0620 · 12.5Identification of ions and gases

Principles of chemical identification

0:00Two solutions can look identical but contain completely different substances. Their appearance alone gives us very little evidence. Chemical tests reveal differences by producing gases, precipitates or colour changes. By the end of this video, you will understand how these observations identify ions and gases, and why the order of a test matters. Think of each test as collecting a chemical clue. One observation may suggest several possibilities, while a second test narrows the choice.

0:27The strongest identification connects a particular reagent with a particular result. That means we need to understand both what we add and what we observe afterwards. An ion is an electrically charged particle. Positive ions are called cations, while negative ions are called anions. A dissolved salt contains both types, dispersed through the water. For example, sodium chloride solution contains sodium ions with a 1+ charge and chloride ions with a 1− charge.

0:56The word aqueous means dissolved in water, not simply wet or liquid. A reagent is a substance added to produce a chemical reaction during a test. Sometimes dissolved ions react to form an insoluble solid called a precipitate. Other tests produce a gas or change the colour of a solution. Keep a precipitate separate from a coloured solution. A precipitate contains solid material suspended in the liquid or settled at the bottom.

1:20A solution has no visible solid, even when it is coloured. If a precipitate dissolves, the solid disappears into solution; it has not simply fallen out of sight. Use small, separate portions of the unknown solution for different tests. Otherwise, a reagent from an earlier investigation can interfere with the next reaction. Record observations before drawing conclusions, and include negative results when relevant. These habits matter in both the Practical Test and the Alternative to Practical paper.

1:49These tests belong in a supervised laboratory, using eye protection and small quantities. Never taste a chemical or identify a gas by deliberately smelling it. Some gases in this topic are harmful, especially chlorine and sulfur dioxide. Gas tests give safer evidence than relying on an odour.

Tests for common gases

2:07Ammonia turns damp red litmus paper blue. Litmus responds to acidity or alkalinity, and ammonia produces an alkaline solution when it dissolves in water. The moisture on the paper allows that process to happen. Dampness is therefore part of the test, rather than an optional detail. To test an unknown gas for ammonia, hold damp red litmus paper near the mouth of the test tube. A change to blue is the required observation.

2:31Saying that the gas is alkaline explains the result, but does not replace it. Your record should name the paper and describe its colour change. Carbon dioxide turns limewater milky. Limewater is an aqueous solution of calcium hydroxide. Carbon dioxide reacts with it to form insoluble calcium carbonate, which scatters light and makes the liquid appear milky. Bubble the gas through the limewater so that it actually contacts the reagent.

2:57Chlorine bleaches damp litmus paper. Bleaching means the paper loses its colour completely. An ordinary acid colour change is not the required result. Moisture allows chlorine to react at the paper. Write that the damp litmus paper is bleached, rather than saying that chlorine simply turns it red. Hydrogen gives a pop with a lighted splint. A small sample reacts rapidly with oxygen from the air when ignited. The sound is the identifying observation.

3:23Oxygen, by contrast, relights a glowing splint because it supports combustion and allows the remaining hot wood to burn more vigorously. A lighted splint has a flame, whereas a glowing splint has a hot tip without a flame. Those starting conditions distinguish the two tests. Do not swap them or write that oxygen itself burns. Oxygen supports combustion, while hydrogen is the gas being burned in its identification test. Sulfur dioxide turns acidified aqueous potassium manganate(VII) from purple to colourless.

3:54This reagent acts as an oxidising agent, and sulfur dioxide reduces it. That reaction changes the coloured manganate ions into a colourless product. The useful observation is the complete change from purple to colourless, not simply that the solution becomes lighter. The carbonate test combines gas production with gas identification. Add dilute acid to the sample and look for effervescence, which means bubbling as gas escapes. Then test the gas with limewater.

4:22If the limewater turns milky, carbon dioxide has been produced, supporting the presence of carbonate ions. Carbonate ions react with hydrogen ions from the acid to form carbon dioxide and water. This explains why adding acid releases a gas from a substance that originally contained no gas bubbles. Effervescence alone is not enough for identification because other reactions also release gases. The limewater result provides the distinguishing evidence.

Tests for anions

4:50Chloride, bromide and iodide ions are tested using the same sequence. First acidify the sample with dilute nitric acid, then add aqueous silver nitrate. Silver ions react with these anions to form insoluble silver salts. The colour of the precipitate tells us which of the three anions is present. Chloride gives a white precipitate. With bromide the solid is cream, and iodide gives a yellow one. These colours describe the solid, not the original solution.

5:17Use the precise colour names because cream and yellow distinguish different ions. A cloudy appearance without a colour description leaves the identification incomplete. Acidifying first helps remove interfering ions such as carbonate, which could also form a precipitate with silver ions. Nitric acid is chosen because it does not introduce chloride, bromide or iodide. Hydrochloric acid would add chloride ions itself. A white precipitate could then come from the reagent rather than the unknown sample.

5:47The sulfate test follows a similar pattern but uses a different precipitating reagent. Acidify the sample with dilute nitric acid, then add aqueous barium nitrate. A white precipitate indicates sulfate ions. Barium ions combine with sulfate ions to form insoluble barium sulfate, making the solid visible. Sulfite ions are tested with acidified aqueous potassium manganate(VII). The solution changes from purple to colourless because sulfite reduces the manganate ions. Notice the connection with the sulfur dioxide gas test, which uses the same reagent.

6:21In both cases, the evidence is a colour change rather than a precipitate. Sulfate and sulfite are different ions despite their similar names. Their formulae on screen show that sulfate contains four oxygen atoms, while sulfite contains three. Both have a 2− charge, but their reactions differ. Identify sulfate through the barium nitrate precipitate and sulfite through the manganate colour change. For nitrate ions, add aqueous sodium hydroxide and aluminium foil, then warm gently.

6:51Test any ammonia produced using damp red litmus paper. The paper turns blue if ammonia is present. This test transforms the nitrate into a recognisable gas instead of producing a characteristic precipitate. Aluminium acts as a reducing agent in the alkaline mixture, converting nitrate into ammonia. Sodium hydroxide provides the alkaline conditions, and warming helps the reaction proceed. You do not need a complicated equation to describe this test. You do need all the reagents, the warming step and the ammonia result.

7:23Ammonium ions can also release ammonia when warmed with sodium hydroxide, so they can complicate the nitrate test. If both are possible, test a separate portion without aluminium first. Ammonia from that portion shows ammonium is present. Do not automatically attribute every positive ammonia result to nitrate.

Cation precipitate tests

7:40Most of our cation tests use aqueous sodium hydroxide or aqueous ammonia. Add a small amount dropwise, observe any precipitate, then add the same reagent in excess. Excess means enough additional reagent to test whether the solid dissolves. Repeat the other reagent test using a fresh portion of the sample. These reagents provide hydroxide ions, which can combine with dissolved metal ions to form insoluble metal hydroxides. Different metal ions produce different precipitate colours.

8:09Some hydroxides dissolve when more reagent is added because further reactions produce soluble species. That second stage often separates ions with the same initial appearance. Aqueous ammonia is not interchangeable with sodium hydroxide. It provides hydroxide ions through its reaction with water, but can also interact with certain metal ions differently. We therefore compare the results from both reagents. The useful fingerprint includes the initial colour and whether the precipitate remains in excess.

8:38Copper(II) ions give a light blue precipitate with aqueous sodium hydroxide. It is insoluble in excess, so adding more sodium hydroxide does not remove the solid. Aqueous ammonia also produces a light blue precipitate initially. In excess ammonia, however, it dissolves to give a dark blue solution. That dark blue solution is a particularly useful clue for copper(II) ions. Excess ammonia allows a soluble copper species to form, so the solid disappears.

9:05Describe both changes: a light blue precipitate forms, then dissolves in excess to give a dark blue solution. Simply writing blue misses the most distinctive stage. Iron(II) ions give a green precipitate with either aqueous sodium hydroxide or aqueous ammonia. In both tests, it is insoluble in excess. Iron(III) ions instead give a red-brown precipitate with either reagent, again insoluble in excess. The charge matters because these are different ions of the same element.

9:35Chromium(III) ions also produce a green precipitate, so green does not identify iron(II) by itself. With aqueous sodium hydroxide, the chromium precipitate dissolves in excess. With aqueous ammonia, it remains insoluble in excess. Comparing that behaviour distinguishes chromium(III) from iron(II), whose precipitate remains with both reagents. This is why the excess stage cannot be skipped. An initial green precipitate gives us a group of possibilities rather than a complete identification.

10:05If it dissolves in excess sodium hydroxide, chromium(III) is supported. When it remains insoluble with both reagents, iron(II) matches the expected results. Aluminium, zinc and calcium can all produce white precipitates with aqueous sodium hydroxide. The colour alone therefore cannot separate them. We need to observe the effect of excess reagent and compare the ammonia test. These tests work together like successive clues, each eliminating a different possibility. Aluminium ions give a white precipitate with aqueous sodium hydroxide.

10:36It dissolves in excess to give a colourless solution. With aqueous ammonia, a white precipitate also forms, but it is insoluble in excess. The contrast between the two reagents is the identifying pattern. Zinc ions give a white precipitate with either reagent. In excess sodium hydroxide, it dissolves to give a colourless solution. The solid also disappears in excess aqueous ammonia, and again the liquid left is colourless. Unlike aluminium, zinc therefore shows dissolution in both tests rather than only with sodium hydroxide.

11:10Calcium ions give a white precipitate with aqueous sodium hydroxide, insoluble in excess. With aqueous ammonia, there is no precipitate or a very slight white precipitate. Ammonia provides a lower concentration of hydroxide ions under these test conditions. That difference explains why the calcium result is much less pronounced. The white precipitate patterns become easier when organised as a decision. A solid that remains in excess sodium hydroxide fits calcium among these three possibilities.

11:38If it dissolves, consider aluminium or zinc. Test a fresh portion with ammonia to distinguish them, because only the zinc precipitate dissolves in excess. Aluminium and zinc hydroxides are amphoteric, meaning they can react with both acids and bases. Their reaction with excess sodium hydroxide produces soluble species, explaining the disappearance of the precipitate. We do not need to memorise the structures of those species here. The required evidence is dissolution and the resulting colourless solution.

12:07Ammonium ions behave differently from the metal cations in these tests. Add aqueous sodium hydroxide and warm gently. Ammonia is given off and turns damp red litmus paper blue. Hydroxide ions remove a hydrogen ion from ammonium, producing ammonia and water. Do not use aqueous ammonia to identify ammonium ions. That reagent already contains the very substance you would be trying to detect. The required ammonium test uses sodium hydroxide and warming, without aluminium foil.

12:38Aluminium belongs to the nitrate reduction test, where it has a different job.

Flame tests and unknown salts

12:41Some cations can be identified by the colour they produce in a flame. Clean a suitable wire loop with hydrochloric acid and heat it until it gives no flame colour. Place a small amount of sample on the loop. Then hold it in a non-luminous Bunsen flame and observe the colour. Heating gives energy to electrons in the metal ions, and light is emitted as those electrons lose that energy.

13:03Each metal gives a characteristic colour because its electrons have their own set of energy levels. You do not need detailed atomic energy calculations here. The essential connection is that the emitted light provides an identifying clue. Lithium gives a red flame, sodium gives yellow, and potassium gives lilac. Calcium produces orange-red, while barium gives light green. The flame colour for copper(II) is blue-green. Use these precise names because similar descriptions can otherwise make different cations sound indistinguishable.

13:33Cleaning the loop prevents material from a previous sample affecting the result. Sodium contamination is especially troublesome because its yellow emission can obscure another flame colour. Repeat the test with a clean loop if the result is uncertain. A flame test identifies a cation, not the complete salt or its anion. Suppose an unknown salt solution gives the three results shown. We first acidify one portion with dilute nitric acid and add aqueous silver nitrate.

14:00A white precipitate forms, supporting chloride ions. We have identified the anion, but we still need evidence about the cation before naming the salt. A fresh portion gives a white precipitate with aqueous sodium hydroxide. The solid dissolves in excess, leaving a colourless solution. This fits aluminium or zinc, but not calcium. Calcium is excluded because its precipitate would remain insoluble in excess sodium hydroxide. Another fresh portion gives a white precipitate with aqueous ammonia.

14:29This time, the solid remains in excess. Aluminium matches that pattern, while zinc does not because its precipitate would dissolve. Combining aluminium ions with chloride ions identifies the unknown salt as aluminium chloride. Notice how the conclusion depends on the complete pattern rather than one isolated observation. A white precipitate with sodium hydroxide was not enough to identify aluminium. The excess stage narrowed the possibilities, and ammonia separated the remaining candidates.

14:59Using fresh portions also kept each result independent of the earlier reagents. Here is a second unknown, this time a blue solution. A flame test gives a blue-green flame, which suggests copper(II) ions. Aqueous sodium hydroxide gives a light blue precipitate that is insoluble in excess. Aqueous ammonia gives a light blue precipitate that dissolves in excess to give a dark blue solution. Those three cation results all point to copper(II) ions.

15:27A separate portion is acidified with dilute nitric acid, and aqueous barium nitrate is added. A white precipitate forms, so sulfate ions are present. Putting the two ions together, the salt is copper(II) sulfate. Notice that the blue colour of the original solution was only a hint. Many exam answers lose marks by naming copper from that colour alone. The precipitate tests and the flame test supply the real evidence.

15:52The anion test was still needed, because the cation tests say nothing about sulfate. Copper(II) chloride, for example, would give the same cation results but no precipitate with barium nitrate.

Exam answers and recap

16:04An examination may ask you to describe a test, state an observation or identify an ion from given results. Match your answer to the command word. When asked to describe a test, give the reagent, any condition such as warming, and the positive result. When asked to state an observation, report what is seen or heard rather than only naming the substance. Consider the instruction to describe a test for nitrate ions in an aqueous solution.

16:28Add aqueous sodium hydroxide and aluminium foil to the sample. Warm the mixture gently. Test the gas with damp red litmus paper, which turns blue because ammonia is produced. For a sulfate test, the answer must specify dilute nitric acid followed by aqueous barium nitrate. A white precipitate is the positive observation. An answer that leaves out the acidification step is incomplete. Using silver nitrate instead would test a different group of anions.

16:56Another question might ask you to explain why a green precipitate does not prove iron(II) ions are present. Chromium(III) also gives a green precipitate with sodium hydroxide. Add the reagent in excess to distinguish them. The chromium precipitate dissolves, while the iron(II) precipitate remains insoluble. Gas tests are often asked as short describe questions. Each answer needs a method and a result. Hold a lighted splint at the mouth of the tube, and hydrogen gives a pop.

17:24A glowing splint placed in oxygen relights. Chlorine is shown by holding damp litmus paper in the gas, where the paper is bleached. Avoid confusing observations with conclusions in practical records. A white precipitate is an observation. The presence of chloride ions is a conclusion drawn from that observation in the correct test. Similarly, bubbling does not prove carbon dioxide without the limewater result. Complete evidence includes the reagent, any necessary condition and the characteristic change.

17:53Several small omissions cause large errors in these tests. Dry litmus leaves out the moisture needed for the gas test. Missing the word excess can hide the distinction between aluminium and zinc. Swapping a glowing splint for a lighted one changes the method, while confusing sulfate with sulfite selects the wrong reagent. Chemical identification works by converting invisible differences into visible evidence. Anion tests may produce gases, insoluble salts or changes in reagent colour.

18:22Cation tests compare precipitate colour and behaviour in excess, while flame tests use characteristic light. The gas tests then identify products through litmus, limewater, splints or acidified potassium manganate(VII). Remember the logic behind the sequence, not just a collection of colours. Acidification removes interference, fresh portions keep tests independent, and excess reagent reveals additional chemical behaviour. Use a confirming test whenever an observation could have several explanations. Together, these ideas let you identify an unknown substance from a coherent set of chemical clues.