S338 – Unit 11 Redox

S338 – Unit 11
GER
LEO
Redox
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20-1 Practice Problems
1. What is the oxidation number of each element in: 6. What is the oxidation number of each element in:
a. H 2 SO 4 a. Mg 3 N 2
b. H 2 S 2 O 7 b. FeCl 3
c. SO 3 c. BaSO 4
2. What is the oxidation number of each element in: 7. What is the oxidation number of each element in:
a. H 3 PO 4 a. NaH 2 PO 4
b. P 4 O 6 b. Na 3 PO 4
c. KH 2 PO 4 c. H 3 BO 3
3. What is the oxidation number of each element in: 8. What is the oxidation number of each element in:
a. Ca(H 2 PO 4 ) 2 a. ZnCO 3
b. CaSO 4 b. Ag 2 CO 3
c. KMnO 4 c. BaF 2
4. What is the oxidation number of each element in: 9. What is the oxidation number of each element in:
a. Be(OH) 2 a. Mg(OH) 2
b. B(OH) 3 b. MgCl 2
c. Si(OH) 4 c. CrCl 3
5. What is the oxidation number of each element in: 10. What is the oxidation number of each element in:
a. NH 3 a. Co(OH) 2
b. SCl 2 b. HAuCl 4
c. Sr(OH) 2 c. HgSO 4
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11. Identify the oxidizing agent and the reducing 16. Identify the element that is oxidized and the
agent in the following reaction:
element that is reduced in the following
reaction:
C + 2Cl 2 CCl 4 2Al + 3Br 2 2AlBr 3
12. Identify the element that is oxidized and the 17. Identify the oxidizing agent and the reducing
element that is reduced in the following
agent in the following reaction:
reaction:
H 2 + Cl 2 2HCI Pb + 2HCl PbCl 2 + H 2
13. Identify the oxidizing agent and the reducing 18. Identify the element that is oxidized and the
agent in the following reaction:
element that is reduced in the following
reaction:
2P + 3Cl 2 2PCI 3
SiO 2 + 2C Si + 2CO
14. Identify the element that is oxidized and the 19. Identify the oxidizing agent and the reducing
element that is reduced in the following
agent in the following reaction:
reaction:
C + H 2 O CO + H 2 CO 2 + 2Mg 2MgO + C
15. Identify the oxidizing agent and the reducing 20. Identify the element that is oxidized and the
agent in the following reaction:
element that is reduced in the following
reaction:
2Fe + 3Cl 2 2FeCl 3 H 2 SO 4 + Zn ZnSO 4 + H 2
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Balancing Redox Equations
Balance the equation using half reactions and oxidation numbers.
1. ___ HNO 3 + ____HI ____NO + ____I 2 + _____H 2 O
2. ____HNO 3 + ____I 2 ____HIO 3 + _____NO 2 + ____H 2 O
3. ____S + ____HNO 3 ____SO 2 + ____NO + _____H 2 O
4. ____SO 3 -2 + _____MnO 4 + _____H + ____SO 4 -2 + _____Mn +2 + _____H 2 O
5. ____NO 3 -1 + ____I 2 + ____H + ____IO 3 -1 + ____NO 2 + ____H 2 O
6. ____HCl + ____KMnO 4 + ____H 2 C 2 O 4 ___CO 2 + ____MnCl 2 + ___KCl + __H 2 O
7. ____H + + ____MnO 4 -1 + ____H 2 C 2 O 4 ____CO 2 + ____Mn +2 + ____H 2 O
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Name____________________
Activity Series Lab
Oxidation-Reduction reactions, also known as redox reactions, make up an extremely important area of study in
chemistry. These reactions provide energy in electrochemical cells, and have a variety of commercial uses in
batteries and solar cells. In addition, oxidation-reduction reactions are involved in rusting, cleaning agents,
photography and many other processes.
Redox reactions can be thought of as reactions in which electrons are exchanged between one species that is
oxidized and another that is reduced. Reduction involves the gain of electrons, while oxidation involves the loss
of electrons. A reduction reaction for a metal ion can be written in the following manner:
+n 0
M + ne - M
The oxidation reaction is written as:
0 +n
M M + ne -
+n
M is typically a metal with +n oxidation state, and n represents the number of electrons (e - )
0
gained. M is the metal in its reduced form.
An activity series of metals is a list that puts in order, metals, according to their relative activity. Metals at the
top of the list are oxidized most readily to their metallic ion. They
+n
tend to remain in ionic form (M). Conversely, metals at the bottom of the list are oxidized with difficulty. Their
metallic ion is relatively active so they tend to be in their
0
reduced form (M). In this investigation, you will establish an activity series based on the ease with which
metallic ions are oxidized. The ions to be studied are Cu +2 , Mg +2 , Zn +2 , Pb +2 , and Ag +1 . The metals to be studied
are Cu, Mg, Zn, Pb, and Ag.
Procedure:
1. In your notebook, state your purpose for this lab activity. Will you support the activity series as it is? Or will
you be bold enough to state that you will somehow find flaws in the existing activity series?
2. Put on your goggles (and lab apron if you’d like).
3. Place a paper towel on the microscope platform, then turn on the microscope.
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4. Take a strip of zinc and polish one entire side of it with a piece of steel wool. Then, lay it on the paper towel.
5. Take the Magnesium solution contained in the dropper bottle and place one drop of the solution onto the Zinc
strip. Note in your data table if a reaction occurs (write “Y” if a reactions occurs, write “NR” if no reaction
occurs).
6. Repeat step #5 with the remaining solutions (Mg +2 , Cu +2 , Pb +2 , Ag +1 ) making sure that the drops of solution
do not come in contact together.
7. Rinse the Zinc strip in the sink, dry it and set it aside.
8. Repeat steps 4 through 7, but perform the procedure with the following metal and solution combinations…
Mg strip with Zn, Cu, Pb and Ag solutions
Cu strip with Zn, Mg, Pb and Ag solutions
Pb strip with Zn, Mg, Cu and Ag solutions
Ag strip with Zn, Mg, Cu and Pb solutions
9. Once you are done with all the metal strips and the solutions, clean up your table, wash your hands, then
return the Silver strip to your teacher.
Data Table – Activity Series Lab
Zn
Mg
Cu
Pb
Ag
Zn +2 Mg +2 Cu +2 Pb +2 Ag +1
X
X
X
X
X
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1. Purpose: State your purpose for this lab.
Activity Series Lab – Lab Report Details – what to include
(typed):
2. Data: Present your data in an organized form (re-create your data table with your results written in it. This
can be easily done by using the table function in MS Word).
3. Analysis:
a. Go through step-by-step where each atom should be located in the activity series. Place the atoms in
order from most easily oxidized to least easily oxidized (hint: the most easily oxidized atom is the one
that likes to go into its ionic state. Example: Mg Mg +2 + 2e - Another hint: you need to look at
reactivities to determine the order). This list will be your experimental activity series.
b. Provide THOROUGH evidence for why you placed your atoms in the order you did for part 3a. For
example – let’s say you placed Mg above all the other elements in your activity series. How do you
know this? What experimental evidence do you have to support your answer?
c. Write out a balanced redox equation for each of the reactions that took place. Remember that only
half of the reactions will work, so you should have ten equations written.
4. Conclusion:
a. Restate your purpose.
b. Restate your experimental activity series (the list you formulated from your data).
c. Examine your experimental activity series and compare it to the actual activity series (this was given
to you on a hand-out when we were studying types of reactions, remember? If you can’t find yours, then
you can find it in your chemistry textbook). How do they compare?
d. Did your experiment support your purpose of either confirming or contradicting the existing activity
series? How so?
e. Would you change anything about the procedure to improve your results? If so – what would you
change and why? If not – why wouldn’t you change anything?
(over) →
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5. Post Lab Questions:
a. Would it be a good idea to store a copper metal ion solution (Cu(NO 3 ) 2 (aq), for example) in an
aluminum can? Why or why not?
b. Would it be a good idea to store an aluminum metal ion solution (Al(NO 3 ) 3 (aq), for example) in
a copper can? Why or why not?
5c. Although silverware is beautiful, it tarnishes with time. Tarnish results from a chemical reaction in which
silver reacts with hydrogen sulfide and oxygen in the air to form silver sulfide:
Ag (s) + H 2 S (g) + O 2(g) Ag 2 S (s) + H 2 O (l)
Balance the equation using half-reactions and oxidation numbers. Label the half-reactions as either
oxidation or reduction.
Silver will also react with various sulfides found in rubber, egg yolks, and other substances, including some
vegetables. There are always traces of hydrogen sulfide in the atmosphere. Thus, it is important you protect
silver items by wrapping them in protective cloth or placing them in an environment containing substances that
will absorb the hydrogen sulfide. If silverware does become tarnished, how can it be restored?
• Get a glass baking dish. Line it with a piece of aluminum foil. Place your tarnished silver on top of
the aluminum foil. It must be in contact with the foil for this to work!
• Boil some water on a stove. Once in comes to a rolling boil, add about 1 cup of baking soda for every
gallon of water. Be careful, though! The water tends to boil over quickly so add the baking soda slowly!
• Pour the hot water/baking soda solution into the pan containing the aluminum foil and tarnished silver.
Moderately tarnished pieces will be clean within seconds, while badly tarnished pieces might take over an hour
to work.
How does it work?
The silverware is cleaned by an redox reaction in which silver sulfide produces elemental silver according to
the following equation:
Al (s) + Ag 2 S (s) + H 2 O (l) Ag (s) + Al(OH) 3(aq) + H 2 S (g)
Balance the equation using half-reactions and oxidation numbers. Label the half-reactions as either
oxidation or reduction.
In the electrochemical process, the silver, surrounded by the electrolyte (NaHCO 3 , baking soda), forms
one plate of an electric cell and the aluminum forms the other. The equation suggests that hydrogen sulfide gas
(the gas produced in rotting eggs) is evolved in small quantities. If this reaction occurs, then you may be able to
detect the odor of rotten eggs by removing part of the aluminum from a piece of silver as the reaction is
proceeding.
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Name ______________________________________
Voltaic Cells Lab
Data and Analysis
A. For each cell below, label the following on the diagram:
1 & 2. The identity of the metal
3. The reading on the voltmeter
4. The direction of the electron flow through the wires of the voltmeter
5. Label the metal that is the cathode and the metal that is the anode.
6. Label the metal that is oxidizing and the metal that is being reduced
7. Draw the direction in which each ion in the salt bridge moves.
8. Calculate the electrical potential of the cell (show work)
9. Write the half reactions and a balanced redox reaction for the cell.
Zn/Zn +2 – Cu/Cu +2 Cell
1
.
2
.
.
3
.
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Zn/Zn +2 – Pb/Pb +2 Cell
1
.
2
.
3
.
Zn/Zn +2 – Mg/Mg +2 Cell
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1
.
2
.
Pb/Pb +2 – Mg/Mg +2 Cell
3
.
1
.
2
.
3
.
Mg/Mg +2 – Cu/Cu +2 Cell
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1
.
2
.
3
.
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Voltaic Cells Worksheet
On each of the following diagrams, label: 1. the anode, 2. the cathode, 3. Direction of electron
flow, 4. Direction of anion and cation flow through the salt bridge, and 5. Which electrode
should be increasing in mass. Then write out the half-reactions, the overall reaction and
calculate the net voltage of the cell.
1)
A = Cr B = Cr +3
E = Fe D = Fe +2
2)
A = Sn B = Sn +2
E = Al D = Al +3
3)
A = Ag B = Ag +1
E = Cu D = Cu +2
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