Experiment 18
Determination of an Equilibrium Constant: Solubility of Ca(OH)2
GOALS:
This experiment should help you learn more about equilibrium concepts and how these can be applied to
solubility. In addition, you will gain more experience with quantitative measurements and the use of the
Atomic Absorbance Spectrophotometer.
INTRODUCTION:
Calcium hydroxide is slightly soluble in water. This dissolution process can be represented by an
equilibrium equation, Eqn.1.
Ca(OH)2(s) ↔ Ca2+(aq) + 2 OH1-(aq)
Eqn. 1
The equilibrium constant for this equation, called Ksp or the solubility product, is then
Ksp = [Ca2+] [OH1-]2
Eqn. 2
We will measure [Ca2+] and [OH1-] in solutions at equilibrium and then use these values to calculate the
value of Ksp, the equilibrium constant. We will do this first in a solution at room temperature. The values
for equilibrium constants that you find in reference books are usually measured at 25 °C, roughly room
temperature. Later, you will prepare an equilibrium solution at a temperature near 100 °C and measure
the new value of the equilibrium constant. In general, the value of an equilibrium constant changes when
temperature changes. Recall, however, that varying ion concentrations or amount of solid used, will not
affect the value of the equilibrium constant. Unlike most solids, calcium hydroxide actually becomes less
soluble as temperature increases.
We will use Atomic Absorption, AA, to determine the [Ca2+]. Since concentration is proportional to
absorption, we will use the absorptions of solutions with known concentrations to prepare a calibration
curve, and then use that curve to find [Ca2+] in our solutions. We will use an acid/base titration to find
[OH1-] via stoichiometric calculations. Once we have [Ca2+] and [OH1-] at a given temperature, we will
use Eqn. 2 to find the value of the equilibrium constant.
PRE-LAB ASSIGNMENT:
In lab you will prepare 250.0 mL of 0.0100 M HCl from a 1.00 M HCl solution. Calculate the volume
that you will need. Show this calculation in your pre-lab notebook entries. See section 4.4 of your
textbook for help.
HAZARDS:
All the solutions used in this experiment may go down the drain since they are dilute acids and bases and
contain no hazardous metal ions. Look up the MSDSs for calcium hydroxide and hydrochloric acid
(http://hazard.com/msds/index.php, or Trexler 464). In your pre-lab notebook entry, record summaries of
the potential health effects for each (Hazard Identification section). Now, consider how this information
should be applied in this experiment. In particular, we are using small volumes of 0.01 M HCl as
compared to concentrated HCl, which is 12 M. Discuss how this affects the potential hazard of our
experiment.
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LABORATORY OBSERVATIONS AND DATA:
As always, be sure that your notebook entries include what you do, what you observe, and clearly labeled
numerical data with units and appropriate sig figs.
You will need two different tables of titration data: one for the Part I room temperature data and one for
the Part II high temperature data. Each of these tables should include the measured solution temperature
and exact concentration of the HCl as read from the bottle. Each table needs columns for three trials and
rows labeled: volume of Ca(OH)2 solution, initial buret reading, and final buret reading. Include
appropriate units.
You will also need a table for the AA data. This single table will need 6 rows labeled: distilled water,
standard 1, standard 2, standard 3, room temperature solution, and hot solution. The columns should be
labeled [Ca2+], and absorbance.
PREPARATION OF HCl SOLUTION:
1. Confirm with your instructor that you have correctly calculated the volume of 1.00 M HCl needed
to prepare 250.0 mL of 0.0100 M HCl
2. Use a pipet, 250 mL volumetric flask, and some parafilm to prepare your solution. Mix well.
PROCEDURE FOR PART I: Measuring the value of Ksp at room temperature
1. Set up a buret filled with the 0.0100 M HCl you just prepared. Recall that you must use a small
amount of the solution to rinse the buret before filling. Be sure that solution fills the volume
below the stopcock without any bubbles.
2. Get 50-60 mL of saturated aqueous Ca(OH)2 solution. Filter this solution to remove any
suspended solid using a clean, dry funnel and filter paper. [Note that we normally wet the filter
paper to help it stick to the funnel. Here we don’t want any added water to dilute our solution.]
Record the temperature of this solution. Collect the filtrate in a clean, dry beaker or flask.
3. Rinse your 1 mL pipet with some of the clear solution. Pipet 1 mL of the solution into a 100 mL
volumetric flask. Carefully dilute to the mark with distilled water. Seal the flask with a piece of
parafilm and then invert several times to mix. Label the flask with your initials and “RT” for
room temperature. Save this flask for AA analysis later (Step 13).
4. Rinse your 10 mL pipet with some of the filtered solution. Pipet 10 mL of this solution into a 125
mL Erlenmeyer flask. Add three drops of bromothymol blue indicator. Your solution should be
a medium blue.
5. Record the initial buret reading (± 0.01 mL) and then titrate the calcium hydroxide solution until
the last green color just disappears and the solution is yellow. Record your final buret reading.
Discard the titrated solution down the drain with excess water. Be sure that you are recording
observations as you do your titrations in addition to your numerical data.
6. Repeat the titration Steps 4 and 5 twice more. You will need to refill the buret with HCl.
7. Once you have all three titrations done, you may discard any remaining Ca(OH)2 solution, but
keep the HCl in the buret for Part II.
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PROCEDURE FOR PART II: Measuring the value of Ksp at high temperature
8. To prepare a high temperature saturated solution of Ca(OH)2, use a Bunsen burner to bring 100
mL of distilled water to a boil in a 250 mL beaker. After the water has been boiling for several
minutes, add about 0.2 g of solid Ca(OH)2. Gently boil the solution for an additional 3 minutes,
stirring occasionally. The mixture should remain cloudy.
9. As you did in Step 2, this solution must be filtered to remove any suspended solid. So that we
can filter it before it has a chance to cool, we will use an unusually large funnel and filter paper.
Get a 10 cm funnel from the supply bench and fit it with a dry piece of 18 cm filter paper. While
your Ca(OH)2 mixture continues to boil, place a clean dry beaker (at least 100 mL size) beneath
the funnel to receive the solution.
10. Borrow one of the insulated gloves and a thermometer from the supply bench. Turn off the
Bunsen burner, record the solution temperature, and immediately pour 50-60 mL of the hot
solution through the filter. We need to filter the solution before its temperature can change
significantly. Solution that cannot be filtered quickly will be discarded because, as its
temperature changes, so does the Ca(OH)2 solubility and the value of Ksp. Once the solution has
passed through the filter, however, you may allow it to cool to room temperature without
affecting the experimental results. Return the gloves and thermometer to the supply bench so that
others may use them.
11. Rinse your 1 mL pipet with the hot filtered solution. Pipet 1 mL of this solution into a clean 100
mL volumetric flask. Dilute to the mark with distilled water, seal with parafilm, and invert
several times to mix. Label the flask with your initials and “hot.” You may take your two
volumetric flasks to the AA either now or later as directed by your instructor (see Step 13).
12. Titrate 10 mL aliquots of your filtered solution following Steps 4 and 5 as before. Be sure to
rinse the 10 mL pipet with some of the hot solution before measuring out samples to titrate. Get 3
trials. Watch the level of HCl in your buret and refill as needed. Be sure that you have initial and
final buret readings for each trial.
CALCIUM CONCENTRATION VIA AA:
13. Take your two volumetric flasks to the AA room.
14. As directed by your instructor, record the [Ca2+] concentration and absorbance of the distilled
water and three standard solutions that your instructor has prepared.
15. Now record the absorbance of each of the solutions in your volumetric flasks. Record
observations made at the AA and your numerical data. (Note that you won’t have values for the
concentrations of the RT or hot solutions. These values are calculated later.)
CLEAN UP:
16. Add the HCl remaining in your buret to the residue of solid Ca(OH)2 remaining in the beaker you
boiled. Most of the solid should dissolve. Rinse this down the drain with excess water.
17. All of your remaining solutions can go down the drain with excess water.
18. Rinse your buret with distilled water.
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RESULTS:
Prepare a spreadsheet in Excel for your AA data. Prepare a well labeled table that includes your
concentration and absorbance data for distilled water and standards 1-3 (but not your RT and hot
solutions). Graph this data using an XY scatter plot. Be sure that absorbance goes on the y-axis and
concentration (molarity) goes on the x-axis. Add a linear trendline and the equation for this line. See Exp
13 for a reminder of how to do this. Be sure that your graph includes properly labeled axes and a title.
Electronically copy/paste your graph into your lab report. Save your Excel file with the experiment
number and your initials in its file name. Email this file to your instructor as part of your report.
In addition to the graph, include in your lab report a Table of Calculated Results for Part I, and a Table of
Calculated Results for Part II. You do NOT need to do the calculations in Excel unless it is your
preference. You may do the calculations by hand if you prefer. Each table should look like the sample
below with temperature (°C), [HCl] (M), and then 3 columns with rows labeled: volume of Ca(OH)2
solution (mL), volume of HCl added (mL), moles of HCl, moles of OH1-, molarity of OH1-, average
molarity of OH1-,[Ca2+] in AA soln, [Ca2+] in equil soln, and Ksp.
Entries for the first several rows (temperature, [HCl] and volume
of Ca(OH)2) are just data you collected in lab. Calculate the
volume of HCl used from the initial and final buret reading for
each trial. The moles of HCl is calculated from the volume and
molarity of the HCl. Moles of OH1- is equal to moles HCl since
the titration is a one-to-one mole ratio reaction. [OH1-] for each
trial is calculated from the moles of OH1- and volume of
Ca(OH)2 used in that trial. Now average the molarities
calculated from the three columns to give one average [OH1-]
value.
Temperature, oC
[HCl], M
Trial
1 2 3
Volume Ca(OH)2, mL
Volume HCl, mL
Moles HCl
Moles OH1[OH1-], M
The [Ca2+] in the AA soln is calculated from the calibration
Average [OH1-], M
curve you prepared in Excel. For your Part I Results table,
[Ca2+] in AA soln, M
substitute the AA absorbance of your RT solution into the
equation for your calibration line and solve for concentration.
[Ca2+] in equilibrium soln, M
(Recall that absorbance substitutes for y, and then you solve for
Ksp
x.) To compensate for the dilution you did with the volumetric
flask, multiply the molarity of Ca2+ in the AA soln by 100 to get
the molarity of Ca2+ in the equilibrium solution. Finally, calculate the value of Ksp from the average
molarity of OH1- and the molarity of Ca2+ in the equilibrium soln using Eqn 2. For your Part II Results
table, you follow the same logic, but using the data collected at the high temperature. Be sure that you
have included units and correct significant figures in both of your tables.
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DISCUSSION:
In a couple of paragraphs, explain your understanding of the equilibrium process. (The Introduction to
this experiment and textbook sections 16.1, 16.2, and 16.5 will be helpful to you.)
 Make general statements about what is happening to the molecules or ions in an equilibrium
solution. Discuss both the forward and reverse reactions.
 Refer to specifics of the Ca(OH)2 system (Eqn 1) to illustrate your general points.
 Continue your discussion to include how the concentrations or amounts of the molecules and ions
relate to each other.
 Describe what does and does not affect the value of an equilibrium constant. Again, illustrate
your general points by referring to specifics of the Ca(OH)2 system you studied.
 Finally describe how the value of Ksp is related to the solubility of a salt. That is, does a larger
value of Ksp indicate a more or less soluble salt? Again, use the specifics of the Ca(OH)2 system
and Eqns 1 and 2 to illustrate your points and clearly explain why Ksp and solubility are related in
this way.
QUESTIONS:
1. In some equilibria, the value of K increases as temperature increases. In other equilibria, the value of
K decreases as temperature increases. Explain how the endothermic or exothermic nature of a reaction
determines which relationship we will see. {See your textbook section on LeChatelier’s Principle for
help.} Is our Ca(OH)2 equilibrium endothermic or exothermic as written in Eqn 1? Explain.
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