LC AP Chem Labs – ctrl +left click to open
Ch 1 Indirect Measurements
Ch 1 Qualitative Analysis
Ch 1 Measurements: Precision and Accuracy
Ch 2 Synthesis of a Complex Iron Salt
Ch 2 Determination of the Charge of an Unknown Anion
Ch 2 Identification of Cations and Anions by Qualitative Analysis
Ch 3 The Chemistry of Copper
Ch 3 Determination of the % Oxalate in a Complex Salt
Ch 3 Determination of an Empirical Formula
Ch 3 Determining Mole Ratios in a Chemical Reaction
Ch 3 Determination of the Formula of an Iron Chloride Using Stoichiometry
Ch 4 Double Displacement Reactions
Ch 4 Acid-Base Titrations
Ch 4 Determination of Phosphate in Beverages
Ch 4 Analysis of Hydrogen Peroxide
Ch 4 Determination of the Concentration of Acetic Acid in Vinegar
Ch 5 Additivity of Heats of Reactions
Ch 5 Determination of the Caloric Content of Various Nuts
Ch 6 The Atomic Spectrum of Hydrogen
Ch 6 Determination of Iron by Visible Spectrophotometry
Ch 10 The Gas Law Constant
Ch 10 Identification of Two Unknown Metals through the Production of H2
Ch 11 Intermolecular Attractions
Ch 11 Determination of a Heat of Vaporization
Ch 11 Identification of an Unknown Liquid by Vapor Density
Ch 13 Molar Mass from Freezing Point Depression
Ch 13 Identification of an Unknown Solid by Freezing Point Depression
Ch 13 Determination of the % Water in an Iron Oxalato Complex Salt
Ch 13 Determination of the Percentage of Water in Copper(II) Sulfate
Ch 14 Determination of a Rate Law and Activation Energy using Crystal Violet
and Sodium Hydroxide
Ch 14 Determination of a Rate Law and Activation Energy using Thiosulfate
Ions and HCl
Ch 14 Detrmination of Aspirin by Visible Spectrophotometry
Ch 14 Determining the Order of a Reaction using UV-Vis Spectroscopy
Ch 15 Determination of an Equilibrium Constant
Ch 15 Determination of an Equilibrium Constant using Spectrophotometry
Ch 16 Dissociation Constants for Weak Acids
Ch 16 Titration Curves & Indicators
Ch 16 Identification of an Unknown Acid by Titration
Ch 17 Solubility of Calcium Iodate
Ch 17 Using Conductivity to Find an Equivalence Point and Ksp Value
Ch 17 Determination of the Ksp for Silver Acetate using the Mohr Method
Ch 20 Determination of the Percentage of Potassium & Iron by Ion Exchange
and Potentiometric Titration of an Iron Oxalato Complex Salt
Ch 20 Standardization of KMnO4
Ch 20 Electrochemical Cells and Thermodynamics
Ch 20 Determination of Thermodynamic Data from Standard Cell Potentials
Ch 20 Oxidizing & Reducing Agents
Ch 20 Reduction Potential Series
TLClab Synthesis of Aspirin
TLClab One Tube Reactions
TLClab Copper - Silver Nitrate REACTION
TLClab The 12 Bottle Problem
TLClab The Synthesis & Analysis of Alum
TLClab Synthesis, Purification, and Characterization of Aspirin
TLClab Isolation and purification of nicotine from tobacco
Experiment 1: Indirect Measurements
At times, you will be expected to develop parts or all of a procedure to follow while you are in the
laboratory. This is one of those times.
There are instances when a particular dimension of a "piece" of matter is needed, but it is too small
to measure directly. In those instances, one must approach the problem in an indirect manner. That is
your task in this experiment.
Part I: Place a small quantity of zinc granules into one well of a spot plate
and a small amount of iron filings into another well. Using a
dropper, place several drops of 6 M HCl on both the metallic zinc
and iron. (A spot plate is ceramic plate with small wells or
depressions built into it, often used for qualitative analysis. The
wells (or spots) are used to perform reactions on a very small amount
of materials.) Which metal reacts significantly faster?
Part II: Galvanized iron is made by placing a thin coating of zinc on the
surface of iron. Obtain a sample of the galvanized iron. Using the
results from part I, devise a procedure which will allow you to
“separate” the zinc and iron. The data you collect should allow
you to indirectly determine the thickness of the zinc on one side
of the iron. (The theoretical density of Zn is 7.13 g/cm3.) You
may also use a Vernier caliper and analytical balance. Enter your
data onto the computer, so class results can be analyzed.
Zn
Fe
Calculations: Determine the thickness of the zinc on one side of the galvanized iron in µm. Use the 4d
Rule to eliminate any outlying results. Calculate your precision and the class precision after using the 4d
Rule. Zinc atoms have an atomic radius of 135 pm. Determine how many
layers of zinc atoms are on one side of the galvanized iron (a) if they are
stacked in a simple cubic arrangement and (b) if they are stacked in a facecentered cubic arrangement (consult your text to learn about these cubic
arrangements). You do NOT need to do any statistics on the calculations
involving the number of layer of zinc atoms.
Summary: Report all results. Discuss the individual and class precision. What
assumption(s) about the galvanized iron must be made in this determination?
What are some possible errors in the experiment and what effect would they
have on the results?
Vernier Caliper
Synthesis of a Complex Iron Salt
Introduction: In this experiment a complex compound containing the elements potassium, iron, carbon,
hydrogen, and oxygen will be synthesized. Carbon and oxygen will be present in the compound as the
oxalate ion (C2O42-) whereas hydrogen and oxygen will be present as water. The final product, consisting
of emerald green crystals, may be given the empirical formula KwFex(C2O4)y.zH2O, where the zH2O is
called the water of hydration. This will be the first of a series of experiments in which the complex salt
will be synthesized, and then its simplest formula (i.e. w, x, y, z) will be determined, using a variety of
analytical techniques.
One important factor in any chemical synthesis is the actual quantity of desired product obtained compared
to the theoretical amount predicted on the basis of the stoichiometry of the reaction. The ratio of the mass
of the product obtained to the theoretical quantity, expressed as a percentage, is referred to as the 'percent
yield' or more simply the 'yield'.
There are many reasons why the actual yields are not 100%. Possibly the reaction reaches equilibrium
before going to completion. Maybe the reactants are involved in reactions other than the one that produces
the desired product. Probably some product is lost in crystallizing and separating crystals from the
supernatant liquid, etc.
In this experiment an aqueous solution containing about 4 grams of FeCl3.6H2O will be reacted with an
aqueous solution containing excess K2C2O4.H2O to produce KwFex(C2O4)y.zH2O.
Questions:
(1) IF all of the iron (Fe) originally in FeCl3.6H2O ends up in the product, KwFex(C2O4)y.zH2O, how
many moles of product should be obtained?
(2) What additional information would be needed to calculate a 'percent yield'?
Part A: Synthesis of a Complex Iron Salt
Objective: To prepare several grams of pure emerald green crystals of KwFex(C2O4)y.zH2O.
Chemicals: FeCl3.6H2O solution, solid K2C2O4.H2O, ice, acetone
Safety, Environmental, and Economic Concerns:
1. Waste solutions may safely be discarded down the drain. Flush with excess water.
2. Avoid overheating since overheating may result in violent boiling when heating liquids.
3. ACETONE IS A FLAMMABLE SOLVENT! Extinguish all flames and turn off hot plates in your
work area before using acetone. Make sure that the acetone bottle is covered when not in use. If a
spillage occurs, make sure all burners and hot plates in the area are turned off, wipe up the spilled liquid
with a dry paper towel, the use a wet paper towel to wash down the area.
Notes on Experimental Procedure:
1. When a desired product is formed by crystallization from a reaction mixture containing excess
reactants and other products, the crystals are likely to be relatively impure. The crystals can be
separated from the impure solution (often called the "mother liquor") by filtration or decantation. The
mother liquor clinging to the crystals can be removed by washing with an appropriate solvent.
However, washing will not remove impurities occluded within the crystals.
A standard method for purifying a crystalline product is recrystallization. For crystals that are more
soluble in hot solvent than in cold solvent, the recrystallization can be done by dissolving the crystals
in a minimum quantity of hot solvent and then cooling it in an ice bath. The purified crystals can then
be "harvested" by filtration.
A second "crop" of crystals can be obtained from the filtrate by evaporating a fraction of the solvent by
heating, followed by cooling the remaining solution in an ice bath. The second crop of recrystallized
product is generally less pure than the first.
2. The wet crystals dry very slowly. The purpose of the acetone is to wash the water off the crystals. The
acetone, which has a high vapor pressure, evaporates quickly, thus leaving the crystals dry.
Since acetone is flammable, make sure there are no flames on your work bench when you do your
acetone washes.
3. The product will slowly decompose when exposed to light. Hence the crystals should be stored in the
dark, or in a brown amber bottle.
Procedure:
Day 1
1. Obtain (in a clean, dry small beaker) 10.00 ml of the stock solution of iron. (Record the precise
concentration from the stock bottle).
2. Weigh about 12 grams of potassium oxalate monohydrate (K2C2O4.H2O) into another clean, dry small
beaker, using the Mettler balance (record to 0.0001 gram). Add about 20 ml of distilled water to
dissolve the salt (K2C2O4.H2O). Heat and stir to dissolve the salt (take care not to heat too strongly).
3. Using beaker tongs to handle the hot beaker, pour the hot solution into the beaker containing the
iron(III) chloride solution and stir.
4. Cool the solution for 30 - 45 minutes by placing the beaker in a larger beaker containing ice and water.
Crystals should form during this time. Take care that the beaker does not sink into the ice water.
5. After giving the crystals ample time to form, carefully pour off and discard the solvent without
removing any crystals - a process called decanting. [Decant into a clean, dry vessel. If any material
leaves the beaker that shouldn’t, you can combine everything and start over.] Add 20 ml distilled
water to the crystals. Heat gently with stirring to completely dissolve the crystals. If some dark
residue remains undissolved, carefully decant the clear solution into another beaker and discard the
residue.
6. Label the beaker with your name(s) and lab time. Cover the beaker with a watch glass and set it in a
drawer or cabinet until the next lab period in order to allow the crystals to form. If the crystals are
allowed to form slowly without being disturbed, large crystals will be obtained. If the solution is
moved, stirred, or disturbed while the crystals are forming, smaller crystals will result. At some point,
clean and dry a small brown bottle (your own prescription bottle?). Have it available for later use.
Day 2
7. After the crystals are formed, set up a vacuum filtration system
and filter the crystals using a Buchner funnel and a clean filter
flask. Make sure the filter paper is properly "seated".
8. Wash the crystals twice with ice water (distilled). [chill the
water in a graduated cylinder in an ice bath] Use less than 5 ml
of ice water for each wash and work quickly to avoid dissolving
the product in the wash water. Finally rinse the crystals twice
with 5 ml portions (aliquots) of acetone.
9. Spread the crystals in the bottom of a clean, dry, labeled 250-ml
beaker and set it aside to air dry in a drawer or cabinet until the
next lab period.
Day 3
10. Using the Mettler balance, weigh to the nearest 0.0001 gram the clean, dry brown bottle.
11. Place the dry crystals in the pre-weighed brown bottle and weigh again to the nearest 0.1 milligram.
Store these crystals in the capped amber bottle for use in future experiments. A minimum of 3.5 grams
of product will be needed in subsequent experiments. If your yield is less than 3.5 grams, consult the
instructor. Do NOT put the bottle of green crystals in a desiccator or in the drying ovens for further
drying! [Question: Why would this be a bad technique?]
The Chemistry of Copper
In this experiment you will qualitative study a series of reactions of copper. The series begins and ends
with solid copper.
Safety: This experiment uses concentrated nitric acid. To contain the fumes, assemble an individual fume
hood consisting of an inverted funnel connected to the water aspirator with a piece of rubber tubing.
Remember that concentrated acids are dangerous, especially to skin and the respiratory system. The gas
generated in the reaction is also very dangerous if inhaled. Minimize breathing any of the vapors in this
experiment.
PROCEDURE
[The instructor will do the following as a demonstration. Take notes. The quantities given are for
a pair of students in the lab.] Measure about one gram of solid copper on a pan balance and place the Cu
in an Erlenmeyer flask. The copper should be in relatively small pieces. Place the flask in the fume hood.
Add a maximum of 5 mL of nitric acid.
Measure about 10-mL of the prepared solution into a 125-mL Erlenmeyer flask. Place the flask in
a cool tap water bath. Add 6 M NaOH slowly to the solution and watch the reaction. Heat may be
evolved, so control the temperature to keep the reaction under control. Add NaOH until no more solid
forms. Sufficient NaOH (< 5 mL) must be added to both neutralize excess acid and form the hydroxide
precipitate. The reaction is complete when no more solid forms. Mix the contents of the flask by
SLOWLY swirling the flask. Wash down any solid that clings to the inside of the flask with about 2-3 mL
of distilled water using your distilled water bottle.
Place the flask and contents into a hot tap water bath and heat the water to boiling. Stir while
heating. This dehydrates the previous precipitate. Heat until you see a separation between the solid and
solution. Cool the contents of the flask and allow the solid to settle.
Decant the liquid from the solid and wash the solid with 10-15 mL of distilled water. Decant the
wash water. Repeat the wash two more times. Add 6 M HCl slowly to the solid while stirring. HCl
should be added until the solid just disappears. Then add 1 mL in excess.
Add about 1 g of aluminum to the contents of the flask. How can you tell if sufficient Al is
“available?”
The reactions are now complete. Place several paper towels in the sink and carefully pour the
contents of the flask onto the towel, collecting the solid. Place the towels and solid in the waste container.
The experiment is finished!
Your report must include a balanced ‘molecular’ equation for each and every reaction that
occurred, including side reactions. Correlate all observations with the chemical equations [this typically
means you are to write a descriptive phrase under each formula in the equation (not water)].
Double Displacement Reactions
In the table below are the 12 solutions with which you will experiment. You are to combine all
possible combinations of pairs of solutions, making and recording observations. Initially, use just a few
drops of each solution into a microplate for each combination. You do not need to add a chemical to itself.
Do add the chemicals to each other both ways (i.e. Add LiOH(aq) to KSCN(aq) and add KSCN(aq) to
LiOH(aq). The apparent reaction may depend on sequence! You will observe that some combinations
produce precipitates, some produce a gas, and others appear to do nothing (Make special note of these
combinations for use in an experiment second semester.). If a precipitate forms, you need to record the
color & type of precipitate. Some precipitates dissolve as more chemical is added (a complex ion may be
formed). A matrix is handy for recording data in this type of experiment. [Make sure you have sufficient
room to record your observations in the matrix.]
Make a table and record the color of each original solution. Also, conduct and record the color of
the flame test for each original solution. [Soak a cotton swab in the solution and place it at the tip of the
blue cone in the flame. Use a new swab for each solution. You might want to look at
http://chemistry.about.com/library/weekly/aa110401a.htm for reference.]. Determine the pH of each
original solution. [Tear about 1 cm of pH paper off the roll and place on a paper towel. Using a glass
stirring rod, place a drop of solution on the pH paper. Record the pH to which the color of the wet paper
indicates. Use a new piece of pH paper for each solution.] You will use this information in an experiment
second semester.
1. FeCl3
2. Ba(NO3)2 3. LiOH 4. KSCN 5. H2SO4
6. AgNO3
7. Al(NO3)3 8. CuSO4
9. SrCl2 10. NH3* 11.Na2CO3 12. CoCl2
*NH3(aq) is the proper designation for NH4OH(aq) which is better to use in equations
You may use information in Chapter 4 of the text and the Ultimate Chemical Equations Handbook
to assist you. For each combination that appeared to react, write a net ionic equation. Include a short
observation under each chemical species written in the net ionic equations. [To decrease the total work,
you may write one (1) equation for several reactions, if it is appropriate. However, please state what
combination of molecular reactants are being presented.] For combinations that appear not to react, write
the ionic reactants  NAR [no apparent reaction]. [Suggestion: for organizational purposes in the
summary, you might want to categorize the various combinations according to the commonality of
results.] From these results, devise a set of solubility rules for the cations and anions.
AP Chemistry
FeCl3
DOUBLE DISPLACEMENT REACTIONS
Ba(NO3)2
LiOH
KSCN
H2SO4
AgNO3 Al(N03)3 CuSO4
SrCl2
NH3
FeCl3
Ba(NO3)2
LiOH
KSCN
H2SO4
AgNO3
Al(N03)3
CuSO4
SrCl2
NH3(aq)
Na2CO3
CoCl2
The matrix above represents a “data” table where 0.1 M solutions of each chemical are mixed, two at a
time. Using the Solubility Tables and other resources in the “Ultimate Chemical Equations Handbook,”
predict what product, if any, would result. Write the formula of the product(s) in the blank. If no reaction
would occur, write NR. For every combination where a reaction occurred, write a net ionic equation
(NIE). Do NOT write duplicate NIEs. Before a NIE, write the formulas of the pair of chemicals that
produce the predicted results.
Your score on this activity will be entered in the Lab portion of the grades. These predictions will be
useful in one of the experiments in the spring semester.
Na2C
Acid-Base Titrations
Parts A & B of this experiment, both acid-base titrations, requires a solution of sodium hydroxide
whose molarity is known accurately. Such a solution cannot be prepared by adding a carefully measured
mass of NaOH to an appropriate quantity of water because solid NaOH is very hygroscopic, meaning it
absorbs water directly from the air. As a result, a sample of this substance will be absorbing water as you
attempt to measure its mass. The recorded mass, no matter how carefully it was measured, will be a
combination of an unknown mass of NaOH and an unknown mass of water.
A.
Standardization of a base.
The problem of the water absorption by NaOH can be overcome by making the NaOH solution to
an approximate concentration and titrating the solution against an acid solution of known concentration or
a known mass of a primary standard acid. The solid primary standard used in this experiment is potassium
hydrogen phthalate, KHC8H4O4, which will be referred to throughout the experiment by the abbreviation
KHP. KHP is a monoprotic acid meaning that it provides one H+, (H3O+) ion per molecule or reacts in a
1:1 mole ratio with a monohydroxyl base, such as NaOH.
O
KHP(s) H

 K+(aq) + HPhthalate-(aq)
HPhthalate (aq) + OH-(aq)  H2O + Phthalate2-(aq)
2
Using the stoichiometry suggested by the equations above, the known mass of KHP used, and the
measured volume of NaOH solution, we will be able to calculate the molarity of the NaOH solution.
B.
Determining the unknown concentration of an acid
The standardized NaOH solution (a secondary standard) will now be used to determine the
unknown concentration of a sulfuric acid solution. The stoichiometric point in the reaction between the
acid and base solutions is approximated using phenolphthalein indicator. Phenolphthalein turns from
colorless in acid solution to pink in basic solution at a point that is close enough to the equivalence point to
be within experimental error. The use of an indicator will be discussed in more detail later in the course.
At this point be assured that the error introduced is small enough to be acceptable.
Preparation of the buret & pipet:
Wash the buret with a little soap and lots of tap water using a buret brush to scrub it clean. Do not
forget to run some soapy water and a lot of rinse water through the buret tip. Rinse the buret twice with
distilled water and then with 2-4 mL of the solution to be used in the buret during the experiment. You
may now fill the buret with the appropriate solution.
Pipets should never be washed with soap. It is very difficult to rinse soap out of a pipet. Pipets
should be rinsed with distilled water and the solution being measured in a similar fashion to burets.
NEVER pipet by mouth.
The two parts to the experiment are similar, but not identical. There are preparation steps that apply
to both. All glassware must be carefully washed, rinsed with distilled water, and drained before use. ALL
titrations will be done individually
A.
Standardization of the Base:
Prior to coming to the lab, determine the mass of KHP [molar mass of KHP is 204.23 g/mol.] that
would be needed to do a titration using the maximum capacity of the buret [50 mL], assuming the [NaOH]
is 0.1 M (show work). For each titration, use ⅔ to ¾ of this mass. A solution of approximately 0.1 M
NaOH will be provided. There is a limit of 200 mL per student.
Measure an appropriate mass of KHP on the centigram balance, then determine the precise mass on
the analytical balance. Add the KHP to a clean Erlenmeyer flask. Add about 50 mL distilled water. Be
sure the KHP is completely dissolved. KHP is a little slow getting the idea of how to dissolve. Be patient.
Add several drops of phenolphthalein to each flask and titrate (add the base from the buret slowly to acid
in the flask) until a PERMANENT PALE PINK color is observed. A GOOD titration is VERY pale!
Swirl the flask gently between additions of base to mix thoroughly. Record the volume (to the nearest
0.01 mL) of solution in the buret before and after the addition of the base. It is not necessary to have the
buret on zero before you start. Be careful to not overfill the buret. NEVER refill a buret during a titration.
This leads to cumulative reading errors. Three trials within 3% are sufficient. Record the mass of your
KHP used for your standard solution and the initial and final volumes of NaOH solution used for each
titration on the class spreadsheet. Determine the molarity of the NaOH for each trial. After analysis of the
class results, use the 4d rule to determine if any of the values should be rejected. Determine individual and
class precision of the results.
B.
Determination of the Molarity of a Sulfuric Acid Solution.
Prior to coming to the lab, write the balanced molecular equation for the reaction between sodium
hydroxide and sulfuric acid. Pipet 10.00 mL of the sulfuric acid solution into each of two 125-mL
Erlenmeyer flasks. Add some distilled water to the acid solution for volume to see. Add phenolphthalein
and titrate the base into the acid in the same manner described above for the standardization of the base.
[You need not calculate the molarity of the base to collect your data.] Since you pipetted a precise amount
of acid into each flask, you may test your reproducibility by comparing the volumes of base added to
achieve the end point. If your volumes disagree by more than 3%, perform a third trial. . Record the
initial and final volumes of NaOH solution used for each titration on the class spreadsheet. Rinse all
equipment with distilled water before putting it away. Use the class average [NaOH] to determine the
molarity of the sulfuric acid for each trial. After analysis of the class results, use the 4d rule to determine
if any of the values should be rejected. Determine individual and class precision of the results.
Additivity of Heats of Reaction
In this experiment you will build a calorimeter from two Styrofoam coffee cups. A model will be
shown to you in lab. Be sure to return the cups at the end of the experiment so they can be used by other
students. Temperatures will be taken with a thermometer.
Traditionally, the experiment consists of two parts. In the first part, the calorimeter constant or the
amount of heat taken or given by the calorimeter per degree of temperature change is determined. The
calorimeter constant is measured in the temperature range in which the experiment is to be conducted. The
calorimeter constant can be measured by mixing known masses of two samples of water that are at
different temperatures. Even though our “coffee cup” calorimeter is not perfect, it is really quite good.
Thus we will NOT determine a calorimeter constant, but let’s ASSUME the calorimeter constant = 11.2
J.deg-1.
The experiment may be performed in pairs. This handout merely outlines the experiment. It will
be necessary for you and your partner to prepare together, since you are expected to arrive in lab with a
procedure worked out to complete the experiment. It might be helpful to consult your textbook in deciding
what data are needed and how the experiment would best be performed.
You will use the coffee cup calorimeter to measure the heat released by three reactions. One of the
reactions is equal to the sum of the other two. Therefore, according to Hess' Law, the heat of the one
reaction should equal the sum of the heats of the other two. The three reactions are: the solution of solid
sodium hydroxide in water; the neutralization of aqueous sodium hydroxide by aqueous hydrochloric acid;
and the neutralization of solid sodium hydroxide by aqueous hydrochloric acid.
Before you come to lab, you should write the three “molecular” equations, determine the
theoretical heat of reaction for each reaction, and show the additivity you are going to study in the
experiment.
You must design the detailed procedure for conducting this experiment. Here are the parameters
under which you will work.
Equipment available
two Styrofoam coffee cups to make a calorimeter; a corrugated cardboard cover; thermometer
[a 400-mL beaker works well as a calorimeter holder]
Maximum quantities of chemicals available per pair of students:
3 g NaOH(s) 60 mL 1.0 M NaOH 110 mL 1.0M HCl
For the two neutralization reactions, the NaOH and HCl should be mixed in stoichiometric
quantities. All solutions should be 1 M since heats of formation of aqueous solutions are determined at
that concentration. Prior to coming to the lab, determine the quantities of chemicals you will use in each
reaction. After mixing, make sure you stir the solution continuously until a constant temperature is
achieved.
Don’t forget that precise masses of the reactants are required for heat calculations. Determine the
Heat of Reaction per mole NaOH for each system. You are expected to include a error analysis in your
report. The analysis should be in two parts: Internal Error (consistency), which is the change in enthalpy
for the net reaction compared to the sum of enthalpy changes for the other two reactions; and External
Error which is a comparison of your experimental values with the accepted values of each system.
The Atomic Spectrum of Hydrogen
From the text, you are aware that after electrons are excited from the ground state (by an external
source of energy), they naturally fall back to lower energy levels emitting energy in the process. It has
been determined that electrons that fall to the first energy level emit energy in the range of ultraviolet light
[Lyman Series], visible light is emitted if they fall to the second energy level [Balmer Series], and infrared
light is emitted if they fall to the third energy level [Paschen Series]. If one uses appropriate equipment, it
is possible to view the visible emission lines of substances and make appropriate measurements. Using
these measurements, one can determine the orbital number from which the electron fell.
At the left is a hydrogen spectral tube excited by
a 5000 volt transformer. The three prominent hydrogen
lines are shown at the right of the image. In this
demonstration, the instructor will show the visible
spectral lines of hydrogen as collected by a Vernier
spectrometer. You are to record the wavelength of 4-6
of these visible spectral lines. From this data, you are
to determine the number of the energy level from which
the electrons fell to produce each line using the
equations: c = , E = h,
 1
1 
 E   2.178  1018 J  2  2  ,
 n2 n1 
where n2 = 2 for visible light. [Remember that the light
is being emitted.] A sample calculation should be
shown for the determination of the energy emitted of one emitted line and the energy level from which the
electrons fell. Present a table which includes the wavelength, energy, and initial energy level for all
emitted lines (including the decimal equivalent before rounding to the integer energy level). Discuss the
relationships among these three entities.
Oxidation-Reduction Reactions
Now you will carry on a series of reactions and identify the oxidizing and reducing agents on the
basis of your observations. To help you identify products you will first note the characteristic colors of
solutions and perform a series of identifications tests. Remember, it is as important to record the initial
colors of chemicals as it is the color after a combination of chemicals are mixed. All quantities given are
only approximations and are not to be measured precisely. [In a 13 x 100 test tube, a depth of about 1 cm
is about 1 mL.] Keep in mind that this is a qualitative, not a quantitative, lab.
Procedure
1.
Observe the colors of all the provided solutions. You might need to use a spot plate. Use these and
previous observations to deduce the color of each cation and anion.
2.
The following reactions are intended to help you identify, or confirm, some of the previous ions
which may appear as products of redox reactions in subsequent procedures. In a small test tube,
mix each of the following combinations thoroughly.
a.
Add a drop of potassium thiocyanate, KSCN, to several drops of FeCl3.
b.
Add a drop of potassium ferricyanide [potassium hexacyanoferrate(III), K3Fe(CN)6] to
several drops of FeSO4.
c.
Add a couple of drops of BaCl2 to ~ 1 mL of AgNO3.
d.
Add a couple of drops of BaCl2 to a couple of drops of Na2SO3. Observe. Then add several
drops of 6 M HCl. [is the ppt soluble in HCl?]
e.
Add a couple of drops of BaCl2 to a couple of drops of H2SO4. Observe. Then add several
drops of 6 M HCl. [is the ppt soluble in HCl?]
f.
To ~ 1 mL of K2CrO4 add drops of HCl until a change occurs. To this, add NaOH until a
change again occurs.
Write a balanced net ionic equation for each reaction that occurred in step 2. Justify each reaction
written by relating to the observations.
3.
In each of the following reactions, the observations you make will “tell” you what chemical species
is present, based upon your observations in 1 and 2 above.
a.
Place ~ 2 mL 0.1 M Na2SO3 in a small test tube. Add ~ 0.5 mL of 6M HCl to make the
solution acidic. Add several drops 0.1% KMnO4. Observe. Add a few drops of BaCl2 .
b.
Repeat procedure 3a using a few drops of 0.1 M K2Cr2O7 instead of KMnO4.
c.
Repeat procedure 3a using a few drops of chlorine water, Cl2(aq), instead of KMnO4.
Transfer half of the solution to another test tube. To one test tube add a few drops of BaCl2
and to the other test tube add a few drops of AgNO3.
d.
Repeat procedure 3a using a few drops of concentrated HNO3 instead of KMnO4.
Write a balanced net ionic equation for each reaction that occurred in step 3. Justify each reaction
written by relating to the observations. For each reaction, identify the element oxidized, the element
reduced, the oxidizing agent, and the reducing agent.
4.
Similar to step 3, in each of the following reactions, the observations you make will “tell” you what
chemical species is present, based upon your observations in 1 and 2.
a.
To ~1 mL of 0.1 M CrCl3, add ~ 3 mL of 6 M NaOH drop wise. Stir the solution while
adding 6% H2O2 drop wise until a change occurs. If no color change is noted, gently warm
the solution. It may be necessary to add a little more H2O2 .
b.
To ~ 1 mL of 0.1 M K2CrO4 add 6 M HCl until the solution is orange. Then add a mL of
HCl in excess. Add 6% H2O2 drop wise with stirring and record any changes.
c.
Add 2 drops of 3 M H2SO4 and 4 drops of 0.4% KMnO4 to 2 mL of freshly prepared 0.1 M
FeSO4. Test for the presence of Fe(III) ions by adding one drop of KSCN.
d.
Add 6 drops of concentrated HNO3 to 1 mL of freshly prepared 0.1 M FeSO4. Test for the
presence of Fe(III) ions by adding one drop of KSCN.
e.
Add 1 mL of 0.1 M KI to 2 mL of 0.1 M FeCl3. Test for the presence of Fe(II) by adding a
drop of potassium ferricyanide solution. Note any evidence of I2. [To test for I2, add several
drops of chlorine water, add ~ 1 mL CH2Cl2 and observe.]
f.
To ~ 1 mL chlorine water, add several drops 6 M NaOH. To this, add a few drops
AgNO3(aq).
g.
To ~ 1 mL KI, add ~ 0.5 mL 6 M HCl to make the solution acidic. Then add several drops
0.1% KMnO4. Observe. Note any evidence of I2. [To test for I2, add several drops of
chlorine water, add ~ 1 mL CH2Cl2 and observe.]
Write a balanced net ionic equation for each reaction that occurred in step 4. Justify each reaction
written by relating to the observations. For each reaction, identify the element oxidized, the element
reduced, the oxidizing agent, and the reducing agent.
Don’t forget, colors & characteristics always go with equations. For Parts 3 & 4 in this lab, ID
what is reduced, what is oxidized, oxidizing agent, & reducing agent with the equations. In the summary,
justify the reasons for your decisions about the products in parts 3 & 4 as concisely as you can referring
back to Parts 1 & 2. Also discuss common oxidizing agents and reducing agents used in this experiment
and their products. Make sure you include the effect of the acidity of the solution.
Standardization of KMnO4
Background
An oxidation-reduction titration is a process used to determine the concentration of an ion in an unknown
solution by reacting it with another ion in a solution having a known concentration. The equivalence point
is reached when the total number of electrons lost in the oxidation reaction is equal to the total number of
electrons gained in the reduction reaction. In this experiment, a purple-colored solution of potassium
permanganate, with an approximate concentration of 0.025 M, will be added to a solution containing Fe2+
ions. The permanganate ions (MnO4-) are a strong oxidizing agent which causes the iron to be oxidized to
Fe3+ ions. The manganese is reduced from a 7+ oxidation state in the permanganate ion to form colorless
Mn2+ ions. The equivalence point is indicated at the point when all of the Fe2+ ions in solution are
oxidized and the colorless mixture retains a purple tint. (This color may be more orange in appearance,
depending upon the concentration of the Fe3+(aq), which has a yellow tint. A few drops of concentrated
phosphoric acid can be added to form a complex with the iron and minimize this color. Sulfuric acid is
added to increase the concentration of hydrogen ions in the solution.
Caution: Potassium permanganate and sulfuric acid can cause chemical burns. The KMnO4 will stain
skin and clothing. Avoid skin contact with these chemicals.
Pre-Lab: Write the balanced net ionic equation for the reaction occurring during the titration.
Procedure:
1.
2.
3.
4.
5.
6.
Rinse a clean a buret several times with a few mL of distilled water. Rinse the buret with a few mL of
the potassium permanganate and discard the washings into a waste beaker. Fill the buret and tip with
the KMnO4.
Use a clean Erlenmeyer for the titration. Place approximately 1 g of iron (II) ammonium sulfate
hexahydrate, Fe(NH4)2(SO4)2•6H2O (abbreviated FAS), in a weighing boat and mass with an
analytical balance to ±0.0001 g. Transfer the FAS to the Erlenmeyer flask. Add 25 mL of distilled
water, 15 mL of 3M H2SO4, and a few drops of concentrated H3PO4(aq) to the flask and swirl to
dissolve the FAS.
Place about 50 mL of water in a beaker and add 1 drop of the permanganate solution. This is the color
standard for the reaction. When the equivalence point is reached, the color intensity of the mixture
should match this standard.
Record the initial volume reading in the buret to ± 0.01 mL. Add the permanganate to the FAS
solution in the Erlenmeyer flask until the equivalence point is reached. Record the final volume
reading in the buret.
Repeat the titration process. Wash the mixtures in the flasks down the drain, then, rinse the flasks
with distilled water. The two trials should be within 1% of each other. If not, do another trial and
consult with the instructor.
Enter your data from each trial into the spreadsheet.
After analysis of the class results, use the 4d rule to determine if any of the values should be rejected.
Carry out all appropriate statistical calculations.
Determination of the % Oxalate in an Iron Oxalato Complex Salt
Introduction: In a previous experiment a green crystalline product having the formula
KwFex(C2O4)y.zH2O was prepared. The percentage oxalate in KwFex(C2O4)y.zH2O will now be determined
by titrating a solution containing a known mass of green salt with a standardized solution of KMnO4. The
mass and percentage of oxalate in the sample can be determined by measuring the volume of KMnO4, of
molarity M, required to completely oxidize the oxalate ion (C2O42-) in acid solution to give CO2 gas as a
product.
Safety, Environmental, and Economic Concerns:
The excess KMnO4 you have should be poured in the waste container under the fume hood. Pour the
solution resulting from the titration down the sink with plenty of tap water.
Notes on Experimental Procedures:
1. The oxidation state of iron in KwFex(C2O4)y.zH2O is +3, normally the highest value for iron. Thus
the KMnO4 does not oxidize the iron in this experiment. However, the presence of the ferric ion
imparts a yellow color to the solution. The H3PO4 forms a colorless phosphate complex with the
iron, making it easier to detect the color change which occurs at the equivalence point.
2. The reaction of the permanganate ion with the oxalate ion is rather slow at room temperature. The
purpose of heating the solution is to increase the rate of the reaction.
3. Prior to coming to the lab, write the balanced net ionic equation for the reaction.
Experimental Procedure:
1. Weigh two samples to the nearest 0.0001-gram, each of about 0.125 grams of the green salt crystal.
2. Transfer, completely, the green crystals to 250-ml Erlenmeyer flasks and make an aqueous solution.
Add about 6-ml 6 M H2SO4 and 1 ml of 85% (concentrated) H3PO4 to each sample.
3. Heat one of the two solutions to just below the boiling point.
4. While the first solution is being heated, prepare the buret with the standardized potassium
permanganate (KMnO4) solution.
5. Remove the Erlenmeyer flask from the heat source and titrate the green salt solution with the
KMnO4 solution. Titrate quickly so that the reaction occurs at a nearly constant temperature.
6. Repeat the titration procedure with the second solution (be sure to heat this solution as well).
7. The equivalence point is reached once the solution has turned a very faint pink color. Wash the
mixtures in the flasks down the drain, then, rinse the flasks with distilled water. The two trials
should be within 1% of each other. If not, do another trial and consult with the instructor.
8. Enter your data from each trial into the spreadsheet.
After analysis of the class results, use the 4d rule to determine if any of the values should be rejected.
Carry out all appropriate statistical calculations.
The Gas Law Constant
The Ideal Gas Law, PV = nRT, contains a constant, R, called the universal gas constant. In this
experiment, you will experimentally determine R and compare it with the theoretical value. You will also
study the stoichiometry of the reaction.
The reaction used is: Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
The gas will be collected in a eudiometer (gas collecting tube). Combining the volume of gas with the
room temperature, barometric pressure and moles produced from the stoichiometry, the value of R can be
calculated. You will need to report an analysis of both accuracy and precision, so repeated measures
should be considered.
The technique is simple. Prior to coming to lab, calculate the mass of magnesium needed to
generate more than 40 mL but less than 50 of hydrogen gas at STP (show work). The magnesium will be
provided in ribbon form. The precise mass of the ribbon can be measured on the analytical balance. Fold
the Mg ribbon into a small ball and tie off with string (5-6 cm). About 4 mL of concentrated HCl are
placed in a 50-ml eudiometer. Remember that concentrated HCl can be dangerous to skin and
clothing. Layer distilled water on the acid so there is minimum mixing of the acid and water. Fill the
eudiometer to overflowing. Immerse the Mg in the water in the buret. Insert a 1-hole 00 rubber stopper to
capture the string and invert the buret. Place the end in water in a large beaker and clamp to a ring stand.
Allow the gas to generate until the reaction stops. The reaction is slightly exothermic, so let the gas sit for
several minutes to come to thermal equilibrium with the room. Equalize the pressure of the gas inside the
tube with the atmospheric pressure before reading the volume. The experimental technique will be
discussed in lab, so do not worry if you are confused. Record on the computer: the dry pressure of the gas,
the volume of gas, the temperature at which the gas was collected, and the mass of Mg used.
After analysis of the class results, use the 4d rule to determine if any of the values should be rejected.
Carry out all appropriate statistical calculations.
Report the average value of R, in
L  atm
mol  K
individual & class. Sources of error? Effects?
, and an analysis of precision and accuracy, both
The Synthesis of Aspirin
Aspirin, the ubiquitous pain reliever, goes by the chemical name acetylsalicylic acid. One of the
compounds used in the synthesis of aspirin is salicylic acid, which is itself a pain reliever that was known
to many ancient cultures, including the Native Americans who extracted it from willow tree bark. Salicylic
acid is extremely bitter tasting, and frequent use can cause severe stomach irritation. The search for a
milder form of this pain reliever led to the successful synthesis of acetylsalicylic acid by the German
chemist Felix Hoffmann in 1893.
There is more than one way to synthesize aspirin; in this experiment, you will react acetic
anhydride with salicylic acid in the presence of phosphoric acid (which acts as a catalyst). The reaction is
shown below. In this experiment, you will synthesize a sample of acetylsalicylic acid (aspirin) and
calculate the percent yield of your synthesis.
PROCEDURE
1. This reaction should be conducted in a fume hood or a well-ventilated area of the room.
2. Measure 2-3 grams of salicylic acid into a small Erlenmeyer flask using a pan balance. Record your
precise mass. Measure about 5 mL of acetic anhydride.
3. Add the acetic anhydride and 5 drops of 85% phosphoric acid to the salicylic acid. [Make sure you
take appropriate data to determine the mass of the acetic anhydride.] Swirl the mixture. CAUTION:
Handle the phosphoric acid and acetic anhydride with care. Both substances can cause painful burns
if they come in contact with the skin.
4. Heat the mixture in a boiling water bath for 15 minutes or when the mixture ceases releasing vapors.
Stir the mixture occasionally during heating. After about 10 minutes, add 2 mL of distilled water to the
flask. Set up a Büchner funnel and filter flask so that you are ready to filter the reaction mixture after it
has cooled.
5. When you are confident that the reaction has reached completion (no vapors appearing), carefully
remove the flask from the hot plate and add 20 mL of distilled water. Allow the mixture to cool to near
room temperature. Transfer the flask to an ice bath for about five minutes. As the mixture cools,
crystals of aspirin should form in the flask. Add about 10 mL of cold distilled water.
6. Determine the mass of the filter paper to use in the Büchner funnel. ID, with a pencil, your filter paper
and place it in the Büchner funnel. Turn on the aspirator. Wet the filter paper in the Büchner funnel
with distilled water. Transfer the contents of the cooled flask to the Büchner funnel. When most of the
liquid has been drawn through the funnel, turn off the suction and wash the crystals with 5 mL of cold,
distilled water. After about 15 seconds, turn the suction back on. Wash the crystals with cold, distilled
water twice more in this manner.
7. Remove the filter paper with the aspirin crystals from the Büchner funnel and set aside the aspirin in a safe place to dry.
8. Determine the mass of the filter paper and aspirin.
9. Store the remaining crude aspirin for later experimentation in a vial which you have labeled.
Determine the limiting reagent and calculate the theoretical yield. Assume 100% purity calculate the percent yield of your reaction. Turn this
in for 10 lab points.
Intermolecular Attractions
The polarity and size of molecules determines the extent to which the molecules are attracted to
each other. These intermolecular attractions determine many of the chemical and physical properties of
molecules. You are to examine two perspectives of this phenomenon, chromatography and evaporation.
Part A: Chromatography
The word chromatography means color-writing. The name was chosen at the beginning of this
century when the method was first used to separate colored components from plant leaves.
Chromatography in its various forms is one of the most important methods of chemical analysis of
mixtures.
A single spot of the unknown to be analyzed is applied about half an inch from the end of the thin
layer or paper chromatogram. The chromatogram is then placed vertically in a shallow layer of solvent
mixture in a jar or beaker. Since the chromatogram absorbs liquids, the solvent begins rising by capillary
action. As the solvent rises to the level at which the spot of mixture was applied, various effects can occur,
depending on the composition of the spot. The components of the spot that are completely soluble in the
solvent will be swept along with the solvent front as it rises. Those components that are not at all soluble
will be left behind at the original location of the spot. Most components of the unknown spot mixture will
take an intermediate approach as the solvent passes. Components in the spot that are somewhat soluble
will be swept along by the solvent front, but to different degrees, reflecting their specific solubilities. By
this means, the original spot of mixture is spread out into a series of spots or bands, each representing one
component of the original mixture. The separation of a mixture by chromatography is not solely a
function of the solubility of the components in the solvent. The chromatogram itself consists of molecules
that may interact with the molecules of the mixture. Each component is likely to have a specific
interaction.
To place chromatographic separation on a quantitative basis, a mathematical function called the
retention factor, Rf, is defined: Rf = distance traveled by spot/distance traveled by solvent. The retention
factor depends on what solvent is used for the separation and on the specific composition of the
chromatogram used. Because the retention factors depend on conditions of the analysis, a known sample
of each possible component of a mixture is run on the same chromatogram as the mixture. Identification
of the components of the mixture must be verified by retention factor values, not only visual comparison.
Thin-layer chromatography (abbreviated TLC) uses a thin coating of aluminum oxide (alumina) or
silica gel on a glass slide or plastic sheet to which the mixture to be resolved is applied. In paper
chromatography, filter paper is commonly used as the support medium. The chromatogram can be run
either vertically on a strip or horizontally in a circle. The vertical method is recommended here because of
the similarity to thin layer in calculating the Rf values. Liquid chromatography (LC) is one type of
chromatography that is very useful in research and in industry. High performance liquid chromatography
(HPLC) has become an almost indispensable tool for scientists. Liquid chromatography is similar to paper
or thin lay chromatography in that there is a stationary support medium which attracts the components of a
mixture as a liquid phase passes. In liquid chromatography, the stationary support is a column packed with
a fine, granular solid. The mixture to be separated is placed in the column and clings to the solid. A
solvent is added which washes the mixture through the column. The substances that are more soluble in
the solvent travel more quickly through the column, and emerge early. Those substances that are more
strongly attracted to the stationary support move slowly, and emerge later.
Procedure for Paper Chromatography [work individually]
Cut a piece of filter paper into a rectangle so that it can be used in a clean 250-mL beaker which
will serve in developing the chromatogram. Cut a square of plastic wrap for use as a cover for the beaker.
Select four (4) felt-tip pen with water soluble ink. Draw a light pencil line across the paper about 1.5 cm
from one end and lightly mark four small circles spread across the line. On each circle apply a single
small spot from a different pen. Allow the spot to dry completely. Record the original color of each ink.
Place a stirring rod across the top of the beaker, through the lip of the beaker. Fold the end of the
chromatogram, opposite the spots, so that it can hang on the rod without touching the bottom of the beaker.
Add distilled water to the beaker to a depth of about 1 cm. Carefully lower the chromatogram into the
water in the beaker. Make certain not to wet the spots and do not move the beaker to avoid sloshing water
onto the spots. Cover the beaker with the plastic wrap.
Allow the water to rise on the chromatogram until it nears the rod. Then carefully remove the
chromatogram and set it on a clean paper towel. Immediately mark the solvent front with a pencil – the
solvent front will continue to move slightly after removal.. After the chromatogram has dried, make
appropriate measurements and determine the Rf for each color of ink in each color of pen. Discuss the
similarities and differences of the inks in the different pens and the reliability of the Rf values for
identification purposes. Include your chromatogram as part of the lab report.
Part B: Evaporation [work in pairs]
In order for a substance to change from the liquid to the gaseous state, the molecules must increase
their own kinetic energy. The source of the energy for the change is the environment of the molecules.
Therefore, the temperature of the surroundings goes down as energy flows into the liquid.
In this experiment, temperature probes are placed in various liquids. Evaporation occurs when the
end of the probe is removed from the liquid. Evaporation results in a decrease in temperature of the liquid
remaining on the probe. The rate of evaporation is an indicator of the force of attraction between
molecules of the liquid and can be used to compare the relative intermolecular attractions for the liquids.
Prior to coming to class, draw the molecular structure of each liquid used in the experiment.
Briefly discuss the type of intermolecular attractions of each liquid.
The liquids you will test are:
1. n-pentane
2. n-heptane
3. acetone
4. methanol
5. 1-propanol
6. 1-pentanol (amyl alcohol)
7. ethylene glycol
A temperature probe will be will be used with a computer for quick and precise temperature
measurements. The length of time should be 3:00 minutes. Place the probe in a liquid sample. [Hold the
vial by the neck. Do not hold onto the sample.] Allow the probe to come to thermal equilibrium with the
liquid before starting the collection of data. After the computer begins to collect data, allow about 5
seconds to establish the initial temperature of the liquid. With as little motion as possible, lay the probe on
the counter top, with the tip extending about 2 inches over the edge of the desk. The tip must not touch the
counter or be left to rest on the counter. Monitor the temperature for the remainder of the time or until the
temperature starts to rise. Save or export the data to your desktop. All liquids must be tested. Repeat
trials may be run. All samples may be saved in one file using the “Save & Continue” feature. How are
you going to identify each trial (data column)?
Report:
Prepare four graphs, each containing the cooling curve of two or more liquids. Graph 1 shows the
evaporation of liquids 1 & 2. Graph 2 shows the evaporation of liquids 1 & 6. Graph 3 shows the
evaporation of liquids 4-7. Graph 4 shows the evaporation of liquids 3 & 5. Look up and report the boiling
point of each liquid. Using the “Linear Fit” function of Logger Pro (see “Helpful Hints for Graphing in
LoggerPro”), determine an initial rate of evaporation of each liquid in each graph. (This would be the
slope of the of the fast rate of temperature decline.) Compare and contrast the relative rates of evaporation
of the liquids within each group specified. Discuss any trends of the boiling points and the initial rate of
evaporation based upon the type of intermolecular attractions.
Determination of a Heat of Vaporization
The Clausius-Clapeyron equation is used by chemists to describe the change in vapor pressure of a
 H vap 1
liquid as temperature changes. Stated mathematically: ln P =  C
R
T
In this experiment, you will saturate an air sample with water vapor by trapping a sample of air in
an inverted graduated cylinder and heating it in a water bath. As the gas cools, the amount of water vapor
in the air decreases while the molar quantity of gas remains constant. By calculating the moles of air in the
cylinder, the partial pressure of air at each temperature can be calculated. The vapor pressure of the water
in the cylinder can be determined by reference to the room barometric pressure and some basic math. You
can determine the moles of air at a temperature near 0C where the water vapor in air is less than one
percent. Although this amount is negligible, it will affect the accuracy of your calculations. You must
manipulate the data you gather and determine the enthalpy of vaporization for water in the fashion as
described by Clausius and Clapeyron. [NOTE: There is a methodological error that you cannot avoid. It is
worth consideration and discussion.]
Procedure
Fill a 10-mL graduated cylinder with distilled water such
that when inverted you have trapped about 3 mL of air. Fill a 1L beaker nearly full with tap water. Covering the top of the
graduated cylinder with your finger, quickly invert it and lower
it into the beaker of water. For best results, you want to trap
about 3 mL of air in the cylinder. Make sure the trapped air in
the cylinder is completely covered by the water in the beaker.
Using a Bunsen burner, heat the beaker of water, carefully
observing the volume of trapped air. When the air level expands
just beyond the scale on the cylinder, remove the beaker from
the burner and place it in a plastic tray on the magnetic stirrer.
Using a stir bar, stir the water to maintain an even temperature
throughout the system. Place the thermometer alongside the
graduated cylinder for the best data. When the air bubble
contracts to a volume that can be read, start collecting data.
Collect sufficient data to insure reliable results. Experience has
shown that data collected below 50C is not useful. After the
temperature drops below 50C, cool the water to near 0C by
adding large amounts of ice. Record the air volume at the
lowest stable temperature.
Calculations
While it might appear that the height difference in the water levels might contribute a significant
error, it does not. Why? We will simply assume that: Patm = Pwet air = Pdry air + Pwater(g). The quantity of dry
air [which is a constant] can be best approximated by calculations with the volume of trapped air at the
lowest temperature achieved, using the assumption that the quantity of water vapor at this low temperature
is negligible. Once the quantity of dry air is determined, the pressure the dry air exerts at each temperature
can be determined. Thus, the pressure the water vapor exerts at each temperature can then be calculated.
[Make sure you show a sample of each kind of calculation you do.]
Plot an appropriate graph, using the Clausius-Clapeyron equation as a basis, and determine the
enthalpy of vaporization of water. In general, how does your experimental value compare with the
theoretical value? Sources of error? Effects of errors?
Molar Mass from Freezing Point Depression
The addition of a solute to a solvent, in general, lowers the freezing point of the solvent. For a
given solvent, the freezing point lowering is directly proportional to the concentration of particles
dissolved in it. For naphthalene, the solvent used in this experiment, the freezing point is lowered by 6.9oC
for each mole of solute particles in 1 kg of naphthalene. In this experiment, the molar mass of elemental
sulfur will be determined by observing the freezing points of pure naphthalene and a solution that contains
known masses of sulfur and naphthalene.
Hints in Designing the Procedure:
 Errors are controlled in this experiment by measuring the freezing point on the same sample of
pure solvent that will be used in making the solution.
 A 25 x 150 mm test tube ½ full of naphthalene is a convenient amount of solvent with about 1
gram of powdered sulfur to make the solution.
 A temperature probe will be used to measure temperature. [After you have all the interfacing
connected, call up LoggerPro. Select “Experiments” in the drop down menu. Select “Data
Collection”. In the “Collection” tab, change the Length of time to 300 seconds and the Sampling
Rate to 2 seconds/sample (the same as 0.5 samples/second). When you are ready, Collect. When
you feel confident you have sufficient data to make your analysis, Stop. You probably won’t need
the full 5 minutes.]
 The solid(s) will be melted with the test tube immersed in a boiling water bath.
 The freezing points will be determined with the test tube suspended in the air. The cooling is more
rapid than what would be optimal, but with continual stirring with the temperature probe, error can
be reduced. One trial with the pure naphthalene is generally sufficient.
 Do not try to remove the probe) after it becomes frozen in the solid.
 A clear yellow homogeneous solution should be made after the sulfur is mixed with the
naphthalene. Do one melting-cooling w/stirring without collecting any data to help achieve
homogeneity. Collect data on two trials.
 For disposal, melt the solution and pour it quickly onto a couple of paper towels in the waste
basket. Return the dirty test tube to the teacher for cleaning with a special solvent.
 Your report should contain the cooling curves and analysis for both the solvent and mixture on the
same graph [make it look good!]. An example of this analysis can be found in LoggerPro [open –
Experiments – Sample Data – Chemistry – freezing point depression]. Calculations of the molar
mass and molecular formula of sulfur and all appropriate error analyses should also be included.
Rate Law Crystal Violet
An understanding of chemical reactions must include answers to certain basic questions:
1.
2.
3.
Will two materials react (transform into another substance) when placed in contact with
each other?
If they do react, what determines how quickly the reactants transform into the products?
Why do some reactions fail to go to completion, and stop when an equilibrium mixture of
reactants and products is formed?
The first and third questions are examined in the area of chemistry called thermodynamics; the
second question - the concern of this experiment - is probed in the area known as reaction kinetics.
We look closely at the question of what determines how quickly reactants change into products.
Several factors, discussed in lecture and the text, determine the speed of a chemical reaction. These
include 1) nature of reactants, 2) contact area between reactants, 3) temperature, and 4) concentration of
reactants. The description of the kinetics of a chemical reaction is made easier by a careful definition of
the term "reaction rate." Obviously, the amounts of reactants decrease while the amounts of products
increase during the course of the reaction. The speed or rate of the reaction is given by the change in
amount of reactant or product per unit time. This rate can be defined either through the disappearance rate
of the reactants or through the appearance rate of the products.
The rate at the beginning of the reaction is known as the initial rate; for most reactions the rate is
largest at this time. The reactions generally slow down as the reactants are depleted and as their collision
probabilities decrease. It is not possible to predict the rate of a reaction from the balanced overall reaction.
Detailed information about the reaction process called the pathway or mechanism is needed for this task.
But extensive investigations in the laboratory have shown that the rates of many reactions obey a common
empirical equation known as the experimental rate law: rate = k(T) · [A]a[B]b[C]c...where [A], [B], [C] ...
represent the molar concentrations of substances affecting the reaction rate (usually reactants) and a, b, c...
are experimentally determined exponents. These exponents may be integers or non-integers (positive or
negative), and have no direct connection with the stoichiometric coefficients of the reaction. The constant,
which can be written as k(T), is called the rate constant, and its numerical value generally increases with
temperature. A larger value for the rate constant indicates a faster reaction because the rate is proportional
to k. The overall reaction order is given by the sum of the exponents (a + b + c ...). The reaction is also
said to be a-order in substance A, b-order in substance B, and so on. These exponents are often called
partial orders. We can derive the partial orders in the laboratory, but the trick is to change only one thing
at a time. We can avoid difficult concentration measurements by measuring the initial rate. The
concentrations at the time of the initial rate are simply the initial concentrations. This is the initial rate
method.
[The following is an adaptation of “Rate Law Determination of the Crystal Violet Reaction” from
Chemistry with Computers, (2000), Holmquist, Randall, & Volz; Vernier Software & Technology,
Beaverton, OR.]
In this experiment, you will observe the reaction between crystal violet and sodium hydroxide. The
equation for the reaction is as follows:
N(CH3)2
N(CH3)2
OH
+
C
N(CH3)2
+ OH-
N(CH3)2
C
N(CH3)2
N(CH3)2
A simplified version of the equation is:
CV+
+
OH-  CVOH
(crystal violet) (hydroxide ion)
The rate law for this reaction is in the form: rate = k[CV+]a[OH-]b, where k is the rate constant for
the reaction, a is the order with respect to crystal violet (CV+), and b is the order with respect to the
hydroxide ion. [Assume that the order of the reaction with respect to the [OH-] is zero.] This experiment
examines the quantitative dependence of the speed of a particular chemical reaction upon changes in
temperature and changes in reactant concentrations. Once the order of the reaction with respect to crystal
violet has been determined, you will also find the rate constant, k, and the half-life for this reaction. In
addition, you will determine the activation energy of the reaction.
As the reaction proceeds, a violet-colored reactant will be slowly changing to a colorless product.
Using the green (565nm) light source of a Vernier Colorimeter, you will monitor the absorbance of the
crystal violet solution with time. We will assume that absorbance is proportional to the concentration of
crystal violet (Beer’s Law). Absorbance will be used in place of concentration in plotting the appropriate
graphs.
Procedure:
In Part I you will observe the rate of reaction between given concentrations of crystal violet and
sodium hydroxide. In Part II you will hold the crystal violet concentration constant but vary the
temperature.
You will have available to use:
75 mL of 2.5 x 10-5 M crystal violet
75 mL 0.10 M sodium hydroxide
a Vernier Colorimeter w/cuvetts
a Vernier interface
Connect the computer, LabPro and Colorimeter as instructed
After opening the LabPro file:
Under the pull-down Experiment, select
Data Collection
Mode
Time Based
Length
240 seconds
2 seconds/sample
Set the Colorimeter at 565 nm by pressing the < or > button to select the correct
wavelength
Fill a cuvette about ¾ with distilled water for the blank (100% transmittance or 0
absorbance). Wipe with sides of the cuvette with a soft cloth. Insert the
cuvette into the Colorimeter. Important: Line up one of the clear sides of
the cuvette with the arrow at the top of the cuvette slot. Close the
Colorimeter lid. Next, press the CAL button to begin the calibration process.
Release the CAL button when the red LED begins to flash. When the LED
stops flashing, the calibration is complete and the unit is ready to collect
data.
Part I: Crystal Violet Concentration Dependence
Using two different 10-mL graduated cylinders fill one with 10.0 mL of 2.5 x 10-5 crystal violet
and the other with 0.10 M sodium hydroxide. Simultaneously pour the crystal violet and sodium
hydroxide solutions into a small beaker. Carefully stir the reaction mixture with a thermometer and
record the temperature as precisely as possible [this trial will also provide data for Part III].
Quickly rinse a clean cuvette with ~1 mL of the reaction mixture, dispose of the rinse, and then fill
the cuvette about ¾ full. Place the cuvette properly into the Colorimeter. COLLECT the data,
observing the solution in the beaker as it continues to react. After the data collection is complete,
save the data to the computer Desk Top.
Part II: Temperature Dependence
Repeat Part I at one lower temperature and three higher temperatures, but no higher than 450C.
The experimental temperatures should be in a reasonable “spread.” Using a large beaker to
immerse the graduated cylinders containing the solutions seems to be an appropriate way to heat &
cool the solutions. Complete the mixing and transfer to the cuvettes as quickly as you can. [While
temperatures will change somewhat during the experiment, this still fairly represents the concept.]
Use clean cuvettes.
Make sure you remove the cuvette from the colorimeter. Clean all glassware and cuvettes which contained
crystal violet with HCl(aq) and rinse with distilled water. Place the cuvettes in the box with the
colorimeter and return all materials to the appropriate location.
CALCULATIONS:
Part I: Dependence of Concentration of Crystal Violet
In this entire experiment, we can observe that the changes in the rate of reaction are due solely to
variations in the crystal violet concentration. Since we are assuming that the reaction is independent of the
hydroxide ion, the experimental rate law can be rewritten as follows: rate = k*[CV+]a where k* is a pseudo
rate constant. The rate then should simply be proportional to some power (equal to a) of [CV+]. Graphing
the data would assist us in determining the partial order of this reaction with respect to crystal violet.
However, you recorded the reaction time rather than the reaction rate. As presented in lecture, one
can determine the order of a reaction by graphing concentration and time in various forms to determine
which one gives the most linear graph. In this experiment, we could not directly measure the change in
concentration of crystal violet. Using the Colorimeter, we observed changes in the absorbance of the
solution. The Beer-Lambert Law, A = lc, [where A = absorbance & c = concentration] shows that
absorbance is directly proportional to concentration. [Prior to coming to the lab, search for Beer-Lambert
Law in Wikipedia. Familiarize yourself with the information under Equations (through the development
of the equation) and Prerequisites.] Thus, we will use absorbance, rather than concentration, in our graphs.
Thus: (1) Prepare three separate graphs and determine the partial order of this reaction with respect to
[CV+]; (2) Upon determination of the partial order of this reaction with respect to [CV+], determine the
pseudo rate constant, k*; (3) Determine the half-life of the reaction.
Part II: Temperature Dependence
Arrhenius discovered that the reaction rates often change with temperature according to the

Ea
RT
equation: k  Ae
. Note that the rate constant k is related to the reaction temperature. Using the
appropriate plot from Part 1, construct one graph with all four Abs-time data and find the rate constant at
each temperature. After modifying the Arrhenius equation to a slope-intercept form, use these results to
determine the activation energy of the reaction.
Determination of a Rate Law and Activation Energy
An understanding of chemical reactions must include answers to certain basic questions:
1.
Will two materials react (transform into another substance) when placed in contact with
each other?
2.
If they do react, what determines how quickly the reactants transform into the products?
3.
Why do some reactions fail to go to completion, and stop when an equilibrium mixture of
reactants and products is formed?
4.
The first and third questions are examined in the area of chemistry called thermodynamics; the
second question - the concern of this experiment - is probed in the area known as reaction kinetics.
We look closely at the question of what determines how quickly reactants change into products.
Several factors, discussed in lecture and the text, determine the speed of a chemical reaction. These
include 1) nature of reactants, 2) contact area between reactants, 3) temperature, and 4) concentration of
reactants. The description of the kinetics of a chemical reaction is made easier by a careful definition of
the term "reaction rate." Obviously, the amounts of reactants decrease while the amounts of products
increase during the course of the reaction. The speed or rate of the reaction is given by the change in
amount of reactant or product per unit time. This rate can be defined either through the disappearance rate
of the reactants or through the appearance rate of the products.
The rate at the beginning of the reaction is known as the initial rate; for most reactions the rate is
largest at this time. The reactions generally slow down as the reactants are depleted and as their
collision probabilities decrease. It is not possible to predict the rate of a reaction from the balanced overall
reaction. Detailed information about the reaction process called the pathway or mechanism is needed for
this task. But extensive investigations in the laboratory have shown that the rates of many reactions obey a
common empirical equation known as the experimental rate law: rate = k(T).[A]a[B]b[C]c , where [A], [B],
[C] ... represent the molar concentrations of substances affecting the reaction rate (usually reactants) and a,
b, c are experimentally determined exponents. These exponents may be integers or non-integers (positive
or negative), and have no direct connection with the stoichiometric coefficients of the reaction. The
constant, which can be written as k(T), is called the rate constant, and its numerical value generally
increases with temperature. A larger value for rate constant indicates a faster reaction because the rate is
proportional to k. The overall reaction order is given by the sum of the exponents (a + b + c ...). The
reaction is also said to be a-order in substance A, b-order in substance B, and so on. These exponents
are often called partial orders. We can derive the partial orders in the laboratory, but the trick is to
change only one thing at a time. We can avoid difficult concentration measurements by measuring the
initial rate. The concentrations at the time of the initial rate are simply the initial concentrations. This is the
initial rate method.
In this experiment, you will measure the initial rate of the reaction between thiosulfate ions
and hydronium ions (supplied by hydrochloric acid) according to
S2O32-(aq) + 2 H3O+(aq)  S(s) + SO2(g) + 3 H2O(l)
The sulfur product is a solid that becomes clearly visible as the reaction proceeds. The experimental rate
law has the form
rate = k[S2O32-]m[H3O+]n
29
and you will find the values of m and n. This experiment examines the quantitative dependence of the
speed of a particular chemical reaction upon changes in temperature and changes in reactant
concentrations.
Procedure:
In Part I you will record the time required to produce a fixed amount of sulfur at various initial
concentrations of the thiosulfate ion. The hydronium ion concentration and the temperature do not
change. In Part II you hold the thiosulfate ion concentration constant but vary the hydronium ion
concentration. Again, the times required to produce a fixed amount of sulfur are recorded and the
temperature is held constant. In Part III you hold both the thiosulfate and hydronium ion concentrations
constant but vary the temperature. The times required to produce a fixed amount of sulfur is again
recorded. In each case you look through the reaction mixture at a black cross on a white background.
You will see the following changes as the solid sulfur forms and its particle size increases: clear and
colorless  milky and light blue (skim milk)  milky and white (milk)  opaque yellow. Somewhere
along the way the mixture will become sufficiently opaque for the black cross to disappear from view.
The set amount of sulfur has then been produced and the time is stopped.
You will have available to use:
175 mL 0.25M sodium thiosulfate
125 mL 2.0M hydrochloric acid
50-mL beaker for the reaction
other equipment as needed
Add the HC1 solution to the beaker containing the thiosulfate solution. Stir, and begin timing
immediately. Stop stirring after a few seconds and record the time as precisely as possible when the cross
disappears from view. You will find it helpful to dry the outside of the beaker occasionally. If the
colloidal sulfur sticks to the beaker after a few trials, simply wipe the film away with a towel. (HINT: you
might find it useful to keep the total volume of solution constant within each Part.)
PART I: Thiosulfate Ion Concentration Dependence
Do 5 trials of varying [Na2S2O3]. The most concentrated trial should not use more than 25 mL or
less than 5 mL of the 0.25 M sodium thiosulfate. Use 5.0 mL of the 2 M HC1 for each trial.
Part II: Hydronium Ion Concentration Dependence
Do 5 trials of varying [HC1]. The most concentrated trial should not use more than 20 mL or less
than 10 mL of the 2 M hydrochloric acid. Use 25 mL of 0.1 M Na2S2O3 for each trial.
PART III: Temperature Dependence
For each trial, use 25 mL of 0.04 M sodium thiosulfate and 5 mL of 2 M hydrochloric acid. Do
trials from approximately 20° - 60°C in 10° increments. (Make sure the thermometer bulb does
not touch the side or bottom of the reaction beaker.) Record the temperature as precisely as
possible at the point of mixing and the point of ending. Use the average temperature for your
analysi
CALCULATIONS:
Parts I and II: Concentration Dependence
The hydronium ion concentrations are held constant during Part I, and the changes in the
reaction times are due solely to variations in the thiosulfate ion concentrations. In this case the
experimental rate law can be rewritten as follows:
Rate (wrt [S2O32-]) = k[S2O32-]m[H3O+]n = k*[S2O32-]m since the [H3O+] is constant
where k* is a “pseudo” constant. The rate then should simply be proportional to some power (equal to
m) of [S2O32-]. Some graph which includes the rate and the [S2O32-]would assist us in determining the
partial order of this reaction with respect to thiosulfate ion.
However, you recorded the reaction time rather than the reaction rate. As presented in
lecture, one can determine the order of a reaction by graphing concentration and time in various
forms to determine which one gives the most linear graph. Since we are dealing with initial
concentrations, the equations must be rearranged. For a zero-order reaction, the plot is of the form
Ao = kt + A; for first-order, ln Ao = kt - ln A, and for second order:
1
Ao
= −kt +
1
A
. Thus, using
appropriate data:
1. prepare three separate graphs and determine the partial order of this reaction with respect
to [S2O32-].
2. make similar graphs and determine the partial order of this reaction with respect to
+
[H3O ].
3. combine the results to obtain the overall reaction order.
4. write the rate law for this chemical system.
Part III: Temperature Dependence
In this part, you kept the concentration of the chemicals constant and the time varied as the
temperature changed. How are reaction time and reaction rate related? The rate is always defined
∆(𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛)
as 𝑟𝑎𝑡𝑒 =
. Rate is also generically defined as rate = k[Ao]. Using the
∆(𝑡𝑖𝑚𝑒)
concentrations of reactants, you have measured the time required to produce a fixed amount of sulfur,
which is the amount of sulfur required to make the black cross disappear from view. Accordingly,
(concentration) in the preceding equation is constant and the rate is directly proportional to the
1
reciprocal of the reaction time: 𝑟𝑎𝑡𝑒 = ∆(𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛) ∙ 𝑡 = 𝑘[𝐴𝑜 (𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡)]. Note that in both
equations, the concentration is a constant, then k is directly proportional to the reciprocal of the
reaction time.
Arrhenius discovered that the reaction rates often change with temperature according to the
E
− a
equation: k = Ae
RT .
Note that in the Arrhenius equation the rate constant k is related to the
reaction temperature. Thus,
1
t
Ea
= Ae−RT . Modify this variation of the Arrhenius equation to a
slope-intercept form. Using your data, prepare an appropriate graph and determine the
activation energy of the reaction.
In a previous experiment, you prepared aspirin by combining acetic acid and salicylic acid.
Your synthesis converted most, but not all, of the salicylic acid into acetylsalicylic acid (aspirin).
While there are ways in which the purity of your acetylsalicylic acid can be determined directly, you
will use an indirect method. As can be seen by the chemical system presented, when iron (III) nitrate
is mixed with salicylic acid, a bluish-purple color complex ion is formed. You will analyze a sample
of your crude aspirin to determine the amount of salicylic acid in it. You can use this information to
calculate the purity of your aspirin sample if you assume that the unreacted salicylic acid is the only
impurity,.
+1
H2O
H
H
OH2
Fe
O
O
OH2
H2O
+ Fe3+(aq)
O
O
C
C
O
O
Visible Spectrophotometry is a method of measuring the concentration of colored solutions by
the amount of light absorbed by or transmitted through a colored solution. The absorbance of white
light by a solution containing a colored compound is directly proportional to the concentration of the
colored compound. The constant of proportionality contains the path length of the sample through
which the light passes and a constant that is determined by the color of the solution. This information
produces Beer's Law: A = bc, where A is absorbance,  is the molar absorptivity of the colored
solution, b is the inside diameter of the cell, and c is the molar concentration of the solution. In this
 nd b. Rather we will us a Beer’s
Law plot as a calibration curve or a conversion curve. From the calibration curve we will be able to
determine the concentration of an unknown solution. For a discussion of the spectroscopic theory of
this experiment, search “Beer’s Law” in Wikipedia. You will use a Vernier Visible Spectrophotometer
in this experiment
Procedure
1.
Prior to coming to the lab, calculate how much 0.25 M Fe(NO3)3 is needed to make
approximately 700 mL of a 0.025 M Fe(NO3)3 solution.
2.
You must prepare a standard salicylic acid solution of known concentration. Measure on the
analytical balance approximately 0.08 g of pure salicylic acid and record the precise mass.
Transfer the salicylic acid to a 100-mL volumetric flask and add about 10 mL of 95% ethanol.
Swirl the flask to dissolve the acid. Add some distilled water to the flask and mix. Add more
distilled water to fill the flask to the 100.00 mL mark. Mix the solution thoroughly. Transfer
this stock standard salicylic acid solution a clean, dry beaker Calculate the precise
concentration of your stock solution in mg/mL.
3.
Prepare approximately 700 mL of a 0.025 M Fe(NO3)3 solution from the stock Fe(NO3)3
solution which is 0.25 M.
4.
Use the stock salicylic acid solution to prepare five diluted samples. Pipet 2.00 mL, 4.00 mL,
6.00 mL, 8.00 mL, and 10.00 mL of stock solution into five labeled (1-5) 100-mL volumetric
flasks. To each flask add 0.025 M Fe(NO3)3 to make precisely 100.00 mL. Determine the
concentration (mg/mL) of salicylic acid in each flask before proceeding.
5.
Prepare your crude aspirin sample. Measure out about 0.08 gram of your crude aspirin and
record the precise mass that you use. Transfer the crude aspirin to a 50-mL volumetric flask
and add about 10 mL of 95% ethanol. Swirl the flask to dissolve the acid. Add distilled water
to the flask with mixing. Fill the flask to the 50.00 mL mark with distilled water. Mix the
solution thoroughly. Transfer the stock crude aspirin solution to a clean, dry beaker.
6.
Prepare two (2) samples for your crude aspirin for analysis. Transfer 10.00 mL of the crude
aspirin solution to a clean 50-mL volumetric flask and also to a clean 100-mL volumetric flask.
Add 0.025 M Fe(NO3)3 solution to each flask precisely to the line. Mix each solution
thoroughly.
The use of the visible spectrophotometer is in three parts:
Part 1: Determination of the optimum wavelength
1. Using a USB cable, connect a Vernier Spectrometer to a computer.
2. Start the Logger Pro 3.x program on your computer.
3. The spectrometer should be set up at this point. If not, open the Experiment menu and select
connect Interface . Spectrometer . Scan for Spectrometers.
4. Calibrate the spectrometer.
a. Prepare a blank. Rinse a cuvette with the 0.025 M Fe(NO3)3 solution and fill it ¾ full with this
solution.
b. Open the Experiment menu and select Calibrate . (Spectrometer). The following message appears
in the Calibrate dialog box: “Waiting … seconds for the device to warm up.” After 60 seconds or
so, the message changes to: “Warmup complete.”
c. Place the blank in the cuvette holder of the spectrometer. Align the cuvette so that the clear sides
are facing the light source of the spectrometer. Click “Finish Calibration”, and then click OK.
5. Determine the maximum wavelength for your standard salicylic acid solution and set up the data
collection mode.
a. Select five (5) more curvettes. Rinse each, in turn, with one of the labeled (1-5) standard salicylic
acid solutions. Fill each cuvette ¾ full with the corresponding standard solution.
b. Wipe cuvette #3 with a tissue and place it in the cuvette holder of the spectrometer.
c. Click the Collect button. A full spectrum graph of the standard salicylic acid solution will be
displayed. Note that one area of the graph contains a peak absorbance. Click the Stop button to
complete the analysis. Examine the peak absorbance.
d. To save your graph of absorbance vs. wavelength, select Store Latest Run from the Experiment
menu.
e. Remove cuvette #3 from the spectrometer.
Part 2: Determination of the Beer’s Law plot
6. Click the Configure Spectrometer Data Collection icon, on the toolbar. A dialog box will appear.
a. Select Absorbance vs. Concentration under Set Collection Mode. The peak absorbance should be
automatically selected. If not, check the appropriate box.
b. Change the Units from mol/L to mg/mL.
7. Collect absorbance-concentration data for the five standard solutions.
a. Wipe cuvette #1 with a tissue and place it in the spectrometer cuvette holder. Click the Collect
button. When the absorbance reading stabilizes, click the Keep button. Enter your calculated
concentration of solution #1 and click OK.
b. Repeat Step 7b for the remaining solutions.
c. When you have finished testing the standard solutions, click the Stop button.
8. To determine the best-fit line equation for the standard salicylic acid solutions, click the linear fit
button on the toolbar. Write down the equation for the standard solutions in your lab book.
Part 3: Determination of the Concentration of the Salicylic Acid in the Crude Aspirin
9. Determine the concentration of the salicylic acid in your crude aspirin.
a. Rinse a cuvette twice with a crude aspirin solution and fill it about ¾ full. Wipe the outside of the
cuvette and place it into the spectrometer.
b. Select Interpolation Calculator, from the Analyze menu. A dialog box will appear that displays
the concentration of your unknown at the measured absorbance.
c. Click OK. Write down the concentration of the unknown in your lab book.
d. Remove the curvette from the spectrometer.
e. Repeat 9a-d with the other crude aspirin solution.
f. Dispose of all solution down the drain. Rinse all volumetric flasks and curvettes twice with
distilled water and put away.
Data Analysis
1. Print the graph of the visible spectrum of the Fe3+/salicylic acid complex. Reset both axes to
maximize the plot.
2. Print the graph showing the data and linear-regression equation of the Beer’s Law plot for the
standard solutions.
3. Determine the mass of the salicylic acid in the crude aspirin.
4. Determine the percent purity of your crude aspirin.
In addition to the regular components included in the summary, describe an alternate method for
determining the molar concentration the salicylic acid in the crude aspirin using the standard data.
Determination of an Equilibrium Constant
In this experiment, you will determine the equilibrium constant for the esterification reaction
between n-propyl alcohol and acetic acid.
O
O
CH3
HO
C
CH2CH2CH3
CH2CH2CH3
+ H2O
O
OH
Acetic Acid
(HAc)
CH3C
n-propyl alcohol
(PrOH)
n-propyl acetate
(PrAc)
water
The reaction is set up in such a way that the initial concentrations of PrOH and of HAc can be
determined. The reaction will be allowed to stand for a week, or more, to come to equilibrium. As
HAc reacts with PrOH, the acidity of the mixture will decrease, reaching a minimum once the system
reaches equilibrium. The quantity of acid present in the system will be determined by titration with a
standard base solution. The reaction is catalyzed by 6 M H2SO4. The concentration of the catalyst will
not change during the reaction, but its concentration must be determined. This can be done by
assembling a "blank,” which is a solution with the same amount of 6 M sulfuric acid and the same total
volume as the reaction mixture, but with no HAc or PrOH. Water is used as the solvent. Samples of a
blank are titrated with standard base to determine the quantity of acid catalyst added to the reaction.
The sample provided contains 10.00 mL of glacial acetic acid, 15.00 mL of n-propyl alcohol,
and 1.00 mL of 6 M H2SO4. The sample was prepared well in advance, and it is assumed that it has
achieved equilibrium.
Prepare a blank with water and the acid catalyst only to constitute the same total volume as the
samples. Titrate 1.00 mL of the sample [with the NaOH standardized by the class in a previous
experiment] and separately titrate 1.00 mL of the blank to determine the quantity of acid catalyst that is
present. You should have three repetitive titrations of each. After the titration results are adjusted for
the catalyst, you are to use the stoichiometry of the reaction and the initial concentrations of the
reactants to determine the concentrations of all chemical species at equilibrium in the reaction vessel.
This allows you to determine the equilibrium constant for the reaction.
The density of HAc is 1.0492 g/mL and the density of PrOH is 0.8037 g/mL. Prior to coming
to the lab, calculate the initial concentration of the HAc and ProH in the reaction mixtures.
Using an ICE table, determine the equilibrium constant at room temperature. Possible sources
of error?
Dissociation Constants of Weak Acids
Dilution effects, common ion effect, and heat of neutralization all suggest that weak acids are
incompletely dissociated and exist at an equilibrium condition between the molecules and the ions. The
equilibrium may be expressed
A  H3O
HA(aq) + H2O  H3O+(aq) + A-(aq)
Ka 
HA 
During a titration process the pH of the solution is constantly changing. The pH of the solution
may be monitored throughout the titration as the base is added and reacts with the acid. You will
measure the pH by use of pH probes calibrated at pH 4 and pH 10. A titration curve is obtained by
plotting the volume of standard solution added against the corresponding pH. A titration curve
contains a relatively vertical section which represents a rapid change of pH with a small volume of
standard solution. The midpoint of the vertical section [point of inflection] of the titration curve
corresponds to the equivalence point for the reaction. The volume of standardized base at this
equivalence point can then be used to determine the concentration of the unknown acid. In this
experiment, though, it will be used for another purpose.
 

If we look at the equilibrium system for acetic acid, we have
CH3COOH(aq) + H2O  H3O+(aq) + CH3COO-(aq)
H O CH COOH

CH COO 

Ka
3
3

3
Keep in mind that equilibrium is constantly being re-established during the titration, according
the chemical system H3O+(aq) + OH-(aq)  2 HOH. Thus, at one-half the volume of the
equivalence point, the [CH3COOH] = [CH3COO-], thus Ka = [H3O+], or pH = pKa. Experimentally,
what could be done to more precisely determine Ka?
Procedure
Logger Pro will be used to monitor the pH of the solution during titration. This process is
referred to as a potentiometric titration. After the pH probe has been attached, open “24a Acid-Base
Titration” in the folder Chemistry with Vernier. Change the maximum volume on the x-axis to 50 mL
[click on the current maximum volume and type 50 - Enter]. You might want to modify the label of
pH to take into account the different acids used. [Double click on the column label and modify
“Name”] The pH probe must be calibrated at pH 4 and pH 10. Save the calibration just in case the
computer goes down. You may then simply call up the calibration file rather than having to recalibrate. The bulb of the electrode on the probe must be covered with solution in order to operate
properly. The titration should be done in a beaker to provide space for the buret tip and the electrode.
Magnetic stirrers will also be available. Take no more base than is appropriate for your two titrations.
Acetic Acid
Measure about 20 mL of about 0.1 M CH3COOH into a 150 mL beaker and add about 20 mL
water. Place a stir bar into the beaker and put the beaker onto a magnetic stirrer. Fill a properly
prepared buret with standardized base [which you did last semester] and set it up to titrate the acetic
acid. This Logger Pro experiment is set up for you to record “Events with Entry.” [After you hit
“collect” a window will open where you enter the volume titrant added – then press enter.] Take the
pH of the solution before any base is added and then at 1.0 mL increments to 45-50 mL. Enter the
total volume added after each 1 mL increment. Save the data to disk or the DeskTop. Determine the
equivalence point of the titration curve by examining the graphs of the first and second derivative of
the titration curve. From these results, determine the Ka for acetic acid. In general, how does this
compare with the theoretical value?
Phosphoric Acid
Repeat the procedure as presented above, using about 20 mL of ~ 0.07 M H3PO4. Determine
the Ka1 and Ka2 for the phosphoric acid in a manner similar to the process you used for the acetic acid.
In general, how do these compare with the theoretical values? Be careful about the determination of
Ka2 for the phosphoric acid. The volume you should be examining is the net volume needed to
neutralize the second hydrogen ion. Keep track of what chemical species is present at various points
on the titration curve.
What part on the procedure introduces a probable error in the determination of the dissociation
constant(s) in this experiment? What could be done experimentally to decrease this error?
Solubility Product of Calcium Iodate
The purpose of this experiment is to determine the solubility product constant of calcium iodate
at the temperature of the laboratory. A saturated solution of calcium iodate contains this equilibrium:
Ca(IO3)2(s)  Ca2+(aq) + 2IO3-(aq)
with the solubility product constant given by: Ksp = [Ca2+] [IO3-]2
In this experiment you will measure the concentration of the iodate ion by titration with a
solution of sodium thiosulfate (Na2S2O3) using starch as an indicator. The concentration of the calcium
ion is obtained from the stoichiometry.
The iodate ions are reduced by I- ions from excess KI added to the titration mixture shown in
the following oxidation-reduction reaction in acid environment:
IO3-(aq) + 5I-(aq) + 6H+(aq)  3I2(aq) + 3H2O(l)
The reaction that occurs during the titration is given by:
2S2O32-(aq) + I2(aq)  S4O62-(aq) + 2I-(aq)
Starch is used as an indicator because it reacts with I2 to form a dark color. The dark color will
fade during the titration as I2 is consumed. The end point occurs when one drop of the Na2S2O3
solution causes the disappearance of the last trace of I2 and the solution changes from dark to colorless.
Solution Preparation
A saturated calcium iodate solution has been prepared by mixing equal volumes of 0.2M KIO3
and Ca(NO3)2. The supernate was decanted and the solid Ca(IO3)2 was washed with distilled water.
A quantity of distilled water was added to the washed solid and the mixture stored in a flask to come to
equilibrium. This is your saturated Ca(IO3)2 solution.
Procedure
Decant the saturated Ca(IO3)2 solution into a clean, dry beaker. Titrate the solution with the
approximate 0.025 M Na2S2O3 solution using appropriate iodometric techniques. Two trials, within
3% of each other, should be done. Each titration should include:
10.00 mL of saturated Ca(IO3)2 solution
0.5 g solid KI
1 mL 2M HCl
2 mL starch solution (added when the solution is a pale straw color)
distilled water to make a convenient volume
Record your data on the class spread sheet. Using the given value for the concentration of the
sodium thiosulfate, determine the Ksp of Ca(IO3)2 for each student trial. Then determine the class
average, standard deviation, and class precision. In general (no statistics), how does this compare with
the theoretical value? [Ksp Ca(IO3)2 = 6.5 x 10-6]. If it is assumed that you made no errors in
pipetting or titrating, what other factors in the experiment would lead to errors? Effects of those
errors?
Determination of K & Fe in a Complex Salt
Introduction: In a previous experiment, the percentage oxalate ion in the green iron oxalato
complex salt, KwFex(C2O4)y.zH2O, was determined. This experiment involves determining both the
percentage of potassium (K+) and of iron (Fe3+). A titration will be performed with a sample of a
solution that has passed through an ion exchange column. This solution will contain a known mass of
the iron complex salt.
Ion Exchange: Certain materials called ion exchange resins consist of rather large molecules which
contain ionizable groups. The resins are solids - insoluble in water and usually granular in nature which, when added to water, swell to form a slurry. The ionizable group on the resin ionizes
(exchanges ions) in the presence of water. This process is shown by the following equation for a resin
containing a sulfonic acid group (-SO3-H):
R-SO3-H + H2O ↔ R-SO3- H3O+ (the acidic for of the resin)
where 'R' represents the large insoluble resin molecule to which the sulfonic acid group is chemically
bonded, and H3O+ represents the hydronium ion bound to the resin sulfonate ion. This particular type
of resin is called a cation exchange resin, and the chemical form of the resin shown in the reaction
above is called an acid form resin.
A slurry of resin in water is poured into a vertical glass or plastic column equipped with a
porous plug at the bottom to trap the resin. Excess water is allowed to flow out, and the column
becomes filled with the water-soaked resin.
If an aqueous solution of a salt such as KCl (K+ and Cl- ions) is poured into the resin-filled
column, the KCl solution will displace the solution surrounding the resin and a volume of liquid equal
to the volume of KCl solution added will elute (be washed out) from the bottom of the column. In the
process, as the KCl solution passes down the column, potassium (K+) ions displace (exchange with)
hydronium ions, and aqueous HCl (H3O+ + Cl-) elutes from the column as shown by the following
reaction:
R-SO3- H3O+ + K+ + Cl- ↔ R-SO3- K+ + H3O+ + ClThus the solution coming out of the column (the eluate) will contain a quantity of hydronium (H3O+)
ions equal to the number of potassium (K+) ions that were added to the column. If enough potassium
ions are added, all of the acid form of the resin will be converted to the potassium form and at that
point the resin will become incapable of exchanging any more hydronium ions for potassium ions.
The resin is said to be “exhausted.”
However, it is possible to restore the resin completely to its acid form by pouring an aqueous
solution of HCl into the K+ saturated column. As the HCl solution passes down the column, the
hydronium ion displaces the bound potassium ions, and the reaction represented by the equation above
is reversed, with aqueous KCl eluting from the column.
Other cations such as Na+, Li+, Ca+2, etc. will exchange with the resin-bound H3O+ ions in a
manner similar to that of K+ - hence the term cation exchange resin. Anion exchange resins are also
available, but will not be used in this experiment.
Ion exchange resins are widely used in industry and in research laboratories to selectively
remove certain ions from solution. Home water-softening units, for example, are packed with a
sodium (Na) form of cation exchange resin which removes cations such as Ca+2, Mg+2, and Fe+2 that
cause water "hardness". These resins can be regenerated after they become saturated with the above
ions by passing an aqueous solution of NaCl through the unit.
Determining the Percentage of Potassium in KwFex(C2O4)y.zH2O Using Cation Exchange
When a weighed quantity of KwFex(C2O4)y.zH2O is dissolved in water, the salt dissociates
into ions according to the following equation:
KwFex(C2O4)y.zH2O  w K+ + Fex(C2O4)y-w (aq) + z H3O+(aq)
If this solution is passed down a column containing a cation exchange resin in the acid form, the K+
ions will replace the resin bound H3O+ ions according to the following equation:
x R-SO3- H3O+ + w K+ + Fe(C2O4)y(aq)-w  R-SO3- K+ + w H3O+ + Fex(C2O4)y(aq)-w
Thus for each mole of potassium ion (K+) added to the column, one mole of hydronium ion (H3O+ )
elutes from the column. If the eluted solution is titrated with a standardized NaOH solution, the moles
of H3O+ in the solution eluted from the column (hence the moles of K+ added to the column) can be
determined. Once the number of moles of K+ in the weighed sample of green salt has been
determined, the mass of K+ can be calculated. Since the mass of the green salt sample used in the
experiment is known, the percentage of potassium in the salt can be determined.
The ion exchange equation indicates that the solution which elutes from the column contains
the acid, w H3O+, and the anion, Fex(C2O4)y-w. When this acid is titrated with standardized NaOH,
the reaction that occurs first is given by the following equation:
H3O+ + OH-  2 H2O
After all of the acid is neutralized in the titration, further addition of NaOH results in the reaction
represented as:
Fex(C2O4)y-w + 3 OH-  Fe(OH)3 (ppt) + y C2O4-2
The ferric hydroxide precipitates from the solution as a reddish-brown, gelatinous precipitate when the
Ksp of Fe(OH)3 is exceeded. The above equation indicates that three moles of hydroxide ion are
required to react with each mole of iron ion in the salt. Thus the percentage of iron (Fe) in the sample
is found.
The Titration Curve
When the eluate from the ion exchange column is titrated with standard NaOH using a pH probe or
meter to follow the course of the reaction, a titration curve is obtained.
Two titration endpoints are obtained: the first, after the addition of V1 ml of NaOH; and the second,
after V2 ml of NaOH have been added. The first endpoint represents the completion of the
neutralization of the hydronium ion (H3O+), and the second endpoint represents the completion of the
precipitation of ferric hydroxide, Fe(OH)3. Thus V1 represents the OH- necessary to neutralize the
hydronium ion (H3O+) eluted and the quantity of V3 (V2 - V1) represents the OH- necessary to
completely precipitate the Fe(OH)3. Thus from a single pH titration curve of a weighed sample of
KwFex(C2O4)y.zH2O that has been passed through a cation exchange resin, it is possible to determine
both the % K and the % Fe in the compound.
Objectives:
1. To learn the principles and practice of using an ion exchange column.
2. To determine the % K and % Fe in KwFex(C2O4)y.zH2O using an ion exchange column.
Apparatus:
Ion exchange column (Bio Rad Econo-Column, cat. # 737-1011) packed with about 2 grams of BioRad AG 50 W-X2 100-200 mesh analytical grade cation exchange resin (cat. # 142-1241)
Chemicals:
NaOH – standardized (~0.1 M), buffer solutions (pH = 4 and pH = 10) for pH meter or probe
standardization, 1.0 M HCl(aq), pH Hydrion paper, student prepared green crystals of the complex
salt
Experimental Procedures:
1. The column has been prepared in advance according to the following: Weigh out about 2
grams of the solid resin into a weigh boat. Transfer the resin into the column. Fill the column
with distilled water and permit the resin to settle.
2.
Let the water level in the column drop until it is just above the top of the resin. Pour ~8 ml of 1.0
M HCl solution into the column and collect the eluate. Repeat this procedure twice more.
3.
Take your ion exchange column and mount it on a ring stand. It is important to make sure that
the resin bed is filled with liquid at all times. Using a clean 10-ml graduated cylinder (OR
using the volume markings on the column), rinse the column by pouring a 5-ml aliquot of
distilled water into the column and collect the liquid that elutes from the column in a small beaker.
Using a piece of wide-range pH paper, test the pH of the solution that first elutes from the column
to make sure that it is distinctly acid (pH << 7) If it is not acidic, immediately inform the teacher.
Allow the level of rinse liquid in the column to fall just to the top of the resin.
4.
Repeat the above rinse procedure two or three more times. When the level of water in the third
(fourth?) rinse has dropped to the level of the top of the resin, test a drop of eluate with pH paper
to confirm that the pH is about that of distilled water. If the eluate is still strongly acidic, continue
to rinse until the eluate is about the pH of distilled water. This assures that the only H3O+ in the
column is bound to the resin, and none is in solution.
5.
There will be some waiting time accompanying the rinse procedure. Make good use of this time
by looking ahead and preparing everything that will be needed for the day's work. Specifically,
setting up for the titration can be done between rinses.
6.
Using the electronic top-loader balance, weigh between 0.15 - 0.16 grams of the green crystal salt
in a weigh boat. Use unheated green crystals - not the anhydrous sample from the first analysis
experiment. Make sure that the sample mass does not exceed 0.165 grams! WHY? Take the
weighed sample and a clean, dry 50-ml beaker to the analytical balance. Tare the beaker. [OR
weigh the empty beaker to the nearest 0.0001 gram]. Remove the beaker from the balance.
[Remember that no chemical is ever added or removed from a container while the container is in
the balance compartment. This technique helps avoid damaging the balance pan as well as
preventing chemical spills on the weighing stage.] and pour the pre-weighed sample into the
beaker. Replace the beaker onto the analytical balance and record all appropriate masses.
7.
Using a 10-ml graduated cylinder, add a 5-ml aliquot of distilled water to the sample of green salt
and gently swirl the beaker until the salt is completely dissolved.
8.
Place a clean 150-ml beaker under the ion exchange column and quantitatively transfer the
solution of green salt to the column. The 150-ml beaker is to collect the eluate.
9.
Rinse the 50-ml beaker that held the green salt solution with 5 ml of water and when the level of
liquid in the column has dropped to the top of the resin, pour this rinse water into the column.
[The 150-ml beaker should remain under the column during this entire rinse process in order to
collect the eluate.] Repeat this rinse procedure with two more 5-ml aliquots of distilled water,
waiting each time until the liquid level in the column has dropped to the top of the resin before
adding more liquid to the column.
10. Set up a buret that has been thoroughly cleaned. Set up the pH meter assembly, including
magnetic stirrer, stir bar, and distilled water bottle. Open “24a Acid-Base Titration” in the folder
Chemistry with Vernier.
11. Standardize the pH probe with buffers of pH = 4 and pH = 10.
12. Obtain about 50 ml of the standardized 0.1 M NaOH solution [done in an earlier experiment] in a
clean, dry beaker. Record the precise concentration of the NaOH. Prepare the buret for titration.
13. After the last rinse has eluted from the ion exchange column, transfer the beaker containing the
eluate to the surface of the magnetic stirrer to be used in the pH meter titration.
14. Adjust the level of liquid in the beaker by adding distilled water so that the bulb of the pH
electrode is completely immersed in the solution. You may add distilled water as needed.
15. Titrate the solution with the standardized 0.1 M NaOH. Start with 1.0-ml increments (depending
on how much the pH jumps), then reduce the volume as appropriate. [Watch the pH change with
each increment. Initially, the pH change will be small. When the pH change exceeds 0.3 units,
decrease the increment added to 0.2 mL. Then as the pH change is again less than 0.3 units,
increase the increments to 1 mL again. Since you are looking for two equivalence points, this
incremental change will be repeated.] Prudence and patience is advised! Continue to titrate to
about 7 ml past the second equivalence point.
16. To regenerate the column so that it will be ready for the next time, pour about 10 mL of 1.0 M
HCl solution into the ion exchange column and place a beaker under the column to collect the
eluate. Pour about 10 mL of distilled water into the column. Place a cap on the column
17. From the graphs of the 1st and 2nd derivatives, determine V1 and V2.
18. Enter your mass of the complex green salt, V1, and V3 into the spreadsheet. Determine the
percentage of potassium and iron in the oxalato complex salt. After analysis of the class results,
use the 4d rule to determine if any of the values should be rejected. Carry out all appropriate
statistical calculations.
19. In a previous experiment, you determined the % oxalate ion in the complex compound. In this
experiment, you determined the % K+ and the % Fe3+. IF it is assumed that the remaining %
composition of the compound is water of hydration, determine the empirical formula of the green
crystals.
Electrical Cells & Thermodynamics
This investigation focuses on the reaction: Zn(s) + CuSO4  ZnSO4(aq) + Cu(s)
Using calorimetry, one can determine the heat of reaction (HR) for this system. An
electrochemical method offers another simple and accurate means for the determination of
thermodynamic quantities. A simple electrochemical cell Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) is
constructed as shown and the overall galvanic cell reaction is essentially the same as that which will be
taking place in the calorimeter.
The quantity of the electrical energy, E, produced or consumed during the electrochemical
reaction can be measured accurately. The free energy change, G, of an electrochemical reaction is
related to the voltage, E, of the electrochemical cell. By measuring the voltage, E, of our
electrochemical cell at several temperatures, one can obtain a plot of the free energy versus
temperature. Assuming that H and ΔS remain constant over a small temperature range, one can then
determine S and H from the graph, and determine G0.
Remember that an ideal calorimeter is a perfectly insulated vessel which contains a large
known mass of solution in perfect thermal contact with an accurate thermometer and a small reaction
tube. When measured quantities of reactants are introduced into the reaction tube, the heat of reaction
changes the temperature of the calorimeter solution. If we assume the specific heat of the CuSO4
solution is the same as for water, then the heat of reaction can be determined using the temperature
change and the mass of solution in the calorimeter.
Procedure
Part 1: Electrochemistry
1.
Pour sufficient 0.5 M ZnSO4 solution into a clean
porous cup to make it about half-full. Place the
porous cup into a clean, dry 100-mL beaker.
2.
Pour the 0.5 M CuSO4 solution into the 100-mL
beaker until the level of the CuSO4 solution in the
beaker is slightly above the level of the ZnSO4
solution in the porous cup.
3.
Fold and place a paper towel in the bottom of a
250-mL beaker. Place the 100-mL beaker into the
400-mL beaker. Add tap water [careful not to
contaminate any solution] to the 400-mL beaker
until the water is about the same as the level of the
CuSO4 solution in the 100-mL beaker.
4.
5.
6.
7.
8.
9.
10.
Obtain a zinc metal strip and sand it to remove any oxide coating [place the strips on a paper
towel so as to not scratch the lab bench]. Place it in the ZnSO4 solution. Repeat the procedure
with the copper metal strip and place it in the CuSO4 solution. [You might find it handy to
bend the top of each metal strip in an "L" near the top of its container.]
Place the entire assembly on a ring stand. The ring should be adjusted at a sufficient height for
heating with a Bunsen burner.
Place a thermometer in the 100-mL beaker of the CuSO4 solution. Support the thermometer if
necessary.
With a Bunsen burner, heat the water bath until the thermometer reads above 750C. Stop
heating.
While waiting for the temperature to stop rising, make sure the multi-meter is set to read in the
2 V (2000mV) range. Attach the alligator clips to the metal strips and observe. If a negative
number appears, reverse the wires. Disconnect one of the wires immediately.
Allow the bath assembly to cool naturally. Read and record 5-6 voltage - temperature readings
periodically in the 70-500C range. COMPLETE THE CIRCUIT ONLY AT THE TIMES OF
READING THE VOLTAGE.
Return solutions to the appropriate WASTE bottles.
Part II: Calorimetry
1.
Determine and record the mass of a coffee cup calorimeter. Using a graduated cylinder, add
about 50 mL of 0.5 M CuSO4 solution. Mass the calorimeter a second time.
2.
Let the apparatus stand so that the components attain the same temperature. Record the initial
temperature.
3.
On a pan balance, determine the mass of 0.5 - 0.6 g of zinc powder as precisely as you can.
4.
Add the zinc powder to the copper solution. Swirl thoroughly and observe the temperature.
Periodically use the thermometer to break up the copper-coated zinc powder. Record the
maximum constant temperature. Be patient, this is NOT as rapid a reaction as we have
observed before in calorimetry.
Analysis:
1.
Graphically, determine H and S. Calculate G0.
2.
Determine the HR per mole of zinc from your calorimetry data. Don’t forget the calorimeter
constant. How does this compare [quantitatively] with H as determined from the
electrochemical experiment? Discuss possible reasons for any difference.
3.
Calculate E0 from the experimental data. Quantitatively, how does your experimental standard
state potential compares with the theoretical value?
4.
Calculate Keq at 298K.
5.
Discuss the implications of the sign and magnitude of each of the five properties determined in
this experiment.
The 12 Bottle Problem
You are to determine the identity of 8 of 12 aqueous solutions containing relatively common
laboratory chemicals. The other 4 can be used as additional test solutions. Each sample contains only
one compound, and each sample contains a different compound (no duplicates – although some anions
repeat). The list below contains all the possible compounds in the 8 samples. The solution in bold are
available as additional test solutions and none of these is one of the eight unknowns.
You will be provided with samples of 8 solution.. Make sure you record the ID of your unknown, A,
B, C, etc. Your unknowns will be labeled 1-8. Refer back to the previous experiment “Chemical
Reactions: Part A – Double Displacement Reactions” for the procedure. Prior to coming to the lab: 1.
transfer the Double Displacement data to a new, neat table; 2. prepare an appropriate new data table
for your unknown; 3. predict and defend the acidity of the solution after the hydrolysis of each salt.
You will not be allowed refills so be sure you conduct your testing wisely. For the observations of
reactions, you may only use the “additional test solutions” and your eight unknown solutions. In
addition to the solutions themselves, you will be provided with cotton swabs for flame-tests and pH
paper.
The final lab report should include the identity of each solution in each of your 8 unknown clear,
colorless samples. You must defend the identification of each sample. You must include at least two
(2) unique chemical confirmation tests (a white ppt of AgCl from multiple combinations is ONE
chemical test). If it is relevant, you must also include the pH of the solution and the color of the flame
test. Appropriate net ionic equations are required along with the color of precipitates and complex
ions. The physical attributes alone are insufficient defense.
unknown
LiOH
KSCN H2SO4
AgNO3
solutions
SrCl2
NH3 Na2CO3 Ba(NO3)2
add’l test soln’s Al(NO3)3 FeCl3 CuSO4
CoCl2
12-Bottle Problem matrix
Test Tube
1
2
3
4
5
6
7
8
CoCl2
Al(NO3)3
CuSO4
FeCl3
A
B
C
Using Conductivity to Find an Equivalence Point and Ksp Value
In this experiment, you will monitor conductivity during the reaction between sulfuric acid, H2SO4 and
a saturated solution of barium hydroxide, Ba(OH)2, in order to determine the equivalence point. From
this information, you can find the concentration of the Ba(OH)2 solution, and then calculate the
solubility-product constant, Ksp, for the saturated solution . You will also see the effect of ions,
precipitates, and water on conductivity. The equation for the reaction in this experiment is:
Ba2+(aq) + 2 OH–(aq) + 2 H+(aq) + SO42–(aq) 
 BaSO4(s) + H2O(l)
Before reacting, Ba(OH)2 and H2SO4 are almost completely dissociated into their respective ions.
Neither of the reaction products, however, is significantly dissociated. Barium sulfate is a precipitate
and water is predominantly molecular.
As standardized H2SO4 (~0.30 M) is slowly added to Ba(OH)2 of unknown concentration, changes in
the conductivity of the solution will be monitored using a Conductivity Probe. When the probe is
placed in a solution that contains ions, and thus has the ability to conduct electricity, an electrical
circuit is completed across the electrodes that are located on either side of the hole near the bottom of
the probe body (see Figure 1). This results in a conductivity value that can be read by the interface.
The unit of conductivity used in this experiment is the microsiemens, or µS.
Figure 1
Prior to doing the experiment, it is very important for you to hypothesize about the conductivity of the
solution at various stages during the reaction. Would you expect the conductivity reading to be high or
low, and increasing or decreasing, in each of these situations?

When the Conductivity Probe is placed in Ba(OH)2, prior to the addition of H2SO4.
 As H2SO4 is slowly added, producing BaSO4 and H2O.
 When the moles of H2SO4 added equal the moles of BaSO4 originally present.
 As excess H2SO4 is added beyond the equivalence point.
MATERIALS
Power Macintosh or Windows PC
Vernier computer interface
Logger Pro
Vernier Conductivity Probe
60 mL of standardized (`~0.30 M) H2SO4
60 mL of saturated Ba(OH)2, unknown M
50-mL graduated cylinder
50-mL buret
stirring rod or magnetic stirrer
2 utility clamps
ring stand
two 250-mL beakers
Phenolphthalein (optional)
PROCEDURE
1. Measure out approximately 60 mL of ~0.30 M H2SO4 into a 250-mL beaker. Record the precise
H2SO4 concentration in your data table. CAUTION: H2SO4 is a strong acid, and should be
handled with care. Obtain a 50-mL buret and rinse the buret with a few mL of the H2SO4 solution.
Use a utility clamp to attach the buret to the ring stand as shown in Figure 1. Fill the buret a little
above the 0.00-mL level of the buret. Drain a small amount of H2SO4 solution so it fills the buret
tip and leaves the H2SO4 at the 0.00-mL level of the buret. Dispose of the waste solution from this
step as directed by your instructor.
2. Prepare the computer for data collection by opening the file in the Experiment 26 folder of
Chemistry with Computers. Set the selection switch on the amplifier box of the probe to the 020000 µS range. The vertical axis of the graph has conductivity scaled from 0 to 20000 µS.
Manually increase the y-axis scale to 35000 µS. The horizontal axis has volume scaled from 0 to
25 mL. Place the probe in a clean, dry 250-mL beaker containing 100 mL of distilled water.
Under the Experiment pull down menu, select Zero... in order to re-zero the probe’s conductivity
values that appear in the “Meter Window” display.
3. Use a Buchner funnel and vacuum filtration to remove any undissolved Ba(OH)2 from the saturated
solution. Then measure out 50.0 mL of the filtered Ba(OH)2 of unknown concentration using a 50mL graduated cylinder. Transfer the solution to clean, 250 mL beaker. Rinse the cylinder three
times with 10 mL of distilled water, and add the washings into the beaker. Caution: Ba(OH)2 is
toxic. Handle it with care.
4. Arrange the buret, Conductivity Probe, beaker containing Ba(OH)2, and stirring bar as shown in
Figure 1. The Conductivity Probe should extend down into the Ba(OH)2 solution to just above the
stirring bar.
5. Before adding H2SO4 titrant, click COLLECT and monitor the displayed conductivity value (in
µS). Once the conductivity has stabilized ( 100 µS), click KEEP. In the edit box, type “0”, the
current buret reading in mL. Press ENTER to store the first data pair for this experiment.
6. You are now ready to begin the titration. This process goes faster if one person manipulates and
reads the buret while another person operates the computer and enters volumes.
a. Add about 1.0 mL of ~0.30 M H2SO4 to the beaker. When the conductivity stabilizes, again
click KEEP. In the edit box, type the current buret reading to 0.05 mL. Press ENTER. You
have now saved the second data pair for the experiment.
b. Continue adding 1.0 mL increments of H2SO4, each time entering the buret reading, until the
conductivity has dropped below 5000 µS. Begin adding 0.5-mL increments, and enter each
buret reading, until the conductivity has dropped below 1000 µS..
c. After the conductivity has dropped below 1000 µS, use 2-drop increments (~0.1 mL) until the
minimum conductivity has been reached at the equivalence point. Enter the volume after each
2-drop addition. When you have passed the equivalence point, continue using 2-drop
increments until the conductivity is greater than 1000 µS again.
d. Now use 1.0-mL increments until the conductivity reaches about 20000 µS, or an additional 15
mL of H2SO4 solution has been added beyond the equivalence point, whichever comes first.
7. When you have finished collecting data, click STOP. Dispose of the beaker contents in a heavy
metal waste container as directed by your teacher.
8. Move the mouse to position the cursor at the initial data point where the conductivity begin
decreasing dramatically. Click and hold the mouse button as you drag across the linear potion of
the curve where the conductivity was decreasing. Release the mouse button to highlight this
section of the curve. Under the Analyze pull down menu, select Linear Fit in order to draw the
best-fit line for the decrease in conductivity.
9. Repeat step 9, except highlight the linear portion of the curve where the conductivity was
increasing again. Where these two best-fit lines intersect is the equivalence point for the titration.
Under the Analyze pull down menu, select Interpolate, then move the cursor to align it with this
intersection. The x value for the mL of H2SO4 needed to reach the equivalence point will be
displayed. Record this value.
10. Sketch or print out a copy of the Graph window with your name(s) entered on the graph.
PROCESSING THE DATA
1. From the graph that you printed, determine the volume of H2SO4 added at the equivalence point.
Record the volume of H2SO4.
2. Calculate the moles of H2SO4 added at the equivalence point, using the molarity, M, of the H2SO4
and its volume, in L.
3. Calculate the moles of Ba(OH)2 at the equivalence point. Use your answer in the previous step and
the ratio of moles of Ba(OH)2 and H2SO4 in the balanced.
4. From the moles and volume of Ba(OH)2, calculate the concentration of Ba(OH)2, in mol/L.
5. From the concentration of Ba(OH)2, calculate the concentration, in mol/L, of the Ba2+ and OH1ions in the saturated solution.
6. From the the concentration of the Ba2+ and OH1-, calculate the solubility-product constant, Ksp, of
Ba(OH)2.
DATA TABLE
Molarity of Standardized H2SO4
M
Volume of H2SO4 at Equivalence Point
mL =
L
mL =
L
Volume of Ba(OH)2
Moles of H2SO4
mol
Moles of Ba(OH)2
mol
Molarity of Ba(OH)2
M
Molarity of Ba2+
M
Molarity of OH1-
M
Ksp of Ba(OH)2
Teacher Notes: Typical data for this experiment can be seen on the graph below. The
“Linear Fit” lines draw for the portions of the curve where the conductivity was first
decreasing, and then increasing, intersect when the volume of sulfuric acid added was
20.8 mL. Based upon this equivalence point:



4 



0.295 moleH 2 SO4 
-3
x moles H 2 SO4 = 20.8 mL H 2 SO
mole H 2SO4
 = 6.14 x 10

1000 mL H 2 SO4 
Since 1 mole H 2 SO4 reacts with 1 mole Ba(OH)
x M Ba(OH) 2






2,
6.14 x 10 -3 mole Ba(OH) 2 reacted.

6.14 x 10 -3 mole Ba(OH) 2 
-1
=
M Ba(OH) 2
 = 1.23 x 10

0.0500 L Ba(OH) 2

x M Ba 2+ = M Ba(OH) 2 = 1.23 x 10 -1 M Ba 2+

x M OH
1-

1 2 moles OH

-1
 = 2.46 x 10 -1 M OH1= 1.23 x 10 M Ba(OH) 2 

1 mole Ba(OH) 

2 
Ksp = [Ba 2+ ] [OH1- ]2 = [1.23 x 10 -1 M Ba 2 + ][2.46 x 10 -1 M OH1- ]2 = 7.44 x 10 -3
The accepted value for the Ksp of Ba(OH)2 varies, depending upon reference source. The CRC lists the
value @ 25ºC as 2.55 x 10-4, whereas the Kotz & Treichel and Brown & LeMay textbooks list it as
being 5.0 x 10-3.
Determination of Phosphate in Beverages
Objective: In this experiment, a beverage containing phosphate (in the form of phosphoric acid) will
undergo
two different tests to determine the concentration of phosphate in it. One of the tests will
be a
titration, and the other will be done using the spectrophotometer and analyzing the
concentration
by the absorbance of light. The values obtained by each method will be compared.
Chemicals Used: ammonium metavanadate, ammonium molybdate, nitric acid, sodium hydroxide,
dibasic potassium phosphate
Preparation of Sample:
Each student should examine a different phosphate containing beverage (Coke, Pepsi, DietPepsi, Dr. Pepper, etc.). Cover the sample, about 150 mL should be sufficient, with a watchglass and
heat. Boil the sample gently for about 20 minutes to remove CO2. 25.00 mL of sample will be used
for each titration and a
5.0 mL aliquot will be used for the spectrophotometric determination.
Procedure:
Part 1: Titration of Phosphoric Acid
Titrate a 25.00 mL portion (measured by pipet) of your degassed sample with 0.02 M NaOH. The
endpoints
will be determined from a plot of pH vs. mL of titrant.
Part 2: Spectrophotometric Determination of Phosphate
In dilute solution, ammonium metavanadate (NH4VO3), molybdate (MoO42) and phosphate
form a yellow colored compound with the formula (NH4)3PO4 NH4VO3 16MoO4. The intensity
of the color of a solution containing metavanadate, molybdate and phosphate is directly proportional
to the concentration of phosphate; Beer’s Law is followed. A method based on this reaction can be
used to determine phosphate concentrations as low as 1x105M (10 ppb), Both the reagent and the
sample have measurable absorbances at 400 nm, the wavelength of peak absorbance of the
compound formed in the reaction between the reagent and phosphate. Thus “blanks” will be used to
determine these absorbances. Beer’s Law predicts a linear relationship between absorbance and
concentration and also a simple additivity of absorbances.
Prepare a diluted sample of your beverage by diluting 5.0 mL (measured by pipet) of the
degassed sample to 250.0 mL in a volumetric flask.
Prepare a calibration curve at 400 nm using the spectrophotometer. Six standard solutions
are provided. Determine the absorbance of solutions prepared by adding 5.0 mL of the ammonium
vanadomolybdate reagent to 10.0 mL of each standard solution. Also measure the absorbance of a
reagent blank (5.0 mL of reagent and 10.0 mL of deionized water) and a sample blank (5.0 mL of
deionized water and 10.0 mL of diluted beverage sample). You will find it helpful to plot the data for
the reagent blank and the standards in Excel and add a trendline. Use this “curve” to determine the
concentration of the phosphate in your sample. Remember to subtract the absorbance of the sample
blank from the absorbance of each of the beverage sample determinations performed below.
Triplicate analyses of the beverage sample should be performed. Again, 5.0 mL of ammonium
vanadomolybdate reagent is added to 10.0 mL of the diluted beverage sample before the absorbance
(PO43)
is noted.
Clean Up: All waste containing ammonium vanadomolybdate is to be placed in the waste beaker
labeled for
this waste. Do NOT pour it into the sink.
Lab Report:
Summarize your results for this laboratory exercise in your laboratory notebook, separate from your
data. Provide average values for the concentration of phosphate in your beverage as determined by
titration and spectroscopically (don’t forget that the beverage was diluted for this latter analysis). Your
calibration curve for the spectroscopic determination, with tabulated absorbance values must be
included.

Determination of an Empirical Formula
From previous experience in chemistry, you are aware that elements combine in whole number
ratios to form compounds [Law of Definite Proportions]. Some elements even have the capability to
combine in different ratios to form different compounds [Law of Multiple Proportions]. The formula
which contains the simplest whole number ratio of elements in a compound is referred to as an
empirical formula.
In this experiment, you will react tin with nitric acid and dehydrate the product to an anhydrous
oxide. Tin exists in binary compounds as either tin(II) or tin(IV).
Safety: This experiment uses concentrated nitric acid. To contain
the fumes, assemble an individual fume hood consisting of an
inverted funnel connected to the water aspirator with a piece of
rubber tubing. Concentrated acids are dangerous! They will
“burn” the skin and the vapors can damage the respiratory system.
The gas generated in this experiment can also be quite damaging to
the respiratory system, if inhaled. Minimize breathing any of the
vapors in this experiment.
PROCEDURE
Using a pan balance, measure and record precisely the mass
of an evaporating dish covered with a watch glass. Measure 1-1.5
g of solid tin granules into the dish assembly and record the mass precisely. Prepare to heat the
evaporating dish with a Bunsen burner. Above the evaporating dish, assemble an individual fume
hood consisting of an inverted funnel connected to the water aspirator with a piece of rubber tubing.
Add 15 M HNO3 (6-7 mL), very carefully and slowly,
dropwise, through the lip of evaporating dish, to the tin until the
reaction is complete. Heat the dish and contents slowly until the
excess HNO3 is eliminated, then heat strongly. The desired
product will be light yellow in color. When your compound
appears dry, crush the product to uniform small granules. Heat
to constant mass.
Dispose of the compound in the waste basket
Enter your data into the class spreadsheet.
Determine the empirical formula of the compound in the evaporating dish. Evaluate and discuss the
results. Sources and effects of errors?
Determining Mole Ratios in a Chemical Reaction
Background:
A balanced chemical reaction equation gives the mole ratios of the reactants and the products as
coefficients. When some of the chemical formulas are not known, an experiment must be conducted to
help determine the mole ratios.
This experiment uses two common substances as the reactants: hypochlorite ion (ClO–) from household
bleach (NaClO) and thiosulfate ion (S2O32–), the active ingredient (Na2S2O3) in a photographic “fixer”
solution used to develop black & white film. In the reaction, hypochlorite ions reacts with the
thiosulfate ions according to the unbalanced and incomplete reaction equation below.
a ClO– + b S2O32– → products
It is possible to identify the coefficients, a and b, for the reactants, without knowing the products of the
reaction. The process that you will use to determine the coefficients is called continuous variation. You
will prepare a series of mixtures of the two reactants. Each mixture will have the same total volume
and the same total number of moles of reactants. The reaction is exothermic, thus the mixture that
generates the most heat energy will be the reaction that completely consumes both the hypochlorite
and the thiosulfate ions. You will use this mixture to establish the coefficients, and therefore the mole
ratio, for the reaction.
Objective:
Determine the stoichiometry of an oxidation-reduction reaction in which the reactants are known, but
the products are unknown, using the method of continuous variations.
mol solute 
Molarity  M 
 is the most common unit of concentration in chemistry. As part of your

L sol' n

pre-lab, have calculated the mass of Na2S2O3.5H2O to prepare 100.00 mL of a 0.5 M solution and also
have calculated the mass of NaOH to prepare 100.00-mL of a 0.2 M NaOH solution.
Procedural Guidelines:
1. Prepare the solution of 0.50 M and 0.2 M NaOH in a 100mL volumetric flask. Make sure the solution is
homogeneous.
2. Prepare the NaClO solution (0.5 M) by adding 150.0 mL of
6% bleach to 90.0 mL distilled water. Make sure the
solution is homogeneous.
3. Use a double-cup calorimeter and thermometer. Also use
two different graduated cylinders to measure the quantity of
chemicals as precisely as you can.
Volume
Volume
NaClO(aq)
Na2S2O3(aq)
(mL)
(mL)
55
50
45
40
35
4. Complete one (1) series of various ratios of the two solutions
keeping the total volume of each experiment at 60.0 mL. (Vary the volume ratios of ClO- and
S2O32- in a systematic manner and record the initial and final temperature.) Prior to coming to
lab, prepare a data table in your lab notebook, similar to the one shown, in which to record your
data. Obtain data from two other lab groups.
Safety Considerations:
Sodium thiosulfate is slightly toxic by ingestion and a body tissue irritant.
Sodium hypochlorite solution is a corrosive liquid which causes skin burns and reacts with acid to
evolve chlorine gas and evolves chlorine when heated; moderately toxic by ingestion and
inhalation; avoid contact with organic materials. [The solution we are using is a slightly diluted
solution of common bleach. Do be careful.]
Sodium hydroxide is a corrosive solid and skin burns are possible; extremely hazardous to eyes –
wear splash goggles when using this substance. [The solution we are using will be slippery and
possibly make a hangnail “tingle”.]
Materials:
Graduated cylinders; 0.50 M NaClO solution; 0.50 M Na2S2O3 solution in 0.2 M NaOH;
Styrofoam cups; Thermometer
Analysis:
1. Using LoggerPro, make a graph (points only) of volume of hypochlorite solution used (x) vs
temperature change (y) of data obtained. Add regression lines to the distinct linear regions of
the graph. Find the point of intersection of the two lines. From the volume of NaClO at this
intersect, determine the whole number ratio of the volume NaClO : volume Na2S2O3 which,
which in this case, represents the coefficients of the reactants in the chemical equation.
In addition to summarizing the results:
2. Comment/explain the two “best fit lines” on the plot.
3. The molarities of the reactant solutions were approximately equal in this experiment. Is this
necessary? Why/why not?
Reduction Potential Series
Oxidation-reduction reactions involve a change in the oxidation number of some of the
elements involved in the reaction. Oxidizing agents and reducing agents differ in their strength. In
general, the strongest oxidizing agents have the greatest tendency to take electrons. The strongest
reducing agents have the least tendency to hold on to their own electrons. In one way, an oxidationreduction reaction is analogous to an acid-base reaction. In the latter, two substances are competing
for a proton. The strongest bases have the greatest tendency to take protons. In a redox reaction, two
substances are competing for an electron. Like acid and base strength, redox agents can be ranked in
order of their oxidizing or reducing strength. The table can then be used to predict the tendency for a
given oxidizing and reducing agent to react.
In this experiment you will interpret a series of reactions to determine the relative strengths of
several oxidizing and reducing agents. For example, you have already observed that zinc will react
with HCl to produce hydrogen gas and zinc chloride.
Zn(s) + 2 H+  H2(g) + Zn2+
However, bubbling H2 gas through a solution of zinc chloride does not produce an observable reaction.
On the basis of these observations, you would conclude that H+ is capable of removing electrons from
zinc atoms but that zinc ions cannot remove electrons from H2 molecules. Therefore, you would
classify H+ ion as a stronger oxidizing agent than zinc ions, and zinc metal a stronger reducing agent
that H2 gas. The two half-reactions can be expressed as reductions and listed in decreasing order of
their relative strengths as an oxidizing agent:
2 H+ + 2 e-  H2
Zn2+ + 2 e-  Zn
In this arrangement the strongest oxidizing agent (H+) is at the upper left, and the strongest reducing
agent (Zn) is at the lower right. From another perspective, any chemical below and to the right can
“replace” any chemical above and to the left in a chemical reaction. By interpreting the results of a
number of reactions you will be able to prepare a short table of metals and their ions in order of their
relative reducing and oxidizing strengths. You will do the same for aqueous solutions of the halogens
and the halide ions.
Materials:
spot plate; 13 x 100 test tubes; small pieces of Mg, Cu, Zn, Pb; dropper bottles of 6 M
HCl, 0.1 M AgNO3, Mg(NO3)2, Cu(NO3)2, Zn(NO3)2, Pb(NO3)2, NaCl, NaBr, NaI;
aqueous solutions of Cl2, Br2, I2; CH2Cl2
Procedure
Suggestion: a matrix is handy for recording data in this type of experiment
1.
Observe and record the color of each solution used in the experiment.
Demo by instructor
NOTE: the dichloromethane (CH2Cl2) is a solvent only. It does not enter into any chemical reactions
with the other chemicals.
2.
Put a few drops of aqueous Cl2, Br2, and I2 into separate aliquots (~ 1 mL) of CH2Cl2. Record
observations. [Special attention must be given to the disposal of the CH2Cl2 mixture.]
3.
Put a few drops of aqueous NaCl, NaBr, NaI into separate aliquots (~ 1 mL) of CH2Cl2.
Record observations.
4.
Look for evidence of a reaction between each halogen with each halide solution provided by
adding a few drops of the halogen to about 1 mL of the halide solution. Record observations.
Add about 1 mL of CH2Cl2 to each mixture, shake, observe and record.
Write balanced net ionic equations for each & every reaction that occurred. Prepare a
reduction potential series, as per the example, in decreasing order for the halogens and the halide ions.
Exp by student
5.
Using a spot plate, look for evidence of a reaction between each metal and HCl. Also, test each
metal with each metallic salt solution provided. Be patient!
Write balanced net ionic equations for each & every reaction that occurred. Prepare a
reduction potential series, as per the example, in decreasing order for the metals and metal ions
provided, including hydrogen.
Analysis of Hydrogen Peroxide
Most commercial solutions of hydrogen peroxide are approximately 3% by mass. In this
experiment you will analyze, by titration with potassium permanganate, one of these products. In acid
solution MnO4- oxidizes H2O2 to form oxygen gas and colorless Mn2+. The end point is reached when
the first appearance of excess permanganate causes a pale pink color. Thus, the permanganate ion is a
self indicator.
Procedure
Prepare 100.00 mL of about 0.12 M KMnO4 [determine the mass of KMnO4 to make this
solution as part of the pre-lab] such that the concentration is known to four significant figures.
Volumetric flasks will be available to prepare the solution. Do not take the time to try to measure that
exact mass, but record the precise mass you do measure. This mass/concentration of KMnO4 goes into
solution with difficulty. Be patient. “Listen” for undissolved crystals.
Precisely pipet a 10.00-mL sample of hydrogen peroxide into an appropriately sized
Erlenmeyer flask. 15 mL of 6M sulfuric acid is added to each peroxide sample to make the solution
acidic. Add 10-15 mL of distilled water to each sample to make an appropriate volume to view. Three
good titrations are required for each individual. Record the mass of your KMnO4 used for your
standard solution and the initial and final volumes of KMnO4 solution used for each titration on the
class spreadsheet.
Report
Provide the balanced redox equation for the reaction between permanganate and peroxide in
acid solution and the percent by mass of hydrogen peroxide in each titration. Complete appropriate
statistical analyses. An accuracy calculation is not appropriate since the manufacture’s statement of
concentration is only an approximation.
The Synthesis & Analysis of Alum
The term alum is a general family name for a crystalline substance composed of cations with 1+ and
3+ charges. In this experiment, you will synthesize a type of alum called potassium aluminum sulfate
xxxahydrate, KAl(SO4)y•x H2O. You will synthesize this compound by placing the appropriate ions in
one container in aqueous solution, form the alum crystals, and determine the number of sulfate ions
and waters of hydration.
This particular compound has been chosen because it is relatively simple to prepare a pure sample. The
process of synthesizing this compound is interesting in that it involves both chemical and physical
reactions. Chemically, aluminum is oxidized from aluminum foil to prepare the Al3+ ions. Physically,
as the solution that contains the mixture of ions evaporates, crystals will form which contain a number
of waters of hydration bonded to the aluminum ion and a number of waters bonded to the potassium
ion.
Aluminum is considered a reactive metal, but because its surface is usually protected by a thin film of
aluminum oxide, it reacts slowly with acids. It does, however, dissolve quickly in basic solutions.
Excess hydroxide ion converts the aluminum to the tetrahydroxoaluminate (Al(OH)4-) precipitates.
Continued addition of acid causes the hydroxide ions to be completely neutralized, and the aluminum
exists in solution as the hydrated ion. Aluminum hydroxide is considered to be an amphoteric
hydroxide because it dissolves in both acids and bases.
OBJECTIVES
In this experiment, you will



Synthesize a sample of alum and determine its formula
Observe and record the process of synthesizing a compound.
Calculate the percent yield of your synthesis.
PROCEDURE
Part I: Synthesis of Alum
1. Obtain a piece of aluminum foil and measure its mass. For best
results, you should have about 1.00 g of aluminum. Tear the foil
into small pieces and place the pieces in a 250 mL beaker.
2. Set up a Büchner funnel and filter flask so that you are ready to
filter the reaction mixture that will be produced in Step 4.
3.
Conduct the first part of the synthesis. CAUTION: Potassium hydroxide solution is caustic. Avoid spilling
it on your skin or clothing.
Use a graduated cylinder to measure out 25 mL of 3 M KOH solution.
Slowly add the KOH solution to the beaker of aluminum pieces. Notice that the reaction is
exothermic. Allow the reaction to proceed until all of the foil is dissolved.
Carefully pour the reaction mixture through your Büchner funnel and filter flask setup, and rinse
the filter paper with a small amount of distilled water. Note: The reaction mixture contains three
ions: K+, [Al(OH)4–], and excess OH–.
Rinse the beaker with distilled water, and pour the filtered liquid back into the beaker.
4. Cool the solution to near room temperature.
5. Clean the Büchner funnel and filter flask, and prepare it for more filtering.
6. Complete the synthesis.
a. Use a graduated cylinder to measure out 35 mL of 3 M H2SO4 solution. CAUTION: The
reaction mixture must be cooled to room temperature before proceeding. Handle the H2SO4
solution with care. It can cause painful burns if it comes in contact with the skin.
b. After the reaction mixture has cooled, slowly add the sulfuric acid solution to the beaker of
liquid. Stir the mixture constantly. The reaction is strongly exothermic, so be careful as you stir
the mixture. Note that aluminum hydroxide will precipitate initially. It will dissolve as more
sulfuric acid is added.
If there is some solid remaining in the beaker after the 35 mL of sulfuric acid has been added, pour
the mixture through the Büchner funnel and filter flask to separate the undissolved solid from
the mixture.
7. Gently boil your mixture until you have about 50 mL of liquid in the beaker.
8. Cool the beaker of solution. Prepare an ice bath. Place your beaker of solution, uncovered, in the
ice bath. Do not move the ice bath or the beaker. After about fifteen minutes, crystals of alum will
appear in the beaker. If there are no crystals after fifteen minutes, scratch the bottom of the beaker
with a glass stirring rod to create a rough spot for crystal growth. You may also heat the solution to
evaporate more water and cool the solution again.
9. Collect your alum crystals by pouring them onto the clean Büchner funnel and filter-flask setup. Use about
50 mL of a 50% aqueous ethanol solution to transfer and wash the alum crystals in the vacuum filtration
system. The crystals will not dissolve in this solution.
10. Remove the filter and crystals from the Büchner funnel and allow the crystals to dry at room temperature.
Measure and record the mass of your sample of the dry alum.
Part II: Determine the Percent Water in Alum
11. Clean and paper-dry a crucible and cover. Heat the crucible and cover over a burner flame until it is red
hot. Allow the crucible to cool, and then measure the total mass of the crucible and cover. Handle the
crucible with tongs or forceps to avoid getting fingerprints on it.
12. Place about 2 g of your alum crystals in the crucible, and then measure the mass of the crucible, cover, and
alum.
13. Set up a ring, ring stand and triangle over a lab burner. Use tongs or forceps to set the crucible at an angle
on the triangle and place the cover loosely on the crucible. Use a lab burner to very gently heat the
crucible of alum until you can see no vapor escaping from the crucible. It is important that the vapor does
not carry any alum with it. After the vapor is gone, heat the crucible more strongly for five minutes, and
then cool the crucible.
14. Measure and record the mass of the crucible, cover, and alum.
15. Reheat the crucible and alum sample until a constant mass is obtained.
Part III: Determine the Percent Sulfate in Alum
16. Label a 250-mL beaker and determine its mass. Measure about one gram of your alum sample into the
beaker. Add about 50 mL of distilled water to the beaker of alum and stir the mixture to dissolve the
sample.
17. Slowly add about 80 mL of 0.20 M Ba(NO3)2 solution to the beaker of alum solution. Stir the mixture to
ensure complete mixing of the reagents. CAUTION: Handle the barium nitrate solution with care. This
solution is toxic if ingested.
18. Set up a ring stand, ring and wire gauze for heating over a lab burner. Place a watch glass over the beaker
and heat the beaker of the reaction mixture to near boiling for about 20 minutes. This step helps collect
the particles of precipitate to a larger size and eases the filtering process.
19. Allow the mixture in the beaker to cool overnight. The next day, decant carefully.
20. Wash the BaSO4 with 15 mL distilled water. Cool overnight. Decant. Repeat one more time.
21. Place the beaker and precipitate in the drying oven overnight. After the precipitate is dry and cool,
measure and record mass. Discard the dry BaSO4 into the waste basket and clean the beaker.
Results
Describe the alum crystals. Once you have determined the percent water in alum, and the percent sulfate ion
in alum, you can determine the formula of alum. Using the aluminum foil as the limiting reagent, determine
the percent yield of your alum crystals. Possible errors and affects of these errors?
Write the balanced net ionic equations for the following: (a) aluminum and potassium hydroxide, yielding the
tetrahydroxoaluminate ion and hydrogen gas; (b) hydrogen ions and the tetrahydroxoaluminate ion yielding
aluminum hydroxide; (c) aluminum hydroxide and hydrogen ions, yielding hexaquoaluminum ions; and (d) the
formation of alum from potassium ions, sulfate ions, hexaquoaluminum ions and water.
Experiment 11: Determination of the % Water in an Iron Oxalato Complex Salt
Introduction: In this experiment you will determine the mass of the water of hydration in an
Iron Oxalato Complex Salt.
In a previous experiment a green crystalline product having the formula KwFex(C2O4)y.zH2O was
prepared. The percentage water in KwFex(C2O4)y.zH2O will now be determined.
The green iron oxalato complex is one of a number of solid chemicals that are classified as
"hydrates". A hydrate contains water chemically bound in the solid state so that it is present in
the compound in stoichiometric amounts. Familiar examples of hydrates are Plaster of Paris
(CaSO4.1/2H2O), gypsum (CaSO4.2H2O), and alum [KAl(SO4)2.12H2O]. The water of
hydration of many hydrates can be removed as a gas by heating the hydrate to a temperature
above 100oC for a period of time. The following reaction involving barium chloride dihydrate
(BaCl2.2 H2O) occurs when the solid is heated above 100oC.
BaCl2.2H2O(s)  BaCl2(s) + 2 H2O(g)
The percentage of water of hydration in KwFex(C2O4)y.zH2O will be determined in this
experiment by heating a weighed sample of green hydrate in an open container until all of the
water of hydration has been driven off.
KwFex(C2O4)y.zH2O -------------->
KwFex(C2O4)y + z H2O(g)
The loss in weight will be set equal to the mass of the water of hydration.
Safety, Environmental, and Economic Concerns:
Waste chemicals from this experiment may be safely discarded in the solid waste receptacle in
the lab.
...........................................................................................................................
Notes on Experimental Procedures:
1. The mass of the green crystals must be accurately measured both before heating as well as
after heating. Continue to dry, cool, and weigh the sample until 2 consecutive masses agree
to within 0.0010 gram of each other. [This process is referred to as “heating to constant
mass.”
2. Always have crucible lids at least slightly open while they are being heated. The crucible
should be completely covered while they cool. After they have cooled, the crucibles &
covers should not be touched by fingers until the experiment is completed due to the
moisture and oils your hands may impart. Handle the crucible & cover with crucible tongs.
3. The most frequent cause of erratic balance readings is failure to have objects at room
temperature when they are being weighed. Convection currents in a closed balance
compartment can have a surprising effect. Be sure that the glass windows on the balance are
closed on both sides, by the way.
4. A small inaccuracy in the calibration of a balance is canceled out in the final results when the
mass of the empty container and the mass of the container plus the substance of interest are
measured using the same balance. Using the same balance for successive weighings when
attempting to attain constant mass is important for that same reason.
Materials
Student prepared green crystals
Crucibles w/lids
Experimental Procedures:
1. Clean and paper-dry a crucible and cover. Heat the crucible and cover over a burner flame
until it is red hot. Allow the crucible to cool, and then measure the total mass of the crucible
and cover on the analytical balance. Handle the crucible with tongs or forceps to avoid
getting fingerprints on it.
2. Place about 1.0 gram of the green crystals into the crucible, and then measure the mass of the
crucible, cover, and green crystals.
3. Place the crucible, cover, & green crystals in an oven set to around 1100C. Use tongs or
forceps to set the crucible at an angle on the triangle and place the cover loosely on the
crucible. Heat the apparatus overnight. Allow the crucible and contents to cool to room
temperature.
4. Measure and record the mass of the crucible, cover, and green crystals.
5. Repeat steps 3-4 until a constant mass is obtained.
6. After the final weighing, place the amber bottles with the remaining crystals (capped) back
into the lab drawer/cabinet.
7.
Enter your data into the spreadsheet.
After analysis of the class results, use the 4d rule to determine if any of the values should be
rejected. Carry out all appropriate statistical calculations.
Determination of Iron by Visible Spectrophotometry
Visible Spectrophotometry is a method of measuring the concentration of colored
solutions by the amount of light absorbed by or transmitted through a colored solution. The
absorbance of white light by a solution containing a colored compound is directly proportional to
the concentration of the colored compound. The constant of proportionality contains the path
length of the sample through which the light passes and a constant that is determined by the color
of the solution. This information produces Beer's Law: A = bc, where A is absorbance,  is the
molar absorptivity of the colored solution, b is the inside diameter of the cell, and c is the molar
concentration of the solution. In this experiment we will not attempt to determine the values for 
and b. Rather we will draw a graph of absorbance, A, determined from the spectrophotometer
vs. c and use it as a calibration curve or a conversion curve. From the calibration curve we will
be able to determine the concentration of an unknown solution.
For a complete discussion of the spectroscopic theory of this experiment, see Appendix Three of
Zumdahl.
The purpose of this experiment is for you to
analyze a sample containing an unknown amount of
iron. The iron in the sample will be quantitatively
converted to a colored complex so that the
spectrophotometer can be used. The complex is
represented by the diagram to the right.
You will be provided with a standard iron
solution that contains acidified ferrous ammonium
sulfate such that 1 mL = 50.0 g Fe. Any iron(III) is
reduced to iron(II) by the hydroxylamine hydrochloride.
The calibration curve is determined from the following
solutions:
N
N
N
Fe2+
N
N
N
Place 2 mL, 4 mL 6 mL, 8 mL, and 10 mL of standard iron solution in five properly
labeled 100 mL volumetric flasks. To each flask add, in sequence:
2 mL of 1M ammonium acetate
2 mL of 10% hydroxylamine hydrochloride
20 mL of 0.30% o-phenanthroline solution (the complexing agent)
dilute to volume with distilled water.
Mix well and allow the color to develop for 45 minutes.
The experiment is in three parts:
Part 1: Determination of the optimum wavelength
Wavelength
Sample
Compartment
The operation of the Spec 20 will be
demonstrated
in the laboratory. Fundamentally,
Control
a lamp providing white light is separated into
component colors by a diffraction grating. It is
focused and passes through a sample where some
of it is absorbed. The remaining light is detected
by a photo tube and the quantity of light
absorbed (or transmitted) is read directly from
the meter.
Transmittance/
Power Switch/
Absorbance
SPECTROPHOTOMETER
Control
Zero Control
Experimentally, it has been found that
lamp
the quantity of light absorbed is directly
proportional to the molar concentration of the
METER
phototube
solution. This is represented by Beer’s Law,
lens
diffraction
A = bc , where A is the absorbance,  is the
slit
lens slit lens grating lens
molar absorbtivity, b is the path length the
sample
light travels through the solution, and c is the
mirror
molar concentration of the solution. If the
mirror
quantity of light that enters the solution is

represented by I0 and the quantity of light
leaving the solution is represented by I, then
the internal transmittance of the solution can be represented by T, and that the relationship
between the quantity of light absorbed A and the quantity transmitted T can be expressed as:
1
I
and A  log 10
.
T
T
I0
On the Spec-20 meter, transmittance is a linear scale, while absorbance is a logarithmic
scale. Log scales are difficult to read, thus, for increased precision, transmittance is normally
read and converted to absorbance.
Measure the transmittance at the associated wavelengths at 15 nm intervals for the range
of the instrument using one of the intermediate solutions to determine an approximate optimum
wavelength. For the report, prepare a graph (smooth curve of “connect-a-dot”) of wavelength (of
the entire spectrum) vs. absorbance to show the wavelength at which maximum absorbance
occurs.
Part 2: Calibration Curve
Set the instrument at the optimum wavelength for the solution. Measure the
transmittance for each prepared solution. Graph absorbance vs. concentration (g/mL) to
prepare the calibration curve. This should be a xy-plot [data points only] on which the
regression line is superimposed. Include the equation of the regression line and the correlation
coefficient.
Part 3: Concentration of an unknown solution
Transfer 25.00 mL well water into a 100-ml volumetric flask. Add ~1 mL 6 M
H2SO4and the same quantity of the other chemicals as instructed for the standard solutions.
Dilute to volume. After appropriate measurements, calculate the concentration g/mL) of the
unknown solution by using the equation of the regression line. Don’t forget that you diluted the
well water.
Titration Curves & Indicators
Several titrations have been performed in this course using indicators to determine the
equivalence point of a reaction. The indicator endpoints which were observed were relatively
sharp so that it was not difficult to locate them. For solutions that are highly colored or turbid the
use of visual indicators may not be possible or practical. During the titration process the pH of
the solution is constantly changing. A pH probe can be used to monitor the pH of the solution as
base is added and reacts with the acid. A titration curve is obtained by plotting the volume of
standard solution added against the corresponding pH. A titration curve contains a vertical or
almost vertical section. This section represents a rapid change of pH with a small volume of
standard solution. The midpoint of the vertical section (point of inflection) of the titration curve
corresponds to the equivalence point for the reaction.
In this experiment you will titrate a strong acid with a strong base, a weak acid with a
strong base [done in a previous experiment], a strong acid with a weak base, a weak acid with a
weak base and a polyprotic acid with a strong base [done in a previous experiment]. You will
examine various acid-base indicators to evaluate which indicator might be appropriate for the
various acid-base systems.
HCl-NaOH
HCl-NH4OH
CH3COOH-NH4OH
thymol blue
methyl orange
bromothymol blue
Logger Pro will be used to monitor the pH of the solution during titration. After the pH
probe has been attached, open “07a Acid-Base” in the folder Advanced Chemistry w Vernier.
Change the maximum volume on the x-axis to 50 mL [click on the current maximum volume and
type 50 - Enter]. You might want to modify the label of pH to take into account the different
acids used. [Double click on the column label and modify “Name”] Since you are interested
only in the titration curves (not the derivative curves), delete the columns d1 and d2. [click on
column title – Data tab – Delete Column – select column to delete] The pH probe must be
calibrated at pH 4 and pH 10. Save the calibration just in case the computer goes down. You
may then simply call up the calibration file rather than having to re-calibrate. The bulb of the
electrode on the probe must be covered with solution in order to operate properly. The titration
should be done in a beaker to provide space for the buret tip and the electrode. Magnetic stirrers
will also be available. Take no more base than is appropriate for your titrations.
All titrations are to be done with about 20 mL of acid with about 20 of distilled water
added. One-mL quantities of titrant are to be added incrementally so each titration curve is
‘symmetric’. Since the titration will be monitored continuously, only one trial is needed for each
titration. Add sufficient indicator to give a reasonable coloration to the solution. During the
titration, make sure you record the color of the indicator at each pH. Save the data to disk or
desk-top.
Report
Plot all five titration curves [including the HC2H3O2 - NaOH and H3PO4 - NaOH] for
each acid-base system on one graph, with the appropriate regions marked as to all theoretical
indicator color changes [including phenolphthalein]. Discuss the pH range you recorded with the
theoretical range for each of the indicators. Discuss the similarities and differences of the
titration curves. Using the theoretical pH ranges, discuss the feasibility of each indicator for
each acid-base system.
Materials
per pair
250 mL ~0.1 M NaOH
25 mL each:
~0.1 M HCl [A:W = 1:120]
~0.1 M CH3COOH [A:W = 1:170]
~0.1 M NH4OH [B:W = 1:140]
buffers: pH 4 & 10 for probe standardization
bromothymol blue
methyl orange
thymol blue
pH meter/probe
buffers: pH 4 & 10 for probe standardization
One Tube Reactions
Background Information:
In this experiment, you will take home a set of chemicals and materials which you will set up and
observe for a period of days. A written record to be handed in at the beginning of the next
laboratory period will include two sections: (1) a physical description of all chemicals in their
present state (2) a daily description of all changes which occurred from the moment that the
procedure was completed until the next lab period.
Safety:
• Be sure to follow transportation and set-up instructions and precautions as discussed by your
instructor!
• While none of these chemicals are dangerous in the given quantities, they could cause stains if
spilled, or possible skin irritation if handled with your bare hands.
• Select a site for this experiment which is safe from your roommate and others and does not
present a potential problem for staining furniture.
Procedure:
1.
Obtain the following: 1 test tube; ParafilmTM; 1 iron nail (sanded);
copper sulfate (blue crystals); sodium chloride (white crystals).
2.
Slide copper sulfate crystals to the bottom of the test tube with a paper
funnel as demonstrated in class. The tube should be about 1/3 full.
3.
Using a pencil, push some tissue paper into the test tube so as to cover
the blue crystals.
4.
Slowly, and with as little disturbance as possible, add enough water to
just cover the blue crystals and the tissue paper. This is best
accomplished if the test tube is held at an angle under a slowly dripping
water faucet.
5.
Hold the test tube at an angle and slowly add the white crystals using a
paper funnel filling the tube another 1/3 full.
6.
Using a pencil, push some tissue paper into the test tube so as to cover
the white crystals.
7.
Add enough water (as in procedure 4) to cover the white crystals).
8.
Slide the nail into the test tube and cover completely with water.
9.
Cover the test tube with 2 layers of ParafilmTM and record your
observations as required.
IMPORTANT: Bring the test tube experiment to the laboratory at the
assigned time and place it in the designated test tube rack.
COPPER -- SILVER NITRATE REACTION
INTRODUCTION
In this experiment you will mass solid silver nitrate (AgNO3) and prepare a water solution of it. You
will also mass a piece of coiled copper wire, place it in the silver nitrate solution and observe the
reaction. By massing the copper wire at the end of the experiment, you will be able to determine the
amount of copper reacted. Using these and other measurements, you will be able to determine a
quantitative relation between reactants and products. What are the quantitative relations between
reactants and products?
WEAR SAFETY GOGGLES AND DO NOT TOUCH THE SILVER NITRATE.
WASH YOUR HANDS AT THE END OF THIS LAB ACTIVITY.
PURPOSE: To determine the quantitative relation between reactants and products.
PROCEDURE
1. Obtain a 30 cm length of copper wire. Coil the copper wire by wrapping it around a pencil. Measure
and record the mass of the copper wire. Record your balance number and use the same balance for
all mass measurements.
2. Label a large test tube with your name and period number. Add about 20 ml of dH2O to the test
tube. Mass a piece of filter paper and set the balance 2.00 g heavier. The instructor will add AgNO3
to the filter paper until the balance pan moves down. Determine the mass of the AgNO3 and filter
paper to the nearest 0.01g. Add the AgNO3 to the test tube stir to dissolve. CAUTION: Silver
nitrate will stain your skin and clothing, so be sure not to get any solid or solution on them.
3. Place the test tube containing the silver nitrate solution in a test tube rack. Add the copper wire to the
silver nitrate soltuion in the test tube, leaving enough copper wire above the solution to make it easy
to remove the copper from the test tube.
4. Record your observations of the reaction in the observation table. Allow the reaction to continue
overnight.
5. Label a clean, dry 150 ml beaker with your name and period number. Measure the mass of the
beaker.
6. Remove the copper wire and shake the crystals from the wire into the 150 ml beaker. Rinse the
copper wire with distilled water from a wash bottle into the 150 ml beaker. Allow the copper to dry
and mass it.
7. Empty the contents of the test tube into the 150 ml beaker. Allow the silver crystals in the beaker to
settle. Decant the liquid. Repeat this process 3 to 4 times.
8. Allow the silver to dry overnight in the assigned place. Measure the mass of the silver and beaker.
OBSERVATION TABLE: Draw a 15 cm x 15 cm square. Divide the square in half with a vertical
line. Label one side OBSERVATIONS OF REACTIONS and the other side EXPLANATIONS OF
REACTIONS. Complete the table.
DATA TABLE
Mass of Cu wire before reaction
Mass of filter paper
Mass of filter paper plus AgNO3
Mass of 150 ml beaker
Mass of Cu wire after reaction
Mass of beaker plus Ag crystals
ANALYSIS TABLE
Mass of Cu wire reacted (lost)
Mols of Cu reacted
Mass of AgNO3
Mols of AgNO3
Mass of Ag formed
Mols of Ag formed
Mols Ag/mols Cu
Ag atoms
Cu atoms
Atoms Ag/atoms Cu
Expected Result: mols Ag/ mols Cu
Percent Error Ag/Cu
Mols Ag/mols AgNO3
Expected Result: mols Ag/ AgNO3
Percent Error Ag/AgNO3
One mole of Cu plus _____ mole (s) AgNO3, in solution produces
_____ mole(s) Ag plus one mole of Cu(NO3)2 in solution.
Qualitative Analysis
Purpose:
In this experiment, you will devise procedures that will allow you to separate mixtures of
Ag+, Al3+, Ba2+, Ca2+, Fe3+, and Pb2+ ions and to identify each ion after the separation. You will
use your procedures to analyze an unknown for the presence of these ions. Each student will
receive an individual unknown and it may contain from two to five ions. Unknowns will be
selected completely at random.
Background
Qualitative analysis involves the identification of the contents of a mixture. When
chemical methods are used, the substances (in this case aqueous solutions of metal ions) are
usually separated before identification can be made. After they have been separated,
identification is made on the observation of a characteristic chemical reaction.
This experiment deals with the separation and identification of the six ions mentioned
above. Relative solubility, reactions affected by pH environment, and complex ion formations
are some of the concepts used in the process of separation and identification.
Complete descriptions of all observations are helpful. In completing the unknown,
chemical equations for all reactions should be written in AP format.
Dividing the Ions into General Classes:
The approach to separating and identifying a mixture of six cations will begin by dividing
the ions into three general classes. These classes will be based on the solubilities of the chlorides,
hydroxides, and carbonates of the cations. After the ions are separated into general classes, you
will separate them within the classes. After a solution contains only one ion, it is subjected to a
confirmatory test for the presence of that ion.
Prior to coming to the lab, create a table of relative solubilities of the compounds of the
cations with the anion of each reagent. In addition to the anions used to separate the cations into
general classes (Cl-, OH-, and CO32_), the reagents in this experiment will contain nitrate, acetate,
sulfate, thiocyanate, chromate, and oxalate ions. To create this table you might consult a table of
solubility rules, a Ksp table, or the Table of Physical Constants for Inorganic Compounds in the
CRC Handbook of Chemistry and Physics. In addition to the solubility rules there is another
important one: if a sparingly solubility soluble substance contains an anion that is the conjugate
base of a weak acid, the solubility of that substance will be affected by pH. Since the anion will
react with the H+ ion from an acid, the solubility usually increases in the presence of an acid. On
the other hand, if the anion is derived from a strong acid, the solubility will not usually be
affected by pH.
Other important equilibria in this experiment are:
Ag+(aq) + 2NH3(aq)  Ag(NH3)2+(aq)
Al(OH)3(s) + OH-(aq)  Al(OH)4-(aq)
Pb(OH)2(s) + 2 OH-(aq)  Pb(OH)42-(aq)
NH4+(aq)  NH3(aq) + H+(aq)
Note that Al(OH)3 and Pb(OH)2 are amphoteric substances. The tendency for the formation of
Pb(OH)42- is so great that the sparingly soluble substance, PbCrO4 will dissolve in a solution
containing NaOH.
A WARNING: Be very careful with some of the solutions used in this experiment. Nitric
acid, hydrochloric acid, sodium hydroxide, acetic acid, and ammonia can cause chemical burns
in addition to ruining your clothes. Silver nitrate produces a black stain on clothing and skin.
Procedure
Set up a water bath before you start. If you do not know how to operate the centrifuge,
get instruction before using it. It is very wise to carefully label all test tubes throughout the
experiment.
Tests to Establish the Ions in Each General Class
To 0.5 mL (10 drops) of each known solution add 4.5 mL distilled water. To each, add 5
drops of 6M HCl. Mix the contents thoroughly.
To fresh 0.5 mL samples (diluted with 4.5 ml distilled water) of all known solutions, add
10 drops of buffered ammonia solution (a mixture of 6M NH3 and 6 M NH4NO3) and mix. Since
the buffer produces a controlled amount of hydroxide ions, sparingly soluble hydroxides can
precipitate in this step. Soluble complexes of the types M(NH3)nz+ can also form.
To fresh 0.5 mL samples (diluted with 4.5 ml distilled water) of all known solutions, add
10 drops of 3M (NH4)2CO3 and mix. If there is a precipitate in any of the test tubes, save it for
the next step. A precipitate that forms in this step can be either a carbonate or a hydroxide (due to
the hydrolysis of CO32- ions). Dissociation of NH4+ ions produces ammonia that can form
complex ions of the types described in the preceding step.
To each precipitate formed from the addition of (NH4)2CO3, centrifuge each test
tube for about 1 minute. Decant the solution and discard it. Wash each precipitate by
adding about 20 drops of distilled water, stirring until the precipitate is suspended in the
water and centrifuge again. Each precipitate should be washed three times, discarding
wash water each time. It is essential that all of the ammonium carbonate is removed or it
will provide a false positive for the presence of CO32-.
Treat each washed precipitate with a few drops of 6M HNO3. If a gas is formed
(CO2) the precipitate was a carbonate, if no gas was formed, the precipitate was a
hydroxide.
Complete Flowchart I, copied into your lab notebook, prior to starting your unknown.
The formula of each chemical species added and each chemical species formed must be
presented in the appropriate spot. As part of the report, write net ionic equations for each
reaction that occurred. With each equation, provide a cryptic description of your observations.
If no reaction occurred, no equation can be written.
From the information here, devise a scheme to separate the cations into the general
classes, assuming a solution contains cations from all three classes.
Examining the Chloride Class
0.5 mL samples (diluted with 4.5 ml distilled water) of the known solutions containing
the cations in the chloride class will be sufficient to complete this part of the experiment.
Add ~1 mL of 6M HCl to each test tube. Centrifuge and test for complete precipitation
by adding one more drop of the HCl to the solution. If the precipitation is not complete add
several drops of the HCl, centrifuge, and test for complete precipitation again.
After the precipitation is complete, decant the solutions and add about 4 mL of distilled
water to each precipitate. Place the test tubes in a hot (almost boiling) water bath. Precipitates
must be stirred with a clean stirring rod.
To any test tubes in which the precipitate dissolved in hot water, add 3 drops of 1M
K2CrO4. A yellow precipitate that forms and then dissolves when 6M NaOH is added confirms
the presence of the cation.
Decant the water from the precipitates that did not dissolve. Add 2 mL of 6M NH3 (this is
NOT the buffered ammonia solution). Stir vigorously. The presence of the cation is confirmed by
the disappearance of the white precipitate. The presence of the cation can be confirmed by the
reappearance of a white precipitate when 2-3 mL of 6M HNO3 is added to the solution. The
precipitate that disappears is the soluble ammonia complex and the precipitate that reappears is
the original chloride.
Examining the Hydroxide Class
Again, 0.5 mL samples (diluted with 4.5 ml distilled water) of known cations of the
hydroxide class [that have NOT been tested in the chloride class] are sufficient. Add drops of
the buffered ammonia solution to the known solutions until a drop on the tip of a stirring rod
turns red litmus blue. Then add 2mL more of the buffered ammonia solution. Centrifuge and test
for complete precipitation.
After the precipitations are complete, decant the solutions. Add a sufficient quantity of
6M NaOH to each precipitate and stir. To each test tube in which the precipitate dissolved, add
drops of 6M HNO3 until a drop turns blue litmus paper red. Add 2 drops of 1% aluminon reagent
and drops of 6M NH3 until the solution is basic. A flocculent red precipitate, which is a
hydroxide that is stained with aluminon reagent, confirms the presence of the cation.
Centrifuge the test tubes containing precipitates that did not dissolve in sodium
hydroxide. Wash each precipitate with 2 mL of distilled water. Centrifuge and discard the wash
water. Dissolve the precipitate with a few drops of 6M HNO3. Add 2 mL of distilled water and 3
drops of 0.1M KSCN (potassium thiocyanate). A red color confirms the presence of the cation.
The red color is a thiocyanate complex of the cation.
Examining the Carbonate Class
To 0.5 mL samples (diluted with 4.5 ml distilled water) containing the cations of the
carbonate class [that have NOT been tested in the preceding classes], add drops of 3M
(NH4)2CO3 solution until each solution turns red litmus blue. Then add 1 mL more of the
ammonium carbonate. Centrifuge and test for complete precipitation. Decant the solutions.
Add sufficient 6M acetic acid to dissolve each precipitate. Then add 5 drops of 3M
NH4C2H3O2 (ammonium acetate) to buffer each solution. Add 5 drops of 1M K2CrO4. Centrifuge
and test for complete precipitation.
Wash the precipitate with 2mL of distilled water and centrifuge. Dissolve the precipitate
by adding 3 drops of 6M HCl. Add 2 mL of 0.1M Na2SO4. A precipitate should form. Centrifuge
and decant the solution, wash the precipitate with 2 mL of distilled water. A white precipitate of
the sulfate confirms the presence of the cation.
Add 1 mL of 1M K2C2O4 (potassium oxalate) to the solutions that did not precipitate with
potassium chromate. Add drops of 6M ammonia until the solution turns pink litmus paper blue.
A white precipitate confirms the presence of the cation. The precipitate is the oxalate of the
cation.
Complete Flowchart II-IV, copied into your lab notebook, prior to starting your
unknown. The formula of each chemical species added and each chemical species formed must
be presented in the appropriate spot. [This is your “ticket” into lab] For the report, write net
ionic equations for reactions that each metal ion goes through from the beginning of the process
through confirmation, in sequence. With each equation, provide a cryptic description of your
observations. There will be several equations repeated from the “separation of classes”
equations. If no reaction occurred, no equation can be written.
If you wish to practice your schemes or if you run into trouble during analysis of the
unknown, control mixtures may be made of the solutions containing the cations.
Unknown
Obtain an unknown and proceed to separate and identify which of the six cations are
present. Make sure you recognize which steps are used to separate the classes and which steps
separate the ions within each class. You may confer with the instructor about the contents of the
unknown. Be open minded. The unknown can contain from two to five cations and it may be
colored to cover the presence of cations that form colored solutions.
Flowchart I: General Classes
All 6 cations
|
HCl: |
|
precipitate
___________
___________
___________
___________
___________
___________
|||
All 6 cations
|
|
|
buffered NH3: |
|
|
|
|
|
|
|
|
|
precipitate
||| All 6 cations
___________
___________
___________
(NH4)2CO3:
___________
___________
___________
|
|
|
|
|
|
|
|
|
|
|
|
precipitate
then add HNO3
____________  __________
____________
 __________
____________
 __________
____________
 __________
____________
 __________
____________
 __________
Flowchart II: Chloride Class
Cations
|
| HCl
|
precipitate
|
|
solution
|
|
cations of other
classes
___________
|
|
| Hot H2O
|
|
precipitate
|
|
solution
|
|
____________
|
|
| NH3
|
____________
|
|
| K2CrO4
|
____________
____________
|
|
| HNO3
|
_____________
|
|
| NaOH
|
____________
Flowchart III: Hydroxide Class
precipitate
|
|
|
precipitate
|
|
|
Cations
|
|
| buffered NH3
|
|
solution
|
|
cations of other
classes
_____________
|
|
| Excess NaOH
|
solution
|
|
|
____________
|
|
| HNO3
_____________
|
|
| HNO3
____________
|
|
| KSCN
|
_____________
|
|
_____________
_____________
| Aluminon, NH3
|
Flowchart IV: Carbonate Class
Cations
|
|
| (NH4)2CO3
|
precipitate
|
|
|
|
solution
|
|
|
|
cations of other
classes
_____________
|
|
| acetic acid
_____________
|
|
| K2CrO4
|
precipitate
|
|
___________
|
|
| HCl
|
___________
|
|
| Na2SO4
|
____________
solution
|
|
___________
|
|
|K2C2O4, NH3
|
____________
Identification of Cations and Anions by Qualitative Analysis
Purpose:
To identify mixtures of six cations and six anions using a series of chemical tests and
qualitative observations.
Introduction:
Qualitative analysis has long been a fundamental practice in research chemistry. Entire
books have been written to detail the various experiments and tests that can be used to identify
the presence of certain cations, anions, and even types of organic molecules. In this experiment,
you will be given the opportunity to develop your own procedures to identify cations and anions
in a series of aqueous mixtures.
This experiment will be conducted over a period of two days. On the first day, you will
be given a series of ionic solutions that you will use to develop your identification procedure.
On the second day, you will be given six test tubes, each containing one cation and one anion.
Your task will be to identify which cation and anion are present in each tube. Some of the tests
that you will use to identify the various anions are described below, others you will have to
develop on your own through experimentation. The following is a list of the ions you will be
working with:
Cations: Na+, K+, NH4+, Ba2+, H+, and Fe2+.
Anions: OH–, Cl–, SO42–, I–, NO3–, and CO32–.
In addition to solutions containing these ions, you will also have access to pH paper,
concentrated sulfuric acid, conc. FeSO4(aq) and 1% hydrogen peroxide (H2O2) in water.
The following is a suggested list of tests. Which tests you and your labmates choose to
use is up to you. Not all the tests are needed in order to identify the ions.

Brown-ring test: This test is used to identify the presence of nitrate, NO3–. Add a small
amount of the test solution to ~1 mL of conc. iron(II) sulfate in a small test tube. Then,
using a transfer pipet, slowly add concentrated sulfuric acid to the test tube. The sulfuric
acid will form a second layer (it is far more dense than the FeSO4 solution), and at the
interface of the two layers, the appearance of a brown ring signifies the presence of NO3–.
The brown ring is actually trapped NO(g), which is produced through an oxidationreduction reaction with the Fe2+:
NO3–(aq) + 4 H+(aq) + 3 Fe2+(aq)  NO(g) + 3 Fe3+(aq) + 2 H2O(l)
The solution should also turn slightly yellow, as Fe3+ complexes with water to form
yellow Fe(OH2)63+.

pH Test: You will have access to pH paper to test the pH of the various solutions.

Flame Test: When metal ions are heated in a flame, they give off a characteristic color due to
the excitement of certain electron transitions. By dipping a nichrome wire into a solution
and placing it in a Bunsen burner flame, you can observe the colors.

Ammonium Test: One can test for the presence of ammonia by adding sodium hydroxide to the
solution in question. The hydroxide will pull a hydrogen ion off of the ammonium to form
ammonia:
OH–(aq) + NH4+(aq)  NH3(g) + H2O(l).
The formation of ammonia can either be detected by its pungent smell, or by holding a
piece of damp acidic (red) pH paper above the solution.

Solubility Tests: One of the best ways to identify the presence of certain cations and
anions is to look for the formation of precipitates as certain combinations of cation and
anion are mixed. There are two ways for you to determine which tests to perform. One
is to randomly test all combinations of cations and anions, the other is to use a solubility
chart from either your book or the web to narrow down the tests to specific combinations
of ions that might assist in their identification.

Iodine Test: In acidic solution, hydrogen peroxide can oxidize I– to form I2, which then
reacts with another I– to form I3–, which has a yellow-brown color:
2 I–(aq) + H2O2(aq) + 2 H+(aq)  I2(aq) + 2 H2O(l)
I–(aq) + I2(aq) 
I3–(aq)
This test can be performed by adding a some HCl to the solution in question and then
adding some H2O2 dropwise with a transfer pipet.
You are by no means limited to using the tests described above to identify your solutions. If you
know of a technique (other than cheating) for determining the identity of one or more of the ions
using the reagents given, feel free to use it.
Prelab:
1. Before you come to the lab on the first day, you should get together with your labmates and
develop a strategy for the tests you will do on the first day. Also, you will need to look up
the solubilities of various combinations of the cations and anions to determine which
combinations might be useful in identification. Remember, nitrate, ammonium, and
hydrogen ions will not form precipitates with anything.
2. Before coming to lab on the second day, you should prepare a procedure for identifying the
cations and anions based on the first day’s tests.
Materials:
well plate
test tubes
transfer pipets
pH paper
niochrome wire
Bunsen burner
barium iodide solution
iron(II) sulfate solution
dilute nitric acid solution
sodium carbonate solution
potassium hydroxide solution
ammonium chloride solution
1% hydrogen peroxide solution
concentrated sulfuric acid solution
6 unknown solutions (day 2 only)
concentrated iron(II) sulfate solution
Procedure:
In this experiment, you will be using a variety of not-so-pleasant chemicals and an open
flame. The utmost care should be taken at all times, and gloves, goggles and aprons should be
worn throughout the two days. Waste should always be disposed of in the beakers labeled
“Aqueous Waste” and “Solid Waste” in the case of the pH paper. On the first day, you will be
allowed to converse openly with the other groups about your tests and your ideas, but on the
second day, all conversations must cease and individual lab groups must work alone.
The rules for this lab are simple. On the first day, you will have access to all of the
chemicals listed in the materials except for the unknown solutions. It is entirely possible to come
into class on the first day already knowing how to identify all of the cations and anions.
Nevertheless, you are strongly urged to go through the entire procedure on the first day just so
you can recognize what you are looking for. Practice working with small volumes of material,
because on the second day, there will be no refills.
On the second day, you will have six test tubes labeled 1 through 6, each containing an
unknown cation and anion. The cations and anions will be the same as they were the first day,
accept not in the same pairings. Every cation will be paired with a different anion, and you can
use this fact to help with your identification. In fact, for some cations, this very piece of
information drastically limits the possible anions they can be paired with (hint, hint!). In
addition to the six unknown solutions, you will have access to Litmus paper (blue and red),
concentrated sulfuric acid, concentrated iron(II) sulfate, and a 1% hydrogen peroxide solution.
At the end of the second day, you will need to submit a sheet with your lab group number
and the identity of all of the ions in the six solutions.
Assignment:
There is no assignment for this lab. All you need to do is turn in a piece of paper that
lists the cation and anion for each of the six unknown solutions.
Post-lab:
The following questions should be answered in your lab notebook:
1. Hydrogen ions will never register any color in a flame test. Why is this?
2. If you look at the list of cations and anions that form precipitates, two striking features are
clear. First, it is very rare to find a precipitate formed by a +1 cation and a –1 anion (Ag+ and
Hg22+ being the exceptions). Second, as you move up the periodic table in a group (say the
alkaline earth metals, Ba through Mg) the precipitates they form become less and less
soluble. What is the explanation for this?
3. Even though sulfate is the conjugate base of a weak acid (HSO4–), it does not register as basic
in a pH test. There are two very good explanations for this (one mathematical, the other
practical). What are both explanations?
Determination of the charge of an unknown anion
Purpose:
To determine the charge of an anion by monitoring the mass of solid produced in a
precipitation reaction with Ca2+ as a function of the moles of anion and Ca2+ used.
Introduction:
Our current understanding of atoms and molecules is due in large part to John Dalton’s
famous work entitled A New System of Chemical Philosophy, in which he set forth the four
principles that made up his “atomic” theory of matter. The basis for Dalton’s atomic theory was
three fundamental laws that had been developed through careful experimentation and
observation. These laws can be summarized as follows:

Law of Conservation of Mass: This law states that mass cannot be created or destroyed. In
modern times, this law has been amended to state that mass cannot be created or
destroyed in a chemical reaction, because we know now that it can be created or
destroyed in a nuclear reaction. This law was first published by Antoine Lavoisier (the
father of modern chemistry) in his 1789 work Elementary Treatise on Chemistry.

Law of Definite Proportions: This law, originally called Proust’s law after its discoverer,
John Proust, states that a compound will always have the same proportion of elements by
weight. For example, regardless of its state or quantity, CO2 is always 27.3% carbon and
72.7% oxygen.

Law of Multiple Proportions: This law (formulated by Dalton) states that when two
elements form a series of different compounds (like C and O form CO and CO2), the
masses of the second element (oxygen) that react with a set quantity of the first element
(carbon) can be reduced to whole number ratios. For instance, suppose 3 grams of carbon
and 4 grams of oxygen react to form CO and 3 grams of carbon and 8 grams of oxygen
react to form CO2. The ratio of the masses of oxygen in the two compounds is 2:1.
It is this final law of chemistry that forms the basis for stoichiometry as we know it today. What
Dalton had discovered with his law of multiple proportions was that molecules are made up of
discrete atoms, and so the whole number mass ratios that are observed in different binary
compounds made of the same elements simply reflect the whole number atom ratios that exist in
the molecules. The extra 4 grams of oxygen needed to make CO2 instead of CO corresponds to
the additional 0.25 moles of oxygen atoms needed to make 0.25 moles of CO2 from 0.25 moles
of CO. This mole/mass relationship between products and reactants is what makes the study of
stoichiometry possible.
In this experiment, you will be using this mole/mass relationship to determine the charge
of an unknown anion. This will be accomplished by determining the formula of the ionic
compound formed by Ca2+ and the unknown anion. If one knows the charges of the ions that
make up an ionic compound, it is relatively simple to determine its formula. For instance, Al 3+
and O2– combine to form a compound with the formula Al2O3. We know this because the
simplest way to balance out the +3 and –2 charges is to find their least common multiple, which
is 6. Two Al3+ ions (totaling 6+) exactly balance three O2– ions (totaling 6–) giving Al2O3. This
process can also be reversed to determine the charge of an ion. For instance, manganese and
oxygen form the compound MnO2. Since we know that oxygen almost always forms O2– ions,
the charge on the manganese can be determined by balancing it with the charge from the two
oxide ions:
x  2(2)  0; x  4 ,
The cation therefore must be Mn4+. In this experiment you will use a technique call Job’s method,
named for its creator Peter Job, to determine the formula of the compound formed by Ca2+ and the
unknown anion. You will use this formula to determine the charge of the anion.
Job’s method involves mixing two reactants of equal concentrations in a series of
different quantities and plotting the amount of product that results as a function of the amounts of
reactants used. The ratio of reactants that results in the most product must reflect the correct
ratio of the reactants in the final product. For instance, consider the reaction of silver nitrate,
AgNO3(aq), with sodium chromate, Na2CrO4(aq), to form silver chromate, Ag2CrO4(s):
2 Ag+(aq) + CrO42–(aq)  Ag2CrO4(s).
Suppose we were to different volumes of 1.0 M AgNO3 with different volumes of 1.0 M
Na2CrO4. Using stoichiometry, it is fairly simple to calculate the mass of Ag2CrO4(s) produced
for each combination:
mL of 1.0 M Ag+
mL of 1.0 M CrO42–
mg of Ag2CrO4(s)
1.0
6.0
166
2.0
5.0
332
3.0
4.0
498
4.0
3.0
664
5.0
2.0
664
6.0
1.0
332
If we plot the mass of Ag2CrO4(s) as a function of
volumes of reactants used, we get what is known
as a Job-plot. The line that connects the six
points forms a “V”, where the highest point
(which was extrapolated for this example)
represents the point at which the most precipitate
would form. This is the point at which neither
reactant is limiting, so the ratio of the reactants
must reflect the ratios of the ions in the
compound. You will notice that at the maximum,
4.67 mL of 1.0 M Ag+ and 2.33 mL of 1.0 M
CrO42– would have been added. This gives an
Ag+ to CrO42– ratio of 4.67 to 2.33 or 2 to1,
corresponding to the compound Ag2CrO4. In this lab, by determining a Job-plot for the reaction
of Ca2+ with the unknown anion, you will be able to determine the ratio in which the two ions
mix to form the precipitate and hence the charge of the anion.
Pre-lab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table that will hold all of the data that you will be collecting. You won’t be able
to effectively remove the precipitate from the filter paper, so be sure to leave a column to
record the mass of the filter paper.
Materials:
0.50 M CaCl2 solution
0.50 M solution of the unknown
5 test tubes
filter paper
Büchner funnel
2 beakers (100 mL)
5 glass stirring rods
ice
vacuum flask
2 graduated pipets (10 mL)
ethanol
flat spatula
test-tube rack
10 mL graduated cylinder
1 L beaker
Procedure:
1. Fill the 1L beaker about half-full with ice and water and place one of the distilled water
bottles in it.
2. Using labeling tape and a pen, label the two beakers as “Ca2+” and “anion” and fill them with
~30-40 mL of the stock solutions found in the hood. Also, label the five test tubes “1”
through “5”.
3. Using one of the graduated pipets, transfer 1.0 mL of 0.50 M CaCl2 solution into the first test
tube, followed by 2.0 mL in the second test tube, up to 5.0 mL in the fifth test tube.
4. Using the second graduated pipet, transfer 5.0 mL of the unknown anion solution into the
first test tube, followed by 4.0 mL into the second test tube, down to 1.0 mL in the fifth test
tube. Use the glass rods to stir each of the five solutions.
5. Label the edge of a piece of filter paper with the test tube number and record
its mass. Assemble the vacuum filtration setup as shown in the diagram to the
right and place the filter paper into the funnel. Pour the contents of the first
test tube into the funnel. Rinse any remaining solid out of the test tube into
the Büchner funnel using the chilled distilled water. Finally, rinse the solid
with ~5 mL of ethanol. This will help facilitate drying.
6. Carefully remove the filter paper using the flat spatula and set it on a paper towel to dry
overnight. Place your name and the name of your labmates on the paper towel.
7. Repeat steps 4 and 5 for the remaining four test tubes. Allow the five precipitates to dry
overnight.
8. Once the precipitates have dried sufficiently, record the mass of the precipitates and their
filter papers. If possible, use the same balance as before.
9. Once all of the precipitates have been weighed, dispose of them (along with the filter paper)
in the beaker labeled “Solid Waste”.
Assignment:
For your report you must:
1. Create a job plot analogous to the one on the bottom of page 6.
2. Determine the charge of the anion using the job plot.
3. Determine (however you wish) the maximum mass of precipitate.
4. Using the mass of Ca2+ added at the maximum point, the ratio of Ca2+ to anion, and the
maximum mass of the precipitate, determine the molar mass of the anion.
5. Write a conclusion as if writing a formal lab report. A guide to writing lab reports along
with a grading rubric can be found in the “General Documents” folder.
Post-lab:
The following questions should be answered in your lab notebook:
1. Suppose you were to be told that the unknown anion was an oxoanion. Given the charge and
molar mass that you calculated, what is the identity of the anion?
2. One major source of error in this experiment could be the incomplete drying of the
precipitate. How might this affect your determination of the anion’s charge (if at all)?
3. Although it is stated in this lab that the unknown anion could have a –1, –2, or –3 charge, the
lab would not work using Ca2+ if the anion had a –1 charge. Using basic solubility rules
found in your textbook, state why this is the case.
Determination of the formula of an iron chloride using stoichiometry
Purpose:
To determine the formula of an iron chloride produced through an electrochemical reaction
with copper(II) chloride.
Introduction:
One of the most important classes of reactions in chemistry and biology are
electrochemical reactions, also called oxidation-reduction or redox reactions. These reactions
involve the exchange of electrons from one atom to another, thus altering their “oxidation states”
(charges). An atom or ion is said to be oxidized if it loses electrons. For instance, Fe (s) can be
oxidized by O2(g) to form iron(III) oxide, Fe2O3(s), also known as rust:
2 Fe(s) + 3 O2(g)  2 Fe2O3(s).
In this case, each iron atom loses three electrons to become a Fe3+ ion:
Fe 
Fe3+ + 3 e–
Reduction is the gaining of electrons. In the above reaction, the oxygen atoms in O2(g) go from
being neutral (no charge) to having a –2 charge, i.e. they each gain two electrons and are
therefore reduced:
O2
+ 4 e– 
2 O2–
Note that reduction corresponds to the reduction of the charge—in this case a reduction from 0 to
–2. Despite the name, not all oxidization-reduction reactions require oxygen. Indeed, there are
many metal ions that could stand in for O2(g) in the role of electron acceptor. For instance, Au3+
ions want very much to be reduced to Au(s), which is why gold is so expensive (because it is hard
to oxidize or corrode). If mixed with an aqueous solution of AuCl3, Fe(s) would be oxidized to
Fe3+ and the Au3+ would be reduced to Au(s).
In the equation given above, Fe(s) was oxidized to its +3 oxidation state by the presence
of oxygen. However, like most transition metals, iron can adopt more than one oxidation state.
For iron, the +2 oxidation state is also very common. Indeed, hemoglobin in your blood uses
Fe2+ ions to bind to oxygen for transport through your bloodstream. It is not always obvious
which oxidation state a metal ion adopts when it is oxidized in an electrochemical reaction. For
instance, the reaction of Fe(s) with Au3+ could take one of two forms:
Fe(s) + Au3+(aq) 
3 Fe(s) + 2 Au3+(aq)

Fe3+(aq) + Au(s)
three electrons exchanged
3 Fe2+(aq) + 2 Au(s)
six electrons exchanged
In this lab exercise, you will be given a sample of iron in the form of two iron nails. By
reacting the nails with a sample of copper(II) chloride and measuring both the mass of iron
consumed and the mass of copper produced, you will be able to determine the Cu:Fe mole ratio,
and hence the oxidation state of the iron resulting from the reaction.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table that will hold all of the mass data that you will be collecting (mass of the
nails before and after, mass of CuCl2, etc.).
Materials:
anhydrous CuCl2
iron nails
1 M HCl
2 beakers (250 mL)
100 mL graduated cylinder
vacuum flask
tweezers
watch glass
drying oven
distilled water
Büchner funnel
filter paper
sandpaper
glass rod
ring stand with clamps
Procedure:
1. Obtain two iron nails and clean them with the sandpaper until they are clean and shiny. Weigh
the nails together and record their mass to at least two decimal places.
2. Weigh out approximately 8.0 grams of copper(II) chloride and place it in the 250 mL beaker.
Be sure to record the mass to at least two decimal places. Add ~50 mL of distilled water to the
beaker and swirl until all of the CuCl2 is dissolved.
3. Place the two nails into the beaker so that they are completely covered by the aqueous CuCl2
solution. Let the nails sit there undisturbed for approximately 30 minutes.
4. After the 30 minutes is complete, remove the nails from the solution using the tweezers. While
holding the nails above the solution, squirt the nails with distilled water to dislodge any copper
metal that is stuck to the nails. If needed, you can use the glass rod to help dislodge stuck
copper from the nails.
5. Set the cleaned nails on a paper towel and allow them to dry.
6. Using the glass rod, decant the solution off of the solid copper in
first beaker into the second 250 mL beaker. If any Cu(s) does get
the second beaker, pour everything back and try again. You need
remove all of the liquid, just as much as you can.
the
into
not
The following steps will remove any aqueous ions that might be
sticking to the copper.
7. Add ~25 mL of distilled water to the copper and swirl it around in the beaker. Decant off the
added water and repeat this step two more times.
8. Add 25 mL of 1.0 M HCl to the copper sample, swirl, and decant. Then, rinse the copper
sample with an additional 25 mL of distilled water but do not decant this time.
9. Set up a vacuum filtration system and label a piece of filter paper with your initials and place it
in the Büchner funnel. Pour the copper/water mixture into the funnel and allow the vacuum to
suck the water off for ~5 minute. You can use distilled water to rinse any residual copper from
the beaker into the funnel. Place the filter paper on a watch glass and give it to your instructor.
The copper will be dried overnight in an oven, after which, you will measure and record the
mass of the copper.
Assignment:
Before leaving class on the second day you must:
1. Determine the Cu:Fe mole ratio and submit it to your teacher.
For your report you must:
1. Download the class data from Blackboard and determine an average Cu:Fe ration with a
standard deviation.
2. Determine the oxidation state of the iron ions produced by the reaction and compare it to
what is expected using standard reduction potentials.
3. Write a results and discussion section as if writing a formal lab report.
Post-lab:
The following questions should be answered in your lab notebook:
1. In this lab, you went through a lot of trouble to ensure that spurious ions were not sticking to
your copper sample. If you had not done this and the copper mass was artificially high, how
would it have affected your results?
2. Suppose you did not successfully scrape all of the copper off the nails. How would this have
affected your results?
3. The reaction time of 30 minutes is admittedly an arbitrary choice of times and could have just
as easily been much longer or much shorter. Which would be more beneficial to your
accuracy and precision, a longer reaction time or a shorter reaction time? Why?
Measurements: Precision and Accuracy
Purpose:
The purpose of this experiment is to become familiar with measurement equipment
commonly found in the chemistry laboratory, to learn the concepts of accuracy, precision and significant
figures, and to learn the degree of precision and accuracy of these measurement devices.
Apparatus:
-
top loading balance
500 ml (or greater) beaker
100 ml beakers
200 ml (or greater) beaker [with volume marking lines]
thermometer
100 ml graduated cylinder
glass pipets to deliver 1, 5 and 10 ml
deionized or distilled water
a weight from a certified, calibrated set
Definitions:
Accuracy – how close a measurement is to the known value of some accepted standard; e.g. you
weighed a metal object which had been certified by NIST (National Institute of Science and Technology)
to have a mass of 10.000 grams, and your result was 9.999 grams. Your measurement was 9.999/10.0000
or 99.99% accurate.
x
more accurate
x
less accurate
Precision – how close measurements are to one another, if measured multiple times; e.g. you
weighed the same metal object three times and got results of 9.999, 9.997, and 10.001 grams – the
average or mean was 9.999 (99.99% accurate), but the measurement ranged from –0.002 to +0.002 of the
average.
x
x
x x
x
x
x x more precise
x
x less precise
A measure of precision is a mathematical term called the standard deviation and it is defined as follows:
Standard deviation = (sum(xave-xi)2/(n-1))1/2
(Equation 1)
Where
and
xave is the average of the measured values
xi is an individual measurement
n is the number of measurements made
For the example above, the standard deviation equals:
[((9.999-9.997)2+(9.999-9.999)2+(9.999-10.001)2)/(3-1)]1/2 =
[((0.002)2+(0.000)2+(-0.002)2)/(2)] 1/2 =
[(0.00004+0.00004)/2] 1/2 = [(0.00008)/2] 1/2 = [0.00004] ½ = +/-0.002
Usually a result of several measurements is expressed as:
Average of measurements +/- standard deviation
e.g. 9.999+/-0.002
Often it is useful to express the precision of a measurement in percentage terms; i.e. what percent of the
average is variable or less precise. This percent precision term is called the relative standard deviation.
Relative standard deviation = standard deviation/average x 100
(Equation 2)
e.g. relative standard deviation = 0.002/9.999 x 100 = 0.02%
Significant figures – the value of a measurement wherein the last digit reflects where uncertainty
lay; e.g. for the 9.999 gram weighing the third number after the decimal point is the last significant figure.
This number has 4 significant figures. It makes no sense to write a weight of 9.99900000 because the
measurement is meaningless beyond the last 9.
Procedure:
1) Obtain a calibrated weight from your teacher and record its calibration value (mass) in grams on
your worksheet. Note: avoid handling the weight with your fingers (use gloves, paper towels or
tongs) because natural oils from the skin will add weight to this standard. Also store it on clean
paper to avoid having the weight contaminated by other materials that could wet or stick to the
weight.
2) Each student should “zero” the balance with the weighing pan empty, then transfer the weight to
the pan and record the balance reading on the worksheet. Remove the weight, make sure the
balance is zeroed, and add the weight back again. This should be repeated a total of four times,
and each time a reading should be recorded. Calculate the average, standard deviation, relative
standard deviation, and percent accuracy for these four mass measurements.
3) Add lab deionized water to a 500 ml (or greater) beaker to make it about 80% full and insert a
thermometer. Record the temperature of the water before you transfer it as given in the steps 4, 6,
and 7 below.
4) Take the 200 ml (or greater) beaker with volume marking lines, find its mass and record on the
worksheet. Then carefully add deionized water to the beaker from the larger beaker until the
meniscus just touches the line for 50 ml. Find the mass of the beaker again. Discard the water
and add fresh water once again to the 50 ml line and weigh. Do this a total of four times. Use
paper towels to blot any water on the outside of the beaker and above the 50 ml line. Subtract the
mass of the empty beaker from each of the four weighings of beaker plus water to get the mass of
the four samples of water. Calculate the average, standard deviation, relative standard deviation,
and percent accuracy for these four measurements.
5) The density of water is the mass of water for a given volume or
density = mass (g)/ vol (ml)
(Equation 3)
Density, however, will change as a function of temperature, and so it is important to know the
temperature along with the mass and volume. Following is a table of densities for water at
different temperatures.
o
C 0.0
0.1
20 0.998203 0.998183
21 0.997992 0.997970
22 0.997770 0.997747
23 0.997538 0.997514
24 0.997296 0.997271
25 0.997044 0.997018
26 0.996783 0.996756
27 0.996512 0.996485
Density of Water as a Function of Temperature
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
0.998162
0.998141
0.998120
0.998099
0.998078
0.998056
0.998035
0.99813
0.997948
0.997926
0.997904
0.997882
0.997860
0.997837
0.997815
0.997792
0.997724
0.997701
0.997678
0.997655
0.997632
0.997608
0.997585
0.997561
0.997490
0.997466
0.997442
0.997418
0.997394
0.997369
0.997345
0.997320
0.997246
0.997221
0.997196
0.997171
0.997146
0.997120
0.997095
0.997069
0.996992
0.996967
0.996941
0.996914
0.996888
0.996862
0.996836
0.996809
0.996729
0.996703
0.996676
0.996649
0.996621
0.996594
0.996567
0.996540
0.996457
0.996429
0.996401
0.996373
0.996345
0.996317
0.996289
0.996261
Densities (g/ml)
Now in step 4 above, you found the mass of some water and you obtained the temperature of that
water. Knowing the temperature you can find the density of water from the table above. Once
you know the density and the mass of water you can calculate the volume of water you added to
the beaker. Nominally it should be 50 ml. So how accurate was your delivery of 50 ml to the
beaker? How precise?
6)
Now take an empty 100 ml graduated cylinder and find its mass. Add water to the 25 ml
graduated line on the cylinder and find the mass. Empty and repeat three more times. Subtract
the mass of the empty cylinder from each water weighing to calculate the mass of water for all
four measurements. Determine the density of the water used and compute the water volume
from each weighing. Calculate the average water volume, the accuracy of measuring out 25 ml,
the standard deviation, and the relative standard deviation.
7)
Using a 10 ml glass pipet, transfer 10 ml to a clean, dry 100 ml beaker that has been weighed.
Repeat 3 more times. Then use a 5 ml glass pipet to transfer 5 ml and later 1 ml glass pipet to
transfer 1 ml, obtaining four measurements for each volume. Calculate the volume of water
transferred for each weighing, the accuracy of volume delivered, the standard deviation, and the
relative standard deviation.
Determination of the percentage of water in copper(II) sulfate
Purpose:
To determine the percentage of water present in CuSO4∙xH2O through thermal
dehydration and gravimetric analysis.
Introduction:
Gravimetric analysis is a technique by which the composition of a compound can be
studied through measurements of mass. Although this may sound a bit simplistic, the
information that you can obtain from simple gravimetric analysis is not. Indeed, the law of
definite proportions, the law of multiple proportions, and the law of conservation of mass are all
based on experiments performed using gravimetric analysis.
There are two methods of gravimetric analysis that are commonly employed. The first is
the precipitation method, which typically involves the analysis of an ionic compound by forming
an insoluble precipitate with one of the ions from the compound. For instance, if you wanted to
determine the percentage of chloride in a sample of magnesium chloride, you could dissolve the
sample in water, and precipitate the chloride using silver nitrate, AgNO3:
Ag+(aq) + Cl–(aq)  AgCl(s)
By weighing the dried sample of AgCl(s), it is possible to determine the mass of chloride present
in the precipitate:
mass of Cl  mass of AgCl 
1mole AgCl 1mole Cl
35.5 g Cl
.


143 g AgCl 1mole AgCl 1mole Cl
The percentage of Cl present in the original sample is then calculated as follows:
% Cl 
mass of Cl
100% .
mass of MgCl 2
The second common method of gravimetric analysis is the volatilization method, and
involves monitoring the change in mass of a substance as it is heated. For instance, if one were
to heat calcium carbonate for an extended period of time, CO2(g) would be evolved leaving
calcium oxide behind:
CaCO3(s) 
CaO(s) + CO2(g)
Whatever mass is lost by the solid must be the mass of the CO2(g), and so it is possible to
determine the percentage of calcium, carbon, and oxygen in the sample in a manner analogous to
the calculations shown above.
In this lab exercise you will be using the volatilization method of gravimetric analysis to
determine the percent by mass of water in a sample of hydrated copper(II) sulfate. It is not
uncommon for ionic crystals to contain high water content. In fact, ionic crystals obtained from
an aqueous solution always have some water molecules present in the crystal (although
sometimes in very small numbers). These water molecules help to stabilize and define the
crystal structure of the compound. By heating a hydrated ionic compound, it is possible to drive
the water molecules from the ionic crystal and form an anhydrous (water free) ionic salt:
CuSO4∙xH2O(s) 
CuSO4(s) + x H2O(g)
The percentage of water present in the compound can then be determined by measuring the
amount of mass lost when the water molecules were driven off. Over time, most anhydrous ionic
solids will rehydrate by capturing water molecules from the atmosphere (especially right after
heating). It is therefore very important to minimize exposure of anhydrous compounds to humid
air.
Pre-lab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table to hold all of the mass data that you will be collecting. You will probably
only have time to dehydrate one sample, so you need not leave space for multiple trials.
3. Using the internet, look up and print out a copy of the MSDS (material safety data sheet) for
“copper(II) sulfate” and list at least two items from the “Health Hazard Data” section. Make
note of the actual number of waters found in hydrated copper(II) sulfate.
Materials:
crucible with cover
Bunsen burner
analytical balance
crucible tongs
dessicator with fresh dessicant
hydrated copper(II) sulfate
ring stand
clay triangle
Procedure:
1. Assemble the ring stand, Bunsen burner, clay triangle, and
crucible as shown in the illustration to the right. Make sure that
crucible is clean and has no visible cracks. The crucible cover
should always be slightly ajar, leaving a space through which
can escape. It is a good idea to loop and tie a piece of wire through
crucible cover so that it can be more easily lifted with the tongs.
the
gases
the
2. Heat the crucible for ~3 minutes and place it in the dessicator
using
the crucible tongs. After the crucible has cooled for ~5 minutes,
weigh the crucible (with the cover) to three decimal places using
an
analytical balance. (The heating is meant to drive off any water adsorbed to the ceramic
crucible.)
3. Place approximately 1.0 gram of CuSO4∙xH2O into the crucible and record the mass of the
crucible and its contents to three decimal places. Make sure that you use the same analytical
balance for all your measurements to minimize the possibility of error.
4. Gently heat the crucible and its contents using the Bunsen burner. Be sure to move the
burner around to avoid forming hot spots. The CuSO4∙xH2O should turn from a bluish color
in its hydrated state to a whitish/gray color. If the CuSO4∙xH2O turns black, you’ve
overheated the sample. At too high a temperature, anhydrous CuSO4 will decompose to form
copper(II) oxide and poisonous sulfur trioxide gas:
CuSO4(s) 
CuO(s) + SO3(g).
5. After 2-4 minutes, the dehydration should be complete. You can tell when the dehydration is
complete by the lack of any blue (hydrated) crystals. Place the crucible into the dessicator
using the crucible tongs and allow the crucible to cool for ~5 minutes. Once cooling is
complete, weigh the crucible and its contents and record the mass.
6. Heat the crucible and its contents for an additional 2-3 minutes and let the sample cool in the
dessicator. Once cool, weigh the crucible and its contents and record the mass. Repeat this
procedure (called “drying to a constant weight”) until you either get a consistent mass or the
lab period ends.
7. Once finished, dispose of the anhydrous copper(II) sulfate waste in the beaker labeled “Solid
Waste”.
Assignment:
Before leaving lab you must:
1. Calculate the experimental percent mass of water in copper(II) sulfate using your data.
2. Calculate the theoretical percent mass of water in copper(II) sulfate using the formula you
found for the prelab assignment.
3. Calculate the percent error for your data.
For your report you must:
1. Hand in the above calculations.
2. Write an abstract as if writing a formal lab report. A guide to writing lab reports along with
a grading rubric can be found is available in the “General Documents” folder.
Post-lab:
The following questions should be answered in your lab notebook:
1. Anhydrous ionic salts tend to absorb water from the atmosphere, especially when they are
cooling down. This is why the cooling is done in a dessicator. If the sample were not cooled
in a dessicator, how might this affect your final results?
2. As the water molecules begin to leave the heated copper(II) sulfate, the sample will begin to
liquefy a bit before the water molecules vaporize and leave behind anhydrous CuSO4. This
can sometimes lead to spattering and loss of material. How would this affect your final
results?
3. It was mentioned above that SO3(g) is poisonous. This is because a reaction occurs when
sulfur trioxide encounters the moist tissues inside the esophagus and lungs that causes tissue
damage. What do you think is produced when you react SO3(g) with water, and why is it
necessarily harmful?
Identification of two unknown metals through the production of H2
Purpose:
To determine the molecular mass of two unknown metals using both stoichiometry and gas
laws.
Introduction:
In the presence of a strong acid (H+), most metals will react to form metal ions and H2(g).
This reaction is electrochemical in nature, because it involves the exchange of electrons from the
neutral metal atoms to the positively charged hydrogen ions, thus creating positively charged
metal ions and neutral hydrogen molecules.
The stoichiometry of acid/metal reactions is linked primarily to the charge of the metal
ion. For instance, the alkali metal lithium forms a +1 charge, and so will react with a single H +
ion to form a neutral H atom (which then merges with another neutral H atom to form H2(g).
2 Li(s) + 2 H+(aq)

H2(g) + 2 Li+(aq)
Note: Cl- ions are present, but do not react.
In contrast, aluminum (a group III element) forms Al3+ when reacted with acid, and so it will
give up three electrons to three hydrogen ions:
2 Al(s) + 6 H+(aq)

3 H2(g) + 2 Al3+(aq)
In this experiment, the identities of two unknown metals, both of which only form metal
ions with a +2 charge, will be determined. The two metals will be identified based on their
molar masses, which can be calculated from the mass of the metal that you react with the acid
and from the number of moles of H2(g) produced. To determine the number of moles of H2(g)
produced, you will be collecting the H2(g) in a eudiometer and using the ideal gas law (PV =
nRT) to determine the number of moles of gas present.
From the ideal gas law it is clear that determining the moles of
present in the eudiometer requires three pieces of information: the
temperature, which you will measure with a thermometer; the volume,
which can easily be read from the eudiometer; and the pressure of the gas
inside the tube. The pressure is the one complicated value to determine.
shown in the illustration to the right, the atmospheric pressure pushes the
column up the tube. This is why the water does not run out when the
is inverted in the beaker. As H2(g) is created, it pushes down on the water
column against the pressure of the atmosphere. The difference between
and Patm is equal to the height of the water column (in mm of H2O), and
zero when the level of the water in the tube and the beaker are the same.
measure this pressure difference, you need to first determine the height of
water column, h, and then convert from mm of H2O to mm of Hg, and
to atmospheres.
H2(g)
As
water
tube
Pgas
is
To
the
then
The conversion from mm of H2O(l) to mm of Hg(l) is accomplished using the densities of
the two liquids. We know that the force exerted by the column of water on the gas in the tube is
equal to the mass of the liquid times the acceleration due to gravity (F = ma). Therefore, equal
masses of water and mercury will exert the same force. The trick is then to determine what
height of mercury has the same mass as the height of water in the tube. We will start by finding
the volume of the water:
VH 2O  hH 2O  A ,
where A is the area of the circular cross-section of the tube (the same for both water and
mercury). We next find the mass using the density:
mH 2O  VH 2O  d H 2O ,
where m is the mass and d is the density. The volume of the same mass of mercury can then be
found by dividing by the density of mercury:
VHg 
mH 2O
.
d Hg
Finally, the height of the column of mercury can be found by dividing by the area:
hHg 
VHg
A
.
Putting this all together, we get:
hHg 
1 1

 d H 2O  hH 2O  A
A d Hg
 hH 2O
d H 2O
d Hg
.
Thus, if you know the height of the water column in mm, then you can convert this to mm of Hg
by multiplying by a ratio of the two densities. Be aware that the pressure you are calculating is
Patm − Pgas, which makes sense given that the difference in pressure is zero when there is no
water column present.
Pre-lab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create data tables for each metal that will hold all the data that you will be collecting. We
will probably only have time for one run for each metal, so you need not leave space for
multiple trials.
3. Write a balanced net ionic equation for the reaction of some metal (M) with acid to form a
metal ion with a +2 charge (M2+).
Materials:
eudiometer
copper wire
rubber stopper (with hole)
ring stand with buret clamp
ruler
400 mL beaker
6 M HCl
unknown metal X
unknown metal Y
water
Procedure:
The following steps should be followed for each of the two unknown metals.
1. Obtain a piece of one of the two unknown metals and record its mass. The size of the metal
piece has been controlled so that the gas produced is within the volume of the eudiometer, so
do not use more than one piece of metal at a time.
2. Using the copper wire, securely wrap the metal piece in a small cage, as shown in the
illustration below. Then, insert the long end of the copper wire through the narrow end of
rubber stopper and bend it around the stopper so that the cage is secured.
The copper cage need not be spherical, just enough to hold the
metal strips secure for most of the reaction. Even if the metal
eventually escapes, it will still dissolve after a while. Be sure
when making the cage that it is small enough to fit within your
eudiometer!
3. Fill the 400 mL beaker ~3/4 full with distilled water.
4. Carefully add 10 to 15 mL of 6 M HCl to the eudiometer, and fill the remainder of the tube
with distilled water from the beaker. Although this is not proper procedure (acids should be
added to water, not vice versa) in this lab it is necessary, and because of the length of the
tube, fairly safe. Warning: HCl is a strong acid, and care should be taken in its handling.
Gloves should be worn at all times during this lab.
5. Place the rubber stopper with the unknown metal attached to it into the end of the collection
tube. Quickly invert the collection tube and place it into the beaker of distilled water. If any
gas is present before the reaction begins, make note of its volume.
6. Once the unknown metal has completely dissolved, record the volume of H2 gas produced.
Also, make note of the height of the water column above the water level in the beaker using
the ruler.
7. Empty the collection tube into the beaker labeled “Metal/Acid Waste” and repeat the entire
procedure for the other unknown metal.
Assignment:
For your report you must:
1. Determine the moles of H2(g) produced by each metal using the vapor pressure, the volume of
gas, the temperature, and the height of the water column.
2. Determine the molar mass of each metal.
3. Write an introduction and a results and discussion section as if writing a formal lab report.
A rough draft will be due initially, followed by a revised version.
Post-lab:
The following questions should be answered in your lab notebook:
1. Were you to forget to include the vapor pressure in your calculations, how would this have
affected your results?
2. As you will probably notice when performing these reactions, one of the metals contains
trace amounts of carbon in it. How will the presence of this impurity affect the results of
your calculations?
3. In using the ideal gas law to calculate the moles of H2(g), you make an important
assumption: that hydrogen gas can be treated as an ideal gas. Given the van der Waal
constants (a and b) for H2, is this a legitimate assumption?
Experiment 4: Identification of an unknown solid by freezing point depression
Purpose:
To determine the molecular formula of an unknown solid using freezing-point depression
and data from elemental analysis.
Introduction:
When a solute is dissolved in a solvent, even at small concentrations, there can be very
large changes in the properties of the solvent. Among the most common changes are an increase
in the boiling point, a decrease in the freezing point, and a decrease in the solvent’s vapor
pressure. These properties, collectively known as colligative properties, are independent of the
nature of the solute, but are strongly affected by the concentration of the solute particles (ions or
molecules).
For freezing-point depression and boiling-point elevation, the change in temperature is
proportional to the molality (moles of solute divided by kg of solvent) of the solution. Molality
is used because it is independent of temperature, unlike molarity, which changes as the
temperature alters the density of the solution. The equation for freezing-point depression is as
follows:
T f  m K f ,
where Tf is the change in the freezing point of the solvent, m is the molality, and Kf is the
freezing-point depression constant (9.80 °C/m for benzophenone).
Because colligative properties depend only on the concentration of the solute and not on
its identity, they are sometimes used by scientists to determine the molar mass of an unknown
substance. Indeed, before the recent advent of electrospray mass spectrometers, the molar masses
of proteins were commonly determined by measuring the osmotic pressure (another colligative
property) of a solution of the protein.
In this experiment, we will use freezing-point depression to help determine the identity of
an unknown solid. The solvent for this experiment will be benzophenone, an organic compound
used in a variety of household products. The main advantage to using benzophenone (besides the
fact that the unknown is soluble in it) is that benzophenone is a solid at room temperature
(Tf = 48.1°C), so one need only heat the solution to make it liquid, and let the air slowly cool the
solution back to the freezing point.
Benzophenone (also called diphenyl ketone) belongs to a class of
compounds known as ketones, which are characterized by the presence of
the C=O group flanked on either side by carbon atoms—in this case two
benzene rings (also referred to as phenyl groups).
By dissolving the unknown solid in liquid benzophenone and observing the concomitant
decrease in the freezing point, it should be possible to determine the concentration of the
unknown solid and hence the number of moles added.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a pair of data tables in your lab notebook: one for the freezing point data of the pure
benzophenone, and one to contain the masses of the unknown used and changes in the
freezing point. Be sure to include space for multiple trials.
3. Determine the empirical formula of the unknown using the following data showing the
percent composition by mass: 93.708% carbon and 6.292% hydrogen.
Materials:
thermometer secured in a rubber stopper
hot plate with stir bar
400 mL beaker filled with water
ring stand with two clamps
water
test tube
benzophenone
unknown solid
wire stirrer
Procedure:
Note: Benzophenone is listed as a minor skin and eye irritant, so gloves should be worn
throughout the duration of the lab period.
1. Using an analytical balance, weigh out approximately 8 grams of benzophenone and record
the mass. Transfer the benzophenone to the test tube and secure the tube in the hot-water
bath using a ring stand and one of the clamps.
2. Weight out between 0.40 and 0.50 grams of the unknown solid and set it aside until it is
needed.
3. Lower the test tube containing the benzophenone until at least half
the tube is submerged in the water bath. Once the benzophenone
completely melted, place the thermometer and the stirring wire into
tube as shown in the diagram to the right
of
has
the
4. Once the thermometer and stirrer are in place, remove the test tube
the hot water bath. This will be a bit awkward with the thermometer
place, but can be accomplished lifting the entire ring stand up and
gently setting it away from the hot water bath.
from
in
5. With the test tube removed from the water bath, begin watching the benzophenone. As
crystals begin to form, you may see the temperature jump a bit due to supercooling (look
supercooling up in your textbook). You can accelerate this increase a bit by gently stirring
the solution with the wire. Wait until the temperature has leveled off and record that
temperature as the freezing point.
6. Once the benzophenone has solidified, return the test tube to the hot water bath so that the
benzophenone returns to the liquid state. The ring stirrer can be used to facilitate the melting
process, but be careful not to splash the benzophenone up to the top of the test tube. Once
the benzophenone has melted, add about half of the unknown solid to the test tube and stir
until the solid completely dissolves.
7. Repeat steps 4 and 5 for the solution of benzophenone and the unknown.
8. Once the benzophenone has solidified, return the test tube to the hot water bath so that the
benzophenone returns to the liquid state. Weigh the remaining unknown and add it to the test
tube. Repeat the cooling process and record the new freezing point.
9. The thermometer and test tube can be cleaned using acetone (located in the colored squirt
bottles). All waste should be placed in the beaker labeled “Benzophenone Waste”.
Warning: acetone is a highly-flammable liquid, and should not be used near the hot plates.
Assignment:
For your report you must:
1. Determine the molalities of each of the two benzophenone/unknown solutions using the
freezing-point depression.
2. Determine the molar mass of the unknown.
3. Determine the molecular formula of the unknown using the molar mass and the empirical
formula you calculated in the prelab.
4. Write an experimental section as if writing a formal lab report. Here are some tips on
writing a good experimental section:

Remember that a good experimental section provides enough information for the reader
to duplicate the experiment, but should not be so detailed as to describe things that are
obvious (like how to use a balance).

Avoid using first person and do not treat the experimental section like a procedure. If
you used 8.05 grams of benzophenone, do not say that you used “approximately 8
grams of benzophenone.”

The experimental section should be written in paragraph form, not just as a list of steps
that you followed.
Post-lab:
The following questions should be answered in your lab notebook:
1. Suppose some of the benzophenone were to splash out of the test tube before the unknown is
added. How would this affect the molar mass that you determined?
2. Sketch rough cooling curves for benzophenone both with and without the presence of the
unknown. How could the measurement of cooling curves for your experiments decreased the
potential for errors in your results?
3. What might some of the drawbacks be to using a solvent that has a melting point below room
temperature?
Determining the order of a reaction using UV-Vis spectroscopy
Purpose:
To determine the rate constant, order, and activation energy for the oxidation of formic
acid by liquid bromine.
Introduction:
Kinetics plays an extremely important role in all aspects of chemistry. Whether it is an
organic chemist in a lab trying to synthesize the next cure-all drug or an industrial chemist trying
to determine how to maximize the efficiency of a chemical plant, the need for a clear
understanding of how a reaction occurs and at what rate is extremely important. The first step in
answering both of these questions is knowledge of the basic rate law for the chemical reaction.
Consider the following oxidation/reduction reaction between liquid bromine (dissolved in
water) and formic acid:
HCOOH(aq) + Br2(aq) 
2 Br–(aq) + 2 H+(aq) + CO2(g).
A general rate law for this reaction can be written as follows:
Rate  k  HCOOH   Br2  ,
x
y
where k is the rate constant, and x and y are the reaction orders for HCOOH(aq) and Br2(aq),
respectively. One way to determine the values of x, y, and k is to measure the initial rates of the
reaction ( [Br2 ]/ t ) for three experiments that start with different initial concentrations of
bromine and formic acid. However, such a technique is experimentally problematic, because for
anything other than a zeroth order reaction, it is difficult to measure an initial rate, as shown
below.
Because the rate (the slope of the line) is not constant for
1st and 2nd order reactions, it is difficult to determine the
initial rate. Although it is possible to make extrapolations
of the initial rate by assuming a small section of the plot
to be linear, this is often problematic because data taken
at the beginning of the reaction is not accurate due to
incomplete mixing of the reactants.
In this experiment, we will use an alternate method for determining the kinetic data for
this reaction. By using integrated rate laws (often called linearized rated laws) that relate the
concentration of a reactant to the reaction time, t, it is possible to create a series of plots that
identify both the order and rate constant for the reaction. The integrated rate laws for zeroth,
first, and second order reactions are as follows:
0th order:
[A]  kt  [A]0
1st order:
ln[A]  kt  ln[A]0
2nd order:
1
1
 kt 
[A]
[A]0
As you will notice, all of these equations are for one-reactant (A) reactions. Although it is
possible to create integrated rate laws for multi-reactant systems, they are mathematically
complicated and not easy to work with. Instead, we can make the reaction appear like a single
component reaction by employing a simple trick. If we use a concentration of HCOOH that is
much greater than that of the Br2, then only a small percentage of HCOOH would be consumed
in the reaction. If only a small amount of HCOOH is reacted, then the formic acid concentration
is approximately constant, and we can re-write the rate law for this reaction using a pseudo rate
constant, k ' , where:
k '  k  HCOOH  , and
x
Rate  k '  Br2  .
y
Thus, the concentration of HCOOH is effectively removed from our rate law and the reaction
becomes a pseudo single-component reaction.
All that remains now is to figure out how to measure the concentration of liquid bromine
in the solution as a function of time. To accomplish this, we will be using ultraviolet-visible
(UV-Vis) spectroscopy to monitor the concentration of Br2(aq). In a UV-Vis spectrometer, light
at a certain wavelength of light is shined through a solution. If the wavelength of light is
absorbed by the solute in the solution, then the amount of light that exits the solution will be less
than the amount of light entering the solution. The amount of light that is absorbed is called the
absorbance, and as demonstrated by Beer’s law is proportional to the concentration of the solute:
A = lc,
where A is the absorbance, l is the path length (how wide is the sample cell),  is the extinction
coefficient (a constant that relates how much a particular solution absorbs light at a given
wavelength) and c is the concentration. As you can see above, the absorbance is also
proportional to the length of solution that the light must pass through, which should make sense
because as the path length increases, so does the number of solute molecules that the light can
interact with.
The light source for a UV-Vis spectrometer is usually a
combination of two bulbs: a deuterium discharge
bulb for the UV region and a tungsten-halogen bulb
for the visible region. A specific wavelength of light is
selected by using a very fine diffraction grating.
The reddish-brown Br2 strongly absorbs light at 393 nm. Therefore, we will use the
absorbance of light at 393 nm to monitor the disappearance of Br2(aq) from solution as it reacts
with the formic acid. The quantity l needs to be determined for bromine at 393 nm before it
will be possible to determine the concentration. This is accomplished by measuring the
absorbance of a series of Br2(aq) solutions with known concentrations. A plot of absorbance
versus concentration, should give a straight line with a slope equal to l.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Each time you mix the Br2, water, and formic acid, you will be diluting the concentrations of
the two reactants. Calculate the concentrations of the formic acid and Br2 once they are
mixed with the water in the cuvettes (M1V1 = M2V2).
3. Calculate the diluted Br2 concentrations for the three samples used to determine the Beer’s
Law plot.
4. Create data tables to collect the absorbance data for the Beer’s law plot and the temperature,
absorbance, and time data from the three kinetics runs. There is no need to actually draw out
the table for the kinetics runs since you will not know how many data points will be
collected.
Materials:
7 cuvettes
1000 L Pipetteman
UV-Vis spectrometer
stopwatch
capillary tube
distilled water
0.50 M formic acid
0.040 M Br2
2 test tubes with rack
Procedure:
Although the Br2(aq) will absorb the vast majority of the light in this lab, some light will be
absorbed by the water and by the cuvette. To account for any effect this might have on your
results, it is important to first “blank” the spectrometer.
1. Set the Pipetteman to 500 L (0.5 mL). Place a pipet tip onto the Pipetteman and pipet three
0.5 mL aliquots of distilled water into a clean, dry cuvette.
2. Place the cuvette into the spectrometer so that the clear window lines up with the arrow on
the right of the cuvette holder. Use the spectrometer display, typically: nm  or nm 
buttons, to select the correct wavelength, 393 nm. Use the spectrometer display to clear the
absorbance reading, and wait for the spectrometer to blank itself. The spectrometer should
now read 0.00 A.
Now that the spectrometer has been blanked, the next step is to create a series of Br 2(aq)
solutions and measure their absorbance.
3. Pipet two 0.5 mL aliquots of distilled water into a clean, dry cuvette. Next, pipet 0.5 mL of
0.04 M Br2 solution into the cuvette. Place the cuvette into the spectrometer and measure the
absorbance. Record this number in your lab notebook.
4. Repeat step 3, first using 1.2 mL of distilled water and 0.3 mL of 0.04 M Br2 and then again
using 1.4 mL of distilled water and 0.1 mL of 0.04 M Br2.
Now that you have all of the data you will need for your Beer’s Law plot, you can proceed with
the kinetics experiment.
5. First, record the temperature around the spectrometer using the thermometer. Pipet one 0.8
mL of distilled water into a clean, dry cuvette followed by 0.5 mL of the 0.04 M Br2 solution.
Place the cuvette into the spectrometer. There is little point in measuring the absorbance at
this point since you will dilute the solution in a moment.
6. Using a clean pipet tip, carefully pipet 0.2 mL of formic acid solution into the cuvette and
immediately start the clock. Once the run has begun, you should stir the solution with the
capillary tube for about the first 6 seconds. Then quickly close the lid and measure the
absorbance at 10 sec intervals (starting at 10 seconds) for at least 3 min.
7. Repeat steps 5 and 6 two more times with a fresh cuvette, and fresh pipet tips.
8. Dispose of all your formic acid and Br2 waste in the beaker labeled “Formic Acid/Br2
Waste”. The cuvettes and the pipet tips can be disposed of in the garbage.
Assignment:
For your report you must:
1. Create a Beer’s Law plot of absorbance versus concentration using the bromine solutions from
step 4 of the procedure. Use this plot to determine the value of l.
2. For all three runs, determine the order of the reaction with respect to Br2 using the
concentration versus time data. Although this will take 9 plots, if you are skilled at using
Microsoft Excel, this can be done very quickly with a bit of cutting and pasting.
3. Determine both the pseudo rate constant (k’) and the real rate constant (k) for the reaction
of bromine with formic acid. Report these as an average with a standard deviation.
Post-lab:
The following questions should be answered in your lab notebook:
1. Even though the second run went much faster than the first, the rate constants for the two runs
should be very similar (although not the pseudo rate constants). Why?
2. For the rate constants that you measured to be accurate, the concentration of formic acid has
to remain constant. Given the data you collected for the depletion of Br2(aq), is it reasonable
to assume that [HCOOH] remained constant in either run?
3. How would the addition of a catalyst affect the rate of these experiments? Use a sketch of
potential energy versus reaction progress to illustrate your answer.
Determination of an equilibrium constant using spectrophotometry
Purpose:
To determine the equilibrium constant for an aqueous chemical reaction using UV-Vis
spectroscopy.
Introduction:
As you saw in the previous lab experiment, UV-Vis spectroscopy is an excellent way to
measure the concentration of a reactant or product in a chemical reaction. In this lab, we will
again employ UV-Vis spectroscopy to measure the concentration, except this time we will use
this data to determine the equilibrium constant for the reaction:
Fe(OH2)63+(aq) + SCN–(aq)  [Fe(OH2)5SCN]2+(aq) + H2O(l)
When iron(III) nitrate, Fe(NO3)3, is added to water, the solid dissociates to form Fe3+ and NO3–
ions. However, transition-metal ions like Fe3+ rarely exist as ions in water. Instead, they use
their unoccupied valence orbitals (s, p, and d) to bonds with the lone pairs from water molecules.
Each of these water-metal bonds is coordinate covalent because the lone pair supplies both
electrons. Hence these complex ions are known as coordination complexes, and the molecules
that attach to the metal ions (water in this case) are known as ligands.
Fe(NO3)3(s) + 6 H2O(l)

Fe(OH2)63+(aq) + 3 NO3–(aq)
As one might expect from VSEPR theory, this hexaaquairon(III) ion forms an octahedral
complex with the iron in the middle bonded loosely to the six oxygen atoms from the water
molecules, as seen to the below. As you may recall from hybridization, an octahedral
arrangement implies a d2sp3 hybridization. As a result of this orbital mixing, two of the dorbitals have their energy slightly raised, allowing transitions between the unaffected orbitals and
these two higher energy orbitals. The energy difference for metal ion complexes invariably falls
within the visible region, resulting in the striking colors that accompany most transition metal ion
solutions (orange in the case of Fe(OH2)63+). Obvious exceptions are transition metal ions with no
d-electrons (like Ti4+ and Sc3+) or transition metal ions with full d-orbitals (like Zn2+ and Cd2+).
When potassium thiocyanide, KSCN, is added to an aqueous solution of Fe(OH2)63+, the
thiocyanide anion also can act as a ligand, with a bond to the iron that is a bit stronger than the
Fe3+-water bond.
Fe(OH2)63+(aq) + SCN–(aq)  [Fe(OH2)5SCN]2+(aq) + H2O(l)
By changing one of the six ligands that is attached to the Fe3+ ion, the energy levels of the
d-orbitals shift slightly, altering the color of the solution and causing it to turn red. However,
because the reaction does not go to completion, the amount of [Fe(OH2)5SCN]2+(aq) formed (and
hence the intensity of the red color) depends on the initial concentrations of Fe(OH2)63+ and
SCN–, as well as the equilibrium constant Kc.
In this lab, we will use UV-Vis spectroscopy and Beer’s law to monitor the concentration
of [Fe(OH2)5SCN]2+(aq) in a variety of solutions. Based on the measured concentration and the
initial concentrations of Fe3+ and SCN–, the equilibrium constant for this reaction will be
determined. In this lab, it is not possible to use one component of the reaction to calibrate the
spectrometer since Fe(OH2)5SCN2+(aq) exists in equilibrium with Fe(OH2)63+(aq) and SCN–(aq).
One trick to get around this is to add so much Fe(OH2)63+(aq) that the equilibrium lies way to the
right—so much so that the reaction essentially goes to completion. Under those conditions, the
final concentration of Fe(OH2)5SCN2+(aq) will match the initial concentration of the limiting
reactant, SCN–(aq), so that [Fe(OH2)5SCN2+]eq = [SCN–]0.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. For each of the five solutions that you will use to find l, calculate the initial concentration of
SCN–(aq) given the dilution that is performed.
3. For each of the five solutions that you will use to determine the equilibrium constant,
calculate the initial concentration of Fe(OH 2)63+(aq) and the initial concentration of SCN –(aq)
given the dilution that will occur.
4. Create a data table to hold the 10 absorbencies you will measure in this experiment.
Materials:
2 burets
2 volumetric pipets (5 mL)
5 beakers (100 mL)
0.00200 M KSCN
0.200 M Fe(NO3)3
0.00200 M Fe(NO3)3
0.050 M HNO3
10 cuvettes
10 transfer pipets
distilled water
Procedure:
1. As was the case in the kinetics lab, before any measurement of the concentration of
[Fe(OH2)5SCN]2+(aq) can be made, the UV-Vis spectrometer must be first blanked and
calibrated. The wavelength that we will be using for this lab is 477 nm.
2. Add 0.00200 M KSCN to one buret, 0.050 M HNO3 to the other, and attach both burets to
the ring stand. Write down the concentrations of these two solutions if they differ from what
is written here.
3. Take 5 of the clean beakers and, using the burets, add the appropriate amounts of 0.00200 M
KSCN to each (see the list below). Next, using a 5 mL volumetric pipet, add 5 mL of 0.200
M Fe(NO3)3 to each beaker. Finally, add the appropriate amount of 0.050 M HNO3 to each
reaction. Note: the HNO3 is a catalysts and so it does not affect the position of the
equilibrium.
Beaker
0.00200 M KSCN
0.200 M Fe(NO3)3
0.050 M HNO3
1
5.0 mL
5.0 mL
15.0 mL
2
4.0 mL
5.0 mL
16.0 mL
3
3.0 mL
5.0 mL
17.0 mL
4
2.0 mL
5.0 mL
18.0 mL
5
1.0 mL
5.0 mL
19.0 mL
4. Using the transfer pipet, place ~1.5 mL of each solution in a clean cuvette and measure the
absorbencies of each solution.
5. Pour the waste from the five beakers and the cuvettes in one of the beakers labeled “Iron
Thiocyanide Waste”. Clean the beakers using distilled water and wipe them dry with a paper
towel.
6. Using the 5 mL volumetric pipet, transfer 5.0 mL of 0.00200 M Fe(NO3)3 into the five clean
beakers. Then, dispense the appropriate amounts of 0.00200 M KSCN and 0.050 M HNO3
(see the list below) using the two burets. Be sure to swirl each mixture to ensure proper
mixing and wait at least 15 minutes before you begin measuring the concentrations (perhaps
longer for the mixture with 1 mL of HNO3).
Beaker
0.00200 M KSCN
0.00200 M Fe(NO3)3
0.050 M HNO3
1
4.0 mL
5.0 mL
1.0 mL
2
3.5 mL
5.0 mL
1.5 mL
3
3.0 mL
5.0 mL
2.0 mL
4
2.5 mL
5.0 mL
2.5 mL
5
2.0 mL
5.0 mL
3.0 mL
7. Using the transfer pipets, place ~1.5 mL of each solution in a clean cuvette and measure the
absorbencies of each solution.
8. Dispose of your waste in the “Iron Thiocyanide Waste” beaker. Cuvettes can be thrown in the
garbage can, and the transfer pipets can be thrown in the glass disposal bin.
Assignment:
For your report you must:
1. You will first need to create a Beer’s Law plot using the solutions from step 3 in order to find
l. Then use l to determine the equilibrium concentrations of Fe(OH2)5SCN2+ measured in
the five solutions from step 6.
2. You need to calculate the equilibrium concentrations of Fe(OH2)63+(aq) and SCN–(aq). Since you
know what the concentrations were initially (prelab assignment) and you know how much of
them got used up (because you know the concentration of Fe(OH2)5SCN2+), this is pretty easy.
3. For each solution in step 6, you need to calculate an equilibrium constant. Since Kc should
be the same for these solutions you should calculate a mean and standard deviation for the
equilibrium constant. Don’t be alarmed if the deviation seems large. Remember, you have
seen equilibrium constants that range from 1050 to 10−50, so your bull’s-eye is a lot bigger
than you might think.
Post-lab:
The following questions should be answered in your lab notebook:
1. Based on the values you determined for Kc, was it reasonable to assume that the reaction
went to completion when you used 0.200 M Fe(NO3)3? Do a quick calculation for one of the
solutions to justify your answer.
2. If the reaction with the 0.200 M Fe(NO3)3 did not go to completion, how would this affect
your calculations for the equilibrium constant?
3. If the color of the solution is red, why is its point of strongest absorption at a wavelength of
477 nm, which corresponds to the green portion of the spectrum?
Determination of the concentration of acetic acid in vinegar
Purpose:
To measure the concentration of acetic acid in a sample of vinegar to make sure that it
conforms to FDA regulations.
Introduction:
Plain vinegar is a product of apple juice that is created through a two step process
involving microscopic organisms. The first step involves fermentation, in which glucose present
in the apple juice is converted into ethanol by yeast:
As glucose is oxidized into two pyruvic acid molecules, it gives up four protons to an enzyme
called NAD+ (nicotinamide adenine dinucleotide) and in the process, provides the energy to form
two ATP molecules from ADP and phosphate ions (not shown). In the absence of oxygen
(which is what makes this a fermentation process), each pyruvic acid molecule loses a CO2 to
form acetaldehyde, which then picks up the four protons from the NADH to form two ethanol
molecules. At this stage, you have what is often referred to as hard cider.
In the second step of vinegar formation, a form of aerobic bacteria called acetobacter
oxidizes the ethanol formed by the fermentation into acetic acid:
CH3CH2OH + O2
 CH3COOH + H2O.
According to the U.S. Food and Drug Administration (FDA), for a manufacturer of vinegar to
use the word “vinegar” in its labeling of the product, it must contain a minimum of 4.0 grams of
acetic acid per 100 mL of solution, which corresponds to a concentration of 0.67 M. The rest of
the solution is primarily water, with trace amounts of alcohol, phosphoric acid, sugar, and
glycerol present. In this experiment, you will test the concentration of acetic acid in a sample of
vinegar. The technique that will be used to determine the concentration of acid in the sample is
called titration.
In an acid-base titration, the concentration of
an
acid can be determined by reacting it with a strong
base
that has a known concentration (or visa versa). At
the
equivalence point of an acid-base titration, when an
equal
number of moles of acid and base are mixed, the pH
of
the solution will change drastically, as the solution
is
being flooded with excess strong base or strong acid
(depending on what is being added). This is
illustrated to the right for the case of a weak acid
titrated with a strong base. By monitoring for this
sudden fluctuation in the pH, it is possible to determine when the moles of acid and base are
equal, and hence the concentration of the acid.
There are a couple of methods that can be used to monitor the pH change in a solution.
One way is to use a pH meter to constantly measure and record the pH of the solution. Although
usually pretty accurate, titration using a pH meter is fairly tedious, and as it turns out
unnecessary. In this experiment you will use an
indicator to monitor the pH change. An indicator
is usually a weak acid or base that changes color
when the pH reaches a certain point. Different
indicators change colors at different pH values, so
you must always be careful to select an indicator
that will change colors at the desired pH. Some
examples of indicators and the pH at which they
undergo a color change are shown to the left. In
this lab, we will be using phenolphthalein as our
indicator. Phenolphthalein changes color at a pH of ~9.0, which is perfect for our experiment
since the titration of a weak acid with a strong base will always result in a slightly basic solution
at the equivalence point.
For your titration of acetic acid (a weak acid with Ka = 1.810-5), you will be using 0.500
M NaOH (a strong base). It should be noted that most strong base solutions need to be
standardized (have their concentrations determined) before use. Although KOH and NaOH are
both solids that can be weighed out on a balance and dissolved in water, the concentrations of the
resulting solutions are not that accurate because it is difficult to accurately determine the mass of
solid KOH or NaOH. This is because both of these solids are hygroscopic, which means they
readily absorb water. Therefore, the mass of the solid usually includes some water. The 0.500
M NaOH solution that is used in this lab has already been standardized by titrating it with a
solution made from solid oxalic acid, H2C2O4, which is fortunately not hygroscopic.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. For each of the three sets of three titrations, create data tables in your lab notebook to record
the data you will collect.
3. Calculate what the maximum pH of vinegar can be given the FDA’s 4 grams acetic acid per
100 mL of solution rule.
Materials:
125 mL Erlenmeyer flask
10 mL volumetric pipet
buret
100 mL graduated cylinder
2% phenolphthalein solution
0.500 M NaOH
vinegar
distilled water
Procedure:
The key to any titration experiment is repetition. The more times you measure the concentration
of the vinegar, the more accurate your answer will be (unless you happen to always overshoot
the endpoint). How many times you perform the titration is up to you, but before you leave the
lab you should have at least three titrations that are within 0.3 mL of each other.
1. Use the 10 mL volumetric pipet to precisely measure out 10.00 mL of vinegar into the 125
mL Erlenmeyer flask. Then, using the graduated cylinder, add ~30 mL of distilled water to
the flask and mix the solution by swirling. Finally, add 2 drops of phenolphthalein to the
solution.
2. Place ~10 mL of 0.500 M NaOH into the buret and swirl the liquid around in the tube so that
it comes into contact with all of the inside wall. Then, pour the base into abeaker labeled
“Titration Waste”. This rinsing step will neutralize any acid that may be lurking in the tube
and need only be done once.
3. Fill the buret with the 0.500 M NaOH solution and secure the buret in the buret clamp. Since
there is no way to know when the titration will reach the equivalence point (it could be after
1.0 mL, it could be after 100 mL), it is usually a good idea to do a test run. Let the NaOH
from the buret flow into the flask at a relatively fast pace while constantly swirling the
solution. When the solution turns pink, stop the titration and make note of how much NaOH
was added.
4. Repeat step 1 to prepare a solution for the second titration. Be sure to use vinegar from the
same bottle as before! If you need more NaOH, be sure to refill the buret. For the next
titration, you have a pretty good idea of where the endpoint is going to be, so you can quickly
add the NaOH to the acid until you get to ~3-5 mL from the endpoint. How close you get
depends on how much you trust your first titration. Once you get within 3-5 mL of the
endpoint, slow the flow of the NaOH solution so that it is flowing dropwise into the HCl
solution. You should be constantly swirling the solution as the NaOH is added to ensure that
the NaOH is totally dispersed and the color change is visible. As you approach the endpoint,
you will start to see the solution turn pink around where the base hits the solution. At this
point, you might want to slow the titration even further. The endpoint has been reached
when the pink color can no longer be eliminated by swirling the solution. It should only
take 1 drop of base for the color to persist. In other words, how pink the solution gets is
irrelevant. What is important is that the solution stays pink, even with swirling.
5. You will need to repeat steps 1 and 4 until you are satisfied with your data, you run out of
reagents, or you run out of time (whichever comes first). Be sure to clean your flasks
between runs, and dispose of all waste in the waste beakers, not in the sink!
Assignment:
For your report you must write a formal lab report. The entire point of this lab was to
determine if your vinegar sample contained enough acetic acid to be called vinegar according to
FDA standards. Use the report to explain the task, provide background information, present your
findings, and consider the results. Be scientific! That means using significant figures in your
calculations, employing statistics to analyze your data, and considering any shortcomings in the
procedure. Good luck!
Post-lab:
The following questions should be answered in your lab notebook:
1. Why is the pH at the endpoint of a weak acid/strong base titration is always greater than 7?
2. Suppose that the NaOH solution is not standardized and the solid NaOH used to make the
solution actually contained some water making the mass of the NaOH artificially high. How
would this error affect the concentration of acetic acid that you measured?
3. In step 1 of the procedure, you arbitrarily added ~30 mL of water to the 10 mL of vinegar,
thus decreasing its concentration by a factor of 4. If it was so important to accurately
measure the volume of acetic acid used, why isn’t it important to accurately measure the
volume of water measured?
Identification of an unknown acid by titration
Purpose:
To identify an known acid by measuring its acid dissociation constant (Ka) and molar mass
through a titration with NaOH.
Introduction:
Buffer solutions play an absolutely crucial role in sustaining life here on earth. The
presence of bicarbonate (HCO3–) in most lakes maintains the pH at levels that are suitable for the
life that lives within those lakes. Bicarbonate also acts in conjunction with carbonic acid
(H2CO3) to keep our blood pH at ~7.4. Maintaining this pH is absolutely essential to our
existence, because almost all bodily functions are pH sensitive. Below a pH of 7.4, a condition
known as acidosis is induced, and if a human’s blood pH drops below 6.8 (or goes higher than
7.8), the condition is fatal. What makes the presence of a buffer so essential in human blood is
not the fact that it tunes the pH to such a precise level, but rather that the buffer is capable of
maintaining that pH level, despite the fact that the body is constantly acting to lower blood pH by
the production of CO2 (which forms H2CO3 in water) and lactic acid whenever physical exertion
occurs.
A buffer solution is created whenever a weak acid and its conjugate base are present in
approximately equal amounts. Typically, a solution is considered to be buffered when the ratio of
weak acid to conjugate base is between 0.1 and 10.
10 
[acid]
 0.1
[base]
This is often referred to as the buffer region. As stated above, the main function of a buffer is to
prevent any drastic changes in the pH of a solution, either by dilution with water or by the
addition of a strong acid or base. It is relatively simple to understand why a solution that contains
equal amounts of an acid and its conjugate base would be resistant to the addition of strong acids
or bases if we look at the Henderson-Hasselbalch equation:
 [acid] 
pH  pK a  log 
.
 [base] 
Over the entire range of the buffer region, the pH of a solution will only change by 2, which is
the difference of log(10) and log(0.1). So what does this have to do with the addition of a strong
acid? Consider the ionization equation for acetic acid:
HC2H3O2(aq) + H2O(l)  H3O+(aq) + C2H3O2–(aq)
Ka = 1.810–5
Acetate, C2H3O2–(aq), is the conjugate base of acetic acid, and so a buffer containing 1.0 M
concentrations of acetate and acetic acid would have a pH of:
 [1.0 M] 
pH  pK a  log 
  pH  4.74  log 1  4.74 .
 [1.0 M] 
Now consider what would happen if a strong acid (H3O+) were added to the solution. The
reverse of the ionization equation shown above would occur:
H3O+(aq) + C2H3O2–(aq)  HC2H3O2(aq) + H2O(l)
K a1 = 5.6104
The size of the equilibrium constant clearly indicates that this reaction will go nearly to
completion, thus decreasing the amount of acetate present in the solution, and increasing the
amount of acetic acid present. Let us assume that 0.8 moles of strong acid are added to a liter of
the buffer solution. A 0.8 M H3O+ solution would normally have a pH of ~0.1. In the buffer
solution, the acid reacts with the acetate to form 0.8 moles of acetic acid, while decreasing the
amount of acetate to 0.2 moles:
H3O+(aq) + C2H3O2–(aq)  HC2H3O2(aq) + H2O(l)
Initial

Final
0.8 moles
–x
0.0 moles
1.0 moles
–x
0.2 moles
1.0 moles
+x
1.8 moles
The ratio of acid to base is now 1.8 to 0.2, so using the Henderson-Hasselbalch equation, the new
pH is:
 [1.8] 
pH  pK a  log 
  pH  4.74  0.95  3.79
 [0.2] 
So a quantity of acid that would normally reduce the pH of a solution down to 0.1, only lowers
the pH of this buffer system by ~1.0.
There are two important things to note about this calculation. First, the ability of a buffer
to withstand the addition of a strong acid is dependent on its concentration. If we had started
with a solution that was 0.1 M in both acetic acid and acetate, 0.8 moles of strong acid would
have pushed the solution way out of the buffer region, since there would be 0.7 moles of strong
acid remaining after the acetate is completely reacted. Second, it is not necessary to worry about
changes in the volume of buffer solution due to addition of the strong acid. Since the
Henderson-Hasselbalch equation is based on a ratio of weak acid to conjugate base
concentrations, an increase in the volume will not affect the pH, since the increase in volume
affects the concentrations of acid and base equally. Indeed, this is why buffers are resistant to
pH changes due to dilution. The ratio of acid to base is invariant to changes in the volume of
water, so long as the concentrations of acid and base don’t get too dilute. Eventually, the buffer
will loose its effectiveness as either the acid or base concentration begins to approach the
magnitude of either Ka or Kb.
In this experiment, you will be using the properties of a buffer to determine the identity
of an unknown weak acid. By measuring the pH of the weak acid solution as it is titrated with a
known solution of NaOH, it is possible to draw a titration curve. Among the things that can be
derived from the titration curve are the number of moles of the acid present, the pH of the
solution at the equivalence point, and most importantly, the pH at the half-equivalence point.
Since at the half-equivalence point the pH = pKa it is relatively trivial to determine the Ka of the
acid. Furthermore, since you will know the mass of acid added and the moles of acid present,
you can also calculate its molar mass.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table that will hold the pH of the acid solution as a function of the volume of
base added.
Materials:
0.1 M NaOH solution
unknown acid
150 mL beaker
pH calibration solutions
phenolphthalein
hotplate with stir bar
buret
pH meter
ring stand with clamps
100 mL graduated cylinder
Procedure:
1. The first step in this procedure will be to make the acid solution. Weigh out ~0.3 grams of
the unknown acid and record the mass. Place the acid in the 150 mL beaker and add ~30 mL
of distilled water and a few drops of phenolphthalein. Place the stir bar in the beaker and let
the solution stir until all of the solid is dissolved. While you are making your solution, your
instructor will come around to each lab group and show you how to calibrate the pH meter
using the calibration solutions.
2. Take the distilled water squirt bottle and rinse the
of the pH probe. Submerge it into the buffer solution
secure the probe with a clamp as shown in the figure to
right.
bottom
and
the
3. Take the buret and rinse it with ~10 mL of 0.1 M
Then, fill the buret with the NaOH solution and secure
the ring stand with a buret clamp. Begin stirring the
solution, but only gently so that a vortex does not form.
the initial volume in the buret and the initial pH.
NaOH.
it
to
Record
4. Begin adding the 0.1 M NaOH in ~1.0 mL increments.
After
each increment, record the pH of the solution and the volume of NaOH added. If you want to
use a smaller increment (giving a smoother curve and more accurate results) you are free to do
so.
5. When the acidic equivalence point is reached (the pH should start increasing rapidly), stop
the titration and record the final pH. If you wish to go a little past the equivalence point, that
is fine too. Once your acid titration is complete, pour any excess NaOH into the beaker
labeled “Acid/Base Waste” and place the used buret in the dirty glassware bin.
Calculations:
There are two separate tasks that you need to complete for this lab. First, you need to
determine the Ka of the acid. This can be accomplished by plotting the pH as a function of
volume of NaOH added and finding the pH at the half-equivalence point. The second thing you
need to do is determine the molar mass of the acid. Since this is a monoprotic acid, you can find
the moles of acid used by determining the moles of base added at the equivalence point.
Questions:
The following questions should be answered in your lab notebook:
1. Most ionic compounds are hygroscopic, that is they absorb water readily. If the unknown
acid were left out in a humid room, how would the absorption of water by the acid affect
your determination of the Ka? The molar mass?
2. In addition to the pH meter, phenolphthalein is used in this lab as an indicator because it has
a color change at around pH = 9, which is around the equivalence point for this acid. How
would your results be affected if you had used an indicator like bromthymol blue (color
change at pH = 7) to determine the endpoint?
3. The pKa of carbonic acid is 6.4, which is in the fatal range for blood pH. Given that the pH
of human blood is normally around 7.4, what is the ratio of [H2CO3] to [HCO3–] in human
blood?
Determination of the caloric content of various nuts
Purpose:
To accurately measure the caloric content of peanuts, cashews, and macadamia nuts by
measuring their enthalpies of combustion in a bomb calorimeter.
Introduction:

The enthalpy of combustion, H comb
, is a measure of the energy released or absorbed
when a certain amount of a substance is combusted in oxygen. Depending on its composition, a
variety of gaseous products can be formed by the complete combustion of a substance—nitrogen
is converted into NO2(g), sulfur is converted into SO2(g), carbon is converted into CO2(g),
hydrogen is converted into H2O(g), and any metals present are converted into metal oxides.
Because of the inherent stability of the bonds in all of these gases, the heat released from
combustion reactions is usually quite substantial, and given that the heated products will expand
quite substantially, the containing such a combustion reaction can often be tricky.
In this experiment, we will be determining the caloric content of three different nuts by
measuring the enthalpy of combustion for each nut. Most nuts have a similar composition by
mass. They are ~55% fat (usually in the form of oil), ~20% protein, ~20% carbohydrates (both
simple sugars and starches) and ~5% water. As is evident
Peanut oil, shown to the left, has the formula
C57H104O6. It can be classified a triester because of the
three groups where a COO group is flanked on either
side by carbon atoms. These groups are called ester
groups, and are common in many organic molecules.
A fiber of polyester is comprised of a series of
molecules linked together by ester groups.
from the nutritional label of any nut, the vast majority of its Calories (usually ~75%) come from
the fatty oils present in the nut. Fat molecules typically consist of very long hydrocarbon chains.
Saturated fats consist of hydrocarbon chains in which all or almost all of the carbon atoms are
completely saturated with hydrogens, i.e. no carbon-carbon double bonds are present.
Unsaturated fats usually have double or even triple bonds connecting some of the carbon atoms.
For instance, peanut oil, with the formula C57H104O6, is an unsaturated fat, as can be seen from its
structure shown above.
To determine the enthalpy of combustion for any substance, it is necessary to monitor
how much heat flows out of the system (the combusting nut in this case) and into the
surroundings. The study of heat flow is called calorimetry, and the device used to measure the
flow of heat is called a calorimeter. The theory behind the workings of a calorimeter is relatively
simple. According to the 1st law of thermodynamics, the energy of the universe is constant,
Euniv  0 . Therefore the amount of energy that flows out of a system, Esys , must be equal to
the energy flowing into the surroundings, Esurr , since
Euniv  Esys  Esurr  0
Since it is impossible to measure the amount of energy flowing out of a system and into all of the
universe except the system, it is advantageous to create a little “mini-universe” that is thermally
isolated from the rest of the universe (usually with insulation). This is the basic function of a
calorimeter, which usually consists of a reaction chamber that is in thermal contact with a bath
containing water or some other substance with a known specific heat, with the bath insulated
from the outside world. As the reaction or process being studied proceeds, heat is transferred
either to or from the bath depending on whether the reaction is exothermic or endothermic.
In this lab you will be using what is known as a bomb calorimeter. Bomb calorimeters
are specially designed for studying combustion reactions because they permit the use of high
O2(g) pressures and they are capable of containing the resulting explosion as the gases produced
by the combustion reaction expand. Because of the complex design of a high-precision bomb
calorimeter, there are many things that need to be accounted for when calculating the enthalpy of
combustion for a substance. There are two sources of heat in a typical bomb calorimeter—the
tungsten ignition wire will combust to form tungsten(IV) oxide and the substance that you are
combusting. Determining the energy released by the combustion of the tungsten wire is quite
trivial—2.3 calories of energy are released per cm of wire that is combusted. The enthalpy of
combustion of the substance is the unknown in this experiment, and can be determined by
measuring how much heat is absorbed by the calorimeter and the bath.
For this experiment, the bath will consist of 2.000 kg of water. However, this is not the
only thing that will be absorbing heat as the reaction proceeds. The large stainless-steel bomb
will also absorb some of the heat, as will the stainless-steel bucket that holds the water. You
could account for these by determining their combined mass and looking up the heat capacity of
stainless steel, but a much more accurate way of measuring the amount of heat absorbed by the
calorimeter is to measure the heat capacity of the entire calorimeter. This can be accomplished
by performing a test run of the calorimeter using a substance whose enthalpy of combustion is
well documented. For our experiment, this substance will be benzoic acid (C7H6O2).
If we assume that the calorimeter is fairly well insulated and the pressure remains
constant (so no work is being done), it is possible to equate the exo- and endothermic processes:
qendothermic  qexothermic ,
so
qcalorimeter  qwater  [qwire  qrxn ] .
On the endothermic side, the heat flowing into the calorimeter is
qcalorimeter  ccalorimeter T ,
where ccalorimeter is the calorimeter heat capacity, and the heat flowing into the water is
qwater  mcwater T ,
where cwater is the specific heat of water. On the exothermic side, qwire is given above, and the

heat flowing out of the reaction is equal to the enthalpy of combustion, H comb
.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
Materials:
bomb calorimeter
benzoic acid pellet
calcium chloride
peanuts, macadamia nuts, or cashews
Procedure:
A. Determining the calorimeter heat capacity
1. Begin by placing the top of the bomb on the ring stand and
inserting the sample cup into the slot. Measure out a 10.0 cm
of ignition wire and set it aside.
2. Obtain a benzoic acid pellet and measure its mass. Use the
analytical balance in room 2013… you will need the
significant figures. Place the pellet in the sample cup and
attach the ignition wire as shown in the diagram to the right.
might want to bend the wire into a V-shape before attaching it
the bomb, and make sure that the bottom of the wire contacts
benzoic acid pellet.
piece
You
to
the
3. Gently insert the top of the bomb into the bomb shell and screw on the flange that holds the
bomb together. Then take the bomb over to the oxygen cylinder where your instructor will
fill the bomb with 40 atm of O2(g).
4. Using the large scale, measure out exactly 2000 grams of cold water into the metal bucket
and put the bucket into the insulated shell of the calorimeter. The pegs in the bottom of the
shell should fit into the indentations in the bottom of the bucket.
5. Using the tongs, carefully lower the bomb onto the round indentation in the bucket. Without
putting your hands into the water, attach the wire leads to the bomb. If you see any bubbles
flowing from the bomb, get your instructor!
6. Place the cover on the calorimeter and attach the stirring wheel to the drive wheel with the
rubber o-ring. Turn the motor on and insert the thermometer. When the temperature
equilibrates, record the initial temperature of the water. To help read the thermometer, a
magnifying glass that attaches to the thermometer is included.
7. Attach the leads leaving the calorimeter to the ignition box, one to the “common” port and
the other to the “10 cm” port. When you are ready, press the ignition button and hold it for 5
seconds. You should see the red light come on and go off (although hold the switch for all 5
seconds. Within a short period of time, you should see the temperature begin to rise. When
it stops rising, record the final temperature.
The thermometer you will use for this lab is far
more precise than any you have used in the past.
With a range of ~10ºC, the major tick marks come
at 1º increments, the minor tick marks give the
temp-erature to 0.1º, and the smaller tick marks
come at intervals of 0.02º. If you include another
uncertain digit, you can measure T to 4 significant
digits!
8. To disassemble the bomb, remove the thermometer and put it back in its protective sheath.
Remove the lid, disconnect the leads from the bomb and remove the bomb using the tongs.
BEFORE OPENING THE BOMB, OPEN THE GAS RELEASE VALVE AND LET
THE PRESSURE EQUILIBRATE!! Once the excess gas has vented, unscrew the cap and
remove the bomb lid. There should be nothing left in the bomb except for some condensed
water (from the combustion reaction) that you need to dry out. Also dry off the lid of the
bomb and remove an excess ignition wire still attached to the lid.
B. Determining the heat capacity of a nut
9. Your group will be assigned a nut—either a peanut, a cashew, or a macadamia nut.
Following the same procedure as you did for the benzoic acid, measure the enthalpy of
combustion of the nut. You will need to refill the bucket with cold water since (1) the
temperature went up a bit and (2) you undoubtedly lost water removing the bomb.
10. When you have finished combusting your nut, you may do a second trial with the same nut if
you have time. Otherwise, cleanup and dry off all the equipment. There is no waste for this
lab… combustion is complete!
Assignment:
Before leaving lab you must:
1. Calculate the specific heat capacity of the calorimeter (including the water).
2. Calculate the heat evolved from the nut you used.
For your report you must:
1. Hand in the above calculations.
2. Write an abstract and introduction as if writing a formal lab report.
Post-lab:
The following questions should be answered in your lab notebook:
1. If you failed to account for the heat produced by the combusting ignition wire, how would
the enthalpy of combustion be affected?
2. If the value you measured for the caloric content of the nut deviates from the value reported
on the jar, what might be the cause of the deviation?
Identification of cations and anions by qualitative analysis
Purpose:
To identify mixtures of six cations and six anions using a series of chemical tests and
qualitative observations.
Introduction:
Qualitative analysis has long been a fundamental practice in research chemistry. Entire
books have been written to detail the various experiments and tests that can be used to identify
the presence of certain cations, anions, and even types of organic molecules. In this experiment,
you will be given the opportunity to develop your own procedures to identify cations and anions
in a series of aqueous mixtures.
This experiment will be conducted over a period of two days. On the first day, you will
be given a series of ionic solutions that you will use to develop your identification procedure.
On the second day, you will be given six test tubes, each containing one cation and one anion.
Your task will be to identify which cation and anion are present in each tube. Some of the tests
that you will use to identify the various anions are described below, others you will have to
develop on your own through experimentation. The following is a list of the ions you will be
working with:
Cations: Na+, K+, NH4+, Ba2+, H+, and Fe2+.
Anions: OH–, Cl–, SO42–, I–, NO3–, and CO32–.
In addition to solutions containing these ions, you will also have access to pH paper,
concentrated sulfuric acid, conc. FeSO4(aq) and 1% hydrogen peroxide (H2O2) in water.
The following is a suggested list of tests. Which tests you and your labmates choose to
use is up to you. Not all the tests are needed in order to identify the ions.

Brown-ring test: This test is used to identify the presence of nitrate, NO3–. Add a small
amount of the test solution to ~1 mL of conc. iron(II) sulfate in a small test tube. Then,
using a transfer pipet, slowly add concentrated sulfuric acid to the test tube. The sulfuric
acid will form a second layer (it is far more dense than the FeSO4 solution), and at the
interface of the two layers, the appearance of a brown ring signifies the presence of NO3–.
The brown ring is actually trapped NO(g), which is produced through an oxidationreduction reaction with the Fe2+:
NO3–(aq) + 4 H+(aq) + 3 Fe2+(aq)  NO(g) + 3 Fe3+(aq) + 2 H2O(l)
The solution should also turn slightly yellow, as Fe3+ complexes with water to form
yellow Fe(OH2)63+.

pH Test: You will have access to pH paper to test the pH of the various solutions.

Flame Test: When metal ions are heated in a flame, they give off a characteristic color due to
the excitement of certain electron transitions. By dipping a nichrome wire into a solution
and placing it in a Bunsen burner flame, you can observe the colors.

Ammonium Test: One can test for the presence of ammonia by adding sodium hydroxide to the
solution in question. The hydroxide will pull a hydrogen ion off of the ammonium to form
ammonia:
OH–(aq) + NH4+(aq)  NH3(g) + H2O(l).
The formation of ammonia can either be detected by its pungent smell, or by holding a
piece of damp acidic (red) pH paper above the solution.

Solubility Tests: One of the best ways to identify the presence of certain cations and
anions is to look for the formation of precipitates as certain combinations of cation and
anion are mixed. There are two ways for you to determine which tests to perform. One
is to randomly test all combinations of cations and anions, the other is to use a solubility
chart from either your book or the web to narrow down the tests to specific combinations
of ions that might assist in their identification.

Iodine Test: In acidic solution, hydrogen peroxide can oxidize I– to form I2, which then
reacts with another I– to form I3–, which has a yellow-brown color:
2 I–(aq) + H2O2(aq) + 2 H+(aq)  I2(aq) + 2 H2O(l)
I–(aq) + I2(aq) 
I3–(aq)
This test can be performed by adding a some HCl to the solution in question and then
adding some H2O2 dropwise with a transfer pipet.
You are by no means limited to using the tests described above to identify your solutions. If you
know of a technique (other than cheating) for determining the identity of one or more of the ions
using the reagents given, feel free to use it.
Prelab:
1. Before you come to the lab on the first day, you should get together with your labmates and
develop a strategy for the tests you will do on the first day. Also, you will need to look up
the solubilities of various combinations of the cations and anions to determine which
combinations might be useful in identification. Remember, nitrate, ammonium, and
hydrogen ions will not form precipitates with anything.
2. Before coming to lab on the second day, you should prepare a procedure for identifying the
cations and anions based on the first day’s tests.
Materials:
well plate
test tubes
transfer pipets
pH paper
niochrome wire
Bunsen burner
barium iodide solution
iron(II) sulfate solution
dilute nitric acid solution
sodium carbonate solution
potassium hydroxide solution
ammonium chloride solution
1% hydrogen peroxide solution
concentrated sulfuric acid solution
6 unknown solutions (day 2 only)
concentrated iron(II) sulfate solution
Procedure:
In this experiment, you will be using a variety of not-so-pleasant chemicals and an open
flame. The utmost care should be taken at all times, and gloves, goggles and aprons should be
worn throughout the two days. Waste should always be disposed of in the beakers labeled
“Aqueous Waste” and “Solid Waste” in the case of the pH paper. On the first day, you will be
allowed to converse openly with the other groups about your tests and your ideas, but on the
second day, all conversations must cease and individual lab groups must work alone.
The rules for this lab are simple. On the first day, you will have access to all of the
chemicals listed in the materials except for the unknown solutions. It is entirely possible to come
into class on the first day already knowing how to identify all of the cations and anions.
Nevertheless, you are strongly urged to go through the entire procedure on the first day just so
you can recognize what you are looking for. Practice working with small volumes of material,
because on the second day, there will be no refills.
On the second day, you will have six test tubes labeled 1 through 6, each containing an
unknown cation and anion. The cations and anions will be the same as they were the first day,
accept not in the same pairings. Every cation will be paired with a different anion, and you can
use this fact to help with your identification. In fact, for some cations, this very piece of
information drastically limits the possible anions they can be paired with (hint, hint!). In
addition to the six unknown solutions, you will have access to Litmus paper (blue and red),
concentrated sulfuric acid, concentrated iron(II) sulfate, and a 1% hydrogen peroxide solution.
At the end of the second day, you will need to submit a sheet with your lab group number
and the identity of all of the ions in the six solutions.
Assignment:
There is no assignment for this lab. All you need to do is turn in a piece of paper that
lists the cation and anion for each of the six unknown solutions.
Post-lab:
The following questions should be answered in your lab notebook:
1. Hydrogen ions will never register any color in a flame test. Why is this?
2. If you look at the list of cations and anions that form precipitates, two striking features are
clear. First, it is very rare to find a precipitate formed by a +1 cation and a –1 anion (Ag+ and
Hg22+ being the exceptions). Second, as you move up the periodic table in a group (say the
alkaline earth metals, Ba through Mg) the precipitates they form become less and less
soluble. What is the explanation for this?
3. Even though sulfate is the conjugate base of a weak acid (HSO4–), it does not register as basic
in a pH test. There are two very good explanations for this (one mathematical, the other
practical). What are both explanations?
Determination of thermodynamic data from standard cell potentials
Purpose:
To determine the standard enthalpy and entropy changes for the oxidation of zinc metal
by copper(II) ions.
Introduction:
Although it is relatively simple to measure the enthalpy change of a system using
calorimetry, determining other thermodynamic data, like G° and S°, is not so trivial. The most
common method to determine G° for a reaction is to measure the equilibrium constant for that
reaction and convert it to a free energy change using the equation:
G  RT ln K ,
where R is the universal gas constant (8.31 J/mol∙K), T is the temperature, and K is the
equilibrium constant for that reaction. Then, using data for H° taken from calorimetry
experiments, it is possible to determine S° from the equation:
G  H   T S  .
In the late 1800’s, a German scientist by the name of Jacob V’ant Hoff developed a more
straightforward way of determining the enthalpy, entropy, and free energy of a reaction. By
combining the two equations above,
RT ln K  H   T S ,
and rearranging the terms a bit,
ln K  
H   1  S 
,
 
R T 
R
V’ant Hoff realized that if you plotted the natural log of the equilibrium constant as a function of
T–1, you would get a straight line. Furthermore, the slope (m) and y-intercept (b) of that line
would be functions of the enthalpy and entropy, respectively:
m
H 
R
and
b
S 
.
R
The V’ant Hoff plot shown to the left is for the dimerization
of NO2(g) to N2O4(g), with K determined at various
temperatures. The slope of the fitted line is 6980 K–1, which
when multiplied by –R, gives a H° = –5.8104 J/mol. The yintercept of the fitted line is –21.2 (unitless), which when
multiplied by R, gives a S° = –176 J/mol∙K. It should be
noted that for the Van’t Hoff equation to be valid, H° and
S° must be independent of temperature over a wide
range. Fortunately this is the case for most reactions.
V’ant Hoff’s contribution of the above equation (known as the V’ant Hoff equation) along with
his discovery of the relationship between concentration and osmotic pressure (=MRT) won him
the very first Nobel prize in chemistry, awarded in 1901.
In this experiment we will be using V’ant Hoff’s method to determine H°, S°, and G°
for the oxidation of Zn(s) by Cu2+(aq) in a Galvanic cell. However, we will not be measuring the
equilibrium constant as V’ant Hoff did. Rather, we will determine the free energy change by
measuring the cell potential of our Galvanic cell using a multimeter (which is a fancy version of
a voltmeter). The standard free-energy change for a reaction can be related to the cell potential
using the following equation:
G  nFE ,
where n is the number of electrons being exchanged (two for the oxidation of zinc), E° is the cell
potential in Volts, and F is Faraday’s constant,
F  96, 487
C
.
mol of e
Thus, we can write the equation for H° and S° in terms of E°:
nFE  H   T S  ,
and rearrange to be a linear equation of the form y = b + mx:
E  
H  S 

T,
nF
nF
where the y-intercept and slope are simply:
b
H 
nF
and
m
S 
.
nF
Therefore, to find H° and S° we need only to measure the cell voltage at a variety of
temperatures, plot the voltage as a function of temperature, and calculate S° and H° from the
slope and y-intercept, respectively.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table to record the cell voltage at a minimum of five temperatures.
3. Using a list of reduction potentials, calculate the cell potential for this
Zn|Zn2+(0.5 M)||Cu2+(0.5 M)|Cu Galvanic cell. Note: you will not need to use the Nernst
equation to calculate the potential. As long as the concentrations of both ions are the same,
the fact that they are not 1.0 M (standard conditions) will not matter for this reaction, since:
E  E 
RT  [Zn 2+ ] 
RT
ln 
 E 
ln 1  E .
2+ 
nF  [Cu ] 
nF
Materials:
large glass dish
zinc strip
copper strip
0.5 M Cu(NO3)2 solution
0.5 M Zn(NO3)2 solution
2.0 M KNO3 solution
2 thermometers
2 beakers (1 L)
sandpaper
2 beakers (250 mL)
hot water bath
multimeter
water
ice
2 glass stirring rods
salt bridge
Procedure:
1. The first step in this lab is to assemble the Galvanic cell. Fill each of the two beakers about
half way with the Zn2+ and Cu2+ solutions.
2. Using the sandpaper, thoroughly buff both sides of the two metal strips. Dip the Cu strip into
the Cu(NO3)2 solution and the Zn strip into the Zn(NO3)2 solution. Attach both electrodes to
the multimeter using the alligator clips.
The cell you assemble should look something like the one
shown to the left. You may wish to bend the metal strips
over the lips of the beakers to so that they are a bit more
secure. A rubber band can be used to further secure the
metal strips if you are having problems. While heating or
cooling the solutions, it is important to stir the two solutions
with the stirring rods provided so that the entire solution is at
the same temperature.
3. Using the thermometer, measure the temperature of one of the solutions (they should be the
same). Record the temperature and then hook insert the U-tube to complete the circuit.
Record the voltage registered by the multimeter once it settles a bit. Then, remove the Utube to break the circuit.
4. Make a pair of ice baths using the 1 L beakers, ice, and some water. Place the two half cells
into the ice baths and wait for the temperature to equilibrate (0ºC). Reinsert the U-tube and
measure the voltage. Then, remove the U-tube to break the circuit.
5. Place the two beakers in the hot water bath and raise the temperature to ~35ºC. Measure the
temperature and the voltage, and then repeat this step at successively higher temperatures
until you run out of time.
6. When you have completed measurements at a minimum of five temperatures, discard the
solutions in the 1 L beaker labeled “Cu2+/Zn2+ Waste”. Dispose of the metal strips and the
glass wool in the beaker labeled “Solid Waste”.
Assignment:
For your report you must:
1. Make a plot of voltage versus temperature in Microsoft® Excel and fit the plot using a
linear equation.
2. Determine H° and S° using the slope and intercept from the V’ant Hoff plot.
3. Compare your values for H°, S°, and G° that you calculated in the prelab. Calculate a
percent deviation for each.
4. Write a conclusion as if writing a formal lab report.
Post-lab:
The following questions should be answered in your lab notebook:
1. In the cell that you created, which beaker was the anode and which was the cathode?
2. Suppose that the Zn2+ solution was mislabeled, and the actual concentration was 1.0 M. How
would this affect your calculations of H° and S°? Is there a way that this can be accounted
for in your calculations?
3. The graphing technique that you used to find H° and S° in this lab tends to give much
more accurate numbers for S° than it does for H°, unless the experiment can be carried out
at much lower temperatures. Why is this the case? Hint: the exact opposite is true of V’ant
Hoff’s method—H° is more accurate than S° unless you measure K at high temperatures.
Identification of an unknown liquid by vapor density
Purpose:
To identify an unknown volatile liquid by determining its molar mass through
measurements of its vapor density.
Introduction:
At this point, you should understand that gravimetric analysis, a technique such as Job’s
method, can be used to determine the empirical formula of a sparingly soluble salt. Although
obtaining the empirical formula is always a good first step in identifying any compound, it is
never enough because an empirical formula only indicates the ratios of the atoms, not their
numbers in a single molecule. For instance, the empirical formula CH2 is common to an almost
infinite number of compounds called alkenes (hydrocarbon chains that possess C–C double
bonds), so determining that a compound has a hydrogen to carbon ratio of 2:1 is not sufficient to
identify the compound.
The next step in determining the identity of a compound is to determine its molecular
formula, and although there are a lot of methods for accomplishing this task, all of them basically
aim for the same piece of information—the molar mass of the compound. Since the empirical
formula gives the simplest ratio of atoms present in the molecule, all that must be determined to
learn the molecular formula is how many of those simple units make up the compound; the molar
mass provides that information. For instance, take the example of a molecule with an empirical
formula of CH2. The molar mass of this simple unit is ~14 g/mol. If it is determined that the
molar mass of the compound is 56 g/mol, then there are exactly four CH2 units present in the
compound, giving a molecular formula of C4H8.
56 g/mol
 4 CH 2 units
14 g/mol
It is important to understand that just because you know that the formula is C4H8 does not mean
that you know the structure of the compound. As can be seen in the figure below, many
compounds can have the same molecular formula, and some form of spectroscopy is usually
needed to distinguish between these different molecular structures known as isomers.
In this experiment you will use the volatization method to determine the molar mass of an
unknown gas. The volatilization method works for any liquid or solid that is easily vaporized,
which typically limits use of the method to organic compounds. The method works by utilizing
the ideal gas law, PV  nRT , to determine the molar mass of the volatized gas based on its
density at a given pressure and temperature. Since the molar mass (M) of a compound is just its
mass (m) divided by the number of moles (n), the ideal gas law can be easily rewritten in terms of
M:
m
n
PV  nRT  m   RT 
RT .
M
m
Rearranging the terms to isolate the molar mass gives:
M
mRT
.
PV
A ratio of mass per unit volume is just the density (d), and so this equation is often written as:
 m  RT dRT
.
M  

P
V  P
Thus, if we know the density, pressure, and temperature of an ideal gas, then it is relatively
trivial to determine its molar mass. In this experiment, you will be using a very common
technique for determining the vapor density of a volatile liquid in order to determine its molar
mass and hence its molecular formula.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table that will hold all the data that you will be collecting (mass of the flask and
condensed gas, mass of the flask without the condensed gas, etc.). You will be doing at least
three trials, so be sure to make space for them all in your table.
3. Using elemental analysis, it is possible to determine the elemental composition of an
unknown sample. The unknown sample in this lab is 85.62% carbon and 14.38% hydrogen.
Based on this information, determine the empirical formula of the assigned compound.
Materials:
125 bottle with septum
600 mL beaker
1000 mL beaker
thermometer
hot plate
10 mL graduated cylinder
Procedure:
ice
methylene blue
unknown liquid
syringe needles
balance
water
1. Place a small amount of the methylene blue dye (~1 scoop) in the 125 mL bottle and cap the
bottle with the septum. Carefully insert two needles through the center of the septum.
Record the mass of the bottle, septum, and needles.
2. Carefully remove the septum and place ~2 mL of the unknown liquid into the septum bottle.
Tightly re-cover the bottle with the septum.
3. Using the 600 mL beaker and hot plate, assemble a hot water bath and submerge the bottle
into the bath (it will actually float on top, but that is ok). Slowly raise the temperature of the
hot water bath to ~85°C and wait for the liquid to completely vaporize. You can see this best
by watching the behavior of the methylene blue—when you no longer see it moving, the
solution has dried out. Record the temperature of the hot water bath at the time when the
liquid finishes vaporizing.
4. Remove the bottle from the hot water bath and plunge it into the ice bath (made using the 1 L
beaker, ice, and some water). Once all of the gas has condensed into the bottle remove it
from the bath and dry the outside of the bottle. Record the mass of the bottle with the
condensed liquid inside (as well as the septum and needles).
5. You will need to repeat steps 3 and 4 at least two more times. It is not necessary to clean the
flask out in between runs, however you may want to add a bit more of the unknown volatile
liquid before each additional run. When you are finished, dispose of the remaining volatile
liquid and dye in the beaker labeled “Organic Waste”.
6. Once you are satisfied with your results, you will need to determine the volume of the flask.
Assignment:
For your report you must:
1. Determine the molar mass of the unknown liquid and report it as an average with a
standard deviation.
2. Using the empirical formula you determined in the prelab and the molar mass that you
calculated, determine the molecular formula of the unknown liquid.
3. Draw three possible isomers for the molecular formula you determined.
4. Write an abstract and conclusion as if writing a formal lab report.
Post-lab:
The following questions should be answered in your lab notebook:
1. It is usually a good idea to work in the hood when working with volatile gases like we do in
this lab. Why is this a bad idea for this particular experiment? How would the molar mass
that you determined be affected if you had done the experiment in the hood?
2. The purpose of the nonvolatile dye is to make it easier to determine when all of the liquid has
vaporized. Suppose you neglected to place the dye in the flask and not all of the liquid
evaporated. How would this affect the molar mass that you calculated?
3. The fact that the unknown liquid is a liquid at room temperature is a pretty good indication
that its vapor is not the best example of an ideal gas. However, under the conditions that this
experiment is performed, it is a fairly good approximation to say that the gas is behaving ideally.
Under what conditions do non-ideal gases behave ideally?
Determination of the Ksp for silver acetate using the Mohr method
Purpose:
To determine the solubility product constant (Ksp) for silver acetate through titrations of
potassium chloride with a variety of silver solutions.
Introduction:
In experiment #7, you determined the concentration of acetic acid present in a sample of
vinegar by titrating the vinegar with a known concentration of sodium hydroxide until the
endpoint was reached. In this experiment, we will again call on titration to determine the
concentration of an aqueous solute, but in a slightly different manner. Instead of watching for
the complete neutralization of a weak acid by a strong base, we will instead perform a
precipitation titration, in which the endpoint of the titration is achieved when the solute in
question is completely precipitated from the solution.
Predicting the solubility of a salt is not a trivial task; a variety of factors determine the
degree to which a salt will dissolve in a polar solvent (like water), such as the lattice energy of
the salt, the favorable enthalpy change associated with the creation of many ion-dipole
interactions, and the unfavorable enthalpy change that occurs when the natural interactions of the
solvent molecules are disturbed. Add to that the entropy changes associated with the dissolution
and what you have is a very complicated problem. Fortunately, determining the solubility of a
salt in a solvent is fairly trivial—add the salt to the solvent until the solution is saturated, then
determine the concentration of the solution that was created. There are a variety of ways in
which the concentration of a saturated solution can be determined, such as:

Measuring the mass of solid added and the resulting volume of the solution (not always
accurate because it is hard to hit the exact point of saturation);

Add an excess of the solid and use a conductivity probe to determine the concentrations
of the ions present; and

Create a saturated solution and measure the concentration of one of the ions using a
precipitation titration.
In this experiment, you will be using precipitation titration to determine the solubility product
constant, Ksp, for silver acetate using a specific precipitation titration technique known as the
Mohr method.
Silver acetate, AgC2H3O2, is an ionic compound that is somewhat soluble in water (much
more so than most silver salts). When placed in water, AgC2H3O2 dissociates according to the
following reaction:
AgC2H3O2(s)  Ag+(aq) + C2H3O2–(aq)
Remembering that Ksp does not depend on the concentrations of
solids or liquids, the equilibrium expression for this reaction is
simply:
Ksp = [Ag+][C2H3O2–].
Therefore, once equilibrium is achieved (also referred to as
saturation), the concentrations of Ag+(aq) and C2H3O2–(aq) do not
change, as shown to the left. As AgC2H3O2(s) is added to water,
the concentration of Ag+(aq) increases until it reaches its
saturation point, after which, the concentration of Ag+(aq)
cannot be increased regardless of the amount of solid added. If we can measure the
concentration of Ag+(aq) in a saturated solution of silver acetate, then we can use that
concentration to calculate the Ksp.
For this experiment, you will use a precipitation titration with potassium chloride to
determine the concentration of Ag+(aq) in a saturated solution of silver acetate. By titrating a
known volume of standardized KCl solution with a saturated solution of silver acetate, it will be
possible to determine how much Ag+(aq) is present in the saturated solution based on how much
of the solution it takes to completely precipitate all of the Cl– as AgCl:
AgCl(s)  Ag+(aq) + Cl–(aq)
Ksp = 1.610-10
The only real trick to the whole procedure is figuring out when the Cl– is completely precipitated.
This is actually a very famous problem, and a variety of techniques have been developed to solve
it:

The Volhard Method: Titrate Cl– with an excess amount of Ag+ solution, then titrate with
SCN– in the presence of Fe3+. The SCN– will first complex with the silver ions, forming
[Ag(SCN)2]–. Once all the silver is consumed, the endpoint will be signaled by the
formation of the red [Fe(OH2)5SCN]2+ ion.

The Fajans Method: Titrate Cl– with the saturated Ag+ solution in the presence of an
indicator called dichlorofluorescein. Once all of the AgCl(s) precipitate has formed, any
excess Ag+ binds to the indicator, which then becomes fluorescent.

The Mohr Method: Titrate Cl– with the saturated Ag+ solution in the presence of
potassium chromate, K2CrO4. Once all of the AgCl(s) precipitate has formed, any excess
Ag+ binds with the chromate ions to form a brick-red Ag2CrO4 precipitate.
Because it is the most straightforward technique, we will be using the Mohr method to determine
the Ag+(aq) concentration in our saturated silver acetate solution. The behavior of the Ag+ ions in
the presence of the two anions (Cl– and CrO42–) is worth examining. One might be tempted to
assume that Ag2CrO4(s) would precipitate before AgCl(s) because its Ksp (1.110-12) is smaller
than that of AgCl. However, a simple calculation shows us that this is not true. Suppose we had
a solution with equal concentrations of Cl– and CrO42–, both 0.10 M. The maximum
concentration of Ag+(aq) that could be present in the solution without an AgCl precipitate forming
is:
[Ag + ]=
K sp
[Cl ]

1.6 1010
 1.6 109 M .
0.10
Since the equilibrium expression for the dissolution of silver chromate is:
Ksp = [Ag+]2[CrO42–],
the maximum concentration of Ag+(aq) that could be present in the solution without an Ag2CrO4
precipitate forming is:
[Ag + ]=
K sp
2
[CrO 4 ]

1.11012
 3.3 106 M
0.10
Thus, because the concentration of Ag+(aq) is squared in the equilibrium expression, the
concentration of silver ions needed to form a silver chromate precipitate is much higher than that
needed to form a silver chloride precipitate, despite the fact that Ag2CrO4(s) has a smaller Ksp.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. For each of the three sets of two titrations, create data tables in your lab notebook to record
the volume of KCl used, the initial volume of silver acetate, the final volume of silver
acetate, and the concentration of KCl.
Materials:
0.0500 M KCl solution (standardized)
saturated silver acetate in distilled H2O
saturated silver acetate in 0.100 M KNO3
saturated silver acetate in 0.100 M NaC2H3O2
distilled water
filter paper
10 mL graduated cylinder
5% potassium chromate in H2O
3 Erlenmeyer flasks (125 mL)
10 mL volumetric pipet
25 mL buret
Büchner funnel
vacuum flask
Procedure:
In this experiment, you will be performing a total of six titrations: two titrations of KCl with
saturated silver acetate in distilled H2O, two of saturated silver acetate in 0.100 M KNO3, and
two of saturated silver acetate in 0.100 M NaC2H3O2. You will be able to use all three
experiments to determine the Ksp of silver acetate.
Part 1: Titration of KCl with saturated silver acetate in distilled H2O
1. For the first part of this lab, you will titrate 10 mL of 0.0500 M KCl solution with a saturated
solution of silver acetate. You should dilute the KCl with ~15 mL of water before titrating
and add ~1 mL of 5% K2CrO4 solution to the KCl solution as an indicator. Don’t forget to
rinse the buret out with about 10 mL of saturated silver acetate before beginning the titration.
Once you are done with two titrations, transfer all waste to the large beakers labeled
“Silver/Chromium waste”.
Part 2: Titration of KCl with saturated silver acetate in 0.100 M KNO3
2. Repeat step 1, except this time use saturated silver acetate in 0.100 M KNO3. This will test
the Ksp under the influence of the “uncommon-ion effect” since neither K+ or NO3− will
precipitate with anything present in these solutions.
Part 3: Titration of KCl with saturated silver acetate in 0.100 M NaC2H3O2
3. Repeat step 1, except this time use saturated silver acetate in 0.100 M NaC 2H3O2. This will
test the Ksp under the influence of the “common-ion effect”. Of course the common-ion
effect theoretically should not affect the Ksp since the concentration of C2H3O2− can be
accounted for in the equilibrium expression, but the presence of 0.1 M Na+ will make the
AgCl a bit more soluble (again, the uncommon-ion effect).
Assignment:
For your report you must:
1. Calculate the concentration of Ag+ in each of the six titrations using the volume of
saturated AgC2H3O2 needed to precipitate out all the Cl− in the flask.
2. Calculate the Ksp’s for each of the six titrations, and report the Ksp’s for each part as an
average of the three runs with a standard deviation.
3. In a couple of paragraphs, explain the differences in the Ksp’s for the three parts using your
knowledge of solubility equilibrium and intermolecular forces. Be sure to compare the
Ksp’s you determined to the known Ksp value for silver acetate.
Post-lab:
The following questions should be answered in your lab notebook:
1. Suppose the saturated silver acetate solutions had not been filtered before the lab and and
some residual solid silver acetate made it into your buret. How would this affect the value of
the Ksp that you calculated?
2. Why was so much more of the saturated silver acetate in 0.100 M NaC2H3O2 solution needed
to completely titrate the KCl solution?
Isolation and purification of nicotine from tobacco
Purpose:
To utilize various techniques in organic chemistry to extract and purify nicotine from
tobacco leaves.
Introduction:
In this laboratory experiment, you are going to be using a variety of techniques that are
commonly used in organic synthesis. You will begin by leeching the nicotine from the tobacco
leaves using a strongly basic solution. You will then use an organic technique known as
extraction to selectively pull the nicotine out of the basic solution and into a solution of diethyl
ether, a polar organic liquid that is immiscible in water. Because nicotine is far more soluble in
diethyl ether than it is in the strongly basic solution, by mixing the ether with the nicotine/base
solution, the nicotine should diffuse (with a bit of help) into the ether.
The purpose of the extraction step is to pull the nicotine out of the aqueous phase that
contains all sorts of compounds leeched from the tobacco (as well as all that NaOH) and put it
into a relatively pure organic liquid that can be easily removed by evaporation. After the solvent
is evaporated you will use another technique, called derivitization to isolate the nicotine as a
solid.
In its natural state, nicotine is a viscous oil that is very hard to work with. One way to
make it more manageable is to react the oil with a derivitizing agent to turn it into an ionic
compound that can be precipitated out as an ionic solid. Since nicotine has two amine groups on
it (essentially weak base functional groups), it is possible to react it with a weak acid to form an
ionic compound. The acid you will use for this is called picric acid. Although it does not look
like a typical organic acid (protons on –OH groups are rarely acidic unless they are in carboxylic
acid groups), the presence of those –NO2 groups pulls so much electron density away from the
O–H bond that it is relatively easy to remove the proton.
Once the nicotine has been derivitized to form nicotine picrateThe , the last step is a
purification step through recrystallization. For any compound to form a crystalline solid, the
solid must form a very precise lattice of molecules, each positioned to maximize the
intermolecular forces that hold the solid together. Because the arrangement of molecules has to
be so precise, impurities are naturally excluded from the formation of the crystal since they
would interrupt to formation of the lattice. Therefore, one way to purify a solid that contains
impurities is through recrystallization.
As you did with the aspirin lab, you will recrystallize the nicotine picrate by raising the
temperature of an ethanol/water solution to the point where the nicotine picrate is barely soluble,
and then slowly lower the temperature to decrease the solubility and induce crystallization. If all
goes well, the nicotine picrate should form very long needle-like crystals.
Day 1 Materials:
cigar
beaker (400 mL)
non-absorbent cotton
100 mL beaker
white paper
5% sodium hydroxide
Büchner funnel
250 mL beaker
100 mL graduated cylinder
glass rod
filter flask
Day 2 Materials:
250 mL separatory funnel
100 mL graduated cylinder
ring stand with medium ring 150 mL beaker
glass rod
cork stand for RB flask
diethyl ether
150 mL round bottom flask
250 mL Erlenmeyer flask
Day 3 Materials:
ring stand with clamp
10 mL graduated cylinder
cotton
Büchner funnel
methalolic picric acid
glass rod
100 mL beaker
1 L beaker
filter paper
600 mL beaker
funnel
methanol
filter flask
50% ethanol solution
transfer pipet
400 mL beaker
filter paper
filter flask
Day 4 Materials:
50 mL Erlenmeyer flask
10 mL graduated cylinder
Day 5 Materials:
Büchner funnel
Procedure Day 1:
1. Over a sheet of paper, cut apart one cigar and remove all the tobacco inside. Using your
hands (with gloves on), break up the tobacco into pieces and record their mass.
2. Place the shredded tobacco in a 400 mL beaker. Carefully add 100 mL of 5% sodium
hydroxide solution to the beaker.
3. Stir the tobacco/base suspension for ~15 minutes using a glass rod. While one person is
stirring, the other should set up the vacuum filtration apparatus. Instead of filter paper, use a
thin layer of non-absorbent cotton instead. Strong bases don’t pass very well through filter
paper.
4. Once stirring is complete, filter the solution through the Büchner funnel. You may wish to
use a 100 mL beaker to press the remaining liquid out of the leaves. Wash the tobacco leaves
with about 20 mL of distilled water and press it again with the beaker.
5. Dispose of the tobacco leaves and glass wool in the solid waste beaker. Transfer the brown
filtrate to a 250 mL beaker and cover the top with parafilm.
Procedure Day 2:
6. Begin the day by carefully decanting the liquid from yesterday into the 250 mL Erlenmeyer
flask. Try to leave behind as much of the solid particulate as you can.
7. Place the separatory funnel (sep funnel) into the ring on the ring stand and pour in your
solution from day 1. Add 25 mL of diethyl ether and place the stopper in the top of the flask.
8. Give the flask a gentle swirl or two, then invert the flask with your finger holding the glass
stopper and open the stopcock to relieve the resulting pressure. Close the stopcock and
gently shake the funnel.
9. Place the sep funnel back on the ring stand and wait for the layers to separate. Then remove
the glass stopper and let the aqueous layer (the bottom layer since diethyl ether is not as
dense as water) drain out the bottom into your original 250 mL flask. If you have particulate
in the bottom of the funnel, it may not drain well. Try rotating the stopcock again and again
to free the clog. You may end up having to pour out the entire mixture and clean the sep
funnel with warm water (followed by distilled water) before retrying.
10. Once you have drained off the aqueous layer, POUR the organic layer out of the top of the
sep funnel and into the 150 mL beaker. You never drain both layers out of the stopcock…
things just get messy that way.
11. Repeat the extraction two more times using two more 25 mL quantities of diethyl ether.
Each time, drain the aqueous layer into the 250 mL beaker and add the organic layer to the
150 mL beaker.
12. Using the glass rod, decant the organic layer into the 150 mL round bottom flask, being
careful to exclude any remaining aqueous layer or oil that remains. Cover the round bottom
flask with parafilm and set it aside for the next day.
Procedure Day 3:
13. Using a hot water bath and a ring stand to secure the flask, evaporate off the diethyl ether
until all that remains is an oily solid. Be sure to place a stirring rod into the round bottom
flask while you are heating the solution to prevent bumbing.
14. Remove the round bottom flask from the hot water bath and add 1.0 mL of water. Swirl the
flask to dissolve the solid residue. Then add 4.0 mL of methanol and warm the mixture on
the hot water bath for a few minutes.
15. Place a small plug of cotton into the bottom of the funnel and place the funnel into a 100
mL beaker. Filter the solution through the funnel and into the 100 mL beaker. Rinse out
the round bottom flask with ~5 mL of methanol and filter that into the 100 mL beaker.
16. Add 10 mL of methanolic picric acid to the solution. You should see a fluffy yellow
precipitate form. Pre-mass a piece of filter paper and filter the precipitate out using the
Büchner funnel and set it aside to dry.
Procedure Day 4:
17. Mass the dried precipitate and filter paper. Then transfer the precipitate to a 50 mL
Erlenmeyer flask.
18. Add ~4 mL of the 50% ethanol solution and heat the mixture on the hot water bath. Then,
using the transfer pipet, continue to add small quantities of 50% ethanol solution until the
solid eventually dissolves (you may have to wait a bit in between additions for the solution
to come back up to temperature).
19. Once all the solid is dissolved, cover the flask with a piece of paper towel and let the
crystals form overnight.
Procedure Day 5:
20. Cool the mixture from day 4 on an ice bath and filter it using a Büchner funnel onto a piece
of pre-massed filter paper. Set the filter paper aside to dry.
Procedure Day 6:
21.
Mass your product!!
Synthesis, purification, and characterization of aspirin
Purpose:
To synthesize acetylsalicylic acid from salicylic acid and acetic anhydride and to purify
and characterize the resulting product.
Introduction:
Since the time of Hippocrates in 400 B.C., man has known that salicylates like salicylic
acid relieve pain and reduce fever. However, most salicylates tend to be rather poisonous,
inducing a condition known as “salicylism” that is characterized by abdominal pain, vomiting,
increased respiration, and mental disturbances. In 1897, a German chemist by the name of Felix
Hoffman developed a derivative of salicylic acid called acetylsalicylic acid. This derivatized
salicylate was much less harmful than natural salicylate and was quickly patented by Hoffman
and his employer, Friedrich Bayer, as aspirin.
In this laboratory experiment, you will be synthesizing aspirin from salicylic acid and a
organic liquid called acetic anhydride. The products of this reaction are acetylsalicylic acid and
acetic acid, as shown below.
In addition to synthesizing the aspirin you will also be learning two common techniques in
organic chemistry: recrystallization, which is a method for purifying solids, and thin-layer
chromatography, which we will use to test the purity of the product.
Thin-layer chromatography (TLC) is a technique for measuring the purity of a sample by
separating its components based on their solubilities in a particular solvent. The primary tool for
performing TLC is the TLC plate, which consists of a plastic plate covered with a substance
called silica gel. A small amount of the substance that is being characterized is dissolved in a
volatile solvent and is spotted onto the bottom of the TLC plate. The bottom of the plate is then
submerged in some solvent and is allowed to sit for a few minutes. The solvent will begin to
creep up the TLC plate and, as it goes, it will carry the sample with it. The greater the solubility
of the solid in the solvent, the faster it will move up the plate.
Since TLC is usually used to characterize the purity of a sample, it is common practice to
simultaneously run pure examples of the starting materials and the desired product. In this way
you can compare the distances traveled by the spots and
possibly determine the identity of any impurities.
Although the distance the spots travel depends on the
time that the TLC plate spends in the solvent, the ratio
of the distance traveled by the spot and the distance
traveled by the solvent line is independent of time, and
can thus be used to identify a substance. This value is
commonly referred to as the retention factor, Rf.
The procedure for this lab will be split into two days. The first day, you will synthesize a
sample of acetylsalicylic acid and purify the sample by recrystallization. The sample will then
be dried over night so that you will be able to obtain an accurate yield. On the second day, you
will characterize the purity of your product by TLC.
Prelab:
1. In your lab notebook, write a brief outline of the procedure.
2. Create a data table to record the amounts of the reactants used, the mass of the final product,
and the percent yield.
Materials:
600 mL beaker
hot plate
2 Erlenmeyer flasks (125 mL)
10 mL graduated cylinder
salicylic acid
85% H3PO4
acetic anhydride
spatula
100 mL graduated cylinder
1000 mL beaker
distilled water
glass stirring rod
Büchner funnel
filter paper
vacuum flask
2 large test tubes
ethanol
thermometer
3 small test tubes
ruler
acetonitrile
commercial aspirin
TLC plate
eluting solution
eluting chamber
iodine crystals
developing chamber
ice
mechanical pencil
transfer pipet
Procedure: Day 1
Warning: Acetic anhydride is a rather nasty chemical (especially if you inhale it) as is
phosphoric acid, so the entire reaction must be carried out in a fume hood with the sash pulled
down as low as possible. A hood will be assigned to your lab group at the start of the lab period.
1. Prepare a hot water bath by filling a 600 mL beaker to about the 350 mL line with water. Place
the water bath on the hot plate and raise the temperature to between 70 and 80°C.
2. Weigh out approximately 3.0 grams of salicylic acid and place it in the Erlenmeyer flask. Be
sure to record the mass of the salicylic acid. Next, add approximately 6 mL of acetic anhydride
to the flask, followed by 10 drops of 85% H3PO4. Place the flask in the hot water bath and
gently swirl the flask until all of the salicylic acid dissolves.
3. Let the reaction mixture rest in the hot water bath for ~20 minutes, swirling occasionally.
While you are waiting, prepare an ice bath using the 1 L beaker, some ice, and some tap water.
4. Remove the flask from the hot water bath and slowly add 20 drops of distilled water using the
transfer pipet. Then quickly add 25 mL of distilled water to the mixture and scratch the walls
of the flask with a glass rod to induce crystallization. Then place the entire mixture in an ice
bath. Fill the 2 large test tubes with distilled water and place them in the ice bath as well.
5. Isolate the crystals of acetylsalicylic acid by pouring the contents of the flask through the
Büchner funnel. You can rinse any crystals remaining on the inside of the Erlenmeyer flask
into the funnel using the chilled distilled water in the test tubes.
Procedure: Day 2
6. Place the crude aspirin in a clean 125 mL Erlenmeyer flask and add approximately 3 mL of
ethanol. Heat the mixture in the water bath, and if the solid does not dissolve, add some more
ethanol.
7. Once the solid has dissolved, cool the solution to room temperature and slowly start adding
distilled water while swirling the solution. Once you see the solution turn cloudy, stop
adding water and swirl for about a minute. If the cloudiness disappears, add more water and
try again. When the cloudiness persists with swirling for one minute, place the flask back
into the hot water bath and swirl until the solution turns clear. Then, let the solution cool to
room temperature on the bench top and then to 0 °C in the ice bath. If crystals do not form,
they can be seeded using a tiny bit of commercial aspirin.
8.
Filter the solution using the Büchner funnel and give the product to your instructor to dry
over night.
Procedure: Day 3
1. Retrieve your aspirin sample and scrape it into a tared weigh boat using the spatula. Record
the mass of your purified product and calculate a percent yield. Percent yield is based on
the theoretical yield of aspirin, not the initial mass of salicylic acid!
2. In each of the three small test tubes place 1 mL of acetonitrile. Add ~20 mg of your
synthesized aspirin to the first tube, ~20 mg of commercial aspirin to the second tube, and
~20 mg of salicylic acid to the third tube. If necessary, shake the test tubes to help dissolve
the solids.
3. Using a mechanical pencil, lightly draw a line about 1 cm above the bottom of the TLC plate.
Dip a clean capillary tube into the first solution and tap it on the line you have drawn.
Additional solid can be deposited by tapping the same spot after the initial deposit dries.
Repeat this process for the other two solutions (in different places on your line). Make sure
you record the order in which you placed the dots.
4. Place the TLC plate in the eluting chamber and close the lid. Watch the progress of the
solvent line, and when it nears the top of the plate, remove the plate from the eluting
chamber. Mark the solvent line with the mechanical pencil.
5. Once the eluting solution has evaporated from the plate, place it in the developing chamber
and close the lid. Once all the spots are visible, remove the plate from the developing
chamber and close the lid.
Assignment:
For your report you must write a formal lab report. The entire point of this lab was to
synthesize, purify, and characterize a sample of acetylsalicylic acid. Use the report to explain the
task, provide background information, present your findings, and consider the results. Talk
about the effectiveness of the procedure, and how you determined the purity of your sample. Be
sure to include a yield based on the amount of salicylic acid you started with.
Post-lab:
The following questions should be answered in your lab notebook:
1. Why is it important to always cool the solution in an ice bath before filtering through the
Büchner funnel? How would your percent yield be affected if you did not cool the solution
before filtering?
2. The concentrated sulfuric acid does not occur anywhere in the balanced equation for this
reaction. Why do you think it is needed in this reaction?
3. If everything goes well, you should notice that the commercial acetylsalicylic acid spot on
the TLC plate contains two compounds, acetylsalicylic acid and salicylic acid, while your
compound (after the recrystallization) is a bit more pure. Why is this? You might have to do
some research to figure this one out.
Analysis of Aspirin by Visible Spectrophotometry
In a previous experiment, you prepared aspirin by combining acetic acid and salicylic
acid. Your synthesis converted most, but not all, of the salicylic acid into acetylsalicylic acid
(aspirin). While there are ways in which the purity of your acetylsalicylic acid can be
determined directly, you will use an indirect method. As can be seen by the chemical system
presented, when iron (III) nitrate is mixed with salicylic acid, a bluish-purple color complex ion
is formed. You will analyze a sample of your crude aspirin to determine the amount of salicylic
acid in it. You can use this information to calculate the purity of your aspirin sample if you
assume that the unreacted salicylic acid is the only impurity,.
+1
H2O
H
H
OH2
Fe
O
O
OH2
H2O
+ Fe3+(aq)
O
O
C
C
O
O
Visible Spectrophotometry is a method of measuring the concentration of colored
solutions by the amount of light absorbed by or transmitted through a colored solution. The
absorbance of white light by a solution containing a colored compound is directly proportional to
the concentration of the colored compound. The constant of proportionality contains the path
length of the sample through which the light passes and a constant that is determined by the color
of the solution. This information produces Beer's Law: A = bc, where A is absorbance,  is the
molar absorptivity of the colored solution, b is the inside diameter of the cell, and c is the molar
concentration of the solution. In this experiment, we will not attempt to determine the values
 nd b. Rather we will us a Beer’s Law plot as a calibration curve or a conversion curve.
From the calibration curve we will be able to determine the concentration of an unknown
solution. For a discussion of the spectroscopic theory of this experiment, search “Beer’s Law” in
Wikipedia. You will use a Vernier Visible Spectrophotometer in this experiment
Procedure
1.
Prior to coming to the lab, calculate how much 0.25 M Fe(NO3)3 is needed to make
approximately 700 mL of a 0.025 M Fe(NO3)3 solution
2.
You must prepare a standard salicylic acid solution of known concentration. Measure on
the analytical balance approximately 0.08 g of pure salicylic acid and record the precise
mass. Transfer the salicylic acid to a 100-mL volumetric flask and add about 10 mL of
95% ethanol. Swirl the flask to dissolve the acid. Add some distilled water to the flask
and mix. Add more distilled water to fill the flask to the 100.00 mL mark. Mix the
solution thoroughly. Transfer this stock standard salicylic acid solution a clean, dry
beaker Calculate the precise concentration of your stock solution in mg/mL.
3.
Prepare approximately 700 mL of a 0.025 M Fe(NO3)3 solution from the stock Fe(NO3)3
solution which is 0.25 M.
4.
Use the stock salicylic acid solution to prepare five diluted samples. Pipet 2.00 mL, 4.00
mL, 6.00 mL, 8.00 mL, and 10.00 mL of stock solution into five labeled (1-5) 100-mL
volumetric flasks. To each flask add 0.025 M Fe(NO3)3 to make precisely 100.00 mL.
Determine the concentration (mg/mL) of salicylic acid in each flask before proceeding.
5.
Prepare your crude aspirin sample. Measure out about 0.08 gram of your crude aspirin
and record the precise mass that you use. Transfer the crude aspirin to a 50-mL
volumetric flask and add about 10 mL of 95% ethanol. Swirl the flask to dissolve the
acid. Add distilled water to the flask with mixing. Fill the flask to the 50.00 mL mark
with distilled water. Mix the solution thoroughly. Transfer the stock crude aspirin
solution to a clean, dry beaker.
6.
Prepare two (2) samples for your crude aspirin for analysis. Transfer 10.00 mL of the
crude aspirin solution to a clean 50-mL volumetric flask and also to a clean 100-mL
volumetric flask. Add 0.025 M Fe(NO3)3 solution to each flask precisely to the line. Mix
each solution thoroughly.
The use of the visible spectrophotometer is in three parts:
Part 1: Determination of the optimum wavelength
1. Using a USB cable, connect a Vernier Spectrometer to a computer.
2. Start the Logger Pro 3.x program on your computer.
3. The spectrometer should be set up at this point. If not, open the Experiment menu and select
connect Interface . Spectrometer . Scan for Spectrometers.
4. Calibrate the spectrometer.
a. Prepare a blank. Rinse a cuvette with the 0.025 M Fe(NO3)3 solution and fill it ¾ full with
this solution.
b. Open the Experiment menu and select Calibrate . (Spectrometer). The following message
appears in the Calibrate dialog box: “Waiting … seconds for the device to warm up.” After
60 seconds or so, the message changes to: “Warmup complete.”
c. Place the blank in the cuvette holder of the spectrometer. Align the cuvette so that the clear
sides are facing the light source of the spectrometer. Click “Finish Calibration”, and then
click OK.
5. Determine the maximum wavelength for your standard salicylic acid solution and set up the
data collection mode.
a. Select five (5) more curvettes. Rinse each, in turn, with one of the labeled (1-5) standard
salicylic acid solutions. Fill each cuvette ¾ full with the corresponding standard solution.
b. Wipe cuvette #3 with a tissue and place it in the cuvette holder of the spectrometer.
c. Click the Collect button. A full spectrum graph of the standard salicylic acid solution will
be displayed. Note that one area of the graph contains a peak absorbance. Click the Stop
button to complete the analysis. Examine the peak absorbance.
d. To save your graph of absorbance vs. wavelength, select Store Latest Run from the
Experiment menu.
e. Remove cuvette #3 from the spectrometer.
Part 2: Determination of the Beer’s Law plot
6. Click the Configure Spectrometer Data Collection icon, on the toolbar. A dialog box will
appear.
a. Select Absorbance vs. Concentration under Set Collection Mode. The peak absorbance
should be automatically selected. If not, check the appropriate box.
b. Change the Units from mol/L to mg/mL.
7. Collect absorbance-concentration data for the five standard solutions.
a. Wipe cuvette #1 with a tissue and place it in the spectrometer cuvette holder. Click the
Collect button. When the absorbance reading stabilizes, click the Keep button. Enter your
calculated concentration of solution #1 and click OK.
b. Repeat Step 7b for the remaining solutions.
c. When you have finished testing the standard solutions, click the Stop button.
8. To determine the best-fit line equation for the standard salicylic acid solutions, click the linear
fit button on the toolbar. Write down the equation for the standard solutions in your lab book.
Part 3: Determination of the Concentration of the Salicylic Acid in the Crude Aspirin
9. Determine the concentration of the salicylic acid in your crude aspirin.
a. Rinse a cuvette twice with a crude aspirin solution and fill it about ¾ full. Wipe the outside
of the cuvette and place it into the spectrometer.
b. Select Interpolation Calculator, from the Analyze menu. A dialog box will appear that
displays the concentration of your unknown at the measured absorbance.
c. Click OK. Write down the concentration of the unknown in your lab book.
d. Remove the curvette from the spectrometer.
e. Repeat 9a-d with the other crude aspirin solution.
f. Dispose of all solution down the drain. Rinse all volumetric flasks and curvettes twice with
distilled water and put away.
Data Analysis
1. Print the graph of the visible spectrum of the Fe3+/salicylic acid complex. Reset both axes to
maximize the plot.
2. Print the graph showing the data and linear-regression equation of the Beer’s Law plot for the
standard solutions.
3. Determine the mass of the salicylic acid in the crude aspirin.
4. Determine the percent purity of your crude aspirin.
In addition to the regular components included in the summary, describe an alternate method for
determining the molar concentration the salicylic acid in the crude aspirin using the standard
data.