*Oxidation Numbers
*Unit 3 Lecture 2
Oxidation-Reduction Reactions
Aka
Redox Rxns
Chemical
rxns that transfer electrons (e-)
OXIDATION – the LOSS of e-’s
REDUCTION – the GAIN of e-’s
LEO the lion goes GER
Loss of Electrons = Oxidation
Gain of Electrons = Reduction
*Example 1
*A single piece of iron (II) is immersed in a solution of copper
(II) sulfate.
* This is a single replacement reaction…
Fe (s) + CuSO4 (aq) FeSO4 (aq) + Cu (s)
Make a net ionic equation….
Fe (s) + Cu+2 (aq) + SO42- (aq) Fe+2 (aq) + SO42- (aq)
+ Cu (s)
Fe (s) + Cu+2 (aq) Fe+2 (aq) + Cu (s)
Now you can figure out what has been oxidized and reduced…
*Iron is oxidized
*Copper is reduced
*What about this reaction?
*How can you tell what gets oxidized and reduced?
*You must assign oxidation numbers!
*Oxidation # Rules
*Yes, you must memorize these… page
137
Oxidation #’s
• We must know the oxidation # of each element involved to fully understand redox rxns.
• Charges work a lot of the time, but oxidation #’s work EVERY time.
Rules for Assigning Oxidation
Numbers (O#)
1. The oxidation # of an atom of an element is zero.
Example 1 Fe = 0
Example 2 O2 = 0
2. The O# of a monatomic ion is equal to the charge on the ion.
Example 3 Ca+2 = +2
Example 4 H+ = +1
Rules for Assigning Oxidation
Numbers (O#)
3. Fluorine (most EN) is always -1!!!!!
4. The O# of oxygen is always -2 EXCEPT:
in peroxides, HxOy, Oxygen’s O# = -1 when bonded to fluorine Oxygen’s O# = +2
5. The O# of hydrogen is +1 EXCEPT when bonded to:
5. alkali metals (1A Group but not H)
6. alkaline earth metals (2A Group)
7. Al.
Rules for Assigning Oxidation
Numbers (O#)
6. The O# of the more electronegative (EN) atom in a molecule is the same as the charge it would have if it were a monatomic ion.
Hint 1: Molecules = covalent compounds…
(they start with and only have nonmetals) without charges ANYWHERE.
Hint 2: The closer to _F_, the more electronegative ▫ Example 5 PbS S= -2
Rules for Assigning Oxidation
Numbers (O#)
7. Alkali metals, alkaline earth