H₂SO₄ — Sulfuric Acid

Sulfuric acid (H₂SO₄) is one of the most important and widely used industrial chemicals. Known as the king of chemicals, it is a dense, oily, and colorless to slightly yellow liquid that is both strongly acidic and highly corrosive. Sulfuric acid plays a vital role in numerous industries, including fertilizer manufacturing, petroleum refining, wastewater treatment, and chemical synthesis. Its versatile nature allows it to act as a dehydrating agent, oxidizing agent, and strong acid in various reactions.

Because of its extensive use and economic importance, the global production of sulfuric acid is often considered an indicator of a country's industrial strength. It is also an essential component in laboratories and is used in the manufacture of other acids such as hydrochloric, nitric, and phosphoric acid.

The molecular structure of sulfuric acid consists of a central sulfur atom covalently bonded to four oxygen atoms. Two of these oxygen atoms are double-bonded to sulfur, while the other two are bonded via hydroxyl groups (-OH). The molecular geometry of H₂SO₄ is tetrahedral, and the hybridization of sulfur is sp³.

In aqueous solutions, sulfuric acid dissociates in two stages:

The first dissociation is almost complete, making it a strong acid, while the second is partial, classifying sulfuric acid as a diprotic acid. The high polarity of the O–H bonds and extensive hydrogen bonding give sulfuric acid its high boiling point and viscosity.

The industrial production of sulfuric acid is primarily carried out by the Contact Process. This process involves the catalytic oxidation of sulfur dioxide (SO₂) to sulfur trioxide (SO₃), followed by absorption in concentrated sulfuric acid to produce more H₂SO₄. The process can be summarized in the following steps:

1. Burning of Sulfur or Sulfide Ores: Sulfur or metal sulfides like pyrite (FeS₂) are burned in air to produce sulfur dioxide.

2. Oxidation of SO₂ to SO₃: In the presence of a vanadium(V) oxide (V₂O₅) catalyst at 450°C and 2 atm, SO₂ is oxidized to SO₃.

3. Formation of Sulfuric Acid: Sulfur trioxide is absorbed in concentrated H₂SO₄ to form oleum (H₂S₂O₇), which is then diluted with water to produce sulfuric acid.

This method yields highly pure and concentrated sulfuric acid with 98% concentration typically available in industries.

• Highly Hygroscopic: Sulfuric acid readily absorbs water vapor from the air, which makes it an excellent drying agent for gases.
• Dehydrating Agent: It removes water molecules from organic compounds, converting sugars and carbohydrates into carbon.

\(C_{12}H_{22}O_{11} \xrightarrow{H_2SO_4} 12C + 11H_2O\)

• Reaction with Metals: Reacts with active metals to produce hydrogen gas and corresponding metal sulfate.

\(Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2\)

• Oxidizing Agent: Hot concentrated sulfuric acid acts as an oxidizing agent, converting Cu, S, and C into oxides or sulfur dioxide.

\(Cu + 2H_2SO_4 (conc.) \rightarrow CuSO_4 + SO_2 + 2H_2O\)

• Reaction with Bases: Neutralizes alkalis to form sulfates and water.

\(2NaOH + H_2SO_4 \rightarrow Na_2SO_4 + 2H_2O\)

• Fertilizer Industry: Sulfuric acid is used in the manufacture of phosphate fertilizers such as superphosphate and ammonium sulfate.
• Chemical Manufacturing: Used in producing hydrochloric acid, nitric acid, phosphoric acid, and other chemicals.
• Petroleum Refining: Employed in refining crude oil and removing impurities from gasoline and kerosene.
• Battery Acid: The electrolyte in lead-acid batteries is dilute sulfuric acid, which facilitates redox reactions.
• Metallurgy: Used in pickling (cleaning) of metals like iron and steel before galvanization or electroplating.
• Dehydrating Agent: Acts as a powerful desiccant in laboratories and industrial setups.
• Textile and Paper Industry: Used for bleaching fabrics and processing paper pulp.

Concentrated sulfuric acid is extremely corrosive and can cause severe chemical burns on skin contact. Its reaction with water is highly exothermic, which can lead to splattering and explosions if not handled carefully. Always add acid to water, not water to acid, to minimize heat buildup.

Inhalation of fumes can damage respiratory passages, and ingestion is fatal. Protective equipment such as acid-resistant gloves, goggles, and lab coats must be used while handling the compound. Spills should be neutralized with a weak base like sodium bicarbonate before cleanup.

Despite its hazards, sulfuric acid remains indispensable in modern industry and research due to its versatility and chemical reactivity.