Laboratory Manual
Name______________________________
Honors Chemistry
Teacher________________Period_______
Contents
Page
Lab Notebook Format……………………………………………………………………………….3
Experiments:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
Metric Measurements………………………………………………………………5
Density……………………………………………………………………….…….7
Law of Conservation of Matter …………………………………………………...9
Physical and Chemical Changes…………………………………………………..11
Determining the Specific Heat of Iron…………………………………………....13
Temperature of a Bunsen Burner Flame…………………………………………..14
Law of Definite Composition……………………………………………………..15
Al Foil Lab………………………………………………………………………...17
Absolute Zero Lab…………………………………………………………………19
Precipitation Reactions: Formula Writing…………………………………….…...21
Heat of Fusion of Ice……………………………………………………………....23
Qualitative Analysis 1: Relative Solubilities……………………………………...25
The Mole Lab……………………………………………………………………...29
Hydrates…………………………………………………………………………...31
Determining The Value of R.……………………………………………………...33
Single and Double Replacement Reactions……………………………………….35
Predicting Products………………………………………………………………..37
Qualitative Analysis 2: Ag+, Hg22+, and Pb2+…………………………………….41
Molar Volume of a Gas …………………………………………………………...45
Quantitative Study of Chemical Reactions………………………………………..47
Titration 1…………………………………………………………………….…….51
Titration 2…………………………………………………………………….…….53
Determining Solution Concentration: Using a Spectrophotometer………………..55
Periodic Law……………...………………………………………………………..57
Building Molecular Models…………………………………………………….….61
Solutions Lab……………………………………………………………………....67
Heat of Reaction………………………….………………………………….…….69
Heat of Formation of Magnesium Oxide………...…………………………….…..71
Stoichiometry Lab………………………………………………………………….73
Moles of Iron and Copper…………………………………………………….……75
2
3
Lab Reporting and Assessment
Laboratory Notebook:

You are required have a marble composition book for lab work with your name and period number
labeled on the front cover.

Designate the first two pages of the book (front and back of 1 page is fine) as a table of contents.
Complete the table of contents as you do the labs throughout the year. Number the pages of your
notebook in the top, right corner beginning with the table of contents page (page 1). Number the
odd numbered pages only (even numbered pages will be the back of each page). You may write
on one or both sides of the pages at your discretion. Be sure to include the page number that each
lab starts in your table of contents.

All information entered into lab notebooks must be neat and logically organized.

It is highly recommended that lab notebooks be completed in pencil.
Pre-Lab Requirement: Before coming to class on the day of a lab, you must read the entire experiment
and complete the following in your lab notebook:
 Title of Lab (Be sure to include in the table of contents as well)

Date that the lab was performed

Pre-Lab Score: __________ Post-Lab Score:___________
The score will be filled in by the teacher, but please have the labels and lines prepared.

Objective(s): State the purpose(s) of the experiment.

Procedure: Write, in a paragraph, a brief synopsis of the procedure conducted. Please note that
this is a summary of the procedure. Do not simply copy the procedure. The intention is to provide
a big picture understanding of the experiment. Do not include specific details.

Data and Results: Data are measurements and observations recorded during an experiment and
results are calculations performed with these data. Prepare tables, in advance, for the recording of
the necessary data to be collected during the experiment and the results to be calculated. The data
and results sections may be combined into one table or separated at your discretion. Data and/or
observations are to be recorded throughout the experiment in these tables prepared in advance.

Error Discussion Table: A table must be set up to discuss possible errors for each procedure step
of the lab. The table should have 3 columns with the following headings: Procedure Step,
Possible Error, and Effect on Objective (you should restate the objective here.)
Point of emphasis: The number one cause of laboratory accidents, mistakes, and
misunderstandings of experiments is not being prepared in advance. Therefore, the pre-lab
requirement explained above is mandatory. Students who do not have their lab notebooks on the
day of the lab or have not completed these requirements will not be permitted to do the
experiment at that time and will not receive credit for the pre-lab requirements. The lab will have
4
to be made up after school. The pre-lab requirements will still be in effect, but points will not be
awarded. This is a lab safety issue, so exceptions to this policy will not be made. Students
making up lab work after school due to absences will still be able to earn the credit for the pre-lab
requirements.
Post-Lab Requirement: There will usually be some time at the conclusion of an experiment to begin the
following:

Graph: (If applicable) Graphs must be done on graph paper and attention must be paid to
neatness and precision. Be sure to include a title and properly label the axes. Cut the graph paper
down to size and attach it to the appropriate page in your lab notebook.

Sample Calculations: Show the work for one of each calculation type. Solutions to calculations
must be included in the results section. This section is simply for showing work. There should
not be any important information in this section that does not appear in another section.

Questions: It is not necessary to recopy the question, but be sure that they are properly numbered.
Work must be shown for calculation based questions.

Conclusion: State whether or not the objectives of the experiment were met. Support your
conclusions with details from the experiment. For example, if the objective of an experiment is to
determine the molar mass of an unknown, the determined molar mass must be stated in the
conclusion. Also, it is appropriate to mention sources of error in the conclusion except when there
is a specific question regarding error.
Lab Assessment


Periodically, a lab test or quiz will be given which will assess all of the labs performed since the
last lab test. Students will sometimes be permitted to use their laboratory notebooks for these
tests, but may not use the lab handouts. Forgetting to bring your lab notebook on the day of the
assessment does not excuse you from taking the test.
Lab notebooks will be checked at the beginning of each lab. At this time, notebooks will be
checked for completion and thoroughness of the pre-lab requirement of the current lab. The postlab requirement due date will vary for each lab. You will be awarded points for each of these
requirements and these points earned will be factored into the lab test to follow.
The lab quizzes will vary in points depending on the labs assessed and the number of questions. Points
for the pre-lab and post-lab discussion will be added to the assessment grade.
5
Metric Measurements Lab
Name: ______________________
Copy this table in your notebook. Work for all bolded items must be shown in your notebook.
Measurement
(with unit)
Convert to:
1
Your height (in)
2
The mass of a #4 rubber stopper
cg
3
The diameter of a 250 mL beaker
µm
4
5
The volume of a chemistry textbook
(cm3) (record 3 measurements)
The area of the seat of a student's desk
(cm2) (record 2 measurements)
km
mm3
ft2
6
Your weight (lb)
7
The height of a lab bench
Gm
8
The mass of a clay triangle
mg
9
The circumference of your wrist (cm)
10
The mass of a pair of safety goggles
11
The volume of liquid in the 1 L
Graduated Cylinder on the front bench
dL
12
The length of your stirring rod
nm
13
The width of your lab drawer
Mm
14
The volume of an evaporating dish
pts
15
The mass of your crucible tongs
hg
16
The volume of 43 drops of water
μL
17
The length of the teacher's lab desk
yds
18
The area of the glass plate (cm2)
in2
19
The length of your foot (wearing shoe)
nm
20
The mass of your Erlenmeyer flask
oz.
(record 2 measurements)
hg
ft
dag
6
7
Density
Name___________________________
Period______Date_________________
Procedure:
1. Determine the mass and volume (water displacement) of 5 different quantities of either copper,
lead, chromium, silicon, iron or aluminum. Do not weigh the metal when it is wet.
Strive for a 1 mL or larger difference between all volume measurements.
2. Return the metal to the beaker in the front of the room so that it can be dried and reused.
Graph:
1. Prepare a mass vs. volume graph, using proper graphing techniques outlined below. Do not
connect the plotted points. Draw a line of best fit.
2. Determine the slope of the line (show work). This is your experimental value for density. Include
a unit.
Elements of a good graph:
 Constructed on graph paper with a ruler.
 Title that reflects the meaning of the graph. A title needs to explain the graph.
 Axes are labeled with the property being measured and the unit
 Each axis is numbered with equal intervals that begin at the origin with 0. The intervals
are chosen in a fashion that includes all of the data and maximizes the size of the graph.
The axes do not need to be numbered with the same intervals.
Questions:
1. Determine the percent error for your experiment. Use the slope for your density! Accepted values
in book and on periodic table.
Experimental  Accepted
% Error 
 100
Accepted
2. List two sources of error for this experiment (be specific – do not include human error, or
incorrect reading of the balance or graduated cylinder).
3. Based on your graph, describe the precision of your experiment.
4. Describe the difference in appearance of a graph of (A) precise, accurate data and a graph of (B)
precise, inaccurate data.
5. Describe the difference in appearance of a graph of (A) precise, accurate data and a graph of (B)
imprecise, inaccurate data.
6. Using your experimental value for the density of your metal, determine the weight, in pounds, of
a 0.856 ft3 block of the metal. Show all work.
8
9
Law of Conservation of Matter
Name___________________________
Period______Date_________________
Introduction:
In part I of this lab, you will be reacting separate solutions of two compounds to produce two different
compounds. One of these products will be insoluble in water and will therefore precipitate out of
solution. The other product will remain dissolved. If the precipitate is filtered out of the solution and
dried, its mass can be determined. If the remaining filtrate is heated to dryness, the other compound will
be left as a residue and its mass can be determined. Knowing the masses of initial substances reacted and
the masses of the substances produced, the law of conservation of matter can be observed. In part II, a
different reaction will be performed, but the reactant and product masses will be compared as in part I.
Procedure:
Part I
1. Weigh 2.12 g of strontium nitrate — Sr(NO3)2 — and place it in a small beaker. Remember to
weigh the paper first. Similarly, weigh 1.06 g of sodium carbonate − Na2CO3 − and place it into
another small beaker. Add 10 mL of deionized water to each and stir until dissolved.
2. Combine the two solutions.
3. Weigh a piece of filter paper. Set up a filtering apparatus using a long stem funnel, the folded
filter paper, and an Erlenmeyer flask. Gently swirl the beaker containing the products in order to
mix up the precipitate and begin to slowly pour it into the filter. Do not allow the contents to rise
higher than the filter paper in the funnel. Continue adding the product to the filter (while swirling
the beaker) until the beaker is empty. Add a small amount of water to the beaker to rinse it and
pour that into the filter as well.
4. After all of the liquid has passed through the filter, carefully remove the filter paper (use forceps)
and lay it flat on a watch glass. Place the watch glass under the heat lamp in the fume hood.
5. Weigh a 250 mL beaker. Pour the filtrate into this beaker and heat it with a Bunsen burner (ring
stand, etc.) to dryness.
6. Once dry, weigh the filter paper containing the precipitate and the beaker containing the residue.
Part II
1. Add 20 mL of dilute hydrochloric acid – HCl – to an Erlenmeyer flask. Weigh the flask and
contents.
2. Weigh 1.00 g of calcium carbonate – CaCO3 – and add this to the Erlenmeyer flask.
3. When the reaction is complete, weigh the flask and contents again.
Part I Calculations: (Show work and units)
 Total mass of reactants
____________

Mass of precipitate
____________

Mass of residue
____________

Total mass of products (precipitate and residue)
____________

Amount of mass gained or lost
____________
10
Part II Calculations: (Show work and units)

Total mass before mixing
____________

Total Mass after Mixing
____________

Amount of mass gained or lost
____________
Questions:
1. State the Law of Conservation of Mass.
2. Clearly define the following terms:
a. filtrate
b. residue
c. precipitate
d. solution
3. Interpret the data from Part I of the lab and indicate if your results are consistent with the Law of
Conservation of Mass (within a reasonable margin of error). Explain.
4. List some specific sources of error for part I of this experiment. Sources of error need to be
specific from this lab. Generic responses such as “human error”, “incorrect measurements”, and
“inaccurate balances” are not acceptable.
5. Interpret the data from Part II of the lab and indicate if your results are consistent with the Law of
Conservation of Mass (within a reasonable margin of error). Explain.
11
Physical and Chemical Changes
Name___________________________
Period______Date_________________
***Measurement tip: Most liquid measures in this experiment are 2 mL. To approximate this volume,
measure 2 mL of water with a graduated cylinder and pour it into a test tube. Note the height of the water
in the test tube. Fill test tubes to that height throughout the lab when a 2 mL is called for.
Procedure:
1. Grind several crystals of copper(II) sulfate pentahydrate (CuSO4 ∙ 5H2O) with the mortar and pestle.
Record observations. Transfer the powder into a test tube and heat it in a Bunsen burner. Note
changes in the appearance of the solid and any changes at the mouth of the test tube. Allow the test
tube to cool for several minutes and then add tap water one drop at a time until a noticeable change
occurs. Disposal: add water and heat over Bunsen burner. Discard in drain with water.
2. Obtain a 5 cm piece of magnesium ribbon and cut it into two equal sized pieces. Grip one piece of
magnesium with crucible tongs and hold it in the flame until it ignites. Hold the burning Mg over an
evaporating dish and do not look directly at it while it is burning. When it stops burning, allow the
ashes to fall into the evaporating dish and observe. Disposal: metal trash can.
Add 2 mL of 3.0 M HCl (hydrochloric acid) to two clean test tubes and place them in a test tube rack.
Add the other piece of magnesium to one of the test tubes and the ash from the burned magnesium to
the other. Observe. Disposal: HCl waste beaker in classroom.
3. Add 2 mL of 1.0 M NaOH (sodium hydroxide) to one test tube and 2 mL of 1.0 M NH4Cl
(ammonium chloride) to another. Using proper wafting techniques, observe the odor of each solution.
Combine the contents of each test tube in a small beaker. Gently swirl the beaker and observe the
odor. Disposal: discard in drain with water.
4. Pour 15 mL of tap water into a beaker. Add 3 scoops of copper (II) chloride (CuCl2) to the water. Stir
to dissolve. Measure the temperature of the solution. Obtain a piece of aluminum foil and roll it up in
a ball. Add the foil to the beaker and insert the thermometer. Record the highest temperature reached.
Disposal: solids in metal trash can, liquid in drain with water.
5. Measure 1.0 g of solid sodium sulfate (Na2SO4) and put it in a beaker. Add 50 mL of deionized water.
Stir until dissolved. Add 2 mL of this solution to a test tube and 20 mL of it to an evaporating dish.
The rest can be discarded. Place the evaporating dish on a ring stand with a wire gauze and heat the
evaporating dish to dryness. While it is heating, add 2 mL of 0.10 M Ba(NO3)2 (barium nitrate) to a
new test tube and observe. Pour the test tubes containing the barium nitrate and the sodium sulfate
into a beaker and observe. Disposal: barium waste beaker in classroom.
6. Add 2 mL of 0.10 M HCl to one test tube and 2 mL of 0.10 M NaOH (sodium hydroxide) to another.
Measure the pH of each solution by dipping a clean stirring rod into the solution and touching a piece
of pH paper with the stirring rod. Compare the color of the paper to the color chart to obtain the pH
value. After the pH has been measured for each solution, combine the solutions in a beaker. Gently
swirl the beaker and measure the pH of the resulting solution. Disposal: HCl waste beaker in
classroom.
7. Measure 0.20 g of iron filings and 0.20 g of powdered sulfur. Combine the two solids in a small
beaker and mix them thoroughly. Carefully pour the mixture into a small, disposable test tube. Run a
magnet along the outside of the test tube and record observations. In the fume hood, heat the test tube
in Bunsen burner for several minutes until the mixture glows. Allow it to cool for several minutes.
Run a magnet along the outside of the test tube again. Observe. Disposal: broken glass container.
12
Analysis:
Determine whether each procedure step was a physical or chemical change and explain why you
made this determination. (Hint: Most procedure steps will have more than one change.)
Questions:
1. List three pieces of evidence from this experiment that can be used to determine if a physical
change has taken place.
2. List three pieces of evidence from this experiment that can be used to determine if a chemical
change has taken place.
3. Imagine that two people are sitting around a campfire heating marshmallows. How can one
person cause a physical change to occur with the marshmallow while the other person causes a
chemical change?
13
Name___________________________
Period______Date_________________
Determining the Specific Heat of Iron
1. Note: There are two data tables provided below. One pertains to deionized water and the second
pertains to iron. Be sure to record the data in the appropriate place.
2. Determine the mass of a piece of iron. Use nichrome wire to attach the iron to your ring stand, as
shown by your teacher.
3. Measure and add 100 mL of deionized water to a Styrofoam cup. Density of water = 1.00g/mL.
4. Record the temperature of the water.
5. Fill a 400 mL beaker two-thirds full of tap water, and bring it to a boil on the ring-stand set-up. Once
the water is boiling, record the temperature of the boiling water.
6. Submerge the piece of iron in the boiling water for 5 minutes.
7. Turn off the flame and carefully lift the iron from the boiling water and quickly submerge it in the
water in the Styrofoam cup. Record the highest temperature reached by the water.
8. Remove the iron from the water. Dry the iron and leave it set-up for the next portion of the lab.
9. Use your collected data to determine the specific heat of iron.
Calculations:
1. Show calculations for the determination of the specific heat of iron. Be sure to use units.
2. The accepted value for the specific heat of iron is 0.449 J/g •°C. Determine the % error for your
results.
14
Name__________________________
Period______Date________________
Temperature of a Bunsen Burner Flame
Procedure:
1. Fill the cup half way with water. Measure the volume of water using a graduated cylinder and
note the total volume of water in the cup. Record the mass of water (density = 1.00 g/mL).
2. Measure and record the temperature of the water.
3. Turn on the Bunsen Burner and position it so that the metal is in the hottest part of the flame (the
inner blue cone).
4. Allow the metal to heat in the flame for 5 minutes.
5. Turn off and remove the Bunsen Burner.
6. Slowly lift the cup of water so that the metal is submerged. Be prepared for a loud sizzling noise.
Do not be startled by it.
***DO NOT ALLOW THE METAL TO TOUCH THE SIDE OF THE CUP***
7. Allow the metal to remain in the water for 1 minute while gently stirring. After the 1 minute,
record the temperature.
Analysis: Using the data and your experimental value for the specific heat of iron, determine the
temperature of the Bunsen burner flame. Show all work.
Assume:
The initial temperature of the iron is the same temperature as the flame
Heat lost by metal = Heat gained by water
Specific Heat of water = 4.184 J/goC
15
Law of Definite Composition
Name___________________________
Period______Date_________________
Introduction:
Elements are a kind of matter that cannot be decomposed by ordinary chemical means. Compounds are
chemical combinations of elements. The law of definite composition (a.k.a. law of definite proportions)
states that the elements forming a compound always combine in the same proportion by mass. Water,
H2O, is always a chemical combination of hydrogen and oxygen in a 1:8 ratio by mass (1 g of hydrogen
for every 8 g of oxygen). If a mixture of hydrogen and oxygen were reacted in some other mass ratio, for
example 1:2, water would still be formed but some hydrogen would remain unreacted.
In this experiment, you will examine the reaction between magnesium metal and oxygen gas. When
magnesium burns, it chemically combines with oxygen to form magnesium oxide. You will determine
the mass ratio of magnesium to oxygen.
Procedure:
1. Coordinate with another lab group in the class. Each of the two groups is to obtain a piece of
magnesium (15 cm – 25 cm) of different lengths, complete the procedure, and exchange data
with the other group.
2. Determine the mass of an empty crucible and cover. Coil (spiral) the magnesium loosely, add it to
crucible and reweigh it with its cover.
3. Set up a ring stand with an iron ring, clay triangle and Bunsen Burner. Support the covered
crucible in the clay triangle and begin to heat it gently. Periodically, lift the lid slightly to allow
oxygen to enter and to determine if the magnesium is reacting (it will glow strongly and produce
smoke). It is important to minimize the amount of smoke that escapes.
4. Once the magnesium is reacting, continue heating the crucible and slightly lifting the lid until no
more smoke is present. At that point, rest the lid on the crucible and triangle so that it is only
partly on and heat the crucible intensely for 5 minutes. Allow the crucible to cool for several
minutes.
5. In addition to magnesium oxide forming in the crucible, magnesium nitride may also form. To
remove the nitride, add ten drops of distilled water (this will convert the magnesium nitride to
magnesium oxide and will release ammonia gas – see if you can smell the ammonia being released
by using proper wafting techniques). Heat the crucible intensely again for five minutes with the
lid off.
6. Allow the crucible to cool for several minutes and determine the mass of the crucible, cover and
product. Consider this mass as the mass after the first heating.
7. Reheat the crucible for another 3 minutes without the lid. Allow it to cool and reweigh it for the
second time. It is necessary to have two consecutive masses within 0.03 g of one another. If the
second mass is not within 0.03 g of the first, then a third heating and massing will be necessary.
Use the final mass for all calculations.
16
Calculations:
Trial 1
Trial 2
Mass of magnesium reacted
Mass of magnesium oxide
produced
Mass of oxygen reacted
Ratio of the mass of magnesium
to the mass of oxygen
Percent Error
(accepted ratio = 1.52)
% Error 
Experimental  Accepted
 100
Accepted
Questions:
1. In the procedure, you were asked to reheat the crucible repeatedly until the last two masses agreed
within 0.03 g. What is the purpose of this reheating?
2. If you were to combine 80.0 g of oxygen with some hydrogen, how much hydrogen would you
need to completely use up all the oxygen? The hydrogen to oxygen ratio is 1:8 by mass. Show
work.
3. Suppose a compound of sodium and chlorine is formed in the ratio of 1.54g of chlorine for each
gram of sodium. How much sodium would you need to completely react 45.0 g of chlorine?
4. How would your results for this experiment be affected if all of the magnesium did not react? (Be
sure to explain the effect on the final ratio)
5. Explain the significance of analyzing data obtained from the combustion of 2 pieces of magnesium
of different lengths.
6. Suppose you tried to combine 42.0 g of magnesium with 45.0 g of oxygen.
a. Which of these two substances would have been left over after the reaction and how much?
(Show work)
b. How much magnesium oxide would be formed? (show work)
17
Name__________________________
Period______Date________________
Thickness of Aluminum Foil
Purpose: To determine how many aluminum atoms make up the thickness of a piece of aluminum foil?
Imagine that there are aluminum atoms stacked on top of each other. How many aluminum atoms would
be required to make up the thickness of a piece of aluminum foil?
Materials:






7 – 10 pieces (shots) of aluminum metal (You MUST use all of them)
Rectangular piece of aluminum foil (No bigger than 10 cm x 12 cm)
10 mL graduated cylinder
Balance
Ruler
Water
Given Information:
D = m/V
V=lxwxh
Required work:




Prediction
Data/Results
Calculations
Conclusion w/source of error
radius of Al atom = 1.43 Å
18
19
Name___________________________
Period______Date_________________
Absolute Zero
Procedure:
1) Wear goggles and apron at all times.
2) Measure the total volume of the Erlenmeyer flask by first
placing the stopper in the top of the flask and drawing a line
indicating the bottom of the stopper with your grease pencil. Fill
the Erlenmeyer flask with water up to the line and pour the water
into a graduated cylinder to measure it.
3) Take the 1000 mL beaker and fill it with approximately 400
mL of water. Place a one-hole stopper fitted with glass
tubing in the top of the Erlenmeyer flask, and place the flask
in the beaker of water. Make sure there is space between the
flask and sides of the beaker to allow room for the steam to
escape. Heat the water to boiling. Continue heating for 3-5
minutes. (see figure A)
4) Measure the temperature of the boiling water and assume this
temperature to be the same as the temperature of the gas in
the flask. Remove the flask from the beaker. Protect your
hand with a towel while placing your finger firmly over the
end of the glass tubing. CAUTION: FLASK IS HOT!
Submerge the flask upside down in the bin of water in the
sink. (Be careful to not allow air to enter the flask while
transferring.)
5) Remove your finger from the glass tubing and hold the flask
under the water (with open end down) until the flask has
cooled and the water ceases to enter. Raise the flask until the
water level outside the flask is equal to the water level inside
the flask. The pressure is now equal to the atmospheric
pressure. (see figure B)
6) Place your finger over the glass tubing while the outside levels are equal. Remove the
flask from the water and place it upright on the lab counter. Measure the volume of the
water in the flask by pouring the water in a graduated cylinder.
7) Measure the temperature of the water bath and assume this to be the gas temperature.
8) Repeat the procedure for Trial 2 and 3 using different water temperatures. (The bins in the other
two sinks.)
20
Analysis:
1) Determine the Gas Volume and Temperature in K for the initial reading and all trials.
2) Prepare a graph of Gas Volume (y) vs. Gas Temp in Kelvin (x)
3) Draw a line of best fit and extrapolate the graph to determine the x-intercept (experimental value
of absolute zero)
4) Determine the percent error.
5) Convert 37oC to Kelvin.
6) Convert 398 K to oC.
7) What is the volume of a gas at 295 K if the same gas has a volume of 42.9 mL at 357 K?
8) A gas has a volume of 50.8 mL at 25oC. What is the volume of the gas at 50oC.
a. Calculate first using Celsius temperatures
b. Calculate again using Kelvin temperatures
c. Which is correct? Why?
21
Precipitation Reactions: Formula Writing
Name___________________________
Period______Date_________________
Procedure
1. Obtain 6 disposable pipettes and label them 1 through 6.
2. Place each pipette upside down in a beaker.
3. Fill each pipette with the solution that corresponds to its number and keep the pipette in its beaker
upside down. Do not empty the contents of the pipette into the beaker
1.
2.
3.
4.
5.
6.

barium chloride
magnesium nitrate
sodium chromate
aluminum sulfate
potassium chromate
silver nitrate
Combine each solution with every other solution by combining 1 drop of each on a transparency.
(Combine solution 1 with each of solutions 2 through 6. After that set is complete, combine
solution 2 with each of the solutions 3 through 6, etc.) Record all combinations that produce a
precipitate.
Data: Record all combinations that produced a precipitate.
Analysis: For each precipitate reaction, write the formula and name for each reactant on the space
provided. Write the names and formulas for the two compounds produced. (This can be determined by
switching the metals in the reactants). Consult a table of solubility rules to determine which product is the
precipitate.
Note: There is space provided for 10 precipitates. There may be less.
Example: A yellow precipitate is produced when solutions of lead(II) nitrate and sodium iodide are
mixed.
The two products (by swapping metals) would be sodium nitrate and lead(II) iodide. According to the
solubility rules, sodium nitrate is soluble. Therefore, lead(II) iodide must be the precipitate.
22
Reacting Chemicals
formula
name
formula
Chemicals produced
name
Pb(NO3)2
lead(II) nitrate
NaNO3
sodium nitrate
no
NaI
sodium iodide
PbI2
lead(II) iodide
yes
PPT?
Example
1
2
3
4
5
6
7
8
9
10
23
Name___________________________
Period______Date_________________
Heat of Fusion of Ice
Procedure:
1. Warm about 125 mL of water to about 50-60oC.
2. Measure 100. mL of this warm water into a Styrofoam cup. Record the volume and temperature.
3. Obtain several ice cubes (6-7 pieces). Shake excess water from them and dry them with a paper
towel. Place the ice in the warm water and stir the mixture until the temperature is about 0oC.
Add more ice if needed to cool the water. Record the lowest temperature reached (Tf).
4. Remove the unmelted ice using crucible tongs or forceps. Be sure to drain back as much water as
possible into the cup when you remove the ice.
5. Measure the volume of water remaining in the calorimeter.
Calculations:
Mass of water cooled
Mass of ice melted
ΔT of water
Moles of ice melted
Heat lost (q) by water
Heat gained by ice (q)
Heat gained per gram of ice
Heat gained per mole of ice
Heat of fusion of ice in KJ/mol
Actual Heat of Fusion:
Percent Error:
6.008 kJ/mol
________________
24
25
Name___________________________
Period______Date_________________
Qualitative Analysis 1: Relative Solubilities
The elements of the second column of the periodic table all form ions with a +2 charge. The
chemistry of these elements is so similar that they are difficult to separate. Many of their compounds are
only slightly soluble; however, it is possible – by choosing the proper anion – to find differences in
solubility, which will permit you to differentiate between the cations of these metals.
In this experiment, you will study the effect of adding reagents containing specific anions to solutions
containing the cations of the metals in the second column. After a systematic study of the relative
solubilities of their carbonates, chromates, sulfates, oxalates, and hydroxides you should be able to make a
qualitative analysis of an unknown solution containing one or more of these cations.
Materials:
Solutions
0.1 M Ba(NO3)2
0.1 M Sr(NO3)2
0.1 M Ca(NO3)2
0.1 M Mg(NO3)2
Reagents
2 M (NH4)2CO3
0.5 M K2CrO4
0.2 M (NH4)2C2O4
1 M (NH4)2SO4
9 M NH3 (aq)
Cations Used
Ba+2 barium
Sr+2 strontium
Ca+2 calcium
Mg+2 magnesium
Anions Used
CO3-2 carbonate
CrO4-2 chromate
C2O4-2 oxalate
SO4-2 sulfate
OH-1 hydroxide
Data: Add 1 drop of each cation and each anion on transparency as shown on the table below.
(NH4)2CO3
CO3 -2
Ba(NO3)2
Ba+2
Sr(NO3)2
Sr+2
Ca(NO3)2
Ca+2
Mg(NO3)2
Mg+2
K2CrO4
CrO4-2
(NH4)2C2O4
C2O4 -2
(NH4)2SO4
SO4-2
NH3
OH-1
26
Analysis of Data:
1. a. Which carbonate of the above positive metal cations has the greatest solubility? _____________
b. Which cations have similar solubilities?_____________________________________
c. Describe how you could separate one of these cations from the other three.
2. a. Which chromate of the metals is the least soluble? _________________________
b. How can this difference in solubility be used in an analytical separation of a solution containing
both Sr+2 and Ba+2?
3. With which of the anions does the magnesium ion have the lowest solubility? _________________
4. Which oxalate of these metals is the most soluble? ___________________________
Flow Chart of Data for Knowns
Mg+2, Ca+2, Ba+2, Sr+2
↓
add CrO4-2
ppt:
soluble:____________________________
↓
add SO4-2
ppt:
soluble:____________________
↓
ppt:
add CO3-2 or C2O4-2
soluble:__________________
↓
add NH3
ppt:
____soluble:________
27
Unknowns Procedure:
1. Obtain your unknown set # from your teacher.
2. Test each of your unknowns with the anions listed in the flow chart below and record your results in
the appropriate spaces
3. On transparency, add 1 drop of unknown with CrO4-2 and 1 drop of unknown with NH3.
4. If precipitate forms with CrO4-2, add 20 drops of unknown with 20 drops of CrO4-2. This is to ensure
that all the barium is removed from the unknown solution. Make sure that a slight yellow color
remains in solution. This indicates that chromate is in excess and all barium is removed.
5. Centrifuge test tube to separate solid from solution.
6. If no precipitate forms on the transparency, move to the next step. Add 20 drops of unknown with 20
drops of SO4-2.
 ANY time a precipitate forms, centrifuge test tube for at least 10 seconds. Remove liquid with a
pipette. The liquid must be clear. Continue to centrifuge until it is clear. Once clear solution is
removed and placed in a new test tube, add 20 drops of the next anion to the solution.
 If no precipitate forms, move on to the next step adding 20 drops of the next anion.
 If a precipitate formed with NH3, it is unnecessary to do the last step of adding 20 drops of NH3.
If no precipitate formed, double check work at the end by adding 20 drops of NH3 to see if a
precipitate forms.
7. Fill in the ions present in both of your unknowns on the slip provided to you by your instructor and
hand in by the end of the period.
The flow chart lists all the possible metal ions that may be present in your unknown at the top. When you
test your unknown with the first anion, CrO4-2, a reaction will occur. Record whether or not a precipitate
was formed as well as a description of it. Then refer to your data table above that contains your knowns
and compare the results. For example, if there is a match with the reaction between Mg+2 and CrO4-2, then
Mg+2 must be present in your unknown. Complete the analysis for both of your unknown solutions.
28
SCHEME of ANALYSIS (to be used for analysis of unknowns) = FLOW CHART
Mg+2, Ca+2, Ba+2, Sr+2
↓
add CrO4-2
ppt:
soluble:____________________________
↓
add SO4-2
ppt:
soluble:____________________
↓
add CO3-2 or C2O4-2
ppt:
soluble:__________________
↓
add NH3
ppt:
____soluble:________
Mg+2, Ca+2, Ba+2, Sr+2
↓
add CrO4-2
ppt:
soluble:____________________________
↓
add SO4-2
ppt:
soluble:____________________
↓
add CO3-2 or C2O4-2
ppt:
soluble:__________________
↓
add NH3
ppt:
____soluble:________
29
The Mole Lab
Name___________________________
Period______Date_________________
Procedure:
1.
2.
3.
4.
Determine the mass of a packet of sugar.
Determine the mass of a scoop of salt.
Determine the mass of a nickel.
Determine the mass of chalk used to write your signature. Write your signature 3 times on the board.
Data and Calculations: You MUST show ALL work to get full credit!
1. Measured mass of sugar (C12H22O11):
______________________
a. Determine the molar mass of sugar:
____________
b. Calculate the number of moles of sugar:
____________
c. Calculate the number of and kind of particles of sugar:
____________
d. Calculate the total number of hydrogen atoms present:
____________
2. Measured mass of table salt:
________________________
a. Determine the molar mass of salt:
____________
b. Calculate the number of moles of salt:
____________
c. Calculate the number of and kind of particles of salt:
____________
30
3. Measured mass of the nickel:
___________________
Assume a nickel to be 25.0% nickel and 75.0% copper.
a. Calculate the mass of nickel in the coin.
____________
b. Determine the molar mass of nickel.
____________
c. Calculate the number of moles of nickel present
____________
d. Calculate the number of and kind of particles of nickel.
____________
e. Calculate the mass of copper in the coin.
____________
f. Determine the molar mass of copper.
____________
g. Calculate the number of moles of copper present
____________
h. Calculate the number of particles of copper. Identify the particle.
____________
4. Measured mass of chalk used in your signature:
_________________________
Chalk is calcium carbonate
a. Formula of calcium carbonate:
____________
b. Determine the molar mass of the chalk
____________
c. Calculate the number of moles of chalk used.
____________
d. Calculate the number of and kind of particles of chalk used
____________
e. Calculate the total number of oxygen atoms present.
____________
31
Name___________________________
Period______Date_________________
Hydrates
Introduction:
Many ionic compounds (salts) have one or more water molecules loosely bonded to it which can be easily
removed by heating to produce the anhydrous salt. For example, the ionic compound copper(II) chloride
is hydrated with two water molecules and is properly named copper(II) chloride dihydrate. The chemical
formula is CuCl2∙2H2O. Mathematically, one can conclude that every mole of copper(II) chloride
dihydrate contains one mole of anhydrous copper(II) chloride and two moles of water.
In this lab, two hydrated salts will be heated to drive off the water molecules to produce the anhydrous
salt. Based on the mass differential, you can determine the mass and moles of water present in the
compound. From these data, the number of water molecules in the formula is to be determined.
Hydrated Salts:
MgSO4∙×H2O
MnSO4∙×H2O
CuSO4∙×H2O
Na2CO3∙×H2O
BaCl2 ∙×H2O
Procedure:
1. Using a ring stand, iron ring, clay triangle and Bunsen burner, heat an empty crucible (no lid)
intensely for two minutes to burn off any impurities on the crucible that may ultimately affect the
mass. For the remainder of the experiment, the crucible is only to be handled with crucible tongs.
Allow the crucible to cool for a few minutes and determine its mass.
2. Add one of the hydrated salts to the crucible until it is approximately one-third full and determine
the combined mass.
3. Heat the crucible intensely for 10 minutes.
4. Allow it to cool for a few minutes and then determine the mass.
5. Heat the crucible again for about 2 minutes. Allow it to cool and remass. If the mass is not within
0.03 grams of the previous mass, it must be heated again until there are two consecutive masses
within 0.03 grams. Remember to use the final mass in all calculations.
6. Repeat the above procedure for the other hydrated salt.
32
Calculations: Show all work and be attentive to significant figures.
Hydrate 1
Hydrate:2
Mass of hydrated salt
Mass of anhydrous salt
Moles of anhydrous salt
Mass of water
Moles of water
Moles of water for
every 1 mol of
anhydrous salt
Chemical Formula of
hydrate
Actual number of moles
of water in compound
Percent error
Questions: Show work for all calculations.
1. Suppose one of the hydrates in this experiment was not heated long enough. What effect would
that have on the determined chemical formula? Explain.
2. Calculate the percent water by mass for magnesium nitrate hexahydrate.
3. Calculate the mass of water in 2.89 g of sodium sulfate decahydrate.
4. Calculate the total number of oxygen atoms in 15.99 g of calcium nitrate tetrahydrate.
5. A sample of a hydrated salt is analyzed and is determined to be composed of 3.097 g of iron,
5.910 g of chlorine and 5.993 g of water. Determine the chemical formula for this hydrate.
33
Name___________________________
Period______Date_________________
Determining the Value of R
Procedure:
1. Obtain a butane lighter and determine the mass.
2. Fill a pneumatic trough and a 500 mL Erlenmeyer flask with very warm water. Be sure that the
flask is completely filled with water. Place a glass plate over the mouth of the flask and carefully
invert the flask into the trough. Place a thermometer in the trough, but do not yet measure the
temperature. Do not submerge entire thermometer.
3. While holding the lighter in the water under the mouth of the flask, depress the gas release button.
You will observe the flask filling with butane as it displaces the water. You must collect enough
butane so that the water level inside the flask (while touching the bottom of the trough) is at or
below the water level in the trough.
4. Shake off any water from the lighter and dry it thoroughly with a paper towel. Allow it to sit for a
few minutes to ensure that it is dry and then reweigh it.
5. Equilibrate the pressure in the flask with the atmosphere. While holding the flask at this position,
slide the glass plate under the mouth and quickly remove the flask from the trough turning it right
side up without losing any water. Keep the glass plate over the mouth and carry the flask to the
fume hood to release the gas.
6. Measure and record the water temperature, barometric pressure, volume of water remaining in the
flask and the volume of water required to completely fill the flask
Results: (work must be shown for bold items)
Volume of gas collected _____________
Mass of butane in flask _____________
Vapor Pressure of water ___________
Partial pressure of butane ____________
Experimental value of R _____________
Actual value of R _____________
Percent Error ________________
Questions:
1. Calculate the volume that your butane sample would occupy at STP.
2. Assuming the same temperature, pressure and R value determined in this experiment, what
volume would the gas have occupied had it been methane instead of butane?
34
35
Name___________________________
Period______Date_________________
Single and Double Replacement Reactions
Single Replacement Procedure:
1) Clean and dry a spot plate.
2) Place 4 pieces of each metal in the spot plate. Place 5 drops of each solution (HCl, CuCl2, MgCl2,
FeCl3) on the metals. Note if a reaction occurs. Dispose of the chemicals according to your
teacher’s instructions. Thoroughly clean the spot plate and use it for the double replacement
reactions.
Double Replacement Procedure:
1) React each solution labeled Reactant 1 with each solution in the Reactant 2 column.
Reactant 1
AgNO3
Reactant 2
Na3PO4
Cu(NO3)2
Na2SO4
Fe(NO3)3
NaOH
NaCl
2) Mix 5 drops of each solution in a clean spot plate. Note if a precipitate forms.
Analysis: Single Replacement
1. Write a balanced chemical equation for each reaction that occurred.
2. Prepare an activity series of the three metals tested by ranking them from most active to least
active.
3. Based on your observations of reactions between the metals and hydrochloric acid, where would
hydrogen fit in your activity series?
4. The Statue of Liberty is made of copper. Use your investigation results to explain why copper is
a better material for a statue than magnesium or iron.
5. Gold does not react with any of the solutions used in this investigation. What does this tell you
about gold’s chemical activity.
6. How does the chemical activity of gold account for its use in jewelry?
36
7. Lead is less active than zinc but more active than copper. Predict the results if lead metal is put
into separate solutions of zinc nitrate and copper (II) nitrate.
Analysis: Double Replacement
1. Write a balanced chemical equation for each reaction that occurred.
2. Refer to a list of solubility rules and indicate if the following types of compounds are typically
soluble or insoluble in water.
a. sodium compounds:______________ b. nitrates_______________
3. In each of the equations written in question #1, underline the formula for the product that was the
precipitate formed.
4. In each of the following, solutions of the two indicated compounds are mixed. Determine if a
precipitate will form. If so, write the formula for the precipitate.
a. tin (II) nitrate and sodium hydroxide
b. sodium nitrate and potassium carbonate
c. sodium bicarbonate and sodium chromate
d. silver acetate and calcium chloride
37
Predicting Products Lab
Name:
Date:
Period:
Pre-lab Discussion:
1. What constitutes a positive test for each of the following gases?
a. oxygen (O2):
b. hydrogen (H2):
c. water vapor (H2O):
d. ammonia (NH3):
e. carbon dioxide (CO2)
and/or
2. What is the proper way to smell a substance in the lab? _____________________________________
3. What is the role of a catalyst in a reaction? How can you tell if when a substance serves as a catalyst?
__________________________________________________________________________________
Procedure:
Read the procedure for each reaction in its entirety before doing it so that you have the necessary
materials to complete the procedure. Wear goggles and an apron throughout the entire lab. Be sure
to tie back long hair. Write detailed observations in the Chart. Follow disposal instructions.
Reaction 1: Cut a 3 cm piece of magnesium ribbon. Clamp the magnesium with a pair of crucible tongs.
Ignite the magnesium in a Bunsen burner flame. Do not look directly at burning Mg. After it
burns, collect the ashes on a watch glass. Disposal: discard ashes in trash can after it has
cooled down.
Reaction 2: Obtain a piece of magnesium. Add 5-10 mL of hydrochloric acid to a test tube and place the
test tube in a rack. Add the magnesium to the acid. Using a test tube clamp, invert a second
test tube (as shown in the figure below) over the mouth of the reaction test tube and collect the
gas being produced. Keep the test tube inverted and test the collected gas by inserting a
flaming splint. Disposal: after all of the Mg has reacted completely with the HCl, the solution
can go down the sink with lots of water.
38
Reaction 3: Mass a scoop of copper powder in a crucible. Place crucible in a clay triangle on ring stand
and heat over a Bunsen burner. Mass the copper once it cools. Disposal: after the copper
cools, discard in the trash can.
Reaction 4: Obtain one scoop of ammonium carbonate and place it into a small, dry test tube. Clamp the
test tube and heat it in a Bunsen burner – use proper technique. While the test tube is heating,
touch a piece of cobalt chloride paper to the mouth of the test tube and remove it. Smell the
test tube using the proper wafting technique. Hold a flaming splint in the test tube. Record
the results of each test. Disposal: any solid remaining in test tube, discard in trash can.
Reaction 5: Measure 10 mL of hydrogen peroxide and add to a test tube. Add a very small quantity (tip
of spatula) of manganese dioxide, MnO2, (catalyst) to the test tube. Place a glowing splint
inside the mouth of the test tube. Hold a piece of cobalt chloride paper to the mouth of the test
tube. Disposal: discard solution in trash can.
Reaction 6: Add 2 drops of potassium iodide and 2 drops of lead (II) nitrate to a watch glass. Record the
results. Disposal: wipe the watch glass with a paper towel and discard the paper towel in
trash can. Do not put it into the sink. Wash hands well with soap and water.
Reaction 7: Place two scoops of copper (II) carbonate in a large test tube. Insert a stopper with a glass
bend in the test tube. Clamp the test tube to a ring stand on an angle so that the flame will
touch the bottom of the test tube only. Fill a small test tube approximately half way with
limewater. Position this test tube so that the end of the glass bend is in the test tube and is
submerged in the limewater. Light the Bunsen burner and heat the solid in the large test tube.
Observe any changes in the limewater and the solid. Disposal: discard any solid in the trash,
the limewater solution may be put down the sink.
Stopper in large
test tube with the
CuCO3
Rubber tubing placed
into the limewater
Please clean your lab area and lab equipment! Wash your hands!!!
39
Questions:
1. Write a balanced equation for every reaction using the test results. Include states in your equations.
Also, indicate the reaction type for each reaction.
2. Write complete, balanced equations for each of the following:
a. When potassium bromate is heated, it decomposes into potassium bromide and a gas that reignites a
glowing splint.

Reaction type:
b.Sodium metal reacts violently with water to produce sodium hydroxide and a gas that pops in the
presence of a flame.

Reaction type:
c. When calcium hydroxide is heated, it makes calcium oxide and a substance that will turn cobalt
chloride paper pink.

Reaction type:
d.When CH4 is burned in the presence of oxygen. It produces a substance that will extinguish a
flaming splint and a substance that turns cobalt chloride paper pink.

Reaction type:
40
41
Name___________________________
Period______Date_________________
+
2+
Qualitative Analysis 2: Ag , Hg2 , and Pb
2+
Background Information:
Most common metal ions form soluble chlorides except Ag+, Hg22+, and Pb2+. Because of this it is
possible to separate these ions from other cations as chloride precipitates. Identification can then be made
on the basis of distinguishing reactions of each chloride. The mercury (I) ion is a bit unusual in that is
occurs in pairs: hence Hg22+ rather than Hg+.
In the first part of this experiment you will become familiar with a few reactions used to identify the
silver, mercurous, and lead ions in aqueous solution. From these observations you should be able to
devise a method by which you can analyze an unknown solution and determine the presence or absence of
Ag+, Hg22+, and Pb2+.
Procedure:
1. Prepare precipitates of AgCl, Hg2Cl2, and PbCl2 in separate, labeled test tubes by adding about 5
drops of 6 M hydrochloric acid, HCl, to about 1 mL of each of the test solutions, AgNO3, Hg2(NO3)2,
and Pb(NO3)2. Make a record of your observations for each test.
2. Allow the chloride precipitates to settle to the bottom of the test tube. Centrifuge. Decant and discard
the solution. Add about 2 mL of distilled water to each precipitate. Stir/shake the tests tubes to allow
the precipitate to dissolve (not all of the precipitate will dissolve).
3. Centrifuge. Remove about 5 drops of the clear solution from each solution and transfer the drops to
clean test tubes. Add about 5 drops of 0.1 M potassium chromate solution, K2CrO4, to each of the 3
test tubes. Determine and record which of the cations forms a precipitate with the chromate ion.
4. Decant and discard the solutions from the test tubes in step number 2. Save the precipitate for the next
step.
5. To test the solubility of the chloride precipitates in aqueous ammonia, add about 3 mL of 6M
ammonia solution, NH3 (aq), to each of the precipitates.
6. Shake the test tubes to dissolve as much of the solid chloride as possible. For each, decant about 1
mL of clear solution into a clean test tube. Add about 2-3 mL of 6 M nitric acid, HNO3, to each tube
of decanted solution. Which chloride re-precipitated?
7. Examine your record and complete the flow chart which summarizes the steps you would use to
analyze and unknown containing all three cations.
8. Analyze the unknowns starting with step #1. (Note: Make sure to react only 1 mL of your unknown
with the HCl.)
42
Data Table:
Record your observations in the following flow charts as you complete the procedure. The flow charts are
designed to correspond to the procedural steps.
Pb+2 Pb(NO3)2
Ag+1 AgNO3
Hg2+2 Hg2(NO3)2
6 M HCl
6 M HCl
6 M HCl
H2O
heat
H2O
heat
H2O
heat
Ppt
Soln
0.1 M
K2CrO4
Ppt
Soln
0.1 M
K2CrO4
6M
NH3 (aq)
Soln
6 M HNO3
Ppt
Soln
0.1 M
K2CrO4
6M
NH3 (aq)
Soln
6 M HNO3
6M
NH3 (aq)
Soln
6 M HNO3
43
Data Table for Unknown # _______
Unknown # ___________
(The unknown could contain Ag+1, Pb+2, and/or Hg2+2)
1 mL of your unknown reacts with 6M HCl
To 5 drops of the clear solution,
add 5 drops of K2CrO4
To the precipitate, add 3 mL of NH3
To less than 1 mL of the clear solution,
add 2-3 mL of HNO3
44
Data Table for Unknown # _______
Unknown # ___________
(The unknown could contain Ag+1, Pb+2, and/or Hg2+2)
1 mL of your unknown reacts with 6M HCl
To 5 drops of the clear solution,
add 5 drops of K2CrO4
To the precipitate, add 3 mL of NH3
To less than 1 mL of the clear solution,
add 2-3 mL of HNO3
45
Name___________________________
Period______Date_________________
The Molar Volume of a Gas
Procedure:
1. Measure the length of a piece of magnesium ribbon no more than 2.5 cm. Cut the Mg square and
measure accurately. Using the grams per meter conversion provided, calculate the mass of the
magnesium. Return the Mg to its container immediately after cutting. Do not set aside on table.
2. Cut and wrap a piece of copper wire (approximately 15 cm) to form a “cage” around the
magnesium. Thread one end of the copper wire through the hole in the rubber stopper so that the
magnesium is suspended from the narrower end of the stopper.
3. Using a funnel, pour approximately 10 mL of 6.0 M HCl into the gas measuring tube.
4. Using a beaker or the distilled wash bottle, slowly add distilled water into the tube. Hold the
eudiometer at an angle while filling to minimize the amount of water that mixes with the acid. Be
sure to fill the tube to the very top so that there is no air in the system.
5. Place the stopper with the magnesium in the copper cage into the eudiometer tube so that the
magnesium ribbon is inside the tube. Fill the holes of the stopper with distilled water.
6. Invert the tube into a 600 mL beaker and clamp the tube to a ring stand. The eudiometer should
not sit on the bottom of the beaker. Gently tap the eudiometer to release any trapped gas.
7. Move the eudiometer from the beaker to the graduated cylinder.
8. When the reaction stops, adjust the tube so that the water level in the tube is the same as the level
in the large graduated cylinder and record the gas volume at the point. Equilibrate the pressure.
9. Measure the temperature of the water (assume this to be the gas temp) and the barometric
pressure.
10. Repeat the procedure for a second trial.
46
Note: Calculations using R are not acceptable for this experiment. The value of R is derived from
the actual molar volume. Since determining the molar volume is the purpose of this experiment,
using the value of R is not mathematically appropriate.
Analysis: Work must be shown for Trial 1 only.
1. Write a complete balanced equation for the reaction that took place. Include all physical states.
2. Determine the mass of magnesium ribbon used.
Trial 1:____________
Trial 2:____________
3. Determine the number of moles of magnesium reacted.
Trial 1:____________
Trial 2:____________
4. Determine the number of moles of hydrogen produced.
Trial 1:____________
Trial 2:____________
5. Determine the partial pressure of hydrogen for each trial.
Trial 1:____________
Trial 2:____________
6. Given the volume, temperature and pressure of the hydrogen collected, determine the volume (L)
of hydrogen, for each trial, at STP.
Trial 1:____________
Trial 2:____________
7. Determine what the volume of hydrogen gas produced at STP would be for every 1 mole of
hydrogen.
Trial 1:____________
Trial 2:____________
8. Average these values to get your experimental value for the molar volume.
Molar Volume (experimental): _________________
9. Determine your percent error for this experiment.
Percent Error: __________________
Problems:
1. What would the volume of hydrogen collected in trial 1 of your experiment be at 115.0 kPa and
87.5oC?
2. Using your experimental value for the molar volume (#8), calculate the volume of hydrogen that
would be produced at STP if a 10.0 g piece of magnesium is reacted completely.
47
Name___________________________
Period______Date_________________
Quantitative Study of Chemical Reactions
Equations are balanced in accordance with the law of conservation of mass-energy. In other words, the
number of atoms of each element in the reactants must equal the number of atoms for those same
elements in the products.
Describing a chemical reaction in terms of the number of atoms and molecules involved is not always
practical. Chemists are forced to work with large quantities of atoms and molecules, which are measured
in grams, kg, and moles.
Objectives: In the following experiment, you will…
 React copper and silver nitrate to obtain silver and copper(II) nitrate
 Measures the masses of the reactants and products
 Determine the coefficients for the balanced equation of the reaction based on your mass data.
 Calculate the percent reclaimed silver following additional reactions.
Your success in this experiment depends heavily upon your careful lab technique, your clear labeling of
all beakers and your use of the same balance throughout this experiment.
Part A: Reaction of silver nitrate and copper
Balanced Equation: _________________________________________________________________
1. Obtain a vial of silver nitrate and determine its mass.
2. Empty the contents into a large test tube. Fill the test tube approximately one-half full with
distilled water. Stir the solution until all is dissolved. Support the test tube in a beaker or a test
tube rack.
3. Replace the cap on the vial and determine its mass empty.
4. Obtain a copper wire (25 – 30 cm) and determine its mass.
5. Coil the copper around a pencil. The copper will be suspended into the silver nitrate solution.
Compress the coils so that all of the coils will be submerged in the solution.
6. Insert the copper wire into the test tube. Bend the end of the copper wire and hook it to the rim of
the test tube. The copper will need to remain in the test tube for at least 30 minutes. After 15
minutes, gently agitate the copper wire to shake off some of the silver being produced to expose
more of the copper to the silver nitrate. If the test tube is going to be left over night to react, be
sure to agitate the copper wire before leaving the lab.
7. Check to see if the reaction is complete by first shaking the copper wire to remove the silver and
then by observing the wire for a minute to see if any more silver forms. If the reaction is
complete, proceed to the next step. If not, wait a few more minutes and repeat this procedure.
48
8. Obtain a clean, dry 100 mL beaker and label it with your initials and period number. Determine
the mass of the beaker.
9. While carefully removing the copper wire from the test tube, rinse it using a wash bottle into the
test tube. Be sure that all of the silver is removed from the copper and is in the test tube. Take the
copper wire to the front desk and dip it in the bottle of acetone (facilitates drying). Place the
copper on a paper towel at your station and allow it to dry. Once dry, determine its mass.
10. Pour the contents of the test tube into the preweighed 100 mL beaker. Use a wash bottle to flush
all of the silver from the test tube into the beaker.
11. Allow the silver to settle into a beaker. Carefully decant the solution into a larger beaker leaving
the silver behind. (Decanting into another beaker will allow you to recover any silver that may
accidentally be poured out.)
12. Wash the silver 3 times with 15 mL (approximately) portions of distilled water. Each time, allow
the silver to settle and decant the solution into the larger beaker. It is okay to leave some liquid in
the beaker. On the third washing, decant a small portion of the wash solution onto a watch glass.
Add 1 drop of the 6M NH3 solution. If a blue color is present, there are still copper ions in the
solutions and additional washings will be necessary. If it is colorless, there is no longer any
copper solution left and you can proceed to the next step.
13. Place your beaker of silver in the fume hood under the heat lamps to dry. Once dry, determine the
mass of the beaker with the silver.
Part B: Reaction of silver and nitric acid to produce silver nitrate, nitrogen dioxide and water
Balanced Equation: _________________________________________________________________
1. This procedure will be performed by your teacher. After you have measured the mass of the dry
silver crystals, 10 mL of 6M nitric acid will be added to the beaker in the fume hood. A reddishbrown gas, nitrogen dioxide, which is toxic will be formed and the crystals will dissolve. The
beaker will then be left to dry under the heat lamps.
2. Determine the mass of the beaker containing silver nitrate.
Part C: Reaction of silver nitrate and sodium chloride
Balanced Equation: _________________________________________________________________
1. Add 25 mL of distilled water to the beaker of solid silver nitrate. Stir until the silver nitrate
dissolves.
2. Mass out 2 grams of sodium chloride and add it to the silver nitrate solution. A white precipitate
forms.
3. Place the beaker on a hot plate or ring stand and heat it until it begins to boil (do not allow the
solution to boil continuously). Slowly heat the solution to boiling or until the liquid becomes clear
and the precipitate has settled to the bottom of the beaker. While heating, the solution can be
gently stirred occasionally.
4. Determine the mass of a piece of filter paper.
49
5. Fold the filter paper and place it in a long stem funnel. Support the funnel in an Erlenmeyer flask.
6. Using beaker tongs, slowly pour the mixture into the filter apparatus to filter the precipitate. Use a
wash bottle to rinse any remaining precipitate onto the filter paper.
7. Wash the precipitate with distilled water and allow it to drain. Squirt a liberal amount of acetone
on the precipitate and allow it to drain.
8. Using forceps, carefully remove the filter paper from the funnel and lay it flat on a watch glass.
Squirt more acetone on the precipitate and filter paper. Place the watch glass in the hood under
the heat lamp to dry. Put a piece of scrap paper with your name on it under the watch glass.
9. When dry, determine the mass of the filter paper with the precipitate.
Quantitative Study of Chemical Reactions - Report Sheet
Part A:
1. Write a complete balanced equation for the reaction that took place in part A. Assume the copper
ion to be +2. Include physical states of all reactants and products.
2. Using the mass of silver nitrate used, calculate the mass of copper metal that should have reacted.
3. Calculate the percent error for the reacted copper.
4. Calculate the theoretical yield of silver for this reaction.
5. Calculate the percent yield of silver.
Part B:
1. Write a complete balanced equation for the reaction that occurred in part B (see procedure for the
products). Include physical states of all reactants and products.
2. Using the mass of silver recovered from part A, calculate the theoretical yield of silver nitrate for
this reaction.
3. Calculate the percent yield of silver nitrate.
Part C:
1. Write a complete balanced equation for the reaction that occurred in Part C. Include physical
states of all reactants and products.
2. Using the mass of silver nitrate recovered in part B, calculate the theoretical yield of silver
chloride.
3. Calculate the percent yield of silver chloride.
50
Questions:
1. Calculate the mass of silver in the silver nitrate that was initially reacted in part A.
2. Calculate the mass of silver in the silver chloride that was recovered in part C.
3. Calculate the percentage of silver that remained at the end of the experiment.
4. List one source of error that would cause there to be less than 100% of the silver at the end of the
experiment (answers such as “human error” and “inaccurate measurements” are not acceptable).
5. List one source of error that would that would cause there to be more than 100% of the silver at
the end of the experiment (same conditions apply).
6. A student repeats this entire experiment and react 6.89 g of copper metal. Assuming a 100%
yield, calculate the mass of silver chloride that is recovered at the end of the experiment.
51
Name___________________________
Period______Date_________________
Titration 1
Objective: To determine the molar concentration of a hydrochloric acid solution by titrating it with a
sodium hydroxide solution with a molarity of _______M.
General Titration Instructions
Buret Preparation:
1. Obtain the buret(s) and clamp it (them) to a buret clamp attached to a ring stand.
2. Open the spout of the buret. When the spout is parallel to the tip, it is open.
3. Add distilled water to the buret and allow the water to flow through the tip. This is to both rinse
the buret and to make sure that it is working properly.
4. Position the buret so that the top of it is below eye level.
5. Using a long stem funnel, add 5 – 10 mL of the solution to be added to the buret and allow that to
flow out.
6. Close the spout and fill the buret with the solution. It is not necessary to fill it exactly to the zero
mark. Open the spout briefly to allow some of the solution to flow out of the buret to insure that
there is no air in the tip.
7. If using two burets, be sure to label them.
Titration:
1. Record the initial buret readings.
2. Deliver a known volume (14-20mL) of one solution (the acid for this lab) to an Erlenmeyer flask
and add two drops of the indicator, phenolphthalein.
3. Open the spout of the other solution (base) and begin to deliver this solution to the flask while
constantly swirling it.
4. As the endpoint nears, you will notice the color changing briefly and then returning to the original
color. The closer the endpoint is, the longer the color remains. When close, the spout can be
repositioned so that the solution is delivered dropwise. Continue adding the solution until the
color change is permanent. The color must change due to the addition of a single drop. If more
than one drop was added, add a small amount of the other solution to return it to the original color
and repeat.
5. Once finished, record the final buret readings.
6. A titration should be done at least three times.
Clean-Up:
The burets are to be emptied of the solutions. They should be cleaned with a dilute soap solution and
thoroughly rinsed. Tap water can be used for this process. Once cleaned, the burets should be rinsed
with distilled water and stored upside down with the spouts open.
Analysis: Write a complete balanced equation for the reaction. Show your work for each molarity of
acid calculation. Then determine the average molarity of your trials.
52
Questions:
1. Obtain the actual molarity of the HCl solution from your teacher and calculate your percent error.
2. What happens in the titration during the color change of the indicator? Why is it important that
the color change occur by the addition of a single drop?
3. Why is an Erlenmeyer flask preferred over a beaker when performing a titration?
4. Suppose calcium hydroxide was used in this experiment instead of sodium hydroxide. Assuming
all experimental data was the same, what would be the molarity of the acid? Explain using
calculations.
5. Suppose there was a large air bubble in the tip of the acid buret. Explain the effect that this would
have on the calculated value of the acid’s molarity. Use calculations to show work.
6. A student titrated 30.00 mL of a hydrochloric acid solution with 1.00 M sodium hydroxide. The
sodium hydroxide was initially filled to the 2.30 mL mark on the buret. Once the level in the buret
reached 50 mL, the student was forced to refill the buret and did so to the 5.06 mL mark. The
endpoint was finally reached when the sodium hydroxide level dropped to 32.09 mL on the buret.
What is the molarity of the acid solution? Show work.
7. How many milliliters of a 0.250 M sulfuric acid solution would be required to completely react
with 2.398 grams of solid aluminum hydroxide? Work must be shown.
53
Name___________________________
Period______Date_________________
Titration 2









Determine the molarity of the HCl solution by titrating it with sodium carbonate.
The indicator to use is bromocresol green (approximately 6 drops per trial).
A minimum of 3 trials is required.
The unknown letter for the HCl must be included.
Accuracy counts.
Prepare sodium carbonate solution using a mass of sodium carbonate between 7 and 9 grams. Use
a 250.0 mL volumetric flask.
Note: The sodium carbonate used in this experiment may be anhydrous or it
may be a monohydrate. Refer to the label on the jar to determine which it is because it must
be considered in the molar mass determination. This must be included in the report!
Start with 20-25 mL of the base in the Erlenmeyer flask. Vary the amount for each trial.
ALL data for sodium carbonate molarity and HCl titrations must be included in lab report.
Questions/Problems to be included in report. The answers need to be typed in the appropriate section of
your lab report. The work must be shown in the sample calculations.
1. What is the minimum mass of sodium bicarbonate that must be sprinkled on 115.9 mL of the
hydrochloric acid solution (your average molarity) that you titrated in order to completely
neutralize it?
2. What volume of your sodium carbonate solution would need to be combined with 250.0 mL of
3.00 M nitric acid to completely neutralize it?
3. A 0.5425 g sample of an unknown monoprotic acid requires 12.97 mL of a 0.1025 molar solution
of calcium hydroxide for neutralization in a titration. Determine the molar mass of the unknown
acid.
54
55
Name___________________________
Period______Date_________________
Determining Solution Concentration: Using a Spectrophotometer
Introduction
The spectrophotometer is a powerful tool which can be used for colorimetric determination of
concentration. The process is based on the fact that colored ions absorb light from the visible spectrum.
The greater the concentration of the ions, the greater the absorbance (A) of light.
Conversely, the more light that is absorbed by the ions in solution, the less light that is transmitted
through the solution. Thus, the inverse of absorbance is percent transmittance (%T). Therefore,
concentration can be measured using absorbance or percent transmittance.
To determine the concentration of a colored ion in solution, a set of carefully prepared solutions of
known concentration must be first measured for absorbance or percent transmittance. The
spectrophotometer should have the wavelength setting at a previously determined maximum absorbance
(or %T) value for the ion in question while testing the standards. A graph plotting the absorbance (or %T)
against concentration data for the standards gives a calibration curve. Then, the absorbance (or %T) can
be measured for a solution of unknown concentration, then matched with the calibration curve to
determine the concentration.
The cobalt ion, Co2+, concentration will be determined in this experiment and should be expressed
in units of molarity. The absorption maximum (%T minimum) for Co2+ occurs at a wavelength setting of
510 nm. The percent transmittance, %T, scale will be used to produce a calibration curve.
Procedure
1. Make 50.0 mL of a 0.10 M Co(NO3)2 · 6 H2O standard stock solution. Show your calculations for
this step.
2. Using the stock solution and a small graduated cylinder, prepare the solutions indicated below in test
tubes labeled 1 – 5. Each solution should have a total volume of 10. mL.
Test Tube #1:
Test Tube #2:
Test Tube #3:
Test Tube #4:
Test Tube #5:
0.02 M
0.04 M
0.06 M
0.08 M
0.1 M
Show work for the dilution calculations for test tubes 1 – 4
56
3. Obtain 1 disposable test tube, your 5 prepared solutions, a wash bottle and a paper towel and proceed
to a spectrophotometer in the room. It is convenient to stand your prepared solutions in a beaker.
4. The spectrophotometer should be warmed up and the wavelength set to 510 nm. The device needs to
be zeroed. With the sample compartment empty and closed, adjust the %T (transmittance) to 0 with
the left front knob.
5. Place a disposable test tube filled with distilled water, into the sample compartment and adjust the
%T to 100% T with the right front knob (this means all of the light travels through the totally clear
water). Be sure to wipe all liquid and smudges off of the tubes before placing in
spectrophotometer! In addition, no bubbles should be evident int eh test tube.
6. Remove the distilled water tube and empty it. Fill it approximately ¾ full with the solution in Test Tube
#1. Wipe the outside of the test tube and insert it into the compartment. Record the absorbance. Repeat
for Test Tubes 2–5, but be sure to rinse the disposable test tube with distilled water, then with the next
prepared solution between uses. Start with the lowest concentration and work up to the highest
concentration.
7. Construct a calibration graph plotting absorbance on the y-axis and molarity on the x-axis
8. Once the graph is complete, determine the absorbance for unknown A and unknown B. It will be
necessary to recalibrate the machine before doing this by following procedure 4. Using your graph,
determine the molarities for the two unknowns.
Conclusions
1. The actual concentrations of the unknown solutions were A = 0.075 M and B = 0.050 M. Calculate the
percent error of your results for A and B.
2. Using your calibration graph, determine the absorbance of a 0.036 M and 0.072 M solution of Co2+.
57
Periodic Law
Name: _________________________
1. Use the regular periodic table to draw or highlight the zigzag line (staircase).
2. Place the atomic numbers in the lower left hand corner and the atomic masses to the tenths in the
upper right hand corner.
3. The following sets of elements appear as GROUPS: O, G, CC, II, QQ ; D, KK, X, L, GG ;
U, J, Z, S, K, HH ;
A, PP, I, Y, AA ;
T, DD, FF, LL, E ;
R, V, EE, F, N, B ;
C, OO, M, W, JJ ;
MM, H, P, Q, BB ;
_________________________________________________________________________
1) EE is an alkali metal.
2) Y has an outer electron configuration of 4s2.
3) U is a noble gas.
4) T has an atomic mass of 32 amu.
5) B is a gas at room temperature.
6) I has the smallest atomic radius in its group.
7) K has an electron configuration that ends in 6p6.
8) C is an element that is used in fertilizer.
9) R has a lower first ionization energy value than N and V but a higher value than F.
10) Y is the metallic component of limestone.
11) D is the most electronegative element in the group.
12) LL is a metalloid.
13) OO has the smallest electron affinity of the group.
14) KK is a halogen that sublimates.
15) Q has the most energy levels in the group.
16) CC has one more proton than I.
17) V has an outer subshell configuration of 3s1.
18) J has a smaller radius than Z, but a larger radius than S.
19) L is a dark-red liquid.
20) G has a larger electronegativity than QQ but a smaller electronegativity than O.
21) H has two allotropic forms called diamond and graphite.
22) F is a liquid at room temperature.
23) A has a larger first ionization energy than PP and AA.
24) FF is the most abundant element in the earth’s crust.
25) II is the least metallic metal.
26) C is especially stable because of a half-filled “p” subshell.
27) Both P and BB are metalloids but BB has a lower ionization energy.
28) AA has the largest atomic size in the set.
29) Z has a total of 18 protons.
30) W is the most abundant element in the atmosphere.
31) N is the first alkali metal with a completed 3p.
32) HH has one more proton than L.
33) E is the most metallic in this group.
34) Both JJ and M are metalloids, but M has more metallic properties.
35) O has a greater first ionization energy than G, but a lower first ionization energy than II.
36) MM comes from the Latin word stannous.
37) PP has a lower electronegativity than Y and A, but a higher electronegativity than AA.
38) GG is radioactive.
39) DD has two half-filled “p” orbitals in the fourth energy level.
40) C has one less electron than T.
58
Name___________________________
Period______Date_________________
1A
8A
1
2A
2
3
4
5
6
3A
4A
5A
6A
7A
59
QUESTIONS:
1. Which has a larger radius?
R or F
C or X
2. Which has the highest electronegativity?
EE or D
X or KK
3. Which has the lowest ionization energy?
B or F
V or A
4. Which has a high 1st I.E. due to a filled subshell?
J or W
V or A
5. Which probably has the highest electron affinity?
D or KK
N or L
6. Identify the make-believe element that corresponds to:
_______a)
Atomic number 12
_______b)
Atomic mass 84
_______c)
3s23p3
_______d)
4s24p1
_______e)
10 electrons
7. Which make believe element is:
______a)
the most metallic in group 2A
______b)
the most metallic in period 1
______c)
the most nonmetallic in group 3A
______d)
the most nonmetallic in period 3
8. What charge are you most likely going to find on the following ions?:
N _____ W ______ KK ______ EE ______ Q ______ T ______
9. Identify the following formulas as either CORRECT (C) OR INCORRECT (I):
AX2 ____
QT2 ______
VW3 ______
RX _____
10. Finish writing and balance the following double replacement equations:
a)
ID2
b)
RX
+ (EE)3C ------->
+ Y(LL) ------->
11. Write balanced composition reaction equations for the following:
a)
b)
R + X2 ----->
V + LL ------>
N3M ______
60
61
Building Molecular Models
molecular
formula
H2
HBr
PF3
CH4
N2
Lewis Structure
(Dot Diagram)
Name__________________________
Structural Formula
(drawn to shape)
Molecular
Geometry
around Central
Atom(s)
Bond
Angles(s)
Polar?
Teacher's
Initials
62
molecular
formula
CH3NH2
CO2
H2CO3
C2H2
CH3Cl
Lewis Structure
(Dot Diagram)
Structural Formula
(drawn to shape)
Molecular
Geometry
around Central
Atom(s)
Bond
Angles(s)
Polar?
Teacher's
Initials
63
molecular
formula
HCOOH
HCN
H 2O 2
CH2Cl2
H2CO
Lewis Structure
(Dot Diagram)
Structural Formula
(drawn to shape)
Molecular
Geometry
around Central
Atom(s)
Bond
Angles(s)
Polar?
Teacher's
Initials
64
molecular
formula
C2H4
O2
NF3
CH3CH2CH3
CH2CHCH3
Lewis Structure
(Dot Diagram)
Structural Formula
(drawn to shape)
Molecular
Geometry
around Central
Atom(s)
Bond
Angles(s)
Polar?
Teacher's
Initials
65
molecular
formula
CH3OH
SCl2
CF4
Cl2CO
ClNO
Lewis Structure
(Dot Diagram)
Structural Formula
(drawn to shape)
Molecular
Geometry
around Central
Atom(s)
Bond
Angles(s)
Polar?
Teacher's
Initials
66
67
Name___________________________
Period______Date_________________
Solutions Lab
1. Carefully weigh out 2.2 grams of anhydrous sodium sulfate and add it to a test tube. Using a
graduated cylinder, measure and add 10.0 mL of distilled water to the test tube. Stopper the test
tube and shake.
a. Describe the contents of the test tube:
b. Is the solution saturated, unsaturated or supersaturated?
c. How do you know?
2. Remove the stopper from the test tube. Clamp the test tube and gently heat it in a Bunsen burner.
While heating, stir the solution frequently with a stirring rod until all of the solute is dissolved.
Add a very small amount of solid solute to the test tube and stir.
a. Describe the contents of the test tube:
b. What happened to the added solute?
c. Is the solution saturated, unsaturated or supersaturated?
d. How do you know?
68
3. Place the test tube in a beaker of ice water for five minutes. Do not disturb the solution during this
process. Gently remove the test tube and place it in a test tube rack. Add a very small amount of
solid solute to the test tube. Do not stir or shake the tube. Observe the test tube for several
minutes.
a. Describe the contents of the test tube before the extra solute was added.
b. What happened when the extra solute was added?
c. Describe the contents of the test tube after the extra solute was added.
d. Was the solution before adding the crystal saturated, unsaturated or supersaturated?
e. How do you know?
f. Was the solution after adding the crystal saturated, unsaturated or supersaturated?
g. How do you know?
4. Explain one simple test that will determine whether a solution is saturate, unsaturated or
supersaturated. Explain how to interpret the test result.
69
Name______________________
Period_______Date___________
Heat of Reaction
Procedure:
1. Measure 50 mL of the hydrochloric acid and pour it into the calorimeter.
2. Measure and record the temperature of the acid solution.
3. Weigh approximately 2g of sodium hydroxide pellets. DO NOT TOUCH THE PELLETS!
Record the actual mass.
4. Carefully place the pellets into the acid and insert the thermometer.
5. Gently swirl the cup to mix the contents. Record the highest temperature reached for
mixture.
the
6. Repeat the procedure using 4.0g of NaOH.
***Assume the density and specific heat of the solution are the same as water’s values.***
Analysis:
1. Write a complete balanced equation for the reaction that occurred.
2. Calculate the heat of reaction (∆H) in kJ per mole of NaOH for each trial using the temperature
data. Work must be shown.
∆H (trial 1) = ____________________
∆H (trial 2) = ____________________
3. Determine the average value of ∆H for your two trials. This is your experimental value.
∆H (average) = ____________________
4. Determine the ∆H for the reaction using the heats of formation. This is your accepted value.
Work must be shown.
5. Calculate percent error.
70
71
Name___________________________
Period______Date_________________
Determining the Heat of Formation of Magnesium Oxide
Procedure:
1. Measure 100 mL of 1.0 M HCl (assume the mass to be 100 g) to a graduated cylinder, pour it into
a Styrofoam cup and measure its temperature.
2. Cut a 15 cm piece of magnesium ribbon and determine its mass.
3. Add the magnesium to the acid solution. Record the highest temperature reached.
4. Clean the Styrofoam cup and repeat the above procedure using 0.50 g of magnesium oxide instead
of magnesium.
Data and Calculations:
Reaction 1
Complete Balanced
Equation
Mg +
HCl →
Reaction 2
MgO
+
HCl →
Mass of HCl
Mass Solid Added
Mass of Solution
Initial Temperature
Final Temperature
Δt
Show Work
Show Work
Show Work
Show Work
qsolution
(assume the
solution has the
same specific
heat as water)
Calculate moles of
solid added
72
Show Work
Show Work
Calculate Enthalpy
Change of the
Reaction in KJ/mol
Analysis:
1. Complete the first two reactions and transfer the ΔHrxn values for each from the data table.
Reaction 1:
Mg
+
HCl →
ΔHrxn = _________KJ
Reaction 2:
MgO +
HCl →
ΔHrxn = _________KJ
Reaction 3:
H2
+ ½ O2 → H2O(l)
ΔHrxn =
-285.8 KJ
2. Apply Hess’s Law to the above 3 equations and determine the ΔHrxn for:
Mg + ½ O2 → MgO
(note: ΔHrxn for this reaction is the heat of formation)
3. What is the published value (accepted value) for the heat of formation of magnesium oxide?
4. Determine the percent error for this experiment.
% Error 
observed  accepted
 100
accepted
5. Write the thermochemical equation for reaction 1.
6. Draw the enthalpy diagram for reaction 2.
7. If reaction 3 was performed independently of the other two and liberated 1593.64 KJ of thermal
energy, how many liters of oxygen at STP would have been consumed?
73
Name___________________________
Period______Date_________________
Stoichiometry
Procedure:
1. Weigh approximately 0.3 g of sodium carbonate monohydrate and 0.3g of copper(II) chloride
dehydrate and place them in separate beakers. Record the actual mass of each reactant. Add 20
mL of distilled water to each beaker. Stir each until dissolved.
2. Reaction 1: Combine the contents of the two beakers and filter the precipitate formed. Rinse the
precipitate several times with distilled water. Rinse beaker to ensure all precipitate is in funnel.
3. Reaction 2: Using forceps, carefully remove and fold the filter paper and place it in a preweighed
crucible (no lid). Gently heat the crucible with a Bunsen burner until the filter paper is charred.
Then, heat the crucible intensely for 15 minutes. Allow the crucible to cool and determine its
mass.
Questions: Work must be shown for all calculations. Be attentive to units and significant figures in
answers.
1. Write a complete balanced equation for the first reaction that occurred. Include all physical states
for all reactants and products.
2. Determine the limiting reagent.
3. Calculate the theoretical mass of the precipitate formed.
4. Calculate the theoretical mass of the product that remained dissolved.
5. Write a complete balanced equation for the second reaction that occurred. Include all physical
states and write the color of any solids under its formula.
6. Identify the product that remained in the crucible after the reaction and calculate its theoretical
mass.
7. Calculate the percent yield for this reaction.
8. Given the percent yield of the reaction, calculate the number of molecules of the other product that
was actually produced in this reaction.
74
75
Name___________________________
Period______Date_________________
Moles of Iron and Copper
1. Mass a beaker. Add approximately 2g of copper(II) chloride and dissolve into15 mL of distilled
water.
2. Weigh 2 clean iron nails and add them to the solution. Allow the iron to react for 20 minutes.
3. Rinse the nails into the beaker so that the maximum amount of produced copper is in the beaker.
4. Dry and weigh the nails to determine the mass of reacted iron.
5. Decant and rinse (multiple times) and dry the copper to determine the mass of copper produced.
Data and Observations: create a data table with all measurements and observations taken in the lab.
Analysis: Determine if the iron ions in solution are iron(II) or iron(III). Proof is required. Show all
possible balanced chemical equations and calculations.
Questions:
1. Write a balanced chemical equation for the reaction (according to your results).
2. Determine the molarity of the iron (?) chloride solution after the reaction was ended.
3. Determine the mass of iron(?) chloride that would be produced if 3.87 g of CuCl2 reacted with
excess iron.