Catalyst Demonstration
The concept of a catalyst is fundamental to chemical reactions but not applied very much in introductory classes. In fact, the term catalyst is probably more commonly used by the general public than the chemistry teacher. The problem is simple, a catalyst is relatively uncommon at the general chemistry level.

There are a few common applications, like the decomposition of hydrogen peroxide in the presence of manganese dioxide. This decomposition can also be related to the bubbling observed when hydrogen peroxide is used to clean a minor wound. But it is difficult to see how either of these examples illustrates the definition appearing in most textbooks. Most books give a definition similar to "a catalyst is a substance that greatly accelerates the rate of a chemical reaction without itself being permanently changed". The part about increasing the rate of the reaction is easy but the rest of the definition is elusive.

There is a great reaction available that demonstrates the full action of a catalyst using readily available chemicals. I can certainly recommend it even for general chemistry classes.


Materials Needed:

• 250 ml beaker
• Laboratory grade thermometer (110^oC)
• Laboratory burner
• Ring stand, iron ring, and wire gauze
• Glass stirring rod
• Graduated cylinder
• 0.2 g laboratory grade cobalt(II) chloride, CoCl₂ dissolved in 5 mL water
• 20 mL 6% hydrogen peroxide
• 5 g sodium potassium tartrate, NaKC₄H₄O₆ (also called Rochelle's Salt)
• Large tray to contain the setup

The chemical equation for the reaction is:

5 H₂O₂(aq) + NaKC₄H₄O₆(aq) → 4 CO₂(g) + NaOH(aq) +
KOH(aq) + 6 H₂O(aq)

Dissolve the 5 g of NaKC₄H₄O₆ in 60 mL of deionized (or distilled) water in the 250 mL beaker then pour in the 20 mL of 6% H₂O₂. Place the beaker on the ring stand and heat to 70^oC. Observe the very slow rate of the reaction at this temperature then remove the burner and shut off the flame.

Now add the pink catalyst solution (0.2 g CoCl₂ dissolved in 20 mL of water). Stir once or twice and stand back. Observe the pink color of the Co⁺² disappear and a green solution form with CO₂ gas being evolved very rapidly. Have the setup sitting in the tray since the effervescence may cause some of the solution to overflow the beaker. The evolution of CO₂ gas ceases after about 30 seconds and the original pink color of the Co⁺² ion reappears. The temporary appearance of the green color allows you to discuss the possibility of an intermediate species or activated complex.

This demonstration is easy to run and very impressive to see. It is certain to open the door to a lively discussion and change the topic of a catalyst from a dull definition to an exciting concept.

More Teaching Tips from Quantum coming soon!