Important Compounds: Bleaching Powder, Baking Soda & Plaster of Paris - Tenth Grade Science
Common salt, sodium chloride, serves as the primary chemical feedstock for a vast array of industrial compounds that underpin public sanitation, manufacturing, food chemistry, and medical orthopedics. Through chlor-alkali electrolysis and subsequent chemical syntheses, basic brine is transformed into essential commodities including calcium hypochlorite bleaching powder, sodium hydrogen carbonate baking soda, sodium carbonate washing soda, and hydrated calcium sulfate plaster of paris. Each of these compounds possesses unique crystalline architectures, thermal decomposition profiles, and chemical functionalities tailored for specialized domestic and industrial applications. In this chapter, you will investigate the chlor-alkali manufacturing process, trace the chemical synthesis, properties, and uses of bleaching powder, baking soda, and washing soda, and explore the crystallization hydration chemistry governing gypsum and plaster of paris.
The Chlor-Alkali Industrial Feedstock Foundation
The chlor-alkali process is the electrolytic decomposition of an aqueous solution of sodium chloride, known as brine, producing three industrially vital chemical commodities: chlorine gas, hydrogen gas, and sodium hydroxide.
+-----------------------------------------------------------------------------------+
| CHLOR-ALKALI ELECTROLYTIC CELL |
+-----------------------------------------------------------------------------------+
| |
| 2NaCl(aq) + 2H2O(l) --[Electricity]--> 2NaOH(aq) + Cl2(g) + H2(g) |
| |
| At the ANODE (Positive Electrode): At the CATHODE (Negative Electrode): |
| 2Cl-(aq) -----> Cl2(g) + 2e- 2H2O(l) + 2e- -----> H2(g) + 2OH-(aq) |
| (Chlorine gas evolved) (Hydrogen gas evolved) |
| |
| In the Solution (Cathode Compartment): |
| Na+(aq) + OH-(aq) -----> NaOH(aq) (Caustic soda solution concentrated) |
+-----------------------------------------------------------------------------------+
Electrolysis of Brine
When a direct electric current passes through concentrated aqueous sodium chloride, sodium ions ($\text{Na}^+$), chloride ions ($\text{Cl}^-$), hydronium ions ($\text{H}^+$), and hydroxide ions ($\text{OH}^-$) migrate toward oppositely charged electrodes:
At the positive anode, chloride ions are oxidized in preference to hydroxide ions due to overpotential dynamics, evolving chlorine gas ($\text{Cl}_2$):
$$2\text{Cl}^-\text{(aq)} \longrightarrow \text{Cl}_2\text{(g)} + 2\text{e}^-$$
At the negative cathode, water molecules are reduced in preference to sodium cations because hydrogen has a more favorable reduction potential, liberating hydrogen gas ($\text{H}_2$):
$$2\text{H}_2\text{O(l)} + 2\text{e}^- \longrightarrow \text{H}_2\text{(g)} + 2\text{OH}^-\text{(aq)}$$
The remaining sodium cations ($\text{Na}^+$) pair with newly generated hydroxide anions ($\text{OH}^-$) in the cathodic compartment, producing concentrated sodium hydroxide ($\text{NaOH}$), commonly called caustic soda. The three products are isolated separately to prevent chlorine gas from reacting with sodium hydroxide.
Bleaching Powder: Calcium Oxychloride ($\text{CaOCl}_2$)
Bleaching powder is chemically designated as calcium oxychloride, possessing the empirical formula $\text{CaOCl}_2$. It serves as a potent oxidizing agent, disinfectant, and industrial bleaching reagent.
+-----------------------------------------------------------------------------------+
| BLEACHING POWDER MANUFACTURING |
+-----------------------------------------------------------------------------------+
| |
| Ca(OH)2(s) + Cl2(g) --[40 deg C]--> CaOCl2(s) + H2O(l) |
| (Dry Slaked Lime) (Chlorine) (Bleaching Powder) |
| |
| Chemical Structure: Mixed salt containing Ca2+, Cl-, and Hypochlorite (OCl-) |
| Available Chlorine: Percentage of active chlorine released upon acidification |
+-----------------------------------------------------------------------------------+
Industrial Preparation and Structural Identity
Bleaching powder is manufactured by passing chlorine gas over dry slaked lime, calcium hydroxide, maintained at approximately forty degrees Celsius in Hasenclever plants:
Ca(OH)2(s) + Cl2(g) -> CaOCl2(s) + H2O(l)
The resulting product is a pale yellowish-white powder with a strong odor of chlorine. Structurally, bleaching powder is not a simple double salt; it is a complex compound containing calcium cations ($\text{Ca}^{2+}$), chloride anions ($\text{Cl}^-$), and hypochlorite anions ($\text{OCl}^-$), often represented as $\text{Ca(OCl)Cl}$.
Chemical Properties and Available Chlorine
When bleaching powder is exposed to ambient atmospheric air, it slowly reacts with atmospheric carbon dioxide, releasing chlorine gas:
CaOCl2(s) + CO2(g) -> CaCO3(s) + Cl2(g)
When treated with dilute mineral acids, such as hydrochloric or sulfuric acid, bleaching powder undergoes complete decomposition, liberating all of its active chlorine content:
CaOCl2(s) + H2SO4(aq) -> CaSO4(aq) + H2O(l) + Cl2(g)
CaOCl2(s) + 2HCl(aq) -> CaCl2(aq) + H2O(l) + Cl2(g)
The chlorine liberated during this reaction is designated as available chlorine. High-grade commercial bleaching powder contains between 35% and 38% available chlorine.
Bleaching action is driven by oxidation. The hypochlorite group liberates nascent oxygen ($[\text{O}]$) in the presence of moisture:
Cl2 + H2O -> 2HCl + [O]
Coloured Organic Matter + [O] -> Colourless Oxidized Product
Because nascent oxygen bleaches by oxidative destruction of chromophores, the bleaching effect is permanent. Bleaching powder is used to bleach wood pulp in paper factories, cotton and linen in textile mills, and to sterilize municipal drinking water supplies by eradicating microbial pathogens.
Baking Soda: Sodium Hydrogen Carbonate ($\text{NaHCO}_3$)
Sodium hydrogen carbonate, commonly called baking soda, is a white crystalline powder that functions as a mild, non-corrosive basic salt with widespread culinary, medical, and fire-safety applications.
+-----------------------------------------------------------------------------------+
| SOLVAY PROCESS FOR BAKING SODA |
+-----------------------------------------------------------------------------------+
| |
| NaCl + H2O + NH3 + CO2 -----> NH4Cl(aq) + NaHCO3(s) (Precipitates out) |
| |
| Thermolysis on Cooking: |
| 2NaHCO3(s) --[Heat]--> Na2CO3(s) + H2O(g) + CO2(g) (Leavens dough) |
+-----------------------------------------------------------------------------------+
Synthesis via the Solvay Process
Industrial manufacturing of baking soda utilizes the Solvay process. Cold, concentrated aqueous brine ($\text{NaCl}$) saturated with ammonia gas ($\text{NH}_3$) is reacted with carbon dioxide gas ($\text{CO}_2$) in a carbonating tower:
NaCl + H2O + NH3 + CO2 -> NH4Cl + NaHCO3
Sodium hydrogen carbonate has a relatively low solubility in water compared to ammonium chloride. Consequently, it precipitates out of the ammoniacal brine as a crystalline solid, which is filtered, washed, and dried.
Thermal Decomposition and Baking Powder Formulation
When sodium hydrogen carbonate is heated during cooking, it undergoes thermal decomposition, releasing carbon dioxide gas, steam, and sodium carbonate:
2NaHCO3(s) --[Heat]--> Na2CO3(s) + H2O(g) + CO2(g)
The escaping bubbles of carbon dioxide gas become trapped in the viscous batter, expanding upon heating to create a porous, spongy texture in bread and cakes. However, unadulterated baking soda cannot be used alone in baking because the residual sodium carbonate ($\text{Na}_2\text{CO}_3$) imparts a bitter alkaline taste and yellow discoloration.
Commercial baking powder solves this problem by combining sodium hydrogen carbonate with a mild edible solid organic acid, typically potassium hydrogen tartrate (cream of tartar) or tartaric acid ($\text{C}_4\text{H}_6\text{O}_6$). When dry baking powder is moistened in batter, the solid acid dissociates, reacting with baking soda:
NaHCO3 + H+ (from tartaric acid) -> Na+ + H2O + CO2(g)
Carbon dioxide is evolved without thermal decomposition, and the resulting sodium tartrate salt has a pleasant, neutral flavor.
Uses in Fire Extinguishers and Antacids
Baking soda is utilized in soda-acid fire extinguishers. When the extinguisher is triggered, concentrated sulfuric acid mixes with a concentrated solution of sodium hydrogen carbonate:
2NaHCO3(aq) + H2SO4(aq) -> Na2SO4(aq) + 2H2O(l) + 2CO2(g)
The reaction generates high volumes of pressurized carbon dioxide gas, which forces a stream of water and gas onto the fire, smothering flames by displacing atmospheric oxygen. In medicine, sodium hydrogen carbonate serves as an oral antacid to neutralize excess gastric hydrochloric acid.
Washing Soda: Sodium Carbonate Decahydrate ($\text{Na}_2\text{CO}_3\cdot 10\text{H}_2\text{O}$)
Washing soda is the decahydrate form of sodium carbonate, synthesized through the thermal calcination of baking soda followed by controlled recrystallization.
+-----------------------------------------------------------------------------------+
| WASHING SODA SYNTHESIS PATHWAY |
+-----------------------------------------------------------------------------------+
| Step 1: Thermal Calcination of Baking Soda |
| 2NaHCO3(s) --[Heat]--> Na2CO3(s) + H2O(g) + CO2(g) (Soda Ash) |
| |
| Step 2: Aqueous Recrystallization |
| Na2CO3(s) + 10H2O(l) -----> Na2CO3.10H2O(s) (Washing Soda Decahydrate) |
+-----------------------------------------------------------------------------------+
Manufacturing and Water Softening Mechanics
Heating baking soda yields anhydrous sodium carbonate, known industrially as soda ash. Dissolving soda ash in hot water and allowing the saturated solution to cool causes sodium carbonate decahydrate crystals to precipitate:
Na2CO3(s) + 10H2O(l) -> Na2CO3.10H2O(s)
Washing soda is an essential industrial chemical used in glass manufacturing, paper production, and domestic laundry detergents. Its primary industrial chemical role is softening permanent hard water containing dissolved calcium and magnesium sulfate or chloride. The carbonate ions precipitate the dissolved divalent metal ions out of solution:
CaSO4(aq) + Na2CO3(aq) -> CaCO3(s) + Na2SO4(aq)
MgCl2(aq) + Na2CO3(aq) -> MgCO3(s) + 2NaCl(aq)
Once converted into insoluble carbonates, calcium and magnesium ions cannot interfere with soap lathering.
Plaster of Paris: Calcium Sulfate Hemihydrate ($\text{CaSO}_4\cdot \frac{1}{2}\text{H}_2\text{O}$)
Plaster of Paris, historically named after gypsum deposits quarried in Montmartre near Paris, is calcium sulfate hemihydrate, possessing the chemical formula $\text{CaSO}_4\cdot \frac{1}{2}\text{H}_2\text{O}$ (or $2\text{CaSO}_4\cdot\text{H}_2\text{O}$).
+-----------------------------------------------------------------------------------+
| PLASTER OF PARIS HYDRATION EQUILIBRIUM |
+-----------------------------------------------------------------------------------+
| |
| CaSO4.2H2O(s) ---[Heat at 373 K (100 C)]---> CaSO4.1/2H2O(s) + 1.5H2O(g) |
| (Gypsum) (Plaster of Paris) |
| | |
| + 1.5 H2O (Setting process) | |
| v |
| CaSO4.2H2O(s) |
| (Hard Interlocking Mass) |
+-----------------------------------------------------------------------------------+
Thermal Calcination of Gypsum
Plaster of Paris is manufactured by carefully heating gypsum—calcium sulfate dihydrate ($\text{CaSO}_4\cdot 2\text{H}_2\text{O}$)—to precisely 373 kelvin (100 degrees Celsius) in a rotary kiln:
CaSO4.2H2O(s) --[373 K]--> CaSO4.1/2H2O(s) + 1.5H2O(g)
The fractional coefficient $\frac{1}{2}\text{H}_2\text{O}$ denotes that two formula units of calcium sulfate share a single coordinate water molecule within the crystal lattice ($2\text{CaSO}_4\cdot\text{H}_2\text{O}$).
Strict temperature control is essential during manufacturing. If the thermal kiln exceeds 373 kelvin (e.g., above 393 K), gypsum loses all of its water of crystallization to become anhydrous calcium sulfate ($\text{CaSO}_4$):
CaSO4.2H2O(s) --[Above 393 K]--> CaSO4(s) + 2H2O(g)
Anhydrous calcium sulfate is termed dead burnt plaster because it has lost its capacity to rehydrate and set into a hard mass upon mixing with water.
The Setting Mechanism of Plaster of Paris
When Plaster of Paris powder is blended with approximately one-third its weight of water, it forms a plastic, workable paste that hydrates back into gypsum:
CaSO4.1/2H2O(s) + 1.5H2O(l) -> CaSO4.2H2O(s) + Heat
The setting process is an exothermic physical transformation occurring in two distinct phases:
First, within five to ten minutes, the paste sets into a stiff semi-solid mass. Second, over the subsequent twenty to thirty minutes, elongated monoclinic gypsum micro-crystals precipitate and interlock into a rigid structural matrix. During setting, the material undergoes a minute volumetric expansion (approximately 0.5% to 1%), enabling the slurry to fill molds and reproduce intricate fine details.
Plaster of Paris is used extensively in orthopedic medicine to immobilize fractured bones in rigid casts, in dentistry for oral impressions, in architecture for ornamental ceiling moldings, and in sculpture for casting fireproof statues.
Chapter Exercises
Exercise 1: An industrial plant manufactures bleaching powder, caustic soda, and baking soda. State the primary chemical feedstock common to all three processes, write the balanced chemical equations demonstrating the production of chlorine gas from this feedstock, and explain how that chlorine gas is converted into bleaching powder.
Exercise 2: A baker prepares a cake using baking soda alone and observes that while the cake rises effectively, it develops an unpleasantly bitter taste and yellowish streaks. Explain the chemical reactions responsible for both the rising and the bitterness, and describe how baking powder prevents this issue.
Exercise 3: Explain the structural meaning of the formula $\text{CaSO}_4\cdot \frac{1}{2}\text{H}_2\text{O}$, explain why Plaster of Paris must be stored in airtight, moisture-proof containers, and describe the consequences of overheating gypsum above 393 kelvin.
Solutions and Explanations
Solution 1: The primary chemical feedstock common to all three processes is sodium chloride ($\text{NaCl}$), obtained as rock salt or seawater brine. In the chlor-alkali process, concentrated aqueous sodium chloride is electrolyzed: $2\text{NaCl(aq)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)} + \text{Cl}_2\text{(g)} + \text{H}_2\text{(g)}$. The chlorine gas evolved at the anode is collected and reacted with dry slaked lime (calcium hydroxide) at approximately forty degrees Celsius: $\text{Ca(OH)}_2\text{(s)} + \text{Cl}_2\text{(g)} \rightarrow \text{CaOCl}_2\text{(s)} + \text{H}_2\text{O(l)}$, producing bleaching powder.
Solution 2: The rising of the cake occurs because heating baking soda (sodium hydrogen carbonate) induces thermal decomposition: $2\text{NaHCO}_3\text{(s)} \rightarrow \text{Na}_2\text{CO}_3\text{(s)} + \text{H}_2\text{O(g)} + \text{CO}_2\text{(g)}$. The evolving carbon dioxide gas expands within the dough, creating a porous, leavened matrix. However, this decomposition also yields sodium carbonate ($\text{Na}_2\text{CO}_3$), which is an alkaline, basic salt. The alkalinity of sodium carbonate imparts a sharp, unpleasantly bitter taste and causes the alkaline degradation of flour pigments, creating yellowish streaks. Baking powder resolves this problem by formulating sodium hydrogen carbonate with a solid edible organic acid, such as tartaric acid. When moistened, tartaric acid reacts directly with baking soda without requiring thermal decomposition: $\text{NaHCO}_3 + \text{H}^+ \rightarrow \text{Na}^+ + \text{H}_2\text{O} + \text{CO}_2\text{(g)}$. The resulting sodium tartrate salt is neutral and tasteless, preserving the flavor and appearance of the cake.
Solution 3: In the chemical formula $\text{CaSO}_4\cdot \frac{1}{2}\text{H}_2\text{O}$, a fractional half-molecule of water does not represent a split water molecule, which is physically impossible. Rather, it indicates a crystallographic ratio wherein two formula units of calcium sulfate share a single coordinate water molecule within the unit cell, which can be represented as $2\text{CaSO}_4\cdot\text{H}_2\text{O}$. Plaster of Paris must be stored in airtight, moisture-proof containers because it is hygroscopic; exposure to ambient atmospheric moisture causes it to slowly hydrate into calcium sulfate dihydrate (gypsum): $\text{CaSO}_4\cdot \frac{1}{2}\text{H}_2\text{O} + 1.5\text{H}_2\text{O} \rightarrow \text{CaSO}_4\cdot 2\text{H}_2\text{O}$. This premature hydration sets the powder into a hard, unusable crystalline mass. Overheating gypsum above 393 kelvin strips away all water of crystallization, producing anhydrous calcium sulfate ($\text{CaSO}_4$). This anhydrous form is dead burnt plaster, which lacks the ability to hydrate and set upon the addition of water.