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5.R: Le Chatelier's Principle (Lab Report)

  • Page ID
    127150
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    Name: ____________________________ Lab Partner: ________________________

    Date: ________________________ Lab Section: __________________

    Part A – Equilibrium and an Acid-Base Indicator

    Equilibrium system:

    \[\ce{ HA (aq) <=> H^{+} (aq) + A^{-} (aq)}\]

    Observations

    Record your results upon completing each of the following steps:

    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

    Analysis

    • Complete the following:

    The acidic form of the bromothymol blue indicator, \(\ce{HA}\) (aq), is _______________ in color.
    The basic form of the bromothymol blue indicator, \(\ce{A^{–}}\) (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.

    Part B – Solubility Equilibrium and \(K_{sp}\)

    Equilibrium system:

    \[\ce{PbCl2 (s) <=> Pb^{2+} (aq) + 2 Cl^{–} (aq)}\]

    Observations

    Step 4: Observations upon addition of just 1.0 mL of \(\ce{HCl}\) to the \(\ce{Pb(NO3)3}\) solution
    Step 5: Total volume of \(\ce{HCl}\) 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 \(\ce{PbCl2}\) precipitate mL
    Step 10: Total solution volume upon completion mL

    Analysis

    • Why didn’t any solid \(\ce{PbCl2}\) form immediately upon addition of 1 mL of \(\ce{HCl}\) (aq) in Step 4? What condition must be met by \([\ce{Pb^{2+}}]\) and \([\ce{Cl^{–}}]\) if solid \(\ce{PbCl2}\) is to form?

    • Consider your observation in hot water in Step 6:

    In which direction did the equilibrium shift? ____________________

    Did the value of \(K_{sp}\) get smaller or larger? ____________________

    Is the dissolution of \(\ce{PbCl2}\) (s) exothermic or endothermic? ____________________

    Explain below.


    • Explain why the solid \(\ce{PbCl2}\) dissolved when water was added to it in Step 9. What was the effect of this water on \([\ce{Pb^{2+}}]\), \([\ce{Cl^{–}}]\), and \(Q_{sp}\)? In which direction would such a change drive the equilibrium system?

    • The point at which the \(\ce{PbCl2}\) precipitate just dissolves in Step 9 can be used to determine the value of \(K_{sp}\) for this equilibrium system, where \(K_{sp} = [\ce{Pb^{2+}}][\ce{Cl^{–}}]^{2}\). Calculate \([\ce{Pb^{2+}}]\) and \([\ce{Cl^{–}}]\) in the final solution (consider the “dilution effect”). Then use these equilibrium concentrations to determine the value of \(K_{sp}\) for this system. Show all work below.

    Part C – Complex Ion Equilibria

    Equilibrium system:

    \[\underbrace{\ce{Co(H2O)6^{2+}(aq) }}_{\text{Pink}} + \ce{4Cl^{-} (aq) <=> } \underbrace{\ce{CoCl4^{2-}(aq) }}_{\text{Blue}} + \ce{6 H2O (l) }\]

    Observations

    Step 2: Color of solution in 12 M HCl
    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

    Analysis

    • What form of the complex ion, \(\ce{Co(H2O)6^{2+}}\) (aq) or \(\ce{CoCl4^{2–}}\) (aq), is predominate in:

    The 12 M \(\ce{HCl}\) (aq) __________________

    The diluted solution __________________

    The heated solution __________________

    • Explain why you obtained the observed color in 12 M \(\ce{HCl}\) (aq) (Step 2).
    • 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.

    • 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.


    Part D – Dissolving Insoluble Solids

    Equilibrium system:

    \[\ce{Zn(OH)2 (s) <=> Zn^{2+} (aq) + 2 OH^{-} (aq)}\quad K_{sp} << 1 \nonumber\]

    Observations

    Step 1: Adding 1 drop of \(\ce{NaOH}\) (aq) to \(\ce{Zn(NO3)2}\) (aq)
    Step 2: Tube A: Effect when \(\ce{HCl}\) (aq) is added
    Step 3: Tube B: Effect when \(\ce{NaOH}\) (aq) is added
    Step 4: Tube C: Effect when \(\ce{NH3}\) (aq) is added

    Analysis

    • Explain your observation upon addition of \(\ce{HCl}\) (aq) to the precipitate in Tube A. You must consider the various equilibria that are occurring in solution and the effect of \(\ce{HCl}\) on \([\ce{OH^{–}}]\).

    • Explain your observations upon addition of \(\ce{NaOH}\) (aq) to the precipitate in Tube B. Consider the various equilibria that are occurring in solution and remember that \(\ce{Zn^{2+}}\) forms stable complex ions with \(\ce{OH^{–}}\) at sufficiently high concentrations.

    • Explain your observations upon addition of \(\ce{NH3}\) (aq) to the precipitate in Tube C. Consider the various equilibria that are occurring in solution and remember that \(\ce{Zn^{2+}}\) forms stable complex ions with \(\ce{NH3}\).

    Part E

    Equilibrium system:

    \[\ce{Mg(OH)2 (s) <=> Mg^{2+} (aq) + 2 OH^{-} (aq)}\quad K_{sp} \ll 1\]

    Observations

    Step 1 Adding 1 drop of \(\ce{NaOH}\) (aq) to \(\ce{Mg(NO3)2}\) (aq)
    Step 2: Tube A: Effect when \(\ce{HCl}\) (aq) is added
    Step 3: Tube B: Effect when \(\ce{NaOH}\) (aq) is added
    Step 4: Tube C: Effect when \(\ce{NH3}\) (aq) is added

    Analysis

    • Based on your observations in Steps 3 and 4 do you think that \(\ce{Mg^{2+}}\) forms stable complex ions? Explain your reasoning.

    5.R: Le Chatelier's Principle (Lab Report) is shared under a not declared license and was authored, remixed, and/or curated by Santa Monica College.

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