UHC Gen Chem 2 Manual

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Table of Contents

Introduction to Honors General Chemistry Laboratory II iii

Safety in the Laboratory iv

Experimentation and Data Taking vi

Writing your Laboratory Report vii

Experiment #1: Molecular Mass of a Volatile Liquid 1

Experiment #2: Chemical Kinetics: Reaction Rates, Rate Laws, and ReactionMechanisms 6

Experiment #3: Quantitative Absorption Spectroscopy (Vernier Version) 12

Experiment #4: The Law of Mass Action: The Ionization of Picric Acid (VernierVersion) 18

Experiment #5: pH Measurements I: Acid-Base Equilibria and Titrations (VernierVersion) 23

Experiment #6: pH Measurements II: Buffers, Solubility, and the Common IonEffect (Vernier Version) 29

Experiment #7: Redox Activity Series 34

Experiment #8: Electrochemistry

38

Experiment #9: Reactions of Iron 45

Experiment #10: Equilibrium, Free Energy, and Metal - Ligand Equilibrium in aComplexation Reaction (Vernier Version) 49

Experiment #11: Aqueous Complexes of Nickel (II): Spectroscopy and Stability(Vernier Version) 56

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Introduction to Honors General Chemistry Laboratory II

Laboratory work gives you the kind of hands-on and sensory experience that no lectureor computer-simulated experiment can provide. You can never truly learn chemistryunless you do it. Observation in science is paramount, accurate observation and

manipulation of experiments requires foresight, care, and skill acquired throughpractice. Perhaps more important, good scientific work entails a scrupulous honesty inthe observation and recording of data. If you dont see the result you expect, recordwhat you dosee. The meter reads what it reads, not what it should read. An importantpart of each lab consists in discovering any sources of error: if your result was not whatyou expected, why? It could be that the error is in your calculations!

Each laboratory experiment is designed to illustrate a topic covered concurrently in thelecture. In Honors Chemistry, we are moving away from highly prescribed chemistryexperiments common in an entry-level course, not because we think you knoweverything already, but because you are clever enough to find out. Not every procedure

or reaction will be spelled out in detail, so you may have to do a bit more preparationbefore the lab beyond simply reading the experimental description, and afterexperimenting when it is time to analyze your data. In some of the labs, you will begiven the option of striking out on your own, by inventing based on what you havealready seen, further experiments to test some idea or result. In the second term, thereis more emphasis on instrumentation and graphical analysis of results

This manual is locally produced,making it exquisitely sensitive to your needs; it isrevised after every use. We always value, and usually act on, feedback from you andyour laboratory instructor. Please let your instructor know of any errors, inconsistenciesor other problems; we would also like to know what you liked about the labs, so we can

include more of that.

Dr. Peter E. SiskaSpring 2008

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Safety in the Laboratory

Playing with chemicals is popularly considered to be a somewhat dangerous activitypracticed by nerds. In fact, with a few simple precautions, it can combine learning witha lot of fun. Safe laboratory practice is also a matter of your own enlightened self-

interest, and concern for your fellow students.

Wear eye protectionwhenever you are working at the lab bench. Eyes are the mostvulnerable to permanent damage from chemicals and fumes. Goggles that fit the facesnugly are required. If you ordinarily wear eyeglasses, wear the goggles over yourglasses. Do not wear contact lenses of any type in the lab, even with goggles. Youmay remove your goggles only when away from the bench to do calculations or toconsult with your laboratory instructor.

Come to lab appropriately dressed. Do not wear your best clothes. Wear loose clothingthat minimizes areas of exposed skin. Shorts, skirts, bare midriffs, shoes that do not

cover your entire foot, or other overly brief clothing are forbidden. Long hair must betied back.

Instead of attempting to list all possible hazards here, safety notes will be given at thebeginning of each experimental write-up, and your instructor will remind you of anyspecific precautions when first briefing you. Try to anticipate the hazards as you lookover the procedure before coming to lab. Accidents in the lab are usually due to:

Firedue to volatile solvents and other chemicals.

Spills and splashingof concentrated chemical reagents.

Cuts, usually from broken glassware.

Burns, usually from picking up hot apparatus with bare hands.

Toxic Fumes

Electric shock

Some general safety rulesthat you must be conscious of during everysession are:

1. Do not carry a stock bottle of flammable solvent or reagent to the lab bench.

Take just what you need at the reagent bench when you need it, and return to thelab bench and use all of it immediately. Flammable reagents must be kept awayfrom flames or heat.

2. Do not handle or touch any solid or liquid chemical with your bare hands. Cleanup chemical spills at once; ask your instructor for help.

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3. If any dangerous reagent spills or splashes on you or your clothing, wash itimmediately with plenty of water. For large scale mishaps, remove contaminatedclothing and make use of the safety shower in the corridor. An eyewash stationis also in the corridor; make sure you know where it is and how to use it.

4. Do not exert great force on glassware; it will break and cut you. Do not clean upbroken glassware with bare hands. Dispose of broken glass only in appropriatelylabeled containers.

5. Always handle hot crucibles, beakers, etc. with a dry pad or tongs.

6. Do not hide any accident. Tell your instructor at once and cooperate fully inmaking an official report and obtaining follow-up treatment, if needed.

7. Do not smoke, eat, or drink anywhere in the lab. Clean up your work area andwash your hands after completing the lab.

You will be asked to sign a sheet verifying that you are aware of these and other safetyrules before your first lab session.

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Experimentation and Data Taking

All experimental data must be recorded in a laboratory notebook with the facility formaking a removable carbon copy. The recommended notebook is the Hayden-McNeil50-set notebook, in stock at the Book Center. Each original page of data must be

signed and dated by you, and initialed by your instructor. Carbon copies are to beturned in to your instructor at the conclusion of each lab session.

Each experiment has been planned to fit into a three-hour session, includingcalculations. In order to make this possible, you must read the lab write-up beforecoming to lab and prepare one or more pages of your notebook, with your name and thetitle of the experiment at the top, and labeled spaces for recording observations andnumerical data such as masses, volumes, temperatures, etc. This preparation isrequired before you can begin experimentation. Please observe the following practices:

(a) Recording of data is to be in ink onlyin your notebook; temporary notes on

scraps of paper are forbidden. If you must change an incorrect datum, cross outthe original entry with a single ink line, and enter the new datum beside it.

(b) Each recorded numerical value must reflect the precision of the measurement bymeans of the appropriate number of significant figures.

(c) When weighing out samples precisely, do not waste time trying to adjust themass of the sample to exactly the suggested amount. For example, theexperiment might call for you to weigh 0.5 g of calcium metal to 4 decimalplaces... Since only one significant figure is specified, any mass from 0.45 g to0.55 g will do, and the important fact is the precise mass of that sample (e.g.,

0.4623 g, 0.5127 g, etc.) Your instructor would regard a recorded mass of0.5000 g with suspicion.

In most cases, the success of the experiment hinges on the cleanliness of yourcontainers and the purity of the reagent chemicals. A contaminated reagent caninvalidate an entire sessions work. Be clean and neat. Take care not to contaminatethe reagent bottles, balances, or other equipment, and recap opened bottlesimmediately after use.

If you have made an egregious error that invalidates your results, there will often betime to repeat at least part of the experiment. Your instructor can advise you as to

which sections should be repeated.

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Writing your Laboratory Report

Good laboratory reports are neat, brief, and to-the-point. They should consist of:

(a) Your name, date, title of the experiment, and the name(s) of your lab partner(s).

(b) A neatly organized tabulation of your experimental procedures, data, andobservations.

(c) Any required calculations, including error assessment.

(d) Conclusions

(e) Answers to the questions at the end of the experiment.

A copy of your lab notebook data sheet(s) will already be in the possession of yourlaboratory instructor. Your instructor in this course is generally an advanced graduatestudent with extensive teaching experience; and is given some freedom to requireanything additional, such as a statement of purpose, etc., and to assign credit for thecomponents of your report, including your data sheet. The instructor may also includeassessment of your state of preparation for the lab as part of the report grade. Makenote of other requirements below.

NOTES:

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Experiment #1: Molecular Mass of a Volatile Liquid

Safety

The samples to be used in todays experiment are organic liquids, whose vapors are

toxic. Avoid inhaling their vapors and keep your apparatus under a fume hood.Some of the unknowns are highly flammable. Keep all samples and reagent bottlesaway from open flames. Take care when handling hot glassware.

Background

In the century or so since the ideal gas law has been considered a reliable tool of thechemist, its major use has been in the determination of molecular masses. Longbefore its universal acceptance however, Jean Baptiste Andre Dumas developed amethod for extending the use of the ideal gas law to substances that are not gassesat room temperature. Dumas, the most well known French chemist of the early

nineteenth century, exploited a well-established fact. When a liquid is boiled away todryness in a small-necked flask, its vapor fills the flask and drives out the air. If theflask is then sealed, cooled, and weighed, the mass due to the recondensed liquid isequal to the to the mass of the vapor that occupied the flask when it is hot. Knowingthe temperature and volume of the hot flask, and the external (atmospheric)pressure, one can use the ideal gas law to find the molecular mass of the liquid. Tomaintain a constant high temperature, Dumas employed boiling water, oil, andmolten metal baths. The hotter baths were reserved for less volatile liquids. Intodays experiment, your unknowns will be highly volatile liquids, for which a boilingwater bath is adequate.

Pre-Lab

You should prepare your laboratory notebook before coming to lab using theguidelines on page xx. Be sure to reserve room for mass and volume readings, theatmospheric pressure, and any other data. You should also leave room for anyobservations. What you wish to record is up to you, but in science it often happensthat seemingly minor details and barely noticeable features are important.

You also need to estimate the volume of liquid to use in this experiment. Assumethat the molecular mass of the liquid is 100 g/mol and the density is 1 g/mL. Makesure there is enough liquid to fill the flask with vaporat least three times over,

rounding the volume to the nearest mL. This surplus guarantees that all air will beswept out of the flask when the sample boils away.

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Procedure

Make a small square of aluminum foil with a pinhole in the center to cap the mouthof a 250 mL Erlenmeyer flask. Measure the diameter of the pinhole to the nearest0.1 mm. Precisely weigh the dry flask and foil cap before setting up the experiment.

Set up a water bath by filling an 800 mL beaker ~2/3 with water. Add boiling stonesto the water and place the beaker on a hot plate inside of a fume hood. Heat thewater to boiling while monitoring the temperature with an alcohol thermometer.

Obtain the letter of your unknown liquid from your instructor. Using thoroughly dryglassware, measure out the estimated volume of your unknown and pour it into theflask. Cover the mouth of the flask with the foil, making sure to center the pinhole.Crimp the foil around the rim of the flask to seal. Clamp the rim (not the neck) of theflask with a two-jaw clamp and immerse it as fully in your water bath. Fasten theclamp to a vertical monkey bar inside the fume hood. Take care that water does notenter the flask or wet the aluminum foil.

Observe the liquid sample in the flask. When it has boiled away entirely, wait 30seconds, and then transfer the flask to an 800 mL beaker halfway filed with roomtemperature water. Keep the flask in the room temperature water for about oneminute, do not remove the foil cap. Remove the flask from the water, unclamp, andthoroughly dry the outside of the flask. Take care that no water remains in the foldsof the foil cap. Weigh the flask with the cap in place.

Repeat the experiment twice more with the same unknown. After the last run, obtainthe actual volume of the Erlenmeyer flask by using a graduated cylinder to measurethe precise volume of water needed to fill it to the top.

Further Exploration

After carrying out the above experiment and calculation of molecular mass (MM)three times, you are free to explore further. It would be nice if you could do thisexperiment on something known, like good old H2O! Decide whether the sametechnique, or some modification of it, is usable for water vapor (Hint: A salt-watersolution boils at a higher temperature than pure water). Try a modified experimentwith water as the volatile liquid, and see how close you come to MM= 18.0 g/mol.

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Report

Your data should allow you to calculate MMfrom the ideal gas law. A correction toyour results for the mass of air excluded from the flask by the equilibriumvaporpressureof the unknown should be made to improve the accuracy of the results.

This correction is analogous to that which is made to correct for the presence ofwater vapor in gas collected over water, using Daltons Law of Partial Pressures. Inthe initial weighing, the flask is filled with air, but after vaporization part of the air hasbeen displaced by the vapor in equilibrium with the liquid. Therefore, the finalweighing should be corrected by addingthe mass of air that has been displaced bythe vapor. The vapor pressure at 25C for your unknown can be obtained from yourinstructor. The molecular mass of air is taken to be 29.0 g/mol and the totalpressure in the flask can be assumed to be that of the atmosphere (in a futureexperiment, you will learn to measure the vapor pressure of a liquid).

Calculate a value of MMfrom each run, uncorrected for the vapor pressure of theunknown and average with error. You can calculate the error by first calculating thestandard deviation (sx):

sx=

sPN

i=1(xi x)2

N 1

Nrepresents the number of trials which is 3 for this experiment. To get the error, you

should divide sxbyN. Do the same after making the correction for each run.

Questions

1. A student t-testcan be used to determine if two sets of data are significantlydifferent. To perform a students t-test on your corrected (M1) and uncorrected(M2) sets of molar masses, you will calculate a t value using the equation

tcalculated =|x1 x2|

spooled

r M1M2

M1 +M2

Here, x1and x2are the averages for the for the corrected and uncorrected sets

of molar masses, M1and M2are the number of data points in each set of molarmasses, and spooledis

spooled= s2

1(M1 1) + s22(M2 1)

M1+ M2 2

Note that s1and s2are the standard deviations for the corrected and uncorrectedsets of of molar masses. Compare your calculated value for t to the t values inTable 1 for M1+ M2- 2 degrees of freedom. If your calculated t value is greater

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than the tabulated t value, then your corrected and uncorrected molar massesare considered be significantly different. Was the vapor pressure correctionsignificant?

2. A student inadvertently terminated the heating of the flask before all liquid had

boiled away. How would this affect the value of MMobtained? Show yourreasoning using an appropriate equation.

3. In this experiment, you rely on the rapid diffusion of air through the pinhole inyour foil cap. The kinetic molecular theory formula for effusion (not diffusion)through a hole of area A. The equation:

Z= 14[X]vA

can provide an upper limit to the diffusion rate. For a pressure behind the hole of1.00 atm of air and a temperature of 100C, find the rate Z in molecules of air

per second, and the mass flow rate dm/dtin grams of air per second. Accordingto this result, how long will it take for all of the air in the flask to escape from thepinhole, assuming a constant rate? What is wrong with this calculation and howcould it be improved?

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Explanatory Notes

The early results Dumas obtained with his new method were controversial, andcaused growing doubt over the validity of Avogadros hypothesis, although Dumashimself firmly believed it. For example, when Dumas experimented with bromine

liquid and iodine crystals, he found MM= 160 and 254 g/mol respectively, whilegravimetric analysis of silver bromide and iodide yielded half these values. As wenow know, this was evidence for the diatomic nature of halogen vapor particles, andnot a counterexample to Avogadros idea. Experiments with sulfur and phosphorusgave inconsistent results. Later, mass spectrometric analysis of the vapors of theseelements have shown that they consist of several different molecular species, withthe composition depending on T(e.g. in sulfur vapor there are S, S2, S3, and S8molecules in variable proportions).

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Experiment #2: Chemical Kinetics: Reaction Rates, Rate Laws, andReaction Mechanisms

Background

The law of mass action, which you will explore later this term in laboratory, had itsorigins in observation of reaction rates for forward and reverse reactions. Studies ofthe reaction rate of the hydrolysis of the ester ethyl acetate

(1) CH3COOCH2CH3+ H2O CH3COOH + HOCH2CH3

by Berthelot and St. Gilles in 1860 showed that the rate of reaction dependedlinearly on the concentrations of ethyl acetate and water. This was later expressedby vant Hoff in the form of a rate law.

d[CH3COOCH2CH3]

dt =k

[CH

3COOCH

2CH

3][H

2O

]where the time derivative expresses the rate of loss of water and k is a rateconstant. The reaction is said to be first orderin both ethyl acetate and water, andsecond orderoverall. The order refers to the power of the concentration appearingin the rate law. Berthelot and St. Gilles also studied the reverse reaction and foundits rate to be first order in both CH3COOH and HOCH2CH3. When equilibrium in anequimolar system was reached, the products were found to comprise 34% of thereaction mixture, no matter in which direction the reaction was studied. With theconcept of the rate law, this meant that the reverse reaction was four times fasterthan the forward for given concentrations of starting material, and that, to balance

the rates at equilibrium, the products had to be present in lower concentration.

The reaction was found to be greatly accelerated by increasing the temperature ofthe reaction mixture. The rate law remained unchanged, implying that the rateconstant kdepends strongly on temperature. The form of this dependence was firstproposed by our old ionic friend Svante Arrhenius as

k = AeEa/RT

where Ais called the pre-exponential factor, and Eathe activation energy. Inmodern times we interpret A as an attempt rate which is a measure of how often the

reagents encounter each other. Ea, which appears in the exponential Boltzmannfactor, can be interpreted as an energy hill that the reagents must climb before theycan succeed in reacting. At higher T, the Boltzmann factor grows, increasing thefraction of reaction attempts that succeed. The form of the Arrhenius equation leadsto a vant Hoff - like expression for the dependence of kon T.

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d ln k

d( 1T)=

Ea

R

A plot of ln kversus1

Twill be linear be linear with a slope of

Ea

R .

Vant Hoff himself later studied the acid catalysisof Reaction (1), finding that theoverall reaction rate could be enhanced many - fold by adding minute amounts of astrong acid, but that the composition of the system at equilibrium was unaffected.This led to speculation concerning the mechanismby which the acid catalystperforms its magic. From that day, the question of mechanism has dominateddiscussions of inorganic, organic, and biochemical reaction rates. A mechanism isformally as a sequence of elementary stepsby which the a reaction is thought toactually proceed. That is, each step is imagined as an actual occurrence involvinganywhere between one and three molecules. The steps are themselves reactions,and when summed, they must yield the overall balanced chemical equation. For

example, the mechanism for the acid catalyzed version of Reaction (1) may consistof two steps, the protonation of the ester followed by the a reaction of the protonatedester with water.

Step 1: ester + H+!esterH+

Step 2:esterH++ H2O!acid + alcohol + H+

When these two reacts are added, they give Reaction (1), the H+catalyst and theesterH+reaction intermediatecanceling out. Note that the catalyst H+is notconsumed in the reaction, hence a single proton can catalyze a large batch of

reactions. Often, a catalyst will change the rate law as well as the rate constant.The rate of a multi - step sequence is greatly determined by the slowest, or ratelimitingstep. In the above example, if the protonation step limits the rate, the ratelaw will become first order in H+and zero orderin H2O. Kinetic data, such as whatyou gather today, can rule out incorrect proposed mechanism. However, mereagreement between a reaction mechanism and an observed rate law never provesthe correctness of the mechanism; this is a dictum often neglected in the rush to findexplanations for chemical happenings.

In todays experiment, you will examine the kinetics of an aqueous redox redoxreaction, the oxidation of iodide ion by peroxydisulfate.

(2) 2I-(aq) + S2O82-(aq)!2SO4-2-(aq) + I2(aq)

As in most redox reactions, the position of the equilibrium for Reaction (2) lies far tothe right, so that you do not have to worry about the occurrence of the reversereaction. The easiest molecule to use as a monitor of the rate is I2, since it forms ablue complex with starch indicator. Normally you would see a gradual buildup ofblue color as the I2product accumulates. With a spectrophotometer and fast

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electronics, you can monitor the time profilefor the reaction ([I2] as a function oftime). In the absence of these tools, you can use another, very fast reaction thatconsumes the I2as a chemical timer or clock. In this case, the reduction of I2bythiosulfate (S2O32-), which as you may recall from last term, was the basis ofiodometric redox analysis, serves this purposes nicely.

(3) I2(aq) + 2S2O32-(aq)!2-(aq) + S4O62-(aq)

Reaction (3) will consume the I2as fast as it is made by Reaction (2), therebypreventing the formation of the blue complex. Thus, if you add a very small, knownamount of S2O32-to the reaction mixture, the solution will flash blue after a certaintime when the supply of S2O32-is exhausted. At that time, the moles of I2producedby reaction (2) is known to be half that of S2O82-initially present due to thestoichiometry of of Reaction (3). This gives you a single point on the time profile. Bykeeping [S2O32-]0constant for each trial, you always detect the time interval !tduring which a fixed amount of I2is produced (![I2]). This allows you to find the

rate,

[I2

]t . By keeping [S2O32-]0small, you insure the measurement of the rate

near the very beginning of the reaction, the so-called initial rate.

You will first attempt to determine a rate law for Reaction (2), by assuming thegeneral form

d[I2]

dt = k[I]x[S2O

2

8 ]y

and finding the orders xand yby systematically changing first [I-]0and then [S2O82-]0

and observing their effect on the reaction time and rate. For example, if the reactiontime is halved (or the rate doubles) when you double the [I -]0, x = 1 and the reactionis first order in I-. These measurements then allow you to determine the rateconstant k, and may help you to distinguish between possible proposed mechanismsfor this reaction. You will also observe the effect of varying the temperature, andtesting a possible catalyst for the reaction.

Procedure

A. Finding the rate law and rate constant

Before you begin, measure and record the ambient room temperature.

The following stock solutions will be available to you: 0.20 M KI, 0.0050 MNaS2O3containing 0.4% starch indicator, and 0.10 M K2S2O8. Note that the [KI]and [K2S2O8] are such that a stoichiometric mixture for Reaction (2) will resultfrom equal volumes. To insure uniform conditions for all runs, use a constanttotal volume of 50.0 mL in a 125 mL or 250 mL Erlenmeyer flask. In addition,maintain a [S2O32-]0of 0.0010 M for each run. Use the dilution equation(M1V1= M2V2) assuming additive volumes to determine the volume of each

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solution to add. All solutions should be delivered from burets, and the totalvolume of solution should be maintained with boiled distilled water whennecessary. Since you must time the reaction, you should first make up a solutioncontaining the iodide, thiosulfate, and any water needed, and then add theperoxydisulfate from a separate beaker or flask. Begin timing as soon as the

peroxydisufate is added, using a magnetic stirrer to insure thorough mixing.

You need a standard or reference reaction mixture, so that other runs, in whichinitial conditions are varied, may be compared to it. For this, use a stoichiometricmixture (equal volumes of KI and K2S2O8) with no dilution by water. Performthree runs or more until reproducible reaction times are obtained. These data willalso be used to determine the rate constant at ambient temperature.

To determine the rate law, vary [I-]0and [S2O82-]0one at a time. In each case, usea smaller volume of reagent solution and dilution with water to reduce itsconcentration. Perform (at least) three runs in which the initial concentration of agiven reagent is varied, down to 1/4 of its reference concentration; that is at leastsix runs in all.

B. Temperature dependence of the rate constant

For these runs, use your reference reaction mixture. Before adding theperoxydisulfate, warm the I-/S2O32-using a hot plate, or cool it using an ice waterbath. Measure reaction times for at least three different non-ambienttemperatures. Avoid boiling or very high temperatures since the I2- starchcomplex becomes unstable and difficult to form. Measure the temperature justafter timing the reaction.

C. Testing a possible catalyst

Certain metal ions, which can act as Lewis acids, often make effective catalysts.Try a reference reaction mixture with one drop of 0.1 M CuSO4added just beforeadding the peroxydisulfate. Repeat once more to test reproducibility.

D. Further exploration

Upon showing your data tables to your instructor, and analyzing a subset of datato obtain estimates for x, y, and k, you may explore further. Here are a fewsuggestions.

1. Does the catalyst change the rate law? Do a few more runs to examine thisquestion.

2. Is Cu2+unique to its catalytic activity (if any)? Try a drop of some 0.1 MArrhenius acid, or some other transition metal ion.

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3. How good a clock is the thiosulfate reaction? Does the reaction keep goodtime? Devise a few runs to test the clock.

4. Use your calculated rate constant to predict the reaction time for acombination of initial concentrations that you did not measure. Then carry out

the experiment, thereby testing the predictive power of your experimental ratelaw.

Report

Report your reaction timing data as a table with columns indicating the run number,temperature, volumes used, initial concentrations, reaction times, and reaction rates.Make two plots of log (rate) vs. log [ ]0, one for each set of runs that varied [I -]0and[S2O82-]0. What do the slopes of these plots represent? Find values of x and y forthe rate law from these plots, rounding them to the nearest half-integer to reflectexperience with simple reactions such as Reaction (2). Describe the order ofreaction with respect to each reagent and overall. Use your reference run along withthe values of x and y to complete rate constant with proper units. You shouldcalculate an average kwith error.

Use your temperature dependence data to prepare a plot of ln kversus1

T(K)

(an Arrhenius plot) and determine the activation energy Eafrom it. Also report thefactor by which the Cu2+catalyst affects the reaction rate.

Consider the following three possible mechanisms for the uncatalyzed reaction:

I. Step 1: I-(aq) + S2O82-(aq)!SO42-(aq) + SO4I- (slow)

Step 2:SO4I-(aq) + I-(aq)!I2(aq) + SO42-(aq) (fast)

II. Step 1: 2I-(aq)!I22-(aq) (slow)

Step 2: I22-(aq) + S2O82-(aq)!I2(aq) + 2SO42-(aq) (fast)

III. Step 1: 2I-(aq) I22-(aq) (fast equilibrium)

Step 2: I22-(aq) + S2O82-(aq)!I2(aq) + 2SO42-(aq) (slow)

When the slow step is Step 1, the observed rate law reflects the bimolecularity ofthat step. When the second step is slow, the rate law involves a reactiveintermediate, whose concentration must re-expressed in terms of the first step. Formechanism II, the equilibrium condition for Step 1 may be used to do this . Write therate laws that are consistent with these three mechanisms, and state which of these(if any) is supported by your results.

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Questions

1. Reaction (2) is exothermic with !H = -82.9 kcal/mol. Combine this with yourmeasured activation energy to sketch an energy profile for this reaction (a hillseparating the reagent and product enthalpy levels.

2. An old rule of thumb in chemical kinetics states that the rate of reaction doublesfor every 10C rise in temperature. How well does Reaction (2) conform to thisrule. To what value of Eadoes the rule of thumb correspond?

3. The equilibrium constant for Reaction (2) is 2.6 "1043. On the basis of this,sketch time profiles of all species in the reaction on a common set of labeledaxes for the reference conditions. Label the equilibrium region. Include beneathyour sketch the relationship among the the rate of loss or gain of the four speciesin the reaction. On your sketch, show the location of your clock measurement.What were the rates for the other species at that time?

4. Identify the reaction intermediates in mechanisms I - III and attempt to drawstructures for them.

Explanatory Notes

Chemists have always been fascinated by reaction intermediates, unstable weirdmolecules that do not exist before or after, but only during a chemical reaction. Therole of a catalyst is generally thought to be the creation of a new, more reactive typeof intermediate species. For example, the Cu2+ion may complex with S2O82-,creating a new species in which the negative charge has been neutralized and theSOOS linkage is held in a rigid confirmation, making it more susceptible to attack by

I-.

The exact treatment of a mechanism actually involves the numerical integration ofcoupled rate equations, in which the rate laws refer to elementary steps. The righthanded sides of the rate laws will contain both production and loss terms, and allintermediates are included. This is generally necessary if two steps in themechanism occur with similar rates. As an example, in mechanism II, the rateexpression for I22-would be

d[I22 ]

dt = k1[I

]2 k2[I2

2 ][S2O

2

8 ]

where the first term on the right is a production term from step one with a rateconstant k1, and the second is a loss du to step 2 with the rate constant with a rateconstant k2. If, as in mechanism II, the intermediate never builds up to anappreciable concentration, its rate of change can be set equal to zero. This leads towhat is known as the Steady State Approximation, which you will learn more about ifor when you advance in your chemical adventures.

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Experiment #3: Quantitative Absorption Spectroscopy (Vernier Version)

Background

In the coming weeks, you will often be dealing with colored aqueous solutions. The

color is due to the absorption of a certain range of visible wavelengths (#) by asolution molecule or ion. The wavelengths that are not absorbed, usually thecomplimentary color to that absorbed, are what you see when you view the solutionagainst a light background. In solution, absorption lines are not seen; instead abroad, structureless band of wavelengths is absorbed by the molecule. Absorptionin the visible region of the spectrum is always due to an electronic transition (anelectron hops between orbitals). The breadth of the band aries from simultaneousexcitation of myriad vibrational degrees of freedom of the absorber and fluctuationsin solvation. The vibrational excitation occurs because of changes in the bondinggeometry that accompany the electrons promotion. Most colored solutes alsoabsorb in the ultraviolet region of the spectrum (#< 400 nm), due to excitation of

higher lying states.

Last term you performed some qualitative color/wavelength solution spectroscopyusing hand-held spectroscopes. The spectroscopes will again be used to get anoverview of the spectrum, particularly with regard to the relationship betweenabsorption and perceived color. For quantitative measurements, you will employ theVernier SpectroVis Plus. Our objectives are twofold: to relate concentrations to thethe extent of light absorption, and to relate the wavelength of maximum absorption tosome aspect of the molecular structure of the absorbing species. The relationshipbetween absorption and concentration is given by the Beer-Lambert law.

A = log10

I

I0

= `c

A!is called the absorbanceat wavelength #, I0!is the intensity of a monochromic(single #) beam of light in the absence of absorption (no absorber present), I!is theintensity with solution present, is the path length (in cm} that the light must travel

through the solution, and $!is the molar absorption coefficient. $! is different foreach molecule, reflecting the probability that an electron will hop when light of light ofthe proper #is present. The above equation predicts a direct proportionalitybetween absorbance and concentration. Once $!for a given molecule is known, the

concentration of unknown can be readily determined.

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Procedure

A. Collecting Absorption Spectra

You will measure spectra of a selection of compounds in aqueous solution:CoCl2, NaC6H2O(NO2)3(sodium picrate), Fe(SCN)Cl2, NiSO4, and the acid-baseindicator bromothymol blue. You will use the hand-held spectroscope (forqualitative observation) and the SpectroVis Plus units will be used to measureabsorption as a function of wavelength. The picrate ion will be the subject of nextweeks lab and the iron complex and nickel solutions those of future labs.

You will need to prepare your samples in the following manner:

NaC6H2O(NO2)3. Combine 2.0 mL of 1.0 "10-3M picric acid solution with 2.0mL of 1.0 M NaOH and 6.0 mL distilled water.

Fe3+and FeSCN2+. You will measure a separate spectrum of the stock

0.10 M Fe(NO3)3solution. To synthesize a sample of FeSCN2+, combine 5.0mL of 0.100 M Fe(NO3)3solution, 1.5 mL of 2.00 "10-3M KSCN solution, 2.5mL of 1.0 M HNO3, and 1.0 mL of distilled water.

Ni2+and Ni(NH3)62+. You will measure a separate spectrum of the stock0.10 M NiSO4solution. Then, slowly ad 5 M NH3to the NiSO3solution until acolor change is observed. Mix the solution thoroughly.

Bromothymol Blue. Obtain 10 mL each of 0.10 M NaOH and 0.10 M HCl.Add 10 drops of bromothymol blue to each sample. Also, add 10 drops ofbromothymol blue to 10 mL of a clear pH 7 solution. You will measure the

spectrum of all three solutions.

Obtain a SpectroVis and LabQuest Mini unit. Plug the SpectroVis unit into one ofthe data ports on the LabQuest Mini. Then, connect the LabQuest Mini to a PCusing the USB adapter provided. Open the LoggerPro software on the PC; youwill use the default file that opens up.

You will need to calibrate the SpectroVis unit before collecting any spectra. Youwill use a blank cuvette filled with water for the calibration. Place the cuvetteinto the cuvette holder of the SpectroVis unit. Go to the Experimentmenuandselect Calibrate. Follow the instructions in the dialog box, then click OK. This

completes the calibration of the SpectroVis unit.

To record a spectrum, use a pipette to add a solution to a clean cuvette andplace it in the cuvette holder. Click on the green Collectbutton. The absorptionspectrum will be displayed on the screen and will be refreshed every 2 seconds(this usually does not change the spectrums quality over time). When you aresatisfied with your spectrum, click the red Stopbutton. You should label thespectrum with the name of the solution. To do this, go to the Insert menu and

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select the Text Annotationitem. A text box with a line going to the spectrum willappear. You may adjust the position of the text box and line if you wish. Beforemoving on to the next solution, examine the data table on the lefthand side of thescreen and find the wavelength at which the absorbance reaches a maximumvalue (#max, Amax). Record this data in your lab notebook.

Follow the above procedure to record the spectra for the remaining solutions.When you click on the Collectbutton, a dialog box will open up. Make sure thatyou select Store Latest Runto keep previous spectra. Your new spectrum willbe plotted on the same screen and two new columns will appear in the data tablefor the new spectra.

When you have collected all of your spectra, you may print them. Go to the Filemenu and select Print Graph. In the dialog window, do not check the boxlabeled Print Visible Spectrum on Wavelength Graphs. Select Landscapemode so that the spectra fill most of the sheet. You should include your spectrawith your report.

B. The Beer - Lambert Law

To verify the equation in the Background, prepare samples of 0.10 M, 0.05 M,and 0.02 M CoCl2by carefully diluting small samples of stock solution withdistilled (notdeionized) water. Be sure to mix these solutions thoroughly.

You will need to configure the spectrometer so that you can create aBeer - Lamberts Law plot. First, save your absorption spectra file and open anew file. You will then need to collect a spectrum of CoCl2. Insert the mostconcentrated of your CoCl2solutions into a clean cuvette and insert it into the

cuvette holder. Click the Collectbutton. When your spectrum has stabilized,click the Stopbutton.

Now, click on the Configure Spectrometer button, which is two places to the leftof the Collectbutton. This is an Events with Entrydata collection mode. In thedrop-down list at the foot of the box, select Individual Wavelength. You shouldselect your #maxfor CoCl2from Part A (or a #close to it). Click the appropriatebox in the list or click on the appropriate position on the displayed spectrum.Then click OK. In the dialog box that opens up, select NO, in order to erase thedisplayed spectrum. The screen should now show a table with two columns anda graph of Absorbance vs. Concentration.

You can now measure the absorbance of each of your solution, beginning withthe leastconcentratedsolution. Transfer your first solution to a clean cuvette andplace it in the cuvette holder. Click the Collect button, then let the signalstabilize. Then click the Keepbutton, which appears in the Toolbar. DO NOTclick on STOP!! In the Editbox that opens, type in the molar concentration ofyour CoCl2solution, then click OK. This entry and the absorbance value at your#maxshould be displayed in the data table and the data point plotted on the

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graph. Continue with the other samples, pressing Keepfor each data point.When you have collected absorbances for each solution, click the red Stopbutton.

At this point, the graph should show the data points, but no lines. If this is not the

case, right-click on the graph and select Graph Options. Make sure that thePoint Protectorsbox is checked and the Connect Pointsbox is NOT checked.Click Done. Refer to the Instruction sheet for details.

To draw a best fit line for your plot, click the Linear Fitbutton (which is fourplaces to the left of the Collectbutton) and then click OK. The best straight linethrough the data points should now be displayed on the graph and a box shouldappear with the slope and intercept.

Calculations and Report

In addition to including your printed spectra, give a neatly arranged table that

includes Amaxand "max for each sample in Part A. For each Amax, calculate thepercentage of light with wavelength "maxthat is transmitted, the %T, and include it inyour table.

%T=10A

100%

Based on your values of Amaxand the known concentrations of each sample,calculate $!max for each to add to your table. Briefly discuss your findings as torelative absorbing power for a given concentration. An acid-base indicator itself isan acid HA, which exists as HA in acid but as A-in base. Point out which forms you

are observing in the first two bromothymol blue spectra you collected (in acid and inbase). What can you conclude about the presence of HA and A-in the pH 7 buffer?

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Questions

1. The Beer-Lambert Law results from assuming that a thin band (dx) containingindependent solute molecules (just like an ideal gas) attenuates the light by asmall amount (dI") according to

dI =Icdx

where %!is a proportionality constant, dx is a small distance of travel along thelight beam, and I!and c are intensity and concentration as defined in theBackground section. Integrate this equation to give a relationship of the sameform as the equation in the Background section. Identify $!in terms of %!(thelaw breaks down at large c or $!).

2. From your experimental "maxvalues, find the energy level spacing !E (in eV)between the highest occupied molecular orbital (HOMO) and lowest unoccupied

molecular orbital (LUMO) for each absorber. Tabulate these !E alongside the"max values in your report.

3. The picrate ion has the structure

in which the &bonds are all conjugated (alternating with single bonds).Approximating this planar ion by a 2D square box containing the &electrons (i.e.,Particle in a Box), find the box edge length (L) that reproduces the observed "maxfor the picrate ion. The box energy levels are given by

Enxny =(n2

x+ n2

y)h2

8mL2

, nx, ny = 1, 2, 3, ...,

where his Plancks constant, mis the mass of an electron, and nx, nyarequantum numbers that characterize each of the two dimensions. First, make anenergy level diagram by considering the quantum number combinations (nx, ny) =(1, 1), (1, 2), (2, 1), (2, 2), etc, being careful to show any degeneracies explicitly.Then, count the number of &electrons (twice the number of double bonds in theabove spectrum), and allow them to occupy the levels on your diagram. Make

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sure that your occupation diagram obeys the Aufbau and Pauli exclusionprinciples. The transition you observe occurs between the HOMO and LUMO.

Assign quantum numbers to the transition as (__, __)!(__, __). You can nowfind Lfrom the energy level difference calculated in Question 2.

4. The transition metal ions, Co2+

, Fe3+

, and Ni2+

absorb visible light due to anenergy splitting in their d orbitals induced by their surroundings. In whichdirection did the splitting change in Ni2+when NH3was added? (This observationwill be of interest in the discussion of transition metal chemistry and theaccompanying lab later this term.)

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Experiment #4: The Law of Mass Action: The Ionization of Picric Acid(Vernier Version)

Safety

You will be using high concentrations of HCl and NaOH, both of which are highlydestructive to eyes, skins, and clothing. Please take appropriate precautions.

Background

In the 1860s several scientists, including Guldberg and Waage in Sweden, Berthelotin France, and vant Hoff in the Netherlands, were investigating reactions which donot go to completion, leaving behind some amount of unreacted starting material.They found that, a certain ratio (K) of product to reagent concentrations was alwaysachieved, independent of the initial concentrations of reagents, provided thetemperature was held constant. For example, for an incomplete reaction

aA + bB!cC + dD (the lower case letters are stoichiometric coefficients), it wasfound that

K= [C]c[D]d

[A]a[B]b

One could even run the reaction backwards, and still achieve the same ratio. Thisbecame known as the Law of Mass Action. Its connection with thermodynamicsremained obscure until the late 1870s, when J.W. Gibbs proposed the free energychange (!G) as the criterion for the for determining the direction of a chemicalreaction. !G = 0 indicated a reaction with no preferred direction, a reaction atequilibrium, and Gibbs showed that this condition led directly to K. Khad also beenfound to a strong to be a strong function of T, as revealed mainly in vant Hoffswork. This was also shown by Gibbs to follow temperature dependence of theentropic contribution to the free energy.

In todays experiment you will explore the equilibrium properties of one of theseincomplete reactions, the ionization of picric acid (2, 4, 6-trinitrophenol,HOC6H2(NO2)3or HPic) to the hydronium ion and picric (Pic -) in aqueous solution,

(1) HPic(aq) + H2O(l)Pic-(aq) + H3O+(aq)

In the previous experiment, you have measured the absorption spectrum of thepicrate ion, which is yellow in solution; picric acid itself does not absorb in the visibleregion of the spectrum. By using the Beer - Lambert Law, and the absorptioncoefficient at a fixed visible wavelength, #= 440 nm from that experiment, you canmeasure the equilibrium concentration of Pic-. Knowing the initial concentration ofundissociated acid [HPic]0and the stoichiometry of Reaction (1) allows you to find[H3O+] and [HPic] at equilibrium as well. You could then evaluate the mass action

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constant, usually called the equilibrium constant Kfrom the mass action(equilibrium) condition

K= [H3O

+][Pic]

[HPic]

In general, this ratio of concentrations is called the reaction quotient Q. Toestablish this ratio as a law for this reaction, you can vary the [H3O+] in the solutionby adding large excess of HCl, which can be assumed to be 100% ionized in waterto give H3O+and Cl-. Then, [H3O+] is determined entirely by the added HCl, and[Pic-] and [HPic] must readjust themselves to restore equilibrium if Kis truly aconstant. (You can qualitatively predict what will happen to the concentrations byinvoking LeChteliers Principle, see Chapter 14 in your textbook.) The temperaturedependence of Kcan also be assessed by performing the absorption measurementsat different temperatures.

Procedure

You will be supplied with stock solutions of 1.0 "10-3M picric acid, 1.0 M NaOH, and2.0 M HCl. To insure accurate volume readings, you should use a buret to delvereach solution. You should prepare each solution in a small beaker; each with a totalvolume of 20 mL. You will need to determine the volume of picric acid stock solutionneeded to make each solution 2.0 "10-4M in picric acid (i.e., [HPic]0= 2.0 "10-4M);use the dilution formula (M1V1= M2V2). Add this predetermined volume of picricacid, any other stock solution (HCl or NaOH) and dilute to 20 mL with distilled water.The concentrations of the stock solutions have been carefully chosen to allow allrequired solutions to be made up without exceeding the pre-selected total volume.

In the real world, you would be responsible for this condition as well.

Dont worry about slight deviations in your actual delivered volumes from yourcalculated volumes, simply record the actual volumes used, and dilute with water to20 mL. You can later compute the actual concentrations using the dilution formula.Swirl each solution well to insure complete mixing and attainment of equilibrium.Extra solution may be required if you elect to do temperature dependentmeasurements. Record the laboratory temperature to the nearest 0.1 C using youralcohol thermometer.

A. Calibration

Obtain a SpectroVis and LabQuest Mini unit. Plug the SpectroVis unit into one ofthe data ports on the LabQuest Mini. Then, connect the LabQuest Mini to a PCusing the USB adapter provided. Open the LoggerPro software on the PC; youwill use the default file that opens up.

You will need to calibrate the SpectroVis unit before collecting any spectra. Youwill use a blank cuvette filled with distilled water for the calibration. Place thecuvette into the cuvette holder of the SpectroVis unit. Go to the Experiment

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menuand select Calibrate. Follow the instructions in the dialog box, then clickOK. This completes the calibration of the SpectroVis unit.

Since each SpectroVis Plus unit is slightly different, you should check yourabsorption coefficient measurement from last weeks work. Prepare a solution

with 2.0 "10-4

M HPic and 0.20 M NaOH, diluting your total solution volume to20 mL. Record the actual volume of stock solution used next to your plannedvolumes. The excess NaOH insures that the (Bronstead - Lowry) acid basereaction

(2) HPic(aq)+ OH-(aq)!Pic-(aq)+ H2O(l)

will be driven to completion, making [Pic-] = [HPic]0. To record a spectrum, use apipette to add a solution to a clean cuvette and place it in the cuvette holder.Click on the green Collectbutton. . When you are satisfied with your spectrum,click the red Stopbutton. You should label the spectrum in a way thatdistinguishes it from the others you will collect. To do this, go to the Insert menuand select the Text Annotationitem. A text box with a line going to the spectrumwill appear. You may adjust the position of the text box and line if you wish.

Examine the data table on the lefthand side of the screen and find the maximumabsorbance. Record this value along with the wavelength and a calculatedpercent transmission in your lab notebook. Compare this with your results fromlast week. Save this solution.

B. Establishing the Mass Action Law

Make up five (or more) solutions, each containing 2.0 "10-4M HPic and various

concentrations of HCl ranging from 0.20 M to 1.0 M; dilute each solution to20 mL. Again, record the actual volume of each solution next to your plannedvolumes. Observe each solution by eye against a white background and rate thecolor intensity of the solutions.

Record a spectrum for each solution, making sure to annotate eachappropriately. For each spectrum, find the maximum absorbance at and recordit, along with a calculated percent transmission in your lab notebook. How do theabsorbances compare with your visual observations? After this series ofmeasurements, remeasure the maximum absorbance for the calibration solution.

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C. Further Exploration

After completing the above measurements and calculating at least twoexperimental K values as described below (to make sure that your experimentshave worked well), you may explore further. Here are three suggestions; feel

free to explore your own ideas, checking them with your lab instructor first.

1. Does your Kvalue determined in Part B depend on temperature? Choose onsolution from Part B and use it to determine the temperature of theabsorbance. Use a dry dropper pipet to add this solution to a clean cuvette.You will need to adjust the volume of solution in the cuvette so that the bulb ofyour thermometer can be inserted in it without causing spillage. Then,immerse the cuvette in an ice water bath, monitoring its temperature T (C).Observe the color relative to the parent solution at ambient T. Record aspectrum, then quicklyremeasure T (C). Take the measurementtemperature to be an average of the temperature readings before and afteryou recorded the spectrum. Repeat with water baths at differenttemperatures, both hot and cold, to give 3 - 4 non - room ambient spectra andabsorbances. Report your results using a table to display Kvs. T (C) and

T (K). Also calculate ln K, and1

T(K) , and use Excel to prepare a plot of ln K

vs.1

T(K) . Insert a best fit line through you data. According to the vant Hoff

equation

d lnK

d(1/T)=

H

R

the slope of your best fit line should equal

H

R, where

!H is the heat of Reaction (2). Report a value for !H . Note that here youcan obtain a !Hwithoutdoing calorimetry.

2. According to LeChateliers Principle, adding H3O+to Reaction (1) atequilibrium will cause it to shift to the left, meaning less ionization and lesscolor. Explore this by putting a small portion of picric acid stock solution in atest tube, noting its color, and adding concentrated HCl (12 M) dropwise whileobserving. What can you conclude about the absorption spectrum ofundissociated picric acid?

3. All of the above measurements employed the same initial concentration[HPic]0. How do you predict things would change if [HPic]0were different?Devise some experiments to test your predictions, and carry them out, reportnumerical measurements and comparisons.

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Report

In addition to reporting ambient temperature and all absorbance and percenttransmission values, use your absorbance at #= 440 nm from Part A to obtain the

product ` =A

[Pic], and use it to compute and tabulate [Pic-] for each solution in

Part B. Because HCl is completely ionized and [H3O+] >> [HPic] in all cases, youmay set [H3O+] = [HCl] in the K expression for Reaction (1). Deduce what to use for[HPic]; note that [HPic] '[HPic]0, since some has ionized. For each of thesesolutions, calculate K, tabulate, and calculate an average and error. Assess howwell the Law of Mass Action holds

Questions

1. Use the K obtained from Part B to calculate the standard free energy change!G at laboratory temperature (See Chapter 17, Equation 17.15). Byconsidering the extreme possible values of K, or by using

pebbles (differentials), give an error estimate for !G. What does your valueimply about the spontaneous direction of Reaction (1).

2. Discuss your results on the concentration dependance of the position ofequilibrium in terms of LeChteliers Principle (See Chapter 14).

Explanatory Notes

The ionization of picric acid was chosen not only because it is a well balancedreaction, but it is an example of acid ionization, a topic that will soon occupy us inboth lecture and lab. Note that in reactions of type AB + C + ..., if the initial [A] is

known, only one concentration needs to be measured at equilibrium to completelydetermine the composition of the system. This is worked out in general byintroducing an extent of reaction variable, an analysis you will be learning to use incoming labs and in lecture as well.

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Experiment #5: pH Measurements I:Acid-Base Equilibria and Titrations

(Vernier Version)

Background

Aqueous solutions of compounds we call acids and bases, and their combinations,produce a wide range of [H3O+], typically between 1 M and 10-14M. To expressthese concentrations conveniently as nearly-whole numbers, the Danish chemistSrensen introduced the pH scalein 1909. pHstands for the negative of the powerof the Hydronium ion (H3O+) concentration and is defined as

pH= log[H3O+]

Later this was generalized to include other solute concentrations, as well asequilibrium constants. Thus, we write,

pOH= log10[OH]

pCu = log10[Cu2+]

pKa= log10Ka

for example. The acidic and basic ends of the pH scale are delineated by the pH ofpure water (pH = 7.00, [H3O+] = 1 "10-7M ). Acids have pH < 7 ([H3O+] > 1 "10-7M)and bases have pH > 7 ([H3O+] < 1 "10-7M). In aqueous solution, [H3O+] and [OH-]are related by the water autoionization equilibrium, leading to

pH+pOH = pKW = 14.00

Measurement of such small concentrations was very tedious before the invention ofthe glass electrodein the 1930s. This small, portable electrochemical device madethe measurement of pH rapid and accurate over the entire [H3O+] range.Technological improvements since then in both the electrode and the measuringelectronics allow the regular use of the device, now known simply as a pHmeter, inintroductory laboratory programs, research, and testing laboratories. Today you willbe using the Vernier pH sensor along with the LoggerPro software, which works verymuch like a traditional pH meter. The pH sensor, like the glass electrode isexpensive, delicate, and easily damaged by mishandling. Pleas follow the

instruction sheet on page xx when using it. You are urged to read Chapter 19 inyour textbook for a description of how a pH meter works.

Today, you will be using the pH sensor as a tool to probe the equilibrium behavior ofa variety of acids, bases, and salts, and to determine the changes in pH that occur inthe course of acid - base reactions by means of titrations. The pH varies smoothlyduring a titration, forming what is known as a pH titration curve, a plot of pH vs.moles or volume of titrant added. In next weeks experiment, you will test the

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effectiveness of a buffer in resisting changes in pH and measure the solubility of aslightly soluble metal hydroxide. Further background reading may be found inChapter 16 of your text.

Procedure

A. Set up and Calibration

Obtain a pH sensor and LabQuest Mini unit. Plug the pH sensor into one of thedata ports on the LabQuest Mini. Then, connect the LabQuest Mini to a PCusing the USB adapter provided. Open the LoggerPro software on the PC; youwill use the default file that opens up.

You should carry out an initial two point calibration, first using a pH 7.0 buffer(yellow solution) and secondly a pH 4.0 buffer (pink solution). Instructions for thecalibration will be provided by your lab instructor. If during the experiment yoususpect that the sensor is not working properly, you can repeat this calibration

procedure.

For Part B, you will read the pH from the display in the bottom lefthand corner ofyour screen. For Part C, you will need to open the Acid-Base Titration fileprovided by the Vernier Company. In this file, the data collection mode, Eventswith Entry has already been fully set up. There will be a table with two columns,labeled Volume (mL) and pH on the left and a graph on the right with similarlylabeled axes. The pH will still be displayed in the lower lefthand corner.

B. pH of acid, base, and salt solutions

You will be supplied with the following stock solutions (all 0.100 M): HCl,CH3COOH (acetic acid), NH4Cl, NaCl, CH3COONa (sodium acetate), NHCO3,Na2CO3, NH3, NaOH. Group these into acidic, neutral, and basic sets. First,measure the pH of the acidic and neutral solutions, and tap and deionized wateras well. Using the HCl and CH3COOH stock solutions, prepare solutions dilutedto 0.01 M and 0.001 M, and measure the pH of each. Put 1 - 2 drops of universalindicator into each of your beakers after the measurement and record theobserved color. Next, measure the pH of the basic solutions, again addingindicator after the measurement. Study the effect of dilution for NH3and NaOH isoptional.

C. pH Titration Curves

Your instructor will assign you HCl and acetic acid solutions of unknownconcentration (concentrations may range from 0.05 to 0.2M). Make sure yourecord the unknown letters in your lab notebook. Recalibrate the pH sensor ifnecessary. Titrate the HCl first. Use a 10 mL volumetric pipet to deliver20.00 mL of your unknown solution into a 150 mL beaker. Place the beaker on a

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magnetic stirrer unit and insert a magnetic stir bar, adjusting so that you arestirring at a moderate rate.

Insert the pH sensor, then click on the green Collect button. When the pHreading on the bottom lefthand corner of the screen is steady, click on the Keep

button. In the edit box that opens, enter 0 mL, then click OK. The data pointwill automatically be entered in your data table.

Next, rinse and fill a buret with the standardized 0.10 M NaOH solution to the0.00 mL mark. Begin to add base, stopping to record the volume (to 0.01 mL)and pH after every mL or when the pH changes by 0.5 units or more. For eachmeasurement, click on the Keep button and record the cumulative volumeadded. At the first sign of rapid pH change, begin to add base drop-wise,recording the volume and pH after each drop. After passing through the rapidly-changing region, where the endpoint occurs, you may resume adding largeramounts of base until you are about 5 mL beyond it. When you are finished withyour titration, click the Stop button. Use your endpoint volume to calculate theconcentration of your HCl. To print your titration curve, go to the Filemenu andselect Print Graph. Select Landscapemode so that the plot fills most of thesheet. You should also copy your data table into your lab notebook.

Repeat the above procedure for your acetic acid unknown. You should save yourHCl data file and open a fresh copy of the Acid-Base Titration file beforebeginning.

D. Reactive preparation of a precipitate

You may have seen in class (and will now see again for yourself) that when

calcium metal reacts with water, displacing hydrogen, a milky-white Ca(OH)2precipitate is formed:

(1) Ca(s)+ 2H2O(l)!Ca(OH)2(s)+ H2(g)

In preparation for next weeks experiment, you are asked to carry out thisreaction and store the resulting solutions and precipitates. Add approximately200 mL distilled water to a clean 600 mL beaker. Slowly add Ca(s) to the waterand allow the reaction to proceed until you notice cloudiness, indicating theformation of Ca(OH)2; then add a little more Ca. Stir the solution with a cleanglass stir rod to insure complete reaction. Record your observations on the

reaction and the appearance of the solution. The total mass of Ca should notexceed 0.5 g or so. Monitor the temperature of the solution, and allow it to coolsomewhat with occasional stirring. Now stir the cooled solution to suspend theprecipitate and immediately pour into a 500 mL plastic bottle for storage. Labelthe bottle 0.5 g Ca. Clean and dry the beaker, and carry out the same reaction,but use roughly 1.0 g Ca. This should result in more precipitate being formed.Note any visual differences. After cooling, stir and store the solution as before,labeling the second bottle 1.0 g Ca.

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E. Further exploration

After completing the above experiments, and show your instructor your datatables and titration curves, you are free to explore further. Here are a fewsuggestions; please check with your instructor before pursuing your own ideas.

1. Bring a sample of one or more liquids, or solids that can be dissolved, fromdorm or home to test pH. Report measured pHs along with your estimate ofthe source compound that yields acidity or basicity.

2. Among the exploratory experiments suggested in Experiment #4 was thetemperature dependence of the picric acid equilibrium. Attempt a similar setof measurements on water ionization or any of the weak acid - baseionizations. Report a vant Hoff plot and estimated heat of reaction andcompare with literature values. (You have probably measured !H for areaction that is the reverse of water ionization!)

3. Titration curves are supposed to turn upside down when a weak base istitrated by a strong acid. Check this out by using your (now known) HClsolution to titrate a 0.1 M NH3sample, generate a titration curve, andcompare it to the week acid case, and to equilibrium-based predictions.

Report

For background acid-base equilibrium theory and relationship needed to analyzeyour results, see the text (Chapters 16 and 17).

Present your results of Part B in a table showing identity of solute, concentration,

and pH reading. In a separate column, predict the pH of HCl and NaOH solutionsbased on known concentrations and assumed 100% ionization. Use the pH valueobtained for CH3COOH and NH3to calculate their ionization constants, Kaand Kbrespectively and pKaand pKb, comparing them to their literature values. Then useyour experimental K values to predict the pH of each solution and their conjugatesalts CH3COONa and NH4Cl. Enter these predictions next to your experimentalvalues in the table. From formulas given in the text (Chapter 16), use the pHmeasured for NaHCO3and Na2CO3to calculate Ka1and Ka2for carbonic acid H2CO3and compare them to literature values.

For part C, preset the pH titration curve data in the form of a tabulation (which can

be reproduced from your LoggerPro table) and copies of your printed titrationcurves. Determine the endpoint for each as the point of steepest slope, interpolatingyour graph if necessary, and use the volume of base added to calculate theconcentrations of the solutions. Also, present a short table containing the pH at thestart of each titration, at the exact midpoint, and at the endpoint. Using the HClmolarity you have just determined, predict the pH at the start and midpoint for theHCl titration and enter these values in a separate column next to the measuredvalues. Use the measured pH at the endpoint to estimate the water autoionization

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constant KW. Use the measured pH values for acetic acid and its computed molarityto determine separate estimates of the equilibrium constant Kaat each of the threepoints. Compare these values to each other and to that obtained in Part B. Useyour determinations to calculate the pH at two other points along the acetic acidtitration curve on either side of the midpoint and compare with your measured pH

values.

Questions

1. Use your average value of Kafor acetic acid to compute a value for the standardfree energy change !G for the ionization of acetic acid, with error. Is thereaction spontaneous under standard conditions? Why does it happen underyour conditions?

2. Hesss Law applied to free energies can allow you to predict the equilibriumconstants K and standard free energy change !G for a number of acid-basereactions you studied in the heat of reaction lab from last term. Make such aprediction for the three acid-base reactions you studied in the heat of reactionlab,

HCl(aq) + NaOH(aq)!NaCl(aq) + H2O(l)

HCl(aq) + NH3(aq)!NH4Cl(aq)

NH4Cl(aq) + NaOH(aq)!NaCl(aq) + NH3(aq) + H2O(l)

as well as the reaction

CH3COOH(aq) + NH3(aq)!CH3COO-(aq) + NH4+(aq)

Compare with literature values. What can you say about the spontaneity of thesereactions under standard conditions? What was assumed about the spontaneityand extent of reaction in the !H lab? Are these assumptions valid in light of yourpresent results?

3. A pH meeter functions by measuring the free energy difference between a 1.00MHCl solution behind the glass membrane and the sample solution inside (i.e., itmeasures the free energy change of dilution). Compute this !G for samplesolutions of pH 1.00, 7.00, and 13.00. What is the proportionality factor relating

pH to !G?

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Explanatory Notes

Membranes like the glass membranes that serve as sensors for a variety of specificions are now commercially available. These ion selective electrodesor pIonelectrodesoperate on the same principle as the pH electrode, through the exclusive

transmission of ions of a certain type across a membrane. Gas permeable andbiocatalytic membranes also allow electrochemical detection of neutral moleculesranging from gases such as O2or CO2to small proteins (chains of amino acids).

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Experiment #6: pH Measurements II:Buffers, Solubility, and the Common Ion Effect

(Vernier Version)

Background

The efforts of Bertholett in ca. 1799 - 1803 (not Berthelot, who came much later)toward the solubility of salts, using both heat and compounds with elements incommon, gave us our first inkling of the possibility of equilibrium and reversiblereactions.. Today we understand the effect of adding an ion in common with areaction product on both buffer solutions and precipitates in terms of its influence onthe equilibrium reaction quotient. In todays world, buffers, solubility, andprecipitation play a role in a wide range of phenomena from water hardness andpollution, to antacid and other gastrointestinal medication, to clean-up of nuclearwaste.

Last week you analyzed an acetic acid CH3COOH unknown using pHmeasurements and a standardized NaOH solution. Now you can use your results toprepare and test a buffer solution. You also prepared two-phase mixtures containinga solution and a precipitate by reacting calcium metal with water. This method wasselected (as opposed to mixing a salt solution with a soluble hydroxide) to produce asystem where the only ions present are Ca2+and OH-. As you will see, the pH of thesupernatant liquid of your equilibrated mixtures in considerably greater than 7,indicated the presence of OH-(aq), and therefore incomplete precipitation. The twophase equilibrium involved is

(1) Ca(OH)2(s)Ca2+(aq) + 2OH-(aq)

If the system is at equilibrium, the solution is saturatedwith Ca(OH)2, i.e., the freeenergy change of reaction (1) under the prevailing system conditions is zero, or thefree energy of the precipitate is balanced by that of the solvated ions. The amount(molarity) of the product that remains dissolved is called the solubility of the salt,and may be determined from the solubility product

Ksp= [Ca2+][OH]2

when !GT = (RTlnKspfor reaction (1). Note that the precipitate itself does notappear in the Kspexpression since no work is required to take it from in standardstate to the actual condition in the system. This in turn means that the equilibriumposition established in reaction (1) cannot be influenced by the amount of precipitatepresent as long as there is some. However, if there is a common ion (Ca2+or OH-)already present in appreciable concentration or if such an ion is added to theequilibrium system, the equilibrium position will change in accord with Le ChateliersPrinciple, as quantitatively represented by the Kspexpression, although Kspremainsconstant. If the temperature is changed, however, the value of Kspitself will alter,

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depending on the sign of !H for reaction (1), and shift the equilibrium. In todaysexperiment, you will be able to explore these effects, and learn how to control one ofthe more important and practical types of chemical reaction.

Procedure

A. Set up and Calibration

Obtain a pH sensor and LabQuest Mini unit. Plug the pH sensor into one of thedata ports on the LabQuest Mini. Then, connect the LabQuest Mini to a PCusing the USB adapter provided. Open the LoggerPro software on the PC; youwill use the default file that opens up. You will read the pH from the display in thebottom lefthand corner of your screen.

You should carry out an initial two point calibration, first using a pH 7.0 buffer(yellow solution) and secondly a pH 4.0 buffer (pink solution). Instructions for thecalibration will be provided by your lab instructor. For Parts C-F, you should

repeat the calibration using the pH 7.0 buffer and a pH 10.0 buffer. If during theexperiment you suspect that the sensor is not working properly, you can repeatthis calibration procedure.

B. Making and testing a buffer solution

Using the same acetic acid unknown you analyzed last week, combineappropriate volumes of you unknown and the standardized NaOH solution tomake about 40 - 50 mL of 1:1 acetic acid/sodium acetate buffer in a 150 mLbeaker. Making up the solution involves reacting exactly half the acetic acid inyour sample with NaOH, thereby providing a concentration of acetate ion (the

common ion) equal to that of the unreacted acetic acid. Record the precisevolumes and measure the pH of the buffer using your pH sensor. Add 2-3 dropsof universal indicator, noting the color. Test the resistance of the buffer to strongacid or base by first adding 1.0 mL of 0.1 M NaOH, then adding 2.0 mL of0.1 M HCl and remeasuring the pH. Save your buffer solution if you wish toexplore it further a la suggestion 3 below.

Now take a volume of distilled water the same as that of the buffer in a secondbeaker, add universal indicator, and add 0.001 M HCl dropwise to give a pHabout the same as that of the original buffer. Then repeat the addition of baseand acid as before, noting the pH and color changes.

C. Determination of [OH-] in the supernatant liquid of a precipitation

Take care not to shake or otherwise disturb your equilibrated precipitationmixtures. Using the 0.5 g Ca mixture, pour off a small portion of thesupernatant solution and measure its pH; then add universal indicator and noteits color. The pH measurement allows you to determine [OH-] and and Ksp.

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D. The common ion effect

To 50.0 mL of the 0.5 g Ca supernatant solution, add about0.02 mol CaCl2!2H2O(s), pre-weighed; observe. Note that based on the solubility

rules, CaCl2is highly soluble. Stir the solution thoroughly with a stirring rod, and

allow it to return to ambient temperature. Repeat the pH and indicatorobservations of Part C for this solution.

E. Varying the amount of precipitate

Repeat Part C with the 1.0 g Ca mixture.

F. The effect of temperature

Using supernatant solution from either mixture, transfer 50.0 mL to a clean150 mL Erlenmeyer flask. Clamp the neck of the flask and immerse it in a600 mL beaker of water with a magnetic stir bar. Place the flask and beaker on a

magnetic stirrer and set the stirrer for a moderate speed, Set the heater tomoderate heat (about 3 or so) and heat the assembly, monitoring its temperature.Every 5C, shut off the stirrer (but not the heat), and measure and record the pHof the contents of the flask, up to 50C.

G. Further exploration

Upon completing your experiments and calculating at least one Kspvalue, youare free to explore further. As usual, you may pursue your own ideas or follow asuggestion; check with your instructor before proceeding.

1. Explore the distinction between equilibrium and non-equilibrium conditions.As you observed last week, the Ca(s) + H2O(l) reaction is exothermic, andwere equilibrium notions to apply to the reaction in progress, heating thewater would slow the net reaction down by encouraging the reverse reaction.Design a small scale experiment to test this idea in the reaction of Ca withcold and hot water; report your data and conclusions concerning anydistinction.

2. What should happen if the saturated Ca(OH)2system is cooled? Try takingthe contents of the reaction beaker of Part D and cooling it while stirring in anice bath before letting it settle. Measure the temperature and the pH and

report your conclusions. Are these results consistent with your othermeasurements.

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3. Establish that the solubility of Ca(OH)2is pH dependent (see Question 2below). After you are assured that all experiments above are completed toyour satisfaction, thoroughly shake the 1.0 g Ca bottle, pour out about10 mL into a small beaker, and slowly add the buffer solution from Part A.Record your observations, and perform any relevant calculations.

Report

For Part B, use the measured pH of your buffer to determine pKaand Kafor aceticacid from the Henderson-Hasselbalch equation (See Chapter 17), and compare it tothose obtained from last weeks measurements, Parts 5A, 5B, and 5C. Averageyour various estimates to give a final experimental value of Kawith error, andcompare it to the literature. Using the H-H equation with your own buffer value ofpKato predict the new pH as a result of adding the strong acid and base andcompare it with your measurements. Also predict the outcome of the unbufferedaddition of strong acid and base, and compare it with your experiment.

For Parts C, D, and E, group data and calculations from the separate runs togetherbut compare runs in a table at the end. Compare Kspfrom the pH measurements forall runs. Is Kspfrom Part D consistent with Parts C and E? In analyzing your results

from Part D, dont forget to account for the Ca2+ addedand that present due todissolved Ca(OH)2; otherwise you will miss the point of the exercise! Does theamount of precipitate affect the position of equilibrium? Find an average Kspwitherror at ambient temperature. Also, give experimental values for the solubility ofCa(OH)2, i.e., its molarity in a saturated solution at ambient temperature, and themass (in g) of Ca(OH)2that will dissolve in 100 mL of solution. Compare Kspas afunction of T from Part F, and prepare a vant Hoff plot. Derive values of !H from

the slope, and !S from !G and !H, and discuss these quantities in terms of theinteractions and arrangements of the reagent and products of reaction (1). Discussthe consistency of your results of Parts D, E, and F with Le Chateliers Principle.

Questions

1. What is the reaction quotient for the Ca(s) + 2H2O(l)!Ca(OH)2(s) + H2(g)reaction? Find !G and K for this reactions from tables, and predict what youcan about the the state of the system at equilibrium. Use this information alongwith your experimental data to predict !G and K for the reaction

Ca(s) + 2H2O(l)!Ca2+(aq) + 2OH-(aq) + H2(g).

Compare the degree of spontaneity under standard conditions of the reactionwith that given it above. Give a free energy diagram that illustrates thecomparison.

2. The solubility of hydroxides is strongly pH dependent. At what buffered pH wouldthe solubility of Ca(OH)2, expressed as [Ca2+]eqbe 1.0 "10-4M?

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Explanatory notes

Although Ca(OH)2precipitates readily, as you have seen, it ranks as one of the moresoluble of the metal hydroxides from Group IIA rightward. Some of the transitionmetal hydroxides, such as Fe(OH)3, are so insoluble that the measurement of

OH-

(aq) is hopeless, and electrochemical methods must be used. For a fixed ioniccharge (say 2+), the metal ions tend to become better Lewis acids as one movesfrom left to right on the Periodic Table, strengthening and lending more covalentcharacter to the bonding in the salt. Hydroxide precipitates of the rightmosttransition metal ions, such as Ni2+or Cu2+, can often be taken back from their waterygraves by strong ligands such as ammonia or cyanide to make soluble transitionmetal complexes.

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Experiment #7: Redox Activity Series

Safety

Aside from the use of sulfuric acid as a reagent, the hazards in this experiment are

minimal.

Background

The idea that chemicals have an affinityfor each other is a very old alchemicalconcept, but it was not until the seventeenth century that Gauber, Boyle, and Newton(our physicist turned chemist) began to recognize relative affinitiesin singledisplacement reactions, where an element would take the place of another in itscompound. For example, copper metal would replace silver in nitre of silver (silvernitrate) to make nitre of copper and silver metal. Written in our modern way, thiswould be

Cu(s) + 2AgNO3(aq)!Cu(NO3)2(aq) + 2Ag(s).

This is a one way street; silver metal will not replace copper in its nitrate. The earlychemists believed this to be due to the greater affinity of Cu for nitrate.

The French apothecary E.F. Geoffroy is credited with assembling the first table ofrelative affinities or activitiesin 1718; he called it a Table des rapports. A halfcentury later, when the distinction between elements and compounds had becomeclearer, the Swedish chemist Torbern Bergman, a contemporary of Cavendish,devoted his entire career to an obsession with ranking elements and compounds in

order of their activity toward a given test substance. Metal-acid reactions played amajor role in his work and he recognized for the importance of water -- there wereseparate wet and dry activity rankings. Electrochemical findings by Davy andFaraday in the early nineteenth century cast a new light on the problem, suggestingthat the origin of affinity is electrical in nature. The activity series was finally put on aquantitative basis decades later by Walther Nernst, who measured the electricalpotentials generated between dissimilar metals and their soluble compounds involtaic cells, using Arrheniuss ion theory to interpret his results. A completedescription still had to await the discovery of the electron by Thompson, makingpossible our present-day picture of a single displacement reaction as an electrontransfer or redox reaction that only occurs in a certain direction. We now recognize

Bergmans distinction between wet and dry activities is due to solvation energyliberated when an ion dissolves in water.

Our twentieth century understanding of the structure of the atom and the PeriodicTable has finally allowed us to appreciate the origins of the activity series in terms ofatomic properties (principally ionization energy IE), with exceptions to predictedperiodic trends attributed to differences in solvation energies of the ions formed.

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Electrons always want to roll downhill, and chemists, and now you young uns, candiscover which way is down through clever experimentation.

Procedure

A. Chemical investigation of activity

Your goal will to be to devise a series of reactions with the reagents below thatutilize the minimumnumber of reactions necessary for the activity ranking of thesix elements: copper, hydrogen, iron, magnesium, sodium, hydrogen, and zinc.For reasons having to do with surviving in the lab, elemental sodium andhydrogen will not be available, so your reaction scheme will have to do withoutthem. Some of the reactions may require a few minutes to show appreciableproduct, so you must be patient.

Available Reagents

Elements Compounds(aqueous solutions)

copper copper (II) sulfate

iron iron (II) sulfate

magnesium magnesium sulfate

*** sodium sulfate

*** sulfuric acid

zinc zinc sulfate

For each reaction that is observed to occur, write a balanced net ionic equation.Then compose an activity series, most active element first, based on yourresults. From the activity series, predict the results of a few reactions you did nothave to try, and experimentally test your predictions.

B. Mystery Element

After completing Part A, obtain a sample of an unknown element (call it X) and asolution of its compound from your instructor. By means of chemical tests,decide where it fits in your activity series. Speculate on its identity.

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C. Further exploration

Take a suggestion or strike out on your own

1. In some early studies, the rateat which a metal would evolve hydrogen gaswhen reacted with acid was used as a gauge of its activity. In some cases,this was found to be unreliable. Does it work for our selection of elements?Set up a gas collection apparatus, and devise experiments to test this idea.Be quantitative, and report any conclusions.

2. The more active metals will react directly with water (no acid needed), andwater near the boiling point is much more reactive than at room temperature.Test those elements that are most likely to react with hot water, based on youractivity series. Report results, chemical equations where appropriate, andconclusions.

3. Set up a voltaic pile by sandwiching a paper towel soaked in sodium sulfate

between two pieces of metal. Use a multimeter to measure any voltagegenerated by touching the leads to the metal pieces. Do the pieces have tobe dissimilar metals to get a voltage? How does the sign of the voltagedepend on the way you measure it? When the voltage is measured, a smallvoltage passes through the pile; what is carrying the current? Try shorting thetwo pieces to each other with another piece of metal; observe. Reportnumerical results, and relate the measurements to your activity series.

Report

Clearly indicate your reaction scheme, and the logic behind it, along with a log of

attempted reactions, observations, balanced equations, and conclusion. State youractivity series clearly, along with the location of X in it.

Questions

1. Next to each element as arranged in the activity series you found, tabulate theionization energy IE, electron affinity EA, and electronegativity ). Which of thequantities is arranged most nearly in a monatomic sequence that matches yourseries? Are there any exceptions? What other energetic feature(s) of thedisplacement reactions might be important factor(s)? Speculate on the atomicproperties of X.

2. Choose one of the metal-acid reactions you run and illustrate its mechanismusing Lewis structures for reactants and products, using loops and curlyarrows.

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3. In last terms experiments, you observed a number of single displacement redoxreactions; list as many as you can. Identify the more active element in eachcase, and classify each as wet or dry. Try to incorporate or corroborate theseobservations with your present results.

Explanatory Notes

A precise knowledge of the activity series is important in metallurgy and metalrecovery from native materials and ores. The alkalis and al

UHC Gen Chem 2 Manual

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