The reaction of zinc and sulfur
Updated October 2017. The quantities originally recomended for this demonstration mean it is prohibited under the 2014 explosives regulations. We have removed the kit list and instructions from this article.
There are several ways to demonstrate to students the differences between elements, mixtures and compounds. The difficulty in separating the elements in a compound is usually shown by using the physical and chemical properties of iron and sulfur. However, when teaching this topic I prefer to use a demonstration based on the reaction between zinc and sulfur. This reaction offers the same teaching opportunities as the Fe/S reaction but uses different techniques for identifying the nature of the reactants and products. Electrical conductivity can be used to show the differences between the elements, rather than magnetism, and a mixture of the reactants can be separated by reaction with water or dilute acid (if zinc dust is used). This provides a more detailed analysis of the chemistry involved in the reaction.
The products of this reaction bear little resemblance to the starting elements, pale blue zinc and bright yellow sulfur. Students can see that a new compound has been formed. The pale yellow residue is a mixture of zinc oxide and zinc sulfide. Reducing this residue (by electrical or chemical means) back to zinc demonstrates the chemical differences between mixtures and compounds. A drop of hydrochloric acid on the residue can be a fitting conclusion to the reaction, if done in a fume cupboard. The reaction produces the pungent, foul-smelling gas hydrogen sulfide which students often associate with stink bombs – once smelled never forgotten. Lead ethanoate paper turns black in the presence of the gas.
The reaction between zinc and sulfur
In this demonstration an intimate mixture of zinc and sulfur produces an unusual chemical reaction when heated. A brilliant flash of light, followed by hot sparks, a hissing sound and a mushroom-shaped cloud of white smoke are generated.
Zinc is an essential element – the mineral is required for skin and bone growth, and our bodies use zinc to process food and nutrients. Zinc ions are vital components in several different enzymes found in the body. A pale yellow, odourless, brittle solid, sulfur is essential to life too, occurring in the amino acids cysteine and methionine and therefore in many proteins. It is a minor constituent of fats, body fluids, and skeletal minerals.
The chemical reactions which are occurring in the reaction are:
Zn(s) + S(s) → ZnS(s) ΔH⦵f = –206.0 kJ mol-1
Zn(s) + ½O2(g) → ZnO(s) ΔH⦵f = –348.3 kJ mol-1
S(s) + O2(g) → SO2(g) ΔH⦵f = –297.0 kJ mol-1
The overall reaction needs heat to get started but the heat it produces is enough to sustain the reaction thereafter.
Zinc can be obtained by electrolysis of zinc sulfate or by smelting in a process similar to the production of iron from the blast furnace. First the zinc ore is roasted in air, converting it to zinc oxide:
2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g)
Coke and the roasted ore are fed into the top of the furnace with air blasted in at the bottom. The most important reaction taking place is:
ZnO(s) + CO(g) → Zn(g) + CO2(g)
Unlike the production of iron (mp = 1535°C; bp = 2750°C), zinc (mp = 420°C; bp = 907°C) is produced as a vapour. Cooling the zinc vapour to produce a liquid results in the re-oxidation of the metal. This problem was solved in the 1950s by Imperial Smelting of Bristol. The zinc vapour is sprayed with molten lead. This chills and dissolves the zinc so rapidly that re-oxidation is minimal. Molten lead and zinc are only partially miscible in each other and so, by cooling the solution, zinc separates as a liquid of nearly 99% purity. Vacuum distillation can further refine the liquid to 99.99% purity. This method has the advantage that the charge composition is not critical and mixed Zn/Pb sulfide ores (often found together) will produce both metals simultaneously, the lead being tapped from the bottom of the furnace.
When using this demonstration, it is worth mentioning that zinc is applied in thin layers to iron and steel products to stop them rusting. This process is called galvanising. More than half of the zinc consumed each year is used for galvanising. About 7.7 kg of zinc is used to protect the average car from rust.