Download Chemistry 15E:

Supplementary Material
Note to Reader: This supplement is an edited version of the original laboratory guide
written and developed by Morgan Mihok for the general chemistry laboratory, Chemistry
15E offered at The Pennsylvania State University in the spring of 2002. The section
entitled, “Getting Started,” was omitted from the supplement because it discussed
procedures related to the laboratory system at Penn State and would not be of little
benefit to most readers. Experiments 1,3, 4, 5, and 6 were also omitted because they are
found in their entirety in one of two versions of Chemtrek, by Stephen Thompson as
referenced below and were only slightly adapted for this course. The table of contents
does contain all the sections of the lab guide. A table listing all the experiments and their
references is included in this Supplement.
Chemistry 15E:
Environmentally-focused Chemistry Student Laboratory Manual
Course Objective
This course was designed as a way for a graduating senior to communicate her love of the
environment, chemistry, and society. While I hope that you will learn kinetics, redox
chemistry, acid/base equilibria, and other chemical concepts, I am more interested in
whether you learn to link those concepts to processes occurring in the environment
around you. It is not only important to understand reaction cycles piece by piece, but also
in the context of the entire ecosystem. Kinetics depends partly on temperature, which in
turn determines concentrations of chemicals, the levels of which may be regulated to
different extents depending on the ability of a government to enforce standards (or a
populous to demand them), which in part depends on the economic stability of a region,
and so on. The connections are numerous and incomprehensible. Yet if later in life, you
can recognize one of these links, many of which Rachel Carson points out in her book
Silent Spring, you may be able to remediate environmental problems at a level beyond
which you can currently imagine. For those of you not planning on entering an
environmentally related field, you can hopefully gain a deeper understanding of the world
around you – the ability to make connections between seemingly unrelated concepts
serves anyone well.
Morgan Mihok
Sources of Experiments
The experiments in this course are drawn from a number of sources and are at
different levels of development with respect to the number of times they have been
conducted, evaluated, and revised.
The table on the following page provides a list of the experiments, some of which
are subdivided to illustrate multiple sources for the experiment. The source is listed next
to each experiment, followed by an estimate of the number of students who have
conducted the experiment.
Experiment
Source
Estimated student runs
1. Introduction to the Laboratory
1
Thousands
2. Iron and Alkalinity Determinations
2
Hundreds
3. AA Cation Determination
--
Thousands for Mg and Ca
4. Paper and Liquid Chromatography
1
Thousands
5. Ion Chromatography
3
Hundreds to Thousands
6. Acid-Base Chemistry, Parts 1 and 2
1
Thousands
7. Acid Rain Deposition
4
Hundreds to Thousands
8. Fe remediation: Kinetics
5
Tens to hundreds at Lewis
and Clark University
Fe remediation: Liquid/liquid extraction
--
Thousands
9. Fe remediation: Degradation of 2,4-D
6
35 at Penn State
10. Wastewater treatment
6
35 at Penn State
Sources
1
Thompson, Stephen. Penn State Version of Chemtrek: Small-scale experiments for
General Chemistry. Prentice Hall: Englewood Cliffs, NJ. 2000.
2
Kegley, S.E.; Landrear, D.; Jenkins, D.; Gross, B.; Shomglin, K. Water Treatment:
How Can We Purify Our Water? Student Manual. John Wiley & Sons: New York.
2000.
3
Sinniah, K.; Piers, K. J.Chem.Educ. 2001, 78, 358.
4
Thompson, Stephen. Chemtrek. Prentice Hall: Englewood Cliffs, NJ, 1989. Adapted
for Chem 15 by J.T. Keiser, 17 April 1997.
5
Balko, B.A.; Tratnyek, P.G. J.Chem. Educ., 2001, 78, 1661.
6
Original
-- standard laboratory technique, adapted for Fe remediation laboratory.
1
References
Throughout this lab manual, the reference for the source of each experiment is
given in the introduction, or, if the introductory and experimental references differ, at the
beginning of the corresponding section. Experimental references are also given in the
table on the preceding page. Only Experiment 7 has a comprehensive reference list at the
end of the experiment due to the vast number of source materials for that experiment.
The material presented in the experiments is in no way comprehensive on any of
the topics. A list of useful books, journals and websites has been included on the
following page. The brief description of each resource explains the level of the material
and how useful the resource might be for this particular class. The experiments that were
extracted in their entirety from a particular source were not included in the supplementary
material and include: Experiments #1, 3, 4, 5, and 6. You can view them directly from
the source documents.
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Suggested Book Resources
Hemond, H.F.; and Fechner-Levy, E.J. Chemical Fate and Transport in the
Environment. 2nd ed. Academic Press: New York, 2000. TD193.H46
This book covers partitioning, acid-base chemistry, redox and kinetics. It is easy to read,
and almost fun if you like chemistry. It is also divided up well – you can pick up and
read any section of it. Overall, it is quite a useful book for explaining basic chemical
concepts through an environmental lens.
Pradyot, P. Handbook of Environmental Analysis: Chemical Pollutants in Air, Water,
Soil, and Solid Wastes. Lewis Publishers, NY: 1997. TD193.P38.
Although the layout is boring, this book is quite good, and covers all the analytical
techniques we will use in this course (at least as the course is now). Furthermore, the
book gives detailed methods for analysis of a variety of environmentally important
compounds and water quality parameters in a step-by-step approach. Finally, this book
also gives detailed problems and solutions regarding how to use the instrumentation
available to get decent numbers – basically, it helps to explain the “black box” and how it
decides what numbers to give you.
Hauser, Barbara A. Practical Manual of Wastewater Chemistry. Ann Arbor Press, Inc:
Chelsea, MI. 1996. TD 735.H38.
This book is designed to enhance wastewater treatment plant operators’ understanding of
the chemistry involved in their analyses of wastewater. Not only are the methods in here
EPA approved, they are described at a level easily understood by someone with only a
basic chemistry background. Hauser also briefly describes the environmental relevance
of each water quality parameter, and gives methods to test for each (this may be useful
for the final experiment/project). This book is so relevant to the course that if there was
one more week of classes, I would have incorporated one of the experiments into the
course.
Radojevic, M.; Bashkin, V.N. Practical Environmental Analysis. Royal Society of
Chemistry, Cambridge: Cambridge, UK. 1999. TD193.R3.
Like Barbara Hauser’s book, this book also covers a huge range of pollutants and water
parameter tests. It will also be quite useful for ideas for the wastewater treatment project.
3
Benjamin, M.M. Water Chemistry. McGraw Hill: Boston. 2002. GB855.B46.
This is a fantastic undergraduate chemistry textbook. It contains explanations for all of
the chemical concepts covered in the course in similar context to that of the course. It’s
likely that if this were a lecture as well as a lab course, that this might be the textbook.
Baird, Colin. Environmental Chemistry. W.H. Freeman and Company: NY, 1999.
TD192.B35.
This book covers all the chemistry you’ll need for this course in environmental detail –
but it’s also fairly easy to read and gives good explanations as to the links between
chemistry and the environment. It was the textbook for Chem 402 (Environmental
Chemistry).
Suggested Journal Resources
Journal of Chemical Education. QD1.J93. Available in the Physical Sciences Library
This is one of the most useful journals for undergraduate chemists because the materials
is explained so that it can be taught – which means that you should be able to learn from
how it is explained. Recent issues of this journal have had fairly sizable sections devoted
to green and environmental chemistry, so it’s a great place to go for information relevant
to this course. You can search for articles in the journal online at
http://jchemed.chem.wisc.edu/Journal/Search.
Environmental Toxicology and Chemistry. QH545.A1E594. Available in the Physical
Sciences Library.
Environmental Science and Technology. Available online through LIAS.
These two journals are combined because they both contain articles that very well may be
of interest to you as students of environmental chemistry. However, the articles
contained in these journals are likely to be more detailed and specific than those in the
Journal of Chemical Education are. If you think you are interested in doing research in
environmental chemistry, the articles in these two journals should point you towards the
top research in the field.
Water, Air, and Soil Pollution. TD172 .W36
Well, as you may guess from the name, this journal specifically addresses pollution of the
water, air, and soil – which are each addressed at some point over the course of this lab.
The review articles in this journal are generally easy to read, and it is fun to pick up the
most recent issue and see how current debates about pollution are playing out.
4
Suggested Web Resources
www.epa.gov/
This website will give you information on water quality standards, pesticide
regulations, and more. There are also a ton of PDF files available on just about any topic
related to the environment that you can imagine. The site is sometimes slow to run, but
when it’s good to go, it’s a great place to go for information.
www.uaja.com/
Hey, look! It’s a bird’s eye view of State College’s wastewater treatment plant!
In all seriousness, this website contains loads more information than it looks like from the
homepage – there are links to numerous PDF files that give detailed information about
the problems surrounding Centre County’s wastewater treatment systems and the Spring
Creek Watershed. To top it off, the people who run this place are friendly and will give
you a tour if you ask for one – just bring noseplugs if you have them.
Note – you will need to visit this site – and read through the material on it – to complete
your final report!
www.springcreekwatershed.org/
This is the website for the Spring Creek Watershed Community, an organization
that is currently monitoring the water quality of the Spring Creek Watershed. It’s a neat
organization, and the water they are monitoring very well may be the water you later
drink at restaurants in town – so it may be good to check out the website.
5
GRADING
Lab Reports
 Experiments 1, 2, 5, 6, and 8 will require reports that can usually be
completed in lab.
 Experiments 3 and 7 will require some calculations and analyses based on
experimental data and these write-ups will be due the following week in class.
 Instead of a lab report for experiments 9 and 10, you will be asked to submit a
formal typewritten report.
Late Lab Reports Policy


Significant late penalties are assessed for lab notebooks that are not handed in on
time. Reports received one week late have a 30-point penalty.
Reports that are more than one week late will not be accepted.
Quizzes/Test
There will be short quizzes given at the beginning of most lab periods according to the
schedule on the inside cover. These cover the major points of the experiment of the day.
Students who arrive late for lab (i.e., after the quiz has been collected) will receive a zero
on that quiz. Contained in this packet are outlines of what you should know for each
quiz.
There will also be a cumulative lab exam, given on the last day of lab. An outline of
what will be covered on this exam is contained in this packet.
Instructor Evaluation
The instructor will assign a grade to each student based on his or her perception of the
student’s overall performance in the lab. This will include the use of the laboratory
notebook, attitude, independence, technique, and CLEAN-UP STYLE!
Lab Monitoring
Each student will be assigned one day for which they are responsible for lab monitoring.
This will normally entail an end of the period clean up of the shared areas such as the
sinks, balances, and the chemical supply area.
THE FINAL GRADE
Lab reports
Quizzes
Final Exam
Instructor Evaluation
Wastewater Treatment Paper
6
40%
15%
20%
10%
15%
Experiment #2:
Det ermination of
[Fe 2+ ], A lkalinity and Met al I on Content of Loc al
Wat er Samples
The background information sections are modified versions of information and fact
sheets for the different water quality parameters, which can be found on the Creek
Connections of Allegheny College’s webpage:
http://creekconnections.allegheny.edu/Chemistry/ChemistryMain.html
Water Quality Parameters
Background Information
pH Information
Background
The pH of water is very important to water quality because it controls the types and rates
of many chemical reactions in water, and aquatic organisms have a specific pH range in which
they can live. Natural, uncontaminated rain water is generally somewhat acidic, with a pH of
about 5.6. This acidity is due to the natural dissolving of carbon dioxide (CO2) in precipitation
(H2 O) to form carbonic acid (H2CO3). The extra hydrogen ions are produced when the carbonic
acid dissociates (breaks apart) producing H+ and bicarbonate HCO3-.
Once precipitation reaches the ground, a variety of organic and inorganic chemical
reactions may take place to alter the pH of water. In the upper parts of the soil, infiltrating water
commonly reacts with organic matter to form organic acids, and eventually lower the value of pH
(more acidic). Reaction with inorganic minerals (in rocks for example) dominate once water
infiltrates beneath the soil; most of these reactions will use free hydrogen ions (buffering the
solution) and therefore cause an increase in pH (more basic). The geology of a region exerts a
strong control on the pH of natural waters. For example, minerals such as calcite (calcium
carbonate – CaCO3), the main component of limestone and the cement that holds sandstone
particles together, are especially effective at causing increases in pH. As calcium carbonate
dissolves, free hydrogen ions are used. This ability of a water sample to act as a base – i.e., to
resist acidification – is called alkalinity. (See the alkalinity information section.)
Once water enters lakes and streams, aquatic life may affect pH. Respiration by plants
and animals and decomposition produce CO2, allowing it to react with water to form carbonic
7
acid and the pH levels of a waterway can decrease. However, during daylight hours, plants
photosynthesize using CO2 and keeping it from forming carbonic acid and extra H+. Under
normal stream conditions, pH levels are usually highest at the end of a day of photosynthesis,
lowest after a night of plant respiration.
All aquatic life has a specific pH range that it can tolerate and to which it is adapted. If
the pH changes even slightly, it will stress the creatures and may even kill them. At extremely
high (9.6) or low (5.0) pH values, the water becomes unsuitable for most organisms. Low pH
causes an imbalance in the sodium and chloride ions in aquatic animals' blood. At low pH,
hydrogen ions may be taken into cells while expelling sodium ions. Higher acidity can increase
the concentration of toxic metal concentrations in a stream, such as aluminum (Al3+) and copper
(Cu2+). These metals were locked up in mineral matter under neutral pH levels, but become
mobile when the pH lowers. Metal can clog fish gills causing breathing complications or cause
deformities to young fish.
Human Impact
Acidic waters have been and continue to be a major environmental concern. Whereas
unpolluted precipitation has a pH of about 5.6, the precipitation in most of the Northeast United
States has a pH of between 4 and 4.5. Air pollution is the cause. Increased amounts of nitrogen
oxides (NOx) and sulfur dioxides (SO2) gases, primarily from the burning of fossil fuels by power
plants and industry and from car exhaust, react with water and are converted to nitric acid
(HNO3) and sulfuric acid (H2 SO4) in the atmosphere. Both of these strong acids lower the pH of
the rain, and the streams that gain this precipitation. Waterways may not be affected by this acidic
rain if the watershed contains a considerable amount of acid-neutralizing rocks such as limestone,
CaCO3 . The State College area’s streams are fairly alkaline due to the large amounts of
limestone in the area. However, a region with low alkalinity can have streams that are damaged
by acid rain. For example, the Adirondack region in New York has rocks and streams that are
unable to neutralize the acid rain; as a result, widespread fish kills have occurred.
Coal mining operations (current and abandoned) can increase the acidity of a waterway
through acid mine drainage (AMD). The waste material of coal mining is called spoils or
overburden and is the discarded soil and crushed rock found above and between coal seams. This
waste contains iron pyrite (fool's gold) and when exposed to air and water, it reacts to form iron
hydroxide (Fe(OH)3) and sulfuric acid (H2 SO4). The acid can dissolve other minerals and metals,
and the water can become very acidic (as low as a pH of 2) as it enters local streams.
8
Water Quality Criteria
As in many chemicals, there is no distinct dividing line between safe and harmful pH
levels. The drinking water standards set by the Environmental Protection Agency (EPA) calls for
a minimum pH of 6.5 and a maximum pH of 8.5. Natural waters should have a pH between 5.0
and 8.5, since lower or higher values are likely to be harmful to fish populations and other aquatic
life.
Alkalinity Information
Background
Alkalinity is a measure of the ability of a water system to resist changes in pH when acid
is added to water. A stream that has a high alkalinity is well buffered so that large inputs of acid
(from acid rain for instance) can be made with little affect on the stream pH. A stream that has a
low alkalinity is poorly buffered and may undergo large, sudden drops in pH in response to acid
inputs.
The amount of carbonate (CO32-) and bicarbonate (HCO3-) in water helps to determine
its alkalinity. The more of these natural buffers that are present, the better chance the water has to
resist a change in pH. Carbonate (CO32-) will react with a free hydrogen ion (H+) to form
bicarbonate (HCO3-). Bicarbonate will react with free hydrogen ions to create carbonic acid
(H2 CO3), which then can dissociate into water and carbon dioxide. During this process, free
hydrogen ions have been locked up, thus keeping the pH from lowering (keep in mind, a low pH
has lots of extra hydrogen ions present). The formula for the above reactions follows (reactions
can also reverse):
CO32- + H+  HCO3HCO3- + H+  H2CO3
H2CO3  H2O + CO2
The reactions are balanced and are able to deal with the free hydrogen ions that are present before
they make the pH level drop. A problem occurs when additional free hydrogen ions are added to
this balanced system. Acids such as sulfuric acid (H2 SO4) and nitric acid (HNO3) (the primary
components of acid rain) provide extra hydrogen ions when they dissociate. Sulfuric acid will
eventually break down yielding 2 hydrogen ions (H2 SO4  2H+ + SO42-).
To combat these additional hydrogen ions, which would lower the pH if left alone,
additional bicarbonate and carbonate need to be added to the water. Carbonic acid (H2CO3) will
9
do this for us. Carbonic acid (H2 CO3) does not have to dissociate into water and carbon dioxide;
instead it can react with carbonate based rocks such as sandstone, limestone, and dolomite as part
of the rock's weathering process. Calcium carbonate (CaCO3) makes up limestone and the cement
that holds sandstone together, while magnesium carbonate (CaMg(CO3)2) makes up dolomite.
Both can react with carbonic acid yielding either calcium bicarbonate Ca(HCO3)2 or magnesium
bicarbonate Mg(HCO3)2:
H2CO3 + XCO3  X(HCO3)2
X = Ca, Mg
The calcium (Ca2+) and magnesium (Mg2+) drop off as a solid to the stream bottom while 2
bicarbonates (HCO3-) remain, each able to react with one free hydrogen (thus maintaining the
pH). This reaction yields carbonic acid again (HCO3- + H+  H2 CO3).
Watersheds with high alkalinity have the sandstones, limestones, and dolomites and the
corresponding calcium carbonates/magnesium carbonates needed to help buffer a stream. They
are able to handle additions of extra hydrogen ions. These rock types exist in Western
Pennsylvania. Watersheds where the bedrock does not consist of sandstones and limestones, but
instead have igneous rocks like granite and basalt, are unable to provide the needed
calcium/magnesium carbonate that rid acidity. Streams in those areas have low alkalinity and a
pH below 5.4. An artificial source of alkalinity is lime (strictly speaking, lime is CaO, but the
term is sometimes used, as it is here, to refer to calcium carbonate), used to neutralize a stream or
even treat acid mine drainage. Lime is also used as a soil amendment to rid acidity in cropland,
gardens, and lawns, and as a large-scale remediation technique in Scandinavia to rid lakes there
of excess acidity.
Human Impact
Alkalinity is an important measure of a stream ability to absorb inputs of acid. Acid rain
and acid mine drainage from coal mining causes a considerable drop in pH of stream water.
Rapid seasonal changes in pH often occur in the spring and fall. In the fall, increased organic
matter can cause greater inputs of organic acids from decaying organic matter. To address this
decrease in pH, bicarbonate and carbonate must be used, removing their availability to react with
hydrogen supplied by organic matter. During the spring, heavy rains and melting snow can result
in a large, sudden input of acid into hydrologic systems, too much to buffer, causing a rapid drop
in pH. In some cases, such an "acid spike" results in fish kills as the pH drops below acceptable
levels for supporting aquatic life.
10
Water Quality Criteria
The EPA has suggested a minimum of 20 mg/L of CaCO3 for freshwater aquatic life
except where natural concentrations are less. Although this criteria has been established, many
problems exist as streams that are acidic or streams that suffer changes in alkalinity through the
year.
Temperature Information
Background
The temperature characteristics of stream water directly and indirectly control aquatic
ecosystems and water quality. Thermal pollution refers to the addition of warmer or colder water
that causes an unstable jump in the temperature of a waterway.
The sun's energy affects water temperature, and every waterway's temperature will
naturally fluctuate from season to season. The more sunlight that hits the water's surface, the
warmer the water will get. Narrow, well-shaded headwater streams are often cooler than wider,
larger streams that are not fully shaded by streamside (riparian) forests.
In addition to shading, the physical dimensions of the waterway will also affect the
temperature. Shallow water will fluctuate in temperature faster than deeper water. Running
water tends to be cooler than stagnant, still water. In a stream, the shallow riffles or rapids are
often cooler than the slow moving, deep pools. The most downstream stretches of creeks and
rivers are often warmer than the upstream sections, and may even have a slight thermal
stratification (temperatures differ at various depths) in these deep, slow sections.
Temperature affects some of the chemical parameters of water, probably the most
important of which is dissolved oxygen. At lower temperatures, more oxygen can be dissolved in
the water because the gas molecules are moving slower and are more compact. At higher
temperatures, dissolved oxygen and other gases in water move faster and spread farther apart,
including out of the water. Also at higher temperatures, the water molecules may move faster and
bump out oxygen. At warmer temperatures, the gas molecules themselves have a greater average
energy, which means that a larger fraction will have enough energy to be in the gas phase. Think
of how the gases (carbon dioxide or the fizz) in soda pop eventually escape as it warms up.
There is a natural fluctuation of waterway temperature from season to season, even day
and night, and aquatic life can cope with these natural changes. When humans alter the
temperature of waterways, it may harm aquatic life; a thermal change of 2°C or more is harmful
to stream organisms. All species have a specific range of temperature in which they are adapted.
Fishermen know that trout like cold water streams, while other fish like carp and bluegills can
11
tolerate warmer waters. If a stream changes temperature, organisms that cannot tolerate the
change are stressed and must either reduce activity, move somewhere else, or in extreme cases,
perish. Many life cycles of fish and aquatic insects are tied to water temperature. These creatures
use temperature cues to determine when to spawn, lay eggs, when the eggs will hatch, and when
insect larvae will emerge from a stream to fly away. Thermal pollution can disrupt the timing of
the life cycles, possibly causing eggs to hatch before sufficient food resources are available or
larva to emerge when it is too cold atmospherically.
Human Impact
Humans can alter natural temperature characteristics of a stream by direct actions to the
waterway or indirectly through alterations to the watershed. Industries and power plants discharge
warm water that was used in the manufacturing process (boilers) or to cool machinery and
turbines. When industries and community water authorities withdraw water from a stream, it may
decrease the water depth. Since shallower water heats up more readily than deeper water, water
withdrawal may increase stream temperatures. Water released from dammed lakes can also alter
temperatures because it is often withdrawn from near the lake bottom and is often cooler than the
stream temperature in the summer, warmer in the winter. The shock of these rapid temperature
changes can be too much for aquatic life to handle.
Water Quality Standard
The EPA has established a formula for two important temperature extremes for streams:
upper temperature limit and a weekly average. The upper temperature limit, or short term
maximum, is set at 30.6°C for the area from the southern shore of Long Island, New York to
Cape Hatteras, North Carolina. The weekly maximum is set at 27.8°C for this same area. No
regulations are established for the zone containing Western Pennsylvania.
Hardness Information
Background
Hardness is one of the most commonly tested parameters of drinking water and is
defined as the sum of the polyvalent cations present in the water. The minerals of calcium (Ca2+)
and magnesium (Mg2+) are usually the predominant cations responsible for hardness levels. Other
ions, such as iron (Fe2+), manganese (Mn2+), and aluminum (Al3+), may contribute to hardness,
but in natural waters these other ions are usually found in insignificant amounts. Just like total
dissolved solids (TDS), hardness is a parameter that somewhat summarizes the amount of various
12
substances that may be in the water. Though methodology for hardness tests can vary to account
for different ions, most simple tests focus on just calcium and magnesium. Hardness test kits
often express results in parts per million (ppm) of CaCO3 (calcium carbonate).
Calcium and magnesium may be added to a natural water system as it passes through
soil and rock containing large amounts of these elements in mineral deposits. Hard water is
usually derived from the drainage through calcareous (calcite-rich) sediments and rock, such as
limestones, sandstones, and siltstones. Dolomites are rich in magnesium. These rocks are found in
Western Pennsylvania, thus affecting the hardness levels in our water. In this area, if water has
had the opportunity to interact with bedrock, and soils for a long time (such as groundwater), it
will be hard. If the cations responsible for making water hard are not calcium and magnesium, but
are iron, sulfate, chloride, manganese, or aluminum, etc. instead, this is considered to be "noncarbonate hardness".
Water that has entered waterways directly without soaking into the ground will be
significantly softer. Collected rainwater is usually soft because it has not interacted with any
geological sources of the cations. Soft water is also derived from the drainage of igneous rocks,
because these rocks don't weather very easily, don't release many cations, and don't always
contain calcium and magnesium.
Hardness in water can have some biological impacts on waterways. Calcium is an
important component of aquatic plant cell walls, and the shells and bones of many aquatic
organisms. Magnesium is an essential nutrient for plants and is a component of the chlorophyll
molecule. If there is very little calcium in a waterway (less than 10 mg/L), only sparse plant and
animal life can be supported because this waterway does not usually contain enough organic
matter and nutrients. Hardness is also helpful in limiting metal toxicity for fish because calcium
and magnesium keep fish from absorbing metals such as lead, arsenic, and cadmium (which are
other polyvalent cations) into their bloodstream through their gills. The greater the hardness, the
harder it is for toxic metals to be absorbed through gills. In addition, hard water is usually also
high in alkalinity, which can help maintain pH levels that aquatic life need in order to survive.
Hard water can also affect fish osmoregulation, the process that controls the concentration of
internal body fluids. Hard water with more ion concentration is closer to body fluid levels,
making the job of osmoregulation a little easier for fish. Soft water or very hard water will disrupt
this balance and fish have to adapt their osmoregulation process.
13
Human Impact
Hardness is a characteristic of water that is not considered a pollutant and in most cases
not considered a major health-related concern. Although beneficial because calcium and
magnesium are essential minerals to a healthy human diet, hard water is generally considered a
nuisance rather than health benefit or threat.
Even though hardness can cause plumbing nightmares, hardness is also desirable
because it reduces corrosion rates in our pipes. This reduces the amount of lead (from lead
sodder), copper, zinc, and other metals from plumbing that may enter our drinking water. Unlike
hard water, soft water with few positive ions is more reactive to picking up cations such as metals
from pipes or the surrounding environment. The calcium coating on the inside of the pipes can
also help reduce corrosion.
Humans can increase the amount of hardness in waterways in a few ways. Drainage
from operating or abandoned mine sites can add calcium, magnesium, iron, manganese, and other
cations from the newly exposed rocks or overburden (soil and crushed rock removed from mining
operations). Some industrial discharges can be high in calcium and metals. Wastewater from
homes is often high in cations from household cleaning agents, food residue, human waste, and
from rinsing out the trapped calcium and magnesium from an ion exchange filter system when
changing the salt. All of these cations cannot always be removed by water treatment facilities
before discharging to a stream.
Water Quality Criteria
Because hardness is a characteristic of water and not a pollutant, there are no published
standards for overall hardness, calcium, or magnesium from the EPA. Government agencies do
hope that a level of hardness does exist so pipes are less corrosive, limiting the release of metals
that do need to be regulated.
Iron Information
Background Information
Iron is a very abundant element, and measurable concentrations of iron often exist in
natural waters, particularly in well water. A variety of common minerals contain iron, including
hematite (Fe2 O3), pyrite (FeS2), and some of the silicate minerals such as olivine (FeMg)SiO4,
and fayalite, Fe2SiO3. Although a high concentration of iron in drinking water will cause no real
health problems, communities often wish to remove it from the water supply for aesthetic
14
reasons, since high concentrations of iron can make water appear murky, smell and taste bad, and
stain plumbing fixtures.
Iron in water can occur in two oxidation states, Fe2+ or Fe3+. Iron (III) is the form of iron
found in ordinary rust (iron oxide, Fe2 O3). This form of iron is quite stable and will not react with
oxygen in the air, NaCl, or water. In contrast, iron (II) is very reactive towards oxygen in the air.
If a solution of Fe2+ is left exposed to air for several hours, much of the Fe2+ will be oxidized to
Fe3+ and a rusty-colored, gelatinous precipitate of iron (III) hydroxide, Fe(OH)3 will form.
Both oxidation states of iron form very insoluble hydroxides, with Fe(OH)3 being much
less soluble than Fe(OH)2. Because of the extreme insolubility of Fe(OH)3, very little iron Fe3+
will be dissolved in natural waters when the pH is in the 7.0-8.5 range. Most of the natural waters
that contain high levels of iron are groundwater with high concentrations of the fairly soluble
Fe2+. Since groundwater is protected from exposure to oxygen, any Fe2+ that dissolves from
surrounding minerals is not oxidized to Fe3+. Because Fe2+ is so much more soluble than Fe3+,
more iron can remain in solution.
15
The Background Chemistry and Iron and Alkalinity Determination Experimental Procedures are
modified from experiments in: Kegley, S.E.; Landrear, D.; Jenkins, D.; Gross, B.; Shomglin, K.
Water Treatment: How Can We Purify Our Water? Student Manual. John Wiley & Sons: New
York. 2000. p 35—45.
Background Chemistry
Determination of Iron in Water
The method used for iron determination makes use of coordination chemistry, or the
ability of transition metals like iron to bond to a substance that has an unshared pair of electrons
available. In many bonds, each atom contributes one electron. However, transition metals can
act as Lewis acids (electron pair acceptors) and tend to form bonds with compounds that are
Lewis bases (electron pair donors), species that contain an unshared pair of electrons. Examples
of Lewis bases include water (H2O), chloride (Cl-), and ammonia (NH3). These molecules or ions
can donate both electrons of an unshared pair of electrons to an empty orbital on the metal.
Groups that bind to metals in this fashion are called ligands.
This type of bond is characteristic of the transition metals. In fact, it is rare to find a
“bare” transition metal that does not have some other ligands bound to it because bare metal ions
are quite unstable and very reactive. Clearly, it is energetically favorable for a positively charged
16
ion to interact with a molecule or ion that has loosely-held electrons (non-bonding electrons). For
example, in aqueous solution, iron ions for bonds to six water molecules to make [Fe(H2 O)6]2+ or
[Fe(H2 O)6]3+. The abbreviation for these species, Fe2+(aq) or Fe3+(aq), indicates that the iron ion is
surrounded by water molecules, or solvated.
One of the most fascinating aspects of coordination compounds is their varied and
beautiful colors. The color usually depends on the ions or molecules bound directly to the metal
ion. For example, an aqueous solution of Fe(II) is a very pale yellow or green, but if certain
nitrogen-containing compounds are added, a deeply colored purple solution results. We can use
this property of coordinate bonds to measure the concentration of a metal ion in solution. If the
sample is treated with an appropriate ligand that will transform it into a colored species, we can
measure the intensity of the color and use this as an indication of how much iron is in solution. In
general, the more concentrated the solution is in iron, the more intense the color. In order to
obtain quantitative results for the amount of iron in solution, a technique called spectroscopy is
used, described in more detail in the following section.
One reagent that coordinates to iron to form a colored complex is an organic molecule
that has the trade name ferrozine. This molecule has two nitrogen atoms in a position to donate
unshared electron pairs to the iron to form a purple iron-ferrozine complex.
Ferrozine contains nitrogen atoms with unshared pairs of electrons that
will bind to Fe(II) ions in solution. The two nitrogens that are boxed in
the figure are the ones that bind to iron.
17
Ferrozine will not react with iron (III), so in the ferrozine solution, a mild reducing agent,
hydroxylamine (NH2 OH), is added to transform all iron (III) in solution to iron (II). A pH 5.5
buffer is also added to ensure that the solution stays within the optimum pH range for reaction of
ferrozine reagent with iron (II). The ferrozine solution is concentrated enough to ensure that
essentially all the iron (II) reacts to form the iron(II)-ferrozine complex.
Colorimetric Methods of Analysis
As you learned in the previous section, when ferrozine is added to a solution containing
iron (II), an iron (III)-ferrozine complex forms and the solution turns a deep purple color.
Because the intensity of this color is related to the amount of iron (II) present, we can measure the
concentration of iron (II) in different solutions based on their color. Chemists call this analytical
technique colorimetric analysis. Colorimetric analyses can be used to measure the concentration
of many different constituents in water by reacting the chemical species of interest with a specific
reagent that produces a colored compound in solution.
18
Colorimetric analysis is carried out using a spectrophotometer, an instrument that
measures the amount of light transmitted through a solution.
incident
light, I0
transmitted
light, It
lamp
detector
sample
monochromator
The spectrophotometer consists of a visible light source with a
monochromator that allows the analyst to select a small range of
wavelengths for the analysis. Light passing through the sample is
absorbed by the species of interest and the intensity of the transmitted
light (It) is detected.
When light is passed through any solution, some of the light is absorbed or reflected, and the
remainder is transmitted through the solution. The amount of light transmitted through a solution
depends on many factors, including the wavelength of the incident light, the color of the solution
(and hence the concentration of the substance of interest), and the path length of the light.
For a simple analogy, consider how a flashlight beam might behave if you were to shine
it through a glass of water. Clean water does not significantly absorb or scatter light, and you
perceive the light beam to be nearly as intense after passing through the water as it was before.
Now imagine adding a small amount of blue food coloring to the water. The solution would turn
blue, and the same beam of light transmitted through the solution would be dimmer than the light
coming out of the flashlight itself. The transmittance, T, of a solution is the ratio of the intensity
of the transmitted light, It, to the incident light, Io:
T=
It
I0
The Beer-Lambert Law
A spectrophotometer can be used to measure either the transmittance or the absorbance (A) of a
solution. As its name implies, absorbance is a measure of the amount of light absorbed by a
solution. The relationship between absorbance and transmittance is logarithmic:
19
A = -log T
The absorbance of a solution is directly proportional to the concentration of the substance of
interest. The relationship between absorbance and concentration is given by the Beer-Lambert
Law:
A = εbc
Where A is the absorbance (a unitless number determined by the spectrophotometer), b is the path
length of the light through the sample (typically 1 centimeter for a standard spectrophotometer), c
is the concentration of the substance of interest in the solution and ε is a constant that is specific
for the substance being analyzed. In practice, ε is determined experimentally for each chemical
species.
In general, the higher the concentration of the substance of interest in solution, the greater
the absorbance measured by the spectrophotometer, a plot of absorbance versus concentration,
also known as a standard curve, must be constructed for the chemical species. The slope of the
resulting line is εb. Using this slope value and the Beer-Lambert Law, the concentration of the
species of interest in any solution can be determined simply by measuring its absorbance.
20
Determination of Total Alkalinity
The ions that determine the total alkalinity of a water sample, CO22-, HCO3-, and OH-,
are all bases, and as such, react readily with acid to form water and neutral salts, thus reducing the
concentration of acid in water, as follows:
Carbonate:
H3 O+(aq) + CO32-(aq)  HCO3-(aq) + H2 O (l)
Bicarbonate:
H3 O+(aq) + HCO3-(aq)  H2CO3 (aq) + H2 O (l)
Hydroxide:
H3 O+(aq) + OH- (aq)  2H2O (l)
Reactions of this type are called neutralization reactions. We will take advantage of the
reactions just shown for the analysis of total alkalinity. Because we know the chemical equations
describing these reactions and therefore, the stoichiometry of the reacting species, we can
determine the alkalinity of a sample by determining how much acid can be added without making
the sample acidic. This can be accomplished through titrating the sample to a known endpoint.
For alkalinity determination, it has been shown that a pH of 4.5 best describes the situation when
all of the carbonate and bicarbonate anions in a solution are used up. Since these two anions are
the predominant alkalinity ions in natural waters, the endpoint has been defined to reflect this
reality.
Experimentally, total alkalinity is determined by titrating a known volume of sample with
acid, usually H2SO4 or HCl, in the presence of bromocresol green indicator. For example,
titration of the carbonate ions resulting from dissolution of calcium carbonate can be described by
the following neutralization reaction:
CaCO3(aq) + H2SO4 (aq)  H2 CO3 (aq) + CaSO4(aq)
The indicator bromocresol green is blue above pH 5.4 and yellow below pH 3.8. At the endpoint
pH of 4.5, the solution will be green.
The reporting of total alkalinity is usually simplified by assuming all acid neutralizing
capacity is due to the presence of CaCO3, and is thus reported as mg CaCO3 per liter.
(Alkalinity is frequently caused by substances other than CaCO3, so this reporting procedure is
arbitrary. However, the net effect in terms of acid neutralizing capacity is essentially the same).
Expected Total Alkalinity Values
The range of total alkalinity values for natural waters is typically between 30 and 5000 mg of
CaCO 3/L, with higher values occurring in regions that have alkaline soils. Rainwater usually has
very little total alkalinity (<10 mg/L) since it has contact with few minerals. Surface waters
21
generally have total alkalinities less than 200 mg/L, while groundwater total alkalinities are
frequently much higher, sometimes over 1,000 mg/L due to higher partial pressure of CO2 in the
subsurface from microbial degradation of organic matter underground. The underground CO2
reacts with water to form bicarbonate and acid:
CO2 (g) + 2H2O (l)  HCO3- (aq) + H3 O+ (aq)
The acid (H3 O+) formed in this reaction reacts with calcium or magnesium carbonate in
the surrounding rock to dissolve these compounds and produce bicarbonate, which is the main
contributor to total alkalinity.
H3O+ (aq) + CaCO3 (s)  H2 O (l) + HCO3- (aq) + Ca2+ (aq)
22
Fe2+ Determination
Chemistry learning goals:
Lewis acids/Lewis bases
Solvation
Visible spectroscopy
Colorimetric Analysis
Coordination Chemistry
Beer-Lambert law






Prelab questions:
1) What chemical compounds are associated with iron in water? With what minerals are
these associated?
2) Define coordination chemistry, incorporating the terms Lewis acid, Lewis base, metal
ions, and ligands into your answer. Draw the reaction to form a final ferrozine-Fe
complex (you may use N N to represent ferrozine). Remember to include charge!
3) Label the below graph as would be appropriate to determine the concentration of an
unknown using spectrophotometry and the Beer-Lambert law. What is the BeerLambert law? Define your labels for the graph.
_________________
Determination of Unknown Concentration by
Spectrophotometry
0.6
0.5
0.4
0.3
0.2
0.1
0
0
2
4
6
_________________
23
8
10
Determination of Alkalinity
Learning goals/Chemistry objectives



Review of titration chemistry
Neutralization reactions
Indicator chemistry
Prelab Questions
1) What minerals (give chemical formulas) are the main source of alkalinity in natural
waters?
2) Streams that have low total alkalinity are particularly susceptible to damage from
what anthropogenic (man-made) pollutant?
3) Consider the carbonate ion, CO32-, a major contributor to alkalinity in natural waters.
Is this ion a Brønsted acid or a Brønsted base? Explain, using a chemical equation to
show the acid-base reaction.
4) Why is total alkalinity also referred to as “acid neutralizing capacity”? Write the
relevant chemical equations that support your answer.
5) Why is an indicator used in the procedure to determine total alkalinity?
24
Iron and Alkalinity Determinations
Hazards
The hydroxylamine solution used to reduce the iron is concentrated and can cause severe
irritation and burns upon skin contact or inhalation. Prolonged exposure to ferrozine may
cause dermatitis. Students should make sure that their cuvettes are well sealed before
mixing the iron/ferrozine/hydroxylamine/buffer solution, and should wear gloves and
protective eyeware. The titration involves a dilute solution of sulfuric acid, which can
cause burns.
Part A.
Part B.
Part C.
Part D.
Calibration of the Spectrophotometer
Determination of Iron in an Unknown Sample
Alkalinity Determination by Titration of an Unknown Water Sample
AA Analysis of Samples for Mg, Ca, Na, and Fe concentrations
The overall goal of these experiments is to use a variety of methods to determine ion
concentrations of water samples from the Spring Creek Watershed. Furthermore, you
should begin to understand the usefulness of colorimetric, titration, and atomic absorption
analysis in the determination of ions in a water sample.
Section A. Calibration of the Spectrophotometer
Goals:
(1) To become comfortable with the use of a spectrophotometer. (2) To calibrate
the instrument for iron analysis.
Discussion
The spectrophotometer must be calibrated to correlate the absorbed light from a
sample to a particular iron concentration. Thus, you must begin by running four
reference standards with known iron concentrations: 0.1 ppm, 0.5 ppm, 1.0 ppm,
and 3.0 ppm. In addition, the analysis also requires a blank, or a solution with all
of the reagents added except iron. Water and other reagents absorb a small
amount of light, and this blank allows you to measure that amount so that you
can subtract it from the light absorbed by the purple iron-ferrozine compound.
Experimental Steps:
1. Set the spectrophotometer analysis wavelength to 562 nm.
2. Prepare a table in your lab notebook with columns for Sample ID, concentration, and
absorbance.
3. Label five cuvettes with numbers corresponding to the concentration of each standard and the
blank sample. Fill the cuvettes with the appropriate sample and note the color intensity of
each in your notebook.

What trends do you observe?
25
4. Wipe the outside of the blank cuvette with a KimwipeTM and place it in the instrument.
Follow the instructions for using your spectrophotometer to first zero the instrument with the
blank solution.
5. Using the same procedures, place each of the standards in the spectrophotometer and read the
Absorbance. Enter the data in the table in your notebook.

Construct a standard curve by plotting Absorbance (y-value) as a function of
concentration (x-value). Draw the best straight line through the data points and calculate
the slope of the line.
26
Section B. Determination of Iron in an Unknown Sample
Goals:
(1) To prepare a sample for spectrophotometric analysis. (2) To determine the
concentration of iron in your unknown sample
Experimental Steps:
1. Gravity filter your sample.
2. Obtain a clean, dry cuvette.
3. Using a 10.0 mL graduated cylinder, measure out 3.0 mL of your filtered
sample, and add it to a cuvette. Make this measurement as accurately as
possible.
4. Using a calibrated micropipet (make a new one if necessary by drawing out
the pipet, cutting off the tip, and testing how many drops it takes to deliver
1.0 mL of liquid), add 0.1 mL of the ferrozine reagent. Cap the cuvette and
invert it to mix the solution. Allow it to stand for at least 15-30 seconds.
5. Using the calibrated micropipet, add 0.1 mL of pH 5.5 buffer solution to the
cuvette. Cap the cuvette and invert to mix the solution. Wait at least two
minutes for the color to develop.
6. Zero the calibrated spectrophotometer using the blank, then place the sample
cuvette into the spectrophotometer to obtain an Absorbance reading. Record
this value in your laboratory notebook.
7. For more precise measurements, analyze at least three replicate samples.
Section C. Alkalinity Determination by Titration of an Unknown Water Sample
Goal:
To perform a titration of an unknown water supply to determine the total
alkalinity of the sample.
Discussion:
In this experiment, you will determine the concentration of total alkalinity in
your sample by titrating the sample with acid in the presence of an acid-base
indicator.
Experimental Steps:
1. Prepare the sample for the total alkalinity determination by measuring 20 mL
of sample into a clean 250 mL Erlenmeyer flask. Add 25 mL of distilled
water and a drop of bromocresol green indicator. If the pH of your sample is
below 4.5 (indicator is yellow), you do not need to determine alkalinity, since
the sample will have no acid neutralizing capacity. If the pH is above 4.5,
proceed to the next step.
2. Fill a calibrated buret by filling it to the top of the graduated area with 0.0100
M sulfuric acid (H2SO4). Make sure to write down the exact concentration of
the acid from the bottle.
3. Check that the buret tip contains no air bubbles. If it does, allow a few mL of
acid to flow out. You may need to tap the buret to dislodge the air bubbles.
4. Take the initial volume reading from the buret. Always have your eyes at the
same height as the liquid and read the value at the bottom of the meniscus to
one decimal point (e.g., 10.1) and estimate the second decimal point.
5. Swirling the sample gently, titrate the solution with 0.0100 M H2SO4 to the
pH 4.5 endpoint, where the indicator changes color from blue to green. If
27
you pass the endpoint by more than a drop or two, the solution will be yellow
and you should re-do the titration. Note the final volume reading on the
buret when you reach the endpoint.
Calculations:
1. From the volume of acid used to titrate the sample, calculate the total alkalinity in millimoles
of CaCO3 per liter. This is a four-step process:
 Use the concentration of H2SO4 to convert mL of H2SO4 used in the titration to moles
of H2SO4.
 Use the mole ratio of reactants in the reaction of CaCO3 and H2SO4 to determine the
moles of CaCO3 that must have been present initially.
 Use the volume of the sample to obtain moles of CaCO3 per liter of sample.
 Convert moles per liter to millimoles per liter and report your concentration of
alkalinity in millimoles per liter.
2. The convention in water chemistry is to treat alkalinity as if it were all due to CaCO3. To
report your alkalinity values in a way that can be compared to the water quality standards,
convert your alkalinity in millimoles per liter to total alkalinity in mg of CaCO3 per liter.
Section D: Atomic Absorption Determination of Ca, Mg, Na, and Fe
Goals:
(1) To utilize an analytical tool that can selectively measure ions in a mixed
solution. (2) To determine the concentration of Ca2+, Mg2+, Na+, and Fe2+ in
your water sample.
Discussion:
The first two techniques used in this lab analyzed specific components of the
water sample, Fe2+ and CO32-. With AA, you can analyze four (or more) ions
from the same sample solution, thus measuring a number of different water
quality parameters at once. Calcium and magnesium are the key ions that
contribute to the hardness of water, and are the most abundant cations in
solution. Sodium is included for later charge balance analysis. The iron
analysis will enable you to compare your results from the AA test with those
from the Spec 20 measurements.
Experimental Steps:
1. Filter your water samples by gravity filtration and take the samples upstairs
to the AAs.
2. Conduct analyses for Mg2+, Ca2+, Na+, and Fe2+. Use this data in the
determination of total positive charge in your solutions. For a refresher on
how AA works, consult the course website.
Post-lab questions
1) Why is an accurate calibration important? What is the sensitivity of the spectrophotometer?
What role does the buffer play?
2) In your own words, describe how the spectroscopic method of determination works. List an
advantage and a disadvantage of this technique compared to atomic absorption. (Consider
28
3)
4)
5)
6)
7)
8)
instrument sensitivity, ease of sample preparation, time/effort required, repeatability, sources
of error).
A student in a different class comes into the lab to use the spectrophotometer for a phosphate
analysis that produces a yellow solution. When he places his highest standard in the
spectrometer, the absorbance reading is nearly zero. What could be wrong with his
procedure? List the most likely problems.
What is the total positive charge in your water sample? Combine charges from your AA
determinations (Mg, Ca, Na) and Fe experiment. Assume Fe is in the 2+ oxidation state.
Watch for charge number (+1, +2, etc.)! You will need this number for your next lab writeup.
If the only source of alkalinity in a sample is sodium hydroxide (NaOH), how would the
calculation for total alkalinity be different?
What is the total negative charge in the water based on these calculations? You will need this
number for the ion chromatography lab write-up.
List advantages to using a method like AA for determining ion concentrations. Give a
suggestion as to why some metals may not be analyzed using AA (think about concentrations,
absorbance/emittance patterns, and other chemical properties).
Which method of analysis (AA or Spec 20) do you think provides more accurate values?
Give at least one reason why.
29
Experiment #8:
Halocarbon
Remediation with
Re duced Iron
Kinetics and
Liquid/Liquid
Extraction
30
Source s f or this e x per ime nt
n t are refere nce d at the e nd o f the ex per im en t.
KINETICS
Background information
Two main factors govern the extent to which a reaction will proceed: the thermodynamic
and the kinetic properties of the system. Investigating the thermodynamic properties is a static
process, and involves the study of the initial and final states of a chemical system. The study of
kinetics, however, involves the dynamics of a chemical system. It is defined as the study of the
various factors that control the rates of reactions and the study of mechanisms by which reactants
become products.1
THE STUDY OF KINETICS IS HUGELY IMPORTANT AND PERTINENT TO
STUDIES OF ANYTHING THAT GOES ON IN THE ENVIRONMENT. Read that last
sentence again. The rate at which things happen explains the selective toxicity of pesticides, how
ozone is depleted over Antarctica, how chemicals are degraded in the environment, and how
rocks and minerals slowly dissolve into a water supply. Rates of reactions vary widely – while it
takes days for much calcium or magnesium carbonate to dissolve into a supply of water, it takes a
fraction of a second for the bicarbonate produced to undergo an acid-base reaction with nitric acid
in rain.
The rate of a chemical reaction indicates how the concentrations of reactants and
products change over time. Factors that contribute to this include1:

The temperature of the system

The concentrations of the various chemical species present in the system

The presence of catalysts

The rate at which reactants can mix or diffuse together
Those who study kinetics investigate all four of these areas for a given reaction, but these factors
can be used to speculate as to how the rate of a reaction will change when the reaction conditions
change.
Kinetics studies are particularly useful in determining the mechanism of a given reaction.
The mechanism is the actual pathway that molecules follow during a reaction, and most of our
insight into these has come about by trying to match mechanistic theories to kinetic data.1
Understanding the mechanism of a reaction is particularly important in the field of environmental
chemistry and toxicology for two reasons – first, that the toxicity of a chemical is generally based
on how it interacts with other molecules in the body, and second, that the primary way to
31
convince a company of why a substance is dangerous to humans and the ecosystem is to
specifically describe the pathways the chemical takes in the environment.
Background Chemistry – Rate Laws and Activation Energy2
A convenient way of studying the effect of concentration on reaction rate is to determine
the initial rate (at the start of the reaction) as a function of the initial concentration of one reagent
while keeping the concentration of all other reactants constant. The results from a series of
experiments carried out in this way, at some known, constant temperature, are expressed in the
form of a rate law. The rate law can be expressed generally by the following equation:
Rate = k[reactant 1]m[reactant 2]n…
The variable k is the rate constant for this reaction, and is constant for a given set of temperature
and pressure conditions. The exponents m and n are called reaction orders, and the sum of these
individual orders is the overall reaction order. The reaction that you will be conducting in this lab
is a pseudo-first order reaction. In reality, the number of active sites on the iron should factor into
the rate law, but the number of these has been determined to be approximately constant with the
given experimental conditions.
A first-order reaction is the easiest to understand; it is one whose rate depends on the
concentration of only one variable. The rate law for a first-order reaction is:
Rate = k[A]
Through calculus magic, this can be transformed to relate the concentration of A at any point in
time, At, to its initial concentration, A0:
ln[A]t - ln[A]0 = -kt
or
ln([A]t/[A]0) = -kt
where the function “ln” is the natural logarithm.
The above equation is in the form of the equation for a straight line, y = mx + b, with a yintercept, b, of 0.
ln([A]t/[A]0) = -kt
The rate constant, k, can then be determined by plotting ln([A]t/[A]0) versus time, with the slope
of the line equal to –k.
32
Kow Studies and Liquid/Liquid Extraction
Introduction
One of the primary chemical properties that determines the toxicity of a chemical in the
environment is the way the compound partitions between water and organic material, including
the tissues of all living creatures. Partitioning is a chemical characteristic that describes the
extent to which molecules of a compound move preferentially into one type of solvent over
another (usually in an organic/aqueous solvent system). The concepts behind partitioning can be
most easily understood by thinking in terms of salad dressing – if you have an oil layer and a
vinegar layer (organic and aqueous, respectively) and you add a second oil to the mixture, it will
mix primarily with the first oil and not with the vinegar. Similarly, if you add sugar, which has
lots of –OH groups, it will dissolve to a greater extent in the vinegar (where there are more –OH
groups, including water, and similarly small-ish molecules) than in the oils (which have fewer –
OH groups and are composed of much larger molecules).
One of the most dramatic illustrations of partitioning can be seen in the bioaccumulation
of DDT in food chains, resulting eventually in levels high enough to cause horrible reproduction
problems in bald eagles. At first, scientists did not understand how the DDT concentration in
eagle tissue could be so much higher than the concentration in the surrounding lakes and streams.
They then discovered elevated DDT levels in the tissues of the fish the eagles ate, and in smaller
organisms that the fish ate. Eventually, scientists realized that the partitioning coefficient for
DDT is so high that the nontoxic concentrations of DDT in water quickly accumulated to
dangerous levels in fatty tissue.3 It is in part because of this atrocity that the maximum allowable
levels of pesticides are kept extremely low. (Note: If you want to learn more about this and other
effects of pesticides, I highly, highly, highly recommend reading Rachel Carson’s Silent Spring.
Although at the time of the book’s publication many thought that Carson was totally wrong in her
conclusions, many of her insights into the environmental fates of pesticides have been proven
33
true. It’s pretty incredible to read something that was brand new in chemistry forty years ago and
still holds true, with few modifications, today).
The standard method for determining partition coefficients to predict a chemical’s
activity in the environment is to compare a molecule’s solubility in water to its solubility in
octanol:4
Kow = (solubility in octanol)/(solubility in water)
While any solvent system can be used, even (if you could measure it well) oil/vinegar, the
octanol/water system has been shown to closely model the partitioning of chemicals in the
ecosystem.
An organic molecule with a low water solubility that is soluble in an organic solvent, like
DDT, has a high Kow, while one with a high water solubility will have a low Kow. The
partitioning coefficient then serves as a good indicator as to whether a molecule will be
problematic in the ecosystem because it is generally necessary for a molecule to have a mid-range
value of Kow to be toxic (Kow = 4-7)4. Molecules with extremely high Kow values will either not
be soluble enough in water to be a problem (they will stay where they are applied until broken
down) or will be repelled by the ionic, hydrophilic outsides of cell membranes. Molecules with
extremely low Kow values pass straight through organisms systems without being absorbed into
cell membranes. Similar considerations influence drug design – to be effective a drug must be
water soluble enough to be circulated throughout the body in the bloodstream, but lipid (fat)
soluble enough to be absorbed into tissue and kept there for a reasonable amount of time
(Tylenol, for example, lasts about three hours).
Background Chemistry
While determining the value of Kow is important in the characterization of an organic
molecule that is being released into the environment, the methods by which Kow is determined are
too complicated to pursue in this course. If you are interested in the methods used too determine
this, see the references at the end of the lab. We will, however, utilize a technique called
34
liquid/liquid extraction that exploits the partitioning properties of substances. This technique has
numerous applications and is used very broadly throughout chemistry (and by anyone who goes
on to take organic chemistry). We will use it primarily to extract organic molecules from dilute
aqueous solutions of the molecules.
A partition coefficient, K, is defined by the equation5:
K = solubility of species in organic (g/100mL)/solubility of species in water (g/100mL)
When a system has reached equilibrium, the molecules will naturally distribute themselves
preferentially in the solvent in which they are more soluble. By varying the solvent system, one
can change the value of K such that a greater separation is achieved.
Because K is not infinite, not all of a solute will reside in the organic layer after a single
extraction. If K is such that 90% of a solute X can be extracted in a single extraction, 10% of the
solute remains in the water. Repeating the extraction with a new portion of organic solvent
recovers 90% of this 10%, or 9% of the original concentration, for a total of 99% recovery. A
general rule is that many small extractions are more efficient than a single large extraction. This
is proved mathematically by the equation:
fraction extracted into B = (1/(1+ VB/VAnK))n
This gives the fraction of a solute extracted into a solvent B where n is the number of extractions
performed, K is the distribution coefficient, VA is the volume of solvent A and VB is the volume
of solvent B5.
Applications (for this class)
This technique will be used in Labs 7 and 8, Reduced Iron Degradation of Halogenated Organics.
You will have to trust the technique and understand the chemistry behind it because you will not
be able to see any changes occurring. We will be using gas chromatography to determine the
35
extent to which the iron particles successfully degrade the pesticide of choice. To use GC,
however, the sample must be free of all traces of water. Doing a liquid/liquid extraction gives us
a means of isolating the compound of choice in an organic solvent while also drying the sample.
36
Reduced Iron Remediation of Contaminated Water
Taken from: Balko, B.A.; Tratnyek, P.G. “A Discovery-Based Experiment Illustrating how Iron
Metal is Used to Remediate Contaminated Groundwater.” J.Chem.Educ. Online Supplement.6
In this experiment, you will investigate the chemistry behind iron permeable reactive
barriers (iron PRBs), a new technology that is widely used to remediate contaminated
groundwater. Contaminant remediation involving iron PRBs is a redox process: the iron metal
undergoes oxidative dissolution while the contaminant is reduced. The reaction is complicated,
however, by the fact that it involves a surface that changes due to the development of a layer of
rust (iron oxide) on the iron.
It has long been known that the oxidation (i.e., corrosion) of metals such as iron, tin, and
zinc can bring about the reduction of halogenated organics. In the late 1980’s, researchers at the
University of Waterloo rediscovered these redox reactions while investigating groundwater
contaminated with halogenated solvents. The researchers recognized that these reducing metals
could be used for remediating contaminated groundwater by constructing permeable reactive
barriers (PRBs) composed of one of the reducing metals. A PRB is a zone comprised of granular
metal that extends below the water table and intercepts the flow of contaminated groundwater; as
the contaminants pass through the PRB, they are reduced to non-toxic compounds, and, ideally,
the groundwater that emerges is free of hazardous substances. Figure 1 shows a schematic of a
PRB. Iron is the current metal of choice for PRBs because it is readily available, inexpensive,
nontoxic, and a good reducing agent.
37
Figure 1: Water infiltrating soil contaminated with chemical waste forms a plume of contaminated
groundwater. This plume travels in the direction of groundwater flow and spreads through the water table
if left untreated. When contaminated groundwater flows through an iron permeable reactive barrier, the
contaminants are reduced by the oxidation of the iron. The plume can then be considered treated. (taken
from http://www.powellassociates.com/sciserv/3dflow.html, and Balko and Tratnyek [6])
Iron PRBs, colloquially known as iron walls, have many advantages over more traditional
groundwater remediation technologies. First, iron walls are able to remediate many contaminants
in addition to halogenated solvents, such as pesticides, munitions, nitrate and heavy metals
(including chromium and uranium). Second, iron walls are a “passive way” to remove
contaminants from groundwater in that they require little maintenance after installation. because
of these advantages, as well as the prevalence of contaminants that iron walls are capable of
removing, iron PRBs are now widely used in North America and Europe to clean up
contaminated sites.
Chemistry of Permeable Reactive Barriers
The degradation of contaminants by iron metal, Fe0, can be explained in terms of
textbook redox chemistry. Consider the degradation of carbon tetrachloride, CCl4, by Fe0. The
anodic half-reaction is the oxidative dissolution of Fe0 (Equation 1) and the cathodic half-reaction
38
of interest, assuming a hydrogenolysis reaction mechanism, is the reduction of CCl4 to CHCl3
(Equation 2). The E0 for the next cell reaction at pH 7 is
1.11 V.
Fe0  Fe2+ + 2e-
E0 = 0.44 V
CCl4 + 2e- + H+  CHCl3 + Cl-
E0 = 0.67 V
(2)
E0 = 1.11 V
(3)
(1)
_________________________________
CCl4 + H+ + Fe0  CHCl3 + Cl- + Fe2+
It should be noted that the product of the degradation reaction, CHCl3, is a hazardous and
regulated substance. While CHCl3 is not readily degraded by Fe0, it is much more biodegradable
than CCl4 so groundwater contaminated with CCl4 should be cleaned by an iron PRB.
It is possible, however, that iron is not directly responsible for contaminant reduction by
iron walls. Fe2+ and/or H2 will be present in the system due to the reduction of dissolved oxygen
and water by Fe0:
2Fe0 + O2 + H2O  2Fe2+ + 4OH-
(4)
Fe0 + 2H2 O  Fe2+ + H2 + 2OH-
(5)
Thermodynamically, both Fe2+ and H2 are capable of reducing some contaminants.
The degradation of a typical halogenated organic contaminant, RX, by Fe0 typically
obeys the following rate law:
Rate = -d[RX]/dt = -k[Fe active sites][RX]
(6)
This equation can be simplified into a first-order kinetic equation, since in most experimental
systems it has been found that the concentration of iron active sites does not vary significantly
during the course of the degradation reaction.
-d[RX]/dt = -kobs[RX]
(7)
39
Thus, the kinetics of degradation of groundwater contaminants by iron metal generally are pseudo
first-order. It should be emphasized that kobs is not a true first-order rate constant and will depend
on the concentration of active sites on the iron metal. Thus, kobs should be proportional to the
surface area of the iron particles as well as the mass of the iron present. The following
relationship exists between kobs, the specific surface area of the iron, as, the mass concentration of
iron, ρm, and the specific reaction rate constant, kSA.
kobs = kSAasρm
(8)
40
The following information is from various sources as listed, but not from the article cited at the
beginning.
More Background Information and Chemistry
While Fe0 remediation techniques have proved useful enough to be implemented around
the world for remediation chemistry, the rates of reaction are noticeably slow. One solution to
improving these rates has been the use of bimetallic particles. Nickel, palladium, platinum, and
zinc have all been deposited on iron particles and generally show a catalytic effect. Of these
metals, nickel has been shown to have the best combination of effective catalytic properties and
low toxicity, and, thus, much research has been invested into nickel-coated iron particles.7 Penn
State researchers are working to develop Fe nanoparticles that can be suspended in water and
sprayed onto or injected into the soil at a contaminated site.8,9 Because the particles do not
aggregate, they can trickle down through the dirt. This method would eliminate the necessity of
digging up huge plots of land to insert the iron barriers.
The nickel particles on the iron have a catalytic effect on the redox process. Basic
thermodynamics tells us that any reaction has an activation energy that must be overcome for the
reaction to proceed. Simplistically, two reactants, AB and CD, going to form two products, AC
and BD, could have an energy reaction coordinate diagram like the following:
41
The difference in energy between AB and CD and the top of the peak is the activation energy, Ea.
What a catalyst does is reduce the activation energy for such a system without actually having its
chemical composition changed at the beginning and end of the reaction. If we added a catalyst,
X, to the above system, it would just be X on the product side as well. When Fe and
CHCl==CCl2 (trichloroethylene, a common halogenated solvent and problematic groundwater
contaminant), are reactants in the above reaction, the value of Ea is fairly high and thus, the
kinetics of the reaction are fairly slow (even if the energy differences between the products and
reactants are large). Adding nickel to the iron reduces the barrier, thus improving the rate of the
reaction.
42
The first part of this experiment is modified from that described by Balko and Tratnyek6 The rest
has been developed at Penn State.
Experimental Procedure
In this lab, you will first examine the rate at which Fe degrades a dye (a model pesticide). You
will then compare the rates at which plain and Ni-catalyzed Fe0 filings degrade the common
pesticide 2,4-D (2,4-dichlorophenoxy acetic acid). You will be using liquid/liquid extraction
techniques to extract the 2,4-D and its degradation products from the contaminated water supply.
You will the analyze these samples using a gas chromatograph (GC) to determine the ratio of
degraded:undegraded compound in your sample.
Flow Chart of the Experiment:
Day 1:
A) Kinetics studies
B) Preparation of filings and reaction initiation
C) Sample collection
D) Liquid/liquid extraction exercise
E) Isolation of contaminant
Day 2:
A) Collection/Isolation of final samples
B) GC analysis of standard
C) GC analysis of samples
D) Determination of effectiveness of the remediation technique and comparison of the
effectiveness of catalyzed and non-catalyzed particles
Hazards
This experiment utilizes two concentrations hydrochloric acid, which causes burns upon
contact and is toxic upon inhalation. The nickel chloride solution used to prepare the nickelcoated filings is acidic and toxic, may cause allergic skin or respiratory reactions. In addition,
nickel known carcinogen upon prolonged exposure. Methanol is toxic and can cause
43
blindness if ingested in high enough amounts. Dichloromethane is readily absorbed through
the skin and is a potential carcinogen and mutagen. 2,4-D is toxic if swallowed or inhaled,
and is a potential carcinogen and teratogen. It also may cause allergic reactions in sensitive
persons.
44
Part A: Kinetics
The Experiment
This experiment uses many of the same techniques you learned in the iron determination
experiment. Therefore, you will be challenged to develop this experiment on your own from the
information provided below.
There are six questions you will need to answer over the course of this experiment.
1.) What is the wavelength of maximum absorption of indigo carmine dye? (Hint: the visible
spectrum goes from 400-700 nm)
2.) Is Beer’s Law followed by indigo carmine dye? (Think about what you need to plot to
determine whether the relationship holds).
3.) What is the rate at which the iron gets rid of the dye? To do this, you need to plot ln(At/Ao)
vs. t. What does the slope of this line give you?
4.) Everyone will run a test with [standard Fe0 filings]. You must then run two more tests in
which you’ve changed a single experimental variable. You will have available:

iron particles washed in 0.1 M HCl

iron particles washed in 5% H2 O2 solution

finer iron particles

coarser iron particles

a more concentrated dye solution
For the variables you have chosen, determine the new value of the rate constant. Explain
why the rate constant changed the way it did.
5.) What effect would each of the variables available to you that you didn’t choose to use have
on k? At least guess.
6.) Why is it important to prevent the re-oxidization of the dye? In what direction will your
value of k be shifted if the dye is re-oxidized?
45
Experimental Procedures
Step 1: Calibrate the spectrophotometer using available standards.
Step 2: Add ~1/8” iron filings to your cuvette, and fill with the dye solution. Cap the cuvette and
shake it up. Try to have the cuvette as full as possible, because any excess oxygen in the
system may re-oxidize the dye and cause errors in your value of k.
Step 3: Monitor dye concentration at 1-minute intervals (be sure to take an initial reading) for a
10-15 minute period. Be sure to shake the cuvette between intervals.
Step 4: Obtain a fresh sample of dye and iron filings (be sure to change a parameter) and again
monitor the reaction at 1-minute intervals over a period of 10 minutes.
Step 5: Choose another variable, and repeat the measurement process.
Step 6: Plot ln(At/A0) vs. t to determine rate constants for the reactions. Answer questions 1-6.
46
Part B: Preparation of filings and reaction initiation
Goals:
(1) To prepare reactive samples of iron filings and nickel-catalyzed iron filings
for the reduction of 2,4-D. (2) To start the degradation process.
Discussion:
When exposed to air in the presence of water, iron rusts quickly, forming a layer of
Fe2 O3 on the surface of the particles. This iron is in an oxidized state and cannot
be used in the oxidation of halogenated organics. Washing the iron particles
with acid serves to remove this oxide by dissolving it. The particles being
catalyzed are catalyzed through a redox reaction with nickel ions in solution that
you should be able to predict given the appropriate activity series.
Experimental Procedure:
1.) Obtain 2 x 1 g samples of iron filings from the front of the room, and place
each into a 125 mL Erlenmeyer flask.
2.) Into one flask place 35 mL 1.0 M HCl.
3.) Into the other 125 mL flask place 35 mL of NiCl2 solution (careful, it is toxic
and acidic!).
4.) Swirl each flask for about ten minutes and then gravity filter the samples
through the filter paper provided.
5.) While the particles are still on the filter paper, rinse each set of filings with
methanol. Allow them to dry in the air.

Why do you use methanol instead of water to rinse the filings?
6.) Transfer the filings to clean flasks. Add 50 mL of the contaminated water
supply to each flask. Cover each with parafilm, and purge the atmosphere
with nitrogen.
47
Part C: Sample Collection
Goal:
To periodically sample the contaminated water solutions.
Discussion:
To determine the rate at which the reaction progresses, it is necessary to
determine the extent to which the contaminant has been degraded at different
points in time. An analysis of these samples should enable you to calculate
the fraction of the 2,4-D sample that has been degraded at different points in
time.
Experimental procedure:
1.) At 45 minute intervals, take 5 mL samples of the contaminated solution (use
a plastic pipet for this).
Note: While this will increase the ratio of active sites to contaminant
molecules, it is unlikely that the active sites are ever saturated and thus, this
should not effect the rate of degradation. By just removing a portion of the
reaction solution, you will not change the relative concentrations of the
chemicals in solution, thus allowing for more accurate reactivity
measurements.
48
Part D: Liquid/Liquid Extraction
READ THE BACKGROUND INFORMATION ON OCTANOL-WATER PARTITIONING
COEFFICIENTS, PLEASE.
Goal:
To extract the 2,4-D into dichloromethane (CH2Cl2) using liquid/liquid
extraction.
Discussion:
Water is not good for the GC column and prevent an accurate
determination of other molecules in solution. Performing liquid/liquid
extraction enables you to extract the species of interest (2,4-D) from a
water solution into CH2 Cl2. The same partitioning coefficients that
govern how a chemical partitions between biota and water are also
responsible for the success of techniques like liquid/liquid extraction.
Experimental Procedure:
1.) Obtain a small (50 mL) separatory funnel and set up a 2” ring on a ring stand.
2.) Make sure that the funnel is in the closed position. Then add one of your
water samples to the funnel.
3.) Add 10 mL CH2 Cl2 to the funnel.
4.) Swirl the funnel gently for about two minutes. (See TA demonstration).
Every 30 seconds, vent the funnel by pointing the bottom slightly upwards
and towards the hood and turning the valve once to release any built-up
pressure.
5.) Set the funnel in the ring stand and remove the stopper.
6.) If your solution is not an emulsion – that is, if you can distinguish between
the two layers – drain out the bottom layer into a 25 mL flask. THIS
CONTAINS YOUR SAMPLE! Empty the water (top layer) into a 100 mL
beaker – it is now waste, and we will dispose of it all at the end.
49

Why is the CH2Cl2 on the bottom?
7.) Add a microscoopula’s worth of anhydrous sodium sulfate (anhyd Na2SO4)
to the flask with the CH2Cl2 (enough to thinly cover the bottom). Swirl the
flask. As long as some particles of Na2SO4 are freely flowing (as opposed to
clumped together), decant your CH2Cl2 into a 20 mL vial. Try to keep out
any Na2SO4 particles. If the Na2SO4 all clumps immediately, add some more
until you get freely moving particles.

What does the Na2SO4 do in the CH2Cl2 solution?
8.) Add a small amount (just enough to cover the bottom) of Na2SO4 to the 20
mL vial. Label and cap the vial.
9.) Repeat steps 2-8 for each of your samples. Be sure to label the vials
appropriately!
References
1
Thompson, S. Penn State Version of Chemtrek: Small-scale experiments for General
Chemistry. Prentice Hall: Englewood Cliffs, NJ. 2000.
2
Brown, T.L.; LeMay, H.E. Jr.; Bursten, B.E. Chemistry: The Central Science. 7th ed. Prentice
Hall: Upper Saddle River, NJ. 1997.
3
Carson, R. Silent Spring. Houghton Mifflin Company: NY. 1962.
4
Baird, C. Environmental Chemistry. 2nd ed. W.H. Freeman and Company: NY. 1999.
5
Minard, R.D. Lab Guide for Chemistry 431-W: Organic and Inorganic Preparations and
Qualitative Organic Analysis. Penn State University, 2001.
6
Balko, B.; Tratnyek, P.G. J. Chem. Educ., 2001, 78, 1661.
7
Nyer, E.K.; Vance, D.B. Ground Water Monitoring and Remediation. Spring 2001. 41-46.
8
Schrick, B.; Hydutsky, B. W.; Blough, J.L.; Mallouk, T.E. “Delivery Vehicles for Zero-valent
Metal Nanoparticles in Soil and Groundwater.” Chem. Mater. 2004, 16, 2187-2193.
9
Schrick, B.; Blough, J.L.; Jones, A.D.; Mallouk, T.E. “Hydrodechlorination of
Trichloroethylene to Hydrocarbons Using Bimetallic Nickel-Iron Nanoparticles.” Chem. Mater.
2002, 14, 5140-5147.
50
Experiment #9:
Halocarbon
Remediation with
Reduced Iron
Sample Analysis
51
Introduction
Gas chromatography (GC) is a separation technique that operates on principles similar to
those you learned in Experiment 3. It is often used to analyze samples for a variety of organic
compounds. The detection limit for GC is fairly low – you need only a minute amount of a
diluted sample (usually a drop of sample diluted in 1.5 mL CH2Cl2) to determine the presence of
a particular compound. In the undergraduate organic labs here, GC is most commonly used to
determine the purity of student products or to determine isomer ratios. However, its applications
extend to drug tests, the analysis of trace hydrocarbons and other pollutants in air, and the
detection of pheromone sex attractants in insects.1
Background chemistry
Taken from Thompson’s Chemtrek1
A basic GC system consist of a carrier gas, a heated sample injection port, a separating
column, and a detector. Commercially available instruments cost $5,000-$50,000 and often come
with a dedicated computer for data collection, storage, and interpretation. The GC flow
schematic is shown below:
Schematic of a Typical GC
1
Thompson, Stephen. Penn State Version of Chemtrek. Prentice Hall: Englewood Cliffs, NJ.
2000.
52
The carrier gas is usually a pure, inert gas (generally He, Ar, or N2) stored in a pressurized tank.
The flow rate of the mobile phase must be very carefully controlled in GC because the rates of
migration of all components are dependent on it. Various pressure gauges, flow controllers, and
meters accomplish exact carrier gas flow control.
The samples to be analyzed by GC may be gases, liquids, or solids. Solid and liquid
samples must be volatilized; thus, they must be heated as they are introduced into the injection
port. Generally, a very small sample volume is needed – on the order of 0.1-50 µL. The
volatilized sample is swept onto the separating column by a flowing stream of carrier gas. The
two main types of column in general use are shown below:
Two types of GC column
Packed columns are relatively short because of the high pressure required to push the gases
through the stationary phase. These columns are inexpensive and therefore widely used.
Capillary columns are much narrower and can be much longer because of the hole all the way
through the column. Capillary columns are tough to make and are expensive, although the
increase in efficiency is worth the price, particularly for the analysis of very complex samples
(e.g. gasoline).
Both types of columns are available with any one of several hundred different liquid
stationary phases. Selection of the type of liquid stationary phase is based on the type of sample
to be analyzed. The real power and flexibility of GC as a method of analysis rests on the fact that
the stationary phase can almost be tailored at will to fit the separation problem. The choice is
often made on the “like dissolves like” principle, or put in a more sophisticated way, the liquid is
chosen on the basis of polarity index. The column is usually placed in an oven, the temperature
of which can be raised or lowered (and monitored) in any predetermined manner. The separated
components that leave the column are then quantitatively detected by a suitable detector or, in
some instances, may be trapped and recovered.
53
Many commercial GCs have three types of built-in detectors: thermal conductivity
(TCD), flame ionization (FID), and electron capture (ECD). The detector output (signal) is
usually fed to a strip chart recorder or to a dedicated computer. A typical GC chromatogram is
shown below:
A GC Chromatogram
The various components do not have Rf values, in the same sense as in paper
chromatography, because the components actually come out of the gas chromatograph, and the
mobile phase is continuously flowing. In GC the retention parameter is called the retention time
(tR) and is the time that elapses between the injection of the sample and when the center of the
component band is detected by the detector. Almost always, the injection of the sample into a gas
chromatograph results in air being injected. Air components (O2 and N2) are generally unretained
– i.e., have no affinity for the liquid stationary phase – and quickly appear in the detector. The
time between sample injection and detected air peak is called the retention time of the air (even
though it is not retained by the stationary phase). The detectors also produce a concentration
profile that, with a suitable calibration line, can be used to quantitatively measure the amount or
concentration of any sample component.
54
Experimental Procedure
1.) Take one final sample from each of your 2,4-D remediation cells. Perform the same
liquid/liquid extraction method you did on the other samples (steps 2-8 of Experiment #7).
Be sure to label each vial!
2.) Pick up two standard vials, one with just 2,4-D and one with 2,4-D and its degradation
products. Gather your samples and follow the TA to the GC analysis room.
3.) Go through the entire CALIOPE program before touching anything on the instrument! The
GC CALIOPE program explains the instrument and how to use it fairly completely. You will
have to study this program before beginning the analysis of your samples.
4.) The first sample you should run is the 2,4-D standard, followed by the 2,4-D + degradation
products standard. Follow the CALIOPE instructions to set the rate and initial and final
temperatures for the GC run. These conditions need to remain constant throughout the
analysis.

Why is it important to maintain uniform run conditions?
5.) Run each of your samples (it is probably easiest to run them in the order that you took them
for each type of Fe particle.) Label each chromatogram with the sample identification.

What is the large peak at the beginning of each spectrum?
Data Analysis
At the bottom of each chromatogram is a list of retention times with their corresponding
peak areas. Ignoring the solvent peak, add the peak areas of each of the sample peaks together.
The fraction of each compound compared to the total sample can then be determined by dividing
that peak area by the total peak area. For example, for the 2,4-D standard, you should have only
one major peak. Thus, the total peak area = 2,4-D peak area, and the 2,4-D:total sample ratio is
1:1. If you had 2 peaks, component 1 with an area of 5,000 units and component 2 with an area
of 15,000 for a total of 20,000 units, component 1 would represent 5,000/20,000 or 0.25 of the
sample, while component 2 would represent 15,000/20,000 or 75% of the total area. For this
experiment, these ratios will give the fraction of undegraded 2,4-D because the only source of
organic compounds in your sample solution should be 2,4-D.
55
Mini-lab Report
For next week’s class, you will need to prepare a miniature lab report. In this report, you must
provide:
I.
An introduction explaining the history and current uses of iron PRBs (at least one page
double spaced).
II.
A brief explanation of the experimental procedure for the 2,4-D degradation (this can be
done in a paragraph).
III.
A results section where you present graphs of the fraction of 2,4-D versus time for both
the plain iron filings and the nickel-catalyzed iron filings. You should write 1-2
sentences explaining each graph.
IV.
A discussion where you explain your results. Did one kind of filing degrade the 2,4-D
faster than the other? Why? And what chemical processes control the degradation
processes? You can and should refer to literature sources for this information (about one
page double spaced should cover this).
56
Experiment #10:
Mr. Fish
exclaimed, “You
want that to pass
through my
system?!”
Wastewater
Treatment
57
ICK! What is it? And get rid of it for me, will you?
Cleaning Contaminated Waste
Lab Learning Goals:
•
To identify some common sources of pollution, as well as some specific compounds
that may come from each source.
•
To use several methods of analysis to identify the unknown contaminants in their
water samples.
•
To significantly reduce the contamination of their water sample.
•
To remember everything from the previous experiments to help prepare for the exam.
Background information
Learning to treat wastewater prior to its reintroduction into the main water system
has been one of the greatest advances for human society. Treatment facilities today focus
mainly on the removal of biological and organic wastes, although nitrates and phosphates
are also removed. The methods used by the University Area Joint Authority (UAJA), the
largest local wastewater treatment plant, are typical for most treatment plants. A concise
explanation of these methods can be found on UAJA’s website (www.uaja.com), and you
will need to read this to write the introduction to your lab report.
As you may know, the population of the State College area has been increasingly
fairly rapidly over the past several years. Because of this, UAJA has to expand its
facilities. While it may seem logical to just expand the current facilities to accommodate
the increase in wastewater, the EPA has done testing on Spring Creek and determined
that any amount of effluent (treated water) in excess of 6 million gallons per day (MGD)
poses a threat to the wildlife of the stream due to increased temperatures (the effluent is
several degrees warmer than the stream temperature). Therefore, UAJA has developed a
treatment plan known as the Beneficial Reuse Project. Part of the problem with reusing
water is that it essentially must be cleaner than the water we drink. Therefore, even
58
minor contaminants that are not considered problematic for release into the environment
must be removed from the water supply.
This experiment is designed to get you to think about how you would remove
various contaminants from the water supply – after you find out what exactly is
contaminating the water. Agricultural run-off leads to high levels of nitrates and
phosphates, while a factory such as the nearby Cerro Metals might release heavy metals
into the water supply. Organic contaminants include not only various pesticides and
herbicides, but also residues from gasoline, roadwork, and other human activities. While
not all contaminants are highly toxic, we do not yet understand the effects on either the
ecosystem or the human body of long-term exposure to many of these compounds.
Chemistry Background Information
Most of the chemistry that you need to know for this experiment has been covered
in other experiments. One remediation technique that we have not covered in this lab,
however, is activated carbon. Activated carbon is basically charcoal that has been
compressed so that the pore size in the material is so small that only water molecules (and
others smaller than water) can get through. The surface area of activated carbon is also
huge – up to 14,000 square meters per gram! Typically, activated carbon is used to
remove organic molecules and other large particles from a water supply, but it can be
expensive to use since the activated sites may fill up fairly quickly.
Flow of the Experiment
This experiment will consist of four main parts:
1) Creation of a contaminated water sample.
2) Identification of contaminants
3) Remediation of water sample
4) A paper describing the identification and purification processes, as well as the success
of the purification. In the conclusions section, you will explore the environmental
ramifications of your purification process.
59
Schedule for the experiment
1) Creation of a contaminated water sample: to be synthesized during the Fe0
remediation lab.
2) Identification of contaminants: Second to last week of lab before the exam.
3) Remediation of water sample: Last week of lab before the exam.
4) Paper: Due the day of the final exam, or as early as the students are willing to turn it
in. Most of this can be written ahead of time.
During the Fe0 Decontamination Lab
Preparation of the Contaminated Water Sample
During the Fe/Ni halocarbon removal lab (while waiting to take samples) you will work
in pairs to prepare your contaminated water samples. A list the contaminants should be
turned in to the TA at the beginning of class to make sure that the sample will meet the
following requirements:
1) There must be at least two steps necessary to remove the contaminants.
2) Only one of the contaminants should be removable by simple filtration (ie dirt, clay,
sludge).
3) You may only use contaminants on which we have conducted analyses or which can
be identified by the same analytical techniques we have used.
You should add a sufficient quantity of the contaminant such that the other group can
successfully identify the contaminants and can remove some significant quantity of them.
For example, if you add enough contaminant to make a 1 x 10-8 M solution of the
contaminant, it will be extremely difficult to test for and remove the contaminant.
However, if you add 1 x 10-5 M, it will be easier to both identify the contaminant and to
remove enough to notice a decrease in the concentration.
60
Week One
Hazards
The hydroxylamine solution used to reduce the iron is concentrated and can cause severe
irritation and burns upon skin contact or inhalation. Prolonged exposure to ferrozine may
cause dermatitis. Students should make sure that their cuvettes are well sealed before mixing
the iron/ferrozine/hydroxylamine/buffer solution, and should wear gloves and protective
eyeware. The unknown samples containing iron and those with a low pH are highly acidic
and may cause burns.
Identification of Contaminants
Over the duration of this course you have used several instrumental and observational
techniques to analyze various samples. As part of your final report, you should turn in a
list of the techniques you used, how they are used for analysis, and what you can detect
with each. ALL INSTRUMENTATION AND ANALYTICAL METHODS WILL BE
COVERED ON THE EXAM, SO IT IS A GOOD IDEA TO CREATE A CHART FOR
ALL INSTRUMENTS AND JUST STICK WHAT YOU NEED IN THE LAB
REPORT!
(just a hint)
To help you out, a list of the instruments and methods that have been used in this course
is given below:
titration
Spec 20
AA (note: you can use the 4 lamp AA)
IC
GC (after liq/liq extraction)
precipitation and filtration
visualization
Q: Can any of these techniques simultaneously identify more than one contaminant?
You and your partner may wish to divide up by instruments, and possible contaminants
(anions vs. cations, etc). Be sure to filter your samples before running AA, IC, etc.
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You should not only identify your contaminants but also estimate a concentration for
them (within an order of magnitude). Once you have figured out your contaminants,
check with the TA to see if they are correct. If so, you will then need to propose a
method for removing each contaminant (it is feasible to have one method remove two
substances, but no, you can’t do reverse osmosis and get rid of everything). You will
carry out these removal methods for each contaminant, and you will need to discuss the
chemistry behind each in your paper. It is possible to get most of your paper done this
week (which is why it’s due the same day as the final).
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Week 2
Carry out your proposed methods for removing your contaminants. Each partner should
attempt a clean up of the water – just divide your remaining sample in two. You don’t
necessarily need to clean up 150 mL of water, but be sure to do enough that you can
properly analyze it when you’re done. Also, you must decide whether to analyze the
sample after each remediation step, or to wait until the end.
You must do ballpark quantitative analyses. You should have approximate initial
concentrations from last week. Note: be sure to take into consideration the volume of
the sample you used if you are doing an analysis that is sensitive to the amount of a
substance being analyzed.
Things to consider:
•
If you are precipitating out a contaminant, you can prove your success in getting rid
of it by massing the precipitate.
•
There may be some contaminants that are quite tricky to remove chemically – if you
find yourself unsuccessful, you must explain why you were unsuccessful, and find out
how municipalities remove that particular contaminant in large-scale wastewater
treatment plants.
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The Final Report
You must write a final report for this paper that includes information arranged according
to the directions below:
I.
An introduction that includes:
A. Main sources of your chemical contaminants
B. A description of how UAJA treats its water
C. What remediation techniques you chose to remove your contaminants.
II.
A materials and methods section that gives:
A. A list of chemicals you used to identify/purify your sample
B. A list of instruments used in the analysis of your sample (this is usually
detailed for scientific papers – you must give the name, manufacturer, and
manufacturer’s location)
C. Any other random materials used: filter paper, special glassware, etc.
D. A brief (no more than three quick paragraphs) account of how you determined
what was in your sample, how you removed it, and how you proved you
removed your contaminants.
III.
A results section that includes:
A. Identification of your contaminants
B. Approximate initial concentrations of contaminants
C. Approximate final concentrations of contaminants
Note: If you could not do any quantitative analysis, state how you proved that
your contaminants were removed (or how you knew that you were
unsuccessful).
IV.
A discussion section where you should:
A. Explain why your remediation plans were or were not successful. Here, you
should discuss the chemistry behind each technique you used.
B. Discuss how your analytical methods proved the presence (or absence) of the
particular compounds in your sample (i.e. how did you know it was NO3- and
not Cl- or F- or SO42-?)
C. In a single paragraph, offer conclusions as to how successful you were and
how you could have improved your results.
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Post lab questions (to be answered at the end of the lab report)
1.) Imagine your water sample is now 2,000,000 gallons in size and that you have some
cost limitations. How would this change your choice of remediation techniques?
You may want to consider cost, chemical vs. biological remediation techniques, and
to what degree you would remove the contaminating species.
2.) A focus of green chemistry today is the total reduction of chemicals going into and
out of the environment. What chemical species did you add to your water sample to
remediate it? Just for something to think about, consider how much of each chemical
you would need if you were doing large-scale wastewater treatment (UAJA can treat
about 6,000,000 gallons per day). Where would you store the chemicals? This is part
of why wastewater treatment plants use biological and radiation techniques instead of
chemical treatment methods.
3.) Draw a possible map of the area from where your water sample came. It should be
feasible – no huge industrial parks next to farms – and should show:
1.) sources of pollution
2.) how the chemicals are transported to a common body of water
3.) where you would live given the sources of water and sources of pollution.
4.) Should public funds and treatment facilities be used to treat industrial wastewater or
should those companies have to treat their own water?
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Appendix A. Spring Creek Watershed
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Bibliography
Baird, C. Environmental Chemistry. 2nd ed. W.H. Freeman and Company: NY. 1999.
Balko, B.; Tratnyek, P.G. J. Chem. Educ., 2001, 78, 1661.
Brown, T.L.; LeMay, H.E. Jr.; Bursten, B.E. Chemistry: The Central Science. 7th ed.
Prentice Hall: Upper Saddle River, NJ. 1997.
Carson, R. Silent Spring. Houghton Mifflin Company: NY. 1962.
“Creek Connections.” Online. Available: http://creekconnections.alleg.edu/ Accessed
23 April 2002.
“Green Chemistry at the University of Oregon.” Online. Available:
http://www.uoregon.edu/~hutchlab/greenchem/ Accessed 22 April 2002.
J. Chem. Educ. 2000, 77 (12).
J. Chem. Educ. 2001, 78 (12).
Kegley, S.E.; Landrear, D.; Jenkins, D.; Gross, B.; Shomglin, K. Water Treatment: How
Can We Purify Our Water? Student Manual. John Wiley & Sons: New York. 2000.
Minard, R.D. Lab Guide for Chemistry 431-W: Organic and Inorganic Preparations
and Qualitative Organic Analysis. Penn State University, 2001.
Nyer, E.K.; Vance, D.B. Ground Water Monitoring and Remediation. Spring 2001. 4146.
Schrick, B.; Blough, J.L.; Jones, A.D.; Mallouk, T.E. “Hydrodechlorination of
Trichloroethylene to Hydrocarbons Using Bimetallic Nickel-Iron Nanoparticles.” Chem.
Mater. 2002, 14, 5140-5147.
Schrick, B.; Hydutsky, B. W.; Blough, J.L.; Mallouk, T.E. “Delivery Vehicles for Zerovalent Metal Nanoparticles in Soil and Groundwater.” Received from the author, Fall
2001.
Sinniah, K.; Piers, K. J.Chem.Educ. 2001, 78, 358.
Spring Creek Watershed Community Water Resources Monitoring Project. 2000 Annual
Report.
Tabbutt, F.D. J.Chem.Educ. 2000, 77 (12). 1594.
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Thompson, Stephen. Penn State Version of Chemtrek: Small-scale experiments for
General Chemistry. Prentice Hall: Englewood Cliffs, NJ. 2000.
Thompson, Stephen. Chemtrek. Prentice Hall: Englewood Cliffs, NJ, 1989. Adapted
for Chem 15 by J.T. Keiser, 17 April 1997.
University Area Joint Authority. Online. Available: http://www.uaja.com/ Accessed 22
April 2002.
Yarnal, Brent. Personal Communication. 4 April 2002.
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Acknowledgments
The help I have had with this project has been incredible, and I would like to thank the
following people:
Dr. Joseph Keiser, for his trust, advice, support, and insight.
Dr. Jackie Bortiatynski, Andrew Greenberg, Dr. Tom Mallouk, Dr. J.P. Lowe, Dr. A.
Daniel Jones, Sue Swope, Bettina Schrick, Dr. Juliette Lecomte (for attempting to keep
me organized), Katie Ombalski and Dr. Robert Carline of the Spring Creek Watershed
Monitoring Project, Diane Jones, Art Brandt of the University Area Joint Authority, all of
my friends and peers who have supported this effort and given me advice on how they
would have designed the course, particularly Allison Carey, Emily Moriarty, Simon
Lobdell, Soma Kedia, Christina Chong, Lindsey Thorne, Jonathan Dick, and Jason
Thorhauer. Finally, thanks to my parents, Michael and Judy Mihok, brother Zack Mihok,
and family, for the way they teach me to love and rejoice with everyone and everything
around me.
And also, thanks to all of the life around me that I hope to connect in some way through
this project.
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