Name:
Lab Partner:
Period:
Date:
Lab One – Unit 1.3
Physical and Chemical Changes
Introduction:
A good understanding of material things requires and understanding of the
physical and chemical characteristics of matter. Such characteristics are familiar to you,
and physical and chemical changes are part of your everyday experience. However, you
may not yet have a clear idea of the difference between a physical change and a chemical
change. You may not yet know exactly how a chemical change is distinguished from a
physical change. The purpose of this experiment is to clarify these important
distinctions.
The physical properties of a substance are those properties that can be observed
and measured without changing the composition of the substance. Because they depend
on there being no change in composition, physical properties can be sued to describe and
identify substances.
The chemical properties of a substance are those properties that can only be
observed when the substance is undergoing a change in composition.
In a physical change, only temperature, size, or physical state of a sample of
matter is altered. In chemical changes, new substances, of different chemical
composition are produced. Readily observable phenomena include the evolution of gas,
the production of a color change, the formation of a solid, and the evolution of heat/light.
A process in which a chemical change takes place is called a chemical reaction.
Purpose:
Students will observe properties of several substances and decide whether changes in
matter are physical or chemical.
Equipment:
50 mL graduated cylinder
Bunsen burner
Evaporating dish
100 mL beaker
Crucible tongs
Test tubes
Cork stopper
Scissors
Spatula
Hot plate
Materials:
Magnesium ribbon (4 cm)
Copper(II) sulfate crystals
Lead(II) nitrate solution (1 M)
Potassium iodide solution (1 M)
Penny coin
Sheet of paper
Baking soda
Vinegar (5%)
Salt
Aluminum foil
Copper(II) chloride solution (1 M)
Procedure:
Investigation A:
1. Examine a piece of aluminum (Al) foil and identify at least three physical
properties.
2. Measure 20 mL of copper(II) chloride solution (CuCl2) in a small beaker.
Identify some physical properties of the solution.
3. Roll the Al foil into a small loose ball and place it in the CuCl2 solution.
Describe the results.
Investigation B:
1. Obtain a scoop of salt (NaCl) and identify some physical properties of sodium
chloride.
2. Measure 20 mL of distilled water (H2O) in a small beaker and identify some
physical properties of water.
3. Place a small portion of the salt in the water. Describe the results.
4. Transfer about one-half of the salt solution you prepared to an evaporating
dish and place the dish on a hot plate. Allow the water to evaporate
completely. Describe the results.
Investigation C:
1. Examine a post 1982 penny. List some physical properties of the penny.
2. Light a Bunsen burner and adjust the flame so that no yellow appears and that
you observe a small cone inside the flame. (adjust the air intake)
3. Using tongs, hold the penny in the outer portion of the flame until you see a
change occur. Describe the results.
Investigation D:
1. Examine a small piece of magnesium (Mg) ribbon and identify at least three
physical properties.
2. Using the crucible tongs, hold the piece of Mg ribbon in the outer portion of
the Bunsen burner flame (CAUTION!). Describe the results.
Investigation E:
1. Select several small crystals of copper(II) sulfate (CuSO4) and identify some
physical properties.
2. Using a graduated cylinder, measure 10 mL of distilled water (H2O) and place
it in a test tube. Identify some physical properties.
3. Drop the CuSO4 crystals into the water. Stopper the test tube and shake the
contents to promote interaction of particles. Describe the results.
Investigation F:
1. Using a graduated cylinder, measure out 5 mL of lead(II) nitrate solution
(Pb(NO3)2) and place it in a test tube. Describe some physical properties.
2. Using a graduated cylinder, measure out 5 mL potassium iodide solution (KI)
and place it in a test tube. Describe some physical properties.
3. Combine the contents of both test tubes. Describe the results.
Investigation G:
1. Obtain a sheet of typing paper. Examine it and identify some physical
properties
2. Using a pair of scissors cut the paper in such a way that you end up with a
hole in the paper large enough to slip your entire body through (2 bodies?).
Describe the results.
Investigation H:
1. Measure out a 1/2 scoop of baking soda (NaHCO3) on a piece of weighing
paper. Identify some physical properties. Place the baking soda in a small
beaker.
2. Using a graduated cylinder, measure 10 mL of vinegar (HC2H3O2) and
identify some physical properties.
3. Transfer the vinegar to the beaker containing the baking soda and allow them
to mix. Describe the results.
Data Analysis:
1. For each change you observe, indicate whether the change was physical or
chemical in nature. Give reasons for you answer.
Part A: Mixing Al and CuCl2 solution (no heating)
Part B: Dissolving NaCl in H2O (evaporation)
Part C: Heating a penny
Part D: Burning Mg
Part E: Dissolving CuSO4 crystals in water
Part F: Combing Pb(NO3)2 and KI solutions
Part G: Cutting paper
Part H: Combining baking soda and vinegar
Conclusions:
1. State in your own words the difference between physical and chemical properties.
Give an example of each that has not been mentioned in this experiment.
2. State in your own words the difference between a chemical change and a physical
change.
Lab One – Unit 1.3
Physical and Chemical Changes
Description
Students perform several activities involving physical and chemical change. They will
make observations of physical and chemical properties and relate them to the type of
change. Students develop a working definition for properties and change.
Time Frame: 100 minutes (2 class periods)
Materials: See student handout. Physical and Chemical Changes.
Procedures: See student handout. Physical and Chemical Changes.
Teacher Talk: Answers to data analysis: (A) chemical change, new substances produced;
(B) physical change, dissolving and evaporation take place; (C) physical change,
separation of an alloy; (D) chemical change, light, heat and new substance produced; (E)
physical change, dissolving; (F) chemical change, precipitate forms; (G) physical change,
changing shape; (H) chemical change, gas produced.
Answers to conclusion: students’ answers will vary.
Name:
Lab Partner:
Period:
Date:
Lab Two – Unit 2.1
Measurements and Density
Introduction:
Density, a physical property of matter, is the relationship between mass and
volume of matter. Mass is a measurement of the amount of matter in a sample, while
volume is a measurement of the space occupied by a sample of matter.
Measurements of mass are made on balances and different types of balances are
used to meet different measurement requirements. A triple-beam balance is used when
only approximate mass measurements are needed. An electronic balance is used when
greater accuracy is required. For maximum accuracy, an analytical balance is used.
Volume measurements are made in different ways depending upon the physical
state of the sample being measured. The volume of a liquid is commonly measured in a
graduated cylinder. The volume of a solid may be calculated from its dimensions, if the
solid is regular and free of air space. If, on the other hand, the solid is irregular of
contains air space, its volume must be determined in another way, such as by water
displacement. The solid must be completely submerged in the water for this method to
yield accurate result, and all the air bubbles adhering to the submerged solid must be
dislodged. This method is only useful for solids that are insoluble in water.
Purpose:
Students will obtain measurements and calculate densities for objects using mass and
volume.
Equipment:
50 mL graduated cylinder
Triple-beam balance
Ruler
Vernier caliper
Materials:
Metallic cylinders, Al, Fe, Cu
Lead fishing weight
Cork stopper
Procedure:
1. Obtain samples of different substances. Be sure that the samples are clean and
dry, and that you can distinguish between them. Get the mass of each sample on a
balance to the nearest 0.01 gram. Record the masses on the data table.
2. Find the volume of each sample in one of the following ways:
a. Water displacement – Fill a 50 mL graduated cylinder about ½ full with
water. Record the initial volume of water in the cylinder. Tilt the cylinder
and slide one the samples into the water, so that it does not break the
cylinder. Record the final volume of water containing the submerged
sample. Calculate the volume by subtracting the initial volume of water
from the final volume of water. Record the volume on the data table.
b. Direct measurement – Using a ruler or Vernier caliper, obtain the
dimensions of the sample. Using geometric formulas, calculate the
volume of the sample. Record the volume on the data table.
3. Calculate the density for each of the samples. Be sure to include the units in your
calculations. Record the densities on the data table.
Data Analysis:
1.
2.
3.
4.
Mass (g)
Volume of water alone
(mL)
Volume of water +
sample (mL)
Volume of sample
(mL)
Density of substance
Conclusions:
1. What does this experiment demonstrate about the density of a substance?
What does it demonstrate about the densities of different substances?
2. Compare your results with other groups in the class. Do you think that
density can be used to identify a substance? Explain.
Lab Two – Unit 2.1
Measurements and Density
Description
Students perform a series of measurements for mass and volume. The students will use a
triple-beam balance for determining mass. They will use direct measurement or water
displacement to find volume. They will use the measurement to calculate density.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Measurements and Density.
Procedures: See student handout. Measurements and Density.
Teacher Talk: Stress to the students the precision and accuracy of measurements by
practicing the use of significant figures and emphasize the precision of the instruments
used to take measurements. Check for accuracy by comparing student answers to
accepted values (might need to run experiments). Check for correct units.
Answers to conclusion: (1) density is a ratio of mass to volume; densities of different
substances are different; (2) yes, all substances of the same material had identical
densities.
Extensions: Have students devise a way to find the density of liquids or of substances
that are soluble in water.
Name:
Lab Partner:
Period:
Date:
Lab Three – Unit 3.2
Atomic Structure – A Journey into the Atom
Introduction:
Atoms are composed of subatomic particles, such as the protons and the neutrons,
which make up the nucleus of the atom and are similar in mass, and electrons, which are
found orbiting the nucleus in an electron, cloud and have a negligible mass. All atoms
contain the same kinds of particles but may differ in the number of each particle. This
accounts for the presence of isotopes and ions for the different elements.
This activity will allow you to use what you know about the composition of the
atom, as well as isotopes and ions, to describe sixteen atoms. The atoms are contained in
Ziploc bags and the subatomic particles are coded as follows.
Protons – black beans
Neutrons – white beans
Electrons – popcorn
Purpose:
Students will collect data and relate number of subatomic particles to atomic number,
mass number, electrical charge, atomic symbol, and name of element.
Equipment:
Materials:
Ziploc bags representing atoms
Procedure:
Analyze each Ziploc bag (atom) and record its vital statistics in the data table
provided.
Data Analysis:
1. List all sets of isotopes. How do you know they are isotopes?
2. List all sets of ions. How do you know they are ions?
Conclusions:
A nuclear reactor generates a very large amount of energy by splitting a uranium235 atom to produce Barium-139 and Krypton-94. How would each of these atoms be
represented using the coding system used for atoms #1 - 16?
Atomic Structure – A Journey into the Atom
Bag #
# of
# of
# of
Atomic
Mass
Electrical Chemical
Protons Neutrons Electrons Number Number Charge
Symbol
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
Name
Lab Three – Unit 3.2
Atomic Structure – A Journey into the Atom
Description
This activity will allow students to use what they know about the composition of
the atom, as well as isotopes and ions, to describe sixteen atoms. The atoms are
contained in Ziploc bags and the subatomic particles are coded as follows.
Protons – black beans
Neutrons – white beans
Electrons – popcorn
Time Frame: 50 minutes (1 class period)
Materials: Sixteen Ziploc bags representing atoms with different combinations of beans
and popcorn.
Procedures: See student handout. Atomic Structure – A Journey into the Atom.
Teacher Talk: Prepare Ziploc bags as follows
#1: 6 black beans, 6 white beans, 6 popcorn
#2: 1 black beans, 1 white beans, 1 popcorn
#3: 1 black beans, 2 white beans, 1 popcorn
#4: 6 black beans, 8 white beans, 6 popcorn
#5: 7 black beans, 7 white beans, 7 popcorn
#6: 7 black beans, 8 white beans, 7 popcorn
#7: 1 black beans, 1 white beans, 0 popcorn
#8: 7 black beans, 7 white beans, 10 popcorn
#9: 19 black beans, 21 white beans, 19 popcorn
#10: 19 black beans, 19 white beans, 19 popcorn
#11: 19 black beans, 19 white beans, 18 popcorn
#12: 8 black beans, 8 white beans, 8 popcorn
#13: 8 black beans, 8 white beans, 10 popcorn
#14: 15 black beans, 17 white beans, 15 popcorn
#15: 11 black beans, 13 white beans, 11 popcorn
#16: 11 black beans, 13 white beans, 10 popcorn
Extensions: Propose the following question to students.
Sometimes isotopes that are radioactive are used as medical tracers to detect disease.
One of the most useful is iodine-131 which is used to detect abnormalities in the thyroid
gland. The isotope can even be used to treat thyroid cancer since the radioactivity
destroys cancer cells. Cancers that cannot be treated with an internalized radioisotope
may utilize cobalt-60 for external radiotherapy. How would these two very useful
isotopes and their non-radioactive states be represented using the coding system?
Name:
Lab Partner:
Period:
Date:
Lab Four - Unit 3.3
Half-life Simulation
Introduction:
Radioactivity is something that is disconcerting to many people because of
pictures seen in war films or science fiction movies. Many elements have radioactive
isotopes that may be in the foods we eat, the things around us, the air we breathe.
Medical diagnosis and treatment have been improved and society has benefited from
radioactive medicines. Nuclear power plants provide energy to light our homes. Even
the fire alarms that most of us have in our homes function because of radioactivity. Some
isotopes of elements have unstable nuclei. As a result, some of the particles within the
nucleus are lost or emitted. This is known as nuclear decay. The amount of time for half
of the sample of a radioisotope to decay is know as its half-life.
In this experiment, you will use M&M plain candies or Skittles candies to
simulate the relationship between the passage of time and the number of radioactive
nuclei that will decay. As with real nuclei, the passage of time will be measured in halflives.
Purpose:
Simulate radioactive decay of radioactive nuclei using candy
Equipment:
Materials:
160 pieces of candy (M&M plain or Skittles)
Pizza Box (medium)
Graph paper
Procedure:
1. Place 160 pieces of candy in the pizza box. All candies should be marked side
up. Record the number of candies you started with (this is trial #0)
2. Close the container. Shake the box sufficiently so each candy has a chance to
flip several times.
3. Open container and remove the candies which are unmarked (marked side
down). Record in Data Table I the number of candies removed (this is trial
#1)
4. Repeat steps 2 & 3 five more times. At this point you will have simulated six
half-lives. You should have seven numbers in your final column, representing
the number of atoms remaining after zero, one, two, three, four, five and six
half-lives.
5. Following your teachers instructions, pool the class data by finding the total
number of atoms decayed for the whole class after the first half-life, the
second half-life, and so on using Data Table II.
6. Using the pooled data (the totals for each half-life), prepare a graph by
plotting the number of half-lives on the X-axis and the number of decayed
atoms for each half-life on the Y-axis.
Data Analysis:
DATA TABLE I
Half-lives
Trial #0 (start)
Trial #1
Trial #2
Trial #3
Trial #4
Trial #5
Trial #6
Undecayed (marked)
160
Decayed (unmarked)
0
DATA TABLE II
Lab
Pair
1
2
3
4
5
6
7
8
9
10
11
12
Total
1
2
Number of Half-Lives
3
4
5
6
1. Describe the appearance of your graph line. Is it straight or curved? Based on the
characteristics of your graph, why do you think radioactive decay is measured in halflives?
Conclusions:
1. Using the concept illustrated by your graph, determine how many undecayed nuclei
would remain in a sample of 600 after 3 half-lives?
2. Using the concept illustrated by your graph, if 175 undecayed nuclei remain from a
sample of 2800 nuclei, how many half-lives have passed?
3. How many half-lives would it take for a one mole sample of atoms (6.02 x 1023
atoms) to decay to 6.25% of the original number of atoms? After 10 ten half-lives,
would any of the radioactive material remain? Explain.
4. How could you modify this simulation to demonstrate that different isotopes have
different half-lives?
5. In this simulation, is there any way to predict when a specific atom (candy) will
decay? If you could follow the fate of an individual atom in a sample of radioactive
material, could you predict when it would decay? Explain.
Lab Four - Unit 3.3
Half-life Simulation
Description:
Students simulate nuclear decay in this lab. 160 candies are used to represent a
radioactive sample. The candies are flipped by shaking them in a box. All the candies
showing no label mark are removed and counted as decayed nuclei. Candies with label
mark showing are counted as undecayed nuclei and remain in the box. Students should
realize that this single step represents one half-life. Each time the step is repeated it
represents another half-life for that isotope.
Time Frame: 50 minutes (1 class period)
Materials: 2 King-sized bags of candy per group and a pizza box
Procedures: See student handout. Half-life Simulation.
Teacher Talk: Answers to Data Analysis: students should produce a curved line;
representing a natural decay of about ½ of the particles each time.
Extensions: Have students think of any other process that could be described in terms of
half-life? How could they modify this experiment to test their answers to their
conclusions?
Name:
Lab Partner:
Period:
Date:
Lab Five – Unit 4.2
Formula Writing and Chemical Names
Introduction:
A chemical formula is a combination symbols and numerical subscripts that
represents the composti9tion of a compound. The symbols indicate which elements are
present and the numerical subscripts indicate the relative proportion of each element in
the compound. These proportions can be predicted using the oxidation numbers
(charges) of the elements. When atoms acquire a charge they are called ions.
It is important that all scientists use the same system for writing chemical
formulas. This helps to ensure clear and consistent transmission of information.
Therefore, the following rules should be used for writing chemical formulas:
1. In a neutral compounds the sum of the oxidation numbers of the elements
(ions) must equal zero. One positive(+) charge will neutralize one negative(-)
charge.
2. Elements (ions) with a positive oxidation numbers (charges) are written first
3. When the relative proportion of the polyatomic ion in a ternary compound is
greater than one, the symbol for that ion must be enclosed in parenthesis and
followed by a numerical subscript indicating its relative proportion, as in the
ternary compound Aluminum Sulfate whose formula would be Al2(SO4)3.
Purpose:
Students will observe precipitate formation and write chemical formulas for the
precipitate. Students will name binary and ternary compounds formed.
Equipment:
Set of micropipettes
containing solutions of
Ag+, Co2+, Fe3+, Cu2+
Cl-, S2-, CO32-, OH-
Materials:
Plastic sleeve and work page
Procedure:
1. Insert the work page into the plastic sleeve and place on top of work table
2. Combine two drops of cation solution with two drops of anion solution in the
appropriate grid square. Be careful not to let the dropper touch the drops of
the other solutions.
3. Observe the reaction (if any) and record you observations on the
corresponding square of the data table.
4. Repeat two and three until you have combined all sixteen possible reactions.
5. In the data table, you will also find sixteen blanks, write the chemical name
and chemical formula in the corresponding blank for each reaction.
Data Analysis:
Data Table
Cl-
S2-
CO32-
OH-
Ag+
1
2
3
4
Co2+
5
6
7
8
Fe3+
9
10
11
12
Cu2+
13
14
15
16
1. __________________________
9. __________________________
2. __________________________
10. __________________________
3. __________________________
11. __________________________
4. __________________________
12. __________________________
5. __________________________
13. __________________________
6. __________________________
14. __________________________
7. __________________________
15. __________________________
8. __________________________
16. __________________________
Conclusions:
1. What is a chemical formula?
2. What information does a subscript in a chemical formula provide?
3. What is a formula unit?
4. When do you need to use a parenthesis in writing a chemical formula?
5. When do you need to use a roman numeral in the name of a compound?
Formula Writing Work Sheet
Cl-
S2-
CO32-
OH-
+
Ag
1
2
3
4
Co2+
5
6
7
8
Fe3+
9
10
11
12
Cu2+
13
14
15
16
Lab Five – Unit 4.2
Formula Writing and Chemical Names
Description:
In this experiment, students will combine four solutions containing cations with four
solutions containing anions and observe the ionic product formed. They will use
micropipettes containing the ions and a worksheet in a protective sleeve. The students
will then write chemical formulas and names for the precipitates formed.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Formula Writing and Chemical Names.
Procedures: See student handout. Formula Writing and Chemical Names.
Teacher Talk:
Prepare 100 mL of 1 M solutions of silver nitrate, cobalt(II) nitrate, iron(III)
nitrate, copper(II) nitrate, sodium chloride, sodium sulfide, sodium carbonate, and
sodium hydroxide. Label and fill micropipettes with ions to be tested. Provide each
work group with a protective sleeve and work sheet.
Extensions: Have students devise other processes to test for other combinations.
Name:
Lab Partner:
Period:
Date:
Lab Six – Unit 5
Water of Crystallization
Introduction:
Water is an intregal part of many ionic solids and such ionic solids are called
hydrates. The water in these solids is called water of hydration. A familiar example of a
hydrate is plaster of paris, which is the monohydrate of calcium sulfate, CaSO4H2O.
When water is added to plaster of paris and the material is allowed to set, it is gradually
transformed into a hard crystalline compound, calcium sulfate dihydrate, CaSO42H2O.
This is the material of plaster casts. The difference in composition between plaster of
paris and the plaster in casts is directly associated with the different degree of hydration
of the calcium sulfate in the two cases.
The water of hydration is not as tightly bound in the hydrated crystal as the ions
are. The water can usually be driven off by heating the crystals in a burner flame. The
material that remains after the water has been removed is called the anhydrous salt.
Purpose:
Students will observe the affect of heat on a hydrate.
Equipment:
Ring stand and ring
Bunsen burner
Clay triangle
Crucible top and bottom
Electronic scale
Crucible tongs
Scoop
Materials:
Hydrated barium chloride crystals
Procedure:
1. Clean and thoroughly dry a crucible and its cover by heating over a blue
flame. Cool and weigh the crucible and cover accurately to 0.01 g. All
masses are to be recorded in data table.
2. Place about 3 g of hydrated barium chloride crystals in the crucible (include
cover) and again weigh accurately.
3. Support the covered crucible on a clay triangle so adjusted in height that the
bottom of the crucible will be a short distance above the tip of the inner cone
of the Bunsen burner flame. Heat the crucible gently at first. Too rapid
heating may cause water of crystallization to be driven off explosively,
carrying some of the salt along with it. Gradually increase to the full intensity
of the flame and continue to heat strongly for about 10 minutes.
4. Allow to cool, and weigh the covered crucible and its contents.
5. Repeat the heating for an additional two minutes, cool and weigh again. This
repeated operation is called “heating to constant weight.” After the final
weighing complete the data table and determine the percentage of water of
hydration in crystalline barium chloride.
Data Analysis:
1. Mass of covered crucible and barium chloride crystals
__________ g
2. Mass of empty covered crucible
__________ g
3. Mass of crystalline barium chloride used
__________ g
4. Mass of covered crucible and contents, first heating
__________ g
5. Mass of covered crucible and contents, second heating
__________ g
6. Mass of anhydrous barium chloride
__________ g
7. Mass of water lost by heating
__________ g
8. % water =
_ _ _ g(mass of water)
x 100% =
_ _ _ g(mass of sample)
__________%
Conclusions:
Obtain the chemical formula for the crystalline barium chloride tested and calculate the
theoretical percentage of water present in the hydrated compound. Compare your results
to the theoretical value and describe your accuracy. What may account for any
inaccuracies?
Lab Six – Unit 5
Water of Crystallization
Description:
In this experiment, students will heat a hydrated compound to determine the percentage
of water of hydration.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Water of Crystallization.
Procedures: See student handout. Water of Crystallization.
Teacher Talk:
Caution: barium chloride dihydrate is toxic. Students should be aware of proper handling
and disposal of product. Other salts that can be used are magnesium sulfate heptahydrate
and copper(II) sulfate pentahydrate.
Extensions:
Calculate the number of molecules of water of hydration in one formula unit of a barium
chloride crystal, BaCl2xH2O
weight of crystalline
barium chloride )
weight of anhydrous
_ _ _ g ( barium chloride )
___g (
=
formula wt of crystalline
)
barium chloride
formula wt of anhydrous
___g (
)
barium chloride
X g(
1. The calculated formula wt. of crystalline barium chloride (X) ……….
2. The formula wt. of anhydrous barium chloride (BaCl2)
.………
3. Part of the formula wt. due to water ( 1 – 2 )
.………
part of formula
4.
_ _ _ ( wt due to water )
= _____
18 (formula wt of water)
molecules of
water in the formula
Name:
Lab Partner:
Period:
Date:
Lab Seven – Unit 6.2
Mass Relations in a Chemical Reaction
Introduction:
When performing an experiment involving chemical reactions, the scientist can
theoretically determine how much of a product should be produced. He/She will make
use of the fact that the coefficients of the reactants and products in a chemical equation
represent the relative number of moles of each reactant and product involved in the
reaction. From this information, the masses of products produced can be calculated.
In this experiment you will react baking soda with hydrochloric acid solution
converting the baking soda into table salt, water and carbon dioxide gas. Evaporation
will be used to separate the water from the salt. The mass of salt produced will be
determined.
Purpose:
Student will experimentally determine the mass of product produced in a chemical
reaction. Student will compare theoretical values for the reaction to the experimental
values obtained.
Equipment:
Ring stand and ring
Wire gauze
Evaporating dish
Crucible tongs
Bunsen burner
Watch glass
Centigram balance
Graduated cylinder
Materials:
Baking soda
Hydrochloric acid (3 M)
Procedure:
1. Write a balanced equation to describe the reaction.
Solid sodium hydrogen carbonate (baking soda) reacts with hydrochloric acid to
produce carbon dioxide gas, water, and aqueous sodium chloride.
2. In the reaction 3 M hydrochloric acid and 2.00 g of baking soda will be used. Predict
the volume of acid needed to completely react (digest) all of the baking soda.
(Molarity = moles/liter) Record your calculated volume in data table.
3. Weigh a clean, dry evaporating dish with watch glass to the nearest 0.01g. Record
this value
4. Into the dish and watch glass combination, add approximately 2.00 g of sodium
hydrogen carbonate. Weigh and record this new mass in data table. Calculate the
mass of sodium hydrogen carbonate used.
5. React the sodium hydrogen carbonate with the measured volume of HCl by slowing
adding (drop wise) the acid to the sodium hydrogen carbonate in the dish. Keep the
dish covered with the watch glass to prevent splattering. (If the quantity of HCl is not
enough to completely digest all the of the sodium hydrogen carbonate then add more;
no solid is left and no bubbles are formed)
6. When the reaction is complete, evaporate the solution containing the salt by placing
the evaporating dish with the watch glass over a medium burner flame.
7. Allow the evaporating dish to cool and weigh the dish, cover and contents. Record to
data table.
8. Use experimental data to calculate the amount of sodium chloride produced.
Data Analysis:
Data Table
1. Volume of 3 M HCl calculated
2. Mass of dish, cover and baking soda
3. Mass of empty dish and cover
4. Mass of baking soda used
5. Mass of dish, cover and product (salt)
6. Mass of salt produced
______ mL
______ g
______ g
______ g
______ g
______ g
Conclusions:
1. How did your calculated volume of HCl needed to react with the baking soda
hold up? Explain
2. Using the balanced equation for the reaction, calculate the theoretical yield of
sodium chloride from the mass of baking soda you used.
3. Calculate the percent yield of sodium chloride in your experiment
experimental mass
% yield =
x 100
theoretical mass
Lab Seven – Unit 6.2
Mass Relations in a Chemical Reaction
Description:
Students use an equation to make predictions on quantities of reactants needed in a
chemical reaction. They will use those values to perform a mass to mass reaction and
compare their results to calculated theoretical values.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Mass Relations in a Chemical Reaction.
Procedures: See student handout. Mass Relations in a Chemical Reaction.
Teacher Talk: Answers to Conclusions: Results may vary due to differing quantities of
baking soda used. However, accuracy should be maintained and mathematical
computations should be reviewed.
(2 g of baking soda should require less than 10 mL of 3 M HCl)
Extensions: Students could calculate volume of carbon dioxide gas released during
reaction. You may include concepts of limiting reagent.
Name:
Lab Partner:
Period:
Date:
Lab Eight – 7.1
Changes in Physical State
Introduction:
Matter can exist in three different physical states – the solid state, the liquid state,
or the gas state. In a pure substance, changes of physical state take place at discrete
temperatures, which are constant and which are characteristic for each substance. In this
experiment, you will closely examine what happens when a pure substance undergoes a
change in physical state. Specifically, you will investigate the melting and freezing
behavior of a sample of an organic compound called paradichlorobenzene, C6H4Cl2. You
will be concerned chiefly with two questions; first, does the liquid paradichlorobenzene
begin to freeze at the same temperature that solid paradichlorobenzene begins to melt?;
secondly, how does the temperature of the paradichlorobenzene change (if it does
change) between the time melting or freezing just begins and the time that freezing or
melting is complete?
Purpose:
The student will observe and use graphics to explain the behavior of
paradichlorobenzene during melting and freezing. The student will consider what
happens to the energy that is put into or removed from paradichlorobenzene during the
process of melting or freezing.
Equipment:
Ring stand and ring
Bunsen burner
400 mL beaker
Wire gauze
Materials:
Test tube of solid paradichlorobenzene
with an imbedded thermometer
Procedure:
1. Place a test filled with paradichlorobenzene (C6H4Cl2) and a thermometer into a large
beaker 2/3 full of water that has been heated just to the boiling point of water. Begin
to take temperature readings every 15 seconds as the solid melts. DO NOT REMOVE
THERMOMETER. Continue to record temperatures until the solid is completely
melted.
2. Place the test tube containing the liquid paradichlorobenzene and the thermometer
into a large beaker 2/3 full of water, at room temperature. Begin to collect
temperature data immediately, taking readings every 15 seconds as before. DO NOT
REMOVE THERMOMETER. Continue to record temperatures until freezing is
complete.
3. Return the test tube with the solid paradichlorobenzene and thermometer to the
reagent cart and clean up your lab station.
Data Analysis:
1. Prepare a data table to record the time and temperature every 15 seconds for both
melting and freezing. It may take up to 6 minutes (480 seconds) to complete the
change.
2. Make a graph of your data for the melting process. Choose a scale that will fill a full
sheet of graph paper. Plot time on your horizontal axis and temperature on the
vertical axis. Connect the points in a smooth curve. Plot the data for the freezing
process on the same graph. (BE SURE TO LABEL YOUR GRAPH)
3. Determine the point at which the two curves intersect. This point is the melting point
for the solid phase and the freezing point of the liquid phase.
Conclusions:
1. Explain the shape of the graph in terms of the energy changes that are occurring in the
sample as it heats up and melts and as it cools down and freezes.
2. What happens to the temperature of the substance while it is actually melting or
freezing?
3. Explain in your own words what is going on at the molecular level as liquid
paradichlorobenzene cools down and freezes.
4. Compare the value you obtained for the freezing point for paradichlorobenzene with
the values obtained by others in the class. Explain any similarities or differences.
5. How would an increase in the amount of paradichlorobenzene used affect the shape
of the graph? Explain.
Lab Eight – 7.1
Changes in Physical State
Description
In this experiment, students will closely examine what happens when a pure substance
undergoes a change in physical state. Specifically, they will investigate the melting and
freezing behavior of a sample of an organic compound such as paradichlorobenzene,
C6H4Cl2. The students will use a test tube of solid compound with an imbedded
thermometer. They will place it in warm water to observe melting and in room
temperature water to observe freezing.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Changes in Physical State.
Procedures: See student handout. Changes in Physical State.
Teacher Talk:
The test tube containing paradichlorobenzene needs to be prepared prior to the activity.
Melted paradichlorobenzene is poured into a test tube, about ½ full, and a thermometer is
inserted and the substance is allowed to cool.
Extensions:
Name:
Lab Partner:
Period:
Date:
Lab Nine – Unit 7.2
Ideal Gas Law
Introduction:
When the temperature, pressure, and volume of a gas is measured, the ideal gas
law allows the number of moles of the gas to be calculated. If the percent composition is
known, the number of moles allows the molecular formula to be calculated. The ideal
gas equation is the following: PV = nRT
Purpose:
The student will collect a sample of gas by water displacement, measure the volume,
temperature and pressure of the gas and calculate the molecular mass of the gas using the
ideal gas equation.
Equipment:
100 mL graduated cylinder
One gallon paint pail
Thermometer
Centigram balance
Materials:
Disposable butane lighter
Procedure:
1. Half fill a paint pail with water. Record the temperature of the water.
2. Place the lighter under water in the pail. Remove the lighter, shake off the water, and
dry the outside with a towel. Then, mass the lighter to the nearest 0.01 gram. Record
this measurement in the data table.
3. Place the lighter back in the pail. Fill the graduated cylinder completely with water.
Cover it with your hand and carefully invert it into the pail of water. Remove your
hand, keeping the mouth of the graduated cylinder under water.
4. Release the gas from the lighter by pressing the small lever near the flint wheel.
Release the gas under water being careful that all of it is collected in the graduated
cylinder by water displacement. Release enough gas to fill the graduated cylinder to
within 3 mL of its calibrated capacity. DO NOT EXCEED THE SCALE AND DO
NOT LIFT THE CYLINDER OUT OF THE WATER.
5. Allow the gas to reach room temperature (about 2 minutes). Then adjust the level of
the water inside and outside the graduated cylinder until they are the same by raising
and lowering the cylinder in the pail. With the pressure inside and outside the
graduated cylinder the same, read the volume of gas collected using the cylinder’s
calibration. Record this volume on the data table.
6. Remove the lighter from the pail. Shake off any excess water and dry off with a
towel. Measure and record the mass of the lighter.
7. Record the barometric pressure (from your instructor).
Data Analysis:
Data Table
1. Mass of disposable lighter and contents
before experiment
2. Mass of disposable lighter and contents
after experiment
_________g
_________g
3. Mass of gas released from lighter
_________g
4. Volume of gas collected from lighter
_________mL
5. Atmospheric pressure
_________atm
6. Room temperature (temp of water and gas)
7. Water vapor pressure at water temp
(see water vapor pressure table)
_________C
8. Partial pressure of dry gas (Pgas = Patm – PH2Ovapor)
_________atm
_________atm
Conclusions:
1. Use the ideal gas equation to determine the number of moles of the gas collected.
Remember to convert the room temperature from Celsius to Kelvin. Use the partial
pressure of the dry gas in the formula. Use the number of moles collected to
calculate the molecular mass of the gas.
2. Butane is the most common gas found in disposable lighters. Compare your
calculated molecular mass with that of butane (obtain chemical formula from
teacher). Use the difference between these two numbers to calculate percent error.
3. Can the same experimental techniques be used to determine the molecular mass of all
gases? Explain your answer.
Lab Nine – Unit 7.2
Ideal Gas Law
Description
In this experiment, students will experimentally determine the molecular mass of a gas
contained in a disposable lighter. They will determine this value using the ideal gas law.
Volume of gas is determined by water displacement, temperature of gas will be the same
as the temperature of the water, and pressure will be atmospheric pressure in the room.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Ideal Gas Law.
Procedures: See student handout. Ideal Gas Law.
Teacher Talk:
Lighter needs to be dried thoroughly as the mass of the gas collected is small and a large
percent error will result if not. Students need to make sure that the levels of the gas and
water are the same in the cylinder before recording volume. Vapor pressure table is in
CRC handbook. Disposable lighters may be purchased at Dollar General or K-mart.
Butane is C4H10. Not all gases can be used for this experiment since some chemically
react with water or are soluble in water.
Extensions: Students could try several different brands of lighters and compare results.
Students could try the experiment using helium gas from a balloon.
Name:
Lab Partner:
Period:
Date:
Lab Ten – Unit 10.2
Solubility Curve of a Salt
Introduction:
The solubility of a solute is the amount of solute dissolved in a given amount of a
certain solvent at equilibrium, under specified conditions (the ability to dissolve).
Increasing the temperature usually increases the solubility of solids in liquids
(endothermic changes only), and decreasing the temperature has the reverse effect
(exception, gaseous solutions).
Purpose:
Students will construct a solubility curve representing data collected experimentally.
Masses of salt will be varied and temperatures required to dissolve it will be recorded.
Equipment:
Hot plate
Large test tube
400 mL beaker
Thermometer
Centigram balance
Scoop
Graduated cylinder
Materials:
Ammonium chloride crystals
Procedure:
1. Measure out exactly 4.00 grams of the salt (NH4Cl) and place in a large test tube.
2. Add exactly 10 mL of distilled water to the test tube containing the salt. Place the
thermometer in the tube (may be used to stir the solution)
3. Using a hot water bath dissolve the salt. Remove the test tube from the bath and
record the temperature when the first trace of crystals appear in the tube. NOTE: you
may need to place test tube under running tap water to cool it. At this point the
solution is saturated ( to prevent supersaturation, stir the solution with the
thermometer).
4. To the above solution add an additional 1.00 g of the salt an repeat procedure 3. DO
NOT ADD EXTRA WATER.
5. Repeat procedure 4 two more time to obtain a total of four trials.
Data Analysis:
Temp Data
4.00 g _______C
5.00 g _______C
6.00 g _______C
7.00 g _______C
Graph a solubility curve for the salt using the x-axis for the temperature and the y-axis for
the mass used per 10 mL of water.
Conclusions:
1. Using the graph, determine the solubility of NH4Cl at room temperature (25C) and at
60C.
2. From your data, is the additional energy needed to increase the solubility proportional
to the amount of solute added? Explain.
3. Is there a limit to the amount of solute that a solvent can be forced to dissolve?
Explain.
4. What use could be made of a solubility curve for a certain salt?
Lab Ten – Unit 10.2
Solubility Curve of a Salt
Description
In this experiment, students start with a given amount of a salt in 10 mL of water and
increase only the mass by 1.00 gram each trial. The temperature is recorded for each trial
when crystals are observed as the solution cools providing data for a solubility curve.
Time Frame: 50 minutes (1 class period)
Materials: See student handout. Solubility Curve of a Salt.
Procedures: See student handout. Solubility Curve of a Salt.
Teacher Talk:
If students experience problems with crystallization in the first trial, use ice water bath.
Last trial will require boiling water. Remind students not to add additional water to the
test tube.
Extensions:
Potassium dichromate may be used as a solute. If used, start with 2.00 gram sample and
increase the amount in solution by 2.00 g each trial.
Name:
Lab Partner:
Period:
Date:
Lab Eleven – Unit 11.2
Neutralization and Titration with Acid and Base
Introduction:
Neutralization occurs when the hydronium ion from an acid interacts with a
hydroxide ion from a base, on a one to one basis, forming water in the process. A salt is
always a byproduct of this type of reaction. Titration is the progressive addition of an
acid to a base to achieve neutralization. The point at which the acid and base are in
equivalent amounts is called the end point.
Purpose:
Students will explore the processes of neutralization of an acid and a base. Titration will
be used to determine the % of acetic acid in vinegar.
Equipment:
10 mL graduated cylinder
50 mL graduated cylinder
Ring stand
Burette clamp
50 mL burette (2)
Test tube (3)
250 mL flask
150 mL beaker (2)
Eye dropper
Test tube rack
Materials:
Bromthymol blue indicator
Phenolphthalein indicator
1 M NaOH
1 M HCl
0.3 M HCl
Vinegar
Unknown NaOH solution
Procedure:
Investigation A: (Neutralization)
1. In a test tube place 1 mL of 1 M HCl and add 2 drops of Bromthymol blue indicator
and record the color ________________
In a test tube place 1 mL of 1 M NaOH and add 2 drops of Bromthymol blue
indicator and record the color ________________
2. Predict the volume of 1 M NaOH needed to completely neutralization 3 mL of 1 M
HCl (complete neutralization only occurs when equivalent quantities of hydronium
and hydroxide ions are present) ________________
a.
Transfer 3 mL of 1 M HCl to a test tube, add 2 to 3 drops of Bromthymol blue
indicator to the acid solution
b.
Slowly add 1 M NaOH, drop wise, to the acid solution until complete
neutralization occurs. This occurs when the color of the indicator changes.
Record the number of milliliters of NaOH used (1 mL = 20 drops) ___________
c.
How did your findings compare with the predicted volume?…..
Investigation B: (Titration)
1. Determine the molarity of unknown NaOH solution
a.
Clean a 50 mL burette with fresh water. Fill the burette 0.3 M HCl solution.
b.
Clean a 50 mL burette with fresh water. Fill the burette NaOH solution.
(Unknown Molarity)
c.
Drain 15 mL of HCl from the burette into a 250 mL flask. Add 25 mL of
distilled water and 2-3 drops of Phenolphthalein indicator to the flask. Titrate the
HCl solution with the NaOH solution until the endpoint is reached (color change).
Record the volume of the base used in data table. Record the volume of acid used
in data table.
d.
Calculate the strength (Molarity) of the base from your data
Strength (acid) x Volume (acid) = Strength (base) x Volume (base)
1. Determine the strength of vinegar
a.
Repeat the processes, 1.a and 1.b, this time substituting vinegar solution for
the acid. Use the calculated strength of the NaOH as the know concentration and
titrate to find the concentration of the vinegar solution.
b.
Drain 15 mL of vinegar from the burette into a 250 mL flask. Add 25 mL of
distilled water and 2-3 drops of Phenolphthalein indicator to the flask. Titrate the
vinegar solution with the NaOH solution until the endpoint is reached (color
change). Record the volume of the base used in data table. Record the volume of
acid used in data table.
c.
Calculate the strength (Molarity) of the vinegar from your data
Data Analysis:
Volume of HCl used
Data Table B.1
________ mL
Strength of HCl
________ M
Volume of NaOH used
________ mL
Strength of NaOH
________ M (calculated)
Volume of NaOH used
Data Table B.2
________ mL
Strength of NaOH
________ M
Volume of vinegar used
________ mL
Strength of vinegar
________ M (calculated)
Conclusions:
1. A 1 M solution of vinegar (acetic acid) would contain, because of its formula mass,
60 grams per liter of water. To express this value in %, you would use 6 g in 100 mL
of water. To find the percentage of acetic acid in vinegar solution, all you need to do
is to multiply the molarity of the vinegar by 6. What is the % of acid in the vinegar?
Lab Eleven – Unit 11.2
Neutralization and Titration with Acid and Base
Description
Students will explore the characteristics of neutralization and titration of an acid and a
base. They will use experimental data to calculate (compare) strengths of acids and
bases.
Time Frame: 100 minutes (2 class periods)
Materials: See student handout. Neutralization and Titration with Acid and Base.
Procedures: See student handout. Neutralization and Titration with Acid and Base.
Teacher Talk:
Quantities of Solutions needed
1 M NaOH
100 mL per class
1 M HCl
100 mL per class
0.3 M HCl
1 L per class
? M NaOH
2 L per class (unknown base should be 0.5 M)
Vinegar
1 L per class (5% vinegar is 0.83 M)
Extensions:
For neutralization study you might suggest they compare the strength of ammonia water
to the strength of vinegar from the volume of each used (qualitative).