Unit 6 – Redox Reactions and Electrochemistry

Redox reactions involve a transfer of electron density from one atom to another. Two common examples include the formation of salt and the formation of water:  2 Na + Cl 2  2 NaCl & 2H 2 + O 2  2H 2 O The transfer of electron density in the formation of salt is more obvious since we form two ions, however the transfer of electron density in making water is more difficult to see. Redox reactions always contain 2 parts that make this process more obvious and shows the transfer of electron density.

Redox - a reaction that is half reduction and half oxidation  Reduction – a gain of electron density (reactant to product) Oxidation – a loss of electron density (reactant to product) A way to remember : LEO the lion says GER Loss of Electron density is Oxidation, Gain of Electron density is Reduction GER

If we break the first reaction into its half reactions, oxidation and reduction, we see the following: Na  Na+ + e- (oxidation) and Cl 2 + 2e  2Cl (reduction)  Thus sodium is being oxidized and chlorine is being reduced.

No substance is ever oxidized unless another is reduced. The substance that accepts electrons is called the oxidizing agent (it is reduced).  The substance that gains electrons is called the reducing agent (it is oxidized).

For the reaction making salt, the reducing agent is sodium (recall that it is oxidized) and the oxidizing agent is chlorine (recall that it is reduced).

 Examples of redox reactions in everyday life: any battery - this creates a flow of electrons between the reducing agent (the one that is being oxidized or losing electrons) and the oxidizing agent (the that is being reduced or gaining electrons).

Metabolism - these use oxygen to convert carbohydrates and fats into CO 2 and H 2 O Bleach – oxidizes stains so that they are easier to remove by detergents

Oxidation Numbers

 In redox reactions there is always a transfer of electron density. This change can be monitored by observing the oxidation numbers of the atoms and ions involved. Oxidation numbers is the charge that an atom would have if the electrons in each bond belonged entirely to the more electronegative element.

There are several rules that make the assignment of oxidation numbers easier for more complex situations that just HCl. They are always used in order and are as follows: 1. Any element that is not combined with others to form a molecule is given the value 0. (ex. Na , H 2 , Br 2 )

2. Any simple monatomic ion has an oxidation number equal to its charge.

(ex. Cl- has an oxidation number of –1, Mg 2+ has an oxidation number of +2) 3.

 The sum of oxidation numbers for the all the atoms in a formula must equal the overall charge of the formula.

(ex. CaCl 2 – Ca = +2, Cl = -1, overall +2 –1-1 = 0 ) 4. In compounds the oxidation number of any group I metal is ALWAYS +1 5. In compounds the oxidation number of any group II metal is ALWAYS +2 6. In compounds the oxidation number of Aluminum is ALWAYS +3. 7. In binary compounds with metals, the oxidation number of a nonmetal is equal to the charge it would have as a monatomic ion (ex. Na 3 N, Na = +1, N = -3)

8. In compounds, F is ALWAYS –1 9. In compounds, O is almost always –2 10. In compounds, H is almost always +1 (alternative: -1 hydride) 11. In polyatomic ions, once the oxidation number is known it is to be considered the constant so long as the polyatomic ion is unchanged. (ex NH 4 + has N –3, H +1) ‘FUN’ = Q 1-3 p. 653, p. 657 - 659, Q 12-16 p. 659