School Science Lessons
Topic 12G
2018-11-01
Please send comments to: J.Elfick@uq.edu.au
Table of contents
See: Chemicals, (Commercial)
12.1.18 Acids with baking soda
12.1.25 Acids with sodium carbonate
12.1.40 Acids with sodium thiosulfate
12.5.2 Acids with zinc
12.5.3 Alkalis with zinc
12.12.1 Ammonium carbonate decomposition
12.12.4 Ammonium carbonate precipitates
12.12.3 Ammonium carbonate with acids
12.12.2 Ammonium carbonate with alkalis
12.1.6 Ammonium iron (III) sulfate
12.1.44 Bromine water oxidizes sodium thiosulfate to sodium sulfate and sulfur
12.8.2 Burn naphthalene crystals
12.1.11 Butyl chloride rainbow reactions
12.1.43 Chlorine water oxidizes sodium thiosulfate to sodium sulfate and sulfur
12.6.5 Citric acid solubility and temperatures
12.6.6 Citric acid with sodium hydrogen carbonate solution
12.1.41 Copper (II) sulfate reduction to copper sulfide, yellow snowstorm reaction
12.11.2 Decolorize vinegar
12.3.15 Dilute nitric acid with carbonates and bicarbonates
12.3.13 Nitric acid with metals
12.3.14 Dilute nitric acid with metal oxides
12.8.1 Naphthalene evaporation
12.7.6 Glycerine oxidation
12.7.5 Glycerine with borax solution, colour change
12.7.8 Glycerine with cobalt chloride solution
12.6.2 Heat citric acid to form carbon
12.7.1 Heat glucose to form carbon
12.7.7 Heat glycerine with sugar to form carbon
12.1.17 Heat sodium bicarbonate
12.1.24 Heat sodium carbonate to form anhydrous sodium carbonate
12.1.39 Heat sodium thiosulfate
12.7.12 Heat starch to form carbon
12.1.3 Iron (II) sulfate oxidation to iron (III) sulfate
12.2.8 Iron (III) sulfate reduction to iron (II) sulfate
12.1.2 Iron (II) sulfate with ammonia
12.1.1 Iron (II) sulfate with sodium carbonate
12.4.10 Magnesium sulfate with ammonia
12.4.11 Magnesium sulfate with sodium carbonate
12.14.3 Prepare ammonium iron (II) sulfate, Mohr's Salt
12.1.6.1 Prepare ammonium iron (III) sulfate solution
12.14.1 Prepare ammonium sulfate by neutralization
12.1.28 Prepare bath salts
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
12.1.30 Prepare hydrochloric acid gas with sodium chloride
12.6.3 Prepare hydrogen gas with citric acid
12.1.33 Prepare hydroxides by precipitation
12.14.2 Prepare iron (II) ammonium phosphate
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
12.8.3 Prepare naphthalene crystals from naphthalene mothballs
12.8.4 Prepare rock candy crystals, sugar crystals
12.1.8 Prepare salts by neutralization of sodium hydroxide
12.1.37 Prepare silicic acid and pure silica
12.1.27 Prepare sodium bicarbonate with sodium carbonate
12.1.34 Prepare sodium carbonate with caustic soda
12.1.22 Prepare sodium carbonate with baking soda
12.1.32 Prepare sodium chloride crystals
12.6.4 Prepare sodium citrate crystals
12.1.14 Prepare sodium hydroxide
12.1.13 Prepare sodium hypochlorite
12.1.16 Prepare sodium nitrate crystals by neutralization
12.1.35 Prepare sodium silicate
12.1.38 Prepare sodium sulfate crystals
12.2.4 Prepare sulfuric acid with iron (II) sulfate
12.11.1 Prepare verdigris with copper and vinegar
12.1.19 Sodium bicarbonate precipitates
12.1.21 Sodium bicarbonate prevents milk from going sour
12.1.23 Sodium carbonate efflorescence
12.1.26 Sodium carbonate precipitates
12.1.29 Sodium chloride flame tests
12.1.10 Sodium hypochlorite bleach
18.7.2.2.2 Sodium hypochlorite
12.1.46 Sodium thiosulfate decomposes to form sodium sulfate and sodium pentasulfide
12.1.47 Sodium thiosulfate with dilute hydrochloric acid forms sulfur dioxide and sulfur
12.1.45 Sodium thiosulfate with iodine forms sodium tetrathionate and sodium iodide
12.1.48 Sodium thiosulfate with silver chloride or silver bromide
12.7.13 Starch with water, iodine test
12.7.19 Sucrose with borax
12.7.17 Sucrose with sodium hydrogen sulfate
12.7.20 Ferment sucrose with yeast
12.10.3 Tea with dilute sulfuric acid
12.10.1 Tea with iron (II) sulfate
12.10.4 Tea with lime water
12.10.2 Tea with iron (III) salts
12.11.3 Tests for acetic acid in vinegar
12.1.9 Tests for aspirin
12.7.2 Tests for glucose and sucrose, copper hydroxide
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
12.1.20 Tests for sodium bicarbonate in a stomach powder
12.1.15 Tests for sodium hydroxide, soapy feel
12.7.14 Tests for starch in adhesive paste
12.5.1 Zinc combustion
12.5.2 Zinc with acids
12.5.3 Zinc with alkalis
Zinc, Table of Elements

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 green-grey 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 Iron (II) sulfate oxidation to iron (III) 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.

12.1.6.1 Prepare ammonium iron (III) sulfate solution
Dissolve 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 Sodium hypochlorite 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 Butyl chloride rainbow reactions
tert-butyl chloride, t-Butyl chloride, 1-chlorobutane, 2-chloro-2-methylpropane, (CH3)3CCl
Make a pH 12 solution by adding 10 drops of 0.1 M NaOH to 100 mL water, in a 250 mL beaker.
Add universal indicator to produce a distinct colour.
Start with universal indicator.
Use a second 250 mL beaker to mix by pouring the solution back and forth between the two beakers or put a magnetic bar into the
solution and start the stirrer motor at a fast rate.
Add 15 drops of t-butyl chloride (2-chloro-2-methylpropane) to the solution and begin mixing.
Observe any colour changes.
After 40 seconds add universal indicator and observe any colour changes.
The full range of colour changes (purple, blue, cyan, emerald-green, lime-green, yellow, orange, orange-red, take about two minutes.
The changes in the middle are more rapid than the changes at either extreme.
Use different indicators to show different colour changes and different induction times

2. Prepare two solutions: 0.1 M 2-chloro-2 methylpropane (t-butyl chloride) in ethanol (1 g per 100 mL) and 0.01 M sodium hydroxide.
Put 5 mL 0.1 M C4H9Cl in a test-tube.
In another test-tube put 5 mL 0.1 M NaOH, 10 mL water and two drops of any one of the following indicators.
Mix the solutions back and forth once and observe for the colour change that occurs after an induction period.
With equal volumes 0.01 M sodium hydroxide and 0.1 M C4H9Cl the colour changes with universal indicator were: Purple to blue
(on mixing), blue to green (after 12 seconds), green to yellow (after 15 seconds), yellow to orange (after 25 seconds total).
Cooling the solutions greatly slows the reaction, increasing the induction period, e.g. with iced water, the methyl red change took more
than 50 seconds.

12.1.13 Prepare sodium hypochlorite
Prepare sodium chlorate (I) solution 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
The double decomposition reaction: sodium carbonate + calcium hydroxide --> sodium hydroxide + calcium carbonate (precipitate)
1. Make a concentrated sodium carbonate solution and add to it calcium hydroxide solution.
Shake well for three minutes.
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.
Test the sodium hydroxide solution with universal indicator paper.
The universal indicator paper turns blue or violet.

2. 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.

3. Dissolve sodium oxide in water to form soluble sodium hydroxide.
Na2O (s) + H2O (l) --> 2NaOH (aq)
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 then be used for other experiments.
Double decomposition reaction:
calcium hydroxide + sodium carbonate --> calcium carbonate + sodium hydroxide.

12.1.15 Tests for sodium hydroxide, soapy feel
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 sodium bicarbonate
Sodium bicarbonate, sodium hydrogen carbonate, "bicarb", "baking soda", bicarbonate of soda, nahalite, NaHCO3
Put 5 g of sodium bicarbonate into a dry test-tube 1.
Put 2 cm of clear lime water in test-tube 2.
Hld 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 lime water in test-tube 2.
The lime water turns milky, showing that carbon dioxide is being produced.
Chemical action: sodium bicarbonate + heat --> sodium carbonate + water + carbon dioxide
2NaHCO3 (s) --> CO2 (g) + Na2CO3 (s) + H2O (g)
Na2CO3 (s) --> CO2 (g) + Na2O (s)

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 Sodium bicarbonate precipitates
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 lime water.

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 with baking soda
Sodium bicarbonate, sodium hydrogen carbonate, "bicarb", "baking soda", bicarbonate of soda, nahalite, NaHCO3
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.
2NaHCO3 (s) --> CO2 (g) + Na2CO3 (s) + H2O (g)

12.1.23 Sodium carbonate effflorescence
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, and can be called sodium carbonate monohydrate.

12.1.24 Heat sodium carbonate crystals to form anhydrous sodium carbonate
Washing soda, Na2CO3.10H2O, sodium carbonate decahydrate, natron, effloresces to form sodium carbonate monohydrate,
thermonitrite, Na2CO3.H2O
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 as soda ash, has lost all of its water of crystallization.
Na2CO3.10H2O --> Na2CO3 + 10H2O

12.1.25 Acids with sodium carbonate
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.
2CH3COOH (aq) + Na2CO3 (aq) --> 2CH3COONa (aq) + H2O (l) + CO2 (g)
ethanoic acid + sodium carbonate --> sodium ethanoate + water + carbon dioxide

12.1.26 Sodium carbonate precipitates
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 sodium bicarbonate with sodium carbonate
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
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 Sodium chloride flame tests
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 with 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 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 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 with caustic soda
Sodium carbonate crystals, Na2CO3.H2O, are called washing soda.
Use an apparatus to deliver a steady stream of carbon dioxide.
Transfer a test-tube full of dilute sodium hydroxide 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
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 crystals
Sodium sulfate crystals, crystalline sodium sulfate, Na2SO4.10H2O, as Glauber's salt, is used as a laxative.
Sodium sulfate forms supersaturated solutions where crystallization can occur suddenly by adding a crystal of the salt, by exposure to
the air or by touching the solution with a glass rod.
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.
Prepare a supersaturated solution of sodium sulfate.
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 (II) sulfate reduction to copper sulfide, yellow snowstorm reaction
The thiosulfate ion, S2O32-, is a reducing agent.
The cupric ion, Cu2+, is an oxidizing agent.
The blue cupric ion., Cu2+ may be reduced to the colourless cuprous ion, Cu+.
1. 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.
2. Equimolar ratio of sodium thiosulfate solution with copper sulfate solution
Prepare 0.5 M copper (II) sulfate solution and 0.5 M sodium thiosulfate solution and observe their colours.
Mix 2 mL of each solution.
Stir the mixture, observe the sudden colour change to pale green, and leave to stand.
After 15 minutes a precipitate begins to form and becomes a yellow copper-thiosulfate complex, e.g. [Cu(S2O3)3]5, after 24 hours.
Use a centrifuge to separate the precipitate and leave to stand separately the precipitate and the liquid.
After a few days a dark layer forms on top of the yellow precipitate.
Some of the thiosulfate reduces some of the copper (II) to copper (I).
3. Excess of sodium thiosulfate solution with copper sulfate solution
Add 1 mL of 0.5 M copper (II) sulfate solution to 3 mL of 0.5 M sodium thiosulfate solution.
Stir the mixture, observe any colour change, and leave to stand.
A dark precipitate of CuS forms after 48 hours with an iridescent coating on the inside of the test-tube.
Leave for weeks and the iridescent coating becomes blue-black, thicker and darker.
Test the precipitate for sulfide in a fume cupboard.
Dry the precipitate, add a drop of nitric acid, hold a piece of lead acetate paper near it to become black in the presence of H2S.
Some of the thiosulfate reduces all of the copper (II) to copper (I).

12.1.43 Chlorine water oxidizes sodium thiosulfate to sodium sulfate and sulfur
Na2S2O3 + Cl2 + H2O --> 2HCl + Na2SO4 + S
Sodium thiosulfate destroys unreacted chlorine in the process of bleaching, so it acts as an antichlor.

12.1.44 Bromine water oxidizes sodium thiosulfate to sodium sulfate and sulfur
Bromine water, bromine solution, Br2 (aq), Toxic by all routes
Na2S2O3 + Br2 + H2O --> Na2SO4 + 2HBr + S

12.1.45 Sodium thiosulfate with iodine forms sodium tetrathionate and sodium iodide
2Na2S2O3 + I2 --> Na2S4O6 + 2NaI
With this reaction, sodium thiosulfate is used to remove iodine stains from clothes.

12.1.46 Sodium thiosulfate decomposes to form sodium sulfate and sodium pentasulfide
4Na2S2O3 --> 3Na2SO4 + Na2S5

12.1.47 Sodium thiosulfate with dilute hydrochloric acid forms sulfur dioxide and sulfur
Na2S2O3 (aq) + 2HCl (aq) --> 2NaCl (aq) + S (s) + SO2 (g) + H2O (l)
The sulfur precipitate turns the mixture milky yellow.

12.1.48 Sodium thiosulfate with silver chloride or silver bromide
Sodium thiosulfate treated with silver chloride or silver bromide forms soluble complex, sodium silver thiocyanate, sodium disthiosulfate,
argentate (I) complex
2Na2S2O3 + 2AgBr --> Na3Ag(S2O3)2 + NaBr

12.2.4 Prepare sulfuric acid with iron (II) sulfate
An early method of preparing vitriol (sulfuric acid), used for bleaching, was by dry distillation of green vitriol (ferrous sulfate,
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 Iron (III) sulfate reduction 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 coiled iron wire or steel wool or small pieces of zinc or zinc powder
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 brown / yellow colour of the solution vanishes, and when it is pure the solution is light green.
2Fe3+ + Zn --> 2Fe2+ + Zn2+
Test1: Add dilute ammonia solution.
A dirty green precipitate forms to indicate iron (II) sulfate.
The precipitate does not dissolve in excess ammonia solution.
Test 2: Add dilute sodium hydroxide solution.
A dirty green precipitate forms that does not dissolve in excess sodium hydroxide.

12.3.13 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 blue-green solution of copper nitrate is left.
3Cu (s) + 8HNO3 (aq) --> 3Cu(NO3)2 (aq) + 2NO (g) + 4H2O (l)
With concentrated nitric acid, the copper is oxidised to Cu2+ ions and the nitric acid is reduced to nitrogen dioxide.
Cu (s) + 4HNO3 (aq) --> CuNO3)2 (aq) + 2NO2 (g) + 2H2O (l)
The solution is dark brown while the nitrogen dioxide is forming.
When all the copper has reacted, with the addition of water, the solution turns pale blue.
Aluminium and zinc do not, or hardly, react with concentrated nitric acid because an oxidation layer forms.
However, zinc reacts with dilute nitric acid.

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.
CuO + 2HNO3 --> CuNO3 + H2O

12.3.15 Dilute nitric 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 lime water.
CaCO3 (s) + 2HNO3 (aq) --> Ca(NO3)2 (aq) + CO2 (g) + H2O (l)

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 Zinc combustion
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 to form carbon
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 that 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 gas 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.

12.6.6 Citric acid with sodium hydrogen carbonate solution
C6H8O7 + 3NaHCO3 --> C6H5Na3O7 + 3CO2 + 3H2O
citric acid + sodium hydrogen carbonate --> sodium citrate + carbon dioxide + water
Lemon juice with sodium bicarbonate
Do not use a glass container.
Put lemon juice into a plastic or tin plate container that can be sealed with a lid, e.g. coffee powder tin, film canister.
Squash absorbent paper, e.g. toilet tissue, newspaper, into the container so that it is wedged against the sides of the container, but
does not touch the lemon juice.
Sprinkle sodium bicarbonate powder over the paper.
Seal the container then invert it.
The citric acid in the lemon juice reacts with the sodium hydrogen carbonate to form carbon dioxide, which increases the gas pressure
in the container and blows off the lid.
Such reactions have been used to make simple rockets.
If the container is attached to a rocket stick, by Newton's third law the force with which the lid is blown off (down), is equal to
an equivalent upward force up, so the rocket moves up.
Be careful! Do the experiment outside with a small amount of reactants.
Stand well away from the container after inverting it and do not look down from directly over the container.

12.7.1 Heat glucose to form carbon
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, copper hydroxide
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
Glycerine, glycerin, is a product containing at least 95% glycerol, C3H8O3, 1, 2, 3-propanetriol.
The terms glycerine, glycerin, glycerol may have no precise meaning in different countries and in different products.
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 with borax solution, colour change
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 Glycerine oxidation
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 Heat glycerine with sugar to form carbon
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 colour.
Cool the test-tube to see the red colour return.

12.7.10 Prepare lactic acid with milk
1. 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.

2. Lactic acid, CH3-CHOH-COOH, forms in milk due to the action of fungi and bacteria acting on the lactose sugar, especially
Lactobacillus
bacteria.
The presence of lactic acid, produced during the lactic acid fermentation is responsible for the sour taste and for the improved
microbiological stability and safety of the food.
Investigate the factors influencing the rate of formation of lactic acid with starter bacteria, e.g. plain yoghurt.
The possible factors include heat, amount of bacteria added, light, access to air, shape of container, sugar concentration, initial pH,
amount of fat (normal, low fat, skim), degree of agitation.

3. Acidity of fresh milk may be measured by titration with a 0.1 M NaOH solution, and shows the reaction with NaOH up to a
phenolphthalein end point.
Fresh milk contains almost no lactic acid and the reaction with NaOH is used to change the pH of carbon dioxide, citrates, casein,
albumin and phosphates.
This indicates lactic acid concentration about 0.13% of the determination of "acidity" in fresh milk by means of titration with NaOH is
a measure of the buffer action of milk, so it can be called the "initial acidity".
The "titrateable" acidity of fresh milk is not due to lactic acid only.
The acidity, which is the result of bacterial activity producing lactic acid during milk collection, transportation, and processing can be
called the "developed acidity".
Measure "titrateable acidity" (not calling it "lactic acid concentration", subtract the initial "acidity" to get a value for the "developed acidity".
An example, for the titrateable acidity of milk:
Day 1: = 0.0288M (7.10 mL titre), initial acidity
Day 7 = 0.0815M (15.65 mL titre), titrateable acidity.

4. Milk fermentation experiments
4.1 Temperature is the independent variable, e.g. 0oC, 10oC, 20oC, 30oC.
titrateable acidity is the dependent variable.
If measurement is done only after 1 week, "time" is a controlled variable, as are variables initial pH, sunlight, sugar concentration,
aeration, exposed surface area.
So prepare a one line graph plotting "titrateable acidity" on the y-axis and temperature on the x-axis.

4.2 Time is the independent variable.
Temperature remains as an independent variable but measurement of titrateable acidity is done every week at 0, 1, 2 and 3 weeks.
So, if using four different temperatures, prepare a four line graph plotting of titrateable acidity on the y-axis and time on the x-axis to
show the fermentation rate at each temperature.

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 lime water.
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 to form carbon
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, iodine test
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.

12.7.17 Sucrose with sodium hydrogen sulfate
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 sucrose 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 that 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 lime water.
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 lime water 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 Naphthalene evaporation
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 naphthalene 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.
Naphthalene sublimes near 70oC, below its melting point 80.26oC.

12.8.2 Burn naphthalene crystals
Like all hydrocarbons naphthalene burns readily.
Crush a naphthalene 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 from naphthalene 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 that 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.

12.8.4 Prepare rock candy crystals, sugar crystals
Cut a length of string longer than the depth of a tall jar and wind the end around a pencil.
Attach a paper clip to the bottom of the string so that it hangs straight down 2 cm from the bottom of the tall jar.
Wet the string and roll it in granulated sugar and leave it to dry.
Slowly add one cup of sugar to three cups of water in a pan while stirring.
Heat the solution until it boils.
Be careful! The boiling point of a sugar solution < 100oC.
Leave the sugar syrup to cool, then pour it into the tall jar.
Place the pencil with attached sugared string across the top of the tall jar.
In a few hour sugar crystals will start to form and will continue forming for days.

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 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 lime water
Add 2 cm of lime water 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
1. In 1773, Antoine Lavoisier (1743-1794, France), reported that when copper transformed into verdigris the substance gained in
weight.
This observation lead to the conclusion that "something" was being taken form the air.
In 1777, after further experiments, he named the "something" as the gas oxygen, "generator of acid".
2. Verdigris has no simple formula, but it is usually mostly basic copper carbonate, CuCO3CH(OH)2, seen in the green colour of old
coins from copper exposed to moist air.
It is harmful if ingested.
In moist sea air, copper chlorides may form.
The verdigris on bronze is called aerugo.
3. The green to green-blue pigment copper acetate used in old oil paintings is also called verdigris.
4. 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.
Blue-green particles of basic copper acetate 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.
Basic copper acetate is used as a paint pigment, mordant and fungicide.
5. Acetic acid with copper oxide forms green-blue crystals of copper acetate, "crystallized verdigris", used as a pesticide, pigment and
for manufacture of Paris green.

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 Ammonium carbonate decomposition
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 lime water.
Later, none of the white powder remains at the bottom of the test-tube.

12.12.2 Ammonium carbonate with alkalis
Warm a little ammonium carbonate in a test-tube with 2 cm of washing soda solution or lime water.
Ammonia gas can again be detected by smell or by litmus.

12.12.3 Ammonium carbonate with acids
Add vinegar or citric acid to some ammonium carbonate in a test-tube.
Note the effervescence.
Test for carbon dioxide with lime water.

12.12.4 Ammonium carbonate precipitates
Add ammonium carbonate solution to solutions of copper sulfate, ferrous sulfate, magnesium sulfate, zinc sulfate, and to lime water.
Note the colours of the precipitates.

12.14.1 Prepare ammonium sulfate by neutralization
Put half a test-tube of dilute sulfuric 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.

12.14.2 Prepare iron (II) ammonium phosphate
Add iron powder and phosphoric acid to demineralized water and stir the mixture to form a partial solution / suspension.
Heat the mixture until no further hydrogen gas is evolved.
Add ammonia solution, ammonium hydroxide, to the mixture to form iron (II) ammonium phosphate, FeNH4PO4.
Dry the mixture to form a green-grey fine powder.
The powder can be purified by passing through a sieve and magnetic separation unit.
Iron (II) ammonium phosphate, known as ferrous ammonium phosphate, FAP, is used for iron fortification of foods and sports drinks.

12.14.3 Prepare ammonium iron (II) sulfate, Mohr's Salt
Dissolve equimolar solutions of hydrated ferrous sulfate and ammonium sulfate in water with drops of sulfuric acid, then leave to form
light green crystals of ferrous ammonium sulfate.
Ammonium iron (II) sulfate, Mohr's Salt (NH4)2Fe(SO4)26H2O has two cations, Fe2+ and NH4+, so it is a double salt of ferrous
sulfate and ammonium sulfate that dissolves in water to form [Fe(H2O)6]2+.

18.7.2.2.2 Sodium hypochlorite, NaOCl
See: Water testing (Commercial)
Sodium hypochlorite is the liquid chlorine for use in swimming pools.
It usually provides 12% to 15% available chlorine and has a pH of 13.
It is generally cheap, but difficult and dangerous to handle.
It also loses its potency rapidly and is usually only used in large swimming pools.
Sodium hypochlorite is stable in alkaline aqueous solutions and is readily available as 5% solution "laundry bleach", e.g. "Chlorox",
and as 12-15% available chlorine solution, often called "liquid bleach", from pool supply stores.
Sodium hypochlorite, with a pH of 13 is relatively low in available chlorine concentration so more is required to maintain the
disinfectant residual in a swimming pool.
Because of the bulk of its liquid form and its poor storage stability, it must be purchased frequently throughout the swimming pool season.
Although it does not add hardness to the swimming pool, its high pH can contribute to scaling tendencies in hard water areas.
It is generally cheap, but difficult and dangerous to handle.
It also loses its potency rapidly and is usually only used in large swimming pools.
NaOCl + H2O --> Na+ + OCl- + H2O
sodium hypochlorite + water --> sodium ion + hypochlorite ion + water
When dissolved, equilibrium is established between the strong oxidant HClO and the weaker ClO- ion.
HOCl + H2O --> OCl- + H3O+ (Ka = 3 × 10-8)
hypochlorous acid + water --> hypochlorite ion + hydronium ion
HOCl is much more effective in oxidizing the cell contents of bacteria because the negative charge on OCR- may prevent passage
through bacterial walls.
At about pH 6 almost all free chlorine is converted to HOCl.
At pH < 2, chlorine gas forms to corrode metals and pool tile grouting
CaSO4 (s) --> Ca2+ (aq) + SO42- (aq)
calcium sulfate (in Portland cement) --> calcium ion + sulfate ion
SO42- (aq) + H+ (aq) --> HSO4- (aq)
sulfate ion + hydrogen ion --> hydrogen sulfate ion
However, this can be minimized by adding calcium ions as calcium chloride, (CaCl2), to the pool water that shift the equilibrium to the
left, i.e. back to solid calcium sulfate, CaSO4.
Both forms have a high pH, and may require frequent additions of acid to maintain swimming pool water in the proper pH range for
chlorine sanitizing efficiency, equipment longevity and swimmer comfort.
The hypochlorite ions also establish equilibrium with hydrogen ions, depending on the pH.
The same relative amounts of HOCl and OCl- exist at equilibrium at a given pH if either chlorine gas or hypochlorites are used.
Chlorine decreases the initial pH, and hypochlorites increase the initial pH.
Neither product provides protection against the destructive effects of sunlight on a chlorine residual, so frequent chemical additions and
adjustments are necessary to maintain satisfactory water quality in outdoor swimming pools.