DEMONSTRATIONS
Demonstration 1: Carbon as a Component of Sugar
Purpose
Carbon, oxygen, and hydrogen are combined in the presence of chlorophyll and sunlight in the plant to produce carbohydrates.
In this demonstration, you will use sulfuric acid to dehydrate the sugar, leaving carbon.
C 12 H 22 O 11 (s) H 2 SO 4 (conc) 12C(s) + 11H 2O(g)
Materials
Small beaker, 150-mL (It will be difficult to clean afterwards, so you might want to use an old one.)
Beaker, 400-mL
Sucrose, C 12 H 22 O 11 (granulated table sugar)
Sulfuric acid, H 2 SO 4 , concentrated, 10 mL
Beaker, 600-mL
Safety
This demonstration will generate heat, which would possibly crack the beaker, and copious noxious fumes consisting of water mixed with sulfuric acid. It must be done in a hood.
Handle concentrated sulfuric acid with extreme care.
Use gloves, safety goggles, and a safety shield.
Procedure
1. Fill the small beaker with “30 mL” sugar and place this beaker inside a 600-mL beaker containing about 100 mL H 2 O.
2. Measure the temperature of the H 2 O in the larger beaker.
3. Carefully pour sulfuric acid onto the sugar until there is a layer about l/8 of an inch deep (about 10 mL). Do not stir.
4. Set the beaker system aside in the hood and watch for a reaction to occur as evidenced by a color change where the acid contacts the sugar. The whole mixture will “puff up” dramatically.
5. After the reaction is completed, measure the temperature of the water again. The highly exothermic reaction should increase the water temperature. [10 mL conc H 2 SO 4 produces about a 10 ºC rise in temperature.]
Remarks
This reaction is a dehydration reaction represented by the following equation: A secondary reaction explains why the carbon forms a solid foam in the beaker. The escaping gases cause the foaming.
C(s) + 2H 2 SO 4 (l) ® CO 2 (g) + 2SO 2 (g) + 2H 2 O(g)
Demonstration 2: The Solubility of Carbon Dioxide and Ammonia
Purpose
A major part of carbon and oxygen cycles involves dissolution of carbon dioxide gas in aqueous solution. In fact, the huge amount of carbon dioxide dissolved in the oceans provides a valuable “sink,” or reserve supply, of carbon dioxide.
Ammonia, one of the more important compounds of nitrogen is also part of the nitrogen cycle. Two demonstrations are suggested here to show the extent of solubility of these two gases and the effect of temperature on their solubility.
A. Carbon Dioxide
Materials
Fresh bottle of soda pop, preferably chilled (10 oz bottle is sufficient)
Empty plastic soda bottle (10 oz)
One-hole rubber stopper (size to fit soda bottle)
Glass tubing
Rubber tubing, 60 cm
Pneumatic trough
Large beaker filled with warm water
Limewater, 25 mL (optional; see Activity 1 for preparation)
Safety
No special safety precautions are required.
Procedure
1. Prepare a one-hole rubber stopper to fit the mouth of a 10 oz plastic bottle of soft drink, like “Sprite” or seltzer water. Put a glass tube in the stopper hole and connect a 60-cm piece of rubber tubing to it. Prepare a water trough to collect gas by water displacement. Fill an empty 10 oz soft drink bottle with water, place it inverted in the water trough. Position the tube through the mouth and into the bottle.
2. Open a fresh bottle of precooled soda pop, immediately insert the rubber stopper, and gently shake the bottle and warm it by placing it in a large beaker of warm water.
3. Collect the carbon dioxide by displacing the water in the inverted bottle. (Students will be amazed at the amount of gas evolved and collected. You will be able to completely fill the second bottle with gas.)
Extensions
To prove that this gas is indeed carbon dioxide, remove the bottle containing the gas and immediately add about 25 mL limewater. A cloudy white precipitate of calcium carbonate is a positive test for carbon dioxide.
CO 2 (g) + Ca(OH)2 (aq) ® CaCO 3 (s) + H 2 O (l)
Remarks
At 20ºC, 0.9 L of CO 2 will dissolve in 1 L of water.
B. Ammonia (The Ammonia Fountain)
Purpose
To test the solubility of ammonia gas in hot and cold water.
Materials
Aqueous ammonia, NH 3 , concentrated, 10 mL
1000-mL beaker of hot water
1000-mL beaker of cold water
Florence flask, 500-mL
One-hole rubber stopper to fit Florence flask
Burner
Glass tubing
Phenolphthalein solution, few drops (optional)
Vinegar, few drops (optional)
Safety
Students should avoid inhaling ammonia vapor and should not handle hot glassware with their hands.
Procedure
1. Prepare a one-hole rubber stopper with 20-cm piece of glass tubing to fit a large (500-mL) Florence flask.
2. Place about 10ÊmL concentrated aqueous ammonia in the flask and insert the stopper. Have a large beaker of hot water and a large beaker of cold water nearby (Figure 6).
3. Gently heat the flask with a burner for 10-15 s and immediately invert it and place the stoppered mouth into the beaker of hot water. (Students will observe that no water from the beaker enters the flask.) Figure 6. Ammonia fountain demonstration. Remove the flask, heat it again for a few seconds, and invert it in the beaker of cold water. (The water in the beaker will enter and fill the flask.)
Extensions
You might want to add a few drops of phenolphthalein indicator to the water in the beakers and a few drops of vinegar. When the dilute acid containing phenolphthalein enters the flask, it will turn pink.
Remarks
Ammonia readily dissolves in cold water, but its solubility in warm water is less. Ammonia in solution is called “aqueous ammonia.” It is incorrect to call this solution “ammonium hydroxide,” since such a molecule has never been shown to exist
(see Acids and Bases module).
Demonstration 3: Formation of NO2 and N2O4 Gases
Description
This demonstration illustrates the NO 2 – N 2 O 4 equilibrium.
Materials
Nitric acid, HNO 3 , concentrated, 10 mL
Copper penny, pre-1982 or 1-2 g Cu wire
Side-arm flask, small
Rubber tubing
Two test-tubes
Rubber stoppers, 3 (one to fit the side arm flask, two for the test-tubes)
Beaker of boiling water
Beaker with ice bath
Strapping tape
Safety
Concentrated nitric acid is corrosive. You must wear gloves and a face shield when using this acid. NO 2 is a very toxic gas. It must be generated in a hood. The tubes should be stoppered tightly and sealed with tape. It has been estimated that as few as five pennies could generate enough NO 2 to reach a toxic level in a school laboratory.
Procedure
A. Preparing the Equilibrium Tubes
1. Prepare a gas generator by attaching a rubber tube to a small side-arm flask.
2. Working in a hood, add about 10ÊmL concentrated nitric acid to the flask. Drop in a pre-1982 penny. A reddish-brown gas, NO 2 , will form immediately. Allow enough of this gas to be formed to displace the air in the flask and stopper the flask.
3. Fill two test-tubes with gas from the generator and immediately stopper the tubes tightly. Color intensity should be the same in both tubes initially.
4. Stop the reaction in the flask by filling it with water. (Note the blue color of the solution formed.)
B. Using the Equilibrium Tubes
1. Place one of the tubes in a beaker of boiling water. Notice the change in color of the gas in the tube.
2. Place the second test-tube in an ice bath. Notice the change in color of the gas in this tube.
Remarks
1. When copper reacts with concentrated nitric acid, reddish-brown, toxic NO 2 (g) is formed.
Cu(s) + 4H + (aq) + 2NO 3 – (aq) ® Cu 2+ (aq) + 2NO 2 (g) + 2H 2 O(l)
The equilibrium mixture in the tubes consists of NO 2 and N 2 O 4 gases, often referred to an NO x . They react according to the equation: 2NO 2 (reddish-brown) N 2 O 4 (colorless) + Heat When the mixture is heated, the equilibrium shifts toward the formation of brown NO 2 . When it is cooled, equilibrium shifts toward the formation of colorless N 2 O 4 . When dry NO 2 is collected, as in this demonstration, oxygen in the air maintains the nitrogen in the N(IV) state. When NO 2 is collected over water, a different set of reactions occurs (see Industrial Inorganic Chemistry,
Demonstration 2).
GROUP AND DISCUSSION ACTIVITIES
Other Demonstration Ideas
Refer to the Simple Chemical Reactions module for additional demonstrations.
Key Questions
1. What is the difference between a dynamic equilibrium and a steady state equilibrium? A dynamic equilibrium occurs within a closed system only. It involves changes in which the same atoms move back and forth, and is affected as Le Chatelier’s Principle predicts. A steady state equilibrium occurs in an open system only. It involves different atoms that enter and leave at the same rate with no net change, and does not usually behave as LeÊChatelier’s Principle would predict.
2. We have all heard of “nitrogen-fixation.” Is it possible to “fix” or “convert to a more useful condensed state” other gases? Certain bacteria “fix” atmospheric nitrogen as more useful solid nitrates (usually in solution) in the roots of legumes. Photosynthesis, on the other hand, takes atmospheric carbon dioxide and “fixes” it as carbohydrates (usually in solution).
3. What source produces the greatest amount of carbon dioxide released to the atmosphere? Contrary to popular belief, respiration or combustion processes do not release the largest amount of CO 2 to the atmosphere. Aerobic fungal and bacterial decomposition of dead organic matter is the major source of atmospheric CO 2 .
4. Compare the geologic processes that change limestone into marble with those that change peat into anthracite coal. Heat and pressure, created by deep burial or volcanic upheaval, literally melt sedimentary limestone to allow it to become metamorphic marble. Individual particles become fused into one mass. In coal formation, the same events create pressure and heat that, coupled with the initial action of aerobic and anaerobic bacteria, volatilize impurities and fuse the carbon particles together more and more over time. Bituminous, or soft coal, is “dirty” and rubs off easily on a white glove. Anthracite, or hard coal, that has been exposed to the greatest amount of metamorphism, is shiny and“ clean,” not producing much coal dust since the carbon particles have been melted and fused. That is why its carbon will not easily rub off on a white glove. Removal of volatile impurities gradually allows the coal to become purer carbon, forming a more perfect graphite ring structure.
5. Even if there were no industrial pollution, rainwater would still be weak “acid rain.” Why? As rainwater falls through the atmosphere, CO 2 dissolves into it. Since CO 2 is an acid anhydride, it reacts with water to form weak carbonic acid, H 2 CO3. Small amounts of the stronger nitric acid, HNO 3 , may be produced during lightning storms as N 2 is changed to NO and then to NO 2 , that reacts with water to form nitric acid.] (See Industrial Inorganic Chemistry, Demonstration 2.)
6. What role has photosynthesis played in creating the reservoir of oxygen in the earth’s present atmosphere? O 2 is released by plants as they undergo photosynthesis. This oxygen, in turn, should be used during respiration to oxidize carbohydrates back to CO 2 , which is then taken in by plants for photosynthesis (with a release of more O 2 ), completing the cycle. If carbon is trapped in an oxygen-poor environment, as it was in coal and petroleum formation, the oxygen released originally is not used in respiration, causing the reservoir (surplus) of O 2 to build up in the atmosphere.
7. Why is the “greenhouse effect” somewhat inaccurate in describing atmospheric warming as certain gases build up in the atmosphere? In a true greenhouse the heating effect is caused by the fact that a barrier (glass) will not allow the heat to be transferred by convection. In the atmosphere, there is absorption of certain infrared frequencies, which heat the atmosphere. There is no physical barrier preventing convection of the “heated atmosphere.”
8. What role does CO 2 play in the transfer of insoluble limestone (CaCO 3 ) under-ground in order to form stalactites and stalagmites?
A. CO 2 (g, excess) + H 2 O(l) ® H 2 CO 3 (aq)
B. H 2 CO 3 (aq) + CaCO 3 (s) ® Ca(HCO 3 ) 2 (aq)
C. Ca(HCO 3 ) 2 (aq) ® CaCO 3 (s) + H 2 O(l) + CO 2 (aq)
In Equation A, an excess of CO 2 dissolves in rainwater forming carbonic acid. This weak acid reacts with insoluble limestone above or in the ground to form soluble calcium bicarbonate or “temporary” hard water, shown in Equation B. The Ca(HCO 3 ) 2 (aq) moves through the soil in various ways, sometimes as droplets hanging from the ceiling of a cave. As conditions favor drying, the droplet evaporates, shifting the equilibrium and leaving behind a layer of limestone (CaCO 3 ) on the cave ceiling or floor and releasing carbon dioxide and water as shown in Equation C. 9. “The nitrogen cycle is dependent at each step upon microorganisms and could not proceed without them.” Support or refute this statement. [Bacteria are necessary in each of the four steps in the nitrogen cycle. N 2 is initially “fixed” in the form of ammonia (NH 3 ) or ammonium compounds by soil bacteria, such as Rhizobium japonicum, or certain blue green algae. In Step 2 of the cycle, nitrification occurs, where chemosynthetic bacteria oxidize these ammonia based compounds into nitrites and nitrates. Plants use the latter to produce amino acids, which are used by animals to produce animal protein. At death, both plants and animals are decomposed by bacteria, releasing ammonia (ammonium compounds) to the soil. In the final step of the cycle, anaerobic bacteria break down nitrates and ammonium compounds to free nitrogen and the cycle begins again.
Counterintuitive Examples
1. The solubility of ammonia decreases (rather than increases) as the temperature of the water increases. Unlike solidsdissolved in liquids, gases normally become more soluble in water as the temperature drops.
2. When burning a candle over limewater, the rise of limewater into the bottle is mistakenly attributed to the removal of O 2 from the air inside the bottle by the burning candle. In fact, limewater enters the bottle because of the removal of some CO 2 produced by the burning candle as it reacts with limewater and the contraction of CO 2 as it cools. This volume decrease creates less pressure inside the bottle, allowing atmospheric pressure to force additional limewater into the bottle.
Metaphors and Analogies
1. “Spaceship Earth.” Like a spaceship, the earth is finite in size and capacity, has limited oxygen, water, and food supply aboard, and does not possess inexhaustible energy supply.
2. “Greenhouse Effect.” In a greenhouse, glass transmits visible light and absorbs infrared radiation from the sun. This is changed to heat and prevented from leaving the greenhouse because it cannot penetrate the glass. Thus, the inside temperature of the greenhouse increases. (The same thing happens on a hot day when the windows of an automobile are up.) In the “atmospheric” greenhouse, there is no glass to absorb infrared rays, but there is a layer of several gases (primarily CO 2 and H 2 O) that serve the same purpose. These gases trap heat and warm the earth without a physical barrier present.
3. Propose a familiar cycle (water in an automobile cooling system, etc.). Ask students to suggest things that might interfere with this cycle. Extend this analogy to cycles of carbon, oxygen, and nitrogen.
4. Consider the human circulatory system as similar to a biogeochemical cycle. What happens when an artery that serves a certain organ (heart, for example) is blocked? [Some damage to the heart may result, and an alternate pathway develops due to collateral circulation.] Compare this to events that might interfere with normal cycles of oxygen, carbon, and nitrogen.