Chapter 20 - Oxidation-Reduction Reactions

Chemistry Regents
Mr. Markic
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Chapter 20 - Oxidation-Reduction Reactions
Aka Redox reactions
Composed of:
• Oxidation reaction
• Reduction reaction
Oxidation Reactions
Combination of a substance with oxygen
• Iron is being oxidized
• Complete or partial loss of electrons
Mg + S  Mg2+ + S2Magnesium is being oxidized
As magnesium loses electrons, what happens to sulfur?
What happens to the charge on sulfur?
Reduction Reactions
1. Loss of oxygen from a substance
2. Complete or partial gain of electrons
Mg + S  Mg2+ + S2Sulfur is being reduced
Redox Reactions
 When one atom loses e-…
 Another atoms gains e Can never occur alone!
Mnemonic Devices
 LEO the lion says GER
o LEO = Loss of Electrons is Oxidation
o GER = Gain of Electrons is Reduction
o OIL = Oxidation Is Loss
o RIG = Reduction Is Gain
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Mr. Markic
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Sample Exercises
Identify the following processes as either oxidation or reduction:
a) Al3+ + 3e- → Al
d) Sr → Sr2+ + 2e-
b) S2- → S + 2e-
e) 2I- → I2 + 2e-
c) 2Cl- → Cl2 + 2e-
f) Zn2+ + 2e- → Zn
Assigning Oxidation Numbers
 Oxidation Number - positive or negative number assigned to an atom to indicate its
degree of oxidation or reduction
Oxidation Rules
1. Each uncombined element has an oxidation number of zero.
Ex. Potassium metal K =
Nitrogen gas
N2 =
2. Monatomic ions have an oxidation number equal to their ionic charge.
Ex. Be1- =
Fe3+ =
3. Hydrogen is +1 in compounds unless it is combined with a metal, in which case it is -1.
Ex. H in HCl =
H in LiH =
4. Oxygen is usually -2 in compounds, except in peroxides, where it is -1, and with fluorine
(which is more electronegative), where it is +2.
Ex. O in H2O =
O in H2O2 =
O in OF2 =
5. The metals of Group 1 are always +1 in compounds, and the metals of Group 2 are always
+2 in compounds.
Ex. Na in NaCl =
Mg in MgBr2 =
6. Fluorine is always -1 in compounds. The ‘other’ halogens are also -1 when they are the
most electronegative element.
Ex. F in HF =
Cl in HCl =
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7. For a polyatomic ion, the sum of the oxidation numbers must equal the charge on the
Ex. NO3 =
PO4 =
8. The sum of the oxidation numbers in all compounds must equal zero.
Ex. HCl =
Mg(NO3)2 =
Sample Exercises
What is the oxidation number of each kind of atom in the following ions and compounds?
a) SO2
b) CO32-
c) Na2SO4
d) (NH4)2S
e) Al(NO3)3
f) Na4SiO4
g) Cr2O72-
h) H3PO4
Determine what is oxidized and what is reduced in each reaction:
a) 2Na(s) + S(s) → Na2S(s)
2Na + S → 2Na + S
2Na -
b) 2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag(s)
2Ag + 2NO3 + Cu → Cu + 2NO3 + 2Ag
Cu -
2Ag+ -
c) 4Al(s) + 3O2(g) → 2Al2O3(s)
4Al + 6O → 4Al + 6O
4Al -
3O2 -
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Mr. Markic
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Redox Agents
 Reducing Agent o the substance that loses eo the substance that is oxidized
Oxidizing Agent –
o the substance that gains eo the substance that is reduced
Sample Exercises
Identify the reducing agent and the oxidizing agent for each reaction:
a) 2Na(s) + Br2(l) → 2NaBr(s)
b) Cl2(g) + S(s) → SCl2(g)
2Na + Br2 → 2Na + 2Br
Cl2 + S → S + 2Cl
2Na -
Br2 –
Cl2 –
2Na -
Br2 –
Cl2 –
c) Mg(s) + Cu(NO3)2(aq) → Mg(NO3)2(aq) + Cu(s)
Mg + Cu + 2NO3 → Mg + 2NO3 + Cu
Mg -
Cu -
Mg -
Cu –
d) MnO2(s) + 4HCl(aq) → MnCl2(aq) + Cl2(g) + 2H2O(l)
Mn + 2O + 4H + 4Cl → Mn + 2Cl + 2Cl + 4H + 2O
Cl - (Cl2)
Mn -
Cl -
Mn -
Identifying Redox Reactions
 Not all reactions are redox reactions
 1st assign oxidation numbers to each atom on both the reactant & product side
 If there is a change in oxidation number of an element, then the reaction is redox
Quick Tips:
o Redox: uncombined element appears on 1side of an equation, and is in a compound on
the other side
o Not Redox: double replacement, acid-base reaction
Sample Exercises
Assign oxidation numbers, to identify which of the following are redox reactions:
a) Mg(s) + Br2(l)  MgBr2(s)
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Mr. Markic
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b) H2CO3(aq)  H2O(l) + CO2(g)
c) Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(g)
d) 4NaOH(aq) + 2H2SO4(aq) → 2Na2SO4(aq) + 4H2O(l)
e) CuO(s) + H2(g)  Cu(s) + H2O(l)
• Shows either the oxidation or reduction portion of a redox reaction
• Shows the e- gained or lost
• Follows the Law of Conservation of Matter
• Charge is conserved (although it may not = 0)
Sample Exercises
Write the half reactions for each of the following redox reactions:
a) Cu + 2AgNO3 → Cu(NO3)2 + 2Ag
b) F2 + CaCl2 → CaF2 + Cl2
Oxidation Half:
Oxidation Half:
Reduction Half:
Reduction Half:
c) N2 + 3H2 → 2H3N
d) Zn + 2HCl → ZnCl2 + H2
Oxidation Half:
Oxidation Half:
Reduction Half:
Reduction Half:
Balancing Equations Using Half-Reactions
Cu + HNO3 → Cu(NO3)2 + NO + H2O
1. Rewrite the equation in ion form with oxidation #’s
Cu + HNO3 → Cu(NO3)2 + NO + H2O
2. Identify the elements that are oxidized & reduced
Chemistry Regents
Mr. Markic
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3. Write the oxidation & reduction half rxs
Oxidation Half:
Reduction Half:
4. Multiply each half-rx by an appropriate # so that the # of e- lost = # e- gained
3 • (Cu0 → Cu2+ +2e-) =
2 • (N5+ + 3e- → N2+) =
5. Add the two ‘unbalanced’ half-rx, cancel the e-’s
3Cu0 → 3Cu2+ + 6e2N5+ + 6e- → 2N2+
3Cu0 + 2N5+ → 3Cu2+ + 2N2+
6. Insert the new coefficients into the original equation
Exception: Do not insert a coefficient for any element that appears in more than one place in
the equation!
__Cu + __ HNO3 → __Cu(NO3)2 + __NO + __ H2O
7. Balance the rest of the equation by inspection
3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O
Sample Exercises
Balance the following redox equations using the half-reaction method:
a) Cu + HNO3 → Cu(NO3)2 + NO2 + H2O
b) HNO3 + I2 → HIO3 + NO2 + H2O
Chemistry Regents
Mr. Markic
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Chapter 21 – Electrochemistry
Spontaneous Redox Reactions
• When zinc metal is placed in a copper(II) sulfate solution, zinc becomes copper plated
• Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
• Movement of e- produces an electric current
Zn0(s)  Zn2+(aq) + 2e-
Cu2+(aq) + 2e-  Cu0(s)
Electrochemical Processes
• Any conversion between chemical energy and electrical energy
• Electrochemical cell – any device that converts chemical energy into electrical energy or
vise versa
• All involve redox reactions
o Must be a spontaneous redox reaction
o Table J
More active metal = more readily oxidized
Less active metal = more easily reduced
Sample Exercise
For each pair of metals listed below, decide which metal is most easily oxidized:
a) Ag, Cu
d) Sn, Ag
b) Ca, Al
e) Pb, Zn
c) Ni, Mg
f) Cu, Al
Voltaic Cells
• Electrochemical cells used to convert chemical
energy into electrical energy
• Electrical energy is produced by spontaneous redox
• Components:
o 2 half-cells
o Salt bridge
o Electrodes
Chemistry Regents
Mr. Markic
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Voltaic Cell – Components
• 2 Half-Cells
o Where oxidation or reduction occurs
o A piece of metal immersed in a solution of its ions
Salt Bridge
o Connects the 2 half-cells
o Allows for the migration of ions
o Tube containing a strong electrolyte (K2SO4)
o Recall… What are electrolytes?
Compounds that conduct an electric current when in aqueous solutions or the molten state
E.x. Ionic compounds, salts, acids & bases
Electrodes (anode & cathode)
o Conductor that carries electrons to or from a substance
o Metal rods
o Anode
o Cathode
 Where oxidation occurs
 Where reduction occurs
 Electrons are produced
 Electrons are consumed
 Gets a negative sign
 Gets a positive sign
 Metal becomes lighter
 Metal becomes heavier
Mnemonic Device
o REDuction occurs at the CAThode
o ANode is the site of OXidation
Non-Spontaneous Redox Reactions
• Electrolytic cell – absorbs electrical energy to make a
non-spontaneous redox reaction occur
Electrolysis – process in which electrical energy is used
to force a non-spontaneous chemical reaction to occur
Similarities Between Electrolytic & Voltaic Cells
• Both use redox reactions
• Anode is site of oxidation
• Cathode is site of reduction
• E- flow through a wire from anode to cathode
Differences Between Electrolytic & Voltaic Cells
• Electrolytic
o Absorbs electrical energy
o Uses a battery to produce a flow of
electrons to force a nonspontaneous reaction to occur
o Anode is positive
o Cathode is negative
o Creates electrical energy
o Uses spontaneous reaction to
produce a flow of electrons
o Anode is negative
o Cathode is positive