CHEM 1111 ACC Flame Tests and Atomic Emission Spectra Lab Report
Austin Community College CHEM 1111 – General Chemistry I Lab
Experiment 1
Measurements
LEARNING OBJECTIVES
•
•
•
•
•
•
•
•
Measure mass accurately using an analytical balance
Measure volume of a liquid using various types of glassware
Operate a pipet and a pipet bulb
Read a meniscus correctly
Measure temperature using a thermometer
Understand relationships between mass, volume and density
Calculate absolute error and percent error for a measured quantity
Apply rules for significant figures in calculations
INTRODUCTION
Every measurement has error associated with it. Error is defined to be the difference
between a measured (or calculated) value and the actual (or accepted) value. In this
experiment, we will examine the error associated with the use of some common chemistry
glassware: a beaker, a graduated cylinder, a volumetric flask and a volumetric pipet
(Figure 1-‐1).
Figure 1-‐1: Glassware to be examined in this experiment
We will consider two cases. First, we will answer the question: Which piece of glassware
more accurately contains a reported volume of liquid? Second, we will answer: Which
piece of glassware more accurately delivers a reported volume of liquid?
To answer the first question, we will compare a 150 mL beaker, a 100 mL graduated
cylinder and a 100 mL volumetric flask. As accurately as possible, we will measure 100 mL
of water using each piece of glassware. Because the density of water is well known, we can
1-‐1
Austin Community College
CHEM 1111 – General Chemistry I Lab
use the mass of water along with its density to calculate the volume of water contained in
each piece of glassware.
To answer the second question, we will compare a 10 mL graduated cylinder and a 10 mL
volumetric pipette. Here, we will measure 10 mL of water as accurately as possible, and
then transfer the water into a pre-‐weighed beaker. As before, the mass of water along with
its density will be used to calculate the volume of water delivered by each piece of
glassware.
PROCEDURE
Part 1: Which piece of glassware most accurately contains a stated volume?
Measure and record the mass of a 150 mL beaker, a 100 mL graduated cylinder and a 100 mL
volumetric flask. Make sure that all glassware is clean and dry before measuring the mass.
Be sure to record the entire mass given by the balance (i.e. record all of the digits displayed;
DO NOT ROUND, even if the last digit is a zero).
Fill each piece of glassware with 100 mL of water as accurately as possible. A transfer pipet
should be used to carefully adjust the volume of water in each piece of glassware. When
water is placed in a glass container the surface of the water forms a curved surface called a
meniscus. The volume of water should be determined by looking at the bottom of the
meniscus (Figure 1-‐2).
Figure 1-‐2: Reading a meniscus
Measure and record the mass of each piece of glassware containing water. Once again, bee
sure to record the entire mass displayed – DO NOT ROUND. Use a thermometer to measure
the temperature of the water in each container to the nearest degree.
Part 2: Which piece of glassware most accurately delivers a stated volume?
Measure and record the mass of a clean, dry 150 mL beaker. As accurately as possible,
measure 10 mL of water into a 10 mL graduated cylinder. [Note: It is not necessary to
measure the mass of the empty graduated cylinder.] Pour the water from the graduated
cylinder into the pre-‐weighed beaker. An aliquot is defined to be a small portion of a total
amount of liquid. The water transferred from the graduated cylinder to the pre-‐weighed
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Austin Community College
CHEM 1111 – General Chemistry I Lab
beaker represents your 1st aliquot. Reweigh the beaker containing 10 mL of water and
record the mass.
Measure another 10 mL of water into the graduated cylinder and pour the water into the
beaker; this water is your 2nd aliquot. Reweigh and record the mass of the beaker containing
20 mL of water.
Repeat with a third transfer of 10 mL of water to the beaker (3rd aliquot) and record the mass
of the beaker containing 30 mL of water. Use a thermometer to measure the temperature of
the water to the nearest degree.
Measure and record the mass of a clean, dry 150 mL beaker. As accurately as possible,
measure 10 mL of water using a 10 mL volumetric pipet. [Note: It is not necessary to
measure the mass of the empty pipet.] Your instructor will demonstrate the proper use of a
volumetric pipet and pipet bulb. When the bottom of the meniscus of water lies on the fill
line, carefully drain the pipet. The method by which a pipet is properly drained can depend
on the type of pipet bulb used. One method that works regardless of the kind of bulb used is
to remove the bulb, place the tip of the pipet against the glass of the beaker, and allow the
liquid to drain naturally by gravity. When properly drained, a volumetric pipet will retain
some liquid in the tip. Blowing this liquid out will deliver more water than expected and
should never be done. Deliver the water from the pipet to the beaker (this is your 1st
aliquot), and reweigh the beaker containing the 10 mL of water.
Measure another 10 mL of water into the volumetric pipet and transfer this water into the
beaker (this is the 2nd aliquot). Reweigh and record the mass of the beaker containing 20 mL
of water.
Repeat with a third transfer of 10 mL of water from the pipet to the beaker (3rd aliquot) and
record the mass of the beaker containing 30 mL of water. Use a thermometer to measure the
temperature of the water to the nearest degree.
When finished collecting data, dry all glassware and return all glassware and thermometers
to their original locations. Proceed to the data analysis portion of the experiment.
DATA ANALYSIS
Part 1: Which piece of glassware most accurately contains a stated volume?
Show your work for all calculations in the space provided on the data sheet. Use the back
of the page if you need more room.
Calculate the mass of water by subtracting the mass of the empty glassware from the mass
of the glassware containing water. Locate the density of water at the measured
temperature of water (refer to Table 1-‐1 below). Use the appropriate density of water to
calculate the volume of water contained in each piece of glassware.
The absolute error of a measurement is defined to be the magnitude of the difference
between the accepted value and the measured value:
𝐴𝑏𝑠𝑜𝑙𝑢𝑡𝑒 𝑒𝑟𝑟𝑜𝑟 = | 𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑣𝑎𝑙𝑢𝑒 − 𝑀𝑒𝑎𝑠𝑢𝑟𝑒𝑑 𝑣𝑎𝑙𝑢𝑒 |
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(1)
Austin Community College
CHEM 1111 – General Chemistry I Lab
The percent error in a measurement is the magnitude of the absolute error divided by the
accepted value, multiplied by 100:
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑒𝑟𝑟𝑜𝑟 =
!”#$%&'( !””#”
!””#$%#& !”#$%
× 100
(2)
The vertical bars in Equations (1) and (2) represent absolute value.
Calculate the absolute error and the percent error for each piece of glassware.
Table 1-‐1: Density of Water at Various Temperatures
Temperature (°C)
Density (g/mL)
16
0.9989
17
0.9988
18
0.9986
19
0.9984
20
0.9982
21
0.9980
22
0.9978
23
0.9975
24
0.9973
25
0.9970
26
0.9968
27
0.9965
28
0.9962
Rank the glassware according to accuracy. The most accurate measurement has the
smallest percent error, while the least accurate measurement has the largest percent error.
Part 2: Which piece of glassware most accurately delivers a stated volume?
As in Part 1, show work for all calculations in the space provided on the data sheet, or on
the back of the page.
Calculate the mass of the first aliquot of water by subtracting the mass of the empty beaker
(into which the water was transferred) from the mass of the beaker containing 10 mL of
water. Determine the mass of water in the second aliquot by subtracting the mass of the
beaker containing 10 mL of water from the mass of the beaker containing 20 mL of water.
Determine the mass of water in the third aliquot by subtracting the mass of the beaker
containing 20 mL of water from the beaker containing 30 mL of water. As before, use the
density of water at the temperature you measured to calculate the volume of water in each
aliquot.
Once the volume of each aliquot has been determined, calculate the average volume of an
aliquot delivered by that type of glassware. Compare the average aliquot volumes for each
measuring device to determine which one more accurately delivers 10 mL of water.
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Austin Community College
CHEM 1111 – General Chemistry I Lab
SAMPLE CALCULATIONS
A student weighs an empty 150 mL beaker and finds the mass to be 74.56 g. She then adds
100 mL of water and finds its mass to be 176.23 g. The temperature of the water is 23°C.
What is the calculated volume of water?
(1) Find the density of water from Table 1.
At 23°C, the density of water is 0.9975 g/mL
(2) Calculate the mass of water.
Mass of water = Mass of water plus beaker – Mass of empty beaker
Mass of water = 176.23 g – 74.56 g = 101.67 g
(3) Calculate the volume of water.
!
Use the equation 𝑉 = (volume = mass divided by density)
!
𝑉=
101.67 𝑔
= 101.9 𝑚𝐿
𝑔
0.9975 𝑚𝐿
WHY IS THIS LAB GREEN?
Only water was used in this experiment.
REFERENCE
Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Ed.). Boca Raton
(FL) : CRC Press
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CHEM 1111
Experiment 1
Name:
Section:
Score:
Experiment 1 Prelab Assignment
A student collects the following data when measuring the volume of a 100 mL graduated
cylinder:
Mass of empty graduated cylinder = 45.446 g
Mass of graduated cylinder + 100 mL water = 145.397 g
Temperature of water = 21.8°C
1. What is the density of the water? (Round the temperature to the nearest degree and
look up the density in Table 1.
2. (a) What is the mass of the water? Report your answer to the correct number of
significant figures.
(b) Justify the number of significant figures in your answer to Part (a).
3. (a) What is the volume of the water? Report your answer to the correct number of
significant figures.
(b) Justify the number of significant figures in your answer to Part (a).
4. Calculate the percent error in this measurement.
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CHEM 1111
Experiment 1
Name:
Section:
Score:
Experiment 1 Report Sheet
Part 1: Which piece of glassware most accurately contains a stated volume?
Measured Data
Glassware:
150 mL Beaker
100 mL
Graduated
Cylinder
100 mL
Volumetric Flask
A
Mass of empty glassware (grams)
B
Mass of glassware containing 100
mL water (grams)
C
Temperature of water (°C)
D
Density of water
Calculated Results
E
Mass of water (grams)
F
Volume of water (mL)
G
Accepted volume (mL)
H
Absolute error
I
Percent error
100.
100.0
100.0
Show work for all calculations in the space below and/or on the back of this page:
Rank the glassware from MOST accurate to LEAST accurate:
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CHEM 1111
Experiment 1
Name:
Section:
Part 2: Which piece of glassware most accurately delivers a stated volume?
Measured Data
Glassware used to Measure
10 mL Graduated Cylinder
10 mL Aliquots:
10 mL Volumetric Pipet
J
Mass of empty beaker (grams)
K
Mass of beaker containing Aliquot
1 (grams)
L
Mass of beaker containing Aliquot
1 and Aliquot 2 (grams)
M
Mass of beaker containing Aliquots
1, 2, and 3 (grams)
N
Temperature of water (°C)
O
Density of water
Calculated Results
P
Mass of Aliquot 1 (grams)
Q
Mass of Aliquot 2 (grams)
R
Mass of Aliquot 3 (grams)
S
Volume of Aliquot 1 (mL)
T
Volume of Aliquot 2 (mL)
U
Volume of Aliquot 3 (mL)
V
Average volume of each aliquot
(mL)
Show work for all calculations in the space below and/or on the back of this page:
Which piece of glassware more accurately delivered 10 mL of water?
1-‐8
Austin Community College
CHEM 1111 – General Chemistry I Lab
Experiment 2
The Density of Liquids and Solids
LEARNING OBJECTIVES
•
•
Reinforce skills from Experiment 1 (measurement of mass and temperature)
Understand relationships between mass, volume and density
INTRODUCTION
Part 1: Density of a liquid
The density (d) of a substance is the ratio of its mass (m) to its volume (V):
!
𝑑 =
!
(1)
Based on this definition, it is clear that both the mass and the volume must be known
precisely to obtain a good measure of the density of a substance.
A pycnometer is an instrument that is used to determine the density of a liquid with a high
degree of accuracy. It is usually made out of glass and is equipped with a stopper that has a
small hole through the center. This small hole allows air bubbles to escape and produces a
repeatable measure of volume. To determine the density of a liquid accurately, both the
mass and volume must be known accurately. In this experiment, a stoppered 50 mL
Erlenmeyer flask will be used as a simple pycnometer. By difference, it is possible to
determine precisely the mass of the liquid contained within the stoppered flask when it is
filled. Determination of the volume occupied by the liquid is somewhat less
straightforward. The full volume of the flask IS NOT 50 mL; the markings on an
Erlenmeyer flask do not provide a very accurate measure of volume.
Determining the volume of the flask:
Because the density of water at a particular temperature is known to a high degree of
accuracy, we will use water to determine the volume of the flask. Initially, the mass of the
empty dry stoppered flask will be measured. The flask will then be filled completely with
water, and the mass of the filled stoppered flask will be measured. The mass of water
contained in the flask may then easily be calculated by difference:
𝑀𝑎𝑠𝑠 𝐻! 𝑂 = 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑓𝑖𝑙𝑙𝑒𝑑 𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 − 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑒𝑚𝑝𝑡𝑦 𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘
(2)
The temperature of the water will be measured (to the nearest degree), and Table 2-‐1 will
be used to find the density of water at the measured temperature. The volume of water
contained within the filled stoppered flask can be determined by rearranging Equation (1),
using the mass of the water and the density at the measured temperature. We will assume
that the volume of water contained within the filled stoppered flask is equal to the
overall volume of the flask. This volume will be constant throughout the remainder of the
experiment.
2-‐1
Austin Community College
CHEM 1111 – General Chemistry I Lab
Table 2-‐1: Density of Water at Various Temperatures
Temperature (°C)
Density (g/mL)
16
0.9989
17
0.9988
18
0.9986
19
0.9984
20
0.9982
21
0.9980
22
0.9978
23
0.9975
24
0.9973
25
0.9970
26
0.9968
27
0.9965
28
0.9962
Determining the density of a liquid:
To calculate the density of an unknown liquid, we will fill the flask completely with the
liquid and determine the mass.
𝑀𝑎𝑠𝑠 𝑙𝑖𝑞𝑢𝑖𝑑 = 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑓𝑖𝑙𝑙𝑒𝑑 𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 − 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑒𝑚𝑝𝑡𝑦 𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 (3)
The volume of the liquid is equal to the overall volume of the flask. The density of the
liquid can then be calculated from Equation (1).
Part 2: Density of a solid
Because a liquid can completely fill a container of any shape, it is easy to determine
accurately a liquid’s volume. In this experiment, the volume of the liquid is simply equal to
the overall volume of the flask. Determining the volume of a solid is slightly more difficult
because the flask cannot be completely filled with the solid (air would be trapped in the
flask between the solid pieces). To determine the volume of an irregularly shaped solid,
the concept of displacement will be used. Any solid object, when placed in water, will
displace a volume of water that is equal to the volume of the solid. We will place a solid in
the flask, and then fill the remainder of the flask completely with water. The overall
volume of the flask (known) will be equal to the sum of the volume occupied by the solid
and the volume of the added water.
Determining the mass of the solid:
After partially filling the flask with a solid and recording the mass of the flask containing
the solid, the mass of the solid may be easily calculated:
𝑀𝑎𝑠𝑠 𝑠𝑜𝑙𝑖𝑑 = 𝑀𝑎𝑠𝑠 𝑜𝑓𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 𝑝𝑙𝑢𝑠 𝑠𝑜𝑙𝑖𝑑 − 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑒𝑚𝑝𝑡𝑦 𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 (4)
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Austin Community College
CHEM 1111 – General Chemistry I Lab
Determining the volume of the solid:
To the flask partially filled with a solid, water will be added until the flask is filled and no
air bubbles are present. Now, the volume of the flask is occupied with both solid and
water. In other words,
𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑓𝑙𝑎𝑠𝑘 = 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑖𝑑 + 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝐻! 𝑂 (5)
The volume of the water in the flask can be determined from the mass of water and its
density, using Equation (1). The mass of the water is calculated by difference:
𝑀𝑎𝑠𝑠 𝐻! 𝑂 = 𝑀𝑎𝑠𝑠 𝑜𝑓𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 + 𝑠𝑜𝑙𝑖𝑑 + 𝐻! 𝑂 − 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑠𝑡𝑜𝑝𝑝𝑒𝑟𝑒𝑑 𝑓𝑙𝑎𝑠𝑘 𝑝𝑙𝑢𝑠 𝑠𝑜𝑙𝑖𝑑
(6)
Once the volume of water is determined, the volume of the solid can be easily calculated by
rearranging Equation (5):
𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑖𝑑 = 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑓𝑙𝑎𝑠𝑘 − 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝐻! 𝑂 (7)
The density of the solid is calculated by substituting the mass of the solid and the volume of
the solid into Equation (1).
Part 3: Compare two methods of density determination
In this portion of the experiment, a 50 mL graduated cylinder will be used to measure the
density of both a liquid and a solid.
Determining the density of a liquid (Method 2)
The volume of a liquid can be measured directly using a graduated cylinder. The mass of
the liquid is calculated by subtracting the mass of the empty graduated cylinder from the
mass of the cylinder containing the unknown liquid.
𝑀𝑎𝑠𝑠 𝑙𝑖𝑞𝑢𝑖𝑑 = 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑔𝑟𝑎𝑑 𝑐𝑦𝑙𝑖𝑛𝑑𝑒𝑟 𝑤𝑖𝑡ℎ 𝑙𝑖𝑞𝑢𝑖𝑑 – 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑒𝑚𝑝𝑡𝑦 𝑔𝑟𝑎𝑑 𝑐𝑦𝑙𝑖𝑛𝑑𝑒𝑟 (8)
Once the mass of the liquid has been calculated, the density can be determined from
Equation (1).
Determining the density of a solid (Method 2)
The mass of a solid can be measured directly using an analytical balance. If the solid is then
added to a known volume of water in a graduated cylinder, the new volume reading after
displacement allows for the calculate of the volume of the solid alone:
𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑖𝑑 = 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑖𝑑 𝑎𝑛𝑑 𝑤𝑎𝑡𝑒𝑟 – 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑤𝑎𝑡𝑒𝑟
Equation (1) can be used to calculate the density of the solid after the volume has been
determined.
2-‐3
(9)
Austin Community College
CHEM 1111 – General Chemistry I Lab
PROCEDURE
Part 1: Density of a liquid
Determine the mass of a clean dry 50 mL Erlenmeyer flask and stopper. Fill the flask
completely with water. Use a thermometer to measure the temperature of the water to the
nearest degree. Place a stopper in the flask such that no air is trapped. If you see any air
bubbles, remove the stopper and add more water. Make sure that the outside of the flask is
dry, and record the mass of the filled, stoppered flask. Empty and dry the flask and stopper.
Your instructor should provide instructions as to how to rapidly dry the flask. It is
important to use the same flask and stopper throughout the experiment. Refill the flask
with an unknown liquid (be sure to record the unknown number) and stopper the flask as
before so that no air is trapped. Dry off the outside of the flask, and record the mass of the
stoppered flask containing the unknown liquid. Pour the unknown liquid back into its
original bottle. Rinse the 50 mL Erlenmeyer flask with water before drying the flask and
stopper according to your instructor’s directions.
Part 2: Density of a solid
Half-‐fill the clean, dry 50 mL Erlenmeyer flask with an unknown solid (be sure to record
the unknown number). Place the stopper in the flask and record the mass of the stoppered
flask containing the solid. Slowly add water to the flask containing the solid. As the water
is added be sure to swirl the flask around to eliminate any air bubbles trapped between the
solid pieces. Continue adding water until the flask is completely filled and there is no air
trapped when the flask is stoppered. Measure and record the mass of the stoppered flask
containing water and solid. Carefully pour the water out of the flask, and return the solid to
be dried as directed by your instructor. Do not mix wet solids with dry solids!
Part 3: Compare two methods of density determination
Determining the density of a liquid (Method 2)
Weigh and record the mass of a clean dry 50 mL graduated cylinder. Fill the graduated
cylinder as accurately as possible with 50 mL of the same unknown liquid that you used
earlier. Weigh and record the mass of the graduated cylinder plus the liquid. Be sure to
use the same unknown liquid used in the first part of this experiment. Return the
unknown liquid to the original bottle for reuse.
Determining the density of a solid (Method 2)
Place about 25 mL of water in a clean dry 50 mL graduated cylinder. Record the volume of
the water to one decimal place. Place a weigh boat on a balance and re-‐zero (or tare) the
balance. Add a small amount (try for around 10 -‐ 20 mL) of solid to the weigh boat and
record the mass. Carefully place the solid in the graduated cylinder containing water and
record the volume to which the water rises. Carefully pour the water out of the graduated
cylinder and return the solid to your instructor as directed. Clean and dry any glassware
and return all glassware and thermometers to their original locations.
2-‐4
Austin Community College
CHEM 1111 – General Chemistry I Lab
SAMPLE CALCULATIONS
A student weighs an empty stoppered Erlenmeyer flask and determines the mass to be
41.626 g. When the flask is filled completely with water at a temperature of 25°C, the mass
of the filled stoppered flask is 96.720 g. What is the overall volume of the flask?
(1) Find the density of water from Table 1.
At 25°C, the density of water is 0.9970 g/mL.
(2) Calculate the mass of water.
Mass of water = Mass of filled stoppered flask – Mass of empty stoppered flask
Mass of water = 96.720 g – 41.626 g = 55.094 g
(3) Calculate the volume of water which equals the overall volume of the flask.
𝑉 = 55.094 𝑔 ×
1 𝑚𝐿
= 55.26 𝑚𝐿
0.9970 𝑔
The same flask is then filled completely with a different unknown liquid. The mass of the
filled stoppered flask is 85.381 g. What is the density of the unknown liquid?
(4) Calculate the mass of the liquid.
Mass of liquid = Mass of filled stoppered flask – Mass of empty stoppered flask
Mass of liquid = 85.318 g – 41.626 g = 43.755 g
(5) Calculate the density of the liquid.
𝑚
43.755 𝑔
𝑔
=
= 0.7918 𝑚𝐿
𝑉
55.26 𝑚𝐿
The same flask is partially filled with an irregularly shaped solid. The mass of the
stoppered flask plus the solid is 200.16 g. Enough water is added to fill the flask
completely. The mass of the stoppered flask containing the solid and water is 235.14 g.
What is the density of the solid?
𝑑 =
(6) Calculate the mass of the solid.
Mass of solid = Mass of stoppered flask plus solid – Mass of empty stoppered flask
Mass of solid = 200.16 g – 41.626 g = 158.53 g
(7) Calculate the mass of the added water.
Mass of H2O = Mass of stoppered flask + solid + H2O – Mass of stoppered flask + solid
Mass of H2O = 235.14 g – 200.16 g = 34.98 g
(8) Calculate the volume of the added water.
𝑉 = 34.98 𝑔 ×
1 𝑚𝐿
= 35.09 𝑚𝐿
0.9970 𝑔
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Austin Community College
CHEM 1111 – General Chemistry I Lab
(9) Calculate the volume of the solid.
Volume of solid = Volume of flask – Volume of water
Volume of solid = 55.26 mL – 35.09 mL = 20.17 mL
(10) Calculate the density of the solid.
𝑑 =
𝑚
158.53 𝑔
𝑔
=
= 7.860 𝑚𝐿
𝑉
20.17 𝑚𝐿
WHY IS THIS LAB GREEN?
All the liquids used in this experiment are recycled through reuse in our labs. The solids are
reusable and only need to be dried between lab classes.
REFERENCE
Lide, D. R. (Ed.) (1990). CRC Handbook of Chemistry and Physics (70th Ed.). Boca Raton (FL)
: CRC Press
2-‐6
CHEM 1111
Experiment 2
Name:
Section:
Score:
Experiment 2 Prelab Assignment
1. A student finds that the mass of a stoppered 50 mL Erlenmeyer flask is 46.976 g. After
filling the flask with water and replacing the stopper, the mass is found to be 101.813 g.
The temperature of the water is 23°C. Calculate the volume of water contained in the
flask.
2. The same flask used in Question 1 is emptied, dried, and then filled with an unknown
liquid. The mass of the stoppered flask containing the liquid is found to be 105.184 g.
Determine the density of the unknown liquid.
3. Consider the situation described in Question 2, but suppose that an air bubble is
present in the flask that was not noticed before the stoppered flask containing the
liquid was weighed. Compared to the “true” value, will the calculated density be too
low or too high as a result of this error? Explain your reasoning.
4. In this experiment, when the density of the solid is determined, the volume of the solid
is found by subtracting the volume of water (in a flask filled with solid and water) from
the overall volume of the flask. What principle is being used here to determine the
volume of the solid?
2-‐7
CHEM 1111
Experiment 2
Name:
Section:
Score:
Experiment 2 Report Sheet
Calculated Data and Results: SHOW YOUR WORK for all
calculations!!! Be sure to include UNITS for all results.
Volume of the flask:
Collected Data: Be sure to include UNITS for all measurements.
A
Mass of empty flask and stopper
B
Mass of stoppered flask filled with
water
C
Temperature of water
D
Density of water (refer to Table 1)
E
Mass of stoppered flask filled with
unknown liquid
F
Mass of stoppered flask half-‐filled
with unknown solid
G
Mass of stoppered flask + solid +
water
Unknown Solid (Part 2)
Mass of water
I
Volume of water = volume of the flask
Density of unknown liquid:
J
Mass of unknown liquid
K
Density of unknown liquid
Density of unknown solid:
Record your unknown codes below!
Unknown Liquid (Part 1)
H
L
Mass of unknown solid
M
Mass of water added to unknown solid
N
Volume of water added to unknown
solid
O
Volume of unknown solid
P
Density of unknown solid
2-‐8
CHEM 1111
Experiment 2
Name:
Section:
Record your unknown codes below!
Collected Data: Be sure to include UNITS for all measurements.
Q
Mass of empty graduated cylinder
R
Mass of graduated cylinder
containing unknown liquid
S
Mass of solid used
T
Volume of water in graduated
cylinder
Volume of water and solid in
graduated cylinder
U
Unknown Liquid
Unknown Solid
Your instructor will provide you with the densities of the
unknown liquids and solids. Calculate your percent error
for the density of the liquid and the solid from each of the
two methods.
50 mL
Erlenmeyer
flask
50 mL
Graduated
cylinder
Calculated Data and Results: SHOW YOUR WORK for all
calculations!!! Be sure to include UNITS for all results.
Liquid (percent error)
V
Mass of the liquid
W
Density of the liquid
X
Volume of the solid
Y
Density of the solid
Which method did you find produced more accurate
results?
Do you feel that the 50 mL Erlenmeyer flask was effective
at approximating the use of a pycnometer?
Solid (percent error)
2-‐9
Austin Community College
CHEM 1111 – General Chemistry I Lab
Experiment 3
How Much Sugar is in a Can of Soda?
LEARNING OBJECTIVES
•
•
•
Determine the mass of sugar in a can of soda
Prepare solutions quantitatively using a volumetric flask
Prepare and use a calibration curve
INTRODUCTION
If you were to measure out identical volumes of regular soda and diet soda, you would find
that the two liquids have different masses. This difference in the mass of the two liquids is
best understood by looking at the mass per unit volume (or density) of the two liquids:
!
𝑑 =
!
(1)
Density is a convenient quantity to use for comparative purposes because it is independent
of the volume of liquid used. In other words, density is an intensive property. An
intensive property is one that is independent of the amount of substance. For this reason,
the densities of two different solutions can be compared without needing to have identical
solution volumes.
When comparing regular soda and diet soda, it is found that regular soda is more dense
than its sugar-‐free relative. To understand why, a molecular view of the two substances
must be examined. Obviously, the main difference between the two is the presence of
dissolved sugar in regular soda, which is absent in diet soda. Since diet soda lacks sugar, its
density is somewhat lower than that of regular soda. To a first approximation, soda can be
represented as a solution of sugar dissolved in water. As the amount of sugar dissolved in a
given volume of water increases, so does the density of the resulting solution. This
provides an ideal means by which the mass of sugar in a soda can be determined.
In this experiment, the relationship between the mass of dissolved sugar and density of
several sugar water solutions will be determined through the use of a calibration curve. A
calibration curve is constructed using known quantities. In this case, you will prepare
solutions of known volume using a known mass of dissolved sugar. After determining the
density of each solution, you will prepare a graph of density vs mass of dissolved sugar.
You will then graphically determine the relationship between the two quantities. After you
determine the density of a sample of soda, you will use your calibration curve to estimate
the mass of sugar present in your sample, and also in a 12 fluid ounce serving of the soda.
Scientific Graphs:
This experiment will also serve to introduce you to scientific graphing. Whenever you are
asked to produce a graph from laboratory data, all of the following criteria must be met:
1. All graphs must have a descriptive title
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Austin Community College
CHEM 1111 – General Chemistry I Lab
2. Axes must be labeled with an appropriate name and units
3. The data should occupy the full space of the graph
4. The minimum size of a graph should be half the size of a standard sheet of paper
5. When plotting a best fit line to the data, the line and equation (if applicable) should
be shown explicitly on the graph
6. The independent variable is plotted along the x-‐axis, and the dependent variable is
plotted on the y-‐axis
Figure 3–1 shows an example of an acceptable scientific graph of raw data. Figure 3–2
demonstrates the proper way to represent a linear fit on a graph.
Figure 3-‐1: Scientific graph of raw data
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Austin Community College
CHEM 1111 – General Chemistry I Lab
Figure 3-‐2: Graph including best-‐fit straight line
Graphing Using Microsoft Excel:
The basic steps required to graph data and perform linear regressions with Excel are
outlined below. These instructions apply to Excel 2013.
Basic Graphing:
1. With the program open, enter the data to be graphed in the cells. Enter the data to
be plotted along the x-‐axis in one column, followed by the associated data to be
plotted along the y-‐axis in the adjacent column.
2. Click and drag the mouse to highlight all the data to be graphed.
3. Click on the “Insert” tab at the top of the screen
4. Select the “Scatter” icon, followed by the unconnected points icon for the chart sub-‐
type. At this point a graph will appear on your screen.
5. Select the “Format” tab under “Chart Tools”. Use the drop-‐down menu at the upper
left of the ribbon to choose “Chart Title” or the appropriate axis. To label an axis,
click on the “+” icon to the right of the graph and be sure that “Axis Titles” box is
checked.
3-‐3
Austin Community College
CHEM 1111 – General Chemistry I Lab
Adding a Linear Trendline to a Graph:
1. Click on the “+” icon to the right of the graph, and check the “Trendline” box. Click
on the arrow to the right to select the “Linear” option. A best fit line will now
appear on your graph.
2. To display the equation of the line (and to access other options related to the
trendline) right click on the Trendline on the graph and select “Format Trendline”.
Click on the “Display Equation on Chart” box and then hit close.
Graphing By Hand:
Using your collected data, construct a graph of density (y) vs. mass of dissolved sugar (x).
Use the graph paper provided with the report sheet (last page of this handout). Remember,
you want your data to fill as much of the graph paper as possible! If your instructor does
not specify the smallest division to use on each axis, you use the following strategy:
1. Determine the range of values along the axis by subtracting the smallest value from
the largest value.
Suppose your minimum mass of sugar used was 0 g, and your maximum mass of
sugar used was 15.770 g. The range of values along the x-‐axis would be:
15.770 𝑔 − 0 𝑔 = 15.770 𝑔
If the minimum density was 0.9951 g/mL and the maximum density was 1.056
g/mL, the range of values along the y-‐axis would be:
1.056
𝑔
𝑚𝐿 − 0.9951
𝑔
𝑚𝐿 = 0.061
𝑔
𝑚𝐿
2. Count the number of divisions along each axis.
In the graph paper provided with this handout, there are 40 divisions along the x-‐
axis, and 40 divisions along the y-‐axis.
3. Calculate the size of the smallest division on each axis by taking the range of values,
and dividing by the number of divisions. Round this quotient UP to a convenient
(i.e. ends in 0 or 5) value!
!”.!!” !
𝑔
The smallest division along the x-‐axis =
= 0.39 𝑑𝑖𝑣
!” !”#
Since 0.39 is not a convenient number, round UP so the smallest division on the x-‐
axis is 0.40 g.
The smallest division along the y-‐axis =
!.!”#
!
!”;
!” !”#
!
= 0.001525 !” 𝑑𝑖𝑣
Round UP so the smallest division on the y-‐axis is 0.0020 g/mL. [Note: If the
smallest division were rounded DOWN to 0.0015 g/mL, the plotted data would not
fit on the graph paper! We always round UP when calculating the smallest division
to insure that the data fits on the graph paper being used.
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Austin Community College
CHEM 1111 – General Chemistry I Lab
After your data is plotted, use a ruler to draw the best fit straight line through your data
points. If your graph includes an origin (x = 0, y = 0), the best-‐fit line should pass through
the origin. Don’t connect the dots! Try to draw a line that passes through the origin, and
also touches or comes close to as many of your plotted data points as possible. Remember
to label both axes with units, and to give a descriptive title to your graph! In this
experiment, your graph represents the relationship between the density of a solution of
sugar water (something that can be determined experimentally) and the mass of sugar
used to prepare the solution (something that cannot be measured directly for an unknown,
like a sample of soda).
PROCEDURE
You will prepare five sugar water solutions. Each solution should have a different mass of
dissolved sugar covering a range from about 2 to about 16 grams of sugar per 100 mL of
solution volume.
To make the solutions in an accurate manner, they must be prepared in a 100 mL
volumetric flask, which is designed for accurate measurement of this volume of solution.
Note that there is always uncertainty associated with volumetric glassware. For significant
figure considerations, we will assume that the volume of the flask (and therefore the
volume of each solution) is accurate to ± 0.1 mL. As a consequence, the volume of each
solution will be considered to be 100.0 mL (4 significant figures).
As shown in Figure 3–3, a volumetric flask is marked with one fill line, somewhere on the
neck. When filled to the marked line, the flask accurately holds the stated volume. To fill a
volumetric flask, it is best to bring the fluid to the line carefully using a transfer pipet to
insure that the volume is accurate. The volume will not be known accurately if the flask is
overfilled via addition of liquid above the fill line. Additionally, if water is present on the
sides of the flask above the fill line, the result will be a solution that is more dilute than
intended.
FILL LINE
Figure 3–3: A volumetric flask
Before preparing any sugar water solutions, you should measure the mass of a clean, dry,
empty volumetric flask with stopper. If you use the same flask and stopper for every
solution, you only need to measure this mass once at the beginning of the experiment.
3-‐5
Austin Community College
CHEM 1111 – General Chemistry I Lab
You need to know the mass of sugar used to prepare each solution. Place a weigh boat on
the balance and tare (or re-‐zero) the balance so that it reads zero with the weigh boat on
the balance. Weigh out (and record) the desired mass of sugar.
Very carefully transfer the sugar to the flask. It is important that all the weighed sugar be
transferred to the flask. Using a funnel to transfer the sugar to the flask may help you avoid
spills. Your instructor will demonstrate the best approach to take for this operation if an
appropriately sized funnel is not available in your lab.
Add water to the flask until the bottom portion (the round part) is approximately half full.
Swirl the flask to dissolve the sugar. Do not shake the flask. Shaking or violent swirling
will cause water to collect above the fill line and will result in inaccurate data.
Once the sugar has completely dissolved, add water carefully until the level is slightly
below the fill line. To avoid adding too much water, add the last portion carefully (drop by
drop) until the bottom of the meniscus just touches the fill line. Insert the stopper and
invert the volumetric flask several times to insure that the solution is thoroughly mixed.
Determine the mass of the stoppered flask containing the solution and record the mass on
the report sheet.
After you measure this mass, you may pour the solution out of the volumetric flask (safe to
pour directly down the drain). Rinse the flask several times with deionized water. It is
now ready for the next mass of sugar that you weigh out. Remember, you only need to
measure the mass of the clean, DRY, stoppered flask once at the beginning of the
experiment. The flask does not need to be completely dried between solution preparations,
since you will be adding water to the sugar anyway to get it dissolved.
Once the data has been collected for all five sugar solutions, you should turn your attention
to the soda that your instructor assigns for study. Use the same volumetric flask and
stopper that you used to prepare the sugar water solutions. Remember, if you use this
same glassware, you already know the mass of the empty, dry stoppered flask. Carefully fill
the flask to the fill line with flat soda. (If you are not provided with flat soda, your
instructor will give instructions for the rapid degassing of carbonated soda.) Weigh and
record the mass of the stoppered flask containing the soda. Your instructor will tell you
what to do with the soda sample once this mass has been determined. Repeat with a
different soda sample if your instructor wants you to analyze more than one soda.
Clean all glassware, and return it to its proper location in the lab.
DATA ANALYSIS
You should calculate the mass of each sugar solution and soda sample by subtracting the
mass of the empty stoppered flask from the mass of the stoppered flask containing the
solution. The density of each solution can then be determined by taking the mass of the
solution and dividing by the solution volume. Remember, when properly filled to the mark,
a 100 mL volumetric flask holds 100.0 mL of solution (4 significant figures).
Using the collected data and calculated results, construct a graph of density (y) vs. mass of
sugar used (x). If plotting on a computer, have the program determine the best straight-‐
3-‐6
Austin Community College
CHEM 1111 – General Chemistry I Lab
line fit for the data, and report the equation for the line on the graph. If plotting by hand,
you will have to “eyeball” the best straight-‐line fit. Your instructor will discuss specific
graphing requirements for your class.
Your graph represents the relationship between the density of the sugar water solution
(which can be determined experimentally) and the mass of dissolved sugar in the solution
(which cannot be measured directly for the soda sample, because you did not add the sugar
yourself). There are two ways you can determine the mass of sugar in 100.0 mL of your
soda from your calibration curve. Your instructor will tell you which approach to use.
1. Use the equation of the best-‐fit straight line, which will be of the form y = mx + b, where
y = the density of the solution, x = the mass of sugar in the solution, m = the slope of the
line, and b is the y-‐intercept. For the soda sample, you know the y-‐value; use the
density of the soda solution that you calculated from your data. Since you did not add
the sugar to the soda, you do not know the x-‐value. Solve for x by rearranging the
equation. This value represents the mass of sugar in 100.0 mL of the soda.
2. If the equation of the best-‐fit straight line is not available, you will need to estimate the
mass of sugar as follows: Find the y-‐value that corresponds to the density of your soda
solution on the y-‐axis of your plot. Using a ruler, draw a straight, horizontal line from
this value to the best-‐fit straight line. At the point where this horizontal line intersects
the best-‐fit straight line, draw a perpendicular (i.e. vertical) line from the intersection to
the x-‐axis (use a ruler). The x-‐value where this vertical line crosses the x-‐axis
represents the mass of sugar in 100.0 mL of the soda.
A single can of soda does not contain 100.0 mL, but rather 12 fluid ounces (fl oz). To find
the mass of sugar in a can of your soda, you will need to perform a unit conversion
calculation. A volume of 12 fl oz is equivalent to 355 mL. Your instructor will tell you the
actual mass of sugar present in a can of each soda that is available for analysis. Compare
your result to the correct value by calculating your percent error.
SAMPLE CALCULATIONS
A student weighs an empty, dry, stoppered 100 mL volumetric flask and finds the mass to
be 65.649 g. After 3.929 g of sugar is weighed out, all of the solid is transferred to the
volumetric flask, enough water is added to completely dissolve the solid, and the flask is
filled to the mark. The mass of the stoppered flask containing the sugar water solution was
found to be 166.89 g. What is the density of the sugar water solution?
(1) Find the mass of the solution.
Mass of solution = Mass of stoppered flask + solution – Mass of empty, dry stoppered flask
Mass of solution = 166.89 g – 65.649 g = 101.24 g
3-‐7
Austin Community College
CHEM 1111 – General Chemistry I Lab
(2) Calculate the density of the solution.
The volume of the solution is 100.0 mL.
𝑑 =
101.24 𝑔
𝑚
𝑔
=
= 1.012 𝑚𝐿
100.0 𝑚𝐿
𝑉
WHY IS THIS LAB GREEN?
All sugar-‐water solutions and soda may be poured directly down the drain.
REFERENCE
J. Chem. Educ., 1999, 76 (10), p 1411 DOI: 10.1021/ed076p1411
3-‐8
CHEM 1111
Experiment 3
Name:
Section: Score:
Experiment 3 Prelab Assignment
1. A student weighs an empty 100 mL volumetric flask and stopper and finds the mass to
be 98.595 g. After adding 4.163 g of sugar to the flask, dissolving it, and filling the flask
to the mark, the stoppered flask is weighed again and the mass is found to be 204.274 g.
Calculate the density of this solution.
2. Explain how to properly prepare a solution in a 100 mL volumetric flask.
3. Suppose that in adding water to a volumetric flask containing dissolved sugar, you
accidentally add too much water, and the level rises above the mark. You decide to
remove the excess water with a pipet, and then you measure the mass of the stoppered
flask plus solution. Explain why the density calculated using this data will not be
accurate.
3-‐9
CHEM 1111
Experiment 3
Name:
Section:
Density (g/mL)
4. A sample of soda occupying a volume of 100.0 mL was found to have a mass of 103.87 g.
Use this data and the calibration curve provided below to determine the mass (in
grams) of sugar in the 100.0 mL sample of soda.
Density of Solution (g/mL) vs. Mass of Sugar (in grams)
1.06
y = 0.0038x + 0.9959
1.05
1.04
1.03
1.02
1.01
1
0.99
0
2
4
6
8
10
Mass of Sugar (g)
3-‐10
12
14
16
18
CHEM 1111
Experiment 3
Name:
Section:
Score:
Experiment 3 Report Sheet
Measured Data: Be sure to include UNITS for all measurements.
Your instructor will tell you how many soda samples you need to
analyze.
A
B
C
Mass of Empty
Mass of Sugar
Mass of Stoppered
Stoppered Flask
Flask with Solution
Calculated Results: SHOW WORK (on the
back or attach a separate page). Include
UNITS.
D
E
Mass of Solution
Density of Solution
(Volume = 100.0 mL when
flask is filled to the mark)
(measure ONCE at
beginning)
Sugar Solution 1
Sugar Solution 2
Sugar Solution 3
Sugar Solution 4
Sugar Solution 5
Soda Sample 1
Soda Sample 2
Soda Sample 3
N/A
N/A
N/A
3-‐11
CHEM 1111
Experiment 3
Name:
Section:
Use the data and results for the sugar solutions to construct a graph of density (y) vs. mass of sugar (x), and fit the data to a
linear relationship. Use the graph paper provided on the next page if plotting by hand, or use a computer. If you do use a
computer, report the equation for the line on the graph. This is your calibration curve. Attach a copy of your curve to this report
form.
Use your calibration curve to determine the mass of sugar in 100.0 mL
of each soda sample that you studied. Report your results in the table to
the right.
F
G
H
I
Grams
Calc.
Actual
Sugar in
Grams
Grams
% Error
100.0 mL
Sugar in
Sugar in
12 fl oz
12 fl oz
Soda
Calculate the mass of sugar in a 12 fl oz can of each soda sample. Report Sample 1
your results in the table to the right. 12 fl oz = 355 mL
Soda
Sample 2
Soda
Sample
3
Your instructor will provide you with the actual sugar content in 12 fl oz
of each soda sample. Use this information and your calculated results to
determine the percent error in your results. Report your answers in the
table to the right.
3-‐12
CHEM 1111
Experiment 3
Name:
Section:
Calibration Curve (Use if Graphing by Hand)
Determine the range of density values (max. value – min. value) =
Determine the value per y-‐axis division (range ÷ 40) rounded UP =
Determine the range of values for mass of sugar (max. value – min value) =
Determine the value per x-‐axis division (range ÷ 40) rounded UP =
(Be sure to label both axes, with UNITS!!!)
Graph Title:
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Austin Community College
CHEM 1111 – General Chemistry I Lab
Experiment 4
Which Alkali Metal Carbonate?
LEARNING OBJECTIVES
•
•
•
Learn how to apply the Law of Conservation of Mass to a chemical reaction, to
determine the mass of a product formed
Learn how to calculate the molar mass of an unknown
Learn how to identify a compound based on its experimentally determined molar
mass
INTRODUCTION
Imagine you are an analytical chemist and have received a sample of a pure alkali metal
carbonate from a newly discovered deposit. Your task is to determine which alkali metal
carbonate is present in the sample. You will determine an experimental molar mass for the
alkali metal in the carbonate and, from that result, you will determine which alkali metal is
present.
To identify your unknown, you will employ the Law of Conservation of Mass. This law
states that in a chemical reaction the total mass of the reactants must equal the total mass
of products.
Alkali metal carbonates (X2CO3) react with HCl according to the following equation (X is
meant to represent the unknown alkali metal cation):
𝑋! 𝐶𝑂! 𝑠 + 2 𝐻𝐶𝑙 𝑎𝑞 → 2 𝑋𝐶𝑙 𝑎𝑞 + 𝐶𝑂! 𝑔 + 𝐻! 𝑂(𝑙) (1)
Note that one of the products (CO2) is a gas. If the reaction is performed in an open vessel,
this gaseous product will escape as the reaction proceeds to completion. The mass of CO2
generated can be calculated if the mass of the reactants used is known accurately, and the
mass of the reaction mixture at the end of the reaction is also known.
Mass of CO2 = [mass of X2CO3(s) + mass of HCl(aq)] – (final mass of reaction mixture) (2)
Using dimensional analysis, the mass of CO2 can be converted to moles of CO2 using the
molar mass (ℳ) of CO2 as a conversion factor. The molar mass of CO2 is calculated by
adding together the atomic weights of all of the atoms in the molecule: 1 C atom (12.01
g/mol) plus 2 O atoms (2 x 16.00 g/mol) = 44.01 g/mol.
𝑔𝑟𝑎𝑚𝑠 𝐶𝑂!
ℳ!”!
𝑚𝑜𝑙𝑒𝑠 𝐶𝑂! (3)
The number of moles of CO2 is related to the number of moles of the alkali metal carbonate
(X2CO3) by the coefficients in the balanced chemical equation (Equation 1, written above).
This equation indicates that 1 mole of CO2 will be produced for every 1 mole of X2CO3
consumed in the reaction.
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Austin Community College
CHEM 1111 – General Chemistry I Lab
𝑚𝑜𝑙𝑒𝑠 𝐶𝑂!
! !”# !”! ! ! !”# !! !”!
𝑚𝑜𝑙𝑒𝑠 𝑋! 𝐶𝑂!
(4)
Once the number of moles of X2CO3 has been determined, the molar mass can be calculated
using the definition of molar mass:
𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 =
!”#$% !” !”#$”%&’
!”#$% !” !”#$”%&’
(5)
Finally, after the molar mass of X2CO3 has been calculated, the molar mass of the alkali
metal atom X can be determined, and the unknown metal can be identified:
𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑋! 𝐶𝑂! = 𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑎𝑙𝑙 𝑎𝑡𝑜𝑚𝑠 𝑖𝑛 𝑡ℎ𝑒 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑
= 2(ℳ of X) + ℳ of C + 3(ℳ of O)
(6)
The “Sample Calculations” section below provides examples of how to use Equations 2
through 6 to complete the calculations that you will need to perform once you collect your
data for Part A of the procedure.
An alternative approach for identifying the unknown requires that the volume of HCl
solution required to react completely with the X2CO3 be measured very accurately. For this
strategy to work, knowledge of the molarity of the HCl solution is also required. Molarity
is a common way to describe the concentration of a solution. It is defined to be equal to
the moles of solute divided by the volume of the solution, expressed in units of liters:
𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝐻𝐶𝑙 𝑎𝑞 =
𝑚𝑜𝑙𝑒𝑠 𝐻𝐶𝑙
𝑙𝑖𝑡𝑒𝑟𝑠 𝑜𝑓 𝐻𝐶𝑙 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
(7)
If the volume of HCl solution of known molarity required to react completely with a sample
of X2CO3 is measured very carefully, rearrangement of Equation (7) allows for the number
of moles of HCl consumed in the reaction to be calculated. The coefficients in the balanced
chemical equation (1) can then be used to determine the number of moles of X2CO3 that
reacted with the HCl:
…
