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The p-Block Elements

The Boron Family & Its Physical and Chemical Properties

Group 14 elements (carbon family): C, Si, Ge, Sn, Pb 

Carbon is the 17th most abundant element in the earth’s crust (by mass).

In elemental state, carbon is available as coal, graphite and diamond.

In combined state, it exists as carbonates, hydrogen carbonates and carbon dioxide in air (0.03%)

Two stable, naturally occurring isotopes: 12C and 13C

14C is a radioactive isotope used for radiocarbon dating.

Silicon is the second most abundant element in the earth’s crust (27.7% by mass).

Silicon is the important component of ceramics, glass and cement.

Germanium is present only in traces.

Tin exists as cassiterite (SnO2) and lead as galena (PbS).

Germanium and silicon (in ultra pure form) are used for making transistors and semiconductor devices.

Atomic Properties of the Elements of Carbon Family

Valence shell electronic configuration is ns2 np2. 

Covalent radius increases from C to Si; after that there is a small increase from Si to Pb.

First ionisation enthalpy of group 14 members is higher than that of group 13 elements.

Electronegativity: The elements of this group are slightly more electronegative than the elements of group 13.

Physical Properties of the Elements of Carbon Family

All the elements are solids.

C is non-metal; Si and Ge are metalloids; Sn and Pb are soft metals.

Melting and boiling points of these elements are higher than those of group 13 elements.

Chemical Properties

Oxidation States:

Common oxidation states are +4 and +2.

Compounds in +4 state are generally covalent. (Since sum of the first four ionisation enthalpies is very high)

C and Si mostly show +4 oxidation state.

Ge in +4 state, forms stable compounds, and in +2 state, forms only a few compounds.

Sn forms compounds in both +2 and +4 oxidation states.

Compounds of lead in +2 state are stable and in +4 state are strong oxidising agents.

Due to the presence of d-orbitals in Si, Ge, Sn and Pb, these elements can exceed covalence more than 4. Thus, the halides of these elements undergo hydrolysis and have tendency to form complexes by accepting electron pairs from donor species.

Examples − (central atom is sp3d2)

Carbon cannot exceed its covalence by more than 4 (due to the absence of d-orbitals)

Reactivity towards Oxygen:

All members form two types of oxides, MO (monoxide) and MO2 (dioxide).

Oxides with higher oxidation states of elements (CO2, SiO2 and GeO2) are more acidic than those in lower oxidation states. SnO2 and PbO2 are amphoteric.

Among monoxides, CO is neutral; GeO is acidic; SnO and PbO are amphoteric.

Reactivity towards Water:

C, Si and Ge are not affected by water.

Sn reacts with steam to form dioxide and dihydrogen gas.

Due to the formation of a protective oxide film, lead is unaffected by water.

Reactivity towards Halogens:

Form halides of formula MX2 and MX4 (X = F, Cl, Br, I)

Most of the MX4 are covalent in nature. (Exceptions − SnF4 and PbF4 are ionic in nature)

The central metal atom in the covalent halides of the form MX4 undergoes sp3 hybridisation, and the molecule is tetrahedral in shape.

Most tetrachlorides are easily hydrolysed by water because of the presence of d-orbital in the central metal atom. d-orbital can accommodate the lone pair of electrons from the oxygen atom of a water molecule.

For example −

Question: Why does PbI4 not exist?

Answer: It does not exist because the Pb-I bond formed initially during the reaction does not release enough energy to un-pair 6s2 electrons, to have four un-paired electrons around the lead atom.

Question: Why does exist while does not?

Answer: Six large chloride ions cannot be accommodated around Si4+ due to its small size. Also, the interaction between the lone pair of chloride ion and Si4+ is not very strong. This is why exists while does not.

Elements of six groups, from 13 to 18 in the periodic table, represent the p-block elements.

Valence shell electronic configuration is ns2np1−6 .

The given table lists the important oxidation states of the p-block elements.

Table: General Electronic Configuration and Oxidation States of p-Block Elements

Group

13

14

15

16

17

18

General

electronic

configuration

ns2np1

ns2np2

ns2np3

ns2np4

ns2np5

ns2np6 (1s2 for He)

First member of the group

B

C

N

O

F

He

Group oxidation state

+3

+4

+5

+6

+7

+8

Other oxidation states

+1

+2, −4

+3, −3

+4, +2, −2

+5, +3, +1, −1

+6, +4, +2

Non-metals and metalloids exist only in the p-block of the periodic table.

Non-metallic character of the elements decreases down the group.

Non-metals have higher ionisation enthalpies and higher electronegativities than the metals.

Non-metals readily form ions.

Group 13 Elements − The Boron Family

Boron (B) is a non-metal; aluminium (Al) is a metal (but shows many chemical similarities to boron). Gallium (Ga), indium (In) and thallium (Tl) are metallic in character.

B occurs as H3BO3, Na2B4O7.10H2O(borax) and (kernite)

Two isotopic forms of boron − and

Aluminium is the most abundant metal. Bauxite and cryolite (Na3AlF6) are the important minerals of aluminium.

Atomic Properties of the Elements of Group 13

Outer electronic configuration of these elements is ns2np1.

On moving down the group, atomic radius is expected to increase. However, the atomic radius of gallium (135 pm) is less than that of aluminium (143 pm).

The order of ionisation enthalpies for these elements is. The sum of these ionisation enthalpies for each of the elements is very high.

Down the group, electronegativity first decreases from B to Al, and then increases marginally.

Physical Properties of Group 13 Elements

Boron is a non-metal and is an extremely hard, black-coloured solid. It has a high melting point due to a very strong crystalline lattice.

Other members are soft metals, with low melting points and high electrical conductivities.

Density of these elements increases down the group.

Chemical Properties of Group 13 Elements

Oxidation State and Trends in Chemical Reactivity:

For aluminium, the sum of the first three ionisation enthalpies is less than that of B. Hence, it forms Al+3 ions. In Ga, In and Tl, both +1 and +3 oxidation states are observed. The relative stability of the +1 oxidation state increases for heavier elements and follows the order- Al < Ga < In < Tl In thallium, +1 oxidation state is predominant and +3 oxidation state has a high oxidising character. In BF3, B is in +3 oxidation state and it contains only six electrons. Hence, it is an electron-deficient molecule. BF3 can accept a pair of electrons to attain a stable electronic configuration. Thus, it behaves as a Lewis acid. The sum of the first three ionisation enthalpies of boron is very high. Hence, it does not form +3 ions, and forms only covalent compounds.

AlCl3 achieves stability by forming a dimer.

Being covalent in trivalent state, most of the compounds are hydrolysed in water. Trichlorides on hydrolysis in water form tetrahedral species, (M is sp3 hybridised). AlCl3 in acidified aqueous solution forms octahedral ion, Al is in sp3d2 hybridised state.

Reactivity Towards Air:

Boron is un-reactive in crystalline form.

Aluminium metal and amorphous boron give Al2O3 and B2O3 respectively (on heating in air)

(E = Element)

B2O3 is acidic in nature; aluminium and gallium oxides are amphoteric; and the oxides of indium and thallium are basic in nature.

Reactivity towards Acids and Alkalies:

Boron does not react with acids and alkalies, even at moderate temperature.

Aluminium dissolves in mineral ac

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