Le Chatelier's principle

Introduction

Using a colour change is one of the best ways to demonstrate chemical equilibrium and Le Chatelier's principle. In fact students will already have met a simple example when they used an acid-base indicator during the titration of an acid with a base although they probably will not have made the connection. I usually demonstrate chemical equilibrium by using a beaker of copper(II) sulfate solution placed on an old overhead projector. I add concentrated ammonia to give the deep blue ammonia complex and then add concentrated ammonia solution. It produces lots of white fumes of hydrogen chloride and looks quite spectacular but eventually the deep blue colour turns to a light green as the [CuCl₄]^2- complex ion is formed. I then add a lot more water to dilute it and the colour turns back to the pale blue copper(II) aqueous ion, Cu(aq)²⁺.

You can also demonstrate with the iron(III) thiocyanate reaction but it is nice to let students do it themselves. You can find lots of versions on the internet - most use petri dishes, test-tubes or beakers (see image on left). I've made it into a microscale experiment to show what can be done with very small quantities (literally one or two drops) of chemicals. If you have special microscale plastic wells then use those (or improvise with empty pill blister packs) but a very simple solution is to use an acetate sheet. It works best if students can put the acetate sheet on an old overhead projector (or use more high tech to link to a beamer or smart board) but this is not at all necessary as a piece of white paper beneath the acetate sheet makes the colour changes easy to see.

Teacher’s notes

The reaction is actually:

[Fe(H₂O)₆]³⁺(aq) + SCN^-(aq) ⇌ [Fe(H₂O)₅SCN]²⁺(aq) + H₂O(l)

but since Standard Level students do not know about complex transition metal ions it is usually shortened to just

Fe³⁺(aq) + SCN^-(aq) ⇌ [FeSCN]²⁺(aq)

The first square will show the red colour due to the formation of the [FeSCN]²⁺(aq) complex. As the concentration of Fe³⁺(aq) is considerably increased the second square will turn a deeper red as the equilibrium shifts to the products side. Similarly increasing the concentration of the other reactant, the thiocyanate ion, will shift the equilibrium towards the products side and cause a deeper red in square 3. In square 4 the colour will lighten. This is because the iron(III) ions will react with hydroxide ions and be removed from the solution as a precipitate of iron(III) hydroxide causing a reduction in the Fe³⁺(aq) concentration and hence the position of equilibrium will shift back to the reactants side . Finally, in the fifth square, the colour will again turn a deeper red as the acid reacts with the hydroxide ions thus releasing more iron(III) ions to form the complex ion.

This practical does not cover any of the mandatory areas so is just to illustrate Le Chatelier's principle. If you have a spectrometer or colorimeter you can adapt this experiment to measure the absorbance and use it to determine the equilibrium constant for the reaction. This could be used as scaffolding to introduce the technique for students to use in their individual scientific investigation. You can find details of how to do this on the web, for example, Determining K_c for [FeSCN]²⁺.

Technician notes

To make a solution of 0.1 mol dm^-3 KSCN dissolve 0.972 g of potassium thiocyanate in distilled water and make the total volume up to 100 cm³. Using a pipette, take 10.0 cm³ of this solution and make up to 100 cm³ with distilled water to make 0.01 mol dm^-3 KSCN.
To make a solution of 0.1 mol dm^-3 iron(III) nitrate dissolve 4.041 g of hydrated iron(III) nitrate, Fe(NO₃)₃.9H₂O, in 25 cm³ of 2 mol dm^-3 nitric acid and make the total volume up to 100 cm³ with distilled water. Dilute ten times as above to make 0.01 mol dm^-3 Fe³⁺(aq).

Student worksheet

USING THE REACTION BETWEEN IRON(III) IONS AND THIOCYANATE IONS TO ILLUSTRATE LE CHATELIER’S PRINCIPLE

INTRODUCTION

Iron(III) ions in aqueous solution react reversibly with thiocyanate ions to form the blood red iron(III) thiocyanate ion:

Fe³⁺(aq) + SCN^-(aq) ⇌ [FeSCN]²⁺(aq)

The aim of this practical is to see how changing the concentrations of the reactants in the mixture will affect the position of equilibrium. The intensity of the red colour is a good indicator of the position of the equilibrium mixture – the greater the concentration of the iron(III) thiocyanate complex ion the more intense the colour.

This practical introduces you to the idea of microscale chemistry. By using just one or two drops of chemicals on an acetate sheet huge savings (both financial and environmental) are made. If you have access to an old-fashioned overhead projector then you can project the image of the reactions onto a screen to see the colours very clearly but this is not actually necessary.

ENVIRONMENTAL CARE:

Since only tiny volumes of solutions are used there is almost no impact upon the environment.

SAFETY:

Avoid getting the chemicals on your skin but there are no particular hazards associated with any of the chemicals as the solutions are so dilute. Wear safety glasses.

PROCEDURE:

Using a felt-tipped pen mark a piece acetate paper into five separate squares of approximately 4 cm by 4 cm. Put the acetate sheet onto a flat white surface (or onto an overhead projector). Place one drop of 0.01 mol dm^-3 potassium thiocyanate solution, KSCN(aq), into the centre of each of the squares then add one drop of 0.01 mol dm^-3 iron(III) nitrate solution, Fe(NO₃)₃(aq), to each drop of the thiocyanate solution. The first square acts as the reference. To the mixtures on the others squares add the following and mix using a small glass rod:
2. one drop of 0.1 mol dm^-3 iron(III) nitrate solution.
3. one drop 0.1 mol dm^-3 potassium thiocyanate solution.
4. one drop of 0.1 mol dm^-3 sodium hydroxide solution.
5. one drop of 0.1 mol dm^-3 sodium hydroxide solution followed by one drop of 0.1 mol dm^-3 hydrochloric acid solution.