Chapter 3: Water
In the late summer of 1854, cholera broke out in the Soho district of London. Within days, hundreds of people living in a few narrow streets were dead. Most doctors of the time blamed bad air rising from the drains.
A physician named John Snow doubted that explanation. He walked the streets, asked at house after house who had fallen ill, and marked every death on a map of the neighbourhood. The marks did not spread evenly outward, as bad air would have. They crowded around one thing: a public water pump on Broad Street, from which almost every affected household drew its drinking water.
Snow paid attention to the exceptions as well, and they mattered just as much. The men working at a nearby brewery drank the beer they made and water from the brewery’s own well, and hardly any of them fell ill.
He carried his map to the local officials and asked them to take the handle off the pump, so that nobody could draw from it. They agreed. The outbreak, already fading by then, came to an end.
Snow never saw the organism that caused the disease. What he had was water that looked and tasted perfectly clean, and evidence that it was killing people. Water hides what it carries — and that is where this chapter begins.
Diagram
Fig. 3.0-1 — Mapping The Broad Street Outbreak Simplified street plan of a crowded nineteenth-century London neighbourhood drawn in plain line style; label a grid of narrow streets, one street marked Broad Street, a public hand pump drawn on it and labelled ‘water pump’, short black bars stacked outside houses to show deaths and clustering thickly near the pump and thinning with distance, and one building set apart labelled ‘brewery with its own well — few cases’.
3.1 Water In Nature
Turn a tap, and water arrives so easily that it is hard to think of it as a chemical substance at all. Yet every drop is a compound with a fixed formula, made of two elements joined in a fixed ratio. This chapter is about that substance — what it dissolves, what it holds, and what it does when it is not pure.
Water On The Earth
About seventy per cent of the Earth’s surface is covered with water. Almost all of it, however, is sea water, which is salty and cannot be drunk. Hardly 2.5 per cent of the world’s water is fresh water, and much of even that is locked up as the frozen water of the polar ice caps and glaciers.
So the water that is actually available to us — in rivers, lakes, wells and rain — is a very small share of an enormous total. That is one reason a chapter on water is also a chapter on using it carefully.
The Formula And What It Names
A molecule of water is two hydrogen atoms joined to one oxygen atom. Each hydrogen atom shares a pair of electrons with the oxygen atom, and the two shared pairs hold the molecule together. Only after picturing those two shared pairs do we write the formula:
\[\\mathrm{H\_2O}\]
Its chemical name is dihydrogen oxide, which simply reads the formula out in words — two hydrogen, one oxide. The sharing of electron pairs is shown in Fig. 3.1-1.
Diagram
Fig. 3.1-1 — Covalent Bonding In A Water Molecule Two side-by-side drawings of one water molecule. Left: the bent structural formula, with O at the apex, an H at the end of each bond and the label ‘water’ beneath. Right: the electron-dot picture — a large circle for oxygen overlapping two smaller circles for hydrogen, oxygen’s outer electrons shown as dots and each hydrogen’s electron shown as a cross, with the two shared pairs sitting in the overlaps; label the shared pairs.
The molecular mass of a substance is found by adding up the atomic masses of all the atoms shown in its formula. For water there are two hydrogen atoms of atomic mass 1 and one oxygen atom of atomic mass 16:
\[1 \\times 2 + 16 \\times 1 = 18\\ \\text{amu}\]
The Three States And The Water Cycle
Water is the one common substance we meet in all three physical states within an ordinary range of temperature. It is solid as ice, liquid as water, and gas as water vapour. Nothing else in the kitchen does this.
Because of that, water is constantly moving between the states. Heat from the Sun evaporates water from oceans, lakes and rivers; plants add more by transpiration. The vapour rises, cools, and condenses into clouds, and returns to the ground as rain or snow, which flows back to the sea. This endless circulation is the water cycle, drawn in Fig. 3.1-2.
Diagram
Fig. 3.1-2 — The Water Cycle Landscape scene with the Sun at the top right. Label, in order around the loop: evaporation from the ocean surface, transpiration from trees, rising vapour, cloud formation, cloud, rain falling on hills, snow on mountain tops, a waterfall and a river returning to the ocean. Draw arrows showing the continuous circulation and mark the ocean, mountains and forest.
Water occurs in two ways in nature. In the free state it is present as water itself — in seas, rivers, rain and ice. In the combined state it is locked inside other substances, such as the water held within the crystals of certain salts, or the water present in the tissues of plants and animals.
Water Inside Living Things
Water is a major part of every living organism — plants, animals and human beings alike. Nearly 70 per cent of our own body weight is water, and every process inside us happens in a watery medium. Digestion, the transport of substances in blood, and the removal of waste all depend on it.
For the same reason, life continues only where water is available. A plant wilts within hours of losing it, and no organism can be kept alive without a regular supply.
Water drawn from a natural source is rarely fit to drink as it stands. Suspended impurities are removed by filtration, dissolved impurities by distillation, and bacteria are destroyed by adding a carefully calculated amount of chlorine. Each of these steps is taken up in detail later in this chapter.
3.2 Physical Properties Of Water
Before asking what water does to other substances, it is worth looking closely at water itself. Several of its physical properties are unusual, and each of those oddities turns out to matter — to cooks, to fish, and to the climate of a coastal city.
3.2.1 Nature, Boiling Point And Freezing Point
What Pure Water Is Like
Pure water is a clear, transparent liquid. It is colourless, odourless and tasteless. If a sample of water has a taste, that taste does not belong to the water; it comes from the gases and solids dissolved in it — that is, from impurities.
Boiling Point
Under normal pressure, which is 760 mm of mercury, pure water boils at 100 °C. A liquid boils when the pressure of its own vapour becomes equal to the pressure of the atmosphere pressing on its surface. That single sentence explains why boiling point is not a fixed number in every situation.
If the outside pressure is raised, the vapour has to push harder before it can escape, so the liquid must be heated further. Greater the pressure, higher the boiling point — and the reverse is equally true. A pressure cooker works on exactly this principle: it traps steam, the pressure inside rises, and the water inside boils above 100 °C, so food cooks faster.
In the hills the atmospheric pressure is low, so water boils below 100 °C — about 70 °C on the top of Mount Everest. Food left in an open pan there simply does not cook properly, however long it is kept on the fire. This is why pressure cookers are so useful in the hills.
The boiling point of water also increases when impurities are dissolved in it. Salted water for cooking boils a little above 100 °C, not below.
Freezing Point
Pure water freezes at 0 °C under normal pressure of one atmosphere. This is the same temperature at which ice melts, so 0 °C is both the freezing point of water and the melting point of ice.
Two things lower it. The freezing point of water decreases with an increase in pressure, and it also decreases when impurities are dissolved in the water. Sea water, carrying a great deal of dissolved salt, freezes below 0 °C.
3.2.2 Density And Anomalous Expansion
Almost every liquid contracts steadily as it is cooled and is densest just before it solidifies. Water does not. On cooling, it first contracts like any other liquid, but below 4 °C it begins to expand again, and goes on expanding until it reaches 0 °C and freezes. This unusual behaviour is called the anomalous expansion of water.
The consequence is that water has its maximum density at 4 °C, where it also occupies its minimum volume. That maximum density is \(1\\ \\text{g cm}^{-3}\), or \(1000\\ \\text{kg m}^{-3}\). The way density changes on either side of 4 °C is shown in Fig. 3.2-1.
Diagram
Fig. 3.2-1 — Water Has Its Maximum Density At 4 °C Line graph on a light dotted grid. X-axis ‘temperature / °C’ running from −10 to 20 in steps of 5; y-axis ‘density / g mL⁻¹’ running from 0.998 to 1.000. Draw a smooth curve rising from the left, peaking at 4 °C, and falling again to the right; mark the peak with a short vertical dashed line labelled 4°.
Water also expands when it freezes, which very few liquids do: 92 volumes of water become 100 volumes of ice. Since the same mass now occupies a larger volume, ice is less dense than water — its relative density is 0.92 — and so ice floats instead of sinking.
Why Fish Survive A Frozen Pond
Put the two facts together and something important follows. As a pond cools in winter, the coldest water rises to the top instead of sinking, and it is the surface that freezes first. Below the ice the water stays liquid, and the deepest layers settle at about 4 °C, where water is densest.
The layer of ice on top then acts as a lid, because ice is a bad conductor of heat and slows down further loss of heat from the water beneath. Fish and other aquatic life survive the winter in that unfrozen water, as shown in Fig. 3.2-2. Had water behaved like ordinary liquids, ponds would have frozen solid from the bottom up.
Diagram
Fig. 3.2-2 — Formation Of Ice At The Top Of A Pond Cross-section of a pond in a shallow bowl-shaped bed. Label the air above as ‘air below 0 °C’, a hatched layer of ice across the surface labelled ‘ice’, and the water beneath labelled ‘water’. Mark horizontal bands of temperature increasing downwards: 0 °C just under the ice, then 1 °C, 2 °C, 3 °C, and 4 °C at the bottom.
3.2.3 Latent Heat And Specific Heat Capacity
Latent Heat Of Fusion Of Ice
Heat a mixture of ice and water and, surprisingly, the thermometer does not move until all the ice has melted. The heat is going in, but the temperature is not rising, because that heat is being used to break the solid structure apart rather than to warm anything. Heat used in this way is called latent heat.
For ice this latent heat of fusion has the specific value 336 J/g, or 80 cal/g. Exactly the same amount of heat is given out when 1 g of water solidifies to form 1 g of ice at 0 °C. Because so much heat has to be removed before water will turn to ice, lakes and rivers do not freeze suddenly.
Latent Heat Of Vaporization Of Water
When water at its boiling point is heated further, it changes into the gaseous state without getting any hotter. The energy needed to change water into its vapour at its boiling point, without any change in temperature, is the latent heat of vaporization of water, and its specific value is 2268 J/g, or 540 cal/g.
This large value explains the cooling effect of sweating. Evaporating sweat takes its latent heat from the body, and the body cools as a result.
The same amount of heat is released when 1 g of steam condenses to form 1 g of water at 100 °C. Steam at 100 °C therefore carries far more energy than water at 100 °C, which is why burns caused by steam are much more serious. It is this sudden release of the latent heat of condensation, on a very large scale, that gives torrential rain its violence.
The two latent heats are compared in Table 3.1.
Specific Heat Capacity
Water is also difficult to heat. It has been found that 1 g of water, when heated through 1 °C, always absorbs 4.2 J, that is, 1 calorie of heat energy.
A substance with a high specific heat capacity soaks up a great deal of heat for a small rise in temperature, and gives it back slowly on cooling. That makes water an excellent coolant, and it is used as one in motor car radiators and in desert coolers.
The same property works on a far larger scale near the sea. A large body of water warms and cools much more slowly than the land beside it, so it keeps nearby land areas warmer in winter and cooler in summer. Land and sea breezes arise from this same difference, as the land heats and cools faster than the water next to it.
3.3 Water As A Universal Solvent
Sugar disappears into tea, salt disappears into dal, and a fizzy drink hisses because a gas has been forced into water. In each case water is doing the same job. It is acting as a solvent.
Water is a remarkable solvent: it dissolves more substances than any other liquid found in nature. Not only solids, but gases and other liquids too dissolve in it, and often to a large extent. A solution produced by dissolving a substance in water is called an aqueous solution.
For this reason water is called a universal solvent. It also follows that absolutely pure water is very difficult to find in nature — whatever water touches, it takes a little of it along.
3.3.1 Why Water Dissolves So Much
The answer lies in the shape of the molecule you met in Fig. 3.1-1. The two hydrogen atoms are not placed opposite each other; they sit on either side of the oxygen atom in a bent, polar arrangement. The oxygen end of the molecule carries a negative charge, and each hydrogen carries a positive charge.
Because the molecule has a positive end and a negative end, it can attach itself to the particles of many other substances — the negative oxygen end turning towards a positive particle, and the positive hydrogen ends turning towards a negative one. This is what allows water molecules to bind with molecules of very different kinds, as Fig. 3.3-1 shows.
Diagram
Fig. 3.3-1 — Why A Polar Molecule Is A Good Solvent Left panel: a single water molecule drawn bent, with a δ− sign marked at the oxygen and a δ+ sign at each hydrogen; label ‘polar molecule’. Right panel: a crystal corner of an ionic solid at the edge of the drawing with a positive ion and a negative ion pulling away from it, each surrounded by a ring of water molecules — the water molecules around the positive ion turning their oxygen ends inwards, and those around the negative ion turning their hydrogen ends inwards. Label ‘positive ion’, ‘negative ion’ and ‘water molecules surrounding the ion’.
Water has one further advantage. It has a high dielectric constant, which means that it greatly weakens the force of attraction between oppositely charged ions placed in it. Once that pull between the ions is reduced, the ions of an electrovalent compound can separate from one another, and the compound dissolves. This is why water dissolves even inorganic compounds, which are usually electrovalent.
The reach of water as a solvent is easy to underestimate. Substances that appear completely insoluble do dissolve in it, in traces. Even a glass vessel gives up an extremely small amount of glass to the water standing in it — which is why distilled water kept in a sealed glass bottle for a long time leaves faint etchings on the inside surface of the glass.
3.3.2 Solids Dissolved In Water
The water that reaches our taps has travelled over soil and rock, and has taken dissolved solids along with it. These dissolved solids cannot be seen, but they can be made to show themselves.
An Experiment With A Watch Glass
Aim: to show that ordinary tap water contains dissolved solids.
Procedure: put some tap water on a clean watch glass and place the watch glass over a beaker containing water, as in Fig. 3.3-2. Boil the water in the beaker. When all the water on the watch glass has evaporated, remove the watch glass from the burner and let it cool. The watch glass is heated over boiling water rather than over the flame directly, so that it warms gently and does not crack.
Observation: hold the cooled watch glass up against the light. A number of concentric rings of solid matter are seen on it.
Conclusion: the rings are the dissolved solids left behind after the water evaporated. Tap water therefore contains dissolved solids.
Diagram
Fig. 3.3-2 — To Show That Water Contains Dissolved Solids Left: a beaker of water standing on a tripod over a burner labelled ‘heat’, with a silica glass tube inside it and a watch glass resting on the mouth of the beaker; label the watch glass ‘tap water’. Right: a separate circular view of the same watch glass after evaporation, showing a set of concentric rings; caption label ‘appearance of the watch glass after evaporation’.
Tap water, river water and well water all contain dissolved solids. Rainwater and distilled water do not, and so they leave no concentric rings behind.
Why The Dissolved Salts Matter
The dissolved solids in water are salts, minerals and impurities, and they are not merely a nuisance.
- Salts and minerals are essential for the growth and development of plants.
- They add taste to water.
- They supply essential minerals needed by our bodies.
3.3.3 Air Dissolved In Water
Air is present in the dissolved state in all natural sources of water. Of the two main components of air, oxygen is more soluble in water than nitrogen is.
That difference has a striking result. Since oxygen dissolves more readily, the air dissolved in water is not the same mixture as the air we breathe. The composition of air dissolved in water is 33 per cent oxygen — against 21 per cent in ordinary air — 66 per cent nitrogen and 1 per cent carbon dioxide.
An Experiment With A Flask And A Graduated Tube
Aim: to show that tap water contains dissolved gases.
Procedure: fill a round-bottom flask completely with tap water and arrange the apparatus as in Fig. 3.3-3. A graduated tube filled completely with water is inverted over the end of the delivery tube in a trough of water. Heat the water in the flask.
Observation: gas bubbles are seen escaping from the water. They collect in the graduated tube by the downward displacement of water. The water vapour carried along with them condenses when it meets the cold water in the tube.
Conclusion: the gas that remains has come out of the air dissolved in the water, and it escapes on heating. The volume collected in the graduated tube is a measure of the dissolved gases present in the tap water.
Diagram
Fig. 3.3-3 — To Show That Water Contains Dissolved Gases Round-bottom flask completely filled with water, clamped on a stand over a burner labelled ‘heat’; label the flask ‘tap water’ and show a two-holed stopper with a delivery tube leading down to the right. The delivery tube ends under an inverted graduated tube standing in a trough of water; mark the graduated tube with the scale readings 20, 40, 60, 80, label it ‘graduated tube’, label the collected gas at its closed top ‘gas’, and label rising bubbles in the flask ‘air bubbles’.
The solubility of a gas is usually stated as the volume of that gas dissolved in a certain volume of the solvent. The solubility of hydrogen in water, for example, is 0.02 volume of hydrogen per unit volume of water at 0 °C, while hydrogen chloride is enormously more soluble — 506 volumes per unit volume of water at 0 °C.
The gases dissolved in water are easily driven out by boiling it. That is why distilled water and boiled water have no taste at all. The pleasant taste of drinking water comes from the dissolved matter in it — air, carbon dioxide and minerals.
Why Dissolved Air Matters
Air dissolved in water is biologically very important.
- Marine life, such as fish, uses the oxygen of the dissolved air for respiration, and aquatic life is sustained by it. One cubic decimetre — that is, one litre — of water contains nearly 40 cm³ of dissolved oxygen.
- Aquatic plants use the dissolved carbon dioxide for photosynthesis, to prepare their food:
\[6,\\mathrm{CO\_2} + 12,\\mathrm{H\_2O} \\xrightarrow{\\ \\text{chlorophyll},\\ \\text{sunlight}\\ } \\mathrm{C\_6H\_{12}O\_6} + 6,\\mathrm{O\_2} + 6,\\mathrm{H\_2O}\]
- Carbon dioxide dissolved in water also attacks limestone. Limestone, which is insoluble, reacts with carbon dioxide and water to form soluble calcium bicarbonate:
limestone + carbon dioxide + water → calcium bicarbonate
\[\\mathrm{CaCO\_3}\\text{(s)} + \\mathrm{CO\_2}\\text{(aq)} + \\mathrm{H\_2O}\\text{(l)} \\rightarrow \\mathrm{Ca(HCO\_3)\_2}\\text{(aq)}\]
Marine animals such as snails and oysters draw on this dissolved calcium bicarbonate, extracting calcium carbonate from it to build their shells. The same reaction returns in Section 3.15, where it turns out to be the reason some water refuses to lather with soap.
3.4 Mixtures
In an earlier chapter you studied pure substances, which are of two kinds — elements and compounds. A pure substance has a fixed composition, and its constituents are chemically joined. Water itself is a compound: its two elements are combined in a fixed ratio and cannot be separated by simply pouring or filtering.
A mixture is a very different thing. Mixtures are formed by simply mixing two or more elements or compounds together, without any chemical combination. Dissolve sodium chloride in water, or sugar in water — both of them pure compounds — and what you get is a mixture.
Two phrases in that definition carry the weight. ‘Physically mixed’ means no new substance has been made, so the components keep their own properties. ‘In any proportion’ means you may add a spoon of sugar or three; a compound would allow no such freedom.
Mixtures are of two kinds, and the difference lies in whether the mixing is uniform, as Fig. 3.4-1 shows.
Diagram
Fig. 3.4-1 — Uniform And Non-Uniform Mixing Two beakers side by side, each drawn with the particles shown as small circles. Left beaker: two kinds of particle, shown as shaded and unshaded circles, spread evenly everywhere; label ‘homogeneous — same throughout’. Right beaker: the same two kinds of particle but separated into an upper layer of one and a lower layer of the other, with a clear line between them; label ‘heterogeneous — visible boundary of separation’.
3.4.1 Homogeneous Mixtures
A mixture is said to be homogeneous if the constituents present in it are uniformly mixed, that is, it has a uniform composition throughout. Take a sample from the top or the bottom and it is the same.
Air is a homogeneous mixture of a number of gases — nitrogen, oxygen, carbon dioxide, water vapour and the inert gases. All the gases present are uniformly mixed throughout, though their proportion may vary from place to place. Sugar solution is another homogeneous mixture.
In a homogeneous mixture the constituent particles cannot be identified or seen without the help of a powerful microscope.
3.4.2 Heterogeneous Mixtures
A mixture is said to be heterogeneous if its composition is not uniform throughout. It has visible boundaries of separation between its constituents.
Oil and water form a heterogeneous mixture, since the two stay in different layers. A mixture of sodium chloride and sand is heterogeneous too, even though both are solids in the same phase, because the particles of sand and of common salt can easily be seen separately and there are clear boundaries between them. Sand mixed with water is heterogeneous as well, since the sand settles down after some time.
In a heterogeneous mixture the constituents can easily be seen with the naked eye. They are either present in different phases, or in the same phase but with clear boundaries of separation. The two kinds of mixture are set side by side in Table 3.2.
3.5 Solutions As Homogeneous Mixtures
Of all the homogeneous mixtures you meet, the commonest by far is the kind made by dissolving one substance in another. That kind has its own name.
The phrase ‘gradually changed’ is worth pausing on. Stir in a little more sugar and the tea is a little sweeter; there is no fixed recipe a solution must obey.
3.5.1 Solute, Solvent And Kinds Of Solution
Every solution has two parts:
\[\\text{solution} = \\text{solute} + \\text{solvent}\]
You already know the solvent as the medium of dissolution. The substance that dissolves in the solvent to form a solution is the solute.
When salt crystals dissolve in water to form saline water, the salt is the solute, the water is the solvent, and the saline water is the solution.
Deciding Which Is Which
Sometimes it is not obvious which component to call the solvent. Two simple rules settle it.
If the components exist in different physical states, the solvent is the component that exists in the same phase, or state, as the final solution. Dissolve sugar in water and the final solution is a liquid; so water, the liquid, is the solvent and sugar the solute.
If the two substances are in the same phase, the component present in excess is termed the solvent.
How Many Components
A solution made up of two components is a binary solution. Those made of three or four components are called ternary and quaternary solutions respectively. Here we shall deal with binary solutions only, composed of one solute and one solvent.
Solutions Of Every Combination
A solution may be formed by mixing:
- a liquid with another liquid — water and alcohol;
- a solid with a liquid — sugar solution, and iodine in alcohol, which is tincture of iodine;
- a liquid with a gas — soda water;
- a solid with another solid — alloys.
A homogeneous solution of a solid in another solid is called a solid solution. Common metal alloys are solid solutions. The gold used in making jewellery is a solid solution of gold with small quantities of copper or silver.
Brass is a solid solution containing 70 per cent copper and 30 per cent zinc. Here copper, being in excess, acts as the solvent, and zinc is the solute — the same rule you have just used, applied to two solids.
Water Is Not The Only Solvent
Water is the most common solvent, but not the only one. Alcohol, petrol, ether, benzene, carbon disulphide and liquid ammonia are some non-aqueous liquid solvents in common use, and the solutions made in them are called non-aqueous solutions. Some familiar solutes and the solvents that dissolve them are listed in Table 3.3.
Table 3.3 — Some common solutes and their solvents
| Solute | Solvent |
|---|---|
| Chlorophyll | Methylated spirit |
| Iodine | Alcohol |
| Paraffin wax | Turpentine oil |
| Sulphur | Carbon disulphide |
| Pigments in paint | Turpentine oil |
| Phosphorus | Carbon disulphide |
| Pigments in nail polish | Acetone |
| Oils in grease | Petrol |
| Bromine | Carbon tetrachloride |
When the solvent is alcohol, the solution is called a tincture. Tincture of iodine is a solution of iodine in alcohol, and is used as an antiseptic on wounds.
3.5.2 True Solutions
Sugar solution looks perfectly clear. Muddy water does not, although both are water with something in it. The difference is the size of the particles that have been added.
A solution in which the size of the solute particles is extremely small — less than \(10^{-9}\\ \\text{m}\), that is, less than 1 nanometre — 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.
Characteristics Of A True Solution
- It is a homogeneous mixture.
- The solute particles are very small, less than \(10^{-9}\\ \\text{m}\).
- It is clear and transparent.
- It does not scatter light.
- Its components cannot be separated by filtration.
- The solute particles in a solution do not settle down.
- Because the solute particles are so small, they do not scatter a beam of light passed through the solution.
Points 4 and 7 say the same thing twice over for a reason: this is the test that will separate a true solution from the two kinds of mixture taken up in the next two sections.
3.5.3 Dilute And Concentrated Solutions
A solution can be described as dilute or concentrated, depending on the proportion of solute dissolved in it.
A dilute solution is one in which the amount of solute dissolved is relatively small. A concentrated solution is one in which the amount of solute dissolved is relatively large.
These two words carry no quantitative information whatsoever. They only compare one solution with another. By adding more solute, a dilute solution becomes more concentrated; by adding more solvent, a concentrated solution becomes dilute.
3.6 Suspensions
Mixtures in which the constituent particles are present in more than one phase are heterogeneous, and they are of two kinds: suspensions and colloidal solutions. Which of the two a given mixture is — and indeed whether it is a solution at all — is decided by one thing only: the size of the particles that have been mixed in.
The three cases are drawn together in Fig. 3.6-1. In suspensions the particle size is greater than 1000 nm, that is, greater than \(10^{-6}\\ \\text{m}\). In colloids it lies between 1 nm and 1000 nm, that is, between \(10^{-9}\) and \(10^{-6}\\ \\text{m}\). In true solutions it is less than 1 nm, or \(10^{-9}\\ \\text{m}\).
Diagram
Fig. 3.6-1 — Particle Size Decides The Kind Of Mixture Three panels in a row. Each panel has a magnified box above and a beaker below, joined by a dotted line. Panel 1, headed ‘size < 1 nm’: box with a dense scatter of very tiny dots; beaker of evenly coloured, clear liquid; label ‘true solution’. Panel 2, headed ‘size between 1 nm and 1000 nm’: box with medium dots; beaker of evenly coloured but slightly cloudy liquid; label ‘colloidal solution’. Panel 3, headed ‘size > 1000 nm’: box with large dots; beaker with large particles visible and some settled at the bottom; label ‘suspension’.
Making A Suspension In A Test Tube
A suspension is easy to make from two clear liquids and a gas. Bubble carbon dioxide through lime water and the lime water turns milky white. Lime water is a solution of calcium hydroxide, and it combines with carbon dioxide to form a precipitate of calcium carbonate and water:
calcium hydroxide + carbon dioxide → calcium carbonate + water
\[\\mathrm{Ca(OH)\_2}\\text{(aq)} + \\mathrm{CO\_2}\\text{(g)} \\rightarrow \\mathrm{CaCO\_3}\\text{(s)} + \\mathrm{H\_2O}\\text{(l)}\]
The milky appearance is caused by the solid particles of calcium carbonate, which are more than 1000 nm across — the size of a suspension. Pour the milky liquid into a beaker and leave it standing, and the solid particles slowly settle at the bottom, exactly as the definition says they should.
Suspensions are formed only by those substances which do not dissolve in water. Muddy water, a chalk–water mixture, some paints and some medicines are all suspensions. This is why the label on such a medicine tells you to shake the bottle before use — the solid has settled since it was last poured.
Characteristics Of A Suspension
- It is a heterogeneous mixture.
- The solute particles are large, more than 1000 nm across.
- The solute particles can be seen with the naked eye.
- The solid particles can easily be separated by ordinary filter paper.
- The particles settle down after some time under the influence of gravity.
Points 4 and 5 follow directly from point 2. Particles that large cannot slip through the pores of a filter paper, and they are heavy enough for gravity to pull them down.
3.7 Colloidal Solutions Or Sols
Between the true solution, whose particles are too small to see, and the suspension, whose particles settle out, lies a third case. Milk looks as uniform as sugar solution, yet it is not one — and it never settles, as muddy water does.
The particles of a colloidal solution are small enough not to settle down under gravity, but large enough to scatter light. Between them, those two facts account for almost everything a colloid does.
A colloidal solution always has two phases. The component present in the smaller amount is the dispersed phase, and the one present in the larger amount is the dispersion medium.
3.7.1 Types Of Colloidal Solutions
Either phase may be a solid, a liquid or a gas, and depending on the pair chosen, colloidal solutions are of eight types. They are set out with their examples in Table 3.4.
Table 3.4 — Types of colloidal solutions
| Name of colloidal solution | Dispersed phase | Dispersion medium | Examples |
|---|---|---|---|
| Emulsion | Liquid | Liquid | Cod liver oil, tonics in liquid form, milk |
| Gel | Liquid | Solid | Butter, cheese, jelly, curd |
| Liquid aerosol | Liquid | Gas | Cloud, fog, mist |
| Foam | Gas | Liquid | Soda water, soap water |
| Solid foam | Gas | Solid | Rubber, bread, pumice stone |
| Solid sol | Solid | Solid | Ruby glass, alloys |
| Sol | Solid | Liquid | Paints, gold sol, starch in water |
| Solid aerosol | Solid | Gas | Dust storm, smoke, fumes, vehicular exhaust |
Emulsions
Colloidal solutions in which the dispersed phase and the dispersion medium are both liquids are known as emulsions. The two liquids are not miscible with each other: in an emulsion one constituent is generally oily or oil-soluble, while the other is water-soluble.
Left to themselves, oil and water form separate layers and do not mix. Certain substances known as emulsifiers, generally proteins, help in forming stable emulsions in which the two stay mixed. Moisturizing cream, milk, cold cream and vanishing cream are all emulsions of this kind.
3.7.2 Characteristics Of Colloidal Solutions
- A colloidal solution is a heterogeneous mixture, though it appears homogeneous. The size of the particles ranges from 1 nm to 1000 nm, which is intermediate between true solutions and suspensions, and they can be seen only under a microscope.
- It is a two-phase system, consisting of a dispersed phase and a dispersion medium.
- Colloidal particles pass through ordinary filter paper, as the particles of a true solution do, but they cannot pass through ultrafilters such as animal membranes, which are semipermeable. A special method of separation known as centrifugation can be used to separate colloidal particles.
- Colloidal particles carry an electric charge. Since all the particles carry the same kind of charge, they repel one another and remain dispersed instead of settling like the particles in a suspension.
- Particles in a colloidal solution move in a continuous zig-zag path, visible under a microscope. This motion is known as Brownian motion, and is drawn in Fig. 3.7-1.
- Colloids scatter a beam of light passing through them. This is known as the Tyndall effect.
Diagram
Fig. 3.7-1 — Brownian Motion Of A Colloidal Particle The path of a single particle drawn as a continuous zig-zag line of straight segments in random directions, with a small solid dot at each turning point, contained within a plain square field. No axes or labels inside the field.
Why The Beam Shows Up
The reason for the Tyndall effect is that the particles of a colloidal solution are big enough to scatter light. The scattered light travels off sideways to your eye, and so the path of the beam itself becomes visible.
You can see it in a room that a beam of light enters through a small hole, where dust and smoke particles in the air scatter the light. You can see it again where sunlight passes through the canopy of a dense forest, the mist there holding tiny droplets of water that act as colloidal particles dispersed in air. Both are shown in Fig. 3.7-2.
Diagram
Fig. 3.7-2 — The Tyndall Effect Two scenes side by side. Left: the inside of a darkened room with a small hole in a shuttered window, a straight shaft of light crossing the room and made visible along its whole length, with fine specks drawn within it; label ‘dust and smoke particles scatter the light’. Right: tall forest trees with slanting shafts of sunlight coming through the canopy into low mist; label ‘droplets of water in mist scatter the light’.
The three kinds of mixture can now be told apart property by property, as in Table 3.5.
3.7.3 Colloids In Everyday Life
Colloids are everywhere. The food we eat — milk, butter, ice cream — our medicines, such as vaccines and ointments, and household products like paints and detergents are all colloids. Among their key uses are the purification of water by coagulation, the stabilising of food products, cosmetics and lubricants, and industrial processes such as smoke precipitation.
Most of these applications depend on one idea. Colloidal particles stay dispersed because they all carry the same charge; take that charge away and they clump together and fall out.
Four Familiar Examples
Alum stops a cut from bleeding. Blood contains haemoglobin, which is present as charged colloidal particles. Alum, a hydrated double sulphate salt, dissolves in water to provide ions; when these come in contact with blood they neutralise the charge on the colloidal particles. The blood becomes thick and the bleeding stops.
Soap cleans dirty clothes. Clothes become dirty mainly through perspiration, and the dirt particles get trapped in oily matter on the cloth. Ordinary water cannot remove oil drops. Soap helps by forming an emulsion between the water and the oil drops carrying the dirt, and the emulsion washes away. How the soap molecule manages this is taken up in Section 3.16.
The sky appears blue. Sunlight is scattered by the colloidal particles in the atmosphere — air molecules, dust and water vapour — by the Tyndall effect. Light of shorter wavelength, which is blue, is scattered much more strongly, and this scattered blue light reaches our eyes from every direction, so the sky looks blue.
Deltas form where a river meets the sea. River water carries suspended colloidal particles of clay, sand and silt. Where the river meets a sea or lake, these particles collide with the salt water and are coagulated by it, settling out. The river slows down, deposits the sediment over a long period, and new land is formed.
3.8 Saturated, Unsaturated And Supersaturated Solutions
Stir sugar into a cup of water and it disappears. Keep stirring in more, and there comes a point at which the last spoonful simply lies at the bottom and refuses to go. The water has taken all it will take at that temperature.
The words ‘at a given temperature’ are not a decoration. The same water at a higher temperature would have taken more, so a solution can only be called saturated with the temperature attached.
3.8.1 Preparing A Saturated Solution
Aim: to prepare a saturated solution of nitre in water.
Procedure: take 100 g of distilled water in a beaker. Add 1 g of nitre and stir with a glass rod till the nitre dissolves. Add one more gram of nitre and stir; it too will dissolve. Continue adding a gram of nitre at a time, stirring vigorously after each addition.
Observation: a stage comes 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.
Conclusion: the solution is now saturated with nitre at that temperature. It cannot dissolve any more of the solute unless the temperature is changed.
Diagram
Fig. 3.8-1 — Reaching Saturation Three beakers in a row, each containing water with a glass rod standing in it and a small heap of solid being added from a spatula above. Beaker 1 label ‘first gram — dissolves completely’; beaker 2 label ‘further additions — still dissolving’; beaker 3 label ‘no more dissolves — undissolved solid stays at the bottom’, drawn with a layer of solid resting on the base. Mark all three ‘at the same temperature’.
This gives a fuller way of describing a solution than the one in Section 3.5: a solution is a homogeneous mixture of a solute in a solvent, and its composition may go on varying until a saturated solution is obtained at a particular temperature. Saturation is the limit at which the composition stops changing.
3.8.2 Effect Of Cooling And Heating A Saturated Solution
On Cooling
The solubility of a solid usually decreases with a fall in temperature. So if the temperature of a saturated solution is lowered, the solution can no longer hold all the solute it is carrying, and a part of the dissolved solute separates out in the form of crystals.
The same thing happens if a solution is left open at a higher temperature for a long time. The solvent slowly vaporises, there is less water to hold the solute, and again crystals separate out. In both cases the liquid left behind is still a saturated solution, but a saturated solution at the lower temperature.
On Heating
If a saturated solution is heated to a higher temperature, it becomes unsaturated, and more solute can be dissolved in it. A saturated solution can also be made unsaturated by adding more solvent to it, since the same solute is now spread through more water.
3.8.3 Unsaturated And Supersaturated Solutions
A solution in which more of the solute can be dissolved at a given temperature is known as an unsaturated solution. Every solution on the way to saturation is an unsaturated one.
A solution that holds more solute than it can theoretically hold at a given temperature is a supersaturated solution.
That sounds impossible, and it is only just possible. To make one, heat a saturated solution and dissolve some more solute in it. Then cool it back to room temperature slowly, without causing any disturbance. The resultant solution now holds more solute than a saturated solution would at that temperature.
Such a solution is metastable, which means it is balanced on a knife edge. Slight disturbances — shaking, stirring, scratching the wall of the container, or dropping in a single crystal of the solute — make the additional solute precipitate out at once, and what is left is an ordinary saturated solution again.
The cooling has to be done with some care. It is possible to bring the temperature of a saturated solution down a little without the excess solute crystallising out, provided that no suspended impurities such as dust particles are present in it, and that the solution is not disturbed physically but is left standing still in its container.
3.9 Concentration Of A Solution
Calling a solution dilute or concentrated compares it with another solution and tells you nothing you can write down. A chemist, a doctor and a factory all need a number instead.
The concentration of a solution is the amount of solute dissolved in a given quantity of that solution. It can be expressed in several ways, of which two are used here.
3.9.1 Mass Percent
Mass percent is defined as the mass of solid solute in grams present in 100 grams of the solution. It is mostly used when the solute is a solid and the solvent is a liquid.
\[\\text{mass percent} = \\frac{\\text{mass of solute}}{\\text{mass of solution (solute} + \\text{solvent)}} \\times 100\]
Notice what sits in the denominator. It is the mass of the whole solution, not the mass of the water alone, and the mass of the solution is found by adding the two together.
For example, if 10 grams of sodium chloride is added to 90 grams of water, the mass of the solution is 100 grams, and the concentration of the solution is
\[\\frac{10}{10 + 90} \\times 100 = \\frac{10}{100} \\times 100 = 10%\]
3.9.2 Volume Percent
Volume percent is defined as the volume of solute in millilitres present in 100 mL of a solution. It is mostly used when the solute and the solvent are both liquids, since liquids are far easier to measure by volume than by mass.
\[\\text{volume percent} = \\frac{\\text{volume of solute}}{\\text{volume of solute} + \\text{volume of solvent}} \\times 100\]
For example, if 30 mL of alcohol is mixed with 70 mL of water to get 100 mL of solution, the volume percent is
\[\\frac{30}{30 + 70} \\times 100 = \\frac{30}{100} \\times 100 = 30%\]
3.10 Solubility
Saying that a solution is saturated tells you that no more will dissolve. It does not tell you how much dissolved before that point was reached. That figure is the solubility, and it is a number you can look up, plot and calculate with.
\[\\text{solubility} = \\frac{\\text{mass of solute}}{\\text{mass of solvent}} \\times 100\]
Put in words, the solubility of a solute in a particular solvent at a given temperature is the maximum number of grams of that solute needed to saturate 100 g of that solvent at that temperature.
Different substances dissolve to very different extents in the same solvent. At 40 °C, 100 g of water dissolves 36.5 g of sodium chloride but 63 g of potassium nitrate, so the solubility of potassium nitrate is clearly the greater of the two.
Soluble, Sparingly Soluble And Insoluble
These three words are only a rough scale of the same quantity. A substance with a negligible solubility is called insoluble — silver chloride, for instance, has a solubility of just 0.000015 g. A substance with more than negligible but less than high solubility is called sparingly soluble; calcium hydroxide, with 0.17 g, is one. A substance with a high solubility, such as sodium chloride, is simply called soluble.
3.10.1 Factors Affecting Solubility
When a solid dissolves, the particles of the solute merely occupy the spaces between the particles of the solvent. No chemical change takes place, and it is this that decides which factors matter and which do not.
(a) Size of the solute particles. The smaller the size of the solute particles, the greater is the total surface area exposed to the solvent, and so the faster the solute dissolves. Powdered sugar disappears into water long before a lump of the same mass does.
(b) Stirring. Stirring brings fresh solvent into contact with the solute and carries away the solution already formed, and so it increases the rate at which the solution forms.
(c) Temperature. This factor changes the solubility itself, and not merely the speed. The solubility of a gas in a liquid always decreases with a rise in temperature. The solubility of most solids in water, however, usually increases with a rise in temperature. The solubility of potassium nitrate in water at 20 °C is 31.6 grams, whereas at 60 °C it is 108 grams.
Some solids behave in the opposite way. The solubility of substances like calcium sulphate decreases as the temperature rises, because their dissolving is an exothermic process.
3.10.2 Solubility Curves
The way the solubility of a substance varies with temperature is best shown as a graph. Solubility is plotted along the ordinate, the y-axis, and temperature along the abscissa, the x-axis.
The solubility curves of some common compounds are shown in Fig. 3.10-1.
Diagram
Fig. 3.10-1 — Variation In The Solubility Of Some Solids With Changing Temperature Line graph. X-axis ‘temperature in °C’ from 0 to 100 in steps of 20; y-axis ‘solubility’ from 0 to 100 in steps of 20. Draw seven curves, each labelled along its own length: potassium nitrate rising steeply from low on the left to off the top of the scale on the right; sodium nitrate rising steadily from about 70 to above 100; potassium bromide rising steadily and crossing the sodium nitrate curve; potassium chloride rising gently from about 28 to about 55; sodium chloride almost flat at about 36 across the whole range; potassium sulphate rising slowly from about 8 to about 24; hydrated calcium sulphate almost flat and very low, marked ‘(hydrated)’.
Three kinds of behaviour can be read off the graph.
- Substances such as sodium nitrate, potassium nitrate and potassium bromide show a considerable increase in their solubilities with a rise in temperature. Their curves climb steeply. Some values are given in Table 3.6.
- The solubility of sodium chloride increases only a little with an increase in temperature, so its curve is nearly flat.
- The curves of substances like hydrated calcium sulphate, \(\\mathrm{CaSO\_4!\\cdot!2H\_2O}\), show a decrease in solubility with further rise in temperature after a certain temperature has been reached. The solubility of anhydrous calcium sulphate, \(\\mathrm{CaSO\_4}\), decreases with a rise in temperature throughout. Cerium(IV) sulphate, \(\\mathrm{Ce(SO\_4)\_2}\), and slaked lime, \(\\mathrm{Ca(OH)\_2}\), behave in the same way.
Table 3.6 — Solubility of potassium nitrate and sodium nitrate at three temperatures
| Temperature | Solubility of \(\\mathrm{KNO\_3}\) in grams | Solubility of \(\\mathrm{NaNO\_3}\) in grams |
|---|---|---|
| 40 °C | 63 | 105 |
| 50 °C | 84 | 114 |
| 60 °C | 108 | 124 |
Why Some Curves Fall
The direction of a curve is decided by whether heat is taken in or given out as the solute dissolves. In an endothermic process, heat is absorbed during dissolution, and the solubility increases with an increase in temperature — as with potassium nitrate and sodium nitrate. In an exothermic process, heat is given out, and the solubility increases on lowering the temperature; the solubility of calcium sulphate and of calcium hydroxide in water therefore decreases as the temperature is raised.
Anomalous Solubility
There are some salts whose solubilities first increase and then decrease with a rise in temperature. Glauber’s salt, \(\\mathrm{Na\_2SO\_4!\\cdot!10H\_2O}\), is the standard example, and its curve is shown in Fig. 3.10-2.
The solubility curve of \(\\mathrm{Na\_2SO\_4!\\cdot!10H\_2O}\) rises till it reaches 32.8 °C, and then it falls slightly. The reason is that the salt itself changes at that temperature: \(\\mathrm{Na\_2SO\_4!\\cdot!10H\_2O}\) is hydrous below 32.8 °C and anhydrous above it, so beyond the peak the curve is really that of a different substance.
Diagram
Fig. 3.10-2 — Solubility Curve For Na₂SO₄·10H₂O Line graph with y-axis labelled ‘solubility’ and x-axis labelled ‘temperature (in °C)’. A single curve rises steeply from the lower left, labelled along its rising part ‘Na₂SO₄·10H₂O’, reaches a sharp peak, and then falls gently to the right, labelled along its falling part ‘Na₂SO₄’. Mark the peak with a vertical dashed line down to the x-axis, labelled 32.8 °C.
What A Solubility Curve Is For
A solubility curve is used to determine the solubility of a substance at a particular temperature, and to find the effect of cooling hot solutions of different substances. If a saturated solution is prepared at 40 °C and then cooled to 15 °C, the curve tells you how much salt the solution can no longer hold, and that much gets crystallised out.
3.10.3 Solubility Of Gases In Water
A gas dissolved in a liquid is held there by the pressure of the gas above the surface. Change that pressure, or change the temperature, and the amount that stays dissolved changes at once.
Effect Of Pressure
An increase in pressure on the surface of water increases the solubility of a gas in the water.
The condition ‘at any given temperature’ is part of the law, not an afterthought: the relation holds only while the temperature is kept constant.
Soda water is the everyday demonstration. The solubility of carbon dioxide in water under ordinary atmospheric pressure is rather low, but when the surface of the water is kept under a much higher pressure, a great deal more carbon dioxide dissolves in it. Open the bottle and the pressure on the surface suddenly drops, so the dissolved gas comes rapidly bubbling out.
Effect Of Temperature
An increase in the temperature of water causes a decrease in the solubility of a gas in it. This is why water loses its taste on boiling: some of that taste is due to the gases dissolved in it, and on boiling those gases escape, leaving the water tasteless.
A decrease in temperature works the other way and increases the solubility of gases in water. That is why a chilled bottle of a soft drink gives out less fizz than a warm one when it is opened.
The two kinds of solute behave quite differently, as Table 3.7 shows.
3.11 Crystals And Crystallisation
Let a saturated solution cool slowly and the solute does not come out as a shapeless powder. It comes out in solids with flat faces, straight edges and sharp corners, and every one of them has the same shape as the next.
Shapes Of Crystals
All the crystals of a pure compound have a similar shape, but crystals of different compounds may be of different shapes. This is useful: the shape is a clue to what the substance is. Some of the shapes in which crystals exist are shown in Fig. 3.11-1.
- Cubic, for example sodium chloride, \(\\mathrm{NaCl}\)
- Rhombohedric, for example copper sulphate, \(\\mathrm{CuSO\_4}\)
- Octahedral, for example iron sulphate, \(\\mathrm{FeSO\_4}\)
- Prismatic, with sharp edges, for example potassium nitrate, \(\\mathrm{KNO\_3}\)
Diagram
Fig. 3.11-1 — Some Shapes Of Crystals Four separate three-dimensional line drawings in a row, lightly shaded to show the faces. A: a cube, labelled ‘cubic’. B: a rhombohedron — a cube leaning over, all faces equal rhombi — labelled ‘rhombohedric’. C: an octahedron, two square pyramids joined base to base, labelled ‘octahedral’. D: a long six-sided prism with slanted, sharply pointed ends, labelled ‘prismatic’.
Crystals form in nature too, and several are recovered from the earth as minerals — common salt, potassium nitrate, ruby, sapphire and diamond among them. In some cases their beautiful colours are due to traces of impurities present within the crystals. Pure corundum is colourless; a trace of chromium in it makes a ruby red.
Obtaining Crystals In The Laboratory
In a laboratory, crystals may be obtained by the following methods.
- By cooling a hot saturated solution.
- By slowly evaporating a saturated solution.
- By cooling a fused mass.
- By sublimation.
The first is the method used most often, and the experiment below shows how to make it yield large crystals rather than a scatter of small ones.
An Experiment To Grow A Large Crystal
Aim: to prepare large crystals of copper sulphate.
Procedure: make a saturated solution of copper sulphate at 80 °C and filter it, so that no dust or undissolved matter is left in it. Allow the filtrate to cool and leave it for a day. Next morning a number of tiny crystals will be seen lying at the bottom of the beaker. Pick up these small crystals and choose one that is well formed.
Now suspend that well-shaped crystal in the cooled saturated solution by means of a thread tied to a glass rod, as in Fig. 3.11-2. Cover the beaker with a piece of paper to prevent dust from getting in, leave the beaker undisturbed, and watch the crystal grow.
Diagram
Fig. 3.11-2 — A Growing Crystal Of CuSO₄ Beaker of solution seen from the side, with a glass rod lying across its mouth. A thread hangs from the rod into the liquid, carrying a single well-formed crystal suspended in mid-solution. Label the rod, the thread, the crystal as ‘growing crystal’, and the liquid as ‘copper sulphate solution’. Show a paper cover resting over the top of the beaker.
Observation: after a few days the crystal will be found to have grown quite big.
The suspended crystal acts as a seed crystal — the solute coming out of solution builds on to it instead of forming fresh crystals elsewhere. This process of inducing crystallisation by adding a crystal of a pure substance into its saturated solution is called seeding.
Caution: keep the solution covered and avoid exposure to dust particles, since crystals will otherwise begin to form around the dust as well, and the growing crystal will be starved.
3.12 Water Of Crystallisation
A crystal of copper sulphate is dry to the touch and can be kept in a bottle for years. Heat it, though, and drops of water appear on the cool glass above it. That water was inside the crystal all along.
Some solids, while crystallising out from their solutions, combine with a definite quantity of water. These loosely held water molecules, present along with the salt, are known as water of crystallisation, or water of hydration.
Two words in that definition matter. ‘Fixed’ means the number of water molecules is definite — five for copper sulphate, never four or six. ‘Integral part of the structure’ means the water is built into the crystal, and is not moisture clinging to the outside.
The water of crystallisation is in loose chemical combination with the salt, and can be driven off as steam by heating the crystals above 100 °C. If the vapours produced are condensed and the liquid formed is tested, it gives a positive test for water. In this process the crystals lose their water and become anhydrous, which means all the water combined in the crystalline material has been removed.
An Experiment With Blue Vitriol
Aim: to show that hydrated copper sulphate crystals, known as blue vitriol, contain water of crystallisation.
Procedure: take some powdered copper sulphate crystals in a clean, dry test tube and clamp it in a tilted position, as in Fig. 3.12-1, so that the condensed water cannot trickle back on to the hot solid. Gently heat the crystals. The test tube is held tilted for a reason of safety as much as of method: water running back on to hot glass would crack the tube.
Diagram
Fig. 3.12-1 — To Show Loss Of Water Of Crystallisation On Heating Test tube clamped at a downward-sloping angle, mouth lower than the base. Label the solid at the closed lower end ‘copper sulphate crystals’, with a burner flame beneath it. Show small droplets forming along the cooler upper part of the tube near the mouth, labelled ‘drops of water’. Include the clamp and stand.
Observation: drops of colourless liquid condense on the cooler parts of the test tube. The residue left behind is anhydrous — without water — and amorphous, that is, non-crystalline, with no definite shape or structure. The blue crystals have turned into a white powder.
Add a few drops of water to that white powder and the blue solid returns. The change runs both ways:
\[\\mathrm{CuSO\_4!\\cdot!5H\_2O}\\ \\text{(blue)} \\rightleftharpoons \\mathrm{CuSO\_4}\\ \\text{(white)} + 5,\\mathrm{H\_2O}\]
The forward change takes place on heating, and the backward change on cooling and adding water. The blue colour, then, belongs not to copper sulphate but to its water of crystallisation.
While preparing crystals of a hydrated salt, one must be careful not to heat them too strongly in order to dry them. Heat them hard and the product is a dehydrated powder, not crystals at all.
Finding How Much Water A Crystal Holds
Since the water is a fixed part of the crystal, its amount can be measured. Take a known weight of crystals of a hydrated salt in a china dish and heat it above 100 °C. Keep heating and weighing until the weight of the residue becomes constant, then stop; a constant weight is the signal that no water is left to be driven off.
Weight of crystals at room temperature = \(a\) g Weight of residue after heating = \(b\) g Weight of water driven off = \((a - b)\) g
Now \(a\) g of crystals contained \((a-b)\) g of water, so
\[%\\ \\text{of water of crystallisation} = \\frac{a - b}{a} \\times 100\]
3.12.1 Hydrated Substances
Substances which contain water molecules along with the salt are hydrated substances. Sodium carbonate decahydrate, \(\\mathrm{Na\_2CO\_3!\\cdot!10H\_2O}\), and copper sulphate pentahydrate, \(\\mathrm{CuSO\_4!\\cdot!5H\_2O}\), are two of them.
The crystals of such salts contain water of crystallisation, and this water gives the crystals their shape. In some cases it also gives them their colour, as the blue of copper sulphate crystals shows. Salts of this kind are known as hydrated salts, and some of the commonest are listed in Table 3.8.
Table 3.8 — Some salts containing water of crystallisation
| Common name | Chemical name | Formula | Colour |
|---|---|---|---|
| Washing soda crystals | Sodium carbonate decahydrate | \(\\mathrm{Na\_2CO\_3!\\cdot!10H\_2O}\) | White |
| Epsom salt | Magnesium sulphate heptahydrate | \(\\mathrm{MgSO\_4!\\cdot!7H\_2O}\) | White |
| Potash alum | Hydrated potassium aluminium sulphate | \(\\mathrm{K\_2SO\_4!\\cdot!Al\_2(SO\_4)\_3!\\cdot!24H\_2O}\) | White |
| Glauber’s salt | Sodium sulphate decahydrate | \(\\mathrm{Na\_2SO\_4!\\cdot!10H\_2O}\) | White |
| Blue vitriol | Copper(II) sulphate pentahydrate | \(\\mathrm{CuSO\_4!\\cdot!5H\_2O}\) | Blue |
| Green vitriol | Iron(II) sulphate heptahydrate | \(\\mathrm{FeSO\_4!\\cdot!7H\_2O}\) | Bluish green |
| Plaster of Paris | Calcium sulphate semihydrate | \(\\mathrm{CaSO\_4!\\cdot!\\tfrac{1}{2}H\_2O}\) | White |
| White vitriol | Zinc sulphate heptahydrate | \(\\mathrm{ZnSO\_4!\\cdot!7H\_2O}\) | White |
| Gypsum | Hydrated calcium sulphate | \(\\mathrm{CaSO\_4!\\cdot!2H\_2O}\) | Colourless |
3.12.2 Anhydrous Substances
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 from which the water of crystallisation has been removed. Sodium chloride, \(\\mathrm{NaCl}\), and potassium nitrate, \(\\mathrm{KNO\_3}\), contain no water along with the salt, and are anhydrous as they stand.
The word is used in one other sense. ‘Anhydrous’ can also refer to the gaseous form of some substances that are commonly met as concentrated solutions — anhydrous ammonia, for instance, to distinguish the gas from its aqueous solution. In the same way, table salt is anhydrous sodium chloride, and gaseous hydrogen chloride is anhydrous.
Anhydrous substances are used to carry out certain chemical reactions that cannot proceed in the presence of water. To obtain them, the water of crystallisation is carefully removed from a hydrated salt, and this can be done by:
- direct heating;
- heating in dry and hot air;
- heating under vacuum;
- using dehydrating or desiccating agents, such as warm concentrated sulphuric acid.
Crystalline Does Not Mean Hydrated
It is tempting to assume that anything with a good crystal shape must owe that shape to water of crystallisation. It is not so. A great many crystalline solids crystallise from water without holding any water of crystallisation at all.
Sodium nitrate, potassium chloride, lead chloride, potassium sulphate, common salt, nitre, sugar, potassium permanganate and ammonium chloride are all crystalline, and none of them contains water of crystallisation.
An Experiment With Two Dry Salts
Aim: to study the action of heat on crystals of potassium chloride and sodium nitrate.
Procedure: gently heat a few crystals of pure potassium chloride in one test tube, and a few crystals of sodium nitrate in another.
Observation: on heating, the crystals of potassium chloride produce a crackling sound, called decrepitation. It is caused by bigger crystals of potassium chloride breaking up into smaller ones.
The crystals of sodium nitrate first melt into a colourless liquid, and then, on strong heating, produce oxygen, which rekindles a glowing splint. A light yellow solid residue of sodium nitrite is left behind:
sodium nitrate → sodium nitrite + oxygen
\[2,\\mathrm{NaNO\_3}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{NaNO\_2}\\text{(s)} + \\mathrm{O\_2}\\text{(g)}\]
Conclusion: neither substance contains water of crystallisation, since no water vapour is produced on heating either of them. Both changes are real changes — one physical, one chemical — but neither gives water.
This gives a simple general test: substances which do not produce water vapour on heating, that is, no tiny droplets of liquid on the cooler part of the test tube, do not contain water of crystallisation.
3.13 Efflorescence, Deliquescence And Hygroscopy
Leave three different salts in three open saucers on the same shelf, and three different things will happen. One will crumble to a dry powder, one will end up as a puddle, and one will simply turn damp. All three changes involve water moving between the salt and the air; what differs is which way it moves, and how far it goes. Fig. 3.13-1 sets them side by side.
Diagram
Fig. 3.13-1 — Three Ways A Salt Reacts To The Air Three panels in a row, each showing a watch glass on a bench with a ‘before’ and ‘after’ pair joined by an arrow. Panel 1, headed ‘efflorescence’: well-shaped crystals → a heap of dry loose powder, with a small arrow leaving the solid labelled ‘water vapour lost to dry air’. Panel 2, headed ‘deliquescence’: crystals → a pool of clear liquid on the watch glass, with an arrow entering the solid labelled ‘moisture absorbed from air’. Panel 3, headed ‘hygroscopy’: a solid → the same solid looking damp but still solid, with an arrow entering it labelled ‘moisture absorbed, but no solution formed’.
3.13.1 Efflorescence
Since the water of crystallisation is what holds the crystal’s shape, losing it results in the loss of the crystalline shape, and finally in the crumbling of the crystals into a powder. Substances showing efflorescence are called efflorescent substances.
Whether a salt effloresces depends on a contest between two vapour pressures. Efflorescence occurs when the vapour pressure inside the hydrated crystals exceeds the vapour pressure of the water vapour in the atmosphere, so the water leaves the crystal. In humid weather the air already holds a great deal of vapour, so efflorescence is minimised.
Examples Of Efflorescent Substances
(a) Washing soda, which is hydrated sodium carbonate, becomes a monohydrate when exposed to dry air — it loses nine of its ten water molecules, not all ten:
\[\\mathrm{Na\_2CO\_3!\\cdot!10H\_2O} \\xrightarrow{\\ \\text{dry air}\\ } \\mathrm{Na\_2CO\_3!\\cdot!H\_2O} + 9,\\mathrm{H\_2O}\]
(b) Glauber’s salt, sodium sulphate decahydrate, becomes a powdery anhydrous sodium sulphate on exposure to air, losing all ten:
\[\\mathrm{Na\_2SO\_4!\\cdot!10H\_2O} \\xrightarrow{\\ \\text{dry air}\\ } \\mathrm{Na\_2SO\_4} + 10,\\mathrm{H\_2O}\]
(c) Epsom salt, magnesium sulphate heptahydrate, becomes a monohydrate when exposed to dry air:
\[\\mathrm{MgSO\_4!\\cdot!7H\_2O} \\xrightarrow{\\ \\text{dry air}\\ } \\mathrm{MgSO\_4!\\cdot!H\_2O} + 6,\\mathrm{H\_2O}\]
The higher the temperature of the air, the greater is the efflorescence. Warm air can hold more water vapour, and warm air of a given moisture content is therefore further from being saturated, so it draws water out of the crystals more readily.
3.13.2 Deliquescence
A substance that behaves in this way is called a deliquescent substance, and the phenomenon is called deliquescency. Notice how far it goes: the substance does not merely get damp, it ends up as a solution of itself in water it has taken from the air.
The contest between vapour pressures runs the other way here. Deliquescence occurs when the vapour pressure inside the crystals is very low compared with the vapour pressure of the atmospheric air, so water moves from the air into the salt. In dry weather, when the air holds little vapour, deliquescence is minimised.
Caustic soda, \(\\mathrm{NaOH}\), caustic potash, \(\\mathrm{KOH}\), magnesium chloride, \(\\mathrm{MgCl\_2}\), zinc chloride, \(\\mathrm{ZnCl\_2}\), calcium chloride, \(\\mathrm{CaCl\_2}\), and ferric chloride, \(\\mathrm{FeCl\_3}\), are all deliquescent. Bottles of these are kept tightly stoppered, and the first two are handled with care in any case, since caustic soda and caustic potash attack the skin.
Why The Salt Cellar Turns Damp
Table salt turns moist and ultimately forms a solution on exposure to air, especially during the rainy season. Yet pure sodium chloride is not deliquescent at all. The commercial version of the salt contains impurities, chiefly magnesium chloride and calcium chloride, and it is these that are deliquescent.
That kind of impurity can be removed. A current of dry hydrogen chloride gas is passed through a saturated solution of the affected salt, and pure sodium chloride comes out as a precipitate. It is then recovered by filtering, and washed first with a little water and finally with alcohol.
3.13.3 Hygroscopy
Such substances are called hygroscopic substances. The stopping point is the whole difference between this and deliquescence: moisture is absorbed, but the substance stays a substance and never dissolves in what it has absorbed.
Some examples of hygroscopic substances are:
- concentrated sulphuric acid, \(\\mathrm{H\_2SO\_4}\);
- phosphorus pentoxide, \(\\mathrm{P\_2O\_5}\);
- quicklime, \(\\mathrm{CaO}\);
- silica gel.
Hygroscopic substances take up moisture from the air, and they may be either solids or liquids — quicklime is a solid, concentrated sulphuric acid a liquid. The three phenomena of this section are compared in Table 3.9.
3.14 Drying Agents And Dehydrating Agents
A gas collected in the laboratory is almost always damp, and many reactions will not work in the presence of water. Two kinds of substance are used to get rid of that water, and although they are often confused, they do quite different jobs.
Drying Agents
Quicklime, \(\\mathrm{CaO}\), anhydrous calcium chloride, anhydrous zinc chloride, phosphorus pentoxide, magnesium sulphate and dry sodium sulphate are all used to absorb water vapour from the air. Drying agents are also called desiccants, or desiccating agents.
Almost all hygroscopic substances are desiccating agents — concentrated sulphuric acid, phosphorus pentoxide, silica gel and quicklime among them. The connection is not a coincidence: a substance that takes moisture from the air will take it just as willingly from a stream of gas passed over it.
Dehydrating Agents
This is a much stronger action. A drying agent takes water that was only clinging to a substance; a dehydrating agent takes water that was part of the substance, so a new substance is left behind.
Concentrated sulphuric acid is the standard example. It removes the chemically combined water molecules from blue vitriol at room temperature — the same change you brought about with a burner in Section 3.12:
\[\\mathrm{CuSO\_4!\\cdot!5H\_2O}\\ \\text{(blue)} \\xrightarrow{\\ \\mathrm{conc.\\ H\_2SO\_4}\\ } \\mathrm{CuSO\_4}\\ \\text{(white)} + 5,\\mathrm{H\_2O}\]
A dehydrating agent can go further still. It can pull out hydrogen and oxygen in the ratio 2 : 1, in the form of water, from a compound that contained no water at all as such. Cane sugar treated with concentrated sulphuric acid turns into a black column of carbon:
cane sugar → carbon + water
\[\\mathrm{C\_{12}H\_{22}O\_{11}} \\xrightarrow{\\ \\mathrm{conc.\\ H\_2SO\_4}\\ } 12,\\mathrm{C} + 11,\\mathrm{H\_2O}\]
There was never any water inside the sugar molecule. The acid has taken the hydrogen and oxygen atoms out in the proportion in which they occur in water, and what is left is sugar charcoal. Alumina, \(\\mathrm{Al\_2O\_3}\), and phosphorus pentoxide, \(\\mathrm{P\_2O\_5}\), are two more dehydrating substances.
The two kinds of agent are compared in Table 3.10.
The second entry in the right-hand column is worth writing out. Formic acid treated with concentrated sulphuric acid loses the elements of water and leaves carbon monoxide:
\[\\mathrm{HCOOH} \\xrightarrow{\\ \\mathrm{conc.\\ H\_2SO\_4}\\ } \\mathrm{CO} + \\mathrm{H\_2O}\]
Concentrated sulphuric acid appears in both columns of Table 3.10, and that is not an error. Which role it plays depends on what you hand it.
Drying Gases, Liquids And Solids
Gases
Gases are dried in three ways.
- By passing the gas through concentrated sulphuric acid in a drying bottle, since the acid has a strong affinity for water and easily extracts it from many substances. Concentrated sulphuric acid is used to dry gases which are acidic in nature, such as hydrogen chloride gas.
- By passing the gas through a drying tower or a U-tube containing quicklime or anhydrous sodium sulphate.
- By passing the gas through a drying bulb containing anhydrous calcium chloride.
The apparatus for all three is shown in Fig. 3.14-1.
Diagram
Fig. 3.14-1 — Apparatus For Drying Gases Four labelled drawings in a row, each with an inlet arrow on the left and an outlet arrow on the right. (a) A drying bottle, half filled with liquid, the inlet tube dipping below the liquid surface and the outlet tube above it; label ‘drying bottle containing concentrated sulphuric acid’. (b) A vertical drying tower packed with lumps, gas entering at the base and leaving at the top; label ‘drying tower’. (c) A U-tube packed with granules, held in a stand; label ‘U-tube’. (d) A round drying bulb packed with granules between two side tubes; label ‘drying bulb’.
The choice of drying agent is not free — it depends on the nature of the gas being dried. Quicklime, being basic in nature, is suitable for drying ammonia, which is a basic gas, because a basic gas will not react with a base. For the same reason, concentrated sulphuric acid must never be used to dry ammonia: the acid would react with the gas instead of merely drying it, and the gas would be lost.
Liquids
Liquids are dried by keeping them overnight over anhydrous sodium sulphate, magnesium sulphate or calcium chloride at room temperature. The solid, which has taken up the water, is then removed by filtration.
Solids
Solids are dried by spreading them on a watch glass or a dish and keeping it in a desiccator for some days.
A desiccator is an air-tight glass vessel with a suitable drying agent, usually anhydrous calcium chloride, placed at the bottom. It is used for drying the solids placed in it, as Fig. 3.14-2 shows. Being air-tight, it keeps the surrounding air out, so the drying agent has only the moisture of the solid to deal with.
Diagram
Fig. 3.14-2 — A Desiccator Thick-walled glass vessel with a heavy ground-glass lid carrying a knob, drawn in cross-section. Inside, a perforated plate rests on a ledge and carries a small dish; below the plate a granular layer covers the base, labelled ‘anhydrous calcium chloride’. Label the lid and mark the joint ‘air-tight’.
3.15 Soft Water And Hard Water
Water is classified as soft or hard by a test anyone can carry out with a cake of soap. The chemistry behind the test comes later; the test itself comes first, because that is how the distinction was found in the first place.
Water is said to be soft if it readily forms lather with soap. Pure water, and water containing sodium salts, gives lather with soap easily. Such a sample is called soft water and is suitable for use. Distilled water and rain water are examples of soft water.
Water is said to be hard when it does not readily form lather with soap. Fig. 3.15-1 shows the difference between the two.
Diagram
Fig. 3.15-1 — Soft And Hard Water Two pairs of vessels side by side. Left pair, headed ‘soft’: a test tube of water labelled ‘rain water’, a plus sign, a small bottle labelled ‘liquid soap’, an equals sign, and a glass showing a thick head of bubbles labelled ‘forms many suds’. Right pair, headed ‘hard’: a test tube labelled ‘river water’, plus the same liquid soap, equals a glass with only a few bubbles and a scum on the surface, labelled ‘forms few suds’.
3.15.1 Causes Of Hardness
The water of some springs, wells and rivers contains dissolved mineral matter, picked up from the rock it has travelled through. If that dissolved matter happens to be the hydrogen carbonates, sulphates or chlorides of calcium and magnesium, the water gets hard.
Hardness in water is due to the presence of bicarbonates, chlorides or sulphates of calcium or magnesium. No other dissolved salt causes it: sodium salts, as you have just seen, leave the water soft.
How The Salts Get Into The Water
Water flowing over beds of limestone and dolomite, which are calcium carbonate and magnesium carbonate, does not simply run over them. Rain water carries dissolved carbon dioxide, and the carbon dioxide slowly converts some of these insoluble carbonates into soluble bicarbonates, which are then carried along with the water and make it hard.
The calcium reaction is the one you met in Section 3.3, where marine animals used the same soluble bicarbonate to build their shells:
\[\\mathrm{CaCO\_3}\\ \\text{(insoluble)} + \\mathrm{CO\_2} + \\mathrm{H\_2O} \\rightarrow \\mathrm{Ca(HCO\_3)\_2}\\ \\text{(soluble)}\]
Magnesium carbonate behaves in exactly the same way:
magnesium carbonate + carbon dioxide + water → magnesium bicarbonate
\[\\mathrm{MgCO\_3}\\ \\text{(insoluble)} + \\mathrm{CO\_2} + \\mathrm{H\_2O} \\rightarrow \\mathrm{Mg(HCO\_3)\_2}\\ \\text{(soluble)}\]
A different mineral gives hardness of another kind. Where water flows over beds of gypsum, \(\\mathrm{CaSO\_4!\\cdot!2H\_2O}\), a little of the gypsum dissolves in the water and makes it hard. Here no carbon dioxide is needed, and the salt that dissolves is a sulphate, not a bicarbonate — a difference that will matter a great deal in the next section.
3.15.2 Stalactites And Stalagmites
The same reaction that hardens water builds some of the strangest structures in nature. In some limestone caves one sees conical, pillar-like objects, some hanging from the roof of the cave and some rising from its floor.
These are formed by water dripping from cracks in the rocks, carrying dissolved calcium hydrogen carbonate. The release of pressure and the warmth inside the cave convert some of that calcium hydrogen carbonate back to calcium carbonate:
calcium hydrogen carbonate → calcium carbonate + carbon dioxide + water
\[\\mathrm{Ca(HCO\_3)\_2} \\rightarrow \\mathrm{CaCO\_3} + \\mathrm{CO\_2} + \\mathrm{H\_2O}\]
This calcium carbonate, little by little, gets deposited both on the roof and on the floor of the cave — on the roof where the drop hangs, and on the floor where it lands.
The conical pillar which grows downwards from the roof is called a stalactite, and the one which grows upward from the floor is called a stalagmite. Given long enough, the two meet and form pillar-like structures joining floor to roof, as in Fig. 3.15-2. The growth is extremely slow, often less than a centimetre in a year, so a large pillar represents a very long stretch of time.
Diagram
Fig. 3.15-2 — Stalactites And Stalagmites Interior of a limestone cave. Draw tapering conical formations hanging point-downwards from the roof, labelled ‘stalactite’ with an arrow, and matching cones rising point-upwards from the floor beneath them, labelled ‘stalagmite’ with an arrow. Show one pair that has joined into a complete pillar, and a single drop of water at the tip of one hanging cone.
3.16 Types Of Hardness And Its Removal
Hardness is not one thing. Two samples of water may both refuse to lather, and yet one of them can be softened by putting the kettle on, while the other will not soften however long you boil it. The difference lies in which salt is dissolved in it.
3.16.1 Temporary And Permanent Hardness
Hardness of water is of two types: temporary hardness and permanent hardness.
It is called temporary hard water because its hardness can be removed simply by boiling. The bicarbonates are unstable to heat, and boiling turns them into carbonates that will not stay dissolved.
These salts are not decomposed by heat, so this hardness cannot be removed by boiling at all, and a chemical has to be added instead. The two kinds are compared in Table 3.11.
3.16.2 Advantages Of Hard Water
Hard water is a nuisance in a boiler and a washtub, but it is not without its uses. Hard water containing some dissolved salts has the following advantages.
- Soft water free from dissolved salts has a very flat taste. The presence of salts in hard water makes it tasty, and hard water is therefore used in the preparation of beverages and wines.
- The calcium and magnesium salts present in small amounts in hard water are essential for the growth of our bones and teeth.
- Hard water checks the poisoning of water carried through lead pipes. When lead pipes are used, some lead salts dissolve in the water and make it poisonous. The calcium sulphate present in hard water forms insoluble lead sulphate, which builds up as a layer inside the lead pipe, and this layer checks the lead poisoning.
3.16.3 Disadvantages Of Hard Water
Hard Water Is Not Suitable For Producing Steam
Steam is usually made in boilers, which are built of a number of narrow copper tubes surrounded by fire. As the cold water enters these tubes it is immediately changed into steam, but the dissolved solids cannot change into vapour, and they get deposited on the inner walls of the tubes.
This goes on, and the bore of the tubes becomes narrower and narrower, as Fig. 3.16-1 shows. Eventually less water flows through the tubes at one time, and less steam is produced. When the bore becomes very narrow, the pressure of steam inside increases so much that at times the boiler itself bursts.
Diagram
Fig. 3.16-1 — Boiler Scale Due To Hard Water Two thick boiler tubes shown end-on, side by side, each with a flanged rim and bolt holes. The left tube shows a clean, wide circular bore. The right tube shows the same bore heavily narrowed by a thick, rough crust lining the inside wall; label the crust ‘boiler scale’ and mark the reduced opening ‘narrowed bore’.
The deposit is called boiler scale, and it does a second kind of damage. Being a bad conductor of heat, it allows only part of the outside heat to go in and convert the water into steam. This results in loss of efficiency of the boiler and wastage of fuel. To minimise the loss, the boiler scale is removed from time to time using a boiler compound, or else the water is softened before it is used. Both cost money.
Furring, Deposits And Damage
The furring of the tea kettle is caused by the same sediment, formed on its walls from boiling hard water. This fur is nothing but the carbonates of calcium and magnesium. Hard water also spoils taps, washing machines and geysers.
The full list of the disadvantages of hard water is worth having in one place:
- it causes corrosion in boilers;
- it causes wastage of soap;
- it gives a dull appearance to clothes;
- it causes limescale to form in pipes, water heaters and appliances such as dishwashers and washing machines, and scale build-up in these appliances decreases their efficiency;
- mineral build-up can cause pipes and plumbing systems to corrode, which can lead to water leaks;
- it leaves mineral deposits on taps, sinks, bathtubs and toilets, which are difficult to remove;
- the minerals left as a residue on skin and hair can leave the skin and scalp dry and irritated, and the hair dull and brittle.
3.16.4 Soap, Detergents And Hard Water
To see why hard water wastes soap, one has to know what soap is. Soap is chemically a sodium salt of stearic acid, an organic acid with the formula \(\\mathrm{C\_{17}H\_{35}COOH}\). Soap therefore has the formula \(\\mathrm{C\_{17}H\_{35}COONa}\), sodium stearate, which for convenience is written as NaSt.
If the water is hard, the calcium and magnesium ions in it combine with the negative ions of the soap to form a slimy precipitate of insoluble calcium and magnesium salts. This precipitate is given the name soap-curd, or scum.
In temporary hard water:
\[2,\\mathrm{NaSt} + \\mathrm{Ca(HCO\_3)\_2} \\rightarrow \\mathrm{CaSt\_2}\\downarrow + 2,\\mathrm{NaHCO\_3}\]
\[2,\\mathrm{NaSt} + \\mathrm{Mg(HCO\_3)\_2} \\rightarrow \\mathrm{MgSt\_2}\\downarrow + 2,\\mathrm{NaHCO\_3}\]
In permanent hard water the formation of soap-curd will go on as long as there are calcium and magnesium ions present:
\[2,\\mathrm{NaSt} + \\mathrm{CaCl\_2} \\rightarrow \\mathrm{CaSt\_2}\\downarrow + 2,\\mathrm{NaCl}\]
\[2,\\mathrm{NaSt} + \\mathrm{MgCl\_2} \\rightarrow \\mathrm{MgSt\_2}\\downarrow + 2,\\mathrm{NaCl}\]
\[2,\\mathrm{NaSt} + \\mathrm{CaSO\_4} \\rightarrow \\mathrm{CaSt\_2}\\downarrow + \\mathrm{Na\_2SO\_4}\]
\[2,\\mathrm{NaSt} + \\mathrm{MgSO\_4} \\rightarrow \\mathrm{MgSt\_2}\\downarrow + \\mathrm{Na\_2SO\_4}\]
As long as the formation of soap-curd continues, no soap lather will be formed, and the cleaning of cloth or of the body is not possible. The soap is being used up in making a precipitate instead of doing its work. Worse, these precipitates are difficult to wash out of the fabric, since they are insoluble in water, and they sometimes leave rusty spots if iron salts are also present in the water.
Synthetic Detergents
This difficulty in cleaning clothes has been solved to a great extent by using synthetic detergents. Detergents such as the familiar washing powders are used in place of soap in domestic and laundry work. They are made by treating long-chain hydrocarbon compounds with sulphuric acid, and then converting the products into their sodium salts using caustic soda.
Detergents are more soluble in water than soap, and they are unaffected by the hardness of water, because their calcium and magnesium salts are soluble in water and so no scum is formed. This is why detergents are also called soapless soaps.
How Soaps And Detergents Clean
When a soap or a detergent is dissolved in water, the molecules gather into clusters called micelles, with their hydrocarbon tails pointing inwards and their charged heads pointing outwards into the water.
In cleansing, the hydrocarbon tail attaches itself to the oil and dirt. When the water is stirred, the oil and dirt tend to lift off the dirty surface and break into fragments, which gives an opportunity to other tails to stick to them as well. The solution now contains small globules of oil and dirt, each surrounded by soap or detergent molecules, as in Fig. 3.16-2.
The negatively charged heads facing the water keep the small globules from coming together and forming larger aggregates. So the globules stay separate and suspended, and the oil and dirt are carried away when the clothes are rinsed with water.
Diagram
Fig. 3.16-2 — Cleansing Action Of Soap Two drawings, one above the other. Upper: a piece of cloth in water with a dark blob on it labelled ‘oily dirt’; several soap molecules drawn as a small open circle (the head) with a wavy line (the tail), their tails buried in the blob and heads out in the water; label one molecule ‘hydrocarbon tail’ and one ‘head’. Lower: a single round globule of dirt lifted off and completely surrounded by soap molecules, all tails pointing inward and all heads pointing outward; label ‘a micelle’.
3.16.5 Removal Of Temporary Hardness
There are three ways of doing it, and which one is chosen depends on how much water has to be treated. The choices for both kinds of hardness are set out in FC 3.1.
Flowchart
FC 3.1 — Choosing A Method Of Softening Hard Water Sample of hard water → test what salts it contains → branch A, only bicarbonates present (temporary hardness) → boiling for small amounts, or slaked lime by Clark’s process, or washing soda → soft water; branch B, sulphates or chlorides present (permanent hardness) → washing soda, with slaked lime added as well if magnesium sulphate is present → soft water; branch C, both kinds present and a large quantity to be treated → permutit process → soft water, and for scientific work → cation and anion exchange resins → completely deionised water
(1) By Boiling
On boiling, the carbon dioxide is driven off and the soluble hydrogen carbonates are converted into insoluble carbonates, which can then be removed by filtration or decantation:
\[\\mathrm{Ca(HCO\_3)\_2} \\xrightarrow{\\ \\text{boil}\\ } \\mathrm{CaCO\_3}\\downarrow + \\mathrm{H\_2O} + \\mathrm{CO\_2}\\uparrow\]
\[\\mathrm{Mg(HCO\_3)\_2} \\xrightarrow{\\ \\text{boil}\\ } \\mathrm{MgCO\_3}\\downarrow + \\mathrm{H\_2O} + \\mathrm{CO\_2}\\uparrow\]
The calcium ions and magnesium ions are thus removed from the water as insoluble calcium carbonate and magnesium carbonate, which are left behind as residue on filtration, and the water becomes soft. This is the same change that builds a stalagmite, and the same change that furs a kettle — here it is being used deliberately.
Boiling is not very practical when large quantities of water need to be softened. It is both costly and slow, and for a town’s water supply it is out of the question.
(2) By Addition Of Lime — Clark’s Process
A more practical method is to add a calculated quantity of slaked lime. The following reactions take place:
\[\\mathrm{Ca(HCO\_3)\_2} + \\mathrm{Ca(OH)\_2} \\rightarrow 2,\\mathrm{CaCO\_3}\\downarrow + 2,\\mathrm{H\_2O}\]
\[\\mathrm{Mg(HCO\_3)\_2} + \\mathrm{Ca(OH)\_2} \\rightarrow \\mathrm{MgCO\_3}\\downarrow + \\mathrm{CaCO\_3}\\downarrow + 2,\\mathrm{H\_2O}\]
The word ‘calculated’ is important. Too little lime leaves the water hard; too much leaves unreacted lime dissolved in it, which is itself an impurity.
The lime is first thoroughly mixed with water in a tank, and then fed into another tank containing the hard water. Revolving paddles mix the two solutions thoroughly. Most of the calcium carbonate settles down, and if there is any solid left over it is removed by filtration. This process goes by the name of Clark’s process.
(3) By Addition Of Washing Soda
When washing soda, or soda ash, which is anhydrous sodium carbonate, is added to hard water, the corresponding insoluble carbonates settle down and can be removed by filtration:
\[\\mathrm{Ca(HCO\_3)\_2} + \\mathrm{Na\_2CO\_3} \\rightarrow \\mathrm{CaCO\_3}\\downarrow + 2,\\mathrm{NaHCO\_3}\]
\[\\mathrm{Mg(HCO\_3)\_2} + \\mathrm{Na\_2CO\_3} \\rightarrow \\mathrm{MgCO\_3}\\downarrow + 2,\\mathrm{NaHCO\_3}\]
The two routes for the two kinds of hardness are summarised in Fig. 3.16-3.
Diagram
Fig. 3.16-3 — Softening Of Hard Water Two sets of three vessels. Upper set, headed ‘temporary hard water’: a beaker labelled ‘temporary hard water — with bicarbonate mineral’, an arrow labelled ‘when boiled’ over a small burner flame, and a beaker labelled ‘becomes soft water’; caption beneath ‘by heating’. Lower set, headed ‘permanent hard water’: a beaker labelled ‘permanently hard water — with sulphate or chloride mineral’, an arrow labelled ‘with washing soda’ over a small heap of powder being added, and a beaker labelled ‘becomes soft water’; caption beneath ‘by adding chemicals’.
3.16.6 Removal Of Permanent Hardness
To remove the permanent hardness of water, it is treated with a calculated quantity of soda ash, when the following reactions take place:
\[\\mathrm{CaSO\_4} + \\mathrm{Na\_2CO\_3} \\rightarrow \\mathrm{CaCO\_3}\\downarrow + \\mathrm{Na\_2SO\_4}\]
\[\\mathrm{CaCl\_2} + \\mathrm{Na\_2CO\_3} \\rightarrow \\mathrm{CaCO\_3}\\downarrow + 2,\\mathrm{NaCl}\]
The sodium sulphate left in solution is not harmful to boilers or to clothing, which is the point of the exercise: the troublesome calcium has been taken out and replaced by a harmless sodium salt.
If, however, some magnesium sulphate is also present in the water, lime has to be used besides soda ash, since magnesium carbonate is not fully precipitated. In that case the reaction is:
\[\\mathrm{MgSO\_4} + \\mathrm{Na\_2CO\_3} + \\mathrm{Ca(OH)\_2} \\rightarrow \\mathrm{Mg(OH)\_2}\\downarrow + \\mathrm{CaCO\_3}\\downarrow + \\mathrm{Na\_2SO\_4}\]
All the methods described so far are suitable for domestic purposes only, where small amounts of water have to be treated.
3.16.7 The Permutit And Ion-Exchange Processes
If the hardness is both temporary and permanent, and particularly when large amounts of water are concerned, the permutit process is used.
Permutit is an artificial zeolite. Chemically it is hydrated sodium aluminium orthosilicate, having the formula \(\\mathrm{Na\_2Al\_2Si\_2O\_8!\\cdot!xH\_2O}\). Since that formula is awkward to write in an equation, it is given the short form \(\\mathrm{Na\_2P}\).
A tall cylinder is loosely filled with lumps of permutit, as in Fig. 3.16-4. When hard water containing calcium and magnesium ions percolates through these lumps, an exchange of ions takes place. The sodium permutit is slowly changed into calcium and magnesium permutit, and since the calcium and magnesium ions have been taken out of the water and sodium ions put in their place, the water becomes soft.
Diagram
Fig. 3.16-4 — Softening Of Hard Water By The Permutit Process Tall closed cylinder drawn in cross-section, standing upright. At the top, an inlet pipe labelled ‘inlet for hard water’; below it a space labelled ‘hard water’; the main body packed with rounded lumps labelled ‘permutit’; near the base a layer labelled ‘glass wool’; at the bottom an outlet pipe labelled ‘outlet for soft water’.
When it is no longer active — that is, when nearly all its sodium has been exchanged away — the permutit is regenerated by running a concentrated solution of brine over it, and the calcium chloride formed is removed by repeated washing:
\[\\mathrm{CaP} + 2,\\mathrm{NaCl} \\rightarrow \\mathrm{Na\_2P} + \\mathrm{CaCl\_2}\]
The same bed of permutit can therefore be used over and over again, as FC 3.2 shows.
Flowchart
FC 3.2 — The Permutit Cycle Sodium permutit in the cylinder → hard water percolates through → calcium and magnesium ions taken up, sodium ions released → soft water drawn off at the outlet → bed gradually becomes calcium and magnesium permutit and stops working → concentrated brine run over the bed → sodium ions restored and calcium chloride washed away → sodium permutit ready again, returning to the start
Zeolite is a naturally formed mineral, found in abundance in large deposits of volcanic tuff. It can easily be dried, and so it too can be used for ion exchange.
Ion Exchange Resins
For laundries and boilers, water softened by the permutit process is quite satisfactory. For scientific purposes, however, even the presence of the positive sodium ion and of the negative bicarbonate and sulphate ions is not desirable.
To remove these ions as well, the water is passed one by one through synthetically prepared cation exchange and anion exchange resins. The sodium ions are replaced by \(\\mathrm{H^{+}}\) ions, and the \(\\mathrm{HCO\_3^{-}}\) and \(\\mathrm{SO\_4^{2-}}\) ions by \(\\mathrm{OH^{-}}\) ions. Since \(\\mathrm{H^{+}}\) and \(\\mathrm{OH^{-}}\) together make water itself, nothing is added to the water at all, and it gets completely deionised.
3.17 Potable Water
Every plant, animal and human being needs water to stay alive. All life processes, from taking in food to getting rid of waste, require it. We depend on water for our daily activities too — for brushing teeth, cooking food and washing dishes at home, in factories to manufacture almost everything from cars to books, and in agriculture to irrigate crops.
But not all water is fit to drink, and clearness is no guarantee, as the opening story of this chapter showed.
Potable water is treated, filtered, and made free from harmful microorganisms, parasites, toxic chemicals and pollutants. It must be clean, clear, free from unpleasant taste or odour, and safe for regular use.
What Drinking Water Must Satisfy
The standards below are those of the Indian drinking water standards specified by the Bureau of Indian Standards, which are broadly aligned with the guidelines issued by the World Health Organization.
Table 3.12 — Characteristics of potable water
| Property | Requirement |
|---|---|
| Appearance | Clear and transparent |
| Colour | Colourless |
| Odour | Odourless |
| Taste | Pleasant, free from any chemical or salty taste |
| Turbidity | Very low or absent |
| pH | Between 6.5 and 8.5 |
| Total dissolved solids (TDS) | Should not exceed about 2000 mg/L |
| Total hardness | Generally within 300–600 mg/L |
| Chlorides | Within 250–1000 mg/L |
| Fluoride | Safe range 1.0–1.5 mg/L |
| Iron | Within 0.3–1.0 mg/L |
Alongside these, three safety requirements must be met: the water must be free from harmful microorganisms; it must be free from toxic substances such as arsenic, lead and other heavy metals; and it must be properly disinfected and safe for regular human consumption.
3.17.1 Making Water Potable
People want drinking water that is clear and without any objectionable taste or odour, and water in its natural state seldom has these qualities. So after water is collected from natural sources, it is usually passed through a treatment plant to make it safe for drinking. The plant may put the water through one process or several, depending on the quality of the untreated water and on the standards of the city concerned. A full plant is shown in Fig. 3.17-1, and the order of the main steps in FC 3.3.
Diagram
Fig. 3.17-1 — Steps For Making Water Potable Wide plant layout drawn left to right and continued on a second row. Row 1: a lake and river labelled ‘surface water’ with an intake pipe labelled ‘water intake’; a small pump house labelled ‘electric pumps’; a borewell labelled ‘groundwater’ with sampling flasks labelled ‘sampling and testing’; a tank with a dosing hopper above it labelled ‘addition of coagulants’ and the tank labelled ‘coagulation’; a long rectangular tank labelled ‘aeration and sedimentation’. Row 2, following on: a bed of layered material labelled ‘activated carbon, sand and gravel filters — filtration’; a set of small dosing units labelled ‘UV light’, ‘pH correction’ and ‘chlorine (Cl)’ over a tank labelled ‘disinfection and final treatment’; an overhead storage tank labelled ‘storage’; and houses at the end labelled ‘distribution’. Join all units with pipes carrying flow arrows.
Flowchart
FC 3.3 — Making Water Potable Raw untreated water → coagulation with alum, forming flocs → settling in a sedimentation basin, suspended impurities removed → filtration through beds of sand, gravel and coal → chlorination to destroy disease-causing bacteria → defluoridation where the fluoride level is above the safe limit → aeration to improve taste and odour → storage → distribution to homes
1. Coagulation And Settling
Raw untreated water is first mixed with chemicals called coagulants. The most commonly used coagulant is alum, a white crystalline hydrated double sulphate salt that has been used for centuries in water purification.
When alum is added to water it forms tiny sticky globs called flocs. Bacteria, mud and other impurities stick to these flocs. The water is then passed into a settling basin, where the flocs gradually become heavier and settle at the bottom. This process removes the suspended impurities from the water — the same coagulation of colloidal particles you met in Section 3.7, put to work on a city scale.
2. Filtration
After settling, the water is passed through certain filters. The filters usually consist of layers of sand, gravel and coal. As the water trickles down through these layers, the remaining particles or impurities are removed. The filtered water is then collected for further treatment, in order to kill microorganisms — for filtration removes what can be strained out, and bacteria cannot.
3. Chlorination
Chlorination is the process in which chlorine is added to water to kill harmful microorganisms. Chlorine is sometimes added before coagulation and settling, and often after filtration. It acts as a disinfectant and destroys disease-causing bacteria. Most cities use chlorination as an important step to ensure that drinking water is germ-free.
The amount added is carefully calculated: enough to destroy the bacteria, and little enough to leave the water safe and pleasant to drink.
4. Defluoridation
Fluoride is naturally present in many water sources. In some areas the level of fluoride in the water is higher than the safe limit, and then it must be brought down. Defluoridation is the process of removing excess fluoride from drinking water to achieve safe fluoride levels and make the water fit for consumption. It prevents the health problems caused by high fluoride levels, such as dental and skeletal fluorosis.
Common techniques include adsorption, using activated alumina or charcoal; chemical precipitation, using alum; and membrane processes such as reverse osmosis and electrodialysis.
Other processes are also used to remove unpleasant tastes or smells, or to give the water special qualities. Aeration, in which air is bubbled through or the water is sprayed, improves the taste and odour of drinking water.
3.17.2 Contamination Of Drinking Water
Contamination means the unwanted entry of disease-causing germs or harmful substances into drinking water. Clean and germ-free drinking water is essential for our health.
Contamination can occur in several ways.
- Sewage, if improperly disposed of, may leak into the soil and contaminate underground water supplies.
- Human waste left near lakes, rivers and ponds, where sanitation is lacking, can pollute water sources.
- Animal wastes and washings from dairies and poultry farms may drain into nearby water bodies.
- Industrial wastes containing toxic chemicals may mix with water and make it unsafe.
- Pesticides and chemical fertilisers washed from agricultural fields can pollute rivers and groundwater.
Some of these are more dangerous than others. Industrial wastes containing mercury can cause serious health problems. Tannery wastes may contain harmful bacteria and chemicals. The excessive use of pesticides can also pollute water and harm human health.
3.17.3 Diseases Caused By Polluted Water
Polluted water is a major cause of many diseases. When drinking water contains harmful microorganisms or toxic chemicals, it becomes unsafe for human consumption and can seriously affect health.
Bacterial Diseases
Many bacterial diseases spread through contaminated water, including diarrhoea, typhoid, cholera and dysentery. These illnesses generally affect the digestive system, and cause symptoms such as stomach pain, fever, vomiting, dehydration and weakness.
They are most common where clean drinking water and proper sanitation are not available. The presence of organic wastes in water allows bacteria to multiply rapidly, which increases the chances of infection.
Diseases Caused By Chemical Impurities
In addition to biological contamination, polluted water may carry harmful chemical substances. Industrial wastes, pesticides and heavy metals dissolve in water and make it toxic. Long-term consumption of such chemically contaminated water may increase the risk of cancer, and excessive intake of polluted water over long periods can also lead to organ damage, particularly affecting the liver, the kidneys and other vital organs.
Specific chemical impurities cause specific problems. High levels of fluoride may lead to tooth and bone disorders, while excess nitrates and heavy metals can disturb normal body functions. Such effects show that water pollution is not only a chemical problem but a serious biological one as well.
Both the microbial purity and the chemical purity of water are therefore essential for public health. It follows that we must avoid polluting water and use it judiciously.
3.18 Summary
- Water is a compound of hydrogen and oxygen, \(\\mathrm{H\_2O}\), chemically named dihydrogen oxide, with a molecular mass of 18 amu. The two hydrogen atoms are each joined to the oxygen atom by a shared pair of electrons.
- About 70 per cent of the Earth’s surface is covered with water, but hardly 2.5 per cent of the world’s water is fresh water, much of it frozen in the polar ice caps and glaciers.
- Water exists in all three physical states — ice, water and water vapour — and circulates between them through the water cycle. It occurs in nature in both the free and the combined states.
- Pure water is a clear, transparent liquid; it is colourless, odourless and tasteless. Any taste in water comes from the gases and solids dissolved in it.
- Pure water boils at 100 °C under a pressure of 760 mm of mercury and freezes at 0 °C at one atmosphere. Boiling point rises with increasing pressure and falls with decreasing pressure, which is why a pressure cooker cooks fast and water on a hill boils below 100 °C.
- Dissolved impurities raise the boiling point of water and lower its freezing point. An increase in pressure also lowers the freezing point.
- Water has its maximum density, \(1\\ \\text{g cm}^{-3}\), and its minimum volume at 4 °C. Below 4 °C it expands on cooling; this is the anomalous expansion of water.
- Water expands on freezing — 92 volumes of water become 100 volumes of ice — so ice, with a relative density of 0.92, floats. A pond therefore freezes from the top downwards, and aquatic life survives in the water below the ice.
- The latent heat of fusion of ice is 336 J/g, or 80 cal/g; the latent heat of vaporization of water is 2268 J/g, or 540 cal/g. The first is why lakes and rivers do not freeze suddenly, the second why steam burns are more severe than burns from boiling water.
- The specific heat capacity of water is high: 1 g of water absorbs 4.2 J, or 1 calorie, for a rise of 1 °C. This makes water a good coolant and allows large bodies of water to moderate the climate of nearby land.
- Water is called a universal solvent because it dissolves more substances than any other liquid. Its molecule is polar and it has a high dielectric constant, which weakens the attraction between oppositely charged ions and lets electrovalent compounds dissolve.
- Tap, river and well water leave concentric rings of solid on evaporation, showing that they contain dissolved solids; rain water and distilled water leave none. Dissolved salts give water its taste and supply minerals to plants and to our bodies.
- Air dissolves in all natural water. Oxygen is more soluble than nitrogen, so dissolved air is 33 per cent oxygen, 66 per cent nitrogen and 1 per cent carbon dioxide. Aquatic animals use the dissolved oxygen, and aquatic plants the dissolved carbon dioxide.
- A mixture is two or more substances physically mixed in any proportion and not chemically combined. It is homogeneous if its composition is uniform throughout, and heterogeneous if it is not and shows visible boundaries of separation.
- A solution is a homogeneous mixture; solution = solute + solvent. Where the components are in different states, the solvent is the one in the same state as the solution; where they are in the same state, the one in excess is the solvent.
- Solutions may be solid in liquid, liquid in liquid, gas in liquid or solid in solid. A solid solution of two metals is an alloy, such as brass, which is 70 per cent copper and 30 per cent zinc. A solution in alcohol is called a tincture.
- Particle size decides the kind of mixture: less than 1 nm in a true solution, 1 nm to 1000 nm in a colloidal solution, and more than 1000 nm in a suspension.
- A suspension is heterogeneous; its particles are visible, can be separated by ordinary filter paper, and settle under gravity.
- A colloidal solution has a dispersed phase and a dispersion medium, and is of eight types. Its particles do not settle, pass through filter paper but not through ultrafilters, carry an electric charge, show Brownian motion, and scatter light, giving the Tyndall effect.
- Neutralising the charge on colloidal particles makes them clump together and settle; this is coagulation. It explains why alum stops bleeding and why deltas form where a river meets the sea.
- A saturated solution can dissolve no more solute at a given temperature; an unsaturated solution can dissolve more; a supersaturated solution holds more solute than its saturated solution at that temperature and is metastable.
- Concentration is expressed as mass percent or volume percent: \[\\text{mass percent} = \\frac{\\text{mass of solute}}{\\text{mass of solution}} \\times 100 \\qquad \\text{volume percent} = \\frac{\\text{volume of solute}}{\\text{volume of solution}} \\times 100\]
- Solubility is the mass of solute, in grams, that saturates 100 g of solvent at a stated temperature: \[\\text{solubility} = \\frac{\\text{mass of solute}}{\\text{mass of solvent}} \\times 100\]
- Powdering and stirring speed up dissolving; only temperature changes the solubility itself. The solubility of most solids increases with rise in temperature, while the solubility of gases always decreases with rise in temperature.
- A solubility curve plots solubility against temperature. Potassium nitrate and sodium nitrate rise steeply, sodium chloride rises only slightly, and hydrated calcium sulphate falls. Glauber’s salt shows anomalous solubility, rising to 32.8 °C and then falling.
- The solubility of a gas increases with pressure. Henry’s law states that at any given temperature the mass of a gas dissolved by a fixed volume of a liquid is directly proportional to the pressure on the surface of the liquid.
- A crystal is a homogeneous solid of definite geometrical shape. Crystallisation is the process of obtaining crystals by cooling a hot saturated solution; crystals may also be got by evaporating a saturated solution, cooling a fused mass, or sublimation. Adding a crystal to induce crystallisation is called seeding.
- Water of crystallisation is the fixed amount of water that forms an integral part of the crystalline structure of certain salts. Salts containing it are hydrated; those without it are anhydrous.
- Heating blue vitriol drives its water off and leaves a white anhydrous powder, and adding water restores the blue: \[\\mathrm{CuSO\_4!\\cdot!5H\_2O}\\ \\text{(blue)} \\rightleftharpoons \\mathrm{CuSO\_4}\\ \\text{(white)} + 5,\\mathrm{H\_2O}\]
- If \(a\) g of hydrated crystals leave \(b\) g of residue on heating to constant weight, \[%\\ \\text{of water of crystallisation} = \\frac{a - b}{a} \\times 100\]
- Efflorescence is the loss of water of crystallisation to dry air, as with washing soda, Glauber’s salt and Epsom salt. Deliquescence is the absorption of moisture from the air until the substance dissolves in it, as with caustic soda, calcium chloride and ferric chloride. Hygroscopy is the absorption of moisture without forming a solution, as with quicklime, silica gel and concentrated sulphuric acid.
- Drying agents, also called desiccants, remove moisture without reacting chemically. Dehydrating agents remove chemically combined water, taking hydrogen and oxygen in the ratio 2 : 1; concentrated sulphuric acid turns cane sugar into sugar charcoal: \[\\mathrm{C\_{12}H\_{22}O\_{11}} \\xrightarrow{\\ \\mathrm{conc.\\ H\_2SO\_4}\\ } 12,\\mathrm{C} + 11,\\mathrm{H\_2O}\]
- The drying agent must suit the gas: concentrated sulphuric acid dries acidic gases such as hydrogen chloride, while quicklime, being basic, is used for ammonia. Solids are dried in a desiccator over anhydrous calcium chloride.
- Water is soft if it readily forms lather with soap, and hard if it does not. Hardness is due to the bicarbonates, chlorides or sulphates of calcium or magnesium, which the water picks up from limestone, dolomite and gypsum.
- Temporary hardness is caused by the bicarbonates alone and is removed by boiling: \[\\mathrm{Ca(HCO\_3)\_2} \\xrightarrow{\\ \\text{boil}\\ } \\mathrm{CaCO\_3}\\downarrow + \\mathrm{H\_2O} + \\mathrm{CO\_2}\\uparrow\] Permanent hardness is caused by the sulphates and chlorides and is not removed by boiling.
- The same reaction that hardens and softens water builds stalactites, which grow down from a cave roof, and stalagmites, which grow up from its floor.
- Hard water gives water a pleasant taste, supplies calcium and magnesium for bones and teeth, and checks lead poisoning from lead pipes. Against this it wastes soap, dulls clothes, furs kettles, and forms boiler scale that lowers efficiency and can burst a boiler.
- Soap, which is sodium stearate, reacts with the calcium and magnesium ions of hard water to give an insoluble scum called soap-curd, and no lather forms until those ions are used up: \[2,\\mathrm{NaSt} + \\mathrm{CaSO\_4} \\rightarrow \\mathrm{CaSt\_2}\\downarrow + \\mathrm{Na\_2SO\_4}\] Synthetic detergents are unaffected by hardness, since their calcium and magnesium salts are soluble.
- Temporary hardness is removed by boiling, by slaked lime in Clark’s process, or by washing soda. Permanent hardness is removed by a calculated quantity of soda ash, with lime added as well if magnesium sulphate is present.
- Where both kinds of hardness are present and large volumes are involved, the permutit process is used. Permutit, an artificial zeolite, exchanges its sodium ions for the calcium and magnesium ions of the water, and is regenerated with brine: \[\\mathrm{CaP} + 2,\\mathrm{NaCl} \\rightarrow \\mathrm{Na\_2P} + \\mathrm{CaCl\_2}\] For scientific work, cation and anion exchange resins remove all the ions and leave the water completely deionised.
- Potable water is water safe for human consumption. It must be clear, colourless, odourless and pleasant to taste, with pH between 6.5 and 8.5, and free from harmful microorganisms and from toxic substances such as arsenic and lead.
- Water is made potable by coagulation with alum and settling, filtration through sand, gravel and coal, chlorination to destroy bacteria, defluoridation where fluoride is above the safe limit, and aeration to improve taste and odour.
- Water is contaminated by sewage, animal waste, industrial effluents, pesticides and fertilisers. Polluted water spreads diarrhoea, typhoid, cholera and dysentery, and the viral diseases hepatitis A and polio; chemically contaminated water taken over long periods may raise the risk of cancer and cause organ damage.
3.19 Key Terms
- Aeration — passing air through water, or spraying it, to improve its taste and odour.
- Anhydrous — containing no water; used of a salt from which the water of crystallisation has been removed, or which never had any.
- Anomalous expansion of water — the expansion of water on cooling below 4 °C, instead of the contraction shown by ordinary liquids.
- Aqueous solution — a solution in which the solvent is water.
- Binary solution — a solution made up of two components, one solute and one solvent.
- Boiler scale — the hard crust of deposited salts that forms inside boiler tubes when hard water is used.
- Brownian motion — the continuous zig-zag movement of colloidal particles, caused by the moving molecules of the medium striking them.
- Centrifugation — a method of separation, using rapid spinning, that can bring down colloidal particles.
- Chlorination — adding a calculated amount of chlorine to water to destroy disease-causing bacteria.
- Coagulant — a chemical, commonly alum, added to raw water to make its suspended impurities clump together and settle.
- Coagulation — the settling of colloidal particles after their charge has been neutralised by an oppositely charged substance.
- Colloidal solution — a heterogeneous mixture whose dispersed particles, of size 1 nm to 1000 nm, neither settle nor can be filtered out.
- Combined state — the state in which a substance occurs as part of another substance rather than by itself.
- Concentrated solution — a solution containing a relatively large amount of solute.
- Concentration — the amount of solute dissolved in a given quantity of solution.
- Crystal — a homogeneous solid of definite geometrical shape, with smooth plane faces meeting in sharp edges.
- Crystallisation — the process of obtaining crystals of a substance by cooling its hot saturated solution.
- Decrepitation — the crackling sound made by certain crystals, such as potassium chloride, when they break up on heating.
- Defluoridation — the removal of excess fluoride from drinking water to bring it within the safe range.
- Dehydrating agent — a substance that removes even chemically combined water from a compound, taking hydrogen and oxygen in the ratio 2 : 1.
- Deliquescence — the absorption of moisture from the air by a soluble substance until it dissolves in the water absorbed, forming a saturated solution.
- Desiccator — an air-tight glass vessel holding a drying agent, used for drying solids kept inside it.
- Dielectric constant — a measure of how far a liquid weakens the force of attraction between oppositely charged ions placed in it.
- Dilute solution — a solution containing a relatively small amount of solute.
- Dispersed phase — the component of a colloid present in the smaller amount, scattered through the other.
- Dispersion medium — the component of a colloid present in the larger amount, through which the dispersed phase is scattered.
- Drying agent — a substance that absorbs moisture from other substances without reacting chemically with them; also called a desiccant.
- Efflorescence — the loss of water of crystallisation by a hydrated salt on exposure to dry air, leaving a powder.
- Emulsion — a colloid in which both the dispersed phase and the dispersion medium are liquids.
- Floc — a tiny sticky mass formed when alum is added to water, to which mud and bacteria stick before settling.
- Free state — the state in which a substance occurs as itself and not combined with anything else.
- Hard water — water that does not readily form lather with soap, because of dissolved calcium and magnesium salts.
- Henry’s law — at a given temperature, the mass of a gas dissolved by a fixed volume of a liquid is directly proportional to the pressure on the surface of the liquid.
- Heterogeneous mixture — a mixture whose composition is not uniform and which shows visible boundaries of separation.
- Homogeneous mixture — a mixture whose composition is uniform throughout.
- Hydrated salt — a salt whose crystals contain a fixed amount of water of crystallisation.
- Hygroscopy — the absorption of moisture from the air by a substance, without enough being taken up to form a solution.
- Latent heat of fusion — the heat needed to change a substance from solid to liquid at its melting point, with no change of temperature.
- Latent heat of vaporization — the heat needed to change a liquid to vapour at its boiling point, with no change of temperature.
- Mass percent — the mass of solute in grams present in 100 grams of solution.
- Mixture — two or more substances physically mixed in any proportion and not chemically combined.
- Molecular mass — the sum of the atomic masses of all the atoms shown in the formula of a substance.
- Non-aqueous solution — a solution in which the solvent is a liquid other than water.
- Permanent hard water — water made hard by the sulphates and chlorides of calcium and magnesium, whose hardness is not removed by boiling.
- Permutit — an artificial zeolite that softens hard water by exchanging its sodium ions for calcium and magnesium ions.
- Potable water — water safe enough for human consumption without risk of immediate or long-term illness.
- Saturated solution — a solution that can dissolve no more of the solute at a given temperature.
- Seed crystal — a small crystal suspended in a saturated solution so that the solute builds on to it and grows it larger.
- Seeding — inducing crystallisation by adding a crystal of the pure substance to its saturated solution.
- Soap-curd — the insoluble scum of calcium and magnesium salts formed when soap is used in hard water.
- Soft water — water that readily forms lather with soap.
- Solid solution — a homogeneous solution of one solid in another, such as an alloy.
- Solubility — the mass of solute in grams that saturates 100 g of a solvent at a stated temperature.
- Solubility curve — a line graph showing how the solubility of a substance changes with temperature.
- Solute — the substance that dissolves in a solvent to form a solution.
- Solution — a homogeneous mixture of two or more components whose composition can be gradually changed.
- Solvent — the substance that acts as the medium in which other substances dissolve.
- Specific heat capacity — the fixed amount of heat absorbed by 1 g of a substance when it is heated through 1 °C.
- Stalactite — a conical deposit of calcium carbonate growing downwards from the roof of a limestone cave.
- Stalagmite — a conical deposit of calcium carbonate growing upwards from the floor of a limestone cave.
- Supersaturated solution — a solution holding more solute than its saturated solution would hold at that temperature.
- Suspension — a heterogeneous mixture whose solid particles, larger than 1000 nm, do not dissolve and settle out on standing.
- Temporary hard water — water made hard by the hydrogen carbonates of calcium and magnesium alone, whose hardness is removed by boiling.
- Tincture — a solution in which the solvent is alcohol.
- True solution — a solution whose solute particles are smaller than 1 nm and which does not scatter light.
- Tyndall effect — the scattering of a beam of light by colloidal particles, which makes the path of the beam visible.
- Universal solvent — a name given to water because it dissolves more substances than any other liquid found in nature.
- Unsaturated solution — a solution in which more solute can still be dissolved at that temperature.
- Volume percent — the volume of solute in millilitres present in 100 mL of solution.
- Water cycle — the continuous circulation of water between the Earth’s surface and the atmosphere by evaporation, transpiration, condensation and rainfall.
- Water of crystallisation — the fixed amount of water that forms an integral part of the crystalline structure of certain salts.
- Water-borne diseases — illnesses spread by polluted water, including diarrhoea, typhoid, cholera, dysentery, hepatitis A and polio.
3.20 Points To Ponder
Ice floats because water expands as it freezes, and a pond therefore freezes downwards from its surface. Suppose water behaved like most other liquids and contracted steadily until it solidified. Which of the properties of water studied in this chapter would still hold, and what would happen each winter to a pond, a river and the creatures living in them?
The very property that makes water essential to life — its ability to dissolve almost anything — is also what allows a factory’s waste to travel a hundred kilometres downstream. Is it fair to describe the same property as both an advantage and a danger, or does the difference lie entirely in what we put into the water?
For potassium nitrate, solubility rises steeply with temperature; for calcium sulphate it falls. Both are solids dissolving in the same solvent. What must be different about the two dissolutions, and how would you set about deciding, from measurements alone, which of the two takes in heat and which gives it out?
Hard water furs kettles, wastes soap and can burst a boiler, yet it also carries calcium and magnesium that our bones and teeth need, and it gives drinking water its taste. If you were advising a town on how far to soften its supply, what would you want to know before deciding, and would you treat the water going to the factories and the water going to the homes in the same way?
Distilled water is chemically pure but is not potable; treated tap water is potable but is far from pure. Does chemistry use the word ‘pure’ in the same sense as everyday speech does? Where else in this chapter has a word carried a stricter meaning than its ordinary one?
Synthetic detergents were introduced because their calcium and magnesium salts are soluble, so they lather even in hard water. Before calling a new substance an improvement on an old one, what questions besides ‘does it work better?’ ought to be asked of it — and how would you go about answering them for detergents?
3.21 Objective Questions
Multiple Choice Questions
Choose the correct answer from the options given below.
- The particle size in a colloidal solution lies between: [1]
- less than 1 nm
- 1 nm and 1000 nm
- 1000 nm and 5000 nm
- more than 5000 nm
- Permanent hardness of water is caused by the presence of: [1]
- calcium bicarbonate
- magnesium bicarbonate
- magnesium sulphate
- sodium chloride
- The specific value of the latent heat of fusion of ice is: [1]
- 80 J/g
- 336 J/g
- 540 J/g
- 2268 J/g
- The solubility of a gas in water is increased by: [1]
- raising the temperature of the water
- increasing the pressure on the surface of the water
- stirring the solution
- powdering the solute
- Quicklime is chosen in preference to concentrated sulphuric acid for drying ammonia because: [1]
- quicklime is acidic and ammonia is basic
- both quicklime and ammonia are basic, so they do not react
- ammonia is insoluble in concentrated sulphuric acid
- quicklime is a dehydrating agent
- Which of the following salts contains no water of crystallisation? [1]
- Epsom salt
- Green vitriol
- Potassium permanganate
- Gypsum
- A colloidal solution can be separated from its dispersion medium by: [1]
- ordinary filtration
- decantation
- centrifugation
- leaving it undisturbed overnight
- The particle size in a colloidal solution lies between: [1]
Assertion–Reason Questions
- For each of the following, choose the answer based on the codes given below.
- Both A and R are true and R is the correct explanation of A.
- Both A and R are true but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
Assertion (A): A hot saturated solution of potassium nitrate deposits crystals as it cools. Reason (R): The solubility of potassium nitrate decreases with a fall in temperature. [1]
Assertion (A): Crystals of washing soda left in dry air crumble into a powder. Reason (R): Washing soda absorbs moisture from the air and dissolves in the water absorbed. [1]
Assertion (A): A beam of light passing through milk is visible along its path, but a beam passing through sugar solution is not. Reason (R): The solute particles of a true solution are too small to scatter light. [1]
Assertion (A): Concentrated sulphuric acid turns cane sugar into a black mass. Reason (R): Cane sugar contains water molecules held as water of crystallisation. [1]
Match The Following
- Match each substance in Column A with the phenomenon it shows in Column B. [4]
| Column A | Column B |
|---|---|
| a. Glauber’s salt | i. Deliquescence |
| b. Ferric chloride | ii. Hygroscopy |
| c. Silica gel | iii. Efflorescence |
| d. Anhydrous copper sulphate | iv. Turns blue on contact with water |
Fill In The Blanks
- Fill in each blank with a suitable word or number. [4]
- Water has its maximum density at ______ °C.
- A solution in which the solvent is alcohol is called a ______.
- The scattering of a beam of light by colloidal particles is called the ______ effect.
- Hardness caused by the ______ of calcium and magnesium can be removed by boiling.
True Or False
- State whether each of the following is true or false. [4]
- A saturated solution is always a concentrated solution.
- Rain water leaves concentric rings of solid when evaporated on a watch glass.
- Detergents form a scum with the calcium salts of hard water.
- Distilled water is soft water.
Name The Following
- Name the following. [5]
- The gas dissolved in pond water that aquatic animals use for respiration.
- The salt added to hard water to remove both temporary and permanent hardness.
- The artificial zeolite used for softening large volumes of hard water.
- The coagulant commonly added to raw water in a treatment plant.
- The type of colloid to which fog belongs.
Give One Word Or Phrase
- Give one word or phrase for each of the following. [4]
- A solution that holds more solute than its saturated solution at the same temperature.
- A substance that removes chemically combined water from a compound.
- The crackling sound made by certain crystals on heating.
- Water that is safe for human consumption without risk of illness.
Odd One Out
- Pick the odd one out in each group and give a reason for your choice. [3]
- Caustic soda, calcium chloride, potassium nitrate, ferric chloride.
- Milk, fog, sugar solution, butter.
- Calcium bicarbonate, magnesium bicarbonate, calcium sulphate, magnesium bicarbonate solution obtained from dolomite.
3.22 Check Your Concepts
Define solubility, and state why the temperature must always be mentioned along with it. [2]
Give a reason why boiled water tastes flat. [2]
State what is meant by the water of crystallisation of a salt, and give one example of a salt that contains it. [2]
Give a reason why a bottle of concentrated sulphuric acid must be kept tightly stoppered. [2]
Differentiate between a drying agent and a dehydrating agent, stating clearly the basis of the difference. [3]
Explain, with an example, why the same soap lathers freely in one sample of water and hardly at all in another. [3]
Why does a pond freeze from its surface downwards and not from its bed upwards? Explain using the properties of water. [3]
Differentiate between deliquescence and hygroscopy, stating the basis of the difference and giving one example of each. [3]
What would happen if a supersaturated solution were shaken, or a single crystal of the solute dropped into it? Explain why. [3]
Explain why hard water is unsuitable for use in a boiler. In your answer, state what forms inside the boiler tubes and give two separate consequences of its formation. [4]
Differentiate between a true solution, a colloidal solution and a suspension on the basis of (i) particle size, (ii) separation by ordinary filtration, and (iii) the Tyndall effect. [4]
Water that is chemically pure is not necessarily potable water, and potable water is not chemically pure. Explain both halves of this statement. [4]
Describe how you would prepare a large, well-formed crystal of copper sulphate in the laboratory, and state two precautions you would take and the reason for each. [5]
Describe the steps by which raw water from a river is made potable in a treatment plant, stating the purpose of each step. [5]
3.23 Equations, Conversions & Identification
Write the balanced chemical equation for the action of heat on a solution of magnesium bicarbonate. [2]
Complete and balance each of the following equations. [3]
- \(\\mathrm{Ca(HCO\_3)\_2} + \\mathrm{Ca(OH)\_2} \\rightarrow\)
- \(\\mathrm{CaCl\_2} + \\mathrm{Na\_2CO\_3} \\rightarrow\)
- \(\\mathrm{MgSO\_4} + \\mathrm{Na\_2CO\_3} + \\mathrm{Ca(OH)\_2} \\rightarrow\)
Crystals of sodium nitrate are heated strongly in a dry test tube. Name the gas evolved, state the test by which you would identify it, and write the balanced equation for the reaction. [3]
State what you would observe in each of the following cases. [4]
- Blue crystals of copper sulphate are gently heated in a dry test tube.
- A few drops of water are added to the residue obtained in (a).
- Carbon dioxide is bubbled through lime water for a short while.
- A pellet of caustic soda is left on a watch glass in the open for a day.
How would you obtain the following? Give the balanced equation in each case where a reaction takes place. [4]
- Soft water from water containing calcium sulphate.
- Anhydrous copper sulphate from blue vitriol, without using heat.
- Sodium permutit from calcium permutit.
- Calcium bicarbonate from limestone.
Identify the substance in each of the following cases and give a reason for your answer. [3]
- A white crystalline solid, left in an open dish, gradually becomes moist and finally forms a clear solution.
- A white powder turns blue when a few drops of a colourless liquid are added to it.
- A blue-green crystalline salt whose chemical name is iron(II) sulphate heptahydrate.
Study Fig. 3.14-2 and answer the following. [3]
- Name the apparatus shown.
- Name the substance placed at the bottom of it and state its purpose.
- State why the vessel must be air-tight.
Study Fig. 3.3-3 and answer the following. [3]
- State what the experiment is intended to show.
- Name the method by which the gas is collected in the graduated tube.
- Explain why the flask is filled completely with water at the start.
3.24 Practice With Numericals
25 g of a salt is dissolved in 475 g of water. Calculate the mass percent of the solution. [2]
45 mL of a liquid solute is present in 300 mL of its solution. Calculate the volume percent of the solution. [2]
40 g of a saturated solution of a salt, when evaporated to dryness, leaves a residue of 8 g. Calculate the solubility of the salt at that temperature. [3]
12.0 g of a hydrated salt is heated to constant weight and leaves 7.2 g of anhydrous residue. Calculate the percentage of water of crystallisation in the salt. [3]
The solubility of a salt at 30 °C is 36 g. Calculate the mass of the salt required to prepare a saturated solution with 250 g of water at this temperature. [3]
Using the values in Table 3.6, calculate the mass of sodium nitrate that would crystallise out when a solution saturated at 60 °C, containing 50 g of water, is cooled to 40 °C. [3]