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

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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. 4Thompson, Stephen. Chemtrek. Prentice Hall: Englewood Cliffs, NJ, 1989. Adapted for Chem 15 by J.T. Keiser, 17 April 1997. 5Balko, B.A.; Tratnyek, P.G. J.Chem. Educ., 2001, 78, 1661. 6Original

-- standard laboratory technique, adapted for Fe remediation laboratory.

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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.

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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.

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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.

6

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 40% Quizzes 15% Final Exam 20% Instructor Evaluation 10% Wastewater Treatment Paper 15%

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Experiment #2:Experiment #2: Determination ofDetermination of

[Fe[Fe 2+2+ ], A lkalinity and Metal Ion Content of Local ], A lkalinity and Metal Ion Content of Local Water SamplesWater 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

(H2O) 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

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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 (H2SO4) 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 (H2SO4). 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.

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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

(H2CO3), 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+ HCO3

-

HCO3- + 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 (H2SO4) 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 (H2SO4 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

10

do this for us. Carbonic acid (H2CO3) 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+ H2CO3).

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.

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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

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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

13

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 "non-

carbonate 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.

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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 (Fe2O3), 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

15

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, Fe2O3). 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.

16

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

17

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(H2O)6]2+ or

[Fe(H2O)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.

18

Ferrozine will not react with iron (III), so in the ferrozine solution, a mild reducing agent,

hydroxylamine (NH2OH), 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.

19

Colorimetric analysis is carried out using a spectrophotometer, an instrument that

measures the amount of light transmitted through a solution.

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:

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:

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.

!

T =It

I0

lamp

monochromator sample

detector

transmitted light, It

incident light, I0

20

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:

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.

A = -log T

A = εbc

21

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:

H3O+(aq) + CO32-(aq) HCO3

-(aq) + H2O (l)

Bicarbonate:

H3O+(aq) + HCO3-(aq) H2CO3 (aq) + H2O (l)

Hydroxide:

H3O+(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) H2CO3 (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

CaCO3/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

22

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) + H3O+ (aq)

The acid (H3O+) 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) H2O (l) + HCO3- (aq) + Ca2+ (aq)

23

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 Beer-Lambert law? Define your labels for the graph.

Determination of Unknown Concentration by

Spectrophotometry

0

0.1

0.2

0.3

0.4

0.5

0.6

0 2 4 6 8 10

_________________

_________________

24

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, CO3

2-, 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?

25

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. Calibration of the Spectrophotometer Part B. Determination of Iron in an Unknown Sample Part C. Alkalinity Determination by Titration of an Unknown Water Sample Part D. 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. 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?

Section A. Calibration of the Spectrophotometer

26

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.

27

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. 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

Section B. Determination of Iron in an Unknown Sample

Section C. Alkalinity Determination by Titration of an Unknown Water Sample

28

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.

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

Section D: Atomic Absorption Determination of Ca, Mg, Na, and Fe

29

instrument sensitivity, ease of sample preparation, time/effort required, repeatability, sources of error).

3) 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.

4) 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 write-up.

5) If the only source of alkalinity in a sample is sodium hydroxide (NaOH), how would the calculation for total alkalinity be different?

6) 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.

7) 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).

8) Which method of analysis (AA or Spec 20) do you think provides more accurate values? Give at least one reason why.

30

Experiment #8:Experiment #8: Halocarbon Halocarbon

Remediation with Remediation with ReReduced Ironduced Iron Kinetics and Kinetics and

Liquid/Liquid Liquid/Liquid ExtractionExtraction

31

Sources for this exper iment are referenced at the end of the exper iment.Sources for this exper iment are referenced at the end of the exper iment.

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

32

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 y-

intercept, 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.

33

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

34

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

35

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

36

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.

37

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.

38

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

39

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 (1)

CCl4 + 2e- + H+ CHCl3 + Cl- E0 = 0.67 V (2)

_________________________________

CCl4 + H+ + Fe0 CHCl3 + Cl- + Fe2+ E0 = 1.11 V (3)

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 + 2H2O 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)

40

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)

41

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:

42

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.

43

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 nickel-coated 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

44

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.

45

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% H2O2 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?

46

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.

47

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

Fe2O3 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.

48

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.

49

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 CH2Cl2. 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 CH2Cl2 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.

50

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.

51

Experiment #9:Experiment #9: Halocarbon Halocarbon

Remediation with Remediation with Reduced IronReduced Iron

Sample AnalysisSample Analysis

52

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

1Thompson, Stephen. Penn State Version of Chemtrek. Prentice Hall: Englewood Cliffs, NJ.

2000.

53

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.

54

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.

55

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.

56

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).

57

Experiment #10:Experiment #10:

Mr. Fish Mr. Fish exclaimed, “You exclaimed, “You want that to pass want that to pass

through my through my system?!” system?!”

Wastewater Wastewater TreatmentTreatment

58

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

59

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.

60

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.

61

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.

62

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).

63

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.

64

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.

65

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. 41-46. 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 Zero-valent 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.