GENERAL PRINCIPLES OF INORGANIC CHEMISTRY

Trends in the Periodic Table

Shells upto ( $\mathrm{n}-1$ ) are completely filled and differentiating electron (last filling electron) enters into $\mathrm{np}$ sub-orbits-elements constitute what we call p-block elements.

1. General electronic configuration is $n s^{2} n p^{1-6}$ ( $n$ varies from 2 to 7 )

Where …

x: covalent/van der Waals’ radius

$y$ : metallic character

$z$ : heat of sublimation

$p$ : electronegativity

$q$ : ionization energy ( $N>0, P>S)$

$r$ : oxidizing power

$s$ : stability of higher oxidation state within a group.

It: stability of lower oxidation state within a group.

Trends in Properties of $p$-Block Elements in the direction of arrow

The highest oxidation state $=($ group number -10$)$ Stability of this oxidation state (0.S.) decreases and that of [(0.S.) - 2] state increases as we move down the group - a case of inert pair effect.

Oxidation states

On descending the group, a lower oxidation state which is two units less than the highest oxidation state becomes more stable in group 13 to 16 . This trend is due to inert pair effect. For example, the highest oxidation state for the elements of group 13 is +3 . However, in addition to + 3 oxidation state, these elements also show +1 oxidation state which becomes more stable than +3 . Similarly, for group 14 , the group oxidation state is +4 , but +2 oxidation state becomes more stable on going down the group. For example, the last element, lead +2 oxidation state is more stable than +4 oxidation state. Similarly, thallium, +1 Oxidation State is more stable than +3 .

This trend of occurrence of oxidation state two units less than the group oxidation state is called inert pair effect and becomes more and more prominent as we move down the group.

The common oxidation states displayed by the p-block elements are given in Table 1.

Table 1. Common oxidation states of p-block elements

Metallic and non-metallic character

The p-block contains metallic and non-metallic elements. It is very interesting to note that the non-metals and metalloids exist only in the p-block of the periodic table. The non-metallic character increases along a period but decreases down a group. In fact the heaviest element in each p-block group is the most metallic in nature. Therefore, the elements with most metallic character are located mostly in the lower left portion while those with most non-metallic character are present at the top right portion of the periodic table. In between these, there are some elements which show characteristics of both metals and non-metals and are called metalloids. The common metalloids in p-block elements are B, Al, Si, Ge, As, Sb, Te, Po, At. This change from non-metallic to metallic brings significant diversity in the chemistry of these elements.

In general, non-metals have higher ionization enthalpies and higher electronegativities than metals. Therefore, in contrast to metals which readily form cations, non-metals readily form anions. The compounds formed by combination of highly reactive non-metals, with highly reactive metals are generally ionic in nature because of large differences in their electronegativities. On the other hand, compounds formed between non-metals themselves are largely covalent in character because of small differences in their electronegativities. It can be understood in terms of their oxides. The oxides of non-metals are acidic or neutral whereas oxides of metals are basic in nature. The oxides of metalloids are amphoteric. Further more, the more electropositive the metal, the more basic is it and the more electronegative the non-metal, the more acidic is its oxide. Therefore, in p-block elements, acidic character of the oxides increases or basic character decreases along a period. Similarly, the basic character of the oxides increases or acidic character decreases down the group.

Differences in behavior of first element of each group

The first member of each group of p-block differs in many respects from its succeeding members (called congeners) of their respective groups. For example, boron shows anomalous behaviour as compared to rest of the members of the 13 group elements. The main reasons for the different behaviour of the first member as compared to other members is because of

(i) small size of the atom and its ion

(ii) high electronegativity and

(iii) absence of d-orbitals in their valence shell

These factors have significant effect on the chemistry of first element as compared to other elements (specially second). For example

(a) Covalence upto four

First member of each group belongs to second period elements and have only four valence orbitals i.e., one $2 s$ and three $2 p$ orbitals available for taking part in chemical combinations. They do not have vacant $d$-orbitals in their valence shell. Therefore, they may have maximum covalence of four (using one $2 s$ and three $2 p$ orbitals). In contrast, the next members belonging to third or higher periods have vacant $\mathrm{d}$-orbitals. For example, the elements of third period of p-block with the electronic configuration $3 s^{2} 3 p^{1-6}$ has vacant $3 d$-orbitals lying between $3 p$ and $4 s$ levels of energy. Therefore, they can easily expand their octets and can show covalence above four. For example,

(i) Boron forms only $\mathrm{BF} _{4}^{-}$(coordination number four) whereas aluminium forms $\mathrm{AlF} _{6}^{3-}$ (coordination number six).

(ii) Carbon can form only tetrahalides ( $\mathrm{CX} _{4}, \mathrm{X}=\mathrm{F}, \mathrm{Cl}, \mathrm{Br}$, I ) whereas other members can form hexahalides,

$\mathrm{SF} _{6}, \mathrm{SiCl} _{6}{ }^{2-}$ etc.

(iii) Nitrogen forms only $\mathrm{NF} _{3}$ (upto octet) while phosphorus forms pentahalides, $\mathrm{PF} _{5}$, $\mathrm{PCl} _{5}$, etc.

(iv) Fluorine does not form $\mathrm{FCl} _{3}$ (more than octet) while chlorine forms $\mathrm{ClF} _{3}$ (extends octet).

(b) Reactivity

Due to availability of d-orbitals of elements of third period, they are more reactive than elements of second period which do not have d-orbitals. For example, tetrahalides of carbon are not hydrolysed by water whereas tetrahalides of other elements of group 14 are readily hydrolysed.

The hydrolysis involves the nucleophilic attack of water molecule.

(c) Tendency to form multiple bonds

Because of the combined effect of smaller size and nonavailability of $\mathrm{d}$-orbitals, the first member of each group shows tendency to form $\mathrm{p} \pi$ p $\pi$ multiple bonds either with itself (such as $C=C, C=C, N=N, 0=0$ ). or with other members of the second period of elements (such $C=0, C=N, N=0$, etc). The other members of the group do not have strong tendency to form $\pi$ - bonding. The heavier elements do form $\pi$ - bonding but they involve $d$-orbitals and form $d \pi-p \pi$ or $d \pi-d \pi$ bonding. For example, the bonds between sulphur and oxygen in oxides of sulphur $\left(\mathrm{SO} _{2}\right.$ and $\mathrm{SO} _{3}$ ) are much shorter, than might be expected for a single bond. In these molecules, in addition to normal $\pi$ bond, a $\pi$ bond is also formed by the sidewise overlap of a filled $2 p-$ orbital of oxygen with a vacant $3 d-$ orbital on the sulphur). This is called $p \pi-d \pi$ bond and results in bringing the two atoms closer and thus accounts for shorter bond length of $S-0$ bond.

Because the $d$-orbitals are of higher energy than $p$-orbitals, they contribute less to the overall Stablity of molecules than does the $p \pi-p \pi$ bonding of second row elements. However, the coordination number in species of heavier elements may be higher in those of first element in the same oxidation state. For example, both nitrogen and phosphorus form ions in +5 oxidation state as $\mathrm{NO} _{3}^{-}$(three coordination with bonding using one $\mathrm{p}$ orbital of $\mathrm{N}$ ) $\mathrm{PO} _{4}{ } _{4}^{3-}$ (having four coordination using s, p and d orbitals contributing to the $\pi$

-bonding).

The first member of 13 group (boron) shows diagonal relationship with silicon (of group 14).

Group 13 and Group 14

Group 13

(elements) B, AI, Ga, In and TI,

The elements belonging to groups 13 to 18 belong to $\mathrm{p}$-block and have the general configuration $\mathrm{ns}^{2}$ $n p^{1-6}$.

B and AI have noble gas core, Ga and In have noble gas plus $10 \mathrm{~d}$-electrons and $\mathrm{TI}$ has noble gas plus 14 f-electrons plus $10 \mathrm{~d}$-electrons cores.

General Trends in Properties

1. General electronic configuration

$n s^{2} n p^{3}$

2.Atomic radius

$\mathrm{B}<\mathrm{Ga}<\mathrm{Al}<\mathrm{In}<\mathrm{T} \mid$

‘Ga’ has smaller atomic size due to the poor shielding effect of the inner 3d electrons.

3.Ionization enthalpy (I.E.)

$\mathrm{B}>\mathrm{T}|>\mathrm{Ga}>| \mathrm{n}>\mathrm{Al}$

This is due to the poor shielding effect of the inner electrons.

The order of ionization enthalpies as expected is $\Delta _{i} H _{1}<\Delta _{i} H _{2}<\Delta _{i} H _{3}$. The sum of these ionization enthalpies for each of the elements is very high.

4.Electro negativity

$\mathrm{B}>\mathrm{TI}>\mathrm{In}>\mathrm{Ga}>\mathrm{Al}$

This is because of the discrepencies of their atomic sizes.

Baron is non metallic in nature. It is extremely hard and black coloured solid.

Density of the elements increases down the group from B to TI.

5. Melting point

$\mathrm{B}>\mathrm{Al}>\mathrm{T}|>| \mathrm{n}>\mathrm{Ga}$

6. Oxidation state

The general oxidation state of the $13^{\text {th }}$ group of elements is +3 , but due to inert pair effect TI shows " +1 " oxidation state. So, thallous compounds are more stable than thallic compounds.

The relative stability of +1 oxidation state progressively increases for heavier elements $\mathrm{Al}<\mathrm{Ga}<\ln <$ TI. The compounds in +1 oxidation state as expected from energy consideration are more ionic than those of +3 oxidation state.

These elements in their trivalent state for electron deficient compounds which act as lewis acid for example $\mathrm{BCl} _{3}$ accepts lone pair easily from ammonia to form $\mathrm{BCl} _{3}$. $\mathrm{NH} _{3}$ while $\mathrm{AlCl} _{3}$ achieves stability by forming a dimer.

In trivalent state most of the compounds being covalent are hydrolysed in water.

Reactivity towards air

$$ 2 \mathrm{E}(\mathrm{s})+3 \mathrm{O} _{2}(\mathrm{~g}) \xrightarrow{\Delta} 2 \mathrm{E} _{2} \mathrm{O} _{3}(\mathrm{~s}) $$

The nature of these oxides vary down the group. $\mathrm{B} _{2} \mathrm{O} _{3}$ is acidic, $\mathrm{Al}$ and $\mathrm{Ga}$ oxides are amphoteric and those of In and $\mathrm{TI}$ are basic in their properties.

Reactivity towards acids and alkalies

B does not react with acids and alkalies. Al dissolves in mineral acids and aqueous alkalies and thus shows amphoteric character

$2 \mathrm{Al}(\mathrm{s})+6 \mathrm{HCl}(\mathrm{q}) \longrightarrow 2 \mathrm{Al}^{3+}(\mathrm{aq})+6 \mathrm{Cl}^{-}(\mathrm{aq})+3 \mathrm{H} _{2}(\mathrm{~g})$

$2 \mathrm{Al}(\mathrm{s})+2 \mathrm{NaOH}(\mathrm{aq})+6 \mathrm{H} _{2} \mathrm{O}(\mathrm{I}) \longrightarrow 2 \mathrm{Na}^{+}\left[\mathrm{Al}(\mathrm{OH}) _{4}\right]^{-}+3 \mathrm{H} _{2}(\mathrm{~g})$

Reactivity towards Halogens

These elements react with halogens to form trihalides (except $\mathrm{TII} _{3}$ )

$2 \mathrm{E}(\mathrm{s})+3 \mathrm{X} _{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{E} _{3}(\mathrm{~s})(\mathrm{X}=\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I})$

Important trends and anomalous properties of boron

The trichlorides, bromides and lodides of all these elements being covalent in nature are hydrolysed in water.

The monomeric trihalides, being electron deficient are strong lewis acids.

Due to absence of $d$ orbitals, the maximum covalence of $B$ is 4 . Since the $d$ orbitals are available with Al and other elements, the maximum covalence can be more than 4 .

Boron

Minerals of boron

1. Borax $\left(\mathrm{Na} _{2} \mathrm{~B} _{4} \mathrm{O} _{7} \cdot 10 \mathrm{H} _{2} \mathrm{O}\right)$

2. Boric acid $\left(\mathrm{H} _{3} \mathrm{BO} _{3}\right)$

3. Colemanite $\left(\mathrm{Ca} _{2} \mathrm{~B} _{6} \mathrm{O} _{11} \cdot 5 \mathrm{H} _{2} \mathrm{O}\right)$

4. Kernite $\left(\mathrm{Na} _{2} \mathrm{~B} _{4} \mathrm{O} _{7} \cdot 2 \mathrm{H} _{2} \mathrm{O}\right)$

Preparation of Boron

1. By thermal decomposition of boron hydrides

Boron can be prepared by the thermal decomposition of its hydrides such as diborane $\left(\mathrm{B} _{2} \mathrm{H} _{6}\right)$

$\mathrm{B} _{2} \mathrm{H} _{6} \xrightarrow{\text { Heat }} \mathrm{B}+3 \mathrm{H} _{2}$

2. Conversion of borax into boron trioxide

Borax is treated with hot and concentrated hydrochloric acid to convert it first into boric acid.

$\mathrm{Na} _{2} \mathrm{~B} _{4} \mathrm{O} _{7}+2 \mathrm{HCl} \longrightarrow 2 \mathrm{NaCl}+\mathrm{H} _{2} \mathrm{~B} _{4} \mathrm{O} _{7}$

$\mathrm{H} _{2} \mathrm{~B} _{4} \mathrm{O} _{7}+5 \mathrm{H} _{2} \mathrm{O} \longrightarrow 4 \mathrm{H} _{3} \mathrm{BO} _{3}$

$\mathrm{B} _{2} \mathrm{O} _{3}+3 \mathrm{Mg} \xrightarrow{\Delta} 2 \mathrm{~B}+3 \mathrm{MgO}$

1. General properties of Boron

i) $\quad 2 \mathrm{~B}+3 \mathrm{O} _{2} \longrightarrow \mathrm{B} _{2} \mathrm{O} _{3}$ (with air)

ii) $\mathrm{B}+3 \mathrm{HNO} _{3} \xrightarrow{\Delta} \mathrm{H} _{3} \mathrm{BO} _{3}+3 \mathrm{NO} _{2}$

iii) $2 \mathrm{~B}+6 \mathrm{KOH} \longrightarrow 2 \mathrm{~K} _{3} \mathrm{BO} _{3}+3 \mathrm{H} _{2}$

2. Hydrides of Boron

Diborane $\mathrm{B} _{2} \mathrm{H} _{6}$

Preparation of $\mathrm{B} _{2} \mathrm{H} _{6}$

i) $2 \mathrm{NaBH} _{4}+\mathrm{I} _{2} \longrightarrow \mathrm{B} _{2} \mathrm{H} _{6}+2 \mathrm{NaI}+\mathrm{H} _{2}$

ii) $4 \mathrm{BCl} _{3}+3 \mathrm{LiAlH} _{4} \longrightarrow 2 \mathrm{~B} _{2} \mathrm{H} _{6}+3 \mathrm{AlCl} _{3}+3 \mathrm{LiCl}$

Properties of diborane

i) It burns in oxygen and air to form $\mathrm{B} _{2} \mathrm{O} _{3}$ with the evolution of heat.

$\mathrm{B} _{2} \mathrm{H} _{6}+3 \mathrm{O} _{2} \longrightarrow \mathrm{B} _{2} \mathrm{O} _{3}+3 \mathrm{H} _{2} \mathrm{O} ; \Delta \mathrm{H}=-2008 \mathrm{~kJ} \mathrm{~mol}^{-1}$

This is why diborane is used as a rocket fuel.

ii) It hydrolyses in water to form boric acid.

$\mathrm{B} _{2} \mathrm{H} _{6}+6 \mathrm{H} _{2} \mathrm{O} \longrightarrow \underset{\text { Boric acid }}{2 \mathrm{H} _{3} \mathrm{BO} _{3}+6 \mathrm{H} _{2}}$

iii) Diborane reacts with ammonia at $450 \mathrm{~K}$ to form borazine.

$$ \begin{aligned} 3 \mathrm{~B} _{2} \mathrm{H} _{6}+6 \mathrm{NH} _{3} \xrightarrow{450 \mathrm{~K}} \underset{3}{ } & 3 \mathrm{~B} _{3} \mathrm{~N} _{3} \mathrm{H} _{6} \\ & \text { Borazine (inorganic benzene) } \end{aligned} $$

Borazine has the following structure which is similar to that of benzene. Hence, borazine is also referred to as inorganic benzene. It is isoelectronic to benzene.

Structure of diborane

$\mathrm{B} _{2} \mathrm{H} _{6}$ is a non-polar molecule in which each $\mathrm{B}$ atom is $\mathrm{sp}^{3}$ hybridized. $\mathrm{B} _{2} \mathrm{H} _{6}$ has four normal $\mathrm{B}$ - $\mathrm{H}$ covalent bonds, i.e., two-centre two-electron ( $2 \mathrm{c}-2 \mathrm{e}$ ) bonds which lie in the same plane and two bridge bonds, B….H…..B, i.e., three-centre two electron bonds ( $3 \mathrm{c}-2 \mathrm{e}$ ) or banana bonds which lie above and below the plane of the four B-H bonds.

Boron Halides

$\mathrm{BI} _{3}>\mathrm{BBr} _{3}>\mathrm{BCl} _{3}>\mathrm{BF} _{3} \quad$ Acidic character

This is due to $p \pi-p \pi$ back bonding which decreases the electron deficiency of the $B$ atom. Since this tendency is maximum for $\mathrm{F}$ to $\mathrm{I}$, therefore, $\mathrm{BI} _{3}$ is the strongest and $\mathrm{BF} _{3}$ is the weakest Lewis acid.

Although both $\mathrm{B}$ and $\mathrm{Al}$ trihalides act as Lewis acids but only aluminium trihalides (i.e., $\mathrm{Al} _{2} \mathrm{Cl} _{6}$ ) exist as dimers. This is due to the reason that boron atom is so small that it cannot accommodate four large sized halogen (except F) atoms around it.

Boron also forms a series of hydridoborates the most important one is the tetrahedral [ $\left.\mathrm{BH} _{4}\right]$ ion. Tetra hydridoborates of several metals are known Li and Na tetrahydridoborates also known as borohydrides are prepared by the reaction of metal hydrides with $\mathrm{B} _{2} \mathrm{H} _{6}$ in diethyl ether.

$\qquad 2 \mathrm{MH}+\mathrm{B} _{2} \mathrm{H} _{6} \longrightarrow 2 \mathrm{M}^{+}\left[\mathrm{Bh} _{4}\right]^{-}(\mathrm{M}=\mathrm{Li}$ or $\mathrm{Na})$

Both $\mathrm{LBH} _{4}$ and $\mathrm{NaBH} _{4}$ are used as reducing agents in organic synthesis.

Boric Acid

Preparation

1. From borax

$\qquad \mathrm{Na} _{2} \mathrm{~B} _{4} \mathrm{O} _{7}+2 \mathrm{HCl}+5 \mathrm{H} _{2} \mathrm{O} \longrightarrow 2 \mathrm{NaCl}+4 \mathrm{H} _{3} \mathrm{BO} _{3}$

2. From colemanite

$\qquad \mathrm{Ca} _{2} \mathrm{~B} _{6} \mathrm{O} _{11}+2 \mathrm{SO} _{2}+9 \mathrm{H} _{2} \mathrm{O} \longrightarrow 2 \mathrm{CaSO} _{3}+6 \mathrm{H} _{3} \mathrm{BO} _{3}$

• Boric acid $\left(\mathrm{H} _{3} \mathrm{BO} _{3}\right)$ contains triangular $\mathrm{BO} _{3}{ }^{3}$ ions in which boron is $\mathrm{sp}^{2}-$ hybridized. In solid state, $\mathrm{B}(\mathrm{OH}) _{3}$ molecules are $\mathrm{H}$-bonded to form a two-dimensional sheet. It is a weak monobasic acid. It does not act as a protonic acid but acts as a Lewis acid by accepting a pair of electrons from $\mathrm{OH}^{-}$ion of water thereby releasing a proton.

• Reaction with ethyl alcohol. Orthoboric acid reacts with ethyl alcohol in presence of conc. $\mathrm{H} _{2} \mathrm{SO} _{4}$ to form triethylborate.

$$ \mathrm{B}(\mathrm{OH}) _{3} \quad+3 \mathrm{C} _{2} \mathrm{H} _{5} \mathrm{OH} \xrightarrow[\Delta]{\text { Conc. } \mathrm{H} _{2} \mathrm{SO} _{4}} \mathrm{~B}\left(\mathrm{OC} _{2} \mathrm{H} _{5}\right) _{3}+3 \mathrm{H} _{2} \mathrm{O} $$

$\qquad \qquad \qquad \qquad $ Orthoboric acid $ \qquad \qquad $ Ethyl alcohol $\qquad \qquad $ Triethylborate

The vapours of triethylborate when ignited burn with a green-edged flame. This forms the basis for detecting borates and boric acid in qualitative analysis.

Heating

Borax $\left(\mathrm{Na} _{2} \mathrm{~B} _{4} \mathrm{O} _{7} \cdot 10 \mathrm{H} _{2} \mathrm{O}\right)$

White crystalline solid

Cyclic structure represented as $\mathrm{Na} _{2}\left[\mathrm{~B} _{4} \mathrm{O} _{5}(\mathrm{OH}) _{4}\right] \cdot 8 \mathrm{H} _{2} \mathrm{O}$

Aqueous solution is slightly alkaline due to hydrolysis

$\mathrm{Na} _{2} \mathrm{~B} _{4} \mathrm{O} _{7}+2 \mathrm{H} _{2} \mathrm{O} \rightleftharpoons 2 \mathrm{NaOH}+\mathrm{H} _{2} \mathrm{~B} _{4} \mathrm{O} _{7}$

On heating borax, itloses water of crystallization and swells into white opaque mass.

Borax bead test

On heating, borax swells up forming a glassy mass of mixture of $\mathrm{NaBO} _{2}+\mathrm{B} _{2} \mathrm{O} _{3}$. The glassy mass on heating with many transition metal salts form coloured metaborate.

$$ \mathrm{CuSO} _{4}+\mathrm{B} _{2} \mathrm{O} _{3} \xrightarrow{\Delta} \mathrm{Cu}\left(\mathrm{BO} _{2}\right) _{2}+\mathrm{SO} _{3} \uparrow $$

Group-14-Carbon Family

Some physical constants of group 14 elements.

1. Atomic Radius

$\mathrm{C}<\mathrm{Si}<\mathrm{Ge}<\mathrm{Sn}<\mathrm{Pb}$

There is a small increase in radius due to presence of completely filled $d$ and $f$ orbitals in heavier members.

2. Ionization energy

$\mathrm{C}>\mathrm{Ge}>\mathrm{Si}>\mathrm{Pb}>\mathrm{Sn}$ Irregular trend is due to inert pair effect.

The influence of inner core electrons is present in group 14 also.

Small decrease in $\Delta _{i} \mathrm{H}$ from Si to Ge to $\mathrm{Sn}$ and slight increase in $\Delta _{i} \mathrm{H}$ from $\mathrm{Sn}$ to $\mathrm{Pb}$ is the consequence of poor shielding effect of intervening $d$ and $f$ orbitals and increase in size of the atom.

3. Electro negativity

$\mathrm{C}>\mathrm{Si} \approx \mathrm{Ge}>\mathrm{Pb}>\mathrm{Sn}$

Due to small size the elements, this group elements are slightly more electronegative than group 13 elements. The electronegativity values for elements from $\mathrm{Si}$ to $\mathrm{Pb}$ are almost same.

4. Melting Point

$\mathrm{C}>\mathrm{Si}>\mathrm{Ge}>\mathrm{Pb}>\mathrm{Sn}$

Oxidation state

The common oxidation states exhibited by these elements are +4 and +2 .

The sum of the first four ionization enthalpies is very high. In heavier members the tendency to show +2 oxidation state increases in the sequence $\mathrm{Ge}<\mathrm{Sn}<\mathrm{Pb}$. It is due to inability of $\mathrm{ns}^{2}$ electrons of valence shell to participate in bonding.

Reactivity towards oxygen

All members form oxides of two types monoxide and dioxide ( $\mathrm{MO}$ and $\mathrm{Mo} _{2}$ ). Oxides in higher oxidation states of elements are generally more acidic than those in lower oxidation states. $\mathrm{CO} _{2}, \mathrm{SiO} _{2}$ and $\mathrm{GeO} _{2}$ are acidic, $\mathrm{SnO} _{2}$ and $\mathrm{PbO} _{2}$ are amphoteric.

Among monoxides $\mathrm{CO}$ is neutral, $\mathrm{GeO}$ is distinctly acidic whereas $\mathrm{SnO}$ and $\mathrm{PbO}$ are amphoteric.

Reactivity towards water

$\mathrm{C} _{1} \mathrm{Si}, \mathrm{Ge}$ are not affected by water. Tin decomposes steam to form dioxide and dihydrogen gas.

$$ \mathrm{Sn}+2 \mathrm{H} _{2} \mathrm{O} \xrightarrow{\Delta} \mathrm{SnO} _{2}+2 \mathrm{H} _{2} $$

Lead is not affected by water.

Reactivity towards halogen

They form halides of formula $M X _{2} \& M X _{4} \cdot M X _{4}$ are covalent. Stability of dihalides increases down the group. Except $\mathrm{CCl} _{4}$ other tetrachloride’s are easily hydrolysed by water because the central atom can accommodate the lone pair of electrons from oxygen atom of water molecule in d orbital.

Allotropes of Carbon

1. Diamond

Three dimensional network of carbon atoms joined through strong covalent bonds.

Each $\mathrm{C}$ is $\mathrm{sp}^{2}$ hybridised and linked tetrahedrally to four neighbouring $\mathrm{C}$ atoms.

Hardest substance known

Very high melting point (3843 K)

Bad conductor of electricity since all valence electrons are involved in bond formation.

Transparent and has high refractive index (2.45)

2. Graphite

Each carbon atom is $\mathrm{sp}^{2}$ hybridised and covalently attached to three neighbouring carbon atoms.

Planar hexagonal rings are formed.

They are held together by weak vanderwaal’s forces.

Sheets can slide over each other

Soft and has lubricating properties.

3. Silica (Silicondioxide)

(a) Structure

Covalentcompound

Si is tetrahedrally surrounded by 4 oxygen atoms

Each oxygen is shared by two silicon atoms

Covalent bonds between $\mathrm{Si}$ and 0 atoms are very strong.

(b) Physical properties

Insoluble in water

Does not react with acids

High melting point

(c) Chemical properties

$\mathrm{SiO} _{2}+2 \mathrm{H} _{2} \mathrm{~F} _{2} \longrightarrow \mathrm{SiF} _{4}+2 \mathrm{H} _{2} \mathrm{O}$

$\mathrm{SiO} _{2}+4 \mathrm{NaOH} \longrightarrow \mathrm{Na} _{4} \mathrm{SiO} _{4}+2 \mathrm{H} _{2} \mathrm{O}$

$\mathrm{SiO} _{2}+\mathrm{Na} _{2} \mathrm{CO} _{3} \longrightarrow \mathrm{Na} _{2} \mathrm{SiO} _{3}+\mathrm{CO} _{2}$

$\mathrm{SiO} _{2}+4 \mathrm{Na} \xrightarrow{\Delta} \mathrm{Si}+2 \mathrm{Na} _{2} \mathrm{O}$

$\mathrm{SiO} _{2}+2 \mathrm{~F} _{2} \longrightarrow \mathrm{SiF} _{4}+\mathrm{O} _{2}$

$\mathrm{SiO} _{2}+\mathrm{Al} _{2} \mathrm{O} _{3} \longrightarrow \mathrm{Al} _{2}\left(\mathrm{SiO} _{3}\right) _{3}$

$\mathrm{SiO}_2+3 \mathrm{C} \longrightarrow \mathrm{SiC}+2 \mathrm{CO}$

4. Silicates

Contain $\mathrm{SiO} _{4}{ }^{4-}$ tetrahedral units.

Classification of silicates

Based on the way $\mathrm{SiO} _{4}{ }^{4-}$ tetrahedral units are linked.

(a) Orthosilicates

Contain single discrete unit of $\mathrm{SiO} _{4}^{4-}$ unit eg. Zircon $\mathrm{ZrSiO} _{4}$

(b) Pyrosilicates

Contain two units of $\mathrm{SiO} _{4}^{4-}$ joined along a corner containing oxygen atom.

Pyrosilicate ion is $\mathrm{Si} _{2} \mathrm{O} _{7}{ }^{6-}$

(c) Cyclic structure - Cyclic or ring silicates have general formula $\left(\mathrm{SiO} _{3}{ } _{3}{ }^{-2}\right) _{0}$ or $\left.\left(\mathrm{SiO} _{3}\right) _{n}\right) _{n}^{2 n-}$. These are formed when two oxygen atoms of each $\mathrm{SiO} _{4}^{4}$ tetrahedron are shared with others. Structures and examples of cyclic silicates containing $\mathrm{Si} _{9}{ } _{9}{ }^{-}$and $\mathrm{Si} _{6} \mathrm{O} _{18}{ }^{12}$ ions are given below.

(d) Chain silicates

If two oxygen atoms per tetrahedron are shared such that a linear single strand chain of the general formula $\left(\mathrm{SiO} _{3}{ }^{2}\right) _{n}$ or $\left(\mathrm{SiO} _{3}\right) _{n}{ }^{2-}$ is formed, then the silicates containing these anions are called chain silicates.

(e) Sheet silicates

Three oxygen atoms of a tetrahedral $\mathrm{SiO} _{4}^{4}$ are shared.

(f) Three dimensional sheet silicates

These silicates involve all four oxygen atoms in sharing with adjacent $\mathrm{SiO} _{4}{ } _{4}^{4}$ tetrahedral. Since, all the oxygen atoms are shared, the silicates are neutral. The common examples are quartz, tridymite and cristobalite (forms of silica).

5. Silicones

Silicones are synthetic organosilicone polymers containing repeated $\mathrm{R} _{2} \mathrm{SiO}$ units held by $\mathrm{Si}-0-\mathrm{Si}$ bonds. These are prepared by hydrolysis of alkyl or aryl substituted chlorosilanes and their subsequent polymerization. For example, hydrolysis of dichlorodimethylsilane followed by polymerization yields straight chain silicone polymer. The chain length of these polymers, can, however, be controlled by adding $\left(\mathrm{CH} _{3}\right) _{3} \mathrm{SiCl}$ which blocks the ends. Silicone polymers are stable towards heat and are chemically inert and are good insulators. Therefore, they are used for making water proof papers, wool, textiles, wood, etc. They are also used as lubricants and in surgical and cosmetic implants.

Linear Chain silicones are obtained by hydrolysis of $\mathrm{R} _{2} \mathrm{SiCl} _{2}$.

If $\mathrm{R} _{2} \mathrm{SiCl} _{2}$ is hydrolysed in limited supply of water, cyclic silicones are obtained.

$\mathrm{R} _{2} \mathrm{SiCl} _{2}+\underset{\text { limited amount }}{\mathrm{H} _{2} \mathrm{O} \longrightarrow}$

Some $\mathrm{R} _{3} \mathrm{SiCl}$ is added to control the molar mass of linear chain silicones. If it is hydrolysed alone, dimeric siloxane is formed :

ii) Cross-linked silicones

Hydrolysis of $\mathrm{RSiCl} _{3}$ gives cross linked silicones.

6. Zeolites

The three-dimensional silicates in which some of the $\mathrm{Si}$ atoms are replaced by $\mathrm{Al}^{3+}$ ions and the negative charge is balanced by cations such as $\mathrm{Na}^{+}, \mathrm{K}^{+}, \mathrm{Ca}^{2+}$, etc. are called feldspar ( $\mathrm{KAlSi} _{3} \mathrm{O} _{8}$ ) and zeolites ( $\left.\mathrm{NaAlSi} _{2} \mathrm{O} _{6} \cdot \mathrm{H} _{2} \mathrm{O}\right)$.

Zeolites are widely used as catalysts in petrochemical industries for cracking of hydrocarbons and

isomerization. Another zeolite called ZSM-5 is used to convert alcohols directly into gasoline. Hydrated zeolites called permutitare used as ion exchangers for softening of hard water.

Solved Examples

1. Which one of the following is the correct statement?

1. Boric acid is a protonic acid.

2. Beryllium exhibits coordination number of six.

3. Chlorides of both beryllium and aluminium have bridged chloride structure in solid phase.

4. $\mathrm{B} _{2} \mathrm{H} _{6} \cdot 2 \mathrm{NH} _{3}$ is known as inorganic benzene.

2. A metal, $M$ forms chlorides in +2 and +4 oxidation states. Which of the following statements about these chlorides is correct?

1. $\mathrm{MCl} _{2}$ is more volatile than $\mathrm{MCl} _{4}$.

2. $\mathrm{MCl} _{2}$ is more soluble in anhydrous ethanol than $\mathrm{MCl} _{4}$.

3. $\mathrm{MCl} _{2}$ is more ionic than $\mathrm{MCl} _{4}$.

4. $\mathrm{MCl} _{2}$ is more easily hydrolysed than $\mathrm{MCl} _{4}$.

3. In silicon dioxide

1. Each silicon atom is surrounded by four oxygen atoms and each oxygen atom is bonded to two silicon atoms.

2. Each silicon atom is surrounded by two oxygen atoms and each oxygen atom is bonded to two silicon atoms.

3. Silicon atom is bonded to two oxygen atoms.

4. There are double bonds between silicon and oxygen atoms.

Practice Questions

1. $\mathrm{H} _{3} \mathrm{BO} _{3}$ is

(a) monobasic acid and weak lewis acid

(b) monobasic and weak bronsted acid

(c) monobasic and strong Lewis acid

(d) tribasic and weak Bronsted acid

2. (Me) $) _{2} \mathrm{SiCl} _{2}$ on hydrolysis will produce

(a) $(\mathrm{Me}) _{2} \mathrm{Si}(\mathrm{OH}) _{2}$

(b) $(\mathrm{Me}) _{2} \mathrm{Si}=0$

(c) $\left[-0-(\mathrm{Me}) _{2} \mathrm{Si}-0-\right] _{n}$

(d) $\mathrm{Me} _{2} \mathrm{SiCl}(\mathrm{OH})$

3. Name of the structure of silicates in which three oxygen atoms of $\left[\mathrm{SiO} _{4}\right]^{4}$ are shared is

(a) pyrosilicate

(b) sheet silicate

(c) linear chain silicate

(d) three dimensional silicate

4. $\mathrm{B}(\mathrm{OH}) _{3}+\mathrm{NaOH} \rightleftharpoons \mathrm{NaBO} _{2}+\mathrm{Na}\left[\mathrm{B}(\mathrm{OH}) _{4}\right]+\mathrm{H} _{2} \mathrm{O}$ How can this reaction be made to proceed in forward direction?

(a) Addition of cis 1,2 diol

(b) Addition of borax

(c) Addition of trans 1,2 diol

(d) Addition of $\mathrm{Na} _{2} \mathrm{HPO} _{4}$

5. The product/s formed when diborane is hydrolysed is/are

(a) $\mathrm{B} _{2} \mathrm{O} _{3}$ and $\mathrm{H} _{3} \mathrm{BO} _{3}$

(b) $\mathrm{B} _{2} \mathrm{O} _{3}$ only

(c) $\mathrm{H} _{3} \mathrm{BO} _{3}$ and $\mathrm{H} _{2}$

(d) $\mathrm{H} _{3} \mathrm{BO} _{3}$ only

6. Reaction of diborane with ammonia gives initially

(a) $\mathrm{B} _{2} \mathrm{H} _{6} \mathrm{NH} _{3}$

(b) Borazine

(c) $\mathrm{B} _{2} \mathrm{H} _{6} \cdot 3 \mathrm{NH} _{3}$

(d) $\left[\mathrm{BH} _{2}\left(\mathrm{NH} _{3}\right) _{2}\right]^{+}\left[\mathrm{BH} _{4}\right]^{-}$

7. The structure of diborane $\left(\mathrm{B} _{2} \mathrm{H} _{6}\right)$ contains

(a) four 2c-2e bonds and two $3 \mathrm{c}-2 \mathrm{e}$ bonds

(b) two 2c-2e bonds and four $3 \mathrm{c}-2 \mathrm{e}$ bonds

(c) two 2c-2e bonds and two $3 \mathrm{c}-3$ e bonds

(d) four 2c-2e bonds and four 3c-2e bonds

8. In borax, the number of $\mathrm{B}-\mathrm{O}-\mathrm{B}$ links and $\mathrm{B}-\mathrm{OH}$ bonds present are, respectively

(a) five and four

(b) four and five

(c) three and four

(d) five and five

9. Which one of the following has highest Lewis acid strength?

(a) $\mathrm{BI} _{3}$

(b) $\mathrm{BBr} _{3}$

(c) $\mathrm{BF} _{3}$

(d) $\mathrm{BCl} _{3}$

10. Which one of the following is the correct statement? (a) $\mathrm{B} _{2} \mathrm{H} _{6} .2 \mathrm{NH} _{3}$ is known as ‘inorganic benzene’

(b) Boric acid is a protonic acid.

(c) Beryllium exhibits coordination number of six.

(d) Chlorides of both beryllium and aluminium have bridged chlorine structure in solid phase.