School Science Lessons
Chemistry
Updated: 2009-10-10
Please send comments to: J.Elfick@uq.edu.au
Table of contents
3.1.0 Bunsen burner
3.2.0 Identify pure substances
3.9.0 Solubility
and solutions
3.18.0 Separate substances from mixtures
3.28.0 Substances that gain mass when
heated,
copper foil
3.30.0 Substances may decompose and
lose mass when heated, thermal decomposition
3.31.0
Hygroscopic, deliquescent and efflorescent chemicals
3.32.0 Prepare, collect
and test gases
3.52.0 Rusting
5.3.4 Acid-base
indicators
3.54.0
Crystal growth
3.55.0 Matter as particles
3.1.0 Conductivity
3.61.0
Construction materials
3.70.0 Chemical reactions
3.80.0 Energy from chemical reactions
3.84.0 Electrical energy from
chemical reactions
3.91.0 Rate of reaction
3.95.0 Breakdown large molecules to
small molecules
3.100.0 Building up molecules
3.1.0 Bunsen burner
3.1.1
Bunsen burner
3.1.1.1 Bunsen burner gases
3.1.2 Lighting a Bunsen burner
3.1.3 Study the Bunsen burner flame
3.1.4 Bunsen burner flame and candle flame
3.2.0 Identify pure
substances
3.2
Melting point of naphthalene
3.3
Melting point of naphthalene with a capillary tube
3.4
Impurities affect the melting point of a substance
3.5
Boiling point of water
3.5.1
Boiling point of sodium chloride solution
3.6
Boiling point of inflammable liquids
3.7
Volatility of different liquids
3.8
Pressure affects the boiling point
3.9.0 Solubility and solutions
3.9
Solubility in water
3.10
Solubility and temperature, solubility of salts in water
3.11
Solubility of a substance in water at a given temperature
3.12
Solubility and particle size
3.13 Solubility and solvents
3.14 Solubility and agitation
3.15 Volume of solutions
3.16 Miscible liquids
3.17 Heat of solution
3.17.1 "Magnetic" sugar cube dissolves
3.18.0 Separate substances from mixtures
3.18
Separate tin from a mixture of tin and carbon (charcoal)
3.19
Separate by sublimation, iodine
3.20
Separate by distillation
3.21
Separate by fractional distillation of crude oil
3.22
Separate salt and sand
3.23
Solvent extraction of oil from nuts
3.24
Separate pigments from green leaves with paper chromatography
3.24.1
Separate mixed inks with paper chromatography
3.25
Gases dissolved in a water sample
3.26
Separate immiscible liquids of different density
3.27
Separate solids using density differences
3.28.0
Substances that gain mass when
heated,
copper foil
3.28.1
Substances that gain mass when heated, prepare magnesium oxide
3.28.2 Substances that gain mass when heated,
prepare lithium oxide
3.28.3 Substances that gain mass when heated,
prepare calcium oxide
3.28.4
Collect and weigh the gaseous products of a burning candle
3.30.0
Substances may decompose and lose mass when heated, thermal
decomposition
3.30.1 Carbonates
3.30.2 Hydrogen carbonates (hydrogencarbonates, bicarbonates)
3.30.3 Hydroxides
3.30.4 Nitrates
3.30.5 Oxides
3.30.6 Sulfates
3.30.7 Sulfites
3.30.8 Salts
with water of crystallization,
hydrous salts
3.30.9 Boric acid, H3BO3
3.30.10 Oxalic acid
3.30.11 Potassium
chlorate, KClO3
3.30.12 Sublimation, iodine, camphor,
naphthalene
3.30.13 Silica sand
3.30.14 Ammonium dichromate
3.30.15 Ammonium chloride
3.31.0
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
3.1.1 Bunsen burner
See diagram 3.1: Bunsen burner flame | See diagram 3.1.1: Bunsen burners
Be careful! Do not turn the gas on
without lighting the Bunsen burner. Gas forms an explosive mixture in
air.
1. Combustion is the burning in oxygen gas of a substance to produce
heat
and
sometimes light energy. A flame appears during combustion when a gas
has such a high temperature that it emits heat and light. A flame
appears only where the burning gas and oxygen gas are in contact.
2.
The Bunsen burner consists of 2.1. a barrel, shaft, 2.2 an air
regulator,
i.e. a sleeve with a hole in it, 2.3 a jet, air mixture valve,
needle valve, 2.4 a base, 2.5. a gas inlet opening.
3. Adjust the flame by
opening or closing the gas tap. When the air regulator is open, the gas
burns with a noisy blue flame that may be nearly invisible in strong
light. If the flame rises up from the burner, turn down the gas supply.
4. When not using the Bunsen burner, either turn off the gas
or close the air regulator to give a safety flame. The flame is
yellow because of the incandescence of carbon particles. It is not as
hot
as the blue flame and leaves black soot deposits on glassware.
5. Regularly inspect gas fittings on the benches and hoses connecting
Bunsen burners to gas turrets to make sure that connections are free of
leaks.
Tests for leaks by dipping the part in soapy water. Be careful!
Do not use
a lighted match.
6. Heat flammable liquids in water baths using electrical
hot plates, not Bunsen burners. Turn the gas off first at the gas
tap, then at the cylinder or main supply tap.
7. Use the Bunsen burner
only in a draught free area. Allow the Bunsen burner to cool before you
move or store it.
8. Do not heat low melting point objects, e.g. plastics,
solder, lead, over the barrel of the burner. Melted pieces may
fall inside the barrel. Hold the burner at an angle. If a match is
blown out, turn gas off, then light the Bunsen burner again.
3.1.1.1 Bunsen burner
gases
See
16.1.1cc:
LPG (liquefied petroleum gas, LP gas)
See
16.1.1a.1: Natural gas
Town gas was previously used on
school laboratories, but nowadays the gas is usually natural gas, i.e.
mostly methane CH4,
or
LPG, bottled gas, mostly propane, C3H8.
In
a laboratory, the pilot light should burn with a 90% blue flame. If the
flame is yellow, the gas may be contaminated with condensates. Do not
use such a gas but immediately inform the local gas authority.
The heating values of fuels: town
gas 88 MJ / kg, natural gas 55.6 MJ / kg, LPG
gas 49 MJ /
kg, diesel fuel oil 38 MJ / L, kerosene 36.7 MJ
/ L, coke or coal 27 MJ / kg, dry split wood 12.5 MJ
/ kg.
CH4 (g) + 2O2 (g) --> CO2 (g) + 2H2O
(g)
+ heat
3.1.2 Lighting a Bunsen
burner
See diagram 3.1.2: Right and wrong ways
to
use a Bunsen burner
1. Close the air regulator, light a match, hold the match flame at the
side of the barrel opening, turn the gas tap on, raise the match flame
to light the gas. The gas burns with a visible yellow flame, a
quiet safety flame. Hold a test-tube just above the flame. Note the
carbon (soot, carbon black) that deposits on the glass. To test
whether unburned carbon gives the yellow colour to the flame, sprinkle
powdered charcoal on the flame and compare the yellow colours.
2. Start to open the air regulator until the gas burns with a
medium blue flame with a light blue inner cone and a pale violet
outer flame with a bushy appearance. The flame has an outer
oxidizing zone where combustion is complete, a middle reducing zone,
and an inner unburned gases zone surrounded by a blue cone. This flame
is the most useful for heating. Fully open the gas regulator until you
get a roaring blue flame
with a light blue triangle in the centre of the blue cone.
3. Open the air regulator. Keep turning down the gas
supply. The gas "blows back", "strikes back". The gas is burning inside
the barrel. Turn the gas fully on and strike the gas supply rubber tube
with a sharp blow from the side of your hand. If the flame does not
reappear, immediately turn the gas off and leave to cool because the
barrel may be hot. Then light the Bunsen burner again.
3.1.3 Study the Bunsen
burner flame
See diagram 3.1.3: Burning the gas in a
cone of flame
1. Hold the end of a glass tube in the
centre of the cone. You can
light the gas coming out of the other end of the glass tube.
2. Hold a piece of wire in different parts
of each kind of
flame, moving it from the bottom to the top. Find the hottest flame and
the hottest place in each flame with a piece of nichrome wire or iron
wire stuck into a cork for a handle. The approximate temperatures and
colours for the wire are as follows: 1. <500oC, wire
gives no
light, flame is non-luminous 2. 500oC to 950oC,
wire becomes red, then dark red, then bright red (red-hot) 3. 950oC
to 1350oC, wire becomes yellow-red then becomes
white 4. >1350oC, wire becomes white (white-hot). The
safety flame has a similar temperature in different parts about 300oC.
It is never used for heating. The medium blue flame has the hottest
point at the tip of the blue cone at about 500oC. The
roaring blue flame has the hottest point at the tip of the cone at
about 700oC.
3. Close the air regulator. Use a wood splint or a taper to test that
parts of the flame support ignition. The wood splint match is set
alight in all positions in the yellow flame where no air mixes with the
gas. Repeat the experiment with the air regulator open. A cone of mixed
air and gas exists in the centre of the cone where the gas is not
burning.
4. Turn off the gas. Push a pin at right angles through a match just
below the chemical on the end of the match. Use the pin to hang the
match in the barrel with the chemical end just above the rim. Open the
air regulator and light the gas again. The match does not ignite inside
the
cone. Move the match to the outer cone of the blue flame. The match
ignites.
5. Close the air regulator and light the gas. Hold a piece
of copper wire gauze with tongs 3 cm above the top of the barrel. Hold
a
lighted match above the gauze. The gas ignites above the gauze.
Lower the gauze until the flame passes through it. Repeat the
experiment with an open air regulator. Light the gas and lower a copper
wire gauze down on the flame. The flame remains below the wire gauze as
the gauze becomes red-hot. Heat is removed from the gas air mixture by
the copper gauze.
3.1.4 Bunsen burner flame
and candle flame
See diagram 3.1.4: Bunsen burner flame
and candle flame
Repeat the above experiments with a candle and a spirit
burner. Just above the wick of a burning candle is a dark region of
unburned gas. Above and around it is a yellow region containing
incandescent particles of carbon undergoing combustion to form carbon
dioxide. Put the candle flame under an evaporating basin. Note
the deposits of carbon, soot, because of insufficient oxygen gas to
complete combustion.
3.2 Melting point of naphthalene
See diagram 3.2: 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 points of naphthalene with a capillary
tube
See diagram 3.3: More accurate way of finding
the melting point
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 Impurities affect the melting point of a
substance
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 3.5: Boiling point of water
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
See diagram 3.6: 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.
3.7 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 affects the boiling point
See diagram 3.8: Decreasing the pressure on
boiling water
1. 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.
2. 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.
3.9 Solubility in water
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
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
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
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
1. Fill a small, narrow-necked 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
Put 10 mL of water in three test-tubes. Add 1 mL of: 1. methylated
spirit 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
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 plaster of Paris mold 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 3.19: 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 3.20.1: Distil ink | See
diagram 3.20.2: Condensing the vapour
Put 10 mL of ink in a flat-bottomed 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 3.21: 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 3.24: A chromatogram
A absorbent paper, B solution of crushed leaves
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. 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. Add more solution to the same place to make a
small concentrated spot. 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 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 Gases dissolved in a water sample
See diagram 3.25: 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 3.26: 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 red-brown. 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 3.29: Gaseous products of burning
candle
Candle wax is a mixture of different alkanes (paraffins) saturated
hydrocarbons with general formula CnH2n+2that
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 Substances may decompose and lose mass when
heated, thermal decomposition
See diagram 3.30: 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 breakdown when heated to a high enough
temperature. Metal compounds higher in the activity series are usually
more
stable
than compounds of metals lower in the activity 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
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
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
Hydroxides
decompose to form water and the metallic oxide
3.30.4
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
Oxides of most metals are stable
Oxides of potassium, sodium, calcium,
magnesium, aluminium, zinc, iron, lead and copper do not decompose.
Grey-black 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
Sulfates if heated very strongle 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
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 Salts
with water of crystallization,
hydrous salts, lose their water of crystallisation, 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 Boric
acid, H3BO3,
loses water until it decomposes to the anhydride, B2O3.
3.30.10
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
Potassium
chlorate, KClO3, decomposes above 368oC into
potassium perchlorate and oxygen gas
Potassium ferricyanide, K2Fe(CN)6,
decomposes before melting. 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
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
Silica sand
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
Ammonium dichromate is an orange coloured
crystalline substance
It starts decomposing with sparks and gives out
heat to form a green fluffy powder chromic oxide, nitrogen and water.
3.30.15
Ammonium chloride
Put ammonium chloride into the bottom of a dry test-tube and warm it
over a small flame. The solid decompses
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)