Electrochemistry

• Oxidation state

  The total number of electrons which have been removed from an element (producing a positive oxidation state) or added to an element (producing a negative oxidation state) to reach its present state
• Reduction


  1. Decrease in the oxidation state

  2. Gain of electron(s)

  3. Gain of hydrogen

  4. Loss of oxygen
• Oxidation

  1. Increase in the oxidation state

  2. Loss of electron(s)

  3. Loss of hydrogen

  4. Gain of oxygen
• Reducing agent

  The compound that is oxidized (causes other particles to be reduced)
• Oxidising agent

  The compound that is reduced (causes other particles to be oxidized)
• Redox reaction


  A chemical reaction involving both oxidation and reduction, in which one reactant is oxidized and another is reduced.
• Balancing redox reactions/equations

  1. Separate and balance the reduction and oxidation half equations

  2. Acidic: add H+ for more hydrogen and H2O to oxygen deficient side

  3. In basic reactions, add H2O for more hydrogen and OH- to oxygen deficient side

  4. Equalize the number of electrons lost and gained

  5. Add the results in half equations to get the overall balanced reaction
• Oxidation


  Loss of electrons
• Reduction


  Gain of electrons
• WRITING IONIC EQUATIONS

  1. Determine the species soluble in water and therefore ionised or dissociated

  2. Assign oxidation number to each element

  3. Remove spectator ions
• Spectator ions


  Those that exist both as a reactant and a product in a chemical equation of an aqueous solution
• The ionic equation is the simplified chemical equation that only shows the ions that are directly involved in the reaction
• Group 1 elements

  • Always +1 oxidation state
• Group 2 elements

  • Always +2 oxidation state
• Fluorine

  Oxidation state is always -1
• Hydrogen


  Oxidation state is +1, EXCEPTION -1 in metal hydrides
• Oxygen

  Oxidation state is -2, EXCEPTION -1 in peroxides and +2 in
• The sum of the oxidation state in a compound is zero
• The sum of the oxidation state in an ion is equal to the charge on the ion
• To determine the oxidation state of an element in a compound or ion, the more electronegative element is given the higher oxidation state
• All (uncombined) elements have an oxidation state of

  Zero
• Oxidation

  The gain of oxygen by a substance or the loss of electrons
• Reduction

  The loss of oxygen by a substance or the gain of electrons
• A reaction in which oxidation and reduction occurs at the same time is called a Redox reaction
• Oxidation number

  Indicates the oxidation state of an element in its free state or in a compound
• Rules for determining oxidation numbers

  • The oxidation state of atoms of an element in its free state is zero
  • The oxidation number of elements that exist as simple ions in ionic compounds is the same as the charge on the ion
  • The sum of all the oxidation numbers of elements in a compound is zero
  • The sum of all the oxidation numbers of the elements in a radical is equal to the charge on the ion
  • Hydrogen is usually +1 except in hydrides where it is -1
  • Oxygen is usually -2 except in peroxides where it is -1
• Oxidizing agent

  A substance that is reduced in the process of oxidizing another substance
• Reducing agent

  A substance that is oxidized in the process of reducing another substance
• Oxidizing agents

  • Acidified potassium manganate (VII) (changes from purple to colourless)
  • Acidified potassium dichromate (VI) (changes from orange to green)
  • Iron(III) salts (changes from yellow to pale green)
  • Concentrated sulfuric acid (produces sulfur dioxide)
  • Nitric acid (produces nitrogen dioxide)
• Reducing agents

  • Potassium iodide (changes from colourless to brown)
  • Iron(II) salts (changes from pale green to yellow)
  • Hydrogen sulfide (produces a yellow precipitate of sulfur)
  • Concentrated hydrochloric acid (produces chlorine gas)
  • Carbon, carbon monoxide, hydrogen
• Substances that can behave as both oxidizing and reducing agents: sulfur dioxide, acidified hydrogen peroxide
• Testing for oxidizing agents

  1. Add potassium iodide (changes from colourless to brown)

  2. Add iron(II) salt (changes from pale green to yellow)

  3. Add hydrogen sulfide
• Testing for reducing agents

  1. Add potassium manganate (VII) (changes from purple to colourless)

  2. Add acidified potassium dichromate (VI) (changes from orange to green)

  3. Add iron(III) salt (changes from yellow to pale green)
• If a substance is being oxidized

  Its oxidation number increases
• If a substance is being reduced

  Its oxidation number decreases
• Identifying oxidized and reduced substances

  • Zn (oxidized from 0 to +2)
  • Fe2+ (oxidized from +2 to +3)
  • Cu2+ (reduced from +2 to 0)
  • N (oxidized from -3 to 0)
• The oxidizing agent

  Is reduced in the reaction
• The reducing agent

  Is oxidized in the reaction