LABORATORY ACTIVITY:
Activity 1: Carbon Dioxide
Introduction
CO 2 plays a major role in the biogeochemical carbon cycle. You will examine some properties of CO 2 and its derivatives to develop an appreciation of the many roles this compound plays in the carbon cycle.
Purpose
To prepare carbon dioxide and study some physical and chemical properties of carbon dioxide and carbonates.
Safety
1. Wear protective goggles throughout the laboratory activity.
2. Immediately wash off any chemical with water if it contacts your skin.
3. Follow all normal precautions associated with the use of strong acids and bases.
4. Waft gases gently toward your nose when testing for odors.
5. Insure that proper care is used when lubricating and inserting glass tubing or thistle tubes into stoppers.
6. Make sure that the thistle tube is below the liquid level in the gas generator to prevent possible loss of CO 2 and spewing of acid.
7. Do not look directly at burning magnesium, since ultraviolet radiation, as well as the brightness, may cause eye injury.
8. Do not discard solids in the sink. Place in container provided by your teacher.
Procedure
1. Set up four test-tubes in a rack. Place approximately 2.5 g of calcium, magnesium and potassium carbonate in each of three tubes and sodium bicarbonate, NaHCO 3 , better known as “baking soda,” in the fourth. Add 3-4 mL of water to each and stir each mixture to effect maximum dissolution.
2. a. Obtain 75 mL limewater, a saturated solution of Ca(OH)2 , in a 250-mL beaker. b. Add a drop or two of 3.0 M hydrochloric acid, HCl, a strong acid, to each test-tube, fit rubber stopper into tube, and bubble any gas evolved into the limewater as shown in Figure 1. This is a standard qualitative test for carbon dioxide, CO 2 . Figure 1. Bubbling CO 2 (g) into limewater.
3. Set up the CO 2 generator as shown in Figure 2. Before adding any acid, be sure to have your instructor check your set-up. Just cover the bottom of the gas generator bottle with marble chips, and carefully add enough (about 10 mL) 3.0 M HCl through the thistle tube to cover the marble. Be sure the thistle tube is below the acid level.
4. Discard the first bottle of gas collected. Adding 10 mL portions of 3.0 M HCl as needed, collect three bottles of CO 2 , and confirm that you are indeed producing carbon dioxide by bubbling the gas into a test-tube of 3-4 mL limewater.
5. Continue to allow CO 2 to bubble into the tube of limewater after the initial confirmation that CO 2 is being produced. What happens now? What you observe occurs because calcium bicarbonate,
Ca(HCO 3 )2 , is much more soluble than calcium carbonate, CaCO 3 . The clear solution is now “temporary” hard water.
6. Shut down your CO 2 generator by unstoppering and filling the generator bottle with water. Carefully decant this diluted acid, and rinse the marble chips several times with water. No solids in the sinks. Place the marble chips in the container provided by your teacher for reuse.
7. Bottle 1: Set up a 250-mL beaker containing a short burning candle inside. Being careful not to drip excess water, “pour” (as you would water) a bottle of CO 2 into the beaker over the flame. What do you observe?
8. Bottle 2: Holding a short piece of magnesium ribbon with small forceps, ignite in a burner flame, and immediately thrust the burning magnesium into a bottle of CO 2 . Do not drop the ribbon as it might break the bottle. What do you observe?
9. Bottle 3: Into the last bottle of CO 2 , add 50 mL of water and, covering the mouth tightly with a stopper, shake vigorously. What evidence of dissolving do you observe?
10. Make a very dilute solution of NH 3 (aq) (one drop 3.0 M NH 3 (aq) in approximately 400 mL water) and add 2-3 drops of phenolphthalein. The solution should be very pale pink. Now add your CO 2 (aq) solution to 25 mL of the basic solution. What happens?
11. Test aqueous solutions of NaHCO 3 and Na 2 CO 3 with pH paper, litmus paper, and phenolphthalein. Describe the results.
12. Thoroughly wash your hands before leaving the laboratory.
Data Analysis and Concept Development
1. Which substances in Step 1 are soluble and insoluble?
2. Is a gas evolved when HCl is added in Step 2?
3. What happens when the gas in Step 2 is bubbled into limewater?
4. Bubbling CO 2 into limewater is a standard qualitative test for CO 2 . State the test in your own words.
5. Write an equation for the reaction of CO 2 with limewater.
6. Why must the thistle tube be below the acid level in the CO 2 generator?
7. Why should you discard the first bottle of CO 2 collected in Step 4?
8. Write one equation for the reactions in Steps 4 and 5 using double arrows.
9. How does Le Chatelier’s Principle apply to Steps 4 and 5.
10. How is the reaction from Question 8 a contributor to the hardness of natural water?
11. What physical and chemical properties of CO 2 make it a good fire extinguisher?
12. Explain your observations in Step 7 in terms of these properties.
13. What are the white powder and the black deposits on the sides of the bottle in Step 8?
14. How could these be formed?
15. Write a balanced equation for this oxidation-reduction reaction.
16. Describe any evidence of dissolving observed in Step 9.
17. What color change occurred in the NH 3 (aq) solution when you added the CO 2 (aq) in Step 10?
18. What does this color change indicate about a solution of CO 2 ?
19. Describe the results of your tests in Step 11.
20. Is the hydronium ion concentration of Na 2 CO 3 greater or smaller than that of water? Is the hydronium ion concentration of NaHCO 3 greater or smaller than that of water?
21. Which compound has the higher pH, Na 2 CO 3 or NaHCO 3 ?
22. Explain your answer to Question 21 in terms of hydrolysis and Le Chatelier’s principle.
Implications and Applications
1. Discuss why CO 2 as a fire extinguishing substance would be appropriate or inappropriate in each of the following:
a. An alcohol fire in a chemistry laboratory.
b. A thermite reaction (burning aluminum and magnesium) gone out of control in the laboratory.
c. An electrical fire in the roof of an elevator containing people stuck between floors.
d. A grease fire in your kitchen.
2. Discuss the baking alternatives: baking soda and yeast. Is there any advantage in using one or the other?
3. Define hard water, distinguishing between the “permanent” and “temporary” kinds. “Temporary” hard water, chiefly composed of calcium and magnesium bicarbonates, can be “softened” by boiling or by the addition of sodium carbonate. Discuss the mechanism in each method that removes the “hard” ions. Which type would you prefer in your community and why?
4. In emergency rooms, a shout of “Bicarb!” is often heard, especially for victims of cardiac arrest or strangulation. What reason would there be for intravenous “bicarb infusion?
Laboratory Activity:
Activity 1: Carbon Dioxide
Major Chemical Concept
Students should develop an appreciation for the many roles CO 2 plays in the biogeo-chemical carbon cycle&emdash;acid former, important constituent in carbonate-bicarbonate equilibria, inorganic carbon source in photosynthesis, and respiration product. Additionally, by observing carbon dioxide’s density, nonflammability, and inability to support combustion, they will more clearly understand its uses in fire fighting, as well as some of its disadvantages and hazards under certain conditions (such as metal fire).
Students apply previously learned concepts in new and practical situations: acid-base theory, hydrolysis of salts, Le Chatelier’s Principle and equilibrium, oxidation-reduction, and writing equations.
Level
This activity is appropriate for any general first-year chemistry class.
Expected Student Background
Since this is an enrichment module, the material will probably be used late in the school year. It presupposes a wide range of conceptual understanding. Students should be familiar with Brønsted-Lowry acid-base theory, hydrolysis, concentration effects in equilibrium (Le Chatelier’s Principle), oxidation-reduction, and how to balance equations. In addition, they should have a practical knowledge of laboratory techniques, specifically the generation and collection of gases.
Time
This activity can be completed in two 50-min periods, but it should be done over three periods if you use extended discussion to show applications and reinforce previously acquired concepts.
Safety
Read the Safety Considerations in the Student Version. Marble chips can be “recycled,” preventing the dumping of solids in the sink as well as stressing conservation of resources and economy. Make sure the chips are rinsed well and then placed on paper towels to dry before putting back into container. Plastic thistle tubes are recommended instead of glass. Students should take care when lighting the short candle in the beaker. Use a splint, and do it quickly. Warn students not to look directly at burning magnesium, since it emits ultraviolet radiation that may harm their eyes. Also remind them not to drop the burning ribbon since it may crack the beaker.
Materials (For 24 students working in pairs)
Nonconsumables
12 Thistle tubes
48 Wide-mouth bottles
12 Pneumatic troughs
12 Graduated cylinders, 10-mL
12 Glass bends
12 Stoppers, 2-hole (sized to fit bottle)
12 Flexible rubber tubing pieces, 45 cm
48 Test-tubes, 25- x 150-mm
12 Test-tube racks
12 Spatulas
12 Beakers, 250-mL
12 Beakers, 400-mL
12 Forceps, small
36 Glass plates, 3" x 3", or large enough to cover the mouths of the collecting bottles
12 Rubber stoppers, solid, for wide mouth bottles
12 Rubber stoppers, 1-hole fitted with glass bends and rubber tubing No. 2
12 Stirring rods
Consumables
Sodium bicarbonate, NaHCO 3 , 30.0 g
The following carbonates are suggested for Step 1. Check safety before using other carbonates:
Potassium carbonate, K 2 CO 3 , 30.0 g
Magnesium carbonate, MgCO 3 , 30.0 g
Calcium carbonate, CaCO 3 , 30.0 g
Marble chips
12 Candles, small (short enough so they do not extend above top of 250-mL beaker); birthday cake candles in clay blob “candle holders” work well
12 Splints
12 Magnesium ribbons, approximately 2 cm long
Litmus paper
pH paper
Phenolphthalein solution, 100 mL (1% in ethanol, divided into several dropping bottles)
3.0 M Hydrochloric acid, HCl, 1.0 L (250 mL 12 M HCl per 1 L solution)
3.0 M Ammonia, NH 3 , 100 mL (20 mL 15 M NH 3 per 100 mL solution)
0.25 M Sodium carbonate, Na 2 CO 3 , 100 mL (2.7 g anhydrous Na 2 CO 3 per 100 mL solution)
0.50 M Sodium bicarbonate, NaHCO 3 , 100ÊmL (4.2 g anhydrous NaHCO 3 per 100 mL solution)
Saturated solution of limewater (Place enough Ca(OH)2 powder or flakes&emdash; approximately 25 g&emdash;in a 1 L bottle deionized water to make the saturated stock solution. Prepare 24 hours ahead.)
Advance Preparation
1. Prepare solutions needed for the activity.
2. Check to be sure the NH 3 (aq)-phenolphthalein solution students will prepare is dilute enough to be decolorized by the weak carbonic acid solution generated. If not, dilute the stock solution of NH3 until you obtain the desired result.
Pre-Laboratory Discussion
1. You may wish to discuss CO 2 in these terms prior to the laboratory activity. CO 2 plays a major role in the biogeochemical carbon cycle. On the geo side its capture from the atmosphere accounts for most of the “sink” of insoluble mineral carbonates in the earth’s crust and ocean floor. It also plays a major role in the production of “hard” water and helps buffer the oceans, maintaining a fairly constant pH for seawater through the production of bicarbonate ions. On the bio side it is essential to plants as the source of inorganic carbon from which organic carbohydrates can be synthesized, storing the energy ofsunlight as chemical energy. CO 2 returns to the atmosphere as a by-product of cellular respiration of carbohydrates and the release of their energy, but not before providing some important buffering activities in the circulatory systems of many organisms. Finally, CO 2 absorbs reflected infrared radiation from the earth’s surface, targeting it as the major cause of global warming and the greenhouse effect. As the product of the burning of vast quantities of carbon-containing fossil fuels (along with water, also a greenhouse gas), the atmospheric concentration of CO 2 has steadily increased during thelast two centuries. Whether or not natural geochemical processes can adapt to keep up with this increased atmospheric loading is debatable.
2. As part of this topic, stress concern for conservation of resources by example. “Recycle” the marble chips by rinsing and saving. This may seem trivial, but it does make an impact on students. Make sure the chips are rinsed well and then placed on paper towels to dry before putting back into container.
3. Caution students not to leave the top off of the limewater container. Why? (CO 2 from the air will react.)
4. Review safety emphasizing:
a. Proper handling and disposal of acids.
b. Proper set-up for gas generation and approval policy. (It should be your policy to inspect generators before the acid is added.)
c. Warning about eye injury from looking directly at burning magnesium strip.
d. Warning about lighting the short candle with care, using a splint.
5. To prevent gas from bubbling out through the thistle tube, emphasize that the tube should almost touch the bottom of the generator bottle so the liquid level completely submerges it. Sometimes an airlock prevents additional acid from going down into the generator. Tapping the bottle on the counter top gently, being careful not to slosh acid out, will break the airlock and allow the acid to flow. Do not allow students to complete this procedure&emdash;only the teacher. (If tapping does not work, briefly loosen the generator stopper; tighten immediately.)
6. Remind students to discard the first bottle of gas generated, as it will be contaminated with air. Three bottles of CO 2 can be collected easily in less than 15 min, although a recharge of acid may be necessary.
7. Several important concepts that have been studied previously should be reviewed here. Density (specific gravity) of gases, factors affecting equilibrium (Le Chatelier’s Principle), Brønsted-Lowry acid-base concept (see respective SourceBook modules) will all play a role in understanding why and how some of the reactions to be observed work the way they do.
You might also wish to review the meaning of acid anhydride, relating this concept to CO 2 the oxide of the nitrogen cycle and acid rain. 8. Emphasize the practicality of the knowledge to be learned and the many applications of these reactions.
Teacher-Student Interaction
First, walk around and check to be sure students understand how to set up the limewater test in Figure 1. More importantly, monitor insertion of glass tubing and thistle tubes into and out of stoppers. (Some teachers like to fit the rubber stoppers with glass tubing and thistle tubes in advance. If you choose to allow the students to prepare their own set-ups, be sure to show them the proper way of inserting glass tubing into a rubber stopper, including lubrication of the tubing with H 2 O and use of a laboratory towel for hand protection. An alternative way to fit the tubing is to insert a well- lubricated&emdash;with silicone grease&emdash; metal cork borer of larger diameter then the glass tubing into the stopper, slip the glass tubing through the borer and remove the borer, leaving the glass tube in place in the stopper. Although this method requires a little more advance planning, it is much safer than the standard methods.) Advise students concerning proper placement of the thistle tube (below liquid level), and remind them to discard the first bottle of gas collected. The reaction, once you have approved each set-up and students have added the acid, is rather rapid. As the rate of bubbling slows, instruct students to add more acid in 10 mL portions.
Remind students to test the gas from the generator also, using limewater, and then stop the reaction. Walk around and make sure they do not leave marble chips reacting in the unstoppered bottles, do not dump solids into the sinks, or forget to rinse the marble and then put it on paper towels to dry. As students begin to use their collected bottles of CO 2 , monitor their work by walking around and asking questions: “Why did the candle go out?”
“Why could you pour the CO 2 ?” “What does a higher pH mean?” Or make comments:
“Be careful not to drop the burning strip.” “That ammonia solution is too pink; dilute it more.”
Anticipated Student Results
Many of the observations will be given in question responses to Data Analysis and
Concept Development (see Answers to Data Analysis and Concept Development). In Step 2, a gas is evolved. This gas is suspected to be CO 2 . At the end of Step 4, CO 2 is bubbled through limewater to produce a cloudy suspension as confirmation that the gas generated is CO 2. However, in Step 5, the cloudiness disappears (see Post-Laboratory Discussion for relevant equations). In Step 7, pouring CO 2 over the flame should extinguish the flame. In Step 8, a black solid (C) and a white solid (MgO) are formed when Mg is burned in CO 2 . In Step 9, suction may be felt as the CO 2 dissolves in water. (Pressure decreases as gas dissolves.)
Answers to Data Analysis and Concept Development
1. Soluble: Na 2 CO 3 , NaHCO 3 , (NH 4 ) 2 CO 3
Insoluble: BaCO 3 , CaCO 3 (MgCO 3 is only slightly soluble.)
2. Yes
3. Limewater turns cloudy.
4. If an unknown odorless, colorless gas is bubbled into limewater and the clear solution becomes cloudy, the gas is CO 2 .
5. CO 2 (g) + Ca(OH)2 (aq) ® CaCO 3 (s) + H 2 O(l)
6. Otherwise, CO 2 bubbles out through the thistle tube instead of going into the collection bottles.
7. The first bottle is contaminated with air that was in the generator. Since it is contaminated with air, you might get false observations.
8. Ca(OH)2 (aq) + 2 CO 2 (g) CaCO 3 (s) + CO 2 (g) + H 2 O(l) Ca(HCO3 ) 2 (aq) [CaCO 3 (s) + CO 2 (g) + H 2 O(l) Ca(HCO 3 ) 2 (aq) C a 2+ (aq) + 2 HCO 3 – (aq) ]
9. As the concentration of CO 2 increases, it drives the reaction beyond CaCO 3 formation to formation of Ca(HCO 3 ) 2 , which is soluble.
10. Hard water contains Ca 2+ (aq), and soluble Ca(HCO 3 ) 2 adds Ca 2+ (aq) to water.
11. Physical property: density; denser than air. Chemical property: does not support combustion.
12. The denser CO 2 settles in the beaker, driving air out. This “smothers” the flame.
13. The white powder is MgO; the black specks are carbon.
14. To continue burning and forming MgO, magnesium literally tears CO 2 apart to remove the oxygen, leaving the carbon behind.
15. 2Mg(s) + CO 2 (g) ® 2MgO(s) + C(s)
16. Suction is felt on the palm.
17. The solution became colorless.
18. It has acidic properties; it neutralizes the base.
19. Litmus turns blue in both solutions, and phenolphthalein turns pink in both. pH paper indicates a higher pH for Na 2 CO 3 than NaHCO 3 .
20. The hydronium ion concentration of both Na 2 CO 3 and NaHCO 3 is smaller than that of H 2 O.
21. Na 2 CO 3 has the higher pH.
22. CO 3 2– is a stronger base than HCO 3 – ; the equilibrium representing CO 3 2– reacting with H 2 O goes farther to the right. The pertinent equations are: HCO 3 2– (aq) + H 2O(l) CO 2 (g) + H 2 O(l) + OH – (aq) CO 3 2– (aq) + H 2 O(l) HCO 3 – (aq) + OH – (aq)
Answers to Implications and Applications
1.
a. A CO 2 fire extinguisher would be appropriate for use in a laboratory alcohol fire because it would smother the fire by cutting off the oxygen needed to feed combustion.
b. A CO 2 fire extinguisher would not be appropriate with a thermite reaction since burning aluminum and magnesium would be hot enough to react with the CO 2 in a reaction similar to the one observed in the Bottle 2 (Step 8) reaction. Covering the burning mixture with sand is the appropriate method of removing oxygen here.
c. A CO 2 fire extinguisher would not be appropriate in this situation since the density of carbon dioxide is so high. It would tend to settle to the floor of the elevator, displacing air, and creating a suffocating pocket of CO 2 that might harm passengers. This would be similar to a common situation faced by miners in coal mines before ventilation requirements became so strict.
d. A CO 2 fire extinguisher would be appropriate in a kitchen grease fire in order to suffocate the combustion process. Water or a soda-acid fire extinguisher would be inappropriate, since the immiscibility of grease and water would result more in spreading the fire than cooling it.
2. Baking soda, NaHCO 3 , dissociates in aqueous medium to Na + (aq) and
HCO 3 – (aq). The HCO 3 – (aq) reacts with H + (aq) in the mixture to produce H 2 CO 3 (aq) that, upon heating, produces CO 2 (g) and H2O(l). As the CO 2 diffuses through the dough, it creates bubbles that cause the dough to rise. After the CO 2 and water vapor leave the dough, the only residue is a sodium salt, which may not be desirable for people with high blood pressure. Note, also, that heating is necessary to cause the dough to rise. In the case of yeast, enzymes (see Enzymes module) accelerate the metabolism of sugar according to the following reaction:
C 6 H 12 O 6 (s) + Yeast ® 2C 2 H 5 OH(l) + 2CO 2 (g)
rendering heat unnecessary above normal temperatures. After the dough has risen because of diffusion of CO 2 , baking stops the fermentation process by killing the yeast and evaporating off the alcohol. (Have you ever eaten rolls that tasted of alcohol? What might have caused this?) The lack of sodium and the rising of the dough without heat can be considered advantages of using yeast instead of baking soda. The disadvantage is the increased time that it takes to produce the bread.
3. Hard water contains dissolved minerals, usually containing calcium, magnesium or aluminum ions, but no bicarbonate anions. It does not lather easily and leaves a soap scum on tile. This scum is the insoluble“salt” of the soap. An example would be the conversion of soluble sodium laurate, a common soap found in some shampoos, into insoluble calcium laurate, which precipitates as a “soap film. In the case of “temporary” hard water, boiling decomposes the soluble bicarbonate into the insoluble carbonate, which removes the “hard”ions, leaving “soft” water. Replacement of bicarbonate with carbonate ions through addition of sodium carbonate accomplishes the same result. The fact that heating causes a precipitate to form makes temporary hard water a nuisance since insoluble carbonates build up in such household appliances as tea kettles, washing machine inlet pipes, and shower heads, and coffee makers.
4. In cardiac arrest or strangulation, the cessation of breathing causes a build up of CO 2 as the living cells continue to respire. This “dumping” of CO2 into the bloodstream causes a drop in blood pH that itself can prove fatal:
CO 2 (g) + 2H 2 O(l) H 3 O + (aq) + HCO 3 – (aq)
At first the blood’s buffering capacity, which is due to blood bicarbonate concentration, attempts to maintain a constant pH by removing hydronium ions produced as CO 2 concentration increases by the reverse process:
HCO 3 – (aq) + H 3 O + (aq) CO 2 (g) + 2H 2 O(l)
As the blood’s bicarbonate concentration becomes depleted by this process, an infusion of “bicarb” is needed to try to maintain this buffering effect in the bloodstream.
Post-Laboratory Discussion
The discussion questions in Implications and Applications can lead to misconceptions or problems arising from the activity. You will particularly want to go through all reaction equations explaining what happened. (The limewater test reactions are given in Question 8 in Answers to Data
Analysis and Concept Development. Hydrolysis reactions are in Question 23 in Answers to Data Analysis and Concept Development.) This is particularly important since equilibrium shifts and dissociation of weak and strong acids in the hydrolysis reactions may not be clearly understood. A possible extension of the post-laboratory discussion can center around CO2 and its role in the “greenhouse effect.”
Assessing Laboratory Learning
Besides the normal laboratory report, including answers to the questions asked, the following laboratory practical can be done. A long candle is set up in a large beaker to which limewater has been added. A bottle of air is lowered quickly so that the mouth of the bottle is below the surface of the limewater, as shown in Figure 3. Account for the following observations in light of the CO 2 activity just completed:
1. The candle slowly goes out. O 2 , which is needed for combustion, is the limiting reactant. Burning ceases when it is depleted.
2. The limewater level rises inside the bottle. A combination of reaction of CO 2 with limewater and rapid cooling of hot gases, with the resultant decrease of volume, reduce gas pressure inside the bottle. Atmospheric pressure is now greater and forces limewater up into the bottle.
3. The limewater inside the bottle becomes cloudy at the surface only. This observation proves the CO 2 is produced by the burning of the candle, and that some of it is being removed by the production of insoluble CaCO 3 .
NOTE: The small amount of CO 2 in the original air could not account for the surface “film.”
