General Chemistry Laboratory Experiments
CHM 165 CHM 166 and CHM 175
Scott Community College
By:
R. Ford and Dr. J. Tayh
2015 Edition
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Student Copy of Physical Sciences Laboratory Safety Contract
General Safety Rules:
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Safety goggles will be worn at all times in the laboratory.
Shoes that cover your entire foot must be worn in the laboratory (NO sandals).
Long hair will be tied back when working with open flames.
Long pants must be worn in the laboratory (no shorts, no skirts, or dresses). Exceptions will be made for
religious beliefs preventing you from wearing pants.
Long, loose, flowing clothing should not be worn.
Contacts are not recommended!! If you must wear contacts, make sure your instructor and lab partner(s)
are aware that you are wearing contacts.
Students are NEVER allowed in the chemical storage areas.
Students are never to be in the laboratory without their instructor or the laboratory assistant present.
Absolutely NO food or drinks allowed in the laboratory.
Never perform unauthorized experiments.
Never remove any chemicals from the laboratory.
Never return chemicals to the reagent (stock) bottles.
Defense Against Accidents:
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I will act in a responsible manner at all times in the laboratory.
I will understand that an unprepared worker is an unsafe worker.
I will always read assignments before coming to class.
I will wear an apron and gloves when instructed to do so.
I may wear an apron and gloves any time I wish.
I will never work alone in the laboratory.
Poisons, Chemical Burns, and Toxicity:
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I will immediately flush any spilled chemical off myself for a minimum of fifteen minutes.
I will flush my eyes for a minimum of fifteen minutes in the event I get any chemicals in them.
I will clean my glassware when finished with it and return it to its proper location.
I will return all laboratory equipment to its proper location. If I am unsure, I will ask.
I will never taste any chemical and will smell chemicals only when instructed to do so.
Fires, Burns, and Hazards:
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I know the location of the fire extinguisher, fire alarms, fire blanket, safety shower, eyewash station, and
first aid kit.
I will know how to use all safety equipment with the exception of the fire extinguisher.
I will only operate the fire extinguisher if trained how to do so.
I will pull the fire alarm and evacuate the building immediately in the event of a fire.
Exception: very small fires may be extinguished by smothering them.
I know that hot and cold items look the same.
I will exercise extreme caution when dealing with glass tubing and rubber stoppers.
I will never heat anything in the microwave without proper approval.
Biological Hazards:
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I will report any and all accidents to my instructor immediately.
I will treat any source of blood as a biohazard and will inform my instructor immediately. My instructor
will tell me how to handle the incident.
This page has been intentionally left blank for printing purposes.
Please Complete and Return to Instructor
Copy of Physical Sciences Laboratory Safety Contract
General Safety Rules:












Safety goggles will be worn at all times in the laboratory.
Shoes that cover your entire foot must be worn in the laboratory (NO sandals).
Long hair will be tied back when working with open flames.
Long pants must be worn in the laboratory (no shorts, no skirts, or dresses). Exceptions will be made for religious beliefs preventing
you from wearing pants.
Long, loose, flowing clothing should not be worn.
Contacts are not recommended!! If you must wear contacts, make sure your instructor and lab partner(s) are aware that you are
wearing contacts.
Students are NEVER allowed in the chemical storage areas.
Students are never to be in the laboratory without their instructor or the laboratory assistant present.
Absolutely NO food or drinks allowed in the laboratory.
Never perform unauthorized experiments.
Never remove any chemicals from the laboratory.
Never return chemicals to the reagent (stock) bottles.
Defense Against Accidents:






I will act in a responsible manner at all times in the laboratory.
I will understand that an unprepared worker is an unsafe worker.
I will always read assignments before coming to class.
I will wear an apron and gloves when instructed to do so.
I may wear an apron and gloves any time I wish.
I will never work alone in the laboratory.
Poisons, Chemical Burns, and Toxicity:





I will immediately flush any spilled chemical off myself for a minimum of fifteen minutes.
I will flush my eyes for a minimum of fifteen minutes in the event I get any chemicals in them.
I will clean my glassware when finished with it and return it to its proper location.
I will return all laboratory equipment to its proper location. If I am unsure, I will ask.
I will never taste any chemical and will smell chemicals only when instructed to do so.
Fires, Burns, and Hazards:








I know the location of the fire extinguisher, fire alarms, fire blanket, safety shower, eyewash station, and first aid kit.
I will know how to use all safety equipment with the exception of the fire extinguisher.
I will only operate the fire extinguisher if trained how to do so.
I will pull the fire alarm and evacuate the building immediately in the event of a fire.
Exception: very small fires may be extinguished by smothering them.
I know that hot and cold items look the same.
I will exercise extreme caution when dealing with glass tubing and rubber stoppers.
I will never heat anything in the microwave without proper approval.
Biological Hazards:


I will report any and all accidents to my instructor immediately.
I will treat any source of blood as a biohazard and will inform my instructor immediately. My instructor will tell me how to handle
the incident.
Acknowledgement
I, ____________________, have read and understand these safety rules. I will abide by these safety rules and any additional safety rules,
written or verbal, provided by my instructor, the physical sciences laboratory assistant, or college administration.
__________________________________
Student Signature
____________________________
Date
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Name____________________________________
Experiment 1: Measurements in Chemistry
Repeat each experiment three times:
1. Weigh 25mL graduated cylinder empty. Fill it with water up to 13mL using 5mL pipette,
and weigh it again. Find the mass of water by the difference. Repeat this measurement
three times. Find the average mass.
Trial one=
Trial two=
Trial three=
Average mass=
2. Measure the volume in cm3 of a block of wood that will be given to you.
Trial one=
Trial two=
Trial three=
Average volume=
3. Measure the volume of a 250mL Erlenmeyer flask when it is filled with water to the top
by using a 100mL graduated cylinder. When the Erlenmeyer flask is filled to the very top,
you cannot tell what volume of water you have directly because Erlenmeyer flasks do
not have much accuracy in their graduations. You must pour the water into a more
accurate measuring device (a graduated cylinder) to figure out the water’s volume.
Repeat this process three times. Average your results.
Trial one=
Trial two=
Trial three=
Average volume=
4. Using a pipet, transfer 10mL of water to a pre-weighted empty beaker. Weigh again,
and then find the mass of water by the difference.
Trial one=
Trial two=
Trial three=
Average mass=
5. Do the following conversion, and report your answer in the scientific notation only, with
the right number of significant figures: For water only you can assume that; (1g = 1mL =
1cm3)
a) 0.90Kg to mg
b) 10.0in to cm
c) 50.0Kg water to L water
d) 25.0cm to in
e) 200.0mL to microliters
f) 55555mg to Kg
g) 0.020KL to microliters
h) 2.5g/cm3 to Kg/L
Name____________________________________
Experiment 2: Density
Your goal is to find the density of the following items:
1.
3.
Salt water solution
2. Vinegar
Trial one
Trial one
Trial two
Trial two
Trial three
Trial three
Average
Average
Irregular shaped piece of metal
Number________
4. Block of wood
Number ________
Trial one
Trial one
Trial two
Trial two
Trial three
Trial three
Average
Average
Significant digits will be important here. Show all of your calculations clearly. Each student
in the group should do one of the trials then give an average of the three trials for each
experiment. Record all measurements to as many digits as the measuring device will allow.
Questions:
1. Express the density of the block of wood (Part 4) in lb/ft3.
2. If the volume of a solution is 0.230L and its mass is 0.210kg, what is the density in
g/mL?
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Name____________________________________
Experiment 3A: Melting Points
1.
For practice, you will accurately determine the melting point of any three known compounds made
available by the instructor. Compare your determined melting points with those in the chart below.
2.
You will determine the melting points of two pure unknowns. Your unknown compounds will be on the
following table:
COMPOUND
MELTING POINT
(Celsius Degrees)
Pimelic Acid
103-105
Catechol
104-106
Azelaic acid
105-106
*Rescorcinol
109-110
*Acetanilide
113-114
Mandelic acid
117-118
Succinic anhydride
118-120
*Benzoic acid
121-122
*2-Naphthol
121-122
Trans-Cinnamic acid
132-133
*Benzamide
132-133
Maleic acid
134-136
Malonic acid
135-137
Benzoin
136-137
Anthranilic acid
145-147
Cholesterol
148-150
*Adipic acid
152-153
*Citric acid
153-155
*Salicylic acid
156-158
Benzanilide
160-161
Itaconic acid
163-165
Sulfanilamide
164-166
To help you determine the identity of an unknown, you may form a mixture of your unknown and a known. For
example, if you suspect your unknown (which experimentally melts at 151-152) is adipic acid, you may mix your
unknown with adipic acid. If the mixture melts at 151-152, your unknown can be assumed to be adipic acid. If
your unknown is not adipic acid, the mixture would be expected to have a lower melting point. (Shut off
thermometers when you finish.)
1) Unkown #_____
is ____________________
2) Unkown #_____ is _____________________
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Name____________________________________
Experiment 3B: Boiling Points
1. Your instructor will give you two unknown liquids. You will determine the boiling point of
each using the electric melting point device (from Part A). Your instructor will explain
how to use the device.
2. Be sure that your electronic thermometers are shut off before you turn them in! Be sure
that your bench top is clean.
UNKNOWN COMPOUND
BOILING POINT
(Celsius Degrees)
Bezaldehyde
Mesitylene
Anisole
1-Butanol
Amyl alcohol
2-Butanol
178.1
165
155
118
136.5
99.5
1) Unkown #_____ has a boiling point of ______ and is ____________________
2) Unkown #_____ has a boiling point of _______ and is _____________________
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Name____________________________________
Experiment 4: Naming Compounds
A. Name the following compounds:
1. MgCl2
2. AlCl3
3. CaS
4. Na2O
5. Fe(OH)3
6. Mg3(PO4)2
7. (NH4)2S
8. Na2Cr2O7
9. SO3
10. N2O4
11. CCl4
B. Write the formula of the following compounds:
1. Copper(II) oxide
2. Boron trichloride
3. Silver cyanide
4. Magnesium carbonate
5. Calcium bicarbonate
6. Ammonium sulfate
7. Sulfur dioxide
8. Barium acetate
Name____________________________________
Experiment 5: Stoichiometry
The synthesis of tetraamminecopper(II) sulfate monohydrate
Reaction:
CuSO4 +
4NH3
+H2O
[Cu(NH3)4]SO4•H2O
Procedures:
1.
Weigh 0.20 g of solid copper sulfate then dissolve it in 1.5mL water in one well of a 24well plate. Draw all of the copper sulfate solution into a pipet bulb.
2. Add 1.0mL of concentrated ammonium hydroxide (source of NH3) provided in the small
bottle to a second well of the 24-well plate. Draw all of this solution into the pipet
containing the copper sulfate solution. Tap the bulb of the pipet with your finger to mix
the solution thoroughly. The solution will turn dark blue due to the formation of
Cu(NH3)42+ ion.
3. Add 1.75mL of ethanol to one of the wells, and then draw all of it into the pipet from
Step 2. Mix the solution and let it stand for 5 minutes. A precipitate should form in the
pipet bulb.
4. Centrifuge the pipet for 1 minute then remove the liquid only. Draw another 1.75 mL of
clean ethanol into the pipet containing the solid. Tap the bulb of the pipet with your
finger and centrifuge for 1 minute. Discard the liquid.
5. Using a pair of scissors, cut the top of the pipet bulb. Weigh a piece of filter paper.
Scrape the solid out of the pipet bulb onto the preweighed filter paper and allow the
solid to air dry or place it in the oven. After it has dried, weigh the solid.
Questions:
1) Determine theoretical yield of the product [Cu(NH3)4]SO4•H2O
2) Determine the percent yield of the product [Cu(NH3)4]SO4•H2O, assuming that the
CuSO4 is the limiting reactant.
Name____________________________________
Experiment 6: Household Chemical Reactions
Note: Different groups may start with different parts of the experiment. It os important to
note, however, that Parts 4 and 8 take the longest and you may want to do them first.
1.
Do the Alka-Seltzer reaction. Add a little citric acid to a little baking soda followed by a few drops of
water. Write down your observations and balance the equation.
C6H8O7
2.
+
NaHCO3
Na3C6H5O7
+
H2CO3
Take 2 pieces of aluminum (one piece of foil and one pop can piece). Put a drop or two of 1M
hydrochloric acid on the piece of foil. Repeat for the inside surface of the piece of pop can. Then, put a
drop or two of 6M hydrochloric acid in a separate spot on the foil. Repeat for the inside surface of the
can. Describe what you see. The 6M is a stronger acid than the 1M. Are there differences in how they
react? Balance the reaction of HCl with Al. Is there a reason why we do not store orange juice in a plain
aluminum container?
Al
+
HCl
AlCl3
+
H2
3.
Oil of wintergreen can be made by adding a pinch of salicylic acid to three mL methanol and then
warming in hot water for a while. Add about 1mL of sulfuric acid (concentrated) little by little during the
heating. If solid forms as you add the acid, allow it to re-dissolve. After about twenty minutes, you should
be able to detect the wintergreen aroma. This is how a commercial food additive can be made.
4.
Make fools gold by adding some of the provided ferrous sulfate to some of the provided sodium sulfide.
You should observe a precipitate. Write a balanced reaction and describe what you see.
Experiment 6 continued
5.
1
Take a piece of drywall (CaSO4• H2O). Drywall is a hydrate. Get it wet. Describe what happens. The
2
product is a dihydrate. Balance the following equation. What happens if you allow the drywall to dry?
𝟏
CaSO4•𝟐H2O
+
H2O
CaSO4•2H2O
6.
Heat the samples of a copper solution in a flame. Use a wire to hold a drop of the solution in the flame.
Describe what you see. Repeat for a lithium solution. Suggest how fireworks are made.
7.
Take a piece of copper metal. Drop it into silver nitrate solution. Observe over a period of twenty
minutes. Describe what you see. Balance the reaction. What is the new solid? Why did the liquid turn
colors?
Cu(s) +
AgNO3(aq)
Cu(NO3)2(aq) +
Ag(s)
Name____________________________________
Experiment 7: Synthesis of Alum
Alum can be found in the supermarket spice section. It has a number of uses in the kitchen.
It was used in pickling. It was used to help keep dyed shirts colorfast. It currently is used
extensively in the paper industry as a “sizing” agent. Raw paper is very porous, and the pores
are filled to make the paper smooth so that the ink will not bleed very much. Alum is used to
help adhere these fillers to the paper.
Alum is made from aluminum:
Step 1: 2Al(s)
Step 2: 2H2O
+
2KOH + 2H2O
+2KAlO2
Step 3: 2Al(OH)3(s)
Step 4: Al2(SO4)3
+
+
H2SO4
3H2SO4
+K2SO4 + 24H2O
2KAlO2
+
3H2(g)
2Al(OH)3(s)
Al2(SO4)3
+
+
K2SO4
6H2O
2KAl(SO4)2•12H2O (alum)
Procedure:
1. Weigh about 0.100 g of aluminum foil. Shred the piece of foil and place it in a perweighed
plastic medicine dose cup and record the actual mass of the aluminum. Fill a polystyrene
cup, almost to the top, with hot water (to be used as a heat source).
2. Add 2 dropper bulbs of 1.4M KOH to the shredded aluminum. Float the small cup in the
hot water bath (swirling the mixture occasionally). What reaction is taking place here?
3. When the aluminum has completely dissolved, filter (gravity filtration) the solution to get
rid of any insoluble materials in the solution. Rinse your medicine cup out with distilled
water. Transfer the filtrate into your plastic medicine cup and return the cup to the hot
water bath.
4. Quickly add half a bulb of 9M H2SO4. Gently swirl until the cloudy solution becomes clear
again. Pour the hot water in your polystyrene cup down the drain. Fill the polystyrene
cup with ice. Add some water to the ice to create an ice bath for cooling purposes. Float
the medicine cup on the ice water to induce crystallization. If crystallization does not start
within 15 minutes, your instructor may provide a seed crystal or other advice.
5. Filter the crystals using suction filtration. Rinse the crystals with distilled water. Weigh a
piece of paper. Place the dried crystals on the paper and weigh them.
Questions:
1. Write an overall reaction for the production of alum from aluminum.
2. Calculate the theoretical yield of alum. Assume aluminum is the limiting reactant.
3. Calculate the percent yield.
Name____________________________________
Experiment 8: Ionic Reactions
A reaction chart follows this page. Run all the reactions on the reaction chart. One drop of
each component is sufficient. Your goal will be to observe any reactions that may occur and to
write down your observations. Some reactions take a long time before they react noticeably.
Questions:
1. Write the formula unit (molecular), total ionic, and net ionic equations for 5 reactions
that react to form a precipitate or a gas.
a)
b)
c)
d)
2. Run all the reactions for one unknown. Based on your observations for all the knowns,
try to identify what cation and what anion is in each of your unknowns (remember to
write down the unknown numbers in your report sheet).
Unkown#
Cation
Anion
Reaction Chart
AlCl3
NH3
BaCl2
CuSO4
FeCl3
HCl
Pb(NO3)2
HNO3
KI
AgNO3
Na2CO3
NaOH
Na3PO4
H2SO4
Na3PO4
NaOH
Na2CO3
AgNO3
KI
HNO3
Five
Unknowns
Pb(NO3)2
HCl
FeCl3
CuSO4
Test
BaCl2
NH3
AlCl3
Single
1of 14
Cation
Anion
Area
H2SO4
General Chemistry (CHM 165 and CHM 175) Name____________________________________
Experiment 9: Degree of Ionization
The degree of ionization can be indicated by the electrical conductivity of a solution.
Procedures:
This is a demonstration experiment by the instructor.
Sample
Distilled water
Tap water
Strength of electrolyte
Ionization equation
XXXXXXXXXXXXXXX
XXXXXXXXXXXXXXX
H2SO4
HCl
HNO3
NaOH
NH4OH
CH3COOH
Mixture of NH4OH and
CH3COOH
Concentrated CH3COOH
NaCl
CuSO4
CaCl2
Sucrose
Questions:
1. Why is tap water a good conductor while distilled water is not?
2. Why is a solid salt a nonconductor?
3. Write a balanced ionization equation for K3PO4.
Name____________________________________
Experiment 10: Oxidation Reduction Reactions
Use the oxidation/reduction chart on the next page. Be sure to cover the chart with a plastic
overlay provided by your instructor.
Perform one trial with each of the appropriate solutions as given on the chart below.
Some of the metals will react (sometimes very slowly). Record your observations below.
H2O
HCl
AgNO3
Cu(NO3)2
Pb(NO3)2
Zn(NO3)2
Zn
Al
Sb
Cu
Mg
Fe
Pb
Questions:
1. Write a balanced chemical equation for the reactions that occur (one equation for each
metal). Identify what is being oxidized, what is being reduced, the oxidizing agent, and
the reducing agent.
2. Arrange all of the metals and hydrogen in an activity series based on your observations.
At the top of your series, you should have the metal that is most active.
3. Does your activity series agree with the one in the textbook? If not, try to explain why.
Name____________________________________
Experiment 10 continued
H2O
Zn
Al
Sb
Cu
Mg
Fe
Pb
HCl
AgNO3
Cu(NO3)2
Pb(NO3)2
Zn(NO3)2
Name____________________________________
Experiment 11: Analysis of Vinegar
Procedure:
1. Rinse the microburet (5mL graduated pipet) with distilled water and then rinse it with
the standardized NaOH solution. Fill the microburet with the standardized NaOH
solution. Record the molarity of the NaOH solution._________________________
2. Transfer 2mL of the vinegar with a pipet to a preweighed plastic cup. Record the mass
of the vinegar and the cup together. Calculate the mass of the vinegar and report it in
the table below. Add one drop of phenolphthalein indicator to the vinegar.
Trial 1
Trial 2
Trial 3
Average Mass
Mass of empty
cup
Mass of vinegar
and cup
Mass of vinegar
3. Titrate the vinegar solution with the standardized NaOH from the microburet while
swirling. The end point of the reaction is indicated by the appearance of a pink color.
Repeat the titration three times. Record the volume of NaOH that has been added to
vinegar.
Trial 1
Trial 2
Trial 3
Average mL of
NaOH added
Questions:
1. Determine the mass percent of acetic acid in the vinegar and show your calculations.
2. How many moles of NaOH are needed to neutralize 20.0mL of 0.15 M HCL solution?
3. You are adding 0.12M hydrobromic acid to calcium hydroxide solution. If it takes 25mL
of hydrobromic acid to neutralize the calcium hydroxide solution, how many grams of
calcium hydroxide were in the solution?
Name____________________________________
Experiment 12: Calorimetry and Hess’s Law
HCl
+
NaOH(aq)
H2O
+
NaCl
Δ HN
NaOH(s)
NaOH(aq)
Δ HS
__________________________________________________
HCl + NaOH(s)
H2O + NaCl
Δ HR
Hess’ Law: ΔHR = ΔHS +
ΔHN
Procedure:
A. Heat of Neutralization ΔHN
1. Measure 100.0 mL of 1.0 M HCl in a graduated cylinder and transfer it to a
Calorimetry cup, then measure its temperature after 3 minutes (Initial Temp.).
____________
2. Measure a 100.0 mL of 1.0 M NaOH solution in a graduated cylinder.
3. Pour the NaOH solution into the Calorimetry cup that has the HCl and record the
highest temperature (Final Temp.). ________
4. When you finish, pour the solutions down the drain.
B. Heat of Solution ΔHS
1. Measure 200.0 mL of distilled water into the Calorimetry cup and record its
temperature after 3 minutes (Initial Temp). ___________
2. Weigh 4.12 g of solid NaOH and dissolve them completely in the 200.0 mL of distilled
water, then record its highest temperature (Final Temp.). _________
3. When you finish, pour the solutions down the drain.
C. Heat of Reaction ΔHR
1. Measure 200.0 mL of 1.0 M HCl solution into the Calorimetry cup and record its
temperature after 3 minutes (Initial Tamp.). __________
2. Weigh 4.12 g of solid NaOH and dissolve them completely in the 200.0 mL HCl
solution, then record the highest temperature (Final Tamp.). __________
3. When you finish, pour the solutions down the drain.
Name____________________________________
Experiment 12 continued
Calculations:
ΔH(q) = (specific heat) X (Mass) X (Ti-Tf)
For Part A and Part C use specific heat = 4.02J/g.c
For Part B use specific heat = 3.93 J/g.c
Questions:
1. Calculate the ΔH for each Part A, B, and C.
2. Apply Hess’ Law to your calculations. Do your results agree with Hess’ Law?
3. Calculate the percent error. percent error = {[(A+B)-C]/C} x 100
Name____________________________________
Experiment 13: Electronic Structure of Atoms
Procedure:
1. Conduct flame tests on all the samples given to you by your instructor. Write down
your observations. With the flame tests you are doing today, you will observe colors
released by the cations. The anions will not release light that we can detect. Also, in
the flame, the ions are converted to neutral atoms. For example, if you are performing
a flame test on NaCl, you observe color released by neutral sodium atoms.
Sample
Observations
2. Draw an energy level diagram for two of the neutralized cations.
Cation 1 is_______________
Cation 2 is________________
Name____________________________________
Experiment 13 continued
3. Give an electron configuration for two of two neutralized cations.
Cation 3 is _________
Cation 4 is___________
4. Explain why each of the samples, based on electrons, gives a colored flame. An electron
energy level diagram would help in your explanation. Use the words “excited” and
“ground state” in your explanation.
5. Explain how you might make green fireworks. How would you make purple fireworks?
Name____________________________________
Experiment 14: Lewis structures
Procedure:
1. Draw Lewis structures for the following:
CH2F2
SeH2
NH2F
N2H4
N2H2
N2 (nitrogen gas)
H2O
C3H8
COCl2
H2CO3
(important in the
blood and in
aqueous systems
exposed to carbon dioxide)
CO2 (carbon dioxide)
SiO2(an ingredient in sand)
C2H6O
AsClFBr
(ethanol ‘found in
beer and gasoline’
or dimethyl ether,
depending on how
you draw it)
Br2
(used as a
disinfectant
in hot tubs)
O2 (oxygen gas)
Name____________________________________
Experiment 14 continued
Procedure:
2.
Draw Lewis structures for the following ions, indicating the formal charge on each atom.
H3O+1
OH-1
C2H5O-1
NO2+1
NH4+1
NO2-1
CH5O+1
O3 (non-cyclic)
Name____________________________________
Experiment 14 continued
3. Draw Lewis structures for the following. All exhibit resonance. Show resonance
structures and delocalized resonance hybrids. Show charges on any atoms that have
charges.
NO2-1
C2H3O2-1
C5H5-1(cyclic)
CNS-1
C2O4-2
-1
C3H3
(hint: no H’s on middle carbon)
4. Draw Lewis structures for the following molecules that violate the octet rule.
AlCl3
BeCl2
SeF4
XeF2
XeF4
BrF5
This page has been intentionally left blank for printing purposes.
Name____________________________________
Experiment 15: Molecular Geometry
Procedure:
Make models of the following molecules or ions. Give names of the electronic and molecular
geometries for each. Which ones are polar?
HCN
CO2
BCl3
CH2S
CCl4
NBr3
H2S
CO3-2
SO2
CH3O-
PF5
SeF6
SCl5+1
BrF5
BrF3
BrF4-1
BrF4+1
CH2FCl
NO2-1
H2O
Name____________________________________
Experiment 15 continued
Procedure:
Make models of the following molecules that have more than one central atom. Give names of
the electronic and molecular geometries for each. Which ones are polar?
These molecules have more than one central atom.
C2H4
C2H6O
C2H6
CH3COOH
Name____________________________________
Experiment 16: Gases (Boyle’s Law)
In this experiment you will measure how the volume of a gas at constant temperature varies
with pressure. [𝑃 ∝ 1/𝑉]
Procedures:
1. Transfer about 10mL of colored water to a medicine cup.
2. Fill a Pasteur plastic pipet completely with the colored water. The bulb should be totally
full, and the water should barely extend into the pipet stem.
3. Seal the pipet tip with a flame.
4. Measure the length (in cm) of the trapped air column in the pipet and record it.
5. Place the pipet on the surface of your table. Set an Aluminum block on the top of the
pipet bulb. Measure the length of the trapped air column and record it.
6. Add another Aluminum block to the top of the first one. Measure the length of the
trapped air column and record it.
7. Repeat step 6 with the third and the fourth Aluminum blocks and record the length.
Number of
Aluminum blocks
0.0
1
2
Length (in
cm)
1/length
(1/cm)
3
4
5
6
Questions:
1.
Calculate (1/V) in units of cm-1. Using the graph paper provided, prepare a graph of
(1/V) on the Y-axis versus relative pressure (Aluminum blocks) on the X-axis. Draw a
straight line that fits through the data points. Extend this line to the zero point.
2.
From your graph, find out the volume of the trapped air column in the pipet stem that
corresponds to 2.5 aluminum blocks. __________________________________
3.
If a gas has expanded from 30mL and 500torr, to 75mL, what is the pressure?
__________________
Name____________________________________
Experiment 16 continued
Name____________________________________
Experiment 17: Solutions
Procedures:
1. Use the plastic surface with the inserted mixing matrix chart to carry out the following
solubility tests.
2. Add one drop of each liquid or one crystal (grain) of the solid. Mix or dissolve
3. Record whether the two liquids or liquid/solid pairs are soluble, partially soluble, or
insoluble in each other.
Acetone
Ethanol
Water
Oil
Heptane
Acetone
Ethanol
Water
Oil
Heptane
Sugar
Salt
Questions:
1. Look up the structural formulas of the compounds used in this experiment in your
textbook. Draw out their structures (except for oil-ask your instructor for help).
Acetone
Ethanol
Water
Oil
Heptane
Sugar
Salt
2. Why does acetone mix with water?
3. Why does heptane not mix water?
4. Why is sugar (sucrose) soluble in water?
Name____________________________________
Experiment 18: Colligative Properties (Freezing Points)
Procedures:
1. In a coffee cup weigh out 25g of distilled water. In a second coffee cup, weigh out 25g
of ice.
2. Store your thermometer in another coffee cup that has a mixture of ice and distilled
water. Equilibrate your thermometer before use.
3. In boats, weigh out the following compounds individually:
a) 1.46g NaCl
b) 8.56g C12H22O11 (sucrose)
c) 2.78g CaCl2
d) 10.1g Fe(NO3)3.9H2O
4. Transfer the solid 1.46 NaCl to the cup that holds the distilled water. Stir until all the
solid has dissolved.
5. Pour the solution into the cup that has the ice. With the equilibrated thermometer, stir
for 30 seconds and then check the temperature. After another 30 seconds, check the
temperature again. If the temperature did not change, record it. If the temperature has
fallen, wait another 30 seconds and check again. Record the coldest stable
temperature.
6. The observed freezing point is the freezing point of water minus the recorded coldest
stable temperature.
7. Dispose of the solution in the sink and rinse the cups with distilled water.
8. Repeat Steps 1-6 for the other three solutes.
Calculations:
∆Tf = Kf •m• i
Kf for water = 1.86 C/m
m = molel concentration of formula units
i = Van’t Hoff factor ( i =
Solute
Mass (g)
NaCl
1.46 g
C12H22O11
(sucrose)
8.56 g
CaCl2
2.78 g
Fe(NO3)3.9H2O
10.1 g
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠 𝑖𝑛 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑
Observed f.p.
)
m
i
Theoretical f.p.
Name____________________________________
Experiment 19: Factors Affecting the Rate of Reactions
Procedures:
1. Make the following mixtures and estimate the rate of reaction for each.
2.
1 piece of Mg
½ spatula of CaCO3
10 drops of NaHCO3
6 drops of 6M HCl
3. Cool one HCl pipet in ice water. Warm another pipet with HCl in hot water. Mix the
following compounds one pair at a time. Record your results.
4.
1 piece of Mg
½ spatula of CaCO3
10 drops of NaHCO3
6 drops of cold HCl
6 drops of warm
HCl
Explain the effect of the temperature on the rate.
5. Make the following mixtures and record your results.
1 piece of Mg
½ spatula of CaCO3
10 drops of HCl
5 drops of HCl and
5 drops of water
1 drop of HCl and
9 drops of water
Explain the effect of concentration on the rate of reaction.
10 drops of NaHCO3
Questions:
1. Which of the following reactions is faster?
a) 1M HCl
+ 1 piece of Mg
b) 1M HCl (hot) + 1 piece of Mg
or
6M HCl
+
1 piece of Mg
or 1M HCl (cold) + 1 piece of Mg
Name____________________________________
Experiment 20: Chemical Equilibrium
Procedure:
A. The effect of the concentration:
Use a plastic surface to carry out the following reactions.
1. Mix 1 drop of bromothymol blue (BTB) and 1 drop 1M HCl. Record your
observations.
_____________________________________________________________________
Now add just enough 1M NaOH to induce another color change.
HBTB (yellow)
↔
H+
+
BTB- (blue)
2. Mix one drop of 6M NH4OH and two drops of Cu(NO)3. Record your observations.
_____________________________________________________________________
Now stir in just enough NH4OH to effect a change. Add HCl with stirring until the
light blue precipitate returns. Add more HCl until the precipitate disappears.
Cu2+
+2OH-
Cu2+
+
4NH3
↔
Cu(OH)2(s) (light blue)
↔
Cu(NH3)42+ (dark blue)
3. Mix one drop of Pb(NO3)2 and one drop of KI. Record your observations.
_____________________________________________________________________
Now stir in just enough NaOH to effect a change. Then add HNO3 until a third
change occurs.
Pb2+
+
2I-
Pb2+
+
2OH-
↔
↔
PbI2(s) (yellow precipitate)
Pb(OH)2(s) (white precipitate)
B. The effect of heat:
Measure 3mL of distilled water in a small (10mL) graduated cylinder and add 4 drops of
Fe+3. Now add 4 drops of SCN- Mix the solution and split it into three small test tubes.
Put one test tube in ice, the second in warm water, and keep the third as a control at
room temperature. Observe the intensity of the red-brown color of the solution in the
Name____________________________________
Experiment 20 continued
first and second test tubes. Compare it with the intensity of the color in your control
tube.
Fe3+
+
SCN-
↔
Fe(SCN)3 (red-brown)
Questions:
1.
Consider the following endothermic equilibrium:
A
How will each of the following affect the equilibrium?
a) More of B is added.
b) Some of C is removed.
c) The reaction is heated.
2.
Is the reaction in Part B exothermic or endothermic?
+
B
↔
C
+
D
Name____________________________________
Experiment 21: PH of Aqueous Solutions
Procedures:
1.
Set up a set of color standards using the universal indicator as follows. In a 13-well strip, place one drop
of universal indicator in each well. Then, in the first well, place three drops of pH 1 solution. In the
second well, place three drops of pH 2 solution, and so on. You now have a color comparison for pH of
one to twelve.
2.
Comparing weak and strong acids
In a 24 well-plate, make up the following solutions:
a) In well-1 place 10 drops of 0.10M HCL
b) In well-2 place 9 drops of distilled water, 1drop of solution from well-1. This is a 10:1 dilution of the
solution in well-1.
c) In well-3 place 9 drops of distilled water, 1 drop of the solution from well-2.This is a 10:1 dilution of
the solution in well-2.
d) In well-4 place 10 drops of 0.10M acetic acid.
e) In well-5 mix a 10:1 dilution of the solution from well-4.
f) In well-6 mix a 10:1 dilution of the solution from well-5.
Place a drop of universal indicator in each of the wells containing solutions or use the paper universal indicator.
3.
pH of common household materials
Use paper universal indicator to check the pH of the household materials available. If the household
material is a solid, dissolve it in water first.
Questions:
1.
What is the measured pH of each solution in Step 2 above?
Well-1
Well-2
Well-3
Well-4
Well-5
Well-6
2.
Looking at HCl, what is the molarity of each solution in Step 2 based on the 10:1 dilutions?
Well-1
Well-2
Well-3
Well-4
***********************
Well-5
***********************
Well-6
***********************
Name____________________________________
Experiment 21 continued
Questions:
3.
Looking at HCL, what is the molarity of H+ in each solution in Step 2 based on the pH?
Well-1
Well-2
Well-3
Well-4
Well-5
4.
5.
************************
************************
Well-6
************************
Compare the molarities in question 2 and question 3. What does this tell you about the acid (strong or
weak) and why?
Looking at acetic acid, what is the molarity of each solution in Step 2 based on the 10:1 dilutions?
Well-1
************************
Well-2
************************
Well-3
************************
Well-4
Well-5
Well-6
6.
Looking at acetic acid, what us the molarity of H+ in each solution of Step 2 based on the pH?
Well-1
************************
Well-2
Well-3
Well-4
************************
************************
Well-5
Well-6
7.
Compare the molarities in question 5 and question 6. What does this tell you about the acid (strong or
weak) and why?
Name____________________________________
Experiment 22: Analysis of Bleach
Reactions:
OCl- + H2O
I2
+
+
2I-
2S2O3-2
2I-
Cl+
+
2OH-
+
I2
S4O6-2
Procedure:
1. Rinse the microburet with distilled water. Then rinse the microburet with standardized
sodium thiosulfate solution. Fill the microburet with the thiosulfate solution.
2. Weigh a plastic cup. Place 2 mL of bleach into the cup and weigh it again. The
difference in the two measurements will be the mass of the bleach.
3. Add 1mL of 6M acetic acid to the plastic cup. Add 7 drops of 0.2M potassium iodide to
the mixture. The solution should turn red-brown.
4. Titrate this solution with the standardized sodium thiosulfate from the microburet,
swirling after each drop. Keep titrating until the red-brown color first changes to yellow
then to a colorless solution. Record the final microburet reading and repeat the
titration 3 more times.
Final microburet reading
First titration
Second titration
Third titration
Average
Questions:
1. Write an overall reaction.
2. Calculate the Molarity and the mass of NaOCl and
3. Calculate the percentage of NaOCl present in the bleach. Show your calculations below.
Name____________________________________
Experiment 23: Ka for Acetic Acid
Procedures:
1.
Pipet 20mL of 0.1M acetic acid solution into a 50mL beaker. Immerse the pH meter in the acid solution.
Fill the buret with the standard 0.1M NaOH solution.
2.
Add the NaOH solution in 1 mL increments (and swirl or stir after each addition) and report the pH and
the total volume of NaOH after each addition. As you get closer to the midpoint of the reaction, add
NaOH one drop at a time, recording the pH after each drop of NaOH is added. Continue repeating this
step until you have reached a steady pH around 10.
3.
Answer the questions below. Repeat the experiment a second time.
Questions:
1.
2.
Plot your data as pH (vertical axis) versus mL of NaOH solution added (horizontal axis) on the graph paper
provided.
From the graph, determine the exact volume of NaOH require to reach the equivalence point.
3.
Determine the half-equivalence point and from there determine the pka, then calculate the ka for the
acid.
4.
Plot your second set of data as pH (vertical axis) versus mL of NaOH solution (horizontal axis) on the graph
paper provided.
From the second graph, determine the exact volume of NaOH required to reach the equivalence point.
5.
6.
From the second set of data, determine the half-equivalence point and from there determine the pka, then
calculate the ka for the acid.
Trial one
Trial 2
Name____________________________________
Experiment 24: Vitamin C in Fruit Juice
Reaction:
L-ascorbic acid
+
I2
L-dehydroascorbic acid
+
2I-
H
HO
C
HO
C
H
O
O
C
C
H
H
C
HO
C
OH
Ascorbic Acid (Vitamin C)
Procedures:
1. The vitamin C (L-ascorbic acid) content of several fruit juices will be determined. Obtain
about 5mL of the juice and place it in a plastic cup. Add ten drops of starch (indicator).
Fill the microburet with 0.01M I2 solution.
2. Titrate the juice sample with the I2 solution, while swirling the plastic cup, until a blue
color is obtained. Record the final buret reading and that will be the volume of I 2
solution used. Discard the contents of the plastic cup and rinse the cup with distilled
water.
3. Repeat steps one and two for the other juices provided.
Questions:
1. Calculate the amount of Vitamin C in mg of each juice sample. Show your calculations.
2. How many mL of juice would be required to get the recommended daily allowance
(RDA) of 60mg of Vitamin C?
3. Once a juice bottle has been opened, the Vitamin C content gradually decreases by the
hour. Try to come up with an explanation for this occurrence.
Name____________________________________
Experiment 25: Ksp for Calcium Hydroxide Ca(OH)2
Reaction:
Ca(OH)2(s) ↔
Ca+2
+
2 OH-
Procedures:
A. Preparation of saturated Ca(OH)2 solution:
1. Transfer about 0.1g of solid calcium hydroxide to a large test tube, then add about
50mL of distilled water
2. Close the test tube with a stopper and shake for approximately 5 minutes.
3. Set the covered tube aside and allow the solution to settle.
B. Preparation and Standardization of HCl solution:
1. Rinse the microburet first with distilled water, then with about 0.1M HCl solution.
2. Fill the microburet with HCl solution.
3. Transfer 3mL of standardized NaOH solution to a plastic cup and add one drop of
Bromocresol purple indicator.
4. Tritrate with HCl until the purple color changes to yellow.
5. Repeat the titration three times.
C.
Titration of Ca(OH)2 solution:
1. Filter about 45mL of the clear solution in the test tube from part A 3-times.
2. Transfer 8mL of the clear filtrate to a plastic cup, add one drop of the indicator and
titrate with standardized HCl solution.
3. Repeat the titration three times.
Questions:
1. Calculate the Ksp for the Ca(OH)2 and compare it to the literature value (5.5 X 10-6).
Calculate the percent error. Show your calculations.
2. Using the literature, calculate the molar solubility of Ca(OH)2.
3. Using the literature, calculate the gram solubility of Ca(OH)2.
Name____________________________________
Experiment 26: Distillation
The purpose of this experiment is to learn how to set up the microscale glassware. We will
then use the glassware to perform a distillation.
Procedure:
A. Set up a distillation assembly with a condenser on top of the Hickman head. Your
instructor will have an example of a fully assembled apparatus on the front desk. Be
very careful to use enough clamps to stabilize the assembly. Glassware should be
greased. The black plastic caps (optional) should not be tightened too much, for they
break easily, especially when the glassware is heated.
B. Distillation
1. Distill about 4mL of acetone/water (50% acetone by volume). The distillate will
collect in the rough in the Hickman head.
2. Occasionally, empty the trough with a disposable pipet, emptying the pipet into a
small graduated cylinder.
3. Stop the distillation when you feel the two liquids have been separated as
completely as possible. Show your distillate to the instructor.
4. By checking the volume of the distillate, verify to yourself that the original solution
was 50% acetone.
Questions:
1. What is a visual indicator showing that essentially all the acetone has been distilled off?
2. Determine the boiling point of your distillate.
This page has been intentionally left blank for printing purposes.
Name____________________________________
Experiment 27: Electrochemistry
Procedure:
PART I:
A. The Zn/Zn+2 – Cu+2/Cu cell:
1. Measure 20mL of 0.1M ZnSO4 solution in a clean graduated cylinder, and pour it into the porous cup.
2. Measure 20mL of 0.1M CuSo4 solution in a clean graduated cylinder and pour it into the main cup.
3. Sandpaper the surface of a Zn and Cu metal strip. Place the Zn strip in the porous cup containing the ZnSO 4
solution. Place the Cu strip in the main cup containing the CuSO4 solution.
4. Place the porous cup inside the main cup as shown in the figure below.
5. Connect one electrical wire to the Zn strip, using an alligator clip. Connect a second electrical wire to the Cu
strip, using an alligator clip. Connect the free end of the wire from the Zn strip to the negative terminal of
the voltmeter. Connect the free end of the Cu wire to the positive terminal of the voltmeter.
6. Read the voltage on the voltmeter and record it as the EMF. ______________
7. Unclip the electrical wire from the Zn strip. Remove the porous cup and pour the ZnSO 4 solution into a
beaker. Save the ZnSO4 and the Zn strip for part C.
Electrical Wires to Voltmeter
Main Cup
Cu Strip
Porous Cup
Zn Strip
B. The Pb/Pb+2-Cu+2/Cu Cell:
1. Repeat the Steps #1-6, in Part A, using a 0.1M Pb(NO3)2 solution and a strip of Pb in place of the ZnSO4
solution and the Zn strip. EMF____________
2. Unclip the electrical wire from the strip. Remove the porous cup and pour the CuSO 4solution into a
beaker. Save the CuSO4 and the Cu strip for Part II.
General Chemistry (CHM 165 and CHM 175) Name____________________________________
Experiment 27(continued)
C. The Zn/Zn+2-Pb+2/Pb Cell:
1. Repeat the steps in Part B, using a 0.1M ZnSO4 solution and a strip of Zn in place of the CuSO4 solution and
the Cu strip in the main cup. Put back the porous cup that has the Pb(NO 3)2 solution and the Pb strip inside
the main cup.
2. Read the voltage on the voltmeter and record it as the EMF._________________
3. Unclip the electrical wire from the Zn strip. Remove the porous cup and pour the Pb(NO 3)2
solution into a beaker. Save the Pb(NO 3)2 and the Pb strip for Part II.
4. Pour the ZnSO4 solution from the main cup into a beaker. Save the ZnSO4 and the Zn strip for Part II.
PART II Studying Cell Reactions:
A.
Zn metal in CuSO4 solution
1. Place the Zn strip in the CuSO4 solution retained from Part I.
2. Observe the solution and the strip for 5 minutes, and then remove the Zn strip from the solution.
Retain the solution for the next step.
3. Write a chemical equation that describes the reaction.
B.
Pb metal in CuSO4 solution:
1. Place the Pb strip in the CuSO4 solution retained from Part I.
2. Observe the solution and the strip for 5 minutes, and then remove the strip from the solution. Pour
the solution into the heavy metal waste. Record all observations.
3. Write a chemical equation that describes the reaction.
C.
Zn metal in Pb(NO3)2 solution:
1. Place the Zn strip in the Pb(NO3)2solution retained from Part I.
2. Observe the solution and the strip for 5 minutes, and then remove the strip from the solution.
3. Pour the solution into the heavy metal waste container. Record all observations.
4. Write a chemical equation that describes the reaction.
Part I
Cell
Zn/Zn+2-Cu+2/Cu
Pb/Pb+2-Cu+2/Cu
Zn/Zn+2-Pb+2/Pb
EMF (v)
Part II
Arrange the three metals in order of their ability to accept electrons:
First_________________________ Second__________________________ Third______________________
Name____________________________________
Experiment 28: Radioactivity
Procedure:
A. Rate of radioactive decay (Cs-137) to (Ba-137)
1. Measure the background counts per minute (cpm) using Geiger counter. __________________
2. Place the isotope of Cs-137 under the Geiger counter tube and measure the counts per minute for the next
5 consecutive minutes.
3. Obtain the corrected cpm by subtracting the background cpm from the observed cpm. _________________
4. Using the graph paper provided, plot the corrected cpm on the y-axis versus the time on the x-axis.
Time (minutes)
Observed (cpm)
Corrected (cpm)
1
2
3
4
5
B. The Penetration power of gamma rays
1. Measure the background counts per minute (cpm) using Geiger counter.
2. Obtain a source for gamma radiation and place it under the tube of Geiger counter.
3. Measure the observed counts per minute.
4. Place one plate of Aluminum metal between the source of gamma rays and the tube of Geiger counter then
measure the observed cpm.
5. Repeat Step #4 by placing a second Aluminum plate and measure the observed cpm.
6. Repeat Step #5 by placing a third Aluminum plate and measure the observed cpm.
7. Repeat Step #6 by placing a fourth Aluminum plate and measure the observed cpm.
8. Repeat Step #7 by placing a fifth Aluminum plate and measure the observed cpm.
9. Obtain the corrected cpm by subtracting the background cpm from the observed cpm.
10. Using graph paper, plot the corrected cpm on the y-axis versus the number of plates on the x-axis.
Number of plates
Observed (cpm)
Corrected (cpm)
0
1
2
3
4
5
Questions:
1. A beam of gamma rays of energy 24k.j. is passed through Aluminum plate 5cm thick. If k for Aluminum is
0.56cm-1, find what fraction of the beam is transmitted by the plate.
I = I0.e-kd
Name____________________________________
Experiment 28 (continued)
Name____________________________________
Experiment 29: Hydrogen Bonding
Procedure 1:
Sodium polyacrylate is a compound found in diapers. It can absorb many times its own weight in water. Water
makes hydrogen bonds to the sodium polyacrylate, therefore sticking to it. Also sodium polyacrylate has been
undergoing testing for use in firefighting. Solutions of it can by sprayed on houses in a region that are having
forest fires.
1.
Take a 100 ml size beaker:
a)
Pour in about 1.5 g of sodium polyacrylate.
b) Add water, little by little, to see how much water can be absorbed (save your product for the
next step).
2.
Take your product:
a)
Smear your product over a piece of wood.
b) Light up a Bunsen burner (remember to use goggles whenever a Bunsen burner is in use).
c)
Try to burn the wood that is covered with the sodium polyacrylate.
Procedure 2:
You are going to add 1 mL heptane (C7H16) and 1 mL water together in a test tube. But before you do this, predict
what you will see and explain why.
1.
In a small test tube add about half mL of heptane to about half mL of water, but pour them together
slowly to see what molecule is less dense and ends up on top (save this solution for the next step).
a)
2.
See if your prediction was right. ________________
Slowly add 1 mL methanol to the water-heptane solution, watching to see what layer (heptane or water)
the methanol goes into. Stir gently (save this solution for the next step).
Name____________________________________
Experiment 29 continued
3.
Draw a diagram showing hydrogen bonds with dotted lines.
5. Add 1 mL hexane to the water-heptane-methanol solution. Gently stir. Which layer did the hexane go
into? Explain.
Procedure 3:
Derivatives of Ethylene glycol are used in the production of synthetic rubbers, synthetic fibers (e.g. Dacron
polyester, Mylar, and Kevlar) resins, paints, adhesives, molded articles, solvents, brake fluid, and even cosmetics.
1.
Place 1 mL ethylene glycol in the bottom of two different test tubes.
2.
In one test tube add 1 mL water.
3.
In the other test tube, add 1 mL hexane.
4.
Gently stir (swirl) the test tubes for a while to see which solvent dissolves the ethylene glycol.
5.
What can you conclude about the ethylene glycol?
Procedure 4:
Vitamin C is known to be a water soluble vitamin. It can easily be excreted out of your body in urine.
1.
Add a small amount of Vitamin C (ascorbic acid) to water. It should dissolve with some swirling.
2.
Draw a picture of Vitamin C (see Vitamin C in your lab packet). Show a bunch of water molecules around
your Vitamin C molecule. Label partial negatives and partial positives on all the molecules and show a
bunch of hydrogen bonds with dotted lines. Because of this hydrogen bonding, water and Vitamin C are
attracted and are soluble in each other.
Name____________________________________
Experiment 30: Intermolecular Forces (Slime Gel)
Reactions:
1.
Na2Br4O7 + 7H2O
2Na+
2. 2Na+ +2B(OH)4- + H2SO4
+
2B(OH)3
+
2B(OH)4-
(Making the Gel)
2Na+ + 2B(OH)3 + (SO4)2-2 + 2H2O
(Killing the Gel)
HHHHH H
Polyvinyl alchol:
-C-C-C-C-C-C-
HOH H OH H H
Procedures:
Making slime gel
1.
2.
3.
Place 50mL of a 4% polyvinyl alcohol solution (PVA) in a 250mL beaker. Add 10 drops of 0.1% methyl red
indicator to the PVA solution. Stirring continuously using a wood splint or glass rod, add 1.0mL of 4%
borax (Na2B4O7) solution drop by drop. Continue vigorous stirring until the mixture forms a homogenous
gel. Note and record the characteristics of the gel. Can you pour the gel from the beaker to another one?
What happens to the gel when you place a lump of it in the center of a clean watch glass?
Now add drop by drop, with stirring, a second 1.0mL portion of 4% borax solution to the gel. Stir until the
gel is homogenous. Note and record any changes in the properties of the gel. Is it more viscous? Or less?
Is it easily stirred?
Finally, add drop by drop, with stirring, a third 1.0mL portion of 4% borax solution to the gel. As before,
note and record any changes in the properties of the gel. Try rolling the gel into a ball, then seeing how
long it takes for the ball to flatten out when placed on clean watch glass. Try pulling on the gel at
different speeds, slowly then very rapidly.
Killing and Resurrecting the Slime Gel
1.
2.
3.
With Stirring, add 3.0M H2So4 solution drop by drop until the color of the gel solution changes to red/pink
and you get a homogenous solution. Note and record any changes in the properties of the gel.
Is the gelation process reversible? To test this, add drop by drop 1.0M NaOH with stirring until you get a
yellow color. Does the gel regain its high viscosity? Note and record your observation.
With stirring, add enough 3.0M H2SO4 solution drop by drop until the color of the gel solution changes to
red/pink and you get a homogenous solution. Discard the solution in the sink.
Name__________________________________
Experiment 30 continued
Questions:
1. Describe what happens when you stir the gel in making the Slime Gel, #1.
2. What happens when you attempt to pour the gel from one beaker to another in
Making the Slime Gel, #1?
3. What happens when a ball of the gel is allowed to stand on the center of the clean
watch glass in Making the Slime Gel, #1?
4. Is the Gel more viscous? Or less in Making the Slime Gel, #2?
Name____________________________________
Experiment 31: Part A: Paper Chromatography
Procedure:
A. Preparing the Chromatogram
1. Fold the filter paper in half, open it and fold it in half again so that the creases from the first fold meet at the
edge.
2. Open the folded paper, there should be an X pattern.
3. Next, fold the filter paper so that the bottom rays of the X meet the top rays.
4. Open the paper and fold it one more time so that the three rays of the bottom half meet the three rays of
the top half.
5. Open the filter paper, there are now eight rays (creases) radiating from the center of the circle.
6. Place the chromatogram on a paper towel.
7. With a No. 2 pencil, mark the point of intersection (this is the center of the circle).
8. Use a ruler and place a pencil mark 1cm from the center along each ray.
9. Fold the paper in half and then in half again, and with a paper punch, clip off the tip of the angle
approximately 1mm from the point. If done properly, the circular hole created will measure approximately
0.5 cm in diameter.
10. With a No. 2 pencil, make a small X between the center and one of the pencil marks to indicate the starting
point (Position 1).
11. Start at Position 1, carefully make a small dot of the color indicated below on each pencil mark, (Do not
hold the marker on the filter paper) let the ink dry and repeat the process one more time.
Position
1
2
3
4
5
6
7
8
Color
Sharpie Permanent Black
Black
Red
Yellow
Blue
Green
Orange
Purple
B. Preparing the Wick
1. Obtain a small filter paper.
2. Fold or roll this filter paper into a small cylindrical shape.
3. Place one end of the filter paper into the hole of the chromatogram. The filter paper will serve as a water
wick.
C. Running the Chromatogram
1. Fill the culture dish with tap water to within 2cm of the top. It is very important to dry the rim of the culture
dish with a paper towel before proceeding any further.
2. Place the chromatogram over the culture dish so that it rests on the rim with the wick extending down into
the water.
3. The water will travel up the cone to the filter paper, wetting the filter paper in a circular pattern. When the
water reaches the ink spots, the radial chromatography pattern begins to develop.
4. Allow your chromatogram to develop until the water reaches a point approximately 1cm from the outer
edge of the filter paper.
5. Remove the filter paper and allow it to dry. Write your name on the dry filter paper and turn it in as your lab
report.
Name__________________________________
Part B: Column Chromatography
Procedure:
1. Obtain an apparatus for column chromatography and carefully remove the upper and bottom caps from the
column (Don’t misplace the two end caps of the column). Let the buffer elute in a beaker.
2. Once the buffer is eluted, put the bottom cap back at the bottom of the column.
3. Using an eye dropper, load 2 drops of the sample dye on the top of the column.
4. Once the sample is loaded into the column, put a microtube under the column to collect the buffer solution.
Remove the bottom cap and slowly load the top of the column with 15mL of elution buffer by using an eye
dropper.
5. As the buffer travels down the column, the sample will separate into two colors. Collect each colored solution
in a separate microtube.
6. Re-equilibrate the columns with 20mL of buffer for the next class. Leave a minimum of 1cm of the buffer in the
column, and replace both the top and the bottom caps of the column.
7. Dispose of all unneeded solutions in the drain.
Questions:
1. Was the color separation distinctive? Would you expect a longer column to more clearly separate the
compounds? Why or why not?
2. Suppose your sample had consisted of a mixture of two compounds. The first compound has a larger molecular
size than the second compound. Which one of the two compounds do you expect to collect first?
Name____________________________________
Experiment 32: Chemical Compounds and Formulas
Formations of a Chemical Compound
We know a chemical reaction occurs if there is a change of color or an alteration of the physical
properties of the reactants. Knowledge about the relative masses of reactants and products also gives
us information about the makeup or formulas of the substance involved in the reaction. In this
procedure you will combust a weighed sample of magnesium metal with oxygen to produce magnesium
oxide (magnesia). You will use the mass of the product and the reactants to determine the formula or
proportion of magnesium and oxygen of the magnesium oxide.
Determining the Formula of a Compound
You will need a crucible and crucible lid, forceps (tong), clay triangle, iron ring and ring stand.
You will be using a Bunsen burner, so be careful with the open flames.
Set up your apparatus as demonstrated.
1. Obtain a 20 cm (8 inch) strip of magnesium ribbon. If it is tarnished, polish it with steel wool
or by scraping it off with a straight edge.
2. Weigh and record the mass of a crucible and lid.
3. Coil the metal strip and put it into the crucible. Weigh and record the mass of the metal, the
crucible and the lid.
4. Heat the crucible and magnesium with a gas flame until the magnesium strip ignites.
Immediately cover the burning magnesium with the crucible lid.
5. Continue to heat, periodically (every 30 seconds or so) lifting the crucible lid to allow oxygen
into the combustion chamber. After the magnesium ribbon no longer flares up as you lift
the lid, remove the lid, and then heat strongly for five minutes to insure complete
combustion of the metal.
6. Remove the crucible from the heat and allow it to cool for at least five minutes.
7. Add 10 to 15 drops of water to the product. Heat the moistened product for 3 to 5 minutes
to drive off any magnesium nitride formed during combustion.
8. Remove the crucible from the heat and allow it to cool for at least five minutes.
9. Weigh and record the mass of the product, the crucible and the lid.
Calculations
I.
II.
a. Mass of metal, crucible and lid
b. Mass of crucible and lid
c. Mass of magnesium metal (a – b)
d. Mass of product, crucible and lid
e. Mass of product (d – b)
f. Mass of oxygen (e – c)
_____g
_____g
_____g
_____g
_____g
_____g
III. Moles of magnesium
g. Mass of Magnesium (c)
IV. Moles of oxygen
h. Mass of Oxygen (f)
X __1mol Mg_ = ________mol Mg
24.3g Mg
X __1mol O__ = ________mol O
16.00 g O
Use the calculations from steps III and IV to determine the formula of magnesium oxide.
Divide the mole values you obtained in steps III and IV by the smaller value of g or h and insert the
rounded value into the subscript boxes of the formula below.
Formula for magnesium oxide:
Mg
O