Acid Bases and Salts Class 10 Notes ICSE — Selina Chapter 3

Acid bases and salts class 10 notes ICSE: this page explains Selina Concise Chemistry Chapter 3, Acids Bases and Salts, in the book’s own order. Each section gets its definition, formula and a worked example pattern. The chapter builds from definitions of acids, bases and salts to pH, salt preparation and water of crystallisation. Keep your textbook open and follow along.

ICSE Class 10 Chemistry Selina Concise Chapter 3 Acids Bases and Salts PDF

What is inside Chapter 3

27 pages · 17 sections · 1 worked examples · 70 exercise questions

• Intext Questions: 14 questions
• Intext Questions: 22 questions
• Intext Questions: 20 questions
• Intext Questions: 9 questions
• EXERCISE: 0 questions
• MISCELLANEOUS QUESTIONS BASED ON ICSE EXAMINATIONS: 5 questions

3.1 INTRODUCTION

Lavoisier and Davy studied acids, bases and salts on scientific lines. Their findings set the definitions this whole chapter uses. Every acid contains hydrogen as a proton \mathrm{H}^{+}. Hydrochloric acid, sulphuric acid and nitric acid are examples. Every base is a metallic oxide, metallic hydroxide or ammonium hydroxide. Copper oxide, sodium hydroxide and ammonium hydroxide fit this definition.

Every salt forms when the hydrogen of an acid is replaced by a metal ion or the ammonium ion \mathrm{NH}_4^{+}. Sodium chloride, calcium carbonate and ammonium phosphate are examples. This replacement idea returns in section 3.11, so hold onto it.

Why it matters: these three findings let you classify any compound. If a compound contains replaceable hydrogen, it is likely an acid. If it is a metallic oxide or hydroxide, it is likely a base. Take sulphuric acid \mathrm{H}_2\mathrm{SO}_4. Its two hydrogens can be replaced by sodium. The product is sodium sulphate \mathrm{Na}_2\mathrm{SO}_4, a salt. That reaction is the whole chapter in one line.

3.2 ACIDS

The word acid comes from Latin acidus, meaning sour. An acid contains one or more hydrogen atoms. Dissolved in water, it produces hydronium ions \mathrm{H}_3\mathrm{O}^{+} as the only positive ions. The bare proton \mathrm{H}^{+} cannot exist alone in water. It combines with a water molecule first.

\mathrm{H}^{+} + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{H}_3\mathrm{O}^{+}

The acidic properties of a solution come from hydronium ions. Four common acids ionise this way: \mathrm{HCl}, \mathrm{H}_2\mathrm{SO}_4, \mathrm{HNO}_3 and \mathrm{CH}_3\mathrm{COOH}. Each produces hydronium ions and a negative ion. For example:

\mathrm{HCl} + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{H}_3\mathrm{O}^{+} + \mathrm{Cl}^{-}

Many foods taste sour because of acids they contain.

Worked pattern: take hydrogen chloride gas. Each molecule donates one proton to a water molecule. The result is one hydronium ion per molecule of acid. That is why the equation for \mathrm{HCl} shows one \mathrm{H}_3\mathrm{O}^{+}.

3.3 CLASSIFICATION OF ACIDS

Acids classify by source, concentration and basicity. By source, organic acids come from plants and contain carbon. Oxalic acid and acetic acid are organic and weak. Inorganic acids come from minerals. Hydrochloric, sulphuric and nitric acids are strong. They ionise fully, giving a high concentration of hydronium ions.

By concentration, a concentrated acid holds little or no water. A dilute acid holds far more water than its own mass. To dilute an acid, pour the acid into water slowly with constant stirring. Never pour water into acid. The process is exothermic. So much heat is released that the solution may splash or crack the container.

Concentration is not strength. Strength depends on the degree of ionisation \alpha:

\alpha = \frac{\text{molecules ionised}}{\text{total molecules in solution}} \times 100

If \alpha exceeds 30%, the acid is strong. Below 30%, it is weak. Worked pattern: suppose 45 of 100 acid molecules ionise in water. Then \alpha = 45\%, which is strong. This is why dilute hydrochloric acid is stronger than highly concentrated acetic acid.

Basicity counts the hydronium ions produced per molecule. Monobasic acids give one, like \mathrm{HCl} and acetic acid. Dibasic acids give two, like \mathrm{H}_2\mathrm{SO}_4 and \mathrm{H}_2\mathrm{CO}_3. Tribasic acids give three, like \mathrm{H}_3\mathrm{PO}_4.

Basicity counts ionisable hydrogens, not total hydrogens. Acetic acid has four hydrogen atoms but only one ionises. So it is monobasic. Similarly, \mathrm{H}_3\mathrm{PO}_3 is dibasic, because only the hydrogens attached to oxygen are replaceable. These two cases are the classic exam traps.

3.4 PREPARATION OF ACIDS

Four laboratory routes make acids. First, by synthesis. Hydrogen combines directly with a non-metal to form a binary acid:

\mathrm{H}_2 + \mathrm{Cl}_2 \rightarrow 2\mathrm{HCl}

Similar reactions give hydrogen bromide and hydrogen sulphide. Second, by the action of water on acidic oxides. Acidic oxides, also called acidic anhydrides, dissolve in water to give oxy-acids. Only oxy-acids form this way.

Third, by oxidation of non-metals. Concentrated nitric acid oxidises sulphur to sulphuric acid:

\mathrm{S} + 6\mathrm{HNO}_3 \rightarrow \mathrm{H}_2\mathrm{SO}_4 + 2\mathrm{H}_2\mathrm{O} + 6\mathrm{NO}_2

Fourth, by displacement. A less volatile acid displaces a more volatile one from its salt. Concentrated sulphuric acid displaces hydrogen chloride from sodium chloride:

\mathrm{NaCl} + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{NaHSO}_4 + \mathrm{HCl}

One oxide breaks the pattern. Nitrogen dioxide is a double acid anhydride. With water it forms two acids at once:

2\mathrm{NO}_2 + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{HNO}_2 + \mathrm{HNO}_3

Worked pattern: carbon dioxide is an acidic oxide. Bubbling it into water gives carbonic acid. That reaction is what makes soft drinks fizzy.

3.5 PROPERTIES OF ACIDS

Acids taste sour. Mineral acids are corrosive, so never taste them. Concentrated sulphuric acid blackens skin. Concentrated nitric acid turns it yellow. Acids change indicator colours.

Acids are electrolytes. They conduct electricity in aqueous solution. Chemically, acids react with active metals to give salt and hydrogen. Zinc with dilute hydrochloric acid:

\mathrm{Zn} + 2\mathrm{HCl} \rightarrow \mathrm{ZnCl}_2 + \mathrm{H}_2 \uparrow

Nitric acid is never used for hydrogen. It is a strong oxidising agent and oxidises the hydrogen to water. Only magnesium and manganese work, with very dilute 1% nitric acid.

Acids neutralise bases to salt and water. They liberate carbon dioxide from carbonates and bicarbonates. They give sulphur dioxide from sulphites and hydrogen sulphide from sulphides. Chlorides react only with concentrated sulphuric acid, giving hydrogen chloride.

Two edge cases matter. Lead carbonate does not react with dilute hydrochloric acid, because lead chloride is insoluble. Calcium carbonate does not react with dilute sulphuric acid, because calcium sulphate coats the solid and stops the reaction.

Acid rain forms when sulphur and nitrogen oxides react with rain water:

2\mathrm{SO}_2 + \mathrm{O}_2 + 2\mathrm{H}_2\mathrm{O} \rightarrow 2\mathrm{H}_2\mathrm{SO}_4

4\mathrm{NO}_2 + \mathrm{O}_2 + 2\mathrm{H}_2\mathrm{O} \rightarrow 4\mathrm{HNO}_3

These acids damage plant leaves, corrode pipes and eat into stone buildings.

3.6 GENERAL USES OF SOME ACIDS

This table answers one-mark and matching questions directly. Citric acid preserves food. Oxalic acid lifts ink stains. Hydrochloric acid cleans metals. Nitric acid makes explosives. Phosphoric acid feeds fertilisers. Cover the use column and test yourself after reading.

For the rest of the chapter list in this book, see the Class 10 Chemistry hub.

3.7 BASES

A base is a metallic oxide, metallic hydroxide or ammonium hydroxide. It reacts with the hydronium ions of an acid to form salt and water only. The word only is deliberate. It excludes compounds that produce anything extra. Copper oxide and magnesium hydroxide are bases:

\mathrm{CuO} + 2\mathrm{HCl} \rightarrow \mathrm{CuCl}_2 + \mathrm{H}_2\mathrm{O}

\mathrm{Mg}(\mathrm{OH})_2 + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{MgSO}_4 + 2\mathrm{H}_2\mathrm{O}

A basic oxide is a metallic oxide containing \mathrm{O}^{2-}. It reacts with acid to give salt and water only. A basic hydroxide is a metallic hydroxide containing \mathrm{OH}^{-}. It also gives salt and water only. Sodium hydroxide is a basic hydroxide:

\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H}_2\mathrm{O}

Lead dioxide \mathrm{PbO}_2 is a metallic oxide. But it is not a base:

\mathrm{PbO}_2 + 4\mathrm{HCl} \rightarrow \mathrm{PbCl}_2 + \mathrm{Cl}_2 + 2\mathrm{H}_2\mathrm{O}

It produces chlorine as well as salt and water. The word only excludes it. This is the edge case examiners repeat. Without the word only, lead dioxide would slip into the definition.

3.8 ALKALIS

An alkali is a base that dissolves in water. It produces hydroxide ions \mathrm{OH}^{-} as the only negative ions:

\mathrm{NaOH} \rightleftharpoons \mathrm{Na}^{+} + \mathrm{OH}^{-}

The four common alkalis are sodium hydroxide (caustic soda), potassium hydroxide (caustic potash), calcium hydroxide (slaked lime) and ammonium hydroxide. The first two are strong alkalis. They ionise completely in water. The last two are weak alkalis. They ionise only slightly.

The rule to memorise: all alkalis are bases, but all bases are not alkalis. Ferric hydroxide \mathrm{Fe}(\mathrm{OH})_3 and cupric hydroxide \mathrm{Cu}(\mathrm{OH})_2 are bases. They are insoluble in water, so they are not alkalis. This rule appears directly in ICSE questions. Learn it with the two insoluble examples attached.

3.9 CLASSIFICATION OF BASES

Strong bases dissociate almost completely in water. Sodium hydroxide and potassium hydroxide are strong. Weak bases dissociate only partly. Ammonium hydroxide and calcium hydroxide are weak:

\mathrm{NaOH} \rightleftharpoons \mathrm{Na}^{+} + \mathrm{OH}^{-}

\mathrm{NH}_4\mathrm{OH} \rightleftharpoons \mathrm{NH}_4^{+} + \mathrm{OH}^{-}

Acidity is the number of hydroxide ions a base produces per molecule. Monoacidic bases give one, like \mathrm{NaOH}. Diacidic bases give two, like \mathrm{Ca}(\mathrm{OH})_2 and \mathrm{Cu}(\mathrm{OH})_2. Triacidic bases give three, like \mathrm{Al}(\mathrm{OH})_3 and \mathrm{Fe}(\mathrm{OH})_3.

Bases form by several routes. Metals burn in oxygen to give basic oxides. Reactive metals react with water, giving alkali and hydrogen. Metal oxides react with water. Double decomposition precipitates insoluble hydroxides. Carbonates and nitrates decompose on heating. Ammonia dissolves in water to give ammonium hydroxide:

\mathrm{NH}_3 + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{NH}_4\mathrm{OH}

Bases taste bitter and feel soapy. Strong alkalis absorb carbon dioxide from air. They neutralise acids. They precipitate insoluble hydroxides from heavy metal salts. Warmed with ammonium salts, they release ammonia gas.

Amphoteric hydroxides dissolve in excess alkali. Zinc hydroxide dissolves in excess sodium hydroxide:

\mathrm{Zn}(\mathrm{OH})_2 + 2\mathrm{NaOH} \rightarrow \mathrm{Na}_2\mathrm{ZnO}_2 + 2\mathrm{H}_2\mathrm{O}

Common uses: sodium hydroxide in soap, magnesium hydroxide as an antacid, aluminium hydroxide in fire extinguishers, calcium hydroxide in bleaching powder.

3.10 STRENGTH OF ACIDIC AND BASIC SOLUTIONS

Pure water ionises only slightly. Its hydrogen and hydroxide ion concentrations are each 1 \times 10^{-7} mol/L at 25°C. Their product is the ionic product of water:

K_w = [katex display="true"]\mathrm{H}_3\mathrm{O}^{+}\mathrm{OH}^{-} = 1 \times 10^{-14}[/katex]

Adding acid raises the hydronium concentration. Adding alkali lowers it. A Danish biochemist devised the pH scale in 1909. The p stands for Potenz, German for power. pH means the power of hydrogen. It is the negative logarithm of the hydrogen ion concentration:

\mathrm{pH} = -\log_{10}[katex display="true"]\mathrm{H}^{+}[/katex]

Worked pattern: take a solution with [katex display="true"]\mathrm{H}^{+} = 10^{-4}[/katex] mol/L.

\mathrm{pH} = -\log_{10}(10^{-4}) = 4

Now take another with [katex display="true"]\mathrm{H}^{+} = 10^{-9}[/katex] mol/L.

\mathrm{pH} = -\log_{10}(10^{-9}) = 9

The first solution is acidic. The second is basic. The scale runs from 0 to 14. pH 7 is neutral. Values below 7 are acidic; the lower the number, the stronger the acid. Values above 7 are alkaline.

A universal indicator gives a colour spectrum. Green means neutral. Blue to violet means alkaline. Yellow to red means acidic. pH matters in daily life. The body works within 7.0 to 7.8. Crops need specific soil pH. Antacids like milk of magnesia fix excess stomach acid. Tooth decay starts near pH 5.5. Baking soda, a base, relieves bee stings. You can verify the syllabus and exam format on the official CISCE board site.

3.11 SALT

A salt forms by partial or total replacement of the ionisable hydrogen of an acid. The replacement is by a metal ion or the ammonium ion. Complete replacement gives a normal salt. Partial replacement gives an acid salt. Sulphuric acid with sodium hydroxide shows both:

\mathrm{H}_2\mathrm{SO}_4 + \mathrm{NaOH} \rightarrow \mathrm{NaHSO}_4 + \mathrm{H}_2\mathrm{O}

\mathrm{H}_2\mathrm{SO}_4 + 2\mathrm{NaOH} \rightarrow \mathrm{Na}_2\mathrm{SO}_4 + 2\mathrm{H}_2\mathrm{O}

There are six kinds of salts:

• Normal salt — \mathrm{NaCl}
• Acid salt — \mathrm{NaHSO}_4
• Basic salt — \mathrm{Pb}(\mathrm{OH})\mathrm{Cl}
• Double salt — potash alum \mathrm{K}_2\mathrm{SO}_4\cdot\mathrm{Al}_2(\mathrm{SO}_4)_3\cdot24\mathrm{H}_2\mathrm{O}
• Mixed salt — bleaching powder \mathrm{CaOCl}_2
• Complex salt — \mathrm{Na}[katex display="true"]\mathrm{Ag}(\mathrm{CN})_2[/katex]

Solubility rules carry exceptions. All sodium, potassium and ammonium compounds dissolve. All nitrates dissolve. Most chlorides dissolve, except \mathrm{AgCl}, \mathrm{PbCl}_2 and \mathrm{Hg}_2\mathrm{Cl}_2. Most sulphates dissolve, except \mathrm{PbSO}_4, \mathrm{BaSO}_4 and \mathrm{CaSO}_4. Carbonates, sulphides, sulphites and phosphates are insoluble except for ammonium, sodium and potassium salts.

Salts react with water by hydrolysis. A salt of a strong base and weak acid gives an alkaline solution. Sodium carbonate does this. A salt of a strong acid and weak base gives an acidic solution. Ferric chloride and copper sulphate do this. Salts of strong acid and strong base, like \mathrm{NaCl}, stay neutral.

Salts form by direct combination, displacement, decomposition, neutralisation and precipitation. Insoluble salts need double decomposition. Lead sulphate is the edge case. You cannot add dilute sulphuric acid directly to lead carbonate. An insoluble crust of lead sulphate forms and stops the reaction. First convert lead carbonate to soluble lead nitrate with nitric acid. Then add sulphuric acid:

\mathrm{Pb}(\mathrm{NO}_3)_2 + \mathrm{H}_2\mathrm{SO}_4 \rightarrow \mathrm{PbSO}_4 \downarrow + 2\mathrm{HNO}_3

3.12 LABORATORY PREPARATION OF SOME SALTS (NORMAL AND ACID SALTS)

Anhydrous iron (III) chloride is made by synthesis. Dry chlorine passes over heated iron in a combustion tube. The reaction is exothermic. Iron turns red hot. Iron (III) chloride volatilises and condenses as a brown scale:

2\mathrm{Fe} + 3\mathrm{Cl}_2 \rightarrow 2\mathrm{FeCl}_3

It is deliquescent, so store it with fused calcium chloride. Never heat hydrated ferric chloride. It gives iron oxide and hydrogen chloride instead of the anhydrous salt.

Copper (II) sulphate pentahydrate comes from dilute sulphuric acid on copper oxide, hydroxide or carbonate. Heat the acid, add the copper compound until excess settles. Filter hot, evaporate to the point of crystallisation, then cool. Bright blue crystals of \mathrm{CuSO}_4\cdot5\mathrm{H}_2\mathrm{O} form.

Zinc sulphate and iron (II) sulphate come from dilute sulphuric acid on the active metal. Zinc gives white vitriol \mathrm{ZnSO}_4\cdot7\mathrm{H}_2\mathrm{O}. Iron gives green vitriol \mathrm{FeSO}_4\cdot7\mathrm{H}_2\mathrm{O}. Add the metal until effervescence stops. That means all the acid is used up. Filter off excess metal, evaporate and crystallise.

Lead chloride and calcium carbonate are made by precipitation. Mix lead nitrate with sodium chloride. White lead chloride precipitates. It dissolves in hot water and recrystallises on cooling. Calcium carbonate comes from calcium chloride and sodium carbonate. Sodium bicarbonate forms by passing carbon dioxide into cold sodium carbonate:

\mathrm{Na}_2\mathrm{CO}_3 + \mathrm{CO}_2 + \mathrm{H}_2\mathrm{O} \rightarrow 2\mathrm{NaHCO}_3

Sodium sulphate crystals use titration. Sodium hydroxide goes into a flask with phenolphthalein. Acid runs in from a burette until the pink colour vanishes. That is the end point. Mix the reactants in that ratio and evaporate. For the acid salt \mathrm{NaHSO}_4, double the acid. One molecule of alkali takes one molecule of acid for the normal salt, but two for the acid salt. This chapter leads into Chapter 4, Analytical Chemistry.

3.13 WATER OF CRYSTALLISATION

Some salts unite with a fixed quantity of water when they crystallise. This is water of crystallisation. It is held in loose chemical combination. Heating the powdered crystals above 100°C drives it off.

Experiment: heat powdered copper sulphate in a tilted test tube. Blue crystals turn to white powder. Water condenses on the cooler parts. Test the liquid three ways. It is neutral to litmus. It turns anhydrous copper sulphate blue. It turns blue cobalt chloride paper pale pink:

\mathrm{CuSO}_4\cdot5\mathrm{H}_2\mathrm{O} \xrightarrow{\text{heat}} \mathrm{CuSO}_4 + 5\mathrm{H}_2\mathrm{O}

Not every crystalline solid holds water. Common salt decrepitates on heating. It crackles and breaks into smaller crystals. Potassium nitrate melts and releases oxygen:

2\mathrm{KNO}_3 \rightarrow 2\mathrm{KNO}_2 + \mathrm{O}_2

3.14 EFFLORESCENCE

Efflorescence is the loss of part or all of the water of crystallisation in dry air. The crystals turn powdery. The drier the air, the faster the loss.

Three examples matter. Washing soda loses nine of its ten water molecules:

\mathrm{Na}_2\mathrm{CO}_3\cdot10\mathrm{H}_2\mathrm{O} \rightarrow \mathrm{Na}_2\mathrm{CO}_3\cdot\mathrm{H}_2\mathrm{O} + 9\mathrm{H}_2\mathrm{O}

Glauber’s salt loses all ten:

\mathrm{Na}_2\mathrm{SO}_4\cdot10\mathrm{H}_2\mathrm{O} \rightarrow \mathrm{Na}_2\mathrm{SO}_4 + 10\mathrm{H}_2\mathrm{O}

Epsom salt loses six of its seven:

\mathrm{MgSO}_4\cdot7\mathrm{H}_2\mathrm{O} \rightarrow \mathrm{MgSO}_4\cdot\mathrm{H}_2\mathrm{O} + 6\mathrm{H}_2\mathrm{O}

Why it happens: dry air absorbs water vapour. Higher temperature and lower moisture increase efflorescence. The crystals crumble because the water that held their structure is gone.

3.15 HYGROSCOPIC SUBSTANCES

Hygroscopic substances absorb moisture from the air without dissolving in it. The phenomenon is hygroscopy. Four examples matter: concentrated sulphuric acid, phosphorus pentoxide, quicklime and silica gel.

In the laboratory, these substances dry gases. They remove moisture from gases passed over them. Almost all hygroscopic substances act as desiccating agents. This connects directly to section 3.17.

The key difference from deliquescence: a hygroscopic substance stays solid. It does not form a solution. It simply holds the absorbed water in its pores or on its surface.

3.16 DELIQUESCENT SUBSTANCES

Deliquescent substances are water-soluble solids. They absorb moisture from the air and dissolve in it. They form a saturated solution. The phenomenon is deliquescence.

The examples include caustic soda, caustic potash, magnesium chloride, zinc chloride, calcium chloride, ferric chloride, zinc nitrate and copper nitrate. Deliquescence happens when the salt’s vapour pressure is much lower than the air’s. Dry conditions reduce deliquescence but increase efflorescence.

The weight rule is quick and useful. Efflorescent substances lose weight in air. Hygroscopic and deliquescent substances gain weight. A question that asks whether a sample gains or loses mass is testing exactly this rule.

Pure sodium chloride is not deliquescent. Commercial table salt turns wet in the rainy season. It contains impurities — magnesium chloride and calcium chloride — which are deliquescent. Remove them by passing dry hydrogen chloride gas through the saturated solution. Pure sodium chloride precipitates out and is recovered by filtering.

3.17 DRYING AND DEHYDRATING AGENTS

A drying agent, or desiccant, removes moisture from other substances. This is a physical change. A dehydrating agent removes the elements of water — hydrogen and oxygen in the 2:1 ratio — from a compound. This is a chemical change.

Concentrated sulphuric acid does both jobs. It dries gases by taking moisture. It dehydrates blue vitriol by removing its water of crystallisation:

\mathrm{CuSO}_4\cdot5\mathrm{H}_2\mathrm{O} \rightarrow \mathrm{CuSO}_4 + 5\mathrm{H}_2\mathrm{O}

So exam answers should call it both a drying and a dehydrating agent. The distinction turns on what is removed. Moisture is free water on a surface. Combined water is locked inside a compound’s structure. Drying removes the first. Dehydration breaks the second.

What to practise from this chapter

Work through these exercise blocks in the book, in order, and time each one.

• Intext Questions after Section 3.6 — 14 questions
• Intext Questions after Section 3.10 — 22 questions
• Intext Questions after Section 3.12 — 20 questions
• Intext Questions after Section 3.17 — 9 questions
• Miscellaneous Questions Based on ICSE Examinations — 5 questions

The miscellaneous set draws on ICSE board questions from 2008 to 2014. It repeatedly tests salt preparation methods, matching salts to their preparation routes, and identifying acid salts. Those three skills carry the most marks in this chapter.

Frequently asked questions

Why is water never added to acid when diluting it?

Dilution is exothermic. Adding water to acid releases so much heat that the solution may splash or crack the container. Pouring acid slowly into water with constant stirring spreads the heat safely.

Why is acetic acid called a monobasic acid even though it has four hydrogen atoms?

Basicity counts ionisable hydrogen atoms, not total hydrogen atoms. Acetic acid produces only one hydronium ion per molecule. Only the hydrogen attached to oxygen is replaceable.

Why is lead dioxide not classified as a base?

The base definition requires salt and water only. Lead dioxide reacts with HCl to produce chlorine as well: \mathrm{PbO}_2 + 4\mathrm{HCl} \rightarrow \mathrm{PbCl}_2 + \mathrm{Cl}_2 + 2\mathrm{H}_2\mathrm{O}. The extra chlorine breaks the only rule.

What is the difference between a base and an alkali?

An alkali is a base that dissolves in water. All alkalis are bases, but insoluble bases like \mathrm{Fe}(\mathrm{OH})_3 and \mathrm{Cu}(\mathrm{OH})_2 are not alkalis.

Why does common salt turn wet during the rainy season?

Commercial salt contains magnesium chloride and calcium chloride impurities. These are deliquescent and absorb moisture. Passing dry hydrogen chloride gas through the saturated solution removes them.

Reference: ICSE Class 10 Chemistry textbook material, chapter 3: Acids Bases and Salts — CISCE-aligned.

More from this book

• Chapter 4: Analytical Chemistry
• Chapter 6: Electrolysis
• Chapter 12: Organic Chemistry