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Week 12, Lab 1 p1
Week 12, Lab 1: 11/22/05 (Activity Series1)
1) Create, understand, and use an activity series.
2) Examine redox reactions of metals at a molecular level.
3) Identify and explain the purpose of the components of electrochemical cells: anode,
cathode, salt bridge & external circuit.
4) Observe that electrons flow from high potential energy to low potential energy.
Reading: 5.5 p188-192; 19.3: 917-921.
Group Roles: A Technician; B Leader; C Recorder
Redox reactions are some of the most common and most useful chemical reactions. They
produce electrical current which can be harnessed to do work. So far this semester, we have
been focused most intently on the s and p blocks of the periodic table. Since transition metals
play a very important role in redox chemistry we are going shift focus during the exploration of
electrochemistry to include the d-block.
1) Create, understand, and use an activity series and 2) Examine redox reactions of metals
at a molecular level.
Questions: Which metals are easily oxidized? Which are good oxidizing agents?
Data/Fact Gathering:
Fill four wells each in a well plate with each of the 0.1 M salt solutions shown below (You
should fill 16 wells.)
1) Which solution contains the most ions? Which the least? Explain.
2) List characteristics that are the same for all of these solutions and characteristics that make
them different.
1
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/redox/home.html
(accessed May, 2005)
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3) Identify the species that account for the color of each of the four solutions. What evidence did
you use to arrive at your conclusions?
Gather four strips each of magnesium, zinc, copper, silver. Use sand paper or steel wool to clean
any oxidation off the metal strips. Place one piece of each metal into each of the four solutions.
(There are 16 possible metal-solution combinations.) Remove the metal after approximately two
minutes. Organize your observations in a table.
Table 1: Observations of metal/metal ion solutions
4) Did any chemical reactions take place? If so:
a) What changes did you observe in the metal and the solution?
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b) What is being oxidized, what is being reduced?
c) Describe the reactions on a molecular level.
d) Write a chemical equation in the table for each reaction you observed.
5) Using the data in Table 1, rank the metal ions in order of reactivity in Table 2a. Rank the
metals in reverse order of reactivity by filling in Table 2b. What criteria did you use for your
rankings?
Table 2: Ranking a) Metal ion and b) Metal Reactivity
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6) How are Tables 2a and 2b related to each other? Write a balanced half reaction relating each
metal ion to its corresponding metal.
7) Pick one of the metals. Using the reactivity data you collected (Table 1), mark the metal ions
in Table 2a that chemically reacted with it. Repeat this process for each of the metals. Write a
one sentence summary of your findings about the relationships of reacting metals and metal ions
in Tables 2a and 2b.
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Week 12, Lab 1 p5
On-line Simulation
There is an online simulation that models the experiments that you just carried out. Goto the
following website2:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/redox/home.html
Click on “Click on Start”.
Click on “Activity 1”
You will see a simulation of the chemical experiments that you just carried out. You can choose
a strip of metal to place in each of the solutions and see the results.
Click on the molecular scale button at the bottom of the screen to view the metal/metal ion
interactions at the submicroscopic level. Follow the instructions in the software, and take a
closer look at one of systems where you observed a chemical reaction and one where you did
not.
8) Relate your submicroscopic observation with your macroscopic observations for the
examples you have investigated.
You have observed four metals. Is this enough data? Use the remaining activities in the
simulation to gather more data with the goal of answering the questions: Which metals are easily
oxidized? Which metals make good oxidizing agents? The following pages include tables and
questions that will help you organize and evaluate your data:
2
Your computer and browser need to have the MacroMedia Shockwave Player (with the Flash
Player) in order to play the simulations. Your browser should be Netscape Communicator V6 or
higher or Microsoft's Explorer V5.2 or higher. To get the Macormedia plug-in and player go to:
http://www.macromedia.com/downloads/
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Week 12, Lab 1 p6
Table 3: Observations from Activity Two
Solutions
Zn2+
Cu2+
Fe2+
Pb2+
Zn
Cu
Metal
Strips
Fe
Pb
9) In Table 3, write balanced chemical equations for the chemical reactions you observed.
10) Rank this set of metal ions and metals as you did in Tables 2 using Tables 4A and 4B.
Table 4: Activity 2: Ranking a) Metal ion and b) Metal Reactivity
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Table 5: Observations from Activity Three
Solutions
Fe2+
Pb2+
Ni2+
Sn2+
Fe
Pb
Metal
Strips
Ni
Sn
11) Write the reaction equations in Table 5 and order the metals and metal ions using Table 6A
and 6B.
Table 6: Activity 3: Ranking a) Metal ion and b) Metal Reactivity
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12) Use the data you collected in Activities One, Two, and Three to rank the eight metal ions and
the eight metals you have studied by listing them in Tables 7A and 7B.
13) This is called an activity series. Use the activity series to predict whether or not reactions a-c
will take place. If a reaction does take place, fill in the products.
a) Ag(aq)+ + Ni(s) 
b) Ni(aq)2+ + Ag(s) 
c) Ag(aq)+ + Sn(s) 
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Week 12, Lab 1 p9
Now see if or how each of the metals react with HCl. Use the simulation to determine whether
or not a reaction occurs.
Metal
Observations
Ag
Cu
Fe
Mg
Ni
Pb
Sn
Zn
14) Write a balanced chemical reaction in each of the appropriate boxes.
15) Where does H+ fit into Table 7a and where does H2 fit into Table 7b.
Evaluation: Which metals are easily oxidized? Which are good oxidizing agents?
16) In 2-3 sentences answer the questions using evidence from your observations.
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_____________________________________________________________________________
3) Identify and explain the purpose of the components of electrochemical cells: anode,
cathode, salt bridge & external circuit & 4) Observe that electrons flow from high potential
energy to low potential energy.
Redox reactions where metals and metal ions exchange electrons can be made more useful by
driving the electrons that are exchanged through an external circuit. An electrochemical cell
(also know as a voltaic or galvanic cell when the chemical redox reaction is spontaneous)
accomplishes external electron transfer by separating the two redox half reactions. The energy
given off in the reaction can now be captured and used to do electrical work such as in a battery.
The Parts of an Electrochemical Cell
Figure 1a shows a diagram of an electrochemical cell, while Figure 1b shows an actual
electrochemical cell. The electrochemical cell consists of four main parts:
1. The anode: the compartment where oxidation occurs.
2. The cathode: the compartment where reduction occurs.
3. External pathway to allow the flow of electrons. (This can include a light bulb,
voltmeter, motor, etc.)
4. Salt bridge or porous barrier: allows ions to flow back and forth so that charge does
not build up.
Figure 1a) Zn/Zn+2/Cu/Cu+2 electrochemical cell with salt bridge; light is produced
as a result of the electric current flowing through the external circuit. 1b)
Zn/Zn+2/Cu/Cu+2 electrochemical cell with voltmeter measuring potential difference
in the external circuit and a porous barrier.
1a)
Images courtesy of Prentice Hall
1b)
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Each metal electrode is in equilibrium with the metal ions of the solution that it is in and with
electrons. The zinc anode is being oxidized to Zn+2 ions (which go into solution) and two
electrons (which remain at the anode.) Zinc ions in solution are also combining with free
electrons at the anode to form zinc metal on the anode.
Zn(metal)
Zn2+(aq) + 2e-(metal)
Zinc is easily oxidized (the forward reaction). When equilibrium is reached there is a net
increase in the number of free electrons at the anode resulting in a net negative charge.
At the cathode a similar equilibrium takes place:
Cu(metal)
Cu2+(aq) + 2e-(metal)
However copper prefers to be reduced (the reactant favored reaction), which creates a net
decrease in the number of free electrons at the cathode and a net positive charge.
Potential
When the anode and the cathode are connected through a wire, the excess free electrons at the
anode will travel through the wire to offset the depletion of free electrons at the cathode. Similar
charges repel each other, therefore if provided a chance electrons will flow away from other
electrons towards a positive charge. This flow of electrons is called a current.
The difference in charge between the cathode and the anode is potential energy. Physicists talk
about potential energy often describing boulders at the top of a hill. If the boulder rolls over the
top of the hill, it will fall due to gravity so it has a potential energy. The higher the hill, the
greater the boulder’s potential energy. Similarly, potential energy can be due to repulsion
between like charges; the charges would disperse if they had somewhere to go (or someone
would push them over the hill.) The more charges that are gathered together the greater the
potential energy. (Figure 2)
Figure 2: Potential in electrochemical cells
Chemists define this potential energy to be
positive on the anode and negative on the
cathode,3 because the electrons will flow
from the anode to the cathode. A difference
in electric potential energy is called a
potential difference, cell potential, a
potential, or a voltage. A volt is defined in
units of energy per charge, which is:
1 Joule
1 Coulomb
3
= 1 Volt
This is the sign convention for electrochemistry; the sign convention is different for other fields. Typically
potential are defined for positive charges, which would be opposite to the flow of negative charges.
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Week 12, Lab 1 p12
Question: So what is the salt bridge for?
Model for Data Gathering: Explore the answer to this question using the salt bridge model.4
Your group has a model that contains a salt bridge. In the bridge, there are K+ (purple) and NO3(green) ions. Electrons are represented by the black beads.
Before working the model, each compartment should contain the following atoms or ions:
Zinc electrode: 5 Zn atoms
Zinc solution: 2 Zn2+ atoms and 4 NO3- atoms
Salt bridge: 9 K+ atoms and 9 NO3- atoms
Copper solution: 5 Cu2+ atoms and 10 NO3- atoms
Copper electrode: 2 Cu atoms
(Notice how the charge in each compartment is balanced.)
The wire holding beads (electrons) should be connected to the holes at the electrode sections of
the box.
Complete the following activities:
Talk your group technician through the running of the cell. You can kick off the reaction
anywhere...for example:
Where does oxidation take place?
What happens when oxidation takes place?
(If you add two electrons to the wire, two have to come off the other end. Oxidation of one
atom results in a series of chain reactions involving all of the compartments. At the end of
the chain reaction, all of the charges in all of your compartments should be balanced.)
Try a second oxidation or a reduction and carry out the chain of events.
4
From Huddle, P. A.; White, M. D.; Rogers, F. J. Chem. Ed. 2000, 77, 104-110.
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17) Why do you think the wire is higher at one electrode than the other? What do you think this
indicates about the potential of the cell?
18) Can you run the cell backwards? What is necessary to run the cell in this manner? Write out
the chemical reactions.
19) What would happen if all of the KNO3 in the salt bridge was replaced with sugar? Would
the cell still work?
20) What would happen if all the KNO3 in the salt bridge was replaced with NaCl? Would the
cell still work?
21) In one sentence, describe the purpose of the salt bridge?
The model that you were working with helps to demonstrate the flow of electrons and ions in an
electrochemical cell. Electrons (and ions) flow in the same manner in all electrochemical cells.
Electrochemical cells differ only in the elements, ions, and molecules used as the anode and
cathode and in the salt bridge.
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Week 12, Lab 1 p14
Model vs. Reality:
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The Styrofoam balls, depending on color, represent zinc, copper, or sulfate ions. In
nature, individual atoms/ions do not have necessarily these or any color - the model color
is used to identify different atoms/ions.
The beads represent electrons.
When two beads are inserted into the indentations in a Styrofoam ball, the "cation" is
converted into an "atom," i.e., Cu+2 ion + 2 e-  Cu atom.
When two beads are removed from the indentations in a Styrofoam ball, the "atom" is
converted into a "cation."
In nature, while the diameter of zinc and copper atoms/ions is roughly the same (~120
pm), the nitrate ion is considerably larger (~240 pm). Also, Zn+2 and Cu+2 ions are
smaller than Zn and Cu atoms.
No water molecules are shown.
In reality, there are a vast number of ions. For simplicity, only a few are shown.
In this model, the cathode is on the right-hand side. It could be on either side - it makes
no difference to the working of the model.