Experiment 4 Electrogravimetry - Determination Of Avogadro

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CHM674 – Advanced Electrochemistry

Experiment 4: Determination of Avogadro’s Number using Electrogravimetry Objective: The objective of this experiment is to make an experimental measurement of Avogadro's number using an electrochemical technique (electrogravimetry). Apparatus:

20 V - Power supply, U-tube, copper electrodes, electrical wires with crocodile clips, retort stand with clamps, emery paper, 4-decimal analytical balance.

Chemicals: 1M CuSO4 and 0.5M H2SO4 PRINCIPLES The atomic mass of an element in grams is equal to one mole of the element. Chemists used this definition of a mole long before they were able to measure the masses of individual atoms or had the means to count atoms. The determination of Avogadro's number, which is the number of particles in a mole, required the development of accurate and suitable measuring devices that were not in existence until the early part of the twentieth century. The mole is considered a fundamental unit has been adopted into the SI system as a basic unit of quantity. In this experiment we will make a careful measurement of electron flow, amperage, and time to obtain the number of electrons passing through the electrochemical cell. The electron flow, in amperes, is usually referred to as the current. The number of atoms in a weighed sample can be related to the number of electrons used and from that the value called Avogadro's number can be calculated. Avogadro's number can be determined in a number of different ways. This experiment will use an electrochemical process called electrolysis via electrogravimetric technique. The experimental setup for this process is called an electrolytic cell. An electrolytic cell is made up of the following: 1. 2.

a source of direct current such as a battery or power supply. insulated wires to carry the electric current.

3.

two electrodes. (In this experiment both electrodes are copper metal. The electrode connected to the negative (-) pin of the power supply is the cathode and the electrode connected to the positive (+) pin of the power supply is the anode.)

4.

a solution of sulfuric acid. (Sulfuric acid in this experiment is the conducting medium in the cell and is called the electrolyte.)

The electrolytic process is used to determine the number of electrons needed to convert one mole of copper atoms to one mole of copper ions, Cu2+. This number divided by two represents the number of atoms converted from copper metal to copper ions: Cu → Cu2+ + 2e 1

CHM674 – Advanced Electrochemistry

This process, which involves the loss of electrons, is called oxidation. The number of copper atoms per mole of copper is Avogadro's number, the value to be determined. The number of electrons consumed in the process is determined by using the charge of an electron and the total charge measured. The charge of an electron was determined in the famous Millikan oil-drop experiment to be 1.60217733 x 10-19 coulombs per electron. The number of coulombs used in the experiment can be calculated from the relationship: one ampere = 1 coulomb/second. An ammeter is used in the experiment to measure the amperage and a clock or stopwatch is used to measure the time in seconds. The mass of copper that reacts can be determined by measuring the mass of the anode and the cathode before and after the electrolysis. In the electrolytic cell contains both copper electrodes and the electrolyte is 1M CuSO 4 + 0.5 M H2SO4. In the course of the electrolysis, the copper electrode (the anode) connected to the positive pin of the power supply loses mass as the copper atoms are converted to copper ions. This loss of mass is visible to the eye after a while as pitting of the surface of the metal electrode. Also the copper ions, Cu2+ , produced, immediately pass into the water solution and deposit on the cathode as the following reaction: Cu2+ + 2 electrons → Cu (solid). It is also possible to measure the copper deposited and use it to calculate Avogadro's number. EXAMPLE Anode mass lost: 0.3554 grams (g) Current(average): 0.601 amperes (amp) Time of electrolysis: 1802 seconds (s) Note:

one ampere = 1 coulomb/second charge of an electron 1.602 x 10-19coulomb

or

Step 1. Find the total charge passed through the circuit.

(0.601 amp)(1coul/1amp-s)(1802 s) = 1083 coul

Step 2. Calculate the number of electrons in the electrolysis. 2

one

amp.s

=

1

coul

CHM674 – Advanced Electrochemistry

(1083 coul)( 1 electron/1.6022 x 10 -19coul) = 6.759 x 10 21electrons Step 3. Determine the number of copper atoms lost from the anode. Recall the electrolysis process consumes two electrons per copper ion formed. Therefore the number of copper(II) ions formed is half the number of electrons. Number of Cu+2 ions = number of electrons measured , namely,

(6.752 x 10 21 electrons)(1 Cu+2 / 2 electrons) = 3.380 x 10 21 Cu+2 ions Step 4. Calculate the number of copper ions per gram of copper from the number of copper ions above and the mass of copper ions produced. The mass of the copper ions produced is equal to the mass loss of the anode. ( The mass of the electrons is so small that it is negligible in this measurement, therefore the mass of the copper (II) ions is the same as copper atoms.) Thus: mass loss of electrode = mass of Cu +2 ions = 0.3554 g

3.380 x 10 21 Cu+2 ions / 0.3544g = 9.510 x 10 21 Cu+2 ions/g = 9.510 x 10 21 Cu atoms/g Step 5. Calculate the number of copper atoms in a mole of copper, 63.546 grams. Cu atoms /mole of Cu

=

(9.510 x 10 21 copper atoms/g copper)(63.546 g/mole copper)

=

6.040 x 10 23 copper atoms/mole of copper

This is the student's measured value of Avogaro's number.

Step 6. Calculate the percent error. Absolute error:

|6.02 x 10 23 - 6.04 x 10 23 |

=

2 x 10 21

Percent error:

(2 x 10 21 / 6.02 x 10 23)(100)

=

0.3 %

PROCEDURE Obtain two copper electrodes. It is may be necessary to first polish and clean the electrodes before any measurements are taken. 3

CHM674 – Advanced Electrochemistry

Do not touch the electrodes with your fingers. Dip the electrodes in a beaker of clean tap water then dip the electrode in the beaker labeled alcohol. Let the electrodes dry on a paper towel. When the electrodes are dry, put a sticker on each electrode before weighing. Weigh them carefully on the analytical balance to the nearest 0.0001 gram. Label the electrode with less weight as anode. The electrolytic solution is 1M CuSO4 in the 250-mL beaker. Before making any connections be sure the power supply is off and unplugged. The power supply must be set at 20 V. The correct sequence requires the positive pole of the power supply be connected to the anode of the first cell. The cathode is next connected to the positive pin of the ammeter. Have your apparatus approved by the instructor before you turn on the power! When the apparatus is approved, plug in the power supply. Make sure the power supply is in the off position. Accurate measurements of the time in seconds and the current in amperes are essential for good results. The amperage should be recorded at 30sec intervals for 10 min. The amperage may vary over the course of the experiment due to changes in the electrolyte solution, temperature, or position of the electrodes. The amperage used in the calculation should be an average of the readings taken. The current should flow a minimum of 10 mins. After 10 mins turn off the power supply record the last amperage value and the time. Now that the electrolysis has stopped you will need to retrieve the anode and cathode from the cell, rinse gently with distilled water. Dry them as before by immersing them in alcohol and allowing them to dry on a tissue paper or towel. DO NOT WIPE THE ANODE AND CATHODE WITH A TISSUE OR TOWEL. If you wipe them you will remove copper from the surface and invalidate your work. Repeat the experiment for second readings. Use the same electrodes, re-polish the electrodes with emery paper and re-weigh.

Repeat the electrolysis process using 0.5M H 2SO4 solution. Record your observations on what occuring at the electrodes and electrolyte solution. Caution-this solution is corrosive and will damage skin and clothing on contact.

ELECTROLYSIS DATA ELECTRODE MEASUREMENTS

4

CHM674 – Advanced Electrochemistry

Trial 1

Trial 2

Mass of anode before electrolysis

__________ g

__________ g

Mass of anode after electrolysis

__________ g

__________ g

Mass loss of anode

__________ g

__________ g

Trial 1

Trial 2

Mass of cathode before electrolysis

__________ g

__________ g

Mass of cathode after electrolysis

__________ g

__________ g

Mass gain at cathode

__________ g

__________ g

Average of weight loss at anode:

_____________g

Average of weight gain at cathode:

______________g

TIME-AMPERAGE MEASUREMENTS

Time (Secs)

Trial 1 Current (A) 5

Trial 2 Current (A)

CHM674 – Advanced Electrochemistry

0 30 60 90 120 150 180 210 240 270 300 330 360 390 420 450 480 510 540 570 600 Average Current SUMMARY OF DATA AND RESULTS Trial 1

Trial 2

Total time of electrolysis (seconds)

__________

__________

Average current during electrolysis (amperes)

__________

__________

Total charge measured (coulombs)

__________

__________

6

CHM674 – Advanced Electrochemistry

Number of electrons passed

__________

Number of Cu2+ ions generated

__________

__________

__________

Number of Cu2+ ions/gram Cu metal (Cu2+/g Cu) __________

__________

Based of weight loss of anode:

Avogadro's number (from your measurements)

__________

__________

Avogadro's number ( true or accepted value)

__________

__________

Absolute error in measured value

__________

__________

Relative % error in measured value

__________

__________

Based of weight gain of cathode: Number of Cu2+ ions/gram Cu metal (Cu2+/g Cu) __________

__________

Avogadro's number (from your measurements)

__________

__________

Avogadro's number ( true or accepted value)

__________

__________

Absolute error in measured value

__________

__________

Relative % error in measured value

__________

__________

Questions: 1. Calculate the charge passed efficiency, Q a/Qc (%) during the electrolysis. 2. Through your observations, differentiate between electrolysis of 0.5M H 2SO4 solution and 1M CuSO4 solution.

7

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