Redox

• Redox Reactions


  A type of chemical reaction that involves a transfer of electrons between two species
• Oxidation


  The loss of electrons
• Reduction


  The gain of electrons
• Common mnemonics for oxidation and reduction

  • LEO the lion goes GER
  • OIL RIG
• Oxidizing agent


  The substance that gains electrons in a redox reaction
• Reducing agent


  The substance that loses electrons in a redox reaction
• Copper is oxidized because its oxidation number increases from 0 to +2
• Silver is reduced because its oxidation number decreases from +1 to 0
• Copper is the reducing agent
• Silver nitrate is the oxidizing agent
• Oxidation number


  The net electric charge that the atoms would have if electron pairs in covalent bonds belong entirely to the more electronegative atom
• Rules for assigning oxidation numbers

  • All pure elements is 0
  • Monoatomic (single atom) ions are the same as its charge
  • Fluorine is ALWAYS -1
  • Hydrogen – most compounds +1, metal hydrides -1
  • Oxygen – most compounds -2, in peroxides -1, in OF2 +2
  • Covalent compound without H, O: More EN element ON equals its ionic charge
  • Zero-Sum Rule: The sum of the oxidation numbers of all the atoms in a compound must equal the overall charge for the compound

  
• Oxidation numbers are denoted by a superscript above the element symbol, with the order of the number and sign being important
• Steps to balance redox reactions using the oxidation numbers method

  1. Write the unbalanced equation and assign oxidation numbers

  2. Identify changes in oxidation numbers as oxidation or reduction

  3. Balance the electrons transferred by multiplying the coefficients

  4. Balance the rest of the equation by inspection. If necessary, balance oxygen by adding water and hydrogen by adding H+

  5. If the redox reaction occurs in a basic solution, add OH- to both sides to neutralise the H+ and then eliminate the water molecules that occur on both sides of the equation

  6. Check that the elements and charges are balanced

  
• Steps to balance redox reactions using the half-reactions method


  1. Write separate equations for the oxidation and reduction half reactions

  2. For each half reaction: balance all the elements except for hydrogen and oxygen, balance oxygen using H2O, balance hydrogen using H+, balance the charge using electrons

  3. If necessary, multiple one or both balanced half-reactions by an integer to equalise the number of electrons transferred in the two half-reactions

  4. Add the half-reactions, and cancel identical species

  5. Check that the elements and charges are balanced

  6. If the redox occurs in a basic solution, add OH- to both sides to neutralise the H+ and eliminate the water molecules that occur on both sides of the equation

  7. Check that the elements and charges are balanced

  
• Activity Series


  Elements can be arranged in order of how easily they are oxidized to form positive ions. Elements higher up the series form positive ions more easily, and are more reactive as they lose their electrons more readily.
• Redox tables


  List common oxidizing and reducing agents in order of the relative strength, typically measured against a reference electrode

  
• As metals lose electrons and form ions in redox reactions, they are oxidized and act as reducing agents. The most reactive metals have the greatest tendency to lose electrons and are strong reducing agents.
• Ions that tend to gain electrons are reduced, and act as oxidizing agents. The order of reactivity of the ions is the order of strength of the oxidizing agents.
• Spontaneity Rule


  A spontaneous reaction occurs only if the reducing agent (RA) transfers electrons to the oxidizing agent (OA)
• IB Data booklets: Dec RA (E0 (V) inc)

  Can displace metals below (consistent with Activity Series)
• OSSD Redox Tables: Inc RA (E0 (V) dec)


  Can displace metals above (opposite pattern than Activity Series)