Redox Reactions
Sections 4.9, 18.2
RW Session ID = MSTCHEM1
Oxidation‐Reduction Reactions
Redox Reactions
• Electrons are transferred from one reactant to another
– Oxidation – loss of electrons
– Reduction – gain of electrons
OIL RIG
Oxidation is
Is
Loss of electrons
LEO GER
Reduction
Is
Gain of electrons
Loss of
Electrons is
Oxidation
Gain of
Electrons is
Reduction
Rules for Assigning Oxidation States
• Free element (neutral and not bonded to another element) → 0
• Monoatomic ion → charge
• In compounds, metals have positive oxidation states:
– Group 1A metals → +1
– Group 2A metals → +2
• In compounds:
Ionic compounds can be broken into the corresponding ions and treated separately
– Fluorine → ‐1
– Hydrogen → +1 (when bonded to nonmetals) or ‐1 (when bonded to metals)
– Oxygen → ‐2 (except w/fluorine or bonded to itself)
• The sum of the oxidation states:
– Neutral molecule → 0
– Polyatomic ion → charge
These rules give you a starting place for assigning oxidation states. You should start at the top and work your way down. All other oxidation states can be determined by knowing that the sum of oxidation states equals the charge on the substance.
Oxidation States
0
K
0
Ca
0
O2
0
Cl2
0
P4
+1
K+
+2
Ca2+
‐2
O2‐
‐1
Clˉ
‐3
P3‐
+2
Hg2+
+1
Hg22+
+3 ‐2
NO2ˉ
+6 ‐2
SO42‐
+6 ‐2
Cr2O72‐
Oxidation States
+1 ‐1
LiBr
+2 +6 ‐2
BaSO4
+2 ‐2
NO
+2 ‐2
CaS
+1 +3 ‐2
Na3PO3
+4 ‐2
NO2
+2 ‐2
CuO
+1 ‐2 +2 ‐3
Cu2O
Mg3N2
+2 ‐2 +1
Fe(OH)2
+3 ‐1
BrF3
‐2 +1
P2H4
+2 +3 ‐2
Sn(BrO2)2
+1 ‐2
S 2O
Redox Reactions
• Oxidation and Reduction
– Loss of electrons and gain of electrons
– Change in oxidation states from reactant to product
• Oxidizing agent (oxidant, oxidizer)
– A substance that causes the oxidation of another substance.
– gains electrons, is reduced
• Reducing agent (reductant, reducer)
– A substance that causes the reduction of another substance
– loses electrons, is oxidized
Determine whether each of the following reactions are redox reactions. If so, identify the oxidant and reductant.
+3
+1
0
+2
Al3+(aq) + Cu+(aq) → Al(s) + Cu2+(aq)
reduction
Cu+ ‐ reductant
Al3+ ‐ oxidant
oxidation
+1 ‐1
+2 ‐2 +1
+2 ‐1
+1 ‐2
HCl(aq) + Ca(OH)2(aq) → CaCl2(aq) + H2O(l)
not a redox reaction
Determine whether each of the following reactions are redox reactions. If so, identify the oxidant and reductant.
0
0
+2 ‐2
Ca(s) + O2(g) → CaO(s)
Ca ‐ reductant
O2 ‐ oxidant
oxidation
reduction
+1 ‐1
+1 +5 ‐2
+1 +5 ‐2
+1 ‐1
NaCl(aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)
not a redox reaction
Determine whether each of the following reactions are redox reactions. If so, identify the oxidant and reductant.
0
+1 ‐1
+1 ‐1
0
NaCl(aq) + F2(g) → NaF(aq) + Cl2(g)
oxidation
NaCl ‐ reductant
F2 ‐ oxidant
reduction
‐2 +1
0
+4 ‐2
+1 ‐2
C3H6(l) + O2(g) → CO2(g) + H2O(g)
oxidation
reduction
C3H6 ‐ reductant
O2 ‐ oxidant
Determine whether each of the following reactions are redox reactions. If so, identify the oxidant and reductant.
+6
+2
‐2
+3
+3
Fe2+(aq) + Cr2O72‐(aq) → Fe3+(aq) + Cr3+(aq)
oxidation
reduction
‐3 +1
+1 ‐1
Fe2+ ‐ reductant
Cr2O72‐ ‐ oxidant
‐3 +1 ‐1
NH3(aq) + HCl(aq) → NH4Cl(aq)
not a redox reaction
Section 18.2 Balancing Oxidation‐Reduction Equations
Half‐Reaction Method of Balancing Redox Reactions
• Assign oxidation states to all atoms
• Separate reaction into oxidation and reduction half‐reactions
• Balance each half‐reaction by mass
– Balance elements other than oxygen and hydrogen
– Balance oxygen with water
– Balance hydrogen with hydrogen ion
• Balance each half‐reaction by charge using electrons
• Multiply each reaction by a factor to make electrons lost and gained equal
• Add the two reactions together and cancel out electrons and any other species necessary
• For basic solution add hydroxide ion to each side (enough to fully react with hydrogen ion and form water)
Al3+(aq) + Cu+(aq) → Al(s) + Cu2+(aq)
Oxidation Half‐reaction
Cu+(aq) → Cu2+(aq)
Cu+(aq) → Cu2+(aq) + 1 eˉ
[ Cu+(aq) → Cu2+(aq) + 1 eˉ ] x 3
Reduction Half‐reaction
Al3+(aq) → Al(s)
Al3+(aq) + 3 eˉ → Al(s)
Complete balanced reaction
Al3+(aq) + 3 Cu+(aq) → Al(s) + 3 Cu2+(aq)
Fe2+(aq) + Cr2O72‐(aq) → Fe3+(aq) + Cr3+(aq)
Oxidation Half‐reaction
Fe2+(aq) → Fe3+(aq)
Fe2+(aq) → Fe3+(aq) + 1 eˉ
[ Fe2+(aq) → Fe3+(aq) + 1 eˉ ] x 6 Reduction Half‐reaction
Cr2O72‐(aq) → 2 Cr3+(aq)
Cr2O72‐(aq) → 2 Cr3+(aq) + 7 H2O(l)
Cr2O72‐(aq) + 14 H+(aq) → 2 Cr3+(aq) + 7 H2O(l)
Cr2O72‐(aq) + 14 H+(aq) + 6 eˉ → 2 Cr3+(aq) + 7 H2O(l)
Complete balanced reaction (in acidic solution)
6 Fe2+(aq) + Cr2O72‐(aq) + 14 H+(aq)
→ 6 Fe3+(aq) + 2 Cr3+(aq) + 7 H2O(l)
Fe2+(aq) + Cr2O72‐(aq) → Fe3+(aq) + Cr3+(aq)
Complete balanced reaction (in acidic solution)
6 Fe2+(aq) + Cr2O72‐(aq) + 14 H+(aq)
→ 6 Fe3+(aq) + 2 Cr3+(aq) + 7 H2O(l)
Complete balanced reaction (in basic solution)
6 Fe2+(aq) + Cr2O72‐(aq) + 14 H+(aq) + 14 OHˉ(aq)
→ 6 Fe3+(aq) + 2 Cr3+(aq) + 7 H2O(l) + 14 OHˉ(aq)
6 Fe2+(aq) + Cr2O72‐(aq) + 14 H2O(l)
→ 6 Fe3+(aq) + 2 Cr3+(aq) + 7 H2O(l) + 14 OHˉ(aq)
6 Fe2+(aq) + Cr2O72‐(aq) + 7 H2O(l)
→ 6 Fe3+(aq) + 2 Cr3+(aq) + 14 OHˉ(aq)
MnO4‐(aq) + Bi3+(aq) → Mn2+(aq) + BiO3‐(aq)
Oxidation Half‐reaction
Bi3+(aq) → BiO3‐(aq)
3 H2O(l) + Bi3+(aq) → BiO3‐(aq) + 6 H+(aq) + 2 eˉ
[ 3 H2O(l) + Bi3+(aq) → BiO3‐(aq) + 6 H+(aq) + 2 eˉ ] x 5 Reduction Half‐reaction
MnO4‐(aq) → Mn2+(aq)
MnO4‐(aq) → Mn2+(aq) + 4 H2O(l)
MnO4‐(aq) + 8 H+(aq) → Mn2+(aq) + 4 H2O(l)
[ MnO4‐(aq) + 8 H+(aq) + 5 eˉ → Mn2+(aq) + 4 H2O(l) ] x 2
Complete balanced reaction
Acidic:
5 Bi3+(aq) + 2 MnO4‐(aq) + 7 H2O(l) → 5 BiO3‐(aq) + 2 Mn2+(aq) + 14 H+(aq)
+ 14 OH‐(aq) →
+ 14 OH‐(aq)
→
+ 14 H2O(l)
Basic:
5 Bi3+(aq) + 2 MnO4‐(aq) + 14 OH‐(aq) → 5 BiO3‐(aq) + 2 Mn2+(aq) + 7 H2O(l)