Chapter 5
EXPERIMENTS WITH NITROGEN OXIDES
The chemicals required for O2 and N2 production as listed in the Master Table (back
cover) are:


0.25 g NaNO2(s)
3 –5 mL acidic ferrous sulfate solution (recipe given below)
These quantities of reagents will produce approximately 60 mL of pure NO. The
production of O2 is relatively fast and it typically takes about 30 seconds to fill a syringe
with NO. The reaction is:
2NaNO2(s) + 2 FeSO4(aq) +3 H2SO4(aq)→ 2 NO(g) + Fe2(SO4)3(aq) + 2 NaHSO4(aq) + 2H2O
The NO gas samples used in these experiments are generated by the General Method
described in chapter 1.
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 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 color. 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 1800 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 an alkaline
solution such as 1 M NaOH.
Suitability. All of these experiments are suited for use as classroom demonstrations.
Experiment 35 is ideally suited for use as a classroom demonstration using an overhead
projector. Except for Experiment 33, which uses dangerous cryogenic materials, these
experiments are also well-suited as laboratory experiments conducted by small groups of
students.
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.
Experiments 26- 29. Well-Plate Reactions with Nitric Oxide. The following reactions
are performed in a 12- or 24-well plate. Prepare the following reagents in separate wells
before generating NO(g).
Well:
1
2
3
4
Contains:
Faintly pink KMnO4(aq)
B r2(aq); (water shaken with Br2 fumes)
Fe +2(aq); (2 mL water + 2 - 3 small crystals FeSO4.7H2O)
I-(aq); (2 mL water + approximately 0.1 g KI
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. 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.
Experiment 26. 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
NO(g) + MnO4-(aq)→ NO3-(aq) + MnO2(s)
EO = + 0.72 v
Experiment 27. NO(g) reduces Br2(aq) to colorless bromide:
2 NO(g) + 3 Br2(aq) + 4 H2O→ 2 NO3-(aq) + 6 Br- (aq) + 8 H+(aq)
EO = + 0.11 v
Experiment 28. 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(I)
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 29. 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(I)
Experiment 30. Quantitative Conversion of Nitric Oxide to Nitrogen Dioxide. 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)
Experiment 31. Quantitative Conversion of Nitric Oxide to Nitrogen Dioxide. Prepare
60 mL of O2(g) in a syringe as described in Chapter 4. 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. Connect the two gas-filled
syringes with a short length of latex tubing as shown in
Fig. 23. By pushing in on the plunger of the O2-syringe,
slowly transfer a volume of O2(g) equal to half of the
volume of the NO(g). Thus 60 mL NO would require 30 mL O2. The balanced reaction
is:
2 NO(g) + O2(g)→2 NO2(g) H = -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.
Experiment 32. Relative Water “Solubilities” of NO and NO2. Prepare 60 mL of O2(g)
in a syringe as described in Chapter 4. Wash the O2(g) as described, and set the syringe
aside for use later. Prepare 60 mL NO(g) using the Fe+2/H2SO4 solution 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 33. Dinitrogen Trioxide is a Blue Liquid. Either liquid nitrogen or a dry
ice/alcohol bath is needed as a source of extreme cold for this experiment.6 (An ice/salt
bath at a temperature below –10OC 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(I)
mp = -102OC
Dinitrogen trioxide partially disproportionates into NO and NO2 at temperatures
above 4 OC according to the equilibrium:
N2O3(I)
NO(g) + NO2(g)
The disproportionate is reversible and becomes expensive as the temperature increases.
The preparation of N2O3 is 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)low temperatures→2 N2O3(I)
To do this, modify Experiment 31 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. Connect the two gasfilled syringes with a short length of latex tubing as shown in Fig. 23. 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), 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 appear.
Experiment 34. Acidic Nature of Nitrogen Oxides. Nitrogen dioxide, NO2(g), is the acid
anhydride of nitric acid. The following disproportionate reaction with water is
instantaneous and is the final step in the Ostwald Process.
6
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.
3 NO2(g) + H2O(I)→ 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. 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.
Experiment 35. Acid Rain Microchemistry7. 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 Fig. 24. 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 Fig. 25. 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 Zip-Lock bag
above the surface of the filled wells. Next, slip the filled
well plate into a Zip-Lock bag, as shown in Fig. 26. Pierce
a small hole through the bag with a sharp pencil. Work a
15-cm length of latex tubing through the hole. (Moistening
the tubing with alcohol helps to facilitate this process.)
Zip the bag shut. Place the assembly on the overhead
projector. Generate 60 mL NO2 as per Experiment 31. Fit the NO2 syringe onto the latex
tubing and slowly 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.
“Chemistry Demonstration Aids That You Can Build!” Bruce Mattson, Mary Alice
Kubovy, Jeff Hepburn, and Joe Lannan, Flinn Scientific Press. 1997
7
Experiment 36. LeChatelier Principle and the NO2/N2O4 Equilibrium. Nitrogen dioxide
is a brown-red gas that exists in equilibrium with its dimer, N2O4(I);
2 NO2(G)
Kc25 = 215
N2O4(g)
ΔH = -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 6, 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 N2O4. 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 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.
Last updated 8-01