School Science Lessons
2017-10-28 SP MF LI
Please send comments to: J.Elfick@uq.edu.au

3.0.0 Chemistry, boiling point
Table of contents
3.1.0 Boiling point
3.30.0 Decomposition reactions
3.13.0 Heat substances
3.17.0 Melting point
3.0.4 Phthalates
3.20.0 Separate substances from mixtures
3.21.0 Solubility and solutions

3.1.0 Boiling point
3.6 Boiling point of inflammable liquids
3.5.1 Boiling point of sodium chloride solution
3.5.0 Boiling point of water
24.2.4 Pressure and boiling point of water
3.7 Volatility of different liquids

3.30.0 Decomposition reactions
3.29 Collect gaseous products of a burning candle
3.30.9 Decomposition of boric acid
3.30.1 Decomposition of carbonates
3.30.11 Decomposition of chlorates
3.30.15 Decomposition of chlorides
3.30.16 Decomposition of dichromates
3.30.12 Decomposition of ferricyanides
3.30.8 Decomposition of hydrates, hydrous salts
3.30.2 Decomposition of hydrogen carbonates
3.30.3 Decomposition of hydroxides
3.30.17 Decomposition of manganates
3.30.14 Decomposition of metals, metallic salts
3.30.4 Decomposition of nitrates
3.30.10 Decomposition of oxalic acid
3.30.5 Decomposition of oxides
3.30.13 Decomposition of phosphates
3.30.6 Decomposition of sulfates
3.30.7 Decomposition of sulfites
3.30.0 Substances decompose and lose mass
8.2.4 Thermal decomposition of acids

3.13.0 Heat substances
3.29 Collect and weigh the gaseous products of a burning candle
3.30.1 Decomposition of carbonates
3.30 Heat substances that decompose and lose mass when heated
3.4.13.5 Storing hygroscopic and deliquescent substances

3.17.0 Melting point
3.4 Impurities affect the melting point of a substance
3.2 Melting point of naphthalene
3.3 Melting points of naphthalene with a capillary tube

3.20.0 Separate substances from mixtures
3.24.1 Separate by chromatography, inks from mixed inks
3.24 Separate by chromatography, pigments from green leaves
3.20 Separate by distillation, distil ink to form water
3.25 Separate by heating, gases dissolved in water
3.18 Separate by melting points, tin from a tin and carbon mixture
3.22 Separate by solubility, salt from salt and sand mixture
3.19 Separate by sublimation, iodine
3.21 Separate crude oil fractions by fractional distillation
3.26 Separate two immiscible liquids of different density
3.27 Separate two solids using density differences
10.12.1 Solvent extraction of oil from peanuts

3.21.0 Solubility and solutions
3.17.0 Heat of solution
3.17.1 "Magnetic" sugar cube dissolves
3.16 Miscible liquids
3.14 Solubility and agitation
3.12 Solubility and particle size
3.13 Solubility and solvents
3.10 Solubility and temperature, solubility of salts in water
3.11 Solubility in water at a given temperature
3.9 Solubility in water of different salts

3.22.0 Shrinking volume
3.15.0 Shrinking volume
10.3.1 Shrinking mixture of liquids, lost volume
10.3.4 Container holds more
10.3.3 Container not leaking

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: A 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), 1,4-dichlorobenzene (deodorizer) (m.p. 53oC),
paraffin wax, (m.p. 45oC to 65oC), butter, soap.

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.

3.5.0 Boiling point of water
See diagram 3.5: The 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.
However, it is reported that the addition of 20 g of salt to 5 litres of water increases the boiling temperature to only 100.04oC.!

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: The boiling point of inflammable liquids
A rubber band, B capillary tube, C test-tube, D inflammable liquid, E thermometer
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 spirits, gasoline, mineral turpentine, kerosene (paraffin oil), household oil,
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.
Push 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 because the water boils at too low a temperature.

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.

Experiment
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
Chemical
Dissolves
Chemical
Dissolves
ammonium chloride
37.2 g
potassium chloride
34.0 g
barium chloride
35.7 g
potassium iodide
144.0 g
calcium chloride
42.7 g
sodium bicarbonate
9.6 g
copper (II) sulfate
20.7 g
sodium chloride
36.0 g
lead nitrate
54.4 g
sodium hydroxide
109.0 g
magnesium sulfate
25.2 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 to 5 g, 10oC to 7 g, 20oC to 12 g, 30oC to 20 g, 40oC to 26 g, 50oC to 34 g, 60oC to 43 g, 70oC to 52 g, 80oC to 61 g,
90oC to 70 g, 100oC to 80 g.
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.

3.11 Solubility in water at a given temperature
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 to w1.
The mass of water = w2 to w1 to w3.
Calculate the solubility of the sodium hydrogen carbonate as weight in grams dissolved in water at room temperature -
(w3 to w1) / (w2 to w1 to w3).
Repeat the experiment using water at a higher temperature.
The solubility of sodium bicarbonate in 100 g of water varies with temperature:
Temperature
Dissolves
Temperature
Dissolves
0oC
to 6.9 g
30oC
to 11.1 g
10oC
to 8.15 g
40oC
to 12.7 g
20oC
to 9.6 g
50oC
to 14.5 g
25oC
to 10.4 g
60oC
to 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 spirits.
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 spirits.

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.

10.3.1 Shrinking mixture of liquids, lost volume
See diagram 10.3.1: Shrinking solution
1. Put about 30 mL sodium chloride crystals into a measuring cylinder.
Add water until the measuring cylinder is exactly full.
After a few minutes, note that as the sodium chloride dissolves the liquid level drops.
2. Repeat the experiment using sucrose crystals instead of sodium chloride crystals.
The liquid level does not drop.
The sodium chloride crystals dissolve to form ions that can fit between the water molecules.
The sucrose crystals dissolve to form sucrose molecules that are much bigger than the sodium ions or chloride ions.
3. Half fill a test-tube with water.
While holding this test-tube at an angle, pour ethanol slowly from a beaker until the test-tube is full.
Hold the test-tube by placing your thumb on the mouth of the tube so that no air bubble is trapped.
The test-tube seems full.
Invert the test-tube several times while keeping a thumb on the opening.
Do not release the pressure.
The liquid level becomes lower.
The alcohol or water did not evaporate and no liquid spilt because of inverting the test-tube.
By inverting the test-tube, mixing of water and alcohol occurs, and the alcohol molecules slip between the water molecules in the
spaces between the molecules, thus making the total volume of the mixture become less.
The spaces between the molecules cannot be seen by the naked eye.

4. Repeat the experiment with methanol or rubbing alcohol (isopropanol).
Not all combinations of liquids give the strange shrinking illusion, but many water / alcohol mixtures do.
5. Make a quantitative estimate of shrinking using two identical, smaller, measuring cylinders, and a larger measuring cylinder with a
volume at least the combined volume of the smaller, measuring cylinders.
Leave the smaller measuring cylinder to stand for some time in a place of no vibration.
Put 3 drops of a blue food colouring in the water.
Use the same procedure to put 3 drops of a yellow food colouring in the methylated spirits.
Observe the spread of the two food colourings.
Put the contents of the first measuring cylinder into the large measuring cylinder followed by the contents of the second measuring
calendar.
Observe the colour change as the food colourings mix and note the final total level of the solution.
Use a pipette to bring the solution up to the total volume of the smaller measuring cylinders.
Calculate the % shrinkage, about 2 vols 1 / vol 2 = 1.9.
The ethanol also reduces the surface tension of the water.

6. Repeat the above experiment at different ambient temperatures.

10.3.3 Container not leaking
Fill a measuring cylinder by 1 / 3 with water softener pellets (Calgon, sodium hexametaphosphate) or with table salt (sodium chloride).
Add water until full and mark the liquid level with a grease pencil or rubber band.
Leave it to stand for 5 minutes and note the liquid level.
The drop in water level is not because water is absorbed by the salt.
Pour out some water before all the salt dissolves
and refill it with fresh water.
The water level drops again because the crystals of the dissolving salt breaks down into ions that can slip in between the water
molecules and make the total volume decrease.
Repeat the experiment with other salts that dissolve in water and sucrose.
The sugar molecule is large and does not ionize when dissolved in water, so that the water level will not drop.

10.3.4 Container holds more
Fill a large beaker with marbles and note the top level of the marbles.
Slowly add sand to the beaker while tapping to make the sand settle between the marbles.
Tip out the sand water and measure its volume.
Replace the sand in the beaker of marbles.
Add water to the marbles and sand.
Tip out the water and measure its volume.

3.15 Shrinking volume
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 spirits, 2. glycerine, and 3. kerosene.
Shake each test-tube.
Miscible liquids can mix in all proportions.
Alcohol and water are miscible.
Glycerine and water are miscible.
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 by melting points, tin from a tin and carbon mixture
Get tin bits by cutting a tin welding rod to pieces.
(66% of a 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), (melting point 232oC), and carbon (crushed charcoal), (melting point
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, iodine
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, distil ink to form water
| See diagram 3.20.1e: 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 3.21: Collect fractions
Crude oil is petroleum distillate oil, unrefined natural petroleum.
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.

2. Arrange the fractions in order of increasing distillation temperature:
2.1 up to 80oC,
2.2 80oC to 120oC,
2.3 120oC to 180oC,
2.4 180oC to 220oC.

3. Examine the different fractions:
3.1 The colour should change from colourless to yellow.
3.2 The viscosity should increase.
3.3 The high temperature fractions should be more difficult to ignite than the low temperature fractions.
3.4 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.
3.5. Note the dark residue remaining in the test-tubes.

3.22 Separate by solubility, salt from salt and sand mixture
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.

3.24 Separate by chromatography, pigments from green leaves
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 spirits 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 spirits solvent and the spot of green solution above the solvent level.
Leave the paper strip in the solvent until the methylated spirits 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 by chromatography, inks from mixed inks
Prepare a mixed solvent from 6 parts of water, 3 parts of methylated spirits, 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 by heating, gases dissolved in water
| See diagram 3.25: Gases dissolved in a water sample
| See diagram 3.2.25: Gas released from water
The higher the temperature of a solutions, the less a gas dissolves, if it does not react with the solvent.
1. 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.

2. Fill a round bottom flask to the top with tap water and insert a stopper with a delivery tube and collecting test-tube completely
filled with water.
Insert the stopper while holding the whole apparatus under water.
Heat the flask with a Bunsen burner.
Bubbles of gas are released from the water and travel into the collecting test-tube.
Continue heating until the water in the flask is boiling.
Collect about half a test-tube of gas from a litre of water, displaced from solution by heating.

3.26 Separate two immiscible liquids of different density
See diagram 3.26: Separation tube: A kerosene, B water, C separating tube, D clip, E rubber tubing
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 two 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, e.g. 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.

Experiments
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, magnesium ribbon
Use magnesium ribbon because magnesium powder is too reactive.
Be careful! Do not heat magnesium powder.
1. Hold a 1 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 + 1/2 O2 --> 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. However, magnesium has density 1.74 g / cm3 and melting point 650oC, but magnesium oxide has density 3.58 g /cm3 and melting
point 2 800oC, because the Mg2+-- O2- chemical bond is stronger than the Mg--Mg bond.

3.29 Collect and weigh the gaseous products of a burning candle
See diagram 3.2.29: Gaseous products of burning candle: A to filter pump, B soda lime
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.

1. 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.

2. The U-tube gains more weight than the candle loses weight, (U1-U2) > (C2-C1) for two reasons:
2.1 The candle wax combines with oxygen in the air to form carbon dioxide gas and water.
2.2 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 Heat substances that decompose and lose mass when heated
See diagram 3.30: Collecting the products of heating copper sulfate crystals
The phrase "Substances that do not decompose when heated" refers to substances that remain stable after heating constantly with a
Bunsen burner flame.
However, all compounds breakdown when heated to a high enough temperature.
1. Most carbonates decompose to form carbon dioxide and a metallic oxide.
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.
2. Hydroxides decompose to form water and a metallic oxide.
3. Nitrates decompose to form oxygen, nitrogen dioxide and a metallic oxide, except potassium nitrate and sodium nitrate that form
the nitrite and oxygen.
Lead nitrate decomposes at 470oC.
4. Nearly all oxides are stable, e.g. zinc oxide, ZnO, m.p. above 1,800oC.
5. Some sulfates decompose to form sulfur trioxide and metallic oxide.
6. Metal compounds higher in the activity series are usually more stable than compounds of metals lower in the activity series.
7. Some salts first lose water of crystallization and then become stable.
8. The 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.
9. Boric acid, H3BO3, loses water until it decomposes to the anhydride, B2O3.
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.
11. Potassium chlorate, KClO3, decomposes above 368oC into potassium perchlorate and oxygen.
Potassium ferricyanide, K2Fe(CN)6, decomposes before melting.
12. Monosodium orthophosphate, NaH2PO4.H2O, and disodium orthophosphate, (disodium hydrogen phosphate (V),
Na2HPO4.12H2O, lose water of crystallization.

Experiments
1. Heat copper sulfate crystals.
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.
Observe 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.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous copper (II) sulfate + water.

2. Prepare test-tubes containing 1 cm of (a) iodine crystals (b) sodium hydrogen carbonate granules or crystals (c) silica sand (d) 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.

3. 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.

4. Heat sodium hydrogen carbonate crystals.
The crystals lose water and carbon dioxide, and at 100oC are converted to sodium carbonate.

5. 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.

6. Heat zinc oxide in a crucible.
The colour changes from white to yellow 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.

3.30.1 Decomposition of carbonates
All the Group 2 carbonates undergo thermal decomposition to give the metal oxide and carbon dioxide gas.
Most carbonates 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
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.
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
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 of very active metals are stable when heated, e.g. sodium hydroxide, potassium hydroxide.
Hydroxides of less active metals decompose with strong heat to form water and the metallic oxide.
Mg(OH)2 (s) --> MgO (s) + H2O (g)

3.30.4 Decomposition of nitrates
All the Group 2 nitrates undergo thermal decomposition to give the metal oxide, nitrogen dioxide and oxygen.
Nitrates decompose to form oxygen gas, nitrogen dioxide and a metallic oxide, except potassium nitrate and sodium nitrate that form
the nitrite and oxygen.
Lead nitrate decomposes at 470oC.
2Ca(NO3)2 (s) --> 2CaO + 4 NO2 (g) + O2 (g)
colourless --> white
2Cu(NO3)2 (s) --> 2CuO + 4NO2 (g) + O2 (g)
blue --> black
2Pb(NO3)2 (s) --> 2PbO + 4NO2 (g) + O2 (g)
colourless --> yellow
Lead nitrate decomposes at 470oC.
2Zn(NO3)2 (s) --> 2ZnO + 4NO2 (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.
2KNO3 (s) --> 2KNO2 (s) + O2 (g)
colourless --> colourless
Sodium nitrate melts as 316oC.
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 above 1,800oC.
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.
Some sulfates decompose to form sulfur trioxide and metallic oxide.
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 hydrates, hydrous salts
Some salts first lose water of crystallization and then become stable.
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.

Experiments
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 chlorates
Potassium chlorate, KClO3, decomposes above 368oC into potassium perchlorate and oxygen gas.

3.30.12 Decomposition of ferricyanides
Potassium ferricyanide, K2Fe(CN)6, decomposes before melting.

3.30.13 Decomposition of phosphates
Monosodium orthophosphate, NaH2PO4.H2O, and disodium orthophosphate [disodium hydrogen phosphate (V)] Na2HPO4.12H2O,
both lose water of crystallization.
10KClO3 <--> 6KClO4 + 4KCl + 3O2

3.30.14 Decomposition of metals, metallic salts
Metal compounds higher in the activity series are usually more stable than compounds of metals lower in the activity series.
The 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.

3.30.15 Decomposition of chlorides
Sodium and magnesium chlorides are solids with high melting points.
The other chlorides are liquids or low melting point solids.

3.30.16 Decomposition of dichromates
Ammonium dichromate decomposes on heating
(NH4)2Cr2O7 + heat --> Cr2O3 + N2 + 4H2O

3.30.17 Decomposition of manganates
Potassium permanganate decomposes into potassium manganate, manganese dioxide and oxygen gas.
2KMnO4 --> K2MnO4 + MnO2 + O2

3.0.4 Phthalates
Phthalates, C6H4(COOR1)(COOR2) are esters of phthalic acid, [C6H4(COOH)2].
Phthalates are used as chemical plasticizers and to make polymers and alkyd resins, and to soften polyvinyl chloride, (PVC).
They are used to make shower curtains, PVC piping and inflatable products.
Alkyd resins are usually made from glycerol and phthalate anhydride for paint, car parts, electric switches and insulators, electronic
components and television parts.
Phthalates are not chemically bound to the plastics they are added to, so they may be continuously released to the air as plastics harden
over time.
So their use is being phased out in some countries because of concern about possible risk to foetuses and young children.
Phthalates include the following
Bisphenol-A, (BPA), Epoxy resin polymer
DPB, dibutyl phthalate | DNOP, di-n-octyl phthalate | DiNP diisononyl phthalate | DEP, diethyl phthalate | BBzP, benzyl butyl phthalate
DEHP, di-2-ethylhexyl phthalate | DiDp, diisodecyl phthalate | DnHP, di-n-hexyl phthalate | DMP, dimethyl phthalate |
DnOP, di-n-octyl phthalate

The Danger of Phthalates by Dr Ameeta Gajjar
"Fitness First Australia" Sept / Oct 2013
See diagram 16.3.5.1: Plastics recycling code
Phthalates are industrial compounds added to different kinds of plastics to make them softer and more durable.
They have found their way into everything from food containers and toys to vinyl flooring, shoes and even cosmetics and shampoos.
Since phthalates are not bonded within the product they are used in, they can leach out and be absorbed orally from food and drinks,
through the skin, via inhalation and even medical injection procedures.
We are only just waking up to the adverse impact of phthalates on our health.
As more links between phthalates and medical problems are established there is increasing feeling that phthalates could be responsible
for one of the great silent health pandemics of our time.
While some phthalates such as DEHP are now banned for certain products, they could still be having an effect as they degrade slowly.
Even very small amounts of phthalates have been shown to interact with other toxins and have cumulative effects.
The common phthalates are DEHP (cheap to use), DBP, DIDP and DINP.
DEHP was banned in Australia in 2011 for certain products, including children's toys, (less than 36 months of age), and eating utensils,
if the concentration of DEHP exceeded 1 percent.
So DEHP can still be present, quite legally in minute amounts in many plastic products.
DEHP, di-2-ethylhexyl phthalate, [dioctyl phthalate, DOP], C6H4(C8H17COO)2, most common phthalate plasticizer
DBP, dibutyl phthalate, C16H22O4, banned in cosmetics in EU, suspected endocrine disruptor
DINP, iisononyl phthalate, C26H42O4, EU restriction
Why you should be concerned
Phthalates have been linked to various medical problems.
1. Cancer
Phthalates have been found to mimic our hormones and with other environmental pollutants, are known as "endocrine
disrupting chemicals."
Studies have shown a link between breast cancer and phthalate exposure.
2. Obesity
Metabolic syndrome and insulin resistance associations have been found between urinary phthalates and increased waist
circumference and insulin resistance in US males and adolescents.
3. Allergies /Asthma
Studies have found an association between allergies in children and phthalates.
There was also an association with asthma.
Other symptoms from phthalates exposure include low birth weight in infants, abnormal foetal development and behavioural and /or
earning difficulties.
How to avoid phthalates
Avoid plastic containers and bottles with the recycling codes 1, 3, 6 or 7.
If they have no number, avoid those by checking the bottom of the product for the number in a triangle.
These are the typical plastics to avoid:
Recycling code 1 are PET, (polyethylene terephthalate): Found in water bottles, soft drink bottles.
Recycling code 3 are PVC, (polyvinyl chloride): Found in cling film, food packaging, cooking oil bottles and toys.
Recycling code 6 are PS, (polystyrene) Found in disposable plates /cups / trays, egg cartons and takeaway containers.
Recycling code 7 are PC, (polycarbonate) and others.
Found in baby bottles, large water containers.
Practical tips on reducing your phthalate exposure
1. Reduce use of plastic water bottles - use glass or stainless steel bottles instead.
2. Do not heat foods and drinks in plastic in microwave ovens, particularly baby bottles.
3. Avoid plastic cookware - opt for Pyrex or ceramic containers, especially for heating and storing.
4. Buy natural, certified organic cosmetics and personal care products, (nearly 900 chemicals used in cosmetics are known to be toxic,
including phthalates, SLS and parabens).
5. Open the car windows!
That wonderful new car smell is partly from phthalates, with their concentration getting worse when the car gets hot.
6. Choose safer plastic toys, (some are mean to be phthalate free), or, just avoid plastic toys.