Vardaan Learning Institute

Litmus Behaviour Difference:

• Acids turn blue litmus solution red, but have no effect on red litmus.
• Bases turn red litmus solution blue, but have no effect on blue litmus.

Three Natural Indicators & Colour Changes:

1. Litmus solution (obtained from Lichens): Remains red/purple in acidic solutions, turns blue in basic solutions.
2. Turmeric: Remains yellow in acidic solutions, turns reddish-brown in basic solutions.
3. Red Cabbage Extract: Turns red/pink in acidic solutions, turns green/yellow in basic solutions.

Is this test sufficient? Yes, because an acidic substance would have turned moist blue litmus paper red, and a basic substance would have turned moist red litmus paper blue. Since neither showed a colour change, the solution is neutral. (Note: The litmus paper must be moist, as dry conditions prevent the dissociation and formation of ions required to show the color change).

Indicator	Colour in Acidic Solution	Colour in Basic Solution
1. Methyl orange	Red / Pink	Yellow
2. Phenolphthalein	Colourless	Pink
3. Red litmus	Red (No change)	Blue
4. Turmeric	Yellow (No change)	Reddish-brown

Therefore, dry $\text{HCl}$ gas cannot change dry litmus paper's colour. However, moist $\text{HCl}$ gas dissolves in the water present on the moist litmus paper, dissociating to form hydronium ions ($\text{H}_3\text{O}^+$), which turn blue litmus paper red: $$\text{HCl}(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$$ Conclusion: This experiment proves that acids exhibit their acidic behaviour and release $\text{H}^+$ ions only in the presence of water.

1. Litmus Paper:
   • If it turns blue litmus paper red $\rightarrow$ Acidic.
   • If it turns red litmus paper blue $\rightarrow$ Basic.
   • If it has no effect on either red or blue litmus paper $\rightarrow$ Neutral.
2. Methyl Orange: Add a drop of methyl orange.
   • If it turns red $\rightarrow$ Acidic.
   • If it turns yellow $\rightarrow$ Basic.
   • If it remains orange $\rightarrow$ Neutral.
3. Phenolphthalein: Add a drop of phenolphthalein.
   • If it remains colourless $\rightarrow$ Acidic or Neutral.
   • If it turns deep pink $\rightarrow$ Basic.

Chemical property: Strong bases (alkalis) like $\text{NaOH}$ hydrolyze the organic esters (fats/oils) to produce soap and glycerol, which gives the soapy/slippery feel.

• Taste: Sour.
• Touch: Corrosive and cause painful burns on the skin.
• Corrosive nature: Acids react vigorously with organic matter and metals, destroying tissues and causing severe chemical burns.

Why never tasted directly? Concentrated acids are highly corrosive and toxic. Tasting them can lead to severe burns in the mouth, tongue, esophagus, and can even be fatal.

Utility: They are extremely useful for visually impaired students who cannot detect colour changes, allowing them to perform acid-base chemistry experiments easily.
Example: Onion extract or vanilla essence.
How it works: Onion extract has a characteristic strong smell. When treated with an acid, it retains its smell. However, when treated with a base (like $\text{NaOH}$ solution), the characteristic onion smell is completely destroyed.

Initially, the phenolphthalein remains colourless in the acid. As the base ($\text{NaOH}$) is added, it neutralises the acid. Once the acid is completely neutralised, the slight excess of base turns the solution basic, turning the phenolphthalein pink.
Type of reaction: This is a Neutralisation reaction.

Gas test: Pass the gas through soap solution. Bubbles will form. Bring a burning candle near a gas-filled soap bubble; the bubble bursts and the gas burns with a characteristic 'pop' sound, confirming it is hydrogen.
Balanced chemical equation: $$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)$$

1. Magnesium ribbon: $$\text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g)$$
2. Aluminium powder: $$2\text{Al}(s) + 6\text{HCl}(aq) \rightarrow 2\text{AlCl}_3(aq) + 3\text{H}_2(g)$$
3. Iron filings: $$\text{Fe}(s) + 2\text{HCl}(aq) \rightarrow \text{FeCl}_2(aq) + \text{H}_2(g)$$

Balanced chemical equation: $$\text{Na}_2\text{CO}_3(s) + 2\text{HCl}(aq) \rightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$ Gas Identification: Pass the evolved gas through lime water (calcium hydroxide solution). The lime water turns milky due to the formation of an insoluble white precipitate of calcium carbonate ($\text{CaCO}_3$): $$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s)\downarrow + \text{H}_2\text{O}(l)$$

Balanced chemical equation: $$2\text{NaHCO}_3(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{Na}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l) + 2\text{CO}_2(g)$$ Products identified: Sodium sulphate ($\text{Na}_2\text{SO}_4$), water ($\text{H}_2\text{O}$), and carbon dioxide ($\text{CO}_2$).

Most metal oxides are basic and react with acids, but amphoteric metal oxides can react with both acids and bases.
Two differences:

1. Basic metal oxides (e.g., $\text{CuO}$, $\text{CaO}$) react with acids to form salt and water, but do not react with bases. $$\text{CuO}(s) + 2\text{HCl}(aq) \rightarrow \text{CuCl}_2(aq) + \text{H}_2\text{O}(l)$$
2. Amphoteric metal oxides (e.g., $\text{Al}_2\text{O}_3$, $\text{ZnO}$) react with both acids and bases to form salt and water.
   • With acid: $\text{ZnO}(s) + 2\text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2\text{O}(l)$
   • With base: $\text{ZnO}(s) + 2\text{NaOH}(aq) \rightarrow \text{Na}_2\text{ZnO}_2(aq) + \text{H}_2\text{O}(l)$ (Sodium zincate is formed).

Equation: $$\text{CuO}(s) + 2\text{HCl}(aq) \rightarrow \text{CuCl}_2(aq) + \text{H}_2\text{O}(l)$$ Colour change and reason: The black copper oxide powder dissolves and the solution turns a beautiful blue-green. This colour change is due to the formation of copper(II) chloride ($\text{CuCl}_2$) in the solution.

Balanced chemical equation: $$2\text{NaOH}(aq) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{Na}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l)$$ Nature of heat: This reaction is highly exothermic, releasing a large amount of heat energy and raising the temperature of the mixture.

1. Dilute hydrochloric acid: $$\text{NaOH}(aq) + \text{HCl}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$
2. Aluminium metal: $$2\text{NaOH}(aq) + 2\text{Al}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{NaAlO}_2(aq) + 3\text{H}_2(g)$$ (Product is Sodium aluminate).
3. Zinc oxide: $$2\text{NaOH}(aq) + \text{ZnO}(s) \rightarrow \text{Na}_2\text{ZnO}_2(aq) + \text{H}_2\text{O}(l)$$ (Product is Sodium zincate).

Examples:

1. Sulphur dioxide in water: Forms sulphurous acid ($\text{H}_2\text{SO}_3$). $$\text{SO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_3(aq)$$
2. Carbon dioxide in water: Forms carbonic acid ($\text{H}_2\text{CO}_3$). $$\text{CO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{CO}_3(aq)$$

Both acids release hydrogen ions ($\text{H}^+$) in water, turning blue litmus red, which proves that non-metallic oxides are acidic.

When excess carbon dioxide is passed through this milky solution, the milkiness disappears. This is because the insoluble calcium carbonate reacts with excess $\text{CO}_2$ and water to form soluble calcium hydrogen carbonate [$\text{Ca(HCO}_3)_2$], which dissolves completely: $$\text{CaCO}_3(s) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \rightarrow \text{Ca(HCO}_3)_2(aq)$$

1. $$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s)\downarrow + \text{H}_2\text{O}(l)$$
2. $$\text{NaOH}(aq) + \text{HCl}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$
3. $$\text{Mg}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{MgSO}_4(aq) + \text{H}_2(g)$$
4. $$\text{Na}_2\text{CO}_3(s) + 2\text{HCl}(aq) \rightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$

1. An aqueous solution of an acid conducts electricity: Acids (like $\text{HCl}$ or $\text{HNO}_3$) completely or partially dissociate in water to produce free-moving hydrogen ions ($\text{H}^+$ or hydronium ions $\text{H}_3\text{O}^+$) along with anions. These mobile ions act as charge carriers, allowing the flow of electric current. $$\text{HCl}(aq) \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)$$
2. A glucose solution does not conduct electricity: Glucose ($\text{C}_6\text{H}_{12}\text{O}_6$) is a covalent compound. Although it contains hydrogen atoms, it does not dissociate or ionise in water to produce free ions. Since there are no mobile ions to carry electric charge, the glucose solution does not conduct electricity.

pH ranges:

1. Strong acid: pH between 0 and 3.
2. Neutral solution: pH exactly equal to 7.
3. Strong base: pH between 12 and 14.

• The solution with pH 3 has a pH < 7, so it is acidic.
• The solution with pH 11 has a pH > 7, so it is basic.

Which is more acidic — pH 2 or pH 5?
The solution with pH 2 is much more acidic than the one with pH 5.
Justification: pH is inversely proportional to hydrogen ion ($\text{H}^+$) concentration. As the pH value decreases, the concentration of $\text{H}^+$ ions increases exponentially. A pH difference of 3 units means the pH 2 solution has $10^3 = 1000$ times higher $\text{H}^+$ ion concentration than the pH 5 solution.

1. pH in our Digestive System: Our stomach produces $\text{HCl}$ (pH ~ 1.5–3.0) to activate pepsin for protein digestion. If the pH drops too low, it causes hyperacidity and irritation.
2. pH Change as the Cause of Tooth Decay: When the pH of our mouth falls below 5.5, the acid begins to corrode the calcium phosphate enamel of our teeth, leading to tooth decay.
3. Soil pH and Plant Growth: Plants require a specific pH range (typically 6.0–7.0) to absorb nutrients effectively. Soil that is too acidic or too basic prevents proper plant growth.

Excess acid (acidity): During indigestion, the stomach produces too much $\text{HCl}$, causing pain, burning sensation, and irritation.
Remedy: Antacids, which are mild bases (such as Magnesium hydroxide $\text{Mg(OH)}_2$ - milk of magnesia, or Sodium hydrogen carbonate $\text{NaHCO}_3$).
Reason: Antacids, being basic, react with the excess hydrochloric acid in the stomach and neutralise it, forming salt and water, which provides immediate relief.

$$\text{SO}_2 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_3 \text{ (Sulphurous acid)}$$ These acids mix with rainwater, lowering its pH below 5.6.
Phenomenon: This is called acid rain. It corrodes buildings, damages forests, and turns water bodies acidic, which is highly toxic to aquatic organisms.

Chemistry involved: These compounds are basic in nature. When added to the soil, they react with and neutralise the excess acids present in the soil, raising the soil's pH back to the optimal neutral range (6.0–7.0) suitable for plant growth: $$\text{CaO}(s) + 2\text{H}^+(aq) \rightarrow \text{Ca}^{2+}(aq) + \text{H}_2\text{O}(l)$$

Remedy: Use toothpaste to brush teeth at least twice a day.
How it helps: Toothpastes are formulated to be basic/alkaline in nature (pH ~ 8.0–9.0). They neutralise the acids produced by oral bacteria, clean the mouth, and prevent the pH from dropping below the critical 5.5 threshold, thereby protecting enamel from corrosion.

Difference from a single indicator: A single indicator like litmus can only tell whether a solution is acidic or basic (binary result). It cannot measure the strength of the acid or base. In contrast, a universal indicator not only identifies if a solution is acidic or basic but also determines its exact pH/strength based on the specific colour shown (e.g., dark red for pH 1, green for pH 7, purple for pH 14).

Classification & Reasons:

1. $\text{NaCl}$ (Neutral Salt): Formed from a strong acid ($\text{HCl}$) and a strong base ($\text{NaOH}$). In water, it does not hydrolyse, and the pH remains 7.
2. $\text{Na}_2\text{CO}_3$ (Basic Salt): Formed from a weak acid ($\text{H}_2\text{CO}_3$) and a strong base ($\text{NaOH}$). The carbonate ion ($\text{CO}_3^{2-}$) hydrolyses in water to release excess $\text{OH}^-$ ions, making the solution basic (pH > 7).
3. $\text{NH}_4\text{Cl}$ (Acidic Salt): Formed from a strong acid ($\text{HCl}$) and a weak base ($\text{NH}_4\text{OH}$). The ammonium ion ($\text{NH}_4^+$) hydrolyses in water to release hydronium ions, making the solution acidic (pH < 7).
4. $\text{CH}_3\text{COONa}$ (Basic Salt): Formed from a weak acid ($\text{CH}_3\text{COOH}$) and a strong base ($\text{NaOH}$). The acetate ion ($\text{CH}_3\text{COO}^-$) hydrolyses to yield excess $\text{OH}^-$ ions, making it basic (pH > 7).

Justification: Neither the cation nor the anion of such a salt undergoes hydrolysis in water, meaning they do not affect the concentration of $\text{H}^+$ or $\text{OH}^-$ ions.
Example: Sodium chloride ($\text{NaCl}$) formed from strong acid $\text{HCl}$ and strong base $\text{NaOH}$: $$\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$$ In aqueous solution, it completely dissociates into $\text{Na}^+$ and $\text{Cl}^-$ ions, which remain spectator ions. The pH remains exactly 7.0.

Balanced chemical equation: $$\text{NaCl}(s) + \text{H}_2\text{SO}_4(conc.) \rightarrow \text{NaHSO}_4(s) + \text{HCl}(g)\uparrow \text{ (at low temperature)}$$ If heated to higher temperatures: $$2\text{NaCl}(s) + \text{H}_2\text{SO}_4(conc.) \xrightarrow{\Delta} \text{Na}_2\text{SO}_4(s) + 2\text{HCl}(g)\uparrow$$ Products: Sodium hydrogen sulphate ($\text{NaHSO}_4$) or Sodium sulphate ($\text{Na}_2\text{SO}_4$), and hydrogen chloride gas ($\text{HCl}$).

Balanced chemical equation: $$2\text{NaCl}(aq) + 2\text{H}_2\text{O}(l) \xrightarrow{\text{Electrolysis}} 2\text{NaOH}(aq) + \text{Cl}_2(g)\uparrow + \text{H}_2(g)\uparrow$$ Products & important uses:

1. Sodium hydroxide ($\text{NaOH}$ - formed near cathode): Used in soap and detergent manufacturing, paper making, and degreasing metals.
2. Chlorine gas ($\text{Cl}_2$ - evolved at anode): Used for water disinfection, PVC manufacturing, CFCs, and bleaching agents.
3. Hydrogen gas ($\text{H}_2$ - evolved at cathode): Used as rocket fuel, in ammonia synthesis for fertilisers, and margarine production.

1. Sodium hydroxide ($\text{NaOH}$):
   • Manufacturing soaps, synthetic detergents, and paper.
   • Degreasing metals and refining petroleum.
   • Preparation of artificial textiles (e.g., rayon).
2. Hydrochloric acid ($\text{HCl}$):
   • Cleaning iron sheets before galvanisation (pickling).
   • Preparation of ammonium chloride and medicines.
   • Used in textile, dyeing, and leather industries.
3. Sulphuric acid ($\text{H}_2\text{SO}_4$):
   • Manufacturing chemical fertilisers (e.g., ammonium sulphate, superphosphate).
   • Used in lead-acid storage batteries for vehicles.
   • Manufacturing paints, plastics, explosives, and synthetic detergents.

Preparation from baking soda ($\text{NaHCO}_3$):

1. Thermal decomposition of Baking Soda: Baking soda is heated to produce anhydrous sodium carbonate (soda ash), water, and carbon dioxide: $$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$
2. Recrystallisation: Anhydrous sodium carbonate ($\text{Na}_2\text{CO}_3$) is dissolved in water and recrystallised to obtain washing soda crystals containing 10 molecules of water of crystallisation: $$\text{Na}_2\text{CO}_3(s) + 10\text{H}_2\text{O}(l) \rightarrow \text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}(s)$$

Substance	Chemical Name	Chemical Formula	Important Use
1. Baking Soda	Sodium hydrogen carbonate	$\text{NaHCO}_3$	Ingredient in baking powder; antacid to neutralise stomach acidity.
2. Washing Soda	Sodium carbonate decahydrate	$\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$	Used in glass, soap, and paper industries; softening hard water.
3. Bleaching Powder	Calcium oxychloride	$\text{CaOCl}_2$	Disinfecting drinking water; bleaching cotton and wood pulp.
4. Plaster of Paris	Calcium sulphate hemihydrate	$\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$	Plastering fractured bones; making toys, casts, and decorative items.

Blue Vitriol:

• IUPAC Name: Copper(II) sulphate pentahydrate.
• Formula: $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$.

Heating effect: When blue vitriol is heated strongly, it loses its 5 molecules of water of crystallisation and turns into a white anhydrous powder. The crystalline shape is lost: $$\text{CuSO}_4 \cdot 5\text{H}_2\text{O}(s) \text{ (Blue)} \xrightarrow{\Delta} \text{CuSO}_4(s) \text{ (White)} + 5\text{H}_2\text{O}(g)$$ If water is added back to this white powder, it absorbs water and regains its blue colour and crystalline shape.

Equation: $$\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}(s) + 1\frac{1}{2}\text{H}_2\text{O}(l) \rightarrow \text{CaSO}_4 \cdot 2\text{H}_2\text{O}(s) \text{ (Gypsum)}$$ Why store in moisture-proof containers? If exposed to atmospheric moisture, Plaster of Paris will slowly absorb water and turn into hard gypsum over time. This makes the plaster completely useless for casting, plastering, or sculpting since it loses its setting property.

Chemical formula: $\text{CaOCl}_2$ (Calcium oxychloride).
Two uses:

1. To disinfect drinking water to make it free of pathogens.
2. For bleaching cotton and linen in textile industries and wood pulp in paper factories.

Reaction with excess carbon dioxide: $$\text{CaOCl}_2(s) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s) + \text{Cl}_2(g)\uparrow$$ This reaction releases active chlorine gas, which is responsible for the bleaching action.

Molecules of water of crystallisation:

1. $\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$ (Washing Soda) $\rightarrow$ 10 molecules of water.
2. $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ (Copper sulphate / Blue vitriol) $\rightarrow$ 5 molecules of water.
3. $\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$ (Plaster of Paris) $\rightarrow$ $\frac{1}{2}$ (half) molecule of water per formula unit (or 1 molecule per two units).

Test results:

1. Red litmus paper turns blue.
2. Phenolphthalein indicator turns deep pink.

Why the solution is basic: In water, sodium carbonate dissociates into sodium ions ($\text{Na}^+$) and carbonate ions ($\text{CO}_3^{2-}$). The carbonate ion undergoes hydrolysis by reacting with water molecules, producing carbonic acid (a weak acid) and hydroxide ions ($\text{OH}^-$): $$\text{CO}_3^{2-}(aq) + 2\text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{CO}_3(aq) + 2\text{OH}^-(aq)$$ The generation of excess hydroxide ($\text{OH}^-$) ions increases the pH of the solution above 7.0, making it alkaline/basic.

First, sodium hydrogen carbonate (baking soda) is thermally decomposed (heated) to yield anhydrous sodium carbonate (soda ash), water, and carbon dioxide: $$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$ Basicity Difference:

• $\text{NaHCO}_3$ (Sodium hydrogen carbonate) is a mildly basic salt (pH ~ 8.3) because it contains a replaceable hydrogen atom and is only slightly alkaline.
• $\text{Na}_2\text{CO}_3$ (Sodium carbonate) is a strongly basic salt (pH ~ 11.5) due to the complete hydrolysis of the carbonate ion, which yields a much higher concentration of hydroxide ($\text{OH}^-$) ions in water.

Thermal decomposition equation: $$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$ Role of Carbon Dioxide: The carbon dioxide ($\text{CO}_2$) gas evolved during heating forms small pockets of gas inside the dough. As the gas expands, it escapes, causing the dough to rise. This makes the bread or cake soft, spongy, and fluffy.

1. Why does he do so? Fresh milk has a pH of about 6.0 (slightly acidic). Over time, bacteria in the milk ferment lactose to produce lactic acid, which lowers the pH, curdles the milk, and turns it sour. By adding a small amount of baking soda ($\text{NaHCO}_3$), the milkman shifts the pH of the fresh milk to be slightly alkaline (above 7). This neutralises any lactic acid as soon as it is formed.
2. How does it affect quality? It prevents the milk from souring quickly, thereby increasing its shelf life. However, this milk takes a longer time to set as curd because lactic acid bacteria must first neutralise the added baking soda before they can lower the pH enough to coagulate the milk proteins.
3. What will happen if excess baking soda is added? The milk will taste soapy or alkaline, and the curdling process will be completely inhibited. When heated, the excess baking soda may decompose, releasing carbon dioxide, which can cause the milk to foam up excessively.

Electrolysis of Brine (Chlor-alkali Process): When concentrated sodium chloride (brine) is electrolysed, water molecules and sodium/chloride ions react to yield sodium hydroxide, chlorine gas, and hydrogen gas: $$2\text{NaCl}(aq) + 2\text{H}_2\text{O}(l) \xrightarrow{\text{Electrolysis}} 2\text{NaOH}(aq) + \text{Cl}_2(g)\uparrow + \text{H}_2(g)\uparrow$$ Products at electrodes:

• At the Anode (+): Chlorine gas ($\text{Cl}_2$) is evolved. $$2\text{Cl}^-(aq) \rightarrow \text{Cl}_2(g) + 2e^-$$
• At the Cathode (-): Hydrogen gas ($\text{H}_2$) is evolved. $$2\text{H}_2\text{O}(l) + 2e^- \rightarrow \text{H}_2(g) + 2\text{OH}^-(aq)$$
• Sodium hydroxide ($\text{NaOH}$) solution forms near the cathode.

Equation for formation: $$\text{H}_2(g) + \text{Cl}_2(g) \rightarrow 2\text{HCl}(g)$$ $$\text{HCl}(g) + \text{H}_2\text{O}(l) \rightarrow \text{HCl}(aq) \text{ (Hydrochloric acid)}$$ Two uses of Hydrochloric acid:

1. For pickling/cleaning steel sheets (removing iron rust) before galvanisation or tin-plating.
2. In the laboratory as a strong acid reagent, and industrially for the synthesis of ammonium chloride, PVC, and medicines.

1. Baking soda is heated: It decomposes into sodium carbonate, water vapour, and carbon dioxide gas. $$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$
2. Washing soda is dissolved in water: It dissociates and the carbonate ions hydrolyse to form a basic solution containing carbonate, sodium, and hydroxide ions. $$\text{Na}_2\text{CO}_3(s) + 10\text{H}_2\text{O}(l) \rightarrow \text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}(aq)$$ $$\text{CO}_3^{2-}(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{HCO}_3^-(aq) + \text{OH}^-(aq)$$
3. Plaster of Paris sets hard: It reacts with water to form a hard crystalline solid mass of gypsum. $$\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}(s) + 1\frac{1}{2}\text{H}_2\text{O}(l) \rightarrow \text{CaSO}_4 \cdot 2\text{H}_2\text{O}(s)$$

1. Balanced equation: $$\text{Mg(OH)}_2(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + 2\text{H}_2\text{O}(l)$$
2. Why basic? General principle: Antacids are weak bases because their function is to neutralise the excess hydrochloric acid produced in the stomach. The general principle is a neutralisation reaction, where the hydroxide ions ($\text{OH}^-$) of the weak base combine with the hydronium ions ($\text{H}_3\text{O}^+$) of the excess acid to form neutral water: $$\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$$ This neutralises the acid and brings the stomach pH back to normal, reducing pain and irritation.
3. Why not in excess? Risks: Taking antacids in excess can make the stomach environment alkaline (pH > 4). This triggers a negative feedback mechanism called acid rebound, where the stomach is stimulated to secrete even more acid to restore its natural acidic environment. Furthermore, high pH in the stomach deactivates pepsin, leading to poor protein digestion and stomach cramps.

1. Why acidic soil is harmful: Highly acidic soil (pH = 4) is toxic to wheat plants. Acidic conditions prevent plants from absorbing essential nutrients (like phosphorus and calcium) from the soil, stunt root growth, and kill beneficial soil micro-organisms, leading to poor crop yield.
2. Quicklime reaction and slaked lime help: Quicklime reacts with soil moisture to form slaked lime: $$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq)$$ Slaked lime is basic. It dissolves in the soil water, releasing calcium and hydroxide ions which neutralise the excess hydronium ions in the acidic soil, thereby raising the soil pH to a healthy range (6.0–7.0).
3. Would excess quicklime always be beneficial? No, excess quicklime will make the soil too basic/alkaline (pH > 8). Highly alkaline soil is just as harmful as acidic soil, as it locks up micronutrients like iron and manganese, preventing plants from absorbing them and causing crop failure. Soil pH must be monitored and kept balanced.

1. Active disinfectant component: The active disinfectant is hypochlorous acid ($\text{HOCl}$) and chlorine gas ($\text{Cl}_2$) released when bleaching powder reacts with water or dissolved $\text{CO}_2$: $$\text{CaOCl}_2(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Cl}_2(g)$$ $$\text{Cl}_2(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{HOCl}(aq) + \text{HCl}(aq)$$ Hypochlorous acid ($\text{HOCl}$) penetrates and kills bacteria and pathogens in the water.
2. Why maintain pH 7.2–7.8?
   • This range is comfortable for human eyes and skin (matches eye fluid pH ~ 7.4).
   • At this pH, the disinfection action of hypochlorous acid is highly effective.
   • If pH drops below 7: The water becomes acidic, causing severe irritation, itching, and redness in the swimmers' eyes and skin. Acidic water also corrodes metal pipes and pumps in the pool filtration system.
3. Why bleaching powder is not suitable for long-term storage: Bleaching powder slowly undergoes autoxidation and reacts with atmospheric moisture and carbon dioxide over time, releasing chlorine gas and converting into calcium carbonate: $$\text{CaOCl}_2(s) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s) + \text{Cl}_2(g)\uparrow$$ This causes the bleaching powder to lose its chlorine content and active disinfecting capacity.

1. Main gases and acid rain formation:
   • Sulphur dioxide ($\text{SO}_2$): $$2\text{SO}_2(g) + \text{O}_2(g) \rightarrow 2\text{SO}_3(g)$$ $$\text{SO}_3(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_4(aq) \text{ (Sulphuric acid)}$$
   • Nitrogen dioxide ($\text{NO}_2$): $$4\text{NO}_2(g) + 2\text{H}_2\text{O}(l) + \text{O}_2(g) \rightarrow 4\text{HNO}_3(aq) \text{ (Nitric acid)}$$
2. Corrosion of marble ($\text{CaCO}_3$) equation: $$\text{CaCO}_3(s) \text{ (Marble)} + \text{H}_2\text{SO}_4(aq) \rightarrow \text{CaSO}_4(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$$ The marble reacts with sulphuric acid to form soluble calcium sulphate, eroding the statue (marble cancer).
3. Two prevention methods:
   1. Install wet scrubbers (containing basic $\text{Ca(OH)}_2$ or $\text{CaCO}_3$ solutions) in factory chimneys to neutralise and trap $\text{SO}_2$ gas before it is released.
   2. Use cleaner, green fuels (like CNG, LPG, solar energy) and fit vehicles with catalytic converters to reduce $\text{NO}_x$ emissions.

1. Chemical name and formula of POP: Calcium sulphate hemihydrate, $\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$.
   Setting reaction equation: $$\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}(s) + 1\frac{1}{2}\text{H}_2\text{O}(l) \rightarrow \text{CaSO}_4 \cdot 2\text{H}_2\text{O}(s) \text{ (Gypsum)}$$
2. Slight expansion & utility: When POP sets into gypsum, it undergoes a slight expansion in volume (about 1%). This expansion ensures that the plaster fits tightly around the fractured leg, filling any small gaps and providing absolute immobilisation to the fractured bones so they can heal in their correct alignment.
3. POP storage moisture warning: POP must be kept in moisture-proof containers because it absorbs atmospheric moisture easily, transforming into hard gypsum. Once it turns into gypsum, it loses its plastering properties and cannot be mixed with water to form paste again.

1. Fizzing indicator & gas: The rapid fizzing indicates a chemical reaction accompanied by the evolution of a gas. The gas produced is carbon dioxide ($\text{CO}_2$).
2. Colour change of blue litmus to red: This confirms that the resulting solution is acidic in nature due to the presence of citric acid (Vitamin C is ascorbic acid) dissolved in water.
3. Common ingredient & reaction: Effervescent tablets contain sodium hydrogen carbonate ($\text{NaHCO}_3$, baking soda) and a solid organic acid (such as citric acid). In dry tablet form, they do not react. When dropped into water, they dissolve, allowing $\text{H}^+$ ions from the acid to react with the bicarbonate ions, producing $\text{CO}_2$ gas: $$\text{H}^+(aq) + \text{HCO}_3^-(aq) \rightarrow \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$$ Or using citric acid: $$3\text{NaHCO}_3(aq) + \text{H}_3\text{C}_6\text{H}_5\text{O}_7(aq) \rightarrow \text{Na}_3\text{C}_6\text{H}_5\text{O}_7(aq) + 3\text{H}_2\text{O}(l) + 3\text{CO}_2(g)\uparrow$$

1. Identification of X, Y, Z:
   • Substance 'X' is Sodium hydrogen carbonate ($\text{NaHCO}_3$, Baking Soda), which is basic and turns moist litmus paper blue.
   • Gas 'Y' is Carbon dioxide ($\text{CO}_2$), which turns lime water milky.
   • Solid 'Z' is Sodium carbonate ($\text{Na}_2\text{CO}_3$).
2. Balanced chemical equation: $$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$
3. Two important uses of 'X' ($\text{NaHCO}_3$):
   1. As an active ingredient in baking powder to make cakes soft and spongy.
   2. As an antacid to neutralise hyperacidity in the stomach, and in soda-acid fire extinguishers.

The given pH values are: Gastric juice (1.2), Lemon juice (2.5), Pure water (7.0), Blood (7.4), Baking soda solution (8.3).
Since pH is inversely related to $\text{H}^+$ concentration, the increasing order of $\text{H}^+$ concentration is: $$\text{Baking soda solution (pH = 8.3)} < \text{Blood (pH = 7.4)} < \text{Pure water (pH = 7.0)} < \text{Lemon juice (pH = 2.5)} < \text{Gastric juice (pH = 1.2)}$$ In other words, the basic baking soda solution has the lowest $\text{H}^+$ ion concentration, while the highly acidic gastric juice has the highest $\text{H}^+$ concentration.

1. Vinegar (weakly acidic, pH ~ 3): Turns orange/yellowish-orange.
2. Ammonia solution (weakly basic, pH ~ 10): Turns blue.
3. Distilled water (neutral, pH = 7): Turns green.
4. Sodium hydroxide solution (strongly basic, pH ~ 14): Turns purple / violet.
5. Orange juice (moderately acidic, pH ~ 3.5): Turns orange-red / orange.

1. Why zinc protects iron better than paint: Paint only acts as a physical barrier. If the paint coat is scratched, the exposed iron rusts quickly. Coating with zinc (galvanisation) protects iron in two ways:
   • It acts as a physical barrier.
   • If the zinc coating is scratched, zinc acts as a sacrificial anode because zinc is more reactive than iron. Zinc oxidises in preference to iron: $$\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$$ This prevents iron from rusting even when exposed directly.
2. Step of paper manufacturing where NaOH is used: Sodium hydroxide is used in the pulping process (cooking step) where wood chips are boiled under pressure with $\text{NaOH}$ and sodium sulphide. This chemical treatment dissolves lignin, which holds the cellulose fibres together, leaving behind pure cellulose wood pulp to be turned into paper.

When acids are dissolved in water (aqueous solution), the polar water molecules pull the acid molecules apart (ionisation), yielding free-moving hydrogen ions (which associate with water to form hydronium ions $\text{H}_3\text{O}^+$) and anions: $$\text{HCl}(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$$ These mobile ions act as charge carriers, allowing electric current to flow.
Ions produced: Acids produce hydronium ions ($\text{H}_3\text{O}^+$ or $\text{H}^+(aq)$) in water.

Property	Strong Acids	Weak Acids
Extent of Ionisation	Completely ionises in aqueous solution.	Partially ionises in aqueous solution.
H? Ion Concentration	Very high concentration of $\text{H}^+$ ions.	Low concentration of $\text{H}^+$ ions.
Reaction Rate	Reacts very rapidly with metals/carbonates.	Reacts slowly with metals/carbonates.
Examples	Hydrochloric acid ($\text{HCl}$), Sulphuric acid ($\text{H}_2\text{SO}_4$).	Acetic acid ($\text{CH}_3\text{COOH}$), Citric acid ($\text{H}_3\text{C}_6\text{H}_5\text{O}_7$).

Which completely ionises? Strong acids completely ionise in water.

Initially, $\text{CO}_2$ gas reacts with sodium hydroxide to form sodium carbonate and water: $$2\text{NaOH}(aq) + \text{CO}_2(g) \rightarrow \text{Na}_2\text{CO}_3(aq) + \text{H}_2\text{O}(l)$$ If CO2 is passed in excess:
The formed sodium carbonate reacts further with excess carbon dioxide and water to produce soluble sodium hydrogen carbonate (baking soda): $$\text{Na}_2\text{CO}_3(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \rightarrow 2\text{NaHCO}_3(aq)$$

1. Preparation of $\text{NaHCO}_3$ from $\text{Na}_2\text{CO}_3$: Pass carbon dioxide gas through a saturated, cold aqueous solution of sodium carbonate: $$\text{Na}_2\text{CO}_3(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \rightarrow 2\text{NaHCO}_3(s)\downarrow$$ Condition: Highly concentrated / saturated cold sodium carbonate solution is required so that the less-soluble baking soda precipitates out as a white crystalline solid.
2. Conversion of $\text{NaHCO}_3$ to $\text{Na}_2\text{CO}_3$: Heat dry solid sodium hydrogen carbonate: $$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$ Condition: Strong direct heating (calcination temperature ~ 100°C or above) is required to drive off water vapour and carbon dioxide.

• Benzene is a non-polar organic solvent. When $\text{HCl}$ gas is dissolved in benzene, it does not ionise or break its covalent bonds. It remains in molecular form ($\text{HCl}$ molecules). Since there are no mobile ions to carry electric charge, the solution does not conduct electricity.
• Water is a highly polar solvent. When $\text{HCl}$ gas is dissolved in water, the strong hydrogen bonding and high dielectric constant of water molecules pull the $\text{HCl}$ molecules apart, causing complete ionisation to form hydronium and chloride ions: $$\text{HCl}(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$$ These mobile ions act as charge carriers, allowing the aqueous solution to conduct electricity.

Oxide	Classification	Chemical Nature
(a) $\text{SO}_2$	Acidic Oxide	Non-metal oxide; reacts with water to form acid.
(b) $\text{Na}_2\text{O}$	Basic Oxide	Metal oxide; reacts with water to form alkali.
(c) $\text{CO}$	Neutral Oxide	Non-metal oxide; does not react with water, acids, or bases.
(d) $\text{CaO}$	Basic Oxide	Metal oxide; reacts with water to form slaked lime.
(e) $\text{NO}_2$	Acidic Oxide	Non-metal oxide; reacts with water to form acid rain.

Equations for reaction of acidic oxides with water:

1. $$\text{SO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{H}_2\text{SO}_3(aq) \text{ (Sulphurous acid)}$$
2. $$2\text{NO}_2(g) + \text{H}_2\text{O}(l) \rightarrow \text{HNO}_3(aq) + \text{HNO}_2(aq) \text{ (Nitric and nitrous acids)}$$

Chemical evaluation: Chemically, the suggestion is reasonable. Gout is caused by hyperuricemia, where insoluble monosodium urate crystals deposit in joints due to high uric acid levels in the blood. Drinking a mild base like dilute baking soda ($\text{NaHCO}_3$) helps neutralise excess uric acid and increases the pH of the urine and blood slightly. This alkalisation increases the solubility of uric acid, facilitating its excretion through the kidneys.
Risks of excess consumption:

1. Metabolic Alkalosis: Consuming excess sodium bicarbonate can raise the blood pH above normal levels (alkalosis), causing muscle twitching, nausea, and confusion.
2. High Sodium Intake (Hypertension): Baking soda contains high levels of sodium. Excess intake can cause fluid retention, swelling, and dramatically raise blood pressure (hypertension), which is highly dangerous for heart patients.