Bleaching Powder

Bleaching powder is made industrially and in the school laboratory by one reaction: dry, powdered slaked lime is spread out and chlorine gas is passed over it.

where: $\text{Ca(OH)}_2$ is slaked lime (calcium hydroxide, a white solid), $\text{Cl}_2$ is chlorine gas, $\text{CaOCl}_2$ is calcium oxychloride — the compound sold and used as bleaching powder, sometimes loosely called "chloride of lime" — and $\text{H}_2\text{O}$ is the water released alongside it. The reaction is a straightforward 1:1:1:1 combination, and the chlorine gas itself usually comes from the chlor-alkali process, the electrolysis of brine that also yields sodium hydroxide. That shared origin is why "Sodium Hydroxide" and "Bleaching Powder" sit back to back in the same syllabus unit. Two things follow from this preparation, and each is a distinct exam angle:

Its usefulness is entirely about giving that chlorine back. Bleaching powder does not bleach, disinfect or oxidise as an inert solid sitting in a beaker. It has to react with a dilute acid — even a weak one like the carbonic acid formed from atmospheric $\text{CO}_2$ and moisture — to liberate chlorine, and it is that liberated chlorine which does the actual work: decolourising dyed cloth, killing microorganisms in water, or oxidising a substrate in an industrial process. This single idea explains all three "main uses" you are asked to list:

• Bleaching — cotton and linen in textile mills, wood pulp in paper factories, and washed clothes in laundries, where the released chlorine strips colour from dyes and stains.
• Disinfecting — drinking water at treatment plants and sanitary surfaces such as toilets, where the released chlorine kills disease-causing microorganisms.
• Oxidising — various manufacturing processes in the chemical industry, where the released chlorine (or the hypochlorite-type species it forms) acts as an oxidising agent.

The same reactivity that makes it useful makes it unstable in storage. Left open to moist air, bleaching powder slowly reacts with atmospheric $\text{CO}_2$ and water vapour exactly as it would with any dilute acid, gradually losing its available chlorine and its bleaching strength, while releasing a chlorine-like smell. This is why it must be kept in closed, air-tight containers, away from moisture and direct sunlight — not an arbitrary storage rule, but a direct consequence of the chemistry that makes it work in the first place.

Question: A plant feeds $148\ \text{g}$ of dry slaked lime into a reactor and passes excess chlorine gas over it until the reaction is complete. Calculate the mass of bleaching powder formed, and the mass of water released alongside it.

Step 1. Convert the given mass of slaked lime to moles using its molar mass of $74\ \text{g/mol}$ (calculated above): $n(\text{Ca(OH)}_2) = \dfrac{148}{74} = 2\ \text{mol}$.

Step 2. Apply the $1:1:1:1$ mole ratio from the balanced equation. $2\ \text{mol}$ of slaked lime reacts with $2\ \text{mol}$ of chlorine gas (present in excess, so it is not limiting) to give $2\ \text{mol}$ of bleaching powder and $2\ \text{mol}$ of water.

Step 3. Convert moles of bleaching powder back to mass using its molar mass of $127\ \text{g/mol}$: $\text{mass}(\text{CaOCl}_2) = 2 \times 127 = 254\ \text{g}$.

Step 4. Convert moles of water back to mass using its molar mass of $18\ \text{g/mol}$ ($2 \times 1 + 16$): $\text{mass}(\text{H}_2\text{O}) = 2 \times 18 = 36\ \text{g}$.

Step 5. Check by mass conservation: the reacted slaked lime plus the chlorine gas consumed should equal the total mass of products. Chlorine consumed is $2\ \text{mol} \times 71\ \text{g/mol} = 142\ \text{g}$, so reactant mass is $148 + 142 = 290\ \text{g}$, and product mass is $254 + 36 = 290\ \text{g}$ — the two match, confirming the arithmetic. The plant obtains $254\ \text{g}$ of bleaching powder and $36\ \text{g}$ of water from $148\ \text{g}$ of slaked lime.

Competitive stretch

To deliver $635\ \text{g}$ of bleaching powder, work backwards: $635 \div 127 = 5\ \text{mol}$ of $\text{CaOCl}_2$ needed, so $5 \times 74 = 370\ \text{g}$ of slaked lime and at least $5 \times 71 = 355\ \text{g}$ of chlorine gas must be fed in.

• Writing the formula as $\text{Ca(OCl)}_2$ or another variant instead of the standard NCERT formula $\text{CaOCl}_2$ used at this level.
• Confusing the reactant "slaked lime" with "quicklime" — slaked lime is $\text{Ca(OH)}_2$ (calcium hydroxide), quicklime is $\text{CaO}$ (calcium oxide); only slaked lime reacts with chlorine gas to give bleaching powder.
• Using the wrong molar mass for chlorine in stoichiometry — chlorine's atomic mass is $35.5$, so $\text{CaOCl}_2$'s molar mass is $127$, not $126$ or $128$; a rounding slip here throws off every mass calculated from it.
• Attributing "removing hardness of water" or "used in glass and soap making" to bleaching powder — those uses belong to washing soda, not bleaching powder; keep the two Class 10 salts and their use-lists separate.

Ask the tutor: want to try the stoichiometry the other way round, starting from a target mass of chlorine gas instead of slaked lime, or see how this preparation connects to the chlor-alkali process a few sections earlier in the chapter?