What Cation Group II Is
In the NCERT qualitative scheme the cations are sorted into six analytical groups plus the zero group, each defined not by chemical family but by a single group reagent that precipitates its members and leaves the rest in solution. Group II follows Group I (lead, precipitated as chloride by dilute HCl). After Group I is removed or shown absent, the original solution is treated with hydrogen sulphide in an acidic medium, and the ions that come down as insoluble sulphides constitute Group II.
Within the NCERT syllabus Group II contains $\ce{Pb^2+}$, $\ce{Cu^2+}$ and $\ce{As^3+}$. In the classical full Group II — useful background for understanding the chemistry — the group also includes $\ce{Bi^3+}$, $\ce{Cd^2+}$, $\ce{Sb^3+}$ and $\ce{Sn^{2+/4+}}$. The group is traditionally split into two sub-groups: Group II-A (the copper group), whose sulphides are insoluble in yellow ammonium sulphide, and Group II-B (the arsenic group), whose sulphides dissolve in it. For NEET, $\ce{Cu^2+}$ is the workhorse ion, so the chemistry below centres on copper while keeping the group picture intact.
| Group | Cations (NCERT) | Group reagent | Precipitate form |
|---|---|---|---|
| Zero | $\ce{NH4+}$ | None (separate test) | — |
| I | $\ce{Pb^2+}$ | Dilute HCl | Chloride |
| II | $\ce{Pb^2+}, \ce{Cu^2+}, \ce{As^3+}$ | H₂S in presence of dilute HCl | Sulphides |
| III | $\ce{Al^3+}, \ce{Fe^3+}$ | $\ce{NH4OH}$ in presence of $\ce{NH4Cl}$ | Hydroxides |
| IV | $\ce{Co^2+}, \ce{Ni^2+}, \ce{Mn^2+}, \ce{Zn^2+}$ | H₂S in presence of $\ce{NH4OH}$ | Sulphides |
| V | $\ce{Ba^2+}, \ce{Sr^2+}, \ce{Ca^2+}$ | $\ce{(NH4)2CO3}$ in presence of $\ce{NH4OH}$ | Carbonates |
| VI | $\ce{Mg^2+}$ | None (separate test) | — |
Notice that Groups II and IV use the same reagent — hydrogen sulphide — but in opposite media. That single difference, acidic versus ammoniacal, is the entire reason two ions as different as $\ce{Cu^2+}$ and $\ce{Zn^2+}$ end up in different analytical groups, and it is examined directly in the NEET matching questions on group numbers.
The Group Reagent: H₂S in Dilute HCl
The procedure (NCERT Experiment 7.1, Analysis of Group-II cations) is precise. After confirming Group I is absent, excess water is added, the solution is warmed, and hydrogen sulphide gas is passed through it for one to two minutes with the test tube shaken. The medium is kept acidic by the dilute HCl already present from the Group I stage. A precipitate at this point signals Group II; more $\ce{H2S}$ is then passed to complete the precipitation.
Hydrogen sulphide is a weak diprotic acid. In water it ionises only slightly, and that ionisation is the source of the sulphide ion that does the precipitating:
$$\ce{H2S <=> 2H+ + S^2-}$$
The defining feature of Group II is that this equilibrium is deliberately pushed to the left by the dilute HCl. The reactions that bring the cations down are simply the sulphides exceeding their solubility products:
$$\ce{Cu^2+ + H2S -> CuS v + 2H+}$$
$$\ce{Pb^2+ + H2S -> PbS v + 2H+}$$
$$\ce{2As^3+ + 3H2S -> As2S3 v + 6H+}$$
The acidic medium is what makes the split clean: only the lowest-Ksp sulphides precipitate, the rest pass through to be caught later in Group IV.
Ksp and the Common-Ion Logic
NCERT names the two principles at the head of the experiment: the solubility product and the common-ion effect. They work together. Precipitation occurs only when the ionic product of a sparingly soluble salt exceeds its solubility product $K_{sp}$. For a metal sulphide:
$$\ce{MS <=> M^2+ + S^2-}, \qquad K_{sp} = [\ce{M^2+}][\ce{S^2-}]$$
The dissolved metal ion concentration $[\ce{M^2+}]$ is fixed by the sample. The lever the analyst controls is $[\ce{S^2-}]$, and that is set by the $\ce{H2S}$ equilibrium:
$$[\ce{S^2-}] = \dfrac{K_a\,[\ce{H2S}]}{[\ce{H+}]^2}$$
Because $[\ce{H+}]$ appears squared in the denominator, raising the acid concentration with dilute HCl crushes the sulphide-ion concentration. This is the common-ion effect in action: $\ce{H+}$ is common to the $\ce{H2S}$ ionisation, and adding it suppresses that ionisation, so only a trace of $\ce{S^2-}$ survives. With $[\ce{S^2-}]$ tiny, only those salts whose $K_{sp}$ is small enough — the Group II sulphides — can still satisfy $[\ce{M^2+}][\ce{S^2-}] > K_{sp}$ and precipitate.
"More acid means more precipitate" — wrong direction
Students often reason that since $\ce{H2S}$ supplies $\ce{S^2-}$, adding more acid should give more precipitate. The opposite is true: dilute HCl suppresses sulphide-ion formation. Its job is to keep $[\ce{S^2-}]$ low so that only Group II comes down. Too much acid can even prevent Group II precipitation; too little lets Group IV slip in early.
In Group II, $\ce{H+}$ is a brake on $[\ce{S^2-}]$, not an accelerator. Low $[\ce{S^2-}]$ = selectivity.
Why Group II Separates from Group IV
This is the single most examined idea of the topic. Group II and Group IV sulphides differ enormously in solubility product. The Group II sulphides ($\ce{CuS}$, $\ce{PbS}$, $\ce{CdS}$, $\ce{Bi2S3}$) have extremely small $K_{sp}$ values; the Group IV sulphides ($\ce{ZnS}$, $\ce{CoS}$, $\ce{NiS}$, $\ce{MnS}$) have much larger ones.
In acidic medium $[\ce{S^2-}]$ is held so low that only the Group II sulphides exceed their $K_{sp}$. The Group IV ions remain dissolved. Later, in Group IV, $\ce{H2S}$ is passed in an ammoniacal medium: $\ce{NH4OH}$ removes $\ce{H+}$, the suppression is lifted, $[\ce{S^2-}]$ rises sharply, and now even the higher-$K_{sp}$ sulphides precipitate. The same gas, two media, two groups.
| Feature | Group II | Group IV |
|---|---|---|
| Medium | Acidic (dilute HCl present) | Ammoniacal ($\ce{NH4OH}$ present) |
| $[\ce{S^2-}]$ available | Very low (suppressed) | High (suppression lifted) |
| $K_{sp}$ of sulphides precipitated | Very small | Larger |
| NCERT cations | $\ce{Pb^2+}, \ce{Cu^2+}, \ce{As^3+}$ | $\ce{Co^2+}, \ce{Ni^2+}, \ce{Mn^2+}, \ce{Zn^2+}$ |
| Example sulphide | $\ce{CuS}$ (black) | $\ce{ZnS}$ (white/dirty white) |
The ammoniacal-medium sulphides come down in Cation Group IV (Co²⁺, Ni²⁺, Mn²⁺, Zn²⁺) — the counterpart that completes the H₂S story.
The Coloured Sulphide Precipitates
A useful diagnostic feature of Group II is that its sulphides are coloured, and the colour itself narrows the field before any confirmatory test. NCERT records that a black precipitate on passing $\ce{H2S}$ in dilute HCl indicates $\ce{Cu^2+}$ or $\ce{Pb^2+}$, while a yellow precipitate indicates $\ce{As^3+}$.
| Cation | Sulphide | Colour | Sub-group |
|---|---|---|---|
| $\ce{Cu^2+}$ | $\ce{CuS}$ | Black | II-A (copper) |
| $\ce{Pb^2+}$ | $\ce{PbS}$ | Black | II-A (copper) |
| $\ce{Cd^2+}$ (classical) | $\ce{CdS}$ | Yellow | II-A (copper) |
| $\ce{Bi^3+}$ (classical) | $\ce{Bi2S3}$ | Brown / black | II-A (copper) |
| $\ce{As^3+}$ | $\ce{As2S3}$ | Yellow | II-B (arsenic) |
The sub-group split is then made with yellow ammonium sulphide, $\ce{(NH4)2S2}$. The Group II precipitate is shaken with excess of it: the copper-group sulphides (II-A) are insoluble and remain as residue, whereas the arsenic-group sulphide (II-B) dissolves as a soluble thio-salt. For arsenic this dissolution is:
$$\ce{As2S3 + 3(NH4)2S2 -> 2(NH4)3AsS4 + S}$$
So insolubility in yellow ammonium sulphide points to II-A (where $\ce{Cu^2+}$ and $\ce{Pb^2+}$ live), and solubility points to II-B (arsenic). Once a black, II-A-confirming precipitate is in hand and lead has been ruled out or removed, the copper confirmatory sequence begins.
Zinc sulphide is not a Group II precipitate
A frequent error is to expect $\ce{ZnS}$ in Group II because zinc forms a sulphide with $\ce{H2S}$. It does not appear here: $\ce{ZnS}$ has too large a $K_{sp}$, and in the acid-suppressed, low-$[\ce{S^2-}]$ medium its ionic product never crosses $K_{sp}$. Zinc waits for Group IV.
Black/yellow sulphide in dilute HCl = Group II. Zinc, cobalt, nickel, manganese = Group IV.
Confirmatory Tests for Cu²⁺
Copper is the representative Group II ion in NCERT, and its confirmation is a two-step classic. First the black copper sulphide is brought into solution; then two independent colour reactions seal the identification.
Step 1 — Dissolve CuS in nitric acid
Boiling the black precipitate with dilute nitric acid dissolves it as copper nitrate, with sulphur set free:
$$\ce{3CuS + 8HNO3 -> 3Cu(NO3)2 + 2NO ^ + 3S + 4H2O}$$
On prolonged heating the liberated sulphur is oxidised to sulphate and the solution turns blue with copper sulphate forming.
Step 2 — Deep-blue ammonia complex
A little $\ce{NH4OH}$ first precipitates a basic copper sulphate, which then redissolves in excess ammonia to give the diagnostic deep-blue tetraamminecopper(II) solution:
$$\ce{Cu(OH)2.CuSO4 + 8NH3 -> 2[Cu(NH3)4]SO4 + 2OH- + SO4^2-}$$
The intense royal-blue colour of $\ce{[Cu(NH3)4]^2+}$ is itself strong evidence for copper.
Step 3 — Chocolate-brown ferrocyanide test
The blue solution is acidified with acetic acid and treated with potassium ferrocyanide, $\ce{K4[Fe(CN)6]}$. A chocolate-brown precipitate of copper(II) hexacyanoferrate(II) confirms $\ce{Cu^2+}$ beyond doubt:
$$\ce{[Cu(NH3)4]SO4 + 4CH3COOH -> CuSO4 + 4CH3COONH4}$$
$$\ce{2CuSO4 + K4[Fe(CN)6] -> Cu2[Fe(CN)6] v + 2K2SO4}$$
Two confirmations in one chain: the deep-blue ammine and the chocolate-brown ferrocyanide both point to Cu²⁺. The second precipitate is copper(II) hexacyanoferrate(II), Cu₂[Fe(CN)₆].
Worked Confirmation of a Copper Salt
A salt solution gives no precipitate with dilute HCl. On warming and passing $\ce{H2S}$ a black precipitate forms, insoluble in yellow ammonium sulphide. Trace the reasoning that confirms the cation.
1. No Group I. No white precipitate with dilute HCl rules out $\ce{Pb^2+}$ in Group I.
2. Group II present. A precipitate on passing $\ce{H2S}$ in the acidic (dilute-HCl) medium places the cation in Group II.
3. Black + insoluble in (NH₄)₂S₂. Black colour means $\ce{Cu^2+}$ or $\ce{Pb^2+}$; insolubility in yellow ammonium sulphide confirms sub-group II-A, not arsenic. With Group I already absent, copper is strongly indicated: $\ce{Cu^2+ + H2S -> CuS v + 2H+}$.
4. Dissolve and test. Boil with dilute $\ce{HNO3}$: $\ce{3CuS + 8HNO3 -> 3Cu(NO3)2 + 2NO + 3S + 4H2O}$. Add excess $\ce{NH4OH}$ — a deep-blue $\ce{[Cu(NH3)4]^2+}$ solution appears.
5. Seal it. Acidify the blue solution with acetic acid and add $\ce{K4[Fe(CN)6]}$: a chocolate-brown precipitate, $\ce{2CuSO4 + K4[Fe(CN)6] -> Cu2[Fe(CN)6] v + 2K2SO4}$, confirms $\ce{Cu^2+}$.
Observation → Inference Table
The whole Group II workflow for copper can be read off as a chain of observations and the inference each one licenses.
| Observation | Inference |
|---|---|
| No precipitate with dilute HCl | Group I ($\ce{Pb^2+}$ as chloride) absent |
| Precipitate on passing $\ce{H2S}$ in dilute HCl | A Group II cation is present |
| Precipitate is black | $\ce{Cu^2+}$ or $\ce{Pb^2+}$ (not $\ce{As^3+}$, which is yellow) |
| Precipitate insoluble in yellow $\ce{(NH4)2S2}$ | Sub-group II-A (copper group), not II-B (arsenic) |
| Black ppt dissolves in dilute $\ce{HNO3}$, solution turns blue on boiling | $\ce{Cu^2+}$ as $\ce{Cu(NO3)2}$ / $\ce{CuSO4}$ |
| Deep-blue solution with excess $\ce{NH4OH}$ | $\ce{[Cu(NH3)4]^2+}$ — copper indicated |
| Chocolate-brown precipitate with $\ce{K4[Fe(CN)6]}$ (acetic-acid medium) | $\ce{Cu^2+}$ confirmed ($\ce{Cu2[Fe(CN)6]}$) |
Group II in one glance
- Group reagent: $\ce{H2S}$ passed in the presence of dilute HCl — acidic medium.
- NCERT cations: $\ce{Pb^2+}, \ce{Cu^2+}, \ce{As^3+}$; classically also $\ce{Bi^3+}, \ce{Cd^2+}, \ce{Sb^3+}, \ce{Sn}$.
- Why acidic: HCl supplies $\ce{H+}$, which suppresses $\ce{H2S}$ ionisation (common-ion effect), keeping $[\ce{S^2-}]$ very low so only low-$K_{sp}$ sulphides precipitate.
- Group II vs IV: same gas, acidic for II (low $[\ce{S^2-}]$, low-$K_{sp}$ sulphides) and ammoniacal for IV (high $[\ce{S^2-}]$, higher-$K_{sp}$ sulphides). $\ce{ZnS}$ belongs to IV.
- Colours: $\ce{CuS}$, $\ce{PbS}$ black; $\ce{CdS}$, $\ce{As2S3}$ yellow. II-A insoluble in yellow $\ce{(NH4)2S2}$; II-B soluble.
- Cu²⁺ confirmation: dissolve $\ce{CuS}$ in $\ce{HNO3}$ → deep-blue $\ce{[Cu(NH3)4]^2+}$ with excess ammonia → chocolate-brown $\ce{Cu2[Fe(CN)6]}$ with $\ce{K4[Fe(CN)6]}$.