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THE P-BLOCK ELEMENTS, The elements in which the last electron enters in the valence p-sub shell are called the p-block, elements. They include elements of the groups 13 to 18. Their general outer electronic configuration is, ns2np1-6 (except He which has 1s2 configuration).They includes metals, non-metals and metalloids., , Group 15 Elements, Group 15 includes nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi). As, we go down the group, the metallic character increases. Nitrogen and phosphorus are non-metals, arsenic, and antimony metalloids and bismuth is a typical metal. The valence shell electronic configuration of these, elements is ns2np3. The s orbital in these elements is completely filled and p orbitals are half-filled, making, their electronic configuration extra stable., Covalent and ionic radii increase down the group. There is a considerable increase in covalent radius, from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the, presence of completely filled d or f orbitals in heavier members. Ionisation enthalpy decreases down the, group due to gradual increase in atomic size. Because of the extra stable half-filled p orbitals and smaller, size, the ionisation enthalpy of the group 15 elements is much greater than that of group 14 elements., Oxidation states and trends in chemical reactivity, The common oxidation states of these elements are –3, +3 and +5. The tendencies to exhibit –3, oxidation state decreases down the group due to increase in size and metallic character. The last member of, the group, bismuth does not form any compound in –3 oxidation state. The stability of +5 oxidation state, decreases and that of +3 state increases (due to inert pair effect) down the group. Nitrogen exhibits + 1, + 2,, and + 4 oxidation states also when it reacts with oxygen. Phosphorus also shows +1 and +4 oxidation states, in some oxoacids. Nitrogen is restricted to a maximum covalency of 4 since only four orbitals (one s and, three p) are available for bonding., Anomalous properties of nitrogen, Nitrogen differs from the rest of the members of this group due to its smaller size, high electro, negativity, high ionisation enthalpy and non-availability of d orbitals. Some of the anomalous properties, shown by nitrogen are:, 1. Nitrogen has the ability to form pπ-pπ multiple bonds with itself and with other elements like C and, O. Other elements of this group do not form pπ-pπ bonds., 2. Nitrogen exists as a diatomic molecule with a triple bond (one s and two p) between the two atoms., So its bond enthalpy is very high. While other elements of this group are poly atomic with single, bonds., 3. The single N–N bond is weak. So the catenation tendency is weaker in nitrogen., 4. Due to the absence of d orbitals in its valence shell, the maximum covalency of nitrogen is four, 5. N cannot form dπ–pπ bond. While Phosphorus and arsenic can form dπ–dπ bond with transition, metals and with C and O., Hydrides of Group 15 Elements, All the elements of Group 15 form hydrides of the type EH3 (where E = N, P, As, Sb or Bi). The, hydrides show regular gradation in their properties. The bond dissociation enthalpy of E – H decreases from, NH3 to BiH3. So the thermal stability decreases from NH3 to BiH3 and the reducing character increases., Ammonia is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the hydrides., Basicity decreases in the order NH3 > PH3 > AsH3 > SbH3 > BiH3. The melting point of these hydrides, increases from top to bottom. This is due to increase in the atomic size of the central atom which increases, the van der Waal’s force of attraction. NH3 has the highest melting and boiling points due to inter molecular, hydrogen bonding. All these hydrides have pyramidal geometry., Q1. Though nitrogen exhibits +5 oxidation state, it does not form pentahalides. Give reason., , The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 1

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Nitrogen with n = 2, has s and p orbitals only. It does not have d orbitals to expand its covalence, beyond four. That is why it does not form pentahalide., Q2. PH3 has lower boiling point than NH3. Why?, Unlike NH3, PH3 molecules are not associated through inter molecular hydrogen bonding in liquid, state. That is why the boiling point of PH3 is lower than NH3., Dinitrogen (N2), Preparation: Dinitrogen is produced commercially by the liquefaction and fractional distillation of air., In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with, sodium nitrite., NH4CI(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq), It can also be obtained by the thermal decomposition of ammonium dichromate., (NH4)2Cr2O7, →Heat N2 + 4H2O + Cr2O3, Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide., Ba(N3)2 → Ba + 3N2, Properties, Dinitrogen is inert at room temperature because of the high bond enthalpy of N≡ N bond. At higher, temperatures, it directly combines with some metals to form ionic nitrides and with non-metals to form, covalent nitrides., 6Li + N2 →Heat 2Li3N, 3Mg + N2 →Heat Mg3N2, It combines with hydrogen at about 773 K in the presence of a catalyst (Haber’s Process) to form ammonia:, N2 +3H2, Fe/ 773K 2NH3, Dinitrogen combines with dioxygen at very high temperature (at about 2000 K) to form nitric oxide, N2 + O2 → 2 NO, Uses: 1. The main use of dinitrogen is in the manufacture of ammonia and other industrial chemicals, containing nitrogen (e.g., calcium cyanamide)., 2. It also used to create an inert atmosphere in metallurgy., 3. Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery., , Ammonia, Preparation: In laboratory, ammonia is obtained by treating ammonium salts with caustic soda (NaOH) or, slaked lime., (NH4)2SO4 + 2NaOH → 2NH3 + 2H2O + Na2SO4, 2NH4Cl + Ca(OH)2 → 2NH3 + 2H2O + CaCl2, On a large scale, ammonia is manufactured by Haber’s process., N2(g) + 3H2(g) → 2NH3(g), In accordance with Le Chatelier’s principle, high pressure of about 200 atm, a temperature of about, 773 K and the catalyst such as iron oxide with small amounts of K2O and Al2O3 are employed to increase, the rate of this reaction., Properties, Ammonia is a colourless gas with pungent smell. It is highly soluble in water because of its ability to, form inter molecular hydrogen bond with water. Liquid ammonia has high melting and boiling points, because of inter molecular hydrogen bonding., , The ammonia molecule has a trigonal pyramidal geometry. It has three bond pairs and one, lone pair of electrons., Its aqueous solution is weakly basic due to the formation of OH– ions., The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 2

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Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic, acid, carbon to carbon dioxide, sulphur to H2SO4, and phosphorus to phosphoric acid., I2 + 10HNO3 → 2HIO3 + 10 NO2 + 4H2O, C + 4HNO3 → CO2 + 2H2O + 4NO2, , S8 + 48HNO3(conc.) → 8H2SO4 + 48NO2 + 16H2O, P4 + 20HNO3(conc.) → 4H3PO4 + 20 NO2 + 4H2O, , Brown Ring Test: It is a test used for the detection of nitrates. The test is carried out by adding dilute ferrous, sulphate solution to an aqueous solution containing nitrate ion, and then carefully adding concentrated, sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and, sulphuric acid layers indicate the presence of nitrate ion in solution., NO3- + 3Fe2+ + 4H+ → NO + 3Fe3+ + 2H2O, [Fe (H2O)6 ]2+ + NO → [Fe (H2O)5 (NO)]2+ + H2O, (brown ring), Uses: It is used i) in the manufacture of ammonium nitrate for fertilizers and other nitrates for use in, explosives and pyrotechnics. ii) for the preparation of nitroglycerin, trinitrotoluene and other organic nitro, compounds.iii) in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels., , Phosphorus, The allotropic forms of phosphorus:, Phosphorus exists mainly in three allotropic forms – white (yellow) phosphorus, red phosphorus and black, phosphorus, 1. White phosphorus: It is a translucent white waxy solid. It is poisonous, insoluble in water but soluble, in carbon disulphide and glows in dark (chemiluminescence). It dissolves in boiling NaOH solution in, an inert atmosphere giving PH3 (phosphine)., P4 + 3NaOH + 3H 2 O → PH3 + 3NaH 2 PO2 (sodium hypophosphite), White phosphorus is less stable and therefore, more reactive. This is because in white phosphorus, the P-P-P, bond angles are only 60°. So it has greater angular strain and highly unstable., It readily catches fire in air to give dense white fumes of P4O10., P4 + 5O2 → P4O10, It consists of discrete tetrahedral P4 molecule, , 2. Red phosphorus: It is obtained by heating white phosphorus at 573K in an inert atmosphere for several, days. Red phosphorus has iron grey lustre. It is odourless, non-poisonous and insoluble in water as well, as in carbon disulphide. Chemically, red phosphorus is much less reactive than white phosphorus. It, does not glow in the dark. It contains polymeric chains of P4 tetrahedra., , 3. Black phosphorus: It has two forms- α-black phosphorus and β-black phosphorus. α-black, phosphorus is formed when red phosphorus is heated in a sealed tube at 803K. It does not oxidise in air., β-Black phosphorus is prepared by heating white phosphorus at 473K under high pressure. It does not, burn in air up to 673K, The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 5

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Structure:, , 5., , H3PO4 [Orthophosphoric Acid], It is obtained by the action of water on phosphorus pentoxide (P4O10), P4O10 + 6 H2O → 4 H3PO4, It is also called Phosphoric acid. It’s a tribasic acid and has a tetrahedral shape., Structure:, , 6., , H4P2O7 [Pyrophosphoric Acid], It is obtained by heating Phosphoric acid at about 2500c., 2 H3PO4 → H4P2O7 . It’s a tetra basic acid., Structure:, , 7., , (HPO3)n [Metaphosphoric acid], It is obtained by heating phosphorus acid with Br2 vapours in a sealed tube., H3PO3 + Br2 → HPO3 + 2HBr, Structure: It exists as a trimer or a polymer as follows:, , The oxoacids of phosphorus in +3 oxidation state undergo disproportionation (i.e. simultaneously oxidised, and reduced). For example, orthophophorous acid (or phosphorous acid) on heating disproportionates to, give orthophosphoric acid (phosphoric acid) and phosphine., 4H3PO3, , 3H3PO4 + PH3, , Group 16 Elements, The members of this group are oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po)., They are also called chalcogens (means ore producing). Oxygen and sulphur are non-metals, selenium and, tellurium are metalloids, while polonium is a radioactive metal., Ionisation enthalpy of these elements decreases down the group. It is due to increase in size., However, the elements of this group have lower ionisation enthalpy values compared to those of Group15, elements. This is due to the fact that Group 15 elements have extra stable half- filled p orbitals electronic, The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 8

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configurations., Oxygen atom has less negative electron gain enthalpy than sulphur because of the compact nature of, its shells due to which the electronic repulsion is greater., .Oxidation states: The elements of Group 16 exhibit a number of oxidation states (-2,+2,+4 & +6). The, stability of -2 oxidation state decreases down the group. Since electronegativity of oxygen is very high, it, shows only –2 oxidation state (except in the case of OF2 where its oxidation state is + 2). Other elements of, the group exhibit + 2, + 4 & + 6 oxidation states. But + 4 and + 6 are more common. Sulphur, selenium and, tellurium usually show + 4 oxidation state in their compounds with oxygen and + 6 with fluorine. Down the, group, the stability of + 6 oxidation state decreases and that of + 4 oxidation state increases (due to inert pair, effect)., Hydrides of 16th group elements, All the elements of Group 16 form hydrides of the type H2E (E = S, Se, Te, Po). Their acidic, character increases from H2O to H2Te. This is due to the decrease in bond (H–E) dissociation enthalpy down, the group. So the thermal stability also decreases down the group. All the hydrides except water possess, reducing property and this character increases from H2S to H2Te., , Dioxygen (O2), Preparation: (i) By heating chlorates, nitrates and permanganates., (ii) By the thermal decomposition of the oxides of metals low in the electrochemical series and higher, oxides of some metals., 2Pb3O4(s) → 6PbO(s) + O2(g), 2Ag2O(s) → 4Ag(s) + O2(g);, 2HgO(s) → 2Hg(l) + O2(g) ;, 2PbO2(s) → 2PbO(s) + O2(g), (iii) By the decomposition of Hydrogen peroxide (H2O2) in presence of manganese dioxide., 2H2O2(aq) → 2H2O(1) + O2(g), (iv) On large scale it can be prepared from water or air. Electrolysis of water leads to the release of, hydrogen at the cathode and oxygen n at the anode. It is also obtained by the fractional distillation of air., Properties:, Dioxygen directly reacts with metals and non-metals (except with some metals like Au, Pt etc and, with some noble gases)., e.g. 2Ca + O2 → 2CaO, P4 + O2 → P4O10, 4 Al + 3O2 → 2Al2O3, C + O2 → CO2, Uses: 1) oxygen is used in oxyacetylene welding, in the manufacture of many metals, particularly steel., , 2) Oxygen cylinders are widely used in hospitals, high altitude flying and in mountaineering., 3) Liquid O2 is used in rocket fuels., , Oxides, Oxides are binary compounds of oxygen with other elements. There are two types of oxides – simple, oxides (e.g., MgO, Al2O3 ) and mixed oxides (Pb3O4, Fe3O4), Simple oxides can be further classified on the basis of their acidic, basic or amphoteric character. An, oxide that combines with water to give an acid is called acidic oxide (e.g., SO2, Cl2O7, CO2, N2O5)., Generally, non-metal oxides are acidic but oxides of some metals in higher oxidation states also have acidic, character (e.g., Mn2O7, CrO3, V2O5 etc.)., The oxide which gives an alkali on dissolved in water is known as basic oxide (e.g., Na2O, CaO,, BaO). Generally, metallic oxides are basic in nature., Some metallic oxides exhibit a dual behaviour. They show the characteristics of both acidic and basic, oxides. Such oxides are known as amphoteric oxides. They react with acids as well as alkalies., E.g.: Al2O3, Ga2O3 etc., There are some oxides which are neither acidic nor basic. Such oxides are known as neutral oxides., Examples of neutral oxides are CO, NO and N2O., The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 9

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Ozone (O3), Ozone is an allotropic form of oxygen., Preparation: When a slow dry stream of oxygen is passed through a silent electric discharge, oxygen is, converted to ozone. The product is known as ozonised oxygen., 3 O2(g)→ 2 O3(g); ∆H= +142 kJ/mol, Since the formation of ozone from oxygen is an endothermic process, a silent electric discharge, should be used, unless the ozone formed undergoes decomposition., Properties: Pure ozone is a pale blue gas, dark blue liquid and violet-black solid. Ozone has a characteristic, smell., Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen, results in the liberation of heat (∆H is negative) and an increase in entropy (∆S is positive). So the Gibbs, energy change (∆G) for this process is always negative (∆G = ∆H – T∆S)., Due to the ease with which it liberates nascent oxygen (O3 → O2 + O), it acts as a powerful, oxidising agent., For e.g., it oxidises lead sulphide to lead sulphate, PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g), Oxides of nitrogen (particularly nitric oxide) combine very rapidly with ozone and deplete it. Thus, nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes, slowly depleting the, concentration of the ozone layer in the upper atmosphere., NO2(g) + O2(g), NO(g) + O3(g), Estimation of ozone: When ozone reacts with an excess of potassium iodide solution buffered with a borate, buffer, iodine is liberated. The liberated iodine can be titrated against a standard solution of sodium, thiosulphate. This is a quantitative method for estimating O3 gas., Structure: O3 has an angular structure. It is a resonance hybrid of the following two forms:, , Uses: It is used as a germicide, disinfectant and for sterilising water. It is also used for bleaching oils,, ivory, flour, starch, etc. It acts as an oxidising agent in the manufacture of potassium permanganate., Allotropes of Sulphur, Sulphur forms a large number of allotropes. Among these yellow rhombic (α-sulphur) and, monoclinic (β -sulphur) forms are the most important. The stable form at room temperature is rhombic, sulphur, which transforms to monoclinic sulphur when heated above 369 K., 1. Rhombic sulphur (α-sulphur), It is prepared by evaporating the solution of roll sulphur in CS2. It is insoluble in water but readily soluble in, CS2., 2. Monoclinic sulphur (β-sulphur), It is prepared by melting rhombic sulphur in a dish and cooling, till a crust is formed. Two holes are made in, the crust and the remaining liquid is poured out. On removing the crust, colourless needle shaped crystals of, β-sulphur are formed. It is stable above 369 K and transforms into α-sulphur below it. At 369 K both the, forms are stable. This temperature is called transition temperature., Both rhombic and monoclinic sulphur have S8 molecules. The S8 ring in both the forms is puckered, and has a crown shape., , The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 10

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(iii) absorption of SO3 in H2SO4 to give Oleum (H2S2O7)., SO3 + H2SO4 → H2S2O7, (iv) Dilution of oleum with water gives H2SO4 of the desired concentration., H2S2O7 + H2O → 2H2SO4, Properties, Sulphuric acid is a colourless, dense, oily liquid. It dissolves in water with the evolution of a large, quantity of heat. Hence, for diluting the acid, the concentrated acid must be added slowly into water with, constant stirring., Chemical properties: The chemical reactions of sulphuric acid are due to the following reasons:, (a) its low volatility, (b) strong acidic character, (c) strong affinity for water and, (d) its ability to act as an oxidising agent., In aqueous solution, sulphuric acid ionises in two steps., H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4HSO4-(aq) + H2O(l) → H3O+(aq) + SO42So it is dibasic and forms two series of salts: normal sulphates and acid sulphates., Because of its low volatility sulphuric acid can be used for the manufacture of more volatile acids from their, corresponding salts., 2 MX + H2SO4 → 2 HX + M2SO4 (where X = F, Cl, NO3 etc. and M is a metal), Concentrated sulphuric acid is a strong dehydrating agent and drying agent. Many wet gases can be dried by, passing them through sulphuric acid. Sulphuric acid removes water from organic compounds, e.g.: C12H22O11 + H2SO4 → 12C + 11H2O, Hot concentrated sulphuric acid is a moderately strong oxidising agent. It oxidises both metals and nonmetals and the acid itself reduces to SO2., Cu + 2 H2SO4(conc.) → CuSO4 + SO 2 + 2H 2O, S + 2H 2SO4(conc.) → 3SO2 + 2H2O, C + 2H2SO4(conc.) → CO2 + 2 SO2 + 2 H2O, Uses: The important uses of Sulphuric acid are:, 1) In the manufacture of fertilizers 2) in petroleum refining 3) in the manufacture of pigments, paints and, dyestuff intermediates 4) in detergent industry 5) in metallurgical applications 6) as electrolyte in storage, batteries 7) in the manufacture of nitrocellulose products and 8) as a laboratory reagent., , Group 17 Elements, Fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) are the members of Group 17., They are collectively known as the halogens (means salt producers). They are highly reactive non-metallic, elements. All these elements have seven electrons in their outermost shell (ns2np5) and so they do not, readily lose their electron. So they have very high ionisation enthalpy., Halogens have maximum negative electron gain enthalpy in the corresponding periods. This is due, to the fact that the atoms of these elements have only one electron less than stable noble gas configurations., Electron gain enthalpy of these elements decreases down the group. However, the negative electron gain, enthalpy of fluorine is less than that of chlorine. It is because, in fluorine the incoming electron goes to the, 2p subshell, but in Cl it enters in to the 3p subshell. Due to the compactness of 2p subshell compared to 3p, subshell, the electron – electron repulsion is greater in fluorine than in chlorine. So F does not easily gains, electron., Halogens have very high electronegativity. The electronegativity decreases down the group., Fluorine is the most electronegative element in the periodic table., All halogens have characteristic colour. For example, F2 has yellow, Cl2-greenish yellow, Br2-red, and I2, violet colour. This is due to absorption of radiations in visible region which results in the excitation, of outer electrons to higher energy level., , The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 12

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The bond dissociation enthalpy of F2 is low. This is due to the relatively large electron-electron, repulsion among the lone pairs in F2 molecule., All the halogens are highly reactive. They react with metals and non-metals to form halides. The, reactivity of the halogens decreases down the group., Halogens are strong oxidising agents since they readily accept electron. F2 is the strongest oxidising, halogen and it oxidises other halide ions in solution or in the solid phase., Oxidation states, All the halogens exhibit –1 oxidation state. Chlorine, bromine and iodine also show + 1, + 3, + 5 and, + 7 oxidation states in their oxides, oxy acids and in inter halogen compounds. Due to the absence of vacant, d orbitals and the maximum electronegativity, fluorine exhibits only –1 oxidation state., Anomalous behavior of fluorine, Due to the small size, highest electronegativity, low F-F bond dissociation enthalpy, and non, availability of d orbitals in valence shell, fluorine shows properties different from other halogens., Some of the anomalous properties of fluorine are:, 1. Ionisation enthalpy, electronegativity, enthalpy of bond dissociation and electrode potentials are higher, for fluorine than expected., 2. Ionic and covalent radii, m.p. and b.p. and electron gain enthalpy are quite lower than expected., 3. Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with, other elements)., 4. F forms only one oxoacid while other halogens form a number of oxoacids., 5. Hydrogen fluoride is a liquid due to strong hydrogen bonding. While the hydrogen halides of other, elements are gases., Hydrides of halogens, Halogens react with hydrogen to give hydrogen halides which dissolve in water to form hydrohalic, acids. The acidic strength of these acids varies in the order: HF < HCl < HBr < HI. The stability of these, halides decreases down the group due to decrease in bond dissociation enthalpy from HF to HI., Chlorine (Cl2), Preparation: It can be prepared by any one of the following methods:, (i) By heating manganese dioxide with concentrated hydrochloric acid., MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O, Conc. HCl can be replaced by a mixture of common salt and concentrated H2SO4, 4NaCl + MnO2 + 4H2SO4 → MnCl2+ 4NaHSO4 + 2H2O + Cl2, (ii) By the action of HCl on potassium permanganate., 2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2, Manufacture of chlorine, (i) Deacon’s process: By oxidation of hydrogen chloride gas by atmospheric oxygen in the presence of, CuCl2 (catalyst) at 723 K., 4HCl+O2 ⎯⎯⎯ CuCl2⎯→2Cl2 +2H2O, (ii) Electrolytic process: Chlorine is obtained by the electrolysis of brine solution (concentrated NaCl, solution). During electrolysis chlorine is liberated at the anode., Properties: It is a greenish yellow gas with pungent and suffocating odour. It is soluble in water., It reacts with a number of metals and non-metals to form chlorides., 2Al + 3Cl2 → 2AlCl3;, P4 + 6Cl2 → 4PCl3, 2Na + Cl2 → 2NaCl;, S8 + 4Cl2 → 4S2Cl2, 2Fe + 3Cl2 → 2FeCl3 ;, With excess ammonia, chlorine gives nitrogen and ammonium chloride whereas with excess chlorine,, nitrogen trichloride (explosive) is formed., 8NH3 + 3Cl2 → 6NH4Cl + N2;, NH3 + 3Cl2 → NCl3 + 3HCl, (excess), (excess), , The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 13

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Uses of noble gases:, Helium is used in filling balloons for meteorological observations. It is also used in gas-cooled, gas, nuclear reactors. Liquid helium is use, used as cryogenic agent for carrying out various experiments at low, temperatures. It is used as a diluent for oxygen in modern diving apparatus because of its very, ver low solubility, in blood., Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes. Neon, bulbs are used in botanical gardens and in green houses., Argon is used to provide an inert atmosphere in high temperature metallurgical processes and for, filling electric bulbs. It is also used in the laboratory for handling substances that are air-sensitive., air, Xenon and Krypton are used in light bulbs designed for special purposes., ¥, ¥, ¥, ¥, ¥, ¥, ¥, ¥, ¥, ¥, ¥, ¥, ¥, , The P Block Elements Anil Kumar K L, GHSS,Ashtamudi,Kollam [Hsslive.in], , Page 17