Electrolysis the faraday and avogadro's number

Student name: Cosumnes River College Due date: Tuesday, Dec 15, 2020, by 11:59 PM Prof. Hoang Chem401 Lab: Faraday’s Constant and Avogadro’s Number INTRODUCTION Zn/Cu Voltaic Cell: In a voltaic cell consisting of Zn and Cu electrodes as shown in Fig. 1, Zn is the stronger reducing agent than Cu and hence will be oxidized as it reduces Cu 2+ ion to metallic Cu. The Zn electrode is the negatively-labelled anode and the Cu electrode is the positively-labelled cathode. Electron current flows spontaneously through an external circuit from the Zn electrode to the Cu electrode. Figure 1 Zn/Cu Electrolytic Cell: When a battery is wired to the external circuit that connects the Zn and Cu electrodes, battery power can be applied to cause a current to flow in reverse from the Cu electrode to the Zn electrode as shown in Fig. 2. This is now an electrolytic cell. e‒ e‒ Zn2+ + 2 e‒→ Zn(s) Figure 2 Cu(s) → Cu2+ + 2 e‒ 1 In an electrolytic cell, the current flow is nonspontaneous and requires a continuous input of an external source of power to occur. Unlike a voltaic cell, the anode of an electrolytic cell is labelled as positive electrode and the cathode as negative electrode. Electrolysis has many applications, including plating of metals on surfaces and production of pure elements from mining ores or other compounds. As in a voltaic cell, oxidation occurs at the anode and reduction occurs at the cathode in an electrolytic cell. The amount of reaction that occurs at the electrodes is directly proportional to the number of electrons transferred. Michael Faraday (1791-1867) determined that the mass of a substance produced or consumed at the electrodes during electrolysis is proportional to the quantity of charge [Charge (C) = Current (A) x Time (sec)] that has passed through the circuit. A faraday is defined as the total charge (in Coulombs) on 1 mole of electrons and is equal to 96,485 C. Hence, the charge on a single electron is 1.602 x 10−19 C. In this electrolysis experiment, you will determine (1) the experimental value of the Faraday’s Constant by measuring the amount of charge required to consume a known mass of copper on an electrode. From the mass of copper consumed (or electrolyzed), you will also be able to determine (2) the experimental value of the Avogadro’s number (NA). You will use a copper strip and a zinc strip as the electrodes. These are placed in a beaker of dilute sulfuric acid (H2SO4) as shown in Fig 3. The strong acid dissociates into ions in solution and thus allows a current to be conducted between the electrodes. A Direct-Current (DC) Power Supply is used as a source of power to drive the electrolytic cell. Using the DC power supply, an electric current of specific amperage can be set. DC power supply e‒ anode cathode Figure 3 At the anode, the following oxidation reaction takes place: Cu(s) → Cu2+(aq) + 2 e‒ By determining the average current used in the reaction, along with the knowledge that all copper ions formed are the 2+ cations, you will calculate the number of atoms in one mole of copper (i.e. experimental Avogadro’s number, NA) and compare this value with the accepted Avogadro’s number (NA = 6.022 x 1023). PROCEDURE A description of the procedure and sample calculations are given below to help you understand how the data are generated and used to achieve the objectives of this lab. 1. Two dry metal strips, one is made of copper and the other of zinc, are cleaned by scrubbing their surface with steel wool to remove any surface oxidation. The Cu dry metal strip is placed in an analytical balance and its mass recorded. 2 2. The metal strips are connected to a DC Power Supply via wires with alligator clips as shown in Fig. 4. The Cu metal strip acts as the anode and is connected to the Power Supply at the positive (+) terminal. The Zn metal strip acts as the cathode and is connected to the Power Supply at the negative (‒) terminal. [Note: In an electrolytic cell, the anode is where electrons are pulled out of the electrode into the Power Supply and is given a (+) sign, and the cathode is where electrons are pushed out of the Power Supply and flow toward]. The sample initial dry masses of the Zn and Cu strips are provided on the diagram, (which I borrowed from an internet source for illustration). The mass of the Zn strip and the voltage shown are irrelevant for this lab. The relevant data are the initial mass of the Cu strip (17.3527 g), the current used (0.4004 A), and the total time (3.00 minutes) of electrolysis. Figure 4 17.8551 g ‒ + 1 M H2SO4 17.3527 g The Electrolytic Cell 3. The metal electrodes are placed in 1 M H2SO4 contained in a large glass beaker and are set far apart from one another to prevent accidental contact. The experiment is performed in 3 trials, each time the electrodes are immersed in the H2SO4 solution to equal depths. The Power Supply is adjusted to deliver a current within the range of 0.2‒0.6 A to the cell for a period of exactly 3.00 minutes. In the sample trial shown in Fig 4, the current recorded as 0.4004 A. 4. After 3 minutes of electrolysis, the power is turned off. The Cu strip is removed and rinsed in distilled H2O using a wash bottle, and gently blotted dry with paper towels. After being allowed to air dry completely, the Cu strip is placed in the analytical balance and its mass recorded. 3 DATA ANALYSIS AND SAMPLE CALCULATIONS During electrolysis, the Cu electrode is the anode and therefore oxidized: Cu(s) → Cu2+ + 2e‒ As a result of oxidation, its mass decreases from the initial 17.3527 g to 17.3247 g.