Chemistry 101 – Laboratory Experiments
Table of Contents
Chemistry 101: Introduction to General Chemistry
Experiment 1:
Introduction to the lab/Measurements…………………………………………2
Experiment 2:
Spectroscopy: Quantitative Analysis with Light……………………………..13
Experiment 3:
Molecular Geometry………………………………………………………….19
Experiment 4:
Inorganic Ions – Onondaga Lake water……………………………………...30
Experiment 5:
Gases: Boyle’s and Charles’ Laws…………………………………………...37
Experiment 6:
Solutions……………………………………………………………………...43
Experiment 7:
Freezing Point Depression Measurements…………………………………...48
Experiment 8:
Factors Affecting Reaction Rate……………………………………………..53
Experiment 9:
Acid/Base Properties of Common Household Products……………………...62
1
Chemistry 101 – Laboratory Experiments
Experiment #1 – Introduction to the Laboratory; Measurements:
length and density
Introduction
The General Chemistry curriculum has two specific goals for students: 1) developing an
understanding of the basic principles that govern chemistry and specifically the behavior of
molecules, and 2) learning how to design and conduct experiments directed at uncovering and
understanding basic principles.
You will be exposed to chemical principles in the text, lecture and laboratory. However, due to
the nature of chemistry and scientific discovery, an understanding of the basic principles is not
enough. In the laboratory part of this course, you will develop laboratory skills, learn about
chemical concepts, and develop computer skills necessary to help you design and carry out
experiments to investigate chemical phenomena.
What to Bring the First Day of Lab
The following items should be brought with you to lab the first day:
-Laboratory Packet
-Eye Protection
-Pen or Pencil
-Calculator
It is also suggested that you bring a notebook of some sort in which you can keep a record of
what you do in the laboratory, as well as record your locker combination.
Check in/out
During the first day of lab you will be assigned a lab drawer. This drawer contains all of the
necessary equipment needed to perform the experiments in this lab manual. Upon check in, you
will be sure all of your equipment is present and in good working condition. You will be
responsible for all of this equipment over the course of the class, be sure to take care of your
glassware and lock your drawer after every lab experiment. At the end of the quarter you will be
checked out of your lab drawer, if there are any items missing you will be responsible for
replacing them at cost.
We will also be using the MicroLAB data interface system with accompanying laptop
computers. These will be assigned to each group. You are responsible for these pieces of
equipment. They are very expensive and brand new, therefore they must last a long time. Take
2
Chemistry 101 – Laboratory Experiments
care of them, do not spill reagents on them, do not work close enough to them for a spill to
occur, if a spill happens clean it up as soon as possible and alert your lab instructor.
If You Miss a Lab
If you find that you must be away from campus or that some other conflict prevents you from
attending a regular lab period, you may attend another section. You must ask permission of the
other laboratory instructor to work in his/her lab, and he/she must sign the data sheet which will
be turned in to your regular lab instructor. Don’t simply “not show up” to lab; let your lab
instructor know so other arrangements can be made.
If you are ill, your lab instructor may excuse you from the lab, provided that you have cleared
your absence with your Dean or the Student Health Service. Notify your lab instructor as soon as
possible.
Experiments must be made up during the week they are offered. If your lab section meets late in
the week, you may want to make up an anticipated absence in advance. It is impossible to make
up an experiment after the labs have been stocked for a new experiment. You must complete
every experiment to pass this course.
Safety in the Laboratory
A chemistry laboratory is a safe place to work if you follow the appropriate safety precautions.
The most important piece of safety equipment is eye protection. Eye protection is required in all
labs unless designated otherwise. Prescription glasses, sunglasses, or contacts are not acceptable.
Safety goggles can be purchased from the Eastern Oregon Chemistry Club. Laboratory
instructors are required to remove students without eye protection from the lab. Be sure you
purchase a comfortable pair of goggles as you will be wearing them a lot. If you need a break
from your goggles you may step out of the lab for a few minutes. Students will not be asked to
wear their goggles, they will simply be asked to leave if they are not being worn which will
result in a “0” for the experiment.
Please observe the following rules when in a chemistry laboratory:
1. Do not eat, drink, chew, or smoke in the lab. After being in the lab, wash your hands
thoroughly before doing any of these things.
2. Shorts, open shoes, or bare feet are not allowed in the labs. Lab aprons or coats are
recommended. Long hair should be tied back or otherwise restrained.
3. Horseplay of any kind is not permitted. The laboratory is a work area and can be dangerous.
4. Clean up all spills immediately and correctly. When in doubt ask your lab instructor.
5. Dispose of all waste in the appropriate containers. We do not put chemicals down the sink!!
3
Chemistry 101 – Laboratory Experiments
6. Read labels on all reagents used. Some chemicals are toxic, some are corrosive, some are
flammable. Handle them appropriately and carefully.
7. Label all containers you use at your workstation.
8. Know where the lab safety features are (i.e. safety shower, eyewash stations, and fire
extinguisher).
9. Flush acids/bases from your skin with lots of water.
10. Follow any additional special safety precautions when necessary.
11. Clean and wipe down (with a damp paper towel) your work station, a clean lab and
workspace is necessary to obtain good results.
Use the following map (where bold outline indicates a table or bench and dashed outline
indicates a fume hood) to locate the following features of the laboratory (you only need to locate
one of each…):
1. Safety shower, 2. Eye wash, 3. Ice maker, 4. Fire extinguisher, 5. Drop box for labs and
homework, 6. Broken glass receptacle, 7. Emergency gas shut off, 8. Ring stands/Bunsen
burners, 9. Deionized water, 10. Iron rings.
4
Chemistry 101 – Laboratory Experiments
Lab Reports
Lab reports will be due immediately following completion of the experiment. These reports will
include data collected, calculations, and answers to post lab questions. They will be turned into
the appropriate section in the lab report box outside of the lab. These will be collected after the
completion of each lab so be sure you remember to turn them in. *IN ORDER TO GET FULL
CREDIT ON YOUR LAB REPORT, IT MUST BE SIGNED BY THE LAB
INSTRUCTOR/ TA PRIOR TO LEAVING THE LAB.
Lab Grades
See syllabus for how lab grades factor into your overall grade for the class.
Introduction to the Computer Interface
A program called Microsoft Windows is the computer's secretary; it keeps track of where
information and data are located on your disks and allows you to retrieve it when you want it.
Although you will not directly encounter the inner details of these systems, they are very
important. They allow the keyboard, monitor, and disk drives to operate together as a system.
They also control the use and execution of the MicroLAB interface program. The MicroLAB
interface program uses these operating systems to manage the data it obtains.
Most of the computer commands that you'll need to know are selected through the MicroLAB
interface software. You will be learning about these commands throughout the course. You will
only be working directly with the computer’s operating system to copy files, copy a disk, or to
see what files are on your disks. Files are simply blocks of information with names, just like a
file folder contains certain papers. To accomplish these tasks, you enter commands
(instructions) through the keyboard which tell the operating system what you want to do.
You will need to know how to make a copy of a disk or to copy files. If you do not already
know, your laboratory instructor will help you learn. The actual sequence of commands will
vary depending upon the exact configuration of your computer. Ask your laboratory instructor if
there is something you don't understand about using operating systems installed on the
computers in your lab.
Using the Computer to Organize and Analyze Data
The computer is an ideal tool for quickly organizing and analyzing data. It accomplishes in a
few seconds what takes much longer to do by hand. Computers are ubiquitous in chemistry (and
all of the sciences for that matter…) laboratories in the 21st century, therefore a certain level of
computer literacy will be obtained from this lab.
5
Chemistry 101 – Laboratory Experiments
Graphing and Spreadsheet Exercises
A Spreadsheet is a two dimensional array of cells which allows calculations, graphing, and
analyzing data obtained from your experiments. You can manually input data into the
Spreadsheet for analysis or import data from a previously generated data file, such as would be
acquired by collecting the data from an experiment into the computer. Manipulation of such data
by graphing and mathematical analysis is integral to all sciences and can significantly improve
the understanding of the experiment.
The MicroLAB software package is a compilation of programs to accomplish a variety of tasks.
It is opened by clicking on the MicroLAB icon, which will open a second window with the
various program options, e.g., Colorimeter, Color Mixer, Half-life and Energy of Light
Experiments, as well as the options to choose pre-designed experiments from the template tabs.
Opening the MicroLAB program:
To create your own experiment  turn on the MicroLAB unit  Double click on the MicroLAB
icon on the desk top  click on MicroLAB experiment from the New template  enter your
name, your partners name and the name of the experiment  click ok.
To open a pre-designed experiment, click on the template tabs to open the type of experiment
you want (e.g. time and temperature, pH titrations, gas laws, etc.). Then choose the pre-designed
experiment you wish to use. For instance, if you want to use a pre-designed Boyle’s law
experiment, choose the gas laws template tab  and then choose Boyle’s law.
Organization of the MicroLAB Interface Software.
The Main Screen of the MicroLAB software consists of six Views, as seen in Figure 1. Moving
from top left clockwise they are: 1. Data Sources/Variables view, where data sources, formulas
and analysis results are listed; 2. Graphing view, 3. Digital Display view, where instantaneous
readouts of the sensors are seen; 4. Spreadsheet view, where each data sampling set is stored for
later reference; 5. Experiment Steps view, where the steps to run the experiment are listed; and
finally, 6. Experiment Control view, where the experiment is started, stopped, repeated and
which has two computer controllable switches activated by clicking on them.
6
Chemistry 101 – Laboratory Experiments
Figure 1. MicroLAB Main Screen, illustrating the six major
views.
7
Chemistry 101 – Laboratory Experiments
Measurements: length and density
Learning Objectives:
-To learn how to use a ruler to make measurements of length
-To learn the difference between metric and English units of length (cm vs. inches)
-To learn how to graph the relationship between cm and inches
-To learn how to use this graph to predict lengths
-To learn how to determine density
-To learn the difference between some common measuring tools of chemists
Materials
MicroLAB unit
Ruler
50 mL Graduated cylinder
Various beakers
Hot plate
Laboratory balance
Pieces of string of different lengths
Experimental Procedure
Part 1. Length measurements
Open the MicroLAB software by double clicking on the MicroLAB icon. Select the “hand enter”
tab, name the experiment with you and your partners names and add “length,” click “ok.”
Obtain 4 different lengths of string. Measure the string in both cm and inches with your ruler. As
you take measurements, add your data to both the table in the data sheet and the spreadsheet in
MicroLAB (you will be prompted to enter labels and units) put inches in the column A (x-axis)
and cm in the column B (y-axis). Next, grab and drag (with the mouse) the inches column to the
x-axis in the graph view, repeat for cm on the y-axis. Choose “accept data” in the lower right
corner, your data points should now be visible in the graph view.
8
Chemistry 101 – Laboratory Experiments
Once you have made your measurements and entered all of your data, press “analysis” and then
“curve fit,” (you will be prompted to ok your graph series, make sure the right data corresponds
to the right axis on the graph) then choose “first-order (linear)” this will generate a straight line
and an equation that represents the relationship between inches and centimeters. This equation
can now be used to make predictions. Label your graph, print and then click “accept and save
this curve fit,” (the equation will now show up in the data sources/variables view under analysis.
Right click on the icon under analysis in the data sources/variables view, choose
“interpolation/extrapolation,” a predict box will appear, using this box determine how many cm
is equal to 1.8 inches, how many cm is equal to 12 inches, and how many cm is equal to 36
inches. Record your answers in the data sheet.
Part 2. Density measurements
In this part of the experiment, you will be determining the density of water using 3 different
means of measuring volume (Erlenmeyer flask, beaker, and graduated cylinder). The tools used
by chemists have very specific purposes, they also have varying degrees of accuracy. You will
determine which of the above measuring devices is most accurate (based on the known density
of water). You will then use the data from that measuring device to predict how much a certain
volume of water weighs and vice versa.
Obtain a clean and dry 50 mL graduated cylinder, 125 mL Erlenmeyer flask and a 100 mL
beaker. Record the mass of each in the table provided in the data sheet. Pour 50 mL of water into
each container as accurately as you can. Re-weigh the measuring devices with water in them,
enter these values into the table provided. Density is equal to mass (in grams) divided by volume
(in mL), calculate the density of water from each measuring device. Which is closest to 1 g/mL?
Using the measuring device that you found to be most accurate, weigh the mass of the following
volumes of water: 10 mL, 20 mL, 30 mL, 40 mL, and 50 mL. Enter your data in the table in the
data sheet.
Open MicroLAB again and choose hand enter (this time the name of your experiment should
include “density”). Click the “add column labels” button in the lower right window. Use your
mouse to select Column A, and label it “volume (mL).” Do the same for column B, labeling it
“mass (g).” Using your keyboard and mouse, enter the data you just acquired.
Next, set up your graph by dragging the “mass(g)” column label to the y-axis of the graph, and
the “volume (mL)” column label to the x-axis.
When all of your data has been entered into the spread sheet, press the “Accept Data” button in
the lower right hand window. The data you entered is now sent to the graph and plotted.
Click on the “Analysis” button in the graph view, then choose “add curve fit,” then add “first
order linear curve”. Record the equation provided from the analysis of your data. Printyou’re
your graph, then choose “accept and save this curve fit.”
Right-click on the density line under analysis in the data sources/variables view, choose
9
Chemistry 101 – Laboratory Experiments
interpolation/extrapolation. Using this function predict what the mass of 100 mL of water should
be. Predict the mass of 374 mL. Predict the mass of 0.5 mL. Record these predictions in the data
sheet.
10
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #1 Measurements: length and density
Name ______________________________ Section ______________ Date _______________
Partner______________________________
Instructor signature___________________
Part 1. Length measurements:
String #
1
2
3
4
Centimeters
Inches
1.8 inches = _______________________cm
12 inches = ________________________cm
36 inches = ________________________cm
Using the equation above and your calculator, determine how many inches are equivalent to
1000 cm (or 1 meter).
Part 2. Density measurements:
Type of
glassware
Mass of
glassware (g)
Mass of
glassware +
water (g)
Mass of water
(g)
Density
Erlenmeyer
Flask
Beaker
Graduated
Cylinder
Which piece of glassware gave a density closest to 1?
Which piece of glassware is the most accurate?
11
Chemistry 101 – Laboratory Experiments
Glassware
Mass of 10
mL (g)
Mass of 20
mL (g)
Mass of 30
mL (g)
Mass of 40
mL (g)
Mass of 50
mL (g)
What is the mass of 100 mL of water?
What is the mass of 374 mL of water?
What is the mass of 0.5 mL of water?
How would you dispose of acids and bases?
How would you dispose of broken glass?
Briefly describe what you would do if your partners sleeve caught on fire.
12
Chemistry 101 – Laboratory Experiments
Experiment #2 – Spectroscopy: the Quantitative Analysis with Light
Learning Objectives
The objectives of this experiment are:
-To observe and record a full spectrum of visible light
-To observe and record three atomic emission spectra
-To observe and record flame emission of at least four different metals
-To observe and record a molecular absorption spectrum using a spectrometer
-To predict and determine the absorption spectrum of food coloring using a spectrophotometer
-To determine the concentration of an unknown using a standard curve (aka calibration curve)
Introduction
Review sections 2.1 – 2.3 in your text prior to coming to lab. We have all observed the rainbow
effect produced when white light (i.e. sunlight) interacts with water droplets in the atmosphere.
As we know, white light from the sun is a mixture of many wavelengths of light (most of which
are invisible to the human eye). The part of these wavelengths that are visible to humans is a
relatively narrow range, from 400 – 700 nm, and is appropriately called the visible spectrum. We
will observe the spectrum of white light from the sun or other source, as well as spectra from gas
discharge tubes and thermally activated metals.
Many people believe that colors are emitted by colored substances. In fact, the opposite is true.
White light is a combination of all of the colors of the visible spectrum (as mentioned above).
The paper on this page appears white because it reflects all visible colors and absorbs none.
Likewise, the ink on this paper appears black because it absorbs all of the visible colors and
reflects none. We see colors because materials all around us absorb some of the colors that make
up white light, but where do these absorbed colors go? We know that light is made up of packets
of energy or photons, when these photons interact with electrons they “excite” them to a different
energy level. If the energies of the photon and the difference in electron energy levels are the
same, the energy in the photon will be transferred or absorbed by the electron. Light is a very
powerful tool which is commonly utilized by chemists both for qualitative and quantitative
measurements. Molecular absorption will be studied in this lab using a spectroscope and
spectrometer. We will also use the spectrometer to generate a standard or calibration curve and
use this curve to determine the concentration of an unknown colored solution.
Materials
Colored pencils
White light source
Power supply and gas discharge tubes
13
Chemistry 101 – Laboratory Experiments
EnVision Optical Spectroscope
Bunsen burner
Wooden splints
Various metal salts
Spectroscope and solution
MicroLAB
Food coloring solutions
Experimental Procedure
Part 1: Observation of white light and gas discharge tube spectra
Looking through your spectroscope, observe the reflection of white light off the chrome in the
laboratory. Select one spectrum to draw. Record your observations using the colored pencils
provided. Draw the spectrum, matching the color of light with the appropriate wavelengths, in
your data sheet (refer to the drawing of the visible spectrum hanging in the laboratory for
assistance).
Observe 3 different gas discharge tubes with and without the spectroscope (DO NOT TOUCH
THE TUBE OR POWER SOURCE WHILE THE CURRENT IS ON. ONLY YOUR
INSTRUCTOR MAY CHANGE THE TUBES FOR YOU). Draw the spectra for each discharge
tube in your data sheet, again matching the color to the approximate wavelength, in your data
sheet.
Part 2: Observation of flame emission spectra
Safety goggles required for part 2, this part works best with the lights off. Light a Bunsen burner
and set it to a relatively low but transparent (light blue) flame. Remove a splint and dip it in one
of the metal salts. Hold the splint in the flame. Extinguish the splint in a beaker of water. Record
the metal salt name and the color that you see. Use a fresh splint and repeat this procedure for all
of the remaining metal salts, recording your observations in the data sheet.
Part 3: Determining molecular absorption spectra using a spectroscope
Arrange a tube of colored solution so it is clamped to the scope over the light slit at the broad end
of the spectroscope. Place the white light (table lamp) close to the tube and switch it on to
illuminate the tube. Check with your instructor to be sure your set up is correct.
Look into the spectroscope and look for the spectrum to the right of the slit. You should also see
a scale to the right of the slit.
Record by drawing, the colors of the spectrum observed. Record the wavelength (nm) at various
points in your sketch. Observe the spectrum of just the white light, note the color that is absorbed
(i.e. missing) from the full spectrum of white light.
14
Chemistry 101 – Laboratory Experiments
Part 4: Predicting and determining molecular absorption spectra
At your work station you will find a set of solutions containing food coloring. The only
difference between each solution is the amount of food coloring. Examine one of the solutions.
Record the color that you see and predict the color that it will most strongly absorb. Remember,
the color you observe is complimentary to the color absorbed.
Next you will use the spectrometer to determine the color (or wavelength) of light that is most
strongly absorbed by your food coloring solution (to see if it matches your prediction).
Open the MicroLAB software, in the “new” tab select spectrometry and open. Before you can
use the spectrometer you will need to generate a baseline. Place a cuvette containing DI water in
the sample slot, choose absorbance under the configure drop down menu (or click the absorbance
bar at the top of the screen), cap the cuvette and click read blank (this solution should not absorb
any wavelength of light; this process takes about 30 seconds). A box prompting you to enter
units will appear, choose concentration/Beers Law. Now place the cuvette containing the food
coloring in the sample slot, cover and choose “add” (you will be prompted to label the sample,
label it the color that it appears).
A spectrum profile will be generated which shows the degree of absorption of each individual
wavelength by the food color solution. Print this profile, record the absorbance at each
wavelength in your data sheet by clicking on the bar in the spectrum profile for each wavelength.
Identify the wavelength that is absorbed the most, indicate this wavelength on your printed
spectrum profile. You will use this wavelength for the measurements you will make in part 5
below.
Part 5: Preparing and using a standard curve
Using the same MicroLAB experiment, you will prepare a calibration curve for the food color
solutions and use this to determine the concentration of an unknown based on its absorbance.
First, highlight the row of your sample from part 4 and choose delete. Also, highlight the
wavelength of most absorbance in the spectrum profile (the bar should be outlined in black when
it is chosen). Next place solution A into the sample holder, cover and choose “add”. You will be
prompted to enter a sample name (“A”) and concentration (4e-5; the software recognizes this as
4 x 10-5 M). The concentration units are Molarity (moles/liter – M), but you don’t need to enter
this as part of your concentration.
Repeat this for solutions B (2e-5) and C (1e-5).
Now choose the “3-Curve” tab in the upper left of the screen and add a linear line fit. Print this
curve.
Next choose the “4-Read” tab and place your unknown sample in the cuvette. Choose “add”, you
will be prompted to add a name to your sample, name it “unknown” and click “ok”. The
absorbance and corresponding concentration of food coloring in your unknown will be added to
your data table once it has been analyzed. Record both of these values in your data sheet.
15
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #2 Spectroscopy: the Quantitative Analysis with
light
Name ______________________________ Section _______
Partner______________________________
Date ____________________
Instructor signature___________________
Part 1. Observation of white light and gas discharge tubes:
White light spectrum:
_______________________________________________
400
500
600
700 (wavelength – nm)
Gas discharge spectra:
Tube 1: Name __________ Color seen by naked eye __________
Spectrum observed by spectroscope:
_______________________________________________
400
500
600
700 (wavelength – nm)
Tube 2: Name __________ Color seen by naked eye __________
Spectrum observed by spectroscope:
_______________________________________________
400
500
600
700 (wavelength – nm)
Tube 3: Name __________ Color seen by naked eye __________
Spectrum observed by spectroscope:
_______________________________________________
400
500
600
700 (wavelength – nm)
16
Chemistry 101 – Laboratory Experiments
Part 2. Observation of flame emission spectra:
Salt 1: Name __________
Color _____________
Salt 2: Name __________
Color _____________
Salt 3: Name __________
Color _____________
Salt 4: Name __________
Color _____________
Salt 5: Name __________
Color _____________
How might this information be useful in determining the metal contaminants in drinking
water?
Where have you observed this type of light display before?
Part 3. Determining molecular absorption spectra using a spectroscope:
Draw the colors of the spectrum observed:
__________________________________________________________________
Wavelength (nm)
The solution color is _____________. The color of light absorbed is ______________.
Part 4. Predicting and determining molecular absorption spectra:
Food color solution:
Observed color ___________ Predicted absorbed color __________
Predicted wavelength of color absorbed _______________
17
Chemistry 101 – Laboratory Experiments
Data table:

360
383
400
430
470
502
525
555
590
635
660
700
740
775
850
940
Abs.
Print and attach the graph you generated to this report.
Indicate in your sketch where the maximum absorbance occurred. How does this compare to
your prediction?
Part 5. Preparing and using a standard curve:
Attach a copy of your calibration curve (be sure the line equation is on the graph).
What is the absorbance of your unknown?
What is the concentration of your unknown in Molarity?
Based on what you learned about how light interacts with matter, why do you think the sky
appears blue?
18
Chemistry 101 – Laboratory Experiments
Experiment #3 – Molecular Geometry
Learning Objectives
The objectives of this experiment are:
-To learn to draw Lewis structure representations of the bonding and valence electrons in
molecules.
-To use the VSEPR model to predict the molecular geometries of molecules.
-To determine whether a molecule is polar or not based on its geometry.
Introduction
Review sections 3.4 – 3.5 in your text prior to coming to lab (it might be good ideas to bring
your text with you for this lab, also, don’t forget your model kit). The physical and chemical
properties of molecules are a direct consequence of the way in which the atoms are bonded
together and the three dimensional geometry or shape of the molecule. For example, water is a
bent molecule, because of this shape it is also a polar molecule. The polarity of a molecule in
turn influences physical properties such as freezing point, boiling point, vapor pressure, and
chemical reactivity. An understanding of molecular geometry is central in determining chemical
and physical properties.
The first step in determining molecular geometry is to write a Lewis structure for the molecule.
This structure gives the arrangements of valence electrons in terms of bonding pairs and lone
pairs of electrons. The Valence Shell Electron Pair Repulsion (VSEPR) model is then applied to
determine the overall geometry of the molecule. The basic theory of VSEPR is that electrons
repel each other and thus will try to minimize interactions with other electrons. In other words,
the electrons will try to get as far away from each other as is physically possible. The VSEPR
theory allows the prediction of molecular geometries and bond angles, which can be used to
predict other properties such as polarity.
Procedure for drawing Lewis structures
1. Add up the total # of valence electrons for all of the atoms in the molecule. Don’t forget to
account for charge. For example a (-) charged molecule will have one extra electron.
2. Draw the molecular skeleton and connect the atoms with single bonds. The central atom is
usually listed first in the molecular formula, otherwise it is the atom with the lowest
electronegativity, but never H.
19
Chemistry 101 – Laboratory Experiments
3. Satisfy the octet rule: distribute remaining valence electrons by adding lone pairs to complete
the octet of the outer atoms first (H can only have 2 electrons). Place any remaining electrons on
the central atom.
4. If there are too few valence electrons to give each atom an octet, multiple bonds may be
required. In this case, convert outer atom lone pairs to bond pairs to form multiple bonds.
5. There are some exceptions to the octet rule: 3rd period or heavier elements may have more
than an octet, Be, B and Al may have less than an octet. In a species with an odd number of
electrons, the least electronegative atom carries the odd electron.
Procedure for applying the VSEPR model
1. Draw the Lewis structure
2. Determine the number of electron pair domains on the central atom. This is equal to the
number of atoms plus the number of lone pairs attached to the central atom (electron domains are
lone pairs of electrons or bonding electrons, multiple bonds count as only one electron domain).
3. Determine the geometrical arrangement of the electron domains (see table in text).
4. Redraw the molecule in the correct electron domain geometry, place the lone pairs in their
proper positions in order to minimize repulsions.
5. Determine the molecular geometry or shape of the molecule.
Prediction of polarities of molecules from VSEPR and electronegativities
A polar covalent bond is formed when a pair of electrons is shared between two atoms with
different electronegativities. The atom with the greater electronegativity exerts a greater
attractive force on the electrons in the bond, thus the electron density is displaced more towards
the more electronegative atom. This induces a partial positive and partial negative region in the
bond, generating a dipole moment.
In molecules containing two or more bonds, the overall polarity depends on the net sum of the
individual bond polarities. In some cases the individual bond polarities will cancel each other out
leading to no net polarity. In the case of water, you have two of the same bonds (O-H), however
due to the bent shape the dipole moments do not cancel and thus water is polar. CO2 on the other
hand is linear, therefore the dipole moments of the individual C=O bonds cancel and the
molecule has no net polarity.
In general, a molecule will be nonpolar when there is a symmetrical arrangement of the same
type of bonds, such that the individual bond polarities cancel each other. The 5 basic electron
domain geometries from VSEPR (linear, trigonal planar, tetrahedral, trigonal bipyramidal, and
octahedral) plus the square planar geometry are symmetrical arrangements. Polar molecules
20
Chemistry 101 – Laboratory Experiments
result when there is an asymmetric arrangement of bonds, such that individual bond polarities do
not exactly cancel one another.
Experimental procedure
During this lab experiment you will determine Lewis structures, molecular geometries, and
predicted polarities for a number of molecular species beginning only with the molecular
formula. Also, you will need to create 3-D models of the molecule with correct bond angles
using chemical model kits.
You are to do the following tasks for each molecule in the list below and write the answers in the
space provided on the worksheet:
1. Determine the # of valence electrons in the molecule or ion.
2. Determine and draw the Lewis structure for the molecule or ion.
3. Determine the molecular geometry using VSEPR.
4. Construct a model of the molecule using your modeling kit.
5. Draw correct 3-D representations of the molecules from your model results.
6. Determine whether the molecule or ion is polar or nonpolar.
Molecular species list
1. H2 (Hydrogen), 2. CH4 (methane), 3. NH3 (ammonia), 4. H2O (water), 5. O2 (oxygen), 6. N2
(nitrogen), 7. CO2 (carbon dioxide), 8. H2O2 (hydrogen peroxide), 9. CO (carbon monoxide), 10.
BCl3 (boron trichloride), 11. BeF2 (beryllium fluoride), 12. O3 (ozone), 13. CCl4 (carbon
tetrachloride), 14. C2H2 (ethene).
21
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #3 Molecular Geometry
Name ______________________________ Section ______________ Date _______________
Partner______________________________
Instructor signature___________________
1. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
2. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
22
Chemistry 101 – Laboratory Experiments
3. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
4. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
23
Chemistry 101 – Laboratory Experiments
5. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
6. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
24
Chemistry 101 – Laboratory Experiments
7. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
8. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
25
Chemistry 101 – Laboratory Experiments
9. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
10. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
26
Chemistry 101 – Laboratory Experiments
11. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
12. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
27
Chemistry 101 – Laboratory Experiments
13. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
14. Formula:
# of valence electrons = __________
Lewis structure:
Molecular geometry = __________
3-D Drawing:
Polar or Nonpolar?
28
Chemistry 101 – Laboratory Experiments
Post-lab questions:
1. Draw and name the shapes of the following molecules:
a) NF3
b) H2S
c) BBr3
d) SiH4
2. If water did not have two unshared electron pairs, what shape do you think it would assume?
3. Do all of the molecules in this laboratory obey the octet rule? Which do not?
29
Chemistry 101 – Laboratory Experiments
Experiment #4 – Inorganic Ions – Onondaga Lake Water
Learning Objectives
The objectives of this experiment are:
-To learn the difference between cations (+ ions) and anions (- ions).
-To learn how to carry out simple tests to determine for the presence of cations and anions.
-To use the tests to determine what ions are present in an unknown water sample.
Introduction
Review sections 3.1 – 3.3 in your text prior to coming to lab. Water is a finite resource and water
pollution is a very serious problem that continues to threaten ecosystems and human life. For
example, Onondaga Lake near Syracuse, New York is a perfect example. Due to industrial
activities, this lake has been severely contaminated with both positively-charged (cations) and
negatively-charged (anions) contaminants.
In this experiment, you will use a variety of chemical tests to determine the presence or absence
of a number of different cations and anions in solution. Anytime you have ions in solution, that
solution will conduct electricity. So, we can use the conductivity of a solution to test for the
presence of ions. We can also perform other tests to distinguish specific types of ions. For
example, certain metal ions emit visible light when they are heated in a flame, this is known as
the flame test. You will also be using several chemical reactions which produce precipitates (i.e.
solid products), or products with color. Because these reactions are unique for different ions, you
will be able to compile results from standards and use this information to determine what ions
are present in an Onondaga Lake sample and in tap water and deionized water samples.
Materials
Sodium chloride
Calcium chloride
Iron(III) chloride
Sodium oxalate
Potassium thiocyanate
Sodium sulfate
Sodium hydrogen phosphate
30
Chemistry 101 – Laboratory Experiments
Silver nitrate
Barium nitrate
24-well well plate
Stirring rods or toothpicks
Flame test wire
Conductivity tester
Bunsen burner and striker
Experimental Procedure
Part 1. Conductivity test
Place a small sample of Onondaga Lake water in a well plate and test its conductivity using the
conductivity tester that is provided. Record your observation.
Repeat with both tap water and deionized water. Record your observations.
Part 2. Flame test for metal ions
In separate, clean wells, place about 10 drops each of NaCl, CaCl2, FeCl3, Onondaga Lake water,
deionized water and tap water.
Obtain a flame test wire and Bunsen burner. Light the Bunsen burner (in the hood) and adjust the
flame so that it is small with a visible light blue inner cone (in a well adjusted Bunsen burner the
hottest part of the flame is just above the blue cone).
Clean the flame test wire by heating the end of it in the flame.
Dip the flame test wire into the CaCl2 solution and immediately put the wire into the flame. The
wire should be put down on the top part of the blue cone. Observe and record the first color you
see.
Clean the flame test wire by putting it in the flame again. Then repeat the step above this time
using NaCl. Compare the color of the sodium test with that of the calcium test in terms of color
and intensity. Record your observations.
Repeat the test using FeCl3, Onondaga Lake water, tap water, and deionized water. Record your
results.
31
Chemistry 101 – Laboratory Experiments
Part 3. Chemical test for cations
Obtain a clean, dry 24-well wellplate and add solutions according to the following scheme:
Na+
+
1
Sodium Oxalate 
Potassium Thiocyanate 
Ca2+
Fe3+
2
3
Lake
water
4
Tap
water
DI
water
5
6
A
B
C
D
Add a few drops (5-10) of 0.1 M NaCl solution to wells A1 and B1.
Add a few drops of 0.1 M CaCl2 solution to wells A2 and B2.
Add a few drops of 0.1 M FeCl3 solution to wells A3 and B3.
Add a few drops of Onondaga Lake water to A4 and B4.
Add a few drops of deionized water to A5 and B5.
Add a few drops of tap water to A6 and B6.
Add a few drops of sodium oxalate to each well in row A. Look carefully for any signs of
reaction (color change, formation of a precipitate). Record your results.
Add a few drops of potassium thiocyanate solution to each well in row B. Look carefully for any
sign of a reaction (a precipitate or color change). You will need a white background to observe
this test. Record your results.
32
Chemistry 101 – Laboratory Experiments
Based on your results, determine which positive ions (cations) are in the three water samples.
Dispose of used solutions in the appropriate waste container and clean and dry your wellplate for
the next set of tests.
Part 4. Chemical test for anions
Once you have cleaned and dried your wellplate, set it up for tests according to the following
scheme:
Silver Nitrate 
A
Barium Nitrate
B
Calcium Chloride 
C
-
SO42-
Cl
42-+
1
2
HPO423
Lake
water
4
Tap
water
DI
water
5
6
D
33
Chemistry 101 – Laboratory Experiments
Place the well plate on a black background. All compounds formed in this test are white.
Add a few (5-10) drops of 0.1 M sodium sulfate (Na2SO4) solution to wells A1, B1, and C1.
Add a few drops of 0.1 M NaCl solution to wells A2, B2, and C2.
Add a few drops of 0.1 M sodium hydrogen phosphate (Na2HPO4) solution to wells A3, B3, and
C3.
Add a few drops of Onondaga Lake water to A4, B4, and C4.
Add a few drops of tap water to A5, B5, and C5.
Add a few drops of deionized water to A6, B6, and C6.
Add a few drops of silver nitrate (AgNO3) solution to each well in row A. Look carefully for any
sign of reaction (color change or precipitate formation). Record your results.
Add a few drops of barium nitrate (Ba(NO3)2) solution to each well in row B. Look carefully for
any sign of reaction (color change or precipitate formation). Record your results.
Add a few drops of calcium chloride (CaCl2) solution to each well in row C. Look carefully for
any sign of reaction (color change or precipitate formation). Record your results.
Based on your results, determine which negative ions (anions) are present in the three water
samples.
Dispose of your solutions in the appropriate waste container.
34
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #4 Inorganic Ions – Onondaga Lake Water
Name ______________________________ Section ______________ Date _______________
Partner _____________________________
Instructor signature___________________
Part 1. Conductivity test:
Sample
Onondaga Lake water
Tap water
Deionized water
Positive or Negative
Part 2. Flame test:
Ion
Sodium (Na+)
Calcium (Ca+2)
Iron (Fe+3)
Onondaga Lake water
Local water
Deionized water
Flame color
Part 3. Test for cations:
Reagents
NaCl
(Na+)
CaCl2
(Ca2+)
FeCl3
(Fe3+)
Onondaga
Lake
water
Tap water
Deionized
water
Sodium
Oxalate
Potassium
Thiocyanate
Part 4. Test for anions:
Reagents
Na2SO4
(SO42-)
NaCl (Cl-)
Na2HPO4
(HPO42-)
Onondaga
Lake
water
Tap water
Deionized
water
Silver
Nitrate
Barium
35
Chemistry 101 – Laboratory Experiments
Nitrate
Calcium
Chloride
Results:
Sample
Onondaga Lake
water
Local water
Deionized water
Ions present
(evidence)
What do the results of testing for conductivity indicate about what is present in Onondaga Lake
water? What about tap and deionized water?
How well do you think flame tests for sodium and calcium will work when both ions are present
at the same time? Explain your answer.
Distillation seems to be a good way to purify water and remove ions. Why then is distillation
seldom used on a large scale for purification of municipal water supplies?
36
Chemistry 101 – Laboratory Experiments
Experiment #5 – Exploring the Gas Laws: Boyle’s and Charles’
Laws
Learning Objectives
The objectives of this experiment are:
-To determine the effect of pressure on the volume of a gas.
-To determine the effect of temperature on the volume of a gas.
-Graphically deduce the relationship between P and V and T and V.
Introduction
Review section 5.1 in your text prior to coming to lab. Of the three states of matter, gases are the
most affected by their environment. As we know, solids tend to maintain their shape and volume
regardless of temperature and pressure, while liquids will maintain their volume but their shape
will change to fit the container. Gases are unique in that their volume and shape depend greatly
on the container, pressure, and temperature. In this experiment we will explore the effects of
pressure and temperature on the volume of a gas.
Materials
Syringe pump
MicroLAB unit
Hot plate
Thermometer
Capillary tube with oil plug
400 mL beaker
Rubber band
37
Chemistry 101 – Laboratory Experiments
Experimental Procedure
You will use the laboratory computer interface to determine the effects of pressure and
temperature on the volume of a gas.
Part 1. Boyle’s Law: Effects of Pressure
In 1662, a British chemist named Robert Boyle discovered the relationship between pressure and
the volume occupied by a gas. He found that “at a constant temperature, a fixed amount of gas
occupies a volume inversely proportional to the pressure exerted on it.” This is called Boyle’s
law.
As the pressure on a gas increases, the volume of that gas decreases (at constant temperature).
Robert Boyle discovered that a plot of pressure vs. 1/volume, resulted in a linear relationship.
Mathematically, we can say that PV = a constant. Since PV equals a constant, we can easily
predict pressure and volume because initial pressure times initial volume is equal to final
pressure times final volume (PinitialVinitial = PfinalVfinal).
Set up an experiment to plot pressure and volume. Gas laws template tab  Boyle’s Law.
The x and y axis should already be set up appropriately. The pressure should read somewhere
between 600 and 700 torr, if not restart the computer and MicroLab unit without the syringe
connected.
Set your syringe to 30.0 mL, then connect the syringe hose to the pressure connection on the
front of the MicroLAB. The hose has been cut to have a volume of 4.0 mL, so if your syringe is
set at 30.0 mL, the actual volume of your gas sample is 34.0 mL.
Take pressure measurements at volume settings from 24 to 54 mL in 5 mL increments (20 to 50
on the syringe, but remember to add the 4 mL to the value you enter into MicroLAB).
Push start experiment, you will then be asked to enter a volume, this value will automatically be
converted to 1/V and added to the spreadsheet along with the pressure. After making all of your
measurements, press stop, a curve will be automatically generated. Choose “accept data.”
Choose analysis  add curve  ok (for data verification)  first order (linear) and print this
graph. Is pressure proportional to 1/V? What generalization can you draw concerning the
relationship between the volume of a gas and the pressure exerted on it? Print your graph and
attach to your data sheet.
38
Chemistry 101 – Laboratory Experiments
Part 2. Charles’ Law: Effects of Temperature
About 100 years after Boyle, a French physicist named Jacques Charles, determined the
relationship between the volume and temperature of a gas. Charles found that at constant
pressure, the volume occupied by a fixed amount of gas is directly proportional to the absolute
temperature of the gas. This is known as Charles’ law. In this case, since they are directly
proportional, as one goes up the other does as well. It also means that V/T is equal to a constant,
so Vinitial/Tinitial = Vfinal/Tfinal. Thus, we can use this relationship to make predictions.
The change in volume of a gas in response to a change in temperature will be measured using a
capillary tube, sealed at one end, with a small liquid plug about ¼ of the way from the open end.
The plug will trap a small sample of air between the plug and the sealed end of the capillary tube.
The constant weight of the plug plus atmospheric pressure will maintain a constant pressure on
the trapped air.
Attach the capillary tube (sealed side down) to your thermometer using small rubber bands.
Align the sealed end with the bulb of the thermometer.
Place the sealed end of the capillary tube into a 400 mL beaker about 2/3 full of ice and 1/3 full
of water, making sure the capillary tube plug is well below the water level (be sure the open end
of the capillary is above the ice-water surface and doesn’t get any liquid in it).
After the temperature has stabilized and is close to 0°C, record both the temperature and the
height of the plug (in cm) by holding the ruler and capillary tube together (not touching the
bottom of the beaker), as close to the wall of the beaker as possible. Record these values in the
data sheet.
Pour the ice water down the sink and repeat these measurements with room temperature water.
Then place the beaker on a hot plate and heat the water, stirring occasionally. Repeat the above
measurements when the temperature gets around 20, 40, 60, 80, and 100°C (the temperature does
not have to be exact, as long as it is close to these values, be sure you record the temperature
accurately though). Do not remove the capillary tube from the hot water, and do not allow the
thermometer to rest on the bottom of the beaker. Be sure the plug in the capillary tube does not
come out of the capillary tube. Record the height and corresponding temperatures in the data
sheet.
Once you have recorded the plug height at the above temperatures, turn off your hot plate,
remove the capillary tube, thermometer and allow the water to cool. Prior to plotting your data
you must convert height to volume and temperature to K.
Use the space provided in the data sheet to calculate corrected height data. You will also need to
convert the temperatures from C  K (K = 273 + C). This time you will correct your height
39
Chemistry 101 – Laboratory Experiments
readings by converting them to volume. The following equation can be used to make this
conversion: V = hr2, where h is the height of the plug in cm, r = 0.09 cm, and  = 3.14.
Graph the data (using the “hand enter” microlab experiment), plotting the corrected volume on
the y-axis and the temperature on the x-axis. Once you have entered your data choose “accept
data” your data should now show in the graph. Next choose “analysis”  add curve  ok (to
accept the appropriate data)  linear curve. Record the correlation coefficient in the data
summary sheet. Print your graph and attach to the data sheet.
40
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #5 Exploring the Gas Laws
Name ______________________________ Section _______
Partner______________________________
Date ____________________
Instructor signature___________________
*For questions 1 and 2 use the following equations: PinitialVinitial = PfinalVfinal; Vinitial/Tinitial =
Vfinal/Tfinal.
1. A gas occupying a volume of 725 mL at a pressure of 0.970 atm is allowed to expand at
constant temperature until its pressure becomes 0.541 atm. What is its final volume?
2. In the Charles’s law portion of the experiment, Dick Xenon observed an air bubble with a
volume of 0.38 mL at 75.1°C. Using Charles’s law, what would you predict for the air bubble
volume at 0.0°C? *you need to convert temperature to K.
3. In your own words describe Boyle’s law.
4. In your own words describe Charles’s law.
41
Chemistry 101 – Laboratory Experiments
Boyle’s law:
Correlation coefficient (P vs. 1/V graph)______________
Print and attach your P vs. 1/V graph
Charles’s law:
Temperature in C
Height in cm
Volume in cm3 (yaxis)
Temperature in K
(x-axis)
Calculations:
Correlation coefficient (V vs. T)_____________
Print and attach your V vs. T graph
42
Chemistry 101 – Laboratory Experiments
Experiment #6 – Solution Preparation
Learning Objectives
The objectives of this lab are:
-To perform calculations relating to solution preparation.
-To prepare three solutions based on the calculations performed.
-To validate the concentration of your solution using density
-To learn how to convert molarity (M) to %(m/v) and vice versa
Introduction
Review sections6.1 – 6.3 in your text prior to coming to class. Knowing the concentration of a
solution is critical in any science lab. If we don’t know these quantities accurately then our
results will be worthless. There are a few different concentration units that are commonly used,
including: %(m/v), Molarity, ppm (parts per million or mg/L). In this lab you will prepare three
different salt solutions using two different concentration units. Your accuracy will be validated
by measuring the relative densities of the solutions you prepared. Commonly used concentration
equations are shown below.
1. % concentration (m/v) = grams of solute/mL of solution x 100%
2. Molarity (M) = moles solute/L of solution
3. M1V1 = M2V2
The first two equations are for preparing solutions from start, while equation #3 can be used to
prepare a solution from a concentrated stock solution (dilution equation).
Materials
100 mL volumetric flask
NaCl
Water
Food coloring
Dropper
Spatula
43
Chemistry 101 – Laboratory Experiments
Beakers
50 mL graduated cylinder
Balance
Thermometer or thermistor
Experimental Procedure
You will prepare three different salt solutions at the concentrations and volumes given below.
Write out in detail your calculations and how you went about preparing your solutions. There
should be enough detail for someone reading it to be able to prepare the solutions themselves.
A. %(m/v) solution: Prepare 40.0 mL of 0.80%(m/v) NaCl solution. Weigh the salt in a small
beaker using the balance (don’t forget to “tare” the balance with the beaker on it to remove the
beaker mass from your measurement). Use a 50 mL graduated cylinder to obtain 40 mL of DI
water, add this to the small beaker and mix thoroughly. Add a few drops of yellow food coloring
to your finished solution.
B. Molarity solution: Prepare 100.0 mL of 0.17 M NaCl solution. Weigh the salt out as above
using a small beaker. Use a 100 mL volumetric flask to obtain 100 mL of DI water, add this
water to the beaker and mix thoroughly. After you use this solution to prepare solution C
(below), add a few drops of green food coloring.
C. Dilution of stock solution: Prepare a dilution of the solution prepared in B. Make 30.0 mL of
0.090 M NaCl solution using a graduated cylinder. In order to prepare this solution you must
calculate the volume of solution B that you need to dilute to 30 mL. This can be calculated using
the dilution equation (M1V1 = M2V2). M1 = the concentration of solution B, V1 = unknown, M2 =
concentration of solution C, and V2 = volume of solution C. Transfer the solution to a small
beaker. Add a few drops of red food coloring to your finished solution.
44
Chemistry 101 – Laboratory Experiments
D. Experimentally determine the relative densities of these solutions. Use the glass part of the
dropper to pick up a portion of each solution. Do this by putting the dropper in the first solution,
place your forefinger over the end and lift up. Keeping your finger in place, dip into the second
solution and then remove your finger so that solution flows into the dropper. Lift up with your
finger back in place on the end of the dropper. Repeat for the next solution. If the three solutions
are arranged by density (most dense at the bottom, least dens at top) then three separate bands of
color will be seen. Record the observation of your successful trial. A drawing will suffice. Be
sure your instructor has seen your dropper. Now measure out 10.0 mL of each of your solutions
using a 10 mL volumetric flask. Weigh each separately and calculate their density (g/mL). After
you have completed the lab in the waste jugs provided.
45
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #6 Solution Preparation
Name ______________________________ Section ______________ Date _______________
Partner_____________________________
Instructor signature___________________
1. Calculate the relative densities of the solutions compared to water. Fill in the table below,
performing the necessary calculations for converting M  %(m/v) and vice versa. Show your
calculations in the space below.
Solution
Water
Solution A (yellow)
Solution B (green)
Solution C (red)
M NaCl
0.00
%(m/v) NaCl
0.80
0.17
0.090
Calculations:
2. Relative densities: List your solutions from least dense  to most dense. Do your
experimental results match your calculated results for relative densities?
3. Calculate the densities (in g/mL) of all three solutions. Which solution is most dense?
46
Chemistry 101 – Laboratory Experiments
Experiment #7 – Freezing Point Depression Measurements
Learning Objectives
The objectives of this experiment are:
-To investigate the effects of solutes on freezing point.
-To measure freezing points and freezing point depression.
-To use the relationship between solute concentration and freezing point depression to determine
the molar mass of anti-freeze.
Introduction
Review sections 6.1 – 6.4 in your text prior to coming to lab. Colligative properties of solutions
are those properties that depend only on the concentration of solute particles (not the nature of
the solute). Boiling point, freezing point and osmotic pressure are all examples of colligative
properties. For example, when you add salt (solute) to water, the boiling point will increase.
Likewise, if you were to freeze a kool-aid solution, which is essentially sugar and water, the
freezing point will decrease. In this experiment, you will prepare different solutions and observe
the change in freezing points. You will also determine the molecular mass of ethylene glycol
(aka anti-freeze) using the relationship between freezing point depression and solute
concentration.
Materials
Thermistor
Test tubes
Salt (NaCl)
Sugar (C12H22O11)
400 mL Beaker
Anti-freeze (ethylene glycol)
25 mL volumetric flask
Stirring rod
MicroLAB data interface and computer
47
Chemistry 101 – Laboratory Experiments
Experimental Procedure
Obtain a 400 mL beaker and fill ¾ full with ice, cover the ice ¼ to ½ inches with salt and stir
with your stirring rod.
Obtain three small beakers, label them “salt,” “sugar” and “anti-freeze.”
Prepare a salt (NaCl) solution by weighing out 1.45 grams of NaCl on a piece of weigh paper.
Transfer the solid to a 25 mL volumetric flask using a funnel, be careful not to lose any solid.
Fill the flask about ½ full with DI water. Cover with a cork and shake the flask to make sure all
of the NaCl dissolves. Fill to the line with DI water. Transfer this solution to your beaker labeled
“salt.” Rinse and clean your volumetric flask.
Calculate the molarity (moles of solute/liters of solution) and molality (moles of solute/kg of
solvent; water = 1g/mL or 1g/.001L) of this solution and record in the data sheet.
Prepare another solution, as above, of sucrose (C12H22O11) by weighing out 8.5 grams of sucrose
and adding that to a 25 mL volumetric flask. Pour the solution into your beaker labeled “sugar.”
Again calculate the molarity and molality of this solution and record in the data sheet.
Prepare a third solution, as above, this time using anti-freeze. Since anti-freeze is a liquid you
will need to tare your 25 mL volumetric flask and dispense the anti-freeze directly into the flask.
Measure out about 3.5 ml of anti-freeze into your volumetric flask. Fill with DI water to the
mark on the 25 mL volumetric flask and then pour into your beaker labeled “anti-freeze.”
Turn on your MicroLab unit and computer. Click the MicroLab icon and open a MicroLab
experiment under the new tab.
Choose data sources, then add sensor from the drop-down menu. Scroll through sensors until
you come to temperature (thermistor). Now choose which port the thermistor is connected to
by clicking on the port in the MicroLab picture. Click next. Now you must calibrate your
thermistor, choose read calibration from file, then open the file called Model 103 Thermistor.
Choose finish.
Drag and drop the thermistor icon, from the data sources/variables view, to the vertical axis (yaxis) on the graph, to column B and to the digital display. If your temperature is around 18-20 C
then you can proceed. If your temperature is highly variable or significantly different from 18-20
then repeat the above procedure.
Once your thermistor has been calibrated and is reading correctly, choose data sources, then
choose add a sensor. Scroll through the drop-down menu until you come to time. Again choose
which timer to use by clicking the timer on the MicroLab picture. Choose next. Set the timer
units to seconds, then click finish.
Drag and drop the time icon, from the data sources/variables view, to the horizontal axis (x-axis)
on the graph, to column A and to the digital display.
Stir the salt-ice mixture using your stirring rod. Now, using your thermistor, make sure the
temperature drops to at least -10 C.
48
Chemistry 101 – Laboratory Experiments
Obtain a test tube and fill it ¼ full with DI water. Place the test tube in the ice bath, making sure
the level of liquid in your test tube is below the level of ice in your ice bath. Dry your calibrated
thermistor carefully with a paper towel and place in the test tube of water. Start the experiment
and stir the water gently with the probe and monitor the temperature change on the computer
screen.
The temperature graph should show steady cooling. It should then level off when the first signs
of ice crystals appear on the side of the test tube, observe the temperature. This should be the
freezing point of the pure water. Continue to collect temperature data for another 10-15 seconds
after the temperature stabilizes. Stop the program, take the test tube out of the ice bath and
remove the temperature probe. Record the final temperature in the data sheet. Print this curve
and attach to the data sheet.
Repeat the above procedure with your salt, sugar and anti-freeze solutions. Record the freezing
points of these solutions in your data sheet. Print the curves obtained and attach to the data sheet.
Dispose of your solutions in the appropriate waste containers and turn off your MicroLab unit.
49
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #7 Freezing Point Depression Measurements
Name ______________________________ Section ______________ Date _______________
Partner_____________________________
Instructor signature___________________
Part 1. Freezing point of pure water:
Print and attach the freezing point graph for pure water, the salt solution and the sugar solution.
What are the M and m of your salt solution and sugar solution?
Salt M =_________
Sugar M = _________
Salt m = _________
Sugar m = __________
What is the freezing point of pure water? _______________C
What is the freezing point of the salt solution? _______________C
What is the freezing point of the sugar solution? ________________C
Using the equation below, to determine the theoretical freezing point of your sugar and salt
solutions. Kf for water = 1.86C/m.
(Kf)(m)
Salt solution = ____________
Sugar solution = _____________
50
Chemistry 101 – Laboratory Experiments
Part 2. Determination of the molar mass of anti-freeze
What is the freezing point of the anti-freeze solution?_______________C
Using the equation below, determine the molar mass of anti-freeze:
Molar mass = [(Kf)(grams of solute)]/[(T)(kg of solvent)]
If you placed a1 m solution of salt and a 1 m solution of sugar in the freezer at the same time,
which would reach its freezing point first? Why?
What is the point of putting anti-freeze in the radiator of your car?
51
Chemistry 101 – Laboratory Experiments
Experiment #8 – Factors Affecting Rates of Reaction
Learning Objectives
The objectives of this experiment are:
-To understand the effect of temperature on reaction rate.
-To understand the effect of a surface area on reaction rate.
-To understand the effect of concentration on reaction rate.
-To understand the effect of a catalyst on reaction rate.
Introduction
Review sections 7.1 – 7.3 in your text prior to coming to lab. This experiment is designed to
investigate the factors that influence the rate of a chemical reaction. Among these factors are:
temperature, surface area, concentration of reactants and the presence/absence of a catalyst. In
this experiment you will investigate each of these factors. Following your experiments, you will
qualitatively describe the relationship between the factors listed above and the reaction rate
based on your data and observations.
Materials
The following materials will be available to each group:
Well plate
100 mL graduated cylinder
Alka Seltzer tablets
HCl solution (1.0 M, 3.0 M, 6.0 M)
Pieces of zinc metal
Powdered zinc
Heat bath
Ice bath
Thermometer
52
Chemistry 101 – Laboratory Experiments
MicroLAB
*Additional materials may be made available upon request.
Part 1. The effect of temperature on reaction rate
Prior to carrying out these experiments, develop a hypothesis for the effects of each of these
factors on rates of chemical reactions. Record these hypotheses in your data sheet.
Obtain 3 test tubes. Fill each about ¼ full with water of different temperatures: add ice to one test
tube, leave one test tube with room temperature water and in the third test tube add water that has
been heated to about 70°C. It is important that you try to have approximately the same volume of
water in each test tube.
Record the temperature of the water in each test tube in your data sheet. Next add about ¼ of an
alka seltzer tablet to each test tube (be sure to record the mass of each tablet in the data sheet).
Record the time it takes for the tablet to completely dissolve in your test tubes.
Part 2. The effect of a surface area (or particle size) on reaction rate.
Obtain a piece of solid zinc metal, record the mass in your data sheet. Add the solid zinc to a
well in a well plate. Next obtain a sample of powdered zinc, make sure this sample is
approximately the same mass as your piece of solid zinc. Place the powdered zinc in a well in
your well plate.
Add 1 mL of 1 M HCl to the wells with the zinc samples (handle the HCl with care, this is a
strong acid). The chemical equation for this reaction is:
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Observe both wells, record your observations in the data sheet.
Part 3. The effect of concentration on reaction rate.
Place 1 mL of each of the three HCl solutions to separate well plates (be sure to keep track of
which solution is in each well; 1 mL = 20 drops). Handle HCl with care. Add one piece of zinc to
each well (record the mass of each in your data sheet).
Record the time you added the zinc to the wells and the time each reaction stops. Also, record
your observations of the wells.
53
Chemistry 101 – Laboratory Experiments
Part 4. The effect of a catalyst on reaction rate.
For this part of the lab you will determine which substance(s) act as a catalyst for the
decomposition of hydrogen peroxide:
2H2O2(l)  2H2O(l) + O2(g)
Dilute the hydrogen peroxide by adding 2 mL of 3% H2O2 to a 10 mL graduated cylinder. Add 8
mL of DI water to obtain 10 mL of diluted hydrogen peroxide (0.6%).
Place 1 mL of your diluted hydrogen peroxide into each of 6 different wells in your well plate.
Add 3 drops of each of the following solutions to separate wells (containing hydrogen peroxide):
0.1 M Fe(NO3)3, 0.1 M KNO3, 0.1 M NaCl, 0.1 M CaCl2, 0.1 M MnCl2
Mix each well with a toothpick (use separate tooth picks for each well). Observe each solution,
noting the production of gas bubbles that may form. Record each reaction as fast, slow, very
slow, or none in your data sheet.
54
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #8 Factors Affecting Rates of Reaction
Name ______________________________ Section _______
Partner______________________________
Date ____________________
Instructor signature___________________
Part 1. The effects of temperature
Water condition
Water Temperature
(C)
Time to Completion
Mass of tablet (g)
Cold
Room temperature
Warm
Hypothesis:
Write a statement to summarize your findings regarding the effect of temperature on reaction
rate.
How will changing the temperature of a reaction affect the rate?
55
Chemistry 101 – Laboratory Experiments
Part 2. The effects of surface area (or particle size)
Substance Tested and mass (g)
Solid Zn
Mass =
Observations
Powdered Zn
Mass =
Hypothesis:
Write a statement to summarize your findings regarding the effect of surface area on reaction
rate.
How will changing the surface area of a reaction affect the rate?
56
Chemistry 101 – Laboratory Experiments
Part 3. The effects of concentration
Acid
concentation
1 M HCl
Start time
Time at
completion
Mass of Zn
Observations
3 M HCl
6 M HCl
Hypothesis:
How will changing the concentration of a reaction affect the rate?
57
Chemistry 101 – Laboratory Experiments
Part 4. The effects of a catalyst
Fe(NO3)3
NaCl
CaCl2
KNO3
MnCl2
Reaction
Rate
Hypothesis:
Write a statement to summarize your findings regarding the effect of a catalyst on reaction rate.
How will adding/removing a catalyst affect the rate of reaction?
Based on your observations, which of your substances would you consider to be a catalyst for the
decomposition of H2O2?
Give two examples, from your own experience, where two different factors investigated today
occur or impact your life.
58
Chemistry 101 – Laboratory Experiments
In your own words, summarize the relationship between the four factors investigated and the
relative rate of a chemical reaction. Be sure to attach any additional graphs, data, etc. that was
used in reaching your conclusions.
59
Chemistry 101 – Laboratory Experiments
Experiment #9 – Acid/Base Properties of Common Household
Products
Learning Objectives
The objectives of this lab are:
-To learn about the pH scale for evaluating acids/bases
-To learn about the pH properties of common household products
-To determine the percentage of HCl in liquid toilet bowl cleaner
Introduction
Review sections 8.1 – 8.3 in your text prior to coming to lab. We all encounter acidic and basic
chemicals on a daily basis. In this lab, you will determine the pH of a number of different
common household and food products and determine what the pH means about the acidic/basic
character of these products.
Materials
Well plate
Samples of all household and food products (including liquid toilet bowl cleaner, save for part 2)
MicroLAB
pH probe
pH buffers (4, 7, and 10)
Experimental Procedure
Part 1. pH of household and food products
CAUTION!! DO NOT MIX ANY OF THE HOUSEHOLD PRODUCTS TOGETHER! VERY
TOXIC GASES CAN BE PRODUCED.
Obtain samples of all of the household and food products. For the liquid samples all you need is
about 1 mL (20 drops), for the solid samples weigh out about 0.500 grams. Place each sample in
a separate well in your well plate, record the location of each sample in the well plate in your
data sheet using the letter/number scheme on the well plate. Describe each in the table provided
60
Chemistry 101 – Laboratory Experiments
in the data sheet. Add about 1 mL of DI water to your solid samples and stir with your stirring
rod (rinsing stirring rod between samples to prevent cross contamination).
Obtain a pH probe, buffer solutions of pH = 4, 7, and 10 (place about 1 mL of each into separate
wells in your well plate. Connect your pH probe to the BNC outlet on the MicroLAB unit. Open
the MicroLAB software and choose “MicroLAB experiment” from the “new” tab, enter a name
for your experiment and then choose “ok.”
From the “Data Sources” drop-down menu, choose “add sensor.” Scroll down the sensor menu
and choose “pH/O.D.,” choose input on the MicroLAB and then choose “next to calibrate.”
Choose “perform new calibration,” then “add calibration point.” Place the pH probe into the well
that contains the buffer of pH = 4, you can rest the probe on the bottom of the well plate, just
don’t put too much pressure on it. Allow the probe to stabilize (about 15-20 seconds), then enter
the actual value and choose “ok.” Repeat with pH 7 and 10 buffers rinsing probe with DI water
between standards.
After you have 3 data points entered (and they are linear) choose “1st order linear,” then “accept
and save data” (units for the calibration should already be in pH, choose “ok”), name your file
and save. Choose “finish.”
From the “Data Sources” drop-down menu choose “add sensor.” Scroll down and add “timer,”
select the timer and choose “ok.” Make sure the units are in seconds, choose “finish.”
Drag and drop the timer, from the sensor screen, to the digital display, column A, and the x-axis.
Drag and drop the pH/O.D. sensor to the digital display, column B, and the y-axis.
Immerse the pH probe into one well at a time and “start” the experiment. Once the pH has
stabilized (15-20 seconds), stop the experiment and record the pH in the data sheet. Rinse off the
pH probe and repeat the experiment (without saving) for each of your products (be sure to rinse
the probe between each experiment). Dispose of your samples in the appropriate waste container.
61
Chemistry 101 – Laboratory Experiments
Data Sheet – Experiment #9 Acid/Base Properties of Common Household
Products
Name ______________________________ Section ______________ Date _______________
Partner_____________________________
Instructor signature___________________
Part 1: pH of household and food products:
Fill in the table below:
Product/Name
Brand
Sample
treatment
(diluted/strai
ght)
Color appearance
Active
Ingredient
and
concentration
(if given)
Function
Well
plate
location
pH
62
Chemistry 101 – Laboratory Experiments
1. List your samples from most acidic to most basic.
2. Which product is the most acidic?
3. Which product is the most basic?
4. What is the [H+] of your most acidic sample? What is the [H+] of your most basic
sample? (pH = -log[H+]).
63