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AP Chemistry Applying Green Chemistry to Purification

Introduction One can characterize a major practice of chemists as harvesting materials from nature, separating the individual substances within them, and putting the pure substances in bottles so that they can make new substances out of them. The practice of designing chemicals is fundamental to all of the sciences that build on knowledge of chemistry. Pharmacists and medical biochemists create, purify, and test the safety of medicines. Analytical chemists use and design instruments to detect the presence and properties of individual substances. Bioorganic chemists design new chemicals for specific purposes and invent and compare the synthetic processes to make them. Chemical engineers manage chemical process industries to produce useful products on larger scales and minimize harm. But even for those students who don't pursue a career in science, understanding some of the basic principles of chemistry is relevant to everyone’s life. Students can use ways of thinking as a chemist to make educated decisions. They will use chemistry principles as they think about the safety of products they buy and which ones are less likely to cause harm, energy reduction in their homes, which containers are safest for cooking or storing food, and how to dispose of all sorts of things in the home and office, such as cleaning agents, unused medicines, and batteries.

This experiment involves a major practice of chemists: separating substances in a mixture by taking advantage of properties of the substances that are unique to each one. In this case, students will rely on the substances’ chemical reactivity upon heating as the property that differs between them. From antiquity, two very important substances in society have been obtained from a salt mixture called natron. Natron has been harvested for thousands of years from dry lake beds. In ancient Egypt, and still today, natron is blended with oil and used as soap. Natron primarily consists of two substances, sodium carbonate (Na2CO3) and sodium bicarbonate (NaHCO3). Each of these substances, when separated, also has important uses. Sodium carbonate is used in the manufacture of glass, as a water softener when doing laundry, as an additive in community swimming pools to raise pH, and as an additive in foods. Sodium bicarbonate has many uses, ranging from cooking and medical uses, to cleaning, pesticide, and fire extinguishing uses.

Every process that a chemist uses involves some kind of transformation and has an efficiency of that transformation associated with it. There are many ways to measure and report efficiency — two of these are discussed here. For the first, the efficiency of a chemical process can be interpreted in terms of its financial cost or in terms of its benefit or detriment to the environment. When considering financial cost, one might compare the amount of the desired product that is actually produced and the amount that could be produced in the ideal situation where all of the starting materials are used up and converted to useful products. Percent yield is such a calculation, as the ratio of the two values.

100xYieldlTheoretica

YieldActualYieldPercent =

The number is usually reported as a percentage (out of 100%) instead of as a decimal, in order to make comparison to the ideal situation easier. Percent yield is a way of comparing actual runs in the

1 Applying Green Chemistry to Purification

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AP Chemistry laboratory of a chemical process, when striving for a maximum yield. If the percent yield is closer to 100%, then maximum financial value is derived because starting materials (which might be expensive) are maximally used, and the end product is maximally produced. In a second way of measuring and reporting efficiency of a process, considering environmental health, one might instead try to minimize undesired products such as wastes or byproducts that are an inevitable outcome of a chemical process. Atom economy is such a calculation, comparing the theoretical yield of the desired product and the total theoretical yield of all products as a ratio.

100xproducedproductsallofMassproductdesiredofMassEconomyAtomPercent =

Various different chemical processes have different atom economies, because the desired product can range from being the only substance produced to being only one of several substances that result from the reaction. If the atom economy is closer to 100%, then the amount of desired product is maximized. Thus, different chemical processes can be compared to each other in terms of which one is more efficient in resulting in the desired product. The process with the largest atom economy is the one that most cleanly produces the substance that is the goal of the process. In modern society, where we now recognize that disposing of unwanted byproducts is expensive and potentially harmful, this method of comparing chemical processes provides a very useful tool in making decisions. This experiment will work with baking soda (sodium bicarbonate). Baking soda decomposes according to the reaction shown below: Sodium bicarbonate (s) Æ sodium carbonate (s) + carbon dioxide (g) + water (g) Pre-Lab Review the 12 Principles of Green Chemistry at http://www.epa.gov/sciencematters/june2011/principles.htm

1. Select one of the principles of green chemistry. Explain in your own words how this principle shifts chemistry toward more environmentally conscious practices. How is this relevant to considering the benefits and risks when making decisions about which of two (or more) possible chemical processes is better?

2. The following two reactions are possible methods for refining copper in the final step of a smelting process, i.e. getting pure copper (Cu) from copper ores found in rocks. Calculate the theoretical atom economy for each reaction.

a. 2CuO(s) + C(s) Æ 2Cu(s) + CO2(g) b. CuO(s) + CO(g) Æ Cu + CO2(g)

3. Use your calculations from the previous question (i.e. 2a and 2b) to answer the following questions.

a. Which of the methods for refining copper ore is greener according to the atom economy principle of green chemistry?


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http://www.epa.gov/sciencematters/june2011/principles.htm

AP Chemistry b. Why is a calculation of atom economy helpful in comparing two chemical reactions to

determine which one is greener? In other words, what does atom economy tell youabout "greenness"?

c. What is another possible consideration from the principles of green chemistry thatcould tell you more about comparing the "greenness" of these two reactions?

Access the following simulation and answer the questions below to strengthen their understanding of percent yield http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/percenttutorial.htm

4. What is a substance? What is a mixture? How are they related?5. What are the general characteristics of substances chemists use to separate mixtures into

individual substances?

Materials • Baking soda• Balance, 0.01 g precision• Bunsen burner• Test Tube• Ring Stand• Universal clamp

• Spatula, micro• Spoon• Weighing paper• Sodium carbonate/Sodium bicarbonate

mixture

Safety • Do not heat covered test tubes.• Do not look directly into the end of a heating test tube.• Hot glass looks the same as cool glass. Be sure to wait for heated objects to cool and always use

tongs.

Procedure Part 1:

1. Measure the mass of an empty test tube.2. Measure approximately 2.00 g of baking soda on weigh paper. Record the exact measurement.3. Pour the baking soda into the test tube, making sure that no mass is left on the weigh paper.4. Assemble the ring stand with the universal clamp. Place the test tube containing the baking

soda in the clamp so that it is almost completely horizontal.5. Heat the test tube and its contents with the Bunsen burner for 10 minutes. Remove the heat

and allow the test tube to cool. Remember, hot glass looks the same as cool glass.6. When the test tube is cool, measure the mass of the test tube and its contents. From this, find

the mass of the solid product.


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http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/percenttutorial.htm

AP Chemistry 7. Place the solid product that remains after heating in the container labeled "Product Made from

Heating Samples".Mass of test tube

Mass of sodium bicarbonate added Mass of test tube and sodium bicarbonate

Mass of test tube and final product Mass of final product

Part 2: Inquiry You will be given a mixture containing both sodium bicarbonate (baking soda) and sodium carbonate. Design a procedure to determine the relative amounts of NaHCO3 and Na2CO3 in this mixture using the principles of green chemistry and stoichiometry. You must also prepare any necessary data tables that you will need for your procedure. Submit your procedure to me for approval at least 2 days before the lab date.

Disposal • Place the solid product that remains after heating in the container labeled "Product Made from

Heating Samples".• Place any mixture that was taken but not heated in the container labeled "Unused Sample" for

the respective part.

Post Lab Analysis Part 1:

1. Determine the theoretical yield of solid product.2. Determine the percent yield of your solid product.3. If the baking soda is not heated long enough for complete decomposition, what effect will this

have on your calculated percent yield? Explain.Part 2:

1. Calculate the percentage of sodium bicarbonate (NaHCO3) in the mixture.2. Epsom salts is a strong laxative used by veterinarians to treat animals, and soaking swollen feet

or skin with a rash in a warm solution of Epsom salts is also sometimes prescribed by doctors torelieve swelling or itchiness. Epsom salts is a hydrate, which means that a specific ratio of watermolecules per unit of salt formula regular repeats in the crystalline structure of Epsom salts.The formula for Epsom salts can be written as MgSO4·xH2O, where the ratio of water moleculesper unit of salt formula is x:1. When hydrated Epsom salts is heated to at least 250°C, all of thewaters of hydration are lost, according to the reaction:

a. When 3.648 g of Epsom salts were heated to constant mass of 250°C, 1.782 g of MgSO4

powder remained. What is the value of x?b. If anhydrous MgSO4 is the desired product, what is the atom economy of this reaction?


4

AP Chemistry What is the Relationship Between the Concentration of a Solution and the Amount of Transmitted

Light Through the Solution?

Spectroscopy is the study of the interaction between matter and light. A spectrophotometer or colorimeter can both be used to measure the amount of a particular wavelength of light absorbed by a substance. The amount of light absorbed is dependent on multiple factors according to Beer's Law. This law can be expressed by the equation abcA = where A is the measured absorbance, a is the molar absorptivity constant, b is the path length, and c is the molar concentration of the substance.

Using a colorimeter, a beam of light of a particular wavelength (equates to color) is sent through a colored sample. A detector on the other side of the solution measures the amount of that wavelength that has passed through the solution. Any light that has not passed through is considered absorbed by the solution. As a general rule, the darker a solution, the more light it will absorb.

Spectroscopy experiments are most valuable when the wavelength (color) of light used is absorbed at a maximum level. A solution appears to be the color of light that it reflects. Thus, a blue light sent through a blue solution will not be absorbed. To establish the best wavelength of light for a particular solution, absorbance values are collected for a sample of the solution at incremental wavelengths. This data is graphed, and the wavelength that produces a maximum absorbance (λmax) is then used for the remainder of the experiment. One such plot is shown below.

Figure 1 Finding Maximum Absorbance Figure 2 Blue Dye #1

Blue dye #1 is found in a variety of food products including Gatorade. The blue dye molecule is shown above. The concentration of this dye affects the intensity of the color of the Gatorade. The maximum absorbance for the blue dye in this experiment occurs at 630 nm. The colorimeter in this experiment should be set at the closest wavelength to this value.

1 Relationship Between Concentration and Transmitted Light

5

AP Chemistry Purpose

• Determine the concentration of blue dye #1 in Gatorade using spectroscopy.• Determine the mass of blue dye #1 in Gatorade.

Safety Precautions

• Be careful when handling glassware. Alert your teacher in the event of broken glass.• Clean up all spills.

Materials

• LabQuest• Colorimeter• Cuvettes• Stock solution of liquid blue dye

concentration 6.4μM (6.4x10- M).• Blue colored Gatorade

• Graduated cylinder (10 mL)• Disposable pipettes• Distilled water• 50 mL beaker• Stirrer

Pre-Lab 1. Read through the procedure. Create a data table with appropriately labeled columns and rows

for the absorbance values of known concentration gathered in the experiment. This will beTable 1 referred to in the procedure.

2. Create a second data table for all other values recorded during the lab. This will be Table 2 inthe procedure.

3. What wavelength should you be sure to set your colorimeter nearest to? Explain why.4. Using a graphing utility such as Microsoft Excel, graph the following concentration and

absorbance values as a scatterplot. Use a linear trendline to find the "line of best fit". Title andlabel your graph well. Make sure your trendline and equation are shown on the graph and printit out.

Concentration (M) Absorbance 6.40x10-6 0.84 5.12 x10-6 0.67 3.84 x10-6 0.57 2.56 x10-6 0.39 1.28 x10-6 0.21

5. Use your trendline from question #4 to determine the concentration of a blue solution that hasa recorded absorbance value of 0.71.

Procedure 1. Obtain approximately 25 mL of the stock solution (6.4 μM) of blue dye #1 in a small beaker. In

another small beaker, obtain approximately 25 mL of distilled water.2. Fill a cuvette with distilled water. Calibrate your colorimeter to zero absorbance using the

water. Make sure your LabQuest is recording Absorbance.


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AP Chemistry 3. Using a pipette, fill another cuvette with the stock solution of blue dye solution. Record the

absorbance value of this solution in Table 1. Keep this cuvette, making sure you know exactlywhat is in it.

4. Using the same pipette, add 4 mL of stock solution to the graduated cylinder. Using a separatepipette, add 1 mL of distilled water to the graduated cylinder with the stock solution. Use astirrer to mix the solution.

5. Pour the 4:1 solution into a cuvette, clean the sides of the cuvette with a towel, and record itsabsorbance in Table 1. Keep this cuvette, keeping track of its contents.

6. Repeat steps 4-5 with the following ratios of stock solution to distilled water:a. 3 mL stock to 2 mL H2Ob. 2 mL stock to 3 mL H2Oc. 1 mL stock to 4 mL H2O

7. Obtain a small sample of the Gatorade. Record the volume of the Gatorade bottle. Fill acuvette and record the absorbance in Table 2.

8. Line up the cuvettes for solutions #1-5 in numerical order. Using your best visual judgment,place the cuvette of Gatorade in line according to its color relative to the known solutions.Make an estimation as to the concentration based on this visual assessment and record in Table2.

9. Dispose of all solutions down the drain. Rinse and dry all equipment and return.

Data

Data tables should be created as given in the Pre-Lab.

Analysis

1. Calculate the concentration of each of the known solutions using the dilution formula or apercentage value.

2. Using a graphing software (Microsoft Excel works), plot the data for concentration vs.absorbance. Using a linear regression (trendline), determine the "line of best fit" for your data.Print out your graph. Make sure it is well titled with axes labels.

3. Using your line of best fit and the absorbance value for the Gatorade sample, determine theconcentration of the Gatorade.

4. Calculate the percent error between your visual estimation of concentration of the Gatoradeand your calculated value.

5. Using your Gatorade concentration and the volume of the bottle it came in, determine thenumber of grams of blue dye #1 in the Gatorade bottle.


7

AP Chemistry

1 Mass Percent of Copper in Brass

How Can Color Be Used to Determine the Mass Percent of Copper in Brass?

Spectrophotometry is an extremely important tool used in forensic science to determine the detailed chemical composition of evidence obtained from a crime scene. It can be used to determine the concentration of either a single chemical species in solution or even the concentration of a species within a mixture of species in solution. For example, it can be used to determine the mass percent of copper in brass shell casings collected by the crime scene investigator (CSI), and then match the brass composition to a particular manufacturer.

Brass is a mixture of copper and zinc. Copper is oxidized by nitric acid according to the following reaction:

Cu(s) + NO3-(aq) Æ Cu2+(aq) + NO(g) in an acidic solution

A second reaction occurs when the colorless NO(g) reacts with oxygen in the air to form the observed brown-orange NO2(g) according to the equation:

2NO(g) + O2(g) Æ 2NO2(g) The Cu2+ ions in the unknown aqueous solution form the complex ion, [Cu(H2O)6]2+, which causes the blue color. This means that when "white" light (all wavelengths) passes through the solution, the dominant emerging color is blue. A spectrophotometer is used to analyze the color intensity of the copper (II) nitrate solution that forms. For this analysis, it will be necessary to determine which color, with a specific wavelength, will be most strongly absorbed by the copper ions. The concentration of the unknown brass solution will then be determined by comparing its absorbance with that of solutions having known concentrations of Cu(NO3)2.

Purpose

x Determine the mass percent of copper in a sample of brass.

Safety Precautions

x Concentrated nitric acid is corrosive and will attack and destroy metals, proteins, and mostplastics. Avoid skin contact and neutralize any spills with baking soda, then rinse with copiousamounts of water. The acid will discolor the skin for days after contact so be sure to weargloves.

x The NO gas that forms quickly oxidizes in air to produce a toxic, reddish-brown gas of NO2. Avoidinhalation.

x Perform the reaction under a fume hood. Be sure to wear splash-proof goggles and an apron atall times.

Materials

x Balancex Beakerx Concentrated nitric acid (15.8 M)x Distilled water

x 100 mL volumetric flaskx 1-2 g sample of brassx Remainder dependent on procedure

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AP Chemistry


Pre-Lab

1. Write the balanced equation for the reaction that occurs between copper metal and nitric acid.2. Read the given procedure. For step #3, you must design your own procedure to gather data to

create a calibration curve of known concentration and absorbance values of copper (II) nitrate.You will have the following materials available to you (you do not have to use all of them) in thelab in order to accomplish this so design your procedure with these items in mind:0.400 M Cu(NO3)2 stock solution 10 mL graduated cylinder LabQuest Distilled water Test tubes Colorimeter Paper Towels Test tube rack Cuvettes

Disposable Pipets Beakers

Use your knowledge of spectroscopy and the process used in Lab #2. This procedure must be submitted to your teacher for approval at least 48 hours prior to this lab. * You must include all directions for any use of the LabQuest technology.** The concentration of the copper in the unknown brass solution should be between 0 and 0.4 M. *** Your procedure must also include any and all necessary calculations that you used to determine amounts of chemicals.

3. Continue your designed procedure to include steps necessary to determine the concentration ofcopper (II) nitrate in the unknown brass solution (step #5).

4. Create any and all necessary data tables according to your designed procedure.

Procedure

1. Determine the mass of a 1-2 gram sample of brass to r0.001. g. Record this value. Place thesample in a small beaker. Label your beaker so that you can identify your group's materialsamong the rest of the class.

2. Assuming your brass sample is 100 percent copper by mass, calculate the minimum volume ofconcentrated 15.8 M HNO3(aq) that needs to be added to react completely with the brass.Under the fume hood, have your teacher add approximately 2mL more than this volume of 15.8M HNO3(aq) so that the acid is in excess, and then your teacher will cover the beaker with awatch glass.

3. While the metal is reacting with the acid, perform your predesigned procedure for gathering acalibration curve for the absorbance values of different known concentrations of copper (II)nitrate.

4. After the metal dissolves completely, add 50 mL of distilled water to the beaker. Then you willremove the beaker from the fume hood and transfer the solution to a 100 mL volumetric flask.Rinse the beaker 3-4 times with 5 mL of distilled water and add the washings to the flask. Diluteto a final volume of 100.0 mL. The excess nitric acid will dissolve the zinc and copper metals inthe brass.

5. Use the remainder of your predesigned procedure to determine the concentration of thecopper(II) nitrate in the unknown brass solution.

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AP Chemistry


6. The remaining brass solution should be neutralized by adding small amounts of solid bakingsoda until the bubbling has subsided and the pH is 5-9. Then dispose of the waste in the wastebeaker provided.

Data

Data tables should be created as given in the Pre-Lab.

Analysis

1. Use the data you collected to create a calibration curve using a graphing utility (Excel). Be sureto label the graph well and include any linear regressions used.

2. Determine the molarity of the Cu(NO3)2(aq) found in 100.0 mL of the brass solution using yourcolorimeter analysis. Support your answer with the appropriate calculations.

3. Determine the mass of Cu dissolved in the brass solution based upon the molarity calculated inStep 2, and use this value to calculate the mass percent of Cu in the brass sample.

4. If a student failed to rinse all of their reacted brass solution from the beaker to the volumetricflask, what effect would this have on the calculated mass percent of copper in the solution?Justify your answer.

10

AP Chemistry

Determination of Concentration by Oxidation-Reduction Titration of Hydrogen Peroxide

Introduction One method of determining the concentration of a hydrogen peroxide, H2O2, solution is by

titration with a solution of potassium permanganate, KMnO4, of known concentration. A titration, as you recall, is a convenient method of learning more about a solution by reacting it with a second solution of known molar concentration. There are a number of ways to measure the progress of a titration. One method used is called a potentiometric titration, in which the electric potential of a reaction is monitored. All acid-base titrations that are measured by a pH probe are potentiometric; thus, this method is not as unusual as it may seem. The reaction is an oxidation-reduction reaction and proceeds as shown below, in net ionic form.

5 H2O2 (aq) + 2 MnO4– (aq) + 6 H+ (aq) 5 O2 (g) + 2 Mn2+ (aq) + 8 H2O (l)

In this experiment, the volume of KMnO4 titrant used at the equivalence point will be used to determine the concentration of the H2O2 solution. Your sample of H2O2 will come from a bottle of ordinary, over-the-counter hydrogen peroxide purchased at a grocery or a drug store. The concentration of this product is labeled as 3% (by mass).

This experiment illustrates the electrical nature of chemical reactions, and offers practice with a process for observing and measuring an oxidation-reduction reaction.

Objectives In this experiment, you will

• Conduct the potentiometric titration of the reaction between commercially available hydrogenperoxide and potassium permanganate.

• Measure the potential change of the reaction.• Determine the concentration of the hydrogen peroxide solution.

Figure 1

Determination of Concentration by Oxidation-Reduction Titration of Hydrogen Peroxide 111

AP Chemistry

Pre-Lab Questions

1. Write the oxidation and reduction half-reactions for the redox reaction taking place in thislaboratory.

2. Identify the oxidizing agent for this reaction. Justify your answer.

3. How many moles of electrons are transferred in the balanced redox reaction? Justify youranswer.

4. Calculate the number of moles of MnO4 in 20.1 mL of 0.020 M potassium permanganatesolution.

5. Calculate the number of moles of hydrogen peroxide needed to react with 20.1 mL of 0.020 Mpotassium permanganate solution.

6. The introduction states that the hydrogen peroxide solution sold in the grocery store is 3% (bymass). Assuming a density of 1.0 g/mL, calculate the molarity and mole fraction of 3.0 % H2O2.

Materials Equipment

250 mL beaker 25 mL graduated cylinder 50 mL buret

Stirring rod or magnetic stirrer Utility clamp Buret clamp

Chemicals

Hydrogen peroxide solution, H2O2, 0.3% Potassium permanganate solution, KMnO4, 0.020 M in H2SO4 Distilled water __________________________________________________________________________________ Safety Precautions

Wear splash goggles and protective apron. Potassium permanganate is a powerful oxidizing agent; can explode on sudden heating; common cause of eye accidents; wear face protection. Strong skin irritant; slightly toxic by ingestion.Wear gloves. Hazard Code: B—Hazardous. Sulfuric acid: Severely corrosive to eyes, skin and other tissue; considerable heat of dilution with water; mixing with water may cause spraying and spattering. Extremely hazardous in contact with finely divided materials, carbides, chlorates, nitrates and other combustible materials. Hazard Code: A—Extremely hazardous. Hydrogen peroxide: 3% solution: Oxidizer and skin irritant. Many substances will cause hydrogen peroxide to decompose into water and oxygen gas. Hazard Code: C—Somewhat hazardous. __________________________________________________________________________________


AP Chemistry

Procedure 1. Measure out precisely 10.00 mL of a 0.3% hydrogen peroxide solution from the dispensing

buret your teacher has prepared into a 250 mL beaker. Should you miss the 10.00 mL mark, just record the volume of your sample exactly in the Data Table.

2. Measure out approximately 10 mL of 4.5 M sulfuric acid, H2SO4, solution from the dispensingburet your teacher has prepared into the beaker containing the 0.3% hydrogen peroxide solution. CAUTION: H2SO4 is a strong acid, and should be handled with care.

3. Use a graduated cylinder to measure out approximately 25 mL of deionized water and add it tothe beaker as well.

4. Place the beaker of H2O2 solution on a magnetic stirrer and add a stirring bar if available. If nomagnetic stirrer is available, stir the mixture with a stirring rod during the titration.

5. Set up a ring stand, a buret clamp, and a buret to conduct the titration (see Figure 1). Rinseand fill the 50 mL buret with 0.0200 M MnO4

– solution. CAUTION: Handle the KMnO4 solution with care; it has been mixed with H2SO4, which can cause painful burns if it comes in contact with the skin.

6. Gently stir the beaker of solution.7. Conduct the titration carefully. When the equivalence point is reached, the solution will

maintain a faint pink color for longer than 30 seconds. Use your first trial to determine theregion of the titration curve near the equivalence point, and not to precisely determine theequivalence point. (Do not record this first trial!)

8. When you have completed the titration, dispose of the reaction mixture as directed.9. Record the volume of MnO4- needed to reach the equivalence point in the Data Table for Trial

1.10. Repeat the necessary steps to conduct a second trial with a new sample of 0.3% H2O2 and acid

solution.11. When you conduct the second titration, carefully add the MnO4

– solution drop by drop in theregion near the equivalence point, so that you can precisely identify the equivalence point ofthe reaction.

12. Precisely determine the equivalence point of the reaction. Record this information in yourdata table for Trial 2.

13. At the direction of your instructor, conduct a third trial.

Data Tables

Trial 1 Trial 2 Trial 3

Volume of H2O2 solution (mL)

Volume of MnO4– solution used (mL)

Calculated number of moles of H2O2


AP Chemistry

Post-Lab Questions and Analysis 1. Calculate the moles of MnO4

– used to reach the equivalence point of the reaction for each trial.

2. Use the number of MnO4– moles to calculate the moles of H2O2 in the sample of solution for

each trial.

3. Calculate the molar concentration of the H2O2 solution for each trial.

4. Calculate the % Error between your experimentally determined molarities for each trial and themolarity value for 3% hydrogen peroxide that you calculated as your answer to Pre-Lab Question6. (Remember that you titrated 0.3% H2O2 solution samples and NOT 3% H2O2 solutionsamples.)

5. Why is hydrogen peroxide stored in a brown bottle?

6. A student fails to dry their 250 mL beaker between trials 2 and 3. What effect does this have onthe calculated molarity of hydrogen peroxide for trial 3?

7. A student titrates three samples of 0.3% H2O2 solution made from a bottle of 3% H2O2 solutiondated 10/31/2007. Would you expect the calculated molarities of the samples to be higher or lower than the value you calculated for your experimental samples? Justify your answer.


AP Chemistry

1 Electrochemical Cells & Electroplating

Measurements Using Electrochemical Cells and Electroplating

The basic counting unit in chemistry, the mole, has a special name, Avogadro’s number, in honor of the Italian scientist Amadeo Avogadro (1776-1856). The commonly accepted definition of Avogadro’s number is the number of atoms in exactly 12 g of the isotope 12C, and the quantity itself is 6.02214199(47) × 1023.1 A bit of information about Avogadro seems appropriate. His full name was Lorenzo Romano Amedeo Carlo Avogadro (almost a mole of letters in his name). He was a practicing lawyer until 1806 when he began his new career teaching physics and math at the University of Turin, where he was later promoted to the chair of physical chemistry. In 1811, Avogadro published a paper in the Journal de Physique, entitled “Essay on a Manner of Determining the Relative Masses of the Elementary Molecules of Bodies, and the Proportions in Which They Enter into These Compounds,” which pretty much says it all. This paper includes the statement that has come to be regarded as Avogadro’s Hypothesis:

The first hypothesis to present itself in this connection, and apparently even the only admissible one, is the supposition that the number of integral molecules in any gases is always the same for equal volumes, or always proportional to the volumes. Indeed, if we were to suppose that the number of molecules contained in a given volume were different for different gases, it would scarcely be possible to conceive that the law regulating the distance of molecules could give in all cases relations as simple as those which the facts just detailed compel us to acknowledge between the volumes and the number of molecules.

In this experiment, you will confirm Avogadro’s number by conducting an electrochemical process called electrolysis. In electrolysis, an external power supply is used to drive an otherwise nonspontaneous reaction. You will use a copper strip and a zinc strip as the electrodes, placed in a beaker of sulfuric acid. You will make the cell electrolytic by using the copper strip as the anode and the zinc strip as the cathode. By determining the average current used in the reaction, along with the knowledge that all of the copper ions formed are the 2+ cations, you will calculate the number of atoms in one mole of copper and compare this value with Avogadro’s number.

Figure 1

1 CRC Handbook of Chemistry and Physics, 82nd edition, p. 1-7.

Ammeter or

Current Probe

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AP Chemistry


OBJECTIVES In this experiment, you will

x Prepare an electrochemical cell to oxidize a copper electrode.x Measure the amount of copper that was deposited in the electroplating process and

determine the average current used.x Calculate a value for Avogadro’s number and compare it to the accepted value.x Calculate the value of a Faraday.

MATERIALS Ammeter 1 M H2SO4 solution 1.5 volt DC power source copper strip (anode) (+)* 4 connecting wires with alligator clips zinc strip (cathode) (−)* steel wool or sandpaper distilled water Analytical balance two 250 mL beakers

* Remember the acronym “EPA”—An Electrolytic (nonspontaneous) cell has a Positive Anode.

PROCEDURE 1. Obtain and wear goggles.

2. Use steel wool or sandpaper to clean a strip of copper, which will be the anode of theelectrochemical cell. Obtain a strip of zinc metal to use as the cathode, and clean it with steelwool if needed.

3. Use an analytical balance to measure the mass of the copper strip. Record the mass in yourdata table.

4. Fill a 250 mL beaker about ¾ full with 1 M H2SO4 solution. CAUTION: Handle the sulfuricacid with care. It can cause painful burns if it comes in contact with the skin.

5. Obtain a DC power supply and either an Ammeter or Current Probe. Use connecting wires,with alligator clips, to connect the DC power supply, ammeter or current probe, and the twometal electrodes to be used in the electrochemical cell as you see in Figure 1. Copper is theanode (+) terminal and zinc is the cathode (−) terminal in this cell because it is an electrolyticcell and the polarities of the electrodes are switched since it is a nonspontaneous process.Note: Do not place the electrodes in the cell until Step 7.

6. Set up the ammeter. You will manually collect data for 3 minutes.

7. Place the electrodes into the 1 M H2SO4 solution in the cell. Make sure that the electrodes areimmersed in the solution to equal depths and as far apart as possible.

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AP Chemistry


8. Turn on the DC power supply and check the current readings. The initial current should be inthe 0.4 – 0.6 amp range. If the current is not in this range, adjust the settings on the powersupply. Once the initial current is in range, turn off the power supply.

9. Begin the data collection and turn on the power source. Data will be collected for 3 minutes.Observe the reaction carefully. Note: Be ready to shut off the power as soon as the datacollection stops.

10. When the data collection is complete, turn off the power supply and carefully remove theelectrodes from the H2SO4 solution. Carefully rinse the copper electrode with distilled water.Dry the copper electrode very carefully.

11. Measure and record the mass of the dry copper electrode.

12. Analyze the graph of your data to determine the average current that was applied during theelectrolysis. Record the mean in your data table as the average current for Trial 1.

13. Repeat the necessary steps to conduct a second trial.

DATA TABLE

Trial 1 Trial 2

Initial mass of copper electrode (g)

Final mass of copper electrode (g)

Average current (A)

Time of current application (s)

PRE-LAB QUESTIONS 1. What is the sign of E°cell for the reaction taking place in this laboratory exercise? Justify

your answer.

2. What is the sign of ΔG° for the reaction taking place in this laboratory exercise?Mathematically justify your answer.

3. In this experiment, which electrode is losing mass? Identify whether this electrode is theanode or the cathode. Justify your answer by writing the half-reaction that occurs at theelectrode.

17

AP Chemistry


4. Answer the following questions about electrolytic cells.(a) Which redox process occurs at the anode? (b) What is the polarity of the anode? (c) Which direction do electrons flow? (d) In which direction do cations migrate?

5. A student conducted an experiment to determine Avogadro’s number using a zinc cathodeand copper anode having an initial mass of 5.100 g. The student applied a current of 0.468amps for a total of three minutes. After data collection was complete, the copper anodewas dried and massed. The final mass of the copper anode was 5.073 g.

(a) Calculate the number of grams lost by the copper anode. (b) Calculate the total quantity of charge in coulombs that passed through the

electrolytic cell. (c) Calculate the number of electrons used to electroplate the zinc. The charge of a

single electron is 1.602 × 10−19coulombs. (d) Calculate the number of copper atoms lost from the copper electrode. (e) Calculate the number of copper atoms per gram of copper lost at the anode. (f) Calculate the number of copper atoms in a mole of copper. (g) Calculate the student’s percent error. (h) Starting with your answer to part (b) above, show one dimensional analysis

expression that allows you to arrive at your answer to part (f).

6. A few seconds pass between the end of data collection and the time the power supply isturned off. What effect does this error have on the calculated value of Avogadro’snumber?

POST-LAB QUESTIONS AND DATA ANALYSIS 1. Calculate Avogadro’s number for each trial.2. Calculate the percent error for each trial3. Which measurement limited the number of significant figures that could be obtained?4. The Faraday is a unit of measure that relates the number of coulombs per mole of electrons.

Use your data to calculate the value of the Faraday for each trial. The accepted value of theFaraday is 96,500 C/ mol e−

.5. A student fails to dry the electrode completely after Trial 1 prior to massing it on the balance.

(a) What effect does this error have on the calculated value of the Faraday? (b) What effect does this error have on the calculated value for Avogadro’s number for Trial

1?

18

AP Chemistry

1 The Hand Warmer Design Challenge

The Hand Warmer Design Challenge: Where Does the Heat Come From?

Hand warmers are small packets that people put inside gloves or mittens on cold days to keep their fingers warm. They are very popular with people who work outside or do winter sports. One type of hand warmer contains water in one section of the packet and a soluble substance in another section. When the packet is squeezed, the water and the soluble substance are mixed, the solid dissolves, and the packet becomes warm. The ideal hand warmer increases in temperature by 20° C (but no more) as quickly as possible, has a volume of about 50 mL, costs as little as possible to make, and uses chemicals that are as safe and environmentally friendly as possible. You will carry out an experiment to determine which substances, in what amounts, to use in order to make a hand warmer that meets these criteria.

Breaking bonds and particulate attractions absorbs energy from the surroundings, while forming new bonds and particulate attractions releases energy to the surroundings. When an ionic solid dissolves in water, ionic bonds between cations and anions in the ionic solid and hydrogen bonds between water molecules are broken, and new attractions between water molecules and anions and water molecules and cations are formed. The amount of energy required to break these bonds and form new ones depends on the chemical properties of the particular anions and cations. Therefore, when some ionic solids dissolve, more energy is required to break the cation-anion bonds than is released in forming the new water-ion attractions, and the overall process absorbs energy in the form of heat. When other ionic compounds dissolve, the converse is true, and the bond making releases more energy than the bond breaking absorbs, and therefore the process overall releases heat. When heat is absorbed, the enthalpy change, q, is endothermic, and the enthalpy change is positive. When heat is released, the change is exothermic, and the value of q is negative. The entropy (disorder) change of solution formation is always positive, regardless of whether it is endothermic or exothermic, because solutions are much more disordered than are the pure solute and solvent from which they are made. This positive entropy change is thermodynamically favorable.

Purpose

x Determine the best substances to create a commercial hand warmer.

Safety Precautions

x The solids and resulting solutions in this lab are potential eye and skin irritants. Calcium chloridecan cause skin burns. Ammonium nitrate is a powerful oxidizer that must be kept away fromignition sources and is quite toxic on ingestion.

x Wear splash proof safety goggles and an apron.x If solutions are spilled on skin, wash with copious amounts of water.

Materials

x LabQuestx Temperature probex Balance

x Coffee cup w/ lidx Magnetic stirrer and stir barx 3 ionic solids (assigned by group)

19

AP Chemistry


Pre-Lab

1. Obtain the MSDS for your three solids. Review each one, making notes about safety concerns,necessary precautions, and disposal.

2. Rank the solids you are given from least to most expensive using the following chart:Substance 2012 Cost per 500 g ($)

NaCl 3.95 CaCl2 6.55

NaC2H3O2 12.90 Na2CO3 6.15

LiCl 32.75 NH4NO3 9.05

3. When sodium chloride is dissolved in water, the temperature of the resulting solution is lowerthan the temperature of the water before the salt dissolves. How can this result be explainedbased on the bond breaking and bond making that is occurring?

4. Summarize the qualities of an ideal hand warmer.

Procedure

LabQuest Notes:

Your LabQuest should be set to record temperatures every second for at least 180 seconds for each of the trials described.

Part 1: Calorimeter Calibration

1. Place 100.0 mL sample of water in a clean, dry 150 mL beaker. Heat with occasional stirring toapproximately 50 °C. Remove the beaker from the hot plate and place on the lab bench.

2. Meanwhile, place exactly 100.0 mL of cool water (approximately 20°C) in the clean drycalorimeter.

3. Measure the temperature of the hot water and the cold water and record in Table 1, thenimmediately pour the entire hot water sample into the calorimeter and quickly put on the cover.Put the temperature probe into the calorimeter through the hole in the cover (make sure theprobe is actually in the water). Record the temperature of the water for approximately 3minutes. [When the temperature readings are continually decreasing, you can stop recording.]Find the maximum temperature of the water and record this in Table 1.

4. If time allows or instructed, repeat this determination.5. Dry the calorimeter for further work.

Part 2: Investigation of Ionic Solids

6. Add 45.0 mL of room temperature water to the clean, dry calorimeter. Measure the initialtemperature of the water and record in Table 2.

7. Measure approximately 5.0 g of one of your ionic solids. Record the exact mass in Table 2.

20

AP Chemistry


8. Add the solid to the water in the calorimeter, quickly cover the solution and insert thetemperature probe. Record the temperature of the solution for 3 minutes or until thetemperature begins decreasing. Record the maximum temperature in Table 2.

9. Dispose of your solution by rinsing it down the drain with copious amounts of water.10. Repeat steps 6-9 for your other two ionic solids.11. Clean and dry your calorimeter. Return all equipment.

Data

Table 1: Calorimeter Calibration

Mass hot water: mhot (g)

Mass cold water: mcold (g)

Initial temp cold, Ticold (°C)

Initial temp hot, Tihot (°C)

Final temp of mixture, Tf (°C)

Table 2: Ionic Solid Investigation

Ionic Solid

Volume of water (mL)

Mass of solid (g)

Initial temperature of water (°C)

Maximum temperature of solution (°C)

∆T

Analysis

Part 1: Calorimeter Calibration

When hot and cold water are mixed, the hot water transfers some of its thermal energy to the cold water. The law of conservation of energy dictates that the amount of thermal energy lost (or the enthalpy change) by the hot water, qhot, is equal to the enthalpy change of the cold water, qcold, but opposite in sign, so qhot=-qcold. The enthalpy change for any substance is directly related to the mass of the substance, m; the specific heat capacity, c; and the temperature change, ∆T. The relationship is expressed mathematically in the equation q = mc∆T. The specific heat capacity of water is 4.184 J/g°C.

1. Calculate the enthalpy change of the water using the equation qhot = mhotc∆Thot. Assumethat the density of water is exactly 1 g/mL. Is this an endothermic or exothermic process?Explain.

21

AP Chemistry


2. Calculate the enthalpy change of the water using the equation qcold = mcoldc∆Tcold. Assume that the density of water is exactly 1 g/mL. Is this an endothermic or exothermic process?Explain.

3. These enthalpy amounts from the hot and cold water are not equal because the calorimeter(coffee cup) absorbs some of the thermal energy transferred by the hot water. Thus, under realconditions observed in the laboratory, the law of conservation of energy equation becomes qhot=-(qcold + qcal), where qcal is the enthalpy change of the calorimeter. Use this equation to calculatethe enthalpy change of the calorimeter.

4. The calorimeter constant, C, is the heat absorbed by the calorimeter per degree of temperaturechange, C = qcal/∆Tcal. Assuming the starting temperature of the calorimeter is the same as thatof the cold water, calculate the calorimeter constant in units of J/°C.

Part 2: Investigation of Ionic Solids

The solid and water, considered together, have a certain amount of internal energy as a function of the bonds that exist in the solid and in the water. The solution that is produced as a result of the dissolving has a different amount of internal energy than the solid and water did because the arrangement of particles and the bonds and attractions between the particles in the solution are different bonds and particulate attractions than the arrangement of particles and the bonds and attractions between the particles in the solid and water. The difference in energy, qsoln, is the reason for the difference in the thermal energy of the two systems (solid and pure water versus solution), with symbol qrxn. Just as with the hot and cold water in the calorimeter constant determination, qsoln and qrxn are equal in magnitude and opposite in sign, qrxn = –qsoln. And just as in that case of the cold and hot water mixing, the calorimeter will also experience an enthalpy change during the solution formation process. To account for this enthalpy change, the relationship is adjusted to qsoln = –(qrxn + CΔT) where C is the calorimeter constant determined above.

This difference in thermal energy of the system before and after solution formation, qsoln, can be calculated using the relationship qrxn = mcΔT, where m is the total mass of the solution and c is the specific heat capacity of the solution and ΔT is the temperature change of the solution. It is important to note that we will assume that the heat capacity of the solutions is the same as pure water but in reality the solutions do not have exactly the same heat capacity, and this assumption affects the accuracy of this determination. Also assume that the density of water is exactly 1 g/mL.

5. Using this information, calculate qsoln and qrxn for all three solids you tested for your handwarmer.

6. By convention, scientists report enthalpy changes for dissolution (and many other processes) inunits of kilojoules per mole of solute dissolved. Using your values of qsoln, calculate the enthalpyin units of kilojoules per mole. This quantity has the symbol ΔHsoln.

7. Report your calculated values for each of the ionic solids in an organized table (you will have tocreate this). This table should include cost and any environmental issues as well as the heatdata.

22

AP Chemistry


8. Based on the cost information provided, and your experimental work and calculations, selectwhich chemical you believe will make the most cost effective hand warmer. The hand warmeryou are designing needs to increase in temperature by 20°C. Calculate the amount of thecompound you selected that would be required for a hand warmer that meets this requirement.

9. Write a paragraph in which you describe all of the factors you considered and explain yourrationale for choosing one chemical and not each of the other chemicals studied in thisexperiment. Your paragraph should start with a claim sentence that clearly states your choiceand the amount of substance to use. The claim should be followed by evidence from yourexperiment and cost and safety analysis. The paragraph should conclude with reasoningexplaining how your evidence supports your claim.

23

AP Chemistry


Group solid assignment -- Students work in pairs (24 and under) or groups of three.

Group # Ionic Solids

1 NH4NO3 Na2CO3 CaCl2 2 NaCl LiCl NaC2H3O2 3 NH4NO3 LiCl NaC2H3O2 4 NaCl Na2CO3 CaCl2 5 NH4NO3 LiCl NaC2H3O2 6 NaCl CaCl2 NaC2H3O2 7 NH4NO3 NaC2H3O2 Na2CO3 8 NaCl LiCl Na2CO3 9 NH4NO3 LiCl CaCl2

10 NaCl NaC2H3O2 Na2CO3 11 NH4NO3 LiCl Na2CO3 12 NaCl LiCl CaCl2

24

AP Chemistry

Determination of the Rate of a Reaction, Its Order, and Its Activation Energy

Reaction kinetics is defined as the study of the rates of chemical reactions and their mechanisms. Reaction rate is simply defined as a change in a measurable quantity divided by the change in time. In chemistry, the “measurable quantity” is usually molar concentration or absorbance. Consider the generalized chemical reaction equation A + B C + D. Symbolically it can be represented in multiple ways:

Note that the units on rate are always M/time which can also be expressed as M time . The negative sign on the first expression indicates that the molar concentration of reactant A will decrease as time goes by. The second expression is simply the differential rate law expression where the rate constant k, and the order of reactant A (the exponent m) must be experimentally determined. Never, ever forget that the value of k is temperature dependent. Since two reactants are present in our example reaction we can write comparable expressions for reactant B, but beware that the order of B will not necessarily be the same as the order for A, so we often use a different variable, such as n, for the exponent on B. The differential rate law can be integrated to link changes in concentration with time as opposed to rate. This sounds way more complicated that it really is! (“Integrated” is a Calculus term you need not worry about in this course—we will linearlize the data to avoid Calculus since it is not a prerequisite for AP Chemistry.) In this experiment you will investigate the reaction of crystal violet with sodium hydroxide. Crystal violet, in aqueous solution, is often used as an indicator in biochemical testing. The reaction of this organic molecule with sodium hydroxide can be simplified by abbreviating the chemical formula for crystal violet as CV.

CV+(aq) + OH–(aq aq) As the reaction proceeds, the violet-colored CV+ reactant will slowly change to a colorless product, following the typical behavior of an indicator. The color change will be precisely measured by a colorimeter (see Figure 1) or spectrophotometer set at 565 nm (green) wavelength. You can assume that absorbance is directly proportional to the molar concentration of crystal violet according to Beer’s law.

Figure 1 The rate law for this reaction is in the form: rate = k[CV+]m[OH ]n, where k is the rate constant for the reaction, m is the order with respect to crystal violet (CV+), and n is the order with respect to the hydroxide ion. Since the hydroxide ion concentration is much more than the concentration of crystal violet, [OH ] will not change appreciably during this experiment. This technique is often referred to as “swamping”. Thus, you will find the order with respect to crystal violet (m), but not the order with respect to hydroxide (n). Therefore, the rate constant you will determine is a pseudo rate constant.

A[A]mRate k

time= =

1 Reaction Rate

25

Determination of the Rate of a Reaction, Its Order and Its Activation Energy

You will use integrated rate law methods to determine the order m and the value of the rate constant by graphing the absorbance and time data that you collect. Set up your axes so that time is always on the x-axis. Plot the absorbance of CV+ on the y-axis of the first graph. Plot the natural log of the absorbance of CV+(ln [CV+], NOT log[CV+]) on the y-axis of the second graph and the reciprocal of the absorbance of CV+ on the y-axis of the third graph. You are in search of the best linear fit. Here comes the elegant part… If you do the set of graphs in this order with the y-axes being “concentration”, “natural log of concentration” and “reciprocal concentration”, the alphabetical order of the y-axis variables leads to orders of and 2 respectively for CV+. You can then quickly derive the integrated rate law equations using y = mx + b.

Another important part of the kinetic analysis of a chemical reaction is to determine the activation energy, Ea. Activation energy can be defined as the energy necessary to initiate an otherwise spontaneous chemical reaction so that it will continue to react without the need for additional energy. An example of activation energy is the combustion of paper. The reaction of cellulose and oxygen is spontaneous, but you need to initiate the combustion by adding activation energy from a lit match. We can use a different graphical analysis method to easily determine the activation energy of a chemical reaction. Each laboratory group will simply repeat the reaction between crystal violet and sodium hydroxide at a temperature other than room temperature, while keeping the initial concentrations of the reactants the same for each trial. Recall that the value of k is temperature dependent. Class data will be collected, graphed and analyzed as follows:

Zero order k = negative slope

First order k = negative slope

Second order k = the slope

ln k

Ea = −R × slope

2 Reaction Rate

26


OBJECTIVES In this experiment, you will

• React solutions of crystal violet and sodium hydroxide at different temperatures.• Graph the concentration-time data and use integrated rate law methods to determine the

order of CVand the value of a pseudo rate constant, k, for the reaction.

• Measure and record the effect of temperature on the reaction rate and rate constant.• Calculate the activation energy, Ea, for the reaction.

Figure 1

MATERIALS Data collection device (computer or handheld) Colorimeter or spectrophotometer Temperature probe or thermometer Cups or beakers

M -5 M crystal violet solution

Plastic cuvettes Pipettes Ice or hot water bath PROCEDURE 1. Obtain and wear goggles.2. Set up the data collection system.

a. Calibrate your spectrophotometer or your colorimeter. We will be collecting data using the565 nm (Green) setting.

b. Connect a temperature probe to your device.c. If using a spectrophotometer, you will need to manually record data every 5 seconds. If using

a colorimeter set the program to generate a time graph with 3 seconds between samples for60 samples.

Do NOT start data collection until Step 3 b. 3. This first trial will be M

NaOH (colorless solution) and a pipette containing an equal quantity of 2.5 –5 M crystal violet (purple solution). a. Simultaneously squirt both solutions into a beaker or cup. Use the tip of the temperature

probe to stir the mixture. Record the temperature of the mixture. b. Rinse the cuvette with the mixture, discard the rinse into the sink, refill the cuvette at least

2/3 full and place it correctly in the colorimeter and start data collection for the first trial. c. Once the trial is finished, discard the reaction mixture into the sink.

3 Reaction Rate

27


4. Analyze your data using graphical methods as explained in the introduction. Determine theorder of CV+ and value of the rate constant, k. Record the value of the rate constant for this trialbefore proceeding to the next step.

5. Repeat Steps 3-4, using the second set of pipettes that have been sitting in the water bath atyour station.

6. Record your data in the class data table that your teacher has displayed. For your own lab report,mark your group number with an asterisk *.

4 Reaction Rate

28


PRE-LAB QUESTIONS Refer back to generalized chemical reaction presented in the introduction and write twocomparable rate expressions for reactant B.

2. Compare the molar concentration of the crystal violet solution to that of the sodium hydroxidesolution. Approximately how much more concentrated is the sodium hydroxide? Justify your answer.

3. Why do we set the spectrophotometer or colorimeter to a wavelength of 565 nm or “green” inorder to measure the absorbance of crystal violet?

4. A student mixes 3.00 mL M M sodiumhydroxide both at 24.5 °C and collects the following data:

Time (min) Absorbance 0.206

6.22 0.072 0.066 0.055 0.055

a. Describe the graphical analysis steps the student should perform in order to determinethe

(i) Order of the reaction with respect to crystal violet (ii) Value of the rate constant, k

b. Use a graphing calculator or computer graphing software to determine the order of thereaction with respect to crystal violet. Justify your answer.

c. Write the law expression for this reaction. Justify your answer.d. Determine the value of k including its units. Justify your answer.e. Calculate the half –life of the reaction. Include units with your answer.f. Determine the absorbance of crystal violet at 3.00 minutes. Justify your answer.g. Determine the time at which the absorbance of crystal violet is equal to 0.060.

5. The student repeats the experiment at 32.5 °C with the same initial quantities and molarities ofcrystal violet and sodium hydroxide. Predict whether the value of the rate constant k will beincrease, decrease or remain unchanged. Justify your answer.

POST-LAB QUESTIONS AND DATA ANALYSIS Graph the class data and calculate the activation energy, Ea, for the crystal violet and sodiumhydroxide reaction.

2. Extrapolate your graph to predict the value of the rate constant k for this reaction at 40 °C.3. A well-known approximation in chemistry states that the rate of a reaction often doubles for every

4. A student failed to fill the cuvette 2/3 full with the reaction mixture. What effect does this errorhave on the measured absorbance values?

5 Reaction Rate

29


CLASS DATA TABLE Trial 1

Room Temperature (°C) Rate constant, k (supply appropriate

units)

Trial 2 Temperature (°C)

Rate constant, k (supply appropriate units)

Group 1

Group 2

Group 3

Group 4

Group 5

Group 6

Group 7

Group 8

Group 9

Group 10

Group 11

Group 12

Group 13

Group 14

Group 15

Group 16

6 Reaction Rate

30

AP Chemistry

The Determination of Keq for FeSCN2+ Introduction

There are many reactions that take place in solution that are equilibrium reactions; that is, they do not go to completion, both the forward and reverse reaction are occurring, and both reactants and products are always present. Examples of this type of reaction include weak acids such as acetic acid dissociating in water, weak bases such as ammonia reacting with water, and the formation of “complex ions” in which a metal ion combines with one or more negative ions. A reaction involving the formation of a complex ion which occurs when solutions of iron (III) are combined with solutions of the negative thiocyanate ion will be studied and the equilibrium constant determined.

Concepts � Chemical equilibrium� Complex-ion reaction

� Equilibrium constant� Colorimetry

Background Chemical reactions are driven to completion by two forces: a decrease in energy (exothermic

reaction), or an increase in entropy. If both an energy decrease and an entropy increase occur in the forward reaction, the reaction will go to completion. An example of this type of reaction is the burning of ethane. The reaction is exothermic and has an increase in entropy, and it goes to completion.

However, when an energy decrease drives a reaction in one direction and an entropy increase drives it in the reverse direction, equilibrium will result. The reaction will not go to completion, but will reach a point where both reactants and products are present in a fixed ratio of concentration. The reaction will continue at the same rate in both forward and reverse directions, and the concentrations of products and reactants will stay constant.

These ideas can be expressed mathematically in the form of the equilibrium constant. Consider the following general equation for a reversible chemical reaction:

Equation 1 The equilibrium constant Keq for this general reaction is given by Equation 2, where the square

brackets refer to the molar concentrations of the reactants and products at equilibrium.

ba

dc

eq BADCK

][][][][

= Equation 2

The equilibrium constant gets its name from the fact that for any reversible chemical reaction, the value of Keq is a constant at a particular temperature. The concentrations of reactants and products at equilibrium vary, depending on the initial amounts of materials present. The special ratio of reactants and products described by Keq is always the same as long as the system has reached equilibrium and the temperature does not change. The value of Keq can be calculated if the concentrations of reactants and products at equilibrium are known.

The reversible chemical reaction of iron (III) ions (Fe3+) with thiocyanate ions (SCN-) provides a convenient example for determining the equilibrium constant of a reaction. As shown in Equation 3, Fe3+ and SCN- ions combine to form a special type of combined or “complex” ion having the formula FeSCN2+.

Fe3+(aq) + SCN- 2+(aq) Equation 3 (Pale yellow) (Colorless) (Blood-red)

The equilibrium constant expression for this reaction is given in Equation 4.

]][[][

3

2

=SCNFe

FeSCNKeq Equation 4

1 Finding Keq

31

The Determination of Keq for FeSCN2+

The value of Keq can be determine experimentally by mixing known concentrations of Fe3+ and SCN- ions and measuring the concentration of FeSCN2+ ions at equilibrium. As noted in Equation 3, the reactant ions are pale yellow and colorless, respectively, while the product ions are blood-red. The concentration of FeSCN2+ complex ions at equilibrium is proportional to the intensity of the red color. Compounds that are colored absorb a part of the visible spectrum of light. If a compound absorbs green light, it will appear red in color, the complementary color to green. A spectrophotometer is an instrument that measures the amount of light of a given wavelength or color that is absorbed by a solution. A visible spectrophotometer consists of several components. First, a source of visible light is needed. Next, a prism or diffraction grating and lenses are used to select light of a given wavelength and pass it through the solution. A sample holder, usually a test tube or cuvette is used to hold the solution. Next a photo tube or device to measure the amount of light that comes through the solution is needed. The last component is a meter to display the amount of light transmitted or absorbed by the solution. To use the spectrophotometer, the desired wavelength must be selected, insert the sample inserted in the light path, and the amount of light absorbed or transmitted read. A colorimeter is a similar device that uses a color filter in place of the prism or diffraction grating. Both a colorimeter and spectrophotometer can determine the concentration of a colored solution using Beer’s law. Beer’s law is a mathematical relation that is stated: In Beer’s law, the term “A” is the absosample. This is a value that indicates how much light of a given wavelength is absorbed by a given substance. If a substance has a very intense color, it will have a high absorptivity. The term “b” is the cell path. This is the width of the cell path (test tube or cuvette) through which the light passes. The term “C” is the molar concentration of the substance absorbing the light. When a substance is analyzed the absorptivity of the sample is constant, and the width of the cell path (test tube or cuvette) does not change. Therefore, the absorption read by the spectrophotometer is directly proportional to the concentration of the sample. A graph of absorption versus concentration is a straight line. Experiment Overview The purpose of this experiment is to calculate the equilibrium constant for the reaction of iron (III) ions with thiocyanate ions. The reaction is tested under different conditions to determine if the equilibrium constant always has the same numerical value. There are two parts to the experiment. In Part 1, a series of reference solutions and test solutions are prepared. The reference solutions are prepared by mixing a large excess of Fe3+ ions with known amounts of SCN- ions. According to LeChatelelier’s Principle, the large excess of iron (III) ions should effectively convert all of the thiocyanate ions to the blood-red FeSCN2+ complex ions. The concentration of FeSCN2+ complex ions in the reference solutions is essentially equal to the initial concentration of SCN- ions. The test solutions are prepared by mixing a constant amount of Fe3+ ions with different amounts of SCN- ions. These solutions contain unknown concentrations of FeSCN2+ ions at equilibrium. In Part 2, the absorbance’s of both the reference solutions and the test solutions are measured by colorimetry. A calibration curve is constructed from the absorption values of the reference solutions. The unknown concentrations of FeSCN2+ in the test solutions are calculated by comparing their absorbance readings to the absorbance values of the calibration curve. These values are used to determine the equilibrium concentrations and the equilibrium constant for the reaction. Pre-Lab Questions

1.

Finding Keq 2 32


2. The reaction for the formation of the diamminesilver ion is as follows: Ag+(aq) + 2NH3(aq) Æ Ag(NH3)2

+(aq) a. Write the equilibrium constant expression for the reaction. b. An experiment was carried out to determine the value of the equilibrium constant Keq

for the reaction Total moles of Ag+ present 3.6x10-3 moles Total moles of NH3 present 6.9x10-3 moles Measured concentration of Ag(NH3)2

+ at equilibrium 3.4x10-2 M Volume of the solution 100 mL i. Calculate the equilibrium concentration of Ag+ (uncomplexed) ii. Calculate the equilibrium concentration of NH3 (uncomplexed) iii. Calculate the value of the equilibrium constant

3. “The equilibrium concentration of FeSCN2+ ions in each reference solution listed below is essentially equal to the concentration of SCN- ions in solution before any reaction occurs.” Use LeChatelier’s Principle to explain why this statement is true.

4. The five reference solutions in Part 1 are prepared by mixing 0.200 M Fe(NO3)3 solution and 0.00020 M KSCN solution in the amounts listed in the following table. Standard Volume of 0.200 M Fe(NO3)3 Solution Volume of 0.00020 M KSCN Solution

Reference Solution #1 8.0 mL 2.0 mL Reference Solution #2 7.0 mL 3.0 mL Reference Solution #3 6.0 mL 4.0 mL Reference Solution #4 5.0 mL 5.0 mL Reference Solution #5 4.0 mL 6.0 mL

The concentration of Fe3+ ions in the first reference solution (M2) before any reaction occurs can be calculated using the “dilution equation” as shown below. M1V1 = M2V2 Dilution Equation

M1 = concentration of solution before mixing = 0.200 M Fe(NO3)3 V1 = volume of solution before mixing = 8.0 mL V2 = final volume of reference solution after mixing = 8.0 mL + 2.0 mL = 10.0 mL

MmL

mLMVVMM 16.0

)0.10()0.8)(200.0(

2

112 ===

Use the dilution equation to calculate the concentration of SCN- ions in the five reference solutions 2+]

5. The following table summarizes the volumes of Fe3+ and SCN- stock solutions that will be mixed together to prepare the test solutions in Part 1. Use the dilution equation to calculate the concentration of Fe3+ and SCN- ions in each test solution before any reaction occurs. Enter the results of these calculations Hint: The final volume (V2) of each test solution is 10.0 mL.

Sample Volume of 0.0020 M Fe(NO3)3 Solution

Volume of 0.0020 M KSCN Solution

Volume of Distilled Water Added

Test solution #6 5.0 mL 1.0 mL 4.0 mL Test solution #7 5.0 mL 2.0 mL 3.0 mL Test solution #8 5.0 mL 3.0 mL 2.0 mL Test solution #9 5.0 mL 4.0 mL 1.0 mL

Test solution #10 5.0 mL 5.0 mL 0 mL

Finding Keq 3 33


Materials Chemicals

Iron(III) nitrate, Fe(NO3)3, 0.200 M, 30 mL† Iron(III) nitrate, Fe(NO3)3, 0.0020 M, 25 mL † Potassium thiocyanate, KSCN, 0.0020 M, 15 mL

Potassium thiocyanate, KSCN, 0.00020 M, 20 mL Water, distilled or deionized

Equipment Beakers or large test tubes, 50 mL, 10 Colorimeter sensor Cuvettes, 6 Labeling or marking pen

or bulb pipets

Stirring rod Thermometer Tissues or lens paper Wash bottle

†Contains 1 M nitric acid as the solvent ________________________________________________________________________________________________________ Safety Precautions Iron(III) nitrate solution contains 1 M nitric acid and is a corrosive liquid; it will also stain skin and clothing. Notify the teacher and clean up all spills immediately. Potassium thiocyanate is toxic by ingestion; it can generate poisonous hydrogen cyanide gas if heated strongly. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory. _____________________________________________________________________________________ Procedure Part 1. Preparing the Solutions

1. Obtain ten 50 mL beakers or large test tubes. 2. Prepare the five reference solution test tubes or beakers listed in the table below. Use a

separate pipet or a dispensing buret to transfer the appropriate volumes of each reagent. Mix each solution using a stirring rod. Rinse the stirring rod and dry it between solutions. Label the test tubes or beakers with the corresponding reference solution number.

Standard Volume of 0.200 M Fe(NO3)3 Solution

Volume of 0.00020 M KSCN Solution

Reference Solution #1 8.0 mL 2.0 mL Reference Solution #2 7.0 mL 3.0 mL Reference Solution #3 6.0 mL 4.0 mL Reference Solution #4 5.0 mL 5.0 mL Reference Solution #5 4.0 mL 6.0 mL

3. Using a separate pipet for each reagent to be added or a dispensing buret, combine the

following volumes of reagents to prepare the test solution. Note: Label the tubes with the corresponding solution numbers 6 through 10. Read the reagent labels carefully before use!

Sample 0.0020 M Fe(NO3)3 0.0020 M KSCN Distilled Water Test solution #6 5.0 mL 1.0 mL 4.0 mL Test solution #7 5.0 mL 2.0 mL 3.0 mL Test solution #8 5.0 mL 3.0 mL 2.0 mL Test solution #9 5.0 mL 4.0 mL 1.0 mL

Test solution #10 5.0 mL 5.0 mL 0 mL

Finding Keq 4 34


4. Mix each solution using a stirring rod. Rinse the stirring rod and dry it between solutions. 5. M

used as the equilibrium temperature for all of the solutions.

Part 2. Colorimetry Measurements

1. Follow the procedure for your colorimetric measurements of the solution as found in the “LabQuest You will need to use light with a wavelength of about 450 nm. Handle cuvettes at the top so no fingerprints are in the light path. Clean cuvettes with a tissue. Calibrate the colorimeter by placing a cuvette which is about 2/3 full of distilled water into the sample holder and making sure the absorbance is zero. Fill a cuvette about 2/3 full of a test solution, place it in the colorimeter and read the absorbance.

2. Measure the absorbance of each of the reference solutions at 450 nm, using distilled water as the zero absorbance reference.

3. Repeat steps 1 and 2 for each of the test solutions. Record the absorbance 4. cuvettes and of the remaining test solutions as directed by your

teacher. Follow your teacher’s directions for rinsing and drying the cuvettes. LabQuest Directions

1. Obtain a SpectroVis Plus and a LabQuest. Plug the SpectroVis Plus into the LabQuest after turning the LabQuest on.

2. Click on “USB:Abs” and find “Change Wavelength”. Enter 450 into the box. 3. Click “Mode: Full Spectrum”, go to the drop down box at the top and press “Time Based”.

Change the rate to 1 and duration to 1. Press OK. 4. Click “USB: Abs” and select “Calibrate”. Calibrate with a clean cuvette and distilled water. 5. Find the Table tab on top. You can put in your samples and test them one at a time. The

LabQuest will record the absorbance values in the right column.

Disposal and Cleanup • all reference and test solutions in the waste beaker provided. • Remove any labeling tape used. • Wash and dry all equipment used. • Store all cords in the correct box for the LabQuest and SpectroVis Plus. • Return all borrowed equipment and store all other equipment.

Finding Keq 5 35


Data Tables Data Table 1 – Reference Solutions Temperature Sample [FeSCN2+] Absorbance Reference Solution #1

Reference Solution #2




Data Table 2 – Test Solutions Temperature Sample [Fe3+]* [SCN-]* Absorbance Test Solution #6

Test Solution #7

Test Solution #8

Test Solution #9

Test Solution #10

* These are the concentrations of ions in solution immediately after mixing and before any reaction has occurred. See the Pre-Lab Questions for calculations. Data Table 3 – Results Sample [FeSCN2+]eq [Fe3+]eq [SCN-]eq Keq Test Solution #6

Test Solution #7

Test Solution #8

Test Solution #9

Test Solution #10

Average value

Average deviation

Finding Keq 6 36


Post-Lab Calculations and Analysis 1.

concentration of FeSCN2+ versus absorbance. Absorbance will be on the vertical axis and molar concentration of FeSCN2+ will be on the horizontal axis. Using a linear regression (trend line), find the line of best fit. Title and appropriately label your graph. Print it out to include the equation for your line of best fit. Include this graph in your lab report.

2. The unknown concentrations of FeSCN2+ ions in each test solution can be determined from the graph and line (y-values), calculate the concentration of FeSCN2+ (x-values).

3. Record the FeSCN2+ 4. Calculate the equilibrium concentration of Fe3+ ions in each test solution #6-10: subtract the

equilibrium concentration of FeSCN2+ ions from the initial concentration of Fe3+ ions. Enter the

3+]eq 3+]initial – 2+]eq 5. Calculate the equilibrium concentration of SCN- ions in each test solution #6-10: subtract the

equilibrium concentration of FeSCN2+ ions from the initial concentration of SCN- ions. Enter the

-]eq -]initial – 2+]eq 6. Use your equilibrium expression for the reaction to calculate the value of the equilibrium

constant Keq for each test solution #6- 7. Calculate the mean (average value) of the equilibrium constant for the five test solutions. 8. Calculate the average deviation for Keq: Find the absolute value of the difference between each

individual value of the equilibrium constant and the mean. The average of these differences for solutions #6-10 is equal to the average deviation.

9. The average deviation describes the precision the equilibrium constant is indeed a “constant” for this reaction? Explain.

10. Post-Lab Questions

1. What does the calculated value of the equilibrium constant, Keq, indicate regarding the degree of completeness of the reaction? In other words, at equilibrium, are there mostly products, reactants, or relatively large amounts of both?

2. What was the color of light chosen for the experiment? What was the color of the FeSCN2+ complex ion? When using a colorimeter, should you set the wavelength of light to be the same color of that as the solution? Explain.

3. What degree of precision (to how many significant figures) can you obtain with the colorimeter that was used? What is the major source of error in the experiment?

Finding Keq 7 37

AP Chemistry

Examining Le Chatelier's Principle

Le Chatelier's Principle states that if an equilibrium system is subjected to a stress, the system will react to remove the stress. To remove a stress, a system can only do one of two things: form more products by using up reactants, or reverse the reaction and form more reactants, using up products. In this experiment, you will examine several equilibrium systems. By putting different stresses on the systems, you will force the system to react in a desired manner.

Acid-base indicators are large organic molecules that can gain and lose hydrogen ions to form substances that have different colors. The reaction of the indicator bromothymol blue can be illustrated as follows: HIn(aq) Æ H+(aq) + In-(aq) (Yellow) (Blue) In this reaction, HIn is the neutral indicator molecule, and IN- is the indicator after the molecule has lost a hydrogen ion. Equilibrium reactions can easily be forced to go in either direction.

An equilibrium solution can be formed with the following ions: Fe3+(aq) + SCN-(aq) Æ FeSCN2+(aq) (colorless) (colorless) (red-brown) The iron ion (Fe3+) and the thiocyanate ion (SCN-) are both colorless; however, the ion that forms from their combination, the FeSCN2+ ion, is a dark red-brown. It is the color of this ion that will indicate how the equilibrium system is being affected. Phosphate ions have the ability to form complex ions with Fe3+, which has the same effect as removing Fe3+ from solution.

The equilibrium between two different complex ions of cobalt will also be investigated. The reaction is endothermic: Co(H2O)6

2+(aq) + 4Cl-(aq) Æ CoCl42-(aq) + 6H2O(l) H = +50 kJ/mol

(Pink) (Blue) Silver and chloride ions combine to form a precipitate of AgCl. Purpose To force an equilibrium system to achieve a desired result by using Le Chatelier's principle. Safety

• HCl is hazardous and should be handled with care. It also has strong vapors that should not be inhaled. Wash spills off with lots of water, and if there are any spills, neutralize it with baking soda.

• NaOH is also hazardous. If spilled, neutralize with vinegar. • Ethanol is flammable. Turn off all flames. • Silver nitrate causes stains on skin and clothing. Wash spills off with soap and water

immediately. • Wear lab safety goggles and a chemical resistant apron.

1 Le Chatelier

38

AP Chemistry Materials

HCl, 12 M and 0.1 M NaCl(s) KSCN, 0.002 M KSCN (s) Bromothymol blue indicator solution NaOH, 0.1 M Fe(NO3)3, 0.2 M AgNO3, 0.1 M Ethanol

Cobalt (II) chloride / CoCl2·6H2O(s) Disodium hydrogen phosphate, Na2HPO4(s) Test tubes, 13x150 mm Test tube rack Beaker, 100 mL Graduated cylinders Funnel, filter paper, and holder for funnel Ice bath Hot water bath

Procedure Your lab group will be given various solutions and asked to complete the tasks listed below. You must develop a procedure to accomplish each using only the equipment that is listed above and the chemicals provided at your table. Your procedure must stay within the confines of the task. You will write the procedure during the lab. If your group tries a method but it does not work to accomplish the task, note this in your observations.

Tasks:

1. Given a saturated solution of sodium chloride in water, precipitate out some sodium chloride. Recover the precipitated sodium chloride. You may not add any additional NaCl to the solution to accomplish this.

2. Given an aqueous solution containing bromothymol blue, change the color of the solution to yellow and then to green.

3. Given 25 mL of 0.002 M KSCN and 5 mL of 0.2 M Fe(NO3)3, a. Create a red-brown solution b. Return the red-brown solution to as colorless as possible.

4. Separate a solution of hydrated cobalt ion into three different test tubes. a. Turn each solution blue using a different method. b. Turn two of the solutions pink using different methods.

Analysis

1. Give all observations in an organized manner in which you describe the system, the stress, and the result. In addition, for each stress, describe why the observed effect occurred using the appropriate chemical reaction.

2. Which of the tasks were the most difficult to achieve? Why? 3. Were there any tasks which you were not able to achieve? If so, describe a method that you

would try in lab in order to achieve this. 4. A student obtained a test tube with a suspension of white, slightly soluble calcium hydroxide in

water. This system was at equilibrium as represented by the following equation: Ca(OH)2(s) Ca2+(aq) + 2OH-(aq)

2 Le Chatelier

39

AP Chemistry

a. Write the equilibrium constant expression for this reaction. b. What would you expect to observe if hydrochloric acid, HCl(aq), was added? Explain

your answer using Le Chatelier's principle. c. What would you expect to observe if calcium nitrate was added? Explain your answer

using Le Chatelier's principle. d. When the solution was placed in an ice bath and cooled, it was observed that more solid

calcium hydroxide was produced. i. Based on this observation, would you expect the reaction to be exothermic or

endothermic? Explain your answer using Le Chatelier's principle. ii. If the solution was placed in a hot water bath and heated, what would you

expect to observe? Explain your answer using Le Chatelier's principle.

3 Le Chatelier

40

AP Chemistry

How Much Acid Is in Fruit Juices and Soft Drinks?

Have you ever wondered how doctors determine the ratios for IV drips during surgery or hospital stays? They use a common laboratory procedure called a titration to calculate specific ratios of different substances using volume measurements. A similar process occurs when someone uses a machine to monitor blood glucose levels, to analyze urine samples, or to conduct a pregnancy test. Pharmacists use titrations when compounding drugs, which allows them to more precisely match a person's drug prescription to their body weight, size, or medical condition. While it may sound like titrations are common practice only in the chemistry lab or medical community, titrations have great practical use. Food scientists use them when testing levels of salt, sugar and vitamins in different foods, and for deciding if wine and cheese are ready for consumption. Others use titrations to test water quality or hardness and in neutralizing the free fatty acids in waste vegetable oil before refining it as biodiesel. In all cases, titrations are used to quantitatively analyze the unknown concentration of a solution or the amount of a substance by comparing it to a solution of known concentration. One of the first titrations you perform is the reaction between a hydrochloric acid solution of unknown concentration and sodium hydroxide. Since we are analyzing HCl to determine its concentration, we call it the analyte. The hydrochloric acid solution would be "titrated" by adding a standard solution of sodium hydroxide drop wise. The sodium hydroxide solution is called the titrant, and is a standard solution because its concentration is accurately known. The point at which enough titrant has been added to react exactly with the analyte is called the equivalence point. Since this point is not visible to the eye, we use a pH indicator to help us detect it. The point at which the indicator changes color is called the endpoint, and assuming it has been selected appropriately, indicates that the equivalence point has been reached. The best indicator for a titration is one in which the endpoint and equivalence point are as close together as possible. In this experiment, you will design your own acid-base titration to determine the acid concentration of fruit juice or carbonated beverages by using a standardized solution of sodium hydroxide, and make a prediction about how knowing the acid content of a certain beverage would be beneficial. Nutritionists cite the dangers of regular consumption of carbonated beverages on bone density and recommend a low-acid diet. Dentists caution patients about the relationship between acid and oral health. Is there cause for concern? Might the knowledge gained in this lab help us better understand their recommendations, and make reasonable decisions about our beverage of choice? Purpose

• Determine how much acid is in common fruit juices and soft drinks.

Safety Precautions

Develop these as part of your procedure.

1 Acid Content of Juices and Soft Drinks

41


The following materials will be available to you during the experiment. You should design your experiment with these options in mind.

• Beakers of varying sizes • 50 mL burets • 125 and 250 mL Erlenmeyer flasks • Graduated cylinders (10 and 50 mL) • Volumetric pipets • Various acid base indicators (methyl

orange, methyl red, bromothymol blue, phenolphthalein, thymol blue)

• pH probes • Lab Quests • 0.10 M acetic acid • Standardized solutions of 0.10 and 0.25

M sodium hydroxide solutions • Distilled water

Pre-Lab

1. Write the complete chemical equation for the reaction of a solution of sodium hydroxide (NaOH) with hydrochloric acid (HCl).

2. How many mL of 0.1 M HCl are required to react completely with 5 mL of 0.1 M NaOH? 3. If equal molar amounts of NaOH and HCl are mixed, when the reaction is complete, what will be

the chemical species in the resulting solution? 4. Will the pH of the mixture in question 3 be acidic, neutral, or basic? Explain. 5. Write the complete chemical equation for the reaction of a 0.1 M solution of acetic acid

(HC2H3O2) with a 0.1 M solution of NaOH. 6. How many mL of the 0.1 M NaOH will be required to react completely with 5 mL of a 0.1 M

acetic acid solution? Explain. 7. When the reaction is complete, what will be the pH (acidic, neutral, basic) of the solution in

question 6? Explain. 8. How is it possible to determine when an acid base reaction is complete when one of the

reactant's concentrations is unknown? 9. Using the table below, explain how indicators are chosen and used during titration.

Indicator pH Range Color Change Methyl orange 3.1-4.4 Orange to yellow

Methyl red 4.2-6.2 Red to yellow Bromothymol blue 6.0-7.6 Yellow to green-blue

Phenolphthalein 8.3-10.0 Colorless to pink Thymol blue 1.2-2.8 Red to yellow

Procedure

1. Identify a research question pertaining to acid concentration in fruit juices and sodas. Your question should involve at least two different beverages. Propose a hypothesis that answers your question, and then outline a tentative procedure. Be sure to clearly identify which pieces of equipment (from the Materials section) you plan to use during your investigation.


42

AP Chemistry

2. Consult MSDS or websites for information regarding safe handling of any chemicals you have chosen to use during your experiment. Describe at least three safety precautions you need to take as you perform your experiment. MSDS information can be found at http://www.ehso.com/msds.php

3. Make a simple sketch showing how you plan to set up your equipment, labeling the following: analyte, titrant, buret, Erlenmeyer flask, volumetric pipette. Show where the indicator solution will be added and comment on your choice of indicator: why did you choose this one and not another?

4. Submit your procedure, material selection, safety precautions, and sketch to your teacher at least 48 hours prior to the lab.

5. As you conduct your experiment, keep detailed written records. Be sure to list all steps taken as you perform your experiment, and all measurements and observations made during the experiments. Create any necessary data tables and put these in your data section. As you are monitoring pH, make sure you save your data for future graphing.

6. Write the balanced chemical equation for the reaction between sodium hydroxide and the primary acid in the fruit juices or carbonated beverages your lab team selected. If you identified different acids in the drinks, write balanced chemical equations that represent each.

Data Data tables will be developed according to your procedure. Analysis

1. Calculate the acid concentration(s) of the fruit juice or sodas tested, being sure to show all work. Include units.

2. Calculate the pH of the fruit juices or sodas you tested. Enter this data on the class Google spreadsheet. Label your data with your group's names.

3. Using your data, graph pH versus volume of base added for all beverages tested. Make a note of where the indicator chosen changes color. Use a graphing utility to make a detailed graph. Print and include in your lab report.

4. Suppose a fellow student chose to measure solution volumes using beakers or graduated cylinders. What effect would this have had on the calculated acid concentration? How might this affect the number of significant figures in your final answer? Explain your answers.

5. A fellow student rinsed the buret with water, but neglected to rinse the buret with titrant before conducting the experiment. What effect would this have on the calculated acid concentration in the juices or sodas? Why?

6. If you did not titrate a cola like Coke, Pepsi, or Dr. Pepper, find another lab group that did so, and ask them to discuss their procedures and results. What step was necessary to determine the endpoint when titrating a cola with a standard sodium hydroxide solution, and why did this step matter?


43

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AP Chemistry

7. Create a diagram that shows the molecular interaction between the acid and base as the titration proceeds. Display this at four points along a titration curve: (a) when 0 mL of base is added, (b) when 5 mL of base is added, (c) at the equivalence point, and (d) when an excess of base has been added, and provide an explanation for what is happening at the molecular level.

8. After conducting this experiment, what recommendation(s) might you make for a person with acid-reflux disease or tooth decay despite the fact that they drink juice or diet sodas? Justify your answer.


44

AP Chemistry

How Do the Structure and Initial Concentration of an Acid and a Base Influence the pH of the Resultant Solution During a Titration?

Many foods taste as they do due to the presence of acidic or basic content. All foods, beverages, pharmaceuticals, biofuels, water in aquariums, drain cleaners, surface cleaners, and vitamins contain acids or bases, or a mixture of acids and bases. The amount of acid, base and the pH of solutions and solids must be maintained at an optimal level. If a solution is too acidic, some base can be added to react with some of the acid. By carefully controlling the amount of base added while doing an acid-base titration, and knowing when to stop adding base by using an indicator or a pH meter, one can determine the amount of acid present in the substance. The food industry uses titrations to determine the amount of sugar, free fatty acid content, and the concentration of vitamin C or E present in products. While there are times when you only need to know if a solution is acidic, basic, or neutral, often the exact concentration is important, such as when making biodiesel fuel from vegetable oil. When vegetable oil degrades it becomes acidic. A base such as lye is added to neutralize the acid. The exact concentration of the acid must be known because if too much base is added, you will get soap instead of biodiesel fuel! A titration of the acid with a base will determine the exact concentration. Besides doing the titration you must be able to analyze the resultant titration curve. Purpose

• Explore the differences in titration curves between strong and weak acids and bases. • Determine how the structure and initial concentration affect the results of a titration.

Safety Precautions

• Acidic and basic solutions can be dangerous. Review the specific hazards according to the MSDS sheets appropriate for your chosen acids/bases.

• Acid base reactions are exothermic. Be sure to always add acid/base to water. • Splash proof goggles and aprons must be worn at all times. If solutions are spilled, inform your

teacher immediately. If solutions get on skin, in should be rinsed with running water for 15 minutes.

Materials

The following materials will be available to you during the experiment. You should design your experiment with these options in mind.

• Hydrochloric acid • Sulfuric acid • Nitric acid • Calcium hydroxide • Ammonia • Acetic acid • Sodium hydroxide

• Graduated cylinders • Erlenmeyer flasks • Beakers • Burets • Magnetic stirrer and stir bar • pH probes • Lab Quests

1 Acid Base Titration

45

AP Chemistry Pre-Lab With your lab group, choose one of the following questions to investigate during this lab.

1. Given 50 mL of 0.10 M HCl and 50 mL of 0.10 M acetic acid, will the amount of 0.10 M NaOH required to neutralize each solution be the same, more, or less?

2. Will the pH at the equivalence point of 50 mL 0.10 M HCl be the same, more, or less as the pH at the equivalence point for 50 mL of 0.10 M acetic acid?

3. What are some structural features that might help us classify an acid as a strong acid or weak acid?

4. Draw a molecular and particulate view of what is happening in the steep part of a general acid-base titration curve.

5. How does the structure of an acid affect the shape of the titration curve? 6. How can a pH titration curve be used to help classify the resultant solution at the endpoint, as

acidic, basic, or neutral? 7. How do the structure and the initial concentration of an acid and a base influence pH of the

resultant solution during a titration?

** No more than three lab groups can choose a particular question. "Claim" to a question will be on a first come first served basis with the completed procedure. Your group will write a procedure that uses different combinations of two acids: one with a known molarity and the other with an unknown molarity, and two bases, one with a known molarity and the other with an unknown molarity. The procedure requires that you perform acid-base titrations to collect data to draw titration curves, calculate the unknown molarities, and answer the question your group is investigating. You procedure must be submitted at least 48 hours prior to the lab. **The known molarities will be given in lab. They are all 0.20 M and below. Procedure This is developed according to your pre-lab.

After collecting your data, make a titration graph. All major points should be identified on the curves. If a sample includes a weak acid and/or base, percent ionization needs to be calculate, Ka and/or Kb needs to be calculate, and the percent error of the calculated Ka or Kb to the accepted value should be calculated.

Disposal: All solutions should be neutralized and the pH tested so that the waste can be safely disposed of down the drain.

Data

Data tables will be developed according to your procedure. Titration curves should be displayed in this section as described above.


46

AP Chemistry Analysis As described in the pre-lab section, you must include the following items in your lab report:

1. Titration curves (labeled) 2. Calculations of the unknown molarities 3. Answer to the pre-lab question you are investigating

In addition, answer the following questions:

4. How do the titration curves drawn vary if the acids or bases are weak or strong? Justify your answer.

5. What would a titration curve look like if an indicator were used to signal when to stop the titration?

6. Commercials about antacids are on television all the time. How would you go about investigating bases like antacids that are solid?

7. How would you investigate which antacid neutralizes the most acid or is the most cost effective? 8. Explain how rinsing the buret with water instead of the titrant before starting the investigation

will affect the calculated unknown molarity of the titrate. 9. What types of data needs to be collected to perform molarity calculations of the unknown? 10. Does the presence of weak or strong acids and weak or strong bases make a difference to when

the equivalence point occurs? Justify your answer.


47

AP Chemistry

The Preparation and Testing of an Effective Buffer: How Do Components Influence a Buffer's pH and Capacity?

Many important biochemical reactions occur only over a small range in pH. Living organisms dependent upon these reactions rely on chemical systems called buffers to maintain a relatively constant pH when acids or bases are added to their environment. The company contracting the students’ chemistry class aims to produce a variety of bacteria designed to destroy harmful living agents. The bacteria must be grown in a medium with a pH similar to that of the environment they will be functioning in. This medium must be able to maintain a pH within plus or minus one unit of the target pH for the bacteria it will support when strong acid or base is added. The students’ job will be to produce a buffer for such a medium. As buffers must neutralize both acids and bases, they must contain both a base and an acid. The question is, how do we prevent the acid and base in the buffer from simply neutralizing one another, thus rendering the buffer useless? Students must also consider how much acid and base their buffer should be able to neutralize. Buffers are most effective when they have been produced so they may neutralize a reasonable amount of either acid or base. The quantity of acid or base that may be added to a buffer while maintaining a relatively constant pH is a function of the buffer’s capacity. You and your classmates are being asked to prepare and test a series of buffer s to be used in an important biochemical project. Each group will be given an assignment card for one of the buffers. Once you have performed the experiment to achieve the assignment on the card, you will share and compare your results with those of your classmates. You will then argue, with evidence, whether you've completed your assignment or need to rely on the product of another group. Scientists often share data and information with other researchers. The need to interact with others is not an indication of failure provided your final argument is sound. Purpose

• Create a buffer solution that will resist the addition of moderate amounts of strong acid and base.

Safety Precautions

• Review the MSDS sheets for all chemicals used in your procedure. • Sodium hydroxide is caustic and hydrochloric acid is corrosive. If spilled on skin, wash with

copious amounts of water. • Splash proof goggles and aprons should be worn at all times.

1 Preparing a Buffer

48


The lab equipment listed will be used for this experiment. Choose the chemicals necessary to make your assigned buffer solution from the options listed below.

2-150 mL beakers 0.1 M acetic acid, Ka = 1.8x10-5 Solid sodium acetate 2-250 mL beakers 0.1 M ammonia Solid ammonium chloride, Ka of

NH4+ = 5.7x10-10

Magnetic stirrer and stir bar 0.10 M sodium dihydrogen phosphate, NaH2PO4, Ka of H2PO4

- = 6.3x10-8

Solid sodium hydrogen phosphate

pH probe 0.1 M citric acid, H3C6H5O7, Ka = 7.1x10-4

Solid sodium dihydrogen citrate, NaH2C6H5O7

Lab Quest 0.1 M sodium monohydrogen citrate, Na2HC6H5O7, Ka = 4.1x10-7

Solid sodium citrate, Na3C6H5O7

Balance 0.2 M sodium hydroxide Solid sodium chloride 2-50 mL burets and stands 0.2 M hydrochloric acid 100 mL graduated cylinder 10 mL graduated cylinder

Pre-Lab 1. How does a buffer solution resist a change in pH? 2. Why would HCl and NaOH be a poor choice for an acid-base pair to make a buffer? 3. Go to the animation at: http://introchem.chem.okstate.edu/DCICLA/pHbuffer20.html

Create a "buffer solution" using equal volumes (100 mL each) of 0.10 M nitric acid and 0.10 M sodium nitrate. Insert the probes and record the pH in the table below. Remove the probes. Then go to Part II of the simulation and add first 0.001 moles of HCl and then the same amount of NaOH. Record the pH in each case. Remove the probes, return to Part I, and try the next combination. Complete the first three lines of the table below. Remake the "buffers" using the same volumes of 1.0 M components (increased molarity by 10x). Add the same 0.001 moles of HCl and NaOH to the "stronger" buffers. Complete the middle three rows of the table. Finally, repeat the process adding 0.011 moles (additional 10x more) of strong acid and base. Complete the last three lines of the table.


49

http://introchem.chem.okstate.edu/DCICLA/pHbuffer20.html

AP Chemistry

Acid-Base Pair pH of "buffer" pH with 0.001 mol HCl added

pH with 0.001 mol NaOH added

HNO3 and NaNO3

HC2H3O2 and NaC2H3O2

NH4Cl and NH3

Increase molarity 10x

HNO3 and NaNO3


NH4Cl and NH3

pH with 0.011 mol HCl added

pH with 0.011 mol NaOH added

HNO3 and NaNO3


NH4Cl and NH3

a. Write a general chemical equation to represent the equilibrium that exists in an aqueous

system of a weak acid, HA, as it ionizes in water. Assume the weak acid to be HC2H3O2 (0.10 M) with a Ka value of 1.8x10-5. What is the pH? What does the addition of NaC2H3O2 do to the equilibrium you just represented? Use these chemical principles to explain why the first pH recorded from the simulation is so much larger than your calculated value for pure 0.10 M acetic acid.

b. How do the pKa values for acetic acid and ammonium ion (Ka = 5.6x10-10) compare to the pH values for the first two buffers in the simulation? Explain this phenomenon.

c. Why don’t the pH values change when the component concentrations are increased? 4. Write a general equation to show how a buffer containing an acid HA and the salt of its

conjugate base, NaA, would respond to the addition of each of the following. a. The strong acid, HCl b. The strong base, NaOH

5. Why are the pH changes so noticeable with the last two additions of strong acid and base in the simulation?

6. Which of the two concentration combinations would be most effective for an antibiological agent's buffer? Explain.

7. Obtain a "Mission Card" from your teacher. Develop a procedure using the chemicals and materials listed above to create the 100 mL of the desired buffer. Submit your procedure to your teacher at least 24 hours prior to the lab.


50

AP Chemistry Procedure The company contracting your class has requested that each group prepare a 100 mL sample of their buffer and test it in the hope that it falls within 0.5 of the target pH stated on the "Mission Card" and that a 50 mL sample can maintain a relatively constant pH (within one pH unit of the initial value) with the addition of up to 20 mL of 0.2 M HCl or NaOH.

1. Neatly record your hypothesis in an “If we combine … then our buffer will …” format. The procedure, including the masses and volumes of all materials used, should be neatly recorded in your lab notebook. Use a data table for all quantitative values. Be sure to show all supporting calculations in an adjacent column or section of your lab report.

2. Once your teacher has given you permission (you may have to modify your hypothesis before you are allowed to proceed), go ahead and test your hypothesis. First check your target pH by placing 50.0 mL of your buffer into a 250 mL beaker and lowering the pH meter’s electrode or probe into the sample. Record the initial pH in the first line of a table with columns headed “Volume of Acid Added” and “pH.”

3. Each lab group must follow the same procedure for capacity testing. Measure a second 50.0 mL portion of your buffer. Why might your total volume of buffer slightly exceed 100 mL? All measurements involve some degree of uncertainly. As the company is allowing you to be within plus or minus 0.5 units of your target pH, do you think a slight increase in volume is significant in this case?

4. The purpose of this step is to generate data to produce as smooth a curve as possible, adding smaller amounts of acid when the pH changes noticeably. Begin your capacity testing by using a buret to add about 5.00 mL (record the actual ongoing volume precisely) increments of 0.20 M HCl to your buffer. Use your table to record the pH after each addition (followed by thorough stirring). Once about 10.00 mL of acid has been added, change to approximately 2.00 mL increments until you reach a total volume of around 16.00 mL of acid added. Now add the acid in approximately 1.00 mL increments until you’ve reached a total volume of near 30.00 mL of acid added. You may now increase the volume of each increment to about 2.00 mL until you reach a total of 40.00 mL added. The last couple of increments may be about 5.00 mL each.

5. Repeat your capacity testing for base addition by following the same procedure with the addition of 0.20 M NaOH to your second 50.0 mL sample of buffer. The volumes added should follow the same pattern as before, but base is being added. These precise volumes and pH values may be recorded in two new columns added to the data table you’ve already prepared.

6. Two graphs will be plotted, one of “pH vs. Volume of Acid Added to a Buffer” and one of “pH vs. Volume of Base Added to a Buffer.”

Data A data table will be developed according to your procedure for creating the buffer. You will also need a data table for Volume of Acid Added vs. pH and Volume of Base Added vs. pH.


51

AP Chemistry Analysis

1. Complete neatly labeled graphs for the data tables you've prepared. 2. Calculate the anticipated pH for the buffer you were assigned. Include neat calculations in your

report. 3. Determine the volume of acid and base your buffer was able to neutralize before allowing a pH

change of more than one unit. Use the volume with the HCl/NaOH molarities to determine the number of moles of acid and base your buffer was able to neutralize.

4. Consider 50.0 mL of the buffer you prepared and calculate the ratio of (moles of added acid neutralize) to (moles of base component in the buffer). Do the same calculation for (moles of added base neutralized) to (moles of acid component in the buffer). State both ratios in the form of x:1. Report the value for x to two significant figures.

5. Upload your graphs to the common Google drive folder. Make sure to label your graphs with appropriate descriptions according to the mission assigned.

6. Find two other groups that attempted the same buffer solution. Make an assessment of the capability and effectiveness of your buffer versus that of the other two. Acting as a "spokesperson" for our class' work, write a paragraph in which you determine which buffer solution would be best to use for the assigned biochemical project. Give adequate justification for your choice.

7. Suppose, during preparation, an additional 10 mL of distilled water was added to your buffer by mistake.

a. What effect would this have on your buffer's pH? Explain. b. Would this affect your buffer's capacity? Explain.

8. Given a solution of hydrocyanic acid (HCN), what additional reagent or reagents is/are needed to prepare a buffer from the hydrocyanic acid solution?

a. Explain how this buffer solution resists a change in pH when moderate amounts of strong acid are added. Use a chemical equation in your explanation.

b. Explain how this buffer solution resists a change in pH when moderate amounts of strong acid are added. Provide an equation with your explanation.


52

AP Chemistry

What's in That Bottle? – An Investigation of Bonding

There are few greater potential hazards around the laboratory than that of unmarked or improperly labeled chemicals. Many schools house unused, unlabeled, and improperly stored chemicals. These chemicals can pose a risk to humans and the environment. All chemicals must have complete identification securely fastened to their containers. Chemicals of unknown stability and those that deteriorate over time should have a preparation date clearly indicated on the label. The chemistry storerooms in schools need to be cleaned out periodically and chemicals properly disposed of. Most chemicals should not be flushed down the drain or thrown into the garbage. Proper disposal of chemicals is costly. If the identity of a substance is not known due to poor labeling or lack of a label, the cost of proper chemical disposal can increase. There are many accidents associated with chemicals that are thrown out and inadvertently mix. Disposal of unlabeled bottles is dangerous and therefore very expensive and closely regulated by law. The purpose of proper labels is to: indicate the source, supplier or manufacturer of the chemicals, the production date, CAS (Chemical Abstract Service) number of the chemical, and to warn of possible hazards. The MSDS provides personnel with procedures for handling and cleaning up each substance in a safe manner along with details on their physical and chemical properties and toxicity. There is a problem in the chemical storeroom. The high humidity in the storeroom caused the labels on some of the chemical bottles to fall off. The labels are lying all over the shelves and it is the student’s job, as a chemistry intern, to design a method that will help identify the chemicals so the labels can be put onto the correct bottles. There are at least four unlabeled bottles that represent at least one of each type of bond. The unlabeled chemicals are all solids but may be ionic compounds, nonpolar or polar covalent compounds, or metals. If the type of substance, or, even better, the identity can be determined, disposal will be less costly to the school. Once the properties of the unknown compounds are determined, students will be given information that can help identify the name of each chemical within the unlabeled bottles. Purpose

• To determine the type of bonding in unlabeled chemicals using physical and chemical properties of substances with ionic, molecular, and metallic bonds.

Safety Precautions • Safety goggles should be worn at all times in the laboratory. • Be cautious of acidic and basic solutions since they can cause skin burns and eye damage. • Liquids and solids are to be disposed of in properly labeled waste containers per MSDS

guidelines.

1 Bonding

53


The list of potential unknown solids is given in the table below:

Ammonium chloride Magnesium oxide Benzoic acid Aluminum Wax/paraffin Magnesium Calcium carbonate Potassium nitrate Urea (NH2)2CO Calcium Iodine (I2) Zinc Copper (II) sulfate anhydrous Sodium carbonate Sucrose Copper Sodium acetate Sodium hydrogen carbonate Copper (II) sulfate pentahydrate

Sodium chloride Salicylic acid, C6H4(OH)COOH

The list of materials available for testing is given below:

95% ethanol Ice Magnifying lens Capillary tube Beaker, 100 mL Test tubes and rack Hexanes Phenolphthalein Ring stand Thermometer clamp Thermometer Conductivity meter (metals and

aqueous only) Distilled water pH paper Wooden splints Rubber band Canned food lid 0.1 M NaOH Magnet Cotton swabs Tongs 0.1 M HCl Hot plate Wire gauze Sandpaper Universal indicator

Pre-Lab Answer questions 1-2 using Table 1.

Compound Observations MP (°C) Solubility in 25°C Water

Types of Elements

Type of Bond

Potassium chloride

White solid 993 Yes Metal/Non-metal

Ionic

Sucrose White solid 186 Yes Nonmetal Polar covalent Iodine Dark gray solid 114 Slightly soluble Nonmetal Nonpolar

covalent Zinc Gray, shiny metal 1535 No Metal Metallic

1. Compare the type of bond with regard to the properties below using Table 1 and explain any

relationships. HINT: Think of what is happening between the bonded atoms as well as what occurs between the particles.

a. Melting point b. Solubility in 25°C water

2 Bonding

54

AP Chemistry

2. Predict the properties of each substance below based on Table 1. Compound Bond Type Relative Melting Point

(High or Low) Solubility in Water

Hexane (C6H6)

Bromobenzene (C6H5Br)

Sodium chloride

Iron

Procedure Your lab group will receive a selection of 4 unknown solids. You must choose at least four different tests to study physical and chemical properties of the substances from the list below:

• Color • Solubility in water • Conductivity of the solid • Conductivity in water • pH of the solution in water • Solubility in ethanol

• Solubility in hexanes • Relative melting point • Reaction with 0.1 M HCl • Reaction with 0.1 M NaOH • Magnetism

Using the four (or more) tests that you have chosen conduct the experiment and record your qualitative and quantitative results in a data table. Data A data table will be developed according to the tests chosen to conduct on the four substances. Analysis

1. Identify the bond type in each unknown. 2. To what extent do you believe the classification of your unknown is reliable? Justify your claims

with evidence. 3. What were the two most significant tests done to identify each of the types of bonds? 4. How do the melting points of ionic compounds compare to those of covalent compounds?

What evidence from the investigation supports your conclusion? 5. When the solids were placed in water were all the results the same? What types of solids

conduct electricity in water? 6. Metal oxides dissolved in water show a pH in what range? In contrast to these metal oxides, do

nonmetal oxides produce the same pH range? 7. Wax is a saturated hydrocarbon, a covalent compound. Wax is not soluble in water yet sugar is

also a covalent compound and is soluble in water. Look at the structure of both compounds and explain what could justify these results.

3 Bonding

55

AP Chemistry Sticky Question: How Do You Separate Molecules That Are Attracted to One Another?

There are two phases in paper chromatography, a stationary phase (the paper) and a mobile phase (the solvent). A molecule can have a greater affinity for either the paper or for the solvent. The filter paper is made of cellulose, a polymer. Cellulose will attract water molecules to the exposed hydroxyl (OH–) groups along the polymer. This interaction makes a thin layer of water on the paper that competes for the attraction of the molecules being separated. Alternately, the molecule can be attracted to the solvent and travel with the solvent up the paper. When doing chromatography, a small amount of solvent is placed in a sealed container. The mixture being separated is put on a piece of paper, the starting point is marked, and the paper is put into the solvent. The container must be sealed so the solvent saturates the paper and does not evaporate first. The level of separation is measured by a ratio that compares the distance that the molecule travels to the distance the solvent travels. This ratio is called the Rf value. To get the Rf value, the experimenter must identify the distance that the solvent traveled on the paper and measure the distance. Second, the experimenter must identify the distance that each molecule traveled and measure that distance. It is best to run the test more than once to reach the best separation values possible. The Rf value is a ratio of the distance of the molecule divided by the distance of the solvent. The greater the distance the molecule travels, the greater its affinity for the solvent and the greater the Rf value. In this experiment, you will have a choice of different solvents to use. When you propose the best solvent for separation of the mixture, you will also need to evaluate it in terms of “greenness.” In modern chemistry, chemists use principles of green chemistry to evaluate the solvents that are used in a chemical process for their level of toxicity to humans and the environment. Solvents are also evaluated in terms of their life cycle or how long the molecule remains in the environment and if the molecule breaks down to become more benign or more toxic. The overall focus of green chemistry is to be more efficient in chemical production, producing less waste, using fewer toxic molecules, and producing waste that biodegrades and does not pose a risk to the environment. [See Tables 1 and 2 in the GSK solvent selection guide, available at http://pubs.rsc.org/en/content/articlelanding/2011/gc/c0gc00918k under “supplementary information,” for an evaluation of solvents based on these guidelines. To understand the differences between solvents, look up the green rating for hexane (found in chromatography solvent) and compare it to 2-propanol.] The food dyes that are in the mixture have their own green chemistry issues. For example, the molecules used may have a life cycle that is longer than previously anticipated and possibly increased toxicity. When scientists evaluate the toxicity of molecules based on the experimental data, efforts to understand the origins of toxicity often look at the molecular structure of the substance. A key strategy for looking at molecular structure is to identify functional groups that are present. There are specific functional groups that are known to create toxic by-products when they are metabolized in the human body, such as acetaminophen that can be converted to N-acetyl-p-benzoquinone imine. All three food dyes used in this lab are azo dyes, which means that they contain a double-bonded nitrogen connecting multiple aromatic carbons. While the molecules resemble one another, only red #40 has been linked to

1 Chromatography

56

AP Chemistry allergic reactions in some people, but the FDA has not found conclusive evidence that such a dye is unsafe. In Europe these food dyes are not used and natural pigments are used instead.

Figure 1: Molecular Structure of Food Dyes

You are working for a crime lab and a chemical residue has been turned in for analysis. To identify the chemicals in the residue, you will need to separate them from the mixture and identify them individually. Another lab technologist has made an attempt to separate the molecules but was not as successful as the boss would like. There was only one molecule separated from the mixture, but the boss suspects that there are at least three different molecules. Science is often a process, where a method is tried and then modified for a second attempt. Your job will be to propose a modification and attempt to improve the separation attained. Purpose

• To develop a method to separate three similar molecules.

Safety Precautions • Safety goggles should be worn at all times in the laboratory. • If solutions are spilled on skin, wash with copious amounts of water.

Materials The list of materials and solvents available for separating the mixture is given below: Materials Solvents Pencil Chromatography solvent

(petroleum ether and acetone mixture) Cylindrical glass containers with lids Distilled water Metric rulers (mm) Ethanol (C2H5OH) 100 mL graduated cylinders 2-propanol ((CH3)2CHOH) Labels Acetone ((CH3)2CO) Chromatography paper

2 Chromatography

57

AP Chemistry Pre-Lab

1. How can molecules attract each other when they are in a mixture? 2. Sketch a picture showing how blue dye #1 might interact with ethanol. (Structure shown

below.) 3. What does the Rf value describe on a microscopic level? Why is this important? 4. If the molecule had a very high affinity for the stationary phase, how would this affect the Rf

value? Explain. 5. Develop the procedure you and your lab group will follow to separate the mixture of food dyes.

Procedure You will design an experiment to test the solvents that you believe will provide the best separation of the three food dyes in Figure 1. There are five solvents listed as available for you to use in this experiment. Develop a hypothesis as to what could lead to an effective separation of the three food dyes in the sample your teacher has provided. Include intermolecular attractions in your hypothesis.

Figure 2: Molecular Structure of Typical Solvents

You will be provided with a maximum of 20 mL of each solvent that you choose. You must test at least two of the solvents provided. The molecular structures of the solvents are given in Figure 2. Record all qualitative observations from your investigation. Also, be sure to indicate necessary measurements for finding Rf values of each dye in your chosen solvents in your procedure. The mixture of food dyes can be disposed of down the sink with the "used" distilled water. The other solvents should be collected and disposed of in an organic waste container Data Create a data table(s) to record values necessary for finding Rf values of each dye in each solvent tested.

3 Chromatography

58

AP Chemistry Analysis

1. Calculate the Rf value of each dye in each solvent chosen. Include these in your data table. 2. Why did you select the solvents that you tested? Did your data support your hypothesis or

disprove your hypothesis? 3. What explanations can you provide for your separation of the three molecules? How was the

choice of the solvent connected to the process? 4. What part of the chromatography setup did the molecules interact with, stationary or mobile

phase? How would you explain this interaction using intermolecular forces? 5. Draw a picture of how the chromatography worked. Explain your picture using the following

terms: stationary phase, mobile phase, and intermolecular forces. 6. Evaluate which solvent is the one with the best "green chemistry" rating. What IMF would this

solvent form with the three molecules in the mixture? 7. Which molecule spent the most time in the stationary phase and why?

4 Chromatography

59

AP Chemistry

Determination of the Percent Water in a Compound and Its Empirical Formula

Introduction The polarity of the water molecule, which makes it a great solvent for ionic compounds, causes

water molecules to cling to the structure of solid substances. When this occurs, the trapped water molecules are called water of hydration and they become an integral part of the crystal structure.

There are many compounds that have a tendency to absorb water vapor from the air. These compounds are said to be hygroscopic, and can be used as moisture-reducing agents. Other compounds absorb such large quantities of water vapor that they will actually dissolve in their own water of hydration, a property known as deliquescence.

In this experiment, you will test a hygroscopic ionic compound containing copper, chlorine, and water molecules locked in the crystal structure of the solid to determine its water of hydration. The general formula for the compound is CuxCly 2

establish the proper chemical formula for this substance. Although the water molecules are securely attached to the ionic solid that you will test, they are susceptible to removal by heat. You will gently heat a sample of the compound to drive off the water of hydration. By measuring the mass of the sample before and after heating, you can determine the amount of water in the sample and calculate its water of hydration.

In the second part of this experiment you will conduct a chemical reaction with the dried sample, which will produce elemental copper. By measuring the mass of copper that forms, you will have the necessary information to determine the moles of copper and chlorine in your sample, and you will be able to establish the proper empirical formula. developed and wrote the modern atomic theory at the turn of the 19th century (documents point to 1803). Proust. A fundamental component of the modern atomic theory is that the mole ratios of elements in a compound will be small whole numbers (law of definite proportions). The whole number mole ratio is commonly referred to as the empirical formula of a compound.

One of the challenges in finding the proper chemical formula for a compound is that there may be more than one plausible mole ratio for the elements in that compound. Dalton called this the law of multiple proportions. For example, if you were testing a compound that contained iron and sulfur, the plausible chemical formula could be FeS or Fe2S3

mass of sulfur present in a given mass of the compound, you will be able to establish the true chemical formula of the compound. Objectives In this experiment, you will

• Carefully heat a measured sample of a hygroscopic ionic compound. • Determine the water of hydration in a copper chloride hydrate sample. • Conduct a reaction between a solution of copper chloride and solid aluminum. • Use the results of the reaction to determine the mass and moles of Cu and Cl in the reaction. • Calculate the empirical formula of the copper chloride compound.

1

Percent Water and Empirical Formula

60


Pre-Lab Questions

1. What is the purpose of preheating the crucible and its cover prior to measuring its mass?

2. Washing soda is a hydrated compound whose formula can be written Na2CO3 · 2O, the number of moles of 2O per mole of Na2CO3. When a 2.123 g sample of washing soda was heated at 130°C, all of the water of hydration was lost, leaving 0.787 g of anhydrous sodium

3. A piece of iron weighing 85.65 g was burned in air. The mass of the iron oxide produced was 118.37 g.

a. Use the molar mass of iron to convert the mass of iron used to moles.

b. According to the law of conservation of mass, what is the mass of oxygen that reacted with the iron?

c. Calculate the number of moles of oxygen in the product.

d. Use the ratio between the number of moles of iron and number of moles of oxygen to calculate the empirical formula of iron oxide. Note: Fractions of atoms do not exist in compounds. In the case where the ratio of atoms is a fractional number, such as ½, the ratio should be simplified by multiplying all the atoms by a constant to give whole number ratios for all the atoms

½ should 2O).

Materials Equipment

Crucible with cover Crucible tongs Spatula Ring stand, ring, and clay triangle Bunsen burner Büchner funnel and filter flask

Filter paper Watch glass Balance Stirring Rod Drying oven Wash bottle

Chemicals Unknown solid copper chloride hydrate Aluminum wire, 20 gauge

M

Ethanol solution, 95% Distilled water

_____________________________________________________________________________________ Safety Precautions Hydrochloric Acid is highly toxic by ingestion or inhalation; severely corrosive to skin and eyes. Hazard Code: A—Extremely hazardous. Copper chloride hydrate is highly toxic by ingestion and inhalation. Hazard Code: B—Hazardous. _____________________________________________________________________________________

2


61


Procedure

1. Obtain and wear goggles.

2. M clean, preheated crucible and cover. Obtain about 1 g of the unknown copper chloride hydrate and place it in the crucible. Use a spatula to break up any large

Mthe exact mass of the crucible with compound and its cover.

3. Set up a ring stand, ring, and clay triangle for heating the sample. Rest the crucible on the clay triangle with the cover slightly tilted so that vapor may escape. Set up a lab burner and ignite the burner away from the crucible. Adjust the burner to get a small flame.

4. gently heat the sample. Do not overheat the compound. Note the color change, from blue-green to brownish, as the water of hydration is driven out of the crystals. When the sample has turned brown, gently heat the crucible for two more minutes.

5. Remove and turn off the burner. Cover the crucible and allow the sample to cool for about ten minutes. M e crucible, crucible lid and sample.

6. Reheat the crucible, crucible lid and sample until constant mass is achieved. Record the final mass.

7. Once the sample has achieved constant mass, tbeaker. Rinse out the crucible with two 8

solution is green as the copper ions are rehydrated.

8. Msolution so that it is completely immersed in the copper chloride solution. Note that the reaction produces a gas, elemental copper is forming on the surface of the aluminum wire, and the color of the solution is fading. The reaction will take about 30 minutes to complete. parts a and b while you wait.

9. When the reaction is done, the so MM

are. It can cause painful burns if it comes in contact with the skin.

10. Use a glass stirring rod to scrape off as much copper as possible from the Al wire. Slide the wire up the wall of the beaker and out of the solution with the glass stirrer and rinse off any remaining copper with distilled water. If any of the copper refuses to wash off the aluminum wire, wash it with

M leaving the solid copper in the beaker.

3


62


11. Collect and wash the copper produced in the reaction.

a. Set up a Büchner funnel for vacuum filtration. b. M

the funnel. Start the vacuum filtration. c. Use small amounts of distilled water to wash all of the copper onto the filter paper on the Büchner funnel. Use the glass stirring rod to break up the larger pieces of copper.

d. Wash the copper twice more with small amounts of distilled water.

12. Turn off the suction on the filter paper and let it sit for about 1 minute. Turn the suction back on and let the vacuum filtration run for about five minutes.

13. M ry watch glass. Transfer the copper to the watch glass. M

14. Dry the watch glass of copper under a heat lamp or in a drying oven for five minutes. Remove the watch glass and allow it to cool. When the watch glass is cool enough to touch, measure the mass of the watch glass plus copper. Repeat the drying and weighing of the copper until you achieve a constant mass. Record the final mass.

15. Dispose of the copper, aluminum wire, and filtered liquid as directed.

Data Tables Data Table 1

M

M

M

M

M

M

M

M

M

4


63


Post-Lab Questions and Analysis

1. Why must objects be cooled before their mass is determined on a sensitive balance?

2. copper chloride hydrate?

3.

4.

5. Write the proper chemical formula and name for the compound that you tested.

6. Use stoichiometry to calculate the theoretical yield of copper in this experiment based on the initial mass of your sample.

7. Calculate the percent yield of copper actually produced in this experiment.

8. A student fails to place the lid on the crucible during the initial heating of the hydrated sample and some of the solid spatters out. What effect does this error have on the calculated mass of the water lost by the hydrate? Justify your answer.

5


64

AP Chemistry

Molar Mass by Freezing Point Depression

Introduction A procedure for determining the molar mass of a substance is very useful to chemists. The molar mass is an important value that must be known in order to identify an unknown substance or to characterize a newly prepared compound. Concepts � Molality � Freezing point depression

� Colligative properties � Molar mass

Background There are a number of ways of determining the molar mass of a substance. One of the simplest involves finding the change in the freezing point of a solvent when an unknown substance is dissolved in it. The change in freezing point is directly proportional to the molality of the solution. This change in freezing point is one of several “colligative” properties of solutions – properties that depend only on the number of dissolved particles in solution, and not on the type of particle. Other colligative properties include changes in boiling point, vapor pressure, and osmotic pressure. Measurements of these properties can be used to find the molar mass of solutes. The molality of a solution, m, is defined as the moles of solute divided by the kilograms of solvent:

(solvent) kg

(solute) moles=m Equation 1

Since the moles of solute is the same as the grams of solute divided by the molar mass of the solute, then:

(solute) massmolar (solvent) kg

(solute) g=m Equation 2

The relation of molarity to change in freezing point is: mfpfp kT = Equation 3

fp is the change in freezing point of the pure substance versus the solution, kfp is the freezing point depression constant for the solvent, and m is the molality of the solution. The value of kfp must be determined for each solvent. Equations 2 and 3 are combined to solve for the molar mass of the solute.

fp

fp

T (solvent) kg(solute) gk

(solute) massMolar = Equation 4

The solvent used in this experiment is a nonpolar solvent with the common name butylated hydroxytoluene. This compound is abbreviated BHT and is frequently used as an antioxidant in foods. The IUPAC name for the compound is 2,6-di-tert-butyl-4-methylphenol. The freezing point of BHT is approximately 70°C. If the freezing points are determined for both the solvent and the solution using a thermometer with minor scale divisions marked every 0.1°C, the freezing points can be estimated in the range: ±0.02°C. Figure 1 shows cooling curves obtained for both a pure solvent and for a solution. Notice that supercooling occurs in both the solvent and the solution. When supercooling occurs, the temperature falls below the freezing point until the first crystal forms. The temperature then rises up and either

1 Molar Mass by Freezing Point Depression

65


stays at the freezing point, in the case of the pure substance, or slowly falls as the solution freezes. The freezing point temperature, Tf of the solution is extrapolated from the graph.

Figure 1 Freezing Point Graph for Pure Solvent and for Solution

Experiment Overview

The purpose of this experiment is to determine the molar mass of an unknown substance by measuring the freezing point depression of a solution of the unknown substance and BHT. The freezing point of BHT is first determined. Even though the freezing point of butylated hydroxytoluene is known, it is necessary to determine it with the thermometer that is used in the experiment. Thermometers can give temperature readings that are slightly different from true values. Even if the thermometer reading is slightly off, the change in temperature should be accurate. It is important that the same thermometer is used to determine both the freezing point temperature of the solvent and that of the solution.

A known amount of stearic acid is then added to a measured quantity of BHT. The freezing point depression of this solution is found and the freezing point depression constant (kfp) is calculated. The unknown is added to BHT, the freezing point depression of this solution is measured, and the molar mass of the unknown is then determined. Pre-Lab Questions

1. The following data were obtained in an experiment designed to find the molar mass of a solute by freezing point depression

Solvent: para-dichlorobenzene Freezing point depression constant: 7.1 °C/m Freezing point of pure solvent: 53.02 °C Mass of para-dichlorobenzene: 24.80 g Mass of unknown substance: 2.04 g Freezing point of solution: 50.78 °C

a. fp. b. Using Equation 4, calculate the molar mass of the unknown substance.

2. The following errors occurred when the above experiment was carried out. How would each

affect the calculated molecular mass of the solute (too high, too low, no effect)? Explain your answers.

a. The thermometer used actually read 1.4°C too high. b. Some of the solvent was spilled before the solute was added. c. Some of the solute was spilled after it was weighed and before it was added to the

solvent. d. Some of the solution was spilled after the solute and solvent were mixed but before the

freezing point was determined.

Molar Mass by Freezing Point Depression 2 66


Materials Chemicals

2,6-di-tert-butyl-4-methylphenol, BHT, 16 g Stearic acid, CH3(CH2)14CH2OH, 1g

Unknown substance, 1g

Equipment

Balance Beaker, 400 mL Hot plate or Bunsen burner Ring clamp Wire gauze Ring stand Split rubber stopper with one hole

Test tube, 18 x 150 mm NOVA Temperature probe Universal clamps, 2 Weighing dish or paper Wire stirrer

______________________________________________________________________________ Safety Precautions 2,6-Di-tert-butyl-4-methylphenol, BHT, is moderately toxic by ingestion and inhalation and is a body tissue irritant. Stearic acid and the unknown substance are slightly toxic by ingestion and are body tissue irritants. Wear chemical splash goggles and a protective apron. Wash hands thoroughly with soap and water before leaving the laboratory. _____________________________________________________________________________________ Procedure

1. Assemble the apparatus as diagrammed in Figure 2. Do not add water to the beaker. Clamp the temperature probe using the split rubber stopper. Do

not seal the test tube with the stopper – it is just to support the temperature probe. Make a stirrer out of wire bent with a circle at the bottom. The test tube is clamped in the beaker so that the solid it contains will be below the level of the water in the beaker. The beaker rests on a hot plate.

2. Disassemble the apparatus by sliding both the thermometer and stirring wire assembly and the test tube clamp off the ring stand. Weigh the test tube on an analytical balance. Record the mass in Data Table 1.

3. Accurately measure about 8 g of BHT into the test tube. Record the combined mass of the test tube and the BHT in Data Table 1.

4. Clamp the test tube in the beaker and insert the thermometer and stirring wire assembly into the test tube and clamp the assembly. Do not force the thermometer into the solid – allow it to rest on top of the solid. Add water to the beaker so that the solid in the test tube is well below the level of the water.

5. Turn on the hot plate and heat the water bath to about 90°C. 6. All the BHT in the test tube to melt.

Figure 2 Diagram of Apparatus for Freezing Point Determination



7. When the temperature of the BHT is 80 °C or above, remove the thermometer and test tube from the water bath. First raise the thermometer clamp so that the temperature probe is slightly higher than the height of the water bath beaker. Next, raise the test tube clamp so that the temperature probe is still positioned correctly in the test tube and the test tube clears the water bath beaker. (See Figure 3.)

8. Record the BHT temperature in Data Table 2 every 20 seconds as the melted BHT cools. It is important to continuously stir the BHT to maintain even cooling. Stirring also helps prevent supercooling. Stir until BHT solidifies.

9. Continue recording temperature values in Data Table 2 until at least 5 values are constant. Make a note of the temperature at which crystals begin to form. Note: After the temperature holds constant, it will then begin to drop again. This drop indicates the substance is frozen and the experiment is over.

10. If instructed, repeat this measurement. 11. Using an analytical balance, accurately measure about 1 g

of stearic acid onto weighing paper and record its exact mass in Data Table 1.

12. Place the stearic acid into the test tube containing BHT. 13. Clamp the test tube in the water bath and insert the temperature probe and stirring wire

assembly into the test tube (see Figure 2). 14. Heat the mixture in the hot water bath until the substances are all melted. Stir well to ensure

the solution is homogeneous. 15. When the solution temperature reaches 80°C or higher, remove the test tube from the hot

water bath. Stir the solution and record the temperature every 20 seconds as the solution cools. Record at least six temperature values after crystals first begin to form.

16. If instructed repeat this measurement. 17. Repeat steps 1-3 and 11-15 using fresh BHT, a clean test tube, thermometer, stirrer, and about 1

g of the unknown compound in place of 1 g of stearic acid. Repeat this measurement, if instructed.

Lab Quest Directions Set the rate to “Every 10 seconds” and the samples to at least 500. Disposal and Cleanup Re-melt all substances in the test tube and dispose of them in the waste container provided. Clean out the test tube and beaker. Return all borrowed equipment and store all other equipment.

Figure 3



Data Tables Data Table 1

Trial 1 Trial 2 Mass of empty test tube #1, g

Mass of test tube #1 plus BHT, g

Mass of BHT, g

Mass of weighing paper, g

Mass of weighing paper plus stearic acid, g

Mass of stearic acid, g

Mass of empty test tube #2, g

Mass of test tube #2 plus BHT, g

Mass of BHT, g

Mass of unknown, g

Data Table 2 Temperature in °C

Time, in seconds Pure BHT BHT + Stearic Acid BHT + Unknown

Data Table 3. Calculation Table Pure BHT BHT + Stearic Acid BHT + Unknown

Freezing Point, °C

ΔTfp, °C --

Kfp, °C/m -- --

Molar mass, g/mol -- --

Post-Lab Calculations and Analysis

1. Graph the cooling data as shown in Figure 1 using a graphing utility (Excel). From your graph, fp.

Calculate the freezing point depression constant from the molal fp of the stearic acid solution. Calculate the molar mass of the unknown solute using the freezing point depression constant, the mass of the unknown solute, and the mass of BHT. Enter these values in the Calculation Table (Data Table 3).

2. What was the least precise measurement in the experiment? How does this limit your significant digits?

3. Did your solutions show any evidence of supercooling? 4. Why is it advantageous to choose a solvent that has a large value for kfp? 5. Explain why the pure solvent shows a level horizontal curve as solidification occurs, but the

curve for the solution slopes downward slightly.


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