Preliminary tests
 

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Colour

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Unless the anion is a complex ion containing a transition metal, main group salts are colourless (white). Significant exceptions are AgBr, cream, and PbI₂ and AgI, both yellow; all three are insoluble in water. Oxides are often brightly coloured; they are insoluble in water. The colour found in oxides and in some other compounds such as AgBr and PbI₂ is due to interactions in the crystal lattice. Lead(II) iodide when dissolved in hot water is colourless, forming shining yellow plates as it crystallises.

Salts of transition elements are usually coloured; however freshly prepared anhydrous CuSO₄ is nearly white, CuI is pale cream – and insoluble in water. Colour in transition metal ions is usually due to electron transitions within the d-shell. Intensely coloured ions with the metal in its highest oxidation state (e.g. Mn(VII), Cr(VI), Fe(VI)) derive the colour from electron transitions between the metal and the oxygen atoms.

The reactions involved in these tests are covered in the cation and anion pages.

Vanadium:

(V) is yellow
(IV) is blue (not quite the same hue as copper(II) ions)
(III) is green
(II) is lavender

Chromium:

(III) is green or purplish green; some crystalline salts are deep purple,  e.g. chromium(III) potassium sulphate-24 water (chrome alum).

The colour of chromium(III) depends on the ligands around the ion. [Cr(H₂O)₆]³⁺ is violet; however production of chromium(III) by reduction of chromium(VI) compounds gives chromium(III) with some of the water ligands replaced with other species, e.g. [Cr(H₂O)₄Cl₂]⁺. Such ions are green, and water displaces the other ligands very slowly indeed.

(VI) is yellow in CrO₄^2-, orange in Cr₂O₇^2-.

Manganese:

(II) is very pale pink, appearing colourless in all but concentrated solutions
(VI) is deep green in MnO₄^2- (stable only in alkaline solutions)
(VII) is purple in MnO₄^- .

Iron:

(II) is pale green
(III) is yellow in solution; some solid salts, e.g. iron(III) ammonium sulphate, are amethyst.

Cobalt:

(II) (hydrated) is red; anhydrous is deep blue. This is the basis of the use of CoCl₂ paper to show the presence of water.
(III) is red, indistinguishable from Co(II) by colour alone.

Nickel:

(II) is green – a more intense colour than that of iron(II)

Copper:

(II) is blue or occasionally bright green in solution (CuCl₄^2-); solids may be either
(I) is nearly white (as CuI) and insoluble in water.

Inorganic analysis introduction  
Cation tests    
Anion tests 

Solubility

The following are insoluble (or nearly so) in water:  

Halides: Pb²⁺, Ag⁺
Sulphates:  Ba²⁺, Sr²⁺, Ca²⁺, Pb²⁺
Carbonates: all except those of the alkali metals. Aluminium carbonate does not exist.

The action of heat on the solid

A little of the compound is heated in a dry ignition tube.

Apparently sublimes:

Ammonium salt, most likely halides. The ammonium salts do not sublime, really; they thermally dissociate into substances that recombine to the ammonium compound on cooling. Ammonium nitrate can explode on heating.

Evolution of carbon dioxide: Carbonate (except those of group 1). Group 2 carbonates decompose only with difficulty, this increasing with increasing atomic mass of the group 2 metal.

Evolution of carbon monoxide and carbon dioxide: Ethanedioate C₂O₄^2-.

Evolution of ammonia: Ammonium salt. The salt thermally decomposes, but the ammonia diffuses more rapidly from the mixture than the hydrogen halide does.

Evolution of O₂: Group 1 nitrate (but see the next point) or, if coloured, higher oxide e.g. PbO₂

Evolution of O₂ and brown NO₂: Nitrate. In practice vigorous heating of a group 1 nitrate will produce brown fumes, but not immediately. The reason is that the nitrate will melt and decompose slowly, giving oxygen and the nitrite. However, the melting temperature of the nitrate is above the decomposition temperature of the nitrite, which will decompose to give NO₂. The liquid is usually somewhat yellow.

Evolution of SO₂: Sulphite or sulphate.

Evolution of water vapour accompanied by a colour change, suggests a hydrated salt of:

blue to white to black: copper(II)
green to yellow to red or black:  iron(II)
yellow to red or black: iron(III)
violet to green: chromium(III)
The colour change may reverse on cooling:
white  to yellow to white: zinc(II)
yellow to  red to yellow:  lead(II)
red to black to red:  iron(III).
 

The flame test

Although used here qualitatively, the use of a flame photometer enables the quantitative measurement of the concentrations of these ions, e.g. in biological fluids, using their flame emissions.

Clean a flame wire (platinum is best, but is also around £60 or $100 a go – so you’ll likely get nichrome) by dipping it repeatedly in concentrated hydrochloric acid and holding it in a roaring Bunsen flame until there is no trace of colour in the flame. Once you have cleaned the wire, don't touch the end of it; you'll put enough salty sweat on it to give an intense sodium flame.

To test a substance, moisten it with a few drops of concentrated hydrochloric acid, pick up the paste on the wire loop, and hold it to the edge of a roaring Bunsen flame.

Bright yellow: sodium. This test is very sensitive, and it may be a good idea to do a blank test (i.e. no salt) on the hydrochloric acid that you use. (The flame is not orange. Sodium street lights are because they also contain neon, which contributes some red to the colour.)

Lilac: potassium. Sodium is nearly always present, but can be filtered out by blue glass, when the potassium flame appears red.

Deep red: strontium or lithium. Lithium and strontium cannot be told apart by a flame test alone.

Apple green: barium.

Brick red or orange-red: calcium.

Green with blue centre: copper.
 

The action of sodium hydroxide solution

To a solution of the test substance, sodium hydroxide solution is added drop by drop until it is in excess (test with red litmus paper). Some solutions may be acidic to begin with (test the original solution with blue litmus paper); in such cases nothing will happen until the acid has all been neutralised. The chemistry of the tests with sodium hydroxide and with ammonia is covered in more detail in the cation analysis page.

White precipitate insoluble in excess NaOH: Mg²⁺, Ca²⁺.

White precipitate soluble in excess NaOH: Al³⁺, Zn²⁺, Pb²⁺.

Blue precipitate turning black on warming: Cu²⁺

Light green precipitate insoluble in excess NaOH: Ni²⁺

Dirty-green precipitate, turning brown on standing: Fe²⁺

Greyish-green precipitate soluble in excess NaOH to a deep green solution: Cr³⁺.

Foxy red (rust coloured) precipitate: Fe³⁺.

Beige precipitate turning brown on standing: Mn²⁺

Blue precipitate (turning grey) from a red solution: Co²⁺
 

The action of aqueous ammonia solution

To a solution of the test substance, ammonia solution is added drop by drop until it is in excess (test with red litmus paper). Some solutions may be acidic to begin with (test the original solution with blue litmus paper); in such cases nothing will happen until the acid has all been neutralised.

White precipitate insoluble in excess ammonia: Mg²⁺, Ca²⁺, Al³⁺, Pb²⁺.

White precipitate soluble in excess ammonia:  Zn²⁺.

Blue precipitate turning to a deep blue solution with excess ammonia: Cu²⁺.

Green precipitate soluble in excess ammonia to give pinkish solution: Ni²⁺.

Dirty-green precipitate, turning brown on standing: Fe²⁺.

Green precipitate insoluble in excess ammonia: Cr³⁺.

Foxy red (rust coloured) precipitate: Fe³⁺.

Beige precipitate turning brown on standing: Mn²⁺.

Blue precipitate from a red solution giving a blue solution with excess ammonia: Co²⁺.
 

The action of dilute hydrochloric acid

A little of the substance is added to about 5cm³ of dilute HCl. If there is no reaction warm gently. Test any gases evolved.

CO₂  evolved with vigorous effervescence: Carbonate or bicarbonate. Note that the only solid bicarbonates are those of group 1 metals.

NO₂ (brown) evolved: Nitrite. The solution will be pale blue.

SO₂ evolved: Sulphite.

Ethanoic acid evolved (smell of vinegar): Ethanoate.

The action of concentrated sulphuric acid

A little of the solid substance is added to about 2cm³ of concentrated sulphuric acid (CARE! Corrosive). If there is no reaction the mixture is warmed cautiously.

HCl evolved as steamy acidic fumes: Chloride.

Brown fumes: HBr + Br₂ + SO₂ from a bromide – these will turn a drop of silver nitrate on a glass
                      rod  cloudy.
                      or NO₂ from a nitrate or nitrite.

Purple fumes, brown mess, smell of bad eggs: HI + I₂ + H₂S from an iodide.

CO and CO₂ evolved: Ethanedioate.

Charring: Ethanedioate or ethanoate.
 

Tests for gases

These require careful technique. If the gas is to be collected it can be done by sucking the gas into a teat pipette.

Hydrogen: Usually ignites with a squeaky pop. A very unlikely product since it is only obtained from reactive metals or from hydrides. The pop is the explosion caused by hydrogen burning in the oxygen which usually contaminates it.

Oxygen: Relights a glowing splint - an example of increased reaction rate due to higher reagent concentration. This sometimes pops but not in the same squeaky way that hydrogen does.

Carbon monoxide: Burns with a blue flame but does not explode.

Carbon dioxide: Turns limewater milky, or, better, gives a white precipitate with limewater.

Sulphur dioxide: When passed into acidified potassium dichromate solution turns it from orange to green. The chromium(VI) is reduced to green Cr(III), and the SO₂ is oxidised to sulphate.

Halogen hydrides (or hydrogen halides): Turn a drop of silver nitrate on a glass rod cloudy: HCl white, HBr cream, HI yellow. Give a white smoke with ammonia fumes.

Chlorine: Turns blue litmus red, then bleaches it. Turns moist starch-iodide paper blue-black (because of oxidation of iodide to iodine). Gives a yellow solution of bromine with aqueous sodium bromide, and a darker yellow or brown solution of iodine with aqueous sodium iodide. On shaking with hexane the solutions colour the hexane brown or purple respectively.

Bromine (brown fumes): Reddens and then slowly bleaches blue litmus paper. Turns fluorescein paper scarlet. Liberates iodine from aqueous sodium iodide.

Iodine (violet fumes): Turns starch-iodide paper blue-black. Gives a cream precipitate with silver nitrate solution because of formation of bromide ions in aqueous bromine.

Nitrogen dioxide (brown fumes): Turns moist starch iodide paper blue. Does not affect silver nitrate solution.

Ammonia: Turns red litmus paper blue. Gives white smoke with concentrated hydrochloric acid vapours.