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Redox Reactions

Redox Reactions. Redox Reactions. Redox is the abbreviated way to write reduction-oxidation reaction. Redox reactions are a very broad type of reaction and many of the types we have already discussed can also be classified as a redox reaction.

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Redox Reactions

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  1. Redox Reactions

  2. Redox Reactions • Redox is the abbreviated way to write reduction-oxidation reaction. • Redox reactions are a very broad type of reaction and many of the types we have already discussed can also be classified as a redox reaction. • In a redox reaction, reduction and oxidation must occur together. • But what is reduction and oxidation? • Reduction – chemical change in which electrons are gained. • Oxidation – a chemical change in electrons are removed.

  3. So, we need to Review Oxidation Numbers • What is the oxidation number? • Book definition – the number of electrons that must be added to or removed from an atom in a combined state to convert the atom into the elemental form. • In other words, the number of electrons gained or lost when a chemical bond is made. • Where can we find possible oxidation numbers? The Periodic Table

  4. Assigning Oxidation Numbers – The steps from page 606 • Identify the chemical formula If no formula is provided, write the formula of the formula or ion from the name. • Assign known oxidation numbers - The oxidation number of an atom of any free element in atomic or molecular form is zero. Examples: Cu, H2, Ag, Na, S8 - The oxidation number of monatomic ion is equal to the charge of the ion

  5. Assigning Oxidation Numbers – Continued - Fluorine always has an oxidation number of -1 since it is the most electronegative element. - In a compound: Group 1 atoms = Group 2 atoms = Aluminum = - Hydrogen = +1, unless it is with a metal, then it is -1 +1 +2 +3

  6. Assigning Oxidation Numbers – Continued • Oxygen is usually -2, exceptions: • If combined with Fluorine, Oxygen’s oxidation number becomes +2 • In peroxides, H2O2, Oxygen’s oxidation number becomes -1 • When comparing nonmetals, the more electronegative atom (page 194) gets the negative charge based on the column it is located in on the periodic table. • Group 17 =, Group 16 = , Group 15 = __ -1 -3 -2

  7. Assigning Oxidation Numbers – Continued • Calculate any remaining numbers, and verify results. • Use the total oxidation number of each element’s atoms = oxidationnumber times subscript • The sum of the oxidation numbers for all the atoms in a molecule is zero. • The sum of the oxidation numbers for all the atoms in a polyatomicion is equal to the charge on that ion.

  8. Examples +2 -2 = 0 +1 -2 -3 +4 = +1 -3 +1 +2 +6 -8= 0 +2 +6 -2 NH4+ PbSO4 Ca(OH)2 Al H2 Fe2(CO3)3 H2O KClO3 H2PO4-1 Pb+2 BF3 S2O7-2 +6 +12 -18 = 0 +3 +4 -2 0 0 +1 +5 -6 = 0 +1 +5 -2 +2 +5 -8 = -1 +1 +5 -2 +2 -2 = 0 +1 -2 +12 -14 = -2 +6 -2 +3 -3 = 0 +3 -1 +2

  9. Now that we reviewed oxidation numbers, what do we do with them? • Remember, in a redox reaction, something has to be oxidized (loses electrons, oxidation number increases) and something else has to be reduced (gains electrons, oxidation number decreases). • If we assign oxidation numbers to all atoms in a chemical equation, we can determine whether it is a redox reaction.

  10. Are the following redox reactions? 2 H3O+ + Zn  H2 + 2H2O + Zn+2 C3H8 + 5O2  3CO2 + 4H2O 2KI + Pb(NO3)2  PbI2 + 2KNO3 2K + Cl2  2KCl

  11. Some other terms • Oxidizing agent – the substance that gains electrons in an redox reaction and is reduced. • Reducing agent - the substance that loses electrons in an redox reaction and is oxidized. • Notice that the oxidation agent oxidizes another substance and becomes reduces in the process and visa versa • Also note that these refer to the whole compound or ion.

  12. Let’s look at the following redox reaction First, write the oxidation number for each atom: 8 H3O+ + MnO4- + 5Fe+2 Mn+2 + 5Fe+3 + 12 H2O What was reduced? What was oxidized? What is the oxidizing agent? What is the reducing agent? Mn Fe MnO4- Fe+2

  13. Homework • Page 611: 6, 7, 8, 9

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