UNIT 11 SOME p-BLOCK ELEMENTS
Syllabus
Group 13 Elements
· General introduction
· Electronic configuration
· Occurrence
· Variation in properties
· Oxidation states
· Trends in chemical reactivity
· Anomalous behaviour of first element of group
· Boron
· Physical and chemical properties
· Some important compounds : Borax, boric acids , boron
hydrides
· Aluminium : uses , reactions with acids and alkalies
Group 14 Elements
· General introduction
· Electronic configuration
· Occurrence
· Variation in properties
· Oxidation states
· Trends in chemical reactivity
· Anomalous behaviour of first element
· Carbon-catenation , Allotropic forms
· Physical and chemical properties
· Uses of some important compounds , Oxides
· Important compounds of silicon and few uses
· Silicon tetrachloride , Silicones
· Silicates and zeolites
In the p-block of the periodic table, the electrons are filled in the valence p-shell of the atoms of elements whereas the valence s-sub-shell is already filled. Thus, the general electronic configuration of the elements is ns2p1-6. Since the p-sub-shell can have a maximum of six electrons, there are six groups in the p-block ( 13 to 18). These are named as :
Boron family (Group 13) : ns2p1
Carbon family ( Group 14) : ns2p2
Nitrogen family (Group 15) : ns2p3
Oxygen family ( Group 16) : ns2p4
Halogen family (Group 17) : ns2p5
Noble gases ( Group 18) : ns2p6
The inner core of the electronic configuration may, however, differ. The difference in inner core of elements greatly influences their physical properties (such as atomic and ionic radii , ionization enthalpy , etc) as well as chemical properties. Consequently a lot of variation in properties of elements in a group of p-block is observed. The maximum oxidation state shown by a p-block element is equal to the total number of valence electrons (i.e., the sum of s- and p-electrons). The number of possible oxidation states increases towards the right of the periodic table. In addition to this group oxidation state, p-block elements may show other oxidation states which normally , but not necessarily , differ from total number of valence electrons by unit of two. The important oxidation states exhibited by p-block elements are shown in the following Table.
Group | 13 | 14 | 15 | 16 | 17 | 18 |
General electronic Configuration | ns2p1 | ns2p2 | ns2p3 | ns2p4 | ns2p5 | ns2p6 |
First member of 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 |
In boron, carbon and nitrogen families the group oxidation state is the most stable for the lighter elements in the group. However, the oxidation state two units less than the group oxidation state becomes progressively more stable for heavier elements in each group. The occurrence of oxidation states two units less than the group oxidation states are some times attributed to the ‘inert pair effect’ . The non-metals and metalloids exist only in p-block of the periodic table. There is a gradual increase in non-metallic character from left to right in the p-block. Down each group, the metallic character tends to increase. This trend is attributed to the change in the main periodic properties of the elements i.e., atomic radii , ionic radii, ionisation enthalpies, electron affinities and electronegativities. In this unit, we shall focus our attention to the chemistry of the first elements in each group (Group 13 to 16) of the p-block. These are all typical non-metals and are also quite reactive in nature.
The first member of each of the groups 13-17 of the p-block elements differs in many respects from the other members of their respective groups. These differences are quite significant in groups 13-16. However , in group 17 , the halogens show much more similarity among themselves in their chemical behaviour than the elements of other groups. Besides its small size and high electronegativity, the first member of each group has only four valence orbitals ( 2s and 2p) available for bonding and seldom forms compounds in which the co-ordination number exceeds four. In other words, there is no expansion of the valence shell to accommodate more than four pairs of electrons. Furthermore, the first member of the group displays a greater ability to form pp-pp multiple bonds to itself ( e.g. C=C, Cº C, Nº N) and to other second row elements ( e.g. C= O, C=N, N=O) compared to subsequent members of the same group. This type of p-bonding is not particularly strong for the elements of the third and subsequent rows of the Periodic Table.
GROUP 13 ELEMENTS : THE BORON FAMILY
Group 13 of the periodic table contains five elements namely, boron, aluminium, gallium, indium and thallium. Boron is a fairly rare element , mainly occurs as orthoboric acid (H3BO3), borax Na2B4O7 10 H2O and kernite Na2B4O7 4 H2O . Boron occurs in two isotopic forms 10B (19%) and 11B (81%). The abundance of boron in the earth’s crust is less than 0.0001% by mass. Aluminium is the most abundant metal and the third most abundant element in earth’s crust (8.3% by mass) after oxygen (45.5%) and silicon (27.7%). Bauxite , Al2O3 2 H2O and cryolite Na3AlF6 are important minerals of aluminium. Gallium, indium and thallium are less abundant elements in nature.
The atomic , physical and chemical properties of these elements are discussed
Atomic and Physical Properties of Group 13 Elements
Electronic configurations
The outer electronic configurations of these elements is ns2np1. An examination of electronic configuration suggests that while boron and aluminium have noble gas core, gallium and indium have noble gas plus 10 d-electrons and thallium has noble gas plus 14 f-electrons plus 10 d-electrons cores.
Atomic Radii
On moving down the group , for each successive member one extra shell of electrons is added and therefore, atomic radius is expected to increase. However a deviation can be seen. Atomic radius of gallium is less than that of aluminium. This can be understood from the variation in the inner core of the electronic configuration. The presence of additional 10 d-electrons offer only poor shielding effect for outer electrons from the increased nuclear charge in gallium. Consequently, the atomic radius of gallium (135 pm) is less than that of aluminium (143 pm)
Ionization Enthalpy
The ionization enthalpy values are expected from the general trends do not decrease smoothly down the group. The decrease from boron to aluminium is associated with increase in size. The observed discontinuity in the ionization enthalpy values between aluminium and gallium and between indium and thalium are due to inability of d- and f-electrons , which have low screening effect, to compensate the increase in nuclear charge.
The order of ionization enthalpies, is DiH1 < DiH2 < DiH3. The sum of the first three ionization enthalpies for each of the elements is very high.
Electronegativity
Down the group , electronegativity first decreases from B to Al and then increases marginally. This is because of the discrepancies in atomic size of the elements.
Physical Properties
Boron is non-metallic in nature. It is extremely hard and black coloured solid. It exists many allotropic forms. Due to very strong crystalline lattice has unusually high melting point. Rest of the members are soft metals with low melting points and high electrical conductivity. Gallium has unusually low melting points (303 K), could exist , could exist in liquid state during summer. Its high boiling point (2676 K) makes it useful material for measuring high temperatures. Density of elements increases down the group from boron to thallium.
Chemical Properties
Oxidation state and trends in chemical reactivity
Due to small size of boron, the sum of its first three ionization enthalpies is very high. This prevents it from +3 ions and forces it to form only covalent compounds. But as we move from B to Al , the sum of the first three ionization enthalpies of Al considerably decreases and is therefore able to form Al3+ ions. In fact Al is highly electropositive metal. However, down the group , due to poor shielding effect of intervening d and f orbitals, the increased effective nuclear charge holds n s electrons tightly (responsible for inert pair effect ) and thereby , restricting their participation in bonding. As a result of this only p-orbital electron may be involved in bonding. In fact in Ga , In and Tℓ, both +1 and +3 oxidation states are observed. The relative stability of +1 oxidation state progressively increases for heavier elements : Al < Ga < In < Tℓ. In thallium +1 oxidation state is predominant whereas the +3 oxidation state is highly oxidizing in character. The compounds in +1 oxidation state , as expected from energy considerations , are more ionic than those in +3 oxidation state.
In trivalent state, the number of electrons around the central atom in a molecule of the compounds of these elements (e.g., boron in BF3) will be only six. Such electron deficient molecules have tendency to accept a pair of electrons to accept a pair of electrons to achieve stable electronic configuration and thus, behave as Lewis acids. The tendency to behave as Lewis acid decreases with the increase in the size down the group. BCl3 easily accepts a lone pair of electrons from ammonia BCl3 NH3.
AlCl3 achieves stability by forming a dimmer.
In trivalent state most of the compounds being covalent are hydrolysed in water. For example, trichlorides on hydrolysis in water form [M(H2O)4]- species. ; the hybridization state of element M is sp3. Aluminium chloride in acidified aqueous solution forms octahedral [M(H2O)6] 3+ ion. In this complex ion, the 3d orbitals of Al are involved and the hybridization state of Al is sp3d2.
Problem
01. Standard electrode potential values , Eө for Al3+/Al is - 1.66 V and that of Tℓ3+/Tℓ is +1.26V. Predict about the formation of M3+ ion in solution and compare the electropositive character of the two metals.
Reactivity towards air
Boron is unreactive in crystalline form. Aluminium forms a very thin oxide layer on the surface which protects the metal rom further attack. Amorphous boron and aluminium metal on heating in air form B2O3 and Al2O3 respectively. With nitrogen at high temperature they form nitrides.
The nature of these oxides varies down the group. Boron trioxide is is acidic and reacts with basic(metallic) oxides forming metal borates. Aluminium and gallium oxides are amphoteric and those of indium and thallium are basic in their properties.
(ii) Reactivity towards acids and alkalies
Boron does not react with acid and alkalies even at moderate temperature ; but aluminium dissolves in mineral acids and aqueous alkalies and thus shows amphoteric character.
Aluminium dissolves in dilute HCl and liberates dihydrogen.
2 Al(s) + 6 HCl(aq) ® 2 Al3+(aq) + 6 Cl-(aq) + 3 H2(g)
However, concentrated nitric acid renders nitric aluminium passive by forming a protective oxide layer on the surface.
Aluminium also reacts with aqueous alkali and liberates dihydrogen.
2 Al(s) + 2 NaOH(aq) + 6H2O(ℓ) ® 2 Na[Al(OH)4]-(aq)
sodium tetrahydroxoaluminate(III)
+3 H2(g)
(iii) Reactivity towards halogen
These elements react with halogens to form trihalides (except TℓI3)
2 M(s) + 3 X2(g) ® 2 MX3(s) ( X = F, Cl, Br , I)
Problem
02. White fumes appear around the bottle of anhydrous aluminium chloride. Give reason.
IMPORTANT TRENDS AND ANOMALOUS PROPERTIES OF BORON
Certain important trends can be observed in the chemical behaviour of group 13 elements. The trichlorides , bromides and iodides of all these elements being covalent in nature are hydrolysed in water. Species like tetrahedral [M(OH)4]- and octahedral [M(OH)6] 3+ except boron exist in aqueous medium.
The monomeric trihalides , being electron deficient , are strong Lewis acids. Boron trifluoride easily react with Lewis bases such as NH3 to complete octet around boron.
F3B + : N H3 ® F3B¬NH3
It is due to the absence of d-orbitals that the maximum covalence of B is 4. Since the d-orbitals are available with Al and other elements, the maximum covalency can be expected beyond 4. Most of other metal halides ( e.g., AlCl3 ) are dimerised through halide bridging (e.g., Al2Cl6). The metal species completes its octet by accepting electrons from halogen in these halogen bridged molecules.
Problem
03. Boron is unable to form BF63- ion. Explain .
SOME IMPORTANT COMPOUNDS OF BORON
Some important compounds of boron are borax, orthoboric acid and diborane.
BORAX (Na2B4O7 .10 H2O)
Preparation
Borax can be prepared by the following methods.
1. From Tincal : Tincal (naturally occurring borax) is present in dried up lakes. It is taken up and boiled with water. The solution is filtered to remove insoluble impurities and concentrated by cooling to get crystals of borax.
2. From Colemanite : The mineral colemanite (Ca2B6O11) is finely powdered and is then boiled with washing soda solution (sodium carbonate).
The solution is filtered to remove the precipitated calcium carbonate. It is then concentrated and cooled to get the crystals of borax. The mother liquor containing sodium metaborate is treated with a current of CO2 gas which converts it into borax .
4 NaBO2 + CO2 ® Na2B4O7 + Na2CO3
3. From Boric acid : Borax can be prepared from boric acid by neutralising with sodium carbonate and upon concentrating and cooling the solution, crystals of borax Na2B4O7 10 H2O separates out.
4 H3BO3 + Na2CO3 ® Na2B4O7 + 6 H2O + CO2
Properties
The important properties of borax are as follows :
1. Borax is a white crystalline solid sparingly soluble in cold water but readily soluble in hot water.
2. Basic nature : An aqueous solution of borax is basic in nature because the base (NaOH) is stronger than the acid (H3BO3) formed as a result of hydrolysis.
3. Action of acids : Borax reacts with HCl or H2SO4 solution to form boric acid.
Na2B4O7 + 2 HCl + 5 H2O ® 2 NaCl + 4 H3BO3 boric acid
Na2B4O7 + H2SO4 + 5 H2O ® Na2SO4 + 4 H3BO3
4. Action with ethyl alcohol and sulphuric acid : When borax is heated with ethyl alcohol and concentrated sulphuric acid vapours of ethyl borate are formed which when ignited burn with green edged flame. The reaction is used to test the presence of BO33- ion in a salt.
Na2B4O7 + H2SO4 + 5 H2O ® Na2SO4 + 4 H3BO3
H3BO3 + 3 C2H5OH ® (C2H5O)3B + 3 H2O
triethyl borate
5. Action of heat : When powdered borax is heated strongly in a flame of bunsen burner , it forms a colourless transparent glassy bead made of sodium metaborate and boric anhydride.
Borax-bead test : Some coloured cations (basic radicals) such as Co2+, Ni2+, Cr3+, Cu2+, Mn2+ etc give specific colours with boric anhydride present in the bead. To perform this test, the bead is initially deposited on the loop of a platinum wire. A few crystals of coloured salt are deposited on the bead and the wire is heated in bunsen flame. The salts get converted to their oxides which then combine with B2O3 to form metaborates with specific colours.
Structure of borax
In borax , two boron atoms are in triangular geometry and two boron atoms are in tetrahedral geometry (Fig).
[B4O5(OH)4]2- ion in borax molecule
The ion [B4O5(OH)4]2- and the remaining eight water molecules are associated with the two sodium ions. Hence the borax contains tetranuclear units [B4O5(OH)4]2- , and therefore is formulated as Na2[[B4O5(OH)4]. 8 H2O
Uses of borax
1. For manufacturing enamels, glazes, optical glass, soaps and drying oils.
2. As a flux for soldering and welding.
3. In the candle industry for stifening the candle wick.
4. In the manufacture of washing powders and soaps .
5. As an antiseptic.
6. An analytical reagent .
BORIC ACID , H3BO3 or B(OH) 3
Preparation
Boric acid can be prepared by the following methods.
1. From borax : A hot concentrated solution containing borax is boiled with hydrochloric acid or sulphuric acid. The solution upon concentration and cooling gives crystals of boric acid.
Na2B4O7 + 2 HCl + 5 H2O ® 4 H3BO3 + 2 NaCl
Borax boric acid
Na2B4O7 + H2SO4 + 5 H2O ® 4 H3BO3 + Na2SO4
2. From Colemanite : Sulphur
Ca2B6O11 + 4 SO2 + 11 H2O ® 2 Ca(HSO3)2 + 6 H3BO3
Colemanite calcium bisulphite boric acid
3. From boron compounds by hydrolysis : Certain boron compounds upon boiling with water give boric acid.
BCl3 + 3 H2O ® H3BO3 + 3 HCl
BN + 3 H2O ® H3BO3 + NH3
Properties of Boric acid
1. Boric acid or orthoboric acid is a white crystalline solid and has a soapy touch.
2. It is less soluble in water but readily dissolves upon heating.
3. Acidic nature : Boric acid is a very weak monobasic acid (Ka = 1 x 10-9 ) . It is a Lewis acid and accepts an electron pair from OH- ion ( in water molecule)
B(OH)3 + H-OH == [B(OH)4]- + H+
4. Action of heat : When heated , boric acid loses H2O molecules in three stages and finally changes to boron trioxide as follows:
Structure of boric acid
The ground state outer electronic configuration of boron is 2s22px1. In the excited state , one of the 2s electrons get promoted to the vacant 2py orbital. The three half-filled atomic orbitals (2s, 2Px and 2Py) thus obtained undergo sp2 hybridisation to give three sp2 hybridised orbitals. Each of these three sp2 orbitals overlaps with 2p-orbitals of O- forming three B-O- bonds.
Therefore , BO33- (borate) ion has trigonal planar structure as shown in Fig.
In boric acid, planar BO33- units are joined by hydrogen bonds to give a layer structure as shown in Fig.
Structure of boric acid
(dotted lines represent hydrogen bond)
Uses of Boric acid
1. Boric acid is used for the manufacture of enamels and glazes for pottery.
2. Boric acid is a preservative for milk and food stuffs.
3. Boric acid is used in making of borosilicate glass.
4. Boric acid can be used as an eye wash because of its antiseptic nature.
Problem
04. Why is boric acid considered as a weak acid ?
Boron hydrides (Boranes)
In view of the trivalency of boron, we would expect to form a simple hydride, BH3, which however does not exist. The simplest boron hydride known is diborane, B2H6. Besides diborane, boron forms a number of higher hydrides known as boranes : most of them have molecular formulae corresponding to BnHn+4 and BnH n+6.
DIBORANE , B2H6
Preparation
1. Treat boron trifluoride with LiAlH4 in diethyl ether.
2. A convenient laboratory method for the preparation involves the oxidation of sodium borohydride with iodine.
3. Diborane is produced on an industrial scale by reaction of BF3 with sodium hydride.
Properties
1. Diborane is a colourless , highly toxic gas with b.p of 180 K.
2. On heating between 373 – 523 K , it changes into higher boranes.
This property of diborane is used for the preparation of higher boranes.
3. It burns in oxygen releasing an enormous amount of energy.
B2H6 + 3 O2 ® B2O3 + 3 H2O : DcH°= - 1976 kJ mol-1
This is why diborane is used as a rocket fuel.
4. Diborane undergoes hydrolysis readily by water to give boric acid.
B2H6(g) + 6 H2O(l) ® 2 B(OH)3(aq) + 6 H2(g)
5. Diboranes undergoes cleavage reactions with Lewis bases (L) to give borane adducts BH3.L
B2H6 + 2 NMe3 ® 2 BH3 . NMe3
B2H6 + 2 CO ® 2 BH3 . CO
6. Diborane reacts with ammonia at 450 K to form borazole (inorganic benzene).
Borazole has the following structure which is similar to that of benzene. Hence borazole is also referred to as inorganic benzene.
Structure of Diborane
The structure of diborane is shown in Fig.
The structure of Diborane
The four terminal hydrogen atoms and the two boron atoms lie in one plane. Above and below this plane, there are two bridging hydrogen atoms. The four terminal B–H bonds are regular two centre-two electron bonds while the two bridge (B–H–B) bonds are different and can be described in terms of three centre-two electron bond as shown in Fig.
Bonding in diborane
Bonding in Diborane
Each B atom uses sp3 hybrides for bonding. Out of the four sp3 hybrides on each boron atom , one without an electron shown in broken lines. The terminal B–H bonds are normal two centre-two electron bonds but the two bridge bonds are three centre- two electron bonds. The three-centre two electron bonds are also referred to as banana bonds.
Uses
1. Diborane is used as a fuel for supersonic rockets.
2. For preparing a number of other boron hydrides such as LiBH4 and NaBH4 etc.
3. As a reducing agent in organic reactions.
Tetrahydridoborates
Tetrahydridoborates of several metals are known. Lithium and sodium tetraborates, also known as borohydrides are prepared by the reaction of metal hydrides with B2H6 in diethyl ether.
2 MH + B2H6 ® 2 M+[BH4]– : M = Li or Na
Both LiBH4 and NaBH4 are used as reducing agents in organic synthesis. They are useful starting materials for preparing other metal borohydrides.
BORON HALIDES
The boron trihalides have the general formula BX3, which are monomeric covalent species. They are planar and bonding can be explained in terms of sp2 hybridisation. They do not dimerise like BH3 because the lone pairs on the halogens can interact with the vacant p-orbitals on boron.
Preparation
1. Boron trifluoride is obtained by heating B2O3 with CaF2 and concentrated sulphuric acid:
B2O3 + 3 CaF2 + 3 H2SO4 ® 2 BF3 + 3 CaSO4 + 3 H2O
2. The other boron halides are prepared by direct union of the elements at elevated temperatures.
Properties
1. Boron halides fume in moist air and are highly sensitive to moisture. Boron trifluoride undergoes slow and partial hydrolysis with water but other halides react violently with water.
BCl3 + 3 H2O ® B(OH)3+ 3 HCl
2. Boron halides are strong Lewis acids and form adducts with bases such as amines, ethers, sulphides or phosphines.
BX3 + NR3 ® X3B.NR3
Uses
Because of its strong Lewis acid character, BF3 is used as a catalyst in several industrial processes.
Uses of boron
1. Metal borides are used in atomic reactors as protective shields and control rods because of the high ability of 10B isotope to absorb neutrons.
2. It is used as a semi-conductor for making electronic devices.
3. It is used in steel industry for increasing the hardness of steel.
4. Boron compounds are becoming increasingly important as rocket fuels because of their high energy/weight ratio.
5. Boron filaments are used in making light composite materials for aircrafts.
6. Boron is an essential element in plant metabolism.
7. Boron carbide fibres are very hard but light and hence are used for making bullet-proof vests.
Uses of Alumnium
1. Aluminum is used extensively in industry and every day life.
2. It forms alloys with Cu, Mn, Mg , Si and Mn. Aluminium and its alloys can be given shapes of pipe, tubes , rods , wires, plates or foils and therefore find uses in packing , utensil making , construction, aeoroplane and transportation industry.
3. The use of aluminium and its compounds for domestic purposes is reduced considerably because of their toxic nature.
GROUP 14 ELEMENTS
THE CARBON FAMILY
The group 14 of the periodic table contains five elements carbon(C), silicon(Si ) , germanium (Ge) , tin (Sn) and lead (Pb). This group is known as carbon family.
Occurrence and uses
All the elements are well known except germanium. Carbon occurs both as free element ( graphite or diamond) and in the combined form (mainly as carbonates of Ca, Mg and other electropositive elements) . It also occurs as CO2, an important constituent of atmosphere. Carbon is seventeenth in the order of abundance in earth’s crust. Silicon is the second most abundant element in the earth’s crust after oxygen as silica (sand or quartz) and silicates. In silicates, [SiO4] unit may occur individual group or linked to form chains, rings , sheets or three-dimensional frameworks.
Germanium occurs rarely. Germanium minerals are extremely rare but the element is distributed in trace amounts in coal and zinc carbonates. Tin occurs as cassiterite, SnO2 and lead is found as galena, PbS. The abundance of the elements in the earth’s crust by weight is shown below :
Element | Abundance in earth’s crust (ppm) |
C | 180 |
Si | 2.72 x 105 |
Ge | 1.5 |
Sn | 2.1 |
Pb | 13 |
Carbon is used extensively in its different forms. Coal is used as a fuel in boilers, engines, furnaces, etc. It is also used for the manufacture of coal gas, water gas, producer gas and synthetic petrol. Charcoal(activated) is used as an excellent adsorbent to purify and decoluorize sugar and other chemicals. It is also used to adsorb poisonous gases in gas masks and for removing offensive odour from the air conditioning processes. Graphite is a good conductor of electricity and is used for making electrodes and carbon rods and in covering moulds for electrodeposition of copper. It is also used in steel making , metal foundries for crucibles, as a lubricant and in pencils etc. It is also used as the moderator in the cores of gas cooled nuclear reactors to slow down neutrons. Diamonds (allotrope of carbon) are cut as gemstones and used in jewellery and other articles. It is also used for industrial purposes mainly for making drills or as an abrasive powder for cutting and polishing.
Silicon is used as n-type or p-type semiconductors when doped with Group 15 or Group 13 element respectively. Silicon and germanium are extensively used in very pure forms in semiconductor devices, which are the basis of the whole electronic industry including computer hardware. Silicon is added to steel or iron as such or more usually in the form of ferrosilicon to increase its resistance to attack by acids. Very pure silicon is used to make computer chips. Its alloys such as silicon bronze possess strength and hardness even greater than steel.
TABLE – 1
Atomic and Physical properties of Group 14 Elements
Property | Carbon | Silicon | Germanium | Tin | Lead |
Atomic number | 6 | 14 | 32 | 50 | 82 |
Atomic mass | 12.01 | 28.09 | 72.60 | 118.71 | 207.2 |
Electronic configuration | [He]2s22p2 | [Ne]3s23p2 | [Ar]3d104s24p2 | [Kr]4d105s25p2 | [Xe]4f145d106s26p2 |
Covalent radius (pm) | 77 | 118 | 122 | 140 | 146 |
Ionic radus M4+ (pm) Ionic radus M2+ (pm) | - - | 40 - | 53 73 | 69 118 | 78 119 |
Ionisation enthalpy(kJ/mol) I II I I I IV | 1086 2352 4620 6220 | 786 1577 3228 4354 | 761 1537 3300 4409 | 708 1411 2942 3929 | 715 1450 3081 4082 |
Electronegativity | 2.5 | 1.8 | 1.8 | 1.8 | 1.9 |
Density (g/cm3) | 3.51 | 2.34 | 5.32 | 7.26 | 11.34 |
Melting point (K) | 4373 | 1693 | 1218 | 505 | 600 |
Boiling point (K) | - | 3550 | 3123 | 2896 | 2024 |
Electrical resistivity (ohm cm) | 1014 – 1016 | 50 | 50 | 10-5 | 2 x 10-5 |
Germanium has largest use in transistor technology, in making transistors and other semiconductor devices. It is transparent to infrared light and therefore is also used for making prisms and lenses and windows in infra red spectrometers and other scientific apparatus.
Because of low strength and high cost of tin, it is rarely used by itself but it used for electroplating and as alloys. Tin plates obtained by electroplating steel with tin are extensively used for making cans for food and drinks. It is used in the preparation of a number of important alloys such as solder (Sn/Pb)
bronze(Pb/Cu/Sn), babbit (Sn/Pb or Cu/Pb/Sn/Sb), pewter (Pb/Sn/Sb/Cu) , type metal(Pb/Sn/Sb), etc. Lead is used for making lead sheets, lead pipes etc. It is also used for making telegraph and telephone wires which are to be burried in earth. It is used in storage batteries, making bullets. It is commonly used in making pigments like chrome yellow, chrome red, red lead, white lead , etc. Lead is also used for making important alloys such as type metal, solder, pewter etc.
General Characteristics of Group 14 Elements
Electronic configurations
The atoms of these elements have four electrons in the outermost shell , two in s and 2 in p-subshell. The general electronic configuration of this group may be expressed as ns2np2. The electronic configuration of the atoms of this group are given in TABLE.
Element | Symbol | Atomic number | Electronic Configuration |
Carbon | C | 6 | [He]2s22p2 |
Silicon | Si | 14 | [Ne]3s23p2 |
Germanium | Ge | 32 | [Ar]3d104s24p2 |
Tin | Sn | 50 | [Kr]4d105s25p2 |
Lead | Pb | 82 | [Xe]4f145d106s26p2 |
Atomic and Physical Properties
The common physical constants of Group 14 elements are shown in TABLE 1.
1. ATOMIC RADII
The atomic radii of Group 14 elements are less than the corresponding elements of Group 13. However, the atomic radii increase down the family.
Explanation :
The decrease in atomic radii is due to increase in effective nuclear charge on going from Group 13 element to Group 14 element within the same period. As a result , the outermost electrons are attracted more strongly towards the nucleus and therefore , atomic radius decreases. Within the group, atomic radii increase on going down the group, due to the increase in the number of electron shells.
2. IONISATION ENTHALPIES
The first ionisation enthalpies of these elements are higher than the corresponding members of Group 13 elements.
Explanation : The higher ionisation enthalpies are due to the higher nuclear charge and smaller size of atoms of Group 14 elements. While moving down the group, the ionisation enthalpies decrease. This is due to the increase in atomic size and screening effect which overweigh the effect of increase in nuclear charge. Therefore, the outermost electron becomes less and less tightly held by the nucleus and ionisation enthalpy decreases.
3. Melting Points
The atoms of this group form covalent bonds with each other and therefore , there are strong binding forces between their atoms in both solid and liquid states. Consequently, the melting and boiling points of Group 14 elements are much higher in comparison to Group 13 elements. On moving down the group, the boiling points decrease.
4. Metallic character
Due to large ionisation enthalpies, the elements of this group are less metallic than the elements of Group 13. On moving down the group , the metallic character increases from carbon to lead. For example, carbon is typical non-metal , silicon also behaves as non-metal, but in certain physical properties , it behaves as semi-metal. Germanium is metalloid while tin and lead are typical metals.
Electronegativity
Due to small size , the elements of this group are slightly more electronegative than group 13 elements. The electronegativity values for elements from Si to Pb are almost the same.
Chemical Properties
The group 14 elements have four electrons in the outermost shell. The common oxidation states exhibited by some elements are +4 and +2. Carbon also exhibits negative oxidation states. Since the sum of first four ionization enthalpies is very high, compounds in +4 oxidation state are generally covalent in nature. In heavier members the tendency to show +2 oxidation state increases in the sequence Ge < Sn < Pb. It is due to the inability of ns2 electrons of valency shell to participate in bonding. The relative stabilities of these oxidation states vary down the group. Carbon and silicon mostly show shown +4 oxidation state. Germanium forms stable compounds in +4 state and only few compounds in +2 state. Tin forms compounds in both oxidation states (Sn in +2state is a reducing agent) . Lead compounds in +2 state are stable and in and in +4 state are strong oxidizing agents. In tetravalent state the number of electrons around the central atom in a molecule (e.g., carbon in CCl4) is eight. Being electron precise molecules, they are normally not expected to act as electron acceptor or electron donor species. Although carbon cannot exceed its covalence more than 4, other elements of the group can do so. It is because of the presence of d-orbital in them. Due to this , their halides undergo hydrolysis and have a tendency to form complexes by accepting electron pairs from donor species. For example, the species like [SiF6]2-, [GeCl6]2- , [Sn(OH)6]2- exist where the hybridization state of central atom is sp3d2.
(i) Reactivity towards oxygen
All members when heated in oxygen form oxides. There are mainly two types of oxides, i.e., monoxide and oxide of formula MO and MO2 respectively. SiO only exists at high temperature. Oxides in higher oxidation states of elements are generally more acidic than those in lower oxidation states. The oxides – CO2 , SiO2 and GeO2 are acidic, whereas SnO2 and PbO2 are amphoteric in nature. Among monoxides , CO is neutral , GeO is distinctly acidic whereas SnO and PbO are amphoteric.
Problem
05. Select the member(s) of group 14 that (i) forms the most acidic dioxide (ii) is commonly found in +2 oxidation state (iii) used as semiconductor.
iii) Reactivity with water
Carbon, silicon and germanium are not affected by water. Tin decomposes steam to form dioxide and dihydrogen gas.
Lead is unaffected by water, probably because of a protective oxide film formation.
iv) Reactivity towards halogen
These elements can form halides of formula MX2 and MX4 (where X = F, Cl, Br , I ). Except carbon , all other members react directly with halogen under suitable condition to make halides. Most of the MX4 are covalent in nature. The central metal atom in these halides undergo sp3 hybridisation and the molecule is tetrahedral in shape. Exceptions are SnF4 and PbF4 which are ionic in nature. PbI4 does not exist because Pb – I bond initially formed during the reaction does not release enough energy to unpair 6s2 electrons and excite one of them to higher orbital to have four unpaired electrons around lead atom. Heavier members from Ge to Pb are able to make halides of formula MX2. Stability of dihalides increases down the group. Considering the thermal and chemical stability , GeX4 is more stable than GeX2 , whereas PbX2 is more stable than PbX4. Except CCl4 , other tetrachlorides 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.
Hydrolysis can be be understood by taking the example of SiCl4 . It undergoes hydrolysis by initially accepting lone pair of electrons from water molecule in d-orbitals of Si , finally leading to the formation of Si(OH)4 as shown below.
Problem
06. [SiF6]2- is known whereas [SiCl6]2-not. Give reasons.
IMPORTANT TRENDS AND ANOMALOUS BEHAVIOUR OF CARBON
Like first member of other groups , carbon also differs from the rest of the members of its group. It is due to its smaller size, higher electronegativity, higher ionization enthalpy and unavailability of d-orbitals.
In carbon, only s and p-orbitals are available for bonding and therefore , it can accommodate only four pairs of electrons around it. This would limit the maximum covalence to four whereas other members can expand their covalence due to the presence of d-orbitals.
Carbon also has unique ability to form pp – pp multiple bonds with itself and with other atoms of smaller size and high electronegativity. Few examples are : C = C , C º C , C = O , C = S and C º N . Heavier elements do not form pp – pp bonds because their atomic orbitals are too large and diffuse to have effective overlapping.
Carbon atoms have a tendency to link with another through covalent bonds to form chains and rings. This property is called catenation. This is because C – C bonds are very strong. Down the group the size increases and electronegativity decreases and thereby , tendency to show catenation decreases. This can be clearly seen from bond enthalpy values. The order of catenation is C >> Si > Ge » Sn.
Lead does not show catenation.
Bond Bond enthalpy kJ mol-1
C – C 348
Si – Si 297
Ge – Ge 260
Sn – Sn 240
Due to property of catenation and pp – pp bond formation , carbon is able to show allotropic forms.
ALLOTROPES OF CARBON
Elemental carbon exists in several crystalline and amorphous forms. The allotropes of carbon are diamond, graphite and fullerenes.
DIAMOND
It occurs in nature. It can also be prepared artificially , but because of high cost and poor quality , diamonds are seldom made artificially.
Structure
In diamond , carbon is sp3 hybridised. Each carbon is tetrahedrally linked to neighbouring carbon atoms through four strong C- C, sp3 hybridised, s - bonds. The net work extends in three-dimensions (Fig ) and is very rigid.
Structure of diamond
Properties
1. Diamond is the purest form of carbon.
2. The C- C bond length in diamond is 154 pm.
3. Since diamond exists as a three dimensional net work solid , it is the hardest substance known with high density and melting point.
4. Since all the electrons are firmly held in C- C s - bonds, there are no free electrons in a diamond crystal. Therefore diamond is a bad conductor of electricity.
5. Because of its high refractive index (2.5) , diamond can reflect and refract light. Therefore diamond is a transparent substance.
Uses
1. Because of its hardness, diamond is used for cutting glass, making borers for rock drilling and for making abrasives.
2. When diamond is cut and polished, brilliant light is refracted from its surfaces. That is why diamond is used for making precious gems and jewellery.
3. It is used for grinding and polishing of hard materials.
4. Diamond is used for making dies for drawing thin wires from metals.
GRAPHITE
It occurs in nature and can also be manufactured artificially by heating coke to 3273 – 3300 K in electric furnace.
Graphite has a layer structure(Fig) in which each carbon atom is bonded to three other carbon atoms to form a hexagonal sheet. Each carbon atom forms three sigma bonds by means of sp2 hybid orbitals.
Structure of graphite
The remaining unhybridised p-orbital of each carbon atom having an electron overlaps with its counterparts on adjacent carbon atoms in the same layer to form p-bonds. The C- C bond distance within the hexagonal layer is 142 pm which corresponds to a C- C bond order of 1.33. The hexagonal sheets are held together by weak van der Waal’s forces. It is easy for one layer to slide over another, graphite is soft and slippery material with excellent lubricating properties. The p-electrons within the layer are free to move and this accounts for the electrical conductivity of graphite. The different structures adopted by diamond and graphite thus nicely explain their contrasting properties. Whereas diamond is hard, transparent and non-conductor of electricity, graphite is soft, opaque and a good conductor of electricity.
Uses of Graphite
1. Graphite is used to make electrodes for electrolytic cells.
2. Being soft and greasy , it is used to lubricate the parts of the machines.
3. Graphite crucibles can withstand very high temperatures and can be used for high melting substances.
4. Graphite is used to moderate the fast moving neutrons in nuclear reactors.
5. Mixed with wax and clay, graphite is used for making cores of lead pencils as it can mark paper black. It is , therefore, often called black lead or plumbago.
FULLERENES
Fullerenes were discovered by H.W Kroto, R.F Curl and R.E. Smalley. The fullerenes consists of hollow cages of carbon atoms. They are large spheroidal molecules of composition Cn ; two important members of this family are C60 and C70 . The 1996 Nobel prize in Chemistry was awarded to the above scientists for the discovery of fullerenes – the new form of carbon.
Fullerenes were originally made by the evaporation of graphite using a laser. A more practical method involves the heating of graphite in an electric arc in the presence of inert gas such as helium or argon. The sooty material formed by condensation of vaporised Cn molecules consists of mainly C60 with smaller quantity of C70 and traces of other fullerenes consisting of even number of carbon atoms upto 350 or above. C60 and C70 can be readily separated from the fullerene soot by extraction in toluene followed by chromatographic separation over alumina. In contrast to graphite or diamond, the fullerenes dissolve in organic solvents to give coloured solutions. A solution of C60 in toluene is purple whereas that of C70 is orange-red. Fullerenes are the only pure form of carbon because they do not have ‘’dangling’’ edge or surface bonds to be attracted to other atoms as is the case with diamond and graphite.
The beautiful structure of C60 is shown in Fig.
The structure of C60, Buckminsterfullerene
It has been named Buckminsterfullerene in honour of American architect Robert Buckminster Fuller who designed goedesic dome structures having hexagonal and pentagonal patterns. The general name FULLERENE refers to the family of spheroidal carbon cage molecules. The shape of C60 resembles that of a soccer ball. It contains twelve five-membered rings and twenty five six-membered rings. Six-membered rings are fused both to other six-membered rings and five membered rings , but the five-membered rings are connected only to six membered rings. All the carbon atoms are equivalent. There are both single and double bonds with C- C distances of 145.3 and 138.3 pm respectively.
Applications of fullerenes
1. They are good lubricants because the balls can roll between surface.
2. Alkali compounds of C60 (A3C60) are super conducting materials even at high temperatures of the order of 10-40 K.
AMORPHOUS FORMS OF CARBON
Carbon also exists in a number of amorphous forms.
1. Coal : It is crude form of carbon. It has been formed in nature as a result of slow decomposition of vegetable matter under the influence of heat, high pressure and absence of air. This process is also called carbonisation. It is found in different forms which represent the stages of transformation of vegetable matter. The forms along with carbon contents are listed : Peat (60%), lignite(67%), bituminous (88%), steam coal(93%), anthracite(95%). Bituminous is the variety of coal commonly used while anthracite is the purest variety which burns with non-smoky flame. When coal is subjected to destructive distillation (heating in absence of air) it loses all the volatile organic matter to leave behind coke as residue. Coal is mainly used as a fuel and also in the large scale preparation of coal gas, coke and synthetic petrol.
2. Coke : Coke is a greyish black hard solid which is left as residue when coal is subjected to destructive distillation, i.e., distilled in the absence of air. It is used as a fuel and also as a reducing agent in metallurgical operations
3. Charcoal : It is black, soft and highly porous substance which exists in the following forms.
(a) Wood charcoal : It is obtained by heating wood or nut-shells strongly in limited supply of air and the volatile gases are allowed to escape. It is mainly elemental carbon (95 to 98%). It is used in gas masks to remove poisonous gases and also as a fuel for domestic purposes.
(b) Animal charcoal : It is also called bone charcoal and is obtained by destructive distillation of bones. It contains 10 to 12% carbon and the rest is calcium phosphate. Animal charcoal is generally used to adsorb colouring matter from sugar cane juice and also to decolourise certain impure compounds that are formed in the laboratory.
(c) Sugar charcoal : Sugar charcoal is formed by the action of concentrated H2SO4 on sucrose (C12H22O11). The acid being a strong dehydrating agent , removes all the hydrogen and oxygen atoms from sugar as H2O molecules leaving behind as black mass called sugar charcoal.
Sugar charcoal is also a very good adsorbent to remove colouring matter.
Activated charcoal : All the three forms of charcoal are highly porous and are used to adsorb gases on their surface. The adsorbing power can be increased by heating the charcoal in contact with steam for some time. The charcoal thus formed is known as activated charcoal and is a better adsorbent than any ordinary form of charcoal.
4. Lamp black or carbon black : It is another amorphous allotropic form which can be obtained by burning certain carbon rich compounds such as kerosene, petroleum turpentine oil etc. in the limited supply of air. The unburnt particles of carbon can be deposited as soot on the coarse wet blankets placed in chambers where combustion is carried. It can be removed later on. It is a pure form of carbon (contains 90-98% carbon) and is used in making black inks, paints and shoe polishes etc.
INORGANIC CARBON COMPOUNDS OXIDES OF CARBON
Carbon forms two main oxides , carbon monoxide(CO)and carbondioxide (CO2) by direct combination with oxygen. In addition to these oxides, carbon suboxide (C3O2) has also been prepared by heating malonic acid with phosphorus pentoxide.
Carbon monoxide
Preparation
1. It is formed by incomplete combustion of carbon and carbon containing fuels.
2 C + O2 ® 2 CO
This type of incomplete combustion occurs during burning of petrol or diesel in automobiles and therefore , CO is always present in automobile exhausts. It is also present in volcanic gases and gases coming out of furnaces.
2. Carbon monoxide is a constituent of water gas (CO + H2) which is formed by passing steam over red hot coke.
Carbon monoxide can be separated by liquefaction.
3. Carbon monoxide can be prepared by passing carbon dioxide through red hot charcoal.
From the mixture, CO2 can be removed by passing the mixture into water under pressure.
4. Carbon monoxide can be obtained by the reduction of oxides of heavy metals with carbon.
ZnO + C ® Zn + CO
Fe2O3 + 3 C ® 2 Fe + 3 CO
5. In the laboratory carbon monoxide can be prepared by the dehydration of formic acid with concentrated sulphuric acid.
6. It can also be prepared by heating potassium ferrocyanide with concentrated sulphuric acid.
K4Fe(CN)6 + 3 H2SO4 + 6 H2O ® 2 K2SO4 + FeSO4
+ 3 NH4(SO4)2 + 6 CO
Structure
In CO molecule, both C and O atoms are sp-hybridised. One sp-hybrid orbital each of C and O overlap to C- O. s-bond while the other sp-orbital on each carbon atom contains lone pair of electrons. The two unhybridised p- orbitals of C and O form two pp - pp bonds. The CO is a linear molecule. The electronic structure of carbon monoxide may be represented as follows:
Due to the presence of a lone pair of electrons on the carbon atom, CO act as a Lewis base or ligand and can form a co-ordinate bond with metals ( M¬Cº O ) to form metal carbonyls.
Carbon monoxide can be represented as a resonance hybrid of the following structures.
The presence of a triple bond between C and O is supported by the following evidences :
i) The carbon-oxygen bond length is just 113 pm which corresponds to a carbon-oxygen triple bond.
ii) The dipole moment of CO is very low due to back donation of a pair of electrons from more electronegative O to the less electronegative C-atom.
Properties
1. It is a neutral oxide.
2. It is a colourless and odourless gas which is only slightly soluble in water.
3. Carbon monoxide is highly poisonous (toxic) in nature. Its toxic nature is due to its ability to form a stable complex with the haemoglobin present in red blood cells to form carboxyhaemoglobin as discussed below :
Haemoglobin + CO ⇌ Carboxyhaemoglobin
In the lungs, haemoglobin combines with molecular oxygen loosely and reversibly to form oxyhaemoglobin.
Haemoglobin + O2 ⇌ Oxyhaemoglobin
Oxyhaemoglobin thus formed in the lungs then travels to all parts of the body through blood stream and delivers O2 to the various tissues of the body. However, CO combines with haemoglobin irreversibly (i.e., forms stronger bonds than O2). Therefore, if CO is present, it will form stable complex with haemoglobin (i.e., carboxyhaemoglobin) which destroys the oxygen carrying capacity of haemoglobin. As a result, haemoglobin does not take up oxygen easily therefore causing suffocation and ultimately death.
4. It burns with blue flame forming carbon dioxide.
2 CO + O2 ® 2 CO2 + 110.5 kJ
5. Since CO can be easily oxidised to CO2 , it acts as a powerful reducing agent. As such it reduces many metal oxides to their respective metals.
ZnO + CO ® Zn + CO2
CuO + CO ® Cu + CO2
Fe2O3 + 3 CO ® 2 Fe + 3 CO2
It also reduces PdCl2 to Pd and I2O5 to I2.
PdCl2 + CO + H2O ® Pd + CO2 + 2 HCl
I2O5 + 5 CO ® I2 + 5 CO2
6. It combines with many transition metals such as iron, cobalt, nickel etc. forming nickel carbonyls. For example,
Nickel carbonyl is volatile. When heated to 440 – 450 K , it decomposes to form pure nickel.
Therefore , nickel carbonyl is used for the purification of nickel by Mond’s process.
7. Carbon monoxide is readily absorbed by a solution of CuCl in Con HCl or NH3 due to the formation of soluble complexes. For example,
CuCl + NH3 + CO ® [Cu(CO)NH3]+ Cl-
Soluble complex
CuCl + HCl + CO ® H+[Cu(CO)Cl2]-
Soluble complex
USES
1. It is an important constituent of two industrial fuels , i.e., water gas and producer gas
Water gas or synthesis gas is a mixture of carbon monoxide and hydrogen and is formed by passing steam over red hot coke.
Producer gas is a mixture of carbon monoxide and nitrogen and is formed when air is passed over red hot coke.
2. It is used in Mond’s process for purification of nickel via its nickel carbonyl.
3. It is used in the manufacture methyl alcohol , synthetic petrol, sodium formate etc.
4. In the metallurgy of iron as a reducing agent.
Fe2O3 + 3 CO ® 2 Fe + 3 CO2
5. Iron carbonyl is used in the manufacture of magnetic tapes for vedioes and tape recorders.
CARBON DIOXIDE
Preparation
1. It is prepared by burning carbon, fossil fuels and other organic compounds in air or oxygen.
C + O2 ® CO2
CH4 + 2 O2 ® CO2 + 2 H2O
2. In the laboratory , it is prepared by the action of acids on carbonates.
CaCO3 + 2 HCl ® CaCl2 + CO2 + H2O
3. On an industrial scale , it is obtained as a by-product in some industrial processes such as manufacture of lime, manufacture of ethyl alcohol etc.
Structure
In CO2 molecule , carbon is sp-hybridised, it forms two s-bonds with two oxygen atoms and two pp- pp multiple bonds. As a result , CO2 is linear, monomeric covalent compound.
The electron dot structure for CO2 may be represented by either formula I or II.
This structure predicts that both the carbon-oxygen bond length in CO2 should be equal and should have a typical bond length of 122 pm. However, experimentally , it has been found that the carbon-oxygen bond length in CO2 is only 115 pm. This can be explained , if carbon dioxide is considered to a resonance hybrid of the following structures:
Due to resonance, carbon-oxygen bond length acquires some triple bond character and hence bond length decreases from 122 pm to 115 pm.
Properties
The physical and chemical properties of CO2 are quite different from those of CO. Some important physical properties of CO and CO2 are given in the TABLE.
Property | CO | CO2 |
Melting pint (K) | 68 | 216.4 at 5.2 atm |
Boiling point (K) | 81.5 | 194.5(sublimes) |
Density (g/L) | 1.250 | 1.977 |
C- O bond length (pm) | 112 | 115 |
Heat of formation DfH (kJ/mol) | - 110.5 | - 393.5 |
1. It is a colourless and odourless gas about 1.5 times heavier than air.
2. Unlike CO, CO2 is not poisonous . However, it does not support life in animals and human beings. In fact animals and human beings die in an atmosphere CO2 due to lack of oxygen.
3. Ordinarily, CO2 is neither combustible nor supporter of combustion. However, certain active metals like Na, K , Mg etc continue burning in it.
2 Mg + CO2 ® 2 MgO + C
4. It is slightly soluble in water. Its solubility in water increases with increase in pressure. Soda water and other aerated soft drinks are, in fact solutions of carbon dioxide in water (containing sugar, some flavouring and colouring agents) under pressure.
5. Acidic nature : Carbon dioxide dissolves in water to some extent to form carbonic acid.
CO2 + H2O ® H2CO3 (carbonic acid)
Carbonic acid being a weak dibasic acid dissolves in alkalies forming two series of salts, i.e., bicarbonates or hydrogen carbonates (HCO3- ) and carbonates (CO32- ). For example, when CO2 is passed through lime water , it turns lime water milky due to the formation of insoluble calcium carbonate.
Ca(OH)2 + CO2 ® CaCO3 + H2O
lime water insoluble
However if carbon dioxide is passed for a longer period, the turbidity disappers due to the formation of soluble calcium bicarbonate.
CaCO3 + H2O + CO2 ® Ca(HCO3)2 (soluble)
6. Photosynthesis : Carbon dioxide is absorbed by plants. In the presence of chlorophyll and sunlight , the absorbed carbon dioxide combines with water to form glucose and starch, which is used as food by plants. This process is called photosynthesis.
By this process plants make food for themselves as well as for animals and human beings. Unlike CO, it is not poisonous. But the increase in combustion of fossil fuels and decomposition of lime stone for cement manufacture increases the CO2 content of the atmosphere. This may lead to increase in green house effect and thus raise the temperature of atmosphere which might have serious consequences.
7. Action of ammonia : When carbon dioxide is reacted with liquid ammonia at 453 – 473 K under a pressure of 220 atmospheres, it first forms ammonium carbamate which subsequently decomposes to urea.
USES
Carbon dioxide is used :
1. in the preparation of aerated waters.
2. As a fire extinguisher because it is a non-supporter of combustion.
3. In the manufacture of washing soda by solvay ammonia process.
4. Solid carbon dioxide is used as a refrigerant under the commercial name dri kold.
5. For artificial respiration ( for victims of CO poisoning) as a mixture of 95% O2 and 5% CO2 under the name carbogen.
6. For the purification of sugar in sugar industry.
DRY ICE
Dry ice is the name given to solid carbon dioxide. It is also called cardice. It is obtained when CO2 is cooled under pressure (50-60 atm). When solid carbon dioxide is allowed to evaporate in air under one atmospheric pressure , it changes directly into the gaseous state without liquefying. As a result, unlike ordinary ice it does not wet the surface on which it melts. Therefore , it is called dry ice.
Dry ice is used for making cold baths in the laboratory by mixing it with volatile organic solvents. For example, a temperature of 196 K is obtained by using a mixture of dry ice and acetone and a temperature of 165 K is obtained by using a mixture of dry ice and ether. It is also used as a coolant for preserving perishable articles in food industry. Dry ice is also used for curing local burns and in hospitals for surgical operation of sores.
Carbonates and bicarbonates
Carbonates of many metals are known but only the bicarbonates of alkali metals exist in the solid state. Bicarbonates of calcium and magnesium are responsible for the temporary hardness of water. Several carbonates such as Na2CO3 10 H2O (washing soda), K2CO3 , CaCO3 and NaHCO3 (baking soda) are commercially important chemicals.
Silicon Dioxide SiO2
95% of earth’s crust is made up of silica and and silicates. Silicon dioxide , commonly known as silica , occurs in several crystallographic forms. Quartz , cristobalite and tridymite are some of the crystalline forms of silica and they are interconvertable at suitable temperature. Silicon dioxide is a covalent , three-dimensional net work solid in which each silicon atom is covalently bonded in a tetrahedral manner to four oxygen atoms. Each oxygen atom in turn covalently bonded to another silicon atoms as shown in diagram.
Three dimensional structure of SiO2
Silica in its normal form is almost non-reactive because of very high Si – O bond enthalpy. It resists the attack by halogens, dihydrogen and most of the acids and metals even at elevated temperatures. However, it is attacked by HF and NaOH.
SiO2 + 2 NaOH ® Na2 SiO3 + H2 O
SiO2 + 4 HF ® SiF4 + 2 H2 O
Quartz is extensively used as piezoelectric material ; it is made possible to develop extremely accurate clocks , modern radio and television broadcasting and mobile radio communications. Silica gel is used as a drying agent and as a support for chromatographic materials and catalysts. Kieselghur an amorphous form of silica is used in filtration plants.
SILICONES
Silicones are polymeric compounds containing Si-O-Si linkages. These are polymers , which contain R2SiO repeating units. These have the general formula (R2SiO)n. These may be linear, cyclic or crossed linked. These have very high thermal stability and are called high temperature polymers ( R may be alkyl or phenyl group). The starting material for the manufacture of silicones is alkyl substituted chlorosilanes. These are obtained by the reaction of alkyl halides with silicon in the presence of metallic copper which acts as catalyst.
The polymers are obtained by the hydrolysis of the above chloroderivative as :
When two molecules of dialkyl silanol combine , we get a dimer with the elemination of a molecule of water.
Since an active –OH group is left on each end of the chain, polymerisation reaction continues and length of the chain increases. It forms linear silicone as :
The hydrolysis of monoalkyltrichloro silanes, RSiCl3 gives cross linked polymers. By regulating the conditions, the condensation can be stopped at any stage and the chains or rings of desired lengths can be obtained.
Properties of silicone polymers
1. Silicone polymers are highly stable towards heat.
2. Low molecular weight silicone polymers are soluble in organic solvents like ether, carbon tetrachloride , benzene etc.
3. They are stable towards chemical reagents. They are not affected by weak acids, alkalies and salt solutions.
4. They are water repellants because of organic side chain.
5. They are good electrical insulators.
6. They are resistant to oxidation. However, when heated in air to 350°C to 400°C, silicones are rapidly oxidised and this leads to cross linking.
Uses of silicon polymers
The important uses of silicone polymers are :
(i) Silicone polymers are used for high temperature oil baths, high vaccum pumps etc.
(ii) Silicone polymers are used as greases, varnishes and these can be used even at low temperatures ( of the order of -40°C).
(iii) Because silicones are water repellants and good insulators, they are used for water proofing and in electrical condensers.
(iv) They are used as lubricants at both high and low temperatures.
(v) Silicone rubbers are very useful because they retain their elasticity at lower temperatures as compared to other rubbers. They are also mixed in paints to make them damp resistant.
(vi) They are used as excellent insulators for electric motors and other electrical appliances.
Silicates
A large number of silicate minerals exist in nature. Some of the examples are feldspar, zeolites , mica and asbestos. The basic structural unit of silicates is SiO44- in which silicon atom is bonded to four oxygen atoms in tetradron fashion.
In silicates either the discrete unit is present or a number of such units are joined together via corners by sharing 1 , 2 , 3 or 4 oxygen atoms per silicate units. When silicate units are linked together, they form chain , ring , sheet or three-dimensional structures. Negative charge on silicate structure is neutralized by positively charged metal ions. If all the four corners are shared with other tetrahedral units , three dimensional network is formed.
Two important man-made silicates are glass and cement.
Zeolites
If aluminium atoms replace few silicon atoms in three dimensional network of silicon dioxide, overall structure known as aluminosilicate , acquires a negative charge. Cations such as Na+, K+ or Ca2+ balance the negative charge. Examples are feldspar and zeolites. Zeolites are widely used as catalyst in petrochemical industries for cracking of hydrocarbons and isomerisation , e.g., ZSM-5 (a type of silicate) used to convert alcohols directly into gasoline. Hydrated zeolites are used as ion exchangers in softening of ‘hard water’.
QUESTIONS
1. Write balanced equations for the reaction of elemental boron with elemental chlorine, oxygen and nitrogen at high temperature.
2. Discuss the pattern of variation in the oxidation state of (i) B to Tℓ and (ii) C to Pb
3. How can you explain higher stability of BCl3 as compared to TiCl3 ?
4. Why does boron trifluoride behave as a Lewis acid ?
5. Consider the compounds , BCl3 and CCl4. How will they behave with water ? Justify.
6. Is boric acid a protic acid ? Explain.
7. Describe the shapes of BF3 and BH4- . Assign the hybridization of boron in these species.
8. Write reactions to justify amphoteric nature of aluminium.
9. Write the resonance structures of CO32- and HCO3-.
10. What is the state of hybridization of carbon in (a) CO32- (b) diamond (c) graphite ?
11. Explain the difference in properties of diamond and graphite on the basis of their structures.
12. Rationalise the given statements and give chemical reactions :
(i) Lead(II) chloride reacts with Cl2 to give PbCl4.
(ii) Lead (IV) chloride is highly unstable towards heat.
(iii) Lead is known not to form an iodide PbI4.
13. Suggest reasons why the B – F bond lengths in BF3 (130 pm) and BF4- (143 pm) differ.
14. If B – Cl bond has a dipole moment , explain why BCl3 molecule has zero dipole moment.
15. Aluminium trifluoride is insoluble in anhydrous HF but dissolves on addition of NaF. Aluminium trifluoride precipitates out of the resulting solution when gaseous BF3 is bubbled through. Give reasons.
16. Suggest the reason why CO is poisonous.
17. How is excessive content of CO2 responsible for global warming ?
18. What happens when :
(a) borax is heated strongly heated.
(b) boric acid is added to water
(c) aluminium is treated with dilute NaOH
(d) BF3 is reacted with ammonia ?
19. Explain the following :
(a) silicon is heated with methyl chloride at high temperature in presence of copper.
(b) silicon dioxide is treated with hydrogen fluoride
(c) CO is heated with ZnO
(d) hydrated alumina is treated with aqueous NaOH solution.
20. Give reasons :
(a) Con HNO3 can be transported in aluminium containers.
(b) A mixture of dil NaOH and aluminium pieces is used to open drain.
(c) graphite is used as a lubricant.
(d) diamond is used as an abrasive
(e) aluminium alloys are used to make aircraft body.
(f) aluminium wire is used to make transmission cables.
21. Explain why there is a phenomenal decrease in ionization enthalpy from carbon to silicon ?
22. What are allotropes ? Sketch the structure of two allotropes of carbon namely diamond and graphite. What is the impact of structure on physical properties of two allotropes ?
23. (a) Classify the following oxides as neutral, acidic , basic or
amphoteric ?
CO, B2O3 , SiO2, CO2, Al2O3, PbO2, Tℓ2O3
(b) Write suitable chemical equations to show their nature.
24. In some reactions thallium resembles aluminium , whereas in others it resembles with group 1 metals. Support this statement by giving some evidence.
25. When a metal X is treated with sodium hydroxide, a white precipitate (A) is obtained, which is soluble in excess of NaOH to give soluble complex(B). Compound (A) is soluble in dilute HCl to form compound (C) . The compound A when heated strongly gives (D) , which is used to extract metal. Identify (X) , (A) , (B) , (C) and (D). Write suitable equations to support their identities.
26. What do you understand by (a) catenation (b) inert pair effect (c) allotropy.
27. A certain salt X gives the following results:
(a) its aqueous solution is alkaline to litmus.
(b) it swells up to glassy material Y on strong heating.
(iii) when conc. H2SO4 is added to a hot solution of X, white
crystal of an acid Z separates out.
Write equations for all the above reactions and identify X, Y
and Z.
28. Write balanced equations for :
(a) BF3 + LiH ®
(b) B2H6 + H2O ®
(c) NaH + B2H6 ®
(d) Al + NaOH ®
(e) B2H6 + NH3 ®
29. Explain why there is a phenomenal decrease in ionization e
30. Why are boron halides and diborane reffred to as ‘’electron deficient’’ compounds ?
31. Write the structures of diborane and explain the nature of bonding in it.
32. What happens when borax solution is acidified ? Write balanced equation for the reaction.
33. Describe what happens when boric acid is heated.
34. Describe how elemental boron can be prepared.
35. Describe the shapes of BF3 and BH4-. What types of hybridisation can we assign to boron in each of these compounds ?
36. Why do boron halides form ‘addition compounds’ with amines ?
37. By means of a balanced chemical equation show how B(OH)3 behaves as an acid in water.
38. Write balanced equations for the following reactions :
(a) Combustion of C4H10 in a limited supply of oxygen to form carbon monoxide and water.
(b) The reaction of calcium carbide with water to form acetylene.
(c) Formation of hydrogen cyanide from methane and ammonia.
39. Write resonance structures of CO32- and HCO3-.
40. What is the oxidation state of carbon in each of the following compounds ?
(a) CO (b) HCN (c) H2CO3 (d) CaC2
41. Explain the differences in the properties of diamond and graphite on the basis of their structures.
42. What is dry ice ? Why is it so called ?
43. What the hybridisation state of carbon in :
(a) CO32- (b) HCN (c) diamond (d) graphite
44. What are fullerenes ? How are they prepared ?
45. Give examples of compounds in which nitrogen exhibits oxidatioin states of - 3, +3 and +5.
46. Discuss the conditions required in the Haber process for manufacture of ammonia.
47. How is dinitrogen prepared in the laboratory ?
48. Write the products of the following reactions (give balanced equations)
(a) Mg3N2 + H2O ®
(b) I2 + HNO3 (con) ®
(c) Cu + HNO3 (Con) ®
(d) Li + N2 ®
(e) HNO3 + P2O5 ®
49. How would you prepare a sample of deuterated ammonia ND3 ?
50. Illustrate how nitrogen compounds provide good examples of multiple bonding and resonance ?
51. Determine the oxidation number of nitrogen in :
(a) N2O (b) NO2 (c) HNO3 (d) NH3
52. Write the resonance Lewis structures for N2O , NO2 and NO3-.
53. Illustrate how a metal such as copper can give different products with nitric acid. Give balanced chemical equations.
54. Write the balanced equations for the manufacture of nitric acid by oxidation of ammonia.
55. Describe the uses of ammonia and nitric acid.
56. What is the shape of ozone molecule ?
57. Write Lewis dot structures for ozone molecule and explain why O-O distances in ozone are equal.
58. What is the importance of ozone for plant life on earth ?
59. Classify the following compounds into acidic, basic and amphoteric oxides :
SrO, SiO2 , N2O5, P4O6 , Cl2O7. Write balanced equations for the reaction of these compounds with water , a base or an acid as the case may be.
60. Why does boron not form B3+ ions ?
61. What is the oxidation state of boron in boric acid ? What is the basicity of the acid ?
62. What is liquid nitrogen used for ?
63. Why does Con. HNO3 turns yellow on standing in sun light ?
64. Why does iron become passive when dipped in Con. HNO3 ?
65. What is calcium cyanamide ? Why is used as a fertiliser ?
66. In the ring test of nitrates what chemical compound is formed ?
67. What is azote ?
68. Phosphorus forms PCl5 but nitrogen does not for NCl5. Why ?
69. Ammonia is a good complexing agent. Explain.
70. Give one reaction in which ammonia acts as a reducing agent.
71. How does sodium hydride react with diborane ?
72. Arrange BF3, BCl3 and BBr3 in decreasing order of their Lewis acid character.
73. Name the anhydride of nitric acid.
74. What happens when lead nitrate is heated ?
75. Who discovered nitrogen ?
76. Ozone is used for purifying air in crowded places . Explain.
77. Why oxide ion is called hard ion ? Explain.
78. N2O supports combustion more vigorously than air. Explain.
79. Write the balanced equations for the reaction of elemental boron with elemental chlorine, oxygen and nitrogen at high temperature.
80. Why are boron halides and diborane referred to as electron deficient compounds ?
81. Write structure of diborane and explain the nature of bonding in it.
82. What happens when borax solution is acidified ? Write balanced equation for the reaction.
83. Describe what happens when boric acid is heated ?
84. Describe the shapes of BF3 and BH4-. What types of hybridisation can be assigned to boron in each of these compounds ?
85. By means of a balanced chemical equation show how B(OH)3 behaves as an acid in water.