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(c) doc bDoc Brown's Chemistry KS4 science GCSE/IGCSE/O Level Chemistry Revision Notes(c) doc b

 Water Cycle - Water as a resource, hard & soft water, colloids, emulsions, salt solubility & water of crystallisation

1. The components of the water cycle are described and explained and the use of water as an important resource, its treatment to make it safe to drink, pollution problems. Colloids (e.g. sol, foam, emulsion) are described with examples. The difference between hard water and soft water is explained and the causes and treatment of hard water * 2. Gas and salt solubility and solubility curves are considered and finally 3. The explanation and calculation of water of crystallisation

1. Water cycle, treatment, pollution, colloids (sol, foam, emulsion) hard/soft water

2. Gas and salt solubility in water and solubility curves

(c) doc b 3. Calculation of water of crystallisation

1. What happens to water on the Earth's Surface?

The water on the Earth's surface is continually being re-cycled. As it falls, rain water contains only dissolved gases but once it reaches the ground water becomes contaminated in various ways.

1a Water Cycle and Resources * 1b Water Treatment-pollution-colloids * 1c Hard & Soft Water


1a The Water Cycle and Water as a Resource

  • Water is the most abundant substance on the surface of our planet and is essential for all life. Water in rivers, lakes and the oceans is evaporated by the heat of the Sun (endothermic). The water vapour formed rises into the atmosphere, cools and forms clouds of condensation (exothermic). Eventually this gives rain and snow 'precipitation' which on melting returns to the rivers, seas and oceans. This is known as the water cycle.
  • Water is an important raw material and has many uses. It is used as a solvent and as a coolant both in the home and in industry. It is used in many important industrial processes including the manufacture of sulphuric acid.
  • Seawater/brine is a valuable resource e.g. large scale evaporation in 'salt pans' (using fuel burning or solar energy) to produce 'sea salt' sodium chloride NaCl, the water also contains lots of other salts including bromides from which the element bromine is extracted.

(c) doc b


1b Water Treatment and pollution - domestic and industrial contexts

  • There are various undesirable materials that need to be removed from water before it is fit for domestic consumption. They include colloidal clay, microscopic organisms, chemicals which cause tastes or odours and Acidic substances:
  • Drinking water is made fit for domestic home consumption by
    • (i) allowing sedimentation to occur, where larger insoluble particles settle out,
    • (ii) passing it through sand filter beds to remove finer solid particles,
    • (iii) treating with chlorine to kill bacteria,
    • (iv) adding small amounts of sulphur dioxide to remove excess toxic chlorine
      • the molecular equation is SO2(aq) + Cl2(aq) + 2H2O(l) ==> 2HCl(aq) + H2SO4(aq)
      • the ionic equation is SO2(aq) + Cl2(aq) + 2H2O(l) ==> 2Cl-(aq) + SO42-(aq) + 4H+(aq)
    • (v) aluminium sulphate is added to coagulate colloidal clay (see colloids below),
    • (vi) carbon slurry absorbs molecules causing 'tastes' and 'odours'.
    • (vii) adding lime slurry to neutralise the water if it is too acid.
  • The use of artificial fertilisers results in many natural waters being contaminated with dissolved nitrate and ammonium ions. Dissolved nitrate ions can have harmful effects on babies and so the levels of nitrate are carefully monitored. Nitrates may be carcinogenic. The ions from this pollution are not easy to remove on a large cost-efficient scale.
  • An ion-exchange filter can remove these and other ions which can cause problems e.g. calcium and magnesium which cause hardness in water and iron compounds (see below).
  • Iron in water is a non-harmful but an aesthetic nuisance impurity:
    • readily soluble iron(II) when exposed to air form rusty brown insoluble iron(III) hydroxide or hydrated iron(III) oxide compounds. These stain yellow/orange/brown washing laundry and white plumbing facilities!
    • The iron(III) ions also form inky black compounds with the tannic acids in tea and giving it a 'metallic' taste.
    • Cooked vegetables turn brown (complex compounds with phenols).
  • Colloidal clay: A colloid consists of one substance (or mixture of substances) very finely dispersed in another substance (or a mixture of substances) without a new true solution forming. So a colloid is a mixture of a dispersed phase and a continuous phase (disperse medium) BUT the dispersed phase is NOT dissolved in the continuous phase.
    • A colloid is NOT a solution, although the colloid particles are not usually seen under a microscope, they are much bigger than molecules, and much bigger than the molecules of the continuous phase (disperse medium e.g. water).
    • In a solution the solvent or solute particles are usually of comparable size and completely mixed at the 'individual particle  level' i.e. completely homogeneous in the same phase.
    • A colloid can be thought of as intermediate between a true solution and a mixture of e.g. a liquid and an insoluble solid. No filtration separation is possible with solutions and filtration is easy and effective with an insoluble solid. Similarly, most colloid particles are too small to be filtered, but separation from truly dissolved substances is possible with a membrane.
    • The colloidal particles of the disperse phase are equivalent to the solute of a solution and the continuous phase is equivalent to the solvent. The mixture is sometimes referred to as the 'colloidal solution'. These descriptors can be somewhat 'blurred' by the intermediate particulate nature of colloidal systems!
    • The particles in a colloid are so small that they remain 'suspended' (the mixture is called a 'suspension') in the disperse medium (e.g. colloidal clay particles in water) with little tendency to settle out. However the colloidal particles are big enough for their surface area properties to be significant (see electrical properties below).
  • Examples of colloids that is the fine dispersion of one substance in another without a new solution forming:
    • A sol is a solid dispersed in a liquid e.g. tiny particles of clay in water. 
    • A foam is a gas dispersed in a liquid e.g. a well shaken soap solution or shaving cream foam.
    • An emulsion is a liquid dispersed or suspended in another liquid ...
      • and is a mixture of two immiscible liquids like oil and water.
      • Emulsions are thicker than either liquid eg the emulsion 'French dressing', is thicker than olive oil or vinegar
      • With time, the two layers settle out, so the less dense oil floats on top of the aqueous/water layer.
      • One way to inhibit the two layers settling out is to use an emulsifier.
        • An emulsifying agent stabilises an emulsion.
        • Two of the most commonly used emulsifiers are lecithin (E322) and the mono- and di-glycerides of fatty acids (E471), and are classified as food additives in the E number system.
        • Egg yolk acts as an emulsifying agent (it contains lecithin).
      • examples of emulsions.
      • (i) milk (aqueous solution + insoluble, but dispersed fats)
      • (ii) French dressing in salads (based on vinegar + olive oil, but these do reform the oil and aqueous layers quite easily which is why they are shaken before use)
      • (iii) Mayonnaise is a mixture of oil, water, emulsifier and other ingredients.
      • (iv) margarines contain emulsifiers to stop the salty water from separating out and mayonnaise also contains an emulsifier to stop the oil and aqueous based components separating out.
    • Emulsifier molecules have a 'water loving'/'oil hating' (hydrophilic) part and a 'water hating'/'oil loving' part (hydrophobic). Therefore they can interact with the different components and keep the different types of molecules dispersed in each other.
      • Diagram A: This diagram represents a true solution where the black dots represent the dissolved individual molecules - they do NOT clump together.
      • Diagram B: This diagram represents an emulsion of oil droplets dispersed in water (oil in water emulsion).
        • Each oil droplet will have millions of oil molecules in it.
        • The oil is the disperse phase and the water is the continuous phase.
        • This is NOT a true solution.
        • Semi-skimmed or full fat milk is like this, droplets of fat (~1-3% oil) are dispersed in water.
        • Single cream (~18% oil), double cream (~50% oil) are oil in water emulsions.
        • Whipped cream and ice cream are oil in water emulsions.
          • Air is whipped or whisked into cream to give it a soft frothy texture to use as a topping.
          • Whipping air into ice cream gives it a softer texture so you can scoop out portions easily.
        • Mayonnaise is an emulsion of sunflower oil or olive oil with vinegar, and these mixtures are used in salad dressings and sauces. A salad dressing coats the salad materials better than either the olive oil or vinegar.
        • Some non-food examples of oil in water emulsions include moisturising creams and other cosmetic lotions.
      • Diagram C: This diagram represents an emulsion of water droplets dispersed in an oil (water in oil emulsion).
        • Each water droplet will have millions of water molecules in it.
        • The water is the disperse phase and the oil is the continuous phase.
        • This is NOT a true solution.
        • Melted margarine is a water in oil emulsion.
      • One of the problems with useful emulsions is that the two main components, the two immiscible liquids, tend to separate out rendering the emulsion useless for its designed purpose.
        • The way round this is to use an emulsifying agent (emulsifier) which inhibits the separation of the emulsion back into two layers.
      • D: This diagram represents the effect of mixing an oil in water emulsion with an emulsifying agent like soap (edible substances like lecithin are used in processed food!).
        • This diagram illustrates the mechanism by which soaps wash oily/greasy clothes or surfaces.
        • The washing process is described and explained below diagram E.
      • Diagrams E1, E2 and S3: Emulsifying molecules like soap have a negative ionic hydrophilic 'head' ('water liking'/'oil hating' end of molecule) and a hydrophobic 'tail' ('water hating'/'oil liking' end of molecule').
        • eg the stearate ion from the soap sodium stearate shown above.
        • When you shake soap with an oily/greasy material (washing clothes or scrubbing a surface), the oil/grease breaks up into tiny droplets or globules. Why? ...
        • The hydrocarbon hydrophobic tail of the soap dissolves in the oil or grease globule and the negative head is on the surface of the globules/droplets.
          • The hydrophobic tail can only interact with oil/grease ie is attracted to oil and grease.
          • The hydrophilic head can only interact with water ie is attracted to water.
          • Two hydrophilic heads cannot interact with each other and tend to repel each other especially if the hydrophilic head carries a negative charge.
          • In effect, the globules of oil/fat get a surface coating of the emulsifier - a general name for these emulsifying molecules is surfactants and includes soaps, detergents and naturally occurring molecules like lecithin found in egg yolk..
        • Because of the negative head of the soap ion on the oil/grease droplet surface, you get repulsion between the globules and an emulsion is formed.
        • The ionic end also strongly interacts with water, again this prevents the oil/grease particles flocking together.
        • So, the oil and grease particles cannot re-clump together to form a separate layer on the clothes or surface being cleaned, and so the emulsion is stabilised.
        • Therefore the oil/grease remains dispersed in the soapy washing water and hence washed away.
        • In other contexts eg food, you use a soap like molecule, but harmless and edible!, to do exactly the same effect, that is, emulsifying the mixture to make a stable emulsion which doesn't separate into two layers.
          • In the food industry emulsifiers are very important for stopping recipe components separating out from emulsions and give processed foods greater stability and longer shelf-life and helps to produce less fatty food and still retain acceptable texture for the consumer. There can be some diet restrictions for some people eg if you are allergic to eggs then any processed food using egg yolk as an emulsifying agent is a no go area! As with any processed food, if you have a sensitive constitution, you must carefully check the ingredients.
          • Incidentally, the emulsifier molecule does not have to be an ionic compound like soap.
          • It can be a non-ionic neutral molecule like lecithin BUT the molecule must have a hydrophilic head that bonds with water and a hydrophobic tail that bonds with oil/grease.
            • The bonds formed are intermolecular bonds (from intermolecular forces of attraction) and NOT chemical bonds like ionic or covalent bonds.
        • Detergents are also emulsifiers, and not just used for washing in the home, they are also used to help disperse oil spilt from tankers into rivers, seas and oceans. Much of the oil spill can be contained by booms and pumped off the surface of the water - but not all unfortunately. Dispersed oil droplets break down (biodegrade) more quickly than large patches of oil, but the process is very slow. Rescued seabirds coated in oil can be washed with detergent to clean them BUT their own natural protective oils are also washed away so their lives are still in danger and the birds need care and rehabilitation.
        • See also ...
    • Colloidal particles may be electrically charged. (Note: So far the discussion has been confined to hydrophobic ('water hating') colloids which do NOT interact strongly with the continuous phase.
      • In contrast 'gels' for example, are hydrophilic ('water liking') colloids, in which the colloid particles are very solvated* and stabilised by the continuous phase). *
      • Solvated means the particle is weakly attracted to layers of surrounding 'solvent' molecules of the dispersal medium e.g. water.
    • Colloidal particles of a sol absorb ions, but not in electrically balanced proportions. Depending on which ion(s) are preferentially absorbed from the water, the net charge on the colloid particle can be positive or negative. The situation is complicated further because the charged colloid particles attract a sheath of oppositely charged ions around them. This is called the electrical double layer effect. This means neighbouring colloid particles have the same 'outer charge' and so are repelled, rather than attracted together. The sol itself is overall electrically neutral like any other solution.
    • Colloids are destroyed when the particles of the disperse phase join together and separate out from the continuous phase. This process is called coagulation. For sols, any disturbance of the double layer can cause coagulation to happen. It can be caused by boiling the sol, the increased random thermal collisions disturb the electrical balance and allows the colloid particles to collect together.
    • Sols are also very sensitive to the presence of ions, so any electrolyte ions present can affect the electrical double layer (the theory is complex but just think of the ions charge as affecting the stability of the double layer). The more highly charged the ion, the greater the electrical field force effect, so the greater its coagulating power. The ions reduce the repulsion between the colloid particles and allow coagulation to occur.
    • Examples of coagulating power:
      • positive cations: Al3+ > Mg2+ > Na+
      • negative anions: [Fe(CN)6]3- > SO42- > Cl- 
      • and this explains why aluminium sulphate Al2(SO4)3 is used to precipitate (coagulate) colloidal clay in water treatment for domestic water supplies.
  • Other onsite references to water pollution:

(c) doc b


1c Hard and Soft Water

  • HARD and SOFT WATER: Many compounds dissolve in water without chemical change but may have a variety of consequences!
    • Water which readily gives a lather with soap (not detergents) is described as soft water.
      • Note: Detergents usually give a good lather with any water.
  • Some of these dissolved substances make the water hard.
    • This means the water does not readily give a good lather with soap and so wastes soap as well as causing a 'scum'! though it does not affect soapless detergents.
    • The 'scum' is due to the formation of insoluble calcium and magnesium salts formed by a reaction between the soap molecules and calcium and magnesium ions.
  • Most hardness is due to water containing dissolved calcium or magnesium compounds.
    • The hard water is formed when natural waters flow over ground or rocks containing calcium or magnesium compounds.
    • e.g. Chalk and limestone, mainly calcium carbonate CaCO3 with some magnesium carbonate too.
    • or Gypsum rock deposits, which are mainly calcium sulphate CaSO4,
    • and magnesium sulphate which was called 'Epsom Salts', formula MgSO4.7H2O, because it crystallised out of evaporated spring water from Epsom on the chalk downs of southern England.
  • Calcium sulphate (slightly soluble) and magnesium sulphate (very soluble) are washed out of rock formations.
  • Insoluble calcium carbonate (in limestone, chalk) and insoluble magnesium carbonate both dissolve in acid rainwater to form soluble hydrogencarbonates
    • e.g. naturally carbonated water (dissolved carbon dioxide makes water acidic so it reacts with the carbonate) ...
    • insoluble calcium carbonate + water + carbon dioxide ==> soluble calcium hydrogencarbonate
    • CaCO3(s) + H2O(l) + CO2(g) ==> Ca(HCO3)2(aq)
  • The simplest test for 'hardness' is to shake the water with an old fashioned 'soapy'* soap. 
    • * The term 'soapy soap' is NOT a joke! e.g. the blocks of 'household' soap based on sodium stearate, sodium palmitate (from palm oil) or sodium oleate (from olive oil).
  • Soft water readily forms a lather with soap but hard water does not.
  • Hard water forms a scum from the dissolved calcium or magnesium compounds.
    • The scum is a precipitate formed from insoluble calcium and magnesium soap salts, instead of a nice frothy lather (see below).
    • Eventually with enough soap, a lather does form, when all the calcium and magnesium ions have been precipitated as a 'scum salt'! However, it does mean a lot of soap is wasted!
  • The amount of hardness in water sample can be estimated by titrating it with soap solution and noting what volume of soap solution is needed to produce a lather.
  • A modern detergent is sometimes called a 'soapless soap', at least when I was a student!, or soapless detergent. Its advantage is that no insoluble salt 'scum' is formed,, because the Ca and Mg salts of it are soluble. So modern detergents e.g. like 'washing up liquids' give a lather with any water which is more acceptable for dish washing.
  • The chemistry of 'scum' formation. Hard water contains dissolved compounds that react with soap to form scum. e.g. with soaps made from the sodium salts of fatty acids, insoluble calcium or magnesium salts of the soap are formed ... 'example of a precipitation reaction' ..
    • calcium sulfate + sodium stearate (a soap) ==>calcium stearate (scum ppt.) + sodium sulfate
    • CaSO4(aq) + 2C17H35COONa(aq) ==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq)
    • or more simply ionically:
      • calcium ion + stearate ion ===> calcium stearate
      • Ca2+(aq) + 2C17H35COO-(aq) ==> (C17H35COO-)2Ca2+(s)
    • A precipitation reaction is generally defined as 'the formation of an insoluble solid on mixing two solutions or a gas bubbled into a solution'.
    • Below are some diagrams of the organic molecules or ions involved
    • Diagram S1: The stearic acid molecule C17H35COOH
    • Diagram S2: The salt sodium stearate C17H35COO-Na+, formed when stearic acid is neutralised with sodium hydroxide
    • Diagrams S3 and E2: The negative stearate anion C17H35COO-, its structure is important in understanding how it forms the calcium salt precipitate, calcium stearate AND explaining how emulsifiers work.
  • Using hard water can increase costs because more soap is needed to make a useful 'washing lather' and hard water often leads to deposits (lime scale) forming in heating systems and kettles which require cleaning at times.
    • The 'lime scale' is usually caused by the thermal decomposition of the dissolved hydrogencarbonates producing insoluble calcium carbonate (so it does remove some of the temporary hardness before washing! and chemically, it is the opposite of the carbonated water dissolving action above) ...
    • Ca(HCO3)2(aq) ==> CaCO3(s) + H2O(l) + CO2(g)  
    • However there is a plus side to the deposition! The coating on the inner surface of the pipe work prevents corrosion and the dissolving of potentially poisonous salts of copper or lead into the water supply.
    • The lime scale can be removed by any acid (hydrogen ion solution) treatment which dissolves the calcium carbonate.
      • ionically this is: CaCO3(aq) + 2H+(aq) ==> Ca2+(aq) + H2O(l) + CO2(aq) 
      • e.g. vinegar contains the weak organic acid ethanoic acid and will dissolve lime scale in kettles but shouldn't react with the steel container or heating element.
      • calcium carbonate + ethanoic acid ==> calcium ethanoate + water + carbon dioxide
      • CaCO3(aq) + 2CH3COOH(aq) ==> Ca2+(CH3COO-)2(aq) + H2O(l) + CO2(aq) 
      • In the school lab. you will doubt at some point you add the 'strong' hydrochloric acid to marble chips, which is essentially a very similar reaction to the one dissolving limescale above.  The reaction is faster if the vinegar is hot because all reactions are speeded by higher temperatures because of the increased kinetic energy of the reactant particles (see rates of reaction for more details) and maybe also because calcium ethanoate is not that soluble in cold water and dissolves more in hot water (not sure of the importance of this 2nd factor?).
      • calcium carbonate + hydrochloric acid ==> calcium chloride + water + carbon dioxide
      • CaCO3(aq) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(aq)
      • or showing the ions involved
      • CaCO3(aq) + 2H+Cl-(aq) ==> Ca2+(Cl-)2(aq) + H2O(l) + CO2(aq)
      • more simply and the more correct ionic equation ...
      • CaCO3(aq) + 2H+(aq) ==> Ca2+(aq) + H2O(l) + CO2(aq)
  • Hard water can be made soft by removing the dissolved calcium and magnesium ions.
    • If due to calcium/magnesium hydrogencarbonates it is removed by boiling (see above).
    • Adding enough 'soapy' soap, see above, but the water is best treated before the washing!, so its not the desired solution with the scum and all that!
    • The addition of sodium carbonate (as 'washing soda' crystals), which dissolves and precipitates out the calcium or magnesium ions as their insoluble carbonates(s) formed.
    • calcium sulphate + sodium carbonate ==> calcium carbonate + sodium sulphate
    • CaSO4(aq) + Na2CO3(aq) ==> CaCO3(s) + Na2SO4(aq) 
    • or more simply ionically: Ca2+(aq) + CO32-(aq) ==> Ca2+CO32-(s) (called an 'ionic equation')
    • Packs of ion exchange resins can hold or release ions in an ion exchange process.
      • Negative polymer resin columns hold hydrogen ions or sodium ions. These can be replaced by calcium and magnesium ions when hard water passes down the column. The calcium or magnesium ions are held on the negatively charged resin. The freed hydrogen or sodium ions do not form a scum with soap.
      • e.g. 2[resin]-H+(s) + Ca2+(aq) ==> [resin]-Ca2+[resin]-(s) + 2H+(aq)
      •  or 2[resin]-Na+(s) + Mg2+(aq) ==> [resin]-Mg2+[resin]-(s) + 2Na+(aq) etc.
    • Extra Note on water purification: You can also use an ion-exchange resin to replace negative ions by using a positively charged resin initially holding hydroxide ions e.g. to remove chloride (Cl-), nitrate (NO3- is potentially harmful) and sulphate ions (SO42-)e.g.
      • [resin]+OH-(s) + Cl-(aq) ==> [resin]+Cl-(s) + OH-(aq)
      • [resin]+OH-(s) + NO3-(aq) ==> [resin]+NO3-(s) + OH-(aq)
      • 2[resin]+OH-(s) + SO42-(aq) ==> [resin]+SO42-[resin]+(s) + 2OH-(aq) etc.
    • Now, by using both a positive and negatively charged resin, you can completely de-ionise water because the released hydrogen ions and hydroxide ions combine to form pure water.
      • H+(aq) + OH-(aq) ==> H2O(l) 
      • However, it will not remove non-ionic substances like organic pesticides etc.
  • Permanently hard water means the hardness cannot be removed by boiling e.g. when caused by dissolved magnesium or calcium sulphate.
  • Temporary hard water means it is softened by boiling e.g. when caused by magnesium hydrogencarbonate or calcium hydrogencarbonate.
  • HOWEVER, a plus point! Hard water contains dissolved compounds that are good for health. Hard water often provides calcium compounds that help the development of strong bones and teeth and help to reduce heart illnesses and also traces of other essential elements like iron and iodine.

(c) doc b


2a. How well do different gases and solids dissolve in water?

  • First, some definitions of words you may encounter in talking about solubility and other water related situations:

    • solute: the material which is to be dissolved in a solvent.

    • solvent: the liquid which dissolves the material (the solute). You will come across water more than any other liquid solvent BUT lots of important organic solvents like hexane (petrol like), ethanol (alcohol) and propanone (acetone) are in common laboratory use.

    • solution: the result of dissolving something in a liquid (solute + solvent => solution).

    • solubility: to what extent a solute material will dissolve.

    • soluble: the material will dissolve in a particular liquid solvent.

    • saturated: means that no more of a substance (the solute) will dissolve in its solution i.e. maximum solubility achieved at a particular temperature.

    • insoluble: not soluble, will not dissolve in a particular liquid (don't assume it means will not dissolve in anything).

    • hydration: means the addition of water to a material.

    • dehydration: means to remove water from a substance.

  • Factors affecting rates of dissolving.

    • heat: heating the mixture to raise the temperature will increase the rate of a substance dissolving - the energy of all the particles involved is increased - increased rate of more energetic collisions between solute and solvent particles speeding up the dissolving process.

    • surface area: if a solid is broken up and crushed into smaller pieces or a powder it will dissolve faster. This breaking down of a solid increases the surface are for the solvent to 'attack' and dissolve the solid.

    • stirring: this increases the rate of dissolving because it prevents 'local' saturation of the solution which will inhibit dissolving.

    • volume of solvent: adding more solvent increases the speed of dissolving, the less concentrated

    • These factors are similar with those affecting the rates of chemical reactions except there is no catalyst that speed up dissolving as far as I know? Also, increasing the volume of the solvent will decrease the rate of reaction because concentrations are reduced.

  • Some gases and solid substances are more soluble in water than others and some are hardly soluble at all.

  • The solubility of gases and solids in water also depends on the temperature of the water:

  • Many gases are soluble in water and the solubility increases as the temperature decreases and as the pressure increases.

  • Carbonated water is produced by dissolving carbon dioxide under high pressure. When the pressure is released the gas bubbles out of the solution. Carbonated water is used to give fizzy drinks a 'tang' to the taste.

    • It is a weakly acid solution, explaining why rainwater containing dissolved carbon dioxide from the air, can very slowly dissolve limestone.

      • The solution of CO2(aq) is sometimes described as 'carbonic acid', H2CO3, but this does not really exist!

        • However, the solution is acidic due to the formation of hydrogen ions.

        • CO2(aq) + H2O(l) (c) doc b H+(aq) + HCO3-(aq) 

        • Note: the equilibrium is almost completely on the left.

  • Thermal Pollution: Dissolved oxygen is essential for aquatic life and the colder the water, the more of it dissolves. Hot water from power stations may be discharged into rivers or lakes. This discharge reduces the amount of oxygen dissolved in the water and this can damage aquatic life and disrupt the natural eco-systems.

  • Chlorine water is made by dissolving Chlorine gas in water and can be a useful chemical reagent, both in the laboratory and industry (e.g. displaces iodine from sea water).

  • Chlorine water is used to bleach materials and kill bacteria.

  • Many ionic compounds are soluble in water and many covalent compounds are insoluble in water (but don't make assumptions!).

  • The solubility of a solute in water, or any other solvent, is usually given in grams of solute per 100 grams of solvent (e.g. water) at that temperature.

  • The solubility of most solid solutes increases as the temperature increases (opposite of gases, but the ambient air pressure has no effect).

  • A saturated solution is one in which no more solute will dissolve at that temperature giving the maximum solubility at that particular temperature.

  • When a hot saturated solution cools some of the solute will separate from the solution (crystallisation). The crystals form because the solubility is lower at the lower temperature.

  • From solubility graphs-data you can calculate how much will dissolve at a given temperature and how much will crystallise out on cooling.

  • Solubility curves:

  • General rules which describe the solubility of common types of compounds in water:

    • All common sodium, potassium and ammonium salts are soluble e.g. NaCl, K2SO4, NH4NO3

    • All nitrate salts are soluble e.g. NaNO3, Mg(NO3)2, Al(NO3)3, NH4NO3

    • Some ethanoate salts are soluble e.g. CH3COONa

    • Common chloride salts are soluble except those of silver and lead e.g.

      • soluble: KCl, CaCl2, AlCl3 or insoluble AgCl, PbCl2

    • Common sulfates are soluble except those of lead, barium and calcium: soluble e.g.

      • soluble: Na2SO4, MgSO4, Al2(SO4)3

      • insoluble: PbSO4, BaSO4, CaSO4 is slightly soluble.

    • Common oxides, hydroxides and carbonates are usually insoluble (e.g. Group 2 and Transition Metals) except those of the Group 1 Alkali Metals sodium, potassium etc. and ammonium:

      • soluble: K2O, KOH, NaOH, NH4OH actually NH3(aq), Na2CO3, (NH4)2CO3  

      • insoluble: MgO, CuO, ZnO, Mg(OH)2, Fe(OH)2, Cu(OH)2, CuCO3, ZnCO3, CaCO3

  • Knowledge of salt solubility is important in deciding which method of salt preparation is employed.

  • See salt preparation methods and details of various methods.

(c) doc b


2b. Solubility curves for selected salts

solubility curves for potassium nitrate, potassium sulphate/sulfate, sodium chloride, copper(II) sulphate

  • Interpretation of graph eg

    • Reading graph: at 38oC the solubility of copper sulphate, CuSO4, is 28g of anhydrous salt per 100g of water.

    • Reading graph: at 84oC the solubility of potassium sulphate, K2SO4, is 22g per 100g of water.

    • Ex Q1: How much potassium nitrate will dissolve in 20g of water at 34oC?

      • At 34oC the solubility is 52g per 100g of water,

      • so scaling down, 52 x 20 / 100 = 10.4g will dissolve in 20g of water

    • Ex Q2: At 25oC 6.9g of copper sulphate dissolved in 30g of water, what is its solubility in g/100cm3 of water?

      • Scaling up, 6.9 x 100 /30 = 23g/100g of water (check on graph, just less than 23g/100g water).

    • Ex Q3: 200 cm3 of saturated copper solution was prepared at a temperature of 90oC. What mass of copper sulphate crystals form if the solution was cooled to 20oC?

      • Solubility of copper sulphate at 90oC is 67g/100g water, and 21g/100g water at 20oC.

      • Therefore for mass of crystals formed = 67 - 21 = 46g (for 100 cm3 of solution).

      • However, 200 cm3 of solution was prepared,

      • so total mass of copper sulphate crystallised = 2 x 46 = 92g

  • Note: The density of water is close to 1.0g/cm3 or ml, so for approximate purposes. the volume in cm3 or ml of just the water is numerically close to the value in g, i.e. 100 cm3 of water or solution is about 100g of water.


Examples of



g salt / 100g water

Salt name

potassium nitrate

potassium sulphate

sodium chloride

hydrated copper(II) sulphate

and formula

Temp. oC




CuSO4 (anhydrous *)

























































* multiply by 1.562 for hydrated crystals CuSO4.5H2O

(c) doc b


3. Water of crystallisation calculations

  • Solubility graphs and data are covered in section 2.

  • How to calculate the theoretical % of water in a hydrated salt

    • eg magnesium sulphate MgSO4.7H2O salt crystals

    • Relative atomic masses: Mg = 24, S = 32, O = 16 and H = 1

    • relative formula mass = 24 + 32 + (4 x 16) + [7 x (1 + 1 + 16)] = 246

    • 7 x 18 = 126 is the mass of water

    • so % water = 126 x 100 / 246 = 51.2%

  • Determination and calculation of salt formula containing 'water of crystallisation'.

    • Some salts, when crystallised from aqueous solution, incorporate water molecules into the structure. This is known as 'water of crystallisation', and the 'hydrated' form of the compound.

    • e.g. magnesium sulphate MgSO4.7H2O. The formula can be determined by a simple experiment (see the copper sulphate example below).

    • A known mass of the hydrated salt is gently heated in a crucible until no further water is driven off and the weight remains constant despite further heating. The mass of the anhydrous salt left is measured. The original mass of hydrated salt and the mass of the anhydrous salt residue can be worked out from the various weighings.

    • The % water of crystallisation and the formula of the salt are calculated as follows:

      • Suppose 6.25g of blue hydrated copper(II) sulphate, CuSO4.xH2O, (x unknown) was gently heated in a crucible until the mass remaining was 4.00g. This is the white anhydrous copper(II) sulphate.

      • The mass of anhydrous salt = 4.00g, mass of water (of crystallisation) driven off = 6.25-4.00 = 2.25g

      • The % water of crystallisation in the crystals  is 2.25 x 100 / 6.25 = 36%

      • [ Ar's Cu=64, S=32, O=16, H=1 ]

      • The mass ratio of CuSO4 : H2O is 4.00 : 2.25

      • To convert from mass ratio to mole ratio, you divide by the molecular mass of each 'species'

      • CuSO4 = 64 + 32 + (4x18) = 160 and H2O = 1+1+16 = 18

      • The mole ratio of CuSO4 : H2O is 4.00/160 : 2.25/18

      • which is 0.025 : 0.125 or 1 : 5, so the formula of the hydrated salt is CuSO4.5H2O

  • All concentration calculations are covered on the on-line CLICK for GCSE Chemical Calculations calculations page, especially sections 7. on molarity, 11. and 12. on molarity and acid-base (alkali) titrations, section 14.3 on dilutions.

  • GCE A level advanced notes on the structure of hydrated salts

Revision KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel GCSE Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  WJEC gcse science chemistry CCEA/CEA gcse science chemistry O Level Chemistry (revise courses equal to US grade 8, grade 9 grade 10)

equation keywords: SO2(aq) + Cl2(aq) + 2H2O(l) ==> 2HCl(aq) + H2SO4(aq) SO2(aq) + Cl2(aq) + 2H2O(l) ==> 2Cl-(aq) + SO42-(aq) + 4H+(aq) CaCO3(s) + H2O(l) + CO2(g) ==> Ca(HCO3)2(aq) CaSO4(aq) + 2C17H35COONa(aq) ==> (C17H35COO)2Ca(s for scum!) + Na2SO4(aq) Ca2+(aq) + 2C17H35COO-(aq) ==> (C17H35COO-)2Ca2+(s) Ca(HCO3)2(aq) ==> CaCO3(s) + H2O(l) + CO2(g)  CaCO3(aq) + 2H+(aq) ==> Ca2+(aq) + H2O(l) + CO2(aq)  CaCO3 (aq) + 2CH3COOH(aq) ==> Ca2+(CH3COO-)2(aq) + H2O(l) + CO2(aq) CaCO3(aq) + 2HCl(aq) ==> CaCl2(aq) + H2O(l) + CO2(aq) CaCO3(aq) + 2H+Cl-(aq) ==> Ca2+(Cl-)2(aq) + H2O(l) + CO2(aq) CaCO3(aq) + 2H+(aq) ==> Ca2+(aq) + H2O(l) + CO2(aq) CaSO4(aq) + Na2CO3(aq) ==> CaCO3 (s) + Na2SO4(aq) Ca2+(aq) + CO32-(aq) ==> Ca2+CO32-(s)  CO2(aq) + H2O(l) <==> H+(aq) + HCO3-(aq)   SO2 + Cl2 + 2H2O ==> 2HCl + H2SO4 SO2 + Cl2 + 2H2O ==> 2Cl- + SO42- + 4H+ CaCO3 + H2O + CO2 ==> Ca(HCO3)2 CaSO4 + 2C17H35COONa ==> (C17H35COO)2Ca + Na2SO4 Ca2+ + 2C17H35COO- ==> (C17H35COO-)2Ca2+ Ca(HCO3)2 ==> CaCO3 + H2O + CO2  CaCO3 + 2H+ ==> Ca2+ + H2O + CO2  CaCO3 + 2CH3COOH ==> Ca2+(CH3COO-)2 + H2O + CO2 CaCO3 + 2HCl ==> CaCl2 + H2O + CO2 CaCO3 + 2H+Cl- ==> Ca2+(Cl-)2 + H2O + CO2 CaCO3 + 2H+ ==> Ca2+ + H2O + CO2 CaSO4 + Na2CO3 ==> CaCO3 + Na2SO4 Ca2+ + CO32- ==> Ca2+CO32-  CO2 + H2O <==> H+ + HCO3- 

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