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Lesson 6. It's a Gas!


Week Two Lessons
6. It's a Gas!
7. How Heavy are They?
8. When You're Hot, You're Hot
9. Watt's Up?
10. Eating Up Energy

Lab producing carbon dioxide and methane | Physical Science (Chemistry)

Links on this page: It’s a Gas!-CO2 Background Information | It’s a Gas!-CO2 Lab Procedure| It’s a Gas!-CO2 Teacher Answer Key | It’s a Gas!-CO2 Student Sheet | It’s a Gas!-CH4 Background Information | It’s a Gas!-Part Two
CH4 Student Lab Procedure
| It’s a Gas!-CH4 Student Sheet |It’s a Gas!-CH4 Teacher Answer Key | It’s a Gas!-Methane Extension Activities

 

National Education Standards Met:

science

 

Goal: Students will investigate the properties of carbon dioxide and methane in order to understand their role in global climate change.

Objectives: Students will…

  • Understand the difference between CO2 and methane
  • Examine the chemical properties of CO2 and methane
  • Use a chemical reaction to trap CO2 and methane

Materials (For class of 30):

  • 3 baby food jars with caps
  • 100 ml of vinegar
  • Sodium acetate
  • Soda lime
  • Bunsen burner
  • 1 #6, two-hole rubber stopper with plastic tubes
  • 1 250-ml flask
  • 1 length rubber tubing, 45 cm long
  • safety glasses
  • 1 250 ml beaker
  • 1 30 ml syringe (no needle)
  • 1 small plastic tub
  • supply of water
  • box of baking soda
  • 50 ml Phenol Red
  • 50 ml Limewater
  • matches
  • straws or rigid plastic tubing
  • Copies of It’s a Gas!-Student Sheet

Time Required: 2, 45-60 minute period

Standards Met: S1, S2, S5, S6

Procedure:

PREP

  • Gather all of the necessary lab materials and run a test lab to be certain of safety procedures.

DAY ONE

  • Explain to students that they will be creating CO2 and Methane (CH4 ) today in class. They will begin to see the different characteristics of each gas in order to understand how they relate to global climate change. Methane and carbon dioxide are considered greenhouse gases that are emitted through various human and natural processes. For more information on this topic, refer to http://www.epa.gov/methane/scientific.html.
  • Give students copies of the It’s a Gas!-CO2 Background Information and review.
  • Divide students into lab groups of 4.
  • Hand out the It’s a Gas!-CO2 Lab Procedure sheets to each group. Review safety precautions and lab procedures with students.
  • Allow students time to complete the lab as they follow steps on their procedure sheets. Remind students to record any observations on their It’s a Gas!- CO2 Lab Procedure sheets.
  • Review clean up procedure with students and give them time to complete a thorough clean up of their lab stations.
  • Hand out It’s a Gas!-Student Sheet and ask students to answer only the questions for Part One. If time allows, review the answers to Part One with students.
  • Prep students for part two where they will be producing methane.

DAY TWO

  • Follow the procedure above only students will produce methane instead of CO2.
  • Be sure to review safety procedures with students prior to beginning the lab.
  • Hand out the It’s a Gas!-CH4 Lab Procedure, review and allow students time to complete the lab.
  • Review clean up procedure with students and give them time to complete a thorough clean up of their lab stations.
  • Hand out It’s a Gas!-Student Sheet and ask students to answer only the questions for Part Two. Review the answers with students.
  • Discuss the implications of methane and carbon dioxide emissions and its relationship to global climate change.
  • Extension activities available.

Note:

Vinegar and baking soda reaction--

  • Vinegar is acetic acid: CH3COOH
  • Baking soda is sodium bicarbonate: NaHCO3
  • Mixing the two is simply an acid base reaction.
  • CH3COOH + NaHCO3 ---> CH3COONa + H2CO3
  • That last product is carbonic acid which quickly decomposes into carbon dioxide and water:

    -- H2CO3 ---> H2O + CO2
    -- The CO2 is what you see foaming and bubbling in this reaction.

Assessment:

  • Completed lab procedures
  • Completed student sheet

 

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It’s a Gas!-CO2 Background Information

Before CO2 gas can be sequestered from power plants or industrial sources, it must be captured as a relatively pure gas. CO2 is routinely separated and captured as a by-product from industrial processes such as synthetic ammonia production, hydrogen production, and limestone calcination.

CHEMISTRY OF LIME
Limestone vs. Lime:
In everyday usage the terms "limestone" and "lime" are used by the general public interchangeably to mean the same material, however there are some significant differences between the two materials. Limestone is a sedimentary rock whereas lime is a manmade chemical that is produced from a sufficiently pure sedimentary rock by heating it to high temperature in a kiln. This process is referred to as "calcining" the limestone.

LIMESTONE: This term refers to a naturally occurring sedimentary rock that is relatively inert, except in the presence of a strong acid. With the proper purity the rock deposit can be used to produce "lime", a manmade chemical. Most often, limestone is found in nature in a mixed form known as "dolomite", which is a blend of calcium carbonate and magnesium carbonate in varying proportions. (In the Shelby County, AL area there are large deposits of limestone, primarily composed of calcium carbonate, which are used as the "raw material" for producing high calcium lime products.)

LIME: This term refers to either "quicklime", the product that is produced by heating the limestone to its dissociation temperature, or "hydrated lime", the product that is produced by the reaction of quicklime with water. (Lime in the form of high calcium quicklime, CaO readily reacts with water to form hydrated lime, which provides a pH of up to 12.454 when in an aqueous solution. Because of elemental differences between magnesium (Mg) and calcium (Ca) the compound magnesium oxide, MgO does not readily react with water at normal temperatures and pressures.

Quicklime Production:
The production of high calcium quicklime (calcium oxide) requires a large amount of heat, which is generated in the kiln environment. The quarried and sized high calcium limestone travels through a rotary kiln and is subjected to these high temperatures where the calcium carbonate begins to dissociate with the resultant formation of calcium oxide. The minimum temperature for the dissociation of calcium carbonate is 1648°F (898°C). For practical production purposes, however, the kiln temperature range is from an initial temperature of about 1750°F (954°C) to a final temperature of about 1950°F (1066°C). These temperatures can vary dependent upon the nature of the limestone being calcined.

"High Calcium" Limestone Calcination:
CaCO3 + Heat ---> CaO + CO2
1750° to 1950°F
954° to 1066°C

"Dolomitic" Limestone Calcination:
CaCO3 • MgCO3 + Heat ---> CaO • MgO + 2CO2

Hydrated Lime Production:
High calcium quicklime readily reacts with water to form hydrated lime. The reaction is highly exothermic and the process is known as "slaking". The reaction is usually carried out in a "slaker" (a specially designed mixer) which, through a process of rigorous mixing, makes certain that all of the quicklime has come into intimate contact with water and no unreacted quicklime remains. From a general viewpoint the hydrated lime produced can be in the form of dry hydrate, putty slurry, or "milk of lime". The exothermic reactions are shown below: (There are various types of slakers available on the market.)

"High Calcium" Quicklime Hydration:
CaO + H2O ---> Ca(OH)2 + Heat

"Dolomitic" Quicklime Hydration:
CaO • MgO + H2O ---> Ca(OH)2 + MgO + Heat

Note: CaO will readily react with water under normal temperatures and pressures, whereas MgO will not. However, existing capture technologies are not cost-effective when considered in the context of CO2 sequestration.

Carbon dioxide capture is generally estimated to represent three-fourths of the total cost of a carbon capture, storage, transport, and sequestration system. The program area will pursue evolutionary improvements in existing CO2 capture systems and also explore revolutionary new capture and sequestration concepts. The most likely options currently identifiable for CO2 separation and capture include the following:

* Absorption (chemical and physical)
* Adsorption (physical and chemical)
* Low-temperature distillation
* Gas separation membranes
* Mineralization and biomineralization

Opportunities for significant cost reductions exist since very little R&D has been devoted to CO2 capture and separation technologies. Several innovative schemes have been proposed that could significantly reduce CO2 capture costs, compared to conventional processes. "One box" concepts that combine CO2 capture with deduction of criteria-pollutant emissions are concepts to be explored.

This activity will introduce students to the gas CO2 (carbon dioxide), how it is formed, and tests to tell that it is present. For this activity the students will produce carbon dioxide from a reaction between vinegar and baking soda.


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It’s a Gas!-CO2 Lab Procedure

Gather the following materials:

  • 1 30ml syringe
  • 100ml of white vinegar
  • 1 cup of baking soda
  • 1, 2-hole rubber stopper with plastic tubes
  • 100ml of tap water
  • 3 baby food jars
  • 3 matches
  • 250 ml flask
  • 1 length rubber
  • Tubing, 45 cm long
  • 1 small plastic tub
  • Supply of water
  • Baking soda to cover
  • bottom of flask
  • 50 ml Phenol Red
  • 50 ml Limewater

generator set up

 

 

  • Assemble the CO2 generator using the drawing above. Make sure that all unions are airtight.
  • Place enough baking soda in the flask to cover the bottom.
  • Pour about 40 ml of vinegar into a 250 ml beaker. Make sure you have your safety glasses on.
  • Put the tip of the 30 ml syringe into the vinegar making sure that the plunger is all the way down. Keep the tip of the syringe below the surface as you pull back on the plunger to fill it to the 30 ml mark. If you get air bubbles in the syringe, empty it, and repeat the procedure again.
  • Put the free end of the rubber tube under the water in the pan. The depth of the water should be enough to completely cover a baby food jar.
  • Place the syringe into the straw on the rubber stopper and slowly add 10 ml of vinegar to the baking soda. Do not have the tubing under the jars at this time.
  • Let the gas bubble from the rubber tube for about 30 seconds before moving on.
  • Place one of the baby food jars into the tank of water and completely fill it with water.
  • Invert it so that the top (open end) is facing down (it must still be completely filled with water…no air pockets).
  • Slip the end of the tube just under the mouth of the baby jar.
  • Slowly add more vinegar to the baking soda until the baby jar fills with gas.
  • Cap the jar tightly while it is still inverted under water.
  • Repeat the procedure with the other two baby jars.
  • Strike a match and quickly add it to one of the jars. Observe what happens. Record your observations below:







  • Add about 10 ml of limewater to one of the jars, cap it quickly and shake.
  • Observe what happens.





  • Add about 10 ml of phenol red to one of the jars, cap it quickly and shake.
  • Observe what happens.





 

  • When you have finished this activity, your teacher will tell you how to clean up your materials and then complete It’s a Gas!-CO2 Student Sheet.

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It’s a Gas!-CO2 Teacher Answer Key

Name: ___________________________________
Date: __________________________________

 

  1. Where did the carbon dioxide that you collected in the baby food jars come from?

    It was produced from the reaction between the vinegar and the baking soda.

  2. Why did the gas push out the water in the baby food jars? Isn’t the water denser than the gas?

    The water is denser than the gas, but the gas built up pressure in the container greater than the pressure of the water in the baby food jars, forcing the water out of the baby food jars.

  3. Why did you let your apparatus bubble for 30 seconds before you began collecting gas in the jars?

    It is important to let the gas bubble for a while to ensure that all of the gas that was in the jar (air) has been removed, and the gas that is collected is entirely a product of the reaction between the vinegar and baking soda.

  4. Does carbon dioxide gas have a color? An odor?

    Carbon dioxide is a colorless, odorless gas.

  5. How can you test for the presence of CO2? Give at least three ways.

    Tests that can be used to confirm the presence of carbon dioxide are: will extinguish a lit match, turns limewater cloudy, and turns phenol red yellow.

  6. How does CO2 differ from normal air?

    “Normal” air is a mixture of several gases (78.084% nitrogen, 20.947% oxygen, 0.934% Argon, 0.033% carbon dioxide and several trace gases). Carbon dioxide is a pure gas produced as a product of respiration and combustion.

  7. If you were to exhale into the rubber tubing and collect the gas in jars, would the tests you performed above have the same results? Explain. (You may want to ask your instructor to try this experiment if time permits.)

    Yes, as a person exhales, a certain percentage of the gas that is emitted from their lungs is carbon dioxide (a product of cellular respiration) and this would give similar result to the experiment. Note: The concentration of the gas (carbon dioxide) emitted from the experiment is greater that what a person would normally exhale.

  8. Write a simple chemical equation for the experiment you did in this activity.

    Vinegar + baking soda carbon dioxide

  9. Why didn’t we produce carbon dioxide by using limestone calcination?

    Limestone calcination requires a great deal of heat (the use of a kiln) and is not feasible for the production of carbon dioxide in the classroom.

 

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It’s a Gas!-CO2 Student Sheet

Name: _____________________________________
Date: _________________________________

 

  1. Where did the carbon dioxide that you collected in the baby food jars come from?


  2. Why did the gas push out the water in the baby food jars? Isn’t the water denser than the gas?


  3. Why did you let your apparatus bubble for 30 seconds before you began collecting gas in the jars?


  4. Does carbon dioxide gas have a color? An odor?


  5. How can you test for the presence of CO2? Give at least three ways.


  6. How does CO2 differ from normal air?


  7. If you were to exhale into the rubber tubing and collect the gas in jars, would the tests you performed above have the same results? Explain. (You may want to ask your instructor to try this experiment if time permits.)


  8. Write a simple chemical equation for the experiment you did in this activity.


  9. Why didn’t we produce carbon dioxide by using limestone calcination?

 

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It’s a Gas!-CH4 Background Information

Methane, CH4, is a colorless, odorless, flammable gas that burns with a faintly blue flame. It is a principal component of natural gas. "Natural gas," used by many North Americans for heating and cooking, is primarily methane (>90%). The 'gas odor' is from an added substance so that gas leaks can be detected. Pure methane has no odor. Methane is relatively non-toxic; it is a simple asphyxiant. It is flammable in air and forms explosive mixtures with air. Explosive mixtures of methane with air contain between 5 - 14 % methane. Mixtures containing more than 14% burn without explosion. The combustion of methane is highly exothermic.

Natural gas occurs in reservoirs beneath the surface of the Earth. It is often found in conjunction with petroleum deposits. Before it is distributed, natural gas usually undergoes some sort of processing. Usually the heavier hydrocarbons (propane and butane) are removed and marketed separately. Non-hydrocarbon gases, such as hydrogen sulfide, must also be removed. Other places Methane is found is by sources such as swamps, livestock, landfills, and some methane is manufactured by the distillation of coal.

Methane is also formed and released to the atmosphere by biological processes occurring in anaerobic environments. Once in the atmosphere, methane absorbs terrestrial infrared radiation that would otherwise escape to space. This property can contribute to the warming of the atmosphere, which is why methane is a greenhouse gas. Methane is about 21 times more powerful at warming the atmosphere than carbon dioxide by weight. It also increases the amount of water vapor in the stratosphere which also traps heat.

See the U.S. Environmental Protection Agency’s website for more information:

http://www.epa.gov/methane/scientific.html

Laboratory preparation of methane gas
In the lab methane is prepared using the following chemical reaction.


methane

 

Sodium acetate is the sodium salt of acetic acid. Sodium hydroxide (NaOH) is obtained from soda lime, which is a combination of NaOH + CaO. Only the NaOH component of the soda lime reacts in the reaction.

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It’s a Gas!-Part Two
CH4 Student Lab Procedure

Gather the Following Materials: (This equipment can be ordered from a variety of vendors including Educational Innovations, Flinn Scientific (US sales only), Micro Mole, and Fisher Scientific. Part numbers and links to their websites are provided.)

several 60-mL plastic syringes with a LuerLOK fitting

Syringe Lubrication:
We recommend lubricating the black rubber diaphragm of the plunger with silicone spray (available from hardware stores) or medium-grade silicone oil (Educational Innovations, $5.95 Part #GAS-150; Fisher Catalog Number S159-500; $34/500 mL.)

  • Latex LuerLOK syringe cap fittings
  • Small plastic weighing boats
  • balance capable of measuring to 0.01 g
  • two pieces, latex tubing, 1/8-inch (3.175 mm) ID, 5 cm lengths
  • two 18 x 150 mm test tubes
  • two-hole #1 stopper fitted with two short lengths (2 cm) of glass tubing
  • pinch clamp or hemostat
  • ring stand and three suitable clamps to hold test tube and syringes
  • small Bunsen burner
  • matches or a lighter
  • ‘permanent’ marker pen

Chemicals.

sodium hydroxide, NaOH

sodium acetate, NaC2H3O2

apparatus

 

 

 

 

 

 

 

Figure 1. Apparatus

 

General Safety Precautions: Always wear safety glasses! Gases in syringes may be under pressure and could spray liquid chemicals. Follow the instructions and only use the quantities suggested!

Procedure:

  • Assemble apparatus as shown in Figure 1. Don’t forget safety goggles!
    NOTE: Syringe plungers should move easily in the barrels. This can be facilitated by applying a drop or two of oil (silicone, vegetable, or glycerin) to the groove in the plunger's rubber seal. A small burner is also needed. Clamp off the syringe labeled CH4.
  • Place a 4-g mixture of 50% (by mass) sodium hydroxide and 50% sodium acetate in the test tube. This will act as the “reagent mixture”.
  • Insert the stopper firmly in order to form an air-tight seal. CAUTION: Do not crimp the latex tubing!
  • Gently begin to heat the test tube using the cool part of the flame. Observe the vapors or fumes. Record your observations below.





  • Continue to heat the test tube until a steady rate of methane gas begins to rise as white fumes or vapor. The plunger of the “Waste” syringe should begin to move. It may be necessary to assist the sliding movement of the plunger up the barrel of the syringe during the reaction.
  • Allow the “waste” syringe to collect gas until the 15-30 mark is reached.
  • Switch clamps so that “waste” is closed off and “Methane” Syringe will begin to move and fill with pure methane. See Step 1 in Figure 2.
  • Collect 50-55 ml of methane in the syringe.
  • Switch to a new clean and dry syringe (2nd) and cap the filled syringe with a latex syringe cap. (DO NOT CLAMP BOTH TUBES AT THE SAME TIME!!) See Step 2 in Figure 2.
  • If you intend to collect multiple syringefuls of methane, replace the Waste syringe with another clean, dry syringe while you are waiting for the gas to accumulate during this step.
  • To fill more syringes switch the “waste” syringe with a fresh empty syringe to collect a 3rd syringe of methane.
  • Switch the pinch clamp back to the tubing connected to the CH4-syringe and remove the heat source. WARNING: Never simultaneously pinch both latex tubes! See Step 3 in Figure 2.
  • It is possible (and probably desirable) to replace the CH4(g) syringe with a clean dry syringe, and repeat steps; numerous syringefuls of methane can be collected. We have collected at least five syringefuls before stopping the reaction.
  • Allow apparatus to cool.
  • Follow the instructions from your teacher for clean-up and disposal of materials then complete It's a Gas! - CH4 Student Sheet

 

step 1 step 2 step 3

Figure 2. Three step procedure for generating methane.

 

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It’s a Gas!-CH4 Student Sheet

 

Name_______________________________________

Date________________________________________

  1. What gases are contained in the “waste” syringe?




  2. What was the purpose of the “waste” syringe?





  3. Why is methane considered a Greenhouse Gas and what is its effect on our climate?





  4. Does methane gas have a color? An odor?





  5. Research two possible ways to reduce methane gas in our atmosphere using the Internet and list below. (Be sure to include the web address with your recommendations.)

 


It’s a Gas!-CH4 Teacher Answer Key

 

  1. What gases are contained in the “waste” syringe?

    The waste syringe has a combination of oxygen, nitrogen and other atmospheric gases along with the start of the first of the methane produced from the reactions.

  2. What was the purpose of the “waste” syringe?

    The reason for removing the first syringe of gas generated was to make sure that the second and any syringe collected there after would be mostly methane produced from the reaction mixture of Sodium acetate and sodium hydroxide.

  3. Why is methane considered a Greenhouse Gas and what is its effect on our climate?

    Methane is a naturally produced gas from the breakdown of many organic and biological systems. Because this gas is so abundant as a byproduct, one of the major effects is the accumulation of this gas in our atmosphere in alarming quantities. The combined pollutants cause an invisible barrier that does not let the heat of the sun to escape. The unfortunate result is global warming.

  4. Does methane gas have a color? An odor?

    Methane in its natural state is both odorless and colorless. This makes this gas extremely dangerous. Coal miners needed canaries to act as their early warning system to signal a methane leak. Methane piped into our homes as a fuel to be burned, has been mixed with a sweet smelling organic “tag” that can warn us of a leak in our home.

  5. Research two possible ways to reduce methane gas in our atmosphere using the Internet and list below. (Be sure to include the web address with your recommendations.)
    There are a number of resources available on the web. Dept of Natural Resources website and website including the EPA, listed earlier, would be very helpful. Proposed limits by congress to reduce the production of methane. One controversial method is to capture the methane in pools either by pumping the methane under ground or under water for later use as a possible fuel source.

 

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It’s a Gas!-Methane Extension Activities

Experiments with Methane.


Most of the experiments given below involve the combustion of methane. The reaction is:
CH4(g) + 2 O2(g) -----> CO2(g) + 2 H2O(g) H = -802.3 kJ

Experiment 1. Products of Combustion.
Equipment:

  • length of latex tubing
  • glass pipet (the tubing fits snugly inside the pipet)
  • screw clamp
  • 125-mL flask
  • matches or a lighter

Chemicals:

CH4(g), 60-mL

Limewater (See: How to prepare and dispense limewater)

Generate a syringeful of methane. Equip the syringe with a length of latex tubing, a glass pipet (the tubing fits snugly inside the pipet), and the screw clamp. Tighten the screw clamp to completely seal the tubing. Using a ring stand and a suitable clamp, clamp the glass pipet in the approximate position shown in Figure 3. Two people are needed for the next part of this experiment. One person should apply continuous, gentle positive pressure on the plunger so that the methane is always under pressure. The second person should open the screw clamp just enough to allow a steady but small flow of methane. Ignite the gas issuing from the pipet. The flame should be no more than 1-cm in height. The screw clamp controls the flow of the gas and should be adjusted as necessary. Position an inverted 125-mL flask over the pipet so that flame is centered inside the flask. Water condensation on the glass will be noted and the flame will go out within seconds due to deprivation of oxygen. Remove the pipet from the flask and close the screw clamp. Test the contents of the flask for CO2(g) by adding 10-mL limewater to the flask and shaking the flask for a few seconds. A cloudy solution indicates the presence of CO2 as a result of the reaction:

Ca(OH)2(aq) + CO2(g) > CaCO3(s) + H2O(l)

 

Figure 3. Screw clamp controls gas flow rate.

Figure 3. Screw clamp controls gas flow rate.

 

Experiment 2. How a Bunsen Burner Works.
Equipment:

  • length of latex tubing
  • glass pipet (the tubing fits snugly inside the pipet)
  • screw clamp
  • matches or a lighter
  • glass tubing (approx. 10 mm inside diameter and 20 cm length)
  • aquarium air pump or a second syringe filled with air
  • ring stand and clamp


Chemicals:

CH4(g), 60-mL

The Bunsen burner works by mixing a hydrocarbon fuel such as methane with air. The principle is simple and can be demonstrated with a simple length of glass or plastic tubing. The same device shown in Figure 3 will be used in this experiment.

Clamp a piece of glass tubing in a vertical position as shown in Figure 4. A source of forced air, such as an aquarium air pump or a second syringe filled with air is optional and is used to create a hotter flame. Generate a syringeful of methane. Open the screw clamp and start the flow of methane through the 'Bunsen burner' tube by applying a continual positive pressure on the syringe plunger. Light the gas at the top of the tube. The flame will be gentle. Start the flow of air. This may blow out the flame if its flow rate is too great. Use a screw clamp on the air delivery tube to reduce the flow of air. When the methane-air mixture is optimal, the flame will be small and sharp and there will be an audible noise. Interestingly, methane prepared as described above will burn with an orange-yellow flame due to trace levels of suspended sodium salts in the gas. These can be removed by washing the methane (suction 5-mL distilled water into methane-filled syringe and shake) after which the methane burns with its characteristic blue flame.

Figure 4. A glass tube Bunsen burner
A glass tube Bunsen burner
Figure 4. A glass tube Bunsen burner  

 

 

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Experiment 3. Flame Chemistry
Equipment:

  • length of latex tubing
  • glass pipet (the tubing fits snugly inside the pipet)
  • screw clamp
  • matches or a lighter
  • glass tubing (approx. 10 mm inside diameter and 20 cm length)
  • piece of glass tubing (5 mm ID x 8 cm length)
  • ring stand and clamp


Chemicals:

CH4(g), 60-mL

Most chemistry textbooks describe the chemistry of the flame, a fascinating subject that was first investigated by Michael Faraday and described in his "The Chemical History of the Candle" lectures which he gave at the Royal Institution during the early and mid-19th century (see info at end of this experiment). Faraday demonstrated that ". . . there are clearly two different kinds of action — one of the production of the vapor, and the other the combustion of it — both of which take place in particular parts of the candle." The former is now called the pyrolysis zone, where the fuel is broken into radicals (such as H atoms and CH
3 groups) and smaller molecules including H2(g). The outer region contains air and is called the combustion zone. In this experiment we will repeat this experiment of Michael Faraday's using methane rather than a candle flame. The general set up uses the Bunsen burner shown in Figure 4. The air pump is not used for this experiment. A smaller piece of glass tubing (5 mm ID x 8 cm length) should be held by a clamp in a 45° position about 2 - 3 cm above the opening of the "Bunsen" burner as shown in Figure 5.

 


Figure 5. Siphoning off the pyrolysis zoneFigure 5. Siphoning off the pyrolysis zone


Prepare several syringefuls of methane. Two people are required to perform this experiment. One person delivers the methane through the main burner in a continuous, steady stream and ignites the gas issuing from the top. The flame should be large enough that the small tube is positioned towards the top of the flame. Gases diverted into the tube are incompletely combusted and can be ignited by the second person as they issue from the opening.

simple bunsen burner
Simple Bunsen burner Flames are yellow from traces of sodium due to reagents Washing gas with water removes sodium and methane burns blue

 


"Faraday's Chemical History of the Candle. Twenty-two Experiments and Six Classic Lectures," Chicago Review Press, Distributed by Independent Publishers Group, ISBN 1-55652-035-2. Material about the life of Michael Faraday is also available at the web site of the Royal Institution of Great Britain: http://www.ri.ac.uk/History/


Experiment 4. Burned Rings in Paper.
Equipment:

  • length of latex tubing
  • glass pipet (the tubing fits snugly inside the pipet)
  • screw clamp
  • matches or a lighter
  • glass tubing (approx. 10 mm inside diameter and 20 cm length)
  • ring stand and clamp
  • heavy-stock paper such as a note card


Chemicals:

CH4(g), 60-mL

This is another experiment described by Faraday for the candle.6 Here we will use methane and the burner (without the air pump) built in Experiment 2. CAUTION! Have a cup of water ready in case the paper used in this experiment catches on fire. While one person operates the burner and methane-filled syringe, a second person holds a piece of heavy-stock paper such as a note card positioned horizontally through the inner cone as shown in Figure 6 — approximately 2 cm above the top of the burner. Within a few seconds, the paper card will begin to burn (turn brown) in a ring. As soon as the brown ring appears, remove the card; do not allow the paper to actually ignite. This experiment reveals the fact that the pyrolysis zone is cool and the combustion zone is hot.

Figure 6. Paper starts to burn near the outside of flame

Figure 6. Paper starts to burn near the outside of the flame.

 

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Experiment 5. Window screen provides thermal insulation.
Equipment:

  • length of latex tubing
  • glass pipet (the tubing fits snugly inside the pipet)
  • screw clamp
  • matches or a lighter
  • glass tubing (approx. 10 mm inside diameter and 20 cm length)
  • window screen, 5 cm x 5 cm, 2 pieces
  • ring stand and clamp


Chemicals:

CH4(g), 60-mL

As a final experiment from Faraday's work with candles, we will investigate how a piece of window screen will affect the flame when it is held in a position similar to that of the paper card in the previous experiment. It works best to hold the screen in position 2-cm above the burner. Do not use the air pump.

Experiment A. While one person discharges the methane-filled syringe through the burner tube, a second person holds the screen and ignites the gases above the screen. Will the flame jump through the screen and start burning below?

Experiment B. While one person discharges the methane-filled syringe through the burner tube, a second person holds the screen and ignites the gases below the screen. Will the flame jump through the screen and start burning above?

Experiment C. Holding two screens 2 and 4 cm above the burner, the gases between the screens can be ignited!


Experiment A Experiment B



The screen's ability to dissipate heat and prevent combustion while allowing flammable mixtures of gases to pass through has been used in practical applications. Sir Humphrey Davy used this principle in his invention of the miner's safety lamp (Figure 7) in 1815. Flammable gases from the mine could pass through the screen and burn in the enclosed flame with a 'colored haze' while the screen prevented the open flame from causing a mine explosion.

Figure 7. Sir Humphrey Davy's Miner's Safety Lamp
Figure 7. Sir Humphrey Davy's Miner's Safety Lamp From the web site (History page) of the Royal Institution of Great Britain.


Experiment 6. Density of Methane: Lighter-than-Air Methane Bubbles.


(Based on "Spectacular Gas Density Demonstration Using Methane Bubbles", R. Snipp, B. Mattson, and W. Hardy, Journal of Chemical Education, 1981, 58, 354.)

Equipment:

  • large bulb polyethylene transfer pipet
  • scissors
  • candle in holder
  • matches or lighter

Chemicals:

CH4(g), 60-mL

3% dish soap solution

Methane is 45% lighter than air, so bubbles of the gas rise. Single bubbles of suitable size are easily generated by the device shown in Figure 8. A large bulb polyethylene transfer pipet is connected to a methane-filled syringe with a 2-cm length of latex tubing. The bulb of the pipet is cut off with a scissors.

 

 

Figure 8. A pipet used as a bubble-maker

Figure 8. A pipet used as a bubble-maker


Making the bubbles: Dip the mouth of the pipet into a 3% dish soap solution 8 a film of soap will cover the opening. Start forming the bubble while directing the pipet's mouth downward (Figure 8, rotated right) so the bubble forms below the device. This allows extra soap solution to gather at the bottom of the bubble as it is forming. While the bubble is still small, a slight shake will dislodge the extra drop which otherwise could make the bubble heavier-than-air. Quickly fill the bubble with the 60-mL gas while tilting the device to a horizontal position (Figure 8). Dislodge the bubble with an abrupt downward flick of the pipet. The bubble may rise, stay suspended in air or slowly drop depending on the amount of methane compared to the mass of the soap film. Bubbles containing 60-mL methane usually rise. The bubbles can be ignited with a candle. They will produce a fireball about 20-cm in diameter and represent about 2 kJ of heat.
USE CAUTION!

 

Experiment 7. Density of Methane: Burning Methane in a Large Test Tube.
Equipment:

  • large test tube (22 x 200 mm)
  • 250-mL beaker or 9-ounce plastic cup
  • candle in holder
  • matches or lighter


Chemicals:

CH4(g), 60-mL

Fill a large test tube with methane using water displacement. The volume of the test tube is 80-mL so two syringefuls will be necessary. Darken the room. Remove the test tube from the water and continue to hold the test tube with its mouth directed downward. Bring a burning candle up to the mouth of the test tube and the gas will begin to burn. In order to maintain the flame and burn all of the gas, the test tube must be rotated to a 45o angle position with open end up so that the lighter-than-air methane can leave the test tube. The gas will burn down the test tube in the form of a narrow, bright blue disk that produces condensation on the glass just above the flame. It takes approximately 15 seconds for the burning disk of methane to burn to the bottom of the tube. Caution: The test tube will become hot, so use a test tube clamp.

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Experiment 8. Explosive Mixture of Methane/Air.
Equipment:

  • 20-ounce (600-mL) plastic soft-drink container
  • aluminum foil, 5 cm x 5 cm
  • ring stand and clamp
  • matches or lighter

Chemicals:

CH4(g), 60-mL

Methane forms explosive mixtures with air in the 5 - 14 % range. This can be demonstrated with the device shown in Figure 9, made from a plastic soft-drink container with the bottom half cut off. Cover the opening with a small piece of aluminum foil. With a sharp pencil, poke a hole of approximately 4-mm diameter in the center of the foil. Clamp the device in the position shown in Figure 9. Set a rubber stopper or similar object over the hole for the moment.


Generate a syringeful of methane and transfer the gas to the device from the bottom. Position the syringe or tube so that most of the gas accumulates near the top of the device. Remove the object covering the hole and immediately ignite the gas. As demonstrated in the previous two experiments, methane is lighter than air and will burn with a large flame as it passes through the hole in the foil. When much of the methane has been consumed and the methane/air mixture falls to 14%, the gas mixture will explode downward into the container. The 'explosion' is quite gentle (unlike hydrogen/air!), but demonstrates an important principle. The demonstration should be repeated in a darkened room.


Figure 9. Pop bottle used for gentle explosion

Figure 9. Pop bottle used for gentle explosion

 

Experiment 9. Bubble Domes.
Equipment:

  • 250-mL beaker
  • strip of cloth
  • 10-cm length of latex tubing
  • match or lighter
  • large plastic weighing boat

Chemicals:

CH4(g), 60-mL

3% dish soap solution

Soap film domes can be made from 3% dish soap solution8 and a strip of cloth. Soak the cloth in the soap solution. Then starting from one side of a 250-mL beaker, slowly drag the cloth across the top of the beaker forming a film of soap. Without drafts, the film will remain intact for as long as a minute. Fill a syringe with methane and equip the syringe with a 10-cm length of latex tubing. Moisten the tubing with the soap solution and insert the tubing through the soap film. When moistened, the tubing will not break the film. Quickly inject the methane; it will cause the film to mound up forming a bubble as shown in Figure 10. Remove the tubing and ignite the bubble with a candle. [Hint: Sometimes an unwanted second bubble forms at the end of the latex tubing while the methane is being injected. To prevent this, initially withdraw the plunger about 5-mL in order to break the film over the end of the tubing.]

Figure 10. A partially inflated soap film dome.
Figure 10. A partially inflated soap film dome.  

 

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