User:Drew.P.Branch/sandbox

Original text

Consider the reaction below:

Cl₂ + 2Fe²⁺ → 2Cl⁻ + 2Fe³⁺

The two elements involved, iron and chlorine, each change oxidation state; iron from 2+ to 3+, chlorine from 0 to 1−. There are then effectively two half-reactions occurring. These changes can be represented in formulas by inserting appropriate electrons into each half-reaction:

Fe²⁺ → Fe³⁺ + e⁻

Cl₂ + 2e⁻ → 2Cl⁻

In the same way given two half-reactions it is possible, with knowledge of appropriate electrode potentials, to arrive at the full (original) reaction.

Revised text

Consider the reaction below:

   Cl2 + 2Fe2+ → 2Cl- + 2Fe3+

The two elements involved, iron and chlorine, each change oxidation state; iron from 2+ to 3+, chlorine from 0 to 1−. There are then effectively two half-reactions occurring. These changes can be represented in formulas by inserting appropriate electrons into each half-reaction:

   Fe2+ → Fe3+ + e-
   Cl2 + 2e- → 2Cl-

Given two half-reactions it is possible, with knowledge of appropriate electrode potentials, to arrive at the full (original) reaction the same way. The decomposition of a reaction into half-reactions is key to understanding a variety of chemical processes. For example, in the above reaction, it can be shown that this is a redox reaction in which Fe is reduced, and Cl is oxidized. Note the transfer of electrons from Fe to Cl. Decomposition is also a way to simplify the balancing of chemical equations. A chemist can atom balance and charge balance one piece of an equation at a time.

For example:

• Fe²⁺ → Fe³⁺ + e^- becomes 2Fe²⁺ → 2Fe³⁺ + 2e^-
• is added to Cl₂ + 2e^- → 2Cl^-
• and finally becomes Cl₂ + 2Fe²⁺ → 2Cl^- + 2Fe³⁺