Bài giảng hoá phân tích electrolytic cells,

Experiment 33Electrolytic Cells, Avogadro’s Number • To identify the reactions occurring at the anode and cathode during the electrolysis of various aqueous salt solutions • To determine

Experiment 33

Electrolytic Cells, Avogadro’s Number

• To identify the reactions occurring at the anode and cathode during the electrolysis

of various aqueous salt solutions

• To determine Avogadro’s number and the Faraday constant

The following techniques are used in the Experimental Procedure:

H2(g)

O2(g)

A 1:2 mole (and volume) ratio of oxygen (left) to hydrogen (right) is produced from the electrolysis of water.

Objectives

Techniques

Introduction

Electrolysis processes are very important in achieving high standards of living The

industrial production of metals such as aluminum and magnesium and nonmetals such

as chlorine and uorine occurs in electrolytic cells The highly re ned copper metal

required for electrical wiring is obtained through an electroplating process.

In an electrolytic cell, the input of an electric current causes an otherwise

nonspon-taneous oxidation–reduction reaction, a nonsponnonspon-taneous transfer of electrons, to occur

For example, sodium metal, a very active metal, and chlorine gas, a very toxic gas, are

prepared industrially by the electrolysis of molten sodium chloride Electrical energy is

supplied to a molten NaCl system (Figure 33.1) by a direct current (dc) power source

(set at an appropriate voltage) across the electrodes of an electrolytic cell

The electrical energy causes the reduction of the sodium ion, Na⫹, at the cathode

and oxidation of the chloride ion, Cl⫺, at the anode Because cations

migrate to the cathode and anions migrate to the anode, the cathode is the

negative electrode (opposite charges attract), and the anode is designated

the positive electrode.1

Electrolysis reactions also occur in aqueous solutions For example,

in the electrolysis of an aqueous copper(II) bromide, CuBr2, solution,

cop-per(II) ions, Cu2⫹, are reduced at the cathode and bromide ions, Br⫺, are

oxidized at the anode (Figure 33.2, page 364)

cell reaction: Cu2⫹(aq) ⫹ 2 Br⫺(aq) l Cu(s) ⫹ Br2(l)

anode (⫹) reaction: 2 Br⫺(aq) l Br2(l) ⫹ 2 e⫺

cathode (⫺) reaction: Cu2⫹(aq) ⫹ 2 e⫺l Cu(s)

anode (cathode (⫹) reaction: 2 Cl⫺) reaction: Na⫺(l) ⫹(l) l Cl⫹ e2 ⫺(g) l Na(l) ⫹ 2 e⫺

Figure 33.1 Schematic diagram of the

electrolysis of molten sodium chloride

Electrolysis: Use of electrical energy

to cause a chemical reaction to occur Electroplating: the use of electrical current to deposit a metal onto an electrode

Electrolytic cell: an apparatus used for an electrolysis reaction

1 Note that the cathode is “ ⫺” in an electrolytic cell but “⫹” in a galvanic cell; the

anode is “ ⫹” in an electrolytic cell but “⫺” in a galvanic cell See Experiment 32.

In an aqueous solution, however, the reduction of water at the cathode (the negative electrode) and the oxidation of water at the anode (the positive electrode) are also pos-sible reactions

(33.1) (33.2)

If the reduction of water occurs at the cathode, hydrogen gas is evolved and the solution near the cathode becomes basic as a result of the production of hydroxide ion If oxidation of water occurs at the anode, oxygen gas is evolved and the solution near the anode becomes acidic The acidity (or basicity) near the respective electrodes can be detected with pH paper or another acid–base indicator

When two or more competing reduction reactions are possible at the

cathode, the reaction that occurs most easily (the one with the higher

reduction potential) is the one that usually occurs Conversely, for two or more competing oxidation reactions at the anode, the reaction that takes

place most easily (the one with the higher oxidation potential or the lower

reduction potential) is the one that usually occurs

In the electrolysis of the aqueous copper(II) bromide solution, Cu2⫹ has a higher reduction potential than H2O and is therefore preferentially reduced at the cathode; Br⫺ has a greater tendency to be oxidized than water, and so Br⫺is oxidized at the anode

In Part A of this experiment, a number of aqueous salt solutions using different electrodes are electrolyzed The anode and cathode are identi ed, and the products that are formed at each electrode are also identi ed

In Part B, a quantitative investigation of the electrolytic oxidation of copper metal is used to determine Avogadro’s number and the Faraday constant:

(33.3)

Two moles of electrons (or 2 faradays) are released for each mole of Cu(s)

oxi-dized; therefore, a mass measurement of the copper anode before and after the electrol-ysis determines the moles of copper that are oxidized This in turn is used to calculate the moles of electrons that pass through the cell:

(33.4)

The actual number of electrons that pass through the cell is calculated from the

elec-trical current, measured in amperes (⫽ coulombs/second), that passes through the cell for

a recorded time period (seconds) The total charge (coulombs, C) that passes through the

cell is

(33.5)

As the charge of one electron equals 1.60 ⫻ 10⫺19C, the number of electrons that pass through the cell can be calculated:

(33.6)

Therefore, since the number of electrons (equation 33.6) and the moles of electrons (equation 33.4) can be separately determined, Avogadro’s number is calculated as

(33.7) Avogadro’s number ⫽ number of electronsmole of electrons

number of electrons ⫽ number of coulombs ⫻ electron

1.60 ⫻ 10⫺19 C number of coulombs ⫽ coulombssecond ⫻ seconds

moles of electrons ⫽ mass Cu ⫻ mol Cu

63.54 g⫻ 2 mol emol Cu⫺

Cu(s) l Cu2⫹(aq) ⫹ 2 e⫺

anode (⫹) reaction for water: 2 H2O(l) l O2(g) ⫹ 4 H⫹(aq) ⫹ 4 e⫺

cathode (⫺) reaction for water: 2 H2O(l) ⫹ 2 e⫺l H2(g) ⫹ 2 OH⫺(aq)

Figure 33.2 Electrolysis of a copper(II)

bromide solution; the anode is on the left

and the cathode is on the right.

1 faraday ⫽ 1 mol e ⫺ ⫽ 96,485

coulombs

Coulomb: SI base unit for electrical

charge

Electrolysis of Aqueous

Solution

Avogadro’s Number and

the Faraday Constant

In addition, the number of coulombs (equation 33.5) per mole of electrons

(equa-tion 33.4) equals the Faraday constant With the available data, the Faraday constant

can also be calculated:

(33.8)

Procedure Overview: The products that result from the electrolysis of various salt

solutions are observed and identi ed; these are qualitative measurements An

experi-mental setup is designed to measure quantitatively the ow of current and consequent

changes in mass of the electrodes in an electrolytic cell; from these data, experimental

constants are calculated

The electrolysis apparatus may be designed differently than the one described in

this experiment Ask your instructor

1 Set up the electrolysis apparatus Connect two wire leads (different colors) attached

to alligator clips to a direct current (dc) power supply.2Clean and mount the glass

U-tube on a ring stand (see Figure 33.3) Connect the alligator clips to the

corre-sponding electrodes, listed in Table 33.1

2 Electrolyze the solutions Fill the U-tube three-fourths full with solution 1 from

Table 33.1 Insert the corresponding electrodes into the solution and electrolyze

for ⬃5 minutes During the electrolysis, watch for any evidence of a reaction in

the anode and cathode chambers

• Does the pH of the solution change at each electrode? Test each chamber with

litmus or pH paper.3Compare the color with a pH test on the original solution

• Is a gas evolved at either or both electrodes?

• Look closely at each electrode Is a metal depositing on the electrode or is the

metal electrode slowly disappearing?

3 Account for your observations Write the equations for the reactions occurring at

the anode and cathode and for the cell reaction Repeat for solutions 2–5

CLEANUP: Rinse the U-tube twice with tap water and twice with

deion-ized water before preparing the next solution for electrolysis Discard each

rinse in the sink

Disposal: Discard the salt solutions into the Waste Salts container

Faraday constant ⫽ number of coulombsmole of electrons

Experimental Procedure

A Electrolysis of Aqueous Salt Solutions

2 The dc power supply can be a 9-V transistor battery.

3 Several drops of universal indicator can be added to the solution in both chambers to detect pH

changes.

Figure 33.3 Electrolysis

apparatus

Table 33.1 Electrolytic Cells for Study

Solution No Solution* Electrodes (Cathode and Anode)

1 2 g NaCl/100 mL Carbon (graphite)

2 2 g NaBr/100 mL Carbon (graphite)

3 2 g KI/100 mL Carbon (graphite)

5 0.1 M CuSO4 Polished copper metal strips

*Try other solutions and electrodes as suggested by your laboratory instructor.

1 Set up the apparatus Refer to Figure 33.4 The U-tube from Part A can again be used The dc power supply must provide 3–5 volts (two or three ashlight batter-ies in serbatter-ies or a lantern battery); the ammeter or multimeter must read from 0.2 to 1.0 A Polish two copper metal strips (to be used as the electrodes) with steel wool

or sandpaper Brie y dip each electrode (use the fume hood) into 6 M HNO3

(Caution: do not allow skin contact) for further cleaning, and then rinse with

deionized water

Add 100 mL of 1.0 M CuSO4(in 0.1 M H2SO4) to the 150-mL beaker (or ll the U-tube)

2 Set the electrodes.Rinse the electrodes with ethanol if available When dry, label the two electrodes because the mass of each will be determined before and after the electrolysis Measure the mass (Ⳳ0.001 g, preferably Ⳳ0.0001 g) of each labeled electrode The copper electrode with the lesser mass is to serve as the anode (⫹ ter-minal), and the other is to serve as the cathode (⫺ terminal) for the electrolytic cell Connect the cathode (through the variable resistor and ammeter/multimeter) to the negative terminal of the dc power supply

Before electrolysis begins, obtain your instructor’s approval of the complete apparatus

3 Electrolyze the CuSO 4 solution.Adjust the variable resistance to its maximum value.4Be ready to start timing (a stopwatch is ideal) Attach the anode to the pos-itive terminal of the dc power supply and START TIME During the electrolysis,

do not move the electrodes; this changes current ow Adjust the current with the

variable resistor to about 0.5 A and, periodically during the course of the electrol-ysis, readjust the current to 0.5 A.5

Discontinue the electrolysis after 20–30 minutes Record the exact time (min-utes and seconds) of the electrolysis process

4 Dry and measure the mass.Carefully remove the electrodes (be careful not to loosen the electroplated copper metal from the cathode); carefully dip each elec-trode into a 400-mL beaker of deionized water to rinse the elecelec-trodes (followed

by ethanol if available) Air-dry, measure the mass (Ⳳ0.001 g, preferably Ⳳ0.0001 g) of each electrode, and record

5 Repeat the electrolysis If time allows, repeat Part B using the same copper

electrodes (with new mass measurements!) and 1.0 M CuSO4solution

CLEANUP: Rinse the beaker or U-tube twice with tap water and twice with deionized water Discard each rinse as directed

by your instructor

Electroplating of metals such as nickel, chromium, silver, and copper is a common industrial process Research a speci c process and design an apparatus and procedure for depositing quantitative amounts of metal to a cathode

Disposal: Discard the copper(II) sulfate solution into the Waste Salts container

Figure 33.4 Setup for

determining Avogadro’s number

and the Faraday constant

4 If a variable resistor is unavailable, record the current at 1-minute intervals and then calculate an average current over the entire electrolysis time period.

5 If the current is greater or less than 0.5 A, vary the time of electrolysis proportionally.

The Next Step

B Determination of

Avogadro’s Number and

the Faraday Constant

Experiment 33 Prelaboratory Assignment

Electrolytic Cells, Avogadro’s Number

Date Lab Sec Name Desk No

1. The standard reduction potential for the Cu2⫹/Cu redox couple is ⫹0.34 V; that for H2O/H2, OH⫺at a pH of 7 is

⫺0.41 V For the electrolysis of a neutral 1.0 M CuSO4solution, write the equation for the half-reaction occurring at the cathode at standard conditions

2. In an electrolytic cell,

a. oxidation occurs at the (name of electrode) _

b. the cathode is the (sign) electrode _

c. cations ow toward the (name of electrode) _

d. electrons ow from the (name of electrode) to (name of electrode) _ _

e. the anode should be connected to the (positive/negative) terminal

3. Experimental Procedure, Part A.2 Describe the proper technique for testing the pH of a solution with

litmus or pH paper

4 a. Identify a chemical test(s) to determine if water is oxidized at the anode in an electrolytic cell

b. Similarly, identify a chemical test(s) to determine if water is reduced at the cathode in an electrolytic cell

5. Very pure copper metal is produced by the electrolytic re ning of

blister (impure) copper In the cell at right, label the anode, the

cath-ode, and the polarity (⫹, ⫺) of each

6 a. When a solution of sodium sulfate, Na2SO4, adjusted to a pH of 7 is electrolyzed, red litmus remains red and a gas

is evolved in the anodic chamber In the cathodic chamber, red litmus turns blue and a gas is also evolved

(i) Write a balanced equation for the half-reaction occurring at the anode

(ii) Write a balanced equation for the half-reaction occurring at the cathode

b. When a solution of nickel(II) sulfate adjusted to a pH of 7 is electrolyzed, the green color of the solution becomes less intense in the cathodic chamber and gas bubbles are detected in the anodic chamber

(i) Write a balanced equation for the half-reaction occurring at the anode

(ii) Write a balanced equation for the half-reaction occurring at the cathode

7. A 1.0 M CuSO4solution was electrolyzed for 28 minutes and 22 seconds using copper electrodes The average current owing through the cell over the time period was 0.622 A The mass of the copper anode before the electrolysis was 2.4852 g and 2.1335 g afterwards

a. Calculate the number of moles of copper oxidized and moles of electrons that passed through the cell (see equation 33.4).

b. Calculate the total charge (coulombs) that passed through the cell (see equation 33.5)

c. How many electrons passed through the cell during the 28.0 minute, 22 second time period (see equation 33.6)?

d. From the data, calculate Avogadro’s number (see equation 33.7)

Experiment 33 Report Sheet

Electrolytic Cells, Avogadro’s Number

Date Lab Sec Name Desk No

A Electrolysis of Aqueous Salt Solutions

Gas Solution Electrodes Litmus Test Evolved? Balanced Equations for Reactions

Cathode _

Cell

Cathode _

Cell

Cathode _

Cell

Cathode _

Cell

Cathode _

Cell

B Determination of Avogadro’s Number and the Faraday Constant

1 Initial mass of copper anode (g)

2 Initial mass of copper cathode (g)

3. Instructor’s approval of apparatus

4. Time of electrolysis (s)

5. Current (or average current) (A)

6 Final mass of copper anode (g)

7 Final mass of copper cathode (g)

Data Analysis

1. Mass of copper oxidized at anode (g)

2. Moles of copper oxidized (mol)

3. Moles of electrons transferred (mol e⫺)

4. Coulombs passed through cell (C)

5. Electrons passed through cell (e⫺)

6. Avogadro’s number (e⫺/mol e⫺)

7. Average value of Avogadro’s number

8. Literature value of Avogadro’s number

10. Faraday constant (C/mol e⫺)

11. Average Faraday constant (C/mol e⫺)

12. Literature value of Faraday constant

Laboratory Questions

Circle the questions that have been assigned

1. Part A.2 If zinc electrodes are used instead of the graphite electrodes, the reaction occurring at the anode may be dif-ferent, but the reaction occurring at the cathode would remain unchanged Explain

2. Part A.2 Nitrate ions, NO3⫺, being anions, migrate to the anode in an electrolytic cell Explain why you would expect

water rather than nitrate ions to be oxidized at the anode Hint: Consider the oxidation state of nitrogen in the nitrate ion.

3. Part B Repeat the calculation of Avogadro’s number, using the mass gain of the cathode instead of the mass loss of

the anode Account for any difference in the calculated values

4. Part B.2 If the current is recorded as being less than it actually is, would Avogadro’s number be calculated as too high

or too low, or would it be unaffected? Explain

5. Part B.4 Because of an impure copper anode (see Prelaboratory Assignment question 5), the measured mass loss is greater

than the actual mass of copper oxidized As a result, will Avogadro’s number be calculated as too high or too low? Explain

*6 The electrolytic re ning of copper involves the oxidaton of impure copper containing such metals as iron and nickel

(oxidized to copper(II), iron(II), and nickel(II) ions) at the anode and then reduction of the copper(II) ion to copper metal at the cathode Explain why the iron(II) and nickel(II) ions are not deposited on the cathode

Appendix B

Appendix B