SCH 4U1 | Equilibrium and Le Châtelier’s Principle | |

Introduction:
Chemical equilibrium is the state of a reaction in which all reactants and products have reached constant concentrations in a closed system (DiGiuseppe, Haberer, Salciccioli, Sanader, & Vavitsas, 2012, p. 420). Chemical reactions will occur until the reaction reaches a point where the concentrations of the products and reactants become constant. Le Châtelier's principle states that chemical systems at equilibrium shift in the direction that opposes the change when a change occurs that disturbs the equilibrium (DiGiuseppe, Haberer, Salciccioli, Sanader, & Vavitsas, 2012, p. 439). In the lab activity performed, the effects of changing temperature and volume of a system were observed and recorded for the endothermic reaction of dinitrogen tetroxide with 57.2 kJ of energy:
N2O4(g) + 57.2 kJ ↔2 NO2(g)
Colourless Reddish-Brown
As well, the effects of modifying the concentrations of the reactants and products of a system were observed and recorded for the reaction of iron (II) thiocyanide:
Fe3+(aq)+SCN1-(aq)↔Fe(SCN)2+(aq)
In each case, as our hypotheses predicted, the reactions shifted in the direction that restored equilibrium to the system.

Discussion:
Concentration: The effects that would be observed on a system at equilibrium all depend on whether products or reactants are added or removed. If products or reactants are added, the equilibrium shifts in the direction that would decrease the concentration of that component. However, if products or reactants are removed, the equilibrium shifts in the direction that would increase the concentration of that component. These observations are supported by the experiments performed in the lab where the reactants were added to iron (II) thiocyanide and the reaction shifts in the corresponding direction in order to restore equilibrium. This shift can be measured through the change in colour of the aqueous solution. For example, when iron (III) chloride was added, the reaction shifted to the right in order to decrease the concentration of colourless iron (III) chloride and as a result more Fe(SCN)2+(aq)was produced. This equilibrium shift was observed through the darkening of the solution because Fe(SCN)2+(aq)is dark red. The results from each experiment were consistent with the hypotheses. Reactions that shifted to the left were more easily noticed because they were distinctly a lighter colour in contrast to reactions that shifted to the right which were harder to see.
Temperature: The effect that temperature has on an equilibrium system depends on whether the reaction is exothermic or endothermic. The reaction that was studied was an endothermic reaction, therefore thermal energy was required in order for this reaction to take place. When the system was placed on ice, the system shifted towards the left, which is shown by the fact that the system turned a light reddish-brown colour. Since there was a decrease in thermal energy in the system, the system responded by trying to increase the concentration of the colourless reactant, causing the system to turn a lighter colour. On the other hand, when the system was placed in a hot water bath, it turned a dark reddish-brown colour. The increase in thermal energy in the system resulted in the system’s attempt to oppose the change by increasing the amount of product. This caused the system to turn darker, because the concentration of the product, which is reddish-brown, was increased. The results of this experiment were also consistent with the hypotheses pertaining to the result of modifying the system’s temperature.
Volume: A decrease in volume in the system results in the pressure in the system increasing; as a result, more molecules of reactant are made (thus shifting the reaction to the left). When the container size is increased, the pressure in the system is decreased; as a result, more molecules of product are made (thus shifting the reaction to the right). These conclusions were supported by the results of the experiment, since when the volume of a container that contained both N2O4(g) and NO2(g) decreased, the solution turned darker and then slightly lighter, proving that the reaction shifted to the left. As well, as the volume increased, the solution turned lighter and then slightly darker, showing that the reaction shifted to the right.

This experiment was successful in illustrating Le Châtelier’s principle. The results were generally consistent with what Le Châtelier’s principle predicted. The first part of this lab activity was more effective than the second portion at demonstrating this principle, because the results were much easier to see and understand. The visual change in the system when the temperature or pressure was altered was much more evident for this experiment than for the second part of the lab, when certain substances were added into the system.