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lab 1

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by: Alexis Notetaker

lab 1 Chem 1415

Alexis Notetaker
GPA 3.67

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About this Document

General Chemistry (Continued)
Dennis Awasabisah
Class Notes
25 ?




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"You can bet I'll be grabbing Alexis studyguide for finals. Couldn't have made it this week without your help!"
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This 8 page Class Notes was uploaded by Alexis Notetaker on Saturday January 2, 2016. The Class Notes belongs to Chem 1415 at University of Oklahoma taught by Dennis Awasabisah in Summer 2015. Since its upload, it has received 10 views. For similar materials see General Chemistry (Continued) in Chemistry at University of Oklahoma.


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Date Created: 01/02/16
Alexis Penny th September 16 , 2015 Experiment #03 Determining Avogadro’s Number CHEM 1315-077 Penny 1 Purpose of the Experiment & Techniques: The purpose of this experiment is to confirm Avogadro’s number through electrolysis. A copper strip and a zinc strip will be used as the electrodes, the copper strip as the anode and the zinc strip as the cathode, and put into a beaker of sulfuric acid to make the cell electrolytic. By determining the average current used in the reaction, along with the knowledge that all of the copper ions formed are the 2+ cations, the number of atoms in one mole of copper can be calculated and compared with Avogadro’s number. A coulomb is the SI unit of electric charge. It is the amount of electricity transferred in one second by a current of one ampere. To calculate the charge in coulombs, the formula C=As would be used. It is known that C is the charge in coulombs, A is the current in amperes, and s is the time in seconds. Percent error can be calculated using the formula: {(theoretical value-experimental value)/(theoretical value)} X 100. When an experiment is being used to prove a known value, the percent error formula can be used to test how precise your calculations were. Percent error tells you how accurate your results are compared the known value. 2 Penny Materials: LabQuest 1 M sulfuric acid, H2SO4, solution Vernier Current Probe Copper strip (anode) 1.5 volt DC power source Zinc strip (cathode) 2 connecting wires with alligator clips Distilled water Sand paper Ohm Resistor Analytical balance Precautions: Handle the sulfuric acid with care. It can cause burns if it contacts skin. Procedure: First, put on safety goggles. Be sure to use the Chemicals Utilized table to keep track of materials used during the experiment. Clean a strip of copper, which will be the anode of the electrochemical cell using sandpaper. Once clean use an analytical balance to measure its mass and record this value in the data table. Obtain a strip of zinc metal to use as the cathode, and clean it with steel wool if needed. Next, fill a 250 mL beaker about ¾ full with 1 M H2SO4 solution. Obtain a DC power supply, an ohm resistor, a Vernier Current Probe, and two connecting wires. Use one connecting wire to connect the Penny 3 copper electrode to the black lead on the Vernier Current Probe. Connect the power supply (red painted end or black cord) to the red lead on the Vernier Current Probe. Connect the power supply (black end or black/white striped cord) to the black lead on the ohm resistor. Connect the zinc electrode to the red ohm resistor lead using a connecting wire. Next, connect the current probe to the LabQuest. To set up the LabQuest, choose New under the File menu and change the data-collection rate to 0.2 samples/second, interval 5 seconds/sample, and the duration to 180 seconds. Then select the graph icon in the upper right hand corner. Plug in the DC power supply and place the electrodes into the 1 M H2SO4 solution in the cell. Quickly make sure that the electrodes are sumberged in the solution to equal depths and as far apart as possible. The initial current should be in the 0.2–0.6 amp range. If the current is not in this range, quickly adjust the settings on the ohm resistor. Once the initial current is in range, quickly move to the next step of beginning data collection. Start by pressing the green arrow in the bottom left hand corner. Allow the data to collect for 3 minutes and observe the reaction carefully. You should be ready to unplug the power supply as soon as the data collection stops. When the data collection is complete, turn off the power supply and remove the electrodes from the H2SO4 solution. Next, rinse the copper and zinc electrodes with distilled water in a separate beaker and dry the copper electrode very carefully. Measure 4 Penny and record the mass of the dry copper electrode. To determine the average current applied during the experiment, start by choosing the Analyze menu at the top middle of the LabQuest screen. Select statistics under the Analyze menu and then record the average current in the Data Table. Reconnect the electrodes, start a new file, and repeat Steps 6–10 for a second and third trial. The initial mass of the copper electrode for the next trial will be the final mass from the previous trial. Observe the color of your 1 M sulfuric acid after your third trial and record your observations in question 1. Dispose of all chemicals properly and clean all glassware and equipment used. Data and Analysis: Data Table Trial 1 Trial 2 Trial 3 1 Initial mass of copper electrode (g) 0.703 0.689 0.671 2 Final mass of copper electrode (g) 0.689 0.671 0.654 3 Average current (A) 0.239 0.250 0.252 180 180 180 4 Time of current application (s) Chemicals Utilized Chemical Name Amount Waste Type 1 M Sulfuric Acid 100mL Liquid Copper Electrode 0.651g Solid Electrode Zinc Electrode 0.822g Solid Electrode Data Analysis Trial 1 Trial 2 Trial 3 Penny 5 5 Change in mass of copper (g) 0.0140 0.0180 0.0170 6 Change in moles of copper (mol) 2.20E-4 2.80E-4 2.70E-4 43.1 45.0 45.4 7 Total Charge (C) 8 Electrons lost 2.69E20 2.81E20 2.84E20 9 Copper atoms lost 1.35E20 1.41E20 1.42E20 10 Copper atoms per mole of copper lost 6.14E23 5.04E23 5.26E23 Discussion: The final color of the 1 M sulfuric acid solution was clear with a pale blue tint. The final balanced equation is Cu + 22 SO  2CuS4 + 2H 4 2. The product copper sulfate is blue in its dry powder state, which was most likely the cause of the resulting solution’s color. I determined the grams of copper lost in each trial by subtracting the initial mass of copper from the final mass of copper resulting in 0.0140g copper lost for the first trial, 0.0180g copper for the second, and O.0170g copper for the third. To convert these values to moles, I divided each value by the atomic mass of copper, 63.546 because there are this many grams of Cu in one mole of Cu. Trial one results in 2.20E-4 moles Cu, trial two results in 2.80E-4 moles Cu, and trial three results in 2.70E-4 moles Cu. To find the total charge in coulombs I multiplied the average current by the amount of time of application for each trial (180 seconds). I 6 Penny calculated that 43.1 coulombs passed through the electrolytic cell in trial one, 45.0 coulombs in trial two, and 45.4 coulombs in trial 3. Next, to find the number of electrons in the electrolysis, I divided 1 by the charge of an electron (1 electron/ 1.602E-19 coulomb) and then multiplied this by the calculated total charge for each trial. The results were 2.69E20 electrons for trial 1, 2.81E20 electrons for trial 2, and 2.84E20 electrons for trial 3. I determined the number of copper atoms lost from the anode by dividing the number of electrons measured in half. The number of Cu 2+ions equals ½ the number of 2+ electrons measured. Therefore, trial one lost 1.35E20 Cu ions, trial two lost 1.41E20 Cu 2+ions, and trial three lost 2.84E20 Cu 2+ions. To determine the number of copper ions in a mole, I took the number of Cu 2+ions for each trial and divided them by the number of moles of copper lost that was calculated for each trial. This is possible because the mass of copper ions produced is equal to the mass loss of the anode because the mass of the electron is so small it is negligible. Trial one is calculated to have 6.14E23 copper atoms/mole, trial two 5.04E23 copper atoms/mole, and trial three 5.26E23 atoms/mole. The average for these values is 5.48E23 atoms/mole. The percent error with my value and Avogadro’s number is found by dividing the experimental value by the theoretical value times 100 (5.48E23/6.02E23 X 100). The calculated percent error is 91.0%. The possible errors that could have occurred are inputting incorrect Penny 7 conditions in the labquest and failing to start the data collection as soon as the reaction begins along with other human errors. 8 Penny


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