Survey Chemistry
Lab Manual
CHEM 106L
revised 2011
Table of Contents
page
Principles for Safety in the Chemical Laboratory................................................................1
Laboratory Notebooks......................................................................................................... 5
Lab 1 Calculator and Measurements...................................................................................9
Lab 2 Glassware and Lab Equipment.................................................................................12
Lab 3 Density........................................................................................ ............................15
Lab 4 Stoichiometry............................................................................................................20
Lab 5 Identification of Cations by Flame Test...................................................................25
Lab 6 Carbon Monoxide in Auto Exhaust...........................................................................30
Lab 7 Using the Library to locate scientific papers.............................................................37
Lab 8 Thermochemistry.....................................................................................................39
Lab 9 Natural pH Indicators................................................................................................45
Appendix I.: Notes on Significant Figures...........................................................................53
Principles for Safety in the
Chemical Laboratory
Safe practices in the chemical laboratory are of prime importance. A student should
consider it an essential part of his or her educational experience to develop safe and efficient
methods of operation in a lab. To do this, one must acquire a basic knowledge of properties
of materials present in the lab, and one should realize the types of hazards that exist and the
accidents and injuries that can result from ignorance or irresponsibility on the part of the
student or a neighbor.
Regulations
1.
Wear safety goggles at all times while in the laboratory.
2.
Report all accidents to the instructor or lab assistant immediately.
3.
NEVER eat, drink, chew, or smoke in the laboratory.
4.
NEVER leave an experiment unattended. Inform the instructor or lab assistant if you
must leave the lab.
5.
After the experiment is completed, turn all equipment off, making sure it is properly
stored, and clean your area.
Failure to comply with these regulations is cause for immediate dismissal from lab.
Precautions
1.
Approach the laboratory with a serious awareness of personal responsibility and
consideration for others in the lab.
2.
Become familiar with the location of safety equipment, such as acid-base neutralizing
agents, eye wash, fire extinguisher, and emergency shower.
3.
Pay strict attention to all instructions presented by the instructor. If something is not
clear, do not hesitate to ask the instructor or lab assistant.
4.
Clean up all chemical spills immediately.
5.
Be aware of all activities occurring within a reasonable proximity of yourself since
you are always subject to the actions of others.
6.
To avoid contamination of community supplies, do not use personal equipment such
as spatulas in shared chemicals and replace all lids after use.
1
7.
Avoid unnecessary physical contact with chemicals; their toxic properties may result
in skin irritation.
8.
Use all electrical and heating equipment carefully to prevent shocks and burns.
9.
NEVER handle broken glassware with your hands; use a broom and a dust pan.
10.
Wash your hands at the end of the laboratory.
Personal Attire
Choice of clothing for the laboratory is mainly left to the discretion of the student. Because
of the corrosive nature of chemicals, it is in your best interest to wear comfortable, practical
clothing. Long, floppy sleeves can easily come into contact with chemicals. A lab coat is
suggested to help keep clothes protected and close to the body. Accessories also need
consideration. Jewelry can be ruined by contact with chemicals. Open toed shoes do not
adequately protect one against chemical spills. If hair is long enough to interfere with motion
or observation, it should be tied back. Remember that your clothes are worn to protect you.
Assembling Equipment
Equipment should be assembled in the most secure and convenient manner. Utility clamps
are provided to fasten flasks, etc., to ring stands. This keeps top-heavy or bulky equipment
away from the edge of the bench where it can be knocked easily off . Consider the safe
location of the hot plate. Keep it away from the bench edge to minimize chances of contact
with the body. If the aspirator is being used, locate your apparatus near the sink or center
trough for convenience.
Handling Glassware
Laboratory glassware is usually fragile, and if it is not properly handled, serious injuries may
result. Do not force glass tubing or thermometers into a rubber stopper. Lubricate the tubing
or thermometer with glycerol or water, wrap it in a towel, and gently insert it into the stopper
by using pressure in a lengthwise direction while rotating it. Always grasp the tubing near
the stopper. When removing the tubing, remember to protect your hands with a towel. If
there are difficulties with this procedure, ask for the instructor's assistance. Apparatus that
can roll should be placed between two immobile objects away from the edge of the bench.
Chipped or broken glassware cannot be used. There are special receptacles near each bench
for these waste materials. After the experiment is completed, all glassware should be
emptied, rinsed, cleaned and returned to your drawer. The equipment in your drawer is used
by several groups during the week. Each group must pay attention to cleaning and returning
equipment to the drawer or the next group will not have necessary items to complete the
experiments.
2
Acids and Bases
In this lab sequence, you will come in contact with several acids and bases. As with all
chemicals, caution must be taken to prevent contact with the skin. When handling these
chemicals, keep hands away from the eyes and face until they have been thoroughly washed.
If an acid or base comes in contact with your skin, flush the area with large quantities of
clean, cold water. Eyes are extremely sensitive. Use the eye wash provided in the laboratory,
or wash with water for at least 10 minutes. Again, the instructor must be notified
immediately. To insure your safety, neutralize acid or base spills before cleaning them up.
Boric acid is available to neutralize base spills, and carbonate powder is provided to
neutralize acids.
Before leaving lab on the first day please 1) sign the provided sheet (page
4) indicating that you have read and understand the lab safety guidelines
and 2) sketch the lab indicating all of the safety features and their
locations 3) have your instructor check and initial your sheet before you
turn it in and 4) sign and hand in the waiver/release form (page 8)
indicating you understand ; the dangers involved in wearing contact lens in
a chemistry lab, release Black Hills State University from liability, and
consent to receive medical treatment in the event of an injury during this
lab.
3
Survey Chemistry Lab CHEM 106L
I have read and understand the safety guidelines for working in the chemistry lab.
Name
Date
Signature
Please sketch the laboratory below and indicate the location of all safety items in the
lab. Also indicate first aid kit and phone locations in relation to the lab. Then have
your instructor check and initial before turning in.
Instructor initials
4
Laboratory Notebooks
You are required to use a notebook in lab to record all primary data and observations. ( a
spiral notebook without perforated pages works well) You should prepare your notebook
each week before coming to lab by writing the title of the experiment on a new numbered
page, summarizing relevant equations from the lab manual, and starting calculations
involving molar masses, etc. Take note of theoretical ideas and special instructions given
by your instructor at the start of each experiment. Your notebook should be a complete
record of your work in lab. You or other chemists should be able to understand the notes
in the future, not just during the current experiment Good note taking in lab is a valuable
skill that you can learn with a little effort and practice.
Guidelines to be Followed:
1. Always bring your notebook with you to lab. You will be graded on the completeness
of your previous note taking and your preparation for the current experiment. You may
use your notebook during a lab quiz.
2. Number the pages sequentially and reserve space at the beginning for a table of
contents.
3. Take your notebook to the balance room, etc. and record values directly in it - not on
loose scraps of paper.
4. Specify each measured quantity by name and include the units.
5. If you make a mistake in your notebook, simply draw a solid line through the error and
write the correction nearby.
6. Tables greatly simplify data entry; they should be set up before coming to lab.
7. Write down all observations such as color and phase changes - don't rely on your
memory.
Before coming to lab this notebook should have entered in it •
a lab title and page reference (where in the note book this lab is written up)
on a table of contents page
•
numbered pages
•
Lab Title, date and lab partners name/s at top of first page for that lab
•
objective for the lab
•
answers to pre-lab questions (showing work)
•
a short procedure for the lab (you can cut and paste from the lab handout)
•
any reaction equations
•
chemical formulas and calculation of molecular masses/weights which will
5
•
be needed for calculating the results from lab
any work that can be done ahead of time will greatly improve your ability to
get through the lab in a timely manner
After lab is finished the notebook should contain •
raw data from experiment (this includes copies of all graphs!)
•
set ups for calculations
•
all final answers
•
answers to the post lab questions (showing work)
•
any and all observations made during the lab - observations can be: color,
lack of color, temperature, physical state (solid, liquid, gas), a mistake you
made in the procedure and how you dealt with it.
•
all information needed for your report sheets should be in your notebook
first and then transferred to the report sheet. That way report sheets are a bit
neater and easier to grade.
•
last of all you need to write up a short conclusion for the lab - your
conclusion should address the objectives of the lab - did you achieve the
objective? Was it easy to do the lab? What parts were difficult? What
exactly did you do to achieve the goals? Include final answers/values that
were determined in the lab.
If your lab notebook is kept up to date and everything gets written in it you will find
quizzes and the final exam much easier as both are open lab notebook!!
6
Attention:
Students are advised against wearing contact lenses while
observing or participating in science laboratory activities. While
hard contact lenses do not seem to aggravate chemical splash
injuries, soft contact lenses absorb vapors and may aggravate
some chemical exposures, particularly if worn for extended
periods.
Please take your contact lenses out prior to entering the
laboratory.
Contact Lens Administrative Policy and Waiver Form
Students are advised against wearing contact lenses while
observing or participating in science laboratory activities.
While hard contact lenses do not seem to aggravate chemical
splash injuries, soft contact lenses absorb vapors and may
aggravate some chemical exposures, particularly if worn for
extended periods. You are asked to please remove your
contact lenses prior to entering the laboratory.
Please fill out and sign the waiver/medical release form on
the next page indicating that you understand and accept
responsibility for wearing contacts in lab. If you do not use
contact lens, this form is also; a liability release and a release
form for allowing medical treatment in cases of injury.
7
Waiver of Liability, Indemnification and Medical Release
I am aware of the dangers involved in wearing contact lenses in a science
laboratory setting. On behalf of myself, my executors, administrators, heirs,
next of kin, successors, and assigns, I hereby:
waive, release and discharge from any and all liability for my personal
a.
injury, property damage, or actions of any kind, which may hereafter,
accrue to me and my estate, the State of South Dakota, and its officers,
agents and employees; and
indemnify and hold harmless the State of South Dakota, and its
b.
officers, agents and employees from and against any and all liabilities
and claims made by other individuals or entities as a result of any of
my actions during this laboratory.
I hereby consent to receive any medical treatment, which may be deemed
advisable in the event of injury during this laboratory.
This release and waiver shall be construed broadly to provide a release and
waiver to the maximum extent permissible under applicable law.
I, the undersigned participant, acknowledge that I have read and understand
the above release.
Name ______________________________________ Age _____________
Signature ___________________________________ Date ____________
8
Measurements and the Calculator
Objective:
<
To become familiar with taking a variety of measurements
<
Practice converting measurements from one type of unit to another
<
Learn how to efficiently use your calculator to obtain the correct answer
Exercises: (all calculations must be included in your notebook)
A.
Length:
Practice Problems:
1.
Convert your height in inches to mm.
2.
Convert your partner’s height in inches to kilofeet.
Using the designated measuring tape determine the length of the whiteboard. Convert the
length of the whiteboard to miles, km and nanoinches.
B.
The Balance:
Practice Problems:
1.
Convert 350 mg to Gg.
2.
Convert 3 tons to kg
Using the designated balance, determine the mass of a 250mL beaker. Convert the mass
of the beaker from its current units to nanograms, megagrams and ounces.
C.
The Graduated Cylinder:
Practice Problems:
1.
Convert 27 mL to dL
2.
Convert 4 pints to cL
9
Using your graduated cylinder and designated balance, determine the volume occupied by
25 grams of water. (Be careful to clean up any water that you spill on or near the
balance!) Convert the volume of water from its original units to liters, microliters and in 3 .
(Note 1 cm 3 = 1 mL).
Post-Lab Questions:
1.
Define density. Did you use the concept of density in any of the above
exercises? If so, which one?
2.
Since you have calculated that 25 grams of water is equal to so many mL of
water, you now have a conversion factor. Using this conversion factor,
calculate how many grams of water you would have, if you had a can of
water? (To answer this question you need to know that a can of water
contains 355 mL of water.)
3
Perform the following calculations on your calculator. If your answer does
not match the answer given, consult your laboratory instructor.
A.
2+3x2=8
B.
4 x 10 – 2 / 5 = 39.6
C.
5 x 10 -2 = 0.05
D.
5 x 10 -2 – 2 = -1.95
E.
(5 x 10 -2) 2 = 2.5 x 10 -3
F.
2 x 3 x10 3 + 2 / 8 = 6000 or 6 x 10 3
G.
3 x (9 x 10 -3)5 = 1.77x10 -10
10
Names
Measurements and the Calculator Report Sheet
(Show all calculations!!)
1.
inches
Length of the Whiteboard
_________________ miles
km
nanoinches
2.
_________________ g
Mass of the beaker
nanograms
megagrams
ounces
3.
Volume of water that weighs 25grams
_________________ ml
liters
microliters (ìl)
in 3 (cubic inches)
IV.
What is the density of water from your experimental data?
11
Laboratory Glassware and Equipment
Objectives:
<
learn to identify the various pieces of glassware and equipment used in a
chemistry lab
<
understand the uses and limitations of the these items
In today’s lab you will take a self guided mini-tour of chemistry lab glassware and
equipment. You will find several stations set up around the lab with information on the
items found at each station. You may do the stations in any order. Make sure you write
down the information in your lab notebook and draw pictures/diagrams of the various
special set-ups so you will know how to set them up later in the semester.
G
Station #1
Beakers and Erlenmeyer Flasks
G
Station #2
Graduate Cylinders
G
Station #3
Burets and Volumetric Pipets
G
Station #4
Mortar & Pestle, Crucible
G
Station #5
Special Set-ups for Filtration
G
Filtering for a pure liquid
G
Filtering for a pure solid
G
Station #6
Water Baths
G
Ice Bath
G
Boiling Water Bath
G
Station #7
Uses of Watch Glasses, glass stirring rods, metal scoops and eye
droppers
G
Station #8
Propane Torch Set-Up
G
Stations #9 & #10 Using the Balance
G
Balance in Lab
G
Balance in Hallway with Lid
Make sure that all of the information necessary to answer the questions on the report
sheet gets into your notebook first, then answer the questions on the report sheet and
turn it in!!
12
Names
Section ______
Glassware & Lab Equipment Report Sheet
1.
What are the two most common uses of beakers and erlenmeyer flasks in the
chemistry lab?
2.
What piece of glassware should you use if you need to measure out exactly 75.0
milliliters of a solution?
3.
What piece of glassware could you use to measure out about 50 milliliters of a
liquid if the volume didn’t need to be exactly 50ml, (if + or - 5ml would be
accurate enough) and you just needed to have the same volume for all reactions?
4.
If you need to know exactly how much of a solution you added to a reaction
mixture for all the sample to react what is the best piece of glassware to use?
5.
Draw and label the filtration Set-Up that would be used if you were interested in
obtaining a clean, pure solid.
13
6.
What type of water should you use in a boiling water bath and what is the purpose
of the rock chips in the bottom of the beaker?
7.
Explain how you would measure the pH of a solution using pH paper.
8.
What is the name of the piece of equipment that is used to light a propane torch
that does not have a built in igniter?
9.
How do you use the balance to weigh out 5grams of a crystalline solid like sugar?
10.
Do the balances have any limitations? Explain!!
14
Density
Objective:
The objective of this experiment is to develop the skills and knowledge necessary
to determine the density of a variety of objects.
Background Information:
Density is the relationship that the mass of an object has with its volume.
Mathematically, density is the ratio of the mass of an object to its volume. Density is an
intensive property. This means that a particular liquid (or solid) will have the same
density regardless of the amount you use to calculate density.
Density = (mass of an object) / (volume of that object)
or
Density = mass / volume
For example: if 25 mL of gasoline had a mass of 22 grams, the density would be
calculated as follows:
Density of gasoline = (mass of the gasoline) / (volume of the gasoline)
Density of gasoline = (22 grams of gasoline) / (25 mL of gasoline)
Density of gasoline = 0.88 g / mL
Specific gravity is defined as the density of an object divided by the density of
pure water. Typically, the temperature used for measuring the density of water is 4 o C.
Other temperatures may be used, but the temperature must be referenced. The density of
pure water at 4 oC is 1.00 g/mL. The specific gravity of gasoline would then be:
Specific Gravity of gasoline = (density of gasoline) / (density of water at 4 oC)
Specific Gravity of gasoline = (0.88 g/mL) / (1.00 g/mL) = 0.88
Note that Specific Gravity is a unit-less number
15
Procedure:
Today you will be determining density of a variety of liquids and solids
experimentally. You will need to determine the volume occupied by the liquid or solid
either, experimentally, or by direct calculation. The volume of liquids can be determined
by using a graduated cylinder. For some of the solid objects the volume can be
determined directly by calculation if you know the formula for its volume. (example: the
volume of a rectangular solid = height X length X width) For objects with irregular
shape the volume can be determined experimentally by displacement of water. (Example:
If you place a known volume of water in a measuring cup, say half a cup, you can
measure out shortening by placing the shortening in the cup of water so that all the
shortening is under water. If the water volume now comes up to the one cup mark you
known that you have half a cup of shortening. Final volume of water minus initial volume
of water = volume of object being measured.) Keeping this example in mind devise a
method to measure the volume of the irregular objects using the equipment on your
equipment list.
To calculate the density you will also need to determine the mass for each volume
of liquid or solid. The solids are simple, just weigh them on the balance and record their
mass. You will need to determine the mass of the liquids by difference. This means you
will need to know the mass of the graduated cylinder you are using to measure the
volume of the liquid, and the mass of the graduated cylinder with the liquid in it. Using
these two pieces of data you can determine the mass of the liquid by difference.
Mass of cylinder with liquid in it - mass of empty cylinder = mass of liquid
In short you need to determine volume and mass for each item you wish to
calculate density for then set up the ratio of mass to volume and calculate the density.
Remember to include units with all numbers. And pay attention to how and what you are
measuring, volume and mass!! As you do these calculation make sure they match up with
the real world!! You all know that some objects will float on water while others sink.
Given that the density of water is ~1.00g/ml if the object you are determining density for
floats, its calculated density had better not be greater than 1g/ml and if it sinks it better
not be less than 1g/ml!!
You should be able to write out a brief procedure for what measurements you need
to do and how you will do them for each of the items you will calculate density for.
The liquids will all follow the same procedure. But, the solids will have two
procedures for measuring volume, direct calculation or volume by difference!
16
Pre-Laboratory Questions: (use the factor-label method when possible)
1.
3.25 grams of acetone occupies a volume of 2.57 mL. What is the density of
acetone?
2.
What is the density of acetone in the units oz / ft3?
3.
What is the volume, in terms of mL, occupied by 17 grams of mercury?
(The density of mercury is 13.6 g / mL.)
4.
What is the mass, in terms of grams, of 15 in 3 of gold? (Gold has a density
of 19.3 g / mL)
Equipment Needed:
1 – 100 mL graduated cylinder
2 – 10 mL graduated cylinders
1 – metric ruler
1 – stirring rod
1 – test tube
Experimental Determination of Density
I.
Liquids
Do not add liquids to the graduated cylinder while the graduated cylinder is on the
balance!!!
Using a balance, determine the weight of each individual 10ml graduated cylinder.
(Keep track of the cylinders as they do not all weigh the same amount!) Add
somewhere between 7 and 9 mL of the liquid you wish to determine density of, to the
graduated cylinder (recording the exact volume) and determine the weight of the
graduated cylinder with the added liquid. The difference in the weight of the graduated
cylinder and the graduated cylinder plus the liquid, is the weight of the liquid. Determine
the density for the following items:
1.
Antifreeze
2.
Power Steering Fluid
3.
Molasses or Corn Syrup
4.
Laundry Detergent or Fabric Softener
17
II.
Solids - determine the density for the following solids
1.
Block of white pine
2.
Rubber stopper
3.
Cork
4.
One penny
(Note: the rubber stopper, cork, and penny are irregular shapes and the volume cannot
be determined through a simple volume calculation.)
III.
Estimation of the density of Liquids
Gently add small amounts of 30w oil, antifreeze, power-steering fluid and
transmission fluid, to a test tube. Keep track of the liquids by observing their colors and
location in the test tube. Estimate the density of the transmission fluid and the 30W oil by
referring back to the densities you have calculated for antifreeze and power steering fluid
from exercise I. Explain your answer. Add a couple of drops of water to the test tube.
What is the result?
Post-Laboratory Questions:
1. What would happen if I added acetone to the mixture in part III of this experiment?
(Remember you calculated the density of acetone in the pre-lab questions)
2. Determine the specific gravity of molasses or corn syrup.
3. I have a giant marble whose diameter is 14 inches and weighs of 3 pounds. What is
its density in g/mL? (What information do you need to solve this?? Look this information
up while in lab!)
18
Names ______________________________________Section ______
Density Report Sheet
Report your answers using the proper number of significant figures. Remember
answers must have the correct units.
What was the order of the layers from Part III. starting with the heaviest and
ending with the lightest for the four automotive fluids you layered into the test tube?
Provide estimates for the unknown densities.
Weight
Heaviest
Heavy
Light
Lightest
Fluid Name
Provide the following densities:
Antifreeze
____________________
Power Steering fluid
Molasses or Corn Syrup
Density
____________________
____________________
(circle one)
Laundry Detergent or Fabric Softener
(circle one)
____________________
Block of white pine
____________________
Rubber stopper
____________________
Cork
____________________
One-penny
____________________
How many significant decimal places can be used when using the following:
10 mL graduated cylinder
100 mL graduated cylinder
Balance
Metric ruler
19
Stoichiometry
Objectives:
<
Become familiar with the technique involved in an acid/base titration.
<
Calculate the amount of sodium bicarbonate in an unknown solution.
Background Information:
Today we will be using stoichiometry to calculate the amount of sodium
bicarbonate in an unknown solution of sodium bicarbonate. Stoichiometry is the
calculation of chemical quantities based on balanced chemical equations.
Today’s balanced reaction equation is:
HCl + NaHCO 3 à H 2 O + CO 2 + NaCl
By knowing Avogadro’s number, 6.022 X 10 23, which is the number of individual
units contained in one mole of that substance and by knowing the molecular mass of the
substance you will be able to use these numbers to calculate masses of substances from
their concentration and the number of molecules present from their mass.
Another calculation you can do is the percent by mass, the ratio of the mass of
solute (material dissolved in the solution) to the mass of the solution multiplied by 100.
Today we will be experimentally determining the amount of sodium bicarbonate in
an unknown solution using a technique called an acid/base titration. In a titration you
use a buret, a special piece of glassware designed to deliver variable volumes accurately.
In a titration you add titrant (today our titrant is HCl) to your unknown solution until all
of your unknown has reacted. In today’s reaction we will determine the endpoint of the
titration (the point at which all of the unknown has reacted and a very small amount of
HCl has been added in excess) by using an acid/base indicator made up of a combination
of methyl red and bromcresol green. This indicator, methyl red/bromcresol green, is a
different color in an acidic solution than it is in a basic solution and will change color
when the first drop of excess acid is added to your reaction flask signaling the endpoint of
the reaction.
You will be titrating your sample of the unknown solution of sodium bicarbonate
with a solution of hydrochloric acid of known concentration. The three pieces of data
you will need to calculate the mass percent of sodium bicarbonate in the unknown are:
1.
Mass percent of the stock HCl solution. This will be written on the label
of the stock solution.
2.
Mass of stock HCl solution used to titrate the sample of unknown
20
solution. This can be calculated by determining the total volume of stock
HCl solution needed to react all of the sodium bicarbonate in your
unknown sample, final volume minus initial volume of stock HCl solution in
buret. Then set up the conversion to mass in grams using the density of the
HCl stock solution, which is 1.00g/mL.
3.
Mass of the sample of the unknown solution you titrated. This can easily
be obtained by weighing your reaction flask empty and weighing it again
after you have added your unknown solution. The difference between these
two masses will be the mass of your sample of the unknown solution.
You will be gathering experimental data on; the volume of HCl used, from reading
initial and final volumes on the buret, and mass of your unknown sample by difference,
mass of sample & flask minus mass of flask. (These measurements are the raw data)
You will need to have the molecular masses for the reactants and products of the
reaction calculated ahead of time in order to do the calculations. Make sure they are in
your lab notebook.
Pre-Laboratory Exercises: (use the factor-label method when possible)
1.
What is the mass, in grams, of 1.2 moles of carbon?
2.
Calculate the number of moles of aluminum in 1.50 kg of Al.
3.
Calculate the number of moles of CaCO 3 in 3.20 grams of calcium
carbonate.
4.
From the following reaction:
How many grams CO 2 would be produced if 2.6 grams of NaHCO 3 reacted
completely with an excess amount of HCl?
Equipment Needed:
1 - 50mL burette
1 - 250mL beaker (to store your HCl titrant in)
1 - watch glass
(to cover your stored titrant)
1 - 50 or 100ml beaker (to use to add titrant to buret)
1 - 250 or 300mL Erlenmeyer flask (to run reaction in)
21
Procedure:
G
Using a clean, dry 250mL beaker obtain ~125mL of the HCl stock solution. (The
density of the HCl solution is 1.00 g/mL). Cover with a watch glass and place on
your bench where it will not be knocked over. Record the concentration of the
stock solution.
G
Set up your 50mL burette.
rinse burette with 5-10mL of your stock HCl solution make sure the
stopcock on the burette is closed and carefully fill the burette with HCl
stock solution
run some solution out of the burette to fill the tip with solution and clear
any bubbles present there
add more HCl stock if needed, and record the initial volume
G
Obtain your unknown sample solution by carefully pre-weighing a 250 or 300 mL
Erlenmeyer flask then have your instructor add some of the unknown solution to
your flask. Accurately weigh your flask + unknown sample. The difference in
masses of the flask with unknown sample and the empty flask will give you the
mass of your unknown solution.
G
Add a 5-8 drops of the methyl red/bromocresol green indicator to your flask, mix
by swirling gently and record the color of the solution. (Use same amount in each
titration)
G
Carefully add HCl solution to your flask, from the buret, while swirling the flask
gently. Remember to record the inital volume of HCl in your buret before
beginning your titration. When you see the indicator change color, stop the
addition of the HCl solution to your Erlenmeyer flask. All of the sodium
bicarbonate has reacted. Record the final volume of HCl solution in your buret.
The difference between the final volume and the initial volume will give you the
volume of HCl you used to react with all of your unknown solution.
G
Refill your burette and record initial volume.
G
Clean out your reaction flask.(you don’t need to dry the inside, just the outside)
Re-weigh and obtain a second sample (preferably slightly heavier than your first
sample so it will take a bit more HCl to react with it. Why??)
22
G
Repeat the experiment.
G
Using the data from the titration that was the closest to the endpoint, (usually the
second one), calculate the mass of sodium bicarbonate in the sample you titrated,
and the mass percent of sodium bicarbonate in the unknown solution.
G
When finished with experiment make sure to clean any glassware you used and put
it back in your drawer. And rinse the buret 2-3 times with de-ionized water and
store tip up and stopcock open, on the buret clamp for the next lab
Post-Laboratory Questions (show all work):
I.
How many grams of NaHCO 3 are in a 25gram sample of the unknown solution if it
takes 35grams of a 2.5% HCl solution to titrate it?
2.
What was the percent by mass of NaHCO 3 in the original unknown solution for
question # 1 above? (The NaHCO 3 is the solute and mixture of the water and
NaHCO 3 is the solution.)
3.
How many grams of water were produced from the reaction in #1 above?
4.
How many atoms of Na were present during the reaction in #1 above?
5.
What gas was produced from the reaction?
23
Names __________________________________ section _____
Stoichiometry Report Sheet
Report your answers with the proper number of significant figures. Remember answers
must have the correct units.
Final buret reading
Initial buret reading
Volume of HCl solution used in titration
Concentration of HCl stock solution
Mass of the sodium bicarbonate solution
sample you titrated
What was the percent by mass of sodium bicarbonate in the original unknown?
Show all calculations, for mass of sodium bicarbonate in unknown sample and
% NaHCO 3 in unknown sample
What was the initial color of the bicarbonate solution once the methyl red/ bromcresol
green was added?
What was the final color of the solution once the reaction was complete?
24
Flame Test and Emission Spectra
Objectives:
<
<
<
<
Identify cations by flame test
Determine the differences in energy of the light emitted by the cations
Using a spectroscope obtain the emission spectra for hydrogen
Identify an unknown element from its emission spectra
Background Information:
Many compounds that you use everyday are classified as ionic compounds; for
example sodium chloride, also known as table salt, is an ionic compound made up of a
sodium cations (positively charges ions) and chloride anions (negatively charged ions).
When sodium chloride is placed in water it dissociates to provide a homogeneous mixture
as shown in the equation below.
As mentioned above, there are many different ionic compounds that exist in
nature. One method for determining which cation is present in a solution is through a
flame test. Each cation produces a unique color when subjected to a flame. The flame
(energy source) excites electrons to higher energy states (levels). When the electrons
return to a lower energy state, energy is emitted in the form of visible light unique to that
cation.
In this experiment, a flame loop (wire loop) will be used to capture a small amount
of the aqueous solution containing a certain cation. The loop will then be placed in the
flame of a propane torch and color of the flame noted. Focus will be placed on sodium,
potassium, lithium, barium and copper ions.
In addition to the flame test, a second exercise will focus on using a spectroscope
to identify elements used in elemental lamps by examining their emission spectra. A
spectroscope is a device that splits light emitted by various light sources, today elemental
lamps, into spectral lines in the visible range. This type of spectra is called a line spectra
and is unique to the element emitting the light.
Light in the visible spectrum runs from the long wavelength red light, through;
orange, yellow, green, blue, and indigo to the shorter wavelength violet light. Beyond the
red end of the visible spectrum you find the long wavelength light referred to as infrared.
Likewise beyond the violet light of the visible spectrum you find the short wavelength
light referred to as ultraviolet . The shorter the wavelength of the light the more energy it
has, thus ultraviolet is high energy light, and infrared is low energy light. Wavelengths of
light are usually measured in nanometers, 10 -9 meters, so when you draw in the spectral
lines on the sheets provided to obtain the emission spectra for the elemental lamps you
will see that they are labeled in nm or ë. But, some of the spectroscopes we are using are
calibrated in angstroms, . You will need to convert from angstroms to nanometers
before drawing the spectral lines.
25
Pre-Laboratory Exercises:
1.
What is absorbance? (use dictionary)
2.
What is transmittance?
3.
What is an angstrom?
4.
What is the conversion factor for converting angstroms to nanometers?
5.
What is the color and energy make-up of the visible spectrum?
6.
How does a color’s wavelength relate to its energy?
Equipment Needed:
1 - Propane torch
1 - Striker
1 - Flame loop
1 - Rack of test tubes containing cation solutions (found on top of your bench)
1 - 100mL or 150mL beaker
Cautions for this lab: Barium salts are toxic and should not be dumped into the sinks..
26
Experiment I.
Using the flame test, observe the color of the visible light that is emitted by each of
the known cation solutions, and by comparison determine the cation present in an
unknown solution.
Procedure:
G
With the instructor’s supervision, light a porpane torch and adjust to obtain a blue
flame.
G
Observe the color emitted by each of the known salt solutions as follows:
a.
Dip your flame loop into a test tube containing a salt solution. (Note: all of
the solutions are referred to as salt solutions as they contain ionic
compounds that dissociate into cations and anions in water solution.)
b.
Place the flame loop into the flame of the propane torch (the top of the blue
cone of flame is the best place) and observe the color emitted.
c.
Record the cation contained in the solution, and, the color it emitted when
placed in the flame, in your lab notebook.
d.
Clean off your wire loop by dipping it into a small beaker of de-ionized
water and flaming it to dry it.
e.
Repeat this exercise to obtain the color produced for each cation in solution.
G
Obtain an unknown sample and record the number in your lab notebook.
G
Using the flame test determine the cation in your unknown sample. Record this in
your lab notebook.
Experiment II.
G
Using the elemental lamps and spectroscopes, obtain the emission spectra for
hydrogen. Draw the lines on the sheets provided and paste it into your notebook
along with any observations you make.
G
Examine the emission spectra for the unknown element. (Remember to record the
unknown number.) Draw it on the sheet provided, compare it to the emission
spectra provided for mercury, helium, and neon to determine the identity of the
unknown element. Paste this spectrum in your lab notebook and indicate the
identity of the unknown.
27
Post-Laboratory Questions:
1.
Which cation emits light with the lowest energy?
2.
Which cation emits light with the highest energy?
3.
Is the color we see due to transmittance or absorbance?
4.
Name two ways to excite electrons to higher energy states?
5.
In the hydrogen line spectra, which line/wavelength represents the highest energy?
6.
Using the symbols for each element from the periodic table indicate what the
cation in aqueous solution (dissolved in water) would be, and its charge, for the
following ionic compounds.
(Example: sodium chloride, NaCl, would have the cation Na + in aqueous solution)
potassium chloride
barium dichloride
copper dichloride
lithium chloride
28
Names
Section
Flame Test and Emission Spectra
Report Sheet
Questions:
1.
What was the color emitted by sodium?
2.
What was the color emitted by potassium?
3.
What was the color emitted by lithium?
4.
What was the color emitted by barium?
5
What was the color emitted by copper?
6.
What was your unknown number?
7.
What was the color emitted by your unknown cation?
8.
What is your unknown cation?
9.
In the hydrogen emission spectra, what color lines did you observe and what were
their wavelengths?
10.
In the hydrogen emission spectra, which line represents the highest energy?
For Bonus Points (2pts if correct)
unknown lamp number
ID of element in lamp
29
Using Infrared light to determine the level of
CO in auto exhaust
Objectives:
<
<
<
<
become familiar with the technique of collecting auto exhaust samples
learn how to load and run gas samples on the IR spectrophotometer
determine what gases are present in the sample
estimate the amount of CO and/or CO 2 in the exhaust sample
Introduction
Electromagnetic radiation is a continuum of wavelengths and energies. The part of
the electromagnetic spectrum you are most familiar with is the visible spectrum. Here,
the wavelength of violet light is about 400 nm, green about 500nm, yellow about 580 nm
and red from 620 to 750 nm. You also learned that the energy of light is inversely
proportional to the wavelength, that is violet light with the shortest wavelength has the
highest energy, while red light with the longest wavelength has the lowest energy.
When your eye perceives light, for instance the color green, it is seeing an excess
of light in the 500-550 nm region, while the light at most other wavelengths is being
absorbed. Since your eye is not the most sensitive light measuring device, scientists rely
on machines called spectrophotometers to measure the amount of light absorbed at
different wavelength accurately. Below is the spectrum of a green dye solution as
measured in a Visible light spectrophotometer.
Notice how this green compound has maximum absorbance at 410 and 630 nm,
indicating that light of these wavelength is being absorbed, while there is 0 absorbance in
the 500-550 nm region, showing how this light is transmitted so you can see it. This is
why the solution appears green. All compounds absorb electromagnetic radiation at
different wavelengths, and these different absorption spectra can be used both to identify
and calculate amount of each different chemical present.
30
The amount of light absorbed is directly proportional to the amount of the
chemical in the light path. Below is the absorption spectrum of the same green compound
from page 1, only now two absorption spectra are plotted on the same graph, one for the
original solution as shown on page 1 and one for a solution ½ as concentrated as the
original solution. Notice how the absorbance peaks are at exactly the same wavelengths,
but the sample absorbs exactly ½ as much light as before. This relationship between
absorbance and concentration allows us to calculate the amount of a substance in a
sample by making a calibration curve by plotting absorbance vs. concentration, for known
concentrations of that substance on graph paper. Where the plot is linear, a straight line,
the values will be most accurate.
Not all chemicals absorb light in the visible region of the spectrum. Many
compounds will absorb light in the ultraviolet region (10-400 nm, higher energy) While
most compounds will also absorb light in the infrared region (800-100,000 nm, much
lower energy)
In today’s lab we will be using an infrared spectrophotometer to measure the light
that passes through a sample of auto exhaust gas. Air itself contains about 80% N 2 and
20% O 2 and these gases have no absorbance in the infrared region. In a well tuned engine
with a properly working catalytic converter, the only exhaust gases you should have are
CO 2 (carbon dioxide) and H 2 O. If something is wrong with the engine or the exhaust
system, other gases like CO (carbon monoxide), NO (nitrogen monoxide) and NO 2
(nitrogen dioxide) will appear. All of these gases have absorbances in the infrared, so an
infrared spectrum of exhaust gases can be used to diagnose engine efficiency.
31
.Reference Spectrum 1 : CO gas
Reference Spectrum 2: CO 2 gas
Pre-Lab Questions:
1.
If you have a red solution, at what wavelengths should it absorb light, at what
wavelengths should it NOT absorb light?
2.
What would the absorbance at 630 nm be if the sample graphed on page 30 was
2x as concentrated? (see graph of page 30)
3.
If you have peaks in your exhaust sample, other than those shown in the above
reference spectra, what could they be?
32
Part I.
Procedure
For this lab we will divide into groups of 3-4 . Depending on the cars available,
each group may sample a different car, or some groups may take different samples from
the same car. The sampling device is simply a garbage bag with a cardboard tube taped
to each end, and an injection septum for a hypodermic needle mounted in the middle.
Collecting the Sample
Start the car, if it is an automatic, make sure it is park, if it is a manual, make sure
it is in neutral and the parking brake is set. Allow the engine to warm up for 5 minutes.
Insert one cardboard tube, the long one, a short distance up the exhaust pipe, and allow
the gas to run through the sample bag for 1 minute. During this time make sure the gas
flows easily through the bag, and try to avoid breathing the exhaust gas as much a
possible. At the end of a minute, spin the cardboard tubes at each end of the bag to seal
up the bag and remove the tube from the exhaust pipe. Use wire twist ties to hold bag
closed. Turn off the engine, and bring the sample bag back to the lab, and place it in a
hood.
Filling the sample cell
MUST BE DONE IN THE HOOD
First insert needles into the septa on sample bag and sample cell. Then use a 50 ml
syringe with a needle to remove a 50 ml sample of exhaust gas from the bag. The IR
sample cell contains two septa. You will need to stick a needle through each septum to
allow gas to move through the cell, as you inject your sample into one septum gas will be
pushed out the other. You should flush the cell 2-3 times with exhaust gas before filling
it with your final sample. (About 200mL total volume) It is easier if you leave the needles
in the septa, and detach the syringe for each flush, rather than pushing a needle through
the septum for every flush. Once it is filled, remove both needles from the sample cell, to
keep the sample inside the cell. Also remove the needle from the sample bag septum and
store all needles in their protective covers. At this point you may take the sample cell to
the Infrared spectrophotometer and obtain a spectrum for your sample.
You should vent all remaining exhaust gas out of you sample bag by squeezing the
air out of the bag into the back of the hood, with the hood on, so the excess gases are
removed from the lab.
Obtain a spectrum of exhaust gas (Using the directions on the Infrared Spectrometer.)
Identify the peaks (by comparing to the reference spectra)
Indicate any unknown peaks (identify if you can)
Calculate the percent carbon monoxide and/or carbon dioxide in your sample (see
page 34 and 35)
33
Part II.
Determination of CO in exhaust gas
As you saw on page 31 the amount of material in a sample correlates with how
much the material absorbs light. One way we can measure the light absorbance of a
sample is to measure the area ‘under’ an absorbance peak. In a separate experiment we
have already measured the area ‘under’ the CO peak that runs from roughly 4450 nm to
4925 nm. Below is our data. Plot this data on the graph and find the line that best fits
this data to make a calibration curve of area vs. concentration. (y vs x)
% CO
Area
1%
4.19
5%
12.12
10%
18.98
12.5%
22.86
Questions to Answer:
Using the directions on the IR machine, determine the area under the CO peak for
your vehicle’s exhaust gas. Record the area of the peak and the wavelengths over which
it was measured in your notebook. What were the actual wavelengths you used for your
area? What was your area?
Find your area on the calibration curve you just plotted. What concentration of CO
does your area correspond to?
34
Part III.
Determination of CO 2 in exhaust gas
As you saw on page 31 the amount of material in a sample correlates with how
much the material absorbs light. One way we can measure the light absorbance of a
sample is to measure the area ‘under’ an absorbance peak. In a separate experiment we
have already measured the area ‘under’ the CO 2 peak that runs from roughly 2650 nm to
2850nm. Below is our data. Plot this data on the graph and find the line that best fits this
data to make a calibration curve of area vs. concentration. (y vs x)
% CO 2
Area
25%
25.9
50%
52.0
75%
69.8
100%
98.8
Questions to Answer:
Using the directions on the IR machine, determine the area under the CO 2 peak for
your vehicle’s exhaust gas. What were the actual wavelengths you used for your area?
What was your area?
Find your area on the calibration curve you just plotted. What concentration of
CO 2 does your area correspond to?
35
Names: ________________________________________Section
CO in Auto Exhaust Report Sheet
Questions to be answered:
1. Using the reference spectra given on page 32, identify all major absorption peaks
found in your exhaust gas sample. Label these peaks on your IR Spectra. Do you have
any unknown peaks? If there were any unknowns peaks, even if you didn’t have any,
what compounds might be causing them?
2. What was the area under your CO and/or CO 2 peak? For each compound between
what two wavelengths was this area measured?
3. What concentration of CO and/or CO 2 do these peak areas, from number 2 above,
correspond to?
Make a copies of; your calibration curves and your IR Spectrum, attach
them to this report sheet.
(HINT: You should also have a copies of these in your lab notebook!!)
36
Library Lab
Objectives:
<
to become familiar with the data bases available at the library
<
learn how to do a literature search using the data bases to find articles on
a specific topic or by a specific author
For this lab we will meet in the library at the regular lab time, instead of in the
chem lab. Attendance will be taken and those attending will receive 5 points. The report
sheet questions will be worth another 5 points, bringing the total for this lab to 10 points.
One of the librarians will be giving a talk on how to do a literature search for
scientific papers using the data bases available at the library. For this talk they use a large
screen and actually show you exactly how to go about doing a search. The talk itself
usually takes 45 minutes to an hour. Remember to treat this as a regular lab and take
notes in your lab notebook.
BONUS POINT OPPORTUNITY!!
With the second half of the lab you will be able to go to the library computers and
conduct a search for papers authored by Mrs. Betty R. Zehfus. (Note: On one of the
papers they have her middle initial incorrect) Print out the article title, authors, journal,
etc., to document each different paper you find and you will receive one bonus point per
paper found. (Hints: she has not published since the late 80's and has between 5 and 10
papers in print.)
Please turn in:
<
your report sheet
<
documentation of your literature search for bonus points
37
Name
Section
Library Lab Report Sheet
1.
What was the name of the librarian that talked to your lab?
2.
Name two of the data bases which contain scientific literature.
3.
Name two specific types of searches that can be conducted.
4.
How can you access “pubmed” from the library site?
5.
How do you access the library data bases if you are not at one of the computers on
campus?
38
THERMOCHEMISTRY
Objectives:
<
use a calorimeter to measure the temperature change which occurs
during a chemical reaction
<
calculate the heat energy evolved by the reaction
Background
When a chemical reaction occurs in a system at constant pressure under
conditions such that no energy is lost to or gained from the surroundings, it is generally
found that the temperature of the system either increases or decreases. When the system is
in thermal contact with the surroundings, energy will be lost as heat if the temperature
rises due to reaction and the reaction is said to be exothermic. If there is a decrease in
temperature of the system due to reaction, energy will be gained from the surroundings as
heat and the reaction is said to be endothermic.
Heat changes that take place in a system at constant pressure are related to a
property of the system called enthalpy, denoted by the symbol H. The SI unit of enthalpy
is the same as that of energy, joule (J) Enthalpy is a state function, and can be related to
heat changes at constant pressure by:
heat absorbed = increase in enthalpy = ÄH = H final - H initial
(1)
When a system is configured such that no heat is gained or lost, the system is said
to be under adiabatic conditions. Thus for adiabatic processes at constant pressure, the
change in the enthalpy of the system equals the heat absorbed by the system, which is
zero. Since enthalpy can be changed by a change in temperature ( )T ) or by a chemical
reaction or by both, a chemical reaction carried out under adiabatic conditions would
result in the following relationships:
)H
system
= )H
)H
ÄT
+ )H
due to reaction
due to reaction
= - )H ÄT
=0
(2)
(3)
The change in enthalpy due to temperature changes can be calculated by:
)H
)T
= specific heat capacity H mass H )T
(4)
where )T = T f - T i and can be positive (for exothermic reactions), or negative (for
endothermic reactions). If the change in temperature is measured for a chemical reaction
39
carried out under adiabatic conditions, the enthalpy change due to the reaction can be
determined from the enthalpy change causing the temperature change.
Enthalpy is also an extensive property, thus depending on the amount of reaction
that occurs. To obtain an intensive property for comparing reactions, we normally relate
the enthalpy change to the amount of one of the reactants or products (e.g., species A).
Thus a reaction enthalpy is given as:
(5)
The SI unit of )H rxn is J mol-1 , although reaction enthalpies are typically reported
in kJ mol-1 . Since we need the amount of A reacted, it is useful to select the limiting
reagent as species A.
We have seen in lecture that reaction enthalpies can be calculated from tabularized
values of standard molar enthalpies of formation, from tabularized combustion enthalpies,
and from experimental calorimetry data. It is this last method which will be used in this
experiment using the principles derived above.
The Calorimeter
A calorimeter is a device with which we can measure the enthalpy change in a
system during a reaction by monitoring the temperature increase (for an exothermic
reaction) or decrease (for an endothermic reaction). Several types of calorimeters have
been demonstrated in the text and in lecture. In this experiment, the calorimeter used is a
styrofoam capped cup. The foam between the liquid and surroundings acts as a thermal
barrier to prevent heat loss from the reaction to the surroundings. The reactions will be
carried out in aqueous solution in the cup, and the temperature will be monitored with a
standard laboratory thermometer. We will make two major assumptions about the
calorimeter and system. First, the device will be considered an ideal calorimeter,
operating adiabatically at constant pressure; i.e., all of the heat evolved by the reaction
goes to raising the temperature of the reaction solution, and none is lost to the beaker and
air space. Second, we assume that the final solution is sufficiently dilute so that its density
and specific heat capacity will be those of pure water (1.00 g mL -1 and 4.184 J g –1 K -1,
respectively). Both of these assumptions will introduce some error.
The enthalpy change associated with the temperature change is calculated from equation
(4) above. From this result, the enthalpy change due to the reaction can be obtained from
equation (3). This result can be converted to kJ and then, after the amount of limiting
reagent is determined, the ÄH rxn can be determined from equation (5).
40
Chemistry of the Reaction
The reaction enthalpy for the following reaction will be determined:
HCl (aq) + NaOH
(aq)
6
NaCl (aq) + H 20
(1)
Note that HCl and NaOH are strong (acid and base, respectively). Because they are
strong you know that they will ionize totally in aqueous solution. All of the HCl will
ionize to give H + ions and Cl- ions in aqueous solution. Likewise all of the NaOH will
ionize to give Na + ions and OH - ions in aqueous solution. Of the four ions found in this
aqueous reaction solution only the H + and OH - ions are actually involved in the reaction.
The Na + and Cl- ions are still present after the reaction is complete. Since they are present
on both sides of the reaction arrow they are considered to be spectators.
To understand what is actually occurring in the reaction today a different way of
writing the reaction can be used. Since we know that HCl and NaOH both ionize totally
in aqueous solution we can write the reaction as a complete ionic reaction like this:
complete ionic equation:
H +(aq) + Cl-(aq) + Na +(aq) + OH -(aq)
6 Cl
(aq)
+ Na +(aq) + H 2 O (l)
Then from the complete ionic equation you can easily write the net ionic equation
by simply removing the spectator ions from the equation, leaving only the actual
participants of the reaction. Like this:
net ionic equation:
H + (aq) + OH -(aq)
6
H 2 O (l)
Written like this it is much easier to see what is actually occurring in the reaction.
This reaction is a Bronsted-Lowry acid-base reaction, it involves a proton transfer
from the acid to the base. More stable bonds are formed in these reactions, so the
reactions are, therefore, exothermic.
Pre-Lab Questions:
1.
For a reaction you measure out 55.0ml of solution A and add it to 55.0ml of
solution B to run the reaction, what is the total mass of the reaction mixture, given
the density of both solutions is 1.00g/ml?
2.
If your initial temperature for a reaction is 24.0 C and the final temperature is
42.0 C , what will )T ( )T = T f - T i ) for the reaction be in Kelvin (K)?
41
3.
Using the mass of reaction mixture from question 1 and the )T from question 2
calculate the change in enthalpy due to change in temperature for this reaction
()H )T ) given that the specific heat capacity for this reaction mixture is
4.1841 J g -1 K -1 . (use equation 4)
4.
You calculated )H ÄT for the reaction in question 3, what is )H
reaction? (Use equation 3)
5.
Using the )H due to reaction from question 4, if you had 0.10 moles of the limiting
reactant in the reaction, what would the ÄH rxn be? (Use equation 5)
due to reaction
for this
Procedure
Arrange your styrofoam cups, thermometer and cap as shown by the instructor.
Clean, dry, and label one 50 mL and one 100 mL graduated cylinders for the reactant
solutions to be used in this experiment (this will prevent cross-mixing of solutions).
Use the following procedure to measure )T:
G
Obtain 45.0 mL of the 2.0M HCl and 45.0mL of the 2.0M NaOH needed for the
reaction in the designated graduated cylinders.
G
Add the acid to the calorimeter, and then measure and record the initial
temperature (T initial ).
G
Rapidly pour the base into the calorimeter, replace the cover, and stir the reaction
solution with calorimeter stirrer.
G
Observe the temperature of the reaction mixture while continuously stirring, and
record the highest reading as the final temperature (T final ).
G
When finished with a measurement, pour out the solution, rinse and dry the cup,
stirrer, and thermometer, and perform the next run.
You need to do four runs of the reaction. Use the first run as a shake down cruise
to clean out the calorimeter and get everyone familiar with who needs to hold the cup and
stir, how to add solutions, and who will keep track of the temperature. Then do three
more runs to gather experimental data for your calculations.
42
Calculations and Report
In the following calculations, assume that the specific heat capacity for the
reaction mixture is 4.184 J g -1 K -1 and its density is 1.00 g mL -1. For each run, calculate
)T, change in enthalpy due to the reaction (in kJ), amount of limiting reagent, and the
reaction enthalpy (in kJ/mol limiting reagent). Show the calculation setup for run 1 of the
experiment as indicated on the report sheet. Calculate the average reaction enthalpy for
the experiment and then fill out the report sheet and answer the post lab questions.
Post lab Questions:
1.
HNO 3 is a strong acid. If you use 45.0 ml of a 2.0 M HNO 3 solution how many
moles of H + ions will it provide for reaction?
2.
HCl is a strong acid also. If you use 45.0 ml of a 2.0 M HCl solution how many
moles of H + ions will it provide for reaction?
3.
Would you get a different )H rxn for the experiment if you used 45.0 mL of 2.0 M
HNO 3 instead of HCl? Explain why or why not. (Hint: what do HCl and HNO 3
supply to the reaction?)
43
Names:
THERMOCHEMISTRY REPORT SHEET
Experimental Data (HCl/NaOH)
Run 1
Run 2
Run 3
_______
_______
_______
_______
_______
_______
_______
_______
_______
Amount limiting reagent (mol)
_______
_______
_______
)H rxn (kJ/mol)
_______
_______
_______
Initial Temperature, T i(o C)
Final Temperature, T f(oC)
)T(K)
)H ÄT (kJ)
Average )H rxn __________________kJ/mol
In the space below show the sequence of all the calculations for run 1:
44
Evaluating and Using Natural pH Indicators
Purpose:
<
compare several different natural pH indicators to assess their
effectiveness in determining pH of solutions
<
use the best one of these natural pH indicators to estimate the pH of
several common household products.
Background Information:
The acidity of solutions can be reported in two ways.
1.
as the concentration in moles per liter, molarity, of H 3O + . (H 3O + , the
hydronium ion, is the species that is found in aqueous solutions of acids as
H + + H 2 O 6 H 3 O +(aq).)
2.
or in pH units, pH is the negative log of the concentration of H 3 O + ,
pH= -log [H 3 O + ], and the pH scale covers a range from pH 0, with an H 3 O +
concentration of 1 M; to pH 14, with an H 3 O + concentration of 1x10 -14 M.
A solution is considered to be neutral if its pH is pH 7, while pH values less than
seven are acidic and greater than seven are basic.
pH
H 3 O + concentration
(M)
Interpretation
1
10 -1
very acidic
3
10 -3
acidic
5
10 -5
acidic
7
10 -7
neutral
9
10 -9
basic
-11
11
10
13
10 -13
basic
very basic
A pH indicator is a compound that changes color depending on the acidity of a
solution. Some have only two colors and the change of color occurs at a certain
concentration of H 3 O + in the solution. While others exhibit a range of colors each
occurring at different H 3O + concentrations. One way chemists use pH indicators is to
track the end points of acid/base titrations. They will select an indicator which changes
color at the pH that the solution will be at when the endpoint or equivalence point of the
reaction has been reached, giving them an easily seen color change to indicate the
endpoint of the reaction. Another frequently used type of pH indicator is litmus which is
blue in basic solution and red in acidic. Litmus is a naturally occurring pH indicator
found in plants. The most common use of this indicator is in the form of litmus paper
45
which comes in two types; blue litmus, paper impregnated with the basic form of litmus,
and red litmus, paper impregnated with the acidic form of litmus. These litmus papers can
be used to determine if a solution is acidic or basic, but only positive tests have any value.
For example - blue litmus will turn red if a drop of acidic solution is placed on it, but if
there is no color change the solution could be basic or neutral. Likewise red litmus will
turn blue if a drop of base is placed on it, but if there is no color change the solution could
be acidic or neutral. Even though its application is limited it does work very well in
separating acidic and basic solutions.
Another chemical class of compounds that produce a more varied reaction to
acidity are the anthocyanins. The anthocyanins contain several OH groups attached to
aromatic groups. Because these OH groups are slightly acidic as the concentration of
H 3O + decreases more H + ions are removed from the cyanin which in turn changes the
wavelength of light absorbed by the molecule, changing the color of the solution. These
compounds are found in; purple cabbage, beets, and black grape skins.
Pre-lab Questions
1.
A chemist is setting up an acid/base titration of NaOH with HCl. The endpoint
will occur at a pH of 7. Which of the following indicators would be the best at
indicating the endpoint?
Bromophenol blue which is yellow at pH 3 and blue at pH 4.5
Phenol red which is red at pH 8 and yellow at pH 7
2.
Three solutions are tested with blue litmus paper
solution #1 turns blue litmus red
solution #2 no change
solution #3 no change
From these results what can you say about these solutions? Acidic? Basic?
Neutral?
3.
The same three solutions from problem number two above are tested with red
litmus paper and the results are:
solution #1 no change
solution #2 no change
solution #3 turns red litmus blue
From these results what can you say about these solutions? Acidic? Basic?
Neutral?
4.
If you combine the results from number two and three what can you now deduce
about solution #2?
46
5.
If the concentration of H 3 O + in a solution is 1 x 10 -6 M. What is the pH of the
solution?
Equipment list
test tube rack
50ml beaker
eye dropper
7 test tubes
watch glass
glass stirring rod
red litmus paper
blue litmus paper
pH paper (pH1 - pH13 range)
In today’s lab you will be comparing and evaluating natural pH indicators. Three
made up from boiling water extracts of purple cabbage, beets and black grape skins and
you will also test litmus paper. The purple cabbage and beets were purchased from the
grocery store and the indicator of interest was extracted from them by boiling each in
deionized water until the solution turned dark. Then the indicator extract solution was
filtered through a coffee filter to remove the solid material. The black grape skins were
frozen (left over from last year’s grape jelly) once thawed they were extracted using the
same procedure used for the beets and cabbage.
CAUTIONS - HCl and NaOH, used for known pH solutions, and drain cleaner are
caustic
47
I.
Evaluating the Natural pH Indicators
Procedure
G
Clean seven test tubes and label them - pH 1, pH 3, pH 5, pH 7, pH 9, pH 11, and
pH 13
G
Pour about 1 - 2 ml of known pH test solutions into the tubes (make sure you have
the same volume in all tubes.)
G
Test red and blue litmus paper, lay pieces of red and blue litmus on a watch glass
then using a glass stirring rod, dip the rod into the pH 1 solution and place a drop
of the solution onto each. Record the results - red
blue, blue
red, or no change.
Clean off the stirring rod in deionized water and repeat the test for the next pH.
Repeat for all known pH solutions.
G
After you have recorded the results for the litmus papers pour about 10ml of the
beet extract into a 50 ml beaker and take it to your bench.
G
Add 3-5 drops of the beet extract to each of the known pH solutions. Mix them
gently by tapping with your finger. ( If the colors are too faint add a couple more
drops to each tube.)
G
Record the color produced by each of the known pH solutions. Be specific with
how you describe the colors and indicate if any of the tubes are the same color.
Save this set of tubes, it is a color key for beet extract.
G
Set up another set of tubes. Pour about 1 - 2 ml of known pH test solutions into the
tubes (make sure you have the same volume in all tubes.)
G
Pour about 10ml of the grape skin extract into a 50 ml beaker and take it to your
bench. Add 3-5 drops of the grape skin extract to each of the known pH solutions.
Mix them gently by tapping with your finger. ( If the colors are too faint add a
couple more drops to each tube.)
G
Record the color produced by each of the known pH solutions. Be specific with
how you describe the colors and indicate if any of the tubes are the same color.
This set of tubes, it is a color key for grape skin extract. Compare the color key
tubes from beet extract to those of grape skin extract and save the best indicator To determine the “best indicator”, look at the range of different colors produced
by each indicator, are there lots of different colors? Or only a couple? Are there
several pH levels that produce the same color? You want to save the one with the
most and clearest color changes.(the one that will be easiest to estimate pH over
the range ph 1- pH 13)
G
Empty the test tubes from the “worst” indicator clean them out and set up a third
set of known pH solutions. Use this set to test the purple cabbage extract. This set
of tubes, it is a color key for purple cabbage extract. Compare the color key tubes
48
from your best indicator of the first two to the purple cabbage and save the best
indicator - the one with the most and clearest color changes.(the one that will be
easiest to estimate pH over the range ph 1- pH 13)You will use this set of test
tubes as a direct color key for determining pH in Part II.
II.
Estimating the pH of Several Common Household Items
Procedure
G
Select the best natural pH indicator to use, write down your choice in your
notebook. This will be the one with the most and clearest color changes between
pH 1 and pH 13. You should have saved the known pH test tubes with this
indicator in them as a visual color key to compare to, as this will make estimating
pH easier.
G
Obtain 1-2 ml of each of the common substances available for testing (test 6 total)
G
Using pH paper test each substance. Use watch glass and glass stirring rod, record
color of paper and pH that it indicates.
G
Add 3-5 drops of your indicator to each tube and mix gently
G
Record the colors
G
Estimate the pH of each by going back to your table of colors shown by the
indicator at the known pH levels, or by comparing directly to the tubes of known
pH with your best indicator in them. (Remember if your indicator gives the same
color for pH 5 and pH 7 and one of your samples turns that color you need to
estimate the pH as being between pH 5 and pH 7.)
49
Post Lab Questions
1.
If you have one liter of a pH 5.0 solution and one liter of a pH 7.0 solution and you
mix them together what will be the pH of the resulting solution? (Hint: start by
figuring out how many moles of H 3 O + will be present in the final solution.)
2.
Calculate the pH of a solution that has
a. [H 3 O + ] = 1.0 x 10 -8 M
b.
[H 3 O + ] = 1.0 x 10 -10 M
3.
How many times more acidic is a pH 3 solution compared to a pH 5 solution?
4.
How many times more basic is a pH 11 solution compared to a pH 4?
5.
Read the labels on the products you estimated the pH for and see if you can figure
out what ingredient is responsible for the pH of the product.
Product
pH
Ingredient responsible for pH
50
Names ___________________________________________________ section _____
Report Sheet for Evaluating and Using Natural pH
Indicators
I.
Evaluating Natural pH Indicators:
Indicator Colors
Red Litmus
Blue Litmus
Beet Extract
Black Grape
Skin Extract
Purple
Cabbage
Extract
pH 1
pH 3
pH 5
pH 7
pH 9
pH 11
pH 13
Which of these natural pH indicators will you use to estimate the pH of common
substances?
Why did you choose this indicator over the others?
51
II.
Using A Natural pH Indicator to Estimate pH of Common Household
Substances
Name of Natural Indicator Used:
Substance
tested
Natural Indicator
color
pH
pH paper
color
Were all of the substances tested easy to read the pH value for?
Do your estimated pH values using the natural indicator agree with the pH values
obtained using the pH paper?
52
pH
Appendix I:
Notes on Keeping track of Significant Figures
Keeping track of significant figures can be extremely confusing on the best of days
for those unfamiliar with the process. Hopefully the following information will help you
to understand their use and importance.
Definition of Significant Figures: Significant figures include all of the certain digits of a
number and the first “uncertain” digit.
Whenever you see published numbers, like concentrations, volumes, or masses;
the last digit is the first uncertain digit; and it indicates the accuracy of the number.
For example if a concentration is given as 2.558 M, this indicates the solution is
2.558 M + or - 0.001M, as the uncertainty of the measurement would be in the last
digit.
In the laboratory the first “uncertain” digit is the one where you must estimate
between calibration markings when reading glassware, rulers, scales on meters, or
gridlines on graphs. We often need to read “between the lines”, whether it be reading the
volume on a buret, a graduated cylinder, or using a ruler to measure the length on an
object. When you read any value from calibrated markings the rule of thumb for figuring
out the uncertainty of the measurement is:
Divide the amount represented by the smallest calibration marks by 10 and this
will be the uncertainty for that particular item.
The above rule of thumb is based on the assumption that you will be able to
visually divide the space between the smallest calibration marks into tenths. Please note
that this uncertainty is a measure of how precisely you can read the equipment, not how
accurate the equipment is. If you use an accurate piece of equipment to measure volumes,
for example a graduated cylinder, and read it precisely then the volume measured will be
more accurate than the volume would be if you used a beaker, as beakers are not designed
to measure accurate volumes.
53
How many Significant Figures Should you have in the final answer?
In an experiment you need to make sure you record your data with the appropriate
significant figures, and keep track of the significant figures for stock solution
concentrations, densities, molecular masses needed in the final calculation.
For example: In an experiment you are gathering data to calculate the
concentration of a base in an unknown solution. You perform a titration with 2.00M Acid
on a 25.00ml sample, and your raw data from the titration gives you an initial volume of
acid in your buret of 0.25ml and a final buret volume of 34.37ml. After you subtract the
initial volume from the final volume of acid in the buret you have a total volume of
34.12ml. This number reflects four significant figures and is correct as when subtracting
or adding you keep the least number of decimal places. So, from your experimental
volume data you get 4 sig. figs. , from the sample volume of 25.00ml you have 4 sig.figs.,
and from the stock acid solution concentration of 2.00M you get 3 sig. figs. In your final
answer for the calculations you can’t have more significant figures than the than the
number with least significant figures used in the calculation. For this example your
answer would have to have three significant figures.
How to read numbers and determine significant figures
Some simple rules:
1.
In a written value all non-zero digits are significant
Example: 12,345.75 pounds has 7 significant figures and the
uncertain digit is found in the hundredths
2.
Trapped zeros are significant
Example: 2005 miles has 4 significant figures and the first
uncertain digit is found in the units.
3.
Leading or trailing zeros are not significant, these are zeros that
place the other digits.
Examples:
1120 has only three significant figures and the
uncertain digit is in the tens
0.0055 has only two significant figures and the
uncertain digit is in the ten thousandths
4.
Zeros after the decimal point are significant if there are non-zero
digits to the left of them
Example:
25.00ml has 4 sig. figs. and the uncertain digit is
in the hundredths
54
5.
Zeros with a decimal point after them are significant.
Example:
1250. has four sig. figs. and the uncertain digit
is the units
How many significant figures to retain after mathematical operations:
In addition and subtraction you are interested in decimal places. You keep the least
number of decimal places.
Example:
25.450 - 0.25 = 25.20
In multiplication and division you keep the number of significant figures found in
in number with the least number of significant figures.
Example:
3.00 x 225.55 = 676.65 so with sig. figs it would be 677
Writing numbers to indicate significant figures:
The best way to write very large or very small numbers to easily indicate
significant figures is scientific notation. The convention for writing a number in scientific
notation is to write a number with one digit to the left of the decimal and any other
significant digits to the right of the decimal and multiply this number by 10 raised to a
power.
Example 1:
57,750,000 would be written 5.775x10 7 , here you can easily see that
the number contains 4 significant figures
Example 2:
0.0000000000654 would be written 6.54x10 -11 , making it easy to see
that there are 3 significant figures present in the number
55