Solubilities – reactions following dissolution
Aqueous Geochemistry (part 2)
2
e.g. CO2 dissolution
• Solutions – solubilities continued
CO2(g) ↔ CO2(aq)
• Effects of additional reactions
• after dissolution
• during dissolution
CO2(aq) + H2O ↔ H2CO3(aq)
• pH, buffering, and alkalinity
+
H2CO3(aq)
( q) ↔ HCO3 ((aq)
q) + H ((aq)
q)
• Oxidation/reduction
HCO3-(aq) ↔ CO3-(aq) + H+(aq)
• Eh, oxidation potential
• Supercritical fluids
In a cases like this, the Ksp of the initial dissolution step does not accurately reflect
the total solubility. The additional reactions increase the overall solubility.
• Additional characteristics of natural waters
Note also that when CO2 dissolves, several different forms (or species) can occur.
The relative abundances of these different forms is referred to as the speciation of
CO2 in the system. Here we have four different reactions, each with its own
equilibrium constant.
4
Solubilities – reactions following dissolution
the formation of complexes and complex ions
e.g. dissolved Cu
a complex (no charge):
a complex ion (has charge):
CuSO4
Cu(H2O)62+(aq)
(Cu2+ surrounded by 6 water molecules)
So, when Cu dissolves in solution, the resulting Cu2+ ions may react further
to form a complex or complex ion thereby reducing the amount of Cu2+ ion
in the solution and ultimately increasing the total amount of Cu that the
solution may contain.
5
6
Solubilities – reactions during dissolution
e.g. hydrolysis, one type of chemical weathering reaction
pH
pH = -log[H+]
where [H+] is the activity of H+
in the solution
MgSiO3(s) + 3 H2O ↔ Mg2+(aq) + H4SiO4(aq) + 2 OH-(aq)
or MgOH+(aq) + H4SiO4(aq) + OH-(aq)
hydrolysis of enstatite, an orthopyroxene
reactions like this are also affected by the pH of the solution
In this case, MgSiO3(s) doesn’t readily dissolve directly into Mg2+ and
SiO32-. Instead, it undergoes a hydrolysis reaction with the water to form
other ionic species. Thus, the solubility is controlled primarily by the
hydrolysis.
Can also rearrange as: [H+] = 10-pH
Activity is essentially the effective concentration rather than
the actual concentration. For sufficiently dilute solutions,
activity ≈ concentration.
This can be thought of as the availability of H+ in the system, although
hydrogen tends not to exist as free protons. In actuality, the H+
becomes associated with another water molecule to form H3O+.
1
7
CO2 dissolution to form carbonic acids
For pure water, the following occurs
these increase the
availability of H+
and therefore lower
the pH
CO2(g) ↔ CO2(aq)
H2O ↔ H+ + OHK = [H+][OH-] = 10-14 (at 25oC)
CO2(aq) + H2O ↔ H2CO3(aq)
For pure water, [H+] = [OH-], so
+
H2CO3(aq)
( q) ↔ HCO3 ((aq)
q) + H ((aq)
q)
K = [H+][OH-] = [H+][H+] = [H+]2
Thus,
[H+]2 = 10-14
8
pH
pH
HCO3-(aq) ↔ CO3-(aq) + H+(aq)
[H+] = 10-7
and
(= 0.0000001)
pH = -log[H+] = -log[10-7] = 7
As a result of these reactions, pure water in equilibrium with
atmospheric CO2 has a pH of 5.6 at 25oC.
Therefore, the pH of a neutral solution is 7 at 25oC
Since K varies with temperature, the neutral pH also varies with temperature.
9
Å greater H+ activity
neutral
H+ and OH- concentrations over a range of pH values at 25oC
greater OH- activity Æ
10
Note that the sum of the exponents on the H+ and OH- concentrations
is always the same, regardless of the pH. Why? K = [H+][OH-] = 10-14
[H+]
pH
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
1
0.1
0.01
0.001
0.0001
0.00001
0.000001
0.0000001
0.00000001
0.000000001
0.0000000001
0.00000000001
0.000000000001
0.0000000000001
0.00000000000001
11
pH
Buffering – tendency to maintain a near-constant pH in the presence of
certain materials
[OH-]
100
1x
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-4
1 x 10-5
1 x 10-6
1 x 10-7
1 x 10-8
1 x 10-9
1 x 10-10
1 x 10-11
1 x 10-12
1 x 10-13
1 x 10-14
0.00000000000001
0.0000000000001
0.000000000001
0.00000000001
0.0000000001
0.000000001
0.00000001
0.0000001
0.000001
0.00001
0.0001
0.001
0.01
0.1
1
Description
10-14
1x
1 x 10-13
1 x 10-12
1 x 10-11
1 x 10-10
1 x 10-9
1 x 10-8
1 x 10-7
1 x 10-6
1 x 10-5
1 x 10-4
1 x 10-3
1 x 10-2
1 x 10-1
1 x 100
very acidic
moderately acidic
mildly acidic
pure water, neutral
mildly basic
moderately basic
very basic
pH
12
Buffering
e.g. carbonic acid and calcium carbonate
H2CO3(aq) ↔ HCO3-(aq) + H+(aq)
e.g. carbonic acid
H2CO3(aq) ↔
K=
HCO3-(aq)
+
H+
(aq)
[HCO3- ][ H + ]
= 10−6.4
[H 2CO3 ]
If we add an acid (lots of H+), the reaction will go to the left, largely
eliminating any affect on pH. If we add a base, this will reduce the
available H+, and the reaction will go to the right, again largely eliminating
any affect on pH from the addition. The pH tends to stay close to 6 due to
the buffering effect of having both H2CO3 and HCO3- present.
CaCO3(s) + H2CO3(aq) ↔ Ca2+(aq) + 2 HCO3-(aq)
Here some of the products and some of the reactants are the same in both
reactions Thus,
reactions.
Thus the equilibria are linked together
together. The combination tends
to buffer a solution at close to a pH of 8.
The buffering effects of carbonic acid alone and carbonic acid together with
calcium carbonate control the pH of many natural waters.
Ocean acidification: As CO2 increases in the atmosphere, more CO2
dissolves in the oceans. This ultimately produces carbonic acid. As the
acidity increases, CaCO3 tends to dissolve more, and this creates problems
for organisms that produce CaCO3 structures (e.g. corals).
2
13
pH
H2CO3(aq) ↔ HCO3-(aq) + H+(aq)
carbonic
acid
(and CO2)
HCO3-(aq) ↔ CO3-(aq) + H+(aq)
bicarbonate
ion
14
pH
H2CO3(aq) ↔ HCO3-(aq) + H+(aq)
HCO3-(aq) ↔ CO3-(aq) + H+(aq)
carbonate
ion
White's Geochemistry
15
pH
pH
Alkalinity
Alkalinity
• Capacity of a solution to neutralize acid (i.e. react with H+ ions)
• Alkalinity can also be defined in terms of the amount of acid that
must be added
• Any ion in the solution that can react with the H+ contributes to the
total alkalinity
2-)
16
• This is directly related to a key way that alkalinity is measured:
titration
-)
• Carbonate (CO3 and bicarbonate (HCO3 ions are major
contributors to total alkalinity, so it is common to refer also to
carbonate
b
t alkalinity.
lk li it
• Titration refers to the addition of a strong acid (or base) in very
small
ll increments
i
t until
til a specific
ifi pH
H is
i reached.
h d In
I this
thi case, the
th total
t t l
amount of acid (e.g. HCl) needed is an indication of the alkalinity.
• Carbonate alkalinity and total alkalinity values are often similar. So,
sometimes the total alkalinity is used as an indication of the amount
of carbonate and bicarbonate in the solution.
17
Oxidation/Reduction Reactions (Redox)
Equivalence points 1 & 2 (EP1, EP2) on this graph represent common
target pH values for the titration.
18
Oxidation:
• Originally defined as combination with oxygen (e.g.
combustion), hence the terminology
• More completely defined as reactions involving the removal of
electrons
• Occurs in the presence of an oxidizing agent which accepts the
removed electrons
Reduction:
• Defined as reactions involving the addition of electrons.
• Occurs in the presence of a reducing agent which provides
(gives up) the needed electrons
White's Geochemistry
3
19
Oxidation/Reduction Reactions (Redox)
Fe occurs in two common oxidation states:
Oxidation/Reduction Reactions (Redox)
Fe occurs in two common oxidation states:
removal of 1 e- (oxidation)
for example:
Fe2+
FeO
→
20
removal of 1 e- (oxidation)
Fe3+
Fe2+
→
Fe3+
for example:
Fe2+ & O2-
addition of 1 e- (reduction)
less oxygen, so
more reduced
Fe2+
←
ferrous iron
“reduced” Fe
(relatively soluble
in H2O)
addition of 1 e- (reduction)
Fe2O3
Fe3+
ferric iron
“oxidized” Fe
more oxygen, so
more oxidized
(relatively insoluble
in H2O, except at
low pH)
Oxidation/Reduction Reactions (Redox)
21
Fe3+
For example:
For example:
FeO
Fe2O3
Fe2+ & O2-
2 Fe3+ & 3 O2-
less oxygen, so
more reduced
more oxygen, so
more oxidized
Oxidation/Reduction Reactions (Redox)
22
• Reflects the overall oxidizing/reducing properties of a solution
half reaction B
• Typically measured by inserting a set of electrodes which compare
the solution to a reference half reaction (H2 → 2 H+ + 2 e-)
flow of electrons in wire
Zn(s) → Zn2+ + 2 e-
←
Eh - oxidation potential (or redox potential)
Batteries work through oxidation-reduction reactions. For example:
half reaction A
Fe2+
2 Fe3+ & 3 O2-
Cu2+ + 2 e- → Cu(s)
• Eh has a large + value if oxidizing agents are abundant (i.e. the
solution has strong oxidizing properties)
flow of ions in solution
• Eh has a large – value if reducing agents are abundant
Cu gains electrons
Therefore it becomes
“oxidized”
Therefore it becomes
“reduced”
This happens at the anode (+)
This happens at the cathode (–)
In the solution, anions will
move toward the anode
In the solution, cations will
move toward the cathode
Oxidation/Reduction Reactions (Redox)
Major controls on Eh and pH in natural systems:
1) Biological processes (photosynthesis, respiration, decay)
• Can be thought of as a measure of the availability of electrons
• Oxidizing agents don’t have to be oxygen, so a large + Eh may not
indicate abundant dissolved oxygen
• Overall solution composition controls Eh, not the oxygen content
alone
23
24
Eh and pH
oxidizing
Zn loses electrons
2) Redox reactions involving:
• S
- e.g. sulfide weathering
• N - e.g. N2O, NO2- (nitrite), NO3- (nitrate), NH4+, organic N
• C - e.g. decay of organic matter
reducing
• Fe - Fe2+ vs. Fe3+
4
25
Eh and pH
Oxidation/Reduction Reactions (Redox)
26
Examples involving Fe
two half reactions:
FeCO3 + H2O → Fe3O4 + 3 CO2 + 2 H+ + 2 esiderite
magnetite
FeCO3 + 3 H2O → Fe(OH)3 + HCO3 + 2 H+ + 2 esiderite
iron hydroxide
What is the oxidation state (charge) of Fe on each side?
What do we need to complete these reactions?
27
28
Eh and pH
reducing
reducing
oxidizing
oxidizing
Eh
less CO2
Dissolved oxygen
The dissolved oxygen content is a commonly measured parameter in
surface waters because it has significant biological implications.
29
30
Dissolved oxygen
Note how the (maximum) solubility varies with temperature
Aerobic – oxygen available, organic matter decays releasing CO2
Anaerobic – oxygen not available
available, organic carbon is more likely to
accumulate in sediments since decay is limited
These two situations produce different microbial populations.
Fe and S content of the water and sediments also will be different due
to the different redox conditions.
5
31
32
Compare the arsenic, sulfide,
and CH4 with the oxygen
content of the water.
Where is the lake water aerobic?
Where is it anaerobic?
Why does this pattern exist?
Mon Lake in July, 1999. From Oreemland et al. (2000).
Dissolved oxygen
Mon Lake in July, 1999. From Oremland et al. (2000).
Supercritical fluids
34
Under magmatic and deep
metamorphic conditions, water
behaves differently than it does near
the Earth’s surface.
At sufficiently high pressures and
temperatures (beyond the critical
point), the distinction between liquid
and vapor vanishes
vanishes. The resulting
phase isn’t exactly a liquid or a gas,
rather it is a supercritical fluid.
For solutions, this transition to a
supercritical fluid takes place at
different P and T than for pure water.
Density and solubility vary
substantially with P close to the
critical point.
Some typical compositions of natural waters
35
Some typical
compositions of
natural waters
36
6
Some typical compositions of natural waters
37
7