Doc Brown's Chemistry

 Water Cycle - including water as a resource, hard and soft water, colloids, salt solubility and water of crystallisation

Revision notes for KS4 Science GCSE/IGCSE/O level Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  (revise courses equal to US grades 9-10)

Advanced Level Notes on pH and Acid Base Theory etc.

1. Water: cycle, resource, treatment, pollution, colloids (sol, foam, emulsion), hard/soft water * 2. Gas and salt solubility and solubility curves * 3. Acid-base proton-hydrogen ion (H⁺) theory of weak/strong acids/bases (alkalis) * 4. Methods of making salts * 5. Calculation of water of crystallisation * Basic stuff on pH, Acids, Bases, Neutralisation and Salts (separate page) * Selected salt solubility data and simple calculations (separate page)

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.

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 SO_2(aq) + Cl_2(aq) + 2H₂O_(l) ==> 2HCl_(aq) + H₂SO_4(aq)
    • the ionic equation is SO_2(aq) + Cl_2(aq) + 2H₂O_(l) ==> 2Cl^-_(aq) + SO₄^2-_(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 in another liquid e.g. (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) and (iii) 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.
  • 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: Al³⁺ > Mg²⁺ > Na⁺ or [Fe(CN)₆]^3- > SO₄^2- > Cl^- 
    • and this explains why aluminium sulphate Al₂(SO₄)₃ is used to precipitate (coagulate) colloidal clay in water treatment.
• Other onsite references to water pollution:
  • Acid rain is mentioned on the Oil Products page.
  • Fertiliser pollution mentioned on the Equilibrium and Ammonia page.

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
  • CaCO_3(s) + H₂O_(l) + CO_2(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.
  • It is a simple and effective way of comparing the 'hardness' in water samples.
  • The apparatus is the same as that used in salt preparation method (a), but no indicator is used, the end-point is detected by the appearance of decent froth!
• 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' ..
  • CaSO_4(aq) + 2C₁₇H₃₅COONa_(aq) ==> (C₁₇H₃₅COO)₂Ca_{(s for scum!)} + Na₂SO_4(aq)
  • or more simply ionically: 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'.
• 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) ==> CaCO_3(s) + H₂O_(l) + CO_2(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: CaCO_3(aq) + 2H⁺_(aq) ==> Ca²⁺_(aq) + H₂O_(l) + CO_2(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
    • CaCO_3(aq) + 2CH₃COOH_(aq) ==> Ca²⁺(CH₃COO^-)_2(aq) + H₂O_(l) + CO_2(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
    • CaCO_3(aq) + 2HCl_(aq) ==> CaCl_2(aq) + H₂O_(l) + CO_2(aq)
    • or showing the ions involved
    • CaCO_3(aq) + 2H⁺Cl^-_(aq) ==> Ca²⁺(Cl^-)_2(aq) + H₂O_(l) + CO_2(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
  • CaSO_4(aq) + Na₂CO_3(aq) ==> CaCO_3(s) + Na₂SO_4(aq) 
  • or more simply ionically: Ca²⁺_(aq) + CO₃^2-_(aq) ==> Ca²⁺CO₃^2-_(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 (SO₄^2-)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) + SO₄^2-_(aq) ==> [resin]⁺SO₄^2-[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.

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 CO_2(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.

      • CO_2(aq) + H₂O_(l) 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:

  • A solubility curve is the graph of how much of a substance dissolves at a given temperature versus the the temperature.

    • See section 2b. for Solubility data and 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 NH_3(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.

2b. Solubility curves for selected salts

 3. Acid-base proton-hydrogen ion (H⁺) theory of weak/strong acids/bases (alkalis)

4. Methods of making salts * Both these sections are now on new pages

5. How can we work out and use the concentration of solutions?

and water of crystallisation calculations

• Solubility graphs and data are covered in section 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%

    • [ A_r'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 calculations page, especially sections 7. on molarity, 11. and 12. on molarity and acid-base (alkali) titrations, section 14.3 on dilutions.

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