In this series, Colin Baker of Bedford School provides spectacular demonstrations, designed to capture the student's imagination. The demonstrations are easy to prepare, safe to dispose of and they work.
Nothing tends to imprint chemical facts upon the mind so much as the exhibition of interesting experiments - Samuel Parkes, 1816
A catalyst increases the rate of a chemical reaction without itself being permanently changed by the reaction. The emphasis on permanently changed is important when discussing homogeneous catalysis because the catalyst does undergo chemical change in providing an alternative reaction pathway of lower energy. The ability to change oxidation state for transition metals makes them suitable as homogenous catalysts. A new activated complex is formed with the catalyst because transition metals can form stable compounds in more than one oxidation state and the transition metal ions can therefore readily move between oxidation states. During the catalysed reaction the transition metal (TM) ion is oxidised by one reactant to a higher oxidation state. This is then reduced back to the original form by reaction with the other reactant. The reactants are therefore converted to the same products as are formed without the catalyst. The only difference is that the reactants are converted into products more quickly.
Reactant 1 + TM ion in low oxidation state Product + TM ion in high oxidation state
Reactant 2 + TM ion in high oxidation state Product + TM ion in low oxidation state
The traffic light reaction .
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.
C4H4O62- (aq) + 3H2O2(aq) 2CO2(g) + 2HCO2-(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.
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), 20cm3;
Cobalt(II) chloride, CoCl2.6H2O , 0.25g.
Dissolve ca 1g of potassium sodium 2,3-dihydroxybutanedioate in ca 50cm3 of water in a beaker. Heat the solution to about 70°C. Add ca 20cm3 hydrogen peroxide solution and reheat to 70°C. 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. 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.
Potassium sodium 2,3-dihydroxybutanedioate is an irritant. Hydrogen peroxide is corrosive. Hydrated cobalt(II) chloride is toxic. The dust may be irritating and, in larger doses, severely damaging to the respiratory tract. Skin contact, inhalation or ingestion should be avoided.
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.
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- C4H4O62- (aq) + 2H2O(l) E= +0.20V
H2O2(aq) + 2H+(aq) + 2e- 2H2O(l) E= +1.77V
The reaction is energetically favourable because E 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)
C4H4O62-(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 E 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.
1. A. Lainchbury et al, ILPAC, 2nd edn, vol 11. London: John Murray, 1996.