Teori tentang perubahan sifat asam basa air laut

1. Asam Basa Laut

2. Acids
HCL  H+ + Cl-
Proton donor, i.e., they donate H+ ions
HCl is a strong acid with a pH 1-2

3. Na+ + OH-  NaOH
NH3 + H+  NH4
OH - + H+  H2O
HCO3
- +H+  H2CO3
Proton acceptor, i.e., they take up H+ ions
NaOH is a strong base ~pH 12
Bases

4. Neutralization-
HCl + NaOH H2O + NaCl

5. Buffer- resists dramatic changes
in pH; ex. tums, rolaids…buffers
stomach acid

6. Acidic 0-6
Neutral 7
Basic (alkaline) 8-14
Type of
Solution pH Value
0-14

7. Logarithmic scale
blood

8. © 2011 Pearson Education, Inc.
Carbonate Buffering System
• Ocean pH averages 8.1 and ranges from 8.0 to 8.3.
• Buffering keeps the ocean from becoming too acidic or
too basic.
• Precipitation or dissolution of calcium carbonate, CaCO3,
buffers ocean pH.
• Oceans can absorb CO2 from the atmosphere without
much change in pH.

9. Carbon Dioxide System in the Ocean
Respiration
Photosynthesis
C6H12O6 +6O2 6CO2 + 6H2O
6CO2 + 6H2O C6H12O6 +6O2
Air
Water
CO2 gas

10. CO2 + H2O ↔ H2CO3 ↔ HCO3
- + H + ↔ CO3
2- + 2H+
By-product of respiration
carbonic acid
bicarbonate
carbonate
The addition of CO2 makes water acidic
The effects of CO2 in an ocean system

11. Bicarbonate buffer
Seawater too basic:
H2CO3 HCO3
- + H + pH drops
Seawater too acidic:
HCO3
- + H + H2CO3 pH rises

12. Fig. 6.16

13. Ocean Acidity

14. Global Ocean Acidity

15. How much CO2 can the ocean
absorb?
• The total amount of any gas seawater can absorb
depends on temperature and salinity
• Salinity is a measure
of the dissolved salt
content of water
Remember this relationship!
Temperature
or
Salinity
Amount of gas
seawater can absorb
15
Increases in temperature and salinity can decrease the
amount of gas seawater can absorb.

16. Ocean acidification
• CO2 is corrosive to the shells and skeletons of
many marine organisms
http://www.biol.tsukuba.ac.jp/~inouyePhoto: Missouri Botanical Gardens
Corals Calcareous plankton

17. Ocean Acidification
CO2 + H20 HCO3
- + H+
Water becomes
more acidic.
(ACID)
Remains in the
atmosphere
(greenhouse gas)
Dissolves in
sea water
CO2
CO2
Over the last 200 years,
about 50% of all CO2
produced on earth has been
absorbed by the ocean.GlobalWarming:TheGreatestThreat©2006DeborahL.Williams
(Royal Society 6/05)

18. Ocean Acidification
 Since 1850,
ocean pH has
decreased by
about 0.1 unit
(30% increase in
acidity). (Royal Society
2006)
 At present rate
of CO2 emission,
pH predicted to
increase by
0.4 units (3-fold
increase in H
ions) by 2100.
 Carbonate ion
concentrations
decrease.
Feely, Sabine and Fabry, 2006
Historical and Projected pH and Dissolved CO2
pH
Dissolved
CO2
Lower pH = MORE ACID
1850 2000 2100

19. Consequences of Ocean Acidity
Animals with CaCO3 skeletons affected
• Plankton
• Corals
• Mollusks
• Fish
http://news.bbc.c
o.uk/2/hi/science
/nature/7933589
.stm
Fisheries

20. In a high CO2 world, the ocean will be…
• More acidic
• More stratified
• More oligotrophic, but better
light conditions
• Less oxygenated
Consequences of Ocean Acidity

21. Acid Rain in Marine Environment
• reduces ability of marine
organisms to utilize calcium
carbonate
• Coral calcification
rate reduced
15-20%
• Skeletal density
decreased,
branches thinner

22. Fig. 6.18

23. Fig. 6.19

24. Fig. 6.20

25. Fig. 6.21

26. Fig. 6.22

27. Page 109

28. Solubility of CO2 and
Carbonate Equilibrium
CO2(g) ↔ CO2(aq) ↔ H2CO3 ↔ HCO3
-
↔ CO3
2-
CO2(g) ↔ CO2(aq) + H2O ↔ H2CO3
H2CO3 ↔ HCO3
-
+ H+
HCO3
-
↔ CO3
2-
+ H+
H2O ↔ H+
+ OH-
“Carbonic acid”

29. Solubility of CO2 and
Carbonate Equilibrium
1. KH = [H2CO3]/PCO2 = 3 x 10-2 M atm-1 = 10-1.5 M atm-1
2. K1 = [HCO3
-][H+]/[H2CO3] = 9 x 10-7 M = 10-6.1 M
3. K2 = [CO3
2-][H+]/[HCO3
-] = 4.5 x 10-11 M = 10-10.3 M
4. Kw = [H+][OH-] = 10-14 M2
Note: Values of these equil. constants are sensitive to temperature and ionic strength of
the solution; these values are appropriate to seawater.
CO2(g) ↔ CO2(aq) ↔ H2CO3 ↔ HCO3
- ↔ CO3
2-
KH K1
K2

30. CO2 Partitioning
Atmosphere - Ocean
pH of Natural Waters
So we want to understand what controls pH…
KKC Box Fig. 8-2

31. pH of Natural Waters
General concept…
At least six unknowns:
• H+, OH-
• PCO2
• H2CO3, HCO3
-, CO3
2-
Need at least six equations:
• Equilibrium expressions 1 - 4
• Typically, constraint on either PCO2 (“open system”) or total
moles carbon (“closed system”)
• Charge balance; Sn[in+] = Sm[jm-]

32. pH of Natural Waters
Pure water in contact with atmosphere
• Six unknowns
• Equilibrium expressions (4 equations)
• PCO2 = 3.5 x 10-4 atm (1 more equation)
• Charge balance (6th equation):
[H+] = [HCO3
-] + 2[CO3
2-] + [OH-]
• Strategy: Rewrite charge balance equation in terms of [H+] and
known quantities…

33. pH of Natural Waters
Example: Pure water in contact with atmosphere
[H+] = K1KHPCO2/[H+] + 2K1K2KHPCO2/[H+]2 + Kw/[H+]
Can solve rigorously for [H+]. Alternatively, make a simplifying assumption:
[CO3
2-] << [HCO3
-]
In this case:
[H+] = K1KHPCO2/[H+] + Kw/[H+]
or
[H+]2 = K1KHPCO2 + Kw
This is easily solved:
For pure water @25°C: K1 = 4.45 x 10-7 M; KH = 3.39 x 10-2 M/atm
[H+] = 2.4 x 10-6 M
pH = 5.62
“Acid rain” is a term applied to pH < 5

34. pH of Natural Waters
Assess simplifying assumption…
Is it fair to assume [CO3
2-] << [HCO3
-]?
K2 = [CO3
2-][H+]/[HCO3
-] = 10-10.33 (pure water, 25°C)
So: [CO3
2-]/[HCO3
-] = 10-10.33/[H+]
Clearly, [CO3
2-]/[HCO3
-] << 1 as long as [H+] >> 10-10.33
i.e., as long as pH << 10.33
This is true in most natural waters

35. pH of Natural Waters
Alkalinity
Consider charge balance:
[H+] = [HCO3
-] + 2[CO3
2-] + [OH-]
By definition, at pH = 7, [H+] = [OH-]
So, if [HCO3
-] or [CO3
2-] are at all comparable to [OH], [OH-] must be less
than [H+].
This would imply pH < 7, not pH > 7.
So how does this work… How do we get pH > 7???
Must add cations…

36. pH of Natural Waters
Alkalinity
• Imagine we dissolve some CaCO3 in the system
• Now: 2[Ca2+] + [H+] = [HCO3
-] + 2[CO3
2-] + [OH-]
• In this case, [H+] is free to have lower values (pH > 7) as long
as [Ca2+] is present

37. Consider CO2 Dissolution (again)
PCO2 = [H2CO3]/KH
= [H+][HCO3
-]/(K1KH)
= (K2[HCO3
-]/[CO3
2-])([HCO3
-]/(K1KH))
= K3[HCO3
-]2/[CO3
2-]; (K3 =K2/(K1KH))
So what?

38. CO2 Dissolution, HCO3- and CO32-
PCO2 = K3[HCO3
-
]2/[CO3
2-
]
This is the equilibrium expression for the reaction:
CO2 + CO3
2-
+ H2O ↔ 2HCO3
-
i.e., dissolution of CO2 consumes CO3
2-
, produces HCO3
-
;
Note that dissolution of CO2 itself does not affect Alkalinity
(gain 2 moles HCO3
-
for every CO3
2-
)
Note also that capacity for CO2 uptake determined by [CO3
2-
]

39. CO2 Dissolution, HCO3
- and CO3
2-
Another perspective
CO2 + H2O ↔ H2CO3
H2CO3  H+ + HCO3
- Net direction if pH >~ 6
H+ + CO3
2-  HCO3
- Net direction if pH <~ 10
CO2 + CO3
2- + H2O  2HCO3
-

40. CO2 Dissolution, HCO3
- and CO3
2-
Alk ~ [HCO3
-] + 2[CO3
2-]
ΣCO2 = [H2CO3] + [HCO3
-] + [CO3
2-]
~ [HCO3
-] + [CO3
2-]
PCO2 = K3[HCO3
-]2/[CO3
2-]
Algebra…
[CO3
2-] = Alk - ΣCO2
[HCO3
-] = 2 ΣCO2 - Alk
PCO2 = K3(2 ΣCO2 - Alk)2/(Alk - ΣCO2)
i.e., PCO2 is controlled by Alk and ΣCO2

41. CO3
2- in Ocean
Didn’t we say [HCO3
-] >> [CO3
2-]?
Yes, but that doesn’t mean CO3
2- is irrelevant:
Today: [CO3
2-] ~ 1/50 [HCO3
-]
As long as there is excess CO3
2-, ocean can take up CO2 without change
in alkalinity (or pH).
Figure 11-3 Distribution of dissolved carbon
species in seawater as a function of pH. Average
oceanic pH is about 8.2. The distribution is
calculated for temperature of 15o
C and a salinity
of 35‰. The equilibrium constants are from
Mehrbach et al. (1973).

42. * Reactions are quite congruent, structure not attacked; simple reactions.
* Acid Precipitation:
2CaCO3 + 2H+ + SO4
=
 2Ca + + + SO4
=
+ 2HCO3
-
* There are natural acids
ie. In rain CO2 in the atmosphere (pH = 5.6)
CO2 + H2O  H2CO3  H + + HCO3-
PCO2 ~ 350 ppm  pH = 5.6
* Rate of carbonate weathering is about 100 times the rate of silicate
weathering.
• ie. Florida- formation of karst due to the dissolving of calcite- wouldn't
see that with silica.

43. pH of Natural Waters
Alkalinity
Not all cation sources will work this way…
• Imagine we dissolve some NaCl into the system
• Now: [Na+] + [H+] = [HCO3
-] + 2[CO3
2-] + [Cl-] + [OH-]
But: [Na+] = [Cl-], right?
• So, no effect on charge balance equation
• To cope: Distinguish between “conservative” and
“nonconservative” ions…

44. pH of Natural Waters
Alkalinity
• Conservative ions: Ions whose concentrations are not affected by pH (or
pressure or temperature; not important variables here)
• Examples: Ca2+, Na+, NO3
-, K+, Cl-, etc.
• Nonconservative ions: Ions whose concentrations are affected by pH
• Examples: CO3
2-, HCO3
-, NH4
+, B(OH)4
-, H+, OH-
• “Alkalinity” ≡ Sn[in+] - Sm[jm-] where i and j are only conservative ions;
alkalinity is what’s left over after these are accounted for.
• Units: equiv./liter

45. pH of Natural Waters
Alkalinity
• If we consider only HCO3
-
, CO3
2-
, OH-
and H+
and conservative ions,
then we may write:
• Sn[in+
] + [H+
] = [HCO3
-
] + 2[CO3
2-
] + [OH-
] + Sm[jm
-]
• Sn[in+
] - Sm[jm-
] = [HCO3
-
] + 2[CO3
2-
] + [OH-
] - [H+
]
• Alkalinity = [HCO3
-
] + 2[CO3
2-
] + [OH-
] - [H+
]
• Typically, Alkalinity ~ [HCO3
-
] + 2[CO3
2-
] ≡ Alkcarb
• For seawater, Alkalinity ~ 2.3 x 10-3
equiv/liter

46. pH of Natural Waters
Alkalinity
• Alkalinity of seawater allows it to dissolve more CO2
• Higher alkalinity leads to lower [H+] and higher pH
• Any reaction that introduces [H+] lowers alkalinity
• e.g., NH4
+ + 2O2 → NO3
- + H2O + 2H+
• Any reaction that raises [CO3
2-] or [HCO3
-] raises alkalinity
• – e.g., CaCO3 + H2O + CO2 → Ca2+ + 2HCO3
-

47. Distribution of major species of dissolved
inorganic carbon at 20o
C
Figure 9.1 Fetter, Applied Hydrogeology 4th Edition

48. Organic acids (e.g.. Oxalic acid)
* Organic material breaks down and releases acids ( pH
~ 5 ). These natural acids play an important role in
weathering in absence of human activity; behaves very
much like carbonic acid
4H2C2O4 (oxalic acid) + 2O2  8CO2 + 4H2O
H2CO3
- -

49. Carbonate Weathering
Simplified reactions (removes H+ ions)
CaCO3 + H2CO3  Ca++ + 2HCO3
-  pH ~ 7 - 7.5 of water with
bicarbonate; about neutral
(Calcite) (Carbonic acid) (bicarbonate)
•large amounts of calcite is the reason, in general,
why places such as Rochester and the Mid-Eastern
states don't have problems with acid rain.
•Problem of acid rain: Determined basically by
the presence of absence of calcite- Tells how
sensitive an area will be to acid.
Calcite reacting
to acid
http://lewis.up.edu/mcs/wasowski/sci110/TnTkCh02/sld001.htm

50. Distribution of major species of dissolved
inorganic carbon at 20o
C
Figure 9.1 Fetter, Applied Hydrogeology 4th Edition

51. Concentration of calcium in equilibrium with calcite as a
function of PC02
in the system of CaCO3-CO2-H20 at 25o
C
and 1 atm total pressure.
Figure 3-4. Drever, The Geochemistry of Natural Waters

52. Concentration of calcium in equilibrium
with calcite as a function PC02
and Na+
concentration in the system of CaCO3-
Na2CO3-CO2-H20 at 25o
C and 1 atm
total pressure.
Concentration of calcium in equilibrium
with calcite as a function PC02
and Cl-
concentration in the system of CaCO3-
CaCl2-CO2-H20 at 25o
C and 1 atm total
pressure.
Figure 3-5,6. Drever, The Geochemistry of Natural Waters

53. Changes in composition of carbonated water as it equilibrates with
calcite when the system is either open or closed to exchange of CO2
gas. Initial PC02
values of 10-2 and 10-1 atm.
Figure 3-8. Drever, The Geochemistry of Natural Waters

54. http://coastal.er.usgs.gov/publications/ofr/00-180/intro/fig9.html
Florida karst
formation

55. Solution and collapse features of
karst and karren topography
http://coastal.er.usgs.gov/publications/ofr/00-180/intro/karst.html