# How can water below a pH of 7 still be alkaline?

Alkalinity is a measure of the water's ability to neutralize acidity and is expressed in levels of bicarbonates, carbonates, and hydroxides.

How can water that is slightly acidic, e.g. 5.5-6.5, still boast any alkalinity? Or put another way, is there a maximum level of alkalinity for water at any given acidic pH?

This is based on my understanding that acidity is, put simply, when there are more H+ than OH- ions. So why don't all carbonates react with the surplus H+ ions at pH ranges 5.5-6.5 until either alkalinity is 0 or the pH has reached a neutral 7?

• From Wikipedia: "Alkalinity … should not be confused with basicity which is an absolute measurement on the pH scale". Looks like that's exactly what you are trying to do here. – andselisk Dec 22 '20 at 9:07
• Thanks you two! So, alkalinity has nothing directly to do with basicity. Instead, it has to do with compounds that will react with (buffer) acids, correct? What confuses me is perhaps what an acid is. I thought an acid is simply a solution with surplus H+ that, subsequently, just raises the H+ surplus when added to water. Which will subsequently "agitate" the buffer. – Malte Dec 22 '20 at 9:41
• For your purpose, an acid is a compound able to release H+, a base is a compound able to either release OH- or to bound H+. Note that a free base can exist in acidic solutions, if it is a weak enough base for given pH. Similarly for acids. A very weak acid can exist in its acidic form even in a concentrated solution of a hydroxide. – Poutnik Dec 22 '20 at 9:47
• A free base can exist in acidic solutions if it is sufficiently weak. Would you expect that to change if the solution was agitated, e.g. stirred and or heated? – Malte Dec 22 '20 at 9:55
• No, it is an equilibrium condition. If e.g. pKa of acid is 13 and pH=11, then about 99% of the acid is at the equilibrium conditions in its acidic form and 1% in its conjugate base form ( being quite a strong base ). – Poutnik Dec 22 '20 at 10:03

Classically, alkalinity is determined by $$\ce{HCl}$$ titration with methyl orange as indicator ( $$\mathrm{pH}$$ transition range $$3.1 - 4.4$$ ). IF some $$\ce{HCl}$$ is spent by lowering $$\mathrm{pH}$$ from $$5.5-6.6$$ to $$\mathrm{pH}$$ of noticing red colour, there is some nonzero alkalinity.
Note that at $$\mathrm{pH} \approx 6.3$$, the ratio $$\dfrac{[\ce{CO2}]}{[\ce{HCO3-}]}$$ is 1 : 1. (*)
Generaly, there is no direct relation between $$\ce{pH}$$ and maximal alkalinity, similarly as there is no direct relation between $$\ce{pH}$$ and $$\mathrm{pH}$$ buffer capacity.
But for typical water composition, containing $$\ce{CO2/HCO3-}$$ $$\mathrm{pH}$$ buffer, the alkalinity can be estimated from $$\ce{pH}$$ and expected $$\ce{CO2(aq)}$$ concentration, which is supposed to be in equilibrium with gaseous $$\ce{CO2}$$.
$$[\ce{HCO3-(aq)}] = K_{a1,\ce{CO2}}^{*} \cdot \frac {[\ce{CO2(aq)}]}{[\ce{H+(aq)}]}$$
(*) The Wikipedia page for $$\ce{H2CO3}$$ has been recently misedited, confusing $$K_\mathrm{a2}$$ and $$K_\mathrm{a1}^*$$(assuming 100% conversion of $$\ce{CO2(aq)}$$ to $$\ce{H2CO3}$$.