Chemistry of Group 15 Elements PDF

1 CHEMISTRY OF GROUP 15 ELEMENTS The group 15 elements are nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi) are also called the pnictogens. The rationalization of the properties of the group 15 elements and their compounds is difficult, despite...

1 CHEMISTRY OF GROUP 15 ELEMENTS The group 15 elements are nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi) are also called the pnictogens. The rationalization of the properties of the group 15 elements and their compounds is difficult, despite there being some general similarities in trends of the group 13, 14 and 15 elements, e.g. increase in metallic character and stabilities of lower oxidation states on descending the group. Very little of the chemistry of the group 15 elements is that of simple ions. Although metal nitrides and phosphides that react with water are usually considered to contain N3- and P3- ions, electrostatic considerations make it doubtful whether these ionic formulations are correct. The only definite case of a simple cation in a chemical environment is that of Bi3+, and nearly all the chemistry of the group 15 elements involves covalently bonded compounds. The thermochemical basis of the chemistry of such species is much harder to establish than that of ionic compounds. In addition, they are much more likely to be kinetically inert, both to substitution reactions (e.g. NF3 to hydrolysis, [H2PO2] to deuteration), and to oxidation or reduction when these processes involve making or breaking covalent bonds, as well as the transfer of electrons. Nitrogen, for example, forms a range of oxoacids and oxoanions, and in aqueous media can exist in all oxidation states from +5 to -3, e.g. [NO3]-, N2O4, [NO2]-, NO, N2O, N2, NH2OH, N2H4, NH3. Much the same is true about the chemistry of phosphorus. The chemistry of the first two members of group 15 is far more extensive than that of As, Sb and Bi, and we can mention only a small fraction of the known inorganic compounds of N and P. In our discussions, Their general outer- shell electronic configuration can be represented as ns2np3, with exactly half- filled np orbitals, leading to their extra stability and lesser reactivity. The penultimate shell of nitrogen has 2 electrons, phosphorus has 8, arsenic and antimony has 18 and bismuth has 18 electrons also. This slight difference in their electronic configuration leads to the difference in their properties. Element and Electronic structure Nitrogen (7N) [He] 2s22p3 Phosphorus (15P) [Ne] 3s23p3 Arsenic (33As) [Ar] 3d104s24p3 Antimony (51Sb) [Kr] 4d105s25p3 Bismith (83Bi) [Xe] 4f145d106s26p3  2 General Properties of Group 15 Elements 1. Physical State Nitrogen exists as a gas, phosphorus as a volatile solid, while the other elements are solids of varying melting points. Nitrogen can be liquefied only at a very low temperature. 2. Atomic and Ionic Size: Atomic size of Group 15 elements is smaller as compared to Group 14 elements and increases down the group in an irregular manner. The presence of 10d-electrons in As, Sb and Bi (with poor shielding effect) increases the effective nuclear charge and the increase in size is not as expected. In case of Bi, presence of more poorly shielding 14f electrons further affect the atomic size. Ionic size also increases down the group, as expected. 3. Ionisation Energy: The IE decreases regularly on moving down the group with increase in atomic size. The IE of Group 15 elements is higher as compared to the corresponding Group 14 elements in their respective periods due to their stable exactly half-filled orbitals. 4. Atomic Volume Density: The atomic volume and density increase down the group, as expected. 5. Metallic Character, Melting Point and Boiling Point: The metallic character increases down the group, as in case of Group 14 elements. Thus, N and P are typical nonmetals, As and Sb are metalloids and Bi is a metal. The melting points and boiling points increase down the group, but Bi has exceptionally low melting point. 6. Electronegativity: The electronegativity decreases gradually on moving down the group supporting for the gradual change in nonmetallic to metallic character. 7. Oxidation States: Due to ns2np3 configuration, these elements are expected to show +3 and +5 oxidation states. However, due to high electronegativity, smaller size and high IE, N rather shows -3 oxidation states. Covalency of five is not exhibited by N as it does not have empty d orbitals to expand its octet. However, it also shows +1 (N2O) and +2 (NO) oxidation. The stability of -3 oxidation state decreases down the group with decrease in electronegativity of the elements. Except N, other elements exhibit both +3 and +5 oxidation states in their compounds. The stability of +5 oxidation state decreases while that of +3 state increases down the group due to inert-pair effect. N and P also show the ability of lone-pair donation from their compounds to form adducts with Lewis acids and exhibit the oxidation state of +4. 8. Nature of Bonds: Due to ns2np3 configuration, these elements are expected to gain three elements (electrons) to attain noble-gas configuration and form M3- ions. However, due to energy constraints, only N and P are known to form nitrides and phosphides respectively. Further, due to high IE, ionic compounds are not formed but mostly covalent bonding is seen. In case of Sb and  3 Bi, the trivalent ions are formed but are rapidly hydrolysed to give SbO+ and BiO+ ions respectively in water. 9. Catenation: The tendency to catenate is very less as compared to Group 14 elements. This is due to very small bond energies of N-N (163.1 kJ/mol), P-P (200.8 kJ/mol) and As-As (146.8 kJ/mol) as compared to C-C (353.9 kJ/mol). Thus, N shows catenation up to three atoms only as in case of N2H4, N3-, etc., and P has considerable extent of catenation. Chemical Properties of Group 15 Elements Group 15 elements differ appreciably from each other in their chemical properties. Nitrogen is almost inert due to its very high dissociation energy (945 kJ/mol). Phosphorus is extremely reactive in its white form while the other elements are comparatively less reactive. The comparative account of their chemical properties can be discussed as follows: 1. Thermal Stability of Trihydrides The hydrides of Group 15 elements are of the composition, MH3. The thermal stabilities of these trihydrides decreases down the group. As the size of the central atom increases, the M-H bond length increases and consequently, the M-H bond strength decreases. As a result, the trihydrides become progressively less stable, so that BiH3 is unstable above -450C. This also results in increase in reducing character. Thus, except NH3, other trihydrides are good reducing agents. 2. Melting Points and Boiling Points of Hydrides The melting points of NH3 are comparatively much higher than other trihydrides while an increase in values is seen on moving from PH3 to BiH3. Likewise, the boiling point of NH3 is higher than that of PH3 and after that increases from PH3 to BiH3. This is due to the highly associated nature of NH3 with strong intermolecular H-bonding. However, from PH3 to BiH3, the increase in atomic size results in increase in Van der Waal’s forces of attraction resulting in an increase in melting and boiling points. 3. Basic Character of the Trihydrides The trihydrides of Group 15 elements have one lone pair of electrons on the central atom. Hence, these compounds can act as Lewis bases. Thus, NH3 is a strong Lewis base and forms adducts with Lewis acids. However, down the group as the size of the central atom increases, the tendency to donate electrons decreases and hence the basic character goes on decreasing; that is why PH3 is a weaker base and AsH3, SbH3 and BiH3 are negligibly basic.  4 4. Bond Angles of the Trihydrides The central atom in the trihydrides is sp3 hybridised and pyramidal in shape with one lone pair of electrons. In case of NH3, the bond angle is 107.3o, smaller than the normal tetrahedral angle. This is due to greater lone pair bond pair repulsions as compared to bond pair-bond pair repulsions, which results in decrease of the bond angle. However, as the size of the central atom increases and the electronegativity decreases, the electron pairs shift away from the central atom and bond pair- bond pair repulsions decrease. As a result, the bond angles decrease from NH3 to SbH3. Structure of NH3 molecule 5. Hydrolysis of Trihalides The trichlorides of Group 15 elements are pyramidal in shape and are predominantly covalent. However, the halides of bismuith and SbF3 are partly ionic. NF3 is not hydrolysed, but trichlorides of these elements are readily hydrolysed. This is due to the reason that N and F cannot expand their octet while Cl and other Group 15 elements can expand their octets. NCl3 hydrolyses to liberate ammonia and hypoclorous acid is formed. NCl3 + 3H2O → NH3 + 3HClO PCl3 and AsCl3 are hydrolysed irreversibly and completely to form hydrochloric acid and the oxide or the oxoacid as shown below: PCl3 + 3H2O → 3HCl + H3PO3 2AsCl3 + 3H2O → 6HCl + As2O3 On the other hand, SbCl3 and BiCl3 are hydrolysed reversibly and partly to form hydrochloric acid and the oxochlorides. SbCl3 + H2O →SbOCl + 2HCl BiCl3 + H2O →BiOCl + 2HCl 6. Formation of Pentachlorides Nitrogen and bismuth do not form pentahalides, while P, As and Sb form pentahalides. This is due to the reason that N cannot expand its octet due to absence of the vacant d-orbitals while P, As and Sb undergo sp3d hybridisation due to the presence of vacant d-orbitals. In case of Bi, +5 state is comparatively less stable than the +3 state. Thus, pentavalent compounds of Bi are unstable.  5 Analogous Behaviour of Nitrogen Nitrogen differs significantly from the rest of the Group 15 elements due to the following characteristics: (a) small size (b) high electronegativity, (c) absence of vacant d-orbitals, and (d) tendency to form multiple bonds. Some of the important points are given here: 1. Nitrogen exists as a gas whereas the other elements exist as solids in allotropic forms. 2. Nitrogen is diatomic and involves two strongly bonded nitrogen atoms by sharing of three electron pairs. :N≡N: Other elements are tetratomic (P4, As4 and Sb4) and bismuth possesses a layered structure in its metallic form. 3. N shows intermolecular H-bonding in its hydrogen compounds while the other elements do not show H-bonding due to their low electronegativities. 4. Nitrogen forms a number of oxides with a large number of oxidation states varying from (+5 to -3). These oxides are monomeric whereas the other elements do not form such a large number of oxides and their oxides are dimeric. 5. The trihalides of nitrogen (except NF3) are unstable and highly explosive whereas the trihalides of other elements are fairly stable. 6. Hydrides of N are stable while of other elements are not so stable. 7. Nitrogen cannot expand its octet and hence does not form complexes, while the other elements form complexes due to the presence of vacant d-orbitals.  6 Occurrence of Group 15 Elements Naturally occurring N2 makes up 78% (by volume) of the Earth’s atmosphere. It also occurs in the combined state as ammonium and nitrates. It is an essential constituent of all animal and vegetable proteins. It is found in fertile soil in the form of nitrates. Phosphorus is an essential constituent of plant and animal tissue. Calcium phosphate occurs in bones and teeth, and phosphate esters of nucleotides (e.g. DNA,) are of immense biological significance. Phosphorus occurs naturally in the form of apatites, Ca5X(PO4)3, the important minerals being fluorapatite (X=F), chlorapatite (X=Cl) and hydroxyapatite (X=OH). Major deposits of the apatite-containing ore phosphate rock occur in North Africa, North America, Asia and the Middle East. Although arsenic occurs in the elemental form, commercial sources of the element are mispickel (arsenopyrite, FeAsS), realgar (As4S4) and orpiment (As2S3). Native antimony is rare and the only commercial ore is stibnite (Sb2S3). Bismuth occurs as the element, and as the ores bismuthinite (Bi2S3) and bismite (Bi2O3). Extraction of Group 15 Elements Nitrogen can be prepared in the laboratory by the following methods: Pure N2 can be obtained by heating sodium nitride. 2NaN3 2Na + 3N2 Nitrogen can be manufacture on a commercial scale by fractional evaporation of liquid air containing mostly liquid nitrogen and liquid oxygen. Nitrogen being more volatile (b.pt. 77.2 K) is obtained first. Mining of phosphate rock takes place on a vast scale (in 2008, 161 Mt was mined worldwide), with the majority destined for the production of fertilizers and animal feed supplements. Elemental phosphorus is extracted from phosphate rock (which approximates in composition to Ca3(PO4)2) by heating with sand and coke in an electric furnace ; phosphorus vapour distils out and is condensed under water to yield white phosphorus. 2Ca3(PO4)2 + 6SiO2 + 10C P4 + 6CaSiO3 + 10CO The principal source of As is FeAsS, and the element is extracted by heating and condensing the As sublimate. Another method is air-oxidation of arsenic sulfide ores to give As2O3, which is then reduced by C. As2O3, is also recovered on a large scale from flue dusts in Cu and Pb smelters. FeAsS In absence of air FeS + As  7 Antimony is obtained from stibnite by reduction using scrap iron or by conversion to Sb2O3 followed by reduction with C. Sb2S3 + 3Fe 3FeS + 2Sb The extraction of Bi from its sulfide or oxide ores involves reduction with carbon (via the oxide when the ore is Bi2S3), but the metal is also obtained as a by-product of Pb, Cu, Sn, Ag and Au refining processes. Chemistry of Nitrogen Physical Properties 1. Nitrogen is a colourless, tasteless and odourless gas. 2. It is slightly soluble in water, 100 volumes of water can dissolve only 2.3 volumes of the gas at 273 K. 3. It can be liquified to give a colourless liquid boiling at 77.2 K and can be solidified to give a colourless solid melting at 62.5 K by suddenly releasing the pressure. 4. It is neither poisonous nor a supporter of combustion. Chemical Properties Nitrogen is an inert gas due to very strong N≡N bond with shorter bond length (1.095 Å) and high heat of dissociation so that there is no appreciable dissociation even at 3000oC. N2(g) 2N(g) ∆H = +945 kJ. As result, free nitrogen is highly unreactive. Some of the important reactions of nitrogen are as follows: 1. Nitrogen reacts with hydrogen under high pressure and in the presence of finely divided iron and molybdenum to give NH3 (Haber’s process). 2. It reacts with oxygen under the influence of lightning discharge or in the presence of electric arc via an endothermic reaction. N2 + O2 → 2NO ∆H = +180.7 kJ 3. It combines with calcium carbide at about 1000oC to give calcium cyanamide, an important fertiliser. CaC2 + N2 → CaNCN + C 4. It readily combines with highly electropositive metals (alkali metals) to give the nitrides containing N3- ion. 6Li + N2 → 2Li3N  8 The nitrides are ionic, crystalline compounds with high melting points. This reaction takes place even at room temperature. 5. It also combines with less electropositive metals to give nitrides. The nitrides of Group 2 elements are ionic and are formed at red heat. The nitrides of groups 13 and 14 are covalent and are formed at white heat. The nitrides of transition metals (Fe, Mn, Mo, W, etc.) are true interestial compounds with nitrogen occupying the interstices of the metal lattice. Active Nitrogen Active nitrogen is the highly reactive and unstable form of nitrogen obtained by passing a condensed induction discharge through nitrogen at a very low pressure. It produces a brilliant luminiscence, which persists for some time often after the discharge is stopped. After some time the normal form is gradually re-obtained. The process is catalaysed by the trace amount of oxygen, carbon dioxide or carbon monoxide. The intensity of the glow increases on cooling by liquid air but decreases on heating. Active nitrogen shows many reactions that are not shown by the ordinary nitrogen. It readily reacts with a number of metals (Na, Hg, Zn, Cd) and nonmetals (S, P, I). However, it does not react with hydrogen or oxygen. It converts nitric oxide to nitrogen sesquioxide and molecular nitrogen. N + 3NO → N2O3 + N2 It also decomposes many organic compounds to form cyanogen. C2H2 + 2N → C2N2 + H2 2CHCl3 + 2N → C2N2 + 2Cl2 + 2HCl However, the nature of active nitrogen has not been fully understood. Uses of Nitrogen and its Compound Gaseous N2 is widely used to provide inert atmospheres, both industrially (e.g. in the electronics industry during the production of transistors, etc.) and in laboratories. Liquid N2 (b.p 77 K) is an important coolant with applications in some freezing processes. Nitrogen-based chemicals are extremely important, and include nitrogenous fertilizers, nitric acid and nitrate salts, explosives such as nitroglycerine and trinitrotoluene (TNT), nitrite salts (e.g. in the curing of meat where they prevent discoloration by inhibiting oxidation of blood), cyanides and azides (e.g. in motor vehicle airbags where decomposition produces N2 to inflate the airbag).  9 Compounds of Nitrogen 1. Nitrogen forms NH3 (ammonia), N2H4 (hydrazine) and HN3 (hydrazoic acid). Ammonia is produced in huge quantities, and it is by far the most common and important compound of nitrogen and hydrogen. Approximately 30 billion pounds of NH3 are used annually with a large portion being used as fertilizer or in the production of nitric acid. Ammonia is produced by the Haber process, which can be shown as N2 + 3H2 300 atm, 450oC, cat. 2NH3 Although the reaction proceeds faster at high temperature, NH3 has a heat of formation of -46 kJ/mol so it becomes less stable. The decomposition of organic compounds that contain nitrogen during the heating of coal to produce coke also produces ammonia. For the preparation of small amounts of ammonia, a convenient reaction involves heating an ammonium salt with a strong base. (NH4)2SO4 + 2NaOH → Na2SO4 + 2H2O + 2NH3 Ammonia is a colorless gas (m.p. -77.8°C and b.p. -33.35°C) with a characteristic odor. Because of the polarity of the N-H bonds, there is extensive hydrogen bonding in the liquid and solid states. Although it is often convenient to use the formula NH4OH, it does not represent a stable molecule. Ammonia is extremely soluble in water, but it ionizes only slightly, NH3 + H2O → NH4+ + OH- Hydrazine (H2N-NH2) is another hydride of nitrogen (in -2 O.S) obtained by boiling a conc. aqueous solution of ammonia and sodium hypochlorite. 2NH4 + NaOCl → NH2NH2 + NaCl + H2O It is a colourless liquid boiling at 386 K soluble in water and alcohol. It is decomposed on heating to give a mixture of nitrogen and ammonia. 3NH2NH2 → N2 + 4NH3 It is oxidised on heating in air or by treatment in O3, H2O2 or halogens. NH2NH2 + O2 → N2 + 2H2O ΔH = -622 kJ 3NH2NH2 + 5O3 → 2N3H +5H2O + 5O2 It also acts as a powerful reducing agent to reduce salts of gold, copper, silver and platinum to their metallic states. NH2NH2 + 4Au3+ → 4Au + 12H+ + 3N2 Hydrogen azide (or hydrazoic acid), HN3, is a volatile compound (m.p. -80°C, b.p. 37°C), and it is a weak acid. It is dangerously explosive (it is 98% nitrogen), and it is also highly toxic. In general, covalent azides or those that are substantially covalent are also explosive. Azides such as Pb(N3)2 and AgN3 are also sensitive to shock, so they have been used as primary explosives (detonators). In contrast, ionic azides are stable and decompose only slowly upon heating.  10 2. Oxides of Nitrogen The oxides of nitrogen are listed on the table below: Formula Name Colour Remarks N2 O Nitrous oxide Colourless Rather unreactive NO Nitric oxide Colourless Moderately reactive N2O3 Dinitrogen trioxide Dark blue Extensively dissociated as gas NO2 Nitrogen dioxide Brown Rather reactive N2O4 Dinitrogen tetroxide Colourless Extensively dissociated to NO2 as gas and partly as liquid N2O5 Dinitrogen Colourless Unstable as gas; ionic solid pentoxide NO3; N2O6 - - Not well characterized and quite unstable Dinitrogen Oxide (Nitrous Oxide) This oxide is obtained in the laboratory scale by heating NH4NO3 in the melt at 250-260oC when water is lost; NH4NO3 2H2O + N2O NO are formed as contaminant, which can be removed by passage through FeSO4 solution. Although N2O has, in the past, been regarded as an inert molecule, this is no longer strictly true, although N2O does not react with halogens, alkali metals, and ozone at room temperature. It is kinetically though not thermodynamically (free energy of formation, ∆GO = +104.18 kJ/mol), inert in the absence of a transition metal center. It has been known for some time as a potent oxygen transfer reagent. N2 O N2 + O Nitric Oxide This is formed in many reactions involving reduction of nitric acid and solutions of nitrates and nitrites. For example, with 8M nitric acid: 8HNO3 + 2Cu 2Cu(NO3)2 + 2NO + 4H2O Reasonably pure NO is obtained by the aqueous reactions 2NaNO2 + 2NaI + 4H2SO4 I2 + 4NaHSO4 + 2H2O + 2NO 2NaNO2 + 2FeSO4 + 3H2SO4 Fe2(SO4)3 + 2NaHSO4 + 2H2O + 2NO Commercially it is obtained by catalytic oxidation of ammonia  11 Nitric oxide is thermodynamically unstable at 25oC and 1atm, under pressure readily decomposes in the range 30 to 50oC, and reacts chemically as NO2 or N2O3: 3NO N2O + NO2 Over Cu2+-ZSM-5 zeolite disproportionation to N2 and O2 occurs The NO molecule has the electron configuration The unpaired π antibonding electron renders the molecule paramagnetic and partly cancels the effect of the π-bonding electrons. Dinitrogen Trioxide The molecule N2O3 is obtained by interaction of stoichiometric amounts of NO and O2 or NO and N2O4. It exists pure only in the pale blue solid (mp ca. -1000C) because of the ready dissociation to give NO and NO2. Nitrogen Dioxide and Dinitrogen Tetroxide These two oxides exist in a strongly temperature-dependent equilibrium: N2 O4 2NO2 K=1.4 x 10-5 mol/dm3 (303 K) The dissociation energy of N2O4 in the gas phase is 57 kJ/mol. In the solid, the oxide is entirely the colorless diamagnetic molecule N2O4. Partial dissociation occurs in the liquid; it is pale yellow at the freezing point and contains 0.01% of NO2, which increases to 0.1% in the deep red-brown liquid at the boiling point, 21.150C. In the vapor at 100oC the composition is NO2 90%, N2O4 10%, and dissociation is complete above 140oC. The monomer NO2 has an unpaired electron and its properties, red-brown color, and ready dimerization to colorless and diamagnetic N2O4, are not unexpected for such a radical. Nitrogen dioxide can also lose its odd electron fairly readily to give NO+, the nitronium ion. Nitrogen Pentoxide It is a strongly acidic oxide of nitrogen in +5 O.S. It is considered an anhydride of nitric acid as it is prepared by the dehydration of HNO3 with phosphorus pentoxide in presence of ozonised oxygen (to avoid decomposition of N2O5). It can also be prepared by the action of Cl2 on dry AgNO3 or by passing ozone through liquid N2O4. It is a colourless solid melting at 303 K, which sublimes reality on heating. It explodes on heating and undergoes decomposition. 2N2O5 → 2N2O4 + O2 Thus, it acts as a powerful oxidising agent. It dissolves in water with a hissing sound and gives nitrates with alkalis. 3. Oxoacids and anion of Nitrogen  12 Nitrogen forms a number of oxoacids such as hyponitrous acid (H2N2O2), nitroxylic acid (H4N2O4), nitrous acid (HNO2), nitric acid (HNO3), peroxonitrous acid (HOONO) and peroxonitric acid (HNO4) Hyponitrate ion: The sodium salt, Na2N2O2, is made by Na reduction of NO in 1,2- dimethoxyethane containing benzophenone or quantitatively by absorption of N2O into Na2O. It is soluble in water and recrystallizable from ethanol. Trioxodinitrate Ion: The interaction of NH2OH and an alkyl nitrate in methanol containing NaOMe at 0oC gives the salt Na2N2O3, which is stable but decomposes in neutral or alkaline media. Nitrous Acid (HNO2) Nitrous acid exists only in the solution form and is very unstable. It is obtained either by dissolving nitrogen trioxide in water or by addition of calculated amounts of cold sulphuric acid to the well-cooled solution of barium nitrite. The precipitates of barium sulphate are filtered off. HNO2 is very unstable and undergoes auto-oxidation on standing: 3HNO2 → HNO3 + 2NO + H2O 2NO + O2 → 2NO2 The decomposition takes place rapidly on boiling to liberate brown fumes in air and nitric acid is obtained. However, in the vapor phase, the following equilibrium state is attained. 2HNO2 → NO + NO2 + H2O Nitrous acid can act both as an oxidising agent as well as a reducing agent. Nitric Acid (HNO3) Formerly known as aqua fortis (strong water), HNO3 is a highly corrosive acid. It is prepared in the laboratory by heating sodium nitrate or potassium nitrate, in presence of conc. sulphuric acid in a glass retort. NaNO3 + H2SO4 → NaHSO4 + HNO3 The vapours of nitric acid are condensed in a receiver, cooled under cold water to give a brown liquid, which is redistilled to remove the dissolved oxides. 4. Nitrides: At high temperature, metals will react in an atmosphere of nitrogen to produce metal nitrides. An example of this behavior is 3Mg + N2 → Mg3N2 Depending on the metallic element, the nitrides may designated as ionic, covalent, or interstitial, as is the case with hydrides, carbides, and borides. 5. Nitrogen Halides Nitrogen halides having the general formula NX3 are known, but unlike the case of phosphorus the pentavalent compounds NX5, are not stable. Some mixed halides such  13 as NF2Cl have been studied. Nitrogen also forms N2F4 and N2F2 that are fluorine analogs of hydrazine and diimine, respectively. Oxyhalides 6. A number of gaseous oxyhalides of nitrogen are known, including the types XNO (nitrosonium or nitrosyl halides) with X=F, Cl, or Br, and XNO2 ( nitryl halides) with X=F or Cl. Nitrosonium halides are prepared by the reactions of halogens and NO. 2NO + X2 → 2 XNO Nitryl chloride can be prepared by the reaction ClSO3H + HNO3 → ClNO2 + H2SO4 Chemistry of Phosphorus Several allotropic forms of phosphorus are known, the most common of which are the white, red, and black forms. 1. White Phosphorus The phosphorus obtained as above is white or yellow phosphorus, called so because the freshly obtained white, transluscent waxy solid acquires a pale yellow colour on standing (formation of surface films of red variety). It is highly poisonous. It can be easily cut with a knife and can be melted at 317 K but only under water due to its low ignition temperature (308 K). Hence, it is always kept under water. It can be boiled at 553.5 K in the absence of air. It is insoluble in water but is readily soluble in turpentine oil, ether and CS2. Its solution in CS2 on evaporation gives octahedral crystals of phosphorus. Its molecular formula is known to be P4, i.e. the four phosphorus atoms occupy the corners of a regular tetrahedron (Fig. 14.16) and are linked to three other phosphorus atoms by covalent bonds. Structure of P4 It glows in the dark with slow oxidation. The main product formed is phosphorus trioxide. This phenomenon is known as phosphorescence. Due to its low ignition temperature, it ignites spontaneously in air to give white fumes of phosphorus pentaoxide. P4 + 5O2 → P4O10 During oxidation, it gives a garlic smell. It reacts with boiling caustic soda, in an inert atmosphere, to liberate phosphine and sodium hydrophosphite is formed (disproportionation). It directly  14 combines with a number of metals such as Na, Mg, Fe, etc., and halogens to form phosphites and halides respectively. P4 + 12Na → 4Na3P P4 + 6Mg → 2Mg3P2 P4 + 6Cl2 → 4PCl3 It combines explosively with sulphur to form a number of sulphides, P2S3, P2S5, P4S3 and P4S7. It forms explosive mixtures with oxidixing agents such as potassium chlorate or potassium nitrate. It is a powerful reducing agent due to its rapid oxidation. Thus, it reduces nitric acid to nitrogen dioxide and sulphuric acid to sulphur dioxide. P4 + 20HNO3 → 4H3PO4 + 20NO2 + 2H2O It also reduces some metallic salts (salt of Cu, Ag and Au) to their corresponding metals. P4 + 8CuSO4 + 14H2O →8Cu + 8H2SO4 + 3H3PO3 + 2H3PO4 White phosphorus is highly poisonous. The persons coming in regular contact with white phosphorus get infected with a disease known as phossy jaw. In this disease, the jaw bones of the infected persons decay and prolonged contact can cause death. 2. Red Phosphorus It is prepared by heating white phosphorus in a cast-iron egg-shaped vessel at 253oC for about eight days. The air present inside the vessel is replaced by coal gas or carbon dioxide and a trace of iodine is added to accelerate the reaction. The product is boiled with caustic soda so as to eliminate any white phosphorus left unreacted in the vessel (red form does not react with NaOH). It is a dark violet-red powdery substance with minute crystalline structure. It is not poisonous and is odourless unlike white phosphorus. It is not soluble in water, caustic soda and Cl2. It is the most stable form of phosphorus and chemically much less reactive. Its ignition temperature is quite high (533 K). Thus, it does not catch fire at room temperature but on heating above its ignition temp. burns to give phosphorus pentoxide. It can be sublimed by heating at 565 K in the absence of air and on boiling, it is converted to the white form. It combines with metals, halogens and sulphur only on heating. Its structure is considered to be polymeric consisting of chains of P4 tetrahedra linked together Structure of red phosphorus  15 3. Black Phosphorus It is obtained by heating white phosphorus at 474 K under 4000 atm. The reaction can be catalysed by mercury. It is the crystalline form of phosphorus and consists of corrugated sheets formed by covalently bonded phosphorus atom linked to three neighbouring phosphorus atoms with P-P distance equal to 28 pm. The adjacent layers are more apart from each other (at a distance of 368 pm) and give flaky crystals. It melts at 587oC and does not burn in air even at 400oC. Structure of black phosphorus Uses of Phosphorus By far the most important application of phosphorus is in phosphate fertilizers. Bone ash (calcium phosphate) is used in the manufacture of bone china. Most white phosphorus is converted to H3PO4, or to compounds such as P4O10, P4S10, PCl3 and POCl3. Phosphoric acid is industrially very important and is used on a large scale in the production of fertilizers, detergents and food additives. It is responsible for the sharp taste of many soft drinks, and is used to remove oxide and scale from the surfaces of iron and steel. Phosphorus trichloride is also manufactured on a large scale. It is a precursor to many organophosphorus compounds, including nerve agents flame retardants and insecticides. Phosphorus is important in steel manufacture and phosphor bronzes. Red phosphorus is used in safety matches and in the generation of smoke (e.g. fireworks, smoke bombs). Chemistry of Arsenic Arsenic also exhibits allotropy and exists in three allotropic forms, namely grey arsenic, yellow arsenic and black arsenic. However, the stable form of arsenic is the gray or metallic form although other forms are known. Cooling the vapor rapidly produces yellow arsenic, and an orthorhombic form (black) is obtained if the vapor is condensed in the presence of mercury. Grey Arsenic: It is brittle due to layered structure of hexagonal rhombohedral crystals in which each atom is bonded to three other atoms. It can be melt by heating at about 773 K under pressure but sublimes at 790 K without melting. It is a good conductor of heat but a fair conductor of electricity. The vapour density studies indicate its formula as As2 at 1975 K and As4 at 917 K.  16 Uses of Arsenic Arsenic salts and arsines are extremely toxic, and uses of arsenic compounds in weedkillers, sheep- and cattle-dips, and poisons against vermin are less widespread than was once the case Arsenic is a doping agent in semiconductors and GaAs has widespread uses in solid state devices and semiconductors. Other uses of As include those in alloys (e.g. it increases the strength of Pb) and in batteries. Chemistry of Antimony Antimony exists in three allotropic forms, namely metallic form, yellow form and black form. 1. Mettalic Form: It is the most common form of antimony having a silver-white lustre. It is very brittle with a layered structure. It melts at 903.7 K in presence of CO and boils at 1653 K. It is a poor conductor of heat and electricity. 2. Yellow Form or α-antimony: It is obtained by passing air or ozonized oxygen on liquid stibine at 183 K. 4SbH3 + 3O2 →4Sb + 6H2O It is yellowish in colour and converts to the metallic form slowly. It is soluble in carbon disulphide. 3. Black or β-antimony: It is the amorphous and explosive form of antimony, which is obtained by the slow electrolysis of antimony trichloride in hydrochloric acid. It is considered an intermediate form of α and metallic antimony. Antimony is quite inert and is not attacked by air and water at ordinary temperatures but is slowly oxidised in moist air. It burns with a blue flame when heated in air. 4Sb + 3O2 → Sb4O6 4Sb + 4O2 → Sb4O8 It is affected only by concentrated acids as dilute acids have no action on its surface. 3Sb + 6H2SO4 → Sb2(SO4)3 + 3SO2 + 6H2O Sb + 5HNO3 → H2SbO4 + 5NO2 + H2O Uses of Antimony Antimony compounds are less toxic, but large doses result in liver damage. Potassium antimony tartrate was used medicinally as an emetic and expectorant but has now been replaced by less toxic reagents. Sb2O3 is used in paints, adhesives and plastics, and as a flame retardant. In addition, Sb2O3 is used in photoelectric devices and electrophotographic recording materials.  17 Chemistry of Bismuth It exists only in one form with a layered structure (m.pt. 544 K and b.pt. 1837 K). It is a white crystalline metal with a reddish tinge and metallic lusture. The molten metal cools to give a solid mass of crystals with slight expansion. It is a fair conductor of heat and electricity and is strongly diamagnetic. It is not attacked by dry air or cold air-free water but it burns in air and decomposes steam at red heat. 4Bi + 3O2 → 2Bi2O3 2Bi + 3H2O → Bi2O3 + 3H2 It is readily dissolved in dilute and conc. acids as well in aquaregia. Uses of Bismuth Bismuth is one of the less toxic heavy metals and compounds, such as the subcarbonate (BiO)2CO3, find use in stomach remedies including treatments for peptic ulcers. COMPOUNDS OF GROUP 15 ELEMENTS 1. Hydride The elements in-group VA (15) of the periodic table form binary compounds with hydrogen, some of which are analogous to the hydrogen compounds of nitrogen (NH3, N2H4, and HN3). However, the hydrides of the heavier elements are much less basic and less stable than NH3. In nitrogen compounds, the unshared pair of electrons function as a hard base, and they are good proton acceptors. In PH3, PR3, AsH3, AsR3, and similar compounds, the unshared pair of electrons resides in a larger orbital, so they are soft bases that have lower basicity toward proton donors. Accordingly, PH3 is a much weaker Brønsted base than is NH3. Formation of stable phosphonium salts requires that the acid be strong and the anion be large so there is a close match in the size of anion and cation. These conditions are met in the reaction with HI. PH3 + HI → PH4I However, with soft electron pair acceptors such as Pt2+, Ag+ , and Ir+ , phosphines are stronger Lewis bases than are NH 3 and amines, so phosphines and arsines interact better with class B metals than do amines. Generally, phosphines and arsines form stable complexes with second- and third-row transition metals in low oxidation states. Phosphine is a less stable compound than NH3 because orbital overlap between hydrogen and phosphorus is less effective than between hydrogen and nitrogen. The thermal stability of the hydrogen compounds of P, As, Sb, and Bi (named as phosphine, arsine, stibine, and bismuthine)  18 decreases in that order, and SbH3 and BiH3 are unstable at room temperature. Similar trends apply to the hydrogen compounds of the group IVA, VIA, and VIIA elements. The hydrazine analogs P2H4, As2H4, and Sb2H4 are toxic and unstable, as indicated by the fact that P2H4 is spontaneously flammable in air, and phosphine burns readily. 4PH3 + 8O2 → P4O10 + 6H2O The extremely toxic trihydrides of the heavier atoms in-group VA are not generally prepared by reaction of the elements. They are usually prepared by making a metal compound of the group VA element and then hydrolyzing it. For example, 6Ca + P4 → 2Ca3P2 Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3 Phosphine also results from the reaction of phosphorus with a hot, strongly basic solution. P4 + 3NaOH + 3H2O → PH3 + 3NaH2PO2 Sodium arsenide can be prepared by the reaction of the elements, and it reacts with water to produce arsine according to the following equation: Na3As + 3H2O → 3NaOH + AsH3 The hydrides of arsenic, antimony, and bismuth are unstable at elevated temperature. The Marsh test for arsenic depends on this instability when an arsenic mirror forms as arsine is passed through a heated tube: 2AsH3 → 2As + 3H2 2. Oxides Phosphorus P4 can be oxidized to yield phosphorus (III) oxide. P4 + 3O2 → P4O6 A tetrahedral arrangement of phosphorus atoms is retained in the P4O6 molecule. This oxide (m.p. 23.9°C, b.p. 175.4°C) reacts with cold water to give phosphorous acid, H3PO3, in which the arrangement of atoms is HP(O)(OH)2. Disproportionation occurs when the oxide reacts with hot water, and phosphine, phosphorus, and phosphoric acid are produced. When heated above its boiling point, P4O6 decomposes into phosphorous and complex oxides that are described by the general formula PnO2n. The oxides As4O6 and Sb4O6, which have structures like that of P4O6, are obtained by burning the elements in air. The +5 oxides of As and Sb can be obtained, but bismuth forms only Bi2O3. Solid As4O6 and Sb4O6 exist in several forms having oxygen atoms in bridging positions.  19 Although P4O10 is an important material, the +5 oxides of arsenic, antimony, and bismuth are much less important. The molecular formula for the +5 oxide of phosphorus is P4O10 molecule, not P2O5, although it is frequently convenient to use the empirical formula. Phosphorus(V) oxide or tetraphosphorus decoxide, P4O10, is the anhydride of the series of phosphoric acids. It is produced in the first step of the manufacture of H3PO4 by burning phosphorus, P4 + 5O2 → P4O10 Arsenic(V) oxide and antimony(V) oxide are produced when the elements react with concentrated nitric acid, 4As + 20HNO3 → As4O10 + 20NO2 + 10 H2O Some compounds of antimony appear to involve the +4 oxidation state. For example, an oxide of antimony is known that has the formula Sb2O4, but this oxide actually contains equal numbers of Sb(III) and Sb(V) atoms. 3. Oxoacids When considering the acids of the group VA elements, the first acid that comes to mind is probably phosphoric acid, H3PO4, which is as it should be. Phosphoric acid is one of the chemicals produced in enormous quantities, and it is used in many industrial processes. However, other acids contain the group VA elements, even though none is very important in comparison to phosphoric acid. In many ways, arsenic acid, H3AsO4, is similar to phosphoric acid, but the similarities are all but nonexistent when bismuth is considered. The acid that bismuth forms is H3BiO3, which can also be written as Bi(OH)3 , showing that it is a very weak acid. Phosphorus(III) is also contained in an acid, H3PO3, but the structure of the molecule is OP(H)(OH)2. Phosphorous acid, H3PO3, is produced when P4O6 reacts with water. P4O6 + 6H2O → 4H3PO3 It is also produced when PCl3 is hydrolyzed. PCl3 + 3H2O → H3PO3 + 3HCl 4. Halides Both the +3 and +5 halogen compounds of the group VA elements contain reactive nonmetal- halogen bonds. As a result, they can be used as starting materials for preparing many other compounds. Halogen compounds of the group VA elements having the formula E2X4 are also known. All the trihalides of phosphorus, arsenic, antimony are known. while bismuth don’t usually formed trihalide with bromine and iodine.  20 Fertilizer Production Feeding a world population of over 6 billion requires the use of tools of all types. Not only is the Machinery of agriculture important, but so are the chemicals of agriculture. The use of effective fertilizers is essential to increase food production from a tillable land mass that is shrinking. Phosphates are an important constituent in many types of fertilizers, and their production involves primarily inorganic chemistry. Naturally occurring calcium phosphate is the primary resource utilized in producing fertilizer. It is found in many places, and it is available in almost unlimited quantity. Converting this insoluble material into a soluble form involves changing the phosphate to some other compound that has greater solubility. In Ca3(PO4)2, the ions have charges of +2 and +3, so the lattice energy is high, and such compounds are generally not soluble in water. Because the phosphate ion is basic, Ca3(PO4)2 will react with an acid to produce H2PO4- (as the process is carried out). With the charges being lower, the compound is more soluble than calcium phosphate. For this conversion, the strong acid having the lowest cost is utilized, and that acid is sulfuric acid. The reaction can be shown as Ca3(PO4)2 + 2H2SO4 + 4H2O → Ca(H2PO4)2 + 2CaSO4.2H2O The product, Ca(H2PO4)2, is more soluble than the phosphate. Sulfuric acid is produced in the largest quantity of any compound, with production that approaches 100 billion pounds annually. Approximately two-thirds of this amount is used in the production of fertilizers. The mixture containing calcium dihydrogen phosphate and calcium sulfate (gypsum) is known as superphosphate of lime, and it contains a higher percent of phosphorus than does calcium phosphate. Plants require nutrients that also contain nitrogen. Ammonium nitrate is an important fertilizer, and it can be produced by the reaction HNO3 + NH3 → NH4NO3 An enormous quantity of ammonium nitrate is produced annually primarily for use as fertilizer and also as an explosive. Ammonium sulfate, ammonium phosphate, and urea are also used as nitrogen containing fertilizers. They are produced by the reactions 2NH3 + H2SO4 → (NH4)2SO4 2NH3 + H3PO4 → (NH4)3PO4 2NH3 + CO2 → (NH2)2CO + H2O A fertilizer that provides both phosphorus and nitrogen is prepared by the reaction of calcium phosphate and nitric acid.  21 Ca3(PO4)2 + 4HNO3 → Ca(H2PO4)2 + 2Ca(NO3)2 Inorganic chemistry is vital to production of food. The materials described are produced on an enormous scale, and they are vital to our way of life and well-being. It is safe to assume that given a projected world population of perhaps 12 billion by the year 2030, there will not be a decrease in this importance.