Spectrophotometric Determination of an Equilibrium … Manual Project/223 Part 1... · Web viewCI3...

CI1 Kinetic Factors-I Worksheet

ChemInquiry1 Name__________________________

Kinetic Factors - IWhat 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 Analysis1. What is the total volume (drops) in each of the 3 runs?

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

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

4. 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)?

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

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 6. What is the total volume (drops) in each of the 3 runs?

7. 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.

8. 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 5

2C 15 5 15 10

3C 15 0 15 15


10. 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 Analysis12. What is the total volume in each of the runs?

13. 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.


16. 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 ReactionsPre-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.

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Reference


RECORD your observations in the box above

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.

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

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

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


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!

[H3O+] =

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Pre-lab Score: ____________/5


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 8.2In−: 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:

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HIn + H2O In− + H3O+

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.

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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 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.

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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!)

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.

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pH vs. log[ In- ][HIn]


-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

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!

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log[ In- ][HIn]

y-intercept = −log Kc

Prep

are

and

Shar

e Ph

enol

phth

alei

n-Bu

ffer

Solu

tions


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.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.

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Prep

are

Blan

k an

d UN

KNOW

N So

lutio

n


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.

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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Take Note!!


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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CI4 What Factors Affect the Solubility of Ions?


What Factors Affect the Solubility of Ions?

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


I. Background Preparation Read this experiment thoughtfully FIRST:


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


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.

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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?

Record all data and observations in blue or black ink pen

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.

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http://hbcpnetbase.com/


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 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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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

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