Oxidation-Reduction Reactions and Single Replacement Reactions

LP 3 - Oxidation-Reduction Reactions (redox)

  • Objectives:
    • Define oxidation and reduction in terms of electron movement and changes in oxidation numbers.
    • Identify the oxidation number of every species in an equation to determine which species undergo oxidation and reduction.
    • Use an activity series to determine if a single replacement reaction will take place.
    • Predict the products of and balance single replacement reactions.
  • LEO goes GER: Loss of Electrons is Oxidation, Gain of Electrons is Reduction.
  • OIL RIG: Oxidation Is Loss, Reduction Is Gain.
  • In some chemical reactions, electrons are permanently transferred from one species to another.
  • One species gives electrons away (oxidation), and another species picks them up (reduction).
  • This process always happens in tandem because electrons are matter and cannot be created or destroyed during a chemical reaction, just like whole atoms.

Oxidation Number Rules

  • Oxidation numbers are hypothetical numbers assigned to atoms in a reaction to help determine which species are being oxidized and reduced and how many electrons are moving during the reaction.
  • The only way to know the oxidation numbers for different species is to memorize a complex set of rules.
  • The following rules will cover MOST situations encountered in a general chemistry course:
    1. Elemental species have oxidation numbers equal to zero.
    2. Monatomic ions have oxidation numbers that are the same magnitude and sign as their ionic charge.
    3. Hydrogen, oxygen, and halogens have special rules:
      • The oxidation number of hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals.
        • Example: In H_2O, hydrogen has an oxidation number of +1.
        • Example: In NaH, hydrogen has an oxidation number of -1.
      • Oxygen has an oxidation number equal to -2 except when it is found as peroxide ion (O_2^{2-}). Oxygen in peroxide ion has an oxidation number of -1.
        • Example: In H_2O, oxygen has an oxidation number of -2.
        • Example: In H2O2, oxygen has an oxidation number of -1.
      • Fluorine has an oxidation number of -1.
      • Other halogens (Cl, Br, I) are usually -1 when found in binary molecular compounds, but must be calculated using rule #4 when they are not.
    4. The sum of the oxidation numbers in a species will be equal to the charge of the species. Unknown oxidation numbers can be calculated using this fact.
      • Example: Determine the oxidation number of sulfur in SO_4^{2-}.
        • Let x be the oxidation number of sulfur.
        • x + 4(-2) = -2
        • x - 8 = -2
        • x = +6
        • Therefore, the oxidation number of sulfur in SO_4^{2-} is +6.

Oxidation Number Flowchart

  • Flowchart to determine oxidation numbers:
    • Is it elemental? If yes, oxidation number is 0.
    • If not elemental: Is it an ionic compound?
      • If yes: Split it. Are any of the ions monatomic?
        • If yes, oxidation number is equal to charge of the ion.
    • If not an ionic compound: Is it a standalone ion?
      • If yes, oxidation number is equal to the charge of the ion.
    • If not a standalone ion: Is oxygen present?
      • If yes: Is the oxygen part of peroxide?
        • If yes, oxidation number is -1.
        • If no, oxidation number is -2.
    • If oxygen is not present: Is hydrogen present?
      • If yes: Is the hydrogen bonded to a nonmetal?
        • If yes, oxidation number is +1.
        • If no, oxidation number is -1 (bonded to metal).
    • If hydrogen is not present: Is fluorine present?
      • If yes, oxidation number is -1.
    • If fluorine is not present: Are chlorine, bromine, or iodine present?
      • If yes: Is the compound binary?
        • If yes, oxidation number is -1.
        • If no, write an equation, set unknown = x, set equation equal to charge of species, and solve for x.

Single Replacement Reactions: The REALLY Bad Prom Date

  • General form: AX + B \rightarrow BX + A
  • Example: Zinc is dropped into a test tube of copper (II) sulfate solution.
  • Example: A sample of iron is submerged in acetic acid.
  • Example: Manganese is added to a beaker of lead (II) nitrate solution.
  • Example: Solid iodine is mixed with a solution of magnesium chloride.

The Activity Series

  • Whether or not single replacement reactions take place is determined by the activity series.
  • The species that undergoes oxidation must be higher on the table than the species that undergoes reduction.
  • Example: Lead pipes are exposed to drain cleaner containing hydrochloric acid vs. copper pipes exposed to the same cleaner.
    • The more active metal will replace the less active metal in a compound.