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Oxidation and Reduction
Chapter 20
Types of Chemical Reactions
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Type I: ions or molecules react with no apparent change
in the electronic structure of the particles.
Type II: ions or atoms undergo changes of electronic
structure. Electrons may be transferred from one particle
to another. On the other hand, the sharing of the
electrons may be somewhat changed.
Type II reactions involving electron changes are called
oxidation-reduction reactions.
It is these "redox" reactions which we will now discuss.
Before we indicate what oxidation-reduction reactions
are, we will briefly indicate what they are not.
What Redox is NOT
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In the BaSO4 reaction in Table 26-1, the substances are all ionic.
Since there is no change in the charge of these ions in the reaction, there
are no electron changes.
This reaction is not an oxidation-reduction reaction.
The production of a (BaS04) is nearly always a result of a non-redox
reaction.
Most acid-base reactions are also the non-redox type.
Since nearly every other kind of reaction is an oxidation-reduction reaction,
redox reactions are important in the laboratory.
They are also important in life processes and in industry.
Oxidation
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The term oxidation was first applied to the combining of oxygen
with other elements.
There were many known instances of this behavior:
– Iron rusts
– Carbon burns
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In rusting, oxygen combines slowly with iron to form Fe2O3.
In burning, oxygen unites rapidly with carbon to form CO2.
Observation of these reactions gave rise to the terms "slow" and
"rapid" oxidation.
Chemists recognize, however, that other nonmetallic elements unite
with substances in a manner similar to that of oxygen.
– Hydrogen, antimony, and sodium all burn in chlorine, and iron will burn
in fluorine.
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Since these reactions were similar, chemists formed a more general
definition of oxidation:
– Electrons were removed from each free element by the reactants O2 or
Cl2.
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Thus oxidation is defined as the process by which electrons
are apparently removed from an atom or ion.
Reduction
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A reduction reaction was originally limited to the type of
reaction in which ores were "reduced" from their oxides.
– Iron oxide was "reduced" to iron by carbon monoxide.
– Copper(II) oxide could be "reduced" to copper by hydrogen.
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In these reactions, oxygen is removed, and the free
element is produced.
The free element can be produced in other ways:
– An iron nail dropped into a copper(II) sulfate solution causes a
reaction which produces free copper.
– An electric current passing through molten sodium chloride
produces free sodium.
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The similarity between oxidation and reduction reactions
led chemists to formulate a more generalized definition
of reduction.
By definition, reduction is the process by which
electrons are apparently added to atoms or ions.
OIL RIG—the Texas Definition
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Oxidation is Loss (of electrons), Reduction
is Gain (of electrons)
Oxidizing and Reducing Agents
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In an oxidation-reduction reaction, electrons are transferred.
–
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All the electrons exchanged in an oxidation-reduction reaction must be accounted for.
It seems reasonable, therefore, that both oxidation and reduction must occur
at the same time in a reaction.
Electrons are lost and gained at the same time and the number lost must
equal the number gained.
The substance in the reaction which gives up electrons is called the
reducing agent. The reducing agent contains the atoms which are oxidized
(the atoms which lose electrons).
–
Zinc is a good example of a reducing agent. It is oxidized to the zinc ion, Zn2+
–
Dichromate ion, Cr2072-, is a good example of an oxidizing agent. It is reduced to the
chromium ion, Cr3+
The substance in the reaction which gains electrons is called the oxidizing
agent. It contains the atoms which are reduced (the atoms which gain
electrons).
If a substance gives up electrons readily, it is said to be a strong reducing agent. Its
oxidized form, however, is normally a poor oxidizing agent.
If a substance gains electrons readily, it is said to be a strong oxidizing agent. Its
reduced form is a weak reducing agent.
Redox of nails and copper
Oxidation Numbers
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How is it possible to determine whether an oxidation-reduction reaction has
taken place?
We do so by determining whether any electron shifts have taken place
during the reaction.
To indicate electron changes, we look at the oxidation numbers of the
atoms in the reaction.
The oxidation number is the charge an atom appears to have when
we assign a certain number of electrons to given atoms or ions.
Any change of oxidation numbers in the course of a reaction
indicates an oxidation-reduction reaction has taken place.
Oxidation numbers are assigned according to the apparent charge
of the element (aka, valence!)
For example, suppose iron, as a reactant in a reaction, has an oxidation
number of 2+.
– If iron appears as a product with an oxidation number other than 2+, say 3+, or
0, then a redox reaction has taken place.
Determining Oxidation Numbers
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For all compounds, whether covalent, polar covalent, or ionic, we
treat as ionic for counting electrons and for oxidation-reduction
reactions.
Rule 1: Sum of the oxidation numbers of all the atoms in the
chemical species equals the charge on the species.
Neutral compounds: Sum of oxidation numbers = 0
Ionic species: Sum of oxidation numbers = charge of the ion
Rule 2: In Binary Compounds, the more Electronegative (EN)
element is assigned to have a negative oxidation number. (See EN
trends.)
Rule 3: Atoms may have only certain oxidation numbers. The range
is:
Maximum oxidation number possible = + Group number.
Minimum oxidation number possible = (Group number - 8) (this
number will be negative)
Determining Oxidation Numbers
(Cont)
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Atoms which will have known oxidation numbers are:
Atoms as Elements: Ex. H2, O2, P4, Fe
Oxidation number = 0
Monoatomic Ions:
Cations: Ex. Na+, Al3+ (main group metals)
Oxidation number = + Group Number
Anions: Cl-, O2Oxidation number = Group Number - 8
Hydrogen
Combined with Nonmetals: Ex. NH3, H2O, HCl
Oxidation number = +1
Combined with Metals: Ex. NaH, CaH2 (hydrides)
Oxidation number = -1
Oxygen (Unless O22-, peroxide)
Oxidation number = -2
CO: (Sum will equal 0 since it is a neutral molecule)
O will have a -2 ox. number.
1C+1O=0
(C?) + (-2) = 0
C? = +2
 Oxidation number of C in CO is +2
Oxidation number of O in CO is -2 (known)
 Check ox. number to see if it falls within range:
+2 is in between the maximum value of C, +4,
(Gr#) and the minimum value of C, - 4, (Gr# 8).
So okay.
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Cr2O72-: (Sum of all oxidation numbers will equal -2 since
it is an ion.)
2 Cr + 7 O = -2
2(Cr?) + 7(-2) = -2
2(Cr?) + (-14) = -2
2(Cr?) = +12
Cr? = +6
 Oxidation number of each Cr in Cr2O72- is +6
Oxidation number of each O in Cr2O72- is -2
(known)
 Check ox. number to see if it falls within range:
+6 is the maximum value that Cr can have
(Gr#). So okay.
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CS2: (Sum will equal 0 since it is a neutral molecule)
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C will have the positive oxidation number since it is less
EN than S
S will have a -2 charge since it is Gr # 6, (6 - 8 = -2)
C+2S=0
(C?) + 2 (-2) = 0
(C?) + (-4) = 0
C? = +4
Oxidation number of C in CS2 is +4
Oxidation number of each S in CS2 is -2 (known)
Check ox. number to see if it falls within range:
+4 is the maximum value that C can have, (Gr#). So
okay.
HNO3(aq) + H3AsO3(aq)  NO(g) + H3AsO4(aq) + H2O(l)
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Step #1: Try to balance the atoms by inspection.
The H and O atoms are difficult to balance in this equation. You might arrive at the
correct balanced equation using a “trial and error” technique, but if you do not
discover the correct coefficients fairly quickly, proceed to Step #2.
Step #2: Is the reaction redox?
The N atoms change from +5 to +2, so they are reduced. This information is enough
to tell us that the reaction is redox. (The As atoms, which change from +3 to +5, are
oxidized.)
Step #3: Determine the net increase in oxidation number for the element that is
oxidized and the net decrease in oxidation number for the element that is reduced.
As +3 to +5
Net Change = +2
N +5 to +2
Net Change = -3
Step #4: Determine a ratio of oxidized to reduced atoms that would yield a net
increase in oxidation number equal to the net decrease in oxidation number.
As atoms would yield a net increase in oxidation number of +6. (Six electrons would
be lost by three arsenic atoms.) 2 N atoms would yield a net decrease of -6. (Two
nitrogen atoms would gain six electrons.) Thus the ratio of As atoms to N atoms is
3:2.
Step #5: To get the ratio identified in Step 5, add coefficients to the formulas which
contain the elements whose oxidation number is changing.
2HNO3(aq) + 3H3AsO3(aq)  NO(g) + H3AsO4(aq) + H2O(l)
Step #6: Balance the rest of the equation by inspection.
2HNO3(aq) + 3H3AsO3(aq)  2NO(g) + 3H3AsO4(aq) + H2O(l)
Cu(s) + HNO3(aq)  Cu(NO3)2(aq) + NO(g) + H2O(l)
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The nitrogen atoms and the oxygen atoms are difficult to balance by
inspection, so we will go to Step #3.
The copper atoms and some of the nitrogen atoms change their oxidation
numbers. These changes indicate that this reaction is a redox reaction. We
next determine the changes in oxidation number for the atoms oxidized and
reduced.
Cu 0 to +2
Net Change = +2
Some N +5 to +2
Net Change = -3
We need three Cu atoms (net change of +6) for every 2 nitrogen atoms
that change (net change of -6). Although the numbers for the ratio
determined in Step #5 are usually put in front of reactant formulas, this
equation is somewhat different. Because some of the nitrogen atoms are
changing and some are not, we need to be careful to put the 2 in front of a
formula in which all of the nitrogen atoms are changing or have changed.
We therefore place the 2 in front of the NO(g) on the product side. The 3
for the copper atoms can be placed in front of the Cu(s).
3Cu(s) + HNO3(aq)  Cu(NO3)2(aq) + 2NO(g) + H2O(l)
We balance the rest of the atoms, being careful to keep the ratio of Cu to
NO 3:2.
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)
Summary
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1. An oxidation-reduction reaction involves an apparent transfer of electrons from one
particle to another.
2. Oxidation is the process by which electrons are apparently removed from an atom or
group of atoms.
3. Reduction is the process by which electrons are apparently added to atoms or groups
of atoms.
3. Any substance in a reaction which loses electrons is a reducing agent.
4. Any substance in a reaction which gains electrons is an oxidizing agent.
5. If a substance gives up electrons readily, it is a strong reducing agent. Its oxidized
form is usually a poor oxidizing agent.
6. If a substance acquires electrons readily, it is a strong oxidizing agent. Its reduced
form is usually a poor reducing agent.
7. Oxidation number is the charge an atom appears to have when we assign a certain
number of electrons to that atom.
8. Six rules for assigning oxidation numbers:
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a. The oxidation number of any free element is O.
b. The oxidation number of any single-atom ion is equal to the that ion.
c. The oxidation number of hydrogen is usually 1+.
d. The oxidation number of oxygen in most compounds is 2-.
e. The sum of the oxidation numbers of all the atoms in a particle equal the apparent charge of
that particle.
f. In compounds, elements of Group IA and Group IIA have an oxidation number numerically
equal to their group in the periodic table.
9. In all chemical reactions, charge, number and kind of atoms, and number of electrons
are conserved. Knowing these quantities, you can do a redox equation.
10. Redox reactions are more easily balanced by splitting the equation into half-reactions.