Download It`s Easy Being a Green Chemist

It’s Easy Being a Green Chemist
High School Science Standards-Based
Green Chemistry Curriculum
Conceived, adapted, and compiled by San Diego Region Teachers
Funded with SB 70 from the California Community Colleges
Chancellor’s Office
Version 1.0 May 2012
Acknowledgements
On behalf of the Southern California Biotechnology Center @ Miramar College, we would like to acknowledge
the teachers who made this manual possible.
In 2008, state officials in California launched a "green initiative" with the goal of knowing the environmental
footprint of every consumer product. At about the same time, the Warner-Babcock Foundation group, Beyond
Benign, created and disseminated the first curriculum for "Green Chemistry". Meanwhile, the pharmaceutical
industry giant, Pfizer began efforts to "green" their processes.
So, the time was right to start to think about creating standards-based curriculum that both "greened" the
teaching lab and introduced the concept that sustainable/green practices were a necessary part of the creation
process to students.
Funding from CA SB 70 Community Colleges Chancellor’s office Career Technical Education made this effort
possible. The process of developing this curriculum was a lesson in collaboration. It began with two teachers,
Dr. Linda Woods from San Diego Miramar College and Mr. Russ Davidson from San Dieguito High School,
attending a Beyond Benign Green Chemistry Symposium. They brought back to our education community a few
labs that they thought should be incorporated into their curriculum. These labs were disseminated in a teacher
workshop. The attendees of this workshop formed the core of the group that created standards based labs that
refined their current labs with both green practices and green concepts, hence "greening" the lab activity and
students simultaneously. These labs were "field tested" at a workshop for teachers. The final step of this
process was to develop a lab manual to broadly disseminate these "best practices".
Lastly, we want to acknowledge the drivers of this product. Ms. Ruth Trejo, while an adjunct faculty member at
San Diego Miramar College, coordinated the creation of this curriculum. Many thanks also to those whose labs
you will find in the manual. Finally, we want to thank Ms. Laurie Batinica Aud for editing and compiling the
curriculum into this “It’s Easy Being Green” manual.
Should you have any questions or comments, please feel free to contact us.
Sandra Slivka Ph.D.
Director, Southern California Biotechnology Center @ Miramar College
10440 Black Mountain Rd
San Diego CA 92129
[email protected] 619-388-7490 www.matsc.org
www.scbcmiramar.org
Table of Contents
Introduction
What is Green Chemistry? ..........................................................................................................................................5
The 12 Principles of Green Chemistry ........................................................................................................................7
Green Chemistry Labs
Limiting Reagent Stoichiometry: What Stops It? .......................................................................................................8
A Series of Copper Reactions and Percent Yield ..................................................................................................... 21
Chemical Reactions ................................................................................................................................................. 25
Eco-Friendly Equilibrium (Teacher/TA Edition) ....................................................................................................... 38
Eco-Friendly Equilibrium (Student Version) ............................................................................................................ 44
The Colors of Le Chatelier........................................................................................................................................ 55
Gas Laws .................................................................................................................................................................. 62
Reaction of Magnesium and Water in the Presence of a Catalyst (Flameless Ration Heater) .............................. 73
Column and Thin-Layer Chromatography ............................................................................................................... 76
Fruit Basket of Esters - Synthesis of Esters .............................................................................................................. 83
Analysis of Fats in the Diet and Isolation of Fat and Olestra in Food Items ............................................................ 90
Classification and Recycling of Plastics .................................................................................................................. 102
Additional Resources
Green Chemistry Resources .................................................................................................................................. 114
What is Green Chemistry?
Green Chemistry is the philosophy originally defined in 1998 by Paul Anastas and John Warner in their book
Green Chemistry: Theory and Practice. Green Chemistry is a revolutionary, yet common sense approach to the
way products are made. It is a science that aims to reduce or eliminate the use and/or generation of hazardous
substances in the design phase of materials development. Considered a subset of sustainable chemistry, Green
Chemistry is defined as:
The utilization of a set of principles that reduces or eliminates the use or generation of hazardous substances
in the design, manufacture, and application of chemical products.
One major goal of Green Chemistry as described by John Warner is to "change the way we make chemists." The
traditional process of designing a chemical product typically does not consider the lifecycle of the product or the
long-term downstream ramifications of the waste it generates. Too often it is discovered that the chemicals
used in everyday products, or the waste generated during the manufacture of these chemicals, cause cancer,
birth defects, or a host of other negative consequences for organisms and the environment.
It is important to make the distinction that what is perceived as “green” may not align with the principles of
Green Chemistry. Many efforts in recent years to "go green" have been intended to save energy and to make
more environmentally friendly products. However, the use and/or generation of toxic chemicals during the
manufacturing of a “green” product may outweigh the benefits to the environment. For example, solar paneling
(photovoltaic) converts solar energy into electricity and produces no green house gas emissions while providing
a sustainable alternative to fossil fuels. On the downside, solar panels are manufactured with silicon-based
products and contain heavy metals, such as lead (neurotoxin) and cadmium (carcinogen), just like computers,
cell phones, and most other electronic devices. The production lifecycle of photovoltaics (extracting, refining,
purifying of raw materials, production, and disposing or recycling) involves environmentally unfriendly and
unsustainable practices. Since photovoltaics have a useful life of no more than 20 years, they end up as a
significant source of hazardous waste in landfills.
Another similar example involves the fluorescent bulbs that are quickly replacing incandescent bulbs as an
energy-saving lighting alternative. Fluorescent light bulbs contain a seemingly insignificant amount of mercury
(3.5-15 mg). The U.S. Environmental Protection Agency estimates 670 million fluorescent light bulbs are thrown
away every year, releasing 2 to 4 tons of mercury into the environment.
Green Chemistry is unique in that it directs the chemists inventing molecules to think about the entire process of
creating a chemical from synthesis to application, to waste and disposal before making a molecule. For a
technology to be considered Green Chemistry, it must accomplish three things:
•
It must be more environmentally benign than existing alternatives.
•
It must be more economically viable than existing alternatives.
•
It must be functionally equivalent
Green Chemistry technology is currently developing replacements for many existing products, including
alternative solar paneling. Refer to the Resources page at the end of the manual for links and articles providing
more information.
To make Green Chemistry a commonplace practice and to replace traditional processes that exist today,
chemists need "a new set of tools (molecules) in their toolbox" as described by John Warner, steering away from
the more hazardous set of starting materials used for chemical syntheses. Educators play a central role in
introducing the practice of Green Chemistry. John Warner recognized the importance of educating K-12 and
college students and teachers and founded the non-profit organization, Beyond Benign. This organization is the
cornerstone of Green Chemistry curriculum and the most comprehensive resource for teachers. Follow their link
for curriculum mapping: http://www.beyondbenign.org/K12education/highschool.html. Currently in the U.S.,
Green Chemistry coursework or a demonstration of knowledge regarding toxicity or environmental impact is not
5
required in chemistry education. The only full Green Chemistry program currently available in the U.S. is
through the Warner Babcock Institute. The goal of the institute is to have 500 universities and colleges offering
programs in Green Chemistry within five years (by the year 2018). As a global perspective, India and China are
two countries with a large manufacturing base whose governments are putting Green Chemistry into practice.
India currently requires one year of Green Chemistry coursework for chemistry majors, and China has 15
national research labs dedicated to the field.
The demand for qualified green chemists in the chemical industry and in product manufacturing is very high and
unmet, giving these candidates a distinct competitive edge when seeking employment. The reason for this
demand is that manufacturers have been known to pay more for regulated hazardous waste and disposal of
chemicals produced as a side product during manufacturing than for producing the chemical or product itself.
Needless to say, the chemical industry is one of the greatest proponents of Green Chemistry, in order to
improve sustainability of their processes and reduce their costs.
There are 12 guiding principles that make up the backbone of Green Chemistry. The Twelve Principles of Green
Chemistry, listed on the next page, provide a framework for scientists and engineers to use when designing new
materials, products, processes, and systems. The Green Chemistry labs presented in this manual apply these
principles and are meant to replace less eco-friendly labs that are traditionally used to teach current high school
chemistry content standards for grades 9-12. Each lab lists the pertinent high school chemistry standards
covered and the appropriate green principle(s). Note: Inherently, the labs fulfill part or all of the Investigation
and Experimentation section of Grades 9-12 standards. The following is the link to the California Science
Content Standards: http://www.cde.ca.gov/be/st/ss/documents/sciencestnd.pdf
Through Green Chemistry, environmentally benign alternatives to current materials and technologies can be
systematically introduced across all types of manufacturing to promote a more environmentally and
economically sustainable future. Teachers play the most important role in introducing this technology and
creating change.
For more Green Chemistry information and curriculum resources, refer to the Resources page and articles at
the end of this manual.
6
The Twelve Principles of Green Chemistry
1) Prevention
It is better to prevent waste than to treat or clean up waste after it has been created.
2) Atom Economy
Synthetic methods should be designed to maximize the incorporation of all materials used in the process into
the final product.
3) Less Hazardous Chemical Syntheses
Wherever practicable, synthetic methods should be designed to use and generate substances that possess little
or no toxicity to human health and the environment.
4) Designing Safer Chemicals
Chemical products should be designed to affect their desired function while minimizing their toxicity.
5) Safer Solvents and Auxiliaries
The use of auxiliary substances (e.g., solvents, separation agents, etc.) should be made unnecessary wherever
possible and innocuous when used.
6) Design for Energy Efficiency
Energy requirements of chemical processes should be recognized for their environmental and economic impacts
and should be minimized. If possible, synthetic methods should be conducted at ambient temperature and
pressure.
7) Use of Renewable Feedstocks
A raw material or feedstock should be renewable rather than depleting whenever technically and economically
practicable.
8) Reduce Derivatives
Unnecessary derivatization (use of blocking groups, protection/ deprotection, temporary modification of
physical/chemical processes) should be minimized or avoided if possible, because such steps require additional
reagents and can generate waste.
9) Catalysis
Catalytic reagents (as selective as possible) are superior to stoichiometric reagents.
10) Design for Degradation
Chemical products should be designed so that at the end of their function they break down into innocuous
degradation products and do not persist in the environment.
11)Real-time analysis for Pollution Prevention
Analytical methodologies need to be further developed to allow for real-time, in-process monitoring and control
prior to the formation of hazardous substances.
12) Inherently Safer Chemistry for Accident Prevention
Substances and the form of a substance used in a chemical process should be chosen to minimize the potential
for chemical accidents, including releases, explosions, and fires.
7
Limiting Reagent Stoichiometry: What Stops It?
Developed/Adapted by Diana Vance, Grossmont College and Ruth Trejo, Miramar College
This protocol was developed with California SB 70 funding in an effort to bring sustainability practices to
teaching laboratories and to increase awareness of sustainability in the next generation of the California
workforce. In other words, these protocols “green” our labs and “green” our students.
California Science Standards, Chemistry Grades 9-12
3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the
ability to calculate the mass of products and reactants. As a basis for understanding this concept:
a. Students know how to describe chemical reactions by writing balanced equations.
b. Students know the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass
of exactly 12 grams.
d. Students know how to determine the molar mass of a molecule from its chemical formula and a table
of atomic masses and how to convert the mass of a molecular substance to moles, number of
particles, or volume of gas at standard temperature and pressure.
e. Students know how to calculate the masses of reactants and products in a chemical reaction from the
mass of one of the reactants or products and the relevant atomic masses.
f. * Students know how to calculate percent yield in a chemical reaction.
Green Chemistry
A widely used limiting reagent lab involves aspirin synthesis. The reagents are harmful to humans and the
environment. Salicylic acid is a toxic and irritant, acetic acid a corrosive lachrymator, and concentrated sulfuric
acid is a toxic, corrosive oxidant that causes severe burns. The aspirin product must be disposed of in a proper
waste container, and the liquid filtrate must be neutralized with a strong base.
In this "green" lab, calcium carbonate will be synthesized. The calcium carbonate product is essentially chalk
and can be disposed of in the regular trash can. The other product is aqueous table salt and can go down the
sink.
Green Chemistry Principles
Principle #3 - Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methodologies should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #4 - Designing Safer Chemicals: Chemical products should be designed to preserve efficacy of
function while reducing toxicity.
Purpose
The purpose of this experiment is to explore the concept of limiting reactant in a chemical equation.
Theory
When two or more reactants react chemically, they do so in a certain mole ratio (the stoichiometric ratio).
This can be seen from the equation:
2 NaOH (aq) + H2SO4 (aq)
Na2SO4 (aq) + 2 H2O (l)
rxn 1
where two moles of sodium hydroxide react with one mole of sulfuric acid forming one mole of sodium sulfate
and two moles of water.
8
If one of the reactants is present in an amount less than the stoichiometric proportion to that of the other (or
those of the others), it will run out and limit the theoretical product yield. Thus, if two moles of sodium
hydroxide and one-half mole of sulfuric acid are allowed to react, only one mole of the sodium hydroxide will be
consumed with the half mole of acid, and only one-half mole of sodium sulfate and one mole of water will be
produced. In this example, the sodium hydroxide is said to be present in excess, and sulfuric acid is called the
limiting reactant (or limiting reagent).
Example:
2 NaOH (aq) + H2SO4 (aq)
Na2SO4 (aq) + 2 H2O (l)
What is the amount of water produced when 2.00g NaOH reacts with 2.00g H2SO4?
2.00g NaOH x (1 mol NaOH ) x (2 mol H2O) = 0.05001mol H2O
(39.99g NaOH) (2 mol NaOH)
2.00g H2SO4 x (1 mol H2SO4 ) x (2 mol H2O) = 0.04078mol H2O
(98.086g H2SO4) (1 mol H2SO4)
Less Mols of Product
= LR!!
0.0478mol H2O x (18.016g H2O) = 0.735g H2O theoretically produced
(1mol H2O)
The theoretical yield is the maximum theoretical amount of a given product that can be formed when the
limiting reactant is completely consumed. The actual yield is the amount of product actually isolated from a
reaction. Percent Yield is a comparison of the actual yield to the theoretical yield. See Equation below.
Percent Yield = Actual Yield x 100
Theoretical Yield
A Percent Yield can be used to measure productivity of the reaction with specific experimental conditions
and/or the efficiency of laboratory technique.
Read the experiment and review the material on limiting reagent stoichiometry in your lecture textbook. When
completing the prelaboratory exercise, the following explanations should be helpful.
Sample Calculation: 25.00 grams of calcium chloride dihydrate reacts with 25.00 grams of sodium carbonate to
form calcium carbonate, sodium chloride and water. What is the mass of calcium carbonate formed? What is
the percent yield of the reaction if 15.50g CaCO3 was produced in the laboratory?
a. Balance the following reaction. * This is not balanced….Please balance it!
____ CaCl2 · 2H2O + ____ Na2CO3
____ CaCO3 + ____ NaCl + _____ H2O
b. First convert grams of both reagents into moles of one of the products.
25.00g CaCl2 · 2H2O x (1 mol CaCl2 · 2H2O) x
(147.012g CaCl2 · 2H2O)
25.00g Na2CO3 x
(1mol CaCO3) = 0.17005mol CaCO3
(1mol CaCl2 · 2H2O)
(1mol Na2CO3) x (1mol CaCO3) = 0.23587mol CaCO3
(105.99g Na2CO3) (1mol Na2CO3)
c. Determine the Limiting Reactant (Produces less moles of the product).
CaCl2 · 2H2O produces less mols of CaCO3 product. Therefore it is the LR!
Limiting Reagent Stoichiometry
9
d. Using the lesser amount of CaCO3 mols, calculate the mass of CaCO3 theoretically produced.
0.17005mol CaCO3 x (100.09g CaCO3) = 17.02g CaCO3 Theoretically Produced
(1mol CaCO3)
e. Calculate the percent yield of the reaction.
Percent Yield = Actual yield x 100 = 15.50g CaCO3 x 100 = 91.07% yield
Theoretical yield
17.02g CaCO3
Method
In this experiment you will work with the chemical reaction between calcium chloride dihydrate and anhydrous
sodium carbonate, which produces insoluble calcium carbonate, soluble sodium chloride, and water. From the
balanced reaction equation you will determine the correct stoichiometric ratio of reactants. You will then
determine the actual yield of crystallized calcium carbonate. Lastly, you will compare the actual and theoretical
yields, in order to determine the percent yield.
Materials
Equipment (per group):
• 2 125 mL Erlenmeyer flasks
• 1 Rubber policeman
• 2 Stirring rod
• 2 Paper filter paper
• 2 Weigh Boats
• 1 Büchner funnel
• 1 250mL filter flask
• 1 250mL Erlenmeyer flask
• 1 Y-tube
Equipment (per group):
• 1 Short tubing (8 cm)
• 1 Pinch clamp
• 1 Ring stand
• 2 Ring stand clamps
• 2 2” rubber hose
Solid Reagents:
• Na2CO3 sodium carbonate anhydrous
• CaCl2 • 2H2O calcium chloride dihydrate
Procedure
1. Obtain an Unknown Letter from your instructor. This letter
Vacuum Filtration
corresponds to a certain amount of each of the two
Setup:
reactants. Record the unknown letter in your Data Sheet.
2. For calcium chloride dihydrate compound: Tare a weigh
boat on the analytical scale, and using a spatula, transfer
the correct amount of solid onto the weigh boat. Record
the exact mass on your Data Sheet. Transfer the calcium
chloride dihydrate in the 125mL Erlenmeyer flask, labeled
“calcium chloride dihydrate.” Make sure the entire
reagent is transferred into the flask. Describe the calcium
chloride dihydrate compound in your Data Sheet.
3. For sodium carbonate compound: Tare a new weigh boat
on the analytical scale, and using a spatula, transfer the
correct amount of solid onto the weigh boat. Record the
exact mass on your Data Sheet. Transfer the sodium
carbonate in the 125mL Erlenmeyer flask, labeled “sodium carbonate.” Make sure the entire reagent is
transferred into the flask. Describe the sodium carbonate compound in your Data Sheet.
4. To each 125 mL Erlenmeyer flask, add 50mL of DI water and dissolve both reagents.
Limiting Reagent Stoichiometry
10
5. With a stirring rod, gently pour and mix the calcium chloride dihydrate solution into the sodium carbonate
solution. Use your stir rod, remember “glass-to-glass-to-glass.”
6. Rinse the flask with a small amount of DI water to ensure a quantitative transfer of calcium chloride
dihydrate solution.
7. Gently stir the mixture, and then let it settle for about ten minutes.
8. While you are waiting, clean up glassware and set up a vacuum filtration with a water trap. See setup
above.
9. Weigh TWO pieces of filter paper and record the mass in your Data Sheet.
10. Place the filter papers into the Büchner funnel and dampen it with DI water to make it stick.
11. Slowly turn on the vacuum line (turn on the water).
12. Pour most of the liquid (supernatant) into the funnel, without transferring solid. Hold the funnel down to
ensure better suction. Allow for this liquid to drain into the filter flask.
13. Gently swirl the flask and then pour the rest of contents IN SMALL AMOUNTS into the funnel so that the
suspended precipitate goes onto the center part of the filter. If you can keep the precipitate back from the
edge, it will be much easier to remove the filter paper, dry it and weigh it without loss of any solid.
14. Remove the remaining precipitate from the flask by using three 10 mL portions of DI water and the rubber
policeman to scrub the solid calcium carbonate, CaCO3, from the sides of the beaker.
15. Hold the funnel down and allow the suction to dry the precipitate for a minute or so; then turn off the
vacuum. Describe the product.
16. While waiting for the precipitate to dry, weigh your evaporating dish on the analytical scale. Record this
mass on your Data Sheet. **After weighing the mass of the dish, you must not touch the evaporating dish
with your hands. Oils from your hands will affect the mass of your compound. Use Crucible Tongs!
17. Carefully remove the filter papers and precipitate, losing as little as possible, and place on your evaporating
dish. Using tongs, place the evaporating dish in your drawer to allow to dry until next period.
18. Clean up all glassware and dispose of the liquid filtrate in the sink with plenty of water.
19. Complete Calculations (#1-5) and Post-Lab questions during the lab period.
20. After one week, weigh the evaporating dish, filter paper and compound on the analytical scale. Remember
to use tongs! Calculate the mass of compound produced and complete percent yield calculation (#6)
Results (Calculations)
Determine the theoretical yield of product from the reaction equation and your actual initial mass of your
reactants. Compare the theoretical yield with your actual yield and calculate the percent yield of your reaction.
Limiting Reagent Stoichiometry
11
Limiting Reagent DATA SHEET
Name: ______________________
Assigned Letter
CaCl2 · 2H2O reagent:
Mass of CaCl2 · 2H2O used
=
g
=
g
CaCO3 Product:
Mass of two filter papers
=
g
Mass of evaporating dish
=
g
Mass of two filter papers, evaporating dish =
and precipitate after 1 week
g
Mass of Calcium Carbonate Product
g
Describe the compound:
Na2CO3 reagent:
Mass of Na2CO3 used
Describe the compound:
Describe the compound:
=
Describe the Final Product:
Limiting Reagent Stoichiometry
12
Limiting Reagent RESULTS (calculations)
Name: _______________________
1. Write the balanced equation for this reaction.
2. Calculate the moles of CaCO3 produced from the mass of CaCl2 · 2H2O used in the experiment.
3. Calculate the moles of CaCO3 produced from the mass of Na2CO3 used in the experiment.
4. Which reactant is the limiting reactant?
Which reactant is in excess?
_______________________
________________________
5. Calculate the theoretical yield of CaCO3 from the limiting reactant.
6. What is the percent yield of the reaction?
Limiting Reagent Stoichiometry
13
Limiting Reagent POST-LAB QUESTIONS
Name: _______________________
1. Describe the effect on the actual yield AND percent yield (Increase or Decrease) if the following
errors occurred. Explain why:
a) The filter papers were still wet when the final product was weighed:
b) Some product fell through the filter papers and was seen in the filter flask during the vacuum
filtration:
c) Some product fell on the floor and was scraped up with some dust particles. It was then
weighed for the final mass.
2. The following questions are based on the reaction below:
2 KCl (aq) + 2 MnCl2 (aq) + 5 PbO2 (aq) + 4 HCl (aq)
2 KMnO4 (aq) + 5 PbCl2 (s) + 2 H2O (l)
a) 25.00g of MnCl2 is reacted with 100.0g of PbO2 and excess KCl and HCl. How many grams of
KMnO4 can be produced by this reaction?
b) If 21.42 g of KMnO4 is actually produced, what is the percent yield? (If the grams of KMnO4
were not obtained in (a), use 35.03g KMnO4)
Limiting Reagent Stoichiometry
14
Name
Section
Date
Limiting Reagent
PRE-LABORATORY EXERCISE
Read the pre-laboratory notes, the experiment, and then complete the following as directed.
1. Define the following terms:
a. Limiting Reactant
b. Theoretical Yield
c. Percent Yield
2. a. Balance the equation and include the states of matter:
____ CaCl2 · 2H2O (aq) + ____ Na2CO3 (aq)
____ CaCO3 ( ____ ) + ____ NaCl ( ____ ) + _____ H2O (
)
b. The mol ratio between CaCl2 · 2H2O and CaCO3 is… _____mol CaCl2 · 2H2O = ____ mol CaCO3
c. How many mols of CaCO3 would be produced from 3.00 mols of CaCl2 · 2H2O? Show your work.
d. The mol ratio between Na2CO3 and CaCO3 is…
_____mol Na2CO3 = ____ mol CaCO3
e. How many mols of CaCO3 would be produced from 2.00mols of Na2CO3? Show your work.
f. What is the mass of CaCO3 produced from the limiting reactant?
Limiting Reagent Stoichiometry
15
Limiting Reagent DATA SHEET--KEY
Name: ______________________
Assigned Letter
CaCl2 · 2H2O reagent:
Mass of CaCl2 · 2H2O used
=
g
=
g
CaCO3 Product:
Mass of two filter papers
=
g
Mass of evaporating dish
=
g
Mass of two filter papers, evaporating dish =
and precipitate after 1 week
g
Describe the compound:
White Solid
Na2CO3 reagent:
Mass of Na2CO3 used
Describe the compound:
White Solid
Describe the compound:
White Solid
Mass of Calcium Carbonate Product
=________________________ g
Describe the Final Product:
Limiting Reagent Stoichiometry
16
Limiting Reagent RESULTS (calculations)--KEY
1. Write the balanced equation for this reaction.
CaCl2 · 2H2O
(aq)
+ Na2CO3 (aq)
CaCO3 ( s ) + 2 NaCl ( aq ) + 2 H2O
(l)
2. Calculate the moles of CaCO3 produced from the mass of CaCl2 · 2H2O used in the experiment.
______g CaCl2 · 2H2O x
(1 mol CaCl2 · 2H2O) x
(147.012g CaCl2 · 2H2O)
(1mol CaCO3) = _______mol CaCO3
(1mol CaCl2 · 2H2O)
3. Calculate the moles of CaCO3 produced from the mass of Na2CO3 used in the experiment.
_______g Na2CO3 x (1mol Na2CO3) x (1mol CaCO3) = _________ mol CaCO3
(105.99g Na2CO3) (1mol Na2CO3)
4. Which reactant is the limiting reactant?
Which reactant is in excess?
_______________________
________________________
5. Calculate the theoretical yield of CaCO3 from the limiting reactant.
________mol CaCO3 x (100.09g CaCO3) = ________g CaCO3 Theoretically Produced
(1mol CaCO3)
Letter
A, L
B, K
Grams of CaCl2 · 2H2O
2.500
2.000
Grams of Na2CO3
1.000
1.000
C, J
2.600
1.100
1.039g
D, I
2.400
1.050
0.9916g LR: Na2CO3
E, H
2.000
1.500
1.362g
LR: CaCl2 · 2H2O
F, G
1.400
1.250
0.9532g
LR: CaCl2 · 2H2O
Theoretical Yield
0.9443g LR:Na2CO3
0.9443g LR: CaCl2 · 2H2O
LR: Na2CO3
6. What is the percent yield of the reaction?
Percent Yield = Actual yield x 100 = XXXXg CaCO3 x 100 = _____% yield
Theoretical yield
XXXXg CaCO3
Limiting Reagent Stoichiometry
17
Limiting Reagent POST-LAB QUESTIONS--KEY
1. Describe the effect on the actual yield AND percent yield (Increase or Decrease) if the following
errors occurred. Explain why:
a) The filter papers were still wet when the final product was weighed:
Actual Yield: Increase the mass because extra water mass
Percent Yield: Increase the actual yield, therefore increase percent yield
b) Some product fell through the filter papers and was seen in the filter flask during the vacuum
filtration:
Actual Yield: Decrease the mass because lost in filtrate
Percent Yield: Decrease the actual yield, therefore decrease percent yield
c) Some product fell on the floor and was scraped up with some dust particles. It was then
weighed for the final mass.
Actual Yield: Increase the mass because extra dust mass
Percent Yield: Increase the actual yield, therefore increase percent yield
2. The following questions are based on the reaction below:
2 KC1 (aq) + 2 MnCl2 (aq) + 5 PbO2 (aq) + 4 HC1 (aq)
2 KMnO4 (aq) + 5 PbC12 (s) + 2 H2O (l)
a) 25.00g of MnC12 is reacted with 100.0g of PbO2 and excess KCl and HCl. How many grams of
KMnO4 can be produced by this reaction?
25.00g MnCl2 x (1 mol MnCl2 / 125.85g MnCl2) x (2mol KMnO4 / 2 mol MnCl2) = 0.198649mol KMnO4
100.0g PbO2 x (1mol PbO2 / 239.20g PbO2) x (2mol KMnO4 / 5mol PbO2) = 0.16722mol KMnO4
LR!!
0.16722mol KMnO4 x (158.04g KMnO4 / 1mol KMnO4) = 26.43g KMnO4
b) If 21.42 g of KMnO4 is actually produced, what is the percent yield? (If the grams of KMnO4
were not obtained in (a), use 35.03g KMnO4)
Percent Yield = 21.42g KMnO4 x 100 = 81.04% yield
26.43g KMnO4
Limiting Reagent Stoichiometry
35.03g  61.15%
18
Name
Section
Date
Limiting Reagent
PRE-LABORATORY EXERCISE--KEY
Read the prelaboratory notes, the experiment, and then complete the following as directed.
1. Define the following terms.
a. Limiting Reactant
Reactant that runs out first, determines the theoretical yield
b. Theoretical Yield
The maximum theoretical amount of a given product that can be formed when the limiting
reactant is completely consumed
c. Percent Yield
Percent Yield is a comparison of the actual yield to the theoretical yield.
2. a. Balance the equation and include the states of matter:
____ CaCl2 · 2H2O (aq) + ____ Na2CO3 (aq)
____ CaCO3 ( s ) + __2__ NaCl ( aq ) + ___2__ H2O ( l )
b. The mol ratio between CaCl2 · 2H2O and CaCO3 is… 1 mol CaCl2 · 2H2O = 1 mol CaCO3
c. How many mols of CaCO3 would be produced from 3.00mols of Ca(NO3)2 · 2H2O? Show your
work.
3.00mols CaCl2 · 2H2O x (1mol CaCO3 / 1mol CaCl2 · 2H2O) = 3.00 mols CaCO3
d. The mol ratio between Na2CO3 and CaCO3 is…
1 mol Na2CO3 = 1 mol CaCO3
e. How many mols of CaCO3 would be produced from 2.00mols of Na2CO3? Show your work.
2.00mols Na2CO3 x (1mol CaCO3 / 1mol Na2CO3) = 2.00 mols CaCO3
f. What is the mass of CaCO3 produced from the limiting reactant?
2.00 mols CaCO3 x (100.09g CaCO3 / 1mol CaCO3) = 200.g CaCO3
Limiting Reagent Stoichiometry
19
Limiting Reagent---Laboratory Prep Sheet
Materials
SOLID REAGENTS:
• Na2CO3
• CaCl2 • 2H2O
sodium carbonate anhydrous
calcium chloride dihydrate
EQUIPMENT:
• 125 mL Erlenmeyer flasks
• Rubber policeman
• Stirring rod
• Paper filter paper
• Weigh Boats
• Büchner funnel
• 250mL filter flask
• 250mL Erlenmeyer flask
• Y-tube
• Short tubing (8 cm)
• Pinch clamp
• Ring stand
• Ring stand clamps
• 2” rubber hose
~1 - 2 grams/group
~1 - 2 grams/group
2 per group
1 per group
2 per group
2 per group (minimum)
2 per group (minimum)
1 per group
1 per group
1 per group
1 per group
1 per group
1 per group
1 per group
2 per group
2 per group
OTHER INSTRUCTIONS:
• Büchner funnel and water trap demonstration setup for instructors
• For each analytical balance, have a small container of CaCl2 · 2H2O and a small container of Na2CO3, so
that students can weigh out their own chemicals.
Answer Key for Unknowns
Letter
A, L
B, K
Grams of CaCl2 · 2H2O
2.500
2.000
Grams of Na2CO3
1.000
1.000
C, J
2.600
1.100
1.039g
D, I
2.400
1.050
0.9916g LR: Na2CO3
E, H
2.000
1.500
1.362g
LR: CaCl2 · 2H2O
F, G
1.400
1.250
0.9532g
LR: CaCl2 · 2H2O
Limiting Reagent Stoichiometry
Theoretical Yield
0.9443g LR:Na2CO3
0.9443g LR:Na2CO3
LR: Na2CO3
20
A Series of Copper Reactions and Percent Yield
Developed by Kathy Pickham, Miramar College
California Science Standard, Chemistry Grades 9-12
3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the
ability to calculate the mass of products and reactants. As a basis for understanding this concept:
a. Students know how to describe chemical reactions by writing balanced equations.
d. Students know how to determine the molar mass of a molecule from its chemical formula and a table
of atomic masses and how to convert the mass of a molecular substance to moles, number of
particles, or volume of gas at standard temperature and pressure.
e. Students know how to calculate the masses of reactants and products in a chemical reaction from the
mass of one of the reactants or products and the relevant atomic masses.
f. * Students know how to calculate percent yield in a chemical reaction.
g. * Students know how to identify reactions that involve oxidation and reduction and how to balance
oxidation-reduction reactions.
Green Chemistry
This lab starts with less hazardous materials, and skips an environmentally unfriendly and hazardous procedure
step. Traditionally, students start with solid Copper and react it with concentrated nitric acid to form copper (II)
nitrate, smog(NO2), and water. In this "green" lab replacement, Copper (II) nitrate trihydrate is used as the
starting material, avoiding hazardous reagents and generation of hazardous products.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #3 - Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methodologies should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #4 - Designing Safer Chemicals: Chemical products should be designed to preserve efficacy of function
while reducing toxicity.
Objective
The goal of the experiment is to become familiar with basic laboratory procedures, to be exposed to a typical
transition element, and to understand the concept of percent yield.
Materials
Equipment and Supplies:
• Iron ring and ring stand
• Drying oven
• 250 mL beakers (2)
• Stirring rod
• Evaporating dish
• Balance
• Red litmus paper
• spatula
Equipment and Supplies (cont'd):
• 10- and 50 mL grad cylinder
• Boiling chips
Reagents (Solids and Solutions):
• Cu(NO3)2▪3H2O
• Granular zinc
• M H2SO4
• M HCl
• 3.0 M NaOH
21
Reagent Setup:
Safety
Many of the solutions used in this lab are hazardous. Do not allow any of them to come in contact with your
skin or clothing. In particular, H2SO4(aq), HCl(aq) and NaOH(aq) are very corrosive. Treat acid spills with sodium
bicarbonate (NaHCO3). If skin is exposed to base, wash thoroughly with water.
Waste
Please be sure to use the proper waste container for your waste. See picture below.
A Series of Copper Reactions and Percent Yield
22
Introduction
Most chemical syntheses involve separation and purification of the desired product from unwanted side
products. Some methods of separation, such as filtration, sedimentation, decantation, extraction, and
sublimation are useful to separate substances in a mixture. This experiment is designed as an introduction to
carrying out some of these operations. At the same time you will become more acquainted with two
fundamental types of chemical reactions. By means of these reactions you will carry out several chemical
transformations involving copper ions and you will finally obtain pure copper solid with maximum efficiency.
The chemical reactions copper will undergo are the following:
1. Cu(NO3)2(aq) + 2 NaOH(aq)  Cu(OH)2(s) + 2 NaNO3(aq)
2. Cu(OH)2(s)  CuO(s) + H2O(g)
3. CuO(s) + H2SO4(aq)  CuSO4(aq) + H2O(l)
4. CuSO4(aq) + Zn(s)  ZnSO4(aq) + Cu(s)
Each of these reactions proceeds to completion. Methathesis reactions proceed to completion whenever one of
the components is removed from the solution, such as in the formation of a gas or an insoluble precipitate. This
is the case in some of the reactions above. In one of the redox equations above, the driving force for
completion is that one of the metal ions has a lower ionization energy or oxidation potential than the other
metal.
The objective in this experiment is to obtain copper solid in analytically pure form. This is a test of your
laboratory skills. The percent yield of the copper can be expressed as the ratio of the actual yield to theoretical
yield multiplied by 100:
%Yield =
actual yield
x100
theoretical yield
Procedure
(Record all your observations in your lab notebook as you proceed with this experiment. Watch carefully for all
signs of a chemical reaction)
1. Carefully weigh 1.5 g of Cu(NO3)2▪3H2O into a 250-mL beaker. Add approximately 100 mL distilled water to
the beaker with the Cu(NO3)2▪3H2O and stir to dissolve.
2. Add 6.0 mL of 3.0 M NaOH (Caution: corrosive) dropwise to the Cu(NO3)2 solution and watch for
precipitation of Cu(OH)2. Test the solution with litmus paper by dipping a glass stir rod into the solution and
then touching it to a piece of red litmus paper. If it turns blue it is basic, if it stays red the solution is acidic
so add another few drops of 3.0 M NaOH until the solution turns basic.
3. Add a boiling chip to the solution and place on a hot plate at a setting of 6 - 7. Heat the solution to just
boiling while stirring constantly to avoid bumping of the solution. Remove the solution from the hot plate
and set on the lab bench.
4. Set a 250 mL beaker with ~150 mL of distilled water on the hot plate to heat. Allow the black solid CuO to
settle, then carefully decant the supernatant liquid into a waste beaker.
5. Add ~100 mL of the heated distilled water to the CuO in the beaker, stir with a glass rod and then allow the
solid to settle again. Decant the liquid into a waste beaker and remove the boiling chip using a spatula.
6. Add 15 mL of 6.0 M sulfuric acid (Caution: corrosive) to the solid. Swirl to mix.
A Series of Copper Reactions and Percent Yield
23
7. In a fume hood, add 2.0 g of granular zinc all at once. Stir until the supernatant liquid is colorless and the
hydrogen gas evolution has become slow. Decant the liquid into a waste beaker.
8. Still in the fume hood, add 10 mL of 6.0 M HCl (Caution: corrosive) to react with the excess zinc metal.
When hydrogen gas evolution becomes very slow and the zinc metal is gone, remove from the hood.
9. While the zinc is reacting, label a clean, dry evaporating dish with your name and weigh the dish.
10. Decant the supernatant liquid into a waste beaker and add 10 mL of distilled water to the copper solid and
then transfer the solid to a pre-weighed, labeled evaporating dish.
11. Decant the supernatant liquid and then rinse the copper solid 2 times with 10 mL of distilled water,
decanting the liquid into a waste beaker each time. After the final rinse, remove as much water as possible
using a disposable pipet.
12. Place the evaporating dish in the drying oven (120°C) for 30 min.
13. Carefully remove the hot evaporating dish from the drying oven and let cool on the lab bench. When cool,
weigh the evaporating dish with the copper.
14. Return the evaporating dish to the oven for 5 additional minutes, then cool and weigh again. Repeat until
the weighings are within a few milligrams of each other.
Results and Calculations
You should have detailed observations for each step of the experiment and you must write balanced chemical
equations for the reactions which take place in steps 2, 3, 6 and 7. This is best done as you are performing the
experiment. Use the formula given in the introduction to determine the percent yield.
Discussion and Conclusion
1. Explain, using atomic theory, why the amount of copper metal present in step 14 should be the same as the
amount of copper present in the starting material.
2. What is the purpose of heating the copper(II)hydroxide in step 3?
3. Explain why nitric acid cannot be used in steps 6 and 8 in place of sulfuric and hydrochloric acid,
respectively.
4. Discuss the sources of error present in the experiment which could result in the percent recovery obtained
being too high or too low.
Post-Lab Questions
1. How many milliliters of 3.0 M of H2SO4 are required to react with 0.80 g of CuO according to eqn 3?
2. If 3.00 g of Zn is allowed to react with 1.75 g of CuSO4 according to equation 4, how many grams of Zn will
remain after the reaction is complete?
A Series of Copper Reactions and Percent Yield
24
Chemical Reactions
Developed by Lisa Selchau, Miramar College
California Science Standards, Chemistry Grades 9-12
2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from
electrostatic forces between electrons and protons and between atoms and molecules. As a basis for
understanding this concept:
a. Students know atoms combine to form molecules by sharing electrons to form covalent or metallic
bonds or by exchanging electrons to form ionic bonds.
3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the
ability to calculate the mass of products and reactants. As a basis for understanding this concept:
a. Students know how to describe chemical reactions by writing balanced equations.
e. Students know how to calculate the masses of reactants and products in a chemical reaction from the
mass of one of the reactants or products and the relevant atomic masses.
g. * Students know how to identify reactions that involve oxidation and reduction and how to balance
oxidation-reduction reactions.
5. Acids, bases, and salts are three classes of compounds that form ions in water solutions. As a basis for
understanding this concept:
a. Students know the observable properties of acids, bases, and salt solutions.
Green Chemistry
In order to make this lab kinder to our environment, we have cut down on the amount of chemicals we use by
using well-plates to see double displacement reactions instead of test tubes. We are using some everyday
chemicals such as baking soda that may be disposed of in the sink. The amount of solids and solutions used has
been cut in half for most steps. Steps which produced hazardous chemicals such as sulfur dioxide have been
deleted.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Principle #3: Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methods should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #12: Inherently Safer Chemistry for Accident Prevention: Substances and the form of a substance
used in a chemical process should be chosen to minimize the potential for chemical accidents, including
releases, explosions, and fires.
Purpose
•
To understand and observe the following chemical reactions: combustion, combination (synthesis),
decomposition, single displacement, double displacement, and acid/base reactions.
•
To identify the products of each reaction.
•
To write a balanced equation for each reaction.
25
Introduction
Chemists have learned that a chemical change involves a rearrangement of the ways in which atoms are
grouped together. We refer to this as a chemical reaction. We represent a chemical reaction by writing a
chemical equation in which the chemicals present before the reaction (the reactants) are shown to the left of
the arrow and the chemicals formed by the reaction (the products) are shown to the right of the arrow. The
arrow indicates the direction of the change and is read as “yields” or “produces”. Chemical reactions are
categorized into the following basic types: Combustion, combination (synthesis), decomposition, single
displacement, double displacement, and acid/base reactions. In this lab, you will observe different examples of
these reaction types, and identify the products of each reaction. You will write balanced chemical equations for
each reaction. While observing these reactions, keep in mind the changes indicating that a reaction has
occurred. These changes are either: 1-color change, 2-formation of a precipitate, 3-temperature change, or 4formation of a gas.
Background
1. A Combustion reaction is when a substance is reacted with oxygen (O2) gas and produces carbon dioxide
(CO2) gas, water vapor and energy. For example:
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O(g) + energy
2. A Combination (or synthesis) reaction is the combination of two substances to make a single substance.
Often, but not always, this involves the reaction of two elements to form a compound. In general, a
combination reaction can be represented by A + B → C. For example:
2H2 (g) + O2 (g) → 2H2O(l)
3. A Decomposition reaction is the opposite of a combination reaction. A substance is broken down into
simpler substances. In general, a decomposition reaction can be represented as X → Y + Z. Some examples
are:
H2CO3 (aq) → H2O (l) + CO2 (g)
NH4OH (aq) → NH3 (g) + H2O (l)
H2SO3 (aq) → H2O (l) + SO2 (g)
(NH4)2S (aq) → NH3 (g) + H2S (g)
4. A Single Displacement reaction occurs when an ionic compound reacts with an element. In generalized
form, the reaction is: XY + A → AY + X. For example:
2HCl(aq) + Zn (s) → ZnCl2 (aq) + H2 (g)
5. A Double Displacement or ion exchange reaction occurs when two aqueous solutions, which both contain
ions, are mixed together and one of the following occurs: (1) a precipitate is formed or (2) a gas is produced
because one of the product further undergoes decomposition. For example:
K2CrO4 (aq) + Ba(NO3)2 (aq) → BaCrO4 (s) + 2KNO3 (aq)
(1)
NH4Cl (aq) + Na2S (aq) → NaCl (aq) + (NH4)2S (aq) → NaCl (aq) + NH3 (g) + H2S (g)
(2)
Chemical Reactions
26
These two reactions are examples of molecular reactions, where all the compounds are written in their neutral
molecular formula despite the fact that they are dissolved in water (aq). The driving force for a double
displacement reaction is usually the formation of a precipitate. These solubility rules allow you to predict if a
product of a molecular reaction is aqueous (soluble in water) or a solid precipitate.
SOLUBILITY RULES FOR IONIC COMPOUNDS
1.
2.
3.
4.
5.
All compounds containing Li+, K+, Na+, NH4+, NO3–, CH3COO– are soluble in water
Most of the halide ions (X–), Cl–, Br–, I– are soluble except with Ag+, Hg22+, and Pb2+.
Most of the SO42– are soluble except with Ca2+, Sr2+, Ba2+ and Pb2+..
Most of CO32–, PO43–, and CrO42–are insoluble except rule 1.
Most of S2– and OH– are insoluble except rule 1 and with Ca2+, Sr2+, and Ba2+.
Solid compounds are often observed as cloudy solutions. The solids eventually settle out, leaving a clear
“supernate” (or supernatant solution) above the precipitate. Gaseous compounds will form from the
decomposition of certain compounds on the product side of the double displacement reaction (look on previous
page under decomposition).
6. Acid and Base Reactions (Neutralization)
This is a highly specialized type of double displacement reaction where one of the products formed is water.
When the reactants are an acid (e.g., hydrochloric acid or anything that ionizes to form H+ in solution) and a base
(e.g., sodium hydroxide or anything that ionizes to form OH− in solution), the “insoluble” product of the double
displacement reaction is liquid water (the product of H+ reacting with OH−). We will use indicators to help us
verify that the acid/base reaction has occurred. Acid-base neutralization reactions often generate heat, which
can be detected by carefully touching the outside of the test tube as you mix the two reactants. An example of
an acid/base reaction is:
Molecular Equation:
H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2H2O (l)
Complete Ionic Equation:
2H+(aq) + SO42– (aq) + 2Na+ (aq) + 2 (OH–) (aq) → 2Na+ (aq) + SO42– (aq) + 2H2O (l)
Find spectator ions (ions that are IDENTICAL on both side of the reaction), remove to write the
Net Ionic Equation:
2H+ (aq) + 2(OH–) (aq) → 2H2O (l)
Chemical Reactions
27
Materials
Supplies:
• Test tubes
• Tongs
• Crucibles
• Bunsen burner
• Ring stands
• Clamp with o-ring
• Well plate
• Stirring rod
• Waste container
• Wooden splint
Reagents (Solids and Solutions):
• 6.0M Hydrochloride acid (HCl)
• Magnesium (approx 1-cm in length)
• 0.10 M Silver nitrate
• Copper wire (small piece)
• Zinc (small piece of Zinc)
• 0.10 M Potassium iodide
• 0.10 M Sodium phosphate
• 0.10 M Ammonium hydroxide
• 0.10 M Nitric acid
• 0.10 M Sodium hydroxide
• 0.10 M Sulfuric acid
• Phenolphathalein indicator
Safety Concerns
Please wear your safety goggles. If your hands come in contact with a chemical, you should wash them in water
for 10 min and follow with soap and water.
Remember, if you are heating a test tube over a flame, point the tube AWAY from yourself and others.
Procedure
For each reaction below, it is not critical to use the chemicals in any precise amounts. Once you see a reaction
happening, it is not necessary to add any more of the reactant. For example, when you add a drop of silver
nitrate to a solution of potassium chloride, you will immediately see a precipitate (silver iodide) form as the
drop hits the solution. There is no need to add any more solution. It will just form more precipitate and it won’t
look any different.
Part A: Combination Reaction
1. Hold a piece of magnesium (approx 4 cm in length) with crucible tongs and ignite the metal in a Bunsen
burner flame. CAUTION: Do NOT look directly at the burning magnesium! Observe out of the corner of
your eye. Throw the solid product directly into the trash when it is completely cooled.
Part B: Decomposition Reactions
1. Cobalt(II) chloride hexahydrate and heat: Place a few crystals (about 1 cm high) of solid cobalt(II) chloride
hexahydrate into a small test tube. Holding the test tube with your test tube holder, heat strongly with a
Bunsen burner flame. Note the change in color and texture and observe the inside wall of the test tube.
Allow the tube to cool, then add a few drops of water. Record your observations. Dispose of in chemical
waste container.
2. Sodium bicarbonate and heat: Add sodium bicarbonate (baking soda) into a 250 ml Erlenmeyer flask so as
to sparsely cover the bottom. Support the flask with a utility clamp on a ring stand with wire gauze and Oring.
a. Hold a flaming wooden splint inside the flask and observe how long it burns. (A flaming wooden
splint will gradually go out in air, burn more brightly in the presence of oxygen gas, and go out
quickly in the presence of carbon dioxide.)
Chemical Reactions
28
b. Heat the flask strongly with a Bunsen burner until moisture collects at the neck of the flask; again
hold a flaming wooden splint inside the flask and observe how long it burns.
c. Dispose of in the sink.
Part C: Single Displacement Reactions
1. Place about 40 drops of 6M HCl in a test tube. CAUTION: Hold the test tube with a test tube holder, NOT
your bare hand. Drop in a small piece of magnesium. Hold a flaming wooden splint over the mouth of the
test tube immediatly after dropping the magnesium in it. (A pop indicates hydrogen gas being produced.)
2. Put about 40 drops of silver nitrate into a test tube and add a small piece of copper wire. Allow a few
minutes for the reaction to take place while you start step C3 and then record your observations.
3. Place about 40 drops of 6M HCl in a test tube. CAUTION: Hold the test tube with a test tube holder, NOT
your bare hand. Drop in a small piece of zinc. Hold a flaming wooden splint over the mouth of the test tube
immediately after dropping the zinc into it. (A pop indicates hydrogen gas being produced.)
Dispose of all waste from part C into chemical waste container.
For the following double displacement reactions, use the well plates provided on the cart in the back of the
classroom.
Part D: Double Displacement Reactions
1. Place 5-10 drops of silver nitrate into a well on the CLEAN well plate. Add 5-10 drops of potassium iodide.
Mix well and observe the reaction.
2. Place 5-10 drops of silver nitrate into a well on the CLEAN well plate. Add 5-10 drops of sodium phosphate.
Mix well and observe the reaction.
3. Place 5-10 drops of iron (III) chloride into a well on the CLEAN well plate. Add 5-10 drops of ammonium
hydroxide. Mix well and observe the reaction.
Dispose of all waste from part D into chemical waste container.
Part E: Neutralization Reactions
1. Put 20 drops of 0.1 M nitric acid into a test tube along with 2 drops of phenolphthalein indicator. Add about
5 drops of 0.1 M sodium hydroxide and mix well with a stirring rod. Continue to add approximately 5 drops
of 0.1M NaOH at a time (counting the drops) and mix with a stirring rod until a permanent color change
throughout the whole solution is observed. Record the total number of drops needed to observe the
permanent color change. Note: Phenolphthalein is an indicator, which is colorless in acidic conditions and
pink under basic conditions.
2. Put 20 drops of 0.1 M sulfuric acid into a test tube along with 2 drops of phenolphthalein indicator. Add
about 5 drops of 0.1 M sodium hydroxide and mix well with a stirring rod. Continue to add approximately 5
drops of 0.1M NaOH at a time (counting the drops) and mix with a stirring rod until a permanent color
change throughout the whole solution is observed. Record the total number of drops needed to observe
the permanent color change.
Dispose of all waste from part E into chemical waste container.
Chemical Reactions
29
Name:
Pre-Lab Questions
1. For each reaction below:
Identify each reaction as (A) combination, (B) decomposition, (C) single displacement,
or (D) double displacement.
Write the upper-case letter for each equation in the right margin of this page.
Also, balance each equation, if necessary.
KClO3 →
KCl
+
O2
ZnBr2 +
K3PO4
K
+
H2O
→
KOH
Sn
+
HCl
→
SnCl2 +
Ba(OH)2
+
H2SO4
→
BaSO4
+
H2O
(NH4)2Cr2O7
+
Co(NO3)3
→
Co2(Cr2O7)3
+
NH4NO3
→
Zn3(PO4)2
+
+
KBr
H2
H2
2. Look for a YOUTUBE video of sodium reacting with water and potassium reacting with water. Write down
the two urls and write the observation you make of the two reactions in the box below.
URL of the
YouTube Video
Sodium Reacting with Water
Potassium Reacting with Water
Observation of
the reaction
3. True or False? A soluble compound dissolves in water.
4. True or False? The symbol “(aq)” next to a reactant or product in a chemical equation means that a
substance is insoluble.
5. True or False? The symbol “(s)” next to a reactant or product in a chemical equation means that a substance
is soluble.
Chemical Reactions
30
Pre-Lab Questions (cont'd)
Name:
6. Consider the reaction in part D-1 between aqueous silver nitrate and aqueous potassium iodide to form
aqueous potassium nitrate and solid silver iodide. Write the balanced molecular equation, the complete
ionic equation, and the net ionic equation for this reaction.
Balanced Molecular Equation:
Complete Ionic Equation:
Net Ionic Equation:
Chemical Reactions
31
Data and Results
Part A: Combination (Synthesis) Reaction
1. Observation:
Name:
Partner:
Part B: Decomposition Reactions
1. Observation:
2. a. How long did the splint burn?
b. How long did the splint burn?
Part C: Single Displacement Reactions
1. Observation:
2. Observation:
3. Observation:
Part D: Double Displacement Reactions
1. Observation:
2. Observation:
3. Observation:
Chemical Reactions
32
Data and Results (cont'd)
Part E: Neutralization Reactions (specific type of Double Displacement)
1. Observation:
Name:
How many drops of NaOH required?
2. Observation:
How many drops of NaOH required?
For parts A, B, and C translate each word equation into a balanced chemical equation.
A. Magnesium metal(s) + oxygen (g) → magnesium oxide(s)
B. 1. Cobalt (II) chloride hexahydrate(s) → Cobalt (II) chloride(s) + water(l)
B. 2. Sodium bicarbonate(s) → sodium carbonate(s), + water(l) + carbon dioxide(g)
C. 1. Magnesium metal(s) + hydrochloric acid(aq) → magnesium chloride(aq) + hydrogen(g)
C. 2. Copper metal(s) + silver nitrate(aq) → copper (II) nitrate(aq) + silver(s)
C. 3. Zinc metal(s) + hydrochloric acid(aq) → zinc chloride(aq) + hydrogen(g)
Chemical Reactions
33
Data and Results (cont'd)
Name:
For parts D and E translate each word equation into a balanced chemical equation, write the complete ionic
equation, and the net ionic equation.
D. 1. Silver nitrate(aq) + potassium iodide(aq) → silver iodide(s) + potassium nitrate(aq)
Complete ionic equation:
Net ionic equation:
D. 2. Silver nitrate(aq) + sodium phosphate(aq) → silver phosphate(s) + sodium nitrate(aq)
Complete ionic equation:
Net ionic equation:
D. 3. Iron (III) chloride(aq) + ammonium hydroxide(aq) → iron (III) hydroxide(s) + ammonium chloride(aq)
Complete ionic equation:
Net ionic equation:
Chemical Reactions
34
Data and Results (cont'd)
Name:
E. 1. Nitric acid(aq) + sodium hydroxide(aq) → sodium nitrate(aq) + water(l)
Complete ionic equation:
Net ionic equation:
E. 2. Sulfuric acid(aq) + sodium hydroxide(aq) → sodium sulfate(aq) + water(l)
Complete ionic equation:
Net ionic equation:
Chemical Reactions
35
Post-Lab Questions
1. Note the differences in length of time the splint burned in Part B2 with and without heating the sodium
bicarbonate. Explain why, in terms of the chemical reaction.
2. In part C, did your observations concerning the flaming wooden splint agree with the reactions that you
wrote for this section? Explain.
3. In parts E1 and E2 you start with equal amounts of the 0.1M acids. Looking at the balanced chemical
equations for these parts, why did the sulfuric acid need more sodium hydroxide to turn the phenolphthalein
pink than the nitric acid?
4. Predict the products for and balance each of the following double displacement reactions. If no reaction
occurs, write “NR”.
AgNO3 (aq) + NaCl (aq) →
NaNO3 (aq) + KI (aq)
Chemical Reactions
→
36
Chemical Reactions Lab Materials Sheet
General Supplies:
Test tubes, tongs, crucibles, bunsen burners, ring stands, clamp
with o-ring, well plate, stirring rod
Reagents and Additional Supplies
Quantity/Semester
Magnesium (approx 4cm in length)
12-piece
Cobalt (II) chloride hexahydrate
20g
Sodium bicarbonate (baking soda)
25g
Wooden splint
24 piece
6.0M Hydrochloride acid (HCl)
80-mL
Magnesium (approx 1-cm in length)
12-piece
0.10M Silver nitrate
5-mL
Copper wire (small piece)
12
Zinc (small piece of Zinc)
25g
Well plate
12
0.10M Potassium iodide
25-mL
0.10M Sodium phosphate
25-mL
0.10M Ammonium hydroxide
25-mL
0.10M Nitric acid
25-mL
Phenolphathalein indicator
4-vials
0.10M Sodium hydroxide
100-mL
0.10M Sulfuric acid
25-mL
Waste container
1
Chemical Reactions
37
Eco-Friendly Equilibrium (Teacher/TA Edition)
Developed by Michael Sixtus (retired), Bonita Vista High School
This protocol was developed with California SB 70 funding in an effort to bring sustainability practices to
teaching laboratories and to increase awareness of sustainability in the next generation of the California
workforce. In other words, these protocols ‘green’ our labs and ‘green’ our students.
California Science Standards, Chemistry Grades 9-12:
9. Chemical Equilibrium is a dynamic process at the molecular level. As a basis for understanding this concept:
a. Students know how to use Le Chatelier’s principle to predict the effect of changes in concentration,
temperature and pressure.
b. Students know equilibrium is established when forward and reverse reaction rates are equal.
c. * Students know how to write and calculate an equilibrium constant expression for a reaction.
Green Chemistry
The traditional labs for dynamic equilibrium and Le Chatelier's principle involve the use of more toxic substances
and larger quantities of materials than this green lab replacement. More specifically:
•
The solubility equilibrium for NaCL is a microscaled version of the traditional lab, minimizing the volume
of concentrated HCl used, and thus the volume of solution to be neutralized at the end of the activity.
•
The acid-base equilibrium activity is a microscaled version of the traditional lab.
•
The pressure effects on gas equilibrium is an original activity that requires no hazardous materials or
techniques and only relatively inexpensive and reusable equipment to demonstrate a frequently
overlooked equilibrium situation.
•
The temperature effects on equilibrium activity, while making use of a less-than-desirable reagent
(Cobalt Chloride), minimizes both the concentration used and possible student exposure to the material
by having it in a sealed pipette. As the containers are sealed, they are also reusable from class to class
and year to year, minimizing the overall quantities used in the lab.
•
The use of zinc nitrate and copper sulfate in the solubility and complex ion equilibria and reversible
reactions activity replaces the more common ferrocyanates used in other sources, thus lowering the
toxicity of the reagents used. The microscale approach also reduces volumes, waste and exposure.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Principle #3: Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methods should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #12: Inherently Safer Chemistry for Accident Prevention: Substances and the form of a substance
used in a chemical process should be chosen to minimize the potential for chemical accidents, including
releases, explosions, and fires.
38
Objectives
In this series of experiments you will:
•
demonstrate the basics of dynamic equilibrium, graph the results, and discuss the establishment of the
equilibrium system;
•
demonstrate various tests of Le Chatelier’s Principle using various equilibrium systems and stresses.
Introduction
In the study of chemical reactions, the assumption has been made that most reactions proceed from reactants
to products until at least one of the reactants was completely consumed.
In some reactions, however, the products remain in contact with each other and if the energy requirement is
low enough, the products react to form the reactants again. For example, if hydrogen and iodine molecules are
placed in a sealed container, they will react to form hydrogen iodide molecules.
H2(g) + I2(g)  2HI(g)
The HI molecules also collide and react to form hydrogen and iodine again.
2HI(g)  H2(g) + I2(g)
When the rates of the forward and reverse reactants are equal, a state of chemical equilibrium is said to exist.
H2(g) + I2(g)  2HI(g)
Chemical equilibrium is a dynamic equilibrium, in which the forward and reverse reactions continue at equal
rates so that there is no net change in the quantities of reactants or products. The concentrations of H2, I2, and
HI become constant at a given temperature and do not change until conditions are changed.
According to Le Chatelier’s Principle, any equilibrium system subjected to a stress tends to change so as to
relieve the stress. If the concentration of H2 is increased, a stress would be created favoring the formation of
more HI. The final effect on the equilibrium concentrations would be a decreased concentration for I2 and an
increased concentration for HI and H2.
STRESS
H2 increased
Concentration change
H2(g)
Inc.
RESULT
+
I2(g)
Dec.

2HI(g)
Inc.
Another stress can be caused by decreasing the concentration of one of the substances in the equilibrium
system. For example, decreasing the concentration of I2 would favor a shift to replace the I2, thus decreasing the
concentration of HI and increasing the concentration of H2.
STRESS
I2 decreased
Concentration change
Eco-Friendly Equilibrium (Teacher)
H2(g)
Inc.
RESULT
+
I2(g)
Dec.

2HI(g)
Dec.
39
Materials (per Lab group/station)
Equipment:
• Hot plate for water bath
Supplies:
• 1 100ml graduated cylinder
• 1 25ml graduated cylinder
• 2 pairs of straws/pipettes of different diameters
• 13 Beral pipettes
• 60-ml syringe plus luer-lok cap
• 1 24-well microplate
• Ice for cold-water bath
• Waste disposal container
Solutions:
Note: You will not need more than 200ml of any solution per period of 30-32 students
•
•
•
•
•
•
•
•
•
•
•
•
•
•
•
1 2L bottle of club soda (other clear carbonated beverages will work but are sticky)
Saturated NaCl (I just add a very large amount of salt to a 500ml bottle, add water, shake well and let it
stand overnight)
12M HCl (Standard concentrated HCl)
0.1M HCl (8.2ml of 12M in 1L H2O)
Distilled H2O
Bromothymol Blue solution (0.04%; standard solution concentration)
0.1M NaOH (4g in 1L)
0.1M CoCl2 in sealed Beral pipettes (I use normal ethanol as solvent; the colors are more vivid. Make
them up ahead of time and seal them using a pair of needle-nose pliers heated in a Bunsen burner; 23.8g
in 1L EtOH)
0.1M Zn(NO3)2 (29.8g in 1L H2O)
6M HCl (492ml in 1L H2O)
6M NaOH (240g in 1 L H2O)
6M NH3 (NH4OH) (402ml 0f 14.8M in 1L H2O)
0.1M CuSO4 (6.25g in 1L H2O)
3M H2SO4 (168ml of 18M in 1L H2O)
0.1M Mg(NO3)2 (This is only if you plan on the option for this experiment)
Eco-Friendly Equilibrium (Teacher)
40
Procedure
Note: You might considers setting up the various experiments as stations and have the students rotate
through them. Obviously this is contingent on the behavior of the students, the layout of your lab, etc.
A. Dynamic Equilibrium
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry
1. Label a 100ml graduated cylinder “A-reactants” and fill to the 50.0 ml mark with tap water. Label a 25ml
cylinder “B-products” and leave empty.
2. You and your partner are to transfer water simultaneously from one cylinder to the other, using
straws/pipettes of equal diameters. Lower the straws into your respective cylinders, and when each
straw touches bottom or as far as you can reach, place your index finger over the end of the straw.
Transfer the water collected to the other cylinder and allow the straw to drain. If you spill, remember
that A + B = 50ml.
3. Remove each straw and record the volume in each cylinder, being careful to read the meniscus to the
nearest 0.1ml.
4. Return the straws to their original cylinders and repeat the process, recording the data after each
transfer.
5. After three successive transfers result in no change in volume, stop and go on to step 6.
6. Repeat steps 1-5 using a larger diameter straw for “A” and a smaller one for “B”.
a. Add 20ml of water to “B” and repeat steps 1-5.
b. Add 20ml of water to “B” and repeat step 6.
In this experiment, which is an analogy for dynamic equilibrium, there is a variety of modifications that you
might consider assigning randomly to individual groups and then discussing their results as a class. These
include:
a. Using different diameter straws/pipettes for reactants and products
b. Adding more “reactants” after equilibrium is reached
c. Removing “reactants” after equilibrium is reached
d. Adding more “products” after equilibrium is reached
e. Removing products after equilibrium is reached
f. Use colored water as product in “e” and observe color change as well as volume
B. Le Chatelier’s Principle (Put on safety goggles and aprons)
1. Solubility Equilibrium for NaCl
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry; atom economy
a. Add 25 drops of saturated NaCl to one well of a 24-well microplate.
b. Add 5 drops of concentrated HCl to the well, one at a time, and record your observations.
c. Discard the solution in the sink and rinse the tray with plenty of water.
2. Acid-Base Equilibrium
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry; atom economy
a. Add 25 drops of distilled water to one well of a 24-well microplate. Add 5 drops of bromothymol
blue
b. Add 0.1M HCl one drop at a time until a color change occurs, then add 2 more drops. Mix by
shaking gently.
c. Now add 0.1M NaOH a drop at a time until a change occurs. Note the change.
d. Now add 0.1M HCl a drop at a time until a color change occurs. Note the change.
e. Discard the solution in waste container and rinse the tray with water.
Eco-Friendly Equilibrium (Teacher)
41
3. Pressure Effects on Gas Equilibrium
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry;
a. Remove plunger from syringe and attach cap. Add approx 15 ml of club soda, carefully reinsert the
syringe, invert the syringe, remove the cap, and depress the plunger to remove the air from the
syringe. Replace the cap.
b. Observe the bubbles in the syringe; they were not evident before the bottle was uncapped due to
the internal pressure that kept the equilibrium of the reaction:
H2CO3 (aq) H2O (l)+ CO2 (g)
favoring the right-hand side.
c. Increase the pressure in the syringe by pressing down on the plunger while resting the cap on the
table so it won’t pop off. Observe the change in the bubbles.
d. Decrease the pressure in the syringe by gently pulling on the plunger without totally removing it
from the syringe. Observe the change in the bubbles.
4. Temperature Effects on Equilibrium
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry; atom economy
a. Obtain 2 sealed Beral pipettes containing 0.1M CoCl2 from your teacher. Observe the color of the
solution.
b. Place one of the pipettes into a hot water bath for 1-2 minutes. Remove and compare the color
with that of the unheated tube. Record your observations. Continue to observe the unheated tube
until it returns to room temperature.
c. Place one of the pipettes into a cold water/ice bath for 2-3 minutes. Remove and compare the color
with that of the uncooled tube. Record your observations. Continue to observe the unheated tube
until it returns to room temperature.
d. Return the 2 pipettes to your teacher.
5. Solubility and Complex Ion Equilibria
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry; atom economy
a. Add 25 drops of 0.1M Zn(NO3)2 to each of 3 wells in a 24-well microplate. Add 1 drop of 6M NaOH
to each well and mix well. Record your observations.
b. To the first well, add 6M HCl drop by drop with mixing. Observe changes and stop adding when no
more changes are seen.
c. To the second well add 6M NaOH drop by drop with mixing and proceed as in “b”.
d. To the third well add 6M NH3 (NH4OH)drop by drop with mixing and proceed as in “b”.
e. Discard the solutions in the waste container and rinse the tray with water.
f. You may repeat the entire sequence with 0.1M Mg(NO3)2 at your teacher’s discretion.
6. Reversible Reactions with Copper (II) Sulfate.
Green Principles: Prevention of waste; Inherent safety; Safe chemicals; safe chemistry; atom economy
a. Add 25 drops of 0.1M CuSO4 to two wells in a 24-well microplate. One will be for reacting and the
other for comparison.
b. Add 6M NH3 drop by drop (with mixing between drops) until there is a definite color or appearance
change. Record your observations.
c. Continue to add NH3 until there is another color or appearance change and record your
observations.
d. Now add 3M H2SO4 dropwise to the solution until the original color is restored.
e. Discard the solutions in the waste container and rinse the plate with water.
Eco-Friendly Equilibrium (Teacher)
42
Safety Concerns:
All standard safety practices for a high school Chemistry lab should be followed at all times, including use of
goggles and aprons, hand washing after chemical use, proper disposal of waste materials, and proper cleanup
of lab equipment and spaces.
Overall Waste Disposal:
Using microscale techniques and less toxic chemicals reduces both the volume and hazard of the waste
produced. The liquid waste can be neutralized, diluted and poured down the sink without major concerns.
Toxicity of Chemicals Used (Oral Rat LD50, per kilogram):
Chemical
Oral Rat LD50, per kilogram
Cobalt chloride
766 mg
Copper(II)sulfate
300 mg
Ammonium hydroxide
350 mg
Zinc nitrate
1190 mg
Magnesium nitrate
5440 mg
Hydrochloric acid
Highly toxic/corrosive
Sulfuric acid
Severely corrosive
Sodium hydroxide
Corrosive solid/liquid
Sodium chloride
n/a
Computer-based Equilibrium Simulations:
1. Virtual Lab
http://www.capital.net/com/vcl/equil/equil.htm
Provides a good introduction to Equilibrium; pre-lab questions; simulations of Cobalt, Chromate,
Ammonium, Nitrogen Dioxide, Iron Thiocyanate, and Copper Sulfate systems; and post-lab questions.
No waste, no cost, and all green.
2. Gas Equilibrium simulation
http://www.chm.davidson.edu/ronutt/che115/equkin/equkin.htm
Provides a data-generating simulation of the effect of temperature on the equilibrium of a 2-gas system.
3. Simulation of Ca(OH)2 equilibrium system
http://www.chm.davidson.edu/vce/Equilibria/EquilibriumConstant.html
Students are able to determine the Keq of this experimental system.
4. Virtual CoCl2 equilibrium lab
http://www.chemcollective.org/vlab/vlab.php
Virtual lab simulation of the CoCl2 equilibrium system from CarnegieMellon.
5. Equilibrium
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animationsindex.htm
Bromine, Cobalt, and NO2 equilibrium animations.
Developed by Michael Sixtus, Retired AP/IB Science Instructor.
[email protected]
Eco-Friendly Equilibrium (Teacher)
43
Eco-Friendly Equilibrium (Student Version)
Developed by Michael Sixtus (retired), Bonita Vista High School
California Science Standards, Chemistry Grades 9-12
9. Chemical Equilibrium is a dynamic process at the molecular level. As a basis for understanding this concept:
a. Students know how to use Le Chatelier’s principle to predict the effect of changes in concentration,
temperature and pressure.
b. Students know equilibrium is established when forward and reverse reaction rates are equal.
c. * Students know how to write and calculate an equilibrium constant expression for a reaction.
Green Chemistry
The traditional labs for dynamic equilibrium and Le Chatelier's principle involve the use of more toxic substances
and larger quantities of materials than this green lab replacement. More specifically:
•
The solubility equilibrium for NaCL is a microscaled version of the traditional lab, minimizing the volume
of concentrated HCl used, and thus the volume of solution to be neutralized at the end of the activity.
•
The acid-base equilibrium activity is a microscaled version of the traditional lab.
•
The pressure effects on gas equilibrium is an original activity that requires no hazardous materials or
techniques and only relatively inexpensive and reusable equipment to demonstrate a frequently
overlooked equilibrium situation.
•
The temperature effects on equilibrium activity, while making use of a less-than-desirable reagent
(Cobalt Chloride), minimizes both the concentration used and possible student exposure to the material
by having it in a sealed pipette. As the containers are sealed, they are also reusable from class to class
and year to year, minimizing the overall quantities used in the lab.
•
The use of zinc nitrate and copper sulfate in the solubility and complex ion equilibria and reversible
reactions activity replaces the more common ferrocyanates used in other sources, thus lowering the
toxicity of the reagents used. The microscale approach also reduces volumes, waste and exposure.
Green Principles Addressed
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Principle #3: Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methods should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #12: Inherently Safer Chemistry for Accident Prevention: Substances and the form of a substance
used in a chemical process should be chosen to minimize the potential for chemical accidents, including
releases, explosions, and fires.
Objectives
In this series of experiments you will:
• demonstrate the basics of dynamic equilibrium, graph the results, and discuss the establishment of the
equilibrium system;
• demonstrate various tests of Le Chatelier’s Principle using various equilibrium systems and stresses.
44
Introduction
In the study of chemical reactions, the assumption has been made that most reactions proceed from reactants
to products until at least one of the reactants was completely consumed.
In some reactions, however, the products remain in contact with each other and if the energy requirement is
low enough, the products react to form the reactants again. For example, if hydrogen and iodine molecules are
placed in a sealed container, they will react to form hydrogen iodide molecules.
H2(g) + I2(g)  2HI(g)
The HI molecules also collide and react to form hydrogen and iodine again.
2HI(g)  H2(g) + I2(g)
When the rates of the forward and reverse reactants are equal, a state of chemical equilibrium is said to exist.
H2(g) + I2(g)  2HI(g)
Chemical equilibrium is a dynamic equilibrium, in which the forward and reverse reactions continue at equal
rates so that there is no net change in the quantities of reactants or products. The concentrations of H2, I2, and
HI become constant at a given temperature and do not change until conditions are changed.
According to Le Chatelier’s Principle, any equilibrium system subjected to a stress tends to change so as to
relieve the stress. If the concentration of H2 is increased, a stress would be created favoring the formation of
more HI. The final effect on the equilibrium concentrations would be a decreased concentration for I2 and an
increased concentration for HI and H2.
STRESS
RESULT
H2 increased
H2(g) +
I2(g)

2HI(g)
Concentration change
Inc.
Dec.
Inc.
Another stress can be caused by decreasing the concentration of one of the substances in the equilibrium
system. For example, decreasing the concentration of I2 would favor a shift to replace the I2, thus decreasing the
concentration of HI and increasing the concentration of H2.
STRESS
I2 decreased
Concentration change
Materials
Equipment:
• Hot plate for water bath
Supplies:
• 1 100ml graduated cylinder
• 1 25ml graduated cylinder
• 2 pairs of straws/pipettes of different
diameters
• 13 Beral pipettes
• 60-ml syringe plus luer-lok cap
• 1 24-well microplate
• Ice for cold-water bath
• Waste disposal container
Eco-Friendly Equilibrium (Student)
H2(g)
Inc.
RESULT
+
I2(g)
Dec.

2HI(g)
Dec.
Solutions:
• 1 2L bottle of club soda
• Saturated NaCl
• 12M HCl
• 0.1M HCl
• Distilled H2O
• Bromothymol Blue solution
• 0.1M NaOH
• 0.1M CoCl2 in sealed Beral pipettes
• 0.1M Zn(NO3)2
• 6M HCl
• 6M NaOH
• 6M NH3 (NH4OH)
• 0.1M CuSO4
• 3M H2SO4
• 0.1M Mg(NO3)2
45
Procedures
A. Dynamic Equilibrium
1. Label a 100ml graduated cylinder “A-reactants” and fill to the 50.0 ml mark with tap water. Label a 25ml
cylinder “B-products” and leave empty.
2. You and your partner are to transfer water simultaneously from one cylinder to the other, using
straws/pipettes of equal diameters. Lower the straws into your respective cylinders, and when each straw
touches bottom or as far as you can reach, place your index finger over the end of the straw. Transfer the
water collected to the other cylinder and allow the straw to drain. If you spill, remember that A + B = 50ml.
3. Remove each straw and record the volume in each cylinder, being careful to read the meniscus to the
nearest 0.1ml.
4. Return the straws to their original cylinders and repeat the process, recording the data after each transfer.
5. After three successive transfers result in no change in volume, stop and go on to step 6.
6. Repeat steps 1-5 using a larger diameter straw for “A” and a smaller one for “B”.
a. Add 20ml of water to “B” and repeat steps 1-5 (Test #1).
b. Add 20ml of water to “B” and repeat step 6 (Test #2).
B. Le Chatelier’s Principle (Put on safety goggles and aprons)
1. Solubility Equilibrium for NaCl
a. Add 25 drops of saturated NaCl to one well of a 24-well microplate
b. Add 5 drops of concentrated HCl to the well, one at a time, and record your observations
c. Discard the solution in the sink and rinse the tray with plenty of water.
2. Acid-Base Equilibrium
a. Add 25 drops of distilled water to one well of a 24-well microplate. Add 5 drops of bromothymol blue
b. Add 0.1M HCl one drop at a time until a color change occurs, then add 2 more drops. Mix by shaking
gently.
c. Now add 0.1M NaOH a drop at a time until a change occurs. Note the change
d. Now add 0.1M HCl a drop at a time until a color change occurs. Note the change.
e. Discard the solution in waste container and rinse the tray with water.
3. Pressure Effects on Gas Equilibrium
a. Remove plunger from syringe and attach cap. Add approx 15 ml of club soda, carefully reinsert the
syringe, invert the syringe, remove the cap, and depress the plunger to remove the air from the syringe.
Replace the cap.
b. Observe the bubbles in the syringe; they were not evident before the bottle was uncapped due to the
internal pressure that kept the equilibrium of the reaction:
H2CO3 (aq) H2O (l)+ CO2 (g)
favoring the right-hand side.
c. Increase the pressure in the syringe by pressing down on the plunger while resting the cap on the table
so it won’t pop off. Observe the change in the bubbles.
d. Decrease the pressure in the syringe by gently pulling on the plunger without totally removing it from
the syringe. Observe the change in the bubbles.
Eco-Friendly Equilibrium (Student)
46
4. Temperature Effects on Equilibrium
a. Obtain 2 sealed Beral pipettes containing 0.1M CoCl2 from your teacher. Observe the color of the
solution.
b. Place one of the pipettes into a hot water bath for 1-2 minutes. Remove and compare the color with
that of the unheated tube. Record your observations. Continue to observe the unheated tube until it
returns to room temperature.
c. Place one of the pipettes into a cold water/ice bath for 2-3 minutes. Remove and compare the color
with that of the uncooled tube. Record your observations. Continue to observe the unheated tube until
it returns to room temperature.
d. Return the 2 pipettes to your teacher.
5. Solubility and Complex Ion Equilibria
a. Add 25 drops of 0.1M Zn(NO3)2 to each of 3 wells in a 24-well microplate. Add 1 drop of 6M NaOH to
each well and mix well. Record your observations.
b. To the first well, add 6M HCl drop by drop with mixing. Observe changes and stop adding when no more
changes are seen.
c. To the second well add 6M NaOH drop by drop with mixing and proceed as in “b”.
d. To the third well add 6M NH3 (NH4OH)drop by drop with mixing and proceed as in “b”.
e. Discard the solutions in the waste container and rinse the tray with water.
f. You may repeat the entire sequence with 0.1M Mg(NO3)2 at your teacher’s discretion.
6. Reversible Reactions with Copper (II) Sulfate.
a. Add 25 drops of 0.1M CuSO4 to two wells in a 24-well microplate. One will be for reacting and the other
for comparison.
b. Add 6M NH3 drop by drop (with mixing between drops) until there is a definite color or appearance
change. Record your observations.
c. Continue to add NH3 until there is another color or appearance change and record your observations.
d. Now add 3M H2SO4 dropwise to the solution until the original color is restored.
e. Discard the solutions in the waste container and rinse the plate with water.
Eco-Friendly Equilibrium (Student)
47
Data Sheet
A. Dynamic Equilibrium
Test #1
Test #2
Trial #
Volume A
Volume B
Trial #
Start
50
0
Start
1
Volume A
Volume B
1
Draw other data tables as necessary for additional tests.
Eco-Friendly Equilibrium (Student)
48
B. Le Chatelier’s Principle
1. Solubility Equilibrium for NaCl
Observations:
2. Acid-Base Equilibrium
a. BTB + HCl
b. BTB + HCl + NaOH
c. BTB + HCl + NaOH + HCl
3. Pressure Effects on Gas Equilibrium
a. Observations under increased pressure:
b. Observations under decreased pressure:
4. Temperature Effects on Equilibrium
a. Observations at room temperature:
b. Observations when heated:
c. Observations when cooled:
Eco-Friendly Equilibrium (Student)
49
5. Solubility and Complex Ion Equilibria
a. Zn(NO3)2 + NaOH
b. Zn(NO3)2 + HCl
c. Zn(NO3)2 + NaOH
d. Zn(NO3)2 + NH3
6. Reversible Reactions with Copper (II) Sulfate
a. + NH3
b. + excess NH3
c. + H2SO4
Eco-Friendly Equilibrium (Student)
50
Analysis of Observations:
A. Dynamic Equilibrium
Graph the data from both sets of experiments on the same set of axes for comparison: Trial number on X
axis and Volume on Y axis.
Eco-Friendly Equilibrium (Student)
51
1. Describe the curves for Trial 1 based on your understanding of dynamic equilibrium in chemical reactions.
2. Describe the curves for Trial 2 based on your understanding of dynamic equilibrium in chemical reactions.
3. What difference in the number of transfers necessary to achieve equilibrium did you notice between trials 1
and 2. Explain.
B. Le Chatelier’s Principle
1. Solubility Equilibrium for NaCl
What evidence is there for a shift in equilibrium?
What caused the shift in equilibrium?
2. Acid-Base Equilibrium
Explain the effect observed in terms of acid-base equilibrium
Eco-Friendly Equilibrium (Student)
52
3. Pressure Effects on Gas Equilibrium
Explain the effect observed in terms of gas equilibrium.
4. Temperature Effects on Equilibrium
Co(H2O)62+ (aq) + 4Cl- (aq) ⇔ CoCl42- (aq) + 6H2O (l)
Pink
blue
Explain the changes observed in terms of the equilibrium expression given
5. Solubility and Complex Ion Equilibria
Zn(NO3)2 (aq) + 2NaOH (aq)  Zn(OH)2 (s) + 2NaNO3 (aq)
Write complete and net ionic equations for the above reaction
Explain what effect the addition of the following compounds had on this reaction:
a. HCl
b. excess NaOH
c. NH3
Eco-Friendly Equilibrium (Student)
53
6. Reversible Reactions with Copper (II) Sulfate
The copper ion in CuSO4 forms a complex with water in aqueous solution, resulting in the light blue color
characteristic of Copper solutions
Cu(H2O)42+ (aq) + 4NH3(aq) ⇔ Cu(OH)2 ⇔ [Cu(NH3)4}2+ (aq) + 4H2O
Light blue clear
cloudy blue
deep blue/purple clear
a. What was the evidence for the first shift in equilibrium when the NH3(aq) was added?
b. Explain how adding more NH3(aq) caused the equilibrium to shift again.
c. Explain how adding 3M H2SO4 caused the equilibria to shift back again
Eco-Friendly Equilibrium (Student)
54
The Colors of Le Chatelier
Adapted by Mary Haus from World of Chemistry by Houghton Mifflin Co.
This protocol was developed with California SB 70 funding in an effort to bring sustainability practices to
teaching laboratories and to increase awareness of sustainability in the next generation of the California
workforce. In other words, these protocols "green" our labs and "green" our students.
California Science Standards, Chemistry Grades 9-12
9. Chemical Equilibrium is a dynamic process at the molecular level. As a basis for understanding this concept:
a. Students know how to use Le Chatelier’s principle to predict the effect of changes in concentration,
temperature and pressure.
b. Students know equilibrium is established when forward and reverse reaction rates are equal.
c. * Students know how to write and calculate an equilibrium constant expression for a reaction.
Green Chemistry
The traditional labs for Le Chatelier's principle involve the use of larger quantities of toxic substances, such as
cobalt and AgNO3. Overall waste is approximately 30 mLs/group.
This "green" lab is a microscale version of the traditional lab: Experiment 75 Le Chatelier’s Principle from World
of Chemistry, Houghton Mifflin Co. This lab does not require hot plates, and uses minute drops of toxics,
creating only 5 mL of waste per group.
Green Principles Addressed
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Purpose
To observe how a system at equilibrium will shift to accommodate changing conditions and to use Le Chatelier’s
Principle to explain these changes.
Pre-Lab
Review your notes, text, and discussions. Answer the pre-lab questions in the worksheet that follows.
Materials
Supplies:
• 50 mL beaker
• Clear well plate or vials
• Plastic Pipet
• Glass Stir rod
• Sealed pipet with sample
• Hot water cup/beaker
• Ice water cup/beaker
Reagents:
• Cobalt chloride hexahydrate (CoCl2.6H2O)
• Silver nitrate solution (AgNO3(aq), 0.10 M)
• Calcium chloride pellets (CaCl2(s))
• 12 M Hydrochloric Acid (HCl(aq))
• Acetone
• Distilled water in dropper bottle
• Ethanol
55
Procedure
Safety
Safety goggles and lab aprons must be worn at all times. Due to some minor toxicity of certain chemicals and
corrosive properties of others, you must wash your hands thoroughly before leaving the laboratory.
1. Make sure your beaker and well plate are thoroughly dried. Pour 5 mLs of ethanol into your beaker.
2. Examine the solid Cobalt (II) chloride and record your observations about its color in the data section of your
lab sheet. Add a small amount of the solid (about the size of 5 grains of sand), into your beaker of ethanol
and stir with your stir rod. Record your observations on your lab sheet.
3. Using your plastic pipet, place 10 drops of the solution you just made into each of 5 wells in your plate.
4. Place 5 drops of distilled water into 4 of the same 5 wells, noting what happens as you add the drops.
Record your observations on your lab sheet. The last well remaining without water should be your control
5. Take your well plate to the fume hood and into one of the wells from step 3 CAREFULLY add 3-4 drops of the
hydrochloric acid. Record your observations on your lab sheet.
6. To a second well from step 3, add a small pellet of solid calcium chloride. Record your observations on your
lab sheet.
7. To the third well, add enough drops of acetone to see a change. Record your observations on your lab
sheet.
8. To your fourth well, add 3 drops of 0.1 M silver nitrate, AgNO3, one drop at a time. Record your
observations on your lab sheet.
9. Using a prepared sample in the sealed pipet, place it into the cup of hot water and hold it for a few
moments, (until a color change occurs). Remove it, and record your observations on your lab sheet.
10. Using the same sealed pipet, place it next into the cup of cold water for a few moments, (until a color
change occurs). Remove it, and record your observations on your lab sheet.
Cleaning Up
1. All solutions left in your beaker and in well plates should be dumped into the marked waste container, then
washed and dried thoroughly.
2. Return all materials where you found them and make sure everything is capped.
3. Wash your hands with soap and water before leaving the lab.
Analysis and Conclusions
Complete Analysis/Conclusion section in sheet below.
The Colors of Le Chatelier
56
The Colors of LeChatelier Lab Worksheet
Pre-Lab Questions:
.
1) The formula for solid Cobalt(II) chloride is CoCl2 6H2O. What name do we give to compounds which have
water molecules bound to them?
2) a. Write the equation for dissolving calcium chloride in water.
b. Use Le Chatelier’s Principle to predict the effect of adding solid calcium chloride to a solution
containing both of the cobalt complexes.
3) a. Write the equation for dissolving silver nitrate in water.
b. Write the equation for the precipitation reaction that you would expect when a solution containing
silver ions is added to a solution containing chloride ions.
Data/Observations
Initial appearance of the solid cobalt(II) chloride hexahydrate:
.
Color of solution produced by adding solid CoCl2 6H2O to alcohol:
Effect of adding water to alcoholic solution of cobalt(II) chloride:
The Colors of Le Chatelier
57
Describe the changes caused in each of the wells in steps 5-8:
Well 1 - Addition of 12 M HCl(aq) :
Well 2 – Addition of CaCl2(s) :
Well 3 – Addition of acetone, (CH3)2CO :
Well 4 – Addition of AgNO3(aq) :
Describe the effects of heating and cooling the prepared mixture from step 9:
Analysis and Conclusions
The net ionic equation for the equilibrium reaction you have been investigating is:
Co(H2O)62+(aq) + 4Cl-(aq) == CoCl42-(aq) + 6H2O(l)
(pink)
(blue)
1. a. Which cobalt complex was favored by addition of water to the solution of cobalt(II) chloride in
alcohol?
b. Use Le Chatelier’s Principle to explain the color change you observed.
2. a. Which cobalt complex was favored in both step 5 and step 6?
b. What ion is common to both of the reagents you used to bring about the color changes in these two
steps?
c. Use Le Chatelier’s Principle to explain why the color change occurs in each case.
3. Acetone absorbs water. Use this fact and Le Chatelier’s Principle to explain the color change that you
saw when you added the acetone to the third well in step 7.
The Colors of Le Chatelier
58
4. Silver chloride, AgCl, is a white solid. The equilibrium constant is K= 6 x 109 for:
Ag+(aq) + Cl-(aq) == AgCl(s)
a. At equilibrium, would you expect to have mostly silver and chloride ions in solution, or mostly solid
silver chloride? Explain.
b. What color was the solid you formed in step 8? What must it have been?
c. What color did the liquid in the vial turn? Which complex of cobalt was favored? Explain.
5. a. Which cobalt complex was favored by the addition of energy as heat? Which complex was favored by
cooling?
b. Rewrite the equation for the reaction, including the energy term in the equation. The value of delta
H for the process is +50kJ/mol.
c. Use Le Chatelier’s Principle and the equation from 5b to explain the color changes that resulted from
the heating and cooling.
The Colors of Le Chatelier
59
The Colors of LeChatelier Lab Worksheet--Teacher Prep Notes and Suggested ANSWER KEY
Teacher Prep Notes:
•
Cobalt chloride hexahydrate crystals dissolved in ethanol. Measure about 25-30 ml of ethanol and a
small sample of the solid cobalt chloride hexahydrate (about the volume of two drops of water). Use
student procedure to mix. Seal pipets by warming with fireplace lighter or bunsen burner. Crimp with
needle-nose pliers.
PreLab Questions:
.
1) The formula for solid Cobalt(II) chloride is CoCl2 6H2O. What name do we give to compounds which have
water molecules bound to them?
Hydrates
2) a. Write the equation for dissolving calcium chloride in water.
CaCl2(s)  Ca2+ (aq) + 2 Cl- (aq)
b. Use Le Chatelier’s Principle to predict the effect of the adding solid calcium
chloride to a solution containing both of the cobalt complexes.
Equilibrium would shift towards the CoCl4 side of the reaction
3) a. Write the equation for dissolving silver nitrate in water.
AgNO3(s)  Ag+ (aq) + NO3 - (aq)
b. Write the equation for the precipitation reaction that you would expect when a
solution containing silver ions is added to a solution containing chloride ions.
Ag+ (aq) + Cl- (aq)  AgCl(s)
Data/Observations:
Initial appearance of the solid cobalt(II) chloride hexahydrate:
Maroon solid crystal
.
Color of solution produced by adding solid CoCl2 6H2O to alcohol:
Bluish/clear
Effect of adding water to alcoholic solution of cobalt(II) chloride:
Turns pink
Describe the changes caused in each of the vials in step 6:
Vial 1 - Addition of 12 M HCl(aq) :
Turns blue
Vial 2 – Addition of CaCl2(s) :
Blue where the pellets are
Vial 3 – Addition of acetone, (CH3)2CO :
Blue at top where acetone sat
Vial 4 – Addition of AgNO3(aq) :
Pink liquid with white precipitate
Describe the effects of heating and cooling the prepared mixture from step 7:
Hot water – turns blue, cold water – turns pink
The Colors of Le Chatelier
60
Analysis and Conclusions:
The net ionic equation for the equilibrium reaction you have been investigating is:
Co(H2O)62+(aq) + 4Cl-(aq) == CoCl42-(aq) + 6H2O(l)
(pink)
(blue)
1. a. Which cobalt complex was favored by addition of water to the solution of
cobalt(II) chloride in alcohol?
Co(H2O)62+
b. Use Le Chatelier’s Principle to explain the color change you observed.
since water is a product, by adding more, LeChatelier would predict the
equilibrium shift would be towards the rectant side
2. a. Which cobalt complex was favored in both step 5 and step 6?
CoCl42b. What ion is common to both of the reagents you used to bring about the color
changes in these two steps?
Chloride ions
c. Use Le Chatelier’s Principle to explain why the color change occurs in each
case.
Since chlorine is a reactant, by adding more, LeChatelier would predict
the equilibrium shift would be towards the product side.
3. Acetone absorbs water. Use this fact and Le Chatelier’s Principle to explain the
color change that you saw when you added the acetone to the third well in step 7.
In the removal of water, a product, LeChatelier would predict the
equilibrium shift would be towards the product side to replace what was
lost.
4. Silver chloride, AgCl, is a white solid. The equilibrium constant is K= 6 x 109 for:
Ag+(aq) + Cl-(aq) == AgCl(s)
a. At equilibrium, would you expect to have mostly silver and chloride ions in
solution, or mostly solid silver chloride? Explain.
Mostly AgCl. The high K value indicates equilibrium lies strongly to the
right.
b. What color was the solid you formed in step 8? What must it have been?
White, AgCl
c. What color did the liquid in the vial turn? Which complex of cobalt was
favored? Explain.
Pink, Co(H2O)62+ , Chloride ions were being removed, therefore, equilibrium
shifted to replace them, or towards the reactant side.
5. a. Which cobalt complex was favored by the addition of energy as heat? Which
complex was favored by cooling?
Heating - CoCl42- , Cooling - Co(H2O)62+
b. Rewrite the equation for the reaction, including the energy term in the equation.
The value of delta H for the process is +50kJ/mol
50 kJ + Co(H2O)62+(aq) + 4Cl-(aq) == CoCl42-(aq) + 6H2O(l)
c. Use Le Chatelier’s Principle and the equation from 5b to explain the color
changes that resulted from the heating and cooling.
Heating – shifted equilibrium in the endothermic direction; Cooling – back towards
the exothermic direction
The Colors of Le Chatelier
61
Gas Laws
Developed by Lisa Selchau, Miramar College
California Science Standards, Chemistry Grades 9-12
4. The kinetic molecular theory describes the motion of atoms and molecules and explains the properties of
gases. As a basis for understanding this concept:
a. Students know the random motion of molecules and their collisions with a surface create the
observable pressure on that surface.
b. Students know the random motion of molecules explains the diffusion of gases.
c. Students know how to apply the gas laws to relations between the pressure, temperature, and
volume of any amount of an ideal gas or any mixture of ideal gases.
d. Students know the values and meanings of standard temperature and pressure (STP).
e. Students know how to convert between the Celsius and Kelvin temperature scales.
f. Students know there is no temperature lower than 0 Kelvin.
g. * Students know the kinetic theory of gases relates the absolute temperature of a gas to the average
kinetic energy of its molecules or atoms.
h. * Students know how to solve problems by using the ideal gas law in the form PV = nRT.
7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for
understanding this concept:
a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or
atoms).
b. Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal
energy.
8. Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules. As
a basis for understanding this concept:
a. Students know the rate of reaction is the decrease in concentration of reactants or the increase in
concentration of products with time.
b. Students know how reaction rates depend on such factors as concentration, temperature, and
pressure.
c. Students know the role a catalyst plays in increasing the reaction rate.
Green Chemistry
Oxygen gas can be collected from the decomposition of a variety of oxygen-containing compounds. Some of
these compounds are toxic or hazardous, including mercury (II) oxide (HgO), lead (IV) oxide (PbO2), and
potassium chlorate (KClO3). Oxygen can also be collected from less toxic chemicals, such as potassium nitrate
(KNO3), hydrogen peroxide (H2O2), and water (H2O). In this "green" lab, we will be decomposing non-toxic
hydrogen peroxide, rather than using the more toxic compounds previously described.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #3: Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methods should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #12: Inherently Safer Chemistry for Accident Prevention: Substances and the form of a substance
used in a chemical process should be chosen to minimize the potential for chemical accidents, including
releases, explosions, and fires.
62
Objectives
In this experiment, oxygen is produced by decomposing hydrogen peroxide (H2O2). Five flasks of oxygen gas will
be collected by the downward displacement of water. After collection, some of the physical and chemical
properties of oxygen will be observed.
Introduction
Oxygen is the most abundant and widespread of all the elements in the earth’s crust. It occurs both as free
oxygen gas and combined in compounds with other elements. Free oxygen gas is diatomic and has the formula
O2. Oxygen is found combined with more elements than any other single element, and it will combine with all
the elements except some of the noble gasses. Water is 88.9 percent oxygen by mass, and the atmosphere is
about 21 percent oxygen by volume. Oxygen is a colorless and odorless gas, and is only very slightly soluble in
water, a property important to its collection in this experiment.
Oxygen may be obtained by decomposing a variety of oxygen-containing compounds. Some of these are
mercury (II) oxide (HgO), lead (IV) oxide (PbO2), potassium chlorate (KClO3), potassium nitrate (KNO3), hydrogen
peroxide (H2O2), and water (H2O).
A. Decomposition of Hydrogen Peroxide to Generate Oxygen
Hydrogen peroxide decomposes very slowly at room temperature. The rate of decomposition is greatly
increased by adding a catalyst, manganese dioxide. A catalyst speeds up a reaction without getting used up in
the reaction. Manganese dioxide contains oxygen but it is not decomposed to produce oxygen in this
experiment. These equations represent the changes that occur in the generator flask.
MnO2
Word Equation: Hydrogen Peroxide
→
Water + Oxygen
MnO2
Formula Equation:
2H2O2(aq)
→
2H2O(l) + O2(g)
B. Collection of Oxygen
The oxygen is collected by a method known as the downward displacement of water. The gas is conducted from
a generator into a 250 mL Erlenmeyer flask filled with water and inverted into a water-filled plastic trough. The
first collection flask will contain some air mixed with oxygen, but the subsequent bottles will contain pure
oxygen gas. The air that was in the generator flask is pushed into the first collection flask. The oxygen, which is
only very slightly soluble in the water, rises into the flask and pushes the water down and out. Because oxygen
is heavier than air, a watch glass is used to cover the opening of the flask while it is inverted to a right-side-up
position and placed on the bench top until it is time to test it.
C. Properties of Oxygen
Like all kinds of matter, oxygen has both physical and chemical properties and you will observe both in this
experiment. One outstanding and important chemical property of oxygen is its ability of support combustion.
During combustion oxygen is consumed but does not burn and this ability to support combustion is one test for
oxygen. Other substances (a wooden splint or a candle, for example) burn in oxygen producing a visible flame
and heat. Compounds containing oxygen and one other element are known as oxides. Thus, when elements
such as sulfur, hydrogen, carbon, and magnesium burn in air or oxygen, they form sulfur dioxide, hydrogen oxide
Gas Laws
63
(water), carbon dioxide, and magnesium oxide, respectively. These chemical reactions may be represented by
equations, for example:
Word Equation:
Formula Equation:
Materials
Sulfur + Oxygen → Sulfur dioxide
S(s) + O2(g )
→
SO2(g)
Equipment and Supplies:
• Ring Stands
• Small o-ring clamps
• Wood splints
• Candles
• Deflagration spoon
• Plastic trough
• 20-25 cm length rubber tubing, 25 X 200 mm
• Ignition tube
• 5 250 mL Erlenmeyer flasks
• 5 Watch glasses
• 90° bend glass tubing
• 135°bend glass tubing
• 100 mL beaker containing sand (for candle)
Reagents (solids and solutions):
• Magnesium (Mg) strips
• manganese dioxide (MnO2)
• Fine steel wool (Fe)
• Sulfur (S)
• 9 percent hydrogen peroxide (H2O2)
Procedure
Safety Concerns
Wear protective goggles. If some hydrogen peroxide gets on your skin, it may burn. Wash off with soap and
water immediately.
A. and B. Generation and Collection of Oxygen from Hydrogen Peroxide.
1. Obtain the generator flask shown in figure 1.1. It consists of a 250 mL Erlenmeyer flask, two-hole stopper,
thistle tube, glass right-angle bend, glass delivery tube with 135 degree bend and a 20-25 cm length piece of
rubber tubing. The thistle tube should have about a 3 mm clearance between the end of the tube and the
bottom of the flask with the rubber stopper in place.
Figure 1.1 Generator Flask
Gas Laws
64
2. Fill a plastic trough with tap water until the water level is approximately two inches from the top.
Completely fill five 250 ml Erlenmeyer flasks with tap water by placing the mouth of each under the water
level in the trough. Invert each bottle and stabilize them each with an o-ring attached to a ring stand. You
want to have about one inch of space between the mouth of the flask and the bottom of the trough. You
will have to slide the o-ring under the water and up onto the flask, then attach the o-ring clamp to the ring
stand. See figure 1.2.
3. Using a spatula put a pea-sized quantity of manganese dioxide (MnO2) in the generator flask. Replace the
stopper, stabilize the flask on the ring stand with a clamp, and make sure that all glass-rubber connections
are tight. Slowly add 25 mL of DI water to the generator flask through the thistle tube. Make sure that the
end of the thistle tube is covered with water so that the oxygen does not escape out the thistle tube. Place
the glass tubing from the generator flask under the water and into the first collection flask.
4. Using a 50 mL graduated cylinder, measure out 45 mL of 9 percent hydrogen peroxide solution. To start the
generation of oxygen, slowly pour 15 mL of the peroxide solution into the thistle tube. The oxygen gas will
now be generated and bubble into the first Erlenmeyer flask. Once the flask is full of gas, the oxygen will
bubble outside of the flask. When the first flask fills with gas, move the glass tubing into the next flask of
water and so on until all five flasks are filled with gas. Continue generating oxygen by adding additional 15
mL portions of peroxide whenever the rate of gas production slows down markedly. You may not need all of
the 45 mLs of peroxide.
Figure 1.2 Preparation of oxygen by decomposition of hydrogen peroxide.
5. Cover the mouth of each gas-filled flask with a watch glass before removing it from the water. Store each
flask mouth upward without removing the watch glass. The oxygen will not readily escape since it is
slightly more dense than air. Remove all five flasks from the water and note which flask of gas was
collected first.
Gas Laws
65
6. Allow the reaction to go to completion while you continue with the testing of oxygen that you
collected. If you have any unreacted H2O2 remaining in your graduated cylinder, return it to the special
container in the waste hood labeled “unreacted 9% H2O2”.
C. Properties of oxygen
Each of the following tests (except C.6) is conducted with a flask of oxygen and, for comparison, with a flask
of air. Record your observations on the data pages.
1. The glowing splint test is often used to verify the identity of oxygen. Ignite a wood splint, blow out the
flame, and insert the still-glowing splint into the first flask of oxygen collected. Repeat with a flask of air. To
ensure having a flask of air, fill the flask with water and then empty it, thus washing out any other gasses
that may be present.
2. (Instructor Demo) Ask your instructor to help you when you are ready for this step. Take a small bottle of
sulfur, a deflagration spoon, a beaker of water, a flask of oxygen and a flask of air to the fume hood. Turn
the light in the fume hood off. Light a Bunsen burner in the fume hood and hold the deflagration spoon
filled with sulfur directly over the flame. First the sulfur gets dark and melts, then it begins to burn with a
blue flame that is barely visible. Once all the sulfur has melted and is burning, lower the burning sulfur
alternately into a flask of oxygen and a flask of air and compare the differing combustions. Quench the
excess burning sulfur in a beaker of water.
3. Obtain a 100 mL beaker containing sand and a candle from the cart. Lower a flask of oxygen over the
burning candle, placing the mouth of the flask slightly into the sand. Using the provided stopwatch,
measure and record the time, in seconds, that the candle continues to burn. Repeat with a flask of air (you
may need a new candle). Note also the difference in the brilliance of the candle flame in oxygen and in air.
Return the unused portion of the candle to the cart.
4. Invert a flask of oxygen, covered with a watch glass, and place it mouth to mouth over a flask of air. Remove
the watch glass from between the bottles and allow them to stand mouth to mouth for 3 minutes. (You may
need to hold them so they don’t fall over.) Cover each flask with a watch glass and set the flasks down,
mouths upward. Test the contents of each flask by inserting a still-glowing splint. Using the stopwatch,
time how long it takes for the splint to burn out in each.
5. Pour 25 mL of DI water into the fifth bottle of oxygen and replace the cover. Take it to a hood and turn the
hood light off. Place the flask of oxygen close to (within 5 or 6 cm) a lit Bunsen burner. Prepare a loose
(plenty of air spaces in-between the metal fibers), 3 cm in diameter wad of steel wool (iron). Hold it in the
crucible tongs and momentarily heat it in the burner flame until some of the steel wool first begins to glow
red hot. Immediately (while it is still red hot) lower the glowing metal into the flask of oxygen. It is essential
that the iron is still glowing red hot when you quickly place it in the flask. Repeat using a flask of air.
6. A small strip of magnesium ribbon will be burned next. Do not put burning magnesium into a bottle of
oxygen. There is enough oxygen in the air for this reaction to proceed vigorously.
∗Do not look directly at the burning magnesium ribbon. It is very bright and the light includes considerable
ultraviolet light, which can cause damage to the retina of the eye.
Take a 2 to 5 cm strip of magnesium metal in some crucible tongs and ignite it by heating it in the Bunsen
burner flame. Have ready your white evaporating dish on the desktop to drop the burnt magnesium strip
into. Compare the look of the product to the magnesium strip that you started with.
Gas Laws
66
Name:
Data
A and B. Generation and Collection of Oxygen
1. What evidence did you observe that oxygen is not very soluble in water?
2. What is the source of oxygen in this procedure?
Name
Formula
3. What purpose does the manganese dioxide serve in this preparation of oxygen?
4. What gas was in the apparatus before you started generating oxygen? Where did it go?
5. What is different about the composition of the first bottle of gas collected compared to the other four?
6. Why are the bottles of oxygen stored with the mouth up?
7. a) What is the symbol for the element oxygen?
b) What is the formula for oxygen gas?
Gas Laws
67
8. Write the word and formula equations for the preparation of oxygen from hydrogen peroxide.
Word Equation:
Formula Equation:
9. What substances, other than oxygen, are in the generator flask when the decomposition of H2O2 is complete?
C. Properties of Oxygen
1. Write word equations for the chemical reactions that occurred.
C.1. Combustion of wood. Assume carbon is the combustible material.
C.2. Combustion of sulfur.
C.5. Combustion of steel wool (iron). (Call the product iron oxide.)
C. 6. Combustion of magnesium.
2. Write the balanced formula equations for these same combustion reactions.
C.1. (CO2 is the formula for the oxide of carbon that is formed.)
C.2. (SO2 is the formula for the oxide of sulfur that is formed.)
C.5. (Fe2O3 is the formula for the oxide of iron that is formed.)
C.6. By looking at the periodic table you should be able to predict the formula for magnesium oxide.
Gas Laws
68
3. Combustion of a candle.
a) Number of seconds that the candle burned in the flask of oxygen
b) Number of seconds that the candle burned in the flask of air
c) Explain this difference in combustion time.
d) Is it scientifically sound to conclude that all the oxygen in the flask was reacted when the candle
stopped burning? Explain your answer.
4. What were the results of the experiment in which a flask of oxygen was placed over a flask of air? Explain
your results.
5. a) Describe the material that is formed when magnesium is burned in air.
b) What elements are in this product?
Gas Laws
69
6. a)
What is your conclusion about the rate or speed of a chemical reaction with respect to the concentration
of the reactants – for example, a combustion in a high concentration of oxygen (pure oxygen) compared
to a combustion in a low concentration of oxygen (air)?
b) What evidence did you observe in the burning of sulfur to confirm your conclusion in question 6a?
Gas Laws
70
Post-Lab Questions
1. The pressure of hydrogen gas in a constant-volume cylinder is 4.25 atm at 10.0°C. What will the pressure be
if the temperature is raised to 80.0oC?
2. A sample of Nitrogen gas, N2, occupies 3.0 L at a pressure of 3.0 atm. What volume will it occupy when the
pressure is changed to 0.50 atm and the temperature remains constant?
3. A sample of methane gas, CH4, occupies 4.50 L at a temperature of 20.0oC. If the pressure is held constant,
what will be the volume of gas at 100.0oC?
4. A 325 mL sample of air is at 720. torr and 30.0oC. What volume will this gas occupy at 800.torr and 75.oC?
5. A quantity of oxygen occupies a volume of 19.2 L at STP (0oC, 1atm). How many moles of oxygen are present?
6. What volume would 10.5 g of nitrogen gas, N2, occupy at 200.0 K and 2.02 atm?
Gas Laws
71
Name:
Pre-Lab Questions
1. Why is it important to make sure that the end of the thistle tube is covered with water in step A.3 of the
procedure?
2. In what way is the composition of the first bottle of oxygen gas that you collect going to differ from the other
four bottles of oxygen gas that you collect?
3. Oxygen is collected over water at a temperature of 25°C and 672 torr. What is the partial pressure of the
oxygen if the vapor pressure of water is 24 torr at 25°C?
4. A given amount of oxygen gas is in a closed, expandable container. Assuming that the oxygen acts as an
ideal gas, answer the following questions.
a. What will happen to the pressure of the gas if the volume is tripled and the temperature remains
constant?
b. What will happen to the volume if the temperature doubles and the pressure remains constant?
c. What will happen to the pressure if the volume is decreased by half and the temperature is doubled?
d. What will have to happen to the volume if the temperature triples, and the pressure remains
constant?
Gas Laws
72
Reaction of Magnesium and Water in the Presence of a Catalyst
(Flameless Ration Heater)
Developed by Russ Davidson, San Dieguito Academy High School
This protocol was developed with California SB 70 funding in an effort to bring sustainability practices to teaching
laboratories and to increase awareness of sustainability in the next generation of the California workforce. In
other words, these protocols "green" our labs and "green" our students.
California Science Standards, Chemistry Grades 9-12
7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for
understanding this concept:
a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or
atoms).
b. Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal
energy.
c. Students know how to solve problems involving heat flow and temperature changes, using known
values of specific heat and latent heat of phase change.
8. Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules. As
a basis for understanding this concept:
a. Students know the rate of reaction is the decrease in concentration of reactants or the increase in
concentration of products with time.
b. Students know how reaction rates depend on such factors as concentration, temperature, and
pressure.
Green Chemistry
A traditionally used exothermic reaction experiment involves the use of Calcium Oxide (quicklime) and water,
requiring contact with a strong base and the disposal concerns associated with CaO.
This "green" exothermic reaction protocol involves the use of non-hazardous materials: magnesium, iron, and
sodium chloride.
Green Chemistry Principles
Principle #3 - Less Hazardous Chemical Syntheses: Wherever practicable, synthetic methodologies should be
designed to use and generate substances that possess little or no toxicity to human health and the
environment.
Principle #4 - Designing Safer Chemicals: Chemical products should be designed to preserve efficacy of function
while reducing toxicity.
Principle #7 - Use of Renewable Feedstocks: A raw material or feedstock should be renewable rather than
depleting wherever technically and economically practicable.
Principle #10 - Design for Degradation: Chemical products should be designed so that at the end of their
function they do not persist in the environment and break down into innocuous degradation products.
Purpose
To investigate an exothermic reaction and to discover the effect of surface area (due to particle size) on the rate
of a reaction.
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Background
A flameless ration heater, or FRH, is a water-activated exothermic chemical heater included with "Meal, Readyto-Eat" (MREs) used to heat the food. U.S. military specifications for the heater require that it be capable of
raising the temperature of an eight ounce entrée by 100°C in twelve minutes, and that it display no visible
flame.
Magnesium heaters are in use with canned foods in Japan, Heater Meals, the Mountain House’s Mountain Oven,
Hot Pack rations for NASA, and U.S. military and paramilitary MREs. They provide a simple and safe method for
heating food, with little risk of fire or explosion.
The most popular design uses heating pads made from a super-corrodible magnesium/iron alloy in a porous
matrix formed from polymeric powders with sodium chloride incorporated with it, or in a separate tablet. To
heat, the contents are mixed with water which dissolved the sodium chloride into an electrolyte solution causing
magnesium and iron to function as an anode and cathode, respectively. An exothermic oxidation-reduction
reaction occurs between the magnesium-iron alloy and water to produce magnesium hydroxide, hydrogen gas,
and heat.
Mg + 2 H2O  Mg(OH)2 + H2 + heat
To ensure that reaction does not prematurely occur while in storage, the heating pad must be kept in a
watertight container protected from liquid water and moist air. If the heater matrix is exposed to moist air at
any time, its surface may become coated with a thin layer of insoluble magnesium hydroxide, which will hinder
future chemical reactions. This greatly reduces the effectiveness of the heater. Heater pads that become wet
in storage release hydrogen gas and can create a serious fire or explosion hazard that may ignite hot burning
magnesium. Refer to the following website for a graphical description of MREs:
http://www.mreinfo.com/us/mre/frh.html
In this experiment, you will use the raw materials that work in a flameless ration heater: magnesium, iron,
sodium chloride and water. During the course of this experiment, the physical size of the magnesium reagent
will be altered and finally “alloyed’ with the iron reagent to witness the overall effect of surface area on the
reaction rate and heat evolved.
Materials
Equipment:
• Laboratory balance
Supplies:
• 3 Test tubes
• Test tube rack
• 1-5mL calibrated transfer pipet
• Mortar and pestle
• Thermometer
• 10 mL graduated cylinder
• Small-mouth funnel
Reagents:
• Iron powder (100 mesh)
• Sodium chloride
• Magnesium powder (100 mesh)
• Magnesium turnings
• Mg/Fe alloy powder from commercial MRE
heater
Safety
Caution: Always wear safety goggles (all steps).
Caution: Avoid agitating and kicking up excessive volumes of iron or magnesium dust. They may become
an inhalation hazard
Caution: Finely divided metallic powders represent a fire hazard when exposed to sparks or open flame.
Do not have any sources of ignition near the reactants in this experiment.
Flameless Ration Heater
74
Caution: In the case of an accidental ignition of magnesium powder, water, and carbon dioxide, fire
extinguishers will not work and will actually increase the rate of combustion of the powdered magnesium.
To extinguish any magnesium powder that may have ignited accidentally have a container of sand or clay
powder on hand. Pour the sand over the magnesium to deprive it of oxygen and extinguish the fire.
Caution: Hydrogen gas will be liberated during the course of the experiment. Hydrogen is an explosive gas
when exposed to open flames or sparks. Do not have any sources of ignition near the products in this
experiment.
Procedure
Part 1
1.
2.
3.
4.
5.
6.
7.
8.
Label four test tubes 1-4.
Mass 0.10 g of magnesium powder and add to test tube #1. Record mass of Mg.
Mass 0.10 g of magnesium turnings and add to test tube #2. Record mass of Mg.
Mass 0.05 g of iron powder and add to test tube #1. Repeat for test tube #2.
Mass 0.10 g of NaCl and add to test tube #1. Repeat for test tube #2.
Fill a small beaker with 10-20 mL of distilled water. Measure and record the temperature of the water.
Using a graduated cylinder or transfer pipette, add 2 mL of water to test tube #1.
Measure the temperature of the reaction with the thermometer. Record the highest temperature
reached as well as any observations about the rate of the reaction.
9. Set tube 1 aside, wash off the thermometer and repeat steps 7-8 using the magnesium turnings.
Part 2
10. Mass 0.10 g of magnesium powder and add to a clean, dry mortar. Record mass of Mg.
11. Mass 0.05 g of iron powder and add to the Mg powder in the mortar.
12. Using the pestle, grind the two powders until a uniform light brown coating is visible on the magnesium
powder. This observation is an approximation showing that relatively even distribution of the reagents
has been achieved.
Note: Grinding two powders in a mortar and pestle is often hazardous due to the heat released from
friction. This heat, in many cases, can be enough to initiate a chemical reaction at an unwanted time.
In this experiment, two metal powders are being ground together. Under normal conditions, metal
powders should not react from this process. Do not attempt this process with other powdered
reagents as an extreme safety hazard may occur.
13. Add the contents of the mortar to tube #3.
14. Mass 0.10 g of NaCl and add to test tube #3.
15. Using a graduated cylinder or transfer pipette, add 2 mL of water to test tube #3.
16. Wait for the reaction to initiate, subside and then measure and record the temperature of the reaction.
Part 3
17.
18.
19.
20.
Mass 0.15 g of MRE alloy powder and add powder to test tube #4.
Mass 0.10 g of NaCl and add to test tube #4.
Using a graduated cylinder or transfer pipette, add 2 mL of water to test tube #4.
Wait for the reaction to initiate, subside and then, using a thermometer, record the temperature of the
reaction.
Flameless Ration Heater
75
Column and Thin-Layer Chromatography
Developed by Matt Buller, Miramar College
California Science Standard, Chemistry Grades 9-12
6. Solutions are homogeneous mixtures of two or more substances. As a basis for understanding this concept:
a. Students know the definitions of solute and solvent.
f. *Students know how molecules in a solution are separated or purified by the methods of
chromatography and distillation.
Green Chemistry
The Column and TLC Lab replace a traditional lab that involved extracting spinach with relatively large amounts
of solvent. Instead of doing two extractions with 20 mL each of ethanol and dichloromethane, no extraction is
performed for the new experiment. The only solvents used now are small amounts of ethyl acetate and hexanes
(<10 mL total) for the column and TLC, which replaces the use of dichloromethane for the column. In this lab,
not only is solvent use reduced, chlorinated solvents are avoided completely.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #5: Safer Solvents and Auxiliaries: The use of auxiliary substances (e.g., solvents, separation agents,
etc.) should be made unnecessary wherever possible and innocuous when used.
Principle #7: Use of Renewable Feedstocks: A raw material or feedstock should be renewable rather than
depleting whenever technically and economically practicable.
Purpose
To purify and analyze a mixture of 1-octanol and octanal using column and thin-layer chromatography.
Background
Column chromatography is a purification technique that is often used when a mixture cannot be sufficiently
separated by extraction and/or recrystallization, both of which have significant drawbacks. Extraction is most
useful when one compound in a mixture can be made ionic through acid/base chemistry, while recrystallization
can only be performed when the desired compound is a solid. Chromatography is not limited by either of these
cases, and instead only requires a moderate (or even small) difference in polarity between compounds in order
to have them distribute differently between a solid stationary phase and a liquid mobile phase.
In “normal-phase” column chromatography, the stationary phase is a highly polar, finely ground solid, such as
silica gel (SiO2) or alumina (Al2O3), which is packed into a glass column. The mixture to be separated is adsorbed
onto the top of the stationary phase, and then the mobile phase, usually an organic solvent such as hexane or
ethyl acetate, is added and filters down through the column, pulling the mixture along with it. Compounds of
differing polarity “stick” to the stationary phase to varying degrees, and the result is that they flow through the
column at different rates. The liquid that empties from the column is collected in test tubes, and the contents
are then analyzed using a related technique: thin-layer chromatography.
Thin-layer chromatography (TLC) uses the same principles as column chromatography, but for TLC, a thin layer
of the stationary phase is adsorbed onto a rectangular plate and the mobile phase creeps up the plate through
capillary action. Because only a small amount of compound can placed on the plate, this technique is only used
for analysis. In this experiment, you will purify and analyze a 50/50 (v/v) mixture of 1-octanol and octanal.
76
Tip: It is a good idea to watch videos on these!
Materials
Equipment/Supplies:
• Column Chromatography – [O-21]
• Thin-Layer Chromatography – [O-22]
• 2 micropipets (set to 100 μL)
• Cotton
• Pasteur pipets (5 ¾”)
• Pipet bulbs
• Pipet clamps
• TLC plates
• 120 ml TLC jars
Equipment/Supplies (cont'd):
• Filter paper for TLC jars
• Rulers
• Iodine chambers
Reagents:
• 1:1 octanal/1-octanol (2 x 15 mL)
• 9:1 hexanes/ethyl acetate
• 4:1 hexanes/ethyl acetate
• Silica gel (60 Mesh)
• Sand
Procedure
1. Use a micropipet to place 100 μL of the 1-octanol/octanal mixture in a test tube
2. Obtain the materials needed to make a Pasteur-pipet column (see the following pages for instructions).
Make a column with enough dry silica gel so that the stationary phase is about 5 cm in length.
3. After adding sand to the top, add enough of the eluent, 9:1 hexanes/ethyl acetate, to moisten the entire
length of the column without having extra solvent on top.
4. Dissolve your mixture in a few drops of the eluent, transfer it to the column using a pipet, and allow it to
settle until the level of liquid is at the top of the silica gel. Rinse the sides of the column down with a few
more drops of eluent and allow it to settle to the same level.
5. Elute the column using enough 9:1 hexanes/ethyl acetate to collect fifteen test tubes with 0.5 mL of eluent
in each. Be careful to never let your column run dry!
You are now ready to analyze the fifteen “fractions” to determine how well the mixture was separated.
6. Obtain three TLC plates and mark them for five fractions each.
7. Using a capillary tube, spot a tiny amount of liquid from each fraction onto the correct marks, and then
develop the TLC plates in a beaker with 4:1 hexanes/ethyl acetate as the eluent. Make sure to cover your
beaker with a watchglass.
8. When the solvent line is about 0.5 cm from the top of the plate, remove the plate and mark where the
solvent line is. Allow the plates to dry and then visualize them under a UV lamp and then in an iodine
chamber. Use additional plate to co-spot with knowns.
9. Combine the test tubes that contain pure 1-octane into one test tube, and combine the test tubes that
contain pure octanal into another test tube.
10. Evaporate off the solvent (BP’s around 70oC) in a warm water bath, and obtain IR of the two compounds for
identification.
Column and TL Chromatography
77
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Fruit Basket of Esters
79
Fruit Basket of Esters
80
Fruit Basket of Esters
81
Fruit Basket of Esters
82
Fruit Basket of Esters - Synthesis of Esters
Modified by Ruth Trejo & Lina Farah, Miramar College
Originally Developed by The U.S. Naval Academy
California Science Standard, Chemistry Grades 9-12
8. Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules. As
a basis for understanding this concept:
a. Students know the rate of reaction is the decrease in concentration of reactants or the increase in
concentration of products with time.
b. Students know how reaction rates depend on such factors as concentration, temperature, and
pressure.
c. Students know the role a catalyst plays in increasing the reaction rate.
d. * Students know the definition and role of activation energy in a chemical reaction.
10. The bonding characteristics of carbon allow the formation of many different organic molecules of varied
sizes, shapes, and chemical properties and provide the biochemical basis of life. As a basis for understanding
this concept:
a. Students know large molecules (polymers), such as proteins, nucleic acids, and starch, are formed by
repetitive combinations of simple subunits.
b. Students know the bonding characteristics of carbon that result in the formation of a large variety of
structures ranging from simple hydrocarbons to complex polymers and biological molecules.
c. Students know amino acids are the building blocks of proteins.
d. * Students know the system for naming the ten simplest linear hydrocarbons and isomers that
contain single bonds, simple hydrocarbons with double and triple bonds, and simple molecules that
contain a benzene ring.
e. * Students know how to identify the functional groups that form the basis of alcohols, ketones,
ethers, amines, esters, aldehydes, and organic acids.
f. * Students know the R-group structure of amino acids and know how they combine to form the
polypeptide backbone structure of proteins.
Green Chemistry
This ester lab replaces the "GlueP" lab where students make slime and apply scientific method to change the
relative amounts of chemicals to make bouncy balls. The GlueP lab required a large quantity of chemicals. This
"green" lab replacement generates approximately only 2.5 mLs of waste per group, while the GlueP lab
generated about 10 times that amount of waste as a result of being a scientific method experiment.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Purpose
In this experiment you will produce esters that smell like different fruits and flowers. Additionally, you will
identify the structure of the ester, which produced the specific scent.
Fruit Basket of Esters
83
Acidic Esterification Reaction
A carboxylic acid and an alcohol are combined under acidic conditions to form an ester and water. The following
is the general esterification reaction, under acidic conditions:
O
O
+
R
OH + HO
H catalyst
R1
R1
R
O
H 2O
+
Pre-Laboratory Assignment
1. Draw the skeletal structure of these carboxylic acids, which will be used in this experiment.
Formic Acid
Acetic Acid
Propanoic Acid
Salicylic Acid
2. Draw the skeletal structure of these alcohols, which will be used in this experiment.
Methanol
1-Propanol
Isopentanol
1-Octanol
Benzyl Alcohol
3. Name the ester made from the following carboxylic acid and alcohol combinations.
a) Methanol and Formic Acid
_________________________________________
b) Ethanol and Acetic Acid
_________________________________________
Fruit Basket of Esters
84
c) 1-Propanol and Benzoic Acid
_________________________________________
d) Benzyl alcohol and Salicylic Acid
_________________________________________
4. Using skeletal structures, write a balanced chemical equation for the reaction of 2-methylbutanoic acid and
2-propanol (isopropanol) with an acid catalyst.
5. Name and describe the procedure for safely detecting the odor of chemicals.
6. What is a hazard of concentrated sulfuric acid?
Materials
Equipment/Supplies:
• Stopper (000)
• Thermometer
• Boiling chips
• Disposable pipettes
• Waste container (main; used stopper, boiling
chips)
Fruit Basket of Esters
Reagents:
• Methanol (CH3OH) (1ml/group)
• 1-Propanol (C3H7OH) (1ml/group)
• Isopentyl alcohol (C5H11OH) (1ml/group)
• 1-Octanol (C8H17OH) (1ml/group)
• Benzyl alcohol (C6H5CH2OH) (1ml/group)
• Formic acid (HCOOH) (0.5 ml/group)
• Acetic acid (CH3COOH) (0.5 ml/group)
• Propionic acid (C2H5COOH) (0.5m l/group)
• Salicylic acid (C6H4(OH)COOH) (0.2 g/group)
• Conc. Sulfuric acid (H2SO4) (5 drops/group)
85
Procedure
Caution:
Wear your goggles at all times.
Glacial acetic acid and concentrated sulfuric acid are corrosive!
1. Assemble a hot water bath by half filling a large beaker with tap water, adding 2-3
boiling chips and heating it on a hot plate. Maintain a temperature of 80–85 oC.
Refer to figure on the right.
2. In a test tube measure either 0.2 g (solid) or 10 drops (liquid) of your assigned
carboxylic acid, 20 drops of your assigned alcohol, and 5 drops of concentrated
sulfuric acid. Swirl to mix well. Place in the hot water bath for 30 minutes. Continue
to mix with a stirring rod until the solution is homogeneous (no solid remains).
Refer to figure on the right.
3. Remove test tube from the water bath. Pour ester into a small beaker containing
about 2 mL of distilled water. Cautiously smell the ester.
4. After all twelve esters have been synthesized by the class, cautiously smell each one
and record detailed observations in the data sheet. Record the appearance of the
ester.
5. Dispose of your ester in the hazardous waste bucket, and clean your glassware.
Experimental Data and Observations
a. Using skeletal structures, write the balanced equation of the ester you synthesized:
b. Write the structure of the ester that is synthesized between the following carboxylic acid and alcohol
combinations.
c. Name the ester.
d. Describe its odor.
Formic Acid & Benzyl Alcohol
Formic Acid & Isopentanol
Name of Ester:
Name of Ester:
Odor:
Odor:
Fruit Basket of Esters
86
Formic Acid & 1-octanol
Acetic Acid & Benzyl Alcohol
Name of Ester:
Name of Ester:
Odor:
Odor:
Acetic Acid & Isopentanol
Acetic Acid & 1-octanol
Name of Ester:
Name of Ester:
Odor:
Odor:
Propanoic Acid & 1-octanol
Propanoic Acid & Benzyl Alcohol
Name of Ester:
Name of Ester:
Odor:
Odor:
Fruit Basket of Esters
87
Propanoic Acid & Isopentanol
Propanoic Acid & 1-propanol
Name of Ester:
Name of Ester:
Odor:
Odor:
Salicylic Acid & Methanol
Salicylic Acid & 1-propanol
Name of Ester:
Name of Ester:
Odor:
Odor:
Fruit Basket of Esters
88
Synthesis of Esters Materials List
Equipment Needed:
Amount
Stopper (000)
01 box
Thermometer
12
Boiling chips
02 jars
Disposable pipettes
01 box
Waste container (main; used stopper, boiling chips)
03
Chemicals needed:
Methanol (CH3OH) (1ml/ea group)
02bottles
1-Propanol (C3H7OH) (1ml/ea group)
02 bottles
Isopentyl alcohol (C5H11OH) (1ml/ea group)
02 bottles
1-Octanol (C8H17OH) (1ml/ea group)
02 bottles
Benzyl alcohol (C6H5CH2OH) (1ml/ea group)
02 bottles
Formic acid (HCOOH) (0.5ml/ea group)
02 bottles
Acetic acid (CH3COOH) (0.5ml/ea group)
02 bottles
Propionic acid (C2H5COOH) (0.5ml/ea group)
02 bottles
Salicylic acid (C6H4(OH)COOH) (0.2g/ea group)
02 jars
Conc. Sulfuric acid (H2SO4) (5 drops/ea)
02 bottles
Calculations are based upon a class work in 12 groups.
Fruit Basket of Esters
89
Analysis of Fats in the Diet and Isolation of Fat and Olestra in Food Items
From Exploring Chemistry by Julie Peller, Contributed by Lina Farah, Miramar College
California Chemistry Standards, Grades 9-12
7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for
understanding this concept:
c. Students know energy is released when a material condenses or freezes and is absorbed when a
material evaporates or melts.
10. The bonding characteristics of carbon allow the formation of many different organic molecules of varied
sizes, shapes, and chemical properties and provide the biochemical basis of life. As a basis for understanding
this concept:
a. Students know large molecules (polymers), such as proteins, nucleic acids, and starch, are formed by
repetitive combinations of simple subunits.
b. Students know the bonding characteristics of carbon that result in the formation of a large variety of
structures ranging from simple hydrocarbons to complex polymers and biological molecules.
e. * Students know how to identify the functional groups that form the basis of alcohols, ketones,
ethers, amines, esters, aldehydes, and organic acids.
Green Chemistry
This "green" lab involves the use of non-hazardous/toxic substances with the exception of a small volume of
diethyl ether (6 mL/group) used to extract fat and olestra from potato chips.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Principle #5: Safer Solvents and Auxiliaries: The use of auxiliary substances (e.g., solvents, separation agents,
etc.) should be made unnecessary wherever possible and innocuous when used.
Principle #12: Inherently Safer Chemistry for Accident Prevention: Substances and the form of a substance
used in a chemical process should be chosen to minimize the potential for chemical accidents, including
releases, explosions, and fires.
90
91
92
93
94
95
96
97
98
99
100
101
Classification and Recycling of Plastics
From Exploring Chemistry by Julie Peller, Contributed by Lina Farah, Miramar College
California Chemistry Standards, Grades 9-12
6. Solutions are homogeneous mixtures of two or more substances. As a basis for understanding this concept:
a. Students know the definitions of solute and solvent.
b. Students know how to describe the dissolving process at the molecular level by using the concept of
random molecular motion.
c. Students know temperature, pressure, and surface area affect the dissolving process.
7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for
understanding this concept:
c. Students know energy is released when a material condenses or freezes and is absorbed when a
material evaporates or melts.
10. The bonding characteristics of carbon allow the formation of many different organic molecules of varied
sizes, shapes, and chemical properties and provide the biochemical basis of life. As a basis for understanding
this concept:
a. Students know large molecules (polymers), such as proteins, nucleic acids, and starch, are formed by
repetitive combinations of simple subunits.
b. Students know the bonding characteristics of carbon that result in the formation of a large variety of
structures ranging from simple hydrocarbons to complex polymers and biological molecules.
d. * Students know the system for naming the ten simplest linear hydrocarbons and isomers that
contain single bonds, simple hydrocarbons with double and triple bonds, and simple molecules that
contain a benzene ring.
e. * Students know how to identify the functional groups that form the basis of alcohols, ketones,
ethers, amines, esters, aldehydes, and organic acids.
Green Chemistry
Density determination of plastics is determined using water and small volumes of methanol, and considered to
be relatively non-hazardous. Small volumes of toluene and acetone are used to determine the solubility of
plastics in the lab. Acetone is considered to be only slightly toxic. Toluene is considered toxic, but is used in
small amounts and replaces the more toxic benzene solvent.
Green Chemistry Principles
Principle #1: Prevention: It is better to prevent waste than to treat or clean up waste after it has been created.
Principle #2: Atom Economy: Synthetic methods should be designed to maximize the incorporation of all
materials used in the process into the final product.
Principle #5: Safer Solvents and Auxiliaries: The use of auxiliary substances (e.g., solvents, separation agents,
etc.) should be made unnecessary wherever possible and innocuous when used.
Principle #12: Inherently Safer Chemistry for Accident Prevention: Substances and the form of a substance
used in a chemical process should be chosen to minimize the potential for chemical accidents, including
releases, explosions, and fires.
102
103
104
105
106
107
108
109
110
111
112
113
Green Chemistry Resources
The list below supplies a wealth of online information pertaining to Green Chemistry. Following this page of
resources is a set of Green Chemistry articles that provide a good introduction to familiarize yourself and your
students with Green Chemistry.
1. Beyond Benign website created by Dr. John Warner, one of the founders of Green Chemistry and co-creator
of Green Chemistry Principles. Beyond Benign provides Green Chemistry curriculum, outreach programs,
and educator training.
The following is a link to the High School Chemistry Curriculum. Click on ‘curriculum mapping PDF’ for a
great introduction to approaching Green Chemistry in the high school classroom. Be sure to watch the John
Warner introductory video.
http://www.beyondbenign.org/K12education/highschool.html
2. Warner Babcock Institute for Green Chemistry. Provides professional green chemistry training and
collaborative research for public and private sectors.
http://www.warnerbabcock.com/
3. Environmental Protection Agency Green Chemistry Site. Provides educational resources, partnerships,
grants, and fellowships.
Homesite: http://www.epa.gov/greenchemistry/
EPA YouTube Video (6.5 min) http://www.youtube.com/watch?v=rIE4T2HLW7c
4. American Chemical Society Green Chemistry Institute. Provides educational resources, Green Chemistry
summer school, links to books and other resources.
http://portal.acs.org/portal/acs/corg/content?_nfpb=true&_pageLabel=PP_TRANSITIONMAIN&node_id=83
0&use_sec=false&sec_url_var=region1&__uuid=1e551db0-801b-4778-be0f-1f924f03aa3a
5. Green Chemistry Network. UK-based organization. Provides education and training for public and private
sectors. Provides Green Chemistry activities for educators:
http://www.greenchemistrynetwork.org/education.htm
6. Green Chemistry Resource List. Provides links to many useful educational resources, including some sites
listed above. Generated by beyondbenign.org.
http://greenchemistrycommitment.org/resources.html
7. Green Chemistry Certificate Program (example)
UC Berkley Green Chemistry Certificate Program: http://extension.berkeley.edu/cert/greenchem.html
114
Green Chemistry: Scientists Devise New
"Benign by Design" Drugs, Paints,
Pesticides and More
Chemists are usually asked to invent a solution, but without considering hazardous
by-products. Green chemists now are doing both with success, but will it take
regulations to enforce the approach broadly?
By Emily Laber-Warren | May 28, 2010 Scientific American http://www.scientificamerican.com/article.cfm?id=greenchemistry-benign-by-design&SID=mail&sc=emailfriend
PROCESS OF ELIMINATION: Green chemists use all the tools and
training of traditional chemistry to create industrial processes that avert
hazard problems--called "benign by design". Image:
ISTOCKPHOTO/Sandralise
Back in the days when better living through chemistry was a
promise, not a bitter irony, nylon stockings replaced silk,
refrigerators edged out iceboxes, and Americans became
increasingly dependent on man-made materials. Today
nearly everything we touch—clothing, furniture, carpeting,
cabinets, lightbulbs, paper, toothpaste, baby teethers,
iPhones, you name it—is synthetic. The harmful side effects of industrialization—smoggy air,
Superfund sites, mercury-tainted fish, and on and on—have often seemed a necessary tradeoff.
But in the early 1990s a small group of scientists began to think differently. Why, they asked, do we rely
on hazardous substances for so many manufacturing processes? After all, chemical reactions happen
continuously in nature, thousands of them within our own bodies, without any nasty by-products.
Maybe, these scientists concluded, the problem was that chemists are not trained to think about the
impacts of their inventions. Perhaps chemistry was toxic simply because no one had tried to make it
otherwise. They called this new philosophy "green chemistry."
Green chemists use all the tools and training of traditional chemistry, but instead of ending up with
toxins that must be treated and contained after the fact, they aim to create industrial processes that avert
hazard problems altogether. The catch phrase is "benign by design".
Progress without pollution may sound utterly unrealistic, but businesses are putting green chemistry into
practice. Buying, storing, and disposing of hazardous chemicals is expensive, so using safer alternatives
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makes sense. Big corporations—Monsanto, Dow, Merck, Pfizer, DuPont—along with scrappy start-ups
are already applying green chemistry techniques. There have been hundreds of innovations, from safer
latex paints, household cleaning products and Saran Wrap to textiles made from cornstarch, and
pesticides that work selectively, by disrupting the life cycles of troublesome insects. Investigators have
also developed cleaner ways of decaffeinating coffee, dry-cleaning clothes, making Styrofoam egg
cartons, and producing drugs like Advil, Zoloft and Lipitor.
Over the past 15 years, green chemistry inventions have reduced hazardous chemical use by more than
500 million kilograms. Which sounds great, until you consider that every day the U.S. produces or
imports about 33.5 billion kilograms of chemicals. The annals of green chemistry are full of crazy,
fascinating stories, like a plan to turn the unmarketable potatoes from Maine's annual harvest into
biodegradable plastics. Still, a decade after the phrase was coined, green chemistry patents made up less
than 1 percent of patents in chemical-heavy industries.
What will it take for green chemistry to be more than the proverbial drop in the bucket, a bucket full of
toxic sludge? Some experts believe that the answer is government intervention—not only laws that ban
harmful chemicals, but laws that simply require chemical manufacturers to reveal safety data and let the
market do the rest. "Right now, companies that make chairs or cars or lipstick don't know which of the
chemicals they incorporate into their products are safe," says Michael Wilson, an environmental health
scientist at the University of California, Berkeley. "Once that information becomes available, there will
be a demand for less toxic ingredients."
That question—to regulate or not to regulate—has split the community of green chemistry advocates.
Some oppose making green chemistry mandatory: its principles are so sensible and cost-effective, they
believe, that industry will implement them voluntarily. Others, such as Wilson, disagree. The key, he
asserts, is "fundamental chemicals policy reform in the U.S."
Now is a critical time: After decades of inaction, the U.S. government is finally examining more
aggressively the health effects of common chemicals. The ambitious Safe Chemicals Act, unveiled last
month in the U.S. Senate, would require all industrial chemicals to be proved safe, creating a strong
incentive for the development of less harmful alternatives. And the President's Cancer Panel released a
landmark report earlier this month decrying the "grievous harm" done by cancer-causing chemicals such
as bisphenol A in food and household products.
The stakes are high, higher than most people realize. The companies that make the 80,000 chemicals
that circulate in our world are rarely required to do safety testing, and government agencies are
relatively powerless. "This is pretty shocking, since most people assume that someone is checking
what's on the market. The ingredients in my shampoo? The ingredients in my child's toys? No one's on
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the job? And that's the answer: By and large, no one's on the job," says Daryl Ditz, a senior policy
adviser at the Center for International Environmental Law (CIEL) in Washington, D.C.
"If we're going to continue on as an industrial society that's based on synthetic chemicals, we've got to
figure out a way around this stuff. There's really no question about that," says Jody Roberts, an
environmental policy expert at the Chemical Heritage Foundation in Philadelphia, Pa. "I think that's
where the frustration for some people is, that it needs to be happening faster."
Green chemistry's beginnings
Perhaps no one has gambled more on green chemistry than John Warner. Along with Paul Anastas, the
co-founder of green chemistry and now the assistant administrator for the EPA's Office of Research and
Development, he helped create a federal awards program that brought the field into the mainstream. And
with Anastas he literally wrote the book: Green Chemistry: Theory and Practice, what Warner calls "a
how-to guide at the molecular level." In it they establish 12 guiding principles for chemists, concepts
like preventing waste by incorporating as much of the materials used into the final product, and
choosing the least complicated reaction.
A dozen years ago Warner, 47, left a lucrative job at Polaroid to found the nation's first doctoral program
in green chemistry. In 2007, tired of lecturing that green chemistry is the wave of the future, he decided
to prove it, founding a start-up, the Warner Babcock Institute for Green Chemistry, in Wilmington,
Mass. His firm, staffed by two dozen bright young scientists, is an ingenuity factory. They are working
on all kinds of projects: a less energy-intensive way to make solar panels, a cheap water purification
device for the developing world, and materials that mimic eye and liver tissue to substitute for live
animals in toxicity testing.
Some of the work is basic research. One of Warner's core technologies is based on thymine, one of the
four bases of DNA. When exposed to light, thymine molecules attach to one another; because this
reaction can be harmful (think: skin cancer) many organisms possess enzymes tasked with breaking
those bonds. If you put thymine in a substance and expose it to light, it hardens; apply enzymes and it
softens again. No toxicity, many potential applications. A scientist in Warner's lab is using this
technology to perm hair without caustic chemicals—simply by coating curled strands with a thyminebased polymer then shining light to freeze them in place. The technology could also act as a masking
technique during the manufacture of printed circuit boards. Or imagine truly recyclable plastics that
could be returned to their raw materials after the user throws them away.
That practical vision is a product of Warner's upbringing. He grew up in Quincy, Mass., a tough
working-class town south of Boston, and he hasn't shed the local dialect. "I am a chemist. I make
molecules," he says, as if he could just as easily be building a house or an engine. In his plaid shirt and
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scuffed sneakers, he comes across more like the kind of guy you might bring your car to when it makes
a funny rattle. Warner's uncles, Sicilian immigrants, worked in construction and stone cutting, and he
sees no disconnect between his blue-collar beginnings and his current gig running a 40,000-square-foot
high-tech lab. "I had uncles with half fingers. I respect that—doing things with your hands, creating
things," he says. "I feel that I'm working with my hands, but just in a different kind of way."
To a chemist, atoms are like so many Lego blocks to arrange and rearrange at will. Add a hydroxyl here,
a phosphate there, and react with various other chemicals to get the desired color, hardness, transparency
or other properties. "If we can draw a molecule, if it doesn't violate some fundamental law, we probably
can make it," Warner says.
But chemistry was invented at a time when people weren't thinking about the environmental impacts.
Raw materials are typically derived from fossil fuels. Turning them into the desired product can be a
multistep process involving hazardous reagents (chemicals that react with the target material) and
solvents (liquids or gases that provide an environment for the reaction to take place). Reactions often
generate more unwanted than wanted chemicals. In making pharmaceuticals, for example, it is not
uncommon to end up with 25 to 100 pounds of waste for every pound of medication.
Green chemistry starts with renewable resources such as plants or microorganisms, recycles its reagents,
uses less hazardous solvents, and streamlines complicated processes. For example, in 2006 Pfizer
changed the way it makes its nerve-pain drug Lyrica, substituting two plant-based enzymes for a
common metallic catalyst called Raney nickel. The process now occurs at room temperature and in
water, takes four instead of 10 steps, and has slashed waste and energy use by more than 80 percent.
Why so slow?
Green chemistry is elegant. It's sensible. It has the potential to improve public health and enhance the
economy. But if everyone loves green chemistry—scientists, environmentalists, politicians, corporate
leaders—then why hasn't it been more successful? After 15 years of innovation, the chemical industry is
as toxic as ever. The politicians who lavish funding on nanotech dole out pathetically little to green
chemistry. The universities that train chemists still do not require students to take a single course in
toxicology. And green chemistry is far from becoming a household phrase.
To many observers, the answer is clear: What's needed is more regulation. "One way to think about it is
to ask yourself: 'What is the purpose of government? Why isn't everything done by voluntary exchange
among willing buyers and sellers?' The answer is, of course, that a lot of important things that need
doing won't be done voluntarily," says Edward Woodhouse, a political scientist at Rensselaer
Polytechnic Institute in Troy, N.Y. "It does require stick as well as carrot." Wilson and his Berkeley
colleagues have acted on that principle; they helped craft the nation's first green-chemistry laws, enacted
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in 2008 in California. These laws require the state to identify, prioritize and take action on chemicals of
concern, to encourage safer alternatives, and to make hazard information available to the state's
businesses and to the public.
Warner is all for transparency, but being a chemist himself, he knows how his colleagues think, and he's
concerned that if green chemistry becomes mandatory, industrial chemists will misunderstand it, writing
it off as a policy-wonk proposal when in fact it is solid science, built on the core principles of traditional
chemistry. Warner favors the "build a better mousetrap" philosophy: Do green chemistry by making
alternatives that are not only safer but effective and economical, and chemical companies will eagerly
adopt them.
But others insist that until heavyweights like Dow and ExxonMobil are forced to own up to the dangers
of their chemicals, smaller companies developing clean alternatives won't be able to compete. "Some
academics say, 'If we had enough students and research dollars, then wonderful new substances would
flow from our labs and the world would beat a path to our door,'" CIEL's Ditz says. "But if no one can
distinguish between a green molecule and a toxic molecule, it is almost impossible for safer products to
break into the market."
No shift this big happens without conflict, and outrage, Woodhouse says. Average people need to know
and care enough about chemical hazards to pressure business and political leaders for change. "Most
people have no idea that many of the things in their houses are a danger to them," he says. "I don't think
that the urgent need for a benign chemical transformation has been put out very effectively."
In addition, even when scientists come up with nontoxic, cost-saving technologies, they don't always see
the light of day. The up-front expense of redesigning factories often eclipses the potential long-term
savings. "Your plants are set up to run nonstop. Any downtime, even if it's going to save you a million
dollars later, is costing you money now," Chemical Heritage's Roberts says.
Warner's concern is that when government gets ahead of science, the effort often backfires. "The ban
will say, 'Use the best available technology.' If the best available technology is nasty, the ban becomes a
license to use that technology," he says. "You can't legislate an invention, only encourage it."
The other side of the coin, however, is that sometimes when government gets ahead of science, science
rushes to catch up. That happened in the mid-1990s, when the chemical company Rohm and Haas
learned that a ban on tin-based marine paints was in the works. Tin-based paints had been used on ships'
hulls for years because they discouraged the growth of barnacles, algae, bacteria and other unwanted
hitchhikers. But tin is toxic and it was accumulating in fish, seabirds and other animals. Japan banned
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tin-based paints in 1992 and other nations were poised to follow suit. Rohm and Haas had never made
ingredients for marine paint—and without the pending ban it would not have tried, because tin-based
paint manufacturers dominated the market. But Rohm and Haas already had a mildew-fighting chemical
t hat acted as a wood preservative . By adapting that active ingredient, company scientists developed
Sea-Nine, a chemical that kills marine organisms by reacting with their own chemistry, breaking down
into nonhazardous components in the process.
However it happens, changing worldviews takes time. It took two decades or more for global warming
to gain any serious traction. Now it is seen as an opportunity to develop a whole new sector of the
economy: alternative energy. The same could happen for green chemistry, as a demand for cleaner
products drives innovation.
What everyone agrees on is that, ultimately, green chemistry principles must become so integrated into
mainstream chemistry that the term loses its meaning. Ironically, we'll know that green chemistry has
succeeded when it disappears. "The day that everyone from kindergarten students on up gets it, we don't
need the field of green chemistry anymore," Warner says. "That is my goal, for it to be just the way
everybody sees science."
To explain the goals of green chemistry, John Warner uses the metaphor of the toolbox. Rather
than wrenches, nuts and bolts, the drawers in the chemical industry's "toolbox" contain
commonly used processes, such as ways to make carbon compounds or oxidation-reduction
reactions. Most of these processes involve hazardous chemicals. Green chemists aim to create a
new toolbox filled with less harmful alternatives, so that in the future when chemists set out to
design a molecule, they'll be able to put their hands on benign tools to get the job done.
GREEN CHEMISTY PRIMER
Here are some promising new technologies destined for the green-chemistry toolbox.
TAMLs: There's no pretty way to say it—TAML is short for tetra-amido macrocyclic ligand—but these
apparently harmless chemicals break down a variety of stubborn pollutants, including pesticides, dyes
and industrial runoff. Developed by Terrence Collins, a chemist at Carnegie Mellon University in
Pittsburgh, TAMLs mimic the enzymes in our bodies that have evolved to fight off toxic assaults.
Collins and his team worked for two decades to develop these smaller, easy-to-build versions of
biological enzymes. When combined with hydrogen peroxide, TAMLs neutralize many contaminants by
breaking their chemical bonds.
Noncovalent derivatization: A longtime passion of Warner's (his license plate reads "NCD"),
noncovalent derivatization is chemistry with a light touch. Covalent bonds are the strong connections
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between atoms that hold molecules together. Normally, when chemists are dissatisfied with some aspect
of a molecule they are creating, they alter its structure by breaking or adding covalent bonds. Such
changes can involve multiple steps and hazardous ingredients. Warner's breakthrough was to posit that
sometimes there's no need to create a new molecule. Simply combine the existing molecule with another
substance that interacts with it, and the transient forces between them can effect the desired change.
"With no energy they find each other and form," he says. "Why does a bunch of lipids fold up to form a
cell membrane? Why does DNA form a double helix? It's always these weak molecular structures."
Liquid CO2: Most of us know carbon dioxide as a gas (we exhale it) or a solid (think: dry ice in fog
machines). But when you put carbon dioxide under pressure, it becomes a liquid. Liquid CO2 is a
benign substitute for the nasty solvents typically used to decaffeinate coffee. Just mix it with green
coffee beans, then take the pressure off. The carbon dioxide evaporates, leaving behind a pile of white
powder—caffeine. Do the same thing to dirty clothes and you extract oils and grime without using
perchloroethylene, the notorious dry cleaning chemical.
Green Chemistry Might Revive Science
Training
Universities find that environmentally friendly chemistry draws more student interest-and could have an outsized impact on industry
By Sara Goodman | March 25, 2009 | Scientific American
http://www.scientificamerican.com/article.cfm?id=green-chemistry-revives-training
GREEN CHEM CLASS? There's a growing demand for chemists
who can help companies meet consumer demands for more
sustainable products, and schools are racing to keep up. Image:
FLICKR/HYKU
Slowly, the chemical industry is going green.
Many companies are starting to emphasize reducing
or eliminating hazardous substances to save money,
reduce inefficiencies and promote their brands to
consumers who favor eco-friendly products.
"Industry really sees the value of 'green chemistry,'"
said Julie Haack, assistant head of the University of Oregon's chemistry department. "If you
want to recruit the best chemists, wouldn't it make sense to promote the opportunity to work in
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an environment where they can align their interest in the environment with their passion,
which is chemistry?"
Having employees concerned about sustainability leads to more innovative, long-term
solutions, said Neil Hawkins, vice president for sustainability at Dow Chemical Co.
"It's very important to us to have a pipeline of the best and the brightest students in science
and technology, but that also have a broader perspective, as well, so they can understand the
tradeoffs," Hawkins said. "This means balancing environmental, social and economic
decisions."
Many universities are responding by creating a green-chemistry curriculum. Their efforts
require addressing what green chemistry advocates call a fundamental problem in chemistry
education: a lack of toxicology training.
"Students can earn a doctoral degree in chemistry in nearly every university in the country and
not have to demonstrate a basic understanding of toxicology or eco-toxicology – how to design
a molecule that doesn't disrupt the endocrine in some way," said Michael Wilson, assistant
research scientist at the University of California, Berkeley.
But students, faculty and industry are starting to change that by pushing for programs and
courses about alternate design principles, slowly shifting chemistry education.
The University of Oregon – a leader in the movement – began an outreach program nine years
ago that teaches professors nationwide about integrating green chemistry into a curriculum.
Haack said that effort has driven up demand for green-chemistry courses nationwide and has
led to changes in how students and faculty approach chemistry.
"We've seen subtle shifts," Haack said in an interview. "Instead of students questioning the
mechanics of something, now they're thinking about chemistry as a tool for sustainability.
They're excited about the possibility of designing out hazards.
'Breaking down walls'
Colleges and universities are eager to tap this enthusiasm to revitalize interest in science, which
has has been flagging for years. Oregon, for example, offers courses in other disciplines that
address green chemistry, including ethics, marketing and public policy development.
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The university has an online database that collects education materials focused on green
chemistry and allows professors and other interested parties to share insights and questions
about teaching techniques and field studies. The database currently lists 240 individuals from
around the country who have expressed professional interest in green chemistry.
Some see green chemistry as a vehicle for crossing lines between academic disciplines. The
University of Oregon, for example, has 11 faculty members doing green chemistry research in
some capacity in various departments.
"There has been far too much fragmentation in academia," said Paul Anastas, director of the
Center for Green Chemistry and Green Engineering at Yale University. "That is the tradition
which things like green chemistry and sustainability are battling against by bringing together
different perspectives and breaking down walls."
Dow has also been working to promote a cross-disciplinary approach to education, Hawkins
said. The company in 2007 launched a collaboration with the University of California,
Berkeley, to create a university-wide program to address a wide range of issues – from
expanding water supplies to measuring the environmental impacts of products' supply chains.
"In the area of sustainable chemistry, it's really important that we look at all of the different
dimensions, because technology alone is not the answer," Hawkins said. "Having a broad
experience, not just in the technical field but looking at the business aspects as well as the
social aspects, is extremely important for the executives, managers and leaders of tomorrow."
The company last year also began the Dow Sustainability Innovation Student Challenge award,
which brings together six universities – the University of Michigan, Northwestern University,
Tufts University, the University of Cambridge, Brazil's University of São Paulo, and China's
Peking University – to recognize innovative student projects.
"If you look at the quality of the projects, you'll see that young people today are on fire to take
their skills and know-how and make a difference in the world," Hawkins said. "As we go out
and recruit, to virtually any place on the planet, these issues of sustainability are very
important to students – students want to know what a company does in this space."
Market forces
The far-reaching implications of green chemistry and design – spreading across disciplines and
applications – suggest that the field has tremendous potential for growth, Yale's Anastas said.
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Green chemistry applications make up 1 percent of the total chemical market share, Anastas
estimated.
Hawkins predicted that the drive to expand green chemistry will come from the market, not
regulations.
"Back 20 years, the world was dealing with these issues more in a regulatory way," Hawkins
said. "What we see today, with cap and trade or other systems, are really economic models for
helping the private sector become a part of the solutions in a sustainable way. Technological
and innovative business models can help fuel a more sustainable economy."
Ultimately, UC Berkeley's Wilson said, market forces rule.
"You have to get the market working properly," Wilson said. "Once you do that, demand for
trained chemists who understand green chemistry and toxicology will ramp up. Universities
will respond, as will research."
Reprinted from Greenwire with permission from Environment & Energy Publishing, LLC.
www.eenews.net, 202-628-6500
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