School Science Lessons
Chemistry experiments
Please send comments to: J.Elfick@uq.edu.au
Updated: 2009-09-16
Table of contents
3.32.0 Prepare, collect and test gases
3.52.0 Rusting
3.54.0 Crystal growth
3.55.0 Matter as particles
3.59.0 Conductors of electricity
3.61.0 Construction materials
15.1.0 Electroplating
15.5.0 Electrolysis
3.32.0 Prepare, collect and test gases
3.44 Prepare nitric oxide (nitrogen monoxide, NO)
3.44.1 Catalytic conversion of nitric oxide (nitrogen monoxide)
3.45 Prepare nitrous oxide (dinitrogen oxide, N2O)
3.45.1 Tests for dinitrogen oxide (nitrous oxide)
3.46 Prepare nitrogen, N2
3.47 Prepare nitrogen dioxide, NO2
3.47.1 Pass nitrogen dioxide through water
3.48 Acid rain and nitrogen oxides, NOx
3.49 Prepare oxygen gas with hydrogen peroxide, O2
3.49a Hydrogen peroxide concentration and storage
3.49.1 Tests for oxygen gas
3.50 Ozone, O3, Prepare ozone
3.50.1 Ozone and photochemical smog
3.51 Prepare sulfur dioxide, SO2
3.51.1 Tests for sulfur dioxide
3.51.2 Reduce potassium manganate (VII) with sulfur dioxide
3.51.3 Reduce iron (III) chloride with sulfur dioxide
3.51.4 Bleach flowers with sulfur dioxide
3.52.0 Rusting
3.52 Conditions necessary for rusting
3.52.1 The mass of iron and its temperature increases during rusting
3.52.2 Oxygen gas combines with iron during rusting
3.52.3 Metals can prevent rusting
3.54.0 Crystal growth
3.54 Prepare crystals from solutions
3.54.01 Prepare double salt crystals
3.54.1 Prepare sea water crystals
3.54.2 Prepare crystals from a melt
3.54.3 Prepare crystals with different shapes
3.54.4 Prepare crystals from a mixture of salts
3.54.5 Prepare large crystals
3.54.6 Prepare clusters of crystals
3.54.7 Prepare split crystals
3.54.8 Prepare stalactite crystals
4.14 Prepare crystals (Primary)
3.1.4 Prepare aspirin crystals
3.1.10 Prepare sucrose crystals from brown sugar
3.1.10.1 Prepare sucrose crystals from sugar cane juice
12.18.6.1 Prepare sodium thiosulfate crystals, "hypo", Na2S2O3.5H2O,
12.18.6.2 Reactions of sodium thiosulfate, Na2S2O3.5H2O
12.10.0 Prepare boric acid crystals
3.55.0 Matter as particles
3.55 Brownian movement
3.55.1 Diffusion of heavier than air gas, carbon dioxide
10.1.2 Diffusion of ammonia and hydrogen chloride gases
3.55.3 Diffusion of liquids
3.56 Particles of matter and dilution
3.57 Size of a molecule
3.58 Clay soil suspension
3.61.0 Construction materials
3.61 Tin-lead alloys
3.62 Tin-lead alloys and pure metals, hardness
3.63 Melting points of metals and alloys
3.64 Heat treatment of steel needles, annealing, quenching, tempering
3.65 Strengths of mud, clay and sand bricks
3.66 Make bricks with cement
3.67 Strength of plaster of Paris
3.66 Make bricks with cement
3.66.1 Change in weight of setting cement
3.66.2 Strength of cement with changing water content
3.66.3 Alkalinity of concrete
3.44 Prepare nitric oxide (nitrogen monoxide, NO)
See diagram 3.44: Prepare nitrogen monoxide | See diagram 1.13a: Simple fume hood
Nitrogen monoxide is a colourless gas that reacts immediately with oxygen gas to form nitrogen dioxide forms brown fumes of nitrogen dioxide, NO2 in air. Nitric oxide has many physiological functions in the body, e.g. dilation of blood vessels.
1. Do this experiment in a fume cupboard, fume hood or near an open window. Add drops of dilute nitric acid to copper. Nitrogen oxide forms that immediately reacts with oxygen in the air to form nitrogen dioxide.
2. Do this experiment in a fume cupboard, fume hood. Fit each end of a glass tube, diameter 2 cm and length 15 cm, and with two one-hole stoppers fitted with short delivery tubes. Connect one delivery tube to a short rubber tube with a pinch tap on it. Fix a 2 cm rubber gasket with holes drilled in it in the middle part of the glass tube. Drop pieces of copper on the rubber gasket. Remove the stopper from the other end of the glass tube and add dilute nitric acid, i.e. 1 to 3 ratio of concentrated nitric acid to water. Add enough diluted nitric acid so that no space is left for an air bubble after replacing the stopper. Invert the apparatus and clamp it vertically with a container under the glass tube. During the reaction the colourless gas produced gradually presses the solution out of the tube. When the copper pieces are no longer in contact with the nitric acid solution the reaction stops leaving nitrogen oxide in the upper portion of the glass tube and a blue solution in the small container. Open the pinch tap to let air enter the tube and note the oxidation to form brown fumes of nitrogen dioxide. The level of the solution rises to show that nitrogen dioxide dissolves in water. To prevent air contamination, absorb the nitrogen dioxide in the tube with an alkali solution.
3Cu (s) + 8HNO3 (aq) --> 3Cu(NO3)2 (aq) + 4H2O (l) + 2NO (g)
2NO (g) + O2 (g) --> 2NO2 (g) [The O2 comes from the air.]
3. Do this experiment in a fume cupboard, fume hood. The experiment may be done in a syringe. Add 5 mL of 1.2 M solution of FeSO4 in 1.8 M H2SO4 to 0.27 g of solid NaNO3. The reaction mixture turns black and and nitric acid gas is produced. Wash the gas through deionized water.
NO2- + Fe2+ (aq) + 2H+ --> NO (g) + Fe3+ (aq) + H2O
3.44.1 Catalytic conversion of nitric oxide (nitrogen monoxide)
When the petrol and air mixture ignites in motor car engines, nitrogen and oxygen gas combine to form poisonous nitrogen oxide. In a three-way catalytic converter fitted to a high performance motor car engine using unleaded petrol, carbon monoxide and unburned hydrocarbons (Hcarb) combine with nitrogen monoxide over a platinum rhodium catalyst to form harmless gases.
2CO (g) + 2NO (g) --> 2CO2 (g) + N2 (g)
Hcarb (g) + NO (g) --> CO2 (g) + H2O (g) + N2 (g)
Hcarb (g) + O2 --> CO2 (g) + H2O (g)
3.45 Prepare dinitrogen oxide (nitrous oxide, N2O, nitrogen (I) oxide, dinitrogen monoxide)
Dinitrogen oxide is a colourless gas that is soluble in water, with a sweet smell that does not change the colour of moist litmus. It was previously used as an anaesthetic in dentistry ("laughing gas") and as a propellant in aerosol sprays. Be careful! Do NOT inhale nitrous oxide. It can cause death.
Be careful! If you heat ammonium nitrate to dryness, it may decompose with an explosion.
Put 2 g of ammonium nitrate in a dry large test-tube. Fit a stopper with a delivery tube in the test-tube. Heat the test-tube slowly and pass the gas formed through 5% iron (II) sulfate solution to remove any nitrogen monoxide. Collect the gas in a receiving test-tube.
3.45.1 Tests for dinitrogen oxide (nitrous oxide) N2O
Put a glowing splint down in the receiving test-tube and note the gas relighting the splint.
3.46 Prepare nitrogen, N2
 Atmospheric nitrogen cannot be used directly by the body. Liquid nitrogen is used to freeze tissues for microscopic examination.
Nitrogen gas is colourless, odourless, tasteless, neutral and unreactive. Nitrogen does not support combustion but magnesium and calcium will continue to burn in nitrogen to form nitrides. Nitrogen is manufactured by fractional distillation of air. Air contains about 78% of nitrogen.
Ammonium nitrite is unstable so the reaction of saturated solutions of sodium nitrite with ammonium chloride can be used to prepare nitrogen.
Be careful! This reaction can explode without warning.
Heat gently a test-tube containing 5 mL of saturated ammonium chloride solution. Add saturated sodium nitrite solution drop by drop. The reaction is exothermic, so stop heating when some gas forms. The ammonium nitrite breaks down into nitrogen and water.
NH4NO2 (s) --> N2 (g) + 2H2O (l)
3.47 Prepare nitrogen dioxide, NO2, nitrogen (IV) oxide
Nitrogen dioxide is a brown gas with a choking smell that may irritate the lungs and lead to death.
 See diagram 1.13a: Simple fume hood
Be careful! Do this experiment in a fume cupboard, fume hood.
Liquid dinitrogen tetroxide, N2O4, b.p. 21.2oC, dissociates as a gas to form nitrogen dioxide, NO2.
Pour drops of concentrated nitric acid on pieces of copper in a test-tube. Fix a stopper in the test-tube immediately because brown noxious gas nitrogen dioxide forms with a pungent irritating odour. The nitric acid acts as an oxidizing agent and is reduced to nitrogen dioxide and water.
Be careful! The reaction is exothermic.
Cu (s) + 4HNO3 (aq) --> Cu(NO3)2 (aq) + 2H2O (l) + 2NO2 (g)
3.47.1 Pass nitrogen dioxide through water
Nitrogen dioxide is decomposed by water to form a mixture of nitric acid (HNO3) and nitrous acid (HNO2).
Note the colour and odour of the water. Test the solution with litmus paper.
2NO2 + H2O --> HNO3 + HNO2
3.48 Acid rain and nitrogen oxides, NOx
Nitrogen oxides, abbreviated as "NOx", are mainly 1. colourless nitrogen oxide, NO (nitric oxide) 2. red-brown nitrogen dioxide NO2 [nitrogen (IV) oxide] in equilibrium with dinitrogen tetroxide, N2O4 (2NO2 < = > N2O4) with a pungent irritating odour, and 3. colourless dinitrogen oxide, N2O (nitrous oxide). Nitrogen oxides are produced naturally by decomposition, bacterial nitrification and lightning. They are produced artificially by coal-fired power stations from coal and by oil-fired power stations. When fuel burns at high temperatures in internal combustion engines, the nitrogen and oxygen gas of the air combine to form nitrogen oxides. These oxides dissolve in water to produce the dilute nitric acid that is one component of acid rain. Rain water pH = 5.6, acid rain pH < 5.
N2 (g) + O2 (g) --> 2NO (g) nitric oxide
2NO (g) + O2 (g) --> 2NO2 (g)
4NO (aq) + O2 (aq) + 2H2O --> 4HNO3 (aq)
2NO2 (g) + H2O (l) --> HNO2 (aq) + HNO3 (aq)
N2 (g) + 2O2 (g) --> 2NO2 (g) nitrogen dioxide
2NO2 (g) + H2O (l) --> HNO2 (aq) + HNO3 (aq) nitrous acid + nitric acid
HNO2 (aq) + heat --> HNO3 + 2NO + H2O (l)
4NO (g) + 3O2 (g) + 2H2O --> 4HNO3
3.49 Prepare oxygen gas, O2
See diagram 3.49: Prepare oxygen, a holder for burning substances | See: Saturation vapour pressure over water
Oxygen gas is a colourless odourless diatomic gas that supports combustion and is essential for aerobic respiration. Oxygen gas reacts with metals to form basic oxides. Oxygen gas reacts with non-metals to form acidic oxides. Oxygen does not change the colour of moist litmus.
The safest method to prepare oxygen is by decomposition of hydrogen peroxide solution.
3.49a Hydrogen peroxide concentration and storage
Hydrogen peroxide may be sold in two strengths as follows:
1. 3% aqueous solution w/v, 3 g / mL, 10 volume, 10 vols, (10% volume)
2. 6% aqueous solution w/v, 6 g / mL, 20 volume = 20 vols. (20% volume)
The commercial strength of an aqueous solution is represented by the the volumes of oxygen gas that 100 cm3 of the liquid solution will give on decomposition. So a 20 volume concentration means that when 1 volume of hydrogen peroxide solution is decomposed, it produces 20 volumes of oxygen gas. As a hydrogen peroxide 20 volume solution contains 6% H2O2, 60g of H2O2 in 1000g of solution, atomic mass of H2O2 (H atomic mass 1.008, oxygen gas atomic mass 16) = 34.016, so the concentration of a 20 volume solution is 60 / 34.016 = 1.76 m. The smaller concentration, 3% w/v H2O2, is less stable and decomposes faster at room temperatures so the actual concentration is probably less than 3%. Protect hydrogen peroxide solution from light and store in a cool place. Keep it in a brown glass bottle closed with a glass stopper, paraffined cork or plastic screw-cap. Airlines may not be allowed to carry hydrogen peroxide as freight. Hydrogen peroxide test strips are available that contain a peroxide reagent to detect production of hydrogen peroxide by certain bacteria, e.g. Streptococcus pneumoniae.
1. Pour some 20 volumes (vols) hydrogen peroxide into a test-tube containing manganese (IV) oxide granules. Collect oxygen gas in receiver test-tubes over water and apply stoppers to the test-tubes. Store test-tubes in a test-tube rack and remove the stoppers just before inserting the burning element.
2H2O2 (aq) --> O2 (g) + 2H2O (l) [with MnO2 as catalyst]
2. Put 1 cm depth of hydrogen peroxide solution in a test-tube. Add a drop of iron sulfate solution (FeSO4.7H2O, ferrous sulfate, green vitriol). The contents froth vigorously. Test for oxygen gas with the glowing splint test.
3. Prepare oxygen gas with household bleach, sodium hypochlorite, or bleaching powder
Liquid household bleach is usually 5% sodium hypochlorite, NaOCl. Commercial bleach is made by passing chlorine gas through sodium hydroxide solution until neutral pH then diluted to 5%. 1. Heat 1 cm depth of bleaching powder in a dry test-tube. 2. Put 1 cm depth of concentrated household bleach solution (NaOCl, bleaching fluid) or bleaching powder solution in a test-tube and add drops of concentrated cobalt chloride solution. A black precipitate forms. Heat the test-tube until frothing starts. 3. Add 1 cm depth of bleaching powder to 2 cm depth of water in a test-tube. Heat the test-tube. No oxygen gas forms. Add drops of copper (II) sulfate solution and heat again. Little oxygen gas forms. Add drops of iron sulfate solution (FeSO4.7H2O, ferrous sulfate, green vitriol). The solution effervesces strongly because much oxygen gas forms.
4. Prepare oxygen gas with potassium manganate (VII)
Wear eye protection
Put a two fingers depth of potassium manganate (VII) in a Pyrex test-tube. To control "spitting" put a loose plug of ceramic wool in the mouth of the test-tube. Heat the test-tube slowly and hold a glowing splint over the mouth of the test-tube to detect oxygen gas.
3.49.1 Tests for oxygen gas
1. Glowing splint test
Light a splint of wood. Blow out the flame then hold the glowing splint in the test-tube full of oxygen gas. The splint relights.
2. Steel wool test
Collect oxygen gas in test-tubes with stoppers. Use an L-shaped piece of nichrome wire with a shield to fit on the top to protect your hand. Fix steel wool into a loop in the lower end of the Nichrome wire. Heat the steel wool to red heat in a Bunsen burner flame then insert it quickly into a test-tube of oxygen gas. The steel wool burns with bright sparkles to form grey-black iron oxide, Fe3O4 (FeO.Fe2O3). Sprinkle iron filings into a Bunsen burner flame. A shower of sparks occurs, as in some fireworks.
6Fe + 4O2 --> 2Fe3O4
3. Charcoal test
Fix charcoal into the loop in the lower end of the Nichrome wire or use a combustion spoon. Heat the charcoal in a Bunsen burner flame until it has a red glow, then quickly insert it into a test-tube of oxygen gas. The charcoal glows much more.
C + O2 --> CO2 (g)
4. Magnesium ribbon test
Be careful! Do NOT look at the bright flame.
Wrap a 3 cm piece of magnesium ribbon around the loop at the end of a wire. Ignite it in a Bunsen burner and put it quickly in the oxygen gas. Magnesium burns with a very bright flame.
2Mg + O2 --> MgO (s)
3.50 Ozone, O3
See diagram 3.50: Prepare ozone
1. Ozone can be prepared in the laboratory using a high voltage induction coil, spark coil. However, the experiment is dangerous. Do not attempt it in a school laboratory.
3O2 (g) --> 2O3 (g)
2. In the upper stratosphere, ozone forms when ultraviolet (UV) light splits an oxygen gas molecule, O2, into two atoms of oxygen, O.
O2 + UV --> O + O.
The oxygen atoms can react with other oxygen molecules to produce ozone that sinks down to the lower stratosphere, between 20 and 40 km above the Earth.
O + O2 --> O3
Also, the oxygen atoms can react with ozone to produce oxygen molecules again.
O + O3 --> O2 + O2
When ozone absorbs UV light, the ozone breaks into an oxygen molecule and an oxygen atom again.
O3 + UV --> O2 + O
Ozone is an allotrope of oxygen, a poisonous blue gas, a powerful oxidizing agent, a bleach and a germicide. The stratosphere holds 90% of the ozone in the atmosphere. The ozone layer in the stratosphere, 15 to 50 km above the Earth, absorbs most of the high frequency ultraviolet light from the sun, which can cause genetic damage to DNA in plant and animal cells, sunburn (erythma) and skin cancer. Ozone is produced with oxides of nitrogen by reactions of car exhaust gases with unburned fuel and sunlight to produce photochemical smog. Small amounts of ozone are produced by electrical discharges. After a thunderstorm, the refreshing smell is ozone and nitrogen oxides. Ozone in the total atmospheric column is measured in Dobson units.
3.50.1 Ozone and photochemical smog
A nitrogen dioxide molecule can be dissociated by absorbing photon, hv, of sunlight.
NO2 (g) + hv -->NO (g) + O (g) [reactive oxygen atom]
O (g) + O2 (g) --> O3 (g)
2NO (g) + O2 (g) --> 2NO2 (g)
3.51 Prepare sulfur dioxide, SO2
See diagram 3.51.1: Prepare sulfur dioxide by burning | See diagram 3.51.2: Sulfur dioxide generator | See diagram 1.13a: Simple fume hood
Sulfur dioxide is a colourless gas that irritates the lungs. Sulfur dioxide dissolves in water to form mainly, sulfurous acid, H2SO3. Sulfur dioxide is one component of acid rain.
SO2 (g) + H2O (g) --> H2SO3 (l)
Be careful! Do not inhale the gas. Do the following preparations in a fume cupboard, fume hood.
1. Ignite sulfur in an evaporating basin and collect the sulfur dioxide formed under a filter funnel. Connect the filter funnel to a receiving container with a two-holes stopper containing water. Use a filter pump to suck the sulfur dioxide into the receiving container.
S + O2 --> SO2
2. Pour some water into a screw cap container. Heat a small amount of sulfur in a combustion spoon over a Bunsen burner until it melts then ignites. Immediately put the burning sulfur into the screw cap container and block the opening. When the burning stops, screw on the screw cap.
3. Add dilute sulfuric acid or hydrochloric acid to sodium sulfite crystals. Collect the gas by upward displacement of air.
Na2SO3 (s) + H2SO4 (l) --> Na2SO4 (aq) + H2O (l) + SO2 (g)
4. Be careful! Add hot concentrated sulfuric acid to copper to form copper (II) sulfate, water, and sulfur dioxide.
Cu (s) + 2H2SO4 (l) --> CuSO4 (aq) + 2H2O (l) + SO2 (g)
3.51.1 Tests for sulfur dioxide
1. Smell
Notice its choking smell.
2. Litmus test
Sulfur dioxide dissolves in water to turn blue litmus paper red. Shake water in a container of sulfur dioxide to form sulfurous acid solution.
3. Burning splint test
Sulfur dioxide extinguishes a burning splint.
4. Potassium permanganate test
Pass sulfur dioxide through a dilute solution of potassium permanganate until it loses its colour.
3.51.2 Reduce potassium manganate (VII) with sulfur dioxide
Add 10 mL of 0.1 M potassium manganate (VII) solution and 10 mL of 3 M sulfuric acid solution to 200 mL of water containing sulfur dioxide. The colour of the manganate ion is lost as the sulfurous acid is oxidized to sulfuric acid. Add a 0.25 M barium chloride solution. The solution becomes milky because of the formation of barium sulfate.
SO2 (g) + H2O (l) --> H2SO3 (aq)
2MnO4- (aq) + 6H+ (aq) + 5SO32- (aq) --> 2Mn2+ (aq) + 3H2O (l) + 5SO42- (aq)
3.51.3 Reduce iron (III) chloride with sulfur dioxide
Pass sulfur dioxide gas through yellow-brown iron (III) chloride solution. Sulfur dioxide reduces iron (III) chloride to Fe2+. The reaction turns red. Pour off some solution and boil. The solution turns green. Add drops of sodium hydroxide solution. The reaction forms a red-brown precipitate of iron (III) hydroxide.
FeCl3 (aq) + 3NaOH (aq) --> Fe(OH)3 (s) + NaCl (aq)
3.51.4 Bleach flowers with sulfur dioxide
See diagram 3.51.4: Bleaching flowers
Add coloured flowers or fruit peel to a solution of sulfur dioxide in water. Replace the stopper and shake. The colour is bleached. Restore the colour by placing the plant material in a dilute solution of hydrogen peroxide.
3.52 Conditions necessary for rusting
Rusting of iron is a multistep process where metallic iron changes to Fe(OH)3.xH2O. Prepare clean nails and put them in test-tubes as follows:
Test-tube 1: Put the nail in the test-tube and half cover with demineralized water. The nail is in contact with air and water.
Test-tube 2: Put anhydrous calcium chloride or silica gel in the bottom of the test-tube. Put a nail in the test-tube and fix put a plug of cotton wool at the mouth. The nail is in contact with air, but not with moisture.
Test-tube 3: Boil water for several minutes to expel dissolved air then pour the water into the test-tube while still hot. Put the nail in the test-tube and half cover with hot water. Put petroleum jelly or olive oil on the surface of the hot water. The petroleum jelly melts and forms an airtight layer then solidifies as the water cools. Half the nail is in contact with water but not with air.
Test-tube 4: Dissolve sodium chloride crystals in water in the test-tube. Put the nail in the test-tube and half cover with salt solution. The nail is in contact with air, water and sodium chloride.
Leave the test-tubes in a rack for several days and note the conditions for rusting. Most rust forms in test-tube 4, then test-tube 1, then test-tube 2 or test-tube 3. Air and water are necessary for rusting. Sodium chloride increases rusting.
3.52.1 The mass of iron and its temperature increases during rusting
1. Balance a piece of iron or steel wool on a knife edge with a brass weight. Leave in moist air for a few days and note the effect of rusting on the weight of the iron. During rusting, metallic iron, Fe, changes to Fe(OH)3.xH2O.
2. Weigh a piece of steel wool. Allow it to rust. Dry it carefully and weigh the steel wool again. The rusted steel wool is heavier than the unrusted steel wool.
3. Invert a measuring cylinder and place it below the surface of the water. Use a bent rubber tube to withdraw air from the cylinder until the volume of air is exactly 100 mL when the water level inside and outside the cylinder are the same. Push some steel wool up to the closed end of the cylinder. Leave the apparatus for a few days and note how much gas is left in the cylinder. Push a lighted taper into the gas and note what happens. Remove the steel wool and examine it. When iron rusts, an increase in weight occurs. The iron combines with oxygen gas from the air. The remaining gas extinguishes a lighted taper placed in it.
4. Almost fill a thermos flask with loosely-packed, damp, steel wool. Place a one-hole stopper fitted with a thermometer into the mouth of the flask. The bulb of the thermometer bulb must be touching the steel wool. Record the initial temperature of the apparatus and note the daily temperatures for five days. The rusting of steel wool is an exothermic reaction.
3.52.2 Oxygen gas combines with iron during rusting
See diagram: 3.52.2: Iron filings rusting
Moisten inside a test-tube with water. Sprinkle iron filings into the test-tube then rotate it horizontally to make the iron filings spread and stick to the walls. Invert the test-tube in a container one third full of water. Support the test-tube so that the water levels inside and outside the test-tube are the same. Mark the levels on the test-tube with a grease pencil and leave for a few days. The iron will rust and the water level will rise inside the test-tube, finally becoming steady. Add water to the container until the levels inside and outside the test-tube are the same and mark the new level. About one fifth of the air volume has been used up, suggesting that oxygen gas has been used up in the rusting of iron. Test the remaining gases with a lighted splint. The lighted splint. is extinguished.
3.52.3 Metals can prevent rusting
Test-tube 1: Put a nail in the test-tube and half cover with demineralized water.
Test-tube 2: Put a nail in the test-tube and half cover with tap water.
Test-tube 3: Wrap a piece of zinc foil around one end of a nail. Put the nail in the test-tube and half cover with tap water.
Test-tube 4: Wrap a piece of tin foil around one end of a nail. Put the nail in the test-tube and half cover with tap water.
Test-tube 5: Wrap a piece of copper wire around one end of a nail. Put the nail in the test-tube and half cover it with tap water.
Note whether zinc, copper or tin best prevents rusting.
3.54 Prepare crystals from solutions
See diagram 3.54: Crystal in supersaturated solution
Seed crystals are very important in the sugar industry. Only experts know how to reduce the crushed sugar cane mixture to a syrup solution then add about 100 g of seed crystals to produce tonnes of sugar crystals.
1. Add sodium thiosulfate crystals, Na2S2O3.5H2O, to 2 cm of water in a small beaker. Heat the beaker and keep adding sodium thiosulfate crystals until no more dissolves. This solution is now a supersaturated solution. Let the beaker cool and note whether crystals form.
2. If no crystals form, add one seed crystal to help crystallization.
3. Tie a cotton thread to a paper clip and suspend the paper clip in the solution. Leave the solution in a warm place. As water evaporates from the solution during the next few days, crystals form, first on the rough edges of the paper clip.
4. Use a magnifying glass to note whether the crystals in the test-tube have the same shape and size as the crystals added to the water in the test-tube.
5. Repeat the experiment by gently heating dry crystals in a test-tube. The crystals dissolve in their own water of hydration to form a saturated solution. Add a seed crystal for recrystallization.
6. Repeat the experiment with crystals of sugar (sucrose), glucose, aluminium potassium sulfate, Al2(SO4)3.K2(SO4).24H2O, ammonium chloride, NH4Cl, ammonium iron (II) sulfate (ferrous ammonium sulfate) (NH4)2SO4.FeSO4.6H2O, iron (II) sulfate, FeSO4.7H2O (add drops of sulfuric acid to prevent the solution turning brown), magnesium sulfate, MgSO4.7H2O, sodium carbonate, Na2CO3.10H2O, sodium sulfate, Na2SO4.10H2O.
3.54.01 Prepare double salt crystals
See: Potassium sodium tartrate
1. Prepare crystals of alum
Add 3.0 grams of alum to 10 ml of warm water and heat with a small flame. Pour the hot solution into an evaporating dish. Observe the crystals forming after about 30 minutes. Leave the solution to evaporate. Wipe off any tiny crystals and keep the large crystals. Suspend the best crystal in a supersaturated solution by a coarse thread. Large crystals of hydrated aluminium potassium sulfate, KAl(SO4)2.12H2O, form.
2. Prepare crystals of ammonia alum
Dissolve 1 cm of ammonium sulfate in 1 cm of water by shaking. In another test-tube dissolve 5 cm of aluminium sulfate in 5 cm of water. Mix the two solutions in an evaporating dish and leave the mixed solutions to evaporate. Diamond-shaped white crystals of the double salt of aluminium, NH4Al(SO4)2·12H2O, deposit in the dish.
3. Prepare crystals of ammonium iron (II) sulfate (ferrous ammonium sulfate)
Dissolve 1 cm of ammonium sulfate in 1 cm of water by shaking. In another test-tube dissolve 2 cm of iron (II) sulfate in 2 cm of cold water. Mix the two solutions in an evaporating dish and leave the mixed solutions to evaporate. Green crystals of (NH4)2SO4 FeSO4.6H2O deposit on the bottom of the dish. The flat square crystals of ammonium iron (II) sulfate do not oxidize or turn yellow when exposed to air. Large crystals can be grown.
4. Prepare crystals of copper ammonium sulfate
Dissolve 1 cm of ammonium sulfate in 1 cm of water by shaking. In another test-tube dissolve 2 cm of copper sulfate in 3 cm of cold water. Mix the two solutions in an evaporating dish and leave the mixed solutions to evaporate. Light blue crystals of CuSO4.(NH4)2SO4.6H2O deposit in the dish.
3.54.1 Prepare sea water crystals
Salinity varies from 32-38 parts per thousand. Sea water contains mainly ions of Na and Cl, also Ca, Mg and K, and dissolved gases, O2, N2 and CO2.
Heat sea water to just below boiling point. When the liquid has decreased by one tenth of its volume, add more sea water and continue heating. When the crystals begin to appear, stir with a glass rod. Push back into the liquid the crystals that appear on the sides of the beaker. The mixture becomes a paste. Heat the paste gently and stir until all the water evaporates and white crystals of salt remain. The salt contains mainly sodium chloride with magnesium chloride and other salts in small quantities. Dry the sea salt and leave it exposed to the air. The salt becomes damp again because the magnesium chloride is deliquescent and attracts water from the atmosphere. Also, the magnesium chloride gives sea salt a slightly bitter taste. Compare the crystals from sea water with crystals from a packet of table salt.
3.54.2 Prepare crystals from a melt
1. Put crystals of naphthalene on a microscope slide. Hold the slide over a flame until the crystals melt. Put a coverslip over the liquid and leave to cool. Watch the crystals grow using a magnifying glass or a microscope. Note whether crystals will grow from several points simultaneously to make boundaries where they meet. Draw the shape of the boundary between the forming crystals and the melt. View the crystals through Polaroid sunglasses.
2. Repeat the experiment with stearic acid (octadecanoic acid).
3.54.3 Prepare crystals with different shapes
See diagram 3.54.3: Different shapes of crystals
Note these crystal shapes by putting 2 drops of the warm concentrated solution on to a microscope slide and view with a magnifying glass or microscope.
Cubic crystals: sodium chloride and potassium chloride
Tetragonal crystals: nickel sulfates, potassium nitrate, zinc sulfate
Monoclinic crystals: a potassium chlorate, sodium sulfate
Triclinic crystals: copper (II) sulfate
Octahedral crystals: sodium chloride crystallizes from alkaline urea or ammonia solution
Funnel-shaped crystals: mixture of sodium chloride and alum solution.
3.54.4 Prepare crystals from a mixture of salts
Dissolve sodium hydrogen sulfate, ammonium iron (II) sulfate, magnesium sulfate and cobalt (II) chloride in just enough water so that the salts dissolve completely only after heating. Let the solution stand undisturbed for several days. Examine the separate crystals that grow from the mixture.
3.54.5 Prepare large crystals
See diagram 3.54.5: Seed crystal in a saturated solution
1. Prepare a saturated solution of potassium alum crystals [Al2(SO4)3.K2(SO4).24H2O] by leaving crystals to dissolve over hours in a warm place. Stir the solution during this time so that it becomes saturated. Pour off the warm saturated solution into a beaker. When the saturated alum solution is cold, some original crystal should remain undissolved. You can also add drops of sulfuric acid or sodium bisulfite solution to keep the solution clear. Suspend a heavy object by cotton in the liquid. Diamond-shaped alum crystals form on the cotton. The crystals have 8 sides but some "corners" may be missing. Detach some well-shaped crystals from the cotton and leave them overnight in the saturated solution to grow. Select the largest crystal, tie cotton around it and suspend the crystal in the saturated solution. After some weeks, the selected crystal can grow too big for the beaker.
2. Repeat the experiment with copper (II) sulfate crystals.
3.54.6 Prepare clusters of crystals
1. Soak pieces of charcoal or brick or unglazed porcelain in a saturated solution of sodium chloride for two weeks. Add ink or dyes then heat the solution to evaporate to dryness. "Blossoms" of crystals form
2. Prepare a saturated solution of potassium dichromate. Immerse cardboard shapes, e.g. a crown, and let the solvent water evaporate. Red crystals form on the shapes.
3. Tie a light weight to a piece of cotton thread and hang the thread in a saturated solution of sodium chloride. Remove the thread and hang it up to dry. Light the hanging thread with a match. The thread burns, but the sodium chloride is left as an ash strong enough to support the light weight.
4. Grow crystal gardens. Drop crystals of metallic salts in a jar of water. Leave the jar undisturbed for weeks and note the growth of "trees" of crystals. Use aluminium potassium sulfate Al2(SO4)3.K2(SO4).24H2O, chromium (III) potassium sulfate (VI)-12-water, copper (II) sulfate-5-water and iron (II) sulfate-7-water.
3.54.7 Prepare split crystals
See diagram 3.54.7: Splitting a crystal with a razor blade
Split crystals of calcite (calcium carbonate) or sodium chloride along their plane of cleavage with a one-sided razor blade or a solid scalpel, i.e. not with a detachable blade. Cleavage is an important identifying property of minerals. Put the edge of the blade on the crystal while holding the face of the blade vertical and parallel to the planes of the two opposite sides of the crystal. Tap with a small hammer on the top of the blade. If the least possible force is used, the crystal splits down the plane of cleavage. When the blade is not directed correctly, the crystal crumbles instead of splitting into two parts. The mineral galena (lead sulfide, PbS) occurs as cubic crystals that are easily split along the three cleavage planes at right angles to each other. Mica has one obvious cleavage, so it can be split into very thin flexible sheets.
3.54.8 Prepare stalactite crystals
See diagram 3.54.8: Stalactite crystals
Half fill two beakers with concentrated sugar solution. Put the beakers on the table about 10 cm apart. Cut a piece of woollen thread about 30 cm long and attach a paper clip to each end of the thread. Put each end of the woollen thread in a beaker with the ends of the thread under the solution. Adjust the distance between the beakers so that the thread dips down between the beakers without touching the table. After some days crystals appear on the lowest part of the thread between the beakers. The saturated solution has moved along the thread by capillarity until reaching the lowest point on the thread. Here water is lost by evaporation and crystals like stalactites form.
3.55 Brownian movement
This observation was named after the Scottish naturalist Robert Brown, the botanist employed by Sir Joseph Banks on the voyage of the Investigator to Australia in 1801. He published his observations in 1827. He had noticed the jittery random movement of pollen grains thought to be caused by the pollen grains themselves. However, he showed that similar random movement could also be seen if any solid material of similar size was observed. The movement is caused by an instantaneous imbalance of combined forces of solution molecules on much larger irregularly-shaped solute particles. In 1905 Albert Einstein was fascinated by this movement. He wrote: "The existence of a never diminishing motion seems contrary to all experience. This difficulty was splendidly clarified by the kinetic theory of matter." His statistical calculations of the mass and number of molecules involved lead to the final acceptance of the atomic theory by scientists.
1. Put a drop of toothpaste on a microscope slide and stir water into it until it is almost colourless. Put a coverslip over the drop and examine it with a microscope under high power with the stage illuminated from the side. Pay attention to one particle. At first it appears to stay in one place. Later you can see that it is moving with an irregular jerky movement in all directions. The irregular movement is caused by molecules of water hitting unevenly on the sides of the particle. This experiment does not work with the modern "gel" type of toothpaste.
2. Make a model using a tray containing many small, light beads and one large marble. The small beads represent molecules of water and the large marble represents a particle of suspended graphite. Shake the tray and note how the small beads hit the marble from all directions. The forces tend to cancel so the large marble may just make very small irregular movements but returns to the same place.
3. Observe dirty water in sunlight. Fill a beaker with tap water and focus sunlight into the beaker with a magnifying glass. You may see the motions of suspended particles of solid matter.
4. Observe smoke in still air. The smoke particles suspended in the air show Brownian movement.
3.55.1 Diffusion of heavier than air gas, carbon dioxide
See diagram 3.55.1: Diffusion of heavy carbon dioxide gas upwards | See diagram 3.34.1: Limewater test for carbon dioxide
Fill a jar with carbon dioxide and invert it over a similar jar full of air. After a few moments separate the jars, pour a little limewater in the lower one and shake it. The limewater will turn milky indicating that the carbon dioxide has fallen into the lower jar because it is the heavier gas.
Repeat the experiment with the carbon dioxide in the lower jar and invert a jar of air on top of it. If the jars are left for 5 minutes carbon dioxide will be carried into the upper jar by diffusion, in the same way air will be carried into the lower jar. The limewater test will show the presence of carbon dioxide in the upper jar.
3.55.3 Diffusion of liquids
1. Put a crystal of potassium dichromate or ammonium dichromate at the bottom of a beaker of water by dropping a crystal down a delivery tube. Remove the delivery tube. The colour of the dissolving crystal spreads through the water.
2. Fill a small bottle with a potassium manganate (VII) solution. Place the open bottle in a larger container. Carefully fill the larger container by pouring water down the side until the water level is above the top of the small bottle. After a few days the potassium manganate (VII) solution diffuses evenly through the water.
3.56 Particles of matter and dilution
Put a crystal of potassium manganate (VII) in a test-tube. Add 1 mL of water and fix a stopper. Shake the test-tube to dissolve the crystal. Add water to a total volume of 10 mL. This is a "10 times" dilution. Pour this 10 mL of purple solution into a 100 mL container and then fill the container with water. This is now "100 times" dilution. Fill the 10 mL test-tube with this solution and throw the rest away. Dilute this again in the container to 100 mL. It is now a "1 000 times" dilution. Note how often the solution can be diluted by a factor of 10 before the colour is so pale that it is only just visible. The final dilution factor shows that if matter is particulate, the size of the particles must be very small.
3.57 Size of a molecule
See diagram 3.57: Oil layer and powder on surface of water
Oil floats on the surface of water as a one molecule thick layer, if allowed to spread out freely.
1. Use a tray with a glass bottom. Put graph paper under the tray. Put water in the tray. Select a suitable oil to pour on the water, e.g. 1% oleic acid / methylated spirit solution, petroleum distillate. Use a burette to find the volume of 100 drops of oil, then calculate the volume of 1 drop of oil. Sprinkle the surface of the water with a very fine light powder, e.g. talcum powder or Lycopodium powder. Let 1 drop of oil fall from the burette onto the water. The oil spreads out over the surface of the water but must not touch the sides of the tray. The oil on the water pushes the powder aside so you can easily see the area covered by the oil. Look down on the graph paper to measure the approximate area over which it spreads. The volume of oil put on the water = area of the oil on the water X thickness of the oil layer. The approximate dimension of a single molecule of oil is 1 X 10-7 m.
2. Another way to calculate the volume of a drop of oil is to let the drop fall on a piece of flat plastic. Touch the oil with the point of a glass rod then touch the water surface. Oil leaves the glass rod and spreads over the water. Use graph paper to measure the approximate area over which the oil spreads. To calculate the volume of oil placed on the water, keep using the glass rod to remove successive fractions from the flat plastic until no oil remains.
3.58 Clay soil suspension
Shake clay soil with water and leave it to settle. Note the humus layer at the surface and particles of rock and mineral at the bottom. Filter the liquid. The filtrate is still cloudy because clay particles have passed through the filter paper. Suspension particles may take days to settle but they will settle eventually. Colloid particle size is about 10-7to 10-5 cm. Divide the filtrate into two parts in test-tubes. Keep one part as a control. To the other test-tube add drops of barium chloride solution or an aluminium salt solution. Note whether the filtrate remains cloudy. The same effect may occur when a clay suspension in a river meets the salts contained in sea water.
3.61 Tin-lead alloys
See diagram 3.61: Casting mould
Make a casting mould by drilling out the thread of a nut to leave a smooth hole of about 0.6 cm diameter and then cut the nut into two halves with a hacksaw. Use wire to bind the two halves together for casting then put this caste on sand. Pure tin melts at 232oC and pure lead melts at 327oC. Weigh out pieces of lead and tin to make four alloys so that the percentage of tin by weight is 20% tin, 40% tin, 60% tin and 80% tin. Put each mixture of lead and tin in a crucible or Pyrex test-tube. Cover each mixture with powdered charcoal to prevent oxidation of the metals then heat with a Bunsen burner until they melt. Stir the melt with a wood splint to help the metals dissolve. Pour each mixture of molten metal into the mould until it is full. Be careful! Hold back the carbon from the charcoal with a wooden splint while pouring. When the cast alloy is cool, knock away the two halves of the nut.
3.62 Tin-lead alloys and pure metals, hardness
See diagram 3.62: Hardness test apparatus
Test the hardness of the four lead-tin alloys and two pure metals, lead and tin. Use a metal punch with a pointed end and a 1 metre plastic tube to guide the punch as it falls on to the alloy and makes a small hole. The softer the alloy the larger the hole. Measure the diameters of the holes with vernier callipers and a magnifying glass. The pure metals should be less hard than the alloys. The 60% tin alloy should be the hardest alloy. This test is a kind of dynamic hardness test, e.g. Vickers hardness test. Geologists test the hardness of minerals with a scratch hardness test, Mohs' test.
3.63 Melting points of metals and alloys
See diagram 3.63: Melting point apparatus
Prepare a metal plate from a 12 cm X 12 cm piece of iron, 0.2 to 0.4 cm thick. Draw a hexagon on the metal plate then drill a small equal depth depressions at each corner of the hexagon. Drill holes through the four corners of the metal plate. Thread wire through the four holes and suspend the metal plate horizontally. Pour a few globules of four alloys and the two pure metals into separate porcelain bowls. Be careful! Put one pellet of each alloy or metal into a depression on the metal plate. Heat the middle of the metal plate with a Bunsen burner. Touch the pellets with a wood splint to check when they melt. When all the pellets are all molten, use the wooden splint to remove excess molten metal from bigger pellets so that they are all the same size. Remove the Bunsen burner flame, leave the metal plate to cool and note the time to form crystals. Make a table of time to crystallize and plot the results on graph paper. Pure lead solidifies first, then 20% tin, then 40% tin, then 60% tin. The alloy that takes the longest time to solidify has the lowest melting point. The 60% tin alloy should have the lowest melting point.
3.64 Heat treatment of steel needles, annealing, quenching, tempering
1. Annealing is used to produce a soft state in worked metals. Heat a needle to bright red heat. Hold it vertically in the flame and then take one minute to raise it slowly out of the flame. Leave to cool. Try to bend the needle with a pair of pliers. The needle should now be soft. You can easily bend it around a pencil.
2. Quenching is used to make steel metals harder and non-ferrous metals softer, e.g. copper. Heat a needle to bright red heat and immediately plunge it into cold water. Try to bend the needle with a pair of pliers. The needle should now be brittle. You can easily break it into small pieces.
3. Tempering of steel is reheating after rapid cooling to give extra secondary harness. Heat a needle to bright red heat and immediately plunge it into cold water.
Use 5 cm sewing needles that are tough and springy and difficult to bend. They are made of an alloy of iron with a small proportion of carbon. Clean and shine the surface of the needle with emery cloth. Heat the needle very gently until a deep blue oxide film appears on the surface. This colour indicates the tempering temperature of the needle. Leave to cool. Try to bend the needle with a pair of pliers. The needle is tough and springy again.
3.65 Strengths of mud, clay and sand bricks
See diagram 3.65: Testing the strength of a brick
Find a source of clay soil or mud. If it is dry, it must be mixed with water. To do this, put about 350 mL of water in a suitable container such as a plastic bowl. Crush the dry clay to a powder and then mix it with water until a thick smooth paste forms. Squeeze it through your fingers until no lumps remain. It will have the correct consistency when it is thick and pliable and sticks more to itself than to your fingers. Spread the clay or mud on to a flat surface very evenly to make a slab of 1.5 cm thickness. Use a clean wet knife to cut four bricks, each 10 cm by 5 cm. Dry one under the sun for two or three days and bake another by a fire. Try making a sand brick of the same size. Also, obtain a brick sold by a building contractor. Examine the bricks for cracks. Test whether the surface comes away by rubbing with a dry finger. Test whether the surface comes away by rubbing with a wet finger. Test the strength of the small 5 X 10 X 1.5 cm bricks. Support the two ends of the brick on the edges of two tables. Load the middle of the test brick with weights or attach a bucket into which sand can be poured. Keep loading until the test brick breaks. Suspend the weights and bucket near the floor so that they have almost no distance to fall.
3.66 Make bricks with cement
See diagram 3.66: Cardboard box for cement test
Make 5 boxes out of stiff paper or cardboard 1.5 cm deep, 5 cm wide and 10 cm long. Use adhesive tape or clips to fasten the edges. A cement brick, the same size as the clay bricks, can be cast in these boxes. Smear a little oil or grease around the inside surfaces of the boxes. Obtain some fresh Portland cement from a builder.
1. Cement / water brick: Mix the cement with water to a thick paste and fill the box with it, smoothing off the top surface level with the paper. It should "set" in a few minutes but it will take a few days to "harden". "Setting" is to change from a fluid to a firm rigid material but a mark can still be scratched on the surface with a nail. To "harden" is to become rock hard.
2. Cement / sand / water brick: Mix 1 part of cement powder with 3 parts of clean sand. Work into a thick paste with water. Pour into the paper box, smooth off the surface and leave to set and harden.
3. Cement / sand / gravel / water brick: Make a brick as before using 1 part cement powder, 1 part of sand, 3 parts clean gravel and water. Cast the brick and leave to set and harden. This is a concrete brick.
4. Cement / lime / sand / water brick: A builder buys quicklime and mixes this with water to make calcium hydroxide on the building site just before he uses it. Mix 1 part of cement, 5 parts of builders' lime, calcium hydroxide, and 2 parts of sand and make into a paste with water. Cast a brick as before and leave to harden.
3.66.1 Change in weight of setting cement
See 34.3.01: Portland cement
Weigh 500 g of sand and cement mixture (commercial mortar mix) into a polystyrene drink cup and add 75 grams of water. Mix the contents until all the lumps are gone. Weigh the polystyrene cup and contents again to check the weight of the added amount of water. Fill another polystyrene cup with water to the same level and weigh the cup + water. Leave the polystyrene cups for one day then weigh them again. The loss in weight of the cup + water only shows the loss by evaporation of the cup + cement mixture + water. The rough surface area of the setting concrete does allow water to evaporate faster than in cup + water only. However, the loss by evaporation is negligible. The experiment shows that most of the added water is absorbed in the chemical reaction of the setting cement.
3.66.2 Strength of cement with changing water content
Wrap the waste cement in newspaper then put it in waste containers. Do not pour cement paste down the sink. Use identical cardboard milk cartons for moulds.
1.1 Put 200 mL of a mixture of dry cement and sand in a large beaker. Slowly add 100 mL of water from a measuring cylinder, with stirring, to the mixture until it becomes a thick paste. Pour the paste into cardboard mould 1. Smooth the surface of the cement in the mould. Wipe out the beaker with paper and rinse with water. Record the volume of water used.
1.2 Repeat the experiment with 20% less water. Pour the paste into cardboard mould 2.
1.3 Repeat the experiment with another 20% less in water. Pour the paste into cardboard mould 3.
1.4 Repeat again with 20% more water than the mixture in mould 1. Pour the paste into cardboard mould 4.
1.5 Repeat 1. Pour the paste into cardboard mould 5.
Cover the cardboard moulds 1 to 4 with plastic wrap to prevent evaporation. Leave mould 5 uncovered. Leave all the moulds in a warm place for 2 days.
2. Examine the mixtures:
2.1 Note the surfaces.
2.2 Scratch the surfaces with your fingernail, a nail and the point of a file.
2.3 Drop a steel ball from the same height on the surfaces, while wearing safety glasses, and note the bounce height. The harder the surface, the greater the bounce height.
2.4 Remove the cardboard and use a hammer to hit each mixture with increasing intensity until it breaks. Wear safety glasses when you do this.
2.5. Record the order of surface hardness by both methods and the resistance to breaking. Note the relative hardness and the volume of water used. Note the relative hardness of mould 1 and mould 5.
3. Repeat the experiment with the ratio of sand to cement from 50 mL of sand + 150 mL of cement, to 50 mL of cement + 150 mL of sand.
3.66.3 Alkalinity of concrete
Concrete is an artificial stone used as a building material. It contains cement, sand, water and an aggregate, crushed stone or slag, a mixture of oxides formed during ore smelting and refining. Reinforced concrete uses steel bars, twist bars, or cables to counteract weakness in tension.
Alkaline cement protects steel reinforcing rods in concrete from corrosion. Clean pieces of steel reinforcing or nails with sandpaper and put them into two jars half filled with water. Put broken pieces of concrete in one jar. After a week, note that the steel in the jar without the concrete corrodes faster. Carbon dioxide from the atmosphere slowly penetrates the surface of concrete and reacts with lime, Ca(OH)2, to convert it to limestone, CaCO3, reducing the alkalinity of the concrete touching the steel bars. The steel can form rust containing iron oxides and hydroxides that have a larger volume than iron. This expansion cracks the concrete.
Find a broken piece of old exposed concrete. Break it and wet the new surface with phenolphthalein indicator solution. A pink coloration indicates the high alkalinity inside the concrete with a rim of untinted concrete around the edge.
3.67 Strength of plaster of Paris
Plaster of Paris is hydrated calcium sulfate, CaSO4.2H2O. When mixed into a paste with water it sets quickly and expands. It is used as a fine casting material. Put 4 mL of water into a beaker. Add the powdered plaster of Paris slowly with a spatula. Continue adding the plaster until it just appears above the surface of the water. The plaster absorbs the water and you should finish with a very thin layer of water, about 1 mm, above the plaster. Stir the mixture well. When it begins to thicken, pour it into the paper box. Smooth the surface of the cement in the mould and leave to set for 1 day. Investigate the surface and strength of these bricks as with the mud and clay bricks. Plaster of Paris is not often used as a construction material, but calcium sulfate as gypsum, CaSO4.2H2O, is used to prepare Portland cement.