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2605 Le Chatelier's Principle

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    440622
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    1.0 INTRODUCTION

    1.1 Objectives

    After completing this experiment, the student will be able to:

    • Perturb chemical reactions at equilibrium and observe how they respond.
    • Explain these observations using Le Chatelier’s Principle.
    • Relate Le Chatelier’s Principle to the concept of coupled reactions.

    1.2 Background

    All chemical reactions eventually reach a state in which the rate of the reaction in the forward direction is equal to the rate of the reaction in the reverse direction. When a reaction reaches this state, it is said to be at chemical equilibrium.

    The concentrations of reactants and products at equilibrium are constant as a function of time. Thus, for a homogeneous aqueous system of the form

    aA(aq) + bB(aq) cC(aq) + dD(aq) (1)

    we can express the equilibrium-constant expression for this reaction as,

    Screenshot_20230615_141712.png (2)

    where the values of [A], [B], [C] and [D] correspond to the equilibrium concentrations (or equilibrium positions) of all the aqueous chemical components, and a, b, c and d are their respective stoichiometric coefficients. Note that for a heterogeneous system including pure solids or liquids of the form

    aA(aq) + bB(s) cC(aq) + dD(l) (3)

    the pure liquids and solids do not appear in the equilibrium-constant expression:

    Screenshot_20230615_141734.png (4)

    It has been observed that when a reaction at equilibrium is perturbed by applying a stress, the reaction will respond by shifting its equilibrium position to counteract the effect of the perturbation/stress. In other words, the concentrations of the reactants and products will shift so that the relationship described by Equation (2) is again satisfied. This idea was first proposed by Henri-Louis Le Châtelier and has since been referred to as, “Le Châtelier’s Principle”.

    Note that when a reaction makes more products as a response to the perturbation, we call it a right shift. When a reaction makes more reactants in response to the perturbation, we call it a left shift. We often designate these respective shifts by drawing right and left arrows below the chemical equation.

    For chemical reactions at equilibrium in aqueous solution, the most common types of perturbations include changing the concentration of one of the aqueous solutes, changing the concentrations of all aqueous solutes by changing the total solution volume, or changing the temperature. The general responses of an aqueous system to these particular perturbations are tabulated below.

    Perturbation

    Effect on Equilibrium Position

    Effect on Kc

    Increase in concentration of a single reactant, or, decrease in concentration of a single product.

    Shift to the right

    None

    Decrease in concentration of a single reactant, or, increase in concentration of a single product.

    Shift to the left

    None

    Decrease in all aqueous concentrations due to an increase in solution volume resulting from the addition of solvent

    Shift towards side with more solute particles

    None

    Increase in all aqueous concentrations due to a decrease in solution volume resulting from the removal of solvent (evaporation)

    Shift towards side with less solute particles

    None

    Increase temperature of an exothermic reaction

    Shift to the left

    Decrease

    Decrease temperature of an exothermic reaction

    Shift to the right

    Increase

    Increase temperature of an endothermic reaction

    Shift to the right

    Increase

    Decrease temperature of an endothermic reaction

    Shift to the left

    Decrease

    Addition of an inert substance, catalyst, pure liquid, or pure solid

    None

    None

    Notice that only a temperature change can affect the value of Kc; in all other cases, the value of Kc remains constant.

    In this experiment, you will perturb reactions that have attained equilibrium. You will then observe how each reaction responds to that perturbation in order to restore equilibrium. In your report you describe these changes in terms of Le Chatelier’s Principle.

    1.3 Specific Experimental Examples

    1.3.1 Acid-Base Equilibrium

    Here you will use coupled equilibria to change the equilibrium position of an acid-base

    reaction. In order to understand how coupled equilibria work, consider the reactions described by the chemical equations below:

    A(aq) B(aq) (5)

    B(aq) + C(aq) D(aq) (6)

    Notice that the species B(aq) is common to both reactions. The presence of this common

    species couples these two reactions.

    We can perturb the equilibrium position of Reaction (6) by the addition of some C(aq). The addition of C(aq) will cause the equilibrium position of Reaction (6) to shift right in accordance with Le Chatelier’s Principle. This right shift in the equilibrium position of Reaction (6) will also result a corresponding decrease in the concentration of B(aq). Because B(aq) is also present in Reaction (5), the decrease in the concentration of B(aq) will in turn result in a right shift in the equilibrium position of Reaction (5). Thus, the addition of C(aq) to Reaction (6) actually results in a right shift in the equilibrium position of Reaction (5) because the equilibria are coupled.

    In Part A we will observe the effect of various solutes on an acid-base indicator (a weak acid) at equilibrium. The equilibrium system can be written in the general form

    HA(aq) H+(aq) + A(aq) (7)

    The equilibrium constant expression for this reaction is

    Screenshot_20230615_141804.png (8)

    where we denote the equilibrium constant, K, with a subscript a for acid. In this experiment, HA and A, are the acidic and basic forms of the indicator bromothymol blue. Since the two forms are different colors, you will be able to determine which form is predominant in the equilibrium mixture. In other words, you will be able to determine whether the equilibrium position lies to the left (more reactants and less products) or whether the equilibrium lies to the right (more products and less reactants).

    Your goal will be to find a reagent that will shift the position of this equilibrium to the opposite side, and then another reagent that will shift it back towards its original position. Instead of directly adding HA or A- to the system, you will cause these shifts by adding H+ or OH-. Note that in order to determine the effect of OH- we must consider a second chemical reaction that shares a common species with the Reaction (7). The second reaction is the autoionization of water, which can be described by the equation

    H2O(l) H+(aq) + OH-(aq) (9)

    The equilibrium constant for this reaction is denoted by Kw, where the subscript w stands for water, and the associated equilibrium constant expression is

    Screenshot_20230615_141826.png (10)

    Because Reactions (7) and (9) share a common chemical species (H+), you can use the concept of coupled equilibria to shift the equilibrium position of Reaction (7) by increasing or decreasing the concentration of OH-(aq).

    1.3.2 Solubility Equilibrium

    Here you will test the effects of changing temperature and volume on the solubility of a slightly soluble salt at equilibrium. Some examples of slightly soluble salts are AgCl, Cu(OH)2, PbCl2, and Fe2S3, which you should recall are, “insoluble in water,” according to the solubility rules you learned in Chem 25. In fact, a very small amount of each of these substances does dissolve in aqueous solution, but the amount is so small that we often classify each of these compounds as “insoluble”.

    This type of equilibrium is often called a solubility equilibrium because it is written in the direction of the dissolution of the solid, as shown in the following example:

    AxBy(s) xA+(aq) + yB(aq) (11)

    The equilibrium constant expression for Reaction (11) is

    Screenshot_20230615_141852.png (12)

    where we denote the equilibrium constant, K, with a subscript sp for solubility product.

    Now let’s consider the process of precipitation. In a typical precipitation reaction two aqueous salt solutions are mixed together resulting in the production of an insoluble salt. Notice that this process corresponds to a left shift of Reaction (11), and so Equation (12) can also be used to examine the conditions required for the precipitation of a solid to occur. We can denote the product [A+]x[B]y under arbitrary conditions (not necessarily at equilibrium) as,

    Screenshot_20230615_141913.png (13)

    where Qsp is called the solubility product reaction quotient. Note that, upon mixing two solutions, one containing A and the other containing B , if Qsp < Ksp the system is not at equilibrium, but since no solid AxBy is present, the reaction cannot shift to the right and therefore no reaction will be observed. In contrast, if Qsp > Ksp the solution contains an excess of aqueous species, and Reaction (11) will shift left, forming the solid precipitate AxBy until the system reaches a state of equilibrium where Qsp = Ksp. Thus, we can use the values of Qsp and Ksp to predict the conditions under which a precipitation reaction will occur.

    In Part B, we will study the solubility equilibrium of PbCl2(s). We will observe the effect on this solubility equilibrium of changes in solution volume (quantity of solvent) and temperature. We will express these changes in terms of the respective values of Qsp and Ksp.




    1.3.3 Complex Ion Equilibrium

    Certain metal ions, most often transition metals, exist in solution as complex ions in combination with other ions or molecules, called ligands. Common ligands include H2O, NH3, Cl and OH-. Many of these complex ions exhibit vibrant colors in solution. For example, the Co(H2O)62+(aq) complex ion is pink and the CoCl42-(aq) complex ion is blue. In Part C, you will study the following complex ion formation reaction:

    Co(H2O)62+(aq) + 4 Cl-(aq) CoCl42- (aq) + 6 H2O (l) (14)

    The equilibrium constant expression for Reaction (14) is

    Screenshot_20230615_141943.png (15)

    where we denote the equilibrium constant, K, with a subscript f for complex ion formation.

    Your goal in Part C is to observe how Reaction (14) shifts from its equilibrium position as the result of various perturbations.

    1.3.4 Dissolving Insoluble Solids

    In Part D, you will use coupled equilibria to affect the solubility equilibrium of Zn(OH)2(s). The solubility equilibrium can be described by the equation

    Zn(OH)2(s) Zn2+(aq) + 2 OH-(aq) Ksp = 5 x 10-17 (16)

    Notice that Ksp << 1 for this reaction, demonstrating that Zn(OH)2(s) is only very slightly soluble in aqueous solution.

    Now consider the reactions described by the following chemical equations, each of which shares a common species with the Reaction (16):

    H2O(l) H+(aq) + OH-(aq) Kw = 1 x 10-14 (17)

    Zn2+(aq) + 4 OH-(aq) Zn(OH)42-(aq) Kf = 3 x 1015 (18)

    Zn2+(aq) + 4 NH3(aq) Zn(NH3)42+(aq) Kf = 1 x 109 (19)

    Because Reactions (17), (18), and (19) each share a common species with Reaction (16) they can be coupled together. In Part D of this experiment, you will observe the effect of coupling each of these equilibria on the solubility of Zn(OH)2(s).


    2.0 SAFETY PRECAUTIONS AND WASTE DISPOSAL

    !!Wear your safety goggles!!

    12-M HCl is extremely caustic and great care must be taken to avoid contact with eyes or skin. The bottle should be kept in a plastic tray and not removed from the fume hood. Should any of this solution enter your eyes rinse immediately in the emergency eyewash. Should any of this solution come in contact with your skin rinse with copious amounts of water and apply saturated sodium bicarbonate to the affected area from the stock bottle located on the sink.

    The solutions you will use in Part B contain lead. Be certain that all of these lead-containing solutions are disposed of in the proper waste container and rinse your hands following this procedure.

    Many of the chemicals used in this lab are hazardous to the environment. All metal ion waste (lead, zinc, magnesium, cobalt ions) must be disposed of in the hazardous-waste container in the fume hood. Rinse all glassware directly into the waste container twice using a small squirt bottle to be certain all hazardous waste ends up in the waste container.

    3.0 CHEMICALS AND SolutionS

    Chemical

    Concentration

    Approximate Amount

    Notes

    Laboratory water

    Pure

    25 mL

     

    Bromothymol blue

     

    several drops

     

    NaOH(aq)

    6 M

    3 mL

     

    HCl(aq)

    6 M

    3 mL

     

    Pb(NO3)2(aq)

    0.3 M

    5 mL

     

    HCl(aq)

    0.3 M

    5 mL

     

    HCl(aq)

    12 M

    2 mL

     

    CoCl2 6 H2O(s)

     

    tip of the scoopula

     

    Zn(NO3)2(aq)

    0.1 M

    6 mL

     

    Mg(NO3)2(aq)

    0.1 M

    6 mL

     

    NH3(aq)

    6 M

    3 mL

    (often labeled as NH4OH)

    4.0 GLASSWARE AND APPARATUS

    6 medium test tubes

    test tube rack

    10 mL graduated cylinder

    stirring rod

    20 mL graduated cylinder

    scoopula

    400 mL beaker (waste container)

     

    Bucket of ice from Stockroom

    5.0 PROCEDURE

    5.1 Acid-Base Equilibrium

    1. Add approximately 5 mL of laboratory water to a medium test tube. Add 3 drops of the bromothymol blue indicator solution. Record the color of your solution on your data recording sheet.
    2. Your solution from Step 1 currently contains one form of bromothymol blue (see background). Now predict which of the two 6 M reagents you obtained, the strong acid or the strong base, will cause a color change in your solution by making the bromothymol blue indicator shift to its other form. Add the 6 M reagent of your choice drop-by-drop and if your solution changes color, write the color of the solution and formula of the reagent on your data recording sheet. If the addition of your reagent does not result in a color change, try other reagents until you are successful.
    3. Since equilibrium systems are reversible, it is possible to shift a reaction left or right repeatedly by changing the conditions. Now find another 6 M reagent that will cause your solution from Step 2 to revert back to its original color. Add the 6 M reagent of your choice drop-by-drop and if your solution changes color, write the formula of the reagent on the data recording sheet. If the addition of your second reagent does not result in a color change, try other reagents until you are successful.

    5.2 Solubility Equilibrium

    1. Set up a hot-water bath by filling a 400-mL beaker about half-full with tap water and heating the water on a hot plate while you work on the following steps. The hot-water bath will be used in Step 6 (and also in 5.3).
    2. Measure 5.0 mL of the 0.3 M Pb(NO3)2(aq) solution into a medium test tube using a 10-mL graduated cylinder.
    3. Thoroughly rinse the graduated cylinder with laboratory water, then measure 5.0 mL of the 0.3 M HCl(aq) solution using the 10-mL graduated cylinder.
    4. Add about 1 mL of the 0.3 M HCl(aq) solution to the Pb(NO3)2(aq) solution in the medium test tube. Stir the mixture gently using your stirring-rod and record your observations on your data recording sheet.
    5. Continue to add the 0.3 M HCl(aq) solution to the Pb(NO3)2(aq) solution in the medium test tube in roughly 1 mL increments until you just begin to see white PbCl2 solid appear in your test tube. To confirm that the solid is present, let the test tube sit on the bench for about 3 minutes, allowing all solid to settle to the bottom where it is easier to see. On your data recording sheet, record the total volume of 0.3 M HCl(aq) needed to produce the solid.
    6. Put the test tube containing the solid into your hot-water bath from Step 1. Stir the contents of the test tube gently for a few seconds using the glass stirring rod and record your observations on your data recording sheet. Continue heating and stirring until a change is observed (about 50oC).
    7. Make a cold-water bath by filling a 400-mL beaker half full of tap water and ice. Cool the test tube down by placing it into this cold-water bath. Observe what happens and record your observations. Save this cold-water bath for use in 5.3.
    8. Dispose of the remainder of the 0.3 M HCl(aq) solution in your 10-mL graduated cylinder by pouring it into the appropriate waste container and then rinse the cylinder using laboratory water. Next measure 5.0 mL of laboratory water using this graduated cylinder.
    9. Add about 1 mL of deionized water to the medium test tube containing the PbCl2 solid. Stir the mixture and observe what happens. Continue to add the laboratory water in 1 mL increments until the white PbCl2 solid just dissolves (or disappears). On your data recording sheet, record the volume of water needed to dissolve the solid.
    10. Pour the contents of the medium test tube into your large graduated cylinder and measure the total volume of the solution. Record this volume on your data recording sheet.
    11. Dispose of all solutions used in Part B in the proper waste container. Rinse all glassware twice using a wash bottle of laboratory water to ensure that all of the lead solution has been removed and transferred to the waste container. Save the hot- and cold-water baths for 5.3.

    5.3 Complex Ion Equilibrium

    1. Reheat your hot-water bath from 5.2 to a near boil. The hot-water bath will be used in Step 4.
    2. Put a few small crystals (an amount that fits at the tip of the scoopula) of CoCl2 6H2O(s) in a medium test tube. Working in the fume hood, carefully add 2 mL of 12 M HCl(aq) (this acid is extremely caustic; see cautions in the safety section of this experiment). Stir to dissolve the crystals. Record the color of this solution.
    3. Using a 10-mL graduated cylinder, add laboratory water to the solution in your test tube in approximately 2 mL increments, stirring after each addition, until no further (obvious) color change occurs (make sure you are not just diluting the solution!). Record the new color.
    4. Place the test tube into the hot-water bath from Step 1 and record any color change.
    5. Cool the solution down in the cold-water bath and record any color change.

    5.4 Dissolving Insoluble Solids

    1. Label three medium test tubes A, B, and C. Add about 2 mL of 0.1 M Zn(NO3)2(aq) solution to each test tube. Add one drop of 6 M NaOH(aq) solution to each test tube. Stir each solution and record your observations.
    2. Into test tube A, add at least 20 drops of 6 M HCl(aq) drop-by-drop while stirring. Record all observations.
    3. Into test tube B, add at least 20 drops of 6 M NaOH(aq) drop-by-drop while stirring. Record all observations.
    4. Into test tube C, add at least 20 drops of 6 M NH3(aq) drop-by-drop while stirring. Record all observations.
    5. Label three additional medium test tubes D, E, and F. Add about 2 mL of 0.1 M Mg(NO3)2(aq) solution to each test tube. Add one drop of 6 M NaOH(aq) solution to each test tube. Stir each solution and record your observations.
    6. Into test tube D, add at least 20 drops of 6 M HCl(aq) drop-by-drop while stirring. Record all observations.
    7. Into test tube E, add at least 20 drops of 6 M NaOH(aq) drop-by-drop while stirring. Record all observations.
    8. Into test tube F, add at least 20 drops of 6 M NH3(aq) drop-by-drop while stirring. Record all observations.




























    6.0 DATA RECORDING SHEET

    Last Name

    First Name

     

    Partner Name(s)

    Date

    6.1 Acid-Base Equilibrium

    Equilibrium system: HA(aq) H+(aq) + A(aq)

    Record your observations for each of the following steps in the table below.

    Step 1

    Color of bromothymol blue in distilled water

     

    Step 2

    Name of reagent “A” causing color change when added

     

    Step 3

    Name of reagent “B” causing a return to original color

     

    The acidic form of the bromothymol blue indicator, HA(aq), is _____________ in color.

    The basic form of the bromothymol blue indicator, A(aq), is ______________ in color.

    Explain why reagent A (in Step 2) caused the color change observed.










    Explain why reagent B (in Step 3) caused the color change observed.







    6.2 Solubility Equilibrium

    Equilibrium system: PbCl2(s) Pb2+(aq) + 2 Cl(aq)

    Record your observations for each of the following steps in the table below.

    Step 4

    Observations upon addition of just 1.0 mL of HCl(aq) to the Pb(NO3)3(aq) solution

     

    Step 5

    Total volume of HCl(aq) required for noticeable precipitation (mL)

    Step 6

    Observations upon placing the test tube with precipitate in hot water

     

    Step 7

    Observations upon placing the test tube with precipitate in cold water

     

    Step 9

    Volume of water added to just dissolve PbCl2 precipitate (mL)

    Step 10

    Total solution volume upon completion (mL)

    1. Why didn’t any solid PbCl2 form immediately upon addition of 1 mL of HCl(aq) in Step 4? What condition must be met by [Pb2+] and [Cl] if solid PbCl2 is to form?







    1. Consider your observation in hot water in Step 6:
    • In which direction did the equilibrium shift? _______________
    • Did the value of Ksp get smaller or larger? _______________
    • Is the dissolution of PbCl2 (s) exothermic or endothermic? _______________

    Explain below.






    1. Explain why the solid PbCl2 dissolved when water was added to it in Step 9. What was the effect of this water on [Pb2+], [Cl], and Qsp? In which direction would such a change drive the equilibrium system?









    1. The point at which the PbCl2 precipitate just dissolves in Step 9 can be used to determine the value of Ksp for this equilibrium system, where Ksp = [Pb2+][Cl]2. Calculate [Pb2+] and [Cl] in the final solution (consider the “dilution effect”). Then use these equilibrium concentrations to determine the value of Ksp for this system. Show all work below.












    6.3 Complex Ion Equilibrium

    Equilibrium system: Co(H2O)62+(aq) + 4 Cl-(aq) CoCl42-(aq) + 6 H2O(l)

    pink blue

    Record your observations for each of the following steps in the table below.

    Step 2

    Color of solution in 12 M HCl(aq)

     

    Step 3

    Color of solution upon addition of water

     

    Step 4

    Color of solution in hot water

     

    Step 5

    Color of solution in cold water

     
    1. What form of the complex ion, Co(H2O)62+(aq) or CoCl42-(aq), is predominant in:

    the 12 M HCl(aq) _________________

    the diluted solution _________________

    the heated solution _________________

    1. Explain why you obtained the observed color in 12 M HCl(aq) (Step 2).







    1. Explain the observed color change that occurred when water was added to the solution in Step 3. Consider how water affects the ion concentrations and Q in this system.







    1. Consider your observations in the hot water bath in Step 4.

    In which direction did the equilibrium shift? _________________

    Did the value of K get smaller or larger? _________________

    Is the reaction (as written) exothermic or endothermic? _________________

    Explain.

    6.4 Dissolving Insoluble Solids

    Equilibrium system: Zn(OH)2(s) Zn2+(aq) + 2 OH-(aq) Ksp << 1

    Record your observations for each of the following steps in the table below.

    Step 1

    Adding 1 drop of NaOH(aq) to Zn(NO3)2(aq)

     

    Step 2

    Tube A: Effect when HCl(aq) is added

     

    Step 3

    Tube B: Effect when NaOH(aq) is added

     

    Step 4

    Tube C: Effect when NH3(aq) is added

     
    1. Explain your observation upon addition of HCl(aq) to the precipitate in Tube A. You must consider the various equilibria that are occurring in solution and the effect of HCl(aq) on [OH-].







    1. Explain your observations upon addition of NaOH(aq) to the precipitate in Tube B. Consider the various equilibria that are occurring in solution and remember that Zn2+ forms stable complex ions with OH- at sufficiently high concentrations.






    1. Explain your observations upon addition of NH3(aq) to the precipitate in Tube C. Consider the various equilibria that are occurring in solution and remember that Zn2+ forms stable complex ions with NH3.



    Equilibrium system: Mg(OH)2(s) Mg2+(aq) + 2 OH-(aq) Ksp << 1

    Record your observations for each of the following steps in the table below.

    Step 5

    Adding 1 drop of NaOH(aq) to Mg(NO3)2(aq)

     

    Step 6

    Tube D: Effect when HCl(aq) is added

     

    Step 7

    Tube E: Effect when NaOH(aq) is added

     

    Step 8

    Tube F: Effect when NH3 (aq)is added

     
    1. Explain your observation upon addition of HCl(aq) to the precipitate in Tube D.







    1. Based on your observations in Steps 7 and 8 do you think that Mg2+ forms stable complex ions? Explain your reasoning.




















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