In simple terms, pH is the concentration of acid protons [H+]. On the other hand, the alkalinity of a solution is its ability to neutralize acids. Alkalinity consists of ions that incorporate acid protons into their molecules so that they are not available as a free acid that can lower the pH. This is known as buffering.

For example acid reacts with CO32- to make HCO3, and converts PO42- to HPO4–  .  This makes it so that a significantly higher quantity of acid is required to lower the pH compared to a solution that does not contain these ions. Deionized water can drop from pH 7 to pH 2 with just one drop of acid, while natural well water may require 200 – 300 ppm of acid just to lower pH from 7 to 6.

Alkalinity consists primarily of carbonate, bicarbonate, phosphate, borate, orthosilicate, sulfides, and organic acids. Most people refer to alkalinity as the concentrations of carbonate (CO32-) and bicarbonate (HCO3) ions, which are the buffers that are typically present in the highest concentrations in natural waters. Bicarbonate in particular, is the strongest buffer (largest Ka value) and the effect of other buffers becomes insignificant in its presence.

At very high pH, like pH 12, the hydroxide ion [OH] concentration is so high that it takes a significant amount of acid to neutralize enough of them before the pH drops. For that reason, hydroxides (OH) are considered as contributors to alkalinity above ~pH 10.5. At very low pH, hydronium ion [H+] concentrations are very high, and as a result, a much higher concentration of acid is required to further lower the pH.

Higher temperature shifts the equation to the right, slightly increasing the carbonate to bicarbonate ratio. At the same time, the acid (H+) concentration increases slightly which causes a slight drop in pH. This means that a warmer solution can have better buffering capacity despite a lower pH.

$${HCO_3{^-} \Leftrightarrow CO_3{^2}{^-} + H^+}$$

The opposite happens at lower temperature.

In reverse osmosis (RO) and nanofiltration (NF) membrane systems, bicarbonate is the primary driver behind the increased pH in the concentrate relative to the feed. Bicarbonates increase in concentration as they are rejected by the membrane, while CO2 concentrations remain constant since gases are not rejected. The increased ratio between bicarbonate and carbonic acid results in a shift to the left, absorbing more free acids and thereby increasing the pH.

$${H_2CO_3 \Leftrightarrow HCO_3{^-} + H^+}$$

**H2CO and CO2 are used interchangeably because H2CO3 converts instantly to CO2.

More on alkalinity can be found at the following blog post:https://www.membranechemicals.com/understanding-alkalinity/