Why Potassium Dichromate Matters
The transition metals owe much of their chemistry to the way they form oxides and oxoanions in high oxidation states. As the oxidation number of a metal rises, the ionic character of its oxide falls and its acidic character grows; thus $\ce{CrO3}$ behaves as an acidic oxide and dissolves to give chromic acid, $\ce{H2CrO4}$, and dichromic acid, $\ce{H2Cr2O7}$. The salts of these acids — chromates and dichromates — are the practical chromium(VI) reagents of the laboratory, and potassium dichromate is the most important of them.
$\ce{K2Cr2O7}$ is described by NCERT as "a very important chemical used in the leather industry and as an oxidant for preparation of many azo compounds." It is the reagent of choice as a primary standard in volumetric (redox titration) analysis, precisely because it can be obtained pure and stable. Every facet of the compound — its synthesis, equilibrium behaviour, structure and redox chemistry — appears in NEET-relevant inorganic chemistry, which is why this subtopic rewards careful study.
The chromium oxidation state never changes during interconversion
A frequent error is to assume that chromium is reduced or oxidised when chromate converts to dichromate. It is not. In both $\ce{CrO4^2-}$ and $\ce{Cr2O7^2-}$, chromium is in the +6 state. The interconversion is an acid–base (condensation) equilibrium, not a redox process. Reduction to $\ce{Cr^3+}$ happens only when dichromate acts as an oxidant in acidic medium.
Interconversion → oxidation state stays +6. Oxidising action → Cr drops from +6 to +3.
Preparation from Chromite Ore
Dichromates are not mined directly; they are manufactured from chromite ore, $\ce{FeCr2O4}$ (also written $\ce{FeO.Cr2O3}$), in a sequence of four conceptual steps. The route is a chain of conversions: chromite → sodium chromate → sodium dichromate → potassium dichromate.
Step 1 — Fusion of chromite to sodium chromate
Finely powdered chromite is fused with sodium carbonate in the presence of a free supply of air (excess oxygen). Chromium(III) in the ore is oxidised to chromium(VI), giving the yellow soluble salt sodium chromate:
$$\ce{4 FeCr2O4 + 8 Na2CO3 + 7 O2 -> 8 Na2CrO4 + 2 Fe2O3 + 8 CO2}$$
In the large-scale process a little quicklime is added to keep the roasting mass porous and prevent fusion into a sticky cake. The roasted product is leached with water: the soluble yellow sodium chromate dissolves, leaving insoluble ferric oxide, $\ce{Fe2O3}$, behind to be filtered off.
Step 2 — Acidification to sodium dichromate
The clarified yellow solution of sodium chromate is acidified with sulphuric acid. The chromate ions condense to dichromate, and orange sodium dichromate dihydrate, $\ce{Na2Cr2O7.2H2O}$, can be crystallised:
$$\ce{2 Na2CrO4 + 2 H+ -> Na2Cr2O7 + 2 Na+ + H2O}$$
Step 3 — Conversion to potassium dichromate
Sodium dichromate is far more soluble than the potassium salt — and being deliquescent it cannot be obtained as a clean, fixed-composition solid. Potassium dichromate is therefore made by treating a hot solution of sodium dichromate with potassium chloride:
$$\ce{Na2Cr2O7 + 2 KCl -> K2Cr2O7 + 2 NaCl}$$
Because $\ce{K2Cr2O7}$ is much less soluble in cold water than $\ce{NaCl}$, on cooling the orange crystals of potassium dichromate crystallise out and are separated. It is moderately soluble in cold water (about 100 g L⁻¹ at 298 K) but freely soluble in hot water (about 1000 g L⁻¹ at 373 K), so it is readily purified by recrystallisation.
| Step | Transformation | Key reaction / observation |
|---|---|---|
| 1 | Chromite → sodium chromate | 4FeCr2O4 + 8Na2CO3 + 7O2 → 8Na2CrO4 + 2Fe2O3 + 8CO2; Cr(III)→Cr(VI); yellow solution |
| 2 | Chromate → sodium dichromate | 2Na2CrO4 + 2H+ → Na2Cr2O7 + 2Na+ + H2O; acidify with H₂SO₄; orange |
| 3 | Sodium → potassium dichromate | Na2Cr2O7 + 2KCl → K2Cr2O7 + 2NaCl; less soluble K salt crystallises |
| 4 | Purification | Recrystallise from hot water (1000 g L⁻¹ at 373 K vs 100 g L⁻¹ at 298 K) |
"Free access of air" is doing real work
The fusion of chromite must occur in a free supply of air. The oxygen is the oxidant that lifts chromium from +3 in the ore to +6 in chromate. Without it, no chromate forms. NEET stems sometimes drop or alter this condition to test whether you understand it is an aerial oxidation, not a simple double decomposition.
The Chromate–Dichromate Equilibrium
Chromates and dichromates are interconvertible in aqueous solution depending on the pH. The yellow chromate ion dominates in basic (alkaline) solution; the orange dichromate ion dominates in acidic solution. This pH-driven colour change is one of the most cited equilibria in transition-metal chemistry.
On adding acid (lowering pH), chromate condenses to dichromate and the solution turns from yellow to orange:
$$\ce{2 CrO4^2- + 2H+ <=> Cr2O7^2- + H2O}$$
On adding alkali (raising pH), dichromate reverts to chromate and the solution returns to yellow:
$$\ce{Cr2O7^2- + 2 OH- <=> 2 CrO4^2- + H2O}$$
Throughout this equilibrium chromium remains in the +6 oxidation state in both ions, so it is a condensation (acid–base) equilibrium and not a redox reaction. The colour shift is the visible signature of the position of equilibrium.
Lowering the pH shifts the equilibrium toward orange dichromate; raising it shifts back to yellow chromate. The oxidation state of chromium is unchanged.
Permanganate is the other classic Cr/Mn oxidant. See Potassium Permanganate — Preparation, Properties, Structure for its tetrahedral $\ce{MnO4^-}$ and medium-dependent redox.
Structures of Chromate and Dichromate
The two ions differ only in how their $\ce{CrO4}$ building blocks are joined. The chromate ion, $\ce{CrO4^2-}$, is tetrahedral: a central chromium atom is surrounded by four oxygen atoms at the corners of a tetrahedron. The dichromate ion, $\ce{Cr2O7^2-}$, consists of two such tetrahedra sharing one corner oxygen — the bridging oxygen — through a Cr–O–Cr linkage. The Cr–O–Cr bridge is bent, with a bond angle of about 126°.
Each Cr remains four-coordinate and tetrahedral; in dichromate the two tetrahedra share the single bridging oxygen, giving the bent Cr–O–Cr linkage at ~126°.
| Feature | Chromate, CrO₄²⁻ | Dichromate, Cr₂O₇²⁻ |
|---|---|---|
| Colour in solution | Yellow | Orange |
| Favoured pH | Alkaline / neutral | Acidic |
| Oxidation state of Cr | +6 | +6 |
| Geometry around Cr | Tetrahedral | Two tetrahedra (each Cr tetrahedral) |
| Cr–O–Cr bridge | None | One bridging O; angle ≈ 126° |
| Cr atoms per ion | 1 | 2 |
Physical Properties
Potassium dichromate forms orange-red prismatic crystals. It has a melting point near 696 K and a specific gravity of about 2.68. Crucially, unlike the deliquescent sodium salt, $\ce{K2Cr2O7}$ is not hygroscopic and is obtained in a fixed, definite composition — the property that makes it usable as a primary standard. It is moderately soluble in cold water, far more soluble in hot water, and insoluble in alcohol.
Oxidising Action in Acidic Medium
Sodium and potassium dichromates are strong oxidising agents. In acidic solution the dichromate ion is reduced from chromium(VI) to chromium(III), accepting six electrons per ion:
$$\ce{Cr2O7^2- + 14 H+ + 6 e- -> 2 Cr^3+ + 7 H2O} \qquad (E^\circ = +1.33\,\text{V})$$
The large positive standard electrode potential, $E^\circ = +1.33\ \text{V}$, quantifies just how strongly acidified dichromate pulls electrons. Acidified potassium dichromate therefore oxidises iodides to iodine, sulphides to sulphur, tin(II) to tin(IV) and iron(II) to iron(III). The reducing-agent half-reactions are:
$$\ce{6 I- -> 3 I2 + 6 e-} \qquad \ce{3 Sn^2+ -> 3 Sn^4+ + 6 e-}$$
$$\ce{3 H2S -> 6 H+ + 3 S + 6 e-} \qquad \ce{6 Fe^2+ -> 6 Fe^3+ + 6 e-}$$
Adding the dichromate half-reaction to a reducing-agent half-reaction gives the full ionic equation. For iron(II), the most common titration case:
$$\ce{Cr2O7^2- + 14 H+ + 6 Fe^2+ -> 2 Cr^3+ + 6 Fe^3+ + 7 H2O}$$
In a redox titration, acidified $\ce{K2Cr2O7}$ oxidises $\ce{Fe^2+}$. How many moles of $\ce{Fe^2+}$ does one mole of $\ce{Cr2O7^2-}$ oxidise, and why?
One dichromate ion gains 6 electrons ($\ce{Cr2O7^2- + 14H+ + 6e- -> 2Cr^3+ + 7H2O}$). Each $\ce{Fe^2+ -> Fe^3+}$ releases one electron. To balance the six electrons accepted, six moles of $\ce{Fe^2+}$ are oxidised per mole of dichromate — exactly the 1 : 6 stoichiometry of the balanced equation above.
Acidic medium is mandatory for the +6 → +3 path
The $E^\circ = +1.33\ \text{V}$ half-reaction and the clean reduction to green $\ce{Cr^3+}$ apply in acidic medium (the 14 H⁺ in the equation make this explicit). Questions that omit "acidic" or hand you an alkaline solution are testing whether you remember that the strong-oxidant behaviour and the $\ce{Cr^3+}$ product are an acidic-medium phenomenon.
Why Cr(VI) Is a Strong Oxidant
The strength of dichromate as an oxidant is rooted in a periodic trend within group 6. For chromium, the lower +3 oxidation state is more stable than +6; consequently Cr(VI) readily accepts electrons to revert to the favoured $\ce{Cr^3+}$ state, so dichromate in acidic medium is a strong oxidising agent. By contrast, for the heavier congeners molybdenum and tungsten the highest oxidation state is itself stable, so $\ce{MoO3}$ and $\ce{WO3}$ show little tendency to be reduced and are not strong oxidants. This is a direct application of the rule that the stability of the higher oxidation state increases down a transition group.
Uses of Potassium Dichromate
The compound's combination of stable composition and powerful oxidising power makes it broadly useful:
| Use | Basis |
|---|---|
| Primary standard in volumetric analysis | Pure, non-hygroscopic, fixed composition; stable on storage |
| Oxidant in organic chemistry (e.g. azo-compound preparation) | Strong Cr(VI) → Cr(III) oxidation in acidic medium |
| Leather (chrome) tanning industry | Source of chromium for tanning processes |
| Chromyl chloride / salt-analysis confirmation of chloride | Reactivity of dichromate with chlorides under acidic conditions |
Because the closely related sodium salt is more soluble, sodium dichromate is the workhorse oxidant where solubility matters, whereas potassium dichromate's reliability as a weighable solid keeps it central to quantitative analysis.
Potassium Dichromate at a glance
- Preparation: chromite $\ce{FeCr2O4}$ → (fuse with $\ce{Na2CO3}$ + air) sodium chromate → (acidify, $\ce{H+}$) sodium dichromate → (treat with $\ce{KCl}$) potassium dichromate.
- Equilibrium: $\ce{2CrO4^2- + 2H+ <=> Cr2O7^2- + H2O}$; yellow (alkaline) ⇌ orange (acidic); Cr stays +6.
- Structure: $\ce{CrO4^2-}$ tetrahedral; $\ce{Cr2O7^2-}$ two tetrahedra sharing a corner O, Cr–O–Cr ≈ 126°.
- Oxidising action: $\ce{Cr2O7^2- + 14H+ + 6e- -> 2Cr^3+ + 7H2O}$, $E^\circ = +1.33$ V; oxidises $\ce{Fe^2+}$, $\ce{I-}$, $\ce{H2S}$, $\ce{Sn^2+}$.
- Why a strong oxidant: Cr(VI) is less stable than Cr(III); contrast with stable $\ce{MoO3}$, $\ce{WO3}$.
- Uses: primary standard, organic oxidant, leather tanning.