chemistry pre-lab

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experiment_10___redox_reactions___recovering_ag_s__from_agcl_s_.pdf

experiment_11___redox_reactions___voltaic_cells.pdf

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CHEM 404
Exp 10
Experiment 10:
REDOX REACTIONS – RECOVERING Ag(s) FROM AgCl(s)
Many chemical reactions can be classified as oxidation-reduction, or redox, reactions,
because they involve the oxidation of one species and the reduction of another. Oxidation is the
loss of electron density while reduction is the gain of electron density and the two must occur
together. A redox reaction consists of the transfer of electron density from one species (the one
being oxidized) to another (the one being reduced). Depending on the reaction, one, two, or more
electrons can be transferred between species.
Redox reactions consist of two parts, the reduction half-reaction and the oxidation halfreaction. These half-reactions combine to give the overall reaction. Consider the following
example of two redox half-reactions:
Pb(s) ? Pb2+(aq) + 2e-
an oxidation half-reaction (loss of e-)
Fe3+(aq) + e- ? Fe2+(aq)
a reduction half-reaction (gain of e-)
When combining two half-reactions to give the overall reaction, the number of electrons
transferred must be balanced. Since Pb(s) loses two electrons in the oxidation half-reaction, two
Fe3+ ions must be present to balance the equation. The entire reduction half-reaction is doubled,
then the two half-reactions are combined, and the electrons are cancelled out, giving the overall
redox reaction:
Pb(s) + 2Fe3+(aq) ? Pb2+(aq) + 2Fe2+(aq)
In this experiment, you will be performing a redox reaction involving silver. During an
experiment in CHEM 403, a solution of silver nitrate was used to titrate an unknown sample
containing chloride ions. The silver nitrate mixed with the chloride ions to form solid silver
chloride.
AgNO3(aq) + Cl-(aq) ? AgCl(s) + NO3-(aq)
The silver nitrate used in the experiment is expensive (100 g costs $396.86 through Fisher
Scientific in July 2017), but can be recycled. The first step in this process is to recover the silver
chloride from the experiment and convert it to elemental silver. This can be achieved through a
redox reaction using zinc metal:
2AgCl(s) + Zn(s) ? 2Ag(s) + ZnCl2(aq)
After collecting the solid silver metal from the lab, nitric acid can be added to reform silver
nitrate, which can be used in the next year’s CHEM 403 experiment. Since this second reaction
involves use of a strong acid and the evolution of nitrogen dioxide gas, this reaction must be
performed in a fume hood; this reaction will not be performed in the CHEM 404 lab.
Ag(s) + 2HNO3(aq) ? AgNO3(aq) + NO2(g) + H2O(l)
*Adapted from Slowinski, E. J., Wolsey, W. C. Chemical Principles in the Laboratory 9th ed.
CHEM 404
Exp 10
Procedure
Weigh approximately 0.3-0.4 grams of AgCl(s) and record the mass to 0.01 g. Weigh
approximately 0.15 grams of zinc and add both solids to a 50-mL beaker. Add 5 mL of distilled
water to the beaker, plus five drops of 6 M HCl. The mixture should begin to bubble. Allow the
reaction to continue, mixing the solids occasionally with a glass stirring rod, until the bubbling
ceases.
After the reaction has ceased, most of the AgCl(s) will be reduced to Ag(s). Any excess
zinc can be removed by adding another 5-10 drops of 6 M HCl to the mixture, which will cause
the evolution of H2 gas and complete the oxidation of any remaining zinc.
Transfer the contents of the beaker to a Büchner funnel while gentle suction is being
applied. Use distilled water to aid transfer of the Ag(s) from the beaker to the filter paper. Continue
to dry the solid product while applying suction.
Once the Ag(s) is dry, turn off the suction, and transfer the solid to a preweighed watch
glass. Dry the solid under a heat lamp for 5 minutes. Allow the sample to cool to room temperature,
then record the mass of the Ag(s) with the watch glass.
CHEM 404 – Exp 10
Name:______________________________ Section: _________
Data & Calculations
Mass of AgCl(s)
____________ g
Mass of Zn
____________ g
Moles AgCl(s) in sample
____________ moles
Calculated moles of Ag(s) that can be produced
____________ moles
Calculated mass of Ag(s) that can be produced
(Theoretical Yield)
____________ g
Mass of watch glass
____________ g
Mass of watch glass and Ag(s)
____________ g
Mass of Ag(s) recovered
____________ g
Percent yield of Ag(s)
(mass recovered/theoretical yield)*100%
____________ %
During a typical fall semester, the experiment in CHEM 403 uses approximately 35 grams of
AgNO3. Based on this reaction’s percent yield (assuming the reaction that initially made AgCl and
the one to reform AgNO3 go to completion), how many grams of AgNO3 can be recycled?
The next reaction with nitric acid is also a redox reaction. Is the silver being oxidized or reduced
in the reaction, Ag(s) + 2HNO3(aq) ? AgNO3(aq) + NO2(g) + H2O(l)?
CHEM 404 – Exp 10
Name:______________________________ Section: _________
Prelab Assignment
1. In the reaction of AgCl(s) with zinc, one species is oxidized and the other is reduced, and two
half-reactions can be written. Omitting the chloride ions from the equation can help show the
metal species and ions involved:
2Ag+ + Zn ? 2Ag + Zn2+
a. Which reactant is oxidized (loses e-) in this reaction?
b. Which reactant is reduced (gains e-) in this reaction?
c. Write the two half-reactions that occur in this redox reaction. Be sure to balance each half
reaction with the appropriate number of electrons.
2. How many grams of Ag(s) can be recovered from a 0.362 g sample of AgCl(s)? Show all
calculations.
CHEM 404
Exp 11
Experiment 11:
REDOX REACTIONS – VOLTAIC CELLS
Recall from the previous experiment that redox reactions consist of two parts, the reduction
half-reaction and the oxidation half-reaction. These half-reactions combine to give an overall
reaction. The tendency for a redox reaction to occur can be measured if the two reactions are made
to occur in separate regions connected by a porous barrier. This apparatus is called a voltaic cell.
Figure 1: Diagram of a Voltaic Cell
B
E
–
A
D
+
C
F
In a voltaic cell, two electrodes, the anode (A) and the cathode (C), are connected by a
voltmeter (B). Oxidation occurs at the anode, and electrons flow from the anode through the
voltmeter to the cathode, where reduction occurs. The anode is considered the negative electrode,
as electrons are emitted from that side of the cell; the cathode is considered the positive electrode.
The two solutions (D and F) contain the associated ions for the oxidation and reduction
half-reactions. As electron density builds up on the cathode side, a salt bridge (E) allows the flow
of ions to balance the charge. For example, as a negative charge from electrons builds in solution
F, positively-charged sodium ions (Na+) can flow from solution D to solution F to balance the
charge. The entire voltaic cell creates a circuit.
The voltaic cell you will be using is the lab is pictured in below. One solution is nested
inside the other. The ceramic cup acts as a salt bridge to allow ions to travel between solutions.
Figure 2: A Voltaic Cell Apparatus
B
A
E
C
D
F
*Adapted from Slowinski, E. J., Wolsey, W. C. Chemical Principles in the Laboratory 9th ed.
CHEM 404
Exp 11
The voltmeter between the two electrodes measures the voltage, or the potential, E,
between them. The magnitude of the potential is a direct measure of the driving force or
thermodynamic tendency of the spontaneous oxidation-reduction reaction to occur.
The voltage produced by a voltaic cell is the sum of the potential for the oxidation reaction
and the potential for the reduction reaction. Note, importantly, that half-reaction potentials are not
multiplied when balancing an overall reaction. This differs from Hess’s law with ?H values, for
example.
Ecell = Eoxidation + Ereduction
(1)
Since any cell potential is the sum of two electrode (half-reaction) potentials, it is not
possible, by measuring cell potentials, to determine individual absolute electrode potentials for
each half reaction. However, if a value of potential is arbitrarily assigned to one electrode reaction,
other electrode potentials can be given definite values, based on the assigned value. This has
historically been done with the half-reaction:
2H+(aq) + 2e- ? H2(g)
E° = 0.00 V
The symbol E° above refers to the standard electrode potential, which is measured at 1 atm
of pressure at 298 K with all solutions in 1.0 M concentrations.
If the potential for the following reaction between lead and hydrogen ions (protons) is
measured, the standard electrode potential for lead can be determined relative to the hydrogen halfreaction.
Pb(s) + 2H+(aq) ? Pb2+(aq) + 2H2
E° = 0.13 V
Using Equation 1, the standard electrode potential for the oxidation of solid lead to Pb2+ ions can
be calculated as 0.13 V.
Ecell = Eoxidation + Ereduction; so 0.13 V = Eoxidation + 0.00 V
Through this method, the standard electrode potentials of many half-reactions can be
recorded for reference. By convention, most tables of standard electrode potentials list all halfreactions as reductions. If the potential for an oxidation half-reaction is known, the potential for
the reverse reduction half-reaction is obtained by taking the opposite sign of the voltage. For
example, for the lead electrode:
Pb(s) ? Pb2+(aq) + 2e-
E°oxidation = 0.13 V
Pb2+(aq) + 2e- ? Pb(s)
E°reduction = –0.13 V
In this experiment, you will measure the voltages of several different cells. By arbitrarily
assigning the potential of a particular half-reaction to be 0.00 V, you will be able to calculate the
potentials corresponding to all of the various half-reactions that occur in your cells.
CHEM 404
Exp 11
Procedure
Note that the ceramic cup used in your voltaic cell is expensive and very fragile. Keep it
close to the bench-top, and take care when rinsing it between measurements. Do not hold it away
from the bench or over the floor.
The purpose of this experiment is to measure voltaic cell potentials to allow you to
determine the half-reaction potentials for five of the six electrodes listed below, by comparing
them with the Fe3+(aq), Fe2+(aq), C(graphite) electrode, which will be arbitrarily assigned a
potential of 0.00 V. After measuring each half-reaction in reference to this electrode, you will be
given the measured potential of the Fe3+, Fe2+, C electrode in reference to the H+, H2 electrode in
order to convert your measurements and compare them to a list of standard electrode potentials.
You will be working with the following six half-reactions. Note that each half-reaction is
expressed here as a reduction process.
Cu2+(aq), Cu(s)
Al3+(aq), Al(s)
Pb2+(aq), Pb(s)
Ni2+(aq), Ni(s)
Zn2+(aq), Zn(s)
Fe3+(aq), Fe2+(aq), C(graphite)
The last half reaction listed is your reference half-reaction, and consists of iron in two oxidation
states. Since there is no solid metal in this half-reaction to act as an electrode, an inert electrode is
used to connect between the voltmeter and the solution in which the reaction takes place. In this
case, it is a graphite electrode.
Using the apparatus shown in Figure 2, set up a voltaic cell with the reference electrode,
Fe (aq), Fe2+(aq), C(graphite), in the outer plastic container. Add 140 mL of a 0.1 M Fe3+/0.1 M
Fe2+ solution to the plastic container. Carefully clamp the round graphite rod to one of the clamps.
Attach the positive pole (red wire) of the voltmeter to the clamp. Turn the voltmeter to “volt” and
set it to read “0.00 V”. You are now ready to measure the remaining half-reactions relative to this
reference half-reaction. Leave this half-cell intact throughout the experiment.
3+
To perform a measurement, obtain 25 mL of one of the half-reaction’s solution (for example,
the 0.1 M Cu2+ solution) in a 50-mL beaker and add it to the ceramic cup. Carefully place the cup
into the reference electrode’s solution. Find the appropriate metal strip (in this case, the Cu strip)
and position it with the thumb screw and clamp so it is partially submerged in its solution. Attach
the negative pole (black wire) to the new electrode, wait approximately 30 seconds, then read and
record the voltage (noting if the value is positive or negative) as the “Cell Potential” in Table 1 of
your Data & Calculations.
To proceed to test the next half-reaction, remove the metal strip from the previous
measurement, rinse it with distilled water and dry it with a paper towel. Carefully remove the
ceramic cup, and rinse the outside of the ceramic cup into your waste beaker with distilled water.
Pour the electrode solution back into the 50-mL beaker. Rinse the inside of the cup with distilled
water into your waste beaker. Return the electrode solution to the correct stock bottle, rinse the
beaker, obtain 25 mL of the next half-reaction’s solution, and proceed with the rest of the setup
process for the new half-reaction.
CHEM 404 – Exp 11
Name:______________________________ Section: _________
Data & Calculations
Table 1: Measured Cell Potentials with Fe3+, Fe2+, C Reference Electrode
Voltaic Cell Half-reaction measured Cell Potential (V)
1
Cu2+(aq), Cu(s)
____________
2
Al3+(aq), Al(s)
____________
3
Pb2+(aq), Pb(s)
____________
4
Ni2+(aq), Ni(s)
____________
5
Zn2+(aq), Zn(s)
____________
For each of the voltaic cells, complete Table 2 with the two half-reactions that occur in the cell. If
the measured cell potential is a positive value, the reduction half-reaction is Fe3+ + e- ? Fe2+; if
the cell potential is a negative value, the oxidation half-reaction is Fe2+ ? Fe3+ + e-. Note that the
Fe3+, Fe2+ half-reaction potential will be 0.00 V whether it occurs as an oxidation or a reduction.
Table 2: Half-reactions Occurring in the Voltaic Cell
Voltaic Cell
Half-reactions
Half-reaction Potential (V)
1
Oxidation: _________________________
__________
Reduction: _________________________
__________
Oxidation: _________________________
__________
Reduction: _________________________
__________
Oxidation: _________________________
__________
Reduction: _________________________
__________
Oxidation: _________________________
__________
Reduction: _________________________
__________
Oxidation: _________________________
__________
Reduction: _________________________
__________
2
3
4
5
On the following page, you will convert your half-reaction potentials from the Fe3+, Fe2+, C
reference electrode, to the more commonly used H+, H2 reference electrode. Since the
concentrations you will be using are 0.1 M instead of the standard 1.0 M and your solutions will
not necessarily be at 298 K, do not expect your data to reproduce the voltages listed in a table of
standard electrode potentials. Some half-reactions will be more affected than others.
Table 3: Converting Potentials to H+, H2 Reference Electrode
Voltaic
Cell
Half-reactions
(written as reductions)
Half-reaction Potential
Half-reaction Potential
(Fe3+, Fe2+ reference)
(H+, H2 reference)
1
_________________________
__________
__________
Fe3+ + e- ? Fe2+
__ 0.00 V __
_ +0.77 V _
_________________________
__________
__________
Fe3+ + e- ? Fe2+
__ 0.00 V __
_ +0.77 V _
_________________________
__________
__________
Fe3+ + e- ? Fe2+
__ 0.00 V __
_ +0.77 V _
_________________________
__________
__________
Fe3+ + e- ? Fe2+
__ 0.00 V __
_ +0.77 V _
_________________________
__________
__________
Fe3+ + e- ? Fe2+
__ 0.00 V __
_ +0.77 V _
2
3
4
5
Note to change an oxidation half-reaction potential to a reduction half-reaction potential,
simply switch the sign of the potential. To convert from one reference to the other, add 0.77 V to
each of the half-reaction potentials in the Fe3+, Fe2+ column and enter these values in the H+, H2
reference column.
Once you have calculated the half-reaction potentials in reference to the H+, H2 electrode,
compare your values to those listed in a table of standard electrode potentials available from your
instructor. How well do your measured values match those in the table?
CHEM 404 – Exp 11
Name:______________________________ Section: _________
Prelab Assignment
A student measures the potential of a cell prepared with l M CuSO4 in one solution and l M AgNO3
in the other. There is a Cu electrode in the CuSO4 solution and an Ag electrode in the AgNO3
solution, and the cell is set up as in Figure 2. She finds that the potential, Ecell, or voltage, of the
cell, is 0.45 V, and that the Cu electrode is negative.
1. At which electrode is oxidation occurring?
2. Write the balanced equation for the oxidation reaction.
3. Write the balanced equation for the reduction reaction.
4. Combine the two half-reactions, balancing the number of electrons involved, to give the overall
redox reaction occurring in this cell.
5. In this cell, the Ag+, Ag electrode acts as the reference electrode. Its potential is initially taken
to be 0.00 V.
a. What is the value of the potential for the Cu2+, Cu half-reaction in reference to the Ag+, Ag
electrode? (Ecell = Eoxidation + Ereduction)
____________ V
b. If the potential of the Ag+, Ag electrode is taken to be 0.80 V in reference to the H+, H2
electrode, what is the value of the potential for the Cu2+, Cu half-reaction in this cell?
(Ecell = Eoxidation + Ereduction)
____________ V

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