Oxidation-Reduction Reactions
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Carbonate reactions are acid-base reactions:
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Transfer of protons – H+
Other acid-base systems are similar:
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Sulfuric acid - H2SO4
Phosphoric acid - H2PO3
Nitric Acid HNO3
Oxidation-Reduction Reactions
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Redox reactions are analogous, but are
transfer of electrons rather than protons
Very important class of reactions
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Elements may have variable charges –
number of electrons (valence state)
Valence state controls speciation of elements

Examples of primary valence states of
some elements
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C = +4 or -4
S +6 or -2
N +5 or +3, also +4, +2
Fe +3 or +2
Mn +3 or +2, also +7, +6, +4
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Minor elements also have various valence
states
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V, Cr, As, Mo, V, Se, Sb, W, Cu…
All nasty elements
Important environmental controls – e.g.,
mining

Valence state very important for mobility,
as well as absorption and thus toxicity
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Fe3+ (oxidized) is highly insoluble


Precipitate as Fe-oxide minerals (magnetite,
hematite, goethite, lepidocrocite, limonite)
Fe2+ (reduced) much more soluble – most Fe
in solution is +2 valence

Common precipitates are Fe-sulfides (pyrite,
marcasite)
Assignment of oxidation state

Valence state of oxygen is always -2
except for peroxides, where it is -1.
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E.g., H2O2 and Na2O2
Valence state of hydrogen is +1 in all
compounds except when bonded with
metals where it is -1.
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NaH
NaBH4
LiAlH4
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Valence state of all other elements are
selected to make the compound neutral
Certain elements almost always have the
same oxidation state

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Alkali metals = +1 (left most column)
Alkaline earths = +2 (second column from
left)
Halogens = -1 (2nd column from right)
Examples

What are the oxidation states of N in NO3and NO2-?

3O2- + Nx = NO3-
6- + x = -1
N = +5

2O2- + Nx = NO2-
N = +3
4- + x = -1

What are the oxidation states of H2S and
SO42-?

2H+ + Sx = H2S
2+ + x = 0
S = -2

4O2- + Sx = SO42-
S = +6
8- + x = -2
Oxidation Reactions

Oxidation can be thought of as involving
molecular oxygen

3Fe2O3
(hematite)
All as Fe3+
2Fe3O4 + 1/2O2
(magnetite)
One as Fe2+ + two as Fe3+
High O/Fe ratio
Oxidized
Lower O/Fe ratio
Reduced
In this case, the generation of molecular
oxygen controls the charge imbalance

Also possible to write these reactions in
terms of electrons:
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3Fe2O3 + 2H+ + 2e2Fe3O4+ H2O
LEO – lose electron oxidation – the Fe3+ is
oxidized
GER – gain electron reduction – the Fe2+ is
reduced
OIL – oxidation is loss
RIG – reduction is gain
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Generally easiest to consider reactions as
transfer of electrons
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Many redox reaction do not involve molecular
oxygen
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Problem is that free electrons are not
really defined
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Reactions that consume “free electrons”
represent only half of the reaction
A complementary reaction required to
produce a “free electron”
Concept is two “half reactions”
The half reaction simultaneously create and
consume electrons, so typically not expressed
in reaction
Half Reactions
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Example of redox reaction without
oxygen:
Zn(s) + Cu2+(aq)

Cu(s) + Zn2+(aq)
Here Zn solid releases electron, which is
consumed by dissolved Cu2+.
Physical model of process
Ammeter
eecations
Dissolves
Salt bridge – keeps
charge balance in
solution.
anions
Precipitates
Increases
Decreases

Ammeter shows flow of electrons from Zn
to Cu:
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Zn rod dissolves – Zn2+ increases
Cu rod precipitates – Cu2+ decreases
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At the rod, the reactions are:
Zn = Zn2+(aq) + 2e2e- + Cu2+(aq) = Cu
Zn + Cu2+(aq) = Zn2+(aq) + Cu
Half
reactions
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Benefits of using half reactions:
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Half reactions help balance redox reactions
Used to create framework to compare
strengths of oxidizing and reducing agents
Rules for writing and balancing
half reactions
1.
2.
3.
4.
Identify species being oxidized and
reduced
Write separate half reactions for
oxidation and reduction
Balance reactions using (1) atoms and
(2) electrical charge by adding e- or H+
Combine half reactions to form net
oxidation-reduction reactions
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Consider reaction
H2O2 + I-

I2 + H 2O
First, ID oxidized and reduced species:
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Iodine is being oxidized from -1 to 0 charge
Oxygen in peroxide is being reduced to water
IH2O2
I2
H2O
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Next – balance elements (oxidation half
reaction:
2I-

I2
And charge:
2I-
I2 + 2e-
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Balance reduction half reaction

First balance oxygen, then add H+ to balance
hydrogen, then add electrons for electrical
neutrality:
H2O2
H2O
H2O2
2H2O
2H+ + H2O2
2H2O
2e- + 2H+ + H2O2
2H2O
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Combine two half reactions to get net
reactions:
2I2e- + 2H+ + H2O2
2H+ + 2I- + H2O2
I2 + 2e2H2O
2H2O + I2
Flow of electrons – Oxygen is electron acceptor,
reduced; I- is electron donor, oxidized
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Common reaction in natural waters is
reduction of Fe3+ by organic carbon
4Fe3+ + C + 2H2O

4Fe2+ + CO2 + 4H+
With half reactions:
4Fe3+ + 4e-
4Fe2+
C + 2H2O
CO2 + 4H+ + 4e-
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From thermodynamic conventions, its
impossible to consider a single half
reaction
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There is no thermodynamic data for e-
Practically, half reactions are defined
relative to a standard
The standard is the “Standard Hydrogen
Electrode (SHE)”
SHE
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Platinum electrode in solution containing
H2 gas at P = 1 Atm.
Assign arbitrary values to quantities that
can’t be measured
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Difference in electrical potential between
metal electrode and solution is zero
DGfº of H+ = 0
DGfº of e- = 0
SHE
Half reaction in solution:
H+ + e- = 1/2H2(g)
By definition,
aH+ = 1
Allows electrons
to flow but
chemically inert
Example of how SHE used
E = Potential
Positive or
negative
Fe3+ + e- = Fe2+
If reaction goes to left, wire removes electrons
If reaction goes to right, wire adds electrons
SHE:
H+ + e- = 1/2H2(g)
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In cell A, platinum wire is inert – transfers
electrons to or from solution only.
If wire has no source of electrons
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Define the potential as “activity of
electrons” = ae
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Pt wire develop an electrical potential –
“tendency” for electrons to enter or leave
solution
Not a true activity, really a “tendency”
Define pe = -logae-, similar to pH
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In Cell A solution, Fe is both oxidized and
reduced
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Fe2+ and Fe3+
Reaction is:
Fe3+ + e- = Fe2+
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If reaction goes to left, Fe2+ gives up eIf reaction goes to right, Fe3+ acquires eIf no source or sink of e-, (switch open), volt
meter measures the potential (tendency)
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Since we have a reaction
Fe3+ + e- = Fe2+
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can write an equilibrium constant
Keq =
aFe2+
aFe3+ ae-
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Rearranged:
ae-= Keq-1
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aFe2+
aFe3+
ae- is proportional to the ratio of activity of
the reduced species to activity of oxidized
species
ae- is electrical potential (in volts) caused
by ratio of reduced to oxidized species
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Consider half cell B:
H+ + e- = 1/2H2(g)
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Direction of reaction depends on tendency
for wire to gain or lose electrons
Equilibrium constant
KSHE =
PH21/2
aH+ ae-
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Switch closed – electrons flow from one
half cell to the other

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Electron flow from the side with the highest
activity of electrons to side with lowest
activities
Flow of electrons
Overall reaction:
Fe3+ + 1/2H2(g) =Fe2+ + H+

Direction of reaction depends on which half
cell has highest activity of electrons

Switch open:
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No longer transfer of electrons
Now simply potential (E) generated at Pt wire
By convention, potential of SHE (ESHE) = O
Potential called Eh, i.e. E (electromotive force)
measured relative to SHE (thus the “h”)
Eh > or < O depends on whether ae- is > or <
that of SHE
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Convention
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Eh > 0 if ae- of the half cell < SHE
I.e. if electrons flow from the SHE to the fluid
For thermodynamics:

Fe3+ + 1/2H2(g) =Fe2+ + H+
Is equivalent to:
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Fe3+ + e- =Fe2+
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Expressions for activities of electrons:
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Eh or pe
pe = [F/(2.303RT)]*Eh
@ 25ºC, pe = 16.9 Eh; Eh = 0.059pe
F = Faraday’s constant = 96,485 coul/mol
Coulomb = charge /electron = quantity of
electricity transferred by 1 Amp in 1
second.