1. Electrochemical Reactions
In electrochemical reactions, electrons are transferred from one species to another.
Mg (s) + HCl (aq) →

2. Oxidation Numbers
In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

3. Oxidation and Reduction
A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.

4. Oxidation and Reduction
A species is reduced when it gains electrons.
Here, each of the H+ gains an electron, and they combine to form H2.

5. Oxidation and Reduction
What is reduced is the oxidizing agent.
H+ oxidizes Zn by taking electrons from it.
What is oxidized is the reducing agent.
Zn reduces H+ by giving it electrons.

6. In a spontaneous Oxidation-Reduction (redox) Reaction electrons are transferred and energy is released
as the reaction proceeds:
Zn strip dissolves…
The blue color due to Cu2+ fades.
Copper metal is deposited.

7. Oxidation and Reduction Half-Reactions
A reaction represented by two half-reactions.
Oxidation:
Zn(s) → Zn2+(aq) + 2 e-
Reduction:
Cu2+(aq) + 2 e- → Cu(s)
Overall:
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
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General Chemistry: Chapter 5

8. Assigning Oxidation Numbers, ON
Elements in their elemental form have an oxidation number of 0.
Mg Cl2 Pb
The oxidation number of a monatomic ion is the same as its charge.
Zn2+ F- Na+

9. HF NaH Nonmetals tend to have negative oxidation numbers
In combination with other elements:
Nonmetals tend to have negative oxidation numbers
Oxygen has an oxidation number of −2, except in the peroxide ion, which has an oxidation number of −1.
MgO Na2O
Hydrogen is −1 when bonded to a metal and +1 when bonded to a nonmetal.
HF NaH

10. Assigning Oxidation Numbers
Fluorine always has an oxidation number of −1.
HF NaF
The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.
ZnBr CuCl2

11. Assigning Oxidation Numbers
The sum of the oxidation numbers in a neutral compound is 0.
HCl H2O K2O
5. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.
CO ClO4-

12. Examples
Determine the oxidation number of sulfur in the following compounds:
H2S
SO42- S8
S2O32-

13. Electrochemical Reactions

14. Corrosion and…

15. Figure: 20-20
Title:
Protection of iron against electrolytic corrosion
Caption:
In the anodic reaction, the metal that is more easily oxidized loses electrons to produce metal ions: in (a), this is iron; in (b), it is zinc. In the cathodic reaction, oxygen gas, which is dissolved in a thin film of water on the metal, is reduced to OH-. Rusting of iron occurs in (a), but not in (b). When iron corrodes, Fe2+ and OH- ions from the half-reactions initiate these further reactions:
Fe2+ + 2OH--> Fe(OH)2(s)
4Fe(OH)2(s) + O2 + 2H2O --> 4Fe(OH)3(s)
2Fe(OH)3(s) --> Fe2O3 • H2O(s) + 2H2O(l)

16. Figure: UN
Title:
Magnesium sacrificial anodes
Caption:
The small cylindrical bars of magnesium attached to the steel ship provide cathodic protection against corrosion. Galvanization is another type of cathodic protection, the use of magnesium here is more cost effective.

17. Oxidation and Reduction
O.S. of some element increases in the reaction.
Electrons are on the right of the equation
Reduction
O.S. of some element decreases in the reaction.
Electrons are on the left of the equation.
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General Chemistry: Chapter 5

18. Oxidation State Changes
Assign oxidation states:
Fe2O3(s) + 3 CO(g) → 2 Fe(l) + 3 CO2(g) D
Fe3+ is to metallic iron.
CO(g) is to carbon dioxide.
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General Chemistry: Chapter 5

19. Balancing Oxidation-Reduction Equations
Few can be balanced by inspection.
Systematic approach required.
The Half-Reaction (Ion-Electron) Method
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20. Balancing in Acid Write the equations for the half-reactions.
Balance all atoms except H and O.
Balance oxygen using H2O.
Balance hydrogen using H+.
Balance charge using e-.
Equalize the number of electrons.
Add the half reactions.
Check the balance.

21. SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq)
EXAMPLE 5-6
Balancing the Equation for a Redox Reaction in Acidic Solution. The reaction described below is used to determine the sulfite ion concentration present in wastewater from a papermaking plant. Write the balanced equation for this reaction in acidic solution..
SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq)
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22. EXAMPLE 5-6 Determine the oxidation states: 4+ 6+ 7+ 2+
SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq)
Write the half-reactions:
SO32-(aq) → SO42-(aq) + 2 e-(aq)
5 e-(aq) +MnO4-(aq) → Mn2+(aq)
Balance atoms other than H and O:
Already balanced for elements.
General Chemistry: Chapter 5
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23. EXAMPLE 5-6 Balance O by adding H2O:
H2O(l) + SO32-(aq) → SO42-(aq) + 2 e-(aq)
5 e-(aq) +MnO4-(aq) → Mn2+(aq) + 4 H2O(l)
Balance hydrogen by adding H+:
H2O(l) + SO32-(aq) → SO42-(aq) + 2 e-(aq) + 2 H+(aq)
8 H+(aq) + 5 e-(aq) +MnO4-(aq) → Mn2+(aq) + 4 H2O(l)
Check that the charges are balanced: Add e- if necessary.
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24. EXAMPLE 5-6 Multiply the half-reactions to balance all e-:
5 H2O(l) + 5 SO32-(aq) → 5 SO42-(aq) + 10 e-(aq) + 10 H+(aq)
16 H+(aq) + 10 e-(aq) + 2 MnO4-(aq) → 2 Mn2+(aq) + 8 H2O(l) 6 3
Add both equations and simplify:
5 SO32-(aq) + 2 MnO4-(aq) + 6H+(aq) →
5 SO42-(aq) + 2 Mn2+(aq) + 3 H2O(l)
Check the balance!
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