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SULFURIC ACID 08.07.2020.Pdf SULFURIC ACID Sulfuric acid (alternative spelling sulphuric acid), also known as oil of vitriol, is a mineral acid composed of the elements sulfur, oxygen and hydrogen, with molecular formula H2SO4. It is a colorless, odorless, and viscous liquid that is soluble in water and is synthesized in reactions that are highly exothermic. Its corrosiveness can be mainly ascribed to its strong acidic nature, and, if at a high concentration, its dehydrating and oxidizing properties. It is also hygroscopic, readily absorbing water vapor from the air. Upon contact, sulfuric acid can cause severe chemical burns and even secondary thermal burns; it is very dangerous even at lower concentrations. Formula: H2SO4 Molar mass: 98.079 g/mol IUPAC ID: Sulfuric acid Density: 1.83 g/cm³ Boiling point: 337 °C Melting point: 10 °C Polarity and conductivity Anhydrous H2SO4 is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis. + − 2 H2SO4 ⇌ H3SO 4 + HSO 4 The equilibrium constant for the autoprotolysis is + − −4 Kap (25 °C) = [H3SO 4][HSO 4] = 2.7×10 −14 10 The comparable equilibrium constant for water, Kw is 10 , a factor of 10 (10 billion) smaller. + − In spite of the viscosity of the acid, the effective conductivities of the H3SO 4 and HSO 4 ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions. Chemical Properties:- Reaction with water and dehydrating property- Because the hydration reaction of sulfuric acid is highly exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid.[17] Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent. This reaction is best thought of as the formation of hydronium ions: + − 6 H2SO4 + H2O → H3O + HSO 4 Ka1 = 2.4×10 (strong acid) − + 2− −2 [18] HSO 4 + H2O → H3O + SO 4 Ka2 = 1.0×10 − 2− HSO 4 is the bisulfate anion and SO 4 is the sulfate anion. Ka1 and Ka2 are the acid dissociation constants. Because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent. Concentrated sulfuric acid has a very powerful dehydrating property, removing water (H2O) from other chemical compounds including sugar and other carbohydrates and producing carbon, heat, and steam. In the laboratory, this is often demonstrated by mixing table sugar (sucrose) into sulfuric acid. The sugar changes from white to dark brown and then to black as carbon is formed. A rigid column of black, porous carbon will emerge as well. The carbon will smell strongly of caramel due to the heat generated. Similarly, mixing starch into concentrated sulfuric acid will give elemental carbon and water as absorbed by the sulfuric acid (which becomes slightly diluted). The effect of this can be seen when concentrated sulfuric acid is spilled on paper which is composed of cellulose; the cellulose reacts to give a burnt appearance, the carbon appears much as soot would in a fire. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric. The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystal is changed into white powder as water is removed. Acid-base properties- As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, the blue copper salt copper(II) sulfate, commonly used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid: CuO (s) + H2SO4 (aq) → CuSO4 (aq) + H2O (l) Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH3COOH, and forms sodium bisulfate: H2SO4 + CH3COONa → NaHSO4 + CH3COOH Similarly, reacting sulfuric acid with potassium nitrate can be used to produce nitric acid and a precipitate of potassium bisulfate. When combined with nitric acid, sulfuric acid acts both as an + acid and a dehydrating agent, forming the nitronium ion NO 2, which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols. When allowed to react with superacids, sulfuric acid can act as a base and be protonated, + + forming the [H3SO4] ion. Salt of [H3SO4] have been prepared using the following reaction in liquid HF: + − ((CH3)3SiO)2SO2 + 3 HF + SbF5 → [H3SO4] [SbF6] + 2 (CH3)3SiF The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply HF/SbF5, however, have met with failure, + as pure sulfuric acid undergoes self-ionization to give [H3O] ions, which prevents the + conversion of H2SO4 to [H3SO4] by the HF/SbF5 system: + − 2 H2SO4 ⇌ [H3O] + [HS2O7] Reactions with metals and strong oxidizing property- Dilute sulfuric acid reacts with metals via a single displacement reaction as with other typical acids, producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series) such as iron, aluminium, zinc, manganese, magnesium, and nickel. Fe (s) + H2SO4 (aq) → H2 (g) + FeSO4 (aq) However, concentrated sulfuric acid is a strong oxidizing agent[6] and does not react with metals 2− in the same way as other typical acids. Sulfur dioxide, water and SO 4 ions are evolved instead of the hydrogen and salts. − 2− 2 H2SO4 + 2 e → SO2 + 2 H2O + SO 4 It can oxidize non-active metals such as tin and copper, depending upon the temperature. 2− 2+ Cu + 2 H2SO4 → SO2 + 2 H2O + SO 4 + Cu Lead and tungsten, however, are resistant to sulfuric acid. Reactions with non-metals- Hot concentrated sulfuric acid oxidizes non-metals such as carbon[21] (as bituminous coal) and sulfur. C + 2 H2SO4 → CO2 + 2 SO2 + 2 H2O S + 2 H2SO4 → 3 SO2 + 2 H2O Reaction with sodium chloride- It reacts with sodium chloride, and gives hydrogen chloride gas and sodium bisulfate: NaCl + H2SO4 → NaHSO4 + HCl Occurrence- Pure sulfuric acid is not encountered naturally on Earth in anhydrous form, due to its great affinity for water. Dilute sulfuric acid is a constituent of acid rain, which is formed by atmospheric oxidation of sulfur dioxide in the presence of water – i.e., oxidation of sulfurous acid. When sulfur-containing fuels such as coal or oil are burned, sulfur dioxide is the main byproduct (besides the chief products carbon oxides and water). Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called acid mine drainage (AMD) or acid rock drainage (ARD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic streams. The oxidation of pyrite (iron sulfide) by molecular oxygen produces iron(II), or Fe2+ : 2+ 2− + 2 FeS2 (s) + 7 O2 + 2 H2O → 2 Fe + 4 SO 4 + 4 H The Fe2+ can be further oxidized to Fe3+: 2+ + 3+ 4 Fe + O2 + 4 H → 4 Fe + 2 H2O The Fe3+ produced can be precipitated as the hydroxide or hydrous iron oxide: 3+ + Fe + 3 H2O → Fe(OH)3↓ + 3 H The iron(III) ion ("ferric iron") can also oxidize pyrite: 3+ 2+ 2− + FeS2(s) + 14 Fe + 8 H2O → 15 Fe + 2 SO 4 + 16 H When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process. ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid- neutralization reaction with the minerals. Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales) concentrates sulfuric acid in cell vacuoles. Stratospheric aerosol:- In the stratosphere, the atmosphere's second layer that is generally between 10 and 50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical: • SO2 + HO → HSO3 HSO3 + O2 → SO3 + HO2 SO3 + H2O → H2SO4 Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer. Contact process- In the first step, sulfur is burned to produce sulfur dioxide. S (s) + O2 (g) → SO2 (g) This is then oxidized to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst. This reaction is reversible and the formation of the sulfur trioxide is exothermic. 2 SO2 (g) + O2 (g) ⇌ 2 SO3 (g) (in presence of V2O5) The sulfur trioxide is absorbed into 97–98% H2SO4 to form oleum (H2S2O7), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid. H2SO4 (l) + SO3 (g)→ H2S2O7 (l) H2S2O7 (l) + H2O (l) → 2 H2SO4 (l) Directly dissolving SO3 in water is not practical due to the highly exothermic nature of the reaction between sulfur trioxide and water.
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