School Science Lessons
Topic 12E
2018-11-01
Please send comments to: J.Elfick@uq.edu.au

12E Acid-base indicators
See: Indicators, (Commercial)
Table of contents
12.11.3.10 Confirmation tests, original solution
12.11.4.0 Tests for cations, group analysis
12.11.6.0 Tests for metallic radicals
12.13.0 Tests for phosphorus and phosphates
12.11.7.0 Group tests instructions
12.11.3.0 Tests for an unknown substance
12.11.2.0 Titration, acid-base neutralization

12.11.3.10 Confirmation tests, original solution
See: Chemicals, (Commercial)
12.11.3 10 Tests for aluminium
12.11.3.11 Tests for ammonia, ammonium ions
12.11.3.12 Tests for antimony
12.11.3.13 Tests for barium
12.11.3.14 Tests for bismuth
12.11.3.15 Tests for cadmium
12.11.3.16 Tests for calcium
12.11.3.17 Tests for chromium
12.11.3.18 Tests for cobalt
12.11.3.19 Tests for copper
12.11.3.20 Tests for iron
12.11.3.21 Tests for lead
12.11.3.8 Tests for lead, heat charcoal with mixture
15.8.3 Tests for lead ions
12.11.3.22 Tests for magnesium
12.11.3.23 Tests for manganese
12.11.3.24 Tests for nickel
12.11.3.25 Tests for potassium
12.11.3.26 Tests for silver
12.11.3.27 Tests for sodium
12.11.3.28 Tests for strontium
12.11.3.29 Tests for tin
12.11.3.30 Tests for zinc

12.11.4.0 Tests for cations, prepare a solution
12.11.4.1 Group 1 tests for Ag+, Pb2+ , Hg+
12.11.4.2 Group 2 tests for Bi3+, Cd2+, Cu2+, Sn2+
12.11.4.3 Group 3 tests for Al3+, Cr3+, Fe2+, Fe3+
12.11.4.4 Group 4 tests for Co2+, Mn2+, Ni2+, Zn2+
12.11.4.5 Group 5 tests for Ba2+, Ca2+, Sr2+
12.11.4.6 Group 6 tests for K+, Mg2+, Na+, NH4+.

12.11.6.0 Tests for metallic radicals
12.11.6.0 Group tests for metallic radicals
12.11.6.1 Chemistry of group separations
12.11.6.2 Preliminary experiments before Group I
12.11.6.3 Separation into groups.

12.11.7.0 Group tests instructions
12.11.7.1 Group I Insoluble chlorides, PbCl2, AgCl, [Hg2Cl2 omitted]
12.11.7.2 Group II Sulfides insoluble in dilute hydrochloric acid
12.11.7.2a Group IIa PbS, Bi2S3, CuS, CdS, [HgS omitted]
12.11.7.2b Group IIb As2S3, Sb2S3, SnS, SnS2
12.11.7.3 Group III Insoluble hydroxides, Fe(OH)3, Cr(OH)3, Al(OH)3
12.11.7.4 Group IV Insoluble sulfides precipitated by hydrogen sulfide, ZnS, MnS, CoS, NiS
12.11.7.5 Group V Insoluble carbonates, CaCO3, BaCO3, SrCO3
12.11.7.6 Group VI Magnesium, sodium and potassium, Mg, Na, K

12.13.0 Tests for phosphorus and phosphates
12.13.5 Prepare microcosmic salt, Na.NH4.H.PO4.4H2O
12.13.2 Prepare phosphorus trichloride, PCl3
12.13.3 Prepare phosphorus pentachloride, PCl5
12.13.6 Reactions of phosphites, HPO32-
12.13.1 Reactions of phosphorus and phosphates, P
12.13.4 Water with chlorides of phosphorus, PCl3, PCl5

12.11.3.0 Tests for an unknown substance, qualitative analysis
12.11.3.2.1 Flame tests of salts
12.11.3.2.2 Flame test sprays
12.11.3.4 Heat substances, sublimation, melting, decrepitation
12.11.3.8 Heat substances with charcoal and fusion mixture
12.11.3.4.1 Loss of substance on heating indicates
12.11.3.9 Tests for aluminium compounds in solution
3.33.1 Tests for ammonia
12.11.3.1a Tests for metals with borax beads
12.11.3.2a Tests for metals with flame tests
12.11.3.1.1 Tests for potassium
12.11.3.3 Tests for solubility
12.11.3.5 Tests for substances with dilute hydrochloric acid
12.11.3.6 Tests for gases with hot concentrated sulfuric acid

12.11.2.0 Titration, acid-base neutralization
12.8.0 Acid-base neutralization, acid with base forms a salt and water
12.8.4.3 Acidity of vinegar and wine
12.8.1 Ammonia with sulfuric acid
12.8.5 Carbon dioxide affects acid-base titration
12.8.6 Heat of neutralization titration
12.8.7 Microscale titration, sodium hydroxide with dilute acids
12.8.7.1 Prepare monoprotic acid solution from unknown molarity acid
12.8.3 Simple titration of acids and bases
12.8.2 Sodium hydroxide with hydrochloric acid
12.8.4.0 Titrate dilute hydrochloric acid with sodium hydroxide solution, with a burette
12.8.4.1 Titrate dilute hydrochloric acid with sodium hydroxide solution, with a burette (second method)
12.8.4.2 Titrate dilute sulfuric acid with sodium hydroxide solution and isolate sodium sulfate crystals.

12.8.0 Acid-base neutralization, acid with base forms a salt and water
In neutralization reactions, an acid and a base react in such proportions as to form a neutral solution of a salt and water.
In the home, wool dresses spotted with another colour from acids or bases can be restored to original colour by neutralization.
The reaction is between the hydrogen ions and the hydroxide ions.
H+ (aq) + OH- (aq) --> H2O.

Experiments
1. Add 2 mL nitric acid or tartaric acid to 4 cm water in a test-tube.
Add blue litmus paper, which then turns red.
Add drops of dilute ammonia solution or washing soda solution, Na2CO3.10H2O, but stop when the solution just turns blue again.
The acid is now neutralized by the alkali.

2. Put 3 mL dilute sodium hydroxide solution on a watch glass.
Use a dropper to add dilute hydrochloric acid drop by drop while stirring continuously and testing the mixture with a fresh piece of
litmus paper after each drop is added.
You can get a mixture where the litmus paper is neither red nor blue but a tint midway between these two colours.
This mixture does not have the taste of an acid or the feel of an alkali.
The solution has the properties neither of an acid nor of an alkali.
The acid and alkali have neutralized each other.
Evaporate the solution to dryness by heating the watch glass over a beaker of boiling water.
A small quantity of solid appears on the watch glass.
This crystalline solid is sodium chloride, common salt.
Water is also a product of the reaction of sodium hydroxide with hydrochloric acid.

3. Pour about 5 mL of dilute solutions of sodium hydroxide, potassium hydroxide, calcium hydroxide, magnesium hydroxide and
ammonia solution into separate test-tubes.
Neutralize each alkali solution with dilute hydrochloric acid then evaporate the resulting solutions to dryness.
Do not taste the residues.
Repeat the procedure with dilute sulfuric acid, dilute nitric acid or dilute acetic acid.
HCl (aq) + NaOH (s) --> NaCl (aq) + H2O (l)
HCl (aq) + KOH (s) --> KCl (aq) + H2O (l)
2HCl (aq) + Ca(OH)2 (s) --> CaCl2 (aq) + 2H2O (l)
2HCl (aq) + Mg(OH)2 (s) --> MgCl2 (aq) + 2H2O (l)
HCl (aq) + NH3 (aq) --> NH4Cl (aq) + H2O (l)
[ H2O (l) + NH3 (aq) --> NH4+ (aq) + OH- (aq)].

4. Add magnesium hydroxide in small amounts to dilute sulfuric acid until excess solid is present.
Filter the mixture and test the filtrate with pieces of red and of blue litmus paper.
Evaporate the filtrate to dryness.
Many metallic hydroxides react with acids to produce water and a salt in the same way as alkalis do.
magnesium hydroxide (s) + sulfuric acid (aq) --> water (l) + magnesium sulfate (aq)
Metallic hydroxides that behave in this way with acids and are insoluble in water are called basic hydroxides.
The three classes of compounds (alkalis, basic oxides and basic hydroxides) are representatives of a group of substances called bases.
H2SO4 (aq) + Mg(OH)2 (s) --> MgSO4 (aq) + 2H2O (l).

12.8.1 Ammonia with sulfuric acid
Pass ammonia through sulfuric acid.
The common fertilizer ammonium sulfate or sulfate of ammonia forms.
2NH3 (g) + H2SO4 (aq) --> (NH4)2SO4 (aq) + 2H2O (aq)
H2SO4 (aq) + 2NH4OH (aq) --> (NH4)2SO4 (aq) + 2H2O (l).

12.8.2 Sodium hydroxide with hydrochloric acid
Put 10 drops of dilute sodium hydroxide solution on a watch glass.
Add drops of dilute hydrochloric acid and stir.
Test the mixture with litmus paper after adding each drop of hydrochloric acid.
When the litmus is neither red nor blue, but between the two colours, stop adding drops of acid.
Wet the tip of the finger with the mixture.
Rub the mixture between the fingers.
It does not feel slippery, so the solution is not alkaline.
When the correct quantities of hydrochloric acid and sodium hydroxide are mixed a solution forms that has the properties neither of
the acid nor of the alkali.
The acid and alkali have neutralized each other.
Evaporate the neutralized solution to dryness by heating the watch glass over a beaker of boiling water.
Crystals of sodium chloride appear on the watch glass.
HCl (aq) + NaOH (aq) NaCl (s) + H2O (l)
acid + alkali --> salt + water
Add dilute hydrochloric acid to dilute solutions of: sodium hydroxide, potassium hydroxide, calcium hydroxide aqueous ammonia
solution.
Evaporate to dryness.
Describe the salt formed.
Repeat the experiment with: dilute sulfuric acid, dilute nitric acid, dilute ethanoic acid (acetic acid).

12.8.3 Simple titration of acids and bases
Titration is an experimental method for measuring the concentration of a solution.
Measure the volume of the solution "A" needed to react with a given volume of solution "B".
For HCl and NaOH titration, molarity "A" X volume "A" = molarity "B" X Volume "B".
The end point in a titration occurs when an indicator changes colour.
Use a medicine dropper or a teat pipette as a simple burette.
The drops must always be the same size.
Within experimental error, when the same dropper is used, the same number of drops of alkali is needed to neutralize the same number
of drops of acid.
When the concentration of the acid is known, the concentration of the base can be estimated by comparing the numbers of drops of
acid and drops of base that just react.
Drop 100 drops of water from a medicine dropper into a measuring cylinder.
Calculate the volume of one drop.
Measure 25 mL of 2 M sodium hydroxide solution in the measuring cylinder and pour into an evaporating dish.
Add 2 drops of phenolphthalein solution and note the red colour of the indicator.
Wash the medicine dropper with the 2 M hydrochloric acid to get rid of remaining sodium hydroxide.
Add 2 M hydrochloric acid a drop at a time to the solution in the evaporating dish.
Stir as each drop is added.
Note the number of drops added until the colour just disappears completely.
Calculate the volume of added acid.
Heat the solution until almost dry.
Use gentle heat to avoid spattering.
Describe the appearance of the residue.
NaOH (aq) + HCl (aq) --> NaCl (aq) + H2O (l)
40 + 36.5 --> 58.5 + 18
Weight of NaOH in 25 mL of 2M solution = (25 X 2 X 40 / 1000) g.
The weight of sodium chloride expected in the evaporating dish = 58.5 X (25 X 2 X 40 / 1000) / 40 = 2.925 g.

12.8.4.0 Titrate dilute hydrochloric acid with sodium hydroxide solution, with a burette
See diagram 12.8.4: Titration
Pour a hydrochloric acid solution of known concentration (such as 0.11 mol / L) into a clear, dry burette until the liquid level is above
the "0" line.
Fix the burette vertically with a burette clamp.
Rotate the stopcock carefully to set the lowest point of the liquid meniscus exactly to "0" and to make simultaneously the tapered portion
of the burette full of the acid solution without any air bubbles in it.
Use a pipette to transfer 20 mL of the sodium hydroxide solution to a conical flask.
Add two drops of phenolphthalein to the flask.
The solution immediately turns red.
Stand the flask on a piece of white paper under the burette.
While adding drop by drop the acid solution from the burette, swirl the flask constantly so that mixing of the base and acid solutions is
rapid and thorough.
Note any change in the solution colour.
The neutralization is exactly completed and the end point occurs when half a drop or one drop of the acid solution turns the pale red
solution colourless in the flask immediately after swirling.
Stop the titration and record the burette meniscus reading.
Read the volume of the used hydrochloric acid solution.
Calculate the concentration of the sodium hydroxide solution according to the related chemical equation.
NaOH (aq) + HCl (aq) --> NaCl (aq) + H2O (l).

12.8.4.1 Titrate dilute hydrochloric acid with sodium hydroxide solution, with a burette (second method)
Use a burette containing 50 mL of 0.5 M sodium hydroxide.
Use a pipette to put 10 mL of 0.5 M hydrochloric acid in a beaker under the burette.
Add 2 drops of phenolphthalein to the beaker.
Stand the beaker on white paper under the burette containing sodium hydroxide.
Add a drop at a time of the sodium hydroxide from the burette and stir the beaker with a swirling motion.
Note the colour change when a drop of acid disappears after the solution is swirled.
The end point occurs when the drop does not change colour after swirling.
The solution is now neutral.
Test the neutral solution with litmus paper.
Pour 5 mL of the neutral solution into an evaporating dish.
Heat to dryness and weigh when cool.
NaOH (aq) + HCl (aq) --> NaCl (aq) + H2O (l).

12.8.4.2 Titrate dilute sulfuric acid with sodium hydroxide solution and isolate sodium sulfate crystals
Safety
1. Wear protective gloves when handling corrosive chemicals.
2. Handle acids and alkalis with care.
If any alkali gets into your eyes or onto your skin, report to your teacher immediately, and wash the affected area under running water
for at least three minutes.
3. If any acid gets into your eyes, report to your teacher immediately, and flush your eyes with running water for at least three minutes.
If any acid gets onto your skin, wash the affected area with plenty of water.
4. Eye protection must be worn
5. Sodium hydroxide 2.0 M is corrosive
6. Sulfuric acid 1.0 M is an irritant.

1. Wash a burette as follows:
1.1 Use a filter funnel to fill a burette with 20 mL of water.
1.2 Hold the burette horizontally and rotate it slowly to wash the inner wall.
Open the stopcock to run out all the water into the sink.
Close the stopcock.
Repeat the above procedure with 1.0 M sulfuric acid.

2. Fill the burette as follows:
2.1 Fill the burette too below the zero mark with 1.0 M sulfuric acid.
2.2 Clamp the burette vertically in a stand.
2.3 Open the stopcock for a few seconds to fill the burette jet completely with the acid.
Do not leave any bubbles in the burette jet.
2.4 Take the initial burette reading (V1) = mL (to two decimal places).

3. Wash a pipette as follows:
3.1 Using a pipette filler, suck sufficient water into a pipette to fill part of the bulb.
(Never use your mouth to suck a pipette.)
3.2 Hold the pipette horizontally.
Rotate it slowly so that the water washes the inner wall, up to the graduation mark.
3.3 Allow the water to run out into the sink.
Repeat 3.1 to 3.3 using 2.0 M sodium hydroxide solution.

4. Titration:
4.1 Using the pipette filler and the washed pipette, transfer 25.0 cm of the sodium hydroxide solution into a clean conical flask.
4.2. Add 2 drops of methyl orange indicator to the conical flask.
Note the colour of the solution in the conical flask.
4.3 Run the 1.0 M sulfuric acid into the flask while swirling the flask continuously.
4.4 Continue adding the acid, until the solution in the flask just turns to an orange colour.
4.5 Take the final burette reading (V2) = mL (to two decimal places) .
4.6 The volume of sulfuric acid needed to neutralize 25.0 mL of sodium hydroxide solution is (V2 - V1) mL.

5. Isolation of sodium sulfate crystals from solution:
5.1 Empty the conical flask and rinse it with water.
5.2 Use a pipette filler to pipette 25.0 mL of 2.0 M sodium hydroxide solution into the flask.
Do not add any indicator to allow preparation of pure sodium sulfate crystals.
5.3 From the burette, add the same volume of sulfuric acid as calculated from 4.6, i.e. (V2- V1) mL.
This results in complete neutralization to form a hot concentrated sodium sulfate solution and water.
5.4 Pour the resultant solution into a small beaker.
5.5 Boil the solution gently to concentrate it.
Heat with a small flame, until 1/3 of the solution is left.
Do not evaporate the solution until dry because spitting will occur.
5.6 Every 10 seconds, dip a glass rod into the boiling solution and then take it out.
If the immersed end becomes "cloudy" within 5 or 6 seconds, stop heating.
The solution has now become concentrated enough to deposit crystals on cooling.
5.7 Leave the solution to cool overnight.
5.8 The next day, sodium sulfate crystals should have formed.
5.9 Filter the crystals from the remaining solution.
5.10 Wash the crystals with drops of distilled water from a plastic wash bottle.
5.11 Use a spatula to transfer the crystals onto a piece of filter paper or absorbent paper.
5.12 Dry the crystals by gently pressing them between filter paper of absorbent paper.
5.13 Store the crystals in a labelled enclosed container.

12.8.4.3 Acidity of vinegar and wine
Vinegar must have a minimum of 5% acidity if it is sold.
The percent acidity, % acidity (percent acetic acid) = (grams of acetic acid, CH3COOH / grams of vinegar) x 100
CH3COOH (aq) + NaOH (aq)--> CH3COONa (aq) + H2O (l)
Titrate 5.00 mL of household vinegar with sodium hydroxide solution.
The molarity of the sodium hydroxide solution is stated on the bottle label.
The molar mass of acetic acid (ethanoic acid) = 60.05.
The density of vinegar = 1.00 g / mL.
One mole of acetic acid reacts with one mole of sodium hydroxide, so the number of moles of acetic acid =
molarity of the sodium hydroxide solution, in moles / litre X volume of sodium hydroxide solution, in litres.
Pour 5 mL of vinegar clean 250 mL Erlenmeyer flask and add 20 mL of distilled water and four drops of phenolphthalein indicator
solution.
Add sodium hydroxide solution from the burette while swirling the flask, until the solution is a faint permanent pink and record the volume.
Discards the solution if you overshoot the end point and the solution is red.
Moles of acetic acid = (molarity of the sodium hydroxide solution, in moles / litre) x (volume of sodium hydroxide solution, in litres)
Grams of acetic acid = (moles of acetic acid) x (molar mass)
Grams of vinegar = (volume of vinegar) x (density of vinegar)
% acetic acid = (grams of acetic acid / grams of vinegar) x 100.

Measure acidity using a titration kit
To measure T.A. in wine , use an inexpensive titration or acid test kit.
Test kits can be purchased cheaply and can be used over and over again.
The amount of acid in must or wine is measured by slowly adding a small amount of the reagent base sodium hydroxide until a change
in colour occurs from the indicator phenolphthalein.
Put a 15 cc sample (one cc equals one ml) of must or wine into a test-tube.
Most test tubes that come with the acid test kits are marked with a line indicating this volume, or use a small plastic syringe in the test
kit to measure the desired amount into the test tube.
Then rinse the syringe.
Put 3 drops of the phenolphthalein solution into the test-tube.
Swirl or shake the test tube to mix the indicator with the must or wine.
Use the syringe to draw out 10 cc of the sodium hydroxide reagent, with no bubbles in the liquid.
Be careful! Avoid contact of the sodium hydroxide solution with your skin or eyes.
Add the sodium hydroxide solution to the test-tube, by 0.5 cc at a time.
After each addition of 0.5 cc, swirl or shake the test-tube to mix the contents.
The colour of the liquid will momentarily change upon the addition of the reagent.
If testing white wines, the colour change will be pink.
If testing red wines, the colour change will be grey.
Keep swirling the test-tube until the colour subsides.
If the colour of the must or wine returns to the original colour, repeat adding 0.5 cc until the colour change is permanent.
So when the colour, pink or grey, does not go away, stop and record the amount of reagent used.
To determine the acidity of the must or wine, for each cc of reagent used, equals 0.1 % TA.
For example, if you used 6 cc of sodium hydroxide to react with the must or wine , the titrateable acidity is 0.6 %.
Discard the sample because it is now toxic.
Do not add the sample back into your original must or wine!
Wash and dry the test equipment before storing it.

12.8.5 Carbon dioxide affects acid-base titration
Add 2 drops phenolphthalein to 100 mL of deionized water.
Add 2 drops of 0.1 M sodium hydroxide.
The reaction forms a red colour.
Swirl vigorously for one minute.
The red colour fades because of absorption of carbon dioxide from the air.

12.8.6 Heat of neutralization titration
The end point occurs at maximum temperature.
Use 25 mL of dilute sodium hydroxide solution.
Note the original temperature.
Add 1 mL of 2 M hydrochloric acid, stir with a thermometer and note the temperature.
Continue to add 1 mL of the acid and note the temperature.
Use graph paper to plot temperature rise against volume of acid added.
Read from the graph the maximum temperature rise and volume of acid that neutralized the sodium hydroxide solution.
Calculation: 25 X 1 / 1000 X concentration of sodium hydroxide = volume of HCl X 1 / 1000 X 2.

12.8.7 Microscale titration, sodium hydroxide with dilute acids
See diagram 12.8.7: Microscale titration apparatus
Use a sodium hydroxide solution to titrate a standardized acid solution.
Use the sodium hydroxide solution to titrate an unknown acid.
Calculate the concentration of the unknown acid.
Conventional titration vs microscale titration:
Conventional titration requires burettes, bulb pipettes and litres of solutions.
Ten microscale titrations will use fewer solutions than used for one conventional titration.
Burettes are large and fragile so spillage and breakage of glassware occur sometimes.
When an operator using conventional titration does three measurements, a microscale operator can do six to ten measurements.
Microscale titration does require different hand and finger skills than conventional titration and involves some differences in the
calculation methods.

12.8.7.1 Prepare monoprotic acid solution from unknown molarity acid
1. Use two 2 mL graduated glass pipettes, graduated to 0.01 mL.
Attach disposable pipette tips.
If they fall off the ends of the pipettes, seal with silicone putty.
Attach a 5 mL plastic syringe to each pipette with silicone tubing.
Lubricate the syringe with glycerol.
Transfer 15 mL of acid solution into a wide neck bottle and one drop of phenolphthalein indicator.
Transfer 15 mL sodium hydroxide into another wide neck bottle.
Clamp the two pipettes on a stand with a double clamp.
Rinse then fill the pipettes by drawing solution up into the pipette with the syringe.
Record pipette volumes to 0.001 mL.
Use a 10 mL Erlenmeyer flask containing water for comparison when detecting the faint pink of the endpoint.
Use pipette volumes in excess of 1.000 mL to provide four significant figures in the volume measurements.

2. Pour 1.25 mL of 0.1 M monoprotic acid into a 10 mL Erlenmeyer flask.
Add 0.1 M sodium hydroxide until the colour changes to a faint pink.
Record the final pipette volumes.
Drop volume is less than 0.02 mL.
Pour 1.25 mL of the unknown acid into a 10 mL Erlenmeyer flask.
Add 0.1 M sodium hydroxide until the colour changes to a faint pink.
Record the final pipette volumes For each titration, calculate the ratio of acid volume to alkali volume to allow concordance between
titrations.
Discard discordant values and calculate the mean for accepted values.
Ratio 1 = volume 0.1 M monoprotic acid / volume 0.1 M sodium hydroxide
Ratio 2 = volume unknown acid / volume 0.1 M sodium hydroxide
Concentration of unknown acid = (Ratio 1 / Ratio 2) X concentration 0.1 M monoprotic acid, e.g. If Ratio 1 = 1.158, Ratio 2 = 1.079,
molarity of unknown acid = (1.158 / 1.079) X 0.1 = 0.1073 M.
Addition of a drop of the indicator to the acid solutions being measured dilutes the acids, e.g. 0.02 mL of indicator solution added to
15 mL of acid solution dilutes the acid to affect the measured concentration at the fourth significant figure.
However, this titration is done twice, with the "known" standard acid and with the "unknown" acid, so the errors will cancel at the fourth
significant figure when the acids do not differ greatly in concentration.

12.11.3.1a Tests for metals with borax beads,
The test depend on the colours of metal oxides in metallic salts and minerals when heated.
Coil the end of a platinum wire to prepare a loop big enough to curl around a match stick.
Heat the loop in a Bunsen burner flame then dip the hot loop in borax powder, Na2B4O7.10H2O.
Heat the attached borax in the hottest part of the flame so that it swells to lose water of crystallization then shrinks and fuses to form a
transparent bead filling the loop of the platinum wire.
Dip the bead in water then in a very small amount of the metal salt that sticks to the bead.
Heat the bead and metal salt in either the outer oxidizing flame or inner reducing flame of the Bunsen burner.
Use a white ceramic tile of saucer instead of platinum wire and microsmic salt, [sodium ammonium hydrogen phosphate,
(Na(NH4)HPO4.4H2O) from urine] instead of borax.
Use sodium carbonate, Na2CO3, for the bead test, but the colours may be different.
The colour of the bead depends on whether you use the outer colourless oxidizing flame or the inner blue reducing flame, the type of
flame used, whether you examine the bead hot or cold and the degree of saturation if you are using solutions, the temperature of the
bead when examined, and the concentration of the solution if a solution of salt is used instead of the solid salt.
For most metals not listed below, the flame is colourless.
Table 12.11.3.1
Solid metal salt Oxidizing flame Oxidizing flame Reducing flame Reducing flame
.
Hot Cold Hot Cold
Chromium yellow to emerald-green
green green green
Cobalt deep blue blue blue blue
Copper light blue-green blue .
red-brown
Gold rose-violet
rose-violet
red violet
Iron brown when hot
yellow when cold
yellow green green
Manganese violet amethyst .
.
Nickel violet to brown
red-brown grey grey
Tungsten pale yellow .
green blue


12.11.3.1.1 Tests for potassium, tetraphenylborate test
To 15 drops of the solution, add 5 drops of 1 M NaOH, boil the solution, add 2 drops of 6 M HCl and 15 drops of 1 M sodium
acetate.
Add 3 drops of 3% sodium tetraphenylborate, NaB(C6H5)4.
If a white precipitate forms the test is positive.

12.11.3.2a Tests for metals with flame tests
Cations in an unknown solution can be identified by using flame tests.
Add drops of concentrated hydrochloric acid to the solution.
Dip a clean piece of platinum wire into it then hold it in a Bunsen burner flame.
Dip platinum wire into concentrated hydrochloric acid (12 M) then into powdered solid and heat in a non-luminous edge of a Bunsen
burner flame.
When a salt is heated in the flame, it dissociates into neutral atoms and electrons are excited into a higher energy level then return to the
ground state and emits light of characteristic colour for that atom.
Remember that each person observes colours differently.
The non-metal atoms in anions do emit light but at wavelengths shorter than ultraviolet light, so we cannot see the colours.

Experiments
Check the flame test colours by doing the test for all the cations against a dark background.
Compare with the following list of colours:
Ammonium compounds: green (faint colour).
Antimony: blue-green to light blue (faint colour)
Arsenic: light blue (moistened with hydrochloric acid)
Barium: pale green to yellow-green:
Bismuth: blue
Calcium: red
Calcium compounds: brick-red to yellow (masked by barium)
Copper: blue-green
Copper compounds: green (not halides) (CuBr2 blue-green)
Lead: light blue to blue
Lithium compounds: crimson (masked by barium or sodium)
Molybdenum: yellow-green:
Phosphates: blue-green (if when moistened with sulfuric acid).
Potassium: lilac but crimson through blue glass, violet through cobalt glass
Potassium compounds: pink-lilac to violet (not borates, phosphates, and silicates.) (masked by sodium or lithium).
Selenium: blue
Sodium: strong golden yellow but no colour viewed through blue glass
Sodium compounds: yellow (if the yellow flame persists and is not intensified by adding 1% NaCl to the dry compound.)
Strontium: crimson
Strontium compounds: scarlet (masked by barium)
Zinc: green-white.

12.11.3.2.1 Flame tests of salts
Soak paper in the following salts, leave to dry then ignite
Calcium chloride: orange, Copper (II) chloride: blue, Copper (II) sulfate: green, Lithium chloride: red, Potassium chloride: purple,
Sodium borate, borax: green.
Sodium carbonate: yellow, Sodium chloride: yellow, Strontium chloride: red.
12.1.29 Flame tests for sodium chloride.

12.11.3.2.2 Flame test sprays
Be careful! Wear eye protection.
Use spray bottles, e.g. window cleaners or garden sprays, to spray saturated solutions of metal salts in ethanol onto roaring Bunsen
burner flames in a darkened room.
The salts can include sodium chloride, potassium chloride, lithium chloride, and copper sulfate.
The spray bottles should have a trigger mechanism and not a scent bottle spray pump, which may allow flash back.

12.11.3.3 Tests for solubility, prepare a solution for group analysis
See 3.71.1: Solubility table and solubility rules
Dissolve 1 g of the substance in the first reagent below that can dissolve the substance .

1. Water: Try to dissolve the salt in deionized water.
If the salt does not dissolve, heat it in a test-tube to observe if the salt dissolves in hot water.

2. Dilute hydrochloric acid: If the salt does not dissolve in hot water, add dilute hydrochloric acid to observe if it dissolves.

3. Concentrated hydrochloric acid, 2.0 to 5.0 mL: When all substance is dissolved, dilute solution to five times its bulk, then leave to cool.
If the dilution produces a precipitate because of hydrolysis of chlorides of bismuth, antimony or tin, add drops of concentrated
hydrochloric acid.

4. Dilute nitric acid: It may dissolve compounds of lead, silver and mercury, but avoid using nitric acid because it may oxidize hydrogen sulfide.

5. Aqua regia, Be careful! Heat with concentrated hydrochloric acid, then add a few drops of concentrated nitric acid.
Dilute as in 3. above.

6. Concentrated nitric acid: Warm with 2.0 to 5.0 mL of the acid.
The solution must be evaporated to dryness and the residue dissolved in water or hydrochloric acid.

7. Insoluble residue: Filter it off, wash it, and fuse it in a crucible with four times its bulk of fusion mixture.
Cool, boil with water, filter.
Test filtrate for acid radicals.
Dissolve precipitate of metal carbonates in hydrochloric acid, and analyse separately.
If the solution so obtained gives no precipitate with that obtained by 1. to 6. above, analyse them together.

12.11.3.4 Heat substances, sublimation, melting, decrepitation
1. Sublimation indicates the presence of ammonium halides, other halides, and some oxides.
2. Melting indicates the presence of sodium, potassium or ammonium nitrate, potassium chlorate, and other less common substances.
3. Decrepitation (crackling noise of some heated crystals), indicates the presence of sodium chloride, lead nitrate and potassium chlorate.

12.11.3.4.1 Loss of substance on heating indicates:
1. Loss of water vapour indicates the presence of water of crystallization and some basic hydroxides, basic carbonates, and acid salts.
Also, it may indicate that the substance has absorbed moisture from the atmosphere.
2. Loss of oxygen indicates oxides of silver, peroxides, sodium or potassium nitrate, permanganates and chlorates.
3. Loss of carbon dioxide indicates carbonate or bicarbonate.
4. Loss of ammonia indicates an ammonium compound.
5. Loss of nitrogen dioxide (dinitrogen tetroxide, N2O4), indicates nitrates of heavy metals, e.g. copper, lead, zinc.
6. Loss of sulfur trioxide indicates some sulfates.
7. Loss of halogens indicates oxidation of a halide.

12.11.3.5 Tests for substances with dilute hydrochloric acid
1. Carbon dioxide produced indicates a carbonate or bicarbonate
2. Hydrogen gas produced indicates some free metals
3. Sulfur dioxide produced indicates sulfite or bisulfite
4. Sulfur dioxide and sulfur produced indicates thiosulfate
5. Hydrogen sulfide produced indicates sulfide
6. Nitrogen dioxide (N2O4) produced indicates nitrite
7. Chlorine produced indicates hypochlorite or oxidizing agent.

12.11.3.6 Tests for gases with hot concentrated sulfuric acid
If organic acid present, substance should be ignited, extracted with dilute hydrochloric acid and filtered before proceeding with main
group separation.
Gas evolved indication
1. Hydrogen chloride produced indicates chloride.
2. Nitric acid produced indicates nitrate.
3. Oxygen produced indicates peroxide, permanganate, chromate, dichromate.
4. Chlorine peroxide (yellow-green gas, violent action) produced indicates chlorate.
5. Sulfur dioxide produced indicates sulfite, thiosulfate or reducing agent.
6. Hydrogen bromide, bromine and sulfur produced indicates bromide.
7. Hydrogen iodide, iodine and hydrogen sulfide produced indicates iodide.
8. Carbon monoxide and carbon dioxide produced indicates oxalate.
9. Carbon monoxide only produced indicates formate.
10. Acetic acid produced indicates acetate.

12.11.3.8 Heat substances with charcoal and fusion mixture
Test substances with charcoal, heat charcoal with fusion mixture, note heated metal appearance
1. Aluminium produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a blue mass but this is also caused by fusible phosphates, arsenates,
borates and silicates.
2. Arsenic produces fumes smelling of garlic and forms a white crust if seen at some distance from the flame.
3. Bismuth forms pink globules, becomes brittle and forms a yellow crust.
4. Cadmium forms a brown crust.
5. Copper forms red scales.
6. Lead forms grey-white soft globules and forms a red crust when hot and a yellow crust when cold.
7. Magnesium produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a green mass.
8. Silver has shining metal particles.
9. Tin forms hard white beads.
10. Zinc forms a yellow crust when hot and a white crust when cold.
Zinc also produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a green mass.

12.11.3.11 Tests for ammonia, ammonium ions
Nessler's reagent
1. Tests for ammonia with damp red litmus paper
2. To test for ammonia, shake vigorously and smell.
3. Wet red litmus paper with deionized water and stick it to the concave side of a large watch glass.
Put 1 mL of 0.1 M NH4Cl in a smaller watch glass or a small beaker, heat it then add 6 M NaOH.
Cover the small watch glass or beaker with the large watch glass.
Ammonia vapour will turn the red litmus blue.
4. Nessler's reagent is not permitted in some school systems because it contains mercury.
Add Nessler's reagent to the original solution.
A brown colour indicates the presence of ammonium ion.
The test is very sensitive, so the test solution must not contain any NH4+ added during the analysis.
5. Note: NH3 (aq) is preferred as there does not seem to be any detectable NH4OH present.
Certainly there are some NH4+ ions and OH- ions.
Thus a reaction with ammonia solution is written as follows:
X+ (aq) + OH- (aq) --> XOH (s)
or as molecules:
XCl (aq) + NH3 (aq) + H2O (l) --> XOH (s) + NH4Cl (aq).

12.11.3.12 Tests for antimony
1. Dilute with its own volume of water.
Pass H2S.
An orange-red precipitate of antimony sulfide, Sb2S3, indicates the presence of antimony.
2Sb3+ + 3S2- --> Sb2S3 (s).

2. Organic reagent: Gallocyanine (Fast violet), C15H13ClN2O5, 0.05% in M HCl.
To one drop of antimony solution on filter paper, add one drop of reagent.
A colour change from wine red to blue indicates the presence of antimony.
Use Group IIb precipitate dissolved in concentrated HCl and diluted.

12.11.3.13 Tests for barium
Confirm by flame test: Light green
Barium and strontium
Organic reagent: Rhodizonic acid, C6H2O6, [(CO)4(COH)2],
1, 2-dihydroxycyclohexene-3, 4, 5, 6-tetrone, dihydrate: C6H2O6.2H2O, 0.1% aqueous solution
Put one drop of test liquid on filter paper then add one drop of reagent.
A red-brown spot indicates the presence of Sr and Ba.
When one drop of dilute HCl is added, a barium spot is intensified and a Sr spot disappears.
Use Group V precipitate after solution in dilute acetic acid.
Prepare fresh solution of reagent if it has decolorized.

12.11.3.14 Tests for bismuth
Organic reagent: Thiourea, H2N.CS.NH2, 10% aqueous solution (10 mL Bi solution + 10 mL dilute HNO3 + 1 mL reagent)
A yellow colour indicates the presence of bismuth.

12.11.3.15 Tests for cadmium
1. Ammonium hydroxide gives white precipitate easily soluble in excess.

2. Organic reagent: Diphenyl carbazide, C13H14N4O, CO(NH.NH.C6H5)2, in saturated alcoholic solution
Add few drops of reagent to Cd to give violet coloration.
Use solution in dilute HNO3 in Group separation.
If Cu present as a blue solution, first saturate reagent with KCNS and add crystal KI, then Cu is reduced and does not interfere.

12.11.3.16 Tests for calcium
Confirm by flame test: Brick-red (green through blue glass).

12.11.3.17 Tests for chromium
1. Fuse with sodium carbonate and a little potassium nitrate in a porcelain crucible.
Dissolve in water, add acetic acid and lead acetate solution.
A yellow precipitate forms.
A filtrate may contain chromium and aluminium as sodium chromate and sodium aluminate.
A yellow precipitate indicates the presence of chromium.
Pb2+ + CrO42- --> PbCrO4 (s), [yellow lead chromate].

2. Chromium (as chromate}
Organic reagent: Diphenyl carbazide, C13H14N4O, CO(NH.NH.C6H5)2, 0.2% solution in one part glacial acetic acid and nine parts
methylated spirit.
Make the chromate solution acidic with acetic acid or sulfuric acid.
Add reagent.
A deep violet-red colour indicates the presence of chromate.
Use in Group III when in form of chromate.

12.11.3.18 Tests for cobalt
Organic reagent: Nitroso-beta-naphthol, 1-nitroso-2-naphthol, C10H7NO2, 1 g in 50 mL acetic acid
Dilute to 100 mL Add reagent to neutral or slightly acid solution.
A brown colour indicates the presence of cobalt.
Use in Group IV when in solution after treatment with KClO3 and acid, or use the solution after Group III.
Cu, Fe, Sn, Ag, Cr, Bi, all interfere with the test.

12.11.3.19 Tests for copper
1. Ammonium hydroxide gives a pale blue precipitate that dissolves in excess to give a deep blue solution.

2. Organic reagent: Rubeanic acid (ethanedithioamide, dithiooxamide), NH2.CS.CS.NH2, saturated 0.5% alcoholic solution
Use 10 mL of neutral Cu solution + 1 mL 5M CH3COOH + drops of reagent.
A green-black precipitate forms.
Test with Group II precipitate.
Ni and Co may interfere with the test.
Dissolve CuS in dilute HNO3 and neutralize with NaOH solution.

12.11.3.20 Tests for iron
1. Dissolve ammonium thiocyanate in water and heat the solution.
The solution turns a characteristic red colour with iron (III) compounds, ferric compounds.

2. Organic reagent: Cupferron
NH4[C6H5N(O)NO] (ammonium salt of N-nitroso-N-phenylhydroxylamine), 5% aqueous solution
Filter if reagent is turbid.
Add reagent to strongly acidic HCl solution.
A red-brown compound indicates the presence of iron.
The reagents is unstable over long periods but decomposition may be delayed by a piece of solid ammonium carbonate added to the reagent.

12.11.3.21 Tests for lead
Add potassium iodide solution to solutions of lead salts to form a yellow precipitate that is soluble in boiling water.
Organic reagent: Rhodizonic acid, C6H2O6, (CO)4(COH)2, dihydrate: C6H2O6.2H2O, sodium salt (CO-CO.C.ONa)2, 0.1% aqueous solution
Add two drops of reagent to a sample of Group I precipitate still wet with acid.
A violet colour indicates the presence of lead.
Make a fresh solution of the reagent, if it has decolorized.

12.11.3.22 Tests for magnesium
1. Heat on charcoal with sodium carbonate.
Add a few drops of cobalt nitrate solution and heat again to produce a pink residue.

2. Organic reagent: The complex dye Titan yellow
Na2C28H19S4O6, as 0.1% aqueous solution.
Add 2 mL 1% KOH to 2 drops of test solution.
Boil to remove NH4+ and add 2 drops of Titan yellow.
A red colour or red precipitate indicates the presence of magnesium.
Tests for Mg in a Group VI solution.
Ammonium ions interfere with the test and must be removed.

12.11.3.23 Tests for manganese
1. Fuse manganese with sodium carbonate and some potassium nitrate in a crucible to form a blue-green mass.
2Mn(OH)2 + 5O --> 2MnO4- + 2H+ + H2O, [5O from oxidizing agents], [MnO4- = purple permanganate ion].

2. Organic reagent: Benzidine
C12H12N2, 0.05% solution in 10% acetic acid
To one drop of solution on filter paper add one drop 0.05% NaOH then one drop of reagent.
Use in Group IV when in solution in dilute acid.
Dissolve Group IV precipitate in very dilute acid and use the solution, rejecting any undissolved solid.

12.11.3.24 Tests for nickel
Organic reagent: Dimethylglyoxime, DMG (CH3C(NOH)C(NOH)CH3), 1% solution in methylated spirit, Toxic if ingested.
Warm a slightly acid test solution, add reagent then ammonium hydroxide until solution is alkaline.
A bright red precipitate indicates the presence of nickel.
Bismuth interferes with the test.

12.11.3.25 Tests for potassium
Reagent: Sodium perchlorate, 20% solution in equal parts of water and alcohol
Add reagent to equal volume of test solution.
A white precipitate of KClO4 indicates the presence of potassium.
Test in Group VI solution concentrated by evaporation and let cool.
Do the test on a glass plate above a black background.

12.11.3.26 Tests for silver
1. Add potassium chromate solution to neutral solution of a silver salt to form a brick-red precipitate.

2. Organic reagent: 5-(4-dimethylaminobenzylidene) rhodanine
C12H12N2OS2, 0.03% in acetone
The reagent detects AgCl in solution in water.
A red colour indicates the presence of silver.
Use a Group I precipitate.

12.11.3.27 Tests for sodium
Organic reagent: Uranyl magnesium acetate.
UO2(CH3COO)2.Mg(CH3COO)2, as a saturated aqueous solution
Add reagent to cold solution.
A yellow precipitate indicates the presence of sodium.
Test in Group VI solution concentrated by evaporation and let cool.
Do the test on a glass plate above a black background.

12.11.3.28 Tests for strontium
1. Confirm by flame test: Crimson
Sr2+ + SO42- --> SrSO4 (s).

2. Strontium and barium
Organic reagent: Rhodizonic acid (CO-CO.CONa)2, 0.1% aqueous solution
Put one drop of test liquid on filter paper then add one drop of reagent.
A red-brown spot indicates the presence of Sr and Ba.
When one drop of dilute HCl is added, a barium spot is intensified and a Sr spot disappears.
Use Group V precipitate after solution in dilute acetic acid.
Prepare fresh solution of reagent if it has decolorized.

12.11.3.29 Tests for tin
1. Use a borax bead containing some copper (II) sulfate.
Add a sample of the original solid and heat again to produce a red bead.

2. Organic reagent: Cacotheline
C21H21O7N3, saturated solution in water
Tin must be as Sn (II) in M HCl.
Add drops of reagent.
A violet colour indicates the presence of tin.
Stability of reagent is about 14 days.
Cu, Ni, Co, Cr and Fe interfere with the reaction.
Do the test with Group IIb solution when as Sn (II).

12.11.3.30 Tests for zinc
Filter and dissolve the precipitate in concentrated nitric acid.
Add a little cobalt nitrate solution, evaporate to concentrate, and soak a filter paper in the mixture.
Ignite the filter paper.
A green ash (Rimann's green) indicates the presence of zinc.
The green ash is a compound of zinc and cobalt oxides.
2ZnO22- + 8H+ + 2S2- --> 2ZnS (s) + 4H2O.

12.11.4.1 Group 1 tests for Ag+, Pb2+, Hg+
The insoluble chloride group
1. These metals form very insoluble chlorides as white precipitates that are easily filtered out of solution.
First test: Add 2 drops of dilute hydrochloric acid to 5 drops of the original solution.
A white precipitate identifies Ag+ or Pb2+or Hg+.
If no precipitate forms, go to (Group 2 test).
Pb2+ + 2Cl- --> PbCl2 (s)
Lead chloride dissolves in hot water and form a bright yellow precipitate with potassium chromate solution.
Ag+ + Cl- --> AgCl (s)
Silver chloride dissolves in ammonia but precipitates again when dilute nitric acid is added to the solution.
This precipitate turns black in the light.
Hg+ + Cl- --> HgCl (s)
Add ammonium hydroxide solution to make the mercurous chloride, mercury (I) chloride, turn black.

2. Second test: Add 2 to 5 drops of K2CrO4 solution to 5 drops of the original solution.
A red precipitate identifies Ag+.
A yellow precipitate identifies Pb2+.
Pb2+ + CrO42- --> PbCrO4 (s).

12.11.4.2 Group 2 tests for Bi3+, Cd2+, Cu2+, Sn2+
1. First test: Add 10 drops of H2S solution.
A yellow precipitate identifies Cd2+.
A brown-black precipitate identifies Cu2+ or Bi3+ or Sn2+.
Then add 10 drops of NaOH solution and heat.
A dissolving precipitate identifies Sn2+.
Cd2+ + S2- --> PbS (s)
2Bi3+ + 3S2- --> Bi2S3 (s).

2. Second test: Add 10 drops of NH3 solution to 5 drops of the original solution.
A deep blue colour identifies Cu2+.
A white precipitate that then dissolves identifies Cd2+.
Add water to the original solution.
A milky precipitate identifies Bi3+.
Cu2+ + 4NH4OH --> [Cu(NH3)4]2+ + 4H2O (deep blue solution)
Cd2+ + 4NH4OH --> [Cd(NH3)4]2+ + 4H2O
Bi3+ + 3OH- --> Bi(OH)3 (s).

12.11.4.3 Group 3 tests for Al3+, Cr3+, Fe2+, Fe3+
1. First test: Add 8 drops of NH4Cl solution to 5 drops of the original solution.
Test the solution with litmus paper.
Add enough drops of dilute NH3 solution to turn red litmus blue.
A green precipitate identifies Fe2+ or Cr3+.
A red-brown precipitate identifies Fe3+.
A white glassy precipitate identifies Al3+.
If no precipitate forms, go to (Group 4 test).

2. Second test: Add to 5 drops of the original solution 6 drops of NaOH then 2 drops of NaHClO solution boil then add 2 drops of
lead (II) ethanoate (lead (II) acetate solution).

3. Use lead acetate test paper.
A yellow precipitate indicates Cr3+.
BE CAREFUL! WARNING!
The following common test is too dangerous to be used in school science experiments:
Potassium hexacyanoferrate (III) (potassium ferricyanide) reacts with strong mineral acids to release toxic potassium cyanide.

4. Add to 5 drops of the original solution 5 drops of K3Fe(CN)6 solution.
A deep blue precipitate identifies Fe2+.
A green-brown precipitate identifies Fe3+.

12.11.4.4 Group 4 tests for Co2+, Mn2+, Ni2+, Zn2+
1. First test: Following Group 3 test, add to 5 drops of original solution 10 drops of H2S solution.
A black precipitate indicates Co2+ or Ni2+.
A dirty white precipitate indicates Zn2+. A pink precipitate indicates Mn2+.
Co2+ + S2- --> CoS (s)
Ni2+ + S2- --> NiS (s)
Zn2+ + S2- --> ZnS (s)
Mn2+ + S2- --> MnS (s).

See 16.3.9: Diacetyl, butanedione
2. Second test: Following Group 3 test, add to 5 drops of original solution 2 drops of sodium hydroxide solution and heat.
A blue to pink precipitate indicates Co2+.
Add to 5 drops of original solution 2 drops of butanedione (dimethylglyoxal).
A bright red precipitate indicates Ni2+.

3. Third test: Following Group 3 test, add to 5 drops of original solution drops of dilute NH3 solution until litmus paper test indicates
it as basic.
A white glassy precipitate that dissolves in excess NH3 solution indicates Zn2+.

4. Fourth test: Following Group 3 test, add to 5 drops of original solution 2 drops of NaOH solution then 2 drops of NaClO solution.
A dirty white precipitate indicates Mn2+.

12.11.4.5 Group 5 tests for Ba2+, Ca2+, Sr2+
1. First test: Add 8 drops of NH4Cl solution to 5 drops of original solution.
Test with litmus paper.
Add enough drops of dilute NH3 solution to turn blue litmus red.
Add 10 drops of H2S solution - no precipitate forms.
Add 8 drops of NH4Cl solution to 5 drops of original solution.
Add 2 drops of concentrated aqueous ammonia solution then 2 drops of ammonium carbonate solution.
A white precipitate indicates Ba2+ or Ca2+ or Sr2+.

2. Second test: Add to 5 drops of original solution 2 drops of saturated CaSO4 solution.
No precipitate indicates Ca2+.
Immediate white precipitate indicates Ba2+.
Slow precipitate on heating indicates Sr2+.

12.11.4.6 Group 6 tests for K+, Mg2+, Na+, NH4+
1. There is no test for Group 6.

2. Verify K+ and Na+ with flame tests.
To verify cations, add to 5 drops of original solution 3 drops of NaOH and boil.
The sharp odour of NH3 identifies NH4+.
NH3 turns red litmus blue.
Add 4 drops concentrated NH4Cl solution to 5 drops of the original solution.
Add enough aqueous ammonia solution to prepare the solution test alkaline.
Add 4 drops of sodium phosphate (trisodium phosphate (V)-12-water) solution then stir.
A white precipitate indicates Mg2+.
Mg2+ + HPO42- + NH4+ + OH- --> MgNH4PO4 (s) (magnesium ammonium phosphate).

12.11.6.0 Group tests for metallic radicals
See 17.5.7.0: Explanation of group analysis
In this version of qualitative analysis, all reference to the use of mercury salts has been deleted because these salts are not allowed in
school science experiments.

12.11.6.1 Chemistry of group separations.
This system of qualitative analysis is based on differences between the properties of 23 metallic radicals.
The first separation into Groups is made by selecting a reagent, which precipitates a few of the metals but leaves the remainder to be
precipitated by a different reagent or reagents.
A precipitate obtained in a given group may contain one or more of the metals of that group and it is then that use is made of specific
differences of properties of the metals likely to be present.
Before each Group separation show how each metal in the group behaves towards the reagents used.

12.11.6.2 Preliminary experiments before the separation of Group I metals, silver and lead.
1. Add dilute hydrochloric acid to solutions of silver nitrate and lead acetate.
The white precipitates are the insoluble metal chlorides, AgCl and PbCl2.
Ag+ + Cl- --> AgCl (s)
Pb2+ + 2Cl- --> PbCl2.

2. Decant the liquid from each test-tube, then half fill with deionized water.
Heat the solutions to boiling to show that lead chloride is soluble in boiling water, but silver chloride is not soluble.
(If the solution is not boiled, some chromium may be present as a pale purple amine solution, Cr(NH3)6(OH)3, that decomposes to
chromium hydroxide on boiling.
An amine is a coordination complex in which the molecules that donate a pair of electrons to a metal are ammonia molecules.)
Add drops of potassium chromate to the clear lead chloride solution to form the yellow precipitate of lead chromate.
Pb2+ + CrO42- --> PbCrO4.

3. Decant the liquid from the precipitate of silver chloride then add dilute ammonium hydroxide solution and shake the test-tube.
Silver chloride dissolves because of the formation of a soluble silver amine complex ion.
AgCl + 2NH4OH --> Ag(NH3)2Cl + 2H2O
or AgCl + 2NH4+ + 2OH- --> Ag(NH3)2+ + Cl- + 2H2O.

12.11.6.3 Separation into groups
1. Add dilute hydrochloric acid to the prepared solution then filter when cold.
However, if effervescence of a gas occurs, pass the gas into lime water.
Turbidity indicates the presence of a carbonate or bicarbonate.

2a. The precipitate may contain AgCl, PbCl2 or Hg2Cl2 as a white precipitate.
See 12.11.7.1: Observe for Group I - Insoluble chlorides
2b. The filtrate should give no precipitate on adding one drop of HCl.
Heat the filtrate then pass hydrogen sulfide until no further precipitate forms.
Filter.
Dilute some filtrate five times with water, then pass more hydrogen sulfide through the filtrate.
If more precipitate appears, dilute the filtrate and continue passing more hydrogen sulfide through the filtrate until no more precipitate
forms.
Filter through the same filter paper.
If no precipitate occurs on dilution, discard the diluted portion.

3a. The precipitate may contain HgS, PbS, Bi2S3, CuS, CdS, As2S3, Sb2S3, SnS, SnS2.
See 12.11.7.2: Observe for Group II - Sulfides insoluble in dilute hydrochloric acid
3b. Put the filtrate into an evaporating dish and boil off the H2S and allow solution to concentrate.
Add dilute nitric acid and excess ammonium molybdate solution to a sample of the solution.
Heat the solution but do not boil, then leave to stand.
Discard any yellow precipitate that indicates phosphate.
To the rest of the main solution add 2 mL of concentrated nitric acid and boil to oxidize iron (II) iron to iron (III) iron.
Transfer to a boiling tube, add ammonium chloride solution and ammonium hydroxide solutions in excess.
Boil and filter the solution.

4a. The precipitate may contain Fe(OH)3, Al(OH)3, Cr(OH)3.
See 12.11.7.3: Observe by table for Group III - Insoluble hydroxides, Fe(OH)3, Cr(OH)3, Al(OH)3
4b. The filtrate should contain excess ammonia.
To test for ammonia, shake vigorously and smell.
Pass hydrogen sulfide.
Boil, filter.
If the filtrate is brown, nickel is present in the solution.
Boil the filtrate until no more ammonia is evolved.
Filter.
The filtrate should not now be brown.

5a. The precipitate may contain ZnS, MnS, NiS, CoS.
See 12.11.7.4: Group IV - Insoluble sulfides precipitated by hydrogen sulfide, ZnS, MnS, CoS, NiS
5b. To the filtrate add ammonium carbonate solution.
Heat then filter.

6a. The precipitate may contain CaCO3, BaCO3, SrCO3.
See 12.11.7.5: Group V - Insoluble carbonates, CaCO3, BaCO3, SrCO3
6b. Observe the filtrate by table for Group VI.
See 12.11.7.6: Group VI - Magnesium, sodium and potassium, Mg, Na, K
If a test portion gives no precipitate on passing hydrogen sulfide, discard and treat remainder for later groups.

12.11.7.1 Group I Insoluble chlorides, PbCl2, AgCl, [Hg2Cl2 omitted]
1. Wash the precipitates twice with cold water and discard the washings.
Make a hole in a filter paper with a pointed glass rod, and wash the precipitate through into a beaker.
Boil with water and filter while hot.

2a. Residue: Wash the residue with hot water then pour warm ammonium hydroxide solution through the filter paper.
2b. Filtrate: The filtrate contains PbCl2 although white crystals may separate on cooling.
Add potassium chromate solution.
A yellow precipitate of lead chromate indicates the presence of lead.
Pb2+ + 2Cl- --> PbCl2
Pb2+ + CrO42- --> PbCrO4(s)
3a Residue: If the residue is black, it may be because of finely divided metallic mercury.
3b Filtrate: Acidify with nitric acid to form a white the precipitate of silver chloride that turns violet on exposure to light to indicate the
presence of silver.
Ag+ + Cl- --> AgCl (s)
Action of ammonia on silver chloride
AgCl + 2NH4OH + Ag(NH3)2Cl + 2H2O
[Ag(NH3)2Cl = soluble amine]
Precipitation of silver chloride by nitric acid.
Ag(NH3)2Cl + 2HNO3 --> 2NH4NO3 + AgCl (s)
See 12.12.11.3.10: Confirmation tests with original solution or solid, silver, lead.

12.11.7.2 Group II Sulfides insoluble in dilute hydrochloric acid
Keep the precipitate covered by a watch glass over the funnel to minimize oxidation of copper sulfide to copper (II) sulfate, which
would be washed out.
Wash the precipitate well with hot water.
Make a hole in the filter paper with a glass rod, and wash the precipitate through into evaporating dish.
Add caustic soda solution and a drops of yellow ammonium sulfide.
Heat briefly then filter.
Residue:
See 12.11.7.2a: Group IIa PbS, Bi2S3, CuS, CdS, [HgS omitted]
Filtrate:
See 12.11.7.2b: Group IIb As2S3, Sb2S3, SnS, SnS2.

12.11.7.2a Group IIa PbS, Bi2S3, CuS, CdS, [HgS omitted]
1. Wash the precipitate with hot water.
Make a hole in the filter paper and wash the precipitate into an evaporating basin.
Add dilute nitric acid and boil.
Sulfur will usually remain here.
Transfer the whole to a boiling tube, add dilute sulfuric acid and alcohol to complete the precipitation of lead sulfate.
Omit this step 1. if lead was not found present in Group I.
Leave to stand then filter.
--> 2a residue, 2b filtrate.

2a. The residue may contain lead sulfate, and sulfur.
Wash with hot water then wash into a boiling tube and boil with ammonium acetate solution.
Filter --> 2a.1 residue, 2a.2 filtrate
2a.1 The residue may contain mercury sulfide.
However, mercury salts are too dangerous for school science experiments.
2a.2 The filtrate contains lead.
Add potassium chromate solution.
A yellow precipitate of lead chromate, PbCrO4, indicates the presence of lead.
Pb2+ + CrO42- --> PbCrO4 (s)
2b. The filtrate may contain bismuth, copper and cadmium nitrates.
Add ammonia in excess.
Warm and filter.
2b.1 The residue is white bismuth hydroxide.
Wash the residue and dissolve it by pouring warm dilute HCl through the filter paper.
Pour filtrate into beaker nearly full of water.
Turbidity because of bismuth oxychloride, BiOCl, indicates the presence of bismuth.
Bi(OH)3 + 3H+ --> Bi3+ + 3H2O
Bi3+ + H2O + Cl- --> BiOCl (s) + 2H+
The filtrate contains bismuth.
2b.2 Divide the filtrate into two parts: 2b.2.1, 2b.2.2
2b.2.1 Part I.
If the liquid is colourless, omit this step.
Acidify with dilute acetic acid, add potassium ferrocyanide solution.
A brown colour or a precipitate because of copper ferrocyanide indicates the presence of copper.
[Cu(NH3)4]2+ + 4H+ --> Cu2+ + 4NH4+
2Cu2+ + Fe(CN)64- --> Cu2Fe(CN)6
copper + ferrocyanide ion --> copper ferrocyanide, [brown precipitate]
The filtrate contains copper.

2b.2.2 Part II.
If coloured, the filtrate may also contain cadmium.
The test for cadmium at this stage is too dangerous for school science experiments.

See 12.12.11.3.10: Confirmation tests with original solution or solid, bismuth, cadmium, copper
Precipitation of sulfides
Action of hydrogen sulfide on sulfides
Pb2+ + S2- --> PbS (s)
The red precipitate PbS is decomposed by more of the H2S gas to black lead sulfide, [similarly Cu2+, Cd2+]
2Bi3+ + 3S- --> Bi2S3 (s)
Action of dilute nitric acid on sulfides
SnS + 2H+ --> Sn2+ + H2S (g)
CuS + 2H+ --> Cu2+ + H2S (g), [Similarly CdS, PbS]
Bi2S3 + 6H+ --> 2Bi3+ + 3H2S
Oxidation of sulfides to sulfate
CuS + 8HNO3 --> CuSO4 + 8NO2 + 4H2O
Action of ammonia on bismuth, copper and cadmium nitrate solutions.
Bi3+ + 3OH- --> Bi(OH)3 (s)
Cu2+ + 4NH4OH --> [Cu(NH3)4]2+ + 4H2O, [deep blue ion]
Cd2+ + 4NH4OH --> [Cd(NH3)4]2+ + 4H2O.

12.11.7.2b Group IIb As2S3, Sb2S3, SnS, SnS2
1. The filtrate may contain antimony, arsenic, stannous and stannic sulfides.
Acidify with dilute hydrochloric acid the filtrate from Group II after treatment with caustic soda and ammonium sulfide.
If Group IIb is present, it will precipitate now, and if absent only white sulfur is seen.
If a precipitate appears, pass H2S to complete the precipitation, heat, filter then discard the filtrate.
Wash the precipitate with hot water, make a hole in the filter paper, and wash the precipitate through into an evaporating basin.
Add a piece of solid ammonium carbonate, warm for a few minutes, then filter the solution.

2a. The residue may contain antimony and tin sulfides.
Wash the residue.
Make a hole in the filter paper and wash through with concentrated hydrochloric acid into an evaporating basin.
Boil then divide the solution into two parts, Part 1 and Part 2. Leave to cool.

3a. Part 1. Dilute with its own volume of water.
Pass H2S.
An orange-red precipitate of antimony sulfide, Sb2S3, indicates the presence of antimony.
2Sb3+ + 3S2- --> Sb2S3(s)
3b Part 2. Add zinc foil in an evaporating basin until effervescence stops. Any tin present will be as a grey precipitate around the rim
of the evaporating basin.
Sn2+ + S2- --> SnS (s)
Sn4+ + 2S2- --> SnS2 (s).

2b.
The filtrate may contain arsenic.
Add dilute HCl carefully in a dish until effervescence stops.
Pass H2S and boil.
A yellow precipitate of arsenic sulfide indicates the presence of arsenic.
2As+ + 3S2- --> As2S3 (s)
See 12.12.11.3.10: Confirmation tests with original solution or solid, antimony, arsenic, tin.

4. Action of caustic soda and ammonium sulfide
The sulfur of the polysulfide, ammonium sulfide, oxidizes the lower sulfide and forms a thio salt.
As2S3 + 3S2- + 2S --> 2AsS43- , [thioarsenate ion], [2S from polysulfide]
The addition of acid precipitates the higher sulfide, arsenic sulfide, As2S5.
2AsS43- + 6H+ --> As2S5 + 3H2S (g).

5. Action of ammonium carbonate
As2S5 + 5(NH4)2CO3 + 3H2O --> 2H3AsO4 + 5(NH4)2S + 5CO2 (g)
[H3AsO4 = arsenic acid]
With continued subsequent passage of H2S
2H3AsO4 + 5H2S --> As2S5 + 8H2O.

6. Action of concentrated hydrochloric acid on boiling with antimony and tin sulfides
Sb2S3 + 6H+ -->2Sb3+ + 3H2S (g).

12.11.7.3 Group III Insoluble hydroxides, Fe(OH)3, Cr(OH)3, Al(OH)3
1. The precipitate may contain iron (III) hydroxide, chromium (III) hydroxide and aluminium hydroxide.
Wash well with hot water.
Make a hole in the filter paper with a glass rod, and wash the precipitate through into a wide mouthed boiling tube.
Add caustic soda solution and hydrogen peroxide solution, boil and filter.

2a. The residue is brown iron (III) hydroxide.
Dissolve in dilute HCl, add potassium ferrocyanide solution.
A blue the precipitate indicates the presence of iron.
To the original solution add:
2.1 Potassium ferricyanide solution.
A blue precipitate indicates the presence of iron (II) iron, Fe (II).
2.2. Potassium thiocyanate (KCNS) solution.
A red coloration indicates the presence of ferric iron, iron (III).
Some manganese hydroxide may occur here but it will appear in Group IV.
Fe3+ + 3CNS- --> Fe(CNS)3 [iron (III) thiocyanate]
2b. The filtrate may contain chromium and aluminium as sodium chromate and sodium aluminate.
A yellow colour indicates the presence of chromium.
Divide into two parts.

3a Filtrate Part 1: If the solution is not yellow, omit this step Add acetic acid in excess, and then lead acetate solution.
A yellow precipitate of lead chromate indicates the presence of chromium.
3b.
Filtrate Part 2: Add litmus indicator solution.
Add dilute HCl in excess, then NH4OH just in excess, then shake and leave to stand.
A blue lake of aluminium hydroxide, Al(OH)3, and litmus indicates the presence of aluminium
AlO2- + 4H+ --> Al3+ + 2H2O
Al3+ + 3OH- --> Al(OH)3 (s)
See 12.11.3.10: Confirmation tests with original solution or solid, aluminium, chromium, iron
Precipitation of hydroxides
Fe3+ + 3OH- --> Fe(OH)3 (s)
Cr3+ + 3OH- --> Cr(OH)3 (s)
Al3+ + 3OH- --> AI(OH)3 (s).

4. Action of caustic soda and hydrogen peroxide.
iron (III) hydroxide unchanged
2Cr(OH)3 + 3O + 4OH- --> 2CrO42- + 5H2O
[3O from oxidizing agent] [chromate ion = CrO42-]
Al(OH)3 + OH- --> AlO2- + 2H2O
[AlO2- = meta-aluminate ion].

12.11.7.4 Group IV Insoluble sulfides precipitated by hydrogen sulfide, ZnS, MnS, CoS, NiS
1. The precipitate may contain zinc, manganese, cobalt and nickel sulfides and sulfur.
If it is not black, cobalt and nickel are both absent.
Wash the precipitate.
Dilute the dilute hydrochloric acid with five times its own volume of water and pour this through the filter paper.

2. The residue may contain cobalt or nickel as sulfides.
Transfer the residue to an evaporating basin, add concentrated hydrochloric acid and a crystal of potassium chlorate.
Heat until all substances are dissolved then heat to evaporate until nearly dry.

3a. The nickel solution is yellow-green with a yellow crystal deposit.
Add an alkaline solution of dimethyl glyoxime.
A red colour or a red precipitate indicates the presence of nickel.
3b. A cobalt solution is pink and deposits blue crystals.
For confirmation apply the borax bead test.
Blue bead: cobalt.
Brown bead: nickel.

1a. The filtrate may contain zinc and manganese chlorides.
Boil the filtrate in an evaporating basin dish to remove hydrogen sulfide.
If the liquid is still turbid, finely divided sulfur is suspended in it.
Add potassium chlorate, boil until clear then leave to cool.
Add excess sodium hydroxide solution then filter.

4a. The residue is manganese hydroxide that turns brown on filter paper.
Wash into a boiling tube, leave to settle, pour off the water.
Add concentrated nitric acid and lead dioxide then boil.
Dilute and leave to settle.
A crimson colour because of permanganic acid indicates the presence of manganese.
4b. The filtrate contains zinc as sodium zincate.
Pass hydrogen sulfide.
A white precipitate, often discoloured, of zinc sulfide indicates the presence of zinc.
See 12.11.3.10: Confirmation tests with original solution or solid, cobalt, manganese, nickel, zinc.

5. Precipitation of sulfides.
Zn2+ + S2- --> ZnS (s)
Mn2+ + S2- --> MnS (s)
Co2+ + S2- --> CoS (s)
Ni2+ + S2- --> NiS (s)
If ammonium sulfide is used, some sulfur may also occur because ammonium sulfide also contains also polysulfides of the type NH4HSx.
A typical reaction:
Zn2+ + S22- --> ZnS (s) + S (s)
Action of very dilute hydrochloric acid.
Cobalt and nickel sulfides are unchanged.
ZnS + 2H+ --> Zn2+ + H2S (g)
MnS + 2H+ --> Mn2+ + H2S (g).

6. Action of excess sodium hydroxide solution on zinc and manganese chlorides.
Zn2+ + 2OH- --> Zn(OH)2 (s)
Zn(OH)2 + 2OH- --> ZnO22- + 2H2O
[ZnO22- = soluble zincate ion]
Mn2+ + 2OH- --> Mn(OH)2 (s)
On exposure to air the manganese hydroxide turns into brown manganese hydroxide.
4Mn(OH)2 + O2 --> 4MnO.OH + 2H2O
[MnO.OH = hydrated manganese sesquioxide].

12.11.7.5 Group V Insoluble carbonates, CaCO3, BaCO3, SrCO3
1. The precipitate may contain calcium, strontium and barium carbonates.
Wash well with hot water.
Pour through the filter paper some warm dilute acetic acid.
To a small portion of the filtrate add potassium chromate solution and boil.
If there is a the precipitate add potassium chromate solution to the whole and boil.
If there is no the precipitate discard sample and treat whole as filtrate.
Filter.

2a. The residue is pale yellow barium chromate that indicates the presence of barium,
12.11.3.2.1: Confirm by flame test.
2b. The filtrate may contain calcium and strontium as acetates.
Divide into two parts: 3a. Part I and 3b. Part II

3a. Part I. Add calcium sulfate solution and boil.
A faint white precipitate of strontium sulfate indicates the presence of strontium.
3b. Part II.
3b.1. If strontium is absent, add excess ammonium hydroxide and ammonium oxalate solution.
A white precipitate of calcium oxalate indicates the presence of calcium.
3b.2. If strontium is present, add dilute sulfuric acid, boil, filter, and reject the precipitate of strontium sulfate.
Add excess ammonium hydroxide and ammonium oxalate solution to the filtrate.
A white precipitate of calcium oxalate, CaC2O4, indicates the presence of calcium.
Ca2+ + C2O42- --> CaC2O4 (s) calcium oxalate [C2O42- = oxalate ion]
See 12.12.11.3.10: Confirmation tests with original solution or solid, barium, calcium, strontium
Precipitation of carbonates
Ba2+ + CO32- --> BaCO3 (s)
Sr2+ + CO32- --> SrCO3 (s)
Ca2+ + CO32- --> CaCO3 (s).

4. Action of dilute acetic acid
BaCO3 + 2H+ --> Ba2+ + H2O + CO2 (g)
SrCO3 + 2H+ --> Sr2+ + H2O + CO2 (g)
CaCO3 + 2H+ --> Ca2+ + H2O + CO2 (g).

5. Action of potassium chromate
Ba2+ + CrO42- --> BaCrO (s) barium chromate.
Calcium and strontium chromates are soluble in acetic acid, so do not precipitate.

12.11.7.6 Group VI Magnesium, sodium and potassium, Mg, Na, K
1. If calcium was found in Group V, add to the filtrate ammonium oxalate solution and boil.
Filter, and reject the precipitate of calcium oxalate.
If calcium was absent from Group V, omit this step.
Divide the filtrate into two parts: 2a. Part I and 2b. Part II.

2a. Part I Add ammonia and sodium phosphate solution.
Shake well and leave to stand.
A white crystalline precipitate of magnesium ammonium phosphate indicates the presence of magnesium.
Mg2+ + HPO42- + NH4+ + OH- --> MgNH4PO4 (s) + H2O
[MgNH4PO4 = magnesium ammonium phosphate].

2b. Part II Evaporate to dryness in an evaporating basin or platinum foil).
Heat until no more fumes from dissociating ammonium compounds are seen.
Observe residue by flame test.
A persistent gold yellow flame indicates the presence of sodium.
A lilac flame indicates the presence of potassium.
If sodium is present, examine the flame through blue glass for potassium.
See 12.11.3.10: Confirmation tests with original solution or solid, magnesium, potassium, sodium.