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Redox Reactions

Redox Reactions. Redox is short for “ oxidation-reduction .” Plating and combustion are redox reactions. Consider the following redox reaction: Fe (s) + Ni(NO 3 ) 2 (aq)  Fe(NO 3 ) 2 (aq) + Ni (s). O xidation I s L oss of electrons R eduction I s G ain of electrons.

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Redox Reactions

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  1. Redox Reactions • Redox is short for “oxidation-reduction.” • Plating and combustion are redox reactions. • Consider the following redox reaction: • Fe(s) + Ni(NO3)2 (aq) Fe(NO3)2(aq)+ Ni (s) Oxidation Is Loss of electrons Reduction Is Gain of electrons • Fe goes from an element to an ion: Fe  Fe2+ • Fe got oxidized • Ni goes from an ion to an element: Ni2+  Ni • Ni2+ got reduced

  2. Oxidation Numbers • We determine whether a chemical species has been oxidized or reduced by looking at the change in its oxidation number. • Fe(s) + Ni(NO3)2 (aq) Fe(NO3)2(aq)+ Ni (s) • ox # 0 +2 +2 0 • The oxidation number of an element is 0. • The oxidation number of a monatomic ion is its charge. • The ox # for Fe went up. It was oxidized. • The ox # for Ni2+ went down. It was reduced (reduced = down). Oxidation Is Loss of electrons Reduction Is Gain of electrons

  3. Oxidation Numbers - Rules • The oxidation number of an element is 0. • The oxidation number of a monatomic ion is its charge. • The oxidation number of oxygen in a compound is usually -2 (exception is peroxide, O22-, where the oxidation number is -1). • The oxidation number of hydrogen in a compound is 1 when bonded to nonmetals and -1 when bonded to metals. • The oxidation number of fluoride is always -1. • The sum of the oxidation numbers in a neutral compound is zero. • The sum of the oxidation numbers in a polyatomic ion is the charge of the ion.

  4. Oxidation Numbers - Examples Species oxidation Species oxidation number number Fe 0 Na+ +1 H2 0 H+ +1 O2 0 O2- -2 Al3+ +3 Cl- -1 Fe2+ +2 N3- -3 S in SO42- : overall charge of ion is -2 = ox # of S + 4(ox # of O) -2 = ox # of S + 4(-2) -2 = ox # of S – 8 -2+8 = 6 = ox # of S Cr in Cr2O72- +6 N in NO3- +5 S in SO32- +4 N in NO2- +3 P in PO43- +5 C in CO32- +4

  5. Redox Reactions To see if a redox reaction has occurred, check to see if the oxidation numbers of any element involved in the reaction have changed. Mg(s) + Fe(NO3)2 (aq) Mg(NO3)2(aq)+ Fe (s) ox #0 +2 +5 -2 +2 +5 -2 0 • Mg goes from an element to an ion: Mg  Mg2+ • Mg got oxidized. • Fe2+ was the oxidizing agent. • Fe goes from an ion to an element: Fe2+  Fe • Fe2+ got reduced. • Mg was the reducing agent. Oxidation Is Loss of electrons Reduction Is Gain of electrons

  6. Redox Reactions - Example Molecular equation: Mg(s) + Fe(NO3)2 (aq) Mg(NO3)2(aq)+ Fe (s) Total Ionic equation: Mg(s) + Fe2+(aq) + 2NO3-(aq) Mg2+(aq)+ Fe(s) + 2NO3-(aq) Net Ionic equation: Mg(s) + Fe2+ (aq) Mg2+(aq)+ Fe (s) Mg got oxidized. Fe2+ was the oxidizing agent. Fe2+ got reduced. Mg was the reducing agent.

  7. Redox Reactions - Example When Zn dissolves in sulfuric acid, the molecular equation is Zn(s) + H2SO4 (aq) ZnSO4(aq)+ H2 (g), the total ionic equation is Zn(s) + 2H+(aq) + SO42-(aq) Zn2+(aq) + SO42-(aq) + H2 (g), and the net ionic equation is: Zn(s) + 2H+(aq) Zn2+(aq) + H2 (g) Zn got oxidized. H+ was the oxidizing agent. H+ got reduced. Zn was the reducing agent.

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