According to Le Châtelier's Principle, if the equilibrium of a system is disturbed by a stress, the system will shift to compensate.
When a chemical system is at equilibrium, there is no net change in the concentration of its reactants or products. If any parameter, such as concentration or temperature, is altered, the equilibrium will be disturbed.
The system readjusts by shifting the direction of the reaction until a new equilibrium is reached.
This video will demonstrate Le Châtelier's Principle by showing the influence of concentration and temperature on chemical reactions at equilibrium.
Reversible chemical reactions consist of two competing processes: the forward reaction, and the reverse reaction. When these two processes occur at the same rate, the system is at equilibrium. Le Châtelier's Principle states that, when a system at equilibrium is stressed, it will shift to counteract the disturbance.
For instance, if the concentration of a reactant species in an equilibrium solution is increased, the equilibrium will shift towards the products, increasing the rate of the forward reaction. Eventually, the system will reach a new equilibrium.
Temperature can also be thought of as a reaction component. In exothermic reactions, heat is released, making it a product. In endothermic reactions, heat is absorbed from the surroundings, making it a reactant. Thus, adding or removing heat will disturb the equilibrium, and the system will adjust.
This experiment will look at the ionic reaction of iron (III) with thiocyanate to form an iron (III) thiocyanate complex. The product is red, while the reactants are yellow or colorless, allowing for shifts in equilibrium to be observed visually.
The concentrations of these components will be altered by either directly adding ions to solution, or by selectively removing them through the formation of insoluble salts. The effect of a temperature change on this solution will also be observed.
Now that you understand Le Châtelier's Principle, you are ready to begin the procedure.
To begin the procedure, place one drop of 1 M iron nitrate solution into a test tube. Place one drop of 1 M potassium thiocyanate solution in a second test tube. Dilute each with 2 mL of water. These two tubes will serve as controls for the remainder of the experiment.
Next, in a new tube, add a drop of each solution. Add 16 mL of water, and mix thoroughly. Record any observations.
Divide this mixture into 2 mL portions in seven labeled test tubes. Set the initial tube aside as an iron thiocyanate control.
Next, add reactants to tubes 1 – 6 according to Table 2 below. Shake to mix every time a species is added, and record any observations.
Place test tube 7 into a hot water bath for 1 – 2 min. Compare the warm solution to the iron thiocyanate control, and record any observations.
In solutions 1 and 2, the red color intensified as the concentration of the reactants was increased. This indicates that the equilibrium shifted to the right, leading to the production of more iron (III) thiocyanate.
The solutions that received silver nitrate became colorless and formed a precipitate. The addition of thiocyanate ion caused the red color to reappear. The red color did not reappear when iron ion was added. From these observations, it can be concluded that thiocyanate ion was selectively removed from solution in the precipitate. As its concentration decreased, the equilibrium shifted to the left. Adding thiocyanate ion back into solution caused the equilibrium to shift back to the right.
The solutions that received potassium phosphate were observed to fade and become yellow. When the iron ion concentration was increased, the red color reappeared and the solution became cloudy. Increasing the thiocyanate ion concentration had no effect. Thus, it can be deduced that iron was selectively removed from solution to form an iron phosphate salt, causing the equilibrium to shift to the left. The iron phosphate salt eventually precipitated out of solution when more iron was added, and the equilibrium shifted back to the right.
The red color of Solution 7 faded to orange as temperature increased. This equilibrium shift to the left suggests that the reaction is exothermic, and that heat is generated when the iron thiocyanate product is formed.
The concept of equilibrium shifting has several applications in a wide range of scientific fields.
Le Châtelier's Principle explains why buffer solutions resist pH change. In this example, a sodium acetate buffer solution was used to maintain a nearly constant pH.
In aqueous solution, acid dissociation is a reversible reaction where the anions dissociate from the hydrogen ions. Buffer solutions are often an equilibrium mixture of dissociated hydrogen ions, a weak acid, and its anion — also known as its conjugate base.
If a strong acid is added, it will dissociate completely, increasing the concentration of the hydrogen ions in solution. The equilibrium of the weak acid reaction shifts to the left in response, reducing the concentration of hydrogen ions until it reaches a new equilibrium. Because of this, buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications.
Polymerization, the process of reacting molecules together to form polymer chains, is essential for bacterial cell division. In this example, Le Châtlelier's Principle was observed by performing FtsZ sedimentation assays under various conditions. Nine buffers were created, each with unique compositions and pH values. Polymerization was induced, then monitored by 90° angle light scattering. It was found that both the pH and the buffer composition affected polymerization, as each provided a stressor that shifted the reaction's equilibrium.
Finally, Le Châtlelier's Principle can be used in the production and recovery of materials in organic reactions. In this example, ammonium was recovered from nitrogen-rich streams.
The stream was passed through an electrochemical system, oxidizing the water and allowing for the separation of ammonium ions. These ions were then subjected to high pH, shifting their equilibrium, and driving the conversion of ammonium to volatile ammonia.
This captured ammonia was then passed through a stripping and absorption column to trap the ammonia in an acidic medium, shifting the equilibrium in the other direction.
You've just watched JoVE's introduction to the influence of temperature and concentration on reactions according to Le Châtelier's Principle. You should now understand the concept of equilibrium, how changes in concentration will cause shifts, and that heat can be considered a reaction component.
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