topic12

School Science Lessons
Topic 12 Acids, chemical reactions, chemical equations
2012-02-04 SP
Please send comments to: J.Elfick@uq.edu.au
See: Interesting websites

Table of contents
12.3.0 Acids
12.1.0 Chemical equations, conservation of mass
3.70 Chemical reactions
12.2.0 Chemical reactions, types of chemical reactions

12.3.0 Acids
12.3.15 Acids with salts
12.3.0.1 Amphoteric substances
12.10.0 Boric acid
12.3.12 Concentrated acids with metals, nitric acid with copper
12.3.13 Concentrated acids with metals, sulfuric acid with copper
12.3.14 Concentrated acids with non-metals, carbon, sulfur
12.3.5.01 Copper sulfate from copper oxide
12.3.8 Dilute acids with acidic oxides (non-metal oxides), carbon dioxide, sulfur dioxide
12.3.5 Dilute acids with basic oxides (metal oxides), copper (II) oxide
12.3.5.1 Dilute acids with basic oxides, (metal oxides) zinc oxide
12.3.9.0 Dilute acids with carbonates, common carbonates
12.3.7.2 Dilute hydrochloric acid with hydroxides
12.3.6 Dilute acids with hydroxides, magnesium hydroxide
12.3.7 Dilute acids with hydroxides, sodium hydroxide
12.3.2 Dilute acids with metals, hydrochloric acid
12.3.2.1 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
12.3.3.1 Dilute acids with metals, sulfuric acid with aluminium
12.3.3 Dilute acids with metals, sulfuric acid with iron
12.3.4 Dilute acids with non-metals, carbon, sulfur
12.3.10.1 Dilute acids with calcium hydrogen carbonate
12.3.10.0 Dilute acids with sodium hydrogen carbonate
12.3.7.1 Dilute acids with sodium hydroxide
12.3.9.1 Dilute hydrochloric acid with calcium carbonate
12.3.9.5 Dilute hydrochloric acid with calcium carbonate
12.3.9.2 Dilute hydrochloric acid with sodium carbonate
12.3.11.0 Dilute nitric acid with copper
12.3.9.6 Dilute sulfuric acid with calcium carbonate
12.3.9.4 Dilute tartaric acid with egg shell, soil, wood ash
12.3.9.3 Dilute tartaric acid with sodium carbonate
12.4.0 Hydrochloric acid
12.3.0.5 Ionization reaction of carbonic acid
12.3.3.3 Iron with sodium hydrogen sulfate
12.3.3.2 Magnesium with sodium hydrogen sulfate
12.5.0 Nitric acid
12.7.0 Nitrous acid
12.3.0.4 pH
12.9.0 Phosphoric acid
12.3.0.2 Polyprotic acids
12.3.0 Properties of acids
12.3.0.3 Strong acids and weak acids
12.18.5 Sulfuric acid
12.3.1 Taste of acids, solid acids in the home

12.1.0 Chemical equations, conservation of mass
12.1.0 Chemical equations and ionic equations, conservation of mass
12.1.1 Conservation of mass with effervescent tablets, health salts, sodium bicarbonate (baking soda)
12.1.2 Conservation of mass with burning steel wool
12.1.3 Conservation of mass in a cycle of copper reactions, the copper cycle experiment

12.2.0 Chemical reactions, types of chemical reactions
12.2.5.0 Acid-base reactions
12.2.3.0 Decomposition reactions
12.2.4.0 Displacement reaction (substitution reaction)
12.2.1.0 Precipitation reactions, double decomposition reactions, double exchange, metathesis
12.2.7.0 Polymerization reactions
12.2.6.0 Redox reactions (oxidation-reduction reactions, electron transfer reactions)
12.2.6.1 Redox reactions, Disproportionation
12.2.2.0 Synthesis reactions (combination reactions), mixtures and compounds

12.2.3.0 Decomposition reactions
12.2.3.1 Decomposition reactions, Examples of decomposition reactions
12.2.3.2 Decomposition of copper carbonate to prepare copper oxide
12.2.3.3 Decomposition of zinc carbonate to prepare zinc oxide

12.2.4.0 Displacement reaction (substitution reaction)
12.2.4.1 Displacement, Iron displaces copper in copper sulfate solution
12.2.4.2 Displacement, Prepare zinc sulfate crystals
12.2.4.3 Displacement, Magnesium displaces hydrogen in ethanoic acid

12.2.1.0 Precipitation reactions, reactions of ions in solutions, double decomposition reactions, double exchange, metathesis
12.2.1.2 Cobalt chloride with calcium hydroxide
12.2.1.1 Sodium carbonate with calcium hydroxide
12.2.1.3 Sodium carbonate with zinc sulfate
12.2.1.4 Sodium carbonate with magnesium sulfate
12.2.1.11 Sodium chloride with sulfuric acid
12.2.1.5 Sodium hydroxide with copper sulfate
12.2.1.6 Sodium hydroxide with cobalt chloride
12.2.1.7 Sodium hydroxide with iron sulfate
12.2.1.8 Sodium hydroxide with alum (Al₂(SO₄)₃.K₂(SO₄).24H₂O, aluminium potassium sulfate)
12.2.1.10 Prepare soda lime
12.2.1.9 Prepare sodium hydroxide solution

12.2.2.0 Synthesis reactions (combination reactions), mixtures and compounds
12.2.2.1 Heat iron with sulfur
12.2.2.2 Heat copper with iodine
12.2.2.3 Heat iron with iodine
12.2.2.4 Zinc with iodine
12.2.2.5 Combination reaction to form a hydrate

12.4.0 Hydrochloric acid
12.4.0 Hydrochloric acid
3.42 Prepare hydrogen chloride, HCl
3.42.01 Prepare hydrochloric acid
3.42.1.0 Tests for hydrogen chloride

12.5.0 Nitric acid
12.5.0 Nitric acid
12.5.0a Grades of nitric acid
12.5.1 Prepare nitric acid, sulfuric acid with sodium nitrate
12.5.2 Nitric acid with copper
12.3.11.1 Nitric acid with metals
12.5.3 Nitric acid with sulfuric acid
12.5.3 Nitric acid with sulfuric acid

12.7.0 Nitrous acid
12.7.3 Heat nitrous acid
12.5.01 Ionization of nitrous acid
12.7.1 Prepare nitrous acid
12.7.01 Nitration
15.2.5.1 Nitrous acid as oxidizing agent or a reducing agent
12.7.2 Nitrous acid with amines

12.1.0 Chemical equations and ionic equations, conservation of mass
Chemical reactions are influenced by the conditions under which they take place and, being reversible, may reach a state of equilibrium. Chemical reactions occur at different rates and changing the nature of the reactants, temperature, or concentration, or introducing a catalyst, may alter these. Chemical reactions may be reversible. Reversible chemical reactions may reach a state of dynamic balance known as equilibrium which, when disturbed, will be re-established.
Specific criteria can be used to classify chemical reactions. Redox reactions involve a transfer of electrons and a change in oxidation number. Precipitation reactions result in the appearance of a solid from reactants in aqueous solution. Acid-base reactions involve transfer of protons from donors to acceptors. Polymerization reactions produce large molecules with repeating units. Chemical reactions involve energy changes. All chemical reactions involve energy transformations. The spontaneous directions of chemical reactions are towards lower energy and greater randomness. Chemical reactions, but not nuclear reactions, obey the law of conservation of mass. The mass of the products is equal to the mass of the reactants. In a balanced chemical equation, the number of atoms of each element in the reactants is equal to the number of atoms in the product. Ionic equations show soluble ionic compounds in solution as separate ions and the sum of charges on the product side is equal to the sum of charges on the reactant side.
Chemical equation: HCl (aq) + NaOH (aq) --> NaCl (aq) + H₂O (l)
Ionic equation: H⁺ (aq) + Cl^- (aq) + Na⁺ (aq) + OH^- (aq) --> H₂O (l) + Cl^- (aq) + Na⁺ (aq)
The Cl^- and Na⁺ ions are called spectator ions because they are on both sides of the equation and do not react. Cancel the spectator ions to leave the net ionic equation:
H⁺ + OH^---> H₂O

12.1.1 Conservation of mass with effervescent tablets, health salts, sodium bicarbonate (baking soda)
See diagram 12.1.1: Conservation of matter | 13.7.7 Prepare carbon dioxide, heat hydrogen carbonates | See 19.2.13: Prepare fruit salts
When a chemical reaction occurs, matter is neither created nor destroyed. The mass of the reactants = the mass of the products. Effervescent tablets or "fruit salts" contain sodium hydrogen carbonate, and dry citric acid or tartaric acid
Put the tablet or fruit salts in water in a test-tube. Carbon dioxide forms as bubbles and any other substance in the tablet or fruit salts dissolves easily. Tests for carbon dioxide by holding a test-tube containing limewater at an angle near the mouth of the test-tube containing the effervescent tablet. The carbon dioxide given off by effervescence is heavier than air and will roll into the limewater test-tube where the limewater will turn milky because of the presence of carbon dioxide. Weigh an effervescent tablet or fruit salts and put into the bottom corner of a small plastic bag. Twist the bag above the corner and tie around the twist with thin string or wire. Weigh the plastic bag + string + effervescent substance. Pour a known amount of water into the open bag then tie string tightly to close the bag so that no liquid or gas can escape. The weight of the water + bag + string + effervescent substance is now known. Undo the string around the twisted part of the bag and untwist the bag. The acid and sodium hydrogen carbonate dissolves in the water and react to produce a salt and carbon dioxide. Weigh the bag and products of the reaction. The weight is the same.
2. Place a small amount of water in a plastic cup and place on the pan of the electronic scale.
Cover the top with a piece of paper or metal foil. Place two Alka Seltzer tablets, (ant-acid tablets), on top of the cover. Record the initial mass. Tilt the cover so that the tablets drop into the water and immediately replace the cover so water droplets will not escape. Record the mass reading until it is constant. Repeat this experiment with the cup enclosed in a sealed container.
3. Place a small amount of water in a large 2 litre plastic drink bottle. Break two Alka Seltzer tablets in pieces that will fit in the bottle. Weigh the bottle. bottle cap, and Alka Seltzer tablets together. Drop the Alka Seltzer tablet pieces into the bottle and quickly replace the bottle cap tightly and place back on the scale. Record the mass reading until it is constant. The results differ from the experiment carried out in an open container. Remove the bottle from the scale, loosen the bottle cap, and measure again the weight of the bottle and its cap.

12.1.2 Conservation of mass with burning steel wool
Roll some steel wool into a ball in your hands and weigh it. Hold it with tongs over a sheet of paper. Heat the steel wool until red-hot; remove the flame and blow gently on the red-hot steel until it stops burning. When cold, weigh the steel wool and any fragments fallen onto the sheet of paper on the balance. When iron burns the product formed, iron oxide, is heavier than the iron.

12.1.3 Conservation of mass in a cycle of copper reactions, the copper cycle experiment
See diagram 1.13a: Simple fume hood
Step 1. Convert copper metal to copper nitrate
1. Weigh 1.000 g of copper wire. It must be clean, bright and shiny. Twist the wire into a flat spiral and put it in a beaker in a fume hood. Slowly add 4 .0 mL of 16 M nitric acid. Be careful! Note the brown fumes of nitric oxide (NO) nitrogen dioxide (NO₂) and dinitrogen tetroxide (N₂O₄), i.e. NOx. When all the copper is dissolved, add 100 mL of deionized water.
Cu (s) + 8HNO₃ (aq) --> 3Cu(NO₃)₂ (aq) + 4H₂O (l) + 2NO (g) [dilute nitric acid]
Cu (s) + 4HNO₃ (aq) --> Cu(NO₃)₂ (aq) + 2H₂O (l) + 2NO₂ (g) [concentrated nitric acid]
N₂O₄ <--> 2NO₂ [in equilibrium]
Step 2 Convert copper nitrate to copper hydroxide
Add 30.0 mL of 3.0 M sodium hydroxide to the solution while stirring. If red litmus paper does not turn blue in the solution, add more sodium hydroxide. Note the precipitate of copper hydroxide, an ionic solid.
Cu(NO₃)₂ (aq) + 2NaOH (aq) --> Cu(OH)₂ (s) + 2NaNO₃ (aq)
Step 3 Convert copper hydroxide to copper oxide
Heat the solution on a hot plate while continually stirring to prevent bumping caused by steam bubbles. Note the precipitate changing to a black solid, copper oxide, CuO. Carefully decant the liquid, add deionized water and decant again.
Heat the precipitate until it becomes a firm mass.
Cu(OH)₂ (s) + heat --> CuO (s) + H₂O (l)
Step 4 Convert copper oxide to copper sulfate
Reaction 4: Converting copper oxide to copper sulfate
Add 15 mL of 6.0 M sulfuric acid to the copper oxide while swirling, not stirring, the copper oxide to help it dissolve.
CuO (s) + H₂SO₄ (aq) --> CuSO₄ (aq) + H₂O (l)
Step 5 Convert copper sulfate to copper metal
CuSO₄ (aq) + Zn (s) --> ZnSO₄ (aq) + Cu (s)
In the fume hood add 2.0 g of zinc metal and keep stirring until the solution becomes colourless. The zinc is oxidized as it reduces the copper.
Add one drop of the solution to 1 mL of concentrated ammonia solution in a test-tube. If the ammonia turns deep blue, some unreduced copper is till in the solution so the reaction is not finished. When all the copper is reduced, decant the liquid and add 20.0 mL of 6M hydrochloric acid to dissolve excess zinc. Note the bubbles of hydrogen gas until the reaction is complete. Cool the beaker and observe the metallic copper settling on the bottom. Carefully decant the solution and wash the copper metal with deionized water and decant again. Transfer the copper to an evaporating dish using a small amount of deionized water. Decant excess water evaporating dish then wash the precipitate with methanol. Decant the liquid the gently heat the precipitate on a hot plate. If you heat the copper precipitate too strongly it will oxidize to copper oxide. Transfer the dried copper metal to a preweighed beaker and calculate the mass of recovered copper.
Percentage recovery = mass of recovered copper / initial mass of copper X 100.

12.2.0 Chemical reactions, types of chemical reactions
Chemical reactions involve energy changes. All chemical reactions involve energy transformations. The spontaneous directions of chemical reactions are towards lower energy and greater randomness. In a chemical reaction, a chemical change occurs where elements or compounds (reactants) form new substances (products). Specific criteria can be used to classify chemical reactions. The main types of chemical reactions are precipitation reactions, synthesis reactions, decomposition reactions, displacement reactions, acid-base reactions, redox reactions and polymerization reactions that produce large molecules with repeating units.

12.2.1.0 Precipitation reactions, reactions of ions in solutions, double decomposition reactions, double exchange, metathesis
Precipitation reactions result in the appearance of a solid from reactants in aqueous solution. Salts can be made by the action of acids with alkalis, carbonates, metals, metal oxides and by replacement and double decomposition reactions. A salt contains a metal and part of an acid, e. g. copper sulfate from sulfuric acid, sodium chloride from hydrochloric acid. A salt is a compound formed when the hydrogen of an acid is replaced by a metal. For example, when zinc reacts with hydrochloric acid it replaces the hydrogen and forms the salt, zinc chloride. The hydrogen comes away as hydrogen gas. In a double decomposition reaction the radicals are exchanged. A double decomposition reaction can occur if one of the products (substances formed) is insoluble or is a gas.
Reactions of ions in solutions, double decomposition reaction
A reaction between ions is shown by precipitation of an insoluble salt as a solid. Use (aq) to show a solution and use (s) to show a precipitate, solid. A chemical reaction between two substances that produces two other substances is called a double decomposition reaction (double replacement reaction, precipitation reaction, metathesis).
AB + CD ---> AC + BD
1. When sodium chloride solution is added to silver nitrate solution an insoluble solution of of silver chloride forms.
AgNO₃ (aq) + NaCl (aq) --> AgCl (s) + NaNO₃ (aq)
silver nitrate + sodium chloride --> silver chloride + sodium nitrate
The silver chloride precipitate can be separated from the sodium nitrate solution.
AgNO₃ (aq) + KCl (aq) --> KCl (s) + KNO₃ (aq)
silver nitrate + potassium chloride --> silver chloride + potassium nitrate
Ionic equations
The ionic equation that shows all the substances
Ag⁺ (aq) + NO₃^- (aq) + Na⁺ (aq) + Cl^- (aq) --> AgCl (s) + Na⁺ (aq) + NO₃^- (aq)
The net ionic equation that does not contain the "spectator ions" that appear on both sides of the equation but do not form a precipitate
Ag⁺ (aq) + Cl^- (aq) --> AgCl (s)
The above reaction may also be called a double decomposition reaction (metathesis) because the positive and negative parts of two compounds swap partners, i.e. exchange radicals. In general: AB + CD --> AD + CB
2. silver nitrate + potassium chloride ---> potassium nitrate + silver chloride (white precipitate)
AgNO₃ (aq) + KCl (aq) ---> KNO₃(aq) + AgCl (s)
Ag⁺ (aq) + Cl^- (aq) ---> AgCl (s)
3. calcium chloride + sodium carbonate --> calcium carbonate + sodium chloride
CaCl₂ (aq) + Na₂CO₃(aq) --> CaCO₃ (s) + 2NaCl (aq)
Ca²⁺ (aq) + CO₃^2- (aq) --> CaCO₃(s)
4. A reaction that forms a coloured precipitate is a good way to show double decomposition reactions.
4.1 lead (II) nitrate (aq) + potassium dichromate (aq) ---> lead (II) chromate (IV) (s) (yellow precipitate, chrome yellow)
Pb²⁺ (aq) + CrO₄^2-(aq) ---> PbCrO₄(s)
4.2 silver nitrate (aq) + potassium chromate (aq) ---> silver chromate (s) (red precipitate)
Ag⁺ (aq) + CrO₄^2- (aq) ---> AgCrO₄ (s)
4.3 lead nitrate (aq) + potassium iodide (aq) ---> lead iodide (s) (yellow precipitate)
Pb²⁺ (aq) + 2I^- (aq) ---> PbI₂ (s)
4.4 copper (II) sulfate (aq) + sodium carbonate (aq) ---> copper carbonate (s) (green precipitate)
Cu²⁺ (aq) + CO₃^2- (aq) ---> CuCO₃ (s)

12.2.1.1 Sodium carbonate with calcium hydroxide, double decomposition
To half a test-tube of sodium carbonate solution add limewater (calcium hydroxide solution). Describe the reaction which occurs. A white solid (precipitate) of calcium carbonate forms.
sodium carbonate + calcium hydroxide --> calcium carbonate + sodium hydroxide.

12.2.1.2 Cobalt chloride with calcium hydroxide, double decomposition
To half a test-tube of cobalt chloride solution add limewater (calcium hydroxide solution). Describe the reaction which occurs. A pink / mauve precipitate of cobalt hydroxide forms.
cobalt chloride + calcium hydroxide --> cobalt hydroxide + calcium chloride

12.2.1.3 Sodium carbonate with zinc sulfate, double decomposition
Add sodium carbonate solution to the zinc sulfate solution. Filter the mixture to obtain the precipitate of zinc carbonate. Allow the filter paper, unfolded, to dry, and scrape off the white powder.
sodium carbonate + zinc sulfate --> sodium sulfate + zinc carbonate

12.2.1.4 Sodium carbonate with magnesium sulfate, double decomposition
Add magnesium sulfate solution to sodium carbonate solution. Filter the solution to obtain the white precipitate of insoluble magnesium carbonate.
magnesium sulfate + sodium carbonate --> magnesium carbonate + sodium sulfate

12.2.1.5 Sodium hydroxide with copper sulfate
Add 5 mL of sodium hydroxide solution to the same volume of copper sulfate solution. Describe what you see. A blue precipitate of copper hydroxide forms. sodium hydroxide + copper sulfate –> copper hydroxide + sodium sulfate.

12.2.1.6 Sodium hydroxide with cobalt chloride
Add 5 mL of sodium hydroxide solution to the same volume of cobalt chloride solution. Describe what you see. A blue-green precipitate of cobalt hydroxide forms.
Sodium hydroxide + cobalt chloride –> cobalt hydroxide + sodium chloride.
Repeat experiment 153, substituting a dilute solution of cobalt chloride for the copper sulfate solution.

12.2.1.7 Sodium hydroxide with iron sulfate
Add 5 mL of sodium hydroxide solution to a dilute solution of iron sulfate. Stopper the test-tube, shake well, and leave to stand. Describe what you see. A green precipitate forms that turns brown. The green iron hydroxide first forms, but it soon reacts with oxygen gas to form a different type of iron hydroxide which is brown.

12.2.1.8 Sodium hydroxide with alum (aluminium potassium sulfate, Al₂(SO₄)₃.K₂(SO₄).24H₂O)
Add 5 mL of sodium hydroxide solution to a dilute solution of alum.. Describe what you see. A faint white precipitate forms. The part of alum that reacts with the alkali is aluminium sulfate.

12.2.1.9 Prepare sodium hydroxide solution
Make a concentrated sodium carbonate solution and add to it calcium hydroxide solution. Shake well for three minutes. Describe what you see. A white precipitate forms. Filter into a clean bottle and label it. This double decomposition reaction between solutions of sodium carbonate and calcium hydroxide occurs because one of the products, calcium carbonate, is insoluble. The filtrate is sodium hydroxide solution.
sodium carbonate + calcium hydroxide --> sodium hydroxide + calcium carbonate (precipitate)
Test the sodium hydroxide solution with universal indicator paper. The universal indicator paper turns blue or violet.

12.2.1.10 Prepare soda lime
Dissolve calcium oxide in sodium hydroxide solution and leave to dry so that a mixture of sodium hydroxide and calcium oxide forms. Soda lime is toxic if ingested and is corrosive.

12.2.1.11 Sodium chloride with sulfuric acid
See 3.71.1: Solubility table and solubility rules | See also 3.71: Reactions of ions in solutions
1. Concentrated sulfuric acid with solid sodium chloride BE CAREFUL!
The reactions contain no water. Two reactions occur and both go to completion if heated. The reactions occur because hydrogen chloride has a lower boiling point than sulfuric acid.
2NaCl (s) + H₂SO₄ (l) --> Na₂SO₄ (aq) + 2HCl (g)
NaCl (s) + H₂SO₄ (l) --> NaHSO₄ (aq) + HCl (g)
2. Dilute sulfuric acid with solid sodium chloride: The reaction does not go to completion because the hydrochloric acid dissolves in the water. One product of the reaction is a slightly ionized substance, e.g. water. In neutralization reactions HOH is forming, so the reaction can almost go to completion. One product of the reaction is a precipitate. An insoluble substance leaves the solution. The solubility rules state that all chlorides are soluble except Ag⁺, Hg²⁺ and Pb²⁺(slightly). Predict whether the following reaction occurs. The reaction occurs because insoluble silver chloride precipitates.
NaCl (aq) + AgNO₃ (aq) --> NaNO₃ (aq) + AgCl (s)
NaOH (aq) + HCl (l) --> NaCl (aq) + H₂O (l)

12.2.2.0 Synthesis reactions (combination reactions)
Reactions of mixtures of two elements, iron with sulfur, copper with sulfur, zinc with sulfur, zinc with iodine to form compounds.
Be careful! The following reactions are vigorous. Do not use large quantities of the chemicals. Use eye protection. Do not get close to the fumes from the reaction.
Elements or simple molecules combine to form a new compound
1. Reactions of zinc with sulfur
Heat a mixture of sulfur in the bottom of a test-tube and a strip of zinc half way up the test-tube to form the compound zinc sulfide
Zn (s) + S (s) --> ZnS (s)
zinc + sulfur --> zinc sulfide
2. Reactions of iron with copper
Heat a mixture of iron filings and copper turnings in a hard glass test-tube. You should not observe any reaction of iron filings with copper because no new compound forms.

12.2.2.1 Heat iron with sulfur
See diagram 12.2.1: FeS
S₈ (s) + 8Fe (s) --> 8FeS (s)
Be careful! The reaction or iron (II) sulfide with hydrochloric acid will form the poisonous gas, hydrogen sulfide, with an odour of rotten eggs.
1. Mix half a metal bottle top of powdered sulfur with the same volume of iron filings. Heat a small portion of the mixture on the metal bottle top with the cork removed or in a hard glass test-tube. When the reaction begins, i.e. the mixture starts to glow, stop heating by moving the Bunsen burner to the side. If the glow stops, heat the test-tube again. The reaction of a mixture of iron with sulfur gives out so much heat that the mixture becomes red-hot. Note the following different properties of powdered sulfur, iron filings and the product, the compound iron (II) sulfide: 1.1 appearance, 1.2 colour, 1.3 hardness, 1.4 magnetism.
Iron is magnetic so is easily removed from a mixture of iron and sulfur but iron (II) sulfide is not magnetic.
Fe + S ---> FeS (s)
Repeat the experiment using the same volumes of powdered copper with sulfur, nickel with sulfur, lead with sulfur, and zinc with sulfur. The compounds formed are called binary compounds.
2. Use < 10 g total material iron with sulfur in a fume cupboard. Heat the mixture to start the reaction. However, be aware that unreacted sulfur may catch fire and produce sulfur dioxide gas to irritate the lungs
3. Mix uniformly reduced iron powder and powdered sulfur in a weight ratio of seven to four. Carve the word "FeS" on a red coloured brick with a knife. Spread the iron sulfur mixture throughout the word groove and press the powdered mixture solid. Heat one tip of a glass rod until red-hot with an alcohol burner and then immediately dig the hot tip into the mixture at one end of the word groove. A chemical reaction is starts immediately. The reaction continues violently to release a large amount of heat and meanwhile to develop rapidly a red glow, which looks like a small "fiery dragon". The heat lost by the reaction is more than the heat needed to start the reaction. The reaction produces a new black solid substance, iron (II) sulfide, that has different properties from the two reactants, iron and sulfur. Compare iron powder, powdered sulfur and iron (II) sulfide. Note their appearance. Test them respectively with a magnet. Add in drops hydrochloric acid solution to them respectively.
4. Mix equal amounts of iron filings and powdered sulfur. Heat the mixture in a crucible or a small tin with sand in the bottom. The sand prevents the bottom of the tin from melting by spreading the heat. Heat the mixture strongly until you see a red glow spreading through the mass. The heat lost by the chemical reaction is more than the heat needed to start the reaction. The reaction forms a new substance iron (II) sulfide that has different properties from the two elements used to make it. Compare iron filings, powdered sulfur, and iron (II) sulfide. Note their appearance. Test with a magnet. Add drops of hydrochloric acid.
5. Make a mixture of 7 parts of iron filings with 4 parts of sulfur powder in a sealed plastic bag. Hold a magnet over the plastic bag to show that the iron filings can be easily separated from the mixture. Quarter fill an ignition tube with the mixture. Near an open window or in a fume cupboard heat the end of the ignition tube with a Bunsen burner. When the mixture glows move the Bunsen burner away but when the glow stops move the Bunsen burn back again until all the mixture reacts. Leave the ignition tube to cool then move the magnet near it. The magnet can no longer attract iron filings or the iron sulfide in the ignition tube.

12.2.2.2 Heat copper with iodine
Mix equal amounts and heat gently in a test-tube.
BE CAREFUL! THE IODINE MAY STAIN THE SKIN. REMOVE STAINS WITH SODIUM THIOSULFATE SOLUTION.
Stop heating when you hear a hissing noise. Heat again to make sure all the copper reacts with the iodine Excess iodine sublimes and solidifies up the tube. Let the tube cool then scrape out the product of the reaction. Compare the crushed product with the reactants copper and iodine. The reaction forms a new substance, copper iodide.

12.2.2.3 Heat iron with iodine
Put iodine crystals in a test-tube and then push in a plug of steel wool. Clamp the test-tube at an angle and heat the steel wool with a Bunsen burner. The steel wool glows red and the iodine evaporates. A new substance forms by direct synthesis

12.2.2.4 Zinc with iodine
1. Do not use a Bunsen burner in this experiment. Dissolve iodine crystals in 5 mL of ethanol then note the temperature of the solution. Add zinc powder to the solution then stir it with the thermometer and note any rise in temperature. When the solution becomes clear, filter it into a beaker leaving behind any excess zinc. Place a watch glass over a beaker of hot water from an electric jug. Pour the filtrate into the watch glass and observe the zinc iodide, a white salt formed by direct synthesis as the alcohol evaporates.
zinc + iodine --> zinc iodide.
Zn + I₂ --> ZnI₂
2. Dissolve 2 g of iodine crystals in 25 mL of ethanol and note the temperature of the solution. Add 2 g of zinc powder to the solution, stir and note the temperature. The colour of the iodine is lost during the reaction and some of the zinc remains. Filter off the unreacted zinc and heat the filtrate in an evaporating basin on an electric hot plate, not over a Bunsen burner. A white odourless crystals of zinc iodide, (ZnI₂), remain. Anhydrous zinc iodide is a white hygroscopic substance used as an X-ray opaque substance to improve the fine details on X-ray plates.
Zn + I₂ --> ZnI₂Repeat the experiment using antimony with iodine.

12.2.2.5 Combination reaction to form a hydrate
Prepare white copper sulfate by heating a finger width of blue copper sulfate crystals in an evaporating basin, while stirring with a glass rod. Do not overheat, and stop heating as soon as the chemical has turned white. Leave to cool. Add drops of water until the powder is blue but still dry. The white compound has combined chemically with the water to form the blue compound, because all the added water has disappeared. This reaction is the combination of an anhydrous salt with water to form a hydrate.

12.2.3.0 Decomposition reactions
A compound breaks down into simpler compounds or into elements, usually caused by heat, the opposite of a synthesis reaction. All compounds decompose on heating to a high enough temperature to form elements or simple molecules. Compounds of metals higher in the reactivity series are harder to decompose by heating than compounds lower in the reactivity series, so copper compounds break up more readily on heating than sodium compounds.
Decomposition of calcium carbonate
CaCO₃ (s) --> CaO (s) + CO₂ (g)
calcium carbonate --> calcium oxide + carbon dioxide

12.2.3.1 Examples of decomposition reactions
1. Place small quantities of zinc oxide and copper oxide in separate small test-tubes and heat carefully. Identify any gases produced. Note any new substance formed in the test-tube.
2. Repeat using sodium or potassium nitrate, lead nitrate, copper nitrate. If a brown gas is produced, it is nitrogen dioxide.
potassium nitrate (s) --> potassium nitrite (s) + oxygen (g)
copper nitrate (s) --> copper oxide (s) + nitrogen dioxide (g) + oxygen (g)
3. Repeat using sodium carbonate, magnesium carbonate, calcium carbonate, lead carbonate.
4. Repeat using sodium sulfate, magnesium sulfate, zinc sulfate and copper (II) sulfate.
zinc sulfate (s) --> zinc oxide (s) + sulfur trioxide (g)
copper (II) sulfate (s) --> copper oxide (s) + sulfur trioxide (g)
5. Repeat using magnesium carbonate
magnesium carbonate (s) --> magnesium oxide (s) + carbon dioxide (g)

12.2.3.2 Decomposition of copper carbonate to prepare copper oxide
1. Prepare copper carbonate by mixing sodium carbonate solution and copper sulfate solution. Pour off the liquid when the copper carbonate has settled in the test-tube. Heat to evaporate remaining liquid and heat more strongly to form the oxide. The oxide could be purified by washing with water, using a filtration apparatus. Most carbonates decompose to form a metal oxide and carbon dioxide, e.g. copper carbonate.
2. Prepare copper carbonate by adding half a test-tube of copper sulfate solution to the same quantity of sodium carbonate solution. Leave the insoluble copper carbonate to settle then pour off the liquid above the precipitate. Gently heat the carbonate just enough to drive off the remaining water as steam. Fit a stopper and bent tube to the test-tube. With the end of the bent glass delivery tube dipping into the limewater solution, heat the copper carbonate more strongly. The copper carbonate turns black and the limewater turns milky. The copper carbonate has been decomposed by the heat into black copper oxide and carbon dioxide gas which turns limewater milky.

12.2.3.3 Decomposition of zinc carbonate to prepare zinc oxide
Heat part of the zinc carbonate, from experiment 163, in an evaporating basin, or stand fairly strongly, and until the white powder turns yellow. You now have zinc oxide. Allow it to cool. Heat again. Note the colour changes. Zinc oxide is yellow when hot, white when cold. The carbonate decomposed on heating to form zinc oxide and carbon dioxide. Most carbonates decompose to form a metal oxide and carbon dioxide, e.g. zinc carbonate.

12.2.4.0 Displacement reactions (substitution reactions)
The reactants are an element and a compound. The element replaces part of the compound with the same valence and same sign, e.g. displacement of Cu²⁺ by zinc
Zn (s) + CuSO₄ (aq) --> ZnSO₄ (aq) + Cu (s)
zinc + copper sulfate --> zinc sulfate + copper
The copper precipitates as the element and the zinc metal goes into solution as zinc ions.
See 3.71 Reactions of ions in solutions | See 12.14.1: Zinc displaces lead from lead nitrate solution

12.2.4.1 Iron displaces copper in copper sulfate solution
1. Prepare iron sulfate crystals. Add iron filings to sodium hydrogen sulfate solution in a test-tube, and heat. When there is no further reaction, filter the mixture and pour a finger width of sulfate solution of the filtrate into a evaporating basin. Describe what you see. Pale green crystals of iron sulfate form.
2. Add iron filings to half a test-tube of copper sulfate solution. Leave until the colour of the solution changes from blue to pale green. Filter the solution. Pour part of the filtrate into an evaporating basin. Describe what you see. Pale green crystals of iron sulfate form no different from the crystals formed in the previous experiment.
3. Place copper sulfate crystals in the test-tube and add a quarter of a test-tube of water. Shake to get a blue solution. Drop in the iron nail, which must not be rusty. Leave for ten minutes. Take out the nail. Describe what you see. The iron nail has turned a pinkish colour, due to a deposit of copper on it.
Iron + copper sulfate –> copper + iron sulfate
Repeat the experiment but use a finger width of iron filings instead of the nail. Leave the iron nail test-tube and contents for a few hours. Describe what you see. The blue liquid turns a pale green sulfate colour. The iron has completely displaced the copper in the copper sulfate, forming a solution of iron sulfate that is pale green.
Repeat the experiment using a centimetre of magnesium ribbon instead of the iron nail. Observe the metal a few moments after it has been in the copper sulfate solution. Describe what you see. The metal turns pink or copper coloured. Note whether iron or magnesium metal reacts more quickly. Magnesium is a more reactive metal than iron.

12.2.4.2 Prepare zinc sulfate crystals
Add pieces of granulated zinc to sodium hydrogen sulfate solution in a test-tube, and heat. When there is no further reaction all the blue colour has disappeared, filter the mixture. The filtrate is zinc sulfate solution.

12.2.4.3 Magnesium displaces hydrogen in ethanoic acid
All acids contain hydrogen, and many metals can displace it, thus setting the hydrogen gas free. The acid in vinegar is ethanoic acid. Add 3 cm of magnesium to half a test-tube of vinegar. As soon as bubbles of hydrogen gas are coming well, hold your thumb or finger over the mouth of the test-tube for half a minute to trap a quantity of hydrogen gas. Quickly hold the open test-tube to the spirit burner flame. Describe what happens. A small pop or squeak occurs. It is a minor explosion. If you got no result, repeat the procedure of trapping the gas and igniting it in the flame. When hydrogen gas mixes with air it explodes, i.e. combines extremely rapidly with the oxygen gas in the air.
Repeat the experiment using tartaric acid (CHOHCOOH)₂ ) instead of ethanoic acid.
Magnesium displaces hydrogen in tartaric acid
Repeat the experiment with iron filings instead of magnesium. You may have to heat the mixture.

12.2.5.0 Acid-base reactions
Acid-base reactions involve transfer of protons from donors to acceptors. An acid dissociates in water to produce positive hydrogen ions, H⁺, that is solvated to produce hydronium ions (hydroxonium ions, oxonium ions) H₃O⁺, by transferring a proton (H⁺) to a water molecule.
HCl (g) + H₂O (l) --> H₃O⁺ (aq) + Cl^- (aq)
A base dissociates in water to produce negative hydroxide ions, OH^-.
NaOH --> Na⁺ (aq) + OH^- (aq)
Acids react with bases to from salts and water. The products are neither acidic nor basic so this reaction is called neutralization.
HCl + NaOH --> NaCl + H₂O
hydrochloric acid + sodium hydroxide --> sodium chloride + water
The ionic equation that shows all the substances
H₃O⁺ (aq) + Cl^- (aq) + Na⁺ (aq) + OH^- (aq) --> 2H₂O + Na⁺ (aq) + Cl^- (aq)
The net ionic equation
H₃O⁺ (aq) + OH^- (aq) --> 2H₂O

12.2.6.0 Redox reactions (oxidation-reduction reactions, electron transfer reactions)
Redox reactions involve a transfer of electrons and a change in oxidation number. Electrons move from one atom to another. Oxidation is loss of electrons. Reduction is gain of electrons. Oxidation reactions and reductions reactions must occur together. The same number of electrons are gained in the reduction as are lost in the oxidation.
In the reaction of dilute hydrochloric acid on magnesium ribbon, each magnesium atom loses two electrons to two hydrogen atoms.
Mg (s) + HCl (aq) --> MgCl₂ (aq) + H₂ (g)
Mg (s) + 2H₃O⁺ (aq) + 2Cl^- (aq) --> H₂ (g) + Mg²⁺ (aq) + 2Cl^- (aq) + 2H₂O
In reactions where no ions form, use the idea of oxidation number (oxidation state) to show the "apparent charge" on an atom. In the reaction between gases:
2SO₂ (g) + O₂ (g) --> 2SO₃ (g)
Give the oxygen atom a net charge of -2, but give O₂ a net charge of zero because the oxygen atom is in the elemental form. Then the sulfur atom in SO₂ has an oxidation number +4 and the sulfur atom in SO₃ has an oxidation number +6. The sulfur atoms have been oxidized because the oxidation number has increased and the oxygen gas atoms in O₂ have been reduced because the oxidation number has decreased.
Similarly, in the the following equation:
NH₃ + CuO --> Cu + H₂O + N₂
The oxidation number of hydrogen atom in NH₃ is +1 and in H₂ is zero, because the hydrogen atom is in the elemental form. The oxidation number of the nitrogen atom has increased from -3 in NH₃ to 0 in N₂, because N in N₂ is in the elemental form. The oxidation number of the copper has decreased from +2 in CuO to zero in Cu, because the Cu atom is in elemental form. The nitrogen atom has been oxidized and the copper atom has been reduced.
However, when chlorine has dissolves in water a disproportionation occurs because the the chlorine becomes both oxidized, when HClO is formed, and reduced, when HCl is formed.
Cl₂ (g) + H₂O (l) <--> HClO (aq) + HCl (aq)

12.2.6.1 Disproportionation
The disproportionation process is when a chemical species is oxidized and reduced simultaneously.
2Cu⁺(aq) --> Cu²⁺(aq) + Cu (s)
Cl₂ (g) + H₂O (l) <--> HClO (aq) + HCl (aq)
(Chlorine dissolving in water)

12.2.7.0 Polymerization reactions
See 16.2.4.2.1: Cyanamides, inorganic, (CN₂^2-), ionization reaction of methylamine, cyanic acid, melamine
Polymerization reactions produce large molecules, polymers with repeating units, monomers. The physical properties of addition and condensation polymers are related to their structure described by the terms thermoset, thermoplastic, elastomer, vulcanization, amorphous, crystalline. Polymer properties depend on the chain length, side branches and cross-linking.

12.3.0 Properties of acids
Acids are good electrolytes, react with active metals, turn blue litmus red, and have a sour taste. Dilute acids contain hydrogen ions in aqueous solution. You can represent the hydrogen ion, which is really a proton, in different ways to show how it is related to the water molecules in the solution. You can show it as the hydrated hydrogen ion, [proton, H⁺ (aq)] or as the hydronium ion [oxonium ion, H₃O⁺ (aq)] but, for convenience, use H⁺ (aq). Concentrated sulfuric acid exists mainly as H₂SO₄molecules. Hydrochloric acid and nitric acid dissociate into ions even in concentrated solution. Weak acids, e.g. ethanoic acid (acetic acid, CH₃COOH) carbonic acid and sulfurous acid dissociate very little in aqueous solution, but their salts, e.g. potassium acetate (CH₃COOK) are completely dissociated into ions.
Using the Bronsted-Lowry definition of acids and bases an acid donates a proton (H+) to another substance and a base accepts a proton from another substance. When sulfuric acid dissociates in water it donates a proton (H⁺) to the water molecule. So in this reaction the water molecule acts as a base.
H₂SO₄ + H₂O --> HSO₄^- + H₃O⁺When ammonia dissolves in water, ammonia accepts a proton and so it is the base. So in this reaction the water molecule acts as an acid.
NH₃ + H₂O <--> NH₄⁺ + OH^-

12.3.0.1 Amphoteric substances can act as an acid or a base. In the above reactions water is acting as a base with sulfuric acid and is acting as an acid with ammonia. Similarly, bicarbonate ion can act as an acid to donate a proton to form carbonate ion: HCO₃^- + H₂O <--> CO₃^2- + H₃O⁺Also, bicarbonate ion can act as a base to accept a proton to form carbonic acid: HCO₃^- + H₂O <--> H₂CO₃ + OH^-

12.3.0.2 Polyprotic acids can donate more than one proton, e.g. carbonic acid.
H₂CO₃ + H₂O <--> HCO₃^- + H₃O⁺ (The first proton to be donated to a water molecule.)
HCO₃^- + H₂O <--> CO₃^2- + H₃O⁺ (The second proton to be donated to a water molecule.)

12.3.0.3 Strong acids and weak acids
A strong acid completely dissociates into ions, e.g. nitric acid has almost complete dissociation, 93%
HNO₃ (aq) + H₂O --> H₃O⁺(aq) + NO₃^- (aq) +
A weak acid only partly dissociates into ions, e.g. acetic acid.
CH₃COOH + H₂O <--> CH₃COO^- + H₃O⁺So describing acids and bases as strong or weak only refers to their reaction with water and has nothing to do with concentration or the number of moles in a volume. The strong acids are perchloric acid (HClO₄), hydrochloric acid (HCl), hydrobromic acid( HBr), hydroiodic acid( HI), nitric acid (HNO₃), and sulfuric acid (H₂SO4). Any other acid is a weak acid because it does not completely dissociate in water.

Strongest acid	HClO₄ + H₂O --> H₃O⁺ + ClO^4-	Weakest base
.	HCl + H₂O --> H₃O⁺ + Cl^-	.
.	H₃PO₄ + H₂O <--> H₃O⁺ + H₂PO₄^-	.
Weakest acid	CH₃COOH + H₂O <--> H₃O⁺ + CH₃COO^-	Strongest base
.	NH₄⁺ + H₂O <--> H₃O⁺ + NH₃	.
.	H₂O + H₂O <--> H₃O⁺ + OH^-	.

12.3.0.4 pH
Water can transfer a proton from one molecule to another, autionization.
2H₂O <--> H₃O⁺ + OH^-
and
H₂O <--> H⁺ + OH^-
The product of hydrogen ion concentration, [H+] and hydroxide ion concentration, [OH^-] = the constant, K_w
K_{w =}[H⁺] X [OH^-] = 1.00 X 10^-14So [H⁺] = 10^-7 and [OH^-] = 10^-7The hydrogen ion concentration is very small in pure water so the concentration is describes in terms of its negative logpH is the negative log of the hydrogen ion concentration, pH = -log[H⁺], so hydrogen ion concentration, [H⁺] = 10^-pH.
So acidic solutions have a high [H⁺] and low pH values . Basic solutions have low [H⁺] and high pH values.
A solution that is neither acidic nor basic, a neutral solution, has [H⁺] = [OH^-], so pH = 7. A more acid solution has pH approaching 1. A more basic solution has pH approaching 14.

12.3.0.5 Ionization reaction of carbonic acid
H₂O (l) <--> H⁺ (aq) + OH^- (aq)
2H⁺ (aq) + CO₃^2- (aq) <--> H₂CO₃ (aq) carbonic acid
CO₂ + H₂O <--> H₃O⁺ + HCO₃^-, K1 = 4.4 X 10^-7
HCO₃^- + H₂O <--> H₃O⁺ + CO₃^2-, Ka = 4.7 X 10^-1112.3.1 Taste of acids, solid acids in the home
BE CAREFUL! NEVER TASTE ACIDS IN THE LABORATORY!
See appendix: Citric acid | See appendix: Acetic acid
Do NOT taste these acids in the laboratory. Each acid has a sour taste that is a characteristic of acids.
Lemon juice contains the white crystalline citric acid. Vinegar contains ethanoic acid (acetic acid, CH₃COOH).
Moisten your finger with a very dilute solution of hydrochloric acid. Rub your fingers together and then lick them. Repeat the procedure with very dilute solutions of acetic acid and citric acid. Do not taste any other acids because they may damage living tissues.

12.3.2 Dilute acids with metals, hydrochloric acid
Reactions of acids with metals are exothermic. The higher the metal is in the reactivity series the greater the heat liberated. Dilute hydrochloric acid with zinc:
Zn (s) + 2HCl (l) --> H₂ (g) + ZnCl₂ (aq)
The order of activity of metals with acids is similar to the order of activity with water.
Evolution of hydrogen occurs

Metal	2M Hydrochloric acid	2M Sulfuric acid
Magnesium	very rapid	rapid
Aluminium	slight	none
Zinc	moderate	Slight
Iron	very slight	very slight
Tin	none	none
Lead	none	none
Copper	none	none

1. Use different cleaned metals, e.g. calcium pieces, iron nail, lead sinker, magnesium ribbon, copper wire, aluminium sheet and zinc granules. Rub them with emery paper to make surfaces clean of oxides.
Put each metal into a separate test-tube. Add 10 mL of 2 M hydrochloric acid to test-tubes. Observe the properties of any gas liberated and name it. Test it with moist pieces of red and of blue litmus paper, with a drop of limewater hanging from a glass rod and with a lighted splint. Compare the rate at which hydrogen gas evolves by noting the rate and size of the hydrogen gas bubbles from the reaction. Describe the rate of reaction as nil, very slow, slow, moderately fast, very fast, and whether energy, in the form of heat, is produced (exothermic) or absorbed (endothermic). List the acids in order of their activity towards metals and state whether the same gas was liberated during each reaction and whether a salt may be isolated when the acids react with a metal.
2. Make up a reactivity series by listing the elements in approximate order of their activity with respect to acids, from the most active to the least active. Compare the results with the table of the reactivity series of some metals. The order of activity of the metals used, from the most active to the least active, is: magnesium, aluminium, zinc, iron with lead and copper displaying no noticeable reaction. When reaction did occur, the gas liberated was hydrogen gas. The reactions of these acids with metals are exothermic. The order of activity of the acids is that dilute hydrochloric and dilute sulfuric acids are about equal in activity but that they are more reactive than acetic acid. The order of activity of the metals with respect to acids is similar to that with respect to water. Magnesium ribbon forms most rapid bubbles of hydrogen gas then zinc then iron. Tin forms few bubbles of hydrogen gas. Copper forms no bubbles of hydrogen gas. Lead forms some lead chloride precipitate on the surface of the lead. Aluminium develops a layer of aluminium oxide that obstructs further chemical reactions.
3. Note the properties of any gas that forms. Test the gas with moist litmus paper a lighted splint and a hanging drop of limewater on a glass rod.
4. Feel the test-tube to note whether heat energy is released or absorbed. The reactions of these acids with metals are exothermic. 4. List the elements in approximate order of their activity with respect to hydrochloric acid from the most active to the least active. The order of activity is: magnesium, aluminium, zinc, iron, lead (no noticeable reaction), copper (no noticeable reaction).

12.3.2.1 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
Dilute hydrochloric and dilute sulfuric acids are about equal in activity, but that they are more reactive than ethanoic acid (acetic acid). Note the slower production of hydrogen gas with the weak acetic acid.
The reaction with sulfuric acid forms insoluble sulfates on the surface of calcium and lead that obstructs or stops reactions. List the acids in order of their activity on metals.

12.3.3 Dilute acids with metals, sulfuric acid with iron
Add dilute sulfuric acid to steel wool in a test-tube. Test the gas that forms with a lighted taper. BE CAREFUL! THE GAS IS HYDROGEN GAS!
Heat the mixture in a beaker of hot water until all the steel wool has dissolved. Add more acid when necessary. Filter the hot solution then leave it to cool. Crystals form on cooling. If no crystals form, add alcohol because the salt is less soluble in it. Dry the green crystals of iron (II) sulfate-7-water between absorbent paper.
Fe (s) + H₂SO₄ (aq) --> H₂ (g) + FeSO₄ (aq)

12.3.3.1 Dilute acids with metals, sulfuric acid with aluminium
Heat dilute sulfuric acid with pieces of aluminium foil in a test-tube. Some effervescence occurs but sometimes not enough to test for hydrogen gas with a lighted taper. After heating for 5 minutes, decant the solution that contains aluminium sulfate into another test-tube and add ammonia solution. A white jelly-like precipitate of aluminium hydroxide forms.

12.3.3.2 Magnesium with sodium hydrogen sulfate
Add 3 cm of magnesium ribbon to 3 cm of sodium hydrogen sulfate solution in a test-tube. The metal reacts with the sulfuric acid in the solution. Describe what you see. Tests for hydrogen gas. Remove any magnesium which has not reacted from the solution, pour part of the liquid into an evaporating basin and leave for magnesium sulfate crystals to form.

12.3.3.3 Iron with sodium hydrogen sulfate
Add a finger width of iron filings to a finger width of sodium hydrogen sulfate solution in a test-tube. Heat the mixture to speed up the reaction. The metal reacts with the sulfuric acid in the solution. Describe what you see. Tests for hydrogen gas. Leave to stand until all bubbles have ceased to appear. Pour part of the liquid into an evaporating basin and leave for magnesium sulfate crystals to form.
Test the liquid with universal indicator paper (The indicator changes colour to red, orange, or yellow for acids and green, or violet for alkalis.) Pale green is the colour for neutral substances. Before testing, make the paper this colour by dipping it into neutral tap water for a few moments. The Universal indicator turns yellow indicating the presence of an acid. Filter the liquid, and pour part of the clear solution into the evaporating basin and leave for pale green crystals of iron sulfate to form (FeSO₄.7H₂O, green vitriol). A solution of a salt is not necessarily neutral because some salts, like iron sulfate, form acids when dissolved in water.

12.3.4 Dilute acids with non-metals, carbon, sulfur
Add a piece carbon and sulfur to dilute hydrochloric acid, dilute sulfuric acid and dilute ethanoic acid (acetic acid) in separate test-tubes. Heat the test-tubes. No reaction occurs. Non-metals do not react with dilute acids.

12.3.5 Dilute acids with basic oxides, metal oxides, copper (II) oxide
1. Heated dilute acids react with metal oxides to form a salt and water: Pour dilute sulfuric acid into a Pyrex test-tube and heat in a beaker of boiling water until the sulfuric acid is nearly boiling. BE CAREFUL!
Add pieces of copper (II) oxide one by one while stirring until some remains unreacted with the acid. Filter the undissolved copper oxide from the hot solution. Leave the filtrate in a watch glass to cool and form crystals. Blue crystals of copper (II) sulfate-5-water form with water. Remove the crystals and dry them by pressing between absorbent paper.
H₂SO₄ (aq) + CuO (s) --> CuSO₄ (aq) + H₂O (l)
acid + basic oxide ---> salt + water
2. Repeat the experiment with dilute nitric acid.
2HNO₃ (aq) + CuO (s) --> Cu(NO₃)₂ (aq) + H₂O (l)

12.3.5.01 Copper sulfate from copper oxide
Add half a test-tube of sodium hydrogen sulfate solution to copper oxide in a test-tube. Heat the solution slowly until it turns blue. Be careful of spurting from the test-tube. Some copper oxide may remain after the reaction. Filter the solution obtain the filtrate of copper sulfate solution.

12.3.5.1 Dilute acids with basic oxides, (metal oxides) zinc oxide
Oxides of Sn, Al, Zn, Pb, and Sb are amphoteric. Amphoteric oxides react with bases to form a salt + water. Amphoteric oxides react with acids to form a salt + water.
Add dilute hydrochloric acid to zinc oxide.
2HCl (aq) + ZnO (s) --> ZnCl₂ (aq) + H₂O (l)
2NaOH (aq) + ZnO (s) --> Na₂ZnO₂ (aq) + H₂O (l)

12.3.6 Dilute acids with hydroxides, magnesium hydroxide
Basic hydroxides are insoluble in water and react with acids to form a salt and water. Many metallic hydroxides react with acids to form a salt and water.
Add magnesium hydroxide to dilute sulfuric acid until the reaction stops. Filter the mixture. Test the filtrate with litmus paper. Evaporate the filtrate to dryness so that crystals form.
Mg(OH)₂ (s) + H₂SO₄ (aq) --> MgSO₄ (aq) + H₂O (l)

12.3.7 Dilute acids with hydroxides, sodium hydroxide
Acids react with (neutralize) alkalis to form a salt and water.
Pour 5 mL of dilute sodium hydroxide solution into a watch glass. Test with litmus paper. Red litmus turns blue. Add dilute hydrochloric acid drop by drop. Stir as each drop is added. Test the mixture with the litmus paper until the litmus paper is neither red nor blue, but between these colours. Evaporate the solution to dryness by heating the watch glass over a beaker of boiling water. Crystals of sodium chloride (common salt) form.

12.3.7.1 Dilute acids with sodium hydroxide
Repeat the previous experiment with: dilute sulfuric acid, dilute nitric acid, ethanoic acid (acetic acid).
HCl (aq) + NaOH (aq) --> NaCl (aq) + H₂O (l)
hydrochloric acid + sodium hydroxide --> sodium chloride + water

12.3.7.2 Dilute hydrochloric acid with hydroxides
[NH₃ (aq) is used because while "NH₄⁺" ions and "OH^-" ions can be detected, "NH₄OH" cannot be detected, so ammonia solution is shown as "NH₃ (aq) + H₂O (l)"]
Repeat the experiment with dilute solutions of: potassium hydroxide, calcium hydroxide, aqueous ammonia solution.
acid + (base) alkali --> salt + water
HCl (aq) + NaOH (aq) --> NaCl (aq) + H₂O (l)
HNO₃ (aq) + NaOH (aq) --> NaNO₃ (aq) + H₂O (l)
HCl (aq) + KOH (aq) --> KCl (aq) + H₂O (l)
HCl (aq) + NH₃ (aq) + H₂O (l) --> NH₄Cl (aq) + H₂O (l)

12.3.8 Dilute acids with acidic oxides (non-metal oxides), carbon dioxide, sulfur dioxide
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Note any reaction for five minutes then evaporate to dryness. In each case, no reaction occurs. In each experiment there is no precipitate. If you evaporate a sample of a remaining solution to dryness in a fume cupboard, no residue remains. Pass carbon dioxide through hydrochloric acid or ethanoic acid (acetic acid) solution. Pass sulfur dioxide through hydrochloric acid or ethanoic acid (acetic acid) solution.

12.3.9.0 Dilute acids with carbonates, common carbonates
Dilute acids react with metal carbonates to form a salt, carbon dioxide and water. Geologists use this reaction to identify calcium carbonate in rock. Drops of hydrochloric acid cause bubbles to form.
1. Make a chemical egg peeler. Put an egg in vinegar (contains acetic acid, ethanoic acid). Note the bubbles forming on the outside of the egg. Leave overnight then, the next day, pick up the egg with your fingers. The egg has become soft. Leave to stand for a few days and the egg shell disappears completely. You can now see through the raw egg.
The acetic acid in the vinegar + calcium carbonate in the egg shell --> calcium acetate in solution + bubbles of carbon dioxide + water
2. Add 5 mL vinegar or dilute hydrochloric acid or dilute sulfuric or dilute nitric acid to pea size amounts of finely divided common carbonates: sodium hydrogen carbonate, sodium carbonate, calcium carbonate, magnesium carbonate, nickel carbonate, limestone, lime, oyster shells, egg shell, snail shell, coral. Continue to add the solid until no further reaction occurs. Filter and evaporate the filtrate to dryness. Note any visible changes. Test any gas liberated by inserting in the mouth of the tube first damp pieces of red and of blue litmus paper then a drop of limewater hanging on the tip of a glass rod and finally a burning splinter. In each case the gas is carbon dioxide.

12.3.9.1 Dilute hydrochloric acid with calcium carbonate
See diagram 3.34.1: Limewater tests for carbon dioxide
1. Put calcium carbonate in a test-tube. Add 2 mL 1.0 M hydrochloric acid. Tilt the test-tube so that its mouth is touching a second test-tube containing 5 mL of limewater. The surface of the limewater turns milky. Shake the test-tube containing the limewater. The milky colour on the surface disappears.
CaCO₃ (s) + 2HCl (aq) --> CO₂ (g) + CaCl₂ (aq) + H₂O (l)
carbonate + acid --> carbon dioxide + salt + water
2. Cover a chicken bone with vinegar or dilute acid hydrochloric acid. Change the solutions each day for 2 days. You can now tie a knot in the bone!
3. Cover a fresh egg with vinegar or dilute acid hydrochloric acid. Change the solutions each day for 7 days. Pick up the decalcified egg and drop it to show that it will bounce and not beak.

12.3.9.2 Dilute hydrochloric acid with sodium carbonate
1. Put sodium carbonate in a test-tube and add drops of dilute hydrochloric acid.
Test any gases formed from the reaction with moist litmus paper, a lighted splint, and a drop of limewater on a glass rod. The reaction forms carbon dioxide. Add more carbonate until no more reaction occurs. Filter and evaporate the filtrate to dryness. Repeat the experiment with dilute nitric acid. Repeat the experiment with magnesium carbonate.
Na₂CO₃ (s) + 2HCl (aq) --> 2NaCl (aq) + H₂O (l) + CO₂ (g)
Na₂CO₃ (s) + 2HNO₃ (aq) --> 2NaNO₃ (aq) + H₂O (l) + CO₂ (g)
2. Shake different solid acids in separate test-tubes half filled with water. Divide the solutions in the test-tubes into three different test-tubes. Test-tube A: Add small pieces of red and of blue litmus paper. Test-tube B: Add three drops of methyl orange solution. Test-tube C: Add three drops of phenolphthalein solution. Observe any changes in the solutions. Add solid sodium carbonate to each acid solution. Observe any changes in the solutions. Pass some gas given off into a test-tube containing limewater. Shake the test-tube for thorough mixing. Note how milky the solution is because carbon dioxide was produced when the acids reacted with sodium carbonate.

12.3.9.3 Dilute tartaric acid with sodium carbonate
Put 5 g of sodium carbonate and the same quantity of tartaric acid in a test-tube fitted with a one-hole stopper and delivery tube With the end of the delivery tube dipping into a second test-tube of limewater add water to the first test-tube and quickly replace the stopper. Describe what you see. The limewater turns milky. The acid reacts with sodium carbonate to form a salt, carbon dioxide, and water.
tartaric acid + sodium carbonate --> sodium tartrate + carbon dioxide + water

12.3.9.4 Dilute tartaric acid with egg shell, soil, wood ash
Many common substances, such as mortar, egg shell, most soils, contain calcium carbonate. and wood ashes contain potassium carbonate. Observe the action of tartaric acid on these substances in a test-tube. Tests for carbon dioxide by holding a drop of limewater, at the end of a glass tube, in the mouth of the test-tube.

12.3.9.5 Dilute hydrochloric acid with calcium carbonate
Put 5 g of marble chips (calcium carbonate) and the same quantity of dilute hydrochloric acid in a test-tube fitted with a one-hole stopper and delivery tube With the end of the delivery tube dipping into a second test-tube of limewater add water to the first test-tube and quickly replace the stopper. Describe what you see. The limewater turns milky. The acid reacts with calcium carbonate to form a salt, carbon dioxide, and water.
hydrochloric acid + calcium carbonate --> calcium chloride + carbon dioxide + water

12.3.9.6 Dilute sulfuric acid with calcium carbonate
Put 5 g of marble chips (calcium carbonate) and the same quantity of dilute sulfuric acid in a test-tube fitted with a one-hole stopper and delivery tube With the end of the delivery tube dipping into a second test-tube of limewater add water to the first test-tube and quickly replace the stopper. Describe what you see. The limewater turns milky. The acid reacts with calcium carbonate to form a salt, carbon dioxide, and water. The reaction of sulfuric acid with calcium carbonate proceeds only for a few moments because the salt formed, calcium sulfate, is only slightly soluble and deposits on the carbonate, preventing this compound from reacting with the acid. So the reaction with hydrochloric acid above is much better.
sulfuric acid + calcium carbonate --> calcium sulfate + carbon dioxide + water

12.3.10.0 Dilute acids with sodium hydrogen carbonate
See 3.34.6: Soda-acid fire extinguisher
Order online: Mighty Seltzer Rocket, uses Alka-Seltzer tablets, Be careful!
The only stable hydrogen carbonates are KHCO₃and NaHCO₃. Sodium hydrogen carbonate, bicarbonate of soda, is used in baking soda, baking powder, self raising flour, effervescent fruit salts, e.g. Alka-Seltzer, and soda acid fire extinguishers and treatment for acid burns. Some people swallow sodium hydrogen carbonate to counteract excess acid in the stomach but using magnesium oxide or magnesium hydroxide that does not react with acids to produce carbon dioxide is better.
1. Add sodium hydrogen carbonate, or other hydrogen carbonates, to acids to form carbon dioxide, water and a salt.
NaHCO₃ + HCl --> CO₂+ H₂O + NaCl
hydrogen carbonate + acid --> carbon dioxide + water + salt
2. Mix vinegar with bicarbonate of soda in a glass jar. Drop some naphthalene mothballs into the solution. The carbon dioxide formed by the reaction of the vinegar (acetic acid) with the sodium hydrogen carbonate forms bubbles of carbon dioxide on the mothballs in the bottom of the jars. The mothballs rise to the surface, lose the bubbles and sink again.
NaHCO₃ + CH₃COOH --> CH₃COONa + H₂O + CO₂ (g)

12.3.10.1 Dilute acids with calcium hydrogen carbonate
Put powdered calcium carbonate into a test-tube containing about 10 mL of water. Bubble carbon dioxide through the suspension until no further change takes place. Soluble calcium hydrogen carbonate forms. Boil the mixture for 10 minutes. Add acids to form carbon dioxide, water and a salt.

12.3.11.0 Dilute nitric acid with copper
Very dilute nitric acid may react with very active metals, e.g. magnesium to form hydrogen gas. When nitric acid reacts with most metals, it oxidizes the hydrogen to water.
Add drops of dilute nitric acid to copper. Nitrogen monoxide forms which immediately reacts with oxygen gas in the air to form nitrogen dioxide.
3Cu (s) + 8HNO₃ (aq) --> 3Cu(NO₃)₂ (aq) + 4H₂O (l) + 2NO (g)
2NO (g) + O₂ (g) --> 2NO₂ (g)

12.3.11.1 Nitric acid with metals
Add slowly small pieces of copper, magnesium and zinc to small amounts of dilute nitric acid in separate test-tubes. If no change is taking place, gently heat the mixture. Now repeat the procedure 1. with concentrated nitric acid 2. with concentrated sulfuric acid and 3. with concentrated hydrochloric acid. The reactions of metals with nitric acid and concentrated sulfuric acid are different from reactions of metals with hydrochloric acid, dilute sulfuric acid and dilute acetic acid. Although copper does not react with dilute acids or with concentrated hydrochloric acid, it does react with both dilute and concentrated nitric acids and with hot concentrated sulfuric acid but does not produce hydrogen gas in reaction with them. The residual mixtures contain solutions of salts but writing equations for the reactions is difficult because more than one reaction can occur simultaneously between copper or magnesium or zinc and nitric acid. For example when zinc reacts with nitric acid the reaction may produce nitrogen dioxide, nitric oxide, nitrous oxide, zinc nitrate and ammonium nitrate.

12.3.12 Concentrated acids with metals, nitric acid with copper
Nitric acid reacts with metals above platinum in the reactivity series, but does not form hydrogen gas.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Pour drops of concentrated nitric acid on pieces of copper in a test-tube. Put a stopper on the test-tube immediately because brown nitrogen dioxide gas forms. The nitric acid acts as an oxidizing agent and is reduced to nitrogen dioxide and water. The reaction is exothermic.
Cu (s) + 4HNO₃ (aq) --> Cu(NO₃)₂ (aq) + 2H₂O (l) + 2NO₂ (g)

12.3.13 Concentrated acids with metals, sulfuric acid with copper
Concentrated acids should be handled only by experienced science teachers. Concentrated sulfuric acid reacts with metals above platinum in the reactivity series, but does not form hydrogen gas.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Add hot concentrated sulfuric acid to a piece of copper foil. Brown nitrogen dioxide gas forms. The sulfuric acid acts as an oxidizing agent.
Cu (s) + 2H₂SO₄ (aq) --> CuSO₄ (aq) + 2H₂O (l) + SO₂ (g)

12.3.14 Concentrated acids with non-metals, carbon, sulfur
DO NOT DEMONSTRATE THIS EXPERIMENT!
Hot sulfuric acid and nitric acid can react as oxidizing agents with carbon and sulfur. Carbon is oxidized to carbon dioxide and nitric acid is reduced to nitrogen dioxide and water.
C (s) + 4HNO₃ (aq) --> CO₂ (g) + 4NO₂ (g) + 2H₂O (l)

12.3.15 Acids with salts
1. Add small quantities of sodium chloride, sodium nitrate, sodium acetate, sodium sulfite and iron sulfide to about 5 mL of dilute hydrochloric acid in separate test-tubes. Observe what happens when the mixtures are cold and when they are warmed.
2. Repeat the procedure using dilute sulfuric acid and then concentrated sulfuric acid.
3. Dilute acids do not react with chlorides, nitrates, sulfates, or acetates unless the metal ions in the salt can form an insoluble salt with the ions in the acid. 4. Acids react with sulfites to produce sulfur dioxide, water and a salt.
5. Acids react with sulfides to produce hydrogen sulfide (rotten egg gas) and a salt.
6. Concentrated sulfuric acid reacts with chlorides to produce hydrogen chloride and a sulfate.
7. Concentrated sulfuric acid reacts with nitrates to produce nitric acid and a sulfate.
8. Concentrated sulfuric acid reacts with acetates to produce acetic acid and a sulfate.

12.4.0 Hydrochloric acid
Hydrochloric acid is an aqueous solution of hydrogen chloride gas. Hydrochloric acid dissolves most metals to form chlorides and hydrogen gas. Hydrochloric acid is available as: 1. 5.0 M, 4.0 M, 2.0 M, 1.0 M and 0.5 M volumetric solutions 2. minimum assay 36% solution density 1.17 g cm^-3 at 20^oC 3. 36% "ANALAR" solution (d) commercial solution called muriatic acid for use in the building trades.

12.5.0 Nitric acid, aqua fortis, spirits of nitre
Nitric acid (concentrated), HNO₃, is a yellow fuming corrosive liquid that dissociates into ions even in concentrated solution. Concentrated nitric acid is a strong oxidizing agent. Dilute nitric acid is a strong acid that reacts with metals as an acid or an oxidizing agent. Nitric acid is a clear, colourless, fuming, unstable, corrosive, poisonous liquid). For nitric acid, use only glass equipment, and attach corks and rubber tubing rapidly. Nitric acid is both a strong acid and a powerful oxidizing agent. It produces severe acid burns, and can explode when mixed with many combustible materials. The vapour of concentrated nitric acid is strongly irritating to the lungs. Spillage of concentrated nitric acid outside a fume cupboard will require evacuation of the area until fumes dissipate. Use sodium bicarbonate to neutralize spills. Avoid skin contact with nitric acid. Prolonged contact with nitric acid causes severe chemical burns. Nitric acid turns the skin yellow because of the nitration of aromatic amino acids, tyrosine, tryptophan and phenylalanine, in the protein of the skin. After a few days, the yellow layer peels off. Wear a face shield when doing any experiments, e.g. pouring, which could result in splashing of the liquid. Concentrated nitric acid can explode when mixed with methylated spirit, ethanol, methanol and other combustible materials. The reaction of nitric acid with alcohols is especially violent. Concentrated nitric acid boils and decomposes at 83^oC. The product of the reaction between concentrated nitric acid and cellulose (cotton, paper, wood) is very explosive. Store concentrated nitric acid in the acid bay of the store room. Dilute nitric acid may be stored safely in preparation rooms. Do not store it near acetic acid. The molar mass of HNO₃ is 63.01 g / mol.

12.5.0a Grades of nitric acid
1. Anhydrous nitric acid
100% nitric acid, density 1.513 g / mL, approx. 24 M, colourless liquid, 1.5222 g / cm³, b.p. 83^oC

2. White fuming nitric acid, WFNA, (called "100% nitric acid")
Concentration close to anhydrous nitric acid, < 2% water, < 0.5% dissolved NO₂. It does not contain N₂O₄.
(Red fuming nitric acid, RFNA, contains 13% dinitrogen tetroxide, N₂O₄, and 3% water. It was used as a rocket propellant.)

3. Fuming nitric acid, > 95%
Relative density 1.5, b.p. 83^oC, miscible with water, strong oxidizing agent, yellow colour from the atmosphere, Light causes photolytic decomposition of the colourless pure acid to form dinitrogen tetroxide that colours the acid yellow-brown. As the acid darkens in light it may be stored in dark glass bottles.
> 86% are commonly called "fuming nitric acid"
Boiling nitric acid
4HNO₃ --> 2H₂O + 4NO₂ + O₂ (72^oC)
NO₂ dissolves in the nitric acid to form fuming nitric acid, 16M, the most concentrated grade at STP

4. Reagent grades, 70% reagent grades, e.g. "70% ANALAR" solution"
< 16 M, oxidizing acid, minimum assay 69%, relative density 1.41
Concentrated acid 70% by weight is very hygroscopic. Use it to reacts with copper to form brown nitrogen dioxide.
If concentrated nitric acid of commerce has a relative density of 1.41 and mass per cent HNO₃ = 69%
Compute the molal concentration of HNO₃
Mass of one litre of concentrated nitric acid = 1410 g, of which (0.69)(1410 g) = 972.9 g HNO₃ / L, the balance is water.
So 972.9/63.01 = 15.44 mol / L is the molar concentration

5. Commercial grade concentrated nitric acid, (Technical grade), 68%
Approx 15 molar, relative density 1.42, the maximum boiling azeotrope (constant boiling mixture), 68% nitric acid, 32% water, at 120.5^oC
Strong oxidizing agent, so metals dissolve in it and nitrogen oxides are formed. Nonmetals react to form the oxo acid or oxide.

6. Dilute nitric acid, < 4 M (24%)
2.0 M and 1.0 M volumetric solutions
Use 1% (> 2M) acid to demonstrate metal activity and acid on carbonates. Prepare nitric acid by action of sulfuric acid on sodium nitrate.
Prepare dilute acids
Use safety goggles and nitrile chemical-resistant gloves.
If commercial nitric acid (70% w/w) has a molarity = 15.4 moles / L
1 litre contains 15.4 moles, so 1 mole is contained within = 1 litre / 15.4 moles.
1 mole = 0.0649 Litres = 64.9 mL
Mass of 64.9 mL of commercial nitric acid:
= 64.9 mL x 1.41g / mL
= 91.5 g
where 1.41g / mL is the density of the commercial nitric acid, 70% w/w
Add 64.9 mL, or 91.5 g, of commercial nitric acid to water, make up to 1L, to prepare 1 M solution of nitric acid.
Nitric acid 2 M, Dilute 129.8 mL of 15.4 M acid (70% w/w, 1.41 g / mL) to 1 litre of water
Nitric acid 3 M: Dilute 194.7 mL of 15.4 M acid (70% w/w, 1.41 g / mL) to 1 litre of water

7. Dilute nitric acid < 0.2 M (1%) Not hazardous
Commercial solution for use by tradesman are very weak solutions.

12.5.01 Ionization of nitrous acid, Ka = 4.5 X 10^-4
HNO₂ + H₂O <--> H₃O⁺ + NO₂^-Nitrosamines, produced by nitrous acid with secondary amines, can be formed in the gut when nitrites react with amino acids.

12.5.1 Prepare nitric acid, sulfuric acid with sodium nitrate
Add concentrated sulfuric acid to sodium nitrate.
Be Careful! Heat gently, nitric acid vapour forms.
NaNO₃ (s) + H₂SO₄ (l) --> HNO₃ (l) + NaHSO₄ (s)
Ionization of hydrogen sulfate ion
HSO₄^- + H₂O <--> H₃O⁺ + SO₄^2-

12.5.2 Nitric acid with copper
Nitric acid dissolves most metals, usually producing copious brown fumes of toxic nitrogen dioxide. However, the composition of the gases produced depends on the concentration of the nitric acid. Reaction of copper with dilute nitric acid produces colourless nitric oxide, NO, that reacts with air at the top of a test-tube to form brown nitrogen dioxide, NO₂. With concentrated nitric acid, copper forms mostly brown NO₂ gas. Nitrogen dioxide is a powerful lung irritant and all reactions which result in the formation of nitrogen dioxide should be done in a fume cupboard or a well-ventilated area.

12.5.3 Nitric acid with sulfuric acid
The stronger conc. sulfuric acid donates a proton, H⁺, to the weaker conc. nitric acid, which then loses a water molecule leaving the electrophoile NO₂⁺.
HNO₃ + H₂SO₄ --> NO₂⁺ + H₂O +2HSO₄^-

12.7.01 Nitration
Substitution of -NO2 group for a hydrogen atom in an arene ring, e.g. bromine

12.7.1 Prepare nitrous acid
Weak acid prepared by acids on nitrites.
barium nitrate + sulfuric acid --> barium sulfate + nitrous acid
Ba(NO₂)₂ + H₂SO₄ --> BaSO₄ + 2HNO₂sodium nitrite + dilute hydrochloric acid --> sodium chloride + nitrous acid
NaNO₃ (aq) + HCl (aq) --> NaCl (aq) + HNO₂ (aq)

12.7.2 Nitrous acid with amines
nitrous acid + amines --> ionic salt in aqueous solution
nitrous acid + amines --> diazonium ion + chloride ion
HNO₂ + R-NH₂ --> R-N⁺=-N + Cl^-
nitrous acid + phenylamine --> benzenediazonium chloride + Cl^-
HNO₂ + C₆H₅NH₂ --> C₆H₅N₂Cl + 2H₂O
Diazonium ion, C₆H₅N₂⁺, forms unstable salts. Alkyl diazonium ions are unstable and rapidly decompose but aryldiazonium ions are more stable if kept in solution below 10^oC, but as solids they are explosive.

12.7.3 Heat nitrous acid
Heated nitrous acid decomposes to form nitric acid and nitrogen monoxide (nitric oxide).
nitrous acid --> nitric acid + nitrogen monoxide2HNO₂ --> HNO₃ + NO

12.9.0 Phosphoric acid
Ionization reaction
H₃PO₄ + H₂O <--> H₃O⁺ + H₂PO₄^-
H₂PO₄^- + H₂O <--> H₃O⁺ + HPO₄^2-
HPO₄^2-+ H₂O <--> H₃O⁺ + PO₄^3-

12.10.0 Boric acid, ionization reaction
1. Orthoboric acid, trioxoboric acid (III) acid, boracic acid, sassolite, H₃BO₃ is a weak acid. White to colourless triclinic crystals, m.p. 169^oC, occurs in volcanic steam vents, slightly soluble in cold water, used to make borosilicate glass, used in buffer solutions, detergents and in pharmacy, e.g. "boracic powder" for eye infections.
Action of continuous heat: boric acid, H₃BO₃--> metaboric acid + water, H₂B₄O₄ --> tetraboric acid (pyroboric acid) H₂B₄O₇ --> boric oxide (anhydrous boron (III) oxide) B₂O₃.
Boric oxide is an intermediate oxide, as is aluminium oxide, with weak acidic and basic properties. Borax is hydrated sodium borate. When heated it fuses to form clear glass that can dissolve metal oxides to give characteristic colours of the borax bead test.
Ionization reaction, Ka = 6.0 X 10^-10
H₃BO₃+ H₂O <--> H₃O⁺ + H₂BO₃^-H₃BO₃<--> H⁺ + H₂BO₃^-H₂BO₃- <--> H⁺ + HBO₃^2-
HBO₃^2- <--> H⁺ + BO₃^2-

12.10.1 Prepare boric acid crystals
Use 5 g of boric acid crystals. Pour some into 2 cm boiling water in a test-tube and leave to dissolve. Continue adding crystals and heat to boiling until all crystals dissolved. Leave to cool to see fine white crystals form.