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Transcript
Electrochemistry
Unit 12 Pages 1-3
Learning Target:
I can determine the oxidation state for any atom in an element, ion, or compound.
I can explain what happens in an oxidation-reduction, or redox, reaction.
Criteria for Success:
I can determine the oxidation state for an atom in an element.
I can determine the oxidation state for an atom in an ion.
I can determine the oxidation state for an atom in a compound.
I can identify the species being oxidized in a redox reaction.
I can identify the species being reduced in a redox reaction.
Introduction to Oxidation-Reduction Reactions
A. Any chemical process in which electrons are transferred from one atom to another is an _________-__________
reaction.
1. The name for this type of reaction is often shortened to what is called a ________ reaction.
2. A species _____ _________ when _______ (LEO).
A species _____ ________ when _______ (GER).
3. ___________________________ (OIL).
____________________________(RIG).
Oxidation Numbers/States
A. Chemists assign a number to each element in a reaction called an _________________ state that allows him/her to
determine the electron flow in the reaction.
1. Even though they look like them, oxidation states are not _________ charges!
2. Oxidation numbers can be assigned to each atom in an element, ion, or compound…whether the compound is
____________ or ___________________!
3. Oxidation states are imaginary charges assigned based on a set of rules simply used to determine
_________________ flow.
Rules for Assigning Oxidation Numbers (In order of priority):
1. The oxidation number of any pure element is _________.
2. The oxidation number of a monatomic ion is __________ to its charge.
3. The ______ of the oxidation numbers in a compound is zero if ____________, or equal to the ___________ if a
polyatomic ion.
4. The oxidation number of alkali metals in compounds is _____, and that of alkaline earths in compounds is
______. The oxidation number of F is ______ in all its compounds.
5. The oxidation number of H is _____ in most compounds. Exceptions are H2 (where H = 0) and the ionic
hydrides, such as NaH (where H = -1).
6. The oxidation number of oxygen (O) is ______ in most compounds. Exceptions are O2 (where O = 0) and
peroxides, such as H2O2 or Na2O2, where O = -1.
7. For other elements, you can usually use rule (3) to solve for the unknown oxidation number.
Examples:
N2(g) elemental state, so N = 0.
N3-(g) monatomic ion, so N = -3.
NO(g) has O = -2, so N = +2.
NO2(g) has O = -2, so N = +4.
SO42- has O = -2. Thus x + 4(-2) = -2. Solving the equation gives x = -2 + 8 = +6.
K2Cr2O7 has K = +1 and O = -2. Thus 2(+1) + 2 x + 7(-2) = 0; 2 x = 12; Cr = +6.
B. In general, redox reactions will be any type of reaction except ___________________________________.
1. However, there are exceptions so you should always check if electrons have been_____________.
Identifying Oxidation-Reduction Reactions
A. In order to identify a REDOX reaction, write the _________ number for each ________ in the reaction as outlined in the
rules above.
1. A substance that has the element that has been ________ (LOST electrons) will have an oxidation number that
becomes more ________.
2. A substance that has the element that has been ________ (GAINED electrons) will have an oxidation number that
becomes more ________.
B. ___________ are produced in oxidation and acquired in reduction. Therefore, for reduction to occur, oxidation must
occur _____________. Each process is referred to as a ___________.
Oxidation Half-Reaction
Zn(s) → Zn2+(aq) + 2e-
Reduction Half-Reaction
Fe2+(aq)+ 2e- → Fe(s)
Guided/ Independent Practice
Directions: In the following questions, give the oxidation number of the indicated atoms/ion.
1. N in N2O3
12. Fe in Fe(ClO2)3
2. S in H2SO4
13. N in NO3-
3. C
14. Cu2+
4. C in CO
15. Zn2+ 16. C in CH4
5. Na in NaCl
17. Mn in MnO2
6. H in H2O
18. S in SO32-
7. Ba in BaCl2
19. Mg2+
8. N in NO2-
20. Cl-
9. S in Al2S3
21. O2
10. S in HSO4-
22. P4
11. Cl in Fe(ClO2)3
23. Na in Na2S
24. S in H2S
28. Mn in KMnO4
25. Ca2+
29. I in Mg(IO3)2
26. C in CN30. C in C2O42-
27. H in OH
Directions: In each of the following equations, label the oxidation state of each atom before and after the process
occurs. Indicate the element that has been oxidized and the one that has been reduced. Write the oxidation and
reduction half reactions.
1. 2 Na + FeCl2 → 2 NaCl + Fe
2. 2 C2H2 + 5 O2 → 4 CO2 + 2 H2O
3. 2 PbS + 3 O2 → 2 SO2 + 2 PbO
4. 2 H2 + O2 → 2 H2O
5. Cu + HNO3 → CuNO3 + H2
6. AgNO3 + Cu → CuNO3 + Ag