Le Châtelier's Principle

General Chemistry

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Overview

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 way as to maintain the equilibrium. In 1888, Henri-Lewis Le Châtelier described this phenomenon in a principle that states, "When a change in temperature, pressure, or concentration disturbs a system in chemical equilibrium, the change will be counteracted by an alteration in the equilibrium composition."

This experiment demonstrates Le Châtelier's principle at work in a reversible reaction between iron(III) ion and thiocyanate ion, which produces iron(III) thiocyante ion:

Fe3+(aq) + SCN- (aq) Reversibly Equals FeSCN2+ (aq)

The concentration of one of the ions is altered either by directly adding a quantity of one ion to the solution or by selectively removing an ion from the solution through formation of an insoluble salt. Observations of color changes indicate whether the equilibrium has shifted to favor formation of the products or the reactants. In addition, the effect of a temperature change on the solution at equilibrium can be observed, which leads to the ability to conclude whether the reaction is exothermic or endothermic.

Cite this Video

JoVE Science Education Database. General Chemistry. Le Châtelier's Principle. JoVE, Cambridge, MA, (2017).

Principles

To fully understand Le Châtelier's Principle, a reversible reaction of the sort expressed by the following chemical equation is considered:

aA + bB Reversibly Equals cC + dD

This reaction actually consists of two competing processes: the forward reaction, in which the products C and D are formed from the reactants, and the reverse reaction, in which the reactants A and B are formed from the products. When the rates of these two processes equal each other, there is no net change in the concentration of either the products or the reactants, and the reaction is said to be at equilibrium. The ratio of the equilibrium concentrations of the products to the equilibrium concentrations of the reactants is a constant, as shown by the following equation:

Generic Kc Equation

where Kc is the equilibrium constant. The brackets signify the concentrations of the various species, and the lowercase letters represent the number of moles of each substance involved in the balanced equation. In the case of the reaction between iron(III) and thiocyanate ions shown previously, the equilibrium constant is:

Specific Kc Equation

When the concentration of either a reactant or a product in an equilibrium solution is altered, the concentrations of the other species must change in order to maintain the constant ratio of products to reactants. These changes are referred to as "shifts" in the equilibrium. The equilibrium can either shift to the left, meaning it proceeds in the reverse direction and the concentrations of the reactants increase, or shift to the right, meaning it proceeds in the forward direction and the concentrations of the products increase. In the reaction between iron(III) and thiocyanate ions, a shift to the left would mean formation of more iron(III) and thiocyanate ions, while a shift to the right would mean formation of more iron(III) thiocyanate ions.

The equilibrium constant is dependent upon the temperature; thus, a change in the temperature of an equilibrium solution can also result in a shift to the right or left, depending on whether the reaction is exothermic or endothermic. For an exothermic reaction, the heat generated by the reaction can be represented as residing on the product side of the equation, since heat is produced along with the products:

aA + bB Reversibly Equals cC + dD + heat

If heat is added to the system by increasing the temperature, the equilibrium shifts to the left, and the concentrations of the reactants increase. For an endothermic reaction, the addition of heat would result in a shift to the right.

aA + bB + heat Reversibly Equals cC + dD

In this case, the concentrations of the reactants would increase with an increase in temperature.

Procedure

1. Preparation of the Iron(III) Thiocyanate Equilibrium Solutions

  1. Place 1 drop of 1 M Fe(NO3)3 solution in a test 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.
  2. Place 1 drop of 1 M Fe(NO3)3 solution in a test tube.
  3. Add 1 drop of 1 M KSCN to the test tube.
  4. Add 16 mL of water to the test tube and thoroughly mix the contents.
  5. Record any observations.
  6. Divide the mixture into 2 mL portions in 8 test 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

  1. To test tube 1, add 1 drop of 1 M Fe(NO3)3 solution.
  2. Shake to mix and record any observations.
  3. To test tube 2, add 1 drop of 1 KSCN solution.
  4. Shake to mix and record any observations.

3. Addition of Silver Nitrate to the Equilibrium Solution

  1. To test tube 3, add 3 drops of 0.1 M AgNO3 solution.
  2. Shake to mix and record any observations.
  3. Add 3 drops of 1 M Fe(NO3)3 to the test tube.
  4. Shake to mix and record any observations.
  5. To test tube 4, add 3 drops of 0.1 M AgNO3 solution.
  6. Shake to mix and record any observations.
  7. Add 3 drops of 1 M KSCN to the test tube.
  8. Shake to mix and record any observations.

4. Addition of Potassium Phosphate to the Equilibrium Solution

  1. To test tube 5, add 3 drops of 0.5 M K3PO4 solution.
  2. Shake to mix and record any observations.
  3. Add 3 drops of 1 M Fe(NO3)3 to the test tube.
  4. Shake to mix and record any observations.
  5. To test tube 6, add 3 drops of 0.5 M K3PO4 solution.
  6. Shake to mix and record any observations.
  7. Add 3 drops of 1 M KSCN to the test tube.
  8. Shake to mix and record any observations.

5. Changing the Temperature of the Equilibrium Solution

  1. Place test tube 7 in a 70–80 °C water bath for 1–2 min.
  2. Compare the warm solution to the solution in the unheated test tube (the FeSCN2+ control), and record any observations.
  3. 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 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.

Thanks for watching!

Results

Observations of the initial solutions and the mixture of the two solutions can be seen in Table 1.

Observations of the equilibrium mixtures upon addition of various reagents can be seen in Table 2.

Observation when the temperature is changed: In test tube 7, the solution turns more orange in color (less red, more yellow) when heated.

In test tubes 1 and 2, when iron(III) nitrate, which contains a reactant, was added to the equilibrium solution, the red color of the solution intensified. This observation indicates that the equilibrium shifted to the right as concentration of the product, iron(III) thiocyanate ion, increased. Similarly, when potassium thiocyanate, which contains the other reactant, was added to the equilibrium solution, the red color of the solution intensified. This observation also indicates that the equilibrium shifted to the right as concentration of the product increased.

In test tubes 3 and 4, when silver nitrate (AgNO3) was added to the equilibrium solution, the red color of the product faded and the solution became colorless. This observation indicates that the equilibrium shifted to the left as the concentration of reactants increased. In addition, a precipitate was observed. The red color reappeared upon addition of thiocyanate ion (SCN-). This observation indicates that the equilibrium shifted to the right as the concentration of the product increased. The red color did not reappear when iron(III) ion (Fe3+) was added.

From these observations, it can be concluded that silver thiocyanate (AgSCN) was the precipitate that formed when silver nitrate was added to the equilibrium solution. The formation of this solid is responsible for the cloudiness observed in both test tubes. When the thiocyanate ion was removed from the solution by precipitation, the equilibrium shifted to the left, because the concentration of one of the reactants had been reduced. When more thiocyanate ion was then added, the equilibrium shifted back to the right to re-establish the equilibrium ratio of concentrations by re-forming iron(III) thiocyanate. The addition of more iron(III) ion did not shift the equilibrium back to the right, because the thiocyanate ion had been removed from the solution as silver thiocyanate precipitate and was no longer available to react with iron(III) to form the iron(III) thiocyanate ion.

In test tubes 5 and 6, when potassium phosphate ion (K3PO4) was added to the equilibrium solution, the red color of the products faded and the solution became yellow. This observation indicates that the equilibrium shifted to the left as the concentration of reactants increased. The red color reappeared upon addition of iron(III) ion (Fe3+). This observation indicates that the equilibrium shifted to the right as the concentration of the product increased. In addition, a precipitate was observed. The red color did not reappear when the thiocyanate ion (SCN-) was added.

From these observations, it can be concluded that iron(III) phosphate (FePO4) salt was formed when potassium phosphate was added to the equilibrium solution. When the iron(III) ion was removed from the solution by formation of this salt, the equilibrium shifted to the left, because the concentration of one of the reactants had been reduced. When more iron(III) ion was then added, the equilibrium shifted back to the right to re-establish the equilibrium ratio of concentrations by re-forming iron(III) thiocyante. Although no cloudiness was detected by eyesight when the phosphate ion was initially added, a cloudiness did appear when the iron(III) ion was subsequently added, which is the solid iron(III) phosphate salt. The addition of more thiocyanate ion did not shift the equilibrium back to the right, because the iron(III) ion had been removed from the solution as iron(III) phosphate salt and was no longer available to react with the thiocyanate ion to form the iron(III) thiocyanate ion.

In test tube 7, as the temperature increased, the red color of the products faded, indicating an equilibrium shift to the left as more reactants were formed. This observation leads to the conclusion that the reaction is exothermic. For an exothermic reaction, the heat generated by the reaction resides on the product side of the equation:

Fe3+ + SCN- Reversibly Equals FeSCN2+ + heat

When heat was added to the system (by increasing the temperature), the equilibrium shifted to the left.

Solution Observation
Fe(NO3)3 Yellow, clear
KSCN Colorless, clear
Fe(SCN)2+ Orange-red, clear

Table 1. Observations of the initial solutions and the mixture of the two solutions.

Test Tube # First Reagent Observation of Equilibrium Solution Second Reagent Observation of Equilibrium Solution
1 Fe(NO3)3

Red, clear
2 KSCN

Red, clear
3 AgNO3 (colorless, clear) Colorless (white), cloudy Fe(NO3)3 Yellow, still cloudy
4 AgNO3 Colorless (white), cloudy KSCN Orange-red, still cloudy
5 K3PO4 (colorless, clear) Yellow, clear Fe(NO3)3 Orange-red, cloudy
6 K3PO4 Yellow, clear KSCN Yellow, still clear

Table 2. Observations of the equilibrium mixtures upon addition of various reagents.

Applications and Summary

Le Châtelier's principle is at work in human bodies. Oxygen is transported from the lungs to muscle and other tissues by a protein called hemoglobin (Hb) that is found in the blood. The oxygen molecule binds to this protein in a reversible reaction that can be described by an equilibrium equation:

Hb + 4 O2 Reversibly Equals Hb(O2)4

In the lungs, the partial pressure of oxygen gas is high (on the order of 100 torr). The equilibrium shifts to the right in this environment, and the oxygen molecules bind to hemoglobin molecules until the protein is saturated with oxygen. When this saturated hemoglobin reaches the cells of muscle tissue, where the pressure of oxygen is much lower, the equilibrium shifts to the left, and the oxygen is released. If the muscle is at rest, the oxygen pressure is about 30 torr, and approximately 40% of the oxygen is released. When the muscle is active, the oxygen pressure ranges from 3 to 18 torr, and about 85% of the oxygen is released to satisfy the increased metabolic demand.

Another physiological example of an equilibrium system involves the regulation of blood pH. Carbon dioxide in the blood reacts reversibly with water to produce carbonic acid, which dissociates to produce hydronium and bicarbonate ions:

CO2 (aq) + H2O (lReversibly Equals H2CO3 (aq) Reversibly Equals H3O+ (aq) + HCO3-(aq)

During strenuous exercise, the amount of carbon dioxide produced by the cells increases as a result of high metabolic activity. The increased concentration of carbon dioxide in the blood causes a shift to the right in this equilibrium to produce more carbonic acid. When this happens, the pH level of the blood decreases as hydronium ion concentration increases. One of the body's responses to this imbalance in blood pH is to increase the rate of breathing so more carbon dioxide gas is exhaled from the lungs, thus shifting the equilibrium back to the left and raising the pH back to normal levels.

Le Châtelier's principle must also be taken into account in many industrial processes. Ammonia is an important chemical used in fertilizers, cleaning agents, and as a building block in synthetic organic reactions. The industrial production of ammonia is accomplished using the Haber process, which relies on the reversible reaction between hydrogen and nitrogen:

3 H2 (g) + N2 (g) Reversibly Equals 2 NH3 (g)

In order to optimize the production of ammonia, the reaction is run at high pressure, usually around 200 atm. There are 4 moles of gas on the left-hand side of the equation and 2 moles of gas on the right-hand side. Le Châtelier's principle dictates that an increase of pressure on the system shifts the equilibrium to the right, because the volume of 2 moles of gas is smaller than the volume of 4 moles of gas. Since volume and pressure are directly proportional, a shift to reduce volume also reduces pressure, and the system returns to equilibrium. In addition, the process involves liquefying the ammonia gas in a condenser, so it is removed from the reaction chamber. This decrease in ammonia also shifts the equilibrium to the right, maximizing the amount of ammonia produced.

1. Preparation of the Iron(III) Thiocyanate Equilibrium Solutions

  1. Place 1 drop of 1 M Fe(NO3)3 solution in a test 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.
  2. Place 1 drop of 1 M Fe(NO3)3 solution in a test tube.
  3. Add 1 drop of 1 M KSCN to the test tube.
  4. Add 16 mL of water to the test tube and thoroughly mix the contents.
  5. Record any observations.
  6. Divide the mixture into 2 mL portions in 8 test 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

  1. To test tube 1, add 1 drop of 1 M Fe(NO3)3 solution.
  2. Shake to mix and record any observations.
  3. To test tube 2, add 1 drop of 1 KSCN solution.
  4. Shake to mix and record any observations.

3. Addition of Silver Nitrate to the Equilibrium Solution

  1. To test tube 3, add 3 drops of 0.1 M AgNO3 solution.
  2. Shake to mix and record any observations.
  3. Add 3 drops of 1 M Fe(NO3)3 to the test tube.
  4. Shake to mix and record any observations.
  5. To test tube 4, add 3 drops of 0.1 M AgNO3 solution.
  6. Shake to mix and record any observations.
  7. Add 3 drops of 1 M KSCN to the test tube.
  8. Shake to mix and record any observations.

4. Addition of Potassium Phosphate to the Equilibrium Solution

  1. To test tube 5, add 3 drops of 0.5 M K3PO4 solution.
  2. Shake to mix and record any observations.
  3. Add 3 drops of 1 M Fe(NO3)3 to the test tube.
  4. Shake to mix and record any observations.
  5. To test tube 6, add 3 drops of 0.5 M K3PO4 solution.
  6. Shake to mix and record any observations.
  7. Add 3 drops of 1 M KSCN to the test tube.
  8. Shake to mix and record any observations.

5. Changing the Temperature of the Equilibrium Solution

  1. Place test tube 7 in a 70–80 °C water bath for 1–2 min.
  2. Compare the warm solution to the solution in the unheated test tube (the FeSCN2+ control), and record any observations.
  3. 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 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.

Thanks for watching!

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