Chapter
Chapter
27
Hydrogen Sulfide
Hydrogen sulfide is a colorless gas with an offensive stench and
is said to smell like rotten eggs. The gas can be detected at a
level of 2 parts per billion. To put this into perspective, 1 mL of
the gas distributed evenly in a 100-seat lecture hall is about 20
ppb.
Kipp generator, 1844
Hydrogen sulfide has been known since early times. The chemistry
of H2S has been studied since the 1600s. In the 19th century,
Petrus Johannes Kipp, a Dutch pharmacist, invented a convenient
device for the generation of a variety of gases in which a liquid
and solid were the reagents. The Kipp generator was especially
useful for the generation of hydrogen sulfide and hydrogen. The
device shown at right was one of the earliest and would not be
familiar to chemists who remember using the Kipp generator in
chemistry lab. More information on the use of this device is given
on our gas chemistry web site.
Hydrogen sulfide has a bent structure similar to that of water.
This is where the similarity ends, however. Sulfur is not as
electronegative as oxygen, so hydrogen sulfide is not nearly as
polar as water. Because of this, comparatively weak intermolecular
forces exist for H2S, and the melting and boiling points are much
lower than they are in water. Hydrogen sulfide and water boil at
-60.7 oC and +100.0 oC, respectively. The melting point of hydrogen
sulfide is only 25 degrees lower than its boiling point, -85.5
oC.
Hydrogen sulfide dissolves in water to make a solution that is
weakly acidic. At 0 oC, 437 mL H2S(g) will dissolve in 100 mL H2O,
producing a solution that is about 0.2 M. However, the solution
process is fairly slow. The solution equilibrium is
H2S(g) H2S(aq)
Natural gas from wells contains up to several percent H2S(g) and
are called “sour gas wells” due to their offensive stench.
Volcanoes also discharge hydrogen sulfide. Anaerobic decay aided by
bacteria produces hydrogen sulfide, which in turn, produces sulfur.
This process accounts for much of the native sulfur found in
nature.
Commercially, hydrogen sulfide is obtained from sour gas natural
gas wells. Hydrogen sulfide has few important commercial uses.
However, it is used to produce sulfur which is one of the most
commercially important elements. About 25% of all sulfur is
obtained from natural gas and crude oil by conversion of 1/3 of the
H2S to SO2 followed by the reaction between H2S and SO2:
2 H2S(g) + 3 O2(g) 2 SO2(g) + 2 H2O(g)
16 H2S(g) + 8 SO2(g) 3 S8(g) + 16 H2O(g)
Hydrogen sulfide has been used for well over a century as a
method of qualitative analysis of metal ions.
Suitability
We recommend that these experiments be presented as classroom
demonstrations rather than student laboratory activities. Advanced
high school or college-level students could conduct these
experiments providing that adequate fume hood facilities are
available. Because hydrogen sulfide has such a foul smell and is so
toxic, only individuals who are experienced with the techniques of
gas syringe manipulation should attempt these experiments. The
following experiments are included in this chapter:
Experiment 1. Hydrogen sulfide is slowly oxidized
Experiment 2. Hydrogen sulfide is a weak acid
Experiment 3. Reaction between hydrogen sulfide and aqueous
sodium hydroxide
Experiment 4. Hydrogen sulfide burns in oxygen with a howling
blue flame
Experiment 5. Reaction between hydrogen sulfide and sulfur
dioxide yields elemental sulfur
Experiment 6. Metal sulfide precipitation reactions
Experiment 7. Oxidation of metal sulfides
Experiment 1, “Hydrogen sulfide is slowly oxidized” provides an
interesting aspect of H2S that is not widely known: it oxidizes in
the presence of air. Experiment 2 shows that, despite the
similarity in the formulas H2S and H2O, the acid/base properties of
the two are quite different. Experiment 3 continues exploring the
acid-base chemistry of H2S in a reaction with aqueous sodium
hydroxide.
Experiments 4 and 5 work well as a pair. Experiment 4 “Hydrogen
sulfide burns in oxygen with a howling blue flame”, is an excellent
example of an important industrial process whereby H2S, commonly
found in natural gas deposits, is burned to prevent its release
into the environment and to provide an alternative to releasing
such a foul-smelling gas. Unfortunately, if the resulting SO2 is
not trapped, the consequences to the environment are only
marginally better. In Experiment 5, H2S and SO2 are brought
together where they react to form useful elemental sulfur.
Experiments 6 and 7 are also a pair. In Experiment 6, metal ions
are precipitated as metal sulfides reminiscent of old qualitative
analysis laboratory procedures. The odor of H2S is largely, if not
completely, avoided by conducting the reaction in a sealable
plastic food storage bag. In Experiment 7, the metal sulfides are
oxidized to form metal sulfates.
Background skills required
Students should be able to:
· generate a gas as learned in Chapter 1.
· know how to prevent accidental/unintentional discharge of
gas.
· understand fundamental concepts of high school chemistry so
that observations can be interpreted.
Time required
These experiments require more than one lecture period if all
are being presented. Furthermore, two of the experiments take
overnight to complete. The others take only minutes to perform,
excluding set-up and preparation time.
Website
This chapter is available on the web at website:
http://mattson.creighton.edu/Microscale_Gas_Chemistry.html
Instructions for your students
For classroom use by teachers. Copies of all or part of this
document may be made for your students without further permission.
Please attribute credit to Professors Bruce Mattson and Mike
Anderson of Creighton University and this website.
Preparation of hydrogen sulfide[footnoteRef:1] [1: Content for
this chapter first appeared as “Microscale Gas Chemistry, Part 8.
Experiments with Hydrogen Sulfide” Mattson, B. M.; Anderson, M.;
Nguyen J; Lannan, J., Chem13 News, 258, May, 1997.]
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. Hydrogen sulfide is
extremely toxic. Individuals who have inadequate experience with
the techniques of gas syringe manipulation should not perform these
experiments. Although the syringe method minimizes the risk of
accidental exposure to the gases generated, as a precaution the
gas-generation and gas-washing steps should be performed in a
working fume hood or outdoors. Hydrogen sulfide is also a flammable
gas.
Use a fume hood or work outdoors
The gas-generation and gas-washing steps should be carried out
inside a working fume hood or outdoors.
Toxicity
Hydrogen sulfide is extremely toxic. H2S(g) has the familiar
smell of rotten eggs. Its odor can be detected at 2 ppb. Low level
exposure can cause headache, dizziness and nausea. Inhaling higher
concentrations of the gas can cause collapse, coma, and death from
respiratory failure. The odor of H2S does not increase in
proportion with its concentration, so higher concentrations of the
gas do not smell worse than low levels.
Equipment
Microscale Gas Chemistry Kit (Chapter 1)
250 mL flask for neutralization solution
additional plastic cup for neutralization solution
Chemicals
4 g NaOH
0.22 g solid ZnS (powdered)
3 - 5 mL 6 M HCl(aq)
This quantity of zinc sulfide will produce approximately 50 - 55
mL of H2S. Under no circumstances should more than 0.22 g ZnS be
used! The production of H2S is relatively fast, taking typically 15
seconds to fill a syringe. The reaction is:
ZnS(s) + 2 HCl(aq) H2S(g) + ZnCl2(aq)
The In-Syringe Method, described in Chapter 1 and summarized
here is used to generate the H2S gas samples used in these
experiments. But first, we must make a neutralization solution:
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 “Neutralization Solution, 1 M NaOH”. This solution will be
used to neutralize excess reagents in the experiments. Sodium
hydroxide solutions can cause severe chemical burns!
Preventing unwanted discharges of hydrogen sulfide
Hydrogen sulfide is noxious and must not be discharged into
breathable air. The use of syringes to generate such gas samples
works exceptionally well and far better than any other method in
preventing undesired discharges. There are two simple
considerations to keep in mind whenever handling noxious gases:
(1) When opening the syringe (by removing the syringe cap), do
so with the plunger slightly withdrawn so the contents are under
reduced pressure. Use your thumb to maintain the plunger in this
position as shown in the drawing. This will allow some air to enter
the syringe but no noxious gas to escape.
(2) After the gas sample has been generated, discharge the used
reagents into a cup containing the Neutralization Solution.
Step-by-step instructions for the preparation of hydrogen
sulfide
Have the neutralization solution ready to go before starting the
preparation of hydrogen sulfide. Pour it into a 250 mL plastic cup
and label the solution.
It may be practical to simultaneously generate three syringes
full of H2S(g) — enough to perform all seven of these experiments.
Experiments 1, 2 and 3 collectively require one syringe full of
H2S(g), Experiment 4 requires one syringe full of H2S(g), and
Experiments 5, 6 and 7 collectively require one syringe full of
H2S(g).
1. Wear your safety glasses!
2. Make sure the syringe plunger moves easily in the syringe
barrel. If it does not, try another combination of plunger and
barrel.
3. Measure out 0.22 g solid ZnS. Place the solid directly into
the vial cap to prevent loss.
4. Fill the syringe barrel with water. Place your finger over
the hole to form a seal.
5. Float the vial cap containing the solid reagent on the water
surface.
6. Lower the cap by flotation. Release the seal made by finger
to lower the cap into the syringe barrel without spilling its
contents.
7. Install the plunger while maintaining the syringe in a
vertical position, supported by the wide-mouth beverage bottle or
flask.
8. Fill the weighing dish with 6 M HCl(aq). Draw 3 - 5 mL of
this solution into the syringe.
9. Push the syringe fitting into the syringe cap.
10. Shake the device up and down in order to mix the reagents.
Gently help the plunger move up the barrel. Upon mixing the
reagents in the syringe with vigorous shaking, gaseous H2S is
produced.
11. After gas generation has stopped, pull the plunger further
outward an additional 5 mL in order to create a slightly reduced
pressure inside the syringe. While working inside the fume hood,
remove the syringe cap while it is directed upwards. Rotate the
syringe 180o in order to discharge the reaction mixture into the
beaker containing the Neutralization Solution. Immediately recap
the syringe.
12. Syringe-to-Syringe Transfer (instead of washing). The
gas-filled syringe is not “washed” in order to remove traces of
unwanted chemicals from the inside surfaces of the syringe before
the gases can be used in experiments. Instead, use a 3 cm piece of
tubing to connect the H2S-filled syringe to a clean dry syringe.
Hold the two syringes in a vertical position with the clean, dry
syringe on top. Transfer the hydrogen sulfide to the clean dry
syringe by simultaneously pushing and pulling on the two plungers
in 10 mL increments. Do not transfer any of the liquid reagent.
After transfer is complete, pull the plungers outward by 3- 5 mL to
assure reduced pressure in the syringes. Remove the connector
tubing and cap the syringes.
Disposal of H2S(g) samples
Unwanted samples of hydrogen sulfide should be destroyed. This
is accomplished most efficiently by drawing some of the
Neutralization Solution into the syringe and reacting with as much
of the gas as possible (the syringe will likely contain some air as
well). Glassware and syringes should be washed inside the hood
before they are removed.
Teaching tips
1. Hydrogen sulfide solutions are safe and relatively odorless
at high pH, but will release the gas under acidic conditions.
2. Remove nothing from the hood until it has been rinsed with
the Neutralization Solution.
3. The Neutralization Solution can be stored in the flask as
instructed, but is transferred to a cup for use during the
generation of hydrogen sulfide and experiments with the gas. When
not in use, store the Neutralization Solution in the stoppered
flask.
4. Hydrogen sulfide can be detected at levels far below
dangerous levels; but one’s nose cannot distinguish between low and
high levels of hydrogen sulfide. Usually, students will get an
accidental whiff of the gas and this is not a problem.
Questions
1. Calculate the amount of zinc sulfide used (in moles).
2. Given that zinc sulfide is the limiting reagent, what
quantity, in moles, of hydrogen sulfide is expected?
3. Use the ideal gas law to convert your answer from Question 2
into a volume, expressed in mL at 298 K and standard pressure.
4. How does the calculated volume of nitric oxide compare with
the experimentally obtained volume? Determine the percent
yield.
5. Use the ideal gas law to determine the density of hydrogen
sulfide. Repeat the calculation for air, assuming a molar mass of
29 g/mol. Determine the ratio of densities, densityhydrogen
sulfide/densityair. Complete the sentence: Hydrogen sulfide is ___
% (heavier/lighter) than air.
6. At 0 oC 437 mL H2S(g) will dissolve in 100 mL H2O. Convert
this to moles H2S per kg water.
Experiments with hydrogen sulfide
Universal Indicator/pH 8 Solution
Experiments 2 and 4 require a slightly basic universal indicator
solution. Prepare a solution by mixing 50 mL distilled water plus 5
mL universal indicator solution. Raise the pH to 8 by bubbling
through the solution a pipet full of gaseous ammonia taken from the
vapors above a solution of concentrated ammonium hydroxide
solution. It may be necessary to repeat this transfer in order to
achieve the desired pH.
Indicator Colors
pH
Universal
Red Cabbage
4.0
Red
Red
5.0
Orange Red
Purple
6.0
Yellow Orange
Purple
7.0
Dark Green
Purple
8.0
Light Green
Blue
9.0
Blue
Blue-Green
10.0
Reddish Violet
Green
11.0
Violet
Green
12.0
Violet
Green
13.0
Violet
Green-Yellow
14.0
Violet
Yellow
Experiment 1. Hydrogen sulfide is slowly oxidized
Equipment
Microscale Gas Chemistry Kit
suitable stopper for the 18 x 150 mm test tube
Chemicals
H2S(g), 15 mL
Suitability
university lab and classroom demonstration
Applications, Topics, Purpose
homogeneous/heterogeneous solutions, chemical formulas, writing
balanced chemical equations, chemical reactivity of hydrogen
sulfide, rates of chemical reactions (chemical kinetics),
precipitation reactions, oxidation-reduction reactions
Instructions
Hydrogen sulfide is fairly soluble in water; 100 mL water at 0
oC will dissolve up to 437 mL H2S(g), producing a solution that is
about 0.2 M. However, the solution process is fairly slow. Fresh
solutions of H2S(g) are clear and colorless but become cloudy white
upon standing. The white suspension of elemental sulfur begins to
appear within an hour and is produced from the reaction between
H2S(g) and dissolved oxygen in water:
2 H2S(aq) + O2(aq) 2 S(s) + 2 H2O(l)
Place 10 mL distilled water in a 18 x 150 mm test tube (capacity
30 mL). Stopper the test tube with a rubber stopper. Prepare a
syringe full of H2S. Before removing the syringe cap, pull the
plunger outward by 5 mL, thus creating slightly reduced pressure
within the syringe. Replace the syringe cap with a 15 cm length of
tubing and bubble 10 mL of the gas below the surface of the water
in the test tube. Remove the syringe/tubing assembly and pull about
5 mL air into the syringe to remove most of the H2S from the
tubing. Replace the tubing with the syringe cap and set the
H2S-syringe aside for use in Experiments 2 and 3. Stopper the test
tube and shake the test tube vigorously to dissolve some of the
H2S(g). Set the stoppered test tube aside and observe it over the
next several hours. After 24 hours most of the H2S will have been
oxidized to elemental sulfur. Discard the resulting solution by
adding it to the Neutralization Solution.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide section, page 411.
2. The sulfur produced often appears creamy white, but not
yellow. This is because of the extremely small particle size of the
solid sulfur suspension.
Questions
1. Why is this reaction slow? What factors make some reactions
slow and others fast? Do you think the reaction would have been
faster if the test tube contained oxygen instead of air?
2. What is the milky-colored precipitate formed? What accounts
for its unusual appearance? What would you have expected the
product to look like?
3. Hydrogen sulfide is produced by anaerobic decay of biomass.
What happens to the hydrogen sulfide after it enters the atmosphere
or natural waters?
4. What steps would be necessary to convert the sulfur in the
test tube into the more familiar sulfur?
Experiment 2. Hydrogen sulfide is a weak acid
Equipment
Microscale Gas Chemistry Kit
Chemicals
H2S(g), 30 mL
20 mL of the universal indicator/pH 8 solution
Suitability
university lab and classroom demonstration
Applications, Topics, Purpose
properties of hydrogen sulfide, solution equilibrium, acids and
bases, weak acids, oxidation-reduction reactions
Instructions
Hydrogen sulfide is a weak acid with a dissociation constant
that is considerably larger than that of water:
H2S(aq) + H2O(l) H3O+(aq) + HS-(aq)Ka = 1 x 10-7
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)Ka = 1 x 10-14
Thus, a 0.01 M H2S solution will have a pH = 4.5. If necessary,
prepare a syringe full of H2S. Transfer the gas to a clean syringe.
Before removing the syringe cap, pull the plunger outward by 5 mL,
thus creating slightly reduced pressure within the syringe. Remove
the syringe cap, draw 20 mL of the universal indicator/pH 8
solution into the H2S-filled syringe, replace the cap and shake to
mix the reagents. The pH of the solution will drop from 8 to 4 as
the H2S dissolves in the solution. Discard the resulting solution
(bur not the gas) by adding it to the Neutralization Solution. Keep
the syringe capped when not in use. Remaining H2S(g) can be used in
Experiment 3.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide section, above.
2. Relatively little of the gas will be used in this experiment;
the remainder can be used in the next experiment.
Questions
1. In this case is the pH change caused by a chemical reaction
or a physical property?
2. Fill in the blanks: (a) When comparing two weak acids, the
one with the ____ Ka is the stronger of the two and will be ___
dissociated than the weaker acid. (b) When comparing two weak
acids, the one with the ____ pKa is the stronger of the two. (c) If
solutions of identical concentration were made of these two weak
acids, the one with the larger Ka would exhibit the ___ pH. (d) The
weak acid with the larger Ka will have the ___ pKa.
3. Identify the conjugate base for hydrogen sulfide.
4. Given that the acid dissociation constant for hydrogen
sulfide is Ka = 1 x 10-7, calculate the pH of a saturated solution
(assume to be 0.2 M) of hydrogen sulfide. Did you have to use the
quadratic equation to solve this problem?
5. Would you expect hydrogen selenide, H2Se(aq) to be acidic or
basic?
Experiment 3. Reaction between hydrogen sulfide and aqueous
sodium hydroxide
Equipment
Microscale Gas Chemistry Kit
250 mL beaker
Chemicals
H2S(g), 50+ mL
25 mL 6 M NaOH
Suitability
university lab and classroom demonstration
Applications, Topics, Purpose
physical and chemical changes, chemical formulas, writing
balanced chemical equations, classifying chemical changes, chemical
reactivity of hydrogen sulfide, solutions, the dissolving process,
solution equilibrium, acids and bases
Instructions
Hydrogen sulfide reacts readily with 6 M NaOH. The reaction
is:
H2S(g) + NaOH(aq) NaHS(aq) + H2O(l)
Pour the 25 mL 6 M NaOH into a plastic cup or leave it in the
beaker. Use the H2S(g) that remains from Experiments 1 and 2 or
prepare a fresh syringe full of H2S. It is unnecessary to transfer
the gas to a clean syringe for this experiment. Pull the plunger
back by 5 mL in order to create a reduced pressure in the
H2S-filled syringe, then remove the syringe cap and draw a 5 mL
NaOH(aq) into the syringe. Hydrogen sulfide reacts instantaneously
with the NaOH(aq). The plunger may move rapidly inward and/or the
NaOH solution will be drawn rapidly into the syringe. The reaction
is very rapid and could be surprising. The cup is used because its
walls will contain any splashed NaOH(aq). Discard the resulting
solution by adding it to the Neutralization Solution.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide section, above.
2. The 6 M NaOH used should be handled with caution.
Questions
1. What observations did you make to indicate that a chemical
reaction has taken place?
2. The reaction given in the Instructions above shows NaHS(aq)
as the product. Considering the excess amount of NaOH(aq) present,
NaHS(aq) probably reacts with excess NaOH(aq) to produce another
substance. What is it? Balance the chemical equation for this
reaction.
3. Why is H2S a stronger acid than H2O? Is it the same rationale
that is used to explain why HCl is a stronger acid than HF?
4. Write the chemical equation, given in the instructions, in
words, “Aqueous hydrogen sulfide and…”
Experiment 4. Hydrogen sulfide burns in oxygen with a howling
blue flame
Equipment
Microscale Gas Chemistry Kit
glass Pasteur pipet
500 mL flask with suitable stopper
birthday candle supported in a one-holed rubber stopper
matches or lighter
Chemicals
H2S(g), 30 mL
25 mL 6% H2O2(aq)
0.10 g MnO2(s)
25 mL of the universal indicator/pH 8 solution
Suitability
university lab and classroom demonstration
Applications, Topics, Purpose
combustion, chemical properties of hydrogen sulfide, energy and
chemical change, chemical formulas, chemical formulas, chemical
reactions, writing balanced chemical equations, chemical reactivity
of hydrogen sulfide, oxidation-reduction reactions
Instructions
Fit a 15 cm piece of tubing into the end of a glass Pasteur
pipet. It should make a snug fit. Enrich a 500 mL flask with O2(g)
by decomposing 25 mL 6% H2O2(aq) with 0.10 g MnO2(s) inside the
flask. (Do not drain the reagents from the flask.) Stopper the
flask to minimize O2 loss. Prepare a syringe full of H2S as
described above. Equip a one-holed rubber stopper with a birthday
candle. Set the candle a safe distance away from the syringe and
light the candle.
The general arrangement of the experimental apparatus is shown
in the figure. Fit the syringe with the pipet/tubing assembly. Two
people are needed to complete this procedure. Move the tip of the
pipet into the vicinity of the candle flame. Slowly discharge the
H2S into the candle flame at a rate suitable to ignite the gas and
maintain a flame. Move the burning pipet into the O2-filled flask.
The H2S will burn with a hotter flame and the characteristic blue
flame will be evident. An audible “roar” will be heard coming from
the mouth of the flask. Important! Never withdraw the plunger while
the pipet is lit! H2S forms explosive mixtures with air.
The combustion of H2S(g) in O2(g) produces SO2(g) as
follows:
2 H2S(g) + 3 O2(g) 2 SO2(g) + 2 H2O(g) H = -1036 kJ
Sulfur dioxide is an acidic oxide. In order to test for the
presence of SO2, add 25 mL of the slightly basic aqueous solution
of universal indicator to the SO2/O2-filled flask and it will turn
to its acidic color. Excess H2S(g) in the syringe can be used in
other experiments or can be destroyed as described in the Disposal
section.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide section, above.
2. H2S forms explosive mixtures with air.
Questions
1. Describe the reaction in terms of appearance and the sound
emitted. How did the combustion reaction change between burning in
air (when the gas issuing from the pipet was first ignited) and
burning in oxygen?
2. Why does H2S burn? Why does H2O not burn? How are these two
compounds different? One way to answer these questions would be to
consider the common oxidation states for oxygen and sulfur.
3. The combustion of H2S to form SO2 is a possible first step in
the production of sulfuric acid. In a more common first step,
elemental sulfur is burned to form sulfur dioxide. Write the
balanced chemical equation for this reaction.
4. Refer again to Question 3. In the second step in the
production of sulfuric acid, sulfur dioxide is further oxidized
with oxygen in the presence of a catalyst to produce sulfur
trioxide. In the third and final step, sulfur trioxide and water
react to form sulfuric acid. Write balanced chemical equations for
these two reactions.
5. Sulfur dioxide and sulfur trioxide are called “acid
anhydrides” because they form acidic compounds with water. Oxides
can be categorized as “acid anhydrides”, “base anhydrides” or
neither. Alkali metal oxides and alkaline earth oxides can be
placed in one of the first two categories. Which is it? Hint: Use
sodium oxide as an example. Non-metal oxides are often acid
anhydrides. Think of at least two other non-metals that form oxides
that are acid anhydrides.
Experiment 5. Reaction between hydrogen sulfide and sulfur
dioxide yields elemental sulfur
Equipment
Microscale Gas Chemistry Kit
large test tube, 25 x 250 mm with suitable stopper
ring stand and clamp
tape (electricians tape)
Chemicals
H2S(g), 20 mL
SO2(g), 30 mL (See Chapter 15)
Suitability
university lab and classroom demonstration
Applications, Topics, Purpose
matter, chemical formulas, mole calculations, chemical
reactions, writing balanced chemical equations, classifying
chemical changes, chemical reactivity of hydrogen sulfide,
molecular structure, rates of chemical reactions (chemical
kinetics), catalysts, volume-volume relationships (law of combining
volume), precipitation reactions, oxidation-reduction reactions,
industrial processes
Instructions
Elemental sulfur is produced from hydrogen sulfide gas obtained
from gas wells. In the first step, some of the H2S is burned to
produce SO2 as we did in Experiment 4. The SO2 is then reacted with
more H2S to produce elemental sulfur. The two steps are given as
follows:
Step 1. Combustion of H2S:2 H2S(g) + 3 O2(g) 2 SO2(g) + 2
H2O(g)
Step 2. Redox Combination:16 H2S(g) + 8 SO2(g) 3 S8(s) + 16
H2O(l)
In this experiment we will demonstrate Step 2 of the sequence.
Produce 30 mL of SO2 as described in Chapter 15 and 20 mL H2S as
described in this chapter. Fill a large test tube with water and
then pour the water out. Moisture catalyzes the reaction and MUST
be present. Position the test tube vertically with a ring stand and
clamp. Tape the end of two lengths of tubing together near the open
end. Connect a length of tubing to each of the two gas-filled
syringes. Place the open ends of the tubing near the bottom of the
test tube. Simultaneously transfer 10 mL incremental amounts of SO2
and H2S to the test tube. Although the stoichiometry calls for 2 mL
H2S for every 1 mL SO2, it is best to keep the H2S as the limiting
reagent. Soon after the gases come in contact, canary yellow sulfur
will completely line the inside of the test tube.
Slowly add Neutralization solution to the test tube until it is
1/3-full. Rest an oversized stopper (suitably sized stopper placed
upside down) over the test tube opening in order to minimize gas
dispersion.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide section.
2. Review the toxicity warnings in for sulfur dioxide, Chapter
15.
Questions
1. Describe your observations for the chemical reaction that
took place.
2. The wet glass serves as a catalyst for the reaction. Describe
how the droplets were involved in the reaction.
3. Assign oxidation numbers for every element in the reaction
between hydrogen sulfide and sulfur dioxide (“Step 2. Redox
Combination”, above)
4. Why was the experiment designed so that hydrogen sulfide was
the limiting reagent?
5. Why do the gases stay in the test tube?
6. Hydrogen sulfide is often present in natural gas wells. It
can be separated from natural gas by cooling. Devise a scheme
whereby the hydrogen sulfide collected could be turned into
elemental sulfur.
Experiment 6. Metal sulfide precipitation reactions
Equipment
Microscale Gas Chemistry Kit
12-well plate
six test tubes, 18 x 150 mm, with suitable stoppers
plastic disposable pipet
gallon (4 L) sealable food storage bag
filter funnel and filter paper (optional — See Bi+3 solution
below)
Chemicals
H2S(g), 40 mL
Cd+2: 0.1 g Cd(NO3)2.4 H2O in 5 mL H2O
Cu+2: 0.1 g CuSO4.5 H2O in 5 mL H2O
Pb+2: 0.1 g Pb(NO3)2 in 5 mL H2O
Bi+3: 0.1 g Bi(NO3)3.5 H2O in 20 mL H2O; stir for 30 minutes and
filter or allow to settle overnight; you will use the clear aqueous
portion
10 mL 6 M NaOH(aq): 2.4 g NaOH
10 mL 30% H2O2(aq) (only necessary if Experiment 6 is being
performed)
Suitability
university lab, and classroom demonstration
Applications, Topics, Purpose
physical and chemical changes and properties, chemical formulas,
chemical formulas, chemical reactions, writing balanced chemical
equations, properties of hydrogen sulfide, chemical reactivity of
hydrogen sulfide, solutions, the dissolving process, solution
equilibrium, precipitation reactions
Instructions
Prepare stock solutions of Cd+2, Cu+2, Pb+2, Bi+3, NaOH, and
H2O2 as described in the Chemicals list. Transfer 3 mL of the Cd+2
solution to each of two wells. Repeat with the Cu+2, Pb+2, and Bi+3
solutions. Thus, eight wells will contain two samples of four
different solutions. Transfer 5 mL of 6 M NaOH to each of two wells
and 5 mL of 30% H2O2(aq) to each of two wells. (If Experiment 6 is
not being performed, only one well of each metal ion is required
and the wells of H2O2(aq) are not required.) Write a key to the
contents of each of the wells for future reference.
Prepare a syringe full of H2S(g) and wash the gas. Place a 6 cm
length of a plastic pipet stem between the four middle wells in
order to prop up the plastic bag above the surface of the filled
wells. Pierce a small hole through the bag with a sharp pencil and
work the tubing through the hole as shown in figure below.
(Moistening the tubing with soapy water facilitates this process.)
Next, slip the filled well plate and a plastic pipet into a plastic
bag and zip shut.
Dispense all of the H2S(g) just above the surface of the eight
wells containing metal ion solutions. (Avoid dispensing it over the
NaOH and the H2O2 solutions.) An immediate reaction will be noted
for each of the metal ions. Blue Cu+2(aq) will produce a brown
web-like film of CuS(s) on the surface. Colorless Cd+2(aq) will
produce a distinctive yellow precipitate of CdS(s). Colorless
Pb+2(aq) will produce a spectacular silvery mirror of PbS(s) on the
surface. Colorless Bi+3(aq) will produce a black/metallic bronze
precipitate of Bi2S3(s) on the surface. The reactions between the
various metal ions and H2S(g) are similar; the reaction for
Cd+2(aq) is:
Cd+2(aq) + H2S(aq) + 2 H2O(l) CdS(s) + 2 H3O+(aq)
Allow the reactions to proceed for at least 5 minutes before
going on to Experiment 7. Do NOT open the plastic bag. The two
wells of NaOH(aq) will absorb the excess H2S(g) overnight.
Clean-up
No clean-up is necessary at this point if Experiment 7 is being
performed. If Experiment 7 is not being performed, follow the
Clean-up procedure at the end of Experiment 7.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide section.
2. Heavy metals should be handled with care, avoiding dermal
contact.
3. Heavy metals should be disposed of properly according to
local regulations. Under no circumstances should heavy metal ions
be poured down the drain or onto the ground.
Questions
1. Describe the reactions that took place between the various
metal ion solutions and hydrogen sulfide.
2. Why did the reactions take place at the surface?
3. The balanced chemical equation for the reaction between
Cd+2(aq) and H2S(aq) was provided in the Instructions. Balance the
chemical equations for the other three reactions.
4. Although the pH of each solution was not determined before
and after the reaction with hydrogen sulfide, what change in pH
would you predict according to the chemical equation?
5. Most transition metal sulfides are extremely insoluble. Look
up the Ksp values for the sulfides produced in this experiment.
Convert these values to molar solubilites.
Experiment 7. Oxidation of metal sulfides
Equipment
(continuation of Experiment 6)
Chemicals
(continuation of Experiment 6)
Suitability
university lab and classroom demonstration
Applications, Topics, Purpose
physical and chemical changes and properties, chemical formulas,
chemical formulas, chemical reactions, writing balanced chemical
equations, solutions, the dissolving process, solution equilibrium,
precipitation reactions, oxidation-reduction reactions
Instructions
Without opening the plastic bag, use the plastic disposable
pipet to transfer at least 3 mL of H2O2(aq) to one of each pair of
wells for each metal sulfide. Within a few minutes, bubbles will
appear in the well containing CuS. Within 40 minutes the solutions
containing CdS and PbS will have returned to clear. In all cases,
the sulfide anion has been oxidized to the sulfate ion. For
example:
CdS(s) + 4 H2O2(aq) Cd+2(aq) + SO42-(aq) + 4 H2O(l)
Within 2 - 3 hours the dark color of bismuth sulfide will be
replaced with white, insoluble bismuth sulfate:
Bi2S3(s) + 12 H2O2(aq) Bi2(SO4)3(s) + 12 H2O(l)
Clean-up: Allow the bag to stand overnight. The NaOH(aq) will
react with excess H2S(g). Without opening the plastic bag, draw a
few mL of the NaOH(aq) into the syringe to remove traces of H2S(g).
Wear gloves to avoid contact with unreacted H2O2(aq). It is now
safe to open the bag indoors. Remove the contents carefully and
discard the bag and pipet in the trash. Discard metal ions
according to local regulations. Wash the syringe contents
(NaSH(aq)) down the drain with plenty of water.
Teaching tips
1. Review the four teaching tips provided in the Preparation of
Hydrogen Sulfide.
2. There are four reactions taking place here. The fastest one
takes a few minutes, but the slowest one takes three hours or more.
As a demonstration, start early enough in the class period to see
the initial changes and revisit the experiment the next day.
3. Skunk scent contains trans-2-butene-1-thiol, CH3CHCHCH2SH,
(or “TBT”) which is somewhat related to H2S in that both molecules
possess the thiol group, SH. As with the oxidations studied in this
experiment, TBT can be oxidized to the odorless sulfonic acid
trans-CH3CHCHCH2SO3H. This chemistry is used in a “home skunk
remedy” for treating pets who have been sprayed by
skunks.[footnoteRef:2]1 [2: 1 “Skunk Non-scents,” Nancy Touchette,
Chem Matters, page 7, October, 1996. ]
Questions
1. Describe the reactions that took place between the various
metal sulfides and hydrogen peroxide. Address what happened to
precipitates present, if anything, as well as any color changes
observed.
2. Write the balanced chemical equation for the reaction that
takes place between hydrogen peroxide and (a) copper(II) sulfide;
and (b) lead(II) sulfide.
3. This experiment is titled “Oxidation of metal sulfides”. What
exactly was oxidized? The metal ion?
4. What happens to the excess hydrogen sulfide as it stands
overnight? Write a balanced chemical equation for this.
5. Metal sulfides are insoluble as noted in Question 5 of the
previous experiment. What are some reasons why one would want to
oxidize metal sulfides?
Clean-up
At the end of the experiments, destroy any unused hydrogen
sulfide by drawing 20 – 30 mL 1 M NaOH(aq) solution into the
syringe. Discharge the solution, and rinse syringe with water.
Place the syringe in a plastic bag and discard in the trash.
Summary of Materials and Chemicals Needed for
Chapter 27. Experiments with Hydrogen Sulfide
Equipment required
Item
For Demo
Microscale Gas Chemistry Kit
1
250 mL beaker
1
glass Pasteur pipet
1
500 mL flask with suitable stopper
1
large test tube, 25 x 250 mm with suitable stopper
1
ring stand and clamp
1
12-well plate
1
test tubes, 18 x 150 mm, with suitable stoppers
7
plastic disposable pipet
1
filter funnel and filter paper (optional)
1
Materials required
Item
For Demo
1-gallon (4 L) sealable plastic food storage bag
1
birthday candle supported in a one-holed rubber stopper
1
matches or lighter
1
tape (electricians tape)
1
Chemicals required
Item
For Demo
sodium hydroxide, NaOH
5 g
sodium hydroxide, 6 M NaOH
40 mL
zinc sulfide, ZnS, powder
1 g
hydrochloric acid, 6 M HCl
40 mL
universal indicator
5 mL
concentrated ammonium hydroxide
a
hydrogen peroxide, 6% H2O2(aq)
25 mL
hydrogen peroxide, 30% H2O2(aq)*
10 mL
manganese dioxide, MnO2(s)
0.1 g
sodium bisulfite, NaHSO3(s)
2 g
cadmium(II) nitrate tetrahydrate, Cd(NO3)2.4 H2O
0.1 g
copper(II) nitrate pentahydrate, CuSO4.5 H2O
0.1 g
lead(II) nitrate, Pb(NO3)2
0.1 g
bismuth(III) nitrate pentahydrate, Bi(NO3)3.5 H2O
0.1 g
* (only necessary if Experiment 6 is being performed)
a. only the NH3 fumes from the concentrated ammonium hydroxide
solution will be used
464 Microscale Gas Chemistry© 2017
Chapter 27. Experiments with Hydrogen Sulfide 463
¬ ¾ ¾
¾ ® ¾ ¾
