oxidation number

Examples of oxidation number in the following topics:

• Oxidation Numbers of Metals in Coordination Compounds

  • O2- and S2- have oxidation numbers of -2.
  • In a molecule or compound, the oxidation number is the sum of the oxidation numbers of its constituent atoms.
  • Solution: The oxidation number of C is -3.
  • The oxidation number of H is +1 (H+ has an oxidation number of +1).
  • What is the oxidation number of chromium?

• Electronegativity and Oxidation Number

  • Six rules can be used when assigning oxidation numbers:
  • In compounds with nonmetals, the oxidation number of hydrogen is +1.
  • Oxygen is assigned an oxidation number of -2 in most compounds.
  • In oxygen difluoride (OF2), the oxidation number of oxygen is +2, while in dioxygen difluoride (O2F2), oxygen is assigned an oxidation number of +1 because fluorine is the more electronegative element in these compounds, so it is assigned an oxidation number of -1.
  • Apply the rules for assigning oxidation numbers to atoms in compounds

• Naming Coordination Compounds

  • Transition-metal and coordination compounds are named using a set of rules that describe oxidation numbers and anion and cation composition.
  • The oxidation number appears as a Roman numeral in parenthesis after the cation.
  • For metals, the oxidation number is the same as the charge.
  • Therefore, the platinum oxidation number is +4.
  • However, the brackets as well as the different oxidation number of the platinum result in a very different name.

• Manganese

  • The most common oxidation states of the metal manganese are +2, +3, +4, +6, and +7; the +2 oxidation state is the most stable.
  • Manganese compounds where manganese is in oxidation state of 7+ are powerful oxidizing agents.
  • Compounds with oxidation states 5+ (blue) and 6+ (green) are strong oxidizing agents.
  • The most stable oxidation state (oxidation number) for manganese is 2+, which has a pale pink color, and many manganese(II) compounds are common, such as manganese(II) sulfate (MnSO4) and manganese(II) chloride (MnCl2).
  • Predict the oxidation or reduction propensity of a manganese species given its formula or oxidation state.

• Oxidations & Reductions

  • To determine whether a carbon atom has undergone a redox change during a reaction we simply note any changes in the number of bonds to hydrogen and the number of bonds to more electronegative atoms such as O, N, F, Cl, Br, I, & S that has occurred.
  • If the number of hydrogen atoms bonded to a carbon increases, and/or if the number of bonds to more electronegative atoms decreases, the carbon in question has been reduced (i.e. it is in a lower oxidation state).
  • If the number of hydrogen atoms bonded to a carbon decreases, and/or if the number of bonds to more electronegative atoms increases, the carbon in question has been oxidized (i.e. it is in a higher oxidation state).
  • If there has been no change in the number of such bonds, then the carbon in question has not changed its oxidation state.
  • Peracid epoxidation and addition of bromine oxidize both carbon atoms, so these are termed oxidation reactions.

• Oxides

  • Metal oxides typically contain an anion of oxygen in the oxidation state of −2.
  • Most of the Earth's crust consists of solid oxides, the result of elements being oxidized by the oxygen in air or water.
  • In these oxides, the coordination number of the oxide ligand is 2 for most electronegative elements, and 3–6 for most metals.
  • Although most metal oxides are polymeric, some oxides are monomeric molecules.
  • Those attacked only by acids are basic oxides; those attacked only by bases are acidic oxides.

• Oxidation States

  • For example, Cl- has an oxidation state of -1.
  • When present in most compounds, hydrogen has an oxidation state of +1 and oxygen an oxidation state of −2.
  • Generally, the oxidation state for most common elements can be determined from their group number on the periodic table.
  • As stated in rule number four above, the sum of the oxidation states for all atoms in a molecule or polyatomic ion is equal to the charge of the molecule or ion.
  • Predict the oxidation states of common elements by their group number.

• Phosphorus Compounds as Reducing Agents

  • Trivalent phosphorus is easily oxidized.
  • The triphenylphosphine oxide produced in reactions 1 & 3 is a very stable polar compound, and in most cases it is easily removed from the other products.
  • Triphenylphosphine is also oxidized by halogens, and with bromine yields dibromotriphenylphosphorane, a crystalline salt-like compound, useful for converting alcohols to alkyl bromides.
  • As in a number of earlier examples, the formation of triphenylphosphine oxide in the irreversible SN2 step provides a thermodynamic driving force for the reaction.

• Oxidation of Reduced Sulfur Compounds

  • Sulfur oxidation involves the oxidation of reduced sulfur compounds, inorganic sulfur, and thiosulfate to form sulfuric acid.
  • Sulfur oxidation involves the oxidation of reduced sulfur compounds such as sulfide (H2S), inorganic sulfur (S0), and thiosulfate (S2O2−3) to form sulfuric acid (H2SO4).
  • An example of a sulfur-oxidizing bacterium is Paracoccus.
  • The two step process occurs because sulfide is a better electron donor than inorganic sulfur or thiosulfate; this allows a greater number of protons to be translocated across the membrane.
  • In addition to aerobic sulfur oxidation, some organisms (e.g.

• Chromium

  • A large number of chromium(III) compounds are known.
  • It is dehydrated by heating to form the green chromium(III) oxide (Cr2O3), which is the stable oxide with a crystal structure identical to that of corundum.
  • Chromium(VI) compounds are powerful oxidants at low or neutral pH.
  • Both the chromate and dichromate anions are strong oxidizing reagents at low pH.
  • The oxidation state +5 is only realized in few compounds but are intermediates in many reactions involving oxidations by chromate.