Laboratory Manual Introductory Experiments and Procedures

Lab # 17 The Iron Chemist

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Lab # 17 The Iron Chemist

Objective - To produce 0.750 grams of iron (III) oxide from iron (III) chloride hexahydrate.


FeCl3 + NH4OH  Fe(OH)3 + NH4Cl

Fe(OH)3  Fe2O3 + H2O
Materials: Iron (III) chloride hexahydrate, 250 mL beaker, 3.0 M HCl, 3.0 M NH4OH, distilled water, filter paper, funnel, clay triangle, crucible, crucible tongs, and Bunsen burner.

  1. Based on the reactions, determine the mass of iron chloride hexahydrate needed to produce 0.750 grams of iron (III) oxide.

  2. Dissolve the iron compound in a minimal amount of 3.0 M HCl.

  3. This is a very important step! Add 3.0 M NH4OH to the solution until it turns basic. Test with red litmus paper. Stir thoroughly and let the solution sit for a minute and test with litmus again. If it is not blue add several drops of 7 M NH4OH.

  4. Gently boil the solution for approximately five minutes.

  5. Fold a piece of filter paper and place in a funnel. Transfer all of the precipitate to the filter paper by washing with distilled water and filter. Make sure to transfer all of the precipitate to the filter paper. The effluent should be colorless and clear. If is not then go back to step 3.

  6. Mass a clean, dry crucible. Record the mass.

  7. Remove the filter paper from the funnel and gently squeeze out the water.

  8. Place the wet filter into the crucible and gently heat to char the paper. Once the paper has begun to burn, place the lid on the crucible and increase the heat. Oxygen can be added by gently and carefully blowing into the crucible.

  9. Once the paper has completely burned, heat strongly for 10 more minutes.

  10. Place the crucible on a wire gauze and allow to cool.

  11. Once cool, mass the crucible and its contents. Record the mass.

  12. Determine the mass of the iron (III) oxide by subtracting the mass of the crucible.



  1. Calculate the percent yield.

Experimental mass of Fe2O3

---------------------------------- x 100% = % yield

Theoretical mass of Fe2O3

  1. What are some possible sources of error in this lab? Be specific and describe how these errors would affect you results.

  1. What chemical reaction is happening in procedure step 10 as the contents of the crucible are cooling? Refer to your observations to help answer this question.


Write a short paragraph that addresses the following;

Did you meet your objective for the lab? How was this accomplished?

What types of reactions did you observe?

What mass of iron (III) oxide did you produce?

What was your percent yield? How can you improve your yield?

Lab # 18 Copper Cycle Lab


In this experiment you will take a sample of pure copper, follow it through a series of reactions, and recover the copper in the end. You will need to record your observations in detail as you attempt to identify the products as well as the type of reaction that occurred at each step. You will then determine the percent of the original copper that you recovered.


You must wear safety goggles at all times while in the laboratory. Several corrosive and dangerous chemicals will be used. Handle concentrated nitric acid (HNO3), 3M sodium hydroxide (NaOH), 3M sulfuric acid (H2SO4), and 3M hydrochloric acid (HCl) with care. If you spill any chemicals on your hands, wash them thoroughly. The gas produced when copper reacts with nitric acid is toxic – perform this reaction under the fume hood. Read the procedure for each section through completely before beginning that section!


Part A Disappearing Copper

1. Obtain a sheet of copper and mass it to the nearest milligram.

2. Place the copper in a clean 250 mL beaker. Label the beaker with your initials.

3. UNDER DIRECT SUPERVISION OF THE TEACHER, place the beaker in the fume hood and add 10 mL of concentrated nitric acid. DO NOT BREATHE THE FUMES! Leave the beaker in the hood until the reaction is complete.

4. Record all your observations.

Part B Basic Blue Goo

  1. Be aware that the contents of the beaker are still very acidic!

  2. Set the beaker in an ice bath.

  3. Very slowly, add about 20 mL of 3M NaOH to the beaker, stirring constantly.

  4. Use a glass stirring rod to place a small drop of the liquid from the beaker onto a piece of red litmus paper. (Do not place the litmus paper directly on the lab bench.) If the litmus paper does not turn dark blue, then continue to add NaOH several drops at a time until the litmus paper turns dark blue. Note: Some of the light blue precipitate may stick to the litmus paper. This does not mean that the litmus paper has turned blue. A positive result can be confirmed when the paper that absorbs the liquid has turned dark blue.

Part C Muddying the Water (Safety Note: Potential Explosion Hazard!)

  1. Set up a ring stand to hold the beaker over the Bunsen burner. Use the small iron ring to hold the wire gauze on which you will place the beaker. Use the large iron ring to place around the beaker to keep it from tipping over.

  2. Use a small, light blue flame to gently heat the contents of the beaker.

  3. Continue gentle heating until the reaction is complete. Aggressive heating will cause the contents to spatter out of the beaker in a violent manner. You should periodically remove the Bunsen burner from beneath the beaker to slow the heating process.

Part D Clearing Things Up

  1. Allow the product from part C to settle.

  2. Carefully decant and discard the clear portion down the drain with lots of water. Be careful not to lose any of the solid.

  3. Add 25 mL of 3M H2SO4.

  4. Stir until the reaction is complete. The change is very obvious.

Part E Completing the Cycle

  1. Add a few pieces of mossy zinc to the product of part D. Do not breathe the fumes given off during this reaction. The gas is a product of the reaction between the zinc and the excess sulfuric acid from part D.

  2. This reaction is complete when all of the blue color has been removed from the solution. If the reaction appears to stop before it is complete, your teacher will add some powdered zinc to the beaker to speed up the reaction.

  3. This reaction will produce a lot of heat which may convert the free copper back to an oxide, an undesirable result. Therefore, as the reaction nears completion, place the beaker in a cool water bath.

  4. After the reaction is complete, remove any large chunks of un-reacted zinc. Before you remove the zinc from the beaker, use a distilled water bottle to squirt off any copper that is sticking to the zinc back into the beaker.

Part F Mopping Up

  1. Add 20 mL of 3M HCl to the beaker. The purpose of the HCl is to use up the excess zinc that could not be removed by hand. The HCl reacts readily with zinc, but not with copper.

  2. When the reaction is complete (no more bubbles), decant the liquid carefully, so as not to lose any of the copper.

  3. Rinse the product with distilled water at least 3 times, decanting each time.

  4. Mass a piece of filter paper, fold it, and place it in a plastic funnel. Wet the filter paper to keep it in place. Place the funnel in a 250 mL Erlenmeyer flask.

  5. Add about 50 mL of distilled water to the copper in the beaker. Swirl the beaker to stir up the copper, and carefully pour it into the funnel.

  6. Add more water to the copper as needed to transfer it to the funnel.

  7. Use the rubber policeman to transfer any remaining copper to the filter paper.

  8. After the water has dripped through, carefully remove the filter paper, unfold it, place it on a paper towel, and allow it to dry overnight.

  9. Mass the dry copper and filter paper.

Data and Observations:
Table 1 - Mass Data:

Mass of copper metal (part A) ________

Mass of filter paper + recovered copper ________

Mass of filter paper ________

Mass of recovered copper (part F) ________

Percent recovery of impure copper ________

Part A

Balanced Equation




Chemical name

State of matter

(s, l, g, aq)


Molar mass




Type of reaction: _________________________ + _________________________

(combination of 2 types)

Calculations: Show a sample calculation for each type of calculation. Show all of your work, including units and significant digits.

Part B

Balanced Equation



Chemical name

State of matter

(s, l, g, aq)


Molar mass




Type of reaction: _______________________

Part C

Balanced Equation


Chemical name

State of matter

(s, l, g, aq)


Molar mass




Type of reaction: ________________________

Part D

Balanced Equation



Chemical name

State of matter


Molar mass




Type of reaction: ___________________________

Part E

Balanced Equation



Chemical name

State of matter


Molar mass




Type of reaction: ___________________________

Part F

Balanced Equation



Chemical name

State of matter


Type of reaction: ____________________________


1. Show a sample calculation for each of the different calculations that you had to perform. Be sure to include all of the proper units and report your answers to the proper number of significant digits.

2. Calculate the percent yield for the recovered copper.
% yield = (mass of recovered copper/mass of original copper) x 100

3. Evaluate the relative purity of your recovered copper and discuss at least three specific substances that are likely candidates that could be contaminating your recovered copper, and describe at which part of the lab these contaminants originated.

4. Create a detailed diagram showing each step of this experiment. Be sure to illustrate the cyclic nature of the entire process. This diagram should be drawn neatly on a separate sheet of paper / poster, and it should be done in color.
Grading Rubric: Mass Data (2) ____

Balanced equations (6) ____

Chemical names (3) ____

Calculations (15) ____

Observations (9) ____

Percent yield (2) ____

Purity evaluation (3) ____

Diagram* (10) ____

Total (50) ____

Lab # 19 Serial Dilutions
Introduction: Making accurate dilutions of an original stock solution is a very important laboratory skill. Many systematic errors that occur during the course of an experiment occur while making dilutions. Since the original solution is usually much more concentrated than the diluted solution you are making, any excess from the original solution will have a profound impact on the concentration of the new solution. Therefore, it is imperative that you work carefully and accurately. This is especially true when you are making “serial dilutions”. It is a good practice to use a new pipette each time you are making a new dilution. If this is not possible, then you should thoroughly rinse the pipette several times with distilled water. Also, note that if you dilute a solution in a volumetric flask above the graduation line, it is impossible to reverse that mistake by removing the excess solution, since some mixing has inevitably occurred.
Objectives: You will make a stock solution of known concentration*. You will then perform a series of dilutions of this solution as accurately as possible. You will be graded on technique and accuracy.


Read step 1 of the procedure, and calculate the mass of KMnO4 that you will need to weigh in order to create a 0.0150 M solution using a 50.00 mL volumetric flask. You must show the instructor your calculation before you may begin the lab.

Procedure: 1. Make a 0.01500 M solution of KMnO4. (Toxic! Do not touch!)

  1. You will first need to calculate the mass of KMnO4 that you will need to place in a total volume of 50.00 mL to create a 0.01500 M solution. Use a plastic weigh boat to mass this amount of solute.

  2. Refer to the figure at the bottom of p. 463 in your textbook, Holt, for the proper technique for making solutions.

  3. This first solution is referred to a the Stock solution.

  1. Make a ten-fold (1:10) dilution of stock solution.

  1. Transfer the stock solution to a clean and dry beaker.

  2. Use a 10.00 mL pipet to transfer 10.00 mL of the stock solution from the beaker into a 100.00 mL volumetric flask.

  3. Be certain that all of the solute is completely dissolved before proceeding to the next step.

  1. Make a 12.5-fold (1:12.5) dilution of this new solution.

a. Transfer the solution from the 100.00 mL volumetric flask to a clean and dry 250 mL beaker.

b. Use a 5.00 mL pipet to transfer 2.00 mL of the solution from the beaker into a 25.00 mL volumetric flask.

c. Be sure to shake the solution with the stopper on to ensure proper mixing.

  1. Place approximately 5 mL of this final dilution into a clean cuvette and bring this solution to your instructor.

  2. Record the absorbance from the Spec 20 of your final solution.


  1. Create a standard concentration curve of Absorbance vs. [KMnO4] from the standard data provided by your teacher. From your standard curve, determine the molarity of your final solution.

  2. Compare your actual molarity with the calculated molarity, and calculate the percent error of your dilution.

  3. Calculate the mass of KMnO4 that can be obtained upon the evaporation of the water from your final solution based on the concentration determined from the standard curve.

  4. Discuss the specific sources of possible errors that may have contributed to the inaccuracy of your results. Indicate whether these sources would have caused the measured concentration to be too high or too low.

[The accuracy of your solution will be determined spectrophotometrically.]

Grading: Pre-Lab calculation (2) ____ +/- 2% = 10/10

Lab technique and clean-up (6) ____ +/- 5% = 8/10

Standard Concentration curve (3) ____ +/- 8% = 7/10

Calculation of final molarity (2) ____ +/- 12%= 6/10

Calculation of mass of KMnO4 (2) ____ +/- 15%= 5/10

Accuracy of final solution (10) ____ +/- 20%= 4/10

Sources of error (5) ____ +/- 25%= 2/10

Total (30) ____ >25%=1/10

* When massing a solid solute to make a solution, you must determine if the solid is anhydrous, a hydrate or if it is hygroscopic. If it is a hydrate, you must take into account the additional mass from the water. If the solid is hygroscopic, such as sodium hydroxide, you can only make a solution of approximate concentration. To determine the concentration more precisely, you would need to standardize the solution by titrating it with a solution of known concentration. KMnO4 is anhydrous.

Lab #20 Acid – Base Properties of Common Substances


Predict A/B/N


Rank (1-16)

Red Litmus

Blue Litmus

pH paper

pH meter

conducts? Y/N/S




Baking soda




Distilled water


Orange juice

Rain water



Tap water



0.10M HCl

0.10M NaOH

Analysis Questions:
1. Which substance is the most acidic? ________________
2. Which substance is the most basic? ________________
3. Rank all 16 substances from most acidic (1) to most basic (16) to the left of the table.
4. What is the purpose of red litmus paper?
5. What is the purpose of blue litmus paper?
6. List at least 4 properties of acids.
7. List at least 4 properties of bases.
8. When you take the pH of something, what do you think it is actually measuring?
9. How is it possible for us to consume substances with such a wide range of pH’s?
10. What may account for the differences in properties between tap water, rain water, and distilled water?

Lab #21 Acid – Base Titration

Objective: To determine the concentration of an acid.
Pre-Lab: Look up definitions for the following terms: (See pp. 550-555, Holt)
Titration –
End point –
Equivalence point –
Titrant –
Standard solution –

Safety Precautions: Wear your safety goggles at all times while in the lab!!
Procedure: [See pp. 552-553 in your textbook (Holt)]
1. Fill each buret with the appropriate solution. Open the valve to allow a small amount of solution to flow out into a waste beaker. This will allow each solution to fill the tip of the buret. Be sure to accurately record the starting volume of both the acid and the base. Remember that a buret is calibrated to measure the volume of liquid that has been dispensed.

2. Place approximately 20 mL of HCl into a clean 125 mL Erlenmeyer flask. Be sure to record the precise starting and final volumes. Also record the precise concentration of the NaOH.

3. Place several drops of phenolphthalein into the flask with the HCl and swirl the flask.

4. Begin adding NaOH to the flask approximately 1 mL at a time. Note the rate of the color change. When the change starts to take longer to return to its original color, begin adding base at a slower rate. Your goal is to add one drop of base that just causes the solution to form a faint pink color. The color should persist for at least one minute.

5. If you add too much base and the solution turns a dark pink color, you over shot the equivalence point. Therefore, you will have to add more acid to return to the equivalence point.

6. Record the final volumes of both the acid and the base added.

7. Repeat this titration two more times if time permits.

8. You may dispose of these solutions down the drain with lots of water.


Reported concentration of NaOH __________ (M)

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