creact1 Mr. Charles Gimbel
Holy Trinity Episcopal Academy Chemistry I
Here's the whole lab...


Investigating Typical Chemical Reactions

Teacher Authored Experimentation

by Charles D. Gimbel


I N T R O D U C T I O N

Many scientific disciplines rely upon indirect evidence.  Chemists often rely on indirect evidence for atomic, ionic, and molecular activity.  We can study concepts, like CHEMICAL REACTIONS and rely upon some concrete evidence that reactions have occurred.  In many instances the evidence is the production of heat, effervescence, formation of a precipitate, production of smoke, release of light, production of water, or change in color.  Other evidences are less obvious and may require instruments to detect.  Numbers in parentheses, like (4) refer to section in CHEMICAL REACTIONS.


P R O C E D U R E

Part A.

The purpose of this part is to help the student become familiar with chemical indicators and reagents that will provide some evidence that a specific reaction has occurred.

According to the Arrhenius theory, an acid is a compound that makes hydrogen ions available in solutions and a base is a compound that makes hydroxide ions available in solutions.  Since one hydrogen could "connect" to one hydroxide and make HOH (another way to write water), an acid and a base can neutralize one another (4) to form a salt and water.

Some chemicals can change color when they come in contact with an acid or a base.  Such compounds are known as indicators.  We will use a few indicators that change color in the range that will be useful to us.  Another specific reagent we will use is limewater.  Limewater is made from lime, which is calcium oxide, and water.  Limewater is actually calcium hydroxide because a base is formed when some metal oxides plus water produce a base  (2c).  When carbon dioxide is bubbled through limewater it turns milky.  Bases plus carbon dioxide form carbonates.  Calcium carbonate is insoluble and the small granules of calcium carbonate that form create a milky appearance.  The use of limewater is the chemist's specific test for the presence of carbon dioxide.

1.  Place ten drops of water in the eight wells of a microchem plate in rows 1 & 6.

2.  Place a drop of 0.1M HCl in each of the four wells of row 1.  (The designation 0.1M refers to the concentration of the solution.  You'll learn about solution concentrations later.)

3.  Place a drop of 0.1M NaOH in each of the four wells of row 6.

4.  Place 2 drops of bromothymol blue (BThB) in wells 1A & 6A.  Place 2 drops of phenol red (PhRed) in wells 1B & 6B.  Place 1 drop of phenolphthalein (PhTh) in wells 1C & 6C.  Place 2 drops of universal indicator (UI) in wells 1D & 6D.  This procedure is to help you become familiar with the colors of four indicators in dilute acid and dilute base, so record the results as you may need this information in the coming ten labs.  Before you clean the plate, add one more drop of acid and base to the respective wells with universal indicator.

5.  Place about 3 mL (estimate) of limewater in a 13 x 100 test tube.  Make a splash guard by cutting a square from a 3 X 5 card and punch a hole in the middle of the square.  Gently, blow your breath into the tube using a small straw or tube provided for this purpose.  The square should be small, but large enough to protect against having the liquid splash you in the face.  Limewater could damage your eyes.  The tube should just barely be below the surface of the limewater - don't stick it way down in the liquid.  Use a dark background for the tube and note the cloudy appearance of the limewater.  If you blow too long the cloudy appearance will come and leave without seeing that it happened!  This disappearing phenomenon should be of interest to you from a chemical standpoint.  Refer to (7d).  The cloudy appearance, as explained in the introduction, is due to the presence of calcium carbonate.  The calcium carbonate will react with an acid as described in (7d).  Where did the acid come from?  As you blow  into the limewater, the carbon dioxide (a nonmetal oxide) dissolves in water.  See (3b).  Your breath plus the water creates carbonic acid which reacts with the calcium carbonate that was formed by the limewater and carbon dioxide in your breath.  Pause on this point until you understand the chemistry of this statement.

6.  Place 3 mL of distilled water in each of two test tubes.  Add three drops of phenol red indicator to each tube.  Add one drop of 0.1M HCl to one tube, and one drop of 0.1M NaOH to the other tube.  Make your observations.

To these same tubes add two drops of the acid to the tube that contains the base, and two drops of the base to the tube that contains the acid.  When these opposite acid or base are added they neutralize one another (4).  The first drop neutralized the solution, and the second drop reversed the solution.

7.  Repeat step #6 using bromothymol blue indicator.

8.  Repeat step #6 using bromocresol green "Hot Lips" indicator.

9.  You may explore with other indicators if your instructor approves of this and has provided them for your use.  You should ask before you proceed as factors such as available solution, time, and experimental value can only be assessed by your teacher.


Part B.

Metals plus sulfur yield metal sulfides.  The sulfides will react with an acid solution to produce hydrogen sulfide and a salt of the acid.  (1a & 1b)

1.  Place a small amount (your teacher will demo how much) of sulfur in a special 13 X 100 test tube provided for this purpose.

2.  Place a galvanized nail, head down, in the test tube.  Using a test tube holder, heat the bottom of the test tube strongly, in the fume hood, until it appears that the sulfur has vaporized.  Put the hot tube on a fiber-ceramic hot pad until it cools.  Leave the test tube holder on it.

3.  Place the nail in a clean tube, head down. Add about ten drops of 3M HCl to the solid.  Waft the vapors from the effervescence toward you and politely describe the fragrance.  Stop the reaction within one minute by adding some tap water to the tube.  Then pour the solution and the nail into an evaporating dish.  Add a small scoop (about the size of a pea) of sodium bicarbonate to the evaporating dish and observe what happens.  Place the contents of the evaporating dish in  the DESIGNATED CONTAINER.


Part C.

Metal oxides may be reduced by thermal decomposition in the presence of a reducing agent. (2b)  You will be using methyl alcohol which is toxic and flammable.  Keep the tube containing the methyl alcohol at least two feet from the burner.

1.  Make a short spiral of copper wire with a handle adequate to hold the copper in the flame of a laboratory burner; the spiral must be able to fit in a 13 x 100 test tube.  Imbed the loose end of the wire in a cork or use a wire holder handle so that the heat will not burn your hand.

2.  Warm a test tube in a small low temperature flame for no more than ten seconds.  The tube must not be hot.   Add fifteen drops of methyl alcohol to the warm test tube.

3.  Hold the copper spiral in a quiet flame of the burner and note the black color caused by the copper reacting with the oxygen in the air.

4.  While the spiral is hot, but not glowing, place it in the vapor portion of the test tube and note the change in the appearance of the copper wire.  Sometimes one has to plunge the hot wire into the alcohol to produce concentrated vapors in the upper part of the tube and repeat the initial part of this step.  Keep the tube away from the flame.  Note the changes in the appearance of the copper.

5.  Dispose of the alcohol in the specially designated container.


Part D.

Metals plus oxygen produce metal oxides which, when added to water act as basic anhydrides and produce a base (hydroxide) of the metal. (2a & 2c).  CAUTION:  Do not look directly at burning magnesium.

1.  Place a watch glass on the black counter top near the laboratory burner.  Use a pair of forceps to hold a the tip of a 3 cm piece of magnesium ribbon horizontally in the burner flame, just until it ignites.  Do not burn the forceps. CAUTION: Do not look directly at the flame; it can be viewed adequately with peripheral vision ("out of the corner of your eye - not sure? ask!).  If you look directly at the flame, you will be cited for a safety violation and your teacher will dismiss you from lab with appropriate penalties.

2.  Place the ash in the watch glass and note its appearance.  Place a piece of white paper under the watch glass and add about 1 mL of distilled water.  Slosh the liquid on the watch glass to assure good mixing, and stir or crush the ash with a wood splint if needed.

3.  Place a piece of red, and a piece of blue litmus paper at the edge of the liquid and interpret any color changes.

4.  Wipe out the watch glass with a paper towel and dispose of the towel in the trash can.  Rinse and dry the watch glass.

Part E.

Some nonmetals plus oxygen produce nonmetal oxides which, when added to water act as acid anhydrides and produce an acid of the nonmetal. (3a & 3c)

In this step, you will be burning sulfur.  The products is sulfur dioxide.  Sulfur dioxide is toxic, irritating, and some people have allergic reactions to it.  If you are known to have allergic reactions to salad bars in restaurants due to a bisulfite reaction you'd better not get involved in this part of the experiment.  To conduct this step, we will use an environmental control known as a fume hood.  

1.Place a small amount of sulfur in a deflagrating spoon.

2.Using the fume hood with the fan turned on, hold the spoon in the burner flame until the sulfur ignites.  You need to hold it in the flame for only about five to ten seconds, then remove it from the flame to see if it is burning.  The flame is a light blue color, and is almost invisible.  If it is burning, you will see smoke.  If it isn't burning, put it back in the flame again.

3.Hold the deflagrating spoon in a 250 mL Erlenmeyer flask so the smoke will fill the flask.  After the flask has filled with smoke, remove the deflagrating spoon and place it in the flower pot.

4. Immediately, pour approximately 8 mL of distilled water into the flask, insert a tight fitting stopper, and shake the flask outside the fume hood for approximately two minutes.

5.  Return to the fume hood to conduct the remainder of this step.  Pour the liquid from the flask to a small plastic cup.  Fill the flask completely with tap water, pour the liquid into the fume hood sink, and rinse the stopper by dangling it in the large beaker of water.  After the rinsing, put the flask and stopper in the drying area.  Return to your lab area.

6.  Place a watch glass on a piece of white paper.

7.  Pour about 1 mL of the solution into the watch glass and test it with both red and blue litmus papers.

8.  Pour about 3 mL of the solution into each of two test tubes.  Add a drop of phenol red to one tube and 1 drop of bromothymol blue to the other test tube.

9.  If enough solution is available, you may wish to try other available indicators.


Part F.

When an acid and a base react, WATER and  SALT are always the products.  (4)

1.  Obtain a watch glass and ensure that it is clean and dry.  The black counter top or a piece of black paper will make an excellent background for making observations.

2.  Place six drops of 1M HCl in the center of the watch glass.  Add five drops of 1M NaOH to the drops of HCl.

3.  Place the watch glass toward the back of the table top for the duration of the period.

4.  At the end of the lab period, place the watch glass in your drawer in such a manner that it will allow the liquid to evaporate, but not be sloshed and spill.

5.  At the beginning of the next lab period, open your drawer very carefully in case the liquid has not evaporated.  If the liquid has not evaporated, go back to step 3.

6.  If the liquid has evaporated completely and the presence of a solid is evident, observe the solid with a hand lens or a dissecting microscope.  Draw the crystalline structure that you can see.


Part G.

Thermal decomposition of chlorates, bromates, or iodates yield a binary salt and oxygen. (6)

There are risks associated with this step that preclude the use of these compounds consistent with the school district's adopted science safety manual.  This chemical reaction will be illustrated by use of a demonstration viewed on the laser disk, or it may have to be omitted.

Part H.

Many carbonates may be thermally decomposed to produce their metal oxide & carbon dioxide. (7a)

When copper (II) carbonate is not available, your teacher may substitute another carbonate or delete this part.  If a white carbonate is used, dramatic color changes should not be expected.  Do note the texture change, however.



1.  Obtain a clean 18 x 150 test tube.  Place a scoop of copper (II) carbonate into the test tube.  The copper (II) carbonate should measure about two centimeters (about three quarters of a "thumb-knuckle" inch) in the tube.

2.  Observe and record the color and texture of the copper (II) carbonate.

3.  Mount the large test tube in a utility clamp on a ring stand.  Ensure that the clamp is placed at the mouth of the tube and that the tube is tipped so that the mouth is about 1 cm lower than the closed end.  If the carbonate is not completely dry some moisture that is not associated with the reaction might be produced.  If condensed water is allowed to roll back on the hot glass, it is likely that the glass will break.

4.  Place a delivery tube in the test tube containing the carbonate.  Pour about 3 mL of limewater into a 13 x 100 mL test tube.

5.  Heat the test tube and carbonate over a medium flame while holding the limewater in contact with the end of the delivery tube so that any gases produced will bubble into the limewater.

6.  If copper (II) carbonate or some other colored carbonate was used, refer to the CRC Handbook to find the colors of the oxide and the carbonate.  Compare the colors to those you observed.

7.  Test a small amount of the original carbonate and a small amount of the product in the test tube with acid and base indicators.  Remember that an oxide will be produced and what happens to an oxide when placed in water.


Part I.

Many metallic bicarbonates, when gently heated, will decompose to produce their metallic hydroxide and carbon dioxide.  Strong thermal decomposition of bicarbonates further decompose the hydroxide produced.  This results in a two step reaction. (7b, 7c)

Repeat the steps from Part H using sodium bicarbonate.


Part J

Carbonates and bicarbonates react with most acids to produce carbon dioxide, water, and a salt of the acid. (7d)

1.  Mount an 18 x 150 test tube at an angle of approximately 45 degrees, closed end downward; use a utility clamp and ringstand.

2.  Place 3 mL of lime water and 4 drops of phenol red into a clean, 13 x 100 test tube.

3.  Place 5 mL of 3M HCl into the clamped tube.  Slide 3 marble chips into the tube containing the HCl.  Place the stopper containing the delivery tube into the mouth of the clamped tube.

4.  Bubble the resulting gas into the tube containing the limewater and phenol red.  Observe the changes in the lime water and phenol red mixture.

DISPOSAL:  Place the contents of the tube in the disposal beaker labeled PART J disposal.

Part K.

Metals from families IA and IIA will generally displace hydrogen from water and form a hydroxide. The elements Be & Mg require steam to so behave. (10)

There are risks associated with this step that preclude the use of most of these elements consistent with the school district's adopted science safety manual. This chemical reaction will be illustrated by use of a demonstration viewed on the laser disk.

Part L.

A metal with a higher electrode potential will displace a metal with a lower electrode potential from an aqueous solution containing ions of the element with the lower electrode potential. (8)

1. Obtain a clean, 13 x 100 test tube. Place 3 mL of 0.5M copper (II) sulfate solution into the tube.

2. Obtain a small amount of iron filings, about the volume of a pea.

3. Place the iron filings into the tube containing the copper (II) sulfate. Observe any changes that occur within the first 5 minutes. Put the tube aside and observe it again toward the end of the period.

4. Obtain a clean, 13 x 100 test tube. Place 3 mL of zinc sulfate solution into the tube.

5. Place 1 piece of copper shot or small piece of copper wire into the tube containing the zinc sulfate. Observe any changes that occur within the first 5 minutes. Put the tube aside and observe it again toward the end of the period.

DISPOSAL: Put solids and liquids in the designated beaker. Rinse with water into the beaker to ensure that no metals go down the sink.

Part M.

Metals with an electrode potential greater than 0.00 volts will displace hydrogen from an acid. (9)

1. Place one very small piece of each of the following metals in separate large wells of a microchem plate:  Zn, Pb, Ca, Mg, Fe, Cu, Al.  Your teacher may substitute some other metals or add other metals to the list.  Be sure to have a sheet labeled with the well locations.

2.  Place 5 drops of 6M HCl into the first well.

4.  After one minute, add acid to the second metal.  Observe that well for one minute.  You may want to use a hand lens.  You might also want to use a microthermometer if they are available.  Some reactions are quite exothermic.

5.  Repeat this process for each of the remaining metals.

6.  Note that some reactions proceed well even though they take some time to get started.

7. Develop a comparison of the rate of reaction.  Which reacted rapidly and which reacted slowly or not at all.

8.  Compare your basis of reaction rate to that of the electrode potential of the metals used.

DISPOSAL:  Metals must not be allowed to go down the sink as the drains will become clogged.  Dilute the acid by flooding the plate and allowing the liquid to overflow into a large beaker or pail supplied for this purpose.  If a strainer is available, strain the liquid so that the metals can be wrapped in a paper towel and disposed in the trash can.

CLEANING The microchem plate:  Weak detergent solution and a Q-tip must be used to clean the plate.  Leave the plate clean and as dry as possible on a white piece of paper for inspection.  YOU WILL NEED TO BRING Q-TIPS.

Part N.

N.  A free halogen atom will displace a less active halogen (in an ionic form) from a salt solution of the less active halogen. (11)

1.  Obtain three clean 13 x 100 test tubes, and place 1 mL  of  water in each tube

2.  Add a very small amount of crystalline sodium chloride to the first tube, sodium bromide to the second tube, and sodium iodide to the third tube.  Or use 1 mL of 0.1M solutions of these compounds if your instructor directs you to do this.  If the sodium salts are not available, then you can use potassium salts instead.  Do you know why?  In this part, look at the demonstration tubes set up by your teacher to see what is meant by the term a very small amount or about.

3.  Place about 1 mL  of chlorine water in each tube.

4.  Place about 0.5 mL   of water immiscible solvent in each tube.

5.  Place a cleaned #00 rubber stopper in the first tube.  Hold the tube so that your thumb can be placed on top the stopper to hold it in, and shake the tube for about ten seconds.  Shake the tube with adequate vigor to cause the liberated halogen, if any, to dissolve in the solvent.

6.  Rinse the stopper well and repeat the process with the other two tubes.  Note the color imparted to the solvent layer each time you perform this shaking.


DISPOSAL:  Place the contents of the tubes in the beaker marked HALOGEN DISPOSAL.
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