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Qualitative Analysis - Tests for Anions

  • Page ID
    34970
  • Detecting Sulfides

    Sodium Nitroprusside

    Use Sodium Nitroprusside (\(Na_2[Fe(CN)_5NO]\)). In the presence of sulfide ions a deep violet coloration is formed. \([Fe(CN)_5NOS]^{4-}\) complex anion formed in the presence of sulfide ions is the cause for this coloration.

    Silver Nitrate

    If the salt is reacted with Silver Nitrate,it forms a black precipitate of \(Ag_2S\).

    Lead Ethanoate

    If the substance is soluble use lead acetate solution, if solid add dilute hydrochloric acid and test the gas with lead ethanoate paper. If the salt in question is a sulfide an acrid smell of rotten eggs will be observed and the lead ethanoate will turn black. This is due to the reaction: \[(CH_3COO)_2Pb + H_2S \rightarrow PbS + 2(CH_3COOH)\]

    H\(_2\)S

    The following two points do not fall in place here, kindly arrange them wherever they are best fit:

    • When \(H_2S\) is bubbled through a solution of Sodium sulphite or Sodium bisulphite, a white turbidity of sulfur is seen. What happens is, the S atom in sulphite/bisulphite (Oxidation state +5) and in \(H_2S\) (Oxidation state -2) have coproportionated to give molecular sulfur in zero oxidation state. The reaction can be viewed as the reverse of a disproportionation reaction.
    • When \(H_2S\) is bubbled through acidified \(KMnO_4\) solution, it decolorizes and a white turbidity is observed. This is because, the powerful oxidizing agent permanganate ion has oxidized the sulfide ion (-2) to zero oxidation state. In doing so, the permanganate ions themselves reduce to \(Mn^{2+}\) ions, thus the purple/violet/pink color of the solution is diluted and eventually fades away as the manganous ions do not impart any color to the solution.

    Detecting Nitrite ions

    The nitrite ion gives several reactions based on its redox properties.

    \(I_{3}^{-}\)

    When it reacts with the \(I^-_3\) anion, the nitrite ion reduces itself to nitric oxide. The student should check the oxidation states of nitrogen in both the cases.

    \[NO_2^-+2I_3^-+2H^+\rightarrow NO(g)+3I_2(aq)+H_2O(l)\]

    The yellow color of the\(I_3^-\) ions changes to a greenish tinge which later turns violet. This observation is made in the presence of an oxidizing agent which oxidizes iodine from \(-1/3\) to 0. The O.A. in our case is of course, the nitrite ion!

    Before turning violet, a greenish tinge is observed because the increasing concentration of violet color and the decreasing concentration of yellow color during the reaction impart a greenish color.

    Lead/Silver Nitrates

    Most lead and silver anions show much the same characteristics in analysis. However, lead nitrite is soluble while silver nitrite is an insoluble white solid.

    Therefore, to test for the presence of nitrite ions we can use silver and lead nitrate. No precipitate will be observed with lead nitrate, but a white precipitate will form when silver nitrate is added.

    \[2NO_2^- +Pb^{2+}\rightarrow Pb(NO_2)_2(aq)\]

    \[NO_2^- +Ag^+ \rightarrow AgNO_2 (s)\]

    Detecting Cyanide Ions

    Ferric ions

    When ferric ions are added to a solution of cyanide ions, no specific change in color is observed. However, if we add the ferric ions to the solution after adding thiosulfate or \(S_2^{2-}\), we get a blood red coloration of ferric thiocyanate.

    Ferric thiocyanate is \(Fe(SCN)_3\).

    Actually, this is a characteristic test for the presence of ferric ions and for thiocyanate ions. Cyanide ions are nowhere directly involved in the formation of ferric(III) thiocyanate. If we review the test procedure for the detection of cyanide ions, we find that we add, according to our choice, either thiosulfate ions or \(S_2^{2-}\) ions. These ions disproportionate in the presence of cyanide ions, forming thiocyanate ions. It is these thiocyanate ions that eventually react with the added ferric ions to give the blood red complex formation

    \[CN^- +S_2O_3^{2-}\rightarrow SCN^- +SO_3^{2-}\]

    \[CN^- +S_2^{2-}\rightarrow SCN^- +S^{2-}\]

    \[Fe^{3+} +SCN^- \rightarrow Fe(SCN)^{2+}\]

    Metallic copper

    It is a popular fact that most metals react with either water or dilute acid solution to liberate hydrogen. This is a redox reaction in which the metal acts as a reducing agent, reducing the proton or hydrogen in water to hydrogen gas, oxidizing itself to a metal cation.

    A classic example of this type of redox reaction is the reaction of metallic sodium with water. The reaction takes place and liberates enough heat to melt the sodium which then floats on a cushion of hydrogen gas. e.g., \(2Na + 2H_2O \rightarrow H_2 + 2NaOH\). In terms of the ions involved in the reaction \(Na \rightarrow Na^{1+} + e^{1-}\) and then \(2H^{1+} + 2e^{1-} \rightarrow H_2\). The same reaction occurs with any of the alkali metals in exactly the same way.

    However, there are some metals like copper which will not react with either water or dilute acids in the same way. Quantitatively, Standard Electrode Potential is lower than that of hydrogen ions, meaning that more energy is required to form copper ions than the hydrogen ions can provide. [Note to editors --> Is this any better?]

    However, if we add metallic copper to a solution containing cyanide ions, copper gives into the greed of forming that ultrastable cyano complex of itself. That is, equilibrium favours the formation of tetracyanocomplexcuprate(I) ions. This is the point - in presence of cyanide ions, copper oxidizes itself to form the +1 metal complex. But then, if an oxidation has occurred so must have a reduction. What is the reduction reaction? The formation of hydrogen (editors - formation from what? proton or water molecule?) is the involved reduction.

    We conclude that an aqueous solution of KCN dissolves metallic copper, generally administered as copper turnings through, the formation of tetracyanocuprate(I) ions.

    (Note to editors --> Explain this anamlolous behaviour of copper through the quantitative aspects of electrochemistry and chemical equilibrium, short enough to maintain interest.)

    Note on the tetracyano complex of copper

    When \(H_2S\) is bubbled through the cyano complex of ferrous/ferric/cuprous/cupric ions, no sulfide is precipitated. While with most other metal cations, metal sulfide does get precipitated.

    e.g. tetracyanocadmiumate ions give CdS with \(H_2S\) and dicyanoargentate (silver complex) give \(Ag_2S\).

    This can be explained on the basis of equilibrium constants - solubility products and formation constants of the concerned sulfides and cyanocomplexes. (note to editors --> plz explain in a little more detail)

    Detecting Hypochlorite ions

    Adding a solution containing lead ions, such as lead nitrate, to a hypochlorite solution produces a brown precipitate of lead(IV) oxide.

    \[2ClO^- (aq)+Pb^{2+} \rightarrow PbO_2 (s)+2Cl^- (aq)\]

    The azide ion

    On adding triiodide ions, the azide ion oxidizes to nitrogen and decolorisation is observed. . Triiodide ions are reduced to iodide ions, and nitrogen is produced. The colour of solution changes from yellow to colourless and a brisk effervescence is noted.

    \[2N_3^- \rightarrow 3N_2 +2e^-\]

    \[I_3^- +3e^- \rightarrow 3I^-\]

    Detecting Acetate ions

    Dilute hydrochloric acid

    When dilute hydrochloric acid is added to an acetate solution, the acetate ions are acidified with the hydrogen ions of the hydrochloric acid, forming acetic acid which produces an observable smell of vinegar.

    Esterification

    When a drop of concentrated sulfuric acid and ethanol are added, the distinctive smell of an ester can be observed. This applies to all carboxylates, but is particularly useful in this case because the two reactants, acetic acid and ethanol, have distinctive smells which are replaced the completely different smell of the ester; in this case, ethyl acetate, which smells not unlike glue or nail polish remover.

    \[CH_3COOH + CH_3CH_2OH \rightarrow CH_3COOCH_2CH_3 + H_2O\]

    The sulfuric acid is not shown as it is merely a catalyst (specifically a dehydrating agent).

    Ferric Chloride

    When added to neutral \(FeCl_3\) acetates produce a red coloration.

    \[FeCl_3 + 3CH_3COOH \rightarrow (CH_3COO)_3Fe + 3HCl\]

    Neutral \(FeCl_3\) is prepared by taking about an mL of \(FeCl_3\) in a test tube and adding about an mL of \(NH_4OH\) to give a red precipitate. Dil. HCl is added to the test tube till the precipitate has JUST dissolved. The resultant solution is added to the salt solution containing acetate to give a red coloration. The acetate ion is confirmed by adding water to this mixture and warming till formation of red precipitate.

    Thiosulfates

    Thiosulfates are reducing agents. They can reduce \(I_{3}^{-}\) ions to iodide ions which are colorless. However, it should be noted that this is not a confirmatory test as any reducing agent can do this.

    \[2S_2O_3^{2-}\rightarrow S_4O_6^{2-}+2e^-\]

    \[I_3^- +2e^- \rightarrow 3I^-\]

    Hydronium ions

    Adding dilute acid to a solution with thiosulfate ions yields a solution with precipitate, elemental sulfur, and release a burnt rubber smell (gaseous sulfur dioxide)

    \[2H^+ + S_2O_3^{2+} \rightarrow S(s) + SO_2(g) + H_2O(l)\]

    Ferric ions

    It also produces a momentary violet coloration with ferric ions. --->Expand<--

    Copper sulfate

    With copper sulfate, a white ppt. is obtained which dissolves in excess of thiosulfate to produce a colorless solution through the formation of a stable complex.

    \[CuSO_4 + Na_2S_2O_3 \rightarrow CuS_2O_3 + Na_2SO_4\]

    This test has to be carried out in such a way that initially the concentration of thiosulfate is low which then gradually increases enough to form the complex. The solution containing thiosulfate should be poured into the copper sulfate to prevent the changes being missed.

    Silver Nitrate

    This is a very peculiar reaction, because on the surface it seems as if the precipitate obtained is changing colors!

    What actually happens is that when we add silver nitrate, a white ppt of silver thiosulfate is formed. This silver thiosulfate disproportionates to give silver sulfide and sulfuric acid, wherein as we know, silver sulfide is black. That is, on adding silver nitrate a white precipitate is obtained which changes to brown and finally to black.

    \[S_2O_3^{2-} + 2AgNO_3 \rightarrow 2NO_3^- + Ag_2S_2O_3\]

    \[Ag_2S_2O_3(s) + H_2O(l) \rightarrow H_2SO_4(aq) + Ag_2S\]

    Furthermore, in the presence of excess of thiosulfate ions the white ppt of Silver thiosulfate dissolves giving a colorless complex.

    Thiocyanate ions

    \(SCN^-\) ions have been discussed once in the testing of cyanide ions.

    In fact, one of the characteristic tests for thiocyanate ions is the same reaction that we discussed over there.

    Ferricthiocyanate

    With ferric ions, thiocyanate ions give a blood red coloration of \(Fe(SCN)_3\). This test is NOT given by ferrous ions - and hence is confirmatory for the presence of both ferric ions and for thiocyanate ions.

    -->Check lingo<--

    Vogel's test

    If Co2+ ions are added to the thiocyanate ions a blue solution of tettrathiocyanatocobaltate is obtained. This complex on addition mercuric ions gives a Blue crystalline precipitate.

    This ppt is NOT Mercuric tetrathiocyanatocobaltate! It is interesting to note that the actual precipitate is an isomer of the expected ppt - it is Cobalt(II) Tetrathiocyanatomercurate(II), that is the metal cations have switched places.

    If no effervescence is obtained with dilute HCl, heat the salt with Conc. H2SO4

    The following ions belong to this list:

    Detecting Fluorides

    Evolution of HF

    When we add concentrated sulfuric acid to fluoride ions, they are protonated to give hydrogen fluoride. This is because hydrogen fluoride is a weak acid, its anion F- can be protonated back to the gas.

    \[F^- + H_2SO_4 \rightarrow HF + HSO_4^-\]

    The extent to which hydrogen fluoride is produced is dependent on the initial concentration of the sulfuric acid. By Le Chatelier's equilibrium principle, a larger amount of the reactant H2SO4, coming from an increase in concentration, would drive the reaction towards the production of more gaseous hydrogen fluoride.

    The question is, how do we know that HF has evolved? When we hold a glass rod at the mouth of the testtube, a white waxy deposit is obtained on the glass rod. This is because of the reaction of HF with Silica. (Glass rod - is of course nearly all silica!).

    \(HF(g) + SiO_2(s) \rightarrow SiF_4(g) + H_2O(l)\)

    The silicon fluoride gas is easily hydrolyzed by the water released, and we get hydrogen silicofluoride and hydrogen silicate - the mixture which is the white waxy deposit in question.

    \[3SiF_4(g) + 4H_2O(l) \rightarrow 2H_2(SiF_6)(s) + H_4SiO_4(s)\]

    Question: Is is possible to detect HF or F- using Ca(NO3)2?

    \[2 F^- (aq) + Ca(NO_3)_2(aq) \rightarrow CaF_2 (s) + 2 NO_3^- (aq)\]

    Oxidation of fluorides

    In this particular test, We added conc. H2SO4. H2SO4 is an oxidizing agent. When it is added to a fluoride salt, does it oxidize to molecular fluorine? Let us think for a second. Yes, we accept that Conc. H2SO4 is a pretty strong oxidizing agent, but the question is, is it strong enough to oxidize fluoride ions? The product of oxidation of fluoride ions is molecular fluorine - something which is rarer than anything. Molecular fluorine reacts with almost anything. It causes mayhem wherever it goes. So our answer is no. H2SO4 just can't oxidize fluoride ions into molecular fluorine.

    -->Note to editors: This is the purpose of the text - to acquaint the reader with the various facts of chemistry and making them realize their importance in the real life. While this point doesn't exactly point out the high reactivity of fluorine, half a dozen scattered discussions can do enough to hammer this point!

    Possible insertions - the problems involved in the electrolytic extraction of fluorine.<--

    Chlorides

    Conc. H2SO4

    When Conc. H2SO4 is added, chloride salts evolve colorless HCl gas.

    Now the presence of this HCl gas can be detected in two ways. The first is to hold an ammonium hydroxide test tube near the neck of the original solution. The HCl gas, if evolved, will react with NH4OH to give dense white fumes of NH4Cl.

    Otherwise, we can use MnO2. Manganese in +4 state will be reduced to +2, oxidizing the chlorine in HCl to molecular chlorine gas - Cl2. Pale green (pale because of low concentration) Chlorine gas confirms that HCl was evolved.

    The Silver Nitrate test

    AgCl is a precipitate. So, if a silver nitrate solution is added to a sample of the being tested solution, a white ppt. is obtained. This precipitated Agcl is soluble in an excess of Ammonium Hydroxide soltion because of the formation of [Ag(NH3)2]+ complex ions.

    If we note carefully, the ammine complex is the same one that is used in the silver mirror test (Tollen's test) for the detection of aldehydes. The reaction, if we recall, was a redox reaction in which the silver (I) is reduced to metallic (0) silver mirror, and the aldehyde being oxidized to an ate ion.

    The Chromyl Chloride test

    It should be noted that this test is not for a solution containing the solvated ions. Chromyl chloride test is performed on the salt in solid state.

    Solid salt is mixed with Solid Potassium Dichromate and Conc. H2SO4. If the salt is a chloride, on heating deep red vapors of Chromyl Chloride (CrO2Cl2) are evolved.

    Some chemists still bear a doubt, because the deep red vapors bear some resemblance to NO2 and Br2 vapors. They can be distinguished, though.

    If these Chromyl Chloride vapors are passed through a dilute NaOH solution, it turns yellow. Most of us might have already guessed that the color is due to the formation of the Chromate ions. Indeed, the guess is true. This is then how we resolve our doubts and say that the salt is a chloride.

    It is not unwise to note here that most chromates are yellow in color, with the notable exceptions being AgCrO4 and Hg(CrO4)2. Both are red!

    Predominantly covalent chlorides like HgCl2 and alkyl chlorides do not offer this test. The real confirmation is adding acetic acid followed by lead acetate to the yellow solution which gives a pale yellow precipitate of lead chromate.

    Bromides/Iodides

    Conc. H2SO4

    With conc. H2SO4, the florides/chlorides gave HF/HCl respectively. However, the bromides and iodides are oxidized to molecular bromine and iodine respectively!

    So, in case of the reaction of iodides/bromides with Conc. H2SO4, bromides give deep brown vapors of Br2, while iodides give violet vapors of iodine.

    Silver nitrate solution

    With silver nitrate solution, bromides give a pale yellow ppt. of AgBr which is partially soluble in NH4OH. Iodides give a yellow ppt of AgI which is insoluble in ammonium hydroxide.

    -->Explain this point with reference to Solubility Product and formation constant of the concerned compounds. We need not give a numerical explanation! We just want to emphasize that Inorganic Chemistry is not plain mugging up.

    The Chloroform test

    Take the solution and add an equal amount of chloroform (CHCl3) to it. Add chlorine water and shake vigorously. Watch the change in color of the non-aqueous layer.

    If iodide/bromide ions are present, they would be oxidized to Iodine/Bromine respectively. Iodine would impart a violet color to the non-aqueous layer while if bromine is created it will give out a reddish color to the non-aqueous layer.

    Oxalate ions

    Oxalate ions are C2O42-

    Conc. H2SO4

    Concentrated sulfuric acid decomposes oxalate ions into Carbon monoxide and carbondioxide. Readers shouldn't take long to realize that the oxalate ions have decomposed through a disproportionation reaction. If the evolution of the two gases is confirmed, we can say that sulfuric acid reacted with oxalate ions.

    If the evolved gaseous mixture is passed through limewater, it will turn milky and finally colourless. Carbon Monoxide in the mixture will burn with a blue flame to produce carbon dioxide. The two gases are confirmed in this way, confirming the presence of the oxalate ion.

    Acidified Potassium Permanaganate solution

    Acidified KMnO4 is an oxidizing agent. It will oxidize the oxalate ions into carbon dioxide. In the process, it will decolorize because of the reduction of the violet manganate ions into colorless manganous ions.

    Now the question is, how do we know that the decolorization was due to oxalate ions? KMnO4 solution is decolorized each time the managanate ions oxidize ANYTHING.

    The answer lies again, in the detection of the carbondioxide gas. Once the gas is confirmed, oxalate ions have been virtually confirmed.

    Calcium chloride solution

    Aqueous solution of oxalate ions give a white precipitate of calcium oxalate which is insoluble in acetic acid.

    Nitrates

    Decomposition by Heat or Strong Acid

    All nitrate salts are decomposed by heat or dilute sulfuric acid to produce reddish brown fumes of nitrogen dioxide.

    \[2Cu(NO_3)_2(aq) \rightarrow 2CuO(aq) + 4NO_2(g) + O_2(g)\]

    Copper chips

    Aqueous salt solution when added with conc. sulfuric acid, and then copper chips are added, Dark brown solution forms.

    The brown ring test

    When a freshly prepared acidified solution of ferrous sulfate is added to a solution of nitrate ions, in the presence of concentrated H2SO4, a deep brown ring is formed.

    This deep brown ring is due to the formation of the complex sphere [Fe(H2O)5NO]2+. In this form, the Fe ion is in both its +1 and +2 oxidation states, and oscillates between the two. This means that even a slight disturbance in the medium will disturb the complex. It is for this reason that the acidified ferrous sulfate is added drop-wise and allowed to roll down the walls of the test tube, this minimises any disturbance that would destroy the ring. It reacts upon contact with the medium, and forms the mentioned complex, but only at the edges, forming the characteristic brown ring.

    The oscillation occurs in the following two ways

    i) The NO ion loses an electron to the Fe2+ ion forming Fe1+ and NO1+ (\(NO \rightarrow NO^{1+} + e^{1-}\) and \(Fe^{2+} + e^{1-} \rightarrow Fe^{1+}\))

    ii) The Fe1+ returns the electron to the NO1+ forming NO and Fe2+ (\(Fe^{1+} \rightarrow Fe^{2+} + e^{1-}\) and \(NO^{1+} + e^{1-} \rightarrow NO\))

    It is worth noting that it is very rare to find iron in its (I) oxidation state, (II) and (III) are much more common

    Reduction to nitrites

    Reduction to Ammonia

    The NItrate ion can easily be reduced to ammonia with either Devardas Alloy or Aluminium Foil. The aluminium is a very powerful reducing agent, and this combined with heating causes the nitrate ions to form ammonia gas. This can be tested for by holding a piece of damp red litmus paper over the end of the test tube. The ammonia will form alkaline ammonium ions in the water and turn the paper blue.

    \[4NO^-_3(aq) + 6H_2O(l) \rightarrow 4NH_3(g) + 9O_2(g)\]

    Aluminium powder is not shown as it merely catalyses the reaction.

    sulfates

    Barium Chloride

    The addition of a sulfate to a barium chloride solution produces a white precipitate of barium sulfate. This precipitate is very insoluble, not dissolving in any common laboratory reagents.

    Magnesium sulfate + Barium Chloride -> Barium sulfate + Magnesium Chloride

    \[MgSO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + MgCl_2(aq)\]

    Lead Nitrate/Lead Acetate

    A similar reaction to the above, a white precipitate of lead sulfate is formed. Again this is very difficult to dissolve in common reagent.

    \[PbNO_3(aq) + MgSO_4(aq) \rightarrow PbSO_4(s) + MgNO_3(aq)\]

    \[(CH_3COO)_2Pb(aq) + K_2SO_4(aq) \rightarrow PbSO_4(s) + 2CH_3COOK(aq)\]

    Manganate Ion

    H2S test

    The oxidising properties of potassium manganate allow it to oxidise hydrogen sulfide rapidly from oxidation state (II) to 0, reducing itself to manganese ions. These do not impart any colour to the solution, yet the sulfur precipitate forms a white cloud as the purple colouration of manganate ions disappears.

    \[2S^{2-} \rightarrow S_2 + 4e^{-1}\] and \[MnO_4^{2-} + 4e^- \rightarrow Mn^{2+} + 2O_2\]

    Borates

    Ethyl Alcohol

    on adding ethyl alchol we get ethyl borate on expose to flame it gives green colour flame

    Ortho-Phosphates

    Ammonium Molybdate

    When added to phosphate ions, warmed with nitric acid, forms a canary yellow precipitate of ammonium phosphomolybdate. Confirmatory test for presence of phosphates.
    \[PO_4^{3-} + 12(NH_4)_2MoO_4 + 21HNO_3 + 3H^+ \rightarrow (NH_4)_3PO_4·12MoO3↓ + 21HN4NO3 + 12H2O\]

    Carbonates and Bi-Carbonates

    We first react it with HCl, CO2 is emitted, so we know it is carbonate or bicarbonate MgSO4 is added to the salt solution, white precipitate in cold indicates carbonate and white precipitate on heating confirms bicarbonate

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