1. business and industrial
2. energy
3. oil

Name
Period_
Volume 5
Falcon Pre-AP Chemistry Lab Notebook
“The dogmas of the quiet past are inadequate to the stormy present.
The occasion is piled high with difficulty, and we must rise with the
occasion. As our case is new, so we must think anew and act anew.”
∙Abraham Lincoln∙
Table of Contents
Page
Lab
3
5
9
11
13
15
18
21
23
27
29
Penny Change
(Bonding 1)
You Light Up My Life
(Bonding 2)
Intermolecular Forces
(Bonding 3)
The Eyes Tell No Lies
(Bonding 4)
Unsaturated, Saturated, Supersaturated
(Solutions 1)
Solubility Rules Determination
(Solutions 2)
Solubility Curve of KNO3
(Solutions 3)
Serial Dilutions
(Solutions 4)
Properties of Acids and Bases
(Acid & Base 1)
Tie Dye
(Acid & Base 2)
Titration of a Strong Acid and Base
(Acid & Base 3)
Would you like your own pair of goggles?
Order from Flinn Scientific, Models AP3306
and AP3309 are both approved.
Reminders:

Lost notebooks: students may print out additional copies from the website but 5
points will be deducted from lab grades for each occurrence.

Labs turned in one day after the due date will receive a maximum of 75, 2 days
late will receive a maximum of a 50, and any work turned in after the 3rd day (or
not turned in at all) will receive a zero.

Goggles are to be worn at all times, unless instructed otherwise. 20 points will be
deducted from the lab grade for the first goggle reminder and a grade of zero will
be given for a second goggle reminder within the same lab period.
Name
Period_
Penny Change – Bonding 1
Introduction:
Alchemy is a practice which began in the early civilizations of the ancient world and continued through
medieval and Renaissance times. One of the primary focuses of alchemy was the attempt to change
common metals, such as lead or iron, into gold. We now know that it is not possible to convert one
element into another. In this laboratory investigation, you will make changes to a penny that give the
appearance of a “new” metal, and you will learn the explanation for the production of this seemingly new
material.
Background information:
Pennies are mainly composed of zinc, with a thin outer coating of copper. Currently, it costs about 1.2
cents to make a penny, but this cost is mainly offset by the fact that pennies change hands many times
before going out of circulation.
An alloy is a mixture of metals. The metal atoms in alloys are held together by metallic bonding.
Materials/Equipment:
1.0 M ZnCl2 solution
2.2 g granular zinc (20 mesh)
400 mL beaker
Crucible tongs
Watch glass
250 mL beaker for water bath
Hot Hands
8-10 clean pennies
Hot plate
Procedure:
1. Start with clean pennies, or the lab may not yield the desired result. If necessary, pennies may be
cleaned with a mixture of vinegar and baking soda. Let sit for 5 minutes and then wipe/rinse pennies.
2. Set aside one clean penny as your control, and describe appearance in data table.
3. Line the bottom of the 400 mL beaker with granular zinc.* (Teacher will do this)
4. Pour enough 1.0 M ZnCl2 solution into the 400 mL beaker to reach the 50-75 mL mark. Do
not exceed this amount.* (Teacher will do this)
5. Using tongs, carefully add as many clean pennies as possible to the 400 mL beaker to make one
layer (approx. 7 pennies). Do not drop the pennies in, or they may cause splashing of chemicals.
6. Put beaker on hot plate and cover with watch glass. Heat to a full boil, then reduce heat and allow
beaker contents to heat for 1-2 more minutes.
7. Using "hot hands," remove beaker from heat and place on lab table.
8. Fill the 250 mL beaker approximately halfway with tap water. Using tongs, move the pennies into
the beaker with the water to cool and rinse them.
9. Remove pennies with tongs and place on paper towel. Blot dry. Record the appearance of the pennies
in your data table (Penny A).
10. Turn the hot plate on high. Using tongs, place all but 1 of the pennies directly on the hot plate
(save this one for comparison later). Leave pennies on the hot plate for approximately one minute
until a color change is observed, then promptly remove pennies with tongs. DO NOT let pennies
melt! Turn off hot plate.
11. Rinse penny under tap water. Place on paper towel and blot dry. Record the appearance of the
pennies (Penny B) in your data table.
12. Clean up: Cover the 400 mL beaker with the watch glass and set aside. DO NOT POUR THE ZINC
INTO THE SINK!!! Pour the water out of the 250 mL beaker and rinse. Wipe counter/wash hands!
3
Data Table:
Penny
Treatment
Control
None (control)
Penny A
Zn/ZnCl2
Penny B
Zn/ZnCl2 + heat
Appearance
Questions/Conclusions:
1. What element forms the coating on Penny A?
2. What was responsible for causing the change from Penny A to Penny B?
3. What two elements is the metal coating on Penny B made of?
4. The answer to #3 is an alloy. What is the name of this alloy?
5.
List two practical uses or applications of this alloy.
6. What type of bonding exists in the atoms coating the pennies? Circle: ionic, covalent, metallic
7. List 3 characteristics of substances that have this type of bonding (see notes).
4
Name
Period_
You Light Up My Life – Bonding 2
Comparing Conductivity in Ionic and Covalent Compounds and Metals
Purpose
Investigate the properties of substances and group substances according to their properties.
Background
In this experiment, you will study the electrical conductivity of various metals, ionic and covalent
(molecular) compounds, and aqueous solutions. In a solution, the size of the conductivity value
depends on the ability of the aqueous solution to conduct electricity. A solution that conducts
electricity very effectively is known as a strong electrolyte. A solution that does not conduct electricity
well is known as a weak electrolyte. A solution that does not conduct electricity at all is a
nonelectrolyte. You will look for patterns in conductivity and make summarizing statements about
which types of substances conduct electricity, and under what circumstances.
Pre-Lab Questions:
1. If you were to drop a spoonful of salt (NaCl) into a glass of water, what would happen?
2. If you were to drop a gold ring into a glass of water, what would happen?
3. What do you think is different about the atoms of these two substances? Why do you suppose the
gold atoms don’t break apart?
4. Hypothesize whether or not each of these substances would conduct electricity, and give a reason for
your answer:
Salt
Sugar water
Sugar
Copper
Salt water
Wax_
You will check your answers to these questions and refine your explanations during this lab.
5. Label the following compounds as I (ionic) or C (covalent):
C6H12O6
NaCl
C12H22O11
AlCl3
CaCl2
H2O (Distilled)
_____BBr3
Cu(NO3)2
C2H6O
Materials:
Vernier Conductivity Probe
Sand (SiO 2)
Distilled water (H 2O)
Conductivity apparatus
Beakers
Aluminum Chloride (AlCl 3)
Vernier LabQuest 2
Sucrose (C 12H 22O 11)
Copper(II)nitrate, Cu(NO 3) 2
Ethanol (C 2H 6O)
Calcium chloride (CaCl 2)
Aluminum foil (Al)
5
Paraffin wax (C 20H 42)
Ring Stand/Utility Clamp
Wash bottles/rinse beakers
Copper (Cu)
Sodium chloride (NaCl)
Investigation
Test 1: Conductivity Apparatus (light bulb/battery)
1. Assemble your conductivity apparatus according to your teacher’s instructions.
2. Make a prediction for the conductivity of each substance in the data table.
3. Then use the apparatus to test the conductivity of each substance. Enter your results in
the data table. DO NOT TOUCH THE ENDS OF THE WIRES TOGETHER, OR THE BULB
WILL BURN OUT QUICKLY.
Test 2: Conductivity Probe
1. Be sure the probe is clean and dry before beginning the experiment.
2. Connect the Conductivity Probe to the LabQuest CH 1 and choose New from the File menu
(this step may already be done for you). If you have an older sensor that does not autoID, manually set up the sensor.
3.
Set the selector switch on the side of the Conductivity Probe to the 0–20000
µS/cm range
4. Make sure to “zero” the LabQuest. It should read 0.00. To zero, tap on the
screen and touch the “zero” command.
5. Measure the conductivity of each of the solutions.
a. Carefully raise each vial and its contents up around the Conductivity Probe until the
hole near the probe end is completely submerged in the solution being tested.
Important: Since the two electrodes are positioned on either side of the hole, this
part of the probe must be completely submerged.
b. Briefly swirl the vial contents. Monitor the conductivity reading displayed on the screen for
6–8 seconds, and once the conductivity reading has stabilized, then record the value in
your data table.
c. Before testing the next solution, clean the electrodes by surrounding them with a 250 mL
beaker and rinse them with distilled water from a wash bottle. Blot the outside of the
probe end dry with a paper towel. It is not necessary to dry the inside of the hole near the
probe end.
d. DO NOT touch the inside of the probe! CAUTION!!!
e. Test the conductivity of the remaining solutions. Just rinse in between to
prevent contamination!
Test 3: Dissolving (Solubility)
1. Make a prediction for the dissolving of each substance in the data table.
2. Take a very small portion of each substance and try to dissolve it in 25-30 mL of water.
3. Enter your results in the data table.
4. Before proceeding to the next substance, test any substances that did dissolve for
conductivity as well. If a substance did not dissolve, enter N/A in the data table
for “conduct if dissolved.” Dry off tips of wires after tests.
5. Rotate to the next station on your teacher’s directive, to complete tests for all
substances.
6
Name
Period_
Data Table
Substance
You Light up My Life: Conductivity Lab
Predict
Predict
Test
Lightbulb/batt.
Test
Test
Lightbulb/batt.
Conductivity Reading
w/ probe for solutions
Conduct?
Dissolve?
Conduct?
Dissolve?
Give value with unit:
µS/cm
Y/N
Y/N
Y/N
Y/N
Conduct if
dissolved?
Y/N or N/A
Distilled water
H2O
omit
N/A
N/A
0 µS/cm
Aluminum foil
Al
omit
N
N/A
N/A
Silicon dioxide
SiO 2
omit
N
N/A
N/A
Paraffin
C 20 H 42
omit
N
N/A
N/A
Copper
Cu
omit
N
N/A
N/A
Calcium chloride
CaCl 2
Copper(II)nitrate
Cu(NO 3 ) 2
Aluminum
Chloride AlCl 3
Sucrose
C 12 H 22 O 11
Sodium chloride
NaCl
Ethanol
C2H6O
7
Processing the Data:
1. Make a list of those substances that conduct electricity but do not dissolve in water.
What other things do these substances have in common? (Think about their properties and their
chemical formulas.)
2. Now consider those substances that dissolve in water. Divide them into 2 categories, those
that conduct electricity when they are dissolved, and those that do not.
Conduct
Substances that dissolve in water
Do not conduct
Refer to the table you just filled in:
3. What do the substances that conduct electricity once they are dissolved have in common?
4. What do the substances that do not conduct electricity once they are dissolved have in
common?
5. Consider the conductivity readings for sodium chloride, calcium chloride, copper (II) nitrate,
and aluminum chloride solutions. What trend do you observe? Explain.
6. Does distilled water conduct electricity well? Explain.
7. If it is dangerous to use a blow dryer while taking a bath, what also must be true about water
in a bathtub?
8. Along those same lines, why do lifeguards require that you exit the pool on the approach of a
thunderstorm? Would it be as dangerous if the pool were filled with distilled water?
8
Name
Period_
Intermolecular Forces – Station Lab – Bonding 3
Station 1: Intermolecular Forces Part A - Table 1
LDF present? Dipole-Dipole
Molar
Boiling Point
Force present?
(Y/N)
Mass
(oC)
(Y/N)
(g/mol)
CH4
16
- 164
NH3
17
- 33
H2O
18
100
HF
20
- 20
HCl
36.5
- 85
C3H8
44
- 42
1. List TWO reasons why CH4 has the lowest boiling point of all the substances.
Substance
H-bonding
present?
(Y/N)
2. Why is water’s boiling point so high, even though it is not the heaviest molecule?
3. Compare HF to C3H8. Why does HF have a higher boiling point than C3H8, even though HF
is less than half the size of C3H8?
Station 2: Intermolecular Forces part B
4. In the data for groups 5A, 6A, and 7A, the boiling point for the lightest molecule is unusually high.
Why does this take place? Why does CH4 not behave this way?
9
Station 3: Intermolecular Forces part C
Sketch of 4 HCl molecules :
Sketch of HCl
(aq)
Sketch of 4 H2O molecules :
:
How does “like dissolves like” apply in the above drawing of HCl(aq)?
Station 4: Intermolecular Forces part D
Initial volume of water only: Initial
volume of alcohol only:
Volume of water and alcohol combined:
mL
mL
mL
Explain why the combined volume does not equal the sum of the other two volumes.
Station 5: Surface Tension part A
1. Why did the pepper initially sink in alcohol but float on water?
2. What effect did the detergent have on the pepper/water and why?
3. Detergents are large, long molecules that dissolve in polar water, and at the same time, they attract
nonpolar substances such as grease and oil. Based on this dual-nature behavior, what can you infer
about the structure of detergents?
Station 6: Surface Tension part B
Substance
Ethyl alcohol
Water
Predicted # of drops
Actual # of drops
1. Give an explanation for the difference in the actual number of drops that could be placed on the
penny (alcohol vs. water).
Station 7: Evaporation and “Like Dissolves Like”
1. List all IMFs present in each liquid:
2. Which liquid would evaporate faster, acetone or ethyl alcohol? Why?
3. Which liquid is in the test tube with the water? How did you know?
10
Name
Period_
The Eyes Tell No Lies – Bonding 4
Background
There are many different solutions that can be gathered at a crime scene or an autopsy, such as alcoholic beverages,
blood, and even fluid from the vitreous humor of the eye. The vitreous humor is the fluid of the eyeball consisting of over
90% water, and some dissolved proteins and ions. A sample is generally taken during an autopsy to determine if alcohol
is present in an individual.
Alcohol can be detected in blood and urine since it is transferred throughout the body after consumption. Unfortunately,
this evidence can be disputed since increased levels of alcohol can be formed in the body after death from biological
changes, bacterial degradation and/or contamination of human tissue, especially near the stomach. However, the vitreous
humor is isolated and does not have any ingredient for the formation of alcohol. The only way for alcohol to be present in
the vitreous humor is through consumption.
SCENARIO: There has been a fatal vehicle crash with no surviving parties. To determine if the individual responsible for
causing the crash was under the influence of alcohol, a sample of their vitreous humor is taken during autopsy. As a
toxicologist, it is your job to test the sample for the presence of alcohol. Follow the procedures below to test for the
presence of alcohol in the vitreous humor sample.
Materials
5 mL suspect's vitreous humor sample
Bromothymol blue indicator (2 drops)
Pipettes
Test tube rack
0.5 grams of potassium carbonate
1.0 M HCl (1 drop)
2 test tubes
2 rubber stoppers
Procedure
To determine if alcohol is present in the sample of vitreous humor, carry out the following investigation to
separate the ethanol alcohol from the water.
1. Pour 5 mL of the suspect's vitreous humor sample in test tube #1.
2. Add 2 drops of bromothymol blue indicator to test tube #1.
3. Add 1 drop of 1.0 M HCl to test tube #1 and swirl until mixed.
4. Place 0.5 grams of potassium carbonate in test tube #2.
5. Pour the solution from test tube #1 into test tube #2.
6. Place the stopper securely into the mouth of test tube #2 and shake.
7. If alcohol is present, the solution will separate into two phases.
8. Record all of your observations and data.
9. Complete the conclusion and analysis section.
11
Name
Period_
Conclusion and Analysis
1. Based on your investigation, was the individual under the influence of alcohol at the time of the crash? Explain your
answer.
2. Ethanol and water are miscible liquids due to their intermolecular forces. These forces can be easily disrupted when a
salt, such as potassium carbonate, is added. The process is known as "salting out." Based on your knowledge of the
behavior of water, explain why the addition of a salt results in the separation of water and alcohol?
3. What types of forces exist between the water and the alcohol?
4. What types of forces exist between the water and the potassium carbonate?
12
Name
Period_
Unsaturated, Saturated, and Supersaturated – Solutions 1
Solutions are considered unsaturated if more solute can be dissolved at the existing temperature.
If the maximum amount of solute is in equilibrium with undissolved solid, the solution is saturated, the
undissolved solid solute will appear to sit on the bottom as excess.
In a few special substances if a hot, saturated solution is allowed to cool without losing any solute to
crystallization, the solution is supersaturated. This is rare, limited to just a few salts which manage to
stay in solution at an amount which exceeds the solubility at a particular temperature.
Most solutes fall out of solution as the temperature is lowered, but one substance which can supersaturate
is sodium thiosulfate, Na2S2O3∙5H2O. In this lab you will prepare all three types of solutions.
PROCEDURE
1. Weigh out approximately 15 grams of sodium thiosulfate.
2. Pour 3.0 mL of distilled water into a large test tube.
3. Add about 2 grams of the 15 g of sodium thiosulfate to the test tube. If it all dissolved, the
resulting solution is unsaturated.
4. Pour the remaining 13 grams of the sample into the test tube. Feel the outside of the test tube to
note any change in temp. This is a saturated solution.
5. Heat the test tube in a hot water bath until the entire 15 gram sample is in solution with no solute
visible. Do NOT allow solution to boil.
6. Remove the solution from the heat and allow to stand for one minute. Shake. Place the test tube in
a room temperature water bath and allow to stand undisturbed for about 15 minutes until it is at
room temperature.
7. Observe the test tube closely. If it has cooled to room temperature and shows no visible solute, the
solution is supersaturated. [If the solute has crystallized out of solution, the solution is just
saturated and needs to be heated again, mixed thoroughly, and allowed to cool again.]
8. Add a tiny piece of sodium thiosulfate crystal to the supersaturated room temperature solution and
observe. Touch the bottom of the test tube each 15 seconds to note temperature changes until all
crystallization is complete.
9. Reheat test tube in hot water bath to redissolve. Pour into waste beaker. Wash test tube.
13
Name
Period_
QUESTIONS - Answer questions 1, 5, 6, and 8 with complete sentences.
1. How can such a large amount of solute be dissolved in just 3 mL of water?
2. Did the dissolving process itself raise or lower the temperature of the water?
3. What type of dissolution process does this indicate, endothermic or exothermic?
4. Two processes occur in dissolving: solute - solute bond breaking, and solvent - solute bond
forming. Which process involved more energy?
5. Describe the part of the procedure at which the solution was saturated.
6. Describe the part of the procedure at which the solution was supersaturated.
7. Did the process of crystallization absorb or release energy?

Is this crystallization endothermic or exothermic?

Two processes occur in crystallization: solute - solvent bond breaking, and solute - solute
bond forming. Which process involves more energy?
8. Which solution is more stable for long - term storage -- saturated or supersaturated? Why?
14
Name
Period_
Determining Solubility Rules by Combining Ionic Compounds - Solutions 2
Objectives
1. Observe and record chemical changes involving precipitates.
2. Deduce several basic solubility rules based on observations
3. Write balanced equations and identify precipitates for reactions that took place.
Safety Precautions
YOU MUST WEAR YOUR GOGGLES AT ALL TIMES DURING THIS LAB!
BEWARE OF:
AgNO3 - brown spots will develop where it touches skin
NaOH - can burn skin
HCI - can burn skin
Introduction
In this lab we will investigate the products formed by the combination of many different ions,
and look for patterns. When we are finished, we should be able to make rough predictions
about whether or not two chemicals will react to form an insoluble solid, or precipitate.
+
An example of a solubility rule involves the sodium ion (Na ). Many people have observed
over the years that compounds with sodium ions are very soluble in water, so there is a
solubility rule that says "all sodium salts are soluble." If we combine a sodium compound with
another compound and an insoluble solid (precipitate) is produced, we know that the positive
ion in the product is not sodium.
Procedures
1. At your lab station you should find a cup with 8 labeled pipets containing various
solutions of ionic compounds. You should also find a labeled and in a plastic sheet
protector on which you will carry out your reactions. Carry out the reactions in the
white squares, the shaded squares represent duplicate reactions.
2. Carefully drop two drops of one solution into the properly labeled square. Remember
these solutions ions that are free to move throughout the solution. If you provide a
“bridge” between two solutions by letting the dropper touch the “bubble” you are
cross contaminating the solutions with ions, which destroys the solutions. SO,
taking care NOT TO TOUCH THE TIP OF THE DROPPER to the solution already on
your grid, carefully add two drops of the second solution. Let the reaction stand
about 10 seconds and then record your observations on your data sheet--even if
there is no visible reaction.
3. In the next square, repeat the above procedure for another pair of solutions. Record
your observations for each combination, even if there is no visible reaction
on the 3rd page of this lab. Repeat for all combinations listed in your data table.
4. When you have made all observations, carefully clean the sheet protector by
absorbing the chemicals with a dry paper towel. Dispose of the paper towels in the
trash. Be careful not to contaminate your hands with the chemicals. If there is still a
residue on your sheet protector, use a damp paper towel to clean it off. Do NOT use
soap as this will affect the surface tension of the water droplets.
15
Name
Period_
Analysis/Conclusions
1. Look at the reactions involving silver compounds. Are they more often soluble or insoluble?
What general conclusions can we draw about the solubility of silver ions?
2. Look at the reactions involving lead compounds. Are they more often soluble or insoluble?
What general conclusions can we draw about the solubility of lead ions?
3. What other generalizations can you make about solubility given the data you collected?
4. For the double replacement reactions specified by reaction number in the table below, fill in
the table and BALANCE the reaction.
Reactant 1
Rxn# Formula/Name of
Reactant 1
1
+
Reactant 2
Formula/Name
of Reactant 2

Product 1
Formula/Name
of Product 1
+
Product 2
Formula/Name of
Product 2
CuCl2 (aq)
K2SO4 (aq)
2 KCl (aq)
CuSO4 (s)
copper (II) chloride
potassium sulfate
potassium chloride
copper (II) sulfate
15
25
28
35
16
Which product is
precipitate?
CuSO4
K2SO4
BaCl2
NaOH
Pb(NO3)2
AgNO3
K2CO3
Na3PO4
1
9
16
22
27
31
34
2
10
17
23
28
32
35
3
11
18
24
29
33
4
12
19
25
30
5
13
20
26
6
14
21
7
15
HCl
CuCl2
36
CuCl2
HCl
Na3PO4
K2CO3
AgNO3
Pb(NO3)2
NaOH
Name_______________________________________________________
Period_____
8
BaCl2
K2SO4
17
Name
Period_
PRE-LAB DISCUSSION
LAB - Solubility of KNO3– Solutions 3
The solubility of a pure substance in a particular solvent is the quantity of that substance that will dissolve in a given volume of the
solvent. Thus, solubility must be expressed as quantity of solute per quantity of solvent at a specific temperature. For most ionic
solids, especially salts, in water, solubility varies directly with temperature. That is, the higher the temperature of the solvent (water),
the more solute (salt) that will dissolve in it.
In this experiment, you will study the solubility of potassium nitrate (KNO3) in water. You will dissolve different masses of this salt
in a given amount of water at a temperature close to the water’s boiling point. Each solution will be observed as it cools, and the
temperature at which crystallization of the salt occurs will be noted and recorded. The start of the crystallization indicates that the
solution has become saturated. At this temperature, the solution contains the maximum quantity of solute that can be dissolved in that
amount of solvent.
After solubility data for several different quantities of solute have been collected, the data will be plotted on a graph. A solubility
curve for KNO3 will be constructed by connecting the plotted points.
PURPOSE
Collect the data necessary to construct a solubility curve for potassium nitrate (KNO3) in water.
EQUIPMENT & MATERIALS
Balance
Graduated cylinder, 10-mL
Hot plate
Glass stirring rod
Spatula
Thermometer
Test tube
Beaker, large
Test tube clamp
Distilled water
Safety goggles/Aprons
Potassium nitrate (KNO3)
CLASS DATA – PLACE a STAR next to your group
GROUP
Mass of KNO3
(grams in 5 mL
of water)
Mass of KNO3
(grams/100 mL H2O)
y-axis
Trial 1
Crystallization
Temp
Trial 2
Crystallization
Temp
Trial 3
Crystallization
Temp
AVERAGE Temp of
Crystallization (C)
x-axis
GIVEN
0.675 g
13.5 g
-
-
-
0.0 °C
1
1.0 g
2
1.5 g
3
2.0 g
4
2.5 g
5
3.0 g
6
3.5 g
7
4.0 g
8
6.0 g
GIVEN
7.0 g
140.0 g
-
-
-
100.0 °C
18
CALCULATIONS
Using proportions, convert the experimental mass/5mL ratios in column 2 to equivalent mass/100-mL in column 3.
GRAPH
Plot the data and graph the solubility curve. Your graph must have a title, labeled axes with units, and be NEAT AND LEGIBLE.
Ensure your scale appropriately reflects the data.
ANALYSIS & QUESTIONS
Answer in complete sentences. Show work where calculations are needed.
1. According to your graph, how does the solubility of KNO 3 change as the temperature changes?
2. What is the solubility of potassium nitrate in 100g of water at 50C?
19
3. At what temperature is the solubility of potassium nitrate 65g per 100g water?
4. How much potassium nitrate will dissolve in 50g of water at 60C?
5. According to your graph, how would you make 500 mL of a saturated KNO 3 at 25C?
6. What is the solubility of potassium nitrate at:
a) 70C
b) 40C
7. How much potassium nitrate will crystallize out of the solution when a saturated solution in 100g of water is cooled from 70C to
40C?
8. If a saturated solution of potassium nitrate in 25g of water is cooled from 70C to 40C, what mass of solid potassium nitrate will
crystallize out of solution?
9. Using the graph you created, classify the following KNO3 solutions as saturated, unsaturated or supersaturated.
a.
75 g KNO3 / 100mL H2O at 40oC._______________________________________
b.
60 g KNO3 / 100mL H2O at 50oC. _______________________________________
10. How does the solubility of a gas change with increasing temperature?
11. Draw a rough sketch showing the general shape of a solubility curve of GAS. All graph labels should be present.
20
Name
Period_
LAB – Creating Serial Dilutions –Solutions 4
Objective:
The objective of this lab is to dilute a 1 M solution of HCl to a specific concentration.
Background:
Frequently in lab, you may need several concentrations of the same substance. The most
accurate and user-friendly way of making many concentrations of a single solution is to perform
SERIAL DILUTIONS, or making many sequential dilutions from a single stock solution. This is
faster and more accurate than making every solution from scratch.
Pre-Lab Calculations:
 Use the dilution equation to calculate the initial volume of acid needed for each dilution.
 Subtract the initial volume of the acid from the final volume of solution to determine how much
water is needed to dilute the solution to the final molarity.
 Show ALL work in the space provided below the table.
Dilution Equation: M1V1 = M2V2
Initial Molarity
Initial volume of
ACID
Volume of WATER
to ADD
Final Molarity
Final volume of
the Solution
(M1)
(V1)
(V2-V1)
(M2)
(V2)
Beaker #
1
1M
0.1 M
100 mL
2
0.1 M
0.02 M
100 mL
Show ALL of your work here for Beaker #1.
Show ALL of your work here for Beaker #2.
21
All calculations must be complete before beginning the lab!
Create a procedure for making the dilutions you calculated in the Pre-lab. The items you have access
to (do not take from another group) are:
Equipment:
10 mL Graduated Cylinder
Two 250 mL Beakers
100 mL Graduated Cylinder
Stirring Rod
1M HCl (with dye)
*SAFETY TIP: Recall from early in the year, always mix acid into water (not water into acid)
Write your procedure below.
First Dilution:
What is the concentration of the solution that is now in beaker #1? __________
Second Dilution:
What is the concentration of the solution that is now in beaker #2? __________
Post-Lab Questions:
1. In step #2 of the procedure, you created a solution. Describe the contents of the solution in a sentence.
2. What is the difference between beaker 1 and beaker 2?
3. Describe the conductivity of each solution
4. List three factors that influence the rate at which a solid dissolves in water?
5. You are preparing reagents for an upcoming acid titration lab and need 22.5 L of 3.0 M Ba(OH)2 base
to neutralize the acid at each station. What volume of 12.0 M Ba(OH)2 is needed to make this dilute
base? How much water would you add to this starting volume of the base to create your dilute solution?
22
Name
Period_
Properties of Acids and Bases – Acid/Base 1
Objective: To observe and study some typical properties and reactions of acids and bases
Equipment
large test tubes (4/station)
wooden splints
dropper pipets
spot well plates
stirring rods
small beakers
Materials
distilled water
1.0 M hydrochloric acid
3.0 M hydrochloric acid
Zinc
phenolphthalein indicator
blue/red litmus paper
bromothymol blue indicator
pH paper
methyl red indicator
1.0M sodium hydroxide
1.0M calcium carbonate
1.0M ammonia (NH3)
Vernier pH probes
1.0M acetic acid (HC2H3O2)
Two unknown solutions
Safety



Wear safety goggles and an apron for the duration of the lab
Strong acids and bases are components of the lab, take special care to
avoid contact.
Do not eat or taste any reagents used in this lab
Background
Acids ionize in aqueous solution to produce hydronium ions (H3O+). This ionization is the source of
many unique properties for acids. The strength of an acid depends on the degree to which it ionizes –
strong acids ionize almost completely, while weak acids ionize to a lesser degree. Assuming equal
concentrations, this means that the pH of a strong acid will be lower (more extreme) than the pH of a
weak acid. Similarly, the pH of a strong base will be higher (more extreme) than the pH of a weak
base.
Bases will dissociate in aqueous solution to produce hydroxide ions (OH-). This dissociation is the
source of many unique properties of bases. The strength of a hydroxide base depends on its
solubility in water. Soluble metal hydroxide compounds are strong bases. Ammonia is also a very
important base in chemistry, but it is different in nature from the hydroxides, and it is weak.
In this experiment, you will observe the following:
1. The effects of acids and bases on various indicators. An indicator is a substance that will have a
specific color in the presence of acidic or basic solutions. Indicators can be liquid or paper. A pH
probe can also be used for a quantitative measure of pH.
2. The effect of acid on an active metal. The single replacement reaction can be
represented by the general equation
metal + acid  ionic compound (salt) + hydrogen gas
3. The effect of acid on a metal carbonate. The double replacement reaction can be
represented by the general equation:
metal carbonate + acid  ionic compound(salt) + carbon dioxide gas + water
4. The reaction of an acid and a base. Acids neutralize bases and vice versa according to the
general equation:
acid + base  ionic compound(salt) + water
23
Procedure
Part 1
1) Add 5 drops of each of the following to separate wells in your well plate. (Note which is which!)
1.0 M HCl
1.0 M NaOH
1.0 M HC3H3O2
1.0 M NH3
distilled water
2) Using a different piece of clean, dry red litmus paper for each of the solutions, dip the end of
the red litmus paper in each of the solutions. Note results.
3) Using a different piece of clean, dry blue litmus paper for each solution, dip the end of the blue
litmus paper into each solution. Note results.
4) Add 1 drop of bromthymol blue to the wells. Note results.
5) Place another 5 drops of each solution into different clean wells and add 1 drop of methyl
red to each well. Note results.
6) Pour contents of spot plate down sink and rinse well with water 3 times.
7) Following your teacher’s instructions, use the pH probe to measure the pH of each solution.
Part 2
1) Place a small piece of zinc in a test tube.
2) Add about 1 mL (just estimate – about 15-20 drops) of 3M HCl to the test tube. (BE CAREFUL:
3M HCl is corrosive.) Note observations.
3) Cover the mouth of the test tube for a few seconds with another test tube and test the gas
produced with a burning splint. Note results.
4) Place a small scoop/chunk of calcium carbonate in a clean test tube. Add about 1 mL of 3M HCl and
note observations.
5) Use a burning splint to test the gas produced. Insert the burning splint into the mouth of the
test tube.
Part 3
1) Add 10 drops of 1.0M HCl to a small beaker. Add 1 drop of phenolphthalein indicator and swirl;
test with pH paper. Note results.
2) Slowly add 1.0M NaOH drop by drop to the beaker, swirling to mix and counting drops as you
go. Count the total number of drops of NaOH needed to cause a color change. Note results.
3) Once a color change is observed, test the mixture with pH paper and note results. (Note: a
pH of 7 indicates a neutral solution, neither acidic nor basic.)
Part 4
1) Obtain a sample of the 2 unknown solutions. Using your knowledge, determine whether each
unknown is a strong acid, a strong base, a weak acid, a weak base, or neutral solution.
2) Clean all equipment and tidy up your lab station. All liquids may be rinsed down the sink, and
any leftover paper or zinc should be thrown in the trash. Please rinse all glassware thoroughly!
24
DATA
Part 1: Indicator Results
Red Litmus
Blue Litmus
Paper
Paper
Bromothymol
Blue
Methyl
Red
pH reading
from probe
Classification*
HCl
NaOH
HC2H3O2
NH3
H2O
Parts 2-3
Reaction
Observations
Splint Test
(identity of gas?)
Initial color:
pH Paper
Zn + HCl
CaCO3 + HCl
HCl + NaOH
before NaOH:
# of drops of NaOH added:
after NaOH:
Final color:
Part 4
Unknown #
Unknown #
Classification*
Classification*
* strong acid (SA), weak acid (WA), strong base (SB), weak base (WB), neutral (N) Also read
notes and/or “Background” of lab for assistance in categorizing!
(continued)
25
Questions/Analysis
1. Why is either red or blue litmus alone not sufficient for determining whether a solution is acidic, basic,
or neutral?
2. Write a balanced equation for the reaction between the zinc and hydrochloric acid. Be sure to
include states of all substances.
3. Write a balanced equation for the reaction between calcium carbonate and hydrochloric acid. BE
CAREFUL! This is really 1 reaction followed by another, including a special case. Include states of
all substances.
4. Write a balanced equation for the reaction between hydrochloric acid and sodium hydroxide.
Include states of all substances.
5. Determine whether the following is describing an acid, a base, both, or neither.
a)
produces hydroxide ions in solution
b)
reacts with metals to produce hydrogen gas
c)
pink when phenolphthalein is added
d)
red when methyl red is added
e)
reacts with an acid to produce a salt and water
f)
conducts electricity
g)
turns blue litmus paper red
h)
no change when exposed to blue and red litmus paper
Indicator
Color in acid
Transition Color Color in base
pH range
Phenolphthalein
Colorless
Pink
Red
8.2-10.2
Bromothymol blue
(BTB)
Yellow
Green
Blue
6.0-7.6
Methyl red
Red
Orange
Yellow
4.4-6.2
26
Name
Period_
Tie-Dye Lab – A/B #2
From Flinn CHEM FAX Publication #10075
Introduction
The art of dyeing clothing fibers likely originated in Asia around 2500 BC. Most natural dyes came from parts of plants such as bark,
berries, flowers, leaves, and roots. Because these dyes did not have a strong attraction for the fibers being dyed, a process known as mordating
(using salts to bind to the dyes) was used to improve colorfastness. Such natural dyes became less and less important as synthetic dyes that
produced brighter colors were developed. Today logwood black is the only natural dye widely used.
In 1856, William Henry Perkin began the synthetic organic chemical industry by accidentally discovering the purple dye, mauveine.
The synthetic dyes were known as coal tar dyes because they were all derived from coal tar. Congo Red was the first dye discovered with so
great an affinity for cellulose that a mordant was not required.
About 100 years after Perkin’s first discovery, fiber-reactive dyes capable of forming covalent bonds with the fiber were discovered.
Fiber-reactive dyes are wash-fast (dye does not fade much). During dyeing, dye molecules move from the aqueous solution into the fibers.
Fibers such as cotton absorb water readily and are said to be hydrophilic, while fibers such as polyester absorb water with difficulty and are
described as hydrophobic (this is why you were asked to bring 100% cotton clothing). Dyeability is influenced if a fiber can somehow carry an
ionic charge and better interact with the oppositely charged colored ions. To dye cellulose (cotton), a reactive dye must combine with the
hydroxyl (hydroxide) groups in the fiber.
Tie-dying became fashionable in America in the late 1960s and early 1970s as part of hippie culture. It was popularized in the United
States by musicians such as John Sebastian, Janis Joplin and Joe Cocker.
Materials
T-shirt, 100% cotton, pre-washed as below
Sodium carbonate dye-fixing solution (mixed for you)
Various reactive dyes in urea solution (mixed for you)
Rubber or plastic gloves
Ziploc® bags or small trash bags
Paper towels (on the wall)
Pipets and beakers
Rubber bands
Safety
1.
2.
3.
4.
5.
Students should wear old clothes and shoes. Reactive dyes are “wash fast.” Once the reactive dye makes
contact with clothes it will not wash out.
Hands may become stained from the reactive dye. The dyes are not easily washed off and will take about two
days to wear off hands and skin.
Wear appropriate protective clothing, and chemical splash goggles, disposable plastic gloves and chemicalresistant aprons.
Students must not squirt each other with filled pipets of reactive dye solution. Sodium carbonate activator
solution is very basic and thereby caustic. Students found to be violating any safety precautions may be
removed from lab participation and given a zero for the lab.
Be sure to wear rubber or plastic gloves throughout the lab.
Preparation (days prior to the lab)
1. Before the lab, bring in a white shirt that has been machine washed in 1 tablespoon of a mild liquid dish
detergent like Joy®, Dawn®, or Ivory®. Dry it on a hot setting. This will remove any lubricants or surfactants on
the shirt which may limit dying.
2. Place your name on your pre washed shirt on the neck or tag with a SHARPIE, markers fade.
3. You may create your shirt patterns (see pages 2 and 3) before the lab begins, but be take special care to wring
the activator solution out of the shirt before you begin the day of the lab.
27
Procedure (day of lab)
1.
(THIS STEP HAS BEEN DONE FOR YOU)
Wear gloves at all times during this lab. Your shirt has been soaked in sodium carbonate activator solution. (pH
of 12) for a minimum of 20 minutes. The ionization of cellulose (cotton) increases with increasing alkalinity (basicity)
of the solution and above pH 8 there is an adequate number of ionized hydroxyl groups in the fiber for most dyeing
purposes. Soaking the T-shirts for 2 hours will maximize the number of possible bonding sites. After the T-shirt
has soaked, wear gloves and wring the T-shirt out over a sink or plastic bucket.
(THE ABOVE STEP HAS BEEN DONE FOR YOU)
2.
Wear gloves at all times during this lab. When you arrive to class, you will immediately don your protective gear
and find your shirt in your class’s bag and wring it out over the sink or bag until it is just damp. Return to your
station and fold your shirt as you desire (see pages 2-3), securing all folds with rubber bands. You are ready to
dye your shirt.
3.
Dyeing must be done inside your trash bag to catch excess dye. Dyes are applied to the shirt using pipets. Apply
the dye to one side of the shirt by slowly dripping the reactive dye solution onto each section of the shirt as you
desire. Once one side of the shirt has been dyed, turn the shirt over and repeat the dying process on the other side.
Once dyed, shirts should drain for 15 minutes, if time permits.
4.
The amount of reactive dye is not specific as it depends on how strong a color is desired. Colors like yellow will
need more dye. Remember, some of the dye will wash out when the T-shirt is washed in hot water, so make the
color darker than desired as an end product.
5.
The colors can be mixed in varying quantities to develop other colors.
6.
Wrap the shirt in a few dry paper towels and place it in a plastic Ziploc bag or small trash bag—close or tie the
bag to keep the shirt moist and secure. Most colors will have completely reacted after 4 hours but less reactive
colors such as green and yellow will take as long as 24 hours. Be patient. Let the dyes react completely in the
bag. Do not remove the shirt from the bag for 24 hours (to keep the dyes moist and reactive). Take the shirt home
in the bag and follow the below instructions.
®
When you take the shirt home
1.
At home, roughly 24 hours later, remove the shirt from the bag and rinse it in warm tap water (use the bath tub!) to
remove any unreacted dye and sodium carbonate activator. Change the water and continue to rinse. Repeat until
the water remains clear and the shirt does not feel slippery.
2.
Machine wash your shirt! Set your washing machine on the HOT water setting and wash shirt in two tablespoons
of the pre-wash soap, such as Joy® or Dawn® dish soap. Then dry shirts on the hottest dryer setting. The
reactive dye is washfast so it is now safe to wash with other clothes using normal detergents. NEVER bleach your
tie-dyed shirt!
28
Name
Period_
Lab – Titration of a Strong Acid and Base – A/B #3
Introduction:
Titration is a way to determine concentration of a solution by measuring the volume of that solution
needed to react completely with a volume of a solution of known concentration. We will gradually add the
standard solution (titrant) to a measured quantity of the solution of unknown concentration (analyte) until
the same number of gram-equivalent weights of each solute has been used. This point in the titration is
called the equivalence point. An indicator is used to detect the equivalence point in the reaction.
Phenolphthalein is an indicator used in acid-base titrations because it is colorless in acidic solution and
pink in basic solution.
At the equivalence point, or neutralization, the number of moles of acid equals the number of moles of if
the mole ratio in the balanced reaction is one to one.
In this experiment, you will be given the concentration of a standard HCl solution. You will carefully
measure the volume of base required to neutralize a measured volume of acid. The concentration of the
base can be calculated using the relationship.
Materials:
Three 100 mL beakers
Ring stand
125 mL Erlenmeyer flask
Double buret clamp
Wash bottle
Two 50 mL burets
0.5 M hydrochloric acid
Unknown sodium hydroxide
Phenolphthalein indicator
Universal indicator
Distilled water
Graduated cylinder
Safety: Wear safety goggles and an apron. Follow all instructions.
Procedures:
1. The apparatus shown below is set up at your station. A buret clamp holds labeled acid (HCl) and base
(NaOH) filled burets.
2. At your station are a waste beaker, a flask of distilled water, a graduated cylinder, and labeled beakers of
acid (HCl) and base (NaOH). Do not mix the beakers!
3. Record the start volume of each buret to the nearest 0.01 mL, in the data chart on the back of this page.
4. Carefully dispense about 20 mL of the base (the analyte) into a flask. Next add two or three drops of
phenolphthalein to the flask. Swirl to mix the solution.
5. Next, slowly dispense the acid from the acid buret into the flask containing the acidic
solution while swirling. The goal is for the solution to turn a PALE pink color. This
indicates a neutral solution at the equivalence point. If you add too much base, you will
see a clear color (see step 6)
6. If you add too much acid, slowly, drop by drop, add more base. Go back and forth
between acid and base burets as is necessary until your teacher approves the color you
have created.
7. Once you have reached the equivalence point, record the final volumes of your
burets. Pour all waste into the waste beaker.
8. Refill your burets (to any point between 0 and 5mL) and perform trial number two.
Use the universal indicator instead of the phenolphthalein. You are trying to produce a
green color. Red, orange or yellow is acidic. Blue or purple is basic.
29
Name
Period_
Data – all volumes in mL
Trial #
Indicator Used
Hydrochloric Acid (HCl)
Start
Volume
Final
Volume
Volume Used
(Final – Start)
Sodium Hydroxide (NaOH)
Start
Volume
Final
Volume
Volume Used
(Final – Start)
Trial #1
Phenolphthalein
Trial #2
Universal Indicator
HCl molarity (given by teacher) = ______________ NaOH molarity (#3 below) ______________
Post Lab Questions:
1. Write the chemical equation for the reaction used in this titration. Predict the products and balance.
2. What are the two products of every strong acid & strong base titration (if done properly!)?
3. Calculate the molarity of your unknown sodium hydroxide using the given molarity of HCl, the total
volumes from the data chart, and the coefficients from your balanced equation.
4. How many ml of a 0.035M KOH solution will neutralize 25 ml of a 0.60M HCl solution?
5. What volume of a 0.50M HBr solution is required to neutralize a 100 mL solution of 2.7M Mg(OH)2?
6. Complete the following table.
[H3O +]
pH
[OH -]
pOH
3.98 x 10 –4 M
0.75 M
11.3
30
Acid/Base/Neutral
𝑇𝑖𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝐹𝑜𝑟𝑚𝑢𝑙𝑎:
MA VA MB VB
=
molA
molB
𝐷𝑖𝑙𝑢𝑡𝑖𝑜𝑛 𝐹𝑜𝑟𝑚𝑢𝑙𝑎: 𝑀1 𝑉1 = 𝑀2 𝑉2
𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 =
𝑚𝑜𝑙𝑒𝑠
𝐿𝑖𝑡𝑒𝑟 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
Type of Force
London Dispersion
Dipole-dipole
Hydrogen Bonds
Substances that exhibit
the force
All molecules!
(polar and nonpolar)
Neutral polar molecules
H bonded directly to
N,O, or F
Source of the force
Temporary, inducing dipole
due to electron field
shifting
Strength of bonds increases
with increasing molecular
mass. (Tar/Wax)
Properties
Examples
Permanent dipoles (+ and –
) attraction between two
molecules
Attractions increase with
increase in polarity in
molecules of equal size and
mass.
Interaction between H (+)
and (–) parts of two
molecules
Causes H2O’s high surface
tension. Higher melting and
boiling points
Usually have the lowest
melting and boiling points
Holds H2O molecules
together as a liquid at room
temp
Weakest IM Force
Strongest IM force
N2, C3H8, C20H52, CCl4, all
nonpolar molecules.
Present but negligible in all
other matter.
HCl, CH3OH, SO2,, all polar
molecules not 
31
Molecules containing:
H–O, H–F, or H–N