strong acids/ bases strong acids more readily release h+ into water, they more fully dissociate –h...
TRANSCRIPT
Strong Acids/ Bases
• Strong Acids more readily release H+ into water, they more fully dissociate– H2SO4 2 H+ + SO4
2-
• Strong Bases more readily release OH- into water, they more fully dissociate– NaOH Na+ + OH-
Strength DOES NOT EQUAL Concentration!
Acid-base Dissociation• For any acid, describe it’s reaction in water:
– HxA + H2O x H+ + A- + H2O
– Describe this as an equilibrium expression, K (often denotes KA or KB for acids or bases…)
• Strength of an acid or base is then related to the dissociation constant Big K, strong acid/base!
• pK = -log K as before, lower pK=stronger acid/base!
][
]][[
AH
HAK
x
x
pKx?
• Why were there more than one pK for those acids and bases??
• H3PO4 H+ + H2PO4- pK1
• H2PO4- H+ + HPO4
2- pK2
• HPO41- H+ + PO4
3- pK3
• LOTS of reactions are acid-base rxns in the environment!!
• HUGE effect on solubility due to this, most other processes
Geochemical Relevance?
Dissociation of H2O
• H2O H+ + OH-
• Keq = [H+][OH-]• log Keq = -14 = log Kw
• pH = - log [H+]• pOH = - log [OH-]• pK = pOH + pH = 14
• If pH =3, pOH = 11 [H+]=10-3, [OH-]=10-11
][
]][[
2OH
OHHKeq
Definition of pH
pH• Commonly represented as a range between
0 and 14, and most natural waters are between pH 4 and 9
• Remember that pH = - log [H+]– Can pH be negative?– Of course! pH -3 [H+]=103 = 1000 molal?
– But what’s ?? Turns out to be quite small 0.002 or so…
BUFFERING
• When the pH is held ‘steady’ because of the presence of a conjugate acid/base pair, the system is said to be buffered
• In the environment, we must think about more than just one conjugate acid/base pairings in solution
• Many different acid/base pairs in solution, minerals, gases, can act as buffers…
Henderson-Hasselbach Equation:
• When acid or base added to buffered system with a pH near pK (remember that when pH=pK HA and A- are equal), the pH will not change much
• When the pH is further from the pK, additions of acid or base will change the pH a lot
][
][log
HA
ApKpH
Buffering example
• Let’s convince ourselves of what buffering can do…
• Take a base-generating reaction:– Albite + 2 H2O = 4 OH- + Na+ + Al3+ + 3 SiO2(aq)
– What happens to the pH of a solution containing 100 mM HCO3- which starts at pH 5??
– pK1 for H2CO3 = 6.35
• Think of albite dissolution as titrating OH- into solution – dissolve 0.05 mol albite = 0.2 mol OH-
• 0.2 mol OH- pOH = 0.7, pH = 13.3 ??
• What about the buffer??– Write the pH changes via the Henderson-Hasselbach
equation
• 0.1 mol H2CO3(aq), as the pH increases, some of this starts turning into HCO3-
• After 12.5 mmoles albite react (50 mmoles OH-):– pH=6.35+log (HCO3-/H2CO3) = 6.35+log(50/50)
• After 20 mmoles albite react (80 mmoles OH-):– pH=6.35+log(80/20) = 6.35 + 0.6 = 6.95
][
][log
HA
ApKpH
Greg Mon Oct 11 2004
0 10 20 30 40 50 60 70 80 90 1005
5.5
6
6.5
7
7.5
8
8.5
Albite reacted (mmoles)
pH
Bjerrum Plots
• 2 D plots of species activity (y axis) and pH (x axis)
• Useful to look at how conjugate acid-base pairs for many different species behave as pH changes
• At pH=pK the activity of the conjugate acid and base are equal
pH0 2 4 6 8 10 12 14
log
ai
-12
-10
-8
-6
-4
-2H2S
0HS-
S2-
H+OH-
7.0 13.0
Bjerrum plot showing the activities of reduced sulfur species as a function of pH for a value of total reduced sulfur of 10-3 mol L-1.
pH0 2 4 6 8 10 12 14
log
ai
-8
-7
-6
-5
-4
-3
-2
6.35 10.33H2CO3* HCO3- CO3
2-
H+
OH-
Common pHrange in nature
Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10-3 mol L-1.
In most natural waters, bicarbonate is the dominant carbonate species!
THE RELATIONSHIP BETWEEN H2CO3* AND HCO3
-
We can rearrange the expression for K1 to obtain:
This equation shows that, when pH = pK1, the activities of carbonic acid and bicarbonate are equal.
We can also rearrange the expression for K2 to obtain:
This equation shows that, when pH = pK2, the activities of bicarbonate and carbonate ion are equal.
*
1
32
3
COH
HCO
Ha
a
a
K
3
232
HCO
CO
Ha
a
a
K
pH0 2 4 6 8 10 12 14
log
ai
-8
-7
-6
-5
-4
-3
-2
6.35 10.33H2CO3* HCO3- CO3
2-
H+
OH-
Common pHrange in nature
Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10-3 mol L-1.
In most natural waters, bicarbonate is the dominant carbonate species!
THE CO2-H2O SYSTEM - I
Carbonic acid is a weak acid of great importance in natural waters. The first step in its formation is the dissolution of CO2(g) in water according to:
CO2(g) CO2(aq)
At equilibrium we have:
Once in solution, CO2(aq) reacts with water to form carbonic acid:
CO2(aq) + H2O(l) H2CO30
2
2
2
CO
COCO p
aK
THE CO2-H2O SYSTEM - II
In practice, CO2(aq) and H2CO30 are combined and this
combination is denoted as H2CO3*. It’s formation is dictated by the reaction:
CO2(g) + H2O(l) H2CO3*
For which the equilibrium constant at 25°C is:
Most of the dissolved CO2 is actually present as CO2(aq); only a small amount is actually present as true carbonic acid H2CO3
0.
46.1* 102
32
2
CO
COHCO p
aK
THE CO2-H2O SYSTEM - III
Carbonic acid (H2CO3*) is a weak acid that dissociates according to:
H2CO3* HCO3- + H+
For which the dissociation constant at 25°C and 1 bar is:
Bicarbonate then dissociates according to:
HCO3- CO3
2- + H+
35.6
*1 10
32
3
COH
HHCO
a
aaK
33.102 10
3
23
HCO
HCO
a
aaK
pH0 2 4 6 8 10 12 14
log
ai
-8
-7
-6
-5
-4
-3
-2
6.35 10.33H2CO3* HCO3- CO3
2-
H+
OH-
Common pHrange in nature
Bjerrum plot showing the activities of inorganic carbon species as a function of pH for a value of total inorganic carbon of 10-3 mol L-1.
In most natural waters, bicarbonate is the dominant carbonate species!
SPECIATION IN OPEN CO2-H2O SYSTEMS - I
• In an open system, the system is in contact with its surroundings and components such as CO2 can migrate in and out of the system. Therefore, the total carbonate concentration will not be constant.
• Let us consider a natural water open to the atmosphere, for which pCO2
= 10-3.5 atm. We can calculate the
concentration of H2CO3* directly from KCO2:
Note that M H2CO3* is independent of pH!
2
32
2
*
CO
COHCO p
MK
2232 * COCOCOH KpM
2232logloglog * COCOCOH KpM
SPECIATION IN OPEN CO2-H2O SYSTEMS - II
• The concentration of HCO3- as a function of pH is next
calculated from K1:
but we have already calculated M H2CO3*:
so2232 * COCOCOH KpM
*1
32
3
COH
HHCO
M
aMK
H
COH
HCO a
MKM *1 32
3
H
COCO
HCO a
pKKM 22
3
1
pHpKKM COCOHCO
2231loglog
SPECIATION IN OPEN CO2-H2O SYSTEMS - III
• The concentration of CO32- as a function of pH is next
calculated from K2:
but we have already calculated M HCO3- so:
and
H
COCO
HCO a
pKKM 22
3
1
3
23
2
HCO
HCO
M
aMK
H
HCO
CO a
MKM 3
23
2
2
12 2223
H
COCO
CO a
pKKKM
pHpKKKM COCOCO2loglog
2223
12
SPECIATION IN OPEN CO2-H2O SYSTEMS - IV
• The total concentration of carbonate CT is obtained by summing:
23332 * COHCOCOHT MMMC
2
211 2222
22
H
COCO
H
COCOCOCOT a
KpKK
a
KpKKpC
2
2111loglog22
HH
COCOT a
KK
a
KKpC
pH2 3 4 5 6 7 8 9 10 11 12
log
conc
entr
atio
n (m
olar
)
-8
-6
-4
-2
0
CTH+
OH-
H2CO3*
HCO3-
CO32-
pK1 pK2
Plot of log concentrations of inorganic carbon species H+ and OH-, for open-system conditions with a fixed pCO2
= 10-3.5 atm.
pH2 3 4 5 6 7 8 9 10 11 12
log
conc
entr
atio
n (m
olar
)
-8
-6
-4
-2
0
CT
H+
OH-
H2CO3*
HCO3-
CO32-
pK1 pK2
Plot of log concentrations of inorganic carbon species H+ and OH-, for open-system conditions with a fixed pCO2
= 10-2.0 atm.
Calcite Solubility?
• CaCO3 -> Ca2+ + CO32-
• Log K=8.48
• Ca2+ in Ocean = 0.0106 m