ICSE Class 9 Chemistry • Chapter 3 • Comprehensive Chapter Notes
Formula: $\text{H}_2\text{O}$
Chemical name: Dihydrogen oxide
Molecular mass: $2 \times 1 + 16 \times 1 = 18 \text{ amu}$
PDF Image: Please insert Fig. 3.1 from the PDF here.
Water is the most important natural resource; more than 70% of Earth's area is occupied with water. However, hardly 2.5% of this water makes up the world's supply of fresh water, including the frozen water in polar ice caps and glaciers.
Water is a major constituent of all living things (plants, animals and human beings) and of the atmospheric environment in which we live. Nearly 70% of our body weight is water.
Water occurs in both free and combined states.
Water exists in all the three physical states: as solid (ice), as liquid (water) and as gas (water vapour).
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1. Nature: Pure water is a clear, transparent liquid. It is colourless, odourless and tasteless.
The taste in water is due to gases and solids dissolved in it, i.e. impurities present in it.
2. Boiling point: Under normal pressure, pure water boils at $100^\circ\text{C}$. The boiling point of water is affected by pressure because boiling point of a liquid is considered as the temperature when vapour pressure equals the atmospheric pressure. Thus, the greater the pressure, the higher the boiling point, and vice versa. This is the principle by which a pressure cooker works.
In the hills, water boils at a temperature lower than $100^\circ\text{C}$ with atmospheric pressure being low ($70^\circ\text{C}$ on the top of Mount Everest) and so food is not cooked properly there. Therefore, pressure cookers are useful in the hills.
Note: The boiling point of water also increases due to the presence of dissolved impurities in it.
3. Freezing point of water or melting point of ice: Pure water freezes at $0^\circ\text{C}$ under normal pressure (one atmosphere). The freezing point of water decreases with an increase in pressure.
Note: The freezing point of water also decreases due to the presence of dissolved impurities in it.
4. Density: At $4^\circ\text{C}$, water has its maximum density, $1\text{ g/cm}^3$ or $1000\text{ kg/m}^3$, and minimum volume. Water expands on freezing, i.e. 92 volumes of water become 100 volumes of ice. Therefore, with the relative density of ice being 0.92, it floats.
5. Anomalous expansion of water: Water has an unusual physical property. When cooled, it first contracts in volume, as do other liquids, but below $4^\circ\text{C}$, it starts expanding and continues to do so till the temperature reaches $0^\circ\text{C}$, the point at which it freezes into ice.
PDF Image: Please insert Fig. 3.3 from the PDF here.
The property of anomalous expansion of water enables marine life to exist in the colder regions of the world, because even when the water freezes on the surface, it remains a liquid below the ice layer.
PDF Image: Please insert Fig. 3.4 from the PDF here.
6. Latent heat of fusion of ice: The latent heat of fusion is the amount of heat required to change a substance from the solid state to its liquid state at its melting point without any change in temperature.
The amount of heat energy required by ice to change into water, is called latent heat of fusion of ice. Its specific value is $336\text{ J/g}$ or $80\text{ cal/g}$. The same amount of heat is released when $1\text{ g}$ of water solidifies to form $1\text{ g}$ of ice at $0^\circ\text{C}$.
It is on account of high specific latent heat of solidification that lakes and rivers do not freeze suddenly.
7. Latent heat of vaporization of water: When water is boiled, it changes to gaseous state.
The energy required to change water into its vapour at its boiling point without any change in temperature is called latent heat of vaporization of water, and its specific value is $2268\text{ J/g}$ or $540\text{ cal/g}$. The same amount of heat is released when $1\text{ g}$ of steam condenses to form $1\text{ g}$ of water at $100^\circ\text{C}$.
It is on account of high specific latent heat of vaporization that steam causes far more serious burns than water at $100^\circ\text{C}$.
8. Specific heat capacity: It has been found that $1\text{ g}$ of water, when heated through $1^\circ\text{C}$, always absorbs $4.2\text{ J}$ (or 1 calorie) of heat energy. The fixed amount of heat absorbed by $1\text{ g}$ of water, when heated through $1^\circ\text{C}$, is called its specific heat capacity.
Water because of its high specific heat is commonly used as a coolant in motor car radiators, desert coolers, etc.
Due to its high specific heat capacity, the presence of a large amount of water is able to modify the climate of the nearby land areas, making them warmer in winter and cooler in summer. Land and sea breeze are also set up because of this great moderating property of water.
Water is a remarkable solvent. It dissolves many substances, forming aqueous solutions (water solutions). Water has a high dielectric constant. As a result, it reduces the electrostatic force of attraction between positive and negative ions and dissolves even inorganic compounds, which are usually electrovalent. Not only solids, but also gases and other liquids dissolve in water. Hence, it is difficult to find absolutely pure water in nature. For the same reason, water is also called universal solvent.
Substances that are apparently insoluble in water actually dissolve in it in traces. Even when we put water in a glass vessel, an extremely small amount of glass dissolves in it. It is for this reason that, when distilled water is kept in a sealed bottle for a long time, it leaves etchings on the inside surface of glass.
To show that ordinary tap water contains dissolved solids.
Procedure: Put some tap water on a clean watch glass, and place it over a beaker containing water. Boil the water in the beaker. When all the water has evaporated from the watch glass, remove it from the burner and let it cool.
Observation: On looking at the watch glass against light, a number of concentric rings of solid matter are seen. These are the dissolved solids left behind after evaporation of water.
The dissolved solids in water are salts, minerals and impurities.
Note: Tap water, river water and well water contain dissolved solids, but rainwater and distilled water do not contain dissolved solids, and so concentric rings are not formed in their case.
Air is present in dissolved state in all the natural sources of water. Of the two main components of air (nitrogen and oxygen), oxygen is more soluble in water than nitrogen.
Note: Since oxygen is more soluble in water compared to nitrogen, the air dissolved in water differs from ordinary air. It is expressed as $0.02$ volume of hydrogen per unit volume of water, at $0^\circ\text{C}$ or $506$ volumes of hydrogen chloride per unit volume of water, at $0^\circ\text{C}$.
The composition of air dissolved in water is $33\%$ oxygen (compared to $21\%$ in ordinary air), $66\%$ nitrogen and $1\%$ carbon dioxide.
Note: The gases dissolved in water can be easily expelled by boiling the water.
Note: Distilled water and boiled water have no taste. The pleasant taste of drinking water is due to the presence of dissolved matter, like air, carbon dioxide and minerals.
Air dissolved in water is biologically very important.
(1) Marine life: Like fish, use the oxygen of the air dissolved in water for respiration, and thus aquatic life is sustained. $1\text{ dm}^3$ of water contains nearly $40\text{ cm}^3$ of dissolved oxygen.
(2) Carbon dioxide:
A solution is a homogeneous mixture of two or more components whose composition may be gradually changed by changing the relative amounts of the components.
$\text{Solution} = \text{Solute} + \text{Solvent}$
The medium of dissolution that allows one or more component to dissolve in it to form a solution is known as SOLVENT. The substance that dissolves in the solvent to form a solution is known as SOLUTE.
A solution which is made up of two components is a binary solution. Solutions which are made up of three or four components are called ternary and quaternary solutions respectively. Here we shall deal with binary solutions only, composed of a solute and a solvent.
For example: When salt crystals dissolve in water to form saline water, the salt is classified as the solute, the water is classified as the solvent and the saline water is an example of a solution.
If the components exist in different physical states, then the solvent is that component which exists in the same phase (state) as the final solution.
For example, if sugar is dissolved in water, the final solution is a liquid solution. So water (liquid) is the solvent and sugar the solute. If the two substances are in the same phase (state), the component that is in excess is termed the solvent.
A solution may be formed by mixing:
A homogeneous solution of a solid into another solid is called a solid solution. Common metal alloys are solid solutions. The gold used in making jewellery is a solid solution containing gold and small quantities of copper or silver.
Brass is a solid solution containing $70\%$ copper and $30\%$ zinc. Here, copper acts as the solvent and zinc as the solute.
A solution in which the size of the solute particles is about $10^{-10}\text{ m}$ is called a true solution. In a true solution, the solute particles and the solvent molecules cannot be distinguished even under a microscope. In common use, the term solution actually means true solution.
1. Water is the most common solvent (a solution produced by dissolving a substance in water is known as an aqueous solution) but it is not the only solvent. Alcohol, petrol, ether, benzene, carbon disulphide, liquid ammonia, etc. are some non-aqueous liquid solvents in common use. The solutions made in these liquids are known as non-aqueous solutions.
2. When the solvent is alcohol, the solution is called a tincture. Tincture of iodine is a solution of iodine in alcohol.
Dilute solution: It is the solution in which the amount of solute is rather small compared to the mass of the solvent.
Concentrated solution: It is the solution in which the amount of solute is relatively large for a given mass of the solvent.
A solution that cannot dissolve any more of the solute at a given temperature is called a saturated solution.
A solution is a homogeneous mixture of a solute in a solvent; its composition may vary till a saturated solution is obtained at a particular temperature.
Experiment: Take $100\text{ g}$ of distilled water in a beaker. Add $1\text{ g}$ of nitre and stir with a glass rod till the nitre dissolves. Add one gram more of nitre and stir it. It too will dissolve. Continue adding a gram of nitre at a time, and stir vigorously after each addition. A stage will come when no more of the nitre will dissolve at that temperature, no matter how long it is left there or how vigorously it is stirred. The solution is then said to be saturated with nitre at that temperature. This is the saturated solution of nitre in water and it cannot dissolve any more of the solute without changing temperature.
The solubility of a solid usually decreases with fall in temperature. Thus, if the temperature of a saturated solution is lowered, or if a solution is left open at a higher temperature for long, (the solvent slowly vaporizes) and a part of the dissolved solute separates out in the form of crystals. The remaining solution is still a saturated solution but at a lower temperature.
A saturated solution can also be made unsaturated by adding more solvent to it.
A solution in which more of the solute can be dissolved at a given temperature is known as an unsaturated solution.
A solution that holds more solute than it can theoretically do at room temperature is a supersaturated solution. Heat a saturated solution and dissolve some more solute to it. Then cool it back to room temperature; you will get a supersaturated solution since it holds more solute than it can theoretically hold at that temperature.
A supersaturated solution at a particular temperature is one that is more concentrated (contains more solute) than its saturated solution at that temperature.
Preparation of supersaturated solutions: A saturated solution of a solute, say potassium nitrate, is prepared in slightly warm water. On cooling this solution, the excess nitre does not separate out. This solution contains more solute in it than it can hold at room temperature, and therefore it is a supersaturated solution.
It is possible to bring down the temperature of a saturated solution a little without the excess solute crystallizing out, if:
(i) no suspended impurities, like dust particles, are present in it.
(ii) the solution is not disturbed physically, i.e. it is left standing still in its container.
Concentration of a solution is the amount of solute dissolved in a given quantity of that solution. It can be expressed in many ways. For example:
(i) Mass percent
It is defined as the mass of solid solute in grams present in $100\text{ grams}$ of the solution. It is mostly used when solute is solid and solvent is liquid.
$\text{Mass percent} = \frac{\text{Mass of solute}}{\text{Mass of solution (solute + solvent)}} \times 100$
(ii) Volume percent
It is defined as the volume of solute in millilitres present in $100\text{ mL}$ of a solution. It is mostly used when solute and solvent are liquid.
$\text{Volume percent} = \frac{\text{Volume of solute}}{\text{Volume of solute + Volume of solvent}} \times 100$
Different substances dissolve to different extents in the same volume of a given solvent. In other words, they have different solubilities. For example, the masses of sodium chloride and potassium nitrate that dissolve in $100\text{ g}$ of water at $40^\circ\text{C}$ are $36.5\text{ g}$ and $65\text{ g}$ respectively. Obviously, the solubility of potassium nitrate is greater than that of sodium chloride.
The solubility of a solute in a particular solvent at a given temperature is equivalent to the maximum number of grams of the solute necessary to saturate $100\text{ g}$ of that solvent at that temperature.
A substance that has negligible solubility is called insoluble, e.g., silver chloride, which has $0.000015\text{ g}$ as its solubility. On the other hand, a substance that has more than negligible but less than high solubility is called sparingly soluble, e.g. calcium hydroxide, which has $0.17\text{ g}$ as its solubility. A substance having high solubility such as sodium chloride is simply called soluble.
At the noted temperature:
$\text{Solubility} = \frac{\text{Mass of solute}}{\text{Mass of solvent}} \times 100 = \frac{\text{M}_2 - \text{M}}{(\text{M}_1 - \text{M}) - (\text{M}_2 - \text{M})} \times 100$
The rate of dissolution or rate of solubility of a solid in a liquid depends on the following factors:
(a) Size of solute particles: The smaller the size of the solute particles, the greater is its total surface area exposed to the solvent. Therefore, the greater is the solubility of that solute.
(b) Stirring: This brings more of the solvent in contact with the solute and thus increases the rate of formation of solution.
(c) Temperature: The solubility of a gas in a liquid always decreases with rise in temperature. But the solubility of most solids in water usually increases with rise in temperature.
A solubility curve is a line graph that plots changes in solubility of a solute in a solvent against changing temperature.
PDF Image: Please insert Fig. 3.7 from the PDF here.
(i) Solubility curves of substances like calcium sulphate ($\text{CaSO}_4\cdot2\text{H}_2\text{O}$) show that there is a decrease (after attaining a certain temperature) in their solubilities with further rise in temperature.
(ii) Substances like sodium nitrate, potassium nitrate, potassium bromide, etc., show considerable increase in their solubilities with rise in temperature.
(iii) Solubility of sodium chloride increases only a little with increase in temperature.
In an endothermic process the solubility of solute increases with the increase of temperature. For example, solubility of $\text{KNO}_3$ increases with rise in temperature. In an exothermic process, the solubility increases on lowering the temperature. For example, solubility by calcium hydroxide in water decreases on increasing the temperature.
Anomalous solubility: There are some salts whose solubilities first increase, and then decrease, with rise in temperature. For example: $\text{Na}_2\text{SO}_4\cdot10\text{H}_2\text{O}$ (Glauber's salt).
Solubility curve of $\text{Na}_2\text{SO}_4\cdot10\text{H}_2\text{O}$ rises till it reaches $32.8^\circ\text{C}$, and then it falls slightly. This is because $\text{Na}_2\text{SO}_4\cdot10\text{H}_2\text{O}$ is hydrous below $32.8^\circ\text{C}$ and anhydrous above it.
(a) Shape of the curve indicates how the solubility of the given substance in a solvent varies with change in temperature.
(b) Solubilities of different solutes in a solvent at different temperatures can be compared by comparing at their respective solubility curves plotted on the same chart.
(c) The effect of cooling of hot solutions of different substances can be found from the curves. When a saturated solution is prepared at $40^\circ\text{C}$ and then it is cooled to $15^\circ\text{C}$, some salts get precipitated out.
Pressure: An increase in pressure on the surface of water increases the solubility of a gas in water.
At any given temperature, the mass of gas dissolved by a fixed volume of liquid is directly proportional to the pressure on the surface of the liquid. This law is known as Henry's Law.
For example, the solubility of carbon dioxide in water under normal atmospheric pressure is rather low, but when the water surface is subjected to higher pressure a lot more of $\text{CO}_2$ gas dissolves in it, as is seen in the case of soda water. On opening the soda water bottle, the dissolved gas rapidly bubbles out since pressure on the surface of water suddenly decreases.
Temperature: An increase in temperature of water causes a decrease in solubility of a gas in it. Thus, on boiling, water loses its taste. Since some taste of water is due to the gases in it, on boiling, these gases escape from water, leaving it tasteless.
A crystal is a homogeneous solid of definite geometrical shape. It has symmetrically arranged smooth plane surfaces which meet forming the sharp edges.
Crystallization is a process by which crystals of a substance are obtained by cooling a hot saturated solution.
All crystals of a pure compound are of similar shape, but those of different compounds may be of different shapes. Some shapes in which crystals exist are: Cubic (e.g., NaCl), Rhombohedric (e.g., $\text{CuSO}_4$), Octahedral (e.g., $\text{FeSO}_4$), Prismatic, with sharp edges (e.g., $\text{KNO}_3$).
In a laboratory, crystals may be obtained by the following methods:
Seeding: This process of inducing crystallization by adding a crystal of a pure substance into its saturated solution is called seeding. The suspended crystal acts as a seed crystal.
Caution: Avoid addition of dust particles, or crystals will form around them as well.
Substances which contain water molecules along with salts like sodium carbonate decahydrate ($\text{Na}_2\text{CO}_3\cdot10\text{H}_2\text{O}$) and copper sulphate pentahydrate ($\text{CuSO}_4\cdot5\text{H}_2\text{O}$) are hydrated substances.
Substances which do not contain any water along with the salts are anhydrous substances. For example sodium chloride ($\text{NaCl}$), potassium nitrate ($\text{KNO}_3$) etc.
The crystals of some salts contain water of crystallisation. This water gives the crystals their shape. In some cases it also gives them their colour (copper sulphate crystals are blue in colour). Such salts are known as hydrated salts.
The fixed amount of water that is associated with hydrated crystals which is an integral part of the crystal, is called water of crystallisation.
This water of crystallization is in loose chemical combination and can be driven off as steam by heating the crystals above $100^\circ\text{C}$, the vapours produced are condensed and the liquid formed is tested. This liquid gives test of water. In this process, the crystals lose their water and become anhydrous, i.e. all the water combined in the crystalline material is removed.
Example: When blue crystals of hydrated copper sulphate are heated in a test tube, they turn into a white powder, which turns back into a blue solid when a few drops of water are added.
$\text{CuSO}_4\cdot5\text{H}_2\text{O} \text{ (Hydrated blue colour salt)} \rightleftharpoons \text{CuSO}_4 \text{ (Anhydrous white colour salt powder)} + 5\text{H}_2\text{O}$
| Common name | Chemical name | Formula |
|---|---|---|
| Washing soda crystals | Sodium carbonate decahydrate | $\text{Na}_2\text{CO}_3\cdot10\text{H}_2\text{O}$ |
| Epsom salt | Magnesium sulphate heptahydrate | $\text{MgSO}_4\cdot7\text{H}_2\text{O}$ |
| Potash alum | Hydrated potassium aluminium sulphate | $\text{K}_2\text{SO}_4\cdot\text{Al}_2(\text{SO}_4)_3\cdot24\text{H}_2\text{O}$ |
| Glauber's salt | Sodium sulphate decahydrate | $\text{Na}_2\text{SO}_4\cdot10\text{H}_2\text{O}$ |
| Blue vitriol | Copper (II) sulphate pentahydrate | $\text{CuSO}_4\cdot5\text{H}_2\text{O}$ |
| Plaster of Paris | Calcium sulphate hemihydrate | $\text{CaSO}_4\cdot\frac{1}{2}\text{H}_2\text{O}$ |
| White vitriol | Zinc sulphate heptahydrate | $\text{ZnSO}_4\cdot7\text{H}_2\text{O}$ |
| Gypsum | Hydrated calcium sulphate | $\text{CaSO}_4\cdot2\text{H}_2\text{O}$ |
A substance is anhydrous if it contains no water, for example, salts lacking their water of crystallisation. The term 'Anhydrous', is most often applied to crystalline substances when the water of crystallisation is removed.
Anhydrous can also refer to the gaseous form of some concentrated solutions such as ammonia to distinguish it from its aqueous solution form.
Examples of anhydrous substances:
Anhydrous salts can be obtained from hydrated salts by carefully removing the water of crystallisation from them. It can be done by direct heating, heating in dry and hot air, heating under vacuum, or by using dehydrating/desiccating agents such as warm concentrated sulphuric acid.
The crystalline shape of a substance is not necessarily the result of the presence of water of crystallization. In fact, there are a number of crystalline solids that crystallize from water without holding any water of crystallization. Examples of crystalline substance which do not contain water of crystallization: Common salt ($\text{NaCl}$), Nitre ($\text{KNO}_3$), Sugar ($\text{C}_{12}\text{H}_{22}\text{O}_{11}$), Potassium permanganate ($\text{KMnO}_4$), Ammonium chloride ($\text{NH}_4\text{Cl}$).
On heating, crystals of sodium chloride produce a crackling sound, called decrepitation. This is due to the breaking of bigger crystals of sodium chloride into smaller ones.
A phenomenon where a compound loses its water of crystallisation on exposure to dry air is called efflorescency. This results in loss of crystalline shape, and finally the crumbling of crystals into a powder. Substances showing efflorescence are called efflorescent substances.
Efflorescence occurs when vapour pressure in the hydrated crystals exceeds atmospheric vapour pressure. Thus, it is minimized during humid conditions.
Examples of efflorescent substances:
The higher the temperature of the air, the higher the efflorescence. This is because the air absorbs more water with rising temperature and decreasing moisture.
Certain water soluble substances, when exposed to the atmosphere at ordinary temperatures, absorb moisture from atmospheric air, become moist, lose their crystalline form, and ultimately dissolve in the absorbed water, forming a saturated solution. Such a substance is called a deliquescent substance and the phenomenon is called deliquescency.
Deliquescence occurs when vapour pressure inside the crystals is very low compared to vapour pressure in the atmospheric air. Thus, deliquescence is minimized during dry conditions.
Deliquescent substances: Example of caustic soda $\text{NaOH}$, caustic potash $\text{KOH}$, magnesium chloride $\text{MgCl}_2$, zinc chloride $\text{ZnCl}_2$, calcium chloride $\text{CaCl}_2$ and ferric chloride $\text{FeCl}_3$.
Table salt [sodium chloride] turns moist and ultimately forms a solution, on exposure to air [especially during the rainy season]. Though pure sodium chloride is not deliquescent, the commercial version of the salt contains impurities, like magnesium chloride and calcium chloride, which are deliquescent substances.
Certain substances absorb moisture (water vapour) from the atmosphere when they are exposed to it. Such substances are called hygroscopic substances. The phenomenon is known as hygroscopy. Hygroscopic substances have a tendency to take up moisture from the air.
Some examples of hygroscopic substances are:
Drying agents are substances that can readily absorb moisture from other substances without chemically reacting with them.
For example, anhydrous calcium chloride, anhydrous zinc chloride, phosphorus pentoxide, magnesium sulphate, dry sodium sulphate, etc., are used to absorb water vapour from the air.
These substances are also called desiccants or desiccating agents.
Almost all hygroscopic substances are desiccating agents, e.g. conc. sulphuric acid, phosphorus pentoxide, silica gel, quicklime.
Dehydrating agents are substances that can remove even the chemically combined water molecules from compounds. For example, conc. sulphuric acid can remove water molecules from blue vitriol ($\text{CuSO}_4\cdot5\text{H}_2\text{O}$). Therefore, it is a dehydrating agent.
$\text{CuSO}_4\cdot5\text{H}_2\text{O} \text{ (blue)} \xrightarrow{\text{Conc } \text{H}_2\text{SO}_4} \text{CuSO}_4 \text{ (white)} + 5\text{H}_2\text{O}$
PDF Image: Please insert Fig. 3.13 from the PDF here.
| Drying agent | Dehydrating agent |
|---|---|
| (i) They remove moisture from other substances. | (i) They remove chemically combined elements of water in the ratio of 2:1 (hydrogen:oxygen) from a compound. |
| (ii) They are used to dry gases like chlorine, sulphur dioxide, hydrogen chloride, etc. They are also used in desiccators to keep substances dry. | (ii) They prepare substances like carbon monoxide, sugar charcoal, etc. (e.g. $\text{HCOOH} \xrightarrow{\text{warm conc. } \text{H}_2\text{SO}_4} 12\text{C} + \text{H}_2\text{O}$) |
| (iii) They represent physical change. Examples: Phosphorus pentoxide $\text{P}_2\text{O}_5$, fused calcium chloride $\text{CaCl}_2$, calcium oxide $\text{CaO}$, conc. sulphuric acid $\text{H}_2\text{SO}_4$. | (iii) They represent chemical change. For example: Conc. Sulphuric acid $\text{H}_2\text{SO}_4$. |
Water is said to be soft, if it readily forms lather with soap. Pure water or water containing sodium salts easily gives lather with soap. Such a water sample is called 'soft water' and is suitable for use. Distilled water and rain water are example of soft water.
Water is said to be hard when it does not readily form lather with soap.
Water of some springs, wells and rivers contain dissolved mineral matter. If this happens to be the hydrogen carbonates, sulphates or chlorides of calcium and magnesium, the water gets hard.
Hard water contains bicarbonates, chlorides or sulphates of calcium or magnesium.
When water flows over beds of limestone and dolomite ($\text{CaCO}_3, \text{MgCO}_3$) carbon dioxide slowly converts some of these into soluble calcium and magnesium bicarbonates which get mixed with the water and make it hard. The formation of stalactites and stalagmites in caves is due to the same reason.
Stalagmites and stalactites: In some lime-stone caves, one sometimes sees conical pillar like objects, some hanging from the roof of the caves and some rising from their floors. These are formed by water dropping from the cracks in the rocks, containing dissolved calcium hydrogen carbonate. Release of pressure results in the conversion of some hydrogen carbonate to calcium carbonate.
$\text{Ca(HCO}_3\text{)}_2 \rightarrow \text{CaCO}_3 + \text{CO}_2 + \text{H}_2\text{O}$
This calcium carbonate, little by little gets deposited both on the roof and floor of the cave. The conical pillar which grows downwards from the roof is called Stalactite and the one which grows upward from the floor of the cave is called stalagmite. They meet after a time. In a year some grow less than even a centimeter but some as tall as 100 cm.
PDF Image: Please insert Fig. 3.15 from the PDF here.
Hardness of water is of two types: temporary hardness and permanent hardness. Water that contains only hydrogen carbonates of calcium and magnesium is called temporary hard water.
It is called temporary hard water because its hardness can be removed just by boiling.
$\text{Ca(HCO}_3\text{)}_2 \xrightarrow{\text{Boil}} \text{CaCO}_3\downarrow + \text{H}_2\text{O} + \text{CO}_2\uparrow$
$\text{Mg(HCO}_3\text{)}_2 \xrightarrow{\text{Boil}} \text{MgCO}_3\downarrow + \text{H}_2\text{O} + \text{CO}_2\uparrow$
The calcium ions and magnesium ions are thus removed from the water by the formation of insoluble calcium carbonate and hence the water becomes soft.
Water containing sulphates and chloride of magnesium and calcium is called permanent hard water. This hardness can not be removed by boiling.
Note: The removing of temporary hardness of water by boiling is, however, not a practical method; it is costly and slow.
Hard water containing some dissolved salts has the following advantages.
Hard water is not suitable for producing steam
Steam is usually made in boilers which are made up of a number of narrow copper tubes surrounded by fire. As the cold water enters these tubes, it is immediately changed into steam, while the dissolved solids incapable of changing into vapour get deposited on the inner walls of the tubes. This goes on and thus making the bore of the tubes narrower and narrower. The result is that less water flows through the tubes at one time and less steam is produced. When the bore of the tube becomes very narrow, the pressure of steam increases so much that at times the boiler itself bursts. Not only that, the boiler scale formed inside the tube being a bad conductor of heat allows only part of the outside heat to go in and convert the water into steam. This results in further loss of heat.
Hard water is also unfit for washing purposes
For washing purposes soap is invariably used. Soap is chemically a sodium salt of stearic acid (an organic acid, formula $\text{C}_{17}\text{H}_{35}\text{COOH}$) and has the formula $\text{C}_{17}\text{H}_{35}\text{COONa}$. If the water is hard, the calcium and magnesium ions of the water combine with the negative ions of the soap to form a slimy precipitate of insoluble calcium and magnesium usually given the name of soap-curd or scum.
The formation of soap-curd will go on as long as there are calcium and magnesium ions present. Till then no soap lather will be formed and cleaning of cloth or body will not be possible. Moreover these precipitates are difficult to wash out from the fabrics and sometimes form rusty spots if iron salts are also present in water.
Removal of temporary hardness:
(1) By boiling: Carbon dioxide is driven off and the soluble hydrogen carbonates are converted into insoluble carbonates and could be removed by filtration or decantation.
(2) By addition of lime (Clark's Process): A more practical method is to add a calculated quantity of slaked lime. The following reaction takes place.
$\text{Ca(HCO}_3\text{)}_2 + \text{Ca(OH)}_2 \rightarrow 2\text{CaCO}_3\downarrow + 2\text{H}_2\text{O}$
(3) By the addition of washing soda (Removal of permanent hardness):
When washing soda or soda ash is added to hard water, the corresponding insoluble carbonates settle down and can be removed by filtration. It removes both temporary and permanent hardness.
$\text{CaSO}_4 + \text{Na}_2\text{CO}_3 \rightarrow \text{CaCO}_3\downarrow + \text{Na}_2\text{SO}_4$
Softening hard water, using ion exchange resins (Permutit process):
Permutit is an artificial zeolite. Chemically, it is hydrated sodium aluminium orthosilicate, having the formula $\text{Na}_2\text{Al}_2\text{Si}_2\text{O}_8\cdot\text{xH}_2\text{O}$. For the sake of convenience let us give it the formula $\text{Na}_2\text{P}$.
A tall cylinder is loosely filled with lumps of permutit. When hard water containing calcium and magnesium ions percolates through these lumps, exchange of ions takes place. The sodium permutit is slowly changed into calcium and magnesium permutit, and with the removal of calcium and magnesium ions, the water becomes soft.
When no longer active, the permutit is regenerated by running a concentrated solution of brine over it and removing the calcium chloride formed by repeated washing.
$\text{CaP} + 2\text{NaCl} \rightarrow \text{Na}_2\text{P} + \text{CaCl}_2$
