School Science Lessons
2016-03-20 SP MF
Please send comments to: J.Elfick@uq.edu.au

12B Reactions of metallic elements and compounds
Table of contents

12.1.0 Aluminium, Reactions of Al

12.2.1 Antimony, Reactions of Sb

12.2.2 Arsenic, Reactions of As and compounds

12.2.3 Barium, Reactions of Ba compounds

12.2.4 Bismuth, Reactions of Bi compounds

12.3.1 Cadmium, Reactions of Cd compounds

12.4.1 Calcium, Reactions of Ca and compounds

12.5.0 Chromium, Reactions of Cr compounds

12.9.0 Lead, Reactions of Pb

12.9.3 Lithium, Reactions of Li with water

12.10.3 Magnesium with CO2, sparklers

12.10.1 Magnesium with water

12.10.2 Magnesium, Reactions of Mg compounds

12.8.1 Manganese, Reactions of Mn (II) salts

12.8.2 Manganese, Prepare manganates

12.9.4 Nickel, Reactions of Ni compounds

12.13.0 Phosphorus, Reactions of phosphorus, P

12.16.0 Silver, Reactions of silver, Ag

12.17.1 Strontium, Reactions of Sr compounds

12.20.0 Tin, Reactions of Sn

12.21.1 Zinc, Reactions of Zn and compounds
12.19.5.0 CFCs, chlorofluorocarbons, "Freons"

12.19.5.1 "Freons"

12.19.7 Bromo-compounds

12.5.0 Reactions of chromium compounds, Cr
12.5.0 Chromium ions in solution, hexaaquachromium ion, [Cr(H2O)6]3+
12.5.7 Chromic acid, ionization reaction, H2CrO4
Experiments
12.5.5 Oxidize chromium compounds to chromates,CrO42-
12.5.2 Prepare chromium trioxide, CrO3
12.5.6 Prepare potassium dichromate, K2Cr2O7
12.5.1 Reactions of chromium, Cr, and chromium compounds
12.5.3 Reactions of dichromates, Cr2O72-, potassium dichromate
12.5.4 Reactions of chromates, CrO42-

12.9.0 Reactions of lead, Pb
12.9.1 Reactions of lead (II) salts, Pb2+
12.9.2 Reactions of lead (IV) salts, Pb4+
15.8.3 Tests for lead ions in an unknown solution
3.61 Prepare lead-tin alloys in a casting mould
3.62 Tests for hardness of lead, tin, and lead-tin alloys
3.63 Tests for melting point of lead, tin, and lead-tin alloys

12.13.0 Reactions of phosphorus, P
12.13.5 Prepare microcosmic salt, Na.NH4.H.PO4.4H2O
12.13.2 Prepare phosphorus trichloride, PCl3
12.13.3 Prepare phosphorus pentachloride, PCl5
12.13.3.1 Prepare phosphorus pentoxide
12.13.6 Reactions of phosphites, HPO32-
12.13.1 Reactions of phosphorus, P, and phosphates, PO43-
12.13.4 Phosphorus trichloride with water

12.16.0 Reactions of silver, Ag
12.16.1 Reactions of silver compounds
12.16.2 Recycle silver

12.20.0 Reactions of tin, Sn
12.20.2 Prepare tin (IV) chloride
12.20.1 Reactions of tin and tin compounds

12.1.0 Reactions of aluminium
Experiments
12.1.1 Aluminium with acids
12.1.2 Aluminium with sodium hydroxide
12.1.3 Aluminium with iodine
12.1.4 Burn aluminium in oxygen
12.1.5 Thermite reaction, thermit welding
12.1.6 Aluminium with sulfur
12.1.7 Aluminium sulfate reactions
12.1.8 Aluminium chloride with water
12.1.9 Bauxite digestion
12.1.10 Alumina as a catalyst in the cracking process

3.61 Prepare lead-tin alloys in a casting mould
See diagram 3.61: Casting mould from a nut and bolt
Make a casting mould by drilling out the thread of a nut to leave a smooth hole of about 0.6 cm diameter.
Then cut the nut into two halves with a hacksaw.
Use wire to bind the two halves together for casting, then put this caste on sand.
Pure tin melts at 232oC and pure lead melts at 327oC.
Weigh out pieces of lead and tin to make four alloys so that the percentage of tin by weight is 20% tin, 40% tin, 60% tin and 80% tin.
Put each mixture of lead and tin in a crucible or Pyrex test-tube.
Cover each mixture with powdered charcoal to prevent oxidation of the metals, then heat with a Bunsen burner until they melt.
Stir the melt with a wood splint to help the metals dissolve.
Pour each mixture of molten metal into the mould until it is full.
Be careful! Hold back the carbon from the charcoal with a wooden splint while pouring.
When the cast alloy is cool, knock away the two halves of the nut.

3.62 Tests for hardness of lead, tin, and lead-tin alloys
See diagram 3.62: Hardness test apparatus
Test the hardness of the four lead-tin alloys and two pure metals, lead and tin.
Use a metal punch with a pointed end and a 1 metre plastic tube to guide the punch as it falls on to the alloy and makes a small hole.
The softer the alloy the larger the hole.
Measure the diameters of the holes with vernier callipers and a magnifying glass.
The pure metals should be less hard than the alloys.
The 60% tin alloy should be the hardest alloy.
This test is a kind of dynamic hardness test, e.g. Vickers hardness test.
Geologists test the hardness of minerals with a scratch hardness test, Mohs' test.

3.63 Tests for melting point of lead, tin, and lead-tin alloys
See diagram 3.63: Melting point apparatus: A metal plate, B suspended metal plate
Prepare a metal plate from a 12 cm X 12 cm piece of iron, 0.2 to 0.4 cm thick.
Draw a hexagon on the metal plate, then drill a small equal depth depressions at each corner of the hexagon.
Drill holes through the four corners of the metal plate.
Thread wire through the four holes and suspend the metal plate horizontally.
Pour a few globules of four alloys and the two pure metals into separate porcelain bowls.
Be careful! Put one pellet of each alloy or metal into a depression on the metal plate.
Heat the middle of the metal plate with a Bunsen burner.
Touch the pellets with a wood splint to check when they melt.
When all the pellets are all molten, use the wooden splint to remove excess molten metal from bigger pellets so that they are all the same
size.
Remove the Bunsen burner flame, leave the metal plate to cool and note the time to form crystals.
Make a table of time to crystallize and plot the results on graph paper.
Pure lead solidifies first, then 20% tin, then 40% tin, then 60% tin.
The alloy that takes the longest time to solidify has the lowest melting point.
The 60% tin alloy should have the lowest melting point.

12.1.1 Aluminium with acids
Dissolve aluminium in heated dilute hydrochloric acid and note that hydrogen gas forms.
2Al + 6H + --> 2A13+ + 3H2 (g)
Hot concentrated sulfuric acid will attack aluminium with the production of sulfur dioxide.
Dilute or concentrated nitric acid acts only very slowly on aluminium.

12.1.2 Aluminium with sodium hydroxide
1. Use a dropper to put drops of concentrated sodium hydroxide solution onto a sheet of aluminium foil or aluminium powder in a
test-tube.
Wear goggles.
Be careful! Hydrogen gas forms from the very rapid reaction

2. Add 5 g coarse aluminium powder to 20 ml of 40 % sodium hydroxide solution in a test-tube.
Quickly place the test-tube in the bottom of a tall glass beaker before the violent react occurs.
The coarse aluminium powder has a surface layer of aluminium oxide, which is first dissolved by the sodium hydroxide before the main
reaction occurs.
The aluminium completely dissolves and the water acts here too as an acid
(The aluminate ion in an anhydrous compound is shown as AlO2-, and in the hydrated form is shown as Al(OH)4-.) 2Al + 2NaOH + 2H2O --> 2NaAlO2 + 3H2
2Al + 2OH- + 2H2O --> 2AlO2- + 3H2 (g)
2Al(s) + 2NaOH (aq) + 6H2O --> 2Na+ (aq) + 2[Al(OH)4]- + 3H2 (g)
The following equations show the aluminium reacting with water to form the amphoteric aluminium hydroxide Al(OH)3, which later
goes in solution to produce aluminates, [Al(OH)4]-.
2Al + 6H2O --> 2Al(OH)3 + 3H2
Al(OH)3 + NaOH --> Na+ + [Al(OH)4]-
Al2O3 + 2NaOH + 3H2O --> 2Na+ + 2[Al(OH)4]-
Commercial drain cleaners, e.g. Drano, contain sodium hydroxide crystals and pieces of aluminium, which react as above when put in
a wet clogged drain.
The reaction is very hot as the sodium hydroxide reacts with fats to form soaps, the hydrogen applies pressure to move the pieces of
aluminium to whirl around and cut the blockage.

12.1.3 Aluminium with iodine
Fine particles of aluminium react violently with iodine, especially after a drop of water has been added.
A large amount of unreacted iodine is liberated as purple vapour into the air.
This reaction should only be done with < 5 g of materials and in a fume cupboard or outdoors.
Mix the ingredients in a small ceramic mortar and pestle.
All observers must wear eye protection.

12.1.4 Burn aluminium in oxygen
Sprinkle aluminium powder onto a Bunsen burner flame or heat aluminium powder in a crucible, then lower it into a gas jar of oxygen.
he aluminium burns brightly to form the white powder magnesium oxide.
4Al (s) + 3O2 (g) --> 2Al2O3 (s)
aluminium + oxygen --> aluminium oxide
Aluminium oxide is an amphoteric oxide that does not dissolve in water.
Stored aluminium is always coated with aluminium oxide, which protects it from most chemical reactions.
So aluminium can be used for many purposes where an unreacted metal is needed.

12.1.5 Thermite reaction, thermit welding
Be careful! The thermite reaction is a hazardous experiment.
Mix aluminium powder or aluminium turnings with iron oxide and ignite the mixture with a burning magnesium ribbon.
Do the experiment in the open with observers at least ten metres away.
Do not use > 25 g of the reaction mixture.
Prepare the reaction mixture in a cut down aluminium beverage can, suspended above a bucket or trough of sand to contain the molten
iron formed.
The mixture may be difficult to ignite, but burns with white heat, producing molten iron that can be tapped from the bottom of the
container.
Be careful! Burning magnesium ribbon held close to the eyes may cause eye damage.
The mixture may react violently if the aluminium particles are too fine.
Any trace of moisture in the reactants or container may cause violent evolution of steam and ejection of the white hot contents.
2Al + Fe2O3 --> 2Fe + Al2O3
8Al + 3Fe3O4 --> 9Fe + 4Al2O3
Thermit welding is used to weld iron rails together
2Al + 3FeO -->3Fe + Al2O3
2Al + Fe2O3 --> 2Fe + Al2O3
The ends of the rails are encased in thermit putty to hold the molten products of the reaction.
The putty consists of silica sand, bentonite, carboxymethyl cellulose and water.

12.1.6 Aluminium with sulfur
Mix dry aluminium powder with twice its volume of sulfur powder.
Put into a test-tube only enough to cover the bottom of the test-tube.
Be careful! Larger quantities may explode! Set up a safety screen.
Clamp the test-tube vertically and heat with a Bunsen burner.
Note the vigorous action where aluminium sulfide is synthesized.
Leave to cool, then add drops of water.
Hydrogen sulfide forms because of the hydrolysis of the aluminium sulfide.
2Al + 3S --> Al2S3
Al2S3+ 6 H2O --> 2Al(OH)3 (s) + 3H2S (g)

12.1.7 Aluminium sulfate reactions
1. Add ammonia solution, NH3 (aq) ("ammonium hydroxide") to aluminium sulfate solution.
Note the white precipitate of aluminium hydroxide that is insoluble in excess ammonia solution.
Al3+ + 3OH- --> Al(OH)3 (s)

2. Add drops of sodium hydroxide to aluminium sulfate solution.
Note the white precipitate that dissolves in excess sodium hydroxide to form sodium aluminate.
Aluminium hydroxide is amphoteric.
Al(OH)3 + OH- --> AlO2- + 2H2O

3. Add blue litmus solution to aluminium sulfate solution.
The blue litmus turns red.
Add sodium carbonate solution and note the production of carbon dioxide.
Aluminium salts in solution can act as acids because of hydrolysis.
Al3+ + 3H2O --> Al(OH)3 + 3H +

4. Pass hydrogen sulfide through aluminium sulfate solution to produce the hydroxide, not the sulfide.

5. Mix aluminium sulfate with twice its volume of anhydrous sodium carbonate and heat it on a charcoal block.
Note the white infusible mass.
Add cobalt nitrate solution and heat again.
A bright blue solid forms.

12.1.8 Aluminium chloride with water
Be careful! Demonstrate this reaction only to senior students.
Place < 5 g of aluminium chloride in a beaker in a fume cupboard and add water drop-by-drop.
The material will hiss, crackle and release clouds of hydrogen chloride and fine particles.
Anhydrous aluminium chloride, AlCl3, reacts violently with water to form the hydrated salt by hydrolysis, and a solution of H+ ions and
Cl- ions (a solution of hydrochloric acid) and hydrogen chloride gas.
The formation of an acid solution is more typical of a non-metal rather than a metal and this reflects the position of aluminium to the
right of magnesium in the periodic table.
Fine aerosol particles may also be generated.
Both the hydrogen chloride gas and the fine particles are extremely irritant to the lungs.
Aluminium chloride should only be used in a fume cupboard and only in small amounts.
Do not mix aluminium chloride with alkaline materials, e.g. sodium hydroxide, because a violent reaction may occur.
Aluminium chloride is exceedingly hygroscopic so keep it in a tightly-sealed plastic container.
Purchase the material only in small amounts, e.g. 100 g.
Aluminium bromide, AlBr3, has dangerous properties similar to anhydrous aluminium chloride.
AlCl3 (S) + 3H2O (l) --> Al(OH)3 (s) + 3H+ (aq) + 3Cl- (aq)

12.1.9 Bauxite digestion
Wear safety gloves and safety spectacles
Weigh 10 g of ground, dried bauxite and transfer to a reflux flask.
Add 100 mL of 20% sodium hydroxide solution and boil under reflux for 1-2 hours.
Leave to cool, then transfer the solution and residue to a 200 mL volumetric flask.
To make sure all the contents of the reflux are transferred, use demineralized water.
Leave to cool, then make up to volume with demineralized water.
When the muddy residue has settled overnight, remove a 10 mL aliquot to a 250 mL beaker and dilute to 100 mL with demineralized water.
Heat the mixture until it boils, then make acid with 1:1 HCl using methyl red indicator.
Add 2 g of ammonium chloride, then add 1:1 ammonium hydroxide until the yellow end point is reached.
Boil the mixture to coagulate the precipitate, then filter it while still hot and wash the precipitate on the filter paper with hot water.
Leave the precipitate and filter paper overnight to dry to a slightly damp consistency.
Transfer the precipitate to a weighed crucible, e.g. porcelain, silica or platinum.
Dry the precipitate on the edge of a hot plate.
Do not allow any material to be lost by spitting.
Transfer the crucible and contents to a muffle furnace and ignite to a constant weight at 1,0000 to 1,200oC.
Leave the crucible and contents to cool in a desiccator, then weigh them.
Calculate the percentage of alumina in the the bauxite: (weight of crucible and residue - weight of crucible) × 20 × 100 / weight of the
sample.
If the "mud" is left with the solution it contributes to less than 1% of the total volume, an almost insignificant error.
The alumina content can also be determined volumetrically with EDTA.

12.1.10 Alumina as a catalyst in the cracking process
See diagram 12.1.10: Alumina as a catalyst
Large quantities of alumina, aluminium (III) oxide, are used in the cracking processes in oil refineries.
One of the products, ethylene gas, C2H2, is used in the petrochemical industry to produce polyethylene and other polymers.
Activate alumina is very porous and is used as a filter for water treatment, an adsorption desiccant and a catalyst for natural gas and oil
refining processes.

Experiment
Use a fume hood for this experiment.
Hot paraffin oil and ethylene gas are flammable.
Set up the apparatus in diagram 12.1.10. Put 5 cm of paraffin oil and boiling chips in the boiling tube.
Clamp it at a shallow angle, then put 0.5 g of alumina powder half way down the boiling tube.
Replace the stopper.
First heat the alumina strongly with a Bunsen burner, then heat the alumina and paraffin oil alternately until the paraffin oil boils and its
vapour passes over the alumina.
Let the first few bubbles of gas escape under the fume hood, then collect two test-tubes of gas.
Fix stoppers on the test-tubes.
Immediately after collecting the two test-tubes of gas, turn off the Bunsen burner and remove the delivery tube and stopper from under
the boiling tube.
This action prevents water "sucking back" into the boiling tube as it cools.

12.2.1 Reactions of antimony
[Tartar emetic, K(SbO)C4H4O6. H2O, crystalline, poisonous, but used as expectorant and to treat schistosomiasis.]
1. Prepare antimony sulfide colloidal solution.
Put 20 drops of yellow ammonium sulfide into a boiling tube full of water.
Put tartar emetic in another boiling tube and fill with water.
Mix equal volumes of the two solutions to produce the colloidal solution and test it as follows: .
1.1 Add sodium chloride.
Precipitation occurs.
1.2 Add iron (III) hydroxide solution.
Coagulation occurs because the particles in the two solutions have opposite charges.
Iron (III) hydroxide sol is positively charged and antimony sulfide is negatively charged.

2. The effect of alteration of concentration, hydrolysis of antimony chloride.
Put antimony chloride in a test-tube and add 1 mL of water.
Note the white precipitate of antimony oxychloride.
Add drops of concentrated hydrochloric acid until the white precipitate disappears.
Add drops of water until the reappearance of antimony oxychloride, SbOCl.
SbCl3 + H2O <--> SbOCl (s) + 2HCl

3.1 Add 2 mL of starch solution to 2 mL of antimony sulfide solution. Add sodium chloride solution.
The sodium chloride solution has no effect where the solution is protected by the starch.
3.2 Dilute 2 mL of the antimony sulfide solution with 2 mL of water to act as a control.
Add sodium chloride solution.
The sodium chloride solution coagulates the control.

12.2.2 Reactions of arsenic and arsenic compounds
Arsenic is widespread and abundant in the earth.
It is used in dyes, pigments, medicines, lead shot alloy, glass-making, fireworks.
Arsenic and arsenic compounds are not use in school science experiments because these substances are very poisonous.
The two forms are yellow arsenic, S.G. 1.97 and grey arsenic, metallic arsenic S.G. 5.73, respectively.
It has steel grey colour and is a very brittle, crystalline, semi-metallic solid, a metalloid solid.
It tarnishes in air.
When heated it oxidizes to arsenous oxide, which has a garlic odour.
Heated arsenic (III) oxide gives off the garlic smell of arsenic and a black ring of arsenic in the test-tube.
Arsenic (III) oxide is amphoteric and is slightly soluble in water.
Occurs in realgar (As4S4), orpiment (As2S3), arsenolite (As2O3), and arsenopyrite (FeAsS).

12.2.2.1 Wood treated with copper chrome arsenate (CCA)
Copper chrome arsenate is highly toxic but the amount of arsenic in treated wood, timber, is not thought to be toxic because a person
would have to ingest about 20 cm3 of treated timber to be at risk from arsenic poisoning.
However, when CCA-treated wood is burnt it forms arsenic vapour so it should not be burnt but disposed of in a landfill.

12.2.3 Reactions of barium compounds
1. Add calcium sulfate solution to barium chloride solution.
Heat the solution and leave to cool.
Note the white precipitate of barium sulfate that is insoluble in water.
Ba2+ + SO42- --> BaSO4 (s)

2. Add ammonium carbonate solution to barium chloride solution.
Note the white precipitate of barium carbonate.
Ba2+ + CO32- --> BaCO3 (s)

3. Add ammonium oxalate solution to barium chloride solution.
Note the white precipitate of barium oxalate that is soluble in dilute hydrochloric acid but insoluble in acetic acid.
Ba + C2O42- --> BaC2O4 (s)

4. Add potassium chromate solution to barium chloride solution.
Note the yellow precipitate of barium chromate.
Ba2+ + CrO42- --> BaCrO4 (s)

5. Do the flame test on barium compounds and note the flame has flashes of green.

12.2.4 Reactions of bismuth compounds
1. Mix solid bismuth nitrate with anhydrous sodium carbonate and heat it on a charcoal block with a mouth blowpipe.
A pink globule of bismuth forms surrounded by brown bismuth oxide Bi2O3.
Bismuth oxide is used in medical suppository creams.

2. Pass hydrogen sulfide into bismuth nitrate solution acidified with dilute hydrochloric acid.
Note the dark brown precipitate of bismuth sulfide that is insoluble in either yellow ammonium sulfide or in sodium hydroxide.
Filter the precipitate, then wash it into an evaporating basin with dilute nitric acid.
Heat the evaporating basin to dissolve the precipitate.
2Bi3+ + 3S2- --> Bi2S3(s)

3. Dissolve bismuth chloride in dilute hydrochloric acid, then pour it into a boiling tube full of water.
A white precipitate of bismuth oxychloride forms.
Pour some precipitate into a test-tube and add drops of concentrated hydrochloric acid to dissolve the precipitate.
BiCl3 + H2O --> BiOCl (s) + 2HCl

12.3.1 Reactions of cadmium compounds
1. Pass hydrogen sulfide into cadmium sulfate solution.
Note the bright yellow precipitate of cadmium sulfide.
Cd2+ + S2---> CdS (s)

2. Add 3 cm of cadmium sulfate solution in a test-tube an equal volume of 5 M concentrated hydrochloric acid.
Pass hydrogen sulfide through the solution.
No precipitate appears in acid of this concentration.
Repeat the experiment and dilute the solution until the yellow precipitate appears.
Cadmium sulfide precipitates incompletely if the solution is too acidic.
Filter off some of the yellow cadmium sulfide and show that it is soluble in dilute nitric acid.
CdS + 2H + --> Cd2+ + H2S (g)

3. Add sodium hydroxide solution to cadmium sulfate solution.
Note the precipitate of cadmium hydroxide that is insoluble in excess sodium hydroxide.
Cd2+ + 2OH- --> Cd(OH)2 (s)

4. Add drops of ammonia solution, NH3 (aq) ("ammonium hydroxide") to cadmium sulfate solution.
Note the white precipitate of cadmium hydroxide that dissolves in excess "ammonium hydroxide".

12.4.1 Reactions of calcium and calcium compounds
1. Heat a flake of calcium on wire gauze with a Bunsen burner flame.
The calcium burns brilliantly with a red flame and leaves a white residue of calcium oxide.
Add drops of water to the calcium oxide in a test-tube and note the vigorous exothermic reaction.
Test the solution with red litmus paper that turns blue.
Note that calcium oxide is not very soluble in water.
2Ca + O2 --> 2CaO
CaO + H2O --> Ca(OH)2 (s)

1.1 Drop a small piece of calcium (not old stock calcium) into a test-tube a quarter full of dilute hydrochloric acid.
Press your thumb over the mouth of the test-tube and when you can feel the pressure on your thumb test the gas for hydrogen with a
lighted splint.
Repeat the experiment with a small piece of magnesium ribbon.
The same reaction occurs, but the calcium is obviously more reactive because it slower down group 2 of the periodic table.
Usually, the lower an element in the same group of the periodic table, the more reactive it is.
Ca + 2HCl --> CaCl2 + H2
Mg + 2HCl --> MgCl2 + H2

2. Add ammonium carbonate solution to calcium chloride solution.
Note the white precipitate of calcium carbonate.
Ca2+ + CO32- --> CaCO3 (s)

3. Add ammonium oxalate solution to calcium chloride solution.
Note the white precipitate of calcium oxalate that is soluble in dilute hydrochloric acid but insoluble in acetic acid.
Ca.
+ C2O42- --> CaC2O4 (s)

4. Add solutions of calcium salts to potassium chromate solution and to calcium sulfate solution.
No precipitate forms with calcium sulfate solution and barium salts with potassium chromate solution.

5. Add sodium phosphate solution to calcium chloride solution.
Note the white precipitate of calcium phosphate that is soluble in dilute hydrochloric acid, nitric acid or acetic acid.
3Ca + 2PO43- --> Ca3(PO4)2 (s)

6. Add concentrated hydrochloric acid to dry calcium chloride and do the flame test.
Note the brick-red flame and observe the green colour when seen through blue glass.

12.5.0 Chromium ions in solution
The simplest ion is the hexaaquachromium (III) ion, [Cr(H2O)6]3+, usually shown as Cr3+, a complex ion with a violet-blue colour,
but, when produced in a chemical reaction, is often green.

1. The hexaaquachromium (III) ion forms "violet-blue-grey" pH 3 solutions in water when the water molecule pulls a hydrogen ion off
the complex ion.
So the complex ion is acting as an acid because it gives an hydrogen ion to a water molecule.
[Cr(H2O)6]3+ + H2O <--> [Cr(H2O)5(OH)]2+ + H3O+, but usually shown simply as:
[Cr(H2O)6]3+ + H2O <--> [Cr(H2O)5(OH)]2+ + H+ (aq)

2. Heat chromium (III) sulfate solution
The violet-blue chromium (III) sulfate solution turns green.
[Cr(H2O)6]3+ + heat --> [Cr(H2O)5(SO4)]4+
One of the water molecules in the complex ion is replaced by a sulfate ion.
Two positive charges are replaced by two negative charges of the sulfate ion.

3. Heat chromium (III) chloride solution
The violet-blue chromium (III) chloride solution turns green.
[Cr(H2O)6]3+ + heat --> [Cr(H2O)4Cl2]+ green, tetraaquadichlorochromium (III) ion,
hexaaquachromium (III) ion --> tetraaquadichlorochromium (III) ion
Two of the water molecules in the complex ion are replaced by chloride ions.

4. Chromium ion + sodium hydroxide
The violet-blue chromium ion solution forms a gelatinous light blue precipitate, with excess sodium hydroxide redissolves to form a
green solution.
[Cr(H2O)6]3+ + 3OH- sodium hydroxide solution --> 3H2O + [Cr(H2O)3(OH)3] (s)
A hydrogen is removed from three of the water molecules in the complex ion to form a neutral complex precipitate and water.
[Cr(H2O)3(OH)3] (s) + 3OH- excess sodium hydroxide solution --> [Cr(OH)6]3- + 3H2O
The precipitate dissolves again to form a solution of green hexahydroxychromate (II) ions
[Cr(OH)6]3- + H2O2 solution + heat --> CrO42-

5. Chromium ion + sodium hydroxide + hydrogen peroxide
The green hexahydroxychromate (II) ions formed by adding excess sodium hydroxide to chromium ion solution are oxidized by heating
with hydrogen peroxide solution to form a bright yellow solution of chromate (V) ions, i.e. a change from chromium (III) to chromium (VI).

6. Chromium ion + ammonia solution
The violet-blue hexaaquachromium (III) ion solution forms a light blue precipitate.
However, with excess ammonia, most of the precipitate dissolves to form a red-blue solution.
[Cr(H2O)6]3+ + 3NH3 dilute ammonia solution, acting as a base --> [Cr(H2O)3(OH)3] + 3NH4+
A hydrogen removed from three of the water molecules in the complex ion to form a neutral complex precipitate and ammonium ion.
[Cr(H2O)6]3+ + 4NH3 excess concentrated ammonia solution, then left to stand <-->
[Cr(NH3)6]3+ + 6H2O
Ammonia replaces water as a ligand in the complex ion to form hexaamminechromium (III) ions.
This reaction is a ligand and exchange reaction.

7. Chromium ion + carbonate ion
2[Cr(H2O)6]3+ + 3CO32- (aq) --> 2[Cr(H2O)3(OH)3] + 3CO2 bubbles + 3H2O
A hydrogen is removed from three of the water molecules in the complex ion to form a neutral complex precipitate, and carbon dioxide
and water.

12.5.1 Reactions of chromium and chromium compounds
1. Dry reactions of chromium
Heat a chromium compound on a carbon block and note the green residue of chromium (III) oxide, Cr2O3.
Heat the residue in a borax bead.
Note the emerald green colour in both the oxidizing and reducing flame of the Bunsen burner.

2. Reactions of chromium in solution
Prepare 2 cm of chrome alum solution alkaline with ammonia solution, NH3 (aq) ("ammonium hydroxide") solution and boil the solution.
The green-grey precipitate of chromium hydroxide forms that is soluble in dilute acids.
Cr3+ + 3OH - --> Cr(OH)3 (s)

3. Reactions of chromium in solution
Add sodium hydroxide solution to 2 cm of chrome alum solution.
Note the precipitate of chromium hydroxide that it is soluble in excess of the reagent to give a green solution of sodium chromite.
Cr(OH)3 + OH- --> CrO2- + 2H2O

4. Add sodium carbonate solution or ammonium sulfide solution to 2 cm of chrome alum solution.
Note the precipitate of chromium hydroxide.
The carbonate and sulfide of chromium are rapidly hydrolysed in solution.

5. Heat chromium (III) sulfate solution
[Cr(H2O)6]3+ + heat --> [Cr(H2O)5(SO4)]4
One of the water molecules in the complex ion is replaced by a sulfate ion.
6. Chromate (VI) - dichromate (VI) equilibrium
CrO42- yellow solution + H+ --> Cr2O72- orange solution --> + OH- --> CrO42- yellow solution
2 CrO42- + 2H+ <--> Cr2O72- + H2O (Add hydrogen ions, the equilibrium shifts to the right.
Add hydroxide ions and the equilibrium shift to the left as hydroxide ions react with hydrogen ions.)
Add dilute sulfuric acid to form orange dichromate ion.
Add sodium hydroxide solution to form yellow chromate ion.

12.5.2 Prepare chromium trioxide, CrO3
Dissolve 25 g of potassium dichromate in 50 mL of boiling water.
Cool the solution to room temperature and very slowly add 35 mL of concentrated sulfuric acid.
Leave for two hours, then pour off the liquid from the potassium hydrogen sulfate crystals.
Heat the liquid to 85oC and add 25 mL of dilute sulfuric acid.
Evaporate the liquid on a water bath until crystals form on this surface, then set it aside to crystallize.
Filter through glass wool, preferably with suction, and evaporate the filtrate to produce more crystals.
To remove traces of sulfuric acid, wash the crystals while still in the filter with concentrated nitric acid.
Chromium trioxide is not soluble in nitric acid.
Transfer the crystals to a dry evaporating basin and heat in an air oven at 130oC.
K2Cr2O7 + 2H2SO4 --> 2KHSO4 + 2CrO3 (s) + H2O.

12.5.3 Reactions of dichromates, potassium dichromate
1. Add one drop of sodium hydroxide solution to 3 cm of potassium dichromate solution.
Note the change of colour of the solution from orange to yellow because of the formation of the chromate ion.
Cr2O72- + 2OH- --> 2CrO42- + H2O

2. Add drops of dilute sulfuric acid to 3 cm of potassium dichromate solution.
Then pass sulfur dioxide through the solution.
The change of colour to green is because of the reduction of potassium dichromate to chromium sulfate.
The sulfurous acid is oxidized to sulfuric acid.
Cr2O72- + 8H+ + 3SO32- --> 2Cr3+ + 3SO42- + 4H2O
Hydrogen sulfide and also ethanol can reduce acidified solutions of potassium dichromate.
K2Cr2O7 + 4H2SO4 + 3C2H5OH --> K2SO4 + Cr2(SO4)3 + 7H2O + 3CH3.CHO (acetaldehyde)
Cr2O72- + 8H+ + 3X --> 2Cr3+ + 3XO + 4H2O

3. Acidify potassium dichromate solution.
Add a 2 cm deep layer of ether above the solution.
Be Careful! Add a drop of hydrogen peroxide solution and note the blue colour because of perchromic acid, HCrO5.

4. Reduce dichromate (VI) ions with zinc and dilute sulfuric acid or hydrochloric acid Add dilute sulfuric acid or hydrochloric acid to
zinc and potassium dichromate (VI) solution in a test-tube or flask.
Fit cotton wool in the top of the test-tube or flask to allow hydrogen gas to escape but prevent air entering to reoxidize chromium (II)
to chromium (III).
Cr2O72- + 14H+ + 3Zn --> 2Cr3+ + 7H2O + 3Zn2+ (reduction from +6 to +3 oxidation states, potassium dichromate (VI) solution to
chromium (III) ions)
2Cr3+ + Zn --> 2Cr2+ + Zn2+ (reduction from +3 to +2 oxidation states, chromium (III) ions to chromium (II) ions).

12.5.4 Reactions of chromates
1. Add a drop of silver nitrate solution to potassium chromate solution.
Note the bricked precipitate of silver chromate.

2. Add potassium chromate solution to the following solutions: 1. lead acetate and 2. barium chloride to form the chromates of the
metals as precipitates.

3. Pass hydrogen sulfide into acidified potassium chromate solution.
The chromate is reduced to a chromium salt.
2CrO42- + 10H+ + 3H2S --> 2Cr3+ + 8H2O + 3S (s)

4. Pass sulfur dioxide through acidified potassium chromate solution. Sulfurous acid reduces the yellow chromate solution to the green
chromium salt.
2CrO42- + 10H+ + 3SO32- --> 2Cr3+ + 5H2O + 3SO42-

5. Add 3 drops of a dilute acid to yellow potassium chromate solution.
The colour of the solution changes to an orange is because of the formation of the dichromate ion.
2CrO42- + 2H+ --> Cr2O72- + H2O

12.5.5 Oxidize chromium compounds to chromates, CrO42-
Add 1 cc sodium peroxide to a dilute solution of chrome alum, then boil the solution.
The yellow colour of the solution shows the presence of sodium chromate, Na2CrO4.
Tests for the chromate ion by acidifying the solution with acetic acid and add lead acetate solution.

12.5.6 Prepare potassium dichromate
1. Dissolve 15 g of potassium chromate in 50 mL of dilute sulfuric acid and evaporate to half the volume.
Leave the solution to cool so that potassium dichromate crystals form.
Crystallize again from hot water to yield purer crystals.
2K2CrO4 + H2SO4 --> K2SO4 + K2Cr2O7 + H2O

2. Add potassium hydroxide solution to chromium (III) chloride solution to form a grey-green, then dark green precipitate, containing
[Cr(OH)6]3- ions.
[Cr(H2O)6]3+ hexaaquachromium (III) ion + (NaOH solution) --> [Cr(H2O)3(OH)3] grey-green +
(excess NaOH solution) --> [Cr(OH)6]3- dark green hexahydroxochromate (III) ions.
Add hydrogen peroxide solution, then heat the solution to turn yellow as potassium chromate (VI) forms.
[Cr(OH)6]3- + (H2O2 + heat) --> CrO42-
Add dilute sulfuric acid to the yellow solution to form orange dichromate solution
2CrO42- chromate + 2H+ <--> Cr2O72- dichromate + H2O (Add H ions equilibrium to right, add OH ions equilibrium to left)
Boil the solution until no more bubbles of oxygen form to decompose any excess hydrogen peroxide.
Add concentrated ethanoic acid to acidify the solution.
Leave to cool and orange crystals of potassium dichromate form.

12.5.7 Chromic acid, Ionization reactions:
H2CrO4 + H2O <--> H3O+ + HCrO4-, K1 = 2 × 10-1
HCrO4- + H2O <--> H3O+ + CrO42-, K2 = 3.2 × 10-7

12.8.1 Reactions of manganese (II) salts
1. Add drops of yellow ammonium sulfide solution to manganese (II) chloride solution.
Note the pink precipitate.
Mn2+ + S2- --> MnS (s)
This same precipitate occurs if you pass hydrogen sulfide into an alkaline solution of a manganese (II) salt but no precipitate occurs with
an acidic solution.

2. Drop sodium hydroxide solution into manganese (II) chloride solution.
Note the white precipitate of manganese (II) hydroxide that rapidly turns brown due to atmospheric oxidation.
Keep on adding the sodium hydroxide solution and note that the precipitate is not soluble in excess.
Mn2+ + 2OH- --> Mn(OH)2 (s)
2.1. Repeat (2.) using ammonium hydroxide with same observations.
2.2. Repeat (2.) after first adding 2 cc of solid ammonium chloride to the manganese (II) chloride solution.
No precipitate occurs.
The ammonium ion introduced depresses the ionization of the hydroxide.

3. To 1 cc of manganese (II) chloride solution, add 1 mL of sodium hydroxide solution, then 2 mL of bromine water or sodium peroxide
and heat.
The valence 2 oxide or hydroxide is oxidized to the higher valence 4 oxide.
manganese dioxide, that forms a dark brown precipitate.
The permanganate forms if the manganese (II) salt is heated with excess oxidizing agent. Boil some of the manganese (II), chloride
solution with a 2 cc of lead dioxide and 1 mL of concentrated nitric acid.
Dilute with water and filter.
The solution comes through showing the pink permanganate colour.

12.8.2 Prepare manganates
Heat on a crucible lid a piece of potassium hydroxide, crystals of potassium nitrate and some manganese dioxide until the whole mass
has fused.
Leave to cool and add some water and filter.
A deep green solution of potassium manganate forms.
The O2 comes from the KNO3.
4KOH + 2MnO2 + O2 --> 2K2MnO4 + 2H2O
The solution is unstable and is readily hydrolysed by dilute acids and even by largely diluting the solution into a permanganate.
3K2MnO4 + 2H2O --> 2KMnO4 + MnO2 + 4KOH
Dilute the green solution ten times with water and boil.
Note the pink colour of the permanganate on allowing the solution to settle.

12.9.1 Reactions of lead (II) salts, Pb2+
1. Add dilute hydrochloric to lead nitrate solution.
Note the white precipitate of lead chloride.
Wash the precipitate, add four times its volume of water and heat.
The precipitate dissolves and precipitates again cooling.
Pb2+ + 2Cl- --> PbCl2 (s)

2. Add dilute sulfuric acid to lead nitrate solution.
Note the white precipitate of lead sulfate.
Wash the precipitate, concentrated ammonium acetate solution and heat.
The lead sulfate dissolves.
Pb2+ + SO42- --> PbSO4 (s)

3. Add potassium chromate solution to 3 mL of lead nitrate solution.
Note the yellow precipitate of lead chromate.
Pb2+ + CrO42- --> PbCrO4 (s)

4. Add potassium iodide solution to 3 mL of lead nitrate solution.
Note the yellow precipitate of lead iodide that is soluble in hot water.
Pb2+ + 2I- --> PbI2

5. Add drops of sodium hydroxide solution to lead nitrate solution.
Note the white precipitate of lead hydroxide that is soluble in excess sodium hydroxide solution.
Pb2+ + 2OH- --> Pb(OH)2 (s)
2 Pb(OH)2 + 2OH- --> PbO22- + 2HO
(Note: PbO22- = plumbite ion)

6. Pass hydrogen sulfide through lead nitrate solution.
Note the black precipitate of lead sulfide.
Wash the precipitate, transfer to an evaporating dish, add dilute nitric acid and heat the solution until it boils.
Some lead sulfide dissolves forming lead nitrate solution, and some lead sulfide is oxidized to lead sulfate.
Pb2+ + S2- ---> PbS (s)

7. Add drops of dilute sodium hydroxide solution to lead acetate solution until a precipitate forms, then disappears.
Add hydrogen peroxide solution and heat the solution.
Note the brown precipitate of lead dioxide.

8. Add sodium carbonate solution to lead nitrate solution.
Note the white precipitate of basic lead carbonate, Pb(OH)2.2PbCO3.
3Pb2+ + 3CO32- + H2O --> Pb(OH)2.2PbCO3 (s) + CO2 (g)
Add sodium hydrogen carbonate solution to lead nitrate solution.
Note the white precipitate of lead carbonate.
Pb2+ + 2HCO3- --> PbCO3 (s) + CO2 (g) + H2O

12.9.2 Reactions of lead (IV) salts, Pb4+
1. Add 2 cc of red lead to 2 cm with glacial acetic acid.
Heat the mixture and the red lead dissolves.
If a brown precipitate occurs repeat the experiment using less red lead.
Cool under the tap to precipitate white crystals of lead tetraacetate, lead(IV) acetate.
Pb3O4 + 8CH3COOH --> Pb(CH3COO)4 + 2Pb(CH3COO)2 + 4H2O
Add three times the volume of water to the mixture and heat it to hydrolyse the lead tetraacetate, lead(IV) acetate.
Note the brown precipitate of lead dioxide.
Pb(CH3COO)4 + 2H2O --> PbO2 (s) + 4CH3COOH

2. Add 2 cc of lead dioxide to 2 cm of concentrated hydrochloric acid and cool under the tap.
Filter the mixture and note the golden yellow solution containing lead (IV) chloride.
Divide the solution into 3 parts.
PbO2 + 4HCl --> PbCl4 + 2H2O
Heat part A of the yellow lead (IV) chloride solution and test for chlorine.
Cool the remaining solution under the tap and leave to crystallize.
Note the white crystals of lead (II) chloride.
PbCl4 --> PbCl2 + Cl2 (g)
Add drops of part B of the yellow lead (IV) chloride solution to 880 ammonia solution, NH3 (aq) ("ammonium hydroxide").
Note the fine yellow crystals of ammonium chloroplumbate.
PbCl4 + 2NH3 + 2HCl --> (NH4)2PbCl6 (ammonium chloroplumbate)
Add drops of sodium hydroxide solution part B of the yellow lead (IV) chloride solution.
Note the red gelatinous precipitate that on heating forms lead dioxide as a brown powder.
PbCl4 + 2H2O ---> PbO2 (s) + 4HCl

3. Prepare lead dioxide and lead nitrate.
Slowly add 20 g of red lead to 50 mL of dilute nitric acid and boil for 1 minute.
Be careful! Filter the solution while hot.
Leave the filtrate to cool and form lead nitrate crystals.
Wash the residue of lead dioxide twice with hot water and dry it by gentle heating in an evaporating basin.
Pb3O4 + 4HNO3 --> PbO2 + 2Pb(NO3)2 + 2H2O

15.8.3 Tests for lead ions in an unknown solution
Prepare separate solutions of lead nitrate, iron (III) chloride and barium chloride.
Test a small portion of each solution in turn with dilute hydrochloric acid, dilute sulfuric acid and sodium hydroxide solution.
Tabulate your results.
Note that lead nitrate solution always produces a precipitate.
Also, iron (III) chloride solution gives a precipitate only when sodium hydroxide solution is added.
Barium chloride solution gives a precipitate with both sulfuric acid and sodium hydroxide solutions.

12.9.3 Reactions of lithium with water
Lithium reacts violently with water to form corrosive lithium hydroxide and hydrogen gas that, if mixed with air, may explode if ignited.
Lithium reacts vigorously with water and acids and so is usually stored under oil.
Lithium floats on paraffin oil so when returning a piece of lithium to the storage container shake the container to recoat the surface with the oil.
Handle lithium in the same way as you would handle sodium metal.
However, lithium is harder to cut than sodium so used a single piece strong scalpel, not a scalpel with a disposable blade.
Lithium is toxic if ingested and corrosive to the skin.
2Li (s) + 2H2O (l) --> 2LiOH (aq) + H2 (g)

12.9.4 Reactions of nickel compounds
1. Heat some nickel carbonate in a hard glass test-tube and note the green brown residue of nickel (II) oxide.
Heat the nickel (II) oxide in a crucible and black nickel (III) oxide, Ni2O3, forms.
Dissolve nickel (III) oxide in dilute sulfuric acid to give green nickel (II) sulfate.
2Ni2O3 + 4H2SO4 --> 4NiSO4 + 4H2O + O2

2. Add sodium hydroxide solution to nickel sulfate solution.
Note the light green precipitate of nickel (II) hydroxide, that is stable in air and is soluble in ammonium hydroxide to give a blue solution
of a complex ion.
Ni2+ + 2OH- --> Ni(OH)2 (s)

3. Heat a solution of nickel chloride made by dissolving nickel carbonate in hydrochloric acid.
Nickel chloride crystals are stable when heated.

12.10.1 Magnesium with water
1. At room temperature magnesium powder slowly forms hydrogen with water.
Clean a magnesium pencil sharpener of piece of magnesium ribbon with sandpaper to remove the magnesium oxide and add a drop of
water.
Tiny bubbles of hydrogen gas form but this may occur only after a few days
Mg (s) + 2H2O (l) --> Mg(OH)2 (aq) + H2 (g) + energy
magnesium + water --> magnesium hydroxide + hydrogen
The magnesium hydroxide formed is only slightly soluble in water to form an alkaline solution, pH 12-14.
This is a type of redox reaction where the oxidation number of the metal increases.

2. Burning or molten magnesium reacts with water in a violent exothermic reaction to produce flammable hydrogen gas.
This experiment is not allowed in a school science laboratory.
It is a very dangerous experiment that has caused injuries in schools.
Magnesium powder should never be heated and is too reactive for most school experiments.

Mg + 2H2O --> Mg(OH)2 + H2 + energy
The reaction may be so hot that the magnesium can react with nitrogen in the air.
3Mg + N2 --> Mg3N2 + energy
Do not try to use water to control burning magnesium because more explosive hydrogen gas may be formed.
Mg + H2O --> MgO + H2 + energy
Also, the carbon dioxide liberated from a soda acid fire extinguisher may release even more energy
2Mg + CO2 --> 2MgO +C + energy

12.10.3 Magnesium with carbon dioxide, sparkler experiment
Fill a gas jar with carbon dioxide.
Hold a piece of clean magnesium ribbon in a pair of tongs, ignite the magnesium with a Bunsen burner flame and plunge it into the carbon
dioxide gas.
The magnesium continues to burn.
If the magnesium is taking oxygen from the carbon dioxide for burning, than you would find carbon in the gas jar.
Look for carbon specks in the gas jar.
To make the carbon more visible, you can add drops of sulfuric acid to remove the magnesium oxide and any unburned magnesium.

12.10.2 Reactions of magnesium compounds
1. Add ammonium carbonate solution to magnesium sulfate solution. Note the white precipitate of ammonium carbonate.

2. Add ammonia solution, NH3 (aq) ("ammonium hydroxide") to magnesium sulfate solution.
Note the white precipitate of magnesium hydroxide.

3. Add of ammonium chloride to magnesium sulfate solution, then add ammonium carbonate solution or ammonia solution, NH3 (aq),
("ammonium hydroxide") solution.
Note white precipitate of basic carbonate forms because the increased concentration of ammonium ion, from the ammonium chloride,
suppresses the ionization of the ammonia solution, NH3 (aq) ("ammonium hydroxide") to leave
insufficient hydroxyl ions to attain the solubility product of magnesium hydroxide.
NH4OH <--> NH4+ + OH-

4. Add ammonium chloride and ammonia to magnesium sulfate solution. Add disodium hydrogen phosphate solution.
Note the white crystalline precipitate of magnesium ammonium phosphate.
Mg2+ + HPO42- + NH3 --> MgNH4PO4 (s)

5. Heat magnesium sulfate crystals on charcoal and let cool.
Moisten the white mass with cobalt nitrate solution, heat again, then leave to cool.
Note the pink precipitate.

6. Fit a 250 mL flask fitted with a stopper and delivery tube and connect it to a U-tube.
Connect the U-tube to a piece of combustion tube.
Mix 5 cc each of ammonium, chloride and sodium nitrite in the flask and add 30 mL of water.
Put 2 cm of magnesium ribbon loosely in the combustion tube.
Heat the flask slowly until a reaction action begins, then remove the flame, and heat the combustion tube.
The reaction produces nitrogen, which combines with magnesium to form magnesium nitride, Mg3N2.
The U-tube allows the steam to condense steam and prevent it passing into the combustion tube.
Transfer the white nitride to a test-tube, add water and boil.
Test for ammonia with litmus paper.
Mg3N2 + 6H2O --> 2NH3 + 3Mg(OH)2

12.13.1 Reactions of phosphorus and phosphates
1. Add three drops of the sodium phosphate solution to 5 cm of ammonium molybdate acidified with concentrated nitric acid.
The ammonium molybdate must be much in excess.
Heat the solution with the heat of the hand.
Note the blue precipitate of ammonium phosphomolybdate, (NH4)3PMo12O40.
The deeper the blue the greater the amount of phosphate.

2. Add drops of sodium phosphate solution to a neutral solution of silver nitrate.
Note the yellow precipitate of silver phosphate that is soluble in dilute nitric acid and also in ammonia solution, NH3 (aq) ("ammonium
hydroxide")
3Ag+ + PO43- --> Ag3PO4 (s)

3. Add drops of sodium phosphate to a solution containing magnesia mixture (magnesium sulfate, ammonia, and ammonium chloride to
prevent precipitation of magnesium hydroxide) Note the white crystalline precipitate of magnesium ammonium phosphate.
Mg2+ + NH4+ + PO43- --> Mg.NH4.PO4 (s)

4. Add drops of iron (III) chloride solution to sodium phosphate solution.
Note the buff coloured precipitate that is soluble in dilute mineral acids and also in excess of iron (III) chloride solution.
HPO42- + Fe3+ --> FePO4 (s) + H+

5. To convert an orthophosphate to a pyrophosphate, heat 3 cm of disodium hydrogen phosphate to red heat and dissolve the residual
sodium pyrophosphate.
Note the residual sodium pyrophosphate solution forms a white precipitate with silver nitrate solution and a yellow precipitate with
disodium hydrogen
phosphate solution.
2Na2HPO4 --> Na4P2O7 + H2O

6. Prepare orthophosphoric acid.
Use a fume cupboard.
Add 2 mL of concentrated nitric acid to red phosphorus in an evaporating basin. Heat the basin gently and note the vigorous production
of nitrogen dioxide.
Add more nitric acid if any phosphorus remains undissolved and heat again.
The remaining liquid is orthophosphoric acid solution.
Heat the solution to evaporate and form a thick syrup.
P4 + 20HNO3 --> 4H3PO4 + 20NO2 (g) + 4H2O

7. Prepare sodium salts of orthophosphoric acid.
Titrate a dilute solution of phosphoric acid against N sodium hydroxide solution using litmus as an indicator.
Suppose x mL of the acid neutralized 25 mL of the alkali.
Repeat the titration without litmus.
This solution contains mainly disodium hydrogen phosphate from which forms crystals after evaporation to a small volume and leaving
to cool.
Filter off the crystals, wash with cold water and dry between filter papers.
2NaOH + H3PO4 --> Na2HPO4 + 2H2O

8. To prepare sodium dihydrogen phosphate, add x mL of the same phosphoric add solution to 12.5 mL of the sodium hydroxide
solution.
To obtain trisodium phosphate, add x mL of the same phosphoric acid solution to 37.5 mL of the sodium hydroxide solution.
Proceed in both cases to obtain crystals as above.
NaOH + H3PO4 --> NaH3PO4 + H2O
3NaOH + H3PO4 --> Na3PO4 + 3H2O

12.13.2 Prepare phosphorus trichloride
See diagram 12.13.2: Prepare phosphorus trichloride.
1. Use a fume cupboard.
The apparatus must be dry.
Pass carbon dioxide to displace the air.
Remove the delivery tube and put sand, then 10 g of pieces of dry phosphorus, in the retort.
The dry sand protects the retort from cracking.
Pass dry chlorine through the delivery tube.
Spontaneous ignition occurs as the chlorine and phosphorus react to produce phosphorus trichloride.
Further chlorine produces yellow phosphorus pentachloride.
2P + 3Cl2 --> 2PCl3
PCl2 + Cl2 --> PCl5

2. To purify the phosphorus pentachloride, transfer it to a distilling flask with a two-holes stopper fitted with a thermometer and delivery
tube.
Attach the delivery tube to a sloping condenser and use another distilling flask with a calcium chloride guard tube as a receiver.
Warm the liquid in the distilling flask on a water bath and collect the product until the temperature is 76oC.

12.13.3 Prepare phosphorus pentachloride
See diagram 12.13.3: Prepare phosphorus pentachloride.
Dry chlorine by passage through wash bottles containing concentrated sulfuric acid.
Pass a stream of dry chlorine into the flask and allow phosphorus trichloride to drop slowly into the atmosphere of chlorine.
The funnel prevents blocking of the inlet tube by any solid.
Phosphorus pentachloride collects as a yellow crystalline solid on the bottom of the flask.
Transfer the phosphorus pentachloride to a storage bottle.
PCl3 + Cl2 -->- PCl5

12.13.3.1 Prepare phosphorus pentoxide
Ignite < 5 g of red phosphorus on a heat resistant mat in a fume cupboard and observe the formation of phosphorus pentoxide, P4O10.

12.13.4 Phosphorus trichloride with water
1. Add one drop of phosphorus trichloride to 1 cm of water.
Hold a rod moistened with silver nitrate near the mouth of the test-tube.
The hydrolysis is vigorous and hydrogen chloride forms.
PCl3 + 3H2O --> 3HCl (g) + H3PO3.

12.13.5 Prepare microcosmic salt.
Na.NH4.H.PO4.4H2O.
1. Put 14 g of sodium phosphate and 2.2 g of ammonium chloride in separate beakers.
Dissolve each substance in 10 mL of hot water.
Mix the solutions while hot and leave to crystallize.
Crystallize again with a minimum of water.
NaHPO4 + NH4Cl --> Na(NH4)HPO4 + NaCl

2. Heat the microcosmic salt to decompose it into ammonia, water and sodium metaphosphate.
Na(NH4)HPO4 --> NaPO3 + NH3 (g) + H2O

3. Dip a loop of red-hot platinum wire in microcosmic salt.
Heat the loop to obtain a glassy bead of sodium metaphosphate.
Dust the bead with manganese dioxide and heat.
again.
Note the amethyst colour because of the formation of manganese orthophosphate.

12.13.6 Reactions of phosphites
Phosphorous acid, H3PO3, behaves as a dibasic acid.
Add silver nitrate solution to a neutral solution of sodium phosphite, NaHPO3.
Note the white precipitate of silver phosphite that, if heated or allowed to stand, darkens, because of reduction to metallic silver.
HPO32- + 2Ag + + H2O --> 2Ag(s) + HPO42- + 2H+

12.14.3 Prepare iron (III) ammonium alum, (NH4)2SO4.Fe2(SO4)3.24H2O
Dissolve 11. 5 g of iron (II) sulfate in 30 mL of dilute sulfuric acid.
Add 5 mL of concentrated nitric acid and evaporate the solution to 15 mL.
Dissolve 2.7 g of ammonium sulfate in 10 mL of water.
Mix the two solutions and leave to crystallize.
Choose a crystal with a regular shape and allow it to grow in the solution.
Iron (III) alum crystals have an amethyst colour but break down on standing in air due to the formation of basic iron (III) sulfate.

12.15.1 Prepare silica and silicon
1. Add 2 mL of dilute hydrochloric acid to a dilute solution of water glass, then heat the solution.
Note the white precipitate of hydrated silica.
SiO32- + 2H + --> SiO2(s) + H2O
Add sodium hydroxide solution to the precipitate of hydrated silica.
Heat the mixture.
The precipitate dissolves forming sodium silicate in solution.
SiO2 + 2OH- --> SiO32- + H2O

2. Mix 3 g of dry silica and 1 g of dry magnesium powder and put in a dry test-tube clamped at an angle.
Be careful! Do this experiment behind a safety screen! Heat the test-tube slowly with a Bunsen burner.
A violent reaction occurs.
Leave the mixture to cool.
Note the brown pieces of silicon in the exploded mixture.
SiO2 + 2Mg --> 2MgO + Si

3. Put two pieces of silicon in a crucible and heat them from above with a Bunsen burner.
Silicon oxidizes to form silica
Si + O2 --> SiO2

4. Add sodium hydroxide solution to amorphous silicon in a test-tube and heat the mixture.
Hydrogen forms and sodium silicate remains in solution.
Si + H2O + 2OH- --> SiO32- + 2H2 (g)

12.16.1 Reactions of silver compounds
1. Grind solid silver nitrate with twice its volume of anhydrous sodium carbonate in a mortar.
Heat the mixture on charcoal in the reducing flame of a blowpipe.
A white bead of metallic silver forms that will not mark paper but will dissolve in dilute nitric acid.

2. Add drops of concentrated hydrochloric acid to silver nitrate solution.
(Expensive!) Note the white precipitate of silver chloride.
Shake the mixture to coagulate the silver chloride, wash with water and leave to settle.
Ag+ + Cl- --> AgCl (s)
Pour off the water and divide the solid silver chloride into three parts.
Part (i): Expose it to light and it turns violet.
Part (ii) Add ammonium hydroxide and it dissolves.
Part (iii) Heat with concentrated hydrochloric acid and it dissolves.

3. Add drops of potassium chromate solution to silver nitrate solution.
Note the brick red precipitate of silver chromate that is soluble in both dilute nitric acid and sodium hydroxide.
2g+ + CrO42 --> Ag2CrO4 (s)

4. Add sodium phosphate solution to silver nitrate solution.
Note the yellow precipitate of silver phosphate.
3Ag+ + PO43- --> Ag3PO4 (s)

5. Dilute bench ammonium hydroxide solution to five times its volume with water and slowly add to silver nitrate solution.
Note the first formed brown precipitate of silver oxide that dissolves in excess of ammonia to form a complex ion [Ag(NH3)2]+.
2gNO3 + 2NH4OH --> Ag2O (s) + 2NH4NO3 + H2O
Similarly, sodium hydroxide precipitates silver oxide but it is not soluble in excess of the reagent.

12.16.2 Recycle silver
Add solid sodium chloride to silver solutions.
Decant the clear solution above the precipitate and wash it down the sink.
Store the dried precipitate.

12.17.1 Reactions of strontium compounds
1. Add ammonium carbonate solution to strontium nitrate solution. Note the white precipitate of strontium carbonate.
Sr2+ + CO32- --> SrCO3 (s)

2. Add ammonium oxalate solution to strontium chloride solution.
Note the white precipitate of strontium oxalate that is soluble in dilute hydrochloric acid but insoluble in acetic acid.
Sr2+ + C2O42- --> SrC2O4 (s)

3. Add sodium phosphate solution to strontium chloride solution.
Note the white precipitate of strontium phosphate that is soluble in dilute hydrochloric, nitric acid or acetic acid.
3Sr + 2PO43- --> Sr3(PO4)2 (s)

4. Add calcium sulfate solution to strontium nitrate solution.
Heat the solution, then leave to cool.
Note the white precipitate of strontium sulfate that is much more insoluble than calcium sulfate.
Sr2+ + SO42- --> SrSO4 (s)

5. Do the flame test with strontium nitrate.
Note the crimson colour of the flame and observe no change in colour when viewed through blue glass.

12.20.1 Reactions of tin and tin compounds
1. Pass hydrogen sulfide through tin (II) chloride solution.
Note the precipitate that is insoluble in dilute hydrochloric acid.
Sn2+ + S2- --> SnS (s)
Filter off the precipitate and wash with distilled water.
Transfer it to an evaporating basin and add yellow ammonium sulfide solution.
The precipitate dissolves.
Oxidation by the free sulfur in the ammonium sulfide occurs, so the S in the equation come from the ammonium sulfide.
(NH4)2S + SnS + S --> (NH4)2SnS3 (ammonium thiostannate)
Add dilute acid to the ammonium thiostannate to precipitate tin (IV) sulfide, SnS2.

2. Add drops of sodium hydroxide solution to tin (II) chloride solution.
Note the white precipitate of tin (II) hydroxide that dissolves in excess sodium hydroxide to form sodium stannite.
Sn2+ + 2OH- --> Sn(OH)2 (s)
Sn(OH)2 + 2OH- --> SnO22- + 2H2O (stannite ion = SnO22-)

3. Add drops of ammonium hydroxide solution to tin (II) chloride solution.
Note the white precipitate of tin (II) hydroxide that is not soluble in excess of ammonium hydroxide.

4. To show that Tin (II) chloride is a powerful reducing agent, add tin (II) chloride solution to solutions of the following reagents.
Reduction occurs with every reagent.
Iron (III) chloride forms pale green iron (II) ions.
Potassium permanganate forms manganese (II) ions.
Potassium dichromate forms green chromic ions.

12.20.2 Prepare tin (IV) chloride
1. Use a fume cupboard
Put sand into the retort protect the glass during heating, followed by 5 cc of granulated tin.
Insert the delivery tube and connect to a chlorine apparatus.
Heat the retort while chlorine passes over.
Note the ignition of the tin, the fine white crystals in the upper part of the retort, and the yellow distillate of tin (IV) chloride.
The white crystals are SnCl4.5H2O, due to traces of moisture in the apparatus.
Sn + 2Cl2 --> SnCl4

2. Add drops of water to the tin (IV) chloride mixture and heat the mixture.
Test for hydrogen chloride.
The white precipitate is hydrated tin (IV) oxide or a tin (IV) acid.
SnCl4 + 4H2O --> SnO2.2H2O + 4HCl

3. Add ammonium hydroxide solution to the tin (IV) chloride mixture and heat the mixture.
Divide the suspension of hydrated tin (IV) oxide into two parts.
To show the amphoteric nature of hydrated tin (IV) oxide, dissolve part A in sodium hydroxide solution and dissolve part B in
hydrochloric acid.
SnO2.2H2O + 4HCl -->- SnCl4 + 4H2O
SnO22H2O + 2NaOH -->- Na2SnO3 + 3H2O
(Note: Na2SnO3 = sodium stannate)

12.21.1 Reactions of zinc and zinc compounds
1. Hold a piece of zinc foil in the Bunsen burner flame, using tongs.
Note the zinc oxide forms that is yellow when hot and white when cold.

2. Add sodium carbonate solution to zinc sulfate solution.
Observe the white precipitate of basic zinc carbonate, ZnCO3.2Zn(OH)2H2O.

3. Add sodium hydrogen carbonate to zinc sulfate solution.
Note the white precipitate of the normal carbonate, ZnCO3.

4. Add drops of sodium hydroxide solution to zinc sulfate solution.
Observe the white precipitate of zinc hydroxide that dissolves in excess of sodium hydroxide solution to form sodium zincate.
Pass hydrogen sulfide is passed through the sodium zincate solution.
Note the white precipitate of zinc sulfide.
Zinc hydroxide is amphoteric.
ZnSO4 + 2NaOH --> Zn(OH)2 + Na2SO4
Zn(OH)2 + 2NaOH --> Na2ZnO2 + 2H2O
(Note: Na2ZnO2 = sodium zincate)

5. Add drops of ammonium sulfide solution to zinc sulfate solution.
Note the white precipitate of zinc sulfide that may be discoloured.

6. Dip a rolled filter paper into a concentrated solution of zinc sulfate with added cobalt nitrate solution.
Burn the filter paper on wire gauze and note the remaining green ash, Rinmann's green.

7. Add drops of ammonium hydroxide to zinc sulfate solution.
The precipitate of zinc hydroxide dissolves in excess, due to the formation of a complex ion, [Zn(NH3)2]2+.

12.19.5.0 CFCs, chlorofluorocarbons, "Freons"
See 3.50.1: Ozone and photochemical smog
Compounds of fluorine or fluorine and chlorine with ethane or methane are called freons.
Freons were widely used for refrigerating fluids, aerosols and fire extinguishers.
However, scientists believe that chemicals like freons combine with the ozone (O3) that forms a layer of the atmosphere between the
heights of 15 to 30 km.
A depleted ozone layer allows more high energy radiation from the sun to reach the earth and damage living cells.
The Montreal Protocol of November 1992 recommended the stopping of manufacture and consumption of CFCs, including the
following:.
1. Freon 11 (CCl3F, trichlorofluoromethane)
2. Tetrachloromethane (carbon tetrachloride, CCl4, perchloromethane, dry cleaning fluid)
3. 1,1,1-trichloroethane (CH3CCl3, methyl chloroform, electrical equipment cleaner)
4. 1,1,2,2-tetrachloroethane.

Modern aerosols are labelled: "NO CFC OZONE FRIENDLY".
Modern refrigerators are labelled: "CFC DEPLETED" or, better still, "NO CFC".
Two fluorocarbons used as refrigerants (Freons) were also used as aerosol propellants.
They are non-flammable, odourless, non-toxic at low concentrations, and chemically inert:
(1.) CFC-11, CCl3F, was used for spraying hair and the body
(2.) CFC-12, CCl2F2 was used in high pressure sprays for insecticides and paints.
They were also used to replace pentane in the production of the foam plastics polyurethane and polystyrene.
CFC-13 (CCl2FCClF2) was used in the electronics and dry cleaning industries.
In the upper atmosphere, UV radiation breaks up CFCs to produce chlorine atoms, which can combine with ozone, O3, to form ClO
and an oxygen molecule, O2.
Then ClO and an oxygen atom, O, combine to produce another O2 and a free chlorine atom, Cl, again.
The initial ozone is lost, and the free chlorine atom can repeat the process.
The chlorine atom may react with methane to form hydrogen chloride, and contribute to acid rain.
Cl + O3 --> ClO + O2
ClO + O --> O2 + Cl

CFCs are persistent with long half lives.
RODP = the relative ozone depletion potential (RODP)
Table 12.19.5.0
CFC
RODP
Half life
CFC-11 (Freon 11) 1.00 75 years
CFC-12 (Freon 12) 0.86 112 years
CFC-13 .
90 years
CFC-22 0.05 20 years
CFC-113 0.80 .
CFC-114 0.60 .
1,1,1-Trichloroethane 0.15 6.5 years
Carbon tetrachloride 1.11 50 years
Halon-1211 10.00 .
Halon-1301 10.00 .

12.19.5.1 "Freons"
Freon is a registered trademark for non-toxic non-flammable gases invented to avoid danger from leaking refrigerator gases.
It is a name for compounds of ethane or methane with hydrogen atoms substituted by fluorine or chlorine, i.e. CFCs.
The manufacture of Freons is being discontinued because of their ozone-depleting properties.
The term "Freon" is not in favour nowadays.
Examples of Freons include the following:
(1.) Freon 11, CCl3F,.trichlorofluoromethane, (CFC-11)
(2.) Freon 12, CCl2F2, dichlorodifluoromethane, (CFC-12), b.p. -30oC, (most common refrigerant gas, solvent, in fire extinguishers)
(3.) Freon 21, CHCl2F, dichlorofluoromethane
(4.) Freon 114, CClF2CClF2, dichlorotetrafluoroethane
(5.) Freon 142, CH2CClF2, 1-dichloro1:1difluoroethane
HCCl3 + 2HF --> HCF2Cl + 2HCl
chloroform + hydrogen fluoride --> chlorodifluoromethane + hydrogen chloride

12.19.7 Bromo-compounds
Bromoacetic acid, CH2BrCO2H
Bromoacetanilide, CH3CONHC6H4Br, p-bromoacetanilide, 4-bromoacetanilide, para-bromoacetanilide
Bromobenzene, C6H5Br, solution <20%, Not hazardous

Bromobutane, C4H9Br
Bromobutane, butyl bromide, (CH3(CH2)3Br), C4H9Br,
(1.) 1-Bromobutane, (n-Butyl bromide), butyl bromide, CH3(CH2)3Br
(2.) 2-Bromobutane, (sec-butyl bromide), CH3CH2CHBrCH
(3.) 1-Bromo-2-methylpropane , (isobutyl bromide), (CH3)2CHCH2Br
(4.) 2-Bromo-2-methylpropane, 2-methyl-2-bromopropane, (tert-butyl bromide), trimethyl bromomethane
(Agrochemical), peroxidase horseradish, trimethyl bromododecane, (CH3)3CBr

Bromochloromethane, CH2BrCl
Bromochloromethane, CH2BrCl, methylene bromochloride, Halon 1011
Bromocresol green, acid-base indicator, C21H14Br4O5S: 5.0
Bromocresol purple, acid-base indicator, C21H16Br2O5S: 6.0

Bromeosin, Eosin, Eosin Y, Eosin B

Bromoethane, BrCH2CH2Br, 1,2-dibromoethane, BrCH2CH2Br, ethylene dibromide, Toxic by all routes, avoid inhalation
1,2-dibromoethane, Solution < 1%, Not hazardous

Bromoform, CHBr3, tribromomethane, (use < 50 mL), Harmful by all routes, eye / lung Irritant, used to separate minerals
Bromohexane, C6H13
Bromomethane, CH3Br, methyl bromide: 12.19.5.2
Bromomethylbenzene, C7H7Br, benzyl bromide, α-bromotoluene

Bromomethylpropane-1,3-diol, C5H10Br2O2, 2-bromo-2-methyl propane, tert-butylbromide, Highly flammable, irritant, environmental danger

Bromophenol, BrC6H4OH, 4-bromophenol, Toxic by all routes
Bromophenol blue, indicator, C19H10Br4O56: 7.0

Bromophenol red, C19H12Br2O5S, dibromophenolsulfonphthalein, histology dye

Bromopropane
1-bromopropane, CH3CH2CH2Br, C3H7Br, n-propyl bromide, Toxic, Solution < 24%, Not hazardous
2-bromopropane, CH3CH3Br, isopropyl bromide, allyl bromide, iso-propyl bromide

Bromopropene, 3-bromopropene, allyl bromide, CH2:CHCH2Br, 1-bromoprop-2-ene, acrid smell, Toxic
Bromthymol blue, indicator, C27H28Br2O5S: 8.0

Bromotoluene, benzyl bromide, C7H7Br
Catalytic cracking of kerosene: 10.6.4

Dibromomethane, CH2Br2, methylene bromide, methylene dibromide, Toxic, Solution < 12.5%, Not hazardous

Dibromopropane
1,2-dibromopropane, propylene dibromide, CH3CHBrCH2Br, Toxic, avoid inhalation
1,3-Dibromopropane, Br(CH2)3Br
1,4-Dibromobutane, tetramethylene dibromide, Br(CH2)4Br, Irritant

Dinitrobromobenzene, C6H3BrN2O4, 2,4-dinitrobromobenzene
Organic bromo compounds, e.g. benzyl bromide
Reactions of benzene, C6H6: 16.8.1
Reactions of manganese (II) salts: 12.8.1
Reactions of the nitrites, NO2-: 12.11.1, (See 6.)
Tests for unsaturated fats, bromine water test: 16.4.7.1
Tests for unsaturated hydrocarbons, bromine water test: 16.4.7.0
Tetrabromophenolphthalein, C20H10Br4O4, disodium salt dye for X-ray examinations