Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 121
The rate of reduction are found to be:
ClO4- < ClO3- < ClO2- < ClOand
ClO4- < SO42- < HPO42Thus the lower the oxidation state of the central atom, the faster the reaction is found to be. Why? Because the O-E bond is
strongest for the highest oxidation state, and this bond must be broken for the reaction to proceed. Further evidence in favor of
this hypothesis is the effect of size:
ClO4- < BrO4- < IO4The strength of the bond is reduced as the central atom size is increased.
7.4.4 Noncomplementary redox reactions
Very slow kinetics are observed when the change in oxidation states in the oxidizing and reducing agent are not the same,
e.g.:
2 Fe3+ + Tl+ à 2 Fe2+ + Tl3+
If this reaction is to proceed stepwise, then one Fe3+ reacts with one Tl+ to give a very unfavorable Tl2+ion. Alternatively the
reaction must be fully termolecular, requiring an activated complex containing two Fe3+ ions and a Tl+ ion. This is also very
unlikely, so the reaction tends to be slow.
7.5
Pourbaix diagrams
We now wish to discuss the fate of the metallic elements in aqueous solution in somewhat more detail. This ties in with a lot of
the work in the lab so far, and has important implications particularly for environmental chemistry.
7.5.1 Oxidation of the elements by water
The basic reaction we will consider is:
M + H2O à M + + ½ H2 + OHThis is the reaction involved in oxidation of metals by water, i.e. the process commonly known as rusting. The tables of
standard reduction potentials tells us for which metals this reaction goes in 1 M H3O+ . We must allow for the overpotential, ~0.6
V. We must also correct for pH if we want to discuss neutral solutions!
Consider Mg:
E½° = -2.37 V is the standard reduction potential in 1 M acid, i.e.
Mg + H2O à Mg + + OH- + 4 H2
E½° = +2.37
2+
0.591 [ H 2 ][ Mg ]
Now apply the Nernst equation:
E = E0 −
log
2
[ H + ]2
0.591
1
For pH 7 this becomes:
E = 2.37 −
log
= +1.956
2
(10− 7 ) 2
So here is a reaction that has enough voltage to overcome the overpotential, so long as a fresh surface is kept exposed to
the water such that the oxidation can proceed. This latter condition is what protects Al: a tough coating of Al2O3 protects bulk
aluminum from air and water oxidation. This is also why the Mg had to be put in boiling water in the lab for it to react at
appreciable rates. A good rule of thumb then is a metal with a reduction potential of ~ –{0.4 + 0.6} = –1V or greater will oxidize
in water at appreciable rates in the absence of air. NB: dissolved O2 will change this picture, of course! We all know that iron
rusts in aerated water.)
7.5.2 Reduction of elements by water
The redox reaction involved in acid solution is:
O2 + 4 H+ + 4 e- ⇔ 2 H2O E4 = +1.23 V
In basic solution it becomes:
O2 + 2 H2O + 4 e- ⇔ 4 OH- E4 = +0.40 V
Some strong oxidizing agents, e.g. Co 3+, are reduced by water. The overall reaction in acid is:
4 Co 3+ + 2 H2O à 4 Co 2+ + O2 + 4 H+ E = 0.59 V
This reaction is thus at the overpotential boundary. It becomes fully favored in basic solution, with E = 1.42 V
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 122
7.5.3 Stability field of water
We can combine the reduction and oxidation of water with the pH
With overpotential
+1.6
dependence and construct a diagram which represents the stability field
pH = 4
for water. This is shown in the diagram at right.
+1.2
The region boxed in the middle is the normal range found for
O 2/H 2O
natural waters, in which water is not oxidized or reduced, and those are
pH = 9
+0.8
the pH ranges found in the various types of natural waters. Specific
Fresh surface water
E/pH zones for different kinds of environments are indicated by the
+0.4
circles drawn into the stability field diagram.
Ocean water
Well aerated natural waters near the surface contain enough
Bog
(V)
Organic-rich
dissolved oxygen to get close to the reduction of water to O2. Eutrophic
0
water
lake water
Organic-rich
lake water contains sufficient dissolved organic matter to approach
waterlogged
closely the H+ reduction line. Ocean waters are relatively basic, and
-0.4
soils
Organic-rich
may be oxidizing if saturated in dioxygen, or reducing if saturated in
saline water
H 2O/H 2
organic matter (i.e. in stagnant lagoons, etc.)
-0.8
With
Fresh waters are considerably more acidic (because of dissolved
overpotential
carbon dioxide), but again they can be oxidizing if saturated in oxygen,
or reducing if too much organic matter is consuming all the oxygen.
This acidity is greatly enhanced in bogs and organic-laden soils, due to
2
4
6
8
10
high humic acid content. (Humic acids are complex organic acids
pH
occurring in the soil and in bituminous substances formed by the
The stability field for water
decomposition of dead vegetable matter.)
These are strongly reducing conditions, and explains the formation
of CH4 in marshes. Methane was first discovered from this source; Alessandra Volta was one of the first to identify this gas.
Ammonia and hydrogen sulfide as well as the inflammable phosphine, PH3, can also emanate from swamps. (There are many
rumors of eerie glows emanating in misty swamps; such tales are likely rooted in phosphine being released and burning above
the surface of the waters.) If you remember that bogs and marshes release such compounds, in other words very reduced
chemical compounds, it is should help you to remember the acidic, reducing character of bog water.
7.5.4 Pourbaix diagrams
The final type of diagram we want to consider for redox chemistry is a type of combination redox/pH predominance diagram
developed by the French electrochemistry Pourbaix. The diagrams are usually named after him. These diagrams are closely
related to the stability field for water just discussed. Indeed, most Pourbaix diagrams include either the main E/pH lines from the
water diagram, or both the main and overpotential lines. What these diagrams add are the relationships between the redox
activity and the Brønsted acidity of the elements. They are thus related both to Latimer diagrams , and to predominance
diagrams for acid/base reactions.
Consider as an example the Pourbaix diagram for
iron. Note that some of the lines from the stability field
of water are drawn into this diagram, since this diagram
applies to aqueous solutions of iron compounds.
•
•
•
•
•
•
The vertical axis plots the standard reduction
potential, and the horizontal the pH. Remember what
a predominance diagram was: simple vertical
boundaries where the most abundant species altered.
The bottom of the diagram refers to reduced species,
i.e. Fe(s), or to the type of conditions that lead to
reduction.
The top of the diagram to oxidized species and/or
oxidizing conditions.
Vertical lines indicate changes in acid base chemistry
independent of E°, e.g. Fe3+/Fe(OH)3.
Horizontal lines indicate redox changes unaffected by
pH, e.g. Fe/Fe2+ below pH 6.
In the more general case, the lines slope, since both E
and [H+ ] or [OH-] affect the redox process.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 123
To test your ability to read Pourbaix diagrams, see if you can find answers to the following:
(a) The form of iron which is the strongest oxidizing agent: FeO42- at [H+ ] ≤ 1 M
(b) The form of iron which is the strongest reducing agent: Fe(0), elemental iron
(c) The predominant form at pH = 7 and E = 0.0 V:
Fe(OH)3 predominates, but close to Fe2+/Fe(OH)2
23+
(d) E½° for (acid) reduction of FeO4 to Fe :
1 M, 0 pH = 0 so E½° = 2.2 V
(e) E½° for reduction of Fe2+ to Fe(s):
E½° = –0.5 V (This MUST be an acid process. Why?)
On the diagram, dashed lines d and e represent respectively the normal and overpotential for oxygen evolution according to:
2 H2O ⇔ 4 H+ (aq) + O2 + 4 eE° = +1.229 V.
The actual E for hydrogen evolution is given by f, while the overpotential is given by line g:
2H+ + 2e - ⇔ H2
E° = 0.00 V
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page 124
7.5.6 How elements behave in natural waters
We can now take any of the other Pourbaix diagrams, and compare them to the natural water limits, and predict what forms
may
exist
in
various
environments.
Several
lanthanide
and
actinide
element Pourbaix diagrams are
shown in the figure at right.
• What form will Yb take in
natural
lake
water?
Answer: Yb(OH)3
• Can uranium be solubilized
in sea water? Answer: Yes
as UO22+.
• Would you expect to find
cerium metal free in
nature?
Answer:
• Plutonium is highly toxic,
as well as being strongly
radioactive. What would
happen if plutonium was
released into a lake or a
stream?
Answer:
• Imagine that plutonium
oxides from a nuclear
weapons
processing
centre had been dumped
into a small, well aerated
lake where the pH = 6-8
and E = 0.0-0.5 V. Over
the course of 20 years, the
lake converted to a bog,
where the pH = 4 and E =
0.1 V.
Discuss the
environmental concerns in
the initial stages and the
final stages of the lake
"storage",
remembering
that plutonium would be a
serious toxic hazard if it
entered the food stream.
Answer:
Note that there are some
limitations to this approach.
First of all, the concentration of the elements in natural waters is often much lower than standard conditions. Furthermore, these
data are only valid for pure water plus the element in question. For example, on this basis we would say gold cannot exist in sea
water. However, it does, as a chloro complex, which is soluble. This is due to a complexation equilibrium. Nevertheless the
combined E/pH diagrams provide a very comprehensive insight into the behavior of the elements in aqueous solution, and this
is clearly the most important set of conditions relevant to the terrestrial environment.