Witness a spectacular chemical reaction and take some careful measurements to work out the empirical formula of a compound.
Every chemical compound has a chemical formula. In fact, there are several different types of chemical formula for any one compound (figure 1). Perhaps the most familiar type is the molecular formula – such as H2O for water and CO2 for carbon dioxide – which tells us the number of different atoms in each molecule. Structural formulae go a step further by showing how the atoms are linked together within the molecule, which is especially important for organic compounds.
The simplest type of formula – called the empirical formula – shows just the ratio of different atoms. For example, while the molecular formula for glucose is C6H12O6, its empirical formula is CH2O – showing that there are twice as many hydrogen atoms as carbon or oxygen atoms, but not the actual numbers of atoms in a single molecule or how they are arranged. These simple, ratio-based formulae were developed by early chemists in the 18th century. They are known as ‘empirical’ formulae because the ratio between the numbers of atoms in a compound can be found by traditional methods of chemical analysis by experiment.
Today, working out an empirical formula experimentally is an important feature of chemistry courses all over the world. It is also the first step in working out the chemical formula of an unidentified compound, making it a useful tool in chemical analysis. The classic school demonstration involves heating magnesium in a crucible to make magnesium oxide – a dull white powder. In this article, a much more exciting alternative is described: a dramatic reaction between tin and iodine, producing a bold purple vapour and bright orange crystals as the reaction progresses.
The aim of the experiment is to calculate the ratio between tin and iodine atoms in tin iodide. This is done by synthesising this compound and accurately measuring the mass of the reagents at the start of the experiment and the leftover tin at the end. The experiment involves a range of techniques, including setting up and using a reflux condenser and using organic solvents for extraction. As well as covering the practical exercise of deriving the empirical formula, the experiment links to more theoretical chemistry topics, such as the benefits of using a reagent in excess, the synthesis of compounds from their elements, and how bonding can be linked to solubility. It’s also a very clear application of the law of conservation of mass, which is a fundamental principle throughout chemistry (and science generally).
Depending on the number of fume cupboards available in your classroom, students can carry out the experiment themselves, but it is also suitable as a teacher demonstration. The experiment takes about two hours and works best in a double lesson, but it is also possible to split it between two single lessons. It is most suitable for students aged 16–18, but it could also be used as an extension activity for those aged 14–16.
The teacher (or each group of students) will need the following:
Students should wear a lab coat, gloves and safety goggles. Solid iodine is corrosive and can stain the skin, which is why gloves should be worn. The experiment should be carried out in a fume cupboard. As iodine vapour is toxic, ensure that the purple vapour does not rise more than one-third of the way up the reflux condenser when heating. Cyclohexane and propanone are highly flammable, so a heating mantle is required, and care should also be taken to keep both these solvents away from naked flames. Propanone should be disposed of in a solvent residues bottle. In addition, teachers should follow their local health and safety rules.
This experiment involves reacting two substances – tin and iodine – in their elemental form to produce the compound tin iodide. Tin has more than one possible oxidation state, so the reaction could produce either tin(II) iodide (SnI2) or tin(IV) iodide (SnI4). Using the experimental data, we can derive the empirical formula for the product, which will tell us the ratio between tin and iodine. From this, we can work out the identity of the tin compound produced.
Tin used in the reaction = initial mass (step 2) minus the leftover mass (step 13)
This activity can produce good results if carried out carefully, with values that should quite closely round to 1:4 (as the ratio of tin to iodine). This leads to SnI4 as the empirical formula for tin iodide.
In general, values ranging between 1:3.2 and 1:3.8 are often obtained. As the common oxidation states for tin are +2 and +4, an experimental outcome giving a ratio close to 1:3 would not be in agreement. However, such values can open up a discussion about sources of experimental error and the importance of precision.
After the experiment, ask all students to think about possible sources of error. What effect might each of the following have on the final results?
Table 1 summarises the effect of each of these sources of error in the experiment and on the final result – that is, how each changes the value of x in the empirical formula SnIx.
|Error||Effect on: mass of tin left over||Effect on: mass of tin reacted (initial mass minus leftover mass)||Effect on: value of x (in SnIx)|
|Incomplete reaction of iodine||Increases, as not all the tin reacts||Decreases||Increases|
|Loss of leftover tin while washing with propanone||Decreases||Increases||Decreases|
|Incomplete drying of leftover tin||Increases, by adding extra mass from the solvent||Decreases||Increases|
|Poor washing of leftover tin||Increases, as some tin iodide is included in the leftover tin mass||Decreases||Increases|
|Loss of iodine vapour from condenser||Increases, as not all the tin was able to react||Decreases||Increases|
This experiment also offers an opportunity to discuss how bonding is linked to solubility. Iodine and tin iodide both dissolve in non-polar solvents (cyclohexane and propanone) but not in water, whereas tin is a metal and is insoluble in cyclohexane, propanone and water. Using this information, can your students draw conclusions about the likely bonding in these substances?
The author would like to thank Alan Carter, who was the Head of Chemistry at Wellington College (Berkshire, UK) until 2004, and who created the initial resource that inspired this article.