Introduction to Chemistry Techniques Prelab (Week 1) Name Total /10 SHOW ALL WORK NO WORK = NO CREDIT 1. What is the purpose of this experiment? 2. Determine the number of significant figures in each of the following numbers. a) 23.7 b) 1340.1 c) 0.000218 d) 33.50 3. a) Add the numbers from question 2 together and give the answer with the correct number of significant figures. b) Multiply the numbers from question 2 and give the answer with the correct number of significant figures. Write the answer in scientific notation. 4. A cylindrical object has a density of 15.1 g/cm 3. The diameter of this object is 1.10 cm and the height is 2.05 cm. What is the mass of the object? (V = hπr 2 ) 5. Define accuracy and precision. 6. You measure the mass of an object four times and obtain the following values of your measurements: 14.2005 g, 14.2007 g, 14.2003 g and14.2004 g. Calculate the average of these measurements and determine the standard deviation for these measurements. What does the standard deviation tell you about the precision of your measurements? 1
Absorbance Introduction to Chemistry Techniques Prelab (Week 2) Name Total /10 SHOW ALL WORK NO WORK = NO CREDIT 1. The following graph shows absorbance versus concentration of a compound in solution. From the information on the graph given to you: Absorbance vs. Concentration for nickel(ii) sulfate hexaydrate 1.200 1.000 y = 110.71x 0.800 0.600 0.400 0.200 0.000 0.00000 0.00100 0.00200 0.00300 0.00400 0.00500 0.00600 0.00700 0.00800 0.00900 0.01000 concentration (mole/l) a) Determine the molar absorptivity ( ) for this compound. b) Determine the concentration of an unknown solution that has an absorbance value of 0.567. Use the information given to you in the introduction. 2
Introduction to Chemistry Techniques In this experiment you will measure the mass of an object with two different kinds of balances and calculate the volume of liquid that is delivered from a transfer pipet. You will also determine the density of an object. The last part will introduce you to spectroscopy and how to use absorbance of light to determine the concentration of an unknown solution. Special mention goes to Brian Stahl for developing Part IV of this laboratory experiment to be used here in the Behrend chemistry curriculum. Introduction Many experiments require some type of measurement, and are often simple measurements of mass and volume. The validity of an experiment is dependent on the reliability of these measurements. A measurement s reliability is usually considered in terms of its precision. Precision is the closeness of the agreement between successive measurements of the same quantity. Precision is not to be confused with accuracy, the agreement of a particular value with the true value. The dispersion in a set of measurements is usually expressed in terms of the standard deviation, whose symbol is s. You are going to be asked to determine the standard deviation for some of your data. The standard deviation measures how close the data are clustered around the mean. The smaller the standard deviation, the more closely the data are clustered around the mean and the more precise your measurements are. After a quantity has been measured in an experiment, it may be necessary to use that measurement in a subsequent calculation. If you use a hand calculator for the arithmetic, eight or more digits may appear in the answer. It is up to you to decide how many of these digits are significant. IT IS YOUR RESPONSIBILITY TO KNOW THE RULES FOR SIGNIFICANT FIGURES BEFORE YOU START THIS EXPERIMENT. You will have to practice using an analytical and top loader balance and then a transfer pipet so you can gain confidence to perform this experiment. Next, you will measure the mass of a flask four times on each balance. The precision of the measurements will be examined when you determine the correct number of significant figures in the mean mass of the flask. You will then add water to the flask from a filled 5-mL pipet, and then measure the mass of the flask and water. You will repeat this process three more times. After calculating the mass of the water that was delivered each time from the pipet, you will calculate the volume of each addition from the mass and density of water. You will then determine the correct number of significant figures in the mean volume. This number will allow you to appreciate the precision that you have achieved with the pipet. You will have to measure volume using a graduated cylinder. A graduated cylinder is used to measure an approximate volume of a liquid. When water or an aqueous solution (a solution containing water) is added, the upper surface of the liquid in the graduated cylinder will be concave. This concave surface is called the meniscus. The bottom of the meniscus is used for all measurements. To avoid error, your eye should always be level with the meniscus when you are measuring the volume. Using the graduated cylinder gives you the ability to measure the volume to only one decimal place. 3
In the last part of this experiment, you will be using spectroscopy to determine the concentration of an unknown solution. There are many ways to determine concentrations of a substance in solution and using spectroscopy is only one method of determination. Many properties of a solution can change with concentration. The property you will be looking at in this experiment is colour. There is a direct correlation between colour intensity and concentration of a solution. Concentration and absorbance are linearly related. This means that the higher the concentration, the deeper the colour and therefore the higher the concentration of a substance in solution. The opposite is also true. Using colour can be much faster than using wet methods such as titration, especially when you have many samples containing different concentrations of the same substance. When coloured solutions are irradiated with white light, they will selectively absorb light of some wavelengths, but not of others. The wavelength at which absorbance is highest is the wavelength to which the solution is most sensitive to concentration changes. This wavelength is called the maximum wavelength or max. In order to determine the precise amount of a substance present in a solution, a calibration curve must be produced. A calibration curve shows the relationship between the absorbance of light and the concentration of a chemical in a solution. The higher the concentration of a substance in solution, the deeper the colour will be and therefore the greater the absorbance of light. The opposite is also true. You will make four solutions where the concentration of your substance is known. By plotting the absorbance readings on the y-axis of the graph and the substance concentration values on the x-axis, you can use this information to determine the concentration of an unknown solution of that same substance. An example of a calibration plot is given. Once you have produced this calibration plot, you will then use it to determine the concentration of your unknown solution. The information you obtain from the calibration plot will be applied to Beer s Law and the unknown concentration of your solution will be determined. The unit for concentration being used in this experiment is moles/litre. Beer s law is defined by the following equation: A = bc (1) The equation represents three variables that influence the response of a solution to light. They are the concentration (c) of the solution, the pathlength (b) of the light through the solution (also called the cell length) and the molar absorptivity ( ), the sensitivity of the absorbing species at max. The pathlength, unless it is stated differently, is usually fixed at 1.00 cm. The molar absorptivity value is dependent on the solvent used (in this case water) and. The units for molar absorptivity are L/mole-cm for a concentration with units of mole/l. You will make four solutions where the concentration of your substance is known. By plotting the absorbance readings on the y-axis of the graph and the substance concentration values on the x-axis, you can use this information to determine the concentration of an unknown solution of that same substance. An example of a calibration plot is given. The slope of the line from the plot is the molar absorptivity for your chemical in solution. You will then measure the absorbance of a solution of unknown concentration. Using this absorbance value, along with the molar absorptivity determined from the calibration plot, you will be able to determine the concentration of your unknown solution by manipulating Beer s Law: 4
absorbance c = εb A (2) You will use pipets and volumetric flasks to make solutions of known concentrations as accurately as possible and measure the absorbance at max. You will be making dilutions from a stock solution of a given substance. Since you will not be able to calculate dilutions yet, you will use the following equation to calculate your concentrations: C 2 = C1xV V 2 1 (3) where C 1 is the concentration of your stock solution, V 1 is the volume of the stock solution that you are measuring and V 2 is the final volume of the solution. For example, if you measure five milliliters (ml) of your stock solution (V 1 ) of known concentration (C 1 ) and add it to a 100 ml volumetric flask and then added water to the line which indicated 100 ml of solution (V 2 ), you could then find C 2 using equation 3. FIGURE 1: A Calibration Curve of Copper(II) Sulphate Pentahydrate in Solution copper(ii) sulphate pentahydrate 1.4 1.2 1 0.8 0.6 0.4 0.2 0 0 0.05 0.1 0.15 0.2 0.25 0.3 0.35 concentration (moles/litre) 5
Procedure Part I: Using the Balance 1. Obtain about 100 ml of distilled water in a beaker. Allow the beaker and water to sit on the laboratory bench while you are learning to use the balances and the pipet for parts II and III. The water should come to the temperature of the laboratory during that time. 2. Obtain a thermometer and a 50-mL Erlenmeyer flask with a rubber stopper. 3. Use the SAME top loader balance throughout the experiment. 4. Place the rubber stopper in the Erlenmeyer flask. Tare the top loader balance. Measure and record the combined masses of the flask and stopper. 5. Use tissue paper or paper towel to remove the stoppered flask from the pan of the balance. This is used because some balances are sensitive enough to detect the oils from your fingerprints and you will be weighing the flask on the analytical balance in step 9. 6. Bring your balance to the zero position again. Measure and record the mass of the stoppered flask once again. 7. Repeat steps 5 and 6 until you have measured the mass four times. 8. Calculate the mean (average) mass. The differences between the measured masses and the mean should be very small. If you are unsure of your results, consult your laboratory instructor. 9. Repeat steps 3-8 with the analytical balance. Be sure to use the SAME balance throughout the experiment. Part II: Using the Pipet Be sure to use the SAME BALANCE for ALL measurements. 1. Obtain a 5 ml pipet. 2. Practice using the pipet with distilled water (not the water you have set aside) until you are comfortable with the technique. You should plan on using the same analytical balance and the same pipet throughout the experiment. 3. Using the thermometer, note the temperature of the laboratory and of the distilled water that you have set aside. When the temperatures are identical or very nearly identical, you can begin. Record the temperature to the nearest degree. 4. Measure and record the mass of the empty stoppered flask again using the ANALYTICAL BALANCE. Use tissue paper/paper towel as you did before. 6
5. Remove the flask from the balance, using tissue paper. Pipet exactly 5 ml of the room-temperature water into the flask without touching the flask with your fingers. Using tissue paper/paper towel, replace the stopper to prevent evaporation. 6. Bring your balance to the zero position. Measure and record the combined mass of the water and the stoppered flask. 7. Remove the flask from the balance. Do not pour out the first sample. Pipet another 5 ml sample into the flask. The volume of the water in the flask should now be 10 ml. Replace the stopper and repeat step 6. 8. Repeat until four samples of water have been delivered to the flask and the final volume is 20 ml. 9. Calculate the mass of water that was delivered each time from your pipet. These masses should be approximately identical. 10. Calculate the volume of each sample from the mass and density of water. Use Table 1 to find the density that corresponds to your recorded temperature. Calculate the average volume delivered from each sample. Part III: Determining the Density of an Object 1. Using your graduated cylinder, measure 40-50 ml of water from your beaker in Part I and leave it in the graduated cylinder. Record this measurement to the nearest 0.1 ml. 2. Take one of the objects from the samples that are given and find its mass on the analytical balance. 3. Place that same object in the graduated cylinder with the water in it. Make sure that the object is completely submerged in the water. If the object is not completely submerged, repeat steps 1and 2 being sure to dry the object and adding more water in the graduated cylinder. Avoid splashing any water out of the graduated cylinder or cracking the bottom of the graduated cylinder. 4. Record the level of the water with object in the graduated cylinder. The difference between this measurement and the measurement in step 1 is the volume of the object. 5. Using the same object that has been dried well, measure the width and height of the object to the nearest 0.01 cm. 6. Calculate the volume of the object using the measurements determined in step 5. 7. Calculate the density of the object (g/ml and g/cm 3 using the volumes from step 4 and step 6) using the correct number of significant figures. 7
Table 1 Density (g/ml) of Water at Various Temperatures ( o C) Temp. Density Temp. Density Temp. Density 17 0.9988 22 0.9978 27 0.9965 18 0.9986 23 0.9976 28 0.9962 19 0.9984 24 0.9973 29 0.9959 20 0.9982 25 0.9971 30 0.9956 21 0.9980 26 0.9968 31 0.9953 Example I: How to Calculate a Standard Deviation You obtain the following measurements: 15.2654, 15.2657, 15.2658 and 15.2655. In the following calculations, each individual measurement from above will be represented using the symbol x i and the mean (the average of these values) will be given the symbol x. 2 d i s N 1 means the sum of d i = x i - x N = the number of measurements x = 15.2656 x i d i (x i x ) d i 2 15.2654-0.0002 0.00000004 15.2657 0.0001 0.00000001 15.2658 0.0002 0.00000004 15.2655-0.0001 0.00000001 2 0.00000010 = d i 0.00000010 s = 0.000183 = 0.0002 4 1 When you report a value, it is the mean (average) value that is reported along with the standard deviation to show the precision of the measurements which contributed to the mean value. Therefore, the value reported for this example would be 15.2656 0.0002. 8
Part IV: Using Spectroscopy to Determine the Concentration of an Unknown Solution 1. Using a 250.0 ml beaker, obtain approximately 75.0 ml of stock solution of known concentration (write this concentration down). 2. Using a pipet, transfer 5.0 ml of stock solution to a 100.0 ml volumetric flask and fill up to the 100.0 ml mark with distilled water. Mix and then transfer all the solution to a clean dry beaker that is 150 ml or greater. This will be called solution 2. 3. You will do this three more times making solution 3 (10.0 ml of stock solution diluted to 100.0 ml), solution 4 (15.0 ml of stock solution diluted to 100.0 ml), and solution 5 (20.0 ml of stock solution diluted to 100.0 ml). Remember, all dilutions are made with distilled water and remember to place each solution in a beaker after mixing. Calculate the concentrations (C 2 ) of solutions 2-5 using equation 3 on page 5. You will need these values when you make your calibration plot. 4. Once all your solutions are made and concentrations have been calculated, go to the netbook and open up the Chem 111 folder. Select the file labeled Introduction to Chemistry Techniques Fall 2011 and double click on that file. You will then see a table and a graph on the screen. 5. Take distilled water and pour it in the cuvette. Make sure there are no fingerprints and that the light is passing through the smooth surfaces not the ribbed surfaces. Fill cuvette with distilled water to make a blank. 6. On the menu bar, select Experiment, then Calibrate, and then Spectrometer:1. Wait for the lamp to warm up. 7. When the lamp has warmed up, place the blank in the spectrometer. Click Finish Calibration, then click OK. Keep the blank in the spectrometer since this will be your first data point. 8. In the top right hand corner, you will see the collect icon. Click on collect. You should see a very low absorbance value in the bottom left hand corner of the screen. If you do not have a very low value, get help from your instructor. 9. Once you have waited 20-30 seconds, go to the top right hand corner of the screen where it says keep and click. You will now be asked for the concentration. Since this is your blank, the concentration is zero. 10. Repeat step 9 for the rest of the solutions. Remember to wait 20-30 seconds before clicking on keep. Copy all the absorbance values as well as the respective concentrations in your laboratory notebook. 9
11. When you are all done getting the absorbances for solutions 2-5, click on the stop icon. Then go to the linear fit icon and click. A box filled with information will appear and you will find the value of the slope of the graph (m). This is your molar absorptivity value. Record that value. 12. You will now take a sample of an unknown solution and place it in a clean cuvette. Place the cuvette in the spectrometer and read the absorbance value for that solution from the netbook. Now you can calculate the concentration of your unknown solution by using equation 2. NOTE: When closing logger pro, do not save changes. 10
Results Part I Using the Analytical Balance Mass of the stoppered flask (g) Calculation: Mean mass (g) Using the Top Loader Balance Mass of the stoppered flask (g) Mean mass (g) Part II Using the Pipet Temperature ( o C) Density of water (g/ml) Addition No. 1 2 3 4 Mass after addition (g) Mass before addition (g) Mass of water in flask (g) Volume of water delivered each time (ml) Mean volume (ml) Calculations (One sample calculation for one addition in Part II): 11
Part III Determining the Density of an Object Mass of the object (g) Volume after adding object (ml) Volume before adding object (ml) Diameter of object (cm) Radius of object (cm) Volume of object (ml) Volume of object (cm 3 ) (using the equation V = h r 2 ) Height of object (cm) Density of the object (g/ml) Density of the object (g/cm 3 ) Calculations: Questions 1 a. Calculate the standard deviation (look at example on page 8) for the volume of water delivered each time in Part II. b. What does the standard deviation from part a tell you about the precision of your measurements? 12
Part IV Solution Concentration Absorbance 1 0.000 0.00 2 3 4 5 Unknown? Unknown Letter or Number Concentration of Stock Solution Molar Absorptivity of Chemical in Solution Concentration of Unknown Solution Calcuations: 13