by Patrick Gormley
Table of Contents
Laboratory Experiments
Appendix
Sample Lab ReportStatistical Calculations
Relative Error  Median 
Determination of Outliers  Standard Deviation 
Mean or Average  Average Deviation 
Grading Policies
Nine week marking period grades will be based on:
Lab Reports & Homework  20% 
Tests  40% 
Citizenship  15% 
Quizzes  25% 
Semester grades will determined by computing a weighted average of the first marking period (40%), the second marking period (40%), and the semester final exam (20%). Two passing grades are necessary in order to earn a passing semester grade. (In the event that any student is exempted from taking a semester final exam, a passing grade for both marking periods is required in order to pass the semester. A student with a failing grade for one marking period will not earn a passing grade for the semester unless he/she elects to take, and pass, the semester final exam. )
As stated above a portion of your nine week grade will be based on citizenship. At the end of a marking period 10 citizenship points, minus any deductions, will be given. Situations which can result in a loss of points include, but are not limited to: tardies, assignments not completed on time, inattentiveness in class, talking in class, coming to class unprepared, and suspensions.
Regular attendance is important in this class. Lapeer Community Schools has an attendance policy. The details of this policy is covered in your handbook. Basically excessive absences and tardies will result in no credit earned for a class. Tardies are converted into absences as per a formula. I have a policy of marking students tardy that are observed running to this class. Running in the hallways is not only dangerous to you, but to other students as well.
All assignments must be submitted for grading on the date they are due. Late assignments are accepted only with prior approval of the instructor and will be marked down at the rate of 20% per day. Laboratory reports and homework assignments may not be late more than two days without permission. All missed tests and quizzes, due to excused absences, must be made up within five school days, and it is the responsibility of the student to check with the teacher regarding makeup assignments when they return from an absence. Failure to do so will probably mean that these assignments will not be graded for credit since only five days are allowed for such makeup work.
Sample Grade Calculation
Homework= 95% Quiz= 85% Test= 82% Citizenship=60% (2 tardies & unprepared twice)
Marking Period Grade= 95 X 0.20 + 85 X 0.25 + 82 X 0.40 + 60 X 0.15 = 82% or a B
(Note: Had the Participation grade been 100% the final weighted average would have been 88% or a B+. )
Posting of Grades:
Grades are posted every week by student number beginning with the fifth week of each marking period. In addition, your current progress in this class is available on the Internet at http://lehs.lapeer.org/ScienceDept/. After this web page is displayed follow the "Chemistry Grades" link. You will need your student number in order to access your record. (Alternatively, you can go directly to the site by pointing your browser to: http://lehs.lapeer.org/gradebook/search.lasso.) This service will go into effect beginning with the fourth week of the first marking period andcontinue to the end of the school year.
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
The stomach's normal pH range is from 1.0 to 3.0. Acid indigestion and heartburn generally occur at pH values near or lower than 1.0. Various substances are used to neutralize the excess hydrogen ion in the stomach to relieve acid indigestion. These products generally contain one or more of the following: magnesium hydroxide (Mg(OH)_{2}), calcium carbonate (CaCO_{3}), sodium bicarbonate (NaHCO_{3}), or sodium aluminum dihydroxy carbonate (NaAl(OH)_{2}CO_{3}).
When comparing the "neutralizing power" of various antacids it is desirable to have a common denominator for comparison. This is usually accomplished by calculating the equivalent weight of the antacid, which represents the mass of antacid needed to neutralize one mol of hydrogen ions. The analysis in this laboratory exercise will be similar to the procedure used in the lab "Equivalent weight of an Unknown Acid". However, to avoid the possibility of a buffer system being established and because the neutralization process is slow we will use a titration technique known as "back titration".
In this technique an excess of a strong acid, like HNO_{3}, is added to quickly neutralize the antacid and break any buffer system that might be established. The solution is then heated to drive off any carbon dioxide and the excess acid titrated with a strong base of known concentration. The total number of equivalents of antacid in the commercial sample plus the number of equivalents of strong base used in the back titration equals the number of equivalents of the strong acid added. Therefore, the number of equivalents of antacid in the sample is equal to the number of equivalents of acid added minus the number of equivalents of base used in the back titration.
The equivalent weight (grams/equivalent) will be used to compare the effectiveness of various commercial brands and the cost per equivalent will be used to determine the "best buy".
Procedure
1. Mass out to the nearest milligram approximately 800 mG of pulverized antacid and place in a 250 mL Erlenmeyer flask. Pipet 50.0 mL of 0.1 N HNO_{3} into the flask and swirl to dissolve. (The inert ingredients may not dissolve but will probably form a suspension.) Heat the solution to boiling and continue to heat for about a minute to remove any dissolved CO_{2}. Add 6 drops of bromphenol blue indicator. Bromphenol blue is yellow in an acid solution and blue in a basic solution. If the indicator is not the proper color more acid will have to be added. Record the exact amount of acid added.
2. Titrate the excess acid with a standardized solution of NaOH.
3. Repeat steps 1 and 2 with a second sample.
Questions
1. If the CO_{2} is not removed by boiling after the 0.1 N HNO_{3} is added, how does this affect the amount of NaOH required to reach the bromphenol blue endpoint? Explain.
2. If the antacid selected for analysis is known to contain only magnesium hydroxide, how could the procedure be modified to expedite the analysis?
3. Calculate the equivalent weight of the antacid and the cost per equivalent?
4. Which antacid analyzed by the class is the "best buy"?
5. Summarize all calculations in a nice neat table.
Objectives
Discussion
[This laboratory exercise was adapted from an experiment in Martyn Berry & Averil Macdonald's Chemistry Through HydrogenClean Energy for the Future, Heliocentris Energiesysteme GmbH, Berlin, Germany, 2000
The atomic weight system used by chemists is a relative system which indicates how massive an atom is relative to some arbitrary standard. Throughout history there have been three standards: hydrogen, oxygen, and now Carbon12 . Carbon12, C^{12} , the most common isotope of carbon, has arbitrarily been given a value of exactly 12. All other atoms are compared to C^{12} and assigned a value which reflects how much heavier, or lighter, they are relative to carbon12 atoms. Therefore, since magnesium atoms are about twice as heavy as C^{12} , they are given a value of about 24; oxygen atoms, which are about 33% heavier than C^{12} atoms are given a value of 16. For the elements these relative weights are called atomic weights. Molecular weights are calculated by adding atomic weights so they also are relative weights and represent how many times heavier a molecule is relative to C^{12} .
One of the most important consequences of having a system of relative atomic weights is that it can be proven mathematically that 1 gramatomic weight (GAW) of element A contains just as many atoms as 1 GAW of element B. One GAW is defined as that mass of an element which is equal to the atomic weight expressed in grams. Therefore, 1 GAW of oxygen is equal to 16 grams of oxygen while 1 GAW of magnesium is equal to 24 grams of magnesium.
In 1811 the first person to realize this was Amedeo Avogadro and so the number of atoms in one GAW become known as Avogadro's number. It was first measured by the American chemist Robert Millikan at 6.06 X 10^{23}. Today the accepted value of Avogadro's number to six significant digits is 6.02257 X 10^{23} and is known as the mol. Therefore, 12 grams of carbon equals 1 GAW of carbon which contains 6.02257 X 10^{23} atoms of carbon and is called one mol of carbon atoms. Likewise, 24 G Mg = 1 GAW Mg = 6.02257 X 10^{23} atoms Mg =1 mol Mg. In this experiment you will measure the value of the Avogadro constant (Avodadro's number) by determining the number of electrons in one mol of electrons.
The charge on the electron, e, was first measured by Robert Millikan around 1910. Today the accepted value is 6.021892 X 10^{19} Coulombs (C). This translates into 6.241460 X 10^{18} electrons/C. A coulomb is an ampsec. Therefore, the number of coulombs of charge passing through a circuit can be calculated by multiplying the amps passing through a circuit times the time in seconds.
The electrode reactions in the electrolyser are:
2H^{1+ }+2e^{1} > H_{2} and
2H_{2}O > O_{2} + 4H^{1+ }+ 4e^{1}
Procedure
1. Assemble the solar cell, lamp, electrolyser and load measurement box as shown in the diagrm below.
All connections must be correctly made, with the correct polarity. Check with you teacher before proceeding. collect hydrogen with the switch at Short Circuit.
2. Fill the gas storage cylinders of the electrolyzer with distilled water to the 0 mL mark.
3. Adjust the lamp distance from the solar cell so that about 300 mAmps are flowing through the circuit.
4. Record the time needed (in minutes) to collect 10 mL of hydrogen gas. Also, record the volume of oxygen collected at the other electrode.
During the expeiment take a reading of the number of amps passing through the circuit at 1 minute intervals. Complete a chart, similar to the following, for each minute the experiment was run.
5. Record the current atmospheric pressure and temperature.
Questions
1. Determine the charge, that is, the number of Columbs, passing through the circuit for each minute the experiment was run. From these values determine the total charge passing through the circuit. Convert this charge into electrons.
2. Using the ideal gas equation detemine the number of mols of hydrogen collected. Don't forget to correct the atmospheric pressure for the vapor pressure of water.
3. Using the ideal gas equation detemine the number of mols of oxygen collected. Don't forget to correct the atmospheric pressure for the vapor pressure of water.
4. Using the half reaction given in the discussion above for the collection of hydrogen, determine the number electrons in one mol of electrons, that is, determine Avogadro's number, also known as Avogadro's constant.
5. Using the half reaction given in the discussion above for the collection of oxygen, determine the number of electrons in one mol of electrons, that is, determine Avogadro's number, also known as Avogadro's constant.
6. Identify the anode, that is, was it the electrode producing the hydrogen or oxygen?
7. Identify the cathode, that is, was it the electrode producing the hydrogen or oxygen?
8. Average the values obtained for Avogadro's number in questions 4 and 5.
9. Summarize your results in a nice neat table.
10. For students having an XX chromosome. Report the values obtained by the class for Avogadro's number at the anode and calculate the average. Determine if their are any outliers at the 90% confidence level. If any outliers are found, recompute the average without the outliers.
11. For students having an XY chromosome. Report the values obtained by the class for Avogadro's number at the cathode and calculate the average. Determine if their are any outliers at the 90% confidence level. If any outliers are found, recompute the average without the outliers.
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
Commercial bleaching agents usually contain the hypochlorite ion (ClO^{1}) as the active oxidizing ingredient. This ion is generally in the form of either the sodium or calcium salt. The number of equivalents of oxidizing agent per unit volume governs the strength of the bleach. The common denominator used by the industry to quote the strength of the bleach is weight percent of chlorine per unit volume (or mass of powder).
Liquid bleaches are usually about 5.25 % NaOCl by weight and are prepared by the electrolysis of a cold sodium chloride solution, producing Cl_{2} at the anode and OH^{1 }at the cathode. If the solution is kept well stirred the Cl_{2} and OH^{1 }react to form OCl^{1 }. The reactions are given below:
anode: 2Cl^{1 }> Cl_{2} +2e^{}
cathode: 2e^{} > 2OH^{1 }+ H_{2}
Cl_{2} + 2OH^{1 }> OCl^{1 }+ Cl^{} + H_{2}O
net reaction: Cl^{1 }+ H_{2}O > OCl^{1 }+ H_{2}
Bleaching powder is a mixture of CaCl_{2}, Ca(OCl)Cl, and Ca(OCl)_{2} and is prepared by the reaction of Cl_{2} and slaked lime, Ca(OH)_{2}.
In both types of bleaches the OCl^{1} is the oxidizing agent which removes electrons from the substances that make up stains and thereby destroys the substance's ability to absorb light. Since light is no longer absorbed, the material appears cleaner because it is whiter.
The redox analysis in this experiment involves the reaction of excess iodide ion with the hypochlorite ion in a commercial bleach. The reaction is below:
OCl^{1 }+ 2I^{1 }+H_{2}O > I_{2} + Cl^{} + 2OH^{1}
However, since we will be analyzing the bleach in terms of available chlorine we will consider the iodide to have reacted with dissolved Cl_{2} in the following reaction:
Cl_{2} + 2I^{1 }> I_{2} + 2Cl^{1}
The iodine produced is then titrated with a standardized sodium thiosulfate solution, NaS_{2}O_{3}, using starch as an indicator. The endpoint of the titration occurs when the blue color of the starchiodine complex disappears. The reaction is:
I_{2} + 2S_{2}O_{3}^{2} > 2I^{1} + S_{4}O_{6}^{2}
Since substances always react in a one to one equivalent ratio, the number of equivalents of thiosulfate required to reach the endpoint is equal to the number of equivalents of available chlorine.
Procedure
1. Determine the density of liquid bleach to at least three significant digits.
2. Pipet 10.0 mL of bleach into a 100 mL volumetric flask and dilute it to the 100 mL mark with distilled water. Mix thoroughly. Pipet 50 mL of this diluted bleach solution into a 250 mL Erlenmeyer flask and add 2 grams of KI and 10 mL of 3 M H_{2}SO_{4}. A yellow color indicates the presence of free iodine.
3. Immediately titrate the liberated I_{2} with standardized Na_{2}S_{2}O_{3} solution until the yellow I_{2} color is almost gone. Add 2 mL of starch solution to form the deep blue starchiodine complex and finish titrating until the blue color is gone.
4. Repeat steps 2 and 3 of the analysis with a second sample.
Questions
1. Calculate the percent "available Cl_{2}" by weight for each trial and average the results.
2. Determine the cost per equivalent of "available Cl_{2}" for each sample analyzed and determine the "best buy".
3. Determine the percentage of NaOCl in the bleach. (You may assume that NaOCl was the only oxidizing agent in the bleach.)
4. The starch solution can be added at the same time as the KI. However, the complex formed between the I_{2} and starch molecules does not readily dissociate when the thiosulfate titrant is added. What difficulties may you expect to encounter if the starch solution is added earlier?
5. What effect would adding the starch solution with the solid KI have on the calculated weight percent of available Cl_{2}?
6. If an air bubble initially trapped in the buret's tip is released during the titration, will the reported percent available Cl_{2} be too high or too low? Explain.
7. Two grams of KI were added in this investigation to react with the hypochlorite ion. What is the maximum concentration of bleach (in % NaOCl) that could be theoretically analyzed for in this laboratory investigation?
8. From the class values listed on the chalkboard determine for the percentage of NaOCl in the bleach determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the percentage NaOCl in the bleach and the standard deviation.
9. Summarize all calculations in a nice neat table.
Colligative Properties: Formula Weight Determination
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
The addition of a nonvolatile solute to a solvent produces characteristic changes in a solvent's physical properties. For example, ethylene glycol is added to car radiators to "protect" the car's cooling system from freezing. The presence of the ethylene glycol causes the freezing point of the water to be lowered to as much as 84 ^{o}F (64 ^{o}C) for a solution which is 70% ethylene glycol and 30% water. At the same time the boiling point of this solution is increased to 276 ^{o}F (136 ^{o}C). The drop in the freezing or the rise in boiling point for a 1 molal solution is independent of the identity of the solute and are constants unique to a particular solvent. These constants are called the freezing point depression constant and the boiling point elevation constants respectfully and are two colligative properties of solutions. The constants for a few common substances are given in the table below:
In this experiment you will determine the formula weight of a nonvolatile, nonelectrolyte solute in cyclohexane by measuring the freezing point for pure cyclohexane and a solution of an unknown solute dissolved in cyclohexane. Cyclohexane's freezing point and that for the solution are to be obtained from a cooling curvea graph representing the temperature of the system as a function of time. Cyclohexane's cooling curve, like that of most pure substances, reaches a plateau at its freezing point; extrapolation of the plateau to the temperature axis determines its freezing point. The solution's cooling curve does not reach a plateau, but continues to decrease slowly as the solution freezes. Its freezing point is determined at the intersection of two lines drawn tangent to the curves both above and below the freezing point and extrapolating this point to the temperature axis.
Procedure
1. Assemble the apparatus as directed by you instructor. Mass a clean dry 200 mM X 20 mM test tube and add approximately 15 G of cyclohexane to it and remass the test tube and its contents. Prepare about 300 mL of a methanolice slurry in a 400 mL beaker. Place the test tube and cyclohexane in the methanolice slurry and insert a thermometer to measure the temperature.
2. Hold the thermometer in the liquid so that it does not touch the sides or bottom of the test tube. Take temperature readings every 1530 seconds until the liquid is completely frozen. You can tell when the liquid is freezing because the temperature remains virtually constant at the freezing point until solidification is almost complete.
3. When you are completed with your temperaturetime readings remove the test tube from the methanolice bath and allow the contents of the test tube to melt and come to room temperature before removing the thermometer. Note: Do not melt the solid in the test tube with your hand as frostbite may occur! Mass out about 0.5 G of unknown to the nearest milligram and add to the cyclohexane. Stir to dissolve and take timetemperature readings as you did in step 2 while cooling the mixture to below its freezing point.
Questions
1. Determine the molecular weight of the unknown substance. Show all work.
2. If the solution's freezing point is erroneously read 0.2 ^{o}C lower than it should be, will the unknown's calculated formula weight be too high or too low? Explain.
3. How will the freezing point change of cyclohexane be affected by:
a) a nonvolatile solute that dissociates?. Explain
b) a volatile solute that does not dissociate? Explain.
c) two solutes that react as: A + B > C?. Explain.
4. If a thermometer is miscalibrated to read 0.5 ^{o}C higher than the actual temperature over its entire scale, how will it affect the reported molecular weight of the solute? Explain.
5. If some solute adheres to the test tube's wall in step 3, is the freezing point change greater or less than it should be? Explain.
6. How will the error referenced in question 5 above affect the calculated molecular weight of the unknown, that is, would the calculated molecular weight be higher or lower? Explain.
7. If the cyclohexane is initially contaminated with a nonreactive, nonvolatile solute, how (if at all) does this affect the reported formula weight value?
8. Assuming that your unknown is soluble in tbutanol and that tbutanol was used in this experiment instead of cyclohexane. How much would the freezing point been lowered?
9. From the class values listed on the chalkboard for the molecular weight of the solid determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the molecular weight of the unknown and the standard deviation.
10. Summarize all calculations in a nice neat table.
Complexometric Titration of Zinc
Objectives
Discussion
[This laboratory exercise was adapted from "Complexometric Titration of ZincAn Analytical Chemistry Laboratory Experiment" in December 1997 edition of the Journal of Chemical Education]
A complexometric titrations with EDTA to determine the hardness of water was done in Chemistry I. This titration was performed at a high pH, because the formation constants of Mg*EDTA and Ca*EDTA are high. There are several ions that form complexes with EDTA at low pH. Among these is the Zn^{2+ }ion. The indicator being used in this titration is Xylenol Orange. When this indicator is complexed with Zn^{2+ }it is red, in the free state it is yellow. During the titration a competition is set up between EDTA and the indicator. When all of the zinc ions have been complexed with EDTA, only the free indicator remains, causing a change in color from pink to yellow. The EDTA will first complex the free zinc ions, then it will complex the zinc ions attached to the indicator.
The commercial product to be analyzed will either be a cold lozenge, like ColdEeze®, produced by Quigley Corporation or a vitamin supplement.
The reaction is: Zn^{2+ }+ EDTA > Zn*EDTA^{2+ }
Procedure
1. Prepare the following solutions, if they have not already been prepared: 0.01 M EDTA, pH 5.5 buffer solution. The buffer solution is prepared by dissolving 1.35 G of glacial (18M) acetic acid and 10.25 G anhydrous sodium acetate (or 17.01 G sodium acetate trihydrate) in 250 mL of distilled water. Adjust pH to 5.5 with either 6 M NaOH or glacial acetic acid.
2. Dissolve the sample to be analyzed in 50.0 mL of buffer solution. Grinding and heating may be necessary. Cool to room temperature.
3. Prepare a blank solution of buffer and a few drops of indicator to determine the endpoint color.
4. Add a few drops of indicator to the solution to be analyzed and titrate with EDTA to the endpoint.
Questions
1. Determine the number of mols of Zn, as Zn^{2+ }, present in the product.
2. Determine the percentage of Zn present in the product.
3. How would your answer in question 2 be affected if you "overshot" the endpoint? Explain
4. How would your answer in question 2 be affected if your EDTA solution happened to be more concentrated than 0.01 M. Explain.
5. Prior to preparing the EDTA solution the solid chemical should have been dried in an oven for about one hour. Let's assume that it was not dried. How would this error affect your answer in question 2? Explain.
6. If an air bubble initially trapped in the buret's tip is released during the titration, will the reported percent in question 2 be higher or lower? Explain.
7. From the class values listed on the chalkboard determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the percentage of Zn in the commercial product analyzed. Determine the standard deviation.
8. Summarize all calculations in a nice neat table.
Determination of Phosphorous in Fertilizer
Objectives
Discussion
[This laboratory exercise was adapted from "Quantitative Determination of Phosphorous in Plant Food Using Household Chemicals", by Sally Solomon, Alan Lee, and Donald Bates, Journal of Chemical Education, p.410., May 1993]
Plants, like animals, require essential nutrients to live. The three main nutrients contained in plant foods are nitrogen, phosphorous (as P_{2}O_{5}), and potassium (as K_{2}O). For example, a common allpurpose plant food labeled 153015 would contain 15% nitrogen, 30% P_{2}O_{5} , and 15% K_{2}O. However, the percentage of the elements present in this fertilizer would be 15% nitrogen, 13% phosphorous, and 12.5% potassium.
The gravimetric analysis of phosphorous in this experiment is based on the precipitation of phosphorous as MgNH_{4}PO_{4}*6H_{2}O from a solution that contains the monohydrogen phosphate ion (HPO_{4}^{2 }), ammonium ions, and magnesium ions. The balanced reaction is:
5 H_{2}O + HPO_{4}^{2 }+ NH_{4}^{1+ }+ Mg^{2+ }+ OH^{1 }<=====> MgNH_{4}PO_{4}. 6H_{2}O
The precipitate forms only upon the slow neutralization of the solution with ammonia water. In an acid solution the phosphate ions react with hydronium ions to form monohydrogen phosphate ions. The net effect of this is to reduce the concentration of the phospate ions (PO_{4}^{3 }) and prevent the formation of MgNH_{4}PO_{4•}6H_{2}O.
PO_{4}^{3 }+ H_{3}O^{1+ }<=====> HPO_{4}^{2 }+ H_{2}O
The hydroxide ions necessary for the neutralization must be supplied by a weak base such as ammonium hydroxide. A strong base, like NaOH, would cause the precipitation of Mg(OH)_{2} instead of the double salt magnesium ammonium phosphate hexahydrate.
Procedure
1. Mass out to the nearest milligram approximately 10 grams of plant food. Dissolve in about 150 mL of tap water. Plant foods generally contain some insoluble material even though they are advertised as "water soluble". If your plant food contains some insoluble material remove it by filtration.
2. Transfer the dissolved plant food to a 1 L Erlenmeyer flask and add 0.40 M MgSO_{4}^{.}7H_{2}O . The amount of magnesium sulfate added should be about 5 mL/100 mG of P_{2}O_{5} that is believed to be in your sample. Your instructor will tell you the approximate amount of P_{2}O_{5} in your plant food.
3. Add a few drops of phenolphthalein to the solution and then slowly add 1M aqueous NH_{3} (NH_{4}OH) with constant swirling or stirring until a white precipitate forms or until the phenolphthalein endpoint is reached. (If no precipitate forms contact your instructor.) Let the solution stand for at least 15 minutes to insure that precipitation is complete.
4. Filter to collect the precipitate. Any precipitate left behind in the flask can be rinsed on to the filter paper by adding small amounts (50 mL) of 70% isopropanol (2propanol) to the flask. Spread out your filter paper to dry overnight.
5. Carefully scrape the precipitate from the filter paper and collect on a watch glass. Mass the precipitate to the nearest milligram.
Questions
1. From the data obtained determine the percentage of phosphorous (as P_{2}O_{5} ) in the plant food.
2. A large amount of potassium in plant foods could cause the formation of MgKPO_{4} along with MgNH_{4}PO_{4}. What would be the effect on the calculated percentage of P_{2}O_{5} in question 1 if MgKPO_{4} was mixed in with MgNH_{4}PO_{4} in the precipitate? That is, would the percentage of P_{2}O_{5} calculated be higher or lower that if the precipitate contained only MgNH_{4}PO_{4}?
3. How would the calculated percentage of P_{2}O_{5} be affected if the precipitate was not thoroughly dry?
4. Determine the average and standard deviation from the class results.
5. From the class values listed on the chalkboard for the percentage of P_{2}O_{5} in the plant food determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the percentage of P_{2}O_{5} in the plant food and the standard deviation.
6. Summarize your results in a nice neat table.
Possible Research Projects
1. After step 3 the normal procedure is to let the solution stand for several hours instead of 15 minutes. What percentage error (if any) results in letting it stand for only 15 minutes instead of several hours?
2. Is a time interval of less than 15 minutes acceptable?
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
The rate of a chemical reaction can be measured in terms of the rate of disappearance of one of the reactants or in terms of the rate of the appearance of one of the products. For example in the hypothetical reaction, A + 2B > AB_{2} the rate could be measured in terms of the disappearance of either reactant A or B, or in terms of the appearance of the product AB_{2}. The rate is nearly always proportional to each reactant's concentration raised to some power. For the example above this means we would write the following relationship: rate=k[A]^{p }[B]^{q}. This expression is called the rate equation for the reaction. The value of k, called the rate constant, is unique for a specific reaction and is dependant only on the temperature. A study of the kinetics of any reaction involves determining the values of k, p, and q.
In this experiment you will determine the rate equation for the reaction of potassium persulfate, K_{2}S_{2}O_{8}, with potassium iodide. The reaction is: S_{2}O_{8}^{2} + 2I^{1} > 2SO_{4}^{2} + I_{2} . The rate equation for this reaction will be: rate=k[S_{2}O_{8}^{2}]^{p} [I^{1}]^{q}.
In logarithmic form this equation becomes: log(rate)=log(k) + p*log[S_{2}O_{8}^{2}] + q*log[I^{1}]. In order to determine the value of p you will measure the rate of a series of reactions for a constant iodide ion concentration as the persulfate ion concentration is changed and plot log(rate) vs. log[S_{2}O_{8}^{2}]. The value of p is the slope of this plot because at a constant iodide ion concentration the logarithmic form of the rate equation reduces to:
log(rate)= p*log[S_{2}O_{8}^{2}] + C ; where C=k+q*log[I^{1}].
As can be seen this equation now takes the form of the algebraic "slope intercept equation" (Y=mx + b) where p and m are equivalent. The value of q can be determined in a similar manner.
The reaction rate constant (k) will be determined by substituting the values of p, and q back into the rate equation, calculating k for each trial and averaging the results.
The rate of the reaction will be the time it takes to produce a given quantity, 0.1 mmols, of iodine. To aid in determining when 0.1 mmols of iodine has been produced we will react the iodine produce with sodium thiosulfate, Na_{2}S_{2}O_{3}. This reaction is:
2S_{2}O_{3}^{2} + I_{2} > 2I^{} + S_{4}O_{6}^{2}.
The reaction of Na_{2}S_{2}O_{3} with I_{2} is many orders of magnitude faster than the reaction we are studying so its presence will have no effect on our investigation. Exactly 0.2 mmols of Na_{2}S_{2}O_{3} will be added to the reaction mixture for each trial. Therefore, excess iodine will begin to accumulate, and react with a starch indicator, only after all of the sodium thiosulfate has reacted. Since 0.2 mmols Na_{2}S_{2}O_{3} will consume 0.1 mmols of I_{2}, the starch indicator will turn purple when 0.1 mmols of iodine have been produced by the reaction we are studying. The rate of the reaction can be found by dividing 0.1 mmols iodine by the time in seconds it takes for the purple color to appear.
Procedure
The table below summarizes the preparation of the test solutions. Measure the volumes of to ±0.05 mL using either a pipet or a buret. All volumes are in mL.
Solution A

Solution B


Test Soln.  Water  0.3 M KI  starch  0.02 M Na_{2}S_{2}O_{3}  0.1 M K_{2}S_{2}O_{8} 
1  148  10  2  10  30 
2  138  20  2  10  30 
3  128  30  2  10  30 
4  108  30  2  10  50 
5  88  30  2  10  70 
1. Prepare solution A in a 250 mL beaker or 250 mL Erlenmeyer flask. Stir the solution thoroughly and record its temperature.
2. The reaction begins when solution B is poured into solution A, therefore, be prepared to start timing the reaction in seconds with a watch or clock. Place the reaction vessel on a white sheet of paper so that the color change is more easily detected.
Calculations
1. Determine the initial concentration, after mixing, of both persulfate and iodide ions for all trials.
2. Determine the rate for each trial in mmols iodine/sec.
3. Using trials 1,2, and 3 plot log rate vs. log([I^{1}]). Determine q from this plot.
4. Using trials 3, 4, and 5 plot log rate vs. log([S_{2}O_{8}^{2}]). Determine p from this plot.
5. Using the p and q determined in steps 3 and 4 and the measured rate, substitute into the rate equation for each trial and determine k. Calculate the average k.
Questions
1. State the effect of each of the following changes on the reaction rate in this experiment. Explain your answers.
a) substituting 0.001 M acetic acid for the distilled water in solution A.
b) increasing the concentration of persulfate.
c) increasing the concentration of KI.
d) an increase in the Na_{2}S_{2}O_{3} concentration.
e) increasing the concentration of the starch solution.
f) increasing the temperature of the reactants.
2. Two test reactions are the minimum required to obtain values of p and q in this experiment. Explain the advantage that additional test reactions have on determining p and q as you did in this experiment.
3. What would appear in solution if the Na_{2}S_{2}O_{3} solution were omitted?
4. Write the rate law for the reaction investigated in this experiment with the determined values of p, q, and k.
5. One of the errors in this experiment is in the way we calculated the rate of each trial (dividing 0.1 mmols of iodine by the time in seconds). This is known as the average rate. What kind of rate would have to be calculated for each trial to improve the procedure? Explain.
6. Summarize all calculations in a nice neat table.
Objectives
Discussion
The activation energy of a reaction is the amount of energy needed to start the reaction. It represents the minimum energy needed to form an activated complex during a collision between reactants. In slow reactions the fraction of molecules in the system moving fast enough to form an activated complex when a collision occurs is low so that most collisions do not produce a reaction. However, in a fast reaction the fraction is high so that most collisions produce a reaction. For a given reaction the rate constant, k, is related to the temperature of the system by what is known as the Arrhenius equation:
where R is the ideal gas constant (8.314 J mol^{1} °K^{1}), T is the temperature in degrees Kelvin, E_{a} is the activation energy in joules per mol, and A is a constant called the frequency factor, which is related to the fraction of collisions between reactants having the proper orientation to form an activated complex.
In this experiment you will measure the rate constant for a chemical reaction, the same one investigated in "Determination of a Rate Law", at several temperatures, plot log k Vs. 1/T, and determine the activation energy of the reaction from the slope. (It is left to the reader to determine how the slope is related to the activation energy. Review the discussion for the following labs: "Determination of a Rate Law", and "The Nernst Equation and the Copper Electrode".) This lab will probably be done after the lab entitled "Determination of a Rate Law". Review the data collected in that lab to determine the best test solution to used for this investigation and create your own procedure. At least four data points should be collected for a good plot: room temperature, one below room temperature, and two above room temperature.
Procedure
The class will be responsible for mixing up all solutions. Divide this responsibility among you and have your procedure approved by your instructor prior to starting the experiment.
Questions
1. Plot log(k) Vs. 1/T and determine the activation energy from the slope. Place your activation energy on the chalkboard.
2. How would the recorded time for the blue color to appear be affected if the solutions being mixed at the temperatures above room temperature cooled after mixing? Explain
3. How would the error addressed in question 2 above affect the calculated activation energy? Explain.
4. From the class values listed on the chalkboard for the energy of activation determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the energy of activation and the standard deviation.
5. Summarize your findings in a nice neat table.
Equivalent Weight of an Unknown Acid
Objectives
Discussion
By definition one equivalent (or equivalent weight) of a substance is the amount of that substance which supplies or consumes one mol of reactive species. In acidbase chemistry the reactive species is the hydrogen ion (H^{1+}) while in oxidationreduction chemistry the reactive species is the electron. For example, in the following two reactions the equivalent weight of H_{2}SO_{4} would be 49 grams or 0.5 mol in the first reaction but 98 grams or 1 mol in the second. On the other hand, sodium hydroxide has the same equivalent weight in both reactions, one mol or 40 grams.
(1) H_{2}SO_{4} + 2NaOH > Na_{2}SO_{4} + 2H_{2}O
(2) H_{2}SO_{4} + NaOH > NaHSO_{4} + H_{2}O
In the first reaction one mol of H_{2}SO_{4} supplies 2 mols of H^{1+} to NaOH, therefore, onehalf mol of H_{2}SO_{4} or 49 grams is one equivalent. The conditions are different in the second reaction because sulfuric acid only "looses" one hydrogen so the equivalent weight of sulfuric acid is one mol or 98 grams. However, sodium hydroxide behaves the same in both reactions, that is, one mol of sodium hydroxide always "consumes" one mol of H^{1+}, so its equivalent weight remains the same at one mol or 40 grams.
In order to determine the equivalent weight of a substance you must know something about the reaction but it does not have to be balanced. Equivalents can help in the analysis of a substance when the balanced reaction is not known or cannot be written for whatever reason; because one equivalent always reacts with or produces one equivalent. (You should prove this to yourself by calculating how much sodium hydroxide is needed to react with 49 grams of sulfuric acid in each of the two reactions above. Do your calculations using traditional mol relationships and the one to one relationship for equivalents. )
In this lab you will be determining the equivalent weight of an unknown acid by using the sodium hydroxide you standardized from a previous lab. Each partner is to titrate a different unknown acid. Be very careful as part of your grade will be based on accuracy.
Procedure
1. Obtain the unknown acid from your instructor. Write down the code number on the vial. Prepare the sample by grinding it to a fine powder with a mortar and pestle.
2. Dissolve small portions of your unknown acid in distilled water and titrate with the standardized NaOH to a phenolphthalein endpoint.
3. Repeat the titration two more times.
Questions
1. Calculate the equivalent weight of your unknown acid. Report this along with the unknown's code number.
2. Calculate the mass of your unknown acid that would be required to neutralize 100 grams of Al(OH)_{3} assuming that the aluminum (III) hydroxide reacted as follows: Al(OH)_{3}  > Al(OH)^{2+}
3. A better procedure would be to dry the unknown acid prior to titration. Why? What effect would this have on your calculated equivalent weight?
4. Polyprotic acids can undergo neutralization to produce both normal and acid salts. Calculate the equivalent weight of phosphoric acid for the production of each of the three possible phosphate salts.
5. Summarize all calculations in a nice neat table.
Objectives
Discussion
You will be given one of the following problems to solve: determine the molecular formula of hydrated ferrous salt, determine the percentage of iron in a sample, determination of the purity of a iron compound, or determine the concentration of ferrous ion (Fe^{2+}) in a solution. In all cases the main redox reaction will center around the following unbalanced reaction:
Fe^{2+} + MnO_{4}^{1} + H^{1+} > Mn^{2+} + Fe^{3+} + 4H_{2}O
After sample preparation you will titrate the Fe^{2+} ion solution with a KMnO_{4} solution until the color of the purple MnO_{4}^{1} persists. The unbalanced half reactions involved are:
Fe^{2+} > Fe^{3+} + 1e^{}
MnO_{4}^{1} + 5e^{} > Mn^{2+} + H_{2}O
The class will be responsible for mixing up all solutions. Since you will not know the concentration of iron in your sample you will have to make some assumptions in order to mix up the potassium permanganate solution. The assumptions made will depend on the type of problem given to you by your instructor.
Procedure
Have your procedure approved by your instructor prior to starting the experiment.
Questions
Questions will be given to you by your instructor depending on the analysis performed. Your lab report should contain all calculations summarized in a neat table.
Objectives
Discussion
Review the discussion on equivalent weights in the laboratory exercise entitled "Equivalent Weight of an Unknown Acid".
Procedure
At this point we have done enough titration labs that you should be able to create your own procedure. Commercial lime is usually a mixture of CaO and Ca(OH)_{2}. You will need to make some assumptions about its composition and do some preliminary calculations before deciding on a procedure to follow. Good Luck!
Questions
1. Calculate either the equivalent weight of your commercial lime sample or the percentage of CaO in your lime sample.
2. How would your answer in question 1 be affected if the sample was inadvertently contaminated with an inert substance like sand by your lab partner? Explain
3. From the class values listed on the chalkboard for question 1 determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average and the standard deviation.
4. Summarize all calculations in a nice neat table.
The Nernst Equation and a Copper Electrode
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
When a metal is immersed in a solution of its ion an electromotive force (voltage) develops between the ionmetal interface. If this half cell is connected via a salt bridge to another half cell comprising a different metal, a voltage and current will flow between the two cells. A typical cell would look like the following:
The half reactions that would occur with such a cell would be :
Cu^{2+} + 2e^{} > Cu E=0.34 v Zn > Zn^{2+} + 2e^{ }E=0.76 v
If the concentration of all ions were at 1M this cell would produce a potential (voltage) of about 1.1 volts. If the concentration of either the Zn^{2+} or Cu^{2+} is changed the voltage will also change. It has been found that the voltage is related to the concentration by the following equation:
E=E^{o} (0.059/n) Log Q (the Nernst Equation)
The variables E, E^{o}, and n are the calculated potential, the standard potential, and the number of electrons transferred respectfully. The variable Q takes the same form as the equilibrium constant.
In this lab you will collect data on the potential produce in a cell similar to the above when the concentration of the Cu^{2+} ions changes. The concentration of the zinc ion will be kept constant. After constructing a plot of E Vs. Log (Q) you will then be given a solution of Cu^{2+} of unknown concentration and by using your copper electrode determine the concentration of the unknown Cu^{2+} solution. (As you can see a plot of E Vs. Log (Q) should be a straight line because the form of the Nernst equation is the same as the "slope intercept form" you worked with in algebra. What would be the significance of the slope of this plot? What would be the significance of the "Yintercept" of this plot?)
Procedure
1. Assemble the apparatus as directed by your instructor.
2. Fill the porcelain cup threefourths full with 0.1 M Zn(NO_{3})_{2} and place a clean strip of zinc into it. The zinc strip should be connected to the negative terminal of the volt meter.
3. Place about 50 mL of 0.2 M Cu(NO_{3})_{2} in a 150 mL beaker and place a cleaned strip of copper into it. Connect the copper strip to the positive terminal of the volt meter.
4. Lower the porcelain cup into the beaker. After the voltage stabilizes take a reading.
5. Dilute 50 mL of the Cu(NO_{3})_{2} solution to 100 mL with distilled water to create a 0.1M Cu(NO_{3})_{2} solution. Repeat steps 3 and 4. Continue in this fashion until you also have taken readings with Cu(NO_{3})_{2} solutions of 0.05M, and 0.025M.
6. After you have standardized your voltmeter secure the unknown from your instructor and substitute it for the copper solutions used in step 5. Record the voltage and determine the concentration of the Cu^{2+ }ion from your standardization plot.
Questions
1. Make a plot of voltage Vs. log(1/[Cu^{2+}]) or voltage Vs. log([Zn^{2+}]/[Cu^{2+}] whichever you prefer. (You should prove to yourself that both of these plots will be straight lines.)
2. What is the concentration of the unknown copper (II) solution?
3. From the class values listed on the chalkboard for the concentration of the unknown solution determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the concentration of copper in the unknown and the standard deviation.
4. Summarize all calculations in a nice neat table.
Factors Affecting Reaction Rates
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
Chemical kinetics is the study of chemical reaction rates, how they are controlled, and the pathway or mechanism that leads to the reaction's products. The factors that affect the rate of a chemical reaction are: Nature of the Reactants, Subdivision of the Reactants, Temperature, Presence of a Catalyst, and Concentration of the Reactants. In this experiment you will investigate the effect of temperature on the Redox reaction between oxalic acid and potassium permanganate. The reaction rate will be measured by recording the time necessary for the color loss of the permanganate ion. The balanced equation is:
5H_{2}C_{2}O_{4} + 2KMnO_{4} + 3H_{2}SO_{4} > 10CO_{2} + 2MnSO_{4} + 8H_{2}O + K_{2}SO_{4}
The reaction is to be run at four different temperatures and a plot made of temperature (xaxis) versus the time in seconds (yaxis) it takes for the disappearance of the purple color of the permanganate ion. If your plot is not linear try different kinds of plots in an attempt to discover the type of relationship that exists between temperature and reaction rate.
Procedure
1. Measure out 10 mL of 0.33 M oxalic acid into a graduated cylinder. Into a test tube place 2.0 mL of 0.010 M potassium permanganate, which was dissolved in 3 M sulfuric acid. Quickly pour the two solutions together in a beaker, stir, and record the time it takes for the purple color of the permanganate ion to disappear. Record the temperature of the solutions prior to mixing.
2. Repeat step one at three other temperatures. Use a cold water bath to lower the temperature of the reactants to about 10 ^{o}C and a hot water bath to obtain two more runs at about 40 ^{o}C, and 80 ^{o}C.
Questions
1. From your plots of Rate vs. Temperature, how is the rate affected by a temperature change?
2. Which of the following relationships best describes your plot?
R =k T; R =k 1/T ; R =k log(T); R =k log(1/T); log(R) =k T; log(R) =k 1/T
Objectives
Discussion
Many substances occur as hydrates, that is, water molecules are bonded to the ions in the crystalline structure of the salt. The number of mols of water bound in this fashion for a specific substance is usually constant. For example, ferric chloride (Iron (III) chloride) is normally purchased as FeCl_{3}* 6H_{2}O, not as FeCl_{3} and sodium carbonate can be purchased either as Na_{2}CO_{3}*10H_{2}O, or Na_{2}CO_{3}*7H_{2}O, or Na_{2}CO_{3}*H_{2}O. For most hydrates the water is bound loose enough so that it can be removed by gentle heating, but for other hydrates the salt will sometimes decompose before the water can be driven off.
Procedure
You are to complete two trials in this experiment and post them on the chalkboard for use by the entire class in answering some of the questions.
1. Preheat your crucible and its cover with an intense flame for 5 minutes. This will drive off any volatile substances including condensed moisture which would contribute to experimental error.
2. Add at most 2.5 grams of the unknown hydrate to the crucible and mass with the lid.
3. Heat the sample slowly at first gradually intensifying the heat over a 3 minute period. Heat at full flame temperature for five minutes.
4. Cool the crucible and its contents to room temperature and mass.
5. Repeat steps 3 and 4 until the two massings agree to within 2 milligrams.
Questions
1. If some volatile impurities are not burned off in step 1 but are removed by the heating in step 3, is the reported mass of water too high or too low? Explain.
2. What happens to the sample's reported percentage of water if the salt decomposes during step 3 yielding a volatile product?
3. Anhydrous CaCl_{2} removes water vapor from the atmosphere inside a container having an airtight seal; this apparatus is called a desiccator. Explain how CaCl_{2} is able to do this.
4. Determine the percentage of water in your sample for each trial and report the average.
5. Calculate the empirical formula for the hydrate from your average percentage composition.
6. From the class values listed on the chalkboard determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the percentage of water in the hydrate and the standard deviation.
7. Assume that the empirical formula determined in question 5 above is the molecular formula for the salt. Write a properly balanced formula equation for the decomposition of the hydrated salt in this experiment using this formula.
8. How would the percentage of water in the hydrate be affected if during the heating process the anhydrous salt started to react with either atmospheric oxygen or nitrogen to form a substance which was very stable and had a melting point beyond the flame temperture?
9. Summarize all calculations in a nice neat table.
Molecular Weight of a Volatile Liquid
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
The Dumas Method for determining the molecular weight of a volatile liquid, named after John Dumas (18001884) requires the use of the ideal gas law (). The accuracy of the method is therefore dependant on how well the vapors of the volatile liquid emulate an ideal gas at the experimental contitions.
In this experiment a liquid will be vaporized at a measured temperature, T, into the measured volume, V, of an Erlenmeyer flask. After the barometric pressure, P, is recorded, the mols of gas, n, are calculated from the equation. The mass difference between an empty and gas filled flask allows us to calculate the mass of the gas. The molecular weight (MW) is then calculated by the equation: . Alternatively one can substitute for n in the ideal gas law to obtain the following equation:
Procedure
You are to complete two trials in this experiment and post them on the chalkboard for use by the entire class in answering some of the questions.
1. Determine the total mass of a 125 mL Erlenmeyer flask, a rubber band, and a square of aluminum foil.
2. Accurately measure the volume of the Erlenmeyer flask by totally filling the flask with water and transferring the water to a graduated cylinder. (Note: this step can be done at the conclusion of the experiment.)
3. Create a hot water bath by filling a 400 mL beaker half full of water. Heat to boiling. While waiting for the water to boil pour about 56 mL of unknown liquid into the flask. Secure the aluminum foil over the mouth of the flask with the rubber band. Poke a small hole in the foil with a pin to let excess vapor escape during heating.
4. Clamp the flask assembly into the beaker so that flask is as far down as possible in the beaker. Heat at the boiling point of water until liquid is no longer visible in the flask, continue heating for another 10 minutes. Record the boiling point of water to the nearest ±0.1 ^{o}C. Also, record the current barometric pressure.
5. Remove the flask and allow it cool to room temperature. Dry the outside of the flask and mass it along with its contents, the aluminum foil and rubber band.
6. Repeat for the second trial.
Questions
1. Determine the molecular weight of the unknown for each trial. Show all pertinent calculations for trial 1.
2. If the outside of the flask is not dried after vaporizing the liquid, will the unknown's calculated molecular weight be too high or too low? Explain?
3. Suppose the atmospheric pressure is assumed to be 1 atmosphere instead of its actual value. How will this error affect the molecular weight of the unknown? Explain.
4. What is the percent error for the molecular weight if the false assumption in question 3 is made? Show your work.
5. If the vapor's volume is assumed to be 125 mL instead of the measured volume, what is the percent error of the unknown's calculated molecular weight? Show your work.
6. If all of the unknown does not vaporize in the 125 mL Erlenmeyer flask, will the reported molecular weight be too high or low? Explain.
7. Suppose the thermometer is miscalibrated and reads 0.2 ^{o}C higher over its entire temperature range than actual. Does this affect the unknown's molecular weight? Explain.
8. An inherent error in determining the unknown's molecular weight is the design of the apparatus: the boiling water bath does not envelop the upper part of the Erlenmeyer flask. What can you do to minimize the error? How does this error affect the unknown's reported molecular weight?. Explain.
9. Determine the class average and the standard deviation.
10. Obtain the identity of the unknown from your instructor and determine your relative error.
11. The Van der Waals equation (given below) corrects the Ideal Gas Equation for the volume of the gas particles (b) and the intermolecular forces of attraction between the particles (a). Using the density of the liquid and its molecular weight, calculate an approximate value for b. With this value of b and your data solve the Van der Waals equation for the variable a,. What units will be attached to both a and b.
12. From the class values listed on the chalkboard for the molecular weight of the unknown liquid determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the and the standard deviation.
13. Summarize all calculations in a nice neat table.
Objectives
Discussion
[This laboratory exercise was adapted from an article in the March 1996 edition of the Journal of Chemical Education, page 257, by John T. Wood and Roberta M. Eddy.]
The endpoint of an acid base titration is usually signaled by a visual indicator. Two of the most common visual indicators are phenolphthalein and bromthymol blue. However, visual indicators are not the only means to signal the endpoint of a titration. Just as a visual indicator changes color in response to a change in the pH of a solution, an olfactory indicator will change either its odor or odor intensity as the pH changes. Vanillin and eugenol have previously been identified as olfactory indicators. The purpose of this experiment is to investigate the odor of raw onion for its usefulness as an acid base indicator.
The odor and flavor of raw onions are due to sulfur containing compounds. Many compounds are present but the principle ones are: 1propenyl thiosulfinate, propyl thiosulfinate, and methyl thiosulfinate.
The experiment will consist of three parts. The first part will attempt to identify which solution, acid or base, causes the odor of onion to be most intense. The second part will consist of extracting the indicator from onions. The actual titration will be done in the third part.
Procedure part A: Indicator Investigation
1. Contaminate three pieces of cloth with the odor of raw onion. This can be accomplished two ways. One method is to rub each piece of cloth with a fresh raw onion until the odor is very strong. The second method is to place three pieces of cloth in a sealed plastic bag containing chopped onion in a refrigerator overnight.
2. Fill three 400 mL beakers so that one beaker contains 200 mL of 0.1 M NaOH, another contains 200 mL of distilled water, while the third contains 200 mL of 0.1M HCl. Place one of the pieces of cloth prepared in step 1 above into each beaker. Stir with a stirring rod for 3 minutes.
3. With forceps lift out each piece of cloth and rinse with distilled water to remove traces of NaOH and HCl. Smell each piece of cloth and note the presence or lack of the odor of onion.
Procedure part B: Indicator Extraction
1. Place about oneteaspoon of chopped onion in about 30 mL of the base to be titrated (0.1 M NaOH). Let the chopped onions soak for about 1015 minutes with occasional stirring.
2. Filter using a "fast" filter paper like Whatham #1.
Procedure part C: The Titration
1. Fill a buret with 0.1 M HCl. Accurately measure out exactly 10.00 mL of the onionindicatorladen (OIL) base prepared in part B above into a Erlenmeyer flask. Using your observations in part A above, titrate until the correct olfactory endpoint is reached.
2. To a second sample of the OIL base add 2 drops of phenolphthalein. Repeat the titration until the visual endpoint of phenolphthalein has been reached. (Note: Comparisons will be easier if you make sure that the volume of base used in each titration is the same.)
Questions
1. Determine the molar concentration of the base for each titration.
2. Determine the class values listed, determine the average concentration of the base for each indicator.
3. Compare the relative effectiveness of the olfactory indicator to the standard visual indicator, phenolphthalein. (You should mention in your discussion not only whether the olfactory indicator is useful, but the relative error, if any, to be expected in the procedure.)
4. Lets assume that in step 2, part C untreated base was used instead of the OIL base. Why would a comparison now be more difficult?
Objectives
Discussion
When a kernel of popcorn "pops', that is, explodes, the explosion is due to the vaporization of water in the kernel of popcorn The pressure needed to rupture the kernel can be approximated using the ideal gas law: PV=nRT. The pressure can be calculated if we know the volume of a kernel, the mols of water present in a kernel of popcorn, and the temperature at which the kernel "popped", that is exploded. Two assumptions will have to be made in order to calculate this pressure:
1. the water vapor behaves as an ideal gas.
2. the temperature of the air is the temperature needed to "pop" a kernel of corn.
Procedure
1. Count out 100 kernels of popcorn and determine the average mass of a single kernel.
2. Count out a second batch of 100 kernels and determine the average volume of a kernel by water displacement. Do not use this batch of kernels in the following steps.
3. Pop 100 kernels of popcorn in a hot air popper. Measure the temperature of the air as the corn is popping.
4. Mass the kernels that popped and determine the average mass of a popped kernel. The difference between this value and the average mass of an unpopped kernel is the amount of water that vaporized.
Questions
1. Use the Ideal Gas Law to determine the pressure generated when a kernel popped.
2. One of the assumptions of the Ideal Gas Law is that no forces of attraction between gas particles exist. Water is a polar substance and its molecules probably have very high forces of attraction. How would these forces of attraction affect the answer calculated above. That is, would you expect the calculated pressure to be lower or higher than the actual pressure. Explain.
3. Another assumption of the Ideal Gas Law is that the molecules of a gas have no volume, that is, are point masses. At low pressures real gasses approximate ideal behavior. However, the pressure inside of the kernel is very high and the volume of the water molecules will start to cause significant deviation from ideal behavior. How is the calculated pressure affected by the volume of the water molecules. That is, would you expect the calculated pressure to be lower or higher than the actual pressure. Explain.
4. How would the caculated pressure be affected if the temperature of the hot air was higher than the temperature needed to pop the kernels.
5. From the class values listed on the chalkboard determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average and the standard deviation.
Objectives
Discussion
[This laboratory exercise was adapted from an article by Gary W. Rice "Potentiometric and Photometric Methods for Determining the Solubility of Lead Iodide" in the Journal of Chemical Education, May 1990]
Lead (II) iodide is a slightly soluble salt. Its solubility can be determined by measuring either the lead ion concentration or the iodide ion concentration in a saturated solution. In this lab we will use a redox titration to measure the iodide ion concentration. This investigation is more involved than past investigations because not one but two redox reactions will be occurring each with its own endpoint. Data will be collected and a photometric titration curve of absorbance vs mL of reactant plotted in order to determine the endpoint of the titration. The redox reactions are:
2I^{1 }+ 2Ce^{4+ }> I_{2} + 2Ce^{3+}
I_{2} + 2Ce^{4+ }+ 4Cl^{1 }> 2ICl_{2}^{1 }+ 2 Ce^{3+}
As a saturated solution of PbI_{2} is titrated with Ce^{4+} the concentration of free iodine (I_{2}) will first rise and then begin to fall. The change in concentration of the I_{2} can be monitored by measuring the change in absorbance as the concentration of the iodine changes. Absorbance changes will be monitored at a wavelength of 435 nM. Ideally a plot of Absorbance vs mL of Ce^{4+} added should look like the following:
In theory either endpoint can be used in determining the concentration of iodide ion present. However past experience in this class has indicated that only the first endpoint is reliable.
Procedure
1. Prepare solid lead iodide by mixing together about 10 mL of 0.1M Pb(NO_{3})_{2} and 20 mL of 0.1M KI. Centrifuge the precipitate and wash several times to insure that it is free of all electrolytes. Let the solid lead iodide stand in contact with about 150 mL of water for a week before proceeding. The solution should be shaken every day. (Alternatively, you can heat the solution to about 8085 °C to dissolve most of the lead (II) iodide. Then let the solution stand for 24 hours so that the excess PbI_{2} precipitates out. When filtered the solution should now be saturated in PbI_{2} at room temperature.)
2. Extract 50.0 mL of the saturated PbI_{2}, and place in a beaker. Be careful that no solid PbI_{2} is transferred. Filter if necessary. To the PbI_{2} solution add 50.0 mL of 2 N HCl. Titrate the saturated PbI_{2} solution with 0.05 M Ce^{4+} by adding 0.50 mL portions. After adding each portion of the titrant transfer enough of the solution to a cuvette and measure the absorbance at 435 nM. Return the iodide solution to the beaker before adding the next portion of titrant.
Questions
1. Make a plot of absorbance vs. mL of titrant added and determine the two endpoints.
2. Determine the mmols of titrant added and the mmols of iodide ion present in the saturated PbI_{2} solution.
3. Calculate the K_{sp} of PbI_{2}.
4. How would the calculated value of K_{sp} be affected if a small solid piece of PbI_{2} were accidently transferred to the beaker before the titration was done? Explain.
5. Ideally each absorbance reading should have been corrected for the dilution that took place because of the titrant. Extra Credit: Make these corrections and recalculate the K_{sp }. In your opinion does the dilution have much effect on the calculated K_{sp }?
6. From the class values listed on the chalkboard for the K_{sp} of PbI_{2} determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the K_{sp} of PbI_{2} and the standard deviation.
7. Summarize all calculations in a nice neat table.
Spectrophotometric Determination of Mercury
Objectives
Discussion
[This laboratory exercise was adapted from an article by Noel S. Murcia, Eric G. Lundquist, Steven O. Russo, and Dennis G. Peters ("Quincy Meets Perry Mason: An Experience in Chemistry and Law") in the Journal of Chemical Education, July 1990]
Both lead(II) and mercury(II) will react with dithizone (diphenylthiocarbazone) to form complex species which will absorb light in the visible portion of the electromagnetic spectrum. When these complex species are dissolved in methylene dichloride (dichloromethane) they absorb at wavelengths of 522 nM for lead(II) and 470 nM for mercury(II). As in previous labs, the spectrophotometer must be standardized with solutions of known concentration.
Procedure
1. Prepare a stock solution of Hg^{2+ }ions by dissolving 4.06 mg of HgCl_{2}, 35.7 mL of glacial acetic acid, and 13.9 mL of concentrated sulfuric acid in water and dilute to 500 mL volume. This solution contains 6.00 micrograms of Hg^{2+} per milliliter. Also, prepare a stock solution of dithizone by dissolving 10 mG of dithizone per liter of methylene dichloride. (Check with your instructor as these solutions may have already been prepared.)
2. Using a pipet place 15.00 mL of stock solution of the dithizone solution in each of 5 labeled 30 mL screw cap test tubes. Use pipets to deliver appropriate volumes of distilled water and the stock solution of Hg^{2+} into each vial, as indicated in the table below:
3. With each of the test tubes well capped shake vigorously for 30 seconds to ensure that the reaction between species in the two phases in complete. After the phases have separated draw off as much of the bottom methylene dichloride phase as possible and place in separate vials.
4. Add one gram of anhydrous magnesium sulfate to each vial, cap, and shake to dry the methylene dichloride phase. After the partially hydrated magnesium sulfate has settled transfer the clear liquid to a cuvette and measure the absorbance. Use methylene dichloride as the blank.
5. Secure the unknown from your instructor, treat in a similar manner, and measure its absorbance.
Questions
1. Make a calibration plot of absorbance Vs. Concentration.
2. Determine the concentration of Mercury in your unknown.
3. From your observations on the appearance of the methylene dichloride phase before drying How would the results of question 2 be affected if you forgot to dry it with magnesium sulfate after its extraction in step 3? Explain.
4. Summarize your results in a nice neat table.
Spectrophotometric Determination of an Equilibrium Constant
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
The magnitude of an equilibrium constant, K_{eq}, expresses the equilibrium position for a chemical system. The larger the equilibrium constant the more the equilibrium "lies to the right". The value of K_{eq} is constant for a chemical system at a given temperature. The chemical system studied in this experiment is:
Fe(H_{2}O)_{6}^{3+ }+ SCN^{1 }<====>Fe(H_{2}O)_{5}SCN^{2+ }+ H_{2}O
Since the concentration of water is constant in dilute aqueous solutions the above reaction is usually simplified to:
Fe^{3+ }+ SCN^{1 }<======> FeSCN^{2+}
The equilibrium expression of this chemical system is:
Five equilibrium systems will be prepared for the above reaction by mixing known concentrations of Fe^{3+ }and SCN^{1 }. Since the product FeSCN^{2+} is a deep, bloodred complex ion with an absorption maximum of 447 nM, its concentration can be determined spectrophotometrically. By knowing the initial concentrations of each of the reactants and the measured concentration of the product the equilibrium concentration of Fe^{3+} and SCN^{} can be calculated. The equilibrium constant can then be calculated by substituting into the equilibrium expression. The five calculated equilibrium constants will then be averaged to obtain the "best" value.
Before the spectrophotometer can be used to measure the concentration of the FeSCN^{2+} ion it must be calibrated with a set of standard FeSCN^{2+} solutions. These solutions will be prepared in part A by mixing a dilute solution of SCN^{1} with a relatively concentrated solution of Fe^{3+}. The large excess of Fe^{3+} will cause the equilibrium to shift nearly 100% to the right thus enabling you to assume that virtually all of the SCN^{1} reacted. The calibration curve will then be constructed by plotting the Absorbance of the solution at 447 nM. Vs. the concentration of the FeSCN^{2+} ion.
Procedure: Part A
1. Pipet 1, 2, 4, 6, and 8 mL of 0.00200 M KSCN into separate 100 mL volumetric flasks. Add to each flask 25 mL of 0.2 M Fe(NO_{3})_{3} and dilute to the 100 mL mark. Measure the absorbance of each solution at 447 nM using 0.2 M Fe(NO_{3})_{3} in 0.25 M HNO_{3} as a blank.
2. Calculate the equilibrium concentration of FeSCN^{2+} and construct a calibration curve by plotting Absorbance Vs. [FeSCN^{2+}].
Procedure: Part B
1. Prepare the following test solutions and determine the concentration of the FeSCN^{2+} ion by measuring the Absorbance of each solution at 447 nM and consulting the calibration plot you constructed above. The blank for this part should be a solution consisting of 50% 0.0002 M KSCN and 50% 0.25 M HNO_{3}.
Questions
1. What affect does a dirty cuvette have on the absorbance reading for a FeSCN^{2+} solution?
2. How does the error in question 1 affect the reported equilibrium constant?
3. In our calculations the solutions's thickness and the probability of light absorption by the absorbing species (a in the equation A=abc) are not considered. Explain.
4. Absorbance is defined as the negative log of the transmittance (fraction of light transmitted throught the solution). What is the transmittance for a solution that has an absorbance reading of 2?
5. How can the procedure be modified to obtain a more accurate reading for solutions having an absorbance of 2 or greater?
6. Over a period of time the transmittance of the blank solution may drift from its initial calibration. If the percent transmittance of the blank drifts to lower values how does this affect the
a) absorbance readings.
b) the calculated [Fe^{3+}]
c) the calculated FeSCN^{2+}
d) the calculated [SCN^{}]
e) the calculated K_{eq}
7. Why is 0.0002 M KSCN in 0.25 M HNO_{3} used as the blank in part B rather than distilled water or 0.25 M HNO_{3}?
8. From the class values listed on the chalkboard for the K_{eq} being studied determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the and the standard deviation.
9. Summarize all calculations in a nice neat table.
Spectrophotometric Determination of the pK_{a} of an Indicator
Objectives
Discussion
[This lab was adapted from "A Simplified Method for Finding the pK_{a} of an AcidBase Indicator by Spectrophotometry", Journal of Chemical Education, Vol. 76, p 395, March 1999.]
The magnitude of an equilibrium constant, K_{eq}, expresses the equilibrium position for a chemical system. Most acidbase indicators can be considered weak acids and their ionization can be represented as: . The equilibrium expression would be:
It is a common practice to quote the pK_{a} (the log of the K_{a}) of indicators and weak acids. In this experiment you will take a simplified approach to finding the pK_{a} of an indicator. This approach will not require that the spectroscope be standardized and an absorbance vs. concentration plot obtained as in past labs. The pK_{a} will be the yintercept of a plot of pH vs. log ; where A_{HIn} is the absorbance of HIn at , A_{In} is the absorbance of In^{} at , a_{HIn} is the absorbtivity of HIn, and a_{In} is the absorbtivity of In^{} .
Taking the log of equation 1 above gives:
Rearrangement give:
or
According to the BeerLambert Law
Substituting equations (3) and (4) into equation (2) gives:
If we now let equal the absorbance at and respectfully, for solutions where [In]=[HIn]. equations (3) and (4) give:
Dividing equation (6) by equation (7), rearranging , and comparing to equation (5) gives:
Combining equations (8) and (5) gives
It can be seen that a plot of pH vs. will be linear with a slope of one (1) and a yintercept equalt to the pK_{a} of the indicator.
The abosorbance of the indicator at both at and will be measured at five different pH values. A plot of pH vs. will be constructed. The value of k can be calculated by taking and absorbance readings at after the equilibrium has been shifted nearly 100% to the right by the addition of NaOH, and after the equilibrium has been nearly 100% to the left by the addition of HCl.
Procedure:
1.Your instructor will assign an indicator to us. Determine and by taking a sample of the indictor to which a few drops of 1M NaOH has been added, and measuring the absorbance of the solution at wavelengths in the range of 380 nM to 780 nM in 10 nM increments using distilled water as the blank. Construct a plot of absorbance vs wavelength. The wavelength at which maximum absorbance occurs is . In a similar way determine by replacing 1 M NaOH with 1 M HCl.
2. Prepare 250 mL of the indicator solution by consulting the following chart.
3. Transfer 50 mL of the solution to each of five beakers. Measure and record the pH of one of the solutions and adjust the pH of solutions by adding either 1 M NaOH or 1 M HCl to five different pH values in a range given by the following chart.
4. Measure the absorbance of each of the solutions prepared in step 3 above both at and . These wavelengths were determined in step 1.
5. Adjust the pH of one of the samples used in step 4 to 10 or 11 by adding a few drops of 6 M NaOH and measure the absorbance. This will be .
6. Adjust the pH of a second indicator sample in step 4 to 1 or 2 by adding a few drops of 12 M HCl and measure the absorbance. This will be .
7. Compute which will be the value of k in equation (9) above. Construct a plot of pH vs. and determine the pK_{a }. As stated in the discussion above the pK_{a} of your indicator will be the the yintercept. of this plot.
Questions
1. Determine the pK_{a} of your indicator.
2. Consult the article referenced in the discussion. How does your result compare to the accepted pK_{a}? What is the relative error?
Objectives
Discussion
[This laboratory exercise was adapted from one in Beran and Brady's Laboratory Manual for General Chemistry , John Wiley & Sons, 1982]
By definition one equivalent (or equivalent weight) of a substance is the amount of that substance which supplies or consumes one mol of reactive species. In acidbase chemistry the reactive species is the hydrogen ion (H^{1+}) while in oxidationreduction chemistry the reactive species is the electron. For example, in the following two reactions the equivalent weight of H_{2}SO_{4} would be 49 grams or 0.5 mol in the first reaction but 98 grams or 1 mol in the second. On the other hand, sodium hydroxide has the same equivalent weight in both reactions, one mol or 40 grams.
(1) H_{2}SO_{4} + 2NaOH > Na_{2}SO_{4} + 2H_{2}O
(2) H_{2}SO_{4} + NaOH > NaHSO_{4} + H_{2}O
In the first reaction one mol of H_{2}SO_{4} supplies 2 mols of H^{1+} to NaOH, therefore, onehalf mol of H_{2}SO_{4} or 49 grams is one equivalent. The conditions are different in the second reaction because sulfuric acid only "looses" one hydrogen so the equivalent weight of sulfuric acid is one mol or 98 grams. However, sodium hydroxide behaves the same in both reactions, that is, one mol of sodium hydroxide always "consumes" one mol of H^{1+}, so its equivalent weight remains the same at one mol or 40 grams.
In order to determine the equivalent weight of a substance you must know something about the reaction but it does not have to be balanced. Equivalents can help in the analysis of a substance when the balanced reaction is not known or cannot be written for whatever reason; because one equivalent always reacts with or produces one equivalent. (You should prove this to yourself by calculating how much sodium hydroxide is needed to react with 49 grams of sulfuric acid in each of the two reactions above. Do your calculations using traditional mol relationships and the one to one relationship for equivalents. )
In this experiment you will standardize a solution of sodium hydroxide by finding the normality of the solution. Normality is defined as the number of equivalents of solute per liter of solution. In order to standardize the NaOH solution you will react a solution of approximately 0.1 M NaOH with potassium hydrogen phthalate, also known as potassium acid phthalate. The balanced reaction is:
NaOH + KHC_{8}H_{4}O_{4} > NaKC_{8}H_{4}O_{4} + H_{2}O.
This reaction will allow you to find the equivalent weight of KHC_{8}H_{4}O_{4} which will be needed for your calculations.
Your next experiment will be "Equivalent Weight of an Unknown Acid". You will use the sodium hydroxide solution you are standardizing here in that laboratory exercise.
Procedure
1. Mass out about 500 mG. of KHC_{8}H_{4}O_{4} as accurately as possible, place in an Erlenmeyer flask and dissolve in distilled water.
2. Titrate with the NaOH solution to a phenolphthalein endpoint.
3. Repeat steps 1 and 2 so that a total of two trials are done.
Questions
1. Determine the number of equivalents of KHC_{8}H_{4}O_{4} used in each titration.
2. Determine the normality of the NaOH solution and place your values on the chalkboard.
3. From class values determine the average concentration of the NaOH solution and standard deviation.
4. Is it quantitatively acceptable to titrate all KHC_{8}H_{4}O_{4} samples with NaOH to the same dark red endpoint? Explain.
5. If the endpoint in the titration of the KHC_{8}H_{4}O_{4} with NaOH is mistakenly surpassed (too pink), what effect does this have on the calculated normality of the NaOH solution? Explain.
6. If a drop of NaOH solution adheres to the side of the flask during the standardization of the NaOH solution, how does this effect
a) the reported normality of the NaOH
b) If this solution of NaOH was used to titrate an unknown acid, would the calculated equivalent weight of the unknown acid end up being too high or too low?
7. "If 2 drops are good, 20 drops are better". Explain why this reasoning is not logical when adding phenolphthalein indicator for endpoint determination. (Hint: Keep in mind that indicators are either weak acids or weak bases.)
8. The mass of the potassium acid phthalate is measured to the nearest milligram, however the volume of water in which it is dissolved is never a critical concern. Explain why the addition of water to the flask is not critical to the analysis whereas its addition to the NaOH solution is?
9. From the class values listed on the chalkboard for the concentration of the NaOH determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the and the standard deviation.
10. Summarize all calculations in a nice neat table.
Thermometric Titration of an Unknown Acid
Objectives
Discussion
[This laboratory exercise was adapted from "The Stoichiometry of the Neutralization of Citric Acid", Journal of Chemical Education, November 1995]
In this laboratory exercise you will attempt to determine the number of acidic hydrogen's in a weak organic acid (HnA) by monitoring the heat produced during neutralization. Hess's Law states that if a process can be considered to be the sum of several stepwise processes, the enthalpy change for the total process equals the sum of the enthalpy changes for the various steps. For the neutralization of a weak acid the important steps are:
Success of this laboratory investigation rest on the assumption that enthalpies for the successive neutralizations of the unknown acid are similar and highly exothermic. This assumption will be correct if the enthalpies associated with the successive ionizations of the weak acid are similar.
The investigation will be simplified by keeping the total volume of the reactants equal and assuming that the specific heat of all trials of the reaction mixture is essentially the same as the specific heat of water (4.184 Joules per gram degree Celsius). This will allow us to "determine" the heat produced by monitoring the rise in temperature of the reaction mixture. The data collected will be treated by plotting the change in temperature vs either the volume of acid used or the volume of base added.
Procedure
You will mix together various volumes of 1 M NaOH and 1 M acid to which an indicator has been added. One or more trials will be assigned by your instructor to investigate. Your results will be shared with all members of the class and the composite data treated. Record the rise in temperature of the reaction mixtures for each assigned trial. The reactants are to be mixed in a styrofoam cup. If the reactants are not at the same temperature, use a weighted average of each reactant's temperature as the initial temperature of the reaction mixture.
Place a 20 mL sample of your reaction mixture in a stoppered, labeled test tube and display it as per the directions of your instructor.
Questions
1. Plot the rise in temperature vs the volume of base added.
2. Plot the rise in temperature vs the volume of acid added.
3. Plot the rise in temperature vs. the ratio of the volume of base added to the volume of acid added
4. From either of your graphs above, determine the number of hydrogens in each molecule of your unknown acid.
5. Secure the structural formula of the unknown acid from your instructor and determine the following:
a) the equivalent weight of your unknown acid.
b) the acid hydrogens in the molecule.
6. The heat of enthalpy for reaction (2) in the discussion above is of water produced. Using this value and the specific heat of water, , determine the enthalpy for the complete ionization of the unknown acid, that is, the for: ??????.
7. Write the ionization equation, including the enthalpy value, for each step in the complete ionization of the unknown acid.
In 8 A.D. Augustus Caesar exiled the Roman poet Ovid (Publius Ovidius Nasso 43 B.C.  17 A.D.) to the remote Black Sea town of Tomis (modern Constanta, Romania). Not only was Ovid isolated from the political, social, and intellectual center of his world, but also he had to endure a climate much harsher than that of Rome. To lament his exile, he wrote the Tristia, poems that detail his physical and emotional discomforts. In one part of the Tristia Ovid writes:
"and the wines stand stiff, jugless but keeping the shape of their jugs, and the people don't drink draughts of winethey eat pieces of it.
Our modern day temperature scales can be traced back to no earlier than the 17th century, therefore no temperature records from Ovid's time are available. However, chemists will immediately recognize that the phenomenon of the freezingpoint depression can be applied to estimate the temperature of Constanta, Romania, provided that the composition of the wine can be accurately estimated. Throughout the Roman Empire there were many kinds of wine consumed, some of which were diluted with water. However, the exact Latin words used by Ovid to describe the wine he writes about were "vina" and "meri". These words were used to describe undiluted wine.
Lab Report on The Winter of Tomis in 8 A.D.
Procedure:
1. Do some research on the composition of the undiluted wines made and consumed today and fill out the chart to the right. Do not report any "fortified" wines.
2. From the average percent alcohol calculated above determine the temperature that existed in Tomis when Ovid wrote the Tristia by calculating the expected freezing point of the wine.
3. Determine the freezing point of the wine assigned to you in the lab. Complete the chart below.
4. If you find that the absolute error is negative, that is, the wine froze at a lower temperature than you calculated in question 3, assume that the difference is due to the presence of sugar. Determine the molal concentration of sugar and complete the following chart. Support your answers with appropriate calculations on an attached page. Support your answers with the appropriate calculations
5. Why would your answer above for the percentage of sugar be too high/
Appendix
Sample Lab Report
Empirical Formula of a Compound
Purpose:
Discussion of Chemistry
The empirical formula of a compound represents the simplest ratio of atoms present in one formula unit or molecule of a compound. For example, a formula of MgCl_{2} indicates that in each molecule or formula unit of this compound one atom of magnesium will be found for every two atoms of chlorine. Alternatively, one could say that in one mol of magnesium chloride molecules one mol of magnesium atoms would be found for every two mols of chlorine atoms. The magnesiumchlorine ratio is 1:2 whether you are talking about one molecule or one mol of molecules. Therefore, if in a given bulk sample of compound AxBy it is found that the mol ratio of A:B is 3:2 the empirical formula of the compound is then known to be A_{3}B_{2}.
In this experiment the mol ratio was determined by measuring the mass of oxygen that combined with a given mass of magnesium by repeatedly heating the magnesium sample to constant mass. Heating to constant mass insures that all magnesium atoms in the sample have reacted. The increase in mass is equal to the mass of oxygen that combined. The number of mols of each element involved in the reaction is determined by dividing the mass reacted by the elements atomic weight.
Data Table
atomic weight of Mg 24.312
atomic weight of O 15.999
Trial I  Trial II  
mass of crucible (G)  20.00  20.000 
mass of crucible + Mg (G)  25.036  24.100 
mass of Mg (G)  5.036  4.100 
mass of crucible + product (G)  
after first heating  28.490  26.318 
after second heating  28.489  26.829 
after third heating    26.829 
mass of compound (G)  8.489  6.829 
mass of oxygen (G)  3.453  2.729 
mols of Mg  0.210  0.171 
mols of oxygen  0.216  0.171 
mols ratio (Mg:O)  0.972:1  1:1 
empirical formula  MgO  MgO 
Questions and Calculations
1. What is the percentage of magnesium in the sample?
Trial 1: (5.036 G / 8.489 G) X 100 = 59.32%
Trial 2: (4.100 G / 6.829 G) X 100 = 60.04%
2. What is the percentage of oxygen in the sample?
Trial 1: (3.453 G / 8.489 G) X 100 = 40.68 %
Trial 2: (2.729 G / 6.829 G) X 100 = 39.96 %
3. How would the results (empirical formula) be affected if not all of the magnesium had reacted with oxygen?
Since the number of mols of oxygen is directly proportional to the mass of oxygen that reacted, the number of mols of oxygen calculated would be lower which would result in a higher Mg:O ratio.
4. Show the calculations on how you arrived at the empirical formula for one of the trials.
Trial 1:
mols Mg = 5.036 G / 24.312 G per mol = 0.2071 mols
mols O = 3.453 G / 15.999 G per mol = 0.2158 mols
ratio of Mg to O = 0.2071 / 0.2158 = 0.9598 to 1.
This is close enough to a ratio of 1:1 to justify a formula of MgO.
Statistical Calculations
The relative (or percentage error) in a measurement is calculated by the following formula:
In this formula E represents the experimental value and A represents the actual or accepted value. Relative Error is used to by scientists to check methods and procedures. For example, suppose a new procedure had been developed for measuring the atomic weight of helium. The accepted value of He is 4.003. If this new procedure gave a value of 4.001 then the relative error would be 0.050%.
Note: The remaining statistical topics will perform calculations on the following sample data set:
Trial  Value 
1  5.32 
2  5.36 
3  5.41 
4  5.43 
5  5.66 
Whenever several measurements are taken of the same quantity, there is the possibility that certain measurements will lie "outside the normal range of the data". These values are called "Outliers" and once detected should generally not be used to calculate any of the statistical quantities discussed below. One of the methods used to detect outliers involves calculating the experimental Qvalue of the suspected data value and comparing it to a standard table of critical Qvalues. If the experimental Q value is larger the critical Qvalue obtained from the table then the suspected data value can be rejected. The experimental Qvalue is defined as the ratio given by the distance of the suspect data value from its nearest neighbor divided by the range of the data values. Experimental Qvalues are easier to compute if the data is first sorted as in the sample data above. The following formula can be used to determine the Qvalue for a suspect data point:
The values D, X_{H}, and X_{L} represent the distance of the suspect point from its nearest neighbor, the highest data value, and the lowest data value respectfully. For the sample data above the suspect point is 5.66 from trial 5. The experimental Qvalue for this point would be 0.676. To determine if this value is too high we compare it to the critical Qvalues obtained from the following table:
Confidence Limit

P=0.90  P=0.96  P=0.99 
Number of data values  
3  0.941  0.976  0.994 
4  0.765  0.846  0.926 
5  0.642  0.729  0.821 
6  0.560  0.644  0.740 
7  0.507  0.586  0.680 
8  0.468  0.543  0.634 
9  0.437  0.510  0.598 
10  0.412  0.483  0.568 
11  0.542  0.460  0.392 
12  0.522  0.441  0.376 
13  0.503  0.425  0.361 
14  0.488  0.411  0.349 
15  0.475  0.399  0.338 
16  0.463  0.388  0.329 
17  0.452  0.379  0.320 
18  0.442  0.370  0.313 
19  0.433  0.363  0.306 
20  0.425  0.356  0.300 
21  0.418  0.350  0.295 
22  0.411  0.344  0.290 
23  0.404  0.338  0.285 
24  0.399  0.333  0.281 
25  0.393  0.329  0.277 
26  0.388  0.324  0.273 
27  0.384  0.320  0.269 
28  0.380  0.316  0.266 
29  0.376  0.312  0.263 
30  0.372  0.309  0.260 
For a confidence limit of 90% (P=0.90) trial 5 can be rejected since its experimental Qvalue (0.676) is larger than the critical Qvalue for 5 trials (0.642). However, the value could not be rejected if one desired a higher confidence limit since the experimental Qvalue is less than the corresponding critical Qvalue at the higher confidence limits.
Note: The following statistical calculations are going to be done at a confidence limit of 90% (P=0.90) and since trial 5 has been determined to be an "outlier" at this confidence limit it will be rejected from those statistical calculations.
The mean of a set of values, which is denoted by the symbol x, is calculated by the following formula:
In this formula n represents the number of data points, and x_{i} represents a specific data value. Applying this formula to the sample data above one obtains a mean of 5.380. (If trial 5 is not rejected the mean would be 5.436.)
The median value of a series of measurements is the middle value. In order to find the median of a set of values the measurements must be sorted as the sample data above is. The median of the sample data is the value of trial 3 if trial 5 is not being rejected, that is, 5.41. If the median is being determined for an even number of trials the average of the two middle values can be taken. Since we are rejecting trial 5 the median would be the average of trials 2 and 3 which is 5.39.
The standard deviation, sometimes abbreviated with the Greek symbol sigma (s), gives an indication of how well the data is grouped around the mean value. A large standard deviation indicates that several values were significantly above and below the mean indicating low precision data, while a small standard deviation indicates that the values were quite close to the mean indicating high precision data. For example, both plots below show the results of two procedures, both of which attempted to measure the atomic weight of Ne. Both plots have a mean of 10, however, the data from procedure 1 is of higher precision because the standard deviation is only 1.61 vs. 2.45 for procedure 2.
The formula for calculating the standard deviation is given below:
When this formula is applied to the above sample data a standard deviation of 0.050 is obtained to a confidence limit of 90%. This means that 68% of the values lie within ±0.050 units of the mean or within the range of 5.330 to 5.430 (5.380 ± 0.050) for the above sample data. Plus or minus two standard deviations (5.380 ± 0.100) of the mean encloses about 98% of all data points.
The average deviation is easier to compute but less often used than the standard deviation. It is computed by the following formula:
In terms of the error curve the probability that another measurement lies within ±a of the mean is 57%. For the sample data given above the average deviation would be 0.04. This means that the probability of the next measurement falling in the range of 5.34 to 5.42 (5.38 ±0.04) would be 57%.
Questions? Comments?? 