Microscale Gas Chemistry:

Experiments with Nitrogen Oxides

     Link to nitrogen oxides data page including physical properties.

 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.

Toxicity.
    Nitrogen dioxide has an irritating odor and is a poisonous gas.  Concentrations of 100 ppm are dangerous.  Concentrations greater than 200 ppm may be fatal.  To put this in perspective, if the contents of one entire syringe of NO2 (60 mL) were discharged into a volume of 1 m3, the concentration of NO2 would be 60 ppm.  Exercise caution when working with poisonous gases and vacate areas that are contaminated with unintentional discharges of gas.

Suitability.
    All of these experiments are suited for use as classroom demonstrations.  Experiment 10 is ideally suited for use as a classroom demonstration using an overhead projector.  Except for Experiment 8, which uses dangerous cryogenic materials, these experiments are also well-suited as laboratory experiments conducted by small groups of students.

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.)
 

Equipment. (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.)

Chemicals.
  • 0.25 g NaNO2(s)
  • 3 - 5 mL acidic ferrous sulfate solution, prepared as a stock solution using:
    • 13.5 g FeSO4.7 H2O
    • 12 mL 6 M H2SO4
    • 28 mL distilled water
  • 1 M NaOH (4 g NaOH per 100 mL water)
  • O2 is required to convert NO into NO2
  • These quantities of reagents will produce approximately 60 mL of pure NO.  The production of NO is relatively fast and it typically takes about a 30 seconds to fill a syringe with the gas.  The reaction is:

    2 NaNO2(s) + 2 FeSO4(aq) + 3 H2SO4(aq) ---> 2 NO(g) + Fe2(SO4)3(aq) + 2 NaHSO4(aq) + 2 H2O

     The NO gas samples used in these experiments are generated as described in Method A.

    Preparation of Acidic Ferrous Sulfate Solution.
        To prepare enough acidic ferrous sulfate solution to perform the syringe reaction ten times, add 13.5 g FeSO4.7 H2O to 28 mL distilled water.  Add 12 mL 6 M H2SO4 and stir for a few minutes until all of the solid is dissolved.  This produces a solution that is 1.8 M H2SO4 and 1.2 M Fe+2(aq).  (Alternatively, dissolve the ferrous sulfate in 36 mL water and then add 4 mL concentrated H2SO4.)

    Preparation of Neutralization Solution.
        Prepare 100 mL of 1 M NaOH (4 g NaOH in H2O to make 100 mL) in a 250 mL flask.  Keep the flask stoppered when not in use.  Label the flask ‘1 M NaOH for neutralization.’  This solution will be used to neutralized excess excess reagents in the experiments.

    Preparation of Nitric Oxide.
        Each NO(g) preparation uses 0.25 g solid NaNO2 and a minimum of 3 mL of the acidic ferrous sulfate solution.  Upon mixing the reagents in the syringe with vigorous shaking, gaseous NO is quickly produced.  A trace of reddish NO2 is observed at first but soon disappears.  The aqueous solution turns black in colour.  The volume of NO produced is approximately 60 mL.  Care must be taken to stop the gas generation after the syringe is full.  This is done by removing the latex syringe cap while it is directed upwards.  Rotate the syringe 180o in order to discharge the reaction mixture and then recap the syringe.

     Washing the Gases.
        The gas-filled syringe must be "washed" in order to remove traces of unwanted chemicals from the inside surfaces of the syringe before the gases can be used in experiments.  To do this, suction 5 mL distilled water into the syringe without discharging any gas, cap the syringe and shake the water to dissolve the contaminants on the inside of the syringe.  Remove the cap and discharge the water but not any of the gas.  Repeat once or twice.

     Disposal.
        Unwanted NOx(g) can be destroyed by slowly bubbling it through the Neutralization Solution (1 M NaOH.)
     
     


    Experiments with Nitrogen Oxides

    Experiments 1 - 4. Well-Plate Reactions with Nitric Oxide.
    Equipment:

    Chemicals:


        The following reactions are performed in a 12-or 24-well plate.  Prepare the following reagents in separate wells before generating NO(g).  Place the well plate containing the solutions inside a 1-gallon (4-L) sealable plastic bag.
     

    Well:
    #1 contains faintly pink KMnO4(aq)
    #2 contains Br2(aq)
    #3 contains Fe+2(aq); (2 mL water + 2 - 3 small crystals FeSO4.7 H2O)
    #4 contains I-(aq); (2 mL water + approximately 0.1 g KI)
    #5 Rinse water; > 3/4  full
    #6 - 8 Neutralization Solution (1 M NaOH)  Fill these three wells > 3/4  full


    One syringeful of NO(g) should be adequate to complete all four of these experiments.  Generate NO(g) using the Fe+2/H2SO4 solution as described above.  Wash the gas thoroughly.  Replace the latex syringe cap with a 15 cm length of latex tubing.  Place the syringe into the plastic bag and seal the bag.  Into each well slowly discharge enough NO(g) through the solution to achieve the desired results.  Rinse the latex tubing before going on to the next well.

        At the end of these experiments, allow the bag to stand overnight.  The NO2 produced from the reaction of NO with air will all react with the Neutralization Solution in Wells 6 - 8.  Open the bag and discard the contents to the Neutralization Flask.
     

    Experiment 1.  Nitric oxide reacts with KMnO4.
        Nitric oxide reacts with KMnO4(aq) to produce either colorless Mn+2(aq) or a brown MnO2 precipitate, depending on the conditions.  The NO(g) is oxidized to nitrate ion:

    5 NO(g) + 3 MnO4-(aq) + 4 H+(aq)  5 NO3-(aq) + 3 Mn+2(aq) + 2 H2O Eo = + 0.55 v
    or
    NO(g) + MnO4-(aq)  NO3-(aq) + MnO2(s)  Eo = + 0.72 v






    Experiment 2.  NO(g) reduces Br2.
        NO(g) reduces Br2(aq) to colorless bromide:

    2 NO(g) + 3 Br2(aq) + 4 H2 2 NO3-(aq) + 6 Br-(aq) + 8 H+(aq) Eo= + 0.11 v






    Experiment 3.  NO(g) reacts with Fe+2(aq).
        NO(g) reacts with Fe+2(aq) to produce the “brown ring test” compound, [Fe(H2O)5(NO)]+2(aq):

    NO(g) + [Fe(H2O)6]+2(aq)  [Fe(H2O)5(NO)]+2(aq) + H2O(l)

    The [Fe(H2O)5(NO)]+2(aq) appears as a brown solution that darkens to a green hue and eventually precipitates, presumably as a ferric compound.
     

    Experiment 4.  Nitric oxide oxidizes I-(aq).
        Nitric oxide oxidizes I-(aq) to produce a yellow solution of I2(aq).  The nitric oxide is presumably reduced to N2O(g) because no bubbles are observed and N2O(g) is quite soluble in water.  If N2, the other reasonable reduction product, had been produced, bubbles would have been evident.  (It is also known that mild reducing agents such as SO2(g) reduce NO to N2O.)

    2 NO(g) + 2 I-(aq) + 2 H+(aq)  N2O(aq) + I2(aq) + H2O(l)







    Experiment 5. Qualitative Conversion of Nitric Oxide to Nitrogen Dioxide.
    Equipment:

    Chemicals:     Prepare 40 mL NO(g) from 0.17 g solid NaNO2 and 3 - 4 mL of the acidic ferrous sulfate solution as described above.  Wash the gas several times.  In front of a white piece of paper, discharge approximately 5 - 10 mL of the gas which will immediately form red NO2(g):

    2 NO(g) + O2(g)  2 NO2(g)

    Next, pull about 20 mL air into the syringe.  Oxygen from the air will immediately react with NO(g) producing reddish NO2(g).

        Discard the gas by slowly bubbling it through the Neutralization Solution.
     
     
     
    Experiment 6. Quantitative Conversion of Nitric Oxide to Nitrogen Dioxide.

    Equipment:

    • Clean dry syringe
    • 3-cm length of latex tubing
    Chemicals:
    • NO(g), 60-mL
    • O2(g), 30-mL
    • Neutralization Solution

    Figure 1








     

        Prepare 60 mL of O2(g) in a syringe (See: Instructions for preparation of oxygen).  Wash the gas as described.  Set the O2-syringe aside for use later.  Prepare 60 mL NO(g) from 0.25 g solid NaNO2 and 3 - 4 mL of the acidic ferrous sulfate solution as described above.  Wash the NO(g) several times.  Transfer the NO(g) to a clean, dry syringe via a short length of latex tubing as shown in Figure 1.  Simultaneously push and pull on the respective plungers.  (This step is necessary because  water readily dissolves NO2.)  Next, connect the new NO(g)-filled syringe with the O2(g)-filled syringe via a short length of clean, dry latex tubing.  By pushing in on the plunger of the O2-syringe, slowly transfer  a volume of O2(g) equal to half of the volume on the NO(g). Thus 60 mL NO would require 30 mL O2.  The balanced reaction is:

    2 NO(g) + O2(g)  2 NO2(g)
    DH = -56.4 kJ

    Note that the volume of NO2(g) is equal to the initial volume of NO(g) as required by the law of combining volumes.  Also note that the NO/NO2 syringe becomes noticeably warmer during the reaction.   Important: Transfer O2 (g) to NO(g), not the other way around!

         The NO2(g) produced can be used in one of several of the next experiments.  Unwanted NO2(g) can be destroyed by slowly bubbling it through the Neutralization Solution.
     

    Experiment 7. Relative Water “Solubilities” of NO and NO2.
    Equipment:

    Chemicals:     Prepare 60 mL of O2(g) in a syringe (See: Instructions for preparation of oxygen).  Wash the O2(g) as described and set the syringe aside for use later.  Prepare 60 mL NO(g) as described above.  Wash the gas thoroughly.  Half-fill a 400 mL beaker with distilled water.  Remove the syringe cap and suction 5 mL distilled water into the syringe.  Keeping the syringe’s LuerLOK fitting under the surface of the water, carefully shake the water and NO(g) in the syringe.  Because NO(g) is not very soluble or reactive with water, additional water is not suctioned into the syringe by this activity.  Discharge the water and add a half-volume equivalent of O2(g) as described in Experiment 6.  As you did with NO(g), suction 5 mL distilled water into the syringe.  While keeping the syringe’s LuerLOK fitting under the surface of the water, carefully shake the water and NO2(g) in the syringe.  Unlike NO(g), NO2(g) is reactive with water, producing nitric acid.  Additional water is suctioned into the syringe as the red colored gas disappears.  For added effect, use some Universal Indicator in the water.
     
     

    Experiment 8. Dinitrogen Trioxide is a Blue Liquid.
    Equipment:

    Chemicals:     Either liquid nitrogen or a dry ice/alcohol bath is needed as a source of extreme cold for this experiment.   Prepare a dry ice/propanol bath in a 400 mL beaker:  Add 250 mL 2-propanol (ordinary rubbing alcohol) to the beaker and slowly add small chunks of dry ice until the dry ice persists in the solution.  A ice/salt bath at a temperature below -10 oC will also work, although not as well.

        Dinitrogen trioxide is a blue liquid produced by the combination of equal quantities of NO and NO2:

    NO(g) + NO2(g)  N2O3(l) mp = -102oC

      Dinitrogen trioxide partially disproportionates into NO and NO2 at temperatures above 4 oC according to the equilibrium:

    N2O3(l)  NO(g) + NO2(g)

    The disproportionation is reversible and becomes extensive as the temperature increases.

         The preparation of N2O3is conveniently accomplished in one step by combination of 4 volumes of NO(g) with 1 volume of  O2(g) and cooling to low temperatures:

    4 NO(g) + O2(g)  2 N2O3(l)

    To do this, modify Experiment 6 as follows:  Prepare a syringe of O2(g) and set it aside.  In another syringe, prepare 60 mL NO(g) from 0.25 g solid NaNO2 and 3 - 4 mL of the acidic ferrous sulfate solution.  Wash the NO(g) several times and transfer it to a clean, dry sringe.  Connect the NO(g) and O2(g) filled syringes with a short length of latex tubing as shown in Figure 1.  Slowly transfer a volume of O2(g) equal to one fourth the volume on the NO(g).  Thus, 60 mL NO would require 15 mL O2.

         The volume of the N2O3(g) expected from the above chemical equation is half that of the initial volume of NO(g): 4 moles of NO produces 2 moles of N2O3.  Even though the N2O3(g) is actually in an equilibrium mixture with NO(g) + NO2(g), the volume of the products is less than that of the initial NO(g).  As the reaction takes place, you should note that the latex tubing collapses, indicating that the pressure within the assembly of the two joined syringes is less than the external pressure.  Usually the plungers will move inward as well.

         Place the capped syringe of N2O3 (= NO2 + NO) into the cold bath or liquid nitrogen to a depth of about 2 - 3 cm (to the 15-mL mark on the syringe.)   Allow the syringe to remain in the cold until you notice that the plunger is beginning to move inward.  Droplets of blue liquid or solid N2O3 will appear.  Allow the syringe to warm to room temperature.  Reddish NO2(g) will reappear.

        The NOx(g) produced can be destroyed by slowly bubbling it through the Neutralization Solution.
     
     

    Experiment 9. Acidic Nature of Nitrogen Oxides.
    Equipment:

    Chemicals:     Nitrogen dioxide, NO2(g) is the acid anhydride of nitric acid.  The following disproportionation reaction with water is instantaneous and is the final step in the Ostwald Process.

    3 NO2(g) + H2O(l)  2 HNO3(aq) + NO(g)

    Prepare a solution of 100 mL distilled water and 2 mL Universal Indicator in a 250 mL beaker.  Add a trace of NH3(g) (one or more pipetfuls of NH3 vapors taken from the head space above an ammonium hydroxide solution is bubbled through the indicator solution.)  Remove the latex cap and fit the syringe with a 15 cm length of latex tubing.  Place cup and syringe in a plastic bag and seal shut.  Slowly dispense the NO(g) just above the surface of the dilute NH3(aq).  The NO(g) will react with O2(g) from the air and then settle on the surface of the water where it reacts making nitric acid and NO(g) according to the above equation.  The NO(g) reacts with O2(g) from the air and the cycle continues.  The solution becomes acidic near the surface creating layers of colors.

        At the end of these experiments, allow the bag to stand overnight.  The NO2 produced from the reaction of NO with air will all react with the large amount of water.  Open the bag and discard the contents down the drain.
     
     
     
    Experiment 10. Acid Rain Microchemistry
     Equipment:
    • 24-well plate
    • 1-gallon (4-L) sealable plastic food storage bag
    • plastic cup, 9-ounce (250 mL)
    • overhead projector
    • plastic disposable pipet
    Chemicals:
    • NO or NO2(g), 60-mL
    • 5 mL of Universal Indicator
    • 0.1 g of sodium bicarbonate, NaHCO3
       Automobiles produce nitrogen oxides which act to produce acid rain.  In this experiment, a 24-well plate is used to create a series of lakes, six of which are buffered.  The 24-well plate is enclosed in a Zip-Lock bag to create an ecosystem.  The layout of a typical ecosystem is shown in Figure 2.  The "B" marks indicate the six lakes that will be buffered.  Mix 100 mL distilled water with 5 mL of Universal Indicator.  Fill all 18 of the unlabeled wells with this solution.  If this experiment is to be used as a classroom demonstration using the overhead projector, fill the wells so they are slightly overfilled as shown in Figure 3.  Use a Beral pipet to add the final drops to each well.  To the remaining Universal Indicator solution, dissolve 0.1 g of sodium bicarbonate, NaHCO3.  Fill the remaining six lakes with this solution.  Place a 6-cm length of a plastic pipet stem between the four middle wells in order to prop up the ZipLock bag above the surface of the filled wells. 

    Figure 2


    Figure 3 Side views of an underfilled well (left) and a properly filled will (right)


    Figure 4(a)


    Figure 4(b)

        Generate 60 mL NO2 as per Experiment 6.  Next, slip the filled well plate into a Zip-lock bag as shown in Figure 4.  If a smaller bag is used, pierce a small hole through the bag with a sharp pencil and work the latex tubing through the hole as shown in Figure 4(a). (Moistening the tubing with alcohol helps to facilitate this process.)  If a larger bag is used, the syringe can be placed inside as shown in Figure 4(b). Zip the bag shut.  Place the assembly on the overhead projector.  Discharge the gas into the bag.  As the gas drifts across the "landscape," the unbuffered lakes will become acidic.  The buffered lakes will eventually become acidified as well.  The entire acidification process takes 1-2 minutes for the unbuffered “lakes” and 10 minutes for the buffered ones.

        At the end of these experiments, allow the bag to stand overnight.  The NO2 produced from the reaction of NO with air will all react with the large amount of water.  Open the bag and discard the contents down the drain.
     
    Figures at right:
    Top right: NO2-filled syringe just before injection into the bag. The 4 dark colored 'lakes' are buffered and the rest are distilled water.  Middle right: NO2atmosphere moves across the 'lakes'.  Right bottom: Within minutes, the non-buffered 'lakes' become reddish (Universal indicator's color at low pH) while the buffered 'lakes' remain unchanged. 

    Below Top: Prior to reaction (same as top right photo but not in bag)

    Below Bottom: After the reaction and removal from the bag. 


     
     
     

     

     


     

    Experiment 11. LeChatelier Principle.
    Equipment:

    Chemicals:     LeChatelier Principle and the NO2/N2O4 Equilibrium.  Nitrogen dioxide is a brown-red gas that exists in equilibrium with its dimer, N2O4(l):

    2 NO2(g)  N2O4(l)     Kc25 = 215  DH = -57 kJ

    The dimer consists of two monomers joined by a weak nitrogen-nitrogen bond.  The equilibrium can be shifted right by cooling in an ice bath.  At room temperature appreciable dissociation exists.

        Prepare 30 mL NO2(g) using the same conditions as in Experiment 31 except that only 0.15 g NaNO2 is used.  Wash the gas several times.  Remember to scale the reaction mixture accordingly.  It is important to use a 2:1 volume ratio of NO2:O2.  Equip the syringe with a latex syringe cap.  The LeChatelier principle provides that an increase in volume will favor a shift to the left in the above equilibrium.  If this is so, the gas sample should become more reddish, which is counter-intuitive because increasing the volume usually means that the concentration has been “diluted.”  LeChatelier’s principle also predicts that the red color at equilibrium will fade as the syringe volume is decreased as NO2 shifts to make N2O2.  To test these hypotheses, quickly pull the syringe barrel outward to the 60-mL mark and hold it in that position.  The NO2(g) will initially fade due to the decrease in concentration but within a few seconds the red color will intensify due to formation of more NO2 as predicted by the LeChatelier principle.  Next, push the syringe plunger down rapidly and firmly and hold it in that position.  You should observe an instantaneous change which is associated with the new pre-equilibrium conditions and then within a second you will see the results of the new equilibrium.

        The NO2(g) produced can be used in one of several of the next experiments.  Unwanted NO2(g) can be destroyed by slowly bubbling it through the Neutralization Solution.
     
     


    Clean-up and Storage.

      At the end of the experiments, wipe excess lubricant off of rubber diaphragm. Clean all syringe parts (including the diaphragm), caps and tubing with soap and water.  Use plenty of soap to remove oil from the rubber seal.  This extends the life of the plunger.  It may be necessary to use a 3-cm diameter brush to clean the inside of the barrel.  Rinse all parts with distilled water.  Be careful with the small parts because they can easily be lost down the drain. Important: Store plunger out of barrel.



     
     
    This article first appeared in Chem13 News in January, 1997.  The authors of the original Chem13 article are: 

    Bruce Mattson, faculty member, principal investigator,  Department of Chemistry, Creighton University, Omaha, Nebraska 68178 USA

    Joseph Lannan, Blair High School, Blair, NE

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    (This page last updated 29 January 2002)