbon atom. Thus, a diamond is an anon-conductor of electricity. vi] Diamond has the highest known thermal conductivity because its structure distributes thermal motion in three dimensions very effectively.

6. Unlike graphite in which the C-atoms are arranged in different distant layers, the C-atoms in diamond are placed at a covalent bond distance (1.54A). Because of this, the density of diamond is higher than that of graphite

Diamond Uses:

1. Because of its transparency, dazzling lustre and beauty, diamond is extensively used as precious gemstone.
2. Because of its extreme hardness, it is used for cutting glass, polishing hard surfaces and drilling purposes. Black or dark-coloured diamonds are generally used for this purpose.

Graphite

Graphite is available as minerals in Sri Lanka, Mexico, Italy, California (U.S.A), Siberia, Korea, Spain and India. The word graphite originates from the Greek word ‘graph’ which

Preparation of artificial graphite: Acheson process:

• In this process, coke dust mixed with silica is heated to a temperature of 3000-3500°C with the help of electrodes in an electric furnace made of fire bricks for 25-30 hours.
• The mixture is kept covered by sand.
• In the first stage of the reaction, silica reacts with carbon to form silicon carbide (SiC) and carbon monoxide (CO).
• Silicon carbide thus formed decomposes to yield graphite and silicon. Q
• At higher temperatures, silicon, on being vapourised, escapes from the furnace and graphite is left.

SiO₂ + 3C→ SiC + 2CO↑; SiC→Si + C [graphite]

Graphite Physical Properties 

1. Graphite is a dark greyish-coloured opaque, soft and slippery crystalline substance possessing metallic lustre.
2. It is lighter (density 2.25 g-cm-3 ) than diamond.
3. It is a good conductor of heat and electricity.

Graphite Chemical properties

1. Graphite is more reactive than diamond

2. When graphite is heated in air at 700°C, it is oxidised to carbon dioxide:

C + O→CO₂

3. At 500°C, fluorine reacts with graphite to produce carbon tetrafluoride (CF4). The compound is a non-conductor of electricity. It is also called graphite fluoride.

4. Graphite is not attacked by dilute acid or alkali. However, when it is subjected to react with molten sodium carbonate, carbon monoxide is formed.

5. Graphite is oxidised to carbon dioxide with a mixture of K₂Cr₂O. and concentrated H₂SO₄ (chromic acid).

6. When graphite is heated in the presence of a mixture of cones. nitric acid and sulphuric acid containing a small amount of potassium chlorate, greenish-yellow-coloured solid graphitic acid (CnH₄O₅) is obtained. Its exact structural formula is not known yet. Graphite can be identified by this test (diamond does not respond to this test).

7. On complete combustion, graphite produces mellitic acid
[C₆(COOH)₆].

Structure of graphite

1. Each carbon atom in graphite is sp2 -sp2-hybridised and is linked to three other carbon atoms directly in the same plane forming a network of planar hexagons and these two-dimensional layers exist in different parallel planes.

2. In each layer, the C—C bond length is 1.42Å and the distance between two adjacent layers is 3.35Åwhich is greater than the C—C covalent bond distance. So, the layers are supposed to be held together by relatively weak van der Waals forces of attraction

3. As the distance between two parallel layers is sufficiently large, graphite is less dense than diamond.

4.  Since the layers are weakly held together, on application of pressure, one layer easily slides over the other. Thus, graphite is found to be soft and lubricating

5. In the formation of hexagons in a layer of graphite, only three of its four valence electrons are used to form three sigma bonds (Csp²-Csp²). The remaining electrons of each carbon atom present in an unhybridised p -p-orbital are utilised to form n -n-bonds.

6. The n -n-electrons are mobile and can move freely through the graphite crystal. Because of the presence of free mobile electrons, graphite is a good conductor of electricity and heat.

7. Of all the crystalline allotropes of carbon, graphite is thermodynamically the most stable one. Its standard enthalpy of formation (AfH°) is taken as zero.

Graphite Uses:

• Graphite is largely used for lining and making electrodes for electric furnaces.
• When mixed with oil and water, graphite is used as a lubricant in machinery.
• It finds use in making crucibles resistant to high temperatures.
• By mixing with desired quantities ofwax or clay, graphite is used for making cores of lead pencils
• Graphite is used as a moderator in nuclear reactors

Fullerenes

Fullerenes or Buckminsterfullerene (named after the famous American designer of the geodesic dome, Robert Buckminster Fuller) is the latest allotrope of carbon discovered in 1985 collectively by three scientists namely R. E. Smalley, R. F. Curl and H. W. Kroto. It is a crystalline allotrope of carbon in which the carbon atoms exist in a cluster form. It is also known to be the purest form of carbon because, unlike diamond and graphite does not have surface bonds that are to be attracted by other atoms.

Structure of fullerenes C₆₀:

Fullerenes are expressed by the general formula C_n, where n is an even number between 30-600, for example, C₆₀, C₇₀, C₈₀….etc.

All these are cage-like spheroidal molecules having polyhedral geometry containing pentagonal and hexagonal planes. Number of hexagons in a C„ molecule =(n/2- 10),

Example:

In fullerene C₆₀ molecule, number ofhexagons =(60/2-10) = 20. Structure of fullerene C₆₀; The C₆₀ molecule consists of twenty-six-membered rings and twelve five-membered rings of sp² -hybridised carbon atoms fused into each other

Each carbon atom forms three or -bonds with the other three carbon atoms and the remaining electron on each carbon is involved in the formation of the n -bond and as a result, the system is expected to be aromatic. However, it is not aromatic because the molecule is not planar and it does not have (4n + 2) electrons.

It is a non¬ aromatic system. This fusion pattern provides a marvellous symmetry to the structure in which the fused ring system bends around and closes to form a soccer ball-shaped molecule (“buckyball”). Of all the fullerenes, C60 is the most stable one.

Structure of fuUerene C₇₀:

The molecule acquires the shape ofa rugby ball. It consists of twelve five-membered rings and twenty-five six-membered rings and their arrangement is the same as that of a C₆₀ molecule.

C₇₀ Preparation:

• The preparation of fullerenes involves heating of graphite in an electric arc in the presence of an inert gas such as He or Ar.
• A sooty material is recovered which consists mainly of C₆₀ with a small amount of C₇₀ and traces of other with an even number of C atoms up to 350 and above.
• The C₆₀ and C₇₀ fullerenes can further be separated from the sooty material by extraction with benzene or toluene followed by chromatographic separation using alumina (Al₂O₃) as the adsorbent.
• In Russia, America, Canada and New Zealand C₆₀ and C₇₀ fullerenes are isolated from natural sources. Fullerenes of this type are formed by the red giant star Antares.

Properties and applications: 

• Fullerenes are solids with high melting points.
• Being covalent, they are soluble in organic solvents.
• They react with alkali metals to form solid compounds such as K₃C₆₀. This compound acts as a superconductor even at temperatures of the order of 10-40K.
• Because of their spherical shapes, they exhibit wonderful lubricating property

Carbon Nanotubes

Carbon nanotubes are crystalline allotropes of carbon with cylindrical nanostructure. This allotrope was discovered by Sumio Iijima (Japan) in 1991. A carbon nanotube consists of a two-dimensional array of hexagonal rings of carbon just as in a layer of graphite or a chicken wire.

The layer is then rolled. into a cylinder and capped at each end with half of a C₆₀  fullerene

Properties and applications: 

The nanotubes are approximately 50000 times thinner than a human hair. These are very tough, about 100 times as strong as steel. They are electrically conducting along the length of the tube.

These cylindrical carbon molecules having unusual properties are valuable for nanotechnology, electronics, optics and other fields of material science and technology. They are also being used as probe tips for the analysis of DNA and proteins by atomic force microscopy (AFM). Many other applications have been envisioned for them as well, including molecular-size test tubes or capsules for drug delivery

Amorphous Carbon Charcoal

Vegetable charcoal

1. Wood charcoal: When wood is subjected to destructive distillation in an iron retort, the volatile organic compounds present escape and the residue left in the retort is called wood charcoal.
2. Sugar charcoal: It can be prepared by heating pure sugar in a closed vessel or by eliminating water from sugar by reacting it with concentrated sulphuric acid

Animal charcoal

1. Bone charcoal:

• Small pieces of animal bones are treated with superheated steam to remove the adhering fat and marrow.
• The dried bones are then subjected to destructive distillation in an iron retort when volatile substances are distilled out and a black residue containing carbon and about 90% impurities like calcium phosphate and calcium carbonate are left behind.
• This is known as bone black or bone charcoal. These impurities are removed by dissolving the black material in dilute HCl.
• The insoluble deep black powder thus obtained is almost pure charcoal and is called ivory black.

2. Blood charcoal:

• Charcoal obtained by destructive distillation of blood is known as blood charcoal.

Charcoal Physical properties

Charcoal is black, soft and porous. It is a bad conductor of heat and electricity. Its specific gravity lies between 1.4 and 1.9. But because of its porosity, air enters in its pores.

• As a result, its specific gravity gets reduced to 0.2 and hence charcoal floats on water. When porous charcoal b freed from entrapped air, it can retain any other ga In its pores, This phenomenon is known as adsorption.
• The adsorbed gas escapes on heating and is much more reactive than the ordinary gas.

Activated charcoal:

Activated charcoal has high adsorption power as compared to ordinary charcoal.

Activated charcoal Preparation:

When the charcoal obtained by destructive distillation of coconut shell is heated to about 800-900°C in a limited supply of air or steam, activated charcoal is obtained.0 besides this, activated charcoal may also be obtained by the destructive distillation of sawdust soaked In aqueous solution of ZnCl₂or MgCl,

Activated charcoal Properties:

Activated charcoal is not only a good adsorbent for gases but it also has the power of decoloursing a coloured substance and absorbing the taste of a substance. Moreover, a catalyst accelerates the rates of many chemical reactions.

Charcoal Chemical properties

1. At higher temperatures, charcoal burns in air or oxygen to form C02 gas. However, in a limited supply of oxygen, it’s combustion produces carbon monoxide.

C +O₂ → CO₂; 2C + O→2CO

2. Charcoal forms different compounds with sulphur, nitrogen and hydrogen at high temperatures.

2C + N₂ → (CN) (Cyanogen)

2C + H₂ → C₂H₂ (Acetylene)

In the first reaction, two solids (C and S) react together to form a liquid (carbon disulphide

3. Charcoal combines with heated Ca, Al, Fe etc. to form their corresponding carbides.

Ca + 2C→CaC₂; 3Fe + C→Fe₃C; 4Al + 3C→Al₄C₃

4. When a mixture of silica (SiO₂) and coke dust is heated at 1500’2000°Cin an electric furnace, silicon carbide (SiC) is formed. It is a black, bright and very hard solid. It is known as carborundum which is used for polishing metals

SiO₂ +3C → SiC+2CO

5. When steam is passed over white-hot charcoal or coke, a mixture containing equal volumes of CO and H₂, called water gas is produced

6. Reducingproperty: Charcoal is a good reducing agent.

At higher temperatures, various metal oxides are reduced by charcoal to their corresponding metals.

CuO + C→ Cu + CO; PbO + C→ Pb + CO

Fe₂O₃ + 3C→2Fe + 3CO

At higher temperatures, charcoal reduces carbon dioxide to CO and sodium sulphate to sodium sulphide.

CO₂ + C→2CO;

⇒ \(\mathrm{Na}_2 \stackrel{+6}{\mathrm{~S}} \mathrm{O}_4+4 \stackrel{0}{\mathrm{C}} \rightarrow \mathrm{Na}_2S{ }^{-2}+4 \stackrel{+2}{\mathrm{~C}} \mathrm{O}\)

Charcoal in the burning condition is oxidised by concentrated nitric acid or sulphuric acid to CO₂

Charcoal Uses: 

• Wood charcoal is used as fuel and as a reducing agent in the extraction of metals.
• It is used in decolourising sugar syrup and refining oils, fats, glycerine etc.
• It is used in the treatment of drinking water as it adsorbs excess chlorine after chlorination.
• It is used in gas masks as it adsorbs poisonous gases.
• It is also used for preparing gunpowder and black paint (Black Japan).

Lamp Black or carbon black

When organic liquids rich in carbon such as kerosene, petrol, turpentine oil, benzene etc. are subjected to bum in controlled air, a black sooty smoke is produced. This smoke on condensation in a cold container forms soot. This soot is called lamp black or carbon black. It may also be obtained when natural gas (methane) is subjected to albumin-controlled air. It is the purest form of all the amorphous allotropes of carbon.

Lamp Black Properties:

It is amorphous, black and non-conductor of heat and electricity.

Lamp Black Uses:

It is used in the preparation of printing ink, shoe polish and black paints

Coke and gas carbon

When anthracite coal (96% carbon) is subjected to destructive distillation, the solid residue left in the iron retort is called coke. At higher temperatures (1000 -1200°C), hard coke is called coke. At higher temperatures (1000 -1200°C), hard coke is obtained whereas at 600-650°C, we get soft coke. The black hard dense residue deposited on the relatively cooler upper part Qf returns as gas carbon. it possesses thermal water gas is produced. and electrical conductivity and electrical conductivity.

Coke Uses:

• Hard coke is used as fuel and as a reducing agent for domestic fuel
• Gas carbon is widely used as electrodes in batteries, arc lights and during electrolysis.

All the allotropic modifications of carbon consist of the same element:

When a fixed weight of a pure allotropic form of carbon is heated in a long hard combustion tube in the presence of pure oxygen, CO₂ and CO are produced.

CuO, kept in the tube, converts CO into CO₂. CO₂ so obtained is absorbed in a previously weighed potash bulb attached to the exit end of the tube. An increase in the weight of the bulb increases the amount of CO₂ formed.

When the experiment is performed separately with the same fixed weight of any other allotrope, the same amount of CO₂ is obtained in each case. This experiment, thus, proves that the different allotropes consist ofthe same element, i.e., carbon.

Carbon Monoxide

Laboratory preparation:

1.  In the laboratory, carbon monoxide is prepared by dehydrating formic acid or oxalic acid after heating with concentrated sulphuric acid.

2. When potassium ferrocyanide is heated with an excess of cone, sulphuric acid, pure carbon monoxide is obtained

CO cannot be fired by concentrated sulphuric acid:

Concentrated sulphuric acid is a strong oxidising agent. Thus when CO ( a reducing agent )is passed through concentrated H₂SO₄, it is oxidised by sulphuric acid to CO₂

Other methods of preparation

From carbon:

When steam is passed over red hot coke, water gas or synthesis gas (CO + H₂) is obtained. CO is separated from the mixture by liquefaction.

When air is passed over hot coke, producer gas (CO + N₂) is formed. CO is separated by liquefaction

From carbon dioxide:

When CO₂ is passed over red hot carbon, zinc, Iron etc., it is reduced to CO.

CO₂ + C→2CO; CO₂ + Zn→CO + ZnO

CO₂+ Fe→FeO + CO

From metal oxides:

Carbon reduces die oxides of zinc, lead or iron to produce CO.

ZnO + C→ Zn + CO; Fe₂O₃+ 3C→2Fe + 3CO

From nickel tetracarbonyl:

Pure CO is obtained when nickel tetracarbonyl vapour is heated above 150°C

Carbon monoxide Physical properties:

1. Carbon monoxide is a colourless, tasteless, odourless gas which is lighter than air.

2. It is slightly soluble in water. It is a neutral oxide.

3. It is a highly poisonous gas. If only a volume of CO is present in 10,000 volumes of air, then that air is considered to be poisonous. Carbon monoxide molecule.

4.

with a lone pair of electrons on carbon combined with Fe-atom present in the haemoglobin of the blood to form a very stable complex compound named carboxyhaemoglobin.

5. Hb + CO → HbCO; (Hb =Haemoglobin) As CO is almost 100 times more rigidly bonded to Fe-atom than O₂, O₂ can no longer combine with haemoglobin.

6. In other words, haemoglobin fails to act as an oxygen-carrier. As a consequence, the body tissues become slackened due to lack of ofoxygen ultimately causing death

7. In case of CO poisoning, the patient should immediately be taken to an open area and artificial respiration with carbogen (a mixture of oxygen and 5-10% CO₂) should be started.

Carbon monoxide Chemical properties:

1. Combustion:

Carbon monoxide is itself a combustible gas but does not support combustion. It burns in the air with a blue flame and is oxidised to C02. Because of the evolution of a large amount of heat, CO is used as fuel.

2CO + O₂ → 2CO₂ + 135.2 kcal

The two important fuels containing carbon monoxide are water gas and producer gas. Water gas contains 50% of H₂, 40% of CO, 5% of CO₂ and 5% of CH₄ and N₂ while producer gas contains 25% of CO, 4% of CO₂,70% of N2 and traces of H₂, CH₄ and O₂

2. Reducing property:

Carbon monoxide is a powerful reducing agent. The oxidation number of carbonin CO is +2 and the highest oxidation number of carbon is +4. So, CO tends to be oxidised and behaves as a strong reducing agent. Various metal oxides are reduced by CO to the corresponding metal.

CuO + CO→Cu + CO₂; PbO + CO→Pb + CO₂

ZnO + CO→Zn + CO₂; Fe₂O₃ + 3CO→2Fe + 3CO₂

At 90°C, CO reduces iodine pentoxide (I₂O₅) to give violet-coloured iodine. This reaction is called the Ditte reaction.

I₂O₅ + 5CO→I₂ + 5CO₂

3. Reaction with sodium hydroxide:

Being a neutral oxide CO does not react with alkali or base under ordinary conditions. But at 200°C and under high pressure, it reacts with caustic soda solution to yield sodium formate.

4. Absorption of CO:

When CO is passed through an ammoniacal or acidified cuprous chloride solution, it gets absorbed in that solution to give a white crystalline addition compound as a precipitate. CO can be separated from a gas mixture by this process.

Cu₂Cl₂ + 2CO + 4H₂O→2[CuCl.CO. 2H₂O]↓

The addition compound evolves CO on heating.

5. Formation of addition compounds:

1. In the presence of sunlight, CO combines direedy with chlorine gas to form carbonyl chloride or phosgene gas. It is a colourless poisonous gas:

2. CO reacts with sulphur vapour to produce carbonyl sulphide.

3. Combines with many transition metals to form metal carbonyl compounds. For example, CO reacts with nickel powder at 30-40°C under ordinary pressure to form nickel tetracarbonyl. Again, at 200°C and 100 atmosphere pressure, CO reacts with freshly reduced iron to form pentacarbonyl.

Ni + 4CO→Ni(CO)₄; Fe + 5CO→Fe(CO)₅

4. Formation of organic compounds:

Hydrogen reacts with CO at 350°C in the presence of Ni or Pt catalyst to yield methane. If the reaction is carried out at 300°C and 200 atmospheric pressure in the presence of ZnO and Cr₂O₃ catalyst, methyl alcohol is produced. The oxidation number of carbon in CO decreases from +2 to -4 in methane and to -2 in methyl alcohol.

Therefore, in these two cases, CO exhibits its oxidising property.

⇒ \(\stackrel{+2}{\mathrm{CO}}+3 \mathrm{H}_2 \rightarrow \stackrel{-4}{\mathrm{CH}_4}+\mathrm{H}_2 \mathrm{O} ; \stackrel{+2}{\mathrm{CO}}+2 \mathrm{H}_2 \rightarrow \stackrel{-2}{\mathrm{CH}_3^{-}}\mathrm{OH}\)

Identification of carbon monoxide:

1. Carbon monoxide burns in air with a blue flame and the gaseous product turns lime water milky [H₂ also burns with a blue flame but in this case, steam is formed which turns white anhydrous copper sulphate blue.]

2. CO is completely absorbed by the Cu₂Cl₂ solution in a cone. hydrochloric acid or ammonium hydroxide and as a result, a white crystalline addition compound is precipitated.

3. When a filter paper soaked with a solution of platinum or palladium chloride is held in CO gas, the paper turns pink-green or black due to the reduction ofthe metal salts.

PtCl₂ + CO + H₂O→Pt + CO₂ ((pink-green))+ 2HCl

PdCl₂ + CO + H₂O→Pd + CO₂ ((black))+ 2HCl

4. When CO gas is passed through an ammoniacal AgN03 solution, the solution becomes brown

5.  When a dilute solution of blood shaken with CO, is subjected to spectroscopic analysis, the observed band in the spectrum indicates the presence of CO. The presence of traces of the air can be detected by this experiment.

6. The presence of a very small amount of CO in the air can be detected with the help of a halamite tube or colour detector tube. When air containing CO is introduced into this tube I₂ O₅ present in the tube reacts with CO to liberate I₂

Because of the violet colour of evolved I₂, colour of the tube changes and the presence of COin air is indicated

I₂O₅ + 5CO→I₂  (Ditte reaction)+ 5CO₂

Structure of carbon monoxide:

Both the carbon and the oxygen atoms in a CO molecule are sp -hybridised. One of the sp -hybrid orbital of each atom is used to form a C —O cr -bond while the other sp -orbital ofeach contains a lone pair of electrons. The two unhybridised 2p -orbitals of each atom are involved in the formation of two pn-pn bonds. In terms of resonance, the CO molecule can be best represented as a resonance hybrid of the following two I resonance structures(1 and 2 ).

The resonance structure (1) is relatively more stable because of the fulfilment of the octet of both atoms.

Uses Of carbon monoxide: 

• CO is used as fuel in the form of producer gas or water gas.
• It is used as a reducing agent in the extraction of metals.
• It is used for the preparation of pure nickel by Mond’s process.
• It is used for the
• Preparation of methanol, methane, formic acid and synthetic petrol (Fischer-Tropsch process).

Preparation of pure nickel:

Ni(CO)₄ is prepared by the reaction between impure nickel and carbon monoxide. Ni(CO)₄ is then allowed to decompose by heating to 1.50°C to get pure nickel.

Carbon Dioxide

Carbon Dioxide Laboratory preparation:

At ordinary temperature, CO₂ is prepared in the laboratory by the action of dilute HCl on calcium carbonate (CaCO₃) or marble.

CaCO₃ + 2HCl→CaCl₂ + CO₂↑ + H₂O

The gas is collected in the gas jar by the upward displacement of air, as it is 1.5 times heavier than air. Carbon dioxide thus produced contains a small amount of HCl and water vapour. The gas is then passed successively through NaHCO₃  solution and cone, sulphuric acid to remove HCl vapour and water vapour respectively.

Dilute sulphuric acid cannot be used for the preparation of CO₂  from marble or limestone:

This is because sulphuric acid reacts with CaCO₃ to produce insoluble; CaSO₄  which forms a layer of CaCO₃. This insoluble layer prevents CaCO₃  from reacting with the acid and as a result, the evolution of CO₂ ceases within a very short time

CaCO₃ + H₂SO₄→CaSO₄+ CO₂ + H₂O

On the other hand, when dilute hydrochloric acid is, used, highly soluble calcium chloride (CaCl₂) is formed. So, the reaction proceeds without any interruption

CO₂ can be prepared by the action of dilute H₂SO₄ on Na₂CO₃:

The salt, Na₂SO₄ produced soluble in water or dilute H₂SO₄

Na₂CO₃ + H₂SO₄→Na2SO₄ + CO₂ + H₂O

At ordinary temperatures, CO₂ is highly soluble in water. Therefore, it is not collected by the downward displacement of water. The solubility of COz in hot water is very low and hence it can be collected over hot

Other methods of preparation:

1. From carbonate salts:

Except for alkali metal carbonates, all other carbonates undergo thermal decomposition to produce CO₂ and the oxides ofthe corresponding metals.

BaCO₃ decomposes only at very high temperatures.

Calcium carbonate or limestone is thermally decomposed (1000°C) for the preparation of carbon dioxide on a commercial scale.

2. From bicarbonate salts:

Bicarbonates of all the elements decompose on heating with the evolution of CO₂

3. From fermentation:

A large amount of CO₂ is obtained as a by-product during the manufacture of ethyl alcohol by fermentation of sugar

4. From water gas:

Water gas is industrially prepared by passing steam through a bed of white-hot coke at about 100°C. C + H₂O→CO + H₂. When a mixture of water gas and excess steam is passed over (Fe₂O₃+ Cr₂O₃) catalyst heated at 400°C, CO is oxidised to CO₂

(CO + H₂) + H₂O→CO₂ + 2H₂

The gaseous product is then passed through a solution of potassium carbonate when CO₂ is completely absorbed and KHC03 is formed. H2 and unconverted CO pass out. When the resulting KHC03 solution is boiled, CO₂ is obtained.

K₂CO₂+ CO₂ + H₂O→2KHCO₃

Carbon Dioxide Physical properties

1. Carbon dioxide is a colourless, odourless and tasteless gas having slightly acidic properties.

2. CO₂ is 1.5 times heavier than air. So, this gas often accumulates in abandoned wells or pits and because of this, severe breathing problems are caused in such places.

3. By the application of pressure (nearly 40 atmospheric pressure and a temperature < 40°C), CO₂ can be easily liquefied. When liquid CO₂ is allowed to vaporise rapidly by releasing the pressure, it further gets cooled down and freezes like ice. This is called dry ice or cardice.

4. When solid carbon dioxide is allowed to evaporate at atmospheric pressure, it gets converted into the vapour state without passing through the intermediate liquid state. Therefore, unlike ordinary ice, it does not wet the surface of the substance and because of this, it is called dry ice.

5. It is highly soluble in water (1.7 cm3 of CO₂ dissolves in 1 cm3 of water). The solubility increases with an increase in pressure. Aerated waters such as soda water, lemonade etc. contain CO₂ under pressure. When the cork of the bottle of aerated water is opened, the pressure is released and excess CO₂ escapes in the form of bubbles. Its solubility in water, however, decreases with a temperature rise.

Carbon Dioxide Chemical properties

1. Combustion:

Carbon dioxide is neither combustible nor helps in combustion. When it (heavier than air) falls on a binning substance, it removes air from the surface of the substance and thereby the substance can no longer remain in contact with air. As a result, the fire is extinguished. A burning jute stick when inserted into a jar of CO₂, extinguishes.

However, when a burning Mgribbon or metallic sodium is inserted into a CO₂ jar, it continues to bum with the separation of black carbon.

⇒ \(\stackrel{+4}{\mathrm{C}} \mathrm{O}_2+2 \stackrel{0}{\mathrm{Mg}} \rightarrow 2 \stackrel{+2}{\mathrm{MgO}}+\stackrel{0}{\mathrm{C}} ; \stackrel{+4}{\mathrm{C}} \mathrm{O}_2+4 \stackrel{0}{\mathrm{Na}} \rightarrow 2 \stackrel{+1}{\mathrm{Na}}_2 \mathrm{O}+\stackrel{0}{\mathrm{C}}\)

During the burning of such metals, the temperature, due to the liberation of a large amount of heat, is so high that CO₂ decomposes into carbon and O₂ and it is the oxygen which helps in the burning ofthe metals.

In these reactions, CO₂ acts as an oxidising agent and itself gets reduced to carbon. These reactions prove the existence of carbon in C02. It is to be noted that the oxidation number of carbon in CO₂ is +4 and this is its highest state of oxidation.

Thus, there is no possibility of an increase in its oxidation number, i.e., CO₂ cannot be further oxidised. That is why CO₂ cannot exhibit any reducing property. For the same basic reason, C02 is not combustible [CO, on the other hand, is combustible because in this case, the oxidation number of carbon may increase from +2 to +4 ].

2. Acidic property:

Carbon dioxide is an acidic oxide. It dissolves in water forming an unstable dibasic acid called carbonic acid (H₂CO₃). CO₂ is, therefore, regarded as the anhydride of carbonic acid.

CO₃+H₂O→H₂CO₃

H₂CO₃ is known only in solution and when the solution is heated, CO₂ is evolved out The solution turns blue litmus red but it cannot change the colour of methyl orange.

H₂CO₃ forms two types of salts, bicarbonates (HCO₃ ) and carbonates (CO^2-₃). Being an acidic oxide, CO₂ combines directly with strongly basic oxides such as CaO, Na₂O etc. to form their corresponding salts.

CaO + CO₂→CaCO₃; Na₂O + CO₂→Na₂CO₃

Reaction with alkali:

When CO₂ is passed through a strong alkaline solution of NaOH, a carbonate salt is first formed. If the passage of CO₂ is continued for a long time, white crystals of sparingly soluble sodium bicarbonate are precipitated. The bicarbonate salt decomposes on heating to form carbonate salt, CO₂ and water.

2NaOH + CO₂→Na₂CO₂ + H₂O

Na₂CO₃ + CO₂ + H₂O→ 2NaHCO₂

Rection with lime water:

When CO₂ is passed through lime water, the solution becomes milky due to the formation of white insoluble calcium carbonate. However, when excess CO₂ gas is passed through this milky solution, its milkiness disappears as insoluble calcium carbonate gets converted into soluble calcium bicarbonate

Ca(OH)₂ + CO₂→CaCO₃↓ (white) +H₂O

CaCO₃+ CO₂ + H₂O→Ca(HCO₃)₂ (soluble)

On heating, calcium bicarbonate decomposes to form calcium carbonate, CO₂ and water and as a result, the clear solution becomes milky again.

Ca(HCO₃)₂→CaCO₃↓ + CO₂ + H₂O

Manufacture of sodium carbonate:

When CO₂ gas is passed through a concentrated solution of sodium chloride (brine) saturated with ammonia at 30-40°C, white crystals of sodium bicarbonate are precipitated. The reaction occurs in two stages

NH₃ + CO₂ + H₂O ⇌ (NH₄)₂CO₃

CaSO₄ + (NH₄)₂CO₃→CaCO₃↓ + (NH₄)2SO₄

Sodium carbonate is prepared by thermal decomposition of sodium bicarbonate. The Solvay process for the manufacture of sodium carbonate is based on this reaction.

Production of ammonium sulphate:

This is carried out by passing CO, and NH₃ gases through a slurry of powdered gypsum (CaSO₄,2H₂O) in water. At first, NH₃ and CO₂ react together in the presence of water to form ammonium carbonate. It then reacts with calcium sulphate (gypsum) to form calcium carbonate and ammonium sulphate by double decomposition.

2NH₃ + CO₂ + H₂O ⇌  (NH₄)₂CO₃

CaSO₄ + (NH₄)₂CO₃→CaCO₃ + (NH₄)2SO₄

The nitrogenous fertiliser ammonium sulphate is manufactured by using this reaction. In this process, (NH₄)₂SO₄ is produced without using H₂SO₄

Production of urea:

At 200-210°C and 150 atm pressure, CO₂ reacts with ammonia to produce urea.

CO₂ + 2NH₃ ⇌(Ammonium carbamate) NH₄ COONH₂ ⇌(Urea) CO(NH₂ )₂ + H₂ O

The important fertiliser, urea is manufactured on a large scale by using this reaction.

Photosynthesis:

Plants absorb atmospheric carbon dioxide. In the presence of chlorophyll and sunlight, the absorbed CO₂ combines with water (absorbed from the soil) to form glucose, water and oxygen. This process is called photosynthesis. In this process, CO₂ is reduced to carbohydrates by water

Reduction of CO₂ :

When CO₂ is passed over heated C, Fe, Zn etc., it is reduced to CO

Identification of carbon dioxide:

• It extinguishes a burning stick.
• Lime water becomes turbid when CO₂ is passed through it. When excess CO₂ is passed through it, the turbidity disappears but when that clear solution is boiled, the turbidity reappears.

N₂gas also extinguishes burning sticks but it does not turn the water milky. Again, SO₂ gas also turns lime water milky but unlike CO₂, it reacts with an acidified solution of potassium dichromate and changes the colour of the solution from orange to green

Uses Of carbon dioxide:

1. CO₂ is used in the manufacture of sodium carbonate by the Solvay process and also for the manufacture of fertilisers such as urea, ammonium sulphate etc.

2. CO₂ is used in fire extinguishers.  It finds extensive use in the preparation of aerated waters such as soda water, lemonade etc… And baking powder.

3. Solid carbon dioxide i.e., dry ice is used as a refrigerant under the commercial name drikold. Dry ice is also used for making cold baths in the laboratory by mixing it with some volatile organic solvents. It is extensively used as a coolant for preserving perishable articles in the food industry, for curing local burns and for surgical operations of sores.

4. Supercritical CO₂ is used as a. solvent to extract organic compounds from their natural sources, for example, caffeine from coffee beans, perfumes from flowers etc.

5. It is used under the name carbogen (a mixture of 95% O₂ and 5% CO₂) for the artificial respiration of patients suffering from pneumonia and affected by poisonous gases (CO poisoning).

Liquid CO₂ is used as a substitute for chlorofluorocarbons in aerosol propellants.

Fife extinguisher:

It is a specially designed metallic pressure vessel having a nozzle at one end. A glass bottle containing dilute sulphuric acid is placed inside it and the remaining portion of the vessel is filled with a concentrated solution of sodium bicarbonate. When required, the glass bottle can be broken by pressing a knob fitted with the vessel at the other end. When the glass bottle is broken, the add comes in contact with sodium bicarbonate solution and reacts to yield copious C02 gas. The gas, ejected under high pressure through the nozzle, falls on the burning substance and as aresult, the fire gets extinguished

Na₂CO₃ + H₂SO₄→Na₂SO₄ + CO₂↑ + H₂O

Baking powder:

The baking powder is used. Fire extinguisher in the preparation of bread consists of a dry mixture of potassium hydrogen tartrate, NaHCO₃, tartaric acid and -starch. When this mixture comes in contact with water present in the bread, a chemical reaction leading to the formation of CO₂ occurs. The resulting CO₂ gas evolved in the form of bubbles making the bread porous and soft. Moreover, NaHC03 and tartaric acid also produce C02 on thermal decomposition

Structure of carbon dioxide:

In a CO₂ molecule, the carbon atom is sp -hybridised whereas the oxygen atoms are sp² – hybridised. Carbon forms two cr -bonds and two pπ- pπ bonds with two oxygen atoms. The shape of the carbon dioxide molecule is, therefore, linear. The molecule is symmetrical (the two bond moments cancel each other) and hence, it is non-polar. The C —O bond length is 1.15Å. CO₂ can be represented as a resonance hybrid of the following three structures:

Compounds Of Silicon

1. Silicon tetrachloride (SiCI₄ )

Silicon tetrachloride Preparation:

Silicon tetrachloride is prepared by heating either silicon or silicon carbide with chlorine

Silicon tetrachloride Properties and uses:

1. Physical state:  It is a volatile liquid (boiling point: 330.5 K).

2. Hydrolysis:

SiCl₄ undergoes ready hydrolysis to produce silicic acid, Si(OH)₄ which on further heating undergoes partial dehydration to yield silica gel (SiO₂  xH₂O).

Silica gel is an amorphous and very porous solid which contains about 4% of water. It is used as an adsorbent in column chromatography and as a catalyst in the petroleum industry. When the hydrolysis of SiCl₄ is carried out at a much higher temperature, finely powdered silica Is obtained instead of silicic acid

The finely powdered silica thus obtained is used as a thixotropic agent (which reduces viscosity temporarily) in polyester, epoxy paints and resins and as an inert filler in silicon rubber.

3.  Reduction: Reduction of SiCl₄ with H₂ gas gives silicon

Ultrapure silicon used for making transistors, computer chips and solar cells is prepared by this method.

4. Reaction with silicon:

When a mixture of SiCl₄ and Si is pyrolysed, a series of perhalosilanes of the general formula, Si_n Cl_{2n+ 2} where n = 2-6, are obtained.

Chains of perhaloslianes are longer than those of silanes and this is due to the formation of pπ-dπ bonding between a lone pair of electrons present on Cl and the empty d-orbitals of Si

2. Silicones

Silicones e Definition:

The synthetic organosilicon polymers containing repeating R₂SiO units held by Si — O — Si linkages are known as silicones The general formula of these compounds is (R₂SiO)n where R = methyl or aryl group. Commercial silicones are generally methyl derivatives and in some cases phenyl derivatives.

Silicones Preparation

Hydrolysis of dichlorodimethylsilane (obtained by the reaction between methyl chloride and silicon in the presence of Cu as catalyst) followed by polymerisation involving intramolecular dehydration yields straight chain polymers, i.e., silicones.

The length of the polymer can be controlled by the reaction of dimethylsiianol with chlorotrimethylsilane. This blocks the terminal end ofthe polymer as follows—

Silicones Properties:

• Silicones containing short chains are oily liquids; those with medium chains are viscous oils, greases and jellies and those with long chains are rubber-like solids.
• They are stable to heat and are also resistant to oxidation, i.e., they are very inert.
• They are water repellents (hydrophobic) & good electrical insulators.

Silicones Uses:

• Silicones are used for making water-proof cloth and paper.
• These are used as electrical insulators.
• Silicon oils are used in high-temperature baths and vacuum pumps.
• Silicon rubbers are very useful as they can retain their elasticity over a wide range of temperatures.
• These are mixed with paints and enamels to make them resistant to the effects of sunlight, high temperatures and chemicals.
• These are used for preparing vaseline-like greases which are used as lubricants in aeroplanes

3. Silicates

Silicates Definition:

Silicates are compounds in which the anions present are either discrete SiO₄^4- tetrahedral units or several such units joined together by corners, i.e., by sharing one oxygen atom but never by sharing edges The negative charge on the silicate structure is neutralised by positively charged metal ions.

Classification Of silicates:

Depending upon the number of comers (0, 1, 2, 3 or 4) of SiO₄^4- tetrahedral unit shared with another tetrahedral unit through oxygen atoms, silicates are following six types:

1. Orthosllicates:

These are simple silicates which contain discrete SiO₄^{4- tetrahedrons}. Some examples are—zircon: Zr₂[SiO₄] , forsterite: Mg₂[SiO₄] , willemite: Zn₂[SiO₄] and phenacite: Be₂[SiO₄]

2. Pyroslllcates:

When two SiO₄^4- – tetrahedra share one corner (i.e., one oxygen atom), Si₂O₇^6-anion is formed. Silicates containing discrete Si₂O units are called pyrosilicates.

The common examples of phyllosilicates are:

1. Thortveitite: Sc₂[Si₂O₇] and
2. Hemimorphite: Zn₄(OH)₂[Si₂O₇]-H₂O

3. Ring or cyclic silicates:

When two O-atoms per tetrahedron are shared to form closed rings, structures with the general formula, (SiO₃)_n^2n-– are obtained. The silicates containing these anions are called cyclic silicates.

Some common examples are :

1. Wollastonite: Ca₃[Si₃O₉] (containing the cyclic ion,  Si₃O₉^6-  and
2. Beryl: Be₃Al₂[Si₆O₁₈] (containing the cycle ion, [Si₆O₁₈]^-12

4. Chain silicates:

If two oxygen atoms of each tetrahedral unit are so shared that a linear single-strand chain of the general formula, (SiO₃)₂ is formed, then the silicates containing these anions are called chain silicates.

Minerals of this type are called pyroxene and these include:

1. Enstatite: Mg₂[(SiO₃)₂],
2. Diopside: CaMg[(SiO₃)₂] and s
3. Spodumene: LiAl[(SiO₃)₂].

On the other hand, those chain silicates, containing double chain are called amphiboles. Here, two chains are attached through the O-atom. These silicates contain (Si₄O₁₁)_n ^6- ions. Minerals of asbestos are most commonly known as amphiboles.

For example: Crocidolite or blue asbestos [Na₂Fe₅(OH)₂(Si₄O₁₁)₂], amosite or brown asbestos [(Mg, Fe)(OH)₂(Si₄O₁₁)₂

Asbestos is heat and fire-resistant and thus is used as a shed for houses. Fine asbestos fibres, on entering the lungs cause asbestosis which can even result in lung cancer.

5. Sheet silicates:

Sharing of three comers i.e., three O-atoms of each tetrahedron results in the formation of an infinite two-dimensional sheet structure of the formula (Si₄O₅)_n ^2n-. Silicates containing these anions are called sheet silicates.

Some of the common examples are:

1. Kaolinite: [Al₂(OH)₄Si₂O₅] and
2. Alc: [Mg₃(OH)₂Si₄O₁₀] .

Clay also belongs to this class containing (Si₂O₆) ^2-– anions

Three-dimensional silicates:

If all the four comers i.e., all the four O-atoms of each tetrahedron are shared with other tetrahedra, a three-dimensional network structure is obtained. These have a general formula, (SiO₂)_n. Some common examples are quartz, tridymite and cristobalite

When a few silicon atoms in a three-dimensional network of SiO₂ are replaced by Al³⁺ions, the overall structure thus obtained carries a negative charge and is called aluminosilicate. Cations such as Na⁺, K⁺ or Ca²⁺ balance the negative charge. Such three-dimensional aluminosilicates are called zeolites.

A common example is natrolite:

Na₂[Al₂Si₃O₁₈]-2H₂O. Feldspars and ultramarines are two other types of three-dimensional aluminosilicates.

Many open channels of molecular levels are present in the structure of the zeolites. Depending on the shape and size of these open channels, ions or molecules of different shapes and sizes are adsorbed by the zeolites. Thus, zeolites are used as molecular sieves for separating molecules of different sizes.

Other two types of three-dimensional aluminosilicates are:

1. Feldspar example: Orthoclase, KAlSi₃O₈) and
2. Ultramarine example: Ultramarine blue, Na₈(AlSiO₄)₆S₂ )