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OXIDATION – REDUCTION REACTIONS

OXIDATION – REDUCTION REACTIONS. Redox. OXIDATION. Originally referred to oxygen combining with other elements to form oxides Principal sources of energy involve oxidation – combustion of gas, food metabolism in cells, etc.

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OXIDATION – REDUCTION REACTIONS

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  1. OXIDATION – REDUCTION REACTIONS • Redox

  2. OXIDATION • Originally referred to oxygen combining with other elements to form oxides • Principal sources of energy involve oxidation – combustion of gas, food metabolism in cells, etc. • Oxidation now has a broader meaning – refers to more than just “burning” and forming oxides • For example, bleaching and rusting are both oxidation processes

  3. REDUCTION • Opposite of oxidation • Originally meant the loss of oxygen • For Example: Iron ore is reduced in the production of metallic iron by heating it with charcoal– oxygen is removed from the ore

  4. REDOX REACTIONS • Oxidation and reduction occur simultaneously • Can’t have one without the other! • When producing iron from iron ore, the removed oxygen combines with carbon and forms carbon dioxide.

  5. REDOX TODAY • Includes many reactions that do not involve oxygen • Oxidation is redefined as complete or partial loss of electrons OR gain of oxygen • Reduction is complete or partial gain of electrons OR loss of oxygen

  6. LEO the lion goes GER

  7. OIL RIG

  8. Oxidation Complete loss of e- ionic reactions Shift of e- away from atom in covalent bond Gain of oxygen Loss of hydrogen by a covalent compound Increase in oxidation number Reduction Complete gain of e- ionic reactions Shift of e- toward an atom in a covalent bond Loss of oxygen Gain of hydrogen by a covalent compound Decrease in oxidation number Processes leading to oxidation and reduction

  9. Oxidation Numbers • Positive or negative number assigned to an atom • Equal to the ionic charge of monatomic ions • Ex. Br1- has an oxidation number of -1, Fe3+ has an oxidation number of +3 • Hydrogen has a +1 oxidation state except in metal hydrides it is -1 (NaH) • Uncombined atoms (free elements) have an oxidation number of 0

  10. Oxidation Numbers Con’t For neutral compounds sum of oxidation numbers is 0 The sum of the oxidation numbers of the atoms in a polyatomic ion is equal to the charge of the ion In a molecule, the oxidation number of the more electronegative atom is the same as if it were an ion. Ex. PBr5 the bromine is -1 and the phosphorus is +5

  11. Changing oxidation numbers • In polyatomic ions, the sum of the oxidation numbers is equal to the charge on the ion. • EX. SO32- Oxygen always (almost) has an oxidation number of -2 so when 3 oxygens bond with 1 sulfur to form a sulfite ion, the oxidation number of sulfur must be +4. • What is sulfur’s oxidation number in SO42-? • If you said +6 – Well done!

  12. Try these in your notes • Determine the oxidation number for sulfur in each of the following: • H2S Na2SO4 • S2Cl2 • Determine the oxidation number for the bold element in each of the following: • NaClO4 AlPO4 AsO43-

  13. Identifying Redox Reactions • Any type of reaction may be redox • Assign oxidation numbers to all elements to see if there is a transfer of electrons • If no change in oxidation #, then not redox • If there is a change, is redox

  14. Using oxidation # • Use oxidation numbers in an equation to determine if rxn is redox and if so, what is oxidized and reduced 0 +1-1 +1-1 0 • Cl2 + 2HBr  2HCl + Br2 • GER:Chlorineis reduced – • oxidation # changed from 0 to -1 • LEO:Bromineis oxidized – • oxidation # changed from -1 to 0

  15. Oxidation changes in reactions • Increase in oxidation # indicates oxidation (loss of electrons) • Decrease in oxidation # indicates reduction (gain of electrons) • Use oxidation numbers to determine what is oxidized and what is reduced +1 +5 -2 0 +2 +5 -2 0 • 2AgNO3 + Cu  Cu(NO3)2 + 2Ag

  16. Balancing Redox Reactions • Two methods: • Oxidation number change method – balanced by comparing the increases and decreases in oxidation numbers. • 1. assign oxidation # to all atoms in equation • 2. identify which are oxidized and reduced • 3. use a bracket to connect atoms that oxidize and another for those that are reduced. • 4. make the total increase in oxidation # equal to total decrease by using coefficients. This tells you the ratio of each.

  17. Oxidation number change method • Fe2O3 + CO  Fe + CO2unbalanced +3 -2 +2 -2 0 +4 -2 1. Fe2O3+ CO  Fe + CO2 Fe decreases and C increases 2. Draw brackets and show oxidation increase or decrease 3. Use coefficients to make total increase equal to total decrease. 4. Make sure equation is balanced for both atoms and charge.

  18. Half Reactions Method • Balance equation by balancing the oxidation half and the reduction half of reactions 1. Begin with the unbalanced, skeleton equation 2. Identify and write the equations for the oxidation and reduction half reactions. 3. Write separate half reactions for oxidation and reduction. 4. Balance the atoms in the half reactions • If needed, add H+ , and H2O as reactant or product to balance the ½ reaction.

  19. Half Reactions Continued 5. Add electrons to one side of each half reaction to balance charges (as a reactant if reduction and as product if oxidation). 6. Multiply each half reaction by the appropriate number to make number of electrons equal in both 7. Combine half reactions to show overall reaction. (e- should cancel)

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