Redox Reactions

• Chemistry deals with varieties of matter and the transformation of matter from one kind into another through various types of reactions
• Redox reactions are important in pharmaceutical, biological, industrial, metallurgical, and agricultural areas
• Redox reactions are used in burning fuels for energy, electrochemical processes, manufacturing chemical compounds, operating batteries, and corrosion of metals
• The classical idea of redox reactions involves oxidation and reduction reactions
• Originally, oxidation was described as the addition of oxygen to an element or compound
• Oxidation was later reinterpreted to include the removal of hydrogen from a substance
• Reduction was initially considered as the removal of oxygen from a compound
• Reduction has been broadened to include the removal of oxygen/electronegative elements or the addition of hydrogen/electropositive elements to a substance
• Oxidation involves the addition of oxygen/electronegative elements or the removal of hydrogen/electropositive elements from a substance
• Reduction involves the removal of oxygen/electronegative elements or the addition of hydrogen/electropositive elements to a substance
• Redox reactions involve simultaneous oxidation and reduction processes
• Redox reactions can be defined in terms of electron transfer reactions
• In redox reactions, half reactions involving the loss of electrons are oxidation reactions, while half reactions involving the gain of electrons are reduction reactions
• In reactions involving electron transfer, sodium acts as a reducing agent by donating electrons to elements like chlorine, oxygen, and sulfur
• Chlorine, oxygen, and sulfur act as oxidizing agents because they accept electrons from sodium
• Oxidation is the loss of electron(s) by any species, while reduction is the gain of electron(s) by any species
• An oxidizing agent is an acceptor of electron(s), while a reducing agent is a donor of electron(s)
• In the reaction 2 Na(s) + H2(g) → 2 NaH(s), sodium is oxidized and hydrogen is reduced, making it a redox change
• In the reaction between metallic zinc and an aqueous solution of copper nitrate, zinc is oxidized to form Zn2+ ions, while copper ions are reduced to form metallic copper
• The competition for electrons between metals can be used to design Galvanic cells, where chemical reactions become a source of electrical energy
• Oxidation number denotes the oxidation state of an element in a compound, ascertained according to a set of rules where electron pair in a covalent bond belongs entirely to the more electronegative element
• Rules for calculating oxidation numbers include:

  • In elements in the free state, each atom has an oxidation number of zero
  • For ions composed of only one atom, the oxidation number is equal to the charge on the ion
  • The oxidation number of oxygen in most compounds is -2
• Oxidation numbers:

  • All alkaline earth metals have an oxidation number of +2
  • Aluminium has an oxidation number of +3 in all its compounds
• Oxygen's oxidation number in most compounds is -2

  • Exceptions include peroxides and superoxides where each oxygen atom is assigned -1 and -(½) respectively
  • Another exception is when oxygen is bonded to fluorine, resulting in oxidation numbers of +2 and +1
• Hydrogen's oxidation number is +1, except in binary compounds with metals where it is -1
• Fluorine has an oxidation number of -1 in all its compounds

  • Other halogens (Cl, Br, I) also have an oxidation number of -1 when they occur as halide ions
  • When combined with oxygen, halogens have positive oxidation numbers
• The highest oxidation number of a representative element is the group number for the first two groups and the group number minus 10 for other groups
• Types of Redox Reactions:

  1. Combination reactions:

  • A + B → C
  • Involves elements in elemental form

  2. Decomposition reactions:

  • Lead to breakdown of compounds into elements

  3. Displacement reactions:

  • Metal displacement and non-metal displacement
• In a displacement reaction, an ion in a compound is replaced by an ion of another element
• Metal displacement:

  • A metal in a compound is displaced by another metal in the uncombined state
• Problem 7.4:

  • Justify the redox reaction: 2Cu2O(s) + Cu2S(s) → 6Cu(s) + SO2(g)
  • Identify the species oxidised/reduced, oxidant, and reductant
• Stock notation:

  • Represents the oxidation number of a metal in a compound
  • Uses Roman numerals in parenthesis after the metal symbol in the molecular formula
• Redox reactions involve changes in oxidation numbers

  • Oxidation: Increase in oxidation number
  • Reduction: Decrease in oxidation number
  • Oxidising agent: Increases oxidation number
  • Reducing agent: Decreases oxidation number
• The algebraic sum of oxidation numbers in a compound must be zero

  • In polyatomic ions, the sum of oxidation numbers must equal the charge on the ion
• Solution:

  • Copper is reduced from +1 to zero oxidation state
  • Sulphur is oxidised from -2 to +4 oxidation state
  • Copper acts as an oxidant and sulphur as a reductant
• Types of Redox Reactions:

  1. Combination reactions

  2. Decomposition reactions

  3. Displacement reactions
• Displacement reactions involve the replacement of an ion in a compound by an ion of another element
• Metal displacement involves a metal in a compound being displaced by another metal in the uncombined state
• Metal displacement reactions involve a metal in a compound being displaced by another metal in the uncombined state
• These reactions find applications in metallurgical processes to obtain pure metals from their compounds in ores