Le Châtelier's Principle
Source: Laboratory of Dr. Lynne O'Connell — Boston College
When the conditions of a system at equilibrium are altered, the system responds in such a w…
1. Preparation of the Iron(III) Thiocyanate Equilibrium Solutions
- Place 1 drop of 1 M Fe(NO3)3 solution in atest tube and dilute with 2 mL of water. Place 1 drop of 1 M KSCN in another test tube and dilute with 2 mL of water. These two test tubes serve as controls to compare against the other test tubes.
- Place 1 drop of 1 M Fe(NO3)3 solution in atest tube.
- Add 1 drop of 1 M KSCN to the test tube.
- Add 16 mL of water to the test tube and thoroughly mix the contents.
- Record any observations.
- Divide the mixture into 2 mL portions in 8test tubes. One of the test tubes remains untouched and serves as a FeSCN2+ control. Number the other test tubes 1–7.
2. Addition of Iron(III) and Thiocyanate Ions to the Equilibrium Solution
- To test tube 1, add 1 drop of 1 M Fe(NO3)3 solution.
- Shake to mix and record any observations.
- To test tube 2, add 1 drop of 1 KSCN solution.
- Shake to mix and record any observations.
3. Addition of Silver Nitrate to the Equilibrium Solution
- To test tube 3, add 3 drops of 0.1 M AgNO3 solution.
- Shake to mix and record any observations.
- Add 3 drops of 1 M Fe(NO3)3 to the test tube.
- Shake to mix and record any observations.
- To test tube 4, add 3 drops of 0.1 M AgNO3 solution.
- Shake to mix and record any observations.
- Add 3 drops of 1 M KSCN to the test tube.
- Shake to mix and record any observations.
4. Addition of Potassium Phosphate to the Equilibrium Solution
- To test tube 5, add 3 drops of 0.5 M K3PO4 solution.
- Shake to mix and record any observations.
- Add 3 drops of 1 M Fe(NO3)3 to the test tube.
- Shake to mix and record any observations.
- To test tube 6, add 3 drops of 0.5 M K3PO4 solution.
- Shake to mix and record any observations.
- Add 3 drops of 1 M KSCN to the test tube.
- Shake to mix and record any observations.
5. Changing the Temperature of the Equilibrium Solution
- Place test tube 7 in a 70–80 °C water bath for 1–2 min.
- Compare the warm solution to the solution in the unheated test tube (the FeSCN2+ control), and record any observations.
- Collect the contents of test tubes 3 and 4 in the laboratory waste jar labeled "Silver." Pour the contents of all the other test tubes down the drain.
According to Le Ch?telier'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 Ch?telier'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 Ch?telier'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 Ch?telier'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 Ch?telier'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 Ch?tlelier'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 Ch?tlelier'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 Ch?telier'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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Q1: What does Le Châtelier's Principle state about systems at equilibrium?
Le Châtelier's Principle states that when a system at equilibrium is stressed by changes in concentration, temperature, or pressure, the system shifts to counteract the disturbance. This shift continues until a new equilibrium is reached. The principle explains how chemical systems naturally resist changes and maintain balance when external conditions are altered.
Q2: How does increasing reactant concentration affect chemical equilibrium?
When the concentration of a reactant is increased, the equilibrium shifts toward the products to consume the excess reactant. This increases the rate of the forward reaction, producing more product until a new equilibrium is established. Understanding solutions and concentrations molarity molality helps predict these shifts accurately.
Q3: Why is heat considered a reaction component in Le Châtelier's Principle?
In exothermic reactions, heat is released and acts as a product. In endothermic reactions, heat is absorbed and acts as a reactant. Since heat functions as either a product or reactant, adding or removing heat disturbs equilibrium, causing the system to shift and adjust to a new equilibrium state.
Q4: What happens to equilibrium when a reactant is selectively removed from solution?
When a reactant is selectively removed through formation of an insoluble salt, its concentration decreases, and the equilibrium shifts left toward the reactants. This shift attempts to replace the removed reactant. Understanding the solubility rules of ionic compounds helps predict which ions precipitate and shift equilibrium.
Q5: How do buffer solutions use Le Châtelier's Principle to resist pH change?
Buffer solutions contain a weak acid and its conjugate base in equilibrium. When strong acid is added, it increases hydrogen ion concentration, shifting the weak acid equilibrium left and reducing hydrogen ions until a new equilibrium forms. This buffering action maintains nearly constant pH despite acid addition.
Q6: What does the color change in the iron thiocyanate reaction indicate about equilibrium shifts?
The iron thiocyanate complex is red, while reactants are colorless or yellow. When equilibrium shifts right toward products, the solution becomes more intensely red. When equilibrium shifts left toward reactants, the red color fades. These color changes provide visual evidence of how the system responds to concentration and temperature changes.
Q7: How can Le Châtelier's Principle be applied to recover materials in industrial processes?
Le Châtelier's Principle guides material recovery by manipulating equilibrium through pH, temperature, or concentration changes. For example, ammonium recovery involves oxidizing water to separate ammonium ions, then shifting equilibrium at high pH to convert ammonium to volatile ammonia. This captured ammonia is then trapped in acidic medium, shifting equilibrium in the opposite direction.