School Science Lessons
Topic 12f Reactions nitrogen, nitrates, nitrites, ammonia and oxygen
2009-09-17
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites

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Table of contents
12.11 Nitrogen experiments
12.12 Oxygen experiments

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12.11 Nitrogen experiments
3.46.0 Nitrogen, N_2,
3.44.0 Nitrogen monoxide (nitric oxide), NO
3.45.0 Dinitrogen oxide (nitrous oxide), N₂O
3.47.0 Nitrogen dioxide, NO₂
12.11.1 Reactions of the nitrites, NO₂^-
12.11.2 Reactions of the nitrates, NO₃^-
12.11.3 Ammonia, NH_3, and the ammonium ion, NH₄⁺
12.11.3 Ammonia and the ammonium ion, NH_3,NH₄⁺
10.1.1.1 Tests for ammonia

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12.12 Oxygen experiments
Oxygen O₂,
12.12.1 Prepare oxides by direct oxidation, O^2-
12.12.2 Prepare oxides by indirect oxidation, O^2-
12.12.3 Prepare hydrogen peroxide solution
12.12.4 Hydrogen peroxide reacts as an oxidizing agent
12.12.5 Hydrogen peroxide acts as a reducing agent
12.12.6 Tests for hydrogen peroxide, ionization reaction
12.12.7 Prepare oxygen with hydrogen peroxide using catalysts
12.17.1.1 Different oxides and the periodic table, O^2-

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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 -->- 2NO₂
The 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

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12.11.2 Reactions of nitrates
The oxide N₂O₅ is unstable above 0^oC and forms nitrogen dioxide and oxygen.
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)
3. Brown ring test
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.
4. Reduce nitrate to ammonia. 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.
5. Repeat the experiment with Devarda's alloy replacing the aluminium or zinc powder. (Devarda's alloy = 45% Al, 50% Cu, 5% Fe).

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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.
NH⁺ + OH^- --> NH₃(g) + H₂O
By hydrolysis, ammonium salts may react acidic.
NH₄⁺ + H₂O --> NH₄OH + H⁺
Ammonium salts sublime when heated.

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12.12 Oxygen
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.

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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.

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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

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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)

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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.

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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 X 10^-12
H₂O₂ + H₂O <--> H₃O⁺ + HO₂^-

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12.12.7 Prepare oxygen with hydrogen peroxide using catalysts
To 3 mL of 6% hydrogen peroxide solution, add:
1. powdered manganese dioxide, MnO₂,
2. 0.5 mL of 1 M FeCl₃,
3. 20 grains of active dry yeast.
Test the oxygen formed by these catalytic reactions with a glowing splint.
2H₂O₂ --> 2H₂O + O₂(g)
or
H₂O₂(aq) --> H₂O(l) + 1/2O₂(g)

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