AP Chemistry. Applying Green Chemistry to Purification. Introduction

Transcription

1 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 (Na 2 CO 3 ) and sodium bicarbonate (NaHCO 3 ). 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. Actual Yield Percent Yield = x100 Theoretical Yield 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 Applying Green Chemistry to Purification 1 1

2 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. Mass of desired product Percent Atom Economy = x100 Mass of all products produced 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) sodium carbonate (s) + carbon dioxide (g) + water (g) Pre-Lab Review the 12 Principles of Green Chemistry at 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) 2Cu(s) + CO 2 (g) b. CuO(s) + CO(g) Cu + CO 2 (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? Applying Green Chemistry to Purification 2 2

3 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 you about "greenness"? c. What is another possible consideration from the principles of green chemistry that could 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 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. Applying Green Chemistry to Purification 3 3

4 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 NaHCO 3 and Na 2 CO 3 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 (NaHCO 3 ) 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 to relieve swelling or itchiness. Epsom salts is a hydrate, which means that a specific ratio of water molecules per unit of salt formula regular repeats in the crystalline structure of Epsom salts. The formula for Epsom salts can be written as MgSO 4 xh 2 O, where the ratio of water molecules per unit of salt formula is x:1. When hydrated Epsom salts is heated to at least 250 C, all of the waters of hydration are lost, according to the reaction: a. When g of Epsom salts were heated to constant mass of 250 C, g of MgSO 4 powder remained. What is the value of x? b. If anhydrous MgSO 4 is the desired product, what is the atom economy of this reaction? Applying Green Chemistry to Purification 4 4

5 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 A = abc 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 (λ max ) 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. Relationship Between Concentration and Transmitted Light 5 1

6 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 be Table 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 in the 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 and label your graph well. Make sure your trendline and equation are shown on the graph and print it out. Concentration (M) Absorbance 6.40x x x x x Use your trendline from question #4 to determine the concentration of a blue solution that has a recorded absorbance value of 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. Relationship Between Concentration and Transmitted Light 6 2

7 Data 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 exactly what is in it. 4. Using the same pipette, add 4 ml of stock solution to the graduated cylinder. Using a separate pipette, add 1 ml of distilled water to the graduated cylinder with the stock solution. Use a stirrer to mix the solution. 5. Pour the 4:1 solution into a cuvette, clean the sides of the cuvette with a towel, and record its absorbance 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 H 2 O b. 2 ml stock to 3 ml H 2 O c. 1 ml stock to 4 ml H 2 O 7. Obtain a small sample of the Gatorade. Record the volume of the Gatorade bottle. Fill a cuvette and record the absorbance in Table 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 Table Dispose of all solutions down the drain. Rinse and dry all equipment and return. 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 a percentage 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 the concentration of the Gatorade. 4. Calculate the percent error between your visual estimation of concentration of the Gatorade and your calculated value. 5. Using your Gatorade concentration and the volume of the bottle it came in, determine the number of grams of blue dye #1 in the Gatorade bottle. Relationship Between Concentration and Transmitted Light 7 3

8 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) + NO 3 - (aq) Cu 2+ (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 NO 2 (g) according to the equation: 2NO(g) + O 2 (g) 2NO 2 (g) The Cu 2+ ions in the unknown aqueous solution form the complex ion, [Cu(H 2 O) 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(NO 3 ) 2. Purpose Determine the mass percent of copper in a sample of brass. Safety Precautions Concentrated nitric acid is corrosive and will attack and destroy metals, proteins, and most plastics. Avoid skin contact and neutralize any spills with baking soda, then rinse with copious amounts of water. The acid will discolor the skin for days after contact so be sure to wear gloves. The NO gas that forms quickly oxidizes in air to produce a toxic, reddish-brown gas of NO 2. Avoid inhalation. Perform the reaction under a fume hood. Be sure to wear splash-proof goggles and an apron at all times. Materials Balance Beaker Concentrated nitric acid (15.8 M) Distilled water 100 ml volumetric flask 1-2 g sample of brass Remainder dependent on procedure Mass Percent of Copper in Brass 8 1

9 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 the lab in order to accomplish this so design your procedure with these items in mind: M Cu(NO 3 ) 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 of copper (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 g. Record this value. Place the sample in a small beaker. Label your beaker so that you can identify your group's materials among the rest of the class. 2. Assuming your brass sample is 100 percent copper by mass, calculate the minimum volume of concentrated 15.8 M HNO 3 (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 15.8 M HNO 3 (aq) so that the acid is in excess, and then your teacher will cover the beaker with a watch glass. 3. While the metal is reacting with the acid, perform your predesigned procedure for gathering a calibration 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 will remove 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. Dilute to a final volume of ml. The excess nitric acid will dissolve the zinc and copper metals in the brass. 5. Use the remainder of your predesigned procedure to determine the concentration of the copper(ii) nitrate in the unknown brass solution. Mass Percent of Copper in Brass 9 2

10 Data 6. The remaining brass solution should be neutralized by adding small amounts of solid baking soda until the bubbling has subsided and the ph is 5-9. Then dispose of the waste in the waste beaker provided. 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 sure to label the graph well and include any linear regressions used. 2. Determine the molarity of the Cu(NO 3 ) 2 (aq) found in ml of the brass solution using your colorimeter analysis. Support your answer with the appropriate calculations. 3. Determine the mass of Cu dissolved in the brass solution based upon the molarity calculated in Step 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 volumetric flask, what effect would this have on the calculated mass percent of copper in the solution? Justify your answer. Mass Percent of Copper in Brass 10 3

11 Determination of Concentration by Oxidation-Reduction Titration of Hydrogen Peroxide Introduction One method of determining the concentration of a hydrogen peroxide, H 2 O 2, solution is by titration with a solution of potassium permanganate, KMnO 4, 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 H 2 O 2 (aq) + 2 MnO 4 (aq) + 6 H + (aq) 5 O 2 (g) + 2 Mn 2+ (aq) + 8 H 2 O (l) In this experiment, the volume of KMnO 4 titrant used at the equivalence point will be used to determine the concentration of the H 2 O 2 solution. Your sample of H 2 O 2 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 hydrogen peroxide 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 1 11

12 Pre-Lab Questions 1. Write the oxidation and reduction half-reactions for the redox reaction taking place in this laboratory. 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 your answer. 4. Calculate the number of moles of MnO 4 in 20.1 ml of M potassium permanganate solution. 5. Calculate the number of moles of hydrogen peroxide needed to react with 20.1 ml of M potassium permanganate solution. 6. The introduction states that the hydrogen peroxide solution sold in the grocery store is 3% (by mass). Assuming a density of 1.0 g/ml, calculate the molarity and mole fraction of 3.0 % H 2 O 2. 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, H 2 O 2, 0.3% Potassium permanganate solution, KMnO 4, M in H 2 SO 4 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. Determination of Concentration by Oxidation-Reduction Titration of Hydrogen Peroxide 2 12

13 Procedure 1. Measure out precisely 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 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, H 2 SO 4, solution from the dispensing buret your teacher has prepared into the beaker containing the 0.3% hydrogen peroxide solution. CAUTION: H 2 SO 4 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 to the beaker as well. 4. Place the beaker of H 2 O 2 solution on a magnetic stirrer and add a stirring bar if available. If no magnetic 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). Rinse and fill the 50 ml buret with M MnO 4 solution. CAUTION: Handle the KMnO 4 solution with care; it has been mixed with H 2 SO 4, 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 the region of the titration curve near the equivalence point, and not to precisely determine the equivalence 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 Repeat the necessary steps to conduct a second trial with a new sample of 0.3% H 2 O 2 and acid solution. 11. When you conduct the second titration, carefully add the MnO 4 solution drop by drop in the region near the equivalence point, so that you can precisely identify the equivalence point of the reaction. 12. Precisely determine the equivalence point of the reaction. Record this information in your data table for Trial At the direction of your instructor, conduct a third trial. Data Tables Trial 1 Trial 2 Trial 3 Volume of H 2 O 2 solution (ml) Volume of MnO 4 solution used (ml) Calculated number of moles of H 2 O 2 Determination of Concentration by Oxidation-Reduction Titration of Hydrogen Peroxide 3 13

14 Post-Lab Questions and Analysis 1. Calculate the moles of MnO 4 used to reach the equivalence point of the reaction for each trial. 2. Use the number of MnO 4 moles to calculate the moles of H 2 O 2 in the sample of solution for each trial. 3. Calculate the molar concentration of the H 2 O 2 solution for each trial. 4. Calculate the % Error between your experimentally determined molarities for each trial and the molarity value for 3% hydrogen peroxide that you calculated as your answer to Pre-Lab Question 6. (Remember that you titrated 0.3% H 2 O 2 solution samples and NOT 3% H 2 O 2 solution samples.) 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 on the calculated molarity of hydrogen peroxide for trial 3? 7. A student titrates three samples of 0.3% H 2 O 2 solution made from a bottle of 3% H 2 O 2 solution dated 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. Determination of Concentration by Oxidation-Reduction Titration of Hydrogen Peroxide 4 14

15 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 ( ). The commonly accepted definition of Avogadro s number is the number of atoms in exactly 12 g of the isotope 12 C, and the quantity itself is (47) 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. Ammeter or Current Probe Figure 1 1 CRC Handbook of Chemistry and Physics, 82 nd edition, p Electrochemical Cells & Electroplating 15 1

16 OBJECTIVES In this experiment, you will Prepare an electrochemical cell to oxidize a copper electrode. Measure the amount of copper that was deposited in the electroplating process and determine the average current used. Calculate a value for Avogadro s number and compare it to the accepted value. Calculate the value of a Faraday. MATERIALS Ammeter 1 M H 2 SO 4 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 the electrochemical cell. Obtain a strip of zinc metal to use as the cathode, and clean it with steel wool if needed. 3. Use an analytical balance to measure the mass of the copper strip. Record the mass in your data table. 4. Fill a 250 ml beaker about ¾ full with 1 M H 2 SO 4 solution. CAUTION: Handle the sulfuric acid 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 two metal electrodes to be used in the electrochemical cell as you see in Figure 1. Copper is the anode (+) terminal and zinc is the cathode ( ) terminal in this cell because it is an electrolytic cell 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 Set up the ammeter. You will manually collect data for 3 minutes. 7. Place the electrodes into the 1 M H 2 SO 4 solution in the cell. Make sure that the electrodes are immersed in the solution to equal depths and as far apart as possible. Electrochemical Cells & Electroplating 16 2

17 8. Turn on the DC power supply and check the current readings. The initial current should be in the amp range. If the current is not in this range, adjust the settings on the power supply. 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 data collection stops. 10. When the data collection is complete, turn off the power supply and carefully remove the electrodes from the H 2 SO 4 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 the electrolysis. Record the mean in your data table as the average current for Trial 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 the anode or the cathode. Justify your answer by writing the half-reaction that occurs at the electrode. Electrochemical Cells & Electroplating 17 3

18 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 cathode and copper anode having an initial mass of g. The student applied a current of amps for a total of three minutes. After data collection was complete, the copper anode was dried and massed. The final mass of the copper anode was 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 coulombs. (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 is turned off. What effect does this error have on the calculated value of Avogadro s number? POST-LAB QUESTIONS AND DATA ANALYSIS 1. Calculate Avogadro s number for each trial. 2. Calculate the percent error for each trial 3. 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 the Faraday 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? Electrochemical Cells & Electroplating 18 4

19 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 Determine the best substances to create a commercial hand warmer. Safety Precautions The solids and resulting solutions in this lab are potential eye and skin irritants. Calcium chloride can cause skin burns. Ammonium nitrate is a powerful oxidizer that must be kept away from ignition sources and is quite toxic on ingestion. Wear splash proof safety goggles and an apron. If solutions are spilled on skin, wash with copious amounts of water. Materials LabQuest Coffee cup w/ lid Temperature probe Magnetic stirrer and stir bar Balance 3 ionic solids (assigned by group) The Hand Warmer Design Challenge 19 1

20 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 CaCl NaC 2 H 3 O Na 2 CO LiCl NH 4 NO When sodium chloride is dissolved in water, the temperature of the resulting solution is lower than the temperature of the water before the salt dissolves. How can this result be explained based 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 ml sample of water in a clean, dry 150 ml beaker. Heat with occasional stirring to approximately 50 C. Remove the beaker from the hot plate and place on the lab bench. 2. Meanwhile, place exactly ml of cool water (approximately 20 C) in the clean dry calorimeter. 3. Measure the temperature of the hot water and the cold water and record in Table 1, then immediately 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 the probe is actually in the water). Record the temperature of the water for approximately 3 minutes. [When the temperature readings are continually decreasing, you can stop recording.] Find the maximum temperature of the water and record this in Table 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 initial temperature of the water and record in Table Measure approximately 5.0 g of one of your ionic solids. Record the exact mass in Table 2. The Hand Warmer Design Challenge 20 2

21 8. Add the solid to the water in the calorimeter, quickly cover the solution and insert the temperature probe. Record the temperature of the solution for 3 minutes or until the temperature begins decreasing. Record the maximum temperature in Table 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: m hot (g) Mass cold water: m cold (g) Initial temp cold, T icold ( C) Initial temp hot, T ihot ( C) Final temp of mixture, T f ( 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, q hot, is equal to the enthalpy change of the cold water, q cold, but opposite in sign, so q hot =-q cold. 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 J/g C. 1. Calculate the enthalpy change of the water using the equation q hot = m hot c T hot. Assume that the density of water is exactly 1 g/ml. Is this an endothermic or exothermic process? Explain. The Hand Warmer Design Challenge 21 3

22 2. Calculate the enthalpy change of the water using the equation q cold = m cold c T cold. 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 real conditions observed in the laboratory, the law of conservation of energy equation becomes q hot = -(q cold + q cal ), where q cal is the enthalpy change of the calorimeter. Use this equation to calculate the enthalpy change of the calorimeter. 4. The calorimeter constant, C, is the heat absorbed by the calorimeter per degree of temperature change, C = q cal / T cal. Assuming the starting temperature of the calorimeter is the same as that of 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, q soln, is the reason for the difference in the thermal energy of the two systems (solid and pure water versus solution), with symbol q rxn. Just as with the hot and cold water in the calorimeter constant determination, q soln and q rxn are equal in magnitude and opposite in sign, q rxn = q soln. 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 q soln = (q rxn + CΔT) where C is the calorimeter constant determined above. This difference in thermal energy of the system before and after solution formation, q soln, can be calculated using the relationship q rxn = 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 q soln and q rxn for all three solids you tested for your hand warmer. 6. By convention, scientists report enthalpy changes for dissolution (and many other processes) in units of kilojoules per mole of solute dissolved. Using your values of q soln, calculate the enthalpy in units of kilojoules per mole. This quantity has the symbol ΔH soln. 7. Report your calculated values for each of the ionic solids in an organized table (you will have to create this). This table should include cost and any environmental issues as well as the heat data. The Hand Warmer Design Challenge 22 4

23 8. Based on the cost information provided, and your experimental work and calculations, select which chemical you believe will make the most cost effective hand warmer. The hand warmer you are designing needs to increase in temperature by 20 C. Calculate the amount of the compound 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 your rationale for choosing one chemical and not each of the other chemicals studied in this experiment. Your paragraph should start with a claim sentence that clearly states your choice and the amount of substance to use. The claim should be followed by evidence from your experiment and cost and safety analysis. The paragraph should conclude with reasoning explaining how your evidence supports your claim. The Hand Warmer Design Challenge 23 5

24 Group solid assignment -- Students work in pairs (24 and under) or groups of three. Group # Ionic Solids 1 NH 4 NO 3 Na 2 CO 3 CaCl 2 2 NaCl LiCl NaC 2 H 3 O 2 3 NH 4 NO 3 LiCl NaC 2 H 3 O 2 4 NaCl Na 2 CO 3 CaCl 2 5 NH 4 NO 3 LiCl NaC 2 H 3 O 2 6 NaCl CaCl 2 NaC 2 H 3 O 2 7 NH 4 NO 3 NaC 2 H 3 O 2 Na 2 CO 3 8 NaCl LiCl Na 2 CO 3 9 NH 4 NO 3 LiCl CaCl 2 10 NaCl NaC 2 H 3 O 2 Na 2 CO 3 11 NH 4 NO 3 LiCl Na 2 CO 3 12 NaCl LiCl CaCl 2 The Hand Warmer Design Challenge 24 6

25 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: A Rate = = k[a] m time 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. Reaction Rate 25 1

26 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. Zero order k = negative slope First order k = negative slope Second order k = the slope Another important part of the kinetic analysis of a chemical reaction is to determine the activation energy, E a. 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: E a = R slope ln k Reaction Rate 26 2

27 Determination of the Rate of a Reaction, Its Order and Its Activation Energy 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 CV and 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, E a, 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 the 565 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 for 60 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 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. Reaction Rate 27 3