Washing Soda

From sodium carbonate to washing soda. Anhydrous sodium carbonate, $\text{Na}_2\text{CO}_3$, is itself made industrially by heating sodium hydrogencarbonate, the same carbonate family that gives baking soda. When this anhydrous solid is dissolved in hot water and the solution is cooled slowly, the sodium carbonate does not simply come back out unchanged — it crystallises with exactly ten water molecules trapped in a fixed ratio inside the crystal lattice for every formula unit of $\text{Na}_2\text{CO}_3$:

where: the left-hand $\text{Na}_2\text{CO}_3(s)$ is anhydrous (water-free) sodium carbonate, $10\text{H}_2\text{O}(l)$ is the liquid water it recrystallises from, and $\text{Na}_2\text{CO}_3\cdot10\text{H}_2\text{O}(s)$ is washing soda, also called sodium carbonate decahydrate. This is why the process is called recrystallisation rather than simple drying.

Water of crystallisation. The dot in the formula is not decorative — it marks a fixed, whole-number count of water molecules chemically bound within a salt's crystal structure, the same idea used for plaster of Paris and other hydrates. Washing soda's ten water molecules per formula unit are exactly why it is a decahydrate; remove them (by heating) and only the anhydrous $\text{Na}_2\text{CO}_3$ remains, no longer washing soda by name.

Why the solution is basic. Washing soda is the salt of a strong base (NaOH) and a weak acid (carbonic acid, $\text{H}_2\text{CO}_3$), and the Class 10 rule for salts of a strong base and a weak acid is that their aqueous solution is basic — it turns red litmus blue.

Uses that follow from this chemistry. Washing soda is used as a raw material in the glass industry and the paper industry, in manufacturing soap and as a cleansing agent itself, and — because its carbonate ion reacts with dissolved calcium and magnesium salts to form insoluble carbonates — for removing the hardness of water, which is also why it is used directly as a laundry and utensil cleaner.

Hard water contains calcium and magnesium ions dissolved as soluble salts. When soap, itself a sodium or potassium salt of a fatty acid, is added to this water, the calcium and magnesium ions displace the sodium or potassium and form an insoluble scum instead of a lather, wasting soap without cleaning anything. Washing soda works because it supplies a soluble source of carbonate ions: these carbonate ions combine with the dissolved calcium and magnesium to form calcium carbonate and magnesium carbonate, both of which are insoluble and settle out or can be filtered away. With the hardness-causing ions removed, the water becomes soft, and soap added afterwards can form a proper lather instead of scum. The same carbonate chemistry that lets sodium carbonate recrystallise cleanly out of water is what lets it selectively pull calcium and magnesium out of solution — the carbonate ion simply prefers to pair with calcium and magnesium over sodium, precipitating the former while leaving sodium salts dissolved. This is a purely descriptive, ion-exchange style reasoning at Class 10 level; no equilibrium constants or solubility-product calculations are expected, only the statement that soluble carbonate converts soluble hardness salts into insoluble ones.

Question: A student recrystallises $53\ \text{g}$ of anhydrous sodium carbonate (one-half mole, since its molar mass is $106\ \text{g/mol}$) completely into washing soda. (a) Write the balanced equation for the process. (b) Calculate the mass of water that must combine with it as water of crystallisation. (c) Calculate the total mass of washing soda formed.

Step 1. Write the recrystallisation equation, the same one used throughout this article: $\text{Na}_2\text{CO}_3(s) + 10\text{H}_2\text{O}(l) \rightarrow \text{Na}_2\text{CO}_3\cdot10\text{H}_2\text{O}(s)$. One mole of anhydrous sodium carbonate always combines with exactly ten moles of water.

Step 2. Convert the given mass to moles: $53\ \text{g} \div 106\ \text{g/mol} = 0.5\ \text{mol}$ of $\text{Na}_2\text{CO}_3$. By the 1:10 ratio in the equation, this requires $0.5 \times 10 = 5\ \text{mol}$ of water.

Step 3. Convert that water to mass: $5\ \text{mol} \times 18\ \text{g/mol} = 90\ \text{g}$ of water is taken up as water of crystallisation.

Step 4. Find the total mass of washing soda formed by adding the sodium carbonate and the water it incorporates: $53\ \text{g} + 90\ \text{g} = 143\ \text{g}$. Checking against the derivation above, $143\ \text{g}$ is exactly half of $286\ \text{g/mol}$, consistent with starting from half a mole. Answer: 90 g of water combines with the 53 g of sodium carbonate to give 143 g of washing soda.

• Writing washing soda's formula without the water of crystallisation (just $\text{Na}_2\text{CO}_3$) — the examinable named compound is specifically the decahydrate, $\text{Na}_2\text{CO}_3\cdot10\text{H}_2\text{O}$.
• Confusing washing soda with baking soda — washing soda is $\text{Na}_2\text{CO}_3\cdot10\text{H}_2\text{O}$, used industrially and for water softening; baking soda is $\text{NaHCO}_3$, used as an antacid and in baking.
• Forgetting to balance the coefficient of $10$ in front of $\text{H}_2\text{O}$ when writing the recrystallisation equation, or dropping it when converting moles of sodium carbonate to moles of water.
• Assuming the basic nature of washing soda's solution needs a full hydrolysis mechanism — at Class 10 level, stating it is a salt of a strong base and weak acid is sufficient.

Ask the tutor: Want to see how the same water-of-crystallisation idea applies to plaster of Paris, or practice another mole-ratio calculation with a different starting mass of sodium carbonate?