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3. INORGANIC Qualitative TESTS Cations and Acids |
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| TEST FOR | TEST METHOD | OBSERVATIONS | TEST CHEMISTRY-comments | |
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Ammonium
ion NH4+ |
Add COLD sodium hydroxide solution to the suspected ammonium salt and test any gas above the solution with red litmus. |
 Smelly
ammonia released! and red litmus turns blue. |
Ammonia
gas is evolved: NH4+(aq) + OH-(aq) ==> NH3(g) + H2O(l) Gentle warming of the mixture may be needed. |
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Acids
Hydrogen
ion i.e. H+ or H3O+ ion (note:
to completely identify acids you need to test for the anion e.g. chloride
for HCl etc.) |
(i)
Litmus or universal indicator or pH meter.
(ii) Add a little sodium hydrogencarbonate powder. |
(i) Litmus turns red, variety of colours with univ. ind. strong - red, weak
- yellow /orange, depending on strength of acid. (ii) Fizzing with any carbonate - test for CO2 as above. |
(i) A
pH meter reading gives a value of less than 7, the lower the pH number the
stronger the acid, the higher the H+ concentration,
(ii) HCO3-(aq) + H+(aq) ==> H2O(l) + CO2(g) |
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 Positive
metal cations via flame tests.(see also below for NaOH(aq) and NH3(aq) tests for metal ion) and heating carbonates. |
The metal salt or other compound is mixed with concentrated hydrochloric acid and a sample of the mixture is heated strongly in a bunsen flame on the end of a cleaned nichrome wire (platinum if you can afford it!) | Group 1: lithium Li/Li+ crimson | All
colours are due to electronic excitation to a higher level. You see the
light emitted as the electron returns to its lower more stable level.
This is the basis of atomic emission and absorption spectroscopy.
Aluminium, magnesium, iron and zinc do not produce a useful identifying
flame colour.
Other metals in Group 1:rubidium - red and caesium/cesium - blue
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| Group 1: sodium Na/Na+ yellow (can be slightly orangeish) | ||||
| Group 1: potassium K/K+ violet/lilac (crimson through cobalt blue glass) | ||||
| Group 2: calcium Ca/Ca2+ brick (yellowish) red (light green through cobalt blue glass) | ||||
| Group 2: strontium Sr/Sr2+ crimson | ||||
| Group 2: barium Ba/Ba2+ yellowish/apple green | ||||
| Transition Metal: copper(II) Cu/Cu2+ blue (flashes of green too) | ||||
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Positive
metal cations via sodium hydroxide (NaOH) or ammonia (NH3)
solutions. Note: (1) Both are alkalis, giving hydroxide ions, OH-, in their solutions. (2) Aluminium, magnesium, iron and zinc do not produce a useful identifying flame colour. |
Dilute
sodium hydroxide solution is added to a solution containing the
suspected ion. Both the precipitate formed and the effect of excess
alkali are important observations.
All precipitates white, unless otherwise stated and all tend to be gelatinous in nature. The test can be repeated with aqueous ammonia solution ('ammonium hydroxide'). The observations are usually, but not always, similar. ppt. = precipitate. More on some of these precipitates on the 3-d block Transition Metals series pages. |
aluminium
ion: Al3+(aq) + 3OH-(aq)
==> Al(OH)3(s)
White precipitate of aluminum hydroxide, which is not soluble in excess of the weak alkali ammonia, but dissolves in the stronger base/alkali sodium hydroxide (amphoteric) to give a clear colourless solution. Al(OH)3(s) + 3OH-(aq) ==> [Al(OH)6]3-(aq) (amphoteric behaviour because it dissolves in acids too) |
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calcium
ion: Ca2+(aq) + 2OH-(aq)
==> Ca(OH)2(s)
White precipitate of calcium hydroxide with sodium hydroxide IF the concentration of calcium ion is high. It is not soluble in excess of NaOH. No precipitate is formed with ammonia solution. |
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magnesium
ion: Mg2+(aq) + 2OH-(aq)
==> Mg(OH)2(s)
White precipitate of magnesium hydroxide, which is not soluble in excess of either NH3 or NaOH. You could distinguish Mg from Ca with a flame test or ammonia test above. |
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copper(II)
ion: Cu2+(aq) + 2OH-(aq)
==> Cu(OH)2(s)
Blue/turquoise ppt. of copper(II) hydroxide, which dissolves in excess ammonia to give a deep blue solution of an ammine complex, but copper(II) hydroxide is NOT soluble in excess NaOH. Cu(OH)2(s) + 4NH3(aq) ==> [Cu(NH3)4]2+(aq) + 2OH-(aq) |
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iron(II)
ion: Fe2+(aq) + 2OH-(aq)
==> Fe(OH)2(s)
Dark green precipitate of iron(II) hydroxide, which is not soluble in excess of NH3 or NaOH. Darkens in air due to oxidation to Fe(OH)3. |
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iron(III)
ion: Fe3+(aq) + 3OH-(aq)
==> Fe(OH)3(s)
Brown precipitate of iron(III) hydroxide, which is not soluble in excess of NH3 or NaOH. Another test for iron(III) ions is to add a few drops of potassium/ammonium thiocyanate solution and a blood-red coloured compound is formed. |
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zinc
ion: Zn2+(aq) + 2OH-(aq)
==> Zn(OH)2(s)
White precipitate of zinc hydroxide, which dissolves in both excess (i) sodium hydroxide or (ii) ammonia to give a clear colourless solution: (i) Zn(OH)2(s) + 2OH-(aq) ==> [Zn(OH)4]2-(aq) (amphoteric behaviour because zinc hydroxide dissolves in acids too). (ii) Zn(OH)2(s) + 4NH3(aq) ==> [Zn(NH3)4]2+(aq) + 2OH-(aq) (soluble complex ion formation) |
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chromium (III)
ion: Cr3+(aq) + 3OH-(aq)
==> Cr(OH)3(s)
grey-green precipitate of chromium(III) hydroxide, which is soluble in excess of NaOH (amphoteric, dissolves in acids too) but not soluble in excess ammonia NH3. With sodium hydroxide a dark green soluble hexahydroxo-complex ion is formed. Cr(OH)3(s) + 3NaOH(aq) ==> [Cr(OH)6]3-(aq) |
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manganese(II)
ion: Mn2+(aq) + 2OH-(aq)
==> Mn(OH)2(s)
off-white precipitate of manganese(II) hydroxide, which is not soluble in excess of NH3 or NaOH and rapidly turns in air due to oxidation to manganese(III) oxide Mn2O3 (and manganese(IV) oxide, MnO2 too?) |
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lead(II)
ion: Pb2+(aq) + 2OH-(aq)
==> Pb(OH)2(s)
White precipitate of lead(II) hydroxide, which dissolves in excess sodium hydroxide (amphoteric) to give a clear colourless solution but does not dissolve in excess ammonia solution. Pb(OH)2(s) + 2OH-(aq) ==> [Pb(OH)4]2-(aq) (with NaOH amphoteric behaviour because zinc hydroxide dissolves in acids too). |
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| The barium ion, Ba2+(aq) does not give a hydroxide precipitate because barium hydroxide, Ba(OH)2, is too soluble. | ||||
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MISCELLANEOUS
CATION TESTS:
(i) Lead(II) ion |
(i) add potassium iodide solution => yellow precipitate | (i) Pb2+(aq) +2I-(aq) ==>PbI2(s) lead(II) iodide ppt. | ||
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Metal
Carbonates
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Sometimes heating a metal carbonate strongly to decompose it provides some clues to its identity. Adding acid ==> CO2 and the colour of the resulting solution (e.g. blue Cu2+(aq), may also provide clues. The metal ion solution might also give a flame colour or a hydroxide precipitate with sodium hydroxide e.g. copper. |
(i) copper(II)
carbonate==> copper(II) oxide + carbon dioxide CuCO3(s) ==> CuO(s) + CO2(g) [green] ==> [black] + [colourless gas, test with limewater, white precipitate] (ii) zinc carbonate==> zinc oxide + carbon dioxide ZnCO3(s) ==> ZnO(s) + CO2(g) [white] ==> [yellow hot, white cold] +[colourless
gas, test with limewater, white precipitate] |
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