1. Redox Reactions
Reactions in which one or more electrons are transferred.
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2. Reaction of Sodium and Chlorine
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3. Rules for Assigning Oxidation States
Oxidation state of an atom in an element = 0
Oxidation state of monatomic ion = charge of the ion
Oxygen = -2 in covalent compounds (except in peroxides O22-where it = -1)
Hydrogen = +1 in covalent compounds
Fluorine = -1 in compounds
Sum of oxidation states = 0 in compounds
Sum of oxidation states = charge of the ion in ions
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4. EXERCISE!
Find the oxidation states for each of the elements in each of the following compounds:
K2Cr2O7
CO32-
MnO2
PCl5
SF4
K = +1; Cr = +6; O = –2
C = +4; O = –2
Mn = +4; O = –2
P = +5; Cl = –1
S = +4; F = –1
K2Cr2O7; K = +1; Cr = +6; O = -2
CO32-; C = +4; O = -2
MnO2; Mn = +4; O = -2
PCl5; P = +5; Cl = -1
SF4; S = +4; F = -1
Non-integer oxidation states are possible in rare cases (ex. Fe3O4)
Fe = +8/3; O = -2
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5. Redox Characteristics
Transfer of electrons
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6. Redox Characteristics
Transfer of electrons
Transfer may occur to form ions
Oxidation – increase in oxidation state (loss of electrons); reducing agent (reductant)
Reduction – decrease in oxidation state (gain of electrons); oxidizing agent (oxidant)
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7. Which of the following are oxidation-reduction reactions?
CONCEPT CHECK!
Which of the following are oxidation-reduction reactions?
Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)
Cr2O72-(aq) + 2OH-(aq) CrO42-(aq) + H2O(l)
2CuCl(aq) CuCl2(aq) + Cu(s)
Zn – reducing agent; HCl – oxidizing agent
c) CuCl acts as the reducing and oxidizing agent
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8. Balance an equation showing the oxidation of Fe2+ to Fe3+ by dichromate ions in an acidic medium. As a result, the Cr2O72- ions are reduced to Cr 3+ ions.
Write the unbalanced equation for the reaction in ionic form.

9. Separate the equations into two half reactions.

10. Balance each half- reaction for the number and types of atoms and charges. For reactions in an acidic medium, add H2O to balance the O atoms and H+ to balance the H atoms.

11. Balance each half- reaction for the number and types of atoms and charges. For reactions in an acidic medium, add H2O to balance the O atoms and H+ to balance the H atoms.

12. Add the two half- equations together and balance the final equation by inspection. The electrons on both sides must cancel. If the oxidation and reduction half-reactions contain different numbers of electrons, we need to multiply one or both to equalize the number of electrons.

13. Verify that the equations contains the same type and numbers of atoms and the same charges on both sides of the equation.

14. If the reaction occurs in a basic medium, we proceed through step 4 as above. Then, for every H+ ion we add an equal number of OH- to both sides of the equation. Where H+ and OH- appear on the same side of the equation, we combine the ions to give H2O.

15. Write a balanced ionic equation to represent the oxidation of iodide ion (I-) by permanganate ion (MnO4-) in basic solution to yield molecular iodine (I2) and manganese (IV) oxide (MnO2).