UNChem1
School Science Lessons
Chemistry
2012-05-10 SPwp
Please send comments to: J.Elfick@uq.edu.au
Table of contents
3.31.0 Hygroscopic, deliquescent and efflorescent
chemicals
3.2.0 Identify pure substances
3.9.0 Solubility and solutions
3.28.0 Substances that gain mass when heated
3.30.0 Substances that lose mass when heated
3.31.0 Hygroscopic, deliquescent
and efflorescent chemicals
3.31 Hygroscopic, deliquescent and
efflorescent chemicals
3.31.1 Expose different salts
to the air
3.31.2 Expose sodium carbonate decahydrate,
washing soda, to the air
5.3.9 Tests for presence of water,
cobalt (II) chloride paper
3.2.0 Identify pure substances
3.6 Boiling point of inflammable liquids, ethanol,
acetone
3.5.1 Boiling point of sodium chloride solution
3.5 Boiling point of water
3.2 Melting point of naphthalene
3.3 Melting point of naphthalene using a capillary
tube
3.4 Melting points of substances changed by impurities
3.8 Pressure of the atmosphere affects the boiling
point
3.7 Volatility of liquids
3.9.0 Solubility and solutions
3.17 Heat of solution, sodium hydroxide, potassium
hydroxide, ammonium chloride, urea
3.17.1 "Magnetic" sugar cube dissolves
3.16 Miscible liquids, methylated spirit, glycerine,
kerosene
3.9 Solubility in water of different salts
3.10 Solubility and temperature, solubility of
salts in water
3.11 Solubility of substances in water at a given
temperature
3.12 Solubility and particle size, sodium chloride,
copper (II) sulfate
3.13 Solubility and solvents, sodium chloride,
methylated spirit
3.14 Solubility and agitation, cane sugar
3.15 Volume of solutions, sodium chloride
3.28.0 Substances that gain
mass when heated
3.28.0 Heat substances that gain mass when heated, copper foil:
3.28.4 Collect and weigh the gaseous products
of a burning candle
3.28.3 Substances that gain mass when heated,
prepare calcium oxide
3.28.2 Substances that gain mass when heated,
prepare lithium oxide
3.28.1 Substances that gain mass when heated,
prepare magnesium oxide
3.30.0 Substances that lose
mass when heated
3.30.0 Heat substances that decompose and lose mass, thermal decomposition
3.30.15 Decomposition of ammonium chloride
3.30.14 Decomposition of ammonium dichromate
3.30.9 Decomposition of boric acid
3.30.1 Decomposition of carbonates
3.30.2 Decomposition of hydrogen carbonates
(bicarbonates)
3.30.8 Decomposition of hydrous salts
3.30.3 Decomposition of hydroxides
3.30.4 Decomposition of nitrates
3.30.10 Decomposition of oxalic acid
3.30.5 Decomposition of oxides
3.30.11.2 Decomposition of phosphates
3.30.11 Decomposition of potassium chlorate
3.30.11.1 Decomposition of potassium ferricyanide
3.30.6 Decomposition of sulfates
3.30.7 Decomposition of sulfites
3.30.13 Heat silica sand
3.30.12 Sublimation, iodine, camphor, naphthalene
3.2 Melting point of naphthalene
See diagram 3.2.2: An approximate melting
point
Put 2 cm of naphthalene flakes in a test-tube. Hold a thermometer with
its bulb in the naphthalene. Use a small flame to heat the test-tube gently
and watch the thermometer reading. To find the melting range, note the
temperature when the naphthalene melts. Leave to cool and note the temperature
when the naphthalene solidifies. To find the melting point, calculate the
average of these two values. The melting point of pure naphthalene is 80.5oC.
3.3 Melting point of naphthalene
using a capillary tube
See diagram 3.2.3: A more accurate method
Make a capillary tube by drawing out a glass tube over a hot flame.
Put a very small amount of naphthalene in a capillary tube sealed at one
end. Attach a thermometer to the capillary tube, a sealed end down. Put
the thermometer and capillary tube in a container of water and slowly heat
the water while stirring with the thermometer and capillary tube. Do not
let water enter the capillary tube. To find the melting range, note the temperature
when the naphthalene melts, leave to cool, and note the temperature when
the naphthalene solidifies. To find the melting point, calculate the average
of these two values.
Repeat the experiment with stearic acid, m.p. 69oC, palmitic
acid, m.p. 63oC, butter, soap, 1,4-dichlorobenzene (deodorizer)
m.p. 53oC, paraffin wax, m.p. 45oC - 65oC,
sodium thiosulfate pentahydrate 48.3oC.
3.4 Melting point of substances
changed by impurities
Mix stearic acid with the naphthalene to make the naphthalene impure.
Note changes in the melting point. Impurities lower the melting point.
3.5 Boiling point of water
See diagram 3.2.5: Boiling point of water
The standard atmosphere (symbol: atm) 101325 Pa, for many purposes has
been replaced by the bar 100 kilopascals (100 kPa) = 1×105 N /
m2 = 0.987 atm. (1 bar = 100 kPa, 100,000 Pa) The standard boiling
point is now the temperature at which boiling occurs under a pressure of
1 bar. On the top of Mount Everest, at 8,848 m elevation, the pressure is
about 260 mbar (26.39 kPa) and the boiling point of water is 69 °C. The boiling
point decreases 1 °C every 285 m of elevation.
1. Pour water into a test-tube. Hold a thermometer with its bulb just
under the water. Heat very slowly by moving the test-tube in and out of
the flame or add boiling chips, anti-bumping granules. Heat the water gently
until it boils. Record the temperature. Note the same temperature in all
parts of the test-tube. Note any change in the reading if the thermometer
touches the bottom of the test-tube. The water must cover the bulb of the
thermometer and the bulb must not touch the sides of the test-tube.
2. Show that the boiling point of water does not depend on the size
of the container. Repeat the experiment with a large container. Heat the
water quickly. The water first starts to boil near the bottom and sides
of the container. Note the temperature in different parts of the container.
Note any change in the reading if the thermometer touches the bottom of
the container. The boiling point is the same in small and large containers.
3.5.1 Boiling point of sodium
chloride solution
1. A solution of sodium chloride in water boils at a higher temperature
and has a lower freezing point than pure water. Use freezing points and
boiling points to find the purity of substances. Use three test-tubes containing
the same volume of water. Add some sodium chloride to the second test-tube.
Keep adding sodium chloride to the third test-tube until no more dissolves
to produce a saturated solution at that temperature. Join the test-tubes
with an elastic band. Heat the test-tubes equally over a Bunsen burner.
The first test-tube containing only water boils first. The second test-tube
containing some sodium chloride boils next. The third test-tube containing
the saturated solution of sodium chloride boils last.
2. Put a beaker containing demineralized water in a broad pan containing
a concentrated salt solution. Slowly heat the broad pan and note that
the demineralized water boils first.
3.6 Boiling point of inflammable
liquids, ethanol, acetone
See diagram 3.2.6: Boiling point of inflammable
liquids
1. Do not use a Bunsen burner to find the boiling point of inflammable
liquids, e.g. ethanol, b.p. 78.4oC and acetone, b.p. 56oC.
Use an electric hot plate or use the following method. Pour 2 cm of the
inflammable liquid into a test-tube in an empty container. Place a thermometer
in the test-tube with its bulb in the liquid. Boil water in an electric jug
or on an electrical hot plate. Pour the hot water into the container so
that the level is higher than the inflammable liquid in the test-tube. Stir
the inflammable liquid gently with the thermometer and read thermometer when
the inflammable liquid boils. [It is not good practice to stir liquids with
thermometers!]
2. Use a very small test-tube or seal one end of a piece of glass tubing,
8 cm length and 3 cm external diameter. Put the inflammable liquid into
this test-tube. Put a capillary tube, sealed at one end, into the inflammable
liquid with the sealed end up and the open end down in the inflammable
liquid. Use a rubber band to attach the test-tube containing inflammable
and capillary tube to the bulb of a thermometer. Hold the apparatus in
a container of water and heat gently with an electric hot plate. When the
temperature rises, bubbles slowly come out of the capillary tube. At the
boiling point the bubbles suddenly come out as a steady stream. Read the
temperature. Let the water cool and read the temperature again when the
steady stream of bubbles stops. Calculate the boiling point as the average
of the two readings.
3. The boiling point of ethanol is 78oC. Heat water in a
kettle. Turn off the gas or extinguish the flame. Pour 2 cm of methylated
spirit into a test-tube. Put the test-tube in a beaker. Put a thermometer
into the methylated spirit deep enough to cover the bulb. Pour hot water
into the beaker so that the water level is higher than the level in the test-tube.
Stir gently with thermometer. When boiling starts, record the temperature.
4. Boiling point of acetone. Repeat the experiment with propanone (acetone).
BE CAREFUL! Highly flammable! THE FLASH POINT IS -20oC!
Boil the propanone (acetone) by placing the test-tube in hot water. Do
not allow flames in the room. The boiling range of 98% propanone (acetone)
is 55.5oC to 56.5oC. Nail polish remover formerly
contained propanone (acetone) but not in modern manufacture.
3.7 Volatility of different
liquids
Evaporation is the movement of particles from the surface of a liquid
to the gas state, when below the boiling point. Volatile liquids evaporate
readily at room temperatures.
1. Select liquids from the laboratory, e.g. water, methylated spirit,
gasoline, mineral turpentine, kerosene (paraffin oil) household machine
oil, car oil, vinegar, vanilla essence, eucalyptus oil, glycerine. Wet a
5 cm piece of absorbent paper with a liquid, Write the name of the test liquid
in pencil. Attach the piece of paper to a horizontal string. Examine the
paper every ten minutes, every two hours and each day.
2. Repeat the experiment with perfumes. Smell the paper every ten minutes,
every two hours and each day. Some perfumes soon disappear, but others
last for days. Record the relative "person-attracting" capacity for each
perfume.
3.8 Pressure of the atmosphere
affects the boiling point
See diagram 3.2.8: Decreasing the pressure
on boiling water
1. Partially fill a flask with water and heat it until the water boils
for a few minutes. Remove the heat source so that the water stops boiling
then stopper the flask which now contains saturated water vapour water vapour
at atmospheric pressure Pour cold water over the flask. Some water vapour
condenses on the walls of the flask and the water boils again. The cold
water caused the vapour pressure in the flask to be reduced by condensation
to below the saturation pressure of the water in the flask at that temperature.
So the temperature at which water boils is lowered when the pressure is
reduced.
2. Put water in a sidearm test-tube or in a round bottom flask with
a one-hole stopper. Insert a thermometer through a hole in the stopper so
that the bulb of thermometer reaches, but does not touch, the bottom of the
test-tube or flask. Add boiling chips to prevent bumping. Boil the water
and read the temperature. Stop heating. Connect a water pump to the sidearm
or to the second hole of the two-holes stopper. When the water stops boiling,
turn on the water pump to reduce the pressure. Read the temperature, heat
to boiling and read the temperature again.
3. When a liquid is at boiling point, its vapour pressure is equal
to the pressure of the atmosphere. Boil water on a high mountain and note
the boiling point. People who climb Mount Everest complain that they cannot
get a good cup of tea.
4. Some teachers play a silly trick on students. The ask a student to
feel the sides of the flask when the water is first boiling. When the second
boiling occurs at much lower temperature, the teacher spills some water
on the hand of the student who pulls way that hand thinking that the boiling
water is still very hot.
3.9 Solubility in water of different
salts
In this document the word "solution" refers to substances dissolved
in water, i.e. aqueous solutions. A solvent is a liquid that dissolves another
substance, the solute, to form a solution. The three ways to increase the
rate at which a solid dissolves in water are as follows: 1. grinding the
solid until finely divided 2. shaking the solution while the solid is dissolving,
and 3. heating the solution. Try to dissolve 5 g of different salts each
in 15 mL of water in a test-tube. Attach a stopper and shake vigorously.
Solubility is a characteristic of a particular substance. Classify each
salt as soluble or slightly soluble or insoluble. The solubility of a salt
can be expressed as the number of grams able to dissolve in 100 g of water
at 20oC, e.g. ammonium chloride 37.2 g, barium chloride 35.7 g,
calcium chloride 42.7 g, copper (II) sulfate 20.7 g, lead nitrate 54.4 g,
magnesium sulfate 25.2 g, potassium chloride 34.0 g, potassium iodide 144.0
g, sodium hydrogen carbonate (sodium bicarbonate) 9.6 g, sodium chloride
36.0 g, sodium hydroxide 109.0 g, sodium nitrate 87.5 g.
3.10 Solubility and temperature,
solubility of salts in water
The solubility of a potassium dichromate in 100 g of water varies with
temperatures: 0oC - 5 g, 10oC - 7 g, 20oC
- 12 g, 30oC - 20 g, 40oC - 26 g, 50oC
- 34 g, 60oC - 43 g, 70oC - 52 g, 80oC -
61 g, 90oC - 70 g, 100oC - 80 g.
1. Show that a saturated solution contains less dissolved solid at
a lower temperature than at a higher temperature. Make a 50 mL saturated
solution of potassium dichromate or potassium nitrate at 60oC.
Pour the clear solution into a clean container and keep the temperature
of this container at 40oC until crystals stop forming. Pour
the clear solution from this container into another clean container. Do not
pour crystals into the container. Leave to cool and note more crystals forming
as the solution cools.
2. Repeat the experiment with sodium chloride. This salt is unusual
because the solubility hardly changes with change in temperature.
3.11 Solubility of a substance
in water at a given temperature
1. Add a teaspoonful of powdered milk to water with ice floating in
it, water at room temperature, warm water, boiling water. Observe the time
taken for the powdered milk to dissolve in the water at different temperatures.
2. Add sodium hydrogen carbonate (sodium bicarbonate) to 100 g of water
in a container while stirring. Keep adding sodium carbonate until no more
solute will dissolve. Decant the clear saturated solution. Read the temperature
of the saturated solution, i.e. room temperature. Weigh a clean evaporating
dish, w1. Add some clear saturated solution and weigh again, w2. Carefully
evaporate the solution in the evaporating dish to dryness and weigh again,
w3. The mass of the dissolved sodium hydrogen carbonate = w3 - w1. The
mass of water = w2 - w1 - w3. Calculate the solubility of the sodium hydrogen
carbonate as weight in grams dissolved in water at room temperature (w3
- w1) / (w2 - w1 - w3).
Repeat the experiment using water at a higher temperature.
3. The solubility of sodium bicarbonate in 100 g of water varies with
temperature: 0oC - 6.9 g, 10oC - 8.15 g, 20oC
- 9.6 g, 25o- 10.35 g, 30oC - 11.1 g, 40oC
- 12.7 g, 50oC - 14.45 g, 60oC -16.4 g.
3.12 Solubility and particle
size, sodium chloride, copper (II) sulfate
Show that small particles dissolve faster than large particles.
1. Add coarse table salt to a first test-tube half filled with water.
Add the same quantity of fine table salt to a second test-tube that contains
the same volume of water. Shake both test-tubes equally and simultaneously.
Note the amount of undissolved table salt left in each test-tube.
2. Use two equal samples of large crystals of copper (II) sulfate.
Grind one sample into a fine powder. Put both samples into equal quantities
of water in separate test-tubes and shake. Compare the rates at that the
different samples dissolve and cause the water to change colour.
3.13 Solubility and solvents,
sodium chloride, methylated spirit
1. Fill two test-tubes one third full with water and another with methylated
spirit. To each test-tube add 1 g sodium chloride, attach a stopper and
shake. Sodium chloride dissolves readily in water, but not so readily in
methylated spirit.
2. Add sodium chloride crystals to a dilute solution of sugar in water
and note whether the crystals dissolve. Drop crystals of potassium dichromate
into the solution. Note whether the solution changes colour. Colour change
shows that potassium dichromate is also dissolving. The presence of one
dissolved substance does not prevent other substances dissolving in the
solution. Unless the concentrations are high, one solute does not affect
the solubility of other solutes in the solution.
3.14 Solubility and agitation,
cane sugar
Prepare two equal samples of cane sugar. Put one sample of cane sugar
into a test-tube half full of water. Add a stopper and shake the test-tube
until all the sugar dissolves. Put the other sample of cane sugar into a
test-tube. Very slowly add the same volume of water as in the first test-tube.
Leave to stand. Most of the sugar has not dissolved but, if left to stand
for long enough, all the cane sugar will dissolve as in the first test-tube.
3.15 Volume of solutions, sodium
chloride
1. Fill a small, narrow neck flask with water to a level in the neck
and mark this level. Add sodium chloride to the water with continual shaking
until the solution is saturated and no more dissolves. Note the new level
of the liquid. The volume of the solution is only slightly greater than
the original volume of the water.
2. Close one end of a glass delivery tube about 30 cm long. Fix it
upright, half fill it with water and mark the level of the water. Slowly
add alcohol to fill the delivery tube. The water and the alcohol fill
equal lengths in the tube. Shake the tube thoroughly to mix the water and
alcohol. The new level of the solution in the tube shows a slight decrease
in total volume.
3.16 Miscible liquids, methylated
spirit, glycerine, kerosene.
Put 10 mL of water in three test-tubes. Add 1 mL of: 1. methylated
spirits, 2. glycerine, and 3. kerosene. Shake each test-tube. Miscible
liquids can mix in all proportions. 1. Alcohol and water are miscible.
2. Glycerine and water are miscible. 3. Kerosene and water are not miscible,
immiscible.
3.17 Heat of solution, sodium
hydroxide, potassium hydroxide, ammonium chloride, urea
Dissolve some equal quantities of sodium hydroxide, potassium hydroxide,
ammonium chloride and urea in separate test-tubes half full of water.
Feel the test-tubes and note any change in the temperature. Sodium hydroxide
and potassium hydroxides dissolve in water with an increase in temperature.
Ammonium chloride and urea absorb heat from their surroundings when dissolving
in water.
3.17.1 "Magnetic" sugar cube
dissolves
Fill a large dish with water Wait until the water is absolutely still
then lower a matchstick into the centre of the water. Carefully dip a sugar
cube in the water near the edge of the dish. The matchstick moves towards
the dissolving sugar cube. When the sugar dissolves in the surface water,
the surface water becomes heavier and falls to be replaced by surface water
flowing towards the sugar cube, carrying the matchstick with it.
3.28 Heat substances that gain
mass when heated, copper foil
Cleaned copper is brown-red. In moist air the surface turns green due
to oxidation. The green surface is called a patina. It also forms on old
unpolished bronze.
1. Clean a piece of copper foil with steel wool. Hold it in a flame
with a pair of tongs. The black copper (II) oxide looks like carbon. To
test the substance, drop dilute sulfuric acid on it, then heat it. Blue copper
(II) sulfate forms. Test some powdered carbon. No colour change occurs.
2. Show that something is added to the copper from the air. Use a sensitive
balance to weigh the copper before and after heating.
3. Use two identical hard glass test-tubes with one-hole stoppers fitted
with bent delivery tubes. Fix both test-tubes to a stand so that the
test-tubes slope down with the ends of the delivery tubes under water
in a beaker. Put copper foil in the first test-tube and heat with a hot
burner flame. After two minutes, heat the empty second test-tube. Move the
burner regularly between the two test-tubes until no more bubbles come out
of the ends of the delivery tubes. Stop heating both test-tubes. As the test-tubes
cool, they suck water up the delivery tube. The test-tube containing the
copper (II) oxide sucks up more water.
3.28.1 Substances that gain
mass when heated, prepare magnesium oxide
Use magnesium ribbon because magnesium powder is too reactive. Be careful!
Do not heat magnesium powder. Magnesium has density 1.74 g / cm3
and melting point 650oC, but magnesium oxide has density
3.58 g /cm3 and melting point 2800oC because the
Mg2+-- O2- chemical bond is stronger than the Mg
-- Mg bond.
1. Hold a 10 cm strip of magnesium ribbon in a pair of tongs. Place
the ribbon in a Bunsen burner flame until it starts to burn. Be careful!
Magnesium burns with a very bright white light. Magnesium ribbon corrodes
slightly in air and burns with an intense white flame to form a white ash
of magnesium oxide.
Mg + 1 / 2O2 --> MgO
2. Clean 25 cm of magnesium ribbon and cut into pieces 1 cm long. Put
the pieces into a crucible with a lid. Weigh the crucible + lid + contents
= W1. Put the crucible on a pipe clay triangle on a tripod stand. Heat gently
then strongly. Use tongs to raise the lid. The magnesium darkens before
it melts. When the magnesium starts to burn, put the lid back on the crucible
and remove the burner. Every few seconds raise the lid slightly to let more
air enter. Do not let white magnesium oxide smoke escape. When the magnesium
does not burn after you raise the lid, remove the lid and heat the crucible
strongly. Hold the lid ready in case the magnesium starts to burn again.
Let the crucible cool. Again weigh the crucible + lid + contents = W2. Note
W2 > W1. The formation of magnesium oxide causes the increase in weight.
3.28.2 Substances that gain
mass when heated, prepare lithium oxide Heat pieces of lithium metal
shot on a metal spoon (deflagrating spoon). Note the violet glow when
it starts to burn, then put the burning lithium in oxygen gas.
3.28.3 Substances that gain
mass when heated, prepare calcium oxide
Heat a shaving of calcium metal in a crucible and heat it with a Bunsen
burner for 10 -15 minutes because it is difficult to ignite.
3.28.4 Collect and weigh
the gaseous products of a burning candle
See diagram 3.2.29: Gaseous products of
burning candle
Candle wax is a mixture of different alkanes (paraffins) saturated
hydrocarbons with general formula CnH2n+2 that
are solid at room temperature. Soda lime is a grey-white mixture of sodium
hydroxide and calcium hydroxide as granules or powder that absorbs the
products of combustion, carbon dioxide and water. Use soda lime instead
of sodium hydroxide because soda lime is not deliquescent. Weigh a candle,
C1. Weigh a U-tube containing granules of soda lime, U1. Put a candle
under an inverted glass filter funnel connected to one arm of the U-tube.
Attach a filter pump to the other arm to draw air through the U-tube. Light
the candle and turn on the filter pump to draw air over the candle. Let
the candle burn for five minutes. Extinguish the candle and disconnect the
filter pump. Weigh the candle again, C2. The candle has lost weight, C1-C2.
When the U-tube is cool, weigh it again, U2. The U-tube containing the soda
lime has gained weight, U2-U1. The U-tube gains more weight than the candle
loses weight (U1-U2) > (C2-C1) for two reasons: 1. The candle wax combines
with oxygen gas in the air to form carbon dioxide gas and water. 2. The air
sucked in by the filter pump contains some water vapour absorbed by the soda
lime. To measure the weight of water absorbed from the air, in a control
experiment, repeat the experiment for the same period, but without the
candle.
3.30.0 Substances that decompose
and lose mass when heated, thermal decomposition
See diagram 3.2.32: Collecting the products
of heating copper sulfate crystals
Substances that remain stable after heating constantly with a Bunsen
burner flame may be listed under the heading "Substances that do not
decompose when heated". However, all compounds break down when heated to
a high enough temperature. Metal compounds higher in the reactivity series
are usually more stable than compounds of metals lower in the reactivity
series. Salts that remain stable when heated constantly with a Bunsen burner
flame are calcium sulfate, potassium chloride, potassium sulfate, sodium
carbonate, sodium chloride, and sodium sulfate. Ammonium oxalate (NH4COO)2,
and ammonium dichromate (NH4)2Cr2O7,
decompose before melting. Ammonium sulfate (NH4)2SO4,
decomposes above 280oC.
3.30.1 Decomposition
of carbonates
Carbonates mostly decompose to form carbon dioxide and a metallic oxide.
Sodium carbonate and potassium carbonate do not decompose when heated
to a high temperature. CaCO3 (s) --> CaO (s) + CO2
(g)
white calcium carbonate --> white calcium oxide + carbon dioxide
CuCO3 (s) --> CuO (s) + CO2 (g)
blue-green --> black
MgCO3 (s) --> MgO (s) + CO2 (g)
white --> white
PbCO3 (s) --> PbO (s) + CO2 (g)
white --> yellow
ZnCO3 (s) --> ZnO (s) + CO2 (g)
white --> white (cold) or yellow (hot)
Ammonium carbonate may decomposes completely without heating when exposed
to the air, to give ammonia, water and carbon dioxide.
(NH4)2CO3 (s) --> 2NH3
(g) + H2O (vapour) + CO2 (g)
colourless
3.30.2 Decomposition
of hydrogen carbonates (bicarbonates)
Hydrogen carbonates (hydrogen carbonates, bicarbonates) decompose to
form the metal carbonate, water and carbon dioxide.
Calcium bicarbonate and magnesium bicarbonate can exist only as a solution
that on heating form the metal carbonate, water and carbon dioxide.
Sodium hydrogen carbonate, NaHCO3 (sodium bicarbonate) begins
to lose carbon dioxide at 50oC to form sodium carbonate. A
solution of a sodium hydrogen carbonate begins to lose carbon dioxide
at 20oC.
Heat sodium hydrogen carbonate crystals. The crystals lose water and
carbon dioxide, and at 100oC are converted to sodium carbonate.
2NaHCO3 (s) --> Na2CO3 (s) + CO2
(g) + H2O (vapour)
colourless --> colourless
Ca(HCO3)2 (aq) --> CaCO3 (s) +
CO2 (g) + H2O (vapour)
Mg(HCO3)2 (aq) --> MgCO3 (s) +
CO2 (g) + H2O (vapour)
2KHCO3 (s) --> K2CO3 (s) + CO2
(g) + H2O (vapour)
3.30.3 Decomposition
of hydroxides
Hydroxides decompose to form water and the metallic oxide
3.30.4 Decomposition
of nitrates
Nitrates decompose to form oxygen gas, nitrogen dioxide and a metallic
oxide. 2Ca(NO3)2 (s) --> 2CaO + 4 NO2
(g) + O2 (g)
colourless --> white
2Cu(NO3)2 (s) --> 2CuO + 4 NO2
(g) + O2 (g)
blue --> black
2Pb(NO3)2 (s) --> 2PbO + 4 NO2
(g) + O2 (g)
colourless --> yellow
Lead nitrate decomposes at 470oC.
2Zn(NO3)2 (s) --> 2ZnO + 4 NO2
(g) + O2 (g)
colourless --> white (cold), yellow (hot)
Potassium nitrate and sodium nitrate first melt and then decompose
to give the metal nitrite and oxygen gas. Potassium nitrate melts at 336oC.
Sodium nitrate melts as 316oC. 2KNO3 (s) -->
2KNO2 (s) + O2 (g)
colourless --> colourless
2NaNO3 (s) --> 2NaNO2 (s) + O2
(g)
colourless --> yellow
Silver nitrate decomposes to give the metal, nitrogen dioxide and oxygen
gas. 2AgNO3 (s) --> 2Ag (s) + 2NO2 (g) + O2
(g)
colourless --> silver
Ammonium nitrate decomposes to form water vapour and nitrous oxide,
N2O (laughing gas) so the ammonium nitrate disappears.
NH4NO3 (s) --> N2O (g) + H2O
(g)
colourless
3.30.5 Decomposition
of oxides
Oxides of most metals are stable. Oxides of potassium, sodium, calcium,
magnesium, aluminium, zinc, iron, lead and copper do not decompose. Black-grey
silver oxide decomposes into the metal and oxygen gas. 2Ag2O
(s) --> 4Ag (s) + O2 (g)
silver oxide --> silver + oxygen
Heat zinc oxide in a crucible. Zinc oxide becomes yellow when hot and
white when cold but no change in weight occurs. The substance does not decompose
and does not gain anything from the air or lose anything to the air. Zinc
oxide has melting point above 1,800oC.
ZnO (s) <--> ZnO (s)
white (cool) yellow (hot) Thermal decomposition of higher oxides of
lead
2PbO2 (s) --> 2PbO (s) + O2 (g)
brown lead dioxide --> yellow lead oxide + oxygen gas
2Pb3O4 (s) --> 6PbO (s) + O2 (g)
red trilead tetroxide --> yellow lead oxide + oxygen gas
3.30.6 Decomposition
of sulfates
Sulfates if heated very strongly may decompose to form the metallic
oxide, sulfur dioxide and oxygen gas
Put 4 cm of crushed blue copper (II) sulfate crystals in a dry test-tube
fitted with a one-hole stopper and delivery tube. Heat the dry test-tube
and crystals gently. Note whether vapour collects on the cooler parts
of the dry test-tube and whether any liquid collects in the receiving test-tube.
Note any change of colour of the crystals from blue to white. Identity
the liquid in the receiving test-tube by measuring the boiling point.
When all the copper (II) sulfate crystals have become white and the dry
test-tube has cooled, pour the liquid in the receiving test-tube back
on the white crystals. Note whether the blue colour of the crystals is
restored and if any heat is given off.
2CuSO4 (s) --> 2CuO (s) + 2SO2 (g) + O2
(g)
grey white --> black
2PbSO4 (s) --> 2PbO (s) + 2SO2 (g) + O2
(g)
white --> yellow
2ZnSO4 (s) --> 2ZnO (s) +2SO2 (g) + O2
(g)
white --> white (cold) yellow (hot)
3.30.7 Decomposition
of sulfites
Sulfites mostly decompose into the metal oxide and sulfur dioxide
Sulfites of sodium and potassium do not decompose when heated. CaSO3
(s) --> CaO (s) + SO2 (g)
white --> white
MgSO3 (s) --> MgO (s) + SO2 (g)
white --> white
ZnSO3 (s) --> ZnO (s) + SO2 (g)
white --> white (cold) yellow (hot)
3.30.8 Decomposition
of hydrous salts
Salts with water of crystallization, hydrous salts, lose their water
of crystallization, and become anhydrous powders and then become stable
They also lose their crystalline shape and sometimes their colour.
Prepare test-tubes containing 1 cm of 1. iodine crystals 2. sodium hydrogen
carbonate granules or crystals 3. silica sand 4. zinc oxide. Fix a cotton
wool plug in the mouth of each test-tube to prevent loss of solid during
heating, then weigh each test-tube. Heat each test-tube and cotton wool plug
thoroughly and weigh it again. Note any change in weight because of the
loss of water of crystallization.
(blue) copper (II) sulfate crystals + heat < = > (white) anhydrous
copper (II) sulfate + water.
CuSO4.5H2O (s) --> CuSO4 (s) +
5H2O (vapour)
blue --> grey white
Na2CO3.10H2O --> Na2CO3
(s) + 10H2O (vapour)
colourless --> white
3.30.9 Decomposition
of boric acid
Boric acid, H3BO3, loses water until it decomposes
to the anhydride, B2O3.
3.30.10 Decomposition
of oxalic acid
Oxalic acid begins to sublime at 100oC, becomes anhydrous
at 189oC and when heated rapidly decomposes into carbon dioxide,
carbon monoxide, formic acid and water.
3.30.11 Decomposition
of potassium chlorate
Potassium chlorate, KClO3, decomposes above 368oC
into potassium perchlorate and oxygen gas
3.30.11.1 Decomposition
of potassium ferricyanide
Potassium ferricyanide, K2Fe(CN)6, decomposes
before melting.
3.30.11.2 Decomposition
of phosphates
Monosodium orthophosphate, NaH2PO4.H2O,
and disodium orthophosphate [disodium hydrogen phosphate (V)] Na2HPO4.12H2O,
lose water of crystallization.
10KClO3 <--> 6KClO4 + 4KCl + 3O2
3.30.12 Sublimation,
iodine, camphor, naphthalene
See diagram 3.30.12: Sublimation of iodine
Sublimation is when a solid changes directly into a gas without melting.
Iodine, camphor and naphthalene do not decompose when heated but sublime
from the crystal to the vapour form.
Put black, shiny crystals of iodine in an evaporating dish. Cover the
dish with a piece of filter paper and stand a filter funnel upside down
on the dish. Heat the dish gently. Purple vapours rise through the filter
paper. As they cool in the filter funnel, shiny black crystals of iodine
form again.
3.30.13 Heat
silica sand, SiO2
Silica sand consists of pieces of silicon (IV) oxide (SiO2)
crystals. Heat sand in a crucible. The sand particles may break up physically,
but do not break up chemically.
3.30.14 Decomposition
of ammonium dichromate, (NH4)Cr2O7
Ammonium dichromate is an orange coloured crystalline substance
1. Show the decomposition of an ammonium dichromate by heating <
5 g on a fireproof surface in a fume cupboard or outside with observers
positioned up wind of the reaction. When outdoors, hose down area well
to ensure fine dust is not scattered. It starts decomposing with sparks
and gives out heat to form a green fluffy powder chromic oxide, nitrogen
and water.
2. An ammonium dichromate is unusual in that it undergoes an exothermic
decomposition on heating to form green chromium (III) oxide, nitrogen gas
and water. Start the reaction with a lighted match. The reaction continues
because it is exothermic. This reaction is not dangerous if only small
quantities of material are used. The reaction is dangerous if particles
of undecomposed ammonium dichromate are projected into the air and inhaled.
Ammonium dichromate is often used in the volcano experiment.
3.30.15 Decomposition
of ammonium chloride, NH4Cl
Put ammonium chloride into the bottom of a dry test-tube and warm it
over a small flame. The solid decomposes to form ammonia and hydrogen
chloride. Some of it recombines at the cool upper part of the test-tube
as a white layer. Heat the recombined layer again and it moves further
up the test-tube. This process is recombination not sublimation. NH4Cl
(s) <--> NH3 (g) + HCl (g)