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


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:

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 mL mL 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

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

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

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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