Laboratory Activity: Teacher Notes

Activity 1:Drug Identification

Major Chemical Concept

This is intended as an introductory activity involving the concept of dynamic equilibrium. The ideas developed in this activity are: (1) chemical reactions do not always go to completion, and (2) a state of dynamic equilibrium can be established in a chemical system by the simultaneous formation of products from reactants and reactants from products. Dynamic equilibrium does not necessarily (or usually) occur when equal concentrations of reactants and products exist in the chemical system.

Level

While the laboratory activity itself is quite easily performed, the underlying notions are not basic level concepts. The activity fits well in either a general or honors level course.

Expected Student Background

1. Familiarity with observational evidence for chemical reactions.

2. Knowledge of common ions and ionic representations of equations.

3. Experience with double replacement reactions.

4. Some knowledge of kinetic molecular theory.

Time

Completing the activity and answering questions requires one 50-min class period.

Safety

No safety precautions beyond those that apply to normal chemical laboratory work need to be stressed for this activity. Some students may be concerned with the superficial resemblance between SCN^­ and CN^­. Although SCN^­should be treated with the same respect as other chemicals, it does not have the level of toxicity of CN^­. Soluble chemicals should be washed down the drain with plenty of water. Solid chemicals should be disposed of in a solid waste jar.

Materials (For 24 students working in pairs)

60 Test-tubes, 18- x 150-mm

12 Test-tube racks

12 Beakers, 100-mL

12 Medicine droppers

0.2 M Iron(III) nitrate, Fe(NO₃ )₃, small dropping bottle

0.005 M Potassium thiocyanate, KSCN, 50 mL

Potassium thiocyanate, KSCN, solid, 5 g

0.1 M Silver nitrate, AgNO₃, small dropper bottle

Sodium fluoride, NaF, solid, 5 g

Advance Preparation

Fe(NO₃ )₃ solution: 8 g solid Fe(NO₃)₃. 9H₂ O in 100 mL solution. Carefully add 2-3 drops of concentrated nitric acid, HNO₃, to the solution. (NOTE: The acid stabilizes the solution and represses the formation of a yellow color. )

KSCN solution: 0.05 g solid KSCN in 100 mL solution. Include enough concentrated nitric acid (a few drops) to make the solution slightly acidic. (NOTE: This solution may not keep. Check the solution before using. )

AgNO₃solution: 1.7 g solid AgNO₃ in 100 mL solution. Small, centrally located containers of solid KSCN and NaF will be sufficient.

Pre-Laboratory Discussion

Students should be told that although the procedure is short and relatively simple, their detailed observations are essential in answering the questions and in understanding the concept being studied. No procedural instructions are needed except the location of the solutions and the solid KSCN.

Teacher-Student Interaction

While students are performing the activity, walk around the laboratory correcting errant procedures, and asking students for oral interpretations of their observations. This is not the time to explain the observations to students. Your questioning serves to focus studentsÕ thoughts on the system being observed. Assure students that there will be a class discussion of this laboratory activity after they have finished it.

Anticipated Student Results

The colorless KSCN solution and nearly colorless Fe(NO₃ )₃solution should form a brownish-red product. Adding either Fe(NO₃)₃ solution or KSCN crystals deepens the color of the solution. Adding AgNO₃ or NaF causes the color of the solution to fade.

Answers to Data Analysis and Concept Development

1. Students will write their observations in making the comparisons. General consensus should be obtained during the class discussion.

2. The active species are Fe³⁺ and SCN ^­ .

3. Fe³⁺ (aq) + SCN ^­ (aq) FeSCN²⁺ (aq) (NOTE: Students will probably not know to include the double arrow. This symbol can be introduced at the appropriate time in the post-laboratory discussion. )

4. The evidence for a chemical reaction between Fe(NO₃)₃and KSCN is the observed color change.

5. Some FeSCN²⁺ product forms when Fe(NO₃ )₃ solution is added to the KSCN solution. Although the Fe(NO₃)₃ solution may have been slightly colored, the deep color of the mixture of Fe(NO₃)₃ and KSCN solutions cannot be explained by dilution of Fe(NO₃)₃ solution with the colorless KSCN solution.

6. The color change is caused by the production of more FeSCN²⁺ . Formation of more FeSCN²⁺ indicates that SCN^­ was still available in the solution to react with Fe³⁺ from Fe(NO₃)₃.

7. The color changes when KSCN solid is added because SCN ­ from the KSCN reacts with Fe³⁺ still present in the system. (NOTE: The importance of the answers to Questions 5 and 6 is that neither reactantÑFe³⁺ or SCN ­ Ñhad been completely used up in forming product during the original mixing of the solutions. This key set of observations can serve as the focal point for introduction of the idea of dynamic equilibrium.)

8. Two observations will be noted. First, Ag + reacts with SCN ­ to form the precipitate AgSCN, decreasing the SCN ^­ concentration and shifting the equilibrium to the left. The second observation is the fading of the solution color as FeSCN 2+ concentration decreases. Focus class discussion on the fading color and its meaning. (NOTE: The specific role of Ag ⁺ can be demonstrated by adding enough AgNO ₃to make the solution colorless, filtering the solution, and testing separate samples of filtrate with Fe ³⁺ and SCN ^­. The color should return with SCN ^­, but not with Fe ³⁺, thus showing that Ag + removed SCN ^­.) Students can be led to this conclusion by questions, trial, and error.

9. As in Question 8, the color fades, indicating a decrease in FeSCN 2+ concentration. The evidence from the laboratory is that the F ^­ reacts with and removes the Fe ³⁺. This is the result of F ^­ forming a complex with Fe ³⁺to form FeF 6 3­ . Focus on the meaning of the observable evidence in the class discussion. (NOTE: Questions 8 and 9 can be used to develop the idea of reversibility of chemical reactions and the role of concentration in controlling reactions.)

Answers to Implications and Applications

1.

a. Shift the equilibrium toward products

b. Shift the equilibrium toward products

c. Shift the equilibrium toward reactants

d. Shift the equilibrium toward reactants

Post-Laboratory Activities

Part B of Laboratory Activity 2: Temperature Effects can be combined with this activity if there is enough time. If this activity is done, tell students to save any excess solution made in Procedure 2 to use in Laboratory Activity 2 .

After students have had time to prepare answers to the questions, lead a class discussion on the laboratory activity. In this discussion it is important to establish first that all student groups have compatible observations of the systems. Discuss any discrepant observations through questioning to clear up any problems. If students are adamant about observations that differ, do a quick demonstration to gain consensus.

Students may suggest that the reaction is a double displacement (aka., double replacement or metathesis) reaction between potassium thiocyanate and iron(III) nitrate, according to: 3KSCN + Fe(NO₃ )₃ Fe(SCN)₃+ 3 KNO₃ However, both products suggested by this equation would be soluble and ionic, resulting in identical ionic reactants and products (show this using a net ionic equation); thus such a prediction would not, by itself, account for the observations. If a precipitate had formed, there would be an argument in favor of the double displacement reaction. Of course, some sort of chemical reaction did occur as documented by the color change.

The formation of the complex ion FeSCN²⁺ may cause some difficulty for students. It could be explained that the Fe³⁺ ion attracts the SCN ­ ion strongly enough to hold them together as an ion. Actually, the chemical reaction is not the important focus of the activity. In fact, this laboratory activity could be carried out equally well by labeling the reactants as ÔAÕ and ÔBÕ and the colored product as ÔCÕ. All of the same logic can be used to develop the ideas of the activity.

A discussion of the chemistry of the reaction can be done by using the questions as a discussion outline. As stated earlier, answers to Questions 6 and 7 permit you to suggest to students that it is possible for a chemical reaction to occur so that the final system contains all reactants and products simultaneously.

Pose a question about whether it is reasonable to accept the idea that products (in this reaction the product is FeSCN ²⁺ ) could change chemically back to reactants, as well as reactants changing chemically to form products. If students agree to such a possibility, ask them to offer an explanation for such a phenomenon. When a satisfactory explanation has been obtained, you can insert the double arrow in the chemical equation and explain that it designates to chemists that the system is reversible, that is, that it contains all reacting species.

Student laboratory observations can be used to establish the idea that this constantly changing system (on the molecular level) is still able to achieve constant concentrations of reactants and products (as observed by the constancy of the colors of the solutions).

The darkening of the solution when more of either reactant is added can be used (either now or later) as an example of LeChatelierÕs principle. You may decide that this is the best place to introduce the idea of LeChatelierÕs principle, or it might be better to address this principle after completing Laboratory Activity 2 (see Extensions).

Extensions

1. Temperature effects on equilibrium point of a system (see Laboratory Activity 2).

2. Pressure effects on the equilibrium point of a gaseous system.

Assessing Laboratory Learning Questions

1. Sodium chloride, NaCl, has a solubility of approximately 36 g/100 mL water. If 45 g of NaCl were added to 100 mL of water, 36 g would dissolve as Na ⁺and Cl ^­ ions, but 9 g of solid NaCl would remain. Explain how this system could involve a dynamic equilibrium between solid NaCl and dissolved ions. [Because the water molecules and dissolved ions are in constant motion, Na⁺ and Cl ^­ ons could deposit on the NaCl solid, but for every pair of ions depositing on the solid, another pair of ions would go from the solid to the dissolved state.]

2. What effect would adding 5 mL of KNO₃ solution to your ÒstandardÓ test-tube (see Step 3) have on the test-tube solutionÕs color? Note that a water solution of KNO₃ is colorless. [The color of the ÒstandardÓ solution would be less intense, due to dilution, but added KNO₃ would not cause a chemical change. It contains K⁺ and NO ₃ ^­ ions that were ÒspectatorsÓ in the reaction.]

3. Explain why an open beaker with zinc metal, Zn, and hydrochloric acid, HCl, will not establish chemical equilibrium. The products of the reaction of zinc metal and hydrochloric acid are hydrogen gas, H₂ , and zinc chloride, ZnCl ₂, solution. [The system will not achieve equilibrium because a productÑ hydrogen gasÑleaves the system. In order for dynamic equilibrium to be established, all products and reactants must remain in the system.]

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