SOLUTIONS AND SOLUBILITY–BEHAVIOUR OF WATER

chapter

25

SOLUTIONS A N D SOLUBILITY—BEHAVIOUR OF W A T E R

S O L U T I O N S A N D SOLUBILITY You have seen that when a solid is shaken with a liquid it may disappear and distribute itself uniformly throughout the liquid to form a mixture which is called a solution. The same thing can often be observed when a gas is shaken with a liquid or when two liquids are shaken together. This process of solution is so important and widespread that we should learn more about it. In this section we will consider in more detail the changes involved when substances dissolve and explain them in terms of the forces acting between particles of solute and solvent. The properties which all solutions have in common include:— • The solution is homogeneous—the solute and solvent are so thoroughly mixed that all parts of the mixture are the same. • The solute and solvent do not separate when allowed to stand in a sealed container. • The individual particles of the dissolved substances are not visible. A saturated solution at a given temperature is one which will not dissolve any more solute when the solution is in contact with solid solute. The solubility of a substance at a given temperature is the amount of it required to produce a saturated solution in a given amount of the solvent at that temperature. Solubilities are often expressed as: y grams of solute in 100 grams of solvent at x °C, for instance, the solubility of sodium chloride is 34 grams in 100 grams of water at 20 °C. Whenever a substance is described as soluble or insoluble in a solvent it is taken for granted, unless otherwise stated, that the temperature is about room temperature.

Solutions of solids and liquids. In Chapter 16, we saw that acids, alkalis and sugars are soluble in water, that fats and oils are insoluble in water but they are soluble in carbon tetrachloride. It is difficult sometimes to tell whether any of a solid has dissolved in a liquid. If the solid disappears, the substance is obviously soluble. Experiment 25.1 illustrates one method of determining whether solids, which have only slight solubility in water, have dissolved. Experiment 25.1. Shake a small quantity of powdered blackboard chalk or slaked lime with a small quantity of distilled water in a test tube. Do you think any of the solid dissolved in the water? Filter the mixture and collect the filtrate in a clean basin. Now carefully evaporate the water by heating the basin over a beaker of boiling water and examine the inside of the basin. If any residue is seen what does this mean? It was not possible to say whether any of the solid had dissolved just by looking at the mixture. The presence of a residue after evaporation of the filtrate means that some of the solid did dissolve. Why was it necessary to use distilled water in this experiment? The following experiment will help you to answer this question. Experiment 25.2. Place about 20 ml each of tap water and distilled water in thoroughly clean basins and carefully evaporate each to dryness. Examine the inside of each basin. The residue in the basin which contained tap water indicates that it contained dissolved solids. We have already seen that some salts like sodium chloride, potassium nitrate and calcium chloride are soluble in water and that others like 25—1

calcium carbonate and barium carbonate are insoluble. Let us examine some further salts and see which are soluble in water and which are insoluble. Experiment 25.3. Shake a small quantity, about the size of a pea, of each of the salts listed below separately with equal volumes of water— about 10 ml. If the salt dissolves, note whether there is any change in the temperature of the mixture. If all the salt does not dissolve in the water, filter the mixture and evaporate the filtrate to dryness in a basin over a beaker of boiling water. Classify each salt as soluble, slightly soluble or insoluble:— Ammonium chloride, sodium acetate, sodium sulphate, sodium carbonate, barium chloride, barium nitrate, barium sulphate, copper nitrate, copper sulphate, copper carbonate, lead chloride, lead nitrate, lead sulphate, lead carbonate, calcium nitrate, calcium sulphate. It is not possible for you to test all salts to determine their solubility in water. From experiments like the one you have performed, we can build up a summary of the solubility of the more common salts in water. We will call this summary the solubility rules:— • All sodium, potassium and salts are soluble.

ammonium

• All nitrates are soluble. • All acetates are soluble. • All chlorides are soluble except silver chloride and lead chloride—lead chloride is slightly soluble in cold water and is more soluble in hot water. • All sulphates are soluble except lead sulphate and barium sulphate—calcium sulphate is only slightly soluble. • All carbonates are insoluble except those of sodium, potassium and ammonium. Solutions of gases in liquids. Do not think that solutions always consist of a solid dissolved in a liquid. If you gently warm some tap water in a beaker, you will notice bubbles forming on the inside of the beaker. These bubbles must have come from gases dissolved in the water. The gases would be those with which the water has come in contact—gases from the air. When a bottle of soft drink is opened, bubbles of the dissolved gas, carbon dioxide, are seen leaving the liquid. If the soft drink is warmed, more bubbles of the gas leave the liquid because in general, gases are less soluble in hot liquids than in cold liquids.

25—2

Gases are more soluble in liquids when the pressure on the system is increased. When the top is removed from a bottle of soft drink, the pressure on the solution is reduced to atmospheric pressure and the carbon dioxide is less soluble in the liquid than it was at the higher pressure under which the soft drink was placed in the bottle. When deep sea divers work at great depths more gas dissolves in their blood than does at atmospheric pressure. If they return to the surface too quickly, bubbles of gas form in their blood and these cause a serious condition known as "the bends". Nitrogen is very much more soluble in blood at high pressures but the solubility of helium is not nearly so greatly affected by increases in pressure. The risk to divers can be decreased, therefore, by pumping a mixture of oxygen and helium instead of air to them to breathe. Solutions of liquids in liquids. We have found that some solids and gases are soluble in liquids, but the solubility is, in general, limited—a given amount of the liquid will dissolve only a limited amount of the solid or gas. Is the same true of liquids dissolved in liquids? Experiment 25 A. Place about 10 ml of water in each of three test tubes and add to the separate tubes 1 ml of each of the following liquids: alcohol, glycerine and kerosene. Shake each tube. Now add 1 ml of water to separate tubes containing 10 ml of alcohol, glycerine and kerosene. Your results will show that alcohol and glycerine dissolve in water and that water dissolves in alcohol and glycerine. Liquids which mix in all proportions are said to be miscible. The results of Experiment 25.4 indicated also that kerosene and water are not soluble in one another. Liquids which do not dissolve in each other are said to be immiscible. You may have noticed a temporary cloudiness when you shook kerosene in water. The cloudy mixture is known as an emulsion. We shall consider emulsions again later in this chapter. Change in solubility with temperature. We have seen how gases decrease in solubility in liquids as the temperature is increased. We saw in Chapter 5 that the solubilities of sodium nitrate and sodium chloride in water are greater at higher temperatures. Are all solids more soluble in hot water than in cold water? Figures 25.1 and 25.2 show graphs of the solubilities of some solids in water at different temperatures. If you study these graphs carefully you will see that— • Not all solids increase in solubility when the temperature is raised.

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powder. Place both samples into equal quantities of water in separate test tubes and shake. Compare the rates at which the different samples dissolve.

SOLUBILITY (g OF S O L I D IN 100g OF W A T E R )

2. Place equal quantities of sugar into separate equal quantities of water in two test tubes. Shake one tube and leave the other to stand. Compare the rates at which the samples dissolve.

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3. Place equal quantities of potassium nitrate— saltpetre—in equal quantities of water in two test tubes. Shake both test tubes, holding one of them over a flame. Compare the rates at which the samples dissolve.

SODIUM C H L O R I D E

The results of your experiments should confirm that, as a general rule, the rate at which a solid will dissolve in water can be increased in three ways:— • by grinding the solid into a finely divided condition; • by shaking the solution while the solid is dissolving; • by warming the solution. Keeping in mind the fact that a solid can only dissolve at those places where it is touched by solvent you should be able to explain why a solid dissolves more quickly when powdered.

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TEMPERATURE (°C) F i g u r e 25.1 T h e v a r i a t i o n w i t h t e m p e r a t u r e of t h e s o l u b i l i t i e s of s o d i u m n i t r a t e , p o t a s s i u m n i t r a t e , p o t a s s i u m c h l o r i d e a n d s o d i u m c h l o r i d e in w a t e r

• Of those which do increase in solubility, the change in solubility with temperature is not the same in each case. • The change in solubility of a substance with temperature is gradual—there are no sudden changes in direction in the curve. • If equal volumes of hot saturated solutions of potassium nitrate and potassium chloride are cooled to room temperature, a larger quantity of crystals will separate from the solution of potassium nitrate than from the solution of potassium chloride. Rates of solution Experiment 25.5. 1. Take two equal quantities of large crystals of copper sulphate. Grind one quantity into a fine

Solutions containing several solutes. In Chapter 5 we saw that it is not possible to keep dissolving a solid in a given amount of a liquid indefinitely— the solution becomes saturated with the solute. If a solution is saturated with one solute, will it still dissolve another different solid?

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SOLUBILITY (g OF S O L I D IN 100g OF W A T E R )

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I 0

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TEMPERATURE

]— 60

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(°C)

F i g u r e 25.2 T h e v a r i a t i o n w i t h t e m p e r a t u r e of t h e s o l u b i l i t i e s of c a l c i u m h y d r o x i d e a n d c a l c i u m s u l p h a t e in w a t e r

25—3

Experiment 25.6. 1. Dissolve some sugar in a small quantity of water. Add some sodium chloride to this solution to see whether it will also dissolve. Then drop pieces of potassium dichromate into the solution and shake it. Does the solution change colour indicating that the potassium dichromate is being dissolved ? 2. Make a saturated solution of sugar and again shake small quantities of powdered sodium chloride and potassium dichromate in it. Did any of the solids dissolve? From these tests you can see that the presence of one dissolved substance does not prevent other substances dissolving in the solution. As a general rule, unless the concentrations are high, one solute has little effect on the solubility of others. Volume changes during solution Experiment 25.7. Fill a small, narrow-necked flask with water to a level in the neck and mark this level. Add sodium chloride to the water with continual shaking till no more will dissolve. Observe the new level of the liquid. You will observe that the volume of the solution is only slightly greater than the original volume of the water. Experiment 25.8. Close one end of a glass tube about two feet long. Fill it with water approximately to the half-way mark. Add alcohol carefully till it fills an equal length of the tube above the water. Mark the level. Now shake the tube so that the liquids mix and note the new level of the solution.

Heat is not always released when materials dissolve in a liquid. Sodium and potassium hydroxides dissolve in water with a considerable increase in temperature of the mixture. Ammonium chloride, on the other hand, absorbs heat from its surroundings when it dissolves in water—the temperature of the mixture is lowered. Choosing a solvent Experiment 25.10. Mix the following pairs of substances in small quantities and observe whether a solution is formed. Many other pairs could be tested:— • Sodium chloride and kerosene. • Olive oil and water. • Petrol and water. • Petrol and olive oil. • Petrol and kerosene. • Methyl alcohol and copper sulphate crystals. • Ethyl alcohol and copper sulphate crystals. • Kerosene and petroleum jelly. Similar tests with various solvents will confirm the general rules that water is the best solvent for acids, alkalis and salts and that other substances which do not mix with water often dissolve one another. In Experiment 25.10, solutions are formed only between petrol and olive oil, petrol and kerosene, methyl alcohol and copper sulphate, and kerosene and petroleum jelly. Since water dissolves so many substances it has sometimes been called "the universal solvent". This is misleading because there are many more substances which do not dissolve in water than there are which do dissolve in it.

A slight decrease in total volume will be observed when the liquids dissolve.

A N E X P L A N A T I O N OF SOLUTION

Heat of solution. Did you notice that the mixture became warm when the water and alcohol went into solution with each other? You may have noticed bubbles of a gas forming in the solution. This gas is oxygen and it is released because oxygen is not as soluble in the solution as it was in water. You will have noticed too that some of the salts used in Experiment 25.3 caused changes in temperature of the mixture.

It is possible to develop explanations for the observations which you have made about solutions and solubility in terms of the interaction of the solute and solvent particles. Since in all the cases we have considered, it is possible to recover the original solute from the solution, the explanations could apply only to such cases and not to any which involve the formation of new substances. As was described in Chapter 5, the dissolving of a solid may be thought of as a process during which particles of solute are separated from one another and become surrounded by particles of the solvent. In this way a heterogeneous mixture of solute A and solvent B become a homogeneous mixture of particles of A surrounded by particles of B—see Figure 25.3.

Experiment 25.9. Care must be taken when using sodium or potassium hydroxide. Dissolve a small amount of ammonium chloride in a little water. Can you detect any change in the temperature of the liquid? Repeat the procedure with sodium hydroxide or potassium hydroxide. What happens in this case? 25—4

two attractions. The greater between water molecules and the the greater the solubility is likely greater the attraction of the ions the lower the solubility is likely

DISSOLVING

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HETEROGENEOUS -

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NOT U N I F O R M L Y

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MIXED

Figure

25.3

Diagram

HOMOGENEOUS UNIFORMLY MIXED

to represent homogeneous

that

solutions

are

Factors which oppose solution are therefore— • forces of attraction between particles of A; • forces of attraction between particles of B; Factors which help solution are— • forces of attraction between particles of A and particles of B; • the kinetic energy of particles of A and of B. Sodium chloride dissolves in water. We have seen in Chapter 24 that a solid piece of sodium chloride consists of positively charged sodium ions and negatively charged chloride ions arranged in a regular pattern and held together by the attraction between their positive and negative electric charges. In Chapter 5, we indicated that solution occurs when the attractions between particles of the liquid and particles of the solid are strong enough to allow the particles of the solid to move away from their fixed positions and mix with particles of the liquid. In the case of sodium chloride the attractions between particles of solvent and particles of solute result from the attractions between the positive sodium ions and the negative ends of water molecules and between the negative chloride ions and the positive ends of water molecules. When the sodium ions and chloride ions move away from the solid, each has a cluster of water molecules attached to it. The ions are then said to be hydrated. Solubility of salts. Very few substances are completely insoluble in liquids but some dissolve only to a very small extent. For example, calcium carbonate is insoluble in water for all practical purposes. In this case, the attractions of the water molecules for the ions of calcium carbonate are not sufficient to overcome the attractions of the ions for one another. In general, the solubility of a salt is determined by the relative sizes of these

the attractions ions of the salt, to be, while the for one another to be.

Solutions of gases in liquids. The molecules of a gas are, as you have seen, widely separated from one another and are moving rapidly. If the gas is in contact with a liquid, molecules of the gas which strike the surface of the liquid may be attracted to the molecules of the liquid and may enter the liquid to form a solution. When the gas is dissolved in the liquid, the molecules of the gas behave in the same kind of way as molecules of a liquid: they are still moving and, although they are widely separated from other similar molecules in dilute solutions, they are in contact with molecules of the liquid. The solubility of different gases in a liquid at the same temperature and pressure depends on the different attraction of the molecules of the liquid for the molecules of the different gases. The greater these attractions, the more soluble the gas will be. The higher the kinetic energy of the molecules of the gas the less likely they are to be held in the solution. Thus, we find that the solubility of gases in liquids decreases as the temperature increases. The molecules of the gas can only enter the solution when they strike the surface of the liquid. In a gas at higher pressures, there are more molecules in a given volume and, in any given time, there will be more molecules striking the surface of the liquid. Thus gases are more soluble in liquids at high pressures than at low pressures. Change in solubility with temperature. The higher the temperature of a salt the greater is the kinetic energy of its ions. It should, therefore, be easier to move the ions from the positions they occupy in the solid and allow them to mix with the liquid at higher temperatures than at lower temperatures. Thus, we would expect the solubility of a salt to be greater at higher temperatures than at lower temperatures. This is found to be so for most salts; however, because some salts are less soluble at higher temperatures than at lower temperatures, the kinetic energy of the ions cannot be the only factor of importance in the process. Rates of solution. When a salt or other solid is broken into fine pieces, its surface area is very much greater than when it is in larger pieces. When you cut an apple into halves, the surface area of the apple is increased because the large 25—5

new surfaces along the cut are exposed. A powdered salt will dissolve more rapidly than larger pieces of the salt because the larger surface area of the powder allows more ions to come in contact with water molecules which can attract these ions and remove them from the solid. Heating the mixture increases the rate at which a salt or other solid dissolves because, as we discussed above, the increased kinetic energy of the particles of the solid makes it easier to remove them from their positions in the solid. Shaking or stirring increases the rate of solution by allowing better contact between solid and liquid. It also removes the saturated layer which may form near the surface of the solid and so allows more water molecules to move to the surface of the solid. Kerosene will not dissolve common salt but it will dissolve petroleum jelly. The molecules in kerosene and petroleum jelly—are both mixtures of various hydrocarbons, which attract one another strongly and therefore one will dissolve in the other. Water does not attract the molecules of kerosene or petroleum jelly so they do not dissolve in water. Volume changes on solution. You found that there was only a slight increase in volume when water was saturated with sodium chloride. This must mean that the ions of the sodium chloride can almost fit into spaces between the water molecules. The spaces cannot accommodate all the ions and so there is a slight increase in volume. The fact that equal volumes of water and alcohol when mixed together do not occupy twice the volume of either alone must mean that, in this case too, the molecules of one liquid can, at least partly, fit into the spaces between the molecules of the other liquid.

EXOTHERMIC AND E N D O T H E R M I C CHANGES In Chapter 15, we saw that heat is given off when many chemical changes such as combustion take place and that such changes are said to be exothermic. Reactions, in which heat is absorbed during the change, are said to be endothermic. We have just discussed the energy changes involved when solutions are formed; the formation of a solution may be either exothermic or endothermic. The changes which occur in most systems, whether or not new substances are formed, either give out or absorb energy. We saw, in Chapter 16, that the formation of hydrated copper sulphate from anhydrous copper sulphate is an exothermic change and that it was necessary to heat hydrated copper sulphate to 25—6

produce the anhydrous salt—an endothermic change. In general we can say that if a change is exothermic and gives out a certain amount of energy, the reverse change is endothermic and absorbs the same amount of energy. In Chapter 6, we discussed the action of a refrigerator. This actually makes use of an endothermic change to remove heat from the interior of the refrigerator and of an exothermic change to transfer this heat to the surroundings. Both exothermic and endothermic changes are of great importance to us. Exothermic changes are used to supply energy. For instance, the combustion of a fuel is an exothermic reaction and the exothermic reactions involved in respiration supply us with the energy required to keep us warm and do the many things we do. At constant temperature, energy must be supplied to produce endothermic changes. For instance, the evaporation of water, which is an endothermic change, keeps our bodies cool; many of the processes which are used to supply us with useful materials such as iron, copper and chlorine are endothermic—some of these changes are discussed in Chapter 44. We have seen that hydrogen burns in oxygen to produce water—an exothermic reaction—and that water can be converted to hydrogen and oxygen by passing an electric current through it. Thus, an endothermic change has been produced not by supplying energy in the form of heat but by supplying electrical energy. As we mentioned in Chapter 9, the voltaic cell uses a chemical change to produce electrical energy. We have also seen in the previous section that when water molecules attach themselves to ions, heat is produced, thus this also is an exothermic reaction. The energy which is- released during an exothermic change must have been possessed by the reacting substances; the products must have less energy—that is they are more stable—than the reacting substances had. The opposite is true of an endothermic change—the products have more energy—they are less stable—than the reacting substances. This again suggests that a substance must possess "potential" energy—recall Chapter 13. The energy possessed by a substance may be of various kinds. For instance, its particles possess kinetic energy and the substance possesses potential energy due to the forces of attraction between these particles, in the same way as an object possesses potential energy due to gravitational forces between it and the earth. In addition, the particles themselves possess energy which is "locked-up" within the molecules or ions due to the chemical forces operating within the particles.

The energy absorbed or given out in a change may be due to a change in kinetic energy of the particles and/or change in the potential energy due to the forces of attraction between the particles. If the change produces a new substance or substances, different particles are present after the reaction and these particles will have different energy due to the different chemical forces present in them. This difference in energy is either absorbed or given out during the change.

ASBESTOS WOOL

FRESHLY POLISHED METAL I

WET ASBESTOS WOOL

I

WATER A N D ELEMENTS Besides being a solvent for many substances, water reacts chemically with many other substances. Let us examine some of these reactions. Reaction of water with metals. In Chapter 15, the rusting of iron was found to be a reaction which involves water, iron and oxygen. In this chapter we will now consider the interaction of metals with water. These reactions are of particular interest because metals are so widely used. Experiment 25.11. 1. Take a number of test tubes containing small quantities of rain or distilled water. Boil for a few minutes and allow to cool. Now place some small pieces of freshly polished copper, zinc, iron, magnesium and aluminium in the separate test tubes. Leave for 10 minutes. Note any change in the metal or water. Then carefully boil the water for a few minutes. Observe any change. 2. If you notice any bubbles on the metal in any of the tests above, place the metal in a small basin of water and invert a test tube of water over it. Leave for a few days to see whether larger quantities of the gas in the bubbles may be collected. Test any gas collected with litmus, limewater and a lighted splinter. The purpose of boiling the water before placing the metal into it is to remove any dissolved air which might react with the metal.

SMALL "FLAME NOT TOUCHING GLASS

F i g u r e 25.4

—BUNSEN

BURNER

Reaction between metals and

steam

Experiment 25.12. 1. Place a small plug of wet glass wool or asbestos at the bottom of a test tube. Place another small piece of glass wool or asbestos about half-way up the tube. Clean and polish a piece of magnesium ribbon and place it on the upper plug. Insert a rubber stopper carrying a few inches of glass tubing in the mouth of the test tube. Warm the lower glass wool till steam comes off freely and then use a second burner to heat the magnesium ribbon—see Figure 25.4. If any action occurs, place a lighted splinter near the outlet tube. In this experiment you need to observe any reaction between water in the form of steam and magnesium, so wait till the steam is coming off freely before heating the magnesium.

2. Repeat the above procedure using in turn pieces of freshly polished aluminium, copper wire and iron wire. Compare the results with those observed with magnesium. Magnesium, aluminium and iron when heated in steam react with the steam, but copper does not. The reactions may be represented by the equations— Magnesium *> Aluminium s) Iron (

(

(S)

+ + +

Water^ Water Water m

(!7)

—> —> —>

Magnesium oxide Aluminium oxide<«> Iron oxide * (S)

(

}

4- Hydrogen^ + Hydrogen^, + Hydrogen^)

Experiment 25.13. This experiment can be dangerous and should not be attempted by pupils. Support a small length of quarter-inch bore glass tubing vertically with one end about one inch below the surface of some water in a beaker and add a piece of red and a piece of blue litmus paper to the water. The upper part of the tube must be quite dry. Cut a small piece of sodium metal about the size of a grain of rice and drop it down the tube. Observe any change and hold a lighted match near the mouth of the tube. Stir the water in the beaker after removing the tube. Can you identify the substances produced during the reaction ?

25—7

You should have been able to identify the gas as hydrogen and from the change of the red litmus to blue you should know that an alkali was formed in the solution. The alkali which formed is sodium hydroxide. The equation for this reaction is— Sodium^)

+

Watery

—>

Sodium hydroxide

These experiments show that some metals can react chemically with water. Your evidence could be the formation of some obviously new substances, the appearance of deposits or colour changes on metals, the production of light and heat. Remember that steam is still chemically water. Comparing the speed and extent of any observed changes, the ease with which they commence and the degree of heat needed to maintain the reaction, you should be able to draw up a list of metals in order of their vigour in reacting with water. Do your observations allow you to agree with this order, starting with the most active metal and proceeding to the least active:— sodium, magnesium, aluminium, iron and copper? Later on, you will study the action of acids on metals and you should compare these results with those above. Reaction of water with non-metals. We have seen that some metals react with water. Do nonmetals behave in the same way? Experiment 25.14. 1. Shake small quantities of sulphur, red phosphorus, carbon and iodine separately with water. Are there any indications of solution or chemical reaction? Filter each mixture. Test a little of the filtrate from the mixture containing iodine by pouring a little of it on to a piece of starch. Evaporate each filtrate to dryness and examine each container carefully to see if any residue remains. 2. Place pieces of red and blue litmus paper into some water in a test tube through which chlorine has been bubbled. Smell the contents of the tube very carefully by holding the test tube a few inches away from your nose and gently waving your hand over the mouth of the test tube to waft any vapours towards you. The non-metals sulphur, red phosphorus and carbon are insoluble in water. The blue colour produced with starch by the filtrate from the mixture of iodine and water indicates that some iodine dissolved—it is only slightly soluble in water. Chlorine and water together turn blue litmus red and then bleach it. It is of interest to note that although carbon does not react with water at ordinary temperatures, a reaction can be started if the carbon is heated to 1,000 °C and the water is converted to steam 25—8

(fl<7)

+

Hydrogen^)

before it is allowed come into contact with the carbon. This agrees with what you have already learned: namely that chemical reactions can usually be started or accelerated if the reacting mixture is given heat energy. This reaction is used in making water gas which is a useful fuel. It is often mixed with other fuel gases such as coal gas—see Chapter 45. The reaction may be represented by the equation:— Carbonw + W a t e r ^ r ^ ^ ^ ^ ^ H y d r o g e n , , , You have already seen that air is soluble to a small extent in water and air contains amongst other things the two non-metallic elements oxygen and nitrogen. It is important for you to remember that your experiments tell you things about the few substances you tested and nothing else, even though you suspect that other similar substances might react in the same way. In general it is found that non-metals are not very soluble in water and usually do not react with it at ordinary temperatures. Water and acids. As you learned in Chapter 24, pure acids are non-conductors of electricity, but their aqueous solutions are conductors. During the passage of the electric current through these solutions, hydrogen is always produced at the cathode. These facts suggest that water reacts with the molecules of acids, producing ions. Since hydrogen is formed at the negative electrode, it is reasonable to suggest that hydrogen ions are always produced by the action of water on molecules of acids and that they are positively charged ions. Equations to illustrate this process are as follows:— Hydrogen Chloride \om ions ) Sulphate Sulphuric + ions^) m )

+

Wa,e„„^d3

(fl(7

Experiment 25.15. This experiment should not be attempted by pupils. Add concentrated sulphuric acid very slowly to water. Stir the mixture thoroughly each time a small amount of acid is added. Note any change in temperature. Repeat the experiment by bubbling hydrogen chloride gas into water and finally add acetic acid to water.

You should have noticed that acetic acid, a weak acid, produced less heat than the strong acids. When diluting strong acids always add the acid slowly to water with great care. Never add water to the acid. Concentrated sulphuric acid combines so readily with water that it can be used as a dehydrating agent. For instance, it removes the water from hydrated copper sulphate crystals and from other hydrated salts. It reacts with carbohydrates like sugar and cellulose charring them by removing the elements of water from them and leaving a mass of black carbon behind.

SUSPENSIONS A N D EMULSIONS In a solution the particles into which the solute is divided are extremely small, being generally molecules or ions. Hence, the single particles are invisible even under a powerful microscope although they may colour the solution. It is possible for certain substances when shaken with water, to break into particles which contain large numbers of molecules and mix with the water. These particles may be large enough to be seen and can settle on standing or be filtered off. These mixtures, in which large particles containing many molecules of an insoluble substance are suspended in a liquid, are called suspensions— recall Chapter 5. Even though the particles may be relatively small and be practically invisible, these mixtures are not solutions. The smaller the suspended particles, the longer will be the time taken for them to settle when allowed to stand. When droplets of a liquid are suspended in another liquid the mixture is called an emulsion. Emulsions, like other suspensions, are usually cloudy or opaque. The droplets will usually separate on standing, the rate depending on the size of the droplets. Certain substances can be added to the water to keep the droplets suspended for a long time, making what is then called a permanent emulsion—see the discussion on soap in Chapter 27. Substances which can make emulsions permanent—emulsifying agents—are very important in living organisms. Bile salts are able to emulsify fats so that they may be distributed as fat droplets through water solutions containing enzymes and this allows the enzyme to react with the fat. This is discussed in Chapter 28. Salad dressing is an emulsion of olive oil and vinegar, and some garden sprays are emulsions of oils and water. Cream in milk is an emulsion but a rather poor one, for it gradually separates out as a layer on top of the milk. This latter is called skim milk. We shall be discussing forces between particles in an emulsion in Chapter 37.

SURFACE A N D U N D E R G R O U N D WATER Our discussion on the behaviour of water has given us considerable insight into the important role it plays in the many diverse chemical and physical reactions here on earth. So far, however, we have said little about what happens to naturally occurring water which reaches us in the first instance in the form of rain or precipitation—see Chapter 7. Of course, we know some of the obvious things that may happen. We have all seen water running into drains, rivers, swamps, lakes and so forth after heavy rain—recall Chapter 17 on erosion. In many instances, large dams—discussed in Chapter 46—have been and are being built; these allow us to store in reservoirs the large amounts of water flowing in from surrounding catchment areas. As much of this surface water flows rapidly from the catchment areas to the reservoir, it has little opportunity to attack the rocks. Thus, it has a relatively small content of dissolved salts and hence is usually fit for drinking and other household purposes. Underground water. During a rainstorm, water falling on a bitumen road surface quickly runs off into the gutters and drains away under the influence of gravity. Soon after the rain stops the gutters cease to flow and the remaining water soon evaporates. Unlike street gutters, creeks and rivers continue to flow for a long time after rain, often all the year even though rain falls quite infrequently. Many of these flowing streams are fed by springs, that is water flowing out of the ground. Where does this water that flows in streams during dry times or from springs, come from originally ? Does it come from some underground river or reservoir? If so, what forces are at work to bring it to the surface? Actually all this water originally fell as rain. Rain falling on the earth's surface partly runs off over the surface, and may cause soil erosion. As you have read earlier, the rest soaks into the ground and flows slowly through the small spaces between mineral grains in the soil and in the rocks beneath. Water in the top few inches may be evaporated by radiant energy from the sun so that the surface dries out. This is why it is necessary for you to water your garden often during hot dry weather. Water percolating down through the soil and rock finally fills all the open spaces in the r o c k that is, the rock becomes saturated with water. The top of this saturated zone is called the water table—see Figure 25.5. If a hole is dug below the water table, this may eventually fill with water to the height of the top of the saturated zone. 25—9

T H E R E A R E FEW S P A C E S L E F T FOR W A T E R TO F I L L DOWN H E R E

F i g u r e 25.5

A s e c t i o n s h o w i n g t h e p o s i t i o n of t h e w a t e r t a b l e

Unfortunately not all holes dug below the water table fill sufficiently quickly with water so as to provide a useful well with a continuous supply for, say, a farmhouse or farm animals. You will understand why this is so if you put a piece of sandstone, shale, limestone and granite each in a separate container with a little water in the bottom. After a time the sandstone will soak up the water like blotting paper—it is a porous rock. The shale, limestone and granite, however, do not absorb any water—they are impervious rocks. If our well is sunk in porous sandstone below the water table, we would expect a good supply of water. A well sunk in shale or granite would however remain dry. Water obtained from wells is often difficult to wash with, as soap will not lather. The water is said to be hard. The reason for this will be given la*er. In some areas the water table is relatively close to ground surface and abundant supplies of water can easily be obtained. In other places it is deep in the earth's crust and in many places it is difficult to find. From what you have already read, what would you expect to happen if the water table comes to the surface? As in the case of our well, water would seep out at the surface. This outcropping of the water table corresponds to the level of water in a creek, river, lake or swamp. The rate of flow from this ground would depend on the porosity of the rocks. Fractures such as joints may provide a channel so that quite a lot of water may flow from a small area on the surface. This is one way in which springs occur. You can see now that water flowing in streams in dry times is water that has slowly flowed through the rocks from earlier rains. In times of drought 25—10

most of the water contained in the rocks above the stream bed may be drained away so that the water table falls below the bottom of the creek. When this happens the creek will stop flowing and dry up. Artesian water. Would it surprise you to know that water soaking through the rocks may move down thousands of feet into the earth's crust? Proof that it does, can be found in considering Australia's Great Artesian basin. As shown in Figure 25.6, rain falling on the intake beds in the hilly eastern highlands, soaks into porous sandstone underlying vast plains. When the sandstone becomes saturated with water it is known as an aquifer. Pressure of the water from the intake beds in the highlands is sufficient to force the water—known as artesian water—above the surface in some places, when the overlying nonporous beds are penetrated by a bore. If water bores such as these are not under sufficient pressure— see Chapter 23—the water must be pumped, it is called sub-artesian water. A number of geological situations which allow the storage of artesian water occur in Australia and are of great importance because of the shortage of surface water in the dry interior. Figure 25.6 illustrates the kind of geological structure which can give rise to the storage of artesian water. Natural water solutions. We have seen that water falling as rain soaks into the ground and makes its way downward, filling cracks and pore spaces up to the level of the water table. It is not necessarily stagnant but may move slowly through the rocks dissolving the cementing material and causing decomposition of many of the minerals. Being charged with oxygen and carbon dioxide, it attacks the feldspars, converting them to clay and quartz; sodium, potassium and calcium compounds dissolve in the water. The ferromagnesian minerals also decompose, giving rise to

SE

NW M O U N D SPRING

G U L F OF CARPENTARIA

F i g u r e 25.6

A n i d e a l i s e d s e c t i o n f r o m S.E. t o N . W . a c r o s s t h e G r e a t A r t e s i a n B a s i n

iron minerals and soluble magnesium compounds. It is these soluble salts which make water hard— see Chapter 27. When water emerges at the surface in the form of a spring it may deposit portion of the dissolved material on the surface. Some spring waters are valued as drinking water because of their taste and supposed medicinal value. Artesian water, which comes from deep in the earth's crust, is hotter than surface water—see Chapter 6. Having moved for long distances through the rocks it has had the opportunity to TABLE

react with them. Artesian water thus often contains quite high concentrations of dissolved salts. This can be seen by the formation of mound-springs—see Figure 25.6—characteristic of places where it reaches the surface. These mounds consist of the various substances deposited from the cooling of the hot artesian water. Though artesian water, because of its high mineral content, is unfit for agricultural purposes, it is suitable for drinking by animals. This fact makes it possible to move huge stock herds over much of the Australian continent. 25.1

W a t e r and life

Property

Importance to Living Organisms

A "universal" solvent A neutral medium

..

..

..

..

..

Protoplasm is a watery medium which can therefore hold many substances in solution. Life is possible only near neutral point because extremes of acidity and alkalinity interfere with the chemical reactions which take place in protoplasm.

Is able to separate ions in solution

Promotes the chemical changes characteristic of living matter.

Relatively high surface tension

Important in maintaining cell form.

Low viscosity—flows easily

Favours movement; for example, blood, muscle, chloroplasts.

High heat capacity

A buffer to sudden temperature changes.

Rapid heat convection

Distributes heat quickly and uniformly by convection currents.

High latent heat of vaporisation

..

Important for cooling; for example, sweating, transpiration.

Catalytic action

Promotes chemical change.

Lubricating action

Minimises friction.

25-11

T H E IMPORTANCE OF WATER I N LIVING ORGANISMS Among the materials that are essential to life, water has a unique position. It is the chief constituent of protoplasm and most of the important chemical processes of life—see Chapters 28 and 35—take place in water. A glance at Table 25.1 shows that water is a key chemical in the story of life.

OSMOSIS In Chapter 5, you saw a simple experiment which gave an example of a special kind of diffusion called osmosis. You will need to know a little more about this before you make a deeper study of living things and the changes which go on inside them. The important thing to remember is that this kind of diffusion can only occur when two liquids are separated by a differentially permeable membrane—or partition. The cellophane used in Chapter 5 is not a particularly good differentially permeable membrane. Parchment paper or pig's bladder would be better. These membranes look solid but are really porous with very small spaces between their fibres. Experiment 25.17. Repeat Experiment 5.16, with sugar solution in the beaker and water in the thistle funnel. You should notice that the water level now falls in the funnel. Water particles are obviously diffusing between the membrane into the solution of sugar. In osmosis, the diffusion of water molecules is always from pure water to the solution; the result is to dilute the solution. Experiments in which a dilute solution of sugar is separated from a concentrated solution in the apparatus would show that water molecules diffuse from the dilute solution to the concentrated solution making the latter more dilute. No observable changes occur if solutions of equal concentration are used. Water molecules would diffuse through the membrane at equal rates in both directions in this case. Osmotic pressure. If water molecules were diffusing through a differentially permeable membrane it would be possible to exert a pressure on the solution into which they were diffusing and increase this pressure to such a point that it would just prevent further water molecules forcing their way through the membrane. This pressure which, when acting against the water molecules, would prevent osmosis through the membrane, is called the osmotic pressure of the solution. It can be quite large at times—several times larger than atmospheric pressure. 25—12

If there is no opposing pressure then water molecules continue diffusing until they build up a big enough pressure on the solution side of the membrane to prevent further molecules coming through—they build up a pressure equal to the osmotic pressure. Our particle theory suggests the following as a possible explanation. Water molecules and solute molecules—sugar in our case—are in continuous motion. On the pure water side only water molecules are striking the membrane and on the solution side both water molecules and solute molecules are striking it. There are, obviously, more molecules of water striking on the water side than on the solution side. Hence more water molecules must pass through the small spaces from the side where there are the greatest number of collisions. Solute molecules also strike the membrane, but cannot pass through the spaces. Osmosis in living things. Membranes on the outside of cells—Chapter 8—are normally differentially permeable and allow free passage of water molecules and other small molecules and ions through the membrane but restrict or prevent the passage of many dissolved substances— particularly complicated substances with large molecules. Hence water and various necessary dissolved substances can move in or out of cells as required while the more complex substances are kept within the cell. Water diffuses from soil through the walls of the roots of plants by osmosis—see Chapter 35. The ions of many salts are able to pass through the roots as well and so provide the plant cells with materials from which they can produce more complicated substances needed for life to go on. The soil solutions are usually dilute. The following simple tests will show how water passes in or out of living cells making them swell or contract respectively. Remember osmotic diffusion occurs in the direction: dilute solution to more concentrated solution. Experiment 25.18. Carefully remove the shell of an egg by dissolving it in dilute hydrochloric acid—the shell is largely made of calcium carbonate—leaving the egg enclosed in the thin outer skin—a membrane. Now place it in pure water. It will swell because water passes into it by osmosis—the liquid in contact with the inner surface of the membrane is an aqueous solution. 2. Place a similar egg in a concentrated salt solution. It will shrink. Water passes out of the egg solution into the salt solution because the latter is more concentrated.

Experiment 25.19. A demonstration of an osmotic effect is to take a stick of celery, or a stem of flowering stalk of a herbaceous plant like a dandelion. Split the celery or stalk down the centre and place it in pure water. Observe what happens. Can you explain why the stalk curls back the way it does? The epidermal cells have a water-proof covering and do not take up water directly from the external source. The inside cells along the cut surface, however, readily take up water by osmosis and swell and expand. Because the epidermal cells do not expand but those on the other side do, the stalk curls back. If you now place the curled stalk in a concentrated solution of salt or sugar, it will begin to straighten out again. Water has passed out of the cells, across the differentially permeable cell membranes into the external solution. Remember that in all these cases the cells of the specimen contain material in solution, and the overall direction of water movement across the differentially permeable membranes depends upon whether this solution is more concentrated or more dilute than the external liquid. The great importance of this will again be seen in Chapter 35.

SUMMARY In this chapter you have learned— SOLUTION 1. All solutions have the following common properties— • The solution is homogeneous; that is, the composition of any one portion of the solution is the same as that of any other portion of the solution. • Solute and solvent do not separate when allowed to stand in an air-tight container. • The individual particles of the dissolved substance are not visible. 2. A saturated solution at a given temperature is one which will not dissolve any more solute when in contact with the solid solute. SOLUBILITY 1. The solubility of a solid substance at a given temperature is the mass of the substance which can be dissolved in 100 grams of solvent.

2. The solubility of a solid depends on— • The temperature. The solubility of most solids increases with increase in temperature of the solvent. There are exceptions: the solubilities of some solids decrease with rise in temperature of the solvent, for example, calcium sulphate, calcium hydroxide. • The solvent. Salts, alkalis, solid acids and sugar are usually more soluble in water than in other liquids. 3. A solid dissolves more rapidly in water if it is— • In a finely divided—powdered—state. • Shaken while the solid dissolves. • Warmed. 4. The presence of one kind of dissolved solid in a solution does not greatly effect the solubility of another kind of solid in the solution. 5. Solubilities of salts in water vary. Most are soluble, some are slightly soluble and some are insoluble: • All sodium, potassium and ammonium salts are soluble. • All nitrates are soluble. • All acetates are soluble. • All chlorides are soluble with the exception of(a) lead chloride which is slightly soluble in cold water and is soluble in hot water. (b) silver chloride which is insoluble. • All sulphates are soluble with the exception of(a) calcium sulphate and silver sulphate which are slightly soluble. (b) barium sulphate and lead sulphate which are insoluble. % All carbonates are insoluble except sodium, potassium and ammonium carbonates. 6. A suspension is a mixture of a liquid such as water and very fine particles of an insoluble solid. Clay shaken with water forms a suspension. The finer the particles of a suspension the longer they take to settle. 7. The solubility of a gas in a liquid depends on— • The temperature. Gases are more soluble at lower temperatures than at higher temperatures. Warming the solution expels the gas from the solution. • The solvent. Many gases are more soluble in water than in other liquids, e.g., carbon dioxide, ammonia, hydrogen chloride, oxygen and sulphur dioxide. • The pressure. Gases are more soluble at higher pressures than at lower pressures. 25—13

8. Organic compounds such as petroleum jelly, fats and oils do not dissolve in water. They readily dissolve in organic solvents such as kerosene, petrol, carbon tetrachloride and benzene. 9. Liquids may dissolve in liquids to form solutions. • When liquids dissolve in one another they are said to be miscible. • When liquids will not dissolve in one another they are said to be immiscible. • Two immiscible liquids shaken together form an emulsion. The mixture will remain permanently emulsified if an emulsifying agent is added. THEORY OF SOLUTION 1. In any substance forces of attraction exist between its particles: • B dissolves in A if the forces of attraction between particles A and B are stronger than the forces of attraction between particles A and A or B and B. 2. B will not dissolve in A if the forces of attraction between particles of B and B or A and A are stronger than the forces of attraction between particles of A and B. 3. When solutions are formed, volume changes occur as follows: % Solution of a salt in water: The volume of solution obtained is only slightly greater than the volume of the original water. This is because the ions of the salt partly fit into the spaces between the water molecules. • Solution of a liquid such as methylated spirits in water: The volume of the solution obtained is slightly less than the total volume of the two separate liquids. This is because the molecules of methylated spirits partly fit into the spaces between the water molecules. EXOTHERMIC AND CHANGES

ENDOTHERMIC

1. Exothermic changes involve the liberation of energy, usually as heat. There is a loss of potential energy by the reacting substances during the change. The substances produced are more stable than the reacting substances. Combustion is an example of an exothermic reaction.

25—14

2. Endothermic changes require the supply of energy. There is a gain of potential energy by the atoms of elements in an endothermic change. The substances produced are less stable than the reacting substances. Decomposition of mercuric oxide by heat and acidified water by an electric current are examples of endothermic changes. 3. If a change is exothermic and gives out a certain amount of energy, the reverse change is endothermic and absorbs the same amount of energy. WATER AND ELEMENTS 1. Some metals react chemically with water. In each case a metallic hydroxide and hydrogen are formed: • Sodium reacts vigorously with water to form sodium hydroxide and hydrogen. • Magnesium burns in steam. • Aluminium and iron decompose steam when strongly heated. 2. Non-metals, generally, are not very soluble in water and do not react with it at ordinary temperatures. Exceptions are— • Chlorine is soluble in water. m Iodine is very slightly soluble in water. • Carbon at white heat decomposes steam to form carbon monoxide and hydrogen, the gas mixture being known as water gas. SURFACE AND UNDERGROUND WATER 1. The water table is the top of a water-saturated zone of soil and rock. It may be visible in wells or where it comes to the surface in springs, swamps, streams and lakes. 2. Aquifers are porous rocks saturated with water, which are overlaid and underlaid by impervious rocks. They contain water originally absorbed by their intake beds which outcrop in the eastern highlands. The artesian water obtained from these is hot, highly mineralised and hard. It reaches the surface through mound springs or bores. This water is unsuitable for irrigation of crops but is suitable for stock to drink. OSMOSIS I. A differentially permeable membrane is a separating membrane which will allow some molecules to pass through but will not allow others, of larger size, to pass.

2. Osmosis is the diffusion of water from an aqueous solution of low concentration through a differentially permeable membrane into an aqueous solution of high concentration.

3. Living matter depends on osmotic diffusion for movement of water and dissolved material into and out of cells through the differentially permeable cell walls.

Questions 1. What are the essential characteristics of a true solution?

14. Why would it be very dangerous to attempt to store sodium metal either in air or under water? What liquid could be used as a storage substance for sodium?

2. What, in general, is the effect of an increase in temperature upon the solubility of— (a) a solid; (b) a gas? 3. Give (a) a (b) a (c) a

an example solution of a solution of a solution of a

of— gas in water; gas in a gas; liquid in a liquid.

4. Name the solute and the solvent in each of the following solutions:— (a) sugar in water; (b) iodine in alcohol; (c) alcohol in water; (d) dilute sulphuric acid solution. 5. How would you determine whether any of a solid placed in a liquid has dissolved? 6. Why cannot fish live in cold, boiled water? 7. How could you obtain a sample of air— (a) free of oxygen; (b) free of water vapour; (c) free of carbon dioxide? 8. What pieces of evidence can you suggest to show that air is a mixture of gases? 9. Is there any difference in the composition of inhaled and exhaled air? If there is a difference, devise experiments to estimate any difference. 10. Water which has been collected in a galvanised iron tank tastes different from water obtained in a city water supply. Suggest a reason for this difference. 11. Describe the quickest method of making a saturated solution of iron sulphate crystals in water at room temperature. 12. If 60 ml of water and 40 ml of alcohol were mixed would you expect the volume to be 100 ml? Explain your answer in terms of the particle theory. 13. Draw a diagram to represent a positive ion in solution, showing the arrangement of the water molecules around the ion.

15. Why is it usual to use freshly cut or polished metals when investigating their reaction with steam or water? 16. A gas dissolves in water to form a solution which turns litmus red. Must this gas be an acidic oxide? Discuss. 17. When magnesium is burnt in steam, hydrogen gas is produced. What would lead you to expect that the hydrogen came from the steam and not from the magnesium? 18. Would you expect gold to react with water ? Give a reason for your answer. 19. A piece of carbon is burnt in air and the gas produced is dissolved in water; this liquid turns a piece of blue litmus red. Write word equations for the chemical reactions which took place. Do not worry about the effect on litmus as a chemical reaction. 20. A piece of calcium is burnt in air, the product is added to water and shaken; a piece of red litmus when placed in this mixture turns blue. Write word equations for the chemical reactions which take place. Do not worry about the effect on litmus as a chemical reaction. 21. Would you expect to be able to continue dissolving different substances in a fixed volume of water? 22. Can you suggest why soft drink is usually kept in a cold place such as a refrigerator? 23. When a pot plant is watered with salt water, it soon becomes limp. Why is this so? 24. Flowers which have become limp can sometimes be revived by completely submerging them in water. Explain why. 25. Explain what you understand by the term osmosis. How is this related to the apparent improvement in the quality of oysters left in fresh water for a short time? 26. It has been said that, if a match is placed over the top of an opened lemonade bottle, the contents will not go flat. How would you test the truth of this statement? 25—15