Reaction 1: Ca (s) + 2 H2O(l) → Ca(OH)2 (s) + H2 (g)
Reaction 2: acid (aq) + calcium hydroxide (aq) → salt2 (aq) + water (l)
1. Write the balanced chemical equation for Reaction 2 using your assigned acid.
Materials:
- 1.0 g calcium metal
- hydrochloric acid (titrant)
- methyl orange (indicator)
- 50.00 mL glass burette
- plastic burette funnel
- retort stand
- burette clamp
- 3 x 250 mL Erlenmeyer flasks
- 2 x 1000 mL Erlenmeyer flask
- 150 mL beaker (for acid)
- 1000 mL beaker (for waste)
- large glass funnel
- Whatman #1 filter paper
- electronic balance
- 100 mL graduated cylinder
- burette brush
- wash bottle
Procedure:
Preparation of Glassware
1. If glassware is cloudy, rinse with 1.0 mol/L hydrochloric acid to remove scale.
2. Rinse with tap water.
3. Scrub the glassware with soapy water and a brush.
4. Triple rinse the glassware using tap water.
5. Rinse once using deionized (DI) water.
6. Dry the 150 mL beaker (for titrant) and 1000 mL Erlenmeyer flask (for filtrate).
Preparation of the Buret
1. Inspect the burette. If it is cloudy, close the stopcock and add 10 mL of DI water. Scrub the burette using a burette brush. Pour the water into the waste beaker.
2. Rinse the burette with DI water using a wash bottle. Open the stopcock and make sure it drains.
3. Pour about 100 mL of the titrant (acid) solution into the dry 150 mL beaker.
4. Close the stopcock and add 10 mL of the titrant solution (HCl) to the burette. Roll the burette, making sure to wet all sides. Pour the acid solution into the waste beaker.
5. Fill the burette just past the 0.00 mL line. Place the waste beaker below and open the stopcock. Allow enough titrant to pass through to clear all bubbles.
6. Record the actual burette reading in Table to 2 decimal places (x.xx mL). Do not calculate & record the volume that is in the burette! This is a common mistake.
Sample Preparation
1. Weigh approximately 1.0 g of calcium metal on a piece of paper.
2. Place about 500 mL of DI water in a clean 1000 mL Erlenmeyer flask.
3. Add the calcium metal and allow the reaction to occur.
4. Using a glass funnel and filter paper, filter the saturated solution into a clean 1000 mL Erlenmeyer flask.
5. Retain the filtrate (the clear solution) and discard the undissolved calcium hydroxide and filter paper.
Titration
1. Using a graduated cylinder, measure 100.0 mL of the filtrate (saturated calcium hydroxide solution) and place into a 250 mL Erlenmeyer flask. Repeat for the remaining two flasks.
2. Add 5 drops of the indicator to each flask and swirl it to mix evenly.
3. Open the stopcock and add the hydrochloric acid solution while swirling the flask. Add the titrant slowly as a yellow color begins to appear. Stop as soon as a single drop produces a yellow color change that remains for at least 10 seconds.
4. Record the final buret reading.
5. Calculate and record the volume of acid added (Vadded = Vfinal - Vinitial).
6. Repeat steps 6 to 8 at least two more times or until a reproducible result is obtained (± 0.20 mL). The final burette reading becomes the next initial burette reading. Add more titrant if there is not enough to carry out the next trial.
7. Empty the buret and rinse it with DI water. Discard the titration waste and excess acid in the acid waste bin. Discard remaining filtrate in the caustic (basic) waste bin.
8. Store the buret in an open, inverted position on the clamp. Rinse all other glassware thoroughly with tap water and return the equipment to the appropriate cupboard or drawer.
Observations:
Table 1: Titration Data
Trial Initial Burette Reading (mL) Final Burette Reading (mL) Volume of Titrant Added (mL)
1.
2.
3.
Average:
Note: this simulation only records titrant volumes to one decimal place.
1. Calculate the moles of acid required to neutralize the saturated calcium hydroxide solution. Using the stoichiometry of Reaction 2, calculate the moles of calcium hydroxide in the sample. Calculate the concentration of the saturated calcium hydroxide solution sample.
2. a) Write the equilibrium dissociation equation for calcium hydroxide. Using stoichiometry, calculate the concentration of calcium ions and hydroxide ions that are present in the saturated Ca(OH)2 solution.
3. Write the solubility product (Ksp) equilibrium equation for Ca(OH)2. The concentrations determined in questions 4b) are the equilibrium concentrations of OH- and Ca2+ in the saturated solution. Using these values, calculate the experimental Ksp value.
4. Calculate the pH of the saturated calcium hydroxide solution. One published value for the Ksp of Ca(OH)2 is 5.02 x 10-6 at 25oC. Using this Ksp value, calculate the theoretical calcium ion concentration present in a saturated solution.
5. Calculate the percentage error between the experimental and theoretical calcium ion values.
Write a conclusion to restate the purpose and state the results.
During filtration, tiny particles of solid Ca(OH)2 can pass through the pores in the filter paper. Will these particles affect the initial concentration of ions in the sample (filtrate)? Briefly explain.
a) If solid particles were present in the filtrate, what would happen to these solid particles during the titration? Explain using Le Chatelier’s Principle and the equilibrium from question.
b) Do the experimental results suggest that this may have occurred? Explain.