Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013EXPERIMENT 4: Determination of Avogadro’s number using Electrogravimetry OBJECTIVE: To make an experimental measurement of Avogadro’s number using an electrochemical technique (electrogravimetry). INTRODUCTION: The most commonly used basic counting unit chemistry field is Avogadro’s number which was found by an Italian scientist, Amadeo Avogadro (Vernier Software & Technology, 2013). The definition of 13 Avogadro’s number is the number of atoms in exactly 12 g of th e isotope C and the quantity itself is 23 6.02214199 × 10 . In this experiment, the Avogadro’s number was confirmed by conducting an electrochemical process called electrolysis. This process used an external power supply to drive and non spontaneous reaction. A careful measurement of electron flow, amperage, and time to obtain the number of electrons passing through the electrochemical cell will be done in the experiment. The electron flow, in amperes, is usually referred to as the current. the Avogadro’s n umber can be calculated through the number of atoms in a weighed sample which are related to the number of electrons used. There are several ways to determine the Avogadro’s number which in this experiment, the electrogravimetric technique was applied. The experimental setup for this process is called an electrolytic cell. An electrolytic cell is made up of these following components: A direct current source (eg. Battery or power supply) Insulated wires to connect the circuit Two electrodes An electrolyte (sulphuric acid) The electrolytic process was used to determine the number of electron needed to convert one mole of 2+ copper atoms to one mole of copper ions Cu . This value was divided by to and it represents the number of atoms converted from copper metal to copper ions: Cu → Cu 2+ 2e The above reaction is half equation that represents oxidation. The number of copper atoms per mole of copper is Avogadro’s number, the value to be determined. To find out the number of electron consumed in this process is determined by using the charge of an electron and the total charge measured. By using the Millikan oil-drop experiment, the charge of an electron was determined to be -19 1.60217733 x 10 coulombs per electron. By applying the relationship: 1 ampere = 1 coulomb/second, the number of coulombs used in this experiment can be calculated. An ammeter was used in the experiment to measure the amperage and a stopwatch was used to measure the time passed. The mass of copper that reacted can be obtained by measuring the mass of the anode and he cathode before and after the electrolysis. In the electrolytic cell contains both copper electrodes and the electrolyte 1M CuSO4 + 0.5M H2SO4. The copper electrode (anode) connected to the positive pin loses mass as the copper atoms are converted into copper ions as shown in the equation earlier. the loss of mass is visible after a while as 2+ pitting of the surface of the metal electrode. In addition, the copper ions, Cu , produced immediately pass into water solution and deposit on the cathode as the reaction shown below: Cu 2+ 2 e → Cu (solid) Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013 APPARATUS: 20V power supply, U-tube, copper electrodes, electrical wires with crocodile clips, retort stand with clamps, emery paper and 4-decimal analytical balance. CHEMICALS: 1M CuSO4 and 0.5M H2SO4 PROCEDURE: 1. Two copper electrodes were obtained, polished and cleaned before any measurement was taken. The electrodes were dipped in a beaker of clean tap water and then they were dipped in a beaker of alcohol. a sticker was put to the electrode after the electrode was dried. The electrodes were weight carefully and less weight electrode was labelled as anode. The electrolytic solution used was 1M CuSO4 in a 250mL beaker. 2. Based on Figure 1 shown, the circuit was set up by setting the power supply at 20. The positive pole of the power supply was connected to the anode of the first cell. The cathode was connected to the positive pin of the ammeter. The amperage was recorded at 30 seconds intervals for 10 minutes. The average amperage was taken to be used in the calculation. 3. When electrolysis had stopped, the anode and cathode were retrieved, rinsed gently and dried with distilled water. Then, dried them with tissue paper and immersed in the alcohol. Do not wipe the electrode since it will remove the copper from the surface. The anode and cathode were weighted. 4. The same electrodes were used, re-polished with emery paper ad re-weighted. The electrolysis was repeated using 0.5M H2SO4 solution. The observation at electrode and electrolyte were recorded. Figure 1. The circuit set up (Vernier Software & Technology, 2013). Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013 RESULT: Electrode measurements CuSO4 24.9801 g 24.1974 g 0.7827 g 25.3221 g 26.5121 g 1.1900 g 1.6410 g 2.0394 g H2SO4 24.6355 g 22.1363 g 2.4992 g 25.2822 g 28.1709 g 2.8887 g Mass of anode before electrolysis Mass of anode after electrolysis Mass loss of anode Mass of cathode before electrolysis Mass of cathode after electrolysis Mass loss of cathode average weight loss at anode average weight loss at cathode Time-amperage measurements CuSO4 Current (A) 0.00 0.98 2.75 3.20 3.60 4.05 4.54 5.07 5.52 5.98 6.38 6.74 7.01 7.20 8.16 8.53 8.77 8.81 8.71 8.61 8.80 5.88 H2SO4 Current (A) 0.00 10.43 10.97 11.51 11.98 12.57 13.12 13.52 13.83 14.00 14.12 14.29 14.47 14.68 14.34 14.44 14.66 14.77 14.76 12.76 Time (secs) 0 30 60 90 120 150 180 210 240 270 300 330 360 390 420 450 480 510 540 570 600 Average Current Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013 Total time of electrolysis Average current during electrolysis total charge measured (amperes) Total charge measured (coulombs) Number of electrons passed Number of Cu2+ generated Based on weight loss of anode: Number of cu2+ ions/ gram Cu metal (Cu2+/g Cu) Avogadro number (from measurement) Avogadro number (true or accepted value) Absolute error in measured value Relative % error in measured value Based on weight gain of cathode: Number of Cu2+ ions/ gram Cu metal (Cu2+/g Cu) Avogadro number (from measurement) Avogadro number (true or accepted value) Absolute error in measured value Relative % error in measured value CuSO4 600 s 5.88 3528 22 2.202x10 electrons 22 2+ 1.101x10 Cu ions 1.407x10 22 H2SO4 540 s 12.76 6890 coul 22 4.300x10 electrons 22 2+ 2.150x10 Cu ions 8.605x10 Cu atoms/g 23 5.468x10 Cu atoms 23 6.022x10 24 0.554x10 9.20% 8.605x10 Cu atoms/g 23 5.468x10 23 6.022x10 23 6.0201 x 10 99.96% 21 21 Cu atoms/g 8.939 x 10 Cu atoms 23 6.022x10 23 2.919x10 48.49% 9.2521 x 10 Cu atoms/g 23 8.939x10 23 6.022x10 23 6.0129 x 10 48.49% 21 23 CALCULATIONS: Electrolysis of copper sulphate with copper electrode: Anode mass lost: 24.9801 – 24.1974 = 0.7827g Current: 5.88A Time of electrolysis: 600s Electrolysis of sulphuric acid with copper electrode Anode mass lost: 24.6333 – 22.1363 = 2.4986g Current: 12.76A Time of electrolysis: 540s Total charged passed through the circuit: = 5.88A x (1 coul/1 amp/s) x 600s = 3528 coul Total charged passed through the circuit: = 12.76A x(1 coul/1 amp/s)x540s = 6890 coul Number of electrons: -19 =3528 coul x(1 electron/1.6022x10 ) 22 =2.202x10 electrons Number of electrons: -19 6890 coul x(1 electron/1.6022x10 ) 22 = 4.300x10 electrons Number of copper atoms lost from the anode: 22 2.202x10 x (1 Cu2+/2 electrons) 22 2+ =1.101x10 Cu ions Number of copper atoms lost from the anode: 22 4.300x10 x (1 Cu2+/2 electrons) 22 2+ = 2.150x10 Cu ions Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013 Anode Number of copper ions per gram of copper: 22 2+ 1.101x10 Cu ions /0.7827g 22 = 1.407x10 Cu atoms/g Number of copper atoms in a mole of copper, 63.546g/mol 22 = 1.407x10 Cu atoms/gx63.54g/mol 23 = 8.939x10 Number of copper ions per gram of copper: 22 2+ = 2.150x10 Cu ions/ 2.4986g 21 = 8.605x10 Cu atoms/g Number of copper atoms in a mole of copper, 63.546g/mol 21 = 8.605x10 Cu atoms/g x 63.54g/mol 23 = 5.468x10 Percent error: 23 23 Absolute error: 8.939x10 - 6.022x10 = 23 2.919x10 23 Percent error: 2.919x10 x 100% 23 6.022x10 = 48.49% Percent error: 23 23 Absolute error: 5.468x10 - 6.022x10 24 = 0.554x10 24 Percent error: 1.149x10 x 100% 23 6.022x10 = 9.20% Cathode Number of copper ions per gram of copper: 22 2+ = 2.150 x 10 Cu ions/ 2.8887g 21 = 7.4427 x 10 Cu atoms/g Number of copper ions per gram of copper: 22 2+ 1.101 x 10 Cu ions /1.1900g 21 = 9.2521 x 10 Cu atoms/g Number of copper atoms in a mole of copper, 63.546g/mol 22 = 5.7947 x 10 Cu atoms/g x 63.54g/mol 20 = 9.1198 x 10 Number of copper atoms in a mole of copper, 63.546g/mol 21 = 7.4427 x 10 Cu atoms/g x 63.54g/mol 20 = 1.1713 x 10 Percent error: 20 23 Absolute error: 9.1198 x 10 - 6.022x10 23 = -(6.0129 x 10 ) 23 Percent error: 6.0129 x 10 x 100% 23 6.022x10 = 99.85% Percent error: 20 23 Absolute error: 1.1713 x 10 - 6.022x10 23 = -(6.0201 x 10 ) 24 Percent error: 1.149x10 x 100% 23 6.022x10 = 99.96% Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013 DISCUSSIONS “Electrogravimetry is electroanalytical method based on gravimetric determination of metallic elements, which are isolated on the cathode in form of metal or on the anode in form of metal oxide during electrolysis. This method employs two or three electrodes, and either a constant current or a constant potential is applied to the preweighed working electrode.“ (University of Wrocław, no date). The determination of Avogadro’s number was done through electrogravimetric technique. However, percentage error for each electrolyte was found to be high and nearing 100%. The values were 99.85 % and 99.96 % for CuSO4 and H2SO4 respectively. Experimentally, the net loss and gain for the Cu ions were found to be higher in H 2SO4 compared to in CuSO4. It was also found that the cathode thickness in H 2SO4 was found to be thicker compared to the cathode in CuSO4.This can be justified with the half reactions that took place in H2SO4 is shown below; In H2SO4 :- Anode: S2O8 (aq)+2e → 2SO4 (aq) O2 + 4H + 4e → 2H2O Cathode: 2H2O + 2e → H2(g) + 2OH Cu 2+ + 22- (aq)+ 2e → Cu (s) - - At anode, due to the fact that the position of OH in the standard reduction potential (SRP) list is lower 22than the SO4 , therefore, OH was much easier to be discharged compared to SO4 . Whereas at 2+ + Cu was discharged due to its higher ability to be discharged compared to H . Here, it is clearly shown that the Cu ions needed to compete only with H ions, however Cu ions 2+ would definitely be discharged easily. Higher chances of Cu ions to be reduced caused the deposition of Cu to be high, say, more efficient. On the other hand, there was a slight difference that occurred in CuSO4 which can be explained through the half equations as below. In CuSO4 :- Anode: Cu (s) → Cu (aq) + 2e + 2+ - 2+ + 2+ O2 + 4H + 4e → 2H2O Cathode: 2H2O + 2e → H2(g) + 2OH Cu (aq)+2e → Cu(s) 2+ - When a very small external current is applied to the copper electrodes, then the equilibrium between 2+ Cu in the solution and the Cu of the electrodes is disturbed. Copper goes into solution at the anode and an equivalent amount of copper ions are deposited at the cathode (PHYWE, no date). Theoretically, there is therefore no change in the total amount of dissolved copper sulphate, however, experiment showed otherwise. The net loss and gain for the electrodes did not tally. The mass loss by anode was lower compared to the mass gained by the cathode. The possible error was from the weighing of the electrodes. The electrode must be dried completely before being weighed. The contribution from electrolyte might have caused the cathode to gain phantom mass. Siti Mariam Abdul Kadir (2010762233) Submission Date: 24/12/2013 By referring to the aforementioned justifications in H2SO4, the only difference in CuSO4 was the reaction took place in the electrolyte that had the same ions as the electrodes used. With this 2+ 2+ situation, the Cu ions from the electrode needed to compete with the Cu from the electrolyte. Thus, the net gain and loss in this particular electrolyte was lower compared to that of in H 2SO4. The charge effect efficiency was found to be 58.35% and the calculation is shown as the following: m = (M x Q)/nF, where m is the mass of metal deposited, M is the molecular weight, Q (Q= It) is the coulombs, n is the number of electrons and F is the Faraday’s constant. At anode, 1.6410g = (63.54 g/mol x Q)/ (2e x 96500 C/mol) Qa = 3614 C At cathode, 2.0394g = (63.54 g/mol x Q)/ (2e x 96500 C/mol) Qc = 6194 C Charge passed efficiency, Qa/Qc x 100% = 58.35% Several errors might have occurred in setting up the circuit and experimentally, the presence of resistance might have reduced the efficiency. REFERENCES PHYWE (n.d.) Electrogravimetric determination of copper. Retrieved from: http://www.phywe.com/index.php/fuseaction/download/lrn_file/versuchsanleitungen/P3062201/e/P306 2201.pdf . [Accessed 22/10/2013]. University of Wrocław. (n.d.) Electrogravimetric Determination Of Copper In Alloys. Faculty of Chemistry, University of Wrocław, Analytical Chemistry Dept., electrogravimetry. Task 17 - p. 3 Vernier Software & Technology (2013) Determining Avogadro’s Number. Retrieved from: http://www.vernier.com/experiments/chem-a/31/determining_avogadros_number/ [Accessed 22/10/2013].
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