Tests for ions and instrumental analysis

A chemical test is only useful if one substance and no other gives that result. If two ions produced the same colour, a positive result would leave you no wiser than before. That is why the tests are chosen as they are, and why a question about a test may ask you to explain a result that could have come from more than one thing — the answer is always to do a second, different test.

The test for ammonia gas has three parts and all three are needed:

• the gas has a sharp, choking smell • it turns damp red litmus paper blue • the paper must be damp, because the ammonia has to dissolve in the water before it can act as an alkali.

Ammonia is the only common gas that turns damp red litmus blue, which is exactly what makes the test worth having. It is also how the ammonium ion is identified: warm the substance with sodium hydroxide solution, and if ammonium ions are present ammonia gas is given off and can be tested at the mouth of the tube.

Flame tests can be used to identify some metal ions (cations). Lithium, sodium, potassium, calcium and copper compounds produce distinctive colours in flame tests:

• lithium compounds result in a crimson flame • sodium compounds result in a yellow flame • potassium compounds result in a lilac flame • calcium compounds result in an orange-red flame • copper compounds result in a green flame.

If a sample containing a mixture of ions is used, some flame colours can be masked. The yellow of sodium is the usual culprit — it is bright enough to hide a lilac potassium flame completely, which is the main limitation of the method. Flame colours of other metal ions are not required knowledge.

The method: clean a wire loop in concentrated hydrochloric acid and hold it in a hot blue flame until it burns with no colour, then dip it into the sample and return it to the flame. Cleaning matters — a trace of sodium left from a previous test will show yellow whatever the new sample is.

Sodium hydroxide solution can be used to identify some metal ions (cations).

Solutions of aluminium, calcium and magnesium ions form white precipitates when sodium hydroxide solution is added, but only the aluminium hydroxide precipitate dissolves in excess sodium hydroxide solution. That is how the three white precipitates are told apart: keep adding sodium hydroxide, and the one that redissolves is aluminium.

Solutions of copper(II), iron(II) and iron(III) ions form coloured precipitates when sodium hydroxide solution is added. Copper(II) forms a blue precipitate, iron(II) a green precipitate and iron(III) a brown precipitate.

You should be able to write balanced equations for the reactions to produce the insoluble hydroxides. The pattern is always the metal ion joining as many hydroxide ions as its charge requires. You are not expected to write equations for the production of sodium aluminate.

Carbonates react with dilute acids to form carbon dioxide gas. Carbon dioxide can be identified with limewater.

So the test for a carbonate is two steps: add dilute acid and look for fizzing, then bubble the gas through limewater and look for it turning milky. The fizzing alone is not enough, because a metal in acid fizzes too — the limewater is what distinguishes carbon dioxide from hydrogen.

Halide ions in solution produce precipitates with silver nitrate solution in the presence of dilute nitric acid. Silver chloride is white, silver bromide is cream and silver iodide is yellow.

The three colours run in order down Group 7, from white through cream to yellow, which makes them easier to remember than to distinguish in practice — the difference between white and cream is slight, so the tubes are compared side by side.

The dilute nitric acid is added first, to remove carbonate ions. Without it, a carbonate in the sample would also give a white precipitate with silver nitrate and be mistaken for a chloride.

Sulfate ions in solution produce a white precipitate with barium chloride solution in the presence of dilute hydrochloric acid.

Again the acid is added first to remove carbonate ions, which would otherwise give a white precipitate of barium carbonate and be mistaken for a sulfate.

You should be able to identify species from the results of the tests above — that is, to work backwards from a set of observations to the compound.

Elements and compounds can be detected and identified using instrumental methods. Instrumental methods are accurate, sensitive and rapid.

You should be able to state advantages of instrumental methods compared with the chemical tests in this specification. Those three words are the answer, and each is worth expanding: accurate, so the result is not a judgement of whether a precipitate looks white or cream; sensitive, so a very small sample is enough; and rapid, so many samples can be analysed quickly.

Flame emission spectroscopy is an example of an instrumental method used to analyse metal ions in solutions.

The sample is put into a flame and the light given out is passed through a spectroscope. The output is a line spectrum that can be analysed to identify the metal ions in the solution and measure their concentrations.

It is the flame test done properly: instead of one colour judged by eye, the light is split into a pattern of lines unique to each metal ion. Because the lines from different ions appear at different places, a mixture that would mask a flame colour can still be read — and the brightness of the lines gives the concentration as well.

You should be able to interpret an instrumental result given appropriate data in chart or tabular form, when accompanied by a reference set in the same form, limited to flame emission spectroscopy.

The method is matching, not memory. Line up the unknown's spectrum against each reference spectrum and look for a set of lines that appears in both at the same positions. Every reference line for an ion must be present in the unknown before you claim that ion is there; if lines are present that no single reference explains, the sample contains more than one metal ion.