Second Semester Exam Review Lab-Based Problems

Chemistry                                                                                                                                                     June 2005

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 6: 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 5b: Thermodynamics “Less Than Zero”

A group of boys face the challenge of getting the temperature of a reaction of baking soda, NaHCO3, and HCl to drop below 0oC from a starting temperature of 20.5oC.  They need to use the following information:

The reaction that occurs is:

NaHCO3 (s)  +   HCl (aq) à NaCl (aq)  +   H2O (l)  +   CO2 (g)        DH = +28 kJ/mol

Remember: (heat lost by the acid solution) q = m x (DT) x Cp  [Cp = 1.0 cal/g-oC]
Also 4.184 J = 1.00 calorie.

Task 1:  Answer these questions:

a)      Is the forward reaction endo- or exothermic?

b)      Is the DS for the forward reaction positive or negative?  Justify your choice. 

 c)      Is this reaction spontaneous?  How can you tell in the lab?  How can you tell given the thermodynamic quantities of DH, DS and DG? 

Task 2: If the students start by reacting 4.00 grams of baking soda with 25 mL of 2.0 M HCl in a Styrofoam cup and they find the temperature drops to 11.3oC, calculate the following:

a)      The mole quantities of the baking soda and HCl. 

b)      The theoretical heat absorbed by the chemicals. 

c)      The actual heat lost by the water/acid solution. 

d)      The percent efficiency of the heat lost by the solution to the heat absorbed by the chemicals. 

Task 3:  Assuming the same efficiency as above, determine a recipe to get the solution below zero.  Lay out your calculations with annotations.

 

 

Unit 7: Solutions, Equilibrium and Solubility, “Finding [Mg2+] in sea water”

A student is given a 250 mL solution of seawater to analyze.  This seawater contains many ions, but predominantly chloride, sodium, sulfate, potassium, magnesium and calcium.  The task of the student is to determine the concentration of the magnesium ion in the seawater in terms of molarity.  The student needs to know that this sample of seawater is from an area with approximately 2.54 x 109 kg of Mg2+ per km3 of seawater.  The quantities of the other ions is as follows: Cl- = 1.95 x 1010, Na+ = 1.09 x 1010, Ca2+ = 1.95 x 109,  SO42- = 9.1 x 108,  K+ = 3.9 x 108.  (All quantities are in kg/km3.)

Task 1: In detail discuss how you will go about separating the magnesium ions from the rest of the ions in the sea water.  Use sentences and chemical equations as appropriate. 

Task 2: Convert the given concentration of Magnesium ions into molarity. 

Task 3: Discuss specifically how you dealt with the presence of Ca ions not precipitating with the anion that you added to precipitate with Magnesium. 

Task 4:  Calculate the amount of an anhydrous sodium anion salt that you would need to fully precipitate the Magnesium ion. 

Task 5:  Calculate the mass of the magnesium precipitate that would form if you had 100% accuracy. 

Task 6:  A student used an excess of NaOH to the sample of calcium free sea water in order to precipitate the magnesium ions.  If the student obtained 6.01 g filter paper and Mg(OH)2 after thoroughly rinsing and drying the filter paper with the precipitate on it.  (The filter paper before the experiment weighed 2.54 g.)  What is the [Mg2+] from this experimental data?  What is the student’s percent error?  Propose two sources of error. 

Task 7:  A second student adds a double excess of KF-2H2O (first dissolved in water) directly to the 250 mL of sea water.  When he analyzes the precipitate he ends up with 85% error over the predicted molarity of Mg2+.  What suggestion would you give to the student for improving his results? 

 

 

Unit 8: 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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