Sulfur | Properties, Allotropes, Compounds & Uses

Sulfur

Sulfur, also spelled sulphur, is a non-metallic element that has an atomic number of 16 and a relative atomic mass of 32. It is located in group six or sixteen and period three of the periodic table. Sulfur has the electronic configuration of 1s2, 2s2, 2p6, 3s2, 3p4 and it is found in the p-block of the periodic table. Like oxygen, which is the first element in group 6, sulfur has six valance electrons in its outermost shell and requires two electrons to complete its outermost shell.

Some elements can only be found naturally in their combined state but sulfur can be found in both its elemental and combined form. For example, elemental sulfur can be found in volcanic regions around the world, in fact, in countries like Japan and Indonesia, sulfur is mined from volcanic deposits. It is also known that sulfur constitutes about 0.1% of the earth’s crust. In its combined states, sulfur occurs naturally in minerals deposits like pyrites which contain iron sulfide and gypsum salts.

Sulfur also occurs in organic compounds present in vegetables and animals. Onions, garlic, hairs, oil, eggs, mustard, and protein consist of compounds containing sulfur in them. It also occurs as a constituent of petroleum and coal. In its compounds states, sulfur is present in the form of minerals like sulfide and sulfate.

SulfideFormula
GalenaPbS
ZnSZinc blendeZnS
CinnabarHgS
Iron pyriteFeS2
StibniteSb2S3
Copper pyriteCu2S.Fe2S3
Sulfide ores of sulfur
SulfateFormula
GypsumCaSO4
Heavy sparBaSO4
Sulfite ores of sulfur

Physical Properties

  • Sulphur is almost tasteless and odorless.
  • It is the poor conductor of heat and electricity.
  • Atomic Number: 16
  • Atomic Mass: 32
  • Phase: Solid
  • Density: 2.067
  • Melting Point: 113℃, 239.65℉, 392.2K
  • Boiling Point: 445℃, 832.3℉, 717.82K
  • Ion Charges: S2-
  • Atomic Radius: 88
  • Covalent Radius: 102
  • Ionization Potential: 10.36
  • Electronegativity: 2.58
  • Electron Affinity: 200 KJ/mol
  • Oxidation States: 0, -1, -2, +1, +2, +3, +4, +5, +6
  • Year Discovered: ancient
  • Discovered by: Antoine Lavoisier

Allotropic forms of Sulfur

Let me ask you a question, what is allotropy? Allotropy is a phenomenon where a substance exists in two or more different forms like carbon and sulfur also exist in several allotropic forms. These allotropic forms can be divided into two types:

Crystalline Allotropes:

  1. Rhombic sulfur or alpha sulfur
  2. Monoclinic sulpfur or beta sulfur

Amorphous Allotropes:

  1. Plastic sulfur or gamma sulfur
  2. Colloidal sulfur
  3. Milk of sulfur

Rhombic Sulfur

This is the most stable form of sulfur. Atoms in rhombic sulfur are arranged as an eight-member ring. The rings are packed together rigidly to give the crystals of rhombic sulfur. It is obtained by slowly evaporating the solution of roll sulfur in CS2 when octahedral crystals of rhombic sulfur are formed. This is lemon yellow.

  • Rhombic sulfur is also called alpha-sulfur or octahydral sulfur.
  • It is crystalline allotrope.
  • It’s density is 2.06 g/mL.
  • It’s melting point is 112.8℃.
  • It is stable at room temperature.
  • It is pale yellow in color.
  • It exist in nature.
  • It is insoluble in water but soluble in organic solvent such as petrol, benzene etc.
  • It is non-polar.
  • It is bad conductor of heat and electricity.
  • Its formula is S8.

Structure:

Rhombic sulfur consists of eight atoms of sulfur in its molecule. Each S8 molecule has puckered rings. In each puckered ring four sulfur atoms lie in another plan. Each sulfur is linked with the other by a single covalent bond. S-S bond angle is 105 degrees and the S-S bond length is 2.12 A. These puckered rings unite with one another and form a crystal of rhombic sulfur. These rings are held together by Vander Waal’s forces of attraction. Each unit cell consists of 16 puckered rings.

Rhombic and Monoclinic Sulfur

Monoclinic Sulfur

Monoclinic sulfur also consists of eight-member rinds but the rings are loosely packed. Rhombic sulfur can be converted into monoclinic by heating it above 96 (369K) so these two forms can be interchanged. This temperature at which both forms exist in equilibrium is called transition temperature. Monoclinic sulfur is stable above 369K. It has a dull yellow color.

  • Monoclinic sulfur is also called beta-sulfur or prismatic sulfur.
  • It is crystalline allotrope.
  • It’s density is 1.98 g/mL.
  • Its meltiing point is 119℃.
  • It is unstable at room temperature.
  • It’s dark yellow in color.
  • It is insoluble in water but soluble in organic solvent.
  • It is non-polar.
  • It’s bad conductor of heat and electricity.
  • Its formula is S8 and it forms needle like crystals.
  • It’s formed by dissolving alpha-sulfur in CS2 and heating it above 96℃.

Structure

Like rhombic sulfur, monoclinic sulfur will also exist in S8 molecules in puckered rings. The only difference between rhombic and monoclinic sulfur is that the shape of the monoclinic crystal is needle-like and the unit cell of monoclinic sulfur consists of 6 rings.

Plastic Sulfur

Plastic sulfur has long opened coil or zigzag chain. When molten sulfur is heated up to boiling (400) and pouring this hot sulfur into cold water. It turns into elastic rubber-like material which is called plastic sulfur.

  • Platic sulfur is also called gamma-sulfur.
  • It is amorphous allotrope of sulfur.
  • It’s density is 1.92 g/mL.
  • The melting point of plastic sulfus is not sharp.
  • It’s stable at room temperature.
  • It’s broen in color.
  • It is insoluble in water but soluble in organic solvents.
  • It’s a non-polar.
  • It’s bad conductor of heat and electricity.

Structure

It consists of the long chain of S atoms coiled to each other. It is a polymer, so an infinite number of sulfur atoms are arranged in a random manner. Its elasticity is due to the uncoiling of long chains by the release of tension.

Plastic Sulphur

Colloidal Sulfur

Colloidal sulfur is obtained when hydrogen sulfide reacts with nitric acid or by treating sodium thiosulfate solution with dilute HCl.

H2S + 2HNO3 → 2H2O + 2NO2 + S

Na2S3O3 + 2HCl → 2NaCl + SO2 + H2O + S

It may also be obtained by creating hydrogen sulfide with sulfur dioxide.

2H2S + SO2 → 2H2O + 3S

Milk of Sulfur

When flowers of sulfur are boiled with lime water, a mixture of calcium thiosulfate (Na2S3O3) and calcium pentasulphide (CaS5) is obtained. After that, the mixture obtained is treated with hydrochloric acid gives white amorphous precipitates called milk sulfur. Milk sulfur is mostly used in medicines and is soluble in carbon disulfide.

2Ca(OH)2 + 12S → Ca2S3O3 + 2CaS5 + 3H2O

Ca2S3O3 + 2CaS5 + 8HCl → 12S + 4CaCl2 + 4H2O

Engel’s Sulfur

Engel’s sulfur is also called E-sulfur or cyclo-S6. It is highly unstable and contains an S6 ring. At a very high temperature of above 1000 K, it exists as an S2 molecule.

Comparision of Oxygen and Sulfur

Similarities

  1. Both oxygen and sulfur have the same electronic configuration. (6 valance electrons)
  2. Both oxygen and sulfur are found in free combined states on earth.
  3. Both oxygen and sulfur are typical non-metals.
  4. Both oxygen and sulfur are usually divalent.
  5. Both exhibit allotropic forms.
  6. Both oxygen and sulfur combine with metals in the form of O2- and S2- with oxidation state -2.
  7. Both combined with non-metals and form covalent compounds e.g. H2O, H2S, etc.
  8. Both have polyatomic molecules. Sulfur has S2 and S8 while oxygen has diatomic O2.

Dissmililarities

OxygenSulfur
It is gas at ordinary temperature.It is solid at ordinary temperature.
Oxygen is sparingly soluble in water.Sulfur is insoluble in water.
Oxygen is paramagnetic in nature.Sulfur is diamagnetic in nature.
It shows -2 oxidation state.It shows oxidation states of -2, +2, +4, +6.
There are two allotropic forms of sulfur.There are three allotropic forms of sulfur.
Oxygen helps in combustion.Sulfur is itself combustible.
It does not react with water.When the stream is passed through boiling sulfur a little hydrogen sulfide and SO2 are formed.
It does not react with alkalies.It reacts with alkalies solution and forms sulfides and thiosulfate.
It does not react with acids.It is readily oxidized by conc. H2SO4 or HNO3.

Oxides of Sulfur

1. Sulfur Dioxide, SO2

Sulfur dioxide, SO2 is a colorless gas having a pungent and suffocating smell. It can be easily liquified and highly soluble in water. When it is dissolved in water, the solution so obtained is called susphurous acid. Thus, sulfur dioxide is an acidic oxide and is called susphurous acid anhydride. An aqueous solution of sulfur dioxide is a strong reducing agent.

2H2O + SO2 → H2SO4 + 2H

Thus, SO2 gas reduces halogens to their respective halogen acids, iodates to I2 and Fe2(SO4)3 to FeSO4, etc. Sulfur dioxide also acts as an oxidizing agent and hence oxidizes H2S to S, Mg to MgO, CO to CO2, Hg2Cl2 to HgCl2, Fe to FeO, H2 to H2O, SnCl2 to SnCl4, etc.

Preparation of Sulfur Dioxide

  • Lab Method:

In the laboratory, Sulfur dioxide is prepared by heating copper turnings with conc. Sulfuric acid.

Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O

It can also be prepared in the laboratory by treating sulfite with dilute sulfuric acid.

Na2SO3 + H2SO4 → Na2SO4 + SO2 + H2O

  • By heating Iron pyrite:

SO2 is also be obtained in the laboratory as a by-product of the roasting of sulfide ore such as iron pyrite.

4FeS2 + 11O2 → 2Fe2O3 + 8SO2

  • By Burning Sulfur in Air:

S + O2 → SO2

S8 + 😯2 → 8SO2

It may also be prepared by the reaction:

2ZnS + 3O2 → 2ZnO + 2SO2

Chemical Properties of SO2

1. Acidic Nature

Sulfur dioxide is an acidic oxide and dissolves in water to form sulfurous acid, H2SO3. It can also react with base, NaOH to form a salt.

SO2 + H2O → H2SO3

2. Decomposition

SO2 decomposes into SO3 to S at 1200.

3SO2 → 2SO3 + S

3. Oxidizing and Reducing Property

Sulfur dioxide behaves as a reducer as well as an oxidizer. It is because the oxidation state of S in SO2 is +4 which is lesser than the highest oxidation state +6 and higher than the lowest oxidation state -2.

Oxidizing Nature:

2H2S + SO2 → 2H2O + 3S

3Mg + SO2 → 2MgO + MgS

3Fe + SO2 → 2FeO + FeS

Reducing Nature:

Sulfur dioxide acts as a good reducing agent in the presence of moisture due to the evolution of nascent hydrogen.

SO2 + 2H2O → H2SO4 + 2[H]

SO2 reduces halogen to halogen acids.

Cl2 + SO2 + 2H2O → 2HCl + H2SO4

SO2 decolourizes the violet (pink) color of acidified KMnO4 solution.

2KMnO4 + 5SO2 + 2H2O → K2SO4 + 2MnSO4 + 2H2SO4

SO2 turns orange K2Cr2O7 solution into green.

K2Cr2O7 + H2SO4 + 3SO2 → K2SO4 + Cr2(SO4)3 + H2O

Bleaching Action of SO2

Sulfur dioxide acts as a bleaching agent due to the reduction by nascent hydrogen liberated by moisture present.

SO2 + 2H2O → H2SO4 + 2[H]

Colured matter + [H] → Colourless matter

The bleaching action of sulfur dioxide is temporary due to reduction. If we oxidize the colorless matter, the original color can be regained.

Hybradization of SO2

Sulfur dioxide, SO2 is a covalent molecule. In the solid-state as well as in gaseous, it exists as a discrete molecule. It has an angular shape that results in the form of sp2 hybridization of central atom S. On the basis of sp2 hybridization, the structure and shape of the SO2 molecule have given below:

Hybradization of SO2

Uses of Sulfur Dioxide

  1. Sulfur dioxide is used as bleaching agent.
  2. SO2 is used for household fumigation.
  3. SO2 is used for the refining and bleaching for sugar.
  4. It is used in the manufacturing of sulfuric acid or sulfites.
  5. In the liquid form, it is used for purifing petrolium and as a solvent for sulfur, iodine and phosphorous.
  6. As an antichlor for removing excess of chlorine from bleached fabrics.
  7. It is used in the refining of cane sugar in sugar industry.
  8. It is used as refrigerant.

2. Sulfur Trioxide, SO3

Sulfur trioxide exists in liquid form at room temperature which strongly fumes in the air. SO3 exists in three varieties which are alpha-SO3, beta-SO3 and, gamma-SO3. It oxidizes P to P2O5, S to SO2, and KI to I2. It combines with water to form sulfuric acid. It is used as a drying agent for gases and as a solvent for the manufacture of H2SO4.

Preparation of Sulfur Trioxide

  • By passing a mixture of SO2 to O2 over heated platinum or V2O5.

2SO2 + O2 → 2SO3

  • By dehydration of sulfuric acid with P2O5.

H2SO4 + P2O5 → SO3 + 2HPO3

Properties of SO3

  • SO3 is an anhydrous product of sulfuric acid.
  • It is an acidic oxide and dissolves in water.
  • Acidic nature:

CaO + SO3 → CaSO4

  • Reaction with water:

SO3 + H2O → H2SO4

  • Effect of heat:

2SO3 → 2SO2 + O2

  • Oxidizing property: It can oxidize hydrogen bromide into bromine.

SO3 + 2HBr → H2O + Br2 + SO2

Hybradization of SO3

SO3 molecule is a covalent molecule. In a gaseous state, it exists as a discrete molecule. It has trigonal planner geometry which results in form sp2 hybridization of central atom S. On the basis of sp2 hybridization, the structure and shape of the SO3 molecule have given below:

Hybradization of SO3

Resonating Structure of SO3

Resonating structure of SO3

Oxy Acids of Sulfur

Sulfur forms a huge number of oxy-acids. These acids are classified into five groups:

Sulphoxylic acid:

The chemical formula of sulphoxylic acid is H2SO2, which is not known in the free state.

Sulfurous acids:

Examples of sulfurous acids are:

  • Sulphurous acid, H2SO3 which is not known in the free state.
  • Hyposulphurous acid or hydrosulphurous acid or dithionous acid, H2S2O4 which is not known in free state.
  • Thiosulphurous acid, H2S2O2 which is not known in free state.
  • Disulphurous acid or pyrosulphurous acid, H2S2O5 which is not known in free state.

Sulphuric acids:

Examples of sulphuric acids are:

  • Sulphuric acid, H2SO4
  • Thiosulphuric acid, H2S2O3
  • Pyrosulphuric acid or disulphuric acid, H2S2O7

Thionic acids:

Examples of thionic acids are:

  • Dithionic acid, H2S2O6
  • Polythionic acids, H2SnO6 (where n=3, 4, 5, 6)

Peroxy (peroxo or per) sulphuric acids:

Examples of peroxysulphuric acids are:

  • Peroxomonosulphuric acid or permono sulphuric acid, H2SO5. It is also called Caro’s acid.
  • Peroxodisulphuric acid or perdisulphuric acid, H8S2O8. It is also called Marshall’s acid.

Here, we shall discuss only sulfurous acid and sulfuric acid.

1. Sulphurous Acid, H2SO3

Sulfurous acid is known only in solution. It is formed by dissolving SO2 in H2O.

SO2 + H2O → H2SO3

Sulfurous acid acts as a strong oxidizing agent towards strong reducing agents. It is dibasic acid and ionizes as:

H2SO3 → 2H+ + SO32-

Two structures have been suggested for the acid. One of these is unsymmetrical and has a pyramidal shape while the other is symmetrical and has a tetrahedral shape.

Sulphurous acid

The preparation of sulfurous acid by the action of water on thionyl chloride suggests the symmetrical formula.

Sulfurous acid

H2SO3 readily takes up one oxygen atom to gives sulphuric acid, H2SO4.

Sulphurous acid to sulphuric acid

H2SO3 readily takes up one sulfur atom to gives thiosulphuric acid, H2S2O3.

Sulphurous acid to thiosulphuric acid

The lone pair of electrons on the sulfur atom in H2SO3 can easily be shared by another oxygen or sulfur atom to produces sulphuric acid or thiosulphuric acid.

The unsymmetrical structure is shown above explains the reducing property of H2SO3 due to the S-H bond. Now it is believed that the two forms of the acid exist in equilibrium with each other.

Sulfuric Acid, H2SO4

Sulfuric acid is an oxy acid of sulfur. It is a colorless, oily, dense, corrosive liquid. It is produced by the reaction of sulfur trioxide with water. It is used in accumulators and in the manufacturing of dyes, fertilizers, and explosives.

SO3 + H2O → H2SO4

Discovery of Sulfuric Acid

In the 9th century, Islamic physician and chemist Ibn Zakariya-al-Razi discovered sulfuric acid by dry distillation of alcohol (ethanol) vitriol (al-zajat).

In the 17th century, by burning sulfur with saltpeter (Potassium Nitrate-KNO3), German-Dutch chemist Johann Glauber discovered sulfuric acid.

In the 18th century, Joseph Gay-Lussac, John Glover discovered sulfuric acid with the help of the lead chamber process.

In the 19th century, Peregrine Phillips discovered sulfuric acid by contact process.

Structure of Sulfuric Acid

  1. Geometery: Tetrahedral
  2. Bond angle: 109.39 or 109.5 degree
Sulfuric acid

Different Names of Sulfuric Acid

  1. King of compounds
  2. Battery acid
  3. Mattling acid
  4. Dipping acid
  5. Oil of vitriol
  6. Electrolyte acid
  7. Dihydrogen sulfate

Physical Properties of Sulfuric Acid

  • H2SO4 is soluble in water.
  • It is an oily liquid.
  • It is highly corrosive.
  • H2SO4 is non-volatile acid.
  • It is a dense, viscous colorless liquid.
  • It is a diprotic acid. Diprotic mean sulfuric acid has a tendency to lose two protons.
  • Molar mass: 98 g/mol
  • Apperance: Clear, colorless, corrosive, odorless
  • Boiling point: 337
  • Melting point: 10
  • Density: 1.84 g/cm3 at 20
  • Viscosity: 26.7cp at 20
  • Special form: oleum (when high concentration of SO3 is added)
  • Solubility in water: fully miscible (exothermic process)
  • As a dehydrating agent: Sulfuric acid acts as a catalyst in the following reactions:

C12H22O11 + H2SO4 → C12 + 11H2O

Sulfuric acid acts as a dehydrating agent in laboratories to dry gas mixtures that are being analyzed or prepared.

Chemical Properties of Sulfuric acid

1. Reaction with Water

The hydration energy of sulfuric acid is highly exothermic. The dilution should always be performed by adding the acid to water (not water to acid). It is because, in equilibrium, the reaction favors the rapid protonation of water. The addition of acid to water confirms that the acid is a limiting reagent. This reaction is good for the formation of hydronium ions:

H2SO4 + H2O → H3O+ + HSO4

Because the hydration of sulphuric acid is thermodynamically favorable and its contact with water is enough strong. As we discussed above, Sulphuric acid is an excellent dehydrating agent. Conc. H2SO4 is a very powerful dehydrating property, removing water from other compounds including carbohydrates, sugars and producing carbon, stream, and a more dilute acid containing an increased amount of hydronium and bisulfate ions.

C12H22O11 + H2SO4 → C12 + 11H2O + H2SO4 (unreacted remaining H2SO4)

2. Acid-base Properties

Sulfuric acid reacts with most bases to give the corresponding sulfate.

CuO + H2SO4 → CuSO4 + H2O

Sulphuric acid can also be used to displace weaker acids from their salts.

H2SO4 + CH3COONa → NaHSO4 + CH3COOH

3. Reactions with Metals and Strong Oxidizing Property

Dilute sulphuric acid reacts with metal via a single displacement reaction, producing hydrogen gas and salts.

3Fe(s) + H2SO4(aq) → H2(g) + FeSO4(aq)

In reactant. the oxidation state of iron is 0 but in the product, the oxidation state of iron is +2. So, sulfuric acid is the oxidizing agent.

4. Reaction with Sodium Chloride

H2SO4 reacts with salt, NaCl and gives sodium bisulfate and HCl gas.

NaCl + H2SO4 → NaHSO4 + HCl

5. Electrophilic Aromatic Substitution

Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acid.

Electrophilic Aromatic Substitution

Process for Manufacturing of Sulfuric acid

Sulfuric acid is manufactured by the following three industrial processes:

  1. Contact process
  2. Lead chamber process
  3. Wet sulfuric acid process (WSA)

1. Contact Process

Raw Material

  • Sulfur (source sulfide ore)
  • Water
  • Air

Sources of SO2

  • Sulfur burning
  • Metal sulfate roasting
  • Pyrites roasting
  • Combustion of H2S or other sulfur-containing gas
  • Metal sulfide roasting and smelting

Steps Involve in the Formation of H2SO4

Step 1: Burning of Sulfur

S + O2 → SO2

It is an exothermic process.

Burning of sulfur

Process

  • Firstly, 93% of sulfuric acid is added in air drying tower. The main purpose of adding the sulfuric acid
  • Then, sulfur is sprayed into burner from storage
  • Reaction temperature is 2000F.
  • Exothermic reaction must be cooled
  • At the end, stream is recovered.

Treatment of Burner gas

During the burning of sulfur, other gases are also produced.

  • Sulphur dioxide, burner gas contains impurities like N2, As, F, Cl2, CO2, and dust.
  • It also contains moisture which can cause corrosion to equipment.
  • The burner gas is passed through dust filter chamber, washing Tower, and then drying tower to remove all these impurities.
Treatment of Burner gas

Step 2: Catalytic Oxidation of Sulfur Dioxide

Sulfur dioxide is converted to sulfur trioxide with the help of V2O5 as a catalyst in a converter or contact chamber.

2SO2 + O2 → 2SO2 △H = -197 KJ/mol

SO2 is mixed with air and passed through trays containing loosely packed porous pellets of catalysts.

Catalytic Oxidation of Sulfur Dioxide

Why Heat Exchanger is Used in Contact Chamber?

1) To Maintain Heat:

The conversion of sulfur dioxide into sulfur trioxide is an exothermic process. It releases -197 KJ/mol of energy. To remove this extra heat, heat exchangers or coolers are used on the outlets of the reaction bed. The temperature and the pressure in the converter are maintained between 673K to 773K (400 to 500) and close to 1 atm respectively.

2) Equilibrium Yeild

By using Le-Chatelier’s principle, the concentration of SO3 will be increased.

  1. By Decreasing Temperature:

According to Le-Chatelier’s principle, if the reaction is exothermic, so by decreasing the temperature reaction is moved in the forward direction.

2SO2 + O → 2SO3 △H = -197 KJ/mol

It is an exothermic process. So, if we decrease the temperature reaction will move in a forward direction.

2. By Increasing Pressure:

According to Le-Chatelier’s principle, by increasing pressure, the reaction is moved in that direction where a lesser number of moles are.

2SO2 + O → 2SO3 △H = -197 KJ/mol

As we see in the above particular reaction, 3 and 2 moles are present in the reactant and in the product respectively. So, if we increase the pressure, the reaction moves forward.

3. Effect of Concentration:

If we increase the concentration of SO2 and O2, the reaction is moved in the forward direction.

Summary:

  • By decreasing temperature
  • By increasing pressure
  • By increasing the concentration of SO2 ans O2, the reaction is moved forward and more and more concentration os SO3 is produced.

Step 3: Absorption of SO3

The direct reaction of sulfur trioxide with water is highly exothermic so as a result, steam is produced. H2SO4 as gas is very difficult to collect.

SO3 + H2O → H2SO4 △H = -103 KJ/mol

Due to this sulfur trioxide is absorbed in sulfuric acid to produce oleum. The chemical formula of the oleum is H2S2O7. Oleum is a very concentrated form of sulfuric acid.

SO3 + H2SO4 → H2S2O7

Step 4: Dilution Tower

Oleum is react with water to form sulfuric acid.

H2S2O7 + H2O → 2H2SO4

Dilution tower

Advantages & Disadvantages

1. Contact Process

Advantages

  • Large amount of sulfuric acid is manufactured by contact process.
  • High concentration of sulfuric acid is obtained as compared WSA and lead chamber process.
  • It is widely used process over the world.
Manufacturing of sulfuric acid by contact process
Manufacturing of sulfuric acid by contact process

Disadvantages

  • The main problem is that the catalyst vinadium pentaoxide can be poisoned.

2. Wet Sulfuric Acid Process

Advantages

  • By using this process, there are no or little waste by-products.
  • WSA process is the ecnomic way to get rid of sulfurous waste gases.

Disadvantages

  • Low concentration of sulfuric acid is obtained.

3. Lead Contact Process

Advantages

  • The orignal acid to be used can be obtained at any concentration.

Disadvantages

  • 62% and 68% sulfuric acid is in the chamber.
  • This process is not widely used as contact process. This is because this process produces more dilute acid than contact process. Contact process is always produce much more sulfuric acid than lead chamber process.

Uses of Sulfuric Acid

Sulfuric acid is one of the most important industrial chemicals that used in our daily life. A huge amount of sulfuric acid is used in fertilizers. About 75% sulfuric acid is used in fertilizer, 5% in petroleum refining, 5% in metal production, and 15% sulfuric acid is used in the manufacturing of other chemicals.

Sulfuric acid is used in the manufacturing of fertilizers such as superphosphate, ammonium phosphate, ammonium sulfate, and rock phosphate which contains insoluble calcium phosphate. For plants used it needs to be converted into a soluble form.

Ca3(PO4)2 + 2H2SO4 +4H2O → Ca(H2PO4)2 + 2CaSO4.2H2O

Sulphuric acid plays an important role in the development of any country. Many industries over the world use sulfuric acids such as agricultural chemicals, aluminum sulfate, cellophane, detergents, kerosene, explosives, fertilizers, sugar, gasoline, herbicides, iron, and steel pickling, leather, batteries, lubricating oil, yellow pigments, paper, rayon, and rubber, synthetic fibers, Jet fuel, water softener regeneration, medical processes, water treatment, and oil additives.

Sulfuric acid is also used in daily routines such as:

  1. Chemical manufacturing
  2. Metal processing
  3. oil refining
  4. The manufacturing of rayon
  5. Potato harvesting
  6. The manufacturing of medicines
  7. The manufacturing of lead acid type batteries

Halides of Sulfur

Sulfur forms a large number of halides. Out of these, sulfur hexafluoride, SF6 is the most important halide. SF6 is prepared by allowing the sulfur to burn continuously in fluorine, F2. It is a colorless and odorless gas. The melting and boiling point of SF6 is 50.8 and 68.8 respectively. It is chemically inert like nitrogen and hence remains unaffected by heat. SF6 also remains undecomposed by water, alkalies, or acid. Even ignited lead chromate or copper and fused caustic potash have no effect on it. It is decomposed by molten (boiling) Na at 250.

SF6 + 8Na → Na2S + 6NaF

It is also decomposed by H2S.

SF6 + 3H2S → 6HF + 4S

It is also decomposed by hydrogen at 400 and by sparking with O2 and H2. In non-reactivity and stability, SF6 resembles CCl4 and differs from all other non-metallic halides. SF6 is a covalently saturated molecule and due to this reason, it is inert. The existence of this molecule proves the hexacovalancy of the sulfur-atom. Sulfur hexafluoride, SF6 has a regular octahedral shape.

Hybradization of SF6

Oxyhalides of Sulfur

Sulfur forms a number of oxyhalides but, we shall discuss only thionyl chloride and sulfonyl chloride.

1. Thionyl Chloride, SOCl2

Preparation

The chemical formula of thionyl chloride is SOCl2, It is prepared by adding sulfur into Cl2O at -12.

S + Cl2O → SOCl2

Commercially, it is also prepared by the action of SO3 on S2Cl2 at 75-80.

SO3 + S2Cl2 → SOCl2 + SO2 + S

Thionyl chloride reacts with KBr to form thionyl bromide, SOBr2.

SOCl2 + 2KBr → SOBr2 + 2KCl

It also reacts with SbF3 in the presence of SbCl5 (Swart’s reagent) to form thionyl fluoride, SOF2.

3SOCl2 + 2SbF3 → 3SOF2 + 2SbCl3

Uses of SOCl2

SOCl2 is used as a reagent in organic chemistry to replace OH groups present in the organic compounds by Cl atom.

Structure of SOCl2

Sulfur is sp3 hybridized in the SOCl2 molecule. The molecule has a distorted tetrahedral shape i.e. trigonal pyramidal shape due to the presence of one lone pair on sulfur atom.

Structure of SOCl2

Sulfonyl Chloride, SO2Cl2

Preparation

The chemical formula of sulfonyl chloride is SO2Cl2. Sulfonyl chloride is also called sulfuryl chloride. It is formed by the direct combination of SO2 and Cl2 in sunlight or in the presence of catalysts like charcoal or acetic anhydride and camphor.

SO2 + Cl2 → SO2Cl2

Uses of SO2Cl2

SO2Cl2 solution in petrol is used in making the wool unshrinkable. In inorganic chemistry, it is used as a chlorinating agent.

Struture of SO2Cl2

Sulfur is sp3 hybridized in the SO2Cl2 molecule. The molecule has a tetrahedral shape.

Hybradization of SO2Cl2

Uses of Sulfur

  1. Sulfur is used to produce sulfur dioxide which is a bleaching agent.
  2. It is also used in the preparation of carbon disulfide which is a solvent.
  3. Sulfur is used in the manufacturing of phosphorus trisulfate which is used in the matchbox industry.
  4. Calcium sulfide is used as a bleaching agent in the paper industry.
  5. A mixture of sulfur charcoal and potassium nitrate is called gun powder and it is used in the manufacture of crackers and other fireworks.
  6. Sulfur is used for the hardening of rubber and this process is known as vulcanization.
  7. Sulfur is also used in the manufacture of antiseptic ointments and many other drugs.
  8. It is also used in the manufacturing of fungicides sprayed on fruit trees and grape plants.
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