Bicarbonate is in equilibrium with carbonic acid, and carbonic acid is in equilibrium with carbon dioxide.

[1] $${HCO{_3}{^-} \Leftrightarrow CO{_3}{^2}{^-} + H^+}$$

[2] $${H_2CO_3 \Leftrightarrow HCO{_3}{^-} + H^+}$$

[3] $${CO{_2}{_(}{_a}{_q}{_)} + H_2O \Leftrightarrow H_2CO_3}$$

[4] $${{CO{_2}{_(}{_g}{_)} \Leftrightarrow CO{_2}{_(}{_a}{_q}{_)}}}$$

When RO permeate passes through a degasifier, the carbon dioxide gas is displaced by oxygen and nitrogen from the air that is blown through the solution.  The loss of carbon dioxide cause a shift in the equilibrium reaction to the left (see Eq. [1],[2],[3]).  The bicarbonate reacts with acid protons to form more carbon dioxide (Le Chatelier’s Principle), and in turn, carbonate proceeds to react with acid protons to replace bicarbonate.  Since H+ is being consumed, there will be a lower concentration of free [H+], resulting in a higher pH.

This higher pH does not help reduce corrosion.  The higher pH was caused as a result of the equilibrium shifting in favor of forming more CO2, at the expense of alkalinity.  Higher pH normally provides a protective effect because of the formation of ferrous carbonate, copper carbonate and lead carbonate scales on the corresponding metals.  But in this case, the carbonate species are being reduced, not increased.  Furthermore, this higher pH water will not be buffered, and by the time it reaches the point of use (POU), the pH will have decreased.  Therefore, a higher pH from the addition of sodium hydroxide has a very different effect on water chemistry than a higher pH from degasification of CO2.

Where possible, it is best to avoid gassing off CO2, and instead, adding a caustic solution to convert CO2 to alkalinity.

[5] $${H_2CO_3 + OH^- \Leftrightarrow HCO{_3}{^-} + H_2O}$$

[6] $${HCO{_3}{^-} + OH^- + CO{_3}{^2}{^-} + H_2O}$$

Carbonates react with corroding metals to form protective metal carbonates and also provide a pH buffer throughout the distribution system.

See also:

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