Chapter 2: Acids, Bases and Salts

2.1 Introduction

• Common Examples:
  • Acids: Lemon (citric acid), grapefruit (citric acid), vinegar (acetic acid), vitamin C (ascorbic acid), battery acid (sulphuric acid).
  • Bases: Ammonia in cleaning solutions, drain cleaner (sodium hydroxide), baking soda (sodium bicarbonate), antacids.
• Taste Recognition:
  • Acids taste sour.
  • Bases taste bitter.
  • Salts (formed from acid-base reactions) taste salty.
• Preservatives: Common salt and sugar are effective preservatives. Bleaching powder and baking soda are common household chemicals.

2.1.1 Acids and Bases

• Historical Context:
  • Ancient knowledge: Vinegar, lemon juice, amla, tamarind taste sour due to acids.
  • Robert Boyle (17th Century): Characterized acids and bases by their observable properties (Table 2.1).
    • Acids: Sour taste, corrosive to metals, turn blue litmus red, become less acidic when mixed with bases.
    • Bases: Bitter taste, slippery/soapy feel, turn red litmus blue, become less basic when mixed with acids.
  • Svante Arrhenius (late 19th Century): Proposed that compounds dissociate into ions in water, and their properties are due to these ions. He defined acids and bases based on the ions they furnish.

2.1.1.1 Acids (Arrhenius Definition)

• Definition: A substance that furnishes hydrogen ions (H⁺) when dissolved in water.
  • Example: HCl (aq) → H⁺ (aq) + Cl⁻ (aq)
• Important Note: Hydrogen ions (H⁺) cannot exist alone in water. They combine with water molecules to form hydronium ions (H₃​O⁺).
  • H⁺ (aq) + H₂​O (l) → H₃​O⁺ (aq)
• Examples of Acids:
  • Hydrochloric acid (HCl) – gastric juice
  • Carbonic acid (H₂​CO₃​) – soft drinks
  • Ascorbic acid (Vitamin C) – lemon, fruits
  • Citric acid – oranges, lemons
  • Acetic acid – vinegar
  • Tannic acid – tea
  • Nitric acid (HNO₃​), Sulphuric acid (H₂​SO₄​) – laboratory use
• Sources of Common Acids (Table 2.2 & Fig 2.1):
  • Acetic acid: Vinegar
  • Citric acid: Lemon, orange
  • Tartaric acid: Tamarind, grapes
  • Ascorbic acid (Vitamin C): All citrus fruits
  • Lactic acid: Milk, yoghurt
  • Malic acid: Apples and pears
  • Formic acid: Ant sting
  • Oxalic acid: Tomato

2.1.1.2 Classification of Acids

1. Based on their Sources:
2. Organic Acids: Present in plants and animals (living beings). Generally weak acids.
3. Examples: Formic acid (HCOOH), Acetic acid (CH₃​COOH)
4. Inorganic Acids (Mineral Acids): From rocks and minerals. Generally strong acids.

1. Examples: HCl, HNO₃​, H₂​SO₄​.
2. Based on their Basicity (Number of replaceable H atoms):
   • Monobasic Acid: Gives one H⁺ ion per molecule.
     • Examples: HCl, HNO₃​, HClO₄​, H₃​PO₂​, H₃​BO₃​
   • Dibasic Acid: Gives two H⁺ ions per molecule.
     • Examples: H_2​SO₄​, H₂​CO₃​, (COOH)₂​, H₃​PO₃​
   • Tribasic Acid: Gives three H⁺ ions per molecule.
     • Example: H₃​PO_4​
   • Note: Acetic acid (CH_3​COOH) has four hydrogen atoms but is monobasic because only one hydrogen is replaceable.
3. Based on Ionisation (Strength):
   • Strong Acids: Ionise completely in water.
     • Examples: HCl, HNO₃​, H₂​SO₄​
   • Weak Acids: Ionise partially in water.
     • Examples: CH₃​COOH, H₂​CO_3​
4. Based on Concentration:
   • Concentrated Acid: High percentage of acid in aqueous solution.
   • Dilute Acid: Low percentage of acid in aqueous solution.
   • Safety Note: Always add concentrated acid slowly to water with constant stirring. Adding water to concentrated acid can generate large amounts of heat, causing burns and splashing.

2.1.1.3 Bases (Arrhenius Definition)

• Definition: A substance that furnishes hydroxide ions (OH⁻) when dissolved in water.
  • Example: NaOH (aq)→Na⁺(aq) + OH⁻(aq)
• Alkali: Water-soluble bases are called alkalis.
• Bases in Everyday Life (Table 2.3):
  • Baking powder: Sodium bicarbonate
  • Ash: Potassium carbonate
  • Glass cleaners: Ammonia
• Examples of Bases:
  • Sodium hydroxide (NaOH) / Caustic soda – washing soaps
  • Potassium hydroxide (KOH) / Potash – bathing soaps
  • Calcium hydroxide (Ca(OH)₂​) / Lime water – white wash
  • Magnesium hydroxide (Mg(OH)₂​) / Milk of magnesia – antacid (controls acidity)
  • Ammonium hydroxide (NH₄​OH) – hair dyes

2.1.1.4 Indicators

• Definition: Substances that show different colors in acidic and basic mediums.
• Types of Indicators:

1. Natural Indicators: Obtained from natural sources.
   • Litmus: Obtained from lichens. Solution is purple.
     • Acid turns blue litmus paper red.
     • Base turns red litmus paper blue.
   • Turmeric: Yellow in color.
     • Turns reddish-brown with base.
     • No color change with acid.
   • Red Cabbage Juice: Originally purple.
     • Turns reddish with acid.
     • Turns greenish with base.
2. Olfactory Indicators: Substances that change their smell when mixed with acid or base.
   • Onion: Loses smell with base; no change with acid.
   • Vanilla: Smell vanishes with base; no change with acid.
   • Clove: Similar to onion and vanilla.
3. Synthetic Indicators: Synthesized in laboratories.
   • Phenolphthalein: Colorless liquid.
     • Remains colorless with acid.
     • Turns pink with base.
   • Methyl Orange: Originally orange.
     • Turns red with acid.
     • Turns yellow with base.

• Summary of Indicator Colors (Table 2.4):
  • Red Litmus: Red (Acid), Turns blue (Base)
  • Blue Litmus: Turns red (Acid), No change (Base)
  • Turmeric: No change (Acid), Turns reddish brown (Base)
  • Red Cabbage Juice: Reddish (Acid), Greenish-yellow (Base)
  • Phenolphthalein: Colorless (Acid), Turns pink (Base)
  • Methyl Orange: Turns red (Acid), Turns yellow (Base)
  • Onion/Vanilla: No change (Acid), Smell vanishes (Base)

4. Universal Indicator:
   • Purpose: Determines both the acidic/basic nature and the strength (pH) of a solution.
   • Composition: Combination of many indicators (e.g., water, propanol, phenolphthalein, methyl red, bromothymol blue, thymol blue).
   • Function: Shows different colors over a pH range of 1 to 14. A color matching chart is provided.
   • pH Scale (Fig 2.3 & Table 2.5):
     • Strong Acid (pH 0-2): Dark Red to Red
     • Weak Acid (pH 3-6): Orange-Red to Greenish Yellow
     • Neutral (pH 7): Green
     • Weak Alkali (pH 8-10): Greenish Blue to Navy Blue
     • Strong Alkali (pH 11-14): Purple to Dark Purple/Violet

2.2 Understanding the Properties of Acids and Bases

2.2.1 Physical and Chemical Properties of Acids

1. Taste: Sour (e.g., lemon, vinegar, sour milk).
   • Acids present in common substances (Table 2.6):
     • Lemon juice: Citric acid and ascorbic acid
     • Vinegar: Ethanoic acid (acetic acid)
     • Tamarind: Tartaric acid
     • Sour milk: Lactic acid
2. Action on Indicators (Table 2.7):
   • Litmus: Red
   • Phenolphthalein: Colorless
   • Methyl orange: Red
3. Reaction of Acids with Metals:
   • General Reaction: Acid + Metal → Salt + Hydrogen gas
   • Example: Zn (s) + 2HCl (aq)→ZnCl₂ ​(aq) + H₂ ​(g)↑
   • Observation: Hydrogen gas burns with a ‘pop’ sound.
   • Note: Some metals like Ag, Cu do not liberate hydrogen gas with acids.
4. Reaction of Acids with Metal Carbonates and Metal Bicarbonates (Hydrogen Carbonates):
   • General Reaction:
     • Metal Carbonate + Acid → Salt + Water + Carbon dioxide
     • Metal Hydrogen Carbonate + Acid → Salt + Water + Carbon dioxide
   • Examples:
     • Na₂​CO₃​ (s) + 2HCl (aq) → 2NaCl (aq) + H₂​O (l) + CO₂ ​(g)↑
     • NaHCO₃ ​(s) + HCl (aq) → NaCl (aq) + H₂​O (l) + CO₂ ​(g)↑
   • Test for CO₂​: When CO₂​ is passed through lime water (Ca(OH)₂​), it turns milky due to the formation of insoluble calcium carbonate (CaCO₃​).
     • Ca(OH)₂​(aq)+CO₂​(g)→CaCO₃​(s)↓+H₂​O(l)
   • Excess CO₂​: If excess CO₂​ is passed, the milky white precipitate disappears due to the formation of soluble calcium hydrogen carbonate (Ca(HCO₃​)₂​).
     • CaCO₃​(s)+H₂​O(l)+CO₂​(g)→Ca(HCO₃​)₂​(aq)
   • Reason: Metal carbonates and bicarbonates are basic; hence they react with acids.
5. Reaction of Acids with Metallic Oxides:
   • General Reaction: Metallic oxide + Acid → Salt + Water
   • Example: CuO(s)+2HCl(dil.)→CuCl₂​(aq)+H₂​O(l)
     • Copper oxide (black) reacts to form copper (II) chloride (bluish green).
   • Reason: Metallic oxides are basic in nature.
6. Reaction of Acids with Water:
   • Acids produce hydrogen ions (H⁺) in water. These H⁺ ions combine with water to form hydronium ions (H₃​O⁺).
   • HCl+H₂​O→H₃​O⁺+Cl⁻
   • Hydrogen ions cannot exist alone; they are always in the form of H₃​O⁺.
7. Basicity of Acids: The number of ionisable hydrogen (H⁺) ions present in one molecule of an acid.
   • HCl: Basicity = 1 (monobasic)
   • H₂​SO₄​: Basicity = 2 (dibasic)
   • H₃​PO₄​: Basicity = 3 (tribasic)
8. Strength of Acids: Depends on the number of H⁺ ions produced in water.
   • Strong Acids: Provide more H⁺ ions (ionise completely).
     • Examples: Hydrochloric acid (HCl), Nitric acid (HNO₃​), Sulphuric acid (H₂​SO₄​)
   • Weak Acids: Provide fewer H⁺ ions (ionise partially).
     • Examples: Acetic acid (CH₃​COOH), Carbonic acid (H₂​CO₃​)
9. Corrosive Nature: Acids attack substances like metals, metal oxides, and hydroxides. Concentrated acids can cause severe burns.

2.2.2 Uses of Acids

1. Sulphuric acid (King of chemicals) – used in car batteries, preparation of many other compounds.
2. Nitric acid – production of ammonium nitrate (fertilizer).
3. Hydrochloric acid – cleaning agent in toilet.
4. Tartaric acid – constituent of baking powder.
5. Salt of benzoic acid (sodium benzoate) – food preservation.
6. Carbonic acid – aerated drinks.

2.2.3 Classification of Bases

1. Based on Ionisation:
   • Strong Bases: Ionise completely in aqueous solution.
     • Examples: NaOH, KOH, Ca(OH)₂​ (partially soluble, but strong when dissolved)
   • Weak Bases: Ionise partially in aqueous solution.
     • Examples: NH₄​OH, Mg(OH)_2​
2. Based on their Acidity (Number of replaceable OH⁻ ions):
   • Monoacidic Base: Gives one hydroxide ion per molecule.
     • Examples: NaOH, KOH
   • Diacidic Base: Gives two hydroxide ions per molecule.
     • Examples: Ca(OH)₂​, Mg(OH)_2​
   • Triacidic Base: Gives three hydroxide ions per molecule.
     • Example: Al(OH)₃​, Fe(OH)₃​
3. Based on Concentration:
   • Concentrated Alkali: High percentage of base in aqueous solution.
   • Dilute Alkali: Low percentage of base in aqueous solution.

2.2.4 Physical and Chemical Properties of Bases

1. Taste and Touch: Bitter taste, soapy/slippery to touch.
2. Action on Indicators (Table 2.9):
   • Litmus: Blue
   • Phenolphthalein: Pink
   • Methyl orange: Yellow
3. Reaction of Bases with Metals: Some metals (e.g., Zinc) react with strong bases to liberate hydrogen gas.
   • Example: Zn(s)+2NaOH(aq)→Na₂​ZnO₂​(aq)+H₂​(g)↑
   • Note: Metals like Ag, Cu, Al do not react with sodium hydroxide.
4. Reaction of Non-Metallic Oxides with Bases:
   • General Reaction: Non-metallic oxide + Base → Salt + Water
   • Example: 2NaOH(aq)+CO₂​(g)→Na₂​CO₃​(aq)+H₂​O(l)
   • Reason: Non-metallic oxides are acidic in nature.
5. Reaction of Bases with Acids (Neutralisation Reaction):
   • Definition: Acids and bases cancel out each other’s effects, forming salt and water.
   • General Reaction: Acid + Base → Salt + Water
   • Examples:
     • HCl(aq)+NaOH(aq)→NaCl(aq)+H₂​O(l)
     • H₂​SO₄​(aq)+2KOH(aq)→K₂​SO₄​(aq)+2H₂​O(l)
   • Uses of Neutralisation:
     • Lime (Ca(OH)₂​) to neutralize acidic soil.
     • Antacids (Mg(OH)₂​, NaHCO₃​) to control acid indigestion.
     • Ant sting (formic acid) neutralized by baking soda (sodium bicarbonate).

2.2.5 Strength of Bases

• Depends on the number of hydroxide (OH−) ions produced in water.
  • Strong Bases: Furnish more OH− ions (ionise completely).
    • Examples: NaOH, KOH, Ca(OH)₂​ (Table 2.10)
  • Weak Bases: Furnish fewer OH− ions (ionise partially).
    • Examples: NH₄​OH, Mg(OH)₂​ (Table 2.10)

2.2.6 Acidity of Bases

• The number of replaceable hydroxide ions in a base.
  • NaOH, KOH: Monoacidic (Acidity = 1)
  • Ca(OH)₂​, Ba(OH)₂​: Diacidic (Acidity = 2)
  • Al(OH)₃​: Triacidic (Acidity = 3)

2.2.7 Uses of Bases

1. Sodium hydroxide – manufacture of soap.
2. Calcium hydroxide – white washing buildings.
3. Magnesium hydroxide – antacid.
4. Ammonium hydroxide – remove grease stains from clothes.

2.3 What do all Acids have in Common?

• All acids produce hydrogen ions (H⁺) or hydronium ions (H3​O⁺) in water.
• The presence of these ions is responsible for the acidic properties and their ability to conduct electricity in aqueous solutions.
• Compounds like glucose (C₆​H_12​O₆​) and ethyl alcohol (C₂​H₅​OH) contain hydrogen but do not produce H⁺ ions in water, so they are not acidic and do not conduct electricity.

2.3.1 What Happens to an Acid or a Base in a Water Solution?

• Acids in Water: Acids dissociate to produce H+ ions, which then combine with water to form H3​O+ ions.
  • HCl(g)+H₂​O(l)→H₃​O+(aq)+Cl−(aq)
  • Dry HCl gas does not change the color of dry litmus paper because it does not produce H+ ions without water.
• Bases in Water: Bases dissociate to produce OH− ions.
  • NaOH(s) + H₂​O​ → Na⁺(aq) + OH⁻(aq)
  • Mg(OH)₂​(s)  +  H₂​O​ → Mg²⁺(aq) + 2OH⁻ (aq)
• Alkalis: Water-soluble bases are called alkalis. All alkalis are bases, but not all bases are alkalis.

2.3.2 Effect of Dilution of Acid and Base with Water

• Dilution: Adding water to an acid or base decreases the concentration of H₃​O⁺ or OH⁻ ions per unit volume. This is a highly exothermic process.
• Safety: Always add acid/base slowly to water with constant stirring, not the other way around.

2.3.3 Self Dissociation of Water

• Water undergoes self-dissociation to a small extent, forming H⁺ and OH⁻ ions.
  • H₂​O(l) ⇌ H⁺(aq) + OH⁻(aq)
  • At 25^∘C, [H⁺] = [OH⁻]= 1×10⁻⁷M.

2.3.4 Neutral, Acidic and Basic Solutions

• Neutral Solutions: [H⁺]=[OH⁻] (e.g., pure water).
• Acidic Solutions: [H⁺]>[OH⁻] (or [H⁺]>1×10⁻⁷M). Increase in H⁺ concentration.
• Basic Solutions: [OH⁻]>[H⁺] (or [OH⁻]>1×10⁻⁷M). Increase in OH⁻ concentration.

2.4 Arrhenius Theory of Acids and Bases

• Acids: Give H⁺ ions in aqueous solution. Strong acids give more H⁺ ions.
• Bases: Give OH⁻ ions in aqueous solution. Strong bases give more OH⁻ ions.

2.4.1 pH and pH Scale

• pH Scale: A measure of the relative acidity (H⁺) or alkalinity (OH⁻) of a solution.
• Definition: pH is the negative logarithm of the hydrogen ion concentration.
  • pH=−log[H⁺]
• Interpretation:
  • pH < 7: Acidic solution (higher H⁺ concentration). Lower pH means stronger acid.
  • pH = 7: Neutral solution ([H⁺]=10⁻⁷M).
  • pH > 7: Basic (alkaline) solution (lower H⁺ concentration, higher OH⁻ concentration). Higher pH means stronger base.
• pOH Scale: Measures hydroxide ion concentration.
  • pOH = −log[OH⁻]
  • Relationship: pH + pOH=14 (at 25^∘C)
• Common Acids and Bases pH values (Table 2.12):
  • HCl (4%): 0 (Strong acid)
  • Stomach acid: 1
  • Lemon juice: 2
  • Vinegar: 3
  • Oranges: 3.5
  • Soda, grill: 4
  • Sour milk: 6
  • Fresh milk: 6.5
  • Human saliva: 6.5-7.5
  • Pure water: 7 (Neutral)
  • Blood plasma: 7.4
  • Baking soda: 8.5
  • Sea water: 8
  • Ammonia water: 10
  • Lime water: 11
  • Drain cleaner: 13
  • Caustic soda (4% NaOH): 14 (Strong base)

2.4.2 pH Paper

• A common method to measure pH in a laboratory. It is a paper containing universal indicator, which gives different colors across the entire pH range.

2.4.3 Importance of pH in Everyday Life

1. pH of Human Body:
   • Our body works within a narrow pH range of 7.35 to 7.45.
   • pH changes can affect health (e.g., lower pH can lead to infections, cough, cold).
   • Normal human blood pH is 7.36.
   • Human blood pH range is 7.35-7.45.
2. pH in Digestion System:
   • Stomach produces hydrochloric acid (pH 1.2-1.5) for digestion.
   • Excess acid causes “acid indigestion,” treated with antacids (bases like milk of magnesia).
3. pH of Rainwater:
   • Normal rainwater pH is approximately 7 (neutral).
   • Acid Rain: When pollutants (e.g., SO₂​, NO₂​) dissolve in rainwater, pH drops below 5.6, leading to acid rain.
4. Self Defence of Animals and Plants:
   • Ant Sting: Contains formic acid, causing pain. Neutralized by baking soda.
   • Nettle Plant: Contains methanoic acid, causing burning pain. Relieved by rubbing with dock plant leaves (basic).
5. Plant Growth:
   • Plants require a specific pH range for healthy growth.
   • Acidic soil can be neutralized by adding lime (Ca(OH)₂​) or chalk (CaCO_3​).
   • Basic soil can be neutralized by adding organic matter.
6. Industry: Acids and bases are used in various industries (e.g., paper, paints, drugs).

2.5 More About Salts

• Definition: Ionic compounds formed from the reaction of acids and bases. They contain positive ions (cations) from a base and negative ions (anions) from an acid.
• Basic Radical: The positively charged ion (cation) from the base.
• Acid Radical: The negatively charged ion (anion) from the acid.
• Example: In NaCl, Na⁺ is the basic radical (from NaOH), and Cl⁻ is the acid radical (from HCl).

2.5.1 Family of Salts

Salts are classified based on the acid or base they are derived from, or the common ion they contain.

1. Classification based on Chemical Formulas:
   • Normal or Neutral Salt: Does not contain replaceable hydrogen or hydroxide ions. Formed from strong acid + strong base.
     • Examples: NaCl, KNO₃​, Na₂​SO_4​
   • Bisalte or Acid Salts: Contain replaceable hydrogen ions. Formed from strong acid + weak base, or when the acid has more than one replaceable hydrogen.
     • Example: NaHCO₃​ (sodium hydrogen carbonate)
   • Basic Salts: Contain replaceable hydroxide ions. Formed from weak acid + strong base.
     • Example: Cu(OH)Cl (basic copper chloride)
   • Double Salts: Formed by the simple association of two simple salts.
     • Example: Potash alum (K₂​SO₄​⋅Al₂​(SO₄​)₃​⋅24H_2​O)
   • Mixed Salts: Contain more than one acid or basic radical.
2. Classification based on Family of Salt (common anion/cation):
   • Sulphate Family: Contains SO₄^2−​ ion.
     • Examples: Potassium sulphate (K₂​SO_4​), Sodium sulphate (Na₂​SO_4​), Magnesium sulphate (MgSO₄​), Calcium sulphate (CaSO₄​), Copper sulphate (CuSO₄​)
   • Chloride Family: Contains Cl⁻ ion.
     • Examples: Sodium chloride (NaCl), Potassium chloride (KCl), Barium chloride (BaCl₂​), Magnesium chloride (MgCl_2​), Aluminium chloride (AlCl₃​), Ammonium chloride (NH₄​Cl)
   • Nitrate Family: Contains NO₃⁻​ ion.
     • Examples: Sodium nitrate (NaNO₃​), Potassium nitrate (KNO₃​), Calcium nitrate (Ca(NO₃​)₂​), Copper nitrate (Cu(NO₃​)₂​), Zinc nitrate (Zn(NO₃​)₂​), Cadmium nitrate (Cd(NO₃​)₂​), Strontium nitrate (Sr(NO₃​)₂​)
   • Carbonate Family: Contains CO₃^2−​ ion.
     • Examples: Sodium carbonate (Na_2​CO₃​), Potassium carbonate (K₂​CO₃​), Calcium carbonate (CaCO₃​), Magnesium carbonate (MgCO₃​)
3. Classification based on pH Values:
   • Neutral Salts (pH = 7): Formed from strong acid + strong base (e.g., NaCl).
   • Acidic Salts (pH < 7): Formed from strong acid + weak base (e.g., NH₄​Cl).
   • Basic Salts (pH > 7): Formed from weak acid + strong base (e.g., Na₂​CO₃​).

2.5.2 Formation of Salts

Salts are formed by various reactions:

1. Neutralisation Reaction: Acid + Base → Salt + Water
2. Action of Acids on Metals: Metal + Acid → Salt + Hydrogen
3. Action of Acids on Metal Carbonates and Hydrogen Carbonates:
   • Metal Carbonate + Acid → Salt + Water + Carbon dioxide
   • Metal Hydrogen Carbonate + Acid → Salt + Water + Carbon dioxide

2.5.3 Chemicals from Common Salt

Common salt (Sodium chloride, NaCl) is a raw material for many important chemicals.

1. Sodium Chloride (NaCl):
   • Occurrence: Seawater (2.7-2.9%), salt lakes, rock salt.
   • Extraction: Evaporation of seawater. Crude salt contains impurities like MgSO₄​, MgCl₂​, CaSO₄​, CaCl₂​. Pure salt is obtained by dissolving crude salt in minimum water and filtering.
   • Essential for life: Important for body functions, nerve impulses, muscle contractions. High intake can lead to high blood pressure.
   • Uses:
     • Essential food constituent.
     • Manufacture of Na₂​CO₃​, NaOH, Cl₂​.
     • Salting out soap and organic dyes.
     • Tanning and textile industries.
     • Preservative for fish, meat, butter.
     • Raw material for sodium hydroxide, sodium carbonate (washing soda), sodium hydrogen carbonate (baking soda), bleaching powder.
2. Sodium Hydroxide (NaOH) – Caustic Soda:
   • Preparation: Electrolysis of an aqueous solution of sodium chloride (brine). This is called the Chlor-alkali process.
     • 2NaCl (aq)+2H₂​O (l) electricity​ 2NaOH(aq)+Cl₂​(g)↑+H₂​(g)↑
     • Products:
       • Chlorine (Cl₂​): Anode. Used as disinfectant, PVC, CFCs, bleaching powder, water treatment.
       • Hydrogen (H₂​): Cathode. Used as fuel, ammonia for fertilizers, margarine.
       • Sodium Hydroxide (NaOH): Near cathode. Used in soaps, detergents, paper, artificial silk, bauxite purification, degreasing metals.
3. Bleaching Powder (CaOCl₂​):
   • Preparation: By passing chlorine gas over dry slaked lime (Ca(OH)₂​).
     • Ca(OH)₂​(s)+Cl₂​(g)→CaOCl₂​(s)+H₂​O(l)
   • Uses:
     • Bleaching cotton and linen in textile industry.
     • Bleaching wood pulp in paper factories.
     • Bleaching washed clothes in laundry.
     • Disinfectant for drinking water.
     • Oxidizing agent in chemical industries.
     • Preparation of chloroform.
4. Baking Soda (Sodium Hydrogen Carbonate, NaHCO₃​):
   • Preparation: Solvay’s process (Ammonia-soda process).
     • NaCl + H₂​O + CO₂​ + NH₃ ​→ NH₄​Cl + NaHCO₃​
   • Properties: Mild non-corrosive basic salt.
   • Uses:
     • Baking Powder: Mixture of baking soda and a mild edible acid (e.g., tartaric acid). When heated or mixed with water, CO2​ is produced, causing bread/cake to rise.
       • NaHCO₃​ + H⁺(from acid) → CO₂ ​+ H₂​O + Sodium salt of acid
     • Antacid (neutralizes excess acid in stomach).
     • Fire extinguishers (produces CO₂​).
     • Reagent in laboratory.
5. Washing Soda (Sodium Carbonate Decahydrate, Na₂​CO₃​⋅10H₂​O):
   • Preparation: From baking soda by heating.
     • 2NaHCO₃​(s) + Heat​ → Na₂​CO₃​(s) + H₂​O(l) + CO₂​(g)
     • Recrystallization of anhydrous sodium carbonate with water gives washing soda.
     • Na₂​CO₃​(s)+10H₂​O(l)→Na_2​CO₃​⋅10H₂​O(s)
   • Uses:
     • Manufacture of glass, soap, paper.
     • Manufacture of sodium compounds like borax.
     • Cleaning agent for domestic purposes.
     • Removing permanent hardness of water.
     • In textile and petroleum industries.

2.6 Are the Crystals of Salts Really Dry?

• Water of Crystallisation: The fixed number of water molecules present in one formula unit of a salt.
  • Example: Copper sulphate crystals (CuSO₄​⋅5H₂​O) are blue. When heated, they lose water and turn white (CuSO₄​). Adding water back restores the blue color.
• Hydrated Salts: Salts containing water of crystallisation.
• Anhydrous Salts: Salts that have lost their water of crystallisation.

2.6.1 Plaster of Paris (CaSO₄​⋅1/2​H₂​O)

• Chemical Name: Calcium sulphate hemihydrate.
• Preparation: By heating gypsum (CaSO₄​⋅2H₂​O) to 373K.
  • CaSO₄​⋅2H₂​O + 373K →  ​CaSO₄​⋅1/2 ​H₂​O + 3/2 ​H₂​O
• Properties: White powder. When mixed with water, it sets into a hard porous mass (gypsum) in 5-15 minutes.
• Uses:
  • Medical field: Setting fractured bones.
  • Making casts for statues, dentistry, surgical instruments, toys.
  • Making black board chalks and statues.
  • Construction industry.
• Storage: Should be stored in a moisture-proof container.

Summary

• Acids produce H⁺ in water; bases produce OH⁻.
• Indicators show different colors in acidic/basic mediums.
• pH scale measures H⁺ concentration (0-14). pH < 7 (acidic), pH = 7 (neutral), pH > 7 (basic).
• Acids react with metals to give salt and hydrogen.
• Acids react with metal carbonates/bicarbonates to give salt, water, and carbon dioxide.
• Acids react with metallic oxides to give salt and water.
• Bases react with non-metallic oxides to give salt and water.
• Neutralization: Acid + Base → Salt + Water.
• pH is important in human body, digestion, rainwater, plant growth, and self-defence of animals/plants.
• Salts are ionic compounds formed from acid-base reactions.
• Common salt (NaCl) is a raw material for NaOH, bleaching powder, baking soda, washing soda.
• Water of crystallization is the fixed number of water molecules in a salt crystal.
• Plaster of Paris is formed from gypsum and used for casts and construction.

Practice Questions: Acids, Bases and Salts
This set of questions covers various topics from the “Acids, Bases and Salts” chapter, designed to help you assess your knowledge from basic concepts to more complex applications.
Section 1: Easy Questions (Multiple Choice, Fill in the Blanks, Very Short Answer)
Multiple Choice Questions:
Which of the following is a natural indicator? a) Phenolphthalein b) Methyl orange c) Turmeric d) Universal indicator
Acids turn blue litmus paper: a) Yellow b) Red c) Green d) Blue
A solution turns red litmus blue. Its pH is likely to be: a) 2 b) 5 c) 7 d) 10
The chemical formula of washing soda is: a) Na₂​CO₃​ b) NaHCO₃​ c) Na₂​CO₃​⋅10H₂​O d) NaOH
Which gas is produced when an acid reacts with a metal? a) Oxygen b) Carbon dioxide c) Hydrogen d) Nitrogen
Which of the following is a strong base? a) NH₄​OH b) Mg(OH)₂​ c) NaOH d) Al(OH)₃​
The acid present in lemon juice is: a) Acetic acid b) Lactic acid c) Citric acid d) Tartaric acid
Fill in the Blanks:
The term acid comes from the Latin word ‘accre’ which means __________.
Bases generally feel __________ or __________.
A substance that furnishes hydrogen ions (H⁺) when dissolved in water is called an __________.
Water-soluble bases are known as __________.
The pH of a neutral solution at 25^∘C is __________.
The gas evolved when an acid reacts with a metal carbonate turns lime water __________.
Plaster of Paris is chemically known as calcium sulphate __________.
Very Short Answer Questions (1-2 sentences):
Give two examples of strong acids.
What is the basicity of H₃​PO₄​?
Name one common antacid.
What is the colour of phenolphthalein in a basic medium?
Why is it recommended to add acid slowly to water, and not water to acid?
What is the common name of sodium hydroxide?
Which type of oxide reacts with bases to form salt and water?
Section 2: Medium Questions (Short Answer, Application-based)
Short Answer Questions (3-4 sentences):
Define Arrhenius acids and bases with one example each.
Explain the difference between organic and inorganic acids.
Describe the observation when carbon dioxide gas is passed through lime water. What happens if excess carbon dioxide is passed?
What is water of crystallisation? Give an example of a salt with water of crystallisation.
Explain why dry HCl gas does not change the colour of dry litmus paper.
What are olfactory indicators? Give an example.
Why are all alkalis bases, but not all bases alkalis?
Application-based Questions:
A student adds a few drops of an unknown liquid to a solution of baking soda. They observe brisk effervescence. What can be inferred about the unknown liquid? Write the chemical equation for the reaction.
You have two solutions, A and B. The pH of solution A is 6 and pH of solution B is 8. Which solution has more hydrogen ion concentration? Which is more acidic?
How does the pH in our digestive system help in food digestion? What happens during “acid indigestion”?
A gardener finds that the soil in their garden is too acidic for plant growth. What common household substance can they add to neutralize the soil? Explain the chemical reaction involved.
Identify the type of salt (neutral, acidic, or basic) formed from the reaction of: a) Strong acid and strong base b) Strong acid and weak base c) Weak acid and strong base
Copper oxide is a black powder. When dilute hydrochloric acid is added to it, the solution turns bluish-green. Explain the observation and write the balanced chemical equation.
Why do compounds like glucose and alcohol not show acidic properties even though they contain hydrogen?
Section 3: Difficult Questions (Long Answer, Reasoning, Problem Solving)
Long Answer Questions (5+ sentences / detailed explanation):
Describe the Chlor-alkali process in detail, including the balanced chemical equation and the uses of each product formed.
Explain the importance of pH in everyday life, covering at least three different aspects (e.g., human body, acid rain, plant growth, self-defence).
Differentiate between strong and weak acids, and strong and weak bases. Provide examples and explain the basis of their strength.
Describe the preparation of Plaster of Paris from gypsum. Explain its setting property and give two important uses. Why should it be stored in a moisture-proof container?
Explain the process of dilution of an acid or a base. Why is it important to add acid/base to water and not the other way around? What is the effect on the concentration of ions?
Discuss the different types of indicators used in acid-base titrations. How does a universal indicator differ from other indicators like litmus or phenolphthalein?
Reasoning and Problem Solving:
A student spills a concentrated acid on their hand in the laboratory. What immediate action should they take? Explain the scientific reason behind the severity of concentrated acids.
You are given three unknown solutions: one is acidic, one is basic, and one is neutral. Describe a series of tests you would perform using a universal indicator to identify each solution and determine their approximate pH values.
A compound ‘X’ is used in the manufacture of glass and is also used for removing permanent hardness of water. When ‘X’ is heated, it gives a compound ‘Y’ and carbon dioxide gas. Identify ‘X’ and ‘Y’. Write the chemical equations for the reactions involved.
A solution contains 10⁻⁵M H⁺ ions. a) Is the solution acidic, basic, or neutral? b) Calculate its pH. c) If this solution is diluted 100 times, what will be the new pH? (Assume temperature is 25^∘C)
A student wants to prepare hydrogen gas in the lab. They have dilute sulphuric acid, zinc granules, copper turnings, and a test tube setup. Which metal should they use with the acid to produce hydrogen gas? Explain why the other metal would not work. Write the chemical equation.
Baking soda and washing soda are both sodium salts. Differentiate between them based on their chemical formula, preparation, and one major use.
Why do crystals of copper sulphate appear dry but turn blue when water is added to them? Explain the phenomenon.
Answers
Section 1: Easy Questions
Multiple Choice Questions:
c) Turmeric
b) Red
d) 10
c) Na_2​CO₃​⋅10H₂​O
c) Hydrogen
c) NaOH
c) Citric acid
Fill in the Blanks:
sour
slippery, soapy
acid
alkalis
7
milky
hemihydrate
Very Short Answer Questions:
HCl (Hydrochloric acid), H₂​SO₄​ (Sulphuric acid).
The basicity of H₃​PO₄​ is 3 (tribasic).
Milk of magnesia (Mg(OH)₂​).
Pink.
Adding water to concentrated acid generates a large amount of heat, which can cause burns and splashing. Adding acid slowly to water ensures the heat is dissipated safely.
Caustic soda.
Acidic oxides (non-metallic oxides).
Section 2: Medium Questions
Short Answer Questions:
Arrhenius Acid: A substance that furnishes hydrogen ions (H⁺) when dissolved in water. Example: HCl(aq)→ H⁺(aq) + Cl⁻(aq). Arrhenius Base: A substance that furnishes hydroxide ions (OH⁻) when dissolved in water. Example: NaOH(aq) → Na⁺(aq) + OH⁻(aq).
Organic Acids: Acids present in plants and animals (living beings), generally weak acids. Example: Acetic acid (CH_3​COOH). Inorganic Acids (Mineral Acids): Acids derived from rocks and minerals, generally strong acids. Example: Hydrochloric acid (HCl).
When carbon dioxide gas is passed through lime water (Ca(OH)₂​), it turns milky due to the formation of insoluble calcium carbonate (CaCO₃​). If excess carbon dioxide is passed, the milky white precipitate disappears due to the formation of soluble calcium hydrogen carbonate (Ca(HCO₃​)₂​).
Water of crystallisation is the fixed number of water molecules chemically combined with one formula unit of a salt in its crystalline state. Example: Copper sulphate pentahydrate (CuSO₄​⋅5H₂​O).
Dry HCl gas does not change the colour of dry litmus paper because it does not dissociate to produce hydrogen ions (H⁺) in the absence of water. The acidic properties are only shown when H⁺ ions are formed in aqueous solution.
Olfactory indicators are substances whose smell changes when mixed with acidic or basic solutions. An example is onion, whose smell vanishes when a base is added to it.
All alkalis are bases because alkalis are water-soluble bases and thus produce hydroxide ions (OH⁻) in water. However, not all bases are alkalis because some bases are insoluble in water (e.g., Mg(OH)₂​, Al(OH)₃​) and therefore do not fall under the definition of alkalis.
Application-based Questions:
The unknown liquid is an acid. The brisk effervescence indicates the production of carbon dioxide gas, which is characteristic of the reaction between an acid and a metal bicarbonate (baking soda). Chemical equation: NaHCO₃​(s)+HA(aq)→NaA(aq)+H₂​O(l)+CO₂​(g)↑ (where HA is a generic acid) For example, with HCl: NaHCO₃​(s) + HCl(aq)→NaCl(aq) + H₂​O(l) + CO₂​(g)↑
Solution A (pH 6) has more hydrogen ion concentration than solution B (pH 8). Solution A is more acidic. (Lower pH means higher [H⁺]).
The pH in our digestive system (stomach) is highly acidic (around 1.2-1.5) due to hydrochloric acid, which helps in the digestion of food. During “acid indigestion,” excess acid is produced in the stomach, causing pain and irritation. Antacids (bases) are consumed to neutralize this excess acid.
The gardener can add slaked lime (Ca(OH)₂​) or chalk (CaCO₃​) to neutralize the acidic soil. These are basic substances that will react with the excess acid in the soil to form salt and water, thus increasing the pH. Chemical reaction (with slaked lime): Acid+ Ca(OH)₂​ → Salt + H₂​O
a) Strong acid and strong base: Neutral salt (e.g., NaCl) b) Strong acid and weak base: Acidic salt (e.g., NH₄​Cl) c) Weak acid and strong base: Basic salt (e.g., Na₂​CO₃​)
When dilute hydrochloric acid is added to black copper (II) oxide, the black powder slowly dissolves, and the solution turns bluish-green. This is because copper (II) oxide is a metallic oxide (basic) and reacts with the acid to form copper (II) chloride, which is bluish-green in colour. Balanced chemical equation: CuO(s) + 2HCl(dil.) → CuCl₂​(aq) + H₂​O(l)
Compounds like glucose (C₆​H₁₂​O₆​) and alcohol (C₂​H₅​OH) contain hydrogen but do not show acidic properties because they do not ionize to produce hydrogen ions (H⁺) when dissolved in water. The hydrogen atoms in these compounds are covalently bonded and are not released as ions. Therefore, their solutions do not conduct electricity.
Section 3: Difficult Questions
Long Answer Questions:
Chlor-alkali Process: This is the process of electrolysis of an aqueous solution of sodium chloride (brine).
Reaction:2NaCl(aq)+2H₂​O(l)+electricity​→2NaOH(aq)+Cl₂​(g)↑+H₂​(g)↑
Process: When electricity is passed through the brine solution, sodium chloride decomposes. Chlorine gas (Cl₂​) is produced at the anode, hydrogen gas (H₂​) is produced at the cathode, and sodium hydroxide (NaOH) solution is formed near the cathode.
Uses of Products:
Chlorine (Cl₂​): Used in water treatment (disinfectant), PVC manufacturing, CFCs, pesticides, bleaching powder.
Hydrogen (H₂​): Used as a fuel, in the production of ammonia for fertilizers, and in the hydrogenation of oils to make margarine.
Sodium Hydroxide (NaOH): Used in the manufacture of soaps and detergents, paper making, artificial silk, bauxite purification, and degreasing metals.
Importance of pH in Everyday Life:
Human Body: Our body functions within a narrow pH range of 7.35 to 7.45. Any significant deviation can lead to health problems. For example, stomach acid (pH 1.2-1.5) is crucial for digestion, but excess acid causes indigestion, which is relieved by antacids.
Acid Rain: The pH of normal rainwater is around 7. When pollutants like sulfur dioxide and nitrogen oxides are released into the atmosphere, they react with water to form acids, leading to acid rain (pH below 5.6). Acid rain damages buildings, historical monuments, and aquatic life.
Plant Growth: Plants require a specific pH range for optimal growth. If the soil is too acidic or too basic, it can hinder nutrient absorption. Farmers often add basic substances like slaked lime to acidic soil or organic matter to basic soil to adjust the pH.
Self-defence of Animals and Plants: Many organisms use acids or bases for self-defence. For instance, an ant sting injects formic acid, causing pain, which can be neutralized by rubbing baking soda (a base). Nettle leaves inject methanoic acid, causing a burning sensation, which can be relieved by rubbing dock plant leaves (basic).
Strong vs. Weak Acids and Bases:
Strong Acids: Acids that ionize completely in water, producing a high concentration of hydrogen ions (H⁺) or hydronium ions (H₃​O⁺). They are highly corrosive. Examples: Hydrochloric acid (HCl), Sulphuric acid (H_2​SO₄​), Nitric acid (HNO₃​).
Weak Acids: Acids that ionize only partially in water, producing a low concentration of hydrogen ions (H⁺). They are less corrosive. Examples: Acetic acid (CH_3​COOH), Carbonic acid (H₂​CO₃​).
Strong Bases: Bases that dissociate completely in water, producing a high concentration of hydroxide ions (OH⁻). They feel slippery and are corrosive. Examples: Sodium hydroxide (NaOH), Potassium hydroxide (KOH).
Weak Bases: Bases that dissociate only partially in water, producing a low concentration of hydroxide ions (OH⁻). Examples: Ammonium hydroxide (NH₄​OH), Magnesium hydroxide (Mg(OH)₂​).
Basis of Strength: The strength of an acid or base depends on the extent of their ionization/dissociation in water. Complete ionization leads to strong acids/bases, while partial ionization leads to weak acids/bases.
Preparation of Plaster of Paris (PoP):
Preparation: Plaster of Paris (CaSO₄​⋅1/2​ H_2​O) is prepared by heating gypsum (CaSO₄​⋅2H₂​O) at 373K (or 100∘C). When gypsum is heated at this specific temperature, it loses three-fourths of its water of crystallization.
CaSO₄​⋅2H₂​O(s) + 373K → CaSO₄​⋅1/2​ H_2​O(s) + 3/2​ H₂​O(g)
Setting Property: When Plaster of Paris is mixed with water, it rehydrates and sets into a hard, porous mass of gypsum within 5 to 15 minutes. This property makes it useful for various applications.
Uses:
Used by doctors for setting fractured bones in the correct position.
Used for making casts for statues, toys, and decorative materials.
Used as a fire-proofing material and for making surfaces smooth before painting.
Storage: Plaster of Paris should be stored in a moisture-proof container because it readily absorbs moisture from the atmosphere. If it comes into contact with water, it undergoes setting prematurely, forming a hard mass of gypsum, which cannot be used for its intended purposes.
Dilution of Acid/Base and Safety:
Process: Dilution involves adding water to a concentrated acid or base to decrease the concentration of H₃​O⁺ or OH⁻ ions per unit volume.
Safety: It is crucial to add concentrated acid or base slowly to water with constant stirring, and not water to the concentrated acid/base. This is because the dissolution of concentrated acids and bases in water is a highly exothermic process, releasing a large amount of heat. If water is added to a concentrated acid, the large amount of heat generated can cause the mixture to splash out due to rapid boiling, leading to severe burns. Adding the acid/base slowly to water allows the heat to be dissipated gradually.
Effect on Concentration: Upon dilution, the concentration of H₃​O⁺ ions (in acids) or OH⁻ ions (in bases) per unit volume decreases. This means the strength of the acid or base solution decreases.
Types of Indicators and Universal Indicator:
Indicators: Substances that show different colours in acidic and basic media.
Litmus: A natural indicator, turns blue litmus red in acid and red litmus blue in base. It only tells if a solution is acidic or basic.
Phenolphthalein: A synthetic indicator, colourless in acid and pink in base. Useful for strong acid-strong base titrations.
Methyl Orange: A synthetic indicator, red in acid and yellow in base. Useful for strong acid-weak base titrations.
Universal Indicator: A mixture of several different indicators that exhibits different colours over a wide range of pH values (from 0 to 14).
Difference: Unlike litmus or phenolphthalein, which only indicate whether a solution is acidic or basic, a universal indicator provides an approximate pH value of the solution by displaying a specific colour corresponding to a pH value on a colour chart. This allows for a more precise determination of the strength of an acid or a base.
Reasoning and Problem Solving:
Immediate Action: If a student spill concentrated acid on their hand, the immediate action should be to wash the affected area thoroughly with a large amount of running water for several minutes. After washing, a mild base like baking soda solution can be applied to neutralize any remaining acid, followed by medical attention. Scientific Reason: Concentrated acids are highly corrosive and cause severe chemical burns. They are strong dehydrating agents and can rapidly destroy tissues. Washing with water helps to dilute the acid and remove it from the skin, minimizing contact time and preventing further damage. The reaction of acid with water is exothermic, so using a large amount of water helps dissipate the heat.
Identifying Unknown Solutions with Universal Indicator:
Preparation: Take small samples of each unknown solution in separate test tubes.
Test: Add a few drops of universal indicator solution to each test tube, or dip a universal indicator paper strip into each solution.
Observation & Conclusion:
Acidic Solution: The solution will turn a colour corresponding to pH < 7 (e.g., red, orange, yellow-orange). The exact colour will indicate the approximate pH (e.g., bright red for strong acid, orange for weak acid).
Basic Solution: The solution will turn a colour corresponding to pH > 7 (e.g., blue, indigo, violet). The exact colour will indicate the approximate pH (e.g., dark blue/violet for strong base, light blue/greenish-blue for weak base).
Neutral Solution: The solution will turn green, indicating a pH of 7. By comparing the observed colours with the universal indicator’s colour chart, the nature (acidic, basic, neutral) and approximate pH of each solution can be determined.
Compound ‘X’ and ‘Y’:
Compound ‘X’ is Sodium Carbonate (Na₂​CO₃​), also known as washing soda. It is used in glass manufacturing and for removing permanent hardness of water.
When sodium carbonate is heated, it decomposes to give sodium oxide and carbon dioxide gas. So, compound ‘Y’ is Sodium Oxide (Na₂​O). (However, the question states it gives compound Y and carbon dioxide gas, implying Y is the solid residue. In the context of washing soda, it’s usually the anhydrous form, Na₂​CO₃​, that is heated, and it’s stable. If “X” is referring to Baking Soda, it would be NaHCO₃​. Let’s assume X is Washing Soda, and the heating refers to its formation from baking soda, or simply its decomposition if it were a less stable carbonate.)
Revisiting the prompt’s context: “A compound ‘X’ is used in the manufacture of glass and is also used for removing permanent hardness of water.” This strongly points to Sodium Carbonate (Na₂​CO₃​). The decomposition part “When ‘X’ is heated, it gives a compound ‘Y’ and carbon dioxide gas” is a bit tricky for Na₂​CO₃​ as it’s quite stable to heat. However, if we consider the preparation of washing soda from baking soda, heating baking soda (NaHCO₃​) yields Na₂​CO₃​ and CO₂​. But Na₂​CO₃​ itself doesn’t easily decompose to CO₂​ and another compound ‘Y’ (like an oxide) under typical heating conditions in a school lab. Let’s assume the question implies a general carbonate decomposition.
Let’s assume X is Calcium Carbonate (CaCO₃​), which also yields CO₂​ on heating and is used in glass manufacture (as a source of calcium oxide).
X = Calcium Carbonate (CaCO₃​)
Y = Calcium Oxide (CaO)
Chemical Equation: CaCO₃​(s) + Heat → ​CaO(s) + CO₂​(g)↑
Self-correction: The initial prompt states “used for removing permanent hardness of water”, which is a key use of sodium carbonate, not calcium carbonate. So, the original identification of X as Sodium Carbonate is more accurate for the first part of the description. The heating part is the problematic one for sodium carbonate. Given the context of a school chapter, it’s more likely the question intends for X to be Sodium Carbonate and the heating part is either a simplification or refers to a different reaction related to its production, or perhaps a less stable carbonate. However, if Na₂​CO₃​ were to decompose, it would be to Na₂​O and CO₂​, but this requires very high temperatures.
Let’s stick to the most direct interpretation of all given facts. Sodium carbonate (Na₂​CO₃​) is used in glass and water softening. Its decomposition to CO₂​ is not typical under “heating” in a school context, but if forced, it would be Na₂​CO₃​→Na₂​O+CO₂​. So Y would be Na₂​O.
Revised Answer for X and Y based on all clues:
X = Sodium Carbonate (Na₂​CO₃​). (Used in glass manufacture and for removing permanent hardness of water).
Y = Sodium Oxide (Na₂​O).
Chemical Equation: Na₂​CO₃​(s) +Heat → ​Na₂​O(s)+CO₂​(g)↑ (Note: This decomposition requires very high temperatures and is not as common as CaCO3​ decomposition in typical school chemistry contexts, but it fits the description if ‘Y’ is an oxide and CO2​ is evolved).
Solution pH and Dilution: a) A solution containing 10⁻⁵M H⁺ ions is acidic because its hydrogen ion concentration (10⁻⁵M) is greater than that of a neutral solution (10⁻⁷M). b) pH= −log[H⁺] = −log(10⁻⁵) = 5. c) If this solution is diluted 100 times, the new volume will be 100 times the original, so the concentration of H⁺ ions will decrease by 100 times. New [H⁺]=10010⁻⁵M​=10⁻⁵×10⁻²M=10⁻⁷M. New pH=−log(10⁻⁷)=7. The solution becomes neutral.
Hydrogen Gas Preparation:
The student should use zinc granules with dilute sulphuric acid to produce hydrogen gas.
Explanation: Reactive metals like zinc (which are above hydrogen in the reactivity series) displace hydrogen from dilute acids. Copper, being less reactive than hydrogen, will not react with dilute sulphuric acid to produce hydrogen gas.
Chemical Equation: Zn(s) + H₂​SO₄​(aq)→ZnSO₄​(aq) + H₂​(g)↑
Baking Soda vs. Washing Soda:
Baking Soda:
Chemical Formula: Sodium Hydrogen Carbonate (NaHCO₃​)
Preparation: Produced by Solvay’s process:
NaCl+H₂​O+CO₂​+NH₃​→NH_4​Cl+NaHCO₃​
Major Use: Used as an ingredient in baking powder (for leavening in cakes/breads) and as an antacid.
Washing Soda:
Chemical Formula: Sodium Carbonate Decahydrate (Na₂​CO₃​⋅10H₂​O)
Preparation: Formed by heating baking soda to produce anhydrous sodium carbonate, which is then recrystallized from water: 2NaHCO₃​(s) + Heat → ​Na₂​CO₃​(s) + H₂​O(l) + CO₂​(g);
then Na₂​CO₃​(s) + 10H₂​O(l) → Na₂​CO₃​⋅10H₂​O(s)
Major Use: Used as a cleaning agent for domestic purposes and for removing permanent hardness of water.
Copper Sulphate Crystals:
Copper sulphate crystals (CuSO₄​⋅5H₂​O) appear dry even though they contain water because the water molecules are chemically bound within the crystal lattice as water of crystallisation. This water is an integral part of the crystal structure and does not make the salt feel wet.
When water is added to the anhydrous copper sulphate (CuSO₄​), it rehydrates and forms the hydrated copper sulphate crystals (CuSO₄​⋅5H₂​O), which are blue. This phenomenon demonstrates that the water of crystallisation is essential for the specific crystalline structure and colour of the salt.