Magnesium nitrate, with the chemical formula Mg(NO₃)₂, is an inorganic salt composed of magnesium cations (Mg²⁺) and nitrate anions (NO₃⁻). It is a white, crystalline solid that is highly hygroscopic, meaning it readily absorbs moisture from the air, often forming hydrates such as magnesium nitrate hexahydrate (Mg(NO₃)₂·6H₂O).
This compound is commonly encountered in its hydrated form due to its affinity for water. Magnesium nitrate is widely used in agriculture, industry, and chemical synthesis due to its solubility, stability, and ability to provide magnesium and nitrogen, both essential nutrients for plants. It is non-toxic in moderate amounts but requires careful handling due to its oxidizing properties, which can enhance combustion in certain conditions.
Structure
Magnesium nitrate features an ionic structure typical of metal nitrates. The magnesium ion (Mg²⁺) is a divalent cation with a small ionic radius, forming strong electrostatic bonds with two nitrate anions (NO₃⁻). Each nitrate anion consists of one nitrogen atom covalently bonded to three oxygen atoms in a trigonal planar arrangement, with a negative charge delocalized across the oxygen atoms due to resonance.
In its anhydrous form, magnesium nitrate adopts a crystalline lattice where magnesium ions are coordinated with nitrate groups. In the hexahydrate form (Mg(NO₃)₂·6H₂O), the magnesium ion is coordinated with six water molecules in an octahedral geometry, with nitrate ions positioned outside this coordination sphere, stabilized by hydrogen bonding with the water molecules. This hydrated structure contributes to its high solubility and hygroscopic nature.
Properties
- Physical Properties: Magnesium nitrate appears as colorless to white crystals or granules in its anhydrous form, while the hexahydrate forms colorless, prismatic crystals. It has a molar mass of 148.31 g/mol (anhydrous) or 256.41 g/mol (hexahydrate).
- Solubility: Highly soluble in water (71 g/100 mL at 20°C for the hexahydrate) and soluble in ethanol and ammonia, making it versatile for liquid applications.
- Melting and Boiling Points: The anhydrous form decomposes at high temperatures (around 330°C) rather than melting, while the hexahydrate melts at approximately 89°C.
- Density: Approximately 1.64 g/cm³ for the hexahydrate.
- Hygroscopicity: Strongly hygroscopic, readily absorbing moisture to form hydrates, requiring storage in airtight containers.
- Chemical Properties: Acts as an oxidizing agent due to the nitrate ion, potentially supporting combustion. It is stable under normal conditions but can release toxic nitrogen oxides when heated to decomposition.
- pH: Aqueous solutions are slightly acidic due to partial hydrolysis of the Mg²⁺ ion.
Preparation
Magnesium nitrate is typically prepared through the reaction of magnesium-containing compounds with nitric acid or nitrate sources. Common methods include:
- Reaction with Nitric Acid: Magnesium oxide (MgO), magnesium hydroxide (Mg(OH)₂), or magnesium carbonate (MgCO₃) is reacted with dilute nitric acid (HNO₃). For example: MgO + 2HNO�/CD>₃ → Mg(NO₃)₂ + H₂O or MgCO₃ + 2HNO₃ → Mg(NO₃)₂ + CO₂ + H₂O. The resulting solution is evaporated to crystallize magnesium nitrate, often as the hexahydrate.
- Reaction with Magnesium Metal: Magnesium metal can be dissolved in nitric acid, producing magnesium nitrate and hydrogen gas: Mg + 2HNO₃ → Mg(NO₃)₂ + H₂. This method is less common due to the reactivity of magnesium metal and the need for controlled conditions.
- Industrial Synthesis: Large-scale production involves neutralizing nitric acid with magnesium compounds, followed by purification and crystallization. The hexahydrate form is typically obtained due to its stability in moist environments.
Uses
Magnesium nitrate has diverse applications across multiple fields:
- Agriculture: A key component in fertilizers, providing magnesium and nitrogen, essential for plant growth. It is used in foliar sprays and soil applications to correct magnesium deficiencies, particularly in crops like tomatoes, peppers, and citrus.
- Pyrotechnics: Acts as an oxidizer in fireworks and flares, enhancing combustion and producing bright effects.
- Chemical Synthesis: Serves as a source of magnesium in laboratory reactions and as a catalyst or dehydrating agent in organic synthesis.
- Industrial Applications: Used in the production of magnesium salts, ceramics, and as a component in explosives (e.g., in mining).
- Water Treatment: Occasionally used in wastewater treatment to precipitate phosphates or control odors.
- Medical Research: Employed in biochemical studies and as a magnesium source in experimental settings, though not a primary medical supplement.
Additional Notes:
- Safety: Magnesium nitrate is generally safe but should be handled with care due to its oxidizing nature, which can intensify fires. Avoid inhalation of dust or fumes from decomposition, and store away from flammable materials.
- Environmental Impact: When used in fertilizers, excessive application can lead to nitrate runoff, potentially causing water pollution. Proper dosing is critical.
- Storage: Store in airtight containers in a cool, dry place to prevent moisture absorption and caking.