Name__________________________________________ Section_______________________ Partner(s)______________________________________________________________________ Chemistry 104 Laboratory University of Massachusetts Boston

AN ELECTROCHEMISTRY EXPERIMENT PRELAB ASSIGNMENT 1. Calculate equivalent masses for each metal in Table 1, and write these into that table. 2. Fill in all the blanks in the Prelab Data Sheet on page 5.

In this experiment we assemble an electrolytic cell in which one electrode is a strip of unknown metal, M, the other electrode is an inert metal (copper), and the electrolyte is an acidic solution. As electrolysis proceeds, H+(aq) ions from the solution acquire electrons, forming H2(g) at the inert copper electrode, and M(s) atoms in the strip of unknown metal lose electrons, thereby entering the solution as Mp+(aq) ions, where p+ is the ionic charge. The overall cell reaction is M(s) + p H+(aq) 6 Mp+(aq) + (p/2) H2(g) where the stoichiometric coefficients on both hydrogen species involve the integer p, equal to the number of electrons in the metal’s oxidation half-reaction, M(s) º Mp+(aq) + p e–. In our experiment, the metal M(s) will be one of those listed in Table 1 on page 2. From the listed reduction half-reactions in Table 1, you can see that p must be 1, 2, or 3. The number of electrons in the balanced reduction or oxidation half-reaction of a species relates to the definition of its equivalent mass. The equivalent mass is the mass of a redox species that reacts or is formed when exactly one mole of electrons is passed through an electrochemical cell. The amount of charge carried by one mole of electrons is defined as the faraday (ö). For example, one mole of H2(g), having a mass of 2.016 g, is oxidized to H+(aq) in the half-reaction H2(g) 6 2 H+(aq) + 2 e– by the transfer of 2 moles of electrons (2 ö). Thus the equivalent mass of H2(g) is 2.016 g/2 ö = 1.008 g/ö. The amount of charge, Q, in coulombs that is passed in an operating electrochemical cell can be calculated using the relationship Q = it (1) where i is the constant current in amperes (A) passed for a time t in seconds. In 1960 the National Bureau of Standards experimentally determined the chemical-scale faraday to be equivalent to 96,489 " 2 coulombs (C). For our purposes, the charge in coulombs calculated by Q = it can be converted to faradays, using the relationship 1 ö = 96,500 C. 1

Table 1. Electrochemical Data for Selected Metals Reduction Half-Reaction

EE (V)

at. wt. (u)

Al3+(aq) + 3 e– º Al(s)

–1.66

26.98

Zn (aq) + 2 e º Zn(s)

–0.76

65.39

Cr3+(aq) + 3 e– º Cr(s)

–0.74

52.00

Ni2+(aq) + 2 e– º Ni(s)

–0.23

58.69

Sn2+(aq) + 2 e– º Sn(s)

–0.14

118.7

Pb2+(aq) + 2 e– º Pb(s)

–0.13

207.2

Bi3+(aq) + 3 e– º Bi(s)

+0.20

209.0

Ag+(aq) + e– º Ag(s)

+0.80

107.9

2+

–

Equivalent mass (g/ö)

IN THE LABORATORY We will attempt to identify the metal M by calculating its equivalent mass from the mass of M(s) converted into Mp+(aq) ions, and from the number of faradays determined from the amount of H2(g) generated. As shown in the sketch, the apparatus consists of ! a small beaker ! an inverted buret ! an inert copper electrode connected to an electrical lead ! an alligator clip mounted on a wood dowel and connected to another lead ! a strip of metal, M, to be held in the alligator clip ! a plug and receptacle connected to a direct current power source Does the buret leak? The first objective is to test the buret for an air leak. Proceed as follows: G Obtain at least 50 mL of water in a small beaker. G Mount the buret upside down in the buret stand with the open end in the water, just above the bottom of the beaker. G Use a pipet bulb to draw water up into the buret. G Close the stopcock and observe the meniscus level. G If the level drops within a couple of minutes, consult your instructor.

2

An Electrochemistry Experiment Set up for electrolysis G Rinse a 150-mL beaker, the buret, and the copper electrode with deionized water. G Use a graduated cylinder to add 100 mL of deionized water to the beaker. Then add 10 mL of sulfuric acid reagent from the dispenser. Record the volume and molarity of the sulfuric acid and the volume of water on the data sheet. WARNING If you get any sulfuric acid on your skin, clothes, or the benchtop, IMMEDIATELY wash or sponge this off with plenty of water. Inform your instructor about the spill. G Insert the copper electrode into the buret so that at least 1 inch of lead wire is up inside. (See sketch on page 2.) G Mount the buret upside down in the buret stand. Position the beaker under the buret and lower the buret almost to the bottom of the beaker. G Use the pipet bulb to draw sulfuric acid solution up into the buret so that the meniscus is within 1 mL of the top graduation mark. Read and record this initial meniscus level to 0.01 mL. G Obtain a strip of the unknown metal, M. This must be clean and dry. If necessary, polish both sides of the strip with steel wool. G Weigh the strip on an analytical balance and record the mass on the data sheet. G Attach the strip to the very tip of the alligator clip, and lower this assembly into the beaker so that the bottom of the strip is about 1/2" below the surface of the liquid. Starting the electrolysis You may wish to ask your instructor to check your apparatus, but try to proceed with starting the electrolysis as soon as possible after immersing the electrodes in the electrolyte solution. Prepare to time the electrolysis with a watch that reads in seconds or the wall clock in the laboratory. G Record the laboratory’s barometric pressure and temperature (posted on the board). G Connect the leads to the power source and begin timing. G Stop the electrolysis when the liquid level in the buret is near but above the lip of the beaker by disconnecting from the power source. Record the time. G When all the bubbles have risen, record the final meniscus level in the buret. Recovering the remaining metal, disposing the sulfuric acid solution, and cleaning up G Open the stopcock to release the remaining liquid in the buret. G Raise the buret and electrodes above the liquid level in the beaker. G Obtain an empty 150-mL beaker (or larger). Carefully remove the filled beaker and replace it with the empty beaker under the apparatus so as to avoid dripping acid solution on the benchtop. G Thoroughly rinse the metal strip, M, in a stream of deionized water. Then remove it from 3

An Electrochemistry Experiment G G G G

the alligator clip and set it aside on a paper towel. Empty the beaker with the sulfuric acid solution into the designated waste container. Carry the entire apparatus to a deep sink. Wearing protective gloves, remove the copper electrode from the buret. Thoroughly rinse the copper electrode, wires, and buret in a stream of running tap water.

How much metal, M, was converted to M p+? “ Dry the strip and polish it with a paper towel (not steel wool). “ Weigh the clean and dry strip on an analytical balance. Record the final mass. CALCULATIONS “ Calculate the mass of metal, M, lost from the strip. “ Calculate the volume change in the buret from the hydrogen gas generated during the electrolysis. “ Record the vapor pressure of water at the room’s temperature. (If a water vapor pressure table is not posted in the laboratory, consult your textbook.) “ Calculate the moles of hydrogen gas (assuming ideal behavior) generated during the electrolysis as follows: ! Calculate the partial pressure of H2(g) from the barometric pressure in the laboratory and the vapor pressure of water at the room’s temperature. ! The volume of H2(g) generated is the volume change in the buret. ! Assume that the temperature of the gas inside the buret is the room’s temperature. “ Calculate the number of faradays required to generate this amount of hydrogen gas. “ Considering that this same number of faradays converted the metal, M, into Mp+ ions, the equivalent mass of M is the mass lost from the strip divided by the number of faradays. Use this calculated equivalent mass to identify the metal, M, from those listed in Table 1. “ Write the two half-reactions and the overall cell reaction, using proper chemical symbols in place of M, for the electrochemical reaction that took place in the electrolysis. “ Calculate the standard cell potential, Eocell, using the standard reduction potentials, Eo, listed in Table 1.

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An Electrochemistry Experiment “ Calculate the final molar concentrations of the cations in the electrolyte by assuming ! the volume of electrolyte is the sum of the liquid volumes added; ! the moles of Mp+ ions relate to the moles of H2(g) through the balanced overall cell equation; ! only one H+(aq) ion dissociates from the added H2SO4(aq), and some of this is converted by the electrolysis to H2(g). “ Calculate the average electric current, i, in amperes flowing during the electrolysis, using 1 ö = 96,500 C. PRELAB PRACTICE DATA SHEET Electrolyte:

10.

mL of

1.8

M H2SO4 and

Buret reading:

Initial

49.62 mL

Final

2.82

Mass of M:

Initial

0.4845

Final

0.2222

g

Time duration of electrolysis:

962

Barometric pressure:

torr

763.5

Moles of H2(g): Equivalent mass of M:

mL

mL of water Change

g

Loss

Room temperature:

22.0

mL g

s

Vapor pressure of water at this temperature: Partial pressure of H2(g):

100.

EC

torr

torr =

atm

mol

ö

Faradays: Identity of M and Mp+:

g/ö

and

Use the correct chemical symbol for M in the following reactions: Anode half reaction:

____________________________________________________

Cathode half reaction:

____________________________________________________

Overall cell reaction:

____________________________________________________

Standard cell potential, Eocell: Final concentrations: [Mp+] = Average electric current flow:

V [H+] =

M A

5

M

An Electrochemistry Experiment Name__________________________________________ Section_______________________ Partner(s)______________________________________________________________________ EXPERIMENTAL DATA SHEET Electrolyte:

mL of

M H2SO4 and

mL of water

Buret reading:

Initial

mL

Final

mL

Mass of M:

Initial

g

Final

g

Time duration of electrolysis: Barometric pressure:

torr

Partial pressure of H2(g):

Equivalent mass of M:

mL

Loss

g

s EC

Room temperature:

Vapor pressure of water at this temperature:

Moles of H2(g):

Change

torr

torr =

atm

mol

ö

Faradays: Identity of M and Mp+:

g/ö

and

Use the correct chemical symbol for M in the following reactions: Anode half reaction:

____________________________________________________

Cathode half reaction:

____________________________________________________

Overall cell reaction:

____________________________________________________

Standard cell potential, Eocell: Final concentrations: [Mp+] = Average electric current flow:

V [H+] =

M A

6

M