Boric Acid: Structure, Bonding, Preparation, Properties and Uses

Boric Acid (H₃BO₃)

Boric Acid is also known as Hydrogen Borate, Orthoboric acid, Boracic acid, Sassolite, Borofax, Trihydroxyborane, Boranetriol, and Acidum boricum

Boric acid has the molecular formula H₃BO₃ (also written as B(OH)₃).
It is a monobasic Lewis acid, not a protonic (Brønsted) acid.
In solid state, it exists as planar triangular B(OH)₃ units.
These units are linked together by symmetrical hydrogen bonds (O–H···O) forming infinite two-dimensional layers (sheet-like structure).
The layers are held together by weak van der Waals forces.
B–O bond length: ~136 pm; O-H: bond length ~97 pm; O–H···O hydrogen bond length: ~270 pm.
Boron is sp² hybridized with trigonal planar geometry (bond angle ≈ 120°), and C_3h molecular symmetry.

Structure of Boric Acid representing O-B-O Bond Angle, B-O and O-H Bond Length

Structure of Boric Acid representing O-B-O Bond Angle, B-O and O-H Bond Length, and Hydrogen Bonding

Covalent bonding in boric acid involves σ-bonding between boron's sp² hybrid orbitals and oxygen p orbitals, with partial π-character resulting from oxygen lone pair donation into boron's empty p orbital.
Boron has only 6 electrons in valence shell → acts as electron-deficient Lewis acid.
It accepts a lone pair from OH⁻ to form [B(OH)₄]⁻ (tetrahedral).
No B–H bonds (unlike boranes); only B–O and O–H bonds.
The B-O bond energy ~536 kJ/mol, significantly higher than typical B-O single bonds due to the partial double bond character.
Extensive intermolecular hydrogen bonding gives it a layered, flaky crystalline appearance.

Method	Reaction	Conditions
From Borax (Commercial)	Na₂B₄O₇·10H₂O + 2HCl → 4H₃BO₃ + 2NaCl + 5H₂O or Na₂B₄O₇ + H₂SO₄ + 5H₂O → 4H₃BO₃ + Na₂SO₄	Concentrated acid, heating, then cooling to crystallize
From Colemanite	Ca₂B₆O₁₁ + 2H₂SO₄ + 11H₂O → 6H₃BO₃ + 2CaSO₄·2H₂O	SO₂ is passed to precipitate CaSO₄
Hydrolysis of Boron Halides	BCl₃ + 3H₂O → H₃BO₃ + 3HCl	Lab method

Appearance: White, shiny flaky crystals (pearly luster).
Solubility: Sparingly soluble in cold water (~5 g/100 mL at 20°C), highly soluble in hot water (~40 g/100 mL at 100°C).
Taste: Slightly bitter.
Melting point: 170°C (decomposes on strong heating >100°C).
Volatility: Volatile in steam (steam distillation property).

Acidic nature (Lewis acid)
B(OH)₃ + H₂O ⇌ [B(OH)₄]⁻ + H⁺
With ethyl alcohol (Borate ester formation – flame test)
H₃BO₃ + 3C₂H₅OH → B(OC₂H₅)₃ + 3H₂O
Triethyl borate burns with green-edged flame.
Dehydration on heating
H₃BO₃ → HBO₂ (metaboric) → H₂B₄O₇ (tetraboric) → B₂O₃ (boron trioxide)
Heating boric acid at 1700°C gives metaboric acid.
H₃BO₃ → HBO₂ + H₂O
Heating boric acid at 3000°C gives tetraboric acid.
4HBO₂ → H₂B₄O₇ + H₂O
Heating boric acid at 3300°C gives boron trioxide.
H₂B₄O₇ → 2B₂O₃ + H₂O
With NaOH
Forms sodium tetraborate (borax):
4H₃BO₃ + 2NaOH → Na₂B₄O₇ + 7H₂O
With HF + H₂SO₄
Forms volatile BF₃ (used for purification).
With Sulphuric Acid
Easily dissolved in sulphuric acid
B(OH)₃ + 6H₂SO₄ → B(HSO₄)₄^– + 2HSO₄^– + 3H₃O⁺

Antiseptic: Mild antiseptic in eye washes (boric acid solution), talc powders, ointments (e.g., Boroline).
Insecticide: Especially against ants, cockroaches, and termites (disrupts their digestive system).
Glass manufacture: As a flux in borosilicate glass (Pyrex).
Enamels and glazes: In ceramic industry.
Flame retardant: In textiles and wood.
Neutron absorber: In nuclear reactors (due to high neutron capture by ¹⁰B).
Analytical reagent: In qualitative analysis (borax bead test, green flame test).
pH buffer: In swimming pools and cosmetics.
Preservative: In some food items (limited use).

Note: Use as antiseptic has declined due to toxicity concerns (especially in infants); now restricted in many countries.