electrochemistry -...

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L A B O R A T O R Y 7

Electrochemistry

OBJECTIVE

To determine the identity of an unknown metal via electrochemical means.

Reference

Chemistry, 8e, Zumdahl & Zumdahl, Chapter 18.

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49 E L E C T R O C H E M I S T R Y

BEFORE COMING TO LAB

• Read through Lab #7: Electrochemistry.

• Complete the Pre-Lab Assignment on Lon-Capa. This assignment will lead you through

many of the calculations you will need to do in this lab.

• Prepare your laboratory notebook with Title and Date. Make sure to determine the equivalent

masses of the metals given in the Introduction to the lab.

• Reproduce Tables 7-1 and 7-2 in your lab notebook.

INTRODUCTION

In this experiment you will receive an unknown metal from your TA. You will determine its chemical

identity by constructing galvanic (voltaic) and electrolytic cells with this metal and making

measurements.

In electrochemical reactions there is a transfer of electrons from one chemical species to another.

The substance losing electrons is oxidized during the reaction. The species gaining electrons is

reduced. When these two reactions are run together, it is called an electrochemical cell.

An example is:

Cu(s) + 2Ag+(aq) 2Ag(s) + Cu2+(aq)

In this case, the silver(I) ions are being reduced, gaining electrons, to form solid silver. This is the

reduction half-reaction, which consumes electrons as one of the reactants.

Ag+(aq) + e– Ag(s) reduction

The solid copper is being oxidized, losing electrons, to form copper(II) ions. This is called the oxi-

dation half-reaction, the reaction which produces electrons.

Cu(s) Cu2+(aq) + 2e– oxidation

When solid copper is immersed in a 1.0 M Ag+ solution, it dissolves to form Cu2+ ions and solid

silver is precipitated from the Ag+ solution. This happens because the free energy change for this

reaction pair is negative, G < 0.

When working with electrochemical reactions, a different function, the reduction potential, , is

more commonly used.

It is related to free energy by the equation:

G = –nF

where n is the number of electrons transferred in the balanced half-reactions and F is Faraday's con-

stant, the charge on one mole of electrons, approximately 96,500 coulombs (C) per mole of electrons.

The reduction potential, , is a measure, in volts (V), of the driving force of a reduction reaction.

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For the silver half-reaction, on the previous page, the standard reduction potential is 0.80 V.

Ag+(aq) + e– Ag(s) o = +0.80 V reduction

The positive value for o indicates that this is a favorable reaction. Silver(I) has a strong tendency to

be reduced. Ag(s) is more stable than Ag+.

The standard reduction potential of the Cu/Cu2+ system is:

Cu2+(aq) + 2e– Cu(s) o = +0.34 V reduction

The positive o indicates that this is also a favorable reaction, but not as strongly driven as the Ag+

reduction (o = +0.80 V).

This is why the copper half-reaction is reversed, to become an oxidation reaction when paired with

the Ag/Ag+ reaction:

Cu(s) Cu2+(aq) + 2e– o = –0.34 V oxidation

Notice that the sign of o changes when the reduction reaction is converted to an oxidation reaction.

The copper half-reaction is a two electron process.

The silver half-reaction is a one electron process.

The silver half-reaction must be multiplied by a factor of two to equalize the number of electrons

produced and consumed.

2Ag+(aq) + 2e– 2Ag(s) o = +0.80 V reduction

Notice that the reduction potential does not change. It is an intensive property that is independent

of the stoichiometry.

Combining the two half-reactions gives:

Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s) o = 0.80 – 0.34 = 0.46 V

Because the standard cell potential, o, is positive, this pair of reactions will proceed spontaneously

as written.

To verify this, we can calculate the standard free energy change for this reaction:

Go = –n F o = –(2 mol e–) (96,500 C/mole e–) (0.46 V) = –88,780 VC = –88.8 kJ

This standard reaction will proceed spontaneously (Go < 0) because the standard potential

(o) is positive. This is a galvanic or voltaic cell. Both terms describe cells with positive cell

potentials.

Simply placing Cu(s) into a 1.0 M Ag+ solution will allow the reaction to proceed, but will not allow

us to measure, or use, the flow of electrons created by this reaction pair. To do that, the two half-

reactions are combined in an electrochemical cell, in which the two half reactions are carried out

in separate containers.

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The components of an electrochemical cell are:

1. Two compartments connected by a “salt bridge” through which ions can flow. In this

experiment, the salt bridge will be filter paper soaked in KNO3(aq).

2. Two solid electrodes, connected by a wire. The electrodes do not necessarily participate in

the chemical reactions. An electrode that does not itself undergo oxidation or reduction

is called an inert electrode. Inert electrodes are often made of graphite, C(s), or of an

unreactive metal like platinum.

3. Two solutions of electrolytes in which the electrodes are immersed. The ions in solution

may participate in the half-reactions or they may be inert electrolytes which simply

carry charge.

Standard Electrochemical Cells

A standard electrochemical cell using the Ag/Cu half-reactions is diagrammed below. The definition

of standard is that all components, reactants and products, must be present at an activity of 1. This

means that solid Cu and Ag must both be present, and that there must be Cu2+ and Ag+ solutions

at 1.0 M concentrations. If gases were involved, they would need to be present at 1 atm pressure.

salt bridge

Cu anode Ag cathode

1 M Cu2+

1 M Ag+

There is a shorthand notation that can be used to describe this electrochemical cell.

Cu(s) | Cu2+(l M) || Ag+(l M) | Ag(s)

By convention, the oxidation reaction is shown on the left. This is the anode. The element furthest

to the left is the solid electrode, in this case, Cu(s). The single vertical line ( | ) represents a phase

change. In this case, a piece of solid copper is immersed in a 1 M solution of Cu2+. The double line

( || ) represents the salt bridge connecting the two half cells.

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The reduction reaction is shown on the right and this half cell is called the cathode. Again, there is

a solid electrode, Ag(s), immersed in a 1.0 M solution of Ag+(aq).

Under these conditions, the voltmeter would register the standard reduction potential of 0.46 volts:

o o o

cell = red – ox = 0.80 – 0.34 = +0.46 V

Using Electricity to Speed Up the Reaction

When metals (Mo) are exposed to acids, a redox reaction takes place. The metals are oxidized and

the hydrogen ions are reduced. The half reactions are summarized below:

M(s) Mx+(aq) + xe– oxidation

2H+(aq) + 2e– H2(g) reduction

With the more reactive (Group I) metals this reaction proceeds rapidly. With less reactive metals,

like the ones listed in the chart below, this is a slower process.

Reduction o (V) Atomic Mass (g/mol) Equivalent Mass (g/mol e–)

Al3+ + 3e– Alo –1 .66 26 .98

Zn2+ + 2e– Zno –0 .76 65 .39

Fe2+ + 2e– Feo –0 .41 55 .85

Sn2+ + 2e– Sno –0 .14 118 .7

Pb2+ + 2e– Pbo –0 .13 207 .2

Under these experimental conditions, a power supply is used to speed up the redox chemistry.

The conventions for the anode and cathode are the same as for galvanic cells; oxidation occurs at the

anode and reduction occurs at the cathode.

Because we are not trying to measure the spontaneous flow of electrons, all of these components can

be placed in the same container.

The anode of the power supply is connected to the solid metal which is then oxidized, forming Mx+

ions.

An inert (nonreacting) cathode reduces the acidic protons (H+) producing hydrogen gas.

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You will carefully measure the mass of the solid metal before and after the reaction takes place. This

will give you the grams of metal oxidized during the experiment.

You will measure the volume of H2 gas collected during the experiment. Using the Ideal Gas

law this can be converted to moles of H2 and from the stoichiometry of the reaction, to moles of

electrons.

With this information, you can calculate the equivalent mass of the unknown metal. This is defined

as the molar mass of a metal divided by the number of moles of electrons consumed in its oxidation

reaction.

Comparing the experimental value to those in the chart will help to identify the unknown metal. In

some cases, your results may be ambiguous, with more than one possible metal. But by combining

this information with the results of the Standard Galvanic Cell obtained in Part 1, you should be

able to identify your unknown metal sample.

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You will using a Labquest interface voltmeter to measure cell potentials.

Caution: All heavy metals are toxic. You should wear gloves during all parts of this experiment.

Part 1: Galvanic Cell

Standard Cells

The reduction potentials of Cu2+ and Fe2+ are positive, while all of the others are negative. This means that Cu2+ and Fe2+

will be spontaneously reduced, and all of the other metals listed will be spontaneously oxidized, when coupled with

Cu/Cu2+ in a standard galvanic cell .

For a standard cell you will need to construct two half cells.

Procedure:

1. To prepare the cathode, the site of reduction (Cu2+ + 2e– Cu o = 0 .34 V):

a. Obtain approximately 15 mL of 1.0 M solution of Cu2+

b. Place a solid piece of Cu metal in the beaker

2. To prepare the anode, the side of oxidation (M Mx+ + xe– o = ? V):

a. Obtain approximately 15 mL of 1.0 M solution of M*+

b. Place a solid piece of the unknown metal, M, in the beaker

3. Attach the alligator clips to the two metal electrodes. Record the potential reading in your lab notebook.

4. To prepare the salt bridge: soak a piece of filter paper in a solution of KNO3. Then, place the two ends

of the filter paper so that they are touching the solutions in both beakers. This is the salt bridge. It allows

the passage of anions to the anode and cations to the cathode to maintain electroneutrality.

5. Record the potential of your standard cell, in your lab notebook (ocell = V), after the addition of the

salt bridge. The potential should now be positive. If the potential is negative, simply reverse the two

leads.

6. You have now created the standard cell. Fill in the components of this cell in the line notation provided

below. Record this information into your lab notebook.

(s) | ( M) || ( M) | (s)

Anode Cathode

7. Based on the measured standard potential, you should be able to deduce the identity of the unknown

metal. In some cases, it may not be obvious which metal you have. You will use the results of the

Electrolyic Cell experiment (Part 2) to help you decide.


Part 2: Using Electricity to Speed Up the Reaction

The experimental setup is shown to the right and consists of:

a 150 mL beaker with 0 .5 M acetic acid (HC2H3O2 ) and

0 .5 M S o d i u m S u l f a t e ( Na2SO4 ). This solution is

Labeled “Electrochem Solution.”

a buret inverted in the acid solution

a wire with a piece of copper to act as the cathode

the unknown metal electrode to act as the anode

a power supply to speed up the reactions

Procedure to Set-up the Equipment:

1. Rinse a 150 mL beaker, the buret, the copper electrode,

and your unknown metal sample with deionized water.

2. With the power supply unplugged, attach one end of the small

copper electrode to the cathode of the power supply. This should

be the smaller of the two alligator clips.

3. Mount the buret upside down, using the buret clamp. Position the

buret so that it almost touches the bottom of the 150 mL beaker.

4. Fill the beaker with the solution labeled “Electrochem Solution”

5. Open the stopcock of the buret and use the pipet bulb to draw

up the electrochem solution. You may need to close the stopcock,

re-compress the bulb, place it on top of the buret, reopen the

stopcock, and draw up more solution a few times. When you are

finished, the electrochem solution should be within 1 mL of the

top graduation mark on the buret.

6. Record the volume of the buret to the nearest 0.001 g.

7. Record the mass of your clean, dry, unknown metal to the nearest 0.001 g.

8. With the power supply still unplugged, attach your unknown metal strip to the anode of the power supply.

9. Immerse your unknown metal in the electrochemistry solution in the beaker. Make sure the metal is not touching the cathode.

Experiment Procedure:

1. Plug in and turn on the power supply. If you do not observe bubbles forming at the copper electrode, you will need to reverse the leads.

2. Let the reaction proceed until you have collected approximately 25 mL of hydrogen gas in the buret. At this time, you will unplug the power supply and

allow time for all of the bubbles to rise in the buret.

3. Record the volume of gas to the nearest 0.01 mL in your lab notebook.

4. Open the stopcock to transfer the remaining electrochemistry solution form the buret to the beaker. Dispose of the contents of the beaker properly and

clean all of the equipment you have used.

5. Clean and completely dry the unknown metal. Record the mass of the unknown metal to the nearest 0.001 g.

Pipet Bulb

Buret

Cathode (-)

Anode (+)

150 mL beaker

0.5 M Acetic Acid and 0.5 M Sodium Sulfate


DATA

Unknown #

Initial mass of metal anode g

Mass of anode after electrolysis g

Initial buret reading mL

Final buret reading mL

Barometric pressure mb (millibar)

Barometric pressure (1000 mb = 1 atm) atm

Temperature

Vapor pressure of H2O at T*

°C

atm

Total volume of H2 produced L

Pressure of H2 (PH2 = Ptot – PH2O) atm

*see the chart below

Because the H2 gas was produced over an aqueous solution, it is necessary to correct for the vapor

pressure of water present in the buret . The pressure of the gases (atmospheric pressure) is the sum

of the pressure of the H2 (PH2) plus the vapor pressure of water (PH2O), which varies with temperature .

These are tabulated below .

Vapor pressure of water at various temperatures

Temperature (°C) PressureH O (atm) 2

Temperature (°C) PressureH2O (atm)

17 0 .0191 24 0 .0295

18 0 .0204 25 0 .0312

19 0 .0217 26 0 .0332

20 0 .0230 27 0 .0351

21 0 .0245 28 0 .0372

22 0.0261 29 0 .0395

23 0.0276 30 0 .0419


Before Leaving the Lab

• All data and calculations should be recorded in Tables 7-1 and 7-2 in your laboratory

notebook.

• Data and calculations from Tables 7-1 and 7-2 should be recorded in your account in

Lon-Capa .

• Check to make sure the PostLab assignment for Lab #7 open. If it does not, check to make

sure all data have been entered into and accepted by Lon-Capa. If there are problems, see

your TA immediately.

• Turn in a copy of your lab report to your TA. This should include your prelab work, all data

(in tables), and any observations and calculations when appropriate.

• Look through the Analysis questions below. These are the questions you will be asked to do

in the PostLab assignment for Lon-Capa (due by 5 p .m . the day before your next lab). Make

sure you know how to do the calculations and answer the questions. If you have time during

lab, it would be a good idea to work on these questions before leaving.

Analysis

The Lon-Capa PostLab 7 assignment will consist of the following analysis . It is due by 5 p .m . the day

before you come to Lab #5. when entering data and doing calculations, be careful with significant figures

.

You will be asked to identify your unknown metal by determining the standard reduction potential and

equivalent mass of the metal. You will also be asked to determine the percent errors in these values

.

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For Part 2, you will be asked for the following information:

Mass of metal oxidized (g)

Volume of H2 gas produced (L)

Pressure of H2 (atm)

Temperature (K)

Number of moles of hydrogen produced

Number of moles electrons transferred

Equivalent mass of unknown metal

Table 7-1 Galvanic Cell

Unknown number of metal

Measured potential of the Cu/M cell

Standard potential of the unknown metal half-reaction

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Table 7-2 Data for Part 2

Unknown number

Initial mass of metal (g)

Final mass of metal (g)

Initial buret reading (mL)

Final buret reading (mL)

Barometric pressure (mb)

Temperature (°C)

Vapor Pressure of water at given temperature

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