School Science Lessons
Topic 12G Reactions of metals and their salts
2012-05-20 SPPwp
Please send comments to: J.Elfick@uq.edu.au
Table of contents  
12.12.3 Acids with ammonium carbonate
12.1.18 Acids with baking soda
12.1.25 Acids with sodium carbonate (washing soda)
12.1.40 Acids with sodium thiosulfate
12.5.2 Acids with zinc
12.12.2 Alkalis with ammonium carbonate
12.5.3 Alkalis with zinc
12.1.6 Ammonium iron (III) sulfate, NH4Fe(SO4)2.12H2O, iron (III) ammonium sulfate, ferric ammonium sulfate (FAS), iron alum
12.8.2 Burn mothballs (naphthalene crystals)
12.6.5 Citric acid solubility and temperature
12.5.1 Combustion of zinc
12.1.41 Copper sulfate solution with sodium thiosulfate
12.11.2 Decolorize vinegar
12.12.1 Decompose ammonium carbonate
12.3.15 Dilute nitric acid forms carbon dioxide with carbonates and bicarbonates
12.3.13 Dilute nitric acid with metals
12.3.14 Dilute nitric acid with metal oxides
12.1.23 Efflorescence of sodium carbonate (washing soda)
12.8.1 Evaporate mothballs, (naphthalene crystals)
12.7.20 Ferment sugar with yeast
12.1.29 Flame tests for sodium chloride
12.7.5 Glycerine colour change reaction
12.7.8 Glycerine with cobalt chloride solution
12.7.7 Glycerine with sugar
12.6.2 Heat citric acid
12.1.17 Heat decomposes baking soda
12.7.1 Heat glucose
12.1.24 Heat sodium carbonate (washing soda)
12.1.39 Heat sodium thiosulfate, Na2S2O3.5H2O
12.7.12 Heat starch
12.7.16 Heat sucrose
12.1.1 Iron (II) sulfate with sodium carbonate
12.1.2 Iron (II) sulfate with ammonia
12.4.10 Magnesium sulfate with ammonia
12.4.11 Magnesium sulfate with sodium carbonate
12.7.6 Oxidation of glycerine
12.1.3 Oxidize iron (II) sulfate to iron (III) sulfate, (ferric sulfate)
12.12.4 Precipitates with ammonium carbonate
12.1.19 Precipitates with sodium bicarbonate
12.1.26 Precipitates with sodium carbonate (washing soda)
12.14.1 Prepare ammonium sulfate by neutralization
12.1.27 Prepare baking soda with sodium carbonate (washing soda)
12.1.28 Prepare bath salts with sodium carbonate and with sodium sesquicarbonate
12.7.9 Prepare casein with milk
12.1.36 Prepare chemical gardens
12.6.1 Prepare citric acid crystals with lemon juice
12.7.15 Prepare glucose with starch
12.7.18 Prepare glucose with sugar
3.42.01 Prepare hydrochloric acid
12.6.3 Prepare hydrogen with citric acid
12.1.5 Prepare iron (II) sulfate crystals with iron filings
12.1.7 Prepare iron (III) hydroxide and iron (III) oxide
12.7.10 Prepare lactic acid with milk
12.1.33 Prepare metal hydroxides by precipitation
12.8.3 Prepare naphthalene crystals with mothballs
12.1.11 Prepare oxygen gas with bleaching fluid
12.1.8 Prepare salts by neutralization of sodium hydroxide
12.1.37 Prepare silicic acid and pure silica
12.1.22 Prepare sodium carbonate (washing soda) crystals with baking soda
12.1.34 Prepare sodium carbonate (washing soda) crystals with caustic soda
12.1.13 Prepare sodium hypochlorite with bleaching fluid and sodium carbonate
12.1.32 Prepare sodium chloride cubic crystals
12.6.4 Prepare sodium citrate crystals
12.2.1.9 Prepare sodium hydroxide solution by double decomposition
12.1.16 Prepare sodium nitrate crystals by neutralization
12.1.35 Prepare sodium silicate precipitates
12.1.38 Prepare sodium sulfate supersaturated solution, Na2SO4.10H2O
12.18.6.1 Prepare sodium thiosulfate crystals, "hypo"
12.1.42 Prepare sodium thiosulfate supersaturated solution
12.2.4 Prepare sulfuric acid with iron (II) sulfate
12.11.1 Prepare verdigris with copper and vinegar
12.2.8 Reduce iron (III) sulfate to iron (II) sulfate
12.1.21 Sodium bicarbonate prevents milk from going sour
12.7.19 Sucrose with borax
12.7.17 Sucrose with sodium hydrogen sulfate, NaHSO4, (sodium bisulfate)
12.7.13 Starch with water
12.10.1 Tea with iron (II) sulfate
12.10.2 Tea with iron (III) salts
12.10.3 Tea with dilute sulfuric acid
12.10.4 Tea with limewater
12.11.3 Tests for acetic acid in vinegar
12.13.1 Tests for aluminium compounds
12.1.9 Tests for aspirin
12.7.2 Tests for glucose and sucrose
12.7.3 Tests for glucose in apples and sweets
12.7.4 Tests for glycerine
12.7.11 Tests for lactic acid solution
12.4.12 Tests for magnesium in compounds
12.1.15 Tests for soapy feel of sodium hydroxide (caustic soda)
12.1.20 Tests for sodium bicarbonate in a stomach powder
12.7.14 Tests for starch in adhesive paste
12.1.10 Use sodium hydrochlorite bleach
12.1.1 Iron (II) sulfate with sodium carbonate
Add sodium carbonate, (washing soda), solution to iron (II) sulfate solution. A green precipitate of iron (II) carbonate forms
iron (II) sulfate + sodium carbonate –> iron (II) carbonate + sodium sulfate
12.1.2 Iron (II) sulfate with ammonia
Add drops of dilute ammonia solution to 2 cm of iron (II) sulfate solution in a test-tube and shake the test-tube. Observe a grey-green precipitate of iron (II) hydroxide. The precipitate left on the side of the test-tube quickly turns brown, because oxygen from the air turns it into iron (III) hydroxide, (ferric hydroxide).
12.1.3 Oxidize iron (II) sulfate to iron (III) sulfate, (ferric sulfate)
Heat iron (II) sulfate solution with a substance rich in oxygen, e.g. hydrogen peroxide. Boil 2 cm of iron (II) sulfate solution in a test-tube with drops of hydrogen peroxide. The green colour changes to yellow or brown. Cool the test-tube under the tap and test the liquid, iron (III) sulfate solution (ferric sulfate) by adding dilute ammonia solution. A brown precipitate of iron (III) hydroxide, (ferric hydroxide) forms.
12.1.5 Prepare iron (II) sulfate crystals with iron filings
When iron is treated with dilute sulfuric acid hydrogen gas forms and a solution of iron (II) sulfate (iron (II) sulfate) forms. Heat half a test-tube of dilute sulfuric acid over a flame, but do not boil the liquid. Remove the test-tube from the flame and add iron filings on the end of a metal spatula. A vigorous effervescence occurs because of the formation of hydrogen. Test with a glowing splint. When the action dies down, add more iron filings, and then more, until a total of 4 mL is added. Put the test-tube on one side for 15 minutes until effervescence has nearly ceased. Filter the liquid into an evaporating basin. Make sure that acid is left, by testing with blue litmus paper. The presence of acid prevents the solution from oxidizing to brown iron (III) sulfate. To obtain large crystals of iron (II) sulfate, leave the solution undisturbed for a day or two to crystallize out. Small crystals can form more quickly by evaporating the
solution over a flame until only one third of it remains. If a brown colour appears in the liquid during the evaporation add a drop or two of dilute sulfuric acid. When the evaporation finishes, leave the remaining liquid to cool. Many small crystals of iron (II) sulfate are deposited.
12.1.6 Ammonium iron (III) sulfate, NH4Fe(SO4)2.12H2O, iron (III) ammonium sulfate, ferric ammonium sulfate (FAS), iron alum, Toxic if ingested
Ammonium iron (III) sulfate forms violet crystals that dissolve in water to form a brown acid solution. Prepare a solution by dissolving 2 mL of the powdered crystals in two thirds of a test-tube of water. Be careful! Do not to spill the solution on clothes because it causes a brown stain of iron mould, followed by rotting of the cloth.
12.1.7 Prepare iron (III) hydroxide and iron (III) oxide
Add 4 cm of dilute ammonia solution to 2 cm of ammonium iron (III) sulfate solution in a test-tube. Shake the test-tube to mix the liquids. A brown jelly-like precipitate of iron (III) hydroxide forms. Filter off the precipitate of iron (III) hydroxide. Put of the red-brown jelly left in the filter paper on to a clean metal lid or into a metal screw-cap. Hold the lid or cap in a pair of pliers and heat it carefully above a flame. Steam is produced and a red powder formed. This is iron (III) oxide, a very pure form of rust. The same
substance forms by heating iron (II) sulfate crystals.
12.1.8 Prepare salts by neutralization of sodium hydroxide
As sodium hydroxide is an alkali it can be used to neutralize an acid. The substance resulting from neutralization is a salt. Chemical reaction: Alkali + acid –> salt + water For this experiment a solution of citric acid, dilute nitric acid, or dilute sulfuric acid can be used. The salt obtained is sodium citrate, sodium nitrate, or sodium sulfate. The details of the preparation are similar to those described to make ammonium sulfate.
12.1.9 Tests for aspirin
Crush half an aspirin tablet and dissolve the powder by heating it with sodium carbonate, (washing soda), solution in a test-tube. Cool the test-tube and make the liquid slightly acid by adding dilute sulfuric acid or sodium hydrogen sulfate, (sodium bisulfate), solution. Add drops of ammonium iron (III) sulfate solution. The liquid turns a mauve or violet colour.
12.1.10 Use sodium hydrochlorite bleach
1. Wet the cork of the bottle of bleaching fluid with the strong solution and touch a piece of red litmus paper. The red litmus paper first turns blue and then becomes bleached. Put one drop of blue ink in a test-tube then half fill it with water. Add one drop of bleaching fluid then shake the test-tube. The colour from the ink disappears.
2. Wet a strip of coloured cotton cloth then squeeze it nearly dry. Put the cloth in a beaker then add half a test-tube of the strong bleaching fluid and two test-tubes of water. Leave the cloth in the liquid for twenty-four hours. If by then the colour has not completely disappeared, renew the bleaching solution and put the cloth in it again. The cloth is rotted if too strong a bleaching solution is used.
12.1.11 Prepare oxygen gas with bleaching fluid
Put 2 cm of bleaching fluid in a test-tube and add a crystal of cobalt chloride. The contents of the test-tube turn black. Heat the test-tube over a flame. Effervescence begins and oxygen forms. Test for oxygen by putting a glowing wood splint into the mouth of the test-tube.
12.1.13 Prepare sodium chlorate (I) (sodium hypochlorite) with bleaching fluid and sodium carbonate (washing soda)
Prepare a solution of sodium carbonate, (washing soda), by dissolving 10 mL of the powdered crystals in 100 mL of water in a beaker. Leave the solution to cool. Put 10 mL of bleaching powder into a beaker and add the cold solution of washing soda while stirring with a glass rod. When all the washing soda is added leave the white precipitate to settle. Pour off the liquid above the solid then filter it. The filtrate is a solution of sodium chloride and sodium hypochlorite.
12.1.14 Prepare sodium hydroxide solution with calcium hydroxide and sodium carbonate
Fill a beaker half full of water and heat it on a gauze and tripod until the water boils. Add 15 mL of powdered sodium carbonate (washing soda) and stir with a glass rod until it has dissolved. Add 7 mL calcium hydroxide (slaked lime) crystals and continue the boiling. After five minutes, filter drops of the hot liquid into a test-tube then test the filtrate by adding drops of an acid solution. If no effervescence occurs, the reaction is complete. If effervescence occurs, add another 5 mL of calcium hydroxide to the beaker and continue the boiling for another five minutes. Tests the filtrate again by adding drops of an acid solution. When the testing shows no sodium carbonate in the filtrate leave the beaker to cool so that a white residue of chalk settles on the bottom of the beaker. When the beaker is cool, pour off the clear solution of sodium hydroxide into a stock bottle, but do not filter it. Add an equal amount of water to this strong solution which could be used for other experiments.
Double decomposition reaction:
calcium hydroxide + sodium carbonate --> calcium carbonate + sodium hydroxide.
12.1.15 Tests soapy feel of sodium hydroxide (caustic soda)
Moisten the tip of a finger with a drop of dilute sodium hydroxide solution and rub the fingers together. The soapy feel is a typical property of strong alkalis. Wash the fingers immediately after the test.
12.1.16 Prepare sodium nitrate crystals by neutralization
Put half a test-tube of dilute nitric acid into an evaporating basin and add a small piece of red litmus paper. Add dilute sodium carbonate or sodium hydroxide solution, a drop or two at a time. Stir the solution after each addition. Continue adding the solution and stirring until the litmus paper turns blue. Remove the litmus paper and heat the evaporating basin gently to boil away about two thirds of the liquid before leaving it to stand. White crystals of the salt sodium nitrate form.
Acid + base --> salt + water.
12.1.17 Heat decomposes baking soda
Put 5 g of sodium bicarbonate into a dry test-tube 1. Put 2 cm of clear limewater in test-tube 2. Hold test-tube 1 in a paper holder in the left hand so that the test-tube slopes down. Hold test-tube 2 in the other hand so that the two test-tubes are mouth to mouth. Heat test-tube 1 over a very flame. Observe the moisture depositing on the cooler part of the test-tube. After heating for a few minutes shake up the limewater in test-tube 2. Th limewater turns milky, showing that carbon dioxide is being produced. Chemical action: sodium bicarbonate + heat --> sodium carbonate + water + carbon dioxide
12.1.18 Acids with baking soda
Any acid solution causes carbon dioxide to be formed sodium bicarbonate. Put a 5 mL of baking soda in a test-tube and add vinegar or any other dilute acid. Dip a lighted taper into the test-tube and it is extinguished by the carbon dioxide formed.
12.1.19 Precipitates with sodium bicarbonate
A solution of sodium bicarbonate forms precipitates with metal salts. These precipitates are not bicarbonates but the ordinary carbonates of the metals. Put 2 cm of baking soda solution into a test-tubes and adding drops of the following solutions: copper sulfate, iron (II) sulfate, magnesium sulfate, zinc sulfate.
12.1.20 Tests for sodium bicarbonate in a stomach powder
This mixture may contain calcium carbonate, magnesium carbonate, bismuth carbonate, and sodium bicarbonate, but only sodium bicarbonate is soluble. Shake 5 mL of the powder with water in a test-tube for a few minutes. Filter the milky liquid. The filtrate is a colourless solution of sodium bicarbonate. Dip the end of a wood splint into the liquid and hold it in the edge of a non-luminous Bunsen burner flame. The flame turns an intense yellow colour, showing the presence of sodium. Heat the rest of the solution. Carbon dioxide forms as shown by testing with limewater.
12.1.21 Sodium bicarbonate prevents milk from going sour
In hot weather milk may turn sour because bacteria change the milk sugar into lactic acid. So milk that is beginning to “go off” can be saved by adding sodium bicarbonate solution to it.
12.1.22 Prepare sodium carbonate (washing soda) crystals with baking soda
Put 10 g of baking soda crystals in a beaker about one third full of water. Stir the powder round then pour off a few drops of the liquid into a test-tube and test it with phenolphthalein. Sodium bicarbonate solution is too weak an alkali to give the usual red colour with this indicator. Heat the beaker on a gauze and tripod and keep the liquid gently boiling for ten minutes. The sodium bicarbonate crystals soon disappear because they decompose into the more soluble sodium carbonate, water, and carbon dioxide. Let the beaker stand to cool. Tests drops of the liquid again with a drop of phenolphthalein. A rose red colour forms showing that sodium carbonate is a stronger alkali than sodium bicarbonate. Colourless crystals of sodium carbonate may deposit in the beaker when cold. Alternatively, transfer the solution to an evaporating basin to heat the solution and increase the concentration, then leave to cool.
12.1.23 Efflorescence of sodium carbonate (washing soda)
Put large, clear crystal of sodium carbonate on a watch glass and leave it for an hour. It forms a white powdery coating that can easily be scraped off with a knife. This change, called “efflorescence,” is caused by the crystal losing its water of crystallization to the air. The white powder has the chemical formula Na2CO3. H2O, can be called sodium carbonate monohydrate.
12.1.24 Heat sodium carbonate (washing soda)
Heat a crystal of sodium carbonate, gently on a metal lid. The crystal soon melts. Continue heating to form clouds of steam. The liquid dries up and a white powder remains. The anhydrous sodium carbonate, known commercially as soda ash, has lost all of its water of crystallization.
12.1.25 Acids with sodium carbonate (washing soda)
Sodium carbonate with acids produces effervescence because of formation of carbon dioxide, and the sodium carbonate dissolves. Observe the reaction of vinegar, citric acid, and dilute sulfuric acid, on sodium carbonate, washing soda, crystals.
12.1.26 Precipitates with sodium carbonate (washing soda)
1. Add sodium carbonate solution to solutions of copper sulfate, iron (II) sulfate, magnesium sulfate, and zinc sulfate. In each case double decomposition occurs.
Reaction for zinc sulfate:
zinc sulfate + sodium carbonate --> zinc carbonate + sodium sulfate.
2. Add sodium carbonate solution to alum (hydrated potassium aluminium sulfate), or aluminium sulfate solution. The precipitate is aluminium hydroxide not aluminium carbonate, because aluminium carbonate is very unstable and is decomposed by the water to form the hydroxide.
12.1.27 Prepare baking soda with sodium carbonate (washing soda)
Make a strong solution of sodium carbonate by heating 10 g of powdered sodium carbonate crystals in a test-tube of water in a beaker. Half fill a test-tube with the strong solution and cool it under the tap. Pass a slow stream of carbon dioxide through the liquid for ten minutes. A white precipitate of sodium bicarbonate forms. Filter the precipitate and leave the opened filter paper to dry out on a piece of newspaper. Tests the precipitate for sodium bicarbonate. If phenolphthalein solution is added to the sodium carbonate solution in the test-tube, the red colour disappears when all the sodium carbonate is converted into sodium bicarbonate.
12.1.28 Prepare bath salts with sodium carbonate and with sodium sesquicarbonate
1. Bath salts crystals may be added to bath water to “soften” it. Sodium carbonate, (washing soda), crystals coloured by dyes may be sold as a cheap form of bath salts. When these bath salts are exposed to the air, they betray their identity by forming a white coating because of efflorescence.
2. Sodium sesquicarbonate (trisodium hydrogen carbonate), Na3H(CO3)2, is a double salt of sodium bicarbonate and sodium carbonate that naturally occurs as Na2CO3.NaHCO3.2H2O, the dihydrate evaporate mineral trona found in evaporation pools. Better quality bath salts containing of crystals of sodium sesquicarbonate are not so alkaline as washing soda bath salts and are kinder to the skin. These bath salts are manufactured by combining sodium carbonate and baking soda together, with colouring matter and perfume. To make bath salt crystals, dissolve 5 g of baking soda in a beaker of hot water. When the baking soda has dissolved, add four times as much powdered sodium carbonate and stir to dissolve all the powder. Transfer the solution to an evaporating basin and leave in a hot place until needle-like crystals form. Unlike washing soda, these crystals are not efflorescent. Sodium sesquicarbonate can also be used as a water softener and to remove copper chloride verdigris from old copper vessels.
12.1.29 Flame test for sodium chloride
Sodium chloride causes a brilliant gold-yellow colour to a flame. This is the chemical test for sodium. “Sodium vapour” lamps are used for lighting roads in towns. The yellow light of sodium makes a person's face appear bloodless and ghost-like.
12.1.30 Prepare hydrochloric acid gas from sodium chloride
Mix together 2 g of sodium chloride and 2 g of powdered alum (hydrated potassium aluminium sulfate), or sodium hydrogen sulfate (sodium bisulfate). Put the mixture into a dry test-tube and have ready a damp blue litmus paper and a bottle of strong ammonia. Heat the mixture over a medium Bunsen burner flame, holding the test-tube in a paper holder and moving the test-tube in the flame. Hydrochloric acid gas, or hydrogen chloride forms as steam-like fumes. Sniff the gas cautiously and put the blue litmus paper into
the fumes. Tests the gas, also, by removing the stopper from the bottle of strong ammonia and blowing the steamy fumes across the top of the bottle. A dense white smoke forms. The white smoke consists of ammonium chloride.
12.1.32 Prepare sodium chloride cubic crystals
Prepare a saturated solution of sodium chloride by heating 2 cm of sodium chloride with half a test-tube of water with a drop of dilute sulfuric acid. Cool the test-tube under the tap. Filter a third of the solution into a beaker. Leave the beaker for twenty four hours in a hot place. Some cubic crystals of sodium chloride are left at the bottom of the beaker. Shine an electric torch up through the bottom of the beaker to see the cubic crystals.
12.1.33 Prepare metal hydroxides by precipitation
Put 2 cm of the following solutions into test-tubes and add dilute sodium hydroxide to each solution: alum (hydrated potassium aluminium sulfate), copper sulfate, iron (II) sulfate, iron (III) chloride, iron alum (iron (III) ammonium sulfate), magnesium sulfate, zinc sulfate. Observe the colours of the precipitates. The hydroxides of aluminium and zinc dissolve if more sodium hydroxide is added and the test-tube is shaken.
12.1.34 Prepare sodium carbonate (washing soda) crystals with caustic soda
Use an apparatus to deliver a steady stream of carbon dioxide. Transfer a test-tube full of dilute caustic soda solution to a boiling tube. Pass a slow, steady stream of carbon dioxide into the solution for five minutes. Make sure that the delivery tube is at the bottom of the liquid. Some gas is absorbed by the solution. Put the remaining solution into an evaporating basin and heat over a flame until one third of it remains. Leave the solution to crystallize. Large colourless crystals of washing soda (sodium carbonate) form.
12.1.35 Prepare sodium silicate precipitates
When a solution of water glass (sodium silicate) is added to solutions of some metal salts, precipitates of metal silicates form. Put into test-tubes 5 mL of the following solutions: alum (hydrated potassium aluminium sulfate), copper sulfate, iron (II) sulfate, iron (III) chloride, iron alum (iron (III) ammonium sulfate), cobalt chloride, nickel sulfate. Prepare a solution of sodium silicate by stirring 5 mL of water glass (sodium silicate) in half a beaker of hot water. Add 5 mL of this solution to each test-tube. Observe the different colours of the precipitates formed.
12.1.36 Prepare chemical gardens
Prepare a strong solution of sodium silicate by dissolving 30 mL of water glass (sodium silicate) in a beaker of hot water. Transfer the solution to a beaker then leave to cool. Drop into the beaker small crystals of any of the chemicals listed in the previous experiment. Leave the beaker where it will not be disturbed, but inspect the crystals in it every day. Later the crystals appear to be sprouting. Growths from the crystals rise through the liquid to appear like plants growing in the solution.
12.1.37 Prepare silicic acid and pure silica
White sand is almost pure silica. However, white silica can be made from water glass (sodium silicate). Prepare strong solutions of sodium silicate and sodium hydrogen sulfate, (sodium bisulfate). Dissolve 5 mL of water glass (sodium silicate) 50 mL of hot water. Dissolve 5 mL of sodium hydrogen sulfate, (sodium bisulfate), in 2 cm of water in a test-tube. Mix the two solutions together in a beaker. A jelly-like precipitate of silicic acid forms. Filter off the gelatinous precipitate then wash it by running hot water through the filter paper. Use a spoon to scrape the precipitate out of the filter paper on to a metal lid. Hold the metal lid in a pair of pliers and heat it over a Bunsen burner. The fine white powder left is pure silica, SiO2.
12.1.38 Prepare sodium sulfate supersaturated solution, Na2SO4.10H2O
Prepare a supersaturated solution of Glauber's salt (sodium sulfate). A supersaturated solution is a solution that contains more of the dissolved substance than is needed to prepare a saturated solution at the given temperature. Half fill a beaker with sodium sulfate crystals then add water until the crystals are covered with 0.5 cm of water. Put the beaker in a larger beaker or metal pot as a water bath with 2 cm of cold water. Heat the larger beaker with a Bunsen burner. When the water starts to boil, turn the flame down, but keep the water boiling until the crystals dissolve. Ignore any white specks in the solution. Cover the smaller beaker containing the sodium sulfate solution to exclude dust then leave it to cool. When cold, the contents of the beaker remain liquid, a supersaturated solution. Drop a small crystal of Glauber's salt (sodium sulfate) into the beaker. Long spiky crystals grow in all directions from the added crystal and the solution becomes almost solid.
12.1.39 Heat sodium thiosulfate, Na2S2O3.5H2O
Use a paper holder to heat crystals of sodium thiosulfate in a dry test-tube. The crystals first melt then lose their water of crystallization. With more heating a yellow deposit of sulfur forms in the cooler part of the test-tube and a smell of hydrogen sulfide gas is noted. Leave the test-tube to cool. Add drops of an acid solution to the remaining substance, sodium sulfide, and note the strong smell of hydrogen sulfide.
12.1.40 Acids with sodium thiosulfate
Do this experiment near an open window or in a fume hood. Heat crystals of sodium thiosulfate in a test-tube with citric acid solution or dilute sulfuric acid solution. Note the sharp, choking smell of sulfur dioxide gas and a milky precipitate of sulfur. A piece of damp blue litmus paper held at the open end of the test-tube turns red because sulfur dioxide is an acid gas.
12.1.41 Copper sulfate solution with sodium thiosulfate
Prepare a strong solution of sodium thiosulfate by heating a dozen crystals with 2 cm of water in a test-tube. Cool the solution under the tap. Add 2 cm of copper sulfate solution drop by drop. The blue colour of the copper sulfate solution fades as it mixes with the sodium thiosulfate. Heat the mixture until it begins to boil. Remove the test-tube from the flame and observe the liquid turning yellow, then brown, and finally a heavy black precipitate of copper sulfide forms.
12.1.42 Prepare sodium thiosulfate supersaturated solution
Add two drops of water to 2 cm of sodium thiosulfate crystals in a test-tube then heat the test-tube gently over a flame. When the crystals have dissolved, plug the mouth of the test-tube with cotton wool then leave it to cool. The contents of the test-tube remain liquid. Remove the cotton wool and pour the liquid onto a sheet of clean glass. Drop a tiny crystal of sodium thiosulfate into the liquid on the clean glass. Crystals form in all directions from the tiny crystal and the liquid becomes completely solid.
12.2.4 Prepare sulfuric acid with iron (II) sulfate
Put 2 cm of powdered iron (II) sulfate crystals into a hard glass test-tube. Fit it with a stopper through which passes a right angle piece of glass tubing. The other end of the tubing dips into a test-tube. Clamp the hard glass test-tube loosely in a stand or hold it in a paper holder. Keep the test-tube sloping downward or moisture condenses in the cooler part of the test-tube, runs back on to the hot glass, and cracks the test-tube. Heat gently at first, moving the flame, then more strongly. Water of crystallization is produced at first and the substance changes to white anhydrous iron (II) sulfate that decomposes with stronger heat. A thick white vapour appears and condenses in the test-tube as a colourless liquid. When no more vapour is produced, let the apparatus cool. The liquid collected in the test-tube is a weak solution of sulfuric acid. Tests it with blue litmus paper and a crystal of sodium carbonate (washing soda). Carbon dioxide gas forms. The red substance left in the hard glass test-tube is red iron oxide, called jewellers' rouge, because jewellers use it for polishing gold and silver.
12.2.8 Reduce iron (III) sulfate to iron (II) sulfate
Put 2 cm of ammonium iron (III) sulfate solution in a test-tube with an equal amount of dilute sulfuric acid or sodium hydrogen sulfate solution. Add 2 mL of iron filings or small pieces of zinc. Heat the test-tube until effervescence starts because of the formation of hydrogen. The hydrogen is the reducing agent. Leave the test-tube to stand. The colour of the solution appears to vanish, but it is a very light green. Test the solution by adding dilute ammonia solution. A dirty green precipitate forms to indicate iron (II) sulfate.
12.3.13 Dilute nitric acid with metals
Put strands of copper wire from an old electric light flex into a test-tube with 2 cm of dilute nitric acid. Heat the test-tube gently until effervescence starts and then stand the test-tube in the test-tube rack or a beaker. The effervescence is caused by the formation of oxides of nitrogen, principally nitric oxide, NO, and nitrogen peroxide, NO2, the first of these gases is colourless, while the second consists of brown fumes. In minutes the copper has dissolved in the acid and a green-blue solution of copper nitrate is left. In a similar way iron filings and zinc can be dissolved in dilute nitric acid to give solutions of iron (II) nitrate and zinc nitrate.
12.3.14 Dilute nitric acid with metal oxides
Heat black copper oxide with 2 cm of dilute nitric acid in a test-tube. The copper oxide disappears and a blue solution of copper nitrate forms.
12.3.15 Dilute nitric acid forms carbon dioxide with carbonates and bicarbonates
Put 2 cm of chalk (calcium carbonate), washing soda (sodium carbonate), or baking soda (sodium hydrogen carbonate, bicarbonate) into a test-tube and add drops of dilute nitric acid. Observe the vigorous effervescence because of the formation of carbon dioxide. Tests the gas with limewater.
12.4.10 Magnesium sulfate with ammonia
Add dilute ammonia solution to a solution of magnesium sulfate. A white precipitate of magnesium hydroxide forms.
12.4.11 Magnesium sulfate with sodium carbonate
Add 5 g of sodium carbonate solution to 5g of magnesium sulfate solution in a test-tube. A white precipitate of magnesium carbonate forms.
The double decomposition reaction:
magnesium sulfate + sodium carbonate --> magnesium carbonate + sodium sulfate.
The precipitate is readily soluble in acids. Add drops of any acid solution to the test-tube to make the magnesium carbonate precipitate disappear.
12.4.12 Tests for magnesium in compounds
Pour drops of magnesium sulfate solution on a filter paper, then drops of cobalt chloride solution. Heat the wet paper over a flame until dry and then use the flame to ignite it over a watch glass or saucer to catch any ash. The ash has a pink colour. The same result occurs for any solution that contains magnesium.
12.5.1 Combustion of zinc
1. Sprinkle the powder into a non-luminous Bunsen burner flame. The particles burn with green flashes.
2. Put 5 mL of zinc powder into an evaporating basin and stir it with two drops of sodium hydroxide solution so that a thick paste forms. Transfer the paste with a spoon to a filter paper placed on a newspaper and press the paste between the paper to squeeze out most of the moisture. Put the remaining damp cake on a metal lid and leave it in the air. Clouds of steam rise up and the mass turns yellow as the zinc combines with oxygen from the air to form zinc oxide. The zinc ignites after 5 minutes or more. The zinc oxide turns white as it cools.
12.5.2 Acids with zinc
Zinc causes hydrogen to be produced from most acids. Add granulated zinc to 2 cm of the following acids in a test-tube: vinegar, citric acid, tartaric acid, dilute sulfuric acid. Test for hydrogen with a glowing splint. The reaction with powdered zinc is more vigorous. If using zinc foil, heat the acid. The reaction is more vigorous if a drop of copper sulfate solution is added.
12.5.3 Alkalis with zinc
Heat small pieces of zinc foil with 2 cm of sodium hydroxide or sodium carbonate solution in a test-tube. Test for hydrogen with a glowing splint. With powdered zinc, the reaction may be violent, so use only drops of the alkali solution.
12.6.1 Prepare citric acid crystals with lemon juice
Squeeze a lemon with the help of a lemon squeezer and transfer the juice to a beaker. Add an equal amount of water and boil the liquid gently over a gauze and tripod for quarter of an hour. Filter the liquid while still hot into an evaporating basin. The filtrate is a solution of citric acid. Evaporate the solution on the gauze and tripod until only one third of it remains. Then put it on one side to cool and crystallize. The crystals of citric acid obtained in this way are not very pure. They are usually coloured brown. To obtain purer crystals, dissolve the impure substance in water and boil the solution with powdered charcoal, (preferably “decolorizing” charcoal), for ten minutes. Then filter the liquid and leave it in a hot place to evaporate and crystallize.
12.6.2 Heat citric acid
Put 2 mL of citric acid crystals in a metal spoon. Hold the end of the spoon in a paper holder and heat the citric acid crystals over a flame. The crystals first melt and then give off a vapour which ignites and burns with a yellow flame. A black deposit of carbon remains on the spoon after the reaction, showing that citric acid is an organic substance. The carbon can be burned off the spoon by continuing the heating.
12.6.3 Prepare hydrogen with citric acid
Put 2 mL of citric acid crystals and an equal amount of iron filings in a test-tube. Add drops of water then heat the test-tube. When the effervescence becomes brisk, remove the test-tube from the flame. To test for hydrogen, put the thumb over the end of the test-tube, count to five, then apply a glowing splint to the end of the test-tube. Repeat the experiment with small pieces of zinc instead of iron filings. In both experiments the formation of the hydrogen is increased by adding drops of copper sulfate solution.
12.6.4 Prepare sodium citrate crystals
Heat 2 cm of citric acid crystals in a test-tube with 2 cm of water. When the crystals have dissolved put the hot solution into an evaporating basin and add powdered sodium carbonate, (washing soda), to the dish. Effervescence occurs because of the formation of carbon dioxide. Continue adding powdered sodium carbonate, (washing soda), in small amounts until effervescence ceases, showing that the acid has been used up. Evaporate the remaining solution on a gauze and tripod until one third of the liquid remains.
When the solution is left to cool, white crystals of sodium citrate form.
12.6.5 Citric acid solubility and temperature
The solubility of solid substances in water usually increases with rise of temperature. However, calcium citrate is less soluble in hot water than in cold water. Put 2 mL of calcium carbonate in a test-tube and add drops of citric acid. Carbon dioxide forms. When the reaction stops, add more drops of the acid solution until the calcium carbonate has completely dissolved. The clear liquid is a solution of calcium citrate. Hold the test-tube in a paper holder and heat the solution over a flame. A white precipitate of calcium citrate forms in the liquid because of the lower solubility of calcium citrate at higher temperatures.
12.7.1 Heat glucose
Use a small piece of barley sugar or Glucodin. Heat the glucose on a metal lid held in a pair of pliers. The substance melts and turns brown then black. Note the smell of burnt sugar. The final black residue is carbon, sugar charcoal, a very pure form of carbon.
12.7.2 Tests for glucose and sucrose
Heat a glucose solution with copper hydroxide solution. A precipitate of copper (I) oxide, (Cu2O, cuprous oxide), forms that gradually turns to red. This test reaction does not occur with sucrose.
12.7.3 Tests for glucose in apples and sweets
Cut four pea-size pieces of apple and put them into a test-tube with 5 mL of sodium carbonate, (washing soda), crystals. Add 2 cm of water. Hold the test-tube in a paper holder and heat it over a flame. When the liquid begins to boil, continue the heating for four minutes. A red-brown solution gradually forms that then gradually darkens until it is almost black. Note the faint smell of burnt sugar. This test reaction does not occur with sucrose. Repeat the test with barley sugar and Glucodin.
12.7.4 Tests for glycerine
Heat drops of glycerine in a dry test-tube with powdered sodium hydrogen sulfate, (sodium bisulfate). Acrolein, (C3H4O, ethylene aldehyde), vapour forms with a penetrating acrid smell of burnt fat in burning cooking oil.
12.7.5 Glycerine colour change reaction
In test-tube 1, dissolve 2 mL of borax in half a test-tube of water and add two drops of phenolphthalein solution. The liquid turns a rose red colour. In test-tube 2 dissolve one drop of glycerine in 2 cm of water. Add the test-tube 2 glycerine solution, drop by drop, to the test-tube 1 borax solution until the red colour disappears. Heat test-tube 1 and the red colour reappears. Cool test-tube 1 and the red colour disappears. This heating and cooling effect can be repeated. Repeat the experiment using sugar solution instead of glycerine.
12.7.6 Oxidation of glycerine
Add drops of hydrogen peroxide and a small crystal of iron (II) sulfate to 5 g of glycerine in a test-tube. Hold the test-tube in a paper holder and heat the mixture over a flame. A vigorous reaction occurs as the glycerine is oxidized to glyceraldehyde. Add a blue precipitate from the reaction of copper hydroxide with copper sulfate solution and sodium hydroxide solution. Heat the test-tube and a yellow precipitate that turns red of copper (I) oxide (Cu2O, cuprous oxide) forms.
12.7.7 Glycerine with sugar
Put into a dry test-tube 5 g of glycerine and an equal amount of sucrose sugar. Heat the test-tube over a flame. A black mass of carbon forms in the test-tube.
12.7.8 Glycerine with cobalt chloride solution
Dissolve 5 g of cobalt chloride crystals in 2 cm of water in a test-tube then add 5 g of glycerine. Heat the test-tube to see the red colour of the liquid turning to a violet-blue colour. Cool the test-tube to see the red colour return.
12.7.9 Prepare casein with milk
1. Use acid to make milk curdle. Use skimmed milk or dried milk that has been reconverted to milk by mixing with hot water. Leave fresh milk to stand and later pour off the top layer of cream. Heat half a beaker of the milk, but stop heating before the beaker becomes too hot to hold in the hand. Pour a test-tube of vinegar into the beaker while stirring the milk with a glass rod, The milk curdles. Remove the curd with a spoon onto a newspaper and squeeze out the moisture by folding and pressing the newspaper
to leave a spongy mass left of casein.
2. Milk also curdles when it turns sour. This reaction is caused by bacteria that change the milk sugar into lactic acid, C3H6O3, that precipitates the casein as a curd. Pure lactic acid is a white crystalline substance.
12.7.10 Prepare lactic acid with milk
Leave a beaker of milk on a shelf for days until it has gone sour. Pour off the watery liquid, leaving as much as possible of the white curd behind. Filter the liquid and boil it in a beaker for a few minutes to precipitate any finely suspended matter. Leave the liquid to cool and again filter it. The filtrate is an almost colourless solution of lactic acid. Tests it with blue litmus paper. Form crystals by boiling the solution in an evaporating basin on a gauze and tripod until the volume of the solution has been reduced to one quarter.
Complete the evaporation by leaving the evaporating dish in a hot place.
12.7.11 Tests for lactic acid solution
1. Add lactic acid solution to the solution to sodium hydrogen carbonate (baking soda) in a test-tube. Note the effervescence because of the formation of carbon dioxide gas. Test for carbon dioxide with limewater.
2. Heat lactic acid solution with iron filings. Note the effervescence because of the formation of hydrogen gas. Increase the reaction by adding drops of copper sulfate solution. However, it is difficult to obtain sufficient hydrogen to test by explosion with a glowing splint.
3. Boil 5 ml of lactic acid solution with two drops of dilute sulfuric acid. Leave the solution to cool then add it to a copper hydroxide precipitate from the reaction of copper sulfate with sodium hydroxide. Heat the solution and observe a yellow precipitate which turns red as copper (I) oxide (cuprous oxide) forms.
12.7.12 Heat starch
Heat starch on a metal lid. Decomposition occurs and inflammable gases, which smell like burning leather, forms. A black residue of carbon remains on the lid.
12.7.13 Starch with water
Shake a small pinch of powdered starch with half a test-tube of water. The starch does not dissolve. Tests the liquid by adding one drop of iodine solution. No reaction is given. Boil a small pinch of starch with half a test-tube of water for seconds. The starch dissolves. Cool the test-tube under the tap and add one drop of iodine solution. A deep blue-black liquid forms.
12.7.14 Tests for starch in adhesive paste
Shake a drop of the paste with water in a test-tube and add a drop of iodine solution. If the paste contains starch the contents of the test-tube turns blue-black. If a red colour forms the paste contains a chemical called dextrin, that is made from starch.
12.7.15 Prepare glucose with starch.
Boil a 2 mL of starch with half a test-tube of dilute sulfuric acid or sodium hydrogen sulfate, (sodium bisulfate), solution for two minutes. Then pour off drops of the liquid into another test-tube, cool it under the tap, and test it by adding a drop of iodine solution. The liquid turns red. The starch has been changed into dextrin, a substance which, like starch, is given the same formula. In this case, however, z (number of amylose units) is supposed to be a much smaller number than for starch. Continue boiling the remainder of the liquid for another three minutes to convert the dextrin into glucose. To show that glucose has been formed, cool the liquid and
neutralize it in an evaporating basin with dilute sodium carbonate or sodium hydroxide solution (test with litmus paper). Add 2 cm of the neutralized solution to a precipitate of copper hydroxide made from solutions of copper sulfate and sodium hydroxide. If the test-tube is warmed a yellow precipitate, which turns red, of cuprous oxide forms. Dry heat breaks down starch to dextrins, ("pyrodextrins") to give the brown colour of toast.
12.7.16 Heat sucrose
Heat sugar on a metal lid. The substance melt to form a liquid which soon turns brown. If it is cooled at this stage the brown solid obtained is called caramel. When heated more strongly the sugar decomposes, giving off inflammable vapours and leaving a black mass of sugar charcoal on the lid.
12.7.17 Sucrose with sodium hydrogen sulfate, NaHSO4, (sodium bisulfate),
Mix together a 2 mL of sugar and a 2 mL of powdered sodium hydrogen sulfate, (sodium bisulfate). Heat the mixture in a dry test-tube. The contents of the test-tube swells up forming a black puffy mass of carbon.
12.7.18 Prepare glucose with sugar
Boil a 2 mL of sugar in a test-tube with 2 cm of dilute sulfuric acid or sodium hydrogen sulfate, (sodium bisulfate), solution. Keep the liquid boiling for two or three minutes and then cool the test-tube under the tap. To show that the solution now contains glucose, first neutralize the remaining acid with sodium carbonate or dilute sodium hydroxide solution (test with litmus paper). Then apply the copper hydroxide test.
12.7.19 Sucrose with borax
A colour reaction. Sugar has a similar reaction to glycerine on borax and phenolphthalein but different strengths of solutions are needed. Dissolve half a 2 mL of borax in a test-tube nearly full of water and add one or two drops of phenolphthalein to obtain a rose red liquid. Add solid sugar, at a time, and shake the test-tube. The colour disappears. Heat the test-tube, the colour reappears, only to vanish again when the test-tube is cooled under the tap.
12.7.20 Ferment sugar with yeast
Fermentation is a chemical reaction caused by lowly forms of life, such as bacteria and moulds. We have already seen an example of this in the turning sour of milk. Baker's yeast is a simple form of plant life which, under suitable conditions, is able to turn sugar (and starch) into alcohol and carbon dioxide. In bread-making it is the carbon dioxide gas which puffs up the dough and makes the bread light. Dissolve 5 mL of sugar in a beaker of water and put the solution into a flask that is fitted with a delivery tube dipping into limewater. Add 5 mL of baker's yeast and leave the apparatus in a hot place, such as a shelf of the airing cupboard. In an hour or two the contents of the flask begins to froth and the limewater turn milky, showing that carbon dioxide is being produced. If the flask is left for two or three days and the contents are then filtered, alcohol can form from the filtrate by distillation. It is illegal to distil alcohol, and an offender is liable to a long term of imprisonment, besides having his apparatus confiscated! when the distillation is done in a laboratory the first few drops of the liquid formed by distillation burn with a blue flame, that is characteristic of alcohol.
12.8.1 Evaporate mothballs (naphthalene crystals)
It is most unusual for a solid substance to evaporate away when left out in the air. This, however, happens with naphthalene. Leave one or two mothballs on a watch glass or saucer outside in the open air. Examine the mothballs some days later and note that they are smaller because of evaporation. In a week they may disappear completely.
12.8.2 Burn mothballs (naphthalene crystals)
Like all hydrocarbons naphthalene burns readily. Crush a mothball and put of the powder into a metal screw cap. Hold the latter in a pair of pliers or pincers and heat it over a flame. The white powder melts then ignites. It burns with a very smoky flame because of the high percentage of carbon in naphthalene, (94%).
12.8.3 Prepare naphthalene crystals with mothballs
Put four or five mothballs into a beaker and place a funnel in the top of the jar. Stand the beaker in a saucepan containing water ( 2 cm deep). Heat the water until it begins to boil and then remove the saucepan from the flame. fn minutes beautiful starry crystals are deposited on the sides of the beaker and funnel. If you remove the funnel and blow it gently, the crystals float off into the air and glisten brightly as they slowly fall. Instead of a funnel use a glass plate over the jar. In this experiment the mothballs do not melt but change straight into vapour which condenses again on the cooler part of the beaker and funnel this change is called sublimation. It is similar to the subliming of ammonium chloride. On a hot summer day mothballs can be made to sublime merely by leaving them in a beaker on a window sill where sunlight falls on the jar. Starry crystals again form. Another method of making naphthalene crystals, but of a different shape, is to dissolve two salt spoonfuls of the powdered substance in half a test-tube of methylated spirit by shaking (do not heat methylated spirits). Leave the solution in an evaporating basin on a shelf to evaporate. Naphthalene crystals form as pearly plates.
12.10.1 Tea with iron (II) sulfate
Wash a small crystal of iron (II) sulfate in a test-tube with water and pour away the water to remove any surface layer of iron (III) sulfate. Dissolve the crystal in half a test-tube of cold water and add drops of cold tea. A violet liquid forms.
12.10.2 Tea with iron (III) salts
Add drops of cold tea to a solution of ammonium iron (III) sulfate or iron (III) chloride. A black precipitate forms that can be the basis of most types of blue-black ink.
12.10.3 Tea acid with dilute sulfuric acid
Add dilute sulfuric acid or sodium hydrogen sulfate, (sodium bisulfate), solution to half a test-tube of cold tea. The liquid is cloudy, because of the precipitation from the tea. Test the liquid a iron (III) salt solution. No black precipitate forms as in the previous experiment.
12.10.4 Tea acid with limewater
Add 2 cm of limewater to an equal amount of cold tea in a test-tube and boil the mixture. A red-brown precipitate of the calcium salt forms.
12.11.1 Prepare verdigris with copper and vinegar
Put a piece of copper coin in a watch glass or saucer and pour drops of vinegar on to the surface of the copper. Leave the copper undisturbed for hours, until the liquid has evaporated. Green-blue particles are left on the surface of the coin. Scrape these off on to a piece of white paper and wash the coin. If the coin is an old, black copper oxide previously on the surface is removed, but black copper sulfide remains.
12.11.2 Decolorize vinegar
Boil vinegar with decolorizing charcoal. Ordinary charcoal is not very successful. Neutralize vinegar with ammonia then half fill an evaporating basin with it and add a piece of litmus paper. Add dilute ammonia to the evaporating basin then stir the solution until the litmus paper turns blue. Transfer the liquid, now a solution of ammonium acetate, to a test-tube, add drops of hydrogen peroxide then heat the mixture until it boils. The brown colour of the solution disappears, destroyed by the oxygen from the hydrogen peroxide. Use the remaining liquid for the next experiment.
12.11.3 Tests for acetic acid in vinegar
Cool the neutralized vinegar remaining from the previous experiment under the tap. Add to it drops of ammonium iron (III) sulfate or iron (III) chloride (ferric chloride) in solution. The liquid turns a bright red colour.
12.12.1 Decompose ammonium carbonate
When heated ammonium carbonate decomposes completely into three gases or vapours, steam, ammonia, and carbon dioxide, so that the substance disappears. Heat ammonium carbonate in a dry test-tube. Holding the test-tube in a paper holder so that it slopes down slightly. Observe the steam and drops of water in the cooler part of the tube. Test for ammonia gas by smell and by holding a piece of damp red litmus paper at the mouth of the tube. The litmus paper turns blue. Test for carbon dioxide with limewater. Later, none of the white powder remains at the bottom of the test-tube.
12.12.2 Alkalis with ammonium carbonate
Warm a little ammonium carbonate in a test-tube with 2 cm of washing soda solution or limewater. Ammonia gas can again be detected by smell or by litmus.
12.12.3 Acids with ammonium carbonate
Add vinegar or citric acid to some ammonium carbonate in a test-tube. Note the effervescence. Test for carbon dioxide with limewater.
12.12.4 Precipitates with ammonium carbonate
Add ammonium carbonate solution to solutions of copper sulfate, ferrous sulfate, magnesium sulfate, zinc sulfate, and to limewater. Note the colours of the precipitates.
12.13.1 Tests for aluminium compounds
Put two drops of red cobalt chloride solution on to a piece of filter paper. Add two drops of aluminium sulfate solution. Dry the paper by holding it over a flame and then ignite it over a saucer. The ash is blue. This is a test for all aluminium compounds in solution.
12.14.1 Prepare ammonium sulfate by neutralization
Put half a test-tube of dilute sulphuric acid into an evaporating basin and add red litmus paper. Add dilute ammonia solution, drop by drop. stirring the liquid with a wood splint after each addition drop. Continue adding the ammonia solution and stirring until the litmus paper turns blue. Remove the litmus paper and leave the evaporating basin in a hot place to evaporate or evaporate two thirds of the liquid by heating before leaving it. White crystals of ammonium sulfate form in the dish.