Download AP Chemistry Lab Manual

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AP Chemistry Lab Manual
Lab Notebook Guidelines ......................................................................................................... 2
QRS lab ..................................................................................................................................... 5
Qualitative Analysis of the Group III Cations .......................................................................... 6
How Much Zinc is in a penny? ............................................................................................... 11
Predicting Products of Chemical Reactions ............................................................................ 12
Redox Titration: The Standardization of Potassium Permanganate ...................................... 13
The EMF Activity Series ........................................................................................................ 15
Heat of Fusion for Ice ............................................................................................................. 19
Additivity of Heats of Reaction: Hess’s Law ........................................................................ 21
Heat of Combustion of a metal-an inquiry based approach .................................................... 23
VSEPR and Molecular Geometry ........................................................................................... 25
Formation of a Coordination Complex of Copper (II) ............................................................ 26
Kinetics of a Reaction -- An Iodine Clock............................................................................. 28
Where did the Crystal Violet go?............................................................................................ 30
Chemical Equilibrium: Finding a Constant, Keq or Kc .......................................................... 35
Entropy of a Reaction ............................................................................................................. 37
Catalytic Converter—Hot Copper Catalysis ........................................................................... 38
Equilibrium and Le Châtelier's Principle ................................................................................ 40
Strong Acid Strong Base Titration .......................................................................................... 43
Titration of a weak acid .......................................................................................................... 45
Determination of the Ka of Weak Acids................................................................................. 46
Determination of the Ksp of an Ionic Compound ................................................................... 49
Buffer Laptop Palooza ............................................................................................................ 51
Preparation and Properties of Buffer Solutions ...................................................................... 53
Corrosion Cells ....................................................................................................................... 56
Polyatomic Ions ...................................................................................................................... 59
Molecular Geometry ............................................................................................................... 60
Rules of Writing Equations ..................................................................................................... 62
AP Chemistry Syllabus ........................................................................................................... 64
Class Rules .............................................................................................................................. 66
Description of Content Covered ............................................................................................. 67
End of Year Review ................................................................................................................ 70
Solution Practice ..................................................................................................................... 79
Redox Practice ........................................................................................................................ 80
Thermochemistry: Standard Heats of Formation Worksheet ................................................. 82
Gas laws practice. ................................................................................................................... 83
8 and 9 Practice worksheet ..................................................................................................... 85
Equilibrium and Entropy Practice ........................................................................................... 89
Ch. 10 questions ...................................................................................................................... 92
Chapter 11 Practice ................................................................................................................. 92
Kinetics part I.......................................................................................................................... 94
Kinetics part II ........................................................................................................................ 97
Accessing Prior Knowledge Acids and Bases ...................................................................... 100
pH PRACTICE ..................................................................................................................... 101
Ka and Kb practice................................................................................................................ 102
Titration Curve Practice ........................................................................................................ 103
Chapter 14 and 15 practice: .................................................................................................. 110
Ksp practice. Keep me but put all answers in your notes. ................................................... 112
Electrochem practice ............................................................................................................. 114
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Lab Notebook Guidelines
You must have a composition notebook or the notebook from Chem I as a lab notebook.
A lab notebook should be used to explain lab procedures, record all lab data, and show how
calculations are made. You may also use the notebook to discuss the results of an
experiment and to explain the theories involved.
A record of lab work is an important document which will show the quality of the lab
work that you have done. You may need to show your notebook and your lab reports to the
Chemistry Department at a college or university in order to obtain credit for the lab part of an
AP Chemistry class. As you record information in your notebook, keep in mind that
someone who is unfamiliar with your work may be using this notebook to evaluate your lab
experience in chemistry. When you explain your work, list your data, calculate values and
answer questions, be sure that the meaning will be obvious to anyone who reads your
notebook.
Guidelines for the notebook:
1. Write your name and class on the front cover.
2. In black or blue ink, number all the right hand pages on the lower right corner if they are
not already numbered.
3. Save the first 2 pages for a Table of Contents. This should be kept current as you
proceed. Each time you write up a lab, place the title and page numbers where the lab
report begins in the Table of Contents.
4. Write in ink. Use only the right hand pages.
5. If you make a mistake, DO NOT ERASE OR SCRIBBLE. Just draw ONE LINE through
your error, and continue. It is expected that some errors will occur. A lab notebook is a
working document, not a perfect, error-free, polished product. Errors should be corrected
by drawing one line through the mistake, and then proceeding with the new data.
6. Do not use the first person or include personal comments.
Prelab Instructions:
1. On most every lab you will have prelab instructions. If it has you read, read carefully
as there will often times be a quiz over that content. If there are questions you are
supposed to answer, do them on a separate sheet of paper and hand them in as your
ticket into lab. If there is a code word in the procedure or weird instructions be
prepared to follow them.
2. Some labs will be full write-ups and some will be data and calculations only. You
must always answer questions if they are in the lab manual. You must get your data
stamped in your lab notebook before you leave the lab.
Lab Reports (Lab reports will be worth 50 points)
Include the following information in your lab reports. Label each section
1. Title – The title should be descriptive. Experiment 5 is not a descriptive title.
2. Date and lab station – This is the date you performed the experiment and lab station.
3. Purpose – A brief statement of what you are attempting to do. Must be a sentence.
4. Procedure – A shortened description of the method you are using. You may refer to
the lab manual for specific instructions, but you should include a brief statement of
the method. Do not include lengthy, detailed directions. A person who understands
chemistry should be able to read this section and know what you are doing.
5. Reactions: Write a balanced reaction including states of matter for any reactions. If
there are no reactions omit this section.
6. Data- Record all your data directly into your lab notebook on the right-hand pages.
Organize your data in a neat, orderly form. Label all data very clearly. Use correct
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sig figs and always include proper units. Underline, use capital letters or use any
device you choose to help organize this section well. Space things out – don’t try to
cram everything on one page. A data table must have a label and a title. e.g. – Table
1: Density Values for Sugar Solutions.
7. Calculations and Graphs- You should show how calculations are carried out. Give
the equation used and show how your values are substituted into it. Give the
calculated values. If graphs are included, make the graphs an appropriate size. Label
all axes and give each graph a title. If experiments are not quantitative, this section
may be omitted.
8. Conclusions – Make a simple statement concerning what you conclude from the
experiment. This is not a place to give your opinion of the lab and whether or not it
was “fun”. It is not your job to review the lab like you would if you saw a movie.
9. Experimental error – If there is a known value for something you are doing in lab,
calculate the experimental error.
10. Error Analysis – What are some specific sources of error, and how do they influence
the data? Do they make the values obtained larger or smaller than they should be?
Which measurement was the least precise? Instrumental error and human error exist
in all experiments, and should not be mentioned as a source of error unless they cause
a significant fault. Significant digits and mistakes in calculations are NOT a valid
source of error. In writing this section it is sometimes helpful to ask yourself what
you would do differently if you were to repeat the experiment and wanted to obtain
better precision.
11. Questions – Answer any questions included in the lab directions. Answer in such a
way that the meaning of the question is obvious from your answer.
Reporting Lab Data
Graphing Data
1. All graphs should have a descriptive title (“Graph” is not a title) and a label. e.g. –
Graph A: Density of Solutions with Varying Sugar Concentrations.
2. Both the vertical and horizontal axes should both have labels and units clearly
marked. Use a ruler to draw the axes.
3. The scales chosen should reflect the precision of the measurements. For example, if
temperature is known to be ±0.1ºC, you should be able to plot the value this closely.
Don’t have each block of the graph equal to 10ºC.
4. There should be a table in which the data values are listed. Don’t put data in a graph
unless you have first listed it in a table.
5. The controlled or independent variable is placed on the horizontal axis. The
dependent variable is graphed on the vertical axis.
6. There should be an obvious small point on the graph for each experimental value. It
is not necessary to include the coordinates of each point since they will be in the data
table.
7. A smooth line should be drawn that lies as close as possible to most of the points. Do
NOT draw a line connecting one point to the next as in a dot-to-dot drawing. If the
line is a straight line, use a ruler to draw it.
8. If a computer program is used to draw the graph, the rules still apply.
Accuracy
Accuracy is a measure of how close an experimental value is to a value which is
accepted as correct. The measure of the accuracy of an experimental value is reported as
Percent Error.
Experimental  Accepted
%error 
Accepted
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Data Tables
1. All data tables must be neatly organized. Numbers should be aligned by decimal
point. Appropriate units must always be used.
2. Data should be appropriately spaced out so that there is room for corrections or
annotations about the data.
3. All data must be in your lab notebook and initialed by the teacher before you leave
the lab.
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QRS lab
Reactions and explanations only, no full writeup.
In this lab, there are three flasks labeled Q, R, and S. Each flask contains one of the following
solutions: 0.1 M Pb(NO3)2, 0.1 M NaCl, or 0.10 M K2CO3. Two other flasks are labeled X
and Y. One of these flasks contains 0.1 M AgNO3 and the other contains 0.1 M BaCl2.
Mix each of the solutions with each of the other and record all observations. For all
precipitates which form, you must write a balanced equation and net ionic equation and
identify the precipitate. You will need to wait until you have identified the solutions to write
the equations. As you carry out the reactions you must use as little solution as possible. Part
of your grade is the way in which you are observed performing the reactions. Frugality is
key. You must explain how you reasoned out the solution’s identity.
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Qualitative Analysis of the Group III Cations
No writeup. Only data sheet.
Discussion
A known solution of the Group III cations (Cr3+, Al3+, Fe3+, Mn2+) and an unknown solution
containing some combination of ions will be analyzed.
Group III contains those cations whose hydroxides do not precipitate under highly acidic
conditions. Under basic conditions, however, the Group III cations will precipitate as the
hydroxides.
Cr3+ (aq) + 3OH-(aq) Cr(OH)3(s)
Al3+(aq) + 3OH-(aq)  Al(OH)3(s)
Fe3+(aq) + 3OH-(aq)  Fe(OH)3(s)
Mn2+(aq) + 2OH-(aq)  Mn(OH)2(s)
The pH is then raised with NaOH and hydrogen peroxide is added to further oxidize the
precipitates. The Fe(OH)3 (s) remains the same, the Mn(OH)2(s) becomes MnO2(s), the
Cr(OH)3(s) oxidizes to the chromate ion CrO42-(aq) and the Al(OH)3(s) complexes with more
hydroxide ion to form Al(OH)4-(aq).
MSDS
2M NH4Cl
solution
6M
Ammonia
6M NaOH
3% H2O2
3M H2SO4
6M HNO3
0.5M KSCN
PbO2
6M HC2H3O2
0.1 M
Pb(C2H3O2)2
Irritating to body tissues. Avoid all body tissue contact.
Liquid and vapor are strongly irritating to skin, eyes, and mucous
membranes. Vapor extremely irritating to eyes. May cause blindness.
Toxic by ingestion or inhalation. When heated to decomposition, emits
toxic fumes of NH3 and NOx.
Moderately toxic by ingestion and skin absorption. Corrosive to body
tissues. Causes severe eye burns. Avoid all body tissue contact.
Slightly toxic by ingestion or inhalation. Irritant to skin, eyes and
respiratory tract. Avoid prolong body contact. Hydrogen peroxide will
decompose rapidly when exposed to almost any substance.
Moderately toxic by ingestion. Corrosive to eye, skin, and all other body
tissues. Avoid all body tissue contact. Very considerable heat generated
when diluted with water.
Corrosive; will cause severe damage to eyes, skin and mucous
membranes. Moderately toxic by ingestion and inhalation. Strong
oxidizer. Avoid contact with acetic acid and readily oxidized substances.
Slightly toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact. Contact with acids or heat may liberate poisonous
hydrogen cyanide gas.
Moderately toxic by ingestion or inhalation. Irritating to body tissues.
Avoid all body contact. Oxidizer. Lead and lead compounds are
possible carcinogens.
Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated.
Moderately toxic by ingestion and skin absorption. Eye and skin irritant.
Possible carcinogen. Avoid ingestion, inhalation and skin absorption.
Chronic exposure to inorganic lead via inhalation or ingestion can result
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in accumulation in and damage to the soft tissues and bones.
Procedure
Obtain a Group III known sample which contains all of the ions. You will get a Group III
unknown solution (which may contain any or all of these cations) after you have completed
the known. All glassware should be cleaned, rinsed and rinsed with distilled water before
starting the lab.
Step
Notes
1.
Place 1-2 mL of solution to be tested in a small test
tube.
2.
Add 1 mL of 2 M ammonium chloride solution to the
sample in the test tube and stir. Add 6 M aqueous ammonia to
the sample dropwise until the solution is just barely basic
(remove a drop of the solution with a stirring rod and touch
the drop to a strip of pH test paper).
3.
Add about 3 mL of distilled water to wash the
precipitate. Mix thoroughly, centrifuge, decant and discard
water.
4.
Add ~ 2mL 6M NaOH to the residue and mix.
5.
Add 10 drops 3% H2O2 and mix immediately. Boil
several minutes to remove excess H2O2. If solution is green,
add more H2O2. (yellow is okay) Look for separation of
precipitate and supernatant. Centrifuge, decant and obtain
residue 2 for step 6, and decantate 2 for step 12. Do not
discard. If necessary, recentrifuge decantate 2 until
absolutely clear or filter into another test tube.
6.
Add 1-2 mL 3M H2SO4 to residue 2. Mix.
7.
Add 5-6 drops 3%H2O2 to hasten the process. Mix
thoroughly. If it doesn’t dissolve, heat for a few minutes until
all the solid dissolves.
8.
Dilute solution to a total of 4 mL with distilled water
and divide the solution into 2 parts to be tested for Fe3+ ions
in step 9 and Mn2+ ions in step 10.
9.
Add 1-2 drops 0.5M KSCN to one test tube from part
6. A blood red solution indicates the presence of Fe3+ ions.
10.
To the other half of the solution from step 6, add ~
1mL 6M HNO3 and mix.
11.
Add solid PbO2 equivalent to 1/10 the volume of the
liquid. Mix well and Boil for 2-3 minutes and let stand for 3
minutes. If Mn2+ ions are present, the solution will turn a
pink to dark purple color. If test is negative, add another
small portion of acid, mix and Boil solution. Centrifuge.
12.
Dilute decantate 2 from step 4 to ~ 4mL with distilled
water and divide into two parts to be tested for Al3+ in step 13
and CrO42- in step 14.
13.
Add ~ 2mL 2M NH4Cl to the first half of decantate
from step 12. DO NOT MIX. Place in Boiling water for 5
minutes. Look closely for a fluffy, translucent solid in the top
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layer which indicates the presence of Al3+ ions. If uncertain
about the aluminum test, centrifuge. If other ions are present,
the decantate may not be clear and a halo effect may be seen
around the precipitate. The iron and manganese hydroxides
will spin down first because of the greater densities and the
aluminum hydroxide will be on top.
14.
Add 6M HC2H3O2 to second half of decantate from
step 9 until acidic to litmus paper. Add 1-2 drops 0.1 M
Pb(C2H3O2)2. Let stand for 3 minutes. A white or yellow
precipitate is a positive test for CrO42- ions.
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Unknown #__________
Station #______________
Name(s)________________________________________________________
Presence of Aluminum ion
Yes
No
Presence of Chromium ion
Yes
No
Presence of Iron ion
Yes
No
Presence of Manganese ion
Yes
No
Group III Cation Analysis
Al3+, Cr3+, Fe3+, Mn2+
NH4Cl
NH4OH
Residue 1
Decantate 1
discard
Fe(OH)3, Mn(OH)2, Cr(OH)3, Al(OH)3
NaOH(aq)
H2O2
Residue 2
Decantate 2
CrO42-, Al(OH)4-
H2SO4(aq)
H2O2
3+
Fe
KSCN
Fe(SCN)2+
blood red
2+
Mn
HNO3(aq)
PbO2
MnO4Pink/purple
HC2H3O2
Pb(C2H3O2)2
PbCrO4
yellow or
white
NH4Cl
Al(OH)3
Translucent
flecks
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How Much Zinc is in a penny?
Full writeup
Discussion
Up until 1982 pennies were made with mostly copper. The history of the penny
shows a plethora of changes in composition and sizes but I digress. If you are interested,
google it. In an effort to save money the treasury started putting a zinc core with a thin
copper skin around it. In this lab you will scratch away the thin copper on an edge to expose
the zinc. You’ll then soak it in acid overnight to dissolve the zinc. You will then choose a
salt that you can use to precipitate out the zinc and calculate the mass of zinc in a penny.
Procedure
Write your own. Pick an acid, 6 M HCl, 6M H2SO4, or 6M HC2H3O2.
Write out a procedure to dissolve the zinc and then choose a salt from the storeroom that you
can dissolve in water and add to the dissolved zinc to precipitate it out.
You must get your procedure approved by your teacher.
Collect your data and look up the actual value online. Cite your source when you report the
expected mass of the zinc. Calculate a percent error.
Questions:
1. Why not use iron? Iron’s cheap.
2. Name 2 other salts that you could have used to precipitate out the zinc.
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Predicting Products of Chemical Reactions
Only reactions and data.
Discussion:
It is not always easy to predict the product of a chemical reaction. Often, a reaction must be
carried out and the products analyzed in order to determine what formed. However,
frequently a very good prediction can be made if one analyzes the type of reaction that
occurs.
When writing ionic equations, soluble ionic compounds that are strong electrolytes are
written as separated ions. Since strong acids and bases ionize totally in water, their formulas
are also written as separated ions. Weak acids and bases are written as molecular
compounds. Even though a weak acid such as acetic acid does ionize slightly in water, the
percent of ionization is very small. Only the formulas of the major species are shown in the
net ionic equation. Solids, gases, and nonionic liquids, whether dissolved or nor, are written
as neutral, molecular formulas. See your notes about types of reactions.
Procedure:
Carry out the reactions as described. Create a data table in which you describe the reactants,
products, and any indication that a reaction has occurred. Identify the type of reaction.
Write a net ionic equation for the reaction. It is not necessary to balance the reaction.
1. Mix 1 mL of 0.1M sodium chloride with 1 mL of 0.1M silver nitrate. Save the product
for step 2
2. Centrifuge the precipitate from step 1. Decant and discard the decantate. Add 1 mL of
6M ammonia to the precipitate and agitate.
3. Add 2 mL 6M hydrochloric acid to the solution from 3. Test with litmus to be sure the
solution is acidic. If not acidic, add 1 more mL of the acid.
4. Add a tiny bit of calcium oxide to distilled water. Test with litmus.
5. Put 10 mL of 0.1M potassium iodide solution in a 50 mL beaker. Using a 9 volt battery
as a power source and graphite electrodes, allow an electric current to pass through the
solution for two minutes. Test the solution near the electrodes with litmus.
6. Place a small amount of sodium hydrogen carbonate in a test tube. Add 1 mL of 0.1 M
acetic acid.
7. Place 2 mL of oxalic acid solution in a test tube. Make a loose ball of a small piece of
aluminum foil (2 cm square) and drop it into the solution. Use a stirring rod to push it
under the solution. Wait five minutes and observe.
8. Place 1 mL of 0.1M sodium phosphate in a test tube, add 5 mL of 0.1 M hydrochloric
acid. Feel the outside of the tube for evidence of a reaction.
9. Place 2 mL of 0.1 M sulfuric acid in a test tube. Add a small spatula of solid sodium
bicarbonate.
10. Place 2 mL of 0.1M iron III nitrate in a test tube. Add several drops of 0.1M potassium
thiocyanate solution.
11. Place 2.0 mL of 0.1 M aluminum nitrate in a test tube. Dropwise, with agitation, add 2
mL of 1.0M sodium hydroxide.
12. Place 1 mL of 0.1 M copper II sulfate in a test tube with 3 mL of 0.1 M sodium
hydroxide solution.
13. Place 2.0 mL of 0.1M iron III nitrate in a test tube. Add 2.0 mL of 0.1M ammonia.
14. Place 1.0 mL of 0.1 M barium hydroxide in a test tube. Add 5 mL of 0.1 M sulfuric acid.
15. Place 5 mL of 0.1M silver nitrate in a test tube. Add a small coil of copper wire.
Observe after 5 minutes.
16. Place 5 mL of 3% hydrogen peroxide in a test tube. Add a small amount of solid
manganese IV oxide.
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Redox Titration: The Standardization of Potassium
Permanganate
Full writeup
Discussion:
Oxalate is commonly used as a primary standard for determining the concentration of many
strong oxidizers used in oxidation/reduction analyses. In this experiment, oxalic acid will
be used to determine the concentration of potassium permanganate.
Oxalic acid is a strong electrolyte that dissociates completely in water. The oxalate ion,
C2O42-, will react quantitatively with permanganate ion, MnO4-, in the presence of strong
acid according to the following equation:
MnO4- (aq) + C2O42- (aq)  Mn2+ (aq) + CO2 (g)
The numerical value of Keq for the equilibrium is very large. (The reaction goes totally
toward the right, very little to the left) When reaction conditions are anywhere near
optimum (in terms of pH and temperature) the reaction can be considered to be
quantitative.
Prelab:
1. Balance the equation above.
2. Identify what is oxidized and what is reduced.
3. Identify the oxidizing agent and the reducing agent.
MSDS:
KMnO4
Sodium oxalate
Sulfuric acid
Irritating to body tissues. Avoid all body tissue contact.
Moderately toxic by ingestion and inhalation. Corrosive to body
tissues. Avoid contact with all body tissues.
Moderately toxic by ingestion. Corrosive to eye, skin, and all other
body tissues. Avoid all body tissue contact. Very considerable heat
generated when diluted with water.
Procedure
1. Transfer 200mL of water to a 400mL (or larger) beaker. Place the beaker on a hot plate.
2. Add 40mL of 6.0M sulfuric acid to the water while stirring with a glass rod.
3. Turn on the hot plate. Monitor the temperature with a thermometer. You will need to heat
the acid solution to about 80oC to 90oC. Continue with the procedure while the acid
solution is heating. (This removes CO2 from the water)
4. Use an analytical balance to mass 0.134g to 0.149g (NO MORE!) of reagent grade sodium
oxalate in a weigh boat. Record the mass of sodium oxalate to the nearest 0.001g.
5. Transfer the sodium oxalate sample to the acid solution. You may use a water bottle to
facilitate the transfer if necessary. Stir until dissolved.
6. Rinse a buret with a few milliliters of the potassium permanganate solution. Dispose of the
rinsing down the sink.
7. Fill the buret with the permanganate solution above the zero mark. Drain the buret until the
liquid level is below the zero mark and to clear most of the air from the tip of the buret.
Mount the buret in a buret clamp on a ring stand and position the buret over the
acid/oxalate solution.
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8. Once the temperature of the acid/oxalate solution has reached the desired 80oC to 90oC
range turn off the hot plate.
9. Rapidly add ~5mL of the permanganate titrant and stir the solution to stir until the purple
color disappears.
10. Continue the titration dropwise with constant stirring. Position the buret’s stopcock so that
the permanganate titrant drips at a moderate pace. Do not allow the titrant to pour rapidly
beyond the first five mLs.
11. As the titration approaches its equivalence point the purple color from the permanganate
will persist longer and longer before it disappears.
12. Slow the rate of titration as the equivalence point is approached.
13. Immediately stop the titration when you believe the equivalence point is reached. The
equivalence point is defined as the point in the titration where the color from the titrant
persists for 30 seconds. A perfect titration will yield a faint pink solution that persists for
30 seconds.
14. If the color fades before the 30 second requirement add more titrant one drop at a time until
the 30 second requirement is reached.
15. Once you are satisfied the equivalence point has been reached record the volume of titrant.
16. Run another trial mixing another sample of oxalic acid to verify your results.
17. Calculate the molarity of the permanganate solution. Be sure to show all calculations in the
calculations area.
18. Post your molarity on the class result sheet hanging in the lab.
19. Your percent error is done using the average molarity for the class as the accepted value.
Clean-up:
1. Pour the remainder of the permanganate back into the stock bottle.
2. Rinse the buret with some tap water, rinse again with a hydrogen peroxide solution
located in the fume hood and then rinse twice with distilled water.
3. Wash the beaker with soap, rinse, rinse and distilled rinse. Put on a paper towel at your
station to dry.
Conclusion:
In your conclusion describe why it was necessary to add the sulfuric acid. Research some
redox reactions that happen in your body and discuss the value of redox reactions in life.
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The EMF Activity Series
Full writeup
Discussion:
The activity series of metals is a table of metals arranged in the order of their decreasing
tendency to lose electrons and enter into chemical reactions. The table of metals is arranged
in the order of decreasing electropositive character. These reactions are redox reactions
where one metal ion is reduced to its elemental state and another more active metal is
oxidized to an ionic state. For example, 2Na + Fe+2  2Na+ + Fe.
Although hydrogen has many physical and chemical properties that are similar to nonmetals,
it frequently functions chemically as a metal, and for this reason, it is included in the activity
series of metals. Its placement indicates that the metals preceding it will displace it from nonoxidizing acids. The metals that are found uncombined in nature in large amounts are those
that are less active than hydrogen, whereas those metals that are more active than hydrogen
are not usually found in the free state. Two metals that are exceptions are metallic iron and
nickel found in meteorites.
The normal test of the chemical activity of an element is its displacing power. If the metal
can displace another metal from a compound, it may be considered to be more chemically
active than the metal it displaces. The relative activity of a metal may be determined by
observing: 1 - metal reactivity with water: cold, warm, or hot; 2 - metal reactivity with acids:
hydrogen producing acids and non-hydrogen producing acids: 3 - metal activity with bases;
and 4 - metal reactivity with salt solutions.
It should be noted that the more finely divided the state of the metals-powdered form rather
than large lumps-the more surface area is exposed, and the greater the activity of the metal. If
the reaction is heated, the reactivity of the elements and compounds tend to increase.
Halogens can also be organized according to their ability to displace other halogens. In the
reaction between a free halogen X2 and a halide ion Y-, the free halogen gains electrons is
reduced to its halide ion X2 + 2e-  2X-. The original halide ion is oxidized to the free
halogen state. 2Y-  Y2 + 2e-. The most reactive halogen is the one most easily reduced
(most hungry for the electron). To determine if a reaction occurs, a method is needed to
identify which halogen is present. Halogens dissolve in the nonpolar solvent mineral oil
forming different colored solutions. Mineral oil does not dissolve in water, but when shaken
with an aqueous halogen solution, the halogen is extracted from the water into the mineral
oil. The color of the mineral oil indicates which halogen is present.
Prelab:
Read the entire discussion and procedure.
MSDS:
0.1 M Fe(NO3)3
0.1 M AgNO3
0.1 M CuSO4
0.1 M Zn(NO3)2
Copper
Lead
Corrosive to body tissues by contact and inhalation. Avoid contact with
skin, eyes and mucous membranes.
Moderately toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact.
Mildly toxic by ingestion. Irritant to skin, eyes and mucous
membranes. Avoid contact with body tissues.
Slightly toxic by ingestion. Corrosive to body tissues. Avoid all body
tissue contact.
Irritant to body tissues as dust. Avoid contact with nitric acid, emits
toxic fumes of nitrogen oxides.
Lead as a powder or dust is toxic by ingestion or inhalation. Lead and
lead compounds are possible carcinogens. Avoid ingestion and
6/23/2017
Magnesium
Zinc
Calcium
Sodium
6M HCl
NaBr solution
NaCl solution
KI solution
Cl2 water
Br2 water
Page 16 of 114
inhalation. Emits highly toxic fumes of Pb when heated. Chronic
exposure to inorganic lead via inhalation or ingestion can result in
accumulation in and damage to the soft tissues and bones.
Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated.
Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated. Inhalation of zinc
dust may cause lung irritations. Zinc dust can spontaneously combust
when in contact with moisture.
Corrosive solid. Avoid body tissue contact. Violent reaction with water
may evolve explosive hydrogen gas. Flammable solid.
Highly corrosive solid, avoid all body tissue contact. Will severely burn
skin, eyes, or internal tissues. Reacts violently with water releasing
hydrogen gas, which will ignite and explode in air.
Toxic by inhalation and ingestion. Severe corrosive to all body tissues,
especially skin and eyes. Avoid all body contact.
Possible body tissue irritant.
Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated.
Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated.
Toxic by inhalation and ingestion. Very irritating to mucous
membranes. This is a weak solution of chlorine gas and water. Chlorine
gas will slowly leave solution.
Highly toxic by ingestion and inhalation. Severe skin irritant; may
cause burns and irreversible eye damage. Strong oxidizer, heat of
reaction may ignite combustibles on contact. Will react with water or
steam to produce toxic and corrosive fumes. Extremely hazardous
substance.
Your lab report must include net ionic equations for each single replacement reaction that
occurs. Leave plenty of room before the data section.
Procedures:
1. For each section below, follow the instructions and perform the reactions.
2. Describe each reactant and product.
3. If no reaction occurs, write N.R.
4. Arrange test tubes in the test tube rack in order as in table 1.
Part 1
Reactions of metals with salt solutions
1. Place 1 to 2 mL of each of the following solutions into the designated number of test
tubes: iron (III) nitrate - 4, silver nitrate - 4, copper (II) sulfate – 4, and zinc nitrate –
4.
2. Obtain the number of strips of each metal as indicated: copper -4, lead - 4,
magnesium - 4, and zinc - 4. Place a strip of metal into each test tube as indicated by
the table. Allow the solution with metal to sit for at least 15 minutes. At the end of 15
minutes, record all changes in the solution and metal in each of the test tubes. If no
reaction occurs within 20 minutes, write N. R. rather than a description.
3. While waiting for the reactions in the salt solutions to occur, set up the reactions for
metals in acid solutions and metals in water.
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Reactions of metals with acid
Place 1 -2 mL of 6M HCl in each of the designated test tubes. Place a strip of each metal into
separate test tubes of acid. Observe immediately for any sign of a reaction occurring. Record
your observations of the changes.
Reactions of metals with water
Place 1-2 mL of water into 3 separate test tubes along with 1 drop of phenolphthalein. Place a
strip of magnesium into one of the test tubes, a chunk of calcium into one of the test tubes
and a piece of sodium into the other test tube. At Record observations of any reactions and
the changes in the color. A pink color indicates a base meaning a reaction has occurred..
Table 1
HCl (aq)
Cu
Pb
Mg
Zn
Na
Ca
Mg
Water
Fe(NO3)3 (aq)
AgNO3(aq)
CuSO4(aq)
Zn(NO3)2
Part 2 Activity Series for some halogens
If you notice a really strong smell that bothers you, do the following steps in the fume hood.
1. As a reference, a rack of test tubes has been set up for you at the front of the room. They
indicate what color the mineral oil changes when combined with either each halogen or
the halide ions.
2. Set up 6 test tubes in a test tube rack according to the following table.
Tube
Contents
1.
NaBr + Cl2
2
KI + Cl2
3
NaCl + Br2
4
KI + Br2
5
NaCl + I2
6
NaBr + I2
3. Each tube gets a mL of each solution. Cork each and shake to mix. Add 1 mL of mineral
oil to each, cork and shake to mix.
4. When the mineral oil layer has separated, determine its color and whether a reaction has
occurred. For example in tubes 1 and 2, if the color of the chlorine appears in the mineral
oil layer than no reaction has occurred. If either the bromine or iodine color appears in
the mineral oil layer, then there was a reaction. Record your observations.
Clean Up
After all observations are made, empty the test tubes into the corner of the sink, rinse with
water, remove the metal to a paper towel at your lab station, wash each test tube, rinse with
tap water and do a final rinse with distilled water. Turn the clean test tubes upside down in
the test tube rack and leave for the next class.
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CONCLUSIONS:
Page 18 of 114
1. According to the tests done in lab, list the metals and hydrogen from most active to least
active. Be sure to also include those metals and hydrogen that were in ionic form in the
solutions used. In your conclusion give evidence from your results for your conclusion.
2. According to the tests done in lab, list the halogens from most active to least active. In
your conclusion give evidence from your results for your conclusion.
3. Were there any tests for the activities of the metals and hydrogen that did not agree with
the order of activity according to the EMF activity series? If there were disagreements,
why do you think that happened?
4. No error analysis or % error
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Page 19 of 114
Heat of Fusion for Ice
Data and Calculations and question only
Discussion:
Melting and freezing behavior are among the characteristic properties that give a pure
substance its unique identity. As energy is added, pure solid water (ice) at 0°C changes to
liquid water at 0°C.
The equation we will use is q=mct where q is heat measure in joules.
In this experiment, you will determine the energy (in joules) required to melt ice. You will
then determine an experimental value for the molar heat of fusion for ice (in kJ/mol) and
compare it to the accepted value. Excess ice will be added to warm water, at a known
temperature, in a Styrofoam cup. Heat from the warm water will be used to warm the ice to
its melting point, 0oC, to melt the ice, and then to warm the resulting liquid to some
temperature above 0oC.
The heat balance for the system can be defined by the following equation:
(c(ice) •t •mice)
+ (miceHf)
+ (c(water) •t •m(water from melt))
c(water) •t •m(hot water)
Heat absorbed by the ice
Heat it takes to melt the ice
Heat absorbed by the melted ice
Heat lost by the hot water
where c is specific heat capacity (one value for liquid water another for ice) , m is mass in
grams, Hf is the heat of fusion for ice, and t is the change in temperature (each t should
be unique) . The minus sign is used on the left because heat is being lost. For liquid water, c
is 4.18 J/g°C. For ice, c is 2.03J/goC
Prelab:
1. Calculate the amount of heat necessary to heat 15.00 g ice from -20 to 0 °C.
2. Calculate the mass of water when 7700. J of heat is added to water at an initial
temperature of 29° which increases 69 degrees.
MSDS:
Ice
It’s cold
Procedure
1. Fill a 250mL beaker about 2/3 full and place it on a hotplate. Turn the hotplate control to
full. The water will need to heat to above 70oC.
2. Obtain a Styrofoam cup for use as a calorimeter and obtain a second Styrofoam cup filled
with ice.
3. Use a digital thermometer to measure the initial temperature of the ice. This value can
range from –20oC to 0oC depending on how long the ice has been out of the freezer.
Record this value on the data sheet as “initial ice temperature.”
4. Once the water is hot, use a 100-mL graduated cylinder to transfer 100.0 mL of the hot
water to the styrofoam cup. Use a balance to determine the mass of the water. Use
beaker tongs and/or “hothand” glove to make the transfer. Place the thermometer in the
Styrofoam cup containing the hot water and allow the temperature reading to stabilize.
Record the temperature reading as “initial hot water temperature.”
5. When ready, start adding ice to the hot water. Stir the mixture continuously with the
thermometer. Continue to add ice as the ice melts to maintain a significant excess of ice.
6. Once the temperature goes below 10oC quickly remove any remaining ice from the
Styrofoam cup. Record the minimum observed temperature as the “final system
temperature.”
7. Record the mass of the water after the melt.
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Data and Calculations
Be sure to collect the following data.
Initial ice temperature
Initial hot water temperature, t1
Final system temperature, t2
Change in temperature of the hot water, t
Change in temperature of melted water, t
Change in temperature of ice before melting, t
Final water mass
Initial water mass
Mass of melt
SHOW YOUR CALCULATIONS for the following:
Mass of ice melted
Heat released by the hot water as it cooled
Heat gained by solid ice as it warmed from its initial temp to its melt temp
Heat gained by liquid melt as it warmed from its melt temp to final system temp
J/g ice melted (heat of fusion)
kJ/mol ice melted (molar heat of fusion)
Percent error (6.03 kJ/mol is the accepted value)
Questions:
1. Research the heat of fusion of paraffin wax. How would it be useful for a
homeowner’s wall to be filled with blocks of paraffin wax?
2. Go to this website, what kinds of PCM’s are used in gloves?
http://www.textileworld.com/Articles/2004/March/Features/Phase_Change_Materials
.html
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Page 21 of 114
Additivity of Heats of Reaction: Hess’s Law
Full writeup
In this experiment, you will use a Styrofoam-cup calorimeter to measure the heat released by
three reactions. One of the reactions is the same as the combination of the other two
reactions. Therefore, according to Hess’s Law, the heat of reaction of the one reaction should
be equal to the sum of the heats of reaction for the other two. This concept is sometimes
referred to as the additivity of heats of reaction. The primary objective of this experiment is
to confirm this law. The reactions we will use in this experiment are: (Write equations for
each one including state of matter for Prelab)
(1) The dissolution of solid sodium hydroxide. H1 = ?
(2) Solid sodium hydroxide reacts with aqueous hydrochloric acid.
H2 = ?
3) Solutions of aqueous sodium hydroxide and hydrochloric acid react.
H3 = ?
You will use a Styrofoam cup as a calorimeter. For purposes of this experiment, you may
assume that the heat loss to the calorimeter and the surrounding air is negligible. Even if heat
is lost to either of these, it is a fairly constant factor in each part of the experiment, and has
little effect on the final results.
MSDS:
NaOH
0.50 M and 1.0 M HCl
Highly toxic by ingestion, inhalation, or skin absorption.
Extremely corrosive to body tissues. Causes severe eye burns.
Avoid all body tissue contact.
Toxic by inhalation and ingestion. Severe corrosive to all body
tissues, especially skin and eyes. Avoid all body contact.
Procedure
Reaction 1
1. Measure out 100.0 mL of water into the Styrofoam cup. Place the thermometer into the
solution. Allow the temperature reading to stabilize and record the temperature as the
“initial temperature.” Be sure to record the mass of the water.
2. Weigh out about 2 grams of solid sodium hydroxide, NaOH, and record the mass to the
nearest 0.001 g. Since sodium hydroxide readily picks up moisture from the air, it is
necessary to weigh it and proceed to the next step without delay. Caution: Handle the
NaOH and resulting solution with care.
3. Add the NaOH to the water in the Styrofoam cup and swirl the cup to aid dissolution.
Monitor the temperature increase. Record the maximum temperature as the “final
temperature.”
4. Rinse and dry the thermometer and Styrofoam cup. Dispose of the solution in the beaker
in the fume hood.
Reaction 2
5. Repeat the steps for reaction 1, using 50.0 mL of 1.0 M hydrochloric acid added to 50.0
mL of distilled water instead of water. Use approximately the same amount of solid
NaOH as before. Caution: Handle the HCl solution and NaOH solid with care. Record
the mass of the water before you add the solid NaOH.
Reaction 3
6. Run another trial using 50.0 mL 1.0 M HCl in place of water and 50.0 mL 1.0 M NaOH
in place of solid NaOH. Record the initial temperature of the HCl and the final
temperature of the mixture. Record the mass of the two solutions together.
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Processing Data
Page 22 of 114
Determine the temperature change, t, for each reaction.
Calculate the heat released by each reaction, q,(Convert joules to kJ in your final answer.)
Calculate moles of NaOH used in each reaction.
Use the results of the Step 4 and Step 5 calculations to determine H/mol NaOH in each
of the three reactions.
5. Use Hess’s law. 1 and 2 should add up to 3. Use 3 as the accepted value. Calculate
percent error.
6. Record your value on the class data sheet in the lab. Report your data and the class
average data.
1.
2.
3.
4.
Conclusion:
Using terms of KMT, describe why temperatures of the solutions increased. Describe why it
doesn’t matter how much reactant you use when calculating H. Describe why Hess’s Law
is useful. Compare your data to the class average.
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Heat of Combustion of a metal-an inquiry based approach
Data and Calculations and questions only
Discussion:
In this experiment you will use Hess’ Law of the additivity of reaction heats to determine the
heat of reaction for a reaction that is difficult to measure directly – the combustion of either
magnesium or zinc. The combustion of a metal with a 2+ oxidation state is represented by
the following equation:
M (s) + ½ O2 (g)  MO (s)
This equation can be obtained by creatively combining equations (1), (2), and (3).
(1)
MO (s) + 2HCl (aq) 
(2)
M (s) + 2HCl (aq) 
(3) H2 (g) + ½ O2 (g)  H2O (l)
The heats of reaction for equations (1) and (2) will be determined in this experiment.
H/mol for reaction (3) is –285.8 kJ/mol.

Prelab: Find the H f for each reactant in equations 1, 2 and 3. Read all lab procedures.
MSDS:
MgO Inhalation may cause respiratory irritation. Slight eye irritant.
Highly toxic by inhalation and ingestion. Severe corrosive to all body tissues,
HCl
especially skin and eyes. Avoid all body contact.
Flammable solid. Substance not considered hazardous. However, not all health
Mg
aspects of this substance have been thoroughly investigated.
Inhalation of zinc dust may cause lung irritations. Zinc dust can spontaneously
Zn
combust when in contact with moisture.
Moderately toxic by ingestion and inhalation. Body tissue irritant. Avoid all body
ZnO
tissue contact.
PROCEDURE
Part I MO + HCl
1. Look up the specific heat of 6M HCl. You will need approx. 1 ±0.5g g of MO and
approximately 100 g ± 50 g of 6M HCl
2. Determine the mass you will use of the MO and the HCl. Write a procedure that will
determine the H for the reaction #1 above per mole of MO.
3. Do the stoichiometry necessary to find the limiting reactant and amount of excess
reagent remaining.
4. Have your teacher approve both the stoichiometry and the procedure you wrote. After
approval perform your experiment and collect the data.
Part II Hydrochloric Acid Plus M
1) You will need approx. 0.5 ±0.25g g of M and approximately 100 g ± 25 g of 6M HCl
2) Determine the mass you will use of the M and the HCl. Write a procedure that will
determine the H for the reaction #2 above per mole of M.
3) Do the stoichiometry necessary to find the limiting reactant and amount of excess reagent
remaining.
4) Have your teacher approve both the stoichiometry and the procedure you wrote. After
approval perform your experiment and collect the data.
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Page 24 of 114
Processing the Data
1. Calculate T for each reaction.
2. Calculate the heat using mcat
3. Calculate the moles of MO and M used.
4. Calculate the H per mole of MO for the first reaction and the H per mole of M for the
second reaction. Record these values on the data sheet.
5. Determine the heat of combustion for the metal by rearranging the equations 1-3 above
using Hess’s law. Show the correct rearrangement as part of the calculations.
6. Determine the percent error for your experimental result. Look up the heat of reaction for
oxidation of zinc or magnesium and use that as the theoretical.
7. Share your data with your classmates and do a 2nd percent error using the class average
for your metal as the expected.
Questions:
1. Research magnesium decoy flares and describe how modern Stinger missile systems can
tell the difference between aircraft engines and these flares.
2. Describe the difference between the H for each metal and describe why one metal
would have a higher H than the other.
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Page 25 of 114
VSEPR and Molecular Geometry
Copy this table in your lab notebook, no formal writeup
Design the Lewis structure and its molecular geometry on a scratch piece of paper and then
draw it in the correct orientation in your lab notebook. Decide whether the molecule is polar
or non polar. Not the bond but the molecule.
Formula
Lewis Structure
Bonding
Lone Pairs
Molecular
Polar?
Pairs
Geometry
(Y/N)
1
CH4
2
NH3
3
H2O
4
CO2
5
SO3
6
SO327
H2CO
8
HCN
9
BrF3
10
I311
C2H4
12
XeF2
13
SO4214
PBr3
15
CH3+
16
BH417
AsCl5
18
BrF5
19
SF6
20
ClO321
O3
22 Thiocyanate
23
IF4+
24
NO
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Page 26 of 114
Formation of a Coordination Complex of Copper (II)
Written observations in lab notebook
Discussion
The class of substances referred to as coordination compounds generally contains a central
metal atom, to which a fixed number of molecules or ions (called ligands) are coordinately
covalently bonded in a characteristic geometry. Coordination complexes are very important
in both inorganic and biological systems. For example, the molecule heme in the oxygenbearing protein hemoglobin contains coordinated iron atoms. Chlorophyll, the molecule that
enables plants to carry on photosynthesis, is a coordination compound of magnesium.
In general, coordination compounds contain a central metal ion that is bound to
several other molecules or ions by coordinate covalent bonds. For example, the compound to
be synthesized in this experiment, tetramminecopper (II) sulfate, consists of a copper (II) ion
surrounded by four coordinated ammonia molecules. The unshared pair of electrons on the
nitrogen atom of the ammonia molecule is used in forming the coordinate bond to the copper
(II) ion. In other words, the ammonia acts as a lewis base to the copper’s lewis acid.
One property of most transition metal coordination compounds that is especially
striking is their color. Generally the coordinate bond between the metal ion ad the ligand is
formed using empty low-lying d-orbitals of the metal ion. Transitions of electrons within the
d orbitals correspond to wavelengths of visible light, and generally these transitions are very
intense. Coordination complexes of metal ions are some of the most beautifully colored
chemical substances known; frequently they are used as pigments as paints.
Prelab: Read the procedure; write the formula for tetraaminecopper (II) sulfate.
MSDS:
CuSO4
NH4OH
Et-OH
Mildly toxic by ingestion. Irritant to skin, eyes and mucous membranes. Avoid
contact with body tissues.
Liquid and vapor are strongly irritating to skin, eyes, and mucous membranes.
Vapor extremely irritating to eyes. May cause blindness. Moderately toxic by
ingestion or inhalation.
Toxic by ingestion and inhalation. Body tissue irritant. Avoid all body tissue
contact. Denatured with isopropanol and methanol. Not for human consumption.
Flammable liquid.
Procedure:
1. Weigh out approximately 1.0 g of copper II sulfate pentahydrate. Record the mass.
2. Dissolve the copper salt in approximately 10 mL of distilled water in a beaker or
flask. Stir thoroughly to make certain that all the copper salt has dissolved before
proceeding. Record the color of the solution at this point.
3. Transfer the copper solution to the exhaust hood and with constant stirring; slowly
add 5 mL of concentrated ammonia. The first portion of ammonia added will cause a
light blue precipitate of copper II hydroxide to form. But upon adding more
ammonia, this precipitate will dissolve as the ammonia complex forms. Record the
color of the mixture after all of the ammonia has been added.
4. To decrease the solubility of the tetramminecopper (II) complex, add approximately
10 mL of ethyl alcohol with stirring. A deep blue solid should precipitate. Code word
“cinnabon”
5. Allow the precipitate to stand for 5 minutes and then filter the precipitate.
6. While in the filter paper, wash the precipitate with 2-5 mL portions of alcohol and
stir.
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7. Remove the filter paper and allow the filter paper to dry.
Page 27 of 114
Clean up: Filtrate can be put down the sink with excess water. Dried precipitate can be
thrown into the garbage can. Clean all glassware with soap and water, rinse, rinse distilled
rinse. Wash hands before leaving the lab.
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Page 28 of 114
Kinetics of a Reaction -- An Iodine Clock
Full writeup
Discussion:
Bisulfite ion and iodate ion react to form elemental iodine and sulfate ion according to the
following equation:
HSO3- + IO3-  I2 + SO42- + H+ + H2O
This reaction is believed to have a two step mechanism as follows:
Step 1
IO3- + HSO3-  I- + SO4-2 + H+
Step 2
I- + H+ + IO3-  I2 + H2O
In this kinetics experiment the student will study the effect of reactant concentration of the
rate of the reaction. The student will measure the time required for the reaction to produce
the product iodine. The experiment will be conducted in three trials and each trial will be
duplicated. The first trial will serve as a baseline for comparison. In the second trial, the
student will reduce the iodate concentration by half leaving the bisulfate concentration the
same as in the baseline trial and determine the effect on reaction time. In the third trial, the
student will reduce the bisulfate concentration by half leaving the iodate concentration the
same as in the baseline trials. The student will use the data to compute the reaction order for
each of the two reactants. If the order of each reactant is known the student can deduce
which of the steps in the mechanism is the slowest; that is, which step is the rate-determining
step.
The rate of this reaction can be defined by the following equation:
Rate = k[IO3-]x[HSO3-]y
Where:
k is a temperature dependent rate constant
[IO3-] is the initial concentration of iodate
[HSO3-] is the initial concentration of bisulfite
and x and y are the reaction orders of iodate and bisulfite respectively
The reaction order values for a set of reactants are totally unrelated to the stoichiometric
coefficients. In all cases reaction order values must be determined experimentally. On
edmodo there is a link to a video and a paper handout of how to solve for the exponents if
you don’t remember how to do logs.
Prelab: Read all the procedures and be ready to hand me a slip of paper 3 cm square with
your name on it in order to get into the lab.
MSDS:
KIO3
Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated.
NaHSO3
Mild body tissue irritant. Avoid contact with body tissues.
Starch
Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated.
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Page 29 of 114
Procedure
1. Obtain three 100 or 150mL beakers and label them “KIO3”, “NaHSO3”, and “Starch”.
2. Obtain ~75mL of the 0.04M KIO3 solution, ~75mL of the 0.04M NaHSO3 in 0.04M
H2SO4 solution, and about 50mL of the starch solution in the labeled beakers.
3. Follow the trial table below to prepare the reaction mixture for each trial. Use syringes to
measure out the correct amount of each solution. A clean 100mL graduated cylinder can
be used for water and starch measurements. The water and starch may be combined in
the 100mL graduated cylinder during measurements. Add the contents of the
graduated cylinder to an Erlenmeyer flask and then add the KIO3. Always add the
KIO3 last when you start the timer. Notice the total volume for each trial is 100mL
4. In each trial the water, starch, and NaHSO3 should be added to a beaker. One student
must then start the timer when the KIO3 is added. . For each trial record the time in
seconds required for the solution to change after the KIO3 is added.
5. After each trial, dispose of the reaction mixture in the sink and rinse the beaker
thoroughly with tap water.
6. You need to calculate the concentration of each chemical for each of the three trials, the
exponents for the rate law (they will be decimals, don’t round) and the value of k.
Cleanup: Take all the syringes apart, rinse with water and lay on a paper towel at your
station to dry. Rinse all the glassware twice with tap water and leave at your station to dry.
Water
1a
75mL
Starch NaHSO3 in
H2SO4
5mL
10mL
1b
75mL
5mL
10mL
10mL
2a
80mL
5mL
10mL
5mL
2b
80mL
5mL
10mL
5mL
3a
80mL
5mL
5mL
10mL
3b
80mL
5mL
5mL
10mL
Trial
KIO3 (always add last)
10mL
Conclusion:
1. Which of the steps in the proposed mechanism is the rate determining step? Why?
2. Given your experimentally determined values for the reaction orders x and y, what
would be the expected reaction time if the following solution were timed: 20mL KIO3,
5mL NaHSO3, 5mL starch, and 70mL water?
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Page 30 of 114
Where did the Crystal Violet go?
Full Writeup
Central Challenge
The purpose of this laboratory activity is to determine the rate law for the reaction
of crystal violet (CV) and sodium hydroxide (NaOH). In Part 1 of the investigation, you will
prepare dilutions of a stock CV solution to generate a Beer’s law calibration curve for CV.
In Part 2 of the investigation, you will perform a reaction of CV with NaOH while
monitoring in real time the concentration of CV remaining. This laboratory investigation will
illustrate a variety of science concepts because determining the rate law for the reaction of
CV with NaOH requires you to use graphical analysis and a simplifying approximation that
leads to a pseudo-rate law while also integrating prior chemistry knowledge involving
spectroscopy, Beer’s law, solution dilution, calibration curves, and chemical kinetics.
Context for This Investigation
If you’re making something, you might think making it to last would always be a good thing.
But what if you’re making a pesticide with known detrimental impacts on human health?
Then you may only want it to stay intact for a few days after it has been applied to crops
before it decomposes into what often are less harmful products. If its molecules stay intact
for too long, the pesticide can persist in the environment and build up in drinking water. In
2000, over 20 million kilograms of the pesticide 1,3-dichloropropene (1,3-D) were applied to
crops in the United States. Scientists investigated the rate of decomposition of 1,3-D in
acidic, basic, and neutral solutions as well as in soil. For each case, they generated plots of
the amount of intact 1,3-D persisting versus time and found that the reaction could be
characterized as pseudo first-order. Knowing the order of the reaction allowed them to
determine the half-life of intact 1,3-D. In acidic media, they found that the half-life for the
decomposition of 1,3-D was about eight days, but in the presence of excess NaOH the halflife was reduced to about four days. Experimentally determined data like this is vital to the
ability of society to use chemicals wisely in improving food production, while not
endangering the end consumers or the people who work with the chemicals during the
growing process. The Beer’s law employs the use of a colorimeter (or spectrophotometer) to
obtain a calibration curve that is used to convert raw absorption data from a colorimeter (or
spectrophotometer) to molar concentration of a chemical in solution. In this investigation,
you will first use a colorimeter (or spectrophotometer) to generate a calibration curve for a
chemical (CV) and then use the colorimeter (or spectrophotometer) to follow the change in
the concentration of CV as it reacts with NaOH. By recording these changes through time
and analyzing them graphically, you will be able to obtain the rate law of the reaction, which
may be used to predict the behavior of the system under different experimental conditions
without doing the actual experiments.
PreLab questions Day 1
1. Answer the following questions about the selection of a wavelength for your experiment.
a. Based on the absorption spectrum of 25 μM crystal violet in Figure 1 and taking into
account the considerations that follow, what wavelength should you use for the Beer’s
law calibration curve and subsequent reaction of CV with NaOH? Please explain your
answer.
b. Simulate the instrument readings you will get in Part 1 of the experiment by doing the
following: Trace Figure 1 onto your own paper. Draw a vertical line at the wavelength
you have chosen, intersecting the absorbance curve at that wavelength. Where your
vertical line intersects the absorbance curve is the absorbance value your instrument
should read for the stock 25 μM CV solution. Keeping in mind Beer’s law from
Equation 1, and being mindful that the wavelength and path length are fixed, draw X’s
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on your vertical line where you expect the absorbance values will be for the diluted
solutions you prepare in Question 2. Use appropriate ratios of concentrations to
determine where on the vertical line to make your marks.
2. A calibration curve requires the preparation of a set of known concentrations of CV, which
are usually prepared by diluting a stock solution whose concentration is known. Describe
how to prepare 10. mL of a 5-, 10-, 15-, and 20- μM CV solution using a 25 μM CV stock
solution.
3. During the reaction of CV with NaOH, do you expect the colorimeter’s (or
spectrophotometer’s) absorbance reading to change? How do you expect it to change if
such a change is anticipated (i.e., increase, decrease, or no change) as the reaction
proceeds? Explain your reasoning.
4. Answer the following questions for a reaction of CV with NaOH in these two scenarios: a
solution with a 1:1 NaOH:CV mole ratio and a solution similar to what you will be using
with a 1000:1 NaOH:CV mole ratio.
a. Using your prior knowledge of reaction stoichiometry, what is the final percentage of
each reactant remaining if each reaction went to completion? Show work and reasoning
to justify your answer.
5. Using the kinetics chapter in your textbook and websites like “Chemical Kinetics –
Integrated rate laws”
http://www.chm.davidson.edu/vce/kinetics/IntegratedRateLaws.html, describe the
graphical analysis that can be done to determine the order (considering only 0th, 1st, or
2nd order) and the value of the pseudo-rate constant, k*, of a chemical reaction from
concentration data collected through time.
6. Based on your answer to Questions 3–5, design an experiment for the reaction of CV with
NaOH and describe the subsequent data analysis to accomplish the Central Challenge, the
determination of the value of (i) w, the order with respect to CV and (ii) k*, the pseudorate constant found in the rate law in Equation 3. For simplicity, use 10. mL for the
combined volume of CV and NaOH because it is a bit more than enough to fill cuvettes
appropriately.
7. Answer the following questions after examining Figure 3 to address the issue of when to
stop collecting data.
a. For early parts of the three different reactions in Figure 3, all three curves seem
relatively linear with different slopes. But as the reactions progress through time, at
roughly what concentration level would you say some graphs start to look nonlinear?
b. Given that you don’t yet know the order of the reaction of CV with NaOH, how might
Figure 3 help you to decide when to stop collecting data? Hint: Think in terms of
percent completion instead of concentration.
Explanation to help your understanding: Read before you come to class.
For a fixed concentration of solute and a fixed path length (e.g., fixed cuvette width), the
amount of light absorbed by a solution varies directly with the absorptivity constant of the
solute. Figure 1 below shows the visible light absorbance spectrum of CV for a fixed, 25 μM,
concentration of CV and a fixed, 1.0 cm, path length. Because concentration and path length
are both kept constant, Figure 1 reveals how the absorptivity constant for CV varies with the
wavelength of light passing through the solution. Figure 1 was generated by a
spectrophotometer.
A colorimeter is an instrument which, like a spectrophotometer, measures how much light is
absorbed when passed through a sample but does so for only a few, predetermined
wavelengths of light set by the manufacturer.
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Figure 1. The visible spectrum of a 25 μM CV solution
If we still keep the path length fixed, but now choose only one particular wavelength of light to
pass through the solution, thereby fixing the absorptivity constant, students can then observe how
the absorbance of light at that wavelength changes as they change the concentration of CV.
Under these conditions, Beer’s law describes a straight-line relationship for a graph of
absorbance versus solute concentration whose slope is simply the product of the molar
absorptivity constant and path length.
In the reaction of CV and sodium hydroxide (see Figure 2), the dye’s color will fade as it reacts
with sodium hydroxide. A colorimeter (or spectrophotometer) will be used to follow the
disappearance through time of CV by measuring the absorbance of a solution of CV during its
reaction with NaOH. The raw absorbance measurements from the colorimeter (or
spectrophotometer) can be transformed to molar concentration of CV via the use of a Beer’s law
calibration curve.
Figure 2. Chemical structures in the reaction in this laboratory activity
The net ionic equation for the reaction can be written as
CV+ (aq) + OH–(aq) → CVOH (aq)
rate = k [CV+]w [OH–]z
Equation 2
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where k is the rate constant while w and z are the order of the reaction with respect to CV+ and
OH-, respectively. Under certain experimental conditions (see prelab Question 4), the rate law in
Equation 2 simplifies to the following equation:
rate = k* [CV+]w
Equation 3
where k* = k [OH–]z
Equation 4
and k* is the pseudo-rate constant. Equation 3 is referred to as the pseudo-rate law, since it is an
approximation to Equation 2, the actual rate law, and significantly simplifies the analysis.
A differential rate law describes the rate of a chemical reaction as a function of the concentration
of the reactants, while an integrated rate law describes the concentration of a reactant as a
function of time; both types of rate laws are related to each other by the use of calculus. Equation
3 is a differential rate law, in which a graphical analysis of the corresponding integrated rate law
can be used to determine the value of the parameters in Equation 3 using least-squares linear
regression analysis. The degree or extent of linear fit may be evaluated using the coefficient of
determination (or square of the correlation coefficient), i.e., it may be used to identify the graph
that has a linear relationship.
Figure 3 shows concentration data plotted versus time for three different hypothetical chemical
reactions. From plots like these and knowledge of integrated rate laws found in your text or at
online resources, one can determine the exponents in the rate law equation.
Figure 3. Concentration data plotted versus time for three different hypothetical chemical
reactions — 0th order (blue line), 1st order (red line), and 2nd order (yellow line)
All reactions have the same numerical value for their initial reactant concentration and the rate
constant.
Procedure
The prelab questions guide you to consider different factors involved in designing an
experiment to address the Central Challenge, the determination of the rate law given in
Equation 2, rate = k [CV+]w [OH-]z, for the reaction between CV and NaOH. Prelab Question
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+ w
4 addresses the issue of how to ensure that Equation 3 (rate = k* [CV ] ) is a valid
approximation to Equation 2.
The analysis described in prelab Question 6 requires that you know the concentration of CV
throughout the course of the reaction. The concentration of CV can be obtained from raw
absorbance data by applying the Beer’s law calibration curve formula you obtained
previously.
Prelab Question 6 asks you to design an experiment to determine the value of w and k* found
in Equations 3 and 4. Both w and k* can be determined by making appropriate plots of your
data from the reaction of CV with NaOH and checking for linear relationships.
Use your answer to prelab Question 8 to decide during the experiment when to stop
collecting absorbance data to get the clearest distinction between 0th-, 1st-, and 2nd-order
reactions during your postlab graphical analysis.
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Chemical Equilibrium: Finding a Constant, Keq or Kc
Full writeup
Introduction
The purpose of this lab is to experimentally determine the equilibrium constant, Kc, for the
following chemical reaction:
Fe3+(aq) +
iron(III)
 FeSCN2+(aq) (red)
SCN–(aq) 
thiocyanate
Iron III thiocyanate complex
When Fe3+ and SCN- are combined, equilibrium is established between these two ions and
the FeSCN2+ ion. In order to calculate Kc for the reaction, it is necessary to know the
concentrations of all ions at equilibrium: [FeSCN2+]eq, [SCN–]eq, and [Fe3+]eq. Since the
complex is intensely colored, its concentration is conveniently measured using a
spectrophotometer. You will prepare several mixtures containing different initial amounts of
iron III nitrate and potassium thiocyanate and will then use a spectrophotometer to measure
the absorbance as an indication of the concentration of the red product. From the initial
concentrations of the reactants taken in each mixture and the concentration of the product
present at equilibrium in each mixture, you can calculate the concentration of the reactants at
equilibrium and thus the Kc.
Prelab: Read the entire lab procedure, and use a graphing program to build and print a
calibration curve as the procedure directs you to do. To enter the lab you must place your
printed graph on top your head and whisper “Doctor Who” to get in
MSDS:
Fe(NO3)3 Corrosive to body tissues by contact and inhalation. Avoid contact with skin,
eyes and mucous membranes. Acidified with Nitric acid.
KSCN
Slightly toxic by ingestion. Irritating to body tissues. Avoid all body tissue
contact. Contact with acids or heat may liberate poisonous hydrogen cyanide
gas.
Procedure
Clean your test tubes, beakers and syringes and rinse with distilled water. Label the test
tubes with a sharpie #1-5 and the beakers Fe3+ and SCN-. Obtain approx 35 mL of the
0.002M iron III nitrate solution and 25 ml of the 0.002M potassium thiocyanate solution in
the beakers.
Fill your test tubes according to the following table: Add Fe3+ and SCN-, rinse grad cylinders
well and then use syringes to add water to tubes.
Test
tube
mL Fe3+
mL SCN-
mL distilled
water
1
5.00
1.50
3.50
2
5.00
2.00
3.00
3
5.00
2.50
2.50
4
5.00
3.00
2.00
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5
5.00
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3.50
1.50
Calibrate the spectrophotometer as you have done previously using distilled water as a blank.
Record the absorbance for each test tube at a wavelength of 450 nm, transferring the sample
to the cuvette and rinsing the cuvette between readings.
Prepare a standard solution with 9 mL of 0.200 M Fe(NO3)3(aq) (note this is different than
before) and 1 mL 0.002 M KSCN. Record the absorbance of that solution.
Using the following data, construct a calibration curve. It must be a scatter plot with a line of
best fit. There is a podcast that illustrates the process using excel.
[Iron III thiocyanate complex] M
Absorbance
3.088 x 10-5
0.152
6.176 x 10-5
0.307
9.264 x 10-5
0.443
1.235 x 10-4
0.587
1.544 x 10-4
0.752
1.853 x 10-4
0.891
1. Find the concentration of the Iron III thiocyanate complex for tubes 1-5 using the
calibration curve or the following equation:
Abs(sample) x 0.000200 M = [complex ion]
Abs(standard)
2. Calculate the initial concentrations of [Fe3+] and [SCN-] present in each of the 5
tubes.
3. Calculate the equilibrium concentrations of [Fe3+] and [SCN-] present in each of the 5
tubes.
Calculate a Keq for each tube; calculate an average Keq and the standard deviation.
Record your values for Keq for each tube on the class data sheet.
Cleanup: Throw away any disposable pipets. Wash all glassware with soap, rinse 2X with
tap water and once with distilled water. Put all glassware at your lab station to dry.
Conclusion: Compare your data to the class average data. Discuss conditions under which
the Keq would be different than the ones you found in lab using the same quantities as
identified in the lab.
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Entropy of a Reaction
NO WRITEUP REQUIRED
Discussion:
An endothermic process or reaction absorbs energy in the form of heat (endergonic processes
or reactions absorb energy, not necessarily as heat). Examples of endothermic processes
include the melting of ice and the depressurization of a pressurized can. In both processes,
heat is absorbed from the environment. You could record the temperature change using a
thermometer or by feeling the reaction with your hand. The reaction between citric acid and
baking soda is a highly safe example of an endothermic reaction, commonly used as a
chemistry demonstration. Do you want a colder reaction? Solid barium hydroxide
octahydrate reacted with solid ammonium thiocyanate produces barium thiocyanate,
ammonia gas, and liquid water. This reaction gets down to -20°C or -30°C, which is more
than cold enough to freeze water. It's also cold enough to give you frostbite, so be careful!
Here's what you need to use this reaction:
15g barium hydroxide octahydrate
9g ammonium thiocyanate (or could use ammonium nitrate or ammonium chloride)
50-ml beaker
Neutral litmus paper
stirring rod
Prelab:
1. Read the discussion and procedure. Be prepared to answer these questions.
2. Explain how this reaction increases entropy.
3. Is the change in enthalpy for the reaction going to be positive or negative? Explain
your reasoning.
4. Why are you using litmus?
Procedure
1. Pour the barium hydroxide and ammonium thiocyanate into the beaker.
2. Stir the mixture.
3. The odor of ammonia should become evident within about 30 seconds. If you hold a
piece of dampened litmus paper over the reaction you can watch a color change showing
that the gas produced by the reaction is basic.
4. Liquid will be produced, which will freeze into a slush as the reaction proceeds.
5. If you set the flask on a damp block of wood or piece of cardboard while performing the
reaction you can freeze the bottom of the flask to the wood or paper. You can touch the
outside of the flask, but don't hold it in your hand while performing the reaction.
6. After the demonstration is completed, the contents of the flask can be washed down the
drain with water. Do not drink the contents of the flask. Avoid skin contact. If you get
any solution on your skin, rinse it off with water
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Catalytic Converter—Hot Copper Catalysis
Discussion
Catalytic converters are used in automobiles to reduce emissions of unburned hydrocarbons. These unburned
hydrocarbons result from incomplete combustion of the gasoline fuel. Catalytic converters contain rare
metals such as platinum and rhodium that catalyze the combustion of hydrocarbons with the oxygen that
remains in the engine’s exhaust stream. This combustion generates heat that, in turn, makes the continued
combustion of the unburned hydrocarbons more efficient.
In this demonstration you will observe the catalyzed combustion of a hydrocarbon, namely 2-propanol,
sometimes called isopropyl alcohol, on the surface of a metal catalyst. A hot copper penny is used to provide
the reaction surface.
2-propanol can react with oxygen to form acetone and water according to the following equation:
2CH3CHOHCH3 + O2  2CH3COCH3 + H2O
(equation 1)
This reaction is thermodynamically permitted but very slow even at elevated temperatures. To make the
reaction go a catalyst is needed. Hot copper(II)oxide will do the trick. When a pre-1982 penny (almost
entirely pure copper) is heated in a burner flame copper(II)oxide forms on the surface of the penny according
to the following equation:
2Cu + O2 + heat  2CuO
(equation 2)
When the hot copper(II)oxide is introduced into 2-propanol vapor with oxygen present the following
exothermic reaction occurs:
CH3CHOHCH3 + CuO  CH3COCH3 + Cu + H2O + heat
(equation 3)
The liberated copper metal reacts further with oxygen to make copper(II)oxide which reacts further with 2propanol and so forth and so on. The heat generated by the reaction described by equation 3 is adequate to
provide the heat needed by the reaction described by equation 2. This series of reactions proceed at an
acceptable rate as long as the temperature of the penny is high enough.
Pre-Lab Questions
1. Many states conduct periodic emissions tests on vehicles. These tests should not be conducted when the
engine is cold but only after the engine has been started and allowed to warm to normal operating
temperature. Speculate on why this is so.
2. One of the first devices used to combat automobile pollution was an “air pump.” This pump would
deliver outside air directly into the exhaust manifold of a running engine. Speculate on how this helped to
reduce the emission of unburned fuel molecules.
Materials
15mL 2-propanol (off-the-shelf rubbing alcohol), 250mL beaker and a watch glass to cover, Bunsen burner,
striker, pre-drilled copper penny (pre-1982), 6” 14 gauge copper wire, hotplate with ceramic top,
thermometer, wire mesh, ruler.
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Procedure
Safety note: A fire is very unlikely if the 2-propanol is kept away from open flames. Just to be
safe, fill a 250mL or larger beaker with water to use as an emergency fire extinguisher if flames
break out.
1. Prepare the reaction chamber according to the following diagram. You will need the 250mL beaker, the
penny, the copper wire, the wire mesh, and a ruler.
The penny is to be suspended 2cm above the bottom of the beaker. The copper wire is directed thru the wire
mesh and bent at a 90o angle to hold the penny in the proper location.
2. Remove the wire mesh (with the penny and copper wire attached) from the beaker. Put the beaker on a
hot plate and add 15mL of 2-propanol to the beaker. Turn the hot plate on a very low setting to allow the
2-propanol to warm. Cover with the watch glass and monitor the temperature with a thermometer. Do not
allow the liquid to exceed 50oC.
3. While the 2-propanol is warming light a Bunsen burner and begin heating the penny. Use tongs to hold
mesh-wire-penny apparatus in place during the heating. The goal is to heat the penny red hot.
4. Once the penny is red hot AND the 2-propanol is heated to 45o to 50oC, remove the beaker from the
hotplate and suspend the penny in the vapor space over the warmed 2-propanol.
5. The reaction should begin immediately. Record all observations. The reaction is more obvious if
observed in a darkened room. The reactions produce enough heat to keep the penny red hot for 2 to 3
minutes.
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Equilibrium and Le Châtelier's Principle
Full writeup
Discussion:
Le Châtelier's Principle states that: If an equilibrium system is subjected to a stress, the
system will react to remove the stress. To remove a stress, a system can only do one of two
things: form more products using up reactants, or reverse the reaction and form more
reactants, using up products. In this experiment you will form several equilibrium systems.
Then, by putting different stresses on the systems, you will observe how equilibrium systems
react to a stress.
Before you carry out each section, write an equation and predict which way you think the
equilibrium will shift when a given step is performed. Record your predictions in a data
table before entering lab. Then, carry out the reaction to verify your prediction. Be very
precise when you record your observations.
Prelab: Read all lab directions before coming to lab. Write a purpose, procedure and build
a data table including your reactions and prediction for each part of the procedure IN YOUR
LAB NOTEBOOK.
MSDS:
NaCl
Very slightly toxic by ingestion. Dust may cause minor irritation to
mucous membranes upon inhalation.
HCl
Highly toxic by inhalation and ingestion. Severe corrosive to all body
tissues, especially skin and eyes. Avoid all body contact.
Bromothymol blue Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated.
NaOH
Highly toxic by ingestion, inhalation, or skin absorption. Extremely
corrosive to body tissues. Causes severe eye burns. Avoid all body
tissue contact.
Corrosive to body tissues by contact and inhalation. Avoid contact with
Fe(NO3)3
skin, eyes and mucous membranes.
KSCN
Slightly toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact. Contact with acids or heat may liberate poisonous
hydrogen cyanide gas.
Ethanol
Toxic by ingestion and inhalation. Body tissue irritant. Avoid all body
tissue contact. Denatured with isopropanol and methanol. Not for
human consumption.
Flammable liquid.
Irritant to body tissues. Moderately toxic by ingestion. Prolonged
exposure may cease production of red blood cells. Avoid ingestion,
inhalation and skin absorption. Cobalt and cobalt compounds are
possible carcinogens.
Moderately toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact.
CoCl2
AgNO3
Procedure
1. Equilibrium in a Saturated Solution
You will investigate the equilibrium in saturated sodium chloride solution:
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1. Obtain some saturated NaCl solution from the solution in the beaker on the stirrer in
the fume hood. To this saturated solution of NaCl, add some Cl- ions in the form of
concentrated HCl Record the results.
2. An Acid-Base Indicator Equilibrium
Acid-base indicators are large organic molecules that can gain and lose hydrogen ions to
form substances that have different colors. The reaction of the indicator bromothymol blue
can be Illustrated as follows:
HIn(aq)  H+(aq ) + In-(aq)
yellow
blue
In this reaction HIn is the neutral indicator molecule, and In- is the indicator ion after the
molecule has lost a hydrogen ion. Equilibrium reactions can easily be forced to go in either
direction. At this point you have hopefully been reading along. The code word is “monkey”
Reactions like this are said to be reversible.
1. Fill a small test tube about half-full of distilled water. Add several drops of
bromothymol blue indicator solution. Add 5 drops of 0.1 M HCl and stir. Note the
color of the indicator.
2. Next add 0.1 M NaOH drop by drop with stirring until no further color change
occurs. Again, note the color. See if you can add the right amount of acid to this test
tube to cause the solution to be green in color after it is stirred (half of the indicator is
blue and half is yellow).
3. A Complex Ion Equilibrium
An equilibrium system can be formed in solution with the following ions:
Fe3+(aq) + SCN-(aq)  FeSCN2+(aq)
colorless colorless
red-brown
3+
The iron ion (Fe ) and the thiocyanate ion (SCN-) are both colorless; however, the ion that
forms from their combination, the FeSCN2+ ion, is colored a dark red-brown. It is the color of
this ion that will indicate how the equilibrium system is being affected.
1. Pour about 25 mL of 0.0020 M KSCN solution (a source of SCN- ion) into a beaker.
Add 25 mL of distilled water and 5 drops of 0.20 M Fe(N03)3 solution. Swirl the
solution and note the following: the color of the KSCN solution, the color of the
Fe(NO3)3 solution, and the color of the resulting complex ion.
2. You will stress the equilibrium system that has resulted in several ways. Pour equal
amounts of the solution from the beaker into four test tubes. The solution in the first
test tube will be the reference solution.
3. To the second test tube add 2-3 crystals of solid KSCN. Describe the results.
4. To the third test tube add 6 drops of Fe(N03)3 solution. Stir and describe the results.
5. To the fourth test tube add drops of 0.1M NaOH, a few at a time. Stir and note the
results.
4. An Equilibrium with Cobalt Complex Ions
In this section we will investigate the equilibrium between two different complex ions
of cobalt. The reaction is endothermic:
Co(H2O)62+(aq) + 4 Cl-(aq) CoCl42-(aq) + 6 H20(l)  H = +50 kJ/mol
Pink
blue
1. Measure about 10 mL of ethanol into a beaker.
2. Examine solid cobalt (II) chloride, noting both its color and the formula of the
compound. Dissolve a small amount of cobalt (II) chloride (about half the size of a
pea) in the beaker of ethanol. The solution should be purple. If it is pink, add a little
concentrated HCl until it is purple.
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3. Put about 2 mL of the alcoholic cobalt solution into each of three small test tubes. To
one of the test tubes, add 3 drops of distilled water, one drop at a time with stirring,
noting what happens with each drop. Add 3 drops of distilled water to each of the
other two test tubes. Make a note of the effect of this stress on the system.
4. The first test tube is the control. To the second test tube, add 5 drops of concentrated
HCl, 12 M, one drop at a time with stirring. Note the results.
5. To the third test tube add a few crystals of solid sodium chloride. Stir and note the
results.
6. Put the remainder of alcoholic cobalt solution from the beaker into the fourth and fifth
test tube. To the fourth test tube add 10 drops of 0.1 M silver nitrate solution, one
drop at a time. Silver and chloride ions combine to form a precipitate of AgCl. Note
the color of the solution as the chloride ions precipitate. You may wish to let the
precipitate settle to observe the solution color more easily.
7. Immerse the fifth test tube in some hot water (about 60° C) and record any color
change.
8. Lastly, chill the fifth test tube in an ice bath to see if the color change in the previous
step is reversible.
Conclusion:
Explain each of your observed results in terms of Le Chatelier’s principle and
equilibrium.
Explain the effect of the temperature change on the cobalt equilibrium taking into
account the H listed.
No % error or error analysis.
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Strong Acid Strong Base Titration
Data only
Discussion:
One method a chemist can use to investigate acid-base reactions is a titration. A pH
titration is performed by either adding small, accurate amounts of standard base to an acid of
unknown concentration or adding small, accurate amounts of standard acid to a base of
unknown concentration. In this lab, you will add standard base to an unknown solution of
hydrochloric acid to determine its molarity. The pH is recorded methodically and is plotted
versus the volume of base added to the acid solution. The equivalence point occurs when the
acid and base in solution are stoichiometrically equivalent. This equivalence point will be
very useful in determining the concentration of the acid.
Prelab1. Read the lab procedure.
2. Build a data table to collect data.
3. On a separate sheet of paper which will be your ticket into lab calculate the pH of a
solution of 20.0 mL of 1.5 M HCl and 15.0 mL of 1.0 M NaOH to which 65 mL of
water has been added. Show your work. Define end point, equivalence point and
strong acid.
MSDS:
HCl
Highly toxic by inhalation and ingestion. Severe corrosive to all body tissues,
especially skin and eyes. Avoid all body contact.
NaOH
Moderately toxic by ingestion and skin absorption. Corrosive to body tissues.
Causes severe eye burns. Avoid all body tissue contact.
Procedure:
1. Turn the pH meter on
2. (Optional) Standardize the pH meter with pH 4 and pH 7 buffers. Your instructor can
show you how this is done.
3. Using a 25 mL volumetric pipet, transfer exactly 25mL of unknown HCl solution into a
250 mL beaker and dilute it with approximately 75 mL of distilled water.
4. Wash a 50mL buret by rinsing three times with about 5 mL of the 0.100 M NaOH
solution.
5. Fill the buret past the 0.00 mL mark with the 0.100 M NaOH solution. Drain the excess
into a waste beaker until the buret reads below 0. Be sure all air is removed from the tip
of the buret.
6. Clamp the buret in place using a buret clamp and ring stand.
7. Immerse the tip of the pH probe in the HCl solution and stir the solution. Allow the pH
reading to stabilize.
8. Record the pH of the solution at the initial volume in the first row of the data table.
9. Begin adding the NaOH solution from the buret in ~1 mL intervals. After each addition
of base, allow the pH reading to stabilize, and record the pH and the volume from the
buret on the data sheet. As the titration approaches the equivalence point the pH will
begin to increase more rapidly. Once you observe this reduce the volume of the NaOH
additions to 0.5mL then 0.2mL until you reach a pH of about 8 to 9. Above pH 9, return
to the 1 mL additions.
10. Continue the titration until all 50.0 mL of NaOH has been added or a pH of 12 is
obtained.
11. To clean up, dump all solutions down the drain with lots of water
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12. Prepare a graph using excel on the laptops of pH vs. mL of NaOH. You will save this
on your U drive for the next lab.
13. From the graph, determine the equivalence point.
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Titration of a weak acid
Data, Graph and Calculations Only
Introduction:
Vinegar is composed mostly of acetic acid and water. In this experiment, the amount of
acetic acid will be found by means of a titration. In a titration, the concentration of an
unknown substance can be found by comparing it to another substance of known
concentration. In this titration, a solution of sodium hydroxide will be added to a vinegar
solution to determine the percent of acetic acid in the vinegar.
Often a titration endpoint is marked by the change in color of an indicator. Although this
method is widely used, it is not always necessary. In many cases, a pH meter can be used to
determine the endpoint. The pH of the solution is monitored as titrant is added. When the
pH has passed through a region of rapid change and then leveled off, the data collected is
graphed. The end point is determined graphically, and other information may also be read
from the shape of the graph.
MSDS:
Vinegar
Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated. Not for human consumption.
NaOH
Moderately toxic by ingestion and skin absorption. Corrosive to body tissues.
Causes severe eye burns. Avoid all body tissue contact.
Procedure:
1.
The strong acid strong base titration was a good warmup to this lab. It is critical you get
the concentration correct. Assume all of your glassware is dirty and needs to be washed
and rinsed. Clean a 50 mL buret and rinse three times with about 5 mL of the 1.0 M
NaOH solution.
2. Fill the buret past the 0.00 mL mark with the NaOH solution and deliver the excess into
a waste beaker until the volume is at or below the 0.00 mL mark. Make sure that the
buret tip has no air bubbles.
3. Pipet 25.00 mL of vinegar into a 250 mL beaker and add 40 mL of distilled water from a
graduated cylinder.
4. Stir the vinegar/water solution with the pH meter and allow the reading to stabilize.
5. Record the pH of the solution at the initial volume (0.00 mL of 1.0M NaOH) in the first
row of the data table.
6. Begin adding the NaOH solution slowly from the buret, stopping at frequent intervals to
record the buret reading and the pH on the data sheet. Collect data at intervals of ~0.2
pH units or 2 mL, whichever occurs first. Be particularly alert as you pass pH 5.3 as the
pH begins rising quite rapidly just beyond this point.
7. Continue the titration until about 50.00 mL of NaOH has been added from the buret OR
the pH has leveled off to a very slow increase (above pH=11).
8. Prepare a graph (same graph as before) using excel on the laptops of pH vs mL of
NaOH. It should be the same graph you generated for titration of a strong acid. You
will end up with two lines, one for the strong acid and one for the weak acid. Be sure
each line is a different color. Submit your graph to edmodo. Be sure to submit the
graph for each person. You must include the graph in your lab writeup.
9. From the graph, determine the equivalence point. Using the mL of NaOH at the
equivalence point, use the equation MAVA=MBVB..
10. Calculate the pKa of acetic acid using your graph. Look up the accepted pKa of acetic
acid and calculate a percent error.
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Determination of the Ka of Weak Acids
Full writeup
Discussion:
Acids are substances that donate hydrogen ions. They vary in their ability to ionize from
strong acids which ionize 100% to extremely weak acids which hardly ionize at all. The
dissociation constant is the value of the equilibrium constant which indicates the acid’s
strength.
The modern Bronsted definition of an acid relies on the ability of the compound to donate H+
to other substances. When an acid dissolves in water, it donated hydrogen ions to water
molecules to form the hydronium ion. It is shown by the general equation
HA(aq) + H2O(l) ↔A-(aq) + H3O+(aq).
The Keq for an acid is called the Ka. The equation for Ka is as follows.
[ A  ][ H 3O  ]
Ka 
[ HA]
Not all acids are created equal. The Ka for a strong acid is extremely large. The Ka for a
weak acid is much less than one. Polyprotic acids contain more than one ionizable hydrogen.
Ionization of a polyprotic acid occurs in a stepwise manner, where each step is characterized
by it own equilibrium constant (Ka1, Ka2, Ka3). The second removal of a hydrogen always
occurs to a much smaller extent than the first and so Ka2 is always significantly smaller than
Ka1.
Acid
Formula
Ka1
Ka2
pKa1
pKa2
Iodic
HIO3
1.7 x 10-1
0.77
Sulfurous
H2SO3
1.7 X 10-2
6.4 x 10-8
1.77
7.19
-5
Acetic
HC2H3O2
1.8x10
4.74
Carbonic
H2CO3
4.3 x 10-7
5.6x10-11
6.37
10.25
-8
Hypochlorous HClO
3.0x10
7.52
Hydrocyanic HCN
4.9x10-10
9.31
Table 1
The Ka can be determined experimentally by measuring the hydronium ion concentration in
a dilute solution of the weak acid. This procedure is most accurate when the solution
contains equal molar amounts of the weak acid and its conjugate base. Consider acetic acid
as an example, acetic acid and the acetate anion represent a conjugate acid-base pair. Write
the Ka expression for acetic acid below.
If the acetic acid concentration and acetate ion are equal, they can be canceled out in the
equation. Write the expression for Ka with the canceled terms removed below.
In this experiment, solutions are prepared in which the molar concentrations of an unknown
acid and its conjugate base are equal. The pH of these solutions are then equal to the pKa.
Most of the unknowns are salts of polyprotic acids that still contain an ionizable hydrogen.
Sodium bisulfate, for example, is a weak acid salt. The HSO4- ion is a weak acid. The Keq
for ionization of HSO4- corresponds to Ka2 for sulfuric acid.
Write the equilibrium expression for both deprotonations of sulfuric acid below as well as the
reversible equations.
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Prelab: Phosphoric acid is triprotic. Ka1 = 7.5 x10-3, Ka2=6.2x10-8 and Ka3=4.2x10-13.
1. Write the equation for the first ionization reaction of phosphoric acid with water.
2. Write the equilibrium constant expression, Ka1, for this reaction.
3. What would the pH of a solution be when [H3PO4] = [H2PO4-]?
4. Phenolphthalein would not be an appropriate indicator to use to determine Ka1 for
phosphoric acid. Why not? Choose a suitable indicator from the following color
chart.
pH
Indicator
1
2
3
4
5
6
7
8
9
10
11
Phenolphthalein Colorless
Pink
Red
Methyl Red
Red
Orange
Yellow
Orange IV
Orange
Peach
5. What would be the pH of a solution prepared by combining equal quantities of
sodium dihydrogen phosphate and sodium hydrogen phosphate?
6. Sufficient strong acid is added to a solution containing Na2HPO4 to neutralize one
half of it. What will be the pH of this solution? Explain.
NOTE: The lab manual must be filled in to this point and shown to the instructor before you
are allowed in the lab.
Procedure
1. Label two weigh boats #1 and #2.
2. Obtain an unknown weak acid and record the unknown designation in your data table.
3. Measure out a small quantity (0.15 – 0.20g) of the unknown into each weigh boat. It
is not necessary to know the exact mass.
4. Using a graduated cylinder, precisely measure 50.0 mL of distilled water into a 150
mL beaker.
5. Transfer the sample from weigh boat #1 to the water in the beaker and stir to dissolve.
6. Using a graduated cylinder, precisely transfer 25.0 mL of the acid solution prepared
in step 5 into an 125 mL Erlenmeyer flask.
7. Add 3 drops of phenolphthalein to the acid solution.
8. Using a pipet, add 0.1 M sodium hydroxide solution dropwise to the flask. Gently
swirl the flask while adding the sodium hydroxide solution.
9. Continue adding sodium hydroxide dropwise and swirling the solution until a faint
pink color persists throughout the solution for at least 5 seconds. This is called the
endpoint. At this point the solution in the beaker contains ½ the original amount of
acid, essentially all of which is in the acid form, HA. The Erlenmeyer flask contains
an equal amount of the conjugate base A-.
10. Add the contents of the beaker to the contents of the flask. Now the acid and
conjugate base concentration are the same.Record the pH using the pH meter.
11. Pour the contents down the sink. Rinse, rinse and distilled rinse the glassware.
12. Repeat the procedure with the contents of weigh boat #2.
13. Obtain two samples of a second, different unknown weak acid and determine the pH
using the procedure above.
Calculations:
Average the pH readings for trials #1 and #2 for each unknown and calculate the average
pKa.
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The following table lists the identities of the unknowns in this experiment. Complete the
table by calculating the pKa for each acid.
Weak acid
Formula
Ka
pKa
-8
Potassium dihydrogen
KH2PO4
6.2 x 10
phosphate
Potassium hydrogen sulfate
KHSO4
1.0 x 10-2
Potassium hydrogen phthalate
KHC8H4O4
3.9 x 10-6
Potassium hydrogen tartrate
KHC4H4O6
4.6 x 10-5
Acetylsalicylic acid
2-CH3CO2C6H4COOH
3.2 x 10-4
Conclusion:
Identify the unknown acids based on pKa. Comment on the precision of the pKa
determinations.
Questions
1. Why was it not necessary to know the exact mass of each acid sample?
2. Why was it not necessary to know the exact concentration of the sodium hydroxide
solution?
3. Why was it necessary to measure the exact volume of distilled water used to dissolve
the acid as well as the exact volume of solution transferred from the beaker to the
Erlenmeyer flask?
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Determination of the Ksp of an Ionic Compound
Data and Calculations and error comparison
Discussion:
An equilibrium constant is a measure of how far a reaction proceeds to completion. A Ksp is
an equilibrium constant that permits the calculation of the amount of a slightly soluble ionic
compound that will dissolve in water. The equilibrium exists between the aqueous ions and
the undissolved solid. At this point it is considered saturated with that particular solute.
Take lead II chloride as an example. The equation for lead chloride dissolving is
PbCl2(s) ↔Pb2+(aq) + 2Cl-(aq).
Write the solubility product expression here:
Knowledge of the Ksp of a salt is useful in determining the concentration of ions of the
compound in a saturated solution. This allows you to control a solution so that precipitation
of a compound will not occur, or to find the concentration needed to cause a precipitate to
form. In this experiment, you will determine the Ksp of calcium hydroxide.
Prelab:
The compound silver chromate is not very soluble in water.
1. Write the equation for the dissociation of silver chromate.
2. Write the Ksp expression
3. If one has a solution of 0.10 M silver nitrate and it is diluted by a factor of two, what is
the new concentration?
4. The dilution of 0.10 M silver nitrate by a factor of two is carried out 5 times. What is the
concentration now?
5. The value for the Ksp of silver chromate is 1.1 x 10-12. In a saturated solution of silver
chromate, the [Ag+] is found to be 2.5 x 10-4 M. What must the chromate ion
concentration be? Show your work.
Procedure:
1. Arrange a microplate so that you have 12 wells across from left to right.
2. Put 5 drops 0.10 M calcium nitrate in well #1 in the first row. Hold the pipet vertically
when dispensing. Make sure no bubbles are in the pipet. Discard the first drop as it may
contain an air bubble.
3. Place 5 drops of distilled water in each of the wells #2 through #12 in the first row.
4. Next, add 5 drops of 0.10 M calcium nitrate to well #2.
5. Use an empty transfer pipet to mix the solution by sucking in the solution and blowing it
out several times. The solution in this well, #2, is now 0.050 M in Ca2+ ion.
6. Use your empty pipet to remove the solution from well #2 and put 5 drops in well #3.
7. Put the remaining solution in the pipet back in well #2.
8. Mix the solution in well #3 as before.
9. Continue this serial dilution procedure, adding 5 drops of the previous solution to the 5
drops of distilled water in each well down the row until you fill the last one, #12.
10. Mix the solution in well #12 and discard five drops.
11. Determine the concentration of solution in each well.
12. Place 5 drops of 0.10 M NaOH in each of the wells. When the NaOH is added to each
well, the initial concentrations of the reactants are halved, as each solution dilutes the
other.
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13. Use an empty pipet to mix each of these combined solutions by drawing each solution
into the pipet and squirting it back into the well.
14. Wait 4 minutes for ppt to form.
15. At one point the concentration of both ions becomes too low to have any ppt form.
Assume that the first well with no precipitate is a saturated solution.
16. Repeat the steps except using a serial dilution of 0.10 M NaOH in each well and then
adding 5 drops of 0.10 M Ca(NO3)2 to each well.
Find the first well with no precipitate and record that well as the [Ca2+]. The [OH-] will be
2x that of the calcium ion. Plug those values into the Ksp expression and find a value for
Ksp of calcium hydroxide.
Find the first well with no precipitate from step 16 and record that as the [OH-]. The [Ca2+]
will be half of the [OH-]. Plug those into the Ksp expression and find the Ksp of calcium
hydroxide. Compare the two results.
Error analysis. Search the web for the accepted Ksp value of calcium hydroxide and
compare to your value.
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Buffer Laptop Palooza
All work and answers on your own paper.
Follow the link on edmodo OR
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/acidbasepH/
pHbuffer20.html
Activity 1. What is the pH of a 0.10 M HC3H5O3(aq), lactic acid, solution? Ka = 1.4 x 10-4.
pH = __________. (Show ice table)
Write a chemical equation that illustrates what happens when pure HC3H5O3 is placed in
water.
What is the pH of a 0.10 M NaC3H5O3(aq), sodium lactate, solution? pH =
_________________.
Write a chemical equation that illustrates what happens when solid NaC3H5O3 is placed in
water.
Without using the computer simulation , predict what happens to the pH of 0.10 M
HC3H5O3(aq) solution when enough NaC3H5O3(s) is added so that the initial concentration of
the NaC3H5O3(aq) is 0.10 M?
The pH of the solution [increase decrease no change] ?
Use the computer simulation to check your prediction. Was your prediction correct?
Explain why the pH changes when NaC3H5O3(aq) is added.
Use the computer simulation to add different concentrations and amounts of NaC3H5O3(aq)
and HC3H5O3(aq) to each other. Make at least two other solutions. Record what solutions you
mixed and the resultant pH.
Activity 2. Use the computer simulation to mix the following solutions. Compare the pH of
the solutions.
a. 100.0 mL 0.500 M HC3H5O3(aq) / 100.0 mL 0.500 M NaC3H5O3(aq) pH =
_____________
b. 100.0 mL 0.100 M HC3H5O3(aq) / 100.0 mL 0.500 M NaC3H5O3(aq) pH =
_____________
c. 100.0 mL 0.500 M HC3H5O3(aq) / 100.0 mL 0.100 M NaC3H5O3(aq) pH =
_____________
Activity 3. Using the computer simulation, choose two solutions that when mixed will create
a 1.0L buffer solution with the designated pH. Record what two solutions you mixed, what
the concentrations were, and the amounts.
a. pH = 4.74 (using acetic acid as one of the components) Ka = 1.8 x 10-5
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b. pH = 5.00 (using acetic acid as one of the components)
c. pH = 9.25. (using ammonia as one of the components) Kb = 1.8 x 10-5
d. pH = 8.00 (using ammonia as one of the components).
Activity 4. Choose one of the solutions you created in Activity 3. Test this solution to see if
it is a buffer solution by going to "Part II" of the program. How will you know that it is a
buffer?
Write an equation that shows what happens when acid is added to your solution.
Write an equation that shows what happens when base is added to your solution.
Activity 5. Calculate the pH of a solution created by mixing 200.0 mL of 0.400 M acetic
acid and 200.0 mL of 1.00 M sodium acetate. Hint: The instant the two solutions are mixed
what are the initial concentrations of each? Use the computer simulation to confirm your
calculation. Build an ice table below to support your calculation.
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Preparation and Properties of Buffer Solutions
Fill in the blank spots in the discussion and prelab prior to lab.
Calculations and sample provide the basis for your grade. No writeup.
Discussion:
Buffered solutions are solutions that change pH only slightly with the addition of an acid or
base. An important example of a buffered solution is human blood. Many biochemical
processes take place within a narrow range of pH values, so it is important to maintain a
constant blood pH. In blood, this pH range is between 7.35 and 7.45. The main buffer
system in blood is the carbonic acid –hydrogen carbonate buffer system. Carbonic acid,
H2CO3, and the hydrogen carbonate ion, HCO3-, are an acid base pair. Fill in the missing
parts of the equation.
H+(aq) + ___________ ↔H2CO3(aq)↔ H2O(l) + __________
When the body removes CO2 by respiration, the equilibrium shifts to the right and
hydronium ions are absorbed. In normal blood the ratio of bicarbonate to carbonic acid is
20:1. This means that the buffer has a high capability to neutralize additional acid, but little
ability to neutralize additional base.
In order to accomplish the feat of acting as an acid or a base, a buffer must contain both an
acidic and a basic component. These two components should not neutralize each other, but
be available to neutralize hydrogen or hydroxide ions from other sources. One way to carry
this out is to combine a weak acid-base conjugate pair, such as acetic acid and acetate ions,
or ammonium ions and ammonia.
Acetic acid-acetate ion buffer can be prepared in several ways. One can combine a solution
of acetic acid and sodium acetate; one can start with a solution of acetic acid and neutralize
some of it with sodium hydroxide or you can begin with a solution of sodium acetate and
partially neutralize it with hydrochloric acid. By varying the type of weak acid or base and
changing the concentration ratio of the conjugate acid-base pairs, buffers can be made for
any pH value.
Write the equation for the dissociation of a weak acid HA below.
Write the Ka expression for the equation.
Solve for [H+] in the Ka expression.
The expression shows that the [H+] is dependent on Ka and the ratio of acid and conjugate
base pair. It is possible to calculate quantities needed to prepare a solution of a known pH by
choosing an acid whose dissociation constant is somewhere near the desired [H+] and by
solving the equation to find the correct ratio of acid/conjugate base. If the concentration of
acid and conjugate base are equal then pH = pKa.
The Henderson Hasselbalch equation can be our friend here.
[base]
pH  pKa  log
[acid ]
Since Ka x Kb = 1.0 x 10-14, pKa + pKb=14
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Substitute into the Henderson-Hasselbach equation for pKa in the expression above and write
the new equation for the pH of a basic buffer.
Prelab:
1. If the pH of a solution is 3.50, calculate the [H+].
2. Fill in below with the missing conjugate acid or base.
Conjugate acid
Conjugate Base
HC2H3O2
CNHSO4
CO32-5
3. Acetic acid has a Ka of 1.8 x 10 . How many grams of sodium acetate would have to
be added to 100. mL 0.10M acetic acid to prepare a buffer with a pH of 4.50?
Procedure:
Part I
1. Place 20 mL of distilled water in a 100 mL beaker. Test the pH using a pH meter and
record.
2. Add one drop of 0.1 M HCl, stir and test the pH and record.
3. Repeat with a second and third drop of HCl.
4. Repeat steps 1-3 with 20 mL of 0.1 M NaCl instead of water.
5. Repeat steps 1-3 with 20 mL of distilled water and 0.1 M NaOH instead of HCl.
6. Repeat steps 1-3 with 20 mL of 0.1 M NaCl instead of water and 0.1M NaOH instead of
HCl.
Part II
7. To prepare a buffer use a graduated cylinder to add 10 mL of 0.1 M acetic acid and 10
mL of 0.1M sodium acetate to a 100 mL beaker.
8. Measure the pH of the buffer solution and record.
9. Add one drop of 0.1 M HCl, stir and test the pH and record.
10. Repeat with a second and third drop of HCl.
11. Prepare a fresh sample of the buffer as in step 7.
12. Repeat using drops 0.1 M sodium hydroxide.
13. Prepare a buffer by combining 10 mL of 0.1 M ammonia and 10 mL of 0.1 M ammonium
chloride in a 100 mL beaker.
14. Repeat the process of drops of HCl and NaOH as described above preparing fresh buffer
each time.
Part III
15. Weigh out 1.8 to 2.0 g of a solid acid sample given to you.
16. Dissolve the acid in 150 mL of distilled water in a 250 mL Erlenmeyer flask.
17. Pour 75 mL of this solution into a second 250 mL Erlenmeyer flask, add 2 drops
phenolphthalein and titrate with 0.2 M NaOH. Record the volume of the titrant used.
18. In the first flask you have a weak acid and in the second flask you have a solution of the
sodium salt of the weak acid (i.e. the conjugate base). Make the concentrations of the
two solutions the same by adding the same volume of distilled water to flask one as the
volume of NaOH you added to flask two.
19. Combine 10 mL of the weak acid in flask 1 with 10 mL of the conjugate base in flask 2,
mix and measure and record the pH. Calculate the pKa of the acid.
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20. I will assign you the pH of a buffer to prepare. Calculate the volume of weak acid and
conjugate base that you need to prepare 50 mL of a buffer of the assigned pH. Assume
the concentrations of the weak acid and conjugate base are equal. Thus we can assume
[base] Volumeofconjugatebase 50  x 


[acid ]
Volumeofacid
x
Write down your calculations and bring your sample to me to be tested.
Cleanup:
All glassware gets washed, rinsed.
All solutions can go down the sink with water.
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Corrosion Cells
Data only
Discussion:
When two dissimilar metals are placed into electrical contact with each other and with a
conducting solution a galvanic corrosion cell is formed. Electrons flow from the most active
(least noble) metal to the least active (most noble) metal. The most active metal is said to be
oxidized. This oxidation process results in the metal’s atoms being converted to cations
which enter the conducting solution. If the cell is allowed to continue to exist the most active
metal will eventually dissolve away due to oxidation. This uncontrolled oxidation of the
most active metal is generally referred to as galvanic corrosion. The least active metal serves
as a surface for a reduction reaction to take place. The nature of this reduction reaction is
difficult to predict without detailed knowledge of the composition of the conducting solution
and the identity of the least active metal. As the electrons flow from the most active metal to
the least active metal an electrical potential is generated. The potential generated by metals
of similar activity will be quite small while the potential generated by metals of strongly
divergent activity can be quite large. Given a voltmeter to supply the electrical connection
between the most active and least active metal we can measure the potential generated by the
corrosion cell and derive an activity series based upon experimental results. The voltmeter
will display a positive reading when the black lead is connected to the most active metal and
the red lead is connected to the least active metal. Conversely, the voltmeter will display a
negative reading when the black lead is connected to the least active metal and the red lead is
connected to the most active metal.
Procedure
Assembling and Testing the CBL Voltmeter
1. Plug in the power adapter to a 110VAC outlet and to the power connection on the CBL.
2. Connect the voltage probe into the CH1 slot on the CBL.
3. Depress the ON/HALT key on the CBL to turn on the system.
4. Depress the MODE key to initialize real time sampling on the CBL. The CBL should be
flashing a potential reading of about 2.0 volts. If this is not the case there is a problem
with the CBL system. Consult with your teacher before proceeding.
5. Touch the voltage probe leads together. The reading should drop to just a few millivolts.
If this is not the case there is a problem with the CBL system. Consult with your teacher
before proceeding.
Preparing the Voltaic Cell Kit for Measurements
6. Obtain a voltaic cell kit, remove the connector collar, and add KCl solution to the glass
container to a depth of about 1cm.
7. Obtain a set of 7 metal specimens. Note each specimen is stamped for identification: CO
(copper), brass, iron, lead, zinc, Ni (nickel), AL (aluminum)
8. Clamp the nickel specimen to one of the posts on the connector collar and set the
connector collar on the glass container. Nickel will be used as a reference for all
measurements in part 1.
9. Connect the black tip of the voltage probe to the post attached to the nickel specimen and
the red tip to the other post. The system is now ready for measurements.
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Measurements, Part 1
10. Attach the brass specimen to the post where the red voltage probe is connected. Be sure
the bottom of the brass specimen reaches the KCl solution.
11. Allow the voltage reading to stabilize and record the value on the data sheet.
12. Repeat steps 10 and 11 for each of the remaining metal specimens. Record each voltage
reading on the data sheet.
13. Based on the observed potential readings draw up an activity series, including nickel,
arranged in order from least active to most active. Write down the potential reading
associated with each metal. The value for nickel will be zero since it was used as a
reference.
Part 2, Predictions and Measurements
14. Based on the activity series developed in Part 1, predict the potential readings that would
result from lead being used as the reference coupled with each of the other metals. Write
your predictions in the table located on the data sheet.
15. Perform a second series of measurements with Lead as the reference and record the
results on the table.
Cleanup
1. Once all measurements have been completed dispose of the KCl solution and rinse and
dry the metal specimens.
2. Disassemble the CBL system and re-package the components.
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Part 1 Experimental Activity Series
Red
Black
Brass
Aluminum
Copper
Zinc
Iron
Lead
Nickel
Nickel
Nickel
Nickel
Nickel
Nickel
Page 58 of 114
Potential
Reading
Activity
Least
Active
Most
Active
Part 2
Red
Black
Brass
Lead
Aluminum Lead
Copper
Lead
Zinc
Iron
Nickel
Lead
Lead
Lead
Predicted
Potential
Reading
Measured
Potential
Reading
Metal
Potential
Reading
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Polyatomic Ions
Ammonium
Acetate
Arsenate
Borate
Bromate
Carbonate
Chlorate
Chromate
Cyanide
Dichromate
Dihydrogen phosphate
Hydrogen carbonate (bicarbonate)
Hydrogen phosphate
Hydrogen sulfate (bisulfate)
Hydroxide
Iodate
Nitrate
Oxalate
Permanganate
Peroxide
Phosphate
Silicate
Sulfate
Thiocyanate
Thiosulfate
NH4+
C2H3O2- or CH3COOAsO43BO33BrO3CO32ClO3CrO42CNCr2O72H2PO4HCO3HPO42HSO4OHIO3NO3C2O4-2
MnO4O22PO43SiO32SO42SCNS2O32-
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Molecular Geometry
Using Lewis Structures, the VSEPR and Valence Bond Theories
Physical and Chemical properties depend on the geometry of a molecule.
BP = Bonding Pairs, LP = Lone Pairs
Electron Pair Geometry: AX2(2 BP)
Molecular Geometry: Linear
Electron Pair Geometry: AX3 (3BP or 2BP + 1LP)
Molecular Geometry: AX3 Trigonal Planar
Molecular Geometry: AX2E1 Bent/Angular
Electron Pair Geometry: AX4 [4BP or (3BP + 1LP) or (2BP + 2LP)]
Molecular Geometry: AX4 Molecular Geometry: AX3E1 Molecular Geometry: AX2E2
Tetrahedral
Trigonal Pyramidal
Bent/Angular
Electron Pair Geometry: AX5 [5BP or (4BP + 1 LP) or (3BP + 2LP) or (2BP + 3LP)]
Molecular Geometry:
Molecular Geometry:
AX5
AX4E1
Trigonal
See-saw
Bipyramidal
Molecular
Geometry:
AX3E2
T-structure
Molecular
Geometry:
AX2E3
Linear
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In AX5: More electronegative atoms in the axial positions, (bond will be a bit longer), and
lone pairs and double bonds in the equatorial position, see see-saw, T- and linear-structures.
Electron Pair Geometry AX6 [6 BP or (5BP + 1 LP) or (4BP + 2LP)]
Molecular Geometry: AX6 Molecular Geometry: AX5E1 Molecular Geometry: AX4E2
Octahedral
Pyramidal Planar
Square Planar
To predict molecular geometry, find the nuclei of the atoms in three dimensional space,
after defining the electron pair geometry from eg AX3E2 as AX3+2 = AX5, remember the
distortions.
Distortions in bond angles are influenced by (1) the lone pairs on the central atom and
(2) the size of atoms, eg
H2 O
OF2
OCl2
HOH 104.5 °
FOF 103 °
ClOCl 111 °
With lone pairs on the central atom, the bond angle will not be the AX4 109.5 °.
The HOH bond will be smaller than the standard 109.5 °, because of the larger volume of the
two lone pairs on the oxygen atom, but in OF2 the more electronegative F atoms will draw
the lone pairs closer to the OF single bonds, influencing the bond angle more. The ClOCl
bond angle is measured as 111 °, larger than the expected 109 °, because of the larger
chlorine atoms, they move away from each other, repulsion of the electronic charges on the
large atoms.
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Rules of Writing Equations
Synthesis (Use Gas Ion Chart)
1. Combination of elements.
2Ag (s) + Cl2(g)  2AgCl(s)
2. A metal oxide plus water yields a base.
Na2O(s) + H2O(l)  2NaOH(aq)
3. A non-metal oxide plus water yields an acid.
SO3(g) + H2O(l)  H2SO4(aq)
N2O5(g) + H2O(l)  2HNO3(aq)
4. A non-metal oxide plus metal oxide yields a salt.
MgO(s) + CO2(g) MgCO3(s)
Decomposition (Use Gas Ion Chart)
1. A base when heated decomposes into a metal oxide plus water.

2LiOH(aq)  Li2O(s) + H2O(l)
2. An acid when heated decomposes into a non-metal oxide plus water.

2H3PO4(aq)  P2O5(s) + 3H2O(l)
3. Metallic carbonates decompose into a metal oxide and carbon dioxide.

CaCO3(s)  CaO(s) + CO2(g)
4. Metallic chlorates decompose into a metallic chloride and oxygen.
2KClO3(s)  2KCl(s) + 3O2(g)
5. Some compounds decompose with electricity or just simply decompose into their
basic elements.
electricity
2H2O(l)  2H2(g) + O2(g)
Single Replacement Reactions (Use activity series)
1. A metal will replace a less active metal in a compound.
Al(s) + 3AgNO3(aq) Al(NO3)3(aq) + 3Ag(s)
2. Some metals will replace the H in water to produce a metallic hydroxide and
hydrogen gas.
2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)
3. Some metals will replace the H in acid to produce a salt and hydrogen gas.
Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)
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4. A halogen (VIIA) will replace a less active halogen in a compound.
F2(g) + CaBr2(aq) Br2(g) + CaF2(aq)
Double Replacement Reaction States of matter are important. Use the solubility rules.
1. An acid and a base yield a salt and water.
H2SO4(aq) + 2NaOH(aq)Na2SO4(aq) + 2H2O(l)
2. A salt and an acid yield a different salt and a different acid.
HNO3(aq) + NaCl(aq) NaNO3(aq) + HCl(aq)
3. A salt and a salt yield a salt and a salt.
CuSO4(aq) + NaClO4(aq)Na2SO4(aq) + Cu(ClO4)2(aq)
4. Some compounds decompose when made in double replacement reactions
 If carbonic acid is made it decomposes into water and carbon dioxide gas.
 If ammonium hydroxide is made it decomposes into water and ammonia (NH3)
gas.
 If sulfurous acid is made it decomposes into water and sulfur dioxide gas.
Gas Ion Chart
Gas
Ion
SO2
SO32SO3
SO42CO2
CO32N2O3
NO2N2O5
NO3P2O3
PO33P2O5
PO43H2O
OHNH3
NH4+
Solubility Rules
All sodium, potassium, ammonium, and nitrate salts are soluble in water.
Activity Series of Metals
Most Active
Least Active
Li Rb K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Hg Ag Pt Au
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AP Chemistry Syllabus
Room N216
Email: [email protected]
School voicemail: 772-2547 ext 651
Home phone: 772-8044 (Please call only if absolutely necessary, no calls past 9pm)
Course Description
Chemistry is the study of the properties, composition, structure, and behavior of matter. A
detailed list of topics to be studied is described below in the "Course Content Outline." The
instructor will use a variety of teaching methods including lecture, demonstrations,
laboratory exercises, and written assignments (both in-class and homework).
Pre-entry Standards and Expectations
In order to achieve a minimum passing grade for this course, a student must be able to follow
and apply basic safety requirements, collect and analyze data, manipulate laboratory
equipment and apparatus, perform laboratory work, prepare and interpret graphs, perform
mathematical calculations using algebra, prepare written reports, communicate effectively
both orally and in writing, solve problems, read, study, and complete assignments. Students
must be prepared to devote at least as much time to study and completing assignments
outside of class as the time spent in class.
Many students find Chemistry to be the first class in which they’ve ever had to study. Most
students cannot expect to make A’s or B’s without going home nightly, looking over the
material covered and doing some self diagnostics asking themselves “Do I understand this
concept” or “Can I work this problem”. If you will identify weaknesses in your
understanding and ask the next day, you will stay current. It is also necessary to access
edmodo and take notes from the podcasts so we can make better use of time.
I am assuming you are knowledgeable about the concepts covered in Chemistry I and I will
expect you to use those skills. I will only be doing minor review on the content covered in
Chemistry I.
Supplementary Materials:
It is suggested that you purchase an AP Chemistry prep manual printed by Princeton Review
or Cliff’s. Be sure to purchase one for the 2014 redesigned curriculum. Use this as a
supplement of problems and concepts.
There will also be podcasts posted to edmodo and/or mediacast and/or youtube and/or
dropbox for you to download and watch. You can also obtain the vodcasts on a DVD to take
home and copy over to your computer or on a DVD to play in a DVD player. We will be
employing a flipped classroom approach this year. Students will be responsible for watching
video content as their homework and the majority of the time in class will be spent discussing
lab results, working problems and building mastery. This approach makes the student
responsible for their own education and moves the teacher from someone who pours
knowledge into their head to a coach or facilitator who is available to help the student in their
quest for success.
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Grading
You will be graded on chapter questions, worksheets, written laboratory activities, assigned
reports, projects, tests and in-class activities. At times, you will earn grades for doing
assignments in class. If you choose to do the assignment as instructed, you will earn a good
grade. If not, you will earn a bad grade.
 No late work or unexcused absence work will be accepted.
 Tests will be a combination of multiple choice and free response to mimic the AP exam.
A large part of Chemistry is problem solving and you will be expected to show a
proficiency in that skill.
 Your grade each nine weeks will be calculated as follows:


30% Daily Grades

Labs- lab quizzes, lab reports and conduct in lab. Labs must be done in a lab
notebook approved by the instructor. See the lab manual in print or online for
information on lab grading and format.

Effort- Being prepared, listening, participation, etc.

Homework and classwork – Your homework will be taking notes and looking at
content on podcasts. Students will be quizzed on their notes and content and
problems will be graded for accuracy in class.
70% Tests Tests will be timed. The AP exam is a timed exam and it is important for
students to develop time management skills.
 Test Corrections
In order to reinforce concepts missed on tests, students will be allowed to
come in before and after school and make corrections to the test. Students may
come in the morning and/or afternoon communicated to the class by the teacher.
Students who come in the morning may come at 7:20 and must get a pass from Mr.
Elegante the day before. Students staying after school may stay from 3:30 – 4:20.
There will not be makeup days for test corrections. In order to do test corrections,
students must come with their notes from the podcasts. A student who has not taken
the time to watch the podcasts and take notes will not get to do test corrections.
Students must provide the correct answer, a description of why the other choices are
incorrect and an explanation of why the student missed that particular question. If the
work is incomplete or is not done to the standards set by Mr. Elegante, points will not
be awarded back to the student. Students can earn ¼ of the points missed back on
their test grade. No points will be earned back on free response questions, only the
multiple choice and mathematical computations parts of a test. The student may use
their text and are encouraged to seek help from Mr. Elegante and classmates when
doing corrections. No corrections will be done at home.
 You may come see me any time before or after school to see about your grades.
 Grades will be entered into the INOW system allowing parents and students to access
grades using the internet.
 Make-ups. You need to schedule a makeup time with Mr. Elegante before or after
school. Makeups will begin at 7:15 AM before school or directly after school. It is the
responsibility of the student to ask for a hall pass for make-up tests. If a test is missed
due to an excused absence, the student will have two weeks after the absence to make it
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up. Students who arrive late for a make-up test may NOT receive extra time to finish the
test and any uncompleted sections will be counted as incorrect.
 End of Term Exam. The end of term exam is a comprehensive test given at the end of the
semester, which counts for 20% of the course grade. The final exam is a scaled down
version of an AP exam with multiple choice and free response portions as well as a lab
component.
Class Rules
In addition to the Student Code of Conduct and academic conduct policies found in the BJHS
Student Handbook, please observe the following rules:
1. Always have a notebook, a pen or pencil, an issued textbook, and a calculator.
2. Always be attentive and follow directions.
3. Always go immediately to your seat when the tardy bell rings, listen for directions,
and be ready to begin work.
4. Never create distractions when someone is addressing the class or the teacher, or
during a test or assignment.
5. Never touch the person or belongings of others.
6. Never throw anything.
7. Never abuse restroom breaks by using them unnecessarily.
8. For some assignments, such as tests, a grade of "zero" (0) will be assigned, and
parents/legal guardians will be notified for the following unauthorized behavior:
a. Any form of communication between students (talking, passing notes, using
signals, etc.),
b. Having access to any unauthorized materials (notes, books, etc.),
c. Submitting and attempting to receive credit for work performed by someone
other than yourself (by copying, plagiarizing, etc.).
The National AP Chemistry Exam
Students will take the national AP Chemistry Exam on Monday May 5, 2014. Most, but not
all colleges will grant some form of credit or advanced placement for a minimum score on
the exam. Student should contact the individual institutions they are interested in attending
to find out their most current policies regarding their acceptance or use of AP grades. There
is a link on edmodo that helps you find this information.
Bob Jones has joined the A+ College Ready Program. This program provides professional
development for AP instructors, funding to purchase lab equipment, $100 gift card to any
student who makes a qualifying score (3, 4 or 5) on an AP exam in Science, Math or English.
The program also provides student study session on January 25, March 8 and April 26 at
Huntsville High School from 8:30 – 2:30. These sessions bring in top notch AP teachers to
review material with the students, provide lunch to the students and door prizes. I encourage
all students to attend these Saturday sessions. I’ll be there if I am not doing a review session
for another school system.
Text
Chemistry by Zumdahl and Zumdahl, 6th ed., Houghton Mifflin Company, 2003.
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Description of Content Covered
Introductory and review concepts (½ week)
Primarily completed during Chemistry I or PreAP Chemistry. It is assumed you know
the following topics in depth.
1. Measurement topics
2. Atomic theory
3. Symbols and formulas
4. Periodic table
5. Ionic and covalent bonds
6. Nomenclature
7. Reactions
7.1. Types of Reactions
7.2. Solubility Rules
7.3. Balancing equations
8. Stoichiometry
8.1. Percent composition
8.2. Empirical formulas
8.3. Solutions
8.4. Mole relationships
8.4.1. percent (%) yield
8.4.2. Limiting reagents
9. Gas Laws
9.1. Ideal gases
9.2. Boyle’s law, Charles’ law and Gay-Lussac’s Law
9.3. Dalton’s law of partial pressure
9.4. Ideal Gas Law and Combined Gas Law
10. Electronic Structure
10.1.
Evidence for the atomic theory
10.2.
Atomic masses
10.3.
Atomic number and mass number
10.4.
Electron energy levels: atomic spectra, quantum numbers, atomic orbitals,
hybridized orbitals
10.5.
Periodic relationships
10.6.
Lewis structures
11. Arrhenius theory
11.1.
Properties of acids and bases
11.2.
Acid base neutralization
12. Lowry-Brønsted theory
12.1.
Amphiprotic species
12.2.
Relative strengths of acids and bases
12.3.
Polyprotic acids
13. Lewis acids and bases. Comparison of all three definitions.
Types of Chemical Reactions and Solution Stoichiometry
(1 week)
Chapter 4
1. Oxidation reduction reactions
1.1. Oxidation number
1.2. Electron transport
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2. Stoichiometry
2.1. Net ionic equations
2.2. Balancing equations including redox
2.3. Mass-volume relationships with emphasis on the mole
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Thermochemistry (1.5 weeks)
Chapter 6
3. Thermal energy, heat, and temperature
4. Calorimetry
5. Enthalpy changes
6. Hess’s Law
Chemical Thermodynamics (1 ½ weeks)
Chapter 16
1. State functions
2. Laws of thermodynamics
3. Relationship of change of free energy to equilibrium constants
Electrochemistry ( 1 ½ weeks)
Chapter 17
1. Galvanic cells and cell potentials
2. Electrolytic cells
3. Electrochemistry: electrolytic and galvanic cells; Faraday’s laws; standard half-cell
potentials; Nernst equation; prediction of the direction of redox reactions
The Kinetic-Molecular Theory and States of Matter (1 week)
Chapters 5, 10
1. Gas Laws
1.1. RMS velocity
1.2. Graham’s law
2. Kinetic-Molecular theory
2.1. Avogadro’s hypothesis and the mole concept
2.2. Kinetic energy of molecules
2.3. Deviations from ideality
3. Liquids and solids
3.1. Liquids and solids comparisons
3.2. Changes of state
3.3. Structure of solids including lattice energies
Bonding and Molecular Structure (1 ½ weeks)
Chapters 8 and 9
1. Binding forces
1.1. ionic
1.2. covalent
1.3. metallic including alloys
1.4. hydrogen bonding
1.5. Van der Waals
2. Relationships to states, structure, and properties of matter
3. Polarity of bonds, Electronegativities
4. VSEPR
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4.1. Geometry of molecules and ions
4.2. Structural, geometric, optical, and conformational isomerism of:
4.2.1. Organic molecules
4.2.2. Coordination complexes
5. Polarity of molecules
6. Relation of molecular structure to physical properties
7. Complex Ions
7.1. Names and structures of complex ions
7.2. Bonding in coordination systems
7.3. Formation of complex ions (reactions).
7.4. Practical applications
Solutions and Colloids (1 week)
Chapter 11
1. Types of solutions
2. Factors affecting solubility
3. Raoult’s law
4. Nonideality of solutions
Chemical Kinetics (1 week)
Chapter 12
1. Rate of reaction
2. Order of the reaction
3. Factors that change the rate of the reaction
3.1. Temperature
3.2. Concentration
3.3. Nature of substance
3.4. Catalysts
4. Relationship between the rate-determining step and the reaction mechanism
Equilibrium (1 week)
Chapter 13
1. Concept of dynamic equilibrium including Le Chatelier’s principle
2. Equilibrium constants and the law of mass action
Weak Ionic Equilibrium (1 weeks)
Chapter 15
1. Weak acids and bases
1.1. pH
1.2. pOH
1.3. Buffer systems
1.4. Hydrolysis
2. Solubility Product
2.1. Factors involving dissolution
2.2. Molar solubility
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End of Year Review
Atomic Structure:
Atomic Number = # of protons = # of electrons in a neutral atom
Mass Number = protons and neutrons (Isotopes)
Mass #-238
U
Atomic # 92
Protons=92 Electrons 92
Neutrons=238-92=146
Electron Configurations:
1. Order of filling: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s2
2. Atoms gain or lose electrons to obtain a filled octet
a) when transition metals lose electrons, they lose from the “s” sublevel first
ex. Fe: 1s22s22p63s23p64s23d6
Fe3+ : 1s22s22p63s23p63d5
b) metals get oxidized (lose electrons) to form cations (pos.) while nonmetals get
reduced (gain electrons) to form anions (neg.).
4. Oxidation states vs. group #
Group: 1= +1
2 = +2
3 = +3
(Group # = highest possible oxidation state)
14 = +4 (Except Carbide =C-4)
15 = -3
16 = -2
17 = -1
18 = 0
Transition metals have multiple oxidation states but (+2) is the most common b/c of “s”
sublevel being lost first.
*Hund’s rule:
Pauli Exclusion Principle
Shielding Effects (penetration)
Heisenberg Uncertainty Principle
diamagnetic: No unpaired electrons
paramagnetic: unpaired electrons
ferromagnetic: Fe, Co, Ni
Periodic Trends:
I.
Atomic Size(radius) :
a)
b)
On Periodic Table
Increases ←
↓
decreases left to right b/c electrons are in same energy level therefore they do
not add size but nuclear charge increases pulling electron cloud in more
tightly
increases going down in a group b/c adding more energy levels
Greater difference in size between energy levels 1,2, & 3 than between 4,5,6,
& 7 b/c energy levels are not evenly spaced.
(Transition metals are all nearly the same size)
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II.
Ionization Energy = inversely related to size (the bigger the atom, the lower the
ionization energy)
a)
exceptions to trend occur when electrons are in filled or ½ filled sublevel
Between groups IIA and IIIA - Increased shielding by "s" electrons
b)
Between groups VA and VIA - Increased electron ↔ electron repulsions
because of electrons beginning to pair up in "p" orbitals
c)
ions become very stable (high I.E.) once they obtain a noble gas config.
Intermolecular Forces
4 types of substances
1. Ionic: metal w/ nonmetal
a) High melting & boiling points b/c of high lattice energy binding ions together
1.
Coulumb’s Law: ΔH
=
Kq1q2
(q1 & q2 = charges on ions)
(lattice energy)
r
2.
As the product of the charges increases, the electrostatic attraction increases,
therefore higher melting and boiling points
3.
The bigger the ions, the lower the electrostatic attractions, therefore the lower
melting and boiling points.
b) Do not conduct electricity in the solid state b/c the ions are held rigidly in place.
they do conduct in the liquid (molten) and aqueous states b/c the ions dissociate
and are free to move.
c) Brittle and therefore neither malleable nor ductile
d) Form large crystals-no molecules
2. Metallic: Like metal atoms bonded together
a) conduct electricity in the solid and liquid states b/c of the “sea of free-floating
valence electrons.”
b) Malleable and ductile b/c (see part a)
c) Vary wide variety of melting points and boiling points.
From Hg = liquid to W = solid w/ highest melting point (Generally the melting
point is greater than 300o C and therefore solids at room temperature)
d) Insoluble in H2O
e) Mixture of metals is called an “Alloy”
3. Covalent Network (Large crystals of nonmetals covalently bonded to one another).
a) only Cdiamonds, Cgraphite, SiO2, SiC
\Allotropes/
(Quartz)
b) Do not conduct electricity (except Graphite)
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c) Diamonds = sp3 3d crystal-very hard
Graphite = sp2 2d crystal-forms loose layers
d) insoluble in H2O
e) High melting and boiling points and therefore solids at room temperature
4. Covalent Molecular: nonmetal w/ nonmetal
a) nonmetals bond with each other to form molecules
b) molecules are held together by weak intermolecular forces called van der
Waals forces
1. London Forces: Increases with the number of electrons in the molecule.
(size of the molecule)
a) The weakest of intermolecular forces
b) The only force of attraction between nonpolar molecules
c) Also called: Dispersion, Induced Dipole, Instantaneous Dipole
2. Dipole-Dipole:
The attraction between oppositely charged portions of
polar molecules.
a) The greater the dipole moment (a measurement of the molecule’s
polarity) the stronger the D.P. –D.P. attraction.
b) Stronger than London forces
3. Hydrogen Bonds: H atom must be attached directly to a N, O, or F.
a) Strongest of the 3 intermolecular forces but covalent network and ionic
bonds are still much stronger
Bonding
*Lewis Dot Structure: #valence electrons = group number
*Atoms want to achieve a stable octet noble gas configuration
*exceptions to rule: Be, Al, B - form stable molecules w/ 6 e- around central atom
*Resonance- occurs when molecule can have more than 1 possible structure (always has
double bonds/triple bonds)
Ex.
O=N-O
↔ O-N=O
(bond order = 1½ )
*all the bonds are actually the same length (average of the 2 bonds)
*Double bonds and Triple Bonds
bond length: single(longest) → double → triple(shortest)
Diatomic elements: H2, O2, N2, Cl2, F2, I2, Br2
*all form single bonds except: O2 (double bond) O=O
N2 (triple bond)
N≡N
*Sigma and Pi bonds-
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first bond between any two atoms=sigma bond (strongest)
second and third bonds = pi bonds
bond order = strength of bond (single, double, triple)
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*Polar vs. Nonpolar
polar: bond that acquires positive and negative ends(has lone pair, asymmetrical)
nonpolar: shared pair of electrons is equally shared, no net displacement
(symmetrical- no lone pair on central atom) *No dipole moment
*Likes Dissolve Likes
Polar dissolves only into polar (*Water is Polar)
Ionic dissolves into polar
Nonpolar dissolves only into nonpolar
*Electronegativity-measure of ability of an atom to attract electrons to itself
most electronegative element: F, least: Fr
*structural isomers- appear to have same structure but are different
ex. C2H5OH and CH3OCH3
^
Hydrogen Bond
^
No H bond
*VSEPR theory (valence shell electron pair repulsion)
-bonding and nonbonding pairs repel to obtain the farthest distances between each other
molecular shapes: see handout in this document
*Hybridization: s, p, and d orbitals can be mixed to form new sets of orbitals
Hybrid Orbitals
Required
2
sp
3
sp2
4
sp3
linear
trigonal planar
tetrahedral
AX2
AX3
AX4
(Could Have)
triple, double, or single bonds
double bonds, or single bonds
single bonds
*Add up exponents to see how many bonds
Differences in Boiling Points
-for covalent compounds, always look for Hydrogen bonding
-also look for London Forces (vary w/ # of electrons ex. bigger molecules have higher B.P.)
-For ionic compounds: Coulombs Law ΔH= KQ1Q2
r
lattice energy-energy holding oppositely charged ions together
-Heat of Vaporization-energy needed to vaporize something at its boiling point (Latent heat)
-varies with strength of attractive force
-
Heat of Fusion - energy needed to melt a substance at its melting point. (Latent Heat)
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Gases
Gas Laws:1 atm = 760mm Hg (torr)
1. Boyle’s Law:
P1V1 = P2V2
Pressure & volume are inversely proportional
2. Charles’ Law:
V1/T1 = V2/T2
Volume is directly proportional to temperature
3. Gay Lussac’s Law: P1/T1 = P2/T2
Pressure is directly proportional to temperature (T is in Kelvin)
4. Avogadro’s Law:
N1/V1 = N2/V2
Moles of gas is directly proportional to volume
5. Ideal Gas Law:
PV = nRT
R = universal gas constant = 0.0821 (L*atm)/(k*mol)
Another form: MM = (dRT)/P (dirty P)
(molar mass = MM, density = d)
6. Dalton’s Law of Partial Pressure:
P T = P1 + P2 + P 3 …
Partial Pressure of a gas = X(Ptotal)
(X = mole fraction of gas)(P = total pressure)
PT = XAP˚ A + XBP˚ B
(ideal solution with more than one volatile compound)
STP = 1atm, 273K
1 mol gas = 22.4L gas
R = 0.0821 (L*atm)/(mol*K) = 8.314 Joules/(mol*K)
Ideal Gases versus Real Gases


2 assumptions of ideal gas:
o no attractive forces between molecules
o molecules are infinitely small
Real Gases behave ideally at:
o High temperature
o Low pressure
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Thermodynamics
- thermodynamics – science of heat &
work
- ΔH = enthalpy change (heat transferred)
- ΔH = (Hproducts – Hreactants)
- ΔS = entropy (measure of randomness)
increase in entropy =>
solid →
liquid → gas
(least entropy)
(most entropy)
ΔH
Negative value
Exothermic
Positive value
Endothermic
ΔS
Entropy decrease
Entropy increase
ΔG
Spontaneous
Non Spontaneous
Less than 1
Non Spontaneous
(reactant favored)
Greater than 1
Spontaneous
(product favored)
- ΔG = ΔH – TΔS
When ΔH = TΔS (and both ΔH and
ΔS are positive/negative) Then ΔG = O
(meaning at equilibrium)
- ΔG = - RTlnK
R = 8.314
K = equilibrium constant
K
Kinetics
-
reaction mechanisms – detailed pathways taken by atoms and molecules as a reaction
proceeds
-
factors that affect reaction rates:
1.) concentration of reactants
2.) temperature
3.) catalysts
4.) orientation of molecules
-
Rate Equation (Law): rate = K[A]x[B]y
K = rate constant; ‘x’ and ‘y’ can only be determined by experimental data
Δrate = (Δ[conc]) x
Slow Rates
▪ strong bonds
in reactant
molecules
Fast Rates
▪ catalyst
▪ high temp.
▪ high [reactant]
▪ low activation
energy
◊ Rate Law dependent upon ◊ Slowest Step ◊ of reaction mechanisms
-
Order of Reaction is sum of the exponents (x + y …)
-
If concentration of reactant is increased, the rate is usually increased (unless the
exponent is 0 – not involved in rate law)
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▪
catalyst lowers activation energy
▪ temp increase allows more molecules to overcome EA
▪ 2 criteria for a mechanism:
1.) steps must add up stoichiometricly.
2.) rate law of rate def’n step = rate law.
Equilibrium
▪ K does not change unless temp.
changes
Equilibrium expression for:
aA + bB ↔ cCdD
K = [C]c[D]d
[A]a[B]b
([products]/[reactants])
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▪ only gases and aqueous included in equilibrium expression
Kc = equilibrium constant of concentration
Kp = equilibrium constant of pressure
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Kp = Kc(RT) Δn
Δn = total moles gaseous product – total moles gaseous reactant
R = 8.31
K=1
at equilibrium
when [prod.] = [reactants],
RARE
K>1
product favored
(spontaneous)
K<1
reactant favored
(non spontaneous)
Q = Reaction Quotient
Q > K reactant favored (non spontaneous)
Q < K product favored (spontaneous)
Q = K at equilibrium
Le Chatelier’s Principle – a change in any of the factors that determine equilibrium
conditions of a system cause the system to change in such a manner to counteract the
effect of the change
▪ factors that effect equilibrium:
1. change in concentration
2. change in temp
3. change in volume
4. addition of catalyst
2A + B ↔ C + D
Shifts
1.) Add A / Remove D
→
2.) Remove B / Add C
←
3.) Increase Volume
←
(Shifts to side w/ greater moles of gas)
4.) Decrease Volume
→
(Shifts to side w/ fewer moles of gas)
If endothermic reaction,
X+H→Y
1.) Add heat, Shifts right
2.) Remove heat, Shifts left
If Exothermic
X→Y+H
1.) Add heat, Shifts left
2.) Remove heat, Shifts right
Electrochemistry
Oxidation
- lose electrons
- metals get oxidized
- occurs at ANODE
- cations formed
- has a negative Ecell
Reduction
- gain electrons
- non metals get reduced
- occurs at CATHODE
- anions formed
- has a positive Ecell
E˚ cell = Ered - Eox
˚ indicates standard conditions (25˚C, 1.0 atm)
▪ if Ecell
( + ) spontaneous
( - ) non spontaneous
-
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Common Oxidizing Agents:
KMnO4 → Mn2+
K2Cr2O7 → Cr3+
Concentrated HNO3 → NO2
Dilute HNO3 → NO
H2O2 → H2O + O2
Cathode(-)
Getting reduced →
-
Anode(+)
← Getting oxidized
Salt Bridge – device used for maintaining balance of ion
charges in the cell compartments
1 Coulomb = current (amps) * time (seconds)
1 Faraday = 96,500 coulombs = 1mol eNernst Equation: E = E˚ - [(0.0257V)/n]*lnQ
at 25˚C
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Solution Practice
1. A 20.0 g sample of HF is dissolved in water to give 2.0 x 102 mL of solution. The
concentration is?
2. How many grams of NaCl are contained in 350. mL of a 0.250 M solution of sodium
chloride?
3. What volume of 18.0 M sulfuric acid must be used to prepare 15.5 L of 0.195 M
H2SO4?
4. Write the net ionic equation for the reaction of calcium bromide and sodium
phosphate.
5. Which of the following is NOT a weak acid?
a. HCNO
b. HBr
c. HF
d. HNO2
e. HCN
6. You have exposed electrodes of a light bulb in a solution of sulfuric acid such that the
light bulb is on. You add a dilute solution and the bulb grows dim. Which of the
following could be in the solution? Explain your anwer.
a. Ba(OH)2
b. NaNO3
c. K2SO4
d. Ca(NO3)2
e. Huh?
7. What precipitates when you combine sodium sulfide and copper II chloride solutions?
8. You have 75.0 mL of a 2.50 M solution of sodium chromate. You also have 125 mL
of a 2.50 M solution of silver nitrate.
a. Calculate the concentration of sodium ion when the two are added together.
b. Calculate the concentration of chromate ion.
c. Calculate the concentration of silver ion.
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Redox Practice
Acid
1. I¯ (aq) + ClO¯(aq)
 I3¯(aq) + Cl¯(aq)
2. As2O3 (s) + NO3¯ (aq)  H3AsO4 (aq) + NO (g)
3. Br¯ (aq) + MnO4¯ (aq)  Br2 (l) + Mn2+ (aq)
4. Cr2O72−(aq) + HNO2(aq)
 Cr3+(aq) + NO3−(aq)
5. MnO4–(aq) + C2O42–(aq)
CO2(g) + Mn2+(aq)
Base
1. Al(s) + MnO4¯ (aq)  MnO2(s) + Al(OH)4¯ (aq)
2. NO2¯ (aq) + Al(s)  NH3(aq) + AlO2¯ (aq)
3. Cr(s) + CrO42-(aq)  Cr(OH)3(s)
Note: Cr(OH)3 is found in BOTH half reactions!
1. ... Li3N(s) + ... H2O(l) ...Li+(aq) + ...OH-(aq) + ...NH3(g)
When the equation above is balanced and all coefficients reduced to lowest whole-number terms, the
coefficient for OH-(aq) is
2. …Cr2O72-(aq) + …H2S(g) + …H+(aq)  …Cr3+(aq) + …S(s) + …H2O(l)
When the equation above is correctly balanced and all coefficients are reduced to lowest
whole-number terms, the coefficient for H+(aq) is
3. 2 MnO4-(aq) + 10 Br-(aq) + 16 H+ (aq)  2 Mn2+(aq) + 5Br2(aq) + 8 H2O(l)
How many electrons are transferred in the reaction represented by the balanced equation above?
4. ...H+(aq) + ...NO2-(aq) + ...Cr2O72-(aq) ...Cr3+(aq) + ...NO3-(aq) + ...H2O(l)
When the equation above is balanced and all coefficients are reduced to lowest whole number terms,
the coefficient for H2O(l) is
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AP Calorimetry worksheet
1. It takes 78.2 J of heat to raise the temperature of 45.6 g of lead by 13.3˚C. Calculate the
specific heat of the lead.
2. It takes 585 J of heat to raise the temperature of 125.6 g of mercury from 20˚C to 53.5˚C.
Calculate the specific heat of the mercury.
3. A 30 gram sample of water at 280K is mixed with 50 grams of water at 300K. Calculate
the final temperature of the mixture assuming no heat is lost to the surroundings.
4. A 15 gram sample of nickel metal is heated to 100˚C and then dropped into 55 grams of
water, initially at 23˚C. Assuming that all the heat lost by the nickel is gained by the
water calculate the final temperature of the nickel- water mixture. The specific heat of
nickel is 0.444 J/g˚C.
5. A 28.2 gram sample of a metal is heated to 99.8˚C and then placed in a coffee-cup
calorimeter containing 150.0 grams of water at 23.5˚C. After the metal cools, the final
temperature of the metal-water mixture is 25˚C. Calculate the specific heat of the metal,
assuming no heat escapes to the surroundings or is transferred to the calorimeter.
6. A 46.2 gram sample of copper is heated to 95.4˚C and then placed in a calorimeter
containing 75.0 grams of water at 19.6˚C. The final temperature of the copper-water
mixture is 21.8˚C. Calculate the specific heat of copper, assuming that all the heat lost by
the copper is gained by the metal.
7. A coffee-cup calorimeter initially contains 125 grams of water at 24.2˚C. 10.5 grams of
potassium bromide, also at 24.2˚C is added to the water. After the KBr dissolves, the
final temperature of the mixture is 21.1˚C. Calculate the heat of dissolving the salt in J/g
and kJ/mol. Assume that the specific heat and density of the solution is the same as that
of water, and that no heat is transferred to the surroundings or calorimeter.
8. In a coffee-cup calorimeter, 100 mL of 1.0M NaOH solution and 100 mL of 1.0M HCl
solution are mixed. Both solutions were originally ate 24.6˚C.
After the reaction, the final temperature of the solution is 31.3˚C.
Assuming that the solution has the density and specific heat of water, calculate the
heat of neutralization of HCl by NaOH. Assume no heat is lost to the surroundings or
calorimeter.
9. In a coffee-cup calorimeter, 50 mL of 0.1M AgNO3and 50 mL of 0.1 M HCl are mixed
together. The equation for the net ionic reaction is:
Ag+1 + Cl-1 ----- > AgCl
The two solutions were initially at 22.6˚C and the final temperature is 3.4˚C. Calculate
the heat that accompanies this reaction in kJ/mole of AgCl formed. Assume that the
combined solution has the density and specific heat of water.
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Thermochemistry: Standard Heats of Formation Worksheet
Use the standard molar heat of formation table to calculate H for:
1. 2 NO(g) + O2(g) ---> 2 NO2(g)
2. NaOH(s) + HCl(g) ----> NaCl(s) + H2O(g)
3. 2 CO(g) + O2(g) ---> 2 CO2(g)
4. CH4(g) + 2 O2(g) ---> CO2(g) + 2 H2O(l)
5. 2 H2S(g) + 3 O2(g) ---> 2 H2O(l) + 2 SO2(g)
6. f) SO2(g) + ½ O2(g) ---> SO3(g)
7. CaO(s) + H2O(l) ---> Ca(OH)2(s)
8. N2(g) + 3 H2(g) ---> 2 NH3(g)
9. C6H6(l) + 1½ O2(g) ---> 6 C(s) + 3 H2O(l)
10. NH3(g) + HCl(g) ---> NH4Cl(s)
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Gas laws practice.
You can keep me as your own but you really need to do the work and all on a separate sheet,
don’t you think?
1.
A)
B)
C)
D)
E)
2.
A)
B)
Is a gas in its standard state at 298 K
Lithium
Nickel
Bromine
Uranium
Fluorine
A hot-air balloon rises. Which of the following is the best explanation for this observation?
The pressure on the walls of the balloon increases with increasing temperature.
The differences in temperature between the air inside and outside the balloon produces
convection currents.
C) The cooler air outside the balloon pushes in on the walls of the balloon.
D) The rate of diffusion of cooler air is less than that of warmer air.
E) The air density inside the balloon is less than that of the surrounding air.
3.
A rigid metal tank contains oxygen gas. Which of the following applies to the gas in the tank
when additional oxygen is added at constant temperature?
A) The volume of the gas increases
B) The pressure of the gas decreases
C) The average speed of the gas molecules remains the same
D) The total number of gas molecules remains the same
E) The average distance between the gas molecules increases
4.
W(g) + X(g) Y(g) + Z(g)
Gases W and X react in a closed, rigid vessel to form gases Y and Z according to the equation above.
The initial pressure of W(g) is 1.20 atm and that of X(g) is 1.60 atm. No Y(g) or Z(g) is initially
present. The experiment is carried out at constant temperature. What is the partial pressure of Z(g)
when the partial pressure of W(g) has decreased to 1.0 atm?
5.
Equal numbers of moles of He(g), Ar(g), and Ne(g) are placed in a glass vessel at room
temperature. If the vessel has a pinhole sized leak, which of the following will be true regarding the
relative values of the partial pressures of the gases remaining in the vessel after some of the gas
mixture has effused?
A) PHe < PNe < PAr
B) PHe < PAr < PNe
C) PNe < PAr < PHe
D) PAr < PHe < PNe
E) PHe = PNe = PAr
6.
Which of the following gases deviates most from ideal behavior?
A) SO2
B) Ne
C) CH4
D) N2
E) H2
7.
Consider three 1-L flasks at STP. Flask A contains NH3 gas, flask B contains NO2 gas, and
flask C contains N2 gas.
1. Which contains the largest number of particles?
2. In which flask are the molecuels least polar and therfore most ideal in behavior?
3. In which flask do the molecules have the highest average velocity?
10.
The valve between a 5L tank containing a gas at 9 atm and a 10 L tank containing a gas at 6
atm is opened. Calculate the final pressure in the tanks.
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11.
A 4.40 g piece of solid carbon dioxide is allowed to sublime in a balloon. The final volume
of the balloon is 1.00 L at 300. K. What is the pressure of the gas?
12.
Calculate the root mean square velocity for the oxygen molecules in a sample of oxygen gas
at 25.0C.
13.
It is found that 250. mL of gas at STP has a mass of 1.00 g. What is the molar mass?
14.
Order the following in increasing rate of effusion:
F2, Cl2, NO, NO2, CH4
15.
Which of the following would have a higher rate of effusion than C2H2?
CH4, Cl2, O2, CO2, N2
16. Given the following equation;
Na(s) + O2(g) → Na2O(s)
How many grams of Na are required to react with 2.5L of O2 at 755 torr and 15°C?
17. A sealed flask contains a gas with pressure P. If the number of moles of the gas is
doubled, the absolute temperature is doubled and the volume of the container is
quadrupled, what is the new pressure in terms of P?
18. In the unbalanced equation,
NO(g) + O2(g) → NO2(g)
If 30g each of NO and O2 are reacted to completion in a 10L sealed container at 20°C,
what is the final
pressure in the container? (NOTE: This is a limiting reagent problem!)
19. If a quantity of chlorine gas occupies 12L at STP, what volume will it need to occupy if
the pressure is halved and the temperature is decreased by 30K?
20. 20g each of Helium and an unknown diatomic gas are combined in a 1500mL container.
If the temperature is 298K, and the pressure inside is 86.11atm, what is the unknown gas?
21. If 4 moles of gas are added to a container that already holds 1 mole of gas, how will the
pressure change within the container?
22. A breathing mixture used by deep-sea divers contains helium, oxygen, and carbon
dioxide. What is the partial pressure of oxygen at 1 atmosphere if PHe=609.5 mmHg,
Pcarbon dioxide = 0.5 mmHg?
23. If a balloon containing 1,000L of gas at 50. C and 760 mmHg rises to an altitude where
the pressure is 380mmHg and the temperature is 10. C, what is the new volume of the
balloon (in L)?
24. What is the pressure exerted by 32g of O2 in a 20-L container at 30.0 C?
25. The gaseous product of a reaction is collected in a 25.0L container at 27.0 C. The
pressure in the container is 3.0atm and the gas has a mass of 96.0g. What is the molar
mass of the gas?
26. What volume of H2O gas can be produced from 20.0g of oxygen and excess hydrogen at
150.0 C and .5atm?
27. In an experiment, it takes an unknown gas 1.5 times longer to diffuse than the same
amount of oxygen gas, O2. Find the molar mass of the unknown gas.
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8 and 9 Practice worksheet
1.
In the gaseous phase, which of the following diatomic molecules would be the most
polar?
A) NaF
B) CsCl
C) CsF
D) NaCl
E) LiF
2.
A)
B)
C)
D)
Which of these is an isoelectronic series?
Na+, K+, Rb+, Cs+
K+, Ca2+, Ar, S2
Na+, Mg2+, S2, Cl
Li, Be, B, C
3.
Which of the following statements concerning lattice energy is false?
A) MgO has a larger lattice energy than LiF.
B) MgO has a larger lattice energy than NaF.
C) The lattice energy for a solid with 2+ and 2 ions should be two times that for a
solid with 1+ and 1 ions.
D) It is often defined as the energy released when an ionic solid forms from its ions.
4.
Using the following data reactions
H2(g) + Cl2(g)  2HCl(g) H =-184 kJ/mol
Bond energy of H2 = 432 kJ/mol
Bond energy of Cl2 = 239 kJ/mol
calculate the energy of an H–Cl bond.
5.
A)
B)
C)
D)
This molecule contains a carbon atom with trigonal planar geometry.
CH3CHO
CO2
CH3Cl
C2H6
6.
NCl3 what is the geometry?
7.
XeF5+ what is the geometry?
8.
What type of geometry does the OXeF2 molecule have?
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9.
Which of the following molecules contains a central atom with sp2 hybridization?
A)
B)
C)
D)
E)
10.
Consider the molecule
What is the hybridization of all the atoms (except H)
11.
12.
A)
B)
C)
D)
E)
The hybridization of Se in SeF2 is
Which of the following groups contains no ionic compounds?
HCN, NO2, Ca(NO3)2
KOH, CCl4, SF4
PCl5, LiBr, Zn(OH)2
NaH, CaF2, NaNH2
CH2O, H2S, NH3
13.
The electron pair in a C-F bond could be considered
A) closer to C because carbon has a larger radius and thus exerts greater control over
B)
C)
D)
E)
the shared electron pair.
closer to F because fluorine has a higher electronegativity than carbon.
centrally located directly between the C and F.
closer to C because carbon has a lower electronegativity than fluorine.
an inadequate model since the bond is ionic.
Which of the following ionic compounds has the smallest lattice energy, i.e., the
lattice energy least favorable to a stable lattice?
A) BaO
B) CsI
C) MgO
D) NaCl
E) LiF
14.
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15.
A)
B)
C)
D)
E)
16.
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Which of the following arrangements is in order of increasing size?
S2 < Cl < K+ < Ca2+ < Ga3+
Ga3+ < Ca2+ < S2 < Cl < K+
Ga3+ < S2 < Ca2+ < Cl < K+
Ga3+ < Ca2+ < K+ < Cl < S2
Ga3+ < Ca2+ < S2 < K+ < Cl
Given the following bond energies (the first c=o bond should be a single bond)
C  C 347 kJ / mol
C=C 614 kJ / mol
C=O
358 kJ / mol
C=O
799 kJ / mol
C H
413 kJ / mol
O  H 463 kJ / mol
O  O 146 kJ / mol
estimate H for the reaction H2O2 + CH3OH  H2CO + 2H2O.
17.
A)
B)
C)
D)
This molecule is the most polar.
CH3CHO
CO2
CH3Cl
C2H6
Select the best Lewis structure for acetone, CH3COCH3. What is wrong with the
other four choices?
18.
A)
B)
C)
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D)
19.
In the cyanide ion (CN), the nitrogen has a formal charge of
Choose the electron dot formula that most accurately describes the bonding in CS2.
(Hint: Consider formal charges.)
20.
A)
B)
C)
D)
E)
21.
A)
B)
C)
D)
E)
Which of the following has an incomplete octet in its Lewis structure?
ICl
F2
SO2
NO
CO2
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Equilibrium and Entropy Practice
1.
Which of the following statements is true?
A) Catalysts are an effective means of changing the position of an equilibrium.
B) The concentration of the products equals that of reactants and is constant at
equilibrium.
C) When two opposing processes are proceeding at identical rates, the system is at
equilibrium.
D) An endothermic reaction shifts toward reactants when heat is added to the reaction.
2. Consider the chemical system CO + Cl2
COCl2; K = 4.6  109 L/mol.
How do the equilibrium concentrations of the reactants compare to the equilibrium
concentration of the product?
3.
For the reaction given below, 2.00 moles of A and 3.00 moles of B are placed in a
6.00-L container.
A(g) + 2B(g)
C(g)
At equilibrium, the concentration of A is 0.300 mol/L. What is the concentration of B at
equilibrium?
4.
A 10.0-g sample of solid NH4Cl is heated in a 5.00-L container to 900C. At
equilibrium the pressure of NH3(g) is 1.20 atm.
NH4Cl(s)
NH3(g) + HCl(g)
The equilibrium constant, Kp, for the reaction is:
5.
Consider the following reaction (assume an ideal gas mixture):
NOBr(g)
NO(g) + Br2(g)
A 1.0-liter vessel was initially filled with pure NOBr, at a pressure of 4.0 atm, at 300 K.
After equilibrium was established, the partial pressure of NOBr was 2.5 atm. What is Kp for
the reaction?
6.
A sample of solid NH4NO3 was placed in an evacuated container and then heated so
that it decomposed explosively according to the following equation:
NH4NO3(s)
N2O(g) + 2H2O(g)
At equilibrium the total pressure in the container was found to be 3.20 atm at a temperature
of 500C. Calculate Kp.
7.
Consider the following reaction:
HF(g)
H2(g) + F2(g) (K = 1.00  102)
Given 1.00 mole of HF(g), 0.500 mole of H2(g), and 0.750 mole of F2(g) are mixed in a 5.00L flask, determine the reaction quotient, Q, and the net direction to achieve equilibrium.
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8.
Consider the following equilibrium:
NOCl(g)
NO(g) + Cl2(g)
5

with K = 1.6  10 . 1.00 mole of pure NOCl and 1.00 mole of pure Cl2 are replaced in a
1.00-L container.
Calculate the equilibrium concentration of NO(g).
9.
Consider the following equilibrium:
2H2(g) + X2(g)
2H2X(g) + energy
Addition of argon to the above equilibrium will have what effect on equilibrium?
10.
A)
B)
C)
D)
E)
For which process is S negative?
evaporation of 1 mol of CCl4(l)
mixing 5 mL ethanol with 25 mL water
raising the temperature of 100 g Cu from 275 K to 295 K
grinding a large crystal of KCl to powder
compressing 1 mol Ne at constant temperature from 1.5 atm to 2.5 atm
11.
Which of the following result(s) in an increase in the entropy of the system? (could
be more than one)
I. (See diagram shown.)
II. Br2 (g)  Br2 (l)
III. NaBr(s)  Na + (aq) + Br  (aq)
IV. O 2 (298 K)  O 2 (373 K)
V. NH 3 (1 atm, 298 K)  NH 3 (3 atm, 298 K)
12.
A)
B)
C)
D)
For a nonspontaneous exothermic process, which of the following must be true?
S must be positive.
S must be negative.
G must be positive.
Two of these must be true.
13. For the process CHCl3(s)  CHCl3(l), H = 9.2 kJ/mol and S = 43.9 J/mol/K. What
is the melting point of chloroform?
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14. When a stable diatomic molecule spontaneously forms from its atoms, what are the signs
of H, S, and G?
H 
S 
G
Option A



Option B



Option C



Option D
Option E






15.
The reaction
2H2O(g)  2H2(g) + O2(g)
has a positive value of G. Which of the following statements must be true?
A) The reaction is slow.
B) The reaction will not occur. [When H2O(g) is introduced into a flask, no O2 or H2
ill form even over a long period of time.]
C) The reaction is exothermic.
D) The equilibrium lies far to the right.
E) None of these is true.
16.
Given the following free energies of formation:
Gf
C2 H 2 (g)
209.2 kJ / mol
C2 H 6 (g)
 32.9 kJ / mol
calculate Kp at 298 K for C2H2(g) + 2H2(g)
C2H6(g)
17.
Given the following data, calculate the normal boiling point for formic acid
(HCOOH).
Hf ( kJ / mol)
S  (J / mol K)
18.
HCOOH(l)
 410
130
HCOOH(g)
 363
251

Given that Gf for NH3 = 16.67 kJ/mol, calculate the equilibrium constant for the
following reaction at 298 K:
N2(g) + 3H2(g)
2NH3(g)
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Ch. 10 questions
1. What are Dipole dipole forces? Give an example.
2. What are the 2 factors that decide how strong hydrogen bonds are?
3. Describe London dispersion forces?
4. What is an instantaneous dipole?
5. How is a dipole induced?
6. What governs the amount of surface tension?
7. Distinguish between cohesion and adhesion.
8. What kind of meniscus does mercury have and why?
9. What makes liquids viscous?
10. Explain the difference in viscosity between gasoline and grease.
11. Define lattice and unit cell, amorphous solid and regular solid.
12. What is a diffractometer.
13. Write the Bragg equation and describe what it is used for?
Chapter 11 Practice
Answers in your notes but keep this sheet for reference.
1. A solution of hydrogen peroxide is 30.0% H2O2 by mass and has a density of 1.11 g/cm3. The
molarity of the solution is:
2. The term “proof” is defined as twice the percent by volume of pure ethanol in solution. Thus,
a solution that is 95% (by volume ethanol is 190 proof. What is the molarity of ethanol in a
92 proof ethanol/water solution?
density of ethanol = 0.80 g / cm3
density of water = 1.0 g / cm3
mol. wt. of ethanol = 46
3. What is the effect of adding 1 mol of glucose vs 1 mol of magnesium chloride on the vapor pressure of an
aqueous solution?
4.What is the molality of a solution of 50.0 g of propanol (CH3CH2CH2OH) in 152 mL water, if the
5.
6.
7.
8.
9.
10.
11.
density of water is 1.0 g/mL?
How many milliliters of 18.4 M H2SO4 are needed to prepare 600.0 mL of 0.10 M H2SO4?
A 20.0-g sample of methyl alcohol (CH3OH, molar mass = 32.0 g/mol) was dissolved in 30.0
g of water. The mole fraction of CH3OH is:
Which of the following concentration measures will change in value as the temperature of a
solution changes? Explain why
Molarity, mass percent, molality, mole fraction
What is the mole percent of ethanol (C2H5OH) in 180 proof vodka, which consists of 71.0 g
of ethanol for every 10.0 g of water present?
If 2.00 g of helium gas and 4.00 g of oxygen gas are mixed together what is the mole fraction
of helium in the solution?
A correct statement of Henry’s law is:
the concentration of a gas in a solution is directly proportional to pressure.
the concentration of a gas in solution is independent of pressure.
the concentration of a gas in solution is directly proportional to the mole fraction of solvent.
the concentration of a gas in solution is inversely proportional to temperature.
hexane (C6H14) and chloroform (CHCl3)
12.
13.
14.
15.
16.
17.
18.
19.
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Describe the deviation with respect to Raoult’s Law. Use the following choices
relatively ideal, positive deviation, negative deviation,
acetone (C3H6O) and water
Describe the deviation with respect to Raoult’s Law. Use the following choices
relatively ideal, positive deviation, negative deviation,
hexane (C6H14) and octane (C8H18)
Describe the deviation with respect to Raoult’s Law. Use the following choices
relatively ideal, positive deviation, negative deviation,
The vapor pressure of water at 25.0C is 23.8 torr. Determine the mass of glucose (molar
mass = 180 g/mol) needed to add to 500.0 g of water to change the vapor pressure to 23.1
torr.
150 g of NaCl completely dissolves (producing Na+ and Cl ions) in 1.00 kg of water at
25.0C. The vapor pressure of pure water at this temperature is 23.8 torr. Determine the
vapor pressure of the solution.
At 40C, heptane has a vapor pressure of 92.0 torr and octane has a vapor pressure of 31.2
torr. Assuming ideal behavior, what is the vapor pressure of a solution that contains twice as
many moles of heptane as octane?
A solution is prepared from 31.4 g of a nonvolatile, nondissociating solute and 85.0 g of
water. The vapor pressure of the solution at 60C is (142 torr). The vapor pressure of water
at 60C is 150. torr. What is the molar mass of the solute?
A solution contains 1 mole of liquid A and 3 mol of liquid B. This solution has a vapor
pressure of 314 torr at 25C. At 25C, liquid A has a vapor pressure of 265 torr and liquid B
has a vapor pressure of 355 torr. Which of the following is true?
This solution is ideal.
More information is needed to answer this question.
This solution exhibits a positive deviation from Raoult’s Law.
This solution exhibits a negative deviation from Raoult’s Law.
Solutions of benzene and toluene obey Raoult’s law. The vapor pressures at 20C are:
benzene, 76 torr; toluene, 21 torr What is the mole fraction of benzene in a solution whose
vapor pressure is 50 torr at 20C?
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Kinetics part I
1. The average rate of disappearance of ozone in the reaction 2O3(g)  3O2(g) is found to be
9.0  10-3 atm over a certain interval of time. What is the rate of appearance of O2 during this
interval?
2. Consider the following rate law: Rate = k[A]n[B]m
How are the exponents n and m determined?
3. The following data were obtained for the reaction of NO with O2. Concentrations are in
molecules/cm3 and rates are in molecules/cm3  s.
[NO]0
[O 2 ]0
Initial Rate
1 1018
1 1018
2.0  1016
2 1018 1 1018
8.0  1016
3 1018 1 1018 18.0  1016
1 1018 2 1018 4.0  1016
1 1018 3 1018 6.0  1016
What is the correct rate law?
4. The reaction of (CH3)3CBr with hydroxide ion proceeds with the formation of (CH3)3COH.
(CH3)3CBr(aq) + OH(aq)  (CH3)3COH(aq) + Br(aq)
The following data were obtained at 55C.
[(CH 3 ) 3 CBr]0 [OH  ]0 Initial Rate
Exp.
(mol / L)
(mol / L) (mol / L s)
1
0.10
0.10
1.0 10 3
2
0.20
0.10
2.0 10 3
3
0.10
0.20
1.0 10 3
4
0.30
0.20
?
What will the initial rate (in mol/L  s) be in Experiment 4?
5. For a reaction in which A and B react to form C, the following initial rate data were obtained:
[A]
[B]
(mol / L) (mol / L)
0.10
0.10
Initial Rate of Formation of C
(mol / L  s)
1.00
0.10
0.20
4.00
0.20
0.20
8.00
What is the rate law for the reaction?
What is the value of k?
What is the order with respect to A?
What is the order with respect to B?
What is the overall order of the reaction?
6. Tabulated below are initial rate data for the reaction
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2Fe(CN)63 + 2I  2Fe(CN)64 + I2
3

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4
Run [Fe(CN) 6 ]0 [I ]0 [Fe(CN) 6 ]0 [I 2 ]0
1
0.01
0.01
0.01
0.01
2
0.01
0.02
0.01
0.01
3
0.02
0.02
0.01
0.01
4
0.02
0.02
0.02
0.01
5
0.02
0.02
0.02
0.02
What is the rate law?
What is the overall order of the reaction?
Initial
Rate (M / s)
1  105
2  105
8  105
8  105
8  105
7. A general reaction written as 1A + 2B  C + 2D is studied and yields the following data:
[A]0
[B]0
Initial [C] / t
0.150 M 0.150 M 8.00  10 3 mol / L  s
0.150 M 0.300 M 1.60  10 2 mol / L  s
0.300 M 0.150 M 3.20  10 2 mol / L  s
What is the order of the reaction with respect to B?
8. A general reaction written as 1A + 2B  C + 2D is studied and yields the following data:
[A]0
[B]0
Initial [C] / t
0.150 M 0.150 M 8.00  10 3 mol / L  s
0.150 M 0.300 M 1.60  10 2 mol / L  s
0.300 M 0.150 M 3.20  10 2 mol / L  s
Determine the initial rate of C production if [A] = 0.200 M and [B] = 0.500 M
9. Consider the following data concerning the equation:
H2O2 + 3I + 2H+  I3 + 2H2O
[H 2 O 2 ]
[I  ]
[H + ]
rate
I. 0.100 M 5.00  10 M 1.00  10 M 0.137 M / sec
4
2
II. 0.100 M 1.00  103 M 1.00  10 2 M 0.268 M / sec
III. 0.200 M 1.00  103 M 1.00  10 2 M 0.542 M / sec
IV. 0.400 M 1.00  103 M 2.00  102 M 1.084 M / sec
Write the rate law for the reaction
10. The rate expression for a particular reaction is rate = k[A][B]2. If the initial concentration of B
is increased from 0.1 M to 0.3 M, the initial rate will increase by what factors?
11. The kinetics of the reaction A + 3B  C + 2D were studied and the following results
obtained, where the rate law is
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[A]

= k[A]n [B]m
t
For a run where [A]0 = 1.0  103 M and [B]0 = 5.0 M, a plot of ln [A] versus t was found to
give a straight line with slope = 5.0  102 s1.
For a run where [A]0 = 1.0  103 M and [B]0 = 10.0 M, a plot of ln [A] versus t was found to
give a straight line with slope = 7.1  102 s1.
What is the value of n?
12. For the reaction 2N2O5(g)  4NO2(g) + O2(g), the following data were collected:
t (minutes) [N 2 O5 ] (mol / L)
0
1.24  10  2
10.
0.92  10  2
20.
0.68  10  2
30.
0.50  10  2
40.
0.37  10  2
50.
0.28  10  2
70.
0.15  10  2
The order of this reaction in N2O5 is?
The concentration of O2 at t = 10. minutes is
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Kinetics part II
1. Determine the molecularity of the following elementary reaction: O3  O2 + O.
2. The decomposition of ozone may occur through the two-step mechanism shown:
step 1 O3  O 2 + O
step 2 O3 + O  2O 2
The oxygen atom is considered to be a(n)
3. The following questions refer to the reaction 2A2 + B2  2C. The following mechanism has
been proposed:
step 1 (very slow) A 2 + B2  R + C
step 2 (slow)
A2 + R  C
What is the molecularity of step 2?
4. The following questions refer to the reaction 2A2 + B2  2C. The following mechanism has
been proposed:
step 1 (very slow) A 2 + B2  R + C
step 2 (slow)
A2 + R  C
Which step is “rate determining”?
5. The following questions refer to the reaction 2A2 + B2  2C. The following mechanism has
been proposed:
step 1 (very slow) A 2 + B2  R + C
step 2 (slow)
A2 + R  C
According to the proposed mechanism, what should the overall rate law be?
6. The decomposition of N2O5(g) to NO2(g) and O2(g) obeys first-order kinetics. Assuming the
form of the rate law is
[ N 2 O5 ]
Rate  
 k[ N 2 O5 ]
t
where k = 3.4  105 s1 at 25C.
What is the initial rate of reaction at 25C where [N2O5]0 = 5.0  102 M?
7. If the reaction 2HI  H2 + I2 is second order, what graph will yield a linear plot?
8. The reaction 2NO + O2  2NO2 obeys the rate law
[O2 ]
= kobsd [NO]2 [O2 ].
t
Which of the following mechanisms is consistent with the experimental rate law?
A)
NO + NO  N2O2
(slow)
N2O2 + O2  2NO2
(fast)
B)
O2 + O2  O2 + O2·
(slow)

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O2 + NO  NO2 + O
(fast)
O + NO  NO2
(fast)
C)
2NO
N2O2 (fast equilibrium)
N2O2  NO2 + O
(slow)
NO + O  NO2
(fast)
D)
NO + O2
NO3 (fast equilibrium)
NO3 + NO  2NO2
(slow)
E)
none of these
9. The questions below refer to the following diagram:
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Why is this reaction considered to be exothermic?
10. The questions below refer to the following diagram:
At what point on the graph is the activated complex present?
If the reaction were reversible, would the forward or the reverse reaction have a higher
activation energy?
12. The reaction 2H2O2  2H2O + O2 has the following mechanism?
H 2 O 2 + I   H 2 O + IO 
H 2 O 2 + IO   H 2 O + O 2 + I 
The catalyst in the reaction is:
13. When ethyl chloride, CH3CH2Cl, is dissolved in 1.0 M NaOH, it is converted into ethanol,
CH3CH2OH, by the reaction
CH3CH2Cl + OH  CH3CH2OH + Cl
At 25C the reaction is first order in CH3CH2Cl, and the rate constant is 1.0  103 s1. If
the activation parameters are A = 3.4  1014 s1 and Ea = 100.0 kJ/mol, what will the rate
constant be at 40C?
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14. 2002-54Which of the following must be true for a reaction for which the activation energy is
the same for both the forward and the reverse reactions?
A)
S for the reaction is zero.
B)
A catalyst is present.
C)
The reaction order can be obtained directly from the balanced equation.
D)
H for the reaction is zero.
E)
The reaction order is zero.
15. 2 NO( g )  O2( g )  2 NO2( g )
2002-27A possible mechanism for the overall reaction represented above is the following.
(1) NO( g )  NO( g )  N 2O2( g ) slow
(2) N 2O2( g )  O2( g )  2 NO2( g ) fast
Which of the following rate expressions agrees best with this possible mechanism?
A)
Rate  k [ N 2 O 2 ][O 2 ]
B)
Rate = k[NO]2
C)
[ NO]
Rate  k
[O2 ]
D)
Rate  k [ NO]2[O 2 ]
E)
[ NO]2
Rate  k
[O2]
16.
The graph above is indicative of what type of reaction?
17. The rate law for a reaction is found to be Rate = k[A]2[B]. Which of the following
mechanisms gives this rate law?
I. A + B
E (fast)
E + B  C + D (slow)
II. A + B
E (fast)
E + A  C + D (slow)
III. A + A  E (slow)
E + B  C + D (fast)
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Accessing Prior Knowledge Acids and Bases
1. Define acids and bases from three viewpoints, Arrhenius, Bronsted Lowry and Lewis
2. What do the [H+] and [OH-] always multiply together to equal?
3. What do pH and pOH always add up to?
4. Write the formulas/names for the follwing:
5. Nitric acid
6. Hydrocyanic acid
7. Chloric acid
8. Acetic acid
9. Hydrobromic acid
10. Sulfurous acid
11. Chlorous acid
12. Boric acid
13. Hydrochloric acid
14. Phosphoric acid
15. Nitrous acid
16. Hydrofluoric acid
17. Perchloric acid
18. Hydroiodic acid
19. Phosphorous acid
20. Carbonic acid
21. Sulfuric acid
22. Formic acid (the formate ion is COOH-)
23. Thiocyanic acid
24. HClO4
25. HCOOH
26. H3PO4
27. HCl
28. H3BO3
29. H2SO4
30. HNO2
31. HI
32. CH3COOH
33. HF
34. H3PO3
35. HCN
36. HClO3
37. H2CO3
38. H2SO3
39. HClO2
40. HNO3
41. HBr
42. H2S2O3
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pH PRACTICE
1. Write the pH of each solution above the [H+]’s. pH = -log[H+]
2. Label the “Z” diagram as “Acidic”, “Basic” and “Neutral.”
3. Knowing that the [H+] x [OH-] always equals 1 x 10-14, fill in the [OH-] for each of
the five solutions in the “Z” Diagram. (1 x 10-14 is called the Dissociation Constant
for water, Kw) (1 x 10-14 = 10 x 10-15)
4. Write the “pOH” of each solution below the [OH-]’s.
5. pH + pOH always equals _______.
6. A solution of acid has [H+] = 3.0 x 10-3 M
a. Calculate the [OH-] ________
b. Calculate the pH _________
the pOH _________
7. A solution of base has an [OH ] = 4.25 x 10-5 M
a. Calculate the [H+] _________
b. Calculate the pH _________
the pOH _________
8. Calculate the pH’s of the following solutions:
2.53 x 10-2M HCl
pH =
2.53 x 10-4M HCl
pH =
2.53 x 10-5M HCl
pH =
A pH with 3 significant figures is written with ____ numbers after the decimal place.
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Ka and Kb practice
1. Benzoic acid, HC6H5CO2, is an organic acid whose sodium salt, NaC6H5CO2, has
long been used as a safe food additive to protect beverages and many foods against
harmful yeasts and bacteria. The acid is monoprotic. Write the equation for its Ka.
2. The [H+] of a 0.10 M solution of cyanic acid (HCNO) is found to be 0.0010 M.
Calculate the Ka for cyanic acid. HCNO↔H+ + CNO3. If 1.22 grams of benzoic acid, HC6H5CO2, is dissolved in 1.0 L of water, the [H+] is
found to be 8.0 x 10-4M. Calculate the Ka for benzoic acid. HC6H5CO2↔H+ +
C6H5CO24. A 0.0050 M solution of butyric acid, HC4H7, has a pH =4.0, calculate Ka.
HC4H7O ↔ H+ + C4H7O25.
Determine the [OH-] and the [H+] of a 0.20 M solution of formic acid. The Ka = 1.8
x 10-4. HCOOH ↔ H+ + HCOO6. HCN has an initial molarity of 0.50 M, with a Ka value of 3.7 x 10-8. Calculate its pH
at equilibrium.
7. Ethylamine (C2H5NH3) is a weak Bronsted-Lowry base. If it has an initial molarity of
0.24M and a Kb of 5.6 x 10-4, calculate its pH at equilibrium.
8. A chemist adds 0.75 moles of NH3 to enough water to make 0.50 liters of solution.
-5
.Kb of ammonia is 1.8 x 10 . Determine the pH of this solution at equilibrium.
9. Hydrazine, N2H4, has been used as a rocket fuel. Like ammonia, it is a Bronsted base.
A 0.15 M solution has a pH of 10.70. What is the Kb for hydrazine?
10. Nicotinic acid, HC2H4NO2 is a B vitamin. It is also a weak acid with Ka=1.4 x 10-5.
What is the [H+] and the pH of a 0.010 M solution?
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Titration Curve Practice
1) First, identify the type of titration simply by looking at the titration curve
a. is it a strong acid being titrated with a strong base?
b. is it a strong base being titrated with a strong acid?
c. is it a weak acid being titrated with a strong base?
d. is it a weak base being titrated with a strong acid?
e. is it a weak base being titrated with a weak acid
f. If an acid is being titrated, is it a mono-, di-, or tri-protic acid
2) You need to fill in the blanks at the top of each titration curve in the following manner:
a. fill in the first blank with the type of solution being titrated (what is in the
beaker/flask being analyzed)
b. fill in the second blank with the titrant (the solution being delivered from the buret
Example: if a weak acid was in the beaker and it was being titrated with a strong
base from the buret, you would write: Weak acid/Strong Base
Then…
1)
Analyze each curve to find the equivalence point(s)
2)
Using the equivalence point(s);
a. Determine the Ka(’s) or Kb(‘s), when applicable
b. Use your text to identify an effective indicator that could be used to show the
endpoint in a titration
c. Determine the concentration of the sample being titrated, assuming the titrant is 0.100
M and that 10.0 mL of sample was titrated each time.
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____________ ________ - __________ _______ titration
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____________ ________ - __________ _______ titration
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____________ ________ - __________ _______ titration
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____________ _______ - _________ _______ Titration
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____________ ________ - __________ _______ titration
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____________ ________ - __________ _______ titration
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Chapter 14 and 15 practice:
In your group be sure to show all work associated with the problem. If you would like to see
the answer, you must first show me the work to the problem.
1. 2NH3  NH4+ + NH2In liquid ammonia, the reaction represented above occurs. In the reaction NH4+ acts as
2. At 25C, aqueous solutions with a pH of 8 have a hydroxide ion concentration, [OH-], of
3.
4.
5.
6.
7.
8.
What part of the curve corresponds to the optimum buffer action for the acetic acid/acetate
ion pair?
How can 100. mL of sodium hydroxide solution with a pH of 13.00 be converted to a sodium
hydroxide solution with a pH of 12.00?
Mixtures that would be considered buffers include which of the following?
I.
0.10 M HCl + 0.10M NaCl
II.
0.10MHF + 0.10M NaF
III.
0.10M HBr + 0.10M NaBr
A) I only
B) II only
C) III only
D) I and II
E) II and III
Ascorbic acid, H2C6H6O6(s) is a diprotic acid with Ka1=7.9x10-5 and Ka2=1.6 x10-12. In a
0.005M aqueous solution of ascorbic acid, which of the following species is present in the
lowest concentration?
A) H2O(l)
B) H3O+(aq)
C) H2C6H6O6(aq)
D) HC6H6O6-(aq)
E) C6H6O6-2(aq)
In a saturated solution of Zn(OH)2 at 25°C, the value of [OH-] is 2.0 x 10-6 M. What is the
value of the solubility product constant, Ksp, for Zn(OH)2 at 25°C?
For the stepwise dissociation of aqueous H3PO4, which of the following is not a conjugate
acid–base pair?
A) H2PO4 and PO43
B) HPO42 and PO43
C) H3PO4 and H2PO4
D) H3O+ and H2O
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
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2


E) H2PO4 and HPO4
What is the Ka expression for HOCl?
Calculate the [H+] in a solution that has a pH of 2.30.
Consider the reaction HNO2(aq) + H2O(l)
H3O+(aq) + NO(aq). Which species is the
conjugate base?
Calculate the pH of 0.250 M HNO3(aq).
For weak acid, HX, Ka = 1.0  106. Calculate the pH of a 0.10 M solution of HX.
Determine the molarity of a solution of the weak acid HClO2 (Ka = 1.10  102) if it has a pH
of 1.25.
Calculate the pOH of a 0.10 M solution of Ba(OH)2.
What is the pH of a 0.45 M KCl solution?
solid sodium carbonate (Na2CO3)
Use the following choices to describe an aqueous solution made from this substance.
Acidic, neutral, basic or cannot tell
solid ammonium perchlorate (NH4ClO4)
For NH4+, Ka = 5.6  1010; for ClO4, Kb  1021.
Use the following choices to describe an aqueous solution made from this substance.
Acidic, neutral, basic or cannot tell
The pH of a 1.0 M sodium acetate solution is:
A) 7.0
B) less than 7.0
C) not enough information is given
D) greater than 7.0
A 1.0-liter solution contains 0.25 M HF and 0.60 M NaF (Ka for HF is 7.2  104)
What is the pH of this solution?
Calculate the pH of a solution that is 0.5 M in HF (Ka = 7.2  104) and 0.6 M in NaF.
For a solution equimolar in HCN and NaCN, which statement is false?
A) This is an example of the common ion effect.
B) The [H+] is equal to the Ka.
C) Addition of NaOH will increase [CN and decrease HCN
D) Addition of more NaCN will shift the acid dissociation equilibrium of HCN to the
left.
E) The [H+]is larger than it would be if only the HCN was in solution.
How many moles of solid NaF would have to be added to 1.0 L of 1.90 M HF solution to
achieve a buffer of pH 3.35? Assume there is no volume change. (Ka for HF = 7.2  104)
What is the molarity of a sodium hydroxide solution if 25.0 mL of this solution reacts exactly
with 22.30 mL of 0.253 M sulfuric acid?
A 10-mL sample of tartaric acid is titrated to a phenolphthalein endpoint with 20. mL of 1.0
M NaOH. Assuming tartaric acid is diprotic, what is the molarity of the acid?
If 25 mL of 0.75 M HCl are added to 100 mL of 0.25 NaOH, what is the final pH?
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Ksp practice. Keep me but put all answers in your notes.
1.
2.
3.
4.
How many moles of solid NaF would have to be added to 1.0 L of 1.90 M HF
solution to achieve a buffer of pH 3.35? Assume there is no volume change. (Ka for
HF = 7.2  104)
Calculate the pH of a solution made by mixing 100.0 mL of 0.300 M NH3 with 100.0
mL of 0.100 M HCl. (Kb for NH3 is 1.8  105).
Find the solubility (in mol/L) of lead chloride (PbCl2) at 25C. Ksp = 1.6  105
Solubility Products (Ksp)
BaSO 4 1.5 109
CoS
5.0 10 22
PbSO 4
1.3 108
AgBr
5.0 1013
Which of the compounds is the most soluble (in moles/liter)?
5.
You have a solution consisting of 0.10 M Cl and 0.10 M CrO42. You add 0.10 M
silver nitrate dropwise to this solution. Given that the Ksp for Ag2CrO4 is 9.0  1012,
and that for AgCl is 1.6  1010, which of the following will precipitate first?
A) silver chromate
B) cannot be determined by the information given
C) silver nitrate
D) silver chloride
6.
The solubility of CaSO4 in pure water at 0C is 1 gram per liter. The value of the
solubility product is
7.
The molar solubility of PbI2 is 1.5  103 M. Calculate the value of Ksp for PbI2.
8.
Calculate the concentration of chromate ion, CrO42, in a saturated solution of
CaCrO4. (Ksp = 7.1  104)
9.
Calculate the concentration of the silver ion in a saturated solution of silver chloride,
AgCl (Ksp = 1.6  1010).
10.
The molar solubility of BaCO3 (Ksp = 1.6  109) in 0.10 M BaCl2 solution is:
11.
It is observed that 7.5 mmol of BaF2 will dissolve in 1.0 L of water. Use these data to
calculate the values of Ksp for barium fluoride.
12.
The Ksp of AgI is 1.5  1016. Calculate the solubility in mol/L of AgI in a 0.30 M
NaI solution.
13.
The molar solubility of AgCl (Ksp = 1.6  1010) in 0.0020 M sodium chloride at
25C is:
14.
Silver chromate, Ag2CrO4, has a Ksp of 9.0  1012. Calculate the solubility in mol/L
of silver chromate.
15.
The solubility of mol/L of Ag2CrO4 is 1.3  104 M at 25C. Calculate the Ksp for
this compound.
16.
Calculate the concentration of Al3+ in a saturated aqueous solution of Al(OH)3 (Ksp =
2  1032) at 25C.
17.
The Ksp for PbF2 is 4.0  108. If a 0.050 M NaF solution is saturated with PbF2,
what is the Pb2+. in solution?
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18.
Chromate ion is added to a saturated solution of Ag2CrO4 to reach 0.10 M CrO42.
Calculate the final concentration of silver ion at equilibrium (Ksp for Ag2CrO4 is 9.0
 1012)
19.
In a solution prepared by adding excess PbI2(s) Ksp = 1.4  108. to water, the I. at
equilibrium is:
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Electrochem practice
1. Which energy conversion takes place in a galvinic cell?
2. Which of the following reactions is possible at the anode of a galvinic cell
A) Zn --> Zn2+ + 2eB) Zn2+ + 2e- --> Zn
C) Zn2+ + Cu -->Zn + Cu2+
D) Zn + Cu2+ --> Zn2+ + Cu
3.A block of Pt is immersed in a solution which has 0.50 M Br2 and 0.10 M Br-. A block of chromium is
immersed in a solution of 0.20 M Cr3+. The wires are connected to a voltmeter.
a. What type of cell is this?
b. What is Eo for this cell
c. What is E for this cell at 25oC
d. How many moles of electrons are involved in the reaction?
7.Consider a galvanic cell with a zinc electrode immersed in 1.0 M Zn2+ and a silver electrode immersed
in 1.0 M Ag+.
a. Calculate E0 for this cell
b. Which of the electrodes in the anode?
c. Calculate delta G
d. Calculate K for the cell
11. Consider a galvanic cell with a zinc electrode immersed in 0.050 M Zn2+ and a silver electrode
immersed in 10.00 M Ag+.
Calculate Q for the cell.
E for the cell at these non standard conditions. Assume Temp is 25oC
12. An antique car bumper is to be chrome plated. The bumper is dipped into acid dichromate solution
and the bumper serves as the cathode of the electrolytic cell. The dichromate deposits solid
chromium onto the surface of the bumper. If the current is 10 amperes, how long does it take to
deposit 1.00 x 102 g of Cr(s)
13. What quantity of coulombs is required to reduce 40.0 g of CrCl3 to chromium metal?
14. If a constant current of 5.0 amperes is passed through a cell containing Cr3+ for 1.0 hour, how many
grams of Cr will plate out on the cathode?
15. How many seconds would it take to deposit 21.40 g of silver from a solution of silver nitrate using a
current of 10.00 amperes?