Lab #1 The Milk Lab by EZe40i5

VIEWS: 285 PAGES: 100

									Date __________________             Lab Partner ___________________
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HOW TO WRITE UP YOUR FORMAL LAB REPORTS                               3

LAB EQUIPMENT                                                         3

LAB #1 SAFETY LAB                                                     6

LAB #2 THE MILK LAB                                                   7

LAB #3 MEASURING MASS                                                 9

LAB #4 DENSITY DETERMINATION                                          14

LAB #5 HEAT OF FUSION OF ICE                                          18

LAB #6 HEATING AND COOLING CURVES                                     23

LAB #7 EMISSION SPECTRA AND ENERGY LEVELS                             30

LAB #8 ALIEN NATION LAB                                               32

LAB #9 FLAME TESTS                                                    34

LAB #10 THREE-DIMENSIONAL MODELS OF COVALENT MOLECULES                37

LAB #11 CONSERVATION OF MASS                                          40

LAB #12 DETERMINING AN EMPIRICAL FORMULA                              43

LAB #13 THE PERCENT COMPOSITION OF A HYDRATE                          47

LAB #14 CHEMICAL EQUILIBRIUM                                          50

LAB #15 ACID-BASE TITRATION                                           53

LAB #16 ELECTROCHEMICAL CELLS                                         57

LAB #17 SYNTHESIS OF ESTERS                                           60

LAB #18 TWIZZLER HALF-LIFE                                            64


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Date __________________                Lab Partner ___________________
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LAB #19 ESTABLISHING EQUILIBRIUM                                         67

LAB #20 WHY IS WATER WEIRD?                                              72

LAB #21 FINDING PATTERNS IN MEASUREMENTS                                 76

LAB # 22 TABLE M                                                         81

LAB # 23 TYPES OF REACTIONS                                              84

LAB #24 OBSCERTAINER                                                     87

LAB 25: TYPES OF SOLUTIONS                                               89

LAB #26 IT'S A GAS EXPLORATION                                           91

LAB # 27 MOLAR VOLUME OF A GAS                                           94




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Date __________________                                      Lab Partner ___________________
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       How To write up Your Formal Lab Reports
                                     How to Write a Formal Lab
        Your Name........................Period___..........................Date of Experiment
         Mr. Pepe...................................................................Date Report is Due
                              Lab Partner ___________________

                            ......................TITLE and EXPERIMENT #


OBJECTIVE: The objective can be found in the Purpose and the Pre-Lab Discussion in
your Lab Manual or in the Handouts provided by the teacher.

OBSERVATIONS: A minimum of 5 observations are required. Number each
observation 1,2,3,4,5, etc. Use complete sentences at all times. Make use of a spell
check or dictionary if necessary. Use information from your Lab Notebook.

DATA: All data is to be written in chart or table form. Always label the chart or table with
a TITLE. (Use capital letters and underline).

CALCULATIONS: List all of the formulas used in this experiment. For each formula that
is used in the experiment, give one example (even though you may have more than one
example from your data). Show all work and be careful to label all the units of
measurement and the units in your answer.

QUESTIONS: First write out the questions that are found at the end of each experiment
in your manual or handout, then answer each question using complete sentences. If
one has already answered this question in the conclusion section, just write, ―See
conclusion 3.‖ etc.

APPENDIX: The lab may require a graph. This is attached to the end of your report. All
graphs will have a TITLE, each axis defined, scales used, and all lines labeled.

All reports will have a complete heading, title, objective, observations, conclusions,
questions, and answers to questions. All reports may not have a data table, graph, or
calculation section. Include these sections, and write N/A in the space.
EXAMPLES: DATA: N/A
CALCULATIONS: N/A




                                     Lab Equipment


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                                    Lab #1 SAFETY LAB




Questions:
1. List 3 unsafe activities shown in the illustration and explain why each is unsafe.

2. List 3 correct lab procedures depicted in the illustration.

3. What should Bob do after the accident?

4. What should Sue have done to avoid an accident?

5. Compare Luke and Duke's lab techniques. Who is following the rules?

6. What are three things shown in the lab that should not be there?

7. Compare Joe and Carl's lab techniques. Who is doing it the correct way?

8. What will happen to Ray and Tim when the teacher catches them?




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Date __________________                            Lab Partner ___________________
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                               Lab #2 The Milk Lab
Introduction
        One of the most important activities of any chemistry investigation is making
observations that help lead to acceptable conclusions. Observations make up an
account of what takes place during an investigation. What is occurring qualitatively and
quantitatively? Record these observations as the investigation is proceeding.
Remember that an observation can be anything that you can detect with your senses
that you did not already know. For example, you can observe that the compound,
copper sulfate, is a blue crystalline solid if you have never seen it before. However, you
would not observe that water is a clear liquid.
        Another component of laboratory investigation is explaining how or why
something occurs. In order to make such ―conclusions‖, research is generally
necessary.
This may involve talking with classmates and checking textbooks or library references.
However, many of the ―conclusions‖ you will make in lab will be concepts that you have
previously studied and learned. The lab environment is a practical way to apply these
concepts in the form of conclusions.

Objective
      Carefully observe, compare, and interpret a series of laboratory events. Form
conclusions about what was observed.

Safety
      1. Wear protective goggles and an apron throughout the laboratory activity.
Be very careful when using the food coloring. It will stain your hands and clothes.

Materials
       Milk, skim milk, water, four different food colors, three shallow plastic dishes, a
Q-tip, and the ―unknown‖ chemical.

Procedure
Pour one of each of the following into three dishes:
whole milk
skim milk
water
   1. At the 3, 6, 9, and 12 o’clock positions in the dishes, place one or two drops of
      the
       four different colors of food coloring.
   2. Using a Q-tip, place a small amount of the unknown liquid provided in the center
      of each dish.
   3. Observe each dish. Record your observations. Make some conclusions.
   4. Wash hands before leaving the lab.



Data


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Date __________________                          Lab Partner ___________________
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      No quantitative data for this lab. Write down N/A.

Calculations
      No calculations for this lab. Write down N/A.

Observations
     Make a minimum of five observations.



Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

1. What might account for what was observed?


2. How can the differences in the three dishes be explained?


3. What was the purpose of the food coloring?


4. In terms of this investigation, what were the differences among the water, the whole
milk, and the skim milk?


5. What was the unknown liquid? Support your answer.




                                          SKIM
   WATER                                  MILK
                                                                               WHOLE
                                                                               MILK




                                           8
Date __________________                           Lab Partner ___________________
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                       Lab #3 Measuring Mass
Introduction
         For laboratory work in a chemistry course, three basic types of measurements
using a lab balance should be mastered. These are: measuring mass directly,
―measuring out‖ a substance, and determining mass by difference. Always remember
that before using any balance, the balance pan must be cleaned and dried, and the
balance itself must be ―zeroed‖ or ―tared.‖
                Measuring mass directly is used to determine the mass of a container or
solid object that does not need to be placed in a container. In a direct measurement,
the object with the mass to be measured is placed on the balance pan. The mass of the
object is read directly from the digital balance display
        Measuring out a substance is used to determine a ―known‖ mass of a chemical.
To do this, a container or even a piece of paper is first placed on the balance and
massed directly (as described above). Once the mass of the container is determined,
the desired mass of the chemical that is required is added to the container on the
balance using a clean scoopula or scoop. This can also be done with a liquid. Instead
of a scoopula, a pipette or eyedropper is used to slowly add liquid to a container. For
data collection purposes, the mass of the empty container, as well as the mass of the
container plus the chemical are both recorded. The mass of the empty container is
often called the tare weight. It is important to note the mass of the container used so
that later in the experiment, when a final mass is measured, the mass of the container
can be subtracted out from the final reading. This is how to calculate the mass of a
product. (See below for ―Determining mass by difference.‖)
        Determining mass by difference is used when an unknown mass of a chemical
needs to be determined. As the name suggests, this is done by subtraction. As in
―measuring out a substance,‖ first a container is massed directly. The mass of the
empty container is recorded. Then the chemical being studied is placed into the
container. The new reading from the digital balance display is recorded for the mass of
the container plus the chemical (This is the same thing done in ―Measuring out a
substance.‖) Lastly, to determine the mass of the chemical, the mass of the container
(tare weight) is subtracted from the mass of the container plus the chemical.
        In this experiment, you will learn how to perform all three types of measurements
described above. In addition, you will gain valuable practice in using a laboratory
balance and in handling different materials and apparatus. Lastly, you must be aware
of the accuracy and precision of an instrument. Therefore, you will apply your
knowledge of significant figures to this lab by reporting the correct number of significant
figures for each measurement.




Objective
      Practice the various techniques of measuring masses using the lab balance.
Gain experience in the techniques of handling laboratory materials and equipment.



                                            9
Date __________________                           Lab Partner ___________________
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Become familiar with the number of significant figures that can be reported based on the
limits of the instruments used in a lab.

Safety
Wear protective goggles and an apron throughout the laboratory activity.
Certain chemicals may cause irritation to sensitive skin. As a precaution, always wash
your hands after leaving the lab.


Materials
lab balance, watch glass, beaker (150-mL), beaker (50-mL),weighing boats, lab apron,
graduated cylinder (100-mL), scoopula, safety glasses, sodium chloride (NaCl), a
penny, and a nickel.

Procedure
Part A – Using a Digital lab Balance
1. Before plugging in a digital balance, make sure that the balance pan is clean and
dried.
Do not use too much pressure. Remember that the digital balance is a sensitive piece
of equipment.
2. Press the Zero On button and allow the balance to ―warm up‖
3. It is often a good idea to ―calibrate‖ the balance to make sure it is working properly.
To do this, hold the Zero On button until CAL is displayed the release it. -C- is
momentarily displayed followed by the value of the mass, which must be placed, on the
pan (200 grams).
4. Place the 200 gram mass on the pan and press the Zero On button.
5. When the mass on the pan is displayed, the balance is calibrated. Remove the
mass from the pan. The reading should go down to 0.00 grams. If it does not, call over
your instructor. NEVER try to fix a digital balance yourself.

Part B – Measuring Mass Directly
6. Check to make sure your balance is clean, dry and ―zeroed.‖
7. Place a penny on the balance pan. Record the mass of the penny.
8. Repeat step 7 for the objects listed below. Record the mass of each object
      a - a nickel                        b - a watch glass
      c - a 150-mL beaker          d - a 100-mL graduated cylinder




Part C – Measuring Out a Substance
Place a weighing boat on the balance pan. Record the mass of the weighing boat.




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Date __________________                           Lab Partner ___________________
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Using a scoopula add sodium chloride (NaCl) to the weighing boat until the digital
display reads 7.50 grams more than the original mass of the weighing boat.. Dispose of
the NaCl. Make a note in your observations what 7.50 grams of NaCl looks like.
Place a dry 50-mL beaker on the balance pan. Record this reading.
In a 100-mL graduated cylinder, obtain exactly 30.0 mL of cold tap water. Slowly and
carefully pour water from the graduated cylinder into the beaker on top of the balance
until 22.00 grams more than the reading obtained in step11. (You may need to use a
pipette to get the exact amount of water into the beaker). Record this new mass in your
data table.
Record the volume of water remaining in the graduated cylinder

Part D - Determining Mass by Difference
Measure and record the mass of a new weighing boat.
Add a small amount (a scoop full) of NaCl to the weighing boat. Record the mass.
    Dispose of the NaCl.
Repeat steps 14 and 15 two more times and record each mass separately.
                          Wash hands before leaving the lab.
Observations
      Make a minimum of five observations about what was learned from using the lab
balance and significant figures.

Data
Part B - Direct Mass Data
Mass of a penny                                     _________ grams
Mass of a nickel                                    _________ grams
Mass of a watch glass                               _________ grams
Mass of a 150-mL beaker                        _________ grams
Mass of a 100-mL graduated cylinder            _________ grams




Part C - Measuring Out of a Substance Data
Mass of the weighing boat                              __________ grams
New setting (weighing boat + 7.50 g)                   __________ grams
Mass of 50-mL beaker                                        __________ grams
New Setting (50-mL beaker + 22.00 grams)               __________ grams
Volume of water remaining in graduated cylinder             __________ mL




Part D - Measuring by Difference Data
Weighing boat #1                      __________ grams
Weighing boat #1 + NaCl        __________ grams



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Date __________________                        Lab Partner ___________________
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Weighing boat #2                      __________ grams
Weighing boat #2 + NaCl          __________ grams
Weighing boat #3                      __________ grams
Weighing boat #3 + NaCl          __________ grams




Calculations
Calculate the volume of water added to the beaker in step 12.      __________mL
      (Initial volume in cylinder - volume remaining in cylinder)
Calculate the mass of 1.0 mL of water.                             __________grams
      (mass of 22.00 grams of water divided by volume of water from calculation #1).
Mass of NaCl from Part D - trial #1                     __________grams
Mass of NaCl from Part D - trial #2                     __________grams
Mass of NaCl from Part D - trial #3                     __________grams
Average Mass of NaCl                                    __________grams


 NOTE: when handing in the formal lab, you must write down the questions as well as
          the answers. They must ALWAYS be in complete sentences.

   1. What measurements are required to determine the average mass of a scoop of
      NaCl?



   2. How would you use these measurements to arrive at an average mass?




   3. The difference between the mass of a balloon when inflated and its mass after
      being punctured is not an accurate determination of the mass of the gas inside
      the inflated balloon. Why not?




   4. Suppose you were asked to measure out exactly 5.00 grams of sodium chloride.
      Briefly describe how you would make this measurement.




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Date __________________                       Lab Partner ___________________
                                                          ___________________
   5. Suppose you wanted to know the mass of a quantity of orange juice that was
      poured into a glass. Describe how you would determine this mass.




   6. A beaker contains a quantity of liquid. You want to know the combined mass of
      the beaker and the liquid. Describe how you would make this determination.




   7. Suppose you were asked to compare the mass of a nickel and a sample of NaCl
      crystals. Which method would you use to find the mass of the nickel?




   8. Which method would you use to find the mass of the NaCl crystals?




   9. Which of these two measurements would be more accurate? Explain.




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Date __________________                               Lab Partner ___________________
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                  Lab #4 Density Determination
Introduction
Density is the amount of matter in a given space. Therefore, to calculate density, one
must find the mass of an object and divide it by that objects volume.

                                       D=M/V
The units for mass are grams, and the units for volume are mL or cm3. (mL and cm3
are considered the same unit and are equal). Hence the units for density are grams/mL
or grams/cm3. Determining mass requires the use of a balance. In this lab, the triple-
beam balance can measure exactly to 0.1 grams. Since the last number of a significant
figure can be estimated, one can use a triple-beam balance to mass to the 2nd decimal
place with reasonable accuracy. (or to the 0.01 decimal place).

Volume can be determined either by displacement of water (in mL) or by calculation of
dimensions.
For example, the calculation for the volume of a rectangular solid is:

                                      V=lxwxh
The calculation for the volume of a cylinder is:

                                       V = πr2h
A 10-mL graduated cylinder can measure exactly to the nearest 0.1 mL; so again one
can use a 10-mL graduated cylinder to measure to the 2nd decimal place with
reasonable accuracy. (or to the 0.01 decimal place). A 100-mL graduated cylinder can
only measure exactly to the nearest 1 mL; so one can use a 100-mL graduated cylinder
to measure to the 1st decimal place with reasonable accuracy.
(or to the 0.1 decimal place). A ruler can measure to the nearest 0.1 cm, so one can
use a ruler to measure to the 2nd decimal place with reasonable accuracy. (or to the
0.01 decimal place). You will be expected to use your skills to determine mass,
calculate volume and density, and calculate your % error.


Objective
Learn and practice techniques for determining the volume and density of substances.


Safety
      Always wear safety goggles and a lab apron or coat when working in the lab.

Materials
digital balance      ruler         plastic cylinder


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Date __________________                           Lab Partner ___________________
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250 mL beaker (for holding water)
10 mL graduated cylinder water            unknown liquid
100 mL graduated cylinder glass rod       wooden solid

Procedure

Part A - Solids
1. Measure and record the mass of a wooden rectangular solid to the nearest 0.01
grams.
(NOTE: Be sure to write down the # of your wooden block in your observations).
2. Measure and record the length, width, and height of a wooden rectangular solid to
the nearest 0.01 cm, using a ruler.
3. Calculate and record the volume of your rectangular solid using the volume formula
for a rectangular solid. This will be recorded under ―Calculated-Volume‖. Remember
your significant figures.
4. Calculate the density of your wooden rectangular solid. This will be recorded under
―Calculated-Volume Density‖ Remember your significant figures.
5. Measure and record the mass of a plastic cylindrical solid to the nearest 0.01 grams.
6. Measure and record the radius and height of a plastic cylindrical solid to the nearest
0.01 cm, using a ruler.
7. Calculate and record the volume of your plastic cylindrical solid using the volume
formula for cylinders. Remember your significant figures. (NOTE: This calculation will
be recorded under your ―Calculated-Volume‖).
8. Calculate and record the density of your plastic cylindrical solid using your calculated
volume. Remember your significant figures. (NOTE: This calculation will be recorded
under your ―Calculated-Volume Density‖).
9. Measure and record the volume of the same plastic cylindrical solid to the nearest
0.1 mL, using a 100 mL graduated cylinder. (NOTE: This measurement will be recorded
under your ―Measured-Volume‖).
Calculate and record the density of your plastic cylindrical solid using your ―Measured
Volume‖
Remember your significant figures. (NOTE: This calculation will be recorded under your
―Measured-Volume Density‖).
11. Measure & record the mass of a piece of glass rod to the nearest 0.01 grams.
 Measure and record the volume of a piece of glass rod to the nearest 0.01 mL using a
10 mL graduated cylinder.
13. Calculate the density of your glass rod using your ―Measured-Volume‖ for the glass
rod.
Remember your significant figures.

Part B - Liquids
14. Record the mass of an empty graduated cylinder.
15. Add exactly 10.00 mL of the unknown liquid to your graduated cylinder.
16. Record the mass of the empty graduated cylinder and the 10.00 mL of the unknown
liquid together.
17. Measure and record the mass of the unknown liquid.



                                            15
Date __________________                             Lab Partner ___________________
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18. Repeat this experiment three times. (three trials) and record the average.
19. Calculate and record the density of the unknown liquid for each trial.
 Determine the % error of the unknown liquid based on the accepted value for the
density of the unknown given by your teacher.


Observations - (Minimum of 5)

   1.

   2.



   3.




Data - Density Data
Part A - Solids
       Object      Mass    Leng Widt Heig Radi Calculated Measured
                   (grams) th   h    ht   us   Volume     Volume
                           (cm) (cm) (cm) (cm) (cm3)      (mL)
       Wooden
       block
       Plastic
       cylinder
       Glass rod


        Object        ―Measured -Volume‖             ―Calculated - Volume‖
                      Density                        Density ............(grams/cm3)
                      ............(grams/mL)
        Wooden
        block
        Plastic
        cylinder
        Glass rod



Part B- Liquids
Substance         Mass of graduated       Mass of graduated        Mass of     Volume   Density


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Date __________________                          Lab Partner ___________________
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                cylinder (grams)       cylinder & unknown       unknown    (mL)        (grams/mL)
                                       (grams)                  (grams)
Unknown trial                                                              10.00
#1                                                                         mL
Unknown trial                                                              10.00
#2                                                                         mL
Unknown trial                                                              10.00
#3                                                                         mL
Average                                                                    10.00
                                                                           mL


Calculations
1. Show Calculations for ―calculated volume‖. (See steps 3 and 7 of Part A)



2. Show Calculations for density. (See steps 4,8,10 and 13 on Part A and Step 19 on
part B).



3. Calculate % error from Accepted value given by teacher for Step 20 - Part B




Questions
1. How can one determine if a solid object would float or sink in water based on its
   density?


2. Of the two methods used to calculate density for the cylinder, which do you believe
   is more accurate? Why?


3. If the density of a 30.0-gram object is 8.93 g/mL, what would be the density of the
   same object if it were cut in half? Why?


4. What was your unknown liquid? Why do you believe this?




                                           17
Date __________________                            Lab Partner ___________________
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                   Lab #5 Heat of Fusion of Ice
Introduction
      When a chemical or physical change takes place, heat is either given off or
absorbed. That is, the change is either exothermic or endothermic. It is important for
chemists to be able to measure this heat. Measurements of this sort are made in a
device called a calorimeter. (see figure 4-1) The technique used in making these
measurements is called calorimetry.




             Figure 4-1


        In simplest terms, a calorimeter is an insulated container made up of two
chambers. The outer chamber contains a known mass of water. In the inner chamber,
the experimenter places the materials that are to lose or gain heat while undergoing a
physical or chemical change. The basic principle on which the calorimeter works is
when two bodies at different temperatures are in contact with one another, heat will flow
from the warmer body to the colder body. Thus, the heat lost by one body will be
gained by the other. This exchange of heat will continue until the two bodies are at the
same temperature. In a calorimeter, heat is exchanged between the water and the
materials undergoing change. The experimenter takes a direct measurement of the
temperature change of the water. From this information, the heat gained (or lost) by the
water can be calculated. The experimenter then uses these data to determine the heat
lost (or gained) by the materials undergoing change.
        Unlike most calorimeters, the simple Styrofoam-cup calorimeter used in this
experiment will have only one chamber. The ice will be placed directly into a measured
amount of water. The heat required to melt the ice will be supplied by the water. By
measuring the temperature change (T) of the water, you can calculate the quantity of
heat exchanged between the water and the ice. Using these experimental data, you will
calculate the heat of fusion of ice.

The following relationships will be used in this experiment:

   Heat lost (or
  gained) by the          Original mass            Change in            Specific heat
                          of the water in   +    temperature of          capacity of
   water in the                                                    +
    container             the calorimeter          the water               water


In symbols, this word formula becomes:


                                            18
Date __________________                             Lab Partner ___________________
                                                                ___________________


                                   Q =m x T x Cp

                 heat given off by the water = heat absorbed by the ice

                                        Qlost = Qgained

    heat needed to melt the ice     =    heat of fusion of ice * mass of the melted ice

The specific heat capacity of a substance is the quantity of heat energy needed to raise
the temperature of one gram of the substance by 1 Celsius. The specific heat capacity
of water is 4.184 joules per gram per degree Celsius (4.184J/gC) or in calorie units
(1.00 cal/gC).

Objective
      Using a simple calorimeter, find the heat of fusion of ice.


Safety
1. Handle the thermometer with care. It is fragile and easily broken.
Tie back long hair and secure loose clothing before working with an open flame.
Always wear safety goggles and a lab apron or coat when working in the lab.



Materials
       beaker, 250mL                                 tongs or perforated spoon
        graduated cylinder, 100mL                      safety goggles
        lab burner                                    lab apron or coat
        Styrofoam cup                                 water
        thermometer                                   ice cubes
        ring stand
        iron ring
        wire gauze




Procedure

1. In a 250mL beaker, heat about 125mL of water to a temperature of 50C.



                                              19
Date __________________                           Lab Partner ___________________
                                                              ___________________
2. Measure exactly 100mL of this heated water in a graduated cylinder and pour it into
   a Styrofoam cup. Record this volume of water V1.
3. Measure accurately and record the temperature of the water T1. Immediately add 2-
   3 ice cubes. See Figure 4-2.




                               Figure 4-2


4. Stir the ice-water mixture carefully with the thermometer. (CAUTION: Thermometers
   break easily.) The cup should contain ice at all times. Therefore if the last of the ice
   appears about to melt, add another ice cube. Monitor the temperature of the ice-
   water mixture as you stir. Continue stirring (and adding ice if necessary) until the
   temperature evens off (no longer drops). Record this final temperature T2.
5. Carefully remove the unmelted ice.
6. Allow any water removed to drain back into the cup.
7. Measure and record the amount of water in the cup V2.


Observations
     Make a minimum of five observations about this experiment.

1.
2.
3.
4.
5.




Data - Calorimetry Data


                                            20
Date __________________                               Lab Partner ___________________
                                                                  ___________________
Initial volume of water in cup =V1=                       ________________
Initial temperature of water in cup = T1=            ________________
Initial mass of water in cup =M1=                    ________________ Use D=M/V
Final volume of water in cup = V2=                        ________________
Final temperature of water in cup = T2=              ________________
Final mass of water in cup =M2=                      ________________ Use D=M/V

PROCESSING THE DATA
1. Use the equation t = t2 – t1 to determine t, the change in water temperature.

2. Subtract to determine the volume of ice that was melted (V2 –V1).

3. Find the mass of ice melted using the volume of melt (use 1.00 g/mL as the density of water).

4. Use the equation given in the introduction of this experiment to calculate the energy (in
   joules) released by the 100 g of liquid water as it cooled through t.

5. Now use the results obtained above to determine the heat of fusion—the energy required to
   melt one gram of ice (in J/g H2O).

6. Use your answer to Step 5 and the molar mass of water to calculate the molar heat of fusion
   for ice (in kJ/mol H2O).

7. Find the percent error for the molar heat of fusion value in Step 6. The accepted value for
   molar heat of fusion is 6.01 kJ/mol.




DATA AND CALCULATIONS
  Initial water temperature, t1                                                            °C

  Final water temperature, t2                                                              °C
  Change in water temperature, t                                                          °C




                                               21
Date __________________                               Lab Partner ___________________
                                                                  ___________________

  Final water volume, V2                                                             mL

  Initial water volume, V1                                                           mL

  Volume of melt                                                                     mL

  Mass of ice melted


                                                                                        g

  Heat released by cooling water (q = Cp•m•t)


                                                                                        J

  J/g ice melted (heat of fusion)


                                                                                    J/g

  kJ/mol ice melted (molar heat of fusion)


                                                                                 kJ/mol

  Percent error


                                                                                        %


Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

   1. List possible sources of error in experiment. How might the use of a calorimeter
      such as the one shown in Figure 4-1 reduce some of these errors?

   2. One source of error is the flow of heat between the water in the cup and the
      surroundings. Explain how this error is reduced by starting with water at about
      50C.

   3. In what way does calorimetry make use of the law of conservation of energy?




   4. Define the following terms:
         a. exothermic
         b. endothermic
         c. heat of fusion


                                                 22
Date __________________                          Lab Partner ___________________
                                                             ___________________
          d. specific heat capacity.




   5. Is the process of melting endothermic or exothermic? Give evidence to support
      your answer.




   6. What is the difference between heat and temperature?




            Lab #6 Heating and Cooling Curves
Introduction
        In this experiment, you will observe by direct measurement the effects of cooling
and heating a pure substance. In Part A, you will start with a substance in its solid
phase at a temperature well below its melting point. While heat is added at a constant
rate, temperature readings will be made until the substance is in its liquid phase at a
temperature well above its melting point.




                                           23
Date __________________                          Lab Partner ___________________
                                                             ___________________
        In Part B, a pure substance will be cooled (heat removed) at a constant rate.
Starting with the substance in its liquid phase at a temperature well above its freezing
point, temperature readings will be made at regular intervals until the substance
changes to its solid phase and cools to a temperature well below its freezing point. The
temperature readings will thus show the effects of removing heat from a pure substance
in the liquid phase, during a phase change (liquid to solid), and in the solid phase.
        The data collected in Parts A and B will be used to construct a graph, which will
consist of two curved lines: a cooling curve and a heating curve. When completed, the
graph will show pictorially what happens to a pure substance as its temperature is
raised and lowered over a temperature interval that includes its freezing and melting
points. The graph also will show how the freezing and melting points of a pure
substance are related.

Objective
      The purpose of this lab is to study the effects of heating and cooling a pure
substance through a change of phase. Construct heating and cooling curves of a
pure substance using experimental data. Determine the freezing and melting point
temperatures of the pure substance.


Safety
        Be careful when using the thermometer to stir the sample. The thermometer is
fragile. Immediately report any breakage to your teacher. Tie back long hair and secure
loose clothing when working with an open flame. Always wear safety goggles and a lab
apron or coat when working in the lab.


Materials
      test tube 18x150-mm       test tube clamp          lauric acid (C12H24O2)
      safety goggles                       water
        lab burner                   lab apron or coat
        beaker, 250mL                 two different colored pencils or pens
        thermometer                  stopwatch (or timer with second hand)
        ring stand                     wire gauze
        iron ring


Procedure
        PART A: HEATING CURVE
   1. Set up a ring stand, iron ring and wire gauze as demonstrated by your teacher.
   2. Get a 400-mL beaker and fill it three-quarters full of tap water. Place it on the
      wire gauze and heat it to a temperature of about 60C. Remove the heat.
      (NOTE: at this point you no longer need to measure the temperature of the water
      bath. You will now only measure the temperature of the lauric acid).
   3. Set the time at 0 minutes and immerse the test tube of lauric acid below the
      water level in the hot water bath and clamp it to the ring stand. Immediately



                                           24
Date __________________                           Lab Partner ___________________
                                                              ___________________
        record the temperature of the lauric acid. Read and record the temperature of
        the lauric acid every ½ minute.
   4.   At this point in the experiment, one partner will call out the time every ½ minute
        and will record temperature data in the Data Table. The other partner will read
        off the temperature of the sample at each half-minute interval.
   5.   As soon as the thermometer is free to move, it should be used to stir the solid-
        liquid mixture.
   6.   Continue stirring and recording the temperature at half-minute intervals until the
        temperature of the sample reaches 50C. (NOTE: Be careful when stirring so as
        NOT to break the thermometer).
   7.   Unclamp the test tube from the ring stand, remove it from the water bath.

PART B: COOLING CURVE
  1. Fill a 400-mL beaker three-quarters full of cold tap water.
  2. Heat the test tube of lauric acid gently until it reaches 70C.
  3. Place the 400-mL beaker of ice water on the wire gauze.
  4. Immerse the test tube into the ice water bath and clamp the test tube of lauric
     acid to the ring stand. Immediately record the temperature of the lauric acid. Set
     the time at 0 minutes.
  5. At this point in the experiment, one partner will call out the time every ½ minute
     and will record temperature data in the Data Table. The other partner will read
     off the temperature of the sample at each half-minute interval. The recording
     partner will call out the time every ½ minute. The second partner will use the
     thermometer to stir the sample constantly as long as some liquid remains. At
     each half- minute interval, this partner will read the temperature to the recording
     partner. Continue this procedure until the temperature of the sample reaches
     35C.
  6. Remove the test tube from the water bath.



Observations
     Make a minimum of five observations about this experiment.

Data - (See Heating and Cooling Data Table.)
(NOTE: The data collection time may be greater or less than the time written in the
table provided).

Calculations
         Plot your data from this experiment on a graph. Title the graph Heating and
Cooling Curves for Lauric Acid. Label the X-axis as Time (minutes) and the Y-axis as
Temperature (C). Make sure to plot both curves on the graph. Label each line with a
different color and provide a key to tell which line is which.

Conclusions
     Make a minimum of five conclusions.


                                            25
Date __________________                          Lab Partner ___________________
                                                             ___________________


Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

   1. Referring to your graph, determine the freezing point of lauric acid. How does
      this temperature compare with the melting point temperature of the same
      substance as indicated on the graph?

   2. Explain the diagonal parts of the cooling curve in terms of changes in kinetic and
      potential energy. Do the same for the horizontal portions of the cooling curve.

   3. What phase changes are exothermic? Endothermic?

   4. In which phase of a substance do its particles have the greatest average kinetic
      energy?




                                           26
Date __________________                      Lab Partner ___________________
                                                         ___________________
Heating and Cooling Data Table
TIME (minutes)   TEMPERATURE (C)     TIME (minutes)   TEMPERATURE (C)
                 HEATING   COOLING                     HEATING  COOLING
                 (part A)  (part B)                    (part A) (part B)

0                                     7

½                                     7½

1                                     8

1½                                    8½

2                                     9

2½                                    9½

3                                     10

3½                                    10½


4                                     11

4½                                    11½

5                                     12

5½                                    12½
6                                     13

6 1/2
                                      13½




                                       27
Date __________________        Lab Partner ___________________
                                           ___________________




                          28
Date __________________        Lab Partner ___________________
                                           ___________________




                          29
Date __________________                           Lab Partner ___________________
                                                              ___________________

   Lab #7 - Emission Spectra and Energy Levels
Objectives:
  1. Investigate the visible light emissions of ―general‖ light sources and
  2. Observe the emission spectra of various elements.

Safety:
      Extreme caution is necessary because the high voltage output from the power
       supply could be lethal. To avoid accidents, turn the supply off when not actually
       in use.
      To preserve the tubes, turn the supply off when not actually in use. Because the
       tubes get hot, always use a cloth to touch the sample tubes or you may burn
       yourself.

Procedure:
   To observe the light emitted from each source, use the spectrometer and:
      1. Hold the spectrometer so the diffraction slit is towards the light source and
         the diffraction grating (window) is towards you.
      2. Raise the spectrometer to your eye. Look through the diffraction grating and
         diffraction slit, and center the light source so that it is visible through the
         diffraction slit…then slightly shift the slit away from the source.
      3. You should see a small spectrum of color, which comes from the light source,
         ranging from purple to red, from left to right, on the inside right side of the
         tube.

                                              Diffraction grating
                 Diffraction slit




       4. Draw the spectrum you see for each light source (natural sunlight, fluorescent
          ceiling lights, and the provided elements). Do the best you can to represent
          the colors you see using the example format below. Put work on back of this
          sheet IN ORDER described in class.
        If you see an individual line, draw a single vertical line of the same color.
        If you see many lines or an area solid with color, draw a box and fill it in with
          that color.


   Example:
   Natural Sunlight




                                            30
Date __________________                          Lab Partner ___________________
                                                             ___________________



Post-Lab Questions:
Use your notes and/or textbook to answer the following questions.
   1. Under what 2 conditions will elements in the gaseous state radiate (or emit) light?



   2. What results when the light emitted by a gas is passed through a prism or
      diffraction grating?



   3. Name 2 characteristics that make each element unique (i.e., different from all
      other elements.)


   4. You are studying the emission spectrum of an unknown gas, and you note that
      the spectrum has 3 visible lines: a purple, a green, and an orange-yellow. Using
      your emission spectra above, which element is likely to be the unknown? How
      do you know?




                                           31
Date __________________                           Lab Partner ___________________
                                                              ___________________


                       Lab #8 Alien Nation Lab



Envelope Number      ____

In a nut shell: We have found an alien nation living the hairs between our eyelashes.
We have identified all the aliens but one.
You have to create a sketch of it.

Purpose:
_________________________________________________________________

Background:
     1.    Why is it important to organize things?
     2.    How did Mendeleev prove his system worked?
     3.    What are three organizing principles of our period table?
     4.    Define them?

Materials: Aliens, pencil, paper, smarts

Method:
Organize your critters so that a pattern is clear and you can locate the missing alien.
Then make a sketch of him or her.

Clues:
Each alien differs in two characteristics from the others.
You will have 3 (three) rows when finished, but each row does not have to have the
same number of critters in it.

Questions:
1.    What does your alien look like? Draw in box.
2.    What are the two ways aliens differ?
   a.
   b.
3.    What do the aliens in a row have in common?

4.     What do the aliens in a column have in common?


5.     Comment how this lab is useful in learning about the periodic table?



                                            32
Date __________________        Lab Partner ___________________
                                           ___________________




                          33
Date __________________                            Lab Partner ___________________
                                                               ___________________

                            Lab #9 Flame Tests
Introduction
        The normal electron configuration of atoms or ions of an element is known as the
―ground state.‖ In this most stable energy state, all electrons are in the lowest energy
levels available. When atoms or ions in the ground state are heated to high
temperatures, some electrons may absorb enough energy to allow them to ―jump‖ to
higher energy levels. The element is then said to be in the ―excited state.‖ This excited
configuration is unstable and the electrons ―fall‖ back to their normal positions of lower
energy. As the electrons return to their normal levels, the energy that was absorbed is
emitted in the form of electromagnetic energy. Some of this energy may be in the form
of visible light. The color of this light may be used as a means of identifying the
elements involved. Such crude analyses are known as flame tests.

       Only metals, with their loosely held electrons, are excited in the flame of a
laboratory burner. Thus, flame tests are useful in the identification of metallic ions.
Many metallic ions exhibit characteristic colors when vaporized in the burner flame. In
this experiment, characteristic colors of several different metallic ions will be observed,
and an unidentified ion will be identified by means of its flame test.

Objective
Observe the characteristic colors produced by certain metallic ions when vaporized in a
flame. Identify an unknown metallic ion by means of its flame test.


Safety

Tie back long hair and secure loose clothing when working with an open flame.
Always wear safety goggles and a lab apron or coat when working in the lab.

Materials
              burner                      apron                 wooden splints
              marking pencil              distilled water       unknown metallic solution
              disposable pipettes         safety goggles

M solutions of:
      NaNO3         (sodium nitrate)
      KNO3 (potassium nitrate)
      LiNO3 (lithium nitrate)
      Ca(NO3)2 (calcium nitrate)
      Sr(NO3)2 (strontium nitrate)
      Ba(NO3)2 (barium nitrate)
      Cu(NO3)2 (copper (II) nitrate)


Procedure



                                             34
Date __________________                           Lab Partner ___________________
                                                              ___________________
   1. Dip a wooden splint into the DW test tube. Then, insert the wet wooden splint
      into the flame of a burner. Observe and record any color produced in the flame
      (if any). Note do not let the wooden splint burn. The color you want to observe is
      the color of the solution not the color of the wood burning. If the wood catches
      on fire, douse it in water and try again with a new splint.
   2. Repeat Step 1, except this time using a metallic solution instead of the distilled
      water. Carefully record any color produced in the flame. NOTE: Color
      description is a subjective observation. Try to describe the color that you see as
      accurately as you can. You may ask your lab partners for their opinion, but you
      decide what color the solutions appear to you.
   3. Repeat Step 1, except this time using the unknown metallic solution. Record the
      color produced in the flame. Try to determine which of the other eight solutions
      the unknown is identical to.
   4. Clean up the lab. Do not throw wooden splints in the sink.
   5. Wash hands before leaving the lab.

Observations
     Make a minimum of five observations about this experiment.
1.
2.
3.
4.
5.

Data - Flame Test Data

Metallic Ion                Color in Flame

Distilled water (no ions)   _________________

Na+1                        _________________

K+1                         _________________

Li+1                        _________________

Ca+2                        _________________

Sr+2                        _________________

Ba+2                        _________________

Cu+2                        _________________

Unknown                     _________________




                                             35
Date __________________                          Lab Partner ___________________
                                                             ___________________
Calculations N/A

Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.
   1. What inaccuracies may be involved in using flame tests for identification
      purposes?



   2. Which pairs of ions produce similar colors in the flame tests?




   3. Explain how the colors observed in the flame test are produced?




   4. Define these terms:
         a. quanta

         b. ground state


         c. excited state



   5. What is a spectroscope?




   6. What is observed if the flame tests are observed through a spectroscope?




                                           36
Date __________________                           Lab Partner ___________________
                                                              ___________________

  Lab #10 Three-dimensional Models of Covalent
                   Molecules
Introduction
        A single covalent bond is formed when two atoms share a pair of electrons.
Each atom can provide one of the electrons of the pair (though there are exceptions). If
the two atoms are alike, they will have the same electro negativity. The difference
between their electro negativities will be zero and the bond is considered nonpolar
covalent. If the atoms are different, then one atom will have a higher electro negativity
and exert a stronger pull on the electrons between them. This makes the bond polar
covalent. If more then one pair of electrons is shared between two atoms, a double or
triple bond may be formed.

        A group of atoms held together by covalent bonds is called a molecule.
Molecules can be either polar or nonpolar. If the bonds are nonpolar, then the molecule
is nonpolar. If bonds are polar, molecules can be either polar or nonpolar. If, for
example, the molecule has symmetry (an even distribution of charge resulting in a
shape that appears even on all sides). Polar molecules do not have symmetry. Some
possible shapes for molecules include: linear, bent, tetrahedral, and triangular
pyramidal.

Background:
Most of our learning is in two dimensions. We see pictures in books and on walls and
chalkboards. We often draw representations of molecules on flat paper. Two-
dimensional representations include electron dot structures and structural formulas. In
electron dot structures, a pair of ―dots‖ (which represents a pair of electrons) is used to
represent a single covalent bond. The hydrogen molecule is shown as H:H. In
structural formulas, a single covalent bond is represented as a straight line. The
Hydrogen molecule is H-H. Although such ―models‖ help us in understanding the
structure of molecules, flat models do not give us the three-dimensional view that is
necessary to truly visualize most molecules. In this experiment, you will build three-
dimensional molecular models and then compare them with the corresponding
structural formulas.
      In covalent molecules there are single, double, and triple bonds between atoms. In
some cases, the molecules are in a chainlike arrangement. At other times, the atoms
arrange themselves in a ring-like structure. Still other molecules are in the form of
branched chains.
      Sometimes a group of atoms may form more than one structure. Thus, a given
molecular formula might represent more than one compound. For example, C2H6O
represents both ethyl alcohol and dimethyl ether, compounds with different structural
formulas and quite different properties. Substances that have the same chemical
formula but different structures are called structural isomers.
      Scientists who are responsible for determining the structure of molecules often
start with molecular model kits like the ones we used in this experiment. Complicated
molecules such as DNA (deoxyribonucleic acid) are most often shown in three-



                                            37
Date __________________                          Lab Partner ___________________
                                                             ___________________
dimensional models. Without these models, we would not understand how the atoms of
the molecules interact.
OBJECTIVES:
1. To construct models of some simple and more complicated covalent molecules
2. To draw structural formulas that show the shape of the molecules.
MATERIALS:
Molecular Model Kit


PROCEDURES:
1. Obtain a molecular model kit and remove the contents. Separate the contents
based on the atoms that are represented. The chart below will help you in identifying
the atoms.
                   ATOM                      Number of Bonding Sites (Holes)
 Hydrogen, Chlorine, Bromine, Iodine,                          1
                  Fluorine
             Oxygen, Sulfur                                    2
                  Nitrogen                                     3
                  Carbon                                       4
In most kits, different colors represent different elements. Record on the Report Sheet
the color that corresponds to each element.

2. Make models of the following molecules. There is only one structural isomer for
each one. To confirm this, try making the atoms combine in some other way and you
will find that once you turn the molecule around, it will be identical to the original
structure. When you are satisfied that you have the correct structure, sketch it .
2a. Water, H2O

2b. Methane, CH4

2c. Methanol, CH3OH

2d. Carbon tetrachloride, CCl4

3. For the following molecular formulas, there can be more than one arrangement of
the atoms. For each one, try to find as many different structural isomers as you can.
Draw a structural formula.
3a. Butane, C4H10

3b. Butanol, C4H9OH

3c. Hexane, C6H14

3d. Dichloroethane, C2H4Cl2




                                           38
Date __________________                           Lab Partner ___________________
                                                              ___________________
4. All of the preceding molecules include only single bonds. In the following group of
molecules there are single, double, and triple bonds. Make each structure, and then
determine if any other isomers exist. Finally, draw the structural formula. for each of the
molecules on the Report Sheet.

4a. Carbon dioxide, CO2

4b. Nitrogen, N2

4c. Butene, C4H8

4d. Butyne, C4H6

5. Continue building molecules as your teacher suggests. Sketch each of the models
or draw the structural formulas on the Report Sheet.
              ________________


Calculations N/A


Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

   1. Which molecules were nonpolar because all bonds were nonpolar?




   2. Which molecules had polar covalent bonds but were nonpolar because of
      symmetry?




   3. Which two shapes appeared to produce polar molecules every time?




   4. Name two types o f substances that do not contain molecules with straight
      bonds?




                                            39
Date __________________                           Lab Partner ___________________
                                                              ___________________

                 Lab #11 Conservation of Mass
Introduction
       Matter cannot be created or destroyed by a chemical change. This very
important principle is known as the law of conservation of mass. This law applies to
ordinary chemical reactions (as opposed to nuclear reactions, in which matter can be
changed into energy). During a chemical change (reaction), the atoms of one or more
substances (reactants) simply undergo some ―rearrangements.‖ The result of these
rearrangements is the formation of new, different substances (products). All of the
original atoms are still present. It is because of the law of conservation of mass that we
are able to write balanced equations. Such equations make it possible to predict the
masses of reactants and products that will be involved in a chemical reaction.
       In this experiment, aqueous solutions of three different compounds will be used
to produce two separate and distinct chemical reactions. The fact that changes occur
during each reaction will be readily observable. The balanced chemical equations for
the two reactions are:
        Na2CO3 (aq) + CaCl2 (aq)                2NaCl (aq) + CaCO3 (s) (Eq. 1)
       CaCO3 (aq) + H2SO4 (aq)                 CaSO4 (s) + H2CO3 (aq) (Eq. 2)

                                   H2O (l) + CO2 (g)

Note that a chemical equation can not be observed, but can be a conclusion. Equation
1 shows a double replacement reaction. Equation 2 is also a double replacement
reaction followed immediately by a decomposition reaction of carbonic acid (H2CO3(aq))
into water and carbon dioxide. The combined masses of the three solutions (and their
containers) will be measured before and after each reaction has occurred. This
experiment should give you a better understanding of the law of conservation of mass
and its importance in chemistry.

Objective
      Determine experimentally whether mass is conserved in a particular set of
chemical reactions.

Safety

Always wear safety goggles and a lab apron or coat when working in the lab.
Handle acid with extreme caution. Report all spills to the teacher.

Materials
Erlenmeyer Flask (50-mL) digital balance                1 M Na2CO3(aq) solution
graduated cylinders (10-mL)      safety glasses and apron 1 M CaCl2(aq) solution
2 test tubes (13 x 100-mm) stopper for flask      1 M H2SO4(aq) solution
100-mL graduated cylinder labels and marking pencil 100-mL beaker

Procedure



                                            40
Date __________________                          Lab Partner ___________________
                                                             ___________________
   1. In a graduated cylinder, measure exactly 25.0 mL of sodium carbonate
      (Na2CO3) solution. Pour into a clean, dry 50-mL Erlenmeyer flask. Stopper the
      flask.
   2. Measure exactly 7.50 mL of 1M calcium chloride (CaCl2) solution and pour into a
      clean, dry test tube. Label the tube and place it in a 100-mL beaker so that it
      does not tip over and spill. Rinse the graduated cylinder.
   3. Repeat step 2 with 7.50 mL of 1 M sulfuric acid (H2SO4) solution. CAUTION:
      Handle acid with care. Label the tube and place it in the same 100-mL beaker
      so that it does not tip over and spill.
   4. Place the stoppered flask and the beaker containing the two test tubes on the lab
      balance and record the combined mass in your data table.
   5. Remove the flask from the balance. Pour the CaCl2 solution from the test tube
      into the flask. Swirl the flask to thoroughly mix the two solutions. Record your
      observations.
   6. Restopper the flask and place it back on the balance along with the empty test
      tube. Record this combined mass in your data table.
   7. Once again, remove the flask from the balance. Pour the H2SO4 solution from
      the test tube into the flask. Swirl the flask to thoroughly mix the three solutions.
      Make sure the reaction has stopped before restoppering the flask. Record your
      observations.
   8. Restopper the flask and place it back on the balance along with the empty test
      tube. Record this combined mass in your data table.
   9. Wash hands before leaving the lab area.


Observations
     Make a minimum of five observations about this experiment.

Data - Mass Data
Mass of containers and solutions before mixing                    __________ grams
Mass of containers and solutions after mixing Na2CO3 and CaCl2 __________ grams
Mass of containers and solutions after mixing all three solutions __________ grams

Calculations - N/A




                                           41
Date __________________                          Lab Partner ___________________
                                                             ___________________
Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

   1. What indications were there that a chemical reaction took place during the
      experiment?


   2. Why were you instructed to leave the flask unstoppered after adding the sulfuric
      acid? What might have happened if this were not done?

   3. Compare the three masses from your data. Account for any differences.




   4. In your opinion, does this experiment verify your law of conservation of mass?
      How might this experiment be improved to bring its results more in line with that
      law?




   5. Discuss how the law of conservation of mass relates to the balancing of chemical
      equations.




   6. When you burn a log in a fireplace, the resulting ashes have a mass less than
      that of the original log. Account for the difference in mass. Does this violate the
      law of conservation of mass?




                                           42
Date __________________                          Lab Partner ___________________
                                                             ___________________

      Lab #12 Determining an Empirical Formula
Introduction
       In a sample of a compound, regardless of the size of the sample, the number of
moles of
one element will be directly proportional to the moles of another element in the same
compound. Furthermore, the ratio between the two elements will always be in small
whole numbers. For example, the ratio of the number of moles of hydrogen to moles of
oxygen in a sample of water will always be 2:1, regardless of how much water is in the
sample.
       In this experiment you will determine the mole ratio of magnesium to chlorine in
the compound, magnesium chloride. You will do this by gravimetric analysis. This
means you will find the mass of the magnesium and chlorine in your compound, then
convert those masses into moles. Your mole ratio should work out to be a small whole-
number ratio. This whole-number ratio will define the simplest formula for magnesium
chloride. Another term for the simplest formula of a compound is the empirical formula.

The reaction between magnesium and hydrochloric acid is:

                   Mg (s) + 2 HCl (aq) -------> MgCl2 (aq) + H2 (g)

Objective
        Using mass relationships show that magnesium and chlorine combine in a
definite whole-     number ratio by mass, thereby determining the empirical formula of
magnesium chloride.


Safety
   1. Wear protective goggles and an apron throughout the laboratory activity.
   2. Be very careful when using the hydrochloric acid. It can cause serious chemical
      burns.
   3. Be very careful when working with a burner. Tie back hair, roll up loose clothing,
      and keep all flammables away from the burner.
   4. Never touch a piece of lab equipment after it has been heated until you are sure
      it has cooled down enough so that it will not burn you.


Materials
Ring stand, iron ring, wire gauze, evaporating dish, 125 mL Erlenmeyer flask, burner,
flint striker, scissors, crucible tongs, balance, safety glasses, apron, magnesium ribbon,
sand paper, ruler, watch glass 6M HCl




Procedure


                                            43
Date __________________                           Lab Partner ___________________
                                                              ___________________


1. Set up a ring stand, iron ring and wire gauze as demonstrated by your instructor.
   Place a clean dry evaporating dish on the wire gauze. Heat for five minutes. Allow
   the dish to cool for one minute.

2. Place a clean dry watch glass on a balance. Then, using crucible tongs, pick up the
   evaporating dish from the wire gauze and place it on the watch glass which is on the
   balance. Find the combined mass of the watch glass and evaporating dish. This is
   mass (a) in your data. After recording the mass, use the crucible tongs to place the
   evaporating dish back on the wire gauze set up. Let cool for three more minutes.
   (You may proceed with Step 3 while the dish cools).

3. Record the mass of a clean dry 125 mL Erlenmeyer flask. This is mass (b) in your
   data.

4. Get approximately 70 cm of magnesium ribbon which has been scraped with sand
   paper to remove any magnesium oxide coating. Cut the ribbon into pieces of no
   bigger then 3 cm with a pair of scissors. Place the pieces of ribbon into the 125 mL
   flask. Record the mass. This is mass (c) in your data

5. IN THE HOOD, take the flask with the magnesium pieces and SLOWLY add 10 mL
   of 6M HCl to the flask in 5 mL increments. (WARNING: Flask will get very hot and
   may produce noxious vapors as well as hydrogen gas). Wait for the reaction to be
   completed. If any residual magnesium remains, add a couple more milliliters of acid.
   Wait a few minutes for the reaction to cool before proceeding to step 6. Note
   everything that happens in this reaction in your observations.

6. After the evaporating dish has cooled, carefully pour the acid solution into the
   evaporating dish on the wire gauze (NOTE: The evaporating dish must be have
   been cooling for at least three minutes prior to adding this solution). Place the watch
   glass (curve side down) on top of the evaporating dish.

7. GENTLY heat the evaporating dish with the acid solution until all of the water boils
   away. When the crackling sound has stopped, this indicates that all of the water has
   gone.

8. After all of the water has been evaporated, remove the watch glass using the
   crucible tongs and run the watch glass through the burner flame once or twice, to
   remove any residual water on the watch glass. Note that there will be some salt
   residue still on the watch glass. Do not remove the residue; this is part of your
   product. Place the Evaporating dish, curved side up, on the lab bench to allow it to
   cool.

9. Wait three minutes after you have finished heating the evaporating dish, then, using
   the crucible tongs, place the evaporating dish on the balance. Then place the watch
   glass on top of the evaporating dish that is sitting on the balance. Record the mass



                                            44
Date __________________                           Lab Partner ___________________
                                                              ___________________
   of the dish, the watch glass and the new compound (magnesium chloride). This is
   mass (d) in your data

10. Do the calculations and determine the empirical formula of magnesium chloride.

11. Wash hands before leaving the lab.


Data -         Magnesium Reaction Data

a. Mass of evaporating dish and watch glass:           _______________ grams

b. Mass of Erlenmeyer flask:                           _______________ grams

c. Mass of flask + Mg ribbon:                          _______________ grams

d. Mass of dish, watch glass and product:              _______________ grams


Observations (minimum of 5)


Calculations

1. Mass of magnesium:                            _______________ grams (c. - b.)

2. Mass of magnesium chloride:                   _______________ grams (d. - a.)

3. Mass of chlorine reacted:                     _______________ grams (2. - 1.)

4. Moles of magnesium reacted: __________ moles of Mg = Mass of Mg x 1mole of Mg
                                                                     24.3 grams

5. Moles of chlorine in product:   ___________ moles of Cl = Mass of Cl x 1mole of
Mg (35.5 grams)

6. Ratio of Cl to Mg in product:   ______________      (5. divided by 4.)

7. Empirical formula of magnesium chloride:      _________________




Questions




                                            45
Date __________________                          Lab Partner ___________________
                                                             ___________________
1. a) What was the empirical formula of magnesium chloride based on your
experiment.


   b) Based on the oxidation numbers of Mg and Cl, was this what you expected?
Explain.


2. a) What was the ratio of the mass of Mg used to the mass of Cl reacted?



   b) Why is this ratio not the same as the mole ratio? Relate this to the Law of
multiple proportions.

3. a) In a chemical formula what do the subscripts represent if the formula represents
one molecule?


   b) What do the subscripts represent if the formula represents one mole?



4. The molecular formula of hydrogen peroxide is H2O2. What is its empirical formula?




5. A sample of sulfur having a mass of 1.28 grams combines with oxygen. The new
compound has a total mass of 3.20 grams. What is the empirical formula of this
compound?




                                           46
Date __________________                           Lab Partner ___________________
                                                              ___________________




  Lab #13 The Percent Composition of a Hydrate
Introduction
Hydrates are ionic compounds (salts) that have a definite amount of water bound to
them as part of their molecular structure. The water is chemically combined with the salt
in a definite ratio. Ratios vary in different hydrates but are specific for any given
hydrate. An example of a hydrate is: Na2CO3 .10H2O. This formula expresses how
each molecule of Na2CO3 (sodium carbonate) is surrounded by 10 molecules of water.
The raised dot (.) means that the molecules of water are loosely bound to the salt
molecule. The coefficient (in this case a 10) stands for the number of molecules of
water bound to each molecule of salt. Though the coefficients differ for different
hydrates, the coefficient for any specific hydrate will always be the same. This
illustrates the law of definite composition.
When a hydrate is heated, the ―water of hydration‖ is evaporated. The remaining solid
is called an anhydrous salt. The general reaction for this is:
hydrate  anhydrous salt + water
The percent of water in a hydrate can be found experimentally by accurately
determining the mass of the hydrate and the mass of the anhydrous salt after heating.
The difference between these two masses is the masses of the water that was
evaporated. The percentage of water can then be calculated:

                          % water = mass of water x 100
                                     mass of hydrate

In this experiment, a hydrate of copper sulfate will be studied (CuSO4 . xH2O). Note
that x indicates an unknown number of moles of water that you will have to determine in
this lab. When the hydrate of copper sulfate is heated, the anhydrous form of copper
sulfate is produced. This reaction involves a color changed from blue to almost white:
                          CuSO4 . xH2O  CuSO4 + x H2O
                                 blue           white

     This lab should help you understand hydrates and the law of definite
composition.

Objective
Determine the percentage of water in a hydrate.

Safety
Tie back long hair and secure loose clothing when working with an open flame. Always
wear safety goggles and a lab apron or coat when working in the lab.

      Materials
            balance                      safety glasses             CuSO4 . xH2O


                                           47
Date __________________                              Lab Partner ___________________
                                                                 ___________________
              burner iron ring               ring stand
              evaporating dish               lab apron or coat         scoopula
              crucible tongs                 wire gauze

Procedure
1. Set up a ring stand, iron ring, and wire gauze for the purpose of heating (as
   demonstrated by your teacher). Place a clean evaporating dish on the wire gauze
   and heat for 5 minutes (using a burner) to remove any water that may have been
   absorbed. Allow the dish to cool for three minutes.
2. Using crucible tongs, pick up the evaporating dish from the wire gauze and place it
   on the balance. Find the mass of the evaporating dish. This is mass (a) in your
   data.
3. With the evaporating dish on the balance, measure into it exactly 2.00 grams of
   copper sulfate hydrate. Record the combined mass of the hydrate and the
   evaporating dish as mass (b) in your data.
4. Place the evaporating dish and the hydrate on the wire gauze. Gently heat the
   evaporating dish. Try to avoid any popping or spattering of the substance in the
   dish. You don’t want to lose any of your product). If the reaction becomes to
   violent, remove the heat for a time, then replace the burner.
5. Continue heating until the entire sample loses its color and becomes an off-white or
   light shade of green. Record the color change in your observations. If there are any
   ―caked‖ portions at this time, try o break them up with your scoopula without
   removing any of the sample.
6. Heat strongly for an additional five minutes.
7. Allow the evaporating dish to cool for only a minute. Using crucible tongs, pick up
   the evaporating dish from the wire gauze and place it on the balance. Find the mass
   of the evaporating dish and the anhydrous salt. This is mass (c) in your data

Observations - (minimum of 5)

Data -       Mass Data
evaporating dish                                    _________g

evaporating dish + hydrate                          _________g

evaporating dish + anhydrous salt                   _________g
Calculations

Find the mass of the hydrate used (b) – (a).        _________grams

Find the mass of the water lost (b) – (c).          _________grams

Find the percentage of water in the hydrate:        _________ %
% water = mass of water x 100
             mass of hydrate




                                               48
Date __________________                           Lab Partner ___________________
                                                              ___________________
Questions -
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.
1. The true value for the percentage of water in this hydrate is 36.0%. Calculate your
   experimental error.


2. Why must you allow the evaporating dish to cool before measuring its mass?




3. Why must you measure the mass of the anhydrous salt immediately upon cooling?




4. Were the class’s results similar to yours? If so, explain how your class’s results
   support the law of definite composition.




5. The molar mass of the anhydrous salt (CuSO4) is 160 grams. The molar mass of
   water is 18 grams. Using your calculations from the calculations sections, calculate
   the moles of CuSO4 and moles of water produced in this experiment.




6. Using your answers from question #5, calculate the molar ratio of CuSO4 to water in
   your sample. Finally, determine the exact formula of the hydrate in this experiment.
   Another words, Find the value for the x in CuSO4 . xH2O.




                                            49
Date __________________                            Lab Partner ___________________
                                                               ___________________

                  Lab #14 Chemical Equilibrium
Introduction
In most of the chemical reactions studied in class, the reactants are usually ―used up‖ to
form products. When this happens, the reaction is said to have ―gone to completion.‖
These ―one way‖ reactions are not the only type of reactions that occur. Some
reactions can go backward as well. These reactions are called reversible reactions. In
a reversible reaction the rate of the forward reaction must be equal to the rate of the
reverse reaction. When this occurs, the reaction is said to be in chemical equilibrium.
        Chemical equilibrium can be affected by several factors which include:
temperature, pressure (if gases are involved), and concentration of chemicals.
According to Le Chatelier’s principle, when a stress is placed on a system at
equilibrium, the equilibrium will shift in the direction that tends to relieve that stress.
Equilibrium will then be reestablished at a different point, with different concentrations of
reactants and products.
        In this experiment, we will study how changing the concentration of a reactant will
shift the equilibrium of three simultaneous reactions. First we will study the dissociation
of copper (II) sulfate:

    CuSO4 (aq)  Cu+2 (aq) + SO4-2(aq)                                     (Equation 1)

The addition of ammonium hydroxide (NH4OH) to this reaction causes two
simultaneous equations to occur. After the ammonium hydroxide dissociates
(NH4OH(aq)  NH4+1 (aq) + OH-1(aq) ), The two ions react with the copper(II)
sulfate in the following way:

       Cu+2 (aq) + 2 OH-1(aq)        Cu(OH)2 (s)                       (Equation 2)

    Cu+2 (aq) + 4 NH4+1 (aq)        Cu(NH3)4+2 (aq)      + 4 H+1(aq)      (Equation 3)

The addition of NH4OH will cause the addition of the OH-1 to Equation 1. This will not
only cause the production of a precipitate, Cu(OH)2 (s) (See Equation 2), but will
remove Cu+2 ions from the original solution. This will force the equilibrium of Equation
1 to shift to the right. In other words, the copper (II) sulfate (CuSO4 (aq)) will produce
more Cu+2 (aq) + SO4-2(aq) so that the stress on the system will be relieved.

As more NH4OH as added to the solution, NH4+1 will be added to Equation 1. This will
also force the equilibrium of Equation 1 to shift to the right by removing Cu+2 ions from
the original solution. However, this shift will be more pronounced because the formation
of Cu(NH3)4+2 (See Equation 3), results in a dark blue solution. Again,, the CuSO4 (aq)
will produce more Cu+2 (aq) + SO4-2(aq) so that the stress on the system will be
relieved.
Finally, we will show how the equilibrium can be shifted back to the left in Equation 1 by
adding more sulfate ion, SO4-2, which relieves the stress on the equilibrium by producing
more CuSO4. In order to do this, we will add sulfuric acid (H2SO4) to the solution.
Sulfuric acid dissociates according to the following reaction:



                                             50
Date __________________                          Lab Partner ___________________
                                                             ___________________
           H2SO4(aq)  2 H+1 (aq) + SO4-2(aq)                     (Equation 4)

The addition of an ion already involved in an equilibrium reaction is called the Common
Ion Effect. The sulfate ions from the sulfuric acid add to the right side of Equation 1
causing the equilibrium to shift back to the left.

Objective
       The purpose of this experiment is study equilibrium systems of reversible
reactions and their responses to stress as described by Le Chatelier’s Principle..



Safety
Always wear safety goggles and a lab apron or coat when working in the lab.
Do not smell Ammonia directly. Waft the vapors to your nose.

Materials
4 test tubes (18x150-mm)         safety goggles        2 disposable pipettes
test tube rack                  lab apron or coat             1 M sulfuric acid
 10-mL graduated cylinder 0.1 M copper(II) sulfate
10% ammonium hydroxide

Procedure
1. Clean four test tubes and label them 1, 2, 3 and 4.
2. Using a clean 10-mL graduated cylinder, pour 3.0 mL of 0.1 M CuSO4 to each of
   the test tubes. Observe the color of the solutions.
3. Add 5 drops of ammonium hydroxide solution to test tube #2. Shake the test tube.
   Observe the reaction. Note any changes.
4. Using a clean 10-mL graduated cylinder, add 2.0 mL of ammonium hydroxide
   solution to test tubes #3 and #4. Shake the test tubes. Compare this reaction to that
   of test tube #2.
5. To test tube #4, now add 10 drops of 1M sulfuric acid. Shake the test tube. Observe
   the reaction that occurs.
6. To test tube #4, now add an additional 1.0 mL of 1M sulfuric acid. Observe the
   reaction.

Observations
     Make a minimum of five observations about this experiment.

Data - N/A

Calculations - N/A


Questions




                                           51
Date __________________                            Lab Partner ___________________
                                                               ___________________
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

1. Using Le Chatelier’s principle, explain how the addition of ammonium hydroxide
   affects the concentration of Cu+2 ions in solution. Which way does the equilibrium
   shift when you add the ammonium hydroxide to copper(II) sulfate (Equation 1) and
   why?




2. What is the Common Ion Effect?




3. Explain how the addition of sulfuric acid affects the equilibrium for the dissociation of
   copper(II) sulfate reaction. Which way does the equilibrium shift when you add the
   sulfuric acid to the copper(II) sulfate (Equation 1) and why?




                                             52
Date __________________                              Lab Partner ___________________
                                                                 ___________________


                     Lab #15 Acid-Base Titration
Introduction
In a chemistry lab, a process known as an acid-base titration can determine an
unknown concentration of an acid or base. In this process, a standard solution is used
to neutralize a precise volume of another solution with an unknown concentration. The
standard solution is a solution of a precisely ―known‖ concentration. In order to
determine when a solution has been neutralized, one or two drops of an ―indicator‖ is
needed. When carrying out an acid-base titration, you must be able to recognize when
to stop adding the standard solution (when you have neutralized the other solution).
This is the purpose of the acid-base indicator. A sudden change in color will signal that
neutralization has occurred. At this point, the number of hydronium ions (H3O+) from
the acid will be equal to the number of hydroxide ions (OH-) from the base. The point at
which this occurs is called the endpoint. At the endpoint, the volume of the solution with
the unknown concentration is measured. From that information, the concentration can
then be determined.
The following steps tell how to calculate the unknown concentration of an acid or base:
Write the balanced equation for the reaction. From the coefficients, determine how
many moles of acid reacts with moles of base. (Remember stoichiometry).
If the mole ratio is 1:1, use the following relationship to calculate the unknown
concentration:
                              Ma x Va = Mb x Vb

where Ma = molarity of the acid solution
      Mb = molarity of the base solution
      Va = volume of the acid solution
      Vb = volume of the base solution

If the mole ratio is not 1:1, alter the original equation by multiplying the left side of the
equation by the number of H+ ions that the acid would produce, and multiply the right
side of the equation by the number of OH- ions that the base would produce:
                       (# of H+ ions) x Ma x Va = Mb x Vb x (# of OH- ions)

Most of the lab work you have done in the past has required you to calculate mass
relationships. These types of quantitative experiments are known as gravimetric
analysis. Titrations require the technique of calculating volume relationships. This
technique is called volumetric analysis. This experiment will help you understand the
properties of acids and bases, neutralization reactions, and titrations.

Objective
       Determine the molarity of a NaOH solution of unknown concentration by titrating it
with a standard HCl solution (0.100 M).

Safety




                                               53
Date __________________                           Lab Partner ___________________
                                                              ___________________
       Always wear safety goggles and a lab apron or coat when working in the lab.
Follow all the safety precautions for working with acids and bases. Wash hands before
leaving the lab.

       Materials
       two 50-mL burets            ring stand  double buret clamp
       safety glasses              10-mL graduated cylinder lab apron
       50-mL Erlenmeyer flask      pipet phenolphthalein indicator
       0.100 M HCl solution        NaOH solution (unknown) distilled water




Procedure

1. Before you start the lab, two burets will be set up for you. One of the burets will
   contain 0.100 M HCl solution and it will be labeled as ACID. The other buret will
   have an unknown concentration of base in it and will be labeled as BASE. Your
   teacher will demonstrate how to adjust, use and read the burets properly. Pay very
   close attention to the operational procedure of these burets and make observations
   on their usage. Burets are highly sophisticated pieces of laboratory equipment.
2. Place a 50-mL Erlenmeyer flask under the acid buret. Record the starting volume of
   your HCl in your data table. (NOTE: if this is you first trial then your starting volume
   is 0.00 mL)
3. Carefully turn the valve on the acid buret to the open position and allow EXACTLY
   10.00 mL of 0.100 M HCl to enter the flask. Record your final volume of HCl added
   in your data table. (NOTE: if this is you first trial then your final volume should be
   exactly 10.00 mL)
4. Using a 10-mL graduated cylinder, add 10.0 mL of distilled water to the same flask.
5. Now add 1 drop of phenolphthalein indicator to the flask. Swirl the flask to mix all
   three ingredients.
6. Now place the flask under the base buret and record the starting volume of NaOH..
   (NOTE: if this is you first trial then your starting volume is 0.00 mL)
7. Carefully turn the valve on the base buret to the open position and slowly add the
   NaOH solution drop by drop. As you are adding the base, swirl the flask gently.
   You will notice a pink color appear and quickly disappear each time a drop of base is
   added. Record these observations. You want to continue to add base until only the
   ―faintest pink‖ color remains for about 30 seconds. At that point you have reached
   your endpoint. Record the final volume of NaOH added at that point. (NOTE: if you
   ―overtitrated,‖ record the final volume anyway and try to be more accurate on your
   next trials.
8. After you have reached your endpoint and recorded your final volumes, discard the
   solution in your flask and thoroughly clean it. Prepare for another trial.



                                            54
Date __________________                                Lab Partner ___________________
                                                                   ___________________
9. Repeat steps 2 through 8 at least two more times. Try to get at least three good
   titrations (no significant overtitrations). You do not have to refill the burets each
   time. Simply read and record the initial and final volumes of the solutions in the
   burets before and after each trial.


Observations - (minimum of 5)


Data -            Acid-Base Titration Data

                          Trial 1            Trial 2            Trial 3          Trial 4

                     HCl       NaOH     HCl      NaOH       HCl     NaOH     HCl     NaOH
                     (mL)      (mL)     (mL)     (mL)       (mL)    (mL)     (mL)    (mL)



Initial reading

Final reading

Volume used




Calculations

For each trial, calculate the molarity of the NaOH solution using the given formula.
Determine the average molarity from your trials (disregard any inaccurate trials).
Ma x Va = Mb x Vb


Trial 1: _____________

Trial 2 _____________                        Average Molarity: _____________

Trial 3 _____________

Trial 4 _____________




                                               55
Date __________________                           Lab Partner ___________________
                                                              ___________________
Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.
1. How reproducible were your results?




2. Define these terms: standard solution, titration, end point, gravimetric analysis, and
   volumetric analysis.




3. If 30.0 mL of a 0.500 M solution of KOH was needed to neutralize 10.0 mL of HCl,
   what was the molarity of the HCl?




4. How many mL of 0.100 M NaOH would be needed to titrate 20.0 mL of 0.100 M
   H2SO4? Use a balanced equation to show the neutralization and to explain your
   calculations. (NOTE: H2SO4 gives off 2 H+ ions in solution).



5. People can use white vinegar in preparing foods without danger to their skin or
   internal organs. This is because white vinegar is a ―weak‖ acid. Explain what is
   meant by a weak acid and why white vinegar is safe to use in cooking.




                                            56
Date __________________                            Lab Partner ___________________
                                                               ___________________

                 Lab #16 Electrochemical Cells
Introduction
        Oxidation is when an atom loses electrons. Reduction is when an atom gains
electrons. A chemical reaction where one species loses electrons and another species
gains electrons is called a redox reaction. In such reactions, the substance that gets
oxidized (loses electrons) is called the reducing agent. The substance that gets
reduced (gains electrons) is called the oxidizing agent. In many redox reactions, there
is a complete transfer of electrons from the substance being oxidized to the substance
being reduced. When these electrons travel through a conductor, such as a wire,
electric current can be generated and measured in volts. In this arrangement, there are
two half-reactions. There is an oxidation half-reaction and a reduction half-reaction. In
this lab each reaction takes place in a different container. These containers are called
half-cells. A wire connects the two half-cells through a device that measures current
called a voltmeter. In order to complete the electrical circuit, ions must be free to travel
from one half-cell to another. This is made possible by connecting the two half-cells
with a salt bridge. A salt bridge is a tube (usually U-shaped) that connects the two half-
cells of a redox reaction, and allows ions to migrate from one cell to another.
In each half-cell a conductor is needed to establish the electric current. These
conductors are called electrodes. In the oxidation half-cell, the electrode is called the
anode. The anode is considered to be a negative electrode because electrons are
generated there. In the reduction half-cell, the electrode is called the cathode. The
cathode is considered to be a positive electrode because electrons are attracted to it.
The current generated by an electrochemical cell is called electrochemical potential
(Eo). Electrochemical potential is the difference in ―potential‖ of the two half-reactions
that can generate current. The difference is a comparative measurement (compared to
a standard hydrogen half-cell) and is measured in volts.
In this experiment you will create an electrochemical cell, and observe a redox reaction.
You will also measure electrochemical potential of this redox reaction and compare it to
the accepted value for that redox reaction.

Objective
      Set up and test the voltage of an electrochemical cell.

Safety
      Always wear safety goggles and a lab apron or coat when working in the lab.
Wash hands before leaving the lab.

Materials
black coated wire with alligator clips on end and male plug on the other end
red coated wire with alligator clips on end and male plug on the other end
two beakers (250-mL)        safety glasses       lab apron or coat
glass U-tube                cotton               DC voltmeter
copper strip                zinc strip           pipet
sand paper
0.10 M Solutions of: KNO3, ZnSO4, and CuSO4.



                                             57
Date __________________                            Lab Partner ___________________
                                                               ___________________
Procedure
1. Using sand paper, clean the strips of copper and zinc.
2. Add about 50 mL of a 0.10 M CuSO4 solution to a 250-mL beaker. Place a clean
   copper strip in the beaker.
3. Add about 50 mL of a 0.10 M ZnSO4 solution to a second 250-mL beaker. Place a
   clean zinc strip in that beaker.
4. Attach the alligator clip with the black wire to the piece of zinc and attach the
   alligator clip with the red wire to the piece of copper. Then attach both male ends of
   the wire into the voltmeter. Turn the voltmeter to read 20 DC volts and then turn on
   the voltmeter. Record the voltage you read from the voltmeter.
5. Fill ―to the top‖ a U-tube with 0.10 M KNO3 and stopper the ends with cotton. Invert
   the tube and place the ends of the U-tube into the beakers. Now read the voltmeter
   again and record the voltage.

Observations - (minimum of 5)


Data -        Voltage Data for an Electrochemical Cell

1. Voltage of the electrochemical cell before adding the salt bridge
       _____________ volts

2. Voltage of the electrochemical cell after adding the salt bridge
       _____________ volts


Calculations
Write out the reduction half-reaction for this redox reaction.


Write out the oxidation half-reaction for this redox reaction.


Show the net ionic equation for this redox reaction



Calculate the theoretical voltage for this Cu/Cu2+ and Zn/Zn2+cell using the
―Standard Electrode Potentials‖ Table provided by your teacher.




                                             58
Date __________________                            Lab Partner ___________________
                                                               ___________________
Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.
1. Define these terms: oxidation, reduction, oxidizing agent, reducing agent, electrode,
   cell potential, anode, and cathode.

2. Which electrode was the anode and which electrode was the cathode in this
   experiment? Explain.



3. What is the function of the salt bridge? How did the voltage of the cell compare with
   and without the salt bridge? Explain.


4. How well did your experimental voltage compare with the theoretical voltage
   calculated from the table.


5. Explain why the electrons flow from the anode to the cathode.


6. An electrochemical cell in which the half cells are Ag/Ag+ and Cu/Cu2+ was made.


7. Draw a picture of what this cell would look like. Label the anode, cathode, salt
   bridge, and where the ions are.


8. Name the reducing and the oxidizing agents.


9. Write out the half reactions for the oxidation and reduction half cells.



10. Write out the net ionic equation for this redox reaction.


11. Calculate the net electrode potential for this redox reaction.




                                             59
Date __________________                            Lab Partner ___________________
                                                               ___________________

                   Lab #17 Synthesis of Esters
Introduction
Esters are an important class of organic compounds which are characterized by the
following generic formula.
                             O
                         R   C O R'        R, R' = any alkyl group

Low molecular weight esters have pleasant odors and are responsible for many
distinctive odors in fruits and flavorings. Esters can be readily prepared from a
carboxylic acid and an alcohol. Generally this reaction is catalyzed by strong acids.
                            R, R' = any alkyl grou p

         O                                            O
                                         H 3O+
     R    C OH + R' CH2             OH            R   C O     CH2    R'   + H 2O


Carb ox ylic Acid +      Alco ho l                        Ester           +   W ater

 Sulfuric acid is used as a catalyst for this reaction because it also serves as a
dehydrating agent, shifting the equilibrium toward products. The following structures will
assist in determining which acid and alcohol will produce the ester with the desired
flavor. The group to the left of and including the
                                             O

                                           C O

belonged to the original carboxylic acid, and the group to the right of the
                                           O

                                           C O

belonged to the original alcohol.

Some ester structures are shown on the next page.




                                             60
Date __________________                            Lab Partner ___________________
                                                               ___________________

         O                   CH3                               O
    H    C O CH2             CH CH3                    CH3     C O      CH2      CH2     CH3

               Raspberry                                                  Pear

                              O                                         O
   CH3       CH2       CH2    C O CH3            CH3     CH2   CH2      C O        CH2    CH3

                                                               Pineapple
               Apple

                   O               CH3                 O                         CH3
 CH3     CH2       C O CH2         CH CH3        CH3   C O     CH2      CH2      CH CH3
                 Rum
                                                                   Banana
             O                                             O
    CH3      C O        CH2                        CH3     C O      CH2      (CH 2)6   CH3

             Peach                                                  Orange

                  O                                                 O
                   C O       CH3                                    C O      CH3

           OH                                                  NH 2
          Wintergreen
                                                                        Grape
Objective
To synthesize an ester from the corresponding carboxylic acid and alcohol, and to
identify the characteristic odor of each ester you have made.

Safety
1. Wear protective goggles throughout the laboratory activity.
2. Be very careful when using the concentrated sulfuric acid. It can cause serious
   chemical burns.
3. You should always use a wafting technique to check the odor of an ester; you should
   never directly smell any chemical used or prepared in the laboratory.


Materials
acetic acid, methanol, pentanol, 18M sulfuric acid, and salicylic acid, scoopula, 10-mL
graduated cylinder, 125-mL Erlenmeyer flask, ring stand, iron ring, 400-mL beaker,



                                            61
Date __________________                           Lab Partner ___________________
                                                              ___________________
utility clamp, burner, one-hole rubber stopper with long glass tube, pipet, lab apron,
safety glasses.




Procedure

Group A- Synthesis of methyl salicylate ester
1.   Record the name of the acid and alcohol you will use. In a 125- mL Erlenmeyer
     flask insert about one flat tea spoon of salicylic acid. Add exactly 5 mL of
     methanol using a graduated cylinder. Then, go to the teacher for 8 drops of
     concentrated sulfuric acid. Be very careful with the sulfuric acid. It can cause
     severe burns if spilled on the skin.
2.   Place about 300 mL water in a 400-mL beaker. Then place the 125- mL
     Erlenmeyer flask into the water bath. Secure the flask with a clamp. Insert the
     one-hole rubber stopper (with the long 10-mm glass tubing in it), into the mouth
     of the Erlenmeyer flask.
3.   Heat the water bath until boiling. Then let simmer for 15 - 30 minutes.
4.   Record any changes which take place in the reaction mixture.
5.   After the mixture cools, remove the reaction assembly from the warm water.
6.   Remove the stopper with the long 10-mm glass tubing in it and waft the fumes
     toward your nose. Carefully note and record the odor of the ester you have
     synthesized. Caution-Occasionally, the vapor is too concentrated and it will
     ―overpower‖ your senses.
7.   The other group has prepared a different ester. Observe the odor of the ester the
     other group has prepared and record its smell.
8.   Thoroughly wash your hands before leaving the laboratory.

Group B- Synthesis of amyl acetate ester
1.   Record the name of the acid and alcohol you will use. In a 125- mL Erlenmeyer
     flask insert 5 mL of acetic acid using a graduated cylinder. Add exactly 5 mL of
     n-pentanol using a graduated cylinder. Then, go to the teacher for 8 drops of
     concentrated sulfuric acid. Be very careful with the sulfuric acid. It can cause
     severe burns if spilled on the skin.
2.   Place about 300 mL water in a 400-mL beaker. Then place the 125- mL
     Erlenmeyer flask into the water bath. Secure the flask with a clamp. Insert the
     one-hole rubber stopper (with the long 10-mm glass tubing in it), into the mouth
     of the Erlenmeyer flask.
3.   Heat the water bath until boiling. Then let simmer for 15 - 30 minutes.
4.   Record any changes which take place in the reaction mixture.
5.   After the mixture cools, remove the reaction assembly from the warm water.
6.   Remove the stopper with the long 10-mm glass tubing in it and waft the fumes
     toward your nose. Carefully note and record the odor of the ester you have
     synthesized. Caution-Occasionally, the vapor is too concentrated and it will
     ―overpower‖ your sense.



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                                                               ___________________
7.     The other group has prepared a different ester. Observe the odor of the ester the
       other group has prepared and record.
8.     Thoroughly wash your hands before leaving the laboratory.


Observations - Minimum of 5


Data - Esterification Data

              Name of Carboxylic     Name of         Name of Ester       Aroma
              Acid                   Alcohol         made
Group A
Group B
                                                        See teacher
                                                        for help on
                                                        naming esters

Calculations
In your report, write a balanced equation for the synthesis of both esters.




Questions

1.     What is the role of the sulfuric acid in these reactions?


2.     How does the -OH group of an alcohol differ from the the -OH group of a base?



3.     In what type of products are esters used?




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                                                                 ___________________


                       Lab #18 Twizzler Half-Life
Introduction

The time required for half of the atoms in any given quantity of a radioactive isotope to
decay is the half-life of that isotope. Each particular isotope has its own half-life. For
example, the half-life of 238U is 4.5 billion years. That is, in 4.5 billion years, half of the
238
   U on Earth will have decayed into other elements. In another 4.5 billion years, half of
the remaining 238U will have decayed. The half-life of 14C is 5730 years, thus it is useful
for dating archaeological material. Nuclear half-lives range from tiny fractions of a
second to very long periods of time.

Objectives: The learner will

       1.      Identify an exponential decay curve
       2.      Find half-life and lifetime of an exponential decay
       3.      Given data or a graph, generate the equation of the exponential decay
               curve
       4.      Create a general model for all observed exponential decay phenomena



Materials: 2 twizzlers per lab group, graph paper

Procedure:
1.   Have students prepare axes on the graph paper. The horizontal axis should be
     marked in half-lives (0-10) and the vertical axis labeled ―Length of Twizzler.‖
     (Note: the vertical axis should be long enough to accommodate the length of a
     single Twizzler.)
2.   Given 2 Twizzlers, the students place one of them on the vertical axis with one
     end touching the horizontal axis.
3.   The second Twizzler is cut in half. One of the halves is placed vertically on the
     graph over the 1-half-life mark.
4.   Cut the remaining piece in half. Put one of the halves vertically over the 2-half-life
     mark.
5.   Repeat until the Twizzlers are effectively uncuttable. (Is this a fundamental
     Twizzler Particle?)
6.   Mark the top of each Twizzler; connect the marks with a smooth curve. (Deduct
     at least 90% of credit for dot-to-dot graphing.)

Observations – Minimum of 5

Conclusions – Minimum of 5




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                                                              ___________________




Questions:

1.    How does the Twizzler length change as time (in half-lives) increases?



2.    How does the rate of decrease in Twizzler length change as time increases?



3.    What is the significance of a single half-life in terms of Twizzler length? 2 half-
      lives? 3 half-lives?



4.    Does the choice of starting point affect your answers to question 3?



5.    The curve you drew is an exponential decay curve. List the important
      characteristics of such a curve.



6.    If you had a 100-gram sample of a material the mass of which decayed in the
      manner of your Twizzlers, how much of it would you have after 1 half-life? 2 half
      lives? 3 half lives?




                                            65
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                                           ___________________




                          66
Date __________________                           Lab Partner ___________________
                                                              ___________________

                Lab #19 Establishing Equilibrium
Objective: How can you determine when a system is at equilibrium?

Materials:

2 25 mL graduated cylinders                            graph paper
2 drinking straws with different diameters             Paper and pencil
Food coloring.


Introduction:

Chemical equilibrium applies to reactions that can occur in both directions. In a reaction
such as:

                          CH4(g) + H2O(g) <--> CO(g) + 3H2(g)

The reaction can happen both ways. So after some of the products are created the
products begin to react to form the reactants. At the beginning of the reaction, the rate
that the reactants are changing into the products is higher than the rate that the
products are changing into the reactants. Therefore, the net change is a higher number
of products.
Even though the reactants are constantly forming products and vice-versa the amount
of reactants and products does become steady. When the net change of the products
and reactants is zero the reaction has reached equilibrium. The equilibrium is a dynamic
equilibrium. The definition for a dynamic equilibrium is when the amount of products and
reactants are constant. The reactants and products are not equal but constant. Also,
both reactions are still occurring.

In dynamic equilibrium the forward and backwards reactions continue at equal rates so
the overall effect does not change. On a molecular scale there is continuous change.
On the macroscopic scale nothing appears to be happening. The system needs to be
closed – isolated from the outside world.




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                                                              ___________________
Procedure:

1.    Label one graduated cylinder ―Reactants‖ and the other ―Products.‖ Label one
      straw ―Forward‖ and one straw ―Reverse.‖

2.    Add a few drops of food coloring to the reactant cylinder. Then fill the cylinder
      with water to the 25.0 mL mark.

3.    Place the forward straw into the reactant cylinder so that it touches the bottom.
      Place the reverse straw into the empty product cylinder.

4.    Put your finger over the top of the forward straw so that the liquid is trapped in
      the straw. Transfer the colored water in the straw to the product cylinder.

5.    Now put your finger over the reverse straw and transfer the colored liquid from
      the product cylinder to the reactant cylinder.

6.    Record the volume of each cylinder in the data table on the next page. Each time
      you transfer liquid from the reactant cylinder to the product cylinder and back
      from the product cylinder to the reactant cylinder constitutes one transfer.

7.    Repeat steps 4,5, and 6 until the volumes in the cylinders remain constant for at
      least 5 readings.

8.    Repeat the entire experiment, but this time begin with 25.0 mL of colored liquid in
      the product cylinder and nothing in the reactant cylinder. Use the same straws for
      the forward and reverse reactions.

Observations – Minimum of 5

      1.     Did you use a larger straw for the forward or reverse reaction?
             _______________________________________________

      2.     At equilibrium, was there a greater volume of products or reactants?
             ________________________________________

      3.     What was the effect on equilibrium volumes when you started with no
             reactants and 25.0 mL of products?(Which had the greater volume ?)
             ___________________________________________________
             ___________________________________________________
             ___________________________________________________

      4.     On a sheet graph paper or using Excel, graph your results for both
             experiments by plotting the transfer number on the x-axis and the volumes
             of both products and reactants on the y-axis.




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                                                      ___________________
Data

                          Trail 1    VOLUME (mL)
                                   Reactant Product
                          Transfer cylinder clinder
                             1
                             2
                             3
                             4
                             5
                             6
                             7
                             8
                             9
                             10
                             11
                             12
                             13
                             14
                             15



                          Trail 2    VOLUME (mL)
                                   Reactant Product
                          Transfer cylinder clinder
                             1
                             2
                             3
                             4
                             5
                             6
                             7
                             8
                             9
                             10
                             11
                             12
                             13
                             14
                             15




                                     69
Date __________________        Lab Partner ___________________
                                           ___________________




                          70
Date __________________                          Lab Partner ___________________
                                                             ___________________
Analysis and Conclusions
How do you know that a reaction has taken place?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

How do reactants differ from products?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

Was equilibrium established in both experiments? If so, what can you conclude?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

What variable determined whether there was greater volume of reactants or products at
equilibrium?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

Why could this system be called a dynamic equilibrium?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

What could be sources of error in this experiment that could give incorrect results?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

How would the outcome of the experiment be different if the drinking straws had been of
different sizes?
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________

Comments on lab.
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________




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                                                                 ___________________

                   Lab #20 Why is Water Weird?
Purpose: Water is a unique and wonderful substance. It is the basis of life processes
and the regulator of our environment. All life functions takes place in water and the
water on the Earth’s surface allows life as we know it to exist by regulating the
temperature on the planet.
Water does all of this great stuff, so we should know its chemistry.
We all experience some of the unique properties of water in this lab.

Background: (Information about water can be found anywhere – start in your textbook
…)

 In paragraph form, tell everything there is to know about the chemistry of water. Diagrams can
    be used. (Include – structural formula, electron config, bond angle, bp/mp, specific heat,
                      polarity, bond type, density, anything you can find..)

Define: adhesion, cohesion, surface tension, refraction, hydrogen bonds

Procedure:
1.   Take a clean penny. Using a pipette, see how many drops you can get on a
     penny. What does the penny look like from the side? Draw it. Why does this
     happen? Look at President Lincoln, what does the water do to him?
2.   Fill a plastic cup with water. Carefully add water to fill it over the top. Use a ruler
     to measure in mm how much over the top you can get it? Draw it. Explain why
     you can do this?
3.   Using the same water and cup from #2, float a paperclip or two or three on the
     water’s surface. Why does it float? What direction does the clip move after you
     place it in the middle? Why? Draw a diagram.
4.   Take a cup; fill 2/3 with water. Stick your finger or a pencil or pen in the water.
     How does it look from the side? Draw it. Why does it look this way?
5.   Put a straw in the water in the cup. Does the water level in the straw match the
     water level in the cup? Explain. Draw what you see.
6.   Pour water into a graduated cylinder. What does the top of the water look like?
     What is this called? Why does water do this? Draw and explain.
7.   Draw a picture of solid (crystallized) water at the molecular level. Use your
     drawing to explain why water is denser as a liquid than a solid.
8.   Find a picture of DNA. (USE GOOGLE) Show the area where replication occurs
     – where the two strands are connected. What holds the DNA together and is the
     related theme through this whole lab?



Conclusion Questions:
1.   Explain how the metric system is based on water … temperature, mass, density,
     volume. (USE GOOGLE)



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Date __________________                          Lab Partner ___________________
                                                             ___________________


2.    Why is it a really good thing that most of the hydrogen on the planet is bonded to
      oxygen in water?



3.    Give your two favorite reasons that water is vitally important to life on earth.
      (Mine are: 1. Mallorca is surrounded by it,and 2. you can go scuba diving in it)




4.    Comments on lab. (At least 2)
      1.

      2.




                                           73
Date __________________        Lab Partner ___________________
                                           ___________________
DRAWINGS                   OBSERVATIONS




                          74
Date __________________        Lab Partner ___________________
                                           ___________________
DRAWINGS                   OBSERVATIONS




                          75
Date __________________                             Lab Partner ___________________
                                                                ___________________

             Lab #21 - Finding Patterns in Measurements
                                        Line Graphs

Background Information

For our distant ancestors, recognizing whether they were about to step on grass, off a
cliff, on slippery mud, or into a pond probably was critical information. Perhaps as a
result, the human brain has evolved an ability to recognize visual patterns. Since we
have this ability, it provides us a valuable tool for analyzing other kinds of information.

Recall that Pythagoras proposed that mathematics and numbers revealed the essence
of the universe. Plato taught that mathematics could be used to explain the universe.
Their followers were increasingly successful in measuring and understanding aspects of
the physical world then learning to use that understanding to control the world to their
benefit.

As their society emerged from the Midevil period, some Europeans such as Thomas
Bradwardine (c. 1295-1349) tried to clarify the causes of change such as motion. He
and other fellows at Merton College searched for geometric methods of visualizing
mathematical patterns for speed and other properties that change. They began to draw
bars and histograms of lengths proportional to successive measurements. Nicole
Oresme of Paris used what today would be called graphs to quantify physical qualities
such as speed, displacement, temperature, whiteness, and heaviness, but also
nonphysical qualities such as love, charity and grace. For example, he used the
geometry of a graph to prove a uniformly accelerated object travels the same distance
as it would, had it travelled steadily at the average (Sav) of its initial (Si) and final (Sf)
speeds. (The area of the distance rectangle is equal to the area of the trapezoid
because moving the pink triangle doesn't change its area.)

That search for patterns in the phenomena of our world continues with efforts such as
Stephen Wolfram's A New Kind of Science in which he uses computers and a program
called Mathematica to seek patterns in the apparent chaos of vast collections of related
numbers.
The Task

While histograms, pie charts, and other bar graphs are helpful in visualizing patterns in
one aspect of nature, we are often more interested in how one aspect of nature relates
to another aspect. For example, a fireman needs to know how the height water can be
delivered from a fire hose depends on the water pressure. Such information can be
visually presented by a line graph. Traditionally such a graph is made manually with
pencil and graph paper. But computer programs using spreadsheets such as Excel can
also generate such graphs. Learning to manually draw good graphs with all the
necessary parts is a valuable precursor to manipulating Excel to make similar graphs.
Incidentally, making a pretty graph using a computer program such as Excel often
requires more time than to make a good graph manually!



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                                                              ___________________
At this point, many people believe they have previously learned how to make graphs,
skip further instruction, then construct FLAWED graphs! A word of advice: Don't
assume you already know everything; at least read the links below about making graphs
and search for construction details that make a pretty but useless graph into effective
communication.
How to Manually Construct a Line Graph

  1. Obtain graph paper.
  2. Identify variables
       1. Independent Variable (controlled by he experimenter): Values for this property
are assigned to the horizontal axis ( ↔ often called the x axis)
       2. Dependent Variable (This property should slavishly change depending on the
value chosen for the independent variable.) Values for this property are assigned to the
vertical axis ( ↑ ↓ often called the y axis)
  3. Determine the range of the two variables. For each variable subtract the lowest
data value from the highest data value. If you wish the origin of the graph to start at
zero, use the highest data value for the range.
  4. Determine the scale of the graph. Count the number of squares horizontally and
vertically on the graph paper. Divide the range of the independent variable by the
horizontal number of squares, then round off to the next larger convenient number. (For
example, if dividing gives 34.3, it might be convenient to allow each horizontal square to
have a value of 50.) Determine the scale for the dependent variable by dividing its
range by the number of vertical squares, then rounding up. This procedure should
spread the graph to use MOST of the available space.
  5. Number each axis. Starting with a number just smaller than the smallest
independent data value (or zero if that is desired for your origin) label each LINE along
the bottom of the graph from left to right by successive multiples of the chosen scale.
(For our example, 0, 50, 100, 150...) It is critical that these numbers represent lines and
not boxes. That will allow the space between the lines to represent values in between.
Number upwards the lines along the left side of the vertical axis using the scale chosen
for the dependent variable, starting just below the smallest data (or zero if so chosen for
the origin). If lines are so close together that the numbers will be crowded, omit writing
some of the numbers but continue to space as if all lines were numbered.line graph
  6. Label each axis.
       1. Briefly describe the properties represented by the independent and dependent
variables.
       2. In parentheses abbreviate the measuring units used to measure the data along
each axis.
  7. Plot the data points. For each pair of related data, use the value for the
independent variable to plot horizontal location and the value for the dependent variable
to plot vertical location using the number lines along the axes. Make a tiny dot at the
intersection of the horizontal and vertical locations to accurately show the two values.
To make the approximate location of the dot apparent later, place a marker such as a
small circle around the dot. If other sets of data are also plotted on the same graph for
comparison, use a different marker (square, triangle etc.) for each set of data and use a
legend box to show what each marker represents.



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                                                              ___________________
  8. Draw the line. Most graphs of experimental data are NOT drawn by connect the
dots. Recall that all measurements contain experimental error, so as a result, most dots
are not exactly where they should ideally be located. To compensate for this
experimentally error, try to draw a straight line that compromises, coming closest to
nearly all the dots. Sometimes a straight line does not fit well. In that case, consider
      1. drawing several intersecting straight line segments, each of which fits much
data.
      2. drawing a smooth curve that comes closest to nearly all the data.
  9. Add a title, completion date, and author's signature. The title should be selected to
clearly but briefly tell what the graph is about. Avoid cute teasers that attract attention
but leave out critical information.

Construct a line graph for data such as provided in Table #1 below.
Table 1: Distance from Home and Driving Time
Time                 Distance        Time             Distance
(minutes)            (miles) (minutes)         (miles)
       0                  0              12                    6.7
       1                 0.2             13                    7.6
       2                 0.6             14                    8.5
       3                 1.1             15                    9.3
       4                 1.6             16                   10.3
       5                 2.2             17                   11.3
       6                 2.6             18                   12.1
       7                 3.0             19                   13.0
       8                 3.6             20                   13.9
       9                 4.2             21                   14.7
      10                 4.9             22                   15.5
      11                 5.8             23                   16.4




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                                           ___________________




                          79
Date __________________                          Lab Partner ___________________
                                                             ___________________


How to Construct a Graph Using a Computer

Nearly all computers have applications for constructing spreadsheets. Such
applications usually are capable of constructing graphs from series of data entered into
the spreadsheet. If you are working on a computer that uses an operating system from
Microsoft, you probably have a spreadsheet program called Excel. Jim Askew at Howe
High School in Howe, Oklahoma has created instructions for making graphs using
Excel. To learn more about making effective line graphs using Excel, use this link to
Jim Askew's instructions.

Using a computer spreadsheet such as Excel, construct a line graph for the data in
Table #1 above.




                                           80
Date __________________                           Lab Partner ___________________
                                                              ___________________


                               Lab # 22 – Table M
Objective : ______________________________________________________
                         (Safety goggles must be worn)


Name ________________________
Lab Partner(s) _________________
             _________________

Equipment: safety goggles, apron, 24 well reaction plate, micro pipettes, white piece of
paper (for background), Table ―M‖ on the Reference Table for Physical
Setting/Chemistry.

Discussion: An indicator is a substance that changes color when it loses or gains a
proton (H+). For example, phenolphthalein is a common indicator that is colorless when
it is protonated (has a proton = H+). When a base is gradually added to an acid
containing phenolphthalein., the solution is colorless. Once the acid has been
neutralized ( [H+] = [OH-]) by the addition of the base, the base reacts with the hydrogen
ion of the indicator. As the phenolphthalein loses the hydrogen ion it turns pink. The
color change is why phenolphthalein is an indicator. Other indicators react in a similar
way to phenolphthalein, but each has a unique color change that occurs at a specific pH
range. Table M of the Reference Table for Physical Setting/Chemistry lists several
common indicators, the color changes they undergo, and the pH range over which the
color changes occur.

Procedure:
   1. Put on goggles.
   2. Clean and rinse 24 well reaction plate with water.
   3. Place the reaction of a piece of white paper, with 12 wells across and 8 wells
      down.




   4. In the first row of the top reaction plate, add two drops of indicator to each of the
      seven wells. Follow the same procedure for each row, using a new indicator for
      each row that corresponds to the data table.
   5. Add enough distilled water to each well to cover the bottom of the well.


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Date __________________                            Lab Partner ___________________
                                                               ___________________
   6. Using the reaction plate from top to bottom, in the first column dispense five
      drops of Milk to each of the four wells. Follow the same procedure for each of the
      other six solutions.
   7. Observe and record the color in each well on the data table.
   8. Using the color ranges and the pH values associated with the colors estimate the
      pH of the solutions.
   9. Using a small strip of pH paper, test each beaker by dipping the strip into the
      solution and record the pH.


                     Milk      Apple     Diet 7-      Pine      NaOH      Soap       Unknown
                               Juice       up         Sol
Phenolphthalein

 Bromthymol
    Blue
Methyl Orange

 Thymol Blue

    Litmus

 Bromocresol
     Green
 Univ Indicator

   pH Paper




Questions:

   1. What can you conclude about each solution, is it acidic, basic, or neutral?




   2. Do the color changes you observed confirm the pH paper results? Does this
      match Table M?




                                           82
Date __________________                          Lab Partner ___________________
                                                             ___________________


   3. Using your data, identify the color changes for universal indicator? What color is
      basic, neutral and acidic?




                                           83
Date __________________                               Lab Partner ___________________
                                                                  ___________________



                         Lab # 23 – Types of Reactions
Matter can undergo both physical and chemical changes. Chemical changes
result in the formation of new materials. Observable signs of chemical change
include the release of a gas, a change in color, the formation of a precipitate, and
a change in heat and light. One classification system involves five general types of
reactions: combustion, synthesis, decomposition, single replacement, and double
replacement.


PRE-LAB PREDICTIONS

For each of the following reactions, specify the type(s) of reaction and predict the products
(including physical states) that will form. Record these in your lab notebook and BALANCE
each equation.
A. CaCl2(aq) + Na2CO3(aq) 
B. Mg(s) + HCl(aq) 
C. CuCO3(s)  HINT: Carbonates break down to form CO2 and a metal oxide.
D. Mg(s) + O2(g) 

E. C4H10 (g) + O2(g) 

F. KClO3 (demo)

PROCEDURE
REACTION A
  1. Add about 1 mL (1 mL  width of your pinky finger) of calcium chloride solution, CaCl2, to
     a clean test tube. Next, add about 1 mL of sodium carbonate solution, Na2CO3, to the
     same test tube. Record your observations.
REACTION B
  2. Stand a clean test tube in the test tube rack and add 1-2 mL of 6M hydrochloric acid.
     CAUTION: 6M HCl can cause burns! Carefully drop a few pieces of magnesium into test
     tube.
   3. While the reaction is still occurring, use a test tube holder to hold an empty test tube
      over the top of the first test tube for ~30 seconds. Meanwhile, another group member
      should light the burner and use it to ignite a wood splint. At the end of 30 seconds, hold
      the inverted test tube (keeping it inverted) while the other group member places the
      burning wood splint into the mouth of the inverted test tube to test for the presence of
      hydrogen gas (a loud ―whoop!‖ should be heard).
   4. Record your observations, including the results of the hydrogen test. KEEP THE
      BURNER LIT FOR REACTION C.



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                                                                 ___________________
REACTION C
   5. Place 2 heaping microspatulas of copper(II) carbonate, CuCO3, in a clean test tube.
      Note the appearance of the sample in your observations.
   6. Using a test tube holder, heat the CuCO3 strongly in the burner flame for 2-3 minutes.
      AIM THE MOUTH OF THE TEST TUBE AWAY FROM YOURSELF AND OTHERS.
      After heating, another group member should ignite a wood splint and quickly place the
      burning splint into the test tube to test for the presence of CO2 gas (the flame should go
      out).
   7. Record your observations, including the results of the CO2 test. KEEP THE BURNER
      LIT FOR REACTION D.
REACTION D
  8. Place a watch glass next to the burner. Examine a piece of magnesium ribbon and
     record your observations. Using crucible tongs, hold the ribbon in the burner flame until
     the magnesium starts to burn. DO NOT LOOK DIRECTLY AT THE FLAME. Hold the
     burning magnesium directly over the watch glass. When the ribbon stops burning, place
     the remains on the watch glass. Examine this product thoroughly and record your
     observations. Turn off the burner.

   9. Clean your test tubes and lab station. Wash any liquids down the drain and wrap solids
      in a paper towel before discarding them in the trashcan.

REACTION E            Butane Demo

REACTION F            KClO3 Demo




CONCLUSIONS

Record your observations in your lab notebook using the following data table as a guide.


                                               85
Date __________________                              Lab Partner ___________________
                                                                 ___________________
 REACTION                  BEFORE REACTION                              AFTER REACTION

 A
 B
 C
 D
 E
 F



Write a conclusion for this lab based on your results. For EACH reaction – A, B, C, and D –
address the following:
        Describe whether or not your observations confirmed the products you predicted before
         the lab. (e.g. If a solid was predicted as a product, did you observe a solid being
         formed? Did you confirm the formation of specific substances?)




        Based on each chemical equation, explain how each reaction represents its designated
         reaction type(s).




                                               86
Date __________________                           Lab Partner ___________________
                                                              ___________________

                            Lab #24 – Obscertainer
   How many times have you asked the question, ―What is that made of?‖ That same
question has been asked millions of times in many different languages by many
scientists. In order to make sense of things that cannot be seen, scientists use indirect
observations and create models to explain scientific phenomena.
   This activity requires you to use the techniques of good scientists—making careful
observations and checking the accuracy of your hypothesis. In order to solve this
problem it takes concentration, alertness to detail, ingenuity, and patience!

PROBLEM: What is the configuration (design) of the inside of a closed container
         (Obscertainer)?

THEORY: The closed obscertainers have walls inside and a steel ball that can freely
            move. You will not be able to see or touch the inside of the
            obscertainer, therefore you will be expected to determine the design of
            the inside by indirect means.




PROCEDURE: Move the steel ball around by carefully shaking and tilting the
obscertainer. By the sound and path of the steel ball determine the shape and location
of the          wall(s). You must examine three different obscertainers.




                                            87
Date __________________                            Lab Partner ___________________
                                                               ___________________
        Obscertainer #        Observation          Hypothesis       Actual Design
                            (what you hear)




Compare your observations and hypothesis for the obscertainers you examined with at
least one other student who examined the same obscertainers. Did your designs
match? Use this experience to explain why it is useful to carry out multiple trials of an
experiment.




                                              88
Date __________________                           Lab Partner ___________________
                                                              ___________________

                            Lab 25: Types of Solutions
Introduction
       In this experiment you will investigate unsaturated, saturated, and supersaturated
solutions. The solute used is hydrated sodium acetate, NaC2H3O2 3H2O, which
shortens to NaAc.

Objective
      To prepare a supersaturated solution, and observe its characteristics.

Procedure
     1. Add ¼ tsp NaAc crystals to 5ml of water in a test tube. Swirl until dissolved.
     2. Add ¼ tsp more, swirl to dissolve.
     3. Add approximately ¾ more tsp and swirl. If all dissolves after at least two
         minutes of swirling add about 1/4tsp more.
     4. When some undissolved solute remains on the bottom of the test tube and
         cannot be dissolved, add approx. 4 level teaspoons of additional NaAc.
     5. NOTE ANY TEMPERATURE CHANGE THAT HAS OCCURRED.
     6. Heat the test tube as directed in class until all the crystals dissolve. While
         heating, gently shake the tube to mix the contents thoroughly so that the rate
         of dissolving is increased.
     7. While heating, the other lab partner should fill a beaker with cold water.
         Stand the heated test tube in the beaker for about 5 minutes. Do not disturb it
         during this time. If crystals appear in the solution, heat the test tube until they
         dissolve and again stand the test tube in cold water.
     8. After the solution has cooled to about room temp, hole the test tube in you r
         hand so you can NOTE THE TEMP OF THE SOLUTION. Do not shake the
         tube …no agitation at all.
     9. Add one crystal of NaAc to the test tube while you are holding it. Do not add
         more than one. Observe the crystallization which should slowly occur. NOTE
         ANY TEMP. CHANGE DURING CRYSTALLIZATION.
     10. If crystallization occurs instantly, reheat and do over.
     11. Dispose of the chemical as directed by teacher.




Conclusions
     Write five conclusions.

Questions
      1. What experiment would you do at the end of step #2 to determine if the
         solution produced is saturated?


                                            89
Date __________________                           Lab Partner ___________________
                                                              ___________________




      2. Because of what happened in step #8, do you think that the resulting
         solutions can be described as saturated? Explain




      3. Explain the observed temperature changes as NaAc dissolved and as it is
         crystallized out of solution.




List four different ways of making a saturated solution
    1.
    2.
    3.
    4.




                                            90
Date __________________                          Lab Partner ___________________
                                                             ___________________

                         Lab #26 -It's a Gas Exploration
Problem:

How do the variables of temperature, pressure and volume of a gas affect one another?

Materials:
lab stations

Hazard Warning:

Liquid nitrogen is cold! Avoid skin substance, as it will cause severe burns. Whenever
working with liquid nitrogen, contact with this Wear safety goggles.

Procedure:

A gas is a state of matter that completely fills and responds most noticeably to changes
in pressure, and volume. Its container temperature,
A number of "lab" stations have been positioned around the room. Under a time limit
determined by the instructor, you will cycle through each station manipulating the
apparatus with the materials provided.

At each station your job is to:
   1) Observe and record the changes that the gas is experiencing
   2) Determine and record which of the three variables (temperature, pressure and
      volume) is held constant
   3) Determine the relationship between the variables not held constant.

Prepare a data table with headings as shown below. Record your findings as you
progress from station to station. Write the word "variable" or "constant" under the
temperature, pressure, and volume columns of your table to reflect what you observe at
each station. You may also wish to takes notes as to the relationship, which you
discover between each of the variables at each station.




                                           91
Date __________________                Lab Partner ___________________
                                                   ___________________


Data Table
Station        Observation   Temp    Press   Volume          Relationship
                                                              between
                                                              variables

1




2




3




4




5



6




7




8




Summing Up:



                                92
Date __________________                          Lab Partner ___________________
                                                             ___________________
Compare your results with other groups.
Discuss the temperature, summarizing variables in results with other groups in the class
about relationships among the variables of pressure and volume.

Write a sentence about the relationship between each pair of variables in your table.
   1.

   2.

   3.

   4.

   5.

   6.

   7.

   8.


In those situations in which volume was changing, did this mean the number of particles
was also changing? Explain your answer.




Were there any stations at which three variables were changing simultaneously?
Explain why this can present a problem when interpreting data.




                                           93
Date __________________                         Lab Partner ___________________
                                                            ___________________



                      Lab # 27 – Molar Volume of a Gas

Introduction
       Avogadro’s hypothesis states that equal volumes of all gases contain equal
number of molecules under the same conditions of temperature and pressure.
Therefore, one mole of any gas at standard temperature and pressure (STP) will take
up a standard molar volume. This volume is 22.4 liters per mole. In this experiment,
you will try to prove that the standard molar volume is 22.4 liters per mole.

      The basis for this experiment is to react a known mass of magnesium with
excess hydrochloric acid to produce hydrogen gas, as shown in the reaction below:

                  Mg (s)   + 2 HCl (aq)  MgCl2 (aq) + H2 (g)

       Based on the mass of magnesium used, you will use stoichiometry to determine
the moles of hydrogen produced from this experiment. You will also collect the
hydrogen gas in a gas-measuring tube. This tube is graduated so that you can
measure the exact volume of gas collected. Using the combined gas law (a
combination of Boyle’s, Charles’ and Gay-Lussac’s laws) you will calculate the volume
of hydrogen gas that would have been collected at STP. Finally, using the volume
calculated from the combined gas law and the moles of hydrogen gas calculated from
the above equation, you will determine the standard molar volume of hydrogen gas.
This experiment should help your understanding of the mole concept and the concept of
molar volume of a gas.


Objective
       Determine the volume of 1 mole of hydrogen gas at STP using experimental
data, known mathematical relationships, and a balanced chemical equation.
      Safety
       Be careful when working with the 3 M HCl. Always wear safety goggles
and a lab apron or coat when working in the lab.

Materials
      gas-measuring tube safety glasses                400-mL beaker
      one-hole stopper with copper wire swirl utility clamp       100-mL beaker 10-
      mL graduated cylinder              lab apron or coat metric ruler
      1-liter graduated cylinder              thermometer         ring stand


Procedure - (Note: one partner should prepare step #6 before hand).

    1. Obtain a piece of magnesium ribbon from your teacher. Clean the ribbon with


                                          94
Date __________________                         Lab Partner ___________________
                                                            ___________________
       some sand paper till it is shiny. Measure the length of the ribbon accurately to
       within the first decimal place. Record this length as (a) in your data table. Also
       record the mass of 1 meter of magnesium ribbon as (b) in your data table. This
       mass will by provided by your teacher.
   2. Obtain a one-hole stopper with copper wire swirl. Wrap the magnesium ribbon
       around the copper wire so that it will not easily come off. (Your teacher will
       demonstrate).
   3. Obtain about 10 mL of 3 M HCl from your teacher. (You will need to get a 10-mL
       graduated cylinder for this). CAUTION: Handle acid with care. Carefully pour
       the acid into the gas-measuring tube.
   4. Fill a 100-mL beaker with water. Tilt the gas-measuring tube slightly. While
       keeping the gas-measuring tube at an angle, pour the water into the tube. Try to
       avoid mixing the acid and the water as much as possible. Fill the tube all the
       way to the top with water. (Your teacher will demonstrate).
   5. Insert the one-hole stopper with the magnesium into the tube. (Note: the water
       will overflow).
   6. Add about 350 mL of water to a 400-mL beaker. Set up a ring stand with a utility
       clamp about 10 cm above the beaker of water.
   7. Place your finger over the hole in the rubber stopper and invert the gas
       measuring tube. Without removing your finger from the hole, lower the
       stoppered end of the tube into the 400-mL beaker of water until the stopper is a
       few centimeters under the water level. Clamp the tube securely to the ring stand
       using the utility clamp. Record your observations.
   8. Let the apparatus stand until the reaction appears to stop. Gently tap the sides
       of the tube to dislodge any gas bubbles that appear in the tube. Carefully,
       without letting air into the tube, place your finger over the hole of the rubber
       stopper again. Loosen the clamp and lift up the tube while keeping your finger
       over the hole the entire time. Transfer the tube into a 1-liter graduated cylinder
       filled with water. Once the stopper is under the level of the water, you may
       remove your finger from the hole.
   9. Move the gas tube up and down inside the graduated cylinder to equalize the
       pressure. Lift the gas-measuring tube so that the water level inside the tube is
       at an equal level with the water inside the graduated cylinder. (Your teacher will
       demonstrate). Read the volume of the gases inside the tube and record the
       volume as (c) in your data table.
   10. Record the room temperature, (d), and the barometric pressure, (e). Both will be
       provided by your teacher. Find the water vapor pressure using your reference
       table and record it as (g) in your data table.


Observations - (minimum of 5)




                                          95
Date __________________                          Lab Partner ___________________
                                                             ___________________

Data -      Molar Volume Data
      a. length of Mg ribbon                           _________cm

      b. mass of 1 meter of Mg ribbon                  __________g (Teacher will tell
         you)

      c. volume of gas in gas tube                     _________mL

      d. temperature of room                           _________oC

      e. barometric pressure                           __________ mm Hg

      f. water vapor pressure at room temperature __________ mm Hg

Calculations
      1. Find the mass of the Mg ribbon
           length of your Mg (in meters)         =    x (grams of Mg ribbon)_____
            1.00 meter of Mg ribbon             mass of 1 meter of Mg ribbon

      2. Calculate the number of moles of Mg reacted (this is equal to the moles of
         hydrogen gas produced because the stoichiometry of the reaction is 1:1)

            x moles of Mg = mass of Mg ribbon x 1 mole of Mg_
                                                       24.3 grams of Mg

            ( = moles of H2 gas)

      3. Find the pressure exerted by the H2 gas (PH2) in the tube using Dalton’s Law
         of Partial Pressures:

            PH2 + PH2O = Pbarometric

      4. Convert room temperature from Celsius to Kelvin:
         K = oC + 273

      5. Find the volume of H2 gas at STP: (V2)

           P1V1     =     P2V2
            T1              T2
         P1 = PH2 V1 = volume gas in tube       T1 = room temp. in Kelvin

         P2= 760 mm Hg V2 = volume of H2 at STP T2 = 273 K

      6. Find the volume of 1 mole of H2 gas at STP:


                                           96
Date __________________                            Lab Partner ___________________
                                                               ___________________

             Volume of H2 (V2) =         x    mL
             Moles of H2                 1 mole


      7. Molar volume conversion to liters per mole:
             x mL     x     1 liter       = ________________liters/mole
             1 mole       1000 mL

Conclusions - (minimum of 5)
Questions
NOTE: when handing in the formal lab, you must write down the questions as well as
the answers. They must ALWAYS be in complete sentences.

      1. The accepted value for the molar volume of a gas is 22.4 liters/mole.
         How does value compare with your experimental value? Calculate
         your per cent error.



      2. What are some possible sources of error for this lab?




      3. How many liters would the following number of moles of any gas occupy at
         STP?

         a. 0.25 moles      b. 0.5 moles c. 1 mole      d. 2 moles   e. 2.5 moles




      4. Find the volume of the following masses of gas at STP.

         a. 80 g of O2      b. 10 g of H2 c. 14 g of N2 d. 66 g of CO2


                                             97
Date __________________                        Lab Partner ___________________
                                                           ___________________




      5. Find the moles for each of the following volumes of gas at STP.

         a. 22.4 liters of O2   b. 11.2 liters of H2 c. 44.8 liters of N2




      6. For each of the answers in question #5, calculate the masses for each
         gas.




      7. What was the other product in the experiment? What happened to it?



                          Lab #28 RATE OF REACTION

A clock can be constructed from molecules that react at a rate that allows an
interval of time to elapse between the mixing of the chemicals and the
completion of the reaction. Such "clock reactions" are important regulators of
biological cycles in nature. In this lab, you will make a chemical clock using
chemicals found in the supermarket.

Preparation of a Vitamin C stock solution

1. Place one (1) - 500 mg vitamin C tablets in a 150 ml beaker and pour 30 ml of
distilled water into the beaker with the tablets and let stand for 5 minutes.
2. Remove the tablets from the beaker (do not discard the water) and crush
them using a mortar and pestle.


                                         98
Date __________________                          Lab Partner ___________________
                                                             ___________________
3. Replace the crushed tablets in the beaker with the original 60 ml of water and
stir until completely dissolved.

Experiment 1

1. In separate 250 ml beakers, prepare Solution A and Solution B according to
the attached table.
2. Stir the mixtures with separate stirring rods.
3. Pour solution A into the beaker containing solution B and stir.
4. IMPORTANT: Start timing as soon as the two solutions are mixed.
5. IMPORTANT: Stop timing as soon as a dark color appears.
6. Record the reaction time in your data table.
7. Discard both solutions, wash and dry the beakers and stirring rods.

Experiment 2 & 3

1. Repeat the procedure for experiment 1 using the amounts listed in the
attached table for experiment 2 and repeat for experiment 3.




                          Solutions A & B Preparation Table

               SOLUTION A                                  SOLTUION B

EXP      Stock         DI      2% I2(aq)        EXP    H2O2(aq)    DI H2O     Starch
       Vitamin C      H2O
           mL                     mL                      mL         mL            mL
                      mL
 1         2.5        30          2.5            1        7.5         30           1
 2         2.5        15          2.5            2        7.5         15           1
 3         2.5        45          2.5            3        7.5         45           1

Data Table


                                           99
Date __________________                       Lab Partner ___________________
                                                          ___________________

              EXPERIMENT                            REACTION TIME (sec)




1. What are the three factors needed for a chemical reaction to take place.
     HINT: KMT

      a.

      b.

      c.

2. List 5 factors that affect the Rate of a Chemical Reaction.
       a.

      b.

      c.

      d.

      e.

3. Which factor listed in answer 1 was altered in this lab?
   _________________

Explain in terms of molecular collisions, why increasing reactant concentration
increases the chemical reaction rate.




                                        100

								
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