CHEM3120 · Environmental and Analytical Chemistry
Carbonate & Sulfide Equilibria; Ocean Acidification
Lectures 19-20 apply speciation to the carbonate and sulfide systems in CHEM3120: the H2CO3 / HCO3- / CO3^2- equilibria (pKa1 ~6.3, pKa2 ~10.3), why dissolved [H2CO3] is fixed by atmospheric CO2, and how rising CO2 drives ocean acidification and dissolves CaCO3. This is central to Prof. Kepert's examinable Block 3 list and appears as quantitative Part A short-answer plus MCQ.
What this chapter covers
- 01The carbonate system H2CO3 / HCO3- / CO3^2- with pKa1 ~6.3 and pKa2 ~10.3
- 02[H2CO3] fixed by atmospheric P(CO2) via Henry's law — pH-independent (horizontal on a log-speciation diagram)
- 03Which species dominates in each pH window; the two crossovers at the pKa values
- 04Ocean acidification: rising CO2 lowers pH and shifts speciation
- 05H2CO3 + CO3^2- → 2HCO3-: added CO2 consumes carbonate, lowering [CO3^2-]
- 06CaCO3 (aragonite/calcite) dissolves when [Ca2+][CO3^2-] falls below Ksp
- 07The sulfide system H2S / HS- / S2- and FeS precipitation controlled by Ksp
Carbonate speciation in seawater, and the effect of rising CO2
- +1Locate pH 8.1 on the pKa scale: it lies between pKa1 (6.3) and pKa2 (10.3), so the middle species, bicarbonate HCO3-, is dominant.
- +1Apply Henderson-Hasselbalch to the second dissociation (HCO3- ⇌ H+ + CO3^2-): pH = pKa2 + log([CO3^2-]/[HCO3-]), so log([CO3^2-]/[HCO3-]) = 8.1 - 10.3 = -2.2.
- +1Therefore [CO3^2-]/[HCO3-] = 10^-2.2 = 6.3×10^-3, so the inverse ratio is [HCO3-]/[CO3^2-] = 10^2.2 = 1.6×10^2 ≈ 158.
- +1Effect of rising CO2: added CO2 forms H2CO3, which consumes carbonate (H2CO3 + CO3^2- → 2HCO3-), lowering pH further and reducing [CO3^2-]. As [CO3^2-] falls, the ion product [Ca2+][CO3^2-] drops below Ksp for aragonite/calcite, so shells and skeletons dissolve rather than form.
Key terms
- Carbonate system
- The coupled equilibria H2CO3 ⇌ HCO3- ⇌ CO3^2- with pKa1 ~6.3 and pKa2 ~10.3; it buffers natural waters and controls CaCO3 solubility.
- pKa1 / pKa2 (carbonic acid)
- The two dissociation constants of carbonic acid (~6.3 and ~10.3); HCO3- is the dominant species between them, which covers most natural-water pH.
- [H2CO3] is pH-independent
- In water equilibrated with the atmosphere, dissolved CO2/H2CO3 is fixed by P(CO2) through Henry's law, so it plots as a horizontal line on a log-concentration speciation diagram.
- Ocean acidification
- The fall in seawater pH as the ocean absorbs atmospheric CO2 (about 30% of emissions); it shifts carbonate speciation toward HCO3- and lowers [CO3^2-].
- H2CO3 + CO3^2- → 2HCO3-
- The reaction by which added CO2 consumes carbonate ion, the key step that lowers [CO3^2-] and undersaturates seawater with respect to CaCO3.
- CaCO3 solubility / Ksp
- CaCO3 (as aragonite or calcite) dissolves when [Ca2+][CO3^2-] falls below its Ksp; falling [CO3^2-] from acidification drives this dissolution.
Carbonate & Sulfide Equilibria; Ocean Acidification FAQ
Which carbonate species dominates in natural water?
Bicarbonate, HCO3-, across most of the natural range. Because HCO3- is the middle species and it dominates between pKa1 (~6.3) and pKa2 (~10.3), and most natural waters sit in that window (seawater ~8.1), bicarbonate is the main form of dissolved inorganic carbon. Carbonic acid dominates only below ~6.3 and carbonate ion only above ~10.3, both unusual in ordinary waters.
Why is dissolved H2CO3 drawn as a horizontal line on a speciation diagram?
Because in water open to the atmosphere its concentration is fixed by the CO2 partial pressure through Henry's law, independent of pH. Changing the pH redistributes HCO3- and CO3^2-, but as long as P(CO2) is constant the dissolved CO2/H2CO3 stays the same — so on a log-concentration-versus-pH plot it is a flat horizontal line while the other species rise and fall.
How does rising CO2 cause ocean acidification and dissolve shells?
Extra CO2 dissolves to form carbonic acid, which lowers pH and reacts with carbonate ion, H2CO3 + CO3^2- → 2HCO3-. That consumes CO3^2-, so its concentration falls. Since CaCO3 stays solid only while [Ca2+][CO3^2-] exceeds its Ksp, the drop in carbonate pushes seawater toward undersaturation, and aragonite or calcite shells and coral skeletons begin to dissolve rather than grow.
Can AI help me with carbonate equilibria and ocean acidification in CHEM3120?
Yes. Sia can walk you through identifying the dominant carbonate species, computing a speciation ratio with Henderson-Hasselbalch, and reasoning through the acidification chain to CaCO3 dissolution, step by step. It explains the method and checks your reasoning; it does not complete graded assessment for you, and University of Sydney academic-integrity rules apply.
Exam move
This chapter is prime Block 3 exam territory, so build fluency with the carbonate system. Fix the pKa landmarks (pKa1 ~6.3, pKa2 ~10.3) and the rule that the middle species HCO3- dominates most natural water; practise computing a species ratio with the correct pKa in Henderson-Hasselbalch. Understand why [H2CO3] is pH-independent (Henry's law) and be able to narrate the acidification chain — added CO2 consumes CO3^2-, lowers [CO3^2-], undersaturates CaCO3 — with the reaction H2CO3 + CO3^2- → 2HCO3- at its centre. The parallel sulfide system (Ksp control of FeS) is built the same way, so learn one template and reuse it. Keep this warm across the semester; the closed-book exam supplies a formula sheet. Confirm the exam date and permitted materials on Canvas.
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