Water Class 9 ICSE — Chemistry Chapter 3

Water class 9 icse is the third chapter in your Class 9 Chemistry book. It explains why water matters, how it dissolves so many substances, and what makes hard water different from soft water. The chapter also covers solutions, solubility, crystals, and the agents used for drying. Each section below follows the book’s own order, with formulas and worked patterns.

Download Water Class 9 ICSE Chemistry Chapter 3 PDF

Download the Water Class 9 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 18 pages, 25 sections.

What is inside Chapter 3

18 pages · 25 sections · 6 worked examples

3.1 INTRODUCTION

Water covers more than 70% of Earth’s area, but only about 2.5% of it is fresh water. That fresh supply includes the ice locked in polar caps and glaciers. Water exists in three states: solid ice, liquid water, and gaseous water vapour. Nearly 70% of our body weight is water.

Water occurs free and combined in nature. Under normal pressure of 760 mm Hg, it boils at 100°C and changes into steam. When steam condenses back to water, it releases the latent heat of condensation. This heat equals the latent heat of vaporization. The sudden release of this heat drives the violence of torrential rain. In the Class 9 ICSE syllabus, this idea returns when you study the water cycle.

3.2 PHYSICAL PROPERTIES OF WATER

Pure water is clear, colourless, odourless and tasteless. The taste of ordinary water comes from dissolved gases and solids. Boiling point is 100°C at normal pressure. Higher pressure raises the boiling point, which is why a pressure cooker works. On Mount Everest, low pressure makes water boil near 70°C, so food cooks poorly.

Pure water freezes at 0°C. Its density is 1 g/cm³, with minimum volume at 4°C. Below 4°C, water expands anomalously until it freezes. This keeps ponds liquid below the surface ice, so fish survive winter. The latent heat of fusion of ice is 336 J/g, and the latent heat of vaporization is 2268 J/g. That high value explains why steam burns are worse than boiling water burns.

The specific heat capacity is 4.2 J/g°C, which moderates coastal climates and makes water a good coolant. In this ICSE chapter, these constants are the ones exam questions quote.

3.3 WATER AS UNIVERSAL SOLVENT

Water dissolves many substances, so it is called the universal solvent. It has a high dielectric constant, which weakens the electrostatic pull between ions. This lets electrovalent compounds dissolve. Even glass dissolves in traces when water is stored in it for long.

Two experiments confirm dissolved matter. In the watch glass test, tap water is evaporated and leaves concentric rings of dissolved solids. In the heating test, boiling tap water expels dissolved gases. The air dissolved in water holds 33% oxygen, against 21% in ordinary air. One cubic decimetre of water contains nearly 40 cm³ of dissolved oxygen. Dissolved carbon dioxide drives photosynthesis and converts limestone to calcium bicarbonate:

6\mathrm{CO}_2 + 12\mathrm{H}_2\mathrm{O} \underset{\text{sunlight}}{\stackrel{\text{chlorophyll}}{\rightleftharpoons}} \mathrm{C}_6\mathrm{H}_{12}\mathrm{O}_6 + 6\mathrm{O}_2 + 6\mathrm{H}_2\mathrm{O}

Distilled and boiled water taste flat because their dissolved gases have been removed. This section is often tested in the Class 9 Chemistry book exercises.

3.4 SOLUTIONS AS MIXTURES OF SOLIDS IN WATER

A solution is a homogeneous mixture of two or more components. The solute is the substance that dissolves; the solvent is the medium that dissolves it. A two-component solution is binary, three-component is ternary, and four-component is quaternary. This chapter deals with binary solutions.

Choose the solvent by rule: if the components differ in state, the solvent is the one in the same phase as the final solution. If they share a state, the solvent is the one in excess. Solutions form in four ways: liquid in liquid, solid in liquid, gas in liquid, and solid in solid. Alloys like brass are solid solutions. A true solution has solute particles about 10^{-10} m in size.

It is clear, does not scatter light, and cannot be filtered. When the solvent is alcohol, the solution is called a tincture.

3.5 SATURATED SOLUTION

A saturated solution cannot dissolve more solute at a given temperature. The book shows this with nitre: add one gram at a time and stir. A point arrives where no more nitre dissolves, however long you stir. That solution is saturated at that temperature.

Cooling a saturated solution makes some solute separate as crystals. An unsaturated solution can still take more solute. A supersaturated solution holds more solute than its saturated form at that temperature. You prepare it by dissolving extra solute in warm water and cooling gently. Crystals do not form if the solution is free of dust and is left undisturbed.

3.6 CONCENTRATION OF A SOLUTION

Concentration is the amount of solute in a given quantity of solution. Mass percent is used when the solute is solid and the solvent is liquid:

\text{Mass}\% = \frac{\text{Mass of solute}}{\text{Mass of solution}} \times 100

Volume percent is used when both are liquids:

\text{Volume}\% = \frac{\text{Volume of solute}}{\text{Volume of solution}} \times 100

The denominator is the solution mass, not the solvent mass alone. That is the most common error in this section. Here is a fresh worked pattern. Suppose 25 g of salt dissolves in 175 g of water. The solution mass is 200 g. Then

\text{Mass}\% = \frac{25}{200} \times 100 = 12.5\%

Use your own numbers when practising this ICSE chapter, and always check which denominator the question gives.

3.7 SOLUBILITY

Solubility is the maximum grams of solute needed to saturate 100 g of solvent at a given temperature. Different solids dissolve to different extents. Silver chloride is nearly insoluble, calcium hydroxide is sparingly soluble, and sodium chloride is freely soluble.

The book measures solubility with an evaporating dish. Weigh the empty dish as M, the dish plus saturated solution as M₁, and after evaporation the dish plus residue as M₂. Then

\text{Solubility} = \frac{\text{Mass of solute}}{\text{Mass of solvent}} \times 100 = \frac{\mathrm{M}_2 - \mathrm{M}}{(\mathrm{M}_1 - \mathrm{M}) - (\mathrm{M}_2 - \mathrm{M})} \times 100

3.7.1 Factors affecting solubility

Three factors matter for a solid dissolving in a liquid. Smaller solute particles expose more surface area, so they dissolve faster. Stirring brings fresh solvent into contact with the solute. Temperature usually raises the solubility of solids. Potassium nitrate rises from 31.6 g at 20°C to 63.9 g at 60°C. Gases behave differently: their solubility falls as temperature rises.

3.7.2 Solubility curves

A solubility curve plots solubility against temperature. Solubility sits on the y-axis and temperature on the x-axis. Most curves rise, but not all. Sodium chloride rises only slightly. Calcium sulphate falls after a certain temperature. Sodium nitrate and potassium nitrate rise steeply.

Some salts show anomalous solubility. Sodium sulphate decahydrate (\mathrm{Na}_2\mathrm{SO}_4 \cdot 10\mathrm{H}_2\mathrm{O}) rises until 32.8°C and then falls. The reason is that it is hydrated below 32.8°C and anhydrous above it. In an endothermic process, solubility rises with temperature; in an exothermic one, it falls.

3.7.3 Applications of solubility curves

Solubility curves let you read the solubility of a substance at any temperature. You can compare different solutes on one chart. You can also predict what happens when a hot saturated solution cools. Some salts precipitate; others stay dissolved.

Here is a worked pattern with original numbers. Suppose 15 g of a saturated solution of a salt at 30°C leaves 5 g of residue when dried. The water mass is 10 g. Solubility is (5/10) \times 100 = 50 g per 100 g water.

Common mistakes when reading these curves:

• Mistake: assuming every curve rises with temperature. Correction: calcium sulphate falls after a point, and sodium sulphate decahydrate falls after 32.8°C.
• Mistake: reading solubility as grams per solution. Correction: solubility is grams per 100 g of solvent.
• Mistake: confusing a steep curve with high solubility. Correction: a steep curve means a large change with temperature, not a large absolute value.

3.7.4 Effect of pressure and temperature on solubility of gases in water (liquids)

Pressure raises the solubility of a gas in a liquid. At a fixed temperature, the mass of gas dissolved by a fixed volume of liquid is directly proportional to the pressure on the liquid. This is Henry’s law. Soda water is the classic example. Carbon dioxide dissolves under high pressure and bubbles out when the bottle opens and pressure drops.

Temperature works in the opposite direction. Heating water lowers gas solubility, so boiled water loses its taste. The pleasant taste of drinking water comes from dissolved air, carbon dioxide and minerals.

3.8 CRYSTALS AND CRYSTALLIZATION

A crystal is a homogeneous solid with a definite geometrical shape and smooth plane faces. Crystallization is the process of obtaining crystals by cooling a hot saturated solution. Crystal shapes differ: cubic for sodium chloride, octahedral for iron(II) sulphate, and prismatic for potassium nitrate.

Four laboratory methods make crystals: cooling a hot saturated solution, slow evaporation of a saturated solution, cooling a fused mass, and sublimation. The copper sulphate experiment builds large crystals. A saturated solution is made at 80°C and filtered. A small well-shaped crystal is suspended in the cooled solution as a seed. This is called seeding. Keep dust out, or crystals will form around the dust particles.

3.9 HYDRATED AND ANHYDROUS SUBSTANCES

Substances that carry water molecules within their crystals are hydrated. Sodium carbonate decahydrate (\mathrm{Na}_2\mathrm{CO}_3 \cdot 10\mathrm{H}_2\mathrm{O}) and copper sulphate pentahydrate (\mathrm{CuSO}_4 \cdot 5\mathrm{H}_2\mathrm{O}) are examples. Substances with no such water are anhydrous, such as sodium chloride and potassium nitrate.

3.9.1 Hydrated substances

The fixed water that forms an integral part of a hydrated crystal is water of crystallization. It gives crystals their shape and sometimes their colour. Copper sulphate is blue because of it. Common hydrated salts include washing soda (\mathrm{Na}_2\mathrm{CO}_3 \cdot 10\mathrm{H}_2\mathrm{O}), Epsom salt (\mathrm{MgSO}_4 \cdot 7\mathrm{H}_2\mathrm{O}), blue vitriol (\mathrm{CuSO}_4 \cdot 5\mathrm{H}_2\mathrm{O}), plaster of Paris (\mathrm{CaSO}_4 \cdot \frac{1}{2}\mathrm{H}_2\mathrm{O}), and gypsum (\mathrm{CaSO}_4 \cdot 2\mathrm{H}_2\mathrm{O}).

Heating blue vitriol drives off this water. The blue crystals turn into a white powder. Adding a few drops of water restores the blue colour. To find the percentage of water of crystallization, weigh the crystals before heating (a) and after heating (b). Then

\%\ \text{water of crystallization} = \frac{a - b}{a} \times 100

Dry the crystals gently. Strong heating gives a dehydrated powder, not proper crystals.

3.9.2 Anhydrous substances

An anhydrous substance contains no water. The term often applies to salts whose water of crystallization has been removed. Table salt is anhydrous sodium chloride, and gaseous hydrogen chloride is anhydrous. Anhydrous salts can be made by direct heating, heating in dry air, heating under vacuum, or using a dehydrating agent like warm concentrated sulphuric acid.

Crystalline shape is not always due to water of crystallization. Common salt, nitre, sugar, potassium permanganate and ammonium chloride crystallize from water without holding any. Heating these crystals shows the difference. Sodium chloride cracks with a sound called decrepitation. Potassium nitrate melts, then decomposes:

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

3.10 PROPERTIES

Section 3.10 groups three phenomena where substances react with moisture in the air. Each is covered in turn below.

3.10.1 Efflorescence

Efflorescence is the loss of water of crystallization when a hydrated salt is exposed to dry air. The crystal crumbles into a powder. It happens when the vapour pressure in the crystals exceeds the vapour pressure of the air. Humid air slows it down; warm dry air speeds it up.

Washing soda becomes a monohydrate, losing nine water molecules. Glauber’s salt becomes anhydrous sodium sulphate. Epsom salt becomes a monohydrate. All three are efflorescent substances.

3.10.2 Deliquescence

Deliquescent substances absorb moisture from the air and dissolve in it, forming a solution. This happens when the vapour pressure inside the crystals is much lower than in the air. Dry conditions slow it down. Examples are caustic soda (NaOH), caustic potash (KOH), calcium chloride (CaCl₂) and zinc chloride (ZnCl₂).

Table salt turns sticky in the rainy season because commercial salt contains deliquescent impurities like magnesium chloride and calcium chloride. Pure sodium chloride is not deliquescent.

3.10.3 Hygroscopic Substances

Hygroscopic substances absorb moisture from the air without dissolving in it. Examples are concentrated sulphuric acid, phosphorus pentoxide (\mathrm{P}_2\mathrm{O}_5) and quicklime (CaO). They stay solid or liquid while drawing water vapour.

Mnemonic to recall the groups: efflorescent salts start with WEG — Washing soda, Epsom salt, Glauber’s salt. Deliquescent ones are KCZ — KOH, CaCl₂, ZnCl₂. Hygroscopic ones are HPQ — concentrated H₂SO₄, P₂O₅, quicklime.

3.11 DRYING AND DEHYDRATING AGENTS

Drying agents remove moisture without any chemical reaction. They are also called desiccants. Anhydrous calcium chloride, phosphorus pentoxide and concentrated sulphuric acid are examples. Dehydrating agents remove chemically combined water. Concentrated sulphuric acid removes the water of crystallization from blue vitriol, so it acts as both.

Gases are dried by passing them through concentrated sulphuric acid, a drying tower with anhydrous sodium sulphate, or a drying bulb with calcium chloride. Quicklime suits ammonia because it is basic and will not react with a basic gas. Liquids are kept overnight over anhydrous salts. Solids are dried in a desiccator, an airtight vessel with a drying agent at the bottom.

3.12 SOFT AND HARD WATER

Soft water readily forms lather with soap. Distilled water and rain water are soft. Hard water does not form lather readily. The difference matters for washing and for boilers, which the chapter examines next.

3.12.1 Causes of hardness

Hardness comes from dissolved mineral matter. Bicarbonates, sulphates and chlorides of calcium and magnesium make water hard. When water flows over limestone and dolomite, carbon dioxide converts some into soluble bicarbonates:

\mathrm{CaCO}_3 + \mathrm{CO}_2 + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{Ca(HCO}_3)_2

This reaction also forms stalactites and stalagmites in caves. Drops of water carry dissolved calcium hydrogen carbonate. On release of pressure, it deposits calcium carbonate:

\mathrm{Ca(HCO}_3)_2 \rightarrow \mathrm{CaCO}_3 + \mathrm{CO}_2 + \mathrm{H}_2\mathrm{O}

A stalactite hangs from the roof; a stalagmite rises from the floor. Water flowing over gypsum also picks up a little calcium sulphate and becomes hard.

3.13 TYPES OF HARDNESS TEMPORARY AND PERMANENT HARDNESS

Temporary hardness comes from bicarbonates of calcium and magnesium. Boiling removes it. The bicarbonates break down into insoluble carbonates:

\mathrm{Ca(HCO}_3)_2 \xrightarrow{\text{Boil}} \mathrm{CaCO}_3 \downarrow + \mathrm{H}_2\mathrm{O} + \mathrm{CO}_2 \uparrow

Permanent hardness comes from sulphates and chlorides of calcium and magnesium. Boiling does not remove it. Hard water has advantages: it tastes better, supplies calcium for bones and teeth, and forms an insoluble layer in lead pipes that checks lead poisoning. Hardness of water is the last major topic in this ICSE chapter.

3.13.2 Disadvantages of hard water

Hard water harms boilers and washing. In boiler tubes, dissolved solids deposit as boiler scale, a bad conductor of heat. The scale narrows the tubes, and the steam pressure may rise until the boiler bursts.

For washing, soap reacts with calcium and magnesium ions to form soap-curd, a slimy precipitate. Lather forms only after all these ions are used up:

2\mathrm{NaSt} + \mathrm{Ca(HCO}_3)_2 = \mathrm{CaSt}_2 \downarrow + 2\mathrm{NaHCO}_3

Synthetic detergents solve this. Their calcium salts are soluble in water, so hardness does not stop their cleansing action.

3.13.3 Removal of hardness

Temporary hardness is removed by boiling, which drives off carbon dioxide and precipitates carbonates. Clark’s process adds a calculated quantity of slaked lime:

\mathrm{Ca(HCO}_3)_2 + \mathrm{Ca(OH)}_2 \rightarrow 2\mathrm{CaCO}_3 \downarrow + 2\mathrm{H}_2\mathrm{O}

Washing soda removes both types of hardness by precipitating insoluble carbonates. For permanent hardness, treat with a calculated quantity of soda ash:

\mathrm{CaSO}_4 + \mathrm{Na}_2\mathrm{CO}_3 \rightarrow \mathrm{CaCO}_3 \downarrow + \mathrm{Na}_2\mathrm{SO}_4

For large volumes of water, the permutit process is used. Permutit is hydrated sodium aluminium orthosilicate. Hard water percolates through it, and sodium ions exchange for calcium and magnesium ions. The spent permutit is regenerated with concentrated brine. For scientific work, cation and anion exchange resins remove all ions and give completely deionised water. The official CISCE syllabus lists hardness of water in the Class 9 Chemistry course.

What to practise from Water Class 9 ICSE

Work through the three exercise blocks at the end of this chapter. They test every idea covered above.

• EXERCISE 3(A) — water properties, the universal solvent, and the two experiments for dissolved solids and gases.
• EXERCISE 3(B) — solutions, solubility, solubility curves, crystallization, and the numerical questions on concentration and solubility.
• EXERCISE 3(C) — soft and hard water, causes, disadvantages, and removal of hardness.

For more material from the same set, browse the Class 9 books hub.

Water Class 9 ICSE frequently asked questions

How do I prepare water class 9 icse notes from this chapter?

Group the chapter into five themes: physical properties, universal solvent, solutions and solubility, crystals, and hardness. For each, note the definition, the formula or equation, and one example. Use the three exercise blocks at the end to test yourself.

What is the difference between temporary and permanent hardness of water?

Temporary hardness comes from bicarbonates of calcium and magnesium and is removed by boiling. Permanent hardness comes from sulphates and chlorides and is not removed by boiling.

Why does a saturated solution form crystals when it cools?

Cooling lowers the solubility of the solute. The excess solute can no longer stay dissolved and separates out as crystals.

What is the difference between efflorescence, deliquescence and hygroscopy?

Efflorescence is losing water of crystallization in dry air. Deliquescence is absorbing moisture and dissolving in it. Hygroscopy is absorbing moisture without dissolving.

How do I calculate the mass percent and volume percent of a solution?

Mass percent is mass of solute divided by mass of solution, times 100. Volume percent is volume of solute divided by volume of solution, times 100. Remember the denominator is the whole solution.

Reference: ICSE Class 9 Chemistry textbook material, chapter 3: Water — CISCE-aligned.

More from this book

• Chapter 1: The Language of Chemistry
• Chapter 2: Chemical Changes and Reactions
• Chapter 4: Atomic Structure and Chemical Bonding
• Chapter 5: The Periodic Table
• Chapter 6: Study of the First Element Hydrogen
• Chapter 7: Study of Gas Laws
• Chapter 8: Atmospheric Pollution