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Last updated 8/8/2013
CHEM 100 Lab Manual
Harrisburg Area Community College
2013/2014
2
Table of Contents
Equipment Illustrations ...............................................................................................................................5
Introduction: Measurements and Recording Data ..................................................................................7
Common Lab Equations ..............................................................................................................................9
Measurements: Density of a Saline Solution ..........................................................................................11
Pre-lab Questions ....................................................................................................................................15
Data: ........................................................................................................................................................17
Separation of Mixtures ..............................................................................................................................21
Pre-lab Questions ....................................................................................................................................25
Data ..........................................................................................................................................................27
Nomenclature..............................................................................................................................................31
Pre-lab Questions ....................................................................................................................................33
Stoichiometry: Empirical Formula of a Hydrate ....................................................................................41
Pre-lab Questions ....................................................................................................................................43
Data ..........................................................................................................................................................45
Balancing Chemical Equations .................................................................................................................49
Balancing Chemical Reactions ................................................................................................................51
Conductivity ...............................................................................................................................................55
Pre-lab Questions ....................................................................................................................................57
Data ..........................................................................................................................................................59
Chemical Reactions: Classification and Prediction of Products ..........................................................65
Pre-lab Questions ....................................................................................................................................69
Data ..........................................................................................................................................................71
Double Displacement Reactions ................................................................................................................73
Pre-lab Questions: ...................................................................................................................................77
Data ..........................................................................................................................................................81
Lewis Structures .........................................................................................................................................83
Pre-lab Questions ....................................................................................................................................89
Lewis Structures .......................................................................................................................................91
Gas Laws .....................................................................................................................................................97
Pre-Lab Questions .................................................................................................................................101
Data ........................................................................................................................................................103
Acid-Base Titrations ................................................................................................................................107
Pre-lab Questions: .................................................................................................................................111
Data ........................................................................................................................................................113
3
Spectroscopy: Determination of Concentration Using Beer’s Law....................................................115
Pre-lab Questions ..................................................................................................................................117
Data Table..............................................................................................................................................119
4
Equipment Illustrations
5
6
Introduction: Measurements and Recording
Data
Taking good measurements is one of the most important aspects of science. This is closely
followed in importance by being able to indicate to others how “good” the results are. Generally,
the final results of a lab can be no better than the data or measurement that the results are based
on, so the ability to get “good” results depends on how “good” the original data is.
How “good” are data and results?
The “goodness” (henceforth known as validity) of data relies primarily on two factors:
•
the measuring device
•
how well the experimenter used the device
For example, if I wanted to determine the mass of a paperclip, I would use a digital balance
instead of holding the paperclip in my hand, because the balance is supposedly better. But be
careful; the assumption that I should get a better value from the balance implies a whole bunch of
other factors (such as calibrating the balance and using the balance correctly.
Two general terms used to discuss the “goodness” of data are accuracy and precision. Accuracy
describes how closely a measured value matches a known value, while precision describes how
reproducible or how closely grouped a series of measurements are. “Good” data should be both
accurate and precise.
How can someone look at a value and tell how precise it is?
Scientists record all significant digits of their measurements. This allows them to communicate
the precision of data and results without having to specifically identify all devices used and the
exact procedures. The reader merely looks at the number to see how well it was measured.
An illustration of how this works follows: Two students are asked to count the amount of money
in a bag that’s in another room. Carol answers around $2, while Joan says $1.92. Who do you
think measured the best? If you’re like me, you answered Joan. You didn’t see them count, but
more decimal places usually implies better measuring.
The more significant digits—also referred to as significant figures—in a value, the more
precisely is was measured. So, when you record a value in a lab, BE CAREFUL. You aren’t
just writing down a number; you’re also telling someone how well you measured.
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CORRECTLY RECORDING EXPERIMENTAL DATA
How many decimal places should be recorded for data?
•
•
For digital devices, like a digital balance:
o Record ALL numbers in the digital display, and record the error limit often printed on the
device.
For non-digital devices like a volumetric pipet or a ruler:
o First look on the device to see if there is a +/- error range on it. For example, if you are
measuring volume with a 10 mL volumetric pipet, the pipet has “+/- 0.02 mL” marked on
it. That means that the pipet is designed to measure to within +/- 0.02 mL of the actual
value (assuming you measured properly), so any data you measure from this device should
be recorded to the hundredths place (i.e. 10.00 mL, not 10.0 or 10).
o If there is no error range on the device you will have to determine the least count of the
device.
 The least count is the smallest division on the measuring device. For example, if
the smallest division on a ruler is 1 cm, then the least count of that ruler is 1 cm or
0.01 m.
 Once you identify the least count, the general rule is to record the value of the
measurement to one decimal place smaller than that of the least count. For example
in the ruler above, a measurement would be recorded to the 1/10th cm (12.3 cm).
How many significant figures should be recorded at the end of a calculation?
Review your lecture notes for all of the guidelines. Here is a simple reminder:
•
•
•
For multiplication and division, the answer will have the least number of total significant figures.
For addition and subtraction, the answer will have the least number of decimal places.
Only round the final answer.
Always follow the rules of significant figures when recording values and performing
calculations.
8
Common Lab Equations
Average is the approximate middle value in a series of similar measurements.
Average =
Sum of all data values
Number of data points
Range is the difference between the highest and lowest valued measurement.
Range = | Highest value - Lowest value |
Percent Error is the percent an experimental value differs from a known or “true” value.
% error =
| Known value – Experimental value |
× 100
Known value
Percent Difference is used when a known value is not given.
| Experimental value #1 – Experimental value #2 |
% difference =
× 100
Average of #1 and #2
9
10
Measurements: Density of a Saline Solution
Objectives: In this lab, students will:
•
•
Determine the density of a saline solution by measuring its mass and volume
Compare the effectiveness of using two different devices for measuring volume: a graduated
cylinder and a volumetric pipet.
Skills: Upon completion of this lab, students should have learned:
•
•
•
To measure and record volume of liquids using graduated cylinders and volumetric pipets
To measure and record mass using a digital balance
To record numerical data and calculated values to the appropriate number of significant figures
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
•
•
•
Measurements and significant figures
Significant Figures in Calculations
Converting from One Unit to Another
Density
Classifying Matter
Physical and Chemical Properties
Section 2.3
Section 2.4
Section 2.6
Section 2.9
Section 3.4
Section 3.5
Lab Manual References: Read these sections PRIOR to lab.
•
•
Introduction: Measurements and Recording Data
Common Lab Equations
Introduction:
Measuring data is what differentiates science from other courses. In this lab, you will use
measurements to explore the concept of density. The density of an object is defined as the ratio of its
mass to its volume.
Density =
Mass
Volume
or
d=
m
V
Elements or compounds can often be identified by determining their density. For example, aluminum
and titanium look very similar, but the density of aluminum is 2.7 g/mL and that of titanium is
4.5 g/mL. If you measured the mass and volume of an unknown silvery metal and calculated the
density to be around 4.5 g/mL, the metal is probably titanium.
11
You will use a volumetric pipet to measure out the solution. A suction bulb is used to withdraw air
from the pipet while drawing up the liquid to be measured. Always use the suction bulb and always
hold the pipet with your dominant hand and hold the pipet bulb in your other hand . Never
pipet by mouth.
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Procedure: Determine the density of a saline solution
Part A.
Obtain 150-200 mL of saline solution in a 250 mL beaker. Record the known density.
Part B.
Graduated Cylinder
1. Measure and record the mass of a dry 150 mL beaker on the digital balance. Record
the mass to the correct number of decimal places.
2. Pour a volume of saline solution between 15 mL and 45 mL into a clean and dry 50 mL
graduated cylinder. Record the volume to the correct number of decimal places in
your data table.
3. Pour the saline from the graduated cylinder into the empty beaker. Measure and record
the mass of the beaker with the saline in it. Don’t empty the beaker.
4. Repeat step B2 and B3, two more times, using a different volume between 15 mL and
45 mL each time.
Part C.
Volumetric Pipet - Practice using the volumetric pipet with deionized water until you are
comfortable with it.
1. Measure and record the mass of a dry 50 mL beaker on the digital balance. Record the
mass to the correct number of decimal places.
2. Suction 10 mL of saline into a volumetric pipet (review instructions on p. 12). Record
the volume to the correct number of decimal places in your data table.
3. Release the saline from the volumetric pipet into the empty beaker. Measure and
record the mass of the beaker with the saline in it. Don’t empty the beaker.
4. Repeat step B2 and B3 two more times. Use the correct number of decimal places
for the volumetric pipet.
5. Clean up.
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Name: _________________________________________ Date due: ___________________________
Measurements: Density of a Saline Solution
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.
1. If a measured quantity is written correctly, which digits are certain? Which are uncertain?
2. A new penny has a mass of 2.49 g and a volume of 0.349 cm3. Is the penny pure copper? (dcopper =
8.96 g/cm3) Show your work!
3. Define accuracy and precision.
4. To the correct number of significant figures, what is the volume, in milliliters, of the liquid in the
graduated cylinder?
5.
Refer to the procedure and explain how you can make small adjustments to the volume of liquid in a
graduated cylinder.
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6.
The mass of a piece of copper (known mass = 4.750 g) is measured three times by two different students.
The students’ results follow:
Gloria
4.68 g
4.86 g
4.75 g
Max
4.69 g
4.71 g
4.66 g
a) Calculate the average mass of the copper determined by each student. Show your work
Gloria___________
Max____________
b) Calculate the range of each student’s measurements. Show your work
Gloria___________
Max____________
c) Calculate the percent error for each student’s mass measurement. (Use the calculated average mass as
the experimental value.) Show your work
Gloria___________
Max____________
d) Which student’s measurements were more precise? ___________
e) Which student’s measurements were more accurate? ___________
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Name: _________________________________________ Date lab performed: __________________
Partner(s) name:_________________________________ Date due: ___________________________
Measurements: Density of a Saline Solution
Data: Record all data with correct units and significant figures.
A. Known density of saline solution
_________
B. Graduated Cylinder
1. Volume of saline added each trial
_________
2. Mass of empty beaker
_________
3. Mass of beaker with first addition of saline
_________
4. Mass of beaker with second addition of saline
_________
5. Mass of beaker with third addition of saline
_________
C. Volumetric Pipet
1. Volume of saline added each trial
_________
2. Mass of empty beaker
_________
3. Mass of beaker with first addition of saline
_________
4. Mass of beaker with second addition of saline
_________
5. Mass of beaker with third addition of saline
_________
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Calculations:
•
•
Show neatly labeled and organized work for each of the following calculations below.
Include correct units and significant figures in all calculations.
Using the Graduated Cylinder Data from Part B, calculate the:
1. Mass (g) and density (g/mL) of the first addition of saline
2. Mass and density of the second addition of saline
3. Mass and density of the third addition of saline
4. Average experimental density of the three samples of saline
5. Range of the densities of the three samples of saline
6. Percent error between the known density (from Data A) and the average experimental
density (from Calc 4).
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Using the Volumetric Pipet Data from Part C, calculate the:
7. Mass (g) and density (g/mL) of the first addition of saline
8. Mass and density of the second addition of saline
9. Mass and density of the third addition of saline
10. Average experimental density of the three samples of saline
11. Range of the densities of the three samples of saline
12. Percent error between the known density (from Data A) and the average experimental
density (from Calc 10).
Results Table:
Part B
Part C
Cylinder
Pipet
1. Density of first volume of saline
_________
_________
2. Density of second volume of saline
_________
_________
3. Density of third volume of saline
_________
_________
4. Average density
_________
_________
5. Density range
_________
_________
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6. % error
_________
_________
Conclusion Questions:
1. Which term - range or % error - best describes the precision of data? __________________
2. Which term - range or % error - best describes the accuracy of data? _________________
3. Were your results more precise using the graduated cylinder or volumetric pipet? Explain
using your values.
4. Were your results more accurate using the graduated cylinder or volumetric pipet? Explain
using your values.
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Separation of Mixtures
Objectives: In this lab, students will:
• Separate the components of a mixture using filtration and evaporation.
•
Verify the Conservation of Mass Law.
Skills: Upon completion of this lab, students will have learned:
•
•
•
•
•
•
To use a digital balance to measure mass
To prepare solutions
To separate a mixture using vacuum filtration
To separate a mixture using evaporation
To identify common lab chemicals as mixtures or pure substances
To identify chemical and physical changes
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
•
Classifying Matter
Physical and Chemical Properties
How Matter Changes
Conservation of Mass Law
Section 3.4
Section 3.5
Section 3.6
Section 3.7
Introduction:
Mixtures are classified as either homogeneous or heterogeneous. When a mixture is homogeneous,
techniques such as evaporation and crystallization are used to separate the mixture into its
components. When a mixture is heterogeneous, techniques such as filtration and decantation are
used to separate the mixture into its components.
This experiment will study the reaction between calcium chloride and sodium carbonate, which is
shown below. The product mixture will be separated into pure sodium chloride and pure calcium
carbonate via filtration and evaporation.
CaCl2 (aq) + Na2CO3 (aq)  2 NaCl (aq) + CaCO3 (s)
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Procedure:
Part A.
Preparation of solutions
1. Measure out 0.9–1.1 g of solid anhydrous calcium chloride, CaCl2, using a digital
balance and weighing paper. Examine the solid CaCl2 and record your observations.
2. Transfer the CaCl2 into a large test tube, labeled “A”.
3. Use a graduated cylinder to measure approximately 10 mL of deionized water, and
add the water to the solid. Stir the solution until all of the solid CaCl2 dissolves.
Record your observations.
4. Measure out 0.9−1.1 g of solid anhydrous sodium carbonate, Na2CO3, using a digital
balance and weighing paper. Examine the solid Na2CO3 and record your observations.
5. Transfer the Na2CO3 into a large test tube, labeled “B”.
6. Use a graduated cylinder to measure approximately 10 mL of deionized water, and
add the water to the solid. Stir the solution until all of the solid Na2CO3 dissolves.
Record your observations.
Part B.
Mixing solutions
1. Mix the contents of test tube A and B together in a clean 50 mL beaker. Rinse each
test tube with a small amount of deionized water (~ 1 mL) and add the rinse water to
the beaker. Swirl the contents of the beaker, and record your observations.
Part C.
Vacuum filtration (Buchner filtration)
1. Assemble the vacuum filtration apparatus as indicated by your instructor.
2. Write your initials on a filter paper circle in pencil.
3. Record the mass of the filter paper and a watch glass together.
4. Place the filter paper in the Buchner funnel (initials down) and wet the paper with a
small amount of deionized water (~1 mL).
5. Turn on the vacuum and pour the beaker’s contents into the Buchner funnel. Record
your observations.
6. Use a small amount of deionized water (no more than 5 mL) to rinse the beaker.
With the rubber end of the stirring rod, transfer the rinse water and the remaining
precipitate (solid) from the beaker onto the filter paper.
7. Continue to pull a vacuum on the solid for 5 minutes.
8. Remove the filter paper, and place it on the watch glass (measured in step 8). Place
the watch glass in the oven.
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9. Record the mass of the watch glass with dried material before the end of the lab
period.
Part D.
Evaporation of water
1. Record the mass of a clean, dry evaporating dish.
2. Pour all of the liquid from the receiving flask into the evaporating dish. Rinse the
flask with a small amount of deionized water (less than 5 mL) and add the rinse to the
dish.
3. Place the evaporating dish on the clay triangle on a ring stand (as indicated by your
instructor) and evaporate the liquid using a Bunsen burner. Record your observations
as the liquid evaporates. Reduce the flame near the end of the heating (when most of
the water has evaporated) to prevent the solid from “popping” out of the dish.
Continue to heat for approximately 5 minutes after you see no more steam.
4. Turn off the gas and allow the dish to cool.
5. Record the mass of the dish and remaining solid.
6. Dispose of solids and filter papers in the appropriate container.
7.
Clean up.
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Name: __________________________________________Date due: ___________________________
Separation of Mixtures
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.
1. Define the following terms in your own words:
Pure substance—
Mixture—
Heterogeneous mixture—
Homogeneous mixture—
Physical change--
Chemical change—
2. State the Law of the Conservation of Mass.
3. Using the conservation of mass law, calculate how much NaCl, table salt, forms when 27.4 g Na
reacts with 42.3 g Cl2.
4. What kind of change occurs when NaCl is formed from Na and Cl2? Hint: Examine the images
of elemental sodium and elemental chlorine shown in section 5.1 in your textbook.
25
5. If NaCl was placed in a beaker of water and dissolved, what kind of change would occur?
6. Write the chemical equation for the reaction you will perform in this lab.
7. Refer to the procedure and indicate the technique you will use to isolate NaCl from the mixture.
8. Refer to the procedure and indicate the technique you will use to isolate CaCO3 from the mixture.
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Name: __________________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date : ______________________________
Separation of Mixtures
Data:
Record all data with correct units and significant figures.
A.
Mass of weighing paper and solid CaCl2
___________
Mass of weighing paper
___________
Mass of solid CaCl2
___________
Mass of weighing paper and solid Na2CO3
___________
Mass of weighing paper
___________
Mass of solid Na2CO3
___________
B. Observations:
CaCl2 solid:
CaCl2 solution:
Na2CO3 solid:
Na2CO3 solution:
After mixing solutions A and B:
Vacuum filtration (see procedure, part C, step 5):
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C. Mass of filter paper and watch glass
___________
Mass of dried filter paper, watch glass, and solid
___________
Observations:
D. Mass of evaporating dish
___________
Mass of evaporating dish and solid
___________
Observations:
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Calculations:
•
•
Show your calculations.
Include correct units and significant figures in all calculations.
Calculate:
i.
The total mass of dissolved solids (in Part A).
ii.
The mass of solid from filtration.
iii.
The mass of solid from evaporation.
iv.
The total mass of recovered solids.
v.
The % error between the total mass of dissolved solids (reactants) and the total mass
of recovered solids (products). Use the total mass of the dissolved solids as the
“known” value in the calculation,
Results Table:
total mass of dissolved solids
_________
total mass of recovered solids
_________
% error
_________
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Conclusion Questions:
1. How well do your results support the Conservation of Mass law (very well, OK, not very
well)? Support your answer using your result values. Evaluate your technique and
discuss where procedural errors may have occurred.
2. Identify the following as an element, a compound, a homogenous mixture, or a
heterogeneous mixture.
a. Solid sodium carbonate
_________________________
b. Sodium carbonate solution
_________________________
c. The result of Part B
_________________________
d. The liquid from Part C
_________________________
3. Identify the following as a physical or chemical change.
a. The dissolving of CaCl2 in water
_________________________
b. The formation of the white precipitate in Part B ______________________
c. The separation of solid from solution in Part C _______________________
d. The evaporation of the liquid in Part D
_________________________
4. The solid collected on the filter paper is either NaCl or CaCO3. Explain which one it is.
5. Identify the liquid that is evaporated in Part D. Explain your answer.
6. Which compound remains after evaporation? Explain your answer.
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Nomenclature
Objectives: In this lab, students will learn:
•
•
•
•
•
To identify a compound as ionic, molecular, or an acid from either its name or its formula
To identify cations, anions, and charges
To identify compounds as binary or polyatomic
To correctly name compounds when given the formula
To write the correct chemical formula when given the name
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab,
and bring your textbook with you to lab.
• Chemical Formulas
• A Molecular View of Elements
• Writing Formulas for Ionic Compounds
• Nomenclature: Naming Compounds
• Naming Ionic Compounds
• Naming Molecular Compounds
• Naming Acids
• Nomenclature Summary
Section 5.3
Section 5.4
Section 5.5
Section 5.6
Section 5.7
Section 5.8
Section 5.9
Section 5.10
Introduction:
Names and Formulas are used to identify chemicals. Each different chemical has a unique formula
and a unique IUPAC name. The IUPAC naming system was developed so that everyone names the
compound the same way to avoid confusion. The IUPAC names will be used in class, but many
compounds also have a common name that you already know.
Most of the chemical compounds used in CHEM 100 can be roughly separated into three main
categories – acids, ionic compounds, and molecular compounds.
Acids are easy to identify by their name or formula. The name of an acid always includes the term
“acid” and the formula of an acid usually begins with the chemical symbol for hydrogen, H. There
are two main types of acids – binary acids and oxyacids.
Binary compounds contain only two
different elements. The first example below is a binary acid and the other two are oxyacids.
IUPAC name Formula
Common name or use
hydrochloric acid
HCl
muriatic acid; used in concrete work and welding
acetic acid
HC2H3O2
vinegar
sulfuric acid
H2SO4
battery acid
31
Ionic compounds often include a metal in the formula and name. There are two types of ionic
compounds: binary ionic and polyatomic ionic. The last two examples have polyatomic ions—ions
with more than two different elements. The names of ionic compounds with polyatomic ions in
them will often end in “ate” or “ite”. A table of the most common polyatomic ions is found in Tro
(Table 5.6, p. 141).
Notes:
•
The last example has no metal in it, but is still an ionic compound because it contains ions.
•
The second example below includes a Roman numeral after the name of the cation, iron. A
Roman numeral is used to identify the ionic charge for metals that can have different charges.
(These metals are referred to as Type II metals.) Iron(III) is the name of Fe3+.
IUPAC name
sodium chloride
iron(III) oxide
potassium nitrite
ammonium nitrate
Formula
NaCl
Fe2O3
KNO2
NH4NO3
Common name or use
table salt
rust
used to cure meats
used in instant cold packs
Molecular compounds include only nonmetals in the formula, and they aren’t acids. You will only
be responsible for naming binary molecular compounds like the first three examples below.
Molecular compounds that contain more than two different elements, like the last example, are
organic chemicals and have an entirely different naming scheme which is beyond the scope of the
CHEM 100 course.
IUPAC name
dinitrogen monoxide
carbon dioxide
ammonia
Formula
N2O
CO2
NH3
ethyl alcohol
C2H6O
Common name or use
nitrous oxide or “laughing gas”
one of the green house gases
this compound is so common that the
common name has been adopted as the
IUPAC name, just like water
in beer, wine, and distilled spirit
32
Name: __________________________________________Date due: ___________________________
Nomenclature
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering prelab questions.
1. Define the following terms in your own words:
a. element
b. compound
c. metal
d. non-metal
e. main group metal
f. transition metal
g. acid
h. ion
i. cation
j. anion
k. polyatomic ion
2. Explain how to name ionic compounds containing type II metals. Give two examples.
33
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35
36
Name: __________________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date due: ___________________________
Nomenclature
A.
B.
Look at each formula in the list below and identify each
• As Acid (A), Ionic (I), or Molecular (M) in the first column
• As binary (B) or polyatomic (P) in the second column
• As Type II (T) if it contains a type II metal
Formula
Acid, Ionic,
or Molecular
Binary or
Polyatomic ion
Type II
metal
KI
________
________
________
HCl
________
________
________
BaSO4
________
________
________
SO2
________
________
________
Fe(NO3)2
________
________
________
HgI
________
________
________
H2SO4
________
________
________
Write names for the ions found in each ionic compound and indicate the charge on each
individual ion (not the total charge). Include roman numerals with the cation if necessary.
Name of Cation
Charge
ScCl3
PbS2
K2SO4
Al(C2H3O2)3
Fe2O3
Mn(HCO3)2
CsMnO4
(NH4)2SO3
37
Name of Anion
Charge
Naming compounds when given the formula
C.
1. Name each ionic compound. (Remember to look for type II metals.)
Hg2O
CaF2
Al2(SO4)3
Au2CO3
2. Name each molecular compound.
CCl4
SF6
N2O5
S2O3
3. Name each aqueous acid.
HNO2(aq)
HC2H3O2(aq)
H2SO4(aq)
HCl (aq)
4. Name each compound.
CrCl3
Li3N
K2O
SrS
CO
(NH4)2SO3
Ni(NO2)3
HI (aq)
HF (aq)
H2CO3 (aq)
H3PO4 (aq)
HNO3(aq)
38
D.
Write chemical formulas when given the name.
1. Write the formula for each ionic compound. (Remember to look for type II metals.)
gallium oxide
copper(II) bromide
aluminum nitrite
silver phosphate
2. Write the formula for each molecular compound.
sulfur dioxide
phosphorus trichloride
iodine monobromide
nitrogen monoxide
3. Write the formula for each aqueous acid.
sulfuric acid
hydrobromic acid
perchloric acid
sulfurous acid
4. Write the chemical formula.
water
lithium sulfate
ammonia
ammonium hydrogen sulfate
silver carbonate
sodium bicarbonate
ammonium nitrite
oxygen difluoride
iron(II) nitrate
potassium phosphate
calcium nitride
barium sulfide
sodium phosphide
lead(IV) oxide
zinc acetate
strontium hydroxide
cobalt(III) oxide
aluminum cyanide
39
40
Stoichiometry: Empirical Formula of a
Hydrate
Objectives: In this lab, students will:
•
Determine the empirical formula of a hydrate
Techniques: Upon completion of this lab, students will have learned:
•
•
To determine mass percent composition
To determine empirical formula
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
•
Converting between grams and moles
Mass Percent Composition
Calculating Empirical Formulas
Classifying Chemical Reactions
Section 6.4
Section 6.6
Section 6.8
Section 7.10
Introduction:
Gravimetric analysis involves comparing the mass of a compound before and after a chemical
reaction. Gravimetric analysis is often used in experiments to determine the stoichiometry of a
chemical reaction.
Decomposition reactions involve one chemical breaking down into two or more simpler chemicals.
Some decomposition reactions occur spontaneously at room temperature while others require higher
temperatures.
Hydrates are a common class of compounds that have water molecules loosely bound within the
crystal structure of a solid. The number of bound water molecules depends on the hydrate being
examined. The loosely bound water molecules can be driven off through heating. The waterless
form of the compound is called the anhydrous form of the compound.
In this experiment, you will determine the stoichiometric ratio of water molecules in copper(II)
sulfate hydrate. When copper(II) sulfate hydrate is heated, it gives off water vapor according to the
decomposition reaction below.
CuSO4 • nH2O (s) → CuSO4 (s) + nH2O (g)
blue
white
The value for n, which represents the mole ratio of water molecules to CuSO4 in copper (II) sulfate
hydrate will be calculated from the experimental data.
41
Procedure:
Safety Precautions:
•
The Bunsen burner flame is extremely hot and is nearly invisible. The metal stand, clamp
and the clay triangle look the same whether hot or cold. Use extreme caution.
Procedure:
1. Assemble the ring stand, ring clamp, clay triangle and Bunsen burner as demonstrated by
your instructor.
2. Obtain a clean crucible and record its mass.
3. Add approximately 5 g of copper(II) sulfate hydrate into the crucible. Record the mass of the
crucible and copper (II) sulfate hydrate. Describe the appearance of copper(II) sulfate
hydrate.
4. Heat the crucible GENTLY over the Bunsen burner for 15 to 20 minutes, or until the sample
has completely turned white. Avoid strong heat, because it may trigger the decomposition of
anhydrous copper sulfate: CuSO4 (s, white)  CuO(s, brown) + SO3(g)
5. Remove the crucible from the heat, and allow it to cool completely.
6. Record the mass of the crucible and anhydrous copper(II) sulfate.
7. Repeat heating (5 minutes) and cooling cycles until the mass of the crucible and anhydrous
copper(II) sulfate changes by less than 0.05 g between two consecutive heatings. Record the
appearance of the anhydrous copper(II) sulfate.
8. Place the anhydrous copper sulfate in the waste container, and remove any remaining solid
from the crucible by washing with water.
42
Name: __________________________________________Date due: ___________________________
Stoichiometry: Empirical Formula of Copper(II) Sulfate Hydrate
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering prelab questions.
1. Define the terms hydrate and anhydrous in an inorganic chemistry context, and provide an
example of each.
2. It is very difficult to predict the number of water molecules in a hydrate compound. Go
online to a legitimate source (such as a chemical supplier or a college) and get the actual
formula for copper(II) sulfate • n hydrate. This will tell you n, which is the ratio of water
molecules to copper(II) sulfate units. Reference your source.
3. A 2.241 g sample of nickel combines with oxygen to produce 2.852 g of a metal oxide.
a.
Calculate the number of moles of nickel in the metal oxide.
b. Calculate the number of grams of oxygen in the metal oxide.
c. Calculate the number of moles of oxygen in the metal oxide.
d. What is the empirical formula of the metal oxide?
43
44
Name: __________________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date due: ___________________________
Stoichiometry: Empirical Formula of Copper(II) Sulfate • n
Hydrate
Data:
Record all data with correct units and significant figures.
1. Mass of crucible
_________
2. Mass of crucible and copper(II) sulfate • n hydrate
_________
3. Mass of crucible and anhydrous copper(II) sulfate (1st heating) _________
4. Mass of crucible and anhydrous copper(II) sulfate (2nd heating) _________
5. Mass of crucible and anhydrous copper(II) sulfate(3rd heating) _________
(if needed)
Observations:
Initial appearance of the copper(II) sulfate • n hydrate:
Observed changes during heating:
Final appearance of the anhydrous copper(II) sulfate:
45
Calculations:
•
•
Show calculations and express answers with correct units and significant figures.
Record the results in the results table.
1. Mass of copper(II) sulfate • n hydrate
______________g
2. Mass of anhydrous copper(II) sulfate
_______________g
3. Moles of anhydrous copper(II) sulfate
_______________mol
Molar mass of copper(II) sulfate
_______________g/mol
4. Mass of water in the copper(II) sulfate • n hydrate _______________g
5. Moles of water in the copper(II) sulfate • n hydrate _______________mol
Molar mass of water _______________g/mol
6. Calculate the ratio of moles of water over moles of anhydrous copper(II) sulfate. This number
equals n.
n = _________________
7. Write the experimental formula for copper(II) sulfate • n hydrate in the form CuSO4 ∙ nH2O.
Round n to nearest whole number.
___________________
8. Calculate the number of moles of copper(II) sulfate pentahydrate used in this experiment.
___________________mol
46
9. Compile your data in the table below:
Compound
Formula
Molar Mass
(g/mol)
Experimental
mass (g)
Moles
Copper(II) sulfate
pentahydrate
Copper(II) sulfate
(anhydrous)
Water
Conclusion Questions:
1. How does your experimental empirical formula of copper(II) sulfate • n hydrate compare to
your predicted formula from the pre-lab? Analyze your procedure and identify where your
experimental error could have occurred.
2. Using your experimental data, calculate the percent of water in your copper(II) sulfate • n
hydrate, by mass. Show your work.
47
3. If you used twice as much copper(II) sulfate • n hydrate at the beginning of the lab, would
you expect the empirical formula of the product to be the same or different? Explain.
4.
If not all of the water was removed from the copper(II) sulfate pentahydrate after heating,
would you expect the empirical formula of the product to be the same or different? Explain.
48
Balancing Chemical Equations
Objectives: In this lab, students will:
• Balance equations
•
Classify chemical equations
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
The Chemical Equation
How to Write Balanced Chemical Equations
Section 7.3
Section 7.4
Balancing Equations:
To balance equations, it is necessary to only CHANGE WHOLE NUMBER COEFFICIENTS IN
FRONT OF formulas for elements and compounds! In the reaction, ALL FORMULAS FOR
REACTANTS AND PRODUCTS ARE INCLUDED, and ALL FORMULAS FOR REACTANTS AND
PRODUCTS ARE CORRECT. Subscripts in the chemical formula are NEVER changed.
Easy Rules for Balancing Equations:
1. Balance everything but hydrogen and oxygen.
2. Polyatomic ions can be treated as one unit if they appear in both a reactant and a product.
3. Balance hydrogen.
4. Balance oxygen.
5. Redo rules 1, 2, and 3 again, if necessary.
49
50
Name: __________________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date due: ___________________________
Balancing Chemical Reactions
A.
BALANCING EQUATIONS: Balance each of the chemical equations. Coefficients of 1
shouldn’t be written.
1.
___C6H14
+ ___O2
→ ___CO2
+ ___H2O
2.
___C9H20
+ ___ O2
→ ___CO2
+ ___H2O
3.
___Zn + ___HCl
→ ___ZnCl2
4.
___P4 + ___Cl2
→ ___PCl3
5.
___NaHCO3 → ___Na2CO3
6.
___HNO3
7.
___Fe + ___O2
8.
___CaC2
+ ___H2O
→ ___Ca(OH)2
9.
___Mg3N2
+ ___H2O
→ ___NH3
10.
___CaCO3
+ ___HCl
→ ___CaCl2 + ___CO2
11.
___Zn + ___H3PO4 → ___Zn3(PO4)2
12.
___AgNO3
→ ___NO2
+ ___H2
+ ___CO2
+ ___H2O
+ ___H2O
+ ___O2
→ ___Fe2O3
+ ___CaCl2
+ ___C2H2
+ ___Mg(OH)2
+ ___H2O
+ ___H2
→ ___AgCl + ___Ca(NO3)2
51
+ ___H2SO4 → ___Al2(SO4)3
13.
___Al2O3
+ ___H2O
14.
___Fe + ___Br2
15.
___Al(OH)3
+ ___H2SO4 → ___Al2(SO4)3
16.
___C2H2
+ ___O2
17.
___Li2O + ___H2O
→ ___LiOH
18.
___NH3
→ ___(NH4)2CO3
19.
___C2H5OH + ___O2
20.
___H3PO4 + ___Ca(OH)2 → ___Ca3(PO4)2
21.
___HBr + ___K2SO3 → ___H2O + ___SO2 + ___KBr
22.
___Na
23.
___Al
24.
___KClO3
25.
___(NH4)2SO4 + ___BaCl2 → ___NH4Cl + ___BaSO4
26.
___CH3OH
→ ___FeBr3
+ ___H2CO3
→ ___CO2
→ ___CO2
+ ___H2O
+ ___H2O
+ ___H2O
+ ___H2O
+ ___H2O → ___NaOH + ___H2
+ ___Fe2O3 → ___Al2O3 + ___Fe
→ ___KCl
+ ___O2
+ ___O2
→ ___CO2 + ___H2O
52
B.
Write a balanced chemical equation for each reaction described below (include states):
1. Solid magnesium oxide is produce by heating solid magnesium metal in the presence of oxygen
gas.
2. Solid calcium reacts with nitric acid to form aqueous calcium nitrate and hydrogen gas.
3. Aqueous hydrochloric acid reacts with solid manganese(IV) oxide to produce aqueous
manganese(II) chloride, liquid water, and chlorine gas.
53
54
Conductivity
Objectives:
•
Students will study conductivity as they compare:
o Ionic and covalent compounds
o Solutions of varying concentrations
o Weak acids/bases and strong acids/bases
Techniques: Upon completion of this lab, students will have learned:
•
To predict conductivity of a compound based on chemical formula
•
To classify compounds as non-electrolytes, weak electrolytes, or strong electrolytes based on
formula
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
A Molecular View of Elements and Compound
Section 5.4
•
Aqueous Solutions and Solubility
Section 7.5
•
Strong and Weak Acids and Bases
Section 14.7
Introduction:
You might remember from Section 4.6 of Tro that metals are good conductors of electricity, while
non-metals are poor conductors. A quick electrical conductivity test can identify an element as
either a metal or non-metal. What about compounds? In this lab, you will explore the relationship
between conductivity, chemical formula, and physical state. You will examine the differences in
conductivity for ionic versus molecular compounds, and how acid strength affects conductivity of
acid solutions.
55
Procedure:
Safety Precautions:
• Normal precautions need to be taken with acid and base solutions
• Dispose of all solutions in the appropriate waste container as instructed
•
Part A: Conductivity Meter Operation
1. Clean and dry a well plate.
2. Place approximately 1 mL of solutions SE, WE and NE into 3 separate wells. (SE stands for
strong electrolyte, WE for weak electrolyte, NE for non-electrolyte.)
3. Immerse the electrodes of the meter into the SE solution. Record your observations.
4. Rinse and dry the meter electrodes. Repeat steps 3 and 4 for the WE and NE solutions.
Part B: Conductivity of pure water, solids and solutions
1. Place ~ 1mL (20 drops) of pure water (distilled or deionized) and ~0.1 g NaCl into separate
clean, dry wells. Measure the conductance of each substance (SE, WE, or NE).
2. Mix the water and the solid NaCl in one well. Stir until the solid dissolves. Measure and record
the conductance of this saline solution.
3. Dilute the saline solution by taking 1 drop of it and placing it in a clean well. Use a clean
dropper to add 9 drops of pure water. Stir the mixture, and then measure the conductivity.
4. Make an even more dilute saline solution by diluting 1 drop of the saline solution from step #3
with 9 drops of pure water. Stir the mixture, and then measure the conductivity.
5. Repeat steps 1-4 using table sugar (C12H22O11), potassium nitrate (KNO3), and calcium acetate
[Ca(C2H3O2)2] solids in place of the NaCl.
Part C: Conductivity of liquids
1. Measure the conductivity of approximately 1 mL of ethanol (C2H5OH).
2. Measure the conductivity of approximately 1 mL of tap water.
Part D: Conductivity of acid solutions
1. Measure the conductivity of approximately 1 mL of 0.1 M HC2H3O2, 0.1 M H2SO4, and 0.1 M
HCl in 3 separate clean, dry wells.
56
Name: __________________________________________Date due: ___________________________
Conductivity
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering prelab questions.
1. Describe the major difference between an ionic and a molecular compound.
2. Describe the difference between a strong acid and a weak acid.
3. Identify the following as strong (S) or weak (W) acids.
a. _______ 0.1 M HCl (aq)
b. _______ 0.1 M H2SO4 (aq)
c. _______ 0.1 M HC2H3O2 (aq)
4. Identify the following as ionic (I) or molecular (M) compounds.
a. _______ H2O
b. _______ NaCl
c. _______ C12H22O11
d. _______ Ca(C2H3O2)2
e. _______ C2H5OH
57
58
Name: _________________________________________ Date lab performed: __________________
Partner(s) name:_________________________________ Date due: ___________________________
Conductivity
Data:
Part A: Conductivity Meter Operation.
Solution
SE
Meter Observation
WE
NE
How the meter indicates conductivity.
Part B: Conductivity of pure water, solids and solutions.
Substance
Pure water, H2O (l)
Conductivity
Substance
Pure water
NaCl (s)
C12H22O11 (s)
NaCl (aq)
C12H22O11 (aq)
Diluted NaCl (aq)
Diluted C12H22O11 (aq)
Most Diluted NaCl (aq)
Most Diluted C12H22O11 (aq)
Substance
Conductivity
Substance
Pure water
Pure water
KNO3 (s)
Ca(C2H3O2)2 (s)
KNO3 (aq)
Ca(C2H3O2)2 (aq)
Diluted KNO3 (aq)
Diluted Ca(C2H3O2)2 (aq)
Most Diluted KNO3 (aq)
Most Diluted Ca(C2H3O2)2 (aq)
59
Conductivity
Conductivity
Part C: Conductivity of liquids.
Substance
Part D: Conductivity of acid solutions.
Conductivity
Substance
Ethanol, C2H6O (l)
0.1 M HC2H3O2 (aq)
Tap water
0.1 M H2SO4 (aq)
Conductivity
0.1 M HCl (aq)
Conclusion Questions:
1) Draw a picture of what happens to the formula units in the following aqueous solutions. Do not draw
the water molecules; just the particles present after dissolving the compounds. Assume you have 4
formula units dissolved in the appropriate beaker. See an example on page 215 of text. You may use
simple characters or geometric shapes to represent each ion. For example a circle may represent a
sodium ion and a square the chlorate ion for the compound sodium chlorate. Label one of both shapes in
your drawing. The line on each beaker represents the surface of the solution.
KNO3 (aq)
NaCl (aq)
2) Does each the first solutions prepared with these ionic compounds in part B conduct an electrical
current?
3) What common feature is present in both solutions?
60
4) You tested a solution CaC2H3O2 for electrical conductivity. What might explain the electrical
conductance of CaC2H3O2?
5) Based on your experimental results in part B, describe how dilution of the solutions affected the
conductivity.
6) Hypothesize why tap water is a better conductor of electricity than pure water.
7) Hypothesize why solutions of NaCl conduct electricity, but solid NaCl does not.
8) In aqueous solutions containing molecular compounds the molecules of the dissolved compound are
separated and distributed throughout the water. For example a solution of sugar consists of individual
sugar molecules distributed through the water.
Draw a picture of what happens to the molecules in the following aqueous solution. Do not draw the
water molecules; just the particles present after dissolving the compounds. Assume you have 4 molecules
dissolved in the solution.
C2H6O (aq)
61
9) Draw a picture of what happens to the formula units in the following aqueous solutions. Do not draw
the water molecules; just the particles present after dissolving the compounds. Assume you have 4 acid
molecules dissolved in each beaker. See examples on page 499-500 of text.
HC2H3O2 (aq)
HCl (aq)
10) Is HCl a strong or weak acid?
11) Is HC2H3O2 a strong or weak acid?
12) What is the difference between your drawings of the two acids?
13) Is HCl (aq) a strong, weak or nonelectrolyte?
13) Is HC2H3O2 (aq) a strong, weak or nonelectrolyte?
15) Explain how acid strength affects electrical conductivity.
62
16) What is common among all the solutions that conducted an electrical current?
17) Suppose that each compound below is placed in deionized water. Indicate whether each mixture is a
strong, weak or nonelectrolyte.
MgI2
Al(NO3)3
HC2H3O2
Cs3PO4
HClO4
(NH4)2SO4
NaOH
CH2O
63
64
Chemical Reactions:
Classification and Prediction of Products
Objectives: In this lab, students will:
•
•
Classify reactions
Predict products of reactions
Skills: Upon completion of this lab, students should have learned to
•
•
Perform and observe chemical reactions
Identify chemical changes
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
•
Precipitation Reactions
Acid-Base and Gas Reactions
Oxidation-Reduction Reaction
Classifying Chemical Reaction
Section 7.6
Section 7.8
Section 7.9
Section 7.10
Introduction:
Chemical reactions can be classified in many ways. One way is to classify the reaction by what is
formed during the reaction (precipitation, acid-base, gas evolution, or oxidation-reduction reactions).
Another way, and the one used in this lab, is to classify reactions by what the atoms or groups of
atoms do. In this classification method, reactions are classified as double displacement, single
displacement, synthesis, or decomposition reactions. Another common type of reaction that will be
examined is the combustion reaction.
Reactions occur in these general forms:
Type of Reaction
Generic Equation
Double Displacement
AB + CD  AD + CB
Single Displacement
A + BC  AC + B
Synthesis (or combination)
A + B  AB
Decomposition
AB  A + B
Combustion
Hydrocarbon + O2  CO2 + H2O
The key to classifying a reaction is to look closely at the reactants (and the products, if you know
them). In the table above, whenever a letter is written in the equation by itself, it usually means that
material is an element (like the “A” on the reactant side of the single displacement reaction).
Decomposition reactions have exceptions to this. If the letters are written together (as in “AB” on the
reactant side of the double displacement reaction), it means that material is a compound.
65
Procedure:
Safety Precautions:
• Before you start any reaction
 Carefully read the whole procedure for that reaction, taking note of all safety concerns like
hot glass, a flame, or acids and bases. If you are not sure of the procedure, ask the
instructor BEFORE running the reaction.
 Gather all of the reactants and equipment at your work station. Do NOT perform the
reaction at the side bench.
 Return chemicals to the side bench when you are finished with them.
 Move slowly, and clean-up after each reaction.
 Dispose of all chemicals in the proper waste container.
Special Notes:
• Write observations before, during, and after completion of each reaction.
•
Unless instructed otherwise,
o All reactions are performed in large test tubes.
o Add chemicals in the written order.
o Use graduated pipets to dispense solutions.
o All masses used are approximate and you can use any amount within ± 10% of the
stated value. Do not waste time trying to get exact masses.
o When heating chemicals in a test tube, clamp the test tube to a stand. Angle the test
tube to ensure it is not aimed at anyone. Use the Bunsen burner with a moderately
sized blue flame. (Only #7 will be heated for this lab.)
•
Observe each reaction for a color change, precipitation, gas evolution, and/or temperature
change.
•
In reactions that produce gases, you will test to see if the gas produced is H2, O2, or CO2.
These tests are performed as the reaction is fully underway and the test tube has filled with
the gaseous product. Don’t wait until the reaction stops.
o H2 test - A wooden splint is lit in a Bunsen burner flame and first placed near the
top of the test tube, and then slowly inserted into the test tube. If the gas “pops”
hydrogen gas is present.
o O2 and CO2 test - A wooden splint is lit and allowed to burn for a few seconds and
then blown out. The red ember of the wooden splint is placed inside the test tube.
If the ember glows very brightly or the flame is rekindled, oxygen gas is present. If
the ember quickly goes out, carbon dioxide gas is present.
66
Reactions: Perform the following seven reactions. Record your observations on the data sheet.
Complete all of the reactions and your observations before you try to classify, write, and/or balance the
chemical equations.
Reaction 1:
Mix 2 mL of 0.1 M calcium chloride with 2 mL of 0.1 M sodium phosphate.
Reaction 2:
Add 2 mL of 3 M hydrochloric acid to a small piece of zinc. Test for H2.
Reaction 3:
Mix 2 mL of vinegar (HC2H3O2) with a very small amount of baking soda (NaHCO3).
Test for O2 and CO2.
Reaction 4:
Place 1 g of ammonium chloride and 2 g of strontium hydroxide in a large test tube. Stir
the solid contents with a stirring rod for 3-5 minutes. Gently waft your hand across the top
of the tube, and note the odor. Try to identify the common household chemical that has
this smell.
Reaction 5:
Mix 2 mL of 3 M sulfuric acid with 4 mL of 3 M sodium hydroxide.
Reaction 6:
Mix 2 mL of 0.1 M lead(II) nitrate with 2 mL of 0.1 M potassium iodide.
Reaction 7:
Heat 0.5 g of copper(II) hydroxide.
67
68
Name: _________________________________________ Date due: ___________________________
Chemical Reactions: Classification and Prediction of Products
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.Classify each reaction as either a double displacement (DD), single displacement (SD),
synthesis (SYN), decomposition (DEC), or combustion (COM) reaction.
• Predict and write the correct formulas for the products of each reaction.
• Balance the equation.
• Include the phase of each of the products.
1.
________
KCl(aq)
+
Pb(NO3)2(aq) 
2.
_______
C3H8O(l)
+
O2(g) 
3.
_______
AgNO3(aq)
+
Mg(s) 
4.
________
Ca(s) +
5.
________
H2CO3(aq) 
N2(g) 
69
70
Name: _________________________________________ Date lab performed: __________________
Partner(s) name:_________________________________ Date due: ___________________________
Chemical Reactions: Classification and Prediction of Products
Data: Record your observations and any test results.
Reaction #:
1.
2.
3.
4.
5.
6.
7.
71
Results:
•
•
•
•
Classify each reaction as either a double displacement (DD), single displacement (SD),
synthesis (SYN), decomposition (DEC), or combustion (COM) reaction.
Predict and write the correct formulas for the products of each reaction.
Write and balance the chemical equation.
Identify the physical state of each chemical as (aq), (s), (l), or (g).
1. Classification
_______
Balanced Chemical Equation:
2. Classification
_______
Balanced Chemical Equation:
3. Classification
_______
Balanced Chemical Equation:
4. Classification
_______
Balanced Chemical Equation:
5. Classification
_______
Balanced Chemical Equation:
6. Classification
_______
Balanced Chemical Equation:
7. Classification
_______
Balanced Chemical Equation:
_____Cu(OH)2 (s)  __________( ) + _____ CuO ( )
72
Double Displacement Reactions
Objectives: In this lab, students will
•
Identify six unknown solutions by mixing and observing their reactions
Skills: Upon completion of this lab, students should have learned to
•
Write molecular equations
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) to be read PRIOR to lab.
•
•
•
•
•
Evidence of Chemical Reactions
Aqueous Solutions and Solubility
Precipitation Reactions
Writing Chemical Equations
Acid-Base and Gas Reactions
Section 7.2
Section 7.5
Section 7.6
Section 7.7
Section 7.8
Introduction:
Double Displacement reactions are very common and occur in one of three main ways:
•
precipitation reactions
•
acid-base reactions
•
gas evolution reactions
In this lab, you will be given 6 “unknown” solutions. You will mix different combinations of two
solutions together and observe the results. Using these results and the predictions from the prelaboratory exercises, you will use logic to identify the 6 unknown solutions.
To help in your predictions of reaction outcomes, you will use your knowledge of precipitation and
solubility to help identify reactions. It will also help to know that:
•
Acid-base reactions often produce heat
•
Reactions between an acid and a compound containing CO32- ions will produce CO2 gas
73
Procedure:
You will be working with 6 unknown solutions. Each solution is one of the following:
1.5 M
1.0 M
0.1 M
1.0 M
3.0 M
0.1 M
H2SO4
K3PO4
Mg(NO3)2
Na2CO3
NaOH
SrCl2
Safety Precautions and Special Notes:
• Since you do not initially know the identity of the solutions, treat every solution as potentially
hazardous.
•
Write observations of the reactions.
•
Unless instructed otherwise:
o All reactions are performed in large test tubes.
o Use automatic dispensers to dispense solutions.
1. Place 2 mL (one pump) of solution A in a small test tube, then add 2 mL of solution B. Observe
the reaction and record your results in the data table. Dispose of your waste in the appropriate
waste container.
2. Repeat this process with 2 mL quantities in a small clean test tube for each of the following
combinations of solutions:
•
AB, AC, AD, AE, AF
•
BC, BD, BE, BF
•
CD,CE,CF
•
DE,DF
•
EF
74
75
s
7 x 10-13
16
s,d
10-8
K+
Ag+
Na+
Sr2+
Zn2+
30
0.01
20
0.8
10
28
26
4 x 10-3
ss
17
26
9
0.2
2 x 10-4
43
27
SO42−
i
i
11
6 x 10-4
47
0.02
0.03
1 x 10-5
i
i
i
i
2 x 10-3
i
26
i
PO43−
56
42
47
70
27
41
43
35
46
55
50
64
56
8
66
42
NO3−
83
64
64
3 x 10-6
59
58
62
0.07
1.1
65
s
68
68
63
s,d
I−
0.9
0.03
8
4 x 10-3
7.5
8
45
2 x 10-3
0.04
0.1
1
0.1
0.3
0.02
2
IO3−
4 x 10-4
1
52
d
53
2 x 10-3
11.3
0.02
1 x 10-5
3 x 10-4
3 x 10-4
i
0.16
4
47
1 x 10-4
OH−
i
0.12
47
4 x 10-3
39
42
50
7 x 10-6
i
i
14
4 x 10-4
25
CrO42−
79
35
26.4
2 x 10-4
25
35
45
1
70
42
35
50
43
26
27
31
Cl−
2 x 10-2
1 x 10-3
22
3 x 10-3
52
0.07
1.3
1 x 10-4
i
i
i
6 x 10-3
2 x 10-3
50
CO32−
82
50
48
8 x 10-6
40
50
62
0.8
s
56
54
3
59
51
43
s
Br−
25
27
32
1
70
40
75
31
7
s
20
26
42
60
ss
C2H3O2−
An arbitrary standard for solubility is that a compound is called soluble if greater than 1 gram dissolves in 100 grams of solution.
vs = very soluble, s = soluble, ss = slightly soluble, i = insoluble, d = decomposes
d
Mg2+
2 x 10-4
Cu2+
vs
4 x 10-4
Co2+
Li+
i
Ce3+
9 x 10-5
0.02
Ca2+
Pb2+
d
Ba2+
3 x 10-17
vs
NH4+
Fe3+
d
Al3+
S2-
SOLUBILITIES OF IONIC COMPOUNDS:
APROXIMATE # OF GRAMS OF SOLUTE PER 100 GRAMS OF SOLUTION
76
Name: __________________________________________Date due: ___________________________
Double Displacement Reactions
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.
1. For the three solutions 0.1 M Ba(OH)2, 1 M Na2CO3, and 2 M HCl,
a. Complete and balance each double displacement reaction (include states). .
b. Determine whether each reaction will produce a precipitate (ppt), a gas (gas), heat (heat), or if there
will be no observable reaction (nr), and write the correct abbreviation in the prediction matrix below.

1.
____Ba(OH)2(aq)+ ____Na2CO3(aq)
2.
____Ba(OH)2(aq)+
____HCl(aq) 
3.
____HCl(aq) +
____Na2CO3(aq)

Prediction Matrix:
Na2CO3
HCl
1
2
3
Ba(OH)2
Na2CO3
2. Assume that each of the three solutions, 0.1 M Ba(OH)2, 1 M Na2CO3, and 2 M HCl, is placed in a separate
bottle labeled either A, B, or C. Based on the results observed for the reactions, determine which solution
is in which bottle.
B
heat
A ____________________
C
gas
A
ppt
B
B ____________________
77
C ____________________
3. Predict the products and write the balanced chemical equation for each reaction.
If the reactants and products are all aqueous, no reaction took place, so write nr.
Label the physical states (g), (s), (aq), or (l) of each product. (Use the solubility table.)
_____ H2SO4(aq) + _____ K3PO4(aq) 
_____ Na2CO3(aq) + _____ K3PO4(aq) 
_____ NaOH (aq) + _____ K3PO4 (aq) 
_____ Mg(NO3)2 (aq) + _____ K3PO4(aq) 
_____ SrCl2(aq)
+ _____ K3PO4(aq) 
_____ Na2CO3(aq) + _____ H2SO4(aq) 
_____ NaOH(aq) + _____ H2SO4(aq) 
_____ Mg(NO3)2(aq) + _____ H2SO4(aq) 
_____ SrCl2(aq) + _____ H2SO4
_____ NaOH
(aq)
(aq) 
+ _____ Na2CO3
(aq) 
78
_____ Mg(NO3)2(aq) + _____ Na2CO3(aq) 
_____ SrCl2(aq) + _____ Na2CO3(aq) 
_____ Mg(NO3)2(aq) + _____ NaOH(aq) 
_____ SrCl2(aq) + _____ NaOH(aq) 
_____ SrCl2(aq) + _____ Mg(NO3)2(aq) 
4. For the six solutions you will use in the procedure of this lab, write whether each reaction will produce a
precipitate (ppt), a gas (gas), heat (heat), or if there will be no observable reaction (nr) based on the
reactions in question 3. You should predict 7 reactions will form precipitates, 1 reaction will produce heat,
1 reaction will produce gas, and 6 have no observable reaction. Copy this matrix to the data section.
Prediction Matrix:
H2SO4
Na2CO3
NaOH
Mg(NO3)2
nr
SrCl2
K3PO4
H2SO4
Na2CO3
NaOH
Mg(NO3)2
79
80
Name: __________________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date due: ___________________________
Double Displacement Reactions
Data: Record your observations and any test results.
•
•
Use the abbreviations precipitate (ppt), a gas (gas), heat (heat), no observable reaction (nr) to record
your observations in the matrix below.
Hint: 7 reactions should form precipitates, 1 reaction should produce heat, 1 reaction should produce
gas, and 6 should have no observable reaction.
Observation Matrix:
D
E
B
C
F
ppt
nr
heat
nr
gas
ppt
ppt
nr
ppt
nr
ppt
nr
ppt
nr
ppt
H2SO4
A
B
C
D
E
Prediction Matrix: (copied from the pre-lab page)
Na2CO3
NaOH
Mg(NO3)2
SrCl2
nr
K3PO4
H2SO4
Na2CO3
NaOH
Mg(NO3)2
81
Results: Use the data in the Observation Matrix and the Prediction Matrix to determine the identity of
solutions A through F.
Solution:
A ____________________________
B ____________________________
C ____________________________
D ____________________________
E ____________________________
F ____________________________
82
Lewis Structures
Objectives: In this activity, students will:
•
Draw Lewis dot structures with ionic bonds, covalent bonds, or both ionic and covalent bonds
Skills: Upon completion of this activity, students will have learned
•
•
To identify the number of valence electrons in main group elements
To draw structures of molecules, compounds, and ions
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
•
Electron Configurations and the Periodic Table
Covalent Lewis Structures: Electrons Shared
Writing Lewis Structures for Covalent Compounds
Resonance
Section 9.7
Section 10.4
Section 10.5
Section 10.6
Introduction:
Chemical compounds contain ionic bonds, covalent bonds, or both. Ionic materials, like sodium chloride
and calcium carbonate contain ions and are often easily recognized by the fact that they contain a metal and
a non-metal in their formula. Covalent compounds, like water and ethyl alcohol contain only non-metal
atoms in their formula. (Covalent compounds are also known as molecular compounds.) There are
exceptions to these simple rules for identifying compounds, but we will not focus on them here.
Valence electrons are the electrons that participate in bonding. Ionic bonds result from the gain or loss of
electrons, while covalent bonds are the sharing of electrons by two atoms. The reason for the covalent
sharing of electrons is so that the atoms can fill their valence shells. Useful chemical information can be
gained by looking at how the electrons are shared in covalent compounds. Drawing a Lewis structure is one
of the most common ways to illustrate this sharing.
83
Drawing Lewis Structures
Here are some simple guidelines for drawing correct Lewis structures.
Step 1: Identify Bonding Types
•
Write down the ions present. (Look for metals.) They will have an ionic bond between
them. If the ion is polyatomic, it will be made up of covalent bonds.
•
If no ions are present, all bonds will be covalent. (All atoms will be nonmetals.)
Step 2: Valence Electrons
•
Use the periodic table to count the total number of valence electrons for all atoms in the
molecule or ion. (Work on ions separately.)
•
Add one additional electron for each negative charge of an anion or subtract one for each
positive charge of a cation.
Step 3: Connect Atoms
•
Connect atoms with lines to represent bonds between atoms.
•
Hydrogen is always terminal.
•
The least electronegative atom is usually the central atom (unless otherwise noted).
•
Symmetrical structures are preferred.
Step 4: Assign Electrons to the Terminal Atoms and Fill Their Valence Shells
•
Place lone pairs (non-bonding electrons) of electrons around each terminal atom to
complete each atom’s octet (except for hydrogen).
Step 5: Recount and Adjust
•
If you have extra electrons, place around central atom as pairs.
•
If you don’t have enough electrons to complete an octet around the central atom, make
double and triple bonds by sharing electron pairs from the terminal atoms.
•
Always keep in mind the octet rule.
•
All atoms have an octet (hydrogen has a duet).
•
All valence electrons were used.
Step 6: Check
Step 7: Put Ions Together
•
Put brackets around the ion, and write the charge on the outside of the brackets.
•
Place the ions next to each other so that each positive ion is next to a negative ion.
84
Example #1,
Water, H2O
1. Identify Bonding Types. All of the atoms are nonmetals, so all bonds will be covalent.
2. Valence Electrons.
Atoms
H
H
O
Total
Valence Electrons
1
1
6
8
3. Connect Atoms. Draw simple and symmetrical structures. The possible arrangements of 3 atoms and 2
bonds for H2O are:
H H O
H O H
O H H
Hydrogen is always a terminal atom. Also, H-O-H is the most symmetrical, so use H-O-H.
4. Assign Electrons to the Terminal Atoms and Fill Their Valence Shells. Of the 8 valence electrons
we started with, 4 electrons were used in making the 2 bonds. That leaves 4 electrons to fill the valence
shells. Each hydrogen has 2 electrons in its valence shell (1 bond = 2 electrons) and is filled.
5. Recount and Adjust. The 4 remaining electrons are placed around the oxygen atom as 2 pairs of nonbonding electrons (also known as lone pairs).
..
H O
.. H
6.
Check. Two bonds and two lone pairs were used; the total electrons used were eight. Hydrogen has a
duet and oxygen has an octet.
7. Put Ions Together. No ions are present.
85
Example #2,
Formaldehyde, CH2O
1. Identify Bonding Types. All of the atoms are nonmetals, so all bonds will be covalent.
2. Valence Electrons.
Atoms
C
O
H
H
Total
Valence Electrons
4
6
1
1
12
3. Connect Atoms. Draw simple and symmetrical structures. Possible arrangements are
H C O H
O
C
H C H
H O H
Hydrogens are terminal atoms in all three structures. The last two structures are more symmetrical.
Generally, the atom that has the lowest electronegativity is the central atom. Carbon has a lower
electronegativity than O, so C is the better central atom. That means the middle structure is “better”.
O
H C H
4. Assign Electrons to the Terminal Atoms and Fill Their Valence Shells. Of the 12 valence electrons
we started with, 6 electrons were used in making the 3 bonds. That leaves 6 electrons to fill the valence
shells of the terminal atoms. The hydrogens are filled, so the remaining 6 electrons are placed around
the oxygen atom.
O
H C
5.
H
Recount and Adjust. In the above drawing, the hydrogens and oxygens are filled; the hydrogens have
duets and the oxygen has an octet. However, the carbon does not have an octet. One of the lone pairs
of oxygen is shared covalently to give both oxygen and carbon octets.
O
H C
H
6. Check. All valence shells are filled and the correct number of electrons has been used.
7. Put Ions Together. No ions are present.
86
Example #3,
Sodium Cyanide, NaCN
1. Identify Bonding Types. Sodium is a metal cation, and cyanide is a nonmetal anion. An ionic bond
exists between sodium and cyanide. Cyanide is polyatomic and made up of nonmetals, so there will be
+
covalent bonds within cyanide. Work on the ions separately (as Na and CN ) and then put them
together at the end.
+
2. Valence Electrons. Na
Atoms
Na
Positive ion charge
Total
Valence Electrons
1
-1
0
3. Connect Atoms. Sodium is not covalently bonded to any other atom, and has zero valence electrons
around it, so don’t draw any dots.
Na
There aren’t any covalent bonds, so skip those steps and just put brackets around sodium along with its
charge.
Na+
The sodium cation is finished, so now work on the cyanide anion.
2. Valence Electrons. CN
-
Atoms
C
N
Negative ion charge
totals
Valence Electrons
4
5
1
10
3. Connect Atoms.
C N
N C
The structure is linear, so either one of the above can be used.
87
4. Assign Electrons to the Terminal Atoms and Fill Their Valence Shells. Of the 10 valence electrons
we started with, 2 electrons were used in making the single bond. That leaves 8 electrons to fill the
valence shells. Lone pairs are placed on the carbon and nitrogen until the 8 electrons are used up.
C N
5.
Recount and Adjust. In the above drawing, neither the carbon nor the nitrogen is filled. Each atom
shares a lone pair so that each will have an octet.
C N
6
Check. All valence shells are filled and the correct number of electrons has been used. Brackets are
placed around anions and the overall charge is written outside of the brackets.
[ .. N
C .. ]-
7. Put Ions Together. Sodium cyanide has the following structure.
.
Na+ [ . N
88
C .. ]-
Name: __________________________________________Date due: ___________________________
Lewis Structures
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.
1. Define the following terms in your own words.
a. valence shell
b. valence electron
c. covalent bond
d. octet rule
e. resonance structure
2. Determine the number of valence electrons in an atom of each of these elements.
a. Determine the number of electrons needed to fill the valence shell.
Atom
Valence Electrons
C
O
H
Cl
N
S
89
# electrons needed
Draw Lewis structures for:
CH4
NH4
+
2-
SO4
90
Name: _______________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date due: ___________________________
Lewis Structures
A. Draw Lewis Structures of Molecules and Polyatomic Ions.
Draw the Lewis structure for each molecule or ion below. Follow the same process used in the
examples.
1. F2
2. O2
3. IO2
−
4. CH4
5. CO2
6. NH4
+
91
7. SO3
2−
8. C2H6
9. C2H4
10. ClO4−
92
11. Cl2CO (carbon is the central atom)
12. CH3OH (O-H bond)
13. NO2
+
14. N2
93
B.
Resonance Structures
Resonance structures occur when more than one correct Lewis structure can be drawn for a molecule or ion.
One resonance structure differs from another only by the placement of a double (or triple) bond. The
skeleton structure of atoms does not change. Ozone, O3, has been done as an example. The double arrow
indicates that the two structures are resonance structures.
O
O
O
O
Draw the indicated number of resonance structures for these formulas.
1. SO2
(2 structures)
2. NO2− (2 structures)
3. SeO2 (2 structures)
94
O
O
C.
Ionic Compounds
Draw the Lewis structure for each ionic compound below. Follow the same process used in the
examples.
1. CaO
2. MgSO4
3. Ca(OH)2
95
96
Gas Laws
Objectives: In this experiment, students will:
•
•
Determine the atomic mass of zinc through the combination of gas law concepts and
stoichiometry
Calculate the molar mass of natural gas using the Ideal Gas Law
Skills: Upon completion of this lab, students will have learned to:
•
•
•
•
Apply Dalton’s law of partial pressures to a gaseous mixture
Calculate the number of moles of a gas generated using the Ideal Gas Law
Use stoichiometry to convert moles of a product to moles of a reactant in a chemical reaction
Calculate the atomic mass of an element from an experimental mass and number of moles
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
Dalton’s law of partial pressures
Ideal Gas law
Stoichiometry
Section 11.9
Section 11.8
Section 8.1, Section 11.10
Introduction:
When zinc reacts with hydrochloric acid, the following single displacement reaction occurs:
Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2 (g)
In Part A of this experiment, a known mass of zinc is reacted with an excess of hydrochloric acid in
order to totally consume the zinc. It is possible to calculate the number of moles of hydrogen gas
produced in this reaction by rearranging the Ideal Gas Law.
𝑃𝑉
𝑅𝑇
=𝑛
Where n = number of moles of hydrogen gas produced, P = pressure due to hydrogen gas,
𝐿 𝑎𝑡𝑚
V = volume of hydrogen gas produced, R = ideal gas constant = 0.08206 𝐾 𝑚𝑜𝑙, and
T = temperature of hydrogen gas produced.
97
Because the hydrogen gas is collected over water, it also contains some water vapor. The total pressure equals
the pressure from hydrogen plus the pressure from water. The partial pressure due to hydrogen can be
calculated by rearranging Dalton’s law of partial pressures as follows:
PH2 = PT - PH2O
Where
PH2 = partial pressure due to hydrogen gas
PT = total pressure of gaseous mixture (atmospheric pressure)
PH2O = partial pressure due to water vapor
PH2 should be converted to atm using 1 atm = 760 mm Hg or 1 atm = 29.92 in Hg.)
Water Vapor Pressure at Various Temperatures
Temperature
˚C
Vapor Pressure
mm Hg
Temperature
˚C
Vapor Pressure
mm Hg
Temperature
˚C
Vapor Pressure
mm Hg
10
11
12
13
14
15
16
9.2
9.8
10.5
11.2
12.0
12.8
13.6
17
18
19
20
21
22
23
14.5
15.5
16.5
17.5
18.6
19.8
21.1
24
25
26
27
28
29
30
22.4
23.8
25.2
26.7
28.3
30.0
31.8
The balanced equation above tells us that for every mole of zinc consumed, one mole of hydrogen is
produced. Therefore, the number of moles of hydrogen produced is equal to the number of moles of
zinc consumed. Knowing the mass of zinc (in grams), the atomic mass of zinc can be calculated as
follows:
Molar Mass =
98
𝑚𝑎𝑠𝑠 (𝑔)
𝑚𝑜𝑙𝑒𝑠 (𝑚𝑜𝑙)
Procedure:
Safety Precautions: Hydrogen and air mixtures are extremely explosive. Keep all
open flames away from flasks containing hydrogen.
Determining the atomic mass of zinc. You will complete TWO trials of this procedure
1. Set up the apparatus as shown below. Fill the 500 mL Florence flask to the neck with water as shown
in the diagram, and put 25 mL of 6 M hydrochloric acid in the 250 mL Erlenmeyer flask.
Erlenmeyer Flask
Florence Flask
Collection Beaker
(400 mL or larger)
2. Remove both the stopper from the Erlenmeyer flask and the clamp from the rubber tubing between
the Florence flask and the beaker. Using a rubber bulb, apply compressed air to the tube inserted
through the stopper. When water begins to flow through the tubing, remove the rubber bulb and
clamp the tubing through which water is flowing. Remove the stopper gently from the Florence flask,
careful to keep the long glass tube under the water. Pour the water collected in the beaker back into the
Florence flask, and replace the stopper.
3. Obtain 0.900 g - 1.000 g of zinc. Remove the stopper from the Erlenmeyer flask, and transfer the
zinc into the hydrochloric acid in the flask. Immediately replace the stopper in the Erlenmeyer flask
and remove the clamp from the rubber tubing.
4. After the zinc has been completely consumed by the HCl, allow the gas in the flasks to cool to room
temperature (a few minutes). Raise the flask or the beaker until the levels of water in the flask and the
beaker line-up and then clamp the rubber delivery tubing before gently removing the beaker. Use a
500 mL graduated cylinder to measure the volume of water in the beaker. Use a glass thermometer to
measure the temperature of water in the beaker.
.
5. Dispose of the contents of the Erlenmeyer flask in the appropriate waste container and repeat the
experiments.
99
100
Name: __________________________________________Date due: ___________________________
Gas Laws
Pre-Lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.
1.
Perform each of the following pressure conversions:
a. 1.020 atm to mm Hg
b. 30.65 in Hg to mm Hg
2.
472 mL of H2 gas was collected over water when 1.256 g Zn reacted with excess HCl. The atmospheric
pressure during the experiment was 754 mm Hg and the temperature was 26 oC.
a) Write the balanced chemical equation for this reaction.
b) What is the water vapor pressure at 26 oC? ____________mm Hg
c) What is the partial pressure (in atmospheres) of dry hydrogen gas in the mixture? (Show work.)
d) Calculate the number of moles of H2 produced by this reaction using the ideal gas law. (Show work.)
e) Use the data from the experiment to calculate the experimental molar mass of zinc in g/mol.
101
f) What is the molar mass of zinc from the Periodic Table (known value)? __________ g/mol
g) Calculate the percent error for the experimental molar mass of Zn.
102
Name: __________________________________________Date lab performed: __________________
Partner(s) name: _________________________________Date due: ___________________________
Gas Laws
Data: Report all measurements in the correct number of significant figures and units. For all responses
requiring calculations, the mathematical setup must be shown.
Atomic Mass of Zinc
1. Mass of zinc used
____________ g
____________ g
2. Volume of water collected
____________L
____________L
3. Temperature of water collected
____________ °C
____________ °C
____________ K
____________ K
____________ in Hg
____________ in Hg
____________ mm Hg
____________ mm Hg
____________ mm Hg
____________ mm Hg
____________ mm Hg
____________ mm Hg
____________ atm
____________ atm
____________mol
____________mol
2. Atmospheric pressure
3. Partial pressure of water vapor at
temperature recorded above (See water
vapor pressure table in the introduction.)
4. Partial pressure of dry hydrogen gas
(show work)
7. Moles of hydrogen gas collected
103
8. Write the balanced chemical equation.
____________________________________________
9. Moles of zinc consumed
(Use the mole ratio of hydrogen gas to the
moles of zinc from the chemical equation.)
____________ moles Zn
____________ moles Zn
9. Calculated experimental molar mass of
zinc (MUST show work)
____________
____________
11. Known value for the molar mass of
zinc
____________
____________
____________
____________
12. Percent error (MUST show work)
Conclusion Question:
After starting with 0.343 g Al and excess HCl, the volume of H2 gas collected over water was 472 mL.
The atmospheric pressure was 754 mmHg and the temperature was 21oC.
1. Write the balanced chemical equation for this reaction.
2. Calculate the pressure in atm of dry hydrogen gas in the mixture. (Show work.)
104
3. How many moles of H2 were produced by this reaction? (Show work.)
4. Calculate the experimental molar mass of Al. (Show work.)
5. What is the percent error for the experimental molar mass?
105
106
Acid-Base Titrations
Objectives: In this lab, students will:
•
Determine the concentration of an unknown acid solution.
Skills: Upon completion of this lab, students will have learned to:
•
•
Use volumetric pipets and burets
Perform titrations between strong acids and strong bases
Textbook References: (Tro, Introductory Chemistry, 4th Ed.) Read these sections PRIOR to lab.
•
•
•
Molarity
Solution Stoichiometry
Acid-Base Titration
Section 13.6
Section 13.8
Section 14.6
Introduction:
Titration is a method used to determine the concentration of an acid solution or a base solution. Water
hardness of your city’s public water supply is probably determined using a titration. Often when performing
an acid-base titration, a base solution of known concentration is slowly added to the acid solution of
unknown concentration. A strong base reacts with a strong acid to form water according to the general
reaction below.
HA(aq) + MOH(aq)  H2O(aq) + MA(aq)
This reaction is often called a neutralization reaction because the strong base neutralizes the strong acid.
The chemist knows when all of the acid has been neutralized by either taking the pH of the solution during
the titration or by using an indicator. The indicator changes color as the pH of the solution becomes neutral.
By titrating—adding base—until the indicator changes color, the chemist knows when the equivalence point
is reached. The equivalence point is when the moles of OH added are exactly enough to neutralize the
+
moles of H present in the acid solution.
A common indicator to use is phenolphthalein because it changes from colorless in acidic solution to pink
when the solution becomes basic.
Glassware:
Titrations are performed using a buret. A buret is a long, narrow, calibrated tube which is designed to
deliver (TD) a quantity of a liquid. The stopcock at the bottom controls the flow of the liquid. Due to the
narrow top, a funnel is used to fill a buret.
107
Buret calibrations increase in value from the top to the bottom, the reverse of the graduated cylinder.
Graduated Cylinder
Buret
When reading the buret, you can hold a colored index card or something else behind the buret to help you
see the meniscus better. Read the bottom of the meniscus. Be sure your eye is at the level of meniscus, not
above or below. The readings are recorded to ±0.02mL.
You will use a volumetric
pipet to measure out the
unknown acid. Be sure to
always use two hands (one on
the pipet bulb and your
dominant hand on the pipet
itself).
108
Procedure: The instructor will demonstrate the set-up and use of the buret to perform a titration.
Part A. Determination of acid concentration using titration.
1. Obtain approximately 50 mL of the HCl solution of unknown concentration in a small beaker.
Record the ID#.
2. Use a volumetric pipet to transfer exactly 10.00 mL of the HCl solution into a clean Erlenmeyer
flask. Add 2 drops of phenolphthalein indicator to the HCl solution.
3. Obtain approximately 100 mL of the sodium hydroxide solution. Record its concentration.
4. Clean and set-up a buret according to the directions given by your instructor. Place a funnel on the
top of the buret, and carefully pour 45-50 mL of NaOH into the buret. Remove the funnel.
5. Place a large waste beaker underneath the buret tip, and open the stopcock to let ~ 5 mL of NaOH
run down to fill the buret tip. Make sure air bubbles aren’t trapped around the stopcock.
6. Record the initial buret reading of the NaOH solution to the nearest 0.02 mL. See tips for reading
the buret correctly in the introduction.
7. Perform a “rough” titration to get an idea of what color change to expect at the endpoint and also to
get a rough idea of the volume of NaOH needed to get to the equivalence point by adding
approximately 2 mL aliquots of NaOH to the HCl solution. Swirl the flask to stir. Continue to add
NaOH in 2 mL increments until the phenolphthalein indicator remains pink for more than 20
seconds. Record the final buret reading of NaOH in the buret. Dispose of the contents of the flask in
the sink with lots of water.
8. Measure and transfer another 10.00 mL of HCl to a clean flask. Check to see if you need to refill the
buret. Record the initial buret reading of NaOH and titrate until you have added about 2 mL less
than the total amount added in the rough titration. Now add NaOH drop by drop until the HCl
solution just turns light pink and remains for at least 30 seconds while you swirl the flask.
(Proper technique here is essential. You want one drop of added base to cause the acid solution to
turn from colorless to light pink. If you add too much base and “overshoot” the equivalence point
you will need to start over.)
9. Repeat the titration until you have gotten 5 “good” titrations without overshooting the equivalence
point. (Ask your instructor to help you define “good”.)
10. Mix all acid and base solutions together and pour down the sink with lots of water.
11. Clean up.
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Name: __________________________________________Date due: ___________________________
Acid-Base Titrations
Pre-lab Questions: Read the relevant textbook sections and the entire lab before answering pre-lab
questions.
1. Define “equivalence point” in your own words, and describe how you will know when it has been
reached in this lab.
2. Write the balanced chemical equation for the reaction of a sodium hydroxide solution with:
a. hydrochloric acid
b. sulfuric acid
3. What is the level of the liquid in the buret? Use the correct number of significant figures to reflect the
precision of the instrument?
____________mL
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4. 17.65 mL of a 0.110 M sodium hydroxide solution is needed to titrate 25.00 mL of a hydrochloric
acid solution to the equivalence point.
a. Write the balanced equation for the neutralization reaction.
b. How many moles of sodium hydroxide are used for the titration?
c. How many moles of hydrochloric acid reacted with sodium hydroxide?
d. What is the molar concentration (Molarity) of the hydrochloric acid solution?
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Name: _________________________________________ Date lab performed: __________________
Partner(s) name:_________________________________ Date due: ___________________________
Acid-Base Titrations
Data:
A. Titrations
Unknown acid ID #
_____________
Concentration of NaOH
_____________
Record your data to the correct significant figures.
Rough
Trial 1
Trial 2
Trial 3
Trial 4
Initial buret
reading
(mL)
Final buret
reading
(mL)
Volume of
NaOH added
(mL)
Calculated acid
concentration
(M)
Calculations:
1. Calculate the concentration of acid for each trial. Show one sample calculation below and
remember to use the correct number of significant figures and units.
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Trial 5
2. Calculate the average concentration of the acid. (Use your three closest values to calculate your
average.)
Average acid molar concentration ________________mol/L
3. Calculate the % error. Your instructor will provide the known acid concentration.
Known acid molar concentration __________________mol/L
Percent error = __________ %
Conclusion questions:
1. The equivalence point of a titration is overshot! How will this error affect the calculated
concentration of acid? (Will it be too high or too low?) Explain.
2. Describe at least one modification you would make to improve the accuracy of your experimental
data.
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Spectroscopy: Determination of
Concentration Using Beer’s Law
Objectives: In this lab students will
•
Determine the concentration of copper(II) sulfate unknown
Skills: Upon completion of this lab, students should have learned to:
•
•
•
Prepare solutions of known concentration by dilution from a stock solution.
Use a spectrophotometer to measure the absorbance solutions
Draw a Beer’s Law graph (calibration graph) of absorbance versus concentration
Textbook References: (Tro, Introductory Chemistry, 3rd Ed.) to be read PRIOR to lab.
•
•
Specifying Solution Concentration: Molarity
Solution Dilution
Section 13.6 p. 457
Section 13.7 p. 461
Introduction:
Some chemical solutions appear colored because certain wavelengths of light are absorbed by the
solute. A red solution looks red because it absorbs light at wavelengths other than red and then only
the red wavelengths pass through to your eye. You can easily guess that the darker the red color of
the solution, the higher the concentration of light absorbing material in the solution. This idea of light
absorbance being proportional to concentration is known as Beer’s Law, and written in equation form
is
A = εbc
A = absorbance
ε = is the absorption constant (different for each chemical)
b = the path length the light travels through the solution
c = the concentration
In this lab we are only interested in the relationship between absorbance and concentration (A and c),
and ε and b are constants during our experiment so we can say that the Beer’s Law becomes
A = kc
(Where k is a constant . . . the product of ε x b = k). And now it is obvious that absorbance, A, is
directly proportional to concentration, c. Because absorbance and concentration are directly
proportional, a graph of absorbance versus concentration should be a straight line. The data for the
graph is collected by measuring the absorbance of several solutions with known concentrations. This
type of graph is known as a Beer’s law Plot (or calibration plot). We can use the Beer’s Law plot to
determine the concentration of an unknown solution. One of the many practical uses of this process is
to determine the amount of a contaminant (like nitrate ion, sulfate or lead ion, to name a few) in a
drinking water sample.
Absorbance will be measured using a spectrophotometer.
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A detector within the spectrophotometer measures the relative % of light transmitted through the
sample and converts the % transmittance to an absorbance value using the equation below.
A = 2 – log (% Transmittance)
Procedure: The instructor will demonstrate the use of the spectrophotometer.
Part A. Visible wavelengths Introduction (Optional; done only in class)
1. Insert an empty spectrophotometer tube containing a strip of white paper into the sample
holder; leave the sample holder cover open. Set the wavelength of the spectrophotometer to
one of the wavelengths listed below and look down into the tube and record what color you see
at each of these wavelengths - 660nm, 600 nm, 560 nm, 440, nm, and 400 nm.
Part B. Preparing the CuSO4 solutions of known concentration
2. Turn on the spectrophotometer. The machine takes about 10 minutes to warm up.
3. Obtain about 30 mL of 0.40 M CuSO4 stock solution in a small beaker.
4. Use a 10 mL graduated pipette and a 10 mL volumetric flask to prepare 10 mL of a 0.08M
solution according to the volumes you calculated in the prelab. Transfer this solution to a
clean, dry test tube labeled #1.
5. Repeat for solutions of 0.16M, 0.24M, 0.32M and 0.40 M CuSO4, placed in test tubes # 2-5,
respectively. Thoroughly mix each solution with a stirring rod. Clean and dry the stirring rod
between stirrings.
Part C. Measuring the absorbance of the CuSO4 solutions
6. Set the wavelength of the spectrophotometer to 635 nm.
7. Prepare a “blank” by rinsing a cuvette (a special “test tube” for spectrophotometers) with
approximately 1 mL of deionized water. Repeat the rinse another time. And then cuvette
about 2/3 full with deionized water. Use the “blank” to zero the spectrophotometer according
to your instructor’s directions. Do not pour out the water in this cuvette, you will use this blank
to zero the machine before measuring the absorbance of every sample.
8. Prepare “sample” for measurement by taking another cuvette and rinsing two times with
approximately 1mL of the 0.08 M CuSO4 solution and then filling to approximately 2/3 full.
Zero the machine using the blank and then place the sample cuvette in the machine and
measure and record the absorbance. Repeat for the other four solutions of known
concentration.
9. Obtain an unknown. Record its ID # and the measure and record its absorbance.
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Name:
Date due:
Spectroscopy: Determination of Concentration Using Beer’s Law
Pre-lab Questions: Read the relevant textbook sections and the entire lab. Then complete the
following:
1. Use the internet or library to find two practical uses of copper(II) sulfate. Reference the website.
2. Use the internet or library to find two specific health hazards associated with copper(II) sulfate and
outline the safe handling practices of copper(II) sulfate. Reference the website.
3. Use the dilution equation ( M1V1 = M2V2 ) to calculate the volume of 0.40 M copper(II) sulfate
solution needed to prepare 10.00 mL of these 4 solutions - 0.08M, 0.16M, 0.24M, 0.32M. Show your
work.
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118
Name:
Partner(s) name:
Date lab performed:
Date due:
Spectroscopy: Determination of Concentration Using Beer’s Law
Data Table: Record your observations and any test results.
A. Visible wavelengths Introduction (Optional)
Wavelength (nm)
660
Observed Color
600
560
440
400
Insert an empty spectrophotometer tube containing a strip of white paper into the sample holder; leave
the sample holder cover open. Set the wavelength of the spectrophotometer to one of the wavelengths
listed below and look down into the tube and record what color you see at each of these wavelengths 660nm, 600 nm, 560 nm, 440, nm, and 400 nm.
Solution
Concentration
(mol/L)
1
2
3
4
5
Unknown ID # ______
XXXXXXXXXX
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Absorbance
Calculations:
Make a properly labeled Beer’s Law plot of absorbance versus concentration (absorbance on the yaxis) for the five solutions of known concentration. Draw the best fit line through these five points.
Use this graph to determine the concentration of the unknown. Record your result in the results
section below. Use either graph paper or a computer program to make the graph. Attach the graph to
the back of this lab report.
Results: Unknown ID # _______________ Concentration _______________
Conclusion Questions:
1. Use your graph to predict the absorbance of a 0.20 M CuSO4 solution.
2.
Optional Bonus: Calculate the absorbance of a solution that has a % transmittance of 75%.
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