topic12F

School Science Lessons
Topic 12F Nitrogen, oxygen, hydrogen peroxide, periodic table, sulfur, sulfuric acid
2012-05-20 SPPwp
Please send comments to: J.Elfick@uq.edu.au

Table of contents
12.7.0 Copper experiments
12.11.0 Nitrogen experiments
12.12.0 Oxygen experiments, hydrogen peroxide
12.19.10 Periodic table
12.18.0 Sulfur experiments
12.18.5 Sulfuric acid

12.7.0 Copper experiments
12.7.10 Copper from brass
12.7.11 Flame test of copper wire
12.7.9 Oxidize copper foil
12.7.7 Prepare copper (I) chloride, CuCl
12.7.4 Prepare copper (II) ammonium sulfate, (NH₄)₂SO₄.CuSO₄.6H₂O
12.7.5 Prepare cuprammonium sulfate, Cu(NH₃)₄SO₄.H₂O
12.7.6 Prepare copper (I) oxide, CuO
12.7.1 Reactions of copper (II) oxide, CuO
12.7.2 Reactions of copper (II) ions, Cu²⁺
12.7.3 Reactions of copper (I) compounds, Cu⁺
12.7.8 Recycle copper

12.11.0 Nitrogen experiments
12.11.3 Ammonia gas, NH_3, and the ammonium ion, NH₄⁺12.11.3.10.2 Ammonia, Tests for ammonia, ammonium ions
5.35.1 Deficiency symptoms and fertilizing the soil (See 1. Nitrogen)
12.11.2 Nitrates, Reactions of the nitrates, NO₃^-12.11.2.1 Nitrates, Brown ring test for nitrates
12.11.2.2 Nitrates, Reduce nitrate to ammonia
12.11.1 Nitrites, Reactions of the nitrites, NO₂
13.1.24 Nitrogen gas, N₂
3.46 Prepare nitrogen gas with ammonium nitrite
3.47 Prepare nitrogen dioxide, {nitrogen (IV) oxide, NO₂}
3.44 Prepare nitric oxide (nitrogen monoxide, NO)
3.45 Prepare nitrous oxide (dinitrogen oxide, N₂O)
6.39 Plants need nitrogen, nitrogen cycle

12.12.0 Oxygen experiments, hydrogen peroxide
12.12.8 Bleaching action of hydrogen peroxide
12.12.5 Hydrogen peroxide acts as a reducing agent
12.12.4 Hydrogen peroxide reacts as an oxidizing agent
12.17.1.1 Oxides and the periodic table
13.1.29 Oxygen gas, O₂
12.12.3 Prepare hydrogen peroxide solution
12.12.1 Prepare oxides by direct oxidation, O^2-
12.12.2 Prepare oxides by indirect oxidation, O2^-
3.49.0 Prepare oxygen gas with hydrogen peroxide
12.12.6 Tests for hydrogen peroxide, ionization reaction

12.19.10 Periodic table
12.19.10 Periodic table
12.17.1.1 Oxides and the periodic table

12.18.0 Sulfur experiments
8.2.15 Heat sulfur
12.18.1 Prepare forms of sulfur
12.18.2 Prepare sulfides
12.18.3.1 Prepare sulfur monochloride
12.18.3.2 Prepare thionyl chloride
12.18.4 Properties of sulfur dioxide and sulfites
12.18.5 Sulfuric acid
12.18.6.1 Prepare sodium thiosulfate crystals, "hypo", Na₂S₂O₃.5H₂O
12.18.6.2 Reactions of sodium thiosulfate, Na₂S₂O₃.5H₂O
12.18.7 Reactions of sulfamic acid, NH₂.SO₂OH

12.18.5 Sulfuric acid
12.6.0.1 Formation of acid rain, SOx, by burning sulfur or sulfur compounds
12.6.0.4 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
12.2.4 Prepare sulfuric acid with iron (II) sulfate
12.18.5.4 Reactions of dilute sulfuric acid as an acid
12.18.5.5 Reactions of dilute sulfuric acid as a sulfate
12.3.3.1 Reactions of dilute sulfuric acid with aluminium
12.18.5.0 Reactions of sulfuric acid
12.18.5.6 Reactions of sulfuric acid with sodium chloride
12.18.5.1 Sulfuric acid dehydrates copper (II) sulfate crystals
12.18.5.2 Sulfuric acid acts as a displacer of acids from their salts, sodium chloride, sodium acetate, sodium formate
12.18.5.3 Sulfuric acid acts as a displacer of acids from their salts, potassium bromide, potassium iodide
12.6.1 Sulfuric acid acts as an oxidizing agent
12.6.2 Sulfuric acid dehydrates copper (II) sulfate crystals
12.6.3 Sulfuric acid dehydrates sucrose (cane sugar)
12.6.0.0 Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide
14.1.3 Sulfuric acid with water
12.6.0.3.1 Sulfur dioxide to sulfuric acid 1.
12.6.0.3.2 Sulfur dioxide to sulfuric acid 2.
12.11.3.6 Test substances by action with hot concentrated sulfuric acid
12.6.0.2 The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite)

12.7.1 Reactions of copper (II) oxide, CuO
1. Mix copper (II) oxide with fusion mixture and heat it on a charcoal block in the reducing flame of a blowpipe. Brown scales of copper forms.
2. Add concentrated hydrochloric acid to copper (II) oxide on a watch glass. Dip a platinum wire in the mixture for a flame test. Note the intense blue-green flame.
3. Put a borax bead on a platinum wire. Sprinkle copper (II) oxide on the bead and heat in the oxidizing flame of a Bunsen burner. The bead is blue then green when hot.
4. Prepare copper from copper (II) oxide. Mix 2 mL of copper (II) oxide crystals, 2 mL of sodium hydrogen carbonate or powdered washing soda, and 4 mL of sucrose sugar crystals. Put 2 mL of the mixture into a small metal screw cap or on a metal lid. Heat the mixture with a Bunsen burner flame. The mixture will swell up and form a copper-coloured mass. To obtain the copper from the mass let the screw cap or metal lid to cool and form a copper-coloured residue. Put the residue into a test-tube and heat it for a few minutes with dilute sulfuric acid solution. Pour away the contents of the test-tube leaving a small amount of solid at the bottom of the test-tube. This solid consists of small particles of metallic copper.
5. Prepare copper sulfate crystals from copper oxide. Half fill a test-tube with dilute sulfuric acid. Hold the test-tube with a paper holder in a flame until the liquid nearly begins to boil. Add a small amount of black powder copper (II) oxide and observe it dissolving in the acid. Heat the test-tube again and add further small amounts of copper (II) oxide until a black sediment of copper (II) oxide remains at the bottom of the test-tube, even after warming for a further two minutes. Filter the blue solution into an evaporating dish and leave it to evaporate in a cupboard. Blue crystals of copper sulfate form.

12.7.2 Reactions of copper (II) ions, Cu²⁺
1. Pass hydrogen sulfide into copper (II) sulfate solution. Note a dark brown precipitate of copper (II) sulfide.
Cu²⁺ + S^2- --> CuS (s)
Wash the precipitate and pour off excess water. Add excess of dilute nitric acid and boil in an evaporating basin. The copper (II) sulfide dissolves.
CuS + 2H⁺ --> Cu²⁺ + H₂S (g)
2. Add potassium iodide solution to copper (II) sulfate solution. A precipitate of white copper (I) iodide and iodine forms. Add sodium thiosulfate solution to dissolve the iodine and note the white precipitate of copper (I) iodide.
2Cu²⁺ + 4I^- --> CuI₂ (s) + I₂ (s)
3. Add sodium hydroxide solution to copper (II) sulfate solution. Note the blue jelly-like precipitate of copper (II) hydroxide.
Cu²⁺ + 2OH^- --> Cu(OH)₂ (s)
Pour the jelly-like precipitate into a test-tube, add ammonia solution, NH₃ (aq) ("ammonium hydroxide") and note the blue precipitate dissolving to form a deep blue solution. This solution contains the cuprammonium ion [Cu(NH₃)4]²⁺. Ammonia has a similar reaction with silver, copper (I) and copper (II) compounds. Boil the remaining solution and note the black precipitate of copper (II) oxide.
Cu(OH)₂ --> CuO (s) + H₂O
4. Add potassium ferrocyanide solution to copper (II) sulfate solution. Note the brown precipitate of copper ferrocyanide.
2Cu²⁺ + [Fe(CN)₆]^4- --> Cu₂Fe(CN)₆ (s)

12.7.3 Reactions of copper (I) compounds, Cu⁺
1. Add drops of potassium iodide solution to copper (I) chloride solution in concentrated hydrochloric acid. Note the white precipitate of copper (I) iodide.
Cu₂Cl₂ + 2KI --> Cu₂I₂ (s) + 2KCl
2. Pour some of the solution of copper (I) chloride in concentrated hydrochloric acid into water. Note the white precipitate of copper (I) chloride that is soluble in a high concentration of chloride ions but is insoluble in water.
Cu₂Cl₂ + 4Cl^- <--> 2[CuCl₃]^2-3. Add dilute hydrochloric acid to 2 cc of copper (I) oxide, then heat. Note the white precipitate of copper (I) chloride.
Cu₂O + 2HCl --> Cu₂Cl₂ + H₂O
4. Add dilute sulfuric acid to 2 cc of copper (I) oxide. Note the red precipitate of metallic copper in a blue solution.
Cu₂O + H₂SO₄ --> CuSO₄ + Cu (s) + H₂O
5. Add dilute nitric acid to 2 cc of copper (I) oxide. The copper produced in the reaction reacts with any excess dilute nitric acid to form blue-green copper nitrate solution, nitric oxide that turns brown on exposure to air, and water.
Cu₂O + 2HNO₃ --> Cu(NO₃)₂+ H₂O + Cu (s)
3Cu + 8HNO₃ --> 3Cu(NO₃)₂ + 2NO + 4H₂O

12.7.4 Prepare copper (II) ammonium sulfate (NH₄)₂SO₄.CuSO₄.6H₂O
Dissolve 5 g of copper (II) sulfate in 50 mL of boiling water. Dissolve 2.6 g of ammonium sulfate in 10 mL of water. Mix the solutions and evaporate until crystallization begins, then set aside to cool. The crystals are a double salt.

12.7.5 Prepare cuprammonium sulfate, Cu(NH₃)₄SO₄.H₂O
Dissolve 10 g of copper (II) sulfate by boiling in 50 mL of water in a 200 mL flask and leave to cool. Slowly add concentrated ammonia solution, NH₃ (aq) ("ammonium hydroxide") until any precipitate redissolves, then leave to cool. Be careful!. Add 20 mL of ethanol to form a layer on top of the blue solution. Stopper the flask loosely and leave undisturbed for a week. Filter off the crystals of cuprammonium sulfate and transfer to a container with a stopper. Cuprammonium sulfate is a complex salt in which the copper ion and ammonia form a single divalent ion [Cu(NH)]²⁺] unlike the double salt copper (II) ammonium sulfate, which behaves in solution as would the two sulfates in its molecule.

12.7.6 Prepare copper (I) oxide, CuO
1. Use a 5 cm square of shiny copper foil. Fold in the corners of the square then hammer the corners flat. Use tongs to heat the pieces of copper foil in a Bunsen burner flame. The heated foil turns black as a layer of black copper oxide forms on it. Use pliers to open the folded corners so that you have restores the 5 cm square again. The area of copper foil where the corners were folded over is still shiny because the copper had no access to the oxygen in the air.
2Cu (s) + O₂ (g) --> 2CuO (s)
copper + oxygen --> copper oxide
2. Fill a boiling tube to the depth of 2 cm. with copper (II) sulfate solution and add 2 cc of Rochelle salt, sodium potassium tartrate. When the salt has dissolved, add sodium hydroxide solution. The solution is now Fehling's solution. Add 2 cc of glucose and boil. An orange-red precipitate of copper (I) oxide forms by the reducing action of glucose on the copper (II) copper in solution.
2Cu(OH)₂ - O --> Cu₂O + 2H₂O
Reduction by glucose.
The precipitate is soluble in concentrated hydrochloric acid, but dilute sulfuric acid or nitric acid gives free copper and the copper (II) salt. Excess nitric acid. acts on the copper.
Cu₂O + 2HCl --> CuCl₂ + H₂O
Cu₂O + H₂SO₄ --> Cu + CuSO₄ + H₂O

12.7.7 Prepare copper (I) chloride, CuCl
1. Use enough copper (II) oxide to cover the bottom of a test-tube. Add five times that volume of concentrated hydrochloric acid then heat the solution. Note the green copper (II) chloride solution.
CuO + 2HCl --> CuCl₂ + H₂O
Add copper filings of equal volume to the copper (II) oxide used, and boil for 2 minutes. Filter the mixture through glass wool into a beaker of water. Note the white precipitate of copper (I) chloride.
Cu + CuCl₂ --> Cu₂Cl₂ (s)
Pour off the supernatant liquid into two parts. Use part A to show that the copper (I) chloride is soluble in ammonia solution, NH₃ (aq) ("ammonium hydroxide") because of the formation of a complex ion.
Cu₂Cl₂ + 4NH₃ --> 2[Cu(NH₃)]²⁺ + 2Cl^-
Use part B to show that the copper (I) chloride is soluble in concentrated hydrochloric acid, because of the formation of another complex ion with the chloride ion. This complex ion is unstable and decomposes on dilution with water.
Cu₂Cl₂ + 4Cl^- --> 2[CuCl₃]^2-
2. Prepare dry copper (I) chloride with a filter pump and Buchner funnel to get the white solid. Wash it with sulfurous acid then glacial acetic acid, then dry it by heating on a water bath. Keep the dry solid in a sealed container.

12.7.8 Recycle copper
1. Precipitate insoluble metal salts with sodium carbonate, sodium hydroxide or sodium sulfide. Decant the clear solution above the precipitate and wash it down the sink. Store the dried precipitate.
2. Add iron filings or steel wool or nails or any waste iron to displace copper from solution. Decant the clear solution above the precipitate and wash it down the sink. Store the dried precipitate.

12.7.9 Oxidize copper foil
This experiment can be done with a copper coin but in some countries it is illegal to use the coinage for any purpose except that of currency. Hold copper foil by the edge with pair of pliers and heat it at the tip of a hot flame. The surface of the copper turns black because the copper combines with oxygen in the air to form black copper (II) oxide, cupric oxide, CuO. After heating for a few minutes, leave the copper to cool, then scratch it with a knife point. A red layer below the black layer is copper (I) oxide, cuprous
oxide, Cu₂O. Below the red layer is unchanged copper, Cu. Put the blackened copper in a beaker and warm it with dilute sulfuric acid. The surface of the coin becomes clean and a blue solution of copper sulfate forms. The surfaces of copper coins become black with age because copper combines with oxygen and hydrogen sulfide gases in the atmosphere to form black copper oxide and black copper sulfide. Remove the black coating with dilute nitric acid. Dilute sulfuric acid dissolves the copper oxide, but not the copper sulfide.

12.7.10 Copper from brass
Heat a piece of brass in a test-tube half full of dilute nitric acid. Effervescence begins. Observe the brown-yellow gas in the test-tube. At the same time a green-blue solution of copper nitrate forms. Let the action continue for five minutes and then pour the solution into an eggcup. Put a clean penknife blade into the solution. A deposit of pure copper forms on the blade. The gas is a mixture of colourless nitrogen tetroxide, nitrogen peroxide, N₂O₄ with red-brown nitrogen dioxide, NO₂.
N₂O₄ <--> 2NO₂

12.7.11 Flame test of copper wire
Dip the end of a copper wire into an iodine solution or tincture of iodine, then into the edge of a flame. Observe a blue-green colour of the flame which is quite beautiful appearance in the dark.

12.11.1 Reactions of nitrites
Use a fume cupboard for these experiments. The oxide N₂O₃ is unstable at temperatures above -21^oC and decomposes into nitric oxide and nitrogen dioxide.
N₂O₃ --> NO + NO₂
1. Add dilute hydrochloric acid to sodium nitrite solution. Note the solution turns pale blue with effervescence and brown fumes are given off. Acids react with nitrites to produce unstable nitrous acid that decomposes into nitric oxide and oxygen. The oxygen oxidizes nitrous acid to nitric acid.
3NO₃^- + 2H⁺ --> NO₃^- + H₂O + 2NO
Finally, the nitric oxide finally reacts with the oxygen of the air to form nitrogen dioxide
2NO + O₂ --> 2NO₂
2. Mix iron (II) sulfate solution and sodium nitrite solution. Add drops of dilute sulfuric acid. The solution turns brown because of the loose compound that iron (II) sulfate makes with nitric oxide. This test distinguishes between a nitrite and a nitrate.
3. Add drops of concentrated hydrochloric acid to two mL of potassium iodide solution. Pour the mixture into sodium nitrite solution. Iodine forms as a brown colour or a black precipitate because of the oxidation of potassium iodide to iodine. Iodide ion is oxidized by electron loss and nitrous acid is reduced by electron gain.
2I^- - 2e- --> I₂
2HNO₃ + 2H⁺ + 2e^- --> 2H₂O + 2NO
2HNO₃ + 2H⁺ + 2I^- --> 2H₂O + I₂ + 2NO
Finally, the nitric oxide finally reacts with the oxygen of the air to form nitrogen dioxide
2NO + O₂ --> 2NO₂The nitrous acid acts as an oxidizing agent.
2NO^2- + 4H⁺ + 2e^- --> 2H₂O + 2NO (electron gain)
4. Pass hydrogen sulfide into sodium nitrite solution acidified with dilute hydrochloric acid. Note the rapid reaction, the precipitate of sulfur and the brown fumes. Sulfide ions from the partially ionized hydrogen sulfide lose electrons and are oxidized. Nitrite ions accept these electrons and are reduced.
S^2- - 2e^- --> S
2NO₂^- + 4H⁺ + 2e^- --> 2H₂O + 2NO
2NO₂^- + 4H⁺ + S^2- --> 2H₂O + 2NO + S
The nitrous acid acts as an oxidizing agent.
2NO^2- + 4H⁺ + 2e^- --> 2H₂O + 2NO (electron gain)
5. Add a piece of copper to sodium nitrite solution acidified with dilute sulfuric acid. Note the rapid reaction to attack the copper and the solution turns blue. The nitrous acid acts as an oxidizing agent.
2NO^2- + 4H⁺ + 2e^- --> 2H₂O + 2NO (electron gain)
6. Add sodium nitrite solution to bromine water acidified with dilute sulfuric acid. The bromine water loses its colour because of reduction to hydrobromic acid. The bromine is reduced to bromide ions by accepting electrons, the nitrite ion is oxidized and so acts as a reducing agent by supplying the electrons.
Br₂ + 2e^- --> 2Br^-
NO₂^- + H₂O --> NO^3- + 2H⁺+ 2e^-
NO₂^- + H₂O + Br₂ --> NO₃^-+ 2H⁺ + 2Br^-
The nitrous acid acts as a reducing agent and is oxidized to nitric acid. No gas is given off. If the free nitrous acid is in excess, it will decompose into nitric acid and oxygen.
H₂O + NO₂^- --> NO₃^- + 2H⁺+ 2e^- (electron loss)
If the free nitrous acid is in excess, decomposition occurs.
7. Add sodium nitrite solution to potassium permanganate solution acidified with dilute sulfuric acid. The solution loses its colour. The permanganate ion oxidizes and the nitrite ion reduces and is itself oxidized.
MnO₄^- + 8H⁺ + 5e^- --> Mn²⁺ + 4H₂O (electron gain)
H₂O + NO₂^- - 2e^- --> NO₃^- + 2H⁺ (electron loss)
5NO^3- + 2MnO₄^- + 6H⁺ --> 5NO₃^-+ 2Mn²⁺ + 3H₂O
The nitrous acid acts as a reducing agent and is oxidized to nitric acid. No gas is given off. If the free nitrous acid is in excess, it will decompose into nitric acid and oxygen.
H₂O + NO₂^- --> NO₃^- + 2H⁺+ 2e^- (electron loss)
If the free nitrous acid is in excess, decomposition occurs.
8. Add ammonium chloride solution to sodium nitrite solution. Heat the solution and observe the effervescence and the colourless odourless gas given off. The gas gives negative results for the limewater test, litmus test, and lighted splint test. The gas is nitrogen.
NH₄⁺ + NO₂^- --> 2H₂O + N₂(g)
9. Nitrosamines, produced by nitrous acid with secondary amines, can be formed in the gut when nitrites react with amino acids

12.11.2 Reactions of nitrates
The oxide N₂O₅is unstable above 0^oC and forms nitrogen dioxide.
1. Add concentrated sulfuric acid to sodium nitrate and heat gently. Be careful! Nitric acid vapours form with some decomposition causing brown fumes of nitrogen dioxide. The nitric acid condenses as oily drops on the cooler parts of the test-tube.
NaNO₃ + H₂SO₄ --> NaHSO₄+ HNO₃
4HNO₃ --> 2H₂O + 4NO₂ + O₂
2. Add three small pieces of copper to sodium nitrate solution and just cover with concentrated sulfuric acid. (Be careful!) Heat gently the mixture slowly until brown fumes of nitrogen dioxide form.
Cu + 2NaNO₃ + 3H₂SO₄ --> 2NaHSO₄ + CuSO₄ + 2H₂O + 2NO_{2 (g)}

12.11.2.1 Brown ring test for nitrates
Shake crystals of iron (II) sulfate with 2 cm of water. Add sodium nitrate and shake the mixture again until all dissolves. Pour concentrated sulfuric acid carefully down the side of the test-tube to form a 1 cm deep layer under colloidal iron (II) sulfate solution. Note the brown ring at the junction of the two liquids.

12.11.2.2 Reduce nitrate to ammonia
1. Add 3 small crystals of sodium nitrate to 3 cm of dilute sodium hydroxide solution. After the sodium nitrate dissolves add aluminium or zinc powder. Heat the solution and test for ammonia.
2. Repeat the experiment with Devarda's alloy replacing the aluminium or zinc powder.
(Devarda's alloy = 45% Al, 50% Cu, 5% Fe).

12.11.3 Ammonia and the ammonium ion
Hot ammonia gas is a reducing agent, e.g. hot black copper oxide is reduced to copper by a stream of hot ammonia. Ammonia forms complex ions with many metallic ions, e.g. cuprammonium sulfate. Ammonia dissolves in with water to produce ammonia solution, NH₃ (aq), ("ammonium hydroxide"), which is mainly undissociated, but contains enough hydroxyl ions to behave as a weak alkali.
The ammonium ion, NH₄⁺, is stable and has a metallic character. When heated with a high concentration of hydroxyl ions, the ammonium ion forms ammonia gas.
NH₄⁺ + OH^- --> NH₃(g) + H₂O
By hydrolysis, ammonium salts may react acidic.
NH₄⁺ + H₂O --> NH₄OH + H⁺
Ammonium salts sublime when heated.

12.12.1 Prepare oxides by direct oxidation
Put the following elements into separate test-tubes containing oxygen; calcium, carbon, iron wire, magnesium, oxygen, phosphorus, sulfur. Heat each test-tube just enough to start a chemical reaction. Heat the iron wire until it is red-hot. Test each solution with litmus paper. Note that metal oxides act as bases and non-metal oxides act as acids.

12.12.2 Prepare oxides by indirect oxidation
Add 2 cm of concentrated nitric acid to a small piece of copper foil. Add more acid if the reaction stops before all the copper goes into solution. Transfer the solution to an evaporating dish. Heat the dish slowly to evaporate the nitrate solution until crystals are form then heat strongly to complete the decomposition of the nitrate to oxide. Leave the solution to cool. When cool, show that the oxide is a base by dissolving it in dilute sulfuric acid.
2. Repeat the experiment, using lead instead of copper foil. and dissolve the oxide in acetic acid instead of dilute sulfuric acid.

12.12.3 Prepare hydrogen peroxide solution
Dilute 5 mL of a syrup of phosphoric acid with its own volume of water. (Acid to water!) Slowly add barium peroxide while cooling the test-tube under the tap. Filter the solution to remove the barium phosphate. The hydrogen peroxide solution may contain excess phosphoric acid and barium ions.
2H₃PO₄ + 3BaO₂ --> Ba₃(PO₄)₂ (s) + 3H₂O₂

12.12.4 Hydrogen peroxide reacts as an oxidizing agent
H₂O₂ + 2H⁺ + 2e^- --> 2H₂O (electron gain)
1. Add hydrogen sulfide solution to 5 mL of lead acetate solution. Note the precipitate of black lead sulfide. Pour off the liquid. Add 5 mL of hydrogen peroxide and shake the solution. Add more hydrogen peroxide if needed to oxidize the lead sulfide to white lead sulfate. Hydrogen peroxide acts as an oxidizing agent and accepts electrons. Lead sulfide is oxidized by loss of electrons.
4H₂O₂ + 8H⁺ + 8e^- --> 8H₂O
PbS + 4H₂O - 8e^- --> PbSO₄ + 8H⁺
PbS + 4H₂O₂ --> PbSO₄ + 4H₂O
2. Add drops of hydrogen peroxide solution to 5 mL of potassium iodide solution acidified by dilute sulfuric acid. Test the brown colour of iodine with a drop of starch solution. Hydrogen peroxide as oxidizing agent accepts electrons. Iodide ions is oxidized by loss of electrons.
H₂O₂ + 2H⁺ + 2e^- --> 2H₂O
2I^- - 2e^- --> I₂ (s)
H₂O₂ + 2H⁺ + 2I^- --> 2H₂O + I₂ (s)

12.12.5 Hydrogen peroxide acts as a reducing agent.
H₂O₂ - 2e^- --> 2H⁺ + O₂ (electron loss)
1. Add drops of hydrogen peroxide solution to 5 mL of potassium permanganate solution acidified by dilute sulfuric acid. Note the gas given off and the potassium permanganate loses its colour. Keep the gas given off in the test-tube and test for oxygen with a lighted splint. In this reaction, hydrogen peroxide acting as a reducing agent loses electrons and the permanganate ion is reduced by gain of electrons.
5H₂O₂ - 10e^- --> 10H⁺ + 5O₂ (g)
2MnO₄^- + 16H⁺ + 10e^---> 2Mn²⁺ + 8H₂O
2MnO₄^- + 6H⁺ + 5H₂O₂ --> 2Mn²⁺ + 8H₂O + 5O₂ (g)
2. Add a slight excess of sodium hydroxide solution to 5 mL of silver nitrate solution. Note the brown precipitate of silver oxide. Pour off the supernatant liquid. Add drops of hydrogen peroxide to the silver oxide. Oxygen is given off as the silver oxide is reduced to black metallic silver. Hydrogen peroxide acts as a reducing agent and loses electrons. Silver oxide is reduced and gains electrons.
H₂O₂ - 2e^- --> 2H⁺ + O₂ (g)
Ag₂O + 2H⁺ + 2e^- --> 2Ag (s) + H₂O
H₂O₂ + Ag₂O --> 2Ag (s) + H₂O + O₂
Add more drops of hydrogen peroxide and show that the finely divided silver acts as a catalyst.
3. Add 1 cc of lead dioxide to 5 mL of dilute nitric acid, Add drops of hydrogen peroxide. Note the oxygen given off and how lead dioxide slowly dissolves. Test the solution for lead ions by adding a drop of potassium chromate solution. In this reaction hydrogen peroxide acts as a reducing agent and lead dioxide is reduced by gain of electrons.
PbO₂ + 2H⁺ + 2e^- --> PbO + H₂O
Lead monoxide and nitric acid then form lead nitrate solution.

12.12.6 Tests for hydrogen peroxide, ionization reaction
Acidify 2 cm of potassium dichromate solution with dilute sulfuric acid. Cover the solution with 2 cm of ether. Add a drop of diluted hydrogen peroxide. The blue colour in the ether layer may come from perchromic acid (HCrO₅). This reaction is a test for hydrogen peroxide.
Ionization reaction, Ka = 2.4 × 10^-12
H₂O₂ + H₂O <--> H₃O⁺ + HO₂^-

12.12.8 Bleaching action of hydrogen peroxide
Put a piece of red or blue litmus paper and a strand of dark-coloured hair into a test-tube and shake drops of hydrogen peroxide in the test-tube. Both the paper and hair are bleached.

Copper (I) oxide, Cu₂O, copper oxide, cuprous oxide, brown copper oxide, red copper oxide, red-brown powder, deliquescent
Copper (I) oxide, cuprite, copper ore, "copper oxide", Harmful if ingested
Alkalis with basic oxides, copper oxide: 12.7.5
Reduce copper (I) oxide, (copper oxide) to copper: 10.10.2

Copper (II) carbonate, CuCO₃, cupric carbonate, blue-green powder, Harmful if ingested
Copper (II) carbonate basic, CuCO₃.Cu(OH)₂.H₂O, basic copper carbonate, (shows decomposition of carbonates, at 200^oC)
Copper carbonate, azurite: 35.20.4
Prepare rayon, basic copper carbonate with ammonia solution: 3.4.8.1

Copper (II) chloride, CuCl₂, cupric chloride, brown powder, Harmful, Environment danger
Copper (II) chloride, anhydrous copper chloride, cupric chloride, brown-yellow powder, Harmful if ingested
Copper (II) chloride, Solution < 25%, Not hazardous
Copper (II) chloride dihydrate, copper (II) chloride 2H₂O, copper chloride, cupric chloride, Harmful if ingested
Copper (II) chloride, strong electrolytes: 15.02
Alkalis with salts, sodium hydroxide with copper salts: 12.7.4

Copper (II) nitrate, Cu(NO₃)₂.3H₂O, copper (II) nitrate 3H₂O, copper (II) nitrate trihydrate, cupric nitrate, copper nitrate crystals, Harmful
Copper (II) nitrate hydrate, copper nitrate hydrated, cupric nitrate Std, blue crystal, deliquescent, Harmful if ingested
Copper (II) nitrate, For 0.1 M soln., 29.6 g in 1 L water

Copper (II) oxide, CuO, black copper oxide, cupric oxide, melconite, tenorite, black solid, soluble in dilute acids, (used in craft, deep blue colour in glass).
CuO, basic oxides, copper (II) oxide powder: 12.17.2
Copper (II) oxide, "copper oxide", cupric oxide, Harmful if ingested
Dilute acids with basic oxides, copper (II) oxide: 12.3.5
Heat copper carbonate, CuCO₃.Cu(OH)₂.H₂O, basic copper carbonate: 8.2.05
Heat food with copper (II) oxide: 16.9.2
Heat zinc with copper (II) oxide: 12.17.2.1
Reduce copper (II) oxide to copper with ammonia: 13.6.7
Reduce copper oxide with natural gas, methane: 3.41.4

Copper (II) sulfate, CuSO₄.5H₂O, copper (II) sulfate 5H₂O, copper sulfate, cupric sulfate, copper (II) sulfate hydrated, blue vitriol
1. Copper (II) sulfate solution, 0.5 M, Different authorities state that copper (II) sulfate is Harmful / Slightly toxic / Poisonous if ingested (swallowed), by skin contact, in contact with wounded skin. Use eye and skin protection (safety glasses and gloves) where splashes may occur. Do not breathe in the powder. If swallowed or skin contact occurs immediately flush the eye or skin or wash out the mouth with plenty of water.
2. Some school systems do not allow primary children access to copper (II) sulfate solution. However, hey may be shown the beautiful blue crystals.
3. Copper (II) sulfate is also harmful to organisms in the environment but small amounts of waste solutions may be disposed of down the sink with plenty of water.
Copper (II) sulfate pentahydrate, cupric sulfate pentahydrate, bluestone, copper sulfate
Copper (II) sulfate anhydrous, copper sulfate, cupric sulfate anhydrous, Harmful if ingested
Copper (II) sulfate anhydrous,
Copper (II) sulfate, CuSO₄.5H₂O, For 0.1 M soln., 25 g in 1 L water + 5 mL conc. H₂SO₄
Copper (II) sulfate is insoluble in methylated spirits
Copper (II) sulfate solution with magnesium: 14.1.2
Distil copper (II) sulfate solution: 10.5.2
Electrolysis of copper (II) sulfate solution, Faraday's laws: 15.5.15
Electrolysis of copper (II) sulfate solution with copper and platinum electrodes: 15.5.17
Electrolysis of copper (II) sulfate solution with copper electrodes: 15.5.18
Electrolysis of copper (II) sulfate solution, electrochemical equivalent of copper: 15.5.19
Electrolysis of copper (II) sulfate solution: 15.5.14
Electrolysis of copper (II) sulfate solution, microscale electrolysis: 15.5.16
Exothermic reactions give out heat energy: 3.80 (See 1, 2.)
Heat copper (II) sulfate-5-water crystals: 3.2.1
Heat of solution of anhydrous copper (II) sulfate Heat of solution, anhydrous: 14.1.1
Iron and zinc with copper (II) sulfate solution: 12.14.2.3
Magnesium, or zinc, with copper (II) sulfate solution: 12.14.2.1
Metals in copper (II) sulfate solution: 12.14.2
Prepare copper (II) carbonate and copper (II) oxide 12.15.4
Prepare preserving agent for cut flowers: 19.6.5
Prepare rayon, copper (II) sulfate with ammonia solution: 3.4.8.0
Prepare ink, yellow invisible ink: 12.15.1
Tests for oxidizing agents by change of colour of iron with copper (II) sulfate: 15.2.9
Tests for oxidizing agents by change of colour of zinc with copper (II) sulfate: 15.2.10
Use golden syrup to make red copper (I) oxide: 12.15.5
Zinc with copper (II) sulfate solution, heat from a displacement reaction: 14.1.6
Zinc with lead nitrate solution and iron with copper (II) sulfate solution: 12.14.2.4

12.15.1 Prepare ink, yellow invisible ink
Add a 1 mL of powdered copper sulfate and a 1 mL of ammonium chloride to a test-tube nearly full of water. Dissolve the substances by holding the thumb over the end of the test-tube and shake the test-tube upside. Use a small paint brush to write on a piece of paper with the faint blue solution When the writing has dried, it is invisible. Warm the paper in front of a flame to make the writing appear in yellow letters.

12.15.2 Copper (II) sulfate is insoluble in methylated spirits
Make a strong solution of copper sulfate by heating 1 cm of the powder with 2 cm of water in a test-tube. When the powder has dissolved, cool the test-tube under the tap. Pour drops of methylated spirits into the test-tube. A shower of very small crystals of copper sulfate precipitates in the test-tube because copper sulfate is insoluble in methylated spirit.

12.15.3 Prepare copper from copper (II) sulfate
Shake 2 mL of iron filings with 4 cm of copper sulfate solution in a test-tube for a few minutes. The colour of the solution will fade. If the shaking continued for long enough the colour disappears. Filter the contents of the test-tube. A red-brown powder is left in the filter paper. The chemical action is as follows:
Copper sulfate + iron –> copper + iron (II) sulfate.
The red-brown powder is not pure copper, because as excess
iron filings are used in the experiment it is a mixture of copper and excess iron filings coated with copper. Repeat the experiment with zinc or lead instead of iron filing. After cleaning the lead, rub it with sandpaper and then immerse it in hot copper sulfate solution. The heat is necessary because the action is slow in the cold. A film of copper deposits on the lead.

12.15.4 Prepare copper (II) carbonate and copper (II) oxide
green copper (II) carbonate, copper carbonate, cupric carbonate and black copper (II) oxide, cupric oxide
Add an equal amount of sodium carbonate, washing soda, solution to 2 cm of copper sulfate solution in a test-tube. A blue-green precipitate of copper carbonate forms. The double decomposition reaction is as follows:
copper sulfate + sodium carbonate –> copper carbonate + sodium sulfate.
Filter the precipitate and transfer it to a metal lid. Hold the lid in a pair of pliers and heat it gently over a small flame. The green colour of copper (II) carbonate will change to black because of the formation of black copper (II) oxide, cupric oxide.

12.15.5 Use golden syrup to make red copper oxide
Add sodium hydroxide (caustic soda) solution to 2 cm of copper sulfate solution. A blue jelly-like precipitate of copper hydroxide forms. Stir golden syrup on the end of a spoon into hot water until the syrup has dissolved. Add 2 cm of this solution to the copper hydroxide in the test-tube. Heat the test-tube gently. Soon a yellow precipitate forms in the test-tube. This is copper (I) oxide, cuprous oxide, CuO, appearing in a yellow form. Gradually the colour changes to orange. Filter off the precipitate. When the copper (I) oxide on the filter paper, it will turn red, its usual colour. Repeat the experiment with black treacle, or the sweet sold as “barley sugar.” These materials contain glucose, a form of sugar necessary for the experiment. Its chemical action is used as a test for glucose. Repeat the experiment with an apple. Heat small pieces of the apple with water, and, after filtering the liquid, use it in the same way as described for the golden syrup solution.

12.6.0.0 Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide
See diagram: 12.6.0
Sulfuric acid is a colourless oily liquid available as:
1. 2.0 M (4.0 N) 1.0 M (2.0 N) and 0.5 M (1.0 N) volumetric solutions
2. Minimum assay 97% solution density 1.83 g cm^-3 3. 98% "ANALAR" solution
3. "Battery acid" solution for lead cell accumulators minimum assay 30% density 1.25 g cm^-3 at 20^oC (battery acid).
Sulfuric acid is a strong dibasic acid that forms sulfates and hydrogen sulfates a strong oxidizing agent that dissolves copper and a strong dehydrates agent that can remove water from organic compounds. Sulfuric acid is made by the contact process. Sulfur is burned or the ores zinc sulfide or iron sulfide (pyrites) are heated to form sulfur dioxide. The gases pass over vanadium (V) oxide or platinum catalyst at 450^oC to form sulfur trioxide that combines with water to form sulfuric acid.
Sulfur trioxide is produced by the action of oxygen on sulfur dioxide in the presence of a catalyst, e.g. iron oxide.

12.6.0.1 Formation of acid rain, SOx, by burning sulfur or sulfur compounds
When coal is burnt, the compounds that contain sulfur can form sulfuric acid, as in the equations below, to become components of acid rain (rainwater pH = 5.6, acid rain pH < 5). There may be more than one pathway for the formation of sulfuric acid from sulfur dioxide.

12.6.0.2 The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite).
4 FeS₂ (s) + 11O₂ (g) --> 2Fe₂O₃ (s) + 8SO₂ (g)
S (s) + O₂ (g) --> SO₂ (g) sulfur dioxide
Also, other sulfide ores may produce sulfur dioxide in the atmosphere during smelting to obtain the pure metal.
Lead (II) sulfide (galena)
3PbS + 3O₂ (g) --> 3Pb (s) + 3SO₂ (g)
Copper (II) sulfide, chalcocite (bornite)
Cu₂S (s) + O₂ (g) --> 2Cu (s) + SO₂ (g)

12.6.0.3.1 Sulfur dioxide to sulfuric acid 1.
Sulfur dioxide is oxidized to sulfur trioxide by oxygen gas.
2SO₂ (g) + O₂ (g) <--> 2SO₃ (g)
Sulfur dioxide is oxidized to sulfur trioxide by nitrogen dioxide.
SO₂ (g) + NO₂ (g) --> SO₃ (g) + NO (g)
Sulfur trioxide dissolves in water to form sulfuric acid.
SO₃ (g) + H₂O (l) --> H₂SO₄ (aq) sulfuric acid

12.6.0.3.2 Sulfur dioxide to sulfuric acid 2.
Sulfur dioxide dissolves in water to form sulfurous acid
SO₃ (g) + H₂O (l) --> H₂SO₃ (aq)
Sulfurous acid is oxidized to sulfuric acid by ozone
H₂SO₃ (aq) + O₃ (g) --> H₂SO₄ (aq) + O₂ (g)

12.6.0.4 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
SO₂ (g) + H₂O₂ (l) --> H₂SO₄ (aq)

12.6.1 Sulfuric acid acts as an oxidizing agent
See 12.3.13: Concentrated sulfuric acid with copper
BE CAREFUL! YOU ARE USING HOT CONCENTRATED SULFURIC ACID!
Add hot concentrated sulfuric acid to carbon. The reaction forms carbon dioxide and sulfur dioxide.
C (s) + 2H₂SO₄ (l) --> CO₂ (g) + 2SO₂ (g) + 2H₂O (l)
Add hot concentrated sulfuric acid to sulfur. The reaction forms sulfur dioxide and water.
S (s) + 2H₂SO₄ (l) --> 3SO₂ (g) + 2H₂O (l)
Add hot concentrated sulfuric acid to carbohydrates. The reaction forms carbon dioxide or carbon and water.

12.6.2 Sulfuric acid dehydrates copper (II) sulfate crystals
Add drops of sulfuric acid to blue copper (II) sulfate crystals. The crystals turn white as they lose water. Concentrated sulfuric acid combines so readily with water that it can be used as a dehydrates agent, e.g. removing water from hydrated copper (II) sulfate crystals and from other hydrated salts.
CuSO₄.5H₂O (s) <--> CuSO₄ (s) + 5H₂O (l)

12.6.3 Sulfuric acid dehydrates sucrose (cane sugar)
This favourite experiment forms a large quantity of toxic gases, e.g. carbon monoxide and sulfur dioxide, and the voluminous char (remaining carbon and other substances) should be washed thoroughly with water to remove any remaining acid before handling.
1. Heat a mixture of 0.5 cm of sucrose and 1.0 cm of concentrated sulfuric acid gently for 2 seconds and then leave to stand. Note the vigorous reaction and the colour change from white sugar to black carbon.
C₁₂H₂₂O₁₁ (s) (H₂SO₄ catalyst) --> 12C (s) + 11H₂O (l)
This experiment causes great amusement to children if the sugar is in the shape of a volcano in a deep beaker and the sulfuric acid is poured into the "crater " of the "volcano".
2. Add 35 mL of sulfuric acid to 50 g of sugar in a 100 mL beaker placed on a heat resistant mat.
3. Add 10 mL of sulfuric acid to a large test-tube almost filled with sugar. The sugar is dehydrated to form carbon and steam which causes the material to expand, forming a black porous column which rises out of the beaker or test-tube. Do the experiment in a fume cupboard because foul-smelling vapours are also released. To remove remaining acid, cool and wash the product thoroughly with water before touching it.
4. Put some sucrose (cane sugar) in a tall beaker. Add drops of concentrated acid to the sugar.
BE CAREFUL!
The sugar turns yellow then brown then black and rises in the beaker. It reacts with carbohydrates like sugar and cellulose charring them by removing the elements of water from them and leaving a mass of black carbon behind.
5. Roll paper into a tube and hold it in the middle of a soft plastic container, e.g. ice cream tub. Do not use a glass jar. Fill the container with sugar. Pour just enough water to dampen the sugar down the tube to reach the bottom. Leave to stand for five minutes to allow the water to spread throughout the sugar. Remove the paper tube to leave a hole in the damp sugar.
BE CAREFUL!
Pour 30 mL of concentrated (98%) sulfuric acid down the hole and onto the top of the sugar. The sugar starts to turn brown, and black in patches. After some minutes bubbles of steam form. The reaction became more vigorous as the material in the container expands. A black cylinder rises out of the jar. Jets of steam spurt out. Heat is given out as the cylinder keeps rising. The black steaming cylinder is spongy carbon. Tap with a spatula to show it is hard, like expanded polystyrene packaging. If the carbon solidifies to make a seal over the top of the jar and the reaction continues deeper in the container, below the seal, pressure may build up to cause an explosion and a shower of black crumbling carbon.

12.18.1 Prepare forms of sulfur
See diagram 12.18.1: Sulfur crystals
Almost fill a dry test-tube with sulfur powder and heat slowly to boiling, using a safety holder. Pour the boiling sulfur into a beaker of water. Immerse any floating sulfur with a stirring rod. Remove and examine the plastic sulfur. Note the gradual loss of elasticity as the plastic sulfur changes to rhombic sulfur.

12.18.2 Prepare sulfides
1. Use an ignition tube with 3 cm of powdered sulfur and heat until melted. Hook a strip of copper over the rim of the ignition tube so that its lower edge is just above the surface of the sulfur. Heat to boil the sulfur and note the glow as copper sulfide forms on the copper.
2. Mix equal parts. of iron filings and sulfur. Heat the mixture until a reaction starts. Note the glow of the mixture as iron (II) sulfide forms.
3. Pass hydrogen sulfide into copper (II) sulfate solution. Filter off the precipitated copper (II) sulfide.Cu²⁺ + S^2- --> CuS (s)4. Prepare sulfides of iron, cobalt and nickel. Prepare solutions of iron (II) cobalt and nickel salts. Pass hydrogen sulfide into each solution Note a slight precipitate of dark iron (II) sulfide but no precipitate with the cobalt or nickel salts. Add ammonia solution, NH₃ (aq) ("ammonium hydroxide") and pass more hydrogen sulfide through the three solutions. All three solutions form a black precipitate of the metallic sulfide, however only iron (II) sulfide dissolves in dilute hydrochloric acid. The sulfides of cobalt and
nickel dissolve in concentrated hydrochloric acid in the presence of potassium chlorate or in aqua regia. Transfer the sulfides to evaporating basins, add concentrated hydrochloric acid and a crystal of potassium chlorate. Heat until the crystals dissolve. The cobalt salt becomes is pink in solution. The nickel salt becomes yellow-green.

12.18.3.1 Prepare sulfur monochloride, S₂Cl₂
This experiment may not be allowed in some school systems.
Put 10 cc of sulfur in a distilling flask. Use a one-hole stopper fitted with a delivery tube to reach the level of the sulfur. Connect the delivery tube to a supply of dry chlorine. Heat the sulfur on a gauze and pass in chlorine. Stand the flask in a water bath of cold water. Collect the liquid product of sulfur monochloride, S₂Cl₂, in a dry test-tube. Heat drops of the product in water and tests for sulfur dioxide and hydrochloric acid. Note the deposit of sulfur.
2S₂Cl₂ + 2H₂O --> 4HCl + 3S (s) + SO₂ (g)

12.18.3.2 Prepare thionyl chloride, SOCl₂
This experiment may not be allowed in some school systems.
See diagram 12.18.3.2: Prepare thionyl chloride
1. Use 10 cc of phosphorus pentachloride in a dry distilling flask attached to a sloping condenser. Fit the flask with a stopper and delivery tube that reaches deep into the flask. Stand the flask in a cold water bath. Pass sulfur dioxide into the phosphorus pentachloride until it has completely liquefied. Heat the water bath and collect the distillate of thionyl chloride, SOCl₂. The liquid remaining in the flask is phosphorus oxychloride.
PCl₅ + SO₂ --> POC1₃+ SOCl₂
Add drops of the thionyl chloride to water and tests for sulfurous acid and hydrochloric acid.
SOCl₂ + 2H₂O --> H₂SO₃ + 2HCl
2. Cut a cube of camphor into pieces and place in the dry distilling flask. Pass sulfur dioxide through the flask until the camphor liquefies. Disconnect the source of sulfur dioxide and pass dry chlorine through the flask until it is no longer absorbed. Heat the. water bath and collect the distillate of sulfuryl chloride,
SO₂C1_2.
SO₂ + C1₂ --> SO₂C1₂
Add drops of the distillate to water and show that the resulting solution contains sulfuric acid and hydrochloric acids.
SO₂Cl₂ + 2H₂O --> H₂SO₄+ 2HCl

12.18.4 Properties of sulfur dioxide and sulfites
See 3.51.1: Tests for sulfur dioxide | See 12.11.5.16: Tests for sulfates
1. Heat sulfur in an evaporating basin and test the sulfur dioxide formed by drops reagents on a glass rod. Alternatively, pass sulfur dioxide into 6 cm of water. Show that the solution is acid, potassium permanganate loses its colour, reduces potassium dichromate, and forms a deposit of sulfur with hydrogen sulfide.
2. Prepare sodium sulfite and sodium bisulfite in solution. Saturate 10 mL of sodium hydroxide solution with sulfur dioxide to form sodium bisulfite, NaHSO₃. Add 10 mL of sodium hydroxide solution to form sodium sulfite, Na₂SO₃.
NaOH + SO₂ (g) --> NaHSO₃
NaHSO₃ + NaOH --> Na₂SO₃ + H₂O3.
Hold 2 cm of magnesium ribbon in a pair tongs and heat until it ignites, then hold the burning magnesium in sulfur dioxide. Sulfur dioxide decomposes into sulfur and oxygen. Magnesium oxide forms.
2Mg + SO₂ (g) --> 2MgO (s) + S (s)
SO₂ (g) --> S (s) + O₂ (g)
2Mg (s) + O₂ (g) --> 2MgO (s)
4. Pass air (oxygen gas) through a hot solution of sodium sulfite, Na₂SO₃. Test the solution for sulfate.
2SO₃^2- + O₂ (g) --> 2SO₄^2-
5. Dissolve 5 cc of sodium sulfite crystals in 50 mL of water. Add 2 cc of crushed roll sulfur and boil for an hour. Transfer the mixture to an evaporating basin and heat to a small volume. Test the concentrated solution for sodium thiosulfate, Na₂S₂O_3:1. By addition of iodine solution,
2. by addition of an acid.
Na₂SO₃ + (O) --> Na₂SO₄
Na₂SO₃ + (s) --> Na₂S₂O₃
6. Add dilute hydrochloric acid to crystals of sodium sulfite and heat. Sulfur dioxide forms. Tests for sulfur dioxide.
SO₃^2- + 2H⁺ --> SO₃ (g) + H₂O
7. Add barium chloride solution to a freshly made solution of sodium sulfite. A white precipitate of barium sulfite forms. Unlike barium sulfate, barium sulfite is soluble in dilute hydrochloric acid.
SO₃^2- + Ba²⁺ --> BaSO₃ (s)
8. Add drops of iodine in potassium iodide solution (tincture of iodine) to sodium sulfite solution. The iodine loses its colour. Test the final solution for sulfate ion.
SO₃^2- + H₂O + I₂ --> SO₄^2- + 2I^- + 2H⁺

12.18.5.0 Reactions of sulfuric acid
Sulfuric acid acts as a dehydrating agent, removing water, or the elements of water, from another substance.
1. Add 2 cm of concentrated sulfuric acid to 1 cm of copper sulfate crystals. After ten minutes, note the colour change from blue copper sulfate crystals to white anhydrous copper sulfate.
CuSO₄.5H₂O + (H₂SO₄) --> CuSO₄ + (H₂SO₄.5H₂O)
2. Heat a mixture of 0.5 cm of sucrose and 1.0 cm of concentrated sulfuric acid gently for 2 seconds and then leave to stand.. Note the vigorous reaction and the colour change from white sugar to black carbon.
C₁₂H₂₂O₁₁ + (H₂SO₄) --> 12C + (H₂SO₄.11H₂O)
Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most other acids.
3. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium chloride. Test the fuming gas with silver nitrate solution on a glass rod to form white silver chloride with hydrogen chloride. The less volatile sulfuric acid displaces the hydrogen chloride.
NaCl + H₂SO₄ --> HCl + NaHSO₄
4. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium acetate. Note the smell of the displaced the acetic acid.
5. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium formate. Note the displacement of formic acid followed by dehydration. Sulfuric acid acts as a displacer of acids from their salts, sulfuric acid being much less volatile than most
other acids. Also, sulfuric acid acts as an oxidizing agent.
6. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium bromide. A fuming gas first forms then a brown gas. Hydrogen bromide is displaced then partially oxidized to bromine. Hydrogen bromide turns silver nitrate on a glass rod to pale yellow silver bromide. Potassium permanganate solution on a glass rod decolorizes the sulfur dioxide, formed by reduction of sulfuric acid.
KBr + H₂SO₄ --> HBr + KHSO₄
2HBr + H₂SO₄ --> Br₂ (g) + 2H₂O + SO₂
4H⁺ + 2Br- + SO₄^2- --> Br₂ (g) + 2H₂O + SO₂ (g)
7. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium iodide. A fuming gas first forms then a brown gas. Hydrogen iodide is displaced then oxidized to iodine. Note the greater extent of oxidation compared with the previous experiment. Much of the hydrogen iodide is oxidized to iodine. Heat the test-tube and note the violet vapour of iodine.
4H⁺ + 2I^- + SO₄^2- --> I₂+ 2H₂O + SO₂
8. Reactions of dilute sulfuric acid as an acid. Add 2 cm of dilute sulfuric acid to 1 cc of zinc powder. Close the test-tube with the thumb until enough hydrogen forms to give a mild explosion when the mouth of the test-tube is held in a flame.
2H⁺ + Zn (s) --> Zn²⁺ + H_{2 (g)}
9. Add 2 cm of sodium carbonate to 1 cc of zinc powder. Test for carbon dioxide by passing the gas given off to pass into limewater that turns milky due to the fine precipitate of calcium carbonate.
2H⁺ + CO₃^2- --> H₂O + CO₂ (g)
Ca(OH)₂ + CO₂ (g) --> CaCO₃ (s) + H₂O
10. Reactions of dilute sulfuric acid as a sulfate. Add an equal volume of barium chloride solution to 3 cm of dilute sulfuric acid. Note the white precipitate of barium sulfate. Allow the precipitate to settle, filter, wash and leave to dry.
SO₄^2- + Ba²⁺ --> BaSO₄ (s)

12.18.5.1 Sulfuric acid dehydrates copper (II) sulfate crystals
Sulfuric acid removes water, or the elements of water, from another substance
1. Add 2 cm of concentrated sulfuric acid to 1 cm of copper (II) sulfate crystals. After ten minutes, note the colour change from blue copper (II) sulfate crystals to white anhydrous copper (II) sulfate.
CuSO₄.5H₂O + (H₂SO₄) --> CuSO₄ + (H₂SO₄.5H₂O)

12.18.5.2 Sulfuric acid acts as a displacer of acids from their salts, sodium chloride, sodium acetate, sodium formate
Sulfuric acid is much less volatile than most other acids.
1. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium chloride. Test the fuming gas with silver nitrate solution on a glass rod to form white silver chloride with hydrogen chloride. The less volatile sulfuric acid displaces the hydrogen chloride.
NaCl + H₂SO₄ --> HCl + NaHSO₄
2. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium acetate. Note the smell of the displaced the acetic acid.
3. Add an equal volume of concentrated sulfuric acid to 0.5 cm of sodium formate. Note the displacement of formic acid followed by dehydration.

12.18.5.3 Sulfuric acid acts as a displacer of acids from their salts, potassium bromide, potassium iodide
Sulfuric acid being much less volatile than most other acids. Also, sulfuric acid acts as an oxidizing agent.
1. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium bromide. A fuming gas first forms then a brown gas. Hydrogen bromide is displaced then partially oxidized to bromine. Hydrogen bromide turns silver nitrate on a glass rod to pale yellow silver bromide. Potassium permanganate solution on a glass rod decolorizes the sulfur dioxide, formed by reduction of sulfuric acid.
KBr + H₂SO₄ --> HBr + KHSO₄
2HBr + H₂SO₄ --> Br₂ (g) + 2H₂O + SO₂
4H⁺ + 2Br- + SO₄^2- --> Br₂ (g) + 2H₂O + SO₂ (g)
2. Add an equal volume of concentrated sulfuric acid to 0.5 cm of potassium iodide. A fuming gas first forms then a brown gas. Hydrogen iodide is displaced then oxidized to iodine. Note the greater extent of oxidation compared with the previous experiment. Much of the hydrogen iodide is oxidized to iodine. Heat the test-tube and note the violet vapour of iodine.
4H⁺ + 2I^- + SO₄^2- --> I₂ + 2H₂O + SO₂

12.18.5.4 Reactions of dilute sulfuric acid as an acid
1. Add 2 cm of dilute sulfuric acid to 1 cc of zinc powder. Close the test-tube with the thumb until enough hydrogen gas forms to give a mild explosion when the mouth of the test-tube is held in a flame.
2H⁺ + Zn (s) --> Zn²⁺ + H_{2 (g)}
2. Add 2 cm of sodium carbonate to 1 cc of zinc powder. Tests for carbon dioxide by passing the gas given off to pass into limewater that turns milky because of the fine precipitate of calcium carbonate.
2H⁺ + CO₃^2- --> H₂O + CO₂ (g)
Ca(OH)₂ + CO₂ (g) --> CaCO₃ (s) + H₂O

12.18.5.5 Reactions of dilute sulfuric acid as a sulfate
Add an equal volume of barium chloride solution to 3 cm of dilute sulfuric acid. Note the white precipitate of barium sulfate. Allow the precipitate to settle, filter, wash and leave to dry.
SO₄^2- + Ba²⁺ --> BaSO₄ (s)

12.18.5.6 Reactions of sulfuric acid with sodium chloride
In absence of water
NaCl + H₂SO₄ → NaHSO₄ + HCl (at room temperature)
NaCl + NaHSO₄ → HCl + Na₂SO₄ (if reaction proceeds above 200^oC)

12.18.6.1 Prepare sodium thiosulfate crystals, Na₂S₂O₃.5H₂O, "hypo"
Put 150 mL of water, 30 g of sodium sulfite and 15 g of crushed sulfur in a 250 mL round bottom flask and fit it with a reflux condenser. Heat the flask on a gauze for three hours. Filter the solution and evaporate to 30 mL. Leave to cool and crystallize.
SO₃^2- + S (s) --> S₂O₃^2- (thiosulfite ion = S₂O₃^2-)

12.18.6.2 Reactions of sodium thiosulfate
1. Heat crystals of sodium thiosulfate in a dry test-tube until the test-tube begins to melt. Note water and sulfur as products of the reaction. Leave the mixture to cool. Add dilute hydrochloric acid to the residue and note that hydrogen sulfide is given off.
4Na₂S₂O₃ --> 3Na₂SO₄+ Na₂S₅
Na₂S₅ --> Na₂S + 4S (s)
2. Add iodine solution to sodium thiosulfate solution. The iodine loses its colour and thiosulfate ion is converted to tetrathionate ion.
2S₂O₃^2- + I₂ --> S₄O₆^2- + 2I^- (tetrathionate ion = S₄O₆^2-)
3. Add chlorine water or bromine water in excess to sodium thiosulfate solution and test with barium chloride solution. The products are sodium sulfate and sulfur that may be further oxidized to sulfuric acid.
S₂O₃^2- + Cl₂ + H₂O --> SO₄^2- + S (s) + 2H⁺ + 2Cl^-
4. Add concentrated hydrochloric acid to sodium thiosulfate solution. Sulfur precipitates and sulfur dioxide forms.
S₂O₃^2- + 2H⁺ --> SO₂ (g) + H₂O + S (s)

12.18.7 Reactions of sulfamic acid, NH₂.SO₂OH
(H₃NSO₃) (tautomer: NH₂.SO₂OH) (amidosulfonic acid, amidosulfuric acid, aminosulfonic acid, sulfamidic acid)
Sulfamic acid is used in preparations to clean stainless steel and copper utensils to remove hard water scale. It is also used to make sweeteners.
1. Dissolve 1 cc of sulfamic acid in 2 cm of water. Note the high solubility of the acid. Tests for sulfate ion by adding dilute hydrochloric acid and drops of barium chloride solution. At first there is little action but leave to stand and white suspension of barium sulfate forms. Boil the mixture and the barium sulfate becomes more apparent as the sulfamic acid hydrolyses.
NH₂.SO₂.OH + H₂O --> NH₄HSO₄
2. Dissolve 1 cc of sulfamic acid in 2 cm of water. Dissolve 1 cc of sodium nitrite in 2 cm of water. Mix the solutions. Note the vigorous effervescence as nitric oxide, nitrogen dioxide and nitrogen are given off. Sulfamic acid is a strong fully ionized acid that reacts with the nitrites to give oxides of nitrogen and its -NH₂ group. Sulfamic acid also reacts with the nitrite to give nitrogen.
2H⁺ + 2HNO₂ + 2e^- --> 2H₂O + 2NO
2NO + O₂ --> 2NO₂
NH₂.SO₂.O^- + H⁺ + NO₂^- --> N₂ (g) + HSO₄^- + H₂O
Tests for sulfate ion by adding dilute hydrochloric acid and drops of barium chloride solution. A white suspension of barium sulfate forms.
3. Add sodium hydroxide solution to 1 mL of sulfamic acid to a depth of 2 cm for an excess of sodium hydroxide. Heat the solution and test the gas formed for ammonia with damp red litmus paper.
NH₂.SO₂.O^- + 2OH^- --> NH₃ (g) + SO₄^2- + H₂O
4. Heat 1 cc of sulfamic acid in a dry test-tube. Tests for sulfur dioxide with a spot of potassium permanganate on a filter paper. Test for sulfur trioxide by allowing the white fumes to flow into a test-tube containing barium chloride solution acidified with hydrochloric acid. Note the crystalline sublimate and dissolve the crystals in 2 cm of sodium hydroxide solution. Heat the solution then tests for ammonia with damp red litmus paper. Acidify the remaining solution with hydrochloric acid and add barium chloride solution to tests for sulfate ion.
5. Sulfamic acid solution in water is unstable and forms ammonium bisulfate, however, the colourless crystalline solid is stable, not hygroscopic, and is very soluble in water to form the zwitterion (H3N⁺SO3^-). On heating, the solution produces ammonia.
6. Sulfamic acid reaction with nitrous acid to form nitrogen
HNO₂ + (NH₂)HSO₃ --> H₂SO₄ + N₂ + H₂O
7. Sulfamic acid reacts with nitric acid to form nitrous oxide
HNO₃ + (NH₂)HSO₃ --> H₂SO₄ + N₂O + H₂O
8. Sodium hydroxide solution is standardized by titration with primary standard sulfamic acid solution
NaOH + (NH₂)HSO₃ --> NaNH₂SO₃ + H₂O