Oxidation Reduction Reactions
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Transcript of Oxidation Reduction Reactions
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OXIDATION REDUCTION REACTIONS
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Oxidation Number Rules The oxidation number of any free,
uncombined element is zero. The oxidation number of an element in a
simple (monatomic) ion is the charge on the ion.
In the formula for any compound, the sum of the oxidation numbers of all elements in the compound is zero. In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion.
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Oxidation Number Rules
Hydrogen, H, in combined form, has the oxidation number +1, usually.
Oxygen, O, in combined form, has the oxidation number -2, usually.
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Oxidation Numbers assign oxidation numbers to each
element in these compoundsNaNO3 K2Sn(OH)6 H3PO4
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Oxidation Numbers assign oxidation numbers to each
element in these compoundsNaNO3 K2Sn(OH)6 H3PO4
+1 +5 -2
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Oxidation Numbers assign oxidation numbers to each
element in these compoundsNaNO3 K2Sn(OH)6 H3PO4
+1 +5 -2 +1 +4 -2 +1
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Oxidation Numbers assign oxidation numbers to each
element in these compoundsNaNO3 K2Sn(OH)6 H3PO4
+1 +5 -2 +1 +4 -2 +1 +1 +5 -2
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Oxidation Numbers
assign oxidation numbers to each element in these polyatomic ionsSO3
2- HCO3- Cr2O7
2-
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Oxidation Numbers
assign oxidation numbers to each element in these polyatomic ionsSO3
2- HCO3- Cr2O7
2-
+4 -2
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Oxidation Numbers
assign oxidation numbers to each element in these polyatomic ionsSO3
2- HCO3- Cr2O7
2-
+4 -2 +1+4 -2
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Oxidation Numbers
assign oxidation numbers to each element in these polyatomic ionsSO3
2- HCO3- Cr2O7
2-
+4 -2 +1+4 -2 +6 -2
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Redox Reactions oxidation - loss of electrons
oxidation number increases reduction - gain of electrons
oxidation number decreases or reduces! oxidizing agents - chemical species that
gain electrons and oxidize some other species electron thieves they are reduced
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Redox Reactions reducing agents - chemical species that
lose electrons and reduce some other species victims of electron thieves they are oxidized
Predicting reactions?
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Activity Series Allows us to predict if a certain metal will
be oxidized by an acid or a particular salt. Activity series is arranged in order of
decreasing ease of oxidation Metals at the top lose electrons easily Metals at the bottom are very stable Any metal can be oxidized by the ions of
the substances below it.
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Balancing Redox Reactions
Balancing redox reactions.
We will use the Half-Reaction Method
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Half Reaction Method Balancing redox reactions by half reaction
method rules: Write the unbalanced reaction Break into 2 half reactions -
1 for oxidation 1 for reduction
Mass balance each half reaction by adding appropriate stoichiometric coefficents in acidic solutions we can add H+ or H2O in basic solutions we can add OH- or H2O
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Half Reaction Method
Charge balance the half reactions by adding appropriate numbers of electrons
Multiply each half reaction by a number to make the number of electrons added equal to the number of electrons consumed
Add the two half reactions Eliminate any common terms
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Half Reaction Method
Example: Tin (II) ions are oxidized to tin (IV) by bromine. Use the half reaction method to write and balance the net ionic equation.
reaction starting BrSnBrSn -+42
+2
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Half Reaction Method
Example: Tin (II) ions are oxidized to tin (IV) by bromine. Use the half reaction method to write and balance the net ionic equation.
reaction halfoxidation e2Sn Sn
reaction starting BrSnBrSn-+4+2
-+42
+2
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Half Reaction Method
Example: Tin (II) ions are oxidized to tin (IV) by bromine. Use the half reaction method to write and balance the net ionic equation.
reaction halfreduction Br 2e2Br
reaction halfoxidation e2Sn Sn
reaction starting BrSnBrSn
--2
-+4+2
-+42
+2
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Half Reaction Method
Example: Tin (II) ions are oxidized to tin (IV) by bromine. Use the half reaction method to write and balance the net ionic equation.
Sn Br Sn Br starting reaction
Sn Sn e oxidation half reaction
Br e 2 Br reduction half reaction
Sn Br Sn Br balanced reaction
2+2
4+ -
2+ 4+ -
2- -
2+2
4+ -
2
2
2
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Half Reaction Method Example: Dichromate ions oxidize iron (II) ions to
iron (III) ions and are reduced to chromium (III) ions in acidic solution. Write and balance the net ionic equation for the reaction.
reaction starting FeCrFeOCr +3+3+2272
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Half Reaction Method Example: Dichromate ions oxidize iron (II) ions to
iron (III) ions and are reduced to chromium (III) ions in acidic solution. Write and balance the net ionic equation for the reaction.
reaction halfoxidation 1e+ FeFe
reaction starting FeCrFeOCr-+3+2
+3+3+2272
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Half Reaction Method Example: Dichromate ions oxidize iron (II) ions to
iron (III) ions and are reduced to chromium (III) ions in acidic solution. Write and balance the net ionic equation for the reaction.
red. OH 7Cr 2 e 6H 14OCr
ox. 1e+ FeFe
rxnstart FeCrFeOCr
2+3-+2
72
-+3+2
+3+3+2272
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Half Reaction Method Example: Dichromate ions oxidize iron (II) ions to
iron (III) ions and are reduced to chromium (III) ions in acidic solution. Write and balance the net ionic equation for the reaction.
red. OH 7Cr 2 e 6H 14OCr
ox. 1e+ FeFe 6
rxnstart FeCrFeOCr
2+3-+2
72
-+3+2
+3+3+2272
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Half Reaction Method Example: Dichromate ions oxidize iron (II) ions to
iron (III) ions and are reduced to chromium (III) ions in acidic solution. Write and balance the net ionic equation for the reaction.
OH 7Cr 2Fe 6H 14OCrFe 6
red. OH 7Cr 2 e 6H 14OCr
ox. 1e+ FeFe 6
rxnstart FeCrFeOCr
2+3+3+2
72+2
2+3-+2
72
-+3+2
+3+3+2272
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Half Reaction Method
Example: In basic solution hydrogen peroxide oxidizes chromite ions, Cr(OH)4
-, to chromate ions, CrO42-.
Write and balance the net ionic equation for this reaction.
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OH8CrO 2OH 3OH2Cr(OH) 2
6OH+OH8CrO 2OH 3OH8Cr(OH) 2
OH2e2OH3
3e+OH4CrOOH4Cr(OH)2
solution) (basic CrOOHCr(OH)
22
422-
4
-2
2422
-4
--22
-2
24
-4
24224
Half Reaction Method
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Half Reaction Method
Example: When chlorine is bubbled into basic solution, it forms chlorate ions and chloride ions. Write and balance the net ionic equation.
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Cl 5+ OH 3ClO OH 6Cl 3
Cl 10+ OH 6ClO 2OH 12Cl 6
Cl 2e 2Cl 5
10e + OH 6ClO 2OH 12Cl 1
solution) (basic ClClOCl
23-
2
23-
2
-2
-23
-2
32
Half Reaction Method