Why Purple Permanganate Turns Colourless in Acid

Potassium dichromate and potassium permanganate are the strong oxidising agents whose colour changes signal the end of a titration, while the lanthanoids and actinoids — the two rows at the foot of the periodic table — include the metals in powerful magnets and the uranium in nuclear fuel.

This part covers potassium dichromate, potassium permanganate, the lanthanoids with lanthanoid contraction, and the actinoids.

Potassium dichromate is made by fusing chromite ore with sodium carbonate in air, acidifying the sodium chromate formed and treating the sodium dichromate with KCl; the dichromate ion is two tetrahedra sharing an oxygen, it is a strong oxidising agent in acid, and changing the pH converts orange dichromate and yellow chromate into each other.

Preparation:

$$4{\text{FeCr}}_2{\text{O}}_4+8{\text{Na}}_2{\text{CO}}_3+7{\text{O}}_2\to 8{\text{Na}}_2{\text{CrO}}_4+2{\text{Fe}}_2{\text{O}}_3+8{\text{CO}}_2$$

$$2{\text{Na}}_2{\text{CrO}}_4+2{\text{H}}^{+}\to {\text{Na}}_2{\text{Cr}}_2{\text{O}}_7+2{\text{Na}}^{+}+{\text{H}}_2\text{O}$$

$${\text{Na}}_2{\text{Cr}}_2{\text{O}}_7+2\text{KCl}\to {\text{K}}_2{\text{Cr}}_2{\text{O}}_7+2\text{NaCl}$$

Effect of pH:

$$2{\text{CrO}}_4^{2-}+2{\text{H}}^{+}\rightleftharpoons {\text{Cr}}_2{\text{O}}_7^{2-}+{\text{H}}_2\text{O}$$

Acid pushes it towards orange dichromate; alkali pulls it back to yellow chromate. Chromium stays $+6$ throughout.

Oxidising action in acid:

$${\text{Cr}}_2{\text{O}}_7^{2-}+14{\text{H}}^{+}+6e^{-}\to 2{\text{Cr}}^{3+}+7{\text{H}}_2\text{O}$$

It oxidises Fe${}^{2+}$, I${}^{-}$ and H${}_2$S, and is a primary standard in titrations.

Worked example. One mole of dichromate gains $6$ electrons, so it oxidises $6$ mol of Fe${}^{2+}$. Oxidising $0.012$ mol of Fe${}^{2+}$ needs $0.002$ mol, or $0.002\times 294=0.588$ g of K${}_2$Cr${}_2$O${}_7$.

An everyday example. Older breath-alcohol test tubes relied on orange dichromate turning green as it oxidised ethanol.

The substance. The orange-to-yellow change is not a redox reaction — chromium remains in the $+6$ state.

**Potassium permanganate is made by fusing MnO${}_2$ with KOH and an oxidising agent to form green manganate, which is then converted into purple permanganate; it is a powerful oxidising agent whose product depends on the medium — Mn${}^{2+}$ in acid, MnO${}_2$ in neutral or faintly alkaline solution.

Preparation:**

$$2{\text{MnO}}_2+4\text{KOH}+{\text{O}}_2\to 2{\text{K}}_2{\text{MnO}}_4+2{\text{H}}_2\text{O}$$

$$3{\text{MnO}}_4^{2-}+4{\text{H}}^{+}\to 2{\text{MnO}}_4^{-}+{\text{MnO}}_2+2{\text{H}}_2\text{O}$$

Structure. Permanganate is tetrahedral, with manganese at $+7$. Its intense purple colour comes from charge transfer, not d–d transitions, since Mn(VII) has no d electrons.

Oxidising action:

- Acidic — ${\text{MnO}}_4^{-}+8{\text{H}}^{+}+5e^{-}\to {\text{Mn}}^{2+}+4{\text{H}}_2\text{O}$; oxidises Fe${}^{2+}$, oxalate and I${}^{-}$
- Neutral or faintly alkaline — ${\text{MnO}}_4^{-}+2{\text{H}}_2\text{O}+3e^{-}\to {\text{MnO}}_2+4{\text{OH}}^{-}$; oxidises I${}^{-}$ to IO${}_3^{-}$

Worked example. In acid, $1$ mol of MnO${}_4^{-}$ oxidises $5$ mol of Fe${}^{2+}$. So $20.0$ mL of $0.02$ M KMnO${}_4$ reacts with

$$0.0200\times 0.02\times 5=2.0\times 10^{-3}{\text{ mol of Fe}}^{2+}$$

Uses: titrations, disinfecting water, and bleaching wool, cotton and silk.

An everyday example. **A few pink crystals of potassium permanganate, often called lal dawa, are added to water to rinse vegetables in many Indian homes.

The substance. KMnO${}_4$ acts as its own indicator** — the first drop in excess leaves a lasting pink colour, so no separate indicator is needed.

**Lanthanoids fill the $4f$ orbitals, with general configuration $[\text{Xe}]4f^{1-14}5d^{0-1}6s^2$; their usual oxidation state is $+3$, they are reactive metals, and their radii shrink steadily across the series — lanthanoid contraction — because $4f$ electrons shield the rising nuclear charge poorly.

Oxidation states:

- $+3$ is the most common
- Ce${}^{4+}$** reaches the empty $4f^0$ configuration, so it is an oxidising agent
- **Eu${}^{2+}$ and Yb${}^{2+}$** have $4f^7$ and $4f^{14}$, so they are reducing agents

Reactivity. Lanthanoids tarnish in air, release hydrogen from dilute acids, burn in halogens, form basic hydroxides, and combine with carbon, nitrogen and sulphur when heated.

Consequences of lanthanoid contraction:

- Second and third series radii are nearly equal — zirconium $160$ pm and hafnium $159$ pm
- Lanthanoids are hard to separate, because their sizes and properties are so alike
- Basic strength of the hydroxides falls from La(OH)${}_3$ to Lu(OH)${}_3$

Worked example. Gadolinium ($Z=64$) is $[\text{Xe}]4f^{7}5d^{1}6s^2$. Losing the $5d$ and both $6s$ electrons gives Gd${}^{3+}$ as $4f^7$, with $7$ unpaired electrons:

$$\mu =\sqrt{7\times 9}\approx 7.94\text{ BM}$$

An everyday example. The flint in a gas lighter is misch metal, an alloy of lanthanoids with a little iron, which throws sparks when struck.

The substance. Lanthanoid contraction affects the elements after the series, which is why zirconium and hafnium occur together and are so hard to separate.

**Actinoids fill the $5f$ orbitals, show a wider range of oxidation states than lanthanoids because $5f$, $6d$ and $7s$ energies are close, contract more strongly across their series, are more reactive and are all radioactive; d- and f-block elements supply many vital metals and compounds.

Actinoids versus lanthanoids:

- Configuration** — actinoids fill $5f$; lanthanoids fill $4f$
- Oxidation states — actinoids show many, up to $+7$; lanthanoids mostly $+3$
- Contraction — greater for actinoids, because $5f$ electrons shield even less effectively
- Reactivity — actinoids are more reactive, especially when finely divided
- Radioactivity — every actinoid is radioactive, which makes them hard to study

Applications:

- Titanium — light, strong alloys for aircraft
- Vanadium(V) oxide and nickel — industrial catalysts
- Uranium — nuclear fuel
- Lanthanoid oxides — phosphors in television screens

Worked example. Uranium ($Z=92$) is $[\text{Rn}]5f^{3}6d^{1}7s^2$. Removing all six of these outer electrons gives the $+6$ state found in the uranyl ion, UO${}_2^{2+}$ — far beyond the usual $+3$ of the lanthanoids.

An everyday example. India's nuclear power stations use uranium, an actinoid, as their fuel.

The substance. **The wide range of actinoid oxidation states comes from the small energy gaps between $5f$, $6d$ and $7s$ orbitals**, which let many electrons take part in bonding.