1. Unit 7: Redox & Electrochemistry

2. What’s the point ? REDOX reactions are important in …
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What’s the point ?
REDOX reactions are important in …
C3H8O + CrO3 + H2SO4  Cr2(SO4)3 + C3H6O + H2O
Purifying metals (e.g. Al, Na, Li)
Producing gases (e.g. Cl2, O2, H2)
Electroplating metals
Electrical production (batteries, fuel cells)
Protecting metals from corrosion
Balancing complex chemical equations
Sensors and machines (e.g. pH meter)

3. What is Redox? REDOX stands for REDuction/OXidation
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REDOX stands for REDuction/OXidation
Oxidation is often thought of as a combination of a substance with oxygen (rusting, burning)
Just like with acid/base definitions the definition of oxidation is expanded
Oxidation refers to a loss of electrons
Reduction refers to a gain of electrons
As a mnemonic remember LEO says GER
Loss Electrons = Oxidation
Gain Electrons = Reduction
Read 12.1 (pg )

4. Testing concepts
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Q- What is oxidation? What is reduction? Represent each as a chemical equation.
A- oxidation = loss of e– … X X+ + e–
reduction = gain of e– … X + e– X–
Q- Why are 2Na + Cl2  2NaCl & 2H2 + O2  H2O considered redox reactions?
A- Both involve the transfer of electron density (Na has no charge, the atoms in diatomic molecules have no partial charge. After reaction the atoms have different shares of the electrons because of different EN values)
Q- Is it possible to oxidize a material without reducing something else?
A- No. A lost e– is taken up by something else.

5. Testing concepts Q- Define oxidizing agent, reducing agent.
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Q- Define oxidizing agent, reducing agent.
A- An oxidizing agent causes oxidation by being reduced itself
A reducing agent causes reduction by being oxidized itself
Q- PE 1
A- CaCl2 is an ionic compound with a positive calcium ion and negative chlorine ions
Ca + Cl2  CaCl2
Ca  Ca2+ + 2e–, Cl2 + 2e– 2Cl–.
Thus Ca is losing electrons (oxidation) and Cl is gaining electrons (reduction).
These are called “half reactions”

6. Oxidation numbers
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We will see that there is a simple way to keep track of oxidation and reduction
This is done via “oxidation numbers”
An oxidation number is the charge an atom would have if electrons in its bonds belonged completely to the more electronegative atom
E.g. in HCl, Cl has a higher EN (pg. 255). Thus, oxidation numbers are Cl = -1, H = +1
Notice that oxidation numbers are written as +1 vs. 1+ to distinguish them from charges.
Instead of referring to EN chart, a few rules are followed to assign oxidation numbers
Refer also to study note

7. Rules (and rationale - 12.2)
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1. Any element, when not combined with atoms of a different element, has an oxidation # of zero. (O in O2 is zero)
2. Any simple monatomic ion (one-atom ion) has an oxidation number equal to its charge (Na+ is +1, O2– is –2)
3. The sum of the oxidation numbers of all of the atoms in a formula must equal the charge written for the formula. (if the oxidation number of O is –2, then in CO32– the oxidation number of C is +4)

8. Rules (and rationale - 12.2)
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4. In compounds, the oxidation # of IA metals is +1, IIA is +2, and aluminum (in IIIA) is +3
5. In ionic compounds, the oxidation # of a nonmetal or polyatomic ion is equal to the charge of its associated ion. (CuCl2, Cl is –1)
6. F is always –1, O is always –2 (unless combined with F), H is usually +1
rule
total
Ox.#
or rule 5 6 3 6 4 3 6 3 6 6 3 5 6 3 6 +1 +5 -6 +2
+12
-14 -4 +6 -2 +1 -1 +2 +5 -8 +1 +5 -2 +1 +6 -2 -2 +1 -2 +1 -1 +1 +5 -2
H N O 3
K2Cr2O7
C2H6O
AgI
H2PO4–
PE 2 (450), 12.9, (484), 12.10, 12.11, (484)

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More practice
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PE 2 pg. 450: answers in back of book
12.9, (pg. 484): answers in back of book
12.10: a) Cl=+7, b) Cr=+3, c) Sn=+4, d) Au=+3
12.11: a) Na= +1, H = +1, P= +5, O= -2
b) Ba= +2, Mn= +6, O= -2
c) Na= +1, S= +2.5, O= -2
d) Cl= +3, F= -1
12.13: a) +2, b) +4, c) +3, d) +5, e) -2, f) 0,
g) –1, h) -3 I) –1/3
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