Second Semester Exam Review Lab-Based Problems

Chemistry                                                                                                                                                     May 2008

The following questions are based on the key concepts of the various second semester units.  They are focused around the key labs of each unit.

Unit 7: States of Matter “Getting the Volume Right”

A student is faced with the challenge of filling a 1.00 L graduated cylinder with CO2 gas.  The student will generate the gas by a reaction between solid, anhydrous sodium carbonate and a 1.0 M HCl solution.  The student needs to know that on the day she is performing the experiment, the temperature is 22.0oC and the atmospheric pressure is 78.5 cm Hg.  She also finds out from her teacher that CO2 is slightly soluble in water, meaning that for every 100 mL of reaction solution volume that she uses she will lose 88 mL of CO2 gas.

Task 1: Determine how much HCl and sodium carbonate she should use to produce 1.00 L of gas.  Clearly lay out the chemical equations and calculations.  Explain with a sentence or two what a set of calculations are determining.

Task 2: Simply and clearly describe the procedural steps that the student needs to take to accomplish this task in the lab.  Point out any particular pitfalls that she should avoid.

Unit 6: Stoichiometry – Determining the formula of a New Salt

 A student is analyzing a hydrated salt, containing Cr, Cl and H2O.  She must determine the empirical formula of the salt including the moles of water to salt ratio.  To do so she gets a clean, dry, crucible and weighs it.  She then puts about two grams of the salt into the crucible, weighs it again.  She meticulously heats the salt noting a complete change in color as the heating progresses.  After she is satisfied that she has thoroughly dehydrated the salt, she lets the crucible cool and weighs it.  She then dissolves the salt in a small amount of water and adds a large piece of Aluminum wire.  She allows the wire to completely react and notes that a new metal has formed on the bottom of the beaker and the aluminum wire has shrunk significantly.  She separates the new metal and filters it.  She rinses and allows the filtered metal to dry.  She records the following data during the course of the lab:

             Mass of empty crucible:                                  12.96 g

            Mass of crucible and hydrated salt:                14.97 g

            Mass of crucible and salt after heating:          14.60 g

            Mass of filter paper:                                          1.27 g

            Mass of filter paper and metal:                         1.81 g

Clearly and thoroughly lay out all calculations and steps for your solution to finding the empirical formula. 

Answer the following questions.

1.   Write the chemical equation for the reaction with aluminum.

2.   What effect on the final complete empirical formula would not driving off all the water have had?

 

 Unit 10: Acids and Bases: Turn it Pink!  Acid Solution Stoichiometry

Overview: In this experiment, you will be trying to get the proportions of Ca and HCl just right so that when the reaction is complete the resulting solution’s pH will be a neutral 7 (or above 7 with a small amount of Ca added).  If you have too little Ca or too much acid, the solution will remain clear and colorless.  But when you add too much calcium the solution will turn pink.  You want to get it right the first time, but there are more factors involved than just doing the math and adding the right amounts.  In the end you will not only practice your knowledge of solutions and stoichiometry, but you will also consider some of these other “factors.”

Procedure: 

  1. Write a balanced chemical equation for the reaction of Ca with HCl.
  2. Obtain between 50mL and 100 mL of HCl of known concentration using a graduated cylinder.
  3. Place the HCl solution in an appropriately sized beaker.
  4. Calculate what mass of Ca is required to react 100% with the number of moles of HCl you have in your beaker.
  5. Weigh out that mass of Ca and record this mass.
  6. Add two drops of phenolphthalein indicator to the acid.
  7. Add the Ca metal to the acid.  Observe and gently stir.
  8. When the reaction is finished, measure the pH with a piece of universal pH paper.  Record this pH.
  9. If the solution is still colorless, weigh a very small piece of Ca and add that to the acid.  Stir. (Try to find a piece of Calcium that has a mass below 0.05 g.)
  10. Again when the reaction is complete, measure the pH with a piece of universal pH paper.  Record the pH.
  11. Repeatedly add a small amount of calcium (0.03-0.05 g) to the acid until the solution turns pink.  Then measure the pH of this solution and stop.
  12. Clean up by pouring the solution down the drain.

Questions:

  1. Add up the total mass of calcium you needed just before the solution turned pink.  What percentage of the total calcium mass is the “extra” you added to turn the solution pink? 
  2. Give two reasons why your experimental mass of calcium added didn’t agree with your theoretical mass of calcium calculated. 
  3. The reaction that caused the acid solution to turn pink is a different reaction than the equation that you wrote above in procedural step 1.  Write the equation of the reaction with calcium and whatever else to show this reaction.  In a couple of sentences describe what is happening when the solution turns pink. 

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