Redox Chemistry

Introduction

• Some redox reactions involve the transfer of electrons between reactant species to yield ionic products
•  It is helpful to split the overall reaction into individual equations called half-reactions.
• In the balanced equation, electrons lost must = electrons gained.
• Oxidation: loss of electrons
• Reduction: gain of electrons
• Reducing agent: is oxidized; LOSES electrons, and its charge goes up.
• Oxidizing agent: is reduced; gains electrons, and its charge goes DOWN.
• Oxidation number (or oxidation state) of an element in a compound: the charge its atoms would possess if the compound was ionic. 

Rules for Assigning Oxidation Numbers

• The oxidation number of an atom in an elemental substance is zero.
• The oxidation number of a monatomic ion is equal to the ion’s charge.   ( Group 1 metals are always +1,  group 2 always +2)
• Oxidation numbers for common nonmetals are usually assigned as follows: 
  • Hydrogen: 
  • +1 when combined with nonmetals
  • −1 when combined with metals
  • Oxygen: 
  • −2 in most compounds
  • Sometimes −1 (so-called peroxides, O22−)
  • Very rarely − ½  (so-called superoxides, O2−)
  • Positive values when combined with F (values vary)
  • Halogens: 
  • −1 for F always 
  • −1 for other halogens except when combined with oxygen or other halogens (positive oxidation numbers in these cases, varying values)
• The sum of the oxidation numbers for all atoms in a molecule must be zero.
• The sum of the oxidation number in a polyatomic ion equals the charge on the ion.  
• The oxidation number of an element in a polyatomic ion is always the same, whether that ion is in a compound or not

Types of Redox Reactions

• Single-displacement (replacement) reactions: redox reactions in which an ion in solution is displaced (or replaced) via the oxidation of a metallic element.

Steps for Balancing Redox Reactions

• Write the two individual half-reactions: the oxidation and the reduction. 
• Balance all elements except oxygen and hydrogen. (most often forgotten, leading to tragic results! )
• Balance oxygen atoms by adding H2O molecules.
• Balance hydrogen atoms by adding H+ ions.
• Balance charge by adding electrons.
• If necessary, multiply each half-reaction’s coefficients by the smallest possible integers to yield equal numbers of electrons in each.
• Add the balanced half-reactions together and simplify by removing species that appear on both sides of the equation
• For reactions occurring in basic media (excess hydroxide ions), carry out these additional steps:
  • Add OH− ions to both sides of the equation in numbers equal to the number of H+ ions.
  • On the side of the equation containing both H+ and OH− ions, combine these ions to yield water molecules.
  • Simplify the equation by removing any redundant water molecules.
• Finally, check to see that the number of atoms and the total charge are balanced.

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