chemistry 223 w13 - pcc manual project/223... · web viewci1 kinetic factors-i worksheet ci2...

Table of ContentsKinetic Factors - I................................................................................................................1Shifting Reactions..............................................................................................................15Determination of the Equilibrium Constant of Phenolphthalein Dissociation..................33What Factors Affect the Solubility of Ions?......................................................................53What Factors Affect the Reactivity of Acids?...................................................................77Acid/Base Interactions.......................................................................................................97How Can Chemical Reactions Be Made to Produce Electricity?....................................115How Can Electricity Drive a Chemical Reaction?..........................................................137What is a Polymer?..........................................................................................................165

CI1 Kinetic Factors-I Worksheet

ChemInquiry1 Name__________________________

Kinetic Factors - I

What Factors Affect the Rate of a Reaction?

Background & ProcedureQUESTIONS OF THE DAY

What factors can be used to change the rate of a reaction? Are these factors predictable from the Model (Collision Theory)? Are these factors predictable from the balanced equation?

IntroductionModel: The Collision Theory of Chemical Reactions

Chemists have come up with a model, based on logic and mountains of experimental evidence, that attempts to explain a large number of observations about factors that affect the rate at which chemical reactions occur. This model is called the “Collision Theory of Chemical Reactions.” In the Collision Theory, the following assumptions are made:

Particles are in constant, random motion. The average kinetic energy (speed) of a particle is dependent only on the temperature. Particles must collide in order to react. The probability of a collision is affected by the concentration of the reactants. Greater

concentrations give greater collision probability. More molecules collide than actually react. In other words, only a small fraction of collisions are

successful in producing a reaction.o Steric Factor: Particles must often be oriented correctly when they collide in order to

actually react.o There is a minimum energy requirement for a reaction to occur. This energy is called the

activation energy (Ea).

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Figure 1 The effect of molecular orientation (the steric factor) on the reaction of a particular system, that of NO and O3.

Figure 2 The diagram shows how the energy of this system (NO + O3 NO2 + O2) varies as the reaction proceeds from reactants to products. Note the initial increase in energy required to form the activated complex.

Reacting molecules must have enough energy to overcome electrostatic repulsion, and a minimum amount of energy is required to break chemical bonds so that new ones may be formed. Molecules that collide with less than the threshold energy bounce off one another chemically unchanged, with only their direction of travel and their speed altered by the collision. Molecules that are able to overcome the energy barrier are able to react and form an arrangement of atoms called the activated complex or the transition state of the reaction. The activated complex is not a reaction intermediate; it does not last long enough to be detected readily.

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Figure 3 The potential energy diagrams for a reaction with (a) ΔE < 0 (exothermic) and (b) ΔE > 0 (endothermic) illustrate the change in the potential energy of the system as reactants are converted to products. Ea is always positive. For a reaction such as the one shown in (b), Ea must be greater than ΔE.

For an A + B elementary reaction, all the factors that affect the reaction rate can be summarized in a single series of relationships:

rate = (collision frequency)(steric factor)(fraction of collisions with E > Ea)

You may recall from the kinetic molecular theory (KMT) in CH 222 that for any substance, there is a distribution of kinetic energies of the molecules of a substance. As shown in the diagram below of the Maxwell-Boltzmann distribution, the fraction of molecules with a collision kinetic energy E > E a can be represented by the shaded area underneath the distribution curve.

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Figure 4 Having a KE greater than the Ea for a reaction allows for the possibility of, but does not

guarantee, a collision resulting in a chemical reaction.

The Reaction: Permanganate Ion Reduced by Oxalic AcidIn this experiment you will be studying the reaction between permanganate ion, MnO4

1 and oxalic acid, C2O4H2, in acid solution. The overall balanced reaction equation is:

2 MnO41(aq) + 5 C2O4H2(aq) + 6 H3O+(aq) 2 Mn2+ (aq) + 10 CO2(aq) + 14 H2O(l)

NOTE: H3O+(aq) represents the “acid” in this reaction. In water, H3O+ is the reactive chemical species formed by any substance called an “acid”. In this reaction, the H3O+ reactive species is formed from sulfuric acid, H2SO4(aq) when it reacts with water as shown:

H2SO4 (aq) + 2 H2O (l) 6 H3O+ (aq) + SO42- (aq)

Thus, you may consider H3O+ and H2SO4 (sulfuric acid) as synonymous for this reaction.

Experiment 1:Watching the Reaction1. Using clean, dry, labeled 100-mL beakers, obtain ~15 mL each of the stock solutions

0.0014 M KMnO4

0.20 M C2O4H2

2.0 M H2SO4

2. Put 50 mL of distilled water in a 4th beaker.

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Number of

particles

activation energy

The particles to the right of the activation energy value are the only particles in the sample with enough energy to possibly react when they collide.

The particles in this area of the curve do not have enough energy to react when they collide


3. Put a clean dropper in each beaker. You will be making up various mixtures of these four solutions by counting drops. Be careful that the droppers are not switched between beakers during any of the work.

4. Make up the three solutions as described below. Mix each by flicking (“spanking”) the bottom of the tube while holding its top securely. Observe these solutions for about 10 minutes and RECORD YOUR OBSERVATIONS .

Drops of Stock SolutionsRun # KMnO4 H2O H2SO4 C2O4H2

1A 15 15 15 02A 15 25 0 53A 15 10 15 5

5. Answer the Data Analysis questions on the worksheet.

Experiment 2:Effect of H3O+ (H2SO4 acid) Concentration

In the experimental runs from here on, you will be timing how long it takes for all the MnO 41 ions to

react by watching the color fade. The recommended procedure is as follows: Dump the current contents of your test tubes into your waste beaker and rinse each a few times with about 1 mL of distilled water. Add the stock solutions to each tube as called for on the data table in the worksheet, except for the

C2O4H2. Then quickly add the drops of C2O4H2 to each tube, mix by flicking, and record this as the starting time. Also record the time when the color has faded.

1. Carry out the experiment as described in the previous paragraph, recording the time data on the worksheet.


Experiment 3:Effect of C2O4H2 (oxalic acid) Concentration

Prepare the runs shown on the data table in the worksheet using the procedure of Experiment 2 above.

1. Carry out the experiment, recording the time data on the worksheet.


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Experiment 4:Effect of Temperature

1. Using a hotplate, make a hot water bath in a half-filled 250-mL beaker. Heat the bath until it is 40-50 oC. Maintain this temperature interval by heating or allowing the bath to cool.

2. Make up the runs shown on the worksheet except for the C2O4H2. Put the test tubes in the bath for several minutes to get the solution to temperature. Remove one run from the bath, quickly add the 5 drops of C2O4H2, mix, and return to the bath. When the reaction is done, repeat the experiment on the other run.



Experiment 5:Effect of Copper Metal

1. Prepare the runs as outlined in the table on your worksheet using the procedure in Experiment 2. Cut copper metal turnings in lengths of about 2 and 4 inches.

Bunch each length of copper into a very loose ball that can be submerged in the solution and still have all of the turning exposed to solution.

Again, put everything together first except the C2O4H2, then start counting time with its addition.



Lab Report: Complete the Worksheet and turn it in at the end of this lab session, if there is time.

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Report Sheet

EXPERIMENT 1: Watching the Reaction

Drops of Stock Solutions

Run # KMnO4 H2OH3O+

(= H2SO4)C2O4H2 Observations

1A 15 15 15 0

2A 15 25 0 5

3A 15 10 15 5

Data AnalysisCTQ 7. What is the total volume (drops) in each of the 3 runs?

CTQ 8. Comparing Runs 1A and 3A: Of which reagent are you investigating its effects on the reaction rate?

CTQ 9. Comparing Runs 2A and 3A: Of which chemical are you investigating its effects on the reaction rate?

CTQ 10. Offer explanations for your observations

Did all runs show a change?

If not, why not (i.e. what might have prevented a reaction from occurring)?

Did those that changed do so instantly?

If not, why not (i.e. what might have caused a reaction to take time to react)?

CTQ 11. Based on your observations, is it necessary for all reactants to be present in order for a reaction to occur?

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Experiment 2: Effect of H3O+ (H2SO4 acid) Concentration

Drops of Stock Solutions Times (include units)

Run # KMnO4 H2OH3O+

(= H2SO4)C2O4H2 Starting Ending Total

1B 15 20 5 5

2B 15 10 15 5

3B 15 0 25 5

Data Analysis CTQ 12. What is the total volume (drops) in each of the 3 runs?

CTQ 13. Claim: What effect does increasing the concentration of the H3O+ (H2SO4) reactant have on the rate of the reaction? Answer using grammatically correct complete English sentences.

CTQ 14. Evaluate the Validity of the Collision Model Relative to This Claim: Consider the Collision Theory of Chemical Reactions Model. (p. 1 – Introduction) Based on your interpretation of the Model, how well does the Collision Model explain or support your claim? The “validity” of the model indicates how well, if at all, the Collision Model can explain your claim.

An example might begin like this: The Collision Model is very valid relative to this claim because it can explain it very well. Our claim states that increasing the concentration of a reactant increases the rate of the reaction. According to the Model, increasing the concentration of the reactant particles…

Or, if the Model does not support your claim well:

The Collision Model is NOT valid relative to this claim because it cannot explain it. Our claim states that increasing the concentration of a reactant increases the rate of the reaction. According to the Model…

Answer using grammatically correct complete English sentences.

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Experiment 3: Effect of C2O4H2 (oxalic acid) Concentration


Run # KMnO4 H2OH3O+

(= H2SO4)C2O4H2 Starting Ending Total

1C 15 10 15 52C 15 5 15 103C 15 0 15 15


CTQ 16. Claim: What effect does increasing the concentration of the oxalic acid reactant have on the rate of the reaction? Answer using grammatically correct complete English sentences.



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Experiment 4:Effect of Temperature


Run # KMnO4 H2OH3O+

(= H2SO4)C2O4H2 Temp Starting Ending Total

1E 15 10 15 5 40-50 oC

2E 15 10 15 5 40-50 oC

Data AnalysisCTQ 18. What is the total volume in each of the runs?

CTQ 19. Claim: What effect does increasing the temperature have on the rate of the reaction? Answer using grammatically correct complete English sentences.



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Experiment 5: Effect of Copper Metal


Run # KMnO4 H2O H2SO4 C2O4H2 Cu Length Starting Ending Total

1D 15 10 15 5 0

2D 15 10 15 5 2 in.

3D 15 10 15 5 4 in.


CTQ 22. Claim: What effect does the presence and amount of Cu metal have on the rate of the reaction? Answer using a grammatically correct complete English sentence. (Do NOT attempt to explain how or why you think Cu metal had this effect. Simply state what the effect is.)



(Do NOT attempt to explain how or why you think Cu metal had this effect. Simply analyze if the Model says anything about how something other than a reactant may or may not affect the reaction rate.)

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Summary Writing Assignment 1. Claim 1: What could a chemist do to this reaction to change its rate (i.e. speed it up or slow it

down)? Answer using grammatically correct complete English sentences.

List which factors are more effective.

List which factors are less effective.

2. Evidence: Provide evidence from the lab that supports your claim. Answer using grammatically correct complete English sentences.

3. Claim 2: If you were given only the overall reaction equation for some other reaction in general, could you predict none, some or all of the factors that you found affect the reaction rate?

a. List those factors you think could be predicted if you only had the chemical equation.

b. Answer using grammatically correct complete English sentences.

4. Evidence: Provide evidence from the lab that supports your claim. Answer using grammatically correct complete English sentences.

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5. Conclusion: Assess the overall validity of the Model in explaining the observations of this experiment. That is, on the whole, is the Model complete enough to explain everything you saw in this experiment? Explain why or why not. Answer using grammatically correct complete English sentences.

Does your assessment validate or invalidate the Model?

a. If the Model is not completely successful, what must be done to improve its validity?

Extension ExerciseWhile this experiment was only semi-quantitative, kinetics experiments are generally highly quantitative. In the introduction to this lab, it was stated that the rate of a reaction is dependent on a number of kinetic factors:

rate = (collision frequency)(steric factor)(fraction of collisions with E > Ea)

As stated above, the rate of the reaction is dependent upon the fraction of collisions with E > E a. As the fraction increases, so does the rate. The fraction, f, of collisions with E > Ea depends on both Ea

and the temperature, T. The general mathematical form of this relationship is actually exponential, and looks like:

fraction of collisions with E > Ea = f=e−aR ∙z

Where e is the base of the natural log, R = ideal gas constant, and a and z represent variables, in this case temperature, T and the activation energy, Ea.

Based on the Model and your observations in this experiment, place T and Ea into the correct boxes in the equation below. Recall this basic exponential identity:

e−x = 1ex

Consider how a variable (such as T or Ea) would affect the value of f if placed in either box. For example, if a variable is placed in the lower box, and you increase that variable, work out what would happen to f. Would it increase or decrease? You must correlate this analysis to your observations and predictions of how T and Ea affect the rate of the reaction.

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f = e Justify why you answered this question the way you did. That is, justify why you put T in the

box you did and Ea in the box you did by reference observations from this lab. To get credit, you justification must include references to observations from this lab.

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― R

CI2 Shifting Reactions


Shifting Reactions

Pre-Lab Assignment This pre-lab assignment is worth 5 points.

This part of the pre-lab assignment is due at the beginning of the lab period, and must be done individually before you come to lab!

I. Background Preparation Read this experiment thoughtfully FIRST:

Mentally note any procedural questions and plan how you and your partner will complete all experiments efficiently during the three-hour lab period.

II. Safety Hazards/Precautions1. Complete the following table. Note that KSCN is completed for you to

reference as an example. Use the PCC MSDS online link on your lab web page. Be sure to select the location Sylvania ST and be sure the chemical name matches precisely.

MaterialsGHS Pictograms

(Circle all that apply)

Hazard Statements

(Check and list all that apply)

potassium thiocyanate

KSCN

Corrosive Toxic ___________________ Flammable Reactive ______________ Irritant eye_____________ Other? harmful if swallowed, in contact with skin, if inhaled, to aquatic life

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Pre-lab Score: ____________/5

!


iron (III) nitrate Fe(NO3)3

Corrosive Toxic ___________________ Flammable Reactive _________________ Irritant Other? __________________

concentrated (12M)

hydrochloric acid HCl


Other hazards(equipment, glassware,

etc.)

Identify at least one hazard with the equipment that you will use during this lab.

Identify the precaution(s) you will take during your lab to avoid each hazard identified above.

2. In addition to careful handling and wearing goggles, what other precautions are needed in this experiment? If no further precautions are needed, indicate by writing “N/A.”

3. Workplace/Personal Cleanup Notes (indicate what you will do to clean up yourself and your lab space before you leave the lab):

IV. Work Plan (Procedural Flow Chart or Numbered List) (Attach additional sheet if necessary)

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Background & ProcedureQuestions of the Day

What is the nature of the reaction between KSCN and Fe(NO3)3? How does temperature affect the outcome of this reaction? How do other reagents affect the outcome of this reaction? In general, how can we shift equilibrium reactions backwards and forwards?

Record all data and observations in blue or black ink pen

Experiment 1: Data Collection: Observing a Reaction

A. Measure 50 mL of distilled water with a graduated cylinder. Pour the water into a small beaker.

B. Obtain a dropper bottle of 1 M KSCN (dissolved in 0.3 M HNO3) and a dropper bottle of 1 M Fe(NO3)3 (dissolved in 0.3 M HNO3). Note the color of each solution below:

KSCN Fe(NO3)3

C. Add 4 drops of the KSCN and 4 drops of the Fe(NO3)3 to the water in the beaker. Stir the solution with a glass stirring rod. Record your observations below:

Do you see any evidence that a chemical reaction has occurred?

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KSCN soln


D. The net ionic equation for the reaction you observed is: Fe3+(aq) + SCN(aq) FeSCN2+(aq)

Identify the color of each substance in the reaction. Save the beaker with its contents until the conclusion of the experiment. Record the colors of each species below:

SCN(aq) Fe3+(aq) FeSCN2+(aq)

Experiment 2: Data Collection and Analysis: Investigating the Reaction

A. Fill three SMALL test tubes 1/3 full of the solution from part 1.C above. Retain the remaining solution, which will be your stock reference solution.

B. Refer to this diagram below to help organize your observations, as indicated in parts C and D next.

SCN added Fe3+ added

C. Test Tube #1 (add SCN-) Reference Test Tube #2 D. Test Tube #3 (add Fe3+)

C. Into Test Tube #1 add a few drops of the 1 M SCN(aq) solution.

RECORD your observations in the box above.

COMPARE the color of this test tube to the reference solution.

D. Into Test Tube #3 add a few drops of the 1 M Fe3+(aq) solution.

RECORD your observations in the box above

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Reference


COMPARE the color of this test tube to the reference solution.

E. Dispose of the contents of the Test Tube #1 and #3 into the waste container. Retain your reference test tube solution for further study.

Critical Thinking QuestionsAt the end of this lab is a Report Sheet with CTQs that you are expected to answer DURING LAB as indicated in the procedure.

This is in addition to the making of observations and recording of data on the lab itself.

Please answer CTQs 1-7 on the Report Sheet at this time BEFORE moving on.

Experiment 3: Data Collection and Analysis: Effects of Other ReagentsA. Fill two SMALL test tubes 1/3 full of the solution you saved from Experiment 1. Retain the

remaining stock reference solution from Experiment 2.

You should have 3 identical solutions, one of which will be a reference for the next two tests.

B. Refer to this diagram below to help organize your observations, as indicated in parts C and D next.

KNO3(aq) added Na2HPO4 added

C. Test Tube #1 (add KNO3)Reference

Test Tube #2 D. Test Tube #3 (add Na2HPO4)

C. Into Test Tube #1 dissolve 1-2 drops of KNO3 solution. Mix well by spanking the tube. (Note that KSCN and Fe(NO3)3 are the original compounds reacted. What role do K+ and NO3

− play?)


COMPARE this solution to the reference solution. Do you see evidence of further reaction? In other words, were more or less products present after this addition? Explain your reasoning:

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Reference


D. Into Test Tube #3 add 3-6 drops of the 1M Na2HPO4 solution.


E. COMPARE this solution from 3.D to the reference solution. Test for the presence of a precipitate by shining a laser beam through the mixture. A true solution does not scatter light and the beam cannot be seen (example: test the reference solution with the laser. Do you see a beam?) Tiny particles (precipitate) floating in solution scatter light enabling the beam to be seen.Do you see evidence of further reaction or just the results of dilution? In other words, were more or less products present after this addition?

Explain your reasoning:

F. Refer to this diagram below to help organize your observations, as indicated in parts G and H below.

Na2HPO4 + Na2HPO4 +

SCN added Fe3+ added Reference

G. Test Tube #4 (add SCN-) Reference Test Tube #2 H. Test Tube #3 (add Fe3+)

G. Take the solution remaining in Test Tube #3 from Experiment 3.D above, and pour half of it into a new small clean Test Tube #4. Add a few drops of the 1M SCN(aq) solution.


H. To the other half of the solution in Test Tube #3 from Experiment Part 3.D above, add a few drops of the 1M Fe3+(aq) solution.


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I. Dispose of the contents of the Test Tube #1, #3 and #4 into the waste container. Retain your reference test tube solution for further study.

Critical Thinking Questions

Please answer CTQs 8-10 on the Report Sheet at this time BEFORE moving on.

Experiment 4: Data Collection and Analysis: Effects of TemperatureA. Prepare a hot water bath (not quite boiling) in a beaker in which your SMALL test tube can stand

upright. Also prepare an ice water bath in a similar sized beaker.

B. Fill two SMALL test tubes 1/3 full of the stock reference solution you saved from Experiment 1. Retain the remaining solution. Compare the two tubes and the reference tube as indicated below.

C. Place one test tube into an ice bath, one into a hot water bath, and leave one test tube at room temperature.

Place in hot water Place in ice water

Reference

Test Tube #1 (hot water) Reference Test Tube #2 Test Tube #3 (ice water)

D. After about 10 minutes:

RECORD your observations in the boxes above.

COMPARE the hot and cold tubes with the room temperature tube.

IMPORTANT NOTE: The effect in hot water will be more noticeable than that in ice water. Boiling water is about 80 oC above room temp, while ice water is only about 20 oC below room temp. Your instructor may show an example of the reference solution in a dry ice/acetone slurry that is about 20 oC for comparison.

E. Dispose of the contents of all of the test tubes into the waste container.

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Critical Thinking Questions Please answer CTQs 11-14 on the Report Sheet at this time BEFORE moving on.

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Experiment 5 - Application: Data Collection and Analysis: Cobalt(II) Complex with Water and Chloride IonIt is possible to estimate the position of some equilibria (those that contain distinct colored species) by noting the color of the solution. If more than one species is present in the same solution, then the color of the solution will be a combination of the colors of the various species present. The equilibria studied in this lab contain coordination complexes.

A coordination complex contains a central metal ion that interacts with multiple ions/molecules (called ligands) through coordinate covalent bonds (i.e. a bond where the ligand donates both electrons to form a covalent interaction with a metal). Coordination complexes can be converted to other complexes by exchanging the type of ligands interacting with the metal-this often results in a color change as observed in this lab.

Cobalt(II) ion bonds (complexes) with water and chloride ion as shown in the equilibrium below:

[CoCl4]2 (aq) + 6 H2O (l) [Co(H2O)6]2+ (aq) + 4 Cl (aq) blue pink

A. Fill a small clean, dry test tube two-thirds full with the CoCl2 in ethanol solution from your bench kit. Record your observations of this initial solution:

B. Add deionized water, one drop at a time until you have added about 10 drops or a change is observed. Try to obtain a light purplish-lavender color that is in between the blue of [CoCl4]2

and the pink of [Co(H2O)6]2-. Record your observations:

C. Divide this initial solution into three additional small clean test tubes, giving you four total test tubes with the same volume. One of the tubes will be a reference/control tube.

D. To one of the tubes, add concentrated (6 M) HCl [CAUTION! Concentrated strong acid!] until a color change occurs. Record your observations:

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E. CAREFULLY add about 1 mL (20 drops) of 0.1 M AgNO3 to a second test tube.

Caution: AgNO3 will stain your skin and clothing!Avoid skin contact and wear gloves if available.

Record your observations:

Information: Silver and chloride ions react to form solid silver chloride:

Ag+ (aq) + Cl (aq) AgCl (s)

F. Place a third tube in a hot water bath (60oC to 80oC) for ~5 minutes while shaking once in a while. Record your observations:

G. Remove the test tube from the hot water and place in an ice bath, to see whether the change to step F is reversible. Record your observations:

H. Try to reverse any of the changes above by adding reagents such as H2O, HCl or AgNO3. Perform and record at least 3 experiments and your observations:

Experiment: Observations:

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CI2 Shifting Reactions Report Sheet


Partner’s Name__________________

Report SheetExperiment 2:CTQ 1. What property of the solutions you observed will be used to monitor the concentration of reagents?

CTQ 2. What is the only reagent whose concentration you will be able to monitor effectively?

CTQ 3. In Experiment 2, the only reaction possible is: Fe3+ + SCN− = [FeSCN]2+.

If the solution becomes more blood-red in color, what conclusion can you draw about the amount of [FeSCN]2+ now present (circle one)?

less [FeSCN]2+ the same [FeSCN]2+ more [FeSCN]2+

What is the only source for the production of this additional [FeSCN]2+?

If the solution becomes less blood-red in color, what conclusion can you draw about the amount of [FeSCN]2+ now present (circle one)?

less [FeSCN]2+ the same [FeSCN]2+ more [FeSCN]2+

To where did this additional [FeSCN]2+ go?

CTQ 4. When additional SCN(aq) was added (Step 2.C), what ion had to be present in the solution to account for your observations (i.e. with what did the SCN− react)?

CTQ 5. When additional Fe3+(aq) was added (Step 2.D), what ion had to be present in the solution to account for your observations (i.e. with what did the Fe3+ react)?

Conclusion. Before the addition of extra reactants (Steps 2.C and 2.D), was the reaction complete? In other words, had all of both reactants reacted completely into products? Explain why or why not. That is, how do you know that when the reaction appeared to stop, there were still original reactants remaining unreacted? (Hint: This is NOT a limiting reactant question. Assume that equi-molar amounts of each reactant were mixed, noting the reaction is 1:1 in each reactant.)

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Mental Model. We are going to make a model of the system to help conceptualize what is occurring at the particulate level. Below is a picture at the level of ions of the substances in solution after mixing the original solutions from Part I – your reference solution (water molecules omitted for clarity).

In this model, take note of the amount (number) of each species present (these are representative particles – they could represent moles of each species or some fraction of moles – the key is the proportion).

CTQ 6. Step 2.C: Redraw the diagram above at the level of ions, showing what happens after adding 3 more ions of SCN− to the above mixture. Complete the table next to the box by paying attention to the number of each species you are starting with (see above), and your conclusions from CTQ 1-5.

CTQ 7. Step 2.D: Draw a picture at the level of ions based on the diagram in the Mental Model, showing what happens when additional Fe3+ is added to the above mixture. (Ex: When 3 more Fe3+ are added to the above reference solution). Pay attention to the number of each species you are starting with (see above), and your conclusions from CTQ 1-5. Please create a model and table similar to that in CTQ 6.

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For the “Change” column, choose an arbitrary number, such as “2” or


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Experiment 3:CTQ 8. Based on your observations from Experiments 3.D and 3.G (adding 1M Na2HPO4

followed by the addition of SCN(aq)), and your observations of Experiments 3.D and 3.H (adding 1M Na2HPO4 followed by the addition of Fe3+(aq)), with what would you say the PO4

3 (aq) reacted? Explain your reasoning.

CTQ 9. Consider your observations from the reaction of PO43 ion (from Na2HPO4). Do you see a

precipitate? How would you describe the solubility of phosphate compounds? Provide evidence from this lab to support this claim.

CTQ 10. Now consider the FeSCN2+ compound formed in this experiment. The SCN– ion is known to complex with metal ions. In this case, thiocyanate ion is called a ligand (lig’-and). A metal ion-ligand complex consists of a central metal ion covalently bonded to two or more anions or molecules (ligands). Hydroxide, chloride, and cyanide ions are also some examples of ionic ligands; water, carbon monoxide, and ammonia are some molecular ligands. At right is a picture of the SCN–

complexed to an Fe3+ ion. Many metal complexes can accommodate up to 6 ligands. In FeSCN2+, the other 5 spots are filled with water molecules. So the more precise formula of the complex is [Fe(SCN)(H2O)5]2+ and is called:

pentaaqua(thiocyanato)iron(III) ion!

What evidence is there that the combination of Fe3+ and SCN–

forms a soluble complex ion, FeSCN2+, and not an insoluble compound (precipitate) such as Fe(SCN)3(s)?

Experiment 4:

CTQ 11. What was the effect of increased temperature? Were more or fewer products produced in the hot solution? (Circle your answer). Explain how you know this.

CTQ 12. What was the effect of decreased temperature? Were more or fewer products produced in the cold solution? (Circle your answer). Explain how you know this.

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General Conclusions:CTQ 13. The diagram below shows another model of our reaction system. It is a plot of hypothetical concentrations of each reactant and product over time. The plot starts with a representation of the concentrations of species in the reference solution, and represents the initial equilibrium position of the system (that is, a set of concentrations of reactants and products that are stable and unchanging over time, if left undisturbed).

At the time marked with the dashed line, a little Fe3+ is added to the mixture (as in Experiment 2.D.) That is to say, the equilibrium mixture is “spiked” with Fe3+ at this point.

Draw how the concentrations of EACH species will change as the system tries to establish a new equilibrium position. Use the following considerations to help youo If the system was at equilibrium before it was spiked with Fe3+, could it still be at equilibrium

immediately after the spike?o As far as the equilibrium is concerned, immediately after spiking, the “problem” is too much

Fe3+. How is this problem “solved”? The [Fe3+] must be reduced!o The only way to do so is to REACT away the Fe3+. The only thing in this mixture (Expt. 2)

that Fe3+ reacts with is SCN−.o If Fe3+ and SCN− react, what is produced? (Hint: what happened to the color of the solution

after spiking with Fe3+?)o Note: Will ALL the Fe3+ react away? How will the [Fe3+] at the NEW equilibrium position

produced after the spike compare to before the spike?

CTQ 14. Draw how the concentrations of EACH species will change as the system tries to establish a new equilibrium position after PO4

3− is added. Note that [PO43−] is not part of the equilibrium

system. However, it does have an effect on the system, as described earlier.

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Exercises:

Experiments 1-4: Fe3+(aq) + SCN(aq) [FeSCN]2+(aq)

1. The potassium and nitrate ions do not participate in the reaction you studied. What term(s) describe ions that have this role?

2. Why are these ions (in Exercise 1) necessary in the reaction in the first place?

3. When a reaction is reversible, the products can react to become reactants and the reactants simultaneously react to become products. The reaction does not “go to completion”, and an equilibrium may exist. We show an equilibrium with double arrows:

Fe3+(aq) + SCN(aq) [FeSCN]2+(aq)

Do your observations for the addition of extra reactants and temperature change indicate that this reaction is an equilibrium reaction? Explain clearly how your observations indicate a reaction that does not go to completion.

HINT : Consider the results of Experiment 2 and Experiment 4. How can these observations support your claim?

4. When you do something to a reaction at equilibrium to cause it to consume reactants and produce more products, we say that it shifts to the right. If the reaction you studied exists as an equilibrium, which of the following caused a shift to the right? (circle all that apply)

addition of a reactant addition of KNO3 addition of Na2HPO4

increased heat decreased heat

5. When you do something to a reaction at equilibrium to cause it to consume products and produce more reactants, we say that it shifts to the left. If the reaction you studied exists as an equilibrium, which of the following caused a shift to the left? (circle all that apply)

addition of a reactant addition of KNO3 addition of Na2HPO4

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increased heat decreased heat

6. In this reaction system, PO43− binds with Fe3+ and “removes” it from effectively being part of our

reaction (it’s still there of course, but it is “bound” up). We could “think” of the PO 43− as a

“molecular tweezers” if you like, sequestering the Fe3+ from the reactive system. In general, what is the effect of removing a reactant from a reversible reaction?

7. Which of the following (A or B) correctly describes the placement of “heat” in this reaction? (circle the letter)

A. Heat + Fe3+(aq) + SCN(aq) [FeSCN]2+(aq)

B. Fe3+(aq) + SCN(aq) [FeSCN]2+(aq) + Heat

8. Explain how you arrived at the conclusion in Exercise 7. (Hint: treat “heat” as a reactant or product)

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Experiment 5 – Application [CoCl4]2 (aq) + 6 H2O (l) [Co(H2O)6]2+ (aq) + 4 Cl (aq)

9. What were the major cobalt complex species in the initial reference solution that was light purple-lavender in color? How do you know this? Think carefully by considering the colors of [CoCl 4]2−

and [Co(H2O)6]2+. Did the reference solution match either of these? What does this imply?

10. Which way did the above equilibrium shift when 6 M HCl was added? (circle one)

Left Right No shift

11. Explain in terms of Le Châtelier’s principle your answer to #10 above.

12. Which way did the above equilibrium shift when 0.1 M AgNO3 was added? (circle one)

Left Right No shift

13. Explain in terms of Le Châtelier’s principle your answer to #12 above.

14. Is the reaction above exothermic or endothermic as written? (circle one)

endothermic exothermic neither

15. Explain in terms of Le Châtelier’s principle how you determined your answer to #14 above.

16. Discuss and explain the experiments you carried out in part H. Clearly discuss what you did and what you observed. Were you able to reverse any of your original changes? If you did, what does this imply?

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Summary Writing Assignment (may be typed)Connection:Based on your experience in this lab, draw a connection to something in your everyday life or the world around you (something not mentioned in this lab). This will probably take some thought and discussion and may not be as easy to consider correctly as originally thought.

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CI3 Determination of Kc for Phenolphthalein


Determination of the Equilibrium Constant of

Phenolphthalein Dissociation



I. Background Preparation1. Read this experiment thoughtfully FIRST:


2. Pre-lab questions:

A. Which form of phenolphthalein, HIn or In− is pink? Circle your choice: HIn In−

B. Define an acidic solution.

C. Define a basic solution.

D. What is the pH of a solution whose [H3O+] = 1.58 × 10−9? Show your work!

pH =

E. If the pH of a solution is 8.6, what is [H3O+]? To “undue” the pH log, you take the antilog: [H3O+] = 10−pH. Show your work!

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Pre-lab Score: ____________/5


[H3O+] =

F. Which form of phenolphthalein do we expect to predominate in basic solutions?

Circle your choice: HIn In−

G. How will you experimentally determine [H3O+] in the experiment?

H. How will the pH (i.e. [H3O+]) of each solution be fixed at a constant value?

I. How will the concentration of In− be determined in this experiment?

J. What is the calculation used to determine [HIn]?

K. What are the two methods you will use to determine a value of Kc?

Method 1:

Method 2:

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L. The empty graph below is missing axis labels, a title and identification of the slope and y-intercept. Complete each box with the appropriate variable or value used when using the graphical method of analysis in this experiment.

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y-axis label

x-axis label:

Graph title:y-intercept =

slope =


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ChemInquiry3

Background & ProcedureThe goal of this lab is to determine Kc for the acid-base indicator phenolphthalein. Phenolphthalein (C20H14O4; M = 318.32 gmol−1) is a weak acid. As you will discover later in the course, a weak acid is type of molecule that reacts with water and dissociates into a weakly basic anion and a hydronium ion; it is considered a weak acid in that this dissociation does not go to completion-that is, it establishes an equilibrium:

HIn(aq) + H2O(l) In−(aq) + H3O+(aq) Kc = ¿¿ Reaction 1 and Equation 1Where:

HIn = phenolphthaleinIn− = phenolphthalein after losing an H+

Because H2O is the solvent, its concentration will not change significantly as the reaction proceeds (the amount that reacts with HIn is insignificant compared to its initial concentration ( 55 M). Thus Kc is a “working” equilibrium constant where [H2O] is constant and is folded into the value of Kc

You should recognize phenolphthalein from earlier discussions in lecture. Specifically, we say that phenolphthalein has two forms, the “acidic” form, HIn, also called the “protonated” form, and the “basic” form, In−, also called the “deprotonated form, as shown below:

HIn: pH 0 to 8.2 In−: pH 8.2 to 10 Colorless form Pink form“Protonated form” “Deprotonated form”

Figure 1 Acidic (left) and basic (right) forms of phenolphthalein

When phenolphthalein is dissolved in water, it partially dissociates into ions as noted earlier:

HIn + H2O In− + H3O+

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Since one form is colorless (protonated form, HIn) and the other form is pink (deprotonated form, In−), we can estimate the relative concentrations of the two forms by observing the color of the solution. But what factors might serve to push this equilibrium either toward reactants (colorless) or products (pink)? We will rely on Le Châtelier’s principle to provide insight.Basic Definitions

Aqueous solution: Any solution in which water is the solvent. In ALL aqueous solutions, water molecules react with themselves to a small extent to produce both H3O+ and OH− ions:

H2O + H2O H3O+ + OH− Kc = 1.0 × 10−14 @ 25oC

In pure water these ions are of equal concentration ([H3O+] = [OH−]). Adding an acid or base can change the relative concentration of H3O+ and OH− ions as follows:

Acidic solution: Any aqueous solution in which [H3O+] > [OH−] Basic solution: Any aqueous solution in which [OH−] > [H3O+] pH: pH = −log[H3O+]; @ 25oC, if pH < 7 = acidic solution, if pH > 7 = basic solution

As mentioned, phenolphthalein changes color when its protonation state changes: In solutions with relatively high background [H3O+] (i.e. acidic solutions), the HIn form (“acidic” form) predominates, and this form is colorless; In solutions with relatively high [OH−] (i.e. basic solutions) the In− form (“basic” form) predominates, and this form is pink.

We can understand why phenolphthalein changes color depending on the pH by invoking Le Châtelier’s principle, which was studied in a previous lab:


In acidic solutions (solutions where the background concentration is high in H3O+), Le Châtelier’s principle says that the equilibrium will be shifted to the left (reactants side), and thus we should expect HIn to be the predominant form and the solution to be colorless.

In basic solutions (solutions where the background concentration of OH− is high), the OH− reacts with the H3O+ produced in the phenolphthalein equilibrium to produce water, as shown here:


OH− + H3O+ 2H2O Kc = 1.0 × 1014 @ 25oC

Essentially, this removes H3O+ from the phenolphthalein equilibrium, and Le Châtelier’s principle says that the equilibrium will be shifted to the right (products side) , and thus we should expect In− to be the predominant form and the solution to be pink. This effect is similar to adding PO4

3− to the Fe3+

solution in the previous week’s lab, CI2. A separate chemical species reacts with one of the species in the equilibrium reaction, causing a perturbance and thus an equilibrium shift.

The relative position of the equilibrium (i.e. relative concentrations of HIn and In−) depend on [H3O+]. As the pH of a solution rises above 7.0, the concentration of H3O+ decreases as OH− increases, and this causes the equilibrium to shift further and further to the right. Above a pH = 8.2, we start to see a

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predominance of the pink In− form of phenolphthalein. Thus, we expect the intensity of the pink color to increase as pH rises above 8.2.

In this lab you will prepare a series of phenolphthalein samples at different [H3O+] (i.e. pH) values. Each solution should contain different relative concentrations of HIn and In− in equilibrium and should therefore have different color intensities. Based on equation 1 on the previous page, we can determine an experimental value for Kc by measuring the equilibrium values of [HIn], [In−] and [H3O+]. The value of [H3O+] is measured experimentally as the pH of the solution, where pH is defined as –log [H 3O+]. The pH in each solution will be fixed at a constant value by using a buffer (a buffer is a substance that limits changes in pH).

To determine [HIn] and [In−] at different [H3O+] values (pH values), you will utilize an experimental ICE table:

Reaction HIn + H2O In− + H3O+

I (initial) (HIn)o − (In−)o = 0 (H3O+)o

C (change) −x − +x no change (buffered)

E [equilibrium] [HIn] = (HIn)o − x − [In−] = x [H3O+] = (H3O)o

Initial line in table:(HIn)o is known (prepared stock solution)(In−)o = 0(H3O+)o is known and pre-set by the pH of buffer

Equilibrium line in table:[In−] will be determined by measuring the absorbance of the solution at 550 nm (A550)

Note: By examining the table, note that “x” = [In−][H3O+] = (H3O+)o (the buffer keeps this at a constant value even though H3O+ is generated during the conversion of HIn to In−; that is, any H3O+ generated by HIn is absorbed by the buffer and the concentration of H3O+ stays constant for each trial.)[HIn] = (HIn)o – x*Note that since (In−)o = 0, x = [In−], and thus [HIn] = (HIn)o – [In−]

Once the equilibrium values are determined, one can substitute [HIn], [In−] and [H3O+] into Equation 1 to calculate Kc for that specific [H3O+].

There are two methods to determine a value of Kc for phenolphthalein in this experiment.

Method 1: By varying [H3O+] (i.e. using different pH buffer solutions), you will determine a series of Kc values, which can be averaged together to get a final average Kc value. (Note that we expect Kc to be constant at constant temperature, regardless of starting conditions!)

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Method 2: There is also a graphical method that works well to calculate Kc. If one takes the (–log) of both sides of Equation 1 and rearranges, the result is Equation 2:

−log Kc = −log[H3O+] − log[In- ][ HIn ]

But because −log[H3O+] = pH: −log Kc = pH − log[ In- ][ HIn ]

Rearranging and solving for pH gives:

pH = log[ In- ] [HIn ]

− log Kc

Equation 2

y = m x + b

Equation 2 is the equation of a straight line (y = mx + b), where:y = pH

x = log [In- ][ HIn ]

m (slope) = 1

y-intercept = −log Kc

Thus, a plot of measured pH vs. log[ In- ][ HIn ]

should yield a straight line with slope = 1 and a y-intercept

equal to –log Kc.

-1.4 -1.2 -1 -0.8 -0.6 -0.4 -0.2 08.2

8.4

8.6

8.8

9

9.2

9.4

9.6

9.8

10

pH

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pH vs. log[ In- ][ HIn ]

log[ In- ][ HIn ]

y-intercept = −log Kc


Procedure

I. Preparation of Required Solutions

Work in a group of 4 (2 Teams) to prepare the phenolphthalein buffer solutions as follows:

In clean containers, each Team should obtain the following solutions from the reagent cart:

o ~ 8 mL of the 2.05 × 10−4 M phenolphthalein (phth) solution

o ~ 20 mL of assigned buffer solutions, labeled pH 8.8 through pH 9.6

Divide preparation of phenolphthalein (phth) into two groups: Share buffer solutions in groups of 4 at lab bench Team A prepares phth at pH = 8.8, 9.0, 9.2 Team B prepares phth at pH = 9.4, 9.6, 9.8

o ~10 mL of an unknown solution

Prepare your Team’s 3 buffer solutions:o Add 2.0 mL of phth + 18.0 mL of the appropriate buffer solution using a 25-mL graduated

cylinder as follows: Very carefully fill the graduated cylinder up to the 18.0 mL line with the first buffer

solution. Use a dropper to add the last few drops to get precisely at the 18.0-mL mark. Add phenolphthalein to bring the total volume precisely to the 20.0-mL mark. Use a dropper

to add the last few drops to get precisely at the 20.0-mL mark. NOTE: the graduated cylinder for buffer must be rinsed and PRIMED with each new

solution before measuring out sample volume.

Store each solution in a 24-mL capped vial. Be sure to label the vial with a wax pencil.

BE SURE TO LABEL the pH of each solution with a wax pencil!

NOTE: Each Team (pair of students) measures all six solutions, but shares its 3 prepared solutions with the other Team and vice versa.

Prepare a blank solution. A blank solution contains all chemical species except the species that will be measured with the spectrometer. For the blank, use a pH 9.0 solution with NO phth.

Prepare your Team’s UNKNOWN solution:o Add 1.0 mL of phth + 9.0 mL of the UNKNOWN solution using a 10-mL graduated cylinder

as follows: NOTE: the graduated cylinder for the UNKNOWN must be rinsed and PRIMED with a

small amount of UKNOWN solution before measuring out the sample volume. Very carefully fill the graduated cylinder up to the 9.0 mL line with the UNKNOWN

solution. Use a dropper to add the last few drops to get precisely at the 9.0-mL mark. Add phenolphthalein to bring the total volume precisely to the 10.0-mL mark. Use a dropper

to add the last few drops to get precisely at the 10.0-mL mark. Store UNKNOWN solution in a 24-mL capped vial. Be sure to label the vial with a wax pencil.

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Prep

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II. Measurement of pH of Phenolphthalein Solutions A. Set up the LabPro pH sensor for data collection by connecting the pH Sensor to the LabPro by

inserting the cable into Channel 1. Be sure your LabPro has 3 connections: Power; Probe; Computer connection (USB cable).

B. Start the LoggerPro program. Open the file “CI3 pH” by going to the “Open” menu in LoggerPro CH 223 CI3 pH (Or as directed by your instructor).

C. Raise the pH Sensor from the sensor soaking solution and set the solution aside. Use a wash bottle filled with distilled water to thoroughly rinse the pH Sensor. Catch the rinse water in a 250-mL beaker. Dry the probe by GENTLY blotting it with a Kim-wipe (NOT a brown paper towel).

D. Calibrate the pH meter using 2 buffer solutions as follows:

1. Obtain some pH-7 and pH-10 (also known as the “pH soaking solution”) buffer solution into separate vials.

First Calibration Point

2. Choose Calibrate ► CH1:pH from the Experiment pull-down menu and then click Calibrate Now.

3. For the first calibration point, rinse the pH sensor with distilled water, then place it into a buffer of pH 7.0.

4. Type “7” in the edit box as the pH value.

5. Swirl the sensor, wait until the voltage for Input 1 stabilizes (wait 15 seconds), then click “Keep.”

Second Calibration Point

6. Rinse the pH sensor with distilled water, and place it into a buffer of pH 10.0.

7. Type “10” in the edit box as the pH value for the second calibration point.

8. Swirl the sensor and wait until the voltage for Input 1 stabilizes. Click “Keep”, then click “Done.” This completes the calibration.

E. Measure the pH of the first solution. Measure the pH by immersing the probe into the vial for each solution. Be sure the solution completely covers the glass bulb at the bottom. Gently swirl the probe in the solution and when the pH reading displayed on the screen stabilizes, record the pH value (round to the nearest 0.01 pH unit).

F. Prepare the pH sensor for reuse:

1. Rinse it with distilled water from a wash bottle.2. Place the sensor into the sensor soaking solution and swirl the solution about the sensor

briefly.3. Rinse with distilled water again.

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Take Note!!


G. Determine the pH of all the other solutions. You must clean the pH Sensor between tests, using the procedure above each time. Be sure to gently blot dry the probe with a Kim-wipe.

III. Measuring Absorbance at 550 nm of Phenolphthalein Solutions A. Connect the spectrometer to the USB port on your computer.

B. Open the program LoggerPro by clicking on the icon at the bottom of the screen.

C. Open the file “CI3 Phth Absorbance” by going to the “Open” menu in LoggerPro CH 223 CI3 Phth Absorbance (or as directed by your instructor).

D. Obtain a small rectangular “cuvette.” (pronounced kyoo-vet’) Make sure it is clean.

E. Prime and fill the cuvette 2/3 full with your blank solution.

See: http://www.chem.vt.edu/RVGS/ACT/lab/Experiments/cuvettes.html

F. Calibrate the Spectrometer.

1. Open the Experiment menu and select Calibrate → (Spectrometer). The following message appears in the Calibrate dialog box: “Waiting … seconds for the device to warm up.” After 60 seconds, the message changes to: “Warmup complete.”

2. Place the blank solution in the cuvette holder of the Spectrometer. Align the cuvette so that the clear sides are facing the light source of the Spectrometer. Click “Finish Calibration”, and then when the button is highlighted, click .

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G. Collect absorbance-concentration data for the six phth solutions + unknown (all of these will be individually measured in the same cuvette).

1. Rinse the cuvette 3 times with DI water.2. Rinse the cuvette 3 times with a small portion of the first sample (Don’t waste solution by filling

the cuvette. Simply use a dropper to put a few drops of sample in the cuvette, then swirl/maneuver the cuvette to rinse the inside walls into the waste beaker.)

3. Fill the cuvette 2/3 full with the first phth solution.4. Place the first phth solution cuvette in the spectrophotometer and click the button on the

Toolbar to start data collection. When the absorbance reading stabilizes, record the Absorbance of the solution in the data table.

5. Discard the cuvette contents into a waste beaker. 6. Rinse the cuvette 3 times with DI water.7. Rinse the cuvette 3 times with a small portion of the next sample (Don’t waste solution by filling

the cuvette. Simply use a dropper to put a few drops of sample in the cuvette, then swirl/maneuver the cuvette to rinse the inside walls into the waste beaker.)

8. Fill the cuvette 2/3 full with the next sample. Wipe the cuvette and place it in the spectrometer. When the absorbance reading stabilizes, record the Absorbance of the solution in the data table.

9. Repeat Steps 2-5 for the remaining samples of phth, and the unknown (all to be individually measured in the same cuvette)

10.When finished, hit the STOP button to stop data collection.

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CI3 Determination of Kc for Phenolphthalein Report Sheet

ChemInquiry3Name__________________________

Partner’s Name__________________

Report SheetData Collection and AnalysisII. Measurement of pH of Phenolphthalein Solutions

SampleID Measured pH

8.8

9.0

9.2

9.4

9.6

9.8

III. Measurement of Absorbance at 550 nm of Phenolphthalein Solutions

Record absorbance at 550 nm for the phenolphthalein samples and the unknown:

SampleID A550

8.8

9.0

9.2

9.4

9.6

9.8

Unknown

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Record your Unknown # here:


IV. Prediction: Based on the color intensity of your Unknown sample, predict what you think is the approximate pH of the Unknown:

AnalysisI. Calculate (H3O+)o from the measured pH for each phth sample:

SampleID (H3O+)o (M)

8.8

9.0

9.2

9.4

9.6

9.8

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Predicted approx. pH of Unknown:

Show one sample calculation here:

Example Calculation:

If the measured pH for some sample was 9.23, then:

[H3O]+ = 10−9.23 = 5.89 × 10−10 M


II. Calculate [In−] from the measured A550 using Beer’s law for each phth sample. Beer’s law states that absorbance is proportional to the concentration of a solution via the following relationship:

A = ϵlCWhere

A = absorbance at 550 nm

ϵ = the “molar absorptivity” = 29,300 1

M ∙ cm (at = 550 nm for phth)

l = path length of light through sample cuvette = 1.00 cm for all your samples (= width of cuvette)C = molar concentration of species = [In−]

Thus, for your solutions, this will be:

[In−] = A550

29,300 1M ∙cm

(1.00 cm)

SampleID [In−] (M)

8.8

9.0

9.2

9.4

9.6

9.8

Unknown

III. Calculate [HIn] for each phth sample. A. Recall that each phth solution was diluted 1:10 with buffer (2.0 mL phth to 18.0 mL buffer). The

stock (HIn) = 2.05 × 10−4 M. Based on this 1:10 dilution, compute the diluted (HIn)o.

(HIn)o =

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If the A550 for some sample = 0.048, then:

[In−] = 0.048

29,300 1M ∙cm

(1.00 cm) = 1.64 ×



B. [HIn] is calculated as follows: Reaction HIn + H2O In− + H3O+

I (initial) (HIn)o − 0 (H3O+)o

C (change) − x − + x no change (buffered)E [equilibrium]

[HIn] = (HIn)o − x − [In−] = x [H3O+] = (H3O+)o

Initial line in table:(HIn)o = computed in part A. above(In−)o = 0(H3O+)o is known (from measured pH of buffer)

Equilibrium line in table:[In−] was determined by measuring the absorbance of the solution at 550 nm (A550)[H3O+] = (H3O+)o

[HIn] = (HIn)o – x = (HIn)o – [In−] (because [In−] = x, see Table above)

SampleID [HIn] (M)

8.8

9.0

9.2

9.4

9.6

9.8

Unknown

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If (HIn)o = 2.05 × 10−5 M, and the measured [In−] = 1.64 × 10−6 Mthen:[HIn] = (HIn)o – [In−] = 2.05 × 10−5 M − 1.64 × 10−6 M = 1.89 × 10−5 M




IV. Determine K c by Using Equilibrium Constant Equation

Calculate Kc for each phth sample using the equilibrium constant equation:

Kc = ¿¿

SampleID Kc

8.8

9.0

9.2

9.4

9.6

9.8

Average Kc

V. Determine K c by Using Graphical Method

1. Calculate log[In- ][ HIn ]

for each phth sample:

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[HIn] = 1.89 × 10−5 M[H3O]+ = 5.89 × 10−10 M[In−] = 1.64 × 10−6 M

Kc = ¿¿ = (1.64 × 10−6)(5.89× 10−10)

(1.89 ×1 0−5)

Kc = 5.11 × 10−9

SampleID log[ In- ]

[ HIn ]

8.8

9.0

9.2

9.4

9.6

9.8


2. Plot Measured pH vs. log[ In- ][ HIn ]

. Use MS Excel or LoggerPro to prepare an acceptable graph which

includes the following:1) Graph title2) x- and y-axis labeled with variables, not x or y3) Best-fit linear trendline − re-written with actual variables, not x or y4) Equation of trendline and R2 value5) Attach copy of graph to this Report Sheet

3. Report the equation of the best-fit line, slope and y-intercept:

Equation of line(do NOT use “x” and “y”, use the

variables plotted)

slope of line

y-intercept

4. Report the value of Kc using the results of graphical analysis.

Value of Kc

VI. Determine pH of Unknown Sample

Determine the pH of the unknown sample. Use the equation of your best-fit line to predict the pH of the unknown. Note that you determined [In−] and [HIn] for your unknown. You also now have a value of Kc. Use these values and the equation of the line to predict the value of pH of the unknown. Show all work below that details the method and values used to determine the pH of your unknown.

Unknown #

pH

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Note!


Conceptual Check: Does the pH for your unknown correlate with your predicted pH based on the intensity of the color of the solution? Discuss briefly using grammatically correct English sentences.

Reflection Questions1. Based on the average value of Kc that you determined, classify the phenolphthalein reaction as either

reactant- or product favored. Explain your choice.

2. Assume that for the pH = 9.2 phenolphthalein sample you used half as much phenolphthalein solution (i.e. your sample was prepared with 18.0 mL buffer and 1.0 mL phenolphthalein solution and 1.0 mL water). What should happen to the value of Kc you determine?

3. The value of Kc of phenolphthalein is commonly reported as 5.01 × 10−10. Calculate the percent error for the Kc value that you determined two different ways. %-error is calculated as:

%-error = | actual-experimental |actual x 100%

Show your work and report your %-error for each method of determining Kc.%-error by using averaging the

Kc values method%-error by using the

graphical method

4. Your error may be between 50-100%, roughly. That’s OK, it’s expected. You should reflect on the fact that you measured a very small value of K and came within about 1 order of magnitude of the actual value. For this class, that is very good! Which technique used do you think was the most prone to introducing error into this

experiment?

In what way do you think the above technique affected the error in your determination of K c? In other words, which part of the calculation do you think it affected the most? Why do you think so?

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Which method, taking the average of a series of computed Kc’s, or using the graphical method, gave the most accurate result? Propose a reason why this would be so.

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L54

CI4 What Factors Affect the Solubility of Ions?


What Factors Affect the Solubility of Ions?





II. Safety Hazards/Precautions1. Complete the following table. Use the PCC MSDS online link on your lab web

page. Be sure to select the location Sylvania ST and be sure the chemical name matches precisely.



Hazard Statements


Ba(NO3)2

barium nitrate


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!

Pre-lab Score: ____________/5


Co(NO3)2

cobalt (II) nitrate




Hazard Statements


Ni(NO3)2

nickel (II) nitrate


Cr(NO3)3

chromium (III) nitrate


AgNO3

silver nitrate



etc.)



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H. In addition to careful handling and wearing goggles, what other precautions are needed in this experiment? If no further precautions are needed, indicate by writing “N/A.”

I. Workplace/Personal Cleanup Notes (indicate what you will do to clean up yourself and your lab space before you leave the lab):

J. Pre-lab Question: What are the three ways the chemical equation for reaction in solution can be written?

III. Work Plan (Procedural Flow Chart or Numbered List)

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Background & Procedure

Introduction Suppose you were an employee at a local company with plans to produce test kits for identifying ions in aqueous solutions. You are part of a set of research teams to study the solubility chemistry of ions. Teams are to collaborate and develop a plan for identifying metal ions in aqueous solution based on solubility data. The company also wishes to know what factors affect ion solubility. When you are done with your investigation, you will have a better idea of the factors affecting solubility as they are related to the periodic properties of some of the elements and other factors.

Questions of the Day:

What is the design of the tests and the logic used to determine the identity of reacting ions forming precipitates?

How can we collect and analyze data on the solubility behavior of metal ions in aqueous solution to determine whether metal ion solubility is predictable from metal ion characteristics?

How can we design experiments to determine the identity of metal ions in a sample of water based on collected solubility data?


Procedures

The research plan is to:

1. Become acquainted with the design of tests and the logic used for identifying ions in solution (Part I).

2. Different teams investigate and share data about the solubility of different ions to determine what factors affect ion solubility (Part II).

3. Develop and test a plan to identify metal ions in aqueous solution (Part III).

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Part I: What is the Precipitate Identity?

When the positive metal ion (cation) of a dissolved salt combines with the negative ion (anion) from a different dissolved salt, the recombined ions may either stay in solution or come out of solution in the form of a solid called a ‘precipitate.’ In this part of your inquiry, you will become familiar with the design of tests and the logic used to determine the identity of the reacting ions forming a precipitate. Ions that are present and do not precipitate but remain in solution are called spectator ions.

Part IA Reaction of Calcium Chloride and Sodium Oxalate

You will analyze the reaction that occurs upon mixing solutions of calcium chloride and sodium oxalate: CaCl2 + Na2C2O4 precipitate. Your goal is to determine the identity of the precipitate by conducting tests using other solution mixtures containing three of the four ions in the calcium chloride and sodium oxalate reaction and observing whether the precipitate forms or does not form. The resulting observations will allow you to identify the reactant ions forming the precipitate in the calcium chloride and sodium oxalate reaction.

Caution: Do not dump any of the reagents down the sink. Discard the waste in an appropriate waste container. Do not allow the solutions to come in contact with your skin.

1. Obtain 2 mL of 0.10 M calcium chloride, CaCl2, and 2 mL of 0.10M sodium oxalate, Na2C2O4, and mix in a test tube. In the table below, record the appearance of each individual solution and the combined mixture. Label and save the mixture for reference.

Reagent 1 Reagent 2 New Ion Combos Possible After Mixing?

Observations of mixture

(precipitate?)

Reagent CaCl2 Na2C2O4

Ions When Dissolved

Ca2+(aq) + Cl(aq)

Na+(aq) + C2O4

2(aq)Ca2+ + C2O4

2

Na+ + Cl

Observations

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2. Conduct tests of the reacting system to determine the identity of the ions forming the precipitate. Note that the following test solutions are the same concentration (0.10 M) as that in the reacting system under analysis. In addition, identical volumes (2 mL) are combined for the tests in each test tube. In the table below, write the new ion combinations that are possible and the observations for the resulting mixture in each test.

Test Reagent 1 Ions Reagent 2 Ions New Ion Combos Possible

Observations (precipitate?)

10.10 M CaCl2

Ca2+(aq) + Cl(aq)

0.10 M NaNO3

Na+(aq) + NO3(aq)

20.10 M Ca(NO3)2

Ca2+(aq) + NO3(aq)

0.10 M Na2C2O4

Na+(aq) + C2O42(aq)

30.10 M NaCl

Na+(aq) + Cl(aq)

0.10 M Na2C2O4

Na+(aq) + C2O42(aq)

3. Compare your test results above with the results from the original reaction of calcium chloride and sodium oxalate. In the box below identify and record any ion combinations that stay in solution and thus must be spectator ions in the CaCl2 and Na2C2O4 reaction.

Spectator Ions

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Critical Thinking Questions Answer Critical Thinking Questions #1-2 on the Report Sheet at this time BEFORE moving on.

Part IB Reaction of Calcium Chloride and Sodium Sulfate

Design tests and use your logic to confirm the identity of the precipitate that forms upon mixing solutions of calcium chloride and sodium sulfate: CaCl2(aq) + Na2SO4(aq) precipitate.

1. From the box labeled Part IB, obtain bottles of 15 g/150 mL calcium chloride, CaCl2, and the bottle of 15 g/150 mL sodium sulfate, Na2SO4. Mix 2 mL of each in a test tube and observe the results. You may have to wait 1-2 min and scratch the test tube for a reaction to appear. In the table below, record the appearance of each individual solution and the combined mixture. And write the new ion combinations that are possible in each test similar to that from Part 1A-1. Label and save the mixture for reference.

Reagent 1 Reagent 2 New Ion Combos Possible After Mixing?

Observations of mixture

(precipitate?)

Reagent CaCl2 Na2SO4

Ions When dissolved

Observations

2. Based on your logic and the information gained in Part IA about the solubility of some ion combinations, what is the likely identity of the precipitate? Record your hypothesis in the box below.

Identity of Precipitate

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3. Design and carry out tests to determine and confirm the identity of the ions forming the precipitate. Complete the table below (similar to the table in Part IA) to record your tests, new ion combos, and observations. For example, can you design a test that omits sulfate ion in order to determine whether it is critical to precipitate formation?

Test Reagent 1 and Ions Reagent 2 and Ions New Ion Combos Possible

Observations (precipitate?)

1

2

3

Critical Thinking Questions Answer Critical Thinking Question #3 on the Report Sheet at this time BEFORE moving on.

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Data Analysis and Implications (Part I)The chemical equation for a reaction in solution can be written in three ways.

The overall chemical equation (sometimes called the “molecular equation”) shows all the substances present in their undissociated forms.

The complete (total) ionic equation shows all the substances present in the form in which they actually exist in solution.

The net ionic equation is derived from the complete ionic equation by omitting all spectator ions, ions that occur on both sides of the equation with the same coefficients. Net ionic equations demonstrate that many different combinations of reactants can give the same net chemical reaction.

Answer Data Analysis and Implications Questions #1-2 on the Report Sheet at this time BEFORE moving on.

Part II: What Factors Determine Solubility?

Why do some combinations of ions stay in solution, while others precipitate? Is the solubility of different ion combinations predictable? You will work in teams to collect and share data about the solubility of different ion combinations and to interpret results.

Table 1 and the information below indicate the different cations to be assigned and investigated by teams. Different team data will be compiled and the results used to explore possible links between structure and solubility.

Table 1 Team-Assigned CationsI Na+ Ba2+ Mg2+ Co2+ Ni2+ Cu2+ Al3+

II K+ Ca2+ Sr2+ Cr3+ Fe3+ Zn2+ Ag+

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INFORMATION All teams are to use 0.10 M nitrate salts of the assigned cations in order to ensure that solubility

differences are the effect of the different cations.

All teams are to use 0.10 M sodium salts of the same anions (NO3, Cl, CrO4

2, I, C2O42, CO3

2, SO4

2) in order to ensure that solubility differences are the effect of the different anions.

Solubility and Precipitation

For a salt to dissolve, ionic bonds must be broken and reformed, involving changes in energy. In the solid, the ions are fixed in a rigid lattice, while in solution the ions are mixed with water molecules and free to move about in solution. A waterion attraction cloaks each ion on the surface of the solid with water molecules, and the ions are pulled into the water phase. At the same time the orientation of the water molecules about the ions (Figure 1) due to water ion attraction reduces the water molecules freedom of movement. When a salt precipitates, the process of dissolving is reversed.

Critical Thinking Questions Answer Critical Thinking Questions #4-5 on the Report Sheet at this time BEFORE moving on.

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Figure 1 Orientation of H2O Molecules about Ions


Procedure

1. Table 3, at the end of this lab, provides a full-page grid sheet for conducting the experiment. Alternatively, your instructor may ask you to use cell well plates. If using the Table 3 grid, write the symbols of your team-assigned cations in column 1. Place an acetate sheet over the grid sheet.

2. Add one drop of your first assigned cation solution (the 0.10 M sodium nitrate, NaNO3, or 0.10 M potassium nitrate, KNO3) to each column of the first horizontal row. Add one drop of the appropriate cation solution to each column of the remaining horizontal rows.

Caution: Do not allow your skin to be exposed to the solutions. Silver ion, Ag+, will discolor your skin. Some ions are toxic. Do not discard any of the reagents down the sink. Do discard the waste in an appropriate waste container.

3. Add one drop of the first anion solution (0.10M sodium nitrate, NaNO3) to each row of the first vertical column. Add one drop of the appropriate anion solution to each row of the remaining vertical columns. Take care not to allow the dropper tip to contact the cation solution drop or you will contaminate the anion solution!

4. Record your team data in Table 2 on the next page. Record a (P) if a precipitate formed and indicate the color: (W = white, Y = yellow, O = orange, R = red, B = blue, G = green, Gr = gray, Bl = black, Br = brown). Record an (S) for soluble if there was no precipitate. When you have recorded your data, compare your results with any other team testing the same group of compounds. If necessary, repeat your tests. Do not discard your experiment grid sheet! Save it to refer to while conducting Part III. After completing Part III, rinse the acetate sheet or cell well plates, whichever you used, with water into a waste container as indicated by your instructor.

5. Share and collect the results of the different teams. Compile the different team data into a class data spreadsheet on the computer at the front of the lab room.

Critical Thinking Questions Answer Critical Thinking Question #6 on the Report Sheet at this time BEFORE moving on.

Data Analysis and Implications (Part II) Answer Data Analysis and Implications Questions #3-8 on the Report Sheet as part of your Final

Lab Report. Continue on to Part III of this experiment for now.

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Team DataTable 2 Precipitation of Cations

Anions

Cations NO3– Cl– CrO4

2– I– C2O42– CO3

2– SO42–

Abbreviations used in table:(S) = soluble; (P) = precipitateColor: W= white, Y = yellow, O = orange, R = red, B = blue, G = green, Gr = gray, Bl = black, Br = Brown

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Part III: What is the Ion’s Identity? Your task is to use the results and compiled solubility data of the different teams from Part II to design an experiment to identify ions in solution. You will be given one unknown sample with a cation labeled with a number and the letter “C” and a second unknown sample with an anion labeled with a number and the letter “A”.

Part IIIA What Is the Identity of the Metal Ion in the Well Water?You will be given a unique sample of “well” water. The sample contains one of the following ions: Ba2+, K+, Al3+, Cu2+, Ca2+, or Sr2+. You are to determine experimentally which cation is present. Record your unknown sample ID# below and on page 6 of the report sheet.

Record your procedure and results from your tests in the space below.

Part IIIB What Is the Identity of the Anion in the Solution? Your challenge is to design a reaction procedure to determine which of four anions (CrO4

2, C2O42,

CO32, SO4

2) is present in your aqueous solution. You are to determine experimentally which anion is present. Record your unknown sample ID# below and on page 6 of the report sheet.

Record your procedure and results from your tests in the space below.

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Unknown ID#

Unknown ID#

CI4 What Factors Affect the Solubility of Ions? Report Sheet

Data Analysis and Implications (Part III) Answer Data Analysis and Implications Question #9 on the Report Sheet as part of your Final Lab

Report.

Questions: Extensions and Applications Answer Extensions and Applications Questions #1-4 on the Report Sheet as part of your Final Lab

Report.

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Table 3

Place this grid UNDERNEATH an acetate sheet to help more easily visualize possible precipitates

Anions

Cations NO3– Cl– CrO4

2– I– C2O42– CO3

2– SO42–


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Partner’s Name__________________

Report SheetCritical Thinking QuestionsPart I: What is the Precipitate Identity?

CTQ 1. Based on the tests, and ruling out ion combinations that are soluble (that is, stay in solution), what is the likely identity of the precipitate in the reaction of calcium chloride and sodium oxalate?

CTQ 2. Check the properties of your proposed precipitate at the CRC Handbook of Chemistry and Physics website http://hbcpnetbase.com/. a. Select Section 4: Properties of the Elements and Inorganic Compounds on the left menu bar. b. Select Physical Constants of Inorganic Compoundsc. search for the name of the precipitate you identified above or use the filter option for an

interactive table.

How do the listed properties in the CRC match the observable properties of the precipitate?

Congratulations! The precipitate you have identified forms from solutions in the bodies of people who suffer from kidney stones.

CTQ 3. Check the properties of your proposed precipitate in the Physical Constants of Inorganic Compounds section of the CRC Handbook of Chemistry and Physics as you did for CTQ2.

How do the listed properties in the CRC match the observable properties of the precipitate?

Geologists call this precipitate gypsum. CTQ 4. HYPOTHESIS: Before starting the experiment, study Figure 1 and the information provided

about salt solubility and precipitation. Do you expect solubility to be linked to a particular ion

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http://hbcpnetbase.com/


characteristic (e. g., magnitude of charge, size, or something else)? If so, clearly identify which characteristic(s). Record your hypothesis in the box below .

Hypothesis

CTQ 5. If your hypothesis were to be correct, predict which four of your seven assigned cations are least likely to form precipitates. Write these four cations in the box below.

Cations least likely to form precipitates, based

on my hypothesis.

CTQ 6. HYPOTHESIS TEST. How does your team data on Table 3 validate or contradict your hypothesis regarding the four cations least likely to produce precipitates?

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CI6 What Factors Affect the Reactivity of Acids?

Data Analysis and ImplicationsPart I: What is the Precipitate Identity? 1. For the main reactions investigated in Parts IA and IB (below),

Part IAReaction of Calcium Chloride and Sodium Oxalate

Part IBReaction of Calcium Chloride and Sodium Sulfate

In the boxes below, write each indicated equation.

Be sure to include states of matter, (s) or (aq), for all substances in EACH reaction.

A. An overall chemical equation that represents the reaction, includes both spectator and reactant species, and reflects your data (that is, which are precipitates (s) and which remain in solution (aq)).

B. A complete (total) ionic equation.

C. A net ionic equation that shows only the reacting ions producing the precipitate.

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Part 1A:

Part 1B:

Part 1A:

Part 1B:

Part 1A:

Part 1B:


2. Identify any ions present in both reacting systems that are just spectator ions.

Part II: What Factors Determine Solubility?

3. Is it possible to predict whether a precipitate will likely be white or a color other than white based on the position of the cations element in the periodic table? Refer to a periodic table and the compiled data of the different teams with regard to the precipitate color of the cations that reacted to help draw your conclusion.

Circle one: Yes or No If it is possible to make a prediction, write a general claim statement that discusses the

predictability of precipitate color based on position of the cations in the periodic table.

4. What generalizations, if any, can be made about the solubility of ionic compounds with singly-charged alkali metal cations (Li+, Na+, and K+)? Refer to the compiled solubility data to help draw your conclusions. List specific examples to support your conclusion.

5. What generalizations, if any, can be made about the solubility of ionic compounds with singly-charged anions (Cl and I and NO3

)? Refer to the compiled solubility data to help draw your conclusions. List specific examples to support your conclusion.

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Part 1A:

Part 1B:


6. What generalizations, if any, can be made about the solubility of ionic compounds with multiply-charged anions? To help, compare the solubility data of multiply-charged anions (CrO4

2, C2O42,

CO32, SO4

2 ) with the singly-charged ions investigated in Question 5. List specific examples to support your conclusion.

7. What generalizations, if any, can be made about the solubility of an ionic compound, if both the cation and anion are multiply-charged? Refer to the compiled solubility data to help draw your conclusions. List specific examples to support your conclusion.

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8. Figure 2 gives the ionic radii of some common metal ions.

Figure 2 Ionic Radii (picometers) of Common Metal Ions

What generalizations, if any, can be made about solubility and metal ion size (radius)? To help with this analysis, compare the number of precipitates produced with the different anion combinations for the tested alkaline earth cations (Mg2+, Ca2+, Sr2+, Ba2+) of Group 2 and the cations (Zn2+ and Cd2+) of Group 12. Organize and refer to the data to answer this question.

Part III: What is the Ion’s Identity? 9. Based on your experimental results, identify the cation in your first unknown sample and the anion

in your second unknown sample and write the identity of each ion in the table below.

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Unknown ID # Ion Identity

Cation

Anion


Questions: Extensions and Applications 1. Write net ionic equations based on the compiled data (Part II) for the combination of the cation

reagent, Ba(NO3)2, with the different anion reagents. Net ionic equations are ONLY written for reactions in which a PPT forms. You should have 4 net ionic equations for Ba2+.

2.

Circle the insoluble solid compound within each pair below that is a color other than white, and indicate the reasoning for your decision.

A. Nickel carbonate or lead carbonate

Reason:

B. Cobalt oxalate or strontium oxalate

Reason:

3. Check the ‘solubility rules’ in a chemistry textbook (use the index!). How do the rules compare to the generalizations about solubility made from the compiled team solubility data? Be explicit about the comparison between the two.

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4. An unknown solution is either 0.10 M LiNO3 or 0.10 M Ba(NO3)2 or 0.10 M Cu(NO3)2. When you add 1 mL of 0.10 M K2CO3, a white precipitate forms.

A. Think carefully about the 3 different metal cations, their location on the periodic table, and your data. Circle the compound below that correlates to the identity of the solution.

LiNO3 or Ba(NO3)2 or Cu(NO3)2

B. For the formation of the white precipitate formed from mixing K2CO3(aq) and the solution you chose in Question 4.A, write the following equations:

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Overall Chemical equation

Total Ionic

equation

Net Ionic Equation


ChemInquiry 6 Name________________________

What Factors Affect the Reactivity of Acids?


This part of the pre-lab assignment is due at the beginning

of the lab period, and must be done individually before you

come to lab!



II. Safety Hazards/Precautions4. Complete the following table. Use the PCC MSDS online link on

your lab web page. Be sure to select the location Sylvania ST and be sure the chemical name matches precisely.



Hazard Statements


HCl

hydrochloric acid


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Pre-lab Score: ____________/5


NaC2H3O2

sodium acetate



etc.)






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ChemInquiry 6 Name_______________________


How can differences in acid reactivity be determined? How is a pH probe used? How does pH change with the concentration of acids? How does the pH of an acidic solution change with the addition of a salt?


IntroductionMany common household solutions contain acids and bases. Acid-base indicators, such as litmus (a lichen) and red cabbage juice, turn different colors in acidic and basic solutions. They can, therefore, be used to show if a solution is acidic or basic. The measurement of the acid-base content of a substance can be accomplished in a number of ways: litmus paper, pH paper, indicators, or with a pH meter. An acid is classified in its simplest definition as a substance that reacts with water to produce H3O+ ions:

HA + H2O H3O+ + A

An important measurement of the acidity or basicity of a solution is to measure its pH. As will be shown below, a pH measurement is really a measurement of the concentration of H3O+ ions in solution. The pH of a solution is related to amount of hydrogen ions in the solution by the following:

pH = log[H3O+]or

[H3O+] = 10pH

This second equation shows that pH is simply the negative of the exponent of 10 that will give the equilibrium hydrogen ion concentration. For example, if [H3O+] = 1.0 × 106 M (or more simply, 106

M), then pH = 6 (the negative of the exponent of the 10). Each unit change in pH is a change of a factor of 10 in the concentration of [H3O+] ions.

NOTE: acetic acid, HC2H3O2 or CH3COOH, is designated as “HOAc” acetate ion, C2H3O2

− or CH3COO−, is designated as “OAc−”

The acids and bases used in this experiment, while relatively dilute, will cause eye and skin damage. Wear goggles at all times in today’s lab. In addition, ammonia solution is toxic. Its liquid and vapor are extremely irritating, especially to eyes. Handle these solutions with care. Do not allow the solutions to contact your skin, eyes or clothing.

You may dispose of all solutions in today’s lab by flushing them down the drain with plenty of water. Please remove the solid metals so they don’t go down the drain. Rinse solid metals with water and dispose of solids in the appropriate waste collection container provided.

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CI6 What Factors Affect the Reactivity of Acids?t

Experiment 1:Data Collection: Acid/Base Properties

Specific Tests of ReactivityA. Problem 1: What is the reaction between acids and some metals ?

Obtain 3 eyedropper squirts of 1.0 M HCl into 5 separate LARGE test tubes. IMPORTANT: Be consistent in the size of the “squirts” .

Put 2 small pieces of the following metals into the acid-filled test tubes (use a separate test tube for each metal): Mg, Ca, Cr, Cu, Ag

Loosely cover each tube with a cork so that it is NOT airtight. After waiting several minutes, remove the cork and quickly hold a lighted wooden splint to the top of the test tube. Record your observations below for each reaction, paying special attention to the vigor of each metal-acid reaction.

If no reaction occurs, let sit for 10-15 minutes with the cork on top and repeat the wooden splint test.

RECORD your observations in the boxes below.

Mg Ca Cr Cu Ag

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Critical Thinking QuestionsAt the end of this lab is a Report Sheet with CTQs that you are expected to answer DURING LAB as indicated in the procedure.

This is in addition to the making of observations and recording of data on the lab itself.

Please answer CTQ’s 1-3 on the Report Sheet at this time BEFORE moving on.\

B. Problem 2: What is the reaction between acids and carbonates ? Obtain 3 eyedropper squirts of 1.0 M HCl in one LARGE test tube, Put a small amount of CaCO3 powder into the test tube and loosely cover the tube with a cork so

that it is NOT airtight. After waiting several minutes remove the cork and quickly thrust a lighted wooden splint deep

INTO the test tube. RECORD your observations below.

1.0 M HCl with CaCO3 powder


Please answer CTQ 4 on the Report Sheet at this time BEFORE moving on.

C. Problem 3: Does acid concentration affect reactivity? Obtain 3-eyedropper squirts of 0.10 M HCl in 2 separate LARGE test tubes and 3-eyedropper

squirts of 1.0 M HCl in 2 other separate LARGE test tubes.

Test C1: Using the same procedure outlined above : Put a small piece of your most reactive metal into a 0.10 M HCl test tube and another piece

into a 1.0 M HCl test tube. Test with a lighted wooden splint as before. RECORD your observations below.

0.10 M HCl with most reactive metal

1.0 M HCl with most reactive metal

Test C2: Using the same procedure outlined above :

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Put a small amount of CaCO3 powder into a 0.10 M HCl test tube and a 1.0 M HCl test tube. Test with a lighted wooden splint as before. RECORD your observations below.



Please answer CTQ 5 on the Report Sheet at this time BEFORE moving on.

D. Problem 4: Does acid identity affect reactivity? Obtain 3-eyedropper squirts of 0.10 M HCl in 2 separate LARGE test tubes and 3-eyedropper

squirts of 0.10 M HOAc in 2 other separate LARGE test tubes.

Test D1: Using the same procedure outlined above : Assemble the 0.10 M HCl test tube and the 0.10 M HOAc test tube side-by-side in a test tube

rack. Put a small piece of Mg into each tube at the same time. Observe the vigor of the reactions Test with a lighted wooden splint as before. RECORD your observations below.

0.10 M HCl with Mg

0.10 M HOAc with Mg

Test D2: Using the same procedure outlined above : Assemble the 0.10 M HCl test tube and the 0.10 M HOAc test tube side-by-side in a test tube

rack. Put a small amount of CaCO3 powder into each tube at the same time. Observe the vigor of the reactions Test with a lighted wooden splint as before. RECORD your observations below.


0.10 M HOAc with CaCO3 powder

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Please answer CTQ’s 6-7 on the Report Sheet at this time BEFORE moving on.

Rinse the solid metal pieces with water and place into the waste container to be saved. Dispose of all the liquid contents from Experiment 1 down the drain with plenty of water.

Experiment 2:Data Collection: pH and Concentration

H. Obtain 3 clean and dry large test tubes from your drawer. Label them as follows:0.100 M HOAc*0.0100 M HOAc0.00100 M HOAc

*Record the “Actual ‘0.100 M’ HOAc concentration (M) for term” in the table on the next page as provided on the solution bottle or by your instructor.

I. Obtain two clean disposable droppers.

J. Obtain about 5 mL of 0.100 M acetic acid into a small beaker. Fill one labeled test tube with about 1-2 inches of 0.100 M acetic acid (HOAc).

K. You will need to prepare 0.0100 M and 0.00100 M solutions of acetic acid in the other beakers. You will do this by a process called serial dilution. The 0.0100 M solution is made by obtaining 1.0 mL of 0.100 M acetic acid in a 10.0-mL graduated cylinder and following the directions below:

It is important that you carefully adjust the amount with a dropper so that the meniscus reads exactly 1.00 mL. Then add 9.00 mL of distilled water to yield a total volume of exactly 10.00 mL in the graduated cylinder.

Essentially, you have used your old friend, the dilution equation:

CconcVconc = CdiluteVdilute (where C = concentration, V = volume)

You have simply made a 1:10 dilution of the concentrated solution to obtain a solution that is 10 times more dilute.

L. Fill a second labeled test tube with about 1-2 inches of the 0.0100 M solution you made. Save 1.0 mL of the solution to continue to the next dilution.

M. The 0.00100 M solution is made by following the same procedure – making a 1:10 dilution of the 0.0100 M solution.

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N. Fill a third labeled test tube with about 1-2 inches of the third solution (0.00100 M) you made.

O. Set up the LabPro pH sensor for data collection by connecting the pH Sensor to the LabPro by inserting the cable into Channel 1. Be sure your LabPro has 3 connections: Power; Probe; Computer connection (USB cable).

P. Start the LoggerPro program. Open the file “CI3 pH” by going to the “Open” menu in LoggerPro CH 223 CI3 pH (Or as directed by your instructor).

Q. Raise the pH Sensor from the sensor soaking solution and set the solution aside. Use a wash bottle filled with distilled water to thoroughly rinse the pH Sensor. Catch the rinse water in a 250-mL beaker. Dry the probe by GENTLY blotting it with a Kim-wipe (NOT a paper towel).

R. Calibrate the pH meter using 2 buffer solutions as follows:9. Obtain some pH-4 and pH-7 (also known as the “pH soaking solution”) buffer solutions

that are stored in the scintillation vials.First Calibration Point10. Choose Calibrate ► CH1: pH from the Experiment pull-down menu and then click

Calibrate Now.11. For the first calibration point, rinse the pH sensor with distilled water, then place it into a

buffer of pH 4.0.12. Type “4” in the edit box as the pH value.13. Swirl the sensor, wait until the voltage for Input 1 stabilizes (wait 15 seconds), then click

“Keep.”Second Calibration Point14. Rinse the pH sensor with distilled water, and place it into a buffer of pH 7.0.15. Type “7” in the edit box as the pH value for the second calibration point.16. Swirl the sensor and wait until the voltage for Input 1 stabilizes. Click “Keep”, then click

“Done.” This completes the calibration.

S. Measure the pH of the first solution. Measure the pH by immersing the probe into the test tube for each solution. Be sure the solution completely covers the glass bulb at the bottom. If not, add more of the respective solution. Gently swirl the probe in the solution and when the pH reading displayed on the screen stabilizes, record the pH value in the box below (round to the nearest 0.01 pH unit).

T. Prepare the pH sensor for reuse:4. Rinse it with distilled water from a wash bottle.5. Place the sensor into the sensor soaking solution and swirl the solution about the sensor

briefly.6. Rinse with distilled water again.

U. Determine the pH of the other solutions and record them in the boxes below. You must clean the pH Sensor between tests, using the procedure above each time. Be sure to gently blot dry the probe with a Kim-wipe each time.

V. If your pH values do not come out near (within a few tenths of) the following, you have a bad pH probe or a calibration error. Please see your instructor.

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Take Note!

!


0.1 M; pH ~ 2.9 0.01 M; pH ~ 3.4 0.001 M; pH ~ 3.9

W. When you are done, rinse the sensor with distilled water and return it to the sensor soaking solution.

*Actual “0.100 M” HOAc Concentration (M) for term

HOAc Solution pH

0.100 M

0.0100 M

0.00100 M



Experiment 3:Data Collection: Effects of Adding a Salt

A. Obtain a clean, dry, small 20-mL beaker, and a 10.00 mL graduated cylinder. Put precisely 10.00 mL of 0.100 M acetic acid into the beaker. Measure the pH with the LabPro pH sensor and record the pH value in the box below.

B. Obtain two samples of sodium acetate, one about 0.1 g and the other about 0.3 g. Record the actual mass exactly to three decimal places in the boxes below.

C. Add the 0.1-g sample of sodium acetate to the beaker of acetic acid. Stir well with a stirring rod until dissolved. Measure the pH of the solution using the probe, and record the pH value in the box below.

D. Add the 0.3-g sample of sodium acetate to the same beaker. Stir well with a stirring rod until dissolved. Measure the pH of the solution using the probe, and record the pH value in the box below. Note the TOTAL mass of sodium acetate added to the solution.

Actual mass of sodium acetate sample (g)

HOAc Solution (M)

Total sodium acetate added (g) pH

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An Important Step!


0.100 M + 0.00 g

0.100 M

0.100 M

Dispose of the contents of all of the test tubes and beakers from Experiments 2 and 3 down the drain with plenty of water.



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ChemInquiry 6 Name__________________-_______

Partner’s Name__________________

Report SheetCritical Thinking Questions

Experiment 1:

CTQ 15. Do all the tested metals react with HCl? Yes or No Explain how you know:

CTQ 16. Do all the tested metals react as vigorously with 1.0 M HCl? Yes or No

If not, prepare a list of metal reactivity from most vigorous to least vigorous.

Metal

CTQ 17. Write a chemical equation which represents each reaction from Exp 1.A of an acid with a metal below. If no reaction takes place, write “No reaction”

Hint: These are single replacement reactions. That is, the metal atom becomes an ion, and replaces the H in H-Cl:

M(s) + 2 HCl(aq) (a gas) + MCl2 (aq)What was the gas produced? Three common gases produced in the general chemistry laboratory are:

O2: Supports combustion (does not itself ignite) H2: Highly flammable/explosive gas CO2: Extinguishes combustion/flame

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Most Vigorous

Least Vigorous


CTQ 18. Write a chemical equation which represents the reaction of Exp 1.B in the box below.

Hint: This is a double replacement reaction. CaCO3(aq) + HCl(aq) ___ + ___

Hint: Whenever carbonic acid, H2CO3 is formed in water, it immediately decomposes into a gas and water: H2CO3(aq) gas + H2O(l). What is the gas? _____

Chemical Reaction (be sure to include phases)

CTQ 19. Describe the differences in reactivity between 0.10 M and 1.0 M HCl by using the metal and the carbonate reactions as data. Hypothesize the nature of any differences.

CTQ 20. Does the reactivity of an acid depend on its identity?

Yes No Unsure

CTQ 21. Explain how you determined your answer to CTQ 6.

Critical Thinking Questions tied to Experiment 2 are continued on the next page…

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Chemical Reaction (be sure to include phases)


CTQ 8. This table will be filled out based on the instructions below.

*Be sure to use the “Actual ‘0.100 M’ HOAc concentration (M) for term” for your solution, the dilutions, and all calculations.

a. Record your pH values from Experiment 2 on your data sheet into the third column above.

b. From the pH, calculate the concentration of hydronium ion for each dilution. Enter your values in the fourth column in the table. Recall [H3O+] = 10pH

Show one sample calculation for the 0.100 M solution:

c. Next, determine the concentration of the acetate ion, [OAc−], for each solution (Hint: See the example shown below). By looking at the formula for acetic acid and the dissociation equation, how would the concentration of hydronium ion be related to that of the acetate ion?

d. Finally, determine the equilibrium concentration of acetic acid, [HOAc], that is in the solution (IMPORTANT NOTE: This will NOT be the original concentration you calculated for your dilutions in the first column of the table, because some HOAc will have dissociated into H 3O+ and acetate ion). See the example shown below. Show a set of sample calculations below.

Example calculation: Recall [H3O+] = 10pH. You may set up an ICE table to solve this problem, as demonstrated below. HOAc + H2O H3O+ + OAc

I 0.100 M 0 0C x +x +xE 0.100 x x x (determined from the pH)

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Initial concentration

acetic acidsolution

(HOAc)o

Equilibrium concentration

acetic acid[HOAc]

(M)

“(HOAc)o x”

EquilibriumpH

(as measured)

Hydronium ion[H3O+]

(M)(determined from

pH)

“x = 10pH”

Acetate ion[OAc]

(M)

“x”

0.100 M*

0.0100 M

0.00100 M


Show one sample calculation for the 0.100 M solution:

CTQ 9. Based on your results to CTQ 8, does the ionization reaction of acetic acid go to completion, or is it a partial reaction that represents an equilibrium?

Completion Partial Reaction Unsure

Look at the pH you got and the pH you’d get if it were ALL dissociated. How do they compare?

Explain your answer to the type of reaction you chose above.

CTQ 10. What general pattern do you observe in the concentration of acetic acid vs. pH?

Critical Thinking Questions tied to Experiment 3 are continued on the next page…

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CTQ 11. This table will be filled out based on the instructions below:

EquilibriumpH

(measured)

[H3O+](M)

(determined from pH)

“x” = 10pH

Total moles of acetate ion

added

Initial concentration

acetate ion added(OAc)o

(M)

“Y”

[OAc](M)

Computed using an ICE table

Y + x

[HOAc](M)

Computed using an ICE table

0.100 x

0.00 0.00

a. Record your pH values from Experiment 3 on your data sheet into the first column above.

b. From the pH, calculate the concentration of hydronium ion for each dilution. Enter your values in the second column in the table. Recall [H3O+] = 10pH

CTQ 12. Calculate the total number of moles of acetate ion added from the solid sodium acetate, and determine the initial acetate ion concentration. Enter these values in the third and fourth columns of the table above. SHOW YOUR CALCULATIONS below!

Note : Compute the total # of moles of acetate ion based on the ~0.1 g added (2nd row in table) and also the ~0.1 + 0.3 g added (3rd row in table).

NOTE: Be sure to use the total mass of sodium acetate to determine moles of sodium acetate added. Then moles of acetate ion are determined by observing that moles acetate ion = moles sodium acetate added.

Show ALL calculations:

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CTQ 13. Assume that the acetate ion present at equilibrium comes from both the acetate you added to the solution AND that from the dissociation of acetic acid. Calculate the equilibrium concentration of acetate ion [OAc] from all sources in the solution. You must calculate this using an ICE table, as shown in the example calculation below. Enter these values in the fifth column of the table. Show the work for the ~0.1 g added trial below.

CTQ 14. Calculate the equilibrium concentration of acetic acid, [HOAc], in the solution after mixing in the sodium acetate. You must calculate this using an ICE table, as shown in the example calculation below. Enter these values in the sixth column of the table. Show the work for the ~0.1 g added trial below.

Example calculation: Recall [H3O+] = 10pH. You MUST set up an ICE table to solve this problem, as demonstrated below.

HOAc + H2O H3O+ + OAc

I 0.100 M 0 YC x +x +xE 0.100 x x Y+x (determined from the pH)

Show one sample ICE table for the ~0.1 g added solution (tied to CTQ’s #13-14):

CTQ 15. What general trends do you see in the pH of the solutions compared to the amount of sodium acetate added? As more sodium acetate is added, the pH tends to:

Increase Decrease No effect seen

CTQ 16. Explain the trends in CTQ 15 in terms Le Châtelier’s principle.

CTQ 17. In acid/base terms, what is the relationship between acetic acid and acetate ion?

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concentration determined from the initial mass of sodium acetate added and the volume of the solution (10.00

mL).


CTQ 18. For each of the three trials of varying amounts of acetate ion added, calculate the experimental equilibrium constant, Ka = ¿¿.

Show all calculations and place the final answer for each trial in the table below:

What do you notice about these values? Are they relatively constant?

Hypothesize why this might be so or not so.

Data Analysis questions are found on the following pages…

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Trial # Ka

1

2

3


Data Analysis

Answer these questions using grammatically correct English sentences. You may type up the Claim and Evidence questions.

1. Claim 1: How does acid concentration affect reactivity (= vigor of the reaction)?

2. Evidence: Use your data to provide evidence that explains the difference in the rate of the reaction between the metals or CaCO3 with 1.0 M HCl and with 0.10 M HCl.

3. Mental Model: Draw a PARTICLE VIEW of a solution of 1.0 M HCl and 0.10 M HCl that would help explain the difference in the rate of reaction between these two acid concentrations (you may omit water for clarity). Use the following symbols, as needed, in your drawings:

= HCl

= H+

= Cl−

4. Claim 2: Do different types of acid react similarly to each other (i.e. why can we classify different compounds under the category “acid”)? (Be sure to compare acids of the same concentration.)

5. Evidence: Use your data to provide evidence that supports your claim as to why we can classify different compounds as acids.

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1.0 M HCl 0.10 M HCl


6. Claim 3: What are the differences in reactivity between 0.10 M HCl and 0.10 M HOAc? (If you used concentrations other than 0.10 M, use these in your claim.)

7. Evidence: Use your data to provide evidence that explains the difference in the rate of the reaction between CaCO3 with 0.10 M HCl and with 0.10 M HOAc. Use the results of your tests in Exp. 1.D “Problem 4” to help answer this question.

8. Mental Model: Draw a PARTICLE VIEW of a solution of 0.10 M HCl and 0.10 M HOAc that shows in general how they are different in solution (you may omit water for clarity). Use the following symbols, as needed, in your drawings:

= HCl

= H+

= Cl−

= HOAc

= OAc−

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0.10 M HCl 0.10 M HOAc


9. The accepted value of Ka(acetic acid) = 1.8 105. How do your answers to CTQ 18 compare to this value? Propose reasons why they are or are not similar.

10. What conclusion can be drawn about the effect on the pH of adding a salt of the solution (NaOAc in this experiment)? (In other words, adding a “common ion,” OAc−)

11. Predict what should happen to the pH of an acetic acid solution if a salt that did not contain acetate ions were added (such as NaCl). Explain your answer

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CI7 Acid/Base Interactions

ChemInquiry 7 Name__________________________

Acid/Base Interactions

Titration Curves


This part of the pre-lab assignment is due at the beginning of

the lab period, and must be done individually

before you come to lab!



II. Safety Hazards/Precautions7. Complete the following table. Use the PCC MSDSonline link on




Hazard Statements


NaOH

sodium hydroxide


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Pre-lab Score: ____________/5


CH3COOH

Acetic acid



etc.)






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How do we generate a titration curve for the titration of a weak acid with a strong base? How can we determine the concentration of a weak acid by titration? How can we determine the pKa and identity of an unknown weak acid by titration?


Titrations and pKa

Background Information on titration curves In a titration, a measured amount of one solution, called the titrant, is added to a known quantity of a second solution, called the analyte, whose concentration and/or identity is unknown, until a reaction of known proportion is completed.

Titrations are stoichiometry problems, and in this class we don’t worry about cases where complex equilibrium problems must be solved. We can ignore equilibrium if either the acid or base, or both, are considered strong. If so, this ensures the titration reaction will go to completion and equilibrium analyses are not necessary.

The reaction you will observe today in Part I is between a strong base (NaOH) and a weak acid (HOAc):

HOAc(aq) + NaOH(aq) H2O(l) + NaOAc(aq) acid + base water + salt

net ionic equation: HOAc(aq) + OH−(aq) H2O(l) + OAc−(aq)

The goal is to discover the concentration of HOAc in Part I. In Part II, you will extend the technique to discover the identity of a solid acid powder by performing a titration with NaOH. In either case, you will titrate an unknown concentration of a weak acid analyte (acetic acid, HOAc), against a titrant of strong base of known concentration (sodium hydroxide, NaOH). The titrant is added slowly to the analyte. After the addition of each drop, the above reaction takes place according to the stoichiometry of the reaction (1:1 in our case). After each addition of NaOH, an equal number of HOAc molecules will be transformed into water and the salt, NaOAc. When the NaOH has completely consumed all of the acid in the unknown HOAc solution, you will be able to determine the concentration of the acid.

Traditionally, titrations are performed such that the analyte is added until the number of moles of the titrant added is precisely equal to the number of moles of the analyte in question. At this point, the

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number of moles of the analyte have been discovered (= moles titrant if the reaction is 1:1). Finally, by dividing the newly discovered moles of analyte by the original volume of the analyte, the concentration of the analyte can be determined. The “magic” is determining when the moles titrant = moles analyte and to stop adding the analyte and perform the calculations. This can be accomplished by adding a small amount of an indicator that changes color at the precise pH that results when the titrant has exactly neutralized the analyte. It requires some skill and intuition to pick the correct indicator, but this can be a very precise quantitative procedure.

We can also map out the progress of the titration by plotting the pH of the analyte solution after each drop or volume of titrant is added. An analysis of the curve will (as you will see) very clearly indicate the point where the moles titrant = moles analyte. In an acid-base titration, the curve that is generated when the pH of the analyte solution is plotted versus volume of titrant added is called an acid-base titration curve (see diagram at right).

The stoichiometric point when the moles of base analyte have exactly reacted with the unknown moles of acid is called the equivalence point of the titration. The precise technique to generate a titration curve is to use a pH-probe and to generate the curve as noted above. As mentioned, an analysis of this curve will provide the precise volume of the titrant of known concentration required to react with all the analyte. This then enables a determination of the concentration of the unknown analyte, as described below. In addition, a pH titration curve also enables us to discover something else that a simple titration does not allow: the determination of the pKa of the unknown acid, thus determining or confirming its identity.

There are 3 easily calculable and significant “sign-posts” during an acid/base titration.

A. Before titration starts (Initial pH): The first is the pH at the beginning of the titration before any titrant is added. In this case, the only reacting species in the analyte solution is the analyte itself. Whether the analyte is an acid or a base, what differentiates the difficulty of the calculation is whether the analyte is strong or weak.

1. Case 1: Strong acid or base analyte. In this case, the calculation of pH is straightforward. If we know the concentration we do the following:

[H3O+] = (HA)o

Strong acid: pH = −log[H3O+] Strong base: pOH = −log[OH−]; pH = 14.00 – pOH

2. Case 2: Weak acid or base analyte. In this case, we can justifiably use the weak acid/base approximation equations provided we know the Ka or Kb and initial concentration of the acid or base, respectively:

Weak acid: [H3O+] = √Ka ∙(HA)o

Weak base: [OH−] = √K b ∙(B )o

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For a weak acid analyte, we can consider ~100% of the acid to be in it’s acid (protonated) form, HA.

%-Completion of Titration0%

HA

A

Mol

es

Volume of titrant added (mL)

pH o

f ana

lyte

0

7

14

0

Initial pH


pH o

f ana

lyte

0

7

14

0


B. ½-way volume (pKa): This point is critical in that it can tell us the identity of an unknown acid. To see why, we consider two concepts: stoichiometry and equilibrium.

1. Stoichiometry first. Consider an acid analyte. Half-way to the equivalence point, exactly ½ of the acid is converted into its base form:

HA(aq) + OH−(aq) A−(aq) + H2O(l)

If exactly half of HA is now in the form A−, then the moles HA = moles A−. Since they are in the same solution, then [HA] = [A−].

2. Equilibrium second. Once we know the pH, and the concentrations of HA and A− at the ½-way point, we can use the equilibrium constant expression to generate a formula for the pKa of the acid as follows:

Ka = [ H3O+¿][ A- ]

[HA ]¿

But if [HA] = [A−], they will cancel each other out in the Ka expression, and we are left with the following useful result:

At ½-way point: Ka = [H3O+]

Taking the –log of both sides

−log Ka = −log [H3O+]

The –log of anything is the ‘p’ of that thing, so

pKa = pH½-way

NOTE: pH is not the same “thing” as pKa. pH is a measure of [H3O+] at any moment in a solution pKa is the equilibrium constant, Ka. It does not change and is unique to a particular

acid. At the ½-way point, these two different quantities are merely numerically equal to

each other. However, they are not the same “thing”. It’s important to remember this.

3. By measuring the pH at the ½-way point in a titration, we can discover the pKa of the acid and hence, the identity of the acid. Of course, you have to complete the titration first in order to determine where the ½-way point is!

4. A similar line of reasoning works for weak bases as well.

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HA

A

Mol

es

50%

HA A


pH o

f ana

lyte

0

7

14

0

½-way point pH

½-way point volume

Continued…


C. Equivalence point: At this point, ALL the analyte is converted into its conjugate form. The pH of the solution is determined by the strength of the conjugate species in solution and the total volume of the solution.

1. Strong acid analyte:


In this case, A− is so weak as a base it has no effect on the pH of the solution. For strong acids, the pH at the eq. pt. should = 7.

2. Weak acid analyte (today’s experiment):


In this case, A− is a marginally weak base, and as a result, the pH of the solution at the eq. pt. will be basic. (We won’t worry about how to calculate it at this point, just realize it should be basic.)

3. Similar computations can be made for strong and weak base analytes with acid titrants

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pH o

f ana

lyte

0

7

14

0

equivalencepoint pH

equivalence point volume


HA

A

Mol

es

50%

HA A

A

HA

100%


Concentration of Unknown Analyte

The equivalence point for a 1:1 stoichiometric titration is defined as the point at which moles titrant added = moles analyte original:

mole analyte = mole titrant

For this lab in particular, the analyte is an acid and the titrant is a base:

mole acid = mole base

MacidVacid = MbaseVbase

Macid = Mbase ∙V base

Vacid

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Volume base added to equivalence point

Recall:

M = nV

So n = MV

ORIGINAL volume of acid analyte


Part 1: Titration of Acetic Acid SampleData Collection

Setting up the Titration Apparatus (Work together as a class with the instructor demonstration)

1. Obtain 100 mL of ~0.250 M NaOH standard solution from the stock carboy. Clearly label your beaker for future reference.

2. Obtain about 40 mL of the acetic acid analyte solution from the stock carboy. Clearly label your beaker for future reference.

3. Obtain a LabPro from the lab cart. Also obtain a box containing a drop counter apparatus, pH probe and solution reservoir/drop dispenser. An example of the experimental setup is displayed on your instructor’s desk. Follow the specific directions below.

4. Connect the pH Sensor to CH 1 of the computer interface. Lower the Drop Counter onto a ring stand and connect its cable to DIG/SONIC 1.

5. Obtain the plastic reagent reservoir. Note: The bottom valve will be used to open or close the reservoir, while the top valve with the black knurled knob will be used to finely adjust the flow rate. For now, close both valves by turning the handles all the way clockwise.

Rinse it with a few mL of the ~0.250 M NaOH solution. Record the “Actual NaOH concentration for term” on first page of the report sheet. Use the special reservoir clamp to attach the reagent reservoir to the ring stand, as shown by the demonstration setup. Add the remainder of the NaOH solution to the reagent reservoir.

Drain a small amount of NaOH solution into an empty 250 mL waste beaker so it fills the reservoir’s tip. To do this, first open the bottom valve by turning the handle vertical. Then, slowly open the upper valve by turning it counterclockwise. When a few milliliters have drained out, close both valves again.

6. Open the LoggerPro program. Prepare the computer for data collection by opening the file “Acid-Base Titration Template” from the “Open” “CH 223” folder (or as directed by your instructor).

1) You will probably find a pop-up window asking you to connect an interface. Simply hit “Connect” in the center of the window.

7. Calibrate the pH meter using 2 pH-buffered solutions from your bench kit as follows:

1) Obtain some pH-4 (PINK) and pH-7 (YELLOW) buffer solutions that are stored in the scintillation vials. These solutions are in your bench kit.

First Calibration Point

2) Choose Calibrate ► CH1:pH from the Experiment menu and then click Calibrate Now.

3) For the first calibration point, rinse the pH sensor with distilled water, blot dry with a Kimwipe, then place it into the YELLOW buffer solution of pH 7.0.

4) Type “7” in the edit box as the pH value.

5) Swirl the sensor, wait until the voltage for Input 1 stabilizes, then click “Keep.”

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Second Calibration Point

6) Rinse the pH sensor with distilled water, blot dry with a Kimwipe, and place it into the PINK buffer of pH 4.0.

7) Type “4” in the edit box as the pH value for the second calibration point.

8) Swirl the sensor and wait until the voltage for Input 1 stabilizes. Click “Keep”, then click “Done.” This completes the calibration.

9) Place the pH probe back into its soaking solution. Do not put the pH probe in the drop counter at this time.

8. Calibrate the volume of a drop so that a precise volume of titrant is recorded in units of milliliters.

1) Set up the Drop Counter on the ring stand as indicated in the demonstration setup at the front of the room. Do not insert the pH probe or put a beaker under the Counter at this time.

2) Be sure the tip of the reservoir is aligned so that drops will fall through the opening in the Drop Counter without hitting the sides.

3) Place a 10 mL graduated cylinder directly below the notch in the Drop Counter, lining it up with the tip of the reagent reservoir.

4) Open the bottom valve on the reagent reservoir (vertical). Keep the top valve closed (turned all the way clockwise).

5) Go to the pull-down menu: Experiment Calibrate LabPro: 1 Dig1: Drop Counter.6) In the pop-up window, check the “Automatic” button.7) Click the “Start” button.8) Slowly open the top valve of the reagent reservoir so that drops are released at a slow rate (~a

few drops every second). You should see the drops being counted on the computer screen. Be sure the computer screen is counting drops. If not, you should empty the graduated cylinder, realign your apparatus and start over. Ask for help if still not working.

9) When the volume of NaOH solution in the graduated cylinder is between 9 and 10 mL, close the bottom valve of the reagent reservoir.

10) Enter the precise volume of NaOH (read to the nearest 0.1 mL) in the edit box.11) In the box below record the number of drops/mL displayed on the screen for future use.

Number of drops/mL

12) Click . Discard the NaOH solution from the graduated cylinder into your waste beaker.

9. Prepare the apparatus.

1) Use a 25-mL volumetric pipet to dispense a 25.00-mL sample of the acetic acid (HOAc) analyte into a 250-mL beaker. (Be sure to “prime” the pipet first).

2) Record the precise volume of the acid analyte in the box below.

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(correct sig figs?)


Volume of acid analyte

3) Dilute the analyte to approximately 40 mL in the 250-mL beaker with deionized water. You do not have to be precise in this dilution, but you should have a final volume of about 40 mL.

4) Add 3 drops of phenolphthalein indicator to the analyte. In the data table on page 9, record the color of the indicator before the equivalence point.

5) Fill the Drop Reservoir with the ~ 0.250 M NaOH if not already full.6) Place a waste collection beaker under the drop counter/reservoir. Be sure the black upper knob

is turned tightly all the way clockwise.7) Open the bottom blue butterfly valve – the “on/off” valve (vertical).8) Slowly turn the upper black knob counterclockwise until you obtain a drip rate of about 1 drop

every 2 seconds.9) Close the blue “on/off” valve (horizontal).10) Remove the soaking solution from the pH probe. Using distilled water rinse the probe into the

waste beaker. Blot dry with a Kimwipe.11) Place the magnetic stir plate on the base of the ring stand.12) Insert the pH sensor through the large hole in the Drop Counter.13) Add a stir bar to your titration beaker. Center the stir plate and beaker so the stir bar is

approximately in the center of the beaker. See instructor demo setup for reference.14) Adjust the positions of the Drop Counter and reagent reservoir so that both the drops and the pH

probe are as close to the center of the beaker as is possible. Don’t let the stir bar hit the pH probe!

15) Remove the waste beaker and place your HOAc analyte solution under the probe and drop reservoir.

16) Adjust the reagent reservoir so its tip is in an appropriate position to be counted by the drop counter.

17) Turn on the magnetic stirrer so that the stir bar is stirring at an intermediate fast rate, a setting of about 6 (just make sure the solution isn’t splashing!).

18) TAKE NOTE!! Before you begin the titration, make a note of and prepare for the data you will collect. The required data collection is listed in the table on the next page.

19) Click . No data will be collected until the first drop goes through the Drop Counter slot. Fully open the bottom valve—the top valve should still be adjusted so drops are released at a rate of about 1 drop every 1-2 seconds. Drop rates faster than this will produce poor data. When the first drop passes through the Drop Counter slot, check the data table to see that the first data pair was recorded.

20) Open the bottom blue valve on the drop reservoir. Data will begin to be collected when the first drop falls. Be sure the red light flashes each time on the drop counter as the drops fall, and that pH and volume are being recorded in LoggerPro.

21) Titrate the acid analyte until a few mL past the equivalence point.

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22) Record the following in the data table (if not already done).

23) Print a copy of your titration curve and keep with your lab notes.

Data Analysis

Answer Data Analysis Questions # 1-2 on the Report Sheet at this time BEFORE moving on.

NOTE: You will answer the remaining Data Analysis Questions for your post-lab analysis.

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Indicator Color before the Equivalence Point

Indicator Color at the Equivalence Point

Indicator Color after the Equivalence Point


PART II – Determination of an unknown solid acid sample.General Instructions1. Each pair is assigned a weak acid of unknown identity and pKa.

2. Measure the mass between 0.5-0.7 grams of your assigned acid sample on an analytical balance (or as directed by your instructor). BE SURE TO RECORD THE PRECISE MASS OF YOUR ACID SAMPLE to the correct sig figs!

3. Quantitatively transfer the acid sample into a 250 mL beaker, then dissolve, using about 40 mL of deionized water.

4. Titrate the ~40 mL of weak acid analyte with ~0.250 M NaOH standard to the equivalence point, using the drop counter apparatus to record the titration curve.

5. Print a copy of your titration curve and keep with your lab notes.

Table 1. Selected Physical Data for Weak Acids

Reference Table above taken from Lange”s Handbook of Chemistry, McGraw-Hill Book. Co., 1979.

mandelic acid – C8H8O3

(acid form)CAPS – C9H19NO3S (acid form)

HEPES – C8H18N2O4S (acid form)

Tricine – C6H13NO5 (acid form)

KHP – K+ C8H5O4−

(acid form)

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Acid pKa

Molar Mass

(g/mol)mandelic acid 3.4 152.1

potassium hydrogen phthalate (KHP) 5.41 204.23HEPES 7.55 238.3tricine 8.15 179.2CAPS 10.4 221.3

CI9 How Can Electricity Drive a Chemical Reaction??

ChemInquiry 7 Name_________________________ Partner’s Name ________________

Report SheetPart 1: Titration of Acetic Acid Sample

1. Click on the “examine” button in LoggerPro for your titration curve. Record the following values in the table below:

equivalence point volume and the pH . The equivalence volume can be located roughly using your titration curve; a better estimate can be made using the first derivative plot of the same curve. The derivative is simply the slope of the line pHV. The volume of base added at the equivalence point will correspond to the maximum derivative value (slope). The first derivative curve is plotted in green below the pH curve on your LoggerPro spreadsheet. Use the examine cursor to pinpoint the equivalence point.

pH at ½-equivalence point. Use the examine cursor to pinpoint the ½-way point.

Vol for the range pH+/-1 of the ½-way point . This is found by noting the titrant volume at pH = pHhalfway−1 and the titrant volume at pH = pHhalfway+1. The V is the difference of these two.

2. REPORT the following values to the class spreadsheet at the front of the lab room. Equivalence point pH

½-way point pH

The Vol for the range pH+/-1 at the ½-way point.

3. Based on the pooled class results, do the results confirm our prediction, from the Introduction (Part C, page 4), of the equivalence point pH range for a weak acid/strong base titration?

Circle one: Yes or No

4. Calculate the pKa of HOAc based on the average class value of the pH at the ½-way point. Show your calculation and put your answer in the box.

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Actual NaOH Concentration for term (M)

Equivalence point volume

Equivalence point pH

Vol for the range pH +/-1 at the equivalence point

pH value at ½ -equivalence point

Vol for the range pH +/-1 at half-way point

CI7 Acid/Base Interactions – Report Sheet

Experimental Class DatapKa of HOAc

5. Determine the Ka of HOAc based on the average class value. Show your calculation and put your answer in the box.

Experimental Class Data Ka of HOAc

6. The accepted value of the pKa of HOAc is 4.74. Calculate the accepted Ka of HOAc based on this value and put the value in the following box. Show your calculation and put your answer in the box.

Accepted valueKa of HOAc

7. Determine the %-error in the class determination of the Ka (NOT the pKa) of HOAc. Show your calculation in the space provided and write your % error value in the box.

%-error = | accepted-experimental |accepted × 100%

% error of classKa of HOAc

8. Determine the concentration of the acetic acid HOAc analyte in this titration. The following information is helpful. It is important to note that for 1:1 stoichiometry:

MacidVacid = MbaseVbase

IMPORTANT! Vacid = 25.00 mL…NOT 40 mL!

Show your calculation in the space provided and write your concentration value in the box.

Concentration ofHOAc

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Check all 5 previous numerical answers. Are the following correct for each value? (Worth 1 point)

Correct units? yes no Correct sig figs? yes no


9. The indicator was expected to change color right at or very near the equivalence point of the titration.

A. Comment on how well the indicator’s color transition matched the actual equivalence point.

B. Would this be a good indicator to use to locate the equivalence point if you did not have a pH meter? Explain why or why not.

10. Connection – buffers: VpH±1 represents a volume of titrant that can be added to the solution at some point, in which the pH only changes 2 units. Write your answers to the following questions in the boxes below.

A. Report the class average Vol for the range pH+/-1 at the ½-way point.

B. Report YOUR Vol for the range pH+/-1 found at the equivalence point.

C. Which region best “buffers” the solution from drastic changes in pH per mL of added base? In other words, in which region does the pH change the least per mL of added base?

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Class average Vol for the range pH +/-1 at the ½-way point

YOUR Vol for the range pH +/-1 at the equivalence point

Best “buffer” region (between which 2 pH values?)


Part 2: Determination of an unknown solid acid sample.DATA ANALYSIS11. Report the ID # and mass of your unknown.

Unknown ID#

Mass of unknown

12. Determine the pKa of the unknown acid based on your titration curve.

Experimental pKa value of unknown

13. Determine the molar mass of your unknown acid. Show your calculations in the space provided and write your value in the box below.

Hint: You have the mass of your sample. From the molarity and volume of the titrant, you can calculate the moles of your sample. Use the moles and mass to get the molar mass ratio in g/mol.

Experimental Molar mass value of unknown

14. Determine the identity of the unknown acid based on comparison to the table of weak acid possibilities given in Table 1 on p. 10 of the lab instructions. Draw the structure of your proposed acid and circle the acid functional group.

Identity of unknown (name)

Structure of unknown with circled functional group

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15. Determine the %-error in the determination of the molar mass of your unknown acid. Show your calculation and write your value in the box below.

% error of molar massvalue for unknown sample

16. Acceptable %-error is <5%. Comment on your level of error and possible sources of error in your pKa and molar mass determination. Poor comments include “human error” or “equipment error”. Excellent comments include an analysis of why your data might be high or low.

Could the error be systematic as in a miscalibration of a pH probe or drop volume?

Are there other factors?

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How Can Chemical Reactions Be Made to Produce

Electricity?


This part of the pre-lab assignment is due at the

beginning of the lab period, and must be done individually

before you come to lab!



II. Safety Hazards/Precautions 10. Complete the following table. Use the PCC MSDSonline link on




Hazard Statements


Pb(NO3)2

lead(II) nitrate


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Pre-lab Score: ____________/5

CI9 How Can Electricity Drive a Chemical Reaction?

Cu(NO3)2

copper (II) nitrate




Hazard Statements


Pb

lead metal (solid)


Mg

magnesium metal (solid)



etc.)





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III. Questions

1. How will you clean the metal strips in Experiment 1?

2. How many total potential reactions will you observe in Experiment 1?

3. In Experiment 2 when using the Mult-EChem module, you must make sure that it is filled with solution to a particular level. What is this level and why?

4. Describe how you will make the three 1:10 dilutions of Cu2+ starting with 0.20 M Cu2+(aq).

V. Work Plan (Procedural Flow Chart or Numbered List)

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Lab Partner___________________

Background & ProceduresQuestions of the Day

How do metals and metal ions interact? How can a metal activity series be determined? How can chemical reactions be made to produce electricity?


Introduction

Electrochemistry is a branch of chemistry that deals with the production of electricity by chemical reactions and the changes produced by an electrical current. Perhaps you have seen batteries produced by using a lemon and some strips of metal, or some similar devices that produce electricity via a chemical reaction. Of course, the most common examples of this phenomenon are batteries.

All chemistry in general is electrical in the sense that it involves the behavior of electrons and other charged particles (remember back to the definition of chemistry!). Electrochemistry is generally reserved, however, for the processes that convert chemical energy into electrical energy or vice versa.

The energy output that can be produced from the electrical current produced is measured with a quantity called voltage, which has units of volts (V). In a sense, measuring the voltage of an electrochemical cell is essentially measuring the degree of product favoredness of the reactions involved. In other words, a large voltage equates to a large value of K for the reaction.

Class Discussion (to be discussed in lab as a group)

Your instructor has been researching the relative reactivity of 2 newly discovered metal elements, Jimbonium (Jb) and Schneidoleum (Sh) by means of an electrochemistry experiment. In the lab, a piece of solid Sh was placed in a 0.20 M solution of Jb2+(aq), and in a second experiment, a piece of solid Jb was placed in a 0.20 M solution of Sh2+(aq). The following results were obtained:

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CTQ 1. Which one of the following conclusions is correct based on the observations? (Circle the letter of your answer)

A. Jb2+(aq) + 2 e Jb(s) Sh(s) Sh2+(aq) + 2 e

B. Sh2+(aq) + 2e Sh(s) Jb(s) Jb2+(aq) + 2 e

CTQ 2. Which is the correct net reaction? (Circle the letter of your answer)

A. Sh2+(aq) + Jb(s) Sh(s) + Jb2+(aq)

B. Jb2+(aq) + Sh(s) Jb(s) + Sh2+(aq)

Metal “Activity”

Metal “activity” is synonymous with “reactivity” in the sense of this experiment. When comparing the activity of 2 metals by doing an experiment like the one above,

I. Of the pair of metal elements, the most active metal is the one that forms or remains an ion in the reaction.

II. Of the pair of metal ions, the most active metal ion is the one that forms or remains as a solid metal in the reaction.

Critical Thinking QuestionsCTQ 3. Which is most active metal? Jb Sh

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CTQ 4. Which is most active ion? Jb2+ Sh2+

Electrochemical (Voltaic) CellsIn the last example, nothing much useful other than the “plating out” of solid Jb metal onto a Sh bar happened. However, the observation that the same reaction occurs over and over again without exception leads us to the conclusion that there must be a “driving force” or a “potential” for the reaction

to occur in this way and not the reverse. Analogously, water at the top of a waterfall has a “driving force” or “potential” to fall down and land in the pool at the bottom of the falls. We can take advantage of this potential. In the water analogy, we can stick a turbine or water wheel in the waterfall, and use the potential energy lost when the water falls to drive the turbine or water wheel and do work for us. In a similar way, electrons are “falling” from one metal to another. Falling from a higher potential to a lower one. In the example above, electrons in Sh are at a higher potential than they would be if they were in Jb, and when Sh and Jb2+ are in contact, there is a “driving force” for electrons to get from Sh to Jb/Jb2+. We can’t get any useful work out of it in this example, because when the electrons “fall” from the Sh(s) to the Jb2+(aq) such that Jb(s) is formed, (as well as Sh2+(aq)), there isn’t any place to stick an electrical “water wheel,” so to speak. To get useful work out, we have to separate the half reactions and make the electrons “fall” down a wire from the high potential side to the low potential side. As they travel in the wire, we can put in an electrical device, like a light, or a motor, or an MP3 player, etc. and use the “potential difference” to cause the device to run. This “potential difference” between the two half cells is commonly called a “voltage” and is measured in “volts.” This setup is known as an electrochemical (voltaic) cell, or more commonly, a “battery.”

Critical Thinking QuestionsCTQ 5. In which direction do electrons travel? To left To right

Cell Notation. We simplify the drawing of electrochemical cells by using “cell notation”, which combines the half-cell notation for the oxidation and reduction half cells. The complex drawing above is simplified into a cell notation as follows, using Zn|Cu electrochemical cell as an example:

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CTQ 6. Write the cell notation for the Sh|Jb electrochemical cell from the class discussion:

Experiment 1:Data Collection: Interaction of Metals and Metal Ions

1. Obtain a well plate, 4 rectangular pieces or strips of metal (Cu, Pb, Zn, Mg), and steel wool or sand paper.

2. Clean the 4 metal strips by using the steel wool or sand paper. Be sure to get them very clean and shiny. The metals MUST also be thoroughly recleaned BEFORE and AFTER each use.

3. Place the well plate on a blank sheet of paper on which you can make notes about the contents of the wells. Fill four wells about ¾ full with 0.20 M Cu2+. Do likewise with the other 0.20 M metal ion solutions (Pb2+, Zn2+, Mg2+). Sixteen wells will be about ¾ full when this step is completed.

4. Put one of the four metal strips into each well containing the Cu2+ solution. Allow these to set for a few minutes. A reaction may appear as a black residue on the metal strip.

You can wipe the metal strip on a paper towel to see if there is any residue.

Record your observations in Table 1 below; use NR for no reaction.

Remove the metal strips, clean them as before, and repeat the procedure with the Mg2+

solution, then the Zn2+ solution, and finally the Pb2+ solution.

5. When you have finished, empty the well plate into the waste jar. Rinse the well plate and invert it on a paper towel. Reclean the four metal strips with steel wool or sandpaper.

Table 1. Interaction of Metals and Metal Ions

Cu(s) Mg(s) Zn(s) Pb(s)Cu2+

(aq)

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Mg2+

(aq)Zn2+

(aq)Pb2+

(aq)

Critical Thinking QuestionsCTQ 24. Write the balanced chemical equations for each reaction you observed. Include phases of

matter, (s), (aq) in your equations.

Eqn. 1

Eqn. 2

Eqn. 3

Eqn. 4

Eqn. 5

Eqn. 6

CTQ 25. Choose one of your equations and explain why you chose the particular products you did. (i.e. why those products and not something else? What is your reasoning?)

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Experiment 2:Data Collection: Voltaic Cells

6. Obtain a Mult-EChem module, fill well #1 about ¾ full (so that the lower channel is filled) of the Zn2+ solution. Place the Zn strip into the Zn2+ solution. You have just created a Zn half cell (a metal in a solution of its ions). This cell is indicated by the symbols:

Zn|Zn2+ (0.2 M)

In wells #3, #5 and #7, do the same with each of the 3 metal ion solutions and metal strips to create Pb, Cu, and Mg half cells.

7. Prepare a “salt bridge” by filling the central well with 0.20 M KNO3. A salt bridge provides a “pathway” for positive and negative ions to move between cells, while not contaminating the one half cell with the metal ions of the other half cell.

8. An “electrochemical cell” is a combination of any two of the half cells, connected through the salt bridge. For example, an electrochemical cell is formed between wells 1 and 3, or 1 and 5, or 1 and 7, etc. In this setup, there are 6 possible electrochemical cell combinations (1-3, 1-5, 1-7, 3-5, 3-7, 5-7). The notation for the electrochemical cell which is composed of a combination of the two “half-cells” of Zn|Zn2+ and Cu2+|Cu is:

Zn|Zn2+(0.2 M)||Cu2+(0.2 M)|Cu

9. Follow the instructions for using the digital multimeter as demonstrated by your instructor. Use the alligator clips on the ends of the metal leads to clamp the leads onto the metal strips. You will have to be sure that the metal strips do not fall out of their half cells. Any time you find a negative voltage, simply reverse the leads. This setup is called a voltaic cell. Read the voltage of this voltaic cell. TO GET GOOD READINGS THE LEADS MUST BE CLAMPED TIGHTLY TO THE METALS.

10. Record the voltage of a new combination of half cells in Table 2 below.

11. Continue until you have measured the voltage of all six combinations of the four half cells

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12. Now repeat the process from Step 3 to obtain a second voltage reading for each combination. Calculate the average reading for the voltage for each combination of half cells.

13. When you have finished, empty the Mult-EChem module into the waste jug. Rinse it and invert it on a paper towel. Clean the four metal strips with steel wool or sand paper.

Table 2. Voltage Readings for Various Electrochemical Cells

Half-Cell Combinations

Voltage Reading #1 (V)

Voltage Reading #2 (V)

Average(V)

Zn|Zn2+ & Cu|Cu2+

Zn|Zn2+ & Pb|Pb2+

Zn|Zn2+ & Mg|Mg2+

Mg|Mg2+ & Pb|Pb2+

Mg|Mg2+ & Cu|Cu2+

Pb|Pb2+ & Cu|Cu2+


CTQ 26. Electricity is the flow of electrons. Choose one combination of cells and propose (hypothesize) how the flow of electrons travels by writing a complete cell notation for that cell.

Consider the tests from Experiment 1: did Zn solid form when Cu metal was placed in a Zn2+ solution (Zn2+ + 2 e Zn(s); Cu(s) Cu2+ + 2 e) OR did Cu solid form when Zn metal was placed in a Cu2+ solution (Cu2+ + 2 e Cu(s); Zn(s) Zn2+ + 2 e)?

CTQ 27. Write the reactions occurring at each of the half cells in the cell combination you chose.

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Oxidation Half-Reaction

Reduction Half-Reaction

CTQ 28. Combine the reactions from the two half-cells to produce a net reaction that gives a net flow of electrons.

Experiment 3:Data Collection: Concentration Effects

1. Fill one well on the Mult-EChem module about ¾ full (so that the lower channel is filled) of 0.20 M Zn2+ solution. Place the Zn metal strip in the Zn2+ solution. Fill a different well on the Mult-EChem module about ¾ full of the 0.20 M Cu2+ solution, but do not put in the Cu metal strip. Prepare the “salt bridge” by filling the central well with 1 M KNO3.

2. Dilutions. You will need to prepare solutions for three more Cu half cells. You will need half cells made of 0.020 M Cu2+, 0.0020 M Cu2+, and 0.00020 M Cu2+. You can make these solutions by using serial dilutions (successive 1:10 dilutions) as was done in previous labs. It is quickest to prepare these serial dilutions using a 10-mL graduated cylinder and a plastic pipes. Prepare 10 mL total of each dilution. Fill a well on the Mult-EChem module with each of the Cu2+

solutions until the channel to the central salt bridge is full.

3. Clean your copper metal strip and place it in the 0.00020 M Cu2+ solution.

4. Connect the voltmeter between the Cu|Cu2+ and Zn|Zn2+ half cells and record the voltage in Table 3 below.

5. Remove the Cu strip and rinse in distilled water. Dry the strip.

6. Place the Cu strip into the 0.0020 M Cu2+ solution measure the voltage between this half-cell and the Zn half-cell.

7. Repeat this procedure with the 0.020 M Cu2+ solution and the 0.20 M Cu2+ solution.

8. Now repeat the whole procedure since Step 4 to obtain a second voltage reading for each. Calculate the average reading for the voltage for each combination of half cells.

9. Report your data. Report your data of voltage and Cu2+ concentration for these cells in the class spreadsheet at the front of the room.

Table 3. Voltage Data for Zn|Cu cell at Varying Concentrations of Cu2+

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Net Reactio

n


Voltage Reading

#1 (V)

Voltage Reading

#2 (V)

Average(V)

Zn|Zn2+(0.20M) || Cu2+(0.00020M )|Cu

Zn|Zn2+(0.20M) || Cu2+(0.0020M )|Cu

Zn|Zn2+(0.20M) || Cu2+(0.020M )|Cu

Zn|Zn2+(0.20M) || Cu2+(0.20M )|Cu

Note: (Zn2+)o = 0.20 M in all trials

Data Analysis Answer these questions using completely legible handwriting and grammatically correct English sentences as appropriate.

Experiment 11. Claim 1: Rank the metals in terms of their reactivity from most reactive least reactive.

Most Reactive Least Reactive

2. Claim 2: Rank the metal ions in terms of their reactivity from most reactive least reactive.

Most Reactive Least Reactive

3. Connection: Are there any connections between the trends you observed in Claim 1 and Claim 2?

4. Mental Model: Draw a picture at the atomic level of the reaction between the Zn metal bar and Cu2+ ions in solution.

A. Show electron transfer at the Zn metal surface.

B. Show the oxidation of Zn atoms and the reduction of Cu2+ ions.

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C. Take this question seriously, and carefully consider what happens at the level atoms and ions as this reaction proceeds. Your picture model should essentially speak for itself and be clear as to what is going on. When you are done, show it to someone and ask them, “Do you understand what I am trying to say with this picture?” Adjust as necessary.

Experiment 2 – Voltaic Cells5. Claim 3: Rank the following cells in terms of their voltages from greatest least.

Electrochemical Cell

Rank 1 = greatest

3 = least

Zn|Zn2+||Cu2+|Cu

Zn|Zn2+||Pb2+|Pb

Pb|Pb2+||Cu2+|Cu

6. Analysis1: A. For each of the cells in Question 5 above, write the balanced net chemical equation. For

example, for the cell Zn|Zn2+||Cu2+|Cu you would write:

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) Ecell = ___

Electrochemical Cell Balanced Chemical Equation Ecell

(V)

Example: Zn|Zn2+||Cu2+|Cu Example: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Zn|Zn2+||Pb2+|Pb

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Pb|Pb2+||Cu2+|Cu

B. Add the reactions for the Zn|Zn2+||Pb2+|Pb and the Pb|Pb2+||Cu2+|Cu cells together. What is the

net reaction?

Net Reaction Ecell

(V)

C. Add the Ecell values for these two reactions to give the Ecell for the net reaction and enter in the table above. Compare to the Ecell you measured for the Zn|Zn2+||Cu2+|Cu cell. What do you notice?

7. Analysis 2: Carry out a similar analysis for similarly related cells such as

Mg|Mg2+||Cu2+|Cu Mg|Mg2+||Pb2+|Pb Pb|Pb2+||Cu2+|Cu

A. Add the reactions for two related cells together cells together. Determine the net reaction.

Reaction 1Balanced Equation

Ecell =

Reaction 2 Balanced Equation

Ecell =

Net Reaction

Ecell =

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B. Add the Ecell values for these two reactions to give the Ecell for the net reaction and enter in the rightmost column in the table above. Compare to the Ecell you measured for the cell represented by your net reaction above cell. What do you notice?

8. Synthesis Claim 1: Is voltage (E) a state function? A state function means the value of the function depends only on the current state of the system, not on how it got to this state. An example of a state function is enthalpy, H. You have previously used Hess’s Law to compute the enthalpy change for a reaction by adding different reactions (and their enthalpies) together. Is voltage a state function? To answer this question, you should look at Data Analysis questions 6 and 7. In effect, you are carrying out a Hess’s Law calculation. When you add two related reactions together, do their voltages add up to give the measured voltage of the net reaction? Provide an explanation along with data to support your claim.

9. Mental Model: Choose one of the voltaic cells you constructed in Experiment 2 and draw

an illustration of how the electrons and ions (both positive and negative) move through the cell.

A. Be sure to illustrate all processes taking place at the electrodes and in the solution.

B. Be sure to show electron travel in the electrodes and wire.

C. Take this question seriously, and carefully consider what happens at the level atoms and ions as this reaction proceeds. Your picture model should essentially speak for itself and be clear as to what is going on. When you are done, show it to someone and ask them, “Do you understand what I am trying to say with this picture?” Adjust as necessary.

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10. Synthesis Claim 2: Prepare a claim statement that describes the relationship or connection, if any, between the differences in reactivity you observed in Experiment 1 with the voltages of the different cells you observed in Experiment 2. You might start your claim with something like: “As the difference in reactivity between metal elements increases, the voltage of an electrochemical cell comprised of these metal elements…”

11. If you were manufacturing a battery from the half cells used in this Experiment, describe a design that would give the largest voltage possible from these materials (studied in this experiment).

12. In the diagram to the right is an activity series of metals.

A. What is the activity series? (Use your text if necessary.)

B. Create an activity series for the metals used in this experiment based on your experimental results (be sure your lab results match your series!)

Experiment 3 – The Effect of Concentration13. Summarize the class VOLTAGE data for the first data set of Experiment 3 in Table 4 below:

Table 4. Compiled Class Data Voltage from Exp 3

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Strength as reducing agent (increasing activity)


Ecell (V)(Class

Average)Q

Zn|Zn2+(0.20M) || Cu2+(0.00020M )|Cu

Zn|Zn2+(0.20M) || Cu2+(0.0020M )|Cu

Zn|Zn2+(0.20M) || Cu2+(0.020M )|Cu

Zn|Zn2+(0.20M) || Cu2+(0.20M )|Cu

14. Q = ¿¿ for Experiment 3, so compute Q for each trial and enter them into Table 4.

15. Claim 1: Propose a qualitative claim that explains how the working cell voltage E cell changes as Q K.

Note that K for this reaction is K ≈ 10+37! You might start your claim with “As Q approaches K, the working cell voltage E…”

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16. Determine the mathematical relationship between the working Ecell and Q.

A. Use MS Excel or LoggerPro to graph the following relationships between Ecell and Q (Your instructor will present a brief tutorial on how to use these programs if necessary.)

Ecell vs. Q

Ecell vs. 1/Q

Ecell vs. lnQ

Gradable Expectations for Linear Best-Fit Graph

Graph has appropriate title (Ex: Relationship between Ecell and Q)

Graph axes are labeled

Graph axes have appropriate units (if applicable)

No just connect-the-dots line

Best-fit linear line is displayed

Equation of best-fit line is displayed on the graph

B. Report the best fit line for your graph. Do NOT use variables “x” and “y” but rather use the physical variables studied in this experiment (Ecell and Q)

C.

D. If you are in lab when doing this, have your instructor check and initial your best fit line equation above.

E. ATTACH the graph that produced a linear relationship to the END of this lab report.

17. Nernst Equation: Ecell is a measure of how far the reaction is from equilibrium. As explored above, another measure of how far a reaction is from equilibrium is the reaction quotient, Q. Thus, we should expect there to be a mathematical relationship between Ecell and Q.

Observe your graph of class data from Data Analysis #16. The Zn|Zn2+||Cu2+|Cu cell reaction has been well studied, and the relationship between Ecell and Q (called the Nernst Equation) at 25oC has been determined:

Ecell = (−0.01285 V)lnQ + 1.1 V

y = m x + b

This is the equation of a straight line!

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Equation of Best-Fit Line

Instructor Initials


A. Use the best-fit line from your graph of the class data to predict the expected working voltage (E cell) when Q = Qeq = K. That is, when the cell has reached equilibrium. The accepted value for Keq is:

K = 1.62 × 1037

So, plug in the accepted numerical value of K for Q in “ln Q” in your equation, and compute Ecell

using your curve.

Show all work used to determine answer (no work = 0 points).

ExperimentalEcell =

B. The expected working cell voltage (Ecell) at equilibrium is 0.0V (Note: this is NOTEcello ¿. Compare

the class value with the accepted value and discuss any reasons for the discrepancy. (Hint: What are standard conditions? Did you operate your cells under standard conditions? What were sources of error in your experiment?)

C. The y-intercept value of the Nernst equation represents the cell voltage under standard state conditions, when Q = 1, or lnQ = ln(1) = 0. That is, it is the value of the y-variable (Ecell) when the x-variable (lnQ) = 0. This value of working voltage is called the standard state potential, Ecell

o .

The Nernst equation simplifies, under standard state conditions, to:

Ecell = Ecello = (−0.01285 V)ln(1) + 1.1 V

Eocell

= 1.1 V

This is the accepted value of Eocell for this reaction.

Report the experimentally predicted value of Ecello based on the class results. (i.e. what is your y-

intercept?)

Experimental Ecello =

D. Compute the %-error in the class value of E Ecello . The accepted value of Ecell

o for the Zn|Cu2+ cell is 1.1V. Propose and discuss reasons for any %-error > 5%.


%-error = |accepted-experimental|accepted × 100%

%-error =

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Proposed reasons for %-error > 5%

18. Computation of experimental value of Keq. Use the relation Ecello = RT

nFlnK to compute the

experimental value of K. Use your experimentally determined value of Ecello from your best-fit line and

the following values for the constants:

R = 8.314 Jmol∙K n = 2 mol e− F = 96,485

coulombsmol e−¿¿

T = 298 K


Experimental Keq =

19. Results Summary Table. Prepare a table similar to the one below.

Ecello

(Q = 1)

Ecell at equilibrium (Q = K) Keq

Accepted Value 1.1 V 0.0 V 1.62 × 1037

Experimental Value

%-difference

20. Attach the class data sheet to the end of this lab report.

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ChemInquiry 9 Name_________________________

How Can Electricity Drive a Chemical Reaction?

What is a Fuel Cell?

Pre-Lab Assignment

This pre-lab assignment is worth 5 points. This part of the pre-lab assignment is due at the beginning of

the lab period, and must be done individually before you come to lab!

I. Background Preparation Read this experiment thoughtfully FIRST: Mentally note any procedural questions and plan how you and your partner will complete all experiments efficiently during the three-hour lab period.

II. Safety Hazards/Precautions 1. Complete the following table. Use the PCC MSDSonline link on your lab web

page. Be sure to select the location Sylvania ST and be sure the chemical name matches precisely.

Materials GHS Pictograms


Hazard Statements


HNO3 nitric acid


H2SO4 sulfuric acid

Corrosive Toxic ___________________ Flammable Reactive _________________

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Pre-lab Score: ___________/5


Irritant Other? __________________

Materials GHS Pictograms


Hazard Statements


Pb lead

metal

(solid)


H2 hydrogen

(gas)


Other hazards (equipment,

glassware, etc.)





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III. Questions 1. What safety precautions are needed when using the power supply? 2. How will you clean the copper bar electrode in Experiment 3? 3. What is the range of current (low to high) in amps (A) to be measured in Experiment 3? And

what is the maximum current to not go over? 4. What do you have to make sure not to do at the end of Experiment 3 so you are able to

complete all of the Data Analysis?

IV. Work Plan (Procedural Flow Chart or Numbered List)

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CI9 How Can Chemical Reactions be made to Produce Electricity? What is a Fuel Cell?

ChemInquiry 9 Name___________________________

Lab Partner______________________


How can electrical energy drive a reactant-favored reaction? How can we measure the amount of reaction in an electrolytic cell? How does a H2/O2 fuel cell work?


Introduction

Voltaic cells use a product-favored (spontaneous) chemical reaction to drive an electric current through an external circuit. These cells are an extremely important as this class of oxidation and reduction reactions are used to provide useful electrical energy as batteries. A simple electrochemical cell can be made from copper and zinc metals with solutions of their sulfates. You have studied in detail such cells in the previous ChemInquiry 8 Lab.

Electrochemical cells which generate an electric current are called voltaic cells or galvanic cells, and common batteries consist of one or more such cells.

Figure 1. A schematic of a common voltaic (a.k.a. galvanic) cell.

Electrolytic CellsVoltaic cells in general do work on their surroundings (such as a turning on a light bulb or powering an electronic device) as a result of this driving force for electrons to move from the oxidation half reaction to the reduction half reaction solution. Voltaic cells are not the only kind of electrochemical cells,

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CI9 How Can Chemical Reactions Be Made To Produce Electricity?

however. It is also possible to construct a cell that does work on a chemical system by driving an electric current through the system. These cells are called electrolytic cells. Electrolysis is used to drive an oxidation-reduction reaction in a direction in which it does not occur naturally.

In the electrolytic cell, energy from an applied voltage is used to drive an otherwise reactant-favored (nonspontaneous) reaction. Such a cell could be produced by applying a reverse voltage to a voltaic cell. In other words, using an external voltage supply to cause electrons to travel in the reverse direction in which they would naturally move in such a cell. If a voltage greater than 1.10 volts is applied to the Zn|Cu cell under standard state conditions, then the reaction

Cu(s) + Zn2+(aq) Zn(s) + Cu2+(aq) Eocell

= −1.10 V

will be driven by removing Cu from the copper electrode (Cu(s) Cu2+(aq) + 2e) and plating zinc on the zinc electrode (Zn+2 (aq) + 2 e Zn(s)) (this is the reaction of the electrochemical cell in reverse). Electrolytic processes are very important for the preparation of pure substances like aluminum and chlorine.

If the electrical work is stopped (remove the voltage source), the reaction stops. Such a process is non-spontaneous as written. The value of Eo is negative for the Cu/Zn reaction illustrated above (indicating a reactant-favored process).

Figure 2. Comparison between a galvanic (voltaic) cell (a) and an electrolytic cell (b).

Although at first such processes might not seem very useful they are sometimes the only way to obtain certain elements. Some very useful commercial and ornamental processes such as plating are also electrolytic in nature. Electrolysis can be roughly divided into three categories:

the electrolysis of molten salts electroplating the electrolysis of aqueous solutions

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− +− +e− e−

e−e−


Of the three, perhaps the first is the simplest. Molten salts generally contain cations and anions which can be reduced or oxidized to elements. An example is sodium chloride. Electrolysis of a molten bath of sodium chloride will produce sodium metal and chlorine gas.

Electroplating is a process of covering one metal (or other conductive object) with another. Generally, the metal that is to be plated onto another object is the anode in the cell (recall anode = oxidation, it will dissolve into the mixture). The ions of this metal are then reduced at the cathode to cover it (recall cathode = reduction). Because the cathode, which will be covered with the new solid metal, is negatively charged, it attracts the positively charged ions from the anode (oxidation). The positively charged ions reach the cathode, and electrons flow from the cathode to the cations, reducing the cation to a solid metal (reduction). This solid metal coats the cathode and “plates” out. It is interesting to point out that the charges on the cathode and anode are reversed in the electrolytic cell.

Plating can take place in either aqueous solution or in molten salt baths. As demonstrated in CI8, metals can spontaneously plate out onto other more active metals in a simple aqueous displacement process but the results are typically not attractive unless the process is very slow (dilute solutions) and adherence of the coating is often poor. That is why almost all practical plating is done in electrolytic cells.

H2/O2 Fuel Cells

A fuel cell is a galvanic cell that requires a constant external supply of reactants because the products of the reaction are continuously removed. Unlike a battery, it does not store chemical or electrical energy; a fuel cell allows electrical energy to be extracted directly from a chemical reaction. In principle, this should be a more efficient process than, for example, burning the fuel to drive an internal combustion engine that turns a generator, which is typically less than 40% efficient, and in fact, the efficiency of a fuel cell is generally between 40% and 60%.

Unfortunately, significant cost and reliability problems have hindered the wide-scale adoption of fuel cells. In practice, their use has been restricted to applications in which mass may be a significant cost factor, such as US manned space vehicles. These space vehicles use a hydrogen/oxygen fuel cell that requires a continuous input of H2(g) and O2(g), as illustrated in Figure 3.

Figure 3. A schematic of a H2/O2 fuel cell, similar to the one used in this experiment (left). Image of the fuel cell/electrolyzer used in today’s experiment (right).

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The electrode reactions are as follows:

cathode: O2(g) + 4 H+ + 4e− → 2 H2O(g)anode: 2 H2(g) → 4 H+ + 4e−

overall: 2 H2(g) + O2(g) → 2 H2O(g)The overall reaction represents an essentially pollution-free conversion of hydrogen and oxygen into water, which in space vehicles is collected and used as drinking and other water. Although this type of fuel cell should produce 1.23 V under standard conditions, in practice the device achieves only about 0.9 V. One of the major barriers to achieving greater efficiency is the fact that the four-electron reduction of O2(g) at the cathode is intrinsically rather slow, which limits current that can be achieved. All major automobile manufacturers have major research programs involving fuel cells: one of the most important goals is the development of a better catalyst for the reduction of O2.

Closing the Loop. Renewable and Sustainable Sources of PowerInterestingly, many fuel cells can be used as electrolytic cells (such as the fuel cell you will observe in today’s lab). With a renewable (or other) supply of outside energy (such as from a solar cell, wind or water turbine, battery, etc.), a fuel cell can be run in reverse (as an electrolytic cell) and used to generate H2 and O2. Consider connecting a bank of solar cells to H2/O2 fuel cells. Free energy (apart from the materials used to construct and maintain the solar cells) is captured from sunlight and used to power a fuel cell in electrolytic mode. The solar energy, through the solar cell, converts solar energy into electrical energy (the most efficient form of power we have). The electrical potential is then converted into stored chemical energy (2 H2O 2 H2 + O2, Keq << 1), through the use of a fuel cell. The chemical energy is “stored”, not in the H2 and O2 bonds, but rather in the difference in potential energy (mostly bond enthalpy) of the products compared to the reactants. The following figure should make sense based on your previous knowledge of chemistry!

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from sun

Energy released as fuel cell operates

H2O

H2, O2

Potential Energy

Oxidation of H2 Reduction of O2

Pollution-free/ high-purity drinking water

Proton (H+) exchange membrane

e−e−


When power is necessary, we simply run the fuel cell in the forward direction, converting the stored chemical potential due to the difference in potential energy between H2/O2 and H2O back into electrical energy, which we can use to drive whatever we want! Think about it, free solar energy which might otherwise be “wasted” can be converted through electrical and chemical means (which you can now understand!), into stored chemical potential, ready for us to harness at our whim, even if the sun isn’t shining.

Consider the difference between this and the fossil fuel cycle currently used: free energy from the sun is captured by plants and stored as chemical potential energy (sugar and O2 vs. CO2 and water) through the process of photosynthesis (solar cells are an analog to chloroplasts in plants). The chemical potential energy stored in plants is over millions and millions of years transformed into fossil fuels (the analog of H2/O2 in our fuel cell system). Finally, the chemical potential energy is obtained through combustion in an internal combustion engine, a highly inefficient and highly polluting process (the analog of using a fuel cell)! What a better potential technology the renewable fuel cell cycle is! Water is plentiful on planet earth (although most water is not fresh water). We know much of the background science necessary for the above process. The problem is now one of developing cost-effective technology to implement it. Perhaps you will be involved in the future development of these systems! Make it happen! Be the future!

Experiment I: Data Collection: Lead and copper reactions

1. Obtain approximately 50 mL of 0.2 M solutions of Cu(NO3)2 and Pb(NO3)2 and place each in separate 100 or 150 mL beakers. Clean a copper and lead electrode by sanding lightly one inch from the bottom of the tip. Place the lead electrode in the Cu(NO3)2 solution and the copper electrode in the Pb(NO3)2 solution. Let them remain in the solution for 5 minutes. Record any observations below.

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Critical Thinking Question

CTQ 1. Based on these observations, what is the preferred reaction between lead and copper? Write the reduction and oxidation half reactions, as well as the net reaction.

2. Place a second sanded copper electrode in the Pb(NO3)2 solution separated from the first Cu electrode. (It might help to put a test tube between the electrodes to keep them separated.) Attach the DC Power Supply to the electrodes as shown. Note which electrode is attached to which pole of the source. Your apparatus should look similar to the diagrams below.

3. Plug the Power Supply into a socket on your lab bench, but DO NOT yet turn the power on.

4. Turn the CURRENT knob clockwise about ½-way to maximum.5. Make sure the VOLTAGE knob is fully counterclockwise (OFF).

6. Turn the POWER switch ON.

7. Slowly turn the voltage knob clockwise until the VOLT reading shows about 2.5 V.

8. Run the experiment for approximately 5 minutes. Record your observations below:

9. When finished, turn the VOLTAGE knob OFF (fully counterclockwise) and then turn off the power switch.


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Net Reaction


CTQ 2. Compare your results obtained in Step 1 with your results obtained in (Steps 2-9). Why is the power supply necessary in Steps 2-9 but not in Step 1?

CTQ 3. For Steps 2-9, write the reaction occurring at the reduction () electrode, oxidation (+) electrode, and the net reaction.



Net Reaction

CTQ 4. On which side of the reaction would you write “Energy”? (Circle your choice.)

Reactants (endothermic) Products (exothermic)

CTQ 5. How does the net reaction in CTQ 3 compare to the “preferred” reaction in CTQ 1?

Experiment 2: Data Collection: The copper and H+ reaction - Qualitative

1. Set up a 250 mL beaker as above and fill it with about 120 mL of distilled water. Fill a SMALL test tube with distilled water and invert it into the beaker such that no air is trapped in the test tube. A rubber stopper may be helpful.

2. Carefully add, while stirring, 40 mL of 4 M H2SO4.

CAUTION: 4 M H2SO4 is a concentrated, strong acid. Sulfuric acid can cause severe burns to skin and eyes. Be sure your goggles are on. If sulfuric acid comes into contact with your skin or eyes, flush with plenty of water for 15 minutes.

3. Substitute an S-shaped copper electrode (from the reagent cart) for one of the copper bars and connect it to the NEGATIVE electrode of the power supply. Connect the power supply to the copper electrodes in the same manner as you did in Experiment 1 (except now you have H2SO4 in solution instead of Pb(NO3)2). Be careful to note which electrodes are connected to the positive and negative poles. Follow the same

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NOTE!


procedures from Expt. 1 to power up the power supply to about 1.5 V. Run the reaction for about 5 minutes. Record your observations below.

4. Collect some of the gas being produced by capturing the bubbles in the inverted small test tube. Don’t let the test tube directly rest on the S-hook: acid from the beaker needs to be in contact with the

metal electrode surface. When the test tube is filled with gas, carefully remove the test tube from the beaker, carefully holding it

upside down. Hold a lighted match to its opening. Record any observations below.

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Critical Thinking QuestionsCTQ 6. What is the composition of the gas produced in Experiment 2?

CTQ 7. In the table below, do the following: Write a half-reaction which describes the chemistry at the electrode connected to the negative pole of

the power supply. Use the table of half-reactions supplied as an Appendix to this lab. (The half-reactions can be reversed if that will better describe the reaction being studied.)

Write a half-reaction which describes the chemistry at the electrode connected to the positive pole of the power supply.

Combine the two half-reactions above to describe the net reaction which occurred in the beaker.

Negative Pole Reaction

Positive Pole Reaction

Net Reaction

CTQ 8. Draw a diagram of the apparatus tracing the flow of electrons

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To LabPro

Vernier CurrentProbe


Experiment 3: Data Collection: The copper and H+ reaction - Quantitative

0. The data tables for Experiment 3 are found at the end of this set of instructions on page 10-11.

1. Obtain and clean a 50 mL buret. Use a 10-mL graduated cylinder to measure the volume contained in the 50 mL buret between the 50 mL mark and the stopcock. IMPORTANT: Be sure the buret tip is FULL before recording the volume.

HOWEVER, do not let the water drain out of the tip. Record this value in the blank at the top of Data Table 1 for Experiment 3.

2. Put about 375 mL of distilled water in a 600 mL beaker. Carefully add 125 mL of 4 M H2SO4. The final solution will be 1 M.

CAUTION: 4 M and 1 M H2SO4 is a concentrated, strong acid. Sulfuric acid can cause severe burns to skin and eyes. Be sure your goggles are on. If sulfuric acid comes into contact with your skin or eyes, flush with plenty of water for 15 minutes.

3. Open the stop-cock on the buret. Invert the buret into the solution and carefully fill it by drawing the solution into the buret using a suction bulb. Close the stop-cock when the space below the stop-cock is just filled. Leave the tip empty. Check for leaks.

4. Obtain your S-shaped copper electrode. Bend and place the wire so that one of the stripped ends is entirely inside the buret and the other end is outside the beaker. Use the diagram below to aid in your set-up.

5. Clean another copper bar electrode by sanding, dipping it into 6M nitric acid, rinsing it with water and drying it.

CAUTION: 6 M HNO3

(nitric acid) is a concentrated, strong acid. Nitric acid can cause severe burns to skin and eyes. Be sure your goggles are on. If nitric acid comes into contact with your skin or eyes, flush with plenty of water for 15 minutes.

6. Weigh the electrode as accurately as is possible on an analytical balance. Record the value in your data table.

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Measure this volume


7. Connect the Vernier Current Probe in the circuit as shown. Insert the cable from the Current Probe into Channel 1 of the LabPro. On the computer, open LoggerPro, and open the “Electrolysis” file from the “Open” menu on the computer.

8. Plug the DC power source into the electrical socket on your bench.

9. Turn the VOLTAGE knob clockwise ½-way.

10. Turn the CURRENT knob fully counterclockwise (OFF).

11. Turn the POWER switch ON.

12. Click “Collect” in the LoggerPro program. Immediately adjust the current knob to between 0.10 amp and 0.15 amp, and monitor the current on the computer screen. Be careful not to accidentally adjust the current over 0.60 amps or you will destroy the Current Probe. Record the current in Data Table 2 on the next page. Check the current every seven minutes and adjust if necessary, keeping the readout to between 0.10 – 0.15 amps.

13. Obtain an atmospheric pressure and temperature reading and record these in your data table.

14. Six times during the progress (approximately every 7 minutes) of the reaction, record on Data Table 2 the volume reading and the elapsed time. Be careful when reading the graduations on the buret to note that they are upside down. Determine the actual volume of gas collected and record this value in your Data Table 2.

15. Sometime before the collected gas reaches the zero mark on the buret, turn OFF the CURRENT knob (counterclockwise) and then turn the POWER switch OFF. Also, note the elapsed time from the graph, and the final volume reading. Record these in Data Table 2.

16. Very carefully remove the flat copper bar electrode from the solution. Dip the copper bar electrode into a beaker of distilled water, being careful not to dislodge any material. Then dip/rinse the same electrode in acetone and allow it to air dry until no more moisture remains. Weigh the electrode and record this data in your Data Table 1.

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SAVE YOUR LOGGER PRO GRAPH FOR

LATER DATA ANALYSIS.

DO NOT CLOSE OUT LOGGER

PRO!!


Data Table 1. Experiment 3Volume of buret between 50-mL mark stopcock mL

Initial mass Cu electrode g

Final mass Cu electrode g

Elapsed time of experiment sec

Atmospheric Pressure mmHg

Temperature oC

Data Table 2. Gas Volume Produced in Experiment 3.It is recommended that you read the buret and write down the volume reading. The gas will be evolving rather quickly. Just read the mL marking of the liquid in the buret as it is – upside down. Then, perform the math and determine the volume of gas collected after obtaining the volume reading.

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TIME ELAPSED

(sec)

BURET READING(mL mark)

Current Reading (Amps, A)

VOLUME GAS COLLECTED

(mL)


Experiment 4: The Hydrogen-Oxygen Fuel Cell

Demo: Electrolysis of Water demo

Your instructor will lead you through the setup and functioning of your fuel cell. Answer the questions below as you experiment.

Expt. 1: Quantitative Measurement of Gas Generation RatesThe purpose of this experiment is to measure the rate that hydrogen and oxygen gas are produced by the fuel cell.

1. Follow the assembly procedures as demonstrated by your instructor.

A. Fill the fuel cell with DI water as demonstrated.

B. Connect (be sure it is turned OFF first) the battery pack.

a. NEGATIVE (−) connects to H2

b. POSITIVE (+) connects to O2

C. Be sure the gas collection tubes are prepared and set at 0 mL.

In this experiment, you will need a timer and someone to watch the oxygen tank and someone to watch the hydrogen tank. The tank watchers can also write down their data or have someone else record the results in the data table below.

The timer starts timing when the first bubbles start to enter the tanks. The recorders enter the data under the tank reading column for either hydrogen or oxygen. Be ready as soon as the timer turns on the battery pack switch.

Data Table 3. Gas Production RatesHydrogen Oxygen

Time (sec.)

Hydrogen Tank

Reading (mL)

Time (sec.)

Hydrogen Tank

Reading (mL)

Time (sec.)

Oxygen Tank

Reading (mL)

Time (sec.)

Oxygen Tank

Reading (mL)

15 102 15 102

30 120 30 120

45 135 45 135

60 150 60 150

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SAVE YOUR LOGGER PRO GRAPH FOR LATER DATA ANALYSIS.

DO NOT CLOSE OUT LOGGER PRO!!


75 165 75 165

90 180 90 180

Now graph the data on LoggerPro. Use two colors: one color for hydrogen and a different color for oxygen. Time in seconds is on the x-axis and gas volume in milliliters (mL) is on the y-axis.

Determine the best fit straight line for each curve. Report the equations below.

H2 (reduction)

O2 (oxidation)

The slope of each line will represent the rate of gas production.

Rate of H2 production = mL/sec.

Rate of O2 production = mL/sec.

7. Determine the ratio of H2 production to O2 production. Explain why this should be so.

Ratio of H 2

O2 production =

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Optional Extra Credit Experiments. (0.5 points each for raw data) E.C. 1: Measuring Current and Voltage during Fuel Cell ElectrolysisYour instructor will demonstrate how to measure voltage and current during electrolysis using a multimeter.

Follow the directions carefully and take readings for voltage and current while the fuel cell is producing hydrogen and oxygen. Remember to turn OFF the battery power to stop the electrolysis when you switch from measuring voltage to measuring current.

Voltage reading for fuel cell electrolysis = Volts

Current reading for fuel cell electrolysis = Amps (not mA)

E.C. 2: Experiment Measuring Voltage and Current Output of the Fuel Cell

Our final goal is to find the efficiency of the fuel cell as it is used to produce mechanical energy. The purpose of Experiment 3 is to measure the current and voltage while the fuel cell is “working” and while it is not (no-load).

1. Measure the no-load voltage. Disconnect the fuel cell from battery supply. Make sure the H2 and O2

reservoirs are connected. Set the multimeter to V (voltage) and connect the leads to the terminals on the fuel cell. You may need to use a couple of wire leads to be able to take a reading. Use the small wires that come with the kit to plug into the fuel cell. Connect the voltmeter leads to these wires.

Measured no-load voltage = V

2. Measure the voltage under load. Detach the voltmeter temporarily. Connect a small motor to the leads from the fuel cell. It doesn’t matter which lead goes where on the motor. You might want to put a piece of tape on the motor axle to see it spin. Connect the voltmeter to the same connections on the motor as is the fuel cell leads. Wait until the voltage settles down to a steady reading.

Measured operating voltage = V

3. Change the wiring so that you can measure the short circuit current. Remember that to read current, the multimeter must be in line (i.e. in series) with the circuit. You will NOT connect a motor to

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measure the short circuit current.

Measured short circuit current = A

Data Analysis 20. What is the relationship between the volume of gas and the elapsed time? Develop an algebraic

equation which expresses this relationship. It will be helpful to graph your data and develop the algebraic equation from the plot.

A. Use MS Excel or LoggerPro to graph the volume and time data. Once plotted, determine the equation of the best-fit line.

Gradable Expectations for Linear Best-Fit Graph

Check off that you have completed each criteria below

Graph has appropriate title

Graph axes are labeled

Graph axes have appropriate units (if applicable)

No just connect-the-dots line!

Best-fit curve is displayed

Equation of best-fit curve is displayed on the graph

B. Report the best fit line for your graph. Do NOT use variables “x” and “y” but rather use the physical variables studied in this experiment as well as their UNITS.

C. If you are in lab when doing this, have your instructor check and initial your best fit line equation above. ATTACH the graph that produced a linear relationship to the END of this lab report.

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Instructor Initials

Equation of

Best-Fit Line


21. Why are the volume of the gas and time related in the manner described by your equation? For example, if the relation is linear, why should it be linear and not something else? If the relation is a polynomial, why should it be a polynomial and not something else? Etc. You should consider the following in your discussion:

Recall from earlier in the term, that a graph of concentration (amount) vs. time for a reactant or product can give us the rate of the reaction. The RATE of the reaction (in this case the rate at which the gas is produced) can be determined by looking at the slope of the tangent line in your graph of gas volume vs. time. You can assume that the volume of the gas produced is proportional to the moles of gas produced (because of the ideal gas law; note P and T are assumed to be constant).

Examine the graph you made in Data Analysis 1. Does the slope (rate of production) remain constant, or does it change (increase or decrease with time)?

Examine the stoichiometry of the reduction half reaction: With what are the moles of gas proportional (voltage? current? # e? moles? etc). How is this variable related to production of gas according to the half-reactions that you previously developed; how is this related to your algebraic equation developed above?)

Combine these ideas to explain why your graph of gas volume and time are related the way they are, and not some other way.

22. Moles of gas produced (reduction): Use the ideal gas law to calculate the TOTAL number of moles of gas produced during the experiment. Show your work. IF the gas in the buret fills the buret down to the water line in the beaker, the TOTAL pressure of the gas inside the buret can be approximated to be equal to the atmospheric pressure minus the vapor pressure of water.

A. IMPORTANT: You will have to subtract off the vapor pressure of water at this temperature from atmospheric pressure to obtain the pressure of the gas alone. In other words:

Ptot = Patm = Pgas + Pwater vapor

The vapor pressure of water at various temperatures can be found by doing an internet search.

[Recall that in the ideal gas equation, R = 0.08206 L·atm/mol·K. You will have to make sure your units for T, V, and P match this.]

Work Area (show work-including units, for credit, no work = zero credit)

Moles of gas produced:

23. Moles of copper metal consumed (oxidized):

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Method 1: mass data: From the masses of the Cu electrode before and after the experiment, calculate the number of moles of Cu consumed during the experiment. Show your work.


Mass of Cu consumed (mass data):

Moles of Cu consumed (mass data):

Method 2: current/time data: Calculate the moles of Cu reacted by using the current and the time you ran the experiment. Faraday’s Law states that the amount of a substance consumed or produced at one of the electrodes in an electrolytic cell is directly proportional to the amount of electrical current that passes through the cell.

Current (amperes = A) is defined as the amount of electric charge (Coulombs = C) that flows through a circuit per time (seconds = s). Stated mathematically: 1 ampere = 1 coulomb/second or 1 coulomb = 1 ampere · second 1 A = 1C/s or 1 C = 1 A · s

a) One electron has a charge of 1.602 1019 C. Calculate the charge carried by one mole of electrons. The value you have calculated is called Faraday’s constant, F. It has units of C/mole e.


Charge carried by one mole of electrons:

b) Determine the number of coulombs of charge delivered during this experiment. To do so, use the Integral function of LoggerPro to find the area under your Current vs. Time graph. Click on the button at the top of the screen. The program will compute the area under the curve in this region (Current Time) and display this result in a window on the graph. Note the units will be A·s, and 1 As = 1 C.


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Coulombs of charge delivered:

c) Determine the moles of electrons transferred during this experiment based on your answer to part b).


Moles of electrons transferred:

d) Using the oxidation halfreaction and the moles of electrons transferred, determine the mass then the number of moles of Cu atoms oxidized during the reaction based on current/time data (i.e. counting electrons). Show your work.


Mass of Cu consumed (current/time data):

Moles of Cu consumed (current/time data):

24. Compile all your answers from Data Analysis 3-4 in the Results Table below.

Results Table

Moles H2 gas produced mole H2 (g)

Mass Cu consumed (mass data) g Cu(s)

Moles Cu consumed (mass data) mole Cu(s)

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Charge carried by 1 mole electrons Coulombs

Coulombs of charge delivered Coulombs

Moles of electrons transferred mole e−

Mass of Cu consumed (current/time data) g Cu(s)

Moles of Cu consumed (current/time data) mole Cu(s)

25. Compute the % difference between the mass of copper lost as determined by weighing the electrode and the mass of copper lost determined from the current vs. time data above.

For any two measured values A & B of the same quantity,

%-difference =

Comment on your result and suggest reasons for any discrepancies

26. Compare the moles of H2 and Cu produced or consumed during this reaction. Are they similar? If they are similar, why should this be so? If they are not similar, why should they not be? Answer using grammatically correct, COMPLETE English sentences, using legible handwriting. Failure to do so will result in the loss of points.

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Extra Credit: DATA ANALYSIS Part II - The Hydrogen-Oxygen Fuel Cell (worth up to 2 points for Data Analysis)1. Efficiency of the Fuel Cell at Converting Electrical Energy to Chemical

Energy

You will determine the efficiency of the water electrolysis using the fuel cell. It requires data from the Expt. 1 and 2 (under Experiment 4 on Hydrogen-Oxygen Fuel Cell).

Efficiency is a measure of how good a fuel cell is at storing energy (from battery, or a solar cell, or any other source of electrical energy) as chemical energy in the form of hydrogen gas. 100% would mean that all of the electrical energy from the solar cell produces a chemical change resulting in hydrogen gas. That would be impossible, since the Second Law of Thermodynamics says that in any real process, some energy must be lost as heat to the surroundings and cannot be converted into work.

For a fuel cell, we would really only be interested in the conversion of water into hydrogen, since that is the fuel. Oxygen, which is the oxidizer in the reaction, is freely obtained from the environment.

To determine efficiency, you need the power produced by the fuel cell divided by the power put into the fuel cell (times 100). The power consumed or produced is measured in watts (W).

1 W = 1 Js . It is a measure of the rate at which energy (in J) is consumed.

The efficiency of H2 production for the fuel cell is measured as follows:

Efficiency = output powerinput power × 100%

Output Power (O.P.) = rate of H2 production (mol/sec) × enthalpy of combustion of H2 per mol (J/mol)

To keep life simple, we’ve done the math to convert mL of H2 at STP to mol. This conversion factor is: 4.166 × 10−5 mol H2/mL. The enthalpy of combustion of H2 = 2.37 × 105 J/mol

Thus, compute the output power of your fuel cell as follows. Show your work and place your answer in the box below.

O.P. = Rate of H2 production (mL/s) (Expt. 1) × 4.166 × 10−5 mol H2/mL × 2.37 × 105 J/mol

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O.P. = W

Input Power (I.P.) In an electrical circuit, the formula for the determination of power consumption is as follows:

Power = Voltage (V) × Current (A)

Using the values of voltage and current from E.C. 1, compute the power consumed to produce H2 gas by your fuel cell. Show your work and place your answer in the box below.

I.P. = W

% EfficiencyCalculate the % efficiency of your fuel cell:

% efficiency = O.P.I.P. ×100%

Show your work and place your answer in the box below.

% efficiency =

This tells you how good the fuel cell is at storing electricity as hydrogen gas.

2. Efficiency of the Fuel Cell OutputThe purpose of this analysis is to determine the efficiency of the fuel cell and to compare that efficiency to other power generating devices.

Remember that efficiency shows how much power comes out of a device compared to how much power goes into the device. You will use the equation “power out / power in.”

With the fuel cell there are two conversions that take place: 1) changing the electricity from a solar cell or battery into stored hydrogen, and 2) changing the stored hydrogen into electrical energy.

The overall efficiency of the fuel cell will combine both steps and will use a unique formula to show the efficiency. The result will not be exact, but it will be close enough for getting a rough estimate for our purposes.

Overall Efficiency = operating voltage electrolysis voltage × 100%

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Use the operating voltage from E.C. 2 and the electrolysis voltage from E.C. 1 to calculate the total efficiency of your fuel cell.

Operating voltage (E.C. 2) = V

Electrolysis voltage (E.C. 1) = V

Determine the efficiency of your fuel cell. Show your work and place your answer in the box below.

% efficiency = 3. Efficiency ComparisonsThe chart below gives an approximate efficiency for the variety of machines listed.

1. Comment on how the efficiency of your fuel cell compares to the machines listed in the chart. Can you think of why scientists are so interested in fuel cells?

2. What new questions do you have or experiments would you do?

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Steam engine

Gasoline engine

Steam turbine

Diesel engine

Electric motor

Human body

Bicycle

Efficiency


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CI10 What is a Polymer?


What is a Polymer?

orLet’s Make Polymers and Get the Heck Outta Here!


This part of the pre-lab assignment is due at the beginning

of the lab period, and must be done individually before you

come to lab!



II. Safety Hazards/Precautions1. Complete the following table. Use the PCC MSDSonline link on




Hazard Statements


1,6-hexane-diamine

“hexanediamine”


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Pre-lab Score: ____________/5


adipoyl chloride

“adipic acid chloride”




Hazard Statements


acetone


Other hazards(equipment, glassware, etc.)




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How are polymers created? What are characteristics of polymers that make them different from other chemical compounds? What is a “cross-linked” polymer?

Introduction

Polymers are extremely large molecules (molar masses from 1000 to greater than 106 g/mole!) which result from chemically linking thousands of relatively small molecules called monomers. Dramatic changes in physical properties accompany this process. Monomers, due to their weak intermolecular forces, are either gases, liquids or structurally weak molecular solids. In a polymerization reaction, these monomers are joined together to form larger molecules.

As the attractive forces between the molecules increase, the mixture becomes more viscous. Eventually the molecules become so large that their chains become entangled. It is this entanglement of the individual polymer chains that give polymeric materials their characteristic properties. A classic example of the changes accompanying polymerization is given by the conversion of ethylene to polyethylene (PE). Ethylene, CH2CH2, is a gas at room temperature while polyethylene is a rigid plastic used widely as storage containers.

While chain entanglement is a common structural feature in all polymeric materials, other structural features give different polymers their unique properties, such as those shown below. These include: 1) the flexibility and microstructure of individual polymer chains, 2) the strength of interchain forces and most importantly, 3) the presence of crosslinks (usually covalent bonds) between the polymer chains.

Polymers can be cross-linked if the polymer molecules that were formed still have a reactive place on the molecule, such as a double bond or a reactive side group. With heat or chemical methods, those reactive areas can be bonded together to make a stronger, more rigid, material. Crosslinking restricts the ability of individual polymer chains to slide past each other. The result is a polymer with a three-dimensional network structure. Depending on the degree and

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type of crosslinking, various properties can be obtained. Low crosslink density materials, such as elastomers (e.g., rubber bands), are elastic and deformable. Highly cross-linked materials, such as Bakelite (used in pot handles and billiard balls) or some epoxies, are more rigid and brittle. Temporarily cross-linked materials behave as viscous, liquid-like gels.

Rubber is an example of a cross-linked polymer. Cross-linked polyesters with which you may be familiar include so-called glyptal resins. These are present in adhesives, paints, and serve as binding agents in composite materials.

In this investigation, you will produce some linear and cross-linked polymers. You will produce a linear polyamide called nylon-6,6. You will then produce some temporarily cross-linked polymers called slime and gluep.

Another common class of polymers are called polyesters. Examples of linear polyesters with which you may be familiar include Dacron (fabric), Mylar (transparent sheet plastic) and Lexan (bullet-proof glass). Interestingly, the linear polyester (polyethylene terephthalate - PETE) used to make Dacron fabric is also used to make 2-liter soft drink containers. You may already know that some companies make polyester fleece jackets from recycled 2-L type soft drink bottles.

Because we use so much plastic, recycling is an excellent way to conserve resources. A code for plastics was established based on the most frequently used plastics (see chart below).

No. Name Abbreviation1 Polyethylene terephthalate (PET)2 High-density polyethylene (HDPE)3 Polyvinyl chloride (PVC)4 Low-density polyethylene (LDPE)5 Polypropylene (PP)6 Polystyrene (PS)7 Other (composites and mixtures)

You can explore a lot more interesting concepts about polymers by visiting the Macrogalleria at http://pslc.ws/macrog/level4.htm. It’s very cool!

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Experiment 1: Nylon: A Linear Polyamide

Nylon-6,6 is made from two monomers each of which contain 6 carbon atoms - hence its name. One of the monomers is a 6 carbon acid with a carboxylic acid group (COOH), at each end - hexanedioic acid.

HOOCCH2CH2CH2CH2COOHThe other monomer is a 6 carbon chain with an amino group (-NH2), at each end. This is 1,6diaminohexane (also known as hexane-1,6-diamine).

H2NCH2CH2CH2CH2CH2CH2NH2

When these two compounds polymerize, the amine and acid groups combine, each time with the loss of a molecule of water. This is known as condensation polymerization. Condensation polymerization is the formation of a polymer involving the loss of a small molecule. In this case, the molecule is water, but in other cases different small molecules might be lost.The diagram shows the loss of water between two of the monomers:

monomer A monomer B

This keeps on happening, and so you get a chain which looks like this:

-------- A ---------------- B ----------------A ---------------- B --------In the lab, it is easier to make nylon-6,6 at room temperature using an acyl chloride (acid chloride) rather than an acid. The 1,6-diaminohexane is used just as before, but hexanedioyl dichloride (adipoyl chloride) is used instead of hexanedioic acid.

If you compare the next diagram with the diagram further up the page for the formation of nylon6,6, you will see that the only difference is that molecules of HCl are lost rather than molecules of water.

Safety Note: This reaction uses organic solutions that should not be touched and will produce HCl. Be

cautious!

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1. Wear gloves and goggles!

2. Add 2.0 mL of 20% NaOH solution to a 100/150 mL beaker and then add 10.0 mL of 4% 1,6-hexanediamine solution.

3. Add a few drops of phenolphthalein.

4. Slowly add (DO NOT MIX!!!) 10 mL of 3% (v/v) adipoyl chloride dissolved in cyclohexane on top of the solution in the beaker.

5. Use a cotton swab to slowly lift and pull the film from the layer between the two liquids. Wind it around 2 small wooden sticks, a beaker or some other cylindrical container. See how long of a fiber you can make by continuously pulling the “rope.”

6. Cut the rope and transfer it to a beaker filled with water. Observe what happens to the rope.

7. Remove the rope from the water and dry it between 2 pieces of paper towel.

8. IN THE HOOD, vigorously mix the contents of the beaker used to make your nylon rope. Record your observations of this vigorous reaction.

9. Pour the mixture into a cold water bath and wash it. Dry the nylon between 2 pieces of paper towel as before.

10. All liquid waste must be disposed of in the appropriate waste containers. The nylon can be thrown in the trash or washed in distilled water and taken home.

Critical Thinking Questions (answer in a lab notebook) CTQ 1. What happens to the thickness of the rope when it is put into water? Explain this observation using the

structure of nylon-6,6 and water.

CTQ 2. If white smoke forms when mixing the solutions in Step 8, what is it? (Hint: It's the reason why you did this part of the reaction in the hood and the reason we added NaOH at the beginning of this experiment.)

Experiment 2: A Non-Newtonian Fluid

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Cornstarch (amylose) and water can be considered a colloidal suspension. A colloidal suspension is a two-phase system in which the starch and water are not dissolved but simply mixed into a permanent suspension that will not settle on standing. Other examples of colloids are blood, fog, whipped cream, foams, Jell-O®, and styling gel.

Newtonian and Non-Newtonian Fluids: We use the term “viscosity” to describe the resistance of a liquid to flow. Water, which has a low viscosity, flows easily. Honey, at room temperature, has a higher viscosity and flows more slowly than water. But if you warm honey up, its viscosity drops, and it flows more easily. Most fluids behave like water and honey, in that their viscosity depends only on temperature. We call such fluids “Newtonian,” since their behavior was first described by Isaac Newton. “Non-Newtonian” fluids also have a viscosity that depends on temperature, but their viscosity also depends on the force applied to the liquid or how fast an object is moving through the liquid.

Examples of non-Newtonian fluids include ketchup, silly putty, and quicksand. If you struggle to escape quicksand, you apply pressure to it and it becomes hard, making it more difficult to escape. The recommended way to escape quicksand is to slowly move toward solid ground; you might also lie down on it, thus distributing your weight over a wider area and reducing the pressure. Ketchup exhibits the opposite behavior, and is called a thixotropic fluid: its viscosity decreases under pressure. That’s why shaking a bottle of ketchup makes it easier to pour.

Shear Thickening (Dilatant) Fluids: How can the behavior of quicksand be explained? Think of a busy sidewalk. The easiest way to get through a crowd of people is to move slowly and find a path between people. The people will have time to move out of the way. If you just took a running start and headed straight for the crowd of people, you would quickly slam into someone and you wouldn't get very far. They don’t have time to “react” to your motion.

Perhaps a little more technically correct explanation is to think of a mixture of a polymer and water (cornstarch, which is amylose, is a sugar polymer). If you start with dry polymer chains and slowly add water, at some point you will reach the critical concentration where there are no voids [the polymer molecules are in contact with each other and the water fills the interstitial spaces]. At this point, it cannot be said if you have a 'solid' or a 'liquid.' Remember, polymers are HUGE molecules…so huge you might consider a polymer molecules as particulate themselves.

The application of pressure (shear stress = motion, stress, stirring) disturbs this mixture and the flow characteristics change -- the water is 'squeezed out' (into adjacent areas) leaving the molecules behind to interact. So, when you apply high shear forces (e.g. the rapid application of pressure or force) to the mixture, the polymer chains do not have time to quickly get out of the way, and they get tangled up in each other, effectively thickening the material. If you apply low shear forces (move through the material slowly), then the polymer particles have time to slip past each other like a fluid.

1. Place a few paper towels on the lab bench. 2. Put mixing bowl or pie tin in the middle of the paper towels.

3. Add about 60 cm3 (~1/2 cup) of dry cornstarch to the bowl.

4. Slowly add about 30 mL of water to the corn starch and stir slowly with your hands. Add water slowly to the mixture, with stirring, until all of the powder is wet.

5. Continue to add water until the cornstarch acts like a liquid when you stir it slowly. When you tap on the liquid with your finger, it shouldn't splash, but rather will become hard. If your mixture is too liquid, add

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more cornstarch. Your goal is to create a mixture that feels like a stiff liquid when you stir it slowly, but feels like a solid when you tap on it with your finger or a spoon.

6. Scoop the cornstarch mixture into the palm of your hand, and then slowly work it into a ball. Record your observations.

7. Can you think of other tests you can do with it? Your instructor may provide a hammer to slam into the mixture. Ask to get one, if not on the cart.

8. Observe cool videos of this mixture at the following web addresses! http://www.youtube.com/watch?v=f2XQ97XHjVw http://www.youtube.com/watch?v=WTCkVh9CWT8 Do a search on YouTube

Critical Thinking Questions (answer in a lab notebook)

CTQ 3. Is the cornstarch/water mixture Newtonian or Non-Newtonian? Explain how you know this.CTQ 4. Is the cornstarch/water mixture a dilatant or thixotropic fluid? Explain how you know this.CTQ 5. These types of fluids are being researched by the military as body armor. Explain why this type of

material would make good body armor.

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Experiment 3: Getting Slimed! (An adventure in cross-linked polymers)

Slime is really a polymer of vinyl alcohol monomers polymerized together to form a polyvinyl alcohol (PVA) polymer. Then, the PVA is cross-linked with tetrahydroxyborate, also known as Borax.

The B(OH)4

– ion is believed to cross-link the polymer chains as shown below. Note that the cross-linked bond occurs because of hydrogen bonding between the OH groups on the polymer chain and the OH of the borate ion.

PVA polymer + Borax cross-linked Slime!

Prepare Stock Slime Solution

1. Add 60 mL of a 4% polyvinyl alcohol (PVA) solution to a beaker.

2. Make observations about the appearance of the solution, its thickness, and any other pertinent information.

3. Add 2-3 drops of food coloring to the beaker and a small scoop of phosphorescent glow powder to produce a fluorescent slime that glows under black light.

4. Add approximately 8-12 mL of borax solution to the beaker. (More borax yields thicker slime, less borax yields more watery slime.)

5. Stir the mixture with your glass stirring rod. Observe the mixture.

6. If some of the liquid is not gelling, add more borax solution.

7. Give your slime a name! It will be your little created buddy for the lab.

The Characteristics of Slime

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Once a gel is formed, remove the material from the beaker and examine its physical characteristics (the slime is non-toxic; you can touch it with your hands. In fact, you are encouraged to do so).

I. Mechanical Properties

Perform each of the following tests, and record your observations for each in a lab notebook.

1. Suspend the SLIME mass from your hand. 2. Place the SLIME mass on a flat, unbounded surface (for instance a table top) and leave it at rest for a few

minutes.3. Stick the SLIME mass on a glass window. 4. Place the SLIME mass in a beaker.5. Place the SLIME mass on a table top and squeeze it gently.6. Roll the SLIME mass into a ball and drop it on the floor.7. Slowly pull the SLIME mass apart.8. Place the SLIME mass on a table top and hit it with your hand. 9. (Ask your instructor before doing this test) Throw a hunk of your slime as hard as you can against an

uncluttered wall or the door. Observe what happens immediately upon impact, and then after a few moments when the slime has had time to “rest.”

10. Clean all residue of slime from the wall/door and floor.11. Sharply pull the SLIME mass apart. 12. Transfer the SLIME mass to a 500-mL beaker. With the help of friend, keep the beaker about 60 cm above

the level of the table top and tilt it slightly to allow the fluid to slowly fall from its lip. Stick a small dark piece of paper on the falling stream close to the lip of the beaker in order to follow any change in the direction of the flow. Cut the falling stream with a pair of scissors, coated with Vaseline, a few centimeters below the lip of the beaker. What do you observe?

13. Obtain a 20-mL syringe. Pull the plunger out the back of the syringe, and fill the syringe with your stock SLIME. Re-insert the plunger. A. Very slowly push the plunger until a small droplet of SLIME forms at the end. Continue to push the

plunger slowly. Observe what happens to the diameter of the SLIME stream as it exits the plunger. This is called die swell and is characteristic of dilatant (dilating) fluids.

B. Very quickly and forcefully squeeze out a line of SLIME onto the benchtop using your syringe. Observe the opacity and form of the SLIME immediately after it exits the syringe and then later after it relaxes.

14. First fill a large beaker about ½ full of water. Take a glass stirring rod and either spin it in the water or stir the water. You should notice the usual vortex form in the middle. Next, take a glass stirring rod and insert it about halfway down into your stock SLIME solution. Spin the rod between your hands and watch what happens to the SLIME. If you spin the rod too fast you will break the slime. Try to find just the right speed to get the SLIME to climb the rod.

15. Stop stirring. What do you notice?


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CTQ 6. How is the SLIME polymer different from the poly-vinyl alcohol polymer solution used to make the SLIME?

CTQ 7. If more borax is added to the SLIME, how do the characteristics of the resulting SLIME change? Explain these observations using the concept of crosslinking and the 3-D structure of the cross-linked SLIME polymer.

CTQ 8. Claim: Based on your experiments, is SLIME a Newtonian or Non-Newtonian fluid?

CTQ 9. Evidence: Provide evidence for your claim by referencing actual experiments you did.

CTQ 10. A gel can be considered a material that exhibits both solid and liquid properties. How is the SLIME gel like a liquid?

CTQ 11. How is the SLIME gel like a solid?

CTQ 12. How does the hydrogen bonding in the cross-linking explain the gel’s properties (that is, how it can be easily squeezed, dripped, stretched, etc.)? How can the SLIME so effectively trap water?

II. Chemical Property

1. On a watch glass, put 2 drops each of 3 M H2SO4 and 6 M NaOH (do NOT allow them to mix). To each drop add 1 drop of methyl orange indicator. Record the color of methyl orange in acid and base solutions. (The pKa of methyl orange is 3.7. Recall at which pH value the color of methyl orange will change!)

2. Add 10 mL of fresh polyvinyl alcohol solution to a 100-mL beaker.3. Add 2-3 drops of methyl orange to the PVA solution.4. Add about 2-3 mL of borax solution and stir to make SLIME. Record the color of the indicator. Is your

SLIME pH above or below 3.7?5. Using a dropper, slowly add with stirring 3M H2SO4. Stir after each drop added and record the color of the

methyl orange indicator.6. Continue to add the sulfuric acid until you have “neutralized” the borax in SLIME. What do you notice has

happened to your SLIME?7. Reverse the process by adding, with stirring, 6M NaOH, recording the color of the indicator and the

appearance of the SLIME after each addition. Continue to add NaOH until you reconstitute the slime.

The destruction of the crosslinks in SLIME can be represented by the reaction below:

8. Take a sample of your stock slime polymer and put into a small beaker. Pour a generous amount of table salt, NaCl, into the slime and mix well. What do you notice happens?

Experiment 4: Gluep! (Slime Part II – The Son of Slime!)

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Warning: Gluep is super sticky, so keep all regular paper away from it. Wax paper is okay. It’s fine to touch with your hands to have on the lab bench.

1. Put about 40 mL of a 50% (v/v) polyvinyl acetate (white glue) solution in water into a beaker. Add a few drops of food coloring if desired.

2. Record the appearance and consistency of the product.

3. Add about 10-14 mL of borax. Stir with your hands or a stir rod.

4. Describe how the mixture changes as you stir.

5. Scrape the gluep onto a piece of wax paper or the lab bench. Discard any excess liquid. Knead the gluep with your hands until it is no longer wet and you can roll it into a smooth ball. Describe how it feels.

6. Draw a simple image on wax paper using a water-soluble felt tip marker (overhead pen). Roll the gluep into a ball then press it onto the image. Peel the gluep off carefully. Describe what happens.

7. Repeat the mechanical tests 1-11 (or more if you like). Be sure to clean up your mess.

8. Make more gluep following the instructions, but this time vary the amount of borax solution used. Repeat the above tests. How do these glueps compare to the original?

9. Your instructor may provide you with other types of glue to use.

10. Coming un-glueped: Can gluep be un-crosslinked in the same way the SLIME can? Try the titration procedure with sulfuric acid and sodium hydroxide the same way you did it for SLIME.


CTQ 13. Why is polyvinyl acetate (white glue) so viscous?

CTQ 14. Compare the properties of the cross-linked polyvinyl acetate (white glue)-based polymer with the cross-linked polyvinyl alcohol polymer (slime).

CTQ 15. Claim: Based on your experiments, is gluep a Newtonian or Non-Newtonian fluid?

CTQ 16. Evidence: Provide evidence for your claim by referencing actual experiments you did.

CTQ 17. The repeat unit for gluep, polyvinyl acetate (white glue) is shown at right. In a lab notebook, draw what a “vinyl acetate” monomer would look like.

CTQ 18. In a lab notebook, draw a representation of the polymer chains before and after crosslinking with borax. (Be sure to include some water molecules in your drawing…after all, this material is 96% water!)

Experiment 5. I'm Melting! (Investigating the solubility of polymers in solvents)

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Investigate the solubility of polystyrene peanuts or packing foam in water, acetone, and ethanol. Which of these solvents work well to melt or dissolve them? For the solvent that works the best, dissolve part of a polystyrene in the solvent and mold the "lump" into an interesting shape and allow it to dry on the lab bench. Describe the characteristics of the polystyrene through its transformation from a packing peanut or foam into an interesting shape. Why do you think the polymer peanut/foam was made with air trapped inside initially?

Experiment 6: Polyurethane – a thermosetting polymer

1. Place a few paper towels on the lab bench, as monomer A and B stick to almost everything and is extremely hard to clean up.

2. Be sure to wear gloves for this experiment!3. Mix equal proportions of monomer A and monomer B of polyurethane into a paper cup (about 10-mL of

each). Add food coloring if you wish.4. Stir vigorously with a wooden stick for at least 1 minute. You can cover the cup with a rubber glove and

rubber band to create a shape or figure.5. Sit back and enjoy the results. [Note: Do not touch the polymer directly until it has hardened.]

Lab Report

1. Turn in copies of the lab notebook pages containing your data, observations, and CTQs which you completed while doing this experiment.

2. Congratulations! You’ve survived an entire year of General Chemistry!

3. Now get the heck outta here!

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