Traffic lights

In this simple demonstration, potassium sodium 2,3-dihydroxybutanedioate or potassium sodium tartrate (Rochelle salt) is oxidised by hydrogen peroxide in the presence of cobalt(II) ions. You simply mix the hot reactants and add the catalyst.

C₄H₄O₆^2-(aq) + 3H₂O₂(aq) → 2CO₂(g) + 2HCO^2-(aq)+ 4H₂O(l)

The interesting feature of this reaction is the colour change which occurs as the reaction proceeds. The initial colour is pink which changes to dark green and then back to pink again - hence the traffic light reaction. When you see the coloured intermediate, you can try to stabilise it by cooling the mixture rapidly. The nature of this intermediate can lead to some stimulating discussion about colour, thermodynamic and kinetic stability.

Kit

• Gloves; spatula
• Beaker, 250cm³
• measuring cylinder, 25cm³
• Distilled water
• thermometer (0–100°C)
• Bunsen burner and mat; tripod and gauze
• Stirring rod
• teat pipette
• two test tubes and rack
• Potassium sodium 2,3-dihydroxybutanedioate, CO₂K(CHOH)₂CO₂Na.4H₂O – also known as potassium sodium tartrate (Rochelle salt), 1g 
• Hydrogen peroxide (20 vol), 20 cm³
• Cobalt(II) chloride, CoCl₂.6H₂O, 0.25g (

Procedure

1. Dissolve ca 1g of potassium sodium 2,3-dihydroxybutanedioate in ca 50cm³ of water in a beaker.
2. Heat the solution to about 65°C (avoid going above 70°C).
3. Add ca 20 cm³ hydrogen peroxide solution and reheat to 65°C.
4. Note any sign of a reaction. Dissolve ca 0.25g of cobalt(II) chloride in 5cm³ distilled water in a test tube and add to the hot reaction mixture.There will be an induction period before the reaction proceeds.
5. As soon as the solution appears dark green, quickly transfer a small portion, using a teat pipette, into a test tube which is placed in a salted ice-bath.

Tips

If the reaction gets too hot or the reactants are too concentrated, effervescence from the reaction could cause the mixture to spill out of the beaker.

Teaching goals

This experiment can lead to a full-scale kinetic investigation by changing the concentration of the reactants and the catalyst. It can be used to develop a deeper understanding of non-standard electrode potentials and their use in predicting the feasibility of a reaction:

2HCO₂^-(aq) + 2CO₂(g) + 6H⁺(aq) + 6e^-→ C₄H₄O₆^2-(aq) + 2H₂O(l) E^⦵ = +0.20V

H₂O₂(aq) + 2H⁺(aq) + 2e^-→ 2H₂O(l) E^⦵ = +1.77V

The reaction is energetically favourable because ΔE^⦵ is +1.57V, which is large enough for the reaction to go to completion. However, being thermodynamically feasible does not mean the reaction is kinetically favourable and the reaction is very slow even when heated because of a high kinetic barrier. The catalyst will provide an alternative reaction pathway with a lower activation energy.

H₂O₂(aq) + 2H⁺(aq) + 2Co²⁺(aq) → 2H₂O(l) + 2Co³⁺(aq)

C₄H₄O₆^2-(aq) + 2H₂O(l) + 6Co³⁺(aq) → 2HCO₂^-(aq) + 2CO₂(g) + 6H⁺(aq) + 6Co²⁺(aq)

For this mechanism to work, the standard electrode potential (SEP) for the Co³⁺/Co²⁺ half-cell must lie within a certain range of values as indicated above (+0.20 V to +1.77V). However, the SEP for the half-cell Co³⁺(aq), Co²⁺(aq) | Pt is +1.84V. At first glance, you might think that in non-standard conditions this value is not that far out of the range required, but concentration or temperature changes do not alter ΔE values very much. The main reason is cobalt ions can form complexes with the 2,3-dihydroxybutanedioate ions - a bidentate ligand. Electrode potentials for half-cells involving complexes are often substantially different from those involving simple ions. This is another fruitful area for discussion (crystal and molecular orbital theory).

Unfortunately, the value for this half-cell is not quoted in the literature but even if it were shown to be energetically favourable, there would still be no way of knowing whether the reaction was kinetically favourable. You would have to do this experiment to find out. Finally, you should note that cobalt(II) catalyses the decomposition of hydrogen peroxide into water and oxygen and the equations written here do not take that into account. Clearly, this side reaction will influence any thermodynamic or kinetic discussion with this reaction.

Further reading

1. A. Lainchbury et al, ILPAC, 2nd edn, vol 11. London: John Murray, 1996.

This article was edited in June 2018