1. Those Pesky Electrons are Doing it again!
Redox
Those Pesky Electrons are
Doing it again!

2. Oxidation and Reduction
Oxidation: the loss of electrons or the gain of oxygen
Reduction: the gain of electrons or the loss of oxygen
Early chemists saw oxidation as the process of combining an element with oxygen…this has been expanded since then.

3. Remember the Lion! Leo the lion goes ger!
Loss of Electrons is Oxidation
Gain of Electrons is Reduction
The substance that loses electrons is called the reducing agent (it is oxidized).
The substance that gains electrons is called the oxidizing agent (it is reduced).

4. Oxidation: options Complete loss of electrons (ionic rxn)
Shift of electrons away from an atom in a covalent bond
Gain of oxygen
Loss of hydrogen by a covalent cpd
Increase in oxidation number

5. Reduction: options Complete gain of electrons (ionic rxn)
Shift of electrons toward an atom in a covalent bond
Loss of oxygen
Gain of hydrogen by a covalent cpd
Decrease in oxidation number

6. Practical Application
Corrosion
Coating: preventing the easily oxidized metal from exposure to oxidizing agents
Painting iron/steel
Anodized coatings on aluminum
Galvanizing (zinc plate)
Sacrificial anodes: letting a more easily oxidized metal; Zn & Mg common

7. Oxidation Numbers: rules
The oxidation number of a monatomic ion is equal in magnitude and sign to its ionic charge.
The oxidation number of hydrogen is +1, unless a hydride where -1.
The oxidation number of oxygen is -2, except for a peroxide, where -1, or in compounds with Fluorine, where it is positive.
The oxidation number of an atom in uncombined form is 0.
For any neutral compound, the sum of the oxidation numbers must be 0.
For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.

8. Reaction Types Obviously, we have redox and non-redox reactions.
All reactions are one or the other.
Most double-displacement and acid/base reactions are not redox.
Many single-displacement, synthesis, decomposition, and combustion reactions are redox.

9. The math of redox
Simply enough, the total increase in oxidation number of those species oxidized must be balanced by the decrease in oxidation number of those species reduced.
Once again, everything must be balanced!

10. Balancing Redox reactions
Easiest way: set up equations into half-reactions, one each for oxidation and reduction.
This allows for easier determination of each part.

11. Method for balancing Redox
Step 1: write the unbalanced equation in ionic form
Step 2: write separate half-reactions for ox. & red. Processes
Step 3: balance the atoms in the half-reactions
Step 4: add enough electrons to one side of each half-reaction to balance the charges (you will add these to the opposite side of the reaction later)
Step 5: multiply each half-reaction by an appropriate number to make the numbers of electrons equal in both (think coefficient)
Step 6: add the half-reactions to show an overall equation
Step 7: add the spectator ions and balance the equation