Chemical Equilibrium: Le Châtelier’s Principle

Chemical Equilibrium: Le Châtelier’s Principle

  1. Introduction

Equilibrium is a state where the forward rate of the reaction is equal to the reverse rate. At this state, chemical reaction is not observed making it look as if the reaction has stopped.

However, the equilibrium can be easily changed when there is any stress in the system. The equilibria’s response to these stresses is explained by Le Châtelier’s principle: if equilibria are disturbed by an outside factor, the system will react in a direction to adjust to the new environment, reestablishing the equilibrium. However, the principle does not only state that the equilibrium will shift but it also predicts how it will change. Knowing the factors that disturb the equilibrium helps figuring out whether the equilibrium will move. The factors that are specified in Le Châtelier’s principle are pressure, concentration of the reactant or product, and temperature.

The equilibrium constant (Kc) and the reaction quotient (Qc) are also important in finding which direction the equilibrium will shift to when the system is upset. The equilibrium constant is the concentration of product divided by the concentration of the reactant raised to their orders accordingly, and the constant dependant on temperature when the system is in equilibrium (Kc=K([A]a[B]b)/([C]c[D]d) when cC+dD←→aA+bB). The reaction quotient has the equal way of calculating as the equilibrium constant. However, the reaction quotient can be calculated at any point on the progression of the reaction(Qc= K([A]a[B]b)/([C]c[D]d) when cC+dDaA+bB).

The significance of the equilibrium constant and the reaction quotient is that they decide which direction the equilibrium will shift to according to their comparative values. The reaction quotient being bigger than the equilibrium constant indicates that the product’s concentration is bigger than the ones of the reactant’s shifting the equilibrium towards the left. In exact opposite, the reaction quotient being smaller than the equilibrium constant indicates that the reactant’s concentration is bigger than the ones of the product’s shifting the equilibrium towards the right.

Aim of the experiment is to define Le Châtelier’s principle according to the factors that upset the equilibrium: concentration of reactants and product, and Temperature; to observe the equilibrium using sparingly soluble salts; and to explain the direction of the shifts and their reasons.

2. Materials

    1. Apparatus
  • Chemicals in a Medicine Dropper
  • Test Tubes
  • 50mL Graduated Cylinder
  • 10mL Beaker
  • Clamp Stand, Clamp
  • Heating Plate
  • 100mL Beaker
    1. Chemical
  • 0.1M CuSO4
  • 0.1M NiCl2
  • 1.0M CoCl2
  • 0.1M KI
  • 0.1M Na2CO3
  • 0.1M AgNO3
  • 6M HNO3
  • 0.1M HCl
  • 1M HCl
  • 12M HCl
  • 15M NH3

3. Procedures

    1. Changes in Reactant or Product Concentration
      1. Copper and Nickel Ion
        1. Place 1mL of 0.1M CuSO4 in a 10mL beaker and record the color
        2. Add 15M NH3 dropwise until color change occur and the solution is clear
      • Solids may occur (Cu(OH)2)
        1. Record observations
        2. Mix the solution by tickling it while adding NH3
        3. Add 1M HCl dropwise and mix the solution until the solution is clear
        4. Record observation
        5. Repeat the procedure using 0.1M NiCl2 in place of CuSO4 and record observation
      1. Cobalt Ions
        1. Place about 0.5mL (10 drops) of 1M CoCl2 in a clean 10mL beaker
        2. Record color
        3. Drop 12M HCl to the test tube until a distinct color change occurs
        4. Record observation
        5. Slowly add water to the test tube while mixing
        6. Record observation
    1. Equilibria Involving Sparingly Soluble Salt
        1. To 0.5mL (10 drops) of 1.0M Na2CO3 in a 10mL beaker
        2. Add 1 drop of 0.1M AgNO3
        3. Record observation
        4. Add 6M HNO3 dropwise until you observe a change in appearance
        5. Record observation
        6. To the above solution, add 0.1M HCl dropwise
        7. Record observation
        8. While mixing, add 15M NH3 dropwise until chemical reaction occurs
        9. Add 6M HNO3 dropwise until chemical change is no more change.
        10. Record observation
        11. Continue to add 0.1M KI dropwise until you see evident chemical change
        12. Observe reaction
        13. Dispose the waste accordingly
    2. Effect of Temperature on Equilibria
        1. Heat 75mL of water to boil in a 100mL beaker on a heating plate
        2. Place about 1mL of 1.0M CoCl2 in a small test tube in the boiling content using a clamp stand
        3. Compare the color of the cool cobalt solution to that of the hot solution
        4. Record observation

4. Data

    1. Changes in Reactant or Product Concentration
      1. Copper and Nickel Ions
        1. CuSo4 (aq) – Light Blue
        2. [Cu(NH3)4]2+ (aq) – Dark Blue
        3. After HCl reaction – Extremely Clear, almost transparent but slightly light bluish
        4. NiCl2 (aq) – Light Green, turquoise
        5. [Ni(NH3)6]2+ (aq) – Pale violet
        6. After HCl reaction – Light Blue
      2. Cobalt Ions
        1. CoCl2 (aq) – Rose pink, slightly red
        2. After HCl reaction – Dark Blue, Gas evolve, Temperature increase
        3. After H2O reaction – Rose pink, same as the color before the HCl reaction
    2. Equilibria Involving Sparingly Solution Salts
        1. Na2CO3 (aq) – Transparent, slightly pinkish
        2. After adding AgNO3 (aq) – White precipitate formed
        3. After HNO3 reaction – White precipitate seems to dissolve, pale white
        4. After HCl reaction – Solution turned white, precipitate form
        5. After NH3 (aq) reaction – The solution turned transparent, gas evolved, extremely small amount of precipitate left
        6. After HNO3 reaction – Turned again into milky white, precipitate reappear
        7. After NH3 reaction – Turned transparent, however no gas evolved this time, precipitate disappear
        8. After adding KI – The solution turned milky green, white precipitates remained
    3. The Effect of Temperature on Equilibria
        1. Cool CoCl2 – Rose Pink
        2. Hot [CoCl4]2- – Very dark blackish rose pink

5. Analysis

    1. Changes in Reactant
      1. Copper and Nickel Ions
        1. [Cu(H2O)4]2+(aq) + 4NH3 (aq) ←→ [Cu(NH3)4]2+(aq) + 4H2O (l)
        2. [Ni(H2O)6]2+(aq) + 6NH3 (aq) ←→ [Ni(NH3)6]2+(aq) + 6H2O (l)
      • Addition of NH3 made the equilibrium shift right
          1. H+(aq) + NH3 (aq) ←→ NH4+
      • Reaction with HCl made the equilibrium shift left
      1. Cobalt Ions
        1. [Co(H2O)6]2+(aq) + 4Cl (aq) ←→ [CoCl4]2-(aq) + 6H2O (l)
      • Reaction with HCl made the equilibrium shift right
      • Reaction with H2O made the equilibrium shift left
    1. Equibria Involving Sparingly Soluble Salt
        1. 2Ag+(aq) + CO32-(aq) ←→ Ag2CO3 (s)
          1. 2H+(aq) + CO32-(aq) ←→ H2CO3 (aq)
          2. H2CO3 (aq) CO2 (aq) + H2O (l)
      • Reaction with HNO3 made the equilibrium shift left
        1. Ag+(aq) + Cl(aq) ←→ AgCl (s)
      • Reaction with HCl creates a new precipitate of AgCl and moves the equilibrium to the left
          1. Ag+(aq) + 2NH3 (aq) ←→ [Ag(NH3)2]+(aq)
      • Reaction with NH3 shifts the equilibrium to the left
          1. H+(aq) + NH3 (aq) ←→ NH4+(aq)
      • Reaction with HNO3 moves the equilibrium to the right
          1. I(aq) + Ag+(aq) ←→ AgI (s)
      • Reaction with the iodide ion creates a new precipitate of AgI and shifts the equilibrium to the left
    1. The Effect of Temperature on Equilibria
        1. CoCl2+ΔH Δ [CoCl4]2-
      • When the solution is heated, the equilibrium shifts to the right
      • When the solution is cooled, the equilibrium shifts to the left

6. Discussion

[Cu(H2O)4]2+(aq) + 4NH3 (aq) ←→ [Cu(NH3)4]2+(aq) + 4H2O (l)

[Ni(H2O)6]2+(aq) + 6NH3 (aq) ←→ [Ni(NH3)6]2+(aq) + 6H2O (l)

For above reaction NH3 reacts with both Cu2+ and Ni2+ ions to create [Cu(NH3)4]2+(aq) and [Ni(NH3)6]2+(aq) respectively. Thus, by adding NH3 to the solution, the equilibrium shifted right since while adding NH3 the concentration NH3 increases allowing more NH3 to react with the ions to create [Cu(NH3)4]2+(aq) and [Ni(NH3)6]2+(aq).

H+(aq) + NH3 (aq) ←→ NH4+

However, when HCl is added to neutralize NH3, the concentration of NH3 decreases, shifting the equilibrium left, thus creating less [Cu(NH3)4]2+(aq) and [Ni(NH3)6]2+(aq).

[Co(H2O)6]2+(aq) + 4Cl (aq) ←→ [CoCl4]2-(aq) + 6H2O (l)

Similar to above reaction, adding more HCl, the concentration of Cl(aq) increases shifting the equilibrium right. Through this reaction, more [CoCl4]2-(aq) is being created. However, when more water(H2O) is added, the concentration of H2O increases shifting the equilibrium left, creating less [CoCl4]2-(aq)

2Ag+(aq) + CO32-(aq) ←→ Ag2CO3 (s) …………….(1)

In reaction (1), CO32- reacts with Ag+ to create Ag2CO3 (s). The addition of Na2CO3 shifts the equilibrium to the right.

2H+(aq) + CO32-(aq) ←→ H2CO3 (aq)

H2CO3 (aq) CO2 (aq) + H2O (l)

In the above reaction, H+ reacts with CO32- to form H2CO3 that quickly decomposes to CO2 and water. Thus, addition of HNO3 reduced the concentration of CO32- ions shifting the equilibrium (1) left. Decrease in production of Ag2CO3 (s) is the result of this reaction. This process is observed through decrease in precipitates, Ag2CO3.

Ag+(aq) + Cl(aq) ←→ AgCl (s) …………………….(2)

From equilibrium (1), Cl from HCl reacts with Ag+ to create AgCl(aq). This decreases the concentration of Ag+(aq) shifting the equilibrium left. However, if excess HCl is added until chemical reaction has settled, the reaction has changed to equilibrium (2). This creates a new precipitate of AgCl(s).

Ag+(aq) + 2NH3 (aq) ←→ [Ag(NH3)2]+(aq)…………(3)

NH3 reacts with the silver ion to form [Ag(NH3)2]+(aq) which is a soluble substance. Thus, addition of NH3 reduces the concentration of Ag+(aq) ion shifting the equilibrium (2) left. This leads to less production of AgCl (s) so that the precipitates seem to disappear.

H+(aq) + NH3 (aq) ←→ NH4+(aq)

However, H+ react with NH3 to form NH4+(aq) which does not take part in reaction. This reduces the concentration of NH3 in equilibrium (3) which eventually increases the concentration of Ag+ in equilibrium (2). This shifts the equilibrium (2) right creating more AgCl (s) precipitate.

Process of adding NH3 and HNO3 – in this context H+ – is repeated several times. For all those times the equilibrium shifts for the same reason as explained above.

I(aq) + Ag+(aq) ←→ AgI (s)……………………….(4)

The iodide ion, I, reacts with Ag+ to form AgI(s) precipitate. This decreases the concentration of Ag+ in equilibrium (2) shifting the equilibrium left. However, if excess KI is added to the solution, the reaction completely changes to equilibrium (4).

CoCl2+ΔH Δ [CoCl4]2-

The process of CoCl2←→ [CoCl4]2- is an endothermic reaction. Thus, when heated, heat is given into the system. Thus, letting more heat get absorbed into the system shifts the equilibrium to the right.

On the other hand, if the reaction is cooled down, heat is taken out of the system. Thus, not letting heat take part in reaction shifts the equilibrium to the left.

Thus overall, all the reaction’s shift in equilibrium is explained by Le Châtelier’s principle. The change in concentration was the main reason of shifts in equilibrium. Furthermore, the change in heat demonstrated in the last reaction was also another factor that shifts equilibrium. Thus, Le Châtelier’s principle which explains the shift in equilibrium due to change in concentration and heat is well demonstrated throughout the experiment.

However, some errors that might be present on the process of the experiment are insufficient amount of substances (or even heat) that did not lead the experiment to completion and lack of observation.

One error observed was in the process of observing the color change of CoCl2+ΔH Δ [CoCl4]2- reaction. According to the text book, the color change in CoCl2+ΔH Δ [CoCl4]2- should have been from rose pink to blue. However, the color change observed was from rose pink to very dark blackish pink. One reason to explain this error is that, the amount of heat exerted in the reaction was insufficient to completely shift the equilibrium to its end but made it stop in the middle of shifting the equilibrium.

Another error that was observed was in the process of adding HNO3 to 2Ag+(aq) + CO32-(aq) ←→ Ag2CO3 (s). According to the process and its analysis, CO2 gas is supposed to bubble when HNO3 is added. However, there were no such bubbles observed. Two reasons that explain this error might be that, first; there wasn’t enough amount of solution present for any bubble to be observed and second; the amount of CO2 produced was too less that most of CO2 being produced dissolved in water before being able to form gas. The amount added to the 10mL beaker was 0.5mL of Na2CO3 which equals to about 10 drops, and a drop of AgNO3 which approximates to about 0.05mL. In total, only 0.55mL of solution was present in a 10mL beaker. This might have led to having too much surface area for bubble to be seen thus no significant bubble was observed.

Another error that was observed was that in process of adding highly concentrated acids and bases, gases evolved when it was not supposed to: when adding 12M HCl and when adding 15M NH3. One way to explain this is that HCl and NH3, being a very strong acid and a very strong base, reacts readily with water to form H2 gas for HCl and and NH4+ for NH3.

In order to improve the experiment, more time would have helped observe more definite data. Using definite data, observations of formation of more justifiable substances would have been given leading to more clear shifts of equilibrium. More time in the last experiment would have provided chances for more CoCl2 to form [CoCl4]2- thus changing the color into definite blue. To add on, addition of time would have given opportunities to carry experiments out for more than one time solidifying the argument of equilibrium shifts.

Work Cited

Brown and LeMay the Central Science Lab Manual

Brown, Lemay, Bursten, Murphy, Woodward, Stoltzfus. 2015. Chemistry the Central Science. 13th edition. United States of America: Pearson

John J. Carroll. John D. Slupsky. Alan E. Mather. June 13, 1991. The Solubility of Carbon Dioxide in Water at Low Pressure. Ref. Data, Vol 20, No. 6: University of Alberta

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