Download Post Lab Questions

1
2
College Preparatory
High School Chemistry Laboratory Manual
and
Course Curriculum
10th Edition 2015
“A day without chemistry is a day without sunshine”
- Vicki Collins, Warren Wilson College,
Asheville, NC
“The world is not only stranger than we imagine, but
it is stranger than we can imagine.”
- Haldane, J.S. as per Dean Kahl, Warren Wilson College,
Asheville, NC
Compiled and developed
by
Brian Wright M.Ed.
First edition 2004
10th edition 2015
For the students and faculty of Olympia High School
Copyright 2014 Brian P. Wright
All rights reserved
Olympia High School 1302 North Street
Olympia WA 98501
3
4
Table of Contents:
Number
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
19
20
Section
Course Policies and Procedures
Safety Contract
How to be Successful
How to Read a Science Textbook
Course Syllabus
Lab Notebook Set-up
How to Write a Lab Report
How to Write a Discussion
Experiments
Do Your Own Demo Principled of Chemical Reactions
Techniques and Measurement
Emission Spectroscopy
Rapid Oxidation of Metallic Fuel
Inorganic Nomenclature
Determination of an Empirical Formula
Reaction Rate: Marble Lab
Flame Test
Qualitative Analysis
Chemical Reactions
Stoichiometry
Beer’s Law
Titration
Comparisons of Geometry and Shape
Vapor Pressure of Liquids
Calorimetry
Boyle’s Law
Quantitative Reaction of HCl and Mg
Heat of Neutralization
Model Project
Appendix: Reference Material
Page
6
9
11
12
12
13
14
15
16
23
30
38
42
48
54
58
62
69
79
85
89
98
108
113
116
119
126
132
136
5
Mr. Wright
Chemistry
[email protected]
Grading
Your grade will be determined as follows:
Approximate value
Exams
30%
Quizzes
30%
Problem sets and in class assignments 5%
Laboratory experiment reports
30%
Notebook
5%
A
AB+
B
BC+
93.0%
90.0%
87.0%
82.0%
79.0%
76.0%
C
CD+
D
66.0%
64.0%
60.0%
52.0%
Required Course Materials:
• Approved safety goggles
• Non-graphing Scientific calculator (TI 30X IIS is recommended)
• Notebook sized periodic table (Student Store)
• Black or blue pen for all work, and a red or green pen for corrections only
• Straight edge (clear plastic is best)
• Mini stapler that uses standard staples.
• Notebook paper
• Three-ring notebook with dividers
• Duplicating laboratory notebook (not spiral bound)
• Interactive notebook (spiral notebook)
• Calendar/Planner to record assignments and test dates
• Lab manual
A more detailed description of these items will be provided during the first days of class. Students are expected
to bring a required materials everyday to class. Not having needed materials or not being prepared and ready to
work when the bell rings is counted as a tardy. I will not loan you a calculator for use during a test or quiz.
You will not be allowed to use your phone as a calculator.
Laboratory work:
The only person you may work with on a lab is your lab partner. You will work with this person all
year long. If your lab partner drop the course or is absent you will work alone. Your lab group is
composed of you and your partner and the pair working across from you. Please discuss your lab with
your lab group. Each partnership must complete the lab independently from the lab group.
You will not be allowed to work in the laboratory until you successfully complete the laboratory safety
test. Safety rules will be enforced at all times during laboratory experiments. You will receive one
correction (warning) for laboratory safety violations. On the second correction you will be required
stop working on the experiment. If you are not prepared for laboratory work or not dressed
appropriately for laboratory work, you will not be allowed to participate. During laboratory activities
you are required to wear closed-toe shoes and clothing that completely covers your legs.
Attendance:
This is a college preparatory course. Ninety percent of your future success is just SHOWING UP.
You will have several legitimate reasons to be away from class. Please realize when you are not
here you miss out. You are expected to take responsibility for your learning. If you miss a class for
any reason it is your responsibility to learn what you missed. I want you to succeed and I care about
your learning. Your success in this class is dependent on your consistent attendance. When the bell
6
rings to begin class I expect you to be in your seat and ready to work. If you are not in your seat
you will be counted tardy. My attendance policy is consistent with the policy found in your student
handbook. If you have any questions regarding this I am happy to discuss it with you.
Policy for making up missed work and turning in late assignments:
If you are going to miss class please inform me with an email. Your opportunity to make up work
is dependent on written approval.
In general most missed work is due on the next Friday. For any planned absence work that is due
during the absence must be turned in before your departure. Laboratories, quizzes, and exams
cannot be made up unless prior arrangements in writing have been made.
If you make prior arrangements and miss laboratory work you have one week or until the materials
are put away to complete your work. In many cases it is not possible to make up laboratory work.
The instructor may be able to provide you data so that you can complete the laboratory report. If
you miss an exam you have one week to make it up or until corrected exam are returned. You can
drop up to three quiz grades from your final grade. Quiz make-ups are not generally given
therefore, if you miss a quiz it will count as one of your three dropped quizzes. Making
arrangements to take a quiz before your absence is possible. Make-up quizzes are not generally
given. (If there are fewer then 10 quizzes then only 2 will be dropped)
Classroom Behavior:
Each student is expected to behave in a manner that enhances the learning process. Inappropriate
behavior will be dealt with following the Olympia High School behavioral management program, as
found in your student handbook. During the first week of class we will discuss specific classroom
expectations. Each student is expected to be in class on time and prepared to learn every day.
Academic/Scientific Honesty:
You will often work with other students in this class. However the work you turn in must be
completely yours. No part of your work can be copied from another student. Your work will be
compared to other students work, including your lab partner. If any portion of your work does not
appear to be original it will be deemed to be a violation of the schools policy regarding cheating and
plagiarism. If you are working in a “lab group” each student’s participation must be acknowledged.
You are only allowed to conduct laboratory experiments with your assigned lab partner. If your
partner is absent then you will work alone, you may not work with someone else without permission
in writing. If your partner is not present to conduct the lab you may not share your data with them
without your instructor’s permission.
Classroom Procedures:
At the start of the period
You are required to be in your seat and ready to work when the bell rings. If you are not, you will
be marked tardy. Upon entering the classroom, collect graded work from your period’s box. Unless
otherwise stated, all work (laboratory reports, homework, etc.) is due at the beginning of class.
Backpacks and bags are not allowed to sit on the floor. Hooks are conveniently placed on desks to
hang your backpacks and bags.
During the period
Do not ask to leave the class for personal reasons. Please remain in your seat during class time
unless instructed otherwise. Raise your hand if you wish to ask a question or add a comment. No
use of personal electronic devices such as MP3 players or cell phones is allowed in the classroom
7
and any electronics must remain in your backpack or bag. Permission must be explicitly granted to
text, record sound or images. Calculators are the only approved electronic devices for classroom
use. Do not consume gum, food, or drink nor apply cosmetics in the classroom, it is not safe
because this is also a science laboratory.
At the end of the period
Students will stay in their assigned seats until the class is dismissed. Your instructor will dismiss
the class. The bell does not dismiss the class.
Explanation of Evaluated Course Work:
Your assignments are a reflection of you, your commitment to quality and your interest in the class.
All assignments will be turned in on flat, smooth paper with no tears. Notebook paper will not have
spiral notebook fuzz. All assignments are to be done in ink, blue or black only. Assignments
should have your name, your class ID number, and a heading which includes the date, title of the
assignment, and your period.
Exams: will cover material presented in the current unit. Unless otherwise stated the only materials
students may use to complete midterm exams are a calculator and a provided periodic table.
Most exams will have a practical laboratory component.
Quizzes You will be allowed to drop your two worst quizzes (unless there is 10 or more quizzes
then worst three will be dropped). Some quizzes will be given without prior planning.
Problem sets will be assigned frequently. They will be checked in class on the following day.
They may or may not be collected for credit.
Laboratory reports are due the day after the lab is completed. The only exception is if the
following day is a test or quiz day. Then the lab will be due on the day after the test or quiz.
A complete description of how to write a lab report will be handed out to each student.
Notebooks: You will keep a notebook, which will contain all handouts and returned work.
Periodically and without warning you will be asked to produce three items from your
notebook. Your notebook is for your own use. I would suggest that you organize it in a
manner that allows you to find useful information quickly. A possible scheme for organizing
your notebook follows.
§ Section 1: Course syllabus and laboratory report guidelines at the front of the
notebook.
§ Section 2: Graded exams, quizzes, and problem sets in chronological order.
§ Section 3: Class notes and handouts in chronological order.
§ Section 4: Laboratory instruction handouts in chronological order.
Interactive Notebook: your Interactive notebook (INB) will be kept in a spiral notebook. This
will be a place to keep your class notes, and other work from the class. The INB has a fairly
strict format that will be described in detail.
Final exams will be cumulative. All guidelines that applied to the midterm exams will also apply to
the final exam.
This constitutes one of your first assignments. Please read, understand, and sign
this syllabus, then have Mr. Wright check it off.
Thank you, Brian P. Wright NBCT, M.Ed
Student Signature_____________________________________
Parent or Guardian Signature_____________________________
8
Student Laboratory Safety Contract
Purpose
Science is a hands-on laboratory class. You will be doing many
laboratory activities which require the use of hazardous chemicals.
Safety in the classroom is the #1 priority for students, teachers and
parents. To ensure a safe science classroom, a list of rules has
been developed and provided to you in this student safety contract.
They must be followed at all times. Two copies of the contract are
provided. One copy is to be signed by both you and a
parent/guardian. The other is to be kept in your science notebook
as a constant reminder of safety rules.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
General Guidelines:
Conduct yourself in a responsible manner at all times in the
laboratory.
Follow all written and verbal instructions carefully. If you do
not understand a direction or part of the procedure, ask the
instructor before proceeding.
Never work alone. No student may work in the laboratory
without an instructor present.
When first entering a science room, do not touch any equipment,
chemicals, or materials in the laboratory until you are instructed to
do so.
Do not eat, drink or chew gum. Do not use laboratory glassware
as containers for food or beverage.
Perform only those experiments authorized by your instructor.
Never do anything in the laboratory that is not called for in the
laboratory procedures or by your instructor. Carefully follow all
directions, both written and oral. Unauthorized experiments are
prohibited.
Be prepared for your work in the laboratory. Read all
procedures thoroughly before entering the laboratory. Never fool
around in the laboratory. Horseplay, practical jokes, and pranks
are dangerous and prohibited.
Observe good housekeeping practices. Work areas should be
kept clean and tidy at all times. Bring only your laboratory
instructions, worksheets, and/or reports to the lab area. Other
materials (books, purses, backpacks) should be stored in the
classroom area.
Keep aisles clear and your chair pushed under your desk.
Know the locations and operating procedures of all the safety
equipment including the first aid kit, the eye wash station, safety
shower, fire extinguisher, and fire blanket. Know where the fire
alarm and exits are located.
Always work in a well ventilated area. Use the fume hood when
working with volatile substances or poisonous vapors. Never place
your head into the fume hood.
Be alert and proceed with caution at all times in the laboratory.
Notify the instructor immediately of any unsafe conditions you
observe.
Dispose of all chemical waste properly. Never mix chemicals in
sink drains. Sinks are to be used only for water and solutions
designated by the instructor. Solid chemicals, metals, matches,
filter paper, and all other insoluble materials are to be disposed of
in proper waste containers, not in the sink. Check the label of all
waste containers twice before adding your chemical waste to the
container.
Labels and equipment instructions must be read carefully before
use. Set up and use the prescribed apparatus as directed in the
laboratory instructions or by your instructor.
Keep hands away from face, eyes, mouth, and body while using
chemicals or preserved specimen. Wash your hands with soap and
water after performing all experiments. Clean (with
detergent), rinse, and wipe dry all work surfaces (including
the sink) and apparatuses at the end of the experiment.
Return all equipment clean and in working order to the
proper storage area.
16.
Experiments must be personally monitored at all times.
You will be assigned a laboratory station at which to work.
Do not wander around the room, distract other students, or
interfere with the laboratory experiments of others.
17.
Students are never permitted in the science storage
rooms or preparation areas unless given specific permission
by their instructor.
18.
Know what to do if there is a fire drill during a
laboratory period; containers must be closed, gas valves
turned off, fume hoods turned off, and any electrical
equipment turned off.
19.
Handle all living organisms used in a laboratory activity
in a humane manner. Preserved biological materials are to
be treated with respect and disposed of properly.
20.
When using knives and other sharp instruments, always
carry the points and tips pointing down. Always cut away
from your body.
Never try to catch falling sharp
instruments. Grasp sharp instruments only by the handles.
Clothing
21.
Any time chemicals, heat or glassware are used, students
will wear laboratory goggles. There will be no exceptions!
22.
Contact lenses should not be worn in the laboratory
unless you have permission by your instructor.
23.
Dress properly during laboratory activities. Long hair,
dangling jewelry, and loose/baggy clothing are a hazard in
the laboratory. Long hair must be tied back, and dangling
jewelry and loose/baggy clothing secured. Shoes must
completely cover the foot; no sandals are allowed.
24.
Long pants are required for lab work and natural fibers
are suggested.
Accidents and Injuries
25.
Report any accidents (spill, breakage etc.) or injury (cut,
burn, etc.) to the instructor immediately, no matter how
trivial it may appear.
26.
If you or your lab partner are hurt immediately yell out
“code one, code one” to get the instructor’s attention.
27.
If a chemical should splash in your eyes or on your skin,
immediately flush with running water from the eye wash
station or safety shower for at least 20 minutes. Notify the
instructor immediately. When mercury thermometers are
broken, mercury must not be touched.
Notify the
instructor.
Handling Chemicals
28.
All chemicals in the laboratory are to be considered
dangerous. Do not touch, taste, or smell any chemicals
unless instructed to do so. The proper technique to smell
chemicals will be demonstrated.
9
29.
32.
33.
34.
35.
Check the label on chemical bottles twice before
50.
removing any of the contents. Take only as much as you
need.
30.
Never return unused chemicals to their original
51.
containers.
31.
Never use mouth suction to fill a pipette. Use a rubber
bulb or pipette bulb.
When transferring reagents from one container to another, hold
the containers away from your body.
Acids must be handled with extreme care. You will be shown
the proper method for diluting strong acids. Always add acid to
water, swirl or stir the solution, and be careful of the heat
produced, particularly with sulfuric acid.
Handle flammable hazardous liquids over a pan to contain spills. 1.
Never dispense flammable liquids anywhere near a source of flame 1.
or heat.
2.
Take great care when transferring acids and other chemicals
from one part of the laboratory to another. Hold them securely and
walk carefully.
Do not place hot apparatuses directly on the laboratory desk.
Always use an insulating pad. Allow plenty of time for hot
apparatuses to cool before touching.
When bending glass, allow time for the glass to cool before
handling. Hot and cold glass have the same visual appearance.
Determine if an object is hot by bringing the back of your hand
close to it prior to grasping it.
Questions
Do you wear contact lenses?
yes____ no_____
Are you color blind?
yes_____ no_____
Do you have any allergies?
yes____ no_____
If yes, please list: __________________________
____________________________________________
Agreement
I, __________________(student’s name) have read and agree to
follow all of the safety rules set forth in this contract. I realize that
I must obey these rules to ensure my own safety, and that of my
fellow students and instructors. I will cooperate to the fullest
extent with my instructor and fellow students to maintain a safe lab
environment. I will always closely follow the oral and written
instructions provided by the instructor. I am aware that any
violation of this safety contract that results in unsafe conduct in the
laboratory or misbehaviour on my part, may result in being
removed from the laboratory, study hall, receiving a failing grade,
and/or dismissal from the course.
Handling Glassware and Equipment
36.
Carry glass tubing, especially long pieces, in a vertical position
to minimize the likelihood of breakage or injury.
37.
Never handle broken glass with your bare hands. Use a brush
and a dustpan to clean up broken glass. Place broken or waste
glass in the designated broken glass container.
38.
Inserting and removing glass tubing from rubber stoppers can be
dangerous. Always lubricate glassware (tubing, thistle tubes,
thermometers, etc.) before attempting to insert it into a stopper.
Always protect your hands with towels or cotton gloves when
inserting glass tubing, or removing it from, a rubber stopper. If a
piece of glassware becomes “frozen” in a stopper, take it to your
instructor for removal.
39.
Fill the wash bottles only with distilled water and use only as
intended, ex. rinsing glassware, or adding water to a container.
40.
When removing an electrical plug from its socket, grip the plug,
not the electrical cord. Hands must be completely dry before
touching an electrical switch, plug, or outlet.
41.
Examine the glassware before each use. Never use chipped or
cracked glassware. Never use dirty glassware.
42.
Report damaged electrical equipment immediately. Look for
things such as frayed cords, exposed wires, and loose connections.
Do not use damaged electrical equipment.
43.
If you do not understand how to use a piece of equipment, ask
the instructor for help.
44.
Do not immerse hot glassware in cold water, it may shatter.
_______________________________
student name (print)
_______________________________
student signature
_______________________________
date
Dear Parent or Guardian:
We feel that you should be informed regarding the school’s effort
to create and maintain a safe science classroom/laboratory
environment. With the cooperation of the instructors, parents, and
students, a safety instruction program can eliminate, prevent, and
correct possible hazards. You should be aware of the safety
instructions your son/daughter will receive before engaging in any
laboratory work. Please read the list of safety rules above. No
student will be permitted to perform laboratory activities unless
this contract is signed by both student and parent/guardian and is
on file with the teacher. Your signature on this contract indicates
that you have read this Student Safety Contract, are aware of the
measures taken to insure the safety of your student in the science
laboratory, and will instruct your son/daughter to uphold his/her
agreement to follow these rules and procedures in the laboratory.
Heating Substances
45.
Exercise extreme caution when using a gas burner. Take care
that hair, clothing, and hands are a safe distance from the flame at
all times. Do not put any substance into the flame unless
specifically instructed to do so. Never reach over an exposed
flame. Light gas (or alcohol) burners only as instructed by the
teacher.
46.
Never leave a lit burner unattended. Never leave anything that
is being heated or is visibly reacting unattended. Always turn the
burner or hot plate off when not in use.
47.
You will be instructed in the proper method of heating and
boiling liquids in test tubes. Do not point the open end of a test
tube being heated at yourself or any one else.
48.
Heated metals and glass remain hot for a long time. They
should be set aside to cool and only be picked up with caution. Use
tongs or heat protective gloves if necessary.
49.
Never look into a container that is being heated.
_______________________________
parent/guardian
_______________________________
date
10
How to Get an A
in
High School Chemistry
1. Come to class every day with the proper clothing, materials (calculator, class notebook, lab
notebook, homework, lab manual, etc.) and a positive attitude.
2. Keep an organized planner and notebook
3. Memorize all polyatomic ions and elements that you are asked to memorize. (See labs on
nomenclature and reactions)
4. Learn your naming rules. (See labs on nomenclature and reactions.)
5. Learn how to use significant figures.
6. Know how to solve for moles.
7. Know how to use dimensional analysis.
8. Keep up with your Interactive Notebook and do Left Hand Activities.
9. Make a note card for each quiz.
10. Ask questions regarding topics you do not understand.
11. Develop a study plan for each exam.
12. Turn in all assignments on time.
13. Have fun.
14. To maximize your time in the laboratory complete as much of the lab report as possible before
arriving on a lab day.
15. Read the text book effectively. To do this:
a. Skim the text, look at pictures, diagrams and tables
b. Read side bars and text for pictures, diagrams and tables
c. Read the text. If you do not understand a section, take note of this and keep on reading.
d. Read the section you did not understand and reread it until you do.
e. Take notes while reading. This includes paying special attention to any words that are
bolded. This may include making an outline.
f. Do the problems.
g. Do the reading over many short sessions.
h. Take reading notes.
11
[email protected]
Topics to be covered
History of science*
Measurement*
Physical vs chemical*
Properties of phases*
Precision, accuracy*
Significance**
Safety**
Dimensional analysis**
Know elements 1-38, 46-50,
53-56, 78-83, 86 87**
Atomic structure*
Sub particles
Theories of the atom
Experiments to
Determine structure*
Octet rule
Exceptions
Subatomic particle*
Size
location
Charge
Symbol
Mole*
Molecular weight
AMU
Empirical formula*
Nomenclature**
Acid/Bases
Two non metals
Metal / non metal
Polyatomic
Periodic table*
Development
Trends
Atomic size
Electronegativity
Ionization energy
Ion size
Titration*
Molarity*
M1V1=M2V2*
Solution making (Gramsà)*
Limiting reactants*
Reaction types*
Gas forming
Oxidation reduction
Combustion
Combination
Decomposition
Neutralization
Acid Base*
pH
Electron configuration*
Orbitals
Geometry
Noble gas configuration
Trends
VSEPR
Electron dot config
Bonding*
Polar
Nonpolar
Ionic
Lewis Structure**
Resonance
Octet
Hybridized orbitals
Formal Charge*
Gas Laws*
Combined
Charles’
Boyle’s
Kinetic theory
Ideal
Graham’s
Dalton’s
Diffusion
Effusion
12
Electrochemistry
Organic chemistry
Functional group
IUPAC naming
Substitution
Mechanism
Stereochem
Polymers
Inter- vs. intra-molecular
forces*
Balancing redox reactions
Final Project*
Laboratory activities
Measurement lab*
Reaction Rate*
Density lab*
Qualitative analysis*
Determination of an empirical
formula*
Making borate glass
Copper Brass (Alloy of copper)
Spectroscopy*
Modeling and Lewis structure
model lab*
Titration*
Stoichiometry*
Ice (testing Hypothesis)*
Heat or Reaction*
Heat of fusion*
Heat of neutralization*
End of year project*
Hydrogen rockets*
*Topics covered last year
**Fundamental topic that your
success depends on.
How to Prepare Your Notebook
and
Write a Lab Report
(Front cover)
(Title Page)
Chemistry 1
Olympia High School
2012-2013
Chemistry 1
Olympia High School
2012-2013
Name
Address
Email
Phone #
Name
Address
Email
Phone #
Table of contents
Date
Title
Description:
pg:
Write a Short title
Describe the purpose and results of experiment. This may take
several sentences.
Using various glassware and laboratory techniques, the density
of water will be determined.
Copper coins will be zinc plated and then heated to make
bronze.
Using NaOH of known molarity, the concentration, in
moles/liter, of an unknown diprotic acid will be
determined.
4
Determination
of
the
density of water
Copper to Gold
Titration of an unknown
acid
8
13
15
Skip one page before the first lab entry
Every page in your lab notebook should have the following information:
Name_____________________ Seat no.___ Period/Section_________
Page No. _____
Partner’s Name ______________
Experiment Title________________
Date____
______________________________________________________________________________
13
Brian Wright
Seat no. 32
Per 1
page 05
Dean Kahl
Determination of the density of water
2/22/02
The above sample heading, which includes you and your partner’s names, the date, page
number and the experiment being reported, must appear on every page of your lab notebook.
Read and understand the background information before writing your Pre-Lab. Each lab will need
the following sections.
Title: This is what your experiment is called. This will often be provided, however, if another
appropriate title makes more sense, you are welcome to use it.
part of Pre Lab
***The “Pre-Lab”***
Purpose: This is a short description of why you are doing the lab. What theory are you testing?
What do you hope to learn from the experiment?
Hypothesis: What are the likely results of this experiment? Why is this the likely result?
Pre-Lab Questions: These will be assigned questions. They typically will help prepare you to
succeed in the lab. Write out the question and then answer it. Use sig figs, units, and labels as
appropriate. Show your work and circle the answer.
Materials: This is a list of all equipment and supplies that are needed to conduct the experiment.
This is an excellent place to list chemical formulas and concentrations, as well as appropriate
molecular weights. Trust me on the concentrations and molecular weights. You will someday
wish you had written them down.
Planned Procedure: YOU WILL LOSE POINTS IF YOU ATTEMPT TO COPY THE
WRITTEN PROCEDURE. Instead what I would like you to do is to paraphrase the written
procedure. Read and understand the experimental procedure. Then write in your own words
what it is you plan to do. This is a way for me to gauge your level of understanding before you
do the experiment. Please indicate the source of the actual procedure that you plan to follow.
This may be written in the future tense. Please paraphrase and keep it short, about 4-5
sentences. It may need to be longer if the lab has more then one part.
Procedure: Write down what you actually did. Please note any variations from the planned
procedure. An acceptable procedure may state: “Followed procedure as per page ## in
Laboratory Manual (Wright 2013) with the following revisions…
Data: This is where you organize your data. Data can take the form of observations, which are
qualitative, and measurements, which are quantitative. In all labs I expect a qualitative
observation section (color, time, vigor, sound, smell) and in many labs a quantitative data
section. Any time you make a table you are required to use a straight edge. All labels and
units as well as significant figures must be used when making observations.
****The only place that data is recorded is directly into your lab notebook.***
Calculations: Any time you do any mathematical operations with data you must include it in your
calculations section, even if it is a simple adding or subtracting problem. It is very important to
show every step. Show the correct units and make sure to use significant figures. Remember: do
not round until the end, just keep track of the sig figs as you go. You must show all work that you
do and must keep track of units, labels and sig figs. Label each calculation so that it is very clear
as to which part of the lab the calculation is relevant. Please use dimensional analysis and or the
“4 step method” as appropriate.
14
Results: This is where you can report the objective consequences of your experiment. You can
report what happened. This is not where you interpret what anything means. Your results
should somehow reflect the intentions stated in your purpose. Any time you manipulate data
using mathematical operations you get an answer. These are your results. Report what your
calculations equal. This is not a section for you to report what the data means or what the
results imply. It is just a place for you to present the facts.
Discussion: In every lab report I expect the following issues to be discussed:
• Make a statement regarding how the data, observations and results relate to the stated purpose of
the lab. And whether the results supported or refuted the stated hypothesis (do not write “it
was proved…”)
• What went wrong and what effect might those errors have in your results? I am looking for a
well thought out treatment of how the data may be flawed, and what specific effect this would
have on the results.
• Compare your results with known values or theoretical values. Then determine the % error
between the known value and the experimental value.
• What could be done differently to improve the validity of the experimental results?
• What are appropriate follow up experiments?
Post Lab Questions:
These will be direct questions that I may ask you. Remember to always write out the question and
then answer in complete sentences. If it is a calculation, remember to use sig figs, units and
labels.
Guidelines to Scientific Writing
The general purpose of any scientific paper is to report what happened in enough detail that other
people can read your work, follow what you did, and understand your results. It needs to be clear
and concise.
When writing for a science class (at least Mr. Wright’s) please use the following style conventions.
• Write in the past tense.
• Leave out any affective comments regarding how you feel about the experiment or your work.
• People do not do things in science, things happen. Except in an extraordinary circumstance
leave out I, you, we, he, she, it.
• Be clear and concise.
• Make as few assumptions as possible.
• Write using the objective voice. Avoid using colloquialisms.
• Your grade depends on presentation. Be very neat, use rulers, use pens, and if you make a
mistake put a single line through the mistake instead of scribbling all over it.
• Never destroy your data or work; never tear out pages.
• Take your time; create something you can be proud of. Do it right the first time.
• Make statements that you can support with experimental evidence.
• Use a tentative voice in your writing.
15
Experiment 1
Principals of Chemical Reactions
DYOD’s
Purpose:
The purpose of this activity is to become familiar with working with chemicals and provide context
for the reactions that will be covered thoughout the rest of the year. These experiments will provide
opportunities to learn about laboratory safety and proper chemical handling and disposal.
Hypothesis:
Before each experiment you will be asked to make some kind of prediction about the out come of
the experiment.
Background:
What does it feel like to be a scientist? What causes matter to change state: go from solid to liquid,
liquid to gas or gas to liquid? What happens in chemical reactions? You and your group will
perform the same experiment that have inspired countless people to study science. Before
beginning the experiment, predict what might happen and what might be produced. As your group
proceeds through the activity, you will have many questions about the changes in matter that are
occurring. Record these questions for discussion at the end of the activity.
Detailed description of each part of your lab report:
The following laboratory report format is based on my personal work in both industrial and
academic laboratories. It is also based on what will be expected from you when you take college
science classes. I think the one primary criticism of the following instructions is that it leaves out
any personal reflection. Many scientists use their lab notebook as a journal as well as a place to
record their experiments.
Reading the Background:
Please make sure to read and understand all the information found in the Background. The goal is
that all the information necessary to do the lab should be found in the Background. This includes
formulas, constants and how to do the calculations.
• Writing the Pre-Lab:
It is essential that you know what do in the lab before you arrive to class. To ensure this, you will
write the Pre-Lab. Please carefully follow the directions on how to write a lab report. If your PreLab is not completed you will not be allowed to participate in the laboratory activity. Please
construct all data tables prior to coming to class.
• Procedure:
The procedure will be written in two parts. The first part is what you plan to do. The second part is
comments on what you actually did and how it deviated from what you planned. When writing
your planed procedure please summarize the detailed procedure that you plan to follow in 4-5
sentences. For some procedures it might be a little longer and for some a little shorter. Focus on
what you are doing, why you are doing it. I would encourage you to draw pictures of the
experimental setup and the step you plan to take. Please mention what procedure you are
referencing. The second part of the procedure is completed after you have done the lab. This is
16
where you discuss any revisions you made to the actual procedure you followed. For instance the
lab called for a 250mL beaker and you used a 250mL Erlenmeyer flask.
• Data Table:
The data to be collected must be based on the purpose of the experiment, the hypothesis, and
procedure. Make sure to have it well organized, labeled and use a straight edge for all straight lines.
In many labs it is necessary to collect numeric data or quantitative data. However make sure that
you also collect qualitative data for every experiment.
• Quantitative Observations:
Make sure to record all data to the correct number of significant figures. Be careful about
remembering to estimate the last digit for all non-digital measuring devices. Be sure to list units
and include all labels.
• Qualitative Observations:
When making observations, please consider the following
*Sight* *Smell* *Hearing* *Time* *Manner*
• Calculations:
Your calculation section must be very carefully labeled. It needs to be clear where the numbers
have come from (what part of the experiment). Make sure to use sig figs, labels and units. You will
be expected to use dimensional analysis and/or the 4-step method for any appropriate calculations.
• Results:
Your results are the summarized presentation of the data you have collected. This may take the
form of a graph. You do not need to explain the results. Just present them. Often times the results
represent an answer to the hypothesis and/or purpose. If the question is, “What color is Amy
Pond’s hair?” the answer is “Red.” You do not need to discuss red hair, or discuss if Amy is better
than Rose. The results are just the answer to the question. The results of an error analysis are often
presented here.
• Discussion:
The discussion is where you will write about what happened during the lab. There is a pretty good
description of how to write the discussion in “How to Write a Lab Report.” Please remember to be
succinct and to write in the past tense. Do not use any personal pronouns.
• Post-Lab Questions:
Most of these questions attempt to measure what you learned during the experiment. Often times
the Post-Lab questions will reflect the kind of work done in the calculations section.
Materials:
Sodium bicarbonate
1/3 of 1/4 of a FRH
Wooden splint
0.2 M Copper II sulfate
Salicylic acid
Solid iodine
Beaker
Glycerine
0.2 M HCl
0.50 M Acetic acid
Ammonium chloride
1.0 M HCl
0.2 M NaOH
Ethanol
evaporation dish
Test tubes
Potassium permanganate
Thymol blue
17
Solid Potassium iodide
3% H2O2
Magnesium ribbon
Methanol
Digital thermometer
Ice
Test tube rack
Super saturated sodium acetate
Disposable pipette
Pre-Lab Questions:
1. Provide an example and brief desctiption of each of the following concepts from daily life.
a.
b.
c.
d.
e.
f.
2.
a.
b.
c.
Conservation of mass
Exothermic reaction
Endothermic reaction
A reaction that causes a change in color
A reaction that produces a gas
A reaction that produces a solid from two aqueous solutions
g. A reaction that produces a change in smell
h. Evaporation
i. Sublimation
j. Solidification
k. A reaction between an acid and a base
l. A reaction that produces light
What do you think would happen if a match were placed into a container full of….
Carbon dioxide
Hydrogen gas (assume there is a small amount of oxygen present)
Oxygen
Procedure:
Part 1: Conservation of mass
1. Obtain small Ziploc bag and determine its mass to the 0.01 grams. Record data.
2. Place ~5.0 grams of sodium bicarbonate into Ziploc and record the new mass of the bag. Shake
powder into one corner of bag.
3. Obtain 20 mL of 1.0 M HCl in small beaker. Add 4 drops of Thymol Blue to beaker.
4. Obtain disposable pipette and record the mass of the pipette. Fill pipette with the colored acid.
Record mass of pipette with acid.
5. Carefully place large pipette into Ziplock bag. “Burp” Ziploc bag, being very careful to not spill
any of the acid into the powder already in the bag.
6. With the bag sealed squeeze the pipette to mix acid with the sodium bicarbonate.
7. After the reaction is completed (things stop changing) carefully open bag and lower a burning
splint into the bag. Observe any changes.
8. After the reaction is complete record the mass of the bag and its contents.
9. Check data table for completeness, make requested predictions.
Part 2a: Exothermic reactions
1. Obtain a digital thermometer and Lab Quest. Plug Lab quest into wall outlet. Plug thermometer
into channel 1 on the Lab Quest. Confirm that you are able to read the temperature.
2. Place lab quest into graph mode. Set data collection for 1 sample/s for 600 seconds.
3. Obtain 1/3 (this is actually 1/12th of a full FRH) of a packet of a powder from a FRH. Place
powder into a test tube. Set test tube into a test tube rack.
4. Measure 10mL of water using a graduate cylinder record the temperature of the water.
5. Place thermometer into test tube. Hit the green start button. Then quickly pour the water into
the test tube. Observe graph on Lab Quest. Make sure to record maximum temperature and time
to reach this temperature. .
Part 2b: Endothermic reactions
1. Please repeat the procedure used in the exothermic reaction but instead of using the FRH powder
use 5 grams of ammonium chloride. Please use 5mL of water instead of 10mL.
Part 3: Change in color
1. Obtain ~5mL of 3% hydrogen peroxide and place into a test tube.
2. Obtain 10-15 crystals of potassium iodide. Place this into a second test tube.
3. Carefully observe both chemicals.
4. Obtain a burning splint and shake the flame out until it is smoldering.
5. Pour the liquid into the test tube which contains the crystals. Observe and note any change in
temperature.
18
6. After the reaction has started to slow down seal the top of test tube with a parafilm covered
finger. After a moment uncover and lower the smoldering splint into the test tube.
7. Set reaction aside and observe again after about 10 minutes
Part 4: Production of gas
1. This part is very similar to the above reaction. Use 2.0 mL of 1.0 M HCl in one test tube and
1cm of magnesium ribbon in the other test tube.
2. While the reaction is most vigorous, cover the test tube with parafilm and a finger. After a
moment, use a burning splint.
Part 5: Production of a smell
1. Place 50mL of water into a small beaker and heat on a hot plate.
2. Tightly wrap the top of a test tube with paper towels making sure to leave 1cm of glass at the top
of the test tube exposed. Place one Teflon boiling chip into test tube.
3. Add 2.0 mL of methanol to test tube. Smell test tube using a wafting technique.
4. Place 0.5 grams of salicylic acid into the test tube. Smell test tube using a wafting technique.
5. Add 3 drops of concentrated sulfuric acid. This is a very dangerous chemical please ask your
instructor how to do this safely.
6. Using a pipette wet the paper towel wrapped around the top of the test tube, this will help keep
the top of the test tube cool.
7. Heat test tube in hot water bath for about 15 minutes.
Part 6: Precipitation
1. Place about 5.0 mL of 0.20 M sodium hydroxide into a test tube.
2. Fill a small pipette with 0.20 M copper II sulfate.
3. Add the copper II sulfate to the test tube which contains the sodium hydroxide one drop at a
time. Label and set aside to observe test tube the next day in class.
Part 7: Production of flame (this reaction must be done in the hood with the fan on and the sash lowered)
1. Work in a group of 4 students.
2. Place 1 grams of finely divided potassium permanganate into an evaporation dish.
3. Use a scupula to pile the potassium permanganate into a volcano like cone.
4. Place evaporating dish into the hood. Make sure the fan is turned on.
Stop!!!
Is the evaporating dish in the hood? If it not go back to step 4. This is not a joke.
5. Make sure the evaporating dish with potassium permanganate is in the hood.
6. Using a small pipette obtain less then 1.0mL of glycerin.
7. Place 5-8 drops of glycerin on the center of the pile of potassium permanganate.
8. Measure length of time from the start of the reaction until the end of the reaction. Observe the
reaction.
Part 8: Acid/base Reaction
1. Note any changes in color throughout this experiment.
2. Place 2 mL of 0.2 M HCl into a test tube.
3. Place 3 drops of Thymol Blue into test tube and agitate test tube ask teacher to demonstrate
technique to be used.
4. Using a small pipette carefully add 0.2 M sodium hydroxide drop wise into the test tube. Count
each drop. COntinously agitate using demonstrated technique.
5. Take note of the number of drops added each time the color changes in the test tube.
6. When the Thymol Blue reaches a yellow color, stop adding drops.
7. Measure the total volume of the final mixture.
19
Part 9: Change of state (The procedure for this purposely not developed)
a. Evaporation (drops of ethanol onto paper towel wrapped thermometer)
b. Sublimation (Demo? Iodine heated to cool on watch glass)
c. Solidification (Na acetate in test tube with thermometer)
Data:
You must always write your observations, in ink, into your laboratory notebook. This is a permanent record of the
laboratory work that you do. Do not write any data in your lab manual.
Part 1: Conservation of mass
Mass
Mass of bag
Mass of Sodium bicarb and bag
Mass of Pipette
Mass of pipette and acid
Mass of everything before rxn
Mass of everything after rxn
Predictions
Final mass of bag
Final color of bag
Initial Observations
Sodium carbonate
Acid
Acid and thymol blue
Part 2a: Exothermic reactions
Observations during reaction
Initial temperature of water
Final temperature of water
Please sketch the graph created on the Lab Quest.
Part 2b: Endothermic reactions
Observations during reaction
Initial temperature of water
Final temperature of water
Please sketch the graph created on the Lab Quest.
Part 3: Change in color
Observations
Hydrogen peroxide
Potassium iodide
Chemical reaction
Splint placed in test tube
Final observation
Why do you think the color changed again after the reaction stopped?
Part 4: Production of gas
Observations
Hydrochloric acid
Magnesium ribbon
Chemical reaction
Splint placed above test tube
Why do you think the burning splint caused the change?
20
Part 5: Production of a smell
Observations
Smell of methanol
Smell of methanol and salicylic
acid
Chemical reaction
What do you think it might smell
like when it is done?
Why are boiling chips used?
Why is the wet paper towel used?
Part 6: Precipitation
Observations
Sodium Hydroxide
Copper II sulfate
Chemical reaction
Mixture the next day
Part 7: Production of flame (this reaction must be done in the hood with the fan on and the sash lowered)
Observations
Potassium permanganate
Glycerin
Chemical reaction
Product of chemical reaction
Time to complete reaction in seconds
Part 8: Acid/base Reaction
Observations
Hydrochloric acid
Thymol Blue
Thymol blue and HCl
# of drops to1st color change
# of drops to 2nd color change
# of drops to 3rd color change
Color
Color
Color
Part 9: Change in state (The procedure for this purposely not developed)
a. Evaporation (drops of ethanol onto paper towel wrapped thermometer)
b. Sublimation (Demo? Iodine heated to cool on watch glass)
c. Solidification (Na acetate in test tube with thermometer)
Qualitative Observation
Many students often do not make high quality observations, as a result their qualitative data is not
detailed or complete. It is essential to make qualitative observations while the experiment is being
conducted. It is not possible to make these observations later. Do not create a situation where one
lab partner writes down the observations and the other partner does the lab work. Each person
working on a lab is responsible for making their own observations. When you make the data table
in the laboratory notebook it may be a good idea to leave more space between lines so as to not
limit the quantity of observations made.
Calculations:
Make sure to show all work and use labels. You will also be required to use sig figs (you may have
to learn how to use sig figs first. It is fun so you have this to look forward to). It is also critically
important that the calculations section is organized and it is completely clear and easy to
21
understand. Each time you do any math, even simple adding and subtracting of data, it is essential
that the work is shown and it is clear where the numbers (data) came from and what calculation was
used to process this data.
The calculations for this lab are straightforward and easy to organize. The data table has been
arranged and labeled in such a manner to make labeling the calculations easier to follow. This is
something that you should repeat in the future when you are making your own data tables. Please
note how the equations are listed and what is being solved for is very easy to understand.
Results:
Part 1: Conservation of mass
Initial mass of everything ________ Final mass of everything_______
Part 2a: Exothermic reactions
Part 2b: Endothermic reactions
Total change in temp________
Total change in temp________
Part 3: Change in color
Color before reaction_______ Color during reaction________ Color after reaction is
complete_____
Part 4: Production of gas
What happened when burning splint was placed over test tube _________
Part 5: Production of a smell
What does the product smell like ?__________
Part 6: Precipitation
Color of product________________
Part 7: Production of flame
Time to first flame________Color of the flame___________
Part 8: Acid/base Reaction
Each color observed_________, ___________, ___________, ____________
Discussion:
This discussion will focus exclusively on Part 1 of this lab. The Law of Conservation of Mass is a
fundamental law of chemistry. Matter cannot be created or destroyed due to a chemical reaction.
This means that the mass of the products of a reaction must equal the mass of the reactants. Discuss
if the data in this lab supported or did not support the Law of Conservation of Mass. Discuss why it
did or did not. Make sure to use the data and results to support any statements.
In addition to the above please address all other discussion topics in the section of your lab manual
called “how to write a lab report.”
Post-Lab Questions:
1. Sometimes the best way to understand how something works, or what it is made of, is to
take it apart. Did you take any of the chemicals apart? If so which ones?
2. Provide examples of state changes observed in this experiment.
3. When a chemical is taken apart, what happens to these parts?
4. How is this different than when a chemical experiences a state change?
22
Experiment 2
Techniques and Measurements
Purpose:
The methods and equipment used to measure volume and mass will be studied and used to find
density.
Hypothesis: State a hypothesis regarding which glassware will produce the most accurate
result.
Background:
Volume
The SI unit for volume is 1 cubic meter, 1m3. This, however, is too large a volume for use in most
laboratory settings. Volume may also be measured in cubic centimeters. 1cm3=milliliter. Volume
is usually measured in units of liters (L) or milliliters (mL) (1000mL=1L). Volume may be
measured with various degrees of accuracy depending upon the equipment used.
Pipette
A pipette is a glass tube, usually with a bulb in the middle and an index mark on the upper stem.
It is designed to accurately deliver a specific volume of liquid. Pipettes are available in various
sizes, 1.00, 2.00, 5.00, 10.00, 25.00mL. Pipettes are always to be used with a rubber bulb or
pipette filler. Never use your mouth. They are accurate to the 0.01mL.
↑
Fill line is here
Graduated cylinder
A glass tube standing on the flared
base with uniform calibration marks
along the tube. The TC accuracy of a
10.00mL and 25.00 mL +/- 0.02. If
the Cylinder is used to deliver a
solution the accuracy drops to +/- 5%
23
Burette
A burette is a calibrated glass tube with a stopcock and
tip on its lower end. A burette is designed to measure
how much liquid has been delivered. The amount
delivered is read from the difference in the graduations
on the tube. When used and read correctly, a burette is
accurate to the 0.01mL. Be careful when using a
burette, please note that for instance the water level may
start at 1.00mL when water is drained out of the bottom
it may stop at 4.50mL. In this case 3.50 mL of water
would have been delivered.
Volumetric Flask
A flask with a narrow neck designed to contain a given volume of
liquid up to an index mark. Typical capacities are 100.00mL,
200.00mL, 500.00mL, and 1000.00mL. Accuracy is 0.1%.
The correct glassware for the job
Burettes can be used when the volume needed is not a specific
quantity, but when the actual volume delivered must be known.
An initial volume reading is measured using the bottom of the
meniscus (the curved surface of the liquid) against the lines on the
burette. Some liquid is delivered through the stopcock. A final
volume reading is measured in the same manner as the initial
volume. The difference between the initial and final volumes is the exact quantity delivered.
A pipette or a volumetric flask is used when
a specific quantity of liquid must be
accurately measured. The vessel is filled
until the bottom of the meniscus is level with
the calibration mark.
ß24.52mL
A measuring cylinder is used when an
approximate volume of liquid is needed.
Some beakers and Erlenmeyer flasks have
graduations marked on them. These
graduations are accurate to about 20% and
are only useful for crude estimates.
TC vs. TD
Glassware that is marked TC is only accurate when measuring materials in the container. Think of
TC as “To Contain.” While glassware marked TD only accurate measures volumes delivered.
Think of TD as “To Deliver.”
Mass
24
There are two primary kinds of tools to determine mass. First is a mechanical balance which
determines the mass of an object by “balancing” against the mass of known objects. An example of
this is called the triple beam balance. The other method of determining mass uses an electronic
device referred to as a top loading balance, digital balance, or analytical balance. These balances
use electricity and magnetism to determine the mass of an object.
Triple Beam Balance
It is important that before loading that it is correctly zeroed. Read the needle at eye level to avoid
the effects of parallax. This balance will mass objects to approximately 600g and reads to the
0.01g. It is necessary to estimate the final digit.
Electronic balance
The mass of samples can be determined quickly and accurately using an electronic balance. The
balance can be read to the nearest 0.01g or 0.001 g depending on the model. The last digit of a
digital reading is an estimate. Before loading the balance it is essential that it reads zero. If it does
not, press the “tare” button gently.
Procedure for finding mass
When the mass of a solid object is to be determined, the mass of an empty weighing bottle is first
measured and recorded. The object is then placed into the weighing bottle and the mass of the
bottle and the contents are remeasured. Again the mass is recorded. The difference in the two
masses is the mass of the object. This is called weighing by difference.
When the mass of a chemical substance, a reagent, is needed for an experiment, the technique of
“weighing by difference” is used. The reagent (liquid, solid, or solution) is placed into a container.
The total mass of the container and its contents is measured. The mass is recorded into a notebook.
Some material is tipped out into a beaker or the glassware where it will be used. The mass of the
container and its contents are remeasured. Again, the mass is recorded into a notebook. The
difference in the two masses is the amount of material transferred to the beaker.
Laboratory instructions generally say that the approximate mass of a sample is to be placed into a
container and its mass determined. This means that it is not necessary to measure out exactly the
specific amount stated. The amount used should be within about 0.05g of the amount stated, but the
actual mass of the amount used should be measured to the nearest 0.01 or 0.001g.
The mass of any object should be measured with the object at room temperature, otherwise
convection currents make it impossible to obtain a stable balance reading. The sample should not
be wet or damp. Evaporation makes it difficult to obtain a stable reading. Furthermore, moisture
increases the apparent mass of the sample and leads to error when the degree of moisture is not
known. Moisture can also lead to unwanted corrosion of the balance pan.
The tolerances on the masses in the balance are extremely small; therefore a mass obtained on one
balance will be very close to that found on another balance. Nevertheless, it is best to use the same
balance for all measurements during the course of an experiment.
25
Density
Density is defined as mass per unit volume, d=m/v. It is an important physical property of a
substance. The SI unit of density is the kg/m3. These units are somewhat hard to work with in the
classroom setting, thus the usual units quoted are g/mL or g/cm3. Density was one of the earliest
methods of sample identification. The density of a material can give information about its identity.
Significant figures
The accuracy of the results is limited by the accuracy of the equipment used. The significant
figures show the accuracy of the results of any experiment. The last figure in each measurement is
an estimate thus it is “in doubt.” Each measurement has a range or reliability. Thus, an object that
has a mass of 3.672g when measured on an electronic balance with an accuracy of 1mg has a range
of 3.672g +/- .001g, which is 3.671g to 3.673g
When more than one measurement is used in a calculation, the ranges of all measurements must be
considered and the rules of significant figures must be used in the calculations. For example, in
determination of the density of an object, the mass is found to be 6.45 +/-.01g and volume is 1.43ml
+/-0.1mL. The density range is therefore 4.47g/mL to 4.55 g/mL. (6.44g/1.44mL) to
(6.46g/1.42mL). To obtain the smallest possible density take the smallest mass over the largest
volume. To obtain the largest possible density take the largest mass over the smallest volume.
Range of values for calculated measurements
When finding density there is both uncertainties in the measurement of mass and in volume.
To find the range of volume it is necessary to consider what the greatest possible volume is and
what the least possible volume could be. Vá = Vfinal á - Vintial â and Vâ = Vfinal â - Vintial á . Thus it
is necessary to factor in the uncertainty of the denominator and the numerator. To do this it is
necessary to determine what is the largest and smallest possible densities that could result.
Large mass
ámass
Largest density =
áD=
Small vol
âvol
Small mass
âmass
âD=
Large vol
ávol
Thus if the measured mass is 47.58 grams +/- 0.01 grams the largest possible mass is 47.59 and the
smallest mass is 47.57 grams. If the measured volume is 28.5 mL +/- 0.5mL then the largest
volume is 29.0 mL and the smallest is 28.0 mL.
47.58 g
ámass
47.59 g
Density=
=1.67g/mL
áD=
= 1.70 g/mL
áD = âvol
28.5 mL
28.0 mL
Smallest density=
âD = âmass
47.57 g
âD=
ávol
29.0 mL
Thus the density is------------- 1.67 g/ml with a range of 1.64g/mL to 1.70 g/mL
or it could be reported as ----- 1.67g/ml +/- 0.03 g/mL.
26
= 1.64 g/mL
Error analysis
It is possible to compare the experimental results to that of the literature value (theoretical value).
This is an important calculation for it helps illustrates the accuracy of the experiment. Percent error
is calculated as follows
| theoretical – experimental |
theoretical
* 100 = Percent error
Water density
The density of water is the basis of many measurements. Water is often considered to have a
density of 1g/ml, thus 1000g = 1000mL or 1kg=1liter. Rarely does water have a density of 1g/mL.
The density of water at 25.0 oC is 0.997 g/mL. Only at 3.98 oC is water 1.00 g/mL.
Pre-Lab Questions:
1. How many significant figures are in each of the following measurements? Write both the
measurement and answer.
a. 42.78mL
f. 6,000mL
b. 4.0767g
g. 1010.0g
c. 50.00mL
h. 350mL
d. 0.005g
i. 100.000g
e. 0.0610L
j. 300ml
2. The mass of an object is measured on an electronic balance (accurate to 0.001g). Its mass is
recorded as 34.194g. The volume of the object is measured using the displacement of water
in a 25mL measuring cylinder. The initial volume of water was 10.10 mL (before the object
was submerged). The volume after the object was placed in the cylinder was 22.40mL.
a. What is the range of possible masses for the object?
b. What is the range of possible values for the volume of the object? The cylinder is only
accurate to +/- .05mL.
c. Calculate the density of the object (to the correct number of significant figures) based
upon the reported mass and volume (ignore ranges)
d. Calculate the density of the object at the lowest possible mass and the highest possible
volume
e. Calculate the density of the object at the highest possible mass and the lowest possible
volume.
f. What is the range of densities based on parts d and e?
g. State the density of the object and its range.
3. An iron object was determined experimentally to have a density of 7.781 g/mL. What is the
percent error? Please use the Sargent Welch periodic table to find the theoretical value.
Procedure:
Density of water
1. Determine the mass of a clean, dry weighing container (the smallest beaker in your drawer).
Fill the burette to a volume between 1.00mL and 5.00 mL. Use a burette to measure between
10 and 15 mL of water into the weighing container. Record the initial and final readings from
the burette to the nearest 0.01 mL Calculate the actual volume of the water delivered into the
27
weighing container to the highest number of significant figures possible. Do not attempt to
start the burette at exactly 0.00mL. Also, do not attempt to measure an exact number of mL, it
is far more important to know exactly what volume was delivered. Determine the mass of the
weighing bottle with the water in it (remember to use the same balance throughout the
experiment). Use this data to calculate the density and range of water.
2. Repeat part 1 using a 10.00 mL pipette instead of a burette. Compare your results with part 1,
making note of the number of significant figures obtained by measuring the volume with a
pipette. There are no initial and final readings on a pipette; it can only deliver one volume. Use
this data to calculate the density and range of water.
3. Repeat part 1 using a 10mL measuring cylinder instead of the burette. Compare your results
with part 1. Make note of the number of significant figures obtained by measuring the volume
with a measuring cylinder. When using a measuring cylinder to deliver a volume it is filled to
a certain level and this level is recorded. The cylinder is then completely emptied. Do not
attempt to obtain initial and final readings. Use this data to calculate the density and range of
water. Do not attempt to fill cylinder to 10.00mL.
Density of unknown liquid
4. Obtain 20mL of the unknown liquid in your Erlenmeyer flask. Write down the liquids ID #.
Make observations regarding the liquid. Determine the mass of a clean dry weighing
container. Use a pipette to place 10.00mL of the unknown liquid into the weighing container.
Find the mass of the container and liquid.
5. From data, established in step 4, determine the density and range of the unknown liquid.
Density of unknown object
6. Obtain a “density unknown” from your instructor. Write down the solids ID #. Make
observations regarding the solid. Determine the mass for the object. Using a measuring
cylinder, determine the volume of the object by measuring the volume of water that the object
displaces. To do this, partially fill the cylinder with water. Record the volume. By tilting the
cylinder slowly slide the density unknown into the water. BE CAREFUL NOT TO BREAK
THE BOTTOM OF THE CYLINDER WITH THE OBJECT. TILT THE CYLINDER!
Record the volume again. Use this data to calculate the density and range of unknown object.
Data and Observations:
Your data should show all masses and volume measurements with the correct units and to the correct
number of significant figures. Do not forget to also include any qualitative observations. You will
need data tables for all sample of water, unknown liquid and solid. Example data tables follow.
Water density determined using a burette
Mass of beaker
87.01g
Mass of beaker and water
97.03g
Mass of water
10.02g
Initial burette reading
0.15mL
Final burette reading
10.18mL
Volume of water
10.03mL
Water density determined using a pipette
Mass of beaker
Mass of beaker and water
Mass of water
Volume of water
Calculations:
Calculate the following. Use the information about the accuracies of the equipment to calculate the
ranges. Express the results of each calculation to the appropriate number of significant figures and
show the units in each case.
1. The density of water from part 1 (with range and percent error)
2. The density of water from part 2 (with range and percent error)
28
3.
4.
5.
6.
7.
8.
The density of water from part 3 (with range and percent error)
The density of liquid from part 4 (with range)
The mass of the object from part 6 (with range)
The volume of the object from part 6 (with range)
The density of the object from part 6 (with range)
From the possible metals and their densities, calculate a percent error for your metal
Results:
Make clear statements regarding the results of each density calculation. Show the possible range in
each case and the % error.
Glassware used
Density
Range
% error
Volumetric pipette
Etc.
Possible element
Unknown liquid
Unknown object
Discussion:
In addition to the standard discussion topics discuss the following: The density of water was
measured three times. A different piece of glassware was used each time to measure the volume.
Based on your percent errors and range, which measuring tool is likely to provide the best accuracy
and precision? If one result is different from the other two, explain how the glassware used affected
that result. Explain your answer using the accuracy of the equipment used. The density of water is
1.000 g/mL at 3.98oC. Explain why your results were not exactly the same as the literature value.
Provide a possible identity of your unknown object based both qualitative and quantitative evidence
Post-Lab Questions:
1. What is the density of a sample with a mass of 30.425g and a volume of 16.8mL? Report the
answer to the correct number of significant figures.
2. The mass of an empty weighing bottle is 22.814 g when measured on a digital balance. The mass
of the bottle containing a solid object is 71.115 g. The object’s volume is determined by the
displacement of water in a 100mL (+/-0.1mL) measuring cylinder. The volume reading before the
object is immersed is 49.5mL, and after immersion the reading is 86.0mL.
a. What is the density of the object?
b. How many significant figures should the answer have?
c. What is the range of the answer?
3. A piece of household aluminum foil has a mass of 0.257g and measures 3.96 cm*11.2cm. If the
density of aluminum is 2.70g/cm3, how thick is the aluminum foil? (Hint: the mass and density are
given, so think carefully about how to calculate the volume of the foil.)
References
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102
laboratory manual. Star Publishing Company: Belmont, California
29
Experiment 3
Emission Spectroscopy
Purpose:
To compare predicted emission spectra with experimental emission spectra.
Background Information:
Electrons control the chemistry of the elements. In order to understand the role of the electron it is
necessary to know something about the arrangement of the electrons in the atom. This information
can be obtained by studying the way in which atoms absorb and emit light.
It is first necessary to know something about the properties of light. Light is a form of energy
known as electromagnetic radiation. One of the ways to describe light is to say that it has the
properties of a wave. Light can be thought of as a wave traveling in a straight line (Figure 1).
Figure 1
Direction of travel
Light in a vacuum travels at 2.998x108 meters per second (or about 185,000 miles per second). The
frequency is the number of waves that pass a given point during a given unit of time (usually one
second). Light waves are also characterized by wavelength. Wavelength is the distance between
crests of the waves (Figure 2).
Direction of travel
Figure 2
If the wavelength is short (λ1) then the number of waves which pass a given point per unit of time is
large. This creates a high frequency wave (ν 1). If the wavelength is long (λ2) then the number of
waves which pass a given point per unit of time is small. This creates a low frequency wave (ν 2).
Since the speed of light is independent of wavelength or frequency, it is possible to summarize the
behavior of light by the following equation:
30
λ* ν = c
(Equation 1)
Where λ = wavelength
ν = frequency
c = speed of light (3.00 x 108 m/s)
Equation 1 indicates that when the wavelength is short (λ1), the frequency is high (ν 1). When the
wavelength is long (λ2), the frequency is low (λ2). Equation 1 can be used to calculate the
frequency of light if the wavelength is known. The frequency of light having a wavelength of 651
nm is calculated below. First we solve Equation 1 for frequency (see Equation 2).
(Equation 2)
ν= c
λ
Convert the wavelength from nanometers to meters. Note that 1.00 nm = 1.00 x 10-9 m.
651 nm
1.00 x 10-9m
(Equation 3)
= 6.51 x 10-7 m
1.00 nm
Then substitute in Equation 2 and calculate frequency:
3.00 x 108 m/s
(Equation 4)
ν= c =
= 4.61 x1014 sec-1
-7
6.5100 x 10 m
λ
Note
1
= sec-1
sec
This means that frequency is measured as wave crests per second. In the above example, a crest of
the light wave passes a given point 461,000,000,000,000 times per second (461 trillion times per
second).
It is interesting to note that the only difference between radio waves, visible light and x-rays is the
wavelength of the electromagnetic radiation. This is illustrated in Figure 3.
Radio, TV
106
104
Visible light Ultraviolet
Infrared
X-rays
Radar
102
Gamma rays
Cosmic rays
100
10-2
10-4
10-6
10-8
10-10
10-12
10-14
Low energy
High energy
Wave length (m)
Figure 3
If a narrow beam of white light is directed at a prism or diffraction grating, the white light is
separated into its component parts. A continuous spectrum appears which shows that the white
light is composed of light of all wavelengths.
Low energy
V
380nm
B
450nm
G
495nm
Y
570nm
O
590nm
R
620nm
750nm
Visible Light Spectrum
High energy
31
Schematic diagram of emission spectroscopy apparatus
Figure 4
If a narrow beam of light from a hydrogen discharge tube is directed at a prism or diffraction
grating, a continuous spectrum is not obtained. Instead, a series of lines appear. These lines
correspond to light of a certain wavelength. If a sodium lamp is used instead of a hydrogen
discharge tube, a different series of lines appears. The lines of emission spectra are unique for each
element. The study of emission spectra is called emission spectroscopy (Figure 4).
A very simple experiment illustrates that emission spectra are caused by electrons. The spectra of
He, Li, Be2+ and B3+ are quite similar in spite of the fact that their nuclei are different. These four
species all have two electrons. The fact that the emission spectra will be different when the number
of electrons changes (regardless of the nuclear composition) indicates that the electrons are
responsible for the emission spectra.
In 1913, Niels Bohr proposed a theory which yielded a fairly good explanation for the emission
spectra. In time, this theory was shown to be partially incorrect, but some of Bohr’s original ideas
have been retained in the modern atomic theory.
The current theory of electronic structure assumes that an atom consists of a positively charged
nucleus with one or more negatively charged electrons at a distance from the nucleus. This theory
assumes that the electrons can have only certain energy levels (or specific distances from the
nucleus). We say that the energy of an electron is “quantized” or that electrons can exist only in
certain energy levels. The only way an electron can go to a higher energy level is to absorb energy.
This concept is illustrated in Figure 5.
When an electron absorbs energy,
it is excited and is propelled into a
higher energy level. Higher
energy levels are generally farther
away from the nucleus than lower
energy levels. An exact amount of
energy is needed to move an
electron from the 2nd to 3rd energy
level. A lesser amount of energy
can not raise the electron to the 2½
energy level because there is no
energy level between the 2nd and
3rd.
Electron absorption of energy
Figure 5
32
When an electron drops from a higher to a lower energy level, energy is released from the atom
(Figure 6).
Electron emission of energy
Figure 6
In this case an excited
electron returns to its ground
state (the energy level
normally occupied by the
electron). Energy is released.
Because the energy between
the 2nd and 3rd energy level is
always the same, an electron
which drops from the 2nd to
the 3rd level always releases a
predictable quantity of
energy in every atom.
If an electron in energy level 3
drops to energy level 2, it will emit
energy which is exactly equal to the
difference in energy between levels
2 and 3 (see Figure 7).
Energy emitted =
ΔE = Ef - Ei = E2 - E3 (Equation 5)
Ef = energy of the final level
Ei = energy of the initial level
Electronic energy levels and emission of energy
Figure 7
Please note that in this case ΔE will have a negative value since the energy of the e- at E3 is greater
than the energy of the e- at E2. This indicates that the electron has lost energy, confirming that the
energy is emitted or given off.
The electron could have fallen to energy level 1 and emitted a greater amount of energy (see Fig. 7)
’
ΔE’ = E1 - E3
(Equation 6)
Where ΔE > ΔE
According to this theory it is not possible for the electron to fall to some intermediate level between
E2 and E1. The energy of the electron is quantized. The electron can therefore occupy only certain
levels as dictated by the principal energy level, n.
An electron can also absorb energy and be excited to a higher energy level (see Figure 8).
33
It is interesting to note that the energy absorbed
must exactly numerically equal the difference in
energy of levels 3 and 4.
Energy absorbed =
ΔE’’ = Ef - Ei = E4 - E3
(Equation 7)
ΔE’’ is positive since the energy is absorbed by the
electron. Energy level n=4 is greater than energy
level n=3.
Electronic energy levels and absorption of energy
Figure 8
If the exact amount of energy is not absorbed the electron will not make the transition. It is not
possible for an electron to absorb enough energy to be excited to an intermediate level because there
are no intermediate levels—the energy levels of the electrons are quantized.
With these concepts it is fairly easy to explain why the emission spectrum of hydrogen is a series of
lines. The electron of the hydrogen atom in the discharge tube is subjected to bombardment of
energy when the tube is powered on. When this occurs some of the electrons absorb energy and
these electrons are forced to leave their ground state to occupy a higher energy level (excited state).
When the excited electron falls to a lower level, a certain amount of energy (ΔE) is emitted.
Where
ΔE = Ef - Ei
Ef = energy of the e- at the final level
Ei = energy of the e- at the initial level
(Equation 8)
The energy is emitted by each atom as a quantum of light. In 1900, Max Planck described the
relationship between the energy and frequency of a quantum of light (and the energy and frequency
of any electromagnetic energy).
ΔE = h * ν
(Equation 9)
Where
ΔE = Energy of emitted light (J)
h = Planck’s constant (6.626x10-34 J*sec)
ν = Frequency of light (sec-1)
If Equation 8 and Equation 9 are combined, Equation 10 is obtained.
Ef - Ei = ΔE = h * ν
thus
Ef - Ei = h * ν
(Equation 10)
This mathematical expression means that when an electron drops from a high energy level to a low
energy level, a quantum of light, a flash with a fixed frequency and wavelength, is emitted. Since
only certain energy levels are allowed, only certain electronic transitions can occur, and the
emission spectrum of hydrogen consists only of certain wavelengths. Thus we should see an
emission spectrum which consists of light flashes of a certain wavelengths. Since a tremendous
number of atoms emit flashes simultaneously, we see instead, lines having wavelengths the same as
those of the flashes.
34
Bohr’s great achievement consisted of constructing a theoretical model of the hydrogen atom and
calculating its properties which fitted the experimental data exceedingly well. His model was found
to be excellent for the hydrogen atom. Unfortunately, it failed for any other atom.
Bohr’s model began with fundamental constants of nature such as the speed of light and the charge
and mass of the electron, and the assumptions outlined above. He arrived at the Bohr-Rydberg
equation for the energies associated with electronic transition.
ΔE = Ef - Ei = -RH
(
1
Nf 2
ΔE = -RH
(
1
Nf 2
Where
)
-
1
=h*ν
Ni 2
or in a more simplified version
-
1
Ni 2
)
(Equation 11)
nf = the principal energy level of the final level
ni = the principal energy level of the initial level
RH = Bohr-Rydberg constant, 2.178x10-18 J (calculated from fundamental constants)
A note regarding positive and negative signs:
When the sign on energy is negative that means that the energy is leaving the system. If the sign is
positive that means the energy is being added to the system. So for instance when a match burns
the match is loosing energy and thus from the perspective of the match’s system the energy would
have a negative sign. A negative sign on frequency or wavelength should be avoided there is no
such thing as a negative frequency or a negative wavelength.
Pre-Lab Questions:
1. Find the frequency for the following wavelengths.
1a. 6.55 x10-7 m
1b. 4.84 x10-7 m
1c. 4.32 x 10-7 m
2. Find energy in joules for the following frequencies of light.
2a. 4.58 x 1014 sec-1
2b. 6.20 x 1014 sec-1
2c. 6.94 x 1014 sec-1
3. Find energy emitted or absorbed in Joules for the following electron energy level transfers. Was
energy absorbed or emitted?
3a. ni = 1 to nf = 2
3b. ni = 3 to nf = 4
4. Find the frequency of light given off for the following electron energy level transfers.
4a. ni = 3 to nf = 1
4b. ni = 5 to nf = 3
Materials:
Spectroscopy apparatus, discharge tube and power supply
35
Procedure:
In this laboratory activity you will replicate Bohr’s work and evaluate the Bohr model of the
hydrogen atom. Use the Bohr-Rydberg equation to calculate the energies associated with the
transition for…
ni = 3 to nf = 2
ni = 5 to nf = 2
ni = 4 to nf = 2
ni = 6 to nf = 2
You will make observations of hydrogen’s line spectra. Then from the observed wavelengths you
will calculate the energy associated with each spectral line. You will then compare the line energy
data from the experiment with theoretical values.
To complete this laboratory activity, view into the spectroscope and make careful observations.
Record the data and complete the calculations. Then compare the results of your experimental
calculations with theoretical values.
Data:
Line
Observed Wavelength of spectral line (cm x10-5)
Color of line spectra and other
qualitative data
Red
Green
Blue
Violet
Calculations:
Calculations Table 1: (this is an example of how you may wish to report your calculation, you
will still need to show all of your work)
Line
Convert wavelength (cm
Observed frequency
Observed energy of spectral line
x10-5) to wavelength (m)* of spectral line (sec-1)** (Joules)#
Red
Green
Blue
Violet
* consult Equation 3
** consult Equation 4
# consult Equation 9
The next step is to calculate the theoretical energies associated with the electronic transitions. To
do this, use the Bohr-Rydberg equation (Equation 11).
Calculations Table 2:
(Violet has the highest energy followed by blue, green, yellow, orange, and red has the lowest energy of light)
Electron energy level transition
ni = 3 to nf = 2
ni = 4 to nf = 2
ni = 5 to nf = 2
ni = 6 to nf = 2
Theoretical energy in Joules
36
Likely color
The last step is to match up the observed energies and the theoretical energies and to calculate the
absolute and relative error. Please remember the note regarding negative signs. Determine the
accuracy by calculating the absolute errors and relative errors for the red and violet lines only. The
measured values will be the observed energies from calculation table 1. The theoretical values will
be the theoretical energies from calculation table 2.
Absolute error = E = | m – t |
m = experimental measured value
t = theoretical value
Relative error (accuracy) E/t *100
E = | m – t | (see above)
t = theoretical value
Calculations Table 3: Keep in mind that red light has least energy and violet has the most. You might
want to be careful how you match up your data for error analysis
Line color
Observed energy
Theoretical energy
Observed energy
Relative error
Absolute error
Relative error
Red
Violet
Results:
Line color
Red
Violet
Discussion:
Discuss how this lab met the stated purpose. Discuss how the observed and theoretical energies
matched. If the relative error is less than 5%, the data validates the Bohr model. Discuss one area
for potential follow up experiments. Discuss the relationship between light and energy, frequency
and wavelength. How does this lab support the observation regarding the quantized nature of
electron energy levels?
Post-Lab Questions:
1. Do some research and find the color spectrograph of a fluorescent light. What is the
difference between the light the fluorescent bulb emitted and light emitted by electrons of
excited atoms in a hydrogen discharge bulb?
2. Calculate the wavelength of radio waves transmitted by KMTT. The frequency of the these
waves is 103.7 MHz (103.7 x106 sec-1).
3. Discuss why the orange sidewalk lights are called “Sodium Lamps.” Provide a citation for
your source of information.
References:
Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson
College Press: Swannanoa, NC.
Silberberg, Martin. (1996) Chemistry: The molecular nature of matter and change. Mosby: New
York, NY
37
Experiment 4
Rapid oxidation of metallic fuels with color additives
“Sparklers”
Purpose:
To determine effective procedures to overcome the difficulties in the creation of sparklers,
especially in how to get a good mixture that keeps burning and a composition that binds easily and
effectively to the wire or thin wood stick.
Hypothesis:
Predict what your sparkler should look like before it is ignited. Also propose a recipe for the
unknown sparkler based on your observations of the other sparklers.
Materials:
Supplies
Popsicle stick
Rubber cement
Wax paper
Burner
Chemicals
Oxidizer (provides oxygen for rapid oxidation)
Potassium perchlorate
Potassium nitrate
(other nitrates may also serves as an oxidizer)
(do not grind)
(may need to be ground very carefully in a clean mortar)
Fuel (rapid oxidation; reacts with oxygen and produces sparks)
Magnesium powder
Aluminum powder
Iron Filings
Magnalium
(an alloy of magnesium and aluminum)
Coloring agent (because of the electron configuration these metals burn with very specific colors)
Copper II oxide
Barium nitrate
Strontium nitrate
Sodium chloride
Other (other materials that may be used in various sparkler recipes)
Shellac
Parlon
Charcoal
Sulfur
(grind)
Manganese IV oxide
(grind)
Background:
In the fireworks industry, sparks are probably the oldest pyrotechnically produced effect and
can be generated from many different materials. What makes a spark glow? It’s incandescence—
yes, just like an incandescent bulb. When a solid is heated to a high temperature, it incandesces or
glows. Sparks that you see in fireworks are solid or liquid particles that have been ejected from a
burning surface. This surface has been heated to a high temperature (hotter than 1700 degrees C!)
by the chemical reactions and flame.
One of the factors involved in making sparklers is controlling the color. Manufacturers of
pyrotechnics can produce sparklers with only a limited range of colors: they vary from dim red to
orange to yellow to bright white. No other colors are possible from incandescing objects. You can
see this range of colors if you watch the filament of a clear incandescent bulb as you vary the
voltage delivered to it with a dimmer switch.
Another factor important in sparkler production is the longevity of the spark. We don’t want
the sparks to cool too fast or they’ll be duds: they’ll glow brightly at first, but then cool quickly,
grow dim, and be invisible. We want the sparks to stay bright for a longer time, and to do this, they
38
must continue to react and produce heat energy as they fall through the air. This spark reaction is
oxidation, with the spark as the fuel and the air oxygen as the oxidizer. In the sparkler recipes you’ll
be testing, magnesium and aluminum are the fuels:
4Al (high temp) + 3O2 à 2Al2O3 + heat
2Mg (high temp) + O2 à 2MgO + heat
Pre-Lab Questions:
1. Many metal compounds can be identified by the color that is produced when they are
heated. Do some research and determine what colors will be produced when the following
materials are heated (burned)?
A. Copper salts
B. Strontium salts C. Barium salts
2. Many metals can be identified by what there sparks look like. Do a little research and then
draw and describe what the following sparks will look like.
A. Iron Filings B. Magnesium
C. Aluminum
Safety:
•
•
•
•
•
Do not grind potassium perchlorate, KClO4. If it is necessary to grind the potassium nitrate
do it gently with a clean mortar and pestle. They could detonate. Do not grind any of the
ingredients once they have been mixed together! They could detonate.
Do not make a recipe that makes more than 2.0 grams of powder.
If you have powder left over after your sparkler has been made, please return it to the waste
beaker in the fume hood.
Do not place any of the powder in the trash cans.
Wash your hands thoroughly before leaving the lab.
Colored perchlorate recipes:
Red
KClO4
Sr(NO3)2
Magnalium
Sulfur
Green
21%
46%
20%
13%
KClO4
Ba(NO3)2
Magnalium
Sulfur
Copper II oxide
Lilac *
20%
45%
17%
13%
5%
KClO4
KNO3
Sulfur
Charcoal
Al powder
7%
64%
18%
5%
6%
Other potassium perchlorate recipes:
KClO4 **
Mg powder
Sr(NO3)2
30%
50%
20%
KClO4
Magnalium
Copper II Oxide
50%
39%
11%
KClO4 **
Al Powder
Copper II Oxide
50%
39%
11%
KClO4
Al powder
Copper II oxide
60%
30%
10%
KClO4 **
Al powder
Magnalium
60%
30%
10%
KClO4
Magnalium
Sulfur
Copper oxide
Barium nitrate
Sodium chloride
20%
42%
12%
12%
12%
2%
KClO4
Ba(NO3)2
Magnalium
Titanium
Mg Powder
20%
14%
16.5%
33%
16.5%
KClO4
Al powder
Al 80 Mesh
50%
25%
25%
39
Potassium nitrate recipes:
KNO3 *
CuO
Magnalium
Charcoal
12.5%
25%
37.5%
25%
KNO3
Mg Powder
Charcoal
Sulfur
70%
15%
10%
5%
KNO3
Al powder
Charcoal
Sulfur
47%
18%
24%
11%
KNO3
Charcoal
Sulfur
58%
35%
7%
KNO3 **
Mg Powder
Charcoal
Sulfur
Iron Filings
33.8%
45%
6.7%
4.5%
10%
Ba(NO3)2
Al powder
Fe filings
Boric acid
45%
19%
34%
2%
Other Nitrate based recipes
Sr(NO3)2 **
Mg powder
Fe filings
40%
17%
43%
Ba(NO3)2
Iron Filings
Dextrin
Al powder
Charcoal
Boric acid
50%
30%
10%
8%
0.5%
1.5%
Sr(NO3)2
KNO3
Magnalium
Sulfur
Charcoal
??
Ba(NO3)2
Iron Filings
Al flakes
25%
30%
37.5%
2.5%
2.5%
50%
30%
20%
Sr(NO3)2
86%
Shellac dry
14%
For this to work the shellac must
be ground very fine
Procedure:
1. Decide which ingredients and what amounts you will use to create no more than 2 grams
of sparkler powder. Obtain your instructor’s initials on the recipe, indicating that you have
permission to proceed.
2. Make sure each of the ingredients is a fine powder. Lumps decrease the surface area for
reaction, slowing the process (it’ll be a dud). Use a mortar and pestle if necessary, but DO
NOT grind the potassium perchlorate or the potassium nitrate!!
3. Mix the powders VERY WELL together. Make sure you have a uniform texture.
4. Obtain a piece of wax paper and a Popsicle stick from the materials table.
5. Write your name on one end of the stick.
6. Place your powder onto your wax paper.
7. Gently hold up opposite sides of the wax paper to form a thick line of powder down the
center, then place it down on the lab counter.
8. Holding onto the end marked with your initials, dip the stick into the rubber cement so that
one half of the stick has a medium coat of the glue. Use a clean stick to return the excess
glue to the rubber cement container.
9. Immediately press the gooey end of the stick into the pile of powder. Work the stick,
pressing down and pulling up, coating both sides and the edges with powder. You should
have a nick thick coating on the stick. If you have no powder left on the waxed paper then
you are finished.
10. If you have given your sparkler as many coats as you can and still have leftover powder,
return the powder to the proper waste beaker.
40
11. Carefully carry your sparklers over to the fume hood. Place them down on your wax paper
and leave them to dry. Your sparkler must dry completely.
Data: Do not write on the following data table make your own.
Also remember you will
need to have a data table for at least 2 sparklers and a table for the unknown.
Recipe 1
Name
Mass %
Mass (g)
Observations before ignition
Oxidizer
Metal fuel
Color additive
Total
Predicted Charactersitics
Color
Spark
Observations of burning
Color
sparkler
Spark
Other
Unknown
ID #
Oxidizer
Metal fuel
Color additive
Observations
Name
Results:
Connect the color of sparkler to the chemical used in your sparkler. Connect the type of spark to
the metal used in your sparkler.
For instance
Copper produced a _____________ colored sparkler.
The aluminum produced a ___________ colored spark that had a _____________tail.
For the unknown determine the
• Metal fuel
• Oxidizer
• Color additive
Discussion:
None at this time.
Post-Lab Questions:
None at this time.
References:
Beth Eddy
41
Experiment 5
Inorganic Nomenclature
Purpose:
The names and formulae of common anions and cations will be learned. The formulae of
compounds will be written from their names. The name of compounds will be written from their
formulae.
Background:
In order to communicate with other scientists around the world, chemists use a set of standard rules for
naming compounds. The system used is called the IUPAC (International Union of Pure and Applied
Chemists) system. The older way of naming chemicals, which is still in use, utilized a mixture of
various systems, and such names are now called common or trivial names. IUPAC names will be used
as much as possible in this class, but some very common trivial names will also be used so you should
become familiar with them as well. The set of rules should be learned along with a list of names and
formulae for common ions.
Predicting Charge
Column
1A
2A
3A
4A
5A
6A
7A
8A
1B-8B
Family Name
Alkali metals
Alkaline earth metal
Boron family
Carbon family
Nitrogen family
Oxygen family
Halogens
Noble gas
Transition metals
Charge of Simple Ions
1+
2+
B and Al 3+, others vary
Varies
N can be 3-, others vary
21None
Varies however, 2B typically 2+, 3B typically 3+
Binary Ionic Compounds
Nearly all ions are made of a metal and a nonmetal. The easiest compounds to name are binary ionic
compounds. These contain only metal ions (cations) and nonmetal ions (anions) whose charges can be
predicted from the periodic table. If the cation is a metal with a known charge, it is only necessary to
know the names of the elements in order to name the compound. The metal (cation) is named first and
is given its elemental name. The nonmetal (anion) is named next. To name the nonmetal, the root of the
element name is used but given the suffix –IDE. In these simple cases it is not necessary to use prefixes
or Roman numerals as the charges on all ions are known.
Examples: NaCl is called sodium chloride
K2O is called potassium oxide
MgF2 is called magnesium fluoride
The formula for a binary ionic compound is called a formula
unit. The formula unit describes the lowest whole number ratio
of the atoms in the compound. Ionic compounds are crystals. In
the crystal, regardless of size, there is a specific ratio of metal
ions to nonmetal ions. There is not a discrete molecule. In the
picture to the right there is 1 salt crystal. It is made up of sodium
ions and chloride ions. The formula unit (lowest whole number
ratio) is NaCl.
42
Compounds Made with Variable Charged Metals
A transition metal or a post-transition metal can have more than one possible positive charge as a
cation. It is not sufficient to simply name the metal and the nonmetal. The charge of the metal ion
must be included in the name. It is included as a Roman numeral immediately after the name of the
metal. It is not possible to predict this charge, but it can be deduced if the number and charge of the
nonmetal anions is known. The anions again use the suffix –IDE.
Examples: FeCl2 is called iron II chloride. The iron has a 2+ charge
FeCl3 is called iron III chloride. The iron has a 3+ charge
Binary Molecular Compounds
When two nonmetals form a compound, they must each be named using prefixes. These number
prefixes must be used because these are not ions, charge is not a factor and can not be used in
determining the formula. The more metallic (to the left or further down the periodic table) element
is named first. It is simply the name of the element. It is preceded by a prefix which tells exactly
how many atoms of that type are present. In the case where there is only one atom of the first
element, the prefix mono- is dropped and presumed redundant. The second element is then named.
It is again preceded by a prefix which shows the number of atoms of that element present. Then it
is given the –IDE ending. In the case where the prefix ends in a vowel and the element name begins
with a vowel, the vowel is dropped from the prefix. A list of common prefixes follows.
Prefix
Number of atoms A molecular formula describes the quantity and type of each
mono1
atom in the molecule. Below are 4 water molecules. Each
di2
discrete molecule has the formula of H2O.
tri3
tetra4
penta5
hexa6
hepta7
octa8
Examples: N2Cl4 is called dinitrogen tetrachloride
SF6 is called sulfur hexafluoride (mono- is not used before sulfur)
P2O5 diphosphourous pentoxide (pent- instead of penta- before oxygen)
Polyatomic Ions
Many ionic compounds contain polyatomic ions. It is necessary to learn the names of all the
polyatomic ions. It is also necessary to learn the charges of the polyatomic ions. When the cation is
a metal with a known charge, the metal is named first, followed by the name of the polyatomic ion.
When the cation is a metal with a variable charge, Roman numerals are used to show the charge on
the metal ion. The charge of this cation can be deduced by knowing the charge on the polyatomic
anion. Compounds which contain polyatomic ions are usually named without prefixes.
Examples: • K2SO4 : potassium sulfate.
• Cu2SO4 : Sulfate has a 2- charge. Therefore, the two copper ions must have a net
charge of 2+, since there are two coppers each copper must have a 1+ charge. This is
called copper I sulfate.
• CuSO4 : Sulfate has a 2- charge. Thus, the copper ion must have a 2+ charge. This
is called copper II sulfate.
• Pb(SO4)2 : Sulfate has a 2- charge. Therefore, the sulfate ions must have a net
charge of -4, so the lead ion must have a 4+ charge. This is called lead IV sulfate.
43
Writing Formulae from Names
All of the compounds in this laboratory exercise will have a neutral charge. To write the formula
for a compound from its name, you must use enough ions to balance the charges. Use subscripts to
indicate the number of times an ion occurs in a compound. When polyatomic ions are used, the
entire ion is placed in parentheses and subscripts are added as needed.
Examples:
calcium chloride
Calcium has a 2+ charge and chloride has a 1- charge, therefore each calcium
ion requires two chloride ions to achieve a net charge of zero on the
compound. The formula will be CaCl2. The subscript 2 shows that there are
two chlorides. This CANNOT be written as Ca2Cl.
aluminum oxide
Aluminum has a charge of 3+, oxygen in oxide ion is 2-. The compound
between aluminum and oxygen, Al2O3, has two aluminum ions and three
oxide ions. The net charge of 6+ on the two aluminum ions is balanced by 6on the three oxide ions.
barium nitrate
Ba(NO3)2: Barium has a 2+ charge, nitrate has a 1- charge. It takes two nitrates
to combine with each barium.
barium
Ba3(PO4)2: The three 2+ barium ions combine with two 3- phosphate ions.
phosphate
Note that you must know the charges on the ions before you can write the
formula for the compound.
Simple Anions and Polyatomic Anions
Most of the polyatomic ions in this course will be anions and will consist of one or more elements
combined with oxygen. These polyatomic anions are negatively charged. You will need to
memorize all of these polyatomic anions and their charges. Note: when the charge on an ion is 1+
or 1- it is customary to omit the 1 and simply write + or -.
Anions with a 1- charge
Hypochlorite
ClO
Nitrite
NO2Chlorite
ClO2Nitrate
NO3Chlorate
ClO3
Cyanide
CNPerchlorate
ClO4Cyanate
OCNPermanganate
MnO4
Thiocyanate
SCNAcetate
CH3COOHydroxide
OHHydrogen carbonate HCO3
Hydrogen sulfate (bisulfate)
HSO41(bicarbonate)
Nonmetals from 7A (*.ide)
F-,Cl-, Br-, etc.
2-
Oxide
Sulfide
Sulfite
Sulfate
Thiosulfate
Peroxide
O
S2SO32SO42S2O32O22-
Phosphate
Nitride
PO43N3-
Anions with a 2- charge
Carbonate
Oxalate
Chromate
Dichromate
Hydrogen phosphate
Anions with a 3- charge
Phosphide
44
CO32C2O42CrO42Cr2O72HPO42-
P3-
Cations
Some common metal cations you will be expected to know are given below. When a Roman
numeral is used it indicates that a metal is present which can have more than one possible charge.
Cations which can have more than one possible charge must include the charge along with the name
of the metal. This is not an all-inclusive list.
Cations with a 1+ charge
Copper I
Cu+
Ammonium
NH4+
+
Silver ion
Ag
All metals ions from 1A Li+, Na+, K+, etc.
Proton
H+
(“lithium ion” etc)
Chromium II
Cadmium ion
Copper II
Iron II
Lead II
Manganese II
Chromium III
Iron III
Lead IV
2+
Cr
Cd2+
Cu2+
Fe2+
Pb2+
Mn2+
Cations with a 2+ charge
Nickel II
Tin II
Zinc ion
Mercury I
Mercury II
All metals from 2A
3+
Cations with a 3+ charge
Aluminum ion
Cobalt III
4+
Cations with a 4+ charge
Tin IV
Cr
Fe3+
Pb
Ni2+
Sn2+
Zn2+
Hg22+
Hg2+
Be2+, Mg2+, etc
Al3+
Co3+
Sn4+
Acids
Many acids are compounds which donate H+ ions. When a H+ is attached to any of the polyatomic
anions on the previous page an acid is formed. To name an acid, first look to see if any oxygen
atoms are present.
If no oxygen atoms are present the prefix is HYDRO- and the suffix –IC is added to the root of the
element name. Finally the word acid is added.
Examples: HCl is called Hydrochloric acid (the acid from the chloride ion)
HCN is called Hydrocyanic acid (the acid from cyanide ion)
If any oxygen atoms are present, no prefix is added. The root of the polyatomic anion name is used.
a. If the anion name ends in –ATE, the suffix –IC is added to form the acid name.
b. If the anion name ends in –ITE, the suffix –OUS is added to form the acid name.
Examples: H2SO4 is called sulfuric acid (the acid from the sulfate ion)
H2SO3 is called sulfurous acid (the acid from sulfite ion)
Hydro is never used to name an acid when oxygen is present in the anion.
45
Acid Salts
Acid salts are formed when a metal ion replaces some, but not all, of the hydrogen ions in a
polyprotic acid (an acid with more than one proton i.e. H2SO4). The word hydrogen then appears in
the name of the salt.
Examples: KHSO4 is called potassium hydrogen sulfate.
When there is more than one metal ion or hydrogen ion, prefixes are used to indicate the number of
these ions
Examples: Na2HPO4 is called disodium hydrogen phosphate
NaH2PO4 is called sodium dihydrogen phopshate
Bases
The names of many common bases follow the normal rules for naming salts but the anion involved
is the hydroxide ion, OH-.
Examples: NaOH is sodium hydroxide
Ca(OH)2 is called calcium hydroxide
CuOH is called copper I hydroxide
Ammonia is considered a base because it increases the hydroxide concentration in a solution
because of following.
NH3 + H2O à NH4OH à NH4+ + OH-
Pre-Lab Questions:
1. Look at the ingredients on a tube of toothpaste. Does the toothpaste contain fluorine or
fluoride? What is the difference between fluorine and fluoride? State the brand of
toothpaste that you used for this assignment.
2. Look at a soda pop container. Determine if it is a low sodium beverage (less than 35mg of
sodium in 240mL). What is inaccurate about this statement? What kind of beverage did
you use?
3. People refer to the chemical used in swimming pools as chlorine. Why is this inaccurate?
What chemical is in fact used in swimming pools?
4. Define the following trivial names of ions
a. Ferric
b. Ferrous
c. Stannic
d. Stannous
e. Plumbic
f. Plumbous
5. From the above list state a rule that governs the assignment of the suffix –IC and the suffix –
OUS regarding the trivial naming of variably charged metal ions.
46
Procedure:
Read the data and observation section to know how to arrange data table.
For each question write out the question and its answer in your laboratory notebook in the data
section. Provide a brief explanation which shows your reasoning where appropriate.
Data and Observations: Arrange data table as follows.
1. To which family does the metallic element in each of the following compounds belong?
What is the charge on the metallic element in each of the following compounds?
Question
Answer
Explanation
Alkali metals lose e- to be more like the nearest noble gas.
a. Na3N
Alkali metals, Na+
b. MnO2
a. Na3N
c. MgS
e. Al2(SO4)3
b. MnO2
d. PbCl4
f. Ni(NO3)2
2. To which family does the non-metallic element in each of the following compounds belong?
What is the charge of the non-metallic element in each of the following compounds?
Example Question
a. Na2O
Answer
O is a member of the
oxygen family
b. Ca3P2
a. Na2S
b. CaP2
It has a 2- charge.
c. Li3N
d. BaO
Explanation
e. CaF2
f. KBr
3. Name the following compounds
Example Question
Answer
Explanation
2nickel III
SO4 is called sulfate. The sulfates have a net charge of 6-,
Ni2(SO4)3
sulfate.
a.
b.
c.
d.
NaI
CuSO4
KHCO3
KBr
e.
f.
g.
h.
therefore the two Ni ions must each have a charge of 3+. Nickel is
a transition metal and therefore the name must include the charge.
Na2SO4
N2 O4
SrCO3
Ca3(PO4)2
i. NaOCN
j. LiNO3
k. SF6
l. K2C2O4
m. SCl2
n. Pb(CH3COO)2
o. (NH4)2Cr2O7
p. Sn(ClO4)4
4. Write the formula for the following compounds.
Question
Answer
Explanation
HgI2
Mercury II is Hg2+ and iodide is I-. It will take two iodide ions to
Example
balance the charge on each mercury ion.
Mercury II iodide
a. Iron II sulfate
g. Tin IV oxalate
m. Mercury II nitrate
b. Barium phosphate
h. Chromium III carbonate
n. Calcium perchlorate
c. Sulfur dioxide
i. Dipotassium biphosphate
o. Strontium hydrogen carbonate
d. Sodium oxalate
j. Magnesium nitrite
p. potassium cyanide
e. Potassium permanganate
k. Chromium II bromide
q. mercury I bromide
f. Sodium hydrogen sulfate
l. Silver chromate
r. Nickel II nitrate
47
5. Name the following acids or bases.
Question
Answer
Explanation
Example
chlorous acid this acid comes from a chlorite ion and thus the ending is _ous
HClO2
a. HBr
b. HCN
c. HNO3
d. H2S
e. H2SO4
f. Mg(OH)2
g. H3PO4
h. KOH
i. HNO2
j. HClO3
6. Write the formula for the following compounds.
Question
Answer
Explanation
Example
Ba(OH)2 Barium ion is always Ba2+, Hydroxide is OH-. Two
Barium hydroxide
hydroxides are needed for each barium.
a.
b.
c.
d.
Hydrofluoric acid
Perchloric acid
Aluminum hydroxide
Copper II hydroxide
e. Oxalic acid
f. Hypochlorous acid
g. Ammonia
h. Carbonic acid
i. Potassium hydroxide
j. Acetic acid
k. Sulfurous acid
l. Ammonium hydroxide
Calculations:
Not applicable
Results:
Not applicable
Discussion:
Not applicable
Post-Lab Questions:
Not applicable
References
Collins, V. Kahl, D. (1995) Good stuff from the chemistry laboratory. Warren Wilson College
Press: Asheville, North Carolina
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102
laboratory manual. Star Publishing Company: Belmont, California
Finnegan, M. Place, H. Weissbart, B. (1997) Washington state university chemistry 105-106
laboratory manual. Star Publishing Company: Belmont, California
48
Experiment 6
Determination of an Empirical Formula
Purpose:
To determine the simplest (empirical) formula of a chemical compound through experimentation.
Develop a hypothesis regarding the mass of the material at start and the mass of the material at the
end. (Will the burnt magnesium have more or less mass than the unburned magnesium?)
Background:
Chemical formulae and moles
The subscripts in a simple formula provide the relative number of atoms in a substance. For
example, in the compound CO2 the subscripts show that each carbon dioxide molecule contains one
carbon atom and two oxygen atoms.
A simple formula provides two pieces of information. The subscripts in the simple formula give
both the number of atoms and the number of moles in a substance.
•
For 1 molecule of CO2 there is one atom of carbon and two atoms of oxygen
•
For 10 molecules of CO2 there are ten atoms of carbon and twenty atoms of oxygen
•
In 1 dozen molecules of CO2 there are 1 dozen atoms of carbon and 2 dozen atoms of Oxygen
•
For every 6.02 x1023 molecules of CO2 there are 6.02 x1023 atoms of carbon and 2x6.02 x1023
atoms of oxygen
•
For 1 mole of CO2 there is 1 mole of carbon atoms and 2 moles of oxygen atoms
•
For 0.25 mole of CO2 there are 0.25 moles of carbon and 0.50 moles of oxygen atoms
In the laboratory we can not count individual numbers of atoms as we did in the example above.
Thus, we use the mole concept to determine the formulas because 1.00 mole of any element has
6.02 x1023 atoms and a mass which is equal to the atomic mass. This concept allows us to count
atoms indirectly. If the mass of an element is known, the moles and atoms of that element can be
calculated.
In practical terms the problem is to determine how many grams of each element are present in the
compound. In this experiment we will start with pure magnesium, and then determine the mass of
magnesium. The magnesium will react with oxygen to produce magnesium oxide. Because
magnesium and magnesium oxide have very high boiling temperatures the mass of magnesium will
not change during the experiment. When we produce magnesium oxide, oxygen will combine with
the magnesium; the increase in mass will be due to oxygen that has reacted with the magnesium.
We can find the mass of oxygen in the magnesium oxide by subtracting the mass of magnesium
from the total mass of the compound.
When the masses of magnesium and oxygen in the sample are known we can calculate the moles of
magnesium and oxygen. Once the moles of magnesium and the moles of oxygen are known the
simple whole number ratio of atoms can be calculated. This information will provide the formula.
49
In this experiment we will carefully measure the mass of magnesium to three decimal places using
the analytic balance. We will then combust (burn with oxygen) the magnesium in the presence of
oxygen to form magnesium oxide. When we produce magnesium oxide, oxygen will combine with
magnesium. The oxygen content will be determined by subtracting the mass of magnesium from
the final mass of the magnesium oxide. Once we know the mass of magnesium and the mass of
oxygen we can calculate the moles of magnesium and the moles of oxygen and the formula of the
magnesium oxide.
Error analysis
In this lab you will form magnesium oxide from a given quantity of magnesium metal. The mass of
the magnesium oxide is affected by experimental errors. It is useful to know how much magnesium
oxide should have been formed theoretically and compare that to the quantity that was formed
experimentally. If the values are close it may indicate that experimental error did not significantly
impact your results. If on the other hand the difference is large between the mass produced
experimentally and the mass that you should have obtained theoretically then experimental error did
have an impact on your results.
To calculate how much product should be produced it is necessary to write the balanced equation
then use dimensional analysis to convert grams of reactants to grams of products that should be
produced. For instance:
If 1.035 grams of sulfur reacts with excess oxygen gas to form sulfur monoxide how much sulfur
monoxide could be formed?
Step 1. Write and balance equation from sentence use subscripts to indicate state.
2S(g) + O2 (g) à 2SO (g)
Step 1: Find moles of sulfur
Calculations 6.1
1.035 g O2
1 mol O2
32.06 g O2
= 0.03228 mol molecular oxygen
From the periodic table
Step 2: Convert moles of sulfur to moles of sulfur monoxide
0.03228 mol S
2 mol SO
2 mol 2
= 0.03228 mol SO
From the balanced equation
Step 3: Convert back to grams
0.03228 mol SO
48.07 g SO
1 mol SO
From the periodic table
1 mole of oxygen is 16.00g/mol
1 mole of sulfur is 32.07g/mol
= 1.552 g sulfur monoxide
Or you could combine all the steps
1.035 g S
1 mol S
2 mol SO
32.06 g S 2 mol S
48.07 g SO
1 mol SO
= 1.552 g sulfur monoxide
Once the theoretical mass of the product is
| experimental-theoretical |
calculated this can then be used in the percent error
theoretical
formula to determine the accuracy of your data.
50
*100 = % error
formula 6.1
Pre-Lab Questions:
1. From the periodic table find the atomic weights for:
a. Oxygen
c. Carbon
b. Magnesium
d. Nitrogen
2. Find moles in:
a. 0.101grams of magnesium
c. 1.25 grams elemental hydrogen (H2)
b. 0.102 grams of elemental oxygen (O2)
d. 12.35 grams of water (H2O)
3. Find grams in:
a. 12.1 mole nitrogen
c. 1.087 mole carbon monoxide (CO)
b. 0.00345 mole chlorine
d. 0.00891 mole of magnesium oxide
4. 12.25 grams of hydrogen are combined with excess oxygen to form water.
2H2(g) + O2(g) à 2H2O(g)
a. How many grams of water should be formed?
b. Find the percent error if 99.75 grams of water are formed.
Materials:
0.100 - 0.200 g Mg ribbon
Crucible and lid
Clay triangle, ring and ring stand
Steel wool
Bunsen burner
Wire gauze
Distilled water
Tongs
Procedure:
1. Obtain a crucible and lid. Wipe out crucible with paper towels, do not get the crucible wet. Use
only crucible tongs to handle the crucible and lid for the remainder of the experiment. Support the
crucible on a clay triangle, heat gently at first and then heat with an intense blue flame for 5
minutes. (A yellow flame will deposit carbon on you crucible which negatively effecting your
results). Using tongs (the crucible is HOT!), place them on a wire gauze and allow them to cool to
room temperature. The crucible may appear dirty; in most cases they have permanent stains. If you
are uncertain about the cleanliness of the crucible, check with your instructor.
2. While the crucible is cooling, obtain a sample of magnesium ribbon from your instructor and
polish it with steel wool.
3. Use an analytical balance to determine the mass of the cooled crucible and lid; record the mass
directly into your lab notebook data table. Make certain the balance reading is zero prior to
obtaining your mass.
4. Curl the Mg ribbon to lie in the crucible and determine the combined mass of the crucible, lid
and magnesium, then record this mass directly into your lab notebook. Verify that the mass of your
magnesium is between 0.100 grams and 0.200 grams.
51
5. Place the crucible containing the Mg ribbon and lid
on the clay triangle. Heat slowly, occasionally lifting
the lid to allow oxygen to reach the Mg ribbon. If too
much air comes in contact the Mg ribbon, rapid
oxidation will occur and it burns brightly. This will
negatively impact your results. You do not want this to
happen. Immediately return the lid to the crucible.
6. a. Continue heating the crucible until no visible
change is apparent in the magnesium ash at the
bottom of the crucible when the lid is lifted
b. Remove the lid. Continue to heat the open
crucible and ash for 30 seconds. Remove from
the heat, replace lid and allow the crucible to
Experimental Setup
cool to room temperature.
c. Add 3 drops of water to decompose the magnesium nitride formed during combustion.
Caution: cool water added to a hot crucible will cause it to break. Make sure it has cooled
before adding water.
d. Dry the ash with a low blue flame for a minute or two and then allow it to cool. Allow the
steam to escape by setting the lid slightly to one side leaving a gap.
e. Measure the mass of the crucible, lid and ash on the same balance that was used earlier and
record.
7. a. Reheat the sample for 1 minute
b. Again measure the mass of the crucible, lid and ash. If this second reading is greater than +
or – 0.005 grams
from that recorded for the 1st mass repeat procedure 7. Repeat as required until the change in
mass is less than + or – 0.005 grams
8. Clean up your crucible with water and invert on the drying rack.
Data and Observations:
1.
2.
Mass of crucible and lid
Mass crucible, lid and magnesium
1st mass of crucible lid and magnesium oxide
2nd mass of crucible lid and magnesium oxide
3rd ? you may need 2,3,4 or more repeat mass to drive off all water.
3. Final mass crucible lid and magnesium oxide
Observations
52
Calculations:
1.
2.
3.
4.
5.
Mass of magnesium (data item 2. –data item 1.)
Mass of oxygen
(3. – 2.)
Moles of magnesium
Moles of oxygen
Experimental ratio between moles of magnesium and
moles of oxygen (ratio of 6.:7.). Use significant figures.
6.
7.
Divide both the moles of O and the moles of Mg by the smaller coefficient.
If the ratio in step 6 is within + or - 0.15 units of an integer, you can round off to the closest integer.
If this is not the case, multiply both terms in the ratio by the same number until a whole number for
each is achieved.
The answer may be something
like Mg:O = 1: 1.209
Error analysis
1. Determine how much MgO was produced (3.-1.). (This is the Experimental mass.)
2.
3.
Use the balanced equation and stoichiometry to find out how much magnesium oxide should have
been produced. (This is the Theoretical mass.) (see calculation 6.1 as an example)
Use the % error to evaluate this experiment. (|Theoretical-Experimental|/Theoretical)*100 = % error
Results:
1. Empirical formula of oxide formed
2. Name of oxide formed
3. % error
Discussion:
Summarize results, comment on the obvious sources of error, explaining how they would make
your results too high or low, use the percent error as evidence. A percent error of less than 10% is
acceptable, less than 5% is good, less than 1% is excellent. Address the discussion questions
provided in “How to Write a Lab Report.”
Post-Lab Questions:
1. In a laboratory experiment to determine the formula of the compound formed between Cu
and S, a student obtained the following data.
Mass of crucible
19.732 g
Mass of crucible and Cu
27.304 g
Mass of crucible and compound of Cu and S
29.214 g
What is the formula of the compound?
2. A compound containing nitrogen and oxygen is analyzed. A sample is found to contain
0.483 g N and 1.104 g O. What is the simplest formula of the compound? Show all
calculations, use dimensional analysis, use sig figs.
3. A compound containing carbon and hydrogen is analyzed. A sample is found to contain
4.804 g C and 1.210 g H. What is the empirical formula of this compound? Show all
calculations, use dimensional analysis, use sig figs.
References:
Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson
College Press: Asheville, North Carolina
53
Experiment 7
Reaction Rate
Marble Lab
Purpose:
Graph the reaction rate of Calcium carbonate and Hydrochloric acid. Then determine the changes in
reaction rate due to changes in temperature.
Hypothesis: State a hypothesis regarding reaction rate and temperature (how will temperature
effect reaction rate? Is the correlation negative or positive?).
Materials:
Marble (calcium carbonate) chunks Stop watch
Ice
Thermometer
Balance
200 mL 1.0M HCl
Styrofoam cup
Burette stand & clamp
Background:
Reaction rate
Reaction rate is a measure of the rate at which a chemical reaction proceeds from reactants to
products. The units used are moles/sec.
A useful model for understanding reaction rate is the Collision Theory. The collision theory
proposes that atoms, ions and molecules can form a chemical bond when they collide, provided the
particles have enough kinetic energy. Particles lacking the necessary kinetic energy may collide but
simply bounce apart. The minimum energy that colliding particles must have in order to react is the
activation energy. During a reaction, when the products initially collide, particles that are neither
reactants nor products momentarily form. An activation complex is the arrangement of atoms at
the peak of activation energy. Activation complexes are very unstable and last on the order of 10-13
seconds. Activation complexes are sometimes referred to as the transition state.
Collision theory explains why some spontaneous reactions are imperceptibly slow at room
temperature. The reaction of carbon and oxygen is spontaneous, but it has a high activation energy.
At room temperature the collision of oxygen and carbon molecules are not energetic enough to
overcome the activation energy required to break the O-O and the C-C bonds. These bonds must
first be broken to form the activation complex. Thus the reaction rate of carbon with oxygen at
room temperature is essentially zero.
Factors effecting Reaction Rate
Each of these factors which effect reaction rate can be considered using the collision theory of
chemical reaction. When considering the collision theory it may be useful to think about these
chemicals as clay balls rolling around on a billiards table. For a reaction to take place these clay
balls have to collide with one another with enough energy and the correct orientation to overcome
the activation energy to form a transition state and ultimately chemically recombine.
Temperature
Temperate is a measure of the average kinetic energy of a sample. Thus raising the temperature in
general makes the molecules move faster, increasing their kinetic energy. This has two impacts on
reaction rate. First the particles are moving faster and thus have a higher kinetic energy and thus are
54
more likely to have the energy required to overcome the activation energy. Secondly because the
particles are moving faster this also increases the number of collisions. Thus by increasing
temperature, the particles heat each other with more energy and more often. These two things
increase the reaction rate. As temperature goes up reaction rate often goes up.
Concentration
As more particles are placed in a fixed volume container the concentration increases. If there are
more particles in a given volume the chances of a collision also increase. Imagine trying to
randomly collide two particles on a billiards table. The chance of a collision is less than if there are
100 particles on the billiards table.
Surface area
One of the universal answers to all questions in science is surface area. If you need a vague answer
that is probably correct at some level, surface area is it. In fact surface area effects reaction rate. As
particle size decreases surface area increases dramatically. For instance it is much easier to ignite
some dry leaves than it is to ignite a whole tree. Why? Surface area. One of the ways chemists
often increase surface area is to dissolve materials into aqueous solutions. This increases surface
area. Two very concrete examples of surface area are coal dust and grain silos. Grain itself is not
particularly flammable nor would you consider a chunk of coal as an explosive. Certainly there is a
lot of potential chemical energy stored in these two items. However to access their explosive
characteristics it is only necessary to increase their surface area.
When pouring grain into the top of a silo a very fine powder of cellulose is suspended in the air. If
this is ignited the silo explodes creating a bomb that will certainly snuff out the cigarettes of any
nearby smoking farmers. Coal dust in the mines is notorious for huge levels of destruction. In fact
a safety lamp was created that utilizes a wire mesh to reduce the possibility of igniting the coal dust.
This of course leaves all the coal dust in the air for our ancestral miners to breathe. Coal dust
explosions kill.
Catalysts
A catalyst is a substance that increases the rate of a reaction without being used up in the reaction.
Catalysts reduce the activation energy necessary for a chemical reaction. Think of it as making the
products stickier so that when things collide they are more likely to react. It is probably more
correct to say that a catalyst provides a location for the reaction to take place on. The catalyst can
be thought of as holding the product in an orientation so that they may react more easily. However
catalysts themselves do not take part in the reaction and do not get used in the reaction. Catalysts
do not provide the energy to allow a reaction to take place. Heat, concentration and particle size are
never considered catalysts.
Your body is at around 37 oC. Your temperature can not increase very much without causing
brain damage and ultimately death. Yet without catalysts (in the body they are called enzymes)
very few, if any, of the chemical reactions could take place in our cells at body temperature. What
kind of heat is necessary to break down proteins? Consider the heat necessary to cook an egg.
Cooking eggs is just using heat to break protein bonds (that is why they turn white). Enzymes allow
you to break down the proteins in an egg at body temperature without having to reach very high
temperatures.
When a catalyst is used in a chemical reaction it is not consumed in the reaction. It is neither
a reactant nor a product. In a car’s catalytic converter Pt is used to convert 2
poisonous gases into 2 harmless gasses. Please note how the Pt
Pt
is above the reaction arrow. The Pt is not consumed. The Pt
2CO + 2NO → 2CO2 + N2
never runs out.
55
The Mole
An important concept used in this lab is the mole. This is not the insectivore that lives under your
lawn that is spelled mole. In chemistry, when we talk about a mole it refers to a very specific
number. A mole of something is 6.02x1023. A mole refers to a number. Just like a dozen means
12, a mole means 6.02x1023. Chemists use the concept of the mole to refer to the number of
particles involved in a reaction.
Example A: 1 mole of oxygen reacts with 2 moles of hydrogen.
vs.
Example B: 620,000 times a million, times a million, times a million molecules of oxygen reacts
with 1,240,000 times a million, times a million, times a million molecules of hydrogen.
Pre-Lab Questions:
1. Find grams
a. For one mole of calcium
b. For one mole of carbon
c. For three moles of oxygen
2. How many moles of each element are in one mole of calcium carbonate (CaCO3)?
3. What is the mass of one mole of calcium carbonate in grams?
Procedure:
In this lab it is necessary to establish certain operational definitions. An operational definition is a
statement regarding how an activity is to be conducted. They are very specific and are designed so
another scientist can follow your procedure exactly. Write operational definitions for when the time
starts and when the time stops. Also write any other definitions necessary such that if someone
repeats your experiment they would get the same results.
You will work in your lab groups of 4 students to share a balance at your lab station.
1.
2.
3.
4.
5.
Cover your balance with a layer of plastic wrap (saran wrap©)
Obtain 2 marble chips of similar size and shape. Mass the marble chips and record data.
Mass a Styrofoam cup. Record this mass
Place 100 mL 1.0M HCl in a Styrofoam cup.
Set a thermometer up so that it is suspended in the acid solution without touching the cup.
You will record the temperature every minute. Find the temperature of the acid and record
this data. It should be close to 20 oC. Record the data using 3 sig figs.
6. Place cup and acid onto a balance and record mass.
7. Place one marble chip onto balance next to the Styrofoam cup (not inside the cup).
8. Prepare stopwatch to be used. When you are ready proceed.
9. Place the marble chip into the acid. Record mass every 15 seconds.
10. Record temperature every 1 minute.
11. Continue to record data until 2 concurrent data points are within 0.5 grams. At this point
continue collecting data for 4 more time intervals.
9. Record the final temperature of the experiment.
10. Repeat the above steps, however use the chilled hydrochloric acid.
56
Data: (example data tables follow)
Marble Chip 1
Mass of marble chip
Time
Mass
Temp
Observations
T0
T15
T30
T45
T60
Tetc
Leave room for the entire length of trial (which could last several minutes)
Marble Chip 2
Mass of marble chip
Time
Mass
T0
T15
T30
T45
T60
Tetc
Temp
Observations
Results:
Each experiment should have its own graph.
Graph: Provide a title for each graph. List the dependent variable on the Y axis (Mass) and the
independent variable on the X axis (Time). Please do this graph such that it uses a full sheet of
paper in a landscape orientation. Please use a ruler and equal graduations. Do not draw the graph
with a (0,0) origin. The origin of the graph needs to be appropriate to the data graphed.
Discussion:
Compare the graphs and determine what effect each tested variable had on the reaction rate.
Post-Lab Questions:
1. How does each of the following factors affect reaction rate?
a. temperature
b. particle size
c. concentration
d. catalysts
2. A tablet with a mass of 3.251 grams dissolves in 25 seconds. What is the average reaction rate?
3. Based on your graph, what would the reaction rate be if a tablet dissolves in a 40 degree water
bath?
57
Experiment 8
Qualitative analysis:
Flame test
Purpose:
To conduct a series of experiments and to make qualitative observations
To identify an unknown salt using flame testing
Hypothesis: (Write the hypothesis for part 2 of this experiment.)
Provide a hypothesis regarding the identity of your unknown samples.
Background:
Qualitative observations
In many sciences the primary focus of data collection is quantitative data. Quantitative data uses an
instrument to act as an interface between the observer and the environment which is being observed.
This tool provides a numeric value to quantify observations. The value in this kind of observations
is that in theory anyone, who uses the instrument correctly, will get essentially the same data.
Because science is driven by the concept of repeatability it places quantitative data collection in the
forefront.
Another form of data collection is qualitative observations. Qualitative observations focus on
qualities: hot, cold, white, chunky, etc. These observations have limitations because an
experimental situation can be described differently by different people. However qualitative
observations are no less important.
Because qualitative observations do not rely on instruments there is no need to consider
instrumental error. People have a lifetime of experience describing things. In qualitative analysis
the scientist observes the qualities of the situation or experiment and then, makes detailed notes of
these observations. Then these notes are compared to other observations. In general qualitative
observations can produce very satisfactory and accurate results. One particular weakness of
qualitative observations is the difficulty in statistically analyzing the results.
Flame Test
A spectral line is the result of the emission of a photon of specific energy (thus specific frequency)
when the electron moves from a higher energy state to a lower one. An atomic spectrum appears as
a line rather than as a continuum because the atom’s energy has only certain levels or states. In the
Bohr model the quantum numbers (1,2,3,n) are associated with the radius of the electron’s orbit,
which is directly related to the atom’s energy. When the electron is in the orbit closest to the
nucleus (n=1) the atom is in its lowest energy level, called the ground state. By absorbing energy
equal to the energy necessary to push the electron from the first to second energy level the electron
can move between levels. When the electron is in this higher than normal state it is called the
excited state.
In a flame test you excite the electrons from the ground state to an excited state. When the electron
returns to the ground state it releases energy and this releases a photon of light. This analysis
produces spectral data. The use of spectral data to identify and quantify substances is essential to
58
modern chemistry. The terms spectroscopy, spectrometry and spectrophotometer denote a large
group of instrumental techniques that measure substances atomic and molecular energy levels from
the spectra produced.
The two types of spectra most often obtained are emission and absorption spectra. An emission
spectrum, such as the H atom line is produced when atoms that have been excited to a higher energy
level emit photons characteristic of the element as they return to a lower energy level. Some elements
produce a very intense spectral line (or several closely spaced ones) that serve as a marker of the
element’s presence. Such an intense line is the basis of flame tests, rapid qualitative procedures
performed by placing a granule of an ionic compound or a drop of its solution in a flame.
Some of the colors in fireworks and flares are due to the emission from the same elements as the flame
tests: red from strontium, blue green from copper salts. The characteristic color of sodium vapor
lamps and mercury vapor are due to one or a few prominent lines in their emission spectra.
An absorption spectrum is produced when atoms absorb photons of certain wavelengths and become
excited from lower to higher energy levels. Therefore the absorption spectrum of an element appears
as dark lines against a bright background.
Didymium glass is used to filter out the yellow produced from the heated metal. It also filters out
the sodium flare contimation when testing other metals. Didymium is greek for twin element. The
glass is made with both praseodymium and neodymium. It is used in safety glasses for glassblowing
and blacksmithing, especially when a gas is used for heating. It filters out the yellowish light at 589
nm emitted by the hot sodium in the glass.
Materials:
Spot reaction plate (96 well tray)
Didymium glass
All solutions are 0.2 Molar
Copper nitrate
Strontium chloride
Lithium nitrate
Bunsen burner
Chloride test solution
(Silver nitrate 0.2M)
Sodium chloride
Copper chloride
Potassium chloride
Lithium chloride
Ni-Chrome wire loop
3.0 M HCl copper cleaner
Sodium nitrate
Strontium nitrate
Potassium nitrate
Procedure:
You will use a piece of Ni-chrome wire (nickel chromium alloy) with a loop at the end to place one
drop of a solution you want to test in a Bunsen burner flame. A piece of copper wire will be flame
tested by placing it directly into the flame. You will also test the effects of a test solution to test for
chlorides on each of the known salt solutions. Upon completion of these tests an unknown solution
will be obtained and its identity will be determined.
Part 1: Flame test of known chemicals
Note: Do not exchange wires. For each solution only use the wire that is already in that solution.
After you use the wire, be sure to put it back with the solution from which it came. To reduce the
possibility of mixing up the wires, test only one solution at a time. When you have completed one
test, return the solution to the test tube rack and obtain the next solution.
1. Remove the wire from the solution.
59
2. Place the loop of the wire with the solution on it in the flame. For testing the copper wire,
first clean the wire with hydrochloric acid. Then place the wire directly into the flame using
the tongs.
3. Observe and record the color of the flame.
4. Place the wire back in the solution.
5. Record results in the data Table A in your lab notebook.
6. Repeat steps 2-5 this time obseve flame through didymium glass square.
7. Obtain next sample and repeat steps 1-6.
Part 2: Flame test unknown salt solution
1. Record the unknown’s Identification Number in data Table B. Keep you unknown in your
drawer. Under no circumstance return the unknown. Dispose the unknown after the
satisfactory evaluation of your lab.
2. Repeat above steps with unknown salt solution.
3. After determining the possible identity of the unknown solution, repeat the test with the
known and unknown side by side to confirm the identity. This means having both sample in
the flame at the same time.
Part 2.1: Hypothesis
1. Based on experimental evidence develop a well reasoned hypothesis based on experimental
evidence.
2. Write down the hypothesis.
Part 3: Test for chloride (silver nitrate).
Use a magic test solution to determine the presence of chloride in the unknown salt.
1. Place one drop of unknown solution into one of the wells of the 96 well tray.
2. Place one drop of the test solution into the same well.
3. Observe any changes
4. Record observations in data table.
If chloride is present solid chunks will form when the two solutions are mixed.
If no chloride is present no visible reaction will be observed.
Data table (Write your observations directly into your lab notebook. Never write on scratch
paper or onto a sample data table in your lab manual.)
Known chemical flame tests
Table A
Name of salt solution
Nichrome wire
Sodium chloride
Potassium nitrate
Copper nitrate
Lithium chloride
Strontium nitrate
Copper chloride
Sodium nitrate
Lithium nitrate
Potassium chloride
Strontium chloride
Copper wire
Initial
Flame, no filter
60
Observations
Didymium glass
Initial unknown flame test:
Table B
Initial
Flame, no filter
Didymium glass
Unknown ID #
Side by Side confirmation test
Name or number
Table C
Observations
Chloride test
Unknown ID number
Known salt w/chloride
Known salt w/nitrate
Results:
Unknown:
Unknown chemical’s number ________
Unknown chemical’s identity____________________________
Discussion:
As per “How to write a lab report.” Discuss the evidence that the metal and not the accompanying
anions is responsible for the color of the flame.
Post-Lab Questions:
1. Group the substances based on the color of the flame produced.
2. What patterns do you notice in the groupings?
3. Predict the color of the flame for a substance called strontium sulfate. Explain your
reasoning.
4. What evidence do you have that atoms of certain elements produce a flame with a specific
color?
5. The yellow color of the flame for sodium indicates that the sodium atoms changed in some
way when they were heated. Consider the following possibility that the electron
configuration of sodium changed from [Ne]3s1 to [Ne]4s1. What is the difference between
[Ne]3s1 and [Ne]4s1?
6. Do you think gold can be made by changing the arrangement of electrons in atoms?
Explain.
References:
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102
laboratory manual. Star Publishing Company: Belmont, California
Willbraham, A. Staley, D. Matta,. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park, CA
Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New
York, NY
Stacy, A, Coonrod, J, Claesgens, J (2003) Alchemy: atoms elements, and compounds. Key
Curriculum Press: Emeryville, CA
61
Experiment 9
Qualitative Analysis:
Precipitation testing
Purpose:
To conduct a series of experiments and to make qualitative observations
To identify an unknown salt using qualitative analysis
Background:
Qualitative observations
As in the Flame Test laboratory activity this activity relies on qualitative observations. When
conducting the flame tests it was sufficient to only note the color of the flame. However in this lab
much more detailed observations will be necessary.
Precipitation
When a salt (a material containing a positively charged cation and a negatively charged anion) is placed
in water the general tendency is for the polar nature of water to be attached to the cations and anions.
These interactions are so strong that they pull the cation and anions apart from each other and into the
solution making an aqueous solution. That is, a solution of water with free floating cations and anions.
Here the water has not yet dissolved the salt into an
aqueous solution, note that the water has both a positive
(δ+) and a negative side (δ-).
The salt (+ cations, -anions) are surrounded
by polar water molecules.
Some salts have such strong attractive forces between the positive cations and negative anions that the
partial electrical charges found on the ends of the water molecule are not sufficiently strong to separate
the cations and the anions from the crystalline structure. These salts are not soluble in water. That
means when you place a salt that has very strong internal forces in water, the water molecule does not
dissolve it and the salt remain a solid surrounded by water.
Soluble means the water dissolves the salt and the solution will look clear (though it may have
coloration). A material that dissolves in water is denoted with the symbol (aq), meaning aqueous.
62
Insoluble means that water can not dissolve the salt and the solution will look cloudy and often be a flat
white, yellow, pale blue, etc. These solutions may translucent (Tyndal effect). A material that does not
dissolve in water is denoted with the symbol (s), meaning solid.
Solubility Rules
1. Salts containing Na+, K+, and NH4+ and acids are always soluble (dissolves, aqueous).
2. Salts containing NO3-, ClO3-, ClO4-, and CH3COO- are always soluble.
3. All Cl-, Br-, and I- are soluble except for those of Ag+, Pb2+, and Hg22+, which are insoluble
(will not dissolve).
4. All SO42- are soluble except of those of Sr2+, Ba2+, Hg22+, Hg2+, and Pb2+. The sulfate salts
of Ca2+, Ag+ are moderately soluble.
5. All OH- are insoluble expect those of alkali metals, which are soluble, and hydroxides of
Ca2+, Ba2+, and Sr2+, which are moderately soluble.
6. All SO32-, CO32-, CrO42- and PO43- are insoluble except those of NH4+and alkali metals.
7. All sulfides S3- are insoluble except those of NH4+, the alkali metals and the alkaline earth
metals, which are soluble.
Acids and bases
In the most general of terms, acids are chemicals that contain a proton that can be disassociated
(falls off) when placed in water. This causes the water’s proton concentration to go up.
H2O + HCl + à H2O + H+ +ClThis is most often written this way:
H2O + HCl + à H3O+ +ClThis is true for all hydroxy acids (acids with free protons). These include H2SO4, HNO3 and others.
Bases are chemicals that contain the polyatomic ion hydroxide (OH-). The hydroxide in a base
disassociates in water.
NaOH + H2O à Na+ + OH- + H2O
The hydroxide, once disassociated, will wander around and scavenge any protons it can find to
make water.
H+ + OH- à H2O
Thus any solution that OH- is placed into has a very low concentration of H+ because the OH consumes it all.
Acids and base neutralize each other to form water and a salt.
HCl + NaOH à H2O + NaCl
In summary
63
Acidic solutions have high proton concentrations and a low pH.
Basic solutions have low proton concentrations and a high pH.
Pre-Lab Questions:
When working on of the following questions completely write out the question and indicate the
state of each product and reactant in the answer. Use (aq) aqueous, (s) solid, (l) liquid, (g) gas.
1. Write and balance each of the following reactions.
a. Ammonium hydroxide reacts with lead II nitrate to produce ammonium nitrate and lead II
hydroxide.
b. Iron III chloride reacts with silver nitrate to produce iron III nitrate and silver chloride.
c. Sodium carbonate reacts with hydrochloric acid to produce carbon dioxide, water and
sodium chloride.
2. Determine the products and then write and balance the equation.
a. Sodium carbonate reacts with calcium chloride to produce…
b. Sodium hydroxide reacts with copper II sulfate to produce…
3. Write and balance each reaction, indicate state. Then produce a written description of
the reaction. IN this description indicate number and state of each constituents.
a. CuSO4 + Na3PO4 à Cu3(PO4)2 + Na2SO4
b. CaCl2 + Pb(NO3)2 à Ca(NO3)2 + PbCl2
4. Determine the products and provide a written description for each set of reactants.
a. KI + AgNO3 à
b. Potassium iodide reacts with lead II nitrate to produce…
Materials:
Spot reaction plate (well tray)
Didymium glass
0.2M Ammonium hydroxide
0.2M Hydrochloric acid
0.2M Potassium Iodide
1.0M Sodium Carbonate
0.2M Sulfuric acid
Bunsen burner
Litmus paper, blue and pink
0.2M Calcium chloride
0.2M Iron III chloride
0.2M Silver nitrate
0.2M Sodium hydroxide
0.2M Sodium nitrate
Ni-Chrome wire loop
HCl cleaning solution
0.2M Copper II sulfate
0.2M Lead II nitrate
0.2M Sodium phosphate
Procedure:
Part 1 Standardization
The first part of this lab is to establish known qualitative data that will then be used to compare to
the unknown sample. Here you are finding out what things look like when you know what
everything is. This will help you when you are trying to figure out what unknown chemical you
have.
Before conducting any tests, make observations of what the chemicals look like in the pipette.
These are your pre-observations.
Acid-base test
To conduct this test simply use a glass stirring rod place a very small drop of the chemical on both
the blue litmus paper and a very small drop on the red litmus paper.
64
Red litmus- will turn blue in the presence of a base
Blue litmus- will turn red in the presence of an acid
Also make detailed observations regarding how the solutions absorb into the paper.
Precipitation test
Clean your well tray very well with tap water. Then do a final rinse using distilled water. To
conduct the precipitation test, place the well slide on a dark surface. Determine some method of
establishing what is in each well used. I might suggest making a diagram on a separate piece of
paper that incorporates a drawing of the well slide and the well labels. Following the scheme
proposed in the data table, mix each of the indicated chemicals together. Use 1-2 drops of each
reagent per reaction. As an example, based on the pre-made data table, you might conduct your
trials in the following order:
1. NH4OH and hydrochloric acid
7. NH4OH and sodium hydroxide
2. NH4OH and sodium carbonate
8. NH4OH and sodium nitrate
3. NH4OH and sulfuric acid
9. NH4OH and copper II sulfate
4. NH4OH and calcium chloride
10. NH4OH and lead II nitrate
5. NH4OH and iron III chloride
11. NH4OH and silver nitrate
6. NH4OH and potassium iodide
12. NH4OH and sodium phopshate
Then start on the reactions with hydrochloric acid, etc.
Make very detailed observations of your reactions. It will be crucial in future steps. You will use
this data to help determine the identity of an unknown salt.
You might consider the following abbreviations:
NVR = no visible reaction
(s) yellow = a yellow precipitate formed
gas = bubbles formed
Part 2 Unknown testing
You will identify an unknown salt solution. Collect your first unknown from your instructor.
Record the unknown’s number in your data section immediately. Your unknown is one of the
solutions for which you have already collected data. If you have made careful observations this
activity will not be frustrating and will be rewarding. If you contaminated your samples or did not
make careful observations during the first part of this lab, identifying your unknown salt will be
difficult, if not impossible.
Acid base test
Conduct the litmus paper test using the same technique as before with your unknown. Make
detailed observations. Not only changes in color but how the chemical absorbs into the paper.
Compare this data with the data collected in part 1.
Precipitation
Clean your well tray very well with tap water. Then do a final rinse using distilled water. Place 23 drops of each of the different solutions into the well slide in some predetermined order. Then add
2-3 drops of your unknown to each of the known samples. Record your observations and compare
them to the data collected in part 1.
Part 3 Confirmation Tests
65
Using the data collected in part 1 and 2 you should be able to establish with some level of certainty
what your unknown is. It is now time to confirm your results.
To do this you will perform each of the experiments side by side: the unknown with the most likely
known salt solution.
As an example: For the precipitation test fill two rows of wells, one with the unknown and the other
with the most likely known. Then one at a time add each of the different knowns to these samples
first to the unknown then to the mostly likely known. Compare data. Do this for all the known
chemicals. When you are done you should have two identical rows of mixtures.
Then do the litmus test. Place a drop of the mostly likely known next to a drop of the unknown on
blue then pink litmus. Compare.
Repeat
Repeat above steps for the 2nd unknown
Pre-observation table
Data:
A large portion of your grade in
this lab will be based on the
correct identification of your 2
unknowns.
The only possible way to
identify your unknowns will be
to have made excellent
observations and then record
those observations in a way that
you can use them efficiently.
There is a limited amount of
time and you will have to use it
carefully.
Ammonium hydroxide
Hydrochloric acid
Sodium carbonate
Sulfuric acid
Calcium chloride
Iron III chloride
Potassium iodide
Sodium hydroxide
Sodium nitrate
Copper II sulfate
Lead II nitrate
Silver nitrate
Sodium Phophate
Litmus paper test
Blue Paper
Pink Paper
Ammonium hydroxide
Hydrochloric acid
Nitric acid
Sodium carbonate
Sulfuric acid
Calcium chloride
Iron III chloride
Potassium iodide
Sodium hydroxide
Sodium nitrate
Copper II sulfate
Lead II nitrate
Silver nitrate
66
Other observations
Sodium Phophate
67
Sample Precipitation Data Table
Sodium
Phosphate
Silver
Nitrate
Lead II
nitrate
Copper II
sulfate
Sodium
Nitrate
Sodium
hydroxide
Potassium
Iodide
Iron III
chloride
Calcium
chloride
Sulfuric
acid
Sodium
carbonate
Hydrochloric
acid
Ammonium
hydroxide
Hydrochloric
Acid
Sodium
carbonate
Sulfuric
Acid
Calcium
chloride
Iron III
Chloride
Potassium
Iodide
Sodium
Hydroxide
Procedural Summary
Sodium
nitrate
Step 1
Complete known observations and tests.
Copper II
sulfate
Step 2
Repeat tests with unknown.
Lead II
nitrate
Step 3
Develop a hypothesis regarding unknown.
Silver nitrate
Step 4
Complete side by side confirmation tests of
the unknown and the chemical identified in
the hypothesis.
68
Data: (continued)
Sodium
phosphate
Silver
nitrate
Lead II
nitrate
Copper II
sulfate
Sodium
nitrate
Sodium
Hydroxide
Potassium
iodide
Iron II
chloride
Calcium
Chloride
Sulfuric acid
Sodium
carbonate
Hydrochlori
c acid
Ammonium
Hydroxide
Red litmus
Blue litmus
Unknown #
Results:
Unknown number 1:
Unknown chemical’s number
________
Unknown chemical’s identity____________________________
Unknown number 2 (if a second is assigned):
Unknown chemical’s number
________
Unknown chemical’s identity____________________________
Discussion:
As per “How to write a lab report.”
Post-Lab Questions:
1. Write and balance and indicate phase of the reaction of silver nitrate with each of the other
chemicals used in this lab. Make sure to indicate phase of all compounds.
References:
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102
laboratory manual. Star Publishing Company: Belmont, California
Willbraham, A. Staley, D. Matta, M. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park,
CA
Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New
York, NY
69
Experiment 10
Chemical Reactions
Purpose:
To perform simple chemical reactions
To identify the products of these chemical reactions
To account for the reactions with balanced equations
To name the reactants and products
Background Information:
Evidence of a Chemical Reaction
A chemical reaction has taken place when a substance has changed its chemical composition.
The evidence that a chemical reaction has taken place can be one or more of the following
1. The appearance of gas (bubbles)
5. Production of light
2. The appearance of a solid (precipitate)
6. Flame
3. A change in color
7. A change in smell
4. A change in temperature
Reaction types
It is useful to be able to determine which kind of reaction is taking place for it will make the
prediction of the product that will form more easily.
The primary reaction types discussed in this class are:
1. Neutralization reactions
2. Precipitation reactions
3. Oxidation-reduction reactions (redox reactions)
a. Single replacement
b. Combination reactions
c. Decomposition reactions
Some reactions can be classified as more than one of the reaction types listed above.
Neutralization
A reaction between an acid and a base to produce salt and water is called a neutralization
reaction. In this experiment the only bases used will be those containing the hydroxide anion
OH-. The generalized formula for this kind of neutralization reaction is:
HA(aq)+ COH(aq) à HOH(l) and CA(aq)
This kind of reaction of an acid and a base always yields water (HOH or H2O) and a salt.
Example: Hydrochloric acid and sodium hydroxide react to produce sodium chloride and water.
Since sodium chloride is soluble, its formation cannot be seen, but there is evidence that a
reaction has taken place because the reaction mixture gets warmer.
NaOH(aq) + HCl à NaCl(aq) + HOH(l)
(Please note HOH is another way of writing H2O.)
70
The following reaction is a special type of neutralization reaction. This example illustrates how
any carbonate will react with a strong acid to produce salt, water and carbon dioxide.
Example:
Bubbles of gas are formed when sodium carbonate reacts with hydrochloric acid.
Na2CO3(aq) +HCl(aq) à CO2(g) + NaCl(aq) + H2O(l)
Carbon dioxide and water is produced when any carbonate or bicarbonate reacts
with any acid. This is also considered a neutralization reaction
Precipitation reaction
A reaction in which an insoluble substance is produced when two aqueous salt solutions are
mixed together is called a precipitation reaction. A salt is a compound that consists of a positive
ion, called a cation and a negative ion called an anion. When a salt dissolves in water, it
separates into its constituent cations and anions. This is an aqueous salt solution and will be
denoted with (aq).
A list of names and formulae of the commonly occurring ions follows this section. You are
solely responsible to learn each of the ions. Your success in this class depends heavily on this.
Also following this background section is a set of solubility rules which will help you identify
any insoluble solid precipitates formed during the reactions. You may also want to reference
the nomenclature laboratory section; this may help you during the laboratory activity.
Example:
A white precipitate is formed when solutions of silver nitrate and sodium chloride
are mixed.
Write and balance the equation for the reaction showing the appropriate phases such as (s),
(g), (aq) or (l).
A. Write the formula for materials that will react in this reaction (the reactants).
Silver nitrate is AgNO3
Sodium sulfite is Na2SO3
B. List all of the ions present in a solution of each of the two salts.
Ag+ + NO3- and 2Na+ + SO32C. Write the formulae for the possible products when the salts exchange partner ions.
(make sure that each cation is paired with a new anion)
2Ag+ pairs with SO32- and Na+ pairs with NO3Ag2SO3
and
NaNO3
silver sulfite
sodium nitrate
D. Use the solubility rules to determine which of the possible products is insoluble and precipitates.
The solubility rules (page 59) indicate that nitrates are soluble, so the sodium nitrate is soluble and
does not form a precipitate. Silver sulfate is affected by rule #5: all sulfites are in soluble except
those of ammonia and alkali metals. Silver sulfite forms a precipitate. The equation for the reaction
can now be written and balanced:
AgNO3(aq) + Na2SO3 (aq) à Ag2SO3 (s) + NaNO3(aq)
The (s) indicates a solid. This is a precipitate.
The (aq) indicates a aqueous solution, that is the two ions are dissolved in the water
E. Write and balance the equation for the reaction showing the phases such as (s), (g), (aq) or (l).
Example:
A white precipitate is formed when aqueous solutions of sodium phosphate and
calcium chloride are mixed. The final form of the equation would be:
2Na3PO4(aq) + 3CaCl2(aq) à Ca3(PO4)2(s) + 6NaCl(aq)
71
Oxidation-Reduction Reactions
Any reaction in which electrons are transferred between reactants in the formation of the products is
called an oxidation-reduction reaction (referred to as redox). Many reactions which also fit into
other classifications involve a transfer of electrons and should be classified as oxidation-reduction
reactions as well. Please note that any time an element is present as a reactant but then present in an
ionic form as a product this is a redox reaction.
When white hot sodium metal is placed into a flask filled with chlorine gas a very bright
Example:
reaction takes place. Both reactants are present as uncharged elements. After the
reaction a salt is present which is made up of cations and anions. Each sodium has lost
an electron to become a cation, each chlorine has accepted an electron to become an
anion. Electrons were transferred from the sodium (leaving it with a positive charge) to
the chlorine (making it negative).
2Na(s) + Cl2(g) à 2NaCl(aq)
Redox: Single replacement reactions
A very common form of a single replacement reaction is a reactive elemental metal reacting with a
less reactive metal salt. In this single reaction it is possible to predict which metals can be
substituted based on the metal’s reactivity. How metals react in a single replacement reaction
depend on the metals’ position on the Activity Series.
lithium
potassium
strontium
calcium
sodium
The Activity Series of the metals is an invaluable aid to predicting the
products of replacement reactions. It also can be used as an aid in
predicting products of some other reactions. Pay attention to the notes
below as they are provided to help you make better use of the activity
series than just the list of metals by themselves.
------------------------------------------------------------
magnesium
aluminum
•
CARBON
zinc
chromium
•
-----------------------------------------------------------
iron
cadmium
cobalt
nickel
tin
lead
•
•
----------------------------------------------------------
HYDROGEN
antimony
arsenic
bismuth
copper
--------------------------------------------------------
mercury
silver
paladium
platinum
gold
•
•
Each element on the list replaces from a compound any of the
elements below it. The larger the interval between elements,
the more vigorous the reaction.
The first five elements (lithium - sodium) are known as very
active metals and they react with cold water to produce the
hydroxide ion and hydrogen gas.
The next four metals (magnesium - chromium) are considered
active metals and they will react with very hot water or steam
to form the oxide and hydrogen gas.
The next six metals (iron - lead) replace hydrogen from HCl
and dilute sulfuric and nitric acids. Their oxides undergo
reduction by heating with H2, carbon, and carbon monoxide.
The metals lithium - copper, can combine directly with oxygen
to form the oxide.
The last five metals (mercury - gold) are often found free in
nature, their oxides decompose with mild heating, and they
form oxides only indirectly.
These single replacement reactions can also be completed using an
acid. See example 2.
72
Example:
When Zn metal is reacted with aqueous iron II sulfate, the zinc metal dissolves and an
orange solid forms.
Zn(s) + FeSO4(aq) à ZnSO4(aq) + Fe(s)
Because these two metals are not far apart this reaction is not very energetic.
Example:
Bubbles of gas are formed when solid magnesium reacts with hydrochloric acid.
Mg(aq) +HCl(aq) à MgCl2(aq) + H2(g)
In this reaction the solid magnesium appears to disappear as it dissolves into an
aqueous solution of magnesium chloride. Because magnesium is much higher in the
series than hydrogen this reaction is fairly vigorous.
Note: These single replacement reactions are the base of the electrochemistry of batteries. For
instance batteries use the following materials to produce electricity: lithium and carbon (lithium
ion), nickel cadmium (NiCad), Zinc and Carbon (Heavy Duty).
Redox: Decomposition Reactions
A single substance breaks apart to give two or more new substances. The generalized formula
for this reaction is
A à B + C
Example:
Hydrogen peroxide decomposes to form liquid water and oxygen gas. The hydrogen
peroxide is initially dissolved in water this is called an aqueous solution.
Hydrogen Peroxide (aqueous) à water (liquid) + oxygen (gas)
The balanced equation is
Fe
2H2O2(aq) à 2H2O(l) + O2(g)
Please note the presence of the Fe above the reaction arrow. When something is located above the
reaction arrow it is used as a catalyst in the reaction. Catalysts speed up a reaction but are not
consumed in the reaction.
Redox: Combination reactions
Two or more substances combine together to form one new substance. The generalized formula
for this reaction is
A+BàC
Example:
When charcoal burns, carbon (charcoal) reacts with oxygen to form carbon dioxide
Carbon(solid) + oxygen (gas) à carbon dioxide (gas)
C(s) + O2(s) à CO2(g)
73
Names and Formulae for Commonly Occurring Ions
(I would suggest you make one 3x5 flash cards for each ion and then memorize the cards. The
cards should have the name on one side and the formula on the other)
Cations
1+
Proton
Lithium ion
Sodium ion
Potassium ion
Silver ion
Copper I
Ammonium
+
H
Li+
Na+
K+
Ag+
Cu+
NH4+
Anions
2+
Magnesium ion
Calcium ion
Strontium ion
Barium ion
Mercury I
Mercury II
Copper II
Zinc ion
Iron II
Tin II
Lead II
Cobalt II
Nickel II
Manganese II
2+
Mg
Ca2+
Sr2+
Ba2+
Hg22+
Hg2+
Cu2+
Zn2+
Fe2+
Sn2+
Pb2+
Co2+
Ni2+
Mn2+
Please note that there are other metal ions. For
instance the formula for Manganese III is Mn3+.
3+
Aluminum ion
Chromium III
Iron III
Al3+
Cr3+
Fe3+
2-
-
Hydride H
Oxide
Hydroxide OHPeroxide
Fluoride FSulfide
Chloride Cl
Sulfite
Bromide BrSulfate
Iodide IThiosulfate
Hypochlorite ClOCarbonate
Chlorite ClO2
Chromate
Chlorate ClO3Dichromate
Perchlorate ClO4Oxalate
Cyanide CN
Cyanate CNOThiocyanate SCNNitrate NO3Nitrite NO2Permanganate MnO4Acetate CH3COOHydrogen carbonate HCO3(bicarbonate)
Hydrogen Sulfate HSO4- (bisulfate)
Peroxide O-
4+
3-
Tin IV Sn4+
Lead IV Pb4+
O2O22S2SO32SO42S2O32CO32CrO42Cr2O72C2O42-
Nitride N3Phosphide P3Phosphate PO43-
Common Acids
Hydrochloric acid
Sulfuric acid
Nitric acid
Phosphoric acid
Carbonic acid
Acetic acid
1-
Common Bases
HCl
H2SO4
HNO3
H3PO4
H2CO3
CH3COOH
Sodium hydroxide
Potassium hydroxide
Calcium hydroxide
Ammonia
74
NaOH
KOH
Ca(OH)2
NH3
Solubility Rules (I would suggest you memorize rule 1-3 at least)
1. Salts containing sodium (Na+), potassium (K+), and ammonium (NH4+) and acids are always
soluble.
2. Salts containing nitrates (NO3-), chlorate (ClO3-), perchlorate (ClO4-), and acetate (CH3COO-)
are always soluble.
3. All chlorides (Cl-), bromides (Br-), and iodides (I-) are soluble except for those of silver (Ag+),
lead II (Pb2+), and mercury I (Hg22+), which are insoluble.
4. All sulfates (SO42-) are soluble except of those of strontium (Sr2+), barium(Ba2+), mercury I
(Hg22+), mercury II (Hg2+), and lead (Pb2+). The sulfate salts of calcium (Ca2+), silver (Ag+) are
moderately soluble.
5. All hydroxides (OH-) are insoluble expect those of alkali metals, which are soluble, and
hydroxides of calcium (Ca2+) , barium (Ba2+), and strontium (Sr2+), which are moderately
soluble.
6. All sulfites (SO32-), carbonates (CO32-), chromates (CrO42-) and phosphates (PO43-) are insoluble
except those of ammonium (NH4+) and alkali metals.
7. All sulfides (S2-) are insoluble except those of ammonium (NH4+), the alkali metals and the
alkaline earth metals, which are soluble.
Please note:
§ Soluble means that the solid (s) will completely dissolve in water making an aqueous
solution (aq).
§ Moderately soluble means that some of the material will dissolve, although some portion of
it will remain as a solid.
§ Insoluble means that the solid will not dissolve and remain as a solid.
Pre-Lab Questions:
The stepwise process to writing a balanced reaction from a word description is
a.
b.
c.
d.
Write the formula for materials that will react in this reaction (the reactants).
List all of the ions present in a solution of each of the two salts.
Write the formulae for the possible products when the salts exchange partner ions.
Use the solubility rules to determine which of the possible products is insoluble and
precipitates.
e. Write and balance the equation for the reaction showing the appropriate phases such as (s),
(g), (aq) or (l).
Show the stepwise process towards writing a balanced equation for each reaction below.
1. Sodium phosphate solution will react with calcium chloride solution to form a white
precipitate. (To confirm your balanced equation check the background information.)
2. A yellow precipitate is formed when solutions of silver nitrate and sodium phosphate are
reacted.
3. Bubbles of gas are produced when a solution of potassium hydrogen carbonate (potassium
bicarbonate) reacts with sulfuric acid.
4. A white precipitate appears and the solution gets warm when solutions of barium hydroxide
and phosphoric acid are mixed.
75
Materials:
Bunsen burner
crucible
3 M KI
1.0 M HNO3
0.2 M AgNO3
0.2 M BaCl2
0.2 M NaOH*
ring stand
Magnesium metal
2.0 M NaOH*
NaHCO3 (solid)
0.2 M KBr
0.2 M CuSO4
Saturated Ca(OH)2
clay triangle
3% hydrogen peroxide
1.0 M H2SO4
0.1 M NaHCO3
0.2 M Na2SO4
0.2 M NaCl
Iron Nail
* please note the two different consintrations of NaOH. Do not use the wrong one during your
different trials.
Procedure:
*******Note well*******
In all cases, get material from the hood in labeled test tubes and then return to your laboratory
bench. Make observations of the unreacted reagents, then conduct the experiment at your lab bench
continue to make observations of the reaction.
Combination Reactions
1a. and 1b. Reaction of magnesium and nitrogen and the reaction of magnesium and oxygen.
When you burn magnesium it reacts both with oxygen in the air (≈ 20%) and nitrogen in the air
(≈75%). You will only burn one strip of magnesium. When it burns it will complete 2 reactions.
Set up a Bunsen burner, ring stand, and a clay triangle. Support a dry crucible (dry is far more
important than clean in this experiment) on the triangle. Place 1 inch of loosely coiled magnesium
metal in the crucible and heat the bottom of the crucible until the metal burns. It is important to
make careful observations in this step as you will need them again in later experiments.
!!!DO NOT LOOK DIRECTLY AT THE BURNING MAGNESIUM!!!
Once the magnesium has completely burned, allow the crucible to cool to room temperature.
Disturb the ash by stirring gently with a glass rod and describe the appearance of the ash in your
laboratory notebook
Decomposition Reaction
2. Decomposition of hydrogen peroxide
This is the same reaction covered in the discussion. Without a catalyst this reaction proceeds
extremely slowly. A catalyst is a substance that speeds up a reaction without being consumed by
the reaction. You will use a catalyst to speed up the decomposition of hydrogen peroxide so that the
bubbles of oxygen being produced can be seen.
Place about 2 mL of 3% hydrogen peroxide in a small test tube. Add 3-4 drops of KI solution. Stir
with a glass rod. Note the rate of evolution of bubbles of oxygen. Write careful observations into
your laboratory notebook.
Neutralization reactions
Check each reactant with litmus paper before mixing, and check each mixture with litmus paper
after the reaction has taken place. The proper technique is to dip a clean glass stirring rod into the
solution and then touch the stirring rod to the litmus paper. Write the results of the tests with litmus
paper into your laboratory notebook.
76
3. Sodium hydroxide and sulfuric acid
Place 2mL of aqueous 2.0 M sodium hydroxide in a small test tube. Do not measure this, simply
use 40 drops or approximately 2mL as shown by your laboratory instructor. Carefully add 2mL
dilute 1.0 M sulfuric acid. Stir with a glass stirring rod. Note any temperature changes by touching
the test tube. It is not necessary to record the exact temperature change.
4. Nitric acid and sodium bicarbonate
Place about 2 grams of solid sodium bicarbonate into an evaporating dish and then place about 2mL
of aqueous nitric onto the solid sodium bicarbonate. Please check the pH of the resulting solution
using litmus paper.
Please repeat experiment with aqueous sodium bicarbonate in a test tube.
Precipitation Reactions
Set up a warm water bath using a Bunsen burner, beaker, distilled water, a ring stand and a wire
gauze.
Do not let the water boil.
Set up the small test tube rack containing six clean test tubes. Number each tube either with a
pencil on its label spot or with a grease pencil. As you record the results of each reaction in your
notebook, be sure to include the test tube number. Add about 2mL of each of the two reagents. It is
not necessary to measure the volume of reagents.
Write down all observations including colors of original reagent solution, amount of precipitate
formed and color of precipitates.
5. Silver nitrate and potassium bromide
6. Sodium sulfate and barium chloride
7. Copper II sulfate and sodium hydroxide
8. Sodium bicarbonate and barium chloride
9. Copper II sulfate and sodium chloride
10. Calcium hydroxide and sulfuric acid (test before and after reaction with litmus paper)
After you have run each reaction heat all test tubes in a water bath and make additional
observations. Carefully record any and all changes and events that you observe, including a rough
estimate of the time involved.
!!! NEVER HEAT A TEST TUBE DIRECTLY IN A FLAME!!!
Single replacement reactions
11. Copper II sulfate and iron
Place about 2ml of copper II sulfate into a test tube. Then polish an iron nail with steel wool.
Now place the polished nail into the copper II sulfate. Make careful observation regarding the
rate of reaction and the differences between the submerged and not submerged portion of the
nail.
12. Magnesium with hydrochloric acid
Place a 1 cm strip of magnesium metal into an evaporating dish. Add about 2 mL of dilute
hydrochloric acid. Write all observations into your laboratory notebook. Please recover and
rinse the unused Mg ribbon to be used in experiment 1.
77
Data: This data table is to be completed before you come to class.
I expect four experiments
per page; please align your table as follows (landscape). Use a ruler, be neat and above all make
sure your table is arranged in numerical order (this is not necessarily the order in which you
conduct the experiments). As a safety percaution all of the information in the shaded boxes for all
trials are to be filled out before coming to class.
1a
1b
2
3
Nitrogen
Oxygen
Magnesium
Mg
N2
O2
Mg
Reactant
Formula
Magnesium
Hydrogen
peroxide
Sodium
hydroxide
Sodium
bicarbonate (s)
& Nitric acid
Sulfuric acid
4a
Sodium
bicarbonate (aq)
4b
& Nitric acid
Preobservations
Observations
Product
names
Product
Formula
Balanced equation (Include
state)
78
Calculations:
No calculations for this lab
Results:
See data table for the balanced equations and names of products.
Discussion: (please only address the following)
In this experiment you did not perform any reactions labeled as oxidation-reduction. Explain how
the reaction of magnesium burning in air and the reaction of magnesium with hydrochloric acid
could both be called oxidation-reductions. Also explain why the precipitation reactions could not
have been described as oxidation-reduction reactions.
Post-Lab Questions:
For each pair listed below, write a balanced equation and name all reactants and products. All
reactions except 2c. involve aqueous salts listed above.
1.
a. AgNO3 and KI
b. CaCl2 and Na2SO4
c. H3PO4 and NaOH
2.
a. barium hydroxide and nitric acid
b. calcium hydroxide and hydrochloric acid
c. burning aluminum metal and oxygen
References:
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102
laboratory manual. Star Publishing Company: Belmont, California
79
Experiment 11
Stoichiometry
Purpose:
Determine the limiting reactant and predict mass of products.
Hypothesis:
Develop a hypothesis regarding the conservation of mass in this reaction.
Background Information:
In this experiment you will compare the experimentally measured and the theoretical amount (in grams) of
barium sulfate produced by the reaction of a barium chloride dihydrate solution and a potassium sulfate
solution. The reaction involved is shown below.
K2SO4(aq) + BaCl2·2H2O(aq) à BaSO4(s) + 2KCl(aq) + 2H2O(l)
(11.1)
The barium sulfate forms a precipitate while potassium chloride remains in solution.
The barium sulfate precipitate is collected by gravity filtration because the water and the potassium chloride
will pass through the filter paper (see figure 1).
Figure 1 Gravity Filtration and filter paper preparation
The filter paper and precipitate are placed together in a beaker or on a watch glass. The filter paper and
precipitate can then be air or oven dried. The experimental mass of barium sulfate can then be determined
and compared to the theoretical mass of barium sulfate.
This reaction is a rarity because conversion to product is essentially complete. Most reactions do not yield
complete conversion to product. In this case, however, the experimental and the theoretical yield should be
fairly close. In fact the experimental amount of barium sulfate is often slightly larger than the theoretical
amount of barium sulfate because the precipitate is contaminated with a small amount of the filtrate which
contains potassium chloride. This contaminated filtrate could be eliminated by washing with water. But this
process is very time consuming.
At the end of the experiment you will need to determine the accuracy of the experiment. In this experiment
accuracy is defined as the agreement between the experimental and theoretical values, in grams, of barium
sulfate. We will measure accuracy by the absolute error and the relative error.
Absolute error is defined by equation (11.2).
Absolute error = E = |m-t|
Where E = absolute error
m = experimentally measured value
t = theoretical value
80
(11.2)
% relative error is defined by (equation 11.3).
% relative error = |m-t|/t *100 (11.3)
A related equation determines the % yield (equation 11.4).
m/t*100= % yield (11.4)
Limiting reactant: You will be asked to determine which material limited the reaction. Think of
this as the material that ran out first. For instance in building bicycle if you have 200 bicycle
frames and 300 wheels. Which runs out first, assuming each frame needs two wheels. The answer
may seem obvious. However when you start working with grams of chemicals it can seem harder.
Back to the example, another important question might be how many seats are needed?
Here is a step by step explanation of how to solve a limiting reactant problem
1. Write and balance equation, indicate state.
2. Do dimensional analysis to determine how much product can produced if all of the first
reactant is used.
3. Do dimensional analysis to determine how much product can produced if all of the second
reactant is used.
4. The limiting reactant is the reactant that produces the least amount of product.
Example:
When 2.045 grams of solid sulfur reacts with 1.807 grams of gaseous oxygen react to form gaseous
sulfur dioxide.
Step 1.
2S(s) + O2 (g) à 2SO (g)
Step 2. Determine how much sulfur dioxide is produced if all of the sulfur is consumed.
2.045 g S
1 mol S
2 mol SO
48.07 g SO
= 3.066 g sulfur monoxide
32.06 g S 2 mol S
1 mol SO
Step 3. Detemine how much sulfur dioxide is produced if all of the oxygen gas is consumed.
1.807 g O2 1 mol O2
2 mol SO
48.07 g SO
= 5.429 g sulfur monoxide
32.00 g O2 1 mol O2
1 mol SO
Step 4. When all the sulfur is consumed the least amount of product is produced. Thus the sulfur is
the limiting reactant. (This may seem odd because there is more grams of sulfur this is due to the
balanced equation and the formula weights of each reactant.)
Pre-Lab Questions:
1. Use the provided experimental data to answer the following questions. 0.987g of potassium
sulfate are reacted with 0.897 grams of barium chloride dihydrate.
a. Write and balance the equation, indicate state.
b. Determine how much product is produced when the first reactant is completely consumed.
c. Determine how much product is produced when the second reactant is completely consumed.
d. Indicate which reactant is the limiting reactant and how much insoluble product is produced.
3. In the lab it indicates in step 4 and 8 to add approximately 30 ml water. The water is not taking
part in the reaction, it might seem likely that the amount of water is not important. It turns out that
the amount of water is very important. Why might the volume of water added really make a
difference?
81
4. Make the following table on a whole sheet of paper (please fill up a landscape oriented piece of
paper with the following table.) Fill in the answers.
Effect on
theoretical
value
Effect on
experimental
value
Effects on %
yield < or >
Explanation
(leave the most
room for this)
BaCl2* H2O left on balance pan
BaCl2* H2O left in beaker or on thermometer
BaSO4 left in beaker
BaSO4 runs through filter
K2SO4 Left on filter paper
BaSO4 crumbles and falls out of filter
Filter is not completely dry
Materials:
Potassium sulfate (approximately 1.0g)
Barium chloride dihydrate (approximately 1.0g)
Wooden filtration rack
Filter paper
Wash bottle
Bunsen burner
Wire gauze
Ring stand
Ring clamp
Grease pencil
Barium sulfate waste receptacle
Procedure:
1.
2.
3.
4.
Mass a clean dry 150 ml beaker. Then record mass directly into lab notebook.
Obtain approximately 1.0 grams of K2SO4 and place in beaker.
Record the mass of the beaker and sample directly into the notebook.
Label beaker with your initials and the contents with pencil or tape. Add approximately 30 ml of
distilled water and stir the solution. Bring solution to a gentle boil.
5. Mass another clean dry 150 ml beaker. Then record mass directly into lab notebook.
6. Obtain approx. 1.0 g of barium chloride dihydrate and place in beaker.
7. Record the mass of the beaker and the sample directly into the lab notebook.
8. Add 30 ml of deionized water to beaker and dissolve the barium chloride dihydrate.
9. The precipitate is formed by adding the barium chloride solution very slowly to the hot sulfate
solution. To mix this reactants together very slowly use a disposable pipette or burette.
Consider trying one drop every few seconds. By mixing the solutions slowly it may avoid the
formation of very small barium sulfate particles.
10. Rinse the barium chloride dihydrate solution beaker with another 25 ml of D.I. water to remove
any residual barium chloride dihydrate and transfer this rinse solution to the potassium sulfate
solution beaker also. The white precipitate that forms is barium sulfate, this is what you want to
keep.
11. After the reagents have been mixed. Heat the solution over a Bunsen burner (do not boil), for
about 15 to 20 minutes. After this time the barium sulfate precipitate should be at the bottom of
the beaker under a clear solution. (This may be a good time to stop for the day.)
12. Obtain a sheet of filter paper; put your initials on the paper in pencil. Mass the filter paper to
the nearest 0.001 g on the same analytical balance previously used for your other masses and
record the results directly in your lab notebook.
13. Set up the gravity filtration apparatus described in Figure 1. Have an assistant or instructor show
you how to correctly prepare the filter paper. Make sure that you do not punch a hole in the
filter paper. If there is a hole in the filter paper, the barium sulfate will not be separated from
the water and the other product, potassium chloride.
82
NOTE: There will be a number of times during this experiment when you are not busy. Use this
“dead” time to begin the analysis of the data. Calculate the number of moles of each reagent.
Begin to complete the table in the Results Section. Note that the barium chloride is the dihydrate,
barium chloride* water. Be sure to use the correct molecular weight. Choose the limiting reagent.
On the basis of this result, complete the table in the Results Section. If you can complete this
calculation during the lab period, the subsequent lab report will be simplified greatly.
14. Pour most of the clear, hot solution into the filter. Do not let the liquid level rise above the edge
of the paper. Keep the remaining solution and solid hot while waiting for the filter to drain.
When only a few mL remain in the beaker, agitate to suspend the Barium sulfate and pour the
suspension into the filter. Do not let liquid level in filter rise above the edge of the paper.
15. Rinse the sides of the beaker with a few milliliters of deionized water, using a wash bottle.
Transfer the rinsing into the filter. Rinse the beaker as required until the last of the barium
sulfate is in the filter. This is a stopping point; if you must leave, place your funnel in a beaker
so the filtration can continue to completion until the following lab period.
16. When the filtration is complete, rinse the filter cake with a few milliliters of deionized water.
Do not let liquid level rise above the edge of the paper.
17. Air dry the filter paper & barium sulfate in your drawer until the following lab period then oven
dry overnight or as designated by the instructor.
18. Mass the filter paper & barium sulfate once they are dry on the same analytical balance
previously used. Record the mass directly into your notebook.
Observation
Data:
1. mass of beaker and potassium sulfate
mass of empty beaker
mass of potassium sulfate
g
g
g
2. mass of beaker and barium chloride
mass of empty beaker
mass of barium chloride
g
g
g
3. mass of filter paper and barium sulfate
mass of filter paper
mass of barium sulfate
g
g
g
Calculations:
Show your calculations for the theoretical mass of barium sulfate produced, the mass of potassium
chloride and the mass of the excess reagent. Use unit analysis and use the proper number of
significant figures. Clearly indicate the limiting reagent.
First find theoretical yeild
1. Calculate how much product is produced if all of the first reactant is used
2. Calculate how much product is produced if all of the second reactant is used.
3. Calculate how many grams of unused reactant is present at the end.
83
4. Clear show which material is the limiting reactant and how much insoluble product is formed.
(this is the theoretical yield)
Find Experimental Yield
5. Calculate the experimental yield. This is how much barium sulfate you actually recovered in
your filter.
Error Anlysis
6. Compare the theoretical and experimental masses of barium sulfate produced.
7. Calculate the absolute error. (11.2)
8. Calculate the relative error. (11.3)
9. Calculate the percent yield. (11.4)
Results:
State the theoretical yield. (from Table 1)
State the experimental yield.
Report the absolute error.
Report the relative % error.
Report the % yield.
Discussion:
Briefly summarize your results and discuss your relative errors.
If your percent yield is greater than 100 explain what factors may have caused this error. If your
percent yield is less than 90 percent explain what may have happened to cause this error. Be very
specific in this section. For example discuss the how the results would be effected for instance
losing material before or after the initial massing. What would happen if it was not completely
dried out? It is unsafe to place barium salts into the sink . How do you if it is safe to put the filtrate
down the sink (stuff that goes through filter). Answer the other discussion questions as presented in
your “How to Write a Lab Report.”
Post-Lab Questions:
Soda ash (sodium carbonate) is widely used in the manufacture of glass. Prior to the environmental
movement much of it was produced by the following reaction.
CaCO3 + 2 NaCl à Na2CO3 + CaCl2
Unfortunately, the byproduct calcium chloride is of little use and was dumped into rivers, creating a
pollution problem. As a result of the environmental movement, all of these plants closed. Assume
that 125g of calcium carbonate (100.09 g/mol) and 125 g of sodium chloride (58.44 g/mol) are
allowed to react. Construct a table similar to that shown in the calculation section.
a. Determine how many grams of useful sodium carbonate (105.99 g/mol) will be produced.
b. How many grams of useless calcium chloride (110.98 g/mol) will be produced?
c. You should also determine how many grams of excess reagent are left.
Show all calculations; use dimensional analysis; and use the proper significant figures. Use a table
if you wish to organize your data.
Clearly show how you chose the limiting reactant.
84
Experiment 12
Determining the Concentration of a Solution:
Beer’s Law
The primary objective of this experiment is to determine the concentration of an unknown
nickel (II) sulfate solution. You will be using the Colorimeter shown in Figure 1. In this device, red
light from the LED light source will pass through the solution and strike a photocell. The NiSO4
solution used in this experiment has a deep green color. A higher concentration of the colored
solution absorbs more light (and transmits less) than a solution of lower concentration. The
Colorimeter monitors the light received by the photocell as either an absorbance or a percent
transmittance value.
Figure 1
Figure 2
You are to prepare five nickel sulfate solutions of known concentration (standard solutions). Each is
transferred to a small, rectangular cuvette that is placed into the Colorimeter. The amount of light
that penetrates the solution and strikes the photocell is used to compute the absorbance of each
solution. When a graph of absorbance vs. concentration is plotted for the standard solutions, a direct
relationship should result, as shown in Figure 2. The direct relationship between absorbance and
concentration for a solution is known as Beer’s law.
The concentration of an unknown NiSO4 solution is then determined by measuring its absorbance
with the Colorimeter. By locating the absorbance of the unknown on the vertical axis of the graph,
the corresponding concentration can be found on the horizontal axis (follow the arrows in Figure 2).
The concentration of the unknown can also be found using the slope of the Beer’s law curve.
Objectives:
In this experiment, you will
•
•
•
•
Prepare NiSO4 standard solutions.
Use a Colorimeter to measure the absorbance value of each standard solution.
Find the relationship between absorbance and concentration of a solution.
Use the results of this experiment to determine the unknown concentration of another NiSO4
solution.
85
Materials:
two 100 mL beakers
Vernier computer interface
five 20 Í 150 mm test tubes
Vernier Colorimeter distilled water
one cuvette
50 mL of 0.40 M NiSO4
5 mL of NiSO4 unknown solution
10mL graduated cylinders)
tissues (lense paper)
test tube rack
Procedure:
•
•
•
•
To correctly use a Colorimeter cuvette, remember:
All cuvettes should be wiped clean and dry on the outside with a lense paper tissue.
Handle cuvettes only by the top edge of the ribbed sides.
All solutions should be free of bubbles.
Always postion the cuvette with its reference mark facing toward the white reference mark
at the top of the cuvete slot on the colorimeter.
Part 1: Determine correct LED for experiment
1. Obtain and wear goggles! CAUTION: Be careful not to ingest any NiSO4 solution or spill any
on your skin. Inform your teacher immediately in the event of an accident.
2. Connect the Colorimeter to Lab Quest
3. Fill cuvette ½ full with a stock solution of NiSO4, swirl and discard into a waste container.
4. Fill cuvette ¾ full with a stock solution of NiSO4 and place into the colorimeter.
5. Select the 430nm LED and record absorbance. Continue with the other LEDs (470, 565, 635nm)
and record aborbances for each.
6. The LED setting that provides the highest absorbance is the setting that you will continue to use
during the experiment (because the solution you are using is green the mostly likely LED setting
will produce red light)
Part 2: Calibrate colorimter
1. You are now ready to calibrate the Colorimeter. Prepare a blank by filling a cuvette 3/4 full with
distilled water, swirl and discard in a waste container.
2. Fill the cuvette ¾ full with distilled water.
3. Open the Colorimeter lid. Holding the cuvette by the upper edges, place it in the cuvette slot of
the Colorimeter. Close the lid.
4. Press the < or > button on the Colorimeter to select a wavelength determined from above (most
likely 635 nm). Press the CAL button until the red LED begins to flash. Then release the CAL
button. When the LED stops flashing, the calibration is complete.
Part 3: Serial dilution
1. Obtain about 50 mL of 0.40 M NiSO4 stock solution to a 100 mL beaker. Add about 50 mL of
distilled water to another 100 mL beaker.
2. Set up 5 test tubes in a test tube rack. Label test tubes carefully.
3. For sample #1 carefully measure about 2ml of nickel II sulfate in a 10ml graduated cylinder.
Record the actual volume collected. How much you have is not important but careful
measurements is very important. To this very carefully add distilled water until the volume
approaches 10 mL. Do not go above 10mL. How much you add is not important but a careful
measurement of the final volume is very important. Empty graduated cylinder into labeled test
tube.
4. Please refer to following table to repeat samples #2,3,4,5 each with a different volume.
86
Target volumes. Your actual volumes will vary.
Sample
#
1
2
3
4
5
Estimated vol
Estimated vol
Estimated
Estimated
0.40 M NiSO4 (mL)
Distilled H2O (mL)
Final Vol
Concentration (M)
10
10
10
10
10
0.08
0.16
0.24
0.32
0.40
2
4
6
8
~10
8
6
4
2
0
Part 3: Data collection of know concentrations
1. You are now ready to collect absorbance data for the five standard solutions. Click . Empty the
water from the cuvette. Using the solution in Test Tube 1, rinse the cuvette twice with ~1 mL
amounts and then fill it 3/4 full. Wipe the outside with a tissue and place it in the Colorimeter.
After closing the lid, wait for the absorbance value displayed on the monitor to stabilize.
2. Discard the cuvette contents as directed by your teacher. Rinse the cuvette twice with the Test
Tube 2 solution, 0.16 M NiSO4, and fill the cuvette 3/4 full. Wipe the outside, place it in the
Colorimeter, and close the lid. When the absorbance value stabilizes.
3. Repeat the Step 8 procedure to save and plot the absorbance and concentration values of the
solutions in Test Tube 3 (0.24 M) and Test Tube 4 (0.32 M), as well as the stock 0.40 M NiSO4.
4. In your Data and Calculations table, record the absorbance and concentration data pairs that are
displayed in the table. Using the data pairs create a graph.
5. From the data points on your graph chart a best fit line. From this best fit line use Y=mX+B to
write an equation for the line.
Part 4: Determination of unknown concentration
1. Obtain about 5 mL of the unknown NiSO4 in another clean, dry, test tube. Record the number
of the unknown in the Data and Calculations table. Rinse the cuvette twice with the unknown
solution and fill it about 3/4 full. Wipe the outside of the cuvette, place it into the Colorimeter,
and close the lid. Read the absorbance value displayed in the meter. Record the data.
2. Using the equation for the line and the unknow data find the concentration for the unknown.
Data and Calculations:
Trial
0.40 M NiSO4
(mL)
Final Vol
Distilled H2O
(mL)
Conc (mol/L)
Absorbance
1
2
3
4
5
6
Unknown number ____
Concentration of unknown
mol/L
87
Calculations:
Use M1V1=M2V2 to find the concentration for each trial.
M1 = Molarity intial V1 = Volume final
M2 = Molarity final
V2 = Volume final
Based on your data us Y=mX+b to find the slope of the line.
Rearrange equation to solve for X. Insert the value of the unknown as Y (absorbation) solve for X.
Results:
Graph the known values
Place the unknown material on the graph
Report the unknown concentration and ID number.
Discussion:
Discuss how this technique could be used in a practical application in either an industry, research or
medical application. Propose an experiment where you could use this technique to measure the
concentration of something you are interested in.
88
Experiment 13
Titration
Purpose:
Develop titration skills. Use the titration technique to determine the concentration of an unknown
solution.
Background:
Solid sodium hydroxide has a tendency to absorb water from the atmosphere. The longer sodium
hydroxide is exposed to air the more water it absorbs. This characteristic of sodium gydroxide makes it
impossible to make a standard solution by weight alone.
For instance if you wanted to make a 1 molar solution of sodium hydroxide all that is needed is to
demtermine the molecular weight of Sodium Hydroxide (40.00 grams/mole). Thus you should be able
to just place 40.00 grams of NaOH into a container and then bring the water level up to 1.00 liters.
However this does not work. If you did what was previously outlined a less then 1.00 molar solution is
likely to be created.
Because when 40.00 grams of sodium hydroxide is massed there an unknown quantity of water that has
been absorbed into the NaOH crystals. In what is supposed to be 40.00 grams of NaOH there might
only be 32.00 grams of NaOH and 8 grams of water.
To determine the actual concentration of NaOH in a solution it is necessary to complete a stoichiometic
reaction called a titration. A titration is when a known concentration of a substance is used to determine
an unknown concentration. There are many types of titrations. To determine the concentration of NaOH
an Acid-Base Titration will be completed.
The Nature of Acids and Bases
There are multiple definitions of what constitutes an acid and a base. However the most basic
definition will be sufficient for this laboratory activity.
The Arrhenius Acid-Base definition asserts that:
• An acid contains hydrogen and dissociates in water to yield H2O and H+ (sometimes written
as H3O+).
• A base contains an OH group and dissociates in water to yield OH-.
Acids and bases in general focus on the H+ (proton) concentration in water. This concentration is
expressed in moles/liter. But because this occurs in water it is important to note that a glass of
water just sitting there minding its own business, will undergo the following reversible reaction all
on its own without any help.
H 2O
à
ß
H+ + OH-
Please note that the arrow towards products is very large. This is to indicate that the reaction
from water to its ions is pretty infrequent. But it does happen. It is possible to measure just how
much H+ has dissociated. In any given sample of water the H+ concentration is
1.00 x10-7 mole H+/liter of water
+
That is not a lot of H floating around in the water. It was decided that a scale should be made to
help people understand how acidic or basic things are. It was further decided that on this scale
water should be considered neither acidic, nor basic, it would be neutral. Thus the scale should be
easy to understand.
89
We use the pH scale for this purpose. It goes from around 0 to 14. Small numbers on this scale
have a very high proton concentration and a low hydroxide ion concentration and are called acidic.
Large numbers on this scale have a high hydroxide ion concentration and a low proton
concentration and are called basic.
pH
0
1
2
3
4
5
7
6
8
9
10
11
12
13
14
Protons
A lot of H+
1.0 M H+
Equal to OH1x10-7M H+
Practically no H+
1x10-14M H+
hydroxide
Practically no OH1x10-14M OH-
Equal to H+
1x10-7M OH-
A lot of OH1.0 M OH-
So how do you go from some simple scale of 0-14 to a molar concentration? To do this you need to
understand what pH means.
The p= the –log
The H= the H+ concentration in mol/liter (which is often written as [H+])
pH= -log[H+]
If you want to know what the pH of a solution that has a H+ molarity of 1.0 x10-7 mole/liter (or
[H+]=1.0x10-7 M) all you do is
pH=-log1.0x10-7
The calculator syntax for converting concentration to pH the TI 80 series calculator is
(-)
LOG
1
EE
(-)
7*
ENTER
*The number following the log is the concentration of the solution in moles/liter H+
To calculate the molarity from the pH you just do it in reverse (though it is slightly harder).
Let us find the concentration of H+ in a solution with a pH of 7
pH= -log[H+]
7= -log[H+]
(-1)*7= (-1)*-log[H+]
-7= log[H+]
to get rid of the log you 10x both sides (10log = 1)
10-7 = 10log[H+]
10-7= 1*[H+]
and
1x10-7= [H+]
The calculator syntax for pH to concentration on the TI 80 series calculator is
2nd
*
LOG(10x)
7*
(-)
ENTER
is the pH of any solution you are interested in find the H+ concentration of
Because the pH scale is a logarithmic function the difference in pH between pH 7 and pH 6 is a
factor of 10. That is, a solution with a pH of 6 has 10 times more H+ than a solution of pH 7.
0
1
2
3
4
5
6
7
10,000,000 1,000,000 100,000
10,000
1000
100
10
1 unit H+
90
A solution of pH 0 has 100,000,000 times more H+ as a solution with a pH of 7. The same is
true for numbers greater than seven each pH level increase H+ concentration is reduced by a factor
of ten and the OH- concentration is increased by a factor of 10.
Indicators
An acid base indicator is a weak organic acid whose acid form, which will be denoted as
HA(indicator) is a different color than its base form A-(indicator), with the color change occurring over
a specific pH range. Typically, one or both forms are intensely colored, so only a tiny amount of
indicator is needed and the presence does not affect the pH of the solution. Indicators are used for
approximate pH monitoring in acid base titrations or in reactions.
HA(indicator)
à
ß
H+ + A-(indicator)
The double arrows of this reaction indicate that it is a reversible reaction. A reversible reaction can
go in either direction. The direction that is favored is often based on which side has the highest
concentration. In the above example if there is a lot of H+ then the direction is likely to be moving
toward the reactants. If the H+ concentration is very small the reaction will move in the direction of
products.
Consider phenol red as an indicator.
When it is in the HAindicator form it is yellow.
When it is in the A-indicator it is red.
When you place phenol red in an acid solution there are very large quantities of H+. Because of the
high concentrations of H+ the H+ on the indicator will not dissociate and the reaction moves in the
direction of reactants. Thus there are high concentrations of HAindicator.
Yellow
Red
HAindicator ß H+ + A-indicator
à
When you place phenol red in a basic solution there is a very low concentration of H+. This forces
the reaction to favor products.
Red
Yellow
HAindicator à H+ + A-indicator
ß
When phenol red is in high concentrations of H+ (an acidic solution) the primary species is
HAindicator and the color is yellow.
When phenol red is in low concentrations of H+ (a basic solution) the primary species is A-indicator
and the color is red.
91
Some commonly used indicators:
1
Thymol Blue
2
Red Orange
Methyl red
Bromothymol Blue
3
4
pH
5
6
Yellow
7
8
9
Green
10
11
Blue Green
12
Red Orange dfds Yellow
Yellow sdf
Phenol red
Green
Blue
Yellow sdfs Red Orange
Clearsdf Pink
Phenolphthalein
Acid Base Reactions
A frequently used definition of the term acid and base is attributed to BrØnsted and Lowry
An acid is any substance which donates protons (H+ ions)
A base is any substance which accepts protons
When an acid dissolves in water, some or all of the protons are given up. These bare H+ ions quickly
react with water to form an ion usually written as H3O+ (hydronium ion). Thus when hydrogen chloride
dissolves in water to form Hydrochloric acid the process maybe presented as:
HCl + H2O à Cl- + H3O+
This dissolution process is an acid-base reaction with hydrogen acting as an acid and water acting as a
base. The acid is donating protons (HClàCl- + H+) and the base is accepting protons (H2O + H+ à
H3O+). When this solution is mixed with another base which will accept protons it is the hydronium ion
H3O+, which acts as an acid.
When a weak acid dissolves in water, the proton transfer is not complete. For example, the weak acid
HNO2, dissociates partially. Only a small fraction of the HNO2 molecules transfer their protons to
water.
HNO2 + H2O à NO2- + H3O+
Many bases form the hydroxide ion, OH- in water.
NaOH à Na+ +OHThe reaction of an acid with a base is a proton transfer which converts both reacting species into
water molecules.
H3O+ + OH- à 2H2O
The overall reaction of hydrochloric acid with sodium hydroxide can be written:
H3O+ + Cl-+ Na+ + OH- à 2H2O + Na+ + ClThe formula of the salt produced is NaCl and the name of the salt is sodium chloride.
Similarly, when a weak acid reacts with hydroxide ions, the H+ ions are pulled of the acid. The reaction
is complete when all of the acid has lost its H+ ions or when the base is used up. An example is given
below for the reaction of NaOH and HNO2:
HNO2 + NaOH à H2O + Na+ + NO2-
92
Titration
A titration is a process used to determine the volume of a solution needed to
react with a given amount of another substance. In this experiment, you will
titrate hydrochloric acid solution, HCl, with a basic sodium hydroxide
solution, NaOH. The concentration of the NaOH solution is given and you
will determine the unknown concentration of the HCl. Hydrogen ions from
the HCl react with hydroxide ions from the NaOH in a one-to-one ratio to
produce water in the overall reaction:
H+(aq) + Cl–(aq) + Na+(aq) +OH–(aq) àH2O(l) + Na+(aq) + Cl–(aq)
Typical titration setup
Titration calculation
1. Determine moles of known (acid in this case)
2. Convert moles of known to moles of unknown (acid to moles of base)
3. Convert moles of unknown to Molarity (to find molarity of the base)
For this example assume that 10.00 mL of a 0.200 Molar solution of HCl was
used to neutralize 15.25 mL of sodium hydroxide.
Moles of acid (the known)
The number of moles of acid can be found using the morality and volume data.
Thus:
10.00 mL HCl
1 Liter HCl
0.200 moles HCl
=2.00 x10-3 moles HCl
1000 mL HCl
1 Liter HCl
Equation 1a
Moles of Base (the unknown)
From the balanced equation
NaOH + HCL à NaCl + H2O
please observe that for every one mole of NaOH 1 mole of CH3COOH is neutralized
Thus:
2.00 x10-3 mol HCl
1 mol NaOH
1 mol HCl
Equation 1b
-3
= 2.00x10 mol of NaOH
Convert moles of base to molarity
Now that the moles of base has been determined it is not possible to use the volume of base
measured using the burette to determine the molarity. Keep in mind molarity is a measure of the
number of moles per liter of solution.
Thus:
1000 mL NaOH
2.00x10-3 mol of NaOH
= 0.13114 = 0.131 M NaOH
15.25 mL NaOH
mole
liter
1 liter NaOH
Equation 1c
Or you can just do it all at once
10.00 mL HCl
15.25 mL NaOH
0.200 mol HCl
1000 mL HCl
1 mol NaOH
1 mol HCl
1000 mL NaOH
1 liter NaOH
= 0.131 M NaOH
Equation 1d
From the data in the
lab
Known molarity of
the acid
From the balanced
equation, it is not
always a 1:1
relationship
Converting mL to
Liters is necessary
M= Mole/liter of solution
93
Statistical Analysis
The first is a rough titration this is only to determine the approximate quantities needed. This data
is discarded. You will run at least three experimental trials. Then statistically analyze these three
trials to determine the mean and the average deviation from the mean.
Statistics allow us to determine the quality of the data. One reason random errors occur is due to the
estimated nature of all measurements. Errors occur whenever measurements are made. It is useful
to know how large the errors are. In this experiment precision of the results will be determined.
First perform the experiment three or more times. Then calculate the average or mean as shown in
Equation 3.
_
X1 + X2 + X3
Mean = X
n
=
Where X1 is the value determined from experiment trial 1 etc
n= the number of trials
Equation 3
If an experiment has good precision, all of the results will be clustered closely around the mean or
average. If the experiment has poor precision, the results will be widely scattered around the mean.
Precision can be determined by calculating the deviation from the mean. Deviation is defined in
equation 4 as:
_
i takes on the values of 1,2,3… for each trial
δi = |Xi –X|
Equation 4
_
X = the mean for all trials
δi
= the deviation for measurement i
Xi = the value of measurement i
The average deviation provides a measurement of how closely each of your datas points are to each
other. This is defined below in equation 5
_
|δ 1| + |δ 2| + |δ 3|
Average deviation = δ =
n
Equation 5
Where
|δi| = absolute value of deviation (no algebraic sign is used)
n = number of measurements
Precision is often calculated as the relative average deviation. To determine the overall precision
for the experimental data it is necessary to take the average deviation and divide it by the mean
(found in equation 3) then multiple this number by 100. See equation 6.
_
Equation 6
δ
_____
Precision = relative average deviation = X * 100
94
If the relative average deviation is 5% or less, the data is acceptable. If the relative average
deviation is 2% or less, the precision is good. If the precision is 1% or less, the quality of the data is
outstanding.
Pre-Lab Questions:
Write the balanced equation for each of the following acid base reactions. In each case, you are
expected to predict the formula of the salt produced.
1. Nitric acid reacted with sodium hydroxide
2. Sulfuric acid reacted with potassium hydroxide
3. Acetic acid reacted with calcium hydroxide
Materials:
NaOH unknown concetration ~ 0.200 M
Burette clamp
Pipette pump
Phenolphthalein
150 mL beaker
Burette
10.00 volumetric pipette
Hydrochloric acid known concentration
125 mL Erlenmeyer flask
Ring Stand
Procedure:
1. Obtain a burette, burette clamp and large ring stand from the lab bench. Set up material as per
instructor’s demonstration. Describe this laboratory apparatus in your procedure section.
2. Record the molarity of the hydrochloric acid solution directly into your lab notebook.
3. Obtain in a clean and dry beaker enough sodium hydroxide to fill the burette, Only take what
you need. Do not waste the sodium hydroxide. If the class runs out there will be no more to
complete your experiments.
4. Ask the instructor if the burette needs to be cleaned. If the burette needs cleaning rinse the
burette three times with 10 mL of the solution. After each rinsing, allow the solution to drain
through the stopcock. Then fill the burette so that the liquid level is above the zero line.
5. With the burette in the clamp place the zero line at eye level and the end of the burette over an
empty beaker. Slowly drain the solution until the liquid level is between 1.00 and 2.00 mL.
When finished, wipe the tip of the burette.
6. Transfer 10.00 mL Hydrochloric acid to a clean but not necessarily dry 125 mL Erlenmeyer
flask. Use the technique demonstrated in class. Describe this technique in your lab note book in
the procedure section.
7. Add two to three drops of phenolphthalein indicator to the Erlenmeyer flask. Now add
approximately 50 mL of distilled water to the flask. Make sure to use the same amount of
phenolphthalein and water for each trial.
8. Record the NaOH burette reading as the initial burette volume. Proceed to titrate the sample.
The indicator will change from colorless to pink when the endpoint point has been reached (the
color will change when the moles of H+ are the same as the moles of OH-). When the indicator
changes to pink and the color persists for about 30 seconds, you have completed the titration.
95
Note: Titrate rapidly at first. When the pink color begins to persist, slow the
addition of sodium hydroxide to about one drop per second. As the equivalence
point is approached, add sodium hydroxide so that only one drop of NaOH is
needed to turn the solution pink. If you happen to overshoot the end point, do not
despair, just start over.
9. Record the burette reading. If the volume is less than 10mL, you should read to 3 sig figures;
example, 6.35 mL. If the volume is greater than 10mL you should read to 4 sig figures;
example, 14.31mL. Record the volume directly into your laboratory notebook. If you have
trouble reading the burette to the proper number of significant figures, ask the instructor for
assistance.
10. After the titration is complete, you may pour the contents of the Erlenmeyer flask down the
drain with running water.
11. Repeat the titration as needed until you have a total of at least three titrations. The excess
sodium hydroxide solution and acid should be drained into the sink with running water.
Data:
Stated Molarity of base___________ (this is the theoretical molarity of the base)
Molarity of hydrochloric acid _____________
Trial one
Volume of HCl _____
Final burette vol __________
Initial burette vol __________
Vol of NaOH
___________
Observations:
Trial 2
Volume of HCl _____
Final burette vol __________
Initial burette vol __________
Vol of NaOH
___________
Observations:
Trial 3
Volume of HCl _____
Final burette vol __________
Initial burette vol __________
Vol of NaOH
___________
Observations:
DO NOT RECORD DATA IN YOUR
Include color of when
titration is completed
LABORATORY MANUAL
Calculations:
Trial 1
Determine moles of acid
(see equation 1a)
________
Trial 2
Determine moles of acid
(see equation 1a)
________
Trial 3
Determine moles of acid
(see equation 1a)
________
Determine moles of base
(see equation 1b)
________
Determine moles of base
(see equation 1b)
________
Determine moles of base
(see equation 1b)
________
Determine molarity of
Determine molarity of
Determine molarity of
unknown
unknown
unknown
(See equation 1c) _________ (See equation 1c) _________ (See equation 1c) _________
Statistical analysis
A. Molaritytrial1 ________
Molaritytrial2 _______ molaritytrial3 ________
ß (From calculations
above)
ß (see equation 3)
B. Molarity mean ___________
96
Trial 3 δ3=_______ ß(see
C. Deviation trial 1 δ1= _______
equation 4)
Trial 2 δ2=_________
D. Standard deviation
____________ ß (see equation 5)
E. Relative average deviation
____________ ß (see equation 6)
F. Percent error of the base ______________ ( % error calculations can be found in previous lab)
Results:
1. Molarity of sodium hydroxide
Trial 1 ___________
Trial 2 ___________
Trial 3 ___________
2. Deviation
Trial 1 ___________
Trial 2 ___________
Trial 3 ___________
3. Average Deviation
_________________
4. Relative average deviation _________________
Discussion:
•
•
Discuss your relative average deviation. If it is greater than 1% explain possible sources of
error.
Discuss topics presented in the discussion section of “How to write a lab report”
Post-Lab Questions:
1. 10.00mL of a solution of hydrochloric acid requires 12.81 mL of a solution of 0.1365 M
sodium hydroxide to neutralize it. Give the concentration of the hydrochloric acid in
molarity and mass %. Assume that the hydrochloric acid solution has a density of 1g/mL.
2. You have used 25.38 mL of 0.1073 M potassium permanganate to titrate 9.86 mL of
hydrogen peroxide. Given the balanced equation below and the fact that the molecular
weight of hydrogen peroxide is 34.01 g/mole and the density of hydrogen peroxide is
1.000g/mL, calculate the molarity and the % mass of the hydrogen peroxide. Show all
work; use unit analysis, the proper significant figures, etc.
3H2SO4 + 2 KMnO4 + 5H2O2 à 2MnSO4 + K2SO4 + 5O2 + 8H2O
(Hint: you can disregard all information save the mole relationship between the potassium
permanganate and the hydrogen peroxide.)
References:
Collins, V. Kahl, D. Perry, F. (1996) Good stuff from the chemistry laboratory. Warren Wilson College
press. Swannanoa NC
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102 laboratory
manual.Star Publishing Company: Belmont, California
Silberberg, Martin. 1996. Chemistry the molecular nature of matter and change. Mosby: New York, NY
Willbraham, A. Staley, D. Matta, M. 1995 Chemistry. Addison-Wesley: Menlo Park, CA
97
Experiment 14
Comparison of the geometries and shapes of molecules and ions
Purpose:
To make models of various molecules and ions
To study the geometry and shape of these models
Background:
Valence shell electrons
Valence electrons are the electrons in the outermost energy level. “The number of valence electrons
largely determines the chemical properties of an element” (pg. 374 Willbraham). The periodic table is
arranged such that elements with the same number of valence electrons are arranged in vertical columns.
In the group A elements the number of valence electrons equals the group number thus all elements in
group 1A all have 1 valence electron. The elements in group 2A all have 2 valence electrons, etc. As a
note of interest all transition metals in the B group have 2 valence electrons. That is to say all transition
metals have 2 valence electrons.
Electron dot diagram
The Lewis structure of a single atom, often called a “dot diagram,” consists of a chemical symbol for an
element and a dot for each valence electron. When assigning electrons in an electron dot configuration,
use the following conventions:
1. There are four orbitals represented as four sides surrounding the atom.
2. Place one electron into each side (orbital).
3. After all sides have one electron you may, if necessary, place a second electron in an orbital.
1 valence
electron
2 valence
electrons
3 valence
electrons
4 valence
electrons
5 valence
electrons
6 valence
electrons
7 valence
electrons
8 valence
electrons
When there is only one electron on a side (orbital) the electron is called a free radical. If there are
two electrons on a side it is called a lone pair. Thus: when there are two valence electrons there are
two free radicals. When there are seven valence electrons there are three lone pairs and one free
radical.
Lewis structures
When attempting to draw Lewis structures please use the following easy steps:
1. Find the sum of all valence electrons in all atoms of the molecule
2. Determine the central atom and use a pair of electrons to form a bond between each atom
bound to the central atom.
3. Add remaining electrons in pairs starting with the highest electronegativity (right and up on
the periodic table) to satisfy the octet rule (except, 2 e- for H and 6 e- B).
4. Rearrange electrons to from double and triple bonds to satisfy octet rule (except for H and
boron).
There are exceptions to the rule that all elements needed 8 electrons, for instance hydrogen only
needs 2, boron only needs 6 and phosphorus, sulfur and chlorine can have more than 8 electrons,
because of there large atomic radius and the presence of the “d” subshell.
98
Bond polarity in covalent bonds
In each bond between two nonmetals the electrons are being shared. This shared pair of electrons is
called a covalent bond. There are three types of covalent bonds.
•
•
•
Covalent: electrons are shared equally
Polar covalent: electrons are shared but one atom has more contact than the other atom. This
makes one atom slightly positive and the other atom slightly negative.
Ionic in nature: one atom in the bond is in so much more contact with the electron. That bond
takes on an ionic nature, where one atom maintains a nearly full negative charge and the other
atom maintains a nearly full positive charge.
Each bond in a molecule has a shared pair of electrons. However some atoms do not share equally.
Electronegativity is a measure of how equally the atoms in a bond share the electrons. You will
find the electronegativity value for each atom on the periodic table on the back of this lab manual.
It is in italics in the bottom right corner.
When evaluating the polarity of a bond find the difference between the atoms’ electronegativities.
For instance in an O-H bond, hydrogen has an electronegativity of 2.1and oxygen has an
electronegativity of 3.5. The difference is found by subtracting the two values and taking the
absolute value. Thus in an O-H bond the difference is 1.4. After calculating the difference in
electronegativity consult the table below to determine the type of bond that exists between two
atoms.
Covalent
Polar covalent
Ionic in nature
0.0
0.4
2.0
------------------------------------|---------------------------------------------|----------------------------------From 0.0 up to and
Above 0.4 and less than 2.0
From 2.0 and greater
including 0.4
The above information and table only applies to nonmetal-nonmetal bonds. When considering
combinations of metals and nonmetals the bonds are referred to as ionic.
Geometry
The Lewis structure represents the molecule in two dimensions. The actual arrangement of the
atoms in space occurs in three dimensions. The actual arrangement of atoms within the molecule
can be approximated by considering the arrangements which maximizes the distance between
regions of the negative charge in the central atom. These regions of negative charge are attributes
to the outer, or valence, shell electrons.
After determining the number of electron pairs around an atom, Valence Shell Electron Pair
Repulsion Theory (VSEPR) is used to predict the geometry. The VSEPR theory states that the
geometric arrangement of atoms or groups of atoms, around a central atom, is determined solely by
the repulsion between electron pairs present in the valence shell of the central atom. This means
that the things attached to the central atom will try to stay as far away from each other as possible
and yet stay connected to the central atom. They attempt to stay away from each other because each
electron pair is negatively charged and is repelled by the negative charge on the other electron pairs.
Atoms and lone pairs connected to the central atom are considered electron pairs.
There are only five possible geometries:
99
# of electron groups (domains)
Geometry
2
Linear
3
Trigonal planar
4
Tetrahedral
5
Trigonal bipyramidal
6
Octahedral
Using water as an example, there are four lone pairs of electrons arranged around the central
oxygen and these four pairs of electrons repel each other. The resulting arrangement allows the four
pairs to be as far away from each other as possible. This particular arrangement is called
tetrahedral.
In general, the geometry can be predicted from the number of electron pairs around the central
atom. However the possibility of double and triple bonds must be accounted for. Rather than count
the pairs of electrons around the central atom, the number of groups of electrons (domains) around
the central atom are counted. Thus a double bond with two pairs of electrons contributing to the
bond counts as just one group of electrons. The three pairs of electrons in a triple bond count as
only one group.
Note:
Each lone pair attached to the central atom counts as one group (or domain).
Each atom attached to the central atom counts as one group (or domain).
100
VSEPR Class
Once the Lewis structure has been determined, it is easy to define the VSEPR class. Use an A to
represent the central atom. Represent each atom surrounding the central atom as an X. Then
represent each lone pair of electrons surrounding the central atom as an E. Thus a molecule which
has a central atom with four atoms around it and no lone pairs would be AX4. A molecule with a
central atom that is connected to two atoms and has two lone pairs would be AX2E2.
Geometry
Linear
AXE type
Trigonal planar
AXE type
Tetrahedral
AXE type
0 lone pairs
1 lone pair
2 lone pairs
3 lone pairs
AX2
AXE
AX3
AX2E
AXE2
AX4
AX3E
AX2E2
AXE3
AX5
AX4E
AX3E2
AX2E3
AX6
AX5E
AX4E2
Trigonal Bipyramidal
AXE type
Octahedral
AXE type
Example: H2O has a tetrahedral geometry but the shape is described as bent
Lewis structure
Geometry Tetrahedral
or
Shape
After the geometric arrangement of the electron groups has been determined, it is possible to
describe the shape. The shape describes the three dimensional arrangement of the atoms in a
molecule. If the molecule has no lone pairs, the shape is the same as the geometry. For molecules
with lone pairs, the unshared pairs are used to determine the geometry but are ignored when
describing the shape. The following table illustrates the different geometries and subsequently
possible shapes.
101
Shape
Geometry
No lone pairs
1 lone pair
2 lone pair
3 lone pair
2 electron groups
(domains)
Linear
Linear
Linear
Trigonal Planar
Bent
Linear
Tetrahedral
Trigonal Pyramidal
Bent
Linear
Trigonal Bipyramidal
Seesaw
T-shaped
Linear
Octahedral
Square pyramidal
Square planar
3 electron groups
(domains)
Trigonal
Planar
4 electron groups
(domains)
Tetrahedral
5 electron groups
(domains)
Trigonal
Bipyramidal
6 electron groups
(domains)
Octahedral
Example: H2O has a tetrahedral geometry but because it has two lone pairs the shape is described
as bent. It’s AXE type is AX2E2
Lewis structure
Geometry
Shape
or
Tetrahedral
Consider all electron
domains
102
Bent
Considers the orientation of
only the bound atoms
Molecular Polarity
Once the shape of a molecule has been determined it is now possible to determine if the molecule is
polar. A molecule is polar if there is an unequal distribution of charge. Thus if one side of a
molecule is negative and anther side is positive, the molecule is polar. The distribution of charge in
a molecule is due to the polar bonds within the molecule. But be very careful it is possible to have a
molecule that has polar bonds but the bonds are equal arranged due to the shape of the molecule and
as a result the molecule is not polar.
Experimentally a molecule is polar if the molecular orients itself in space when placed in a charged
field.
If a molecule is polar, the positive and negative areas of the molecule should be shown. There are
two methods of showing this described below.
Example:
Water is a polar molecule
The oxygen side of the molecule has a slight negative charge because when it shares the electrons
with hydrogen in the bond it keeps them in the oxygen orbitals for longer periods of time. This
makes the oxygen side have a negative charge. Because the electrons are pulled more frequently to
the oxygen side it leaves the hydrogen side with a low electron density which makes the hydrogen
side slightly positive.
Bond angle
The bond angles are an important part of the description of the geometry and shape of a molecule in
three dimensions. It is possible to predict the approximate bond angle from the geometric
arrangement of the atoms in a molecule or ion.
Geometry
Linear
Trigonal Planar
Tetrahedral
Trigonal bipyramidal
Octahedral
Approximate bond angle
180o
120o
109.5o
90o, 120o, 180o
90 o
103
Hybridization of the Central Atom
In atoms, the areas where the electrons are likely to be found are called orbitals. In molecules, these
atomic orbitals become hybridized so that the bonding electrons can be located between the atoms.
The hybridization about the central atom can be predicted if the number of electron domains is
counted. The number of electron domains is always the same as the number of orbitals that make
up the hybridization.
Geometry
# of electron groups (domains)
Hybridization of central atom
Linear
2
sp
Trigonal Planar
3
sp2
Tetrahedral
4
sp3
Trigonal bipyramidal
5
sp3d
Octahedral
6
sp3d2
In a “sp” hybridized orbital one orbit from the “s” subshell and 1 orbit from the p subshell are
combined to form the sp hybridized orbital. There are a total of 2 orbits in a sp hybrid. Please note
that there are 3 p orbitals. In the sp hybrid 2 p orbitals have not been used
In a “sp2” hybridized orbital one orbit from the “s” subshell and 2 orbits from the p subshell are
combined to form the sp2 hybridized orbital. There are a total of 3 orbitals in a sp2 hybrid. Please
note that there are 3 p orbitals. In the sp2 hybrid 1 p orbitals have not been used
In a “sp3” hybridized orbital one orbit from the “s” subshell and 3 orbits from the p subshell are
combined to form the sp3 hybridized orbital. There are a total of 4 orbitals in a sp3 hybrid.
In a “sp3d” hybridized orbital one orbit from the “s” subshell, 3 orbits from the p subshell and one
orbit from the d subshell are combined to form the sp3d hybridized orbital. There are a total of 5
orbitals in a sp3d hybrid.
In a “sp3d2” hybridized orbital one orbit from the “s” subshell, 3 orbits from the p subshell and two
orbit from the d subshell are combined to form the sp3d2 hybridized orbital. There are a total of 6
orbitals in a sp3d2 hybrid.
Pre-Lab Questions:
1. When drawing Lewis structures it is important to know the octet rule.
a. State the octet rule.
b. Which atoms are allowed to take less than eight electrons when drawing Lewis structures?
c. Which atoms are allowed to obtain more than eight electrons when drawing Lewis
structures?
104
Procedure:
Atoms
Look at your atom kit and determine what model atom will be used for each of the following
Boron
Hydrogen
Phosphorus
Fluorine
Nitrogen
Sulfur
Carbon
Oxygen
Chlorine
Model kit electron pairs
Parts
Short straight
Paddle
Long Curved
Electron pairs
2 electrons in a single bond
2 electrons as a pair, but not involved in a bond (lone pair)
2 electrons as part of a double or triple bond
For each of the molecules below, draw the Lewis structure in your notebook, build the model, then
answer questions A-J. Finally, do all the comparisons.
Model 1
BF3
(boron trifluoride)
Model 7
SF6
(sulfur hexafluoride)
Model 2
CH4
(methane)
Model 8
CO2
(carbon dioxide)
Model 3
NH3
(ammonia)
Model 9
CO32- (carbonate ion)
Model 4
H2 O
(water)
Model 10
CH3Cl (chloromethane)
Model 5
NH4+
(ammonium)
Model 11
BeCl2 (Beryllium chloride)
Model 6
PCl5
(phosphorus pentachloride) Model 12
ICl5
(Iodine pentachloride)
Questions
Identify the central atom and answer each of the following questions for each model. You must
completely copy the question into your laboratory notebook before working on the first model. You
may find it easiest to make a large table with headings to do this exercise.
A. How many electron pairs (domains) are present
F. What is the shape of the molecule?
around the central atom?
G. Are there any polar bonds?
B. How many of the e pairs are atoms bonded to the
H. Is the molecule polar or nonpolar?
central atom (count double bonds as one bond)?
C. How many electron pairs connected to the central
I.
List all bond angles 180o or less.
atom are lone pairs?
D. What is the VSEPR class (AXmEn)?
J.
Hybridization around the central atom?
E. What is the geometry of the molecule?
K. Make a 3-D sketch of the model.
When making the 3-D sketch please use the following conventions as an example. Let us consider
the chloroform molecule CH2Cl2:
A standard 3-D
sketch
These two lines and
attached H atoms are on
the same plane as the
central atom
This dashed wedge
indicates a bond and Cl
atom that is going away
from you.
105
This wedge indicates a
bond and Cl atom that is
coming toward you
In general keep double bonds on the same plane as the central atom.
Data Table:
Please construct your data table to look like this, (landscape oriented). It will
make organizing and comparing your data easier and will simplify grading. If it is not like this
example it may not be accepted.
Model
Example:
H3 O+
Model #1
BF3
Model
Model #2
CH4
Model #3
NH3
Lewis
structure
Lewis
structure
A.
B.
1
C.
C.
3
B.
4
A.
D.
E.
Tetrahedral
E.
F.
Trigonal
Pyramidal
F.
H+àO
or
δ+H – Oδ-
G.
G.
D.
AX3E
H.
Polar
H.
I.
I.
J.
Sp3
J.
K.
K.
106
H-O
109.5
Results:
Do all of the following comparisons. Be sure to consider your answers to the questions on the
previous page. Always use the formula or name of the compound, not the model number when
making comparisons. Write in complete sentences to answer each question.
1. Compare models for CH4, NH3 and H2O. Are there similarities in the number of electron
pairs, the geometry, the shape, hybridization, etc? Differences? Are any of the molecules
polar?
2. Compare CH4 and NH4+. Are there similarities in the number of electron pairs, the
geometry, the shape, hybridization, etc? Differences? Do the sketches look similar? Is
either molecule polar?
3. Compare the models for NH3 and NH4+. Are there similarities in the number of electron
pairs, the geometry, the shape, hybridization, etc? Differences? DO the sketches look
similar? Is either molecule polar?
4. Compare models for PCl5 and SF6. What do they have in common? Is either molecule
polar?
5. Compare CH4 and CO2. Are there similarities in the number of electron pairs, the geometry,
the shape, hybridization, etc? Differences? What do they have in common? Is either
molecule polar?
6. Compare H2O and CO2. Are there similarities in the number of electron pairs, the geometry,
the shape, hybridization, etc? Differences? Is either molecule polar?
7. Compare models for BF3 and CO32-. Are there similarities in the number of electron pairs,
the geometry, the shape, hybridization, etc? Differences? Is either molecule polar?
8. Compare the models for NH3 and CO32-. Are there similarities in the number of electron
pairs, the geometry, the shape, hybridization, etc? Differences? Is either molecule polar?
9. Compare CH4 and CH3Cl. Are there similarities in the number of electron pairs, the
geometry, the shape, hybridization, etc? Differences? Is either molecule polar?
Discussion:
What information is shown by a model but is not shown in a Lewis structure? Does a Lewis
structure show whether a molecule is polar? Explain your answer using example from the
models.
Post-Lab Questions: create a data table like the one used previously in this laboratory activity to
record the following answers.
1.
2.
3.
4.
Predict the answers to questions A-K for a model of CCl4.
Predict the answers to questions A-K for a model of CHCl3
Predict the answers to questions A-K for a model of CH2Cl2
Compare CH4, CH3Cl, CH2Cl2, CHCl3, CCl4
References:
Finnegan, M. Place, H. Weissbart, B. (2000) Washington state university chemistry 101-102
laboratory manual. Star Publishing Company: Belmont, California
Willbraham, A. Staley, D. Matta, M. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park, CA
Silberberg, Martin. (1996) Chemistry the molecular nature of matter and change. Mosby: New
York, NY
107
Experiment 15
Vapor Pressure of Liquids
In this experiment, you will investigate the relationship between the vapor pressure of a liquid and
its temperature. When a liquid is added to the Erlenmeyer flask shown in Figure 1, it will evaporate
into the air above it in the flask. Eventually, equilibrium is reached between the rate of evaporation
and the rate of condensation. At this point, the vapor pressure of the liquid is equal to the partial
pressure of its vapor in the flask. Pressure and temperature data will be collected using a Gas
Pressure Sensor and a Temperature Probe. The flask will be placed in water baths of different
temperatures to determine the effect of temperature on vapor pressure. You will also compare the
vapor pressure of two different liquids, ethanol and methanol, at the same temperature.
Objectives:
In this experiment, you will
•
•
Investigate the relationship between the vapor pressure of a liquid and its temperature.
Compare the vapor pressure of two different liquids at the same temperature.
Figure 1
Materials:
computer
Vernier computer interface
Logger Pro
Vernier Gas Pressure Sensor
Vernier Temperature Probe
rubber-stopper assembly
plastic tubing with two connectors
20 mL syringe
two 125 mL Erlenmeyer flasks
methanol
ethanol
ice
four 1 liter beakers
108
Procedure:
1. Obtain and wear goggles! CAUTION: The alcohols used in this experiment are flammable and
poisonous. Avoid inhaling their vapors. Avoid contacting them with your skin or clothing. Be
sure there are no open flames in the lab during this experiment. Notify your teacher immediately
if an accident occurs.
2. Use 1 liter beakers to prepare four water baths, one in each of the following temperature ranges:
0 to 5°C, 10 to 15°C, 20 to 25°C (use room temperature water), and 30 to 35°C. For each water
bath, mix varying amounts of warm water, cool water, and ice to obtain a volume of 800 mL in
a 1 L beaker. To save time and beakers, several lab groups can use the same set of water baths.
3. Prepare the Temperature Probe and Gas Pressure Sensor for data collection.
a. Plug the Gas Pressure Sensor into CH1 and the Temperature Probe into CH2 of the computer
interface.
b. Obtain a rubber-stopper assembly with a piece of heavy-wall plastic tubing connected to one
of its two valves. Attach the connector at the free end of the plastic tubing to the open stem of
the Gas Pressure Sensor with a clockwise turn. Leave its two-way
valve on the rubber stopper open (lined up with the valve stem as
shown in Figure 2) until Step 9.
c. Insert the rubber-stopper assembly into a 125 mL Erlenmeyer flask.
Important: Twist the stopper into the neck of the flask to ensure a
tight fit.
Figure 2
Figure 3
4. Follow your instructions on how to calibrate the sensors.
5. The temperature and pressure readings should now be displayed in the meter. While the twoway valve above the rubber stopper is still open, record the value for atmospheric pressure in
your data table (round to the nearest 0.1 kPa).
6. Finish setting up the apparatus shown in Figure 3:
a. Obtain a room-temperature water bath (20–25°C).
b. Place the Temperature Probe in the water bath.
c. Hold the flask in the water bath, with the entire flask covered as
shown in Figure 3.
d. After 30 seconds, close the 2-way valve above the rubber stopper as
shown in Figure 4—do this by turning the white valve handle so it is
perpendicular with the valve stem itself.
open
closed
Figure 4
7. Obtain the methanol container and the syringe. Draw 3 mL of the methanol up into the syringe.
With the two-way valve still closed, screw the syringe onto the two-way valve, as shown in
Figure 3.
109
8. Introduce the methanol into the Erlenmeyer flask.
a. Open the 2-way valve above the rubber stopper—do this by turning the white valve handle so
it is aligned with the valve stem (see Figure 4).
b. Squirt the methanol into the flask by pushing in the plunger of the syringe.
c. Quickly return the plunger of the syringe back to the 3 mL mark of the syringe, then close the
2-way valve by turning the white valve handle so it is perpendicular with the valve stem.
d. Remove the syringe from the 2-way valve with a counter-clockwise turn.
9. To monitor and collect pressure and temperature data:
a. Click
.
b. When the pressure and temperature readings displayed in the meter stabilize, equilibrium
between methanol liquid and vapor has been established. Click
. The first pressuretemperature data pair is now stored.
10. To collect another data pair using the 30–35°C water bath:
a. Place the Erlenmeyer flask assembly and the temperature probe into the 30–35°C water bath.
Make sure the entire flask is covered.
b. When the pressure and temperature readings displayed on the computer monitor stabilize,
click
. The second data pair has now been stored.
11. For Trial 3, repeat the Step-10 procedure, using the 10–15°C water bath. Then repeat the
Step-10 procedure for Trial 4, using the 0–5°C water bath.
12. Click
to end data collection. Record the pressure and temperature values in your data
table, or, if directed by your instructor, print a copy of the table.
13. Gently loosen and remove the Gas Pressure Sensor so the Erlenmeyer flask is open to the
atmosphere. Remove the stopper assembly from the flask and dispose of the methanol as
directed by your teacher.
14. Obtain another clean, dry 125 mL Erlenmeyer flask. Draw air in and out of the syringe enough
times that you are certain that all of the methanol has evaporated from it.
16. Repeat Steps 6–8 to do one trial only using ethanol in the room temperature water bath. When
the pressure stabilizes, record the measured pressure of ethanol displayed in the meter in your
data table.
17. Gently loosen and remove the stopper assembly from the flask and dispose of the ethanol as
directed by your teacher.
110
Processing the Data:
1. Convert each of the Celsius temperatures to Kelvin (K). Write the answer in the space provided.
2. To obtain the vapor pressure of methanol and ethanol, the air pressure must be subtracted from
each of the measured pressure values. However, for Trials 2–4, even if no methanol was present,
the pressure in the flask would have increased due to a higher temperature, or decreased due to a
lower temperature (remember those gas laws?). Therefore, you must convert the atmospheric
pressure at the temperature of the first water bath to a corrected air pressure at the temperature
of the water bath in Trial 2, 3, or 4. To do this, use the gas-law equation (use the Kelvin
temperatures):
P2 P1
=
T2 T1
where P1 and T1 are the atmospheric pressure and the temperature of the Trial 1 (room
temperature) water bath. T2 is the temperature of the water bath in Trial 2, 3, or 4. Solve for P2,
and record this value as the corrected air pressure for Trials 2, 3, and 4. For Trial 1 of methanol
and Trial 1 of ethanol, it is not necessary to make a correction; for these two trials, simply
record the atmospheric pressure value in the blank designated for air pressure.
3. Obtain the vapor pressure by subtracting the corrected air pressure from the measured pressure
in Trials 2-4. Subtract the uncorrected air pressure in Trial 1 of methanol (and Trial 1 of
ethanol) from the measured pressure.
4. Plot a graph of vapor pressure vs. temperature (°C) for the four data pairs you collected for
methanol. Temperature is the independent variable and vapor pressure is the dependent variable.
As directed by your teacher, plot the graph manually, or use Logger Pro. Note: Be sure to plot
the vapor pressure, not the measured pressure.
5. How would you describe the relationship between vapor pressure and temperature, as
represented in the graph you made in the previous step? Explain this relationship using the
concept of kinetic energy of molecules.
6. Which liquid, methanol or ethanol, had the larger vapor pressure value at room temperature?
Explain your answer. Take into account various intermolecular forces in these two liquids.
111
Data and Calculations:
Atmospheric pressure
_______ kPa
Substance
Trial
Methanol
1
2
Ethanol
3
4
1
Temperature
(°C)
°C
°C
°C
°C
°C
Temperature
(K)
K
K
K
K
K
kPa
kPa
kPa
kPa
kPa
Measured
pressure
Air pressure
no correction
corrected
corrected
corrected
no correction
kPa
kPa
kPa
kPa
kPa
kPa
kPa
kPa
kPa
kPa
Vapor pressure
Results:
Draw a graph that reports your data.
Discussion:
What does the graph and the differences in vapor pressure tell you about the strength of
intermolecular forces of the different chemicals? What might account for these differences? How
does temperature effect vapor pressure and why? Propose a possible follow-up experiment.
Post Lab Questions:
5. What kind are the molecular weights for each chemical tested in this lab.
6. Draw Lewis structures for each chemical tested in this lab.
a. Indicate all polar bonds in each Lewis structure.
b. Indicate molecular polarity
7. What strongest type of intermolecular forces are present in each of the chemicals tested in
this lab
8. Provide an explanation based on molecular polarity and molecule weight to account for the
differences in vapor pressure.
112
Experiment 16
Calorimetry: Heat of Fusion of Ice
Purpose:
To determine the energy (in Joules) required to melt one gram of ice and to determine the molar
heat of fusion for ice (in kJ/mol).
Background:
When a chemical or physical change takes place heat is either given off or absorbed. That is, the
change is either exothermic or endothermic. It is important for chemists to be able to measure this
heat. Measurements of this kind are made in a device called a calorimeter (kal rim’əә təәr). The
technique used in making these measurements is called calorimetry.
In simplest terms, a calorimeter is an insulated container made up of two chambers. The outer
chamber contains a known mass of water. In the inner chamber, the experimenter places the
materials that are to lose or gain heat while undergoing a physical or chemical change. The basic
principle on which the calorimeter works is that when two bodies at different temperature are in
contact with one another, heat will flow from the warmer body to the colder body. Thus, heat lost
by one body will be gained by the other. This exchange of heat will continue until the two bodies
are at the same temperature. In a calorimeter heat is exchanged between the water and the materials
undergoing change.
Unlike most calorimeters, the simple Styrofoam-cup calorimeter used in this experiment will have
only one chamber. The ice will be placed directly into a measured amount of water. The heat
required to melt the ice will be supplied by the water. By measuring the temperature change (ΔT)
of the water, you can calculate the quantity of heat exchanged between the water and the ice. Using
these experimental data, you will calculate the heat of fusion of ice.
Melting and freezing behavior are among the characteristic properties that give a pure substance its
unique identity. As energy is added, pure solid water (ice) at 0°C changes to liquid water at 0°C.
In this experiment, you will determine the energy (in joules) required to melt one gram of ice. You
will then determine the molar heat of fusion for ice (in kJ/mol). Excess ice will be added to warm
water, at a known temperature, in a Styrofoam cup. The warm water will be cooled down to a
temperature near 0°C by the ice. The energy required to melt the ice is removed from the warm
water as it cools.
To calculate the heat that flows from the water, you can use the relationship
q = m• Cp •Δt
where q stands for heat flow, Cp is specific heat, m is mass in grams, and Δt is the change in
temperature. For water, Cp is 4.18 J/g°C.
Materials:
100 mL graduated cylinder
Vernier computer interface
warm water
Temperature Probe
Styrofoam cup
ring stand
Logger Pro
tongs
113
stirring rod
ice cubes
utility clamp
250 mL beaker
Procedure:
1.
2.
3.
4.
Connect the probe to the computer interface.
Place a Styrofoam cup into a 250 mL beaker as shown in Figure 1.
Use a utility clamp to suspend the Temperature Probe on a ring stand as shown in Figure 1.
Use a 100 mL graduated cylinder to obtain 100.0 mL of water at about 60°C from your
instructor. Record this as V1.
5. Obtain 7 or 8 large ice cubes.
6. Lower the Temperature Probe into the warm water (to about 1 cm from the bottom).
7. Click
to begin data collection. Wait until the temperature reaches a maximum (it will
take a few seconds for the cold probe to reach the temperature of the warm water). This
maximum will determine the initial temperature, t1, of the water. As soon as this maximum
temperature is reached, fill the Styrofoam cup with ice cubes. Shake excess water from the ice
cubes before adding them (or dry with a paper towel). Record the maximum temperature, t1, in
your data table.
Figure 1
8. Use a stirring rod to stir the mixture as the temperature approaches 0°C. Important: As the ice
melts, add more large ice cubes to keep the mixture full of ice!
9. When the temperature reaches about 4°C, quickly remove the unmelted ice (using tongs).
Continue stirring until the temperature reaches a minimum (and begins to rise). This minimum
temperature is the final temperature, t2, of the water. Record t2 in your data table. Click
when you have finished collecting data.
10. Use the 100 mL graduated cylinder to measure the volume of water remaining in the Styrofoam
cup to the nearest 0.1 mL. Record this as V2.
11. You can confirm your data by clicking the Statistics button, . The minimum temperature (t2)
and maximum temperature (t1) are listed in the floating box on the graph.
Calculations:
1.
2.
3.
4.
Use the equation Δt = t2 – t1 to determine Δt, the change in water temperature.
Subtract to determine the volume of ice that was melted (V2 –V1).
Find the mass of ice melted using the volume of melt (use 1.00 g/mL as the density of water).
Use the equation given in the introduction of this experiment to calculate the energy (in joules)
released by the 100 g of liquid water as it cooled through Δt.
5. Now use the results obtained above to determine the heat of fusion—the energy required to melt
one gram of ice (in J/g H2O).
6. Use your answer to Step 5 and the molar mass of water to calculate the molar heat of fusion for
ice (in kJ/mol H2O).
7. Find the percent error for the molar heat of fusion value in Step 6. The accepted value for molar
heat of fusion is 6.01 kJ/mol.
114
Data and Calculations:
Initial water temperature, t1
°C
Final water temperature, t2
°C
Change in water temperature, Dt
°C
Final water volume, V2
mL
Initial water volume, V1
mL
Volume of melt
mL
Results:
Mass of ice melted
Heat released by cooling water (q = Cp•m•Δt)
g
J/g ice melted (heat of fusion)
J
kJ/mol ice melted (molar heat of fusion)
J/g
Percent error
kJ/mol
Discussion:
Please discuss issues suggested in “how to write a lab report.”
%
Post-Lab Questions:
1. List possible sources of error. Describe the design of a calorimetry experiment that would
reduce some of the error potential.
2. One source of error is the flow of heat between the water in the cup and the surroundings.
Explain how this error is reduced by starting with water at 50 oC.
3. In what way does the calorimetry make use of the law of conservation of energy?
4. Define the following terms: a. exothermic, b. endothermic, c. heat of fusion, d. specific heat
capacity.
5. Is the process of melting exothermic or endothermic? Give evidence to support this answer.
6. What is the difference between heat and temperature?
115
Experiment 17
Boyle’s Law: Pressure-Volume Relationship in Gases
The primary objective of this experiment is to determine the relationship between the pressure and
volume of a confined gas. The gas we use will be air, and it will be confined in a syringe connected
to a Gas Pressure Sensor (see Figure 1). When the volume of the syringe is changed by moving the
piston, a change occurs in the pressure exerted by the confined gas. This pressure change will be
monitored using a Gas Pressure Sensor. It is assumed that temperature will be constant throughout
the experiment. Pressure and volume data pairs will be collected during this experiment and then
analyzed. From the data and graph, you should be able to determine what kind of mathematical
relationship exists between the pressure and volume of the confined gas. Historically, this
relationship was first established by Robert Boyle in 1662 and has since been known as Boyle’s
Law.
Please note that cm3 is often written “cc” in the medical field. Also, keep in mind that cm3 is the
same as mL. You may find it also useful to know that 101.325 kPa equals 1 atm (this relationship
has infinite sig figs).
Objectives:
In this experiment, you will
Use a Gas Pressure Sensor and a gas syringe to measure the pressure of an air sample at
several different volumes.
• Determine the relationship between pressure and volume of the gas.
• Describe the relationship between gas pressure and volume in a mathematical equation.
• Use the results to predict the pressure at other volumes.
•
Figure 1
Materials:
Lab Quest
1.
2.
3.
4.
5.
Vernier Gas Pressure Sensor
20 mL gas syringe
Obtain Lab Quest, Gas Pressure Sensor and a 20mL syringe.
Connect Gas Pressure Sensor to the Lab Quest into port 1.
Turn Lab Quest on.
Make sure Lab Quest is reading pressure in kPa. If it is not in kPa please change units to kPa.
The pressure in kPa should be between around 90 and 120 kPa. If it is not, please ask your
instructor for help. Please do not forget to record both your observations and the quantitative
data.
116
6. Obtain a syringe. DO NOT CONNECT YOUR SYRINGE to the gas pressure sensor. Please
notice the front of the black rubber portion of the plunger is the side that measurements are taken
from (see figure 2)
(figure 2)
This is a measurement of 5.0 mL
7. With your syringe NOT CONNECTED to the gas pressure sensor move the plunger to the 20.0
mL mark.
8. With the syringe at 20.0 mL now CONNECT the syringe to the gas pressure sensor.
9. Record the gas pressure in kPa. Now record the volume. Please add 0.8ml to the volume to
account for the volume inside the gas pressure sensor.
10. Do not disconnect the syringe from the gas pressure sensor. Move the plunger to 15.0 mL and
hold it long enough to record the pressure in kPa. One person should move the plunger and one
person should record the gas pressure. Record the volume. Remember to add 0.8mL to all
volume data readings. Move the plunger back to 20.0 mL. Do not remove syringe from Gas
Pressure Sensor.
11. Repeat step 10. However in this step move plunger to 12.0 mL.
12. Repeat step 10. However in this step move plunger to 10.0 mL.
13. Repeat step 10. However in this step move plunger to 8.0 mL.
14. Graph your data on graph paper. Place the volume on the X axis and the pressure in atm on the
Y axis. Please use all graphing conventions.
Data:
Volume
(mL) add 0.8mL
Pressure
(kPa)
Calculations:
Convert all kPa values into atm, using dimensional analysis.
Calculate P/V for all samples. This may produce a constant value for K
Calculate P*V for all samples. This may produce a constant value for K
117
Report values in a table such as
Pressure in Atm
P/V
P*V
The correct mathematical relationship between gas volume and pressure will
produce a constant value. This value will be labeled as k.
Results:
Provide a graph of the data. Draw a best fit smoothed line. Do not connect the dots with straight
lines.
Report the Value found for k.
Discussion:
Cover topics described in how to write a lab report. Propose an experiment that would demonstrate
Charles’s Law.
Post lab questions:
1. Using the same the best fit line on your graph. What does your data show would happen to the
pressure if the volume is changed from 20.0 mL to 13.5 mL? Show the pressure values in your
answer.
2. From your data and the shape of the curve in the plot of pressure vs. volume, do you think the
relationship between the pressure and volume of a confined gas is direct or inverse? Explain
your answer.
3. Based on your data, what would you expect the pressure to be if the volume of the syringe was
increased to 40.0 mL? Explain or show work to support your answer.
4. Based on your data, what would you expect the pressure to be if the volume of the syringe was
decreased to 2.5 mL? Explain or show work to support your answer.
5. What experimental factors are assumed to be constant in this experiment?
6. How constant were the values for k you obtained in Question 8? Good data may show some
minor variation, but the values for k should be relatively constant.
7. Using P, V, and k, write an equation representing Boyle’s law. Write a verbal statement that
correctly expresses Boyle’s law.
118
Experiment 18
Quantitative Reaction of HCl and Mg
Purpose:
To calculate the molar volume of a gas collected over water
Hypothesis: Develop a hypothesis stating how many liters of gas will be produced for each
mole of gas produced.
Background:
Pertinent Gas Laws
To understand the gas laws at an intuitive level it is valuable to understand the kinetic molecular
model of gas behavior. The kinetic molecular model it makes three basic assumptions:
1. Gas is composed of particles and there is no attractive force between gas particles.
2. The particles of gas move rapidly in straight lines and are in constant random motion.
3. All collisions are perfectly elastic, meaning that kinetic energy is transferred from one
particle to another in a collision but no energy is lost.
Temperature in this model is a measure of how much kinetic energy the speeding particles have.
The equation for kinetic energy is Ekinetic = ½ mass * Velocity2. From the equation you can see that
as temperature goes up particles move faster. While working with gas laws, temperature is always
measured in Kelvin. To convert from Celsius to Kelvin add 273.14 to convert from Kelvin to
Celsius subtract 273.14.
Pressure in this model is how much force is pressed against the container when the speeding gas
particles collide with the surfaces of the container. Pressure is measured in many different units.
The three units commonly used in this class are millimeters mercury (mmHg), and kilopascal (kPa),
atmospheres (atm)
760 mmHg = 101.3kPa = 1atm
Boyle’s Law
Consider the effect of pressure on the volume of a contained gas while temperature is held constant.
As volume decreases the pressure in the container goes up. As volume increases pressure goes
down.
Boyle’s Law states that for a given mass of gases at constant temperature, the volume of gas varies
inversely with the pressure. Mathematically it is expressed as:
P1V1=P2V2
In this equation P1and V1 represent the temperature and pressure at the initial set of conditions and
P2 and V2 are the final set of conditions. Thus if under one set of conditions the pressure is 100 atm
and the volume is one liter the same mass of gas at the same pressure would have a pressure of 50 if
it was contained in a 2 liter container.
119
P1V1=P2V2
100 atm* 1L = 50 atm *2 L
By applying the kinetic model to the behavior of gas the pressure increases in a smaller container (if
temperature is held constant) because as the volume decreases the surface area decreases a lot.
With a smaller surface area each small area of surface is getting hit more often. Imagine 5 pool
balls constantly moving on a pool table. If the pool table is huge, say the size of a soccer field, the
balls infrequently hit the sides. If you keep the balls moving at the same speed but shrink the table
down to the size of your desk it is easy to see that the balls now hit the sides of the table more
frequently. Remember pressure is a measure of how often and how hard the particles hit the
surfaces of the container.
Ideal Gas Law
The Ideal Gas Law allows you to consider the effect of changing the number of molecules present
in a sample. There are different ways to increase pressure, for instance you could speed the
molecules up so they hit the surface of the container harder (increase temperature), or you can
decrease the surface area of the container so that the particles hit more often (decrease volume) or
you can do one other thing to increase pressure. You can put more particles in the container so that
the surfaces of the container are hit more often because there is more stuff inside to hit the walls.
Using the pool table analogy to change pressure you can change the size of the table. Rolls the balls
faster or just put more balls on the table.
The Ideal Gas Law states that in any gas system if you multiply the pressure times the volume and
divide that by the product of the temperature and the moles it will equal a constant value. This is
true for any ideal gas.
Thus
P*V
Where
P = pressure (atm)
V= Volume (L)
=R
T*n
T= temperature (K)
n= moles of gas
Thus
P1 * V1 = R
T1 * n1
R= 0.08206
or expressed in common form:
PV=nRT
atm*L
mol*K
This is an amazingly powerful equation. It is the superman of gas law equations. It can pretty
much do anything.
Dalton’s Partial Pressure Law
Consider the pool table. Pressure is measured by how often and how hard the balls hit the side of
the container. It is easy to picture the balls rolling around in straight lines and colliding with the
bumpers and bouncing back to keep on moving. Dalton’s Partial Pressure Law states that each gas
in a container contributes to the total pressure. So if on the pool table there are blue balls and some
red balls and some green balls the total pressure exerted on the sides of the pool table is equal to the
sum of the pressure exerted from each of the different balls.
Thus Ptotal =P1 +P2 +P3 etc
The contribution each gas in a mixture makes to the total pressure is the partial pressure exerted by
that gas. Dalton’s Law of Partial Pressures: at constant volume and temperature, the total
120
pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component
gases.
So if on the pool table you want to know how much pressure is exerted you need to know how
much pressure is being exerted by all of the different gases in the container.
So in your class room, let us assume that the pressure is 1.00 atm. Let us also assume that the air in
the classroom is 78% nitrogen, 21% oxygen and 1% carbon dioxide. In this situation the total
pressure in the room is
0.78 atm nitrogen + 0.21 atm oxygen + 0.01 atm carbon dioxide = 1 atm air
Calculations
The most challenging part of this lab will be to complete the calculations.
A. To find grams of Magnesium ribbon
1. Convert cm of Mg to grams of Mg
Let us assume that 1 meter of Mg ribbon has a mass of 0.750 grams. If you have 5.03 cm of
magnesium ribbon you can calculate how many grams of Mg as follows:
5.03 cm Mg
1m
100 cm
0.750 grams Mg
1 meter Mg
Sample equation 1
= .0377 g Mg
2. Convert grams of Mg to Moles Mg
Use dimensional analysis and the atomic weight of Mg found on the periodic table.
0.0377 g Mg
1 mol Mg
Sample equation 2
24.30 g Mg
= 0.00155 mol Mg
(Result A)
B. To find theoretical moles of hydrogen gas produced
0.00155 mol Mg
1 mol H2
1 mol Mg
= 0.00155 mol H2
Sample equation 3
(Result B)
C. Find experimental ideal gas constant.
The ideal gas law is PV=nRT. To solve for R rearrange the equation into this form
R=PV/nT
1. To solve for n, convert mole of Mg to moles of hydrogen gas
2HCl(aq) + Mg(s) à H2(g) + MgCl2(aq)
1. Find T
T = the temperature in Kelvin to convert Celsius to Kelvin add 273.14 to the Celsius
measurement
273.14 + 19.0 oC = 292.2 K
3. Find Volume
Read V off of your gas collection tube to 4 sig figs. In this example the value is 37.85mL
The units of this value must be liters so convert mL to liters.
4. Determine pressure of hydrogen gas
121
To do this you will need to find the total pressure for today and then subtract the pressure due
to the water vapor in the gas collection tube. To do this you will use Dalton’s partial pressure
law. Dalton’s partial pressure law states
Vapor Pressure of water at various
that P1 + P2 + Pn = Ptotal. This just means
temperatures
that the partial pressure of each gas in a
Temperature
Pressure
Temp.
Pressure
system contributes to the total pressure in
(K)
(atm)
(K)
(atm)
the container.
288.2
0.0168
296.2
0.0276
In this case you have
289.2
0.0179
297.2
0.0295
Phydrogen + Pwater vapor = Ptotal
290.2
0.0191
298.2
0.0312
291.2
0.0204
299.2
0.0332
The total pressure in this experiment will
300.2
0.0351
292.2
0.0217
equal the room pressure. The water
301.2
0.0372
293.2
0.0230
vapor pressure will be determined based
302.2
0.0395
294.2
0.0245
on room temperature and a table. These
303.2
0.0418
295.2
0.0261
two pieces of information will allow you
to find the partial pressure of the
hydrogen gas.
In this example assume the pressure today is 0.996 atm and the temperature is 19.0 oC which
then is converted to 292.2 K (from Celsius to Kelvin add 273.14).
Thus:
Ptotal - Pwater vapor
= 0.996 atm - .0217 atm
= 0.974 atm Phydrogen
Sample equation 4
5. To solve for R
Using the information from above
R=PV/nT
R= 0.08314atm*L/mol*K
P=0.974 atm
T= 292.2 K
V= 0.03785 liters
n= 0.00155 moles
D. Percent Error
To find the percent error you just use the percent error formula.
Percent
Error
= | Theoretical – Experimental | *100
Theoretical
| 0.08206mol*K/atm*L- 0.08314mol*K/atm*L |
0.08206 L/mol
* 100
122
Sample equation 5
=1.32% error
Sample equation 6
Pre-Lab Questions:
1.
a. Convert 0.879 atm to kPa
b. Convert 787 mmHg to atm
2. The pressure on 2.50 L of gas changes from 0.987 atm to 0.400 atm. What is the new volume if
the temperature is held constant?
3. When a rigid hollow sphere containing 6.80 liters of gas is heated to 601 K the pressure is 180.0
atm. How many moles of helium are in the sphere?
4. Determine the total pressure of a gas mixture that contains oxygen nitrogen and helium if the
partial pressures of the gases are PO2 = 0.20 atm, PN2= 0.467 atm, PHe= 0.267 atm?
Materials:
Magnesium ribbon
Volumetric gas collection tube
Battery jar
Ring stand
Copper wire
400 mL beaker
10.0 mL 3.0 M HCl
Burette clamp
Procedure:
1. Obtain a piece of magnesium ribbon that is 1-2 cm long. Measure the length of the ribbon
carefully and record the value to the nearest 0.01cm.
2. Record the mass of 1 meter of ribbon this information may be posted on the white board.
3. Fold the
magnesium ribbon
so that it can be
encased in a small
spiral cage made of
fine copper wire.
Let enough wire
serve as a handle
so that the cage can
be placed
approximately 2cm
form the stopper.
4. Set up a ring and
burette clamp in
position to hold a
50 mL gas
measuring tube
which has been
fitted with a one or
two hole rubber
stopper as per
figure 2a. Place a
400 mL beaker
about two thirds
full of tap water
near the ring
stand.
Figure 1
5. Tilt the gas
measuring tube
slightly and pour
in 3M HCl to
about the 10mL
mark.
123
Figure 2
Figure 3
6. With the tube still tilted, slowly fill it with tap water from the beaker. While pouring rinse any
acid that may be on the sides of the tube so that the final liquid near the opening of the tube
contains very little acid. Avoid stirring up the acid layer in the bottom. You are attempting to
have a layer of less dense water on top of a layer of more dense acid (See figure 3) .
7. Holding the copper coil by the handle, insert it into the tube until the wire case is positioned
between the 49mL -50mL marks. Hook the copper wire over the edge of the tube and clamp it
there by inserting the rubber stopper. When properly set up, the tube will have no air bubbles
and the water will completely fill the holes in the stoppers as well as the tube. If the tube is not
full, attempt to add water through the stopper holes by gently using a wash bottle. (see figure 3)
Again, be certain that the acid layer stay, at the bottom undisturbed.
8. Cover the holes in the stopper with your
finger and invert the tube in the container
of water. Clamp the tube in place. The
acid being denser than water will begin to
stream down through the water and will
eventually react with the metal.
9. Record observations, feel the gas tube with
your hand and observe and record any
changes in temperature.
10. After the reaction has stopped, wait about 5
minutes to allow the tube and its contest to
come to room temperature. Bubbles
clinging to the sides of the tube can be
dislodged by gently tapping the tube.
11. Cover the hole(s) in the stopper with your
finger and transfer the tube to a large
cylinder or battery jar which is almost
filled with water at room temperature.
Raise or lower the tube until the level of
the liquid inside the tube is the same as the
level outside of the tube. This permits you
Figure 4
to measure the volume of gases in the tube (hydrogen and water vapor) at the air pressure in the
room. Read the volume with your eyes at the same level as the bottom of the meniscus. Record
the volume of gas to the nearest 0.01mL.
12. Place your finger over stopper holes and remove the gas collection tube from the water. Pour
the remaining acid solution down the sink. Rinse the tube with tap water.
13. Record the room temperature, record the air pressure in the room.
124
Data:
Air pressure in room
Air temperature in room
Length of Mg ribbon used in experiment
Length of magnesium standard
Mass of magnesium standard
Volume of gas produced
Observations:
___________
___________
___________
___________ (AS WRITTEN ON BOARD)
___________ (AS WRITTEN ON BOARD)
___________
Calculations:
A. To find grams of Magnesium ribbon
1. Convert cm of Mg to grams of Mg
2. Convert grams of Mg to Moles Mg
3. Calculate temperature in Kelvin
B.
1.
2.
3.
To find theoretical moles of hydrogen gas produced
Find the partial pressure of hydrogen gas
Find experimental ideal gas constant.
Percent Error
C. Determine experimental molar volume
1. Determine the volume in liters that gas would occupy at STP.
2. You may wish to use PV=nRT. STP is defined as 1atm, 273 K
Results:
A.
B.
C.
D.
E.
F.
Grams of Magnesium ribbon
Theoretical Moles of H2 gas produced
Experimental volume of gas produced
Experimental Gas Constant
Percent error
Molar Volume (liters/Mol, at STP)
___________
___________
___________
___________
___________
___________ (STP= temp@273K and 101.2 kPa)
Discussion:
Discuss topics presented in “How to write a lab report”
Post-Lab Questions:
1. In a second experiment, 0.113 grams of zinc reacted with 3.0 Molar HCl. The pressure is 0.995
atm and the temperature is 21 oC. 38.35 mL of gas were collected over water. From this data
calculate the experimental value for R. What is the percent error for the experimental R (the
theoretical value of R is 0.08206 mol*K/atm*L)?
References:
Unknown (unknown). Experiment 10 Reaction of a metal with hydrochloric acid. unknown
Willbraham, A. Staley, D. Matta, M. (1995) Chemistry 4th edition. Addison-Wesley: Menlo Park,
CA
125
Experiment 19
Heat of Neutralization
Purpose:
The purpose of this laboratory activity is to compare the heat of neutralization of two different acids and to
calculate the experimental heat of formation for water.
Hypothesis: To develop the hypothesis, compare the net ionic equations for the reaction of nitric acid
and sodium hydroxide to the reaction of hydrochloric acid and sodium hydroxide. From the net ionic
equation develop a hypothesis which compares the heat of neutralization for nitric acid to the heat of
neutralization for hydrochloric acid.
Background:
Driving Forces
There are two primary reasons that reactions take place: enthalpy and entropy. Some
chemists refer to these as driving forces, as in the question posed in many chemistry classes, "What
is the primary driving force in this reaction…?"
Entropy is the measure of disorder in a system. This is beyond the scope of this laboratory
activity. However it is sufficient to think that in general the universe tends towards disorder. In
terms of energy, generally energy is released as things move toward disorder and it takes energy to
either organize things or to fight off entropy.
Enthalpy is the measure of heat (molecular movement) that is present in a substance. The
universe has a fixed amount of enthalpy. There is only so much energy, and you cannot create or
destroy energy. When a bond is formed the atoms become more stable for each
atom now has each of its orbitals
Energy diagram
filled. As a result when a bond is
formed the atoms have less energy.
Thus when a bond is formed energy
is released to the environment. The
atoms constantly attempt to reach the
lowest possible energy state, just as a
ball rolls down hill to the lowest
possible energy state, a state of rest
with no potential energy. Atoms
react to make bonds releasing energy
in effect rolling down hill and giving
off energy as they do so. For instance
two materials that have very little
energy are water and carbon dioxide.
Carbon dioxide and water are the
products of so many reactions
Figure 1
because they have a low potential
126
energy. Water and carbon dioxide are
like the rock at the bottom of the hill.
Reactions which release energy
are called exothermic (see Figure 2).
Exothermic reactions release energy to
the environment. They feel hot to us; this
is because energy is released when the
new bonds are formed to create the
product bonds.
Exothermic reaction gives off energy (Figure 2)
Some reactions are
Diagram of a endothermic reaction
endothermic (see Figure 3).
A reaction that releases energy to the environment
That is, they require energy
from the environment to
proceed. This indicates that the
product bonds formed release
less energy than is required to
break the reactant bonds. Thus
energy from the environment is
needed to keep breaking bonds.
As energy is repeatedly
removed from the environment
the environment starts to feel
cold. This is because heat
energy is removed from it. As
heat energy is removed the
molecules begin to move more
slowly and this is what is
measured as cold. Endothermic
reactions feel cold to the touch,
because they require energy
Figure 3
from the environment to
proceed. They
get this energy from the surrounding air or any chemist who happens to touch the reaction vessel.
Why would a reaction that constantly requires heat be able to proceed spontaneously? A
spontaneous reaction is a reaction that will take place without further addition of heat energy.
That is, once the activation energy is met (the energy required to start the reaction) the reaction
will proceed without any additional input of energy from the surroundings. If a reaction is
endothermic (meaning it draws energy from its surroundings to proceed) what could possibly be its
driving force? The answer is entropy.
127
Heat:
Heat is the transfer of energy due to differences in temperature. Heat is not the same as
temperature. When heat is released in a chemical reaction the reaction is referred to as an
exothermic reaction. When heat is absorbed the reaction is termed endothermic. The heat released
in a reaction can be measured using a known quantity (mass) of water. As heat is released from the
chemical reaction it is absorbed by the water and the water’s temperature goes up. The amount that
the temperature goes up indicates how much energy was released in the reaction.
There are several different types of reactions that release energy in the form of heat.
Heat of formation:
The amount of heat involved in the production of one mole of a
substance.
Heat of combustion: The amount of heat produced when 1 mole of a material is reacted with
oxygen.
Heat of solution:
The amount of energy produced when one mole of a material is dissolved
in water.
Heat of vaporization: The amount of energy involved when one mole of a substance changes
phase from a liquid to a gas (vaporization) or changes phase from gas to a
liquid (condensation).
Heat of fusion:
The amount of energy involved when one mole of a substance changes
phase from a liquid to a solid (freezing) or changes phase from solid to a
liquid (melting).
Heat of neutralization: The amount of heat involved when on mole of water is created when an
acid reacts with a base, the literature value is -55.90 kJ/mol.
The heat that is released or absorbed in a chemical reaction (which maintains a constant
pressure) is called ∆H. The ∆ is a capital delta (Greek letter), which symbolizes “a change in”. The
∆ indicates a difference in the amount of heat in a system. This is expressed mathematically as
∆H= Hf-Hi
Hf = heat final
Hi = heat intial
If heat is released the ∆H is negative and the reaction is exothermic.
If heat is absorbed the ∆H is positive and the reaction is endothermic.
In the lab you will be using a simple calorimeter.
When acid is neutralized water is produced. This reaction releases heat.
Because both the acid and the base in this reaction are dissolved in water
the heat given off by this reaction will be absorbed by the surrounding
water. If this temperature increase is carefully measured the amount of
energy released can be calculated. Some heat energy may be lost to the
surroundings. This heat loss will be small because of the insulated
containers. The amount of heat lost to the surrounding air will be
accounted for by plotting a cooling rate graph.
128
The heat from the reaction is absorbed
by the water. The temperature of water
increases. However some of the heat
escapes to the surrounding air. Using
the graph a line can be drawn to
extrapolate the temperature at the time
the solutions are mixed.
Temp
final
The greatest difference between the
temperature of the reactants before they
were mixed and the products after they
were mixed is the ∆T, or change in
temperature.
Temp initial
This will be used in the equation
q=mc∆T.
Cooling-Rate Graph
Pre-Lab Questions:
1. Explain the following in regards to enthalpy and entropy.
a. When sodium chloride is dissolved into water the surrounding water gets colder. Is this
physical change endothermic or exothermic? Why is this reaction spontaneous? Provide
evidence for your answer.
b. Many gases are held under enough pressure that the molecules come so close together that
they become liquid. When the pressure is released the compressed gas in liquid phase
returns to its gaseous state. When this happens the surroundings become very cold. This is
in essence how a refrigerator works. Is this physical change endothermic or exothermic?
Why is this reaction spontaneous? Provide evidence for your answer.
2. The temperature of 100.00mL of water was raised 25.5 oC. How much heat in joules was
added to the water? How much heat in calories was this? (the specific heat of water is
4.184 j/g oC or 1.00 cal/g oC)
q=mc∆T
Where
q = energy
m = mass
c = the specific heat
∆T = change in temperature
3. If 50.0mL of 1.80 M HCl was added to 50.0 mL of 1.80 M NaOH, calculate the molarity of
the resulting NaCL solution. (Hint: Calculate the moles of NaCl made. The resulting
solution has a volume of 100mL.)
Materials:
Calorimeter
Thermometer
Watch with second hand
1.0 M NaOH
1.0 M HCl
1.0 M HNO3
129
100.0 mL graduated cylinder
Procedure:
Heat of neutralization: HCl and NaOH
9. Set up calorimeter as shown in illustration. This is the same way it was set up in the
previous experiment. Make sure the thermometer does not touch the bottom or sides of the
cups.
10. Obtain about 50.0mL of 1.0 M HCl with the graduated cylinder. Pour the acid into the
calorimeter. Record actual volume.
11. Measure about 50.0 mL of 1.0 M NaOH with a graduated cylinder. Pour contents into a
beaker. Allow this solution to sit in the beaker while reading the temperature of the HCl in
the calorimeter. Record actual volume.
12. Measure and record the temperature of the acid in the calorimeter to the nearest 0.1 oC. This
it the first temperature, T0. After 30 seconds record the temperature again, this is T1.
Continue recording data points for 2 minutes and 30 seconds. Assume the NaOH solution is
at the same temperature. Is this a valid assumption?
13. At exactly 3 minutes, T6, quickly pour the NaOH into the calorimeter begin the chemical
reaction. You cannot measure the temperature accurately at this time because it will be
changing very rapidly. After you have poured the NaOH into the calorimetric immediately
replace the lid, thermometer and stirrer and gently stir the mixture.
14. 30 seconds after mixing, T7, read the thermometer and record the temperature. Read and
record the temperature at 30 second intervals. Stir the mixture occasionally. Continue
stirring and recording data for 4 more minutes.
15. Draw a graph of temperature (Y-axis) vs. Time (X-axis). Connect the points from T0 to T5.
The line should be nearly horizontal. Use this line to extrapolate what the initial
temperature at T6 was.
16. Draw a best fit line between points T7 -T16. Use this line to extrapolate a second value for
T6. This line is the final temperature.
17. The difference between the two extrapolated lines is the ∆T.
Heat of neutralization: HNO3 (nitric acid) and NaOH.
1. Repeat the above using 50 mL of 1.0 M HNO3 and 50mL of 1.0 M NaOH. Calculate ∆T as
before with extrapolated lines on a graph.
Data:
Trial 1à
Volume
HCl and
NaOH
HCl
NaOH Trial 2 à
Write Data
In
HNO3
Note
NaOH
Book
Temperature at various times…
T0
T1
T2
T3
T4
T5
T6
T7
Write Data In
Lab Note Book base
added
T8
T9
Etc
Record all actual volumes used, use these values for all calculation. Record all observation.
130
Calculations:
1. Write and balance the reaction for aqueous HCl reacting with NaOH.
2. Determine how many moles of water are produced when 1 mole of HCl reacts with 1 mole of
NaOH.
3. For trial one use the volume of HCl or NaOH which ever one is less (it acts as a limiting
reactant) and find the moles.
4. Using dimensional analysis find the moles of water produced based on item 3 above.
5. Using dimensional analysis calculate the mass of solution used in the reaction. Assume that the
solution formed is volume of acid plus vol of base. The density of the salt solution is 0.9982
g/mL.
6. From the mass of the solution (m), the specific heat (c)of the solution and its change in
temperature (∆T). Calculate the amount of heat evolved by the reaction. The specific heat of
this solution is 4.017 J/goC,
Heat absorbed by water=mass of solution * specific heat * ∆T
q= mc∆T
7. Calculate the amount of heat evolved by the reaction. This is the amount of heat absorbed by
the water, multiplied by -1. This ∆H is for the number of moles of water your reaction
produced.
Heat evolved by the reaction = -(heat absorbed by the water) = ∆H
∆H = -q
8. Calculate the value of ∆H for 1 mole of water being produced by this reaction. Refer to the
answers from calculation steps 3 and 7. This value of ∆H is the heat of neutralization. Be
certain that the units are J/mole.
9. Is the reaction exothermic or endothermic? What sign should the ∆H have, - or +.
10. Repeat the above steps for the neutralization of acetic acid. Assume that the sodium acetate
solution produced has the density and specific heat as the sodium chloride solution.
Results:
A.
B.
C.
D.
E.
∆H of water from the HCl reaction.
∆H of water from the acetic acid reaction.
Report the theoretical ∆H of water.
Report % error from the HCl reaction. Please background information for theoretical value.
Report % error from the HNO3 reaction. Please background information for theoretical value.
Discussion: Discuss the results of your experiment in regards to supporting or refuting the
stated hypothesis (do not use the word “prove”). Is the heat of neutralization the same regardless of
which acid was used? Explain your answer using the net ionic equation. Answer the discussion
questions as presented in your “How to Write a Lab Report.”
Post-Lab Questions:
1. When 0.25 mole of NH4Cl is dissolved in 1.00 L of water the temperature drops 12.5 oC.
(Assume 1013g of solution and a specific heat of 4.02 J/g oC). What is the heat of solution for
ammonium chloride?
131
Experiment 20
Molecular Model Mobile Project
I want to get this assignment to you as soon as possible because some of you may be going places
for spring break. I would strongly urge you to read this before spring break arrives so that any
misunderstandings get cleared up before you leave. Essentially you will research a chemical,
determine its molecular structure, and build a model. I would strongly urge you to not spend a lot
of money. Also I am rather partial to space filling models, though a ball and stick model is fine.
Molecule
Four different elements
1
Three different shapes
1
Two different bonds
1
You must recevie 3/3 in this section or no credit for the project will be awarded. Really cool
molecules might earn extra credit. Your model may not be scored if another student is doing
the same chemical. Please check the sign up sheet so that you will receive credit.
Paper
If your paper even appears to be copied from some other source you will be asked to rewrite
the whole document for reduced credit.
Typed, Times New Roman, 1" margins, 12 point font, clean clear paper.
1
Paragraph of the chemical’s history, who discovered it and why. If it is a naturally 1
occurring material who first determined the chemical formula? Trivial info
The societal context of this chemical? (Why is it important? How is it used?)
1
A complete reference section in APA format
1 internet citation 1
2 print book references (textbook, encyclopedia, other books, etc.) 1
Journal citation 1
One reference citation (CRC, Merck) 1
The Model
Your model will hang from the ceiling from no more than 2 points and the lowest
point of model will be less than 75cm (including sign) from the ceiling. If it does
not meet these requirements you will receive no credit.
The model is sturdy; it will not fall apart
The model is not made of food
Bond angles are correct
Relative size of atoms is close to accurate. Things in one energy level are all about
the same size, things in lower energy levels are smaller; things in higher energy
levels are larger.
A sign with the name of the molecule, the MW, and formula in IUPAC notation
Different colors to represent different atoms
Presentation
I am dressed professionally, speak professionally, act professionally
Presentation lasts between 1 1/2 and 3 minutes
The presentation is informative and accurate
132
1
1
1
1
1
1
1
1
1
Molecular Model Mobile Project
SCORE SHEET
TO BE TORN OUT AND TURNED IN WITH PAPER.
PAPER IS DUE BEFORE YOU GIVE YOUR TALK.
Scoring
Molecule
Four different elements
Three different shapes
Two different bonds
1
1
1
You must recevie 3/3 in this section or no credit for the project will be awarded. Really cool
molecules might earn extra credit. Your model may not be scored if another student is doing
the same chemical. Please check the sign up sheet so that you will receive credit.
Paper
If your paper even appears to be copied from some other source you will be asked to rewrite
the whole document for reduced credit.
Typed, Times New Roman, 1" margins, 12 point font, clean clear paper.
1
Paragraph of the chemical’s history, who discovered it and why. If it is a naturally 1
occurring material who first determined the chemical formula? Trivial info
The societal context of this chemical? (Why is it important? How is it used?)
1
A complete reference section in APA format
1 internet citation 1
2 print book references (textbook, encyclopedia, other books, etc.) 1
Journal citation 1
One reference citation (CRC, Merck) 1
The Model
Your model will hang from the ceiling from no more than 2 points and the lowest
point of model will be less than 75cm (including sign) from the ceiling. If it does
not meet these requirements you will receive no credit.
The model is sturdy; it will not fall apart
The model is not made of food
Bond angles are correct
Relative size of atoms is close to accurate. Things in one energy level are all about
the same size, things in lower energy levels are smaller; things in higher energy
levels are larger.
A sign with the name of the molecule, the MW, and formula in IUPAC notation
Different colors to represent different atoms
Presentation
I am dressed professionally, speak professionally, act professionally
Presentation lasts between 1 1/2 and 3 minutes
The presentation is informative and accurate
133
1
1
1
1
1
1
1
1
1
134
Appendix
Solubility Rules:
1.
2.
3.
4.
5.
6.
7.
Salt containing sodium, and potassium ions, ammonium and acids are always soluble.
Salts containing nitrates, chlorate, perchlorate and acetate are always soluble.
All chlorides, bromides and iodides are soluble except for those of silver, lead II, and mercury I, which are
insoluble.
All sulfates are soluble except of those of calcium, silver, strontium, barium, mercury I, mercury II and lead.
All hydroxides are insoluble expect those of alkali metals, which are soluble, and hydroxides of calcium, barium,
and strontium, which are moderately soluble.
All sulfites , carbonates, chromates and phosphates are insoluble expect those of ammonium and alkali metals.
All sulfides and insoluble expect those of ammonium, the alkali metals and the alkaline earth metals, which are
soluble.
Bond
Energy
kJ/mol
H-H
H-F
H-Cl
H-Br
H-I
C-H
C-C
C-Si
C-N
C-O
C-P
C-S
C-F
C-Cl
432
565
427
363
295
413
347
301
305
358
264
259
453
339
Bond
C-Br
C-I
C=C
C=N
C=O
C=O (CO2)
F-F
F-Cl
F-Br
F-I
N-H
N-N
N-P
N-O
Selected Specific heat capacities
[J/g* oC]
H2O(liquid)
4.18
Iron
0.46
Acetone
Benzene
Hofus kJ/mol
5.72
9.87
Energy
kJ/mol
Bond
Energy
kJ/mol
Bond
Energy
kJ/mol
Bond
Energy
kJ/mol
276
216
614
615
745
799
159
193
212
263
391
160
209
201
N-F
N-Cl
N-Br
N-I
O-H
O-P
O-O
O-S
O-F
O-Cl
O-Br
O-I
O2
N=N
272
200
243
159
467
351
204
265
190
203
234
234
498
418
N=O
N=N
S-H
S-S
S-F
S-Cl
S-Br
Si-Si
Si-O
Si-S
Si-Cl
Si-Br
Si-I
P-H
607
945
347
266
327
271
218
226
368
226
381
310
234
320
P-P
P-F
P-Cl
P-Br
P-I
C=C
C=N
C=O
Cl-Cl
Cl-Br
Cl-I
Br-Br
Br-I
I-I
200
490
331
272
184
839
891
1070
243
215
208
193
175
151
H2O(solid)
Aluminum
Hovap kJ/mol
29.1
30.8
1Atm = 101.3 kilo
pascal
1Atm = 760 mmHg
1Atm = 760 Torr
1Atm = 1013 mBar
PV = nRT
P1 V 1 = P2 V 2
(P1V1)/T1 = (P2V2)/T2
Pt = P1+P2+P3+Pn
R = 0.08206
Density of water
Temp oC
g/ml
0 (solid)
0.9150
0 (liquid)
0.9999
4
1.000
20
0.9982
40
0.9922
60
0.9832
80
0.9718
100(steam) 0.0006
% error = 100* (|Exp-Theo|/Theo)
% yield = 100* (Exp/Theo)
Absolute error = |Theo – Exp|
135
[J/g* oC]
2.1
0.90
Water
Methanol
H2O(gas)
ethanol
[J/g* oC]
1.7
2.4
Hofus kJ/mol
6.01
.94
Hovap kJ/mol
40.7
8.2
1cm = 2.54 inches (∞ sig figs)
cm3 = 1 mL
1 mole = 6.022 x1023
Q = mCΔT
λ* ν = c
c=2.998x108 meter per second
Planck’s constant = 6.626x10-34 J*sec
ΔE = h * ν
RH = 2.178x10-18 J
M1V1=M2V2
1
1
2
ΔE = -RH
Nf
Ni 2
(
)