School Science Lessons
Chemistry experiments
Please send comments to: J.Elfick@uq.edu.au
2012-05-20 SPwp
Table of contents
3.66 Cement
3.61 Construction materials
3.55 Matter as particles
3.32 Prepare gases
3.1.0 Prepare crystals
3.61 Construction materials
3.64 Heat treatment of steel needles, annealing, quenching, tempering
3.63 Melting points of metals and alloys
3.68 Putty
3.65 Strength of mud, clay and sand bricks
3.67 Strength of plaster of Paris
3.61 Tin-lead alloys, make with a casting mould
3.62 Tin-lead alloys and pure metals, hardness
3.66.0 Cement
3.66.3 Alkalinity of concrete
3.66.1 Change in weight of setting cement
35.24.2 Make artificial rocks, sedimentary rocks
3.66.5 Make bricks with cement
3.66.4 Make mortar
3.66.6 Portland cement
3.66.2 Strength of cement with changing water content
3.55 Matter as particles
3.55 Brownian movement
3.58 Clay soil suspension
10.1.2 Diffusion of ammonia and hydrogen chloride gases
3.55.1 Diffusion of heavier than air gas, carbon dioxide
3.55.3 Diffusion of liquids
3.56 Particles of matter and dilution
3.57 Size of a molecule
3.32 Prepare gases
3.33.0 Prepare ammonia
6.1.1d Prepare butane
3.34.1 Prepare carbon dioxide
13.4.1.0 Prepare chlorine
16.1.14 Prepare chloroform, (trichloromethane)
10.6.5 Prepare gases by destructive distillation of
coal
10.6.5.1
Prepare gases from wood
3.32.0 Prepare gases with a gas generation apparatus
41.0 Prepare hydrogen gas
3.41.2 Prepare hydrogen gas bubbles
16.1.1a Prepare methane gas
3.44 Prepare nitric oxide (nitrogen monoxide, NO)
3.47 Prepare nitrogen dioxide, {nitrogen (IV) oxide, NO2}
3.46 Prepare nitrogen gas with ammonium nitrite
3.45 Prepare nitrous oxide (dinitrogen oxide, N2O)
3.49.0 Prepare oxygen gas with hydrogen peroxide
12.12.8 Prepare oxygen gas with household bleach, sodium hypochlorite, or bleaching powder
3.51.0 Prepare sulfur dioxide
3.44 Prepare nitric oxide (nitrogen
monoxide, NO)
See diagram 3.44: Prepare nitrogen monoxide |
See diagram 1.13a: Simple fume hood
Do these experiments in a fume cupboard, fume hood, or use small quantities
near an open window or a well-ventilated area. Use eye and skin protection
when splashes are possible. Nitrogen monoxide, NO, is a colourless toxic
gas that reacts immediately with oxygen gas in the air to form brown
fumes of nitrogen dioxide, NO2. Nitric oxide has many physiological
functions in the body, e.g. dilation of blood vessels.
1. Add drops of dilute nitric acid to copper. Nitrogen oxide forms that
immediately reacts with oxygen in the air to form nitrogen dioxide.
2. 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. Repeat the experiment with cold 7M nitric acid.
4. This 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 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 nitrous oxide
(dinitrogen oxide, N2O)
(nitrous oxide, dinitrogen oxide, 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
relights the splint.
3.46 Prepare nitrogen gas with ammonium nitrite
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, {nitrogen (IV) oxide, NO2}
Nitrogen dioxide is a brown gas with a choking smell that may irritate
the lungs and lead to death. It is commonly formed by the reaction of concentrated
nitric acid with copper and the decomposition of metal nitrates and nitrites.
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.
1. 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)
2. Heat lead (II) nitrate crystals. The decomposition may be noisy.
Nitrogen dioxide and oxygen form, leaving yellow lead oxide.
Pb(NO3)2 (s) --> 4NO2 (g) + 2PbO
(s) + O2 (g)
lead (II) nitrate --> nitrogen dioxide + lead oxide + oxygen
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. brown-red 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
3.49.0 Prepare oxygen gas with hydrogen peroxide
See diagram 3.2.35: 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.
1. The safest method to prepare oxygen is by decomposition of hydrogen peroxide
solution.
2H2O2 --> 2H2O + O2 (g)
2. Pour dilute hydrogen peroxide into a measuring cylinder. Add drops of
detergent. Add manganese (IV) oxide (manganese dioxide) powder. The reaction
forms oxygen gas as a foam of bubbles. Use the oxygen foam for combustion
experiments with burning twine, burning iron wire and burning magnesium.
Test the gas in the space above the liquid.
3. To 3 mL of 6% hydrogen peroxide solution, add 1. Powdered manganese
dioxide, MnO2, 2. 0.5 mL of 1 M FeCl3, 3. 20 grains
of active dry yeast.
Test the oxygen formed by these catalytic reactions with a glowing splint.
2H2O2 --> 2H2O + O2 (g)
or
H2O2 (aq) --> H2O (l) + 1/2O2
(g)
4. 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]
5. 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.
6. 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. Hold a glowing splint above the top
of the test-tube while continuing the heating. The glowing splint relights.
12.12.8 Prepare oxygen gas with household bleach, sodium hypochlorite, 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.
3.49a Hydrogen peroxide concentration
and storage
Hydrogen peroxide 50% w/w solution
Hydrogen peroxide, 100 volume, vols, 30% w/w solution
Hydrogen peroxide, 8-20%, Hydrogen peroxide, 40 volume, vols, 12% w/v
solution
Hydrogen peroxide, 5-8%, Hydrogen peroxide, 20 volume, vols, 6% w/v solution
(hair and teeth bleach, antiseptic)
Hydrogen peroxide, < 5%, Hydrogen peroxide, 10 volume, vols, 3% w/v
solution
Hydrogen peroxide may be sold in pharmacies 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
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. Hydrogen
peroxide is used as an antiseptic where catalase enzyme in blood catalyses
the decomposition of hydrogen peroxide to water and oxygen.
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 black-grey
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, Highly
toxic, pale blue gas, with a distinct pungent odour, like weak chlorine.
It is a powerful oxidizing agent that irritates the lungs. Ozone is formed
in the atmosphere by the action of ultraviolet light with oxygen to form
the ozone layer. You may smell ozone in underground train tunnels if electric
sparks had occurred between the power rail and the electrical pickup shoe.
Small concentrations of ozone are used as an air freshener / sanitizer
in public facilities and rest rooms to destroy atmospheric germs and odours.
Ozone may be smelt around photocopiers, e.g. Rank Xerox, and laser printers,
if they are not well-ventilated. However, manufacturers of these
products claim that ozone can be smelt in concentrations as little as one
part in 500,000, while the threshold limit value for short term exposure
is 0.3 parts per million.
See diagram 3.50: Prepare ozone
1. Prepare ozone in the laboratory using a high voltage induction coil,
spark coil (an ozonizer). However, the experiment is dangerous. Do not
attempt it in a school laboratory.
3O2 (g) --> 2O3 (g)
2. Use electrolysis of 4 M sulfuric acid with carbon anodes. Pass 12 volts
through the circuit and note the smell of ozone at the anode.
3. 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
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)
3.51.0 Prepare sulfur dioxide,
SO2
See diagram 3.2.75.1: Prepare sulfur
dioxide by burning | See diagram 3.2.75.2:
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, (Na2SO3),
crystals. Collect the gas by upward displacement of air.
Na2SO3 (s) + H2SO4 (l) -->
Na2SO4 (aq) + H2O (l) + SO2
(g)
Na2SO3 (s) + 2HCl (l) --> 2NaCl (aq) + H2O
(l) + SO2 (g)
4. Add dilute sulfuric acid or hydrochloric acid to sodium metabisulfite,
(Na2S2O5), crystals. Collect the gas by upward displacement
of air.
Na2S2O5 (s) + 2HCl (l) --> 2NaCl (aq) + H2O
(l) + 2SO2 (g)
5. 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)
6. Prepare sulfur dioxide
with sulfuric acid and sodium sulfite
See diagram 13.13.3: Prepare sulfur dioxide
Na2SO3 (s) + H2SO4 (l)
--> Na2SO4 (aq) + H2O (l) + SO2
(g)
7. Prepare sulfur dioxide with sulfuric acid
and copper
Add hot concentrated sulfuric acid to copper to form copper (II)
sulfate, water and sulfur dioxide.
BE CAREFUL!
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 0.1M potassium manganate (VII) solution and 10 mL of 3M dilute
sulfuric acid solution to 200 mL of water containing sulfur dioxide.
The solution will gradually become colourless as the sulfur dioxide reacts
with the potassium permanganate. The colour of the manganate ion is lost
as the sulfurous acid is oxidized to sulfuric acid. Add a 0.25M barium
chloride solution when 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 brown-yellow 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 brown-red 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 3.2.76: 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.
3.54 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 aluminium potassium sulfate
[Al2(SO4)3.K2(SO4).24H2O],
ammonium chloride (NH4Cl), ammonium iron (II) sulfate [ferrous
ammonium sulfate, (NH4)2SO4.FeSO4.6H2O],
glucose (C6H12O6), 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), sucrose (C12H22O11).
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, will
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 seawater crystals, composition of seawater
Salinity varies from 32-38 parts per thousand. seawater contains mainly
ions of Na and Cl, also Ca, Mg and K, and dissolved gases, O2,
N2 and CO2.
Heat seawater to just below boiling point. When the liquid has decreased
by one tenth of its volume, add more seawater 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 seawater with crystals from a packet of table salt.
A typical composition of seawater:
Sodium chloride 77.74%
Magnesium chloride 10.9%
Magnesium sulfate 4.7%
Calcium sulfate 3.6%
Potassium sulfate 2.5%
Calcium carbonate 0.34%
Magnesium bromide 0.22%
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 3.2.47: 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 3.2.49: 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 crystal clusters
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 3.2.51: 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 or baking soda or Epsom
salts solution. Put the beakers on the table about 10 cm apart. Cut a piece
of woollen thread, or any thread that will soak up water, about 30
cm long and attach a paper clip to each end of the thread. Put each end of
the woollen thread in one of the beakers 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. Place a saucer
or Petri dish below the dipped down thread to catch any drips. 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. Leave the experiment in place for weeks. A stalactite
grows down from the thread and a stalagmite grows up from the saucer. After
a long time the stalactite and stalagmite will join to form a column. Remember
that the stalactite must hold "tight" to the thread and the stalagmite "might"
grow up to meet the stalactite to form a column.
3.54.9 Prepare crystal blossoms
Soak pieces of charcoal, brick or unglazed porcelain in a saturated
solution of sodium chloride. Keep the pieces covered by adding more saturated
sodium chloride solution over a period of two weeks. At this point mix some
Prussian blue dye or ink with the sodium chloride and add this to the pieces
of charcoal. Then leave to evaporate to dryness. Blossoms of crystals will
form. A variety of colours may be produced by adding different dye compounds.
3.54.10 Prepare crystal crowns
See diagram 3.2.50: Crystal crown
A crystal crown. Cut a crown from a piece of tin taken from a fruit can.
Fasten it with a piece of wire as shown in the diagram. Wrap the crown with
strips of cotton cloth. Dip the whole crown into a solution of potassium
dichromate and then leave to dry. Seed crystals will form on the cloth. Prepare
a saturated solution of potassium dichromate at 800 C and immerse the crown
for a day or so in this saturated solution. Red crystals should form and
make a beautiful display on the crown. If the crown is small, only a small
amount of potassium dichromate will be required.
3.55 Brownian movement
This observation was named after the Scottish naturalist Robert Brown,
the botanist employed by Sir Joseph Banks (1743-1820) on the voyage of
the Investigator to Australia in 1801. Robert Brown 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.
5. Use a microscope to observe the movement of particles in a smoke cell.
3.55.1 Diffusion of heavier
than air gas, carbon dioxide
See diagram 3.2.53: 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 3.2.57: 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 (the "dragon's breath" used in
Chinese fireworks). 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 × thickness of the oil layer.
The approximate dimension of a single molecule of oil is 1 × 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 10-7
to 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 seawater.
3.61 Tin-lead alloys, make with
a casting mould
See diagram 3.2.61: Casting mould made from
a nut
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 3.2.62: Hardness test apparatus
Test the hardness of the four tin-lead 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 3.2.63: Melting point apparatus
Prepare a metal plate from a 12 cm × 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 Strength of mud, clay
and sand bricks
See diagram 3.2.65: 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 × 10 × 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.
3.66.0 Cement
Cement is any material that binds loose sediment into a rock and may
be ferruginous (containing iron), calcareous (containing calcium) and siliceous
(containing silica). Builders' cement contains calcium and aluminium silicates.
Concrete contains aggregate
(gravel and sand), cement, and water. Concrete can be cast into shape
to become load bearing.
Mortar contains sand, cement and water and is
used for plasters.
Grouts contain cement and water and are used to fill
gaps.
3.66.1 Change in weight of
setting 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.
3.66.4 Make mortar
Use 5 mL of slaked lime and 20 mL of clean sand. Wash sea sand four times
with water to get rid of the salty impurities. Put the slaked lime into an
old cup and make it into a paste with water. Stir in the sand at a time, adding
more water as needed, until a stiff paste forms. Scrape out the paste on
to a tin lid and leave it for a day or two. It will set into a hard mass.
3.66.5 Make bricks with cement
See diagram 3.2.66: 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 purchases 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.
3.66.6 Portland cement
Portland cement hardens as it reacts with water. It was thought to have
the same colour as stone on Isle of Portland, U.K. Portland cement is a fine powder produced by grinding Portland cement
clinker and some gypsum. The raw mixture is mainly chalk or limestone containing
clay or silicon dioxide and other materials, including clay, shale, sand,
iron ore, bauxite, flies ash and slag, i.e., minerals containing calcium
oxide, silicon oxide, calcium aluminate, aluminium oxide, ferric oxide, and
magnesium oxide. Calcium and
silicon form the calcium silicates that give
strength to the concrete. Aluminium and iron compounds produce the liquid
solvent flux in the kiln that helps in the formation of silicates at a conveniently
low temperature. The raw mixture is heated in a cement kiln
at 1400-1450
oC so that the ingredients become sintered, i.e. about one third
melted, but not fused into a molten mass. It cools to become grey-coloured
clinker containing at least two thirds by weight of calcium silicates. Calcium
sulfate as gypsum is added
to the clinker. The gypsum hydrates very rapidly
during the concrete setting reaction and helps to control the initial setting
rate. The mixture is ground to form fine cement powder that can be stored
dry and later mixed with water to form an alkaline cement workable slurry
for casting.
Portland cement powder may contain 50% tricalcium silicate, 3(CaO).SiO2,
25% dicalcium silicate, 2(CaO).SiO2, 10% tricalcium aluminate,
3(CaO).Al2O3, 10% tetracalcium aluminoferrite, 4(CaO)4.Al2O3.Fe2O3,
and 5% gypsum, CaSO4.2H2O. So Portland cement contains
approximately 65%, calcium oxide, CaO 25%, silicon oxide, SiO2,
5% aluminium oxide, Al2O3, 1% ferric oxide, Fe2O3
and 4% calcium sulfate, CaSO4. Different types of cement contain
the same four major compounds that make up at least 90% of the total weight,
but in different proportions.
Tricalcium silicate + Water (yields) --> Calcium silicate hydrate + Calcium
hydroxide + heat
When water is added to concrete powder, hydration occurs and during
this chemical reaction the concrete gradually hardens as calcium silicate
hydrate gel that forms in the first few days at the surface and later deeper
in the pour. The strength of hard
concrete comes from the solid part of the
paste, the calcium silicate hydrate and other crystalline phases. The pores
remaining in hard concrete are filled with water and air and have no strength.
dicalcium silicate and dicalcium silicate + water --> calcium silicate
hydrate + calcium hydroxide + heat
The volume of setting concrete should not change because the added
water should be used up in the hydration process.
So the weight of cement
powder + water + aggregate = weight of the set concrete block (conservation
of mass).
The water-cement ratio (by weight) of completely hydrated cement
is 0.22 to 0.25, excluding evaporable water. So the warning "Do not touch
wet concrete until it dries" is inaccurate because nearly all the water
is lost of the hydration reaction, not by evaporation. The rate of reaction
of the cement with water is proportional to the surface area of the particles.
Cement production requires high energy input and produces large quantities
of carbon dioxide, so it contributes to global warming. However, EMC, Energetically
Modified Cement, uses very finely ground ingredients that have increased
surface area for the chemical reaction and uses less energy to produce it.
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 so 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.
3.68 Putty
Putty, glazier's putty, painter's putty, calcium carbonate paste, (whiting)
+ linseed oil, (+ white lead), filler in glazing, sealing glass into frames
"Silly putty", silicone, bouncing putty (Dow Corning 3179 dilatant compound)
"Tricky Putty" (to amuse young children): 7.2.6