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The cobalt chloride colour equilibrium: Le Chatelier’s principle

In a tube of violet solution, pink [Co(H₂O)₆]²⁺ and blue [CoCl₄]²⁻ are in equilibrium. Add hydrochloric acid, water or silver nitrate, or dip it in hot or iced water; watch the colour change and explain each shift with Q and K.

Difficulty ●○○ ≈ 20 min

Aim Observe how concentration and temperature changes affect [Co(HX2O)X6]X2++4 ClX−⇌[CoClX4]X2−+6 HX2O\ce{[Co(H2O)6]^2+} + \ce{4Cl-} \mathrel{\ce{<=>}} \ce{[CoCl4]^2-} + \ce{6H2O}, and decide whether the forward reaction is endothermic or exothermic.
Real systems · colours computed from absorption spectra

Le Chatelier: disturb it and watch it re-equilibrate

[Co(HX2O)X6]X2+(aq)+4 ClX−(aq)⇌[CoClX4]X2−(aq)+6 HX2O(l)\ce{[Co(H2O)6]^2+(aq)} + \ce{4Cl^-(aq)} \mathrel{\ce{<=>}} \ce{[CoCl4]^2-(aq)} + \ce{6H2O(l)}

Endothermic: blue when hot, pink in ice water. K′ is a conditional constant calibrated to what the demonstration shows.

ControlYour tube · 25 °C
1.5 % of the cobalt is blue [CoCl₄]²⁻
Temperature
25 °C
Kc
5.00 × 10⁻⁶
Q / K
1
[[Co(HX2O)X6]X2+][\ce{[Co(H2O)6]^2+}]
0.165 M
[ClX−][\ce{Cl^-}]
7.49 M
[[CoClX4]X2−][\ce{[CoCl4]^2-}]
0.0026 M

Q = K: back at equilibrium

Concentrations over time (vertical lines = disturbances)

012345678Model time t / s020406080100Share of the cobalt / %
  • [Co(H₂O)₆]²⁺
  • [CoCl₄]²⁻
After each disturbance the concentrations relax by mass-action kinetics (treating the overall reaction as one step) to the new equilibrium; the end point comes from the exact ICE solver, with K(T) from van ’t Hoff at constant ΔH°. Relaxation is scaled to about 1 s so you can watch it; the real systems are far faster (N₂O₄ takes microseconds). Colours come from each species’ absorption bands via Beer–Lambert and CIE colorimetry: the NO₂ band is fitted to measured cross-sections; the cobalt K′ is a conditional constant calibrated to what the demonstration shows; phosphate “binding half the free Fe³⁺” is schematic.

Procedure

  1. The starting solution follows the Royal Society of Chemistry recipe: about 4 g of CoClX2 ⋅ 6 HX2O\ce{CoCl2.6H2O} in 40 mL of water, made up with 60 mL of concentrated hydrochloric acid (about 0.17 M Co(II) and 7.5 M chloride). The control tube on the left is never touched.
  2. Add concentrated hydrochloric acid to your tube and note the colour; then add water a few times and compare with the control.
  3. Press “Start again”, put your tube in hot water (heat twice), then in iced water (cool several times).
  4. Start again and add silver nitrate: white AgCl precipitates and removes chloride. Watch the colour and which curve moves on the concentration–time plot.
  5. Switch to the iron thiocyanate system and add Fe³⁺, SCN⁻ and hydrogen phosphate in turn; compare how the two systems respond.

What to notice

  • Blue when hot, pink when cold: heating raises K, so the forward reaction is endothermic (literature value about +62 kJ/mol).
  • Adding water turns it pink — water is a “product” yet does not appear in K: dilution hits the five solute particles on the left harder than the one on the right, making Q larger than K, so the equilibrium shifts left.
  • After each disturbance the concentrations bounce back part of the way, never all the way.
  • Because [CoCl₄]²⁻ absorbs light more than a hundred times more strongly than [Co(H₂O)₆]²⁺, the solution already looks bluish when only a few per cent of the cobalt is in the blue complex.

In a real lab

  • Concentrated hydrochloric acid is corrosive and fumes: work in a fume cupboard with goggles and gloves.
  • Cobalt chloride is toxic, may cause skin and respiratory sensitisation, and harms aquatic life: collect cobalt waste separately.
  • Heat in a water bath (80–90 °C) rather than directly in a flame, so the tube does not spit.
  • The simulation’s K′ is a conditional constant calibrated to the classroom demonstration; real solutions also contain intermediate complexes such as CoCl⁺ and CoCl₂, so the colour change is more gradual.

Related lessons

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