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I was reading up on the buffering effect of the ocean and I was quite confused about how the buffer is able to do anything. The buffer I will refer to is below: $$\ce{CO3^{-2} +H2O <=> HCO3 + OH^{-}}$$ $$\ce{HCO3^{-} +H2O <=> CO3^{-2} + H3O^{+}} \tag{2}$$

When we have an excess of $\ce{CO2}$ this reaction (between sea water and carbon dioxide) will shift to the right: $$\ce{CO2 +H2O <=> H2CO3}$$ Thus it should also protonate $$\ce{H2CO3 + H2O <=> HCO3^{-} + H3O+} \tag{1}$$ However the products of this ionization produce species that are on the opposite sides of Equation 2. Therefore how can the equilibrium shift away to counteract this increase as both sides would have had an increase? I've heard that what should happen is that Equation 2 shifts to the left and decreases the concentration of carbonate ions but I don't understand why it shifts.

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    $\begingroup$ I took the liberty of numbering your equations so that they can be referred to. Now to the point: equilibrium (2) is not as important as (1). $\endgroup$
    – Ivan Neretin
    Commented Mar 8, 2019 at 7:10
  • $\begingroup$ But without (1) wouldn't there be no buffering action - which would defeat calling it an "ocean buffer"? $\endgroup$
    – John Hon
    Commented Mar 8, 2019 at 10:20
  • $\begingroup$ That's right, (1) is important. (2) isn't. $\endgroup$
    – Ivan Neretin
    Commented Mar 8, 2019 at 10:31

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