School Science Lessons
Chemistry
Updated: 2012-01-24
Please send comments to: J.Elfick@uq.edu.au
Table of contents
3.31.0 Hygroscopic, deliquescent and efflorescent
chemicals
3.2.0 Identify pure substances
3.18.0 Separate substances from mixtures
3.9.0 Solubility and solutions
3.28.0 Substances that gain mass when heated
3.30.0 Substances that lose mass when heated
3.31.0 Hygroscopic, deliquescent
and efflorescent chemicals
3.31 Hygroscopic, deliquescent and efflorescent
chemicals
3.31.1 Expose different salts to
the air
3.31.2 Expose sodium carbonate decahydrate,
washing soda, to the air
3.31.3 Tests for water with cobalt
(II) chloride
3.2.0 Identify pure substances
3.6 Boiling point of inflammable liquids, ethanol, acetone
3.5.1 Boiling point of sodium chloride solution
3.5 Boiling point of water
3.2 Melting point of naphthalene
3.3 Melting point of naphthalene using a capillary tube
3.4 Melting points of substances changed by impurities
3.8 Pressure of the atmosphere affects the boiling point
3.7 Volatility of liquids
3.18.0 Separate substances
from mixtures
3.20 Separate by distillation
3.21 Separate by fractional distillation of crude
oil
3.19 Separate by sublimation, iodine
3.25 Separate gases dissolved in a water sample
3.26 Separate immiscible liquids of different density
3.24.1 Separate mixed inks with paper chromatography
3.24 Separate pigments from green leaves with paper
chromatography
3.22 Separate salt and sand
3.27 Separate solids using density difference
3.18 Separate tin from a mixture of tin and carbon
(charcoal)
3.23 Solvent extraction of oil from nuts
3.9.0 Solubility and solutions
3.17 Heat of solution, sodium hydroxide, potassium
hydroxide, ammonium chloride, urea
3.17.1 "Magnetic" sugar cube dissolves
3.16 Miscible liquids, methylated spirit, glycerine,
kerosene
3.9 Solubility in water of different salts
3.10 Solubility and temperature, solubility of salts
in water
3.11 Solubility of substances in water at a given temperature
3.12 Solubility and particle size, sodium chloride,
copper (II) sulfate
3.13 Solubility and solvents, sodium chloride, methylated
spirit
3.14 Solubility and agitation, cane sugar
3.15 Volume of solutions, sodium chloride
3.28.0 Substances that gain
mass when heated
3.28.0 Substances that gain mass when heated, copper foil
3.28.4 Collect and weigh the gaseous products of
a burning candle
3.28.3 Substances that gain mass when heated, prepare
calcium oxide
3.28.2 Substances that gain mass when heated, prepare
lithium oxide
3.28.1 Substances that gain mass when heated, prepare
magnesium oxide
3.30.0 Substances that lose
mass when heated
3.30.0 Substances that decompose and lose mass when heated, thermal
decomposition
3.30.15 Decomposition of ammonium chloride, NH4Cl
3.30.14 Decomposition of ammonium dichromate, (NH4)Cr2O7
3.30.9 Decomposition of boric acid, H3BO3
3.30.1 Decomposition of carbonates
3.30.2 Decomposition of hydrogen carbonates (bicarbonates)
3.30.8 Decomposition of hydrous salts
3.30.3 Decomposition of hydroxides
3.30.4 Decomposition of nitrates
3.30.10 Decomposition of oxalic acid
3.30.5 Decomposition of oxides
3.30.11.2 Decomposition of phosphates
3.30.11.1 Decomposition of potassium ferricyanide
3.30.6 Decomposition of sulfates
3.30.7 Decomposition of sulfites
3.30.13 Heat silica sand
3.30.11 Potassium chlorate, KClO3
3.30.12 Sublimation, iodine, camphor, naphthalene
3.2 Melting point of naphthalene
See diagram 3.2.2: An approximate melting point
Put 2 cm of naphthalene flakes in a test-tube. Hold a thermometer with
its bulb in the naphthalene. Use a small flame to heat the test-tube gently
and watch the thermometer reading. To find the melting range, note the temperature
when the naphthalene melts. Leave to cool and note the temperature when the
naphthalene solidifies. To find the melting point, calculate the average of
these two values. The melting point of pure naphthalene is 80.5oC.
3.3 Melting point of naphthalene
using a capillary tube
See diagram 3.2.3: A more accurate method
Make a capillary tube by drawing out a glass tube over a hot flame. Put
a very small amount of naphthalene in a capillary tube sealed at one end.
Attach a thermometer to the capillary tube, a sealed end down. Put the thermometer
and capillary tube in a container of water and slowly heat the water while
stirring with the thermometer and capillary tube. Do not let water enter
the capillary tube. To find the melting range, note the temperature when the
naphthalene melts, leave to cool, and note the temperature when the naphthalene
solidifies. To find the melting point, calculate the average of these two
values.
Repeat the experiment with stearic acid, m.p. 69oC, palmitic
acid, m.p. 63oC, butter, soap, 1,4-dichlorobenzene (deodorizer)
m.p. 53oC, paraffin wax, m.p. 45oC - 65oC,
sodium thiosulfate pentahydrate 48.3oC.
3.4 Melting point of substances
changed by impurities
Mix stearic acid with the naphthalene to make the naphthalene impure. Note
changes in the melting point. Impurities lower the melting point.
3.5 Boiling point of water
See diagram 3.2.5: Boiling point of water
The standard atmosphere (symbol: atm) 101325 Pa, for many purposes has
been replaced by the bar 100 kilopascals (100 kPa) = 1×105 N / m2 =
0.987 atm. (1 bar = 100 kPa, 100,000 Pa) The standard boiling point is now
the temperature at which boiling occurs under a pressure of 1 bar. On the
top of Mount Everest, at 8,848 m elevation, the pressure is about 260 mbar
(26.39 kPa) and the boiling point of water is 69 °C. The boiling point decreases
1 °C every 285 m of elevation.
1. Pour water into a test-tube. Hold a thermometer with its bulb just under
the water. Heat very slowly by moving the test-tube in and out of the flame
or add boiling chips, anti-bumping granules. Heat the water gently until
it boils. Record the temperature. Note the same temperature in all parts
of the test-tube. Note any change in the reading if the thermometer touches
the bottom of the test-tube. The water must cover the bulb of the thermometer
and the bulb must not touch the sides of the test-tube.
2. Show that the boiling point of water does not depend on the size of
the container. Repeat the experiment with a large container. Heat the water
quickly. The water first starts to boil near the bottom and sides of the
container. Note the temperature in different parts of the container. Note
any change in the reading if the thermometer touches the bottom of the container.
The boiling point is the same in small and large containers.
3.5.1 Boiling point of sodium
chloride solution
1. A solution of sodium chloride in water boils at a higher temperature
and has a lower freezing point than pure water. Use freezing points and
boiling points to find the purity of substances. Use three test-tubes containing
the same volume of water. Add some sodium chloride to the second test-tube.
Keep adding sodium chloride to the third test-tube until no more dissolves
to produce a saturated solution at that temperature. Join the test-tubes with
an elastic band. Heat the test-tubes equally over a Bunsen burner. The first
test-tube containing only water boils first. The second test-tube containing
some sodium chloride boils next. The third test-tube containing the saturated
solution of sodium chloride boils last.
2. Put a beaker containing demineralized water in a broad pan containing
a concentrated salt solution. Slowly heat the broad pan and note that the
demineralized water boils first.
3.6 Boiling point of inflammable
liquids, ethanol, acetone
See diagram 3.2.6: Boiling point of inflammable
liquids
1. Do not use a Bunsen burner to find the boiling point of inflammable
liquids, e.g. ethanol, b.p. 78.4oC and acetone, b.p. 56oC.
Use an electric hot plate or use the following method. Pour 2 cm of the
inflammable liquid into a test-tube in an empty container. Place a thermometer
in the test-tube with its bulb in the liquid. Boil water in an electric
jug or on an electrical hot plate. Pour the hot water into the container
so that the level is higher than the inflammable liquid in the test-tube.
Stir the inflammable liquid gently with the thermometer and read thermometer
when the inflammable liquid boils. [It is not good practice to stir liquids
with thermometers!]
2. Use a very small test-tube or seal one end of a piece of glass tubing,
8 cm length and 3 cm external diameter. Put the inflammable liquid into this
test-tube. Put a capillary tube, sealed at one end, into the inflammable liquid
with the sealed end up and the open end down in the inflammable liquid. Use
a rubber band to attach the test-tube containing inflammable and capillary
tube to the bulb of a thermometer. Hold the apparatus in a container of water
and heat gently with an electric hot plate. When the temperature rises, bubbles
slowly come out of the capillary tube. At the boiling point the bubbles suddenly
come out as a steady stream. Read the temperature. Let the water cool and
read the temperature again when the steady stream of bubbles stops. Calculate
the boiling point as the average of the two readings.
3. The boiling point of ethanol is 78oC. Heat water in a
kettle. Turn off the gas or extinguish the flame. Pour 2 cm of methylated
spirit into a test-tube. Put the test-tube in a beaker. Put a thermometer
into the methylated spirit deep enough to cover the bulb. Pour hot water
into the beaker so that the water level is higher than the level in the test-tube.
Stir gently with thermometer. When boiling starts, record the temperature.
4. Boiling point of
acetone. Repeat the experiment with propanone (acetone). BE CAREFUL! Highly flammable! THE FLASH POINT IS -20oC!
Boil the propanone (acetone) by placing the test-tube in hot water. Do
not allow flames in the room. The boiling range of 98% propanone (acetone)
is 55.5oC to 56.5oC. Nail polish remover formerly
contained propanone (acetone) but not in modern manufacture.
3.7 Volatility of different liquids
Evaporation is the movement of particles from the surface of a liquid to
the gas state, when below the boiling point. Volatile liquids evaporate readily
at room temperatures.
1. Select liquids from the laboratory, e.g. water, methylated spirit, gasoline,
mineral turpentine, kerosene (paraffin oil) household machine oil, car oil,
vinegar, vanilla essence, eucalyptus oil, glycerine. Wet a 5 cm piece of
absorbent paper with a liquid, Write the name of the test liquid in pencil.
Attach the piece of paper to a horizontal string. Examine the paper every
ten minutes, every two hours and each day.
2. Repeat the experiment with perfumes. Smell the paper every ten minutes,
every two hours and each day. Some perfumes soon disappear, but others last
for days. Record the relative "person-attracting" capacity for each perfume.
3.8 Pressure of the atmosphere
affects the boiling point
See diagram 3.2.8: Decreasing the pressure on
boiling water
1. Partially fill a flask with water and heat it until the water boils for
a few minutes. Remove the heat source so that the water stops boiling then
stopper the flask which now contains saturated water vapour water vapour at
atmospheric pressure Pour cold water over the flask. Some water vapour condenses
on the walls of the flask and the water boils again. The cold water caused
the vapour pressure in the flask to be reduced by condensation to below the
saturation pressure of the water in the flask at that temperature. So the
temperature at which water boils is lowered when the pressure is reduced.
2. Put water in a sidearm test-tube or in a round bottom flask with a
one-hole stopper. Insert a thermometer through a hole in the stopper so that
the bulb of thermometer reaches, but does not touch, the bottom of the test-tube
or flask. Add boiling chips to prevent bumping. Boil the water and read the
temperature. Stop heating. Connect a water pump to the sidearm or to the
second hole of the two-holes stopper. When the water stops boiling, turn
on the water pump to reduce the pressure. Read the temperature, heat to boiling
and read the temperature again.
3. When a liquid is at boiling point, its vapour pressure is equal to the
pressure of the atmosphere. Boil water on a high mountain and note the boiling
point. People who climb Mount Everest complain that they cannot get a good
cup of tea.
4. Some teachers play a silly trick on students. The ask a student to feel
the sides of the flask when the water is first boiling. When the second boiling
occurs at much lower temperature, the teacher spills some water on the hand
of the student who pulls way that hand thinking that the boiling water is
still very hot.
3.9 Solubility in water of different
salts
In this document the word "solution" refers to substances dissolved in
water, i.e. aqueous solutions. A solvent is a liquid that dissolves another
substance, the solute, to form a solution. The three ways to increase the
rate at which a solid dissolves in water are as follows: 1. grinding the
solid until finely divided 2. shaking the solution while the solid is dissolving,
and 3. heating the solution.
Try to dissolve 5 g of different salts each in 15 mL of water in a test-tube.
Attach a stopper and shake vigorously. Solubility is a characteristic of
a particular substance. Classify each salt as soluble or slightly soluble
or insoluble. The solubility of a salt can be expressed as the number of grams
able to dissolve in 100 g of water at 20oC, e.g. ammonium chloride
37.2 g, barium chloride 35.7 g, calcium chloride 42.7 g, copper (II) sulfate
20.7 g, lead nitrate 54.4 g, magnesium sulfate 25.2 g, potassium chloride
34.0 g, potassium iodide 144.0 g, sodium hydrogen carbonate (sodium bicarbonate)
9.6 g, sodium chloride 36.0 g, sodium hydroxide 109.0 g, sodium nitrate 87.5
g.
3.10 Solubility and temperature,
solubility of salts in water
The solubility of a potassium dichromate in 100 g of water varies with
temperatures: 0oC - 5 g, 10oC - 7 g, 20oC
- 12 g, 30oC - 20 g, 40oC - 26 g, 50oC
- 34 g, 60oC - 43 g, 70oC - 52 g, 80oC
- 61 g, 90oC - 70 g, 100oC - 80 g.
1. Show that a saturated solution contains less dissolved solid at a lower
temperature than at a higher temperature. Make a 50 mL saturated solution
of potassium dichromate or potassium nitrate at 60oC. Pour the
clear solution into a clean container and keep the temperature of this container
at 40oC until crystals stop forming. Pour the clear solution from
this container into another clean container. Do not pour crystals into the
container. Leave to cool and note more crystals forming as the solution cools.
2. Repeat the experiment with sodium chloride. This salt is unusual because
the solubility hardly changes with change in temperature.
3.11 Solubility of a substance
in water at a given temperature
1. Add a teaspoonful of powdered milk to water with ice floating in it,
water at room temperature, warm water, boiling water. Observe the time taken
for the powdered milk to dissolve in the water at different temperatures.
2. Add sodium hydrogen carbonate (sodium bicarbonate) to 100 g of water
in a container while stirring. Keep adding sodium carbonate until no more
solute will dissolve. Decant the clear saturated solution. Read the temperature
of the saturated solution, i.e. room temperature. Weigh a clean evaporating
dish, w1. Add some clear saturated solution and weigh again, w2. Carefully
evaporate the solution in the evaporating dish to dryness and weigh again,
w3. The mass of the dissolved sodium hydrogen carbonate = w3 - w1. The mass
of water = w2 - w1 - w3. Calculate the solubility of the sodium hydrogen
carbonate as weight in grams dissolved in water at room temperature (w3 -
w1) / (w2 - w1 - w3).
Repeat the experiment using water at a higher temperature.
3. The solubility of sodium bicarbonate in 100 g of water varies with temperature:
0oC - 6.9 g, 10oC - 8.15 g, 20oC - 9.6 g,
25o- 10.35 g, 30oC - 11.1 g, 40oC - 12.7
g, 50oC - 14.45 g, 60oC -16.4 g.
3.12 Solubility and particle
size, sodium chloride, copper (II) sulfate
Show that small particles dissolve faster than large particles.
1. Add coarse table salt to a first test-tube half filled with water. Add
the same quantity of fine table salt to a second test-tube that contains the
same volume of water. Shake both test-tubes equally and simultaneously. Note
the amount of undissolved table salt left in each test-tube.
2. Use two equal samples of large crystals of copper (II) sulfate. Grind
one sample into a fine powder. Put both samples into equal quantities of
water in separate test-tubes and shake. Compare the rates at that the different
samples dissolve and cause the water to change colour.
3.13 Solubility and solvents,
sodium chloride, methylated spirit
1. Fill two test-tubes one third full with water and another with methylated
spirit. To each test-tube add 1 g sodium chloride, attach a stopper and shake.
Sodium chloride dissolves readily in water, but not so readily in methylated
spirit.
2. Add sodium chloride crystals to a dilute solution of sugar in water
and note whether the crystals dissolve. Drop crystals of potassium dichromate
into the solution. Note whether the solution changes colour. Colour change
shows that potassium dichromate is also dissolving. The presence of one dissolved
substance does not prevent other substances dissolving in the solution. Unless
the concentrations are high, one solute does not affect the solubility of
other solutes in the solution.
3.14 Solubility and agitation,
cane sugar
Prepare two equal samples of cane sugar. Put one sample of cane sugar into
a test-tube half full of water. Add a stopper and shake the test-tube until
all the sugar dissolves. Put the other sample of cane sugar into a test-tube.
Very slowly add the same volume of water as in the first test-tube. Leave
to stand. Most of the sugar has not dissolved but, if left to stand for long
enough, all the cane sugar will dissolve as in the first test-tube.
3.15 Volume of solutions, sodium
chloride
1. Fill a small, narrow neck flask with water to a level in the neck
and mark this level. Add sodium chloride to the water with continual shaking
until the solution is saturated and no more dissolves. Note the new level
of the liquid. The volume of the solution is only slightly greater than the
original volume of the water.
2. Close one end of a glass delivery tube about 30 cm long. Fix it upright,
half fill it with water and mark the level of the water. Slowly add alcohol
to fill the delivery tube. The water and the alcohol fill equal lengths in
the tube. Shake the tube thoroughly to mix the water and alcohol. The new
level of the solution in the tube shows a slight decrease in total volume.
3.16 Miscible liquids, methylated
spirit, glycerine, kerosene.
Put 10 mL of water in three test-tubes. Add 1 mL of: 1. methylated spirits,
2. glycerine, and 3. kerosene. Shake each test-tube. Miscible liquids can
mix in all proportions. 1. Alcohol and water are miscible. 2. Glycerine and
water are miscible. 3. Kerosene and water are not miscible, immiscible.
3.17 Heat of solution, sodium
hydroxide, potassium hydroxide, ammonium chloride, urea
Dissolve some equal quantities of sodium hydroxide, potassium hydroxide,
ammonium chloride and urea in separate test-tubes half full of water. Feel
the test-tubes and note any change in the temperature. Sodium hydroxide and
potassium hydroxides dissolve in water with an increase in temperature. Ammonium
chloride and urea absorb heat from their surroundings when dissolving in
water.
3.17.1 "Magnetic" sugar cube
dissolves
Fill a large dish with water Wait until the water is absolutely still then
lower a matchstick into the centre of the water. Carefully dip a sugar cube
in the water near the edge of the dish. The matchstick moves towards the
dissolving sugar cube. When the sugar dissolves in the surface water, the
surface water becomes heavier and falls to be replaced by surface water flowing
towards the sugar cube, carrying the matchstick with it.
3.18 Separate tin from a mixture
of tin and carbon (charcoal)
Get tin bits by cutting a tin welding rod to pieces because 66% of the
tin welding rod is tin and the rest is lead. Do not use a "tin can" because
it is mostly iron with a thin layer of tin on its surface!
1. Make a mixture of tin (tin filings or small cut pieces of tin) mp 232oC
and carbon (crushed charcoal) mp 3,730oC. Mix the tin bits and
charcoal bits uniformly. Heat the mixture in a crucible. Stir with a splint
until the tin melts and forms a liquid below the charcoal. Pour the tin onto
a plaster of Paris mould or other heat-proof surface. While pouring, hold
back the charcoal in the crucible with a wood splint. Use melting point and
melting point behaviour to identify a substance and decide if it is pure.
Tin solder melts at 250oC. Carbon melts at 3,700oC.
2. Mix solder filings with powdered charcoal. Heat the mixture in a crucible.
Stir with a splint until the solder melts and forms a liquid below the charcoal.
Pour the liquid into a container by holding back the charcoal in the crucible.
3.19 Separate by sublimation
See diagram 3.2.19: Sublimation of iodine
Separate iodine from a mixture of crystals of iodine and sodium chloride.
Heat the mixture in an evaporating dish with a funnel placed over it. The
iodine sublimes on to the cool sides of the funnel.
3.20 Separate by distillation
See diagram 3.2.20.1: Distil ink | See diagram 3.2.20.2: Condensing the vapour
Put 10 mL of ink in a flat-bottom conical flask. Add boiling chips to prevent
bumping. Fit a stopper with a delivery tube reaching half way down a receiving
test-tube or a U-tube, in a container of water. Heat the ink with a Bunsen
burner flame. Drops of a colourless liquid appear in the receiving test-tube.
Identify the liquid as water by its action of turning white anhydrous copper
(II) sulfate to blue hydrated copper (II) sulfate. Do not allow ink to froth
up or splash into the delivery tube.
3.21 Separate crude oil fractions
by fractional distillation
See diagram 3.2.21: Collect fractions
1. Use crude oil or a substitute for crude oil, e.g. a mixture of used
car oil, paraffin, thin lubricating oil, diesel oil and petroleum jelly.
Use a hard glass test-tube, or sidearm test-tube, fixed to a retort stand,
a delivery tube and five small ignition tubes. Use a 0o to 360oC
thermometer. Put 4 mL of crude oil in the test-tube. Add boiling chips to
prevent bumping. Set up five small ignition tubes to collect the fractions.
Heat the oil very gently. Collect 10 drops of distillate in the first ignition
tube, then collect 10 drops of distillate successively in the other ignition
tubes. The boiling point of the remaining oil will become higher as distillation
proceeds and oil will then require more heat from the Bunsen burner. Arrange
the fractions in order of increasing distillation temperature: 1. up to 80oC
2. 80 - 120oC 3. 120 - 180oC 4. 180 - 220oC.
2. Examine the different fractions:
2.1 The colour should change from colourless to yellow.
2.2 The viscosity should increase.
2.3 The high temperature fractions should be more difficult to ignite than
the low temperature fractions.
2.4 The high temperature fractions should burn with more soot in the flame
than low temperature fractions. Burn the fractions in bottle tops with the
cork removed.
2.5 Note the dark residue remaining in the test-tube.
3.22 Separate salt and sand
Prepare a mixture of salt and sand. Put 2 mL of the mixture in a test-tube.
Add 5 mL of water and shake until all the salt has dissolved. Pour the contents
of the tube into a filter paper in a funnel over an evaporating basin. Wash
the test-tube with water and add this to the filter paper. The sand will
remain on the filter paper and may be dried and collected. Recover the salt
from the filtrate by warming the evaporation basin to drive off the water.
3.23 Solvent extraction of oil
from nuts
Put peanuts (groundnuts) or pieces of chopped coconut into a mortar. Add
20 mL of acetone or methylated spirit. Grind the nuts in the solvent as finely
as possible. Pour off the liquid into a test-tube and filter into an evaporating
basin. Warm the evaporating basin for 10 minutes. The solvent evaporates
leaving the oil extracted from the nuts.
3.24 Separate pigments from
green leaves with paper chromatography
See diagram 3.2.24: A chromatogram
1. Collect green leaves and cut them into very small pieces. Use a mortar
and pestle to grind the leaves for five minutes with a small volume of
methylated spirit and clean sand until a deep green solution forms. Draw
a fine pencil line 5 cm from the end of a 1 cm wide strip of absorbent paper (or chromatography paper).
Suspend the absorbent paper in a test-tube without touching the bottom. Use
a fine eye-dropper to put one small drop of the solution on the centre of
the fine pencil line and let it dry. When the drop is dry add more solution to the same place
to make a small concentrated spot of 5 drops. Hang the paper strip with the lower end
in the methylated spirit solvent and the spot of green solution above the
solvent level. Leave the paper strip in the solvent until the methylated
spirit has almost reached the top of the absorbent paper. Capillary attraction
draws up the solvent. Mark the chromatogram on the paper to show a top orange-yellow
band of xanthophyll and a lower green band of chlorophyll. A band of carotene
is visible if the solvent is toluene.
2. Repeat the experiment with other solvents, e.g. toluene, acetone (propanone)
3.24.1 Separate mixed inks
with paper chromatography
Prepare a mixed solvent from 6 parts of water, 3 parts of methylated spirit,
and 1 part of ammonia solution. Put 5 mL of mixed solvent in a test-tube.
Prepare mixed ink from equal quantities of red and blue ink. Put a drop of
the mixed ink near one end of a 2 cm wide paper strip. Lower the paper strip
so that its lower end is in the mixed solvent. Use a stopper to prevent evaporation.
As the solvent moves up the paper strip, the component colours of the ink
separate to form different coloured bands with red above and blue below.
Try other solvents and other inks to obtain good separation of colours.
Repeat the experiment by drawing a line with a ball pen or an ink pen near
the end of the paper strip.
3.25 Separate gases dissolved
in a water sample
See diagram 3.2.25: Gases in water
Stand a container of water in sunlight. Bubbles of air appear. The taste
of boiled water is different from the taste of tap water because boiled water
has lost its dissolved oxygen. Note the temperature of a sample of water.
Boil the water until no more bubbles appear. Collect the gases from the water
in an inverted measuring cylinder.
3.26 Separate immiscible liquids
of different density
See diagram 3.2.26: Separation tube
Separate two immiscible liquids of different density, e.g. kerosene (paraffin
oil) and water. Use a separating funnel or make a separating funnel with
a piece of wide plastic tubing fitted with a one-hole stopper and rubber tubing
with a clip. Shake the mixture thoroughly in a closed container then run
it into the separating funnel. Wait until a clear boundary appears between
the two liquids and then run off the more dense layer into a container below.
3.27 Separate solids using density
differences
In industry, a separator concentrating machine shakes mixed ores to separate
the different ores. Beach sand often consists of quartz particles mixed with
heavier particles such as ilmenite or zircon. Shake a mixture of sand and
iron oxide to make them separate into different layers.
3.28 Substances that gain mass
when heated, copper foil
Cleaned copper is brown-red. In moist air the surface turns green due to
oxidation. The green surface is called a patina. It also forms on old unpolished
bronze.
1. Clean a piece of copper foil with steel wool. Hold it in a flame with
a pair of tongs. The black copper (II) oxide looks like carbon. To test the
substance, drop dilute sulfuric acid on it, then heat it. Blue copper (II)
sulfate forms. Test some powdered carbon. No colour change occurs.
2. Show that something is added to the copper from the air. Use a sensitive
balance to weigh the copper before and after heating.
3. Use two identical hard glass test-tubes with one-hole stoppers fitted
with bent delivery tubes. Fix both test-tubes to a stand so that the test-tubes
slope down with the ends of the delivery tubes under water in a beaker. Put
copper foil in the first test-tube and heat with a hot burner flame. After
two minutes, heat the empty second test-tube. Move the burner regularly between
the two test-tubes until no more bubbles come out of the ends of the delivery
tubes. Stop heating both test-tubes. As the test-tubes cool, they suck water
up the delivery tube. The test-tube containing the copper (II) oxide sucks
up more water.
3.28.1 Substances that gain
mass when heated, prepare magnesium oxide
Use magnesium ribbon because magnesium powder is too reactive. Be careful!
Do not heat magnesium powder.
Magnesium has density 1.74 g / cm3 and melting point 650oC,
but magnesium oxide has density 3.58 g /cm3 and melting point
2800oC because the Mg2+-- O2- chemical bond
is stronger than the Mg -- Mg bond.
1. Hold a 10 cm strip of magnesium ribbon in a pair of tongs. Place the
ribbon in a Bunsen burner flame until it starts to burn. Be careful! Magnesium
burns with a very bright white light. Magnesium ribbon corrodes slightly
in air and burns with an intense white flame to form a white ash of magnesium
oxide.
Mg + 1 / 2O2 --> MgO
2. Clean 25 cm of magnesium ribbon and cut into pieces 1 cm long. Put the
pieces into a crucible with a lid. Weigh the crucible + lid + contents =
W1. Put the crucible on a pipe clay triangle on a tripod stand. Heat gently
then strongly. Use tongs to raise the lid. The magnesium darkens before it
melts. When the magnesium starts to burn, put the lid back on the crucible
and remove the burner. Every few seconds raise the lid slightly to let more
air enter. Do not let white magnesium oxide smoke escape. When the magnesium
does not burn after you raise the lid, remove the lid and heat the crucible
strongly. Hold the lid ready in case the magnesium starts to burn again. Let
the crucible cool. Again weigh the crucible + lid + contents = W2. Note W2
> W1. The formation of magnesium oxide causes the increase in weight.
3.28.2 Substances that gain
mass when heated, prepare lithium oxide
Heat pieces of lithium metal shot on a metal spoon (deflagrating spoon).
Note the violet glow when it starts to burn, then put the burning lithium
in oxygen gas.
3.28.3 Substances that gain
mass when heated, prepare calcium oxide
Heat a shaving of calcium metal in a crucible and heat it with a Bunsen
burner for 10 -15 minutes because it is difficult to ignite.
3.28.4 Collect and weigh the
gaseous products of a burning candle
See diagram 3.2.29: Gaseous products of burning
candle
Candle wax is a mixture of different alkanes (paraffins) saturated hydrocarbons
with general formula CnH2n+2 that are solid at room
temperature. Soda lime is a grey-white mixture of sodium hydroxide and calcium
hydroxide as granules or powder that absorbs the products of combustion,
carbon dioxide and water. Use soda lime instead of sodium hydroxide because
soda lime is not deliquescent. Weigh a candle, C1. Weigh a U-tube containing
granules of soda lime, U1. Put a candle under an inverted glass filter funnel
connected to one arm of the U-tube. Attach a filter pump to the other arm
to draw air through the U-tube. Light the candle and turn on the filter pump
to draw air over the candle. Let the candle burn for five minutes. Extinguish
the candle and disconnect the filter pump. Weigh the candle again, C2. The
candle has lost weight, C1-C2. When the U-tube is cool, weigh it again, U2.
The U-tube containing the soda lime has gained weight, U2-U1. The U-tube gains
more weight than the candle loses weight (U1-U2) > (C2-C1) for two reasons:
1. The candle wax combines with oxygen gas in the air to form carbon dioxide
gas and water. 2. The air sucked in by the filter pump contains some water
vapour absorbed by the soda lime. To measure the weight of water absorbed
from the air, in a control experiment, repeat the experiment for the same
period, but without the candle.
3.30.0 Substances that decompose
and lose mass when heated, thermal decomposition
See diagram 3.2.32: Collecting the products
of heating copper sulfate crystals
Substances that remain stable after heating constantly with a Bunsen burner
flame may be listed under the heading "Substances that do not decompose when
heated". However, all compounds break down when heated to a high enough temperature.
Metal compounds higher in the reactivity series are usually more stable than
compounds of metals lower in the reactivity series. Salts that remain stable
when heated constantly with a Bunsen burner flame are calcium sulfate, potassium
chloride, potassium sulfate, sodium carbonate, sodium chloride, and sodium
sulfate. Ammonium oxalate (NH4COO)2, and ammonium dichromate
(NH4)2Cr2O7, decompose before
melting. Ammonium sulfate (NH4)2SO4, decomposes
above 280oC.
3.30.1 Decomposition
of carbonates
Carbonates mostly decompose to form carbon dioxide and a metallic oxide.
Sodium carbonate and potassium carbonate do not decompose when heated to
a high temperature.
CaCO3 (s) --> CaO (s) + CO2 (g)
white calcium carbonate --> white calcium oxide + carbon dioxide
CuCO3 (s) --> CuO (s) + CO2 (g)
blue-green --> black
MgCO3 (s) --> MgO (s) + CO2 (g)
white --> white
PbCO3 (s) --> PbO (s) + CO2 (g)
white --> yellow
ZnCO3 (s) --> ZnO (s) + CO2 (g)
white --> white (cold) or yellow (hot)
Ammonium carbonate may decomposes completely without heating when exposed
to the air, to give ammonia, water and carbon dioxide.
(NH4)2CO3 (s) --> 2NH3 (g)
+ H2O (vapour) + CO2 (g)
colourless
3.30.2 Decomposition
of hydrogen carbonates (bicarbonates)
Hydrogen carbonates (hydrogencarbonates, bicarbonates) decompose to form
the metal carbonate, water and carbon dioxide.
Calcium bicarbonate and magnesium bicarbonate can exist only as a solution
that on heating form the metal carbonate, water and carbon dioxide.
Sodium hydrogen carbonate, NaHCO3 (sodium bicarbonate) begins
to lose carbon dioxide at 50oC to form sodium carbonate. A solution
of a sodium hydrogen carbonate begins to lose carbon dioxide at 20oC.
Heat sodium hydrogen carbonate crystals. The crystals lose water and carbon
dioxide, and at 100oC are converted to sodium carbonate.
2NaHCO3 (s) --> Na2CO3 (s) + CO2 (g)
+ H2O (vapour)
colourless --> colourless
Ca(HCO3)2 (aq) --> CaCO3 (s) + CO2
(g) + H2O (vapour)
Mg(HCO3)2 (aq) --> MgCO3 (s) + CO2
(g) + H2O (vapour)
2KHCO3 (s) --> K2CO3 (s) + CO2 (g)
+ H2O (vapour)
3.30.3 Decomposition
of hydroxides
Hydroxides decompose to form water and the metallic oxide
3.30.4 Decomposition
of nitrates
Nitrates decompose to form oxygen gas, nitrogen dioxide and a metallic
oxide.
2Ca(NO3)2 (s) --> 2CaO + 4 NO2 (g)
+ O2 (g)
colourless --> white
2Cu(NO3)2 (s) --> 2CuO + 4 NO2 (g)
+ O2 (g)
blue --> black
2Pb(NO3)2 (s) --> 2PbO + 4 NO2 (g)
+ O2 (g)
colourless --> yellow
Lead nitrate decomposes at 470oC.
2Zn(NO3)2 (s) --> 2ZnO + 4 NO2 (g)
+ O2 (g)
colourless --> white (cold), yellow (hot)
Potassium nitrate and sodium nitrate first melt and then decompose to give
the metal nitrite and oxygen gas. Potassium nitrate melts at 336oC.
Sodium nitrate melts as 316oC.
2KNO3 (s) --> 2KNO2 (s) + O2 (g)
colourless --> colourless
2NaNO3 (s) --> 2NaNO2 (s) + O2 (g)
colourless --> yellow
Silver nitrate decomposes to give the metal, nitrogen dioxide and oxygen
gas.
2AgNO3 (s) --> 2Ag (s) + 2NO2 (g) + O2
(g)
colourless --> silver
Ammonium nitrate decomposes to form water vapour and nitrous oxide, N2O
(laughing gas) so the ammonium nitrate disappears.
NH4NO3 (s) --> N2O (g) + H2O
(g)
colourless
3.30.5 Decomposition
of oxides
Oxides of most metals are stable. Oxides of potassium, sodium, calcium,
magnesium, aluminium, zinc, iron, lead and copper do not decompose.
Black-grey silver oxide decomposes into the metal and oxygen gas.
2Ag2O (s) --> 4Ag (s) + O2 (g)
silver oxide --> silver + oxygen
Heat zinc oxide in a crucible. Zinc oxide becomes yellow when hot and white
when cold but no change in weight occurs. The substance does not decompose
and does not gain anything from the air or lose anything to the air. Zinc
oxide has melting point above 1,800oC.
ZnO (s) <--> ZnO (s)
white (cool) yellow (hot)
Thermal decomposition of higher oxides of lead
2PbO2 (s) --> 2PbO (s) + O2 (g)
brown lead dioxide --> yellow lead oxide + oxygen gas
2Pb3O4 (s) --> 6PbO (s) + O2 (g)
red trilead tetroxide --> yellow lead oxide + oxygen gas
3.30.6 Decomposition
of sulfates
Sulfates if heated very strongly may decompose to form the metallic oxide,
sulfur dioxide and oxygen gas
Put 4 cm of crushed blue copper (II) sulfate crystals in a dry test-tube
fitted with a one-hole stopper and delivery tube. Heat the dry test-tube
and crystals gently. Note whether vapour collects on the cooler parts of the
dry test-tube and whether any liquid collects in the receiving test-tube.
Note any change of colour of the crystals from blue to white. Identity the
liquid in the receiving test-tube by measuring the boiling point. When all
the copper (II) sulfate crystals have become white and the dry test-tube
has cooled, pour the liquid in the receiving test-tube back on the white crystals.
Note whether the blue colour of the crystals is restored and if any heat
is given off.
2CuSO4 (s) --> 2CuO (s) + 2SO2 (g) + O2
(g)
grey white --> black
2PbSO4 (s) --> 2PbO (s) + 2SO2 (g) + O2
(g)
white --> yellow
2ZnSO4 (s) --> 2ZnO (s) +2SO2 (g) + O2
(g)
white --> white (cold) yellow (hot)
3.30.7 Decomposition
of sulfites
Sulfites mostly decompose into the metal oxide and sulfur dioxide
Sulfites of sodium and potassium do not decompose when heated.
CaSO3 (s) --> CaO (s) + SO2 (g)
white --> white
MgSO3 (s) --> MgO (s) + SO2 (g)
white --> white
ZnSO3 (s) --> ZnO (s) + SO2 (g)
white --> white (cold) yellow (hot)
3.30.8 Decomposition
of hydrous salts
Salts with water of crystallization, hydrous salts, lose their water of
crystallization, and become anhydrous powders and then become stable
They also lose their crystalline shape and sometimes their colour.
Prepare test-tubes containing 1 cm of 1. iodine crystals 2. sodium hydrogen
carbonate granules or crystals 3. silica sand 4. zinc oxide. Fix a cotton
wool plug in the mouth of each test-tube to prevent loss of solid during
heating, then weigh each test-tube. Heat each test-tube and cotton wool plug
thoroughly and weigh it again. Note any change in weight because of the loss
of water of crystallization.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous
copper (II) sulfate + water.
CuSO4.5H2O (s) --> CuSO4 (s) + 5H2O
(vapour)
blue --> grey white
Na2CO3.10H2O --> Na2CO3
(s) + 10H2O (vapour)
colourless --> white
3.30.9 Decomposition
of boric acid
Boric acid, H3BO3, loses water until it decomposes to the anhydride,
B2O3.
3.30.10 Decomposition
of oxalic acid
Oxalic acid begins to sublime at 100oC, becomes anhydrous at
189oC and when heated rapidly decomposes into carbon dioxide,
carbon monoxide, formic acid and water.
3.30.11 Decomposition
of potassium chlorate
Potassium chlorate, KClO3, decomposes above 368oC
into potassium perchlorate and oxygen gas
3.30.11.1 Decomposition
of potassium ferricyanide
Potassium ferricyanide, K2Fe(CN)6, decomposes before
melting.
3.30.11.2 Decomposition
of phosphates
Monosodium orthophosphate, NaH2PO4.H2O,
and disodium orthophosphate [disodium hydrogen phosphate (V)] Na2HPO4.12H2O,
lose water of crystallization.
10KClO3 <--> 6KClO4 + 4KCl + 3O2
3.30.12 Sublimation,
iodine, camphor, naphthalene
See diagram 3.30.12: Sublimation of iodine
Sublimation is when a solid changes directly into a gas without melting.
Iodine, camphor and naphthalene do not decompose when heated but sublime
from the crystal to the vapour form.
Put black, shiny crystals of iodine in an evaporating dish. Cover the dish
with a piece of filter paper and stand a filter funnel upside down on the
dish. Heat the dish gently. Purple vapours rise through the filter paper.
As they cool in the filter funnel, shiny black crystals of iodine form again.
3.30.13 Heat silica
sand, SiO2
Silica sand consists of pieces of silicon (IV) oxide (SiO2)
crystals. Heat sand in a crucible. The sand particles may break up physically,
but do not break up chemically.
3.30.14 Decomposition
of ammonium dichromate, (NH4)Cr2O7
Ammonium dichromate is an orange coloured crystalline substance
1. Show the decomposition of an ammonium dichromate by heating < 5 g
on a fireproof surface in a fume cupboard or outside with observers positioned
up wind of the reaction. When outdoors, hose down area well to ensure fine
dust is not scattered. It starts decomposing with sparks and gives out heat
to form a green fluffy powder chromic oxide, nitrogen and water.
2. An ammonium dichromate is unusual in that it undergoes an exothermic
decomposition on heating to form green chromium (III) oxide, nitrogen gas
and water. Start the reaction with a lighted match. The reaction continues
because it is exothermic. This reaction is not dangerous if only small quantities
of material are used. The reaction is dangerous if particles of undecomposed
ammonium dichromate are projected into the air and inhaled. Ammonium dichromate
is often used in the volcano experiment.
3.30.15 Decomposition
of ammonium chloride, NH4Cl
Put ammonium chloride into the bottom of a dry test-tube and warm it over
a small flame. The solid decomposes to form ammonia and hydrogen chloride.
Some of it recombines at the cool upper part of the test-tube as a white
layer. Heat the recombined layer again and it moves further up the test-tube.
This process is recombination not sublimation.
NH4Cl (s) <--> NH3 (g) + HCl (g)