Chapter 12: Nitrogen and Sulfur

Nitrogen Introduction • Nitrogen belongs to group 15 of the periodic table. • It's atomic number is 7 amu and atomic mass 14 amu. • It is a non-metal element belongs to p-block of periodic table. Preparation of Nitrogen in laboratory In laboratory settings, nitrogen can be generated by slowly heating a solution of ammonium nitrite. NHANO 2(29) -4 Uses • In industrial processes, nitrogen is typically obtained by cooling air until it becomes a liquid. • Liquid nitrogen is commonly used for rapid cooling purposes. Table 12.1 Physical properties of nitrogen. Atomic number Relative atomic mass 14.007 u Colourless Gas Physical appearance Electronic configuration [He]2s2p? -210 °C Melting Point - 195.8 °C Boiling Point Covalent radius 74 pm Reactivity of Nitrogen (N2) Nitrogen is a significant component of the air, known for its low reactivity due to its small size, symmetrical electronic cloud, and nonpolar triple bond. With an electronic configuration of 1s 2s 2p}; nitrogen requires three electrons to complete its octet, forming a triple bond by sharing three electrons with another nitrogen atom as shown in Figure. This bond has a . High energy is required to break this bond to form new bond enthalpy of +944 kJmol → N2(g) + 2H2O(g) ionic radius (3-) 170 pm 1402 kJ/mol 1st ionization energy 3.0 Electronegativity Electron affinity -7.0 kJ/mol Density 1.25 kg/m Principal oxidation states 3+, 5+ N=N Fig: Bonding in N2 molecule bonds, making N2 very unreactive. The second reason for its lack of reactivity is the non- polarity of its bond. Both the atoms are the same having zero electronegativity difference. This causes equal sharing of the three bonded electrons between the two atoms making the bond nonpolar. Interesting Information! S.Q. What is the role of inertness of nitrogen? Ans. The high concentration of nitrogen in the air serves to dilute oxygen, preventing every spark in our atmosphere from igniting a massive fire. In the case of large shipments of hydrocarbons or edible oils, it is crucial to utilize blankets of Nitrogen or any other inert gas on ships to safeguard them from oxygen and moisture. It is also used in laboratory to carry out the reactions which require inert atmosphere. Ammonia (NH3) Industrial production of Ammonia: Ammonia (NH3) is an important industrial compound of nitrogen, which is mainly used as a fertilizer. It is prepared industrially by Haber-Bosch process. N 2(g) + 31 2(g) = Basicity of Ammonia Ammonia behaves as a Lowry-Bronsted base by accepting a proton (HT) from an acid to form ammonium ion: It dissolves in water to form ammonium hydroxide (NH4OH) and equilibrium is established between ammonia molecules and ammonium ions in the solution. Ammonia solution is a weak base due to the low basicity constant (K) and the equilibrium position being towards the far left side. Structure of Ammonium (NH'*) ion • Ammonia molecule has pyramidal shape due to lone pair of nitrogen. • When nitrogen atom in ammonia utilizes its lone pair of electrons to form ammonium, this ion adopts a tetrahedral shape. • All the bonds are of equal length and strength in ammonium ion as depicted in figure. Synthesis of Ammonia from Ammonium/salts Laboratory preparation: In the laboratory, ammonia gas can be synthesized by heating an ammonium salt such as ammonium chloride (NH4Cl) with a base like calcium hydroxide Ca(OH)2 as shown in figure. 2NH, C1 (s) + Ca (OH) 2(aq) → 2NH 3(g) H \ HI.N H H H H Fig: Pyramidal and tetrahedral shapes of ammonia and ammonium ion → CaC12 (ag) + 2H2O(e) + 2NH3(8) In this acid-base reaction, Acid: NH + acts as an acid by donating H* ions, Base: OH acts as a base by accepting Hions. This reaction displaces from the ammonia gas NH,CI + Ca(OH), ammonium salt and produces salt and water. Uses • It is commonly used to identify ammonium ion salt analysis. • If a gas with a pungent smell is released and turns moistened red litmus paper the it indicates blue, presence of ammonium ion in the compound Quick Check 12.1 (a) Why No gas is used in food packaging? Ans. To protect the packaged food from oxygen and moisture so that the food can be preserved for a longer duration, due to it's inertness and non polarity (b) Both CO and N2 have triple bonds in their molecules. Why do you think CO is more reactive than N2? Ans. Nitrogen has a high bond energy making it less reactive, CO is a slightly polar molecule while N2 is a non-polar molecule Due to polarity CO is more reactive. (c) Ammonium salts such as (NH4)2SO4 or NH4NO3 are commonly used as fertilizers. Why a farmer wouldn't treat a field with an ammonium fertilizer at the same time as using lime? What would be the chemical reactions? Ans. Lime has an alkaline pH which causes ammonium in the fertilizer to become ammonia gas and hence is lost in atmosphere therefore to avoid losses of nitrogen from the fertilizer these are not used simultaneously with lime. The reaction of ammonium sulfate and line. (NH4)≥SO, + Ca(OH≥ CaSO, 2H, 0 + 2NH, OXIDES OF NITROGEN

Oxides of Nitrogen: • Oxides of nitrogen are NO, NO, NO2, N204 and NOs in which oxidation states range from Ito V. • N204 and N2Os decay quickly to other oxides. NO and NO2 are collectively called as NOX. • The structures, properties and uses of these oxides are given in Table 12.2. NH, Air All CaO Fig: Synthesis of ammonia gas Table 12.2 Properties of some common oxides of nitrogen Name and Structure formula of oxide Nitrous oxide oxide Nitrogen gas) (Laughing N iNE Nitic oxide, Dinitrogen dioxide :NE=O NO, N202 tree Nitrogen dioxide/Nitrogen peroxide, Dinitrogen tetraoxide • O •O NO2, N204 N-N // :O: Did you know? S.Q. How the name ammonium is derived? Ans. The recommended name for ammonia by IUPAC is Azane and for ammonium is Azanium. These names have been derived from Azote, a Greek name for nitrogen, meaning "no life". These names are used in naming derivatives of ammonia and ammonium e.g. sodium azide (NaN3). Formal Uses Properties oxidation state Dental Colourless, gas, water-soluble anaesthetic neutral, sweet propellant 1+ for whipped smelling, helps in combustion. ice cream, synthesis of NaN3. Biochemical Colourless gas, messenger slightly water- (Lowers soluble, neutral, NO is blood pressure and paramagnetic role in other 2+ while N202 is body diamagnetic as oxidizing functions), well as synthesis of in nitrosyl reducing nature. carbonyls. iS a Rocket reddish-brown propellant, HNO3 gas, formation by paramagnetic, and reacts with Ostwald 4+ water to form process, and HNO3 explosives. HNO2, N204 is colourless liquid or solid.

Sources of oxides of nitrogen The main categories of NOx sources. (i) Natural Sources Natural sources include lightning, volcanic eruption, forest fires, and denitrifying bacteria in soil. NO is produced when N2 and O2 in the air react during lightning. N2O is produced by microorganism using air N2. (ii) Anthropogenic (Man-made) Sources The main anthropogenic sources of NOx are the combustion of fossil fuels in vehicles and power plants. Other sources include chemical plants, biomass burning, welding etc. Role of NO and NO2 in smog and PAN formation NOx is responsible for numerous harmful éffects on living organisms. Photochemical Smog Formation: Photochemical smog (Los Angeles smog) forms in the atmosphere from NOx and volatile organic compounds (VOC's) in the sunlight. • It is oxidizing in nature. • Photochemical oxidants, such as NO2, ozone, and peroxyacetyl nitrates (PANs) can react and oxidize specific compounds in the atmosphere. • Photochemical smog is becoming more common than classical smog (London smog) due to increasing NOx emissions. Reactions: The formation of photochemical smog involves the following chemical reactions. → 2NO(B) N2(g) +02(B) S.Q. Give composition of Lahore's smog. →→ 2NO 2(g) 2NO (g) + 02(g) Ans. Lahore's smog consists of volatile organic compounds (VOCs), NOx, ground level ozone (O3), (g) + 0 NO2(g) →* > NO particulate matter PM 2.5, CO and SO2. 03(g) O(8) +02(g) RCH = RCH (g) + 03(g) → RCO%) + RCH 2(g) Formation of peroxyacetyl Nitrates (PANs) Starting material: NOx is starting material of PAN formation. Process: NOx take part in a series of reactions leading to the formation of ozone (03), aldehydes, peroxyacetyl nitrates (PANs) and peroxybenzoyl nitrate (PBN). PAN is one of the members of peroxyacetyl nitrates. Did you know? →RC(O)OONO, O + C N - Fig: General structure of PANs Reactive Hydrocarbon ROH, RCH, 0, + HO. NO ple 519 Fig: Mechanism of PAN formation Steps of PAN production • In peroxyacetyl nitrate (PAN) the alkyl (R-) group is methyl (CHz-). Other hydrocarbon chains may also be present. The main component of oxidizing smog is ozone. • Ozone oxidizes hydrocarbon to produce aldehyde. The aldehyde then reacts with hydroxyl radical to produce acyl radical. • The acyl radical reacts with oxygen to produce peroxyacetyl radical. • peroxyacetyl radical finally reacts with nitrogen peroxide to form peroxyacetyl nitrate. Quick Check 12.2 (a) Draw the structures of the following oxides of nitrogen. Also, briefly explain their bonding. Ans. Name Nitrous oxide/ Nitrogen oxide (Laughing gas) N2O a triple bond and a single bond. Nitric oxide, Nitrogen dioxide NO, N2 02 a triple covalent bond in NO. the double bonds between N and O. A single between two N atoms. O, RCH,0. * NO NO YOH RO-O Y NO: R( CONO m Structure : NEN O: Ö :NE - N dioxide/Nitrogen peroxide, Nitrogen Dinitrogen tetraoxide NO2, N2Oz. One single bond between N and O. A double bond between Nand another O. (b) What does PAN stand for? Give its general formula. Ans. PAN stands for peroxy actyle nitrate General formula is RC(O)OONO2 (c) Write down the formulas of the compounds responsible for the formation of PAN. Ans. These formulas of compounds responsible for PAN formation are NOx, (NO, NO2) (d) If magnesium ribbon is ignited and placed in a jar containing NO, it continues to burn brightly, how does the product form in this reaction confirm the structure of N20. • MgO + N2 Ans. Mg + N20 The reaction confirms that in nitrous oxide oxygen is attached to a nitrogen that is bound to another nitrogen and nitrogen bond is more strong than nitrogen oxygen bond as a triple bond. The structure of NO is: :N=N-

Construction • A catalytic converter is a ceramic or metallic monolith with a honeycomb-like structure. • Its inner channels have a layer of alumina to provide a high surface area. • Noble expensive metals such as Pt (Platinum), Pd (Palladium), and Rh (Rhodium) are dispersed on the alumina. Working These metals catalyze three redox reactions to remove the half harmful exhaust gases. The three-way converter converts harmful CO, NO, and hydrocarbons into CO2, N2, and water. These precious metals can also be recycled. Reduction Pt/Rh > N 2(g) + 2C0 2(g) 2NO (B) + 2C0 (g) Oxidation Pt/Pd → 2CO 2(g) 2CO (g) + 02(B) 2С, H4 + 60 2(g) 4C0 2(g) + 4H, 0(8) N N-N :0. :0: washcoat with catalyst inside the H2O, CO,, N, monolith channel HC, CO, NO, Fig: Three-way catalytic converter NITRIFICATION AND DENITRIFICATION Nitrification and denitrification Nitrogen cycle: Nitrification and denitrification are the two phases of the nitrogen cycle. Nitrification: It involves the conversion of ammonium (NH, ) to nitrite (NO,) and nitrate (NO,) Denitrification: It involves the conversion of nitrate (NO, ) to Nitrogen (N2). These two processes are also involved in the wastewater treatment to remove nitrogen. Some differences between nitrification and denitrification are given in the table 12.3. Table 12.3: Differences between nitrification and denitrification Nitrification Definition Ammonia NH3 / NH, is converted into nitrite (NO,) and nitrate (NO,). Conditions Nitrifying bacteria aerobic conditions, pH 6.5 - 8.0, optimum temperature 20 °C- 30°C. Importance Plants absorb these nitrites and nitrates for their nutrition as they cannot assimilate nitrogen directly from the atmosphere. Reactions Oxidation of nitrogen NH: →NO: - → NO: 2NH 4(a9) + 30z (aq) → 2N05 (aq) + 4H (ag) + 2H, O) 2NO2(ag) + 02(ag) → 2NO3(a9) Quick Check 12.3 (a) Write down the reduction and oxidation reactions that occur in the catalytic converter in the vehicle exhausts. Ans. Reduction Pt / Rh → N 2(g) + 2CO 2(g) 2NO (g) + 2CO Oxidation Pt / Pd → 2CO 2 (g) 2C0 (g) + 02(g) Pt / Pd 2C, H4 + 60 2(g) → 4CO 2(g) + 4H, 0(g) Denitrification Definition Nitrite (NO,) and nitrate (NO,) are converted back to N2 that is released into the atmosphere. Conditions Denitrifying bacteria anaerobic conditions, pH 7.0-9.0, optimum temperature 26°C - 38°C Importance It is important in wastewater treatment and useful for aquatic life, to oxidize NH, with NO, to form N2 gas. Reactions Reduction of nitrogen NO: →NO: →NO→N,0→N, 2NO3raq) + 10e + 12H (aq) → N 2(ag) + 6H,0() (b) What is the basic principle of catalytic converter? Describe the role of catalyst in the catalytic converter. Ans. Catalytic converter works on the principle of harmful gases into less toxic gases by redox reaction. The metal catalysts catalyzes these redox reactions to remove the half harmful exhaust gases. The three-way converter converts harmful CO, NO and hydrocarbons into CO2, N2 and water (c) Do hybrid and electric cars have catalytic converters? Explain why or why not Ans. Hybrid cars have catalytic converters to reduce the emission of harmful gases in atmosphere. Electric cars do not need catalytic converters these cars do not emit harmful gases.

Sulfur Sulfur is a member of group 16 which is also called the Chalcogen family. Some physical properties of sulfur are given in table 12.4 Table 12.4 Physical properties of sulfur Atomic number Ionic radius Relative atomic mass 1st ionization energy Physical appearance Tige Electronegativity Electronic configuration Electron affinity Melting point Density Boiling point Common oxidation states Covalent radius Common Bonding Reactivity of sulfur • Sulfur usually forms single bonds with other sulfur atoms instead of double bonds due to poor overlapping of the orbitals. • Sulfur forms larger molecules and structures through a process called catenation. • S& is a crown-like molecule Oxidation states of sulfur • Sulfur exhibits oxidation states of -2, 0, +2, +4, and +6. • The oxidation state is determined by the number of unpaired electrons. 16 184 pm 32.06 a.m.u 1000 kJ/mol Solid Yellow 2.5 [Ne] 3s 3p4 -200 kJ/mol 113 °C (honey yellow) 2.07 g/cm? 445 °C (dark brown) 4+ and 6+ 104 pm Covalent bond S S S S S Fig: Cyclo-octasulfur, (S8) crown molecule • Under standard conditions, sulfur and oxygen react to produce sulfur dioxide (SO2) in which sulphur has an oxidation state of +4. • In formation of sulfur trioxide (SO3), sulfur with an oxidation state of +6. It requires high energy. Electronic configurations of sulfur for attaining different oxidation states is shown in figure. 2s a) Oxygen atom in ground state (2s° 2p*) 3.5 b) Sulfur atom in ground 3p state (3s- 3p*) c) Sulfur atom in lst excited state (35? 3p 3d) d) Sulphur atom in second excited state (35" 3p" 30) Fig: Electronic configurations of sulfur for attaining different oxidation states STABILITY OF OXIDATION STATES OF SULPHUR Sulfur displays a range of oxidation states, principally from -2 to +6. Factor effects stability of oxidation states of sulfur The stability of these states is influenced by the factors like pH, temperature, nature of the compound and the chemical environment. • pH: In general, under acidic conditions, reduced forms (like H2S, - 2 oxidation state) are more stable, while under basic or neutral conditions, oxidized forms (like SO*, +6 oxidation state) become more stable in water. • Thermodynamics and kinetics: Although sulfur (+6), as in SO3, 1s thermodynamically the most stable kinetic limitations can prevent it from forming readily at standard temperatures, making sulfur (+4), as in SO2, the more frequent form. • Nature of the compound: The stability of sulfur's oxidation state can be affected by the elements it is bonded to, and the overall chemical environment. In acidic environment SO, is kinetically stable due to strong O-S bonds. • Catalyst: A catalyst can also enhance the rate of formation of a specific oxidation state, for example vanadium in the contract process increases the rate of formation of SO3. Quick Check 12.4 (a) Determine the oxidation state of S in the following species (i) SOs (i H2SOa (iii) SO:- Ans. (i) SQ3 Oxidation state of sulfur + 3(oxidation state of oxygen) = 0 x +3(-2) =0 x-6=0 X=6 There are no d-orbitals in 2nd shell therefore excitation is not possible. 3d 2+ oxidation state due to two unpaired electrons 4+ oxidation state due to four unpaired electrons 6+ oxidation state due to 1 six unpaired electrons (iv) S,03- (ii) H2SO4 2(oxidation state of hydrogen) + oxidation sate of sulfur + 4 (oxidation state of oxygen) =0 2(+1)+x+4(-2)=0 2+x-8=0 x=8-2 X=6 (iii) (SO?) Oxidation state of sulfur + 4(oxidation state of oxygen) = -2 x+ 4(-2) = -2 X - 8 = -2 x=8-2 X=+6 (iv): (S,05 ) 2(oxidation sate of sulfur) + 3 (oxidation state of oxygen) =-2 2(x)+3(-2)=-2 2x-6=-2 2x=6-2 2x=4 x=2 (b) Which oxidation states of sulfur are the most common? Explain your answer 15J- with reason. +4 and +6 are most common oxidation states of sulfur. In basic or neutral conditions these are the most stable oxidation states. SO2 and SO3 show +4 and +6 respectively. (c) Explain the involvement of d orbital in variable oxidation states of S in its compounds. Ans. The oxidation state is determined by the number of unpaired electrons as shown in figure. Sulfur atom in 1st | I excited state (35? 3p? 3d') Sulphur atom in second excited state (3s 3p3 38)

Reactions of Sulfur Sulfur can combine with many elements to form a wide variety of inorganic and organic compounds. Reactivity of sulfur It is unreactive to water under normal conditions, dilute non-oxidizing acids, and noble gases. Its ability to catenate allows it to form ring structures and linear chains. Some important reactions of sulfur (i) Burning of sulfur Sulfur burns in the air to form SO2 with a blue colour flame. The other main sulphur oxide is SO which requires higher temperature and a catalyst for its formation. S(s) +02(g) SO 2(B) 4+ oxidation state due to four unpaired electrons 6t oxidation state due to six unpaired electrons 22(g) 02(0) - Sulfur can be oxidized by nitric acid to produce SO2 ànd H2SO4 S(s) + 6HNO3(1) →H, SO4(1) + 6NỌ2(g) + 2H, (i (i Sulfur as oxidizing agent When sulfur reacts with elements that have lower electronegativity, it acts as an oxidizing agent and forms their sulfides. It tarnishes Ag, Cu, and Zn by forming a coating of metal sulfide. 2Ag(s) +S (s) → Ag,S(s) Hg() + S(s) HgS(s) 2C4(8) + S(s) → Cu, S(s) Sulfur does not react with Au and Pt (iii) Reaction with cyanate Sulfur converts cyanide into thiocyanate which is also known as pseudohalide Application of reaction It is used to analyze Fe3+ KCN(s) + Ss) >KSCN(8) (iv) Reaction with fluorine Sulfur reacts directly with F2 to form SF4 and SF6. Sulfur hexafluoride (SF6) is a gas and is very unreactive. It is used as an insulator gas in electric devices. Sulfur readily reacts with Ch to form S2 Cl2 (yellow liquid) which further reacts with Cl to form SCl (red liquid). S(s) + 3F2(g) → >SF6(8) 25(s) + C12(g) Cl 2(g) +S, Cl 2(e) → 25Cl 2(e)

Uses of sulfur and its compounds Vulcanization: Sulphur is used as a cross linker for the rubber molecular chains. This is called vulcanization it improves the strength of rubber. Fertilizer Sulfur is an essential nutrient for plant growth. When soils become depleted in sulfate, sulfur can be restored in soil by , N/P fertilizers, or sulfur-coated applying sulfur containing fertilizers such as sulfur-coated urea. Soil components and microbes convert elemental sulfur into soluble forms for the use of plants. Gypsum (CaSO4.5H2O) is also used as a fertilizer. • → 2503(8) Pt Fig: S-S cross-linkages between polymer chains Gun powder Gun powder is a coarse blend of 75% potassium nitrate (KNO3), 15% wood charcoal, and 10% sulfur. Charcoal carbon is the main fuel, nitrate is the oxidizer and sulfur is the additional fuel that burns the powder faster. The following reaction in burning takes place: 10KNO 3(6) + 35(6) + 8C) → > 2K, CO 3(6) + 3K,SO 4(6) + 6C02(g) + 5N2(g) Quick Check 12.5 (a) What is the function of SO2 and sulfite (SO, ) salt in preserving the food? Ans. These compounds of sulfur give antimicrobial and anti oxidative action Antioxidative action prevents browning of food products and prevents oxidation that degrades food quality Anti microbial action prevents food spoilage by inhibit growth of microorganism (b) Give the reaction of S with HCl and NaOH. Ans. Reaction of S with NaOH. 3S+ 6NaOH -2Na,S + Na,S,O, + 3H2O S and Hcl At room temperature there is no reaction between sulfur and dilute hydrogen chloride as sulfur does not allow chlorine to replace it in the compound At higher temperature concentrated hydrochloric acid to sulfur chloride and hydrogen gas. 2HCI+ S - → SCI, + H, Role of Sulfur in organic synthesis Carbon-Sulfur bonds are prevalent in a wide range of compounds with biological, pharmaceutical, and material properties. These bonds form a large number of organic compounds containing a variety of functional groups such as thiols or mercaptans, thioethers, sulfoxides, sultones, etc. Drugs Sulfa drugs are the antibacterial sulfonamides such as penicillins and cephalosporins contain sulfur. The common drug omeprazole, used in GERD (Gastroesophageal reflux disease) contains sulfoxide group Dyes Sulfur dyes are synthesized by the process of thionation or sulfurization of organic compounds that contain nitro or amino groups. These compounds contain sulfur linkages. They generally give black, brown, blue, and green colours. Some examples of sulfur colors are: • Black • Blue • Brilliant green Odorants/Fragrances Mercaptans are used to give odour to natural gas. Some thiols have pleasant odours on high dilution, Example: Thioterpineol is the key. ingredient in the aroma of grapefruit. cis-glabanum oxathiane is a fragrant compound. It is used in fine fragrances, soaps, shampoos and shower gels. Many naturally occurring odorants are produced synthetically and also applied as flavouring agents. SULFURIC ACID

Structure of sulfuric acid It has tetrahedral structure with two S-O and two S=O bonds 142.2 pm S H 97 pm 157.4 pm Fig: Structure of H2SO4 Contact process The major portion of sulfur, around 85% is used for the production of sulphuric acid (H2SO4). Contact process for the industrial production of sulphuric acid. The Contact process can be divided into the following stages, Water Conc. H,SO, spray Impure spray SO, +0, Dry SO, +02 Sulphur Air → Waste Waste Sulphur , acid • water burner Washing and Drying Dust - tower cooling tower precipitator Fig: Contact process for the industrial production of sulfuric acid Sulfur/pyrite Burners The process starts with the combustion of molten sulphur or by heating pyrites such as iron pyrite (FeS) in excess of air to produce sulfur dioxide S02. 2502(g) 25(g) + 02(g) 4FeS(s) + 60 2(g) Conc. H,SO, SOg V,O, quartz Preheater Les Testing box мест Catalytic Oleum converter Arsenic purifier containing gelatinous orated AH =-297.03 kJ/mol → 2Fe, 03(s) + 450 2(g) Purification Unit If pyrite ore is used as a sulfur source, the SOz gas formed may contain contaminants like dust particles, vapors, and arsenic oxide. These contaminants affect the efficiency of the catalyst. Hence, the gas needs to pass through the purification unit. In an arsenic purifier, gelatinous ferric hydroxide Fe(OH) present in horizontal shelves, absorbs arsenic oxide AszO3. As, 03(g) + 2Fe(OH)3(1) - Contact Tower and Heat Exchangers Purified SO2 and air, preheated at 420°C-450°C, are fed to the first converter stage of the contact tower at 1-2 atm pressure. Here, these gases come in contact with vanadium pentoxide (V20s) catalyst. SO 2(g) 2(g) 503(8) AH=-98.98 kJ mol The catalyst works in two steps Oxidation of SO2 (into SO3 by VS+) → SO3(8) + V20 4(5) SO 2(g) + V205(8) Oxidation of V4t back into Vs+ by oxygen (catalyst regeneration) 2 Quick Check 12.6 (a) Why is high temperature and catalyst needed to form SO3? Ans. Catalyst increases the rate of reaction Increasing temperature lowers the yield but increases the rate therefore a temperature compromise is made and yield is increased by adding oxygen. (b) Write down dehydration reactions of conc. H2SO4 with starch and oxalic acid. Ans. (Starch) H2804 >n (6C) + nH,0 starch is a polymer of glucose hence first it is broken down into individual glucose and each glucose is then dehydrated. (Oxalic acid) 124 → CO+CO, + H2O (COOH)2 (c) How does the catalyst VOs function in the conversion of SO to SO3? Ans. The catalyst works in two steps Oxidation of SO2 into SO3 by VS+ → SO 5(s) SO2(g) + 20 Oxidation of Vit back into Vst by oxygen (catalyst regeneration): + V,O 2 → 2FeAsO 3(1) + 3H1, 0() 3(8) + V20 4(s) → V20 s(s) 2(s) Absorption Tower Sulfur trioxide is cooled and can be converted to sulphuric acid by reacting with water H, 0(1) + S03(g) H5O4(1) Mixing SO with water is not feasible because the reaction is extremely exothermic and acidic vapor or mist is produced rather than a liquid solution. Mainly, sulfur trioxide is dissolved in recirculating hot 98.5% sulphuric acid. The term fuming sulphuric acid or oleum is used for the mixtures of sulfur trioxide with 100 percent sulphuric acid. Oleum undergoes a reaction with water to make a highly concentrated solution of H2SO4 whose concentration can be adjusted H,SO 4(1) + SO 3(g) H,S,0 7(1) + H, 0(1) → 2H,SO 4(1)

Physical properties of Sulphuric acid • Solubility Sulphuric acid is soluble in water and hygroscopic in nature. • Absorbs water vapors: It readily absorbs water vapor from the air. • Polarity: Anhydrous H2SO4 is a very polar liquid. Corrosive: It is highly corrosive to various materials. On contacting the skin, it causes chemical burns. Table 12.5 Physical properties of sulfuric acid Molar mass Physical appearance 5- Odour Melting point Boiling point Specific gravity at 15 °C Viscosity vapor pressure Caution Sulphuric acid is a highly corrosive substance. It can badly burn cloths, plastic, rubber and injured the human skin, eyes, etc. Handle it very carefully. Chemical properties (i) Self ionization It self-ionizes or undergoes autoprotolysis as follows. H,SO A() + H,SO 4(1) → H,SO+ The equilibrium constant value is greater than that of water which makes it to be used as a non-aqueous protic solvent. Sulphuric acid is a strong acid as shown by its pKa Value: H, 0(0) + H, SO 4(ag) = H3O (a9) + HSO*(ag) (pK,, = -2) But hydrogen sulfate (HSO' ) is a far weaker acid due to a positive pRaz value: HSO'- AH = -176.6 kJ/mol →H,S, 07(0) 98.08 g/mol colourless viscous liquid Odourless 10 °C 290°C 1.83 g/cm 25.24 centipoise (25°C) 0.001 torr + HSO Ka = 2.7 × 104 4(aq) 4(aq) 4(aq) (pKa2=1.92) (aq) +SO?- (il) Dehydrating agent Concentrated sulphuric acid is a powerful dehydrating agent that removes water from many substances such as sucrose, starch, wood, and paper to produce carbon, steam, and heat. Dehydration of table sugar C2#22011(8) +HSO 4(1) → 12C() + 11H, 0(g) + H,SO 4(34) A black porous carbon mass called carbon snake protrudes out of the apparatus Dehydration of alcohol It also dehydrates ethyl alcohol to ethene or ethoxyethane depending upon the reaction conditions С, н, ОН (e) + H,SO 4(C) - →→ C, 4(g) + H, 0(e) + H,SO 4(ag) Interesting Information! S.Q. How sugar is dehydrated in laboratory? Ans. A common laboratory demonstration is the dehydration of table sugar, where a black porous carbon mass called carbon snake protrudes out of the apparatus. C12H22011(8) + HI,SO A(e) 12C(s) + 11H, 0(g) + H,SO 4(a9) (iii) Reaction with sodium chloride Hydrochloric acid (Hcl) gas, is formed when sulphuric acid reacts with sodium chloride NaC (s) + HI,SO 4(1) → NaHSOAs) + HC (g) (iv) Reaction with metals Reactions of sulphuric acid with metals depend upon the metal, concentration of the acid, and temperature. Metals that are above hydrogen in electrochemical series such as Fe, Al, Zn, Mn, Ni, and Mg react directly with dilute sulphuric acid to produce hydrogen gas and metal sulfates. → H2(g) + Z (24) + SO 3(a4) Zn s) + HI,SO 4(29) But with cold conc. H2SO4, they liberate SO2 and form sulfates. ZM (s) + 2H,SO 4(1) → SO 2(g) + Z (ag) + SO 3(ag) + 2H2O,) Metals like Cu, Ag, and Hg react with hot conc.H2SO4 to form metal sulfates (v) Oxidizing power • Sulphuric acid is not regarded as a typical oxidizing agent due to the stability of SO% anion. This anion is weakly oxidizing. • Hot concentrated sulphuric acid is a moderately strong oxidizing agent due to high temperature, high concentration of protons (H*), and formation of nascent oxygen • Hot concentrated sulphuric acid oxidizes Cu, as given below: CU(s) + 2H50 4(1) - → SO 2(g) + 2H2O,) + SO 3(ag) + Cu(a) Quick Check 12.7 (a) Write down dehydration reactions of conc. H2SO4 with starch and oxalic acid. Ans. Starch n(CH,00s) 6) + nH,SO 40) →п(6C)s) + nH, 0(r) + nH,SO 4(a) starch is a polymer of glucose hence first it is broken down into individual glucose and each glucose is then dehydrated. Oxalic acid (СООН)(8) + ,50 4() → СО(в) + СО 2(g) + H,S0 4(29) + H2O(e) (b) How does conc. HzSO4 react with NaCl? What is the importance of this reaction? Ans. Hydrochloric acid (Hcl) gas, is formed when sulphuric acid reacts with sodium chloride → NaHSO4(s) + HCL(g) Importance It is used to produce hydrogen chloride gas in laboratory (c) What is the difference between the oxidizing power of cold and hot sulphuric acid! Ans. Sulphuric acid is not regarded as a typical oxidizing agent due to the stability of SO, anion. This anion is weakly oxidizing. However hot concentrated sulphuric acid is a moderately strong oxidizing agent due to high temperature, high concentration of protons (HT), and formation of nascent oxygen.

Uses and Industrial Applications Sulfuric acid is considered a king of chemicals and its consumption is an indicator of the industrial progress of a country. • Fertilizers: A major portion of the acid is used in making fertilizers, normally 77.67% 1s used to digest the phosphate rock containing calcium phosphate Ca3(PO4)2. Sulphuric acid also reacts with ammonia to make ammonium sulfate fertilizer • Extraction of metals: It is used in the extraction of metals from ores such as Cu, Ni, steel etc. • As a catalyst: It is utilized as a catalyst in oil and coal refining, polymers, synthetic rubber, and plastic industries. • Industrial usage: It is used in the pulp and paper industry and involved in the production of pesticides, insecticides, herbicides, varnishes, dyes, pharmaceuticals, soaps and detergents. • Explosives: It is used for nitration in making explosives such as trinitrotoluene (TNT), nitroglycerine, picric acid, nitrocellulose, etc • Food industry: It is involved in the food industry for making sugar, starch, and corn syrup. • Paint industry: It is used in the paint industry for making titanium dioxide (TiO2) pigment. • Drying agent: It is used to dry gases in industrial processes. • Batteries: 35.67% acid is used in lead storage batteries. • Reagent: It is used as a laboratory reagent. • Even though it is used in various industries, it is rarely contained in the final product Did you know? S.Q. What are the uses of phosphate? Ans. Calcium phosphate Ca3(PO4)2 and fluorapatite (Cas(PO4)3F) are used to make single superphosphate Ca(H›PO4)2 and phosphoric acid (H3PO4). HзPO4 is further used to make double and triple superphosphates and ammonium phosphate. .