CuSO₄ — Copper(II) Sulfate

Copper(II) Sulfate (CuSO₄) is an inorganic compound composed of copper, sulfur, and oxygen. It exists in both anhydrous (white) and hydrated (blue) forms, the latter being most common and known as copper(II) sulfate pentahydrate (CuSO₄·5H₂O). This compound has been used for centuries in agriculture, medicine, and chemistry laboratories due to its diverse properties and chemical reactivity.

Known commonly as blue vitriol or blue stone, copper(II) sulfate is notable for its striking blue crystals and ability to form complex ions in solution. It serves as a fungicide, algicide, and root killer, as well as a reagent in analytical chemistry. The compound is also an essential educational material for demonstrating crystallization and chemical reactions involving transition metals.

The chemical formula of copper(II) sulfate is \(CuSO_4\). It contains a Cu²⁺ ion (copper ion with a +2 charge) and an SO₄²⁻ ion (sulfate ion with a −2 charge). The bond between copper and the sulfate ion is largely ionic, although the hydrated form involves coordination between the copper ion and water molecules.

In the pentahydrate form (CuSO₄·5H₂O), each copper ion is surrounded by four water molecules in a square planar geometry and one water molecule hydrogen-bonded to the sulfate ion. This arrangement gives rise to the bright blue color due to d–d electronic transitions in the copper ion’s 3d orbital.

The structure of anhydrous copper(II) sulfate is different, where water molecules are absent, leading to a white powder with a more compact crystal lattice.

Copper(II) Sulfate occurs naturally as the mineral chalcanthite, found in arid copper-mining regions. It can form when copper ores like chalcopyrite and bornite are oxidized in the presence of sulfur and oxygen. In nature, chalcanthite crystals are rare because they are water-soluble and often dissolve in moist environments.

Commercially, copper sulfate is produced synthetically by reacting copper metal or copper(II) oxide with sulfuric acid. The crystalline blue form (CuSO₄·5H₂O) is commonly used in laboratories, while the anhydrous white form serves as a drying agent and industrial catalyst.

1. From Copper Metal and Sulfuric Acid:

Copper reacts with hot concentrated sulfuric acid to produce copper(II) sulfate, sulfur dioxide, and water:

This is a redox reaction where copper is oxidized from 0 to +2, and sulfuric acid acts as an oxidizing agent.

2. From Copper(II) Oxide and Sulfuric Acid:

This method is often used in the laboratory to produce pure copper(II) sulfate solution:

The reaction is exothermic and yields blue crystals of CuSO₄·5H₂O upon evaporation.

3. From Copper Carbonate:

When copper(II) carbonate reacts with dilute sulfuric acid, copper(II) sulfate is formed along with carbon dioxide and water:

Physical Properties:

• Blue crystalline solid in hydrated form (CuSO₄·5H₂O) and white powder in anhydrous form.
• Highly soluble in water, forming a bright blue solution.
• Losess water of crystallization on heating, turning white as it becomes anhydrous.
• Exothermic dissolution in water due to hydration of copper ions.

Chemical Properties:

• Thermal Decomposition: On heating, CuSO₄·5H₂O loses water molecules:

The anhydrous CuSO₄ is white, and when exposed to moisture, it rehydrates and turns blue again.

• Reaction with Water: The anhydrous form acts as a drying agent and readily absorbs water to form the hydrated blue compound:

• Reaction with Alkalis: Forms insoluble copper(II) hydroxide when reacted with sodium hydroxide:

• Reaction with Iron: When iron is placed in a solution of copper sulfate, displacement occurs:

This reaction demonstrates the reactivity series where iron is more reactive than copper.

Copper(II) Sulfate is one of the most versatile and commercially valuable copper compounds used across industries:

• Agriculture: Used as a fungicide and pesticide to control fungal and bacterial growth on crops. It is a key ingredient in Bordeaux mixture (CuSO₄ and Ca(OH)₂) used for protecting grapevines and vegetables.
• Analytical Chemistry: Serves as a reagent in Benedict’s and Fehling’s solutions to test for reducing sugars.
• Electroplating and Metal Refining: Used as an electrolyte in copper plating and purification processes.
• Education: Demonstrates crystallization and chemical reactions in school laboratories.
• Medicine: Used as an astringent, antifungal agent, and antiseptic (though in controlled quantities due to toxicity).
• Industry: Used in dyeing textiles, leather tanning, and as a mordant in printing and dyeing fabrics.
• Environmental Applications: Controls algae growth in water bodies and septic systems.

While copper(II) sulfate has many beneficial applications, it must be handled with care. It is toxic if ingested and can cause irritation to the skin, eyes, and respiratory tract. Prolonged exposure may lead to copper accumulation in the body, causing symptoms of poisoning such as nausea, vomiting, and liver damage.

Environmentally, excessive copper sulfate in water bodies can harm aquatic life by disrupting enzyme systems in fish and microorganisms. Therefore, its agricultural and industrial use is regulated to prevent pollution. Proper disposal and controlled application ensure minimal environmental impact.