עקרון לה שאטלייה
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Overview
מקור: המעבדה של ד”ר לין אוקונל — מכללת בוסטון
כאשר התנאים של מערכת בשיווי משקל משתנים, המערכת מגיבה באופן כזה כדי לשמור על שיווי המשקל. בשנת 1888, אנרי-לואיס לה שאטלייה תיאר תופעה זו בעיקרון הקובע, “כאשר שינוי בטמפרטורה, בלחץ או בריכוז מפריע למערכת בשיווי משקל כימי, השינוי יתובל על ידי שינוי בהרכב שיווי המשקל”.
ניסוי זה מדגים את העיקרון של Le Châtelier בעבודה בתגובה הפיכה בין ברזל (III) יון ויון thiocyanate, אשר מייצר ברזל (III) יון thiocyante:
Fe3+(aq) + SCN– (aq) 
FeSCN2+ (aq)
הריכוז של אחד היונים משתנה או על ידי הוספת כמות של יון אחד לפתרון או על ידי הסרה סלקטיבית של יון מהפתרון באמצעות היווצרות מלח בלתי מסיס. תצפיות של שינויי צבע מצביעות אם שיווי המשקל עבר לטובת היווצרות של המוצרים או המגיבים. בנוסף, ניתן לראות את ההשפעה של שינוי טמפרטורה על הפתרון בשיווי משקל, מה שמוביל ליכולת להסיק אם התגובה היא אקסותרמית או אנדותרמית.
Principles
Procedure
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– 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 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 (l) H2CO3 (aq) 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) 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.
Transcript
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!