Experiment# 10 Determining the Ksp of Calcium Hydroxide -3 PDF

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Experiment# 10 Determining the Ksp of Calcium Hydroxide -3 leacture practice lecture practice Experiment# 10 Determining the Ksp of Calcium Hydroxide -3Experiment# 10 Determining the Ksp of Calcium Hydroxide -3...

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Experiment #10

Ksp Determination for Calcium Hydroxide

Learning Outcomes After completing this experiment, you should be able to: • • • • •

Write chemical equations and equilibrium expressions representing solubility equilibria Carry out equilibrium computations involving solubility, equilibrium expressions, and solute concentrations Describe the equilibrium of a saturated solution macroscopically and microscopically with supporting illustrations. Write equilibrium expressions for salts dissolving Calculate K sp from experimental solubility

Background Information From OpenStax Chemistry 2e Solubility equilibria are established when the dissolution and precipitation of a solute species occur at equal rates. These equilibria underlie many natural and technological processes, ranging from tooth decay to water purification. An understanding of the factors affecting compound solubility is, therefore, essential to the effective management of these processes. This section applies previously introduced equilibrium concepts and tools to systems involving dissolution and precipitation. The Solubility Product Recall from studying solutions that the solubility of a substance can vary from essentially zero (insoluble or sparingly soluble) to infinity (miscible). A solute with finite solubility can yield a saturated solution when it is added to a solvent in an amount exceeding its solubility, resulting in a heterogeneous mixture of the saturated solution and the excess, undissolved solute. For example, a saturated solution of silver chloride is one in which the equilibrium shown below has been established. dissolution

⇌

AgCl(s)

Ag+ (aq) + Cl− (aq)

precipitation

In this solution, an excess of solid AgCl dissolves and dissociates to produce aqueous Ag+ and Cl– ions at the same rate that these aqueous ions combine and precipitate to form solid AgCl (Figure 1). Because silver chloride is a sparingly soluble salt, the equilibrium concentration of its dissolved ions in the solution is relatively low. The equilibrium constant for solubility equilibria such as this one is called the solubility product constant, Ksp, in this case: AgCl(s)

⇌

Ag+ (aq) + Cl− (aq)

Ksp=[Ag+][Cl−]

2 Recall that only gases and solutes are represented in equilibrium constant expressions, so the Ksp does not include a term for the undissolved AgCl. A listing of solubility product constants for several sparingly soluble compounds is provided in various appendix materials.

Figure 1 Silver chloride is a sparingly soluble ionic solid. When it is added to water, it dissolves slightly and produces a mixture containing a very dilute solution of Ag+ and Cl- ions in equilibrium with undissolved silver chloride

Refer to a few examples given in Figure 2 for writing solubility product constant (Ksp) expressions

Figure 2 Some examples of writing the solubility equilibrium reaction (dissolution reaction) and the solubility product constant expression

3 Ksp and Solubility The Ksp of a slightly soluble ionic compound may be simply related to its measured solubility provided the dissolution process involves only dissociation and solvation, for example:

For cases such as these, one may derive Ksp values from provided solubilities, or vice-versa. Calculations of this sort are most conveniently performed using a compound’s molar solubility, measured as moles of dissolved solute per liter of saturated solution. Typical calculations from this equilibrium system include calculating Ksp from solubility equilibrium data as well as calculating the solubility from the Ksp constant as given in the examples below.

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Common Ion Effect Compared with pure water, the solubility of an ionic compound is less in aqueous solutions containing a common ion (one also produced by dissolution of the ionic compound). This is an example of a phenomenon known as the common ion effect, which is a consequence of the law of mass action that may be explained using Le ChÂtelier’s principle. Consider the dissolution of silver iodide: AgI(s) ⇋ Ag+(aq) + I-(aq) This solubility equilibrium may be shifted left by the addition of either silver(I) or iodide ions, resulting in the precipitation of AgI and lowered concentrations of dissolved Ag+ and I–. In solutions that already contain either of these ions, less AgI may be dissolved than in solutions without these ions. This effect may also be explained in terms of mass action as represented in the solubility product expression: Ksp = [Ag+][I-] The mathematical product of silver(I) and iodide ion molarities is constant in an equilibrium mixture regardless of the source of the ions, and so an increase in one ion’s concentration must be balanced by a proportional decrease in the other.

5 In this experiment, you will determine the molar solubility and the Ksp of calcium hydroxide, Ca(OH)2 using the following solubility equilibrium equation: Ca(OH)2 (s) ⇋ Ca2+(aq) + 2OH-(aq)

Ksp = [Ca2+][OH-]2

A saturated solution of Ca(OH)2 will be titrated dropwise using hydrochloric acid. From the titration results, the hydroxide ion [OH-] concentration present in the saturated calcium hydroxide solution can be calculated. From the stoichiometry of the solubility reaction the concentration of [Ca2+] as well as the molar solubility and Ksp can also be determined. The titrant in this experiment will be a known concentration of hydrochloric acid, HCl and will be added via a transfer pipet to a sample of the saturated calcium hydroxide solution (aka limewater) until the equivalence point is reached. Remember that the equivalence point in a titration is the point where the moles of added titrant is stoichiometrically equal to the moles of analyte. Hydrochloric acid neutralizes the hydroxide ion from calcium hydroxide according to the following net reaction: H+(aq) + OH-(aq) → H2O(l) The endpoint of this titration will be identified with the use of an indicator known as methyl orange. Methyl orange is a dye that turns red in acid and yellow in base. At the endpoint of this titration which is when the hydroxide ions are completely neutralized, a persistent orange-red color forms. In the second part of the experiment, the common ion effect will be explored. The solubility of the calcium hydroxide in the saturated solution will be lowered by the addition of some calcium chloride which supplies the common ion Ca2+. The additional Ca2+ ions supplied by the calcium chloride will then shift the solubility equilibrium reaction back towards reactants. + Ca2+ (from CaCl2) Ca(OH)2 (s) ⇋ Ca2+(aq) + 2OH-(aq) In the presence of additional [Ca2+], less hydroxide ions would be available in solution for titration. This set of titration data will be used to determine the molar solubility in the presence of a common ion. Safety Precautions Be sure to wear your safety goggles and gloves while performing this experiment. Wear long pants as well as closed toed shoes. Hydrochloric acid and calcium hydroxide are corrosive agents. If either of these contact your skin or eyes, flush the affected area immediately with water for at least 15 minutes. Calcium chloride is a serious eye irritant and can be harmful if swallowed. Keep all food/drink, pets, and children away from these chemicals. Protect your work area and clean up the work area with soap and water when finished. Reference: Paul Flowers, et. al. (2019) Chemistry: Atoms First, 2e. Available at https://openstax.org/details/books/chemistry-atoms-first-2e (Downloaded: 20 October 2020).

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PROCEDURE Part I. Determine the Molar Solubility of Ca(OH)2 in Water 1. Using the balance, measure & record the mass of the 25mL Erlenmeyer flask. 2. Measure 3.00 mL of the saturated Ca(OH)2 (lime water) and transfer to the Erlenmeyer flask. 3. Measure and record the mass of (flask + lime water). 4. Add 5 drops of the methyl orange indicator to the lime water. Solution should be yellow. 5. Obtain the HCl dropper bottle. You will assume for the purposes of this titration that for every drop delivered from the HCl dropper bottle is equivalent to 0.05 mL. Be sure to check that the dropper portion of your HCl bottle is firmly seated before using it. 6. Using a white background to perceive the color changes better, add HCl dropwise while gently swirling the reaction mixture in the Erlenmeyer flask. Continue adding the HCl dropwise until the endpoint of the titration is reached (which is a persistent orange-red color). 7. Take a photo of the reaction mixture at the endpoint of the titration. 8. Repeat the titration process once more using a fresh 3.00 mL sample of saturated Ca(OH)2.

Calculations: • Convert the # drops HCl to mL HCl using the relationship 1 drop = 0.05 mL • Using the given molarity of HCl, calculate the moles of HCl used in the titration. • Using the balanced net titration reaction between H+ and OH-, calculate the moles of OHpresent in the saturated Ca(OH)2 analyzed. • Using the stoichiometric ratio between Ca2+ and OH- in the solubility equilibrium reaction, determine the moles of Ca2+ in the saturated Ca(OH)2 • Accounting for the volume of saturated Ca(OH)2 analyzed, calculate the molarity of Ca2+ in the saturated Ca(OH)2 solution which is also the molar solubility. • Using the molar solubility, calculate the Ksp

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Part II. Determine the Molar Solubility of Ca(OH)2 in solution of CaCl2 1. Remove the tip from the saturated Ca(OH)2 (lime water) bottle and tare out its mass on the electronic balance ( balance should read 0.00 grams). 2. Using a plastic spoon (or knife) transfer approximately 0.03-0.06g of calcium chloride crystals to the lime water. 3. Reinsert firmly, the dropper tip on the lime water bottle, recap, and shake thoroughly until all the calcium chloride dissolves. Let the bottle sit for approximately 20-30 minutes before using it. 4. Measure the mass of an empty 25mL Erlenmeyer flask (Record in Data Table 2) 5. Carefully remove 3.00 mL of the saturated Ca(OH)2 / CaCl2 solution (without transferring any sediment that may be at the bottom of the bottle) and transfer to the Erlenmeyer flask. 6. Measure and record the mass of (flask + calcium hydroxide/calcium chloride solution). 7. Add 5 drops of the methyl orange indicator to the lime water. Solution should be yellow. 8. Obtain the HCl dropper bottle. You will again assume for the purposes of this titration that for every drop delivered from the HCl dropper bottle is equivalent to 0.05 mL. Be sure to check that the dropper portion of your HCl bottle is firmly seated before using it. 9. Using a white background to perceive the color changes better, add HCl dropwise while gently swirling the reaction mixture in the Erlenmeyer flask. Continue adding the HCl dropwise until the endpoint of the titration is reached (which is a persistent orange-red color). 10. Take a photo of the reaction mixture at the endpoint of the titration. 11. Repeat the titration process once more using a fresh 3.00 mL sample of calcium hydroxide/calcium chloride solution.

Calculations: • Convert the # drops HCl to mL HCl using the relationship 1 drop = 0.05 mL • Using the given molarity of HCl, calculate the moles of HCl used in the titration. • Using the balanced net titration reaction between H+ and OH-, calculate the moles of OHpresent in the saturated Ca(OH)2 analyzed. • Using the stoichiometric ratio between Ca2+ and OH- in the solubility equilibrium reaction, determine the moles of Ca2+ in the saturated Ca(OH)2 • Accounting for the volume of saturated Ca(OH)2 analyzed, calculate the molarity of Ca2+ in the saturated Ca(OH)2 solution which is also the molar solubility (in the presence of a common ion). Disposal/Clean Up All titration mixtures can be discarded directly in the sink. Clean glassware and return to kit. Pour remaining reagents into paper towels and discard in the trash.

8 Experiment #10 Ksp Determination for Calcium Hydroxide DATA SHEET Name:_____________________________________________

Date:_________________

Part I. Determination of the Molar Solubility of Ca(OH)2 in Water & Ksp Data Table 1 Trial 1 Mass of Erlenmeyer Flask Mass of Erlenmeyer Flask + Calcium Hydroxide Solution (lime water) Mass of Calcium Hydroxide Solution Volume of Ca(OH)2 Solution, mL (Assume Solution Density = 1.000 g/mL)

Concentration of HCl (M) Drops of HCl needed to reach Endpoint Volume of HCl Delivered

Moles of HCl Delivered

Moles of OH- in Saturated Ca(OH)2 Moles of Ca2+ in Saturated Ca(OH)2 (remember 2:1 stoichiometry) Molar Solubility (M) remember to use L of lime water measured out (analyzed)

Calculated Ksp

Average Calculated Ksp

Trial 2

9 Part II. Determination of the Molar Solubility of Ca(OH)2 in Solution of CaCl2 Name:_____________________________________________

Date:_________________

Data Table 2 Trial 1 Mass of Erlenmeyer Flask Mass of Erlenmeyer Flask + Calcium Hydroxide Solution (lime water) Mass of Calcium Hydroxide Solution Volume of Ca(OH)2 Solution, mL (Assume Solution Density = 1.000 g/mL)

Concentration of HCl (M) Drops of HCl used Volume of HCl Delivered Moles of HCl Delivered Moles of OH- in Sample Moles of Ca2+ in Saturated Ca(OH)2 (remember 2:1 stoichiometry) Molar Solubility (M) remember to use L of lime water measured out (analyzed)

Average molar solubility

Trial 2

10 Name:_____________________________________________

Date:_________________

Photos Photo 1 Insert the photo of the Erlenmeyer flask with lime water and methyl orange indicator at the endpoint of the titration in Part I.

Photo 2 Insert the photo of the Erlenmeyer flask with lime water and methyl orange indicator at the endpoint of the titration in Part II.

11 Experiment #10 Ksp Determination for Calcium Hydroxide POST-LAB QUESTIONS Name:_____________________________________________

Date:_________________

(You must show work for credit) 1. The solubility of silver carbonate (mm = 275.745), is 33.5 mg/L. What is its molar solubility, and what is its Ksp?

2. In Part 2 of the experiment, how was the molar solubility of calcium hydroxide impacted by the addition of calcium chloride? Explain.

3. A student performing the lab for the solubility of calcium hydroxide, titrated a 5.00 mL sample of saturated calcium hydroxide solution with a 0.15 M hydrochloric acid solution. The titration required 28 drops to reach the orange-red endpoint with methyl orange. What was the concentration of the hydroxide ion in the saturated calcium hydroxide solution? What was the Ksp that the student determined for that titration? (Consult the background section and lab calculations section). Assume 1 drop = 0.05 mL...