3_2 Borax and Boric Acid: Preparation, Bead Test, Structure and Uses

Boron's two best-known compounds are borax, sodium tetraborate decahydrate (

\mathrm{Na_2B_4O_7 \cdot 10H_2O}

), and boric acid, also called orthoboric acid (

\mathrm{H_3BO_3}

). Both follow from the boron chemistry of the previous lesson: boron bonds covalently to oxygen and stays hungry for an extra electron pair. That explains why borax solution is alkaline and why boric acid is a weak acid that works in an unusual way.

Making Borax

Borax occurs naturally as tincal in dried-up salt lakes and can be purified simply by dissolving it in water and recrystallizing. Most borax is now made from the calcium borate mineral colemanite. The powdered mineral is boiled with sodium carbonate solution, which precipitates calcium carbonate and leaves borax and sodium metaborate in solution. Borax crystallizes from the cleared liquid. Blowing carbon dioxide through the remaining liquid converts the metaborate into more borax:

\begin{aligned} \mathrm{Ca_2B_6O_{11} + 2Na_2CO_3} &\rightarrow \mathrm{2CaCO_3 + Na_2B_4O_7 + 2NaBO_2} \\ \mathrm{4NaBO_2 + CO_2} &\rightarrow \mathrm{Na_2B_4O_7 + Na_2CO_3} \end{aligned}

Properties of Borax

Borax is a white crystalline solid that is much more soluble in hot water than in cold. The decahydrate crystallizes below about 62 °C and the pentahydrate above it. Its solution is alkaline because the tetraborate ion reacts with water (hydrolysis) to give sodium hydroxide, a strong base, and boric acid, a very weak acid. The strong base dominates, so the solution has a pH of about 9. Adding a strong acid to hot borax solution reverses this and precipitates boric acid on cooling:

\mathrm{Na_2B_4O_7 + 2HCl + 5H_2O \rightarrow 2NaCl + 4H_3BO_3}

On heating, borax loses its water of crystallization and puffs up into a white porous mass. At higher temperature it melts to a clear glass of sodium metaborate and boric anhydride (

\mathrm{B_2O_3}

):

\mathrm{Na_2B_4O_7 \cdot 10H_2O \xrightarrow{\ \Delta\ } Na_2B_4O_7 \xrightarrow{\ \Delta\ } 2NaBO_2 + B_2O_3}

The Borax Bead Test

This glass is the basis of a classic test for metal ions. A loop of platinum wire is heated, dipped in borax and heated again until a clear, colourless bead forms. A few grains of the sample are then fused into the bead. The bead is heated in the hot outer zone of the flame, where there is plenty of oxygen, and again in the inner zone, which is rich in unburnt gas and reduces the metal. The acidic

\mathrm{B_2O_3}

in the bead combines with the metal oxide to form a metal borate, and many of these are strongly coloured. Copper gives a blue bead in the oxidizing flame, for example, while cobalt gives a deep blue bead in both flames:

\mathrm{CuO + B_2O_3 \rightarrow Cu(BO_2)_2}

Boric Acid: Structure and Acidity

Boric acid is made by adding sulphuric acid to hot, concentrated borax solution, or by treating colemanite with sulphur dioxide in boiling water. It forms white, soapy-feeling flakes. In the crystal, flat

\mathrm{B(OH)_3}

units are joined by hydrogen bonds into layers, and the layers slide over each other, which explains the slippery feel. It dissolves only slightly in cold water but much more in hot water, so it is easy to purify by recrystallization.

Boric acid is a very weak, monobasic acid (

pK_a \approx 9.2

), but not because it releases its own hydrogen atoms. Boron in

\mathrm{B(OH)_3}

has only six valence electrons, so it acts as a Lewis acid. It accepts a hydroxide ion from a water molecule, and the proton left behind makes the solution acidic:

OB(O)O.O.O>>O[B-](O)(O)O.[OH3+]

Because it is so weak, boric acid cannot be titrated accurately with alkali on its own. Adding glycerol or another compound with neighbouring OH groups locks the borate into a stable complex, which shifts the equilibrium and lets the acid be titrated with phenolphthalein. Boric acid also reacts with alcohols to form borate esters. Triethyl borate burns with a green flame, which is a simple test for boron:

OB(O)O.CCO.CCO.CCO>>CCOB(OCC)OCC.O.O.O

Strong heating dehydrates boric acid step by step, first to metaboric acid (

\mathrm{HBO_2}

), then to tetraboric acid and finally to glassy

\mathrm{B_2O_3}

.

Uses

• Borosilicate glass: adding

  \mathrm{B_2O_3}

  lowers the glass's thermal expansion, so laboratory and oven glassware survives sudden temperature changes.
• Fluxes: molten borax dissolves metal oxides, so it cleans metal surfaces in welding, brazing and metal refining.
• Cleaning and water softening: borax is a mild alkali, and it ties up calcium ions in laundry products.
• Glazes and enamels: borate glazes melt at lower temperatures than silicate glazes.
• Mild antiseptic: dilute boric acid is used in some eye washes and powders. It is toxic in larger amounts, so its use is now limited.

Tips for remembering

• Boric acid takes OH⁻, it does not give H⁺: B(OH)₃ accepts a hydroxide from water to form B(OH)₄⁻, and the hydrogen ion left behind is what makes the solution acidic.
• Strong base plus weak acid gives an alkaline solution: borax hydrolyses to NaOH and H₃BO₃, and the strong base wins.
• Bead colour comes from the metal borate: molten borax gives up B₂O₃, which combines with the metal oxide, so copper gives blue and cobalt gives deep blue.