Focus Keyword: Class 10 Science Chapter 2 Notes Acids Bases and Salts
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H1 Title: Class 10 Science Chapter 2 Notes: Acids, Bases and Salts (Ultimate 2026-2027 Board Revision Guide)
Navigating through the CBSE Class 10 Science curriculum requires an in-depth understanding of electrolytic dissociation, hydronium and hydroxide dynamics, indicator responses, and industrial salt synthesis. Chapter 2 of Class 10 Science, “Acids, Bases and Salts”, forms the foundation of modern chemical analysis, physiological pH regulation, aqueous solution thermodynamics, and industrial manufacturing. It investigates the ionic behavior of acidic and alkaline solutions; explores the quantitative calibration of acidity via the pH scale; and details the extraction and commercial synthesis of essential salts such as sodium hydroxide, bleaching powder, baking soda, washing soda, and plaster of Paris. To help students master every aspect of this high-weightage chapter, this comprehensive guide offers textbook-accurate, highly structured, and pedagogically sound responses strictly aligned with the latest CBSE evaluation standards.
Every concept presented in the official NCERT textbook—ranging from microscopic chemical definitions and an exhaustive breakdown of all 15 practical laboratory activities to zero-mark-loss balanced chemical equations, authentic Wikimedia visual references, and a master revision cheat sheet—has been compiled with extreme precision. Key scoring terms, state symbols, chromatic shifts, and examiner cautions have been highlighted to ensure students secure maximum marks in their CBSE Board Examinations.
Chapter 2: Acids, Bases and Salts
Master Concept Overview & Thematic Architecture
Acids are substances that dissociate in water to produce hydrogen ions ($\text{H}^+$ or hydronium ions $\text{H}_3\text{O}^+$), turn blue litmus red, and possess a sour taste. Bases dissociate to yield hydroxide ions ($\text{OH}^-$), turn red litmus blue, feel soapy, and taste bitter. Salts are ionic products of neutralisation.
| Chemical Class | Defining Microscopic Species | Characteristic Sensory / Physical Trait | Classical Indicator Response | Prototypical Board Exam Example |
|---|---|---|---|---|
| Acid | Hydronium Ion [$\text{H}_3\text{O}^+(aq)$] | Sour taste, electrolytic conductor in aqueous media | Turns blue litmus red; phenolphthalein remains colourless | Hydrochloric Acid: $\text{HCl}(aq)$ |
| Base | Hydroxide Ion [$\text{OH}^-(aq)$] | Bitter taste, soapy/slippery feel, electrolytic conductor | Turns red litmus blue; phenolphthalein turns deep pink | Sodium Hydroxide: $\text{NaOH}(aq)$ |
| Alkali | Dissolved $\text{OH}^-(aq)$ | Water-soluble base; corrosive to skin in concentrated form | Turns red litmus blue; methyl orange turns yellow | Potassium Hydroxide: $\text{KOH}(aq)$ |
| Neutral Salt | Neither excess $\text{H}_3\text{O}^+$ nor $\text{OH}^-$ | High melting point, crystalline lattice, neutral $pH \approx 7$ | No colour change on red or blue litmus paper | Sodium Chloride: $\text{NaCl}(s)$ |
| Acidic Salt | Contains excess $\text{H}_3\text{O}^+$ from strong acid parent | Hydrolyses in water to generate acidic solution ($pH < 7$) | Turns blue litmus red | Ammonium Chloride: $\text{NH}_4\text{Cl}(aq)$ |
| Basic Salt | Contains excess $\text{OH}^-$ from strong base parent | Hydrolyses in water to generate alkaline solution ($pH > 7$) | Turns red litmus blue | Sodium Carbonate: $\text{Na}_2\text{CO}_3(aq)$ |
🧠 Examiner’s Secret: All alkalis are bases, but not all bases are alkalis. An alkali is strictly a base that dissolves completely in water (e.g., $\text{NaOH}$, $\text{KOH}$, $\text{Ca(OH)}_2$). Insoluble bases like Copper(II) oxide [$\text{CuO}$] or Ferric hydroxide [$\text{Fe(OH)}_3$] are bases, but they are never classified as alkalis.
Concept Map: Structural Overview of Acids, Bases, and Salts
[mermaid]
graph TD
A[Matter in Aqueous Solution] –> B[Acids: Furnish H3O+ Ions]
A –> C[Bases: Furnish OH- Ions]
A –> D[Salts: Product of Neutralisation]
B --> B1[Organic Acids: Weak e.g., CH3COOH]
B --> B2[Mineral Acids: Strong e.g., HCl, H2SO4]
C --> C1[Water Soluble: Alkalis e.g., NaOH, KOH]
C --> C2[Water Insoluble: e.g., CuO, Fe OH 3]
B & C --> E[Neutralisation Reaction]
E --> D
D --> D1[Common Salt: NaCl]
D1 --> F[Chlor-Alkali Process]
F --> F1[NaOH at Cathode]
F --> F2[Cl2 at Anode]
F --> F3[H2 at Cathode]
D1 --> G[Industrial Compounds]
G --> G1[Bleaching Powder: CaOCl2]
G --> G2[Baking Soda: NaHCO3]
G --> G3[Washing Soda: Na2CO3.10H2O]
G --> G4[Plaster of Paris: CaSO4.0.5H2O]
[/mermaid]
Classification of Chemical Indicators
An indicator is a chemical substance or biological pigment that exhibits distinct colour or olfactory changes depending on the acidic or basic nature of the medium.
Indicators provide an immediate qualitative test for the presence of acidic or alkaline species in an unknown solution.
| Indicator Class | Specific Indicator Agent | Original Native Colour | Colour / Sensation in Acidic Medium | Colour / Sensation in Basic Medium |
|---|---|---|---|---|
| Natural Indicator | Litmus Solution (Lichen extract) | Purple | Red | Blue |
| Natural Indicator | Red Cabbage Leaf Extract | Purple | Red / Pinkish-Red | Green / Yellowish-Green |
| Natural Indicator | Turmeric Paste / Extract | Yellow | Yellow (No Change) | Deep Reddish-Brown |
| Synthetic Indicator | Phenolphthalein | Colourless | Colourless | Vibrant Pink / Magenta |
| Synthetic Indicator | Methyl Orange | Orange | Red / Pinkish-Red | Yellow |
| Olfactory Indicator | Vanilla Essence | Sweet Pleasant Odour | Retains characteristic odour | Odour completely destroyed / disappears |
| Olfactory Indicator | Onion Extract | Pungent Sharp Odour | Retains characteristic odour | Odour completely destroyed / disappears |
| Olfactory Indicator | Clove Oil | Characteristic Herbal Odour | Retains characteristic odour | Odour completely destroyed / disappears |
💡 Did You Know?: The bright yellow stain of a curry spill on a white cotton shirt turns reddish-brown when scrubbed with laundry soap. This occurs because commercial soap is alkaline in nature, which causes the turmeric pigment (curcumin) to change color from yellow to reddish-brown. When the garment is rinsed with copious amounts of water, the alkali is diluted, and the yellow color returns.
Chemical Properties of Acids and Bases
Reaction with Metals
Acids react with reactive metals situated above hydrogen in the electrochemical activity series to yield a corresponding metal salt and liberate hydrogen gas.
$$\text{Acid} + \text{Metal} \longrightarrow \text{Salt} + \text{Hydrogen Gas}\uparrow$$
- With Dilute Hydrochloric Acid:
$$\text{Zn}(s) + 2\text{HCl}(aq) \longrightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)\uparrow$$ - With Dilute Sulphuric Acid:
$$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)\uparrow$$
Bases also react with specific amphoteric metals (such as Zinc and Aluminium) to release hydrogen gas while forming complex oxo-salts:
$$\text{Strong Base} + \text{Amphoteric Metal} \longrightarrow \text{Complex Salt} + \text{Hydrogen Gas}\uparrow$$
- With Sodium Hydroxide:
$$2\text{NaOH}(aq) + \text{Zn}(s) \xrightarrow{\Delta} \text{Na}_2\text{ZnO}_2(aq) [\text{Sodium Zincate}] + \text{H}_2(g)\uparrow$$
$$2\text{NaOH}(aq) + 2\text{Al}(s) + 2\text{H}_2\text{O}(l) \longrightarrow 2\text{NaAlO}_2(aq) [\text{Sodium Aluminate}] + 3\text{H}_2(g)\uparrow$$
🧠 Examiner’s Secret: Nitric acid ($\text{HNO}_3$) generally does not liberate hydrogen gas when reacted with metals. Because $\text{HNO}_3$ is a powerful oxidising agent, it oxidises the evolved $\text{H}_2$ gas into liquid water ($\text{H}_2\text{O}$) while itself getting reduced to nitrogen oxides ($\text{NO}_2$, $\text{NO}$, or $\text{N}_2\text{O}$). Only Magnesium ($\text{Mg}$) and Manganese ($\text{Mn}$) react with extremely dilute ($1%$) $\text{HNO}_3$ to liberate $\text{H}_2$ gas:
$$\text{Mg}(s) + 2\text{HNO}_3(\text{very dil.}) \longrightarrow \text{Mg(NO}_3)_2(aq) + \text{H}_2(g)\uparrow$$
Reaction of Acids with Metal Carbonates and Metal Hydrogen Carbonates
Acids react vigorously with metal carbonates and metal hydrogen carbonates to produce a metal salt, liquid water, and brisk effervescence of carbon dioxide gas.
$$\text{Metal Carbonate} + \text{Acid} \longrightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}\uparrow$$
$$\text{Metal Hydrogen Carbonate} + \text{Acid} \longrightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}\uparrow$$
- Reaction with Sodium Carbonate:
$$\text{Na}_2\text{CO}_3(s) + 2\text{HCl}(aq) \longrightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$$ - Reaction with Sodium Hydrogen Carbonate:
$$\text{NaHCO}_3(s) + \text{HCl}(aq) \longrightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$$
The Two-Stage Confirmatory Lime Water Test for Carbon Dioxide:
- Initial Passage of $\text{CO}_2$ Gas (Lime Water Turns Milky):
When carbon dioxide gas is bubbled through clear calcium hydroxide solution (limewater), an insoluble white precipitate of calcium carbonate forms:
$$\text{Ca(OH)}_2(aq) [\text{Clear Limewater}] + \text{CO}_2(g) \longrightarrow \text{CaCO}_3(s)\downarrow [\text{White ppt / Milky}] + \text{H}_2\text{O}(l)$$ - Excess Passage of $\text{CO}_2$ Gas (Milkiness Disappears):
On prolonged bubbling of excess carbon dioxide, the insoluble calcium carbonate converts into water-soluble calcium hydrogen carbonate, causing the milkiness to clear completely:
$$\text{CaCO}_3(s) + \text{H}_2\text{O}(l) + \text{CO}_2(\text{excess}, g) \longrightarrow \text{Ca(HCO}_3)_2(aq) [\text{Soluble / Colourless}]$$
Neutralisation Reaction
A neutralisation reaction is a chemical process where hydronium ions from an acid combine with hydroxide ions from a base to form water and an ionic salt.
$$\text{Acid} + \text{Base} \longrightarrow \text{Salt} + \text{Water} \quad (\Delta H < 0, \text{Exothermic})$$
- Prototypical Molecular Equations:
$$\text{HCl}(aq) + \text{NaOH}(aq) \longrightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$
$$\text{H}_2\text{SO}_4(aq) + 2\text{KOH}(aq) \longrightarrow \text{K}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l)$$ - Net Ionic Essence of Neutralisation:
$$\text{H}^+(aq) + \text{OH}^-(aq) \longrightarrow \text{H}_2\text{O}(l) \quad \text{or} \quad \text{H}_3\text{O}^+(aq) + \text{OH}^-(aq) \longrightarrow 2\text{H}_2\text{O}(l)$$
Reactions of Metallic and Non-Metallic Oxides
Reaction of Metallic Oxides with Acids (Basic Oxides):
Metallic oxides react with dilute mineral acids to form a salt and water, mirroring the behavior of standard bases. This confirms that metallic oxides are basic in nature:
$$\text{Metal Oxide} + \text{Acid} \longrightarrow \text{Salt} + \text{Water}$$
- Reaction of Black Copper(II) Oxide with Hydrochloric Acid:
$$\text{CuO}(s) [\text{Black}] + 2\text{HCl}(aq) \longrightarrow \text{CuCl}_2(aq) [\text{Blue-Green Solution}] + \text{H}_2\text{O}(l)$$
(The black powder dissolves, and the solution turns bluish-green due to the formation of hydrated copper(II) chloride ions).
Reaction of Non-Metallic Oxides with Bases (Acidic Oxides):
Non-metallic oxides react with alkaline solutions to produce a salt and water, mirroring the behavior of standard acids. This confirms that non-metallic oxides are acidic in nature:
$$\text{Non-Metal Oxide} + \text{Base} \longrightarrow \text{Salt} + \text{Water}$$
- Reaction of Carbon Dioxide with Calcium Hydroxide:
$$\text{CO}_2(g) + \text{Ca(OH)}_2(aq) \longrightarrow \text{CaCO}_3(s)\downarrow + \text{H}_2\text{O}(l)$$ - Reaction of Sulphur Dioxide with Sodium Hydroxide:
$$\text{SO}_2(g) + 2\text{NaOH}(aq) \longrightarrow \text{Na}_2\text{SO}_3(aq) [\text{Sodium Sulphite}] + \text{H}_2\text{O}(l)$$
[👉 Also Read: Class 10 Science Chapter 1 Chemical Reactions and Equations Notes]
Exhaustive NCERT Practical Laboratory Activities Vault
Activity 2.1: Testing Acids and Bases with Indicators
- Aim: To observe the colorimetric shifts of common acids and bases using various indicators.
- Setup: Test samples of $\text{HCl}, \text{H}_2\text{SO}_4, \text{HNO}_3, \text{CH}_3\text{COOH}, \text{NaOH}, \text{KOH}, \text{Ca(OH)}_2, \text{Mg(OH)}_2,$ and $\text{NH}_4\text{OH}$ placed in clean watch glasses against red litmus, blue litmus, phenolphthalein, and methyl orange.
- Observations & Master Summary Table:
| Tested Solution | Red Litmus Paper | Blue Litmus Paper | Phenolphthalein Solution | Methyl Orange Solution | Chemical Nature |
|---|---|---|---|---|---|
| $\text{HCl}$ / $\text{H}_2\text{SO}_4$ / $\text{HNO}_3$ | No Change | Turns Red | Remains Colourless | Turns Red | Strong Acid |
| $\text{CH}_3\text{COOH}$ | No Change | Turns Red | Remains Colourless | Turns Pinkish-Red | Weak Acid |
| $\text{NaOH}$ / $\text{KOH}$ | Turns Blue | No Change | Turns Deep Pink | Turns Yellow | Strong Base |
| $\text{Ca(OH)}_2$ / $\text{Mg(OH)}_2$ | Turns Blue | No Change | Turns Pink | Turns Yellow | Mild Base |
| $\text{NH}_4\text{OH}$ | Turns Blue | No Change | Turns Pink | Turns Yellow | Weak Base |
- Board Conclusion: Litmus, phenolphthalein, and methyl orange produce distinct, repeatable color changes that reliably distinguish acidic species from alkaline species.
Activity 2.2: Olfactory Indicator Characterisation
- Aim: To evaluate the effect of acidic and basic solutions on the odor of olfactory indicators.
- Setup: Finely chop onions and place them in a sealed plastic bag along with strips of clean cloth overnight to absorb the onion vapor. Test these cloth strips, as well as vanilla essence and clove oil, against dilute $\text{HCl}$ and dilute $\text{NaOH}$ solutions.
- Observations:
- The cloth strips treated with dilute $\text{HCl}$ maintain their strong, characteristic onion odor.
- The cloth strips treated with dilute $\text{NaOH}$ lose their onion smell completely.
- Vanilla essence and clove oil retain their characteristic scents when mixed with acidic solutions, but their odors are destroyed when mixed with basic solutions.
- Board Conclusion: Olfactory indicators retain their natural scent in acidic media but lose it entirely in alkaline media. This makes them ideal for visually impaired students to distinguish acids from bases.
Activity 2.3: Displacement of Hydrogen from Acids by Zinc
- Aim: To observe the evolution and properties of hydrogen gas produced by the reaction of zinc with dilute sulphuric acid.
- Setup: Place about $5\text{ g}$ of granulated zinc in a test tube and add $5\text{ mL}$ of dilute $\text{H}_2\text{SO}_4$. Fit the tube with a delivery tube directed into a trough filled with liquid soap solution.
- Observations:
- Vigorous effervescence occurs around the surface of the zinc granules in the reaction tube.
- Gas bubbles pass through the delivery tube and emerge in the soap solution, forming light, gas-filled soap bubbles that rise into the air.
- When a burning candle is brought near a floating soap bubble, the bubble bursts and burns with a sharp pop sound.
- Balanced Chemical Equation:
$$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)\uparrow$$ - Board Conclusion: Active metals displace hydrogen from dilute mineral acids. The low density of hydrogen gas causes the soap bubbles to rise, and its flammability produces the characteristic pop sound when ignited.
Activity 2.4: Reaction of Zinc with a Strong Base
- Aim: To observe the liberation of hydrogen gas during the reaction between zinc and sodium hydroxide.
- Setup: Place granulated zinc in a test tube, add $2\text{ mL}$ of concentrated sodium hydroxide solution, and heat the mixture gently over a burner flame.
- Observations:
Bubbles of gas evolve slowly from the surface of the heated zinc granules. Passing this gas through soap solution and bringing a flame near the bubbles confirms the presence of hydrogen by the characteristic pop sound. - Balanced Chemical Equation:
$$2\text{NaOH}(aq) + \text{Zn}(s) \xrightarrow{\Delta} \text{Na}_2\text{ZnO}_2(aq) [\text{Sodium Zincate}] + \text{H}_2(g)\uparrow$$ - Board Conclusion: Zinc is an amphoteric metal that reacts with both acids and strong bases to liberate hydrogen gas. However, such reactions do not occur with all metals.
Activity 2.5: Reaction of Carbonates with Acid and the Lime Water Test
- Aim: To evaluate the gas evolved when metal carbonates and hydrogen carbonates react with hydrochloric acid.
- Setup: Place about $0.5\text{ g}$ of $\text{Na}_2\text{CO}_3$ in Test Tube A and $0.5\text{ g}$ of $\text{NaHCO}_3$ in Test Tube B. Add $2\text{ mL}$ of dilute $\text{HCl}$ to each, and channel the evolved gas through freshly prepared calcium hydroxide solution.
- Observations:
- Rapid, frothing effervescence occurs in both test tubes.
- When the evolved gas enters the limewater, the solution turns milky white.
- When the gas is passed continuously for several minutes, the milky precipitate dissolves, leaving a clear solution.
- Balanced Chemical Equations:
- Tube A: $\text{Na}_2\text{CO}_3(s) + 2\text{HCl}(aq) \longrightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$
- Tube B: $\text{NaHCO}_3(s) + \text{HCl}(aq) \longrightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$
- Limewater Initial: $\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \longrightarrow \text{CaCO}_3(s)\downarrow [\text{White ppt}] + \text{H}_2\text{O}(l)$
- Limewater Excess: $\text{CaCO}_3(s) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \longrightarrow \text{Ca(HCO}_3)_2(aq) [\text{Soluble}]$
- Board Conclusion: All metal carbonates and bicarbonates react with acids to liberate carbon dioxide gas, which turns limewater milky and subsequently clears upon exposure to excess gas.
Activity 2.6: Acid-Base Neutralisation with Phenolphthalein
- Aim: To demonstrate reversible acid-base neutralisation using an indicator.
- Setup: Add $2\text{ mL}$ of dilute $\text{NaOH}$ to a test tube, add two drops of phenolphthalein solution, and then add dilute $\text{HCl}$ dropwise with continuous shaking. Once the color changes, add a few drops of $\text{NaOH}$ again.
- Observations:
- The initial $\text{NaOH}$ solution turns a vibrant deep pink upon adding phenolphthalein.
- Adding dilute $\text{HCl}$ dropwise neutralizes the base, causing the pink color to fade and eventually disappear, leaving a clear solution.
- Adding a few drops of $\text{NaOH}$ restores the pink color.
- Balanced Chemical Equation:
$$\text{NaOH}(aq) + \text{HCl}(aq) \longrightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$ - Board Conclusion: Acids and bases neutralize each other’s characteristic properties. Phenolphthalein acts as a reversible indicator, turning pink in alkaline conditions and remaining colorless in neutral or acidic media.
Activity 2.7: Reaction of Metallic Oxides with Acids
- Aim: To confirm the basic nature of metallic oxides by reacting copper(II) oxide with dilute hydrochloric acid.
- Setup: Place a small amount of black copper(II) oxide in a beaker and slowly add dilute hydrochloric acid while stirring continuously.
- Observations:
The insoluble black powder dissolves, and the solution changes color to a clear blue-green. - Balanced Chemical Equation:
$$\text{CuO}(s) [\text{Black}] + 2\text{HCl}(aq) \longrightarrow \text{CuCl}_2(aq) [\text{Blue-Green}] + \text{H}_2\text{O}(l)$$ - Board Conclusion: Copper(II) oxide reacts with hydrochloric acid to form copper(II) chloride and water, mirroring a standard acid-base reaction and confirming that metallic oxides are basic anhydrides.
Activity 2.8: Electrical Conductivity of Aqueous Solutions
- Aim: To determine whether all hydrogen-containing compounds conduct electricity in aqueous solution.
- Setup: Insert two iron nails into a rubber cork placed inside a $100\text{ mL}$ beaker. Connect the nails to a $6\text{ V}$ battery through a miniature bulb and an on/off switch. Test solutions of glucose ($\text{C}6\text{H}{12}\text{O}_6$), ethyl alcohol ($\text{C}_2\text{H}_5\text{OH}$), dilute hydrochloric acid ($\text{HCl}$), and dilute sulphuric acid ($\text{H}_2\text{SO}_4$).
- Observations & Analytical Comparison:
| Tested Solution | Chemical Nature of Bonding | Free Mobile Ions Present? | Bulb Glows? | Electrical Conduction |
|---|---|---|---|---|
| Dilute $\text{HCl}(aq)$ | Polar Covalent / Ionised | Yes [$\text{H}_3\text{O}^+, \text{Cl}^-$] | Yes, brightly | Conducts electricity |
| Dilute $\text{H}_2\text{SO}_4(aq)$ | Polar Covalent / Ionised | Yes [$\text{H}_3\text{O}^+, \text{SO}_4^{2-}$] | Yes, brightly | Conducts electricity |
| Glucose Solution | Covalent Organic Compound | No (Remains molecular) | No (Bulb remains dark) | Does not conduct |
| Ethyl Alcohol Solution | Covalent Organic Compound | No (Remains molecular) | No (Bulb remains dark) | Does not conduct |
- Board Conclusion: Electrical conduction in liquid solutions requires free, mobile ions. While glucose and alcohol contain hydrogen atoms, they do not ionize in water to yield hydronium ions ($\text{H}_3\text{O}^+$). Therefore, not all hydrogen-containing compounds behave as acids.
Activity 2.9: Production of Dry Hydrogen Chloride Gas and the Role of Water
- Aim: To demonstrate that acids produce hydrogen ions only in the presence of water.
- Setup: Place about $1\text{ g}$ of solid $\text{NaCl}$ in a dry test tube and carefully add $1\text{ mL}$ of concentrated $\text{H}_2\text{SO}_4$. Fit the tube with a dry delivery tube. Test the escaping dry $\text{HCl}$ gas first with completely dry blue litmus paper, and then with moist blue litmus paper.
- Observations:
- Dry blue litmus paper shows no color change when exposed to the dry $\text{HCl}$ gas.
- Moist blue litmus paper turns red immediately upon contact with the gas.
- Balanced Chemical Equations:
- Gas Generation:
$$\text{NaCl}(s) + \text{H}_2\text{SO}_4(\text{conc}.) \longrightarrow \text{NaHSO}_4(s) + \text{HCl}(g)\uparrow$$ - Ionisation in Moisture:
$$\text{HCl}(g) + \text{H}_2\text{O}(l) \longrightarrow \text{H}_3\text{O}^+(aq) + \text{Cl}^-(aq)$$
- Gas Generation:
- Why a Guard Tube with Calcium Chloride ($\text{CaCl}_2$) is Used:
During humid weather, the escaping $\text{HCl}$ gas absorbs atmospheric moisture inside the delivery tube. Passing the gas through a guard tube containing anhydrous calcium chloride absorbs this moisture, ensuring the gas remains completely dry for the test. - Board Conclusion: Hydrogen chloride exhibits acidic properties only after dissolving in water. Liquid water provides the polar environment needed to separate the polar covalent $\text{H}-\text{Cl}$ bond into hydrated hydronium [$\text{H}_3\text{O}^+$] and chloride [$\text{Cl}^-$] ions. Dry $\text{HCl}$ gas contains no free ions and cannot alter the color of dry litmus paper.
Activity 2.10: Dilution Thermodynamics of Concentrated Acids
- Aim: To evaluate the thermal changes that occur during the dilution of concentrated sulphuric acid with water.
- Setup: Measure $10\text{ mL}$ of distilled water into a beaker. Check the initial temperature with a thermometer. Add a few drops of concentrated $\text{H}_2\text{SO}_4$ down the side of the beaker while swirling continuously.
- Observations:
The thermometer reading rises rapidly, and the bottom of the beaker becomes hot to the touch. - Critical Safety Protocol (Standard CBSE Question):
- Always add concentrated acid slowly to water with constant stirring.
- Never add water to concentrated acid.
- Scientific Explanation: The hydration of acid molecules is an extremely exothermic process ($\Delta H \ll 0$). If water is added directly into concentrated acid, the small volume of water boils instantly due to the localized heat release. This causes the mixture to flash into steam, splashing concentrated acid outward onto the user’s face and hands, while the thermal shock can shatter the glass container. When acid is added slowly to a larger volume of water with continuous stirring, the water absorbs and distributes the generated heat safely.
Activity 2.11: Measuring the pH of Common Solutions
- Aim: To measure the pH values of various household and laboratory solutions using universal indicator paper.
- Setup: Spot drops of various solutions onto universal indicator paper strips, observe the resulting color shifts, and match them against the standard colorimetric pH reference chart.
- Observations:
| Solution Sample | Observed Colour on Indicator Paper | Approximate pH Value | Chemical Classification |
|---|---|---|---|
| Dilute $\text{HCl}$ | Deep Red | $1.0$ | Strongly Acidic |
| Dilute $\text{CH}_3\text{COOH}$ | Orange-Red | $3.0$ | Weakly Acidic |
| Lemon Juice | Reddish-Orange | $2.2$ | Moderately Acidic |
| Aerated Soft Drink | Light Orange | $4.0$ | Weakly Acidic |
| Pure Distilled Water | Pure Green | $7.0$ | Strictly Neutral |
| Human Blood | Pale Greenish-Blue | $7.4$ | Slightly Basic / Alkaline |
| Milk of Magnesia [$\text{Mg(OH)}_2$] | Blue | $10.0$ | Moderately Basic |
| Dilute $\text{NaOH}$ | Deep Violet / Purple | $14.0$ | Strongly Basic |
- Board Conclusion: Universal indicator strips provide a semi-quantitative measurement of relative acidity or alkalinity by changing color across the entire $0 – 14$ pH scale.
Activity 2.12: Evaluating the pH of Soil Samples
- Aim: To measure the pH of garden soil to determine its suitability for plant growth.
- Setup: Place about $2\text{ g}$ of soil in a test tube, add $5\text{ mL}$ of distilled water, shake the mixture vigorously, and filter it through filter paper. Test the clear filtrate using universal indicator paper.
- Observations:
Most productive agricultural and garden soils produce a pale green color corresponding to a near-neutral pH range of $6.5 – 7.5$. - Agricultural Remedy:
- If soil is too acidic ($pH < 6.0$): Farmers treat it with bases like Quicklime ($\text{CaO}$), Slaked Lime [$\text{Ca(OH)}_2$], or Chalk ($\text{CaCO}_3$) to neutralize the excess acidity.
- If soil is too basic ($pH > 8.0$): Farmers add decaying organic matter (compost/manure), which releases organic acids to bring the pH back to the optimal range.
Activity 2.13: Water Solubility and Crystal Structure of Common Salts
- Aim: To analyze the chemical formulas and solubility of common salts.
- Setup: Tabulate the chemical formulas, parent acids, and parent bases of common salts including potassium sulphate, sodium sulphate, calcium sulphate, magnesium sulphate, copper sulphate, sodium chloride, sodium nitrate, sodium carbonate, and ammonium chloride.
- Board Conclusion: Salts sharing the same positive metallic cation or negative non-metallic anion belong to the same salt family (e.g., $\text{NaCl}$ and $\text{Na}_2\text{SO}_4$ belong to the family of sodium salts; $\text{NaCl}$ and $\text{KCl}$ belong to the family of chloride salts).
Activity 2.14: Determining the pH of Salt Solutions (Salt Hydrolysis)
- Aim: To evaluate the acidic, basic, or neutral nature of aqueous salt solutions.
- Setup: Dissolve samples of various salts in distilled water and test their pH using universal indicator paper.
- Observations & Salt Hydrolysis Principles:
| Salt Sample | Chemical Formula | Parent Acid [Strength] | Parent Base [Strength] | Solution pH | Net Solution Nature |
|---|---|---|---|---|---|
| Sodium Chloride | $\text{NaCl}$ | $\text{HCl}$ [Strong] | $\text{NaOH}$ [Strong] | $\approx 7.0$ | Neutral |
| Potassium Nitrate | $\text{KNO}_3$ | $\text{HNO}_3$ [Strong] | $\text{KOH}$ [Strong] | $\approx 7.0$ | Neutral |
| Copper Sulphate | $\text{CuSO}_4$ | $\text{H}_2\text{SO}_4$ [Strong] | $\text{Cu(OH)}_2$ [Weak] | $< 7.0$ | Acidic |
| Ammonium Chloride | $\text{NH}_4\text{Cl}$ | $\text{HCl}$ [Strong] | $\text{NH}_4\text{OH}$ [Weak] | $< 7.0$ | Acidic |
| Sodium Carbonate | $\text{Na}_2\text{CO}_3$ | $\text{H}_2\text{CO}_3$ [Weak] | $\text{NaOH}$ [Strong] | $> 7.0$ | Basic |
| Sodium Acetate | $\text{CH}_3\text{COONa}$ | $\text{CH}_3\text{COOH}$ [Weak] | $\text{NaOH}$ [Strong] | $> 7.0$ | Basic |
- Board Rules of Salt Hydrolysis:
- $\text{Strong Acid} + \text{Strong Base} \longrightarrow \text{Neutral Salt Solution } (pH = 7)$
- $\text{Strong Acid} + \text{Weak Base} \longrightarrow \text{Acidic Salt Solution } (pH < 7)$
- $\text{Weak Acid} + \text{Strong Base} \longrightarrow \text{Basic Salt Solution } (pH > 7)$
Activity 2.15: Water of Crystallisation in Copper Sulphate Crystals
- Aim: To demonstrate the presence of water of crystallisation in hydrated copper sulphate and observe its thermal reversibility.
- Setup: Place a few blue crystals of copper sulphate in a dry boiling tube. Heat the tube over a burner flame. After the color changes, allow the tube to cool and add $2 – 3\text{ drops}$ of water onto the remaining white powder.
- Observations:
- Condensation forms as tiny water droplets along the cooler upper inner walls of the boiling tube.
- The brilliant blue color of the crystals fades completely, leaving a dry, powdery white solid.
- When $2 – 3\text{ drops}$ of water are added to the cooled white residue, the brilliant blue color is restored immediately.
- Balanced Chemical Equations:
- Heating (Dehydration):
$$\text{CuSO}_4 \cdot 5\text{H}_2\text{O}(s) [\text{Blue Pentahydrate}] \xrightarrow{\Delta} \text{CuSO}_4(s) [\text{White Anhydrous}] + 5\text{H}_2\text{O}(g)\uparrow$$ - Rehydration:
$$\text{CuSO}_4(s) [\text{White Anhydrous}] + 5\text{H}_2\text{O}(l) \longrightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O}(s) [\text{Blue Restored}]$$
- Heating (Dehydration):
- Board Definition: Water of crystallisation is the fixed number of water molecules chemically bonded within one formula unit of a salt in its crystalline state. Anhydrous copper sulphate is white; its blue color is produced by the coordination of water molecules within the crystal lattice.
Understanding the pH Scale and Everyday Applications
The pH Scale
The pH scale is a logarithmic scale ranging from $0$ to $14$ that measures the concentration of free hydrogen (hydronium) ions present in an aqueous solution.
$$\text{Mathematical Definition: } \text{pH} = -\log_{10}[\text{H}^+] = -\log_{10}[\text{H}_3\text{O}^+]$$
The letter ‘$p$’ stands for the German word potenz, meaning ‘power’, while ‘$H$’ represents hydrogen ion concentration:
- Neutral Solution: $[\text{H}^+] = [\text{OH}^-] = 10^{-7}\text{ M} \implies \text{pH} = 7$
- Acidic Solution: $[\text{H}^+] > 10^{-7}\text{ M} \implies \text{pH} < 7$
- Basic Solution: $[\text{H}^+] < 10^{-7}\text{ M} \implies \text{pH} > 7$
Because the scale is inversely logarithmic, a lower pH indicates a higher concentration of hydronium ions. Each single unit change on the pH scale represents a tenfold ($10\times$) change in hydrogen ion concentration.
Crucial Everyday Roles of pH
1. Physiological Tolerance in Humans and Animals:
Living organisms function efficiently within a narrow physiological pH range of $7.0\text{ to }7.8$. Human blood, tears, and saliva are tightly buffered around $pH \approx 7.4$. Deviations outside this narrow range disrupt enzymatic activity and cellular metabolism.
2. Acid Rain and Aquatic Ecosystems:
Normal rainwater has a slightly acidic pH of approximately $5.6$ due to dissolved atmospheric carbon dioxide forming weak carbonic acid ($\text{H}_2\text{CO}_3$). When industrial air pollutants like sulphur dioxide ($\text{SO}_2$) and nitrogen dioxide ($\text{NO}_2$) dissolve in rainwater, they form sulphuric and nitric acids. When the pH of precipitation drops below $5.6$, it is classified as Acid Rain:
$$\text{SO}_2(g) + \text{H}_2\text{O}(l) \longrightarrow \text{H}_2\text{SO}_3(aq)$$
Acid rain lowers the pH of freshwater lakes and rivers, dissolving toxic aluminium ions from soil into the water and damaging the reproductive cycles and survival of aquatic organisms.
3. Digestion in the Human Stomach:
The stomach lining produces concentrated hydrochloric acid ($\text{HCl}$), establishing a strongly acidic environment ($pH \approx 1.2 – 2.0$). This low pH activates the protein-digesting enzyme pepsin and kills ingested microorganisms. During indigestion, excess acid production causes irritation, heartburn, and pain. This is relieved using mild, non-corrosive basic compounds called Antacids (such as Milk of Magnesia [$\text{Mg(OH)}_2$] or Baking Soda [$\text{NaHCO}_3$]), which neutralize the excess acid:
$$\text{HCl}(aq) + \text{NaHCO}_3(aq) \longrightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$
4. Tooth Decay Prevention:
Tooth enamel is the hardest substance in the human body, composed of a crystalline form of calcium phosphate known as hydroxyapatite [$\text{Ca}_5(\text{PO}_4)_3\text{OH}$]. While insoluble in water, it degrades in acidic media through a process called demineralisation. When oral bacteria ferment residual food particles and sugars, they produce organic acids that lower the oral pH. When oral pH drops below $5.5$, tooth decay begins. Using alkaline toothpastes neutralizes these bacterial acids and promotes enamel remineralisation.
5. Self-Defence Mechanisms in Animals and Plants:
- Honeybee and Ant Stings: Inject Methanoic acid (formic acid, $\text{HCOOH}$), which causes sharp pain, inflammation, and localized burning. Applying a mild base like baking soda paste ($\text{NaHCO}_3$) neutralizes the acid and provides quick relief.
- Wasp Stings: Inject an alkaline liquid. Neutralizing it requires a mild acid, such as vinegar (dilute acetic acid).
- Nettle Plant Leaves: Wild nettle leaves bear stinging hairs that inject methanoic acid upon contact. Nature provides a remedy: the Dock plant, which grows nearby in alkaline soil, has mildly basic sap that neutralizes the sting when rubbed onto the affected area.
[👉 Also Read: Class 10 Science Chapter 3 Metals and Non-Metals Notes]
Chemicals from Common Salt: Industrial Synthesis Vault
The Chlor-Alkali Process
The Chlor-Alkali Process is the industrial electrolysis of a concentrated aqueous solution of sodium chloride (brine) to manufacture sodium hydroxide, chlorine gas, and hydrogen gas.
$$\text{Brine Solution} = \text{Concentrated Aqueous NaCl } (\approx 28% \text{ by mass})$$
- Overall Balanced Chemical Equation:
$$2\text{NaCl}(aq) + 2\text{H}_2\text{O}(l) \xrightarrow{\text{Electric Current}} 2\text{NaOH}(aq) + \text{Cl}_2(g)\uparrow + \text{H}_2(g)\uparrow$$
Spatial Partitioning of Reaction Products:
- At the Anode (Positive Electrode):
Chloride ions lose electrons to form Chlorine Gas ($\text{Cl}_2$):
$$2\text{Cl}^-(aq) \longrightarrow \text{Cl}_2(g)\uparrow + 2e^- \quad (\text{Oxidation})$$ - At the Cathode (Negative Electrode):
Water molecules gain electrons to form Hydrogen Gas ($\text{H}_2$):
$$2\text{H}_2\text{O}(l) + 2e^- \longrightarrow \text{H}_2(g)\uparrow + 2\text{OH}^-(aq) \quad (\text{Reduction})$$ - In the Electrolyte Near the Cathode:
Accumulating sodium cations ($\text{Na}^+$) combine with hydroxide anions ($\text{OH}^-$) to form a concentrated solution of Sodium Hydroxide ($\text{NaOH}$).
Commercial and Industrial Applications of Products:
- Chlorine Gas ($\text{Cl}_2$): Water treatment (disinfectant), swimming pool purification, manufacturing of PVC, CFCs, pesticides, chloroform, and bleaching powder.
- Hydrogen Gas ($\text{H}_2$): Clean fuel for rocketry, hydrogenation of vegetable oils to margarine, and industrial synthesis of ammonia ($\text{NH}_3$) for fertilizers.
- Sodium Hydroxide ($\text{NaOH}$): De-greasing metals, manufacturing soaps and synthetic detergents, paper pulp bleaching, and refining petroleum and bauxite ore.
- Hydrochloric Acid ($\text{HCl}$, from combining $\text{H}_2 + \text{Cl}_2$): Pickling and cleaning steel surfaces before galvanisation, extraction of glue, and pharmaceutical production.
Bleaching Powder ($\text{CaOCl}_2$)
Bleaching powder, chemically known as calcium oxychloride, is a mixed salt synthesized by the action of dry chlorine gas on dry slaked lime.
Chemical Synthesis:
Produced in industrial towers by passing chlorine gas over dry calcium hydroxide:
$$\text{Ca(OH)}_2(s) [\text{Dry Slaked Lime}] + \text{Cl}_2(g) \longrightarrow \text{CaOCl}_2(s) [\text{Bleaching Powder}] + \text{H}_2\text{O}(l)$$
Chemical Nature and Behavior:
Bleaching powder is a pale yellowish-white powder that smells strongly of chlorine because it reacts with atmospheric carbon dioxide:
$$\text{CaOCl}_2(s) + \text{CO}_2(g) \longrightarrow \text{CaCO}_3(s) + \text{Cl}_2(g)\uparrow$$
Commercial Applications:
- Textile and Paper Bleaching: Bleaches cotton and linen fabrics in textile mills, and bleaches wood pulp in paper manufacturing.
- Chemical Synthesis: Acts as an oxidizing agent in many chemical production processes.
- Water Disinfection: Kills pathogens and sterilizes municipal drinking water supplies.
Baking Soda ($\text{NaHCO}_3$)
Baking soda is the common name for sodium hydrogen carbonate (or sodium bicarbonate), a mild, non-corrosive basic salt synthesized industrially via the Solvay process.
Industrial Preparation (The Solvay Process):
Cold brine is saturated with ammonia and reacted with carbon dioxide under pressure:
$$\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) + \text{NH}_3(g) \longrightarrow \text{NH}_4\text{Cl}(aq) [\text{Ammonium Chloride}] + \text{NaHCO}_3(s)\downarrow [\text{Sodium Bicarbonate}]$$
Sodium bicarbonate is sparingly soluble in cold water, so it precipitates out, is filtered, washed, and dried.
Thermal Decomposition on Cooking:
When sodium hydrogen carbonate is heated during cooking, it decomposes to release carbon dioxide gas:
$$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) [\text{Soda Ash}] + \text{H}_2\text{O}(g) + \text{CO}_2(g)\uparrow$$
Distinction Between Baking Soda and Baking Powder:
- Baking Soda: Pure sodium hydrogen carbonate ($\text{NaHCO}_3$). If used alone in baking, the sodium carbonate ($\text{Na}_2\text{CO}_3$) produced upon heating imparts an unpleasantly bitter taste to cakes and breads.
- Baking Powder: A dry blend of baking soda ($\text{NaHCO}_3$) and a mild edible acid, typically Tartaric acid ($\text{C}_4\text{H}_6\text{O}_6$).
- How it works in baking: When mixed with water into dough and heated, the tartaric acid reacts with the baking soda to liberate carbon dioxide gas. The trapped bubbles of $\text{CO}_2$ expand, causing cakes and breads to rise and become light, soft, and spongy, while the tartaric acid neutralizes the basic sodium carbonate to form harmless, pleasant-tasting sodium tartrate:
$$\text{NaHCO}_3(s) + \text{H}^+(aq) [\text{from edible acid}] \longrightarrow \text{Na}^+(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)\uparrow$$
Additional Applications of Baking Soda:
- Active Ingredient in Antacids: Its alkaline nature neutralizes excess stomach acid without damaging tissue.
- Soda-Acid Fire Extinguishers: Contains sodium bicarbonate solution and concentrated sulphuric acid in separate compartments. When triggered, the two mix to produce large volumes of carbon dioxide gas, which blankets the fire and cuts off the oxygen supply:
$$2\text{NaHCO}_3(aq) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{Na}_2\text{SO}_4(aq) + 2\text{H}_2\text{O}(l) + 2\text{CO}_2(g)\uparrow$$
Washing Soda ($\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$)
Washing soda is hydrated sodium carbonate containing ten fixed molecules of water of crystallisation per formula unit, known chemically as sodium carbonate decahydrate.
Chemical Synthesis:
Obtained in two steps:
- Thermal Calcination of Baking Soda:
$$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) [\text{Anhydrous Soda Ash}] + \text{H}_2\text{O}(g) + \text{CO}_2(g)$$ - Recrystallisation from Water:
Dissolving anhydrous soda ash in water followed by controlled cooling yields washing soda crystals:
$$\text{Na}_2\text{CO}_3(s) + 10\text{H}_2\text{O}(l) \longrightarrow \text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}(s)$$
Efflorescence Phenomenon:
When washing soda decahydrate crystals are left exposed to dry air, they lose nine molecules of water of crystallisation from their surface, forming a dull white powder of sodium carbonate monohydrate:
$$\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}(s) \xrightarrow{\text{Dry Air}} \text{Na}_2\text{CO}_3 \cdot \text{H}_2\text{O}(s) + 9\text{H}_2\text{O}(g)$$
Commercial Applications:
- Water Softening: Removes permanent hardness caused by dissolved calcium and magnesium sulphates or chlorides by precipitating them as insoluble carbonates:
$$\text{CaSO}_4(aq) + \text{Na}_2\text{CO}_3(aq) \longrightarrow \text{CaCO}_3(s)\downarrow + \text{Na}_2\text{SO}_4(aq)$$ - Glass, Soap, and Paper Manufacture: Used as an essential flux in industrial glass melting furnaces.
- Manufacture of Borax: Used as a raw material in synthesizing sodium tetraborate decahydrate (borax, $\text{Na}_2\text{B}_4\text{O}_7 \cdot 10\text{H}_2\text{O}$).
- Domestic Cleanser: Used as a household cleaning agent for laundry.
Plaster of Paris ($\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$) and Gypsum
Plaster of Paris (POP) is calcium sulphate hemihydrate, a white powder prepared by heating gypsum under controlled thermal conditions.
Chemical Preparation:
Synthesized by carefully heating crystalline Gypsum ($\text{CaSO}_4 \cdot 2\text{H}_2\text{O}$) to a precise temperature of $373\text{ K}$ ($100^\circ\text{C}$) in industrial kilns:
$$\text{CaSO}_4 \cdot 2\text{H}_2\text{O}(s) [\text{Gypsum}] \xrightarrow{373\text{ K}} \text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}(s) [\text{Plaster of Paris}] + 1\frac{1}{2}\text{H}_2\text{O}(g)$$
Alternatively written without fractional water molecules:
$$2\text{CaSO}_4 \cdot 2\text{H}_2\text{O} \xrightarrow{373\text{ K}} (2\text{CaSO}_4) \cdot \text{H}_2\text{O}(s) + 3\text{H}_2\text{O}(g)$$
Critical Temperature Caution:
The temperature must not exceed $373\text{ K}$. If heated above $373\text{ K}$ ($> 100^\circ\text{C}$), gypsum loses all of its water of crystallisation, forming anhydrous calcium sulphate ($\text{CaSO}_4$), known as dead burnt plaster. Dead burnt plaster lacks the ability to set into a hard mass upon adding water.
Setting Mechanism of Plaster of Paris:
When mixed with roughly one-third its weight of water, Plaster of Paris rehydrates within $10 – 15\text{ minutes}$, forming an interlocking network of gypsum needle crystals that sets into a hard, rigid mass with a slight expansion in volume:
$$\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}(s) + 1\frac{1}{2}\text{H}_2\text{O}(l) \longrightarrow \text{CaSO}_4 \cdot 2\text{H}_2\text{O}(s) [\text{Hard Set Gypsum}]$$
Storage Precaution:
Plaster of Paris must be stored in airtight, moisture-proof containers. Exposure to humid air causes it to slowly absorb moisture and convert into useless, hard gypsum rock inside the package.
Major Applications:
- Orthopaedic Casts: Immobilizes broken or fractured bones in proper alignment while healing.
- Decorative Construction: Used for decorative moldings, false ceilings, and cornices.
- Art and Dentistry: Used for making molds, statues, ornamental casts, and dental impressions.
- Fireproofing: Used as a thermal barrier material in fireproof doors and walls.
Master Formula and Definition Vault
CHEMICAL IDENTITIES AND COMMON FORMULAS REFERENCE:
1. Quicklime: CaO
2. Slaked Lime: Ca(OH)2
3. Limestone / Marble / Chalk: CaCO3
4. Gypsum: CaSO4 · 2H2O
5. Plaster of Paris (POP): CaSO4 · (1/2)H2O
6. Bleaching Powder: CaOCl2
7. Baking Soda: NaHCO3
8. Washing Soda: Na2CO3 · 10H2O
9. Caustic Soda: NaOH
10. Caustic Potash: KOH
11. Brine: Concentrated NaCl(aq)
12. Blue Vitriol: CuSO4 · 5H2O
13. Green Vitriol: FeSO4 · 7H2O
14. White Vitriol: ZnSO4 · 7H2OQuick Revision Cheat Sheet Table
| Chemical Substance | Formal Chemical Name | Commercial / Popular Name | Chemical Formula | Parent Acid & Base | Key Primary Application |
|---|---|---|---|---|---|
| $\text{NaOH}$ | Sodium Hydroxide | Caustic Soda | $\text{NaOH}$ | $\text{HCl} + \text{NaOH}$ | Soap making, paper pulp, degreasing metals |
| $\text{CaOCl}_2$ | Calcium Oxychloride | Bleaching Powder | $\text{CaOCl}_2$ | $\text{HCl} + \text{Ca(OH)}_2$ | Disinfecting drinking water, textile bleaching |
| $\text{NaHCO}_3$ | Sodium Hydrogen Carbonate | Baking Soda | $\text{NaHCO}_3$ | $\text{H}_2\text{CO}_3 + \text{NaOH}$ | Baking powder synthesis, antacid, fire extinguishers |
| $\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$ | Sodium Carbonate Decahydrate | Washing Soda | $\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$ | $\text{H}_2\text{CO}_3 + \text{NaOH}$ | Removing permanent water hardness, glass manufacturing |
| $\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$ | Calcium Sulphate Hemihydrate | Plaster of Paris (POP) | $\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$ | $\text{H}_2\text{SO}_4 + \text{Ca(OH)}_2$ | Bone fracture casting, decorative ceilings, sculpture molds |
| $\text{CaSO}_4 \cdot 2\text{H}_2\text{O}$ | Calcium Sulphate Dihydrate | Gypsum | $\text{CaSO}_4 \cdot 2\text{H}_2\text{O}$ | $\text{H}_2\text{SO}_4 + \text{Ca(OH)}_2$ | Raw material for POP synthesis, cement retarding agent |
| $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ | Copper(II) Sulphate Pentahydrate | Blue Vitriol | $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ | $\text{H}_2\text{SO}_4 + \text{Cu(OH)}_2$ | Fungicide, agricultural sprays, electroplating |