1 Introduction to Chemistry Techniques Prelab (Week 1 ...

Introduction to Chemistry Techniques
Prelab (Week 1)
Name___________________________________________________ Total________/10
SHOW ALL WORK NO WORK = NO CREDIT
1. What is the purpose of this experiment?
2. Determine the number of significant figures in each of the following numbers.
a) 23.7 b) 1340.1 c) 0.000218 d) 33.50
3. a) Add the numbers from question 2 together and give the answer with the correct number of
significant figures.
b) Multiply the numbers from question 2 and give the answer with the correct number of
significant figures. Write the answer in scientific notation.
4. A cylindrical object has a density of 15.1 g/cm 3 . The diameter of this object is 1.10 cm and
the height is 2.05 cm. What is the mass of the object? (V = hπr 2 )
5. Define accuracy and precision.
6. You measure the mass of an object four times and obtain the following values of your
measurements: 14.2005 g, 14.2007 g, 14.2003 g and14.2004 g. Calculate the average of
these measurements and determine the standard deviation for these measurements. What
does the standard deviation tell you about the precision of your measurements?
1

Introduction to Chemistry Techniques
Prelab (Week 2)
Name___________________________________________________ Total________/10
SHOW ALL WORK NO WORK = NO CREDIT
1. The following graph shows absorbance versus concentration of a compound in solution.
From the information on the graph given to you:
Absorbance
1.200
1.000
0.800
0.600
0.400
0.200
Absorbance vs. Concentration for nickel(II) sulfate hexaydrate
y = 110.71x
0.000
0.00000 0.00100 0.00200 0.00300 0.00400 0.00500 0.00600 0.00700 0.00800 0.00900 0.01000
concentration (mole/L)
a) Determine the molar absorptivity (�) for this compound.
b) Determine the concentration of an unknown solution that has an absorbance value of 0.567.
Use the information given to you in the introduction.
2

Introduction to Chemistry Techniques
In this experiment you will measure the mass of an object with two different kinds of balances
and calculate the volume of liquid that is delivered from a transfer pipet. You will also
determine the density of an object. The last part will introduce you to spectroscopy and how to
use absorbance of light to determine the concentration of an unknown solution.
Special mention goes to Brian Stahl for developing Part IV of this laboratory experiment to be
used here in the Behrend chemistry curriculum.
Introduction
Many experiments require some type of measurement, and are often simple measurements of
mass and volume. The validity of an experiment is dependent on the reliability of these
measurements. A measurement’s reliability is usually considered in terms of its precision.
Precision is the closeness of the agreement between successive measurements of the same
quantity. Precision is not to be confused with accuracy, the agreement of a particular value with
the true value. The dispersion in a set of measurements is usually expressed in terms of the
standard deviation, whose symbol is s. You are going to be asked to determine the standard
deviation for some of your data. The standard deviation measures how close the data are
clustered around the mean. The smaller the standard deviation, the more closely the data are
clustered around the mean and the more precise your measurements are. After a quantity has
been measured in an experiment, it may be necessary to use that measurement in a subsequent
calculation. If you use a hand calculator for the arithmetic, eight or more digits may appear in
the answer. It is up to you to decide how many of these digits are significant. IT IS YOUR
RESPONSIBILITY TO KNOW THE RULES FOR SIGNIFICANT FIGURES BEFORE
YOU START THIS EXPERIMENT.
You will have to practice using an analytical and top loader balance and then a transfer pipet so
you can gain confidence to perform this experiment. Next, you will measure the mass of a flask
four times on each balance. The precision of the measurements will be examined when you
determine the correct number of significant figures in the mean mass of the flask. You will then
add water to the flask from a filled 5-mL pipet, and then measure the mass of the flask and water.
You will repeat this process three more times. After calculating the mass of the water that was
delivered each time from the pipet, you will calculate the volume of each addition from the mass
and density of water. You will then determine the correct number of significant figures in the
mean volume. This number will allow you to appreciate the precision that you have achieved
with the pipet.
You will have to measure volume using a graduated cylinder. A graduated cylinder is used to
measure an approximate volume of a liquid. When water or an aqueous solution (a solution
containing water) is added, the upper surface of the liquid in the graduated cylinder will be
concave. This concave surface is called the meniscus. The bottom of the meniscus is used for
all measurements. To avoid error, your eye should always be level with the meniscus when you
are measuring the volume. Using the graduated cylinder gives you the ability to measure the
volume to only one decimal place.
3

In the last part of this experiment, you will be using spectroscopy to determine the concentration
of an unknown solution. There are many ways to determine concentrations of a substance in
solution and using spectroscopy is only one method of determination. Many properties of a
solution can change with concentration. The property you will be looking at in this experiment
is colour. There is a direct correlation between colour intensity and concentration of a solution.
Concentration and absorbance are linearly related. This means that the higher the concentration,
the deeper the colour and therefore the higher the concentration of a substance in solution. The
opposite is also true. Using colour can be much faster than using wet methods such as titration,
especially when you have many samples containing different concentrations of the same
substance. When coloured solutions are irradiated with white light, they will selectively absorb
light of some wavelengths, but not of others. The wavelength at which absorbance is highest is
the wavelength to which the solution is most sensitive to concentration changes. This
wavelength is called the maximum wavelength or �max. In order to determine the precise amount
of a substance present in a solution, a calibration curve must be produced. A calibration curve
shows the relationship between the absorbance of light and the concentration of a chemical in a
solution. The higher the concentration of a substance in solution, the deeper the colour will be
and therefore the greater the absorbance of light. The opposite is also true. You will make four
solutions where the concentration of your substance is known. By plotting the absorbance
readings on the y-axis of the graph and the substance concentration values on the x-axis, you can
use this information to determine the concentration of an unknown solution of that same
substance. An example of a calibration plot is given. Once you have produced this calibration
plot, you will then use it to determine the concentration of your unknown solution. The
information you obtain from the calibration plot will be applied to Beer’s Law and the unknown
concentration of your solution will be determined. The unit for concentration being used in this
experiment is moles/litre.
Beer’s law is defined by the following equation:
A = �bc (1)
The equation represents three variables that influence the response of a solution to light. They
are the concentration (c) of the solution, the pathlength (b) of the light through the solution (also
called the cell length) and the molar absorptivity (�), the sensitivity of the absorbing species at
�max. The pathlength, unless it is stated differently, is usually fixed at 1.00 cm. The molar
absorptivity value is dependent on the solvent used (in this case water) and �. The units for
molar absorptivity are L/mole-cm for a concentration with units of mole/L.
You will make four solutions where the concentration of your substance is known. By plotting
the absorbance readings on the y-axis of the graph and the substance concentration values on the
x-axis, you can use this information to determine the concentration of an unknown solution of
that same substance. An example of a calibration plot is given. The slope of the line from the
plot is the molar absorptivity for your chemical in solution. You will then measure the
absorbance of a solution of unknown concentration. Using this absorbance value, along with the
molar absorptivity determined from the calibration plot, you will be able to determine the
concentration of your unknown solution by manipulating Beer’s Law:
4

A
c = (2)
εb
You will use pipets and volumetric flasks to make solutions of known concentrations as
accurately as possible and measure the absorbance at �max. You will be making dilutions from a
stock solution of a given substance. Since you will not be able to calculate dilutions yet, you will
use the following equation to calculate your concentrations:
C2 =
C1xV
V
2
where C1 is the concentration of your stock solution, V1 is the volume of the stock solution that
you are measuring and V2 is the final volume of the solution. For example, if you measure five
milliliters (mL) of your stock solution (V1) of known concentration (C1) and add it to a 100 mL
volumetric flask and then added water to the line which indicated 100 mL of solution (V2), you
could then find C2 using equation 3.
FIGURE 1: A Calibration Curve of Copper(II) Sulphate Pentahydrate in Solution
absorbance
1.4
1.2
1
0.8
0.6
0.4
0.2
1
copper(II) sulphate pentahydrate
0
0 0.05 0.1 0.15 0.2 0.25 0.3 0.35
concentration (moles/Litre)
5
(3)

Procedure
Part I: Using the Balance
1. Obtain about 100 mL of distilled water in a beaker. Allow the beaker and water to sit
on the laboratory bench while you are learning to use the balances and the pipet for parts II
and III. The water should come to the temperature of the laboratory during that time.
2. Obtain a thermometer and a 50-mL Erlenmeyer flask with a rubber stopper.
3. Use the SAME top loader balance throughout the experiment.
4. Place the rubber stopper in the Erlenmeyer flask. Tare the top loader balance. Measure and
record the combined masses of the flask and stopper.
5. Use tissue paper or paper towel to remove the stoppered flask from the pan of the balance.
This is used because some balances are sensitive enough to detect the oils from your
fingerprints and you will be weighing the flask on the analytical balance in step 9.
6. Bring your balance to the zero position again. Measure and record the mass of the
stoppered flask once again.
7. Repeat steps 5 and 6 until you have measured the mass four times.
8. Calculate the mean (average) mass. The differences between the measured masses and the
mean should be very small. If you are unsure of your results, consult your laboratory
instructor.
9. Repeat steps 3-8 with the analytical balance. Be sure to use the SAME balance throughout
the experiment.
Part II: Using the Pipet
Be sure to use the SAME BALANCE for ALL measurements.
1. Obtain a 5 mL pipet.
2. Practice using the pipet with distilled water (not the water you have set aside) until
you are comfortable with the technique. You should plan on using the same analytical
balance and the same pipet throughout the experiment.
3. Using the thermometer, note the temperature of the laboratory and of the distilled
water that you have set aside. When the temperatures are identical or very nearly
identical, you can begin. Record the temperature to the nearest degree.
4. Measure and record the mass of the empty stoppered flask again using the ANALYTICAL
BALANCE. Use tissue paper/paper towel as you did before.
6

5. Remove the flask from the balance, using tissue paper. Pipet exactly 5 mL of the
room-temperature water into the flask without touching the flask with your fingers.
Using tissue paper/paper towel, replace the stopper to prevent evaporation.
6. Bring your balance to the zero position. Measure and record the combined mass of the
water and the stoppered flask.
7. Remove the flask from the balance. Do not pour out the first sample. Pipet another
5 mL sample into the flask. The volume of the water in the flask should now be 10 mL.
Replace the stopper and repeat step 6.
8. Repeat until four samples of water have been delivered to the flask and the final
volume is 20 mL.
9. Calculate the mass of water that was delivered each time from your pipet. These
masses should be approximately identical.
10. Calculate the volume of each sample from the mass and density of water. Use Table 1
to find the density that corresponds to your recorded temperature. Calculate the average
volume delivered from each sample.
Part III: Determining the Density of an Object
1. Using your graduated cylinder, measure 40-50 mL of water from your beaker in Part I and
leave it in the graduated cylinder. Record this measurement to the nearest 0.1 mL.
2. Take one of the objects from the samples that are given and find its mass on the
analytical balance.
3. Place that same object in the graduated cylinder with the water in it. Make sure that
the object is completely submerged in the water. If the object is not completely
submerged, repeat steps 1and 2 being sure to dry the object and adding more water in
the graduated cylinder. Avoid splashing any water out of the graduated cylinder or cracking
the bottom of the graduated cylinder.
4. Record the level of the water with object in the graduated cylinder. The difference
between this measurement and the measurement in step 1 is the volume of the object.
5. Using the same object that has been dried well, measure the width and height of the object to
the nearest 0.01 cm.
6. Calculate the volume of the object using the measurements determined in step 5.
7. Calculate the density of the object (g/mL and g/cm 3 using the volumes from step 4 and step 6)
using the correct number of significant figures.
7

Table 1 Density (g/mL) of Water at Various Temperatures ( o C)
Temp. Density Temp. Density Temp. Density
17 0.9988 22 0.9978 27 0.9965
18 0.9986 23 0.9976 28 0.9962
19 0.9984 24 0.9973 29 0.9959
20 0.9982 25 0.9971 30 0.9956
21 0.9980 26 0.9968 31 0.9953
Example I: How to Calculate a Standard Deviation
You obtain the following measurements: 15.2654, 15.2657, 15.2658 and 15.2655. In the
following calculations, each individual measurement from above will be represented using the
symbol xi and the mean (the average of these values) will be given the symbol x .
2
d i
s �
N �1
�
� means “the sum of”
di = xi - x
N = the number of measurements
x = 15.2656
xi di (xi – x ) di 2
15.2654 -0.0002 0.00000004
15.2657 0.0001 0.00000001
15.2658 0.0002 0.00000004
15.2655 -0.0001 0.00000001
0.00000010 = � di 2
0.00000010
s � = 0.000183 = 0.0002
4 �1
When you report a value, it is the mean (average) value that is reported along with the standard
deviation to show the precision of the measurements which contributed to the mean value.
Therefore, the value reported for this example would be 15.2656 � 0.0002.
8

Part IV: Using Spectroscopy to Determine the Concentration of an Unknown Solution
1. Using a 250.0 mL beaker, obtain approximately 75.0 mL of stock solution of known
concentration (write this concentration down).
2. Using a pipet, transfer 5.0 mL of stock solution to a 100.0 mL volumetric flask and
fill up to the 100.0 mL mark with distilled water. Mix and then transfer all the solution to a
clean dry beaker that is 150 mL or greater. This will be called solution 2.
3. You will do this three more times making solution 3 (10.0 mL of stock solution diluted to
100.0 mL), solution 4 (15.0 mL of stock solution diluted to 100.0 mL), and solution 5 (20.0
mL of stock solution diluted to 100.0 mL). Remember, all dilutions are made with distilled
water and remember to place each solution in a beaker after mixing. Calculate the
concentrations (C2) of solutions 2-5 using equation 3 on page 5. You will need these values
when you make your calibration plot.
4. Once all your solutions are made and concentrations have been calculated, go to the netbook
and open up the Chem 111 folder. Select the file labeled “Introduction to Chemistry
Techniques Fall 2011” and double click on that file. You will then see a table and a graph on
the screen.
5. Take distilled water and pour it in the cuvette. Make sure there are no fingerprints and that
the light is passing through the smooth surfaces not the ribbed surfaces. Fill cuvette with
distilled water to make a blank.
6. On the menu bar, select Experiment, then Calibrate, and then Spectrometer:1. Wait for the
lamp to warm up.
7. When the lamp has warmed up, place the blank in the spectrometer. Click Finish
Calibration, then click OK. Keep the blank in the spectrometer since this will be your first
data point.
8. In the top right hand corner, you will see the collect icon . Click on collect.
You should see a very low absorbance value in the bottom left hand corner of the screen. If
you do not have a very low value, get help from your instructor. Once you have waited 20-30
seconds, go to the top right hand corner of the screen where it says keep and click.
You will now be asked for the concentration. Since this is your blank, the concentration is
zero.
9. Repeat step 8 for the rest of the solutions. Remember to wait 20-30 seconds before clicking
on keep. Copy all the absorbance values as well as the respective concentrations in your
laboratory notebook.
9

10. When you are all done getting the absorbances for solutions 2-5, click on the stop icon.
Then go to the linear fit icon and click. A box filled with information will appear and
you will find the value of the slope of the graph (m). This is your molar absorptivity value.
Record that value.
11. You will now take a sample of an unknown solution and place it in a clean cuvette. Place
the cuvette in the spectrometer and read the absorbance value for that solution from the
netbook. Now you can calculate the concentration of your unknown solution by using
equation 2.
Results
Part I
Using the Analytical Balance
Mass of the stoppered
flask (g) ____________ ____________ ____________ ____________
Calculation:
Using the Top Loader Balance
Mean mass (g) ____________
Mass of the stoppered
flask (g) ____________ ____________ ____________ ____________
Mean mass (g) ____________
10

Part II
Using the Pipet
Temperature ( o C) __________ Density of water (g/mL) __________
Addition No. 1 2 3 4
Mass after
addition (g) __________ __________ __________ __________
Mass before
addition (g) __________ __________ __________ __________
Mass of water
in flask (g) __________ __________ __________ __________
Volume of water
delivered each
time (mL) __________ __________ __________ __________
Mean volume (mL) __________
Calculations (One sample calculation for one addition in Part II):
Part III
Determining the Density of an Object
Mass of the object (g) __________
Volume after adding object (mL) __________
Volume before adding object (mL) __________ Volume of object (mL) __________
Diameter of object (cm) ____________ Volume of object (cm 3 ) __________
(using the equation V = h�r 2 )
Radius of object (cm) __________
Height of object (cm) __________
Density of the object (g/mL) __________
Density of the object (g/cm 3 ) ____________
Calculations:
11

Part IV
Solution Concentration Absorbance
1 0.000 0.00
2
3
4
5
Unknown ?
Unknown Letter or Number ________
Concentration of Stock Solution __________________
Molar Absorptivity of Chemical in Solution __________________
Concentration of Unknown Solution____________________
Calcuations:
Questions
1 a. Calculate the standard deviation (look at example on page 8) for the volume of water
delivered each time in Part II.
b. What does the standard deviation from part a tell you about the precision of your
measurements?
12